PB86-158607
Arsenic (3) Oxidation and
Removal from Drinking Water
Houston Univ., TX
LIBRARY
U. 5.
^EWYQRK, N.Y. 10007
Prepared for
Environmental Protection Agency, Cincinnati, OH
Feb 86
U.S. DepvtmeRt of Coiranerce
Natimal Techrecd tofarmatwn Service
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PB86-156607
EPA/600/2-86/021
February 1986
ARSENIC(III) OXIDATION AND REMOVAL
FROM DRINKING WATER
by
Phyllis Frank and Dennis Clifford
University of Houston
Houston, Texas 77004
Cooperative Research Agreement No. CE-807939
Project Officer
Thomas J. Sorg
Drinking Water Research Division
Water Engineering Research Laboratory
Cincinnati, Ohio 45268
WATER ENGINEERING RESEARCH LABORATORY
OFFICE OF RESEARCH AND DEVELOPMENT
U.S. ENVIRONMENTAL PROTECTION AGENCY
CINCINNATI, OHIO 45268
NATIONAI TECHNICAL
INFORMATION SERVICE
JitUPARKim Of COJMIRCt
murjuin •» niM
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TECHNICAL REPORT DATA
(Plcatf lead Instructions on thr reitrtt Injure cumplctmgl
. REPORT NO.
EPA/600/2-86/021
2.
3 RECIPIENT'S A
4 TITLE AND SUBTITLE
Arsenic (III) Oxidation and Removal from Drinking Water
5 REPOST DATE
February 1986
B. PERFORMING ORGANIZATION CODE
7 AUTHORIS)
Phyllis Frank
Dennis Clifford
B PERFORMING ORGANIZATION REPORT NO
9. PERFORMING ORGANIZATION NAME AND ADDRESS
Dept. of Civil Engineering
University of Houston, 4800 Calhoun
Houston. TX 77004
10 PROGRAM ELEMENT NO
tl CONTRACTSGHANT NO
CR807939
12 SPONSORING AGENCY NAME AND ADDRESS
Water Engineering Research Laboratory - Cincinnati, OH
Office of Research and Development
U.S. Environmental Protection Agency
Cincinnati, OH 45268
(3. TYPE OF REPORT AND PERIOD COVERED
14. SPONSOPING AGCNCV CODE
EPA/600/14
IS. SUPPLEMENTARY NOTES
Thomas J. Sorg, Project Officer
313-569-7370
16 ABSTRACT„. The oxidative pretreannent of As(III) using chlorine and oxygen was studied
following quantification of As(III) and As(V) removals by activated alumina columns.
Activated alumina removed 100 ug/L As(V) from a typical grounduater at pH 6.0 much
more effectively than As (IT. I). Approximately 23,500 bed volumes of water were treated
before As(V) reached the 0.05-mg/L maximum contaminant level (MCL), whereas only 300
bed volumes could be treated before As(IIT) reached that level.
,'
Variables affecting the oxidation of As(IlI) by chlorine include the pH, chloride
concentration, other ions, chloramine formation and TOC. In artificial grcundwater
containing 100 ug/L As(III), a 1.0 mg/L chlorine dose oxidized 95 percent of the
As(III) to As(V) in less than 5 seconds. Increasing chloride concentration slowed the
reaction slightly, but not significantly for water treatment. The counterion (sodium
or calcium) did not appear to affect the extent of reaction in the artificial ground-
water or in chloride solutions up to 0.010 M. The presence of 5 mg/L TOC substantially
slowed the oxidation kinetics of 100 ug/L As(IlI) by 1.0 mg/L chlorine dosage. Sparging
1 hr with oxygen did not oxidize 100 ug/L As(III) in artificial groundwater, while 100
ug/L As(III) in deionized water was approximately 14 percent oxidized. Monochloramine
was found to be capablo of oxidizing 40 percent of the initial 100 ug/L As(III) in the
pH range of 6.5 to 10.5.
17.
KEY WORDS AND DOCUMENT XNALVSlS
DESCRIPTORS
b IDENTIFIFRS/OPEN ENOED TERMS
C COSATI I'iclli/Gioup
13. DISTRIBUTION STATEMENT
Release to Public
19 SECURITY CLASS ITIllI Rrporl)
Unclassified
O OF PAGES
84
2O SECURITY CLASS nh'tpffe
Unclassified
EPA Form 2220-1 (R«». 4-77) POEVIO.O EDITION is OBSOLETE
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DISCLAIMER
Although the information described In this article has been funded wholly
or In part by the United States Environmental Protection Agency through
assistance agreement number CR-807939 to the University of Houston, it has not
been subjected to the Agency's required peer and administrative revlev and
therefore does not necessarily reflect the views of the Agency and no official
endorsement should be inferred.
11
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FOREWORD
The U.S. Environn* —'ta 1 Protection Agency is charged by Congress with
protecting the Natlo. s land, air, and water systems. Under a mandate of
national environmental laws, the agency strives to formulate and implement
actions leading to a compatible balance between human activities and the
ability of natural systems to support and nurture life. The Clean Water Act,
the Safe Drinking Water Act, and the Toxics Substances Control Act are three
of the major congressional laws that provide the framework for restoring and
maintaining the integrity of our Nation's w&ter, for preserving and enhancing
the water we drink, and for protecting the environment from toxic substances.
These laws direct the EPA to perform research to define our environmental
problems, measure the impacts, and search for solutions.
The Water Engineering Research Laboratory is that component of EPA's
Research and Development program concerned with preventing, treating, and
managing municipal and industrial wastewater discharges; establishing
practices to control and remove contaminants from drinking water and to
prevent its deterioration during storage and distribution; and assessing the
nature and cortrollability of releases of toxic substances to the air, water,
and land from manufacturing processes and subsequent product uses. This
publication is one of the products of that research and provides a vital
communication link between the researcher and the user community.
This research is part of a concerted research effort at the University of
Houston to examine the removal of inorganic contaminants from drinking water
sources. In particular, this research examines the oxidation and removal of
naturally occuring arsenic concentrations in ground water.
Francis T. Mayo, Director
Water Engineering Research Laboratory
111
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ABSTRACT
The concentration of arsenic In drinking water is regulated because of
the known toxicity of arsenic. A survey of unit processes for water treatnent
and our previous research reveal that pentavalent arsenic is more effectively
removed from water than trivalent arsenic: Thus, following quantification of
As(III) and Ae(V) removals by activated alumina columns, the oxidative
pretreatment of trivalent arsenic using chlorine and oxygen was studied.
Arsenic(V) at a concentration of 0.100 mg/L was much more effectively
removed from a typical groundwater by activated alumina at pH 6.0 than
As(III). Approximately 23,500 bed volumes (BV> could be treated before As(V)
reached the 0.05 mg/L MCL whereas only 300 BV could be treated before As(III)
reached the MCL.
Variables affecting the oxidation of As(III) by chlorine include the pH,
chloride concentration, other ions, chloramine formation and TOC. In
artificial ground water containing no amirDnla or TOC vlth 100 ppb As(III)
present initially and 1.0 mg/L chlorine dosage, the reaction reached 95
percent completion in less than our shortest possible observation time of 5
seconds. Thus with 1.0 tig/L chlorine dosage the As(lll) oxidation rate vas
greater than 20 mlcrograms/L second. The extent of oxidation at 30 seconds
was insensitive to pH in the range 6.5 to 9.5, with decreasing extent of
reaction outside this range. Increasing chloride concentration slowed the
reaction slightly, although this effect is not significant for water
treatment. The counterlon (sodium or calcium) did not appear to effect the
extent of reaction in the artificial goundwater or in chloride solutions up to
0.010 molar. Honochloranlne is capable of oxidizing 40 percent of the initial
100 ppb As(III) in the pH range 6.5 to 10.5. The presence of 5 mg/L TOC
substantially slowed the oxidation kinetics of 100 ppb As(IZl) by 1.0 mg/L
chlorine dosage. Although the reaction reached 50 percent completion in less
than 30 seconds. It did not reach 80 percent completion until approximately 30
minutes.
