129 PRIORITY POLLUTANTS
                               Volune I:

     Introduction and Technical Background, Metals and Inorganics,
                          Pesticides and PCBs
                Michael A. Callahan, Michael W. Siimak,
  Norman W.  Gabel, Ira P.  Hay, Charles F.  Fowler,  J.  Randall Freed,
Patricia Jennings, Robert L. Durfee, Prank C« Whitmore, Bruno Maestri,
         Will law R. Mabey, Buford R. Holt, and Co.istance Gould
                      EPA Contract No, 68-01-3852
                      EPA Contract No. 68-01-3867
                  Project Officer.  Donald J. Ehreth
             Monitoring and D»ta Support Division (WH-5S3)
                Office of Water Planning and Standards
                        Washington,  0,C, 20460
                        WASHINGTON, D.C. 20460

     Effective regulatory action for toxic chemicals requires an understanding
of the human and environmental risks associated with the manufacture, use, and
disposal of the chemical.  The assessment of risk requires * scientific judg-
ment about the probability of harm to the environment resulting from known or
potential environmental concentrations.  Environmental concentrations are a
function of (1) the amount and fora of the chemical released into the environ-
ment, (2) the geographic are*, (3) prior acc-jmulation, (4) time nf measure-
ment, and (5) the behavior of th* chemical in the environment.  The behavior,
or face and transport characteristics» of toxic pollutants in the environment
depends on a variety of chemical, physical, and biological processes (e.g.,
photolysis, hydrolysis, volatilization, sorption, biodegradatlon, biotransfor-
mation).  Evaluating these processes for specific compounds and placing each
Interaction into environmental perspective is the basic goal of this report.

     This two-volume report is a comprehensive review of the water-related
environmental fate and transport literature available for 129 chemical com-
pounds and elements, sometimes referred to as the 129 priority pollutants.

                                 Michael W. Slimak, Chief
                                 Ixposure Assessment Section
                                 Vater Quality Analysis Branch
                                 Monitoring and Data Support Division (WH-553)
                                 Office of Water Planning and Standards

This  rep.'rt  ha-;  h#>en  re"l<"W*-d  r>y  the   •'::<>•  >•  Wat*?r
:','>.  KPA, and   sppr .!'/<•> ,s,tr ;  L v  r*-f i.i.-f.r  tl'v:  vi--ws  !..
           PESTICIDES, AND PCB«.

Chapter    Dctcription

          . Section Is  Introduction and TechnicalR«ckground
  1.               Introduction
  2.               Face arid Traaaport.  proeeaaea
  3.               De'ermliiatio'i of Water-Related  Environmental Fate:
                     Procedure, Mathoda, and  Rap; rt  Format
  4.               Conciuaioni and Rectx»a«ndatiotu
           Section II;  Metals and Inorganica


Section IIIi Pesticides
Di eld tin

Endo»alfan and Endoaulfan Sulfata ,
Endrln and Endrln Aldehyde
Haptachlor E pox Ida
Hexachlorocyc lohaxana Ca , 8 , 5

Y-H«xachlorocycloh«xan« (Lindane)
Section IV; PCBa and legated Comjpounda
Polychlorinated Biphanyla
2-CWLor onapht ha lane

                  page blanl


Chapter     Description

            Section V;  Halogenated Aliphatic Hydrocarbons
 38.               Chlorome thane (Metbyl Chloride)
 39.               Dichloronethane (Methylene Chloride)
 40.               frichlorotnethane (Chloroform)
 41.               Tetrachloromethane (Carbon Tetrachloride)
 42.               Chloroethan* (Ethyl Chloride)
 43.               1,1-Dichloroethane (Ethylidlne Chloride)
 44.               1,2-Dichloroethane (Ethylene Dicbloride)
 45.               l,ltl-friehloroethan« (Methyl Chloroform)
 46.               l,l,2-Trichloroethan«
 47.               1,1,2,2-fetrachloroethane
 48.               Hexachloroe thane
 49.               Chloro«than8 (Vinyl Chloride)
 50.               1,1-Oichloroethene (Vinyl idine Chloride)
 51.               1 , 2-trans-Dichloroethene
 52.               frichloroethene
 S3.               Tetrachloroethene (Perchloroethylene)
 54.               1,2-Dichloropropar.e
 55.               1,3-Dichloropropene
 56.               Hexachlorobutadiene
 57.               Hexachlorocyclopentadiene
 58*               Bromorae thane (Methyl Bromide)
 59.               Brooodichloroae thane
 60.               Dibrooochloromethane
 61.               Tribromome thane (Broraofonn)
 62.               Dichlorodifluoroaechane
 63.               Trlchlorofluorome thane
            Stction VI^|  Halogenated Ethers
 64.               Bi«(chloroaethyl)eth«r
 65.               Bi»(2-chloroethyl)eth«r
 66.               ii8(2-chlorolsopropyl)ether
 67.               2-Chloroethyl vinyi ether
 68.               4-Chlorophenyl phenyl ether
 69.               4-Broaophenyl phenyl ether
 70.               Jii(2-chloroethoxy)methane
            Section VII:  Monocyclic
 71.               Benzene
 72.               Chlorobenzene
 73.               1,2-Oichlorobenzene (o-Dichlorobenzene)
 74,               1,3-Dichlorobentene (a-Dichlorobenzene)
 75.               1,4-Dichlorobenzene (j»~Dichlorobenzene)
 76.               1,2,4-Tri^' lorobenzene

            Pgsc cip t ion

            Section VII ;   ( eont 1 iued }
 77.               Hexachlorobenzene
 78,               Ethylbenzene
 79.               Nitrobenzene
 80.               Toluene
 81.               2,4-Dinltroi:ol«ene
 82,               2,6-Dinltrotoluene
 83.               Phenol
 84.               2-Chlorophenol
 85,               2,4-Diehlorophenol
 86.               2,4,6-Triehlorophenol
 87.               Pentachlorophenol
 88.               2-Nltrephenol   '
 89.               4-Nitrophtnol
 90.               2,4-DlnJ.trophenol
 91                2,4-OiMthyl phenol
 92.               jg-Chioro-a-cre«ol
 93.               4,6~Dinttr0-a~cresol
 94.               Phthalat* Esters;  Dimethyl, Dtethyl,
                     Dl-n-butyl, Ol-n-octyl; Il*(2-ethylh**yl),
                     and &ut>l benzyl      >
            gection IX ;   Foly sycl ic Aroaat ic Hydrocarfeon<
 95.               Polycyelic Aromatic Hydrocarbons: Acen«phthen«,
                     Acenaphchylene, Fluorene, and Naphthalene
 96,               Polycycllc Aromatic Hvarocarbons.-  Anthracene,
                     Fluoranthene, and Phenanthrene
 97,               Polycycllc Arotutlc llydrocarbons'  B*n2o(a]«Tthr*c«n«,
                     B«nzo[bJ f luoranthene , Benco(k)f luoranthsne, Chr/»en«,
                     and Pyrene
,98.               Polycyclle Aroaatlc Bydro««rboni;  Seneofghljperylene,
                     Benzol ajpyrene, Dib«nzo(a]anthrac«n», and
                     Indenojl , 2 , 3-cdl pyten*
            Sect ityg •__  ii t rotaalaea «nd Ml •ctilaneou*
                      __ ............
 99.               01«€thyl
100.               Olphenyl nltrosamine
101*               61-a-ptopfi nltroaaainc
102.               Benzidine
103.               3,3'-Di;hlorobenzidine
104.               l,2~Dip<»fiylhydrazine (Hydrazobenzen*)
105.               Acrylocltrlle

          A number of people have actively participated in this atudy, which was
      begun In September of 19??.  The assistance, support, and management guidance
      provided by Martin P. Halper, former Chief of the Water Quality Analysis
      Branch (1976-1978), and Donald J. Ehreth, present Chief of Che Hater Quality
      Analysis Branch, U.S. EPA, is gratefully acknowledged.

          Michael A* Callahan, O.S. EPA, conceptualized and managed the Phase I
      tasks and was the principal technical editor of all fate chapter*.  In addi-
      tion, he performed part of the literature search for the halogenated aliphatic
      hydrocarbons and participated in extensive rewrites of that section.  Michael
      V. Sliaak, U.S. EPA, managed the Phase II tasks, provided additional editorial
      assistance, and completed the final report*  Acknowledgement is also given to
      the staff members of th* Water Quality Analysis Branch.  Specifically, our
      thanks to:  Rod Frederick, Lynn Delpire, Charles Delos, Hark Segal, Charles
      Gentry, Chris Ehret, and Richard Scraydarlan.

          Most of the literature search and the writing of the report was done by
      two contractors, Versar, Inc. and S1I International.  The following individ-
      uals participated as shorn:

           1.  VERSAR IMC. (EPA Contract No. 68-01-3852)
                 Prograa Management:  Robert L. Durfee, Donald H. Sargent, Gayaneh
                 Contos, Bruno Maestri
                 Technical Direction?  Michael W. Slimak (If77-78), Norman W, Gabel
                   *  Metals and Inorganics:  J. Randall Freed, Ira ?. May
                   *  PCBs and Related Compounds:  J. Randall Freed, Frank C.
             ' .          Hhltmore, Charles F. Fowler
                   *  Halogenated Aliphatic Hydrocarbons:  Patricia Jennings,
                        Robert L. Durfee
                   *  Ralogenated Ethers:  Sorman H. Gabel
                   *  Monocyclic Arootatlcs:  Patricia Jennings, Norman W. Gabel,
                        Robert L. Durfee, Charles f« Fowler
                   *  Monocyclic Aromatics - Phenols and Cresols:  Norman W, Gabel
                   *  Phthalate Esters:  Bruno Maestri, Charles F. Fowler
                   *  Polycyclic Aromatic Hydrocarbons:  Michael W. Sliaak,
                        Charles F. Fowler
                   *  Sitrosaaines and Miscellaneous Compounds:  Norman W, Cabel

                 Data Gathering and Automated Literature Searches-;  Theodore French,
                 Michael Keller
                 Editing;  Neman W. Gabel, Robert L. Durfee, Bruno Maestri, Juliet
                     Preceding page blank

         VERSAR_INC._ (continued)

         Typing:   Laura Skiba, Shirley Harrison, Lorraine Douglas, Kathy
         Turnblom, Rebecca Brown, Nancy Downie, Neda Helmaadollat
         Graphics atnl Regrojuction;   Kenneth Ratkiewicz, Joieph Gillette,
         Carol* Craddock, Michael Cairns.

     2.  SRI INTERNATIONAL (EPA Contract So. 68-01-3867)

         Program Management:.  Stephen L. Brown, Oscar Johnson, Willian R,
         Ma bey
         Technical Direction;   William R.  Mabey

            *  Pesticides:  William R. Mabey, Buford Holt, Constance ,

         Automated Literature  Searches:  Shirley Radding, Jerie Etherton
         Editing:  Lee Work
         Typing:   Kathleen Williams, Maria Buyco,
    The general assistance and guidance provided by personnel within EPA's
Athens Environmental Research Laboratory,  especially George Baughmart and hla
staff, is also greatly appreciated.

    Numerous individuals provided comments and suggestions on the draft
report; their efforts are' appreciated.  , A special note of gratitude is due
those who carefully reviewed the draft  report and provided timely and
pertinent comments.  These reviewers Include;

              6. laughman, U.S. EPA-A Chens
              G. Zwaig, U.S. IPA-Washingr  n
              J. lariya, U.S. EPA-Washingtoc
              J. Cohen, U.S. EPA-Cincinnati
              J. Gillette, U.S. EPA-Corvallis
              M. Strier, U.S. EPA-Hashington
              J. Butler, Harvard University
              R. Baker, U.S. Geological Survey
              I. Irgolie, Texas A&M  University
              R, Bailey, Dow Chemical U.S.A.
              W. Dilling, Dow Chemical  D.S.A.

    Special gratitude is due Bruno Maestri and Shirley Harrison,  Versar Inc.,
for their efforts in completing the  final  document,  especially their attention
to detail.


                    Chapters 1-4

                             1.  INTRODUCTION

     The Off .ce of Water Planning and Standards (Monitoring and Data Sup-
port Division) of the U.S. Environmental Protection Agency (EPA) is con-
ducting a program to evaluate exposure and subsequent risk from the pres-
ence of toxic pollutants In our nation's environment.  This program
addresses the goals of tb« Clean Water Act of 1977.

     The environmental fat* processes discussed in this report are a key
component In an exposure assessment.  The goal of an exposure assessment is
an exposure profile which identifies subpopulations (geographic, demo-
graphic, etc.) and associate* how much, and what form, of a chemical comes
In contact with each •depopulation.  Ideally, this exposure profile can be
synthesized by Batching the location and habits of various a ubpopulatlons
with the location and fora of the chemical.  If the characteristics of
environmental release are well defined (location, amount released per unit
tine, for* of chemical, etc.), environmental fate processes can be used to
determine the ultimate location and form of the chemical .  Therefore,
environmental fate and environmental release data are a major part of an
exposure assessment.

     This two- volume report is a comprehensive review of the water-related
environmental fate and transport literature available for 129 chemical
compounds 'and elements, sometimes referred to as the 129 priority pollu-
tants. The pollutants (or In some cases classes of pollutants) are listed
In Section 307(a)(l) of the 1977 Clean Water Act (33 U.S.C. 46i et aeq.;
Committee Print HR. 3199).

         The objectives of this study were to:

          1.  Review and analyze the available information concerning
              significant environmental processes and associated
              kinetics for each of the 129 priority pollutants;

          2.  Identify the probable environmental pathways and fate of
              129 priority pollutants when introduced into surface waters;

          3.  Indicate the degree of confidence for the conclusions

     Volume I of thia report contains an Introduction, a description of
fate and transport processes, conclusions and recommendations, and a brief
discussion about the procedures used for collecting and reviewing the
literature. Volume I also contains chapters describing the fate of metals
and inorganic compounds, pesticides, and polychlorinatfid btphenyls.  Volume
II consists of a discussion of the fate of the halogenated aliphatic hydro-
carbons, halogenated ethers, monocyclic aromatics, phthalate esters, poly-
cyclic aromatic hydrocarbons, nitrosamines, and miscellaneous compounds,

1.2  Approach

     The fate of chemicals in the environment depends on a variety of
chemical, physical and biological processes.  Evaluating these processes
for a specific compound and placing each interaction into environmental
perspective is the basic goal of |
Athens, Ga., and Is described In a number of recent publications Including
the works of Wolfe *t al. (1976), Paris jet al. (1975), Hill «t al.1(1976),
and Smith et al. (1977).  The fundamental premises on which this approach
is based are:  (1) the overall rate of disappearance of a compound from the
aquatic environment is controlled by the dominant transformation and trans-
port processes; (2) these processes (e.g., photolysis, hydrolysis, volati-
lization, sorptlon, biotranaformatlon/biodegradatlon) can be studied In-
dependently in the laboratory; and (3) the laboratory data can be inte-
grated, using a model which simulates the environment, and extrapolated to
environmental conditions to predict exposure levels or concentrations.

    This two-voluae report consists of a detailed review of the literature
which describes the transport and transformation processes for each of 129
priority pollutants and indicates the most probable fate process of each
compound.  It certain case*, the data reviewed could not be analyzed in the
context of the transport and transformation processes, and in many in-
stances the studies reported in the literature were not conducted using
typical environmental conditions.  It should also be recognized that infor-
mation and data available in the literature may not always be suitable for
use in the context of an exposure model, and a critical review of experi-
mental procedures described In the literature is required to obtain reli-
able data for exposure assessments*

1.3  Literature _Cited

Hill, J., H.P. Kollig, 0.F, Paris, N.L. Wolfe and t.G. Zepp.  1976.
  Dynamic behavior of vinyl chloride in aquatic ecosystems.  U.S.
  Environ. Protection Agency, Athena, Ga. EPA-600/3-76-001,

Isenses, A.R., P.C. Kearney, I.A, Woolsen, G.I. Jones, and U.P. Williams.
  1973.  Distribution of alkyl arsenicals in model ecosystem**  Environ.
  Sci. Technol,  7j841-845.

Metcalf, R.L., J.R. Sanborn, P.-Y. Lu and D. Rye.  1976.  Laboratory model
  ecosystem studies of the degradation and fate of radiolabeled tri-,
  tetra-, and pentachlorobiphenyl compound with DDE.  J. Environ. Qual.

Paris, D.P., D.L. Lewis, J.T. Barnett, and G.L. Baughman.  1975.  Microbial
  degradation and accumulation of pesticides in aquatic systems.  U.S.
  Environ. Protection Agency, Athena, Ga. EPA-660/3-75-007.

Sanborn, J.R., R.L. Metcalf, W.N. Bruce, and P.-Y. Lu.  1976. The fate of
  chlordane and toxaphene in a terrestrial-aquatic model ecosystem.
  Environ. Entomol. 5:533.

Smith, J.H., W.R. Mabey, 9. Bohonos, B.I. Holt, S. S. Lee, T.-W. Chou,
  D.C. Bomberger, and T. Mill.  1977.  Environmental pathways of selected
  chemicals in freshwater systems; Part 1:  Background and experimental
  procedures.  U.S. Environ. Protection Agency, Athens, Ga.

Wltherspoon, J.P., I.A. Bondieth, S. Draggan, P. Taub, N. Pearson, and J.R.
  Trabokla.  1976.  State-of-the-art and proposed testing for environmental
  transport of toxic substances*  U.S. Environ. Protection Agency, Wash.,
  D.C.  EPA-560/5-76-001.

Volfe, N.L., R.C. Zepp, G.L* Baughman, B.C. Pincher, and J.A. Gordon.
  1976.  Chemical and photochemical transformation of selected pesticides
  in aquatic systems.  U.S. Environ. Protection Agency, Athena, Ga.

                     2.  FATE AND TRANSPORT, PROCRS_SF.S

2.1  Introduction

     There are a number of physical, chemical and biological processes that
nay be important in affecting the concentration of a chemical in an aquatic
system.  These processes include photolysis, hydrolysis, volatilization,
sorpticm, bioaccumulation, and biotransfonaatlon/biodegradation, and their
relevance and mathematical expression are briefly described below.  It must
be noted, however, that much of the literature reviewed did not report data
in the context of the theoretical discussions.  Each author was responsible
for carefully reviewing and evaluating the data and, where possible, pre-
senting the data in a form or expression similar to the way data are pre-
sented for the processes discussed below.  Often times this was not possi-
ble and qualitative judgments had to be made in order to interpret the re-
sults in light of these specific fate and transport processes.

2.2  Transport Processe_8

     2.2..1  Volatilization

            Volatilization of organic chemicals from water to the atraos-
phere can be an important pathway for chemicals with high vapor pressures
or low solubilities.  Early work reported in the literature often attri-
buted losses of chemicals actually due to volatilization as being due to
chemical or biological transformations.  Recognition of the possible im-
portance of volatilization In laboratory experiments and in the environ-
ment subsequently led to studies of, or at least consideration of, volati-
lization as a discrete process.  Some papers refer to such losses as "co-
distillation" with water, which is not technically correct since the loss
of water and of the, chemical are not interdependent (see Section 25,4.4).
While the importance of the volatilization pathway has been reported for
many chemical*, most of the reported data are difficult to apply to an
environmental assessment because of Inconplete Information on the factors
which influence the volatilization (e,g,» turbulence, temperature, ex-
perimental design).  Recent research has developed a better understanding
of volatilization processes in the aquatic environment.  The following
discussion describes the current understanding of the fundamentals of the
volatilization process, and discusses applications which have been employed
in some exposure assessment models.  This discussion is also useful for
evaluating the deficiencies and limitations of the information reported in
literature, which, in many instances, was never presented in the context of
a theoretical treatment.

         A two-res istaace theory, first proposed by Whitman (1923), can be
used to describe the rate of volatilization of a chemical (Llss and Slater
1974; Mackay and Leinonen 1975).  In general, the volatilization rate,
Rv, is a first-order process and can be described by:
       Rv * volatilization rate of a chemical, C
              (moles liter'1 hr~l);

       Cy » concentration of C in water
              (mole liter""1 - M);

       kv - volatilization rate constant (hr~*);

       L  " depth (ca);

       ^•i * liquid phase mass transport coefficient
             (CM hr'1);
       Hc - Henry** law constant (torr

       k« " gas phase mass transport coefficient
In both the gas and liquid phase ,

                                       55l                                (3)
where D is the diffusion coefficient and $ is the boundary layer thickness.

            There are several approaches which can be used to estimate the
mass transport coefficients for the chemical in the water body of Interest.
One convenient simplification i* based on the observation that if Hc >
3000 torr M"*1, Rv is determined by rhe value of k^ and is Halted by
diffusion through the liquid phase boundary layer*  For these highly vola-
tile compounds, equation (5) should be useful over a wide range of
environmental condition*;
where k  is the volatilization rate constant for the chemical
and ky1 is the oxygen reaeratlon rate constant (hr"*-) in the laboratory
or the environment.  The ratio k§/k§ for benzene has been found to be
independent of turbulence, salt concentration (seawatar), temperature
{4*~50*C}, and the presence of a surface active ioapound (Smith et il.
1979).                                                          "
            Alternatively, if Hc < 10 torr M~» only the second term in
equation (2) ia significant.  Then

                               i  .-JO.               •  .                (
                                *   LBT

and the volatilization rate Is United by diffusion through the gas phase
boundary layer.  If 3000 torr M"1 > Hc > lu torr M""1, both terras in
equation (2) art significant.  In these cases, the mass transport coeffi-
cients of the chemical In the water .body can be estimated from represen-
tative values of the aass transport coefficient for oxygen reaeratlon,
which is liquid phase res.  .tance controlled, and from the mass transport
coefficient of water, which is gas phase resistance controlled.

     2.2.2  Sorption

       ,  ;   The 3orption of chemicals to suspended sediments and bottom
sediment can be an Important process in aquatic environments.  The term
sorption is used in these reports since the other commonly used terms of
adsorption and absorption have mechanistic connotations which cannot be
identified in aost experiments*  In general, the more hydrophobia a
chemical is the more likely it is to be sorbed to sediment.

            Data for sorption of chemicals to particulates are frequently
expressed In terms of the Freundlich isotherm equation,
where Cs and Cy are the concentrations of chemical in particulate and
water phases, K- is a partition coefficient for sorption, and i/n Is an
exponential factor.  At environmentally relevant concentrations of a chemi-
cal in solution which are low compared to the particulars sorption capaci-
ty, the i/n tem is usually equal to unity within experimental error.  It
should be realized also that the measurement of Kp oust allow sufficient
tine for equilibrium between phases to be established; Information In the
literature indicate that tiae» to reach equilibrium range froa a few
minutes to several days.

            For neutral organic chemicals, the degree of sorption to sus-
pended sediments is dominated by Interaction with the organic content of
the particulate; a partition coefficient corrected for organic carbon,
KOC, equal to Up divided by th« fraction of organic carbon,  corre-
lates well with water solubility and octanol/water partition coefficients
(Kow> (Kenaga and Coring 1978).  This relationship between Kp and Koc
Is useful for predicting Kp values for a number of sediments where the
organic carbon content Is known or chosen.  Ic should be realized, however,
Chat neutral compounds art also sorbed by materials with little or no
organic content, such as sands and Inorganic clays, and the Kp data for
sorption on soils or sediments with vary high sand or clay content do not
fit the correlations for K- data for soils/ sediments where the organic
content is greater than about 1 percent.  Therefore, Koc data -should be
used with appropriate limitations.  Another important factor to le con-
sidered in the use of literature data is that concerning the units of Kp
(or Koc); although the units of C, and Cw are usually the sane so
that Kp is unitleis, sorae literature does not conform to this convention
and requires K_ recalculation before use.  Also, If the data were fit to
a Pruendllch equation with the exponent, n» not equal to unity, the value
of Kp with n « 1 at low concentrations of the chemical must  be
recalculated fron the original data.

2.3  ChemicalProcesses

     2.3.1  Photolysis

            Photoiyses of chemicals dissolved in aquatic systems occur at
wavelengths greater than 290 nm since ozone In the stratosphere filters out
light of shorter (higher'energy) wavelengths.  Photochemical transforma-
tions may occur by one or more processes depending on the chemical struc-
ture and substances' In the environment.  "Direct photolysis" processes take
place If tne chemical absorbs light and then undergoes a> transformation're-
action Jrom an excited state by any one of several mechanisms (i.*., re-  .
arrangement, dissociation, oxidation, etc.).  The rate of the reaction t»
dependent on the sunlight photon flu*, the light adsorption coefficients of
the chemical, and the reaction quantum yield; the last 1* the efficiency
for conversion of the absorbed light into chemical reaction.

            In contrast to direct photolysis, "Indirect photolysis" takes
place if substance* naturally present In aquatic environments absorb sun-
light to form excited chemical species or radicals which interact with the
chemical to produce a reaction.  On* type of indirect photolysis is a pho-
tosensitized reaction, In which the light-absorbing substance transfers
excited state energy to the chemical which chen undergoes a reaction, whicn
it»y or may not be similar to the teacrlon of those found in direct photoly-
sis processes.  Although th* literature suggests that such photosensitized
reactions may occur in natural waters*, there la as yet no unambiguous evi-
dence that this aechanism Is actually important In the  environment.  Other
types of photochemical reactions that nay be considered in the Indirect
photolysis class are those In which photolyzed natural substances produce
energetic tr,t/mediate* that react with the ground state of the chemical;
singlet oxygen and oxy radicals are examples of such intermediates which
can result in transformations of chemicals.  Since these intermediates
result from photoreactions of naturally occurring substances and since the
photochemistry of che chemical itself is not involved, It is not strictly
correct to refer to such reactions as photulyses; rather they should be re-
ferred to as photooxygt ration or photo-Initiated free radical reactions
(for iinglet oxygen and oxy radicals, respectively).  In most literature
information, arid frequently In laboratory studies, It Is impossible to
differentiate among the several type* of photochemical mechanisms in
natural waters, and therefore they are often necessarily grouped Into dis-
cussions of photochemistry.

            The rate of loss of chemical (- ^-} by either direct or in-
direct photochemical processes may be expressed by simple first  order
kinetic expressions.  The equation for direct photolysis is
where k_ is a first order rate constant, t is the reaction quantum yield,
and ka is a rate constant for absorption of light by the chemical; the
last is a function oi the photon flux, the distribution of light and the
light absorption coefficient of the chemical (Zepp and Cllne 1977; Mabey
e_£ a_l. 1979),  Th* rate equation for indirect photolysis by any mechanism
•ay be expressed by

                        -f - 4 ww - *; m                  .       uo)
where k.2 is a second order rate constant for reaction of chemical C with
the reactive intermediate X; for a photosensitized reaction the k-
rate constant would be a combined term for concentration of excited state
species and quantum yields (or efficiencies) of energy transfer to and
subsequent reaction of the chemical.  In any estimation of  kp or kp,
values of ka or (X} will either be for instantaneous rat**  (or concentra-
tions), or for those averaged over a specific time interval,  because the
first order rate constants will be dependent on the sunlight  Intensity
which varies with tlmt of day, season, and latitude,

     2,3.2  Oxidation

            As discussed in the photolysis section, oxidation may occur as
a result of oxldants formed during photochemical processes  in natural
waters.  The complex chemistry of free radical reactions in natural waters
has recently been reviewed by Mill *£ tl_,  (19/9), along with  the production
and reaction of singlet oxygen in these systems.  Based on  limited experi-
mental work, the average effective concentrations of the two  oxidant
species, alkylpcroxyl radicals and singlet oxygen, were estimated to be
10"* M and ICT*2 M, respectively.

            Other oxidation .processes have also Neen Investigated In con-
nection with water treataent processes using chemical oxidants such as
chlorine, ozone, or permanganate.   While this information may be relevant
to watei treatment processes, it obviously has little quantitative rele~
vance to aquatic envivonments except for providing a qualitative neasure of
susceptibility to oxidation; a  'heaiital which is inert to thnse oxtdants
will probably be stable to ox!1?'.ion in aquatic systems.

            The kinetic expression for the rate of loss of a chemical by
reaction with an oxidant [QxJ is given by
where k^ Is a second order rat* constant for reaction of the oxidant
with chemical C, and [0*] and [Cj »r« the concentrations of oxidant and
chemical, respectively.  Data for kg^ and the use of such information In
estimating oxidation half-lives has recently been reviewed by Mill (1979)
and Mill et. al. (1979).  As for sunlight variations in photolysis rates,
the value ofToxJ awst b* defined In terns of an average effective oxidant
concentration over a time period.  Since most literature information has
not identified oxldants responsible for los* of a chemical nor separated
out the component photolysis mechanisms that say have contributed to the
loss of a chemical, use of such information, must usually be considered
particular to defined experimental conditions, and therefore not applicable
to more general evaluations of oxidation.

            Hydrolysis of organic compounds usually results with the intro-
duction of • hydroxyl group (-OH) into a chemical structure, cowaonly with
the loss of a leaving group (-X):
                       MX * Hjl — BOM + HX (orM*, X~1.

7h* rat* of the reaction may be promoted by acid (hydronium Ion,  thO
and/or baa« (hydroxyl ion, OH").  Son* cheaical* will also show • pH
dependent elimination reaction;
                     K   X
                     II      M*w    N    -'
                    • C—C—  	--**   t'C   + HX
                     l|      -OH    /    \

Both processes are included in the scope of hydrolysis studies, where the
rate of hydrolysis is given by the equation,

                -'5-" MCI " **lH4f{CJ * h«lOH~J {C] + hM (C)               (13)
                  fft    n      jW          D            *•    '            i

where kj, is the first order rate constant at a specific pH; k^ and kg
are second order rate constants for the acid and base promoted processes,
respectively; and k^ is the first order constant for the pH independent
reaction of a chemical with water.

            Mabey and Mill (1978) h«v« recently reviewed data for hydroly-
sis of a variety of organic chemicals for use in prediction of half-lives
in aquatic systems.  Some chemicals such as alkyl halides have hydrolysis
rates which are independent of pH over che environmental pH range 4-9,
while others such a* carboxylic acid esters are acid and base promoted with
a minimum hydrolysis rate at pH 4-5.  Rate constants kA, kg and k»;
for a large nuaber of hydrolyzable structures can be estimated with reason-
able accuracy fron published data (Mabey and Mill 1978) or froa structure
reactivity correlations for these processes (Mill 1979).

2.4  Biological Processes

     2,4.1  _BijMceupuia;tloti

            BloaccumulatIon of chemicals in various living species has
been shown to result in significant ecological effects, and is especially
impcrtaft for hydrophobic cheaicals which can be partitioned into fat and
lipid tissues*  Bioaccumulation also occurs with Inorganic chemicals being
partitioned into bone narrow, etc.  Although bloaccunulation has been the
subject of ouch valuable research, interpretation and subsequent use of
bioaccutnulatIon data must be performed with careful attention to the ex-
perimental procedures employed.  The bioconcentration factot (8CF) is
usually defined as the concentration of a chealcal in tissue (on a dry-
weight basis) divided by the concentration In water; sone literature data,
however, reported the concentration in tissue on a wet-weight basis, and
therefore the SCF is lower than when reported on a dry-weight basis.  It
should be noted that BCFs frequently «ay be low as a result of insufficient
tine being allowed for the true partitioning equilibrium to be attained in
the systen*  The use of bioconcentration data Is alsr< complicated by the
fact that concentrations of a chemical will usually be higher in fatty
tissues of the species than in leaner tissues.  The rate of uptake and time
for attainment of equilibrium in various organs (and species) will also

 depend on the route of uptake (i.e.,  dietary, skin absorption,' etc.);  for
 obvious reasons,  moat bioconcentratlon data for aquatic systems are from
 fish studies,

             In spite of these problems, bloconcentration/bloaccumulation
,data are an important parameter for evaluating the impact of chemicals in
 an aquatic environment; it is significant to note that most of the above
 problems will lead to estimates of bloconcentratlon valued that are lower
 than the actual values.  Kenaga and Coring (1978) have also provided a
 useful correlation between BCF and octanol/water partition coefficients and
 water solubility  data.  The correlation is useful in assessing a chemical's
 potential for bioaceuniulation; therefore, the log P octanol/water partition
 coefficient is included in the physical properties section of each chapter,

      2.4.2  Biotransformation and Biodegradation

             Biodegradation results from the »nzyme-catalyzed transforma-
 tion of chemicals.  Organisms require energj, carbon, and other fundamental
 inputs fron the environment for their growth and maintenance.  In the pro-
 ce*s, they Manufacture enzyaes to transform many chemicals introduced into'
 the environment*   Because blodegradation processes in aquatic and soil en-
 vironments are carried out primarily by microbes, the effects of macro-
 biota are usually considered insignificant in studying biotransformation
 and biodegradation.                                                       ,

             The biodegradation rate is the function of a nlcrobial blonaas
 and a chemical's  concentration under given environmental conditions.  When
 microorganisms utilize chemical substrates, there are increases in blomass,
 and biodegradation rates will then be a function of cell growth rate.   When
 an organic compound Is utilized by microorganisms as a sole carbon source,
 the specific growth rate of organisms is a function of chemical concentra-
 tion.  The widely used Monod kinetic  equation, describing the relation be-
 tween growth-lira!ting substrate concentration (C) and the specific growth
 rate in a well nixed system is expressed as
 where y is th«  specific growth rate,  X is the biomass per unit volume,  Uffl
 Is the aaxiraum  specific growth rate,  and K8 is the concentration of sub-
 strate supporting  half~«aximu« specific growth rate (0.5 urn).

             Th« rate  of substrate utilization is then

                  -£, m  £2, t f y \11 ****i }- t h  }tt-,—-".-...                  (15)

where k^ la the substrate utilization constant or bi Degradation con-
stant, equal toym/Y, and ¥ Is the bioraass produced from a unit amount of
substrate consumed.  The constants pm  Kg and Y are dependent on the
characteristics of micfobes, pH, tempe'rature, other nutrients, etc.

            When substrate concentration is high and C » Ks, then the
above expression becomes.

                                 dC   .  x
                               ~iT   fc  •                             (16)
The degradation rate is first order with respect Co all biomass concentra
tions and zero order with respect to chemical concentration.

            For many pollutants in the en/ironment, substrate concentra-
tions are very low, such that C « Ka.  Equation (15) then becomes
where kj^ is a second order rate constant.  The degradation rate is then
first order early in cell concentration and in chemical concentration.

            In the environment, where the cell concentration X is relative-
ly large and pollutant concentration is low, alcrobial populations will not
change significantly when the chemical is consumed. The degradation rate
under these conditions Is pseudo first order and can be described by the
                                 «l    •-                               (18)

where k., is a pseudo first order rate constant.  The factor kp is
dependent on cell concentration (Xo)» therefore

                                  »                 '
     '    '               ..        -i.«"l2           '                   -(19)
and kj,2 1s tn* second order rate constant.  The half-life of  the chemical
       at a given X0 will be

           The degradation rate constants shown above are to be used under
those conditions where microorganisms are acclimated to the chemical and
can actively utilize the chemical.  However, when a pollutant is initially
introduced into the environment, there is often a lag period between the
exposure of the chemical to the organisms and the initiation of blodegrada-
tion for some chemicals.  A lag or acclimation period is required to Induce
the organisms to produce necessary enzyme(s) or to develop biodegradation
organisms by mutation.  A lag period may also result when there are too few
degrading microbes initially present in the ays tea.  Under these condi-
tions, no significant substrate consumption is usually detected until bac-
terial cells reach a substantial level.  This acclimation period may also
be caused by diauxic utilization1 of substrates where other readily metabo-
lizable organic compounds are present.

            Monod kinetics is applicable if the lag period Is caused by too
low an initial microbial level; mathematical treatment Is difficult if ac-
climation is caused by physiological adjustment of the oicrobial community.
In most cases, it is possible to1 assume that the environment has already
been exposed to the chemical and that acclimated organisms are already
present ; the biodegradation rate constant is then used for such a calcula-
tion with a given level of microorganisms.  When a chemical is newly re-
leased into an uncontaminated place, the lag period cannot be Ignored and
the tine required to reduce 50 percent of the original concentration
(Tj/^) is the sum of time required to reach acclimation (to) and the
half-life of transformation
            Natural aquatic .environments contain a number of organic com-
pounds of natural and anthropogenic origin.  Some pollutants may be bio-
transformed only when another organic compound is present to serve as a
carbon and energy source; this phenomenon is known as cotnecabolism.
Methods to evaluate the biodegradation of chemicals as a sole carbon source
may then often underestimate the rate of the biodegradation In the natural
environment.  Mathematical treatment of cometabolic transformation of
chemicals is not currently available, but cometabollsm should be considered
in any environmental fate assessment.

2«5  Other Reactions

     The foregoing processes are those which have been the subject of some
research and have beer, described in a quantitative manner suitable for
estimating half-lives o;' for use In various modelling efforts.  Other

processes have also been Infrequently reported as being possible important
pathways for some chemicals In aquatic systems, but have not been quanti-
tatively expressed in ways that are useful for Inclusion in modelling
efforts.  Two of these processes are described below.

    2.5,1  Reduction

           In anaerobic environments, reduction of chemicals by both bio-
logical and non-biological processes can occur.  Most,frequently report-
ed for these reactions are organochlorine chemicals, where a chlorine atom
is replaced by a hydrogen atom (see chapters on DDT, toxaphene.)  It is
expected that further work in this subject will provide knowledge of
specific reducing agents in aquatic environments and kinetic expressions
for the use of this Information In environmental assessments.

    2.5.2  Hydrajtlon

           Another process that may be important in aquatic ayscems for
some chemicals is hydratlon.  Carbony1 compounds are known to fora hy-
drates, which will have different properties than the parent cheaicals.
Therefore both th* transport and transformation pathways of the chemical in
the aquatic environment.will be affected.  The exlstance of hydrated
species must be considered in using experimental kinetic and equilibrium
data for environmental assessments.  Thus, since the hydratlon reaction is
                     unhydrated * H^feshydrated,

the kinetic expression for loss of the chemical may then Include a kinetic
loss term and the hydratlon equilibrium term.

2.6  Evaluationof Processes

     2.6.1  Kinetics (Half-lives)

            The transport and transformation processes, discussed above,
were reviewed and evaluated for each of the 129 priority pollutants to
determine a chemical's most probable aquatic fate by Identifying the pro-
ceases with relatively short half-lives.  A half-life is an estimate of the
environmental persistence of a chemical or the time required for removal of
one-half of th* Initial concentration of the chemical.  The principal ex-
perimental approach to the derivation of chemical "half-lives" is to meas-
ure the rate constants with which a reaction proceeds and the dependence

of the reaction rate constant on the concentration of the reacting species
and physical parameters (e.g., temperature).  The rates of chemical reac-
tions range from those that are completed almost instantaneously to those
that proceed so slowly that the reaction is essentially imperceptible.
Techniques of measurement typically involve the determination of the change
in concentration of a reactant or product as a functioT of time.

            For example, at constant volume, and for a unimolecular reac-
tion such as

     A  __»  Products           or            A  — -  B *- C,

the rate of disappearance of A can be expressed as:
where  C^   * concentration of A, In moles per liter;
       t    • tine, appropriate units;,
       kj   - reaction rate constant for process j, units of inverse
                time; and
      4$A   * rate of change of C^ with respect to time,

            Integrating the above between the limits of Co (Initial tine)
and t produces:
where C*  « initial concentration of A.
            The above is a first order reaction, In which the rate of
reaction of a species depends upon the first power of the chemical con-
centration.  Other reactions may be of higher order, for example the
nominal second order reaction:

                           A 4- §    —»    C + D

for which

Although some reactions of this font can b« treated as first order with
some assumptions (many hydrolysis react ions of pollutants In water, at
fixed pH, for example), this cannot be assumed until experimental data are
available which indicate first order kinetics.

            An estimate of environmental persistence for a process is given
in the form of a "half-life".  For first-order decomposition reactions when
         , the half-life takes th* form,
           Other equations for half-life must be used if the order of the
reaction is other than first order.  Obviously, knowledge of the rate
constant is required for half-life calculations based on first-order
reactions;  knowledge of concentration levels are required in addition to
rate constants if th* order is higher than first order.

           If all the transformation and transport processes have been
expressed as first or pseudo first order rate expressions, it is possible
to calculate an overall or net half-life for the chemical by the equation,
    2.6,2  Mi crocosa Studies, Field Studleg and Modelling

           Field studies and, on a aor« controllable scale,  microcosm
studies, are considered by sone researchers as direct approaches to en-
vironmental fate assessments of chemicals.  The modelling approach, on the
other hand, is a way of synthesizing information on the fate of a chemical

using data on component processes.  Each approach has its merits and dis-
advantages for use in the determination of the fate and pathways of chemi-
cals in aquatic environments.  When available, Information on such studies
has been Included in this literature review.  The discussion Is presented
by Individual process, such as biotranaformation or bioaccumulation, when
that process is clearly dominant over others.

           Field studies and microcosm experiments use real or physically
similar environmental situations to evaluate the rate of disappearance of a
chemical.  Such studies readily measure the bioaccutaulation or movement to
sediment of a chemical, but generally, do not provide information on the
relative importance of some component loss processes (e.g., photolysis,
volatilization, blotransforaatlon) which may be subsequently applied to
evaluations of other aquatic environments.   While microcosms provide note
control and knowledge of the factors Influencing the experiments, they
often must alao depart from reality in scale and in other factors such as
wind speed, sunlight variations and sediment scouring.  In stannary, fj.eld
studies, and to an extent microcosms,' are holistic approaches to predicting
the fat* of chemicals in aquatic systems, but suffer from their specificity
and unknown Interactions which Halt the application of results to other
aquatic systems.

           The modelling approach to environmental fate assessments com-
bines kinetic and equilibrium data on component processes to predict the
transport, transformation, and concentration of chemicals In the environ-
ment. Provided that suitable data for a chemical are available by measure-
ment or estimation methods, the modeJ may be designed with any degree of
sophistication.  An Important advantage of such models is that the fate of
a chemical can be predicted for any environmental situation; one such
scenario may be designed to estimate a conservative half-life or concentra-
tion for general environmental asstjament purposes, whereas a specific
field study may fortuitously provide an unreallstically rapid rate of loss
of the chemical due to dominance of a uniquely rapid process.  The model
may also allow for prediction of only transport and dilution in an environ-
ment where transformation data are unreliable, which could also be used for
a conservative risk assessment.  A major disadvantage of Che modelling ap-
proach with the present state of knowledge is that no chemical degradation
and transport model has been verified by comparison to actual environmental
experience, although work is underway in this area.  It is likely that the
verification of an aquatic fate model will eventually require a series of
fie?«* studies or alcrocorm experiments, with the model being used to design
che sampling strategy.  It Is also important to realize that the model must
be verified on several chemicals, each with different dominant fate pro-
cesses, before a widely applicable model can be established.

2.7  Literature ..... . Cited                                        .

Kenega, E.E. and C.A.I. Goring.  1978.  Relationship between water solu-
  bility, soil-sorption, octanol-water partitioning, and bioconcentration
  of chemicals in biota.  Amer. Soc. Test. Mat.  Third Aquatic Tox. Syap. ,
  Ne»: Orleans, La,  $3 pp.

Lias, P. 3. and P.G. Slater.  1974.  Flux of gases across the air sea
  Interface.  Nature.  247:181-184.

Mabey, W.R., T. Mill, and D.G. Hendry.  1979.  Test protocols for
  environmental processes:  Photolysis in water.  EPA Draft Report.  EPA
  Contract 68-03-2227.

Mabey, W.R. and T. Mill.  1978.  Critical review of hydrolysis of organic
  coapounds In water under environmental conditions.  J. Phys. Che*. Ref.
  Data 7:383.
Mackay, D. and P. J. Leinonen.  1975.  Rate of evaporation of
  solubility contaminants fro* water bodies to atmosphere*  Environ. Scl.
  Technol. 9, -1178-1 180.

Mill, T. , W.R. Mabey, and D.G. Hendry.  1979.  test protocols for
  environmental processes:  Oxidation in water,  EPA Draft Report.  EPA
  Contract 68-03-2227.

Mill, T.  1979.  Structure r*acti/lty conditions for environmental
  reactions.  EPA Report - EPA 560/11-79-012.

Smith, J.H., D.C. Bomber get, and D.M. Haynea.  1979.  Prediction of the
  volatilization rates of high volatility chemicals for natural water
  bodies.  Paper submitted for publication.  Study supported by NSF Grant
  So. PFS78-10691.

Whitman, W.G.  1923.  Preliminary experimental confirmation of the 2-film
  theory of gas absorption.  Chem. Mat all. Eng. 29:146-148.

Zepp, R.G. and D.M. Cline.  1977.  Rates of direct photolysis in aquatic
  environments.  Environ. Sci. Technol.  11:359-366.

3,1  Ident1fication and ColJ.ect ion of^

     3.1.1  Automated Data Searchea

            Coaputerized databases were used as one of the Intitlal steps
in collecting data on priority pollutants and fate-related processes.  The
bulk of the autossaeed data activity was conducted between December 1977 and
March 1978.  The following databases were searched for all priority

                   CHEM ABSTRACTS
                   OCEANIC ABSTRACTS
                \  POLLUTION ABSTRACTS
                ' I  SCISEARCH
                   SSIE CURRENT RESEARCH

            For a.l database* except CHEh ABSTRACTS, evsry citation con-
taining the priority pollutant's name or synonym was selected and com-
bined into a single data set using the Boolean operator "or."  Froa this
data set, specific citations were further evaluated using such keywords as
"volatility, photolysis, adsorption, «tc.," and compiled into a second data
set.  The Boolean operator "and" was used to combine the two data sets and
yield the final data set with one or sore chemical names and one or more
keyword identifiers.

            The CHEM ABSTRACTS database was searched using the combination
of registry nuabers and chemical area (e.g., photolytic reaction mecha-
nisms) to create a data set.  in the case of some chemical groups (such as
polycyclic aromatic hydrocarbons) all data were retrieved without qualifi-

     3.1.2  Manual Da ta Search

            The automated data search was followed by a comprehensive
manual search of the literature.  The scope and depth of the search vdried
for each chemical.  Both published and unpublished data were collected,
reviewed1 and filed for use in developing the individual fate chapters.
References within reviewed papers were also used to identify literature not
found by computer searching methods.  Journal articles published up to
January 1979 were included*
                                                                     ** '
3.2  Estimation of Physical and Chemical Parameters   .               *

     For many of the priority pollutants, basic data on phystcochemical
properties were not available.  Since information on vapor pressure,
oetanol/water partitioning, and solubility in water was necessary to evalu-
ate fate and transport processes , these parameters were estimated when no
literature values were found.  The methods used to estimate these para-
meters are discussed below.

     3.2.1  Vapor Pressure

            Vapor pressure at 25°C was estimated by tising one of four
methods, in descending order of preference:

            1.  Where constant 3 were available in handbooks, vapor
                pressure was calculated.

            2.  Where data for vapor pressure at temperatures bracketing
                25*C (e.g., 20*C and 30"C) was available from the
                literature, an interpolation was used.

            3.  Tables compiled by Driesbach (1952) were consulted.
                These tables relate pressure, volume, and temperature
                for major "Cox Chart'* chemical families.

            4.  Th* Glaus lus-Clapeyron equation was employed.

The above methods are summarized below.

            Calculation/interpolation,  Vapor pressures of many compounds
were reported for temperatures above and below 25*C, but not specifically
for 25* C.  for these compounds, the vapor pressure at 25*C was calculated
using the following formula (Weast 1974):
                             f * (-8J1W A/K) + I

where P Is the "apor pressure expressed in torr, A is the molar heat of
vaporization, K is the temperature in degrees Kalvin, and 8 is a constant..
For a given compound, A and 8 are unique and can1 be assumed constant over a
moderate temperature range.  Values for A and B are tabulated in Weast
(1974) far several of the priority pollutants and can be used to calculate
vapor pressures directly.

            Where two or acre values of vapor pressure were given in the
literature for two or more temperatures bracketing 25JC, and the constants
A and B were not reported, a linear interpolation method was employed.
Since Equation (1) has the general linear form of y • ox + b, where two
sets of ordered pairs (Kj, Pj) and (K2t ?2^ were- known, A and B
could be calculated.  These constants were then substituted into Equation
(1) to calculate P at K - 298° (-25*C).

            Tabular Values for Chemical Families.  If equation (1) could
not be employed because of insufficient data, tables in Driesbach (1952)
were consulted.  These tables were developed for the "Cox Chart" chemical
families (e.g., naphthalenes, halo-benzenes with saturated side chains,
phenols), using Antoine's equation, a variant of equation (1)<  To estimate
vapor pressure at 25SC, it is necessary to know the boiling point at 760
torr and the chemical family to which a compound belongs.  The appropriate
table Is then referenced for an estimated vapor pressure.

            Clausius^*- Clapeyron^Eq_uati.oa»  If neither of the aformentloned
Methods could be used, vapor pressure was calculated using the Clauslus-
Clapeyron Equation:
where P is vapor pressure in torr, dHv is the molar heat of vaporization,
T is teasperature in degrees Kelvin, R is the gas law constant (1.99 cai./
mole °K), and the subscripts 1 and 2 represent two different temperatures.
To solve this equation for vapor pressure at 25*C, the boiling point and
heat of vaporization must be known.  Although the vapor pressure calculated
by this formula is generally not as accurate as the estimates derived from
the other methods, the Clausius-Clapeyron equation is useful for providing
a rough estimate,

     3.2.2  Petanol/Water Partition_Coefitciegt

            Several investigators, notably Hansch e_t_ _al_. (1974), Lea et a I.
(1971), and Tute {1971), have demonstrated that log P (log of the octanol/-

water partition coefficient) can be estimated based on the functional
groups In organic molecules.  The basic approach used in this method ia to
consider a molecule as the s»m of It* functional groups.  These functional
groups can be assumed to cause a certain proportion of the partitioning
between octanol and water.  An index value called the "if" value has been
estimated for many of the common functional groups; this value can be posi-
tive or negative.  Estimated * values for various functional groups are
listed in Tut e (1971).  These values are summed to give an approximation of
log P.  An example of this calculation is given for dichlorobroaome thane:

                               i,_c— a

This molecule can be divided into the following components with correspond-
ing 1* values:

                  m                    o.so
              2 * Cl         2 * 0.39 - 0.7S
                  8r                    0.60
The estimated log F for dlchlorobroooaethan* is thus 1.88.  For further
information on estimation of log P values, the reader is referred to TuKe
(1971/ and leo et al. (1971).

     3.*«3  Aqueous Solubility

            Although aqueous solubility is readily available for most of
the priority pollutants, there were several cases where an estimate had to
be aade.  A number of different techniques are document ed in the litera-
ture, but the one chosen far this study was that of Moriguchi (1975).  This
technique is based on factoring water solubility into' two intrinsic com-
ponents:  free aolecular volume and hydrophillc effect of polar groups.

            Moriguchi (1975) tested six different additive parameters re-
lating to molecular voluoe and concluded that Quayle's parachor (Quayle
1953) was aost satisfactory for predictive purposes.  Like the techniques
used for estimating octanol/water partitioning,' Quayle's parachor is cal-
culated by considering a molecule as the SUM of its functional groups.
Each group is assigned a certain value, based on Interpretation of empiri-
cal data (Quayle 1953).  Quayle's parachor, and other parachors, were de-
veloped as physical parameters which can be readily correlated with the
structure of organic compounds to estimate molecular volumes.

            The second component to Moriguchl's (1975) technique, hydrophi-
He effect of polar groups, accounts for solute-solvent and solute-solute
Interaction*.  By examining eapirlcal data, a "hydrophtlic group effect"
factor waa calculated for various functional groups; factors are listed In
Morlguchl (1975).

            Finally, after calculating the Quayie's parachor for a molecule
and referencing the appropriate hydrophillc group factor, aqueous solubil-
ity can be estimated via the following formula:
where S in th* water solubility in nolal concentration, Pr is Quayle's
parachor, and E^ is the hydrophillc group factor.  This formula appears
to be fairly reliable:  when known aqueous solubilities for 156 compounds
wer» tested against Mtlnated solubility, th« correlation coefficient was
r - 0.962 (Morigtichl 1975).

            For more Information on this method of estimating solubility,
th« r«ad*r it referred to Morlguehi (1975) and Quayle (1953).

3.3  Review and Assessment of Data

                       of Data
            The literature w»s searched, using both automated and manual
methods, by individual* who were assigned specific chemicals or chenlcal
groups.  Th*re was aubatantial interchange of information among the in-
dividual* assigned to th* project (including individuals in EPA and both
contractors), and periodic meetings were held to discuss common problems
and share various interpretations of the fate literature.  The basic Phase
I evaluation lasted for approximately 5 vonths, vith an Interim draft sub-*
alt ted at th* end of th* period.  There were insufficient data to determine
many of th* compounds' most probable aquatic fat**  for others, data were
sufficient to adequately describe the dominant transport and transportation
processes.  The data collected after 5 month* were evaluated and used to
prepare th* interim draft report.

     3.3.2  Ff gpj r at ion o f ........ I nt e r ia &ra ft

            Bach individual assigned a compound or group of compounds was
r*«ponsibl« for writing an interim draft report sujmaarlzing the data
evaluated at the end of S Months.  These reports were carefully reviewed by
the Contractor's task atanager and, whenever necessary, meetings were held

among staff members to review the literature and help ascertain the pre-
dominant transport and transportation process.  IT addition to summarizing
the data reviewed, the Interim draft reports highlighted data gaps and
areas where Information was Insufficient for drawing conclusions.  As a
result, the interim draft reports focused on additional data needs for each
chemical, and provided the direction for additional data analysis. .
               i s                               i
     3.3.3  The PeerReview^Process                        , •

            Additional literature VMS collected during months 7,8, and 9 of
the program and a considerable amount of effort was put Into evaluating the
collected literature.  In many cases, original authors were actually con-
tacted by telephone to clarify their results, and in most cases, literature
was traced back to its primary source or author; secondary references were
used only when no alternatives existed.  As a result of this thorough
literature review and comprehensive evaluation, a more complete picture of
the aquatic fate of many of the compounds was obtained.  Still, for tome
compounds, there was not enough information to adequately determine the
aquatic fate and more work was necessary.

            To satisfy the EPA's standards for the technical accuracy of
each fate chapter and to make each chapter read as if It were written by
the same author, an elaborate Internal peer review process (EPA and the
contractors) was used.  Each scientist prepared a detailed oral presenta-
tion cohering his or her chemicals which was presented at a meeting at-
tended by senior scientists and project staff within the organization and
the EPA task manager and staff members.  During this presentation, data
were summarized, their weaknesses and strengths discussed, and conclusions
reviewed as a group, with general concurrence attained for each chemical.
This internal peer review process brought about a stimulating exchange of
Information and Ideas.  Conclusions reached were baaed on the group's con-
sensus rather than the assessment of one individual.

     3.3.4  Prejgar tlort of Final Draft

            The final draft report (a separate fate chapter for each
chemical) was prepared on the basis of the conclusions reached during the
Internal peer review.  This report was substantially more comprehensive
than the interim draft because it covered more data and was much more
carefully reviewed and evaluated by other individuals in addition'to the
prineipa,*, author.  The report formut was agreed upon for the final draft
and a common expression of data confidence or reliability was  followed by
all authors.  A uniform method of citing literature was adopted, and all
chapters conformed to this method.

            The chapters were typed using computerized word processing
equipment which eventually facilitated final changes.  The final draft re-
port was distributed to approximately 75 individuals and/or organizations
for review and comment; about half of this number were EPA personnel in
other program offices and research laboratories who were not participants
In the Internal peer review.  The remaining half consisted of Industrial or
academic personnel and, individuals with special knowledge and expertise
about specific ehenica] groups.  The distribution of the final draft for
external peer review (March to September 1979) completed the first phase of
the program.

   '  3,3.5  Preparation of	the Final Report

            The final fate chapters, as they appear in this two-volume re-
port, ware based on r.h* draft final report but included the Incorporation
of current literature (since September of 1977) and those comments (that
were considered appropriate) which were submitted by the external
reviewers.  In some cases individual compounds were contained into one
chapter (e.g., phthalate esters, polychlorinated biphenyla) and some
chapters were cospletely revised requiring an extensive technical rewrite.
Exaaples of the latter were the chapters on metals; the draft final did not
adequately consider how polluted environments affect the behavior of the
netals, especially with regard to metal-organic interactions.

            Each author prepared a brief technical letter to EPA describ-
ing the changes that were to be made in the final report based on the
external reviewers' comments. This provided a mechanise for EPA's continued
participation in the peer review process.  In addition to the final fate
chapters, an introductory section and a description of the transport and
transformation processes were prepared and are included in this final

            In summary, this report represents over two years of careful
data collection and review*  It is the product of an extensive exchange of
information and ideas which underwent a thorough peer review.  The result,
we believe, is a document that meets the objectives of the program.

     3.3,6  Con fidenee of Data

            One of th* basic objectives of the study was to indicate the
degree of confidence for the conclusions reached.  Each fate chapter in
this document contains a Cable which summarizes the aquatic fate Informa-
tion.   The last coluam in the table describes the confidence of th« summary
statement w.ade from the data reviewed.  Three somewhat subjective levels of
confidence were used — Low, Median, and High.  A brief description of
these  levels follows.

            High Confidence'.  This usually required that the data reviewed
be quantitative; rate constants and rates were either explicitly described
or could be calculated from the results.  In addition, experiments were
considered relevant if done at typical ambient environmental conditions of
temperature, pH, etc.  Normally, corroborating evidence from an unrelated
associate or experiment (I.e., "a second opinion") was also required to
place data in the "high" category.  In many Instances, the chemical struc-
ture itself and its inertness to certain processes (e.g., no hydrolyzable
groups for hydrolysis) was sufficient to place a summary statement in the
"high" category.

            Medium Confidence.  Qualitative data about a particular chemi-
cal (i.e., no rate data or information from which rates could be derived)
were typically given a "medium" confidence ranking.  In some cases, the
data were collected at somewhat irrelevent conditions (e.g., temperature,
pH) for the purpose of this study or in such a way that only a qualitative
judgment could be made about the results.  Some quantitative data were
placed in this category if there were no other data to corroborate the
results.  Quantitative data reported for a different but structurally
related compound were sometimes given a "medium" ranking if the different
compound was thought to react analogously to the actual compound being

            Low Confidence.  Summary statements made speeulatlvely, or
based on theoretical estimates or calculations, were given a "low" ranking
of confidence*  In many cases, the reviewers believed a particular process
to be important, based on theory; however, a "low" confidence was assigned
to the statement when thera were no actual investigations performed on the
particular chtalcal.  If the quantitative data were characterized by con-
troversy over rates and mechanisms, then usually the confidence was given a
"low" ranking.

3.4  Report Format

     The first four chapters of this report serve as Introductory and
explanatory material for the individual fate chapters that follow.
Chapters 5 through 105 are reports for each chemical o'r group of chemicals
comprising the 129 priority pollutants.

     The fate chapters for each organic priority pollutant (chapters 20
through 105) generally follows uniform format.  For several chemicals
(some well-studied pesticides, for example), however, there are slight
deviations from the format presented below.  For an organic chemical in a
typical chapter "X", the following sections are included:

X.I  Statement of Probable Fate

     This section summarizes the entice chapter.  A statement on
     Che probable transport and aquatic fate mechanisms is given.
     Other processes thought to be possibly important are identified.

X.2  Identification

     This section includes the chemical structure, alternate
     chemical names, the Chemical Abstract Service Registry Number
     (CAS), and the Toxic Substances List (TSL) number (as found la
     the,1977 NIOSH Registry of Toxic Effects of Chemical Substances).

X.3  Physical Properties

     This Includes literature or calculated values for:
     molecular weight, melting point, boiling point, vapor
     pressure, water solubility, and octanol/water partition

>',4  Summary of Fate Data

     X.4.1  ghotolyjis

            Statements are made on direct photolysis in aquatic
            systems and, If volatile, in the atmosphere.   A
            discussion of available spectral data Is also
            included, where relevant.

     X.4,2  Oxidation

            Reactions which involve oxidation processes are

     X.6.3  Hydrolysis

            Data on expected hydrolysis under natural conditions
            (temp. 0"C - 30°C, pH 6-9) are presented.

     X.4.4  Volatilization

            Experimental data or 'information are presented when
            available; if not, inferences are drawn based on vapor
            pressure and solubility data.

     X.4.5  Sqrgtipn

    1        Specific experimental evidence (e.g.» X adsorption on-
            clay) is presented.  Without these data, inferences
            are drawn based on partitioning values and other

     X.4.6  Bloaccumulation

            Same as for sorption.

     X*4.7  Biotranafoniation and_ jiocje^radatlon

            A summary of data on in vivo and ^n vitro
            degradation is presented.

     X.4.8  OtherReactions(Optional,depending on the cheoical)

            All data on processes that cannot be specifically
            characterized in the previous sections (X.4.1 through
            X.4.7) are presented here.

     X.4»9  Microcosm Studies, Field Stucf-ies, and Modelling (Optional,
            dependingonthe chemical)

            Pertinent data from microcosm studies, field studies, and
            modelling evaluations are presented in detail in this

X.5   DataSummary

            A summary matrix is presented with an identification
            of the most probable fate process(es) included.
            Also, all literature or calculated rate constants or
            half-lives are summarized.

X.6   Literature Cited

            All literature cited in the "X" chapter are presented
            using the American Institute of Biological Sciences
            standard format (AIBS,  3rd Edition).

    The format for the aetals chapters (chapters 5-19} is slightly differ-
ent from that shown above.  For metals, the Identification section of the
report Includes a discussion on the geochemistry of the substance.  Also,
the oxidation/reduction and hydrolysis sections were contained and expanded
into a section entitled "chemical speciation," which deals with the en-
vironmental chemistry of each metal.  The data s,umraary matrix was simpli-
fied because, barring radioactive decay, there are no possible half-lives
for metals.

3.5  literature Cited

Dreisbach, R.R.  1952.  Pressure-voluoe-teaperature relationship of
  organic compounds.  Handbook Publishers, Sandusky, Ohio.  p.3-260.

Hanseh, C., A. Vlttorla, C, Slllpo, and P.Y.C. Jow.  1974.  Partition
  coefficients and the structure-activity relationship of the
  anesthetic gases.  J. Med. Chen. 18(6):546-548.

Leo, A,» C, Hansch, and D. Slkins.  1971.  Partition coefficients and
  their uses.  Chen. Rev. 71(6):525-565.

Morlguchi, I.  1975.  Quantitative structure-activity studies.  I.
  Parameters relating to hydrophobicity. Chem. Phana. Bull,

Quayle, 0.R.  1953.  The parachors of organic compounds.  Cheo. Reviews

Tute, M.S.  1971.  Principles and practice of Hansch analysis:  a.
  guide to structure-activity correlation for the medicinal chemist.
  Adv. Drug Res.  621-77.

Weaat, R.C. (ed.).  1974*  CRC Handbook of chemistry and physics, 54th
  edition.  C1C Press, Cleveland, Ohio.  p.D-162 to D-188.

                   *•  COyCLlJSIQES. AFP RECOMMENDATIONS

4,1  Conclusions

     The literature search In support of these efforts, comprising both
automated and manual searches In addition to direct contacts with current
workers In the field and chemical manufacturers, was thorough.  This
conclusion has been corroborated by technical reviewers from Industry,
universities, and the U.S. Environmental Protection Agency.

     In general, the behavior of moat of the metals but only * few of the
organica has been extensively studied in natural surface water systems.
The well-studied organlcs Include PCBs, DDT, chioroethene (vinyl chloride),
pentachlarophenol, and bis(2-«thylhexyl) phthalate.  Despite the relatively
large amount of effort expended on these compounds, quantitative models of
their environmental transport and fate are, at present, somewhat inconclu-
sive.  For the rest of the organic pollutants discussed in these chapters,
data sufficient to quantitatively define the principal transport and fate
processes were not available, with the exception of a few specialized
cases.  The information obtained for most pollutants, however, was
sufficient to qualitatively Identify the importance of the discrete
processes considered.

     Conclusions regarding aspects of environmental behavior of pollutant
groups are listed below:

     1*  Metals.  Transport and fate of metals are, in general, controlled ,
         by sorptlon processes In the sediments.  The metal-organic
         relationships, both in the sediments and in the water column,
         increase in importance as the organic content Increases and
         strongly affect «etal transport in polluted areas, for
         example, industrialized Urban areas.  These metal-organic
         relationships have not been thoroughly studied.

     2,  Chlorlnatfd Fe»cicides.  The fate of most chlorinated pesticides
         in the aquatic environment will be determined by sorptlon,
         volatilization, and blotransfonnation.  Very little reliable
         quantitative data are available for assessing the half-lives of
         these processes, however.  Except for the rapid hydrolysis of
         heptachlor, there is Insufficient and sometimes contradictory
         Information to support a conclusion that chemical transformation
         processes will b* important la aquatic environments.

     3.  Halogena t e d Al1pha11c Hydrocarbons.  Transport and fate of
         halogensted aliphatics are generally dominated by volatilization,
         so that their ultimate fate typically involves atmospheric
         processes.  $ost of the data available on rates of volatilization
         are valid only for comparisons (relative races) between compounds.

     4»  Halogenated Ethers andSelected Monocyclic Aromatlcs.  For several
         groupings of compounds, Including the haloethers, phenols,
         phthalates, and benzene and its derivatives, it is typical that
         only one or two members of each group have been studied exten
         •ively.  Extrapolation of environmental behavior to other members
         of the groupings was performed when feasible, but wide variations
         In chemical and physical properties Unit the accuracy of such
       . extrapolations.

     5.  golycyc11cAromaticHydrocarbons.  The behavior of the polycyclic
         aromatic hydrocarbon* was found to be a function of the number of
         ring* present, and these compounds were grouped accordingly.  In
         general, it appears that the important processes for these
         compounds are adsorption onto particulates, sedimentation, and
         subsequent blodegradation.

     6*  Nitrophenola. Sltrosaalnes,and Miscellaneous Compounds.  For
         several groups of compounds, almost no environmental data were
         found.  Tnese groups include the nitrophenols, the nitrosamines,
         and various other nitrogen-containing coopounds.  In many of these
         cases, environmental scenarios could be developed from research
         studies conducted for purposes other than the determination of
         environmental behavior, and the scenarios could be used Co
         Indicate the relative importance of the processes considered.

    An annotated table of conclusions for each of the 129 priority
pollutants is presented In Section 4.3.

4*2  Recommendations •

     Based on the conclusions presented above, the general recommendation
is Bade that the state of knowledge of environmental transport and fate of
these (and other) pollutants should be improved through studies directed
specifically towards definition of their environmental behavior.  In adui-
tion, unavailable physical constants and process rates should be deter-
mined.  Some of the more Important areas of needed research are listed ,

    1.  Metal-organic relationships and their effects on metal transport in
        surface wafers should be defined.

     2.  Definitive,studies on the Importance of volatilization of
         pollutants from water relative to the other transport and
         fate processes should be conducted for specific pollutants and

     3.  A detera!nation should be made of the inportance of toxjc products
         resulting from reactions of the pollutants in the environment and
         In ^ater treatment facilities.

     4.  Further work is recommended on the development and validation of
         fate models, so that the data generated by the above studies can
         be used effectively In predicting the environmental fate of

4,3  Sunaary of Conclusions

     An annoted table of conclusions for each of the 129 priority pollu-
tants, by chemical group, is presented in Table 4-d.  Two ratings are
presented for each compound and related fate and transport processes.  The
first rating (Yes, No, or Uncertain) is a statement of importance, and the
second Is a numerical rating dealing with the depth of available supportive
data.  These ratings were assigned by the major author of the chapter on
each compound, or group of compounds, and Involved a certain degree of
subjective judgment.  Th« application of these ratings may therefore vary
slightly with each chemical group.

     Table 4-1 condense" the large amount of information reviewed in each
pollutant chapter into a single summary line statement about the water-
related fate and transport of 129 priority pollutants.  For a specific
compound, therefore, the reader is encouraged to consult the chapter de-
voted to that compound before making final judgments.

                                                                          TABLE 4-1

                                                                  A.  Metal* and Inorganic*
Aat tatony
the Process Important
Aquatic Transport?
(See Kay Below)
Volatilisation Sorptlon
YES( 1 )
YES( 1 )
la the Process Important In
Determining Aquatic Fate?
(See Key Below)
YE 5(2)
Special Ion Bloai-C'mulat Ion
YIT3( 1 )
UNCT( 1 )
UNCT( 1 )
Biotransforaat ion/
  For each chcnlcal and related proceaa, two rating* *r«  presented.   The first 1* * statement of laportance and can b« one of the following:   YES,  NO,  or
  t'NCT (for uncertain).  The second 1* a numerical rating, dealing with  available supporting data,  explained below:
  (I).  There are environmental data available to support  this  conclusion.
  (2).  There are no direct conclusive environmental data; sone  laboratoty  data can be extrapolated to aupport conclusions.
  O).  TTitie are no supporting data available; evidence  la  drawn  fro*) theoretical calculations, estlauites, result* for similar chealcals, and  Inferences.

                                                                      TABLE 4-1 (Omt.)
                                          SUNNAKY  Of  CONCLUSIONS fOH THS TKAN$PO*T ANU JAW Of PRKWITY FOU.UTANTS:
                                                         ».   Paetlcldee, Kite, *nl faUted Coapounde





Ac fold In
Endoaulfan aod Endoaultan
Endrln and bidrla Aldehyde)
Heptachlor Epoxlde
(a, 8, S Isowtr*)
t -H« »«^h loroc yc lolxiue«
P«lychlorla*c*d tlplMnyl*
th« Proc**» laparoat
Aquottc Transport?
(S«« by kluu)
VoUUltMtlon SotpClaa


t£S( 1 )



UNCT( 2)



U th« Proce§» laportant In
Duteralnlng Aquatic rat*?
(Sea Kay talow)






Ktoaccuaulat Ion
                                                                      TABLE 4-1 (Cant.)

                                          SIMUJtY OF CONCLUSIONS FOX THE THAHSPOtT AMD FATE OF PtlORlTY  POLLUTANTS:

                                                            C.   Halogen*ted Aliphatic Hydrocarbon*
Dt c h I o roew t haue
Trie ul oromelhane
Te t rachlor»**t han*
1, '-Dlchloroethane
1.1. 1-Trlchloroethan*
1. 1,2-Trlchloroethao*
1, 1,2,2-TetrachloroethM*
1. t-Dlchloroethen*
1 . 2-ttans-DlcMucoethene
1 , 2-Dtchloropropan*
1, }-Olchloroprop»ne
Broavuew t han*
Dlchlorodlf luorosMthao*
Trtchlorotluoroaw thane
tea) Procese Important
Aquatic Transport?
(See Key Below)
Volatilisation Sorptlon
YES (2)
YES (3)
I* the Process Important In
Determining Aquatic Fate?
(See K*y Below)
UNCT( 3)
UNCT( 2 )
Blot ransfonut Ion/
UNCT( 3)
Ui»CT( 3)
  For each c ealcal and related pruce»a.  two  (atln«a  are  prevented.   The  flret  U • atatee*nt of laportance and can be one of the following:  YES,  NO, or
  UNCT (for uncertain).  The eecond la a  auaartcal  rating, dealing with available (upportlnt data, explained below:
  (1).  quantitative (rate conatanta, half-llvea) data are awallrjle  to cupport conelualone.
  (21.  Qualitative description only; there arc no  direct environmental data, however,  aoew laboratory data can be extrapolated to lupport conclusions.
  (3).  Tttere are no eupportlng data available; evidence  la  drawn  troai theoretical calculation*, eeclMtee. results for similar chealcaU. and  Inferences.

                                                                             TABLE 4-1 (Coot.)

                                                SUMHAKY Of CONCLUSIONS rot THE TRANSPORT AND FATE Of HI I OK ITT  POLLUTANTS:

                                                                           D.   H»lo|*n«t«8


81 «( 2-chloroethyl )ether
r i(2~chluroltoprouyl)*tn*r
2~Chloroethyl vinyl ether
4-Chloroph*nyl phenyl ather
4-ftro*ophenyl phenyl ether

the Process laportaot
Aquatic Transport?
(See Kav Below)

Volatilisation Sorptlon
H0( 3)

MHCT( 3)
Is the Process Important In

UNCT( 3)
(See Kt

UNCT( 1 )
NO( 3)
UNCT< 3)
Aquatic Face?
> Below)

Bl uac c uau 1 a t 1 on
UNCT( 1 )
UNCT( ))

Biotranafortaat ion/
Blodegradat Ion
        Kut «acl> >'h«»Uj] and  related  pcocvca,  tuo ratlnga ar« pt««ent«J.  Th*  firm  la  a  atat«**nt  of laportance and cun be j;>* of  the  tolluwlng:  YES,  NO,  or
        UNCT (fur uncertaluj.   Th*  »«cond  Is a  nuawrlcal rating, dealing with avallabi*  lupiiortlng data, captained below:
        (I).  i,idntttatlv*  (rate  constants,  half-llvea) data ar« available to aupport  conclualona.
        (2).  Oualttdt lv« description  only;  there are no direct environmental data, however,  aoai* laboratory data can h« extrapolated  to  support  conclusions.
        (3).  Fhere are no"~3«^portIng  data available; evidence la drtwn  fro« theoretical calculations, e«ttsjat««, teaulta for  slallar  chemtcalii,  and Inferences.

                                                                           TAILS *-l (Coot.)
                                               SUMMARY 0» CONCLUSIONS rot Tttt TRANS PV-KT AMD FATt 0V MtlOKITY  POLLUTANTS:
                                                                        g.  Neiwcyclic  AroaMtica
Ma. CMaical
1 , 2-Dlchlorob*n»«n*
1, l-Ulchiorob*ni«o«
1 , 4-Otchlot aben**o«
1,2. 4-Tr ic hlorobenMM
2, t>~Dinltrotolu«a*
2 , « - 1>U hlurophanol
UHCT( 3)
tloaccuwilat ion
N0( 3)

UNCT< 1)
NU( 1 )
     Key:                                             -                           -
       For  edch vhealcal, Jiid related proceed, two ratings are presented.   The  tir*t  la a *Cat**Mnt of Importance and can be on* of  th«  following:  YES,  NO,  or
       UNCT (lor uncertain).  The second is a nuawrical rating, dealing with available supporting data,  explained below:
       (1).  tjuj'it U4t tve (rate constant*, half-live*) data are avaliabl*  to support conclusion*.
       (2).  ^idlltjtive description only; there ar« no direct environmental data,  however,  suaw laboratory data can b« extrapolated  to  support  conclusions.
       (1)4  There  ire  uo sujiportlng data available; evidence la drawn  fri>«  theoretical calculation*, estiMtvs, result* for aiallar  chemlcdlK,  and  Intevcucea

                                                                      TAU.K  *-l  (Cent.)
                                          SUMMHY Of CONCLUSIONS H* THE TRANSPOKT AND KATE Of MtKMUTY POLLUTANTS:
                                                  K.  Ptithalat* Eitvr* and PolycytUc  AroautU  Hydrocarbon*
Nu. CbcDlcal
94 Phthalate Enter*:
• Ul*ethyl
• DUthyl
• M-n-butyl
• Di-o-vctyl
* 5ts( 2-«)Chylhe»yl )
* Butyl bentyl
Pulycycllc Aro«atlc
Hydroc<*cbon« :
9S • Vrt>naphtheue
• Acenaphthylen*:
• Fluoreix
• Naphthalene
96 • Anthracene
• Fluoranthene
• hi«;n { gh 1 Jpvryleue
a B«Mtzo(
N0( 3)
UNCT( 3)
UNCT( 3)
UNCT( 3)
Procea* Important
auat Ic Traoaport?
(See Key klaw)
Sorpt Ion

YES( 1 )
UNCT( 3)

YKS(2) -
Downat reaa

YEi( 1)

UNCT( 3)
UNCT( 3)
UNCT( 3)
la th« Proceaa Ta^K>rtant In
Ditarailnlng Aquatic Fate?
(See Key Below)


YES( 1 >


N0( 3)
N0( 3)
& i oac c uau 1 a 1 1 on

UNCT(2) "

NO; 2)
!IMCT( 1)
UNi U ))
Blot ran* format ton/
ttiode^raddt iun

UMCT( 2 )
UNCT( 2 )

  for ea< h chenical and related  process,  twu  rat ings are presented.  1>i** first  is a  statement  of  importance and can be one ot the  tul losing;  YKS,  Nt>,  or
  UNCT (tur uncertain).  ft- • second  U  .»  nuawrical  rating, dealing with available supporting data,  explained below:
  ( 1 >.  t^ant Itative (rate com»taiu;., tialf~livti«)  data are available to support conclusions.
  ( i ) .  tjuai 11 Jt ivt de script *,'»ti  onl y; £ here are  no direct environment »1 data , however .  so»e  laboratory data can be cxt rapolated  to  support  conduit ions ,
  (3).  [he-tt are ru» supporting  data available;  evidence i* drawn fro« theoretical calculations,  estimates, results tor similar  chemicals,  and  inferences.

                                                                          TABU 4-1  (Cont.)


                                                             C.   Nitrosamlnes and  Miscellaneous Compound *




Dlnothyl nitrosamlne
Dtiinenyl nit co-dial ne
Di-n-propyl nltrosamine
3, V -Dlchlorobcnildlne
1 , 2 -Di pheny 1 hyd rax 1 ne

NO( 1 )

Is the Process Important
Aquatic Transport?
(See Key Below)

'•"Ion Sorptlon
) N0<3)


Is the Process
Important In

Determining Aquatic Fate?

(See Key


Hydrolysis Bloaccumulatlon

Blot rans format ion/
      For each ch.-ralcal  and  Delated  process,  two  ratlngn  are presented.   The first is a statement of Importance and can be one of the following:   YES,  NO, or
      UNIT  (fur  uncertain)'.   The  second  ts  a  numerical  rating,  dealing with available supporting data, explained below:
      (I).  Quantitative (rate  constants,  half-lives)  data are  available to support conclusions.
      U).  Qudlitdilve  description  only;  there are no direct environmental data, however, some  laboratory data can be extrapolated  to  support  conclusions.
      .  There  ..it no supporting  data available; evidence is drawn from theoretical calculations, estimates, results  for  similar  chemical*,  and inferences.

          Cliapters 5-19

                               5.  ANTIMONY
5-1  Stacement	of Probable Fa ta

    The fate of antimony in the. aquatic environment is determined by a
number of factors including pH, Eh, sorptlve interactions, and biologically
mediated methylation.  Due to the relatively high solubility of the anti-
monlte and antimonate ions, most of the antimony introduced into the aqua-
tic environment is probably transported in solution to the oceans.  Copre-
cipitation with iron and aluminum oxides, adsorption by mineral surfaces.
and bioaccumulation may, however, be responsible for removing some antimony
from solution.  Biologically mediated methylation or reduction to stibine
(SbHj) may occur in reducing environments, resulting in remobilization of
antimony.  The relative importance of each of these orocesses varies widelv
among watersheds; but, in general, transport of dissolved antimony to the
oceans is the most probable dominant fate.

5.2  Identification ^geochemistry of Antimony

    Antimony is an element occurring in concentrations of about 0.2-0.5 ^m
in the earth's crust (A.D. Little, Inc. 1976).  Important minerals of anti-
atony are the native element Sb; stibnite Sb2S3j Kermesite Sb2S20; serarroonite,
SktOjJ Jamesonite, 2PfaS«Sb2S3; boulangerite 5PbS'2Sb2S3; and the sulfanti-
morlds of copper, silver, and nickel.  Important artificial compounds include
stibine, SbH3 (a noxious poisonous gas), th* chlorides SbCl3 and SbCls, the
sulfides Sb->Si and Sb->Se, and the oxides Sb2Ch *«d
    Antimony has chalcophilic properties, and thus combines readily with
sulfides.  Antimony shows no marked preference for mafic or silicic rocks.
Antimony becomes enriched in the early stages of magmatlc differentiation
in sulfldc bodies.  la addition, antimony accumulates in late-stage grani-
tic pegmatites, together with nlobrlum and tantalum oxides, in granodio-
rites, and hydrothermal sulllde deposits.  Antimony is also present in
galena (a lead ore), in amounts up to 1%, where it replaces either the lead
or the sulfur.  Antimony may also substitute for arsenic in many minerals.

    Little is known about the behavior of antimony during weathering.  The
antimony sulfides may be converted to the corresponding oxides, and it
probably occurs in both hydrolysate and oxidate sediments.  Antimony may
accumulate with heavy elements in carbonaceous shales 07 become sorbed on
clays and hydrous cxides.  Thus, it may be enriched in manganese oxide  '
sediments and black shales.

    Antimony has an atomic  number of 51 and  an atomic  weight of  121.75.   In
its compounds, It has a valence of +5,  +3, pr -3.   When in the +3 state  it
has metallic characteristics.   Antimony's chemical  prcperties are analogous
to thosa of arsenic, rfhich  is  directly  above it in  the periodic  table, and
forms compounds with a number  of other  elements,  such  as oxygen,  hydrogen,
sulfur, and the halogens (Weast 1977).

    Antimony's CAS nunber is 7440-36-0; its  TSL number is A167-6664.

5.3  Summary of Fate Data

    5.3.1  Photolysis

         Antimony compounds may undergo photochemical, reactions,  but  none
of these appear to be important in determining aquatic fate.  Stibine
(SbHj) reacts with sulfur at 100'C in the presence  of  light to form anti-
mony aulfide.  finder the same  conditions, stibine reacts with selenium  to
form antimony selenide (A.D. Little,-Inc. 1976).  .Antimony trioxide can  act
as a photocatalyst for the1 oxidation of organic matter by ultraviolet
light, producing organic peroxides (Mar khan  ejt _al.  1958.)

    5.3.2  Chemical Speciation

       ',  The chemical properties of antimony resenble  those of bismuth and
arsenic.  Antimony loses its 2(s) and 3(p)  electrons readily, and"may exist
in the oxidation states -J, +3, +5, and 0, although Che existence of  simple
     or, Sb"1"-* ions is improbable.  In reducing environments, stibine
      ' may be formed.  Stibine is a gas at room temperature, and  it is
quite soluble in water (5,000  ag/1).  It is  not stable, in aerobic waters
and is hydrolyzed to form the  oxide.  The formation of, stibine in bed sedi-
ments (which usually offer  a reducing environment)  may result in  remobili-
zation of antimony which had been removed from solution by adsorption, co-
precipitation, etc.

         Under moderately oxidizing conditions, antimony has a valence
of +3,•&nd it is found in solution as the hydrated  tribxide,
           .  Unlike arsenic ,  which forms arsenious acid
         under mildly oxidizing conditions,  the lower  valence acid  of
antimony is unknown; however,  the antimonites are well defined salts
(Cotton and.Wilkinson 1972),  The higher acid, HjSbC^, does exist,
although it does not appear to dissociate completely to the Sb04~^  ion
even in tUe most alkaline conditions (Cotton and  Wilkinson 1972).  This
form  of antimony, In which the element exhibits  a  valance of +5, is  found
only in highly oxidizing environments.

         Antimony salts, added to aqueous media, are hydrolyzed to the
oxide1 or antimonic acid forms.  Although the hydrolysis products ate usu-
ally less soluble than the original sales, the solubility is still suffi-
cient to keep antimony in solution, except for cases of extremely heavy
loading.  When the system is no longer saturated, any antimony that pre-
cipitated out as the oxide will go back into solution.

         The oxides, i.e., antiaonite and ancimonate forms, which are
stable in the redox range typically observed in natural waters, all have
sufficient solubility to keep antimony in solution at the levels of con-
centration normally found in natural waters.  The ealciura, magnesium, and
sodium salts of the antimonaces and antimonites are probably not a signifi-
cant control on antimony solubility; but, by analogy to arsenic, some of
Che trace metal compounds may exhibit limited solubility.  Thus, one -tould
speculate that the presence of heavy metals (e.g., copper) in solution may
reduce the mobility of antimony.

    5.3.3  Volatilization

         Antimony may be volatilized when in the form of stibine or its
methylated derivatives,  Stibine can be formed by reduction of antimony in
the sediments.  Although bioaethylation of antimony has not been demon-
strated, there are no obvious thermodynamic or kinetic obstacles (Parris
and Brinkman 1975, 1976).  Moreover, the elements Sn, Pb, As, Se» Te, which
surround antimony in the periodic table, are subject to bioaethylation,
suggesting that biomethyiaUon pathways could exist for antimony {Parris
and Brinkman 1975, 1976).  Stibine is rapidly oxidized in air or oxic water
to form Sb203.  It is likely then, that most of the stibine formed in
the sediments reacts in the water column to produce the oxide, resulting in
remobilization of antimony.

         The methylated forms of antimony are also subject to oxidation.
Parris and Brinkman (1976) estimate the rate of oxidation of trimethyl-
stibine as greater than IQ~2 H~^- s~*.  The product of this reactioa,
(CH-j)jSbO, is much more soluble than trimethylstibine; and, therefore,
this oxidation would probably tend to reduce volatility.  The rapid rate of
oxidation implies that, if trlmethylstlbine is formed in natural systems,
auch of It would be oxidized before it volatilized and only a small amount
of the volatile antimony compounds formed by either abiotic or biotic
mechanisms would be liberated to the atmosphere.

    5.3.4  Sorption

         The extent to which sorption reduces the aqueous transport of
antimony is unknown, but it Is clear that sorption processes are normally
the most important mechanisms resulting In the removal of antimony from
solution.  Antimony apparently has an affinity for clay and other mineral

surfaces.  Coprecipication of antimony with hydrous iron, manganese, and
aluminum oxides may exert' a significant control on antimony mobility in
areas where there Is active precipitation of these metals.  Cracelius e_t
al. (1975) studied the oLetal concentrations of PageC Sound, Washington, and
found that, in uncontamlnated areas, most of the antimony In the sediments
was bound to extractabl*: iron and aluminum compounds.   Antimony bound in
such forms would probably be susceptible to remobilization via bioaccuow-
lation, reduction, or bi.omethylatlon.  On the other hand, antimony in
heavily contaminated areas was found aainly in stable, usiextractable forms.
These forms might include the oxide or insoluble metal antimonates or anci-
moriites.  Crecelius re ill, (1975) found that less than 102 of the antimony
in sediment.3 Iron both the contaminated and uncontazninated sediments was
bound to easily oxidizable organic matter.  Since antimony has an'anionlc
character in aqueous solution, it probably has little  affinity for cooplex-
ation with nu»lc and fulvlc acids, which ar« important completing agents
for metals.

         Maxfield «£ £1. (1974), In their study of heavy metals in the
Coeur d' Alene River of Idaho, found that antimony was concentrated and
evenly distributed throughout the sediment.  They suggested that, although
antimony is being adsorbed by all types of particulate matter, it is not
being strongly bound in the sediments.  The high levels of antimony that
are characteristic of this region are due to former smtimony mining acti-
vities; and the antiaony found in the sediments probably entered the sys-
tem, not in the dissolved state, but rather on particulate matter.   Max-
field jet al, (1974) concluded that the results suggest that antimony is
leaving the sediments; but since this is a diffusion controlled process, It
proceeds slowly.

         Strohal «£_•!. (1975) Investigated antimony in the sediments of
the North Adriatic Sea.  They found that the fixation  rate of antimony on
various inorganic particles is rather small.  They found, however, that
antimony could be accumulated by organic matter, especially humic acids.
Unfortunately, sorptlon processes of antiaony have not been studied in en-
ough detail to quantify the role of sorption In its aquatic fate.

    5.3.5  B1oa ccumulatiion

         Antiaony is only slightly bioaccumulated and  has been little
studied in aquatic organisms.  Leatherland JK _al. (1973) found low levels
of antimony in various fishes and Invertebrates collected off the northwest
coast of Africa; antimony was generally present In higher concentrations in
invertebrates than in fish.  Aquatic organisms from the Danube River and
Danube Canal in Vienna, Austria, were found to contain only background

levels of antimony (Rehwoidt e_t a_l_. 1975).  Similar results "were obtained
in clams, mussels, and shrimp by Sertina and Goldberg (1972).  Table 5-1
summarizes known bioconeentration factors for antimony.
                                 Table 5-1

                   Bioconcentration Factors for Antimony

    Taxon               Bioconeen_cration Facto£a         Rffer«nce

Freshwater fish                      40               Chapman e£ al. 1968

Freshwater invertebrates         16,000               Chapman et, al. 1968

Marine fish                          40               Chapman «£ _al. 1968

Marine Invertebrates             16,000               Chapman et al. 1968
a.  Concentration factors are defined by the ratio of the
    concentration of the element in the organism in ppm (wet weight)
    divided by the concentration of the elemenf in water (rpa>)«
    5.3.6  Bio ' t rans f ormat i on

         It has been reported that a species of bacteria, Stibio'bacter
senaraoritii, utilizes the energy released by metabolically induced oxida-
tion of antimony (Lyalikova 1974), but tha distribution and ecological
importance of this organism is unknown.  As previously mentioned, bio-
ttethylatlon of antimony has not been demonstrated, but it is thought that
this process could occur in the environment (Parris and Brinkman 1976).

5.4  Oata Sumnary_

    Although the aquatic fate of antimony has not been comprehensively
studied, it appears that most of the antimony Introduced into aquatic en-
vironments is probably transported ^in solution to the oceans.  Sorption
processes act as temporary sinks for antimony, but bioaccumulation, re-
duction to stibina, and possibly hiomethylation may act ,to remobilize anti-
mony in bed sediments.  There is a possibility that heavy metaJs in solu-
tion could react with antinonite (H^SbQj, H^SbOj", HSbOj
Sb03~3> or anTimonate (H3Sfe04, J^SbQ^-, HSbC^'2) species
to form insoluble compounds, greatly inducing the mobility <"f antimony, but
the importance of such reactions is unknown.  The aquatic fate of antimory
is summarized in Table 5-2.


Pho to 1 ys i s




Biotrans formation3
                                 Table 5-2

                    Summary of Aquatic Pate of Antimony

Not Important:.
Confidence of
Antimony is present as the
soluble oxide or antimc^ite
salt in tnosc natural waters.
In reducing environments, vola~
til» SbH3 say be formed.  Most
species of antimony are soluble
and mobile in the aquatic environ-
ment .

Important where SbHj Is stable.

Antimony is adsorbed by clays
and organic materials.

Very slight,

Biotnethylation may occur.



    All of the noted environmental processes are important; however, their
    relative 'Importance wirh respect to each other Is uncertain for
    determining final fate.

 5«5  Literature Cited

 A.0. Little, Inc.  1976,  literature study of  selected  potential
  environmental contaminants; antimony and Its compounds,  pp.5-13.
  PB 251-438.  147p.

 Bertine, K.R, and 8.0. Goldberg.  1972.  Trace elements in clams, mussels,
  and  shrimp.  Hanoi, Qceanogr.  17(6) :877-884.

 Chapman, W.H., H.L. Fisher, and M.W. Pratt.  1968.   Concentration factors
  of chemical elements in edible  aquatic 'organisms.   Lawrence Radiation
  Laboratory.  Liveraore, Calif.  UCHL-50564.  ,46p.

 Cotton, F.A. and G. Wilkinson,  1972.  Mvanced  inorganic chemistry.
  pp.367-402.  Inter-Science Publishers, New York.   114Sp.

 Crecelius, B.A., M.H. Bothner, and R. Carpenter.   1975.  Geochemistries  of
  arsenic, antimony, mercury, and related elements in sediments of  Puget
  Sound.  Environ. Sci. Technol.  9(4);325-333.

•Laatherland, T.M., J.D. Burton, P. Culki$, M.J.  McCartney, and R»J. Morris.
  1973.  Concentrations of some trace metals in  pelagic organisms and of
  ••rcury In northeast Atlantic Ocean water.   Deep-Sea  Res.  20:679-685.

 Lyaiikova, U.S.  1974,  Stibiobacter senaraontii, a  new microorganism
  oxidizing antimony.  Mikrobiologiya~4Tf6)7441-948.  (Abstract only).   CA
  1975.  82;l35404y.

 Markhan, M.C., M.C. Hannan, L. Liu, C. Coffey, and B. Jones.  i?58,
  Photoch«anical properties of antimony trioxide. J.  Phys. Chem.

 Haxfield, D., J.M. Rodriguez, M.  Buettner, J.  Davis,  L.  Forbes, R.  Kovaes,
  W. Russel, L. Schults, R. Snith, J. Stanton, and C.M. Wai.  1974.  Heavy
  st«tal pollution in the sediments of the Coeur  d' Alene River delta.
  Environ. Pollut. 7:1-6,

 Parris, G.E. and f.E. Brinkaan.   1975.  Reactions which relate to the
  «nvironBental mobility of arsenic and antimony.  1. Quarternization of
  trincthylarsine and triawthylstibln*^  J. Org. Chen.  40:3801-3303.

 Parris, G.E. and F.E. Srlnkaan.   1976.  Reactions which relate to the
  environmental mobility of arsenic and antimony.  II, Oxidation of
  trimethylaraine and trimethylstibine.  Environ.  Sci.  Technol,
  10(12):1128-1134.               /

Rehwoldt, R.,  D. Karlmian-Tehen^ni, and H. Altmanni  1975.  Measurement and
  distribution of various heavv raetals tn the Danube Xiver and Danube Canal
  aquatic communities in the vicinity of Vienna, Austria.  Set. Total
  Environ. 3:341-343.                ^

Strohal, P.,  D. Huljar,-S.  Lubic, and M. Picer.  1975,   Antimony in the
  coastal marine.environment, North Adriatict  'Estuarine and Coastal Marine
  Science 3:119-123.

Weast, R.C. (««(.).  1977.  Handbook, of chemistry and physics, 58th
  edition.  CRC Press, Cleveland, Ohio.  239dp.

                                6.  ARSENIC

6.1  SCateaent of Probaole Fate

    The face of arsenic in the aquatic environment depends largely on pre-
vailing pH and Eh conditions.  Arsenic is extremely mobile in the'aquatic
environment and cycles through the water column, sediments and biota.  Al-
though the equilibrium cnemistry of arsenic has been extensively discussed
in the literature, no information regarding the kinetics of arsenic reac-
tions in the environment was found.  It appears that, in most cases, the
sediments and the oceans are the primary sinks for arsenic in the aquatic

6.2  I den t i f i ea_t i on - Qeocheitist ry o .f_A_r senie

    Arsenic is considered to be a rare but ubiquitous element in the
earth's crust,  the average abundance of crustal arsenic hai been established
at 5 ppm (Weast 1977). .Arsenic is the third member of group V3 of the
periodic system, which also includes nitrogen, phosphorus, antimony, and
bismuth.  In some of its chemical reactions, arsenic behaves ^uch like
phosphorus and antimony.

    In the natural environment, four oxidation states are possible for
mrsenic:  the-3 state, the metallic (0) state, and the +3 and +5 states.
The Metallic state is not uncommon for the element in certain types of
mineral deposits.  The +3 and +5 states are common in a variety of complex
minerals and in dissolved salts in natural waters.  The --3 stae is present
in gaseous AsHj (arsine) which nay fora under some natural conditions.
Because of Its multiple oxidation states and its tendency to form soluble
complexes, the geochemistry of arsaenic is Intricate and not well character-
ized.  The element most commonly associated with arsenic in nature is sul-
fur (Boyle and Jonasson 1973).

    la all, there are 100 or more arsenic bearing minerals known to occur
in nature.  The principal arsenic minerals are arsenopyrlte (FeAsS), nic-
collte (NiA»), cobaltite (CoAsS), tennantite (Cuj^As^Sjj), enargite
           native arsealc (As), orpiaent (AS2S;j), realgar (AsS), prousite
           scorodite ((Fe,Al)(As04)-ZH^O), bendantite (PbFe3(As04)(S04){QH)6)»
ollvenite (CujAsO^OR), oimetite 





l_ _



J f ^xv^




1 t








                  Figure  6-1   The generalized geochenical cycle  of  arsenic.

                              Joiiasson (1973) .
Fr< 11 Boy It unJ

    Arsenic haa no aqueous cationlc chemistry other Chan In organic quar-
teraatry salts (Cotton and Wilkinson 1972;.   The oxo acids, arsenlous acid
(H^AsOy) and arsenic acid (l^AsC^), are the prevalent forms of
arsenic in aerobic waters.  Arsenic can fora complexes with a number of
organic compounds, most of which increase solubility.

    The atomic number of arsenic is 33; its atomic weight is 74,91; density
is 5.72 (20*C); melting point (28 atm.) is, 817"C; and boiling point is
613 *C (Weast 1977).

    The CAS number of arsenic is 7440-38-2; the TSL number Is A-1418-3227.

***''  Summary of Fate Data

    6.3.1  Photolysis

         So evidence was found that photolysis is an important mechanism in
determining the fate of arsenic compounds.

    6.3.2  Chemical Speelation       .                .                      '

         In aquatic systems, arsenic has an unusually cooplex chemistry
with oxiJation-reduetioa, ligand exchange,  precipitation, and adsorption
reactions all taking place.   Arsenic is stable in four oxidation states
(+5, +3, 0, -3) under Eh conditions occurring In aquatic systems.  Arsenic
octal occurs only rarely and the -3 oxidation state Is  stable orly at ex-
tremely low Eh values.  Since the valence state of arsenic is extremely
important in determining toxicity (the +3 state Is ouch more toxic than the
+5 state; National Academy of Sciences 1976) as well as ecwplexation be-
havior, the chemical speclatlon of arsenic  is very important when consider-
ing its aquatic fate,

         Wggetnann (1978) exaained trie typical concentrations of major and
minor Ionic constituents £n freshwater systems in an attempt to find the
possible controls on totii dissolved arsenic in freshwater.  He selected
four metals (3a, Cr, P«, Ca) as possible controlling factors and studied
their metal arsenates aor« closely In the laboratory.  Barium Ion, at typi-
cal freshwater concentrations, was the most likely freshwater constituent
that would be capable of holding total dissolved arsenic to rather low con-
centrations,  ftased on these studies, an Eh-pH diagram (Figure 6-2), which
summarizes theoretical arsenic speciation in freshwater environment a, was

Figure 6~2   Eh-pH diagraa for arsenic at 25*G *nd I atit. pressure, showing
             the fields of stability for the moat important arsenic species
             tn th* presence of 10"" *M of total arsenic, 1Q~%| of total
             sulfur and 2.2 x 10'^.M of total barium,  0«sh«d lines define
             domains for §p«cie3 enclosed in parentheses.  From Uagemann

         Andreae (1978) analyzed seawater from the Southern California
coast and terrestrial watars from several locations in the United States
for four arsenic species:  arsenite, arsenate, monomethylarsonic acid, and
diaethylarslnic acid.  Generally, arsenate was dominant, but in waters of
the photic zone, the other species were found in significant concentra-
tions.  'A positive correlation was evident between the concentrations of
arsenite and methylated arsenicals and biological activity.  These results
indicate that the speclation of arsenic in natural waters is significantly
influenced by the biota.  In the waters below the euphotic zone, arsenate
concentrations increased with depth, suggesting regeneration from biologi-
cal material.  Inasmuch as arsenate is the thermodynamlcally stable form of
arsenic under the conditions prevalent in most natural waters, the non~
equilibrium species reflect the biological activity in natural waters.

         These results were confirmed by the work of Waslenchuk and Windon
(1978) in estuaries and Haslenchuk (1979) in rivers.  Waslenchuk and Windoo
(1978) found that in estuaries the only detectable species was arsenace
which remained in solution as fresh and salt water mixed. ' Complexes oc-
curred between arsenic and low molecular weight dissolved organic matter.
These complexes presumably prevented adsorptive and copreclpitatlve inter-
actions with the sediments and allowed the arsenic to travel to the ocean
in a dissolved form.  Arsenic which enters the estuary associated with
part icuiates, however, apparently remained so and accumulated in the sedi-
ments.                                                  !  ,

         Waslenchuk (1979) found that the levels of dissolved arsenic in
rivers in the southeast of the United Stat»s are controlled by the avail-
ability of arsenic, by rainwater dilution, by the extent of coffiT-lexation
with dissolved organic matter, and perhaps by the metabolic activity of
aquatic plants.  Arsenic complexation by dissolved organic matter prevents
adsorptive interactions between the arsenic and solid-phase organic and
inorganic materials.  Despite high arsenate solubility, arsenate concentra-
tion is limited to levels below saturation, due to raactlons which remove
the free arsenate ion from solution.  The particulate arsenic load may be
as important as the dissolved load with r**p«ct to material transport in
rivers.  Only the dissolved load, however, is delivered to the ocean.  It
appears further that those biologically mediated reactions which result in
arsenic species disequilibrla, in the ocean and lakes, have an insignifi-
cant effect on arsenic speclation in rivers.

    6.3.3  Volatilization

,         Volatilization of arsenic nay be a significant process in extrerae-
ly reducing environaeits where arsine (A»H3) is formed, but under normal
circumstances, it is not an important mechanism in determining the fate of
arsenic after the element's introduction to the aqueous environment.

 Araine la  probably oxidized  rapidly in aerobic  waters or the atmosphere
 (Parris and  Srinkman  1?75,  1976).

         Methylated arsine  derivatives may have acre  potential for volati-
 lization.  Trlmethylarsine  is quite volatile at room,temperature (vapor
 pressure 322 Corr) and Is oxidized to more soluble  products very s owly.
 Parris and  Brinkaan (i975)  reported a rate constant of less than
 10~2M-is"l for oxidation of  (CHj^As by dissolved oxygen.   In the
 gas phase, the rate constant for oxidation by oxygen s»as 10~"M~-''S~^.
 Thus,  trlmethylarsine can travel considerable distances without undergoing
 chemical change,  even in aerobic systems.

