ROBERT A. TAFT WATER RESEARCH CENTER
                         REPORT NO. TWRC-1
   DILUTE SOLUTION  REACTIONS
   OF THE NITRATE ION
   AS APPLIED TO WATER  RECLAMATION
         ADVANCED WASTE TREATMENT LABORATORY-I
    U.S.  DEPARTMENT OF THE INTER/OR
FEDERAL WATER POLLUTION CONTROL ADMINISTRATION
            OHIO BASIN REG/ON
              CINCINNATI, OHIO

-------
  DILUTE SOLUTION REACTIONS OF THE NITRATE ION

                      AS
                  /
         APPLIED TO WATER RECLAMATION
                      by

         Frank C. Gunderloy, Jr., Cliff Y.
         Fujikawa, V. H. Dayan and S. Gird

                      for

The Advanced Waste Treatment Research Laboratory


      Robert A. Taft Water Research Center
          This report is submitted in
          fulfillment of Contract No.
          14-12-52 between the Federal
          Water Pollution Control Ad-
          ministration and Rocketdyne,
          a Division of North American
          Rockwell Corporation.
       U. S. Department of the Interior
Federal Water Pollution Control Administration
               Cincinnati, Ohio
                October, 1968

-------
                         FOREWORD
     In its assigned function as the Nation's principal natural
resources agency, the United States Department of the Interior
bears a special obligation to ensure that our expendable re-
sources are conserved, that renewable resources are managed to
produce optimum yields, and that all resources contribute their
full measure to the progress, prosperity, and security of
America — now and in the future.

     This series of reports has been established to present the
results of intramural and contract research carried out under
the guidance of the technical staff of the FWPCA Robert A. Taft
Water Research Center for the purpose of developing new or im-
proved wastewater treatment methods.  Included is work conducted
under cooperative and contractual agreements with Federal, state,
and local agencies, research institutions, and industrial organi-
zations.  The reports are published essentially as submitted by
the investigators.  The ideas and conclusions presented are,
therefore, those of the investigators and not necessarily those
of the FWPCA.

     Reports in this series will be distributed as supplies per-
mit.  Requests should be sent to the Office of Information, Ohio
Basin Region, Federal Water Pollution Control Administration,
4676 Columbia Parkway, Cincinnati, Ohio 45226.
                            11

-------
                         ,   ACKNOWLEDGEMENT


     This report is submitted to the Department of the  Interior, Federal
Water Pollution Control Administration,  in fulfillment  of Contract  lU-12-52,
Article III, B.  The work was carried out over the period 5 January 1968
through 5 August 1968 by the Chemical and Material Sciences Branch  of
Rocketdyne Research Division.  Dr. B. L.  Tuffly (Manager, Environmental
Sciences and Technology) served as the Program Manager.  Dr.  F. C.  Gunderloy
(Principal Scientist, Inorganic and Organometallic Chemistry) was the
Responsible Scientist.  Members of the Technical Staff  contributing to
this program were Dr. C. Y. Fujikawa, Dr. V. H. Dayan,  Dr. S. R. Webb,
Dr. M. A. Rommel, Dr. J. Foster, Mr. S.  Cohz, and Mr. S.  Gird.

     Dr. R. B. Dean (Ultimate Disposal)  of the Cincinnati Water Research
Laboratory acted as Project Officer for the Federal Water Pollution Control
Administration.  The Rocketdyne staff wishes to express their appreciation
for the interest, expertise, and guidance provided by Dr. Dean throughout
the course of this research.  Thanks are also due to Mr.  F. M. Middleton
and Dr. D. G. Stephan of the FWPCA for time devoted to  several enlightening
discussions held at the Cincinnati Water Laboratory prior to  the  inception
of this program.
                                 111

-------
                               CONTENTS






FOREWORD	ii




ACKNOWLEDGEMENT	iii



CONTENTS	iv




ABSTRACT	vi



INTRODUCTION 	    1




SUMMARY AND CONCLUSIONS  	    3



LITERATURE SURVEY  	    6



DISCUSSION AND RESULTS 	    3



    BACKGROUND 	    3



         Nitrates in Water - Occurrence and Effects  	    8



         Natural Denitrification 	    9



         Nitrate Removal Processes 	    9



    SURVEY OF CANDIDATE AGENTS	10



         Chemicals in Water Reclamation  	  10



         Choice of Reducing Agents 	  10



         Choice of Deammination Agents	13



         Catalysis in Reduction and Deammination	1^



    DILUTE SOLUTION REACTION STUDIES 	  lU



         Screening of Reducing Agents  	  1^



         Deammination Agents 	  16
                                iv

-------
     FERROUS ION AS A DENITRIFICATION AGENT  ............    18
          Initial Screening Studies  ...........              jo

          Attempted Identification of "Missing N" ..........  20
          Effect of pH on Denitrification ..............  22
          Effect of Catalysts on Denitrification  ..........  2h
          Effect of Air on Reduction and Denitrification  ......  24
          Lime for pH Adjustment  ................ \     26
          Effect of Phosphate and Carbonate  Ions  ..........  26
     ANALYTICAL SUPPORT .......................     3Q

          Background and Selection of Methods ............   30
          Analytical Program   .................         31
          Development  of Ultraviolet  Method  .............   31
          Ammonia Distillation Method ................   37
 EXPERIMENTAL DETAILS  ....................
    APPARATUS   ........................        ,g
    REAGENTS    .......................          .
    PROCEDURE

    TEST RESULTS

REFERENCES
             	45
APPENDIX - ANALYTICAL METHOD	                              a.

-------
                               ABSTRACT
     A new and unexpected partial denitrification of dilute nitrate ion
solutions (10 to 50 ppm NO^-N) has been achieved by treatment with 8 moles
of ferrous sulfate per mole of nitrate in unbuffered alkaline reactions.
The nitrogen loss, which probably represents evolution of N2 or N20, has
been as high as 50$.  Total reduction to lost nitrogen plus nitrite and/or
ammonia has approached 100$.  The reduction takes place in the presence of
partially oxidized black iron hydroxides, and requires catalytic quantities
of cupric ion.  Denitrification is suppressed by phosphates, as well as by
several other factors, some as yet unidentified.  Silver ion catalysis or
a 16-fold excess of the ferrous salt permits reduction to ammonia in the
presence of phosphate, but there is no accompanying denitrification.
Keywords :  Nitrates, wastewater, denitrification,  reduction,  ferrous ion,
           catalysis, cupric ion
                                  VI

-------
                              INTRODUCTION
      The presence of nitrate ion in water,  reclaimed or otherwise, presents
 several distinct problems.   In high concentrations,  it can cause  methemo-
 globinemia, a disease of the newborn,  and it serves  as a nutrient for algae
 at any concentration.  Algae bloom becomes  a problem in many aspects  of
 water usage, such as the fouling of reclaimed waters in reservoir storage.

      Although there have been numerous studies of methods of nitrate
 removal and control for use in water reclamation, removal by some direct
 chemical reaction,  other than ion exchange,  does  not seem to have been given
 any really serious  consideration in the past.   This  is not an oversight on
 the part of the interested  scientific  community;  it  is simply a reflection
 of circumstances that are not obviously amenable  to  attack by way of  some
 common chemical reaction.   For instance,  if  one considers a nitrate-N level
 of 10 ppm in water,  then such ordinary reactions  as  quantitative  precipita-
 tion or nitration of an organic compound  cannot be brought to bear on
 removing nitrate from reclaimed waters.   Precipitation requires expensive
 organic precipitants (e.g.,  "Nitron"),  and nitration of organics  is only
 accomplished in concentrated solutions.

      Conversion of nitrate  to a nitrogenous  gas and  denitrification by the
 evolution of such a  gas is  an attractive  concept, and  there  are three  gases
 that can be considered,  although each has drawbacks  (Ref. A1-A3).  These
 are ammonia,  nitrous oxide  (N20)  and nitrogen.  (Such  gases  as N02, NF3,
 NOC1,  NO,  and the like have  to be discarded  on  the basis  of  reactivity,
 water solubility, or the  simple impossibility of  converting N0§ to such a
 gas in any simple reaction or series of reactions.)  Ammonia has high water
 solubility, and requires a physical stripping step.  It cannot simply be
 left in the water, since eventually, as part of the biological nitrogen
 cycle,  it  will be reconverted to  nitrate.  Nitrous oxide, while not nearly
 as  soluble as Nfo, still has  a fairly high water  solubility compared to
 nitrogen.  Nitrous oxide in  solution does oxidize, albeit at a very lov;
 rate (Ref. A-l). Thus,  at concentrations  in the ppm range, N20 could be
 subject  to the same  drawbacks  as  ammonia.

     As  far as chemical characteristics are concerned, nitrogen itself
would be the  ideal choice, since  it has the lowest water solubility and
greatest oxidation resistance of  all the nitrogenous gases.  However,  in
the  oxynitrogen and hydronitrogen series of compounds, nitrogen is unique
in  that  it is  very difficult to obtain by simple reduction (Ref.  A-3).."
While redox potentials often appear favorable, in the case of N0§ there
is a large activation energy that must be overcome,  and very strong
reducing agents lead to NHo in most cases, while weaker ones take the
nitrogen only  to the +3 oxidation state.

     Nitrogen evolution would be an attractive means of ridding water  of
nitrate ion, and the work described in this  report was undertaken with

-------
 Just that objective in mind, even though there appeared to be no simple
 one-step reaction sequence to go from N0§ to N2.  The premise ofVthis
 work was that there might be efficient and economically attractive two-
 step reaction sequences that could be applied, based on reduction of
 nitrate to nitrite ion, and subsequent deammination of a primary amine
 with the nitrite.  A model reaction series for this concept,  using formal-
 dehyde as the reducing agent and urea as the deammination agent is shown
 in equations (l) and (2) .

                 + HCHO •* 2HN02
           2HN02 + COdlH,,),, •» 2N2 * C02 + 3^0                        (2)
 If applicable to dilute solutions of nitrate  ion,  these known reactions
 would be very attractive.   The products  N2, C02, and IfeO are all innocuous,
 and based on bulk prices,  a chemicals cost of 2 to 3 cents per thousand
 gallons  could be projected for removal of NO§-N at the 10-20 ppm level.

      Further consideration of  the model  reaction sequence allows one to
 speculate about  other potential advantages of a direct chemical reaction
 Incorporation of such a sequence  into an existing  plant might be possible
 by simply adding the  appropriate  metering devices  to inject the reagents
 into the stream.  A chemical process  is  easily adjusted to variable nitrate
 levels;  there is no need to design for the maximum level, which would result
 in unused removal capacity during minimum flow periods.  A chemical process
 can be turned on and  off at will;  the plant that experiences seasonal
 nitrification would be  greatly benefited by having  a process that it could
 start and stop on demand at the appropriate time of the year.  Thus, flexi-
 bility and inherent reliability could make the reduction-deammination
 sequence more  attractive than  just simple economics might suggest.

      While many reduction  and  deammination reactions are known and have
 been extensively  studied,  the  conditions for such studies have been almost
 entirely restricted to  relatively concentrated solutions.   Dilute solution
 reaction chemistry of any  type, let alone dilute solution reaction chemistry
 of  specifically nitrate  reduction and deammination, is an area of study
which has been faintly touched at best.  The objective of the present
program was to select groups of candidate reduction and deammination agents
that might possibly be used in water reclamation processes,  and to  test the
feasibility of developing a denitrification process using  these agents  when
the nitrate ion was present at dilutions of 10 to 50 ppm of  NO"-N.

-------
                         SUMMARY AMD CONCLUSIONS


      The objective of this program was  to  demonstrate the  feasibility of
 denitrification by a chemical process.  Although  departing somewhat from
 the original reduction-deammination concept,  such a demonstration was
 achieved in the course of the work.

      After an extensive literature survey,  eight  reducing  agents and
 three deammination agents were  selected for testing the feasibility of
 denitrification by a reduction-deammination process.  The  original plan
 was to study various pairs of these  agents  in a statistically designed
 matrix of experiments,  varying  numerous environmental factors such as pH,
 temperature, etc.   However,  it  soon  became  evident that unknown second-
 order interactions in the design could  defeat the purposes  of such a study.
 Accordingly, these eight reducing  agents were screened under anaerobic
 conditions at high NO^-N concentrations (50 ppm)  and high  temperature (85° F)
 with the agent in  excess, varying  only  pH and catalyst.  On this basis,
 ferrous ion (Fe++),  iron powder, and hydrazine  or its salts (N2Hij, ^1^30^)
 showed appreciable reducing  power.   A very  small  amount of  reduction was
 accomplished with  glucose.   Formaldehyde, carbon,  sulfur dioxide, and
 carbon monoxide were inactive.

      The deammination agents studied were sulfamic acid (HSC^NHg) and urea.
 Glycine was slated for  study, but  later abandoned.  No nitrogen loss was
 detected that could  be  attributable  to  the  deammination agents in any of
 the tests where reducing and deammination agents were studied together.
 Separate studies showed that urea was ineffective, and that sulfamate
 could only deamminate under  acid conditions.

      Of the three  reducing agents that  passed the screening, only ferrous
 ion was economically attractive; it  is  available as crude copperas,
FeSOli'TH^O,  for only four dollars per ton.  From the technical viewpoint,
 ferrous ion was the  only choice, since  the  test series revealed that up
 to  55% denitrification was occurring with ferrous ion alone, apparently
by  direct  reduction  of NO" to either N2 or NgO  ("missing N").

     Accordingly,  the bulk of the experimental program was devoted to
 studies  of the  direct denitrification reaction with ferrous ion.   Ferrous
 sulfate was used as the  source of ferrous ion, but the solution must be
initially  basic, so that the system is actually heterogeneous, with pre-
cipitated  ferrous hydroxide being the reducing agent.   As the  reaction
progresses, the pH drops and the black ferrous-ferric  complex  is formed
 (Te^pii or  the corresponding hydroxide).   Assuming that the "missing N"
evolves as Ng, then the reduction is a five-electron reaction, and since
additional ferrous ion is consumed in forming the complex, the theoretical
requirement for complete reduction would be 7.5  moles  of Fe"*"1"  per mole
of NO"

-------
      At the 10 ppnHDvJI level, using an 8:1 molar Fe++:N<>; ratio, reduction
 (to mixtures of N02, Ah, and "missing N") ranged from 50)Tto looi and
 denitrification ranged from 1(# to UjJ.  These results vere obtained over
 an initial pH range of 7 to 11 using either lime or sodium hydroxide for
 pH adjustment.  Neither reduction nor denitrification was observed under
 acid conditions.  A trend for the amount of reduction to increase with
 increasing pH was noted, but the amount of denitrification did not seem to
 follow a trend.

      Catalysis by either cupric ion (Cu++) or silver ion (Ag*) in 1-5 ppm
 concentration is necessary for reaction to occur, but denitrification was
 observed only when cupric ion was used as a catalyst.  Addition of phosphate
 ion to the solutions interfered with catalysis by cupric ion,  and no reduc-
 tion occurred, even when the catalyst level was increased.   However  at
 increased levels of iron (16:1 and 2k:l), or with silver as the catalyst,
 reduction occurred readily.   As before,  no denitrification occurred when
 Ag  was the catalyst.

      Carbonate ion did not interfere with the reduction reaction,  but,
 again,  no denitrification occurred when C0| was present.  Carbonate buf-
 fered the solutions; this fact,  coupled with other results  where pH was
 held relatively constant by addition of base during  the course of  a run
 indicated that denitrification would not occur unless the pH was dropping
 as the  reactions progressed  - that is, buffering  prevented  denitrification.

      Based on  these results,  development of a denitrification  process based
 on direct  chemical reduction of  nitrate  ion to nitrogen or N2© appears
 feasible.   The denitrification can be carried out with an inexpensive
 reducing agent (copperas) and is worthy  of  further technical evaluation.
 However, much more will have  to  be  known about the basic chemistry of
 this  reaction before a definitive process emerges.

      The effect of pH and buffering on the  reaction needs to be resolved
 since there is no  evident explanation for the  absence of denitrification
 under constant pH  conditions.  Quite possibly, observation of denitrifica-
 tion  has been obscured by the effect  of  some unknown variable.   The fact
 that  the amount of denitrification does not correlate with pH level, and
 that  there has been an occasional failure to denitrify even in the absence
 of buffering, indicates that such a variable may indeed exist.   The hetero-
 geneity of the system could be the source of this inexplicable  behavior.
Minor variations in the mode of precipitation of the ferrous hydroxide,'
 the manner of absorption of catalyst ions and the degree to which this'
precipitate absorbs them, the manner in which the ferrous-ferric complex
 forms, and a number of other factors relative to the solid phase could all
play a role, and consistent denitrification and reduction might wen prove
to be a function of consistent precipitation technique.

-------
      Of  great  importance, too, is a determination of the identity of "missing
N",  and  an investigation of the intermediates that lead from HOo to "missing
K".   Obtaining total denitrification may veil depend on steps in the mech-
anism that are not evident in the present vork.  For instance, if N2 is the
gas  evolved, the denitrification could actually be the result of the oxida-
tion of  N^^h or NHgOH formed at some intermediate stage.  The possibility
of loss  of nitrogen in  some form as a part of the precipitated iron insol-
ubles should also be investigated, even though such a mode of denitrifica-
tion seems unlikely.

      Further studies are also required on catalysis of the reaction.  Cupric
and  silver ions are the classical catalysts for homogeneous reductions, yet
ferrous  ion reduction/denitrification reaction is apparently heterogeneous.
Cupric and silver ions are definitely different in their response to phos-
phate, and apparently different in their ability to lead to denitrification,
although this  last difference could again be confounded vith some other
variable.

      With  the  information now at hand, there are three possible approaches
to development  of a ferrous ion denitrification process as described
below.

      The first  approach, total denitrification with ferrous ion, assumes
that  further studies of the reaction will result in the data necessary to
make  reduction  of HO^ to "missing N" by ferrous ion both consistent and
close to Quantitative.