Sparging 1 hour with oxygen did not oxidize 100 ppb As(III) in artificial
ground water, while 100 ppb As(lII) in DI water was approximately 14 percent
oxidized. However, capped samples of As(III) in DI water and artificial
ground water were completely oxidized after 61 days on the shelf with air In
Che head space.
This study is part of a comprehensive research effort on arsenic removal
froa drinking water being carried out by researchers at the University of
Houston. The related work includes field studies on As(IIl)/As(V) rrooval in
San Ysld.ro, New Mexico and Ae(III) removal in Hanford, California. Previous
iv
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related laboratory studies include work on the removal of As(V) by activated
alumina and anion exchange resins, and development of an analytical method for
arsenic speciation.
This report was submitted in fulfi 1 latent of Crart No. CR-807939 by the
University of Hourton under the sponsorship of the U.S. Environmental
Protection Agency. This report covers the period May 1982 to Hay 1984, and
this work was completed as of December 1984.
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CONTENTS
Abstract iv
Figures ix
Tables xi
Abbreviations and Symbols xii
Acknowledgments ............... xiii
1. Introduction 1
General Chemistry 1
Occurrence 2
Toxlcity 3
Arsenic Removal 4
Removal by Coc&ulation '•
Removal during Filtration 7
Removal by Reverse Osmosis 8
Removal by Electrodialysis 8
Removal by Adsorption 9
Removal by Ion Exchange 10
Summary • ..11
2. Conclusions ;... .12
3. Recommendations ...... .....13
4. Theory **
Aqueous Arsenic Chemistry 1*
Oxidation Studies 17
Aqueous Chlorine Chemistry 17
Free Chlorine 19
Combined Chlorine 19
Chlorine Oxidant Decay 20
Reaction Thermodynamics ' 20
Reaction Mechanisms 22
Oxygen Oxidation 24
5. Experiments 25
Preliminary Column Studies 25
Column Experiment Design 25
Oxidation Studies 26
Objectives 26
Reagents 26
Procedure Development for Chlorine Oxidation 27
vii
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Ammonia quenching .................... 28
Thioleulfate quenching ................. 29
Sodium Sulfite quenching ................ 29
DPD quenching ...................... 30
Chlorine Oxldant Experiments ................. 32
Oxygen Oxidant Experioente ............. .... .33
6. Results and Discussion ................ ...... .36
Results and Discussion of Column Studies ............. 34
Chlorine Oxldant Studies ..................... 62
Kinetics ........................... *2
pH Effect .......................... *5
Effect of Chloride Concentration ............... *7
Effect of Counterion ..................... 69
Effect of Chloramines .................... *'
Effect of pH on Chloramint Oxidation ............. 52
Effect of TOC ........................ 54
Oxygen Oxidant Studies ...................... 56
Sparging Tests ........................ 56
Shelf Life Experiments .................... 56
References
.58
Appendices 63
1. Equations for arsenic Eh-pH diagram 63
2. Equations for chlorine Fh-pH diagram 65
3. Chlorine Oxidant Decay Reactions 6*
4. Conditions Procedure for Activated Alumina 67
5. Arsenic Analysis 68
6. Change in Absorbance 69
viii
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FIGURES
Page
pH-concentration diagram for solubility of AsjO^ IS
Eh-pB diagram cf arsenic-water system at 25 C 16
Eh-pH diagram of chlorine-water system at 25 C 18
Calculated free energy of As(III) oxidation by chlorine
as a function of pH and chloride concentration 23
5 Effluent history from an activated alumina colunn with
As(III) in flouridated feedwater 35
6 Effluent history from an activated alumina column with
As(V) in flouridated feedwater 36
7 Comparison of arsenite and arsenate breakthrough curves
from activated alumina columns 37
8 pC-pH diagram for arsenious acid and its anion 39
9 pC-pH diagram for arsenic acid and its an ions 40
10 Comparison of flouride adsorption by activated alumina
from columns fed As(III) or As(V) 41
11 Arsenic and flouride effluent histories from activated
alumina column at San Ytidro, NM 43
12 Kinetics of reaction in artificial ground water with
DPD quenching 44
13 Effect of pH on extent of reaction In artificial
groundwater 46
14 Comparison of effect of chloride ion with different
eounterion upon extent of As(IZI) oxidation 48
15 Comparison of kinetics of As(III) oxidation in presence
of different counterions 50
ix
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16 Kinetics of chloranine oxidation - -51
17 Effect of pH upon Ae(III) oxidation by ch lor amines 53
18 Compcrlaon of kinetics of Ao(III) oxidation in
artificial ground water and aged tap water 55
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TABLES
Number
1
2
3
4
5
6
7
8
9
10
11
12
13
14
Competition of feedvater to activate4 Alumina columns
Kinetics of AsdlZ} Oxidation in DJ water with no
Results of screening of sodium aulflte quenching . . .
Final concentration of DPD and phosphate used in
Composition of backgound vater used in oxidation
Composition of calcic background vater used in
Comparieon of adsorption capacities from laboratory
Page
. . .25
. . .27
. . .28
. . .29
. . .29
. . .30
. . .31
. . .31
. . .32
. . .32
. . .56
. . . .57
xl
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ABBREVIATIONS AND SYMBOLS
BV Bed volume; volume occupied by adsorbent in a column, or an
equal volume of solution passed through the column.
C Concentration of solute in liquid.
CQ Initial concentration.
DPD N'N diethyl-p-phenylene diamlne oxalate. used to quench the
oxidation reaction with chlorine.
EBCT Empty bed contact time; the time required for a volume of
aolution equal to the bed volume of adsorbent to pass through
a column.
G Cibbs free energy.
C° Standard Cibbs free energy.
Cf° Standard Gibbs free energy of formation.
GFAA Graphite furnace atomic adsorption.
K Equilibrium constant.
K Acidity constant; equilibrium constant for the deprotonation
reaction for an acid.
MCL Maximum contaminant limit.
meq Mil liequlvalent; one-thousandth part of standard gram
equivalent.
NPOC1 Non-purgeable chlorinated organlcs; the non-volatile fraction
of the chlorinated organlcs formed in water upon chlorlnation.
pKa Negative logarithm of acidity constant.
THM TrihaLomethanes.
TOC Total organic carbon; a measure of the organics present In
water.
xii
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ACKNOWLEDGMENTS
We are particularly grateful to the following Individuals for their
important contributions to this study:
Mr. Ton Sorg, the U. S. EPA project officer, for his continuing support,
guidance, and cooperation during this and related studies, and
Dr. Arup Sengupta for his interest, helpful suggestions and encouragement
during the conduct of Che laboratory studies and the preparation of the
report.
xili
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SECTION 1
INTRODUCTION'
This study on arsenic oxidation and removal from drinking voter was
undertaken at the University of Houston as part of a comprehensive research
effort on arsenic removal from drinking water. Related field studies using
Che Ul'/EPA Mobile Drinking Water Treatment Research Facility have been carried
out in San Ysldro, New Mexico and are continuing in Hanford, California.
Related laboratory studies have been done on establishing the As(V) capacity
of activated alumina, regenerating arsenic-spent alumina, establishing the
fundamentals of As(V) uptake by ion-exchange resins, and developing an
analytical method for the separation and determination of As(IIl) and As(V).
Arsenic is a ubiquitous, naturally occurring, semi-metallic element. The
global average concentration of arsenic in the earth's crust is 1.8 ppm.
Certain geologic formations are relatively rich in arsenic, whereas, other
formations contain little or none of the element. Arsenic in various
inorganic or organic forms may be present in the air, soil, or water. Its
presence may arise from natural phenomena such as the weathering or
dissolution of arsenic-bearing rocks or from human activities such as the
manufacture and application of arsenical pesticides, the combustion of
arsenic-bearing coal and oil, and the smelting of various ores, commonly
copper- and lead-bearing ores.