     6.3.4  Socption

         Cycling  of arsenic  in the aquatic environment is dominated by
 adsorption and desorption to sediments.  Arsenic atay be sorbed onto clays,
 aluminum hydroxide, iron oxides, and organic material (Ferguson and Gavis
 1972;  Jackson _e_t  _*!«  1978). >  In areas where phosphate minerals occur, arse-
 nate aay isoatorphously substitute for phosphate (Hem  1970).  Under most
 conditions,  copreci'pitation  or' sorption of arsenic  with hydrous oxides of
 iron is probably  the  prevalent process in the removal of dissolved arsenic.
 The oxyanions of  both arsenic an
uately 55?* of the arsenic entering the lake was removed to the sediments.
About 60% of the arsenic in Che sediments was extractab'le with the Iron-
manganese conpounds, indicating that sorption or coprecipitacion was tht '
primary removal process.  Of the riverine input to the lake, 6i" of the
arsenic was dissolved and 35% was associated with particles.-  Since TSUCK  v
the arsenic entering the lake originated from the nearby smelter and was
probably chemically bound to particles which subsequently settled tc. the
sediments, it is quite possible'that in uncontaminated environments, less
arsenic would be removed to the sediaentS and a, greater proportion of the
arsenic in the sediments would be associated with the extractable iron-
manganese compounds.

         La Peintre (1954) deraonstrated that arsenate species are copreci-
pitated or adsorb onto hyd'rous iron oxide.  Shnyukov (1963) observed that
Iron ores are always enriched with arsenic, owing to the high adsorp-
tive capacity of the hydrous Iron oxides and the fact that ferric arsenate'
Is very insoluble1,  Arsena'te species are adsorbed by aluminum hydroxide and
by clays; however, bauxite and silicates are usually only moderatelv en-
riched in arsenic (Onishl and Sandell 1955),  It appears, therefore, that
adsorption of arsenic by sediments is one of the controlling mechanisms f.jr
its fate in the aquatic environment.

    6.3.5  Bioaceuaulati on

         Although arsenic ,1s toxic, a number of studies have shown chat it
is bioaccumuiated.  Arsenic is accumulated by fish both from water and frra
food, but reported concentration f-ictors for arsenic in aquatic organisms
are generally quite low (Table 6-1).

         Reay (1973) studied the arsenic levels In an arsenic-rich river,
the Waikato (New Zealand), and related bioaccuraulation of arsenic by aqua-
tic plants to the total amount transported by the river.  By estimating
total production (ecological) and the amount of( arsenic transported by :he
river, the author estimated that 'only 3-*% of the annual arsenic input to
the river was bioaccumulated, with auch of the balance being discharged tj
the sea and the rernainder settling out with sediment at impoundments.

         In a talc roc osm experiment, Isensee ££ a_l. (1973) investigated the
bioaccuraulation of two organic arsenicals, cacodyllc acid and dimethyl-
ursine, for a total of 32 days in a model ecosystem that contained algae,
snails, daphnia, and fish.  Fish exhibited the least accumulation, *-i:h a
bloconcentration factor of 21 for cacodyllc acid and "U far dlmethvlarsine'.
Snails accumulated t^e compounds to a greater extent (the bioconc*intrat ion
factor ranged froai 110 to 446), and the two ptanktanic components concen-
trated arsenic the most, with bioconcentration factors ranging from T}-? t>

                                 Table 6-1

                    aioconcencration Factors  for Arsenic

Freshwater Planes

Freshwater Invertebrates.

Freshwater Fish

Marine Planes

Marine Invertebrates

Marine Fish
Bioconcentration Factor3







 Chapman e_t  al. 1968
, Reay 1973

 Chapman e_t  a_l. 1966

 Chaptaan «Jt  _aj.. 1968

 Chapman ££ a1. 1968

 Chapoan ££  al. 1968

 Chapman at  al. 1968
a.  Bioconcestracion factors are the ratio derived  from  the  concentration
of the element in the aquatic organism (in pp« wet  weight) divided  by  the
concentration of the element in water (in ppo).

2175.  Ic was concluded that the arsenic compounds did not show a tendency
to biQBignify (increase in concentration as trophic level increases); and
after 32 days, about 30% of the original arsenic in solution was incorpo-
rated by the biota.

         Sorensen (I976a) exposed green sunfish (Lepogig eyanellus) to
various concentrations of arsenic (as sodium arsenate) in water and mea-
sured the accumulation.  There appeared to be a relationship between ex-
posure concentration and arsenic accumulation, but the data were not
statistically correlated.  In further experiments green sunfish were ex-
posed to sodium arsenate under varying temperature and exposure intervals
{Sorensen 1976b).  Arsenic uptake by liver, gut and muscle increased with
arsenic concentration in water and with temperature and exposure interval.
Dead sunfish did not passively accumulate arsenic and no useful method was
found for confirming arsenic-caused fish kills.  Biological half-life for
arsenic In gut and liver was about seven days.  >

         Gilderhus (196S) observed the arsenic uptake by young and adult
blutgills (Lepetal3 aarcpchj.ru*) placed in ponds that had been treated with
various concentrations of sodium arsenate as a herbicide.  After sixteen   !
weeks exposure, whole adult bluegills contained arsenic levels very simi-
lar to the concentration of arsenic remaining in the pond after that
period.  Immature bluegills attained arsenic concentrations nearly twice
those present in adults.  By the end of the experiment, 20 to 80 percent of
the arsenic applied to the ponds remained in solution.

         Sandhu (1977) neasured arsenic content of fish and water In a pond
accidentally sprayed with an arsenical herbicide.  Arsenic levels in the
pood reached 2.5 mg As/1; fish accumulated up to 12.4 ug As/g in muscle,
representing a concentration factor of only five.  Lake Michigan pltnkton
and benthos were found to contain 6.0 and 6.6 yg As/g, respectively (Seydel
1972).  Lake Superior plankton contained about 30 percent less.  The arse-
nic concentrations present in phytoplankton and zooplankton were similar.

         In general, 'fat contains snore arsenic than other tissue fractions.
Fish muscle tissue also accumulates arsenic; however, the biological half-
life of arsenic is only seven days in green sunfish.  Shellfish concen-
trate arsenic to a much greater extent than fish, and marine organisms
contain more arsenic than ,freshwater species.

    6.3.6'  |iotransfortaation

         Arsenic has been shown to undergo a number of biologically medi-
ated transformations in aquatic environments, most of which involve uethyl-

ation to derivatives of arsine (Johnson 1972; Wood 1975, Zlngaro and
Irgolie 1975).  Arseaic forms stable bonds with sulfur and carbon in
organic compounds, and It is the affinity of crivalent arsenic for sulf- •
hydryl groups, most notably in aair.s a^iij, vhic'.i account* fui' t'ue pi iu»«i£/
mode of arsenic toxiclty (Natianal Academy of Sciences 1976).  Pentavalenc.
arsenic is not reactive with sulfnydryl groups, but since some organisms
are capable of reducing arsertate to arsenite, biological reduction in natu-
ral waters could cause an increase in the ratio of arsenlte to arsenate
(Braman and Foreback 1973),

         In mechylation studies, McBride and Wolfe (1971) demonstrated that
Het hatio tjac te r i urn reduced and aethylated arsenate under anaerobic conditions
to dimethy larsine.  A aethyl donor, methylcobolamin in this case, was
accessary.  Cox and Alexander (1973) demonstrated that three species  of
fungi, found in sewage sludge, could produce trimethylarsine.  Two of the
fungi were able to form trimethylarsine from mono- or dimethylarsine, while
the third was able to produce trimethylarsine from arsenite and arsenate as
well,  la general, more trimethylarsine was produced in acidic .nedia  than
under neutral conditions.  The methylarsinea can be produced by a number of
yeasts, bacteria, and fungi.  This literature has been reviewed by Ferguson
and Gavis (1972) and Woolson (1977).  .

         The biological function of the methylation o" arsenic is not
known, but Braraan and Foreback (1973) speculate that -sethylation say  be a
detoxification mechanism since nose of the organic metabolites are con-
siderably less toxic than arsenite. • Alternatively, methylation of arsenic
could be purely adventitious.  In an anaerobic environment, it way be en-
ergetically preferable for organise^ to transtaechylats metals rather  than  .
to synthesize siethane.  Only aerobic metabolisra has been found t.> yield
methylarsines and aethylation may occur in the aerobic upper i.tyer of the
sediment.  The probable mobility of nethylarslnes from the sediuents  to
solution and to the aquatic food chain plus the increased anthropogenic
discharges of arsenic could bring about ever increasiig arsenic concentra-
tions In the aquatic environment.  This cyclic behavior of arsenic in bio-
logical systens has been summarized in Figure 6-3.

6.4  Data Summary

    Arsenic is extremely -avbile in the aquatic environment. ' Although a
Ruaber of studies have described the equilibrium chemistry of arsenic, the
rites T£ -nest of the^e reactions are unknown.  It is evident chat once in
the aquatic system, arsenic cycles through several components, t.*.,  the
water colusn,,the sediments, the biota, and th*> atmosphere.  Figure. 6-i
summarizes this cyclical nature of arsenic.

                                 h«0 _«/'__ CH



                                I *        " 5
                            CM-,— A.— CH,-  K— «.	 CM1

                          THPW«TKTL*"»t>»« ^k  /OrtlETMVL*

                       \    X     N







rti* 1
FiguT« 6-3  The biologic arsenic cycle  in the aquatic environment.

           from Wood (1975).

                      . (WJ.MC. j
                    J fo,,;..o * f
                    »  4.  4r"j"
j !
siifflfvri ••**
-,f...o:*^ ,
i 1! •! -
Hufe.1 MM»iC»lj
•» I«««T'«II
•-* i«e
Figure 6-4  Cycle  of arsenic through different environnentai conpartnsents.
            From Wool son (1977)'.

    Obviously, the fate of arsenic in the aquatic environment is a cooplex
problem, depending on a nuab«r of factors including Eh, pK, aetai sulfide
and sulfide ion concentrations, presence of phosphorus minerals, Iron con-
centration, .temperature, salinity, and distribution and composition of the
biota.  It appears that, in most cases, the sediment is the major sink for
arsenic, but that mobilization by bacteria and other benthie organists.8 re-
turn* much of this arsenic to the cycle.  Much, if not moat, of the arsenic
introduced to the aquatic ecosystem is eventually transported in solution
to the oceans.  Table 6-2 ftiunmarizes ci:« aquatic fate of arsenic.



             Table 6-2

Sumtaary of Aquatic Fata of Arsenic


    Not an important process.

    laportant in determining
    arsenic distribution and
    sob illty.  Interconversions
    of +3 ,and +5 state and organic '
    compiexation are most important,

    laportant when biological
    activity or highly reducing
    conditions produce AsH^ or

    Sorpticn onto clayi, iron
    oxides, and organic material
    are a controlling mechanism -for
    the fare .of arsenic in the
    aquatic environment.
    Appears to b* aost significant
    in lower trophic levels.
    High toxicity lowers overall
    accumulation by aquatic organises.

    Arsenic is metabolized by a
    number of organisms to organic
    arsenicals, thereby increasing
    araenic mobility in the environment.
Confidence or
    All of the noted environmental processes are laportant; however, their
    relative importance with respect to each other is uncertain for
    determining final fat*.              ,                       .

6.5  Literature Cited

Andreae, M.O.  1978.   Distribution, and speeiation of  arsenic  in natural
  waters and s^ae marine algae.  Deep-Sea Research 2j(4) ;39i-M)2.

Boyle, R.W. and l.R.  Jonasson.  1973.  The geochemistry of  arsenic  and its
  use as an indicator element in geochefflical prospecting,   J»  Georhem.
  Expl. 2:251-296.

Braman, R.S, and, C.C. Forebaek.  1973.  Methylated foras  of arsenic in the
  environaent.  Science 182:1247-1249.

Chapman, W.H., H.L. Fisher, and M.W. Pratt   1968.  Concentration  factors
  of chemical elements in edible aquatic organisms,  Lawrence  Radiation
  Laboratory, Li^ermore, Calif.  UC8L-50564.  46p.

Cotton, F,A. and C. Wilkinson.  1972.  Advanced inorganic  chemistry.
  pp.367-402.  Interscience Publishers, New York.  1145p.

Cox, D.P. and M. Alexander.  1973.  Production of triiaethylariine  gas from
  various arsenic compounds by three sewage fungi.  Bull.  Environ.  Contao.
  Toxicol.  9:84-88.

Crecelius, E.A.  1975.  The geocheaical cycle of arsenic  in Lake Washington
  and its relation to other element?..  Liminbl. Oc*anogr,  20(3):441~451,

Cracelius, E.A., M.H. iothner, and R. Carpenter.  1975.  Geochemistries  of
  arsenic, antimony,  mercury, and related elements in sediments of  Puget
  Sound.  Environ. Sci. Technol.  9(4):325-333.

Ferguson, J.F. and J. Gavis.  1972.  A review of the  arsenic  cycle  in
  natural waters.  Water Res. 6:1259-1274.

Gilderhus, p.A.  1966.  Sane effects of sublethal concentrations of sodium
  arsenate on bluegllls and the aquatic environment.   Trans.  Am. Fish Soc.
  95(3}:289-296.       ,

Gupta, $. and K.Y. Chen.  1978,  Arsenic removal by adsorption. J,
  Water Pollut. Control Fed.  50(3):493-506.

Haa, J.D.  1970,  Study and interpretation of ch* rheiaical  characteristics
  of natural water.  pp.206~2Q7.  U.S.G.S.  Water Supply  Paper 1473*
  Washington, D.C.  363p.

Isensee, A.R., P.C. Kearney, E.A. Woolson, O.E. Jones, and  V'.P. Uilliaas.
  1973.  Distribution of alkyl arsenicals in model ecosystem.   Environ.
  Sci. Tecnnol.  7(9);S41-S45.

Jackson, K.S., I,R. Jonasson, and G.B. Sklppen,  1978,  The nature of
  metals-sediment-water interactions in freshwater bodies, with emphasis
  on the role of organic matter.  Earth-Science Reviews 14:97-146,

Johnson, D.L,  1972.  Bacterial reduction of arsenate in seawater.  Nature

La Peintre, M.  1954.  Solubilization par les eaux naturelles de 1'arsenic
  lie au fer dans lea roches sedlaentaries.  Coot. Rend. Acad, 3ci.
  (France). 239:359-360.

McBrlde, B.C. and R.S. Wolfe,  1971,  Biosynthesis of dimethylarslne by
  jtethanobaeterluai.  Biochemistry 10:4312-4317.

National Academy of Sciences.  1976.  Arsenic,  pp.6-121.  U.S. Environ-
  mental Protection Agency, Research Triangle Park, N.C.  STIS P8 262 167.
  480p.  !

Onlshi, H.  and 1.8. Sara'ell.  1955.  Geochemistry of arsenic,  Geochia,
  Costnochim. Ace a 7.-1-33.

Parris, C.E. and F.E, Brinckaan.  1975.  Reactions which relate to the
  environmental mobility of arsenic and aneiaony 1,  Quarternizatlon of
  trimethylarslne and trlaetnylstibine.  J, Org. Cheta.  40:3801-3803.

Parris, C.E. and F.E. Brinckman.  1976.  Reactions which .relate to the
  environmental aobility of arsenic and antimony II,  Oxidation of
  trlaethylarslne and trioethylstibine.  Environ. Sci. Technol.

Reay, P.F.   1973.  The 'accumulation of arsenic from arsenic-rich natural
  waters by aquatic plants.  J. Appl. Ecology 9{2):557-565.

Sandhu, S.S.  1977.  Study on the post-mortem identification' of pollutants
  In the fish killed by water pollution:  detection of arsenic.  Bull.
  Environ,  Contam. Toxicol, 17(3):373-378.

SeyJel, I.S.  1972.  Distribution and circulation of arsenic through water,
  organisms and sediment i of Lake Michigan. . Arch. Hydrobioi.  7l(l):l7,~30.

ShnyukoV, E.F.  1963,  Arsenic in the Cimmerian iron ores of the Azov-Black
  Sea region.  C«oehemistry (USSR) 163:87-93. .

Sorenaen, E.'M.B.  1976a,  Thermal effects on the accumulation of arsenic
 . in green sunflsh, Lepomls cyanellus.  Arch. Environ. Contain, Toxicol.

Sorensen, E.M.B.  1976b.  Toxieicy and accumulation of arsenic  In green
  sunflsh, Lepoais cyanellug exposed to arsenate ia waf»r.   Bull.
  Environ. Contain. Toxicol.  15(6) = 756-761.

yagemann, R.  1978.  Some theoretical aspects' of stability  and  solubility
  of inorganic arsenic in'the freshwater environment.   Water Res.

Haslenchuk, D.G. and H.L. Windon, 1978.  Factors controlling the estuarine
  chemistry of arsenic.  Estuarine Coastal  Mar,  Sci.  7:455-464.

Waislenehuk, D.G.  1979.  The geocheaieal controls on arsenic
  concentrations in southeastern United States rivers.  Chem. Geol.

Weast, R.C. (ed.).  1977.  Handbook of chemistry and physics, 58th
  edition.  CRC Press, Cleveland, Ohio.  2398p.

Wood, J.M.  1975.  Metabolic cycles for toxic elements in the environment  -
  a study of kinetics and mechanism,  pp.105-112.  in Heavy metals in the
  aquatic environment:.  P.A. Krenkel (ed.).  Pergataon Press, Oxford,
  England.  352p.

Woolaon, E.A.  1977.  Fate of arsenicals in different environmental
  substrates.  Environ. Health Perspect.  19:73-81.

Zingaro, S,A. and K.j. Irgolic.  1975.  The nethylation of  arsenic
  compounds.  Science 187:765.

                             7.  AS8ESTOS

7.1  Statement of ProbableFate

    Asbestos is naineralogicall.y staole .and ts not prone to significant
chemical or biological degradation in che'aquatic environment.  After in-
troduction into surface waters, asbestos remains In suspension until phys-
ical degradation or chemical coagulation allows ic to settle into the sedi-
ment •layer.

7.2  I dene i figa tiun - Geoc'neaia try of As bes t _os

    Astestos is .1 generic term for a variety of hydrated silicate minerals
which ha*1* one cjmtnon attribute, namely, the ability to be separated into
relatively so£t, silky fibers.  Although the'natae is ordinarily associated
with those varieties which have technologic 'importance, it has often been
applied to all minerals which fit the above description.  The term
"asbestos" as used in this report, refers to the definition of asbestos
fibers currently used by the Environmental Protection Agency..  The acre •
general term "asbeatifora minerals" denotes minerals of this type without
reference to size, forts, or usage.

    The definition of asbestos currently used by the EPA ia from the notice
of proposed rule-making for "Occupational Exposure to Asbestos" published
in the Federal Register (Oct. 9, 1975; pp. 47652, 47660>• by the U.S. Occupa-
tional Safety and Health Administration (OSHA).  In this notice, the natu-
rally occurring minerals chrysotile, ansosite, crocidolite, treraolite, antho-
phyllite, and actinolite are classified as "asbestos" if the individual
crystallites or crystal fragments have the following dimensions: length -
greater than 5 sicrometers; maximum diameter - less'than 5 micrometers; and
a length to diameter ratio of 3 or greater.  Any products containing _anv^ of
these minerals in this size range are also designated as "asbestos"
(Campbell *t_ _al. 1977).

    The known varieties of asbesciform minerals can be divided into two
sain classes on the basis of their crystal structures: serpentine and am-
phibole.  The sole member of the serpentine class Is chrysot'ile asbestos,
which is by far the aoat common of the- asbestiform minerals.  It accounts
for more than 95 percent of the asbestos fiber produced today (Sp#il 1974.).
There are five recognized asbestiform varieties of amphibole:  crocido-
lite, aoosit*, anthophyllite, treiuollte, and actinoltte.  Although the *m •
phiboles are common rock~form ing minerals, the asbestiform varieties are
much less abundant than chrysotiie.  The physical and chemical properties

of Che asbestiform minerals which are responsible for their commercial
tapuctance can be directly related £o their crystal struc;-tre and chemical
composition.  Therefore, an attempt will be made in this report to
elucidate the structure and composition of these minerals,

    7.2.1  Mineralogy of Commercial Asbestos

         The commercial deposits of asbestos contain one of the following
talnerals;         •
  chrysotile               Mg;jSi2G54 tetrahedra
chat are coordinated to a second layer of linked Mg02(OH>4 octahedra
through the sharing of oxygen atoms; the composite double layer rolls up,
in the sane aanner as a window( shade, to form long hollow tubes.  The
diameters of the individual tubes are approximately 25 nas; the length-
t,o-diaa«ter ratio can vary from 5 to 10 to well over 10,000 (Ross 1973).
These individual cubes are easily separated because the bonding between the
tubes is weak and is due only to van der Waais forces.

         Tht structures of ' the anphibole minerals (Figure 7-2 }, on the
other hand, «rs composed of strips or ribbons of linked poiyhadra, which
join together to form th« three-dimensional crystals.  The individual
strips are composed of three elements - two double chains of linked
(Si»Al)C4 tetrahedra that forts a sandwich with a strip of linked MgO^
or F*06«  These structures are easily broken into fibers due to :hetr
chain defects, which are also called Wadsley defects (Franco «£ aJL. 1977).
These defects are caused by Che formation of expanded "I-bearas" thac are  •
eoaposed of triple, quadruple, etc., chains of linked SiO^ tetrahedra
cachcr than the double chains found in all other araphibole crystal struc-
tures.  If thes* "l-b*aas'  are expanded indefinitely, the resulting strip
b«eow«s identical with th* single talc layer of composition or
Hg6Sigd2o(OH)4; these expanded "I-beans" intermix with the regular
aaphibole structure to produce the fibrous cleavage planes of the amphibole
asbtsci form minerals.

    7.2.2  F 9. rma c ion
         Modes of origin can be inferred frora the stability relationships
araong talc ,, ant nophyl lite ,  enstatite, iarsceri;e, antigorite, and chr>so-
tlie given by Hemle/ ejc a_L, (1977).  Their mineral stability r.elds .5-  I
kbar HiO, in taras of crystallization. temperature and naiaiity ->f aqaecus
silica, are gi'/en in Figure 7-3.  This  fi^ur-j shows J number of relation-
ships pertinent to the problem of formation of asbestos minerals. 'As the
temperature decreases, rorsterite (Mg-rich olivine) can react ta fora
antigonite or chrysotlle depending on the silica concentration in th» aque-
ous solution co ^hich Che divine-bearing rock is exposed.  Ac silica con-
centrations -near the q'jiartz saturation curve, the- forscerite or anthophyl-
lite can undergo alteration to one of the amphibole asbest i form Minerals
.(depending on the a«tal cation concentrations), -jr directly to talc.  Fi-
gure 7-t» shows the stability fields of  forsterice, enstatite, anthophyl-
lite, talc and the amphibole asbestiforn ainerals.  Th« fibrous nature  of
the amphibole asbescifora minerals can  be explained if the alteration
process of a chain silicate (anthophyllite) to a sheet silicate (talc)
proceeds by 'reforming the double chain  at the unit-cell li»vel.  In  this
manner, the fibrous nacure of amphibole asbesti conn minerals appears tc be
related to the crystal growth mechanism; i.e., the crystallites nav -iacle-
ate at many centers and grow as individual fibers oni> a few tets jf nano-
saeters thick (.tranco _ejt ,»l. 1977).  Thus, the presence of "Wart a ley"' defect s
aay be the result of rapii growth and,  in addition, -nay hinaer growth in a
direction perpendicular to  the fiber axis.

    .".2.3  Chemi ca 1 f.ea c t i _y i t_y

         Asbestos has often been touted, as an indestruc table mineral; In
reality, however this is far from the case.  Experiaents on thermal effects
and the acid leaching of asbestos have been suamarized by Speil and
'Leineweber (1969) and hive only a peripheral relation to rhe aquatic fate
of asbestos, but the following laboratory experienertts are definitely
germane with regard to degradation in ambient surface waters.

         Choi and Soich (1972) observed1 the kinetics of the dissolution of
chrysotiie in water over a temperature  range of 5 to 455C.  A parallelism
was noted between the rate of dissolution of -nagneslum from the chrysotiie
and the rate of pH drift.  The rate of  the dissolution reaction was di-
rectly proportional t.> the  specific ^urface area of the asbestos minerals.
It was noted that magnesium cations say btf continuously liberated from  the
chrysotiie fibers, leaving behind an intact silica structure.  This orig-
inal structure, could then r.eadsorb aetal cations, since it will develop a
highly negative charge.  In general, however, this readsorptlon of metal
cations is not observed; the soalier the particle, the faster the zagnesiuia
is liberated from the asbestos structure.  Moreover, the reaction Is tem-
perature sensitive only in the Initial stages of contact between' chrysotiie
and water.

Figure 7-1 Th* structure of chrysotlle.  From Speil and Lelneweber (1969)

                                  8 OXYGEN^	
                                  '4 SILICON
                                  3 OXYGEN*fHYDROXYL
                             O  CATION   LAYER o o o
                                   SILICA  RIBBON
                             o o ooooo o
Figure 7-2  -The structure of the amphibole asbestiform minerals.  From

          Spefl and Leln*weber .(1969).



          (1 kbar  H2O)
                                  2OO     3OO
5OO   7OO
                                 TC -*
  Figure 7-3  Mineral stability fields in the MgO-SiOo-l^O system at  1

            kbar H20.  Modified from Hemley et_ jd. (1977).

              (1 Rbar  H2O)
      '•*;--QTZ SAI	v^-—-•-•-—i-
        !          TALC
Figure 7-4 Mineral stability fiblds In the MgO-SiO,-R,0 system In the
        region of amphibole asbestos formation. Modified from Hemley
        et al, (1977).

         Hostetler and Christ  (1968),  in a.  laboratory  study  of  the  dissolu-
 tion of chrysotile in water, determined an  activity  product  of  chrysociie
 tn water at 25°C of  1Q~^*0,   These  results suggest  that  cnrysotiie is
 slowly soluble  in water  under  conditions of continuous  extraction.   How
.useful these  results are  in the environment can only be determined  through
 further experimentation.   For  instance, Chowdhury  (1975)  studied  the  leach-
 ing of asbestos An distilled water and at body temperature  (37°C).   He
 found that, for all  practical  purposes, amosite and  erocidolite were  Inert
 under, these conditions.   Nonetheless,-  although he  was  unable  to reach a
 chemical equilibrium after two months  of leaching, a significant  amoun;  of
 Che chrysocl.Ze  had dissolved (1,000 Utaol of Mg/g asbestos had been
 leached).  He found  further that under a dynamic system,  after  the  magne-
 sium had leached out, the  silica skeleton began flaking apart,'  thereby
 eliminating the asbestos  structure.

         Langtr ejr_ aJU (1978)  investigated  the effect  of'milling  (done to
 produce industrial useful  fibers) on the chemical  structure  cf  asbestos.
 The milling altered  the  crystallinity  of the chrysotile whica appeared to
 Include shifts  In interlayer bonding between the magnesium  tnd silica
 sheets and changes in hydroxyl configuration.  The change- would  theoreti-
 cally increase  the leachability of chrysotiie in water  2).  These tailings were  dumped  into  the lake
 at the raie of  about 60,000 -  70,000 metric tons per day  in  a water slurry
 of about 2.4 x  10* ai^/day  (Cook 1973) and have been detected In the
 drinking water  of Duluth, Minnesota  (about  75 miles distance).  Although
 Che asbestos fibers are  traveling great distances  in the  water column, they
 are being coagulated and  *edinented  in the  western part of the  lake near
 the tailing delta (Kramer  1976).  If this process  were not going on,
 according to  the calculations  of Kramer (1976), 3.5 x  106 fibers/liter
 should be found distributed evenly throughout the  volume  of  Lake Superior,
 In actuality, however, only 1  x 10®  fibers/liter are present in the
 eastern part of Laice Superior.  Kramer (1976) found, as well, that  the
 greater the distance from  the  tailings themselves, the  richer in magnesium
 the asbestos  became.  This effect was attributed  to the  magnesium-rich
 asbestos having a more negative zeta potential which would prevent
 coagulation and sedimentation.  It appears,  therefore, that  the fate of
 asbestos in the aquatic environment depends  upon little understood
 geochemical mechanisms, which  are a function of the surface  chemistry and
 th* detailed Mineralogy of the fiber.

         The CAS number for asbestos is 1332-21-4; for chrysotiie, 12001-
29-5; for anthophyllite, 17068-78-9, for actinolite, 13768-00-3; for
aoosite, 12172-73-5; for crocidolite, 12001-28-4; for tremolite, 14567-
73-8.  The TSL number for asbestos is A152-4265 and for asbestos fibers

?«3  Suaaary^ of Fa tie Data

    7.3.1  Photolysis

         Asbestos, as a mineral, is not affected by photolytic processes,
It is possible however, that some of the snore than sixty organic coopounds
observed to bs adsorbed Co asbestos fibers (Hilborn et a1. 1974) might be
Susceptible to photolysis.  The importance of such a process is presently

    7.3.2  Chemical SJpectation

         Asbestos, as a Mineral group, is almost indestruetable in the
aquatic environment, as has been discussed in the identification section.
Asbestos is, therefore, resistant Co the processes of chemical speciation
such as oxidation/reduction, hydrolysis, etc.  Differences in ehenical
sp«ciacion, however, are observed in the amount of trace constituents in
the asbestos fibers.  Cralley e_t 'a_l. (19&d) analyzed asbestos for trace
metals and found up to 1482 Ug/g of nickel, 1378 ug/g chrooiura; 54 yg/g
cobalt and 444 ug/g manganese in chrysotiie fibers.  They found as well
that the degree of solubility of the same metal varied considerably by
source and type of asbestos.

         Lockwood (l'J74) analyzed nine Can'adian chrysotiie samples for
trace metals and found 2-14 Ug/g beryllium, 3-10 US/8 cadmium, 202-771
yg/g chrostiutn, 36-78 ug/g cobalt, 9-26 j,g/g copper, 325-1065 ug/g
manganese, 299-11S7 ug/g nickel and 35-71 ug/g thallium in the asbestos
samples.  Sioiliar results were found by Korda e_t _al_, (1977) in th^ir stiidv
of trace elements in the Reserve Mining Company's taconlte dumpings In Lake

         Hilborn e£ al^. (1974), observed more Chan 60 different organic com-
pounds in reference standard asbestos.  The principal organi: raaterials
were alkanes, polycyclic aromatic hydrocarbons and amino acids.  Haringcon
(1962) investigated the benzopyren* content of asbestos be.c-iua*; of its
activity as a carcinogen.  In a later study, Harington (1965) proposed that
the majority of the organic natter originated froat algae aud bacteria ''.hat
were present during the formation of the asbestos that was being mined.  Ha
also suggested taat the asbestos is filtering organic matter out of '.he
aqueous solutions found in most ore bodies.  It is not clear whether chase
trace constituents are important to the overall aquatic late of

 The  trace  metals  and  organic  material,  however,  may  have  a. great  effect  on
 the  biological  systems  with which  the a'sbestos cones in contact.

     7«3.3   Volatilization

         The  Importance cf  the  transport  of  asbestos from the  surface  of
, aquatic  environments  by wind-activated  aerosol fo'raa,tion  is  presently
 indeterminate.  The mobilization of asbestos  from'the surface of highways
 and  soils  into  the atmosphere by the action  of wind  has been observed  in
 urban air  (Holt and Young 1973;  ,
                          1234547*9  10
figure 7-5  Variation of zeta potential with pH for araosite using che
            streaming potential and electraphoresis techniques,  from
            Praaad and Pooley (1973).
Figure 7-6  Variation of, zeta potential with pH for crociddlite using
            streaming potential and electrophotesis techniques,  from
            FrȤad and Poolay (1973).

         Chowdhury and Kitchener (1975), In a study of the zeta potentials
of natural and synthetic chrysotiles, found a wide variety of zeta poten-
tials.  Strongly positive values were found in samples containing an excess
of magnesium in the form of brucite Mg(0fi>2.  Synthetic chrysotile and
natural samples containing little or no brucite, gave moderately positive
zeta potentials over the pH range of 3-11.  Feebly positive or weakly
negative zeta potentials were found in chrysotiles which had undergone
weathering (due to natural leaching of the brucite layer).  Since the pH
and the ambient concentration of Mg+^ ions near the surface ate the main
controlling factors of the chrysotile zeta potential, and since chrysb-
tile's bruC'.te layer is susceptible to leaching in aqueous solution, the
zeta potential of chrysoti-.e is a constantly changing value.  These results
explain the temporary colloidal stability of dilute suspensions of chryso-
tile in environmental media and the mutual coagulation of ehrysotlle and
amphibole asbestos slurries.
         This effect of the colloidal stability of chrysotile was
described by Nauaann and Dresner (1968).  They found that, due co the posi-
tive zeta potential of chrysocil* in environmental media, low viscosity
suspensions could be prepared by means of the inherent charge of the
chrysotile surfaces.  This charge however, is so small in pure chrysotile
that dispenslon was obtained only with short fibers and low fiber concen-
trations (1 percent),  ly increasing the concentration of certain metal-
lic salts, howevar, it was found that low viscosity suspensions could be
prepared under almost any environaental condition.  These observations
suggest that the presence of trace metals will produce a suspension of
chrysotile asbestos in water which will persist until sufficient saagnesiuo
has leached from the chrysocile structure to degrade the suspension.
Furthermore, it is probable that under certain conditions asbestos will
persist in the water column until its concentration becomes high enough co
destroy the suspension or until leaching of the brucite layer decays tne
zeta potential to a point where it will become negative*

         These suspensions of asbestos might be. susceptible to settling
under certain environmental conditions.  Although no specific data are
available on settling rates of" asbestos, several analytic models of the
physical processes in aquatic environments have been developed la recent
years.  Examples Include calculations of vertical eddy diffusivjty within a
nepheloid layer (Feely 1976), dissolution of diatoms (Lol and lennan 1973),
ar.J suspended sediment transport (Nihoul 1971).  Many of the analytic
models have relied on Stokes Law, in Itself an analytic solution to the
problems of a sphere settling in an unbonded Newtonian fluid ( Sverdrup e_t
£l. 1964).  The work of Neihof and Loeb (1972, 1974) and Chase (1979) on
the electrochemical characteristics of natural particulate^ and their
settling behavior in natural water systems suggests that the Stokesian
assumption may not be appropriate for charged particles.  It appears to be

impossible at present cr describe the settling behavior of asbestos  in
aquatic, systems more definitely than Co say that asbestos will stay  tn
suspension for quite a long time.

    7 .• 3,5  Bioaccumulaeion

         !fo evidence was found regarding the bioaccuaulatlon of asbestos in
aquatic organisms.  Asbestos in city drinking water (Levy _et al.  1976; has
been found to be linked with certain gastrointestinal cancers in  humans,'
and! asbestos in food and water (Cunningham jet _a_l. 1977) has been  found t'o
have a siniliar effect in rats.  The release of metal ions from ingestsd
asbestos (Chowdhury 1975} might lead to an accumulation of the heavy mecais
in  some aquatic organisms.  The importance of these processes is  indeter-
tainate at this tiae.

    7.3.6  Biotransformation

         Asbestos is considered to be non-degradable by aquatic organisms.

7.4  Data Summary

    Although there has been little study of asbestos in aquatic *,nviron-
aental systems, it appears that asbestos is refractory In the aquatic en-
vironment.  Therefore, asbestos, once introduced into the aquatic environ-
ment, will remain In the water column until surface charge coagulation or
changes in flow regime allow it to settle out of the system.  Table  7-1
summarizes the aquatic fate of asbestos.






                                 Table 7-1

                    Summary of Aquatic Fate of Asbestos

Does not occur,

So** dissolution of chrysotile
1* observed; tract raetal and
organic materials in the struc-
ture sight have an effect on
the fate of asbestos in bio-
logical systems.'

Mot a jignifleant process.

The sorptive capacity of
asbestos (as defined by IEP
and 2PC) has great effect on
colloid behavior of asbestos*

Not observed in aquatic organisms.

Doas not occur in aquatic
Confidence of



a.  All of the noted envirotmental processes are iaportant; however, their
    relative importance with respect to each other is uncertain for
    determining final fate.

7.5  Literature deed

Alst.e, J, , 0. Watson, and J. Sagg.  1976.  Airborne asbestos ir> the
  vicinity of a freeway.  Atsros, Environ. 10:553-539.

Campbell, W.J., R.L. 31ake, L.L. Brown, E.E. Cat'-ier, and J.J. Sjoberg.
  1977.  Selected silicate Minerals and their asbescifora varieties:
  mineralogical definitions and identification-characterization.  U.S.
  Sureau of Mines Information Circular 8751, Washington, B.C. 56p.

Chase, R.R.P,  1979.  Settling behavior of natural aquatic particulatas,
  Liamcl. Oceanogr. 24(i):41?-426.

Choi, I. and R.tf, Saith.  1972.  kinetic study of dissolution of asbestos
  fibers in waters.  J. Colloid Interface Sci. 40(2):253-262.

Chowdhury, S.  1975.  Kinetics of leaching of asbestos Minerals at body
  temperatures.  J. Appl. Chen. Biotechnol. 25:347-153.

Chowdhury, S. and J.A, Kitchener.  -1975.  The zeta-potentials of natural
  and synthetic chrysotiles.  Ir.ternat. J. Mineral Processing   2:277-283.

Cook, P.M.  1973.  Distribution of taconlt* tailings in Lake Superior water
  an/ pnblic water supplies.  U.Sii Environmental Protection Agency.
  "v.t .onal Water Quality Laboratory.  Second Progress Report, July, 1973.
  luj,uth, Minnesota, 88p.

Cralley, L.J., E.G. Keenan, a.E. Kupel, R.E. ICinsen, and J.R. Lynch.  1968.
  Characterization and solubility of metals associated with asbestos
  fibers.  Aatr, Ind. Hyg. Asaoc. J.  29:569-573.

Cunningham, H.M., C.A. Hoodie, C.A, Lawrence, and R.D. ?onte£ract.  1977.
  Chronic effects of ingested asbestos in rats.  Arch. Environ, Concani.
  Toxicol. 6:507-513.

Feeley, R.A,  1976,  Evidence for aggregate formation lr.,a nepheloid layer
  and Its'possible role in the sedimentation of parttculate natter.  Marine
  Geol, 20:7-13. '

franco, M.A., J.L Hutchinson, D.A. Jefferson, and, J.M. Thomas.  1977, -
  Structural imperfection and Morphology of crocidolit« (blue asbestos).
  Satur« 266:520-521.

Harington, J.St  1962.  Natural occurrence of oils containing
  3,4-benzopyrene and related substances in asbestos.  Datura 193:-»3.

Harington, J.S.  1965,  Chemical studies of asbestos.  Ann. S.Y.  Acad.
  Sci,  132:31-47,                                   '

Hetaley, J'.J. , J.W. Montoya, D.R. Shaw, and R.W. Luce,  1977.  Mineral
  equilibria In the MgO-SiC^-^O system:  II. 'Talc -antigorice
  -forsterite -anthophyilite - snstatite stability relations and  some
  geologic implications in the system.  Am. J. Sci. 277:353-383.

Hllbora, J. , R.S. Thowas, and R.C. Lao.  1974.  The organic concent of
  international reference samples of asbestos.  Sci. Total Environ.

Holt, P.P. and O.K. Young.  1973.  Asbestos fibres In the air of  towns.
  Atmos. Environ. 7:481-483.

Hosteller, P. 8. and C.L. Christ.  1961,  Studies In the system
  MgQ-Si02-C02-H2G(I);  the activity-product constant of chrysotile.
  Ceochim. Ccsmochim. Acta 32:485-497.

fCorda, R.J., I.E. Henzler, P. A. Helmke, M.M. Jimenez, L.A, hiskin, and S.M.
  Larten.  1977.  Trace elements In saaples of fish, sediment , and taconite
  from Lake Superior.  J. Great Lakes Res. 31(1-2) :149~154.

Kramer, J*R.   1976.  Fibrous cutBnilngtonite in Lake Superior,  Can.
 • Mineral. 14:91-98,

Langer, A.M., M.S. Wolff, A.M. Rohl, and I.J. Selikoff.  1978,  Variation
  of properties of chrysotile asbestos subjected to milling.  J. Toxicol.
  Environ. Health 4; 173-188,                      ' '

Levy, 8.S,, E. Sigurdson, J. Mandel, E. Landon, and J, Pearson.  1976.
  Investigating possible tffects ot asbestos in city water:  surveillance
  of gastrointestinal cancer incidences in Duiuth, Minnesota,  Amer. J.
  Epideraol. 103(4) : 362-363.                     -

Lockwood, T.H.  1974.  The analysis of asbestos for trace metals.  Aaer.
  .Ind. Hyg. Assoc. J. 35:245-251.

Lol, D. and A. Lcnuin.  1975.  Size spectra of blogenic particles in ocaan
        and sediments.  J. Geophys. Res» 80:423-430.
Nautnanri, A.W, and W'.H. Dresher.  1968. -Colloidal suspensions of chrysotile
  asbestos surface charge enhancement.  J. Colloid Interface Sci,

Neihof, R.A, and G.I. Loeb.  1972,  The surface charge of {articulate
  natter in seawater.  Limnol. Oceanogr. 17:7-16.

Neihof, R.A, and G.I. Loeb,  1974.  Dissolved organic matter'in sea water
  aad the electric charge of Immersed surfaces.  J. Mar. tes.  32:5-12.

Nlhoul, J.C.  1977.  Turbulent boundary layer bearing silt in suspension.
  Phy*. Fluids 20:5197-5202.

Parks, G. A.  1967,  Aqueous surface chemistry of oxides and complex oxide
  minerals, pp.121-160.  _in Equilibrium concepts in natural water systems.
  W. Stumm (ed.). Advances in Chemistry Series 67.   American Chemical
  Society, Washington, D.C. 344p.

Prasad, N.A. and F.D. Pooley.  1973.  Characteristics of amphibole asbestos
  dust surfaces in aqueous media with reference to quartz.  J. Appl. Chem.
  Biatechnol. 23:675-687.
Ross, M.  1978.  The "asbestos" minerals; definitions, description, mode of
  formation, physical and chemical properties, and health risk to the
  mining community.  Workshop on asbestos: definitions and measurement
  methods.  National Bureau of Standards, Special Publication 506,
  Gaithersburg, Md. p.49-63.

Speil, S.  1974.  Chrysotile in vater.  Environ. Health Perspect.

Speil, S. and J.I. Leineweber.  1969.  Asbestos minerals in modern
  technology.  Environ. Res. 2:166-208.

Sverdrup, H.V., M.W. Johnson, and R.H. Fleming.  1964k  The oceans:  their
  physics, chemistry, and general biology,  pp.956-958.  Prentice-Hall,
  Inc., Englewood, Cliffs, N.J. 1565p.

                              ,8.  BERYLLIUM

8,1  Statement of Probable Fate

    Beryllium has a very low aqueous solubility and is probably precipi-
tated or adsorbed onto solids soon after introduction to the aqueous en-
viroment.  Complex ing agents may solubilize beryllium, but ambient water
quality data suggest that the concentration of this element In heavily
polluted waters is quit* low*  Apparently, beryllium in natural water sys-
tems is found predominately in particulate rather than dissolved fora.

8.2  Identification -Geochemistry ofBeryllium

    Beryllium is a naturally occuring element, found in the earth's crust
at an average concentration of 2.5 ppm (Weast 19?7).  Beryllium is found
chiefly as the minerals beryl (Be3Al2Si60i8), bromellite (BeO),
ehrysoberyl (BeAljO^ and beryllonlte (NaBeP04>.  The chemical be-
havior of beryllium is siaillar to that of magnesium (Cotton and Wilkinson
1972).  In crystalline compoiiods, beryllium exists as bivalent ions of re-
latively small radius (0.35 A).  Two natural nuclldes of beryllium are
known, 'Be and ^Be, the latter being produced by the action of cosmic
rays in the atmosphere.

    Analyses of surface, ground, and rain waters have shown that, in gener-
al, beryllium concentrations are well below lUg/1.  Meehan and Soythe
(1967) reported that the maximum beryllium concentration in 20 rain water
samples and 56 river water samples (from 5 different Australian rivers) was
0.18yg/l.  In a stud" of beryllium In ground water, drinking water, and
surface water, Rtichert (1973) found that even In the heavily polluted
Rhine and Main Rivers (Germany), the concentration was below 0.02 Ug/1.
Men (1970) estimates that the average concentration of beryllium in fresh
surface waters is less than 1 Ug/1.

    Berylliua is concentrated in silicate minerals relative to sulfides.
In common crystalline rocks, the element is enriched in the feldspar miner-
als relative to ferromagnesiua minerals and apparently replaces the sili-
con ion; 85-98% of the total crustal beryllium may be bound in the feldspar
structures (Beus 1966),  Beryllium is thought to become  concentrated in
the later stages of mag ma tic differentiation.  The greatest known concen-
trations of beryllium are found in certain pegmatite bodies, where crystals
of beryl account for a few percent of the total pegmatite volume, and may
be found in several of .the strata of zoned-dyk«s*  The element is sometimes
concentrated in hydrothermal veins, and some granitic rocks contain suffi-
cient amounts to permit the crystallization of small amounts of b^ryl*
During the weathering of crystalline rocks and during  sedimentation
processes, beryllium appears to follow the course of aluminum, and it be-
comes enriched in some bauxite deposits, clays,  and deep-sea sadlnents.

    Beryllium has a complicated coordination chemistry and  can form com- .
plexes, Gxycarboxy'ates,  and chelates with a variety of materials (Bertin
and Thomas 1971),  In aqueous solution,  beryllium does not  exist as actual
Be'*''- ions, but as hydrated complexes (Cotton and Wilkinson  1972).  Com-
plexiag of beryllium may result 'in soluble beryllium concentrations in ex-
cess of those predicted on the basis of  conventiot»al thermodynamic con-

    Beryllium, atomic number 4, is among the lightest of elements with an
atonic weight of 9.C12,  It is1 a member  of the alkaline eartb elements,
and, except for its taetallie state, always has a valence, of +2,   The metal
has a specific gravity of 1.85,- a melting point  of 127fa3U,  and a boiling
point of ,.2970'C (Weast 1977).

    The CAS number of beryllium is 74*0-41-7, and the TSL number is

8«3  Sumaary ofFateData

    8.3.1  Photolysis

         No data were found regarding photolysis of beryllium compounds.
Photolysis is probably not an important  process  in the aquatic fate of

    8.3.2  ChemicalSpeciation

         Beryllium isr the smallest of the group  II metals - the crystal
radius of the divalent ^on is only 0.31^ .  The  snail ionic radius and the
resultant large surface charge-density ar* dominant influences on the chem-
istry of beryllium.  Thus, beryllium forms stable compounds with small
aniong, such as fluoride and oxide, because unusually close approaches to
their ionic centers are possible.  The highly nydrated state of beryllium
ion in acidic solution and the aaphoteric nature of beryllium are all
further consequences of the small si?e and high surface charge-density of
the beryllium ion (Drury £t £l. 1978).    }

         Soluble beryllium salts are hydrolyzed  to form insoluble beryllium
hydroxide, B*(OH)2 (Cotton and Wilkinson 1972).   The solubility of
Be(OH>2 is quite low in the pH range of  most natural waters,  A pH-log
(species) diagram constructed for a simple system in which  the solubility
of Be(OH>2 is controlled by dissolution  into Be+2 and HB«02~ indi-
cates that beryllium will have minimal solubility at about  pH 7.5 (Figure


Figure 8-1  pH-log (species) diagram for Be(OH)„,Be
            From Cotton and Wilkinson  "4^
and HBeO.

         Formation of complexes with hydroxide ions may increase the
solubility somewhat, especially at higher pH where polynuclear hydroxide
complexes may form (Hem 1970).  It is probable, however, that In saost
natural environments, beryllium is present in particulate rather than
dissolved form (Hem 1970).  This is substantiated by empirical data which
indicate that, even in polluted rivers, dissolved beryllium levels are very
         Although inhalating the airborne dust is the most widely known
hazard associated with beryllium, no evidence was found that biomethylatiort
or other processes result in volatilization of berylliuia from aquatic

    8.3.4  Sorgjtion

         No data ware found on the adsorption of beryllium.  Nonetheless,
due to its geochetnical similarity to aluminum, one would expect that at low
pH, beryllium would tend to be adsorbed onto clay mineral surfaces while at
high pH, it should be cooplexed in some insoluble compounds.  Beryllium
should be present in particulate (either sorbed or precipitated) rather
than in dissolved fora in most natural environments.

    8.3.5  8 ioacc uaml a t i on

         Little work has been completed concerning the biological pathways
of beryllium in water,  Tarzwell and Henderson (1960) have demonstrated
that beryllium is highly toxic to wanawater fishes in soft water.  Slonim
(1973) measured 'fle uptake by guppies (Po eg ilia reticulata) in a static
freshwater system.  Levels were highest in the viscera and intestinal
tract, followed by kidney and ovary.  Uptake was directly related to
beryllium concentration in water, inversely related to fish size, and not
related to fish age.  Water hardness was inversely related to beryllium
toxic ity but did not appear to influence beryllium uptake by guppies.
Thus, the body burden of beryllium is not the controlling factor governing
toxicity.  It was suggested that beryllium, present in a particular target
organ, may be related to tostic response.

         Cowgill (1973), in a study of biogeochemical cycles in a fresh-
water glacial lake in Connecticut, found no evidence of bioaccunulatlon or
food chain magnification.  He found r.ha : beryllium seemed to be concen-
trated in the stalks of aquatic plants, with lower quantities in the
flowers and leaves.

         Although beryllium has a low solubility in water, ic is possible
that benthos could accumulate beryllium from sediment and thereby transfer
the metal to higher organisms via the food chaia.  Although beryllium- tox-
icity decreases with increasing water hardness, beryllium uptake appears to
be unaffected by increasing hardness.  There is a general paucity of infor-
mation (Table 8-1) concerning the accumulation of beryllium by aquatic
plants and animals.

 —    ,                    .      Table 8-1        '          '     '

                  Bioconcentration Factors for Beryllium

       Taxon           Sioeoncantration Factor*        Reference

Freshwater Plants                 100                 Chapman _e_t a 1. 1968
Freshwater Invertebrates          100                 Chapman _££ al. 1968
Freshwater Fish                   100                 Chapman e£ al. 1968
Marine Plants                     100                 Chapman e_t al. 1968
Marine Invertebrates              100                 Chapman jet al. 1968
Marine Fish          .             100                 Chapman jet al. 1968
,a.  iloconcentratlon factors are the ratio derived from the concentration
    of che element in the aquatic organism (in ppro wet weight) divided by
    the concentration of the element in water (in ppm).
    8.3.6  Biotransforaation

         No data were found relative to aquatic fate on biotransformation
of beryllium or its compounds.

8.4  Data Summary                       .•                 '

    Beryllium has a very low aqueous solubility under normal pH conditions
due to the formation of insoluble Be(OH>2.  Adsorption probably further
reduces the concentration of dissolved beryllium. > Formation of soluble
complexes may tend to increase the solubility of'this element, but it ap-
pears that under most circumstances, beryllium is associated with the
particulate rather than the dissolved components of natural water systems.
The aquatic fate of beryllium is summarized in Table 8-2.

                                  fable  8-2
                    Summary of  Aquatic  Fata  of  Beryllium
Confidence of
, Photolysis

 Chemical  Speciation*



Not an important fate.

Beryllium la hydrolyzed
co fora insoluble compounds.
i controlling mechanism for
beryllium in the aquatic

Not an important fate.

Probably adsorbed by clays
and other mineral surfaces
at low pH.

Slight accumulation by
aquatic organisms.  Mo
food chain magnification
in evidence.





     All  of  the  noted  environmental  processes  are  important;  however,  their
     relative  importance  with  respect  to  each  other  ia  uncertain  for
     determining final fate.

8,5  Literature Citad

Bertin, F, and G. Thomas.  1971,  Sur la chimie de coordination du
  beryllium.  Bull. Soc. Chin. France.  10:3467-3498. •

Beua, A.A.  1966.  Distribution of beryllium in granites.  Geochemistry
  (USSR) 5:432-437.

Chapman, W.H. , H..L. Fisher, and M.W. Pratt.  1968.  Concentration factors
  of chemical elements in edible aquatic organisms.  Lawrence Radiation
  Laboratory, Livenaore, Calif.  UCRL-SG564.  46p,

Cotton, F.A. and G. Wilkinson.  1972.  Advanced inorganic chemistry.
  pp.245-254.  Intersciertce Publishers, New York.  1145p.

Cowgill, (J.M.  1973.  Biogeochemical cycles for the chemical elements in
  Nymphata gdorata Ait. and the aphid Rhopalosiphum nymphaeae (L.) living
  in Linsley Pond.  Sci. Total Environ"2:259-303.

Drury, J.S., C.R. Shriver, E.B. Levd.3, L.u. Towill, and A.S. Hanraons.
  1978,  Reviews of the environmental effects of  pollutants:  VI
  Beryllium.   pp.8-38.  Oak Ridge National Laboratory, Oak Ridge, Tenn.
  PB 290-966.  198p.

Hem, J.D.  1970.  Study and interpretation of the chemical characteristics
  of natural water,  p.194.  U.S.G.S. Water Supply Paper 1473.  Washington,
  D.C.  363p.

Meehan, W.R, and L4E. Smythe,  i967,  Occurence of beryllium as a trace-
  element in environmental materials.  Environ. Sci. Technol.
  1(10): 83 9-84 4.

R-ichert, J.K.  1973.  Beryllium, ein toxisches element In der
  mensch-lichen umgebung unter besonder berilcksichtigung seines vorkommens
  in gewassern,  Voa Wasser 41:209-215.

Slonlm, A.R.  1973.  Acute toxicity of beryllium  sulfate to the common
  guppy.  J. Water Pollut. Control fed. 45(10):2110-2121,

Tarzwell, C.M. and C. Henderson.  1960.  Toxicity of less cotnmon raetala to
  fishes,  Ind. Wastes 5:12.

Weast, R.C. (ed.)  1977,  Handbook of chemistry and physics, 58th
  edition.  CRC Press, Cleveland, Ohio.  2398p.

                                9.  CAPMFJM

9.1  Statement of Probable Fate

     Compared to the other heavy metals, cadmium la relatively mobile in
the aquatic environment and may be transported in solution as either hy-
dra ted cations or as organic or inorganic complexes.  In moat natural sur-
face waters, che affinity of complsxlng Uganda for cadmium probably
follows the order; humi? acids > CQ$2~ ^ OH" ;> Cl~ >_ 5Q^2~.
In polluted waters, conplexing with organic materials is the moat important
factor in determining the aquatic fate and transport of cadmium.   Sorption
processes account for the removal of dissolved cadmium to bed sediments and
ars increasirgly effective as pH increases.  Adsorption onto organic
materials, mineral surfaces, co-precipitation with hydrous natal  oxides,
and isomorphous substitution in carbonate minerals can all result in re-
ductions in aqueous cadmium concentration.  Cadmium is strongly accumulated
by organisms at all trophic levels.

9. 2  Iden£j,f icatioi^ -^ Geochemistry; of Cadmium

     Cadmium is a relatively rare element that is concentrated in zinc-
bearing sulflde ores (Zn/Cd ratio Is usually 100 to 200) and, consequently,
is found in all zinc-containing products.  It is found at an average con-
centration of 0.15 ppm in tfce earth's crust (Weast 1977). Most fresh wafers
contain less than 1 ppb cadmium.  The chemistry of cadaiua in surface and
ground waters has been reviewed by Hen (1972), giving calculations of
equilibrium solubilities with the hydroxides or carbonates as solid phases.
Host natural waters are undersaturated with respect to these phases, i.e.,
only 20 percent have cadmium concentrations in reasonable agreement with
calculated solubilities that assume CdCOj Co be the equilibrium solid
phase.  Cadmium levels in sea waters average about 0.15 ppb.

     Cadmium has an atomic weight of 112.41.  Metallic cadmium has a melt-
ing point of 320.0*C, a boiling point of 765*C, specific gravity of 8.642,
and a vapor pressure of X mm at 394*C (Weast 1977),  The geochemistry of
cadmium has been extensively reviewed by Waklta and Schmitt (1970).

     Cadmium's CAS number is 7440-43-9;  its TSL number is A344-224S.

9.3  Summary of fata Data

     9.3.1  Photolysis

          Ho evidence was found indicating that photolysis Is an important
mechanism in determining the fate of cadaiua compounds in the aquatic en-

     9.3.2  Cheaical
          In natural waters cadmium can be found in several chemical forms,
for example, as simple aqua ted ions, as metal-inorganic complexes, or as
metal-organic complexes.  An understanding of the chemical speciation of
cadmiua in any given situation is based upon theoretical calculations of
hydrolysis, oxidation/reduction and organic complexation.  A short pre-
sentation of this material, will be given after which the chemical speci-
ation of cadmium in various aquatic environments will be discussed.

          Cadmium forms complexes with OH' such as CdOH+,
Cd(OH)2<*q.), Cd(OH)3*, and Cd(OH)42~.  A distribution diagram
for c«d«lua hydroxide complexes is shown as Figure 9-1 (Weber and Posselt
1974),  As is evident froa the diagram, almost all of the soluble cadmium
is in the divalent cation form up to about pH 9.  The solubility of cadmium
decreases as pH increases due to formation of solid Cd(OH>2» according to
Che reaction:

          Cd2+ + 2 OH"    ^   >      Cd(OH)2

Patterson et_ _ai» (1977) studied the removal of dissolved cadmium by hydrox-
ide and carbonate precipitation.  A comparison of experimentally determined
Cd(OH>2 solubility with the calculated solubility^ curve is shown as Fig-
ure 9-2.  The diagram shows that, even at the optimal pH for precipitation,
the equilibrium solubility of cadmium is still approximately 1 mg/1.

          Since it seemed possible that carbonate could be a more effective
control on cadmiua solubility than precipitation as the hydroxide,
Patterson jet al_. (1977) investigated the use of a carbonate system as a
possible treatment technology.  The results of adding carbonate ss
N*2C03 with Oj « lO"**2 M and Or • KT2'7 M are shown as
Figures 9-3 and 9-4, respectively.  The treatment time was 4 hours.  Al-
though the dissolved cadmium concentrations are still quite high in the
carbonate systems at pH > 10, they are considerably lower than in the
hydroxide system.

          Cadmium is always found in the 4-2 valence state in water (Cotton
and Wilkinson 1972).  Therefore, redox potential has little direct effect
oa cadmium.  Under reducing conditions and In the presence of sulfur, how-
ever, cadmiua may react to form the insoluble sulfide.  The log Kgp (log
solubil!ty product constant) for cadmium sulfide is 5.73 (Huang ft al.
1977) for the following reaction:

          CdS(s) + 2H+     «  	*   Cd2+ + H2S(aq)

          0«  -
                                                          13    14
Figure 9-1  Distribution diagram  for  cadmium hydroxide complexes.  From
            Weber and Posselt  (1974),
                   S  8-
m  -3-1

9  iji -nm

9  4-
          r    i


                                              10    12   1*
Figure 9-2  Comparison of cadmium hydroxide  solubility -data with
            theoretical solubility curve.  From Patterson et ai. (197?).

                        2    4    •    •   18    12   14

Figure 9-3  Comparison of cadnium  solubility data with theoretical phase

            dlagraa, ('Cf «• 10"^*^ mol/1).   From Patterson _*£ a_i. (1977).
                                         10   12   14
Figure 9-4  Comparison of cadmium carbonate  solubility  data with

            theoretical phase diagram  (Cj  *  10~2*7 aol/1).   From

            Patterson et al. (1977).

Under acidic conditions, CdS is mare soluble.  In the sediments,  in anaero-
bic digestion of waste water, and in other reducing environments  where sul-
fur is available, the solubility cf cadmium may be controlled by formation
of CdS (Holmes rt al. 1974).

          Long and Angino (1977) developed a theoretical model to study the
chemical speciation af cadmium in aqueous solutions and the response of
cadmium to variations In ionic strength and complexation.   Association of
cadmium with the ligands OH~, Cl", C0^~, S0^~ and HCC»3~
were considered at pH values from 3.5 to 11.0 at 25°C in differing sea-
water-freshwater mixtures.  The results are summarized in Figure 9-5.   In'
general, the relative importance of the various ligand-cadnium complexes
can be predicted from a comparison of their stability constants.   Unfortu-
nately, this model does not take into account natal-organic complexes, and
it is, therefore, useful only in unpolluted, relatively organic-free

          Gardiner (I9?4a), in his study of the speciation of cadmium in
natural waters, found that a,substantial portion of the total cadmium in
river and lake water will usually be present as the divalent cadmium ion,
the concentration of which will b« inversely related to the pH and the
concentration of organic material present in the water.  Humic substances
usually account far most of the complexation, followed in importance by the
carbonates.  O'Shea and Mancy (1978), in their study of the effect of pH
and hardness on cadmium speciation, found that the effect of pH and hard-
ness was insignificant in trace metal-inorganic interactions.  Hardness and
pH were quite important, however,  in trace metal-huoic acid interactions.
Increasing the pH increased the exchangeable cadmiua while an increase in
hardness led to a most pronounced decrease in the extent of the huoic acid
interaction.  Metals responsible for hardness apparently inhibit the
exchangeable interactions between metals and huaic materials in ways that
are not yet full; understood.

          Guy and Chakrabarti (1976), in their study of metal-organic in-
teractions in natural waters, found that huiaic acids in solution and other
natural complexing agents can maintain cadmium ions in a bound form at a pH
as low as 3.  The release of cadmium from sediments is, therefore, ap-
parently controlled by a combination of ion exchange and complex formation
whereby the stability of the metal-organic complex determines the amount of
metal solubilized.

          In summation, the transport of cadmium in the aquatic environment
is controlled by the speeiation of the ion.  Although it appears that in
most unpolluted waters cadmium will exist mainly as a divalent cation, or-
ganic material will have a pervasive effect on the chemical fora in which

                 100.0 i
                   I.I *5 S J *.$ 7.J I.J « 5 W,5 11.5
Figure 9-5   Cadmium speciation la  seawater-freshwater mixtures
             (Long and Anglno 1977).

 cadmium will  be  present  in  industrialized urban areas.  In most  fresh water
 systems,  the  affinity of complexing  ligands  for cadmium appears  to follow
 Che  order  of;  hunic acids  > COj2" > OH" >_ Cl~ >_ S042".
 Nonetheless,  there  is as yet Insufficient information  to present a compre-
 hensive model  for  the transport and  sorption of cadmium based upon these
 and  similar  Interactions.

      9.3.3  Volatilization

           Cadmium  Is not known to form volatile compounds In ehe aquatic
 environment.   Unlike mercury, which  is directly below  cadmium in the peri-
 odic table and,  therefore,  chemically similar, biological methylation of
 cadmium has  not  been reported.  The  Instability of alkyl cadmium compounds
 In the presence  of  water or oxygen (Cotton and Wilkinson 1972) probably
 precludes  their  formation  iri aqueous ecosystems.

      9.3.4  Sorptioa

           Although  sorptlon processes affect cadmium to a lesser extent
 than most  of  the other heavy aetal pollutants, sorptlon by mineral sur-
 faces, hydrous metal oxides, and organic materials probably removes more
 cadmium fiom  solution than  does precipitation.  Various studies  place
 different  emphasis  on the  relative contributions of sorptlon to  clays,
 sorptlon  to organic matter, co-precipitation with hydrous Iron,  aluminum,
 and  manganese  oxides, and  isonorphous substitution in  carbonate minerals.
 All  studies  indicate that cadmium concentrations in bed sediments are
 generally at  least  an order of magnitude higher than in overlying- waters
 (Gardiner  I974b;   Kubota _et _al. 1974;  Perhac,1974a, 1974b;  Helz etal.
 1975;  Steele and  Wagner 1975;  Farrah and Pickering 1977;  Schell and
 Nevissi 1977;  Suzuki e£ al. 1979).

           Suzuki «t al.  (1979), in their study of a polluted Japanese
 river, found  that  the cadmium content of the sediment  is directly propor-
 tional to  the ignition loss of the sediment  Indicating that the  organic
 material  is mainly responsible for the accumulation of cadmium in organi-
 cally polluted river sediments.  Moreover, laboratory  studies on the river
.sediments  demonstrated that cadmium  sorption could be  correlated with the
 amount of  organic  matter present in  the sediment by a,  Freundlich-ty'ps
 equilibrium  relation:

                       O.M. « k C

 where OtM. is organic matter as measured by  ignition loss, C is -concentra-
 tion of cadmium  In the aqueous phase, and k  and n are  experimentally deter

mined equilibrium constants.  These results suggest that In the transport
of cadmium, suspended solids of high organic content play a dominant  role
In polluted waters.

          Gardiner (1974b), in a laboratory study of the adsorption of cad-
mium on mud solids, particles of clay and silica, and humlc material, found
that the adsorption of cadmium on mud solids is of major importance In con-
trolling the concentrarion of cadmium in fresh water.  He found that  con-
centration factors for mud varied between 5,000 and 50,000 depending  on the
type of solid, its state of subdivision, 'the concentration of metal ion and
complexing llgands present, as well as the temperature, pH, and hardness of
the water.  It appeared further that humic material was at all times  the
major component of sediment responsible for adsorption.

          In contrast, Perhac (1974b) found that most of the cadmium in the
bottom sediments of an ^unpolluted Tennessee stream was associated with car-
bonates and (to a lesser extent) iron oxides, and therefore hypothesized
that cadmium occurs in cation lattice sites within the carbonate minerals.
Isomorphous substitution of cadmium for calcium in such minerals could pro-
bably occur;  the crystal ionic radius for cadmiura(II) is 0.97 X, and the
radius for c«»lciua{Il) is 0,99 X (Weast 1977).  It is pertinent that
Perhac's (1972;  1974a) studies of metal distribution in unpolluted streams
show that most of the cadmium transported in the water column is in the
dissolved state (77.4-95.4 percent).  Minor amounts are transported with
the coarse particulates (3.5-21.9 percent), and only a small fraction is
transported with colloids (0.5-6.0 percent).  Even though the concentration
in colloids is greatest, the concentration in coarse particulates inter-
mediate, and the concentration as dissolved species is the lowest, the mass
of water transported in the stream so greatly outweighs the mass of sus-
pended particles that the greatest transport (by mass) is in the dissolved
state.  Perhac's (1974b) assertion that the carbonate minerals and iron
oxides can control cadmium mobility is corroborated by Steele and Wagner
(1975), who found that most of the cadmium in Buffalo River (Arkansas)
sediments was in the font of mineral clasts, although a minor amount  was
associated with metal oxides.

          Ramaaoorthy and Rust (1978), in their study of Ottawa River sedi-
ments, found that, although the sediment was composed mainly of well  sorted
sand, It was an efficient sink for heavy metals.  .They discovered that this
was due to the significant amount of organic material added to the sedi-
ments by the commercial use of the river for logging.  The mobility and
persistence of the cadmium itself was in part dependent on the extent of
Its sorption onto the sediments while its redistribution was a function of
desorption processes at the sediment-water interface.  Both sorption  and

desorption were controlled by the nature of total heavy metal loading, th«:
sediment type and the surface water characteristics.

        •  The adsorption of cadmium onto soils and silicon and aluminum
oxides was studied by Huang _et a_l. (1977).  The results of this laboratory
study indicate that adsorption is strongly pH-depeadent, increasing as con-
ditions become more alkaline (see Figure 9-6).  When  the pH is below 6-7,
cadmium is desorbed from these materials.  Cadmium has considerably less
affinity for the absorbents tested than do copper, zinc, and lead, and thus
sight be expected to be more mobile in the environment than these materi-

          Another relevant observation of Huang jet al. (1977) was that
addition of anions to the dissolved cadmium caused an Increase in adsorp-
tion (Figure 9-7).  Hvmic acid was most effective in  this regard, followed
by nitrilotriacetate, tartrate, glycine, and phosphate, respectively.
Huang et al. (1977) suggest that the anions complex the cadmium and that
subsequent adsorption of the organo-cadmium compounds by hydrous solids *
occurs through specific chemical bonds (e.g., "sharing of free electrons
available from the metal-ligand complexes and hydrogen bond").

          Tht mode by which cadmium is sorbed to the  sediments is important
in determining its disposition toward remobilization.  Cadmium found in
association with carbonate minerals, precipitated as  stable solid com-
pounds, or co-precipitated with hydrous iron oxides would be less likely to
be mobilized by rasuspension of sediments or biological activity.  Cadmium
adsorbed to-mineral surfaces (e.g., clay) or organic  materials would be
•ore easily bioaccumulated or released in the dissolved state when sedi-
ments are disturbed, such as during flooding.

          Although It is widely reported chat dissolved cadmiua concentra-
tions decrease with distance from the source and sediment concentrations
art concomitantly enriched, several authors have reported a remobilization
phenomenon in which sorbed cadmium is subsequently released due to salinity
(H«li ji£ al. 1975) and redox effects (Holmes _et al. 1974; Lu and Chen
1977).  Kubota' tt £l. (1974) showed that the cadmiua  concentration of Lake
Cayuga was higher than the concentration in tributary streams, suggesting
remobilizatlon from sediments.  Such remobllization may have occured due to
a reduction in pH in the lake water relative to most  of its tributaries;
unfortunately, pH values for the lake were not reported.

          It appears, therefore, that sorption processes are important in
determining cadmium transport, partitioning, and potential for remobiliza-
tion.  In unpolluted waters, exchange of cadmium for  calcium in the lattice
structure of carbonate minerals can remove cadmium from solution.  Thus,

                        3  4   S  I   ?  •   9

Figure 9-6  Adsorption of cadmium on various  solids  \.froo Huang et ai,

            1977).  Concentration of solids » 5  g/1;   [Cd2+\  «    1

            110 oig/1;  , 3nic strength  -  10~7  M (NaCl).


figure 9-7  Effect of anlona on  the adsorption  of  cadmium by Metapeak soil,
            In addition  co the experlnental  conditions shown in the caption
            of figure 9-6, the anion  concentration was 1Q~*  M except for
            huetic acid which wee added at SO mg/1  (froa Huang e_c a_l. 1977).

eo-precipication with hydrous iron, aluminum, and manganese oxides can be
the controlling factors in these unpolluted waters.  In polluted or or-
ganic-rich waters, however, the adjorption of cadmium by humlc substances
and other organic complexing agents will be the controlling factor in
determining transport, partitioning, and potential for remobillzation.

     9.3,5  B i oac cumu 1at1op

          Cadmium is strongly accumulated by all organisms.  Because of its
chemical kinship to tine, cadmium tnay displace zinc in certain enzymes,
thereby disrupting normal metabolic function.

          Cadmium has been reported to accumulate in the tissue of aquatic
and marine organisms at concentrations hundreds to thousands of times high-
er than in the water column (Sisler et al. 1962;  Friberg et al. 1971;
Huckabee and Blaylock 1973;  Kelao and Frank 1974;  Valiela _et al. 1974;
Kinkade and Erdman 1975).  Fish accumulate cadmium most readily in the
liver, kidneys, and intestines, followed by the gills and the remainder of
the body (Cearley and Coleman 1972;  Huckabee and Slaylock 1973).

          Several surveys of the concentration of cadmium in various marine
and fresh water biota have been completed.  Lovett ejt _aJu (1972) examined
concentration levels in fish from various New York State fresh water lakes
and streams.  Maximum concentration levels exceeded 0.1 ppo Cd but moat
fish contained lass than 0.02 ppm Cd.  Martin and Broenkow (1975) reported
that mixed phytoplankton and zooplankton collected off Baja California near
San Diego averaged 13,2 ppm Cd (dry weight basis);  samples collected from
other coastal areas never exceeded 7.5 ppm Cd.

          Reported microcosm and field studies differ in the relative con-
centration factors for cadmium in biota.  Lu je_t ad, (1975) found that bio-
accunulation of cadmium was strongly correlated with the cation exchange
capacity of test soils in their microcosm.  As cation exchange capacity
Increased, levels of cadmium in the biota decreased.  Fish (Gambusia
fjffinia) accumulated less cadmium than algae, snails, mosquito larvae, or
sorghum.  The invertebrates accumulated more cadmium than algae or sorghum
In two out of three teat soils; but in the soil with the least cation ex-
change capacity, algae accumulated cadmium about 4 times as much as either
of the Invertebrates.  Bioconcentrarlon factors (concentration in organism
* concentration in water) calculated from the data of Lu e_£ al. (1975)
range from about 10* to 10*.

          the influence of hardness on uptake of cadmium by a microcosm
containing an alga, a rooted plant, snails, catfish, and guppies was
studied by Kinkade and Srdman (1975).  They found that initial uptake of

          cadmium was faster In hard than in soft water but that the total concen-
          tration of cadmium was greater in the organisms that were placed In soft
          water.  The relative bioaccumuation factors descended in the following
          order:  rooted piant > alga > guppies > snails > catfish.

                    Pascoe and Mattey (1977) exposed three-splned sticklebacks to a
          range of cadmium levels in water (0.001-100 mg Cd/1) for up to 79 days.
          Sticklebacks accumulated cadmium at all concentrations tested;  however,
          the concentration factor was inversely and linearly related to exposure
          concentration.  Concentration factors ranged from 311 at the lowest ex-
          posure to 0.51 at the highest.! All of the concentrations tested were
          lethal to sticklebacks.

                    Cadmium is readily accumulated through both food and water by
          fresh water organisms, and either source of uptake can result in the de-
          velopment of toxic symptoms by fishes.  Fish tissues appear to reach equil-
          ibirium with respect to cadmium after 8-20 week?' exposure, depending upon
          the water temperature (Phillips and Russo 1978).  Cadmium uptake increases
          with increasing water temperature and decreasing salinity.  There is an In-
          dication that sex may determine the rate of cadmium accumulation in some
          fi»h species due, perhaps, to some sex-related metabolic differences.  Pish
          accumulate highest cadmium concentrations in the kidneys and liver and
          little in the edible portions.  Bioconcentratidn factors for cadmium are
          summarized in Table 9-1.

/              9.3.6  Biotransf orma t ion

I                    Mo evidence was found in the reviewed literature for blomethyla-
1          tion of cadmium.  Biologically produced Uganda aay affect the nobility of
          cadmium in aquatic environments, especially under eutrophic conditions.
          Cadmium also can b« completed in vivo by polydentate ligands that are nor-
          mally Involved in the binding "sites of essential metal ions such as iron,
          manganese, cobalt, zinc and copper (Fulkcrson and Goeller 1973)*

               '"4  DataSummary

               Cadmium is mobile in the aquatic environment relative to most other
          heavy metals.  It occurs as the divalent metal cation in acidic and clr-
          cumneutral water, and it forms complexes with organ{•• material in highly
  •        polluted waters and complexes with carbonate and hydioxide ions at higher
          pH values.  The formation of complexes with huraic substances Is important
          because these complexes are more easily assimilated by the sediments than ,
          the hydrated divalent cation. Sorption processes are the most important
          factor in reducing the aquatic load and transport velocity of cadmium.
          Cadmium is less mobile in alkaline than in acidic waters.  Sorption to

                                 Table 9-1
                   Sioconeantration Factcrs  foe Cadmium
     Taxon                Gonc
organic materials and clay minerals, co-precipitation with hydrous dietal
oxides and substitution in carbonate minerals all affect the distribution
and fate of cadmium.   Cadmium, aJ'hough .highly toxic, is concentrated by
all organisms.   The aquatic fate of cadmium Is summarized in Table 9-2,

                                 Table 9-1

                    S'unmary of Aquatic Fate of Cadmium


Chemical Speciation3
Confidence of

Biotraniformat ion
Not an important process*                High

In moat unpolluted waters the            Medium
majority of the cadmium will exist
aa the hydrated divalent cation.
In polluted waters, eomplexation
with organic material will be most
important.  Affinity of ligands for
cadmium follows the order of humic
acids >C032~X)H~>-C1^S042-.

Not an important process.                High

Various sorption processes re-           High
duce the mobility of cadmium
and result in the enrichment of
suspended and bed sediments re-
lative to the water column.  In
unpolluted waters, sorption onto
clay minerals, and hydrous iron
and manganese oxides are controlling
factors.  In polluted waters, sorption
onto organic materials is the con-
trolling factor.

Biota strongly accumulate Cd             High
with concentration factors rang-
ing frost 1Q^ to 10* or more.
Bioaccumulation is greater in
soft than hard water.

No bioraethylation in evidence.           Medium
Organic ligands of biological
origin may affect solubility and
    All of the noted environmental processes are important; however, their
    relative importance with respect to each other is uncertain for
    determining final fate.

9.5  Literature Cited •  '                                     .

Cearley, J.E. and R, L. Coleman.  1972.  Cadmium toxic icy and bioaceumu-
  lation In largenouth bass and bluegill.  Bull. Environ. Contami Toxicol.

Chapman, W.H., H. L. Fisher, ami M.W. Pratt.  1968.  Concentration factors
  of chemical elements in edible aquatic organisms.  Lawrence Radiation
  Laboratory.  Livurmore, Celif. UCRL-50564.

Cotton, F.A. and G» Wilicersoa.  1972.  Advanced inorganic chemistry.
  pp.503-527.  Inter-Science Publishers, Mew York.  1145p.

Staler, R.,  G.E. Zaroogian, and R.J. Hennekey.  1962.  Cadmium uptake by
  marine organisms,  J. Fish. Res. Bd. Canada.  29:1367-1369.

Farrah, H, and W.F. Pickering,  1977,  Influence of clay solute inter-
  actions on aqueous heavy metal ion levels.  Water Mr Soil Pollut.

Friberg, L., M. Plscatoe, and G. Sordberg.  1971.  Cadmium in the en-
  vironment:  a foxieological and epldemlological appraisal,  pp.4-16 -
  4-41,  U.S. Environmental Protection Agency, Durham, N.C. (NT1S PB It-i
  795) 319p.

Fulkerson, U. and H.E. Goeller. (eds).  1973.  Cadaium:  the dissipated
  element,  pp. 279-323.  Oak Ridge National Laboratory, Oak Ridge,
  Tennessee.   ORNL-NSF-EP-21.  472p.

Gardiner, J.  19?4a«  The chemistry of cadmium in natural waters - I. A
  study of cadmium complex formation using the cadmium spec!fie-ion elec-
  trode.  Mater Research. 8(1):23-30.

Gardiner, J. 1974b.  The chemistry of cadmium in natural waters - It.  The
  adsorption of cadmium on river muds and naturally occurring solids.
  Water Research.  8:157-164,

Oiy, R.D. and C. L. Chakrabarti.  1976.  Studies of aetal-organic inter
  actions in model systems pertaining to natural waters.  Can. J. Chen.

Helz,  kR., R.J. Huggett, and J.M. Hill.  1975.  Behavior of Mn, Fe, Cu,
  Zn, Cd, and Pb discharged from a wasteuater treatment plant into an es-
  tuarine environment.  Water Research.  9:631-636.

Hera, J.D.  1972.  Chemistry and occurrence of cadmium and zinc  In  surface
  and ground waters.  Water Resources Research.  8:661-671.

Holmes, C.W. , S.A. Slade , and C.J. McLerram.  1974.  Migration  and re-
  distribution of zinc and cadmium in a marine estuarine ecosystem.
  Environ. Sci. Technol.  8(3): 255-259.

Huang, C. ?. , H.A. Elliott, and R.M. Ashraead.  1977.  Interfacial reactions
  and the fate of heavy aetals in soil-water systems.  Jour. Water Poll.
  Control Fed. 49(5): 745-756.

Huekabee, J.W. ana B.C. Blaylock,  1973.  Transfer cf mercury and cadmium
  from terrestrial to aq-jatic ecosystems.  Adv. Exp. Med. Eiol. 40:125-160,

Kelso, J.R. and R. Prank.  1974.  Organochlorine residue, mercury, copper,
  and cadmium in yellow perch, white bass, small mouth bass, in Long Point
  Bay, Lake Brie,  Trans. Am. fish, Soc,  1Q3(3):577-581.

Kinkade, M»L, and H»E. Srdman.  1975.  The influence of hardness components
  (Ca4"^ aru$ Mg+2) in water on the uptake and concentration of cadmium
  in a simulated fresh water environment.  Environ. Res. 10(2): 308-313.

Kubota, J. , E.L. Mills, and R.T. Ogleoby.  1974.  Lead, Cd, Zn, Cu, and Co
  in streams and lake waters of Cayuga Lake Basin, New York.  Environ. Sci.
  Techol.  8(3): 243-248.

Long, D.T. and E.S. Angino.  1977.  Chemical speciation of Cd,  Cu, Pb» and
  Zn in mixed fresh water, seawater, and brine solutions,  Geochim.
  CosmochitB. Acta.  41:1183-1191.

Lovett, R.J, , W.H. Gutennann, 1.5, Pakkak, W.0. Youngs, D.J. Lisk, G.S.
  Burdick, and E.J. Harris.  1972,  A survey of the total cadmium content
  of 406 fish from 49 S«w York State fresh waters.  J. Fish Res, id. Can.
  29(9): 1283-1290.

Lu, J. C. S. and K.Y, Chen,  1977,  Migration of trace aetals in interfaces
  of seawater and polluted surficial sediments.  Environ, Sci. Technol.
Lu, P.Y., R.L. Metcalf, R. Furman, R. Vogel, and J. Hassett.  1975.  Model
  ecosystem studies of lead and cadmium and of urban sewage sludge
  containing these elements.  J. Environ. Quality  4(4);505-509.

Martin > J.H. and W.W. Broenkow.  1975,  Cadmium in plankton:  elevated con
  centration* off Saja California, Science. 190:884-885.

O'Stiea, T.A. and K.H. Mancy,  1978.  The effect of pH and hardness metal
  ions on the competitive Interaction between trace metal ions and
  Inorganic and organic completing agents found In natural waters.  Water
  Research. 12;703-711.

Paacoe, D. and D.L. Mattey.  1977.  Studies on the toxicity of cadmium to
  th* three-splned stickleback, Castensteus aculeatus L.  J. Fish. Biol.

Patterson, J.W., H.E. Allen, and J.J. Scala.  1977.  Carbonate precipita-
  tion for heavy metal pollutants.  Jour. Water Poll. Control Fed.

Ptrhac, R,M.  1972.  Distribution of Cd, Co, Cu, Fe, Mn, Mi, Pb, and Zn in
  dissolved arid particulate solids from tvo streams in Tennessee.  Jour.
  Hydrol.  15:177-186.

Perhac, R.M.  1974a.  Hattr transport of heavy metals in solution and by
  different sizes of particulate solids.  Univ. of Tenn. Water Res.
  Research Ctr., Knoxville, Tenn.,  Project No. 023.  41p.

Ptrhac, R.M.  1974b.  Heavy metal distribution in bottom sediments and
 . wtstes in the Tennessee River - ioudon lake Reservoir system.  Univ. of
  Tenn. Hater Res. Research Ctr. Rnoxville, Tenn.  Research Report No. 40.

Phillips, A.R. and R.C. Russo.  1978.  Metal bioaccumulation in fishts and
  aquatic invertebrates:  a literature review,  pp.13-20.  U.S.
  Environmental Protection Agency, Duiuth, Minn.  (EPA 600/3-78-103).

Ramaaoorthy, S. and 8.R. Rust.  1978.  Heavy metal exchange processes in
  •ediaent-water systems.  Environ. Gcol. 2(3}:165-172.

Schell, W.R. and A. Nevissl.  1977.  Heavy metals from waste disposal in
  Central Puget Sound.  Environ. Scl. Technol. 11(9):887-893.

Steele, K.F. and G.H. Wagner.  1975.  Trace metal relationships Jn bottom
  sediments of a fresh «ater stream - The Buffalo River, Arkansas.  Jour.
  Sed. Pet.  45<1}:31CKU9.

Suzuki, H.,T. Yamada, T. Miyazaki, and K. Kawazoe.  1979.  Sorption and ac-
  cumulation of cadaiun in Che sediment of the Tama River.  Water Research.

Valiela, I., M.D. Banus, and J.M. Teal.   1974,   Response of salt marsh
  bivalves to enrichment with metal containing sewage sludge and retention
  of lead, zinc, and cadmium by marsh sediments,,  Environ. Pollut.

Waklta, H. and R.A. Schmitt.  1970.  Cadmium in:  Handbook of geochemistry.
  (K.H. Wedepohl, ed).  Vol. II, Chap. 2,  Springer-Verlag'» New York.   23p.

Weast, R.C. ed«  1977,  CRC Handbook of  chemistry and physics, 58th
  Edition, CRC Press, Cleveland, Ohio.  2398p,

Weber, W.J; and H*S. Posselt.  1974. •  Equilibrium models and precipitation
  reactions for cadmium( 11) In;  Aqueous envlronoental chemistry of metals
  (A.J. Rubin, *d}»  Ann Arbor Science Publishers, Ann Arbor, Mich.

                               10,  CHROMIUM

10.1  Sja_tamenc of Probable Fate  ,

    Chromium1 exists in two oxidation states In aqueous systems:  Cr(lIX)
and Cr(VI).  The hexavalent fora 13 quite soluble, existing In solution as
a complex anlon, and Is not aorbed to any significant degree by clays or
hydrous metal oxides.  It Is, however, sorbed strongly to activated carbon.
Hexavalent chromium is a moderately strong oxidizing agent arid reacts with
reducing materials to form trivalent chromium.  Trivalent chromium reacts
with aqueous hydroxide ion to form the insoluble chromium hydroxide
(Cr(OH)3>.  Precipitation of this material is thought to be the dominant
fate of chromium in natural waters,  Sorption processes also result in re-
moval of dissolved chromium to the bed sediments.  Chromium forms complexes,
with a variety of organic materials.  The importance of these materials in
aolublliziog trivalent chromium is unknown, but is probably not signifi-
cant.  Chromium is bioaccumulated by aquatic organisms and passage of
chromium through the food chain has been demonstrated.

10.2  Identification - Geochemistry, of_Chromium

    Chromium, a transition element, occurs in nature principally as the
triv.alent ion Cr*-% although valence states ranging from -2 to +6 have
been reported.  Chromium is found in concentrations of about 10-100 ppo in
the crust and about 0.001-0.8 ppm in river waters (National Academy of
Sciences 1974).  The principal ehromium-be'arihg minerals belong to the
chromite spinel group with the general formula (Mg, Fe)0(Cr,Al,Fe)203»
Depending on the degree of substitution in the Al, Fe, Cr series, the chro-
sBites contain from 13 to 65 percent C^Qj (Towiil jet _al. 1978).  A
variety of chromium compounds are prepared from these chromites.  Host of
these compounds contain chromium in the stable trivalent and hexavalent
oxidation states.

    The geochemistry of chromium Is dominated by the ability of the tri-
valentoion Cr(III), with a jadius of 0.64 A to substitute for Fe(Itl)
(0,67 A) and Al(tll) (0.56 A) during crystallization.  Chromium typi-
cally is precipitated from magmas at an early stage, either in the chromite
spinels or In silicate minerals, especially cllnopytoxene,  Chromite is
generally resistant to chemical weathering.  Due to its high specific grav-
ity, it may be mechanically concentrated in laterites or heavy mineral
placers.  The chromium-bearing silicates release chromium which is then
incorporated Into shales and schists.  Little chromium becomes solubi-
lized, and thus, geological precipitates and evaporates have a low chromium

    Chromium, atomic number 24, has an atomic weight of 51.996 (Weast
1977). The metal has a melting point of i8,57"C, a boiling point of 26?2°C
and a specific gravity of 7.20 at 20"C (Weast 1977).

    Chromium forma thousands of chromium (III) complexes, almost all of
which are hexacoordinate (Cotton and Wilkinson 1972).  In .aqueous solu-
tions, the principal characteristic of these complexes is their relative
kinetic inertness, even under conditions where they are thermodynamlcally
unstable (Cotton and Wilkinson 1972).  The importance of complexation in
determining aquatic fate is unknown, but it Is probably not significant
relative to oxidation, precipitation, and sorption reactions discussed

    The CAS number of chroniua is 7440-47-3; the TSL number is A 431-4218.

10«3  Suamary of Fate Data

    10.3.1  Photolysis

         No data were found that would indicate that photolysis of chromium
compounds plays an important role in determining aquatic fate.

    10.3.2  Chemical Speciation
         The inorganic chemistry of chroniua has been well studied and
documented; however, its biological and environmental interactions are ob-
scure and poorly characterized.  This dichotomy is the direct result of the
chemical complexity of the element and the extremely low concentrations of
chromium found in the environment.  Chromium occurs in valence states rang-
ing from -2 to +6.  The tripositlve state (the most stable form) exhibits a
strong tendency to form hexacoordinate octahedral complexes with a great
variety of llgands such as water, ammonia, urea, halides, sulfates, ethyl-
enediamine, and organic acids.  In neutral,and basic solutions, trivalent
chromium forms polynuclear compound* in which adjacent chromium atoms are
linked through OH or 0 bridges.  These compounds nay eventually precipitate
as Cr2
         Hexavalent chromium, Cr(VI), is a strong oxidizing agent, and is
always found in aqueous solution as a component of a complex anion.  The
anicmic form varies according to pH, and may be chr ornate (CrQ4~*), !
hyd roc hr ornate (HCrO^"), or dichromate (C^Oj"^).  Dichromate con-
centration is not significant unless pH values are well below those ob-
served in' most natural waters.. ' Thus, hexavalent chromium present in most
natural waters (pH>6.5) will b€ in the form of the chromate ion,
CrQ4~2.  All of the anionic forms are quite soluble, and are thus quite
nobile in the aquatic environment (Towlli ejt al. 1978).

         Schroeder and Lee (1975), in a laboratory study on the transfer-
nation of chromium in natural waters, found that Cr(III) and Cr(VI) are
readily intarconvertible under natural conditions.  Their results indicated
that Cr(VI) can be reduced by Fe(II), dissolved sulfldes, and certain
organic compounds with sulfhydryl groups, while Cr(III) can be oxidized by
a large excess of Mn02 and at a slower rate by 02 under natural water
conditions. Moreover, if aquatic conditions favor Cr(VI), then chromium
will accumulate as soluble forms in waters; if, however, Cr(III) is
favored, then the accumulation will occur in the sediaientst

         This envtroiaaental accumulation of Cr(III) In the s*dinients has
be«n noted by several authors (Nelson and'Hanshild 1970; Bruland et al.
1974; Perhac 1974; Morel jet al. 1975; Rehwoldt jet al. 1975; S-eele and
Wagner 1975; Samminga and Wilhn 1977) and can be explained by the hydroly-
sis of Cr(IH) complexes to insoluble hydroxide forms, especially
         It appears, therefore, that chemical speciatlon plays a dominant
rol« in the fate of chromium ir. the aquatic environment.  Conditions favor
able to Cr(VI) will keep chromium in a soluble form in the water, while
conditions favorable to Cr(III) will lead to precipitation and adsorption
Of chromium in Che sediments.

    10.3.3  Volatilization

         Mo data were found that would indicate that volatilization of
chromium compounds plays an important role in determining aquatic fate.

    10.3.4  Sorpttpn

         Hexavalent chromium is not adsorbed to any significant degree by
clays, ferric hydroxide, or ftrrlc and manganese oxides (Kharkar et al.
1968).  IE is efficiently removed by activated carbon (Linstedt et al.
1971), and thus may have some affinity for organic materials in natural

         The fractional extraction of sediments indicates thar surface
adsorption, which is a relatively weak binding process, does not account
for most of the chromium associated with sediments.  Furthermore, there is
generally a strong inverse correlation between chromium concentration and
sediment grain size (Nelson and Haushild 1970; Perhac 1974;  Sceele and
Wagner 1975).   Besides precipitation, several sorptlon mechanisms have been
postulated to explain these observations.

         Steele and Wagner (1975) noted that there was a good correlation
between extractable chroaiua and extractable iron1 in sediaents from an
Arkansas river, and suggested that Incorporation of chromium into hydrous
iron oxides was probably the reason for this.  Their extraction technique
used aqua regia, which is undoubtedly capable of solubillzing Cr(OH>3» as
well as the chroaiua incorporated with the hydrous Iron oxides.  Also, the
fact that chroaiua is mineraioglcally associated with iron implies that
chromium introduced into the streaa by weathering would be precipitating
out In the same areas as iron.  Thus, another way to explain Che results
would be that the correlation between Cr and Fe is not due to removal of
dissolved Cr by precipitating Fe, but is a result of the fact that
Cr(QH>3 is precipitated in the same areas as Fe(OH>3.  Coprecl?itation
of these materials may increase the speed with which chromium is removed
from solution.

         Perhac (1974) found that very little chromium in sediments was
bound up In iron oxides.  He extracted the iron oxides with sodiuta dithio-
nlte, which reduces the ferric iron to ferrous iron and thus destroys the
hydrous iron oxide coatings,  Cr(OH>3 would probably not be salubilized
by such an extraction procedure.

         A pertinent observation was reported by Griffin e_t a 1. (1977) in
their laboratory study on the effect of pH on the adsorption of chromium by
clay minerals.  Since this study was carried out In soils, direct extrapo-
lation cannot be made to the aquatic environment; however, the physico-
chemical generalizations,should be applicable.  They found that adsorption
of Cr(Vl) decreased as pH increased and that the HCr(>4~ ion was the
Cr(Vl) species predominantly adsorbed.  The adsorption of CrClil), how-
ever, increased as the pH increased.  About 30 to 300 times more Cr(III)
than Cr(VI) was adsorbed by clays, and the amounts of Cr(III) adsorbed
corresponded to cation exchange adsorption of hydrolyzed Cr(IIl) species.
These results suggest that while Cr(VI) is highly mobile, Cr(III) will be
quickly immobilized into the sediments.

         Gibbs (1973) studied transport of trace metals in the Yukon and
Amazon Rivers.  He concluded that chromium was transported primarily in
crystalline sediments, with transport as dissolved species and biological
solids running a distant second and third.  If chromium is indeed trans*

ported to any appreciable degree in crystalline sediments, 1C Is possible
that 1C is due Co isomorphous substitution of Cr(III) for &1(III) and
Fe(III).  The importance of this process to the environmental transport of
chromium is still unclear.

         In summary, it appears that Cr(III) and Cr(VI) are only w«akly
adsorbed into inorganic «olids, although Cr(III) is adsorbed «ore strongly
that Cr(?I).  Sorption of Cr(III) may be ancillary to precipitation of
    10.3.5  Bioaceua ulation

         Chromium is an essential nutrient (National Academy of Sciences
1974), and it it accumulated in aquatic and marine biota to levels much
higher than in krabient water.  Levels in biota, however, are usually lower
than levels in tht sediments,

         Mamlnga and Wilhm (1977} studied heavy octal partitioning between
water, sediments, and chironomid larvae (a benthic invertebrate).  They
found an average chromium concentration of 1, lug/1 in water, 7.&4yg/g in
sediments and 2.96 pg/g in chironomid s.  Bloboncentration factors far
chironomids to water are thus about 3,000, and for chironoraids to sedi-
ments, about 0.39.  Sahwoldt e£ _al. (1975) found similar relationships
aaong water, sediments, and biota in the Danube River.

         Patrick and Loutit {1976} examined the ability of bacteria to
mobilize metals by accumulating them and passing the-a up the food chain.
Tubificid worms (benthie) were f«d bacteria tha.t had accumulated chromium
and retained some of the element.  Xubificids are apparently able to ex-
crete cnromium more effectively than the bacteria, because the concentra-
tions in the worms were lower than concentrations in the bacterial cells.
Nevertheless, the experiment proved that chromium can be passed on through
the food chain.  Accumulation of metals by benthie species may result in
chromium mobilization through the biota.

         Promo and Stokes (1962) found that rainbow trout took 10 days to
reach whole-body equilibrium concentration upon exposure to hexavalent
chromium levels below 0,01 ag Cr/1.  Fish exposed to chromium concen-
trations of 0.05 mg/1 and higher, however, continued to accumulate chrom-
ium linearly in time until the test was terminated after 30 days.  In a
laboratory study, Buhler e_t al. (1977) analyzed two groups of rainbow trout
raised in two natural waters differing in chromium content.  The trout
accumulated chromium rapidly during the first day of exposure hut did not
accumulate much more chromium during further exposure for 22 days.  Appar-
ently, an equilibrium condition was rapidly reached.  The trout contained
chromium levels 'in proportion to -fie chromium in their environment.

         Saptist and Lewis (1969) studied the transfer of radiolabeled
Cr(III) in an estuarlne food chain consisting of phytoplankton,  brine
shrimp, post-larval fish, and aumaichog.  Chromium was transferred through
the food chain through each trophic level,, with concentrations declining as
trophic level increased.  Theoretical calculations indicated that, in gen-
eral, the food chain was a more efficient pathway for uptake of  chromium
than direct uptake from seawater.

         Distribution of chromium in water, sediment, seston (suspended
abiotic and biotic material), phytoplankton, mollusks, annelids, and fish
in Narragansett Bay, R. I. , was studied by Phelps e_t_ a_l. (1975).   The high-
est concentrations of chromium were found in the sediments, followed by the
seston.  Phytoplankton concentrated chromium to a greater extent than other
organisms, with the lowest levels being found in bottom-feeding  fish.

         Some bioconcentration factors reported for chromium are given in
Tale 10-1.  The range in concentration ratios probably reflects not only
differences among, taxa, but also differences in ambient water concentra-
tions of chromium.

    10.3.6  Biotransforffldtion

         No data were collected concerning the importance of biodegradation
of chromium compounds on aquatic fate.  There has been some.speculation
that chromium could- be methylated in reducing environments (Anon. 1977),
but no evidence was found that this process occurs in natural or experi-
mental systems.  Under anaerobic conditions, there is a possibility that
Cr(VT) species such as HCr04~ (hydrochromate) and Cr04~2 (chromate)
could be utilized by bacteria and other anaerobes as an oxygen source
(Adaas «£_al. 1975).  Chemical reduction to Cr(III) with concomitant loss
of oxygen would be indistinguishable from this effect, and certainly
occurs In such environments.  Biogenic complexing agents may have some
effect on chromium distribution, especially in eutrophic systems which
typically have high concentrations of organic material.

10.4  Data Summary

    Most of the trlvalent chromium in the aquatic environment is hydrolyzed
and precipitates as Cr{OH)j.  Sorptlon processes and bioaccumulation will
remove the remaining Cr(III) from solution.  Under certain natural water
conditions, chromium can exist in the hexavalent form,  Cr(Vl) exists as an
oxyanion in aqueous solution and Is quite soluble.  It has little affinity
for clays and other inorganic surfaces, although it is strongly  sorbed by
activated carbon.  It is very toxic to aquatic organisms.  Processes re-
lating to chemical speciation are important in 'determining the aquatic fate
of chromium, which is summarized in Table 10-2.

                                Table 10-1

                   Bioeoncentration Factors for Chromium

Freshwater fish

Freshwater invertebrates

Freshwater plants

Marine fish

Marine invertebrates

Marine plants

Benthie algae



Mollusc viscera

Crustacean ousel*

Fish muscle
iioconcentratlon Factor3













Chapman ejc al. 1968

Chapman jet al. 1968

Chapman e£ al. 1968

Chapman et al. 1968

Chapman e_t al. 1968

Chapman _et al. 1968

National Academy of
   Sciences 1974   '
National Academy o£
   Sciences 1974
National Ac ad eta y of
   Sciences 1974
National Academy of
   Sciences 1974
National Academy of
   Sciences 1974
National Academy of
   Sciences 1974
a.  Concentration factors are defined by the ratio of the concentarion of
    the element in the organism In ppm (wet weight) divided by the con-
    centration of the element in water (ppm).

                                Table 10-2

                    Summary of Aquatic Fata of Chromium





Not an important process.

An important consideration in the
aquatic fat* of chromium.  Con-
trols the intertranaforaation of
Cr(VI) to ,Cr(III).  Cr(VI) remains
soluble, while Cr(III) will hydrolyze
and precipitate as
Confidence of


Mot an important process.

Cr(III) is adsorbed weakly to in-
organic materials.  Cr(VI) may be
adsorbed by organic materials.

As an essential nutrient, chromium
Is bioaccumulatad by a variety of
aquatic organisms.  May be transferred
via the food chain.
Biotransformation   Probably not important.



    All of the noted environmental processes are important; however, their
    relative importance with respect to each other is uncertain for
    determining final fate.

id, 5 Lj. tg_r a t u r e Cited

Adams, C.E., Jr., W. Eckenfelder, Jr., and B.L. Goodman.  1975.  The
  effects and removal of heavy metals In biological treatment:,  pp.
  277-292.  _in Heavy Metals in the Aquatic Environment.  P.A. Krenkel
  (ed.). Pergamoti Press, Oxford, England.  352p,

Anon.  1977.  Tracking trace metals in the biosphere,  Cham. Eng. News

Baptist, J.fr. and C.W. Lewis.  1969.  ' Transfer of 65Zn and_5lCr through
  an estuarine food chain,  pp. 420-430.  Proc. 2nd Nat. Symp. on
  Radloecology.  U.S. Atomic Energy Commission Conf.  670503.  631p.

Bruland, K.W., K.K. Bertine, M. Koide, and E.D. Goldberg,  1974.  History
  of metal pollution in Southern California coastal zone.  Environ* Sci.
  Technol. 8(5):425-432.

Buhler, O.K., R.M. Stokes, and S.S. Caldwell.  1977.   Tissue accumulation
  and enzymatic effects of hexavalent chromium in rainbow trout (Salmo
  gairdneri).  J. Fish. Res. Bd. Can. 34(1}:9-18.

Chapman, W.H., H.L. Fisher, and ,M.W.  Pratt.  1968.  Concentration factors
  of chemical elements in edible aquatic organisms.  Lawrence Radiation
  Laboratory, tivermore, Calif.  UCKL-50564.  46p.