      In the second approach, a reduction-deammination sequence, reduction
and some denitrification is carried out by the ferrous ion, and the pH
falls from some_initially alkaline value to belov 7 as the reaction pro-
gresses.   If HOg vere the major reduction product other than WgO or Ng,
then deammination might still be used to complete the denitrification.
However, carbonate ion would have to be absent for this sequence to take
place, which would severely limit the applicability of this process.

     The third possible process, coupling ferrous ion reduction with
ammonia stripping, might be the easiest to develop.   In some of the
reactions  conducted during this study, 33-^5$ denitrification occurred,
and another 3^6$ of the nitrate was converted to ammonia.  Thus,  a
sequential ferrous reduction-ammonia stripping process has demonstrated
potential  of up to 90$ total denitrification.

-------
                            LITERATURE SURVEY


      Approximately one-third of the effort  on this program was devoted to a
 comprehensive literature search, covering the period from early 1968 wen
 back into the nineteenth century.   Recent references were obtained from
 Keywords (Chemical Abstracts),  Chemical Titles. Current Contents. Water
 Pollution Abstracts,  and on-the-shelf journals and reports.  For thTperiod
 back through 1907, Chemical Abstracts and the annual literature reviews in
 tne Journal Water Pollution Control Federation were prime sources.  Qnelins
 Handbuch der anorganischen  Chemie  revealed  references back to the early -
 1800's.   A patent search was conducted by the North American Rockwell
 Patent Department, and several  pertinent  current references were provided
 by the FWPCA Project  Officer, Dr.  R. B. Dean.

      While the primary objective of the search was selection of suitable
 reduction and deammination  agents  for subsequent laboratory testing, much
 related  information was collected  on water  reclamation in general, and
 nitrified waters  in particular.  Between  kOQ  and 500 references were col-
 lected in original or abstracted form.

      For this report,  slightly more than  200  references to the most
 pertinent and informative articles  are provided.

      The "Discussion and Results"  section, which follows, contains ref-
 erences  to the literature sources which,  for the convenience of the reader
 are  grouped as follows  in the section entitled "References":              '

A.    Nitrates  in Water
              General Chemistry
     5-30     Occurrence and Effects
     31-^8    Natural and Biological Denitrification
     U9-68    Elimination Methods

B.   Water Reclamation

     1-12     Conventional and Tertiary Treatment
     13-25    Iron Salts in Water Treatment
     26-30    Carbon in Water Treatment

C.   Reducing Agents for Nitrate

     1-23     Ferrous Salts
     2lf-28    Carbon
     29-33    Sulfur Dioxide
     35-^7    Formaldehyde
     U3-56    Sugars
     57-61    Powdered Iron
     62-69    Hydrazine and its Salts
     70-85    Miscellaneous

-------
D.   Deaiamination Agents

     1-lU     Sulfainic Acid
     15-21    Urea
     22-21*    Amino Acids
     25-3^    Miscellaneous

E.   Catalysis

     1-6      (Wot subdivided)

F.   Analytical Methods

     1-9      (Not subdivided)

-------
                         DISCUSSION AND RESULTS


 BACKGROUND


 Nitrates in Water-Occurrence and Effects (See References A-5 through A-30)

      Nitrates occur in natural and reclaimed waters in amounts ranging from
 a fraction of a part_per million to several hundred ppm, calculated as
 nitrate-nitrogen (NO^-N).  The sources may be natural, such as leaching
 from nitrate deposits (e.g., guano in limestone cave areas), the  natural
 decay and oxidation of nitrogenous organic matter (protein) as carried out
 by certain microorganisms,  and the fixation of atmospheric  nitrogen as NO
 and NOg from electrical discharges during thunderstorms. The sources may
 also be man-made,  such as leaching from agricultural lands  treated with
 nitrogenous fertilizers, effluent from fertilizer manufacturing plants,
 and effluent from  other chemical and manufacturing  processes that employ
 nitrates in one form or another.

      In water reclamation,  nitrate ion may find its way into the  stream
 initially from any of the above sources.   However,  regardless of  the  initial
 water quality,  secondary treatments based on biological oxidation of  organic
 matter (i.e.,  activated sludge and trickling filter processes)  can them-
 selves introduce additional nitrate.   Hence,  almost all secondary effluent
 contains an appreciable quantity  of nitrate ion.

      There are  two distinct problems associated with nitrates  in water:
 methemoglobinemia  and algae bloom.   Methemoglobinemia is a  serious  and
 often fatal disease of the  new-born,  characterized  by cyanosis, for which
 nitrates have been designated  a causative  factor.  A drinking water standard
 of  10 ppm N03-N has been established by the Public Health Service  (Ref. A-30)
 as  a  preventive measure for this  disease.

      Both ammonium and nitrate  ions,  as well as phosphates, are excellent
 nutrients for plants,  including algae  (Ref. A-lU).  Algae bloom is one
 aspect of a natural process  called  eutrophication, wherein a body of water
 such  as  a lake  is  gradually converted into a swamp and eventually a meadow.
 The beginning and  end  of  this process  are not particularly unpleasant;
 however,  in the intermediate stages, algae and aquatic plants grow in
 abundance as nutrients build up, this  abundant organic matter decays
 depleting oxygen and killing the aquatic fauna.  Unfortunately, man-made
 sources of nutrients accelerate this process, and induce the algae bloom
 stage  in receiving waters where it would not otherwise occur.  Eutrophica-
 tion would be intolerable in a reservoir to be used for reclaimed water for
drinking purposes.   Other water uses,  such as recreation, incorporate a
certain esthetic value which is certainly not enhanced by overgrowths of
algae.

-------
Natural Denitrification  (See References A-31 through A-1*8)

     Countering the natural and man-made nitrification processes are natural
denitrification processes.  The conditions for nitrification and denitrifi-
cation are very similar, with the main difference being that denitrification
occurs only when the water is close to being depleted of oxygen.  Under such
conditions, certain microorganisms will continue the oxidative decradation
of organic matter, using the nitrate ion as the oxygen source, and eliminating
the nitrogen as N2 gas.  Denitrification can occur in water, in soils, and
in accumulated masses of organic matter such as silage (Ref. A-37, A-^U).
Good reviews of natural denitrification processes are presented by the Thames
Survey Committee (Ref. A-l*7) and by Camp (Ref. A-3U).  The latter author,
however, implies that natural denitrification can be partly chemical, and
uses the deaznmination reaction of nitrite ion with urea as an example,
showing the reaction to be thermodynamically favorable.  Chemical reduction
by ferrous salts is also believed to play a role in natural denitrification
in certain soils (Ref. A-31, C-22) and acid tropical waters (Ref. A-Uu).
However, it is generally agreed that most natural denitrification is bio-
logical rather than chemical (Ref. A-33).

     In the vater reclamation field, natural denitrification first came to
attention as a problem (Ref. A-32, A-3b through A-l*2).  The nitrogen evolved,
when natural denitrification occurs in sedimentation basins, causes the
phenomenon known as "rising sludge" or "rising hunus".  That is, bubbles
of nitrogen entrapped in the sludge carry it to the surface rather than
allowing it to settle.  In recent times, natural denitrification has been
turned to good use, and forms the basis for several advanced waste treatment;
processes (Ref. A-36, A-^3, A-l*6, A-55).  These processes are reviewed below.

Nitrate Removal Processes (See References A-l*9 through A-68)

     There have been a number of methods studied for producing denitrified
water.  These include: control of secondary treatment processes in such a
manner that nitrification is minimised; ion exchange; extraction; and
biological denitrification.  Adsorption of nitrate on such substrates as
carbon, alumina, and silica gel has also been noted, but does not seem to
have been studied for the specific purpose of developing a removal process.
An occasional excursion into reaction chemistry has been made (Ref. C-23,
C-58) using ferrous salts and iron powder, but with little or no success.
However, concentrates (primarily radioactive wastes) have been successfully
treated with formaldehyde and sugars as reducing agents, as discussed later
in this section.

     Controlling the secondary treatment processes can be effective, but
suppression of nitrification is often accompanied by some undesirable
effect, such as decreased BOD removal.   (BOD, Biochemical Oxygen Demand,
a measure of the biodegradable organic  content of the water.)   Ion exchange
is very efficient, but costs can be high.  Extraction works very nicely on

-------
 concentrates, but is inefficient for dilute nitrate solutions.  Apparently
 none of these processes are under serious consideration at the present time
 for extensive incorporation into reclamation plants.  A more detailed pic-
 ture of the state-of-the-art of these processes, as well as processes
 designed to remove ammonia, may be found in the recent publication of
 Farrell, Stern and Dean (Ref. A-55).

      Biological denitrification is currently under active study (Ref. A-U6)
 There are two modifications of this process.  The first (Ref. A-36) involves
 a sequence which first nitrifies the water to the greatest extent possible
 then carries the water into an anaerobic chamber where sludge from a previous
 step is used to supply organic food to the denitrifying microorganisms.   The
 second modification has grown from the unexpected denitrification observed
 when activated carbon columns were being studied for tertiary treatment
 (Ref.  A-i*3).  In this case, the denitrifying organisms had established
 themselves in the carbon columns.   Supplying methanol as additional food
 for the bacteria increases the efficiency of this process, and sand has
 been successfully substituted for  the carbon.


 SURVEY OF  CANDIDATE REDUCING AMD DEAMMENATION AGENTS


 Chemicals  in Water Reclamation (See  References  B-l through B-30)

     In choosing agents for the reduction-deammination study, attempts
 were made, wherever possible,  to project the tentative process in terms of
 other  current and future processes to  see where the new process might be
 conveniently incorporated.

     In general,  it appears that flocculation and  coagulation will play
 an  increasing role  in the future, using  either  lime or alum.  (See Ref  B-l
 for a tabulation  of well-developed tertiary treatments.)  Lime flocculation
 can reduce phosphate ion concentration,  so it is reasonable to expect the
 waters  to be highly basic during some stage of reclamation processes aimed
 at  controlling nutrient  content.  The lime is often used with another salt
 as  an additional  coagulant.  Ferrous and ferric salts  (Ref. B-13 through
 B-25) have been used in this manner.  Finally, activated carbon may be
 used in the final stages to remove the last traces of organic matter
 (Ref. B-26 through B-30).  Both ferrous salts and carbon are potentially
 reducing agents for N03 ion, and are discussed further below.

Choice of Reducing Agents

     Eight reducing agents were selected for study as a result of the
literature survey.  The information amassed on each of these is summarized
below.
                                   10

-------
      Ferrous Ion (Ref. C-l through C-23).  Ferrous ion, as either ferrous
 sulfate or ferrous hydroxide, has formed the basis for many analytical
 determination of nitrate ion,Converting the N0§ to NH3.  The first step,
 conversion to nitrite ion (N02), is slow but may be catalyzed by silver or
 cupric ions.  Subsequent reductions proceed to ammonia by attacking N02 and
 NO from the HN02/N02/NO/H20 equilibrium.  A summary of almost an possible
 reaction sequences is given in Ref. C-22.  The reduction is most effective
 under faintly or strongly basic solutions, but can also occur under acid
 conditions.  It has been stated that the concentration of ferrous ion must
 be at least 70 ppm in order to be effective (Ref.  C-1+).

      Ferrous ion was once studied for use in water reclamation,  and was
 shown to be capable of 90$ conversion of NO? to NHo at the 100 ppm NOo-N
 level (Ref. C-23).   This work was abandoned because of the "ferrous and
 ferric hydroxide sludges" that formed.   From an operational viewpoint,  this
 is difficult to understand,  because the  "sludges"  are  heavy,  settle easily,
 and so should be easy to separate from the treated water.   However,  ultimate
 disposal of the "sludges" may pose some  problems (Ref.  B-13),  and will  have
 to be given serious consideration in the development of any process based
 on ferrous  ion.

      As noted earlier,  some  of the natural denitrification and reduction in
 water and soils has been attributed to the presence of ferrous salts (See
 page  9).

      Carbon (Ref. C-2U  through C-20).  Nitrate  can be  reduced  to  nitrite by
 carbon,  with  C02 being  the other  product.   The well-known  "wet-ashing"
 technique for removal of carbonaceous materials  in analytical procedures
 is a  good example.   Coal, carbon  black,  and graphite reduce nitrate.  How-
 ever,  high  concentration and heating are generally required, under which
 conditions  mixtures  of  nitrogen oxides are  evolved.

      Sulfur Dioxide  (References C-29 through C-33).  Sulfur dioxide, if
 shown to  be suitable as  a reducing agent, could play a dual role in water
 reclamation because of  its bacteriostatic properties.  Sulfur dioxide can
 reduce nitrate to various products, including hydroxylamine (Ref. C-33).
 The most  pertinent article (Ref. C-31) claimed that a 10$ NH3 solution
 saturated with S02 until faintly ammoniacal would reduce oxidizing anions
 and completely eliminate nitrite ion as N2.

     Formaldehyde (References C-35 through C-U?).  Formaldehyde has been
 extensively studied as a reducing agent for nitrate concentrates ("Purex
wastes") by the Atomic Energy Commission.  Under these conditions, NO 'and
N02 are evolved, which implies that HNOg  would be the product in dilute
aqueous solution.  There is an induction  period in the reaction which can
be overcome by ferric ion catalysis (Ref. C-U2, C-^3).
                                   11

-------
     One of the few articles dealing with dilute solution reduction of
nitrate ion has some interesting information on formaldehyde (Ref. C-36).
In the photochemical reduction of nitrate to nitrite, with the results
                              V
expressed as a ratio  (R = — ?), it is shown that the R value is increased
                           NOo

by a factor of 3 to k by the addition of formaldehyde.  Working with a
0.05$ KN(>3 solution, the value for R was 60 after 20 minutes of ultraviolet
radiation in the presence of formaldehyde, representing a change in NO§-N
concentration of about 70 ppm down to 1 to 2 ppm.  The data did not explicitly
show, however, that part of the reduction was by direct reaction with for-
maldehyde rather than being totally photochemical in nature.

     Sugars (References C-U8 through C-56) .  Sugars have also been used
by the Atomic Energy Commission to reduce nitrate concentrates.  Sugars
also increase the R value in the photochemical reduction of NO? (see above).
A sugar need not be one of the classical reducing sugars to react readily
with nitrate: sucrose as well as glucose is suitable.

     Powdered Iron (References C-57 through C-6l) .  Powdered iron is similar
to ferrous ion in the manner in which it reduces nitrate.  It was once
examined as a reducing agent for eliminating nitrate ion from drinking water,
but was classed as ineffective because it did not convert NOg directly to N2
(Ref. C-58).  The efficacy of powdered iron as a reducing agent depends to
some extent on its method of manufacture (Ref. C-59).

     Hydrazine and its Salts (References C-62 through C-69).  Hydrazine is
an excellent reducing agent for nitrates, and forms the basis for the Auto
Analyzer now used for analysis of various waters (Ref. C-63, C-6?).  The
reader should note that this reduction reaction, if it is to be of value
in the proposed process, must lead to near-quantitative oxidation of the
hydrazine to Ng.  If this does not occur, then the reaction may introduce
more inorganic nitrogen compounds into the water than are removed by the
overall process.

     Carbon Monoxide.  Carbon monoxide was selected as an agent purely on
the basis of economy and the desirability of having COg as a byproduct.
No references to CO/N03 reactions were found; on the contrary, a very
early reference (1851) clearly states that carbon monoxide and nitric
acid do not react (Ref. C-83).

     Miscellaneous Reducing Agents (References C-70 through C-85).   Other
reducing agents noted during the course of the literature survey were
generally metals and various lower- valent salts of metals (e.g., Al, Ti**,
Cu-Cd, Sn"*"1", etc.).  These were eliminated from consideration on the
basis of undesirable byproducts and/or expense.


                                   12

-------
Choice of Deammination Agents

     Three deanmination agents were selected, based on the information
presented below.

     Sulfamic Acid (References D-l through D-lU).  Based on an article com-
paring sulfamic acid and urea (Ref. D-l^), sulfamic acid was the prime
choice for deEmmination.  This article showed quantitative reaction of the
theoretical amount of sulfamic acid and nitrite ion in two minutes even at
the 2 ppm NOjj-N level, whereas even excess urea could not effect complete
deammination at that nitrite concentration after one hour.  The system
sulfamic acid-nitrite ion has been studied extensively in analytical appli-
cations, in terns of the intermediates formed, and with respect to thermo-
dynamics.  A means for removing nitrite ion from boiler water using sodium
sulfamate has been patented (Ref. D-13).

     It should be noted that sulfamic acid can react with nitrate ion as
well as nitrite, yielding ^0.  However, this reaction is very slow below
60° C (Ref. C-7).

     Urea (References D-15 through D-21).  As noted above, urea is not
nearly as effective as sulfamic acid for deammination, although its use
has been patented for treating waters to remove nitrite ion (Ref. D-15,
D-21).

     Amino Acids (References D-22 through D-2U).  Since amino acids are
present in water during various stages of reclamation processes, they might
provide an in situ source of deammination agent.  Glycine was chosen as the
model originally, since it reacts readily with nitrate, although side
reactions give some C02 and N2<> as well as nitrogen (Ref. D-22).  Compara-
tive studies (Ref. D-23) have shown that relative to alanine (1.00), most
amino acids, glycine included, are deamminated at about the same rate
(0.70 to 1.50), although there are extremes such as cystine (3-12) and
isovaline (
     Miscellaneous Deammination Agents (References D-25 through D-3^).  In
general, all primary organic amines and amides are susceptible to deammina-
tion by NOjj.  Azides also react with nitrite in a similar fashion.  How-
ever, none of these offer any particular advantage over the agents dis-
cussed earlier, and, of course, the organic amines will leave an organic
residue, which is not desirable.

     Ammonia and its salts generally require heating in order to react with
nitrites at appreciable rates.  However, it was tentatively planned to test
ammonia in combination with 502, in view of the data presented earlier (p. 11)
on this particular combination.
                                   13

-------
 Catalysis in Reduction and Deanmination (See References  E-l through E-6)

      Catalysis plays a role in many of the reactions mentioned earlier,
 i.e., iron and ferrous ion reductions are catalyzed by cupric  or  silver
 ions, hydrazine reductions by cupric ion, formaldehyde reductions by ferric
 ion.   Catalysis has a role in certain biological nitrate reductions as
 well  (Ref. E-3).