Specifically, arsenic may enter surface waters such as rivers, lakes, and
estuaries in runoff from agricultural fields treated with arsenical
pesticides, In various industrial wastes, and in runoff from natural rock
formations that contain arsenic. Similarly, arsenic may enter ground water by
a variety of routes. Inorganic arsenic may enter the ground water naturally
from the dissolution of arsenic-bearing geologic formations within a given
aquifer. Ground waters may also be contaminated by leachate from improperly
landfilled fly ash from copper and lead smelters, wastes from arsenic refining
and pesticide manufacture, tailings from ore processing, and other similar
activities. Concern has been expressed about future activities such as in-
situ coal gasification, which may result In the leaching of various
uncharacterized organic arsenlcals along with other retort products 153).
However, most arsenic-bearing ground waters currently arise as a natural
phenomenon.
GENERAL CHEMISTRY
Arsenic is stable In four oxidation states (+5, +3, 0, -3) at the redox
potentials (Eh conditions) occurring in most aquatic systems. The pentavalent
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and trlvalent states of arsenic are those most commonly encountered In
solution [21]. The pentavalent form appears in solution as an anion of the
triprotic arsenic acid, and the trivalenf form appears as arsenious acid or
its anion. The dissociation of arsenic 8 (2)
•= H+ + AsOA3' pKa = 11.5 (3)
The dissociation of arsenious acid (trivalent arsenic) can be written as:
As (OH) 3 = AsO(OH)2" + H+ pKfl - 9.2 (4)
Thus in the pH range of 6 to 8 pentavalent arsenic occurs as either a singly-
or doubly-charged anion, whereastr iva lent arsenic appears as the nonionlc
arsenious acid. Since these species are interconvertable, arsenate should
predominate in surface waters, and arsenite should be the predominant species
in anaerobic ground waters. Ferguson and Gavis [21] reported that the rate of
oxidation of arsenlte by oyxgen is very slow at natural pH values.
OCCURRENCE
The natural occurrence of arsenic in ground water is worldwide. Various
areas of the world must use arsenic-bearing ground waters as a drinking water
supply. Antofagasta, Chile [5], Nova Scotia [25], and certain regions of
Taiwan are three well-known examples. Various areas of the United States,
most notably the southwest and the northwest regions, have a significant
number of arsenic-bearing groundwaters, some of which are used as drinking
water supplies [12].
Arsenic concentrations found in groundwaters vary. Examples from well-
known problem areas are 0.75 mg/L total As In Antofagasta, Chile, 0.85 mg/L
and 1.1 mg/L total As in Taiwan, and 0.63 and 8.0 mg/L total As in Nova Scotia
[33]. In an AWWA survey of inorganic contaminants, the total arsenic
concentration ranged from 0.052 to 0.19 mg/L among reported wells that
exceeded the maximum contaminant limit (MCL) in use for potable water in the
United States [12]. The modal concentration occurred in the 0.075 to 0.100
mg/L concentration bracket. The total arsenic concentration found in surface
waters such as lakes and rivers can often exceed the concentration found in
ground vater. The highest values, often greater than 3 mg/L, are usually
associated with human activity [17]. Naturally contaminated arsenic-bearing
ground waters usually have high pH and high bicarbonate [17]. This phenomenon
presumably is a result of the chemical characteristics of the arsenic-bearing
rock in the aquifer.
The valence state of arsenic in ground water varies with location;
however, data on the speciation of arsenic-bearing ground water is relatively
scarce. A survey of various amenic containing welle around the world
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Indicated that the arsenlte to arsenate ratio varied from 0.007 to 1.3 (33].
These values correspond to 0.7 percent and 56.5 percent of the total arsenic
in the -4-3 oxidation state. Drinking water wells sampled in Hanford,
California, in February 1984 had a range of total arsenic concentration of 30
to 90 ppb, all of which (> 95 percent) was trivalent arsenic |16j. These well
waters also contained sulfides, indicating reducing conditions within the
aquifer. Ground water drawn from an infiltration gallery 4 m (13 ft) deep in
San Ysldro, New Mexico contained 80 ppb total arsenic, of which 40 percent was
trivalent arsenic [16]. Well water from a 13-m (43-ft) depth in this saire
area had a much higher total arsenic concentration (190 ppb), which was
primarily trivalent arsenic. This water also contained other reduced species,
namely sulfide, Fe(II), and Mn(II). From the above reports, it is apparent
that both As(III) and As(V) are important species in arsenic-contaminated
ground water supplies.
In a survey of individual wells at homes in Nova Scotia, Subramanian et
al. [64] reported that 36 out of 50 samples exceeded 50 ppb total arsenic.
Calculations made from their reported data show that 22 percent of the samples
had no detectable As(III). The arsenic(III) in the remainder of the samples
ranged from 0.8 to 91 percent. Hire of these samples contained As(IIl) in
concentrations exceeding 50 ppb. In fact, six of the samples had As(III)
concentrations greater than 100 ppb, witi. fhree samples exceeding 300 ppb
As(IU).
TOXICITY
The toxicity of arsenic is widely known. Arsenic poisoning way be acute
or chronic. Both cases have been fairly well studied — the former because of
the popularity of arsenic as a poison, and the latter because arsenic t.as the
principle ingredient in Fowler's solution used to treat akin disorders and a
common ingredient in various general health tonics. Ingestion of a large
dosage of arsenic results in severe capillary damage that may ultimately lead
to circulatory failure and death. The symptoms of chronic arsenic poisoning
include diarrhea, skin pigmentation, hyperkeratosis, circumscribed edema,
nausea and loss of appetite. Arsenic exerts its effect by reacting with
cellular sulfhydryl groups. Thus the sulfhydryl enzymes (especially pyruvate
oxidase) essential to cellular metabolism are Inhibited [31].
The toxicity of arsenic depends on its valence state. Arsenite reportedly
is 60 times more toxic than arsenate, primarily because arsenite inactivates
the sulfhydryl enzymes. Arsenate exhibits toxicity when it is reduced to
arsenite by the body [31]. Plants and animals detoxify arsenate by
methylation [47]; and one bacterial species is capable of detoxifying arsenite
by «*:.ididation to arsenate. Bacteria are known to be capable of arsenate
efflux through the phosphate transport system in the presence of phosphate
[55, 73]. Enhanced penetration of the cellular membrane by the non-ionic
trivalent arsenic may contribute to its greater toxicity, but this effect has
not been measured. Increased toxicity because of penetration by non-ionized
HOC1 and NH3 are, however, well known, possibly similar examples.
In addition to the toxic effects, it has also been claimed that arsenic is
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a carcinogen. However, this claim is highly controversial. Epidemiological
studies of increased lung cancers attributed to arsenic exposure among smelter
workers in Tacoma, Washington, have been criticized on the grounds that
exposure to SO? and other agents were not excluded [19]. The increased
incidence of skin cancers attributed to arsenic ingestion in drinking water in
Taiwan has also been questioned because that water also contained ergometrlne
and other fluorescent alkaloids that cannot be discounted as carcinogens [24,
56]. Other instances of arsenic ingestion by means of drinking water have
failed to produce an increased incidence of cancers (56]. Furthermore,
experiments to induce cancer in laboratory animals have failed to demonstrate
that arsenic acts as a carcinogen [24]. Those experiments where cancers have
been Induced have such low animal survival rates « 5 percent) as to make the
results meaningless [24]. Therefore, it has not been demonstrated that
arsenic is a carcinogen [24, 44].
In fact, some current research indicates that arsenic may be an essential
trace element. Arsenic in the form of organic arsenicals has been used as a
feed additive for poultry and swine to promote weight gain (2). Such
observations have promoted research into the nutritional value of arsenic.
Recent research Indicates that arsenic may be an essential trace element in
chickens and goats [44, 69] with a relationship between zinc and arsenic
metabolism. However, the metabolic studies seen to deal only with the +5
oxidation state. Although the toxcity of arsenic is well-known, its exact
role and influence in human metabolism remains unknown-- and fraught with
controversy.