Cotton, F.A. and G. Wilkinson.  1972.  Advanced inorganic chemistry,  pp.
  830-845.  Interseience Publishers,1 New York.  114Sp.

Fromm, P.O. and R.M. Stokes.  1962.  Assimilation and metabolism of
  chromium by trout.  J. Water Pollut. Control Fed. 34(11):1151-1155.

Gibba, R.J. 1973.  Mechanisms of metal transport J.n rivers.  Science

Griffin, R.A., A.K, Au, and R.R. Fro«t.  1977.  Effect of pH on adsorption
  of chromium from landfill-leachate  by clay minerals.  J, Environ. Sci.
  Health Al2(8):431-449.

Kharkar.D.P., K.K. Turekiar, and K.K. Bertiue.  1968.  Stream supply of
  dissolved Ag, Mo, Sb, Se, Cr, Co, Rb, and Cs to the oceans.  Geochim.
  Cosaochin. Acta 32:285-298.

Llnstedt, K.D., C.P. Houck, and J.T.  O'Connor.  1971.  Trace-element
  removals in advanced wastewater treatment processes.  J. Water Pollut.
  Control Fed. 43(7)-.1507-1513.

Morel, F.M.M.» J.C-  Westall, C.R. Q'Melia, and J.J. Morgan.  1975.  Face of
  trace aetals In Los Angeles County wastewater discharge.  Environ. Scl.
  Technol. 9(8):756-761.

Namiainga, H. and J. Wilha.  1977,  Heavy metals in water, sediments, and
  chironomids.  J. Water Pollut. Control Fed. 49(7):1725-1731.

National Academy of Sciences.  1974,  Chromium,  pp.  86-89. U.S. Goverranent
  Printing Office, Washington, D.C.  155p.

Mel son, J.L. and W»L« Haushild.  1970.  Accumulation of radionuclides in
  b*d sediments Of the Coluobia liver between the Hanford reactor and
  MeNay Dam.  Water Resources Res. 6(1):130-137.

Patrick, P.M. and M. Loutit,  1976.  Passage of aetals in effluents through
  bacteria to higher organisms.  Water Res* 10:333-335.

Perhac, R.M.  1974*  Heavy metal distribution fn bottom sediments and water
  in the Tennessee River-Loudon Lake Reservoir system.  Univ. of Tenn.
  Water Resources Research Center, Knox^iHe, Ter.n.,  Research Report No.
  40.  22pp.

Phelps, O.K., G. felek, and R,L. Lapan, Jr.  1975.  Evaluation of the
  distribution of heavy metals in the food chain.  Ing. Ambientale
  4(3):321-328. (Italian) (Abstract only). CA 19/o.  84:l31057f.

Rehwoldt, R. , 0. Karimain-Teherania, and H. Alttaann.   1975.  Measurement
  and distribution of various heavy metals in the Danube River and Danube
  Canal:  aquatic communities in the vicinity of Vienna, Austria*   Sci.
  Total Environ.  3:341-348.

Schroeder, D.C. and G.P. Lee.  1975.  Potential transformations of chromium
  in natural waters,  rfatar Air Soil Pollut.  4:355-365.

Steele, K.F. and G.H. Wagner,  1975.  Trace aetal relationships in bottom
  sediments of • freshwater stream.  The Buffalo River, Arkansas.  J.
  S«d. Pet. 45(1):310-319.

Towill, L.E., C.R. Shriner, J.S. Drury, A.S. Hawaons, and J.W. Holleman.
  1978.  Reviews of the environmental effects of pollutants:  III.
  Chrottlun.   pp. 12-36.  Oak Ridge National Laboratory, Oak Ridge, Tenn,
  PS 282 796 285p.

Weast, R.C. (ed.).  1977,  Handbook of chemistry and  physics, 58th
  edition.  CRC Press, Cleveland, Ohio.  2398p.

                               11.   COPPER

11.1  Seatenant _o_f_Probable Fate

    Several processes determine crwt fate of copper in the aquatic  environ-
ment;  'complex formation, especially with huraie substances;   sorpci<.-n to
hydrous metal oxides, clays, and organic materials;   and bloaccumulation.
The formation of complexes with organic  ligands modifies the solubility and
precipitation behavior of copper such that solid copper species, probably do
not precipitate under normal circumstances.  Furthermore, complexed. copper
Is more easily adsorbed by clay and other surfaces than the  free (hydrated)
cation.  The aquatic fate of copper is highly,dependent on such variables
as pH, Eh, concentrations of organic materials and adsorbents,  availability
of precipitating iron and manganese oxides, biological activity, and com-
petition with other heavy metals.

    Sorptlon of copper by precipitating  hydrous Iron and manganese oxides
is an effective control OR dissolved copper concentrations where these
metals are being actively weathered or otherwise introduced  into unpolluted
aquatic environments.  In organic rich environments, typical of polluted
natural waters, the effective control on dissolved copper concentrations
will be the cdmpetttton between organic  coraplexing in solution and sorptlon
onto clay and particulate organic material.

    Copper Is strongly bioaccumulated and is an essential.trace element;
however, high concentrations of Cu(II) ion are toxic to aquatic organisms.
Biological activity, as a source of organic ligands, plays an important
part in determining the aquatic fats of  copoet.

11,2  Identification - Geochemistry of Copper

    Copper is a metallic element and a member of the first transition
series.  It exists in the llthosphere primarily as a sulfide, both as the
simple1 sulfide and as a great variety of complex sulfide minerals that
include other metals.  Sy far the most abundant of the copper minerals is
ehalcopyrite (CuFeS2>, although metallic copper, chalcocite  (Cu2S), and
bornite (CujfeSi,) are also found in economically important deposits.

    Copper is present in concentrations averaging about 4 ppm in lime-
stones, 55 ppm in Igneous rocks, 50 ppm in sandstones and 45 ppm in shales
(Krauskopf 1972).  The marked concentrations of copper In shales and sand-
stones suggest that copper in the lithosphere exists largely as  adsorbed
iona, fine grained particles or as one of many discrete sedimentary copper
minerals.  Generally, these minerals occur only as sparse tiny grains that
are widely disseminated throughout the sedimentary rocks.

    Reactions leading to precipitation of definite copper compounds, how-
ever, are probably not common in most sediments and, almost certainly, are
less effective than adsc.rption as a general mechanism for removing copper
from solution.  Copper is most strongly adsotbed by the surfaces available
In neutral waters.  Because Cu** forms so readily during weathering, and
because it can persist in acidic oxidizing solutions at fairly high concen-
trations, copper la considered to be among the more aiobile of the heavy
•aetals in surface environments.  The distance it can travel is limited
largely by its strong adsorption to many kinds of surfaces.  Ferric hydrox-
ide, for example, is quite an effective adsorbent of Cu4^ provided the
pH is above the isoelectric point of the hydroxide (Hem and Skougstand

    Copper, atonic number 29, has an atomic weight of 63.546 (Weast 1977).
It. fora s selts and complexes wicii valences of +1, +2, and, very rarely, +3
(Cotton and Wilkinson 1972).  The electrochemical properties of copper are
well known.

    The CAS number for copper is 7440-50-8, and its TSL number is

11.3  Suaaary of Fate 'Data

    11.3.1  Photolysis

         Although some copper cotaplexes are photosensitive, no evidence was
found indicating that photolysis is an important mechanism in determining
the aquatic fate of this metal.

   ' l1*3-2  Chemical.^Speciation

         In aqueous solution, copper is present as Cu(ll), since the only
cuprous (valence +•!) compounds stable in oxic waters are those that are
highly insoluble (e.g., CuCl or CuCN)(Cotton and Wilkinson 1972).  Although
aost cupric salts are not considered to be readily water-soluble, there are
several exceptions, including cupric chloride (CuCl2)» cupric nitrate
(00(803)2) and cupric sulfate (UuS(>4).

         Copper has a pronounced tendency to fora complexes with both
organic and inorganic ligands.  Stiff (197la, 197lb) found th*t, at pH
values and Inorganic carbon concentrations characteristic  of natural
waters, most of the copper in solution is present as complexes of cupric
car'aonate rather than sa the thydrated) divalent cupric ion.  Stiff
(1971a), in a laboratory study of unpolluted waters, found that the copper
in these waters would most likely be present as complexes  of cupric car-
bonate.  Stiff (I97lb) extended this study to polluted waters and showed

that the predominant species of soluble copper in polluted environments
would be complexes with cyanide, amino acids, and humic substances as well
as the complex carbonates and the (hydrated) divalent cupric ton.  In a re-
lated field study reported by Stiff (1971fa), It was demonstrated that much
of the copper present in polluted English rivers was associated with sus-
pended solids and that soluble copper consisted almost entirely of con-
pi exed organic forms.

         Sylva (1976) examined the speciation of copper(II) in fresh water
with respect to inorganic and organic cotnplexation and adsorption and pre-
cipitation.  It was found that these processes are capable of reducing the
level of soluble copper to very low values even in the presence of high
levels of total copper.  Hydrolysis and precipitation doolnate the chemis-
try of copper(II) at pH values expected in most natural water systems
wherever there is a limited amount of organic coaspiexing agents (Figure
11-1).  Ih* most significant process by which divalent, hydrated copper(II)
is removed from unpolluted -«ter is the precipitation of malachite
(Cu2(OM>2C03).  The rate of this precipitation, however, is very slow
at low copper levels and Ch« equilibrium situation aay not always be
reached or even approached because of fluctuating conditions.  The effect
of the presence of organic coaplexing agents can change the system to such
an extent as to alter greatly the results plotted in Figure 11-1, especial-
ly at the lower pH values.  Thus, the speciation of ccpper(ll) can vary
considerably from one natural water system to another and also within one
given system over a period of time.

         Long aad Angino (1977) developed a theoretical model to study the
chemical speciation of copper in aqueous solutions and the response of
copper to variations in ionic strength and complexation.  Association of
copper with the ligands OH", Cl~, C03~2, SC>4*2 and HCC>3~
was considered at pH values troo 3,5 to 11.0 at 25*C in differing seawater-
freshwater mixtures.  The results are summarized in Figure 11-2.  In gen-
eral, thct relative importance of the various llgand-copper complexes can be
predicted from * comparison of their stability constants. This model, how-
ever, dots not take into account metal-organic complexes and it is, there-
fore, useful only in unpolluted, relatively organic-free waters.

         In most surface waters, organic materials prevail over Inorganic
ions in coopltxing copper,  Ramamoorthy and Kushner (1975) demonstrated
that almost all of the heavy-metal binding capacity of Ottawa River water
was due to organic substances.  They calculated an empirical equilibirum
constant from the equation:

         K * . 	(.Metal ion bound]		
                [Metal ion unbound][River component unbound}.

Figure 11-1   Speciation of copper(II)(total concentration 2  ppm)  and
              carbonates as a function of pH.  (A) Cu*2.  (B) (^{GH^*2.
              2 (azurite) will precipitate.  (I) pH at
              which Ci*2(QHHc^3 (nalachite) will  precipitate.  From Sylva

                                           100* Fresnwttr
                                   7 iCHi_    5
         The value,of K  for copper was  5.01 +  1.38 x  10  and  the  calcu-
 lated concentration of the river binding component was  2.54 x 10-'M.

         Lopez and Lee (1977),  in their study  of,a heavily copper-
 polluted, organic-depleted lake  in Michigan found that  the predominance  of
 soluble  copper species followed  the order: Cu(QH)+ >  Cu+2 > CuC£>3%
 It  appeared  further that copper  concentration  in this lake was controlled
 by  hydrous oxides of iron and manganese and not by the  solubility of  copper

         Hem (1975) calculated  the predominance of dissolved  copper species
 and the  stability  fields for solid copper compounds In  a system with  total
 dissolved carbon equal to 10"^M  and total dissolved sulfur equal  to
 10~4M.   ,The Eh-pH diagrams for  this system are shown  as Figure 11-3 and

         The strong tendency of  copper  to form complexes has  important ram-
 ifications in  its  precipitation  and sorption behavior and i,s  a most impor-
 tant process .for considering the aquatic fate  of copper.

     11.3.3  Volatilization

!         No  evidence was found  to indicate that volatilization of copper
 compounds is an important aquatic fate.

     11.3.4  Sorption

         Copper has a strong affinity for hydrous iron  and manganese
 oxides,  clays, carbonate minerals, and  organic matter.  Sorption  to these
 materials, both suspended in the water  column  and in  the bed  sediments,  re-
 sults  in relative  enrichment of  the solid phase and reduction in  dissolved

         Hem and Skougsead (1960) demonstrated that copreeipitation of cop-
 per with the hydrous oxides of  iron effectively scavenges copper  from solu-
 tion.  These materials form a coating on solid surfaces in the water, and
 as  they  precipitate, copper and  other metals' tend to  be attracted due to .
 the negative zeta-potential usually exhibited  by the  hydrpus  iron oxide
 (Jenne 1968).  Copper may thus  be incorporated into the lattice structure
 of  the hydrous iron oxide coating, this process being known as copreeipita-
 tion.  Thus, the hydrous iron (and to a lesser extent,  manganese) oxides
 can control  the mobility of copper in natural  waters.   Jenne  (1968)',  and
 more recently  Lee  (1975), have  presented convincing evidence  that such
 hydrous  metal  oxides ate important controls on the mobility of copper' and
 some _>ther metals  in unpolluted  aqueous and soil environments.  In reducing
 or  acidic environments,  such as  in richly organic bed sediments,  these
 oxides can be  dissolved, resulting in remobilizatlon  of sorbed or copreci-,

                 ,,1.   !._  ,
Figure 11-3  Eh-pH diagram showing azeaa of doainance of five
             species (solute) of copper at equilibrium at 25 C
             and 1 atn.   System Cu-H-G-C-S; total dissolved
             C-IO"3'00 moles/1; total dissolved S-1CT4-00 «ol*s/l.
             From Hem (1975).

                        WATf H QX1DJ2ED


Figure 11-4   Eh-pH diagram showing fields or stability of solids
              and total equilibrium activity of dissolved copper
              at  25°C and 1 atm.   System Cu-H20-C-S;  total dissolved
              OiQ-3,00 moies/l;  total dissolved S-IQ-4-00 moles/1.
              Frotn Hem (1973).

pitated metals.  Several Investigations1 have given evidence for this proc-
ess by reporting a high correlation in the sediments of natural streams be-
tween copper content and iron and nanganeae content (Carpenter et al. 1975;
Steele and Wagner 1975;  Collins 1973).  There is substantiation^ there-
for*, that coprecipiCation of copper by hydrous iron and manganese oxides
is an important process for removing copper from solution in some natural

         Copper is adsorbed to clay and mineral surfaces (Huang et al.
1977) and organic materials (Rashid 1974;  Baker-Blocker et _a_l, T97TJ.
Huang et al. (1977) demonstrated that adsorption of copper to soils'and
aluminum and silicon1 oxides is strongly pH-dependent, as shown in Figure
11-5,  Furthermore, the addition of various anions significantly increased
adsorption.  Humtc acid was particularly effective in this regard (see
Figure 11-6).  Huang j!t al. (1977) hypothesized that the enhanced adsorp-
tion due to the anions resulted from formation of a metal-ligand bond,
followed by adsorption to the hydrous solids through specific chemical
bonds such as "sharing of free electrons available from the tnetal-ligand
complexes."  Thus, the ease with which copper forms complexes with organic
and inorganic ligands (as discussed in Section 11.3.2) undoubtedly facili-
tates its adsorption by solids in natural waters.  Payne and Pickering
(1975), in their laboratory study on the removal of Cu(I.l) species from
aqueous solution by kaolinite clay suspensions, found that the extent of
copper removal was increased by the presence of ligands.  They found that
the important processes for determining the extent of copper adsorption
were solution pH, the nature of the ligands present and the order of con-
tact of the species.  They also reported that, in the presence of organic
ligands at a pH>6, there was virtually a total removal of copper.  In the
highly calcareous lake Monona, Wisconsin, 6.8 x 10 ^ k,g (1.5 x 10^ Ibs.)
of copper sulfate has been added over the last 50 years as an algicide.
Sanchez and Lee (1973) showed that most of the copper in bed sediments from
thii lake was found in the crystal structure of carbonate minerals.  These
investigators hypothesized that copper substitutes for calcium and raagnesl-
um atoms in the carbonate lattice structure.

         Ramamoorchy and Rust (1978), in their study of heavy metal ex-
change processes in the organic-rich sediments from the Ottawa River,
Canada, found that- the ability of the sediments to sorb copper ions was,
directly1 related to the amount of organic materials present.  Unless strong
leaching agents (in cnis study Nad and NfA at about 10~% concentration)
are present, the mobility of copper ions is low and they persist in the
sadinents for a considerable period of time.  Therefore, relative to fresh-
water environments, marine ecosystems should be subject to greater desorp-
tion of copper into the aqueous phase because of chloride complexlng and a
reduced degree of bonding to sediment particles.  Nonetheless, because of

Figure 11-5  Adsorption  of  copper on four solid suspensions
             after 24 hour  exposure.

             Concent:ration  of solids - gtn/1
             [CUT-J -  10**%
             Ionic strength - iO""1}* (NaCl)

             The upper  liaiit of the graph, 200 ymol/gm, represents
             10QZ adsorption.  Froia Huang eit al. (1977).

  HUM1C ACID    '
  TKIACITATI  /=///.'
     _^-^^*^  .^r **
Figure 11-6   Effect of anlons on  the  adsorption of copper by
              Mctapeak toll..  In addition to the experimental
              conditions shown in  Figure 11-5, the anion  con-
              centration was 10~^M except for humlc acid,
              which was 50 og/1.   From Huang e_t al_. (1977).

the somewhat 3Catic nature of marine depositional environments, it' is pos-
sible that coloride-complexed copper can remain in interstitial waters and .
escape only slowly into the overlying water column.

         Jackson and Skippen (1978), in a laboratory study of the disper-
sion of heavy metals via organic acids at the sediment-water boundary, dem-
onstrated that the organic acids increased Che solubility of copper in che
presence of clay,   Hence, an influx of soluble organic matter into streaa
water will favor the prolonged dispersion of copper in solution. Further-
more, humic and fulvic acids, when in excess of the copper ions, have the,
potential to retain copper in solution in competition with hydrolysis and
sorption onto clay. , Sense of this effect might be due solely to the
lowering of pH which will decrease copper sorption; however, the greatest
effect is probably that of competition for clay adsorption1 sites and' of
organic-copper complexing reactions.  While these observations appear to
conflict with those discussed earlier, they can be rationalized by recog-
nizing 'the differences in effects that will be mediated by soluble organic
compounds as opposed to particulate organic matter as described by Huang ec
ai. (1977) and Payne and Pickering (1975).  Jackson and Skippen (1978) also
reported that, although huraic and fulvic acids are capable of remobilizing
copper from a clay-sorbed phase and its- associated metal hydroxide precipi-
tates, the aesorptioti of copper is so kinetically inhibited as to be- almost

         In summary, sorption processes are quite active and efficient in
scavenging dissolved copper and in controlling its mobility in natural un-
polluted streams.   In unpolluted waters, the effectiveness ,of these pro-
cesses varies according to pH, Eh, and .the occurrence of potential sorption
surfaces.  In water polluted with soluble organic material, however, sorp-
tion appears toi be rather ineffective, -thus favoring the prolonged disper-
sion of copper in solution.  The presence of organic acids also can lead to
the mobilization of copper from the sediments to solution.

    11.3.3  Sioaccumulation
       ,     	....-MI.	.•..iiii.ii.nna	mimm	n	n	                          j

         As an essential nutrient, copper is accumulated by all plants and
animals.  Table 11-1 lists bioconcentratlon factors (concentration in
organism/concentration in ambient water) for some aquatic and marine

         Since copper is strongly bioaccumulated, ana because biogenic
ligands play such an important role in coraplexing copper (which affects  ,
precipitation and sorption behavior), biological activity is a major factor
in determining the distribution and occurence of copper in the ecosystem,
Kiasball (1973) studied seasonal fluctuations in copper concentration in a
pond ana found that concentrations were higher in fall and winter months

                                Table 11-1

                    Bioconcentration Factors  for  Copper

  Scenedesmus guadclearda
  Anafaa_ena varlabilia
.  Scenedesmus sp.
  Chlorella sp.


Plants,'Marine and Fresh









 Re f e re nee

Khobot'ev jet al. 1976
Khobot'ev et al. 1976
Stokes ££ al. 1973
Stokes jet al. 1973

Patrick and Loutlt 1976

Chapman jet ail. 1968

Chapman £| al. 1968
Chapman e_t_ al. 1968

layaoat 1972

         and Wilhn 1971
Chapman jet, al. 1968
Chapman et al. 1968
    Bioconcentration factor* are the ratio derived  from  the  concentrations
    of the element in the aquatic organism (in ppm  of wet weight)  divided
    by the concentration of the element in water  (in ppm).

than In spring and summer months.  Namminga and Wilhm (1977)  observed  the
same phenomenon In an Oklahoma stream.   Kimball (1973)  concluded that  the
reason for the seasonal fluctuation was chat copper became concentrated  in
vegetation during the growing season, and was released  from leaf licter  and
decaying aquatic plants in the fall.  Another possible  explanation is  that,
In the warmer months, there is a greater rate of  decomposition of organic
material with conoooitant release of humic substances.  As previously dis-
cussed, these substances can either adsorb copper directly or complex  it,
thus making it more available for adsorption on solids.  Probably both of
these hypotheses are operative and complement each other in causing ele-
vated levels of copper in fall and winter as compared co spring and summer.

         In their study of heavy metals in an Oklahoma  stream, N'amminga  and
tfilha (1977) found that chironomid larvae (benthic Insect forms) concen-
trated copper relative to the water column and the sediment matrix which
they inhabit.  The concentration of copper measured in  water, sediments,
and larvae was 4.lug/1, 1.8ug/g, and  1.91ug/g» respectively, yielding
a bioconcentration factor of 546 from water and 1.1 from sediments.

         N«hring (1976) suggested that  it may be possible to detect in-
stances of intermittently acute copper pollution in streams, by monitoring
copper levels in aquatic Insects. Some  stream insects,  Including the mayfly
(Spheaerella grandis) and the stonefly (Pteronarcys californica), were more
resistant to copper toxicity than fish, and copper residue  accumulation
affected the Insects' cupper exposure history.

         In a food chain consisting of  copper-enriched  sediment, bacteria,
and tubificid worta*, copper levels increased with increasing trophic level
(Patrick and Loutit 1976).  Windom e£ al. (1973) found, however, that  for
several North Atlantic fish species, copper level was Inversely related  to
trophic position.  Similarly, Cross ejc  al. (1973) ebserved no increase in
copper content with age among bluefish TPomatones saltatrtx)  and rnorids
(Antinora rostrata) collected off the North Carolina coast.

       ,  Since copper is toxic to aquatic life at high  concentrations, es-
pecially high concentrations of the divalent copper ion and its hydroxy
complexes, Brungs lit al. (1973) measured copper uptake  at several copper
concentrations by th* brown bullhead (Ictalurls neba1osus).  They hoped  to
establish an autopsy technique useful for confirming copper-caused fish
kills.  No useful relationship was found;  moreover, lethal exposure pre-
ceded by subacute expoaura resulted in higher .tissue copper levels than  in
fish hairing experienced only the lethal conditions.  Bullheads accumulated
copper at all water concentrations equalling or exceeding 27  wg Cu/1.
Copper concentrations In liver and gill tissues most accurately reflected
the copper exposure conditions.  Equilibrium concentrations were reached in
these tissues after 30 days'  exposure.

         Data in Table 11-1 also indicate that copper is not  biomagnified;
concentration ratios for fish (higher trophic levels) are lower  than con-
centration levels for algae (primary producers, i.e., lowest  trophic
level).  The apparent lack of biomagnification is not uncommon with the
heavy metals.  Furthermore, since copper is an essential nutrient, all
organisms have active transport mechanisms for it,  and there  is  no reason
to believe that differences in the physiological ability to excrete'copper
should be related to trophic Levvl.

    11,3.6  B jLo t ran 3 f o tma t i o n

         No evidence.was found to indicate that there is any  biotrans forma-
tion process for copper compounds which would have  a significant bearing on
the fate of copper in aquatic environments.

11.4  DataSummary

    Copper exhibits a very complex behavior in the  aquatic environment.
Sorption processes are probably most important ,in controlling copper dis-
tribution and include;  coprecipitation/sorption by hydrous iron and manga-
nese oxides;  ion exchange in the crystal lattice structure of carbonate
minerals;  adsorption to clays and other mineral surfaces; and adsorption
to organic solids.  Sorption appears to be more important than precipita- •
tion in most circumstances.

    Both organic and inorganic ligands complex copper.  Under normal con-
ditions, most of the copper in solution is in cotnplexed fora.  These com-
plexes alter the behavior of copper  to the extent that it is  generally aura
soluble in natural waters than would be predicted by conventional analysis
employing thenaodynamic equilibria,  and it has a greater adsorptive affin-
ity for hydrous solids than unc-omplexed forms.

    Seasonal fluctuations have been  observed in aqueous copper concentra-
tions with higher levels in fall and winter and lower levels  in  spring and
summer.  This probably reflects changes in bioaccumulation patterns':  dur-
ing the growing season, copper is taken up by biota;  during  fall and win-
ter, decomposing leaf litter and aquatic vegetation release copper.  The
availability of biogenlc ligands (e.g., humic and fulvic acids)  is probably
greater during the warmer part of the year, and this siay enhance adsorption
of copper.  At present, it is impossible to estimate how much of the copper
introduced' into the aquatic environment is partitioned into bed  sediaents
and biota and how much is transported by the water  column to  the oceans.
This undoubtedly varies widely with local conditions.

    Table 11-2 summarizes the aquatic fate information described above.



•Chemical  Speeiatlon3

                                 Table  11-2

                      Summary of  Aquatic  Fate of  Copper

Wot an important process.

'In most unpolluted waters, the
majority of copper will exist
as the carbonate complex.  In
polluted waters, coraplexation
kith organic material will be
most important.

Not an important process.
Various sorption processes
reduce the mobility of copper
and result in the enrichment
of suspended and bed sediments
relative to the water column.
In unpolluted waters, sorption
onto clay minerals, and hydrous
iron and manganese oxides are
controlling factors*  In polluted
waters, sorption onto organic
materials is the controlling

Biota strongly accumulate copper.
Copper is apparently not biomagni-

Some copper complexes say be
metabolized.  Organic Uganda
are important in sorption and
complexation processes.
Confidence of


    All of ,the  noted  environmental  processes are'important; however,  :heir
    relative  importance with  respect  to each other  is uncertain for
    determining  final  fate.

 11.5   Literature  Ci ted

 Baker-Blacker,  A.,  E.  Callender,  and  P.D.  Josephson.   1975.   Trace-element
   and  organic  carbon content of  surface  sediments  fron Grand  Traverse Bay,
   Lake Michigan.   Geol.  Soc. Araer.  Bull.   86:1358-1362.

 Brungs,  W.A.,  E.N.  Leonard,  and  J.M.  MCKla,   1973.   Acute and long-term ac-
   cumulation of copper  by the brown bullhead,  Ictaiuris nebulosus.   J.
   Fish.  Res.  Bd.  Can.   30(4):583-S86.

 Carpenter,  R.K. ,  T.A.  Pope,  and  R.L,  Smith.   1975. ,  Fe-Mn oxide coatings in
   stream sediment geochemical surveys.   J.  Geochera,  Explor.  4:349-363.

 Chapman, W.H.,  H.L.  Fisher,  and  M.W,  Pratt.   1968.   Concentration factors
   of chemical  elements  in edible  aquatic  organisms.   Lawrence Radiation
   Laboratory,  Llvennore,  Calif.   UCRL-50564.   46p.

 Collins, B.I.   1973.  The concentration  control  of  soluble copper in a mine
   tailings  stream.   Oochim. Cosmochim.  Acta,   37;69-75.

 Cotton,  F.A. and  G.  Milkerson.  1972.   Advanced  inorganic chemistry.
   pp.903-922.   Inter-Science Publishers,  New  York.   1145p.

 Cross, F.A., L.H, Hardy,  N.Y. Jones,  and  R.T.  Barber.   1973.   Relation be-
   w«en total  body weight and concentrations of manganese, iron, copper,
   zinc,  and mercury in  white muscle of bluefish (PotBatonuis  saj._ea_trix)  and
   a bathyl-demersal  fish Antiaora rostrata.   J.  Fish.  Res. Bd. Can.

.Hera,  J.D. and  M.W.  Skougstand.  I960.   Co-precipitation effects la
   solutions containing  ferrous,  ferric,  and cupric  ions,  p.104.  'U.S.
   Geological,  Survey Water Supply Paper  1459-E, Washington, D.C.  LlOp'.

 Hea,  J.D.  1975.   Discussion of  "Role of  hydrous metal oxides in the trans-
   port of heavy metals  in the environment" by C.F.  Lee. pp.149-153.  in
   Heavy aetals in the aquatic environment.  P.A.  Krenkel (ed.).  Pergamon
   Press, Qxtord,  England.  352p.

 Huaq;, C.P., H.A. Elliott, and R.M. Ashasad.   1977.   Interfacial reactions
   and  the fate of heavy metals in soil-water  systems.   J. Water Pollut.
   Control Fed.   49(5):745-756.

 Jackson, K.S.  and G.B,  Skippen.   1973,   Geoch^ical  dispersion of heavy
   metals via  organic complexing:   a laboratory study of copper, lead, zinc,
   and  nickel  behavior at a simulated  sediaent-water boundary.  J. Geochem,
   Explor. 10:117-138.

Jenn«, E.A.,  1968.  Controls on Mn, Fe, Co, Nl, Cu at«l Zn concentrations in
  soils and waters:  the significant role of hydrous Mn and Fe oxides.
  pp.337-387. In Trace inorganics in water.  R. P. Gould (ed.).  Advances In
  Chea. SeriesT3.  American Chemical Society, Washington, D. C.  396p,

Khobot'ev, V.G., V.I. Sapkov, E.G. Rlkhadze, N.V, Turanina, and N.A.
  Shidlovskaga.  1976.  Copper uptake by algae from copper-containing com"
  pounds and the effect of this process on their salt metabolism.
  Gldiobiol. Zh. 12(l):40-46 (Russian) (Abstract only;.  CA 1976.

KlBball, K.D.  1973.  Seasonal fluctuations of ionic copper in Knight's
  Pond, Massachusetts. Llnirioi. Oceanogr.  18(1): 169-172.

Krauskopf, K.B.  1972.  Geocheaistry of ulcronutrients.  pp. 7-40.   in
  Micronutrients In agriculture.  Mortvedt, J.J., P.M. Giordano, and W.L.
  Lindsay («ds.). Soil Science Society of America, Inc. Madison,
  Wisconsin.  649p,

Lee, C.F.  1975.  Transport of heavy metals in the environment,  pp.
  137-147. in Heavy metals in the aquatic environment.  P.A. Krendel (ed.).
  Pergamon Press Oxford, England. 3S2p.

Long, 0.?. and E.E. Angino.  1977.  Chemical speciation of Cd, Cu,  Pb, and
  Za In mixed freshwater, seawater, and brine solutions.  Geochim.
  Cosaochia, Act*.  41:1183-1191.

Lopez, J.M. and G.F. Lee.  1977.  Environmental chemistry of copper in
  Torch Lake, Michigan.  Water Air Soil Pollut.  8:373-385.

Namminga, H. and J. Wilhm.  1977.  Heavy metals in water, sediments and
  ehironoaids.  J. Water Pollut. Control Fed.  49(7):1725-1731.

Nehring, R.B.  1976.  Aquatic Insects as biological monitors of heavy oetal
  pollution.  Bull. Environ. Contaa. Toxlcol. 15(2>:1*7-154.

Patrick, F.M. and M. Loutit.  1976.  Passage of aetals in effluents
  through bacteria to higher organisms.  Water Re*.  10(4):333-335.

Payne, K. and W.F. Pickering.  1975.  Influence of clay-solute Interactions
  on aqueous copper ion levels.  Water Air Soil  Pollut.  5:63-69.

Ramamoorthy, S. and D.J. Kushner.  1975.   Heavy metal binding  components of
  river water.  J. Fish. Res. Bd. Can.  32:1755-1766.

Ramamoorthy, S. and B.R. Ruse.  1378.  Heavy aetal exchange processes In
  sediment-water systems.  Environ. Geol. 2(3) :165-172.

Rashid, M.A.  1974.  Adsorption of metals on sedimentary and .peat humic
  acids.  Chea. Geol.  13:115-123.

Rayaont, J.E.G.  1972.  Pollution in Southhaapton water.  Proc. Roy. Soc.
  Ser. B. 180(106):4Sl-468.  (Abstract only).  CA 1973.  78;144603f;

Sanchez, I. and G.P, Lee.  1973.  Sorption of copper on Lake Monorna
  sediments - effect of OTA on copper release from sediaencs.  Water Res.
  7:587-593.             .        .

Sceele, K.F. and G.H. Wagner.  1975.  Trace netal relationships in bottom
  sediments of a fresh water stream - The Buffalo River, Arkansas.  J.
  Sed. Pet. 45(1}:310-319.

Stiff, M.J.  197la.  Copper/bicarbonate equilibria in solutions of bicarbo-
  nate ion at concentrations similar to those found in natural waters.
  Water Res. 5:171-176.

Stiff, M.J.  1971b.  The chemical states of copper in polluted fresh water
  and a scheme of analysis to differentiate them,  Water Res. 5:585-599.

Stokes, P.M., T.C. Hutchinson, and K. Kraute.  1973.  Heavy Metal tolerance1
  in algae isolated from polluted lakes near the- Sudbury, Ontario smelters.
  Water Poilut. Res. Can.  8:178-201.  (Abstract only).  CA 1975.

Sylva, R.N.  1976.  The environmental chemistry of copper(II) in aquatic
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  CRC Press, Cleveland, Ohio.  2398p.

Wfndom, H., R. Stickney, R. Smith, D, White, and F, Taylor.  1973.
  Arsenic, cadmium, copper, mercury, and zinc in some species of North
  Atlantic finfish.  J. Fish. Res. Bd. Can. 30(2):275-279.

                               12.   CYANIDES
12.1  Statementof Probable	Fate

    Cyanides are a divers* group of compounds whose fate in the aquatic en-
vironment varies widely.  Hydrogen cyanide, the most common and nost toxic
of the cyanides, aay be destroyed by biodegradation or can be removed from
solution by volatilization or adsorption.   Cyanide ion (QT) can react
with a variety of metals to form insoluble metal cyanides.  If cyanide ion
is present in excess,, complex metallocyanides may be formed.  The latter
compounds are quite soluble and can be transported in solution.  The fate
of low molecular weight organic cyanides (nitriles) is expected to parallel
the fate of hydrogen cyanide.

12.2  Idea t ifieation - Envi ronmental Chemistr y of C yan id e s

    Cyanides are defined as organic or1 inorganic compounds which contain
the -CM group.  Hydrogen cyanide (HCN) is lighter than air and diffuses
rapidly,  Fre* HCN is very reactive and occurs only rarely in nature; it is
usually prepared commercially from ammonia and methane at elevated tempera-
tures with a platinum catalyst.  Hydrogen cyanide is soluble in all propor-
tions in water.  It is quite volatile, having a vapor pressure of 100 torr
*t -178°C; 360 torr at 7*C; 658.7 torr at 21.9'C; and 760 torr at 26.7°C
(boiling point) (Towill «£ al. 1978).  Cyanide ion forms complexes with a
variety of metals, especially those of tne transition series.  Ferrlcyan-
id«s and ferrocyanides have a variety of industrial uses but do not release
cyanide unless exposed to ultraviolet light.  Thus, sunlight can lead to
the mobilization of cyanide in witers containing iron cyanides. Cyanogen
[(CN>2l is a flammable gas of high toxicity which has a vapor pressure of
about 5 a tin. at 20'C (Towill _et al. 1978).  It reacts slowly with water to
produce HCN, cyanic acid, and other compounds. .Cyanates contain the -OCN
radical.  Inorganic cyanates, which are formed industrially by the oxida-
tion Of cyanide salts, hydrolyze in water to form ammonia and bicarbonate
ion.  Alkyl cyanates trlmerize readily (when sufficiently concentrated) to
form cyanurates.  Alkyl isocyanates contain the -NCO radical and are formed
from cyanates| they, too, are readily hydrolyzed. Thiocyanates (-SCN radi-
cal) are formed from cyanides and sulfur-containing materials and are more
•table  than cyanates.  Solutions of thiocyanates form free hydrogen cyanide
in acidic media.  Nitriles are organic compounds that have a cyanide group
at a aubstituent.  The nitriles are generally much less toxic than the free
hydrogen cyanide or the metal cyanides.  Cyanohydrins [R2C(OH)CN] are
toxic compounds which can decompose with the release of HCN or CN~ under
environmental conditions.

    In general, the cyanides occur tn water as (1) free hydrocyanic acid
(HCN), (2) simple'cyanides (alkali and alkaline earth cyanides), (3) easily
decomposable complex cyanides such as 2n(CW)2> and (4) relatively stable
complex cyanides such as [FeCCN)^]"?, [FeCCJOgS^, and Co(CN)4.
The complex nickel and copper cyanides assume an Intermediate position be-
tween the easily decomposable and relatively stable compounds.

    The CAS number for cyanide ion is 57-12-5; its TSL number is A568-9315.

    The CAS number for HCN is 74-90-8j its TSL number is A943-9671.

12.3  Summary _of__Fa_ce Data

    12.3.1  Photolysis

         The photodecopposition of ferrocyanide and ferricyanide solutions
and the resultant cyanide residuals in test solutions were observed by
Baadish and Bass (1922) and Schwarz and Tede (1927).  This effect was
corroborated for river waters at concentration levels of 2 tag/1 with both
potassium ferrocyanide and ferricyanide by Burdick and Lipschuetz (1948),
A 5-hour exposure of 100 ag/1 potassium ferrocyanide to sunlight produced a
cyanide ion level of 6 oig/1.  No rate constants were- calculated which would
help assess the importance of this photodecomposition as an environmental

         Hydrogen cyanide Is very resistant to photolysis by wavelengths of
light reaching the earth's surface (Frank and Bard 1977).  In the presence
of'titanium dioxide (TiOi) powder, however, photocatalytie oxidation of
cyanide ion proceeds at significant rates in both high Intensity artificial
sunlight and unfocused sunlight.-  Frank and Bard (1977) demonstrated that,
with T102 powder present, more than 99% of a 1 nH (26 ag/1) solution of
cyanide ion was oxidized by exposure to sunlight for two days.  In the ab-
sence of T102 powder, little or no oxidation occurred.

         The significance of photolysis on the aquatic fat« of the cyanides
has not been fully investigated, although it Is possible that the photoly-
sis of the metallocyanides could result in the release of cyanide ion
(Sroderius, 1977). , This process could be important In aquatic environments
downstream from metallocyanid* discharges.

    12.3,2  Chemical sPeclation

         Hydrogen cyanide can be oxidlzerf to isocyanic acid (HNCO) in the
presence of strong oxidizing agents.  This material can then be hydrolyze*i
via the following reaction (Towill _et _a_l. 1978):

                     H20          0         H20
     H - N • C - 0  	te-  H2N - C - OH

         Hydrolysis can also result in Che destruction of HCN and che
tiltriles, but it probably occurs so slowly as to be non-competit.ive with
other processes.  HCN is hydrolyzed via the following reactions (Khorkin jat
al. 1967):

The tautonerization of HCN to HCN:  is the rate determining step, with the
subsequent steps occurring rapidly at rooo temperature (Kreible and HcNally
1929;  Kreible and Peiker 1933; Khorkin ee. al. 1967).  Hydrolysis of HCN in
strongly acidic solutions (pH CH3CH2CN
> CH3CM, CH20IKH2CN > (COOH)CH2CN.  HCN was more reactive than all
of the above nitriles.  This inpliet that hydrolysis of nitriles in the
aquatic environment is slow, in aost cases, and is probably not competitive
with othtr processes.

    12.3,3  Volatilization   • •

         Hydrogen cyanide (HCN) is highly volatile, exerting a vapor pres-
sure of 360 torr at 7*C.  In most natural ^waters, almost all of the free
cyanide in solution is present as HCN, with the remainder present as CN"*.
The relationship of pH to percent HCN is shown below (Towill jtt al . 1978).
              -2S                  Percentage of Total Free jSignide ^as ........ HCH
               8                                     93.3
               9                                     58
              10                                     13

         Unpublished data developed by Dr. S.J.•Broderlus of the University
of Minnesota-St. Paul (Broderius 1977) indicate that volatilization of HCN
is a relatively swift process.  Ten 8-liter natural water samples were
Spiked with HCN and left open to the laboratory atmosphere (with no wind)
in battery jars.  Initial cyanide- concentration1and concentration after 6
hours were measured.  The relationship between rate of HO loss and initial
concentration of free cyanide (as HCN) was observed to be first order.
Half-lives of 22 to 111 hours were calculated for the conditions extant
to St. Paul, Minnesota.  When the experiment was performed outdoors, so
that the solutions were exposed to moderate winds, the rate of hCN loss,
increased by a factor of 2 to 2,5.  The concentrations of cyanide in this
experiment ranged from 25 to 200 yg/1.

         In a duplicate experiment, samples from the ten natural waters
were again spiked with HCN but not exposed to the atmosphere.  Cyanide loss
was much smaller indicating that volatilization was the predominant pro-
cess.  It should be noted, however, that such concentrations of cyanide
would probably retard biodegradation, or at least cause a lag in biological
action by organisms capable of metabolizing cyanide.  Nonetheless, these
unpublished results indicate that volatilization is important as a fate o.f
free cyanide (uncowplexed by metals) in the environment.  The rate of vo-
latilization is, of course, affected by a number of parameters including
temperature, pH, mixing characteristics of the water, wind speed and ice

         A more rigorous experiment, but one which applies less directly to
natural aquatic conditions, was performed by Raef e_t al. (1977a).  The fate
of cyanide in aerobic microbial systems (e.g., secondary sewage treatment)
was studied with respect to adsorption, biodegradation, reaction with glu-
cose, and stripping (volatilization effected by air forced through the sys-
tem).  In the stripping experiments, an air-flow of 2 cc min~^ was
passed through 6 liters of 10 mg/1 cyanide solution (pH 7.0,  30*C). After
50 minutes, the amount of cyanide in the reaction vessel had  declined from
60 mg to about 55 ag.  The amount of cyanide continued to decline until,
after 375 minutes, only about 35 ag remained (Figure 12-1).  Addition of
biological solids had little effect on the stripping rate, although there
was an initial decrease when solids war* introduced.  This reduction in
rate was probably due to adsorption onto the solids.  In comparison to the
other processes investigated, stripping was more effective in removing '
cyanide than either adsorption or biodegradation.  It is difficult, how-
*v«r, to generalize these results to natural aquatic systems.

    12.3.4  Sorption

         Cyanides are sorbed by a variety of materials, including clays
(Cruz £t aj.. 1974), biological solids (Raef et, ail_. 197?a), activated carbon
(Dardan and Popa 1939), and sediments (Kordatov and Vastllev  1971).  In


                                            '0* CW-
             '0    50   100   150   ?00   250  iOO   iSO
                            TIME, MINUTES
Figure 12-1  Sttlpptng of cyanide  froa  a  reactor vessel containing
             6 liters of 10 mg/1 cyanide  solution.   From Rael et al

 comparison  to  many  refractory organic  pollutants, hydrogen  cyanide  is  not
 strongly partitioned  into  the sediments  or  suspended  adsorbents,  primarily
 due  to  Its  high solubility in water.

          Cyanides are fairly mobile  in the  soil  environment (Aleaii  and
 Fuller  1976>,  indicating that adsorption is probably  not  a  significant
 control on  nobility in most aquatic  environments where  sorbents are  ouch
 less concentrated.  Alesii and  Puller  (1976)  reportad that  cyanide mobil-
 ity  is  least where  soils exhibit  lew pH,  high concentrations  of free iron
 oxides, and positively charged  particles (e.g.,  kaolin, chlorite, gibb-
 aite).   Mobility is greatest at high pH,  high concentrations  of free
 CaCOj (high negative  ,cbarge) and  low clay concent.

        -,  Cruz  jet _al.  (1974) studied  the  adsorption  of HCN by montmorillon-
 itic clays. The data showed that adsorption is  fairly  weak and is de-
 creased by  the presence of water.  Thus,  in the  aquatic environment,
 adsorption  to  taontmorillonitic  clays is  probably not  an important fate

          Biological solids sorb cyanides, but, as with  the  other  sorbents,
 the  amount  thus bound is probably insignificant  in  comparison to  the
 amounts volatilized or biodegraded.  Raef ££ ajU (1977a)  demonstrated  that,
 with an initial cyanide concentration  of 20 mg/1 and  a  biological solids
 concentration  of 6000 «g/l (dry weight),  a  non-flocculating bacterial  cul-
 ture did not remove any cyanide from solution.   In  a  siailar  experiment,
.addition of 7260 ag/1 (dry weight) of  a  flocculant  culture  of heterogen-
 eous bacteria  reduced cyanide concentration from 16 mg/1  to 14.1 mg/1  after
 one  hour.   The absence of  strong  sorptive effects led the investigators to
 conclude that  adsorption probably plays  an  insignificant  role in  the over-
 all  removal of cyanide observed in biological treatment plants (Saef et al.

          Although it  appears that sorption  is not important in determining
 the  fate of HCN, aor« data are  required  before the  importance of  sorption
 on the  fate of the  metal cyanides and  nitriles can  b« adequately assessed.

     12.3.5  Bioaccumulation

          tr. biological systems, hydrogen cyanide interferes with  the en-
 zymes associated with cellular  oxidation.  It is either quickly metabolized
 or the  organism diet* Thus, there is little potential for bloaccumuiatton
 of hydrogen cyanide.

          Many  plants  synthesize and  accumulate cyanogenlc glycosldes
 (Gewicz et  ajl. 1976), When the tissues  of  these plants are crushed, hy-
 drolytic enzymes arc  released,  which in  turn cause  the  destruction of  the
 cyanogenic  glycosides to evolve

         Broderius (1973) reported the bloaccumulation of metal cyanide
complexes in fish.  Copper cyanide concentrations ranged from undetectabl'e
to 303.9 pg per gran for various bluegill tissues, with the liver/gall
bladder fraction exhibiting greatest bloaccumulation.  Likewise, silver
cyanide was accumulated in coneentations up to 168.4 yg per gram.  It is
difficult to assess the environmental importance of metal cyanide bioaccum-
ulation, other than to note that the metal cyanides are generally less
toxic than HCN.  Chronic toxic effects of the metal cyanides are un-
doubtedly enhanced by such biorccuotulation.

    12.3.6  Blotrangformation

         Hydrogen cyanide, mtetaliocyanide complexes, and tiltriles are all
subject to biodegradation.  Although few data are available on biodegrada-
tion of cyanides In surface water, the literature is replete with refer-
ences to degradation in anaerobic and aerobic sewage treatment.  It has
been demonstrated that activated sludge treatment can result in virtually
complete reooval of cyanide (!u4z«ck and Schaffer 1962; Koatenbalder and
Fleckstelner 1969), but Saef jet _al. (I977b) showed that most of the loss of
cyanide in such systems is due to stripping (volatilization).  Slodegrada-
tion was of secondary importance, and adsorption accounted for a minor
amount of cyanide removal.  The pronounced difference in biological densi-
ties and physico-chemical conditions (nutrient concentrations, dissolved
gases, etc.) between natural surface water systems and sewage treatment
processes make it very difficult to extrapolate the above results to the
aquaticsenvironment.  It is evident, however, that biodegradation of
cyanides does occur in natural waters, and the importance of this process
varies according to such factors as cyanide concentrations, pH, tempera-
ture, concentration of microbes, availability of nutrients, and whether the
microbes art acclimated to cyanide.

         The ability to metabolize hydrogen cyanide or its salts appears to
be nearly universal.  A bacterium has been Isolated that is capable of
using cyanide as its sole source of carbon and nitrogen (Ware and Painter
1955), but most organisms can only tolerate very low concentations.  Typi-
cal metabolic pathways for degradation of cyanide Include production of
thiocyanate, reaction with hydroxocobalamln to fora cyanoeobalatnin, combin-
ation with amino acids, and oxidation to carbon dioxide and formate ion
(Towlll j« aj.. 1978).

         Of the numerous experiments that have been performed regarding the
biochemistry of cyanide metabolism, one of the most salient was that of
Hardy and Knight (1967).  In this experiment, nitrogen-fixing enzymes were
extracted from the bacteria Aaotobacter vtnelandii and C1 ostridLium_j>as-
teurnlanum.  Under reducing conditions, these enzymes reduced HCN to 0*4,
MH3, and possibly small amounts of CH3NH2.  Adenosine triphosphate
(ATP) was required fo'r the reaction, and it was necessary for the enzymes

to be extracted from cells grown In f^, rather than ammonia or urea,
Apparently the same enzymes that catalyze the reduction of ^2 to NH3
and NC»2 to ^2 are responsible for the reduction of HCN (and aside).'
Hydrogen cyanide is synthesized by a variety of organisms including bac-
teria (Lorck 1948; Wissing 1974; Gastric 1975), fungi (Hutchinson 1973);
and a millipede (Duffey and Blum 1977).

         Nitriles can also be biodegraded.  Some of the low molecular
weight nitriles can be used by microbes 33 their sole source of carbon and
nitrogen.  The bacterium Nnoc_a_rdia r^odochro^us can metabolize acetonitnie,
propionltrile, hydroacrylonitrile, butyronitriie, and succinonitrlie (Oi-
Geronino and Antoine 1976).  Biodegradation of nitriles is sometimes a
Cwo-step enzymatieally mediated hydrolysis with an amide as the interme-
diate product, and ammonia and the corresponding carboxylie acid as end
products (Fukuda e_t a 1. 1971; DiGeronimo and Antoine 1976; Mimura e_t a 1.

         *n ^fiudoaonas, the biodegradation process may involve 'Several
additional intermediates.  Firain and Gray (1976) showed that acetonitrile
is metabolized via the following pathway in Pseudomonas:  acetonitrlle—•
acetamide —— acetate—- tricarboxylic acid intermediates.  This pathway
apparently exists in higher plants as well.  The authors speculated that
higher aliphatic nitriles could also 'be metabolized by enzymatically
mediated hydrolysis combined withe*- and 3 - oxidation.

         It is evident from the literature that the cyanides can be bio-
degraded.  The rapidity of this process varies widely in the laboratory,
however, and it is difficult to generalize these results to environmental
conditions.  In natural surface waters, bicdegradation is an important
face, but tao're work needs to be done before the quantitative aspects of
this process can be determined.

12,4  Data Summary

    The data on the fate of cyanides in the aquatic envlronaent are incon-
clusive, and can only be interpreted with caution.  It appears, however,
that volatilization and biodegradation are the dominant processes affecting
HCN and the nitriles.  Adsorption can also result in removal of those
cyanides frow solution.  The simple mecal cyanides are insoluble and prob-
ably accumulate in the bed sediments.  Complex metallocyanides are trans-
ported in solution by the water column.  Changes in- the concentration ratio
of metals to cyanide can alter the behavior of the metal-cyanide compounds.
If the metals become more prevalent, formation of simple mecai cyanides is
favored; if cyanide, becomes acre prevalent, the coraplexed forms occur.
Some of the netallocyanide complexes can be photolyzed by sunlight, pos-
sibly resulting in diurnal fluctuation in free cyanide levels where 'these
complexes are present.  Table 12-1 summarizes the aquatic fate of cyanides.





                                 Table 12-1
                    Summary of Aquatic Fate of Cyanide

Can cause breakdown of some
metallocyanide complexes.

Chemical transformations occur
very slowly In aost aquatic

At pH<10 most of the free cyanide
will be HCN which is quite
volatile.  A most important
process in the aquatic environment,

Cyanides are sorbed by organic
materials and to a lesser extent
clay minerals.  Not an Important

Cyanides are not bioaccumulated.

Cyanides are biodegraded at low
concentrations by almost all
organisms,  A very important
process for the aquatic fate of
Confidence of


    All of the noted environmental processes are important; however, their
    relative importance with respect to each other is uncertain for
    determining final fata.

 12.5  Literature Cited

 Alesii,  B.A.  and W.H.  Fuller.   1976.   the mobility 'of three cyanide forms
   ia soils,   pp. 213-223.   in  Residual Management by Land Disposal.  W.H.
   Fuller (ed.).  Environmental  protection Agency,  Cincinnatti,  Ohio.  PB
   256768  268p,

 Bandish, 0.  and  L.W.  Bass.   1922.   Eisen ab licht-chetnisher katalyzator.
   I. Uber die .zurseczing von ferrocyaftkaliura in tageslicht.  Ber.  Deut,
   Chess,  Gesell.'  55:2698-2706.

 Broderius, 'S.J.   1973.  Determination oi molecular hydrocyanic acid in
   water  and  studies of the  chemistry  and toxiclty to fish of metal-cyanide
   coaplexes.   Ph.D. Dissertation,  Oregon State Univ., Corvallis,  Ore.
   287p,  Availably froa Univ. Microfilms International, No. 73-21299,  Ann
   Arbor, Michigan.

 iroderius, S.J,   1977.  Personal comroutiicatiott concerning the fate of
   cyanides in the aqyatic  environment.  EPA. Grant R805291, Dec. 8,
   1977.  Univ. of Minnesota, St. Paul.

 Burdick, G.E. and M.  Lipschuetz.  1948.  "oxlcity of ferro- and
   ferricyanide solutions to fish and  determinations of the cause  of
               Traas.  Amer.  Fish Sac.  78:192-202.
 Gastric,  P. A.  1975.   Hydrogen cyanide,  a secondary metabolite of
   Pseudoaonas  aerogmoai.   Can, Jour., Microbiol.  21:613-618.

(Cruz,  M. , A. Xai'ser,, P,G.  Rowxhat,  a«l  J.J.  Pripiat.   1974.   Absorption and
   transformation of  HCN on the surface  of copper and  calcium
   •ontiaorillonite.  Clays  Clay Mineral.   22:417-425.

 0»rdan,  0. and S. Popa.  1939.  Influence of alkalinity on the absorbing
   power  of activated carbon for warfare gases:   phosgene,  chlorine,
   chloropicrin, and  HCS.  Compt. Rend. Inse.  Sci. Roumanie  3:675-682
   (Abstract  only).  CA 1941. 36:6695.

 OiCeronifflO,  M.J. and A.D.  Antolne.   1976. Metabolism of acetonitrile and
   propionltrile by Moeardia rhodochrous LL100-21.  Appl. Environ.
   Microbiol. 31(6):900-906.          r"

 Dwffey,  S.S. and M.S. Blua.  1977.   Phenol and  guaiacol:  biosynthesis,
   detoxification, and function in a polydesinld  millipede, Ox id us  '
   grocillus.   Insect. Biochea.  7:57-65.

Firmin, J,L. and 0.0. Gray.  1976.  The biochemical pathway for the
  breakdown of methyl cyanide (acetonitrile) in bacteria, Biochem, J.

Frank, 3.N. and A.J. Sard.  1977.  Heterogenous photocatalyst oxidation of
  cyanide ion in aqueous solutions at titanium dioxide powder. J. Aner.
  Chem. Soc. 99(1):3Q3-3Q4.

Fukuda, Y, , M. Fukui , T. Harrada,, anf
  coke plant weak ammonia liquor.  J. Water Pollut. Control Fed.

Kordakov, V.A. and B.F. Vasillev.  1971,  Physicochemical processes In the
  elimination of heavy metal ions and cyanides from industrial wastewaters,-
  Tor. Naveh.  Is sled. Prackt. Ingt. Okogashch.  Rud. Tsvet. Metal. 7:49->8
  (Russian) (Abstract only).  CA 1973.  78:163698h.

Kreible, V.K. and J.G. McNally.  1929,  The hydrolysis of hydrogen cyanide
  by acids.  J. Aoer. Chem. Soc. 51:3368-3375.

Kreible, V.K. and A.L. Peiker,  1933.  The hydrolysis of hydrogen cyanide  •
  by acids II.  J. Amer. Chens. Soc.  55:2326-2331.

Kreible, V.K, And C.I. Noll.  1939.  The hydrolysis of nitriles with acids.
  J. Aner. Chem. Soc. 61:560-563.

Lorck, H. 1948.  Production of hydrocyanic acid by bacteria.  Physio!.
  Plant. 1:142-146.

 Ludzack,  F,J.  and  E.8.  Schaffer.   1962.   Activated  sludge  treatment  of
   cyanide,  cyanace,  and thiocyanate.   J.  Water  Polluc.  Control  Fed.
   34(4}:32Q-341.                              ,

 Mioura, A.,  T.  Kawano,  and  Y.  Yaoaga.  ,1969,   Application  of aicroorganisrs
   to petrochemical industry.  1. Assimilation  of nltrlle compounds by
   alcroorganisms.   J,  Ferment, technol.  47:631-638.

 Raef,  S.F.,  W.G.  Characklis,  M.A.  Kossleh,  and  C.H.  Ward.   1977a.  Fate  of,
   cyanide and  related  compounds  in aerobic  aicrobiol systems. I.  Chetsical
   reaction with substrate and  physical removal.  Water  Res.  11:477-483.

 Raef,  S.F. ,  W.G.  Characklis,  M.A,  Kosslch,  and  C.H.  Hard.   I9??b.  Fate  of
   cyanide and  related  compounds  in aerobic  microbial systems.  II.
   Hicrobial  degradation.  Water Res.  11:485-492.

 Schwart,  R.  and K. fede.  1927.   Uber die pnotochemie der
   kooplexverblndungen.   III.   Die  hexacyanokomplexe  des dreiwertigtn
   essins, kabalts, chroms,  and aangans.   B«r. Oeut.  Chea.  Gesell.

 Towill, L.E.,  J.S. Drury, B.L. WMtfield, E.B.  Lewis, E.L.  Calyan, and A.S.
   Hansons.   1978.   Reviews  of the  environmental eftects of  pollutants. V.
   Cyanide,  pp. 11-33.   Oak Ridge  National Laboratory, Oak  Ridge,  lenn.  PB
   289  920.  190p.

•Ware,  G.C.  and K.A.  Painter.   1955.   Bacterial  utilization of cyanide.
   Nature  175 (4464):900.

 Wiegand,  C.H.  and  M«  Tremeilinjj.   1972.   The  kinetics and  aechanism  of the
   decomposition of potassium  cyanide  in  aqueous alkaline nediua.
   Hydrolysis of the  simplest  nitrile,  HCH.  J.  Org.  Chem.  37(6};9W-916.

 Wissing,  f.   1974.  Cyanidt formation  from  oxidation of glyclne by a
   gteudononaa  species.   J.  iacterlol.  117:1289-1294.

                                 13.   LEAD
13.1  Statejaent : q iMhr oba b 1 ejF ate

    Sorption processes are effective in reducing the concentration of sol-
uble lead In natural waters and result in enrichment of bed sediments near
the source.  The  equilibrium solubility of lead with carbonate, sulfate,
and sulfide is low.   In severely contaminated areas, precipitation may be
Important in controlling the mobility of this metal, but under most circum-
stances , so ration predominates.  The tendency for lead to fora complexes
with naturally occurring organic materials (e.g., huailc and fulvic acids)
increases its adsorptive affinity for clays and other mineral  surfaces.
Benthic microbes  can methylate lead to fons tetramethyl lead which is
volatile and more toxic than inorganic lead.  Biomethyiation may,  in this
manner, also provide a mechanism for, retaobilization  of lead in the bed
sediments.  Sioaccumulatioa of weakly sorbed lead phases also  nay result in
reacbiliration.  Lead is generally not bioraagnified;  bioconcentration
factors tend to decrease as the trophic level increases.

13.2  Identification ~ Ge_ochemiscry of Lead

    Lead Is a naturally occurring element and is a major constituent of
•ore than 200 identified minerals.   Most of these are very rare, and only
three are found in sufficient abundance to form mineable deposits:  galena
(PbS) the staple  sulfide, angelesite (PbSQ4) the sulfate, and  cerrusite
(PbCOj) the carbonate.  By far the  most abundant is  galena which is the
primary constituent  of the sulfide  ore deposits from which most lead is
presently mined.

    Ores of lead, as well as those  of zinc, are often closely a -sedated in
deposits formed by replacement of limestone or dolomite.,  Lead ore is conr-
•only present together with ores of copper, zinc, silver, arsenic, and
antimony in complex vein deposits,  but lead ore also may occur in a variety
of igneous, metamorphic, and sedimentary rocks.

    Lead, atomic  number 12, atonic  weight 207.19, is a member  of the group
IV elements (Ueast 1977).  Lead exists in three oxidation states,  0, +2,
and *4.  Although neither metallic  lead nor the com^n lead minerals is
classified as soluble in water, they can both be solubilized by some acids;
In contrast, some of the lead coupounds produced industrially s§re consider-
ably water soluble.   Therefore, natural compounds of lead are  not  usually
•obile in normal  ground or surface water because the lead leached  from ores
becomes adsorbed  by ferric hydroxide or tends to combine with  carbonate or
sulfate ions to fora Insoluble compounds (Hem 19?6a).

    The average abundance of lead in the earth's crust Is approximately 15
ppm (Levering 1976) which Is equivalent to one-half ounce' of lead per ton
of rock.  Shales and unconsolidated sediments have a mean lead abundance
close to 6h« erustal average, showing rhe fairly even distribution of lead
in the environment.

    The CAS number for le'ad Is 7439-92-1;  the TSL number is 36A9-0641.

13.3  g umaa r y of _F a c_e Da t a  ,

    13.3.1  Photolysis

         Although no evidence was found concerning the photolysis of
organo-lead complexes in natural waters, photolysis of these compounds in
the atmosphere has a great bearing on the fora of lead which will enter the
aquatic . environment.  For example, Hir^-hler and Gilbert (196«) report that
the chief constituents of the inorganic lead compounds leaving the exhaust
system of autooobiles burning leaded fuels are two forms of PbClBr,
SH4CX-*?bClBr and 2XH4C1' PbClBr.  The species PbClBr (lead bromo-
chlorid-e) appears to be stable at ordinary temperatures and is isoaorphous
with PbCl2 and Pb8r2-  Both PbCl2 ant^ f^Br2 darken on exposure to
sunlight with the release of halogen.  The ultimate products of the pho-
colysis, of these lead compounds in the atmosphere .would be PbO and the
halogens.  Sines the majority of the lead emitted to the environment orig-
inates from the tailpipes of automobiles, these photolytic processes are
quite important.  Also of importance, as Fierrard' (1969) has pointed out,
is that the halogens produced from the photolysis of the lead halidt's -say
be involved in chain reaction mechanisms with such atmospheric pollutants
as CO, SO, and
    13.3.2  Chemical Speeiat ion  '

         An outstanding characteristic of lead Is its tendency to fora com-
plexes of low solubility with the major an ions of natural environmental
systems.  The hydroxide, carbonate, sulfide, and (more rarely) the *ulfaee
of lead may act as solubility controls.  Throughout tnost of the natural en-
viroraaent, th« divalent fora, Pb*2( is the stable ionic species of lead.
The oore oxidized solid Pb(>2, in which lead has a -Hi charge, is stable
only under highly oxidizing conditions, and probably has very little signi-
ficance in the aquatic environment (Cotton and Wilkerson 1972),  If sulfur
activity is very low, metallic lead can b* a stable phase in alkaline or
circumneutral reducing conditions.

         rtuang et al. (197?) calculated the equilibrium solubility of lead
as a. fijnction of pE for a system with total carbonate and total sulfur con-
centrations of 1Q~3 M at pH ?'.  Figure 13-1 shows the solubility of lead

                  *i   j* 2  M

*  __«
                             001 -

and Che controlling solid species for pE from +2'0 to -8.  At pH 7,
controls solubility over much of the p£ range encountered in natural

         Hem (197&b) calculated the fields of stability for solid species
of lead based on the available thermodynamic data; these results are sum-
marized in Figure 13-2 and Figure 13-3.  Although these figures are useful
in depicting equilibrium behavior, they are limited In that they do not
take into account environmental interactions with organic compounds and
other trace elements and, therefore, may be misleading with respect to fate
and transport in normal surface waters.  Hem (1976a) looked at the equili-
brium distribution between lead in solution and lead adsorbed on cation ex-
change sites in sediments.  H* calculated these distributions using *o,ua-
clons representing aelecticities of substrate for lead over H+, C«»r^,
and Na* and the stabilities of lead solute species.  Included in the
calculations were total concentrations of mejoi ions, cation exchange
capacity of the substrate, and pH,  The cation exchange behavior of lead In
natural systems could be predicted with this model if enough supporting
information w«r« available.  The available information for describing
natural stream sediments, however, is inadequate for accurate use of this
model.  In general, Hem's model suggests that In most natural environments,
•orption processes would more effectively scavenge dissolved lead than

         Long end Angino (197?) developed a theoretical model to study the
chemical speciation of lead in aquatic environments and the response of
lead to variations in ionic strength and complexation.  Association of lead
with the ligands OH", Cl"*, €03"^, $0$"^, and HCOj" was
considered at pH values from 3.5 to 11.0 at 25'C in differing seawacer-
freahwater mixtures.  The results are summarized in Figure 13-4.  In gen-
eral, the relative importance of the various ligand-lead complexes can be
predicted from a comparison of their stability constants; however, since
this model does not take into account metal-organic complexes, it is useful
only in unpolluted, relatively organic-free waters.

         Dissolved lead may b« hydrolyzed to fora Pb(OH>2.  Patterson _et_
al. (1977) studied tht formation of Pb(OH>2 versus PbCQj to determine
the feasibility of treating lead-containing waters with carbonates.  They
found that PbC03 controls lead solubility at pK < ll.S.  Even small _on-
centrations of inorganic carbonate due to dissolution of atmospheric CCH
are sufficient to reduce the solubility of lead to concentrations below
those predicted on the basis of hydrolysis alone.  It should be noted that
lead concentrations were reduced nearly to the computed solubility limits
within four hours;  thus, precipitation of .lead carbonate can occur

                                           10    12    »*
Figure 13-2  Fields of stabtiity for solid species and dominant
             solute species in system Pb + H-0 as functions of
             pH and redox potential.  Dissolved lead activity is
             10~8.32 aol/1 at 25 C and 1 fita. pressure.  From
             Hen (1976b).

Figure 13-3  Fields of stability for solids and  solubility  of  lead
             in system Pb '+ CO, + S + BLQ at 25°C and 1 atra. pressure,
             Ionic strength 0.005.  From Hem (1976b).

                                             100* frt'swater
                      15 4.5,5,5 4.5 f.5 S.5 « 5 10.5 11.5
Figure  13-4  Chemical speciatlon  of lead  in seawater-freshwater
              mixtures.   From Long and Angino (1977),

         At the low concentration in which lead is normally found in the
aquatic environment, almost ail of the lead in the dissolved phase say be
complexed fay the ligands of river water.  By using an ion-specific elec-
trode, Raraanoorthy and Kushner (1975) determined chat lead binding capacity
was predominantly due to organic compounds.  Inorganic complexes were not.
important, since evaporating the water samples, ashing the residue, and
reconstituting the ash in water resulted in complete loss of the binding
capacity.  (In waters with a high carbonate concentration, however, binding
by HC03~ or ODj"2 is important).

         Guy and Chakrabacti (1976), In their study of metal-organic inter-
actions in natural waters, found that humic acids in solution and' other
organic complexing agents can maintain lead ions in a bound fora at a pH as
low as 3.  O'Shea and Mancy (1978), in their study of - the effect of pH and
hardness on lead speciation, found that the effects of pH and hardness ne-
tals were insignificant in lead-inorganic interactions.  They wete  impor-
tant, however, in lead-humic acid interactions.  Increasing the pH in-
creased the concentration of exchangeable lead complexes while an increase
in hardness tends to decrease the extent of the hurnie acid-lead interac-
tion.- Metals responsible for hardness apparently inhibit the exchange-
able  interactions between metals and humic materials in ways that are not
fully understood.

         Jackson and Skippen (1978) investigated the behavior of lead and
organic materials at a simulated sediment-water boundary.  The interactions!
Involved sorption by clays, organic complexing, carbonate reactions, hy-
ck>lysis, and desorption of lead from clay and octal hydroxides.  They
found that organic acids decreased the solubility of lead in the presence
of clay, particularly at acidic pH values.  This organic coraplexing is pro-
bably due to colloidal coagulation.  The organic acids, moreover, proved
capable of remobilizlng lead froa the solid phase.  There is, however, a
general kinetic hindrance to this desorption, particularly at basic pH
values.          '                                  '                       •

         In suBBaation, the transport of lead in the aquatic environment is
influenced by the speciation of the ion.  Although lead will exist mainly
as the divalent cation in most unpolluted waters and-become sorted into
partlculace' phases, organic material in polluted waters will have a great
effect on the chemical form in which lead will be present.

    13.3.3  Vola ti 1 i _zat ion

         The relatively volatile tetramethyl lead ((CH3>4Pb) can be  '
produced by microorganisms in lake sedinents from inorganic (Pb(NOj)2
and PbCl2) and organic ((G^^PbOOCCHj) lead compounds (Wong e_t a 1.
1975).-  Analysis of air in flasks that contained anaerobic lake sediments

inoculated with these compounds showed that tetramethyl lead thus produced
can be volatilized.  Under these experimental conditions, addition of 10 sag
of lead as trimethyl lead acetate resulted in volatilization of weekly in-
crements of 125 ug, 642 ug, 550 yg, and 256 yg of tetraaiethyl lead
during the first four weeks (Wong _et aI. 1975), . Nevertheless, the impor-
tance of volatilization of cetranethyl lead is uncertain.  Although the
rate of destruction of tetramethyl lead in aerobic waters is unknown, this
compound is probably not stable in oxidizing environments,  When a layer of
aerobic watar lies between the reducing sediments and the Atmosphere, vola-
tilization taay not be important.

    13.3.4  Sorption

         Sorption processes appear to exert a dominant effect on the dis-
tribution of lead in the environment.  Several investigators have reported
that in aquatic and estoartne systems, lead is removed to the bed sediments
in close proximity to its source, apparently due to sorption onto the sedi-
ments (Belz jejE _al. 1975; Valielm jet _al, 1974).  Different sorption mechan-
isms have been invoked by different investigators, and the relative impor-
tance of these mechanisms varies widely with such parameters as geological
setting, pH, Eh, availability of ligands, dissolved and participate iron
concentration, salinity, composition of suspended and bed sediments, and
initial lead concentration.

         Pita and Hyne (1975) studied the depositional environment of lead
in reservoir sediments and found that almost all of the lead in the sedi-
ments was in the fraction with specific gravity between 2.0 and 2.9,  This
fraction contains the clays.   The authors suggested that formation of
organo-lead complexes may play an important role In adsorption, noting that
"the same type of organic matter (negatively charged or polar) which tends
to fora organo-metallie compounds would also tend to adhere to clay miner-
als and would occur in the 2.0 to 2.9 specific gravity portion."  The pau-
city of lead in sediments with specific gravity less than 2.0 indicated
that adsorption onto organic  material not active in complex formation was
insignificant; the lack of lead in the denser fraction (sp. gr. > 2.9) in-
dicated that precipitation was not important.  The relative dominance of
adsorption over precipitation is corroborated by calculations made by Hem
(I975a), which indicate that  precipitation is important only under rela-
tively alkaline conditions.                          ,

         The adsorption of lead to soils and oxides was studied by Huang ejt
*1, (1977).  The data indicate that adsorption is highly pH-dependent, but
above pH 7, essentially all of the lead is in the solid phase (Figure
13-5),  It should be noted that at low pH, lead la negatively sorbed (re-
pelled from the adsorbent surface).  The addition of organic coraplexing
agents increased the affinity for Adsorption (Figure 13-6).  Therefore, the

                         . . MIT ALCAIC
                             J   4   5   «   7   •
Figure 13-5  Adsorption of lead on various  solids.   The soil-water
             system consisted of 5 gm/1 solid,  10"% Pb, and 0.1M NaCl
             (thus adsorption of 200 uK/ga  -  100% removal).   From
             Huang «t «1.  (1977).

                          MtTAMAJC -«>

                         MUMIC ACIO (HA)
                         MA I1SO
Figure 13-6  Effect of  hunic acid on the adsorption  of lead by.metapealc
             •oil.  Soil «4|*d at 5 pa/1, initial  [Pb2+]  » 10"3M, ionic
             •trength - 10"% (NaO.).  From Huang et al.  1977,

tendency for lead to be adsorbed probably reflects che face that lead is
strongly complexed by organic materials In the aquatic environment
(Ramaraoorthy and Kushner 1975).  Huang et al. (1977) speculate that the
Increased adsorption is due Co Che ability of the oetal-ligand complexes to
share free electrons, thus facilitating sorptton to electrophilic solid

         Similar studies have been carried ou'c in seawater environments.
For example, Patterson _e_t al. (1976), In their study of sewage effluent en-
tering polluted coastal waters, found'that virtually all the lead-in Che
sewage was contained in the partlcuiate phase before it entered the ocean
but that about 11 percent was made freely available within a day by cation
exchange when the sewage was mixed with seawater.  Further exposure of the
sewage Co seawaeer, .however, did not facilitate the release of more lead.
Lu and Chen (1977), in their laboratory study of the migration of trace
metals from polluted sediment into seawater, found that thq release o>f lead
from the sediment increased as the redox conditions became more oxidizing.
Moreover, after long-term incubation under aerobic conditions, lead con-
centrations were far below equilibrium concentration.  This latter observa-
tion is indicative of the substantial sorption processes which lead under-
goes in aquatic sediments.

         Ramamoorthy and Rust (1978) found that the sorption of lead by
Qttawa River bed sediments can be fitted to the linear fora of Langtnuir's
aviations. They reported chat the partition coefficient 06 lead between
sediment and solution is not greatly changed by the presence of other heavy
metals, provided chat che latter has the same order of concentration.  If
the concentration of one cation exceeds the other by more Chan a factor of
10, however, significant desorption of the less concentrated ion takes
place on a mass action basis.

    i     There are significant differences reported not only in the mode of
binding to bed 'sediments, but also in the distribution of lead atao.ng phases
in the water column.  Some authors report that lead is transported predotn-
inantly in the particulate phase rather than the dissolved phase (Kubota _e_t
al. 1974; Schell and Mevissi 1977);  others report that the amount in the
dissolved phase is about equal to'thaC in the particulate phase (Angino _et
ai. 1974);  and still others find that acre lead is transported in the dis-
solved, phase than in the suspended material (Pica and Hyne 1975). , There is
general agreement, however, that residence in lakes and impoundments cause-s
a  reduction in dissolved lead levels even when lead is initially present in
concentrations below calculated solubility limits (Kubota e_t ,a_l. 1974; Pita
and Hyne 1975).  Thus, sorption processes appear to be effective in re-
ducing dissolved lead levels and result in enrichment of bed sediaents.  It
appears chat, under most conditions, adsorption to clay and other mineral
surfaces, eoprecipitation/sorption by hydrous iron oxides, and incorpora-

tion Into eationic lattice sites in crystalline sediments are the important
sorption processes.

         Several authors, notably Jenne (1968), Lee (1975), and Hohl and
Sturam (1976), have hypothesized that the sorption of heavy metals by hy-
drous iron and manganese oxides Is a major control on the mobility of, these
pollutants in the aquatic environment.   On the basis of a high correlation
between the lead, iron and manganese concentration in sediments, Angino e_c
al. (1974) suggested that sorption by iron and manganese oxides is the
dominant sorption process in several Kansas streams.  Caddie and Laicmen
(1973i) demonstrated that hydrous iron oxides have a high sorption capacity'
for lead, socbing aa much as 0.28 moles lead per mole iron at pH 6.  The
ability of hydrous iron oxides to sorb lead increases with increasing pH.
At pH 8.1, 91 percent of the added lead was sorted.  When the pH drops,
howevsr, lead may be desorbed.  Although the relative importance of
individual sorpcion processes yaries widely, it appears that, in most
circumstances, lead is effectively removed to the sediments by sorption .

    13.3.5  Bioac£uaulat ion

         Bioaccumulation of lead has been demonstrated for a variety of
organisms.  Table 13-1 lists bioconcentration • factors reported fay various

      i   Microcosm studies indicate that lead is not bioaagnif led.  Lu e t
al. (1975) studied the fate of lead in three ecosystems differing only in
their soil substrate.  The ecosysteas contained algae, snails, mosquito
larvae, mosquito fish, and microorganisms.  Lead was concentrated most by
the mosquito larvae and least by the fish.  Furthermore, body burdens and
aqueous lead concentiation appeared to be strongly correlated to .the per-
centage of organic matter and cation exchange capacicy'of the soils, indi-
cating that the availability of lead in the systems was controlled by
adsorption to the soilSt  Since pH was the same for all three soils, pre-
cipitation/dissolution of inorganically bound lead was probably not re-
sponsible for the differences in lead availability and uptake.

         Merlini and Pozzi (1977a) measured lead uptake in purapkinseed sun-
              8 * kl P, Su9 ) exposed to'2Q3pj, aj PH 6.0 and 7.5  Fish at the
             _      ,
lower pH accumulated three times as much lead as fish kept at pH 7.5.
Gill, liver, and fin accumulated the most lead and muscle the least.   The
authors attributed the increased lead uptake at low pH to the increasing
concentration of divalent lead with decreasing pH.  In another experiment,
Merlini and Potzi (1977b) found a direct correlation between lead accumu-
lation by purapkinseed sunfish and the concentration of ionic lead in  water
at various concentrations of total lead.  Results suggest that the condi-
tions existing In the majority of natural waters render most lead unavail-
able for accumulation by aquatic animals.

                                 Table  13-1

                      Bioconcentration  Factors for Lead

Freshwater plants

Freshwater invertebrates

freshwater fish

Marine plants

M*rlr.« Invertebrates

Marlae fiah







Chapman ej. a_l. W68

Chapman et al. 19bd

Chapman ejt _al. 196d

Chapman e_t al. 196A

Chapman e_c al. 'l9f>B

Chapman et .al. 1963
a.  Bioconcentration factors  are  the  ratio derived from the concentration
    of the element In  the aquatic  organism (in ppta wet weight) divided by
    the concentration  of the  element  in water (in ppot).

         Patrick and loutit (19?6> studied uptake of lead by'benthlc bac-
teria and Subsequent transfer to tub!field worms.  The concentration factor
for bacteria was approximately 360.   Concentration of lead by tubificlds
was 0.77 ciases the amount fed then tn the bacteirla. Indicating that the
tublficids can clear lead more easily than the. bacteria.  The fact that the
bacteria could concentrate lead indicates that lead in ,the sediaents can be
teaobilized by bioaccumulation.

         Based upon available Information, fish accumulate very little lead
in edible tissues;  however,'oysters and mussels are capable of accumu-
lating high levels of lead.  Decreasing pH increases the availability of
divalent lead, the principal  form accumulated by aquatic animals.

    13.3.6  Biotransformation

         As previously discussed, lead can be sec.hyl.ated by microorganisms
present In lake sediments.  The volatile compound resulting froo biomethyl-
ation, i.e.., tetraaethyl lead, probably leaves the sediments and Is either
oxidized in the water coluon  or enters the atmosphere.  In any event, bio-
nethylatlon represents a process which enables lead In the bed sediments to
be reintroduced to the aqueous or atmospneric environment.  In addition,
biogenic llgands can play a significant role in complexing lead, especially
in polluted waters, and vill  thereby have a significant impact on the aqua-
tic fate of lead.                 «                 ,

13.4  D*ta Summarv

    The djainant mechanism controlling,the fate of lead appears to be sorp-
tion.  Precipitation of PbS<>4, PbC03, and PbS and bloaccumulation may
also be important.  At low pH values, sorption and precipitation are not
nearly as effective in removing lead fron solution, so that lead is much
aore mobile in acidic waters  than at higher pH values.  In alkaline and
cireuaneutral water?, removal of lead by sorption and precipitation aay
occur relatively quickly.  Table 13-2 summarizes the fate data.

                                table 13-2

                      Summary of Aquatic Face of Lead

Spec lit ion"1
                                   Confidence of
                                       Data '

          Summary         <

Important In determining the form
of lead entering the aquatic er.-
virorutent.  Importance within
natural waters is undeterminable.
Determines which solid species         Medium
controls solubility in unpolluted
waters.  Over most of the normal pH
range, PbCQj and PbS04 control solu-
bility in aerobic conditions.  PbS and
Pb control solubility in anaerobic con-
ditions.  In polluted waters, organic
conplexation is most iaortanc.

Probably not important in most   ,      Medium
aquatic environments.

Adsorption to inorganic solids.        High
organic materials, and hydrous iron
and manganese oxides usually controls
Che mobility of lead.

Lead 'a bioaccunulaced by aquatic      High
organisms.  Bloconccntration factors
are wichin the range of 10^ - 103.
Biometh'/lation in sediments can re-
mobilize lead.
    All of the noted environmental processes are important; however, their
    relative importance with respect to each other Is uncertain for
    determining final fate.

^* *  Literature Cited      •                 ,

Anglno, E.E., L.M. Magnuson, and T.C. Maugh.  1974,  Mineralogy of
  suspended sediment and concentration of Fe, Mn, N'i.'Zn, Cc, and ?b  in
  water and Fe, Jto,  and Pb In suspended load of selected Kansas screams.
  Water Resources Res,-  10(to): 1187-1191.

Chapman, W.H., H.L.  Fisher, and M.W. Pratt.  1968.  Concentration factors
  of chemical elements In wiible aquatic organisms.  Laurence Radiation
  Laboratory, Livermore, Calif.  UCRL-50564.  <»6p.

Cotton, F.A.' and G.  Wilkerson,  1972.  Advanced Inorganic chemistry. -
  pp.309-338.  Interscience Publishers, New York.  1145p.

Caddie, R.R and H.A. Laitaen.  1973.  Study of the sorption  of lead by
  hydrous ferric oxide.  Environ, Letters. 5(4):223-235.

Guy, 8.0. and C.L, Chakrabarti.  1976»  Studies of metal-organic
  interactions in model systems pertaining to natural waters.  Can. J.
  Chen. 54:2600-2611.

H«li, G.R., R.J. Huggett, and J.M.  rfill.  1975.   Behavior of Mn, Fa,  Cu,
  Zn, Cd, and Pb discharged frota a  wastewater treatment  plant into an
  tituarine en/lroruaent.  Water Res.  9(7):631-63b.

Hem, J.D.  1976a.  Geochenical controls on lead concentrations in streaa
  water and sediments.  Geochim. Cosaochim. Acta.  4G:599~o09.

Hea, J.D.  1976b.   Inorganic chemistry of lead in water, pp.5-11.  in Lead
  in the environaient.  T.G. Lbvering (ed.). U.S.  Geological  Survey
  Professional Paper 957, Washington, DtC.  90p.

Hir*chler, A. and L.F. Gilbert.  1964.  Jiature of lead  in automobile
  «xh*iiSt gas.  Arch. Lnviron. Health,  8;297-313.

Hohl, H. and W. Stuaa.  1976.  Interaction of Pb^* with  hydrous
  Y-A1203.  J. Colloid Interface Sci.  55(2):281-288.

Huang, C.P., H.A. Elliott, and R.H. Ashmead.  1977.  Interfacial reactions
  and th* fate of heavy aetals in soil-water system.  J. Water Pollut.
  Control fed.  49(5):745-756,

Jackson, K.S. and G.B. Skippen.  1978.  Geochemical dispersion of heavy
  metals via organic complexing:  a laboratory study of  copper, lead, zinc
  and nickel behavior at a simulated sediment-water boundary.  J. Geochem.
  Explor. 10:117-138.

J«nne, E.A.  1968.  Controls on Mn, Fe, Co, Ni-, Cu, and Zn concentrations
  in soils and waters:  the significant role of hydrous Mn and  Fa  jxides.
  pp.337-387 _in trace Inorganics in water.  R.A. Baker (ed.).   Advances  In
 'Chemistry Series 73.  American Chemical Society, Washington,  DC. 396p.

Kubota, J., E.L. Mills, and R.T. Oglesby.  1974.  Pb, Cd, Zn, Cu, and Co
  in streams and waters of Cayuga Lake Basin.  Environ. Sci. technol*

Lee, G,F.  1975.  Hole of hydrous metal oxides in ehe transport of heavy
  metals in the environment,  pp.137-148.  jji Heavy metals in the aquatic
  environment.  P.A. Krsnkel (ed.)-  Pergamon Press, Oxford, England.

Long, D.T. and E.E. Angino.  1977.  Chemical speciation of Cd,  Cu, Pb, and
  Zn in mixed freshwater, seawater, and brine solutions.  Geochim.
  Cosmochim. Act*.  41:1183-1191.

Lovering, T.G. (ed.).  1976.  Lead in the environment.  U.S. Geological
  Survey Professional Paper 957.  Washington, D.C. 90p.

L.U, J.C.S. and K.Y. Chen.  1977.  Migration of trace metals in  interfaces
  of seawater and polluted surficial sediments.  Environ. Scl.  Technol.

La, P.Y., R.L. Metcalf, R. Furaan, R. Vogel, and J. Hassett.  1975.  Model
  ecosystem studies of lead and cadmium and of urban sewage sludge
  containing these elements.  J. Environ. Qual.  4(4);505-509.

Merlini, M. and A. Pozzl.  1977a.  Lead and freshwater fishes:  part 1 -lead
  accumulation and witer pH.  Environ. Pollut.  12:168-172.

Merlini, M. and A. Pozzl.  1977b.  Lead and freshwater fishes:  part 2
  -ionic lead accumulation.  Environ. Pollut.  13:119-126.

O'Shea, T.A. and K.H. Mancy.  1978.  the effect of pH and hardness aetal
  ions on the competitive interaction between trace metal ions and
  inorganic and organic complexing agents found in natural waters.  Water
  Res.  12:703-711.

Patrick, F.M. and M. Lout it.  1976.  Passage of metals la effluents,
  through bacteria to higher organisms.  Water Res.  10:333-335.

Patterson, C., 0. Settle, and 8. Glover.  1976.  Analysis of lead in
  polluted coastal seawacer.  Marine Chen.  4(4):305-319.

Patterson, J.W., H.i. Allen, and J.J. Scala.  1977.  Carbonate
  precipitation for heavy metal pollutants.  J. Water Pollut. Control Fed.

Pierrand, J.M.  1969.  Photochemical decomposition of lead solids from
  automobile exhaust.  Environ. Set. Technol.  3(1);48-51.

Pita, F.W. and M.J. Hyne.  1975.  The depositions! environment of zinc,
  lead, and cadmium in reservoir sediments. Water Res*  9:701-706.

Ramamoorthy, S. and D.JU Kushner.  1975.  Heavy metal binding components of
  river water.  J. Fish. Res. Bd. Can.  32:1755-1766.

fUaamoorthy, S. and l.R. East.  1978.  Heavy octal exchange processes in
  sediment-water systems.  Environ. Geol. 2(3):165-172.

Schell, W.R. and A. Nevissi.  1977.  Heavy netals from waste disposal in
  Central Puget Sound.  Environ. Sci. Techol.  ll(9):887-893. ,

Vallela, I., H.D. Banes, and J.M. Teal,  1974.  Response of salt marsh
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  of Pb, En and Cd by marsh sediments.  Environ. Pollut. 7(2) .-149-157.

(least, 1.C.  (ed.).  1977.  CRC handbook of chemistry and physics.  58th
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Wong, P.T.S., Y.K. Chan, and P.L. Luxon.  1975.  Methylation of lead In the
  environment.  Nature.  253:263-264.

                               14,  MERCURY

14,1  Statement o* Probable Fate

     Mercury's major removal mechanism from A natural water system is ad-
sorption onto the surfaces of particulate phases and subsequent settling to
the bed sediment.  The overwhelming majority of any dissolved mercury Is
removed in this Banner within a relatively short time, generally in the im-
mediate vicinity of the source.  Much smaller portions of the dissolved
mercury are ingested by the aquatic biota or transported by current move-
ment and dilution.  Secondary transformations of mercury in the sediments
can occur;  these include precipitation as HgS and methylation by bacteria.
Since the mercury itself is not destroyed, these inorganic and organic
forms of mercury may then release ionic or metallic mercury into the water
column as part of a recycling process.  Resuspension of sediments by turbu-
lence or the activity of benthic organisms can also release these compounds
of mercury directly into the water column.  The primary sink for mercury
released to the environment is thus the sediments.

14.2  ldentlfication_- Geochemistry of Mercury

     Mercury is able to exist in the natural environment in three oxidation
states:  as the native element itself, in the +1 (mercurous) state, and in
the +2 '(mercuric) state.  The nature of the species which will occur in a
given assemblage, or predominate in solution, depends upon the redox poten-
tial and pH of the environment.

     It is difficult to estimate the abundance of mercury in the earth's
cruse.  Fleischer (1970) reported that concentrations vary between 5 and
1000 ppb in common natural materials.  Considerably higher concentrations
have been measured in specific formations in mercury-rich regions of the
world.  Erickson (1960) estimated that about 10^ metric tons of rock are
weathered each year.  Using an average mercury content for rocks of 80 ppb,
this would mean that about 800 metric tons of mercury are released from
rock every year.  Since typical soils do not contain higher concentrations
of mercury than does its underlying rock, some of this weathered mercury
must reach the aquatic environment.                           '  '     ,
     Mercury has an atomic weight of 200.59, a melting point of -38.37"C,
and a boiling point range of 356-.'*58'>C.  Ac 20°C, its specific gravity is
13.546 and its vapor pressure is 0.0012 torr (Weast 1977).  The solubility
of metallic mercury in pure water has been determined by Sanemasa (1975) to
be 19.2 pg/1 and 81.3 yg/1 at 5SC and 30°C, respectively.

     Superimposed upon the inorganic movements of mercury in the aq'uatic
environment is the biological mercury cycle.   This cycle is summarized in
Figure 14-1.  The-»>2 interconversicns of mercury compounds maintain a dy-
namic system of reversible reactions which lead to a steady-stace concen-
tration of methyl mercury ia sediments and waters.  This process will be
discussed in greater detail later in this r'eport.

     The chemical abstract number of mercury is 7439-97-6;  its TSL number
is S 100-9715.

14,3  Summary of Fate Data

     14,3.1  Photolysis

          Photolysis seems to be of significance to the chemical speciation
of mercury in the atmosphere and perhaps in the aquatic environment.  A
photolytic breakdown of dimethyl mercury in the atmosphere to methyl mer-
cury has been suggested (Williston 1968;  Hola and Cox 1974;  Johnson and
Braman 1974) as well as photodecomposition of phenyl mercury compounds in
both the atmosphere and natural' waters (Zepp ^ejc _al. 19?3}i  Because of the
limited amount of information on this subject, however, it is not clear
what impact this process might have on the overall fate of mercury in the
aquatic environment.

     14.3.2  Chemical Speciation

          Under the usual conditions of temperature and pressure that occur
in the aquatic environment, mercury can be present in any one of three
different oxidation states.  The most reduced of these forms is the metal,
which is a liquid at ordinary temperatures and which has a tendency to  !
vaporize*  The other two forms are the mercurous ion, Hg* , and the
mercuric ion, Hg*^.  On che basis of commonly available thermodynamic
data (cf. Hen 1970;  Gavis and Ferguson 1972) it ran be seen that mercury
forms many solute species.  Some of these complex ions have considerable
aqueous solubility while others are quite insoluble.  Mercury forms many
stable organic complexes and is generally much more soluble in organic
liquids than in water.  Figure 14-2 shows the solid and liquid forms of
mercury that will be stable in the normal range of environmental condi-
tions. i Figure 14-3'shows the aqueous mercury species under the same con-
ditions as Figure 14-2.

          Within a moderately oxidizing environment above pH 5, the predom-
inant mercury species will be elemental mercury.  The solubility of this
material is nearly constant under all conditions in which the liquid metal
is stable.  Mildly reducing conditions, which are likely to occur In many
sediments, can cause the mercury to be precipitated as the sulfide,

Figure 14-1 - The biological  mercury cycle in the aquatic environice"..
             Modified  from Wood,  1974.

Figure 14-2  Fields of stability for solid (c) and liquid (1) mercury
             species at 25°C and atmospheric pressure.  System includes
             water that contains 36 ppm chlorine and 96 ppm sulfur as
             S04"2 (Hem 1970),

Figure 14-3  Fields of stability for aqueous mercury  species  at  25  C
             and 1 atmosphere pressure.  System  is  the  same as used for
             Figure 14-2.  Hem (1970).

cinnabar, which hag an extremely low aqueous solubility.  In aquatic
environments thac are high in chloride, the solubility of mercury in
oxygenated solutions may be gieaCly increased by the formation of charged
mercuric chloride complexes (Garvia and Ferguson•1972).

          Equally as important as the thennodynamic parameters which are
depicted in Figures 14-2 and 14-3 are the processes which produce the
organic complexes of mercury.  Mercury exhibits an affinity for sulfhydryl
groups (-SH) which are present in the many protein,s that contain the aolno ,
acid c'ysteine as a cooponent part.  Mercury also forma complexes with
organic amino groups, which are present in proteins, amino acids, ana their

          Two types of alkylated mercury compounds are formed in the en-
vironment.  In compounds with a single carbon-mercury botid, the compound  i
thus formed acts as a substituted salt and is reasonably water-soluble.   An
example is aethyl mercuric chloride (CiT^KgCl) which becomes Q^Hg*
ion and Cl"" ion In solution.  The other type Involves rovalent attachment
of two carbon atoms to the aercury.  Although they are considered insolu-
ble, dialkyl, covalent mercury compounds may appear in natural waters at
trace levels.  An example is dimethyl mercury (CH3HgCH3> which is
volatile and is undissociated in solution.  The chemistry of methyl mercury
species and equilibria in aqueous solution have been discussed in detail by
Burrows et al. (1974) and Rabenstein et_ ad. (1975).

     14.3.3  Volatilization

          Metallic mercury, with Its uniquely high vapor pressure relative
to other metals, can enter the atmosphera from the aquatic environment as
several different gaseous compounds.  This factor makes volatilization
important for the aquatic fate of mercury.  The rate of vaporization of
aercury and certain of its inorganic compounds decreases in the sequence
Hg > Hg2Cl2 > HgCl2 > HgS > HgO according to the data of Koksay and
Bradshaw (1969).

          Presumably, the oicrobial aethylatlon of aercury would enhance
the evaporative loss of mercury.  Although monomethyl mercury compounds  are
the principal product of biological oethylatlon rather than,the non-ioniz-
able dimethyl1 mercury (Jensen and Jernelov 1969),  a net increase in
volatility should result.  Because of -limited quantitative data available
on the subject of the volatilization of mercury compounds from natural
waters, it is not clear what impact volatilization will have on the overall
fate of mercury in the aquatic environment.

     14.3.4  Sorption

          Mercury shows a tenacious affinity for surfaces of  many types.
The problems of storing dilute aquatic mercury samples in glass vessels

have been well known for years.  In natural samples, a aajor portion of the
total mercury has been found associated with the partlculates (Hinkle and
Learned 1969),  Studies on che addition of mercury to a variety of natural
samples have led to the same conclusion.  Carr and Wtlknlsa (1972) found
that radioactive mercury, wnen added to storad samples, was rapidly
apportioned onto th«i participate phases with half-lives for adsorption of
less Chan one to fifty hours.  This experiment indicated chat the adsorbed
species are probably not methylated mercury compounds•  The work of Kudo et
al. (1977b) supports this contention by demonstrating that there is no
significant isotopic exchange between ^^UgCl2 and CHjHgCl or
CH3203HgCl and HgCl2.

          Loring (1975) has demonstrated that the concentrations of mercury
In the sediments of che Gulf of St. Lawrence are highest immediately adja-
cent to the source rocka with concentrations decreasing as one proceeds
seaward in Saguenay Fjord.  The mercury is enriched, however, in che sedi-
ments relative to the source rocks. ' This distribution indicates a direct
movement from the source through the aquatic environment to the sediments
since physical transfer of source-rock to the sediment would not produce
the observed enrichment.

          Ramamoorthy and Rust (1976) studied mercury sorptior, onto the bed
sediments of the Ottawa liver in a laboratory study.  By varying Hg2"4"
concentrations and pH at a constant temperature, they found that sorption
rates were highest in organic-rich sands, and. it appeared that sediment:
binding capacity was most closely related to organic content.  They found
that mercury sorptioo was little affected by pH.  Desorption rates were
low, «.g., leas than 1 percent Hg was leached from the sediment after 7u
hours agitation in distilled water.  Several other investigators have re-
ported simlliar results with the 'sediments of other areas, e.g., Thomas
(1972) with Lake Ontario, Fillay ejt al. ('1972) with Lake Erie, Kudo et al.
(1977a) on methyl nercury ir. the Ottawa River, and Creceluis jst al. '71975)
ia Puget Sound.  Moreover, the theoretical work of MacNaughton and Janes
(1974), on adsorption of aqueous mercury complexes, indicated that the
adsorption of inorganic mercury on the inorganic oxide components of sedi-
ments is likely to be very small.

          Reimers and Krenkel (1974), in their study of mercury adsorption
and desorption on sediments, reported that at a constant pH, the adsorption
of inorganic mercury is affected by aquatic chloride concentration, with
the percent loss in capacity depending upon the constituents of the sedi-
ment.  The sedimentary . materials studied exhibited a capacity to sorb
methyl mercury that followed the order;

         organics » illlte » aontmorillonlt* » sand

They found, as well, that Inorganic mercury is bound strongly enough  by
sediments to be transported by sedimentary mobilization,

          In summary, it is evident from environmental studies and theore-
tical considerations, that mercury adsorption onto the sediments is  prob-
ably, the most important process 'for determining the fate of mercury in the
aquatic environment.

     14.3.5  Bioaccumulation

          Due to the recent concern over the danger to human health from
the eating of mercury-containing fish, the bioac cumulation of mercury has
been well studied in the aquatic envxro'nment.  Mercury is acquired by
organisms through direct contact in air and/or water and through the food
chain (Phillips and Russo 1978).

          Bacteria common to most natural waters have been proven to be
capable of converting many mercury compounds to methyl mercury (Jensen and
Jernelov 1969; Bisogni and Lawrence 1975).  Therefore, virtually any mer-
cury compound entering the water may become a bioaccumulation hazard if the
environmental conditions are favorable for biomethylation.

          Ramamoorthy et al« (1977.) have measured the uptake of mercury
from water by both bacteria and sediment^.  Bacteria accumulated mercury
much more rapidly than sediment, taking up nearly 20 times as ouch mercury
as seditkent after 72 hours.  Loss of mercury from the system during the ex-
periment was attributed to the bacterial conversion of divalent mercury t^
the volatile metallic mercury.  This loss did not occur in sterilized con-
trol samples.'

          Kudo (1976) exposed) guppies to water over a ~'J3Hg-enriched
sediment bed and measured mercury uptake by the fish.  Mercury uptake by
the puppies was compensated within the system by increased mobilization of
mercury from sediment into water.  The half-life of mercury in the sediment
under these conditions was estimated to be 12-20 years.  O'er one-half of
che mercury present in the fish was organic, suggesting that conditions
were such that mercury was being methylcted;  Kramer and Neidhart (1975)
measured mercury uptake from water in guppies using inorganic mercury and
methyl mercury.  Methyl mercury was more readily accumulated and retained
than inorganic mercury and the rates of uptake increased with exposure
level.  The accumulative half-life for methyl mercury was ?0 days, a much
lower value than that reported by other'workers.  Inese observations sup-i
port the hypothesis that inorganic mercury is not che ».:ijor source of  -
mercury for bio accumulation, by fish in most natural enviroraer.ts, •

          Potter ec al. (1975) reported1 on the Jiercury Content of various
tissues from  fis.'i and invertebrates collected from Lake Powell, Arizona.

Muscle contained the highest mercury level of any tissue for most species
(e.g., largemouth bass and carp) while in- trout, the vital organs all ex-
ceeded the muscular mercury levels.  Apparently, modes of uptake, reten-
tion, and elimination vary among species.  Factors which were believed to
influence the observed levels* of mercury in plants and animals at different
trophic levels included age, surface area, metabolism, habitat, and activi-

          In summary, methyl mercury is the form of mercury present in most
fish tissue, and it is the most readily accumulated and retained fora of
mercury in aquatic biota.  Methyl mercury is readily accumulated by fish
both from their food and through the water.  Although conflicting evidence
exists regarding the relative importance of these two sources of mercury ta
fish, most reports suggest that both sources can be significant.  Upon en-
tering the biological system, methyl mercury is very difficult to elimi-
nate.  Most studies imply that the depurative half-life of methyl mercury
in aquatic organisms is between one and three years (Phillips and Russo
1978).  Bioconcentration factors for mercury are summarized in Table 14-1,

     14.3.6  Siotransformation

          Mercury, as an element, is not intrinsically altered by chemical
reaction, but does take part in biologically mediated reactions.which
drastically alter its mobility and toxieicy.  Iverson _jst aJL. (11)75) showed
that certain bacteria are capable of transforming mercuric ion and phenyl-
mercuric  acetate to volatile elemental mercury, and Spanglar et _§_!. (1973)
hav« described a process whereby methyl mercury i'9 demethylated.

          Bisogni and Lawrence (1975) described the influences of inorganic
mercury concentration and speciation, pH, microbial activity, and redo* po-
tential on mercury methylation rates.  In general, the more Inorganic mer-
cury present, the more methyl mercury will be produced.  At a circumneutral
pH, the primary product of mercury methylation is methyl mercury.  Methyla-
tion can occur under both aerobic and aaaerobic conditions, although more
Mercury oethylation occurs when more bacteria are present.  Iherefore,
highly organic sediments which favor bacterial growth have at higher methyl-
ation potential than inorganic sediments.  Wood (1974) pointed out that all
the mercury in  natural waters could participate in a system of microblally
catalyzed reactions and chemical equilibria to produce steady state concen-
trations of dimethyl inercury, methyl mercuric ion, metallic mercury, mer-
curic ion and mercurous ion.