      Cupric ion and silver ion are  the classic homogeneous  reduction
 catalysts, particularly for homogeneous reductions  with  hydrogen  (Ref. E-U).
 Of interest to the present program  is the fact that catalytic activity of
 these ions may be enhanced more than a hundredfold  in  the presence of
 organic  acids.  Magnesium, cadmium, and zinc ions are  also  reported to be
 active catalysts in the presence of organic  acids.

      An  article published in 1923 (Ref.  E-5)  reports that nitrate ion
 reduction is an autocatalytic process,  and that in  the absence of some
 small initial concentration of nitrite ion,  nitrate cannot be reduced by
 ferrous  ion, formaldehyde, mercurous ion,  or a number  of other agents.
 Nitrite  ion was removed from the test solutions with urea and amino acids.
 If this  phenomena were indeed confirmed,  it would mean that reduction and
 deanmination would have to be sequential operations.   However, the conclusion
 is refuted to some extent by other  work;  the  simultaneous reduction-deammina-
 tion  cited earlier for m^/SO^ solutions is one example.
DILUTE SOLUTION REACTION STUDIES
Screening of Reducing Agents

     The initial laboratory plan in this program called for a series of
statistical test matrices, studying reducing-deammination pairs, with high
and low levels of the variables as given in Table 1.

     Each test solution was to be sampled after reaction times of one hour
and 2h hours.  Each sample was to be analyzed for nitrate, and if reduction
had occurred, for nitrite and ammonia to determine the total nitrogen
balance.

     In selecting the levels of variables, it was assumed that the denitri-
fication treatment would be applied to secondary effluent, and the values
for pH and temperature are believed to be the extremes.  A different pH
range was chosen for ferrous ion, Fe^, since in this case it was assumed
that ferrous treatment would be coupled with lime coagulation, and the
denitrification would thus be carried out in alkaline waters.

     The initial tests within the first matrix, if accepted at face value
would have led to the strange conclusion that ferrous ion could not even '

-------
                                TABLE 1
               REDUCTION-DEAMMINATION REACTION VARIABLES
     Variable
      High Level
     Low Level
NO"-W Concentration

Temperature

PH  ^
 (Fe   excepted)
 (with Fe**)

Reducing Agent
 Concentration

Deanmination
 Agent Concentration

Order of Addition
Atmosphere


Catalyst
        50 ppm

        Ol-O T,
         9
        11
    Threefold excess
    Threefold excess

  Deammination agent added
one hr after reducing agent

 Nitrogen (water deaerated)
    1 to 5 ppia of
    Cu++ or
     10 ppm
      6
      8
Stoichionetric
Stoichiometric

 Agents added
simultaneously

 Air (water
 untreated)

 No catalyst
                                   15

-------
 reduce nitrate ion, let alone participate in a electrification sequence.
 In view of the large amount of literature that had been uncovered on the
 reducing power of ferrous salts, this result appeared to be anomalous.

      In statistical terms, the apparent anomaly arose because the results
 were confounded by a second-order interaction of variables.  The matrix
 design was such that air oxidation of ferrous ion was predominant in cat-
 alyzed tests, and the rate of the uncatalyzed reactions was essentially nil
 at the low concentration levels.

      Since much less information was available on the other seven reducing
 agents, it was decided that statistically designed matrices with a large
 number of variables could not be conducted with any degree  of assurance
 that other unrecognized interactions would not lead to fallacious conclu-
 sions.  Accordingly, a new test series was devised to screen the'reducing
 agents with most of the variables fixed at the high levels.

      Only the effects of pH,  catalyst, and deaamination were studied in
 this new series, with the deammination agent added either initially after
 2k hours, or not at all.  However,  the deammination agents  had no effect
 on these reactions, as discussed later in this section.   The results of
 the new tests,  which amounted to screening the candidate agents  for reducing
 power under the most favorable conditions possible,  are  given in Table  2.

      On the basis of these tests, only ferrous ion (Pe**) appears to be
 a potentially useful agent.   Hydrazine and its salts,  since  they did not
 achieve quantitative reduction,  introduced more inorganic nitrogen than
 they could remove.   Iron powder is  considerably more  expensive than Fe"*"*"
 ion.   The latter is available as crude copperas, at a cost of a  few dollars
 (less than $5.00) per ton, while the most optimistic  estimates for  iron
 powder are in the area of 10  cents  per pound.

 Deammination Agents

      Deammination agents (urea,  or  sulfamic acid initially neutralized to
 avoid additional pH adjustments) were present  in many of the tests cited
 above,  but once  Fe    ion had  emerged as the prime reducing candidate, the
 deammination phase  of this effort was  stopped.  The Fe++ reduction must be
 carried out under initially basic conditions, and sulfamate deamminates
 effectively only under strongly  acid conditions.  This conclusion was
 reached during a separate series of studies carried out as part of the
 supporting analytical  effort  (See p.35).  The poor performance of urea,  as
 cited in the literature  (Ref. D-lU), was also confirmed.  At the 3 ppm NO|-N
 level,  no  deammination could be detected after U8 hours at either pH U.O or
 10.6, even though the urea was present in threefold excess.

     Glycine was slated to be studied as a deammination agent, but was by-
passed as the investigation of ferrous ion proceeded.

                                   16

-------
                             TABLE 2

                    REDUCING AGENT SCREENING
Reducing Agent
                           Conditions Varied
                         (and number of tests)
                                                             Reduction
                                                            Observed,  %
so2


Carbon
CO


Glucose
Fe Powder


Fe   Ion
                        Fe   ,  Cu   catalysts;
                        pH 6 and 9 (k tests)
                        _      _      .  .
                        Fe   ,  Cu   catalysts;
                        pH 6 (U tests)
                        Fe' '  , Cu   catalysts;
                        pH 6 (U tests)

                        Cu   catalyst;
                        pH 6 (2 tests)
                        _ +++  _ ++  ,,+5   .  ,   .
                        Fe   , Cu  , V   catalysts;
                        pH U and 6 (6 tests)
                        Fe   ,  Cu   catalysts;
                        pH 11 (U tests)
                          j, j^
                        Cu   catalyst, pH 11
                        (U tests)
                          I i^
                        Cu   catalyst,
                        pH 6 (3 tests)

                        *Cu   catalyst,
                        pH 11 (6 tests)
                                                            None
                                                            None
                                                            None
                                                            None
                                                            3 to 6
                                                          (with V+5 only)
                                                            10 to UO
                                                          (none with Fe

                                                            10-20


                                                            15-^5


                                                            15-^0
                                                                       +++
Fixed Conditions:

     Temperature:
     Test Time:

     Atmosphere:
     NOl-N Concentration
     Reducing Agent Concentration:
                                        85° F
                                        U8 hr.
                                        Nitrogen

                                        50 ppm
                                        3X (Denotes 3 moles of reducing
                                        agent for each mole of NOZ)
* Tests with Fe   slightly more complex.   Discussed in more detail in the text.
                                   17

-------
 FERROUS ION AS A DENITRIFICATION AGENT


 Initial Screening Studies

      As noted in Table 2, the initial screening tests with ferrous ion were
 slightly more complex than those with the other reducing agents.   The initial
 matrix+tests had indicated that bubbling air through the solutions oxidized
 the Fe   before it could reduce NOo in cases where catalyst was present,
 and that uncatalyzed reactions were not possible at high dilutions.   However
 there was still the possibility that there was an interference from the
 deanmlnation agent, either direct or by suppression of the autocatalytic
 influence of H02 ions, as had been indicated in the literature (Ref  E-5
 See p. 1U for discussion).  Accordingly,  the first screening  tests with Fe*+
 incorporated the sulfamate addition time  and presence or absence of  N05 ion
 as variables.  Test conditions and results are given in Table 3.

      The above data indicate that nitrite does not catalyze the reduction
 and that sulfamate does not interfere with the reduction.   If anything
 sulfamate aids reduction;  more reduction  occurred in Tests 3  and U where
 sulfamate was added initially.   Of course,  under the basic conditions,
 sulfamate did not deamminate;  decreases in nitrite ion occurred only as
 the N02  was  oxidized back  to nitrate.   This was quite evident  in the  blank
 and in some  of the other tests.   Oxidation occurred because of traces of
 oxygen in the house nitrogen used to blanket the  reactions, or possibly
 because  air  was introduced when  the  2k  hour samples were taken.

      The really significant result,  however,  is the total  nitrogen balance
 that resulted at  the end of Test  3-   Of 1*9  ppm of nitrogen initially present,
 only 39-2 ppm could be  accounted  for after  US hours.  These data lead to the
 conclusion that Fe++ can reduce nitrate ion directly to either Ng or N20
which can escape  from the  reactor.   Since the actual constitution of the*
 escaping material has yet  to be determined,  it is generally referred to in
 this  report  as "missing N".

     The precipitated iron insolubles in these reactions were black, which
 indicated that the iron end product was not ferric hydroxide,  but the mixed
Fe   -2Fe     salt; that is, Fe30^, or the corresponding mixed hydroxide,
although the existence of the latter does not seem to have to be definitely
established.  Thus, a balanced equation for the reduction of nitrate to
nitrite  should be written as in Equation (3).

          HgO + 3Fe*+ + NO" •* Fe"H'.2Fe*++ + KO" + 2(OH~)             ('3)

     The denitrification reactions may be as those shown by Equations (U)  and  (5).

          ^e^ + 2KO- + jHgO t MFe^e^) + NgO + lO(OlT)      (k)

          15F6++ + 2NO- + oHgO - 5(Fe++.2Fe+++) + N,, + 12(OlT)       (5)


                                   18

-------
           TABLE  j
FERROUS ION REDUCTIOIi OF NO.
Test
No.
1
2
3
k
5
Conditions Varied
Sulfamate added after
2k hrs. No NO".
Sulfamate added after
2k hrs. NO" added.
Sulfamate added
initially. No NO".
Sulfamate added
initially. NOg added.
Blank (No Fe++) .
NO' added.
Tine, hr.
0
2k
0
2k
0
2k
kQ
0
2k
kQ
0
2k
kQ
Analytical Results, ppm
NO"-N
Uii.3
30.5
33.9
&'
1*9.0
30.2
33.0
20J5
3U.5
50.0
50.9
5^.7
NO--N
—
5.0
6~7
k.k
9-7
U-3
k.2
3-9
0.8
3
i
--
i.e
--
•~
Fixed Conditions:
Temperature: t>5° F
Atmosphere : Nitrogen
pll: 11
NO^-N Concentration: 50 ppm
Fe++ Concentration: 3X < ,_ _x
; (Denotes 3 moles of agent per mole of NO,)
Sulfamate Concentration: 3X ) •*
Catalyst: Cu , 5 ppm
                19

-------
      Finally, some of the nitrate is reduced to MH3, as in Equation (6).

           12Fe+* + NO' + ohVjO  ->  UCFe^-ZPe***) + NH3 + 9 (OH~)         (6)

      These latter reactions (Eq. Jf, 5, 6) have a much greater iron demand than
 the simple 2 electron N03/N02 reduction.  In fact, if it is assumed that the
 pissing N  in Test 3 evolved as N2, then the reduced products (NOg, N2, and
 NH3) account for about 8o> of the ferrous ion that was originally present
 That is, reduction of N03 by Fe~ is a fairly efficient process, b£in
 this first_screening series there simply was not enough of Fe++ to reduce
 an the N03 via the U, 5, and 8 electron reactions shown.

  4^ J^rification was confirmed in additional screening tests, carried out
 with the Fe   ion concentration increased to eight times that of the NO? ion
 on a molar basis (designated as "8X" concentration).   This would be slightly
 more than enough ferrous ion to convert all the N0§ to N2, assuming N2 to be
 the denitrification product.  In this series,  pH and catalyst level were
 varied.  No deammination agent was present in this or any  subsequent studies.

      The data in Table k clearly show that as  much as 50-55$  denitrification
 can occur as a result of direct Fe++ reduction of N03.   The data also  show
 that denitrification does not occur under initially acid conditions   The
 effects of catalyst level and pH are not clearly separated, however, and no
 ready explanation can be given for the  reversal of the  NO^-N  and NOo-N
 levels shown in the "duplicate" 3a and  3b tests.

 Attempted Identification of  "Missing N"

 +v.   ,,P?11?wing,,tJle  above  screening tests,  an attempt was made  to  identify
 the  "missing N"  by  carrying  out the  reaction in a  closed,  evacuated vessel
 using  carefully  degassed  solutions,  and examining  the evolved  gases mass   '
 spectrometrically.  This  experiment  failed because no denitrification
 occurred,  for reasons unknown.  The  conditions  (except for the atmosphere)
 were the  same as those shown for tests  3a  and 3b in Table U.   However,  the
 mass spectrometer detected only trace amounts of nitrogen and  argon (residual
 air) after trapping out the water vapor, and analyses of the solution showed
 N03-N,  1.1 ppm;  N02-N, 21.3 ppm; and NH3-N, 25.8 ppm.  (U8.2 ppm total for
 a nominal 50 ppm N03-N initial  concentration.)

     Identification of "missing N" was planned for some later  stage  once
 the reaction was more fully understood.  However, other aspects of the
 reaction took on a greater priority, and these plans were never carried
 out.  Nonetheless, these results are of importance: they show that occa-
 sionally, for reasons as yet unknown, the denitrification reaction fails
 completely.  Accordingly, results discussed subsequently in this report must
be interpreted with caution since the occasional failure from unknown causes
can be a confounding factor.
                                   20

-------
          TABLE
DENITRIFICATION WITH FERROUS ION



        (Duplicate Tests)
Test No.
1. a
b
2. a
b
3. a
b
pH
k.O
i*.o
7.1
7.1
11.0
11.0
Cu++
ppm
5
5
10
10
5
5
Time, hr.
0
1*8
0
ua
0
2U
W
0
2U
ua
0
21*
48
0
2U
hQ
Nitrogen Balance, ppm
NO"-N
50
50
50
50
50
25.8
6.1
5U
35.2
6.1
50
3.0
2.6
50
15.3
13-3
NOg-N
--
—
0.5
0
16.0
lit.fc
4~3
^.3
NH.-N
--
__
16~5
19
19-5
19.1
"Missing N"
—
—
26.2
2U.9
n.5
11.3
Fixed Conditions:
Temperature: 35° F
Atmosphere: Nitrogen
NO~-N Concentration: 50 ppm
Fe++ Concentration: 8X
             21

-------
     As a  final check, the possibility of "missing N" being removed from the
 system along with the precipitated iron insolubles should be tested, even
 though this mode of denitrification seems very unlikely.  There are no known
 insoluble  nitrates or nitrites that can form in this system, and while metal
 oxides and their gels (e.g., alumina, Ref. A-50, A-6o) have been shown to
 absorb NOg and NOjj ions,  such absorption is extremely inefficient, even in
 concentrates.

     The catalyst ions are also undoubtedly incorporated, at least in part,
 into the insolubles.  However, loss of nitrogen as a cupric ammine, such
 as Cu(NH3)ii  , is also unlikely.  For instance, in tests number 2a and 2b
 (Table U)  the molar ratio of "missing N" to cupric ion approaches 12 to 1.
 Thus,  much more denitrification occurs than can be accounted for by the
 formation  of a cupric ammine.

     Also, as win be discussed subsequently, the pH drops as the reaction
 progresses.  With an initial pH near 7 (as in tests 2a and 2b), the final
 pH will be near 5.  The simple ammine complexes dissociate readily under
 acid conditions.

     There are a few known complexes of N20 with salts, such as KgSOVNgO
 (Ref.  A-U).  These behave much like the ammines; the NgO is liberated from
 such species in dilute acids.  Thus, although complexes of ammonia or N20
 might  well play a role in the mechanism of the reaction during its initial
 alkaline stage, it does not appear that any significant part of the denitri-
 fication can be attributed to complex formation and loss of such complexes
 as part of the insoluble  materials.

 Effect of pH on Denitrification

     An results reported in this section and the remainder of this report
 were obtained at a 10 ppm NO§-N concentration, in order to generate data
 at a level near that of the average secondary effluent.  The effect of
 varying pH is shown in Table 5.

     There is no readily  evident correlation of pH and denitrification from
 these  data, even excluding the apparently anomalous results in Test 5a.
 There  is a discernible trend for increased reduction with increased pH,
particularly if one considers the extremes (U.3 ppm average residual NO§-N
 at pH  7.0, 0.2 average residual at pH 11.0).

     With the exception of Test 7, the predominant reduced products were
 "missing N" and ammonia.  In tests 3 and 6,  these two products accounted
 for 70 to 90$ of the total nitrogen.   A combination of processes,  where
ferrous ion denitrification would be fonowed by ammonia stripping, thus has
the potential for close to complete denitrification.

     It should be noted that pH was not constant during the course of the
tests reported in Table 5.  The reaction system became more acid as the
reduction progressed.   Starting at pH 7, the  final pH was between  5 and  6.


                                   22

-------
            TABLE
EFFECT OF pH ON DEIIITRIFICATIOW
Test No.
1. a
b
2. a
b
3. a
b
k. a
b
5. a
b
6. a
b
7. a
b
Initial
PH
6.0
6.0
7.0
7.0
7.5
7.5
8.G
0.0
8.5
8.5
9.0
9-0
11.0
11.0
Nitrogen Balance, pp:.i
NO"-N
9.5
9-5
3.6
2.1
2.0
3-5
3.3
7.0
3.5
0.0
2.3
0.1
0.2
NO^-N
0
0
0.6
0.7
0
0.9
l.ci
. 0.9
2.7
1.1
0.1
O.U
6.1
5.7
NH.,-N
0.5
3.0
2.2
3^
3.5
3.6
0.5
h.6
k.Q
1.1
1.0
Fixed Conditions:
Initial NOl-N Concentration: 10 ppm
Fe Concentration: CX
Catalyst: Cu , 5 ppm
Reaction Time: 2U hr
Temperature : 85 F
Atmosphere: Nitrogen
Denitrification, %
0
0
28
22
1*5
37
12
22
0
20
33
27
31







                 23

-------
 Starting at 11,  the final pH was between 8 and 9.   An additional experiment
 was carried out  in which the pH was readjusted manually during the course  of
 duplicate runs in order to hold it near a value of 7.  With all other fixed
 conditions the same as those shown in Table 5, the final nitrogen balances
 were as follows:  N03-N, 6.9, 5-7 ppm; N02-H,  1.8, 2.3 ppm;  WH3-N, 0.8,
 l.U ppm.  Thus,  % or less denitrification occurred in these runs where
 pH was held relatively constant.