Despite the lack of knowledge about the consequences of the ingestion of
very small amounts of arsenic, the ultimate toxicity of arsenic is well
recognized. It is for this reason that the concentration of arsenic in
foodstuffs and water is regulated. The USEPA National Interim Primary
Drinking Water Regulations set a Haximun Contaminant Limit (MCL) for arsenic
at 0.05 mg/L.
ARSENIC REMOVAL
Trace amounts of arsenic are usually successfully removed by conventional
water treatment methods [46]. For a number of surface water sources with
significant levels of arsenic, arsenic occurrence in the finished water was
relatively infrequent after conventional water treatment [43]. Experimental
studies have shown that arsenic may be removed by coagulation, filtration,
reverse osmosis, ion-exchange, electrodialysis, and adsorption onto activated
alumina. In most of these processes it is pentavalent arsenic that is removed
most effectively. Only the removal of inorganic arsenic will be discussed
here because organic species such as monosodium methyl arsenic (MSHA) do not
appear to be a problem in ground water used for potable water supply.
Removal by Coagulation
In the environment, arsenic can be removed from solution by
coprecipitation with metal ion precipitation, most notably the iron oxides
[21]. In treatment plants and experimental studies, arsenic has been removed
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with varying degrees of success by coagulation and precipitation, primarily
with aluminum and iron hydroxides. This process is usually followed by
filtration. Lime softening has also been studied for arsenic removal. Of the
various coagulants, ferric chloride has proved the most successful. Arsenic
removal by coagulation appears to be dependent upon initial arsenic
concentration, dosage of coagulant, pH, and the valence of the arsenic species
as will be discussed below.
Aluii--
In early alum precipitation experiments by Buswe 11 [7], some removal of
high initial arsenic concentrations (25 mg/L) at pH 6-7 was achieved.
Rosehart and Lee [51] found alum to be moderately successful in removing high
concentrations of arsenic (132 mg/L) from gold mine wastes. Addition of alum
to these waters in an Al/As ratio of 4:1 at pH 7-8 achieved a maximum 90
percent removal of As(V) to a final concentration of 13.2 mg/L. The removal
of lower concentrations of As(III) was successful. Under the sane optimal
dosage and pH conditions, 500 ppb As(III) was reduced 95 percent to 25 ppb.
However, Shen [57] achieved only a 32 percent removal of 1 mg/L creenic with
20 mg/L alum at pH 6.8. Although the valence state was unspecified, the
ground water presumably contained As(III) as it also contained siilfides and
ammonia, indicating reducing conditions. This observation agrees with what is
known about the poor removal of *.s(III) by activated alumina--dehydrogenated
A1(OH)3.
Nillson [45] also investigated arsenic removal by alum. At an initial
concentration of 21 mg/L As(V) at pH 6.5-7, a 94 percent removal war achieved
with a 10 mg/L alum dosage. However, these same conditions failed to remove
any of 23 mg/L As(III). Under these same conditions of pH and alum dosage, 71
percent removal of an initial concentration of 4.2 mg/L As(V) was achieved.
An examination of these Hata indicates that As(V) is much better removed by
alum than As(III). And, although data is usually presented as percent removal
which tends to obscure the point, as the initial concentration of arsenic
increases, the arsenic remaining after coagulation and precipitation also
increases.
The extensive research by Logsdon [42] supports these trends. Alum
coagulation resulted in a 5-15 percent removal of As(III) which was
considerably less than the removal of As(V) Gulledge and O'Connor [27]
obtained under the same conditions. Gulledge and O'Connor found that greater
than 90 percent removal of arsenic(V) could be acheived with 30 mg/L dosage of
alum at pH 5-7. Logsdon [60] also found that coagulation with 30 mg/L alum
could achieve the arsenic MCL with areenic(V) concentrations of 1.5 mg/L or
less, but the MCL was not obtainable when the arsenic(III) concentration was
0.1 mg/L or higher unless As(IIl) was oxidized to As(V) prior to coagulatf.cn
and precipitation.
The optimal pH and alum dosage for arsenic removal appear to be specific
to each water. Despite this observation, certain general trends are followed.
Pentavalent arsenic is more effectively removed than trivalent arsenic under
the same optimal conditions. This result is an Important consideration in
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drinking water treatment when a certain HCL needs to be achieved. Also, the
higher the initial concentration of arsenic, the higher the concentration
remaining after alum coagulation and precipitation.
Iron—
Numerous coagulation and precipitation experiments with some form of iron,
principally either ferric or t'errous sulfate, or feme chloride, have been
performed — often in conjunction with alum precipitation experiments. One
study by Pierce and Moore [48] obtained very high arsenic removal rates with
amorphous Fe(OH>3. Ferric hydroxide at 4.45 mg/L removed 119 mg/g at 0.53
Bg/L As(V) and 37.5 mg/g at 0.83 mg/L As(III). Buswe 11 [7] observed an 84
percent removal of 25 mg/L As(V) after coagulation, precipitation, and
filtration at pH 5-6 with 250 mg/L of "Ferrisul" ferric sulfate. As(III),
only after oxidation to As(V) with calcium hypochlorite, was removed to 4 mg/L
under the same conditions. Cherkinskli [10] reported a 96 percent removal of
36.2 mg/L arsenious oxide with a ferrous sulfate dosage of 750 ng/1. Rosehart
and Lee 151] obtained 94 percent removal of 132 mg/L As(V) at pH 8 with
ferrous sulfate at a Fe/As ratio of 1.5, but failed to remove any of 0.5 mg/L
As(IIl) under these same conditions. Using ferric chloride as the coagulant
at a higher Fe/As ratio of 4.0, 90 percent removal of 132 mg/LAs(V) at pH 9
and 95 percent removal of 0.5 mg/L As(III) at pH 8 was obtained. With both
ferrous sulfate and ferric chloride, removal of As(III) was more difficult
than removal of As(V). Gull edge and O'Connor 127] demonstrated that ferric
sulfate doses of 30 mg/L at pH 5-8 would remove 0.05 mg/L As(V) almost
completely. Slien tested both ferrous sulfate and ferric chloride. He
preferred the ferric chloride in his treatment plant design because he
achieved greater removal (82 percent versus 24 percent) of 1.0 mg/L jrsenic at
pH 6.7 and 20 mg/L coagulant. Although the valence was not stated, arsenite
is presumed to be the predominant species because the water was anaerobic and
contained other reduced species. Shen also obtained 98.7 percent removal of
0.60 mg/L arsenic after 20 mg/L chlorination —indicating oxidation. Sorg and
Logsdon [60] achieved high removal rates (81 percent or better) of As(V) with
ferric sulfate at pH 6.7 to 8. However, As(lII) was removed only at very lev
initial arsenic concentrations.
Pierce and Moore [48] concluded that the adsorption of arsenic onto ferric
hydroxide was not solely because of electrostatic processes, but included
specific adsorption or the formation of a chemical bond. Furthermore, their
results showed that arsenate adsorption was 20 tines faster than arsenite
adsorption. They concluded that for concentrations found in natural waters
the optimum pU for arsenite removal is 7 and for As(V) removal pH 4 is
optimal. While other researchers, as discussed above, were able to remove
both states at other pU values, As(V) was always removed more effectively then
As(IIl). This result is important when attempting to remove arsenic to meet
the MCL.
Lime softening-
Lime softening removes arsenic probably by co-precipitation of calcium
arsenate on the hydrous magnesium oxide floe. Lime softening followed by
-------
filtration removed arsenic from water in Argentina [66]. Hard, turbid waters
were Created with lime and recarbonated to adjust pH. However, further
details were not given in the reference.
Cherkinekil (10| studied arsenic removal by lime softening along with
coagulation by ferrous eulfate. A dosage of 2250 mg/L of calcium oxide was
required to remove 96 percent of 362 mg/L areenious oxide. However, pH and
the alkalinity, important variables in softening, were not given.