          Upon entering an aqueous system, virtually any mercurial compound
may be microbially converted to methyl mercury.  Conditions reported to en-
hance the methylation process include large amounts of available mercury,
large numbers of bacteria, absence of strong compi exIng agenta such as

                      Table 14-1

         Bioconcentration Factors for Mercury

Marine planes                                 1,000

Ktrine Invertebrates                        100,000

Marine Fish                                   1,670

Fr«»hv*ter Plants                             1,000

freshwater Invertebrates                    100,000
   i                                            -  '
Freshwater Pish                               1,000
Concentration factors are based on the concentration of elements
in the aquatic organism (expressed in ppa of wet weight) divided
by the concentration of the element in water (expressed in ppa);
from Chepman tt al. 1968.

3uifi.de, circuraneutral pH, high temperature, and a moderately aerobic en-
vironment.  Detnethylation processes also occur but apparently only when
methyl mercury levels become excessive (Fageratrom and Jernelov 1972).

          The conversion, In aquatic environments, of inorganic mercury
compounds to methyl mercury implies that recycling of mercury from sediment
to water to air and back could be a rapid process.  Bacteria can act not
only a? mediators of methylation, but can also preferentially aceuirulate
larga acaounts of mercury.  Although the sediment is probably the most
Important sink for mercury, methylation by bacteria could reduce the ne.r-
cury content of overlying waters resulting in the mobilization of inorganic
nercury from the sediments.

14.4  DataSummary

     The environmental fate of mercury has been well reviewed in several
papers (Gavis and Ferguson, 1972; Kothny 1971; Krenkel 1973, 1974; National
Academy of Sciences 1978).
                              t                                      i
     Mercury is strongly sorbed to inorganic and organic particulates.  De-
position of mercury-laden sediments in reducing zones can result in pre-
cipitation of th* sulfide.  liomethylation of mercury in the sediments can
result In reaobllization.  Since dimethyl mercury has a low solubility in
water and is a gas at room temperature, volatilization nay occur.  Mercury
is strongly bioaccunulated.  Table 14-2 summarizes the aq'iatic fate of


                                Table 14-2

                    Summary of Aquatic Face of Mercury

Chemical Speciation8

Important In the breakdown
of airborne mercurials,
might be important in some
aquatic environments.

Controls volatility of metallic
a«rcury by conversion to coo- ,
plexed species.  In reducing
sediments HgS will precipitate
and may constitute a major
chemical sink.

Important to the movement of
mercury compounds in and out
of the aquatic environment.

Sorption processes result in
the strong partitioning of mercury
into suspended and bed sediments.
Sorption is strongest Into organic

Bioaccumulation has been proven
to occur via numerous mechanisms*
Most are connected to methylated
forms of mercury.

Mercury can be metabolized by
bacteria to methyl and dimethyl
fonts which are quite mobile in
the environment.
Confidence of

    All of the noted environmental processes are important; however, their
    relative importance with resptet to each other is uncertain for
    determining final fate*

14,5  Literature Cited .

liaogfii, J.J.  and'A.W, Lawrence.  1975,   Kinetics of mercury methylation in
  aerobic and  anaerobic aquatic  environments.   .lour.  Water Poll.  Control
  Fed, 47{1):135-152.

Burrows, W.D., E.S.  Reimers, and  E.8.  Shin. 1974.   Chemistry of Mercury in
  the environment, pp 251-264 _ln  Mercury;  environmental considerations,   , ,
  part II.  P.A. Krenkel (ed) CRC Critical Reviews In  Environmental
  Control 251-339p,

Carr, R.A. and P.E.  WilknisSi 1973.   Mercury:  short term storage of
  natural Haters,  Environ. Sci.  Techno1.  7:62-63.

Chapman, W.H., H.L.  Fisher, and M.W.  Pratt.  1968.  Concentration
  factors of chemical elements in edible organisms.  Lawrence Radiation
  Laboratory,  Livenaore, Calif.  UCRL-50564.  46p.

Creceluia,.E.A., M.H.  Bothner, and R.  Carpenter, 1975.   Geochemistries of
  arsenic, antimony, narcury, and related  elements in  sediments of Puget
  Sound,  Environ. Sci. fechnoi.  9(4):325-333.

Ericsson, E. 1960,  The yearly circulation of  chloride and sulfur iln na-
  ture.  Tellus. 12:63-109.

Fagerstroa,  I. and A.  Jernelov. 1972.   Some aspects of the quantitative
  ecology of mercury.   Water Res. 6:1193-1202.

Fleischer, M.  1970,   Summary cf the  literaturei in the  Inorganic geochemis-
  try of aercury.  pp.6-13 jln Mercury in the environment.  U.S. Geological
  Survey Professional Paper 713,  Washington, D.C.  67p.

Gavia, J. and  J.F. Ferguson. 1972,  The cycling of mercury through the en-
  vironment.  Water  Res. 6:989-1008.

Hem, J.D, 1970.  Chemical behavior of mercury in aqueous media, pp, 19-24
  J.O Mercury in the  environment.   U.S.  Geological Survey Professional Paper
  713, Washington, O.C. 67p,

Minkle, M.E. and R.E.  Learned. 1969.   Determination of mercury in natural
  watert by collection on silver  screens,   pp.251-255.   U.S. Geological
  Survey Professional Paper 650D, Washington,  D.C. 299p.

Holm, U.W. and M.F,  Cox.   1974,   Mercury in aquatic systews: methylatlon,
  oxidation -reduction, and bioaccunulatlon.  (EPA 660/3-74-021).  U.S.
  Environmental Protection Agency, Washington,  D.C. 38p.

Iverson, W.P., C. Huey, F.E. Brlnckman, K.L. Jewett, and W. Blair. 1975.
  Biological and nonbiological trans format ions of mercury in aquatic
  systems pp.193-195 in Heavy aetals In the aquatic environment, P.A.
  Krenkei (ed.) Pergamon Press, Oxford, England* 352p.

Jensen, S. and A. Jernelov. 1969,  Biological methylation In aquatic
  organisms. Nature, 223:753-754.

Johnson, D.L. and R.S. Branen. 1974.  Distribution of atmospheric mercury
  near ground.  Environ. Sci. Techno!, 8(12):1003-1009.

Roksay, M. and P.M.D. Bradshaw. 1969.  Secondary dispersion of mercury
  from cinnabar and stibnite deposits, West Turk-y. Colorado School Mines
  Quart, 64(l);33-356.

Kothney, E.L. 1973.  The three-phase equilibrium of mercury in nature.
  pp.48-80 la Trace Elements  in the Environment,  i.f. Gould (td.)
  Advances in Chemistry Series No. 123, American Chemical Society.
  Washington, D.C.  24Sp.

Kramer, H.J. and 8. Neidhart, 197S.  The behavior of mercury in the'system
  water-fish.  Bull. Environ. Contaa. Toxicol, 14(6):699-704«

Krenkel, P.A. (ed). 1973.  Mercury: environmental considerations, part I.
  CIC Critical Reviews in Environmental Control,  pp.303-373.

Rrenkal, P.A. (ed). 1974.  Mercury: environmental considerations, part II.
  CRC Critical Reviews in Environmental Control,  pp.251-339.

Kudo, A. 1976.  Mercury transfer from bed sediments to freshwater fish
  (gupples). J. Environ. Ojuai. 5{4)s427-430.

Iydo,,A,, H, Akagi, tJ.C. Mortiicer, and D.R. Miller. 1977«.  Equilibrium
  concentrations of nethylmercury in Ottawa River sediments. Natur«.
 , 270:419-420.

Kudo, A., A. Hirokatsu, D.C. Mortimer, ami D.R. Millar.  1977b.  Isotopic
  organic and inorganic mercury exchange in river water*  Environ. Sci
  Technol. 11:907-908.                          ,

Loring, D.H. 1975.  'Mercury in the sediments of the Gulf of St. Lawrence.
  Can. J. Earth Sci. 12:1219-1237.

MacNaughton, M.G. and R.O. Janes. 1974.  Adsorption of aqueous taercury(II)
  complexes at the oxide/water interface.  J* Colloid Interface Sci

National Academy of Sciences. 1978.  An assessment of mercury in the en-
  vironment.  Washington, D.C. 135p.

Phillips, G.R and R.C. Russo.  1978.  Metal bloaccumulation Is. fishes and
  aquatic invertebrates. - Lf.S, Environmental Protection Agency,
  Environmental Raserach Laboratory, Duluth, Minn,  (EPA 600/3-78-103).

Pillay, .K.K.S., C.C.! Thomas, Jr., J.A, Sondel, and C.M. Hyche. 1972.
  Mercury pollution of Lake Erie ecosphere.  Environ. Res. 5:172-181.

Potter, L. , D.-Kldd, and 0, Standiford. 1975. Mercury levels in Lake
  Powell, bioamplj, f ication of mercury in man-made desert reservoir.
  Environ. Sci. Te'chnol. 9(1); 41-46.

Rabenatein, D.C.» C.A. Evans, M.C. Tourangeau, and M.T, FairhursC. 1975.
  Mechylmercury speries and equilibria in aqueous solution.  Anal. Chem.
                                  !                       :                 •
Ramamcorthy, S. and B.R. Ruat. 1976.  Mercury sorptlon snd desorptlon
  characteristics of some Ottawa River sediments,  Can. J. Earth  Scl.

Ratmaaoorthy, S., S. Sprlngtborpe, and D.J. Kushner. 1977.  Coupe titloa for
  mercury between river sediment and ,bacteria.  BuJl. Environ. Contain, Tox-
  ic ol. 17 (5):505-511.

Reiaers, R.S. and P.A. Krenkcl. 1974.  Kinetics of aercury adsorption and
  desorptlon in sediments.  Jour. Water Poll. Control Fed, 46(2):3S2-365.

Saneaasa, I. 1975.  Th« solubility of elemental mercury vapor In  water.
  Bull. Chem. Soc. Japan. 48:1795-1798.

Spangler, M.J., J,L. Spigarelli, J.M. Rose, and H.M. Miller, 1973.
  Methylmercury; bacterial degradation in lake sediments. Science.

Thomas, R.L. 1972. The distribution of mercury in the sedlnents of LaVe On-
  tario. Can. J. Earth Sci. 9:636-651.

Weaac, R.C. (ed).  1977.  CRC Handbook of Cfcemistry and Physics.  58th
  Edition. CRC Press, Cleveland, Ohio. 2398i .

Williston, 3.H. 1968.  Mercury in the atmosphere. J. Geophysical  Res.

Wood, J.M. 1974.  Biological cycles for tuxic elements In the environment.
  Scie'nce. 183:1049-101,...

Zepp, R.G., N.L. Wolfe,« and J.A. Gordon. 1973.  Photodecomposition of
  phenylmercury compounds in sunlight. Chemosphere. 3:93-39.

                                15.  NICKSL

15.1  Statement of Probable_ Fate

    Nickel appears to be a relatively mobile heavy metal.  Although sorp-
tlon ami precipitation do not appear to be as effective as they are with
•any of the other heavy metals, sorption processes can scavenge nickel tram
solution.  Nickel has an affinity, for organic materials and hydrous iron
and manganese oxides*  The latter materials are probably the dominant con-
trol on the mobility of nickel in the aquatic environment.

    Moat of Che common aqueous ligands form,moderately soluble compounds
with nickel.  An exception is nickel sulfide, which can be formed in re-
ducing environments and is quite insoluble.  Under aerobic conditions, how-
ever, the hydroxide, carbonate, sulfate, and halide compounds are suffi-
ciently soluble Co allow toxic levels of nickel to persist in solution.
Although nickel is bioaccumulated, the concentration ratios reported for
•ost freshwater organisms indicate that partitioning into biota is not a
dominant fat* process.

15.2  Identification - Geochemistry ofNickel

    Nickel it * naturally occurring element and is found in the earth's
crust in average concentrations of 80 ppm (Weast 1977).  Nickel is usually
divalent in It* compounds which are predominantly ionic in character.
Slckel compounds with valences af 0, +1, +3? and +4 have .been reported,  but
these are extremely rare and  not important in a discussion of the chemistry
of nickel in the aquatic environment.  Nickel forms coopound3 with sulfate,
chloride, nitrate, carbonate, oxide, hydroxide, and with organic ligands
(Cotton and Wilkinson 1972).

    Nickel is »lderof»hllic and will alloy itself with metallic iron when-
ever such' a phase is present.  Nickel is only slightly miscible in iron  and
the two phases separate at low temperatures.  The earth's core is thought
tc be a nickel-iron alloy (barysphere) with a Fe/Ni ratio of about. 11:1.
The weathering of nickel-rich bedrock gives rise Co iron-, nickel-, and
silica-rich solutions,  tonic nickel is very stable in aqueous solutions
and is capable of migration over long distances.  The high affinity of
nickel for sulfur accounts for its occurrence in mag a a tic or aetaaorphlc
segregates of sulfide bodies. ' These sulfide segregates constitute the
large nickel ore body at Sudbury, Ontario, which provides well over a third
of the world's mining production of nickel (Corrich 1973).  Nickel is a
transition element, atomic number 28, atonic weight 58.71 (Weast 1977).

    Nickel's CAS number is 007440020; its TSL number is QR 59500.

15.3  Summary of Fate Data
         fto evidence wad found to suggaat that the photochemistry of nickel
compounds affects the aquatic fate of nickel.

    15.3.2  C heal c a 1 S pe c i a t ion

         Nickel is almost always found in the divalent oxidation state in
aquatic systems (Cotton and Wilkinson 1972).  Under reducing conditions and
in the presence of sulfur, the insoluble sulfide is formed.  Under aerobic
conditions and pH below 9, the compounds nickel forms with hydroxide, car-
bonate, sulface, and naturally occurring organic iigands are sufficiently
soluble to maintain aqueous Mi1** concentrations above 1Q~~^M (60ug/l).
Above pH 9, precipitation of the hydroxide or carbonate exhibits some con-
trol on nickel nobility.

         Hydrolysis of aqueous nickel to the hydroxide, Si(OH)2i is sig-
nificant only under basis conditions.  Patterson _e_t _a_l. (1977) compared the
precipitation behavior of nickel carbonate and nickel hydroxide, in the con-
text of treatment of nickel-bearing waste effluents.  Although precipita-
tion as the hydroxide was found to be the more efficient treatment, the
lowest nickel concentration attained at pH values below 9 was 15 mg/1.
This level is quite high' with regard to its toxicity and indicates that
precipitation , is not an effective control on aickel under most conditions.

         Furthermore, huaic • acids in natural waters alter the solubility
and precipitation behavior of 'nickel.  Sashid and Leonard (1973) exposed
nickel carbonate to.humic acid and found, that complexation with huaic acid
soi'Ailized much of the nickel.  Approximately '200 mg of nickel was re-
leased per gram of huaic acid added.  Addition of tuuaic acid did not solu-
bilize nickel sulfide,  When hutnie acid was added to solutions that con-
tained nickel and some natural inorganic Iigands (either sulfide, carbon-
ate, or hydroxide), the quantity of nickel required- to cause precipitation
increased dramatically.  Hustle acids are ubiquitous in natural "aters and
may be expected to increase the solubility of nickel under natural condi-
tions to the point that precipitation ia probably not a significant fate.

    15.3.3  Volatilization

         No evidence was found to suggest that volatilization of nickel
compounds occurs froa the aquatic environment.

   •15.3.4  §or£t_ion

         Sorption of nickel by hydrous iron and manganese oxides and
organic material probably exerts the major control on the mobility of

nickel in the aquatic environment.  Nickel, however, Is a highly mobile
metal and is sorbed only to a small extent.  Lee (1975) presented cogent
evidence for the importance of hydrous iron and manganese oxides in con-
trolling nickel concentrations la aquatic environments.  When these oxides
precipitate, nickel is attracted to the negative zeta potential  they usual-
ly exhibit and can become incorporated into the crystal lattice  structure
of the dewatering oxides.  Some analyses of bottom sediments have corrobo-
rated this hypothesis, showing a strong correlation between nickel concen-
tration and iron and nanganese concentrations (Angino jet al. 1974; Steele
and Wagner 1975),

         Gibbs (1973) found that most of the nickel in the Amazon and Yukon
River systems was associated with suspended particulates» organic material,
or coprecipitated with hydrous iron and manganese oxides.  In contrast to
Gibbs' (1973) results, Perhac (1972, I974a) found that almost all of the
nickel transported by two Tennessee streams was in the dissolved fora.  The
reason for this discrepancy is probably the fact that about 90%  of the
solid? in the streams studied by Perhac were dissolved solids, so that
there were very few suspended particles available for coprecipitation/
sorption reactions.  Suspended solids probably compose a greater fraction
of the total solids In Gibb's study, provi-'lng a substrate for coprecipita-
cion or sorption.  Perhac's (1974b) analyses of bed sediments indicated
tha^ much of the sedimentary nickel was incorporated in the crystalline
structures of carbonate minerals.  Thus, isooorphous substitution of nickal
cor other cations in lattice sites aay fee another process affecting the
distribution and mobility of nickel.

         The partitioning of nickel into dissolved and particulate frac-
tions is undoubtedly related to the abundance of suspended Material, can-
petition with organic material* and concentrations of iron and manganese.
Hydrous iron and manganese oxides precipitate as a coating on suspended
particles, and attract nickel and other metals.

         Suspended organic matter may be a good adsorbent for nickel,
Rashid (1974) used colloidal humic substances to adsorb nickel and found
that of the nickel thus bound, only 262 could be extracted by ammonium ace-
tate*  Iron chloride added to the solution extracted 76?, showing that pre-
cipitating hydrous iron oxides probably attract nickel more strongly than
organic matter.  The iaportance of these competing processes is  strongly
affected by the relative abundance of sorbents (organic natter,  hydrous
metal oxides) in the water column and sediments.  Although organic matter
can adsorb nickel, Adams ejt aJU (1975) reported that sewage treatment by
primary digestion and activated sludge removed less than 452 of  the nickel
that was present.  Organic material obviously abounds in such environments,
and yet apparently little nickel is adsorbed.

         Jackson and Skipped (1978) found, in a. recent laboratory study,
that organic acids Increase the solubility of nickeT  in the presence of
clay,  .Fulvic and humic acids were shown to be capable of remoblltzing
nickel froo all solid phases, although the reaction appears to be kineti-
cally hindered, especially,ac basic pH values.

         Although data on the sorptlon of'nickel in the aquatic environ-
ment are somewhat limited, a few general conclusions can be reached.  In
natural, unpolluted waters, for exanpie, it would appear' that sorption
processes are at lease noderately effective in limiting the mobility of
nickel in the aquatic environment,  in the rao.re organic-rich, polluted
waters, it would appear that little sorption, if any, of nickel takes
place.  In either eas«, the lack of other controls probably oakes incorpor-
ation ineo bed sediments «n inportant fate of nickel in surface waters.  It
would appear, however, that much of the nickel entering the aquatic en-
vironment will be transported to the oceans.

    15.3,5  B ioaccuau1at ion

         Nickel is bioaccunmlated by some aquatic organisms, but most con-'
centration factors, are less than 1CP.  long (1974) showed that nickel
does not bioaccumulate in lake trout, Salvelinug namaycush.  Friedrich and
Filice (1976) studied the accumulation of nickel by the mussel (Mytilkus
gdults) in artificially prepared seawater under static conditions.  No
significant accumulation was noted after four weeks'  exposure to 0.03 mg
Mi/1, but significant uptake was acted at all concentrations exceeding
0,056 mg Nl/1; rates of nickel elimination were not measured.

         In a study of the accumulation of iron, zinc, leaa, copper, and
nickel by algae collected near a zinc smelting plant, it was found that
nickel exhibited the lowest concentration factor for all metals tested
(Trollops and Evans 1976).  Slcaar _e£ _«!. (1974) demonstrated that uptake of
nickel by the diatom ?haeadactyl;~' trlcornutua is strongly dependent on
raetabolic state and is affected by the phosphate concentration in the

         Wright (1976) observed nickel concentratiuns exceeding 7.0 pg Nl/g
in auscle' froo taarine fishes collected fron the northeast coast df. England,
and Romeril anrt Davis (1976) reported that European eel's (Anguilla_anguil-
la) maintained in Treat River water averaged (dry basis) 21ug Nl/g in
auscle and 16 in liver.

         In general, nickel ia not accumulated in significant amounts by
aquatic organisms.  Concentration factors for organisms shown to accumulate
nickel are given in Table 15-1.

                    Bioconcentracion  Factors  .or  Nickel

Freshwater planes

Freshwater invertebrates'

Freshwater fish
Marine planes



Marine plankton

Marine sponges      <

Marine invertebrates

Marine fish

Skipjack tuna





   550 - 2,000

 2,000 - 40,000

   <20 - 8,000





Chapman e_£ alt 1968

Chapman ejt al. 1968

Chapman jet_ al. 1968

Chapman e_t a 1. 1968

Stunm and Morgan 1970

Stunna and Morgan 1970

Stuarn and Morgan 1970

Stumm and Morgan 1970

Chapman e_t al. 1968

Chapman _e_C_ al. 1968

Stumm and Morgan 1970
a.  Concent radon factors are defined  by  the  ratio  of  the  concentration of
    the element in the organism in ppta  (wet weight)  divided by the
    concentration of the element in water (ppo).

    15.3.6 ,Biotransformation

         No data was found In the available literature to suggest that
nickel is involved in any biological trans formation in the aquatic en-

15.4  Data Safflma r y

    The mobility of nickel in the aquatic environment is controlled largely
by the capability of various sorbents to scavenge it from solution.  Al-
though data is limited, it appears that in pristine environments, hydrous
oxides of iron and manganese control nickel's mobility via coprecipitatlon
and sorption.  In polluted environments, the more prevalent organic
material will keep nickel soluble.  In reducing environments, insoluble
nickel sulfide may be formed.  Although nickel is bioaccumulated, the con-
centration factors are such" as to suggest that partitioning into the biota
is not a dominant fate process.  Nickel is one of the most mobile of the
heavy asetals in ch* aquatic environment.  The aquatic fate of nickel is
summarized in Table 15-2.

                                fable 15-2
                     Summary of Aquatic Fate of Nickel
     Process  ,


Chemical Speclation*


Roc an important process.

In aerobic environments below
pH 9, soluble compounds are formed
with hydroxide, carbonate, sulfate
and organic*. Above pH 9, precipi-
tation of the hydroxide or carbo-
nate will occur.  la reducing en-
vironments, NiS will precipitate.
Not a regulating factor in most

Not an important process.
Nickel is the most mobile of the
heavy metals.  Coprecipltation
with hydrous metal oxides, sorption
into organic material, and ion ex-
change with crystalline minerals
are the dominant factors which affect
its mobility.

Reported blocoucentration factors
are on the order of 102-1CH.
Not a dominant process.

Not an important process*
Confidence of


    All of the noted environmental processes are important; however, their
    relative importance witr respect to each other is uncertain for
    determining final fate.

15.5  Liee ra ture_ Cii&A

Adams, C.E., Jr., W.W. Sckenfelder, Jr., and B.L. Goodman.  1975.  The
  effects and removal of aeavy metals in biological treatment.
  pp.277-292. in Heavy metals in the aquatic environment.  P.A. Krenkel
  (ed.).  Pengamon Press, Oxford, England.  3S2p.

Angino, E.E., L.M. Magauson, and T.C. Waugh.  1974.  Mineralogy of
  suspended sediment and concentration of Fe, Mn, Ni ,  Zn, Cu, and Pb in
  water and Fe, Mn, and Pb in suspended load of selected Kansas streams*
  Water Resources Res.  10(6):1187-U91.

Chapman, W.H.,  H.L. Fisher, and M.W. Pratt.  1968.  Concentration factors
  of chemical elements in edible aquatic organisms.  Lawrence Radiation
  Laboratory, Livertnore, Calif.  UCKL-50564.  46p.

Corrich, J.D.  1973.  Nickel. . pp.1-15.  j.n 1973 Bureau of Mines Minerals
  Yearbook..  U.S. Department of the Interior, Washington, D.C.  396p.

Cotton, F.A., and G. Wilkinson.  1972.  Advanced inorganic chemistry.
  pp.890-903.  Interscience Publishers, New York.  1145p.

Friedrich, A.R. and F.P. Filice.  1976.  Uptake and accumulation of the
  nickel ion by Mytilus eduljs.  Bull. Environ. Contam. Toxieol.

Gibbs, S.J.  1973.  Mechanisms of trace metals transport in rivers.
  Science.  180:71-73.

Jackson, K.S. and G.8. Skippen.  1978.  Geochemical dispersion of heavy
  metals via organic coraplexing:  a laboratory study of copper,'lead, zinc,
  and nickel behavior at a simulated sediment - water boundary,
  J. Geochem. Exptor.  10:117-138.

L«e, G.F.  1975.  Role of hydrous metal oxides in the transport of heavy
  metals in the environment,  pp.137-149.  in Heavy metals in the aquatic
  environment.  P.A. Krenkel (ed.).  Pergamon Press, Oxford, Englai.a.

Patterson, J.W., H.E. Allen, and J,J.• Scala.  1977.  Carbonate
  precipitation for heavy taetal pollutants.  J. Water Pollut. Control
  Fed. 49(12):2397-2410.

Perhac, R.M.  1972.  Distribution of Cd, Co, Cu, Fe, Mn, Ni, Pb, and Zn in
  dissolved and particulate solids from two streams in Tennessee.  J. -
  Hydrol. 15:177-186.

Perhac, R.M.  1974a.  Water transport of heavy metals in solution and by
  different sizes of partlculate .solids.  Univ. of Tenn.  Water Resources
 .Res. Ctr.  Project 23.  Knoxvtile, Tenn.  41p.

Perhac, R.M.  1974b.  Heavy metal distributions In bottom sediments and
  water in the Tennessee River - Loudan Lake reservoir system,  Univ. of
  Tenn.  Water Resources Res. Ctr.  Project 40.  Knoxville, Tenn.  18p.

tashid, M.A. , and J.D. Leonard.  1973.  Modifications in the solubility and
  precipitation behavior of various metals as a result of their interaction
  with sedimentary humic acid.  Chea. Geol.  11:89-97.

Rashid, M.A.  1974.  Adsorption of metals on sedimentary and peat humic
  substances.  Cheau Geol. 13:115-123.

Rotneril, M.G. and M.H. Davis*  1976.  Trace metal levels in eels grown in
  power station cooling water.  Aqua culture.  8:139-14';.
Skaar, H, , B. Rystal , and A. Jensen,  1974.  The uptake of    i by the
  diatom P hoe odac t y 1 um ricqrnutua.  Physiol. Plantum,  32:352-358.

Steele, K.F« and G.H* Wagner.  1975.  Trace metal relationships in bottom
  sediments of a freshwater stream - The Buffalo River, Arkansas.  J, Sed.
  Pet. 45(1):310-319.

Stumm, W., and J.J. Morgan.  1970.  Aquatic chemistry,  p. 289.  Wiley
  Interscience » New York.  583p.

Tong, S.S.C.  1974,  Trace metals in Lake Cayuga lake trout (Salvelinas
  namayaush) in relation to age.  J, Fish. Res. Bd. Can.

Trollope, D.R, and B. Evans.  1S76,  Concentrations oz copper, iron, lead,
  nickel and zinc in freshwater algal blooms.  Environ. Pol lot.
  115109-116.                                                  i

Weast, 8.C* (id.).  1977.  Handbook of chemistry and physics, 58th
  edition.  CRC Press, Cleveland, Ohio.  2398p.

Wright, D.A.  1976.  Heavy metals in animals from the northeast coast.
  Mar.  Poliut, Bull.  7(2>:36-38.

                               16.  SELENIUK

16.1  Stateaent of ProbableFate

     In  aerobic waters, selenium  is  present  in'the  selenite
SeQ3~2) or  selenate  (H2SeQ4,HSe04~,Se04~2)  oxidation  state of 4+ or
These chemical species are very  soluble, and most  of ,the selenium  dis-
charged into  the  aquatic environment  is  probably  transported in these forms'
to  the  oceans.  Under reducing conditions,  selenium can fora metal selenides
either  by direct  reaction with oetals or through substitution for  sulfur  in
metal sulfides.   Most of the metal  selenides, however, have a very low solu-
bility  in water.

    Selenium has a sorptive affinity for hydrous metal oxides,  clays  and
organic materials.  Sorptlon by bed sediments or suspended solids can re-
sult in enrichment of selenium concentrations in the bed sediments,  Sorp-
tion or precipitation with hydrous iron oxides is probably the  major  con-
trol on nobility of selenium in aerobic waters.

    Selenium can be methylated by a variety of organisms,  including benthic
microflora.  In a reducing environment, hydrogen selenide 
  .  Geochemically, selenium resembles sulfur.  Sulfide or elemental sulfur
deposits very often contain significant amounts of selenium.  Metal sul-
fides have been found to contain selenium occasionally at levels of over 20
percent (Davidson 1960).  Jarasite and barite, two sulfate minerals, have
also been found Co contain selenium, but at relatively low levels.  Crude
sulfur often contains selenium at well over 1 percent.  Deposits of sulfur-
containing minerals are often geochemically secondary in nature, and when
selenium occurs in them, it has probably been leached from some other
material and redeposited.

    Selenium probably occurs as a free element or, more.likely, as a metal1
selenide, in unweathered rocks.  It is apparently readily oxidized during
the weathering 'of crustal materials.  In areas of acidic soils, selenium is
probably present as the selenlte anion which is firmly bound in iron oxide
colloids; while in alkaline soils, it should oxidize further to the very
soluble selenate anion.  Selenium, atonic number 34, atomic weight 78.96,
has an elemental melting pfiint of 217°C and a boiling point of 684.9"C
(Weast 1977).

    The CAS number of selenium is 7782-49-2; its TSL number is B562-1046.

16-3  Sumaary ojTFate Data

    16.3.1  PhotQlyjis          ^                             . •

         Although the photoconductive and photovoltaic properties of selen-
ium are of prime importance in the commercial applications of selenium,  no
data were found to suggest that photolysis reactions play an important role
in determining the aquatic fate of the selenium present in the aquatic en-

    16.3.2  ChemicalL Speclation

         As previously noted, selenium is stable in four valence states:
-2f 0, +4 , and +6.  The positive, oxidation states are present in oxyanlons
and organo-selenium compounds, the bond types of which exhibit a cjvalent
character (Cotton and Wilkinson 1972).  The inorganic forms govarn the
physical chemistry of selenium in solution.  Coleraan and Delevaux (1957)
calculated the stability of a selenium-selenite-selenate system (total Se
was 10~*M whic.h is equivalent to 80 yg/1) under various Ehr.pH conditions
(Figure 16-1).  As is evident from the stability diagram, elemental selen-
ium is stable in a wide range of redox and pH values, being favored by low
pH and reducing conditions.  In aerobic conditions, selenium is in the form
of the soluble HSeOj", S«Q;j~2, and S«04~2 anions.

         The major features of selenium chemistry that affect its movement,
toxicity, and deficiency in the environment are associated with changes  in

       O *0.4
                                                 10     12     14
Figure 16-1  Stability Field of S*leniuo,  (Temperature 25"C, pressure  1
             atm,  Concentration of S« 10~^M),   From Coleman and
             Delevaux (W57),

Its oxidation state and the resulting difference in chemical properties.
Therefore, during the following discussion, each oxidation state vill be
discussed in turn.

         Properties of aelenate selenium (+6 oxidation state).  Selenic
acid (H2§«04) is a strong acid; its salts exhibit similar solubilities
as the sulfates of the same metals (Rosenfeld and Beath 1964).  Soluble
selenates would be expected in alkaline environments and even though one
would expect selenate to be converted to selenite or eleaental selenium in
acidic environments, this conversion appears Co be kinetically inhibited
(Lakin 1973).  Because of its stability at alkaline pH, its solubility and
its ready availability to plants, selenate appears to be the most dangerous
fora of selenium as far as potential environmental pollution is concerned.

         Properties_of_sjgleni.te ^eleniua	(+4 oxidation__state).  Selenious
acid (825*63) is a weak acid, and any dissolved selenite would be pre-
sent predominantly as the biselenite ion in water between pH 3.5 and 9.0
(Coleman and Delevaux 1957)*  Most selenite salts are less soluble than the
corresponding selenates (fictional Research Council 1976).  Of special rele-
vance, with respect to the aquatic environment, Is th* very low solubility
of ferric selenites (Geering _et al. 1968).  Another characteristic of sele-
nlts (of importance to environmental cycling of selenium) is the property
of selenite to rapidly become reduced to eleaental selenium under acidic
conditions by alld reducing agents, such as ascorbic acid or S02 (Rosen-
feld and Beath 1964).  The probability that selenite will either form in-
soluble compounds, adsorbs tea with ferric oxides, or be reduced to insolu-
ble elemental selenium minimizes the possibility for its transport in the
aquatic environment*

         Properties of elementalselenium (0 oxidation state).  Various
allotroplc forms of elemental selenium are reviewed and their solubility in
various reagents is tabulated in Rosenfeld and Beath (1964).  The elec-
tronic and photoelectric properties of elemental selenium are responsible
for many industrial uses of the element.  The extreme insolubility of ele-
mental selenium in aqueous systems is of primary importance in any dis-
cussion of the fate of selenium in the aquatic environment.  Insoluble
elemental selenium appears to be a major sink for selenium that can be
considered to be inert In the aquatic environment and is quite important
for determining the overall fate of the element.

         Properties ofselenide selenium (-2 oxidation state;.  Hydrogen
selenlde is a fairly strong acid, and it is highly toxic.  This compound
rapidly decomposes, however, to form elemental selenium, and it is thus of
minor importance to the overall fat* of selenium in the aquatic environ-

         Organic cheatst-yof selenium.  Important features of the organic
chemistry of selenium were reviewed In the proceedings of a symposium
edited by Klayman and Gunther (1973).  As far as the aquatic environment Is.
concerned, the imp rtant feature of the organic c'-ieai.stry of seleraium Is
that essentially all of the relevant compounds contain selenium in the -2
oxidation state.  These compounds will decompose in the environment to form
elemental selenium which-will be insoluble and relatively inert1 to furthe-
chemical reaction.

    16.3.3  Volatllijation

         Methylatlon of selenium occurs in mammals, plants, and microbes
(Stadtman 1974).  The selenium compounds thus produced are volatile and may
escape to the atmosphere.  Quantitative data, however, are lacking bn this
process In the aquatic environment.

         Of the studies performed on selenium methylation, the one most
applicable to aquatic fate was reported by Chau jst _al. (1976) who demon-
strated that microbes present in lake sediments in the Sudbury, Ontario re-
gion could methylate organic and Inorganic selenium compounds Including
sodium selenite and sodium selenate..  The ability of aquatic microflora to
produce gaseo'is dimethyl selenide or dimethyl diseienide from additions of
5 mg/1 sodium selenite or sodium selenate was observed In 11 of 12 sediment
samples.  It is noteworthy that 4 of the 12 sediment samples evolved
methylated selenides even without added selenium (sediment concentrations
of selenium for these four samples ranged from O.A8yg/g to 20.48 pg/gm).
Calculations from the data presented by Chau et al. (1976) show that for
the four cultures in which volatile organic selenium was produced from
selenium in the sediments, the highest ratio of volatilized selenium to
Initially sedimented selenium was 1.6 x 10~3,  fhe duration of the exper-
iment was one week.  Clearly, biotnethylation with subsequent volatilization
oay be an important fate process for seleniun.

         Volatile selenium compounds can also be forned by Inorganic means.
For example, volatile &2S€ can ^e formed under reducing conditions; how-
ever, because of the lack of quantitative data, the importance of this
latter process for the overall aquatic fate of selenlun Is difficult to

    16.3.4  Sorptton

         S«l«ni£« is adsorbed by Iron and manganese hydroxides, organic
matter, and Iron sulfides (Rosenfeld and Betth 1964, Kharkar et_ al. 1967).
Kharkar _et al. (1967) used a lyg/1 solution of selenium that contained
radioactive 'Sjg to study the sorption of selenium by various materials.
The hydrous Iron and aanganese oxides have the greatest affinity for selen-
ium on a mass basis, but these materials usually comprise a smaller

proportion of the suspended solids Chan do the clays.  Thus, mcst of the
selenium in aquatic systems is probably transported as the dissolved
species.  Subsequent to determining the adsorption of selenium by various
sorbents in distilled water, Kharkar e_t _a_i. (1967) tested the amount
desorbed on contact with sea water.  Freshly precipitated ferric hydroxide
released the smallest percentage, indicating that selenium Is probably
co-precipitated and incorporated into th* structure of the ferric hy-
droxide.  Selenium, thus bound, would b« less likely to be desorbed.  Based
on the composition of the suspended solids load of typical rivers and the
sorption of common materials normally comprising the solids load, Kharkar
et al. (1967) estimated that selenium is supplied to the oceans predomi-
nately as dissolved species with desorption from solids making up only
about 10Z of the total.

         Although no other systematic study of selenium sorption was found
in the reviewed literature, a recent study by Frost and Griffin (1977) on
the effect of pi! on adsorption of selenium by clay minerals from a landfill
leachate may have aquatic relevance.  They found a pronounced effect of pH
on the amount of selenium adsorbed by clay minerals.  This strong pH depen-
dency of adsorption results from the effect of pH on the distribution of
selenium species present in solution and the inactivatlon of anlon adsorp-
tion sites on clay minerals.  Their results indicate that selenium would be
quite mobile in clays, especially under alkaline conditions.  It appears,
however,  that only in areas of high hydrous iron oxide concentrations,
will selenium be sorbed into the bed sediments.

    16.3.5  BioaccuaulatioQ

         Selenium is bioacctnulated by a number of aquatic organisms (Chau
and Riley 1965; Fowler and Benayoun 1976b; Sandholm e_t aj.. 1973), and in1
some instances, it is strongly correlated with the concentration of mercury
and other heavy metals in th* organism (Koeman «t jil. 1973; Mac Kay et al.
1W5).  The concentration factors for aquatic organisms have been summa-
rized In Table 16-1.  Very little quantitative data are available on the
bioaccumulation and excretion kinetics of selenium in aquatic biota, and,
consequently, very little is known about the parameters Chat control the
biological cycling of selenium in the aquatic environment,
         Sandhola e_t al. (1973) studied selenium uptake in a laboratory
food-chain consisting of water, phytoplankton, zooplankton and fish; both
selenite and selenoaethicnlne were used.  In these experiments, food was
determined to be the moat laportant Source of selenium to fish, and very
little'selenium accumulated from the water.  In a saltwater study, Fowler.
and Benayoun (1976a) measured selenium uptake In a marine shrimp (Ly^slnata
seticandata) exposed to '%e in food and water and in a mussel (Mytllua
|alloprovlnciali s) exposed to 75$« in water.  Shrimp exoskeleton con-
tained 60 to 90 percent of the selenium accumulated when selenium was taken
up via food.  Only 10 percent of the selenium present in the exoskeleton

                                Table  16-1

                   Bioconcentration  Factors  for  Selenium

                            BioCigncen tra t ion
      Taxpn        •             Factor3Reference

?r«»hwater planes                $00                Chapman et  al.  1963

Freshwater Invertebrates         400                Chapman _e_t  al.  1968

Freshwater Fish                  400                Chapman et  al.  1968

Maria* Plants                    800                Chapman et  al.  1968
Marine Invertebrates             400                Chapman _et_  al.  1968

Marine Flah                      400                Chapman et  al.  1968
a.  Bioaccuaulation factors are  the  ratio derived  from  the  concentration of
    the element in the aquatic organism  (in ppm  wet  weight) divided by the
    concentration of  the element  in  water (in  ppa).

 was  l^st  during  molting.   Mussels  accumulated  the  highest  concentration  of
 selenium  in viscera.   Fowler  and Benayoun  (1976c)  found  similar  results
 with euphacids  (Maganyctiphanes n£rvegica).   Dietary selenium  was  retained
 at an efficiency of 66 percent.  The 'biological  half-life  of selenium  was
 37 days,  and  whole body concentration  factors  were estimated at  11,500 to
 7,500 for this  species.

          Selenium has  often been described as  the  "Dr. Jek.yll;and  Mr.  Hyde"
 element because  of living  organisms* need,  for  selenium at  low  concentra-
 tions! and selenium's extreme  toxicity  at high  concentrations (Stadtnan
 1974),  It is also noteworthy, that selenium  in the diet  is known to  exert a
 protective influence against  mercury poisoning (Ganther  e_t aJL.  1972).  Kim
 ££ a_l, (1977) found that creek chubs (3emotilu3  atromaculatus).  immersed in
 water containing 3.0 mg Se'/l  for 48 hours, were  less susceptible to  mercur-
 ic chloride in  water than  untreated individuals.   Furthermore, at  mercury
 concc  rratlons  below 0.07  mg  Hg/1, selenium  treatment increased  mercury
 accumulation, but at mercury  levels above  0.10 mg  Hg/1,  selenium inhibited
 mercury accumulation.   So  speculation  was  offered  on the mechanism of  this
 interaction,  Koeman «t al. (1973)' showed  that mercury and selenium  are
 present In the  livers  of marine mammals  in a 1:1 molar ratio whereas marine
 fishes usually  contain upwards of  40 times more selenium than  mercury.
 Thia result suggests that  selenium and mercury may occur together  in marine
 mammals,  perhaps resulting in seme degree  of protection  against  mercury,

          The  small amount  of  available data  suggest that while dietary
 seleniura  is the most important source  of selenium  to many  marine and fresh-
 water organisms, little biomagnification takes place. In  fact,  in some
 Instances, the  accumulation of selenium by organisms may be beneficial due
 to the protective action which selenium  provides against mercury.  Because
 of selenium's high toxiclty,  however,  the  relationship between seleniura
 toxicity  to aquatic; organisms and  selenium accumulation  is uncertain,

   .16,3.6  Bio t r an sf Q raia 11 on

          The  biochemistry  of  selenium  compounds  has been well  studied  in
 terrestrial plants and animals (Stadtman 1974; , National  Research Council'
 1976).  Such  studies,  however, have not been  been carried out in  the  aqua-
 tic  environment. Selenium blotransformations  have been  studied  in terres-
 trial environments because of the  agricultural  problems  encountered  In the
 high selenium soils of the Rocky Mountain  states.   Selenlte and  selenatc
 are  both  taken  up by plants and, within  the  plant, these fnrras of  selenium
 »re  reduced to  the -2  oxidation state. The  Se""^ ion is  then incorporated
,into soluble  amino acids and/or protein-bound  amino acids.  This reduction
 of selenium,  especially the selenates, within  the  plant  way not  be quanti-
 tative.   Monogastric animals  may reduce  selenate and selenite, but thev
 apparently do not 1. corporate the  reduced  selenium inm  amino  acids.   The

"selenocrisulfides," formed by the reaction of selentte with sulfhydryl
groups of amino acids,  peptides, and proteins, are a probable first- product
of the reduction of physiologic doses of selenice in monogastric animals
(Ganther 1971).  Major  excretory products of selenium metabolism in animals
are trimethyl selenonium ion in the urine and elemental selenium or metal
seslenides in the feces.

         Some of,the chemical changes possibly invdlved in the aovcment of
selenium from soils through plants and animals are diagramed in Figure
16-2,  When, the metabolic pathways of selenium in plants and aniaala are
considered along .with the reactions of selenium in soils, it appears tuat
the conversion of fhe element to inert and insoluble forms may occur in
aquatic systems.  It should be kept in mind that although extrapolation of
biochemical mechanisms  from the terrestrial to the aquatic environment
night be valid, only further study can truly determine the biotrans forma-
tions o,f selenium compounds In the aquatic environment.  Monetheless, pro-
duction of hydrogen selenide or biomethylation by sedimentary mireoflora
could remobilize the selenium and result in significant recycling.

16.4  DataSummary

    Most of the selenium introduced into the aquatic environment is prob~
ably transported in solution to the oceans.  Sorption/coprecipitatiori re-
actions with hydrous iron and manganese oxides, and to a lesser except,
clays and organic  materials, control the mobility of selenium in oxidizing
conditions.  In reducing conditions, insoluble metal selenides can form,
thereby decreasing nobility, or ti^Se can form, increasing mobility.
Biologically mediated tnethylation of selenium results' in production of
volatile compounds, another mode by which sediment-bound selenium can be
rsoobilized.  Table 16~2 summarizes the aquatic fate described above.

Add-$ccrty aerated
Heavy metai<— - *3emefrtaJ
                                     Sciis   ,
                                               Well aerated-Alkaline
' INsJuM
V i
JI leaching

           St mettiicnine
           Se adenosyf
           Se mettiicntne  \

                                      metfryi-Se cysteine
                                  •Se cystathionine
                                   mettiyt setenides-
                                                  to soils
Protein hydrolysis

Se methicnii
S« cysteine

S0 cysteric acid

St'  k
                                                   Methyl setenides

•   Exnalation
To atmoscnera
                                                      To soils
Figure 16-2  Th« cycling of selenium in  th*  terrestrial biologic system.
             Modified froa National  Research Council  (1976).

                                fable 16-2

                    Summary of Aquatic Fate of Selenium
Confidence of



Bioaccumulat ion*

Blot cans fo rraat ion*
Not an important process.

Controls solubility.  Under
anaerobic conditions and/or
low pH, insoluble elemental
selenium is formed.  Under
other condition, soluble
complexes are formed.

Nay occur via bionethylation
or formation of HjSe.

Hydrous octal oxides sorb
selenium strongly.  Clays
and organic materials have
a lesser affinity.

Concentration ratios depend
on chemical form in soils
and organism.

Metabolism may result in
methflation with subsequent

    All of the noted environmental processes are Important; however, their
    relative importance with respect to each other is uncertain for
    determining final fate.

16.5  Literature Cited          ,                      •

Chapoaa, W.H., H.t. Fisher, and M.W. Pratt.  1968.  Concentration factors
  of chemical elements In edible aquatic organisms.  Lawrence Radiation
  Laboratory, Llvermore, Calif. 46p.

Chau, Y.K. and J.P. Riley.  1965.   Th* determination of selenium in sea
  water, silicates, and marine organisms.  Anal. Cheat. Act a 33:36-49.

Chau, Y.R., P.T.S. Wong, 1.A. Silverberg, P.L. Luxon, and C.A. Bengert.
  1976*  Methylttion of selenium In the aquatic environment.  Science

Coleaan, R.A. and H. Delevaux.  1957,  Occurrence of selenium in sulfides
  fro* some sedimentary rocks of Western United States.  ECon. Geol.

Cotton, F.A. and 6. Wilkinson.  1972.  Advanced inorganic chemistry.
  pp.519-558,  tnterscience Publishers, Ittc. New York. 1145p.

Davidson, O.F.  1960.  Selenium in some epithermal deposits of antimony,
  mercury, silver, and gold,  U.S. Geological Survey Bull, 1112-A.
  Washington, D.C. 15p,

Fowler, S.W. and A. Benayoun.  1976a.  Accumulation and distribution of
  seleniua in mussel and shrimp tissues.  Bull. Environ. Contam.  Toxlcol.

Fowler, S.W. and A. Benayoun.  1976b.  Influence of environmental factors
  on selenium flux in two marine invertebrates.  Marine Biol. 37:59-68.

Fowler, S.W. and A* Benayoun.  1976c.  Selenium kinetics in marine
  sooplankton.  Mar. Scl. Coamun.  2(1) J43-67.

frost, R.R. and R.A. Griffin.  1977.  Effect of pH on adsorption of
  arsenic and seleniua from landfill leachate by clay minerals.   Soil Sci.
  Soc. Aner. Journ. 41;53-57,

Canther, H.I.  1971.  Selenium: the biological effects of a highly active
  trace substance*  pp.211-221.  in Trace substances in environmental
  health.  Vol. IV.  O.D. Reaphlll (td.).  Uhiv. of Hissouri Press.,
  Columbia, Hissouri.  547p.

Ganther, H.E., C. Coudie, M.L. Sunde, M.J. Kopecky, P. Wagner, S-H.  Oh,
  and W.G. Hoekstra.  1972.  Selenium: relation to decreased toxicLty of
  methylmercury added to diets containing tuna.  Science 175:1122-1124,

Gearing, H.R., E.E, Gary, L.H.?. Janes, and W.H. Allauay.  1968.
  Solubility and redox criteria for the possible forms of selenium In
  soils.  Soil Sci. Soc.  Araer, Proc. 32:35-40,

Kharkar, O.P., K.K, tarekian, and K,K. Sentiae.  1967.  Scream supply of
  dissolved silver, molybdenum, antimony, selenium, chromium, cobalt,
  rubidium and cesium to the oceans.  Geochia. Cosmochim. Acta.

Kim, J.H., E. Birks, and J.F. Heisingec.  1977,  Protective action of
  selenium against mercury in northern creek chubs.  Bull, Environ, Contain.
  Toxicol. 17:132-136.

(dayman, D.L, and W.H.H. Gunther.  (eds.).  1973.  Organic selenium
  compounds:  their chemistry and biology.  Wiley-lnterseience Publishers,
  New York.  <1188p.

Koeman, J.Hj, W.H.M. Peeters, C.H.M. Koudstaad-Hol, P.S. Tjioe and J.J.M.
  D«Goeij.  1973.  Mercury-selenium correlations in marine mammals.  Nature

Lakin, H.U.  1973.  Selenium in our environment,  pp.96-111.  in Trace
  elements in the environment.  E.I,, Kothny (ed.).  Advances in Chemistry
  Series 123.  American Chemical Society, Washington, D.C. 245 p.

MacKay, M.J., M.N. Kazacos, R.J. Williams, and M.I. Uedow.  1975.
  S*l«nium and heavy metals in black oarlin.  Mar. Follut. Bull. 6:57-61.

National Research Council.  1976.  Selenium,  pp.1-42, 59-72.  National
  Academy of Sciences.  NTIS Springfield, Va. PB-251 318.  310 p.

Rosenfeld, I. and O.A. Beach.  1964.  Selenium: geobotany, biochemistry,
  toxicity and nutrition, pp.d-61.  Academic Press, New York.  411p.

Sandholm, M., H.E. Oksanen, and L. Pesonen,  1973.'  Uptake of seleniun by
  aquatic organisms.  Limnol.  Oceanogr.  18:496-499.

Stadtman, T.C.  1974.  Selenium biochemistry.  Science 183:915-922.

Weast, R.C. ed. 1977.  CRC handbook of chemistry and physics.  58th ed.
  CRC Press, Cleveland, Ohio.  2398 p.

                                17.  SILVER
17. i  Satement of
     Sorption and precipitation processes are effective in reducing the
concentration of dissolved silver and result in higher concentrations in
the bed sediments than in the overlying wacers.  Sorption by manganese di-
oxide and precipitation with halides are probably the dominant controls on
the mobility of silver in the aquatic environment.  Sams silver is also
bio accumulated , and the remainder is transported in solution to the oceans.

17*2  Identification - Geochemistry of Silver

     Silver is a rare element occurring in concentrations of about 0.1 ppm
in the earth's crust (Weast 1977} »  Silver is toxic tto aquatic bacteria,
invertebrates and fish.  Its toxicity ranks second only to mercury among
the heavy metals (Freeman 1977).
     Silver occurs primarily in the, form of the sulflde (argentite
or intimately associated with other taetal sulfides, especially .those of
lead and copper.  Other common silver minerals Include cerargyritc (AgCl),
proustite (3AgS«As2S3>, pyrragyrite (3Ag2S*Sb2S3) , stephanite
(5Ag2S*Sb2S3) and native metallic silver.  Host lead and copper
ores are argentiferous though there are important exceptions.  Recovery of
silver and gold from these .ores constitutes an important part of their
metallurgical treatment.

     Silver is also commonly associated in nature with gold.  Not only does
gold occur with silver in copper and lead ores, but native metallic gold
usually contains silver.  Gold and silver are mutually soluble in each
other in all proportions in the metallic state.

     Silver is usually found in extremely low concentrations in the aquatic
environment due both to its low crustal abundance and the effectiveness of
controls on its mobility in water.  In a study of ten U.S. rivers, Kharkar
»t ail. (1968) detected silver in concentrations ranging from 0.092 to 0.55
yg/1.  Hea (1970) cites studies of public drinking water supplies and riv«r
waters which report median concentrations of 0.23 and 0.09 yg/1, respect-
ively.  The geochemistry of silver has been extensively reviewed by Boyle

     Silver Is a transition metal, atomic number 47, atomic weight 107,9
(Weast 1977).  Other than in complex ions, silver's only stable cationic
state Is Ag* (Cotton and Wilkinson 1972).

      The  CAS  number  of  silver  Is  7440-22-4,  and  its  TSL  number is  .

'17.3   Summary of  Face  Data

      17,3.1   Photolysis

           Although  the  photochemical  properties  of  some  silver cooponnds
 (notably  the  bromides  and ,thiosulfates)  are  well knowrt atid  form the  basis
 of  photographic  chemistry,  no  data  was  found to  Indicate that  these  proces-
 ses are  important In determining  the  aquatic fate of silver.

      17.3.2   Chemical  Speciation

           In  natural waters, silver can  "be found t"  several 'chemical forms,
 for example,  aquated -atlons,  metal-Inorganic complexes, and  metal-organic
 complexes.  An understanding of the chemical speclation  of  silver  in any
 given situation  is  based upon  theoretical  calculations of hydrolysis,
 oxidation/ reduction  and organic coaplexatlon.

           Ionic  silver  is  found in  aqueous systems  as the unlvalent
 species,  although it can form  compounds  In the Ag(II) or Ag(III) states
 (Cotton  and Wilkinson  1972).   Metallic  silver is stable  over  much  of the
 Eh-pH range for water,  and  formation  of  the  metal (which has  very  low
 solublity) may exert a  control on its nobility.

           Redox  potential has  an  indirect  effect on  silver speciation in
 that  it  dictates  the behavior  of  sulfur  ind  manganese species, which are
 important  controls  on  silver.   M.inganese 'dioxide apparently has sore
 affinity  for  silver  than most  of  the  other sorbenta  in aqueous systems.  If
 silver-containing manganese .dioxide Is deposited in  bed  sediments, which
 often exhibit reducing  conditions,  the Hn(lV) in Mn02 »ay be  reduced to
 Mn(II),  causing dissolution and release  of the silver.   Under such condi-
 tions ', however, Ag(I) would probably  be  reduced  co  insoluble  metallic
 silver,  or  it could  combine with  reduced sulfur  to  fora  the sulfide.  This
 compound  has  the  lowest aqueous solubility of a'ny of the silver compounds
 (Ks?,10~50) (Cotto.  and Wilkinson 1972).

           Solutions  of  silver  salts may  undergo  alkaline hydrolysis  to form
 the insoluble oxid»  via the following reaction (Cotton and  Wilkinson 1972);
               log  K  - -7.42  (25°C,  3M,  NaC104)

In strongly alkaline media, argentious oxide may be hydrolyzed to
Ag(QH)2~, which is soluble.  Dissolved carbon dioxide reacts with
A§20 to form argentious carbonate (Cotton and Wilkinson 1972),  The solu-
bility of the oxide and carbonates are sufficient to suggest that neither of
these compounds is act efficient control on dissolved silver,, except
possibly at high pH.

          .Silver halides are quite insoluble.  Hem (1970) speculates that
silver chloride may exert a major control on silver where chloride concen-
tration exceeds 10"^ M (35 ag/1).  In aquatic environments it has been
suggested that the silver complexes with chloride, bromide and iodide ions
control the amount of free silver present, subject to the effects of phys-
ical parameters and other an ions and cations.

           anne _«t _§!,« (1978) developed a computerized chemical model which
allows the prediction of chemical species at equilibrium in abiotic river-
ine and estuarine environments.  This model Indicates that at normal silver
concentrations, AgHS is an important species (Figure 17-1),  In the river
end-member, the activity of AgHS exceeds that of Ag* and AgCl by as much
as 10 fold.  In the marine end-member, only the activity of AgCl2~
exceeds that of AgHS, and Ag* is a trivial quantity.  AgBr, Ag(HS)2~» ,
AgF, AfOH, Agl, AgNC>3, Ag(N02)2~ and AgSQ4~ are all of minor
importance.  The low activity of silver in natural waters would seem to
preclude any solubility control on silver,  Krauskopf (1956) noted that
silver levels in marine waters were below saturation with probable solid
forms «uch as AgCl.

          So information was found *n the reviewed literature on silver-
organic interactions.  Although much is known about trace metal-organic
interaction, it would not be valid to attempt direct extrapolation of that
data to »ilvtr behavior in aquatic environments.  It can be Inferred, how-
ever, that only a minor fraction of aquatic silver will be in the hydra ted,
cationic fora; and chemical speciatlon trill have great effect on the trans-
port, toxicity and bioavailability of silver in the aquatic environment.

     17.3.3  Volatilization

          Volatilization of silver compounds is probably not an important
process In determining aquatic fate.  Formation of volatile biogenic alkyl-
ated compounds of silver probably does not occur, inasmuch as these com-
pounds are unstable in environmental conditions (Cotton and Wilkinson

     17.3.4  Sorption

          Sorptlon appears to bs the dominant process leadi j to partition-
ing of silver into the sediments.  Although there is little data on

                                         • 01
                                                  803 808
Figure 17-1  Variation tn activity of silver  solute  species  with salinity
             gradient from low conductivity river water  to  typical  marine
             water.  From Jenne et al. (1973),

sorpcion of silver, it appears Chat manganese dioxide has a strong affinity
for the metal, followed by ferric hydroxide and clay minerals.   Kharkar et
al« (1968) studied the sorption of several metals on three clays (mont-
morillonite, illite, and kaolin! te) and three hydrous metal oxides
(Fe2<)3, Mn02, and freshly precipitated F«(OH)3).  Various
quantities of these solids were added to a solution of silver (1 ug/1).
Th Fe{OH)3 > montmcrillonlte > 11 lite > kaolinitft >
Fe2<>3,  Additionally, the experiments were designed to determine the
extent to which adsorbed metals would be desorbed on contact with sea
Mater.  Almost all of the silver sorbed to Mn02 was released on contact
with sea water, and significant amounts were released by the other solids
as well*  Thus, much of the silver transported in 'the particulate phai* of
the water column jay be desorbed in the estuarin'e or marine environment.
The importance of precipitation of silver chloride under these  conditions
la uncertain.  The fact that silver did not have a strong affinity for the
more common solids (the clays), led Kharkar et _*!. (1968) to estimate that
90 percent of the stream supply of silver to the oceans is in the dissolved
fora, with desorption from suspended solids making up the remaining 10
          Th« Importance of Hn02 as a control on silver mobility was con-
f iraed by Chao and Anderson (1974).  Sediments from Colorado streams showed
a very high correlation (r » 0.913) between silver and manganese content.
Although sorption by the much more common iron oxides also controls mobil-
ity, iron plays a secondary role to manganese.  This vork confirmed the
earlier laboratory study by Anderson et al. (1973) why found that all solid
forms of manganese oxide sorbed significant quantities of silver*  This
sorption was adequately described by the Langnuir equation and showed a
direct relationship to pH values.

          Dvck (1%8) observed , in his study of silver adsorption on hyd-
rous ferric oxide, that this sorption of silver fit Freundlich adsorption
isotherms and was explainable in term* of hydrogen-site ion exchange.  Ad-
sorption was essentially complete in less than 5 ainutes and was strongly
dependent on pH.  These results would suggest that whenever manganese and/
or Iron oxides are present, silver will be almost Immediately adsorbed and
will be relatively immobile in the water column.

          Organic materials also adsorb silver.  Freeman (1977) found that
silver concentration in sediments of an alpine lake and nearby soils was
highly correlated with organic content.  High silver content also was asso-
ciated with the finer fractions of sediment rather than coarse fractions.
The silver concentrations in sediments were about 1000 times the concen-
tration of overlying waters.

          Tne relationship between silver cooplexation by humic substances
and,adsorption is not clearly understood.  If silver is conplexed by these
substances, as are several of the other metal a, its complexes would prob-
ably have a greater affinity for mineral surfaces than the free'Ag4* ion.
Further study Is required on sorption of silver before these phenomena can
be adequately assessed,

     17.3.5'  B ioaccumula t ion

          Several studies have shown that silver Is accumulated by aquatic
organisms.  Coleaan and Cearley (1974) demonstrated bioaccumulation of sil-
ver by largemouth bass and bluegill.  Concentration in bass was highest in
the internal organs, followed by the gills and remainder of the body.  Bio-
concentation ratios calculated from these data range from 10-100 (Table
17-1).  Cearley (1971) found that juvenile largemouth bass and bluegill,
exposed to 0.01 or 0.0001 tsg Ag/1 for six months accumulated silver for two
months after which the pollutant levels appeared to become static. Internal
organ* contained more silver than did edible auacle tissue.

          Activated sludge organisms also bioaccumulate silver.  Chlu
(1973) found that a suspension of 3 ag/ml of microbes acclimated to heavy
metals could extract almost 20 percent of dissolved silver present at 100
mg/1.  Expressed as a concentration factor, the microbes had a silver con-
centration about 100 times that of the solution.  In the ouch more dilute
silver concentrations normally found in the environment, activated sludge
organisms may exhibit concentration factors on the order of 1(P-10 .

          freeman (1977) studied the ecological kinetics of silver in an
alpine lake, and found that most silver Is partitioned to the sediments,
with the biota serving as a minor reservoir.  Plankton ccwponents of the
ecosystem showed fluctuations in silver concentrations closely correlated
to changes in lake water concentration, while benthic species showed fluc-
tuations more closely correlated to concentrations in the sediments.

          Terhaar e_t a 1. (1972) reported that silver complexes, as they oc-
cur in photographic processing effluent, are at most only slightly toxic to
fish.  Algae, daphina, freshwater mussels and fathead cinnows are all cap-
able of accumulating silver from water;  but the food chain was not an
important route of silver accumulation for animals at the higher trophic
levels (terhaar «t _•!• 1977). • Luooza and Jenne (1977) found that the accum-
ulation of silver by a deposit-feeding clam was dependent upon the physico-
chemical nature of the silver-sediment association.  Silver bioaccumulation
froo both biogenic carbonates and synthetic calcites was greater than sil-
ver uptake from iron oxides or detrital organic compounds.

                                Table 17-1

                   Itoconcentration Factors for Silver

    Taxon                       Factor8Reference

Freshwater plant*                  200           Chapman jet al.  1968

Freshvattr invertebrate*         3,080           Chapoan ej: al.  1968

Fre*hwater fi*h                  3,080           Chapman e_t al.  1968

Marina plane*                      100           Chapman e_t ajl.  1968

Marina invartabrataa             3,330           Chapman «_t al.  1968

Marine fish                      3,330           Chapman «t al.  1968
a.  Bioconcantration factors art the ratio derived  from  the concentration
    of the eleaent In the aquatic organism (In ppa  wet weight) divided  by
    the concentration of the element in water (la ppm).

          It Mould appear, therefore, that silver bioaccu&ulation is pri-
marily a function of sorption/desorption from sediments.   Silver is not
present in aquatic animals at very high concentrations because most of its
compounds are sparingly soluble in water.   Although silver is one of the
metals most toxic to aquatic life, there seeas to be little food-chain aag-
nifieatian, and silver appears to accumulate mainly in the internal organs.
Moreover, silver has a very,shore biological half-lite.

     17.3.6  B i o t r a n s f oraa cion

          No data indicating chat biota affect _thft fate of silver were
found.  Biotransfonaation of silver in sediments to methylated forms is ex-
tremely unlikely due Co the instability of alkyl silver compounds.  Methyl
silver (CH^Ag) is not stable ac temperatures above -30°C (Cotton and
Wilkinson 1972).

1^.4  Data Summary

     Although little data were found in the reviewed literature, it appears
that the main control on silver mobility is sorption by manganese dioxide,
clays, ferric hydroxide, and organic materials.   The available data have
been reviewed by Smith «nd Carson (197?>.   Bioaccumula'ion removes some of
silver ircra solution and is apparently strongly related to habitat (water
column, benthos) and distribution of biota.  Table 17-2 summarizes Che
aquatic fate described above.

En v i r onsen t al
                                Table 17-2

                     Summary of Aquatic'Fate of Silver
Chemical Sp«ciationa

Probably not important
determining fate.
Confidence of

Chloride, bromide and iodide ions
control the levels of hydrated
silver cations.  Crystalline,
metallic silver and silver sulftdes
may precipitate under reducing con-

Not an important fate.

Silver is strongly sorbed by
hydrous manganese and iron oxides,
clay minerals and organics.  A
major controlling mechanism in
determining the fate of silver in
the aquatic environment.

Numerous plants and primary consumer
organisms accumulate silver.  Little
evidence to suggest bionagnification,

Probably not an important fate.

    All of the noted environmental processes are important;  however, their
    relative importance with respect to each other is uncertain for
    detaraining final fate.

17, 5  Literature Cited

Anderson, B.J, ,  E.A. Jenne, and T.T. Chao,  1973.  The sorption of silver
  by poorly crystallized manganese oxides.  Geochim. Cosiochim.  Acta,

Boyle, R.W.  1968,  The geochemistry of silver and its deposits.  Geologi-
  cal Survey of  Canada, Bull.  160, Ottawa, Canada.  264p.

Cearley, J.E.  1971.  Toxieity and bioconcentration of caamium, chromium,
 1 and silver in Hicropteria salmoidas and Lepomuis macrocluiers.  Ph.D.
  Thesis, University of Oklahoma, Norman, Qkla.  84 p.  (Abstract onlv)
 . Oiss. Abstr. 328:5281(1972),

Qiao, T.T.  and B.J. Anderson, 1974.  The scavenging of silver by manganese
  and iron oxides in stream sediments collected from two drainage areas of
  Colorado.  Chem. Geol.  14:159-166.

Chapman, W.H. , R. L. Fisher, and M.W. Pratt.  1968,  Concentration factors
  of chemical elements in edible aquatic organisms,  Lawrence Radiation
  Laboratory, Liveraore, Calif.  UCRL-50564.  46p.

Chiu, Y,  1973.   Recovery of heavy metals by microbes,  pp. 18-20,  Ph.D.
  Thesis, School of Engineering, University of Western Ontario, London,
 . Ontario.   136p.

Coleman, R.L., and J.E. Cearley.  1974.  Silver toxicity and accumulation in
  largemouth bass and bluegill.  Bull. Environ. Contain. Toxicol .
Cotton, F.A. and G. Wilkinson.  1972.  Advanced inorganic chemistry.
  pp. 1044-1052,  Interscience Publishers, S.Y.  1145p.

Dyck, W,  196S.  Adsorption and coprecipitation of silver on hydrous ferric
  oxides.  Can. J.  Chem. 46:1441-1444.

Freeman, R.A,  1977.  The ecological kinetics of silver in an alpine lake
  ecosystem.  Second ASTM Symposium on aquatic toxicology, Oct. 31 to Nov.
  1,1977.  Cleveland, Ohio. {Preprint only).

Hem, J.D.  1970.  Study and interpretation of the chemical characteristics
  of natural waters,  pp. 202-203.  U.S Geological Survey Water-Supply
  ?ap«r 1473,  Washington, D.C.  363p.

Jenne, E.A. ,  D,C. Clrvin, J.W. Ball, and J.M. Birchard.  1978.  Inorganic
  speciation of silver in natural water - fresh to marine,  pp. 41-63.  jln
  Environmental impacts of artificial ice nucleating agent.  D.A. Klein
  (ed.).  Dowden, Hutchlraaon, and Rosa, Inc.  Strdudsburg, Peim.  256p.

Kharkar, D.P. , K.IC. Turekian, and K.K. Bertine.  1968.  Stream supply of
  dissolved silver, molybdenum, antimony, leienium, chromium, cobalt,
  rubidium, and cesiua to the oceans.  Geochlta. Cosraochim. Acta.

Krauskop€ , K. B,  1956.  Factors controlling the concentration of thirteen
  rare metals in sea water.  Geochim. Cosmochim. Acta.  9:1-32.

Luona, S.N. and E.A Jenne.  1977.  The availability of sediment bound
  cobalt, silver, and zinc to a deposit-feeding clam.  pp. 213-230. _in
  Biological implications of metals in the environment.  H. Drucher and
  R. E. Wildung (eds.).  Haaford Life Sciences Symposium, Richland, Wash,
  NT1S CONF ,750929. 682p.

Smith, I.C. and 8.L. Ca'rson.  1977.  Trace metals in the environment.  Vol.
  II.  Silver.  Ann Arbor Science ?ublishers( Inc., Ann Arbor, Michigan.
  490p.                              '

Terhaar, C,J», W.S. Ewell, S.P. Oziuba, and D.W. Fassett.  1972.  Toxicity
  of photographic processing chemicals to fish.  Photograph. Sci.  Eng,

Terhaar, t.J., W.S. Swell, S.P. Dziuba, W.«. White and P.J. Murphy,  1977.
  A laboratory model for evaluating the behavior of heavy metals in an
  aquatic environment.  Water Res,  11{1) :101-11Q.     •
       R.C, (ed.).  1377. CRC handbook of chemistry and physics.  58th ed
  CSC Press, Cleveland, Ohio.  2398p.

                               18.   THALLIUM

18.1  S tateBent of Probable Fate

    The behavior of thallium in natural waters is not well described in the
literature.  Thallium can be removed from solution by adsorption onto clay
minerals, bioaccumulatlon, or (in reducing environments) precipitation of
the s ul fide.  Most of the ligands common to aerobic waters form soluble
salts with thallium, so that precipitation is not important under oxic con-

18.2  Identification - Geochemistry of Thallium

    Thallium, a heavy metal, is a member of the Group III elements.  It Is
not used extensively by industry and is introduced into the environment
primarily as waste from production of other metals (Zitko 1975).  Thallium
exhibits the properties of both a lithophylic and chalcophylic element.  As
A lithophylic element, it is concentrated late during the magmatie crystal-
lisation of potassium minerals such as feldspars and micas. As a chalco-
phylic element, it is found as a component of Independent minerals and as a
substitute element in minerals such as galena.  The average concentration
of thallium in the earth's crust is about 1 ppm (Zitko 1975; Magorian e_t
al. 1974),  The following minerals of thallium, although rare, have been
found to occur in nature;  crookesite (Cu,Tl,Ag)2Se; hutchltisonite
PbS(Tl,Ag)2S'2As2S3; verbtlte Tl3S-3(Aa .Sb^Sj; lorandrite
           and avicennite
    Thallium, atomic number 82, atomic weight 204.87, exists in the ele-
mental form as a bluish white metal (Weast 1977).  In compounds, it has a
valence rf +1 or +3.  The +3 state is ouch less stable ia water than is the
+1 state (Cotton and Wilkinson 1972).  Thallium (III) forms some organo-
aetalllc compounds; however, T1(I) forms relatively few complexes with the
exception of those with halogen, oxygen, and sulfur ligands (Cotton and
Wilkinson 1972).  In this and several other respects, T1(I) has chemical
properties similar to those of the alkali metal cations.

    The behavior of thallium in the environment can be somewhat explained
by a comparison of the ionic and atomic radii of thallium with the radii of
some associated elements in Table 18-1.  (Numerical values- are from Weast'
1977.)  The similarity of the univalent ionic radius of thallium to those
of univalent potassium and rubidium explains the presence of thallium in
late state potash and plagloclase feldspars as well as its accompaniment
with rubidium with which it is concentrated during the late states of crys-
tallization in igneous rocks.  Likewise, the similarity of the metallic and
covalent radii of thallium and lead suggests that the behavior of thallium
in the aquatic environment will be similar to that of lead.

    The CAS number of thallium is 7440-28-0, and its fSL number is

                                Table 18-1
              Radii of Thallium and Some Associated Elements
 Ionic radii
Metallic radii
                      Covalent radii
T1(I)   1.49  |
Rb (I)  1.49  I
K(I)    1.33  *

T1UX1)  1.05   \\
Y(III)   1.05
In(III)  0.92
1.44  i
1.31  A

18.3  Summary of Fate Data

    18.3.1'  gh_ocoljfsis

         Although some thallium compounds such as T1C1 are photosensitive
(Cotton and Wilkinson 1972), there is no evidence that the photochemistry
of thalliuo plays a significant role in determining aquatic :fate.

    18 . 3 . 2  Chemical Speciatlon

         In reducing 'environments, thai HUBS .may be precipitated as , the
metal or, in the presence of sulfur, as the sulfide (Lee 1971; Magorian et
al . 1974.)  In very oxidizing waters, the most oxidized form of thallium,
11(111), may be present.  In other Eh ranges, thallium(I) has a very high
solubility compared to the other environmentally important heavy metals,

         O'Shea and Mancy (1978), in their study of trace oetai ions and
complexing agents at different pH values, found that increasing pH de-
creased thallium- inorganic interactions.  Increases in pH, however, pro-
duced extensive thallium-hufflic acid interaction.  It appears, therefore,
that thallium-organic interactions may be important in most natural water
systems.  Most coaaaercial uses of thallium are in the form of organometal-
lic compounda (pesticides, poisons, etc.).  Further research, however, is
required to determine the importance of these interactions on the aquatic
face of thallium.
         The thallic ion (Tl") is hydrolyzed to fora the colloidal oxide
over the pH range of natural waters (Cotton and Wilkinson 1972).  Depending
on the relative kinetics of reduction compared to hydrolysis, however,   ,
precipitation as ,11(0^)3 may be an effective mechanism for removing
thallium from solution via the following scenario: .thallium (III) precipi-
tates as the oxide or hydroxide and settles to the bed sediments; in the
sediments, the reducing conditions cause reduction to Tl(I)j, which reacts
with sulfide to fora the insoluble compound lljS  (Lee 1971).

    18.3.3  Volatilization

         No evidence was found for formation of volatile thallium compounds
in the environment.

    i8.3.4_  Sprption  •  ' .

         Thallium is strongly adsorbed by montraorillonitic clays.  Magorian
*_t al. (1974) demonstrated that a 1 gm/1 suspension of the clay hectorite
could remove 972 of a 100 \lg/l concentration of thallium within 24 hours.
Similarly, a 1 mg/1 concentration of thallium was reduced to 115 yg/1, and
a 10 yg/1 solution was reduced to below 1 ug/1.  The above values are for

pft 8.1; adsorption is not as effective at pH 4.0.  Experiments with copper
and zinc ions showed that thallium is not adsorbed as strongly as these
metals. Thallium probably has an affinity for hydfQus metal oxides.  Zitko
ejt a_l. (1975) observed a 28% decrease in thallium concentration after pre-
cipitation of heavy metal hydroxides at pH 9.6 and subsequent centrifuga-

         Mathis and Keverti (1975), in a study of heavy metal cycling In a
lake in southwestern Michigan, were able ?o detect thallium only in th*
sediments.  They suggested that their inability to detect thallium in the
water, plankton, and fish may have been due to the sensitivity of their
analytical Methods; however, the levels of thallium in the sediment were an
order of magnitude higher than the minimum sensitivity of the atomic
adsorption spectrophotomecer that was used suggesting that the sediment is
an active sink for thallium. This observation, combined with the high
solubility of most thallium compounds, implies that thallium is actively
being sorbed into the sediments in the aquatic environment.

    18.3.5  jioaceuaulation

         Since thallium Is soluble in most aquatic systems, It is readily
available to aquatic organisms.  Thus, it is not surprising to find that
thallium is quickly bioaccumulated.  Goldfish have a higher rate of uptake
for thallium than for the five most common alkali metals (Zitko 1975).  The
a*Sa' JO-j^hgix Jt£* > was a°^e Co concentrate thallium by a factor of 127 to
220 within one hour;  in comparison, the concentration factors fot 2.7
hours exposure were 114 for lead, 30 for cadmium, 80 for zinc, and 313 for
copper (Magorian e_t _a_l. 1974).  Other bioconcentration factors,are sum-
marized in Table 18-2.

         Overnell (1975), in a study of the effect of thallium and other
heavy metals on photosynthesis in freshwater algae, found that thallium
decreased photosynthesis by 40 percent.  Unfortunately, he did not measure
thallium accumulation by the alga.  He did report, however, that the inhi-
bition of photosynthesis by thallium could not b« relieved by washing of
the alga, suggesting that the thallium may have been irreversibly bound
and, therefore, was available not only for bioaccumulation but food chain
magnification as well.

    18*3.6  8iotrans|orrna t ion

         Although the toxic effects of thallium are well known (Zitko
1975), no evidence was found that blotransformation of thallium compounds
plays an important rcle in determining aquatic fate.  There has been some
speculation that thallium could be methylated under aerobic conditions by
electrophilic attack (Anon., 1977), but this process has not been demon-
strated in the aquatic environment.

                                Table  18-2
                   Bioconcentration Factors  for  Thallium
Bioconeentration Factor3
Freshwater Plants            100,000
Freshwater Invertebrates     150,000
Freshwater Fish              100,000
Marine Plants                100,000
Marine Invertebrates         150,000
Marine Fish                  100,000
Claa (Mya arenia)            17.6-18.6
Hu'ssel (Mytilus edulis)      10.9-12.4
Atlantic saloon              27-1430
                             Chapman eg al.  1968
                             Chapman et_ al.  1968
                             Chapman _e_£ al.  1968
                             Chapman rt ail.  1968
                             Chaptaan _et al.  1968
                             Chapman eJL al.  1968
                             Zitko ejt Bl_,  1975
                             Zicko and Carson 1975
                             Zitko ec al.  1975
a.  Biocoicentration factors are  the  ratio derived  from  the  concentration
    of the eleaient in the aquatic organism (in  ppm  wet weight)  divided by
    the concentration of the element  1.. water (in ppm).

18.4  Data Summary

    In aerobic waters, adsorption and bioaccumulation are the primary
mechanisms for removal of thallium from solution.  In anaerobic environ-
ments, precipitation as the sulfide may be important.  Much of the thallium
Introduced into freshwater svsttms probably remains in solution and is
transported to the oceans.  The aquatic fate information for thallium is
summarized in Table 18-3.






Blot rana focma t ion
            table 18-3   ,

Suiasary of Aquatic Fate of Thallium '


    Not an important mechanists.

    In reducing environments,    '
    T1(I) may precipitate as a
    sulfide; otherwise, ic will
    remain in solution.

    Not an important mechanism.

    Thallium is adsorbed to clay
    minerals and hydrous metal
    oxides.  Probably a very
    important process.

    Thallium is accumulated by
    aquatic organisms.  Probably
    an important process.

    Not an important process.
Confidence of


    Medj am



a.  All of the noted envirbnoental processes are important; however, their
    relative importance with respect to each other is uncertain for
    determining final fate.

18.5  Literature Cited

Anon.  1977.  Tracking trace metals in the biosphere.  Chens. Eng. News

Chapoan, W.H., H.L. Fisher, and M.W. Pratt.  1968,  Concentration factors
  of chemical elements in edible aquatic organisms.  Lawrence Radiation
  Laboratory, Liveraore, Calif.  UCRL-50564.  46p.
                   S                     i
Cotton, F.A. and G. Wilkinson.  1972.  Advanced inorganic chemistry
  pp.453-455.!  Interscience Publishers, New York.  1145p.

Lee, A.G.  1971,  The chemistry of thallium,  pp.1-18, 295-318.  Elsevier
  Publishing Company, Amsterdam.  336p.

Magorian, T.R., K.G. Wood, J.G. Michalovlc, S.L. Pek, and M.M. Van Lier.
  1974.  Water pollution by thalliua and related metals pp.145-160.  MTIS
  PB 253 333. Springfield, Va. 182p.

Mathis, B.J. and N.R. Kevern.  1975.  Distribution of mercury, cadmium,
  lead, and thallluo in an eutrophic lake.  Hydrobiologia 46(2-3);207-222.

O'Shea, I.A. and K.H. Mancy.  1978.  The effect of pH and hardness metal
  ions in the competitive interaction between trace metal ions and
  inorganic and organic conplexirtg agents found in natural waters.  Water
  1*9. 12:703-711.

Overnell, J.  1975.  The effect of son* heavy metal ions of photosynthesis
  in a freshwater alga.  Festic. Biochem. Physiol. 5:19-26.

Weast, R.C. (ed.)  1977.  Handbook of chemistry and physics, 58th Ed. C1C
  Press, Cleveland, Ohio 2398p.

Zitko, V.  1975.  Toxicity and pollution potential of thallium,   Sci.
  Total Environ. 4:185-192.

Zitko, V. and W.V. Carson.  1975.  Accumulation of thallium in clams and
  nusseis.,  Bull. Environ* Contaa. foxicol. 14:530-534.

Zitko, V., W.V. Carson, and tf.G. Carson,  1975,  Thallium: occurrence in
  the environment and toxicity to fish.  Bull* Environ. Contain. Toxicol.

                                 19.   ZINC

19.1  Statementof Probable Fate

    Moat of Che zinc introduced into the aquatic environment is partitioned
into the sediments by sorption onto hydrous iron and manganese oxides, clay
minerals, and orgaiic materials.  Precipitation of the sulfide is an impor-
tant central an the nobility of zinc in reducing environments, and precipi-
tation of the hydroxide, carbonate, or basic sulfate can occur where 2inc
is present in high concentrations.  An essential trace element in nutri-
tion, 'zinc is oioaccumulated in all organisms.  Formation of complexes with
organic and inorganic ligands can increase the solubility of zinc and prob-
ably increases the tendency for zinc to be adsorbed.

19.2  Idantificat^ioa^- Geochemistry_of Zinc

    Zinc is a naturally occuring element found in the earth's crust £t an ,
average concentration of 123 pptn (Weast 1977).   Zinc is found chiefly as
the minerals sphalerite (ZnS), smithsonite (211003), willenite
(Z^SiG^ and zincite (ZnO); it also substitutes to some extent for
magnesium in silicate minerals and is found in most igneous rocks (Cotton
and Wilkinson 1972).

    Zinc (Zn) is a metallic element, atomic number 30, atonic weight 65.38
(Weast 1977).  The chemistry of zinc is similar to that of cadmium, whicB
is directly below it in the periodic table (Cotton and Wilkinson 1972).  In-
aqueous solution, zinc always has a valence of -s-2, and it exhibits ampho-
teric properties, dissolving in acids to form hydrated Zn(II) cations and
in strong bases to form zincate anions (probably Zn(OH)4~2). Compounds
of zinc with the common ligands of surface waters are soluble in neutral
and acidic solutions, so that zinc is readily transported in most natural
waters and is one of the most mobile of th* heavy metals.  The geochemistry
of zinc in surface water has been extensively reviewed by Hem (1972). Since
the divalent zinc ion does substitute to some extent for magnesium in the
silicate minerals of igneous rocks, weathering  of this zinc-containing
bedrock gives rise to Zn"*"2 in solution whereupon the hydrated cation re-
mains dominant to pH values of about 9.  Zinc forms complexes with a
variety .of organic and inorgaaic ligands, but these compounds are suffi-
ciently soluble to prevent their becoming a limiting factor for the solu-
bility of the small concentrations of zinc found in most aquatic environ-
ments.  Adsorption on clay minerals, hydrous oxides, and organic matter is
a more probable limiting mechanism.

    The CAS nunber for zinc is 7440-66-6; its fSL number is B823-4379,

19,3  Sumna ry	of Fat e • Data

    19*3«i  Photolysis

         Mo evidence was found that photolysis of zinc compounds in the
aquatic environment significantly affects its fate.

    19.3.2  gh««ieg1_ JS pec1 a t ion

         lit natural waters, zinc can be found is several chemical forms,
for example, as simple hydrated ions, as metal-inorganic complexes, or as
metal-organic complexes.  An understanding of the chemical speciation of
zinc in any given situation is based upon theoretical calculations of hy-
drolysis, oxidation/reduction and organic complication.  A short presenta-
tion of this material will be given alter which the chemical speciation of
zinc in various aquatic environments will be discussed.

         Zinc always has an oxidation state of +2 in aqueous systems
(Cotton and Hllkinson 1972).  Unlike the transition metals adjacent to it
in the periodic table» zinc is not directly affected by changes in Eh (re-
dox potential);  however, the valence* and reactivity of ligands reacting
with zinc are affected by Eh.  Figure 19-1 shows the equilibrium solubility
and stable solid species of zinc at pH 5 as a function of pC (Huang et al.
1977).  At pri levels sore characteristic of natural waters, the solubility
of the oxide and carbonate are considerably lower.  Although the graph
shows a very low solubility for Zn+2 in oxidizing conditions, measured
zinc concentrations are usually ouch higher, indicating that formation of
the basic sulfate (Zn4(OH)oSQ4) Is not normally an effective control
on the aobility of zinc.

         Precipitation of zinc compounds appears to be Important only in
reducing or highly polluted environments.  Holmes et _al. (1974) concluded
that formation of zinc sulfide controls the mobility of zinc In Corpus
Christ! Harbor (an estuarine system).  Seasonal fluctuations in dissolved
zinc levels were attributed to variations in redox potential:  in the sum-
mer, reducing conditions prevail in the hypollamion due to combined effects
of lower 02 solubility, greater biological oxygen demand, and thermal
stratification; and zinc, consequently, is Incorporated into the sediments
via formation of ZnS*  In the winter, Eh increases, and zinc is released to
Che water column; Ic is then adsorbed by suspended solids and is trans-
ported to Corpus Christ! Bay, where the solids settle out and zinc la again
partitioned Into the sediments.

         Hydrated zinc cations may be hydrolyzed to form 2n(OH>2(a) «*
ZnO(s) (Scutum and Morgan 1970).  Patterson et al. (1977) studied the pre-
etpitatii  , of Zn(OH)2 and ZnC03 and found that 2n(OH)2 precipitation


-10         -1S
 I	I
is kinetically favored over ZnCQ% precipitation.  Figure 19-2 straws the
computed equilibrium solubility of 2n(OH)2 as a function of pH; and
figure 19-3 shows the stability of Zn(OH>2 and ZnCQj in a iO'1'^ M
Carbonate system !as a function of pH.  When Patterson *£ £l» (1977) added
zinc to systems with and without carbonate, they found thac the resulting
solubilities conformed to those expected if zinc hydroxide alone were con-
trolling solubility (Figure 19-4),  Apparently, the zinc, hydroxide and car-
bonate species did not reach thermodynamic equilibrium over the time-span
of that study.  An interesting observation made by the authors was that
precipitation was essentially complete within four hours, with little    ' •
further reduction in soluble zinc even after 264 hours (Figure 19-5).

         In an Impoundment polluted with zinc (400 Ug/1) introduced by
dumping of mine wastes, Weatheriey and Dawson (1973) found that zinc was
precipitated as an amorphous colloidal deposit of basic carbonates and
sulfates.  Under oxidizing conditions, precipitation of these zinc1
compounds is probably important only where high concentrations of zinc

         Long and Angino (1977) developed a theoretical model to study the
chemical speciation of zinc in aqueous solutions and the response of zinc
to variations in ionic strength and complexation.  Association of zinc with
the ligands OH~, Cl~, C0j~^, 50^, and HCC>3~ were con-
sidered at pH values from 3,5 to 11.0 at 25*C in differing seawat'er-fresh-
water mixtures.  The results are summarized in Figure 19-6.  In general,
the relative importance of the various ligand-zinc complexes can be pre-
dicted from a comparison of their stability constants.  This model, how-
ever, does not take into account metal-organic complexes, and it is, there-
fore, useful only in unpolluted relatively organic-free waters,
         Guy and Chakrabarti (1976), in thelf study of metal-organic inter-
actions in natural waters, found that humic acids in solution and other
natural complexing agents can maintain zinc ions in a bound form at a pH as
low as 3.  the release of zinc from sediments is, therefore, apparently
controlled by a combination of ion exchange and complex formation whereby
the stability of the octal-organic complex determines the amount of metal

         In summation, the transport of zinc in the aquatic environment is
controlled by the speciation of the ton.  Although it appears that in most
unpolluted waters, zinc will exist mainly as a divalent cation and be
easily adsorbed, organic material will have an overwhelming effect on the
chemical form in which zinc will b« present in polluted areas.  Unfortu-
nately, there is at present insufficient information to construct a compre-
hensive model for the transport and sorprion of zinc based upon chemical

                                 t   w   ta  14
Figure 19-2  Theoretical zinc hydroxide  solubility curve,
             Proa Patterson  et  al,  (1977).

             C 0

                             f. .I.,.,,, f mi. I in . I
                                     10  12   14
Figure 19-3  Theoretical  zinc  carbonate-zinc hydroxide phase
             diagram (C^  "  10~^*l«ol/l).   From Patterson
             et al. (1977).

              til  0
                 •*• •


                                          12   14
Figura 19-4  Cooparison of  zinc  hydroxide and zinc carbonate
             data with theoretical  hydroxide solubility curve,
             From Patter$on et al.  (1977).

24  4fl  n  M

                                             t«2 I1« 240 2«4
Figure  19-5  Kinetics of  zinc prtcipicati^n.   from Patcaraon

                                         100* FrtttMltr
                    5,5 45 55 6.5 75 15 » 5 105 115
Figure  19-6  Chemical  speciatton of  zinc  In s«awatcr-freshwater
              mixtures.   Proa Long and  Anglno (1977),

    19.3.3  Volatilization

         No evidence was found that volatilization of zinc is an  important
aquatic fate.  Alkyl zinc compounds are unstable to oxygen and water
(Cotton and Wilkinson 1972), and, therefore, volatile methylated  forms an-
alogous to those of mercury, arsenic, antimony, lead, and selenium are
probably not formed in aquatic environments.

    19.3.4  Sorftion

         Sorptioa of zinc by hydrous metal oxides, clay minerals, and or-
ganic materials is probably the dominant fate of zinc In the aquatic en-
vironment.  Concentrations of zinc in suspended and bed sediments .always
exceed concentrations in ambient waters (Nelson and Hauschild 1970; Hem
1972-, Angino et al. 1974- Kubota et al. 1974; Perhac I974a, 1974b; Rehwoldt
•t _al. 1975, Carpenter et al. 1975; Helz et al. 1975; Pita and Hyne» 1975;
fteele and Wagner 1975; Namminga and 4ilhm 1977) and there is an  inverse
correlation between zinc concentration and sediment grain size (Nelson and
Hautchild 1970; Perhac 1974b; Pita and Hyne 1975; Steele and Wagner 1975;
Namminga ard Wilha 1977).  Since smaller grained particles have a higher
surfsee-to-mass ratio, the fact that higher zinc concentrations are asso-
ciated with such particles implies that sorptlon is responsible for, this
phenomenon rather than precipitation.  James and HacNaughton (1977) pre-
sented theoretical models which show that the removal of zinc from aque-
ous solutions can be explained by colloidal and surface chemical  controls,
wherein the presence of insoluble phases, often with high surface areas,
provide sites for adsorption or interfacial reactions.  Their calculations,
baaed on electrical double-layer and ion-exchange models, developed adsorp-
tion isotherms for various pH values, metal ion concentrations, ionic
strengths and mineral substrates.  These models have provided a valid theo-
retical background from which to approach the adsorption of zinc  on inor-
ganic minerals, but the models have not been extended at present  to adsorp-
tion on organic materials.

         Guy jet al. (1975) developed a chemical model to Investigate the
mechanisms controlling the distribution of zinc between soluble and parti-
culate fractions in natural waters*  Using clays, hydrous manganese oxides,
and organic material, they found that zinc sorption onto clays followed a
Freundlich isotherm but a Langmuir isotherm was required for the  other aa-
tcrials.  Guy e_t al. (1975) interpreted these results as indicating that
these zinc distributional ratios can be explained in terms of coordination
chemistry.  While the model presented «ras in qualitative agreement with re-
ported phenomena in natural waters, further studies are needed to provide
more definitive conclusions regarding the predictive possibilities of these
models for the aquatic fate of zinc.

         the composition of the dissolved and' suspended solids load has an
important effect on the mode of transport of zinc.  Where the solids are
primarily dissolved, most of the zinc is transported in solution as the
hydrated cation or complex species (Perhac 1972, 1974a; DeGroot and Allers-
tna 19755.  Where suspended .solii-? make up a high proportion of the total
solids load, most of the zinc transported will be sorbed to the suspended
and colloidal particles (Kubota ee al. 1974; Steele and Wagner 1975).!  A
common observation Is that residence in impoundments reduces the concentra-
tion of dissolved zinc, apparently due to scavenging by suspended solids
and subsequent deposition (Pita and Hyne 1975; Perhac 1974b; Nelson and
Hauschild 1970; Kubota ££ al. 1974).

         Coprecipi tation and sorption of dissolved zinc, by hydrous oxides
of iron and manganese are important controls on the mobility of zinc,
especially where high concentrations of reduced iron and manganese are
introduced into aerobic surface waters (Lee 1975). , As reduced iron and
manganese are oxidized, their hydrous oxides precipitate as coatings or dis-
crete particles.  The negative zeta-potentlal usually exhibited by these
materials attracts zinc and other cations,,and 'the sorted cations are in-
corporated into the crystal lattice structure of the hydrous iron or man-
ganese oxide (Lee 1975).  The slack encrustation often found on submerged
rocks is usually composed of thesse oxides.  Suspended solids can be coated
with these oxides as well, and Angino e£ al_. (1974) found a significant cor-
relation between zinc concentration and manganese and iron concentration of
suspended solids in Kansas streams.  On a mass-per-oass basis, zinc is par-
titioned sore strongly in hydrous Fe-Mn oxides.than in other components of
the sediment.  This has led Carpenter _e_t al. (1975) to suggest that analy-
sis of zinc and other aetals in oxide coatings may be a useful tool in geo-
chetnlcal prospecting.                       •                            '

         Colloidal and suspended organic matter also adsorbs zinc.  Rashid
(1974) reported that about: 26.1 mg of zinc was sorbed per gram of sedimen-
tary organic matter added to a solution of zinc.  Zinc was sorbed more
strongly than Si, Co, 'and J*n, but less strongly than copper.  Jackson and
Ekippen (1978) found that the presence of organic ligands increased the
solubility of zinc in the presence of clays.  The organic acids proved
capable of reraobilizing zinc from solid phases although the reaction is
klneticaliy hindered, especially at basic pH values.

         The tendency of zinc to be sorbed is affected not only by the na-
ture1 and concentration of the sorbent but by pH and salinity as well.  In a
study of heavy metal adsorption by two oxides and two soils, zinc was com-
pletely removed from solution when pH exceeded 7;  below pH 6, little or no
zinc was adsorbed, as shown in Figure 19-7 (Huang ej: a_l. 1977). Addition of
inorganic complexing ligands enhanced the affinity for adsorption (Huang rt
al. 1977).

                             3   4   S   8   7   8   9
Figure 19-7  Adsorption of zinc on various solids after 24 hours.
             Ionic Strength * 0,1 M; solid suspension - 5 got/1;
             Original (2n**"j - 10~^M; therefore, 200 micromol/gai
             corresponds to 100% adsorption; from Huang et al.  (1977),

          Helz _e£ _ail. (1975)  found  Chat line  is desorbed  from  sediments as
 salinity  Increases.  This phenomenon, which  Is exhibited by nany  of  the
 other aetals as well,  is apparently due  to displacement  of the adsorbed
 zinc  ions.by alkali and alkaline earth cations which are abundant  in brack-
 ish and saline waters.  In  summary, sorption is  the dominant  fate  process
 affecting zinc, and it results*in  enrichment of  suspended and bed  sediments
.relative  co  the water  column.   Variables affecting  the mobility of zinc
 include the concentration and'composition of suspended and bed sediments,
 dissolved and particulate iron  and manganese concentrations,  pH,  Eh, salin-
 ity, concentration, of  complexity ligands, and the concentration of zinc.

    19.3.5 Bioaccumulation

          Zinc is  bloaccumulated by all organises.   One noteworthy aspect of
 bioaccumulation is thai tt  occurs  even in the absence of abnormally high
 zinc concentrations since it is an essential nutrient.   Bioconcentration
 factor's are  listed in  Table  19-1.

          Zinc has been extensively studied In the freshwater  environment.
 2ir.c-6S was accumulated much acre  raaoily than &0(k>,  ^~?Cs or 8*gr ^y
 soft  tissues of carp,  snails, tadpoles and clams during  experiments uti-
 lizing  isotopic measurements conducted in ponds  (Brungs  1967).   In labora-
 tory experiments, the  brown  bullhead  (Ictalurus) accumulated  "^Zti rapidly
 for the first seven hours of exposure  followed thereafter by  a reduced
 accumulation rate (joyner 1961).   Gill and viscera  attained the highest
 zinc  concentration of  the tissues  analyzed.  In one  set of experiments, the
 esophagus was plugged  oa  some fish to  determine  the fraction  of accumulated
 zinc attributable to swallowed  water;  this route of uptake was found to
 have  a  negligible contribution  under  the experimental conditions.  Zinc-
 exposed  fish that were then  transferred  to fresh water lost half  of their  •
 accumulated  zinc  after six  days followed thereafter by a much reduced rate
 of  zinc elimination.

          The importance of  bioaccumulation,  or at least  biologically
 mediated  removal  of zinc  from solution,  was  shown by  Adams e_t ail.  (1975),
 who demonstrated  that  when  dissolved  zinc was added td the influent of «
 wastewater treatment plant  at levels  of  2.5  to 20 mg/1,  primary treatment
 removed only about 8-14 percent of the zinc. After activated sludge treat-
 ment, however,  7i~96 percent of the  zinc was removed.  It Is  uncertain
 whether the  zinc  was bioaccutnulated  by the microorganisms, or if  further
 removal of solids by sludge  formation was responsible for the dramatic re-
 duction in zinc concentration.   Nevertheless, it is clear that in the bio-
 density ranges  found in sewage  treatment plants, zinc is effectively re-
 moved  from solution, and  bioaccumulation probably plays  an Important role
 in  such  removal.
                                        19-? 3


Freshwater Plants
             Table 19-1

Biaconcentration Factors for Zinc

Freshwater invertebrates
  Chironooid larvae

Freshwater fish

M«rin« Plant*
  Alga* - Nitzshia sp.
  Seaweeds        • ^

Marine invertebrates
  American oyater-
  Crasseotrea yirginlca

Marine fish
  Yellow fish tuna
  Skipjack tuna






Chapman _e_t al. 1968

Chapman _ejt al. 1968
Namminga and Wllhm 1977

Chapman a^t al. 1968

Chapman'_e_t al. 1968
Chlpmaa e£ ad. 1958
Stunm and Morgan 1970

Chapman e_t al. 1968

Chipoan e_£ al. 1958

Chapman £t al. 1968
Stumm and Morgan 1970
Stumm and Morgan 1970
Scumm and Morgan 1970
    Bioconcentration factors are Che ratio derived  from  the  concentration
    of the «lem*nt in Cha aquatic organism (in ppm  of  wet  weight)  divided
    by the concentration of the element  in water  (in ppm).

          Microcosm studies  generally indicate  that  zinc  is  not  bicwagnl-
 fied.   Patrick and.Loutit, (1975)  exposed  bacteria  to  elevated- levels  of
 zinc.   When the bacteria  were fed to tubificid worms,  the worag concen-
 trated  zinc,  but the  levels in the worms  were  lower than levels in the

          The  chemical form  in which zinc  occurs  has a  profound  effect on
 its  availability fo'r  bioaccumulation.   The  bottom-feeding clam  Macoma
 balthica accumulated  zinc much more readily from biogenic carbonates
 (crushed clamshells)  than from other sediment -bound sources (Luoma and
 Jenre 1977).   Zinc was readily accumulated  from  detrital organic materials
 as well, but  little uptake  was observed when zinc was  coprecipitated  with
 hydrous iron  or manganese oxides.  The  sinks from which  bioac cumulation was
 greatest also exhibited the greatest rate of sediment  to water  desorption.
          Duke  (1967)  studied  the  distribution  of  ^^>1n  in  an  estuarine  en-
 vironment.   He found  that one day after  addition  of  ^Zn  to  a modeJ  eco-
 system,  36  percent  was in the sediments,  5  percent was in biota,  and 59
 percent  was in the  water  column.   After  100 days, 99.4 percent  was in  the
 sediments,  0.6 percent was in biota,  and none  was in the  water.   Maximum
 k^Zn levels in biota  were attained within 2  days; aftsr 100  days,
 scallops contained  30 percent of  their maximum level,  oysters 60  percent,
 and  clams 25 percent.  Zinc-65 concentration was  smaller  in  fiah, crabs and
 marsh grass than in the molluscs.

          Zinc  is readily  .accumulated  by  both marine  and freshwater fish
 from both food and  water, but internal organs  and bones accumulate much
 higher zinc levels  than edible muscle tissue (Phillips and Russo  1978).
, The  time required  for fish to reach threshold  levels of zinc appears to be
 dependent upon species and the chemical  nature of the  environment.   Upon
 entering fish, some zinc  associates with cadmium- bind ing  protein, and  evi-
 dence suggests that a zinc-binding protein  may exist.   The toxicity of zinc
 to aquatic  organisms  has  been shown to decrease with increasing calcuim
 concentration  even  though calcium appears to stimulate zinc  uptake.  In
 marine fishes, cadmium reportedly decreases  zinc  accumulation (Phillips and
 Russo 1978).   In summary, it  seems .evident  that while  zinc is actively
 bioac cumula ted ,  the biota appear  to represent  a relatively minor  sink  when
 compared to the  sediments.

     19.3.6   Bio transformation '

          Zinc  is one  of the most  important  metals in biological systems,
 probably second  only  to iron  among the heavy metals  (Cotton  and Wilkinson
 1972); over 25 zinc-containing enzymes have  been  identified.