 Effect of Catalysts on Denitrification

      The results of a series of tests where catalysts and their concentrations
 were varied are  given in Table 6.

      Silver ion  (Ag ) appears to be a much more effective catalyst for reduc-
 tion than is cupric ion (Cu++)  but essentially no  denitrification occurred
 when silver ion  was the catalyst.

      Silver and  cupric ions are the classical  homogeneous reduction catalysts
 (See p.  Ik for discussion and references)  and  their effect is  reportedly
 enhanced by the  presence of organic acid anions.   In subsequent  sections of
 this report,  the use of silver and cupric  acetate  will be noted, but no effect
 on denitrification can be attributed to this variation.

      Zinc,  cadmium,  and magnesium ions had no  catalytic  effect in  the reduc-
 tion reaction in the absence of organic acids.  The  ability of Zn++, Cd"1"1",
 and Ms"1"*" to catalyze the reduction or denitrification in the presence of an
 organic  acid was not examined because of time  limitations.

      It  is  important to remember that this  reaction occurs in a heterogen-
 eous system,  with precipitated  ferrous  hydroxide being the reducing agent.
 The catalysts, Cu+*  and Ag+  are, however, the  classical  homogeneous catalysts,
 and their mode of action under  the present  conditions  is not readily evident.
 The heterogeneity of this system may  be the source of much of the  seemingly
 erratic behavior in  this reaction  that  is now unexplained.  Minor variations
 in the way Pe(OH>2 is precipitated, the extent to which  catalyst ions are
 absorbed  in the  precipitate,  the manner in which the black ferrous-ferric
 complex forms from the Fe(OH)2, and a number of other factors associated
with the  solid state could all play a role.

Effect of Air on Reduction and Denitrification

     As noted earlier, the bubbling of air through solutions had confounded
our  first test results,  and data were also presented showing reoxidation of
nitrite to nitrate in screening tests (page 19).  Analytical studies on this
latter effect, which are reported subsequently, showed that nitrite reoxida-
tion would not occur if  the solutions remained basic.  To further test the
effect of air, a reaction was conducted at high pH in an open vessel,  with
the reaction mixtured stirred mechanically.  Under these conditions, 100&


                                   2k

-------
                TABLE 6
EFFECT OP CATALYSTS ON DENITRIFICATION
Catalyst
Type
Cu++
Cu++
Cu++
AE+
A6+
AC+
Fixed
Cone . , ppm
1
5
10
1
5
10
Conditions : Ag
Nitrogen Balance, ppm
NO~-N
6.7
0.2
0.3
0.1
0.3
0.2
NO--N
2.8
5-9
5-9
5-9
8.5
5-3
NH -N
0.8
1.0
—
3.0
l.l
5.1
Denitrification, %
None
29
—
(Trace?)
None
None
and Cu as the sulfate and chloride
NO"-N Concentration: 10 ppm
Fe Concentration: 8X
pH: 11 (initial)
Reaction Time: 2U hr
Temperature : 65 F
A tmo sphere : Ni trogen
Values reported are averages of duplicate runs
                   25

-------
 reduction was observed,  which shows  that  no N0§  is  formed by H0j> reoxidation
 under such conditions.   However,  no  denitrification occurred in this test:
 the final nitrogen balance showed no nitrate; N02-N, 7.8 ppm, and HH3-N,
 2.3 ppm (all fixed conditions,  except atmosphere, were the  same as Test 7 in
 Table 5).  It is difficult to attribute a failure to denitrify to the presence
 of air, particularly when reduction  has been quantitative.  The result may
 well be confounded by the "occasional unknown cause" mentioned earlier.

 Lime for pH Adjustment

      In all the work described earlier in this report, sodium hydroxide was
 used to make initial pH  adjustments.   Table 7 shows the effect of using lime
 (CaO) in place of the NaOH.

      Lime obviously does not  interfere with denitrification unless, as in
 Test 3, enough lime is added  so that  the  pH remains constant as the lime
 slowly dissolves.   However, some  decrease in total  reduction may have occurred.
 Compare Tests la and Ib  (average  68%  reduction)  in  Table 7 with Test 6 in
 Table 5-   In the latter  test, with NaOH,  an average of 85# reduction occurred.
     Also  shown in Table 7 is the effect of increased Fe^level.  In Test 2,
this drove the reduction largely to MH3, although some denitrification still
occurred.

Effect of  Phosphate and Carbonate Ions

     The addition of a mixture of phosphate and carbonate ions (as KgCOs and
KHgPOli) inhibits the reduction reaction.  As the duplicate results in Table 8
indicate,  this inhibition is caused by the phosphate, probably by deactivating
the catalyst.

     Phosphate appears to be exceptionally powerful in its ability to inhibit
catalysis  by Cu4*.  In additional tests, increases in Cu1"*" concentration to
20 ppm or  a change from CuCl2 to cupric acetate (CuAcg), again with an increase
to 20 ppm, were unable to overcame the effect.  However, two other methods were
tested which did prove effective: increasing the iron concentration and sub-
stituting Ag+ for CVL++ as the catalyst.  Results are shown in Table 9.

     Bbte  that in all of the tests above, where 003 alone was present, or
where the  effect of P01|~3 was overcome by one means or another, there was
essentially no denitrification, even though reduction reached the 90$+ level.
The carbonate ion buffers these reactions, so that pH remains relatively
constant throughout the reaction.  Similar results (reduction without denitri-
fication)  in the absence of carbonate were noted earlier, when NaOH or lime
was added periodically during the course of a reaction so that the pH remained
at a high level.  Thus, there is an apparent correlation showing that denitri-
fication is suppressed by having a constant pH, although this relationship is
partially confounded by the fact that Ag+, which catalyzed some of the buffered
runs, does not appear to induce denitrification even in the absence of carbonate.

                                   26

-------
                 TABLE 7
DENITRIFICATION IN THE PRESENCE OF LIME
Test
No.
1. a
b
2.
3.
Lime Added
ppm
300
300
700
7^0
_ -H-
Fe
Level
8X
8X
i6x
i6x
pH
Initial
9.0
9.^
10.2
10.5
Final
6.0
6.1*
C.5
10.7
Nitrogen Balance, ppm
NO"-N
3.1
3.3
0
h.o
NOg-N
1.6
l.U
0.3
2.U
NH3-N
3.0
3.0
7.9
3.6
"Mi s sing ' N"
2.3
2.3
1.0
None
Fixed Conditions:
NOl Concentration: 10 ppm
++
Catalyst: Cu , 5-10 ppm
Tempera ture : 6"5° F
Atmosphere: Nitrogen
Reaction Time: 2b hr
                    27

-------
               TABLE 8
EFFECT OF PHOSPHATE AND CARBONATE IONS
pH Adjusted
with
NaOH
NaOH
NaOH
NaOH
Lime
Lime
NaOH
NaOH
NaOH
NaOH
P
Initial
9.5-10
9-5-10
n
11
7.0
7.1
10.0
9.1
10.3
9.0
H
Final
~
• m
7.5
7.5
9-9
9.7
10.2
8.1*
«V3-p,
ppm
10
10
10
10
10
10
10
10
None
None
C03
ppm
100
100
100
100
100
100
None
None
100
100
Fixed Conditions:
Nitrogen Balance, ppm
NO~-N
8.2
7.2
9.3
8.0
7.8
8.1
9.8
8.7
1.8
3.2

NO'-N
1.6
l.U
1.1
l.U
0
0
0
0
U.3
1.8

NH.-N
0.8
0.8
0.3
1.3
1.3
0.7
1.1
3.5
k.k

N0§ Concentration: 10 ppm
Fe** Level: 6X
Temperature: 85° F
Atmosphere: Nitrogen
Reaction Time: 2k hr
Catalyst: Cu'H', 5 ppm
                  28

-------
                    TABLE V
OVERCOMING THE EFFECT OF PHOSPHATE AND CARBONATE
Fe Level
l6x
2Ux
Ox
8x
5x
Catalyst
Cu , 5 ppm
r. •*-«• c
LU , ) ppm
Ag , 5 ppm
Ag+, 5 ppm
AS+» 5 ppm
Ag+, 5 ppm
(AgAc)
t>H Adjusted
with
NaOH
NaOH
Lime
Lime
Lime
Lime
pH
Initial
O.Q
9-9
10.5
10.5
10.5
10. I*
Final
10.2
10.3
10.3
10.3
9.U
9-1*
Nitrogen Balance, ppn
NO^-N
U.7
3.0
1.7
2.0
0.9
1.1
NO'-N
3A
3.3
6. ii
6.8
5-0
5.3
I«3-N
2.1
u.o
1.2
l.U
3.7
3.3
Fixed Conditions:
NO~-N Concentration: 10 ppm
PO^-P Concentration: 10 ppm
COlr Concentration: 100 ppm
Temperature : 85 F
Atmosphere: Nitrogen
Reaction Time: 2k hr
                       29

-------
      There is no immediately evident explanation for the above effect of
 constant pH.  Obviously, pH-related phenomena win have to be investigated in
 much more detail before applying this reaction in an actual water reclamation
 process.
 ANALYTICAL SUPPORT
 Background and Selection of Methods

      Nitrate ion, which represents  the most highly oxidized phase in the
 nitrogen cycle,  generally occurs  in trace quantities in  surface water
 supplies.   Since a  limit of 1*5 mg/1 nitrate (10 ppm NOo-N) has been imposed
 on drinking waters,  a number of analytical chemistry methods are available
 in the  literature to determine NO§-N at the ppm level.  Methods available
 in the  literature for nitrate ion determination were examined and Judged for
 relative merit with  the particular  parameters of the nitrate ion reduction
 effort  in mind.  For example, for any given set of experiments, analyses
 were  also required for nitrite ion  and to complete a material balance, a
 final analysis for ammonia might  be required.  All of these were to be deter-
 mined in the presence of a reducing agent and possibly a deammination agent.
 Accordingly,  methods for nitrite  ion and ammonia were also examined.  A brief
 review  of the methods considered  is given below.

      Nitrate  and nitrite analyses can be performed according to the method of
 Fisher,  Ibert, and Beckman (Hef.  F-6) which utilizes the sulfur-yellow color
 produced by brucine  in sulfuric acid solution.  By varying the concentration
 of sulfuric acid (less than 25$ for nitrites, greater than 5056 for nitrates)
 the two  can be determined on aliquots as small as 15 ml containing 0-1 micro-
 grams of the  anion.  A modification of this procedure has been incorporated
 as  a  standard method for nitrate  in the Standard Methods for the Examination
 of  Water and Wastewater  (Ref. F-l).  Ferrous and ferric ion have been reported
 to  give  slight positive  interferences.  Interference due to nitrite ion is
 eliminated by the use  of sulfanilic acid.   Where nitrate ion alone is desired
 another  common practice  is to destroy the nitrite ion using solid sulfamic
 acid  (NH2S03H).  The reaction with  sulfamic acid (Ref.  F-5) is almost instan-
taneous  and is not interfered with by the species used in the nitrate reduc-
tion studies.  Nitrite ion may also be determined by the coupling of diazotized
 sulfanilic acid with 1-naphthylamine hydrochloride at pH 2.0-2.5 with the
 formation of the reddish purple azo dye (Ref.  F-l).   The method is sensitive
to 0.1 ppm NOg-nitrogen in a 10 ml sample.   Bastion,  et al.,  (Ref. F-3)
reported an ultraviolet  spectrophotometric  method for the determination of
nitrate ion in alkaline earth carbonates.   The method is based on the
absorption of nitrate ion in the 200-220 millimicron region.   The absorp-
 tion maximum is 200 nM, but in the systems  studied, measurements at  210were
found to give optimum results.  Armstrong  (Ref. F-2)  used a modification of
this method to determine nitrate  in sea water.  The  samples are  run  in  50$


                                   30

-------
      and 0.05 M HC1.  At these concentrations, both nitrite and nitrate
have absorption naxima at 22? mu.  An ultraviolet spectrophotometric method
has been recommended in the Standard Methods of Water Examination as useful
for screening large numbers of drinking water samples for nitrate ion.

     The speed, accuracy, precision, and interferences of each of the above
methods were considered in the selection of a candidate procedure for nitrite
and nitrate.  Of the methods reviewed, the direct measurement of ultraviolet
absorbance appeared most desirable with respect to speed and simplicity and
was selected for laboratory evaluation.  Methods for determination of ammonia
were also examined (Ref. F-l, pp 186-19U).  The method of choice was distil-
lation from a strong base in a micro-Kjeldahl distillation flask, followed
by titration with a standard mineral acid.  Since urea was a deammination
agent under consideration and consequently a source of nitrogen, methods for
its determination at the ppm level were reviewed.  Two possible methods appeared
appropriate (Ref. F-7, F-9).  The former is a spectrophotometric procedure
using p-dimethylaminobenzaldehyde and the latter is a conversion to ammonia
by enzymatic hydrolysis.

Analytical Program

     The major part of the analytical program consisted of the evaluation and
development of the ultraviolet spectrophotometric method for nitrate and
nitrite ion by thoroughly checking out possible interferences from the candi-
date reducing agents and deammination agents.  The ammonia distillation method
was examined to determine sensitivity limits and interferences.  Subsequent
to this, the tentative procedures were applied to the experimental program.
From time to tine, analytical anomalies arose as a result of changes in the
experimental plan (new reducing agent-deammination agent pair, new catalysts,
different pH, etc.), or interferences that were not considered at the begin-
ning of the program.  These anomalies were investigated and modification or
improvements in the methods were made.  As the program progressed and ferrous
ion became the most frequently used reducing agent, the analysis scheme became
routine and the analytical efforts consisted entirely of analysis of samples.
A complete description of the ultraviolet spectrophotometric method used
is shown in the Appendix at the end of this report.

Development of the Ultraviolet Method

     Some of the reagents selected to be screened for the reduction of nitrate
and subsequent deammination of nitrite were procured, and solutions of each
were prepared.  Solutions of potassium nitrate and potassium nitrite were also
prepared.  Each solution was subjected to ultraviolet spectrophotometric
scanning with a Gary I1* Recording Spectrophotometer, covering the region in
which nitrate is absorbant.  The results of this examination are shown in
Table 10.
                                   31

-------
                        TABLE 10
UV ABSOKBANCE OP NITRATE, NITRITE,  AND CANDIDATE AGENTS
Compound
Potassium Nitrate
Potassium Nitrite
Sucrose
Urea
Formaldehyde
fydrazine Sulfate
Sulfamic Acid
Nitrite +
Sulfamic Acid
Concentration
ppm
1 (as N)
2 (as N)
100
100
100
100 (as Ngfy)

2 (as N) +
large excess
Wavelength of
Abs. Maximum, mia
205
210
< 190
< 190
< 190
< 190
< 190
< 190
Specific Absorbance,
= Optical Density
g/1 x Cell Thickness
690
375
Not applicable
Not applicable
Not applicable
Not applicable
Not applicable
Only sulfamic acid
absorbance seen
                           32

-------
     Examination of the data in Table 10 reveals that nitrate has an absorp-
tion maximum at 205 EU» and nitrite has a maximum at 210 nti, in good agreement
with ultraviolet spectra recorded elsewhere (Ref. F-2, F-3).  It should be
noted that although Standard Methods (Rcf. F-l) cans for measurements at
220 nil, this is not an absorption maximum.

     The absorbances of the candidate reducing agents at wavelengths below,
but near, 190 flJ apparently do not interfere with the nitrate or nitrite
absorbance measurements.  The solutions vere so concentrated relative to the
nitrate and nitrite solutions studied that there was some overlap of peak
shoulders (or tail) with the 200 rau region.  It is estimated from these
studies of individual solutions that at reducing agent concentrations of 50
times that of the nitrate ion, the overlap would introduce an error of about
yjt in the nitrate ion determination.

     The addition of solid sulfamic acid to a 2 ppm nitrite-N solution, fol-
lowed by immediate spectrophotometrie scan, was found to destroy quantita-
tively and immediately the nitrite ion.  This experiment indicated nitrite
ion can be removed easily and quickly.  Absorbance measurements before and
after nitrite removal will yield values for both ions, by difference calcu-
lation.

     For the initial phase of the program, four reducing agents vere chosen
for study:
                    J.A
     Ferrous ion (Fe  ) as FeSOU
     Sulfur dioxide (802)
     Filtrasorb kOQ carbon
     Formaldehyde

     The UV spectrophotometric method was further examined for interferences
of the above species at stoichiometric and threefold excess.  None of these
species seemed to interfere with the nitrate determination.  Sulfamic acid
at threefold concentrations introduced a small error, but this proved easy
to correct by use of a reagent blank.  Calibration curves were prepared from
stock sodium nitrate and sodium nitrite solutions, respectively, and the
method was used to analyze the first experimental mixtures.  The analytical
procedure involved the measurement of the UV absorbance of a diluted portion
of the sample, addition of solid sulfamic acid to the diluted solution, and
a re-examination of the UV absorbance.  Experience showed that measurements
at 200 mu gave the most reproducible results.  Since sulfamic acid reacts
quantitatively with nitrite ion, the UV absorbance after addition of sulfamic
acid was a measure of the nitrate ion and the difference in UV absorbance
before and after sulfamic addition was a measure of the nitrite ion.

     During the course of the early analyses, deammination evidently was not
accomplished in some cases in which it was expected.  Study of the problem
revealed that incomplete deammination appeared to be  restricted to cases


                                   33

-------
 in which the pH was not highly acidic.  Data from experiments performed on
 known sodium nitrite solutions containing sulfamate ion at various concentra-
 tions are given in Table 11.