Various researchers have described arsenic removal by lime softening from
mining vaste streams. Lagultton [36| wrote that arsenic valence, pH, calcium
activity and phosphate concentration are important variables for naximlzirig
arsenic removal. Koeehart and Lee [51| attained removal of 123 mg/L As(V) to
6.6 mg/L at pH 12 with CaO. The optimal Ca/As ratio was 9.8. Under these
same optimal conditions, a 95 percent removal of 0.50 Bg/L As(III) to 0.025
mg/L was attained.
In his investigation on drinking water treatment, Ntllson [45] was unable
to remove by lime precipitation any of 23 mg/L As(IIl) at pH 9.5 after
pretreatment of tap water vith 6 mg/L phosphate as P. In tap water with 21
mg/L As(V), treatment with lime at pH 9.5 removed arsenic to 9.7 ng/L. At an
arsenic concentration of 4.2 tng/L Ae(V) in "mechanically treated" municipal
wastewater, lice treatment at pH 9.5 yielded a minimum of 0.1 ng/L for a
maximum removal of 98 percent. The nature of the rtechanical treatment was not
specified. Shen [57] obtained only a 20 percent removal of 1.0 mg/L arsenic
after a dosage of 20 mg/L lime. The pK of the raw water was 6.8, but the pH
attained was not given. Also, as noted before, the valence of the arsenic was
not given but was presumably As(III). Scrg and Logsdon 160] conducted pilot
plant lint softening teats at pK 9.5 and 11.3, the latter tests Included
recarbonation and second stage settling. Lime softening at pH 9.5 resulted in
10 percent removal of O.A8 ng/L As(IIl) and 49 percent renrval of 0.42 ng/L
As(V). At pH 11.3, 63 percent removal of 0.34 mg/L As(III) was attained after
first-stage treatment and 69 percent overall removal after the second-stage.
Ai. this pH, 98 percent removal of 0.58 mg/L As(V) was attained after first-
stage treatment. Second-stage settling did not improve removal. Arsenic(III)
removal was always poorer than arsenic (V) removal. They concluded that the
are«nic HCL can be attained only if 0.35 ng/L or less AsfV) IB present
initial ly, and If As(III) concentrations are Leas than 0.1 mg/L, preoxidation
La not necessary to achieve the HCL when eoftenlng is performed at pH 10.7 or
above.
Hone of the researchers have examined all the variables that Laguitton
theorized were Important in maximazing arsenic removal. However, the work to
date shows that the removal of arsenic to meet an HCL depends significantly
upon the valence state and the pH. Higher pR appears to increase removal and
As(V) is more effectively removed than As(III).
Removal during Filtration
Much of r-.he filtration data has been collected in conjunction vith arsenic
co-precipitation experiments. Shen [57] studied filtration vith sand and
-------
anthracite. At slow filtration rate (2-4 vr/m day), a sand column 48 cm long
removed 90 to 95 percent of 1 mg/'L arsenic, vhile an anthracite column 70 cm
long was needed to achieve 95 percent removal. Slow sand filtration yielded
better arsenic removals than rapid sand filtration (168 - 194 m3/m2/day);
however, the filter rune proved too short to be useful For practical
application. Sorg and Logsdon (60) also studied filtration after coagulation
and precipitation. Increased removals were obtained by following coagulation
with dual media filtration. Filtration with granular activated carbon (GAC)
yielded only slightly higher removals than filtration with dual media.
Filtration studies like those above were conducted as polishing steps.
The mechanism of action can be attributed at least partially to adsorption.
Sand is known to absorb both As(III) and As(V) differentially (26). Elution
profiles of 0.180 mg/L As(V) and As(IIl) from sand columns show that As(lll)
elutes much faster r.han As(V) under oxidizing and neutral conditions, and at
approximately the same time under strongly reducing conditions. This behavior
Is attributed to differential adsorption.
Removal by_ Reverse Osmosis
Preliminary studies by Fox {23] showed reverse osmosis to be effective in
removing arsenite and arsenate from water. Both the DuPont (aramfc*) and the
Osmonlcs (cellulose acetate) membranes removed 0.750 mg/L below the detection
limit, probably 0.005 mg/L. Pentavalent arsenic was removed to met the MCL
when the Initial arsenic concentration was 4 ng/L for the aramid membrane and
6 og/L for the cellulose acetate membrane, A maximum removal of 80 percent
was achieved for trivalent arsenic. Studies by Fox [22] for the EPA in
Fairbanks, Alaska and Eugene, Oregon show 96 to 98.6 percent removal of 0.42
to 0.46 mg/L arsenic. The speiiation of arsenic was not available. Point of
use research directed by Sorg [59] at the Drinking Water Research Division
EPA, studied arsenic removal by a household reverse osmosis unit. Pentavalent
arsenic was generally better removed than trivalent arsenic. Removals of
arsenite from well water varied from 43 percent to 81 percent while removals
for arsenate were 97 to 99 percent. Chlorinatlon of the Ae(III) bearing
waters increased removal in two out of three cases. The removal of Aa(V) by
reverse osmosis is more effective as expected because As(V) exists as an anion
while As(lII) is non-ionic in the pH range of interest, and as such more
easily penetrates the membrane.
Remova 1 by Electrodlalyeis
Electrodialyals (ED) is a process that would be expected to remove As(V)
preferentially over As(IIl) since the basis of removal depends upon the
existence of charged species. Since arsenate exists as an anion for pH above
2.2, while Aa(III) exists predominantly in its non-ionized form below pH 9.2,
As(V) should be preferentially removed over As(III) below pH 9. Reversible
electrodialysis (£DR) was studied for arsenic removal in New Mexico using the
University of Houston/EPA Mobile Drinking Water Treatment Research facility
(16). Water from San Ysidro, NM Well Mo. 4 containing 190 ppb total arsenic
which was predominantly As(III) was treated by EDR. Most of the arsenic (140
ppb) remained In the product indicating poor removal of As(lII). However,
8
-------
more research is needed in this area.
Remova\ by Adsorption
Adsorption by a variety of sorbents has been widely investigated because
of the inherent advantages of packed beds. Two of the most commonly used
adsorbents are activated carbon and activated alumina although a variety of
other adsorbents have been studied and utilized in water treatment.
Gupta and Chen [28] investigated the adsorption of arsenic by activated
alumina, activated bauxite, and activated carbon. These studies indicated
that activated carbon was ineffective in removing arsenic from water. This
conclusion was generally supported by Sorg and Logsdon (60].
On the other hand, activated alumina can successfully remove arsenic from
water. Bel lack [3] removed arsenic from ground water in Falion, Nevada using
activated alumina (F-l grade Alcoe, 28 X 48 mesh) in 18.5 mm-ID glass columns
filled with 25 gm at an optimum flov rate of 2.5 to 3.0 gpm/sq. ft. Initial
arsenic concentration was 0.106 mg/L and the runs were terminated when the
effluent arsenic concentration reached 0.01 mg/L.
The previously mentioned studies by Gupta and Chen [28] indicated that
arsenic was strongly absorbed by activated alumina. At 0.300 mg/1 As(V) final
concentration, the equilibrium adsorption capacity was approximately constant
at 7.5 mg As(V)/g alumina for varicu, solution salinities from fresh water
through 0.67 N NaCl and "Fe(OH)3-precipUated" sea water. However, at a low
equilibrium concentration for 0.50 mg/L As(V), the capacity varied from 1.0 mg
Ae(V)/g alumina for sea water to 4.2 mg As(V)/'g alumina for fresh water. No
explanation of this observation was given. He also observed that the rate of
adsorption decreased win Increasing salinity, and that the As(V) capacity of
alumina was a strong function of pH.
However, these experiments by Gupta and Chen indicated that the adsorption
of As(III) onto alumina was not pH dependent in the range 4 to 9. Between pH
4 and the optimum pH 9, a nearly constant 0.2 mg As(III)/g alumina was
absorbed at equilibrium with fresh water. This adsorption capacity is
substantially less than that for As(V).