         A discussion of the aetabolic role of zinc can be found in a num-
ber of sources {e.g., Prasad (1967)).  Since alkyl-2ine compounds are un-
stable to water and oxygen, biomethylation of zinc in aquatic ecosystems
probably doe's not occur. ! Nevertheless, the presence of biogenic ligands,
such as humic acids, affects the precipitation and adsorption behavior of
zinc.  Biologically generated microenvironments can also alter the nobility
of zinc.  Since zinc is actively bioaccumulated, it would not be surprising
to find that it exhibits seasonal fluctuations in concentration such as
those documented for copper (Kimball 1973; Grimshaw _e_t _al, 1976; Namminga
awl Wilho 1977) » in which degradation of organic material in the fall and
winter results in elevated aqueous concentrations relative to the spring
arid summer months when the metal is actively bioaccumulated.
19,4  Data^ Sumaary

    The doninant fate of zinc in aerobic waters is sorption by hydrous iron
and nanganese oxides, clay minerals, and organic material.  The efficiency
of these materials in removing zinc from solution varies according to their
concentrations, pH, Eh, concentrations of ligands, and the concentration of
zinc.  Precipitation of the sulfide is an important control on the mobility
of zinc in reducing environments.  Under aerobic conditions, precipitation
of zinc compounds is probably important only where zinc is present in high
concentrations.  Zinc is bioaccuaiulated , *hich is to be expected in view of
the fact that it is an essential nutrient.  Although the biota appear to be
a ainor reservoir of zinc relative to the sediments, biological activity
can affect the mobility of zinc in the aquatic environment.  The aquatic
fate of zinc is summarized in Table 19-2.



Sp«ciationa  .

Ho trans format ion
                                Table 19-2

                      Summary of Aquatic Fate of Zinc
       Statement   ,  .

Not an important mechanism
Confidence of

In most unpolluted waters,
the majority of zinc will
exist as Che hydrated diva-
lent cation.  In polluted
waters, coraplexation will

Not an important mechanism.

Zinc has a strong affinity
for hydrous metal oxides,
clays, and organ!-, natter.
Adsorption increases with

Zinc is strongly bioaccumu-
lated,  Bioconcentration
factors range from 10' to

No bioaeehylation in evidence,
Organic ligands of biological
origin may affect solubility
and adsorption.

    All of the noted environmental processes are important; however, their
    relative -importance with respect to each other is uncertain for
    determining final fate.

19.5  Literature Cited

Mans, C.E. Jr., W.W. Eckenfelder, Jr. a?d B.I. Goodman.  1975.  The
  effects and removal of heavy metals in'biological treatment, pp.277-292.
  In Heavy octala in the aquatic environment.  P.A. Koenkei (ed.), Pergamon
  Press, Oxford, England.' 352p.

Angino, E.E., L.M. Magnuson, and T.C. Waugh.  1974,  Mineralogy of
  suspended sediment and concentration of Fe, Mn, Mi, Zn, Cu, and Pb in
  water and Fe, Mn, and Pb in suspended load of selected Kansas streams.
  Water Res. 10(6):1187-1191.

Brungs, W.A.  1967.  Distribution of cobalt-60, zinc-65, strontiun-85, and
  cesiuo-137 in a freshwater pond.  Publ. No. 999-RH-24, U.S. Public Health
  Service, Washington, D.C. 52p»

Carpenter, R.H., T.A. Pope, and R.L. Smith;  1975,  Fe-Mn oxide coatings in
  stream sediment geochemical surveys.  J. Geocheo. Explor. 4:349-363.

Chapman, W.H., H.L. Fisher, and M.W. Pratt.  1968.  Concentration factors
  of chemical elements in edible aquatic organisms.  Lawrence Radiatio'n
  Laboratory, Livermore, Calif.  UCRL-50564. 4t>p.

Chipinan, W.A.1, T.R. Rice, and T.J. Price.  1958.  Uptake and accumulation
 'of radioactive zinc by marine plankton, fish, and shellfish, pp.279-291.
  U.S. Fish and Wildlife Service Fishing Bulletin 133, Vol. 58.
  Washington, D.C. 459p.

Cotton, F.A. and A. Wilkinson.  1972^  Mvanced inorganic chemistry.
  pp.600-610. Interscience Publishers. New York. 1145p.

DeGroot, A.J. and E. Allerama.  1975.  Field observations' on the transport
  of heavy metals in sediments, pp. 85-95. _in Heavy metals in the aquatic
  envire-went.  P.A. Krenkel (ed.). Pergamon Press, Oxford, England. 352p,

Duke, T.W.  196'7.  Possible routes of zinc-65 from an experimental
  estuarine environment to oan*  J. Water Pollut. Control Fed.
 , 39(4);536-541.

Griashaw, D.L., J. Lawin, and R. Fuge.  1976.  Seasonal ana short-tera
  variations in the concentration and supply of dissolved zinc to polluted
  aquatic environments.  Environ. Pollut. 11:1-7.

Guy, 8.D., C.L. Chakrabarti, and L.L. Schramm.  1975.  The application of a
  simple chemical model of natural waters to metal fixation in particulate
  matter.  Can. J. Chem, 53:661-669.

Guy, R..D. and C. L. Chakrabarti.  197$. ' Studies of metal-organic
  interactions in aiodel systems pertaining to natural waters.  Can. J.
  Chem. 54:2600-2611.

Helz, G.R., R.J. Huggett, and J.M. Hill.  1975.  Behavior of Mn, Fe, Cu,
  Zn, Cd and Pb discharged from a wastewater treatment plant into an
  estuarine anvirotsnent.  Water Res. 9:631-636.

Hen, J.D.  1972.  Chemistry and occurrence of cadmium and zinc In surface
  water and groundwater.  Water Resource Res. 3(3):661-'679.

Holmes, C.W., £.A. Slade, and C.J. McLarran.  1974.  Migration and
 , redistribution of zinc and cadmium in a aarine-estuarine system.
  Environ. Sci. Technol. 8{3):255-259.

Huang, C.P...H.A. Elliott, and R,M. Ashmead.  1977.  Interfacial reaction
  and the fate of heavy metals in soil-water systems.  J. Water Pollut.
  Control Fed. 49(5)=745-756.   •                 '             '

Jackson, K«S» and G.B. Skippen.  1978.  Geocheaiical dispersion of heavy
  metals via organic completing: a laboratory study of copper, lead, zinc
  and nickel behavior at a simulated sediment-water boundary.  J.
  Ceochea. Explor.  10:117-138.

James, R.O. and M.A. MacNaughton.  1977.  The adsorption of aqueous heavy
  metals on inorganic minerals.  Geochim. Cosoochim. Acta.   41:1549-1555.

Joyner, T.  1961.  Exchange of zinc with environmental solutions by the
  brown bullhead.  Trans. Am. Fish. Soc. 90:444-448.

Kimball, K..D.  1973.  Seasonal fluctuations in ionic copper in Knight's
  Pond, Massachusetts.  Lirnnol. Oceanogr. 18(1):169-172.

Kubota, J., F.,L. Mills, and R.T. Oglesby.  1974.  lead, Cd, Zn,  Cu and Co
  In streams and lake waters of Cayuga Lake Basin, New York.  Environ, 'Sci.
  Technol. 3(3):243-243.

Lee, G.F.  1975.  Role of hydrous metal oxides in the,transport of heavy
  metals in the environment, pp.137-147., in Heavy metals in the aquatic
  environment.  P.A. Krenkel (ed.). Pergaiaon Press, Oxford, England. 352p.

Long, D.T. and E.E, Angina.  1977.  Chemical speciation of Cd, Cu, Pb, and
  Zn in mixed freshwater, seawater and brine solutions.  Geochim.
  Cosmochin. Acta, 41:1183-1191. r   •

Luoaa, S.N. and E.A. Jenne,  1977.  Availability of sediment-bound cobalt,
  silver, and zinc to a deposit-feeding clam, pp.213-231. _io Biological
  implications' of metals in the environment.  H. Ducker and R.E. Wildung
  (eds.). STIS COKF-750929. Springfield, Va. 682p.

Nafflfflinga, H. and J. Wilhm.  1977,  Heavy metals in water, sediments, and
  chironoaids.  J. fcater Pollut. Control Fed. 49(7): 1725-1731.

Nelson, J.L. and W.L. Hauschild,  1970.  Accumulation of radionuciides in
  bed sediments of the Columbia River between the Haaford reactors and
  McNary Dam. Water Resources Res. 6(1):130-137.       .  ,

Patrick, F.M. and M. Loutit.  1975.  Passage of metals In effluents through
  bacteria to higher organisms.  Water Res. 10:333-339.

Patterson, J.W., H.E. Allen, and J.J, Scala.  1977. 'Carbonate
  precipitation for heavy natal pollutants.  J. Water Pollut. Control Fed,

Perhac, R.M. 1972.  Distribution of Cd, Co, Cu, Fe, Mn» Mi, Pb.'and Zn in
  dissolved and partieulate solids from two streams in Tennessee.  J.
  Hydrol, 15:177-186,

Perhac, R.M,  1974a.  Water transport of heavy setals in solution and by
  different sizes of participate solids.  Univ. of Tenn.  Water Resource
  R*B. Ctr., Project No. 023. Knoxvilie, Tenn. 41p.

Perhac, R.M.  1974b.  Heavy metal distribution in bottom sediments and
  water in the Tennessee River-Loudon Lake reservoir  system. Univ. of Tenn.
  Water Resources Res., Ctr. Report No. 40, Kroxville, Tenn. 22p.

Phillips, A.R. and R.C. Russo.  1978.  Metal bloaccufflulation in fishes and
  aquatic invertebrates pp.60-67.  U.S. Environmental Protection Agency,
  (ERL), Duluth, Minn. 115p.  (EPA-6Q0/3-78-1G3).

Pita, F.W. and N.J. Hyne.  1975.  The depositions! environment of zinc,
  lead, and cadmium in reservoir sediments.  Water Res. 9:701-707.

Prasad, A.S,  1967.  Nutritional metabolic role of zinc.   Fed. Am. Sac.
  Exp. Biol. Proc. 26(1};172-185.

Rashid, M.A.  1974.  Adsorption of metals on sedimentary and peat humlc
  acids.  Chen. Geol. 13:115-123.

Rehwoldt, R., D, Kariaian-Teherami, and H. Altmann.  1975.   Measurement and
  distribution of various heavy aetals in the Danube  River and Danube Canal
  aquatic communities in the vicinity of Vienna, Austria.  Sci. Total En'
  viron. 3;341-348.

  Stetde, K. F. and A, H. Wagner.  1975.  Trace metal relationships  in bottom
    sediments of a freshwater1 stream - the Buffalo River, Arkansas.   J.
    Sediment,. Petrol. 45(1):310-319.

,  Scumm, M. and J.J, Morgan.  1970.  Aquatic chemistry. Wilay-Interscience,
    New York. 583p.

  Weast, R.C.  (ed.).  1977.  CRC handbook of chemistry and physics,  58th
    edition.  CtC Press, Cleveland, Ql\io.  2398p.

  Weatherley, A.H, and P. Da-reoii.  1973.  Zinc pollution  in a  freshwater  ays-
    ten:  analysis and proposal solution.  Search.  4(ll-12):471-476.

     Chapters 20-35

                               20,  ACRQLEIM

20" *  Statement of Probable Fate

     Literature Information .Indicates that acrolein will be removed from
aqueous environments, with half-lives usually on the order of less than a
day.  The primary loss process appears to be ah Initial hydratlon (and
possibly some biotransformat ton) to 3-hydroxypropionaldehyde, which Is
then biocransforaed.  Photolysis, oxidation, and volatilization may also b*
Important processes, but no data were available to assess whether these
processes will be faster than the Hydration-biotransfonsation sequence.

20,2  Identification

     This discussion considers only th*j fate of the molecular species of
acrolein, *an
    Boiling point at 760 tort
    (Verschueren 1977)

    Vapor pressure
    (Verschueren 1977)

    (Smith 1962)
220 torr at 20'C
330 torr at 30'C
215 torr at 20*C
269 torr at 25 aC

given'as 20.82
soluble in water

    Solubility in water at 20'C
    (Martin 1972)

    Log octanol/wa-ter partition
    (Sadding e£ al. 1977)

20,4  Summaryof Fate Data

    2&.4.1  Photolysis

         No information on the photolysis of acrolein in aquatic systems
was found,

         Acrotein has a raoderate uv absorption in the solar spectral re-
gion, with the! following data reported ,for acrolein in hexane (uv Atlas
    '»jnax (no)

         duswell e£ al. (1940) reported Chat freshly prepared acrolein
solutions showed a moderate extinction coefficient in water at 320 nm; this
coefficient decreased after the solution was allowed to stand for several
weeks.  Concommitant with this loss of uv absorption at 320 nts was the
significant grswth of a new band at 267 nm.  \'o conclusion was made as to
the process(es) occurring, but the authors suggested that oxidation or
polycondensation of .acrolein in the presence of water,gave a polymeric
material without the i,£-unsaturated carbooyl chromophore (see ,-ilso
Sections 20.4.8).  Oxidation was suggested as the causative process since ,

the presence of hydroquinone slightly retarded the'disappearance of the
320-nm absorption In aqutous solution; the formation of the acrolein
hydration product, S-hvtlroxproprlonaldehyde, was not discussed.

         Studies of the gas phase photolysis of acrolein suggest that it
may be photoreactive in the solar speutral region (> 300 nm) but no infor-
mation on the aqueous solution phase chemistry ,is available.  Coomber and
Pitts (1969) studied the gas-phase photolysis of acrolein at 313 nm; the
product quantum yields measured for the major products, carbon monoxide and
ethylene, were ~s 5 x 10" ^ anf*!• sec~l for reaction of singlet oxygen at the double bond
of acrolein at -10°C in nethanol solvent; Zepp e_£ a_l (1978) have reported a
maximum concentration of 2 x 10~12>j singlet oxygen when several natural
water samples were irradiated in sunlight in the presence of a chemical
model compound known to undergo singlet oxygen reactions.  Using this
concentration of singlet oxygen, and assuming that the rate constant Is
temperature independent between -10°C and an environmental temperature of
about 20"C» a half-lire of about 6 years is calculated.  Using a rate
constant of 0.1 M~'sec~'- for reaction of alkylperoxyl radical with the
hydrogen atom on the aldehyde carbonyl group and an ambient alkylperoxyl
radical concentration of about 10~%, the half-life for this oxidation
process Is calculated to be over 20 years (Mill 1979).  It should be noted
that hydration of either the double bond or carbonyl groups of acrolein
will change the reactivity of the respective oxidation processes, making  !
these oxidations even less important.'

         The half-life for oxidation of acrolein in the atmosphere by
hydroxyl radical is about 2 days; this half-life is based on a hydroxyl
radica'l-acrolein rate constant of 9 x 10^ M~l sec~^ (Raddlng et at.
1977) and an ambient atmospheric radical concentration of lO"""*-*" M.

    20,4.3  Hydrolysis

         Acrolein contains no hydrolyzabie groups that lead  to environ-
mental transformations.  The hydration of acroleln  to form 2-hydroxy-
propanal is not a hydrolysis process since the reaction  Is reversible.

                   CHj-CHCHO* MjO -y...-*" HOCMjCHjCHO

Similarly, the polycondensation of acroleln in the  presence  of water is  not
* hydrolysis process since it is also potentially reversible.
                     CMj * CHCMO *
                        O          HO-CHjCH C H
                        l            	.  1
               HO • CMjCMjC # + CH7 • CHCMO ^pSf CMjCMjCHO
                  '  .                    ff*.  CH.j CM-CMO

   ' 20.4.4  Volatilization

         No specific studies on the  volatilization' of acrolein  from aquatic
systems have been reported.  Bownser  and coworkers  (see Section  20.4.3) have
implicated volatilization as a  possible process in explaining a  ten-fold
greater loss of acrolein in field studies compared to a prediction by
aodelling studies.  Another study (Battelle 1970), however, discounts the
inportance of volatilization in a similar field study.  Neither  study
presents any data to substantiate the significance of volatilization.  The
application of the equations of Hackay and Leinonen (1975) to predict vol-
atilization are not appropriate since acroleln is  very soluble  in water.

   , 20.4.5  SoTptiort          '   ,           '                       -

         No data are available  to assess the importance of acrolein sorp-
tidn on sediments or onco biota in aquatic systems.  Bowmer and  Higgins
(1976) implicate sorption along with volatilization as possible  important
processes to explain the discrepancy between modelling snd field studies,
but no data were given  to support such processes (see Section 20.4.4),  The
high water solubility and low partition coefficient of acrolein  suggest
that sorption is not an  important process in aquatic systems.

    20.4.6  B i oa c cumirla 11 g n

         No infortnatton on bioaccumulation of aerolein was found.  The
small partition coefficients (octanol/water)'indicates Chat bioaccumulatIon
of acrolain will not be significant In aqaatii. systems.

    20-4•7  Biotransformation and BiodegradatIon

         Nc data were found regarding blotransformat ion of acrolein in
aquatic systems.

         Brink (1975) found that phe presence of up t-o 3 mg/ liter acro-
leln in' sewage sludge' did not effect the sludge unit efficiencies or the
BOD, COD,  TDC, pH, and dissolved" oxygen compared to a control system with
acrolein.   Wierzbtckl and Wojclk (1965) have found that activated! sludge
effectively transforms acroletn at less than 2300 ppm, but no other infor-
mation was provided.  Larson (1967) has also stated that aerolein can be
biologically oxidized in refinery systems.

         Experiments of Bowmer and Higgirts (1976) (see Section 20.4.8)
suggest that biotransformation may occur in aquatic systems since the rate
constant for loss of acrolein in supply water samples from an irrigation
area was reduced from 2.37 x 10~2 hr'1'to 1.59 x 10" 2 hr"1 when
thymol was add*»d to suppress biological activity (t^/2 of 29 hr and 43
hrf respectively); the latter half-life was the sane as that observed for
acrolein In buffered distilled water at the same pH.  However, acrolein In
Irrigation drainage water saaples containing thymol had a half-life of
three times longer ~hat in buffered sterile solutions; hence biotransfor-
nations occurring tn these systems are probably variable.  The authors also
indicate they believe that unknown catalysts in.the natural waters are re-
sponsible  for a more rapid loss of acrolein, but no evidence other than the
first-order kinetic data was provided.  Bowtner and Higgins also noted that
the pH dropped more markedly in experiments where higher initial concentra-
tions of acrolein were used (range was 6.0 to 50.5 ppm), and they suggested
thac carboxylic acids' were formed as products.
    20.4.8  Other React io_ns_

         Burczyk cc_ a_l. (1968) studied the kinetics of reversible hydration
of acrolein in distilled water to form 8-hydroxypropionaldehyde, as shown
below.  Using an equilibrium constant K of 21.2 (Smith et al.'1962), these

                   CHj-CHCHO + HjO ^T*^ HOCHjCHjCHO    K « —

authors calculated a. pseudo-first order rate constant [k|[H20]l of
0.032 day"1 for hydration of acroleirv In water; the calculated half-life
for hydratlon of acroleln Is then 21 days.  The equilibrium constant of
21.2 further Indicates that the acroleln will be ~ 95S in the hydrated form
once equilibrium is established.

         Bowraer and Higglns (1976) studied the loss of acrolein in buffered
«ater in the pH region 5 to - 8.5, where acrolein half-lives of 69 and 34
hours, respectively, were measured; presumably the losses were due to  hy-
dration* (see also Section 20.4.9).  In distilled water Bowtner and Higgins
obtained a rate constant for "acrolein decay" of 2.7 x 1Q~3 hr~^, or a
half-life of  about 11 days (no pH reported).

         There is no clear explanation to rationalize the two-fold differ-
ence in hydration rates in distilled water between the data of Burczyk jet_
ILL,'and the data of Bowtner and Higgins.  The more rapid loss of acroleln
from the experiments in buffered water and one experiment in natural water
do suggest that hydration nay be catalyzed by other .tgents, but more defin-
itive experiments are necessary.  The hydration of acrolein has been shown
to be catalyzed at high acid concentrations (Hall and Stern 1950; Pressman
and Lucas 1942) with polycondensatlon catalyzed at high base concentra-
tions (Gilbert and Donleavy 1938),
         It is of interest that while all studies on the hydration of acro-
lein have focused on the addition of water to the double bond of acrolein,
aldehydes are also known to hydrate at the carbonyl group to give geia-diols
(two hydroxyl groups on the same carbon).   Data of Bell (1966) show that
acetaldehyde is 60 percent hydrated in aqueous solution, and that electron
withdrawing groups (such as the acroleln double bond) favor more extensive
*tn all papers by Bowmer and coworkers, the aldehyde in solution was deter-
mined by coloriraetrlc procedure using 2,4-dlnitrophenylhydrazine (DNPH);
the procedure was developed by Shell Chemical Company.  Acrolein concen-
tration was determined by the difference in total aldehyde measurement
between two solutions, one of which had been purged with air.  The acrolein
was considered to hive been removed by purging, leaving a non-volatile
"degraded acrolein" product.  The authors did not characterize this pro-
duct, but is presumably as the hydrated or polycondensed acrolein (or
both), which does maintain all the aldehydic groups intact and susceptible
to the DNPH procedure (see Section 20.4.3).

hydration.  There La no information to assess the relative rates or extent
of hydration of acrolein at the two possible hydraclon positions,  Hydra-
tlon will, however, change the reactivity of acrolein toward oxidation (see
20.4.2), destroy the carbonyi chromophore responsible for any photolysis,
and possibly change the ease of biotransformation,   Hydration will also
change the physical properties of acrolein in water from those expected
based on the structure and properties of acrolein itself.

    20.4.9  mcrogosm Itudiegj jield^ Studies,  and Modelling

         Bartley and Gangstad (1974) monitored the dissipation of acrolein
in a canal as a function of distance.  Initial acrolein concentrations of
100 ppb were reduced to 90 ppb, 50 ppb, and 30 ppb at 5, 10, and 20  raiies
downstream.  No data for specific transformation, transport  or dilution
processes in the canal were given, and the authors state onlv that "very
low levels" of acrolein persist downstreaa of the application site.

         Bowtnar and coworkers reported laboratory and field studies on
acrolein (Bowmer and Higglns 1976; O'Loughlin and Bowmer 1975; Bownter e£
al. 1974).  In the 1974 paper they reported that acrolein in bottles and in
a large tank of water (sterility or quality of water was not described)
disappeared much faster when determined by bioassay or gas^-purging methods
than when the water was analyzed directly by colorimetric methods using the
dinitrophenyl hydrazine [DNPHj procedure (sea Section 20.4.8),  The authors
concluded that acrolein was "degraded" to a nonvolatile, nonphytotoxic
(less than 0.1 tine's as toxic as acrolein) product that gave a positive
DNPH test.  They also reported that aerolein disappearance in their experi-
ments had a first-order rate constant of 0.83 day"* (f-l/2 * 0.83 day).
This loss rate was much faster than that predicted due to hydracion alone
(see Section 20.4,8); furthermore, no equilibrium concentration of acrolein
with hydrated acrolein was found.  No specific loss processes other"than
hydration were discussed,  except that volatilization was determined to be
much -less important than "degradation" in these experiments. The 1975 paper
extended the acrolein studies to evaluations using a one-dimensional model,
with field studies u,sed to measure "decay rate constants" for acrolein of
0.14 to 0.21 hr-1 (ti/2 of 5 to 3.3 hours, respectively). The subse-
quent 1976 paper reported  studies of acrolein loss ia buffered and natural
waters in the laboratory and In field studies.  Although an  equilibrium
concentration of acrolein  was found in the buffered waters (see Section
20.4.8) no equilibrium acrolein concentrations were found in the natural
waters where all acrolein  was eventually lost.  The rate of  loss of a'cro-
letn in these natural waters was 1.5 to 3 times faster than  that'in buf-
fered water.  Addition of  thymol to suppress biological activity did de-
crease the rate of loss of acrolein in the natural  waters so that it was
similar to the rate in sterile buffered water in one experiment; another
natural water experiment containing thymol, however, showed  an acrolein

loss rate that was still three times faster than that in buffered water.
Thus, hydration and biodegradation seem to be important  removal processes
for acrolein.  Bowser and Higgins .further suggest that substances in natur-
al waters may be catalyzing the '"decay" process (presumably referring to
hydracion) so chat it is faster than the reaction in distilled water.  The
papers by Bowmer and coworkers do not, however, attempt  to identify acro-
leln products other than by the 2,4-DNPH test which, is useful only for
analyzing for total carbonyl.

         Bowmer and Higgins '(1976) also monitored the' kinetics of the loss
of the acrolein "degradation product," and found that although it was
stable in buffered waters, the loss in supply waters was consistent with
the "lag phase" character of microbiological processes,  with a lag phase
time of - 100 hours froo addition of acrolein to the water.
20-5  Daja_ Summary                                 !

    Table 20-1 summarizes the data on the aquatic fate of aerolein.

                                                  Table 20-1

                                        Suwury of Aquatic fate of Ac role in
  Process d








May be important  process.

Oxidation  is slow.

Not an important  process.

May be important  process.

Hot an important  process.

Mot an important  process.

Is an important process.
                                                                                  Ufe t
                                                                                 > 6 yra.
                                                                                < 4 days
  of Data




   Hedtuai -



a. Tlitre is insufficient Information in the reviewed literature  to permit assessment  of  a  most  probable  fate.

b. Literature data suggest that acrolein may undergo reversible  hydration of  acrolein to a product  which is  readily
   btotransformed.  Some experiments show a halt-lift of about 3-5 hours for  acrolein.

20.6  Literature Cited

Bartley, T., and E.O. Gangstad.  1974.  Environmental aspects of   aquatic
  plant control.  J, of the Irrigation and Drainage Oivision-ASCE.

Battelle 1970.  Information supplied by EPA.
                                                *                   »

Bell, t.P.  1966.  The reversi'ule hydratlon of carbonyl compounds.  Adv.
  Phys. Org. Chen.  4:1-29.

Bowoer, K.H., and M.L. Higgins.  1976.  Some aspects of the persistence and
  fate of aerolein herbicide in.water.  Arch. Environ. Cantam. Tax.

Bowmer, K.H., A.I.G. Lang, M»L. Higgins, A.R. Pillay, and Y.T. Tchan.  1974.
  Loss of acrolein from water by volatilization and degradation.  Weed Re-
  search 14:325-328.
                  I                        >                i
Brink, R. fl., Jr.  1975.  Studies with chlorophenols, acrolein,
  dithiocarbaoates and dibromonitrilopropionamide in bench-scale  biodegrada-
  tion units.  Proc. Int. Blodegradation Sytnp, , 3rd.  1975 (Pub.  1976)
  785-91. CA87': 90236.

iurczyk, L.» K, Walczyk, and R. Burczyk.  1968.  Kinetics of reaction  of
  water addition to the acrolein double-bond in dilute aqueous solution.
  Przea. Chem.  47(10):625~627.

Buswell, A.M., E.G.  Dunlop, W*H. Rodebush, and J.B, Swartz.  1940.  Change
  of the ultraviolet absorption spectrum of'acrolein with tine.   J. Amer.
  Chea. Sac.  62(2)-,325-8.

C*mler, J.C.» and X. Deglise.  1973.  Kinetic study of che photosensitized
  oxidation of a»S~uasaturated aldehydes in the liqui.d phase at -10°C.
  C.R. Acad. Sci., Ser. C  277{22}:1187-1189.

Coomber, J.W., and J.N. Pitt*, Jr.  1969.  Molecular structure and pho- ,
  tochenlcal rectlvity.  XII.  The vapor-phase photochemistry of  acrolein
  at 3130 A.  J. Aa«r. Chen. Soc.  91{3):547-550.

Gilbert, E. E.,lai«f J.J. Oonleavy.  1938.  The polycondensatton of
  acrolein.  J. Araer. Chea. Soc.  60:1911^1914.

Hall, I.H., and S.S. Stern.  1950.  Acid-catalyzed hydratlon of
  acraldehyde.  Kinetic* of the reaction and isolation of
  S-hydroxypropaldehyde.  J. Chem. Soc.:490-498.

Larson, R.O.  1967.  Biological treatment In water conservation.  Amer,
  Chem. Soc., Diy. Petrol. Chen., Preprints 12(4):A63-A70.

Mabey, W.R., and T. Mill.  1976.  Kinetics of hydrolysis and oxidation of
  organic pollutants In the aquatic environment.  Syrup, on nonbiological
  transport and transformation of pollutants on land rnd water;  Processes
  and critical data required for predictive description.  May 11-13, 1976.

Mackay, D., and P.J. Lelnonen.  1975.   Rate of evaporation of
  low-solubility contaminants from water bodies to atmosphere.  Environ.
  Sci. Tech. 9(13):1178-1180.

Martin, H.  (ed.).' 1972.  Pesticide manual, 3rd edition.  British Crop
  Protection Council, Worcester, Engl.

Mill, T.  1979.  Structure reactivity correlations for environmental
  reactions.  EPA Final Report EPA 560/11-79-012.

O'Loughlin, E.M., and K.H. Bowmer.  1975.  Dilution and decay of aquatic
  herbicides in flowing channels.  J.  Hydrol.  26(3-4);217-35.

Pressman, D., and H.J.  Lucas.  1942,  Hydration of unsaturated compounds,
  XI. Acrolein and acrylic acid.1  J. Amer. Chem. Soc.  64:1953-1957.

tadding, S.B., D.H. Liu, H.L.' Johnson, and T. Hill.  1977.  Review of the
  environmental fate of selected chemicals.  U.S. Environmental Protection
  Agency Report Mo. EPA-560/5-77-003.

Smith, C.W.  (ed.).  1962.  Acrolein.   John Wiley and Sons, Inc., New York.

UV Atlas.  1966.  Plenum Press.

Vefschueren, SC.  1977.   Handbook of environmental data on organic
  chemicals..  Van Nostrand Reinhold Company.  659 pp.

Wlerzblcki, T., and  0. Wojcik.  1965.  Preliminary trials on the de-
  composition of acroiein, ally! alcohol, and giycerol by activated
  sludge.  Zesz. Sauk.  Politech, Slaska, Inz. Sanit.  (8):173-85, via
  Chemical Abstracts 63:1589z (1963).

Zepp, R. G. , N. L. Wolfe, G. L. Baughman,1 and R. C. Hollis.  19?3.   Singlet
  oxygen in natural water.  Jtature 278:421.

                                21.  ALDRIN
21.1  Stateaent^^Prolbable^^Fate

    Blotranaformatlon, volatilization, bioaccunmlatton, and indirect pho-
tolysis all may be important fate processes for aldrin introduced into
aquatic environments.  Volatilization half-lives of less than a few days
are likely in aquatic systems when sorption to biota and subsequent
biotransformacion to dieldrin do not occur rapidly*  Photosensitized and
photooxidation processes may also be important fates for aldrin, but
insufficient information is available to assess how general and reliable
these processes are for environmental assessments.

21.2  Identification

    This chapter considers aldrin as the pure chemical.  The structure,
alternate names, CAS and TSL numbers for aldrin are given below:
                                   Compound 118
  '  Aldrin

    CAS Mo. 309002
    TSL No. 10 21000

21.3  Physical Properties

    The general physical properties of aldrin are given as follows.

         Molecular weight                    365
         Melting point
         (Martin 1972)

         Boiling point at 760 torr

         Vapor pressure
         (Martin 1972)
         (Gunrher and Guneher 1971)

         Solubility in water
         (Park and Bruce 1968)
         CWeil 1974)   ,
         (Biggmr and Rtggs 1974)*

         Log octanol/water partition
         coaf f icier.t
No data found

2.31 x ICT5 torr at 203C
6 K 1Q'6torr at 25°C
27 ppb at 2>~29"C
17 ppb at 25*0
180 ppb at 25°C

No data found
*Particle size < 5.0 yra,

21.4  Suo««ryjLjof far e Oat a

    21.4.1  Photolysis

         No data are available to estimate the direct photolysis rat* con-
stant for aldria in aquatic systems, although two experiments suggest di-
rect photolysis nay b* slow compared to indirect photolysis in aquatic en-

         Ross and Crosby (1975"* txarained the photolysis of aldrin  In aque-
ous solutions containing acetone or acetaldehyde and in a sterile  paddy
water example.  Although photolyses were performed with a sunlaup  with
wavelengths > 300 no ana in sunlight, the information and data presented  ir.
the paper do not specify which Irradiation source was used.  Therefore, no
quantitative estimate of aldrin photolyses in aquatic environments can be
made.  The authors reported that irradiation of aldrin in demineraiized
water for 10 hours produced no los^ of aldrin.  Photolyses of aHrii for  1
hour in water containing 0.1 percent acetone or acetaldehyde gave  22 per-
cent and 24 percent yields, respectively, of dieldrin (I); in the  paddy
water, a 25 percent yield of dieldrin was obtained after, 36 hours  irradi-J-
tiotu  Unlike other photolyses of aldrin (see below), these photolyses did
not giv« photoaldrln (II), a cage-type product.  The authors concluded that
.photochemically generated oxidants were responsible for aldrin conversion
to dieldrin.                                   '

             I                                         II
         Dlel-Jrtn                               •  Phoeoaidrlr
         SlngSiaster *(19") reported that the photolysis of 0,33 ppb  of
aldrin in a sample of San Francisco Bay water gave a half-life of  1. I day
in sunlight.  In contrast to the pr>du<.ts observed by Ross and Crosby
:(19?5), the photolysis product In -Si ngmaster *s study was determined  not  to
be dleldrin; the product corresponded to the retention tine  of photoaldrin
in the glpc analysis.         '                                   '

         Other studies of the photolysis of aidrin have been  reported,  but
the reaction conditions -ised are not relevant to aquatic svsteas,  A brief
discus Q
kcal/aole did show som« sensit Izat ion in aldrin conversion; the product  of
thd reactions apparently was pho«,oaldr in.  A subsequent  paper  (Ivie and
Casida 19?lb) reported that rotenone codeposited  with aldrin  on bean  teayes
and exposed to sunlight was also effective in sensitizing photolysis  of
aldrin; both photoaldrin and photodieldrin were found as products.

         Rosen and Carey (19683 reported that photoaldrin was  formed  in  a
77 perct.it yield when aldrin was photolyzed at 2*»8-35(>.nm in  benzene  solu-
tion in the presence of the senstti^er  benzophenone.  Rosen and Sutherland
(1967) reported the sunlight photolysis of an aldrin film in  a glass  dish.
At the end of one month's photolysis (mid-June to ntd-July) the reaction
mixture contained 2.6 percent aldrin, i.l percent dieldrln, 24.1 percent

photodieidrln, 9,6 percent photjaldrln, and 59.7 percent polymeric materi-
al.   Analysis of a film of aldrin exposed to 3 days of sunlight in earlv
June showed approximately 13 percent aldrin regaining, with 31 percent
dieldrin,  7 percent photodieldrln, 1 percent photoaldrin, and i<5 percent
polymeric  material formed (see Chapter 26 on dleidrin for discussions on

         Henderson and Crosby (1967) found that the dechl^rinated aldrin
is osier III was the riajor photolysis produr*. of aldrin In hexane solvent
Irradiated at 25-> nm; the same product was also obtained in r.ethanol and
eye iohexan.e solvents.  Rosen (196?) found a 24 percent yield of II and  13
percent yield of III when solid aldrin was phoeoly?ed at ^5^ na, with 9
percent aldrin recovered; the remainder of the Material »*as described as j
polytr.er.  Gab et al. (1974)  phocolyzed aldrin in the solid phase using a
high pressure Hg lamp, with pyrex filter (> 300 no) and found small amounts
of photoaldrin, dieldrin, HC1, and CCH; aost of the material recovered
was identified as a polysaer.  Roburn (19^3) found that a film -~>f aldrin
irradiated at 25^ rua gave dieldrin as the aajor product.

         Since.volatilization of aldrin from water bodies appears to be an
important process, Cros'iy and Moilanen (1974) studied the vapor phase pho-
tolyses of aldrin.  Photolysis of sldrin at concentrations below saturation
in the vapor phase for 168 hours with a sunlamp resulted in a 40 percent
racoverv of' aldrin ind a 63 percent yield of dieldrin, with about 2 percent
ytelds each of photoaldrin-and photodieldrin; in a dark control experiment
34 percent of the aldrin was recovered and a yield of 14 percent dieldrin
was taeasured.  The authors attributed the presence of dieldrin in the dark
control as being due to contamination problems.  Other photolysis experi-
ments with aldrin at saturation in the vapor phase also gave photoaldrin
and dieldrin as product,*.  Although these studies indicate that aldrin  pho-

toiysis may occur In the atmosphere, !no data are available to predict how
fast the photoreactlons will occur.

    21. A. 2
         Although photooxidation processes are Important in the formation
of dieldrin frort*aldrin, no data are available t'o assess the half-lives for
oxidation processes in aquatic environments (see Section 21.4; 1).

         In studies on the photooxidation of aldrin to dieldrin, Ross and
Crosby (1975) found that aldrin was not oxidized by singlet oxygen.  They
did find that peraeetlc acid oxidized aldrin to dieldrln in the dark, but
no information is available to predict how fast such peroxide oxidations
might be in aquatic environments,

         Hoffman and Eichelsdoerfer (1971) found that some ozone readily
oxidized aldrin to dieldrln in hexane and in 10 percent acetone in water
solvents.  Saravanja-Sozanie ££ «!» (1977) reported that oxygen atoms oxi^
dlze aldrin to dieldrln, but no data are available to predict the rates at
which this oxidation process aay occur in the ataosphere.

    21.4,3  Hydrolysis    '•

         No Information Is available on the hydrolysis of aldrin in aquatic
systens; hydrolysis of aldrin is not expected to be Important, however,
since the chlorine groups ar« located at bridgehead and vinyiic positionson
the aldrin structure, and these positions are not readily hydroiy~able.
Hydrolysis half-lives of aany years are probable.  Elchelberger %n1
Lichtenberg (1971) found that aldrin was 80 percent converted to aieldrin
in a sample of raw river water in 8 weeks; these results suggest that bio-
transformation Is probably aore rapid than hydrolysis in aquatic environ-
ments,                       :

   •'21. A. 4  Volatilization      •          .-•'-.

         Volatilization of aidrin froa aquatic systems Is an important
process, with half- lives on the order of a few hours to a few days,

         Singmaster (1975) described studies designed to measure relative
races of volatilization of chlorinated pesticide from pure water and
several natural waters.  In these experiments approximately 1 pptr concen-
trations of pesticides in 900 «1 of water in a ,5-liter flask were gently
agitated on a, shaker while air was drawn through th* flask (but not bubbled
through solution) at a rate of 4.5 l/min»,  The half-lives for volatiliza-
tion of aldrin from pure water and waters from San Francisco Bay,  the Amer-
ican River, and Sacramento Riv«r (all in California) were 0.38 hour, 0.59

hour, 0.60 hour, and 0.60 hour,  respectively;  the water loss In these ex-
periments averaged 3.6 * 0.2 g/hr.   Slngmaster concluded that  volatiliza-
tion of aldrin in natural waters would not  be  snore than two times slower
tlian in pure water.   Although the total experiment is difficult to relate
to conditions in aquatic environments (e.g., temperature,  agitation), the.
author noted that the air exchange  in the  flask corresponded to a wind
velocity of about1 10 ra/hr, which is much  lower than the velocity usually
found in the environment.  Based on the wind velocity factor alone and
assuming removal of  aldrin from  the vapor  space is the dominant force in
volatilization' of aldrin, the half-lives  for volatilization of aldrin in
aquatic environments could bt on the order of  a few hours«

         Mackay and  WoIkoff (1973)  and subsequently Mackay  and Leinonen
(1975) described equations for calculating volatilization rates of chemi-
cals based on mass transfer data; the calculated half-life  for volatiliza-
tion of aldrin was 7,7 days.  In their, calculation, Mackay  and coworkers
used 200 ppb as the  water solubility of aldrin, and as seen in Section
21.3, this value is  about an order  of magnitude higher than two other mea-
sured values.  If the lower values  of aldrin solubility are used, the cal-
culated volatilization half-life would be  approximately an  order of magni-
tude shorter than the value calculated by  Mackay and coworkers, or about 1

         Additional  evidence for the rapid volatilization of aldrin has
also been obtained in volatilization studies performed by Ernst (1977)
preparatory to conducting bioaccumulation  measurements.  lu aerated aquaria
containing seawater  and various  pesticides, Ernst observed  a 97 percent
loss of aldrin, compared to losses  of 11 percent for dieldrin  ^nd 67 per-
cent for DDT.

    21.4.5  Sorptlon

         Very few data are' available to evaluate the importance of sorption
of aldrin tc sediments in aquatic environments.  Kenaga and Goring (1978)
cite a Koc value for aldrin of 410, which  suggests that sorption of
aldrin to sediments  will not be  extensive  in aquatic systems.   Sorption
will eventually remove aldrin frora  aquatic systems if biological processes
do not occur, however,

         Leshniowsky e£ ajL,  (1970)  studied sorption of aldrin  by bacterial
floe and by a lake sediment.  The authors  found that aldrin equilibrium be-
tween water and floe was attained in about  20  minutes, with a  concentration
factor of 625; ic was not stated whether the concentration  factor was based
on a wet- or dry-weight basis.  Equilibrium of aldrin between  sediment
(silt) and water was attained in '.0 minutes, but no partition  coefficient
was given.

         Yaron e£ al. (196?) studied the adsorption of aldrln on soils,
clays, and sands.  Less aldrin was adsorbed in sand Chan on clays; more
aldrin was sorbed by a soil with 6 percent organic matter than the same
soil with' th«i 'organic matter removed by oxidation.  Unfortunately, the
authors did.uot analyze and present their data in a, manner that is useful
for estimating sorpcion coefficients for environmental assessment appli-

    21.4,6  Bloaccuaulation

         The results of terrestrial-aquatic microcosm experiments indicate
that bioconcentration factors for aldrin in aquatic systems will be
approximately 10-* - 10 .  Although biouptake may then be an important
process for aldrin in short time periods, significant bioaccumulation of
aldrin through the food chain will probably not occur because it is quite
rapidly converted in dieldrin.

         In %«rrestrial-aquatic microcosm experiments, Metcalf et al.
(1973) measured bioconcentration factors of 3.9 x 10&, 4.5 x 10 , and
3,1 x 10^ for alga, snail, and fish, respectively.  Given the presence of
large amounts of dieldrin in these species (see Section 26.A.6), these
factors may be low, because metabolism of aldrin to dieldrin may also be a
process competing with uptake/depuration in the species studies.

         Metcalf e_£ a 1. (1973) conducted separate experiments in which
Daphnia, mosquito larvae, and fish were exposed to aldrin by water contact
only (no food chain).  Bioconcentration factors for the fish ranged froa
260 to 460 between the first and third days of exposures.  Values for the
Daphnia ranged from 1800 to 9100, and for the larvae from 970 to 1100.

         Schauberger and Wildman (1977) reported aldrin concentrations in
two species of blue-green algae that were 1 x 10^ and 1.3 x 10^ tines
the initial aldrin concentrations" of the solutions in which they were
grown; a third algae failed to accumulate aldrin, but rather effected
•poxidation to give dieldrin.

    21.4.7  gigtransformation and Biodegradation      •      •

         Although biotransformation of aldrin to dieldrin is probably the
dominant transformation process in aquatic systeas, chare are no useful
data for predicting the rate at which aidrin epoxidation may occur in the

         The rapid transformation of mldrin to dieldrin JU» vivo has been
shown by Bowman et_ al. (1964), who introduced aldrin into solutions con-
taining mosquito larvae.  Some aldrin, but no dieldrin, was recovered from

the solution phase after 20 hours at 27"C; however, In che  Larvae dieldrin
predominated over1 aldrin by factors of 10:1 and 1.6:1 for initial aldrin  •
concentrattotis of 24 and 270 ppb, respectively. '

         Eichelberger and Lichter.berg (19?I) exaninec*. che persistence of
aldrin in raw water from th? Little Miami River fjc 8 weeks; ac  the end of
1, 2, 4, and 8 weeks, the percentages' of aldrin remaining were  100, 80, 40,
and 20 percrnt, respectively.  Since dieldrin w
drin if. the tsicrocosw system.  The discussion and data  for  these  experl-
tn^nts are presented in Sections 21.4.6 and 21.4.7.

    Table 21-1 summarizes Che data on the aquatic fate of a^drin.

                                                                        Table 21-1

                                                             Summary of Aquatic Fat« of Aldrin
                               Env ironaental



Direct photolysis  in

Indirect photolysis
•ay be important.

No quantitative
informalion available.

Not an important process.

Probably an iuport<*nt
 otpti .n            O»n  be  an  important

 .ioaccunuiation     Posaibly ait  Import an

 iot ransf ornat ion/  Is a doniitant  pioiea
  Blodegradat f oa    in aOme s-'Stems.
                                                 Rate          Life
                                                                                       One  expt,  I, M  day
                                                                                        Few hours
                                                                                        day a
                                                                                                  to few
  of Dat i

                                                                                                                   Ned i urn
                             a.   There  l» insufficient iuf orin..; ion in tlt« reviewed  literature  to permit asseb^mt-nt ot d mo a
                                  probable fate.

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Bowman, M.C., F* Acree, Jr., C.S, Lofgren, and M. Beroza.  1964,
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Carlson, G.F.  1974.  Epoxidation of aldrin to dieldrin by lobsters.
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Crosby, D.G., and K,W. Moilanen.  1974.  Vapor-phase photodecomposition of
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Sichelberger, J.W., asnd J.J. Lichtenberg.  1971.  Persistence of
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Weil, L., G. Dure , and K.E. ^uentin.  1974,  Solubility in water Of
  insecticide chlorinated hydrocarbons and polychlorinated biphenyls in
  view of water pollution.  Z. Wasser Abwasser Forsch.  7(6) : 169-175.

Yaron, 3., A.R. awpboda, and G.W,  Thomas,  1967.  Aldrin adsorption by
  soils and clays.  J. Agrxc. Food Ch-.-a. 15(4) :67l-675.

                              22.  CHLORDANE

22. 1  Statement of Probable Fate

    Volatilization, sorption Co sediments, and bioaccumulat Ion are impor-
tant fates for the chlordane isoners in aqueous environments.  The chlor-
dane isomers jlso undergo photosensitized isotnerizat ions, but no informa-
tion is' available to determine whether such reactions way occur in aquatic
systems.  Although biotransformations of chlordane may be Important for the
ultimate transf o.rraat ion of ..hlordane, these processes are likely to be very
slow in the environment .

22. 2  Identification

    This chapter considers only the two major components of the mixture
known as technical chlordans.  An approximate composition of technical
chlordane is given below (Brooks 1974, from data of Velsicol Chemical

_ Fract ion __ • _ Present __
Diels-Alder adduct of cyclopentadiene             , 2 jf  i
  pentachlorocyclopentadiene (CIQK-/CI--, )

Chlordene (CioHfeClft); isomer I              .       I +  1
Chlordene isomers 2,                               7.5 + 2
                  3, and A together                13^2

Heptachlor tC^oHsClj}                              10 * 3

Cis_- chlor dane (C[oH^Clg) (3-isoaer)      . ,         19 4- 3

T rans- eh 1 o rda ne (C^QH^Clp) (y-isomer)              2-i 4- 2

Nonachlor '"iQHsClg)                               7 j- 3

Hexachlorocyclopentadicne (CgCl^,)                  > I

Octachlcitocyclopentene (C^Clg)                     I *• 1

C]0«7-8Cl6-7         '

Constituents with shorter GC retention  .
  time than C5Clg (includes hexachloro

Constituents with longer GC retention cimes than
  npna chlor


Sovocool et_ a_l. (1977) analyzed a  technical  chlordane  sample by gas chro-
matography/inass spectromet ry and partially or  completely identified 45 con-
stituents of the mixture.

    Morley e_t_ al_, (1974) reviewed  information  or.  technical chloriarse and
IwS components, and provided an exceiler.t discussion on  t^e  Msccry of the
nomenclature of the ehlordane  isomers.   The Y-chlordane  designation has
been used for both the vicinal dicMoro  sructure  and for a. geminal df,chloro
structure.  Cis~ and t ran s~ ch 1 o rda ne have been called  <"-*•- and Y-chlordane,
respect ively^ by one group of  researchers t|' and -—  and  "-'-chlordane, re-
spectively, hy another group.  This report uses the cis-and  trans de-
signations to avoid Confusion, and any references  in  literature to the
Greek nomenclature are converted to the  corresponding  c_is_/t_jrans_ notation.

    The structures, alternate  names, and CAS and  TSL numbers of the two
chlordane Isonsers are given, below:

                                             Alternate Names
                                             Trans- chlordane
                                              •4 ,7-Methanoindan,  !,2,-i,3,
                                                 6, 7 ,8 ,8-oetachloro- la ,
        Trar.s-ch lordane
'CAS No.
                FB 98000
                                      -^,7-Methanoindan,,  alpha
                                         1, 2,4,5,^,7,8,R-octac
                                         la ,•* ,7 ," i-t et r ^hy-iro

        CAS N'o. 5103-74-2
        tSL 'Mo, PC Ul?50

22.3  PhysicalProperties

    The general physical properties of the chlordane Isomers are -given be-
low, with the data for individual Isomers noted when available.

      Molecular weight     •   •               406

      Melting Point                          107.0-103.38C (cis)
                                             103.0-105.0*C (.trans)

      Boiling point                          1753C at 2 torr
      (Roark 1951)

      Vapor pressure at 25 °C                 1 x 10~^ torr*
      (Martin 1972)

      Solubility in water at 25"C
      (Weil et_ al. 19";- )                     1.85 ppta
      (Sanborn e_t_ al. 1976)                  0.056 ppm

      Log octanol/water                      2.78
      partition coefficient
      (Sanborn et al,
*For "refined product", which was not further defined.
22.4  Summary, jsf_r Fate	Data i

    22.4.1  Phgtolgsis

         .There are no data useful for estimating the half-life for pho-
tolysis of either of the two chlordane isoraers in aquatic environments.
Although no information is available for assessing how the processes may be
relevant to photosensitized reactions in aquatic environments, the chlor-
dane isoraers have been shown to undergo photosensitized reactions with
acetone as sensitizer, with the cjla isomer being more susceptible to pho-
solvsis than the trans isoraer,                                          (

         Several studies on the photolysis of chlordane isomers in acetone
solvent have bsen reported.  Fischler and Korte (1969) reported that cis-
chlordane photolyzed at'254 nm gave the cage product I; trans-chlordane
did not undergo photoisomerizacion.   A subsequent paper by Parlar and Korte
(1973) reported that at wavelengths above 3UO nm cis-chlordane 'photolyzed
to give I, and also gave product II in which a chlorine atom migrated to
the tertiary carbon.  Benson js_t al.  (1971) also found that I was a product
from cis-chloraane photolysis at > 300 nm.  Trans-chlordane also photo-
reacted under these conditions, but no photoisomer was found;  the authors
reported that only 60% of trans-chlordane was reacted in 50 hours, whereas
cia-chlordane was 99% reacted in 27 hours using the same light source.

Onuska and Comba (1975) reported that c_i_s-chlordaae photolyzed at > 300 nm
gave I, and that trans-chlordane gave the two products, ill and IV.  The
saae products were also reported by Ivia et al. (1972) and tCnox et a4.
         Ivie ejc a 1. (1972) also studied t'.ie photolysis of cis- and trans-
chlordane in sunlight on the surface of bean leaves that had been treated
with ro>tenone (a photochemical sensitizer).  In the bean leaf study expo-
sure of the chlordane tsomers to sunlight fot 4 hours resulted ,ln a 70 to
80% loss of cis-chlordane and a 15 to 20% loss of crans-chiordane; no loss
of chlordane from the bean leaves exposed to sunlight occurred in the abs-
ence of' rotenone.  The same phutoisomerization products were found in the
leaf study and in the acetone solvent experiments.

         Vollner et a1. (1971) reported that photolysis of cis- or trans-
chlordane at 254 nm in dioxane-water solvent gave mainly photoreduction at
the vinylic group of the chlordan? structure.  This result indicates that
formations of products I through j'V are probably due1 to acetone-photosensi-
tized processes,.  There is no information available to determine whether
such photosensitized processes may occur with natural substances present, in
aquatic environments.

         Benson et al. (19/1) reported that exposure of a thin film of cis-
chlordane to sunlight fov <*60 hours resulted in a 10% loss, with a 1% yield
of i.  Baker and Applegate (197'4) also found that chlordane films subjected
to irradiation from a lamp with a maximum output at .350 nm showed extensive
loss of chlordane; neither the chlordane tsotner(s) studied nor any products
were reported,  Ginsburg (1953) also reported that a chlordane emulsion
lost toxicity to tnosquitos after 6 days exposure to sunlight; no further
information was given.

  •  22.4.2  Oxidation       '                                ,

         No information is available Co assess the oxidation half-life of
chlordan-e in aquatic environments.

    22.4.3  Hydrolysis

         Eichelberger and Lichcenberg (1971) reported that both Isomers of
chlordane were 100 percent recovered from samples of tittle Miami River
(Ohio) water after 8 weeks at  room temperature; since the recoveries were
rounded off. to the nearest 5 percent, tht loss must have been less than 2.5
percent, which corresponds to  a half-life of at least 4 years for, hydroly-'
sis.  Bevenue and Yeo (1969) also reported that both chlordane isomers were
stable in water for 60 days.

    22.4.4,  Volatilization

         Laboratory experiments by several researchers suggest that vola-
tilization may be an important loss process for the chlordane isomers in
aquatic environments, although Che information available is not uaaful for
estimating volatilization half-lives in aquatic systems.

         Oloffs et_ _§_!. (1972;  1973) and Oloffs and Albright (1974) reported
experiments in which cis- and  trans-chlordane were incubated in flasks con-
tainiitg natural waters and sed^aents.  Chlordane in water without sediments
was readily volatilized,  with  some of the chlordane found in the glass wool
plugs in the neck of the incubation flasks.  Losses of chlordaae isomers
were' approximately 40 to 50% in 6 weeks and about 60% in 12 weeks* presuma-
bly because of volatilization  and possibly biotransformation; no chlordane
metabolites were detected in the aqueous phase, however, when ^G~labeled
chlordane vms. used.  Other experiments with an initial concentration of 25
ppb trans-chlordane in natural water at 9'C for 3' days resulted in 3.41 of
th€ chlordane being recovered  from the glass wool plug and 70% of the
chlordane recovered from the water.  In a similar experiment with'cis-
chiordane, about 24% and 16* of the chlordane was found in the glass wool
and water, respectively.   Flasks sealed with glass stoppers showed no loss
of chlordane from water.   Other experiments where surfactants or sediments
were added to the flasks sh'owed less chlordane loss, apparently because
volatilization was suppressed.

    ;     Bowman et al. (1964)  found that volatilisation of crans-chlordane
was apparently a. significant loss process.  In their experiments, 51 Vg of
trans-chlordane was placed in  a 250 ml solution of water-chlordane contain-
ing mosquito larvae; at tha'end of 20 hours, only 30% of the" chlordane was

recovered from water and larvae.  The authors suggested that volatilization
was responsible for the poor recovery of chlordane.•

    22.4.5  Sorgtion

         No useful data are available for evaluating the aorptlon of ehlor-
dane to sediments in aquatic environments.  Oloffs _«£ a_l. (1972; 1973) and
Oloffs and Albright (1974) reported experiments in which cis- and trans-
chlordane were • placed tn flasks containing natural water and sediaents.
Volatilization was an important loss process in the experiment without
sediment (see Section 22.4.4).  When sediment was present, more than 80% of
the chlordane initially at 25 ppb in solution was recovered from the sedi-
ment after 12 weeks.  The authors stated that these experiments showed the
competition between sediment sofption and volatilization in the fate of

    22.4.6  ii oaccuaula t ion

         Bioaccumulation in bacteria, daphnids, and fish is an important
fate for the chlordane isomers'ln aquatic environments.  Bioconcentratlon
factors for these species are on the order of 1CH to 10*.

         Grimes and Morrison (1975; reported bloconcentration factors for
cis- and trans-chlordanc in bacterial species.  The concentration factors
f'ar the individual isomers in each species did, not vary by snore than 10Z
from the average of the concentration factors, as may be expected of such
similar structures.  Tha bioconcentration factdrs for cis~chlordane varied
from 2.0 x 102 co 5.6 x IO4; four factors Were betweenToO and 900, and
four factors were between 1CH and 104,-with three factors at 1.6 x
104, 2.3 x iO4, and 5.6 x IO4, specifically.

         Moore *£ £l« (1977) reported chlordane bioconcentration factors of
5.6 x IcP and 2.4 x 10  for algae and daphnids, respectively.  The
daphnids were found to readily eliminate chlordane, with 75% lost in 4,9
hours after the daphnids were placed in clears water.   In another experi-
ment, concentration factors in daphnids of 1.6 x IO4 and 2.0 x IO4 were
meseared for gis- and trans-chlordane, respectively.   The authors also re-
port a bioconcencration factor of 162 for goldfish; sfnce this factor is
calculated on a wet weight basis, It is lower than values figured on a dry
weight basis.  When renoved to clean water, the goldfish eliminated 50% of
the absorbed chlordane in 3 days,

         Schimmel _et al. (I976b) reported bioconcentration factors for
trans-chiordane l.i the range of (3.6-6.4) x 10^ for the fi h, spo":, in

72-hour exposure test.  Schlomel et a,I, ,(197oa) also found trans-chlo rda ne
concentration factors of (3.7-14.8) x 103 for spot and (9.0-16,3) X I03
for sheepshead- minnow in 96-hour exposure tests.  Parrish et aJL. (1976) re-
ported that bioconcentration factors for chlordane (no isomers cited) in
marine species in 96-hour exposure tests (1,3-1.9) x 10  for sheepshead
minnow and 3.0-7.5 x 1CH for pinfish; in a 28-day exposure test, the
sheepshead minnow had concentration factors in the rant»e (8.5-12.3) x

         Roberts £t £!• (1977) reported a study of the bioaccumulation of
chiardane iaomers in fish (redhorse and white suckers); a concentration
factor for cls-chlordane in fish was estimated to be 5.5 x 10-%  The
authors also found that the trans-isomer was eliminated from the fish about
1.8 times faster than the eis-isomer, with half-lives for elimination of
cia-and trans-chlordane being 60 and 30 days, respectively.

    22.4.7  B.lotrans|oraation and Biodegradation

         Vety little information is available on the biotransformation of
chlofdane, and no data were found for estimating the bio trans format ion
half-lives of chlordane in aquatic systems.

         Castro and Yoshida (1971) and Watanabe (1973) both reported that
chlordane is persistent in flooded and nonflooded soils and suggest that
chlordane is comparable to dieldrir. in its slow biotransfortnation.

         lyengar and Rao (1973) reported that the fungus, Asgargillua
nlger, could utilize chlordane in nutrient solutions, with more than 90%
loss found in 48 hours for chl >rdane concentrations initially below 38 pptn.
The authors found that chlordane could not serve as a sole carbon source
for growth of the fungus, and unadapted organisms could not utilize the
pesticide.  They did find that the heptachlor-acclimated-fungus could util-
ize chlordane as a substrate.  So metabolites were identified*

         In terrestrial-aquatic microcosm exeriuients, Sanborn ££*_!• (1976)
found some metabolism of chlordana, but they did not Identify any of the
some 20 products detected by tic (see Section 22.4.9).  Morley _et stl.
(1974) reviewed the metabolites found in rabbit experiments and concluded
they were products from hydroxylation at the chlorinated positions of the
cyclopentane ring.  A dominant metabolite found in a rat experiment by
Barnett and Oorough (1974) was oxychlordane V.

    22.4.8'  Other Reactiona

         No processes other than those described have been found to be
important in the fate of chlordane in aquatic environments.

    22.4,9  Microcosm Studies, Field., Stud_iaa__, jnd Modelling

         Sanborn e_t_ _al_. (1976) repotted a terrestrial-aquatic microcosm
study that demonstrates the bioaccuaulation and alow blotransformat ion of
chlordane in aquatic systems.  The ^*C chlordane constituted 94.52,
91.21, 47.62, and' 77.9? of the total l'4C recovered from alga, snail,
mosquito larvae, and fish, respectively.  The ratios of cis:trans isomera
of chlordane In   e organisms were also found to b* different from the 3.04
ratio in the orfjln*l chlord-tne aixture, with the ratios being 4.02 In
water, 3.08 in thi alga,'5.39 in tha snail, and 6.98 in the fish.  The
authors note that th* increase in the cis:trans ratio suggests either more
facile biotransforoation of the trans-isoaer over tha ci§_» or different
accumulation tendencies for the isoners in the organisms.  No metabolites
were identified in the experiment, although some 20 eowpounda were evident
in the tic analysis.  The bioconeeitration factors determined were 9,8 x
lOf for alga, 1.32 x 1C5 for snail, 6.1 x 10^ for mosquito larvae,
and 8.3 x 10^ for fish.

'22,5  Data Suaaary

    Table 22-1 summarizes the data on the aquatic fate of chlordane.

fM  a*
«  u
p«  Ml
jft  *J
4  i
of DM*


•- *


>< e

•0 *.












' JS







i' '
, f




hi l»
ft> 0.







i ••
Jt »
a* &










*„ .

* **
» • «*

X 9*

, ^ *


 22.6   Literature  Cited

 Baker,  R.  D. , and H.  G.  Appltgate.   1974,   Effect  of  ultraviolet radiation
  on  the  persistence  of  pesticides.   Tex.  J.,  Sci,  25( 1-4) : 53-59,

 Barnett,  J,  R.  and H. W.  Oorough.   1974,   Me'tabolisq  of  chlordane in rats.
  J.  Agr.  Food  Chen.  22(*):612~619.

 Benson, W.  R. ,  P. Loobado,  I.  J.  Egry,  R.  D.  Ross, Jr.,  R.  P.  iarron, D.  W.
  Mastbrook, and  E. A,  Hansen.   1971.   Chlordane photoaiteration products;
  Their preparation and  identification.  J, Agric. Food  Chen.
 B*venue,  A.,  and  C.  Y.  Y«o«   1969.  'Gas  chroaatographic  characteristics
   of  chlordafte.   II,   Observed  compositional  changes of  the pesticide in
   aqueous and non-aqueous  environments,   J. Chromatogr,  42:45-52.

 Bowaan,, M. C. » F.  Acr*e, Jr., C.  S.  Lofgren,  and  M.  Beroza.  1964.
   Chlorinated insecticides:   Fate in aqueous  suspensions containing
   mosquito larvae.   Science.  146(3650): 1480-1481.,

 Brooks, G. T.  1*74.   Chlorinated insecticides:   Volume  I:   Technology and
   applications.   CRC Press,  Cleveland, OH., 249 pages.

 Castro, T. F. , and  T.  Yoshida,   1971.  Degradation of organochlorine
 •  Insecticides in flooded  soils in the Philippines.   J.  Agric.  Food  Chem.

• Eichelberger , J.  W.,  and J.  J.  Lichtenberg.   1971.  Persistence of
   pesticides  in  river water.  Environ, Sci. Technol. 5'(6); 541-344.

 Pitchier, M.  H.  and  F.  Kort*.   1969.  Sensitized  and unsensitlzed
   p'wtoisoBerization of cyclodieae insecticides.   Tetrahedron Lett.

 Ginaburg, J.  M.   1953.   Sate  of decomposition of  the newer insecticides
   when exposed to direct sunlight.  Proc. few Jersey Mosquito Extern,
   ASBOC.  40:163-168.

 (.rin«s, D. J. and S.  M. Morrison, 1975,   Bacterial bioconcentration  of
 •  chlorinated hydrocarbon  'insecticides from aqueous systens. Microb. Ecol,
lyengar, L. , A. V. S. P. Rao.  1973.  Metabolism of chlordane and
  heptachlor by aspergillys niger.  J. Gen. Appl. Micr'obiol. i9(4);32l-324.

Knox, J. R. , ' S. Khalifa, G. W. Ivie, and J. E. Casida.  1973.
  Characterization of the phot0iso
Roberts, J. R., A. S. H. De Prietas, ar$d M. A. J., Sidney,  1977,
  Influence of lipid pool size oa bioaccumulati.cn or the insecticide
  vhlordane by northern redhorse suckers (Moxostona macrolepidotum).  J.
  Fish. Kcs. Board Can. 34(1);89-97.

Sanboru, J. R., R. L. Matcalf, W. N, Bruce, and P.-Y, Lu.  1976.  Th€
 ' fate of chlordane and toxaphene In a terrestrial-aquatic model ecosystem,
  Environ. Entomol. 5<3};533-538.

Schimel, S. C., J. M. Patrick, Jr., s>nd J. Forester,  197 6a.
  Heptachlor:  Taxlcity to and uptake by several estuarine organisms.  J.
  foxie. Environ. Health 1:935-965.

Schinael, S. C,, J, Mi Patrick, Jr»» and J. Forester.  1976b.  Heptachlor:
  Uptake, depuration, retention, and metabolism by spot, (Lelostonus
  xanthurus).  J. Toxlcol. Environ, Health 2(1):169-178.

Sovocool, G. W,, R. G» Lewis, R. L. Harieas, N. KP Wilson, and R. D. Zehr.
  1977.  Analysis of technical chlordane by gas chromatography/mass
  spectroaetry.  Anal. Chem. 49(6):734-740,

Vollner, L., H. Parlar, W. Klein and F. Korte.  1971.  Beltrage zur
  okologischen chetaie-XXXI.  Phocoreaktionen der komponenten des
  technischen chlordaas.  Tetrahedron 27:501-509.

Wataaabe, I.  1973.  Decoapostlon of pesticides by soil microorganisms-
  Special emphasis on flooded soil condition.  JARQ.  7(1):15-18.

Weil, U, G. Dure and K. E. Quentin.  1974.  Solubility in water of
  insecticide chlorinated hydrocarbons and polychlorinated blphenyls in
  view of water pollution.  Z. W?.3ser Abwasser Forsch.  7(6): 169-175.

                                 23.  ODD
23.1  Staeement of Frobable Fatc

    The major fate processes for ODD In aquatic environments are bioaccu-
mulation and sorpcion to sediments and biota.  Volatilization-wili also be
an important process for loss of DDD from aquatic systems, with ODD half-
lives on the order of a few days to several weeks*  DDD is quite stable to
chemical transformations in aquatic environments, and biotransformation is
probably the process resulting In the ultimate degradation of ODD in the

23.2  Identiflcatloa
    This chapter considers only the pure DDD; as for technical DOT, the
pp'-isomer is the major component in the technical mixture and most DDD
studies have focused on this Isooer.  Martin (1972) described technical DDD
as having a melting point of 86*C and containing related compounds in small

    The structure, alternate names, and CAS and TSL numbers of DDD are
glvan below:
  Alternate Names

            pj»'-DDD                 .   . '

        CAS Ho. 72548
        TSL No. KI 07000

 23.3  Physical Properties

    The general physical properties of DDD are as follows.

        Molecular weight            '       320
        Melting point
        (Martin 1972)

                           Boiling point                      No data found
f                          Vapor pressure at 30"C             10.2 x 10~7  corr  (pp,')
i                          (Spencer  1975)                     18. f x 10~7  torr  (op1)

»                          Solubility In water at 25 aC
t    •                      (Wtii-etal. 1974}         •        20 ppb Cpp' )
I                          (Blggar and tiggs 1974)*           90 ppb (pp* )
F                                                             100 ppb (op')
i                      "    Log octanol/wattr partition        5.99 (pp1 )
I                          coefficient                        6.08 (op1)
I                          (O'Brien  1974)
                   *Particle  size < 5 ym,

                   23.4   Summary of Fate Data

                       23.4.1  Photolysis

                   .         No data are available  for  estimating  the  photolysis  rat*  of 'ODD in
                   aquatic environments.   Several  papers indicate  that direct  photolysis  of
                   ODD is slowgr than that of DOT, and since  the  half-life  for direct photo-
                   lysis of DDT in water is estimated  to be nore  than 150" years, photolysis of
                   ODD in water should also be very  slow.  No  information is available on the
                   indirect photolysis of  ODD.

                            Hosier e_t al.  (1969) measured a reaction  quantum yield  of 0.04 for
                   photolysis of DDD in hexane solvent at 254  ran;  the quantum  yield for DDT
                   photolysis under this condition was 0.16.   Only one product peak was found
                   by glpc analysis of the DDD photolysis solution and it was-  not identified.
                   Roourn (1963) reported  that DDD on  glass plates Irradiated  for several
                   hours with a geraicidal lamp  (wavelength of 254 nn) gave dechlorination of
                   DDD.  The  lose of DDD In thia experiment was less  than the  losses  of DDE or
                   DDT under  the same photolysis conditions.

         Volatilization of ODD to the atmosphere is probably an important
process for DDT and ODD (Sections 23.4.4 and 25.4.4),  Crosby and Moilanen
(197?) studied the photolysis of DDT in the vapor phase at-wavelengths
greater than 290 nm and found DDD as a minor product (See Section 25.4.1),
the authors stated that under experimental conditions where DDT was about
half photoreacted, DDD "... appeared to be stable to light."  No other in-
fonnation was given on DDD photolysis.

    23,4.2  Oxidation

         So information has been found concerning the oxidation of DDD in
aquatic environments.  Using diphenylmethane as a model for peroxyl radical
oxidation of DDD at the beizylic position with a rate constant of 1.0 M~*'
sec*1 at 30*C, Hendry e_£ «1. (1974), we can calculate a half-life of 22
years using an assumed radical concentration in the aquatic environment of
1Q-9 jjt  Oxidation of DDD by peroxyl radical in the aquatic environment
is then expected to be slow.

    23.4.3  Hydrolysis

         In a study of the hydrolysis of DDT, Wolfe e£ a_l. (1977) estimated
the hydrolysis half-lives for DDD and DDE.  From.structure-reactivity re-
lationships and literature data, they calculated a second-order rate con-
stant of 1,4 x 1Q~3 MT* see"* for hydroxyl-ion-proraoted hydrolysis of
DDD at 27*C,  This value corresponds to a half-life of 570 days at pH 9,
Th«y also estimated a 190-year half-life for hydrolysis of DDD at pH 5 and

         The calculations of Wolfe ejc al. are in good agreement with the
finding of Eichelberger and Lichtenberg (1971), who observed no loss of DDD
(i.e., less, than 2.5%) after 8 weeks for 10-ppb DDD solutions in distilled
water or in raw river water from the Little Miami River in Ohio.

    23,4.4  Volatilization

         Data available for the relative rates of volatilization of DDT and
DDD indicate that DDD is volatilized from aquatic systems at about one-
third the rate of DDT.  Volatilization of DDT from aquatic systems occurs
with half-lives that range from a few hours to several weeks (see Section
25.4.4).  Therefore, volatilization of DDD from aquatic environments will
probably have half-lives ranging from a day to less than a month.

         Singtnaster (1975) described studies designed to measure relative
rates of volatilization of chlorinated pesticides, including DDD from de-
lonlzed water (see also Sections 23*4.4 and 24.4.4).  In one experiment,
about 1 pptr pp'-DDD in 900 ml of water in a 5-liter flask at 25*C was

agitated by bubbling nitrogen through solution at 0,45 1/mln while 4,5
1/ain of air was drawn through the aic apace above Che solution.  The ratio
for relative rates of volatilization of DDE, DDT, and ODD was about 10:3:1,
respectively.  Although the experimental conditions are not directly
applicable to volatilization processes in aquatic systems, they do indicate
that DDD is less volatile than DDT or DDE.   Since there is strong evidence
that volatilization is an important process for loss of DDT and DDE from
water (see Sections 23.4.4 and 24.4.4), volatilization from water is also
probably an important loss process for DDD, with a rate about one-third
that of DOT.

         Ernst (L977) performed experiments to determine the relative vola-
tility of chlorinated orgatiics preparatory  to conducting bioeoncentration
experiments in marine mussels,  la these experiments, aquaria containing
0.5 ppb of chlorinated material were aerated at a rate of 2.5 l/hr for 67
hours*  Although the conditions do not simulate any realistic aquatic
environment, DDD was 81% recovered,, compared with 33% recovery of DDT in
similar experiments; thus, DDD is less volatile from water than DDT. The
ratio of DDT to DDD lost in this experiment was 3.5:1 (ie. , 67% DDT: 19%
DDD), and is in excellent agreement with the 3.3:1 ratio measured by

    '23.4.5  Sorptioa

         Sorption to sediments and biota are important processes for DDD in
aquatic systems.  Although no specific data are available for DDD sorption
to sediment, the importance of sorption processes in the fate of DDT indi-
cates that sorption of DDD to sediment must also be an important fate in
aquatic systems.  In lieu of available data for DDD, sorption data for DDT
will' suffice for some environmental assessment purposes.

         Hoo ejt £l. (1974) examined the DDD content of undisturbed sedi-
ments fton Che Santa Barbara Easia off the  California coast.  By means of
radio dating of th.e sediment layers they determined that DDD began to ap-
pear in the sediments in about 1955 (at 12-ppb concentration) and at-
tained 18-ppb levels by 1976.  These data clearly show that significant
amounts of DDD in aquatic environments must be associated with sediment,

         Although no sediment/water partition coefficients have been re-
ported for DDD, coefficients on the order of 10^ are probable by analogy
to DDT data.  This value la baaed on the correlation between the partition
coefficient for octanol/water and for sorption as described by Kenaga and
Goring (1978).  Since the logarithm of the  octanol/water partition coef-
ficients for DDD and DDT are 5.99 and u.98, respectively (as cited by
Kenaga and Goring), the sorption coefficients should also be similar.
Sorption of DDT is discussed in Section 25.4.5.

    23.4.6 . Bi_oac_cumalacion

         Bioaccumuiation is an important fate process for ODD in aquaCic
systems.  As would be expected, based on the structural similarities be-
tween ODD, DDT and DOE, bioconcentration factors for ODD range from iCH
to 1Q5, similar to those found' for DDE and DDT.  The tragic effects of
bioaccumulation of DDD have also been demonstrated in the Clear Lake inci-
dent described in-Section 23.4.9.

         Ernst (1977) measured a concentration factor of 9120 for upu See of
DDD in marine mussel, but this value is based on a wet weight of Che mus-
sel, and therefore is lower than factors calculated on a dry-weight basis.
When mussel containing 456 ppb (wet weight) DDD was placed in clean water,
only 150 ppb loss was found in 48 hours, indicating a slow elimination from
mussel; other pesticides (i.e., dieldrin, heptachlor .epoxide, endrin) in
similar depuration experiments had elimination half-lives c'iat were two
times faster than DDD.

         Bioaccumulation of DDD has also been studied in terrestrial-aqua-
tic microcosm experiments by Metcalf _et _al. (1971; 1973); data from these
experiments are presented in Section 23.4.9.

    23.4.7  B i o tr anaf ormation and JSiodegradation

  1       i!fo useful data are available for evaluating the biotransformation
of DDD in aquatic systems, although several reviewers of DDT literature
have stated that DDD is more easily metabolized than DDT or DDE.  The per-
sistence of DDD in microcosm studies, and especially in the Clear Lake in-
cident described In Section 23.4.9, does suggest that DDD biotransforma-
tions in aquatic systems are slow.

         In a thorough review of DDT metabolism in mlcrobial systems,
Johnson (1976) also discusses the biotransformation of DDE and DDD.  Where-
as DDE appears to be a rather stable product of DDT fciotransformation, the
formation of DDD from DDT provides for further transformation of the DDT
structure.  Johnsen cites the biological sequence shown on the following
page as a proposed route for DDD conversion to DDCO.

                    H - CHCIj 	». Ar2C - CMC) 	». ATjCM CHjCl

                    ODD          OOMU           DDMS
                 A»2CHCH2OM —-

                     1            DONS           DDNU
                   DO*          OOM            DOBH
                        *r «• p - Cl - C-H              f
                                 8 *           Af2C«0
                                         v     DOCO
Johnsen notes that this scheme Is based on a  synthesis of  information  from
a number of papers, and that no one paper demonstrates this  pathway.   Obvi-
ously, for the above sequence, some cycling of  the DDT residues  through
anaerobic and aerobic systems is required, so that transport  processes
along with the biological variables will determine how fast  the  transfor-
aation sequence occurs.

    23.4.1:  Other Reactions

         No reactions of ODD other than those listed  previously  ire ex-
pected to be important fate processes in aquatic environments,.

    23,4.9  Microcosm Studies, Field Studies, and Modelling

         A f^eld study and Microcosm experiment have  shown that  bioac-
cumuiation of DDD is an important fate process  in aquatic  systems.  The
•icrocosia experiaent also Indicates that ODD  is subject  to some  biotrans-

         fh« pronounced tendency for accumulation of  DDD in  fish and the
western grebe has b««n shown dramatically in  a case wh*re  DDD was applied
to * lake to control gnats (Hunt and Bischoff 1960),  Clear  Lake, a fresh
w»t«r Icke in northern California, was treated with ODD  at levels of 14 ppb
In 1949, and 20 ppfc in 1954 and again in 1957,  Deaths of  numerous grebes
in 1954-1955 and 1957, and the presence of 1600 ppo DDD  in the fatty tissue
of * §r«b«, proopted examination of DDD levels  in the wildlife of the  area.
Analysis of nine fish species In the lake showed concentrations  of DDD in
tdibl« flih flesh ranging from 5.0 to 22-1, ppm, with DDD  Bevels in visceral

fat exceeding 2000 ppm In some samples.   Although no other data on water or
sediment concentrations of UDD were reported, this case history alone is
sufficient Co demonstrate the tendency af ODD to accumulate in wildlife,
and to obvious tragic consequences.

         The fate of DDD in a terrestrial-aquatic microcosm has been re-
ported by Metcalf ej: al, (1971).  Sioconcentration factors of 933, 967, and
6500 were found for DDD in snail, mosquito larvae, and fish; the concentra-
tion in water was 0.4 ppb.  Of the total ^C in these species, fi92, 59%,
and 85%, respectively, was present as DDD; che authors compared these data
to those of microcosm experiments with DOT and D.DE and concluded that DDD
was the most biodegradable of the DDT-1 compounds in their microcosm ex-
periments.  Small amounts of 1,1- bls(p-chlorophenyl)ethane (DDNS, in Sec-
tion 23.4.6) were found as approximately 42 of the total *4C in snail and
fish, along with two unknowns and polar metabolites (as detected by tic,).
Bioconcentration factors in snail, mosquito larvae, and fish of 1.3 x
10*, 3.3 x lO^t and 4.4 x 104, respectively, were also measured for
DDD formed in a microcosm experiment that studied the fate of DDT, but the
authors did not comment on reasons why these data differed from those in
the DDD-microcosm experiments.  Metcalf <*£,£!• (1973) also reported bio-
concentration factors of 8,4 x 10^ and 8.2 x 10-* for fish and snail,
respectively, but the source of these data is unclear.

23.5  Data Summary

    Table 23-1 summarizes the data on the aquatic fate of DDD.

•j   I
•35 i
M «



i *

i ^

J 3

3 M
£ &


















' (e
; «
, a
• I.*
; "3

' u
1 "^
j «M

23.6  Literature Cited

Biggar, J, W., and L. R. Riggs.  1974.  Apparent solubility of
  organochlorine insecticides in water at various temperatures.  Hilgardia

Crosby, D, G,, and K,. W» Moilanen.  1977.  Vapor-phase photodecoraposition
  of DDT.  Chemosphere. 6(4):167-172.

Eichelberger, J. W., and J.(^. Lichtenberg.  1971.,  Persistence of
  pesticides in river water.  Environ. Sci. Technol. 5(6):541-544.

Ernst, W.  1977.  Determination of the bioconcentration potential of marine
  organisms - a steady-state approach.  Chemosphere 6(ll):73i~74Q.

Hendry, 0. G., T. Kill, L. Piszkiewicz, J. A. Howard, and H. K. Eigenraann.
  i974.  A critical review of h-aton transfer in the liquid phase; chlorine
  atom, alkyl, trichlorooethyl, alkoxy, and alkylpero'iy radical.  J. Phys.
  Ch«m. R«f. Data.  3(4):937-978.

Hoo, W., R. ¥. Risebrough, A Soutar, and D  R. Young.  1974.  Deposition
  of DDE and polychlorinated biphenyls in dated sediments of che Santa
  Barbara basin.  Science.  184:1197-1200,

Hunt, E. G.» and A. I. Bischcff.  1960.  Inimical effects on wildlife of
  periodic 01)0 applications to Clear Lake.  Calif. Fish and Came.

•Johnsen, R. E.  1976.  DDT metabolism in taicrobiai systems.  Residue Rev.

Ktnaga, E. E., and C. A. I. Coring.  1978.  Relationship between water
  solubility, soil-sorption, octanol-water partitioning, and
  bioconcentration of chemicals in biota.  Amer. Soc. Test. Mat. Third
  Aquatic Tox. Symp.» Sew Orleans, LA.  63 pp,

Martift, H., ed.  1972.  Pesticide Manual, 3rd, Ed.  British Crop
  Protection Council, Worcester, Engl*

Metca'if, K., I. P. Kapoor, P.-Y. Lu» C. K. Schuth, and P. Sherman.  1973.
  Modal ecosystem studies of th* environmental fate of six organochlorine
  p«»ticides.  Environ. Health Perspect.  4:35-4-*.

Metcalf, R. L., G. K. Sangha, and !. P. Kapoor.  1971.  Model ecosystem for
  th* evaluation of pesticide biodegradibility and ecological
  •signification.  Environ. Sci. Technol.  5(a):709-7' 1.

Hosier, A, R,, W, D. Cueazl, and L. L. Miller.  1969.   Photochemical
  decomposition of DDT by a. free-radical nechaniaa.  Science 164:1083-1085,

O'Brien, R. D,  1974,  Nonenzytnie effects of pesticides on membranes.
  Pages 331-342 in R. Haque and V. H, Freed, eds.   Environmental dynamics
  of pesticides.  Plenum,Press, Sew York, N. Y.

Roburn, J.-  1963.  Effect of sunlight and ulltraviolet radiation on
  chlorinated pesticide residues.  Chem. Ind.  (1963):1555-155C.

Singstaster, J. A., III.  1975.  Environmental behavior of hydrophoblc
  pollutants in aqueous solutions.  Ph.D. Thesis.   Univ.  California, Savis,
  Calif.  I43pp,  University Microfilms, Ann Arbor, Mich., Ordtr Mo. 76-14,
  237.  (Diss. Abstr. Int. B 1976, 36(12, Pt.l):6206-6207).

Spencer, W. F.  1975,  Movement of DOT and its derivatives into the
  atonosphere.  Residue Rev. 59:91-117.

Weil, L., G. Dure, *nd K. E. Quentia,  1974,  Solubility  in water of
  insecticide chlorinated hydrocarbons and polychlorinated tsiphenyla In
  view of water pollution.  Z. Wasser Abw«ts«r Forsch. 7(6):169-175.

Wolfe, N. L., R. C. Zepp, D. F. Paris, G. L. Baughaian, 'and k,  C, Holli*.
  1977,  Methoxychlor and DtT degradation in water: Rates and  products.
  Environ. Sci. lechnol. 11(12):1077-1081.

                                 24.  DDE_

24. 1  Statement of Probably Fate

    The major fate processes for DDE in aquatic environments are bloaccumu-
lation and sorptlort to sediments and biota.  Laboratory studies suggest
that in aquatic environments, DDE may have' volatilization half- lives of
several hours and photolysis half- lives of several days; the observed per-
sistence of DDE In such environments nay be due to ,the fact that DDE is
mainly formed from DOT under biological conditions in which DDE in the
sorbed . stjate is then not available for volatilization or photolysis.

    The ultimate lo*9 of DDE nay be through photolysis in water or in the
ataosphere after desorption or release from biota or sediments; biotrans-
fruTBation may also be an ultimate transformation process, although DDE is
much less susceptible to such processes than DDT or ODD.

24. 2  Identification

    DDE is foraed as a degradation product of DDT and is not manufactured
as a commercial product.  Unless otherwise stated, all references 'to DOS in
this chapter will be for the pp'-DDE isomer, since no data specific to the
op '-DDE isomer were found.

    The structure, alternate names, and CAS and TSL numbers of DDE are as
                                             A It : e mat e names

                                             1, t-Dichloro-2f2-bis
                                               I, l-dichloro*th«ne


          CAS No. 725S9
          TSL No. ?V , '»500

24« 3  PhyjicaJ fropertjeg

    The general physical properties of DDE are a* follows  (isotaer in

         Molecular weight                     313,0

         Melting point              •          88-90'C
         (Leffingwell 1975)

         Boiling .xslnt                        No data  found

         Vapor picsjure at 20°C   •            6.5  x  I'T^  torr  (pp1 )'
         (Spencer 1975)   ,                    6.2'X  10~^  torr  (op')

         Solubility iri water at 25'C
         (Well e_t aj,. 1974)        '           14 ppb (pp')
         (Blggar and Rlggs 1974)*             120  PPb  Cpp')
                                              140  ppb  (op1}
         (Zepp et al. 1977)                   1.3  ppb
         (Chlou et jd, 1977)    ' .             40 ppb (pp1) *i  20*C
         (Metcalf ej. ml.  1973)                1.3  ppb

         Log octanol/water partition          5,69  (pp')
         coefficient                          5.7^  (op')
         (O'Brlan 1974)
*Particl* size < 5 yau

24.4  SjiBMry of Fate Data

    24.4,1  P_hoc_olysls

         Direct photolysis of DDE In aquatic s/stems  will  resul:  fr  half-
lives ranging, fro® about  I day in summer  "o 6  '{ays  in winter;  n^  informa-
tion on indirect photolysis of DDE>  in aq'jatic  environments  has been  found.

         Zepp £££!..  (197ft, 1*77) Measured a q-jantum  yi«ld  of  0.3 for pho-
tolysis of DDE at 313 ma  in water, and c?Lcul=it<»d  that the  direct photo-
lysis half-life of DDE in sunlight at'40"  Jaclfide  ''Ul range  from 0,9 days
in sumoer Co 6.1 days in  winter.  The photolysis  products  Uentit'ied were
l-(4-chlorophenvl)-l-(2,4-dichlarophenyI}-2-rhloroethylene  (o-Ci-L)OMU) ,
1, l-(4-chlorcphenyl)72-chloro«thylene (DDXt),  4-.d  •iichlorohe.izophe-t  >ne
(DDCO); yields of o-Cl-ODMU and DDCO were  :/:  -ind  ;5*, respect i'velv.  The
authors note that the observed persistence of  DDE  in  the environment is
probably due to it* being aorbed to sediments  ,>r  biota where no light for
photolysis is available,  buc  that photolysis of  DOE  is likely  to  be  an
important f*te in the ultimate transf jmat ion  jf  9i"'E  i.? aq'.ntlc syscetss.

                                  H ^  ^Cl

  ^Ac,      c,
     o  Cl  DOMU
         Ocher studies on ODE photolysis h-ive  been, reported.   Singraaster
(1975) measured a sunlight photolysis half-life  of  1.1  days  for 0.84  ppb
DDE In San Francisco Bay water, with only one  unidentified product  de-
tected*^ In contrast to the findings of Zepp e_t_  aJL.  (1977),  Singmaster
found '-"iii'/ a trace of DDCO and the presence  jf -the  other  possible  products
was not determined.  Although Singmaster did not  discuss  whether direct  or
indirect photolysis of DOE was occurring in the  bay  water, the half- life of
1. 1 'day (at an unspecified season) is in good  agreement with  the direct
photolysis half- life calculation of Zepp e_t_ aJL (1977).

         Other studies at higher energy wavelengths  have  also been  re-
ported,  Mosier et_ a_l_. (1969) measured a reaction quantum yield of  0,26  for
photolysis of DDE in hexane at 254 run.  This resullt is in good agreement
with a quantum yield of 0,24 Treasured by Zepp  _et_ ajl_.  (1977^  for DDE photo-
lysis in haxaiie solution at 313 nra.
         The products reported by Zepp ej;_ aj_.  (1977)  have  also been found
by.Kerner ej^ al_. (1972) for photolysis of DDE  at  >  25r<  nm  in various sol-
vents and in che solid and gas phase; ODMU and o-Ci-DDMU were each formed
in yields of approximately 20-40! when DDE was photolyzed  in hexane or di-
oxane-water solvents; the yield'  of DDCO was  about  52.   Photolysis  of solid
DDE for 10 days (with 1 27, DDE remaining) gave  yields  of 40*4  DDCO  and 45%
DDMU with about 1% o-Cl-DDMU; when a solid mixture  of DDE  and chlorophyll
was photolyzed, a 90* yield of o-Cl-DDMU was found.-  Photolysis of DBE in
the gas phase gave yields of 82 • DDMU and 12% o-Cl-DDMU  at  a  20% loss of

         Plimner et_ aJL (197')) reported that photolysis of DOE at  260 n« in
%2"sparged -rethanol solvent gave reductive dechlorl nation  by a 'free radi-
cal mechanise, and with oxvgen present a complex  mixture  of  products was
obtained, including DDCO, DDMU and 3,6-dich lor of luorenone; the latter was
obtained in I'l* yield.  Leffingwell (1975) h-as also reported studies, on the
photolysis of lr.i-op-n suspensions of DDE in air-saturated water at  wave-
lengths below 290 ntn and found a complex mixture  of products.

         Since volatilization of DOE is an import ir.t  fate, the atmspheric
ehesistry of DDE .nay be an important process.   Crosby and  .Voi lanen (1977)

studied the sunlamp photolysis of DOT in the vapor state in air and found
ODE as the major product.  Stjc'ies on DDE under ti-^ same photolysis condi-
tions found'43 of the DDE was photolyzed in 4 days, >/hich was about eight
tines slower than the rate for DDT (see Section 23. •+.!].  This order of re-
activity for photolysis of DDT'and DDE is opposite of Chat found and pre-
dicted for dir'ect photolysis in solution, and ts at present unexplained.
The aajor T'D£ photolysis products were DDCO and DDMv; some products were
tentatively identified as o-Cl-DDMU, and the di-, tri-, and 'te trachlorobi-
phenyls.  In an earlier study on the photolysis of DOT da methanol, Plimmer
jet _ai. (1970) also found aicnloroblphenyl.  As in the case of DOT, the data
provided are not useful for estimating the photolysis half-life of DDE in
the atmosphere,

    24.4.2  Oxidatipn

         No information has been found on the rates of oxidation of DDE in
aquatic environments.  The products reported by Plimmer _et al. (1970) and
l^effingwell (1975) do indicate that oxidation of DDE as a result of photo-
lysis processes Is extensive.

    i1*'**^  Hydrolysis

         Wolf* ejc al. (1977) found DDE was the product of hydrolysis of DDT
for pH values of 3-11.  DDE appears quite stable to hydrolysis undei: the
reaction conditions, with a half-life in excess of 120 years at prt 5 and

         In good agreement with this expected rtability of DDE toward hy-
drolysis, Eichelberger and Lichtenberg (i97i) found no decrease (within
2.52) in concentration after monitoring a 10-ppb DUE solution in distilled
water and in raw river water, respectively, for an eight-week period at
ambient temperatures.

    24,4.4  ¥olatlliz.ation

         Volatilization of DOE Is probably an important loss process in
aquatic systems.  Although only one study of DDE volatilization has been
reportea, this study found that DDE volatilized from distilled water and
natural water so.lutions five times faster than DDT Under identical condi-
tions.  Since volatilisation of OUT from aquatic systems is established as
aa important process (see Section 2^.4,4), volatilization of DDE .nay also
be an important loss process.

         Sittgmaster (1975) described studies designed to measure relative
rites of volatilization of chlorinated pesticide 'from pure water and .sever-
al natural waters.  In these experiments the pesticide at about '  pptr con- •

Generations in 900 ml of water in a 5-liter flask was gently agitated on a
shaker while air was drawn through Che flask (but not bubbled through solu-
tion) at a rate of 4.5 I/rain,  The half-lives for volatilization of pp'-DOE
from pure water and waters from San Fisncisco Bay, the Areericar  B.iver, «-nd
Sacramento River were 0.67, 1.2, 1.4, and 1.9 hour, respectively; the water
loss in these experiments averaged 3.6 + 0.2 g/hr.  Sinamaster concluded
that volatilization of DDE in natural waters would not,he more than two
times slower than in pure water.  Although it is difficult to relate the
experimental results to conditions in aquatic environments (i.e., tempera-
ture, agitation, etc.), the author noted that the air exchange in, the flask
corresponded to a wind velocity of about 10 n/hr, which is much lower than
that usually found in the .environment.  Based on the wind velocity factor
alone, and assuming removal of DDE from the vapor space is the dominant
force in volatilization of DDE, the half-lives for volatilization of DDE in
aquatic environments could be on the order of a few Hours.  In any event,
since the volatilization half-lives of DDT under identical experimental
conditions range4 from 3 to 10 hours, and volatilization of D'DT  is recog-
nized as an important loss process in aquatic systems (see Section 25.4.4),
volatilization of DDE oust also be an important loss process.

   : 24.4.5  Sorgho, on

         Sorptton to sediment and biota are important processes for ODE in
aquatic systems.  Although no specific data are available for DDE sorption
coefficients, information from field studies and, by analogy, to the impor1
tance of sorption processes for DOT, suggest that 'sorptaon -of DDE to se-di-
tnents will be a significant fate in aquatic systems.  In lieu of available
data for DDE, sorption,data for DDT will suffice for some 'environmental
assessment purposes.

         Hamelink and Waybrant (1976) studied the fate of DDE in a flooded
rock quarry; this study, which includes the movement of  'DDE into sediment,
is fully described in Section 24.4.9.  An important pathway for DDE in this
system was sorption to suspended particulates and subsequent deposition in-
to sediment.  The authors also found that DDE in the sediment remained in
the top 1.5-cm layer and did not move into lower sediment layers.

         Hoffl e_c al. (1974) examined the DDE conrant of undisturbed sedi-
ments from the Santa Barbara Basjn off the California coast.  By means of
radio dating of tue sediment layers, they detemined that DDE began to
appear in the sediment in about 1955 (at 24-ppb concentration) ,  and rose to
160 ppb by 1967.  These data cleany show that significant amounts of DUE
in aquatic environments must be associated with sediment.

         Although no sedlaient/water partition coefficients have  been re-
portad for DDE, coefficients on the order of icP are indicated by analogy

co DDT data.  This value is based on the correlation between partition co-
efficient for octanol/wacer and for sorptian as described by Kenaga and
Goring (1978).  Since' the logarithms of the octanol/water partition coef-
ficients for DDE and DDT are 5.77 and 5.98, respectively (,3s cited by
Kenaga and Goring), the sorption coefficients should also be similar,
Sorption of DDT is discussed in Section 25.4.5,

    24.4.6  Bi paccuBiulac i on

         Bioaccutuulac ion is an important fate process for DDE in aquatic
systems; in this respect, DDE is similar to DDT.  Bloconcentration factors
for aquatic species'measured in a field study and in microcosm experiments
are In the range 10-*-10^.

         Haaellnk ee_ a_l. (1977) found the concentration (Y) of DDE in zoo-,
plankton to be directly related to the DDE concentration (X) in water, and
correlated by the equation Y (ppb) = 27.3 * 53.75X (pptr); the data for   ,
this correlation were obtained in the flooded quarry experiment described
in Section 24.4,9, where other data for DDE bioconcentration are also
cited.  Data for bloaccumulation,of DDE obtained in a microcosm experiment
are also presented In Section 24,4.9.

    24.4.7  Biotransfortnation and Siodegradation

         There are no data useful for predicting the rate of biotransforma-
tlon of DDE In aquatic environments.  Several reviews of DOT have described
DDE as being less susceptible for biotransformation than DDT or ODD.  In a
'summary statement in a review of DDT metabolism in microbial systems,
Johnsen (1976) noted that •"...despite some evidence to the contrary...",
DDE ".'..appears to be a stable end product Incapable of being further de-
graded,"  Bohonos.and Francis (1975) reached a similar conclusion, stating
that formation of DDE from DDT was a "dead end pathway" for iDDT btotrans-
formation.  Although the terrestrial-aquatic microcosm experiments of
Metcalf et a 1. (see Section 24.4.9) did rind small amounts of polar metabo-
lites by thin layer chromatography analysis, it is not clear whether the
biotransformation occurred in terrestrial or aqueous phases of the micro-
costs.  Given the rather strong statements of Johnsen and of Bohonos and
Francis regarding the lack of reactivity of DDE, ic appears that DDE will
be very slowly biotransformed in aquatic environments, if at all.  -
    24.4.8  Other Beact ions

         No processes other than those listed above are considered'impor-
tant in evaluating the '"ate of DOE in aquatic environments.

    24.4.9'  Macrocosm..Studies,. Field Studies, and Modelling

         Microcosm experiments and a field study have shown that important
fate processes for DDE la aquatic systems are bioaeeumulatiort and sorption
to sediments.  The microcosm experiments also indicate that DDE is quite
resistant to biotransformation processes.

         Hamelink and Haybrant (1976) reported,a field study on the fate of
DDE Introduced into a flooded quarry where the water, sediment, and biota
were monitored for DDE levels for one year.  The calculated initial DDE
concentration in the epllimnion was 200 pptr, or 50 pptr for the whole
water column.  The DDE in the water column declined rapidly, with 69? loss
occurring in 5 days; the loss was facilitated by sorption to soil after an
influx of soil due to runoff from a rainstorm on the second day of the ex-
ptrinent.  The concentration of DDE In the water column reached an equili-
brium concentration of about 0.5 to 1.0 pptr DDE in four months.  DDE in
th« bottom mud was found to be in the top 1.5 cm, with a maximum of (35 +
27) ppb level found in the 1.5-cm top layer after 81 days.   The equilibrium
of DDE between water and zooplankton reached equilibrium after one day,
with concentration factors of (3-6) x 10^,  ODE reached equilibrium
between water and a resident quarry fish, bluegill, after 60 days, with a
concentration factor of about 1.1 x 10*; the concentration facto/ for
trout 108 days after the intial DDE application wa.f. measured to be 1.8 x

         Based on this study, the authors concluded that sorption to parti-
culate, which subsequently settled, facilitated distribution of DDE
throughout the water column by sorption/desorption processes as well as the
disappearance of DDE from the water.  Sedimented material collected up to
day 21 contained an average of 1.4 ppm DDE, and accounted for 82* of the
DDE lost from the water at this time.   After 173 days, 94% of the DDE found
in their analyses was associated with the bottom murf.  It Is also interest-
ing to not* that at 358 days the DDE concentration In water samples ana
lyzed was roughly twice that at 242 days (i.e.,  1.15-2.73 pptr compared to
0.66-0.96 pptr, respectively).  The authors attributed this result to DDE
release from sediments following decomposition of organic matter during
winter and the higher water temperature, which increased DDE solubility.

         Metcalf et_ a_l. (1971, 1975.) studied the fate of DDE in a terres-
trial-aquatic microcosm experiment.  In the 1971 paper Metcalf et_ £l_.  re-
ported concentration factors for snail, mosquito larvae, and mosquito fish
in the aquatic phase to be 2 x !04, 3 x 10*, and 3 x 10*, respective-
ly; the DDE concentration in water phase was 5.3 ppb.  The metabolites
found in these species were identified as only polar metabolites (by tic),
in the amounts 141, 61, and 3% of the total ^*C in the respective spe-»
cies.  The microcosm experiment was subsequently repeated and reported In
the 1975 paper.  The bioconcentration factors found were 3.6 x 104, 5.9 x

10*, and 1,2 x 10^ for snail, mosquito larvae, and fish; a concentra-
tion factor for algae In this experiment was 1.1 x 10^.  The concentra-
tion of DDE In the water In this latter experiment was 3.8 ppb.  DDE was'
present as 92, 93, 95, and 977, of the total *4C In snail, alga, fish, and
fflosqulto larvae.   Although It Is Impossible to determine how much of the
DDE was aetabolized in the aquatic phase since the DDE was introduced into
the terrestrial phase by several routes' (mainly caterpillar excrement and
leaf fragments),  it is apparent that DDE is not readily biotransformed in
the organisms examined.   The Importance of the food chain in the.bioaccu-
mulation of DDE reported above has also* been suggested by Metcalf et a_l.
(1973), who found concentration factors of only i 10-515 and 300-344 for
mosquito larvae and fish exposed to 0.90 ppb DDE through aquatic exposure
alone In an aquatic microcosm where the food chain did not exist..  An al-
ternative interpretation of these data is that bioaccumulation In a food
chain provides a more rapid uptake of DDE in the mosquito larvae and fish,
and that in th* aquatic  exposure experiment sufficient time was not allowed
for DDE to reach equilibrium between organism and the aqueous phase.

24.5  DataSuamjry                    '

    Table 24-1 summarizes the data on the aquatic fate of ODE.

- 2
t «
«* « 3
*>J ** *
* 5
i f



S *f
1/5 1 i
«*sQ A

*! * t
'» -* fl>

V 1 Ji
u ^2
& H 3*
< *
s • ^ -

'/it ill 2


«* Ml*


11 S 1
II I 1

J s 5 .
» a w
^ -. JK
J 1-5
M* -g M
a — •»
1 i s
> i
a. 9

















J» 3

»* 31
f i
-4 M
a s —
*» 9 1
*• 5. >»
* « *
3 S
«» 3

° I!
A 3t
fc" 3

3 US











24,6  Literature Cited

BIggar,1 J, tf. , and L. R. Rtggs.  1974,  Apparent solubility of
  organochlorine insecticides  in water at various temperatures.   H'lgat'iia

Bohonos , N. , and A, J. Francis.  1975.  Microbiological degradation  of
  military standard pesticide  f armulat ions.   Final  Report.  SRI  :o  the  ;.'.S,
  Army Medical Research Development Command,  Contract -''DADAl?- 73-C-32 14.

Chiou, C. , V. Freed, D. Schmeddtng, and R.  Kohnert.   1977.  Partition
  coefficient and bloaccumulac ton of selected organic chemical';.  Environ.
  Set. Technol.  1 i(5):475-47S.