      The results shown in Table 11 indicated that deammination did not occur
 rapidly until the solution was quite acid, and that neither nitrate nor nitrite
 can be measured by UV absorbance in strongly basic solutions because the
 hydroxyl ion interferes.  Basic solutions were found not to deanuninate com-
 pletely in 1*8 hours, but in acid solution the deammination was very rapid.

      Indication of the interference by hydroxyl ion prompted pH-absorbance
 studies.  Solutions adjusted to cover the pH range were prepared and their
 UV spectra were taken on the Gary, Model lU, spectrophotometer.   The absor-
 bances at 210 uu of the solutions are shown in Table 12.

      It appears that the hydroxyl ion precludes ultraviolet analysis of
 basic solutions, and the analytical procedure was altered to include addition
 of excess acid to the aliquot of sample being diluted for analysis.   The work
 of Bastion (Ref. F-3) showed that perchlorate ion shows no UV absorbance at
 230-200 muj  therefore,  acidification of samples was accomplished with perch-
 loric acid.   Concentrated perchloric acid is a powerful oxidizing agent, but
 when diluted to less than 20$,  HC10U has virtually no oxidizing  power (Ref. F-8)
 The amount of perchloric acid necessary to bring the pH of sample solutions
 below 3 is insufficient to reoxidize nitrite to nitrate during the course
 of the analysis.

      Investigation of sulfite ion interference with UV nitrate analysis was
 undertaken when test results of solutions containing sodium sulfite  showed
 anomalies between sample and reagent blank when sulfamic acid was added to
 such solutions.   The difficulty was  resolved when,  in the  light  of the pH
 studies  described above,  it  was realized that  the addition of sulfamic acid
 to the  sample  solution  during analysis  dramatically changed the pH, and the
 sample and reagent blank may have been  buffered more  or less  by the sulfite
 ion.  The possibility of sulfamic acid  decomposition  in basic  solution to
 yield ammonium ion was  also  considered, but  since the high pH solutions
 could not be analyzed as received, this possibility was  not pursued.

      The  studies  indicated that caution must be exercised  in  interpreting
 analytical results of some reduction experiments, since deammination was
 found to be pH-dependent.  The  effect of pH changes during the analytical
procedure which increase the deammination rate were borne in mind, to prevent
alteration of nitrite ion content during the analysis.

     An examination was made of the air oxidation of standard solutions con-
taining varying concentrations of nitrate and nitrite ion.  Using standard
handling and transfer procedures, reoxidation of significant amounts of N0§
occurred at pH's of 3 or lower when the time of handling exceeded one hour.
When analyzing a series of six or more samples, the total analysis time
frequently exceeded one hour.  At pH's of 7 or above, oxidation did not

-------
                 TABLE 11
ANALYTICAL STUDY OF SULFAMATE DEAMMINATION




            (3 ppm NO~-N Level)
Solution
No.
1. a
b
c
d
2.
3.
k. a
b
Treatment
No sulfamic acid
Increment of sulfamic acid
added, lowering pH
Further increment added
Further increment added
Near stoichiometric amount
of sulfamic acid added
Sulfaraate added, base added
Sulfamate, base added
Solution acidified
pH
H.O
3.5
3.0
3.0
2.3
11
6.9
2.7
Absorbance at
210 nu
1.1U
0.82
0.25
0.02
0.05
OH" ion
interference
1.12
0.10
                     35

-------
          TABLE 12
EFFECT OF pH ON UV ABSORPTION
Solution
NaOH
HC1
HC10U
NaOH + HC1
PH
11
5
3
3.8
Absorbance at 210 ny
very strong
nil
nil
0.03
             36

-------
occur during this time period.  Therefore, during analyses of ferrous ion
reduction tests, oxidation did not occur during the filtration step as long
as the solution was above pH 7.  To prevent any reoxidation of N0j> at low
pH's, the procedure was modified so that each sample was scanned in the
appropriate UV range immediately after the acidification step.

     The procedure described in the Appendix was used to carry out a series
of tests on standard nitrate-nitrite mixtures to establish error limits.
For samples containing 50 ppm total K, the limits of the procedure were found
to be less than 1 ppm for N0§, and less than 2 ppm for NOJ3 when a large
amount of NO" is also present.

Ammonia Distillation Method

     Analysis for ammonia was examined to ascertain the lower limits of
detection by the distillation method, and possible interferences from nitro-
geneous reducing and deammination agents.  Some literature information on
the stability of the sulfamate ion was obtained (Ref. F-lt); however, analysis
of known standards containing ferrous ion and sulfamic acid was considered
necessary.  The distillation method, deemed most reliable for the purpose
of the reduction study, involved the addition of an aliquot of the sample
and a volume of sodium hydroxide to a small still.  Ammonia in the sample
is steam distilled, caught in a boric acid solution, and titrated with dilute
standard acid solution.  The examination of nitrate and reducing agent solu-
tions was performed with 25 ml aliquots of each solution studied.  The results
are shown in Table 13.

     The data of Table 13 show that recovery of ammonia from dilute solutions
of ammonium ion is sufficiently quantitative to monitor formation of ammonia
from nitrate reduction at the levels of interest in this effort.   However,
the recovery of ammonia from solutions prepared with ferrous sulfate and
nitrate, and with sulfamic acid, are potential sources of error.   The sul-
famic acid contribution of ammonia may be the result of slow decomposition
of the reagent.  The ferrous sulfate or nitrate alone yielded no ammonia,
but the combination yielded ammonia, indicating that reduction of nitrate
to ammonia occurred to some extent under the conditions of the distillation.
The error, however, was only about 1 ppm at the 50 ppm nitrate nitrogen level
and the method was used without further modification.
                                   37

-------
                          TABLE 13
AMMONIA DETERMINATIONS ON SYNTHETIC NITRATE REDUCTION SAMPLES



               (25 ml Aliquots in All Cases)
Solution
Ammonjun salt
AmmonJiB salt
50 ppm NOo-N (plus 3 x stoichiometric
amount of FeSOVTIfeO (150 mg) plus 3 x
stoichiometric amount of NILSO-H (26 mg)
50 ppm NO~-N
150 mg FeSO^'THgO
26 mg NHgSO^H
50 ppm NOl-N plus 150 mg FeSO. 'TILO
50 ppm NO"-N plus 200 mg FeSOr -THpO
Ammonia Nitrogen
ppm Added
1
10
0
0
0
0
0
0
ppm Recovered
0.96
9-9
1.2
0
0
0.2
0.9
1.9
                             38

-------
                         EXPERIMEMTAL DETAILS
APPARATUS
     Reactions were conducted in 200 ml volumetric flasks, each containing a
mechanical stirrer and a tube inserted into the neck to provide the appropriate
gaseous atmosphere.  (During the first few reactions, the gas tubes were in-
serted into the liquid, and the gas bubbled through, but this process did not
effectively stir suspended solids, and was abandoned.)  The reactors were
immersed in a thernostatted water bath during the reaction period.  Adjustments
and changes in pH were monitored with a Beckman pH meter (Model G).
REAGENTS
     Reagents and their sources are listed below:
Water, lUO
Potassium Nitrate, KNO.,
Sodium Nitrite, NaN02
Sodium Hydroxide, NaOH
Ferrous Sulfate, FeSOj/THpO
Carbon, C
Sulfur Dioxide, S02
Formaldehyde, CHgO
Glucose, CgHipOg
Iron Powder, Fe
Hydrazine, NpH.
Hydrazine Sulfate, NpHgSCV
Carbon Monoxide, CO
Sulfamic Acid, NHgSO_H
Cupric Chloride, CuCl2
Silver Sulfate, AggSO^
Cupric Acetate, Cu(CpH Ogjp
Silver Acetate, AgCgH,02
- From a laboratory demineralizer
- J. T. Baker Chem. Co.
- J. T. Baker Chem. Co.
- Mallinckrodt Chem. Works
- J. T. Baker Chem. Co.
- Calgon Corporation's Filtrasorb
- The Matheson Co.
- 36.6% Formalin, J. T. Baker Chem.  Co.
- Eastman Organic Chem. Co.
- Mallinckrodt
- Eastman
- Fisher Scientific Co.
- The Matheson Co.
- Eastman
- Mallinckrodt
- J. T. Baker
- J. T. Baker
- Mallinckrodt
                                   39

-------
 Ferric  Chloride,

 Ammonium Vanadate, Utt^MO

 Vanadium Pentoxide, V^

 Lime» CaO
 Potassium Carbonate, KgCO-

 Potassium Phosphate,
 Monobasic,
                                   - J. T. Baker

                                   - Harshaw Scientific

                                   - J. T. Baker
                                   . j. T. Baker

                                   - J. T. Baker


                                   . General Chem. Division, Allied Chem.
 PROCEDURE
Note  ha      an
Note that a blank
                                         the redUCing ***  is  Ascribed below.
                                  water with no NO^)  was  treated  in  the same
                                                  J       *««•«»  m  me same
 manner as described for the sample.

     »/"?? ^les were PreP«ed by adding standard nitrate  solutions (2 mg
         }y meanS °f a Sma11 burette to water  contained  in 200 ml volumetric
 AhS™* f^hC Wa?6r *' initiaU^  Just short <* the desired 200 ml.
 After the addition of the various  reagents, the volume was  adjusted to the 200

   i??J'«re£OUV0n WaS addCd t0 the nitrate  solution in  ^ form of the
 s'Siton ?^™2? (c^S??^Wei8tad  °n the •M^1C1 f*1^).  The catalyst
 solution (a*.,  1 ng Cu++/nil)  was then added by  means of a pipette.  After
 thorough mixing, the pH of the solution  was adjusted to the desired level by
 adding a few drops of a sodium hydroxide solution  (6M).  The pH of the blank
 was first adjusted in order to establish the volume of base necessary to
 reach the desired pH.   Addition of the same volume of base to the test
 samples,  made further adjustments  unnecessary.  Next, the flasks were Blaced
 in a constant temperature bath (85° F) and siirred m^chanicalS unler a blanket
 of nitrogen  gas.  At  the end  of the reaction period, samples^ wUhL-awn for
 pH detennination,  nitrate-nitrite  analyses and ammonia analysis.       "&Wn Ior

     Deammination  experiments  involved urea and sulfamic  acid.  These were
 usually added Just prior to the pH adjustment.  Urea was  added  as the solid

                                                                           '
     For the experiments where phosphate, carbonate,  and lime were used the
order of addition consisted of: adding phosphate solution (2 mg P/ml) to the
nitrate solution, then adding the carbonate solution  (20 mg col/ml)  adding
weighed amounts of lime (with thorough shaking),  then add^g FlsSj  and  ^
adding the catalyst solution at the very end

-------
TEST RESULTS
     With the exception of the reducing agent screening series,  details  of
all reactions have been tabulated in the DISCUSSION AND RESULTS  section  of
this report.  The screening test results are given here in Table lU.
                                   Ul

-------
                               TABLE
                     SCREENING OF REDUCING AGENTS
Fixed Conditions:
     NO" Concentration:
     Temperature:
     Atmosphere:
     Reaction Time:
     Reducing Agent Concentration:
50 ppm
85° F
Nitrogen
1*8 hr
3x
                                             J Denotes 3 moles of agent per
     Deaomination Agent Concentration:  3x   ) mole of NO"
        (Occasional variation in conditions noted in test)
Reducing Agent
so2
so2
so2
so2
Carbon
C
C
C
C
Deammi nation
Agent
(when added)
NaSO-NHg
NaSOoNHo
(21* fir)
HS03NH2
(start)
HSOoNHp
(start)
HS03NHp
(2U hr)
HS03NH2
(2U hr)
HSOoNHg
(Start)
HS03NH2
(Start)
Catalyst
(concentration, ppm]
Fe~+(2.5)
Cu++ (2.5)
Fe+"(2.5)
Cu++(2.5)
Fe+++(5)
Cu^(5)
Fe+"<5)
Cu++(5)
PH
9
9
6
6
6
6
6
6
NO^-N, ppm
After
2k hours
^9-7
51.8
50.7
50.2
50.8
51.1*
1*9.2
50.9
After
1*8 hours
1*8.1*
1*9.9
50.8
50.7
51.1
52.1*
50.3
53.1* (?)
                                  1*2

-------
TABLE 1^4 (CONT'D)


Reduci:ig Agent
Formaldehyde
CH 0
2
CH.O
CHgO
CHgO
Carbon Monoxide
CO
CO
Glucose
C6H12°6

C6H12°6
C6H12°6
C6H12°6
C6H12°6
Dearami nation
Agent
(when added)

HSOoNHj
(stSrtT-
HSO^NH?
( start )
HS03NH2
(2k hr)
HSOoNHg
(21; hr)
None
«<«,>,
CO(NH2)2
(2k hr)
CO(NH2)2
(2U hr)
CO(NH2)2
(start)
None
None

Catalyst
(concentration, ppm)

Fe+++(2.5)

Cu++(2.5)
Fe+~(5)
Cu++(5)
Cu++(5)
Cu++(5)
Fe+++(5)

Cu++(5)
V+5(5)
V+5(10)
v+5(io)


pH

6

6
6
6
6
6
6

6
6
1*
k
]1
NO"-N, ppm
After
2k hours

U9.6

50.2
50.2
50.9
U9-8
51.7
^9-5

U9.U
50.1
_ kj.6
(NOg-N,*3.5)
After
U8 hours

1*9.9

^9-7
50.7
51.2

--
1*9.0

i.9.1
U8.5
(NOg-N, 2.8)
U8.U
(NOg-N, l.U)
U9.8
(NO'-N, 2.6)

-------
TABLE Ik (CONT'D)

Reducing Agent
Hydrazine
Ngfy

Ngfy

NgH.

Hydrazine Sulfat
S»
NoH/rSOi,
(3.75XJ
^HgSOij.
(3.75X)
NoIkSO),
(3.75X)
Iron Powder
Fe
Fe
Fe
Fe
(U.5X)
Fe

Deammination
A .A.
Agent
(when added)

CO(NH2)2
(21* hr)
CO(NH2)2
(21* hr)
CO(NH2)2
(start)
e
None
None

None

None


HSOdnu
(21* hr}
HS03NHo
CO(NHg)2
(start)
None

None


Catalyst
(concentration, ppm)

Fe+"H"(0.7)

Cu**(0.7)

Cu++(0.5)


Cu+*(0.25)
Cu++(0.25)

Cu (l)

Cu++(l)


Fe++(2.5)
Cu^(2.5)
Cu++(2.5)
Cu+*(lo)
» »
Cu++(lO)


PH

n

n

n


11
n

n

11


6
6
6
6

6

NO"-N, ppm
After
2k hours

50.1*

J+6.2
•
1*8.8


_ 1*8.2
1*3.9
(NOg-N, 8.6)
1*8.6
(NOg-N, 3-2)
1*8.0
(NO'-N, !*.!*)

53.6(7)
1*5.3
!*9.5
1*0.1*
(NO'-N, 6.8)
29-9
(NOg-N, ll*.8
After
1*8 hours

M.5

1*3.6

30.1*


(NO'-N, 0.1*)
1*0.1*
(NOg-N, 0)
1*1.8
(NOg-N, 5.2)
Ul.3
(NO--N, 5.U)

59.5(?)
1*3.5
1*8.9
• 27.3
(NOg-N, U.8)
27.1*
(NO'-N, 10.3)
      kk

-------
                               REFERENCES


 A.   NITRATES IN WATER


 General Chemistry

 1.   Jolly,  W.  L.,  The Inorganic Chemistry of Nitrogen. W. A. Benjamin Co.,
      New York,  New  York,
 2.    Mbeller,  T.,  Inorganic Chemistry. John Wiley & Sons, New York, New York.
      1952.

 3.    Szabo, Z. G., and Bartha, L. G., Recent Aspects of the Inorganic
      Chemistry of  Nitrogen. Special Publication No. 10, The Chemical Society,
      London, 1957, pp. 131-136.

 U.    Gehlen, H., "Reactions and Properties of Nitric Oxide and its Compounds.
      II. The Salts of the Nitric Oxide Compound of Sulfurous Acid," Ber.
      6_5_B, 1130-Uo  (1932).

 Occurrence and Effects

 5.   Azad, H.  S., and King, D. L., "Effect of Industrial Waters on Lagoon
     Biota," Purdue Univ., Eng. Bull., Ext. Ser. No. 118, UlO-22 (1965).

 6.   Biffoli, R. ,  "Determination of the Nitrate Ion and its Hygienic Impor-
     tance in Tap Water.  II. Results of Recent Controls in the Province of
     Florence."  Bon. Lab. Chim. Provincial! (Bologna) 16 (5), 558-70 (1965).

 7.   Burden, E.A.W.J., "The Toxicology of Nitrates and Nitrites with Parti-
     cular Reference to the Potability of Water Supplies - a Review,"
     Analyst 86, ^29.33 (1961).

 8.   Darvas, I., "On the Correlation Between the Nitrate Content of Bacterio-
     logically Contaminated Waters and the Incidence of Methemoglobinemia,"
     Egeszsegtudomeny 10 (4), 357-66 (1966).

 9.   DeMarco, J., Kurbiel, J., Symons, J. M., and Robecka, G.,  "Influence of
     Environmental Factors on the Nitrogen Cycle in Water," J.  Amer.  Water
     Works Ass. 5_£ (5),  580-92 (1967).

10.  Fair, G. M., "Protecting the Purity of Inland Waters," J.  Sanit.  Eng.
     Div., Am.  Soc. Civil Engrs.  go (SA6) 1-11 (196U).

11.  Fassett, D.  W.,  "Nitrates and Nitrites," Nat.  Acad.  Sci. - Nat.  Res.
     Counc., Publi. No.  1351*,  250-6 (1966).

-------
  A.    NITRATES IN WATER


  Occurrence and Effects (Continued)


  12.   Ferguson, F. A., "A Nonmyopic Approach to the Problem of Excess Algal
       Growths," Envir. Sci.  Technol.  2, 188-193 (1968).