At arsenic concentrations of about 4 mg/L As(V) and 1 mg/L As(III), silica
at concentrations of 70 mg/L and higher provided significant competition.
This result is what would be expected since silica, especially the ionic form,
is a highly preferred ion by alumina [13].
Further investigation on arsenic adsorption onto alumina was performed by
Rosenblum end Clifford [52]. They found that arsenic adsorption onto alumina
was optimized at pH 6.0 in column tests with an Influent concentration of 1.0
og/L As(V) in nrtifical groundwater. Arsenic adsorption was found to be
significantly reduced by anion competition. At a liquid-phase equilibrium
concentration of 1 mg/L As(V), adsorption of arsenic was reduced by over 50
percent in the presence of 15 meq/L SO^2" und 20 percent In tl.e presence of
15 meq/L Cl" compared to deionlzed water. An As(V) adsorption capacity of
-------
16.1 mg As(V)/g alumina at 1.0 mg As(V)/L was obtained from minlcolumn tests.
Column tests on As(III) adsorption resulted in much higher than expected
maximun adsorption capacities. Upon examination of the arsenic effluent
histories they concluded that oxidation of As(IIl) had occured during the run
probably as a result of mlcroblal action. Further studies on As(IlI)
adsorption need to be conducted.
Removal by_ Ion Exchange
Calmon (9] investigated the use of cation and anion exchangers for the
removal of 68 mg/L arsenic from "field waters". Not suprisingly, cation
exchangers were ineffective; however, removals of 55 to 82 percent were
obtained with various anion exchange resins in the chloride form. Removals of
99 and 100 percent were obtained in column tests. Other pertinent information
was not given in the reference cited.
Shen [57] conducted a column test with the weak-base resin lonac A-260.
After 6.6 bed volumes of synthetic water containing 1 mg/L As(III), only 21
percent removal was occuring. However, well water treated in the same manner
resulted in the complete removal of arsenic. These results are hard to
definitively explain without feedwater pH data, arsenic speciation in the well
water, and the concentrations of competing anions such as sulfate in both test
waters. However, at acidic pH which weak-base resins require in order to act
as ion exchangers, As(III) is non-ionic and would not be expected to be
significantly removed. The well water which presumably contained at least
some As(III) because of the presence of other reduced species may have been
oxidised upon standing prior to ion exchange treatment.
Lee and Rosehart [72] investigated several ion-exchange resins. They
obtained removal of arsenic with various resinei however, lack of information
on the arsenic speciation and the presence of other ions severely limits the
usefulness of their results.
Horng and Clifford [32] studied the removal of arsenate by strong-base ion
exchange. A study of five anion-exchange resins (Amberlite IRA-900, lonac
ASB-1, lonac ASB-2, Dowex 11, Amberlite IRA-458) resulted in the conclusion
that Amberlite IRA-458, an acrylic-amine, microporous resin with quaternary
fuctionaiity, was the most arsenate selective resin tested. All of the
strong-base anion exchange resins tested at arsenate and carbonate
concentrations likely to be encountered in ground water gave the following
anion selectivity sequence:
S042' > HAs042", C032' > Cl" > H2As04~, HCOj"
However, because arsenate is less preferred than sulfate, chrcmatographic
elution of arsenic occurred. This chromatographic peak phenomena will occur
for all species less preferred than the most preferred species present. This
effect is one very serious drawback to the use of strong-base ion-exchange for
the removal of arsenate.
Studies on strong-base ion-exchange resins indicate that As(IIl) is not
10
-------
removed at neutral and acidic pH. In fact, Amber lite IRA-458 was used to
separate As(III) from As(V) by Clifford et al.(15]. Weak-base ion exchange
resins, which must operate at acidic pH, also do not remove As(lll). It is
not suprisin^ that ion exchange is ineffective for the removal of trivalent
arsenic as it exists predominantly as the non-ionized arsenious acid (pK^ •=
9.22) at the pH values exchangers are operated. For this reason, while ion
exchange can remove pentavalent arsenic in the pH range of interest (6-9),
trivalent arsenic is removed ineffectively if it is removed at all.
SUMMARY
In the treatment processes discussed, pentavalent arsenic is, without
exception, more effectively removed from water than trivalent arsenic. Thus
oxidative pretreatment of trivalent arsenic will improve the arsenic removal
from waters containing As(III).
11
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SECTION 2
CONCLUSIONS
Arsenic(V) is much more effectively removed by activated alumina than
trivalent arsenic. An Alcoa F-l activated alumina co lupin operated at pH 6.0
rcitoved 18 g As(III)/ m3 alumina, whereas 1610 g As(V)/m3 alumina was removed
when Che effluent concentration of arsenic reached 0.05 mg/L.
The oxidation of ppb levels of arsenlc(lll) by ppm levels of chlorine is
very rapid in aqueous solution. The oxidation of 100 ppb of As(IlI) by 1.0
mg/L free chlorine in artificial ground water reached completion (95 percent)
within 5 seconds. Thus for water treatment purposes, contact times do not
need to be particularly long. The extent of oxidation of 100 ppb As(III) by
1.0 mg/L free chlorine in artificial ground water was insensitive to pH after
30 seconds of reaction time in the range of 6.5 to 9.5. The oxidation of 100
ppb As(IIl) by 1.0 mg/L free chlorine reached 90 percent completion in the
presence of 0.01 M chloride. The counterion (sodium or calcium) had no
significant effect on either the kinetics or extent of oxidation. For most
groundwaters, oxioative pretreatment with free chlorine may be effective in
enhancing arsenic removal.
In artificial ground water, 100 ppb As(III) was oxidized by 1.0 mg/L
nonociiloramlne in the presence of excess ammonia. Only a third of the As(III)
was oxidized, however. This reaction was also insensitive to pH in the range
6.5 to 10.5. Although oxidation by monochloramine can occur, the As(III)
remaining may exceed the MCL.
Arsenlc(III) oxidation by 1.0 mg/L chlorine in aged tap water (TOC - 5
mg/L) was considerably slowed. The reaction did not reach 95 percent
completion until 60 minutes. The extent of reaction was the same as that for
oxidation in artificial ground water. Presumably, oxidation was slowed by the
chlorine demand of the TOC present in the water. Chlorine contact times that
are sufficient will be Important for efficient water treatment.
Although very slow, oxygen is capable of arsenic oxidation. Sparging for
60 minutes with oxygen oxidized only 3 percent of 100 ppb As(III) in
artificial ground water, but 14 percent of 100 ppb As(III) in DI water was
oxidized. However, samples of As(III) in DI water and artificial ground water
were completely oxidized after 61 days of sitting in bottles with air in the
head space.
12
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SECTION 3
RECOMMENDATIONS
LAB STUDIES
Further studies will prove useful for producing optimum arsenlc(III)
oxidation by chloramine for subsequent arsenic removal. The limitations of
this treatment need to be examined experimentally. The extent of As(III)
oxidation by nonochloranine may vary with the monochloramine concentration and
the amount of excess ammonia present. Experiments under the conditions used in
this work with different monochloramine dosages (i.e. 0.5, 2.0, 5.0, 10.0
mg/L) and different amounts of excess ammonia should help determine whether
free chlorine or combined chlorine is the actual oxidant. Also, further study
of longer reaction times may indicate which is the true oxidant.
PILOT STUDIES
Pilot studies of arsenic(III) oxidation by chlorination of actual
arsenic-contaminated waters may be beneficial. The presence of reduced
species such as eu If ides, Fe(II), and Mn(Il) may adversely affect the ease of
arsenic(III) oxidation. Furthermore, species such as copper are known to
catalyze the destruction of HOC1, which may adversely affect the extent of
reaction at a given chlorire dosage. Although laboratory studies look
promising, it may be difficult to obtain good oxidation in the field at the
low chlorine dosages such as are required for electrodlalysis and reverse
osmosis membranes.
Further investigation of whether mechanical aeration or ponding can
provide adequate oxidation for subsequent arsenic removal may a?.so be useful.