Crosby, D. G. , and K. W. Moilanon.  1977.   Vapor-phase photodecomposit ion
  of DDT.  Chenosphere, 6(4 ): J&7- 1/2,

Eichelberger , J. W. , and J. J. Lichtenberg.   1971,  Persistence  of
  pesticides in river water.    Environ. Sci. Technol.  5(6): 54 1-544.

Hanteiink, J.  L. , R. C, Waybrant, and P. R.  Yant.  1977.  Mechanisms  of
  bioac cumulation of 'mercury and chlorinated  hydrocarbon pesticides  by  fish
  in lencic ecosystems.  Pages 261-268 _i_n I.  M. Suffet, Fata  oC  pollutants
  in the air and water environment.  Part II.  John Wiley and  Sons,  lew
  York, N. Y.
Hamelink, J,  L. , and R. C.  Waybrant,   1976.   ODE and  lindane  in  a 'ir^e-
  scale model lentic ecosystem.  Trans. Am. Fish. Soc.  105(1 ): 12-'4-1.3i.

Horn, W. , R. W. Rlsebrough,  A.Soutar, and D. R. Young.  1974,   Deposition of
  DDE and polychlorlnated biphenyls in dated  sediments of the  Santa
  Barbara basin.  Science  184:1197-1200.

Johnsen, R. E.   1976,  DDT metabolism  in aicrobial  systems.   Residue Rev,

Kenaga, E. E. , and C. A. I. Gorin*.  1973.  Relationship betwen  water
  solubility, soil-sorption,  rtctanol-water  part itioning, !and  hioconcentra-
  ti*>n of ch*»ical« in biota.  Amer. Soc. Test. Mat. Third Aquatic TOK.'
  Syiap. , Mew Orleans, LA.  .63  pp.
Kerner, I., W. Kl«ln, and F. Korte.   1972.  Seitrag* ZMT
  chemie-XXXIII.  Photochemisehe reakt ionene von  1, l-dichlor-2(p,p'~
  diehlorphenyl) athylen (DDE).  Tetrahedron 28:1375-1578.

Leffingwell, J. T.   1975,  The photolysis of DDT  in water.  Ph. p. Thesis.
  University of California, Davis.  124 p.  University Microfilms, Ann
  Arbor, Mich., Order Mo. ?6-14, 218.   (Digs. Abstr, Int. 3

Metcalf, R. L,, I. P. Kapoor, P.-Y. Lu» C. K. Schuth, and P. Shennan.,
  1973*  Model ecosystem studies of the environmental fits of six
  organochlorlne pesticides.  Environ.  Health Perspect.  .4:35-44.

Metealf, R, L., G, K. Sangha, and I. P. Kapoor.  1971.  Model ecosystem for
  the evaluation -of pesticide biodegradability and ecological magnifica-
  tion.  Environ. Sci. Technol. 5{S);709-713.

Metcalf, R. L., J. R. Sanborn, P.-Y. Lu, and 0. Nye.  1975.   Laboratory
  model ecosystem studies of the degradaion and fate of•radio-labeled trl-,
  tetra-, and pentaehlorobiphenyl compared with DDE.  Arch.  Environ,
  Contaa. Toxicol.  3(2):15i-lfc5,

Hosier, A. R., W. D. Guenzi, and L. L. Miller.  1969.  Photochemical
  decomposition of DDT by a free-radical mechanism.  Science 164:1083-1085.

O'Brien, R. D,  1974.  Nonentynlc effects of pesticides on membranes.
  Pages 331-342 _iti R. Haque and V. H,  Freed, eds. Environmental dynamics
  of pesticides.  Plenum Press, N«w York, N. Y.

Plimmer, J. R., U. I. Klingbiel,  and  S. E. Hummer.  1970.  Photooxidation
  of DDT and DDE.  Science.  167:67-79.

Singnaster, J. A,, III.  1975.  Environmental behavior of hvdrophobic
  pollutants in aqueous solutions.  Ph.D. Thesis,  Univ.  California,  Davis,
  Calif,  143pp.  University Microfilms, Ann Arbor, Mich., order No.
  76-14,237.  (Diss, Abstr. Int. B 1976, 36(12^, Pt ,1) ;6206-6207.

Spencer, W. F.  1975.  Movement of DOT and its derivatives into the
  atmosphere.  Residue Rev, 59:91-11?,

Weil, L., G. Dure, and SC. E. Quentin,   1974,  Solubility in water of
  insecticide chlorinated hydrocarbons and polychlorinated biphenyls  in
  view of water pollution.  Z, Wasser  Abwasser Forsch.  7(6):169-175,

Wolfe, S. t., R. G. Zepp, 0. F, Paris, G., L. Baughman, and  R. C. Hollis,
  1977.  Methoxychlor and DDT degradation in water:  Rates and products.
  Environ. Sci, Technol. 11(12):1077-1081.

Zepp, R, G., N. L. Wolfe, L. V. Azarraga, M. H. Coz, and C.  W. Pape.   1977.
  Photochemical transformation of the  DDT and methoxychlor degradation
  products, DDE and DMDE, by sunlight.  Arch. Environ, Contain. Toxicol.

Zepp, R, G., S. L. Wolfe, J. A. Gordon, and R« C. Finchsr.  1976.
  Light-induced transformations of methoxychlor in aquatic systems.  J,
  Agric, Food Chea. 24(4):727-733i

                                 25.  DPT

25.1  S ta t egien tof P robaj> I e _f a_t_e

    There is ample evictecce to demonstrate chat DDT is very persistent in
the 'environment.  The dominant fate prac^sses In aqtiacic environments are
volatilization and sorpcion to bioca and sediments,, with the importance of
sorpcion being'determined by the amount of suspended participate available
'in the water body.  The ultiaate transformation of DDT in the aquatic en-
vironment is probably by b-iotrans formation, although one study indicates
that, indirect photolysis aay also be a significant loss processs for OPT in
a ".atural water, with a photolysis half—life on the order of a week.  Pho-
tolysis of DDT in the ga)» phase has also been reported, but since DDT has
been widely found•throughout the biosphere, atmospheric transformations
appear to be slow.  There is also abundant evidence to demonstrate  that
bioaceumulation of DOT is a significant process in the environment.

    The following evaluation of the fate processes for DOT in aquatic sys-
tems is based on an abbreviated literature search; in several cases, lit-
erature reviews on DDT have been used without i. critical review of  the
original literature, whichi for some proces-"«s is extensive.  The rationale
for this limited search is that experience has already clearly demonstrated
the persistence and bioaccuinulation of DDT in the environment, and  another
review is not neeaed.  Thus, although the literature cited aay be incom-
plete, omitted information is not contradictory Co1 conclusions presented in
this chapter.

25.2  Ideati f iea t ion

    This chapter discusses the fate in the aquatic environment of the indi-
vidual pp'- and op'-DDT isoaers.  Technics* DOT comprises mainly a  mixture
of two DDT isomers; on« reference gives a range of DDT isoners In three
technical mixtures as 70 to 73 percent of the pp'-isomer, 12,to 21  percent
of the 'op'-isomer, and a snail amount (0.012) of the oo'-isom«r (Gtmther
and Gunther 1971).  Cunther (1945) described a technics! DOT preparation PS
containing approximately 702 pp1, 13X op', and 61 oo'.  The structures,
alternate names, and TSL and CAS numbers of the two major isomers are as



 CAS No. 50293
 fSL NO, XJ 33250

CAS Jte. 789-02-6
TSL Mo» Won* assigned
                                         Other Common Haiats


In this chapter. Information on a particular Isoner will be so noted
whenever available; in many Instances, however, the isoaer Is not specified
and will only be identified as HDT.

25.3  Physical^Properties

    Th« general physical properties  of the DDT iscraers are given below:

    Molecular weight                        354,5.

    Melting point                           108,5-109,0"C (ppf}
    (Gunther and Guncher 1971}              ,74-74.5"C (op')

    Boiling Point     '                      185"C (pp1)
    (Gunther and Gunther 1971)

    Vapor pressure
    (Martin 1972)                           1,9 x 1
25.4  Suanaary of Fate/ Data

    25.4, I  Photolysis

         Direct photolysis off DDT in aqueous solution La very slow, wttrt a
half-life of probably greater than  150 years.  Natural substances  in «>or>e
aauatlc environments aay cause indirect photciysis processes to be  Import-
ant for DDT transformation, with half-lives on the order of a few  days or
possibly even hours for DDT loss.  Half-lives for indirect photolysis of
DOT are difficult t^> predict for general environmental assessments  because
of lack of information on the variability of natural waters to produce such
indirect reactions through photosensitized, photo-initiatad free radical,
or ocher reactions.

         Z«pp *_£ a_l. (1976) calculated that DDT will probably have  a direct
photolysis half-life of more than 150 years. ' This estimate was, based on a
measured uv-absorptiort spectrum of DDT and assuming a reaction quantura
yield of unity, the maximim quantum yield value expected for photolysis of
*' chemical in dilute solution.  This calculation was made in conjunction
wich extensive studies on the DDT analogue, snethoxychlor (I).  The  reaction
quantum yields measured for photolysis of I in hexane -ind in water •solvents
w«r« 0.12 and 0.3, respectively,   Tha authors also found that although the
sunlight photolysis of methoxychlor in distilled wat»r had * half-lit? of
over 300 hours, raethoxychlor photolyzed in three natural waters had
half-lives ranging from 2 to 6 hours at midday in Mav; i.i two other natural

waters no photolysis of I was found after 2 hours exposure to Midday sun-
light.  Sunlight photolysis of I In distiHed water containing 20s ppn of a
commercial "humic acid" had a half-life of 7.3 hours.  Since DOT and
tnethGxychlt>r have similar structures and apparently analogous photareac-
tions involving the CH-CClj group (see discussion below), Indirect
photolysis of DDT in natural waters may also be rapid and .ta variable as
that found for methoxychlor,

         Singtnaseer (1975) studied she sunlight photolysis of 0.90 ppb DDT
in distilled water and in water from San Francisco Bay.  After 7 days expo-
sure to sunlight, the DOT Concentration in distilled water was unchanged
where as the DOT in the bay water was 30% lost.  The authors also deter-
mined that n_o DOE, DOO, or the related photoproducts were £orta*d in the
reaction, a finding that conflicts with other literature reports (set be-
low).  They suggested that trana formation in solution nay occur through
photodissociation of the CH-CClj bond.

         B«eaus* indirect photolysis of DOT appears to be * rapid and pos-
sibly Important fate of 00T In aquatic systems, further discussions on the
photochemistry of DDT are useful, especially regarding the formation of the
DOE and DDD photoproducts.  Since similar products have b«en reported to be
formed by direct and indirect photolysis of DOT, both processes will b« re-
viewed; however, only the laf.er is Important In aquatic environments. Sow*
of the compounds formed by DOT photolysis are given b«low with th«ir
acronyms:           <

         Se%'«ral studies  have  provided  Information on the direct photoly-
sis of DDT at wavelengtns less than  290 nm.   A reaction quantum yield of
0.16 was measured  for  photolysis  of  00T at 25^ nm in hexane solvent: (Hosier
££ *JL< 1969),  DDD and HC1 were the  only products identified, and the .yield
of DDD Increased from  5X  in hexane as  the sole solvent Co l*i when 1M n~
butylmereaptan was  present; DDE was  found to undergo photolysis more rapid-
ly than DOT under  the  reaction conditions a.rtd , therefore, DDE was. not found
as a product.  The  authors did find  DDD, DDE, anci DDCO when DDT fiias ex-
posed to air were  paotolyzed at 254  nra. Plitnmer e_t al. (1970) found ODD and
DDMU as the products when DDT  in  ^-sparged methanol solvent was'photo-'
iyzed At 260 na«   The  photolysis  of  DOT at 280 no in methanol with oxygen
present gave a complex produce mixture, In which the aechyl ester of 2,2-
bls(p-chioroph«nyl) acetic »eid was  found.  Leffingwell (1975) found 16
produces when a IG-ppm suspension of DOT in air-saturated water was pho-
tolyzed using a latap with an output  below 290 tan.  Other studies of pho-
tolysis of DDT on  soil and as  films  have found 0DE and DDCO among other
products (Baker and Applegate  19?Q,  1974;  Fleck 1949; Roburn 1963).

         Ail these  results are in accord with a mechanism of initial photo-
dissociation of a  CCl2~Cl bond to give  the radical I, which then may (i)
io«« a chlorine atom to give DDK, (2)  react with a hydrogan atom donor to
give DDD, or (3) react with oxygen to give the peroxyl radical II, which
«ay undergo further• reactions  to  give a complex set of products (Plimmer e_t
al. 1970; Hosier e_t al. 1969),  It- is  interesting to note that Leffingwell
               Id - C1OI jCH , CClj  '•••••»• (0ClO!»CH CC!-- * Ct-
                                   • W5f )	."y.».DOCOI
                            CM OM  OTMtR ^BOOUCTS IMCLUOtNG
                              f 	{9. c»J,CHC . Cl

(1975) also found CC14 and CHClj as DOT photolysis, product 5, although
it was not determined whether thesei prociucts resulted  by  direct  photolysis
of the C-CCi-j bond or by Indirect processes.

         Several examples of indirect photolysis of  DDT are also  known;
however, except for two studies, the reaction condir.io.i3  are difficult  to
relate to environmental aquatic systems.  As d iscussed , e,ir 1 ier ,  Singsaster
found no DDE, ODD, or related products when DDT was  ohotoly^ed  in sunlight
in a bay water.  In apparent contrast to .this observation,  Zepp  et al.
(1976) found ,the dehydrohaloger.ated methoxycnlor analog of  DDE  (II)  as  a
product when methoxycnlor was photol/?ed  in several  natural fresh-waters.
         Leffingwell (1973)  reported, that  the  photolysis  rate  of  a  10-ppm
suspension of DDT In water in sunlight or  a  Hg  lanap  ^as accelerated  by the
presence of tr iphenylataine,  dipnenylanthracene,  or a free radic.il initia-
tor, azo-bis-isobutyronitrile.  So detailed  product  information  -'as  re-
'ported, but he did corataent chat the  products >'e re  rhe  3'-in?e --is  chose  found
for direct photolysis (presumably UDE, HDD,  DDMU,  DDCO, etc.)  Millar and
Xarang (19"0) 'reported earlier that  photolysis of  DDT  in  cyclohexan«j sol-
vent at 310 am occurred only when triphenylamine or  ^,S-diethylaniline
(N'NDA) was present; DOU,  DDE and DDCO were formed  in^ol,  151,  and t-6%
yields -hen NN'DA was at initial concentration  of l-j~^M.   The aatiiors also
determined that triple t-*state-sensitized processes wer-s aot responsible for
the indirect photolysis of DDT in these systems.

         Ivie and Casida  (197la) also found'that triplet  sensitizer  chemi-
cals codeposited with DOT  on silica  gel plates were  generally  ineffective
in promoting the photolysis  of DDT in sunlight.  Although carba^ole  and
triphenyla«i,-ie did promote the pnotolysis  of DOT in  these studies, the ex-
tent of the JDT photolyscs were not  correlated with  the triplet  er.ergies of

Che 21 sensitizers used.  Ivie and Casida (19/"lb) also found that :ripheny-
laoine catalyzed the sunlight photolysis pate of DDT applied to bean
leaves, but only DIE and DDCO were found as che photoproducts,,.  Earlier re-
ports by Linqulst £t al. (1946) and Ginsburg (1953) indicated that DDT in
media such as kerosene, fuel oil, wettable powders, and emulsions was also
degraded In sunlight.

         Although these examples show that indirect photolysis of DDT does
occur, the mechanisms and the relevance of these studies to photolysis of
DDT in aquatic, systems are generally unclear.  Singtnas.ter' § finding that no
ODE, ODD, or related photoproducts were formed in bay water, compared with
the findings of Zepp ejS ai. that the DDE analog was a major photolysis
product of nethoxychlor in several frtshwaters, may only reflect different
transformation mechanisms available in the different natural waters.  Since
both DD0 and DDE are persistent pollutants of concern (see Chapters 23 and
24, respectively), further studies on the indirect photolysis of DDT in
natural waters seen advisable.  In addition to transformations via radicals
I and II, photooxidation of DDT at the benzylic position may also provide
DDCO and the trichlorociethyl radical, the latter leading to OlClj or
CC14 as found by Leffingwell:
                                                    DDCO * CCI.
         Because volatilization of DDT into rhe acaosphere is an important
environmental process (see Section 25.4.4), Crosby and Moilanen (1977)' ex-
amined the photolysis of DOT in the gas phase using light of wavelengths

 greater Chan 290 nm.   After  4  days,  32% of  che DOT had  been  photolyzed  to
 give a 14:t ratio of  DDE:DDD,  with a product  balance  of  96  percent.   Thus,
 although DDE Is  then  expected  to be the male?" product in the gas  phase  pho-
 tolysis of  DDT,  these data are not useful for predicting the photolysis
 half-life in the gas  phase.   Since DOT has  been found in atmospheric  samp-
 ling studies, photolysis  in  th? atmosphere  is apparently slow.

     24.4.2   Oxidation

          No information is available o.1 oxidation of  DDT under  conditions
 relevant to aquatic environments.   Using diphenylraethane as  a model for
 peroxyl radical  oxidation of DDT at v<*i benzylic position with  a  rate con-
 stant of 1.0 JT1 sec  ~l at 30"C,  (Hendry e£ a_L.  1974), we can calcu-
 late a half-life of 22 years using an assumed radical concentration in  the
 aquatic environment of ICT^M.   Oxidation of DDT by peroxyl  radical in the
 aquatic environment is then  expected to b*  slow.  Hoffman an
    25.4.4  Volatilization

         The worldwide atmospheric distribution of DDT clearly indicates
that volatilization of DDT from soils or aquatic systems (or both) is an
important process.   Laboratory studies and field studies have shown, that
volatilization of DDT from water is a rapid processs, although no useful
data are available  for quantitatively evaluating the rate of the volatil-
ization process in  aquatic syr;tc;?s.  Although several authors have sug-
gested that volatilization of DDT will be facilitated if DDT is concen-
trated in the microlayer of a water body, no sound evidence for this pro-
cess has been found.   The information available does indicate, however,
that DDT may be volatilized from water with half-lives of less than a week..

         Singmaster (1975) described studies designed to measure relative
rates of volatilization of chlorinated pesticides from pure water and
several natural waters.  In these experiments about i pptr concentrations
of pesticide in1900 ml of water in a 5 -liter flask were gently agitated on
a shaker while air  was drawn through the flask (but not bubbled through
solution) at a rate of 4,,5 l/min«  The half-lives for volatilization of
pp'-DDT from pure water and waters from San Francisco Bay, the American
River, and Sacramento River were 3.9 hour,,6.5 hour, 6.0 hour, and 10 hour,
respectively; the water loss in these experiments averaged 3.6 HK 0.2
g/hour.  Frop these experiments, Singtnaster concluded that volatilization
of DDT in natural waters would not be more than two times slower than- in
pure water. Although the experiment is difficult to relate to conditions in
aquatic environments (i.e.* temperature, agitation, etc.), the author noted
that the air exchange in the flask corresponded to a wind velocity of about
10 n/hr which is much lower than that usually found over aquatic systems in
the environment.  Based on the wind velocity factor alone and assuming re-
moval of DDT from the vapor space Is the dominant force in volatilization
of DDT, the half-lives for volatilization of DDT in aquatic environments
could then be on the order of a few hours.  Mackay and Molkoff (1973) and
Maekay and, Leinonen (1975) have calculated volatilization half-lives for a
series of chemicals based on equations for mass transfer in an Idealized
aquatic system.  A half-life of 3.1 days was calculated for DDT volatiliza-
tion,, and the autnors note that for DDT the volatilization rate Is deter-
mined by the DDT concentration gradient in the vapor phase (Mackay and
Leinonen 1975).

         Acree e_t_ al. (1963)'reported studies on volatilization of DDT from
aqueous solutions as a follow-up to earlier studies that reported more than
a 50" loss of a 10-ppb DDT suspension in water after the solution re-
mained for 24 hours at room temperature.  Although the authors described
the loss process as "codistillation" of DDT with water, Spencer (1975)
pointed out that the "codistiliation" term is inappropriate, because It
connotes an additive vapor pressure relationship between water and DDT that
                                   25- 10

is only applicable at the boiling point of a solution and not at en-
vironmental temperatures.  Water and DDT, In fact, vaporize from solution
independent of each other.  Acree e_t al. did find, however, that the weight
loss of ,ODT per weight loss of water was on the order of 3 to & percent at
25t>-3G°C for initial DDT concentrations of 100 to 0.36 ppb.  In all experi-
ments, more than 502 of the initial DDT was lost in 24 hours; insufficient
experimental information or discussion is available to determine whether
the losi of DDT was completely due to volatilization or how the data ob-
tained may be relevant to ,environmental aquatic systems.

         Other studies have also implicated volatilization as an important
loss process.  Haaelink et al. (1971) studied the fate of'DDT applied at
concentrations of 5-15 ppb to natural and artificial ponds, and attributed
a 90 percent loss of the applied DDT after 30-40 days Co volatilization
(see Section 25.4.9).  Oloffs and Albright (1974) reported that DDT losses
from solutions incubated for 12 days in two natural waters were 20 t6 50
percent, with some of th« DDT recovered from the glass wool plugs stop-
pering the flasks; no DDT was lest in water samples incubated in glass-
stoppered flasks.  The authors also found no loss of DDT from volatiliza-
tion in water containing sediments, although some DDT was converted to DDD.
They concluded that these experiments with sediment demonstrated the compe-
tition of sorption to sediments versus volatilization that exists in
natural waters.

         A more general discussion of the movement of DDT and its residues
(DDD and DDE) into the atmosphere was given by Spencer (1975); this review
covers a few papers on the volatilization of DDT from water and the more
extensive literature on volatilization of DDT from soil.  In a discussion
of DOT in the atmosphere, Bidleman and Olney (1974) sugggest that the atmo-
spheric DDT is in the vapor phas,e rather than associated with particulate
matter, because the difference' between the amount of some materials in par-
tlculates (e.g., lead) found over Rhode Island and Bermuda was a hundred-
fold greater than the difference in DDT concentrations in the air samples.
Woodwell et al. (1971) estimated that the mean residence time for DDT in
the atmosphere is 4 years.

         With regard to the preceeding paragraph, it should be noted that
if DDT is In the vapor phasft in the atmosphere as suggested by Bidleman and
Olney, DDT should be rapidly oxidized by hydroxyl radicals in the-atmo-'
sphere; data! of Hendry and Kenley (1979) indicate that the half-life should
be on the order of a few days or less.  There is no information to deter-
mine whether this oxidation half-life is .underestimated, whether atmo-
spheric DDT is actually on particulars, or whether the residence time in
the atmosphere of 4 years estimated by Woodwell and his associates is too
long.  Although this discrepancy cannot be resolved at this time, it does
show the problems- and u»«fuilness- of approaching fate evaluations using
several different methodologies, which in the case of DDT unfortunately do
not give the same conclusions.

    25,4,5  Sorj)ticm

         The sorption of DDT to suspended sediments and bottom muds has
b«en well established by analyses of environmental samples;  fe«r quantita-
tive studies of DDT sorption onto suspended partlculates in  water are
available.  Sorption to sediments is an Important process for DDT in
aquatic systems, however, with partition coefficients of 10^ to 10?
found for some soils suspended in aqueous solutions.   Kenaga and Goring
(1978) cited a value for K0c of 2.38 x 105.

         Weil et_ ajL. (1973) measured the socptlon of  DDT by  an aqueous sus-
pension of sodium humaCe and by a'soil containing 1.4* hucntc substances;
Che 1/n and K vlues for the Freundlich isotherm plots were 0.7 and 1.1 x
10*, respectively, for the suspended sodium hunate, and 0.7  and I x
107, respectively, for the soil.

         SMfl et_ aJL, (1970) also measured K values for sortlon of DOT from
aqueous solution to thne tolls; no data were reported for the 1/n
Freundlich parameter.  The K values for sandy loam, clay, and muck soils
were 1.3 x 103, 1.4 x 10*. and I.I x iO5 respectively.  Plcer et_ al.
(1977) measured the sorption Isotherm for sorption of DDT from seawater
onto several oceanic sediments containing 0.84X to 0.51% organic carbon;
th« values of 1/n and K were 2 and 10^, respectively.

         Huang and Liao (1970) reported Freundlich Isotherm  data for sorp-
tion of DDT on three clays? the 1/n parameters they calculated are given
below and the K values have been recalculated to make K unitless (X/n In
yg/gm and C in Ug/ral):

        Sorpt ion System               K	    '      •      1/n

      MontmorilUnite (clay)      9.0 x IO12                6.0

      Kaolinite (clay)  .          1.2x IO10                5.1

      IlHte (clay)               1.9 x IO7                 3.3

The authors did not content: on the unusually high 1/n values measured for
these experiment*; the partition coefficient(s) are exceptionally large,
probably because of the large 1/n value.  The authors did note thac the
equilibrium between water and clay was reached within several hcurs.  \
subsequent paper by Huang (1971) reported that 50 and 300 rag/1 concentra-
tion* of glucose in solution did not affect the amount of DDT taken up by
the aontnorillinite clay.  The authors also noted that the DDT was not
readily deiorbed from monttaorilllnita.  The amount of DDT taken up by the
clays was greater than the uptake of heptachlor or dieldrin  for the re-
spective clays.

    25.^.6  Bipaccuinulat ion

          The  blaaccuir.ulati.on  of  DDT  in  various  species  in the  biosphere is
well established.   Various studies have found  bioconccnrration factors  'for
DDT that  range up  to  10^  in aquatic  systems.   DOT in concentrations  up  to
hundreds  of ppnt have  also been  found in analyses  of  numerous  environmental
samples as a  result of  direct uptake, sorption  to biota,  and  bloaccumula-
tion in food  chains.  Bioaccumulation is undoubtedly an important  fate  of
DDT In aquatic systems.

          The  bioconcentrattoti and  distribution  of DDT in  the  environment
has been  thoroughly reviewed  by  Revenue (1976), Kenaga  (1972), and  Edwards
.(1970.), and need not  be  extensively  reviewed  in this report,   Biaconcentra-
tion factors  (BCF) cited  by Kenaga range up to  10* for  species in  aquatic
systems;  some BCFs in literature appear to  be  low, possibly due to  DOT
being present in the  aqueous  phase as a suspension and  therefore not avail-
able for  true equilibrium between water and organisms.  Terrestrial-
aquatic Microcosm  experiments have found BCFs  for aquatic organisms  ranging
from IG3  to  105 in various species (see Section 25.4.9).   Metcalf  et_
al. (1973) found that the amount of  DDT in  rtosquita  fish  exposed to  DDT via
water and food chain  was  250  times greater  in  a 33-day  microcosm experiment
compared,with a 3-day aquatic exposure  alone,  but this  difference was prob-
a-bly because  the exposure time  in.  the 3-day experiment  was too brief for
maximum DDT uptake and  equilibrium to be established.

    25.4,7  Biotransformation and  Biudegradati'oh

          The  biotransformation  of  DDT,  and  of  its derivatives  DDE and DDD,
has been  extensively  studied  in  a  number of biological  systems,, and  is  the
subject of thorough review by Johnsen (1976).   N?o data  are available, how-
ever, to  reliably  assess  the  rate  of DDT transformation in aquatic  environ-
ments.  Any Quantitative  evaluation  of  DDT  transformation is  made more
difficult by  the finding  of Johnsen  that no microorganism has  been  found  to
utilize DDT as the sole  carbon  source,  and  that only cometabolic trans-
formations of DDT  are knowr.  Vast literature as  well as  the widespread
environmental occurance  of DDT  clearly  indicates  that DDT is  not readily
metabolized in aquatic  environments, although  biotransformation is
undoubtedly an important  process in  t.he ultimate  loss of  DDT  from the
environment.  Biotransformation  of DOT  occurs more readily under anaerobic
conditions than in aerobic systems';  transformation pf. DDT to  DDE is  favored
•in aerobic systems, whereas DDD  is the  major metabolite in anaerobic en-
vironments.   In some  aerobic  experiments, both  DDE and  DDD have b*«»n found
as metabolites.

         It Is difficult to state whether ODE or ODD is the primary DDT
metabolite In the environment since the products formed are obviously In
part determined by the environment where metabolism or chemical transfor-
mation occurs; trie issue Is also complicated because ODD undergoes further
metabrlisn in the environment more rapidly than DDT itself or DDE.  Some1
papers note that the ultimate transformation of DDT to DDCO via ODD re-
quires cycling through anaerobic and aerobic systems,  so that metabolism
and transport via sorption/desorption will be required for total DDT de-,
gradation.  For DDT in aquatic systems, however, sorption of DDT to sus-
pended particulates and subsequent 'deposition into anaerobic sediment sys-
tems appears to be the dominant process.

         Johnsen (1976) thoroughly reviewed DDT metabolism in mierobial
system*.  The review cover* DDT metabolism in undefined microblal popula-
tions in soils, sewage, sediment, silage, water, and digestive systems;
tran*formations in defined microbial populations of bacteria, fungi, and
alga* and mixed microbial populations were also reviewed.t  Johnsen con-
cluded that the major route of DDT metabolism is through DDD, which can be
subsequently degraded to ODCO or bls~(p-chlorophenyl)-metl»ane«  ODD forma-
tion is favored by anaerobic conditions, although DDD has also been found
to occur along with DDE in aerobf.c conditions.   Johnsen further notes that
DDE is very stabl* to further metabolism; there is not evidence to indicate
that DDE is reduced to 000.  He also notes that no microorganism has been
found to utilize DDT as the soie carbon source but rather that cometablism
of DDT occurs.
    25.4,8  Qther Reactions

         Stutter ('977) reviewed studies of the reaction of DDT to DDD with
eaphasls on the role of seta! ions in the biological transformation of
DDTs,  His review includes reactions of iron-containing systems, and cobalt
systems as well as chrotntum- and tine-centered systems.  The studies de-
scribed include in vivo, in vitro, and model studies.   Although most of the
J8«tal-DBf Interactions described for cobalt or iron are part of anaerobic
biological systems, the possibility exists that release of the complexes
from biological systems may result in abiotic reduction of DDT in the
aquatic environments if the reduced metal species are  stable In the aquatic
environment.  No information is available, however, to assess the Impor-
tance of such abiotic processes in aquatic systems compared with anaerobic
biological processes for reduction of DDT.

         Castro (1964) and Miskus e£ ajL. (1965) showed that porphyrlns re-
duce DDT to DDD under anaerobic conditions.  Glass (1972) found that

 ferrous  ion  reduced DDT  to ODD and that  the  reduction of DOT was  facili-
 tated in soils with higher organic content where  ferric ion could  be  rer-
 duced by easily  oxidlzable organic matter,

          In  studies of the reduction of  DDT  by vitamin B»2» Berry  and
 Stotter  (1977) showed 'that hydroxycob(II)ala«in is capable of reducing DDT
 to ODD under anaerobic conditions; the cobalc(II) compound was  foraed by
 reduction of the cobalc(III)  form with carbon monoKide,  The authors  do
 note, however, that while cobalamin can  then effect  reduction of DDT, cyan-
 ocobalamln—th«  form in  which 8^2 ^s usually available—is not  reduc- ,
 ible  to  the  Co(II) state,

          Ross and Biros  (1970) demonstrated  that  DDT acts as an electron
 acceptor molecule to form ff-complexes with several alky laced aromatic
 molecules In chloroform  and carbon tetrachloride  solvents.  Although  such
 interactions nay also occur with natural materials such as humic acid in
 aquatic  environments, there Is no information to  indicate whether  such
 processes are important  in the environment,

     25.4.9  Microcosm'Studies, Field Studies, and Modelling

          Several field studies and a microcosm experiment have  been re-
 ported that  demonstrate  various aspects  of the fate  of DDT in aquatic sys-
,terns. Unfortunately,'these reports do-not provide an integrated study that
 can be used  to evaluate  all the dominant piccesst^ occurring in these
 aquatic  systems.

          Hatnelink et_ aJL.  (1971) studied  the  fate  of  DDT introduced Into a
 natural  pond and several artificial ponds.1  DDT concentration in  the water
 dropped  from 15  ppb in the natural pond  and  5 ppb in the artificial pond to
 fflinimua  detectable  levels of  about 0.02  ppb  in 30 to 40 days.   A material
 balance  at this  cine found that at.least 90% of the  initial DDT was not
 present  In components of the  ponds (water, sediment, algae, invertebrates,
 and fish), and the DDT was presumed to have  been  lost by volatilization.
 The DOT  or DDT residues, designated DDT-R, In the water were present mainly
 as DDT in the first 30 days,  as both DDT and ODD  in  the second  30  days, and
 primarily as ODD in the  last  30 days.  The1authors also determined that
 lass  than 25% of the DDT applied to the  artificial ponds was deposited In  •
 or an the, sandy  bottom of the ponds.

          The accumulation of  DDT-R in  the  other components of the  ponds was
 also determined.  DDT-R  were  accumulated by  the algae present,  and depended
 on the anount of DDT-R present .in the water.  The concentration of DDT-R in
 invertebrates rose  rapidly in initial  stages of the  experiment  and reached

equilibrium within 5 dayg.  The DDT-R concentration then declined,as Che
DDT-R content of the water declined.  The DDT-R content of fish, on che
other hand, rose rapidly and attained levels of about 12 ppm.   A somewhat
similar study by Bridges e_t_ a_l_. (1963) suffered from the uncertainties of
the analytical methods at that time, but tha results followed  the same
trend, except that accumulation of DDT-R in fish was riot as dramatic as in
Hanelink's study.  In Bridges' study, DDT-R levels fell below 1 ppb within
2 v.ays after application of 20 ppb DDT to a farm, pond, with subsequent
DDT-ft levels generally exceeding 1 ppm in pond vegetation for  8 weeks'and
in pond sediments for 3 weeks.  DDT-R levels were also monitored in pond,
fish for 16 months.  DDT levels showed,a general trend for loss of DDT to
trace or undetectable levels after about 9 months; both ODD and DDE were
present at levels varying from 2 to 0.4 ppm, and neither compound appeared
to be the major DDT metabolite/product in the pond.

         The fate of DDT has also been examined in terreatrial-aquatic
microcosms.  Metcalf et_ a_l_, (1971) found that DDT in this system after 30 <
days gave DDE concentrations in water, snail, moscuito larva*, and fish
that exceeded the DDT levels; some ODD was also formed in this system.
Since the pathway of DDT (or DDT-t) in this system involves application to
plants that are ingested, in part, by caterpillar before the DDT-R are
transported into the aquatic phase of the microcosm, we cannot make firm
conclusions on the transformation of DDT in the aquatic phase  alone.  From
the data obtained in this study, bioconcentration factors of 3.4 x 10 ,
8.2 x 10^, and 8.5 x 104 were calculated for snail, mosquito larvae,
and mosquito fish, respectively (Metcalf et_ al. 1973).  In a study in an
identical Microcosm, Booth (1975) found DDT was still the major DDT-R, with
bioconccntraclon factors of  1.1 x 1CP for snail, 930}for mosquito larvae,
2.6 x 104 for fish, 5.9 x 103 for algae, and 6.2 x 10*5 for daphnia.
In Booth's microcosm axperiment the amount of ODD relative to  DDE in the
various organisms was slightly larger than in the experiment by Metcalf et/
ajU , and this difference in products as well as the difference, in biocon-
centration factors is probably due to the difficulty in reproducing such
•icrocosa experiments between laboratories.

25.5  Data Summary                     '

    Table 25-1 summarizes the data on the aquatic fate of DDT.

                                                                          at  /Uiudtlc  Fate of WiT

                                                                                                  Halt-                 umtlJ                iLJ»d

                          fttal*>ly:»i3            IHite^s. p!iolt>lydi*  i;*          -                 >ISO y««r^               HL||t

                                                iyKis Bdiy  b« iivpitrt Jatt,

                          U»lil^£iutt             llut MI t*t|K*rC                                              c«rt«tn *.onaili«»s.                            81  d»y» pit ">

j|^                        Vakat illzAtiua       1» aa  inport^nt               -                 tt:w tiaucu Cu              Law

                          Bl^I tan»t.)(»ut lJ»l/   (a an  iapurl^nl

                                              -  lass of UOT.
                               f  i*  Lnsu(Jtcigent  luJurmJii.Ki la  the  iLwU'wc-J  1H*;tai-.in' to
                               iib it?  I <»t»j,

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  Insecticide chlorinated hydrocarbons and pol/chlorinated biphenyls in
  view of water pollution.  Z. Wasser Abwasser Forsch.  7(6): 169-U5,

Wolfe, N.  L. , R. G. Zepp, D. F. Paris, G. L. Baughman, and R. C. Hollis.
  197?-. • Methoxychlor and DDT degra^adon in water: Rates and products.
  Environ. Sci. Technol. 7(1 10): 1077-1081 .

Wood well, G, M.» P. P. Craig, and H. A. Johnson.  19/i.  DDT  in the
  biosphere:  Where does it go?  Science 174:1101-1107.

Zepp, R. G., 8. L, Wolfe, J. A. Gordon, and k. C. Fincher. . 1976.
  Light-induced transformations of methoxychlor  in aquatic systems.  J.
  Agric. Food Chenu 24(4); 727-733.

                               26.  DIELDRIS
26. 1  Statement of Probable Face

    The literature inforaacion as well as analysis of numerous environment-.
cal samples Indicates that dieldrin is persistent In the environment.   The
important fate processes in aquatic environments are sorption  to  sediment,
bloaceumulation, and vqlatilization; the latter process may have  half-lives
of several hoi * to several days in some aquatic systems.  Although  direct
photolysis of dieldrin in water is slow (t^..1? ~ 2 months), photosen-
sitized processes may result in photolysis if sensitizers  are  available in
aquatic environments.  Altv,ough dieldrin is quite resistant to blotratisfor-
rnatiin, this process will probably be an important fate for dieldrin  'in
sediisent and biota.

26.2  Id_e njci_f lea t ion

    This section considers only the fate of the pure chemical  dialdrin.
The structure, alternate names, and CAS and TSL numbers of dieldrin are
given below:
               Ct                             '            ,    '

    CAS S"o. 60-57-1
    TSL No. 10 17500
   2,3 ,40, lO-Hexachloro-
   6,7-epoxy-l,4,4a ,5,6,7,8,-
   8a-octahydroexo- 1 ,4-endo-
   5 , 3- dime than o- naphtha lene
 Compound -97
 Oc t a 1 ox
 26,3  Physical Properties

    The general physical properties of  dieldrin are  as  follows;

         Molecular weight                     381
         Melting point
         (Martin J972)

         Boiling point
• N'o data found
                                   26- I

         Vapor pressure at 20"C*
         (Martin 1972)
         (Spencer and Cliath 1969)
1.78 x 10"7 torr
2,8 x  IQ~6 'torr
         Solubility in water
         (Park and Bruce 1968)
         (Biggar and Riggs 1974)f
         (Weil et_ aj.. 1974)
         s(Shavnagary and Jayaran 1974)

         Log octanoI/water
         partition coefficient
186 ppb at 25-2
material balance in the reaction  (O.!14 -t-.0.32 « 0.46 ppn),  the  half-life  •
for direct photolysis of dieldrin  Is 2.1 or  1.8 months',  respectively.   Al-
though some of the initial concentration of 0.51  ppm dieldrin must  have
exceeded the solubility and have  been in suspension., these  data  are the
best available for estimating the  direct photolysis'half-life of  dieldrin
in aqueous solution.

         Photodieldrin has also been reported to.  b* -a product frois  photo-
lysis of dieldrin^ in the solid state at wavelengths below 300 nm.
(Robinson ec_ a I. 1%6; Benson 1971; Rosen e£ alt  1966).  Benson  et_  aj..  al: j
reported that I was formed in yields of 7 percent after  3 weeks  and 25  per-
cent after 2 months for solid films of dieldrin exposed  to  Sunlight.  Pho-
tolysis of. dieldrin at < 300 nm in the solid phase under a  current  of
oxygen also gave a small yield of  !• but most of the residue was  unreacted
dieldrin or polymer (Gab e_t_ ajt.  1974),  This information, although  not
relevant to aquatic systems, does  demonstrate that photodieldrirr is a com-
monly found product of direct photolysis of dieldrin,

         Ivie and CasIda (1970,  1971b) also reported that photodieldrin is
a product >f the rotenone-sensitized photolysis of dieldrin on  bean leaves
exposed to sunlight.  Triplet sensitizers codeposited with  dieldrin on
silica gel plates and exposed to sunlight also produced  I (Ivie  and Casida
I97la).  Rosen and Carey (1963) prepared I in 75  percent yield  by the ben-
zophenone-sen«it'ize'd photolysis of dieldrin in benzene solvent.  Although
these studies clearly show that phorodieldrin can be formed by photosensi-
tized processes, there is no information on what  sensitizers are present in
aquatic environments chat could also effect such  reaction,  or what  the
half-lives of' such, processes may  be.

         In several studies of the photolysis of  dieldrin at wavelengths
below 300 nm, products were reported in which dieldrin was  dechlorinated at
the vinylic position (Benson 19.71; N'agl e£ _a_l.  1970 ;• Henderson and  Crosby
1967.)  The latter authors have shown, however, that such dechlorinated
products are not likely to be formed at wavelengths above 300 nm in the
solar region.

         Since volatilization of dieidrin from water and soil is an Impor-
tant process, Crosby and Motlanen  (1974) studied  the photolysis  of  dieldrin
in the vapor phase.   The only product obtained when dieidrin was photolyzed
with -3 sun lamp was photodieldrin,  and no data are available to estimate how
fast dieldrin will be photolyzed  in the atmosphere.

    26.4.2  Oxidation '

         N'o information is available on the oxidation of dieidrin in aqua-
tic systems.  The highly chlorinated dieldrin structures with the bridge-

head positions should, however, be unreactive toward free radical oxidation
under environmental conditions.  Hoffman and Eichelsdoerrer (1971) found
chat dieldrin was unreaccive coward reaction with ozone in hexane and in
water-acetone solvents,

    26.4.3  Hydrolysis

         A study by Eichelberger and Lichtenberg (1971) indicates that hy-
drolysis of dieldrin in aquatic environments is probably very slow.   They
reported that 1U ppb of dieldrin in distilled water and in a sample  of raw
water froa the Little Miami River (Ohio) was 100 percent recovered after 8
weeks at room teaperature; since the recoveries were rounded off to  the
nearest 5*, less than 2.5% loss of dieldrin had occurred.  A 2,5 percent
loss after 8 weeks corresponds to a half "life of 4 years; hence, the hy-
drolysis half-life of dieldrin is greater than 4 years in aquatic environ-

    26.4.4 ' VolaCi jJ^zar. ion

         Volatilization of dieldrin from aquatic systems into the atmos-
phere is probably an .important process with half-lives on the crder  of a
few hours to a few days.

         •Singoaster (1975) described studies ,designed to measure the rela-
tive rate of volatilization of chlorinated pesticides froa pure water and
several natural waters.  In these experiments the pesticides at approxi-
mately 1 'pptr concentrations in 900 ml of water in a 5-liter flask were
gently agitated on a shaker while air was drawn through the flask (but aot
bubbled through solution) at a rate of 4.5 1/min,  The half-lives for
volatilization of dieldrin from pure water and waters froa San Francisco
Bay, the Aaerican River, and Sacramento River (in California) were 7.7 hr,
6,1 hr, 9.0 hr, and 8.5 hr, respectively; the water loss in these experi-
ments averaged 3,6 + 0,2 g/hr.-  Singmaster concluded that volatilization of
dieldrin in natural waters would not be more than two times slower than in
pure water.  Although the total experiment is diffcult to relate to  condi-
tions in aquatic environments (e.g., teoperaturt, agitation), the author
not«d that the air exchange in the flask corresponded to a wind velocity of
about 10 a/hr, which is much lower than that usually found in the environ-
ment,  ttased on the wind velocity factor alone and assuming removal  of
dieldrin from the vapor space Is the dominant force in volatilization of
dieldrin  the half-lives for volatilization of dieldrin in aquatic environ-
ments could be on the order of a few hours Co a few days.

         Hackay and Wolkoff (1973) and Mackay and Leinonen (1975) derived
equations to predict volatilization of Jow-solubility organic chemicals
from aquatic environments.  -Assuming a reasonable value for e-vaperation of

water from an aquatic system and using physical-chemical data for dieldrin
at 25°C, a half-life of about 1.5 years wa's calculated.  The significant
disagreement between this value and that-suggested by the data of
Singoiaster is probably due ta ehe failure of Mackay's calculation to
Include the contribution of mass transfer in the gas phase for removal of

    26.4.,5  Sorption

         Sorption of dteldrin to s^dlnents containing significant amounts
of organic material will probably be an important fate in aquatic systems,
especially since transformation processes for dieldrin a're slow.

         Weil *££!_. (1973) reported Freundlich isotherm data for sorption
of dialdrin to hunic acid and a soil at 15 *C.  The 1/n and K values for the
huaiic acid~dieldrin sorption experiment were 0.79 and 228, respectively;
for Che sorption to soil, the 1/n and K values were 0.64 and 1.6 x  10*,
respectively.  The  latter sorption coefficient suggests that sorption of
dieldrin to sediments will be appreciable.  Boucher and Lee (1972)  studied
the uptake of dieldrin from water on aquifer and silica sands at 5aC; equi-
librium between water and sand was reached within several hours.  A  i
Freundlich isotherm sorption partition coefficient (K) of 2-5' was estimated
from isotherm plots presented in the paper.  Such low coefficients  are not
unexpected for particulates low in organic naterial.

         The sorption of dieldrin to several clays was reported by  Huang
and Liao (1970) and by Huang (1971; 1974).  As discussed in the chapters on
heptachlor and DDT, the Freundlich isotherra parameters 1/n and K in these
studies were unusual, in that 1/n exceeded 3 instead of being unity  as i«
usually found.  In the 1971 paper the 1/n values for sorption of dieldrin
ranged from 8,8 to  11^6; therefore, the corresponding K values are  so large
as to be meaningless in the normal comparison and use of K values where 1/n
- 1.

    26.4.6  B t oac c uau la ti i on

         Bloeoncentration factors for dieldrin in various organisms range
fr'ita 102 to 10^, indicating that dieldrin will show moderate to signi-
fic'ant bioaccumulat ion in various species present in aquatic systems.

         Grin»es and Morrison (1975) have reported dieldrin concentration
factors for 13 bacterial species that ranged from 90 to 2.3 A 10"*;  6
species had  factors between 90 and 700, with the other 1 bacteria  having
factors of 1.1 x 103, 1.2 x 103, 1.5 x 103, .1.0 x 101, i.7 x 1C'3,
1.6 x I04, and 2..8 x 10*.  Meudorf ard Khan (1975) '-sund a
concentration factor of 1,2 x 10^ for a fresh water algd ffxposed to
dieldrin for 3 hours.

         Data front aieroeosta experiments also suggest similar ranges of
bioconcentration factors.  Metcalf et_ al_. ..(1973) found bloconcentration
factors of 4.57 x 102, 6.2 x 104, ,and 2.7 x 103 for an alga, snail,
and fish, respectively. • Sanborn and Yu (1973)'found concentration factors
of 7,5 x 1CH, 1.1 x 10*, and 6.1 x 1(P  for the same species in their
exp.ertmer.ts.  Ernst (1977) found a bioconcentration factor of 1.6 x 10-
for accuznulation of dleldrin by a ialt water mussel, with a half-life for
dleldrln elimination of about 50 hours when the mussel was transferred to
clean water.    '                                                  ' f

    26.4.7  Btotransformatton and Biodegradat_ijon

         Several authors cits dleldrln as being one of the more ronbiode-
gradable chlorinated pesticides, and although several papers have reported
blotransformat Ion of dleldrin, no useful data are available for estimating
blotfansformat ion rates in aquatic environments.

         Bohonos and Francis (I97S) and Sanborn et_ a_l. (1977) reviewed the
literature on biotransformatlon of dleldrin and stated that dieldrin is one
of the more persistent chlorinated pesticides.  In studies comparing the
rates of blotransforaa:Ion of chlorinated pesticides (including OUT and
metabolites, HCH Isoaers, endrin, heptachlor, and raethoxyehlor} in flooded
and nonflooded soils, Watanabe (1973) and Castro and Yoshida (1971) found
both dieldrtn and chlordane to be unique in that they were persistent In
both types of soils.  Hill and McCarty (1967) stated that dieldrin was more
stable than li'ndane, heptachlor, endrin, DDT, aldrin, or heptacblor epoxide
In anaerobic sewage sludge,  Matsumura and Soush (I967) found that the ma-
jority of 577 mleroblal isolates from soil incubated in nutrient solutions
for 30 days showed no capacity for biological transformation of dieldrin.
These .qualitative observations are difficult to relate to aquatic environ-
ssents, but they do indicate that dieldrin blotrans format ions will be very

         The persistence of dleldrin has also been shown in terrestrial-
aquatic aiicrocosta experiments,  Sanborn and Yy O973) found that of the
I C applied as dleldrln in the terrestrial compartment of the microcosm
an average of 97X of the ^C extracted from aigae, clam, crab, daphnia,
mosquito larvae, fish, and snail was present as dleldrin.  Trace amounts of
ff.etaoolites were found, aaong them 9-hydroxydleldrin and 9-ketodleldrin.
Metcalf ft. al_. (I973) also studied the fate of dieldrin in a terrestrial- ,
aquatic microcosm, and found dleldrin1present as 88, 96, and 95% of the
total ^*C recovered from alga, snail, and fish, respectively; the small
amounts of metabolites found were the sane as those found by Sanborn and

    26.4.8  Other Reactions

         No react tons other than chose described above are considered  im-
portant fates for dieldrin in aquatic systems.

    26,4.9  Hicrocosn. Studies, Fie:ldJatudieSj, =and; Model ling

         The microcosm experiments at Metcalf e^_ ajU  (1973) and  Sanborn  and
Yu (1973) have -demonstrated the stability of d'.eldrin toward 'Motransforma-
tion and Its propensity for bioaccumulation ir several aquatic species  (sea
Sections 26.4.6 and 26.4.7).

26.5  Data_ Summary•

    Table 26-1 summarizes the data on the aquatic fate of dleldria.

                                                                              Tufcle ife-l   -

                                                                         ut Aqua11<> F«t* of  Ulalttrln

                    Knv i fuiwfc.iil.il
                    Vol. 11 tliiallon
                    tit k
                                t iuli
                                                                I y
                                                         S I •t
UirvCt  (jtu'totys I 3
be  t«j»ji i jiu .
Hut  411

Nut  an l«|K-it.iu>C |)iuc«;>i.

is at)
                                                         wry slaw,  but Kiy  be
                                                        t^ndlc lw&s»  prucv^s*  tit
                                                                                             K.i 10
                                                                Life t>j

                                                               2  months
                                                                                                                     K'W  l
                                                                                                                                                    Si- J ma.
                 j.  n.fii-  i-.  litsui t li i«nt  tuturiiMt iuit in
                                                                     L' ii'ui-,1
                                                                                          lu permit  4»«H8»eiit at  » muttl  pixl-ut It

26.5  Literature; Cltg_d

Benchmark,  1975,   Draft of Preliminary Summaries of Literature Surveys  of
  Benchmark Pesticides.  George Washington University Medical Center.
  Draft dated October 30,  1975.

Benson, W, R,   1971.  Photolysis of solid and dissolved dieidrin.   J.
  Agrtc. Food Chem. 19(i):66-72.

Bhavnagary, H. M., and M.  Jayaratn.  1974.  Reterainstion  of water
  solubilities of lindane  a-nd dieidrin at different temperatui .... •  BuM.
  Grain Technol.  l2(2):95-99.

Slggar, J. W., and L.  R. Riggs.  1974.  Apparent solubility of
  organoehlorine- insectictdes in water at various temperatures.  '-Ulgardia

Bohoncs, M. , and A. J. Francis.  1975.  Microbiological degradation of
  military standard pesticide formulations.  Final Report.  SRI -to  the
  U.S. Army Medical Research Development Command.  Contract Mo.

Soucher, F. R., and G. f. Lee.  1972.   Adsorption of lindane and dieidrin
  pesticides on unconsolidated aquifer sands.  Environ. Sci, Technol.

Castro, F., and T. Voshida.   1971.  Degradation  of organoehlorine
  insecticides in flooded  soils in the Philippines.  J. Agric. Food Chem.
  l?Co): 1168-1170,

Crosby, D. G., and K.  W. Moilanen,  1974.  Vapor-phase photodecomoosit ion
  of aldrln and dielarin.  Arch.  Environ. Contam. Tuxicol. 2fl}:62-74.

Elchelb*rger,  J.  W., and J, J. Lichtenberg.  1971.  Persistence of
  pesticides  in river water.  Environ.  Sci. Technol. 5(6 ):54I~544.

Ernst, W.   1977.   Determination of the bioooncentration potential o; ,
  .marine organisms - a steady state approach.  Chemosphere 6{11>:731-740,

Gabf S. , H. Parlar, S. Nltz, K. Hustert, and F.  Korte.  1974.  Beitrage
  zur okologischen chemie.  LXXX1. ' Photochemischer abbau von aldrin,
  dieidrin and photodieldrin als festkorper im sauerstoffstront.
  Chenosphere 3(3 )j183-186.

Grimes, D. J., atid S.  W, Morrison.  1975.  Bacterial bioconcentration-of
  chlorinated hydrocarbon  insecticides from aqueous systems.  Microb.  Ecol.

 Henderson,  G.  L.»  and  0,  G,  Crosby.   1968.   The photodecoraposition of
   dieldrin  residues  in water.   Sull.  Environ.  Contaau Tox'col.
   3(3): 131-134.

 Henderson,  G.  L. ,  and  D.  G.  Cros.by.   1967.  .Photodecomposition of dieldrin
   and  aldrln.   J.  Agric.' Food  Chera.  15:888-893.

 Hill,  D.  W.,  and P.  L.  McCarty,   1967,   Anaerobic degradation of selected
   chlorinated hydrocarbon pesticides.   J.  Water Pollut. Control Fed.
   39(8): 1259-1277.

 Hoffaann, J.,  and  0. Elchelsdoerfer.   1971.   Effect of ozone on
   chlorinated-hydrocarbon-group pesticides in water.  Voia Wasser

 Huang,  J.-C.   1971.  Effect  of selected factors on pesticide sorption and
   desorption  in the  aquatic  system.   J.  Water Poilut. Control Fed.
   43(8): 1739-1748.

 Huang,  J.-C.   1974.  Water-sediment  distribution of chlorinated hydrocarbon
   pesticides  in various environmental conditions.  Proc, Int. Conf, Transp.
:   Persistent  Chem. Aquat. Ecosyst.   11:23-30.

 Huang,  J,-C.,  and  C.-S. Llao.   1970.   Adsorption of pesticides by clay
   minerals.   J.  Sanit.  Div.  Am.  Soc.  Civ.  Eng. 96(SA5):1057-1078.

 Ivie,  G.  W. ,  and J.  E.  Casida,  1970.   Enhancement of photoalteration of
   cyclodiene  insecticide  chemical residues by rotenone.  Science.
   167(3295): 1620-1622.

 Ivie,  G.  W.,  and J.  E.  Casida.  197la.   Sensitized photodecoraposition and
   photosensitizer  activity of  pesticide chemicals exposed to sunlight af
   silica  gel  chroaatoplates.   J.  Agric.  Food Chen. 19:405-409.

 Ivie,  G,  W.,  and J,  E,  Casida.  197lb.   Photosensitized for the
   accelerated degradation of chlorinated cyclodienes and other i
 Martin,  H.,  ed.  1972.   Pesticide  Manual,  3rd  Ed.   British  Crop Proteetic-n
   Council, Worcester,  Eagl.             •

 Matsuraura, F,, an,d  G.  H,  Boush.   1967.  Dieldrin:   Degradation by  soil
   ojicroorganistas.   Science  156(3777) :959-96U

 Metcalf,  R,  t.,  I.  P.  Kapoor,  P.-Y,  Lu» C.  K.  Schuth,  and  P.  Sherman.
   1973.   Model ecosystem  studies  of  the environmental  fate of six
   organo-chloriae  pesticides.   Environ1, Health Perspect, 4:35-44,

 Sagl,  H.  G., W.  Klein, and  F.  Xorte.   1970.   Betraige  zur  okologischen
   eheaie-XXVill.  Ube<" das  peaktlonsverhaltcn von  dieldrin in losung1 und  In
   der  gasphase.  Tetrahedron 26;53l9-25»

 N*udorf,  S., and M. A. Q. Khan.   1975.  Pick-up and metabolism of  DDT,
   dieldrin,  and  photoaldrln  by a  freshwater alga  (Ankis_t_rodesaus
   Aaalloides) and  a nicrocrustacean  (Daphnia  Pulex) Bull.  Environ. Contam.

 Park,  K..  S., and W. N. Bruce.   1968,   The determination  of the water
   solubility of  aldrin, dieldrin,  heptachlor  and heptachlor epoxide. J.
   icon.  Eatomol.l6I(3):770-774.

 Robinson, J,, A. Richardson, S. Bush,  and K.  E. Elgar.  1966.  A
   photo-isomerization  product  of  dieldrin.  Bull.  Environ. Contara.
   Toxicol.  H4):12I-132.  •

 Rosen, J. D,, and  W. F, Carey,  1968,   Preparation of  the  photoisooers  of
   aldrin.and dieldrtn. J.  Agric.  Food  Chera..  16(3);536-537,

 Roien, J. D,, D, J, Sutherland, and  G.  ft. Lipton.   1966.   The
   photochemical  isomerization  of  dieldrin and  endrin and effects on
   toxicity.   Bull.  Environ, .Contain.  Toxicol,  1(4); 133-140,

 Sanborn,  J.. R,, 8. M. Francis, and  R.  L. Metcalf.  1977.   Th* degradation
   of itlected pesticides  in  soil:  A review of the published  literature.
   U.S. NTIS, PB  R*p,,  PB-272353,  633pp,

 Sanborn,  J.  I. and  C.  Yu-.   1973.   The  fate  of  dieldrin in  a tsodel
   tcosystem. Bull. Environ. Contain. Toxicol.  10(6);340-346.

 SlngnaSter,  J. A,,  III.  1975.  Environmental  behavior of  hydrophobic
   pollutants in  aqueous solutions.   Ph.D. Thesis.   Univ. California, Davis,
   Calif.   l*3pp.  University Microfilms,  Ann  Arbor, Mich., Order Xo.
•   76-14,237. (DUs. Abstr.  lit.  B .1976,  36(12, Pe.I):620fc-6207).

Spencer, W. F,, and M. M. Cliath.  1969.   Vapor density of  dleldria,
  Environ. Sci. Technpl.  3(7};670~674.

Watanabe, I.  1973.  Decomposition of pesticides by soil microorganisms -
  Special emphasis on flooded soil condition.  JAJRQ 7(1):15-18.

Weil, L. t G. Dure, ami K. E. Quentin.  1973.   Adsorption of chlorinated
  hydrocarbons to organic particles and  soils.  Z. Wasser Abwasser Forsch.

Weil, L., G. Dur'e, and K. E. Quentin.  1974.   Solubility in water of
  insecticide chlorinated hydrocarbons and polychlorinated  biphenyls  in
  view of water pollution,  Z. Wasser Abwasser Forsch.  7(6);169-175.

27,1  Statement of probable Fate

     Data are inconple.e, regarding the  Important  processes  for  determining
the fate of endoaulfan  in aquatic systems.  The hydrolysis half-life  of
endosulfaa at 20 °C Is about a month at pH  7 and about 6  months at  pH  5.?.
Other information suggests that photolysis, oxidation,  biodegradation,
sorpeion, and volat ilziation mav be occuring under some  envlronaental
conditions, but data for predicting the  rates  and  relative Importance of
these processes in aquatic systems are not available.

     ^o data have been  found on the transformation or transport  of  endo-
sulfan sulfate,,  Hydrolysis could be the most  important  process  for endo-
suifan sulfate, but no  hydrolysis data have been found.

27.2  Idtntificatim

     Endosolfan and its derivatives exist  in two sceroisooeric 'forms, the
a- and the 3- forms. '
     The structure, CAS and TSL numbers, and nomenclature  for  endosulfan
and endosulfan sulfate- are as follows:

                                          Alt e_rnat e Jjarnes  (Ct and  8  isomers)

                                          5-Sorbornene~2 , 3-dimethanol-
                                             1,4 ,5,6,7 , 7-hexaeh lor o~ ,
                                             cyclic sulflte
                                          6 , 7 , 8 , 9 , 10 , 10- Hexach lor o- 1 , 5 ,
                                             5a, 6,9, <>a-hexahydr 0-6,9 ,-
                                             oetharjo-2 ,4 ,,3- benzo(e)di-
                                            cyclo-(2.2. l>-
                                            xywethy lene-5 ,6-sulf ate
                                          Beosi t
             a-£ido>sulf 9 n

             CAS S'o.  115-29-7
             TSL Mo.  RB "*2750


CAS So. 115-29-7
TSL Mo. R3 92750
Endosulfan sulfate

CAS No, 1031-07-8 ,.
TSL So. Mone assigned
                                     Alternate Names
                                     Sea page 27-1.
                                     Alternate Names
                                     None found.

 27,3   Physleal  Properties

      The  general  physical  properties  o;  endosulfan (a-  and  3-  isotners)  and
 endosulfan  sulfate  are  as  follows.
 Molecular  weight
Melting point  ,
(Phillips e£ a_l. 1975) Technical      70-100°C
(Ali 1973)             l Endosulfan  .108-110*0
(All 1978)             3 Endosulfan  207-2C9°C
B .lllng point

Vapor pressure
(Phillips et_ al.
(Martens 1972)
                                       Mo data  found
                                                         Endosulian Sulfate
                                                              JJo data found
                                     9xlO"3 torr at BQ'C
                                     IxlO"5 to"r ac 25*C
 Solubility  in water
 (All  19/8)
 (Phillips  1975) at  223C,  pH
'(Phillips  1975) at  203C,  pH  5.5
 (Weil et_ al,  1974)  at  25*C
 (MRI  1977)
0. 164 ppni
0. 15 pptn
0.26 ppm
0.530 ppia
0.6 ppin
0.070 ppm
0.06 ppm
0.10 ppm
0.280 ppa

0.117 ppn

0. 22 pp«
 Log octanol/water  parrHlon  coefficient  -  (s)  3,55
   (All  1973)                               (3)  3.62
                                           (sulfate)  3.66

 27.4   SuBmary  of  Fate  Data

      27.4.1  Photolvsis

           There are no datai useful  for evaluating the  rate  of  photolysis  of
 endosulfan or  its  sulfate  in aquatic  environments.   Endosulfan photolyses
 in solution and solid  state  have  been reported at wavelengths  > 300  nm
 using laboratory  light sources  and  filters.  Since endosulfan  sulfate has
 been  reported  as  a product  from photolysis of  endosulfan, a review of   en-
 dosalfan photochemistry  is  required.  The  uv spectrum  of  endosulfan  re- ,
 ported  in literature (Gore  «_ ,aU  1971)  13  insufficient  to  evaluate  the ab-
 sorption above 290 nre  except that  the absorption coefficients  must b«i less
 than 800

          Several researchers have reported that a-endosulf an is photo-
isomerlzed to s-endosulfan In hexane (Schumacher e£ al. 1971;  Putnam et_
ai. 1975) and in aqueous suspension (MRI 197^).  Since no quantitative data
were 'given and laboratory lamps were used ,in these photaiyses, it' Is
impossible to determine how important this Isomerizat ton will be in the en-
vironment, or whether an equilibrium mixture of the isooers Is retained be-
fore other photoreact ions occur,
          Schumacher et_ a_l_. (1971, 19/4) reported that irradiation of
dosuifan at > 300 n« in various organic, mixed organic, and water-organic
solvents gave different individual or seta of dechlorinated endosuifan
products.  Gas phase photolysis of 8-endosulfan produced endo^ulfan eiher,
diol, sulfate, and la'ctone as well as the dechlorinated ether and a-isoicer
(See section 27.5 for structures).

          Endosuifan diol was reported to be the product when an emulsion  ,
or either endosuifan isoaier in water was irradiated at 250-580 nra (MRI
1977);  dechlorinated diols were also found.

          Schuphan and Ballschaiter (1972) found that diol was formed when
either endosuifan isoner was irradiated in alkaline aqueous methanol.  When
irradiated in neutral solution, however, the isomers gave different (uni-
dentified) products.  It is significant to note that no sol. it ion photoly-
ses of endosulfaa have been reported to give endosuifan sulfate as a pro-

          Harrison et_ ajl. (1967) reported endosuifan sulfate as the only
product detected from sunlight photolysis of endosuifan -on apple leaves.
These authors also conducted experiments to test the variables in the sul-
fate formation and concluded chat uv irradiation and a moisture-containing
substrate were necessary for sulfate formation.

          Archer et_ a_l. (1972) studied the photolysis of thin films of both
tsoraers on glass plates;  GE germicidai laaps with output >.300 nm were
used for irradation.  Endosuifan diol was the major photolysis product,
with minor amount* of tndosulfan ether, lactone, a-hydroxyether , and
several other unidentified products formed.   Endosuifan sulfate was not
formed in these phocolyses and was found to be stable to photplvses under
these conditions.  Srbsenuently, Archer (1973) studied the fate of en-
dosuifan on alfalfa dried in sunlight, uv light, and air (in the absence of
light U  The total residue of endosuifan and products decreased 'rapidly in
the first several days and then decreased slowly thereafter,   in the air-
dried dark control sample, however, endosuifan appeared to be quantita-
tively converted to endosuifan sulfate after the initial decrease.   Koshy
(1972) studied the persistence of endosuifan on glass dishes and sweet
potato leaves exposed to sunlight and assayed for residues by a ^-mortal-
ity test with insects.  For the endosuifan exposed on leaves for 2 and -
days the insect mortality was 15 and OX, respectively; at B and 12 days

the mortality was ^3 and  IS,  respectively,  for the  etidosulfan exposed o;j
she glass dishes.

          In summary,  the  above  Information Is confusing *nd  incorap lets  .is
to what photolysis  processes  and  products raay  be  expected from the environ-
mental photolysis of endosulfan.   Because of  rhe  different  li^ht  sources
and reaction conditions used,  it  is  difficult  to  reliably assess  endosulf*n
photolysis.  The information  does  suggest,  however,  that endosuifan  sulfate
Is taore stable  to photolysis  than  endosulfan  itself.

      27.4.2  Oxj.datJ.on

          Although  oxidation  of  endosulfan  to  endosu If in suit are  has been
frequently reported as occuring  in photochemical,' biological,  and chemical
systeas, the source or the nature  of the oxidants are unclear (see also
Sections 27.4.1 and 27.4.6).   The  only  quantitative data on oxidation of
endosulfan isoiaers  indicate that  in  air-saturated water the oxidacio.i of
endosulfan (presumably by  molecular  oxygen) may have a  half-life  of  about
70 days at 20*C, and in the absence  of  product information this half-life
must  be used ifich caution.

,          Gr*ve and Wit (1971) reported the oxidation r=jte constants shown
below for both  endosulfan  isomers  i'i water  at  two pH values jnd 2!'JC:

                              2 Endosulfan           3-Endosulfan
           pH  7.0           10.4  x  1G~3  days"1     9, 7  x 10" * days"'

           pH  5.5            8.3  x  10"3  davs"1     9.9  x l^"3
 The  r^te  constants  were  determined  by  the  difference between first-order
 rate constants  measured  in  aerobic  and anaerobic  buffered reaction
 solutions.   The authors  concluded that th«  Txida: ion rate was independent
 of pH,  but  they did not  discuss  what species  were  responsible for the §up-
 posed  oxidation reaction:   no  product* were reported,  and it is  therefore
 itapoastble  to  verify that oxidation was actually  beir.^ m
           In  striking  contrast  to the  apparent  reactivity of .endosu If an to
 ward  oxidation  reported  by Creve  and 'Jit,  Hoffman ind Hichslsdoerf er ;19'1
 have  reported that  endosulfan  is  only  slowly' oxidized by nzone .as follows:

       Oj  Concentration        Concentration       *  Endosu If Jin Reacted
           in  Solvent            of Pesticide          _:            f

      T* ag/ 1  in hexane           5 mg/ 1               2             52

      17 ing/ 1  03 in  9:1   .
       water/acetone              2 m^/1               •"'             11

The reaction eorul{cions were effected  by  bubbling  Che .ozone  solution
through  che endosMifan  in solution  for 45 minutes;  heptachior  and  aldrin
were completely reacted under these conditions.
          Thus, although endosu'fati may be  expected to  be  oxidized  to  its
aulfate,  It is difficult to explain why molecular  oxygen should effect  oxi-
dation while the powerful oxidant ozone does  not give significant oxida-
tion;  no conclusion regarding endosulfan oxidation in  the environment  can
therefore be substantiated,  Endosulfan sulfate should  be  mo-re  stable sthan
endosulfan  cowards oxidation.

     27,4.3  Hydrolysis

          Both isoraers  of endosulfan will hydrolyze slowly in aquatic  en-
vironments of pH < 7^and < 209C with half-lives of greather  than I  month.
Although  the hydrolysis of endosulfan will  be faster above pH 7,  literature
information provides no direct quantitative data for a  half-life  estimate
at pH >  7.  However, analysis of the data of  Martens (1976)  and Greve  and
Wit (1971)  indicates that aC'pR S and  20'C  the half-life will be  about  3.5

          Greve and-Vic (1971) measured'the following hydrolysis  rate  con-
stants for  both isomers of endosulfan at  20*C and  two pH values:

                          3-EndosuIfan             3-Endosulfan

  pH 7.0                2,0 x 10~2 days"1        1.9 x 10*2  days"1

  pH 5.3               4,6 x 10" 3 days'1       3.7 x 10" 3  days"1'

Subsequently, Martens  (1976) reported hydrolysis data that were obtained as
controls  in biotransfortnation studies at  27 °C.  Based on one data point
taken after ,10 days forseach pH, the' following losses of endosulfan were

                       pH     4.3      5.5      6.3     7     >S

   .  Sndosulfan lose          <12I      n      8?    2S?     >9'X?  .

The author also reported that endosulfan  diol  was  the only product  formed,
and that  in some biodegradacton studies above  pH § endosulfan was not
hydroiyzed as rapidly as in the control because of sorption  by  the cell
'mass and, therefore, was not available for  alkaline hydrolysis.

          All (1978) reported that 3- and 3-endo$ul.f*n  were 82", and 87^.
recovered after 33 days in water used in a microcosm experiment, with  no pH
or temperature information given; these losses correspond to half-lives of •
115 and  164 days for tfr  respective isomers.   Eichelherger and  Lichtenber.*?
(1971) studied the persistence of endosulfan  at 10 yg,'1 concener it ion  In ,1

sample of raw river water froc the'Little Miami Rtver, Ohio.  They reported
that the peaks of both isomers, as determined by glc, were reduced 70%
within one week;  at the end of 2 weeks, only 52 endosulfan remained,
though identification by glpc was extrenely difficult.  The hydrolysis
product contained no sulfyr;   Infrared spectral dati Indicated that It was
probably endosulfan dial.  The pH of the river water was reported to vary
from 7.3 to 8.0 during the experiment, so 'rhac the rapid hydrolysis of en-
dosulfan found is reasonable.  Since in the same river water samples, dlel-
drin was formed from aldrin, the sterility of the water must be questioned;
biotransformation stay have therefore contributed to the endosulfan trans-
formation race found.

          From the above information, the data at 20°C of Greve and Wit
(1971) are in'raoderate agreement (i.e., within a factor of two) with the
less precise, one-datum point experiments of Martens (1976) at 27°C; the
half-lives (days) calculated from their d*ta are compared below:

                              Greve and Wit               Martens
                                at 20°C                   at 27°C

          '  pH 5,5            150 a, 187 3                  343

            pH 7.0             34 a, ,37 3                    21

The aore rapid rate of hydrolysis of endosulfan reported by*Eichelberger
(tj/2 : •* days) saay be partially explained by the higher reaction pH, but
biotransformatIon Is also a possible explanation.

          The only information on the hydrolytic stability of endosulfan
sulfate is that of Ali (1978),'who reported that endosulfan sulfate was 882
recovered after 33 days in water used in a microcosm experiaent, with no pH
or temperature daca given (Section 27,4,9),  This loss corresponds to a
half-life of 178 days. • It Is interesting to note that two other cyclic
                    L                   '             ir

sulfates, triai-?thylene and ethylene aulfate (I and II, respectively) have
hydrolysis half-lives of 3.1 and 0.31 hours, respectively, at pH 7 and 20"C
(Radding e£ ajL 1977),  The slower rate of hydrolv«ls!of endosulfan sulfate
may reflect the Influence of the «*ven member ring.


     27.4.4  Volatilization                           ,  '

          Limited and somewhat contradictory data.are available on the
volatilization of erdoaulfan from water.  This information suggests that
endosulfan will have a volatilization half-life of more than 11 days, and
possibly more than a /ear.

        •  A theoretical volatilization half-life of 11 days is calculated
for a quies.ent water body u^ing the equations and assumptions of Mackay
and Leinoner (1975); the calculated half-life would be less in more
turbulent water bodies.   •                     (

          Volatilization of endosulfan was also measured as a control in a
bioaccumulation study by Ernst (1977),  Aquaria aerated at 2.5 iiter-hr~^
for 67 hours showed only an 11? loss of eridosulfan;  in this system the
same losses (US) of dieldrtn and heptachlor epoxide were measured.  Since
Mackay and Leinonen (1975) calculated a half-life of 1.5 years for volatil-
ization of dieldrin using the same equations and Assumptions that we used
to calculate an endosulfan volatilization half-life of 11 days, there is
clearly son>e discrepancy in ^hese data.

          Martens (1976) als'o studied volatilizacion as a control in bio-
transformation studies in which the biode'gradation flasks were aerated for
one hour per day for the 10 days allowed for incubation.  Sorption on biota
wist found to reduce the amount of endosulfan volatilized, with less than 12
volatilization loss of endosulfan observed in the presence of fungi; .
volatilization losses were 2 and 20% in the presence of bacteria and
actinomycetes, which showed le*s. sorption of endosulfan than the fungi.

          Although not directly relevant to aquatic systems,, other studies
have provided Information on volatilization from the solid state (glass
plates).  Although not discussed in the text of the paper, as auctt as one
half of the endosulfan isomers and endosulfan suifate applied to the plates
was lost by volatilization during a 7-day uv-lanp irradation (Archer et al..

          Mo information was found regarding the volatilization of endosul-
fan suifate from aqueous systems.

     27.4.5  Sor pji_on '

          Sorption is an important fate for endosulfar in aquatic systems
and sediments «ay be a sink for endosulfan.

          Greve and Wit (1971) analyzed water from the Rhine River and de-
termined the concentration of endosulfan in the supernatant water, in the

water containing slit chat settled on standing, and in the bottom mud.
Although the data presented are crude, andosulfan was shown to be > 75%
associated with-the particulata material (silt or aud),1 Indicating
significant sorption of both it>oiners of endosulfan.

          Richardson and Epstein (1971) studied the adsorption of endosul-
fan on silt loam and found that the greatest retention of endosulfan oc-
curred on the colloidal 3nd,0.08 to 0.5 ure fraction of silt and clay.  The
authors also stated that treatment of the soil with hydrogen peroxide (to
famove oxidizabie organic laterial) reduced the amount of andosulfan re-
tained on the-soil; no data were given,

          No data were found on the sorption of endosulfan sulfate.

     27.4,6  Bjoaccumulation

          Ernst (1977) 'reported ... ptak* and elimination of several pesti-
cides, including'^-endosulfan, in mussels in seawater.  Mo bloaccumula-
tion studies were found in freshwater aquatic systems.  3-Endosulfan was
found to have a concentration factor of 600 (ppb wet weight in mussel vs»
water at steady state with a 2.05-ppb initial concentration in water).
Ernst also gives first-order rate constant data for uptake and elimination
of "S-endosuifan in the mussels, but the half-lives of 3' minutes and 34
hours for uptake and elimination, respectively, are questionable because
the data reported clearly show that more than half of the 84-pp-.b (wet
weight) concentration in the tmisse.l is eliminated after 9 hours.  Based on
Che latter and the concentration factor of 600, it is not likely that
bioaccuraulattor1 of 3-endosulfan (and presumably also 3-endosulfan) is a
significant process.

        ,  Mo .information was obtained on the bioaccumulation of endosulfan
sulfate.                               '               .         .

     27-4,7  Biotransfonuation and Biodegradation

          Creve and Wit (1971) reported that endosulfan can be degraded by
microorganisms in waters of at least pH 7 and wita high dissolved oxygen.
Under these conditions (pH 7, air-saturated, water) and at 20"C and in the •
presence of Pseudogionaa at 10^' cells/liter, the half-life of endosulfa.i
was one week (no products reported).

          Martens (1976) 'studied the transformation of endosulfan by soil
microorganisms in nutrient solutions.  After 6 weeks of incubation, 16 of
28 soil fungi were found Co have degraded more than 30?, of the applied en-

dosulfan;  13 of the 16 fungi gave end03ulfan sulfate as the major produce
(i.e., an amount at least 10 times greater than that for endosulfan diol}.
Martens also found that after 10 days of incubation, 15 of 49 soil Bacteria
harf degraded more than 3QZ of the applied endosulfan with 10 bacteria
giving endosulfan dlol as the major identified product;  endosulfan sulfate
was the major product of the other 5 bacteria.  In similar experiments, 3
of 10 actinomycetes aetabolized too re than 30X of the applied endosulfan
after incubation for 10 days.

          No information ma found regarding biotram,formation of endosul-
fan sulfate in aquatic systems, although the formation and presence of the
sulfate in the- fungi experiments suggest some stability toward biological
hydrolysis.  In cases where the diol was the major product in biotransfor-
aiation there is no information to indicate whether the diol arose directly
froa endosulfan or did indeed proceed via hydrolysis of Che endosulfate
     27***®  Other Reactions

          No reaction* other than those cited above have been reported to
occur with endosulfan or endosulfan sulfate.

     27.4.9  :|ierocosB Studies.  Field Studies,  and Modelling

          Alt (1978) reported results of terrestrial-aquatic microcosm
experiments in which the d and 0-endosulfan isomers and endosulfan sulfate
were each studied in separate experiments;  a mixture of endosulfan isomers
wa» studied in another microcosm experiment.  In the latter study, the
S-lsoBer was reported to be rapidly lost, wich none detected In the aqueous
phase 26 day* after Introduction of the mixture into the nicrosusa.  At the
end of the 33 days, the a-iaomer comprised 16% of the total endosulfaa
material recovered Iron water, with the remainder beng endosulfan sulfate.
In all experiments, *ndosulf«n sulfate was the only metabolite and also the
major coo pound found in water or in the organisms (algae, snail, mosquito,
and fish).  Baaed on the microcosm experiments, All concluded that the
3-isoraer was not oxidized metabolieally to endosulfan sulfate, but was
first isomerized to the »-isom«r, which wa* then oxidized;  in studies on
each isoaer, some small conversion to the other isomer was also found.
Data for bioconcentration factors were also found to depend on the starting

material and on the organism used.  The upper and lower concentration
factor reported for each orggrism and compound are given below:
                     Bloconcentrat ion factor for:
                                                      endosulf an
               o-endosulfan       8 -endosulfan          sulface

Algae'           17-999            44-3863             223-1654
Snail            1336-5763    '     8174-39457          5457-29'* 30
Mosquito •  '      218-831           1245, 1508*         210-763
Pish             30-304            50,388*             935-1741
  Only two values reported; three factors were reported for other
  endosulfan measurements, with four factors for endosulfan sulfate.
All described a comb!nation of processes Involving tsomerization,
oxidation, uptake by organisms, and serption to sediment to explain some of
che results obtained In his experiments; because'of the terrestrial
processes, which in part determined the fate of endosulfan and the sulfate
in the experiment, no detailed conclusions on the fate of these compounds
In aquatic systems can be made.  From the Information presented it appears
however that endosulfan sulfate is generally more persistent" and
bloaccumuiates more than the endosulfan isomers.  It Is of interest to note
that when the three compounds were dissolved In water from the microcosm
and exposed to th* artificial  light used In the experiments, the a-isomer,
S-lso»»r, and endosulfan sulfate were 82%, 87% and 88% recovered, re-
spectively after 33 days, Indicating that, at least for endosulfan,  hy-
drolysis and photolysis are not important relative to biological processes.

        Gr«ve and Wit (1971) calculated an endosulfan half-life of about
two days in a pond, based on data supposedly obtained from a paper by
"Bears and Mare" (actually Beard and Ware 1969), but the reference cited
contains no data from such an experiment.

27-5  Reac 11onProducts

     The following products have b«en  reported  for  endosulfan in various
studies, with the associated processes as  listed:
    Endosulfan sulfate

       Photolysis of solid of gas
    »  Biotransformation
    .  Oxidation (?)
_Endo£uIf art.. J lg 1

   Photolysis  in  solid,
      liquid, and  gas  phase
.  Biotransforraation
     Etulgsulf an gther

     .  Photolysis

27,6  Sat a Suargary
Endosulfan  iactone
     Tables1 27-"! and 27-2 summarize, ch€ data  on  th*  aquatic  fate  of
dosulfan isomers and «'ndosulfan sulf ate, respectively.

I                VuUUlU
                 Surpt luii

                 I lodi 1 '.«u
                                                                         T.blu 27-3
                                                               of  AquAdc F«t« of Eaduaull«a Surface

              btvirumLxitdl                            SLmku~y                                          Half-                   Ctxif iduncu
                Piux's^-1                             Statemant                     tote               Life tH                   ul Uul 
 2 7 , 7 , Li carat'ire  Cited

 All, S.   1978.   Degradation and environmental fate of endosuifan  isomers
   and ^endos'j Lf 30  sulfata  In ^ouse, insect and laboratory "nodel  ecosystem.
   ?h. D.  Thesis.  Univ. Illinois.   ,ln I pt>. '."niv.  Microfilms, Ann Arbor,
   Mich.,  Order  \'o.  7820391.  (Oiss.  «bstr. Int.  3  1973,  39(5):2117).

 Archer,  T.E.  1973.   Endoaulfan residue*; on alfalfa hay exposed to  drying
   by sunlight,  ultraviolet  light  ind air.  Pestic. Sci,  i:59-?8.

 Archer,  T.S., I.K.  Mazer,  and D.G. Crosby.  1972.  Photod-scorapos it ion  of
   endosulfan and  related  products in thin films by ultraviolet  iisjht
   irradiation.   J.  Agric'.  Pood Cham.  20(5): 954-956.

 Beard, J. E., and G.w. Hare.   1969.   Face of  endosulfan on plants and
   glass.  J, Agric.  Food  Chem.   17(2 ):2 16-220.

 Eicheiberger, J.W.,  and J.J.  Lichtenberg.  1971,  Persistence of  pesticides
   in river «ater.   Environ. Sci.  Techn'ol,  5(6): 5il-5i'4,

 Ernst, W.   1977,  Determination of the  biocancentrat ion potential of "narine
   organisras. -  a  steady state approach.'  Chamosphere *5 (11): 731- 740,

 Gore, C. , S.W.  Haanan, 3.C. Patta^ini,  and T.J.  Porro.  1971.   Infrared and
   ultraviolet 3pe\Ltra of  seventy-six pesticides.  J. Assoc. Off.  'Anal.
   Chem,   5->(3): 10*0-1*''82.

 Gr»v», P.A., and  S.'_.  Wit.   i97t.  fndos"lfan in'the Rhine River.   J. Water
   Pollut. Control F»d. -i 3 ( 12 ): 2 3 3*-23^S .,

 Harrison, R.3. , D.C.  Holmes,  1.  Hoburn  and J.  O'G. Tactan.   19^7.   The fate
   or some organochlorine  pestlv-des on leave1*!.   J. Sci. Fcod Agric,

 Hoffmann, J., and'D.  Eichelsdoerfer.   1971.   Effect of ozone in
   chlorir.ated-hvdrocarbon-group pesticides in water.  '. J.TI Vaster

 Koshy, "., V.M. Das,  and M.S.  G.K.  Vair.   1972.   Deter 1-jr.it i--n  of
   insecticides  on glass and on leaf  surface.   Agric. Res. J,  Kerala.
   in(2):I28-132.                                '   •

lackay, D., and ?.J,  Leinonen,   1975.   Rate of evaporation of
   low-solubility  contaminants from water bodies  to atmosphere,.  Srwiron.
   Sci. Tech.  9C'3):1173-U80.

Harteas, 9.  19"".   lecomp-is It ion of  endon>,
   Schrlfter.  Ver. W.isser-,  Boden- , Lu? t he, ,  Ber If n-Dahlem,  (')7 ;  : IT"- IT },

Martens, R.  1976.   Degradation of  (8,9~^CJ  endosalfan by soil
  raicroorganisms.   Appl, Enviroa. Micfobiol.  31('j,); 153-858.


Midwest Kesearch Institute.  1977.  Initial scientific 'review of
  endosulfan,  MKI Report Nov. 1977.

Phillips, W.iL.J.  1975.  [Publ.j   Endosulfan.  Its effects on environmental
  quality.  Nat, Res. Counc. Can., Environ. Seer., 14u98, l(;0pp.

Putnam, T.3., D.D. Bills, and L.M. libbey,  1975.  Identification of en-
  dosulfan based on the products of laboratory photolysis.  Bull. Environ.
  Contaro. Toxicol.  13(6);662-665.

cladding,' 3.8., O.H. Liu, H.L. Johnson, and T. Mill.  1977.  Review of the
  environmental fate of selected chemcials.  U.S. Environmental Protection
  Agency Report *lo'. EPA-5bO/5-77HjJ3.

Richardson, E.M. and E. Epstein.  1971.  Retention of three insecticides on
  different size soil particles suspended in water.  Soil Sci. 'Soc. Am.
Schumacher, G., W. Klein, and F. Korte.  1971.  Photochemical breakdown of
  endosulfans in solution.  Tetrahedron Lett. N'o. 24:222^-2232.

Schumacher, H. G,, H. Parlar, W, Klein and F. Korte.  1974. , Beitrage zur
  okologischen cheatie.  LV*.  Photoch,eralsche reaktion von endosulfan.
  Cheaosphere 3(2}:b5-7U.

Schuphan, I., and K. Sallschmiter ,  1972.  Persistence of hexachlorobicyclo
  [2.2.1 j heptene derivatives.  Fresenius' Z. Anal. JChena.  239( 1 ): 25-28.

Weil, L. , G. Dureti and K.E. Quentin.  1974.  Solubility in water of
  inst-ctiide chlorinated hydrocarbons and polychlorinated biphenyls in
  view of water pgliution.  2. Uasser Abwa^ser Forsrh.  7 { o): 169-175.

                      28.  ES0RTN AND ENDRIN ALDEHYDE
28.1  Statenent of Probable Fate

    Little information Is available for evaluating the fate of endrin in
aquatic systetas.  Photolysis and bloc rans format ion of endrin occur-under
environmental conditions, but no data are available to assess the rates, of
these processes In aquatic environments; blotransforaation will also be af-
fected by the microblal types and popularlona available to utilize endrin.
Mo information OR the sorpcion or volatilization of endrin fron aquatic sys-
tems is available, although bioaccunmlat ion does appear t<, be significant
with concentration factors on the order of 1CH - 10*.

    No Information whatsoever has been found to evaluate the fate of endrin
aldehyde in aquatic systems.  Infannation available on the photolysis, bio-
transforaation, and thermolysis of endrin suggests that endrin'aldehyde
will be only a minor product of these processes.

28.2  .Identification '

    This section discusses only the pure chemicals endrin and endrin alde-
hyde.  A typical technical endrin sample contains 96.6? pure eadrin and at
lease eight impurities, with each usually at less than l"K cancentration;
the impurities include dieldrin, isodrtn, aldrln, heptachloronorbornadiene,
heptachloroitorbornene, endrin aldehyde ard'i-keto endrin (Brooks 1974).

    The structure, alternate names > and CAS and TSL nuosbers for endrin ire
given below.
          Ci  Cl
Alternate Naae_3

Isodrln epoxide

    CAS Mo. "2-20-8
    TSL So. 10 15750

    Sndrtn aldehyde Is a transformation  product  of  sndrin.   The structure
at endrin aldehyde cited by Burton and Pollard (19/4) is given below.
            0 .0
                                          A It ernat e  Xaaes

                                          1,2 ,4-Methenocyc lon«?nt -J ( c ,d )
                                            2 , 2a, 3,3,4,7- hexach iorodecahydro
    Sndrin aldehyde

    CAS N*o. 7421-93-4
    TSL So, ?»one assigned

   3  Fh y s i c a 1 P r Q rg r t i g s
          ner-il physical properties  of  e^drii are as follows.

    Molecular weight                            381.1"1

    belting point                               235 'C,  with decomposition

    Soiling point                               Vo data found          ' ,

    Vat or pressur** at 25 JC                      2 x l->~ '  corr
        t in 1972)
0.25 ppm
    Solubiiitv in Water at  25'C
    (Veil e£ aj..  19 '-» >
    (Bigger and Riggs 1974)*
    *.og" oetano'l/water partition coefficient     5.*> calc.
    (N'eely et al,  1974)
* Particle size <5.0-!n
    ^o physical property data were  found  for endrin aldehyde except for a
melting point of 145'C - 149'3C  (Phillips  et_ a_l.  1962);  the noletular
formula and weight of endrlri aldehyde  are the sane as those of endriru

28.4  S'.imiaary of Fate D-ta,

    28.4,1  Phot3l,vsji3

         Although the photolysis of endrin in the solid state and in hexane
solution has been shown to occur tn sunlight, there are no data to evaluate
the photolysis races of endrin In aquatic systems.  The studies reported do
suggest that endrin aldehyde Is not a aajor photolysis product in sunlight.

         No information was found on the photolysis'of endrin aldehyde,

         Fujita et_ aj_. (1969)- found that II (below) was a product of the
sunlight photolysis of endrin in hexane solution; they presumed II was a
photolysis product of I.  No information regarding the product yield 'Or
photolysis rate of endrin was given.  Zabik £t_ a_l, (197l> also found II to
b« the major photoproduct when endrin was photolyzed in hexane or cyclo-
hexane solvents at wavelengths above and below 290 run and in sunlight, with
II obtained 'tn yields as  large as 80% in some experiments.  The authors
also reported that irradiation of I at 254 not for 26 hours did not give II;
the photoprodycts of I were not determined, however.
          I. C:"-Keto' endrin}
         Burton and Pollard (1974) studied the sunlight photolysis of en-
drin as thin solid layers of glass planchets.  Th« major product was the
endrin photolsoaer I, which is also called' i-keco endrin.  The authors
state that only ainor amounts of other compounds, §uch as endrin aldehyde
III, were fomed.  After 5 and 12 days of exposure to sunlight in June,
                                                     III. Endrin aldehvde

5-keto endrin was found In yields of io and 63%, respecti"eiy,  it is not
clear whether the percentages were based on the initial amount  of endrin or
che amount: reacted.   Burton and Pollard also reported that ''endrin was 50*
isonerized to I in 7 + 2 days in intense summer sun with complete conver-
sion of endrin in' 17 + 2 days," and noted that the thicker films were pho-
colyzed aiore rapidly than thinner films.  Applications of their data Co
reactions' in aqueous solution must therefore be made with caution.  They
also reported that I was 'stable to* sunlight; thus endrin aldehyde is not a '
product of further photolysis of I, and nay be a direct photoproduct of

         Roburn (1963) and Rosen et_ al_. (1966) reported other studies of
tht photolysis of endrin under various conditions.  Although their studies
do not provide any Jaca useful In evaluating the photolysis rates of endrin
in aquatic systems,  they do provide more information on the photolysis of
endrin and th* products formed.  Roburn (1963) found that solid films'of
endrin irradiated at 254 n» for several hours gave one main product and
several Minor products.  Rosen et_ a^L. subsequently identified two of these
photolysis products  as "5-k«to endrtn -I and endrin aldehyde III  in yields of
37* and 9t, respectively.

         Bulla and Edgerly (1968) reported that the photolysis  rate of en-
drin in dilute aqueous solution at 254 run was independent of reaction tem-
perature (20* to 408C),  Baker and Applegate (1974) studied the photolysis
of solid endrin filns using a lamp with a maximum output at 350 nm; the
loss of the endrtn was found to be due to both photolysis and volatiliza-
tion from glass surfaces.

         Ivie and Casida (1971) reported studies of the sunlight photoly-
sis of endrin on bean leaves in the presence of rotenone as a photosensi-
tizer.  After one hour exposure to sunlight, the endrin photoproducts (as
determined by 14€ label) were found in yields of 22, 15%,, and 42%.at
rotenone levels of 1, 10 and 100 ppm, respectively.  Although these data
indicate that endrin photolysis can be sensitized by rotenot.e,  the rele- ,
vance of these solid state, sensitized processes to aquatic systems is

         From the above information, it is concluded that although photoly-
sis of endrin in sunlight does occur, the nature if the products' is uncer-
tain.  Endrin aldehyde does appear to be a minor photolysis product of
endrin.  It is unlikely that the dechlorinaced f-keto endrin iproduct IT
would be formed in aquatic svstems unless suitable hydrogen-atom donating
substrates are'available in aquatic systems.  If photolysis of  endrin is
determined to be an  Important fate of endrin, more relevant product studies
are required.

    28.4.2  Oxidation

         No information on the oxidation of endrin or endttin aldehyde that
is relevant to aquatic systems was obtained.   Leigh (1969) studied the oxi-
dation if endrin with chlorine, permanganate, and persultate and conclude*
that, endrin was unaffected by these oxi^ants  at 50 '0.

    28.4,3  Hydrolysis              <

         Eichelberger and Lichtenberg (lO7!)  examined the persistence of
endtin in j s.imple of raw water from the Little Miami  River in Ohio,   After
8 weeks at room temperature, all of the endrir; was recovered,  indicating
that endrin nas a half-life of at least i years; this  calculation was made
assuming that a 2,5% conversion in the 9 week period could have been  de-

         There1 is no evidence to suggest' that hydrolysis of endrin aldehyde
vtll occur in aquatic systems; by analogy co  endrin, a hydrolysis half-life
of at lesst i years is- probable for erdrin aldehyde.

    28 . •* . 4  Volat ilizatiqn

         Xo data were obtained on the volatilization of endrin or endrin
aldehyde from, aquatic systems.
         So information was found on sorption of endrin or er.drin aldehyde
to sediments or biota.

    28 . 4 . 6  Si^oaccMulat i_on

         Bioaccumulat ion appears to be an important -process for endrin in
aquatic systems, with concentration factors ranging from 10- to 10"* for
species in microcosm experiments.

         Ernst (1977) has measured a concentation factor of 1.9 x 10^ for
uptake of  endrin by mussels.   When the mussel was placed in clean water the
tine required for elimination of about 50* of the endrin was about 24

        Metcalf et_ ajL, (1973) measured the concentrations of endrin in
several elements of a microcosm; 11.56 ppra, 125 ppm, and 3.40 ppa concen-
trations of endrin were found in an algae, snail, and fish, respectively,

compared to 0.0(1)254 ppm In the water.  The  ratios  of  2ndrtn concentrations
in the organisms to the concentration of endrln  in solution were  1300  for
fish, i9QOO for snails, and 4603 for the alga.   When  exposed  to endrin only
Is the aquatic medium  (i.e., no dietary routes), the  ratios were  630 for
Che fish, 310 for mosquitoes, and 330 for Daphnia  after three  days.  The
difference in the ratios for the'fish may result.from several  factors,  in-
cluding insufficient time for fequilibration  In conjunction with the differ-
ent routes for uptake  of the pesticide by the fish.

    28.4.7  Blotransjf_ormation airtd__3Lodegradation
         Although endrln has been found to undergo biotransformat ion in a
microcosm and in studies with microbial isolates,  there are no data useful
for predicting blotransformation rates in aquatic  systems.  It is interest-
ing that the metabolites reported in the two  types of experiments appear to
be different, with isooerization/dechlorination  occurring in the  'jicrobial
isolates and (possibly) hydroxylation occurring  in the microcosm  experi-
ment.  Mo Information  has been found on'the  biotransformat ion  r •.   endrin

         Several-studies on the biotrans format ion  of endrin 'nave  been  re-
ported by Matsumura's  group (Patil p_t_ al. 1970; Matsunmra e_t_ al.   1971;
Patil e_t_ al. 1972).  In the 19!70 paper,  it was  reported that  20  microbla
isolates from soil that were previously found capable of degrading dieldri
in solution were also  found to degrade endrin.  Of  the several metabolites
detected by tic, only  S-keta endrin was identified.  Matsumura et al.
(1971) subsequently reported that 25 of 150 uiicrobial isolates trom soil
.showed activity in transforming endrin in solution in 30 day experiments.
Three major and four minor metabolites were generally found in these stu*-
dies, with 5-keto endrin being obtained in yields  ranging from 5% to 95% in
the various cultures.  Other metabolites were not  specifically identified
although the authors did conclude on the basis of  mass spectral and infra-
red spectral evidence  that the other metabolites were ketones and aldehydes1
with-five or six chlorine atoms per molecule; two  metabolites had infrared
spectra that were identical to that of endrin aldehyde (III) reported  by '
Phillips e_t_ aJL, (1962).  Patil et_ a_l» (1972) also  studied the transforma-
tion of several chlorinated Insecticides, including endri'n, in water and
algae from a marine fish pond. An unknown metabolite of endrln was obtained
in a 36% yield in the water, whereas a 24t yield of 5-keto endrin was  ob-
tained from the algae  collected from a stagnant fish pond.

         Microcosm experiments by Metcalf et_ aJL.  (1973) also revealed  blo-
transforaation of endrin;  23% of the endrin ingested by a caterpillar in

the terrestrial phase of the microcosm was metabolized to unknown products,
and 1C 'was not determined whether any transformation also occurred in the
aquatic phase of the microcosm (see Section 28.4.9),  Of the products de-
tected by tic, 6-keto endrin was determined not to be a product.  The
authors suggested one metabolite was 9-hydroxyendriti IV, which was a neta-
bolita found in rat faces (Baldwin e£ _al. 1970)..  Baldwin found another
hydroxylated endrin in the rat feces as well as 9-k.eto endrin in the fat of
the rat; the 9-keto endria is the oxidation product of 9-hydrogen endrin.
There is no Information other than Meccalt's study to indicate whether hy-
droxyiated endrin metabolites will^be formed in aquatic systems.  It is
also found that about 80% of the '•"'C labelled material recovered -from .the
alga, snail, and fish was In the form of the original endrin,applied in the
terrestrial phase of the microcosm, this suggests that endrin is moderately
stable toward transformation'in vivo (see Section 28,4.9),
         Hill and McCarty (1967) reported that endrin was about 50% reacted
in thick anaerobic sewage sludge in 5-14 days, and they observed four
unidentified products; no other information was given.

    24.4.8  OtherReactions

         Phillips et al, (1962) reported that endrin is theraally isome-
rized at 230aC during gas chromatographic analysis.' The two products
Identified - - the aldehyde III and ketone I — gave peaks of approximately
«qual size in the gipc trace.  A thermal decomposition of endrin was found
to be an exothermic reaction; the total residue contained approximately 15
to 20% aldehyde III, 55 to 60? ketone I, 51 of a "bird-cage" al >hol V, and
15 to 20% volatile, carbonaceous and unidentified decomposition products,
Sarlow (1966) also reported that a sample of pure endrin stored in the dark

for four years underwent spontaneous  isomerization  to  give  a  70S  yield  of
I.  Mo Information is available  to determine whether such  reactions  may be
catalyzed and occur  in aquatic systems.
    28.4.9  Microcosm Studies, Field Studies, and Modelling

       .  Metcalf et_ al.  (1973) investigated  the  fate  of  14C-iabeled
eudrln In a terrestrial-aquatic microcosm.   Of the  total  ^C-labeled
material recovered from  the alga, snail,  and fish,  the endrin  component
eonsrlfuttsd 85?, C3%, and 7f>% of the label,  respectively.  The  biocon-
centration factors for the above organisms were  4.5 x riCH, 4.9  x  10  ,
and 1.3 x lG-%  respectively.  None  of,the major  metabolites  co-chromato-
graphed (tic) with 5-keto endrin.   The authors suggested  one unknown  could
be 9-hydroxyendrin IV, which was a  metabolite found in rat feues  (Baldwin
et_ al. 1970).

'28. S  O.f j-ja Summary

    Table 28-1  summarizes the data  on  the aquatic fate of endrin  and  endrin

                                                                      Table  M-\

                                                      of Aipiatu  Fate ut Endrln ,jj
                       i. t*^,yg»at


                    Onldai ton
Volat llicat ion

Sorpt Ion

May be *

No information  available-.

Not an Important  process.

No information  available.

No information  available.

la an Important  procebs.
                    Bluttauafuiwation/  Hay b« au important  piocesa.

                                                                                                            of II

               a.  No  tutormut ioti  was available tor niiy proce^^ tor eutirin aldehytifc1.

               b.  Tlierc LS  liiHut He lent information  4n the  levlcyed literature to permit  as^e^sraeat of a mo^t piot>abK

28,6  Literature.. Cited

Baker, R. D. , and H. G, Applegate,   1974,  Effect  of ultraviolet  radiation
  on the persistence of pesticides.  Tex, J.  Set.  23(1-4 ) :53-59»

Baldwin, M. K. , J. Robinson, and D.  V. Parke,   1970,  Metabolisa  of  endrin
  in the rat,  J. Agric. Food Chem.  18(6 ): 1117-1123.