  13.   Feth, J. H., "Nitrogen Compounds in Natural Water - a Review "  Water
       Resources Res. 2 (l),  1+1.58  (1966).


  1U.  Flaigg  N. G., and Reid, G. W., "Effects of Nitrogeneous Compounds on
       Stream Conditions," Sew. and Ind. Wastes 26 (9), ll^-jU (195?).
 15'  frr™n«C;-R'J TWutrieS! Budfiet: Rational Analysis of Eutrophication in
      a Connecticut Lake," Envir. Sci. Technol. 1 (5),  1*25-8  (1967).
 16.  g^^-J^*^ ™ Salt and Potable Water," Kiel Meeresforsch.



 17.  Hanson, A.  M., and Flynn, T. F., "Nitrogen Compounds in Sewage," Purdue
 18.  Hollis, M. D., "The Water Ponution Situation  - Aspirations and Realities "
      J.  Water Pollution Control Federation 37,  1-7  (1965).              -""es,


 19.  Imhoff, K., "Nitrates Again.  A Summary, "  Gesundh.-Ing. 6Jt, 632 (19!*!).
 *°*  JSSJ1; «; i" 5"? "Mfl^r, S.  A.,  "Excessive Nitrite Nitrogen in
               Sludee>  J' Wafcer *>^tion control Federation 33, 1286-9
 21.  Jtopanadze, Sh. Kh   "Maximum Permissible Concentration of Nitrates  in
     Water," Gigiena i Sanit.  26,  No. 9, 7-11 (1961).
              .-            ln "^ water>"


23'  SSfitt: s'(8tei3?"J  (^r*" on Mater Use>" J" *"• wster

2U.  McCarty, P. L.,  et al,  "Nutrient^Associated Problems in Water Quality
     and Treatment,   J.  Amer. Water Works Ass. 5J3 (10),  1337-55 (1966).

25.  Nichiporovich, A. A., "Photosynthesis and Mineral Fertilizers "
     Agrokhimiya.  196!*  (l), Uo-52.                              '

26.  O'Brien  J E.,  and Rosenthal, B. L., "Sman Activated Sludge Plants
     Nitrate-Alkalinity-pH Relationship," Sanitalk 9_ (l), 18-21(1960-1) ? '

-------
A.   NITRATES IK WATER

Occurrence and Effects (Continued)

27.  Petr, B., and Schmidt, P., "The Influence of Altered Living Environment
     on Children.  II. Erythrogram and Methemoglobinemia in Blood," Ccsk.
     Pediat. 21 (6), 505-5 (1966).

28.  Kueffer, H., "Nitrification and Denitrification in Waste Water Purifi-
     cation," Vom Wasser 31, 13^-52 (196U).

29.  Sturm, G., and Bibo, F. J., "Nitrate Content of Drinking Water, Especi-
     ally in the Rheingau Region," Gas-Wasserfach 106 (12), 332-U (1965).

30.  U. S. Department of Health, Education and Welfare, Public Health Service
     Drinking Water Standards, Public Health Service Publication No. 956,
     U. S. Government Printing Office, Washington, D. C., 1962, pp. U7-51.


Natural and Biological Denitrification (See also Ref. A-55)

31.  Allison, F. E., "Losses of Gaseous Nitrogen from Soils by Chemical
     Mechanism Involving Nitrous Acid and Nitrites," Soil Sci. 9_6 (6),
     1*6^-9 (1963).

32.  Brandon, T. W., and Grindley, J., "Effect of Nitrates on the Rising of
     Sludge in Sedimentation Basins," Surveyor 10J*, 7-8 (19^5).

33.  Brezonik, P. L., and Lee, G. F., "Sources of Elemental Nitrogen in
     Fermentation Gases," Air and Wat. Pollut. Int. J. 10, 1^5-160 (1966).

3U.  Camp, T. R., Water and Its Impurities. Reinhold Publishing Corporation,
     New York, 1963, pp. 280-2.

35.  Greenwood, D. J., "Nitrification and Nitrate Dissimulation in Soil,"
     Plant and Soil XVII (3), 365-77 (1962).

36.  Johnson, W. K., and Schroepfer, G. J., "Nitrogen Removal by Nitrifi-
     cation and Denitrification," J. Water Pollution Control Federation _3§,
     1015-36  (196*0.

37.  Jones, D. I. H., and Griffith, G. ap, "Reduction of Nitrate to Nitrite
     in Moist Feeds," J. Sci. Food Agr. 16 (12), 721-5 (1965).

38.  Key, A., and Etheridge, W., "Further Studies in the Treatment of Gas-
     Works Liquors in Admixture with Sewage," Inst. Sewage Purif. (Engl.)
     1936, Pt. II, 278-300.

-------
 A.   NITRATES IN WATER

 Natural and Biological Denitrification (Continued)

 39-  Lockett, W. T.,  "The Phenomena of Rising  Sludge  in Relation to the
      Activated Sludge Process,"  Surveyor IQk,  37-8  (19^5).

 kO.  McLachlan,  J.  A., "The  Settlement and Rising of Activated Sludge,"
      Surveyor go, 39-UO (1936).

 *H.  Mountfort,  L.  P., "Some Experiences with  Rising Sludge in Humus Tanks."
      Surveyor iw,  65-6 (19^5).

 te.  O'Shaughnessy, F. R., and Hewitt,  C. H.,  "Phenomena Associated with the
      Role of Nitrogen in Biological Oxidation," J. Soc. Chem. Ind. 5U.
      167-97 (1935).                                               —

 U3.  Parkhurst,  J.  D., Dryden, F. D., McDermott, G. N., and English, J
      "Pomona Activated Carbon Pilot Plant," J. Water Pollution Control
      Federation  32  (10),  R70-81  (1967).

 kh.  Hsia,  Shu-Fang,  "Dark Reduction of Nitrate by Wheat Leaves," Sheng Wu
      Hua Hsueh Yu Sheng Wu Wa Li Hsuch  Pao 2 (2), 131-3 (1962).

 ^5,.  Silver,  W.  S., "Enzymatic and  Non-Enzymatic Reactions of Nitrate in
      Autotrophic and Heterotrophic  Microorganisms," 7th Intern. Congress
      of  Soil  Science  (Madison, Wis.), 592-9 (1960).

 U6.   Robert A. Taft Sanitary Engineering Center, Summary Report.  Advanced
      Waste  Treatment Research Program. July. 196U - July.  1967. Publication
      WP-20-AWTR-19, U.  S. Department of the Interior,  Federal Water Pollution
      Control Administration, Cincinnati, Ohio, 1968, pp.  65,  68-69.

 U7.   Thames Survey  Committee and Water Pollution Research Laboratory,  Water
      Pollution Research Paper No. 11, Effects of Polluting Discharges  on the
      Thames Estuary. Her Majesty's Stationery Office,  London.   19^	
     pp. 247-255, 537.

 U8.  Vialard-Goudou, A., and Richard, C., "Reduction by Iron,  of Nitrate Ions
     to Nitrite and Ammonium Ions in Acidic  Tropical Waters,"  Compt. rend
     21+1, 978-80  (1955).


Elimination Processes. Excluding Chemical Reduction (See also  Ref.  A-36)

^9-  Bringmann, G., "Optimal Nitrogen Removal by Addition  of Nitrated Living
     Sludge and Oxidation-Reduction Control," Gesundh.-Ingr. 81, lUO-2  (1960).

50.  Calvet, E.,  Boivinet, P., Noel, M., Thibon,  H., Maillard, A.,  and
     Tertian, R., "Alumina Gels," Bull. Soc.  Chem. France  1953. 99-108.


                                   U8

-------
A.   NITRATES IN WATER

Elimination Processes (Continued)

51.  Christenson, C. W., Rex, F. H., Webster, W.  M., and Vigil,  F.  A.,
     "Reduction of Nitrate-N by Modified Activated Sludge," U.S. At.  Energy
     Comm. TID-7517 (pt. la) 26U-7 (1956).

52.  Downing, A. L., Tomlinson, T. G., and Truesdale, G. A., "Effect  of
     Inhibitors on Nitrification in the Activated Sludge Process,"  Inst.
     Sewage Purif. J. Proc. 196!* (6), 537-51*.

53.  Dukes, E. K., and Siddall-III, T. H., "Tetrabutylurea as an Extractant
     for Nitric Acid and some Actinide Nitrates," J. Inorg. Nucl. Chem. 28
     (10), 2307-12 (1966).                                             ~

5U.  Faber, F. M., Olson, H. G., and Taylor, W. A., "What is the Life of
     Silica Gel?"  Chem. Met. Eng. 28, 805 (1923).

55.  Farrell, J. B., Stern, G., and Dean, R. B.,  "Nitrogen Removal  from
     Wastewaters," Envir. Sci. Technol., in press.

56.  Fletcher, J. M., and Hardy, C. J., "Extraction of Metal Nitrates by
     Bu-PO^-HNO  " Nucl. Sci. Eng. 16, 1*21-7 (1963).

57.  Fresenius, W., Bibo, F. J., and Schneider, W., "Pilot Plant Results  on
     the Removal of Nitrate Ions from Tap Water by Anion Exchangers."  Gas
     Wasserfach 107 (12), 306-9 (1966).

58.  Gad, G., "Use of Activated Carbon for Determination of Nitrate,  Nitrite
     and Ammonia in Water and Effluents," Gas-u.   Wasserfach 79, 166-7  (1936).

59.  Knoch, W., "Extraction of Nitric Acid with Amines," J. Inorg.  Nucl.
     Chem. 27 (9), 2075-91 (1965).

60.  Kubli, H., "Information on the Separation of Anions by Adsorption  on
     Alumina," Helv. Chim. Acta 30, 1*53-63 (191*?).   •

6l.  Ludzack, F. J., and Ettinger, M. B., "Controlling Operation to Minimize
     Activated Sludge Effluent Nitrogen," J. Water Pollution Control  Federa-
     tion 3{t, 920-31 (1962).

62.  Myrick, N., Busch, A. W., and Dawkins, G. S., "Activated Carbon  Adsorp-
     tion, a Unit Process in Liquid Industrial Wastes Treatment," Proc.
     Ontario Ind. Waste Conf. 10, 193-210 (1963).

63.  Reznik, A. M., Potapov, G. G., Korovin, S. S., and Aprakin, I. A.,
     "Extraction of Nitric Acid in the Presence of Sulfuric Acid by Tributyl
     Phosphate," Zh. Neorgan. Khim. 11 (8), lB^k-6 (1966).

-------
 A.   NITRATES IN WATER

 Elimination Processes (Continued)

 6*.  Hohlich, G. A., "Chemical Methods for Removal of Nitrogen and Phosphorous
      from Sewage-Plant Effluents," Robert A. Taft Sanitary Eng. Center, Tech.
      Kept, wol-3, 130-5 (1961).

 65.  Sigworth, E. A., "Potentiality of Active Carton in the Treatment of
      Industrial Wastes," Proc. Ontario Ind. Waste Conf. 10, 177-92 (1963).

 66.  Tsitovich, I. K., "Concentration of Ions on Chromatographic Grade
      Alumina for Qualitative Microanalysis," Tr. Kubansk.  Sel'skokhoz Inst.
      19P** (9)) 20*1-8.

 67.  Zeitoun, M. A., Davison, R. R., White, F. B., and Hood,  D. W.f  "Solvent
      Extraction of Secondary Waste Water Effluents: Heterogeneous Equilibrium

                                         11 J> Water ponution  Contro1 Federa-
 68.  Zawodna, B., "The Effect of Activated Carbon on the Determination of
      Nitrogen Compounds in Sewage," Przemysl Spozywczy lU, U30-2 (1960).
 B.   WATER RECLAMATION
 Conventional and Tertiary Treatment  (See Also References A-29 thru A-55)


                                             Feasible»" Chem' and En«- News
2.   Cooper, R. B.,  "The Treatment of Waste Water Containing Inorganic
     Materials," Aust. Chern. Process. Eng. lg (10), 36-8, kO-1 (1966).
3.   Gulp, R. L., "Wastewater Reclamation by Tertiary Treatment," J. Water
     Pollution Control Federation 3£, 799-806 (1963) .

k.   Dietrich, K. R., "The Third Step in the Purification of Waste Waters
     to Prevent Eutrophy of Lakes," Chemiker. Ztg. 8? (21), 772-5 (1963).

5.   Eckenfelder, W. W. Jr., and O'Conner, D. J., Biological Waste Treatment.
     Permagon Press, New York, New York, 1961.    - -        '
6>   £Sf8' °* V*' Water greatMnt; A Guide to the Treatment of Water  and
     Effluents Purification. 3rd Ed. London, Technical Press, 1965. -

7.   Malhotra, S. K. , "Nutrient Removal from Secondary Effluent by Alum

                                                                 Ho-  6
                                   50

-------
B.   WATER RECLAMATION

Conventional and Tertiary Treatment (Continued)

8.   Marshall, J. R., "Today's Wastes: Tomorrow's Drinking Water?",  Chem.
     Eng. 6_2, No. 16, 107-10 (1962).

9.   Powell, 3. T., "What Part Does Chemical Coagulation Play in Today's
     Water Treatment Practices?" Power 9_8, No. 1, 80-2, 216,  218, 220 (195*0.

10.  Rand, M. C., "General Principles of Chemical Coagulation," Sewage and
     Ind. Wastes 31, 863-71 (1959).

11.  Stephan, D. G., "Water Renovation, Current Status of the Technology,"
     Proc. Southern Water Resources Pollution Control Conf. lU, 113-22 (1965).

12.  van Vuuren, L.R.J., Stander, G. J., Henzen, M. R., Me i ring, P.G.J., and
     van Blerk, S.H.V., "Advanced Purification of Sewage Works Effluent Using
     a Combined System of Lime Softening and Flotation," Water Research 1,
          U (1967).
Iron Salts in Water Treatment

13.  Dean, R. B., "Ultimate Disposal of Wastewater Concentrates to the Environ-
     ment," Envir. Sci. Technol., in press.

lU.  Dobrynin, F. T., "The Purification of Water with Iron Vitriol,"
     Vodosnabshenie i Sanit. Tekh. 16, No. 5, 50-3 (19^1).

15.  Dodonov, Ya. Ya., and Plekhanova, T. G., "FeSOij. as a Coagulating Agent
     in the Purification of Water," Vodosnabzhenie Sanit. Tekh. IgUO, No.  12,
     UO-3; Khim. Referat. Zhur. U, No. 7-8, 98 (19^1).

16.  Kunzel-Mehner, A., "Ferric Chloride in the Chemical and Mechanical
     Purification of Water," Chem. Tech. 15, 129-35
17.  Mehner, A. K., "Chemical-Mechanical Treatment of Water with Ferric
     Chloride," Chem. Tech. 15, 129 (19^2).

18.  Pirnie, M., "Some Operating Experiences with Iron and Iron Coagulants  in
     Water Treatment," J. New Engl. Water Works Assoc. 5_1, U37-53 (1937).

19.  Schworm, W. B., "Iron Salts for Water and Waste Treatment," Public Works
     2jt (10), 118-20 (1963).

20.  Scouller, W. D., "Effect of Iron on Sewage Purification,"  Surveyor 82,
         (1932).
                                   51

-------
 B.   WATER RECLAMATION

 Iron Salts in Water Treatment (Continued)

 21.  Simmons, P. D., "Ferrous Sulfate as a Coagulant," Proc. 12th Ann. Conf
      Water Purif., in W. Va. Univ., Eng. Expt. Sta. Tech. Bull. No. 11,
      21-3 (1938).


 22.  Streander, P. B., "Sewage Treatment with Ferrous Sulfate and Aeration "
      Public Works 6J>, No. 3, 29 (1933).                                   '

 23.  Vadyukhim, I. I., "Coagulation of Water with Ferrous Sulfate in Com-
      bination with Chlorine," Vodosnabzhenie i Sanit. Tekh.  lU, No. 10
      35-^3 (1939).                                                    '

 2U.  Wolman, A., "The Role of Iron in the Activated Sludge Process " Ens
      News Rec.  £8, 202-1* (1927).                                  '   *'

 25.  Zhuchkova, A. M.,  "Coagulation of Water with Ferrous Sulfate " Teplo-
      Silovoc Khoz. 1^32, No. 7,  51-2;  Khim Referat.  Zhur. 1939. No.  12, 86-7.

 Carbon in Water Treatment (See Also Reference A-U3,  -U6,  -58,  -62, -65,  -68)

 26.  Johnson, R. L.,  Lowes,  F. J.  Jr.,  Smith,  R.  M.,  and  Powers,  T.  J.,
       Evaluation of the  Use  of Activated Carbons  and Chemical  Regenerants in
      Treatment  of Waste  Water,"  AWTR-11,  U.S.  Department  of Health,  Education
      and Welfare,  Public Health  Service Publication No. 999-WP-13, Jfey, 196^.

 27.   Joyce, R.  S.,  and Sukenik,  V.  A.,  "Feasibility of Granular Activated-
      Carbon Adsorption for Waste-Water  Renovation, 2," AWTR-15, U. S
     Department  of Health, Education and Welfare, Public Health Service
     Publication No.  999-WP-28,  October, 1965.

 28.  McGlasson,  W.  G., Thibodeaux,  L. T., and Berger, H. F., "Potential Use
                  Carbon for Waste  Water Renovation," Tappi to  (12), 521-6
29.  Reissaus, K., and Rummel, W., "Use of Activated Carbon in Water Purifi-
     cation,  Wasserwlrt.-Wassertech. 16 (12), U13-16 (1966).

30.  Shane. M. S., "How to Black Out Algae," Water Works Eng. 116 (7)  552-T
     (1963).                                                  -    '     J



C.   REDUCING AGENTS FOR NITRATE



Ferrous Salts (See Also Reference A-U8,  C-79)


1.   Abel, E., "Kinetics of the Oxidation of Ferrous Ion by Nitric Acid "
     Monatsh. 68, 387-93 (1936).                                       '

-------
C.   REDUCING AGENTS FOR NITRATE

Ferrous Salts (Continued)

2.   Abel, E., Schnid. H., and Pollak, F., "Kinetics of the Oxidation of
     Ferrous Ions by Nitrous Acid," Monatsh.  6_Q,  125-1*3 (1936).