The advantage of these methods is that they are relatively low cost and
require little maintenance.
13
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SECTION 4
THEORY
AQUEOUS ARSENIC CHEMISTRY
Arsenic occurs in the +3, +5, 0, and -3 oxidation states; however, under
the pH and Eh (redox potential) conditions found in natural waters, only the
+3 and +5 states are stable {21}. In aqueous systems, pentavalent arsenic
occurs as arsenic acid or its anions. Trivalent arsenic may occur as
arsenious acid, its anion, or arsenious anhydride. The latter species is an
anphoteric soluble solid. Figure 1 shows the influence of pH on the
solubility of As2O^ in aqueous solutions. However, as shown is Figure 1, at
the concentrations generally encountered in natural ground waters, arsenious
anhydride dissolves completely with the formation of primarily undissociated
arsenious acid—^AsOj. Thus, at the concentrations of interest, trivalent
Inorganic aqueous arsenic occurs as arsenious acid or its anion.
Arsenic acid, H^AsO^, is a triprotic acid with pKj • 2.2, pK2 • 6.98, and
pK-j • 11.5, and appears in solution as either the acid or the corresponding
anion. Monoprotic arsenious acid (pK B 9.2) may be written as HAs02«
However, strong evidence suggests that HAs02 and its anion do not actually
appear in solution [58]. Arsenious acid appears to be As(OH)3 with the
following three possible anionic species depending upon the basicity of the
solution: AsO(OH)2~, As02(OH) , and AsO^3". However, the only disocciation
that occurs in the range of interest is:
As(OH)3 = H+ + AsO(OH)2~ pK = 9.2 (1)
Arsenic is not known to form ligands with water; the species As(H2O) n* is
probably not formed. Although As(III)-Cl bonds are known to occur, these
bonds only occur at high UC1 concentrations [58]. Disproportionation
reactions are not known to occur with either of the acids. Thus the arsenic
oxyions are the principle species of environmental interest.
Figure 2 shows the regions of thermodynamic stability of the relavent
trivalent and pentavalent arsenic species at various pH and Eh values. The
equations used to construct the diagram after the method of Stumm and Morgan
(63] are presented In Appendix 1. Aa Figure 2 shows, three different arsenic
oxidation half reactions are possible in the pH range 4 to )1. Thus the
oxidation potential of arsenic half reactions is a function of pH. Figure 2
also shows that under aerobic conditions, puntavalent arsenic is the
thermodynamics 1ly stable species. Therefore, oxygen is capable, at least
theoretically, of oxidizing As(lII) to As(V).
14
-------
I CM
o
V)
CM
O
(0
I
v»x
+
^*
O
w
O)
O
1
0
-1
•2
•3
•4
-5
-6
-7
V
As2O3(Solid)
AsO
Saturation
Point
HAsO,
I
I
| As2O3 Solution^
| CT = 100 ppb ln P"re Water
I
I . i i
AsO.
-2 0 2 4 6 8 10
Figure 1. The Influence of pH on the solubility of arsenlous anhydride.
12
-------
fc. — T
20
16
12
pE 8
4
0
-4
-8
*x»
• H3AsO4
\
•
AsO+
^
\|
XS"^^
H,A»0.-
H*,0\
«
\
HAsOj"
"^•^ HO Nv
* ^
^^^,
V
w
•
-
'v. •
"^
•
AsO3'
x^
1000
900
800
500
200
100
0
-20 2 4 6 8 10 12 14
PH
Figure 2. pE - pH diagram for arsenic-water system at 25°C.
16
-------
OXIDATION STUDIES
Pourbaix [49] reported that the oxidation of arsenlc(III) occured only
under conditions of considerable overpotential. Oxidizing agents listed wore
the halogens or their oxygen compounds, chromic acid, nitric acid, hydrogen
peroxide, or permanganate. Johnson and Bruckenstein [35] studied As(Ill)
oxidation by Ig at pH 8.2 to 9.2 and discovered the rate expression to be:
-d AB(III) - klI3"][H2As03'] + kjjlj] (H2As03"] <2>
dt
where [As(III)] ° [HjAsOj] + [H2As03~}.
Roebuck [58] reported the rate of oxidation in acidic media to be
d As(III) - kf[An(III)Hl-^.]
dt [I]2 [H+]
Liebhafsky has proposed a me chant em Co explain this behavior. Studies on
Asflll) oxidation by bromine showed the reaction to be zero order in bromine
158]. Other studies on oxidation of As(III) In strong HC1 solutions
demonstrated that the kinetics shoved a complicated dependence upon HC1 1 58].
Chlorine has been used as an oxidant by Shen [57], Sorg and Logs don [60],
and other researchers in drinking water treatment; however, studies of
kinetics or the effects of pH and ions in solution have not been reported.
And although As(V) is the thermodynamical ly stable species under aerobic
conditions, the oxidation of As(III) by oxygen is reported to occur very
slowly at neutral pH [21].
AQUEOUS CHLORINE CHEMISTRY
Chlorine is a common disinfectant and oxidizing agent for w&ter treatment.
Consequently, the aqueous chemistry of chlorine has been extensively studied
and reviewed. The various clilorine speciec can participate in the following
types of reactions: hydrolysis, redox, chloramination, substitution,
addition, atom exchange, radical oxidation, photochemical decay, metal
catalyst decomposition, and self decomposition. All of these reactions can
occur under natural water conditions. However, only the environmentally
Significant reactions will be considered for simplicity.
Figure 3 shows the regions of thermodynamlc stability of the most
Significant chlorine species. The equations used to construct the figure are
given in /VP^udix 2. As Figure 3 shows, chloride is the stable species in
vater under all conditions. A comparison of Figures 2 and 3 indicates that
considerable overpotential exists for tne oxidation of As(III). At a typical
17
-------
30
20
10
HOC)
CI2(aq)
H20
ocr
cr
H2O
-20
8 10 12
24 6
PH
Figure 3. The pE - pH diagram for chlorine in water, 25°C.
18
-------
ground water pH of 8.0, the border of thermodynanic stability for Aa(III)
occurs at Eh • -100 volts while the border of thermodynamlc stability for the
chlorine species OC1" occurs at Eh « 1200 volts.
Free Chlorine
Chlorine gas dissolves instaneously in water by rapid and reversible
hydrolysis.
H20 + C12 - HOC1 + H+ + Cl" (4)
Increasing H or Cl" can cause Cl2 to predominate £t equilibrium thereby
decreasing hypochlorons acid (HOC!) formation. This behavior is found to
occur in sea water [29].
Hypochlorous acid is a weak acid which dissociates as follows!
HOC1 = H+ + OC1" pK - 7.5 (5)
At pH < 7.5, HOC1 is the dominant species while at pH > 7.5 OC1" ir. the
dominant species. Each of these species has a different redox potential.
Free chlorine species (HOC1, OC1~, C12, H2OC1+, Cl+) are capable of oxidizing
I" and Br~ to HOI and HOBr resulting in the formation of bromlnatsd and
lodinated trihalomethanes (THM) in chlorinated natural waters. Since chlorine
species are so reactive, other constituents of natural water such as ammonia
and organics effect the performance of chlorine as an oxidizing agent for a
target species.
Combined Chlorine
The tern "combined chlorine" refers to the halogenated nitrogen compounds
such as moTiochlorcamine, dichloroamine, trichloroamine, and organic haloa^ine.
The formation of chloroamines from ammonia occurs in a stepwise manner:
HOC1 + NH3 «= NH,C1 + K20 fast (6)
HOC1 + NH2C1 - NUCL2 + H20 slow (7)
HOC1 + KHC12 - NC13 + H20 slow (6)
The chloramines obtained depend upon pH, reaction time, the relative
concentration of HOCl and NH^ , and temperature [18]. At pH greater than S
and a molar ratio of HOCl to NH^-N of Itl or less, monochlorcaine Is observed.