Barlow, P.  1966.  Spontaneous decomposition  of a  sample  of  pure  endrin.
  Nature 212:505.

Slggar, J. W. , and I. R. iiggs.  1974.  Apparent solubility  of1 organ o-
  chlorine insecticides, In water at  various teraperatures.  Hilgardia

Brooks, G» I.  1974.  Chlorinated Insecticides: Volume  I:Technology  and
  applications.  CSC Press, Cleveland, OH., 249 pages.  '

Sulla, C« D. , III, and  E. Edg2rley.  1968.  Photochemical  degradation  of
  refractory  organic compounds.  J.  Hater Pollut.  Control. Fed.

'Burton, W. B. , and G. E. "ollard.   1974.  Rate  of  photochemical isomeri-
  zacion of endrin in Sunlight.  8ull, Environ. Contain. Toxicol.
Eichelberger, J. W. , and J. J. Lichtenberg.   1971.   Persistence  of  pesti-
  cides in river water.  Environ. Sci. Technol.   3(6 ):54i-544.

Ernst, W.  1977.  Determination of the bioconcentration  potential of  marine
  organisms. - a steady state approach.   Chemosphere 6(11 ):731-740.

Fujita, M. , A. l^hli, and Y. Sakagaoii.   1969.  Photodecoroposition of
  endrin,  J. Hyg. Cheo. 15:9-12.

Hill, D. W. , and P.'L. McCarty.  1967,   Anaerobic  degradation  of selected
  chlorinated hydrocarbon pesticides.  J. Water Pollut.-  Control  Fed,

I'vie, G. W. , and J, E. Casida.  1971.  Sensitized  phot'odecomposition  and
  photosensitizer activity of pesticide  chemicals  exposed to sunlight  of
  silica gel chromat opiates.  J. Agric.  Food  Chem.   19:405-409.

Leigh, G. M. 1969.  Degradation of selected, chlorinated  hydrocarbon
  insecticides.  J. water Follut,, Control Fed. 41(ll)(Pt. 2):R450-R460.

Martin, H. , ed. 1972.  Pesticide Manual,  3rd  Ed.   British Crop1 Protect iq>n
1  Council, Worcester, Engi.

Maesumura, F., V, C, Khanvilkar, K, C. Patil, and G. M, Boush.   1971.
  Metabolism of endrin by certain soil microorganisms.  J. Agric.  rood
  Chem. 19(1}:27-31.  ,                                           !

Metcait, R. L., I. P. Kapoor, P.-Y. Lu, C, K. Schuth, and P. Sherman.
  1973.  .Model ecosystem studies of the environmental fate, of six
  organoehlorine pesticides.  Enviroii. Health Perspect. 4:35-44.

Neely, W. 8., D. R, Branson, and G. E. 31au. ' 1974. , Partition coefficient
  to measure bioconcentratlon potential of organic chemicals in  fish.
  Environ. Sci. Technol. 5(13>:1113-1115.

Patil, K. C., F. Matsuaura, and G. M. Boush.  1970.  Degradation of  endrin,
  aldrit!, and DDT by'soil microorganisms,  'Appl. Micrbblol. 19(5 ):8?9-831,

Patil, K, C., P. Matsumura, and G. M. B~-ush.  1972.  Metabolic transfor-
  oaeion of DDT, dieldrin, aldrin, and endrin by marine microorganisms.
  Environ. Sci. Tech. 6(7):629-632.

Phillips, D. D., G, E. Pollard, and S. B. Soloway.  1962.  Thermal  isoweri-
  tation of endrin and its behavior in gas chromatography.  J. Agric. Food
  Chein. 10(3): 217-221.

Roburn, J.  1963.  Effect of sunlight and ultraviolet radiation  on  chlori-
 , nated pesticide residues.  Chen. Ind.  (1963):1555-1556,

Rosen,  J. D,, D. J. Sutherland, and G. t. Lipton.  1966.  The photochemical
  isomerization of dieldrin and endrin and effects on toxlcity.  Suli,
  Environ. Contam. Toxicol. 1(4 ):133-140.                           ' •

Rosenblatt, 0. H., I. A. Miller, J. C, Dacre, I. Muui, and D. R. Cogley.
  1975.  Appendix M- Endrin.  Preliminary Assessment of Ecological  Hazards
  and Toxicology of Environmental Pollutants at Rockjr Mountain Arsenal,

Well, L.» C, Dure and X. L. Quentin.  1974.  Solubility in water of
  Insecticide chlorinated hydrocarbons and p'olychlorinated biphenyls in
  view of water pollution.  Z. Wasser Abwasser  Forsth,  7(6):169-175.

Zabik, M. J., R, D. Schuetz, W. L. Burton, and  B. E. °ape.  1971.
  Photochemistry of bioactivc compounds.  Studies of ai ,-najor Dhotolytic
  product of endrin.  J. Agric. Food Chew.  19(2 );3C8-313.

                              29,  HEPTACHLOR
-9.1  Statement of ?rcbnb_le Fate'

    The ^ajor fate pf hepra<:hlor ir the solution phase of aquatic systems
     be hydrolysis to give 1-hydroxyehiordene (1-HC) with a half-life of
      1-3 days; i-HC wil! then be biotcansf ormed co give l-hydro-xy-2,3-
chlordene epoxlda (1-HCE).  Although literature information also indicates
Chat heptacnlor photolysis, volatilization, and sorption to sediments nay
also occur in aquatic environments, no data are available to'cos pa re these
processes with the hydrolysis transformation rate.

    T!"e toxic metabolite, heptachlor epoxide (HE), appears to be a minor
product of heptachlor transformations in aquatic systems; the amount of HE
foraed is dependent on uptake by organisms capable of effecting the
epoxidation ot heptachlor.

29.2 • Identification

    This chapter considefs only the heptachlor component of the technical
heptachlor product.  The technical product contains 72% heptachlor and 28%
related compounds, and has a melting range of 46 to 74°C (Martin 1972),

    The structure, alternate names, and, CAS and TSL numbers for heptachlor
are given below.
                                             Alceriace Names

                                             S 3314
                                             Velsicol 104      ' ,   ,

           CAS No. 76-44-8 ,
           TSL No. PC 07000

    '  Physical Properties

    The general  physical  properties of pure hep each 1'.-r  ar-  »p--/en below.
    Molecular weight

    Melting point
    (Martin 1972)

    Boiling point

    Vapor pressure  at  25°G
    (Martin 1972)

    Solubility  in water
    (Siggar and Riggs  1974)*
    (Park and Bruce 1968)

    Log octinol/water  partition


No data found

3 x ir~4 :orr
0.180 ppm at 25 °C
0.056 ppm at 2V29

*particle size  <  5,0 um

29.4  Suanaa£y_ of  Fa_t_e  Data

    29.4. 1  ^ho_toj.vs_is

         Although  licsrature  information indicates t'%-i£ hept^ch lor  rnav
undergo direct  photolysis  in  sunlight  anc! is .->!so su.i<, e^t i Tie  to  photo-
sensitlr.jd reacrlons,  f nsuf f ic •" ent  dan ire avaiUble r-  predict  half-
lives for photolysis in  sunlight.,

         Ehmann (1976) reported tnat  the i!c-;hiorinateil hcptachlor Isomers  '
and II and heptaehlor  epoxide (HE)  were nrod'tcts o* hept-ich Lor  fi!-ns  photo-
ly^ed-^t > 290  nn>.  The  cani*  author reported th.it photolysis of her.tachlor
film1 In sunlight  (June through'Septs-sher > puve the _^ge product ». H  i ri  yi".
               C1  Cl

         The photoisomer III has also feeen reported as  "he  product  of  b«?n-
asophenone-sensitlzed photolysis of heptachlor  in benzene  solution  (Rosen  e_t
al. 1969; Rosen and Slewierski 197C),  The same product was found  L(n  the
acetone-solvent sensitized photolysis of heptacnior (Onus 290 nti  of heptachlor  in  the
solid phase, as a film, in water-methanol solution and  in the  gas-phase
(Voilner e_£ a_i. 1971),  'In an eariler paper Vcilner e_t  a_l.  (1969)  reported
that photolysis of heptachlor fila at 254 nc gave  yields  of 4a%  and 43%  for
ill and an unidentified product, respectively.  Photolysis  of  heptachlor  in
aethanoi-water solution with a Pyr«x-fiitere«i  H« lamp (>  29'j rim)  gave  a
j^;S2 ratio of these products,aad in dioxane-water solvent  the product
ratio was 7;si,

         McGuire ejc aJL. (1972) have aeas 290 no) by .direct photolysis and possibly  by  i-.direct
photolysis if effective naturally occurri ig tiplet sensitizers  are  writ-
able in aquatic systems.  The uv spectrum for heptachlor given  by Oc-e  e_t
al. (1971)- does snow a tailing, aosorbance above 290 n;n, indicating  that
direct photolysis in sunlight can occur.  I'f the cuaTt^m yield  of '',,v23 for
formation of products 1 and I! is
                                                 wa-e-len-'.th  aid  solvent  and
represents the photodissociation of the vinylic O-L!  sond,  che  product  ,
quantum yield of •.;.'-2 3 say represent a-one fraction or a  larger  reaction
quantum yield (tnat is, products other than  I and II  may  nave ">ten  turned
but not detected.,,  Unfortunately, sufficient uv absorption s.eciral  data

are not available to calculate possible direct photolysis race constants
using the quantum yield data.  In aquatic systems the products I and II
should not be formed unless suitable hydrogen atom donors are present, but
the cage compound III may be formed.

    29,4.2  Oxidation

         Singh -(1969)' reported; that heptachlor is oxidized by CrO} in
acetic acid to give a heptachlor epoxide (HE)'With a melting point of
158~16QSC; this isomer is the toxic component that is formed by metabolic
processes.  Cochrane and Forbes (1974) found HE as well as other products
resulting from oxidatlve cleavage of the allylic double bond when hepta-
chlpr was oxidized wich Cr03«  Neither study provides information useful
in evaluating chemical oxidation of heptachlor In aquatic environments.

    29.4.3  Hydrolysis

         Several studies have shown that abiotic hydrolysis of heptachlor
occurs rapidly, with half-lives of 1-3 days under environmentally relevant
conditions.  The hydrolysis product is 1-hvdroxychlordene (1-HC).

         Detnayo {1972} has measured the first-order rate constant for hy-
drolysis of heptachlor at 29.88 +_ 0.03'G in' -mbuffered distilled water; the
rate constant was (3.00 + 0.08) x 10"^ rtr"^, which corresponds to a
half-life of 23.1 hours.  Eichelberger and lichte"nberg (1971) have reported
that of the heptachlor added to a sample of an Ohio river water maintained
at ambient rotw temperature, ordy 252 remained after one week, and no hep-
tachlor remained after 2.weeks; the product formed was identified as l-hy~
droxychlordene (1-HC).  The authors also found a second product evident in
the glpc trace after 2 weeks; this product, which was tentatively tdenti"
fled as heptachlor'epoxide '(.HE), was present after 4 weeks in 40% yield
based on initial heptachlor.  In parallel studies in water from the same
source it was found that dieldrin was formed from aldrln; thus, for hepca-
chlor, biological or other transformations ®av also have been occurring,
although the formation of HE from 1-hydroxychlordene via any biological or
chemical process appears to be without precedent.  Nevertheless, the 75%
loss of heptachlor after one week, when only 1-HC was present, corresponds

to a half-life of 3.5 days.  Given the variable and unspecified, but prob-
ably lower, temperatures of che latter experiment, this 3.5 day half-life
is in good agreement with the 23 hour half-life for hydrolysis of hepta-
chlor at 30°C reported by Demayo.  This information is also in agreement
with the observation of Bevenue and Yeo (1969a) that heptachlor in aqueous
solution at 22-25"C was markedly decreased in 24 hours and had almost
disappeared after 14 days.  Other studies have also reported rapid abiotic
hydrolysis (see Section 29.4.6).

         Although the pH dependence of the hydrolysis rate of heptachlor
has not been reported, it is reasonable to expect that the rate will be pH
independent by analogy to other allylic halides, specifically allyl
chloride (Mabey and Mill 1978) and hexaehlorocyclopentadiene (Zepp ejt al.

    29.4.4  Volatilization
         No studies that quantitatively define the volatilization of hep-
tachlor froa aquatic systems have been found, although several reports on
biotransformation studies of heptachlor have sited volatilization as a
probable loss process (Bourquin et al. 1972; Miles et al.  1969; Leigh

    29.4.5.  Sprpcion

         Only one set of papers on the sorptlon of heptachlor to particu-
late material has been found, and although the data do fit Freundlieh iso-
therm plots, the large deviation from unity of the Freundlich 1/n tern
makes the partition coefficients unreasonably large for a  simple
quantitative assessment of the sorptiott process.  Sorption does appear to
be an important process for heptachlor in aquatic systems, however.

         Huang and Liao (1970), Huang (1971, 1974) studied the adsorption
of heptachlor to three types of clay and on hutnic substances.  The adsorp-
tion of heptachlor onto the clays reached equilibrium concentrations within
two hours, whereas the humic materials required more than  five hours for
tquilibriunt to be attained.  The parameters 1/n and K for  the Freundlich
isotherm equation (x/m * KCl'n) are given below; the parameters have been
calculated from data in Huang (1974) using units where the sediment concen-
tration is-in ug heptachlor per gram absorbent and*C is the equilibrium
solution concentration in yg heptachlor per al solution:

     " Sorption System       •                 K         "  ;,    1/n

      Montraoril Unite (clay)            5.0 x 10& :           3.51
      Illite (clay)                     1.5 x 109            6,05
      Kaolinite (clay)                  1.5 x 108 .           4.49
      Leonacdtte                        4.1 x 105            1.69
      Humic acids (colloidal)           1.8 x 1G8       ,     2.42

The humic acids used were obtained by alkaline extraction of leonardite,
which was described as a coal-like, humus rich substance; the humic acids
were oven dried, pulverized, and then dispersed In colloidal solution, for
sorption experiments.  The authors did not comment on the unusuallv high
1/n values measured for the clay sorption experiments; the partition co-
efficient (K) values are exceptionally large for sorpcian of an organic sub-
strate to particulate and may be considered an artifact of the calcula-
tion owing to the use of large 1/n values.  Although the 1/n parameters f^r
sorption onto the hutaic substances ara not close to unity as would normally
be expected for Freundiich isotherm data, the experiment for Leonardite
where 1/n * 1.69 and K • 4.1 x 1G3, does indicate that sorption to
sedlsnents oay be a significant process.

    29.4.6  jtigacgumttLa t_ip_n

         Heptachlor shows strong tendencies for bioaccumuiation, with bio-
concentration factors on the order of 10  for several organisms.  Up-
take and bioaccumulatlon in biota of heptachlor may then be an important
process for heptachlor.  Low levels of heptachl-. r found in environmental
samples may result from sorption and accumulation in biota-where solution
hydrolysis cannot occur.

         Data obtained from a terrestrial-aquatic microcosm experiment of
Lu _e_t £l. (1975) indicate^eptachlor has bioconcentration factors of 2.1 x
104 for an alga, 3.7 x W4 for a snail, 3.1 x 10^ for mosquito lar-
vae, and 3.8 x 10^ for mosquito fish; the experiment was complicated,
however, by the toxicity of heptachlor to the organisms studied.  In this
paper, the authors also report that in a 3-day, aquatic microcosm axperT-
ment, heptachlor had bioconcentration factors of 1.8 x 10^ and 1.1 x
10-* foe snail and fish, respectively.  The authors did not comment on the
smaller bioconcentration values obtained in the latter experiment compared
with the terrestrial-aquatic microcosm experiment, however because  the ex-
perineats were run for 3 and 71 days, respectively, the difference may be
due to the exposure time and telaced transport/food chain differences in
the experiments.  Lack of sufficient exposure time to raach equilibrium is
also a reason Bowman e_t _al. (1964) did not find extensive uptake of hepta-
chlor by mosquito larvae, finding bitconcentration factors of, - 35 for 20
hour exposures.•

         Schiomel £t__a_l. (1976) reported bioconcentration factors of 3.6 x
10-* in a 72-hour test and 7»4 * 10  if a 96'hoar test with spot, an es-
tu&rine fish species.  They also cite prior studies showing concentration
factors ranging from 2800 to 21,300 for finfish and shellfish.  When the
     containing 3 ppm heptachior were transferred to clean water, about 50%

of the heptachlor war lost in 14 days.  It was not stated whether the loss
was due to excretion or metabolism,  or both; the heptachlor metabolite,
heptachlor epoxide,, was found in all fish, however.
                 (1969) has measured the distribution ratio K - CO/CW,
where C0 and Cy are the concentrations of heptachlor in hexane and
water, respectively (units are Ug/ml),  K was determined to be 1.11 x
10^ with data from five experiment* giving Co* (concentration of
heptaehlor in hexane extract of water phase) - 0.009 + 0.001 Wl, indicating
good precision of data.  The K value for heptachlor was greater than that
for lihdane, aldrin, dieldrin, heptachlor epoxide or DDT.

         Since 1-hydroxychlordene (1-HC) is formed in significant amounts
by hydrolysis of heptachlor, and 1-HC is subsequently metabolized to 1-hy-
droxy chlordene epoxide (1-HCE), data on bioconcentrmtion factors for these
products are of Interest here.  Lu et alt (1975) reported concentrations of
1-HC and 1-HCE in water and the various organisms in their terrestrial-
aquatic microcosm experiments using chlordene and heptachlor.  Except for
one value where the BCF for snail was 690, the 1-HC BCF values ranged from
197 to 57,  Bioconcentration factors for 1-HCS ranged from 119 to 5.7.
These data indicate thac bioac cumulation of 1-HC and 1-HC! will not be
appreciable in aquatic environments.  Bonderaan and Slach (1972) reported
1-HC was present in soil, fish and crops in a farming area, but found no
1-HC in tissue and fluids from human subjects in the area.

    29.<+«7  Bio trans format: ion and Biodegradatioa             :

         Several studies have shown that heptachlor can be bio trans formed
by oxidation to heptachlor epoxide (HE) or by reduction to chlordene.  The
consensus of several biotransfonnatin studies is, however, that abiotic hy-
drolysis of heptachlor in solution is more rapid than bioerans formation,
with the hydrolysis product itself then epoxidized biologically*  Signifi-
cantly, Che toxic heptachlor metabolite, heptachlor epoxide, doe* not
appear to be a major product in aquatic systems studied*

         Lu e_t a_l, (1975) have conducted microcosm studies to determine the
fate of heptachlor; bioconceneraticn data front these experiments are re-
ported in Section 29.4.7.  In an aquatic microcosm experiment where data
for products (or metabolites) were determined as the percentage of total
**C label in the organic fraction of the water sample, heptachlor was
found to decrease from 100% to 10% of total i4C material within one day.
After one djy. l-hydroxy-2»3. -chlordene epoxide (1-HCI) w«» present «s 50%
of the total ^C, rose to 702 at, the second day, and then remained con-
stant for the duration of the 11 day experlent.  ""he heptachlor hydrolysis

product, 1-HC, reached, a maximum of 101 of Che total ^C at one day and
decreased thereafter.  Heptachlor epoxide was'never found co be greater'
Chan 5% of total -^C; unidentified polar products/metabolites as de-1
tected by tic analysis reached a constant value of about iOX of the total
* C for days 5 through' 13.  The authors concluded that the major • pathway
•of heptachlor in aquatic systems is rapid, abiotic hydrolysis of heptaehlor
to 1-HC followed by metabolism Co 1-HCE.
         Although this experiment of Lu £t al. gives information on the
products and metabolites found in the aqueous phase of the microcosm, a
complete material balance cannot be made with these data because the
product information is only based on ^4C in the water sample, and does
not include losses due to soprtion or volatilization.

         Lu _et a_l. (1975) have reported several studies on the metabolic
products of heptachlor in conjunction with terrestrial-aquatic microcosm
studies.  Heptachlor epoxide (HE) was the major metabolite found in the
whole'body horaogenate of a salt marsh caterpillar after ingestlon of hep-
Cachlor-coated leaves; 1-HCE was the major metabolite in the caterpillar
excretions,  the authors also reported that HE was the major metabolite
identified in a sheep liver microsoiaal preparation to which heptachior was

         Lu jgt jd. (1975) also studied heptachlor in a 71-day terrestrial-
aquatic microcosm experiment.  Since the heptachlor was Introduced into the
microcosm by application to terrestrial plants which were then consumed by
a salt marsh caterpillar, the heptachlor probably entered the aquatic phase
of the microcosm by several routes (e.g., ""«af fragments, fecal matter).
A» in the aquatic faicrocosa system, 1-HC- concentrations were higher than
i-HC concentrations in the aqueous phase as well as in alga, snail, mos-
quito larvae, and fish.  The heptachlor metabolite in highest concentration
was HE, with levels of 1-1.6 ppm found in snail, fish, and mosquito larvae;
concentrations of 1-HC and 1-HCE ranged from Q»QVO,08 pp» and 0.27-0.12
ppa, respectively, in these species.  The larger amount of KE than 1-HCE

found In .this experiment, compared with the aquatic microcosm daC3, was
probably due to metabolism of heptachlor to HE in the terrestrial segment
of the cnirrocosn befoce the compound entered the aquatic phase, followed by
accumulation of HE in the various species.
         Miles ££ _al. (1971) studied the transformation of heptachlor in
aqueous solution containing a mixed culture of organises extracted from a
soil; the experiments described appear to have favored anaerobic condi-
Cions.  The only products found were - 7-141 chlordene and 1-7% 1-HCE, with
less than 0.04% HE.  The reduction of HE to 1-HC was also observed in this
system.  Although this latter pathway may be a source of 1-HC and eventu-
ally 1-HCE, the sequence (heptachlor •*•  HE * 1-HCE)  is probably less
important in aerobic systems than to the abiotic hydrolysis of heptachlor
to 1-HC and then transformation to 1-HCE.

         An earlier paper by Miles _e£ aJL. (1969) reports studies on the
aerobic transformations of heptachlor in solution by various organisms
isolated from soil.  Thirty-five of 47 fungi and 26 of 45 bacteria and
actinoaycetes were found to produce HE.   Besides HE, other products found
were 1-HC, chlordene, chlordene epoxide, 1-HCE and one unknown compound.
The highest yield of HE from heptachlor was 6% formed in 6 weeks.  The
chlordene was obtained in variable but small yields by bacterial dechlor-
ination of heptachlor.  The authors also established that 1-HCE was formed
by spoxidation of 1-HC rather than by hydrolysis of HE; of the 45 bacteria

and actinomycetes, 4 were found to epoxidize 1-HC to 1-HCE, whereas 43 of
the 47 fungi wera able to effect this transformation.
         lourquin jtt al. (1972) studied the transformations of heptachlor
in mixed cultures o£ Pseudompnas j»£ and individual isolates.  Chlordene,
1-HCE and HE were Identified as metabolites, nine unknown products were
also detected.  Although no yield data wer* given for metabolites, the
authors stated that the major pathway for loss of heptachlor was abiotic
hydrolysis to 1-HC followed by biotratis formation to 1-HCE.

         Leigh (1969) attempted to study the blotrans formation of hepta-
chlor in a solution of settled primary waste water with added yeast extract.
The average removal of heptachlor found in these experiments after 4 weefcs
was 95.3% compared with 99.5? loss in control solutions (containing ir
wastewater or yeast extract).  The authors concluded that no statement
could be made regarding the biodagradability of heptachlor; in fact, the
data may show that sorption by biota may make some heptach'or unavailable
for hydrolysis in solution.

         Sethunathan and Yoshida '(1973) have reported that under anaerobic
conditions a Clostridium sjs isolated from a flooded soil previously treated
with lindane was able to effect a 36X transformation of heptachlor (initial
concentration 12.8 ppn) in 24 hours.  lyengar and Rao (1973) found that the
fungus Asjpergilius niger was able to transform 12.5 ppm heptachlor to un-
detectable levels in 48 hours under aerobic conditions, but they noted that
the pesticide could not serve as the sole carbon source for the fungus.
The authors also state that unadapted organisms could not utilize
heptachlor, but that chlordane- adapted A._ niger cruld.  Neither of these
1973 publications discussed the possibility of 'he competing hydrolysis
process or identified any metabolites.

         Hlil and McCarty (1967) found that heptachlor was  completely
transformed to an unidentified product within one day in  chick,  biologi-
cally active wastewater sludge at 35°C.  The authors  further  reported  that
this product persisted for at least 42 days but was completely  removed
after 266 days.

         In sunuuary, the studies by Lu «_t al. (1975), Miles e t  al.  (1969,
1971), Leigh (1969), and iourqula a£ al. (1972) concluded that  abiotic hy-
drolys.is of heptaehlir to 1-HC will be more rapid than biological  trans-
formations, and that in aerobic systems, 1-HCE will be fo'rmed from  1-HC by
blotransformation.  HE does not appear to be a major  transformation product
of heptachlor introduced into aquatic systems.  The amount  of HE formed
will depend on tht amount of biota present and the capability of the organ-
isms to effect epoxidation to give HE.

    29.4.8  OChe.£:i:React:iQn3

         $<•» environmental processes other than those  previously discussed
have been reported to be sign»j»_gtf.u_dies. FiaId Studles,  and ^delling .

         Lu «c al» (1975) studied the fate of heptachlor  in aquatic  and
aquatic-terrestrial microcosm experiments.  Information from  these  studies
o« the bloaccunulation and transformation of heptachlor 's  presented in
Sections 29.4.6 and 29.4.7, respectively.

29.:   Reaction Products
    The following produces or metabolites have been repor.ed  for  heptachlor
in various studies.
                                                            Cl C»
       Photolysis (reduction)
      ,  Photolysis
      ,  Biotransfofmation
      .  Oxidation



    .  Biotransfonaation (of 1-HC)
.   Biotransforraation
29.6  Data Summary

    Table 29-1 summarizes the data on the aquatic fate of heptachlot.

                                                  TabU  ^9-t

                                          ry  of  Aquatic  Fate  of Hcptachlor
Envinxnental Sunmury
Proceaa - Statoaent Rate
Photuly&iik Mill photolyze in sunlight
• t undttternlited rate(s).
OxiJdiiut* No information available.
« ' -2-1
Hyutolyfeis RA|»id process iorbepiachlur 3x10 hr
in solution. at 10"C
Vol.n H i^at iuii Hay be ai) import aiitpiotese.
Sor |>t lou Probably an iM^ortant pcut-^ss,
but uo rel table data ava 1 lab It; .
ttl..,in-.iu.-lat iuii Will bU .c.uulaltf H uut hydiu) yzud.
Blot t jiisluruidtioii/ Sluw Lua|)d>r>-il LO liydrol ysi^ .
Half- Confidence
Ufe tH of Data
l-O days In High
a.  Til*  pririlofliiu nit c-nvi ronntfntal pioce^s  uhi<.U t^, thought  tu iktcrmint-  the J<*te of  th

 29.7   Literacare  Cited         •

 3evenue,  A.,  and  C.Y. Yeo.   1969(a).  Gas chromatograpnic characteristics
   of  chlordane,   I.  Effect  of an  aqueous environment, on  the heptachlor
   componenti   Bull. Environ.'Contam. Toxicol.  4(2);&8-76.

 Bevenue,  A.,  and  C.Y. Yeo.   I969(b).  Gas chromatographic characteristics
 .  of  chlordane.   II.  Observed compositional changes of the pesticide  in.
   aqueous and non-aqueous environments.  J. Chromatogr.   42:45-52.

 Biggar, J.W., and  l.R.  Riggs.  1974.  Apparent solubility of organochlorine
   insecticides in  water at various temperatures,   Hilgdrdia
 Bondennan, D.P, and E.  Slach.  1972.  Appearance  of_ 1-hydroxychlordene in
   soil, crops, and fish.  J. Agric. Food Chem.  20(2};32»-33i.
                         ,                                              >
 Bourquin, A.W., S.K. Alexander,  H.K. Speidel, J.E.  Mann,  and J.F. Fair.
   1972.   Microbial interactions  with cyclodiene pesticides,  Dev. Ind.
   Microfaiol.   !3:264-27o.

 Bowman, ?•!.€», F.  Acree, Jr., C.S.  Lofgren,  and M.  Beroza.   1964,
   Chlorinated insecticides;  fate  in aqueous suspensions  containing       ,
   mosquito larvae. Science.   146{3650):1480-14dl.

 Cochrane, W.P., and M.A. Forbes.   1974.  Oxidation products of  heptachlor
   and its rnetabolites-A chemical stud/,  Chemospnere 3(l}:41-46.

 Detiayo.'A.  1972,  Gas  chromatographic determination of the rate constant
   for the hydrolysis of heptachlor.  Bull.  Environ. Contain. Toxicol.
   8(4};234-237.'   ::       :         '            :

 Ehmann, J.L.   197&.  A  aodel system:  the photochemical interaction of
   heptachlor  and  selected epicuticular wax  components of  the tomatoe  fruit;
•;':  and pear leaf.   Ph.D. Thesis,'  Michigan State University.  ]v2 pp.
   University  Microfilms, Ann Arbor, Mich.,  Order  N"o. 76-18, o!5. (Diss.
   Abstr.  Int. B 1976, 37(2):597-598).

 Eichelberger, J.W., and J ,J. Lichtenb«rg.   1971.   Persistence of pesticides
   in  river water.  Environ,  lici. Technol.   5(b) ;541-5«*4,

 Fischler, H.-M, and Korte, F.  1969,  Sensibilizierce und unsensibilisierte
   photoisomeririerung von cvclodien-insektiziden.   Tetrahedron  Lett,
 :  1969.   (32):2793-2796.

Gore, R. C. , W. Hannah,' S.C. Pattacint, and T.J. POTTO.  L9?i.  Infrared,
  and ultraviolet spectra of seventy-six pesticides,  J. Assoc, Anal. Chew.

Hili, D.W.,iand P.L. McCarty,  1967.  Anaerobic degradation of selected
  chlorinated hydrocarbon pesticide*.  J, Mater Pollut. Control Fed.

Huang, J.-C.  1971.  Effect of selected factors on pesticide sorption and
  desorption in the aquatic system.  J. Water Pollut. Control Fed.
Huang, J.-C,  1974.  Water-sediment distribution of 'chlorinated hydrocarbon
  pesticides in Various environmental conditions.  Proc .  Int. Conf . • Transp.
  Persistent Chem. Aquat. Ecosyst.  11:23-30.

Huang, J.-C., and C.-S. Liao.  1970.  Adsorption of pesticides b%> clay
  minerals.  J. Sanit. Eng. Div., Am. Soc. Civ. Eng.  96(SAS) • 1057-1078.

lyengar, L. and A.V.S. P. Rao.  1973.'  Metabolism 'of chlordane and
  .heptachlor by Aspergillus tviger.  J. Gen. Appl. Microbiol.  194:321-324.

Leigh, C.M.  1969.  Degradation of selected chlorinated hydrocarbon
  insecticides.  J. Water Pollut. Control Fed,  41(11) (Pc.  2):R450-R460.

Lu, P.-Y., R.L. Metcalf, A.S. rfir«e, and J.W. ' Williams.  1975.  Evaluation
  of environmental distribution and fate of hexachlorocyclopentad lane ,
  chlordene, heptachlor1, and heptachlor epoxide in a laboratory model
  ecosystem.  J. Agric. Food Chem.  23( 5):967~973.

Mabey, W.R., and T. Mill.  1973.  'Critical review of hydrolysis of
  organic compounds in wacer under environmental conditions.  J* Phys,
  Chem. Ref, Data 7:383.

Martin, H,» ed.  1972.  Pesticide Manual, 3rd Edition.  British Crop
   Protection Council, Worcester, Engl .

McGuire, R.R., M.J. Zabik, R.O. Schuetz , and R.D.^Flotard,   1970,
  Photochemistry of bioactlve compounds.  J. Agric. Food  Chen.

McCuire, R.R, , M.J. Zabik, R.D. Schuetz, and R.D. Flotard.   1972.
  Photochemistry of bioactive compounds.  Photochemical reactions of
  heptachlor.  Kinetics and mechanisms.  J. Agric. Food Chem.

Miles, J.R.W., C.M.Tu, and C.R, Harris.  1969.  Metabolism  of heptachlor
  and its degradation products by soil microorganisms.'  J,  Econ. En cornel.
  62(6). 1334-1338.

Miles, J.R.W., C.M. Tu, and C.R. Harris.  1971.  Degradation of heptachlor
  epoxide and hepcachlor by a mixed culture of soil microorganisms. J.
  Econ. tntomol,'64(4):839-841.

Onuska, F.I., and M.E. Comba.  1975(a).  Isolation and characterization of
  the phocoalteracion products of cis- and transchlordane.  J. Assoc. Off.
  Anal. Cheta. 58(l):6-9.

Onuska, F.I., and M.E. Cooba.  1975(b).  Isolation and characterization of
  some methanoindene photoproducts.  Biomed. Mass Spectrotn.  2(4): 176-18.2.

Park, K.S., and W.S. Bruce.  1968.  The determination of the water
  solubility of aldrin, dieldrin, heptachlor, and heptachlor epoxide.  J.
 , Econ. Entotnol. 61(3) -.770-774.
Rosen, J.D., D.J. Sutherland, and M.A.Q. Khan.  1969.  Properties o'f
  photoisoners of heptachlor and isodrin.  J. Agric. Food Chem.

losen, J.D., and M. Siewlerski.  1979.  Sensitized photolysis of
  heptachlor.  J. Agric. Food Chertu  18:943.

Sehimmel, S.C. , J.M.'Patrick, Jr., and J, Forester.  1976.  Hepatachlor:
  Uptake, depuration', retention and metabolism by spot, (Leisoromus
  ganthurus).  J, Toxicol. Environ. Health 2( 1): 169-178.,

Sethunathan, X. and T. Yoshida.  1973.  Short comartinication:  Degradation
  of chlorinated hydrocarbons by clQ_stridium s_p.  isolated from
  iindane-anended, flooded soil.  Plant Soil.  38(3):663-666.

Singh, J. -1969.  Conversion of heptachlor to its epoxide.  Bui. Environ.
  Contain. Toxicol. 4(2);77~79.

Voetman, S.  1969.  Distribution of ratio of some chlorinated hydrocarbon
  insecticides between hexane and water.  Bull. Environ. Contain. Toxicol.

Vollner, 1., W. Klein and ,F. Korte.  1969.  Ecological chemistry.  XVIII.
  Photo rearrangement of' the components of technical chiordans.
  Tetrahedron Lett.  (34):2967-297Q.

Vollner, L., H. Parlar, W* Klein and F. Korte.  1971.  Beitrage zur
  okologischen-XXXI;  Photoreaktionen derr kontponenten des technischen
  chiordans.  Tetrahedron 27:501-509.

Zepp, E.G., M.L, Wolfe, C.t. Saughman, P.F,. Schlotzhaver, and'J.M.
  MacAllister.  1979,  Dynamics of processes influencing the behavior of
  hexachlorocyclopentadiene in the aquatic environrrtent.  U.S. Environmental
  Protection Agency (ERL), Athens, Ga.  (Paper presented before the
  Division of Environmental Chemistry, American Chemical Society,
  Washington, D.C. September 9-14, 1979.)

30.1  Statsment__ot_ P rgbable__Fate_

     Heptachlor epoxide Is resistant to chemical and biological transfor-
mations in aquatic environments, and half-lives of over several years are  r
probable.  Although sediment sorption and bioaccumulatlon are not appreci-
able, they may ultimately be relatively important processes in "lew of the
stability of heptachlor epoxide in the environment.  PhotosensitiEed reac-
tions and biotransfortnation in anaerobic sediments are possibly important
processes for eventual transformation of heptachlor epoxide in aquatic en-

30.2  Identification                                              '

     Heptachlor epoxide (HE) is known to exist in two isoneric forms (Singh
1969);  a metabolic product of heptachlor melting at 157-159°C and a pro-
duct formed in halohydrin reactions, melting at S6-89°C.  The 157-153*C
Isoiser, shown below, is the toxic chemical and is assume-I to be the domin-
ant isoiter and of concern in this fate assessment.  In sost literature,the
possibility of two isotners have not been recognized,

     The structure, alternate names, and CAS and TSL names of HE are given

                                             Alrernate Names
            Cl  C»
              Heptachlor epoxide
              4 ,7-Methanoi ndan-1,4,5,6,!7,3,
              Velslcoi 53-CS-17
     Heptachlor epoxide

     CAS No.  1024-57-3
     TSL Mo.  PB 94500

30.3,  PhysicalProperties •

     The general physical properties of HE are as follows:
     Molecular weight

     Melting point
     (Singh 1969)       '  '

     Boiling point

     Vapor pressure

     Solubility in water at 25°C
     (Park and Bruce 1968)
     (Weil et ai. 1974)
     (Biggar and Riggs 1974)


Mo data found

?Jo data found
0.350 ppm
0.350 ppra
0.200 ppm
0.110 ppas at 15 *C
  :   Log oetanol/water partition coefficient   Mo data found

30.4 , Sumaary of Jate Data

     30.4.1  Photolysis

          No data are available for estimating the photolysis half-life of
HE in aqueous systems.  Information on the photolysis of HE in the solid
phase suggests that photolysis of HE can occur in sunlight.  Other informa-
tion indicates that the direct and sensitized photolysis of HE can give the
same products.                       :;       :

          Fischler and Korte (1969) reported that HE in acetone solvent was
photoisowerlzed to a bridged chlorinated ketone structure;  the reaction
was attributed to acetone acting as a photosensltizer,  Ivie e£ aJL. (1972)
and Knox et_ a_l. (1973) repeated the photolysis of HE in aceton4 solvent
using light of wavelengths above 280 nm arid determined the photoproduct .as
structure II, which was formed through the intensediacy of photoproduct I.
A 50-60% yield ,of I was formed when HE on bean leaves .treated with the sen-
sltizer rotenone was photolyzed in sunlight for 4 hours.,  In this experi-
ment only 1% LI was formed,(and .more II was formed only after prolonged
exposure to sunlight.  No photoproducts were formed in the absence of ro-
tenone*  The authors further noted that direct photolysis of I on plant
foliage gave "slow" conversion to II, with a small (unspecified) amount of
I remaining after 7 days time.

          Graham et_ al_. (1973) studied the photolysis of HE as a solid
(pressed) desk, as a powder, and as 0.5? RE in a KBr disk.  After 121 hours
of exposure to sunlight in July (in Texas), 93%, 9K arj OZ HE remained In
Che solid disk, powder, and K3r disk, respectively.  Powdered HE exposed to
sunlight from January to nid-Sep^emner was 39^ photoisojierized.  The pro-
ducts isolated wera the same as those Identified by Ivie e_£ aj^. (1972) and
Knox e_t_ al. (1973), but Graham and ccworkers also identified a photoisomer,
III, which is the same product originally suggested by Fischler and Korte
     30,4.2  Oxidation
          Singh. (1969! reported that heptachlor is oxidized to- the toxic HE
Isonwr ">y CrOj in acetic acid in 60% yield;  the fairly high ylglj r>f HE
Suggests that it is not easily further oxi'dlzed under these rather drastic
conditions.  .Hoffmann and Eichelsdoerfer (1971) found that when ozone at
concentrations of 17 and 4 'm?/liter was bubbled through solutions of HE in
10? acetone  in water for 45 siinutes, losses of HE were 2^v> and 6*', respec-
tively.  Under the same conditions aldrin and heptachlor were completely
oxidized.  Although the latter paper suggests some susceptibility of HE to
oxidation by ozone, neither paper provides any tnformat 'on useful for
evaluating the rate of HE oxidation in the aquatic environment.

     30.4.3  Hydrolysis   .

          In a study of the hydrolysis of HE, Eichelberger and Lichtenberg
(1971) reported that HE was u .changed after 3 weeks at roopi temperatures In
a sample of river water (pH 7.3 to 8.0} and in distilled water,  Assuming
an analytical error at + 2.51 (100% recovery cited was rounded off, to the
nearest 5Z), a half-life of at least 4 years is calculated.

     30.4.4  Volatilization

          Mo information on the volatilization of HE from aquatic systems  ,
is available.
     30.4,5  Sorgtion

          Information on sorption of HE on soil and clay indicate that
sorption is not an extensive process^ in aquatic environments, with parti-
tion coefficients on the order of a few hundred.  However, the lack of HE
transformation by chemical and biological processes does indicate that
sorbed HE may ultimately ba transported to sediments where anaerqbic trans-
formation may slowly occur.

          Weil et al. (1973) measured the Freundlich isotherm parameters
for sorption of HE by a humic acid at 15°C and tound 1/n - 0.83 and K •»
209.  Sorption of HE on a soil (pH 6,7, 1.4% humic material) gave
Freundlich parameters of 1/n1- .71 and K » 400.  From the isotherm graph re-
ported by Hill and McCarty (1967) for sorption of HE on bentonlte clay,
Freundlich parameters of 1/n » 1.4 and K « 650 are, calculated, 'These
authors also noted that for the chlorinated pesticides they studied, sorp-
tion on algae was usually one to two orders of magnitude higher than on the
clay;  no specific data r'or sorption of HE on algae were given.

     30.4.6  Sioaccuinulation                 -    .

          Grimes and Morrison (1975) have reported HE bioc-ncentration
factors for 10 bacterial species;  eight factors were between 60 and 900,
with th« other two factors being 1,900 and 15,200.  The authors also re-
ported that HE uptake was,rapid, with near maximum concentrations achieved
in 15 minutes.  Lu e_t_ al. (1975) have found concentration factors of
roughly 2 x 10^t 8 x 105", and 6 x 10^ for alga, snail, and mosquito
fish in microcosm experiments.  In the microcosm, the HE concentrations
killed daphnla snd ajo*qulto larvae throughout the 43 day experiment; hence
the food chain system was perturbed and this disturb: ice may have affected

the uptake and depuration of HE in the system.  Ernst (1977) also measured
a concentration factor of 1700 for mussel and found a half-lif* of about 2
days for HE elimination when the mussel was transferred to clean water,

          The above data1 show that HE may be moderately accumulated in some
biological systems.      '                 »

     30.4. 7  Biotransf ortEatjion and Btodegradation

          HE Is very slowly trar-jformed by 'biological processes and is-
likely to have half-lives of several yea~s in many aqaatic systems.  Al-
though biotransforsuttions are slow, such processes nay be important fates
for RE since other transformation processes are also slow, and HE Tiay ulti-
mately be transported to aquatic sediments where anaerobic transformation

          Lu e_t_ al_. (1975) reported studies related to biot fans format Ion of
HE showing that HE is quite resistant to blotrangformations in aquatic
microcosms.  Thus, HE was present as 91%, 92%, and 70? of the total 14C
extracted from alga, snail and fish, respectively, after a 43 day microcosm
experiment..  The product l-hydroxyl-2,3-epoxychlordene (IV) was also found
as 3.5t, 8.7%, and 191 of the extractable ^C In  alga, snail, and fish,
respectively;  this product was probably due to in vivo blotransformation
since chemical hydrolysis of HE has been found to b« slow (see Section
30,4.3),  Lu e_t_ ad. (1975) also found chat HE was 962 recovered after
ingestion by a caterpillar and 98% recovered when subjected to nicrosomal

          Hill and'McCarty (1967) stbdied'the degradation of several chlor-
inated pesticides, including HE, in sewage sludges and found no HE loss In
aerobic dilute sludge and only slight HE loss in .anaerobic dilute sludge.
About a 50% loss of HE was found,in thick sludge after approximately .60
days.  The authors concluded that HE wa^ similar to dieldrin in Its resis-
tance to anaerobic transformation, with both epoxide* more stable Chan the
other chlorinated pesticides, Including aldrln, ODD, DDT, eidrin, and hep-

          Miles *£, a_l. (1971) examined the degradation of HE in aqueous so-
lution containing a Mixed culture of soil microorganisms and found that HE

was  lost  at  the  race  of  about  1% per week  For  a  12  week  experiment.   The
hydroxycnlordane product V was  Isolated  in this  experiment.'

     30,4.8  Other Reactions

          Mo reactions other  than  those  discussed  above  have  been  con-
sidered for" HE  In aquatic environments.   In  vfew of  the  lack  of  reactivity
of HE  toward biological and chemical  transformation  orocesses, anaerobic
reduction of HE may  be an important fate,  but  no studies on this subject
have been reported,

     30,4.9  Microcosm Studies, Fie Id Studies,_andMod_g lllng

          Lu «C ,al«  (1975) reported on the face of HE  in a terrestrial-
aquatic microcosm and in an aquatic microcosm.  These  experiments  were made
difficult by th* toxlcity of  HE to organisms  In the  microcosm because the
food chain was  interrupted.   Information from  this paper on bioaccumtlation
and blot ran*format ton Is given in  Sections 30.4.6  and  30,4,7.

.30.5   Data Suamary

     Table 30-1 summarizes the data on the aquatic fate  of heptachlor epox-

                                                             JO- 1
                                    Summary of Aquatic Kate of ll«utaclilor





 Sorpt Ion


 Blot rai>^format ion/
   R iod«H rad
 30.6   Literature Cited

 Biggar,  J.W. ,  and 1.8. Rigga.   1974.   Apparent sb.luM Uty  of dfganochlorlne
   insecticides in water ac various temperatures.  Hilgardia 42(10}:383-391.

 Eichelberger,  J.W. ,  and J.J.  Lichtenberg.   1971.  Persistence of pesticides
   in  river water.  Environ,  Sci,  Technoi,  5(6):541-544,

 Ernst, W.   1977.  Determination of the bioconcentration  potential of marine
   organisms -  a steady state approach.  Chetaosphere 6(11):731-740.
        i                                '
 Fischler,  H.~M, and Korte, F.   1969.   Senaibilisierte  and  ursensibi lisierte
   photoisooeririemng von cyclodian-insektiziden.   Tetrahedron Lett.
   1969.  (32): 2793- 2796.

 Graham,  R,Et.  K.R.  Burson, C.F. Hammer,  L.B,  Hansen, and C.T. Kenner.
   1973.   Photochemical Je compos It ion  of  heptachlor epcxide.  J.  Agrlc,
   Food Chem.   21:824-834.

 Grimes,  D.J. and S.M. Morrison.  19?5,  Bacterial  bioconcentration  of
   chlorinated  hydrocarbon insecticides from aqueous sytems.  Microb. Ecol.

 till,  D.W. , and P.L.  McCarty.   19&7.   Anaerobic degradation of selected
   chlorinated  hydrocarbon pesticides.   J.  Water Pollut.  Control Fed.
   39(8); 1259- 1277.

 rtoffraann  J. , and D.  Eichelsdoerfer.   1971. Effect of  ozone on
   chlorinated-hydrocarbon-group pesticides in water.  Vc«  Wasser
   38:197-206.          ,                         •
     ,  CfW.',  J.H,  Knox,  S.  Khalifa,  I.  Ya«namoto and J.E.  Casida.   1972,
   Novel phot opr .ducts  of heptachlor epoxide,  trans- chlordane and
   trans-nonachlor,   Bull.  Environ.  Contao.  Toxicol,  7(6):376-382.

-Knox.'J.R. ,  S. Khalifa, G.M.  Ivle,  «nd J.E'.  Casida.  1973.
   Characterization  of  the  phptoisooers from cis-  and trana-chlqrdanes,
   trans-nonachlor and heptachlor  epoxide.   Tetrahedron.   29:3869-3879.

 Lu,  P.-Y. ,  R.L. Meccalf, A.S.  Hirwe, and J.W.  Williams.'   1975.   Evaluation
   af environmental  distribution and fate  of  hexachlorocyclopentadiene,
   chlordene, heptachlor, and heptachlor epoxile In a laboratory  model
   ecosystem.  J.  Agrlc. Food  Chem.   23(5 ):
 Miles,  J.R.W. ,  C.M,  Tu,  and  C.R.  Harris.   1971,   'Degradation  of  heptachlor
   epoxide  and heptachlor by  a  mixed  culture  of  soil  microorganisms.   J.
   .icon. Hntpmol.   64(4)j839-841.

Park, K.S., and W»M. Bruce.  1968.  The determination of the water
  solubility of aldrin, dieldrln, heptachlor, and heptachlor epoxide,  J._
  Icon, Entoaol.  61(3 };7TO-77i.

Singh, J.  1969.  .Conversion or heptachlor.to its epoxide.  Bull. Environ,
  Contain. Toxicoi.  4(2):77-7°.,

Well, L., G. Dure, and K.E. Quencin.  ]<*71.  Adsorption of chlorinated
  •hydrocarbons to organic particles ar.J soils.  Z. Wasser Abwasser Forsch.

5-'£il, L.,'C. Dure, and K.E. Quentin.  1974.  Solubtltiy In water of
  insecticide chlorlnate'd hydrocarbons and polychlorinated biphenyls  in
  view of water pollution.  Z.  Wasser Abwass«r Forsth.  7(6):169-175.

         31.  HEXACHLCMOCYCLOHEXANE  (:i-,  3-,  and  ^-BHC  isomers)
 31.1  Statement  o£?robablg
      The  fate  of  the  3-,  3-,  and  "5-BHC  isotaers  In aquatic  systems  Is  de-
 termined  by  their availability  to biotrans format ion processes.   Although
 sorption  to  suspended sediment  and biota  is  not  extensive, sorptiun  Is
 probably  an  Important process for ultimately transporting  BHC  to anaerobic
 sediments where transformations occur.  As for  lindane  (see Chapter  32),
 hydrolysis,  oxidation,  and  photolysis are not important  processes  for the
 BHC Isoners  in aquatic  environments,

 31.2  Idejtt 1 f t ca t i on

      Brooks  (1974) and  Cunther  (1971) prdvided  excellent summaries on the
 preparation, separaCion,  structure, and characteristics  of the  1,2,3,4,5,6-
 hexachloroeyelohexane isotners.  According to Deraozay and Marechal  (1972),
 BHC has  17 optical or stereoisomeric forms,  but  only 9 of  these  are  ener-
 getically feasible.   Five of  these isomera are  given below along with the
 approximate  percentage  ranges found In  technical  BHC preparations.   Because
 of  Its acute toxicity,  most  studies on  BHC have  focused  on the  T^HC
 isooer, also called  lindane;  lindane is the  subject of chapter  32.   Lindane
'isomerizes to  the :*-, 3-, and ""-BHC isoraers  by  biological  proceses and
 to  S-BHC  by  photochemical reaction (see Chapter  32).

      The  structure a: d  orientations of  ~-t, 3.  and  5-BHC are given below,
 and the ! CAS  and TSL number,  and alternate names  for the  isomers  are  given
 on  the following  page.

                  :       _                                 Percent  of  Isomer
                             Orientation of                   in  Technical
  :        . '    •.   •          Cl atoms1 on  ring     Is oner     _ |HC_





v( lindane)  10 - 15

Is oner
GV 35000
Gf/ 43750
3 .'9- 86- 8
GV 45500
                               Alternate J^ajnes:

                               Benzene hexaehloride

31.3  Physical Properties

     The physical properties of a-, 3-, and 5-BHC are shown below.

     Molecular weight              291

                                   'a            1           i

     Melting point
     (Gunthsr 1971)              15 7-158'C      .309'C       138-139'C
     (Metcalf 1955)              159-160!>C    309-310'C     138-139'C

     Boiling point !                        So   data   found

     Vapor Pressure (in turrj at 20°C
     (Baiaon 1947)*            2.5 x 10~5   2.8 x KT7     1,7 x 10~5
     (Slade 1945)                2 x 10"2    5 x KT1      2 x 10"2

     Solubility in water (in ppm) at 25°C
     (Kurihara et al. 1973)t    1.21-2.03    0.13-0.20    3.64-15.7
     (Weil «t £l. 1974)            2.00         0.240       31.4
     (Brooks 1974)                 1.63         0.70        21.3

     tog octanol/water partition
       coefficient at 25°C         3.81'        3.80        4.14
     (Kurihara et al. 1973)
*?referred values, based on agreement of data of Balson with data of other
researchers for lindane; s*>e Section 32.3.
fThis reference reported solubility values  at 28°C.

•^*<*  Summary of Face Data

     31.4.1  Photolysis

          Kawahara (1972) reported that the rates of disappearance of the
BHC isomers were in the order ot > y > 3 > * when dissolved in water at
concentrations Q.Ql 'to 5.0 ppm ami exposed to sunlight;  'half-lives ranging
Iron 4 to 6 days for a-BHC to 10-22 days for i-BHC were reported.  These
data are highly suspect,, however, because experinental details were larking
(no controls described) and the graphical presentation of data which shows
a, rapid decrease in BHC concentrations within 10 days followed by a slower
loss oqC to 48 days.  Gins^urg (1953) reported that the toxicity of a BHC
emulsion was lost alter 12 days exposure to sunlight (the toxicity was at-
tributed to ll.idane); as described for llndane (see Section 32.4,1), the
BHC isotners are not expected to photolyze rapidly in sunlight because of
the slight, if any, light absorption coefficients above 290 nra.  Roburn
(1.963) reported that the four isotners of BHC gave no reaction products
after photolysis with a 254 nm light source for 2-3 hours.  Roburn's ex-
perinents and the expected low absorbance of the BHC isomers suggest that
photolysis will not be an important process in the environment;  the re-
ported photolyses of BHC are likely due to adventitious processes such as
volatilization, sorption on glass, or photoreaction caused by impurities in
the SHC used.

     31.4..2  Oxidation

          No information was obtained concerning oxidation of any BHC
isomer in the aquatic environment.  By analogy to limited studies cn^the
oxidation of lindane, the other BHC isomers should be quite stable to oxi-

   '.  31.4.3  Hydrolysis

          No data are available on hydrolysis rates of the individual BMC
iioswrs although Information fro* one paper does indicate that the BHC (all
isoaers together) has a half-life Cor hydrolysis of more than two years.,

          Eichelberger and Lichtenberg (1971) examined the persistence of
SHC for 8 weeks in water samples from the Little Miami River and in
distilled water; the pH of the river water varied froai 7.3 to 3.0 during
the 8 week period.  No change in SHC concentration was found during chis
time.  Assuming a maximum analytical error of 2,5% (recoveries reported
were rounded off to nearest 5%), the half-life for WC under these
conditions must be at least 4 years, indicating that the &HC isomers are
quit* stable to hydrolysis.

          Crlstol (1947) has studied the hydroxide promoted hydrolysis
(elimination of HC1) of the a-, 3-, 7-, and £-BHC isoraers at 20,l°d in a
solvent of 76% ethanol in water.  The second order r'ace constants for Che
elimination of HC1 from the s-, 3 - and v-isoroers (latter is lindane) were
1.7 x 1Q~1, 3 x 10~6, and 4.4 x iO~iM~isec~1,  respectively; HC1
elimination from 5-BHC was too fast to be measured under the experimental
conditions.  Although the data do show the relative reactivities of the
individual BHC isomers, these data are not applicable to an environmental
assessment because of the unknown effect of the high concentration of
ethanol in the reaction solvent.

     31.4,4  Volatilization

          There are no reliable data with which to estimate a half-life for
volatilization of BHC isomers from aquatic environments.  Although loss of
SKC through volatilization has been addressed  by several research groups
investigating biocransformation or bioaccumtilation of BHC isomers, the in-
formation obtained cannot be directly compared or even used to decide
whether volatilization of BHC can be an important process in aquatic

          Tsukano (1973) has reported on the loss of BHC isomers from aque-
ous solution at 25°C (see Section 31.4.5).   Data presented (in graphical
form) showed that 0, 25, 75, and 752 of the ct-, y (lindane)-, --, ard
Q-BHC isotners, respectively, remained after 2  weeks.  No experimental
cetails were reported tor these axperlaents, although the graph did also
show that about 80%  of the water had evaporated after the two week
experiment.  The author suggested volatilization as the process responsible
for the loss, but no data or information useful for actually .demonstrating
volatilization or estimating volatilisation rates in aquatic environments
were reported.  Although lacking any\ specific  data, Xewland e_t a_l. (1969)
also implicated volatilization <_f a-BHC from aquatic systems as the reason
why significant aoounts of 5-BHC are not formed biologically in the
sequence yBHC (lindane) -» *»-BHC * Ct-BHC (see  Section 32.4.6).

          Ernst (1977) reported data seemingly contradictory to this
information that suggests the facile volatilization of 3-3HC.  He aerated
aquaria containing several chlorinated pesticides, including s-BHC and
lindane, and quantitatively recovered both BHC isoners from water after 67
hours.  The experiments were conducted as controls in bioaccumulation
studies.  This information, as well as that cited in the discussion of
lindane volatilization (see Section 32.4.4), would suggest that
volatilization of BHC (at least the x and Y-isomers) is not a facile
process.  Since chemical and biological processes appear to be rather slow
to transform BKC in aquatic environnents, the  question of volatilization
rates of BHC isoners is important and should be pursued.

           31.4,5  SprptJLpjn

               Only one paper  on the  sorption  of  Che  individual BHC  isomers  to,
     s-ediments was found and in agreement with data on  the Y~BHC  isomer  (lindane
     see  Section 32,4,5), this information  indicates  that BHC should not  be
     sorbed  extensively onto biota and sediments.  Because of the  lack of
     chemical and biological transformation  In aerobic  systems, however,  sorp-
     Cion onto particulates with subsequent  deposition  and transformation in
     anaerobic systems may be  the most important  fate for BHC.

               Tsukano (1973)  reported experiments on the trans location  of BHC
     isomers from standing water to a sediment-soil layer and the  partition  of
    /SHC  isomers between soil  and water.  Insufficient  experimental  information
   /was  providea in  the translocation experiments for  use in a quantitative
,,/   environmental assessment  except  to show that equilibrium between soil and
     water was attained in about 7 days, and that the amount of BHC  tsomer
     sorbed  in the soil layer  compared with  the amount  in water was  In the order
     6"  *  8 > a; information on the Y-SHC isoner (lindane) was complicated
     by biotransformation. ,Tsukano also presented Preundlich isotherm plots for
     sorbtion of SHC  Isomers from water to  two soils.   For the soil  containing
     1.91 carbon, values 6f 1/n were  approximately 0.91 and K was  between 10 and
     30 for  the a, 3,,and t5~BHC.  Foir a soil with 5.2%  carbon, 1/n was
     approximately 0.71 to 0.83 and K ranged from 30  to 120 for the  four  BHC
     isomers.  These  rough data, which were  calculated  from the isotherm  plots
     given in the paper, indicate that the  BHC isoaers  will not be extensively
     sorbed  onto sediments in  aquatic environments.

           31.4,6  Bioaccumulation

               Intormation available  on the  bloaccumulation of a-, 8—, and 5-BHC
     is similar eo that for Y-BHC and indicates that  BHC isoraers are not  ex-
     tensively bioaccuraulated  in organisms  (see also  Section 32,4.7).  Concen-
     tration factors  vary among the four BHC isoners  in the ™ange  of about 10 to
     500,  depending on the isomer and organism.

               Schlarael «t_ ai. (1977b) reported that  ,>ink shrimp,  pinfish, and
     oysters accumulated BHC to concentrations that were 80, 480,  and 130 times
     the  concentrations of BHC in water.  Ernst (1977)  reported concentration
     factors in mussels of 106 and 100 for 3-BHC an<*  lindane, respectively.
     This  paper also  presented data showing  that both BHC isomers  are eliminated
     from the mussels when placed in  clean water, although the first-order rate
     constants and approximately 20-hour half-lives given by the author are
     clearly not correct when  compared to the  actual  data showing  the loss of
     the  pesticides from the mussels  as a function of time.  Only  43 ppb  of

a-BV"  of  the  initial 94 ppb  remained  in  the mussel after  2 hours  la  clean
wa'e*, with  12  ppb  remaining after 48 hours.   For  lindane initially  at  124
ppD  i-i Che mussel,  54 ppb and  18 ppb  remained  afcer  2  and 48  hours,

          Suglura et a_l. (19763 studied  the effects  of 3-BHC 'concentra-
tion In an aquatic  microcosm.  The relationship  between the concentration
of 0-BHC  (In  ppm) In the microorganism (bacteria and algae) and the  con-
centration of 3-BHC (in ppm) in the Bedlura was given by the following

          log(Corganlsins) -  (0.35)«log(CmedluJB)  +  2.54.'

This expression can be restated as

Although slightly dependent on the exponential  ractor  of Cme
          Tsukano (1973) reported chat BHC isomers  incubated  in a soil-
water Mixture were degraded in the order Y > a >  3"> 5 ; the amounts
remaining after 56 days were approximately 5*, 10%, 302, and  302,
respectively.  Since the addea presence of sodium azide In the soil-water
medium resulted in very little less of BHC, microbi-.l  processes were
assumed responsible for the observed losses of BHC,  Tsukano  also stated
that in another study in his laboratory, a flooded  soil experiment
tentatively identified 5-3,,4 ,5,6-tetrachlora-l~eyclohexane as a degradation
product of a-BHC.

      ,    Heritage and MacRae (19?7a) reported that wished cell suspension
°^ Closeridiun .si>henojLde3_ degraded ot-SHC and  liadane.  These  authors did
not give any kinetic tnfortnatlon for the transformation except to mention
that; no a-BHC was recovered after Incubation  for 4  hours;  in an identical
experiment, no lindane was recovered after 2  hours  of  incubation.  The pro-
duct from the transformation of a-BHC was 5-3,4,5,6-tetrachloco-l-cy-
clohexane, the same product reported by Tsukano (1973).  Although not im-
mediately relevant to aquatic systems, it 'is  of interest to note that
Stes .iwandter and Schluter (1978) have reported that 2-BHC is  a product of
y-BHC (lindane} metabolism in grass via the intermediacy of Ti-BHC.
Interconversion of BHC isoners may also be occurring in aquatic environ-
ments and may complicate conclusions on the fa:e of individual BHC isotners
in'Such systems,                                 :

     31.4.8  01!:_er React!pns

          So processes other than those listed above have been implicated
as Important in the fate of BHC in aquatic environments.  >3o  information
was found to indicate that isomerization of 8HC l.somers occur spontaneously
In the aquatic environment, other than the biotransfortnation  discussed in
Sections 31.4.6 and 32.4,6,

     31»4»t  Microcpsffl Studies.^	Field Studies, and Mode1 litig

          Ho microcosm or field studies on the at-, S-, or 5-BHC isomers
have been reported.  The chemical and physical properties of  the BHC
isoiuisrs are similar enough, however, so, that  the information  in Section
32.4.9 also roughly applies to 1-, 3- and 5-BHC.

31.5  Data Summary

     Table 31-1 summarizes the data on the aquatic  fate of the, BHC tsomers.

                                                  Table 31-1

                             *r'inrhry of Aquatic  Fate  of  Hexachlorucycloltexaiit:



Hydro1 ysiti

Volaiil izalioit

So r p 11 un

Nut an  i«tyori**iu process.

Not an  Important- process.

Nut an  t»puiI ant process.

IntuiKat iuu  con t r*idlc tory a*  to
itow Important  prates^ it..

Input taut  f *.-r  t ran spur I tu
aiitteruble  st J it ent s ,
L tun        Not an  Iwporlaut ptutebi*.

*at luii/     f«|K»rtaut  pro«:K!*v that  varies
*t lun       with eav Ironmtut.
                                                                                                Clmtt iii'mfe
                                                                                                  ot Duta

a.  There  Is insuf fie lent  inforvat iutt in the  reviewed liter
31.6  Literature Cited

Balson. E. W.  194?.  Studies tn vapour pressure measurement, Part  III. -
  An effusion nanometer sensitive to 5x10"^ millimetres -of mercury:
  /apour pressure of DDT and other slightly volatile substances.  Trans.
  Faraday Soc. 43:54-60.

Brooks, G. T.  1974,  Chlorinated Insecticides:  Volume I:Technology and
  applications.  CRC Press, Cleveland, OH,, 249- pages.

CrisCol, S. J.  1947.  Th* kinetics of the alkaline dehydrochlorination of
  benzene hexaehloridfc isoners.  Th* mechanism of second-order elimination
  reactions.  J. Aaer. Ch«ffl. Soc. 69:338-342,

Deiaozay, 0., and G. Marechal.  1972.  Physical and chemical properties, In
  linda'-et  Monograph of an insecticide, E. Ulnann, pp. 15-21, K.  Schiller,
  Freiburg in Breisgau.

lichelberger, J. W., and J. J. lichtenberg.   1971.  Persistence  of
  pesticide* in river water.  Environ. Sci, Technol.  5(6):541-544.

Ernst, W,  1977.  Determination of bioconcentration potential of  marine
  organisms. - a steady state approach.  Chemosphere 6(11):731-74Q.

Ginsburg, J. M,   1953,  Rate of decomposition of the newer insecticides
  when exposed to direct sunlight.  Proc. Sew Jersey Mosquito Extern,
  A*soc. 40:163-168.  ,

Gunther, F.A.  1971.  Halogen derivatives of aliphatic hydrocarbons.
   Residue Rev. 36:34-77.

Htritage, A. D., and I. C. MacRae.  1977{a).  Identif icatioti of
  internrediates fortsed during the degradation of hexachlorocyclohexanes by
  Clo«_trijium sphenoides.  Appi. Environ. Microbiol. 33(6): 1295-1297.

Heritage, A. 0,, and I. C. MacRae.  1977(b).  Degradation of lindane by
  cell-free preparations of Closcridium sphenotdes.  Appl. Environ,
  Microbiol.  34(2):222-224.

Kawahara, T.  1972.  Studies of organochlorine compound residues  in crops
  and soil.  Part 20.  Disappearance of 8HC isoraers in water.  Bull* Agr,
  Cheia, Inspect, Stn, No, 12:49-51,

Kurihara, M., M, Cehida, T. Fujita, and M. Nakajima.  1973.  Studies on BHC
  isomers and related coapounds V,  Some ,physicochemical properties of 8HC
  isosers.  Pestic. Bior-hetu. Physlol.  2(4):3*1-390.

Metcalf, R. L.  1955.  Organic Insecticides.  Interseience Publishers  lac.,
      York, N.Y.
New land, L. W. , G. Chesters, and G. B. Lee.  1969   Degradation of Y-BHC  in
  simulated lake impoundments as affected by aeration,  J. Water poll,
  Control Fed.  41(5 )Pt. n:R174-R188.

Roburn, J.  1963.  Effect of sunlight and ultraviolet radiation on
  chlorinated pesticide residues.  Chetn. lad.  (1963) ;1555-155r> . ,

Sen inline 1 , S. C, , J. M. Patrick, Jr., and J. Forester.  1977(a).  Uptake and
.  coxlcity of toxaphene in several estuarine • organisms.  Arch. Environ,
  Cqntam.' Tox.  5:353-367.

Schimmel, S. C. , J. M. Patrick, Jr., and J. Forester.  1977(b). Toxlclty
  and bioconcencration of BHC and  lindane in selected estuarine animals.
  Arch. Environ. Contain. Tox. 6:355-363.

Slade, R. S.  1945.  The Y-isoner  of hexachlorocyclohexane (Gatnaexane).   An
  insecticide with outstanding properties.  Chens. Ind. 40:314-319.

Steinwandter , H. , and H. Schluter.  1978.  Experiments on lindane
  metabolism in plants.  IV,  A kinetic investigation.  Bull. Environ.
  Contain. Toxlcol. 20:174-179.

Sugiura, K. , S. Sato, M. Goto, an
                  32.  :Y- HEXACHLOROCYCLOHEXANE (LINDANE)"
32,1  Statement	of Probable Fa_te

    The fate of lindane in aquatic systems will be controlled by the avail-
ability of and to biotransformation processes.  Lindane transformation will
be favored in biologically rich, anaerobic environments.,  Although sorptlon
to suspended sediment and biota is not extensive, sorptlon is an important
process for ultimately transporting lindane to anaerobic sediments where
transformation occurs.  Hydrolysis and oxidation do not appear to be impor-
tant fate processes for lindane; data on the photolysis of lindane are con-
tradictory and confusing. Lindane is only slightly bioaccumulated in organ-

32.2  Identification

    Lindane is the gamma (Y) isomer of 1,2,3,4,5,6-hexachlorocyclohexane
and can be isolated from other BHC Isomers by solvent extractions and re-
crystallization (see also Chapter 31).  The commercial lindane product is
required to contain not less than 991 of the Y"lsomer and to have a melting
point of at least 112°C.

    The structure, nomenclature, and CAS and TSL numbers for lindane are
given below.

                                             Alternate dames
Gammexane       '
    CAS No. 58-89-9
    TSL No. 5V 49000
*The name lindane will be used in this chapter instead of the y-isoner
nomenclature to minimize confusion with other is oners.

32,3  Physical Prcigertiea

    The general physical properties of llndane are given1 below.

    Molecular weight                         291

    Boiling point                            No data found
    Melting point
    (Martin 1972}

    Vapor pressure at 20°C*
    (Benchmark 1975)
    (Martin 1972)
    (Deaozay and Marechal 1972)

    Solubility in water
    (Masterton & Lee 1972)
    (Kur lhara' et, al. 1973 )
    (Biggar a»-'< Riggs 1974)
    (Weil e£ _al. 1974)
    (Bhavnaga.*y and Jayaram 1974)
                                             (3.3-2.1) x 10~4 torr
                                             9.4 x 1Q~6 torr
                                             1.6 x 10~4 torr
                                             7.52 + 0.04 ppo at 25°C
                                             5.75 to 7.40 ppm at 28aC**
                                             2.15 pp« at 15'Ct
                                             6.80 ppa at 25°C
                                            11.4 ppm at 35"C
                                             7.8 ppm at 25'C
                                            12 ppffl at 26.5*C
    Log octanol/water partition coefficient
      at 25'C                                3.72
    (Kurihar.j et_ al, 1973)

*Slade (1945) reported vapor pressures that are several orders of mag-
nitude higher than recently measured data.
**M«asured by several procedures.
tAfter centr'fugation, particle size Is <5yau
32.4  Sumaar y of
                      Pa t .a
    32.4.1  Phptolysia

         Llndane is a saturated, chlorinated cyclic hydrocarbon structure
and should have little, if any, uv absorption above the solar spectral re-

 giott cutoff  at  ^290  nm.   In  spite  of  this  limited  light absorption,
 several  papers  have  reported that  Itndane  is  photolyzed in sunlight.   Lac*
 of  sufficient experimental  information  in  these  papers.makes any conclusion
 on  the•importance  of lindane photolysis  in the environment very tenuous,

          Steinwandter (1976b)  reported  that photolysis of lindane at  >230
 nm  in petroleum ether, acetone,, or water gave small  amounts of 2-BHC,  and
 prolonged photolysis of  lindane in aqueous solution  gave products that
 could not be extracted with petroleum ether.  The  authors suggested that
 such photoisonerization  was responsible  for finding  a-BHC on grass that had
 been coated  with lindane and then  dried  in the sun.   Subsequently, however,
 Steinwandter (1976a) and Stelnwandter and  Schluter (1978) reported that
 a-BHC was a  metabolic product  in grass  (see Section  32.4.6).  Thus the pho-
 toisotnerizatlon of lindane  to  a-BHC under  environmental conditions is not

          ,Zabik  and Leavitt  (1976)  report that a  Japanese researcher has
 found that 1HC  isoners photo react  in  sunlight in the order of n > 3 >
 lindane > o. Ginsburg 11953)  reported  that a lindar.e emulsion lost 50% of
 its toxicity to mosquito larvae after 6  days  of  sunlight exposure and was
 nontoxic after  11  days of exposure.  Roburn (1963)^  however, found that BBC
 isooers did  not give any reaction  products 'on exposure to 254 ran light,
 Thus, the photolysis of  liadane in the  aquatic environment is still in
 question, and available  information is  insufficient  to decide whether pho-
 tolysis is an iaportant  environmental process.   Especially perplexing is
 the almost certainly small  uv  absorptin of iindane in the ..olar region
 which should result  in low  photochemical transformation rates.,

     32.4,2  Oxidation

          No  information  was obtained  concerning  the  oxidation of lindane  in
 aquatic  systems.  Hoffmann  and Eichel^doerfer (1971) report that ozone bub-
 bled through solutions of lindane  in  hexane or water-acetone solvent  gave
..very little  loss of  lindane under  conditions  where heptacWLor and alcirin
•were-completely consumed,                       '  .                 •

          Leigh  (1969) reported attempts to oxidize lindane using the  chemi-
 cal oxidants chlorine, potassium permanganate, and potassium persulfate at
 pH values of about 2 and'6.   Chlorine or permanaganate solutions at 50 mg
 lit«r~*  produced no  reaction with  lindane  after  48 hours at 204C;r the
 sane concentration of persulfate removed only It lindane at pH 2.2 and ab-
 out 10%  lindane at pH 6.0.   No products  were  reported from the persulfate

    32.4.3  Hydrolysis

         No kinetic data have been found with which Co estimate the hydro-
lysis half-life of lindane in aquatic environments.  The available informa-
tion does indicate that lindane will, be reasonably stable in aquatic en-
vironments with half-lives at least greater than several months.

         Eichelberger and Lichtenberg (1971) monitored technical BHC in
samples of raw river water from Little Miami River, Ohio, and in distilled
water at a BHC concentration of 10  yg/1 over a period of 8 weeks.  The pH
of the river*water varied from pH 7 to pH 8 during che experiments.  No de-
crease in concentration was found.  3'HC contains about 132 lindane and this
amount or a significant change in it could have been detected; assuming a
2,5% change in BHC concentration could have been,detected, a half-life of
over 4 years is calculated.

         Gunther (1971) stated that lindane produced only 0.13% HC1 when
heated in steam at 102°C for one hour,  Cristol (194?) studied the alkaline
hydrolysis of BHC isomers in alkaline solutions of 76% ethanol in water at
20.11aC; the second order, base-promoted, hydrolysis rate constant reported
for lindane under these conditions was 4,4 x 10"*  M~l sec"*. Assuming
no solvent affect on the reaction, a hydrolysis half-life at pH 8 is calcu-
lated as 180 days.  Application of this half-life  to aquatic systems is
very tenuous because of the different solvent used.

  '  32.4.4  •Volatilization

         Mackay and Leinonen (1975) have calculated a half-life for volati-
lization of lindane from water of about 200 days using theoretical equa-
tions and assuming mass transfer coefficients' in literature for the ocean.
Spade ejt a_l. (1977) reported that volatilization  was n
    32.4.5  Sgrgtion

         King et_ a_l. (1969) studied the sorpcion of lindane on two species
of algae and three soils of different characteristic's; data were presented
as sorpcion isotherms without calculation of the parameters of the
Freundlich equation, Ss «'K'S^n, where S3 is the weight of lin-
dane aorb«d per unit weight of sorbept, Sw is the weight of lindane in
solution at equilibrium per unit weight of solvent (water), and n and ft are
constants.  From the graphs presented in the paper, a calculated value of
1/tt is' about 0.4, and values of X for the three soils range from 500 to 80.
The authors note that lindane equilibrium between soil and solution was at-
tained in an hour and that greater sorption was found on soils with higher
organic content as clay content.  Sorpticn on algae required 3 days for
equilibrium to be attained with values of K of about 30 and 50 for two
algae and 1/n equal to about 1.25.

         Well et al, (1973) iseasured the Freundlich isotherm parameters for
sorption of lleidane by humlc acid at 15 °C and found 1/n - 0,8 with K » 45.
Boucher arid tee (1972) studied the adsorption of lindane onto a sand aqui-
fer, and found that sorption was rapid during the first four hours, with
slight additional adsorption of lindane over the next 95 hours.  They also
found that less sorption occurred at 403C than at 5"C and that neither pH
(4.8 to 8.9} nor dissolved organic material in the water affected 'lindane
uptake by the sand.  From the graphs presented in the paper, a .sorption
isorhera K value of 0.35 is calculated.  Newland e_t_ al. (1969) studied the
sof?tion and biotransformat ion of 5 ppm lindane in water over a lake sedi-
ment (10" gin sediment in 3 liters of water) and found 44-57% adsorption of'
lindane in 18 hours, with slightly wore lindane sorbed over longer contact
tiaes.  More quantitative analysis of the sorption process was not possible
due to losses from biotransformat ion and volatilization occurring during
the experiment.

      ,   ftemellnk *_£ a_l. (1976), studied the fata of lindane and DOE in a
flooded quarry and found that although  lindane was rapidly distributed
throughout the water column, it was only slightly adsorbed by the bottom
amd of this, artificial lake; 95? of the lindane accounted for after 6
months was in the water layer.  Although the analyses for lindane in mud
layers are uncertain, the amount in the mud layer was definitely low.   In
the same quarry 94% of DDE accounted for was in the mud layer.  (See Sec-
tion 32.4.9 for further information on this iindane experiment.)

    32.4.6  Bioaccupulati on

         Grimes and Morrison (1975) reported lindane concentrations in bac-
teria that were ^ 10 to 300 times the concentations in the supernatant cul-
ture media; 5 of the concentration factors were 20 or less, 5 factors were

between 20 and 100, and 3 factors were greater than 100,  Ernst (1977) com-
puted concentration factors of 130 to 170 for mussels raised in seawacer.

         The fate of ^4C-radiol ibelled iindane in a terrestrial-aquatic
microcosm has been reported by Metcaif e£ al. (1973),  Altar 33 r"jys 92% of
the i"'C accounted for in mosquito fish was present as Itndane.  Of the
A1*C recovered from snail, only 20% was Iindane, with 70% present as '{-
pentachloracycljhexane.  A polar metabolite (i.e., nonmobile in the tic an-
alysis) was found in snail, fish, and alga and which the authors suggested
to be trichloirophenol( s).  lr\ an extension of this microcosm experiment,
Sanborn (19~4) round that added Aroclor 5t6Q caused Iindane concentration
ratios *o be slightly higher; 2,4,6- and 2,4,5~trlchlorophenol were iden-
tilled as iindane metabolites in this study.  These authors also reported
concentration factors of as much as 810 for daphnia, 125 for mosquito
larvae, and 233 for aosquito fish exposed to Iindane in small microcosms.

         Hamelink _e_t al. (1976; 1977) reported that organisms in a flooded
quarry accumulated Iindane, but to a. lesser extent than DDE (see Section
32.4.9).  The zooplankton rapidly accumulated Iindane, seeming to reach
equilibrium after five days.  The Iindane content then declined as the con-
centration in the water declined.  The Iindane concentrations in fish also.
appeared to be in equilibrium with the water after five days; rhe concen-
tration factor was (7,68 •*-_ 4.41) x 10 .  The concentration factor for
Iindane in zooplankton ranged from 170 to 448,

         Gakstatter and Wiess (1967) found that fish exposed to Iindane
reached equilibrium within a few hours, and the Iindane was eliminated in
less than 2 days ^sfter the fish were transferred to clean water.  Based on
this information and their work, Hamel ink _ej: al. (1977) suggest that Iin-
dane asay undergo exchange between fish and water with half-lives of 3 to 6
t^rafLsf ornat icn and 3 iodegradat ion
         TV.- results of numerous diverse studies on the biological trans-
formations of Iindane suggest that Iindane aay be tr'ansformed with half-
lives on the order of several days to more than a year when introduced into
biologically rich, aquatic environments.  Some papers that demonstrate the
varying ease of transformation and products 'formed from Iindane biotrans-
format ton are summarized below.  It should be noted, however, tnat the -
diverse approaches to Iindane studies as. well as the source of' the micro-
organisms make comparison of the experiments difficult and that the pro-
ducts reported are more likely a function of what products were sought, the
organisms present, and when the sample was analyzed.

          Tu (1976) reported that 71 of 147 microorganisms isolated from a
 loamy sand soil were able to utilize lindane in solution as the sole carbon
 source after 6 weeks incubation as evidenced by cell growth and chloride
 ion production.  Thirteen microorganisms were chosen for further studies.
 Of these, four bacterial strains showed adaptation times of less than a day
 whereas another 3 bacterial strains required 3-7 days adaptation.   Among
 the metabolites identified by tic were:

          Y - 2,3,4,5,6-pentachloro-l-cyciohexane (Y-PCCH)

          Y - 3,4,5,6-tetrachloro-l-cyclohaxane Cf-TCCH)

          8 ~ 3»4,5»6-t*trachloro-l-eyclohexane (8-TCCH)

          pentachlorobanzene (PCI)

          1,2,4,5-tetraeMorobenzen* (1,2,4,5-TCi)

          1,2,3,5-tetrachlorobenzene (1,2,3,5-TCl)

 Oxilation of lindane by a PseudoBonaa sj».  was suggested since oxygen con-
 sumption was found during lindane metabolism.  This  experiment also found
 that the gseudoaonas sj« appreciably oxidized Y-PCCH, 3-TCCH, Y-TCCH,
 and 1,2,3,5-TCS but  slowly oxidized Y-TCCH,  PCS, and 1,2,4,5-TCB.   It
 should be noted,  however, that oxygen uptake during  lindane metabolism does
 not necessarily indicate that lindane is being oxidized, but nay be due
 entirely to the noraal  metabolism of nutrients other than lindane  in the
• system,

          Benezet  and Hatsuutura (1973) showed that lindane is transformed to
 Y-TCCH, Y~?CCB, and  che ot-BHC isooer in the  laboratory by a Pseudononas
 culture*  ci-BHC was  also formed in an oceanic sediment treated with
 lindane. Matsumura *jt £l. (1976) also reported that  of 354 bacterial and
 fungal isolates,  S3  enrobes,  and 13 anaerobes metabolized lindane.   In the
 1976 paper, the Matsuotura group further elucidated the work reported in
 their 1973 paper  by  identifying two metabolic pathways of Pseudomnpnas
 £u|ida, the first gave  Y'PCCH as, a aiajor product and the second, more
 complex, NAD dependent  it«tabolic pathway gave Y-TCCH and Y-BHC as  the

          The degradation of  lindane by Eacherichia' coli hag been reported
 to give Y-PCCH (Francis ej.fl..  1975).   Lindane degradation by Clo8_tr_id_tua

sphenoides has been reported to produce Y-TCCH (Heritage and MacRae 1977a;
1977b; MacRae jjt al. 1969} che latter species was reported to give greater
than 902 degradation of llndane In 2 hours under anaerobic conditions   >
(Sethunachan and Yoshida 1973).

         Hill aud MeCarty (1967) found that 1 ppm and 10 ppra lindane In a
thick anaerobic sewage sludge at 35'C was more than 932 transformed after
several days and that anaerobic processes, were more affective than aerobic
processes; the rate of lindane loss was very sensitive'to temperature

         Oloffs and Albright (1974) and Oloffs et aJL, (1973) reported chat
lindane incubated for 3 weeks in samples of river water and sediments was
about 80% degraded.  When this system was sterilized, more than 95% of the
li»dane was recovered after 12 weeks.  Data reported for the unsterile sys-
cem shoved that 2% of the recovered llndane was in the water and approxi-
mately 202 was in the sediment.  The sterile system had about 142' of the
recovered lindane in the water and approximately 80% was in the sediment.*
AH earlier paper by Oloffs _et al. (1972) reported that lindane incubated in
three different river waters (no sediment) for 12 weeks gave no loss of
lindane in two waters and only about a 202 loss in the other water.  The
authors scat? that these experiments show the importance of sediments in
the fate of organochlorlne chemicals in natural waters.

         Htwland _et aJL. (1969) studied the sorption arid biotransformation
of llndane in water over lake sediment.  The experiments were carried out
under both aerobic and anaerobic conditions.  Although biotransformation of
lindane was complicated by adsorption and volatilization in these experi-
ments, the anaerobic conditions clearly degraded lindane more rapidly than
the aerobic conditions.  The authors estimated' that in §7 days, 15,and 90%
degradation of lindane had occurred in the aerobic, and anaerobic experi-
ments, respectively,  the authors also report that a-BHC waa a product In
both aerobic and anaerobic experiments, with 5-BHC also found in the an-
aerobic system.

         Mtthur and Saha (1975) repotted that lindane was only 10% degraded
after 6 weeks incubation in an anaerobic flooded sandy loam soil and that
the major product was Y'3,4|5,6-tetrachlorocyclohexane (y-BTC) in 52 yield
(based on initial llndane added) along with about 1% Y-PCCH and smaller
amounts of trlchloro- and tetrathlorobenzenes.  The authors also noted that
chlirophenolic metabolites found in plant, Insect, and mammal lindane
metabolism studies reported by other researchers were not found In the
flooded toil work.  In microcosm experiments, Metcalf jst1 a_l. (1973.) and •
Sanborn (1974) did find chlorophenols, however.
•Experiments used 100 _ sediment and 150 ml water.

         Seland et_ ajl.  (1976) found  benzene as a metabolic product  of  lin-
dane in anaerobic clay  loan soil and in sewage sludge; a maximum  of 52
yield of Y-TCCH product was obtained (based on lindane added).  The authors
suggest that this 4.3 due  to Y-TCCH being  metabolized  to benzene.   In the
anaerobic soil system about 50% of the lindane was  transformed  in approxi-
mately $ weeks.  Haider and Jagnow (1975) found up  to 90% of  lindane was
degraded within 4 to 5  days when it  was applied to  an anaerobic mixed bac-
terial flora enriched from an arable soil.  As would  be expected  from the
results of other researchers, chloride ion was liberated from the lindane
structure more rapidly,than hydrogen or carbon atom fragments.  No Y-PCCH
was detected, and increasing the oxygen content in  the gas phase  decreased
the rate of lindane metabolism.        ,               '

         Steinwandter and Schluter (1978) and Steinwandter (1976a)  have re-
ported that lindane is  netabolieally Isomerized in  grass in two parallel
processes.  One process gives ct-BHC,  which then is  isomerized to 3-BHC.
The other process gives hexachlorobenzene (HCS) directly from lindane.  In
a control experiment, no HCB was formed f rotn a-BHC  under the  reaction con-
     32.4.8  OtherReactions

         Slddarantappa and Sethunathan  (1975) studied  the transformation  of
.lindane and 3-BHC in five flooded soils  and found  that  the  BHC  isomets were
 rapidly reduced, if the  redox potentials  of the  fooded  soils were  lowered
 to a range of -40 to -100 mv.  Mo discussion of  blotrangformat ion  was pre-

         Several papers  were found that  discussed  electrochemical  reduction
 of lindane as applied to anaerobic environment  (Beland  «t_ al. 1976;  Block
 «£ al_. 1977).  leland e£ a_l. (1976) found Y-TCCH and  benzene as products  in
 both reduction processes.

  ."  14.4.9  Microcosm Studies,, ,Field_ Studies,and  Modelling

         An excellent study of the fate  of lindane in a very oligotrophic,
 letitic lake aquatic system has been extensively  reported (Hamelink e_t_ al.
 1977; Ham*link and Maybrant 1976; 1973; Waybrant 1973).  In this study,
 equal cdneentrations of  lindane and DDE were added in late  May  to  a  flooded
 limestone quarry and the pesticide concentrations  in  the water, sediment,
 and  biota were subsequently monitored  for a year's tine.  The lake was
 thermally stratified during the summer, was intermittently  covered with  Ice
 in winter, and received  a large influx of sediment because  of a rainstorm
 that occurred one day after the pesticides were  introduced  into the  system.
 Some significant points  of  lindane fate  (in some cases  relative to ODE)  are
 listed on the following  page.

         1,  Most of the lindane in the system yas retained  la  the eoi
         nion until. Che fall turnover, 123-14-* days after application.

         2.  Lindane decreased 322 In 102 days and 50% In 123 days (.the
         percentage was averaged over entire water column).  On day 81, 70%
  ,  '     of the lindane accounted for was In the epilainion,  with  15% each
         in the sietallmnion and hypolimnion.  The lindane in the  latter
         region was assumed to have been transported by the  sedi.aierit-ru.ioff
         since tne concentration was the same from day 5 until  the turn-
         over .

         3»  Alter the fall turnover, lindane was fairly homogeneously dis-
1         tributed throughout the water, column.

         4.  Suspended sediment, collected in traps, contained  essentially
         no lindane compared to a high DDE concentration, Indicating that
         sorption to suspended sedinient and sedimentation was not 4 signi-
         ficant process.

         5.  Lindane in. the bottom mud attained a maximum concentration of
         about 2 ppb compared to .the maximum for DDE of 35 ppb.   Additional-
         ly, Che DDE was contained in the top few centimeters of  mud,
         whereas lindane rapidly diffused to the lower mud layers presum-
         ably because of llndane's solubility in the interstitial water in

         6.  Pesticide concentrations!In zooplankton were at a  Maximum at 5
         days and decreased thereafter as the pesticide concentration in
         water declined.  Concentration factors for lindane  in  zooplankton
         ranged from 170-488.  The lindane in fish (bluegill) reached equi-
         librium with the water in 5 days and showed concentration factors
         from (1.47-0.42) x 103.

32,5  Dat_a _Sumjaary

    Table 32-1 summarizes the data on che aquatic face of lindane.

                                                                    Table  12-1

                                                    of JWfMtic Fate of Y  - Nexachlorocyclohexane  (Undone)
  Process *






Bioatcumulat ion

Biotransformat ion/

Probably not an lapotCaiU process.

Probably nut an Important process.

Probably nut an iKwjriant pioceas.

Probably-nut an important process.-

laportaot for transport to anaerobic

Not a significant process.

The moat iBportant process, and th«
•ost rapid in anaerobic environ-
                                                                                                 > 200 days
  of Data







 a.  Ther*  is  Insufticicnt  information in the  reviewed  literature  to permit assessment  of  a
                                                                                                                     probable  fate.

 32.6  L i t e r a t u re_ C ijced_

'Beland,  F. A.,  S.  0.  Farwell,  A.  ?,,  Robocker,  and'R.  0.  Geer.   1976,
   Electrochemical  reduction  and anaerobic degradation of lindane,   J,
   Agri^ Food Chera.  24(4 ):'?53-7 56.

 Benchmark, 1975.   Draft  of  Preliminary Summaries of Literature Surveys of
   Benchmark Pesticides.   George Washington University Medical  Center.
   Draft  dated October 30,  1975.

 Benezet, H. I.  and F.  Matsumura,   1973.   Isoroerization of f-BHt ;o a i-BHC
   in the environment.  Nature  243-:480-481.

 Bhavnagary, H.  M., and H,  Jayaram.   1974.  Determination of water
   solubilities  of  lindane and  dieldrin at different temperatures.   Bull.
   Grain  Technol.  12(2):95-99.

 Blggar,  W. J,,  and R.  L.  Riggs.   1974.  Apparent solubility of organo-
   chlorine insecticides  in water at  various temperatures.  Hilgardia

 Block, A.M., W.  a.  Bhajan,  J.  Estevez, and L.  W. Newland.  1977.   The
   electrochemical  reduction model of anaerobic degradation of  the gamma
   isomer of 1,2,3,4,5,6-hexachlorocyclohexanp  (y-BHC).  J. Water Pollut.
   Control  Fed.  49(55:857-859.

 Boucher, F. R.,  and G. F.  Lee.  1972.   Adsorption of  lindane and dieldrin
   pesticides on unc'onsolidated aquifer sands.   Environ.  Sci. Tech.

 Cristol, S, J.   1947.  The kinetics  of.che alkaline dehydrochlorination of
   benzene hexachloride isooers.   The mechanism of second-order elimination
   reactions.  J.  Amer. Chem. Soc. 69:338-342.

 Demozay, D*, and G. Marechal.   19"T2.  Physical and chemical properties in
   lindane:  Monograph of  an  insecticide, £. (Jlmanrr, pp.  15-21, K.
   Schiller, Freiburg in  Breisszau.

 Eichelberger, J.  W. and  J.  J.  Lichtenberg.  1971.  Persistence of  pesti-
   cides  in river water.   Environ. Sci. Technol.  5(6):541-54&.

 Ernst, W.  1977.   t)rtermination of  the bioconcentration  potential  of marine
   organistas. -  a  steady  state  approach.   Chemosphere  6(11 ):731-74fi.

 Francis, A.'J. ,  R.  J.  Spanggord,  and G.  I. Ouchi.  1975.  Degradation of
   lindane  by gscherichia  epli.  Appl.  Mlcrobiol. 29(4 ):567-568.

Gakstatter,  J. H.,  and  C. M.  Weiss.   1967.   The elimination of'DDT~l4C,
   dieldrin-^C, and lindane-^C  from  fish  following a  single sublethal
   exposure  in aquaria,  Trans. Amer.  Fish.  Soc.  96(3);30l-307.

Ginsburg, J, M.   1953.  Rate  of  decoenposition of the newer insecticides
   when exposed to direct sunlight.  Proc.  N'ew Jersey Mosquito Extern*.
   Asaoc. 40:163-168.

Grimes, D.  J., and  S. M. Morrison.  1975.   Bacterial bioconcentration  of
   hydrocarbon insecticides  from  aqueous  systems.   Microb.  Ecol.  2(1):43~59.

Gunther, F. • A.  1971.   Halogen derivatives  of aliphatic  hydrocarbons'.
   Residue Rev, 36:34-77.

Haider, K.»  and G.  Jagnow.  1975.  Degradation of -carbon-U-.tritium-, and
   chlorine-36-labeled of-hexachlorocyclohecane by  anaerobic soil micro-
   organisms.  Arch. MicrobioJ. 104(2):113-121.

Haaelink, J. L,,  and R, C.  Waybrant.   1973.   Factors controlling the
   dynamics  of non-ionic synthetic organic  chemicals in aquatic environ-
   ments.  Tech. Rep.-Purdue Univ. Water  Resour.  Res. Cent.  44, 68 pp.

Haaelink, J. L.,  and i. C,  Waybrant.   1976.   DDE and lindane in  a large-
   scale model lenttc ecosystem.  Trans.  Am.  Fish.  Soc. 105(1):124-134.

Famelink, J. u.,  R. C.  Waybrant, and  P.,R.  Yant.   1977.  Mechanisms of
   bioaecumulation of mercury  and chlorinated hydrocarbon pesticides by fish
   in  lentic  ecosysteos.  Pages 261-268 _ln_  I.  M.  Suffet,  Fate of  pollutants
   in  the air and  water  environment.   Part  II.  John Wiley and Sons, Ncv
   York, N.Y.

Harper, L.  A., A, W. White, Jr., R. R. Bruce,  A, W. Thomas,  and  R, A.
   Leonard.   1976.   Soil and microclimate effects on trifluralin  volatili-
   zation.   J. Environ.  Qua I.  5(3);236-242.

Heritage, A. D.,  and I. C.  MacRae.  1977 (a)  Identification of  inter-
   Bsediates  formed, during the  degradation of  hexachlorocyclohexanes by
   C_le«t rid turn sphenoides.   Appl. Environ. Microbiol.  33(6); 1295-1297.

Heritage, A. D.,  and I. C.  MacRae.  1977 (b)  Degradation  of lir.dane by
   cell-free  preparations of CIdstridiugi  sphenoides.   Appl.  Environ.
;   Microbiol. 34(2):222-224.

Hill, D. W., and P. L. MrCarty,  1967.  Anaerobic'degradation of selected
  chlorinated hydrocarbon pesticides. • J. Water Pollut. Control Fed.

Hoffmann, J., and D. Eichalsdoerfer.  1971.  Effect of Ozone on
  chlorinated-hydrocarbon-group pesticides in water.  Vom Wasser

King, P. H., H, H. Yah, P. S. Warren, and C. W. Randall.  1969,
  Distribution of pesticides in surface waters.  J. Arner. Water Works
  Aasoe. 61(9)':483-486,

Kurihara,' X., M. Uchida, T, Fujita, and M. Nakajima.  19'3.  Studies on  BHC
  isomers and related compounds, V, Some physicocheK.ical properties of
  BHC'isotaers.  Pestic. Biochem. Physiol. 2(4): 383-390.

Leigh, G. M.  1969.  Degradation of selected cnlorinaced hydrocarbon
  insecticides.  J. Hater Pollut. Control Fed. 41(H)(Pt. 2j:R4SU-k46u.

Mackay,  D.,  and P. J. Leinonen.  1975.  Rate of evaporation  of low-solu-
  bility contaminants from water bodies to atmosphere.  Environ.   Sci. '
  Tech.  9(13):1178-1180.

.MacRae,  I. C., K. Raghu, and £. M. Bautista,  1969.  Anaerobic degradation
  of  the insecticide lindane by Clostridiuci sp.  Nature 2^1:859-ob'J.

Martin,  H.,  ed.  1972.  Pesticide Manual, 3rd Ed.   British  Crop Protection
  Council, Worcester, Engl.

Masterton, W. L., and T. P. Lee.  1972.  Effect of  dissolved salts on water
  solubility of lindane.  Environ. Sci. Technol. 6(lu):919-92i.

Mathur,  S. P., and J. G. Saha.  1975.  Microbial degradation of carbon-
  14-labeled lindane in a flooded sandy loam soil.  Soil Sci,

Matsumura, F., H. J. Benezet, and K.  C. Pa'til.  1976.  Factors affecting
  oicrobial  metabolism of y-BHC.  Nippon Moyaku Hakkaishl 1(1);3-8.  •

Metcalf, R.  L., I. P. Kapoor, P.-Y. Lu, C. K. Schuth, and P., Sherman',
  1973,  Model ecosystem studies of the environmental fate  of six
  organo-chlorine pesticides.   Environ. Health Psrspect. i:3J-i*-».

.Newland, Li  W., G. Chesters, and G. B. Lee,  1969.  L/egr.i'dation of  >-BHC in
  simulated  lake impoundments as affected by aeration.  J.  Water Pollut.
  .Control Fed. 41(5 Pt.11);R174-R188.

 Qloffs,  ?.  C. ,  L,  J,  Albright,  and S.  Y. SzeCo.  1972,  Face and behavior
 "i  of  five chlorinated hydrocarbons in  three natural waters.  Can. J.
   Microbiol.  18(9): 1393-1398.            ,             _            .

 Oloffs,  P.  C. ,  L.  J.  Albright,  S.  Y. Szeto, and J.  lau.  1973.  Factors
   affecting the behavior of  five chlorinated hydrocarboos in two natural
   waters and  their sediments.   J.  Fish.  Res. Board  of Can.
 Oloffs,  i*.C,,  and L.J,  Albright.   1974.  Transport of some. organochlorines
   in B.  C.  waters.  Proc. Int»,Conf» Transp. Persistent Chen.  Aquat,
   Ecosyst., 1:89-92.

 Roburn,  J.   1963.  Effect of sunlight and ultraviolet radiation on
   chlorinated  pesticide residues.  Chera.  Ind.  (1963) :1555-155fi .

 Sanborn, J. K'.  1974.  , The fate of select pesticides in the aquatic
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   EPA-66Q/3-74-Q25, 83  pp.

 SethunaChan, X.  and T.  Yoshida.  1973.  Short communication:  Degradation
   of chlorinated hydrocarbons by ClosjjrijijuQ s_p_. isolated from lindane-
   aaended flooded .soil.  Plaat  Soil 38(3): 663-666.

 Slddaramappa,  R. , and'N. Sethunanthan.  1975.  Persistence of  ganma-*BHO and
   beta-BriC  in  Indian rice soils under flooded conditions.  Pestic. Sci.\
   6(4):395-4U3.                                                          \

 Slade,  R. E.  1945. The y-i.aom&r of hexachlorocyclohexane (Gammexane).  An
   insecticide  ,with outstanding  properties,   Cheia. Ind. 40:314-319,

 Spacie,  A., J, Haielink, and H. Waybrant.  1977.  Application  of an
 .  evaporative  loss aodel to .stimate the  persistence of contaminants in
   lencic environments.   Fate .of Pollutants  in the Air and Water Environ-
•••-  saent,  ASTM,  STP 634.   Part 2:214-227.

 Spencer, W. F.  1975.   Movement of Di3T and  its derivatives into the
   atmosphere.   Residue  Rev. 59:91-117.

 Spencer, W. F.(  and M.  M. Cliath.  1973.   Pesticide volatilization as
   related to water loss from soil.  J. Envlrotu Qua!. 2(2}: 284-289,

 Steinwand cer,  H.  1976(a).  Zucn lindanm«tabolismus an pflanzen,
   Chemo sphere   5(2) ; 119-125.

Steinwandter,  H,  1976(b).   Beitrage' zur umwandlung der HCH-lsoniere durch
  einwirkung von uv-scrahien1.   I.  Isomerisierung des lindans  in a-HCH.
  Chemosphece  5(4):245-248.

Steinwandter,  H., and H,  $chluter,   1978.  Experiments on lindane
  metabolism in plants.   IV. A kinetic investigation.  Bull.  Environ,
  Contam. Toxiciol. 2Q:],7i-179.

Tu, C. M.  1976.  Utilization and degradation of lindane by soil
  microorganisms.  Arch.  Microbiol* 108(3):259~263.

Waybraat, K. C.  19/3.  Factors .controlling the distribution and
  persistence  of lindane  and ,DDE  in lentic environments.  Ph.D Thesis,
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  Mo. 74-15, 256.  (Diss. Abstr.  Int. B 1974, 35(l);367-368).

Weil, L., G. Dur.e, and K. E. Quentin.  1973,  Adsorption of chlorinated
  hydrocarbons to organic particles and soils.  Z. Wasser Abwasser Forsch.

Well, L., G. Dure, and K. E. Quentin.  1974,  Solubility in water of
  insecticide  chlorinated hydrocarbons and polychlorinated biphenyls in
  view of water pollution,   Z. Wasser Abwasser Forsch. 7(G):169-175.

Zabik, M.,  and R. Leavitt. •  1976.  Photochemistry of bioactive compounds, a
  review of pesticide photochemistry.  Annu. Rev. Encotaol,  21-; 61-79.

                              33,  ISOPRORONE
33.1  Statement of Probable Fata

    No quantitative information is available to assess the environmental
fate of Lsophorone.  The moderate water solubility of isophorone Indicates
that it should remain in the water column until transformation reaction'(s)
occur.  Biological and photochemical reactions could be important In re-
moving isophorone from aquatic systems.

33.2  Identification

    The structure, alternate names, and CAS and TSL numbers for isophorone
are as follows:
                                               cyc loh exe ne~ 1- one
                                             Is of or one

    CAS So. 78591        '                   '    ..   '
    TSL No. GW 77000

33.3  Physical Properties

    The general physical properties of Isophorone are as follows:

    Melcing point                            -8°C          .    '
    (Verschueren 1977)

    Boiling point at 760 torr         '       215 *C
    (Verschueren 1977)

    Vapor pressure at 20°C                   0,38 torr
    (Verschueren 1977)

     Solubility  in water .at unspecified        12000  ppsn
     (Verschueren 1977)

     Log  octanol/water partition               1,7
     (Johnson  1978}

 33.4   Sunmary of Fate Data

     33.4.1  Photolysis

          Isophorone  has a moda-ate ultraviolet  absorption  out  to  approxi-
 mately 350 nm,  with  a *max at 312 ma  (£  » 45  M~l ctn~^)  in  the  solar
 spectral  region (uv  Atlas).  However, no information  is  available  to
 quantitatively  assess the photolysis"  rate of  isophorone  in the environment.

          Several papers indicate that tsophorone is photochemically reac-
_tive.  Jennings (1965) reported that  photolysis of  isophorone  in  water  at
 wavelengths > 200 nm gave dimerization products; Chapman ^e_t aL_» (1967)  also
 found  diiaerizatlon products when isophorone was irradiated at  > 3UO nm  in
 organic  solvents.  Such dioerization  products are unlikely in  the  aquatic
 environment under the highly dilute environmental conditions•  Mettee
 (1967) reported that photolysis of isophorone at >  300  nm  in air-saturated,
 carbon tetrachloridf solution resulted in lass  of iscphorone;  phosgene
 (CUCl^)  was the only product reported and ii  probably due  t
    33.4.4  Volatilization

         The high water solubility and moderate vapor pressure of i-;o-
phorone indicate that volatilization from water is not a significant trans-
port process for tsophorone,

    33.4,5  Sorption

        , No Information is available on sorption of Isophorone Co biota or
sediments.  The high water solubility and moderate partition coefficient of
tsophorone Indicate that sorption is noc an important process, and that
Isophorone will probably remain in solution in aquatic systems.

    33.4.6  Biqaccuiaulation

         For the reasons cited in Section 33.4.5, bioaccumulation is not
likely to be an important process for isophorone,

    33.4.7  Biotransforaation and Biodegradation

         No information was found in biotransformat ion of isophorone in
aquatic systems.  Truhaut e_t_ ajL. (1970) reported that the all/lie methyl
group of isophorone was oxidized to a carbcxylic acid group whan industrial
isophorone was administered orally to rabbits; the product (see f, below)
was detected in urine, and .no other, products were identified nor were mass
balances obtained.  This metabolism of Isophorone indicates that biological
oxidation of isophorone may occur in aerobic environments, but no
Information was found to substantiate this possibility.

    33.4,8  Other^Reactions

       1  No processes ->ther than'those listed above are important  in  dete-
raining the fate of isophorone in aquatic systems.

33.5  Data Summary                                                >

    Table 33-1 summarizes Che data on the'aquatic fate of Isophorone.

                                                                               TabU JS-1

                                                                           at  &i|uat ic fc'aifc uf  Isuptioruins

                            fsLjce&i;*                        Stiifettttut                      Kate   ~                L i if IB t4                 t»f_ Udtt_a

                          Ptititoiyaita                 i*uyi»ibly liiiportaut«                  -                        -                       l^j>w

i^           •             Hyvliutysiti                 ENM^K  not itcc-ai .                      -                        -                       High
 I                -         .


                          VuJdt U U.it lull            Hat  iai^oirl^ot (trocvAit,

33.6  Literature Citei

Chaptaan, O.L., P.J. Nelson, K.W. King, D.J, Treckar, and A.A. Griawold.
  1967.  The aechanisrn of isophorone photodinerization.  Sec. Chen. Progr.
  28(3);167-17a.            •   _   .

Jennings, P. • 1965,  Photochemistry of itiOphorone.  Ph.D. Thesis.
  University of Utah.  94pp.  University Microfilms, 'Ann Arbor, Mich.,
  Ord«sr N*o. 65-4898.  (Oiss. Abscr. Inc. B 1965, 26(2);698).
Johnaon, H.  1978.  Calculation of the esciaated partition coefficient for
  isophorone.  Che-nist, Biorganic Chemistry Lab.  SRI-Iacernational.

Mettee, H. D.  1967.  Photochemical studies of 3,2-unsaturated ketones in
  carbon tetrachloride solution.  Can. J. Chen.  45(4) :339-3*»3,
Truhaut, »., H. Dutertre-Catella, ami S.-P. Lich.  1970.  Metabolijjratudy
  of an industrial solvent, isophorone, in the rabbit,
  Ser, D 271(15):1333-1336.    ••    ,
170.  Metabolijj^atud;
:.  C.R.  AcadT Sri.,"
Verschueren, K.  1977.  Handbook of environmental data on organic
  chemicals. Van Noscrand Reinhold Cotapany.  &59pp.

                                 34.  TCDD
34.1  Statement^	cf^^Probable	Fate

    A conclusion as to Che fate of TCDD'In aquatic systems cannot be pro-
vided ae this tins because of Incomplete information.  Photolysis of TCDD
in sunlight may occur' in less than a day if reactive organic substrates are
available, but no information is available on the reacticns of TCDD with
possible substrates in natural waters; other chemical transformation pro-
cesses do not appear to be important for TCDD.  Sorption to sediments and
biota and possibly bioaecumulation appear "o be important fates for TCDD,
with biotransformations having half-lives of more than 1 year in lake water
and sediments.

34,2  Identification

    TCDD is formed as a by-product under the conditions of synthesis of
polychloriuated phenols and products formed from them.  Kearney jet aJL.
(1973) noted that the amount of TC00 in1 the herbicide 2,4,5-T varies with
each batch and with each manufacturer.  Several investigations of the pos-
sible formation of TCOD from photolysis or biodegradation of polychlorln-
ated phenols suggest that the latter are not converted to TCDD under en-
vironmental conditions (Piimmer and Klingbiei 1971; Kearney _e_t aJL. 1972;
Helling ere aJL 1973;  Crosby et_ a_l. 1973).  Both Plimmer and Klingbiei
(1971) and Crosby _e_t al. (1973) have pointed out, however, that in organic
media, TCDD .formed during photolysis of the polychloriaated phenols may
actually be rapidly photolyzed and therefore not detected.  Recent litera-
ture has also cited combustion of fuels' and wastes as sources of TCDD;
while this literature 'does suggest that combustion may produce pol/chlorin-
ated dibenzedioxlis (PCDDs) , the T,:DO isomers are usually found in very
small porportions of the total PCDD.  A thorough material balance on the
sources of PCDDs and TCDDs is clearly needed to determine the important
sources of these pollutants.

    The structure, alternate names and CAS and TSL number for TCDD are as

                           i                  Al_t_ernate Naaes

                                             !  dibenzo-p-dioxin


    CAS No. 1746-01-6
    TSL No. HP 35000     '                    '                         '

34.3  PhysicalProperties

'  '  The physical properties of TCDD are as follows.

    Molecular weighc                         3.22

    Meic ing .point                          .  303-303' C
    (Crummect and -Stehl 1973)

    Boiling point                -            Mo data found

    Vapor pressure .                          Mo data found

    Solubility In water*                     0.2 ppb
    (Crummett and Stehl 1973)

    Log octanol/vacer partition coefficient  So data found
*No tenperature given, presumably about 20-253C.

34.4  Summary of Fate Data

  '  34-4-l  Photolysis

         TCDD has a uv absorption naxitrtuni at 30? am, however, no absorption
coefficients were reported , (Crosby and 'Jong 1977),  Although photochemical
transformations of TCDD have been studied in several laboratories, no
quantitative information is available to estimate a sunlight photolysis
half-life for TCDO in'aquatic systems.  TCDD in a pure state is photo-
cheraically stable, but It will photoiyze in sunlight with half-lives of
less than a few hours If dissolved in an organic film or solution while in
the presence of a hydrogen atom donating substrate (Crosby and Wong 1977).
Although these conditions may be rnet in some environmental situations, in-
sufficient Information is available on such interactions in aquatic systems
for use in predicting a reliable TCDD photolysis half-life fjr an environ-
mental assessment.

         Plinwier £t_ aJL (1973) reported that a TCDD suspension in distilled
water was unchanged when Irradiated with i sun lamp, however, ^h*n TXD*) in
benzene was added to the water and the solution was stabilized by the sur-
factant, Twe«n-80, irradiation with the sunlamp then resulted in loss of,
TCDO.  A similar pattern of reactivity was reported in other studies.   For
example, a thin''dry film of TCDD on a glass plate was'quant It.it! vely re-
covered after exposure, to sunlight for 14 days (Crosby et_ aJL. 1971).  TCDO
on dry and wet soils showed negligible loss after irradiation by sunlaaps
for 96 hours (Crosby et al. 1971).  In contrast, TCDD in methanol solution

has a half-life 'of about three hours in sunlight (Plipiraer ££ al_. ,1973),
Thin films of TCDD on glass plates in Esteron'(a commercial brushkiller)
and in the herbicide Agent Orange showed 50% crana formations at about 4 and
5,5 hours, respectively (Crosby and Wong 1977),  Botre _e£ ail. (19?8)
described a method for detoxification of TCDD by aolubilizacion with sur-
factants and subsequent photolysis by uv light (254-356 ntu); Gebefugi e_e
a I. (1977) reported, that TCDD on silica gel exposed to a Pyrex-f iltered
mercury lamp (> 290 na) was 92% decomposed in seven days. , • Libert! e_t al.
(1978) have reported on the photochemical transformations of TCDD, and
again demonstrated the requirement for hydrogen-atom donors in the photoly-
,sis of TCDD.  These paper* do not provide any information relevant to
aquatic systems useful in an environmental assessment.

    34.4.2  Oxidation

         No inforaation has been obtained on the oxidation of TCDD.  The
electropositive nature of the taolecule as calculated by Miller ejt al.
(1977) suggests that TCDD will be more resistant to oxidation than~non-
chlorinated or  less cnlorinated aromatic compounds.

    34.4.3  Hydrolysis

         Ho information has been obtained on the hydrolysis of TCDD.  Hy-
drolysis of TCDD is not likely under environmental conditions, because the
halogen and ether groups on the aromatic ring are not susceptible to hy-
drolysis except under extreme conditions.

    34.4.4  Volatilisation

         Ho quantitative information la available on the volatilization of
TCOD from aquatic environments.  Several pap*rs mention volatilizations a
possible loss process, but no studies have been reported on the volatiliza-
tion process Itself.   .            .   •       ^    •.

         Ward and Matsumura (1977) reported a study designed to evaluate
the microbial degradation of TCDD,  In an aerobic degradation study using
eutrophic lake  sediments and water, they found that the recovery of TCDD
was directly related to water loss, with gas evolution from the system     '
(consisting of  flasks with cotton plugs) presumably responsible for volati-
lization of water and TCDD.  No useful data were reported for estimating
volatilization  from aquatic environments.

         In studies of the photolysis of solid TCDD on glass plates and
leaves, Crosby  ct ^al. (1971) concluded that volatilization is not an
important loss  process.  As pointed out in the paper of Mackay and Wolkoff
1973), however, the extreme insolubility of some organic* in'witar *ay
enhance the volatility of such chemicals from water bodies.

    34.4,5  Sqrgtion

         Data from microcosm experiments Indicate that TCDD is highly
sorbed to sediments and biota.  Isensee and Jones (1975) examined the "ate
of TCDD introduced into an aquatic microcosm on soil and found 35-^9'i of
the TCDO still remaining in the soil afcar about 30 days; nost of the fCDB
not in soil was foand in the aquatic organisms.  The two experiments in
wnich TCDD could be detected In the »*ter required 4 ind 15 days for equil-
ibration among soil, water, and organisms, and,the concentration in the
organisms was always greater by three orders of magnitude than that in the
water.  Ward and Matsumura ,(1977) also found that TCDD regained bound to
sediments, with usually more than 90% of the radiolabelled TCDO found in
the sediment.  These authors further argued that most of the TCDD in solu-
tion was bound to organic natter and particulates.  Matsumura and Bene/et
(1973) also reported that TCDD was sorbed and bloconcei'trated by aquatic
organisms (See Section 34,4.7).

    34.4.6  Bioaccumulation

         Isefsee and Jones (1975) reported that concentrations of TCbD in
organisms ii taicrocosms, in which TCDD was added in the form of contamin-
ated sediments, were within an order of magnitude of the concentrations in
the .sedtmeacs, although concentrations in the organises *?era 4-26,-CO times
greater than the concentrations in, the water.   Concentrations (dry weight
basis) in snails, mosquito fish, and daphnids were (2-2.6) x 10-* times
the concentration in water.  Concentrations in duckweed, algae, and catfish1
were (4-9) x 10^ times the concentrations in water,   Isensee and. Jones
(1975) noted that some a'quatic species in their microcosm accumulated TCDD
to levels which were 100 to 1000 the LDgi} of mammals and reported that
bit-accumulation occured in the presence of sediments containing 0,! ppb
TCDO.  Isensee (1978) confirmed conclusions previously reported based on
microcosm studies (tsensee and Jones 1975); more recent data from larger,
nodifled microcosms gave bloconcentration factors of 2 to 7000 in the
aquaelo f-rganisms used in previous studies.

         Matsumura and J?enez«t (1971), using shorter exposure times, f-xi:id
that concentrations (wet weight basis) in brine shrimp and fish (silver-
sides) were 1600 and 54 times, respectively, greater than the conceutrat ion
in water.  Mosquito larvae had concentrations- of TCDD that were 28'5'V92'V)
times greater than the concentrations in water in various treatments.

         Young e£ a_l_. (1976) examined a number of organisms subjected to
long exposure to soils heavily contaminated with TCDD at a lirxe terres-
trial test site (used for testing application of herbicides).   The author-;
stated that TCDD'may accumulate in the tissues of rodents, reptiles, birds,

fish, and insects, hue that Levels of TCDD in the tissues did  not  exceed
Che levels found in tfie environment,  the soil  (environmental}  levels were
10 to 1,500 ppcr,

   ..34.4.7'  3 lot rans fo r ma t i on and B i od eg r a da t ion

         Matsumura and Benezet (1973) reported  chat 5 of 100 microbiai
strains known to degrade chlorinated pesticides gave transformation  ot
TCDD, but no experimental details or discussion were provided.  Ward and
Matsuaiufa (1977) reported that in laboratory experiments as much as  62  of
the TCDD added to anaerobic sediraent-water samples from a eutrophic  lake
was metabolized and that most of the radioactive  label remaining i'n  the
aqueous phase after the first day was in the form of metabolites;  however,
they did not identify the metabolites.  No netabolites ware observed in
aerobic cultures and the major portion of the TCDO was in the  sediments.
These authors also report that the half-life of TCDD in lake sediment was
approximately 550-590 days, and  chat in lake water alone about  70'4 of the
TCDD remained after 589 days.  Isensee and Jones  (1975) failed  to  find
evidence of biotransformation of TCDD in aerobic microcosms, despite the
use of a vide range af TCDD, (.oncencric fins with appropriate replication.
Young e_t_ aJL, C1976) infer chit ~CDD may be degraded by soil microorganisms
on the basis of the 'disappearan ;S of TCDO from heavily contaminated  soils;
no degradation products were mentioned.

    14,4,8  Other Re actions

         Miller et_ aJL. (1977) evaluated the electron accepter  properties csf
chlorinated dibenzo-p-dloxins using, molecular orbital calculation  techni-
ques.  T^ie calculations indicate that the more extensive the chlorination
at the 3 positions (see I, below) the stronger  the charge transfer complex,
with TCDD being the most susceptible to such complexes.  They  suggest that
strong charge transfer coraplsxation of TCDD may account for a  number of en-
vironmental fates ar;d effects of TCDD, including binding to soils, uutagen-
iclty, ;and' slow metabolic decomposition.  The authors point out that exper-
imental verification is- needed regarding the charge transfer cnmplexation
of TCDD.

    3*.4.9  Microcosm Studieg,Field Studies., and Modelling

         Microcosm tests, particularly the extensive experiments of Isensee
and Jones  (1975), demonstrate a marked tendency for 71)0 to accumulate in
the sediments (85-99% of the TCDO In the microcosm .was found In the sedi-
ments).   They also found that concentrations in organises were generally
wichin an order of magnitude of "he concentrations in che sediments.
Similar  results were noted by Ward and Matsumura (1977) and Matsumura and
Senezet  (1973).   Ward and Matsumura reported blocransformat ton of the frac-
tion that  'remained in the aqueous1 phase, but Isensee and Jones, in a care-
ful analysis of their own experiments, failed to find ^ny evidence of bio-

34.5  Daca Summary

    Table  34-1 sumnarlzes the data on the aquatic fate of TCOD.

                                                                      Table  14-1

                                                            Suwnary of Aquatic  Fate of TCDD
                                                      Siii» 11 y
                                                                            Llfe t>»
                                                                        Cuti I: idence
                                                                          of  t>al»



VolatUl/ul ion

Surpt io«

BioaccuBtuIcit ion

Biot rau£»tt-n»4t ion/
Will be an  laipoilaui  procet»i>
If [eactive  !.ubstran'a arc

Not an important  piocebi*.

Does not occur.

Probably not  an  i»poiManl.

laporlant pruceb^.

Probably an  important process.

Could b» an  important process
over iong tine periods.
                                                                                                       1  year




Modi tun


                       a.  There  Is  insufficient  inforaatioit in tlie  reviewed literature to per*it assessment  oi a nost proh.ible  iate.

•^• 6  Literature '-3_L£5

Sotre C. ,  A.  Me:noH, dnd  r.  Aihalj-ie.   !•»;>.   ''f '.)J -s.j luo T. \.?. i :.i.<. :\ -\.:.: ;:no-
  todecoraposi t ion in .iqueous soi ,:i i-in.-.'.   Znvj.ro'i,  ~,c i.' T,j.:rmoi .

Crosby,  D.G., K.W. Moil men, -KH!  -.S.  «\., .;-. .    ! > / i,   in-L i .ui-n.-u. .!  :;»'«>-:•«-
  tton  and degradation ot  •iti.-..::i2;/d ioxi.i-i ,ir:c 'd i.b(±nz->t ur,in=.   EL^iroii.
  Health Perspecc.   5:259-265.

Qrosby,  D.G. , and'A. i. We tig.  L97V.   ~nv ir.jrun^iKal  uegrada r'ion  of
  2,3, 7,8-teCrachlorncib«_nzo- ;> -dioxtn  (!(".[»'.>).   Si.::.e..ce' 195 (iJ3-*) : L 3 >7 - i, >3
                                                                             > > '^ •
Crosby,  B.C., A,S, Wong,  J.R,  pilsner,  and •?. , -V.  Wo^ison,  19?i,   Pho-
  todeconposicion of chloriadteU dibeiiia-p-diov;ins.   Science
  171(3998) ;748-7-t9.

Crummett,  W.B., and R.H,  Stetii.  ^li>.   Je term mac ian pr chi;/r i .lated
  dibenzo-p-dioxins and  diben^ofurans  in. ''jrious aaceriais.   Environ.
  Health Perspect. 5:15-25.
Gebefugi,  I., f,f B<*'i™dn"i ,  and f. iCorr.e.   H~7 .   l-'hoiuchtiai sc.'ner  jbbau ;.ofi
  2, 3,7 jS-tacrachlorodibenzo-y-^ d i-">Kir.  (7CD-D}  writer  iiaulierten,  nun
  weltbedingungen.   NaCur:«'i5^cn!>o.riaf ten  6-i(9'K-«*6.

Helling,  C.S., A.R,  loer.see, £,A.  Wooi. son,  P.O./.  fin so r, C.E. .jor..-:5,  * . ^ .
  PH miner,  ana P.O.  Kearaey.  i9<7j, '  Chiurod i ox 'ins  in pesticides, >,.:!/,,
  and  plants,  J. Environ.  Quality  2(2) : 171- i.7ti.

Isensee,  A. 8.  1978.    Sioacc'irnuiution  >;f 1 , 3 , 7 ,3-t»? crdrhl^rod ibs;i^ i- p ;; a •
  dloxln.   Conference  OP.  chl .>ri:;ated  piien^K;.'  acids,  .ind c;>.eir  dioAia.s.
  Ranel,  ed.   Ecol.  Bull.  (Stokholm)  (27 ) : 2^)- 2o2.
          *     i                           i

Isensee ,  A.R., and G,5.  Jones.   19/5.   Disf r ibat iu,:  ,.,t 2 , 3 , 7 , rf- ce ' r.i.' cii-->'
  dib«nzo-p— dioxin (TCDD)  in aquatic  inodei  ecosys.'., ';m.    Environ1. Sol.
  Technol.   9(7 ): -508-172-

Kearney;  P.C., A.R,  Isjrisee, C.S.  Helling,  ...A.  Vu.-j!»on, ind  !.H. PI ;.;;;. .-.••• r .
.  1973.   Envir'unmentaL  signir icance  of  ,:hlorpiii.ixir>si.  Adv. Chem. St.r.
  120:105-111.                  :         '    «

Kearney,  P.C., E.A.  Wool son, and -.!,.?.  EiLir\^t3.,  Tr.  H72. '  Pec sis to ;;<-••••
  and  metabolism of  chlorodiuxiritj  in  soils.   Eru-iror.. 3c i .'  Techno L.

Ltberti,  A.,  D. Btocco,  £.  M U-^riai ,  \.  Ce.ir.it!},  *nd M. Pi>ssan/i.Ti .
  1978.   Sola* and UV  photo decomposition !;>f  2 , ') , 7 , 8 -le tr actil -iroo ibt-.-iz .•
  para-dioxin in the anvirorunent .   Soieru-e  or  the  Tot,-il Kiwiron.
  (10): 97-10*.

Mackay, D., and A.W, Wolkoff.  1973.  Rate of evaporation  of  low  solubility-
  con tarainants from water bodies to atmosphere.  Environ.  Sci. Technol.

Matsumura,, F», and H.J, Benezet.  1973.  Studies on  the  bioaecunulation  and
  mlcroblal degradation of 2,3,7,8-tetrachlorodfbenzo-p-dioxin.   Environ.
  Health Perspect,  .5:253-258. '

Hiller, G., S. Sontum, aad D.G. Crosby.  1977.  Electron-acceptor
  properties of chlorinated dibenzo-p-dioxins.  Bull. Environ, Contain. Tox-
  icol. 18(5):611-616.

Plimmer, J.R., and U.I. Klingbiel.  1971.  Ribollavln photosensitized
  oxidation of 2,4>-dicnlorophenol:  Assessment of possible chlorinated
  diaxin fortaation.  Science 174:407-408.

Pliraraer, J.R., O.I, Klingbiel, D.G. Crosby, and A.S. Wong.  1973.  Pho-
  tochemistry of dlbenzo-p-dioxins.  Adv. Cheia. Ser. 120;44-54.

Ward, C., and F. Hatsunwra.  1977.  Fate of 2,4,5-T  contaminant,
  2,3,7,8-tetrachlorodibenzo-p-dioxin (TCDD) tn aquatic  environments.
  U.S. XTIS, Pi Rep., PB^264187, 22 pp.

Young, A.L., C.E. Thalken, E.L. Arnold, J.M. Cupelio, and L.G. Cockerham.
  1976.  Face of 2,3,7,8-tetrachlorodlbenzo-p-dioxin (TCOO) in the en-
  vironment:  summary and decontamination recommendations.• Summary report,
  department of Chemistry and Biological Sciences, United States  Air Force
  Academy, Colorado.  U.S. STIS, Rep. AD-A033 491, 41 pp.

                              35.  TOXAPHEME

35.1  Scatemetit ofProbab I
    An inclusive assessment of the fate of the pesticide coxaphene in
aquatic environments is complicated because toxaphene is a complex mixture
of polychlorinated camphene derivatives of different physical properties
and environmental behavior.  Toxaphene is very stable to biological and
chemical processes Ih aerobic environmental systeias, but it does undergo
partial reduction (loss of chloride content) if anaerobic environments,  A
dominant process' in aquatic environments is direct sorption on sediments or
sorption onto particulates followed by deposition into sediment where bio-
logical and possibly chemical reduction occurs.  The rate of loss of toxa-
phene froffi aquatic systems will then be partially determined by particu-r
late loading and quality of the water body; shallow, particulate-laden,
eufrrophic waters give maximum transformation rates of toxaphene, with half-
lives on the order of a few months for some components.  The physical pro-
perties and .chlorinated functionality of the individual tpxaphene struc-
tures will govern which -components will be sorbed and then subsequently
reduced.  The fi.nding of some toxaphene components in aquatic sediments and
species after several years indicates that bioaccunulation in the food
chain may occur.  Unless clear evidence proves otherwise, the absence of
acute toxicity effects of toxaphene should not be interpreted as indicating
that all toxaphene has been degraded and chronic toxic effects are absent.

35.2  Identification

    This chapter considers toxaphene as the chlorinated camphene mixture
containing 67-69% chlorine; Holmstead e_t_ al_. (1974) showed that at least,
177 compounds are present in toxathene.  Saleh and Casida (1977) reported
that about 85% of the glpe peak araa (electron capture detector) is ac-
counted f>r by 29 major peaks that individually vary from 1 to 8% of the
total area,

                       in I  »  TOXAPHENE (AVEflAOf FORMULA C1OH1OCI-)


    Holmstead £t, aJU (1974) showed that at least 177 compounds are pre-
sent in the toxaphene mixture; about two thirds of these compounds are of
C10H11C17' C10H10C18» and C10H9C19 formulae.  The remaining,chemicals are

 the  C10H10C16,  C10HpCl6,  C1QH9Clr  C^d,,  C^Cl^,.  and C^H.C^
 chlorinated  derivatives  of^camphene.   Since both ionic and free radical
 reactions  probably'occur during  the  chlorination of camphene,  such a
 complex reaction mixture is  reasonable.   Many  studies on  the identification
 of various components  of toxaphene and their associated toxicities have been
 reported (Chandurkar £t  al_,  1978;  Turner e_t al.  1977; Saleh ejt ad. 19,77;
 Palmer  et  al.  1975;  Turner et.  al.  1975;  Khalifa  «t_ aJL. 1974).   Most com-
 pounds  identified thus far are of  the polychlorinated bornane  struccura;
 Saleh _et ad.  (1977)  stated that  the  octachlorobornane toxaphene co.mpoaer,ts
 A—1  and A-2  (see below)  are  major  contributors to the acute toxicity of

    Although  it would  be  preferable  to focus on the toxic chemicals of the
 mixture,  lack of  environmental  information on these components makes such
 an  approach  futile.  Focus on 'components may also disregard environmental
 studies  that  have been c?nducted using the toxaphene mixture;  this chanter
 identifies ' components  using teras as described In literature (i.e., "higher
 molecular weight  fraction," "less soluble," "higher chlorinated compounds,"
 etc. ).

    leportant ??ote:  As a mixture of variously chlorinated structures, ail
 toxaphene components do not have-similar properties that  provide parallel
 or  even  similar fat#«  (and toxicity}.   Most data on sorption,  solubility,
'bioaccuraulat ion,  and biotransf ormtion were reported as  t_o_c_a_l  _tc
                                                         ____ _ci
 because  of  the  analytical  difficulty  of  separating the individual compon-
.ents.  However,  most  authors  also report  that  marked qualitative  differ-
 ences  occur in  the  glpc  profiles  of  "toxaphene"  after sorption,  oiotrins-
 formation,  reduction,  or bioac cumulation.   Thus, the total  concentration
 toxaphene should be difficult to  correlate  In  an assessment  or  predi,- 1 i .-ir.
 of  a specific environmental or  ecological effect when the  effect  is  pussi
 bly due  to  the  activity  of a  small component  of  the total  toxaphene  me t-
 sured.   Accordingly,  the argument  that a  rapid "detoxification"  of  toxa-
 phene  Minimizes  pollution  hazards  atay be- a  fallacy, since  the abserrce  ->f -
 acute  toxicity  effect  does not  necessarily  correlate with  the absence  of
 chronic  toxic effect..

    Alternate names and CAS and TSL  numbers  for toxaphene are given below.
    (Mo structure is unique  for  toxaphene;
    structures for a few toxaphene  com-
    ponents are given below. )

    CAS No-. 8001-^5-2
    TSL Mo. XW 52500
Alternate NTames

Hercules 3956
Toxakil      , ,
                           CH2C1    CHC12    A-l    Turner et_ al.. 1975
                           CHC12    CH->C1    A-2    Turner et_ al. 1975
                           C"?C1    CH^Cl     3     Oialtfa rt al. 1974
                                   CHC12    Ac     Chandurkar ejt_ a_l. 1978
*The names below are as presented  In  che  respective literature citations,

35.3  Physical _Pr_Q|ierjJ.es

    The general physical properties of  the  toxaphene  mixture  are:
         Molecular weight
         He It ing point
         (Brooks 1974)

         Boiling point
         (Brooks 1974}

         Vapor pressure at 25*C
         (Brooks 1974)

         Solubility in water at 25°C
         (Brooks  1974)
        , (Weil et al. 1974) ac 25*C
         (Paris «t at_. 1977)

         Log octanol/water  •
         partition coefficient
         (Paris et_ al, 1977)

35.4  Summary__of Fate J)a_ta

    35.4.1  Photolysis
343 for C10H10C16
51? for
range 70-95SC

decomp.  > 120°

0. 2 torr to 0.4 torr
about 3 ppw
0.740 pptn
C.500 ppm

         Wolfe et aJL. (1976) reported photolysis studies on several  pesti-
cides and found that the gipc profile of toxaphene was unchanged  on  expo-
sure to uv light filtered through borosllicte glass  (> 290 ntn).   They  also
reported that the relative photoreactivities of toxaphene, walathion,  and
2,4-D-BEE (2,4 dlchlorophenoxyacetlc acid, butoxyethyl eater) were < I,  i,
and 300, 'respectively.   The sunlight photolysis half-life of 2,4-D-BEE was
determined to be 12-14 days.  Assuming that the relative reactivity  of
2,4-D-BEE and toxaphene is at least 300:1 in sunlight, a half-life of  over
10 years is probable for toxaphene photolysis.  Since 2,4~D~BEK surely a.b-
sorbs more light in the solar spectral region than toxaphene, the  10 ve.ir ,
half-life estimate is probably too short b-it it does serve to show that   (  •
photolysis is not an important process- for toxaphene in aqueous systems.
No reliable, environmentally relevant data 01 photolysis studies were  found

in literature to substantiate che often-found statement that sunlight  is
important in detoxification of toxaphene.

    35,4.2  Oxidation

         N" inforr.at ion is -available on the oxidation of toxaphene in  aqua-
tic systens.

    35.4.3  Hydrolysis

         Wolfe ££ aL. (197-3) reported tnat toxaphene was unchanged after  2
days at 65*C in aqueous solutions at ?H values of 3." and 10.0,  Using this
information, and allowing for 'a two-fold decrease in hydrolysis rate far
every 10°C interval with' extrapolation to more soderace pH values, the hy-
drolysis half-life of toxaphene at environmental pH values (5 to 8) and
25°C is greater than 10 years.

    35.4.*  Volatilization

         Mo data were found on the toxsphene volatilization from aquatic
systems.  Calculation of volar i^zaeiort half-lives using che procedure of
Mackay and Leinonen (1975) is difficult to interpret because of the many
polychlorinated structures constituting, tcxaphene.  However, it is possible
that the very low solubility of the higher chlorinated bornar.e structures,
which are highly toxic, say contribute to volatilization of the components
from aquatic systems.  This transport process in addition to the sorption
on part iculate— deposition transport process nay explain the moderate de-
Eoxifl_3tion rates observed -in some shallow laK.es.

    35.4.5  Sorpt ion      !

         Paris et_ al_. (1977) reported the adsorrtion and equilibrium of
toxaphene with bacteria, fungi, and ?lgae. . Distribution coefficients
^d * C^/C^, where C is the concentration of tcxaphene (xg/ag) in
microorganism (C3) and water '(Cw) were as follows:

        Mj.cr_oprga'ni_sa_                       Ed_Z_12."3_
         2 Bacteria                         1.4 * .).?, 5.2 + 0.:

         1 (fungus               ,            17 * 2

         1 Alga                             17 + 1

         1 Field1 sample (algae              
The  time  for  sorption equilibrium was 10 minutes for a'lgae, 30 minutes for
bacteria,  and 2  hours for  fungi;  desorption and equilibrium were attained
in similar times .and  the saae Kj  values were obtained.   Equilibrium ii
the  field 'sanpie was  reached  within  1 hour.

     These authors also found  that the less soluble toxaphene components
(with  longer  retention tine and'higher molecular weights) were preferen-
tially sorbed by the  tiicroorganisnis.   The importance of sorption of toxa-
phene  onto plankton and subsequent deposition into pediment has also been
described by  Veith and Lee (1971).
     35.4.6 Bipaccuau lac ton

          Sorption of  toxaphene on biota is raff id (Parts et_ a_l_. 19??) and
significant uptake occur,-?  in  natural  waters (cf, reviews by Hughes 1970;
and  Ho!t  1977).   The  ratios of the concentration of toxaphene in organisms
to the coxaphene concentration in water obtained in laboratory studies Ky
Sanborn et al.  (1976) and  Schimmel _e£ aJL. (1977) are representative .
Sanborn «£ aJL.  (1977) reported ratios of 6900 for an alga, 9600 #or a
snail, 1900 for  mosquito  larvae,  and  4200 for fish in microcosms.  Schiamel
ec_ a_l.  Deported  ratios of  3100-21,000 for fish and oysters in 96-hr tests
and  -iOO to 1200  for shriap.   Ratios  of'4200 to 60^,000 were obtained (whole-
body  basis)  for fish in 2SO  day  tests.  Terriere e£ al. (1966) also  re-
ported significant uptake  and concentration of toxaphene in aquatic plants
and  invertebrates and fishes  over several years in two lakes, with concen-
tration ratios of >6o for  aquatic planes, (1-2) x 10 ^ for aquatic animals
other  than fish, and  (1-2) x  10"*  for  rainbow trout in a shallow lake rich
In bioloziw.il life.   Terriere, e_t_ a_l. also found that the toxaphene 'cow-
ponent profiles  were  different among  .the commerical toxaphene mixture and
toxaphene residues recovered  frosn aq>tatic plants or trout.
     3'5 . 4 . 7   B t oe r a n ^f o rraa t ioR=^and_ Biodegradatioti

        •  Parr and Smith  (1976) reported' a 50*5 loss of toxaphene in 6 weeks
 due  to biological transformation of  toxaphene in anae robic ,  unstirred,
 flooded soils but found  no  transformation in aerobic sediments.  The resis-
 tance  of  toxaphene  to  blot rans format ions  in aerobic systems  was also found
•in microcosm experiment  reported by  Sanborn et_ al. (1976) and in the single
 culture experiments  of Paris  et_ aJL.  (1977).

          The susceptibility of toxaphene  to reduction (i.e., loss of ch.l-T-
 ride from structure  by hydrog*" substitut ton or elimination) by biochemical
 ind  cheraicil sat hods has  been clearly demons trited.  Casii^  et  'a_l. (^"y")
 found  that  toxaphene was  partially netabolizel in rats and concluded that
 on- the average  about half  of  the C-C1 bonds in toxaphene were met.-ih.j-
 lically  labile;  they specificallv pointed out, however,  that the

components in toxaphene will show different reactivities and extent of re-
duction because of different degrees of chlorination and structures of Che
           themselves.                     '                  i  •
          Khalifa e_t al, (1976) reported that hematin (ferriprocoporphyrin
hydroxide) reduced toxaphene ,  as did a r»t liver microsome-NADPU system.
The reduced roxaphane species  were determined by the reduced gLpc retention
times and reduced sensitivity In detection by the electron capture detec-
tor.  Toxaphene components A (A-l and A-2) and S were reduced by these
systems through dechlor Lnation and dehydrochlor 'nation mechanisms.

         Subsequent work by Saleh and Casida (1978)  reported chat toxa-
phene component B (a heptachlorobornane; see Section 35*2) was reduced in
bovine rumen fluid, in sewage  primary effluent and in vivo in houseflies as
well as in rats.

         The results of these  studies, and those of chemical reduction
studies (see S«ction 35.4.8) and field studies in lakes (see Section
15,-4.9), indicate that toxaphene will be reduced in eutrophic , anaerobic
environments, but chat different toxaphene components and even different
chlorinated sites within a component's molecular structure will oe reduced
at different rates (see Section 3S..2).  It is also significant to note that
the moderate "detoxif ication"  rates of sow« toxic, higher chlorinated toxa-
phene components in eutrophic, shallow lakes are probably partially due to
removal oy the preferential sorsion of these lower solubility components to
particulate and sediment, with subsequent reduction in taese anaerobic

    3i'4-8  Other Reactions

         Williams and Bidleman (1973) reported that toxaphene was reduced
in anaerobic unsterile and sterile wet estuarine sediments and in sterile
3and containing a Fe(II)/Fe(III) couple.  Although they could not qua-sti-
tate the transformation rates, "he authors found marked changes In the
t )xaph«ne gipc profile after' only several days; the results of a. sterile
sand-distilled water control experiment with the Fe(II)/Fe(iI I)' couple
showed that chemical processes were occurring in sterile systems.  It was
not d«>ternuned whether biological transformations were also occurring i'n
the unsterile sediments.

        -Khalifa e_£ ad. (1976) reported reduction of toxaphene as well as
its itdividual toxicants A and B by reduced heraatin (see section 3?.-«.6).
SaLeh and Casida (1978) reported t'.iat toxicant i is also reduced by free
radical-triphenyltin hydride reactions in hexane and by photolysis in
hexane solution with uv light  > 220 nm.

                                 ~ •} ?
    33.4.9'  Microcosm Studiesf ?ie\A Studi es, ajd_ Modelling

         •Sever.il groups reported studies of  toxaphere persistence  i'n  L;
Data from these studies suggest that toxaphene :nay persts:  in  Likes  roi
periods of several months to more than 6 years as measured  by  j^ut-?
ty to fish.   This toxicity criteria obviously measures the  persi-
the acre acutely toxic toxaphene components  and does not  include
components that are possibly more persistent and may accumulate  in the  r ..
chain.  Toxaphene persistence data Indicates that transformation*  (either
chemical or  biological) will occur fastest where the rate of transfer  t >
sediment/anaerobic systems- is most rapid and where biologically  rich  sys-
tem are also present (i.e., shallow, biologically-rich  lakes  -aiv« r\ist»r
toxaphene reduction/detoxification than deep, oligotrophic  lak^s).

         Stringer and McMynn (I960) reviewed the detoxification  of'toxa-
phene in lakes as determined by survival of  fish in cages.  Oetoxificatio-
dines in 14  lakes ranged from 11 .to 43 months with an average  of J9  -nor,eh<
The author*  state that high turbidity and shallow lakes  increase 'detoxifi-
cation of toxaphene.

         Terriere et_ a_l_. (1966) reported the persistence  of toxaphene  ii
two mountain lakes,  tn a shallow lake rich  in aquutic life (Davis Lake),
Initially treated at 88 ppb toxaphene, the concentration  of toxaphene  in
water after  1, 2, and 3 years was 0.63, 0.41, and < 0.2  ppb, respectivulv
feach concentration is an average of 6 samples).  In a deep, biological lv
sparse lake  (Miller Lake) initially treated  at 40 pnb toxaphene, concentr
tions in water of 2.10, 1.20, and 0.84 ppm were found after i, 3,  and  A
years.  Data for uptake and bioaccumulation  in aquatic platts  and  invert•>•
brates fish and bottom mud in the two lakes  showed nlgher concentrat i -irss  •
Miller Lake'species than in Davis Lake species, with average concent rat I :>•
in trout exceeding 2 p,pm in all analyses (also see Section  35.4.'").   Th<»
authors also stated that trout could not be  restocked in Miller  Lake  for
six years because of the toxaphene levels that persisted.

    :     L»e and coworkers reported a series of studies  on  the persistent
of toxaphene in lakes (Johnson et_ al. J966;  Hughes et_ ajl. 1970;  Hughes
1970; Veith and Lee 1971; Hughes and Lee 1973; Lee at_ a_K 1977).   The  !9n<
paper reported that lakes treated with 0.1 ppra f xaphene  3  co  9  year^  be-
fore analysis contained toxaphene as 1-4 ppb in water, 0.2  to  1  pon  in
sediments and 0.05 to 0,4 ppm in aquatic plants.  The paper iis5 stit.»-j
that sorption1 onto particulats is one detoxification mechanic,  and  ^\i^
gested that  the components of the toxaphene  tiixture are  degraded -it  J;r-e-
ent rates and have different toxicittes.  The 1971 paper  reported  tn.n  i-,
toxaphene-treated lake the concentration of  toxaphene in the iej;-en'-  ; n-
creased for 190 days and then decreased' by a factor of 2 for each  suh--e-
quent 120, day period.  Citing work in their  group, the authors f-xnd  a

significant mechanism for accumulation In sediment is sorption of toxaphene
on particulars matter with subsequent deposition in sediment.  They also re-
port that toxaphene accumulated on sediments under the natural lake con-
ditions could not be desorbed from sediment by pure water in the labora-
tory, suggesting that deserptlon from sediment, is not a significant process
in the aquatic*environment..,la the 1977 paper, Lee ani coworkers found
that toxaphene extracted from sediment «as less toxic to fish than commer-
cial •coxaphere preparations,  demonstrating that some toxic components of
toxaphene were transformed in sediments,

         Sanborn et_ a_l.  (1976) explored the fate of toxaphene in small mi-
crocosms, but did not identify products, although they did quantify the
concentrations of 6 unknown components in the water and the organisms.  At
the end of one month, 60S to 3tt2 of the radiolabelled material in the or-
ganisms remained'in the form of toxaphene.

35.5  Da_ta Suaaa_ry

    Table 35-1 summarizes the data on the aquatic fate of toxaphene.

     Table V>-1

of Aquatic Fat*  of  Toxaphcnc

                                                                                                                              of  Data

Ox Illation




           t ion
Biot (ansfoittat ion/
  Blouegrdddt ion
Not au important  proton.

Frobabiy not an  important
Mot an iaportant  protean

May bte an iBportaut  pioc

Itt an inportiint procebs.

Is an lapoitanl procvi>».

Dtf^catied in auueiubic  uy
but not in aerobic  syhti'
                                                                                                     > 10 years
                                                        >  10 years
                                                                                                     < 2 hours foe uptake
                                                                                                       by •Icro

                                                                                                     Uiil depend on trultia-    Low
                                                                                                     port to auaerubtc
a.  There  Is  insufficient  Information in tht: reviewed  literature to permit dtauessiMfnt  oi  u ML>bt  prubdblv f
Brooks,  G.  T.   19'7i.   Chlorinated  insert i'ii les :   Vohi.ne  I : Technology and
  applications.  CSC  Press, Clew Kind ,  <">H, ,  2-i1?  pp.

C:»sida ,' J.  ". ,  R. .!<.  Holrr»ste«'< ,  :',  Khitlf.i,  r.  R.  Knox, and  T.  Ohsawa.
  19?5»   Toxaphene  composition and  -ne^ aho !i sri  In rats,  Snvtron, 0: 5fc i- '5*9.

 "handnrk.i r ,  ?.  ^ . ,  F.  Matsuiivur.'i ,  :iitd T,  Ikeda,   1973,   Identification and
  tuxicicv  ot  c '>x i 'an*:  Ac, '-s  t-)xi>,:  /,:oniron>:"it  ^f  roy".iph*>ne .   Che^nosphere
  :<2 "el-'1 3- IV.

Holmstead,  R.  L,  , S. 'Kh.ilifa, «ind  ,1.  >-.  Casid^.   1974.  Toxaphene
  composition analyzed  by combined  5,13  chroma tography-ciietical  ionlzaCion
  mass  spectromi»try.   1.  Agric,  Food Cheni.  ,22(^1:939-944,
Holt,  3,  R,   1977.  Population exposure  to  toxaphene (chlorinated
            ).   ^r-Iit.  CRFSS  Report  \o    3'j,  Prepared' for  F.nvi ronmenca 1
                   y,  •£? *-\ft^ -;"- 1--» 3 1 i ,  69pp.
 u^he« ,  ?.  A,  ,  :97ri,   Stui-'ic-s -i •.  ''-.o  per^i^tente  of toxapheic  in treated
   lakes.   Ph.D.  Tnesis.   Vr'* '-:  ',. .-xaphtf.?  • n "••irnr-iL  w.Tt-srs, 't;sh,  -ind !:iborn,j-,tf  jr-.d  ocr 'u-ii1: or j r co^icanc frcwi,
  t^ -:r;nt -:,-il  r -x.ip'vne .   J, /j.ii^LC,  Fcod 'him. 22 '- ) ;f> 5 J-r 5 7;,

Kh-siif-i,  c, , ?,  '.,  Holmste^id,  -ind  J.  E. Casida.   39" -.i,   Toxaphene
  degr.'idat L --, kv  tron(ir) -;, r-c ;yi:r;>byr in  >ys: •TS.   j. Axric.  Food CheTi.
                                     33- ! :

Lee, G. F., R. A. Hughes, and G. D, Vetth.   1977.  Evidence  for partial
  degradation of toxaphene in the aqueous environment.- Water, Air,  Soil
  Pollut. 8:478-484.

Mackay, D, , and P. J, Leinonen.  1975,  Rate  of evaporation  of
  low-solubility contaminants from water bodies to .ataosphere.  Environ,
  Sci. Technoi.  9(13):1173-1180.

Paltaer, K. , R. Wong, R. Lundin, S, Khalifa,  and J. Casida.   1975.  Crystal
  and molecular structure of 2,2,5-endo,6,-exo,8,9,10-heptachlorobornane ,
  CigHiiClj, a toxic component  of toxaphene  insecticide.   J. Am.
  Cheta, Soc.  97(2 );4G8-413.
Paris, 0. P., D. L. Lewis, and  J. T. Barnett.  1977.   Biocpncentration ,of
  toxaphene by microorganisms.  Bull. Environ. Contain. Toxicol.

Parr, J. F. , and S. Smith. ' 1976.  Degradation of coxaphene  in select eJ
  anaerobic soil environments.  Soil Sci.  l21(l'):5,2-57.

Saleh, M. A., and J. 5. Casida.  1977.  Consistency of Toxaphene
  Composition analyzed by open  tubular column gas-liquid chrotaatography.
  J. Agric. Food Chera.  25(l):63-68.

Saleh, M. A., W. V. Turner, and ^. E. Casida. '1977.   Polychlorobcrnane
  components of toxaphene:  Structure-toxicity relations and metabolic
  reductive dechlorination.  Science 193:1256-1258.

Saleh, M. A,, and J. E. Casida,  1978.  Reductive chlorination of  ttv;
  toxaphene component 2,2,5-endot6-eKo,9,9,10-heptachlorobornane in  Carious
  chemical, photochemical, and  metabolic systems.  J.  Agrlc. Food  Chem.

Sanborn, J. R. . R. L. Metcalf, W. N. Bruce, and P.-Y.  tu.  1976. ,  Ttie fate
  of chlordane and toxaphene in a terrestrial-aquatic  model ecosystem.
  Environ. Entomol.  5(3 ):533-538.

Sanborn, J. R., 8. M. Francis,  and R, L. Metcalf.  1977.  The. degradation
  of selected pesticides in soil:  A review  of the published literature.
  U.S. OTIS, PB Rep. 'P3-2723S3, 633pp.

Schimmel, S, C., J. M. Patrick, Jr., and J,  Forester.  1977,  Uptake and
  toxicity of toxaphene in several estuarine organisms.  Arch. Environ.
  Contara. Tox,  5:353-367.

Stringer, G, E., and R.. G, McMynn.   i960.   Three  years'  use  of  toxaphene .is
  a fish toxicant in British Columbia.  Can.  Fish Cult  28:37-44.

Terriere, L. C,, U. Kllgenagi, A. R. Gerlach,  and R.  L.  Borovtcka.   1966,
  The persistence of toxaphene in lake water  and  its  uptake  by  aquatic
  plants and animals.  J. Agric. Food Chem.   l;»:*->6-69.

Turner, W. V., J. L. Engel, and J. E. Casida.   1977.   Toxaphene components
  and related compounds:  Preparation and  toxieity of sorae  hepta-,  o<_ta-,
  and nanachlorobornanes, hexa- and  heptachloronenes, and a
  hexachlorobornadlene.  J. Agric. Food Chem.   25(6 ):139^-1 'Ol.

Turner, W. V., S. Khalifa, and J. E. Casida.   1975.   Toxaphene  toxicant A,
  Mixture of 2,2,5-endo,6-exo,8,8,9,10-octachlorobornane and
  2,2,5-gndo,6-exo,8,9,9,10-octachlorobornane.  J,  Agric.  Food  Chera.
  23 (5): 991-994.

Vieth, G. D., and G. F. Lee.  1971.  Water  chenistry  of  toxaphene-role of
  lake sediments.  Environ. Sci. Technol.'  '5(3 ):230-234.

Weil, L., G. tkire and K. E. Quentin.  1974,   Solubility  in water of
  Insecticide chlorinated hydrocarbons and  polychlorinated  biphenyls  in
  view of water pollution.  Z. Wasser Abwasser  Forsch.   7(6}:169-175.

Williams, R. R. , and T. F. Bidleman.  1973.   Toxaphene  degradation  in
 • estuarine sediments,  J. \gric. Food Chem.   26(1):280-282.

Wolfe, N. L., R. C. Zepp, G. L. Baughman,  R.  C. Fincher,  and J.  A.  Gordon.
  1976.  'Chemical and photochemical  trans fbrmation of sel-acted  pesticides
  In aquatic systems.  0. S. Environ. Prot. Agency, Off.  Res.  Dev. ,
  EPA 600/3-76-067, 151pp.


            Chapters 36 & 37

                      36.  POLYCHLORIN4TEO BIPHENYLS
36»1  Statement of Probable Face                        '

    Polychloritiated biphenyls (PCBs) are a family of compounds which vary
widely in physical, chemical, and biological properties.  For those com-
pounds with four or fewer chlorine, acorns per molecule, biodegrada'ion ap-
pears to be the dominant fate process and results in significant destruc-  .
tion and transfomat ion,  Polychlorinated biphenyls with five or more 'chlo-
rine atoms per molecule have been photolyzed in experimental situations but
it is difficult.to extrapolate these results to environmental conditions.

    Non-destructive processes which affect-the distribution and transport
of polychlorinated biphenyls  re adsorption, volatilization, mnd btoaccusnu—
lation.  In natural water systems, the greatest concentration o!f these com-
pounds is sorbed to suspended1 and bed sediments due to the very low solu-
bility in water.  The tendency of polychlorinated biphenyls for adsorption
increases with the degree of chlorination and with che organic' content of
the absorbent.  The biota are another environmental compartment into which
these compounds are strongly partitioned (measured bioconcentration factors
range up to 10^),                                            '  •

    Volatilization and transport as an aerosol followed by fallout with
dust or rain is the probable cause of the ubiquitous distribution of poly-
chlorinated biphenyls. The more highly chlorinated species are less vola-
tile than the lighter species.  The presence of suspended solids tends to
reduce volatilization,• presumably because the solids adsorb the polychlo-
rinated biphenyls and reduce the concentration in solution.

    The available empirical evidence indicates that polychlorinated bi-
phenyls, especially those with four or more chlorines, are persistent in
the environment.  The composition of polychlorinated biphenyls in (the at-
tnosphere is similar to that of Aroclor 1242 or 1016, while those in sur-
face waters (.mostly adsorbed to suspended solids) approach the composition
of Aroclor 1254.  Polychlorinated biphenyls in biota are heavier and more
chlorinated still, and approximate the composition of Aroclar 1260. Thus
the processes controlling distribution are somewhat selective,'with the
lighter species more likely to volatilize and the heavier species aore
likely to be incorporated into sediments and biota.

36.2  Ident^if igatton

    The Aroclors are technical mixtures of a number of the individual poly-
chlorinated biphenyls nade by the partial chlorination of biphenjl in the

 presence  of  a -suitable  catalyst.   In the designation of  the individual
 Aroclors  (Monsanto  TM)  a  set  of  four digits  was used,  the  first  two,  12,  to
 designate chat  the  preparation, is  a  mixture.   The second set of  two numbers
 is  used to denote the approximate  chlorine content by  weight.  Thus,  Aro-
 clor 1242 is a  mixture  having an average chlorine content  of 42  percent.
 When it was  determined  that there  was a  significant environmental  problem
 associated with the aora  heavily chlorinated  species,  Monsanto (the major
 American  manufacturer)  prepared  a  new mixture that was limited primarily  to
 the mono, di and trichloro  isotners.   This product carried  the  designation
 Aroclor 1016.

     Of the total of 209 possible compounds resulting from  the  partial or
 total chlorination  of biphenyl,  some 100 individual -compounds  have been de-
 tected in the various Aroclors (Hutzinger .ec  a_l.  1974).   The structure of a
 typical (2,2'-dichlorobiphenyl) member of the class of polychlorinated bi-
 phenyls is shown below.
                               3'         65

     CAS  SO,:   Each  of  the  polychlorinated  biphenyls  compounds  has been
               given an individual  CAS  number.   The CAS  numbers of the
               Aroclors are listed  in Table 36-2.

     TSL  MO.;   TSL numbers  SZ32800,  N'Z3300, and NZ33500  have  been assigned
               to three polychlorinated biphenyls  mixtures.  .However,  none
               of these mixtures  coincide with  the molecular  composition of
               Aroclor  species  as listed by Monsanto.

     The  appropriate molecular  compositon of the Aroclors  is  shown in  Table

 36,1  PhysicalProperties

     Individual polychlorinated biphenyls vary  widely in their  physical pro-
 perties  according to the degree  and position of chlorination.   However,  all
 have a ve'ry low water  solubility,  low  vapor pressure, and a  high dielectric
 constant.  The properties  that make these  compounds  so  widely  used in in-
 dustrial applications  include  excellent thermal stability, strong resis-
 tance  to both  acidic or basic  hydrolysis,  and  general inertness  (Custafson
'1970).   The environmentally relevant physical  properties of  the  Aroclors
 are  presented  in Table 36-2,   Again, it is important to remember that the
 Aroclors are mixtures  of many  different polychlorinated biphenyls species.
 The. physical properties of a mixture cannot be properly defifted  as

                                   Table  36-1
                  Approximate  Molecular Composition of Ar,>cis>rs
                       iPercenc)  'Huczin^ar et 
                                                                                             i:u-.»r Uii
                                                                                                                 Yullufc! '
                                                                           I. 35
f.'nvt f U it-lit
                                               ')   1  f 4 ,OO x Id"  }     4.Ufc x  JO
                                      12. HI
                                                                                                                 «./! x  10        (t.it'i  x Hi
<»,   All «,»luK!> «u
l»,   Krjckcttd Jjt
t.   f*n"v t«_ at.
J.   l»tj«ii.j  a»J !'4
u,   i I. ion «.j.  a_t .
f.   Paitittun -, ut-
                       L let.*,  et
                                                                                                &tgf»t t IC.IHI  *|UaiH

constants.  The  true  solubility  of  a aixture for example is undefined
because the water concentration  at  equilibrium is controlled by partition-
ing of the Individual  components between water and tne mixture  itself,   "hs
physical properties of  the  individual pelychlorinated biphenyls ar.e  not
Identical ,• and,  consequently ,  selective joiubili/ation of the mere sol;die
lower chlorinated components  will occur.  Conversely, the more  Highly
chlorinated compounds  will  oe  ie L,.4;  as
the other properties  of  the mixture will be .sn average of the properties jr
the individual chlorinated  biphrrr.-is.  Theoret icail1-' , trja a k.iowie'd^^  >'.
the equilibrium  aass1  concentrations, the individual solubilities  can be  de-
termined or visa versa  (Dexter and  ?avlou 1973;  "Jaris «_£ ai« 1371; Talp  and
Hutzinger 1978).  !t  Table  3h-2,  the physical prooertits js ieveral
mixtures are presented.

36.4  Suaaary of Fate  Data

    36.4.1  Photolysis

         Irradiation  -it  poiy.:hlorir>at«d biphenyis with s nor r— wave  ..ii:ra-
violec light has been  shown  to produce partiil .iechiorinatioi*.
phase (Ma ugh 19/3) and,  in  water wuh jzjne, essentnily cooplece  i
Cion (Versar 1976).   Experimental results of 3ur.ce ind Kumar { i^73>
cate that the more hlgnly chlorinated ipecics .-.-.' susceptible :
sis, . result ing in the  formation  of  dehalog«na:ed or substitiit«d prod-jo:*.
These products aay be  *ccom;ianied bv chlorinated r. ,p'r,»T. > lene^ Jn«i  .:::t')r;-
naced dibeuzof arat.s {iaf^ e£  a_..  :9"a.'.  Phot >pr aduct s of poly<:nl->ri-
natad biphenyis  isoaers  are  listed  in Taoie 3o-? ,fr--it! i-i»e ej"  al. i )'""'<,
There is some doubt .-is  to tr.e  applicability jf tnese exrerisents  ;.j  er.v; -
roniaental cpnditions,  however, since such experiment* are carried  lut  i -,
(aedia other than water  due  t:>  the low aqueous su i-.ibi L i ty of polychlor-
inated biph«nyls.

         Safe and co-workers  Cl0'*1'' S'lggest the following taechanisa  f;r
photochemical breakdown  or  transformation of polycnlorinated hiphenvls:

         "The accepted  r-';iite  for the photochemical excitation of  aroffiarics
in the '280-3^.) no regi >n occurs  5y  a trarsici'^n of electrons i;i the -_
ground state to  an excited8  state ("*).  Frora the excited state, «fhi,:h  ."an
be of singlet or triplet nult illicit .•, the carbon-halogen bond  underijot;^
fission giving rise to  an aryl and  a hydrogen radical.  The radicais tr,e:i
abstract hydrogen rr-ins  the  medium jr -iinefi^e ind, in addition, a  hydrogen
haiide can also  be detected.   Prior to bond fission an alternative reaction
between the excited state and  a  nucler.phiiic species ca1". occur  giviig  the
appropriate substitution product -it the C-X bond."

                                f«bl«   36-3
 ftOTOPEODUCfS  OF PCI  ISOME1S  (from Safe et.  al.  1976)
     PCS Isonwr
3.*-dicW  robipnenyi
                                2.*,* -'rxsiorocipnenyi

I.Z .3.3 -t«sr»£nioro6!enenyt
I -i J -Tieflwoesor.enyi
3,3 -
         When the rate of polychlorinated biphenyis photolysis  In  oxygen-
saturated solutions is compared with the rate in anoxlc  solutions,  it  is
evident that ,-jxvgen sopresses photolvsts, apparently by  acting  as  a  free
radical quencher.  This supports the cone-melon that art  Intermediate  trip-
let excited state is involved in the photolytic reaction'(Safe  et  a1.
1976).  Since In natural water:; the photic zone corresponds  to  the  aost
oxygen-rich zone (due to photosynthesis) , dissolved oxygen  is an.inherent
limitation on the rate of aqueous polychlorinated bipheayl  photodegrada-
tlon.        H                        ,                  •

         Sunce and, Kusiar (1978) -neasured photolysis QC several  polychlo-
rinated biphenyl species in a •*:! acetoni.rile-water solution and  calcu-
lated the approximate rate of dechlo'rinacion under environmental  condi-
tions.  For shallow waters and typical environmental PCS concentrations,
they predicted that up -to 5* of the lightly chlorinated  polychlorinated
bipber.yls sclecules night lose'a chlorine atom in otic year,  but that  at •
least .:ne chlorine atom should be lost from each highly  chlorinated  mole-
cule annually.  The authors' stress the limitations of the assumptions  Bade
in this calculation, but if accurate, it is very significant since  the
lass-chlorinated species can be biodegraded while the heavier species  can-
not.  Furthermore, Bur.ce and Kumar point out the fact that  if polychior-
inaced biphenyis concentrate heavily in the surface film, the photoiytic
rates wo u 1-1 be such higher than those calculated.

         Thus i: is possible that photolysis results in  the  breakdown  of
the "sors hi^hlv-chlorlnated polychlorinated biphenyl species.   Eve*n  though
the rate of such breakdown is slow, it Is significant since  .•'one  of  the
other fate processes result la degradation -if the heavier species.

    3 6 . -k. 2  Ox Id at ion                                                   • .

         Pol/chlorinated biphenyis are extremely resistant  to oxidation
(Hutzinger _et aj.. l97.«).  Gustafson '1970) cites a Monsanto  technical  V--1-
letin that "they can be Heaced to 140aC ur.der 260 p.s.i. of  oxygen  pressure
without showing any evidence of oxidation a? judged by development  of  acid-
ity or foraatlon of sludge."  Thus cheta. ?1 oxidation is r.ot important as
an environmental fate process.

    36 . 4. 3  Hl_y ..plysis       •  '     .

         Poiychlorinated biphenyis are strongly resistant to both1 acidic
and basic hydrolysis (Gustafson 19?0;  Hutzinger jejt al.  1974).  Hydrolysis
does not appear to be important In terms of environmental fate.

  ' '• 36.4.4  Volatilization

          Volatilisation and/or aerosol dispersion of  polyehlorinated  bi-
 phetiyls  is thought  to  account  for  their ubiquitous distribution.   Atmo-
 spheric  transport  followed  by  fallout  with du^t  or rain  is  largely respon-
 sible  for the  fact  that polyehlorinated biphenyls have been detected  in air
 over Baltimore,  Maryland  (Paris, e_c  a_l,  1978)  in "rainwater in  England,
 brown  teals off  che coast of Scotland, white-tailed eagles  in Sweden,  cod
 in the Baltic  Sea,1 mussels  in  the  Netherlands, Adelie penguins  in the  An-
 tarctic*, brown tpelican eggs in-Panama,  Arctic terns, shrimp in Florida,
 river  water in Japan (and)  waters  in the Great Lakes" (Gustafson  1970).
 Although these compounds  have  a fairly high molecular weight and  low  vapor
 pressure, they have a  remarkably high  activity coefficient  in water which
 causes 9 high  equilibrium vapor partial pressure (Mackay and Wolnoff  1973).
 As a result, Che rate  of  volatilization is somewhat higher  than might  be

          Mackay  and Wolkoff (1973),  and later Mackay  and Leinonen (1975)
 calculated the evaporative  half-life of Aroclors 1242, I24f, 1254,  and 1260
 in a water column  1 meter deep on  che  basis of an approach  assuming therao-
 dynaaic  equilibrium.   The estimated  half-lives for the Aroclors published
 by Mackay and  Leinonen range from  9.5  hours for  Aroclor  1248 to 12.1  hours
 for  Aroclor 1242 (see  Table 36-4).   The Mackay  and Wolkoff (1973)  model
 suggests that  volatilization results from the entrainnent of solute mole-
 cules  in the stream of evaporating water molecules.   The rate at  which the
 solute molecules are so carried off  is determined by  the effective  surface
 concentration  of  the solute and by the evaporative rate  of  the  water.   The
 model  proposed by  Mackay and Leinonen  (1375) is  based on the assertion tnat
 equilibrium will be reached when the chemical potential  for the solute in
•the  air  and in che  aqueous  phases  are  equal;  che driving force la  the
 actual inequality of these  chemical  potentials.   Here, again, the  con-
 trolling concentration is that of  the  solute at  the surface. Neither  jf
 Ch*g«  models are directly concerned  with these processes that act  la  the
 aqueous  phase  to control  the surface concentration and thus are jased  on
 the  assumption of  sufficiently perfect nixing within  the aqueous  phase so
 that the surface phenonnena are rate-controlling in the  volatilization
 process.  Since  the Ma.kay  and Leinonen (1975) eodel  h«* a  more acceptable
 thermodynamic  basis, the results presented herein at* calculated  from  this
 •ode!  rather than  from Mackay  and Wolkoff (1973),

          It should  be  noted that, «van chough  the Aroclor tnixturec  1016,
 1221 and 1232  have  larger proportions  of the more volatile  (i.e., higher
 vapor  pressure)  mono-, di-  and trichlaro species that does  1242,  they  also
 exhibit  higher ultimate aqueous solubility so  that, without a direct
 measure  of the Henry constant, it  is not possible to  assert chat  these
 species  will volatilize more rapidly than does Aroclor 1242.

                             Table  36-4

 Calculated Volatilization Half-Life at  25°C, in 1 »  water, 1 m deep
    Preparation.          Henry Law Cong cane.             I/j	
                           [Atmos m-'/raoTe 1

    Aroclor 1242*    |        5.7 x 10"4       '            12..1

    Aroclor 1248*            3.5 x 10~3                 ,   9.5

    Aroclor 1254*            2.8 x 10"3 "                  10.3

    Aroclor 1260*            7.1 x 10"3                   10.2

    Aroclor 1016b                -                         9.9
*MacKay and Leinonen 1975.

b?ari* et al. 1978.

         Haque et aJL. (1974) demonstrated that volatilization of Aroclor
1254 Is touch slower from soil than from sand or unadulterated PCS solution.
Apparently adsorption by organic materials in the soil greatly reduces the
concentration of "free" PCB's In solution which can evaporate,

         The reduction of volatilization by adsorption nay explain the fact
tnat polychlorinated biphenyls in most natural water samples evaporate at a
touch slower rate than that predicted by Mackay and Wolkoff (1973) or.Mackay
and Leinonen (1975).  Oloffs ejt al. (1972) showed that .volatilization loss
of Aroclor 1260 from solutions of 100 Pg/1 in river water was only 67%
after 12 weeks.  When sediment was added to the system maximum loss was
only 34% after 12 weeks (Oloffs jet al. 1973).  Tucker £t ai. (1975), de-' -
aonstrated volatilization rates of 4.2 and 3.6£ for Aroclor, 1221 and 1016,
respectively over a 24-hour period in aerated samples exposed to activated

         All of the above measurements indicate volatilization half-lives
considerably longer than those estimated by Mackay and ttolkoff (1973) and
Mackay and Leinonen (1975).  The estimated half-life resulting from the
Oloffs _et «JU (1973) experiment would be of the order of 1260 hours, which
is to be compared to the Mackay and Leinonen (1975) estimate of 10.2 hours.

         Although empirical evidence shows that volatilization is slow
under natural conditions, the lack of destructive processes for the more
highly chlorinated polychlorinated biphenyls indicates that volatilization
Is an Important transport process.  The persistence of these compounds,
along with the transport afforded them by volatilization, is prdbably the
major factor in their widespread distribution.

    36,4.5  Sc-rgtion

         Adsorption to sediments is the major non—destructive process af-
fecting  polychlorinated biphenyls after introduction to the aquatic en-
vironment.  The combination of low water solubility (.0027-15 mg/1) and
high octanol/water partition coefficient (2,8-7.4) (also see Table 36-2)
indicates that polycblorinated biphenyls will have a high affinity for sus-
pended solids, especially those high in organic carbon (Hamelink _et a_l.
1971),  This has been confirmed by a number of experiments which have shown
that polychlorinated biphenyls are quickly adsorbed and that the greatest
amount is usually associated with sediments or soils in soil-water systems
(Oloffs et_ al. 1973; Haque ejt al. 1974;  'Heeling ejt al. 1978;  Moein 1976;
Hoe in jet a_l. 1976;  Paris <£ al. (1978).  Paris et. al, (1973) found the
partition coefficients for polychlorinate-l biphenyl between water and a
variety of bacteria, seston, and sediments to be only an order of magnitude
less than the corresponding octanol/water partition coefficient.

          When contami.itac.L3n becomes sufficiently high, sediments *may ser«e
as a reservoir , for re-solufion of polychiorinated biphenyls (Veich and
Coins Cock 1975).  This fact has Important caaif ications for areas where
polychlprinaced biphenyls are spilled; even after the Initial degradation
in water quality,•release of these compounds by sediments can cause long-
term pollution.  For example, Wilson and forester (1978) report that for 7
years following a  spill, oyster tissue (Crassostrea virgintea) still Con-
tained seasurable  quantities of Aroclor 1254 even though the amount in
water was below detectable levels.

         Polychiorinated biphenyls have an even greater affinity for oil
than for sediaents (SayLar and Colwell 1976}.  'In areas polluted by both
oil and polychlorinated biphenyls, microbes capable of degrading each of
the pollutants separately appear to' be inhibited by the high concentrations
in combination.

         The preferential sorption of polychlorinated biphenyls on the
organic fraction of the suspended solids coupled with the entrance of these
suspended solids into the sediments is thought to be a major mechanism for
the immobilization in aquatic systems.  The persistance of these chemicals,
however, makes re-solution a possibility for years after they have entered
the sediments.

    36.4.6  Bioaccumulation

         Since polychlorinated biphenyls are adsorbed strongly to organic
sediments, it is not surprising to find that they are also strongly bio-
accumulated.  Their resistance to biodegradation implies that they will be
quite persistent in organisms.  This is indeed the case,

         The potential for bioaccutnulation of polychlorinated biphenyls
species is directly related to,the number of chlorines for two reasons:
first, the more highly chlorinated species have a greater octanol/watfer
partition coefficient, and secondly, the heavier polychlorinated biphenyls  •
species are more resistant to biodegradation (Metcaif _et _al. 1975;  EPA

         Numerous  authors have published the concentrations of polychlor-
inated biphenyls in various aquatic and terrestrial organisms.  Although
tne ambient concentrations are no't often reported along with the _ka viv£
concentrations, it is suggested that bloaccumulation factors are usually on
the order of LO^.-lO*5.  Although these factors are. very high, it is
probable that the  total reservoir of polychlorinated biphenyis in the oce-
anic water column is higher than the total reservoir in oceanic biota
(Clayton et al. 1977).            •                                 ,

         Poly-chlorinated biphenyls can be passed along a food chain; - how-
ever, biomangification is apparently not a controlling factor in attaining
the levels found in aquatic organisms (Meccalf et al. 1975; Clayton ej: al.
1977).  Scura and Theiiacker (1977) found that the partition coefficient
for each organism in an aquatic food chain.determined the ultimate level
and not the food chain itself,

    36.4.7  Biot^ansforoation and Biodegradation

         Individual polychiorinated biphenyl species vary widely in their
susceptibility to biodegradatian.  The nono-r di~, and tri-chlorinated
species can'be degraded by an array of organisms ranging from bacteria to
man.  Polychiorinated biphenyis with five or nore chlorines per molecule
are quite resistant to biodegradation.

         The mechanism of polychiorinated biphenyl biodegradation has beeni
investigated thoroughly (Hutzinger _et al. 1972;  Metcalf e_t al. 1975;
Branson _e_t _al. 1975; Berlin _et al. 1975;  Kaiser and Wong 1974;  Wong and
Kaiser 1976; Furakawa et al. 1978).  For the purposes of this study, it
suffices to say that biodegradability of these compounds is a function of
the number of C-H bonds available for hydroxylation by aicrosomal oxida-
tion.  Adjacent unchlorinated carbons allow the formation of arena oxide
intermediates and thus facilitate metabolism.  There are exceptions to this
simple rule however.  For example, in a study of the biodegradation of 31
isot.ers of polychiorinated biphenyls by Alcaligenes and Actinobac_ter, chlo-
rinated biphenyl containing two chlorines on either the ortho position, of a
single ring (i.e., 2.6-} or both rings (i.e., 2,2'-) showed very poor de-
gradability (eurakawa ££ _al_. 1978).  Whether this will be a general pheno-
menon for all organisms is yet to be determined.

         Rates of biodegradation vary widely, depending on trie composition
and distribution of biota, concentration of polychiorinated, biphenyls,
availability of other nutrients, temperature, and other factors.  Tucker et
al. (1975) reported the rates listed below fot degradation of mixtures of
polychlorinated biphenyls in an acclimated serai-continuous activated sludge
experiment wich 48-hour exposure:*

        Aroclor                            % Degradation

         1221                              31+6
         1016      .         '       •        33 + i!4
         1242                              26+1*
         1254   .                           19+38

         Tulp et al, 1978, In contrast, found tnat the- metabolism of f~
single compound 4,4'-dichlorobiphenyl by a mixed culture of bacteri?
laced from activated sludge was almost totally suppressed by-alte*

carbon sources such as glucose, glycerol, peptone, yeast extract, hunie
acid or activated sludge. , Though the interpretation of the latter result
Is trivial It points out the difficulty in extraipolating conclusions drawn
from laboratory studies to the environment.

         Wong and Kaiser (1976) published data on the degradation of 2-
aono-chiorobiphenyl (0.05%) and 4-monochlorobiphettyl (0.052) by a mixed
culture of bacteria (predominately AehroaiQbacter _s£) isolated from lake
waters.  Interpolating from a graph presented in their paper, the half-life
of the forcer was about 100 hours and the latter was 175 hours.

         Metealf and co-workers (1975) studied the behavior of 2,5,2'-ert-
chlorobiphenyi;  2,5,2',5'-tetrachlorobiphenyl;   and 2,^,5,2*,5*,-penta-
chlorobtphenyl in a model ecosystem.  After 33 days, the concentrations of
the parent materials and metabolites in the water column and organisms were
measured.  The trichlorobiphenyl species was degraded considerably, but the
tetra- and penta-chlorinated forms were, for the most part, unchanged and
were strongly bioaccuraulated.

         The polychlorinated biphenyl species composing the heavier Aroclor
mixtures are essentially non-biodegradable.  Oloffs et al.  (1972) showed
that there was no degradation of the polychlorinated biphenyls in Aroclor
1260 over 12 weeks in natural water sample's.  In a study of the distribu-
tion and fate of Aroclor 1254 after a spill of transformer  fluid, Moein et_
al. (1976) concluded that "no detectable reduction in the concentration of
this mixture in the soil has occurred as the result of chemical transfor-
mation or biodegradacion".  This study compared  the polychlorinated bi-
phenyls found in soil samples in 1975 to those found in mixtures  (askarals)
used in the transformer in 1973.

         To • summarize, biodegradation is very likely to be  an important
fate prdcess for the mono-, di-, and tri-chlorinated biphenyls, but does
not have a signficant effect on polychlorinated  biphenyls with five or more
chlorines.  The tetra-chlorinated biphenyls are  intermediate in their sus-
ceptibility to biodegradation.

36.5  Data Summary

    Biodegradation is the only process known to  transform polycrhloriaated
biphenyls under environmental conditions, and only the" lighter compounds
ara measurably biodegraded.  There Ls experimental'evidence that  the
heavier polychlorinated biphenyls (five chlorines or more per raolecuie) can
be photolyzed by ultraviolet light, but there are no data to indicate that
this .process is operative in the environment.

    Volatilization Is probably responsible for the global disperson of
polychlorinated biphenyls, but on the basis of mass, aqueous transport
(aicher adsorbed1 to suspended solids or in "solution") is probably the more
significant process,

    Polychlorinated biphenyls are strongly partitioned to organic solids
and biota.  Their incorporation into deep 'sediments is-an Important sink
resulting in immobilization.  However, resuspension of these sediments
could cause them to be released back into the water column.

    The fate of polyehibrinated biphenyls is summarized in Table 36-5.

                                                                           Table 16-5

                                                      Summary of Aquatic Fate of Polychlorinated Blphenyls
Environmenta!                 Summary
   Process*                 Statement
                                                                                                                                   at  Data
Photolysis            May result in destruction of heavier

OxiUatlc.i     _       PCBs are stable to oxidation.

Hydrolysis            PCBs are stable to hydrolysis

Volatilization        Important mechanism for transport.         Varies widely
                      Volatility depressed "by presence
                      of organic solids.

Sorptlon              PCBs strongly adsorbed by solids.          Typically rapid
                      especially with high organic content.

Bloaccunulatioa       Bioaccumulation factors range from         Typically rapid
                      about 10* - 10* .

                      The only ptoven mode of destruction        Varies widely
                      of PCBs, but only important  for
                      those with fewer than 4- chlorines
                      per Molecule.
                                                                                                            Varies  widely





          a.   Fhere is  insuifit lent  intormation  in the  reviewed  literature  to  permit  assessment  of  a moat  probable fate.

36.6  Literature Cited

Berlin, M.,  J,  Gage and S.  Holm.  1975.  The distribution and metabolism of
  2,4,5,2,' ,5 '-pentachlorobiphenyl.  Arch. Environ, Health  30:141-147.

Branson, D.R.,  G.E. Blau, B.C., Alexander and W.B. Neely.  1975.
  Bioconcentration of 2,2*,4,4*-tetrachlordbiphenyl in rainbow trout as
  measured by an, accelerated test.  Trans. Am. Fish. Soc.  4:785-792.

Bunce, M.H.  and Y. Kumar.  1978.  An assessment of the impact of solar
  degradation of PCB's in the aquatic enviorianent.  Cheraosphere

Chiou, C.I,, V.H. Freed, D.W. Schmedding and R.L. Kohnsrt.  1977.
  Partition coefficients and bioaccumulation of selected organic chemicals.
  Environ. Sci. Technol.  ll(5);475-478.

Clayton, J.R.,  Jr., S.P. Parlov, and N.F. Breitner.  1977.  Polychlorinated
  biphenyls in coastal marine zooplankton:  bioaccumulation by equilibrium
  partitioning.  Environ. Sci. Technol.  il(7):676-682.

Dexter, R.N. and S.?. Pavlou.  1978.  Mass solubility and activity  ,
  coefficients of stable organic chemicals ,in the marine environaent:
  Polychlorinaeed Biphenyls.  Mar. Chen. 6:41-53.

Environmental Protection Agency.  1977.  Review of the environmental fate
  of selected chemicals.  Office of Toxic Substances, EPA 560/5-77-003.
  Washington, D.C. p. 79-82.

Furakawa, K«, K. Tonomura,  and A. Kaoiibayashi.  1978.  Effect of chlorine
  substitution on the biodegradabi'ity of polychlorinated   biphenyls.
  Appl. Environ. Microbiol.  35(2)=223-227.

Gustafson, C.G.  1970.  PCi's -prevalent and persistent.  Environ, Sci.
,  Technol. 4:814-819.

riaoelink, J.L., R.C. Maybrant and R.C. Sail.  1971.  A proposal:  Exchange
  equilibria control the degree chlorinated hydrocarbons are biologically
  magnified  in ientic environments.  Trans. Am. Fish. Soc.  100:207.

Hansch, C.,  A.  Vittoria, C. Silipo, and P. Jow,  1974.  Partition-
  coefficients and the structure - activities relationship of the
  anesthetic gases.  J. Med. Chem.  18(6):546-548.

Haque, R., D.W. Schmedding, and V.H. Freed. -1974.  Aqueous solubility,
  adsorption, and vapor behavior of polychlorinated biphenyl Aroclor 1254.
  Environ. Sci. Technol. 8(2);139-142.

Hetling, L., 2. Horn and J. Tofflemire.   1978.  Hudson River ?CB study
  results,  (unpublished).  N.Y, Scare Dept. of Env. Conservation.  Albany,

•Hutzinger, 0.» D. M. Nash, S. Safe, A.S.W, DeFreitas, R.J. Sorscrom, D.J.
  Wildish, and V. Zltko.  1972.  PCB's:   Metabolic behavior of pure iaomers
  in pigeons, rats, and brook trouc.  Science 178:312-318.

Hutzinger, S., 3. Safe and ?. Zltko.  1974.  The chemistry of PCBs.
  Chemical Rubber Publishing Co., Cleveland, Ohio.

Kaiser, K.L.E. and   P.T.S. Wong.  1974.  Bacterial degradation of PCBs:
  I. Identification of   some metabolic products from Aroclor 1242.  Bull,
  Environ. Contam. 4 lexical.   11(2):291-196.

Mackay, 0. and P.J. Leinonen.  1975.  Rate of evaporation of low-solubility
  contaminants from tracer bodies to atmosphere.  Environ. Sci. Techno!.

Mackay, 0. and A.W. Wolkoff.  15*73.  Rate of evapo-ation of low-solubility
  contaminants front water bodies to atmosphere.  Environ. Sci. Technol.

Ma ugh, T.H.  1973.  DDT, an unrecognized  source of polychlorinated
  biphenyls.  Science 180:578-582.

Metcalf, R., J.R. Sanborn, P.-Y. Lu and D.! Uye.-  1975,  Laboratory model
  ecosystem studies of the degradation and fate of radiolabeled tri, tetra
  and penta chlorofciphe'tiyl c on pa red to DDE.  Arch, of Environ. Contatn. and
  Toxicol.  3(2):l5t~l65.

Moein, G.J.  1976.  Study of distribution and fate of PCSs and benzenes
  after spill af transformer fluid. ! SPA  904/9-76-014.  Washington, D.C.
  128 p.

Moein, G.J., A.J. Smith, Jr«, and P.L, Stewart.  1976.  Follow-up study of
  the distribution and fate of PCBs and benzenes in soil and ground water
  samples after an accidental spill of transformer fluid.  ln< Proc. of 1976
  Sat'l. Conf,  on Coucrol of Haz. Mmt'l. Spills.  Information Transfer,
  Inc. Rockville, Md. p. 368-372.
Monsanto Industrial Chemical Corp.
  306A.  St. Louis, Mo.
1974,   'PCBs-Aroclors Tech.  Bull.   O/PL
Oloffs, P.t., L.J, Albright, and $.¥. Szeto.  1972.  Fate and behavior of
  five chlorinated hydrocarbons in three natural waters.  Can. T.
  Mlcrobiol. 18:1393-1398.

Oloffs,  P.C.,  L.S.  Albright,  S.Y.  Szeto  and  J.  Lau.   1973.   Factors
  affecting  the  behavior  of  five chlorinated hydrocarbons  and  their
  sediments.   J.' Fish,  Res.  Board, 'Canada.   30:1619.

Paris,  O.F., W.C.  Steen,  and  G.L*  daughman.   1978.   Rcle  of tne physico-
  chemical  properties , of  Aroclori 1016 and 1242  in determining  their rate
  and  transport  in aquatic envirooents.   Chemosphere.   7 (4) :319-325.
                      !                                 I
Safe,  5., N.J.  Bunce,  8.  Chictin,  0.  Hutzinger,  and  L.O.  Ruzo.   !976.
  PhotTidecomposition of halogenated  aromatic expounds.  In:
  Identification and analysis of pollutants  in  water.   L.H. Keith (ed,),
  Ann  Arbor  Sci. Publ., Ann  Arbor, Mich. pp. 35-47.

Sayler,  C.S. and R.R.  Col veil.   197?.  Partitioning  of mercury  and
  polychlorinated  biphenyl by oilj water, and suspended sediment.  Environ.
:  Sci.  and  Techno!.  1QU2) : 1142-1145.

Scura,  E.p.  and  G.H. Thellacker.   1977.   Transfer of  the  chlorinated
  hydrocarbon  PCS in a laboratory  food chain.  Mar.  Biol.   40(-») :317-325.

Thonij  M.S.  and  A.R. Agg.   1973.  The breakdown  of synthetic organic
  coopouads  in biological processes.   ?roc.  Royal Soc . London  139:347-357.

Tucker,  E.S.,  V.W.  Saeger, and 0.  Hicks.  1975.   Activated  sludge primary
  oiodegradacion of polychlorinated  bipheayls.   Bull.  Saviran.  Contao.
  Toxicol,  l4f 6 ): 703-713.

Tulp,  M.Ch.M.  and  0. Hutzinger,  1976.  Some thoughts  on  aqueous
  solubilities  and  partition  coefficients nf ?CB, and  the  rsa t hematic a 1
  correlation^  between bioatcuraulaticn and physico-chemical properties.
  C'neao sphere.   7(10) ;849-8feG.
'Veich,  G.D.  and V.M.  Comstock.   1975.   .Ssdiaents ,ad a PCB source to acjua.ic
-systems.-  J.  Fish.  Res.  Board, Canada.   32:1849.

Versar, Inc.  1976.   Assessment  of  wastewater nanagemenc, treatment
   technology and  asssociated  costs  for the abatement of PCS concentrations
   in  industrial effluents.   N'T IS ?3 231-433/AS.   Sprlngfiald,  Va.   256pp.

Wilson, A.J. and  J.  Forever.   1973.   Persistence  af Aroclor 1254  in a
   contaminated  estuary.   Bull. Environ. Cone an.  Toxicol.   19 :6,37-t>4-o.

Wong,, P.T.S. and  K.L.E.  Kaiser.   1976.  Bacterial  degc=*dat ioa  •:. : ?C3s:  II.
   Rati  Studies . -Bull.  Environ. Contam, aa*4 Toxicol.  1 3(2} :249-253.

Yoshiauri, H.  and K.  Yairauioto.   1973.   Metabolic studies  of pclychlorin'ated
   bipienyls.'  1.  Metabolic  fice  of  3,4 ,3' .4'-tetrachloroMphenyl in rats.
       . Pharm.  Bull.  21:1168.

                         37.  2-CHLQRONAFHTHALENE
37 .1  Scaeeffl_en_t of Probable ............ Face

    Very little data specific to 2-chloronaphthalene were found; the aqua-
tic fate of this compound is inferred from data summarized for naphthalene.
The results of the data summary, which includes theoretical and amphiricdi
evidence, suggests that 2-chloronaphthalene, a compound only slightly solu-
ble in watei (6.74 mg/'l), will be adsorbed onto suspended particulars and.
biota and that its transport will be largely determined by the hydrogeo-
logic conditions of the aquatic system.  That portion of 2-chloronaphtha-
lene dissolved in the water column may undergo direct photolysis.  The ul-
timate fate of the 2-chloronaphthalene which accumulates in1' the sediment is
believed to be biodegradation and biotransfortaation by benthic orga-
    2-Chloronaphthalene is present in the environment from anthropogenic
sources.  It is normally not found alone- but as a cotuplex mixture of naph-
thalenes having varying degrees of chlorinatiou,  Commercial preparations
are marketed under the trade name Halawax with only 1000 and 1031 contain-
ing  stonochlorinated 'species (Rover 19755.  As a group, chlorinated naph-
thalenes are not as widely distributed in the environment as the poly-
chlorinated biphenyls.  The survey of organics in water by Shackelford and
Keith (1976) does not show, however, 2-chloronaphthalene to ,be a widespread
pollutant.  Cruatp-Wiesner e_t al. (1973) -report the presence of 2-chloro-
naphtha-lene in sediment samples and Law and Goerlitz (1974) confirm its
presence in the Guadalupe, River of the San Francisco Bay area.

    The chemical structure of 2-chlofonaphtfhalene is shown below.

                                         • - Alternate Names

                                            Halo wax

    CAS No. 91-58-7
    TSL So. OJ 22750

37,3  PhysicalProperties

    The general physical properties of 2-chlaronaphehalene are as follows.

    Molecular weight                         . 162.62
    (Weast 1977) ,

    Melting Point                               61°C
    (WeasC 1977)

    Vapor Pressure at 20"C                     ! 0,017 Corr (calculated)

    Solubility in waCe;r at 25°C                 6.74 ag/1 (calculated)

    Log oetanol/water partition coefficient     4.12
    (Calculated as per Leo et_ a_l_. 1971)

37.4'  3uanaaryof JFat:eMData

    37.4.1  Photolysis

         2-Chloronaphthalene exhibits moderate adsorption in the 300 ma re-
gion and is, therefore, susceptible to direct photolysis or photooxidation
(Radding et al. 1976),

         Recent photolysis studies have shown a potential for photodegra-
dation of polychlorinated naphthalenes in the environment (Ruzo et a1_.
1975).  Experiments with various polychlorinated-naphthalenes in raethanolic
solution irradiated at a peak energy output of 300 nm resulted in dechlori-
aation and dimerization.  They report a 15 percent dechlorination of
2-chloronaphthalene at a rate approximately '10~^sec~^»

    37.4.2  Oxidation    . .'"   .  .  .,

         In natural water the principal oxidizing species are:  (1)  alkyl-
peroxy (R02*> and hydroperoxy (H0£) radicals generated by photolytic
cleavage of traca catbonyl compounds or from enzymatic sources, an-.i (2)
singlet oxygen.  Singlet oxygen is thought Co be the major oxidant species
involved in the direct photolysis of ''organic molecules.

         No data were found concerning the oxidation of,2-chloronaphtha-
lene.  Naphthalene, however, is believed to have a very lon,g half-life to-
ward oxidation by R0.2* radicals.  The chlorine substituent on naphtha--
len<* (i.e., 2-chloronaphthalene) is. thought to make it even less suscep-
tible to free-radical oxidation.  Oxidation is probably, not an important
fate process for 2-chloronaphthalene. ;

    37,4.3  Hydrolysis

         2-Chloronaphthaiane does not contain groups amenable to hydr-ol--
>is.  Hydrolysis, therefore, is noc thought to be a significant fit«s pro-

    37,4,4  V o_la til i z at i o r.

         An actual volatilization rate is necessary to assess the  impor-
tance if this transport process.  Several authors nave suggested ways  13
estiaia;e volatilization rat-es of compounds from water using theoretical
consideration (Macnay and Wolkorf 19"; Mac it ay and Leinouea 1975).  The-;e
methods, however, ate still under development, require a large amount  of
physical data for the compounds, and still may not predict the actual
volatilization rate.  Another method for determining the role of volatili-
zation is that described by Hill e_£ Jut. (1976). which employs the theiry o:-
fared by Tsivoglou (1967).  His theory states that the volatilization  rate
is directly related1to' the ratio of the compound's volatilization,rate
coefficient to the oxygen rea^ration rate constant which is easily mea-
sured.  Hill's sethod still requires a aea~,urement of volatilization for
the compound.  Measured volatilization rates for 2-chla "onapht'twlene were
not fount! in the literature and an accurate assessment of the role of
volatilization is not possible without these rates.  Work by Lee (1975)'
with naphthalene shows, this compound to be rather volatile when present as
part of an oil spill with the rate of volatilization dependent in  air  and
water temperature, wind spe~d, and wave action

    37.4.5  Sorpcion

         The dati reviewed did not reveal specific partition coefficients
of 2-chloronaphthalene onto .suspended particulars natter or biota.  The
calculated log octanol/water partition coefficient for 2-chloronaphthalene
of 4.12 indicates that the compound should moderately adsorb onto  suspended
particulate'Batter, especially participates high in organic natter.

        'Recent work by Lee and Anderson (1977) show that naphthalene will
accumulate in sediments up to two orders ol magnitude greater than th« con-
centration in the overlying water. , They also showed the importance of
raicraoganisns (plankton and bacteria) in adsorbing and removing napr.chdlene
'from water.  While no data specific to 2-chloronaphthaiene We tie found, its
log ? Qccanol/water partition coefficient and the similarity to naphtitaiane
indicate that adsorption could be an important•transport process.

    37.4.6  B1oaccumuIa t i on

         Little specific data on the bioaccuraulation of 2-chloronaphthalene
were found.  la one experiment Green and Neff (1977) measured the uptake
and release by grass shrimp (galaemonetes pugio) of three chlorinated naph-
thalene mixtures containing varying degrees of chlorination.  Halowax 1000,
containing 60% aonochlorinated naphthalene was apparently accumulated much
less than the mixture containing a greater amount of polychlorinated spe-
cies.  The apparent depuration rate (tj_/2 =1- hrs.) was also greater lea-
ding the authors to conclude that mono- and 'dichloto~specl.es are metabo-
lized and excreted faster than polychlorinated species.

       •  Measurements of naphthalene content in zooplankton exposed to high
concentrations show that significant uptake can occur (Lee and Anderson
1977).  Work by Lee _et al. (1972) reveals naphthalene is readily taken-up
by aquatic organisms and concentrated in the liver, where it is rapidly

         Since 2-chloronaphthalene has a calculated log1 octanol/water par-
tition coefficient which is intermediate, the role of bioaccumulation, is
difficult to assess.  It is probably adsorbed by biota to similar levels
reported for naphthalene.  Like naphthalene, bioaccumulation of 2-chloro-
naphthalene is probably short-term.

    37.4.7  Motransformaciqn and Biodggradation

         Walker and Wiltshire (1953) studied the decomposition of 1-chloro-
naphthalene by soil bacteria and found that two species of bacteria, iso-
lated from the soil, would grow in a mineral salts medium with 1-chloro—
naphthalene a« the $ole carbon source.  They report 8-chloro-l,2,-dihydro-
1,2-dihydroxynaphthalene and 3-chlorosalicylic acid as the major bacterial
metabolites.  Similar results were reported for 2-chloranaphthalene by
Canonica e_t al, (1957),                              :

         Okey and Bogan (1965) examined the rate of metabolism of 1 and
2-chloronaphthalene by sewage sludge bacteria that were enriched on un-
aubstititied naphthalene.  The initial concentration of chlorinated sub-
strate was 1 ag/1 with the substrate being the only source of carbon.
Their work showed that naphthalene was much more easily degraded than
2-chloronaphthalene which was aore readily degraded than 1-chloronaph-

         Ruzo e_t al. (1976) report that chloronaphthalenes are rapidly
metabolized in the pig to their corresponding hydroxylated metabolites.
Furthermore, the chloronaphthalenes are•distributed in the various organs

and tissues whereas the metabolites were concentrated in the urine, bile,-
kidney 'and liver.

         Data for naphthalene (Lee and Rvan 1976; Vernberg 1977; Lee and
Anderson 1977) iadicire It to be rapidly degraded by bacteria and secabo-
lixed by aulti-cellelar organisms.  A biodegradation naif-life of 1 day has
been estimated for naphthalene from these data,

         Thus, it appears that biodegradation and biotransformat ion of
2-chloronaphthalene is rapid ^enough to select these processes as the most
probable in determining the aquatic fate of 2-chloronaphthaier.e.

37,5  Data Summary

    Very little data were found for 2-chJoronaphthalene.  It will probably
be adsorbed to suspended participates, although the role of volatilization1
is unknown at this' time.  The 2-chloronaphthalene adsorbed to suspended
sediment will most likely be taken up by benthic organism and metabolized
at a rapid rate.  That portion of 2-chloronaphthalene dissolved in the
water column may undergo photolysis and will be bj.odegraded by bacteria,
Table 37-1 summarizes the information found for 2-chipronaphthalene.


                                           Table J7-1

                                    vf A>)u«ttc F*t* ol
I lotrau* lucnac Ion/
  1 lodegcadat lo«
                              slow process.
                              not cantata
                         Ma y
a.  TtMtn t* tnauff UlMt infonwtloa In th«
                                                      literature to pemlt »«S«»MMBI ol * BOM probable fate.

37,6  Liceracure Cited

Canonica, L.,' A. Fiecchl and V. Trauani.  1957.  Produces of microbial
  oxidation of some substituted naphthalenes.  Rend.

Crump-Wiesner, H. , R. Feltz and M.L. Yates.   1973.  A study of the
  distribution of polychlorinated biphenyls in the aquatic environment.
  U.S. Geol. Sun. J.  of Res.  1:603-607.
     ,  F.A., Jr. and T. M. Neff.  Toxicity, accumulation, and release of 3
  polychlorinated naphthalenes (Halowax 1(00, 1013, 1099) in postlarval , and
  adult grass shciasp, Pa 1 aeggne tea gugi o .  Built Environ. Contan. Toxicol.

Hill, J., H.?. Kollig, D.F. F?:is, N..L. Wolfe and R.G. Zepp.  1976.
  Dynamic' behavior of vinyl oh1'jfide In aquatic ecosystems.  U.S.
  Environmental Protection Agency, (Office of Research and Development)
  Athens, GA. (EPA 600/13-76-OC1-).

Kover,  P.O.  Environmental hazard assessment report: chlorinated
  naphthalenes.  EPA 560/8-75-001, PB-248834.  36p. '

Law,'L.M. and D.F. Goerlic:.  1974,  Selected chlorinated hydrocarbons  in
  bottom material from screams tributary to San Francisco Say.  J.
  Pesticides Monitoring 3(l):33-36.

Lee, R.F., R. Sauerherber, and G.H. Dobbs.  1972,  Uptake, metabolism, and
  discharge of polycyclic aromatic hydrocarbons by marine fish.  Marine
  Biol. 17(3}:201-208.
Lee, R.F. and C. Ryan,  1976.  Blodegradation of petroleum hydrocarbons by
  marine aicrobes, _tn Proceed. Third Inteaat. Conf, on Biodegradation.
  Applied Publishers, Londoiu

Lee, R.Ft and J.W. Anderson.  1977.  Fate and effect of naphthalenes:
  controlled ecosystem pollution experiment.  Bull, Marina Set.  21:i27,

Lee, R. F.  1975.  Fate of petroleum Hydrocarbons in oiari.ie zooplankton.
  Am«rican Pat'roleum Institute Conf.  8:549-553.

Leo, A., C. War.seh, and D. 'Elktns.  1971.  Partition coefficients and their
  uses.  Cheau Rev. 71:525-616.

Mackay, 0. and p.J. Lelnonen,  19751.  Rate of evaporation of low-solubility
  contaminants from water bodies to atmosphere.  Environ. Sci. Technol.

Macka/, D. and A.W. Wolkoff.  1973,  late of evaporation of lov-solubility
  contaminants front water bodies to atmosphere.  Environ. Sci. Technol.

Okay, R.W. and R.H. Sogan.  1965.  Apparent involvement of electronic
  mechanism in limiting microbial metabolism of pesticides.  J. Wat.dt Poll.
  Control Fed.  37:692-712.

tadding, S.S., f. Mill, C.W. Gould, D.H. Llo, H.L. Johnson, D.C. Boiaberger,
  and C.V. Fojo.  1976.  the environmental fate of selected polynuclear
  aromatic hydrocarbons,  U.S. Environaental Protection Agency, (Office of
  Toxic Sub.}, Wash., D.C.  122p.  (EPA 560/5—75-009).

luzo, L.O., N.M. Bunce, S. Safe, and 0. Hutzinger.  1975.  Photodegradation
  of polyehloronaphthaien*s in tnethanol solution.  Bull. Environ. Contain.
  Toxicol.  14(3) :341-345.

Ruzo,. L.O., N.H. Bunc*, S. Safe, and 0. Hutzlnger.  1976.  Uptake and
  discribucioin of chloronaphthalenes and their metabolites in pigs.  Bull.
  Environ. Cont«s, Toxicol.  16(2):233-239.

Shackalford, W.M. , and L.H. Keith.  1976.  Frequency of organic coopounds
  identified in water.  U.S. Environmental Protection Agency, (Office of
  R«March and Development), Attien», GA.  618p.  (EPA 600/4-76-062).

Taivoglou, E.G.  1967.  Measurements of stream reaeratlon.  U.S. Dept.,
  Int., Washington, D.C.                                             ,

¥«rnb«r§, F.J.  («d.>,  1977.  Physiological response* of aarine biota to
  jrollutants.  Academic Press, New York*  pp.3,23-340.

Walker, M. and G.H. Wiltshire.  1933,  Decomposition of 1-chloro and
  l-brcwo-ftafhth*len« by soil bacteria.  J, G«n. Mlcrobiol,  12:478-483.

W«a»t,,R.C. (ed.).  1977.  CRC Handbook of Chemistry and Physics, 58th
  Edition.  CRC Press.  Cleveland, Ohio.

                                   TECHNICAL RgPOBT DATA
                           •fie/at read Instructions an me rtvtrse 'ittott ;o*v
 EPA 440/4-79-029a
 Water-Related  Environmental  Fate of
 129 Priority Pollutants,  Volume I
                                                          .1 Pec ember
' *UTHOS M.  Callahan1 ,  M,  Slimak"' ,  M. Gabel2, I. May",
C. Fowler2,  E.  Freed2,  P.  Jennings", R. Durfee2,          I
F.^Vhitaore2.  B.  Maestri2.	W.	Mabgy . B. Holtj, C.	Gould3!
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                                                           13. TYPE oc agcgsr A NO »S=!iOO
                                                             Final  Report
                                                           114. SK)NSOHI.NC AOINO1 COOS
 This is a combined Versar/SRI/BPA literature search.
  , *8ST»AC.
           This report  is  a literature search and summary of relevant  data for the
  individual fate processes (hydrolysis, biodegradation, photolysis, etc.)  which might
  be expected to occur  if  a pollutant were introduced into an aquatic  system.   The
  report is organized into 101 individual chapters for pollutants or small  groups of
  pollutants, and four  introductory chapters.  Each chapter has  its own  references so
  the chapters can be used iadependently.
           The approach  taken  by this report is to summarize data on the individual
  processes which might  be important in' describ'lng the transport and fate of  pollutants
  introduced at low concentrations (e.g., ppw or less) into aquatic environments.  If
  transport processes will result in  ,sigt.lficant pollutant transfer to  another medium
  '(e.g., air, sediments),  data are included where available to describe  what  happens
  to the pollutant in the  medium to which the pollutant was transferred.
           A list of the literature covered in the search is included.
           Results of the  literature search are that a significant amount of
  information on most pollutants was found, bur. that the information was  more  useful in
  making qualitative judgements about the pollutant transport and fate than for making
  quantitative predictions, of  concentrations in the environment.  Availability of
 . rate constants useful  in mathematical fate models was limited.
                                                    f I6BS/OPSH E^QiD Ti«)MS C.  COSAti Field. Group
Environments                       ienzen*
Pollutants                       " Toluenes
Distribution                       Phenols
Transport   '                       Pesticides
Water Pollution                    Metals
Halogenated Hydrocarbons           Esters
Organic Sitrogen Compoun-ls         Mitriles
                    Hydrocarbons   F.frhetfa   m
                                                Priority Pollutants
                                                Environmental  Fate
                                                                          07B,  Q7C,  07D
                                                                          68D,  63E,  68A
                                              19. SiGij«iTv C'.iiS This i
                                                                         it MO. 3«= JAuSS
                                              20 sieuaiTY C.ASS .