3.   Baudisch, 0., and Meyer, P., "The Reduction of Nitrites and  Nitrates,"
     Biochem. Z. 107. 1-^2 (1920).

I*.   Benner, J. M., and Shaw, K., "Reduction of Nitrate by Ferrous Hydroxide
     Under Various Conditions of Alkalinity," Analyst §0, 626-7  (1955).

5.   Brown, L. L., and Drury, J. S., "Nitrogen Isotope Effects in the Reduction
     of Nitrate, Nitrite, and Hydroxylamine to Ammonia.  I. In Sodium Hydroxide
     Solution with Fe(ll)."  J. Chem. Phys. U6 (7), 2833-7 (1967).

6.   Carsley, S. H., "The Reduction of Alkali Nitrates by Hydrous Ferrous
     Oxide," J. Phys. Chem. 3Jt, 176-87 (1930).

7.   Gottlieb, 0. H., and Magalhaes, M. T., "The Volumetric Determination of
     Nitrate Ions," Anal. Chem. 50, 995-7 (1953).

8.   Karstein, P., and Grabe, C.A.J., "Determination of Nitrate According to
     Cotte and Kahane," Chem. Weekblad UU, 237-8 (19^8).

9.   Kolthoff, I. M., Sandell, E. B., and Moskovitz, B., "Volumetric Deter-
     mination of Nitrates with Ferrous Sulfate as Reducing Agent," J. Am.
     Chem. Soc. 55, lU5**-7 (1933).

10.  Krejci, F., and Kacetl, L., "Determination of Nitrate by Titration with
     Ferrous Sulfate," Chem. and Ind. (London) 1957. 598-

11.  Laccetti, M. A., Semel, S., and Roth, M., "Colorimetric Determination of
     Organic Nitrates and Nitramines," Anal. Chem. 31, 10^9-50 (1959).

12.  Miyamoto, S., "The Reducing Action of Ferrous Hydroxide," Japan J.  Chem.
     1, 57-80 (1922).

13.  Murakami, T., "Rapid Volumetric Determination of Nitric Acid or Nitrate
     by Reduction with Ferrous Salt," Japan Analyst U, 630-3 (1955).

I1*.  Murakami, T., "Photometric Determination of Nitrite by Using FerroUs
     Sulfate and Phosphoric Acid," Kagyo Kagaku Zasshi 63, 1295-8 (1960).

15.  Pappenhagen, J. M., and Looker. J. J., "Suggested Reduction Methods for
     the Determination of Nitrates/1 J. Am. Water Works Assoc. £1, 1039-^5
     (1959)•
                                   53

-------
 C.   REDUCING AGENTS FOR NITRATE

 Ferrous Stilts (Continued)

 16.  Sandonnini, C., and Bezzi,  S.,  "Reduction of Nitrates with Ferrous
      Hydroxide," Gazz.  Chim.  Ital. 6j>,  693-700 (1930).

 17.  Schoer, E., "Kinetics and Mechanism of the Reaction Between the Ferrous
      Ion and Nitrous and Nitric  Acids," Z.  Physik. Chem. A176. 20-1*7 (1936).

 18.  Szabo,  Z. G.,  and  Bartha, L., "Volumetric Determination of Very %aH
      Quantities of  Nitrate,"  Mikrochemie ver.  Mikochim. Acta 38, 1*13-18
      (1951).                                                —

 19-  Szabo,  Z. G.,  and  Bartha, L., "A New Titrimetric Method for the Deter-
      mination of Nitrate Ion," Anal. Chim. Acta £, 33-Uj (1951).

 20.  Szabo,  Z. G.,  and  Bartha, L., "Catalysis  in Analytical Chemistry.
      I.  Silver Catalysts in the  Reduction of Nitrates by Ferrous Hydroxide "
      Acta Chim.  Huag 1,  116-23 (1951).

 21.  Szabo,  Z. G.,  and  Bartha, L., "Alkalimetric Determination of the Nitrate
      Ion by  Means of a  Copper-Catalyzed Reduction," Anal. Chim. Acta 6, U16-1Q
      (1952).

 22.  Chao, Tyng Tsair,  and Kroontje, Wybe, "Inorganic Nitrogen Transforma-
      tions Through  the Oxidation and Reduction of Iron," Soil Sci.  Am
      Proc. 30 (2),  193-6 (1966).

 23.  Young,  G. K.,  Bungay, H.  R., Brown, L. M., and Parsons,  W. A., "Chemical
      Reduction of Nitrate in Water," J. Water Pollution Control Federation
      36,  395-8 (19&).

 Carbon   .

 21*.   Bylo, Z., and Panek, M.,  "The Influence of Oxidation on  the Reaction of
      Hard Coals with Dilute Solutions of Nitric Acid," Arch.  Gornictwa 9 (k)
      383-97  (196U).                                                     ^ ^  ''

 25.  Donnet, J. B.,  and Lahaye, J., "Oxidation of Carbon Black by Nitric  Acid.
      I. Mode of C02 Formation: Kinetic  Aspect of  the Reaction,"  Bun.  Soc
     Chim. France 1966 (If), 1282-5.

 26.  Farrell, J. B., and Haas, P. A., "Oxidation  of  Nuclear-Grade Graphite by
     Nitric Acid and Oxygen," Ind. Eng.  Chem.,  Process Des. Develop. 6  (3)
     277-81 (1967).                                                  ~    '

27.  van Krevelen, D. W., "A Gas Rich in Nitric Oxide by Reduction  of Nitric
     Acid with Carbon,"  Brit.  66l, 90U,  Nov. 28, 1951.

-------
C.   REDUCING AGENTS FOR NITRATE

Carbon (Continued)

28.  Larina, N. K., Khalmukhamedova, R. A., and Tadzhiev,  A.  T.,  "Products
     of Oxidation of the Angrensk Brown Coals by Nitric Acid," Khim.  Klassi-
     fikatsiya Iskop. Uglei, Akad. Nauk. SSSR, Inst.  Goryuch. Iskop.  1966
     98-107.

Sulfur Dioxide

29.  Ivin, K. J., "Reaction of Nitrates with Liquid Sulfur Dioxide,"  Nature
     180. 90 (1957).

30.  Smedslund, T. H., "Continuous Preparation of Nitric Oxide from Nitric
     Acid and S02," Finska Kemist-Samfundets Medd. 5£,  37-9 (1950).

31.  Soibelman, B. J., and Bresler, F., "Detection of Nitrates in Presence of
     Interfering Anions," Zavodskaya Lab. £, 359-60 (19^0).

32.  Veprek-Siska, J., and Uher, L., "Reduction of the  Nitric Acid by Means of
     Sulfur Dioxide," Collection Czech. Chem. Coramun. 31 (11), U363-71*  (1966).

33.  Zeegers, R.N.G., "Hydroxylamine Compounds," U.S. 2,555,667,  June 5, 1951.

Formaldehyde

35.  Adams, W. H., Fowler, E. B., and Christenson, C. W.,  "A  Method for
     Treating Radioactive Nitric Acid Wastes Using Paraformaldehyde," Ind.
     Eng. Chem. 52, 55-6 (1960).

36.  Cultrera, R., and Farrari, E., "Research on the  Photochemical Reduction
     of Nitrate," Ann. Chim (Rome) U7, 1321-36 (1957);  W, lUlO-25 (1958);
     1*2, 176-82 (1959).

37-  Evans, T. F., "Pilot Plant Denitration of Purex  Wastes with  Formalde-
     hyde," U.S. At. Energy Comm. HW-58587 (1959).

38.  Forsman, R. C., and Oberg, G. C., "HCHO Treatment  of  Purex Radioactive
     Waste," U.S. At. Energy Comm. HW-79622 (1963).

39.  Halliday, H. M., and Reade, T. H., "Action of Nitrous Acid on Formal-
     dehyde, " J. Chem. Soc. 19^0. lte-3.

UO.  Healy, T. V., "The Reaction of Nitric Acid with  Formaldehyde and with
     Formic Acid and its Application to the Removal of  Nitric Acid from
     Mixtures," J. Appl. Chem. 8, 553-61 (1958).

Ul.  Healy, T. V., "Concentration of Aqueous Metal Salt Solutions Containing
     Nitric Acid," U.S. 2,835,555, May 20, 1958.
                                   55

-------
 C.   REDUCING AGENTS FOR NITRATE

 Formaldehyde (Continued)

 U2.  Kourim, V., and Konecny,  C.,  "Decomposition of Nitric Acid by Formal-
      dehyde," Chem.  Listy 51,  1376-7 (1957).

 J*3.  Morris, J.  B.,  "The Reaction  of Nitric Acid with Formaldehyde " Eneraie
      Nucleaire 1, 216-2U (1957).                                         e

 UU.  Nemtsov, M. S., and Trenke, K.  M.,  "ape stigat ions in the Field of Acid
      Catalysis.   I.  Kinetics and Mechanism of the Reactions of Formaldehyde
      in Acid Aqueous Solutions," Zhur. Obshchei  Khim. 22, U15-29 (1952).

 U5.  Nemtsov, M. S., and Trenke, K.  M.,  "Investigations in the Field of Acid
      Catalysts.   I.  The  Kinetics and Mechanism of the Reactions of Formal-
      dehyde in Acid  Aqueous Solutions,"  J. Gen.  Chem. USSR 22, ^85-96 (1952).

 U6.  Shtol'ts, A. K.,  "Reaction of Nitric Acid with Formaldehyde, Rongalite
      and Hydrosulfite,"  Izvest. Vysshikh Ucheb.  Zavedenii, Khim. i Khim.


 !*7.  Vanino,  L.,  and Schinner, A., "The  Reaction Between Formaldehyde and
      Nitrous  Acid,"  Z. Anal. Chem. 5J2, 21-6 (1913).

 Sugars  (See Also  Reference C-36)

 U8.  Bray,  L. A.,  and Martin, E. C.,  "Invention Report - Use of Sugar to
      Neutralize Nitric Acid Waste Liquors," U.S. At. Energy Conm.  HW-75565
      (1962) .

 kg.   Bray,  L. A.,  and Martin, E. C.,  "Removal of Nitric Acid and of Nitrite
      and Nitrate  Ions from Radioactive Waste," U.S. 3,158,577, Nov. 2k,
50.  Breit Schneider, R., and Kopriva, B., "Oxidation of Sucrose with Nitric
     Acid," Listy Cukrovar. 82 (9), 215-20 (1966).

51.  Coppinger, F. A., "Pilot Plant Denitration of Purex Water with Sugar "
     AEC Accession No. 352U3, Rept. No. HW-77080, Avail. OTS (1963).

52.  Justat, A., Gorzka, Z., and Janio, K., "Oxidation of D-Glucose with
     Nitric Acid to Oxalic Acid," Chem. Stosowana 7 (3), 1*09-11* (1963).

53.  Kopriva, B., Markova, J., and Breit Schneider, R.,  "Oxidation  of  Sucrose
     with Nitric Acid.  II. Kinetics of Oxidation," Listy Cukrov 83 (2)
     36-9 (1967).                                               ~    '

5^.  Lesquibe, F., "Degradation of Glucose by Oxidation with Aqueous Acid
     (HN03) " J. Rech. Centre Natl. Rech.  Sci.  Lab.  Bellevue (Paris) Ik (62)
     33-71 (1963).                                                  —    "


                                   56

-------
 C.    REDUCING AGENTS FOR NITRATE

 Sugars  (Continued)

 55.   Soltzberg,  S., "Tartaric Acid," U.S. 2,360,196, July 10, 191*5.

 56.   Tang, Teng-Han, and Kao, F. C., "Preparation of Oxalic Acid I," J. Chem.
      Eng. China  16, 32  (1939).

 Powdered Iron

 57.   Babson,  J. A., Burch, W. G. Jr., and Woodis, T. C. Jr., "Critical
      Evaluation  of the  Reduced Iron Method for Reduction of Nitrate,"
      J. Assoc. Offic. Agr. Chemists U6 (U), 599-603 (1963).

 58.   Delius,  I., "Removal of Nitrates from Drinking Water," Gesundheits-Ing.
      00, 181  (1959).

 59.   Gehrke,  C. W., and Johnson, F. J., "Efficiency of Various Iron Powders
      in Seducing Nitrate," J. Assoc. Offic. Agr. Chemists J4£, U6-9 (1962).

 60.   Travers, A., and Diebold, R., "The Action of Nitric Acid on Iron and
      Iron Carbide (Fe_C)," Bull Soc. Chim. (5), £> 690-3 (1938).

 6l.   Vetter,  K. J., "The Active State and the Spontaneous Repassivation of
      Current-Activated  Iron in Nitric Acid," Z. Electrochem. j>6, 106-15 (1952).

 Hydrazine and its Salts

 62.   Bursa, S., and Straszko, J., "Eudiometric Determination of Nitrate Ion in
     Aqueous Nitric Acid," Chem. Anal.  8, 29-Uo (1963).

 63.  Davies, A. W., and Taylor, K., "Application of the Auto-Analyzer in a
      River Authority Laboratory," Technicon Symp., 2nd, N.Y., London 1965,
     29U-300  (Pub. 1966).                                            ^^

 6k.  Dey, B. B., and Sen,  H. K., "Action of Hydrazine Sulfate Upon Nitrites
     and a New Method for Determining Nitrogen in Nitrites," Z.  Anorg.  Chem.
     71, 236-U2 (1911).

65.  Dzhardamalieva, K.  K.,  "Catalytic  Reduction with flydrazine,"  Tr.  Inst.
     Khim.  Nauk, Akad. Nauk Kaz.USSH 8,  150-6 (1962).

66.  Kahn,  L., and Brezenski, F.  T.,  "Determination of Nitrate  in  Estuarine
     Waters.  Comparison of a Hydrazine Reduction and a Brucine  Procedure,
     and Modification of a Brucine Procedure," Environ.  Sci.  Technol. 1
     (6),  U88-91 (1967).
                                   57

-------
 C.   REDUCING AGENTS FOR NITRATE

 Hydrazine and its Salts (Continued)

 67.  Kamphake, L. J., Hannah, S. A., and Cohen,  J.  M.,  "Automated Analysis
      for Nitrate by Hydrazine Reduction," Water  Res.  1  (3),  205-16  (1967).

 68.  Koltunov, V. S., Nikol'akii, V. A., and Azureev, Yu. P.,  "Kinetics of
      Hydrazine Oxidation in the Presence of HNOs in an  Aqueous Solution "
      Kinetiki i Kataliz. 3 (6), 077-81 (1962). *              ^uwon,

 69.  Mullin, J. B., and Riley, J. P., "The Spectrophotometric  Determination
      of Nitrate in Natural Waters, with Particular  Reference to Seawater "
      Anal. Chem. Acta 12,  k6k-Qo (1955).

 Miscellaneous (Salts. Metals, etc.)

 70.  Banerjee,  P. J., "Vanadous Sulfate as a Reducing Agent.   II. Estimation
      of Chlorates, Nitrates and Persulfates," J.  Indian Chem.  Soc. 13, 301-U
      (1936).                                                      —

 71.  Bartow, E., and Rogers,  J. S.,  "Determination  of Nitrates by Reduction
      with Aluminum," Univ.  111.  Bull.,  W.  S. Series 7,  lU-27 (1910).

 72.  Fletcher,  J. M., and Woodhead,  J.  L.,  "The Reaction of Ruthenium (ill)
      with Nitric Acid,"  J.  Inorg.  Nucl.  Chem. 27  (?), 1517-19  (1965).

 73.  Frank,  J.  A.,  and Spence,  J.  T., "The Reduction of Nitrite by Mo (V)."
      J.  Phys. Chem.  68 (8), 2131-5 (196U).

 7b.  Gasser, J.K.R.,  "Substitute Reagent for Titanous Sulfate for Reducing
      Nitrate Nitrogen," Analyst 88,  237-8  (1963).

 75.   Guymon, E.  Park,  and Spence, J. T., "The Reduction of Nitrate by Mo  (V)  "
      J.  Phys. Chem.  70 (6), 196^-9 (1966).

 76.  Haight, G.  P. Jr., Mohilner, P., and Katz, A.,  "The Mechanism of the
      Reduction of Nitrate.  I. Stoichiometry of Molybdate-Catalyzed Reductions
      of Nitrate  and Nitrite with Sn  (II) in Hydrochloric and Sulfuric Acids "
     Acta Chem.  Scand. 16, 221-8 (1962).

77.  Haight, G. P. Jr., and Katz, A., "The Mechanism of  Reduction of Nitrate.
     II. The Kinetics and Mechanism of the Molybdate-Catalyzed  Reduction of
     Nitrate by Sn (II) in Acid Solution," Acta Chem.  Scand.  16,  659-72-(1962).

78.  Kasbekar, G. S., and Nonnand, A. R., "Reaction  Between Nitric Acid and Tin
     in Presence of Catalysts. II," Proc. Indian  Acad. Sci. 1QA.  37.^0 (1939).

79-  Milligan, L. H., and Gillette, G. R.,  "The Reduction of  Free Nitric Acid
     by Means of Ferrous, Stannous or Titanous Salts," J. Phys. Chem. 28.
        -                                                            —

                                   58

-------
C.   REDUCING AGENTS FOR NITRATE

Miscellaneous (Continued)

GO.  Murakami, T., "Volumetric Determination of Nitric Acid and Nitrate by
     Reduction with Stannous Chloride," Bunseki Kagaku 7, 766-71 (1558).

81.  Pozsi-Escot, E., "The Determination of Nitric Nitrogen by Reduction
     with the Aid of Aluminum-Mercury," Compt. Rend., lU£, 1380 (1910).

82.  Pozzi-Escot, E., "The Reduction of Nitrates to Ammonia and a New Method
     of Determining Nitrates," Ann Chim. Anal., lU, UU5-6 (1910).