Usually nitrogen trichloride is the only species detected at pH less than
three [18, 29]. Under certain conditions, the formation of chloramines is
reversible; however, the equilibriun state probably is rarely attained because
the reverse reaction is so slow in comparison to the forward reaction. The
mono- and dichloramines retain some disinfecting power, although much less
than an equivalent concentration of free chlorine. Free chlorine reacts with
organic amines to form the organic haloamineit RNHC1, RNC12> R2NC1, but not
19
-------
much Is known about the mechanisms of these reactions.
Chlorine Oxidant Decay
Hall [29] describes chlorine oxldant decay as any reaction involving
conversion of chlorine Into constituents that are not delectable by the
analytical methods for free chlorine. Decay reactions are summarized in
Appendix 3; however, only those reactions with significance to this research
will be discussed.
The formation of total organic halogen compounds (TOX) that can be further
hydrolyzed to yield non-haloform and THM is perhaps the most important
chlorine decay reaction [43]. Fleischaker and Randtke [20] report that the
substitution and addition reactions to form non-purgable chlorinated organlcs
(MPOC1) are faster than the oxidation of organic matter by chlorine and
perhaps faster than the hydrolysis of chlorinated intermediate byproducts Co
form chloromethanes.
Another important set of reactions are those associated with the oxidation
of ammonia-nitrogen and the phenomenon of breakpoint for residual chlorine.
When the molar ratio of chlorine to nitrogen exceeds 111, the unstable species
NHClj is formed which subsequently decays to N2 and NO," resulting in a loss
of ammonia-nitrogen. However, if the molar ratio exceeds 1.5, NC13 is formed
resulting in a preservation of ammonia-nitrogen (18, 29]. These reactions
occur during chloramination prior to the breakpoint wnere free chlorine begins
to exist.
Other reactions of lesser importance in oxidant decay in water treatment
are the oxidation of inorganic ions and the eelf-decomposition reactions.
Lister [39] reported that Fe(Il) and Mn(II) were oxidized by OC1" with no
catalytic activity. Other inorganic ions are susceptible to attack; these
ions include nitrite, sulfite, selenite and arsenite. Nickel, cobalt, and
copper ions catalyze the decomposition of OC1" to 0, and Cl~. Mechanisms of
Inorganic ion oxidation will be discussed later. Under certain conditions
the self-decomposition of OC1~ and HOC1 to form 02, CIO., CIO*'", CIO," and Cl"
occurs [29].
REACTION THERMODYNAMICS
A study of thermodynamics shows whether a reaction is possible and how
potentially energetic it is. Since the pK values of hypochlorous acid,
arsenious acid, and arsenic acid are known, four reactions are possible
depending upon the pH. The following half reaction potentials are from the
ttiennodynamic daca of La timer [37], Si lien [54], and Pourbalx [49].
HOC1 + H+ + 2 e" " el" + H20 (9)
E • 1.49 volts
OC1" + H20 + 2 e" « Cl" + 2 OH" (10)
E - 0.90 volts
20
-------
HAs02 + 2 H20 « H2As04~ + 3 H+ + 2 e~ (11)
E = 0.666 volts
HAs02 + 2 H20 = HAsOA2' + 4 H+ + 2 e" (12)
E - 0.881 volts
As02' + 2 H20 - HAs042" + 3 H+ + 2 c" (13)
E = 0.609 volts
The four reactions based upon the possible speciation combinations of arsenic
and chlorine are given below.
a. 2.2 < pH < 6.98
HAs02 + HOC1 + H20 - H2As04" + 2 H+ + Cl" (14)
E - 2.156 volts
AC- -416.2 KJ (-99.43 kcaJ)
K - 7.869 X 1072
b. 6.98 < pH < 7.5
HAs02 + HOC1 + H20 = HAsOA2" + Cl" + 3 H+ (15)
E o 2.371 volts
AC - -457.72 kJ (-109.35 kcal)
K " 1.474 X 1080
c. 7.5 < pH < 9.2
HAs02 + OC1" + H20 - HAs042" + Cl" + 2 H+ (16)
E = 1.781 volts
-343.82 kJ (-82.14 kcal)
R - 1.659 X 1060
21
-------
d. 9.2 < pH < 11.5
As02~ + OC1" + H20 - HAs042~ + Cl" + H+ (17)
E » 1.509 voles
AC - -291.34 kJ (-69.60 kcal)
K - 1.063 X 1051
These calculations show that the oxidation of arsenic is potentially verv
energetic at all pH values from 2.2 to 11.5 and that these reactions posess
very large equilibrium constants, K. At concentrations normally encountered,
the reaction is even more favorable. For a chlorine concentration of 1.0
mg/L, chloride concentration of 73 mg/L, an Ae(III) concentration of 100 ppb,
an As(V) concentiation of 0.1 ppb, and at pH equal to 8.0, the potential is
1.72 volts.
A itethod to see the theoretical sensitivity of the oxidation reaction by
chlorine Is to examine the reaction potentials instead of the standard
reaction potentials. The following assumptions were made in order to solve
the equationsi [As(JII)] - [As(V>] , 1.0 mg/L chlorine as either HOC1 of OC1"
, and no attempt to account for the differences in corcentration of the
species as the pH approached the pK values. Figure 4 shows the reaction
potentials and G values as a function of pll for different Cl" concentrations.
As the pH increases, the reaction potentials and consequently the energies of
reaction decrease.
REACTION MECHANISMS
Redox reactions are electron transfer reactions, but these electron
transfers are not accomplished in the same manner for all reactions. Taube
[65] suggests classifying redox reactions into the following categories:
inner-sphere activated complex, outer-sphere activated complex, and a bridge
activated complex. For the inner-sphere mechanism, electron transfer occurs
within a single primary bond system. While for the outer-sphere mechanism,
electron transfer occurs from one primary bond system to another. In this
case, both the oxidizing and the reducing agents are inert to substitution and
the electron transfer takes place between ions of similar geometry. For the
bridged type reaction, at least one reactant undergoes rapid substitution
thereby forming a ligand which bridges the reacting species. The bridging
ligand may facilitate electron transfer in two ways termed "chemical" and
"resonance". The "chemical" mechanism is a process in which either the
oxidizing ion is strong enough to oxidize o- the reducing ion id strong enough
to reduce the shared ligand whereby the electron deficit or excess is passed
to the reducing or oxidizing ion, respectively. In other words, the electron
hops from the reducing center to a specific bound state in the ligand to the
oxidizing center where it remains. In "resonance" transfer the electron never
occupies a well bound state on the ligand. For both types of transfer, the
rate of electron transfer is sensitive to the nature of the ligand. This
.ligand determines the height ot the energy barrier that the electron must
22
-------
X
(E
(U
Assumptions
[AS] = [AS(V>]
CI2 • 1 mg/L
0.01 M Cl
-100
-75
-50
-25
o
0 24 6 8 10 12
PH
Figure 4. The calculated reaction potential of the oxidation of
As(III) by 1.0 mg/L chlorine dosaae as a function of
pH.
23
-------
penetrate; thereby determining the probabllty of transfer.
HOC1 nay act as an clectrophillc teagent with either oxygen or chlorine as
r.he center of reaction. J. E. Draley [18] suggests that HOC1 is an
electrophlle in which the chlorine atom partially assumes the characteristics
of Cl+ and combines with an electron from solution. Either slmultfiieoucly or
subsequently, hydroxyl ion is split off. The transfer of Cl+ Is facilitated
by a negative charge, the basic part of the molecule. The same type of attack
is thought to occur at amlno-nitrogen.
On the other hand, the Cl atom can have such a greater attraction for
electrons that it may be displaced directly as chloride ion. This type of
attack seems to occur with inorganic ions. Anbar and Taube [1] proposed such
a mechanism for the oxidation of nitrite by aqueous chlorine. Work by Lister
[40] supported this mechanism. Halperin and Taube [30] proposed a similar
mechanism for the oxidation of sulfite. The oxidation of Fe(II) and Mn(II)
and the reaction