83.  Stammer, K., ["Reaction of HN03 and CO?" - Actual Title Unknown]
     Pogg. Ann. 82, 137 (1851).  Cited in Gmelins Handbuch der anorganischen
     Chemie, 8th Edition, Syst. No. U, p. 1006 (1955).

8U.  Thomas, M., "Total Determination of the Nitric Nitrogen and of the
     Nitrogen of Nitrated Groups by Titanium Chloride and Gravimetric
     Determination of a Nitrate Derivative in a Mixture with a Nitrate,"
     Mem. Poudres 3Jt, 357-6? (1952).

85.  Wood, E. D., Armstrong, F.A.J., and Richards, F. A., "Determination of
     Nitrate in Sea Water by Cadmium-Copper Reduction to Nitrite," J. Mar.
     Biol. Ass. U.K. V? (1), 23-31 (1967).


D.   DEAMMINATION AGENTS


Sulfamic Acid

1.   Baumgarten, P., "The Effect of Nitric Acid upon Sulfamic Acid.   A
     Simple Method for the Preparation of Nitrous Oxide," Ber. JIB,
     80-1 (3.938).

2.   Baumgarten, P., and Marggraff, I., "The Reaction of Nitrites with
     Amino sulfonic Acids and the Detection and Estimation of Nitrous Acid
     in the Presence of Nitric Acid," Ber. 63, 1019-2U (1930).

3.   Brasted, R. C., "Detection of Nitrite and Sulfamate Ions in Qualitative
     Analysis," J. Chem. Education 28, 592-3 (1951).

U.   Brasted, R. C., "Reaction of Sodium Nitrite and Sulfamic Acid," Anal.
     Chem. 2k, 1111-lU (1952).

5.   Carson, W. N. Jr., "Gasometric Determination of Nitrite and Sulfamate,"
     Anal. Chem. 23, 1016-19 (1951).

6.   Wu, Ching-Hsien, and Hepler, L. G., "Thermochemistry of Sulfamic Acid  and
     Aqueous Sulfamate Ion," J. Chem. Eng. Data 7, Pt. 1, 536-7 (1962).


                                   59

-------
 D.   DEAMMINATION AGENTS

 Sulfamic Acid (Continued)

 7.   Gumming, W. M., and Alexander, W. A., "Use of Aminosulfonic Acid in the
      Determination of Nitrites," Analyst 68,  273-1* (19^3).

 8.   Groh, H. J. Jr., and Russell,  E.  R., "Intermediates Formed in the Reaction
      of Nitrite with Salts of Sulfamic Acid," J.  Inorg.  Nucl.  Chem.  26 (1*).
      665-7 (19610.                                                  —

 9.   Gottfried, J., and Novak, Jiri V. A., "Polarographic Determination of
      Amidosulfonic Acid and Nitrate,"  Chem. Prunysl 7, 1*76-8 (1957).

 10.   Heubel,  J., and Canis, C., "Reaction Between Nitrates and Sulfamates,"
      Compt. Rend. 255. 708-10 (1962).

 11.   Heubel,  J., and Wartel, M.,  "The  Reaction Between Nitrites and Amino
      Sulfonates," Compt.  Rend. 257  (3),  68U-6 (1963).

 12.   Kaloumenos, H. W.,  "Specific Determination of Nitrite," Werkstoffe u.
      Korrosion n,  626 (1960).

 13.   Kostrikin,  Yu M., and Mikhailova, N.  M.,  "Treatment  of Heating Water "
      USSR 139,997,  Appl.  Oct.  8,  1960.

 1U.   Subrahmanyan,  P.V.R.,  Sastry,  C. A.,  and  Pillar, S.  C., "Determination
      of  the Permanganate  Value for  Waters  and  Sewage Effluents Containing
      Nitrite," Analyst 8J*,  731-5  (1959).

 Urea  (See  Also References D-lU,E-5)

 15.  Asendorf, E.,  "Method  of  Decontaminating Aqueous Solutions of Nitrites
      of Alkali Metals  and/or Alkaline Earth Metals," B.P. 1,028,161 (Assigned
      to Water Engineering Limited), 1* May 1966.

 16.  Bonner, W. D., and Bishop, E.  S., "The Hate of Reaction of Nitrous Acid
     and Urea in Dilute Solutions," J. Ind. Eng. Chem.,  £, 13^-6 (1913).

 17.  Burriel, F., and  Suarez Acosta, K., "Analytical Problem of Separating
     Nitrates and Nitrites.  IV. Destruction of Nitrites with Urea and its
     Derivatives," Anales Heal Soc.  Espan. Fis. y. Quim.  U6_B, 1*29-1*0 (1950).

 18.  Gorenbein, E. Ya, and Sukhan, V. V., "Interaction of Urea  with Nitric
     Acid in an Aqueous Medium," Zh. Neorgan.  Khim. 10 (7),  1701-5 (1965).

19.  Sabbe, W. E., and Reed, L. W.,  "Investigation Concerning Nitrogen Loss
     Through Chemical Reactions Involving Urea and Nitrite," Soil Sci.  Soc
     Am.  Proc. 28 (1*), 1*78-81 (196!*).

                                  60

-------
D.   DEAMMINATION AGENTS

Urea (Continued)

20.  Shaw, W.H.R., and Bordeaux, J. J.,  "The Decomposition of Urea in Aqueous
     Media," J. Am. Chem. Soc. 77, u729-33 (1955).

21.  Weston, C. F., "Removing Nitrous Acid from Solutions Such as  Those of
     Sodium Nitrate," U.S. 2,139,11*2, Dec. 6, 1939.

Amino Acids

22.  Austin, A. T., "Deammination of Amino Acids by  Nitrous Acid with Par-
     ticular Reference to Glycine.  The  Chemistry Underlying  the Van  Slyke
     Determination of a-Amino Acids," J. Chem.  Soc.  1950, 1U9-57.

23.  Cristol, P., Benezech, C., and Lissitsky,  S., "Deammination by Nitrous
     Acid.  I. Rate Constant of Deammination of Amino Acids in Aqueous
     Solution," Bull. Soc. Chim. Biol.  31, 150-6 (19^9).

2k.  Cristol, P., Benezech, C., and Lissitsky,  S., "Deammination by Nitrous
     Acid.  II. Influence of Iodine on  the Rate of Deammination of Amino
     Acids.  Deammination of Mixtures of Amino  Acids,"  Bull.  Soc.  Chim.
     Biol. 31, 156-60 (19^9).

Miscellaneous (Ammonia, Amines, Azides)

25.  Adamson, D. W., and Kenner, J., "Decomposition  of  the Nitrites of Some
     Primary Aliphatic Amines," J. Chem. Soc. 193U,  838-M*.

26.  Burriel, F., and Saurez, R., "Analytical Problem of  Separating Nitrates
     and Nitrites.  III. Destruction of  Nitrites with Ammonium Salts,"
     Anales Real Soc. Espan. Fis. y Quim. U5B,  893-910  (19U9).

27.  Kezdy, F. J., Jaz, J., and Bruylants, A.,  "The  Kinetics  of the Effect
     of Nitrous Acid on Amides.  I. General Method," Bull.  Soc. Chim.
     Beiges 6_7, 687-706 (1958).

28.  Huckel, W., and Wilip, E., "The Conversion of Amines with Nitrous Acid,"
     J. Prakt. Chem. 158. 21-32 (19U1).

29.  Mohrig, J. R., "The Synthesis and Nitrous  Acid  Deammination of Some
     Bicyclic Amines.  The Mechanism of  the Deammination  Reaction," Univ.
     Microfilms, Order No. 6U-U369; Dissertation Abstr. 25 (2), 8U2 (196^).

30.  Ridd, J. H., "Nitrosation, Diazotization and Deammination," Quart. Rev.
     (London) 15_ (U), Ul8-»H (1961).

31.  Seel, F., Wolfle, R., and Zwarg, G., "Kinetics  of  the Decomposition of
     Nitrous Acid with Hydrazoic Acid,"  Z. Naturforsch. 13b,  136-7 (1958).

                                   61

-------
 D.    DEAMMINATION AGENTS

 Miscellaneous (Continued)

 32.   Spence,  L.  U., Whitmore, F. C., and Suraatis, J. D., "Action of Methyl-
      amine with  Nitrous  Acid," J. Am. Chem. Soc. 63, 1771 (19!*!).

 33.   Stedman, G., "Mechanism of the Azide-Nitrite Reaction.  Pt. 1." J. Chem.
      Soc.  1959.  29U3-9-

 31*.   Streitwieser, A. Jr.,  "Reaction of Aliphatic Primary Amines with Nitrous
      Acid," J. Org. Chem. 22, 86l-9 (1957).


 E.    CATALYSIS (See Also Individual Entries in Section D; e.g. D-20, -21)


 1.    Azim,  M. A., and Saraf, S. D., "Catalytic Decomposition of Nitrous Acid."
      J. Indian Chem. Soc. 33, 763-1* (1956).

 2.    Azim,  M. A., and Shafi, M., "Kinetics of Nitric Acid Decomposition in
      Liquid Phase," J. Nat. Sci. Math. 5 (2), 223-6 (1965).

 3.    Catalina, L., "Vanadium and Reduction of Nitrates in Plants.  III.  Acti-
      vation of the Nitrates in Presence of Vanadium," An. Edafol. Agrobiol.
      (Madrid) £2 (11-12), 731-5 (1966).

 U.    Halpem, J., "The Catalytic Activation of Hydrogen in Homogeneous,
      Heterogeneous and Biological Systems," Advances in Catalysis. Volume XI,
      Academic Press, Inc., New York, 1959,  pp. 309-11.

 5.    Quartaroli, A., "The Kinetics of Febrile Reactions.  Contribution to
      the Study of Autocatalysis," Gazz Chim. Ital.  53, 31*5-68 (1923).

 6.    Suzawa, T., "Decomposition of Nitrous  Acid in Aqueous Solution,"
      Kagaku to Kogyo (Osaka) 31, 55-60 (1957).


F.   ANALYTICAL METHODS


1.   American Public Health Association, American Water Works Association
     and Water Pollution Control Federation, Standard Methods for the
     Examination of Water and Waste Water  , Publication Office, American
     Public Health Association,  New York, N.Y.,  1965  (12th edition),
     pp. 166-208.

2.   Armstrong, F.A.J.,  "Determination of Nitrate in  Water by Ultraviolet
     Spectrophotometry,"  Anal. Chem.  35, 1292-U  (1963).
                                   62

-------
F.   ANALYTICAL METHODS (Continued)
3.   Bastion, R., Weberling, R., and. Palilla,  P.,  "Ultraviolet Spectro-
     photometric Determination of Nitrate," Anal.  Chem.  2£,  1795-7 (1957).

U.   Dukes, E. K., and Wanace, R. M., "Stability  of Ferrous Sulfamate in
     Nitric Acid Solutions," Contract No.  AT/07-2/1, E.  I. Dupont  DeNemours
     and Co., Savannah, Feb.
5.   Feigl, F., Spot Tests in Inorganic Analysis,  Elsevier  Publishing Co.,
     New York, N.Y., 1956, pp. 32$ »7

6.   Fisher, F. L., Ibert, E. R., and Beckman,  H.  F.,  "Inorganic Nitrate,
     Nitrite, or Nitrate-Nitrite," Anal. Chem.  30, 1972-U (1958).

7.   Snell, F. D., and Snell, C.  T.,  Colorimetric  Methods of Analysis.  IV.
     D. Van Nostrand Company, Inc., Princeton,  N.  J.,  195^,  pp. 317.

8.   Walton, H. F., Principles and Methods of Chemical Analysis. 2nd Edition,
     Prentiss-Hall, Inc., Englewood,  N. J., 1964,  pp.  327-28.

9.   Watt, G. W., and Chrisp, J.  D.,  "Spectrophotometric Method for the
     Determination of Urea," Anal. Chem. 26,  1*52-3 (195*0 .
                                  63

-------
                               APPENDIX


ANALYTICAL PROCEDURE FOR THE SIMULTANEOUS DETERMINATION OF NITRATE AND
NITRITE IONS


     Nitrate and nitrite can be simultaneously determined from a solution con-
taining insoluble matter, either with or without a deammination agent present,
using ultraviolet spectroscopic methods.

Apparatus and Reagents

     Gary lU recording spectrophotometer with a set of matched 1 cm silica
cells.

     Standard volumetric glassware.

     70-72$ Perchloric acidj Baker "analyzed" grade.

     Sulfamic acid; 99^ Eastman white label.

     Sodium nitrate; Baker "analyzed" grade.

     Sodium nitrite; Baker "analysed" grade.

     Sodium nitrate stock solution; prepare by dissolving 0.6071 gm of sodium
nitrate in 1 liter of distilled water to form a stock solution of 100 ppm
nitrogen.

     Sodium nitrite stock solution; prepare by dissolving 0.^929 gra °f sodium
nitrite in 1 liter of distilled water to form a stock solution of 100 ppm
nitrogen.

Preparation of Calibration Curves

     Make up a set of six standard solutions of mixtures of  sodium nitrate
and  sodium nitrite in distilled water such that each solution contains a
total of 50 ptm nitrogen.  Also make up a blank.  Prepare these solutions by
diluting the following volumes of stock nitrate and nitrite  solutions to 50 ml:

     1.  0.00 ml N03               25.00 ml NOg

     2.  5.00 ml NO"               20.00 ml NOg

     3.  10.00 ml NO"              15-00 ml N0~

-------
       U.   20.00 ml NO"              10.00 ml NO"


       5.   20.00 ml NO"              5.00  ml NO"


       6.   25.00 ml NO"              o.OO  ml NOl
                      •5                       2

       The  blank should contain only  distilled water.  The concentrations of
  the  above solutions  are as follows:


       1.   0 ppm NO~-N               50 ppm NO"-N


       2.   10 ppm NO~-N              kO ppm NO"-N


       3.   20 ppm NO^-N              30 ppm NOg-N


      ^.   30 ppm NO~-N             20 ppm NOg-N


      5.  UO ppm NO"-N              10 ppm NO~-N


      6.  50 ppm NO~-N              o ppm NO~-N


      The blank is zero in both.
    ™?\0f each standard into clea* 100 ml volumetric  flasks  and
   -90 ml of distilled water.   Do not dilute to the mark.


      Check the matched silica cells for cleanliness by scanning from  250-1Q.5

 Sii^SrV^J distin:d water in b°th cel^.   If the baseline varies more
 than 0.005 absorbance units,  reclean the cells.
 +   *v,T° th! fi£!* s°J;ution»  add k drops of 70-72$ perchloric acid and dilute
 to the mark.   This should lower the pH to 2.  Immediately scan from 2^0-igl

 millimicrons using distilled water as a reference.  Read at 200 millimicrons

 and call this  absorbance  A.  Add a threefold excess (approximately U mg) of

 ^i^nf11^0 aCld t0 the flask"  m* wen "4 scan this solution from
 250-195 millimicrons.  Read  this at 200 millimicrons and call this absorbance
 B.   Repeat  for each standard solution and the blank.                 soroance
+>, + Si^.fulfamic acid has a slight absorbance at 200 millimicrons (1/500
that of N03) care must be taken not to add too large an excess and to add a

  iri^COnStant am°Unt to both san^le "d blank'  Correct the absorbances A
and B by subtracting the appropriate blanks.                    -wo-ncea A
     Absorbance B is proportional to the N0§ concentration.  Plot ppm nitrate
nitrogen vs absorbance B.  Can this plot curve 1.                    ™.w«e


     Absorbance A minus absorbance B is proportional to N0£ concentration.  Plot
ppm nitrite nitrogen vs absorbance A minus absorbance B.  Call this plot  curve
2.



                                   65

-------
Procedure with Deajamination Agent Absent

     If the sample has insoluble matter, filter about 15 ml through Whatman k2
filter paper.

     Follow the procedure outlined in the calibration curve discussion with
the following variations.  If the sample started out at 50 ppm NO§-N, dilute
5 ml to 25 ml.  A reagent blank is essential and must be treated exactly
like a sample.  The reagent blank should include the reducing agent, catalysts,
be the identical pH of the sample, and treated under the same temperature and
time conditions of the sample.

Calculations

     Read all values at 200 millimicrons.  To get the true value for each
absorbance, the appropriate blank must be subtracted before any calculations.

     Absorbance B is absorbance due to NO"-N.  Calculate ppm nitrate-N from
Curve 1.                                 •*

     Absorbance A minus absorbance B is absorbance due to nitrite-N.  Calcu-
late ppm nitrite-N from Curve 2.

     Multiply by the dilution factor to obtain the original nitrate and
nitrite concentrations.

Interferences

     Most anions interfere somewhat at 200 millimicrons.  However, they are
much weaker absorbers than either N0§ or N0£.  If the reagent blank has been
carefully made up and handled, these interferences can easily be subtracted
out giving rise to very little error.
ALTERNATE PROCEDURE FOR SOLUTIONS CONTAINING A DEAMMHiATION AGENT


Calibration

     A third curve must be prepared for nitrite ion at neutral pH's.

     Scan a series of nitrite solutions containing 0-2 ppm nitrite nitrogen
in distilled water, from 250-195 millimicrons.  Read the absorbance values
at 200 millimicrons.  Plot the absorbance at 200 millimicrons vs ppm  NOl-N.
Can this Curve 3.

Procedure

     Two aliquots of each solution must be taken.  The first aliquot  is
adjusted to pH between 7 and 9, diluted to the mark, and scanned "as  is"

                                   66

-------
from 250-195 millimicrons.  Read the absorbance at 200 millimicrons and call -
this value absorbance C.  Treat the second aliquot exactly as before to
obtain absorbances A and B.  Absorbances A and B are equal if the deammination
agent present is there in sufficient quantity to deamminate the solution as
the pH is lowered to 2.

Calculations

     Absorbance B is due to the absorbance of NO" ion.  Read from Curve 1.

     Absorbance C minus absorbance B is due to the absorbance of NOZ ion.
Read from Curve 3.                                                 2

     Calculate the results in the same manner as shown in the preceding
section.
                                  67

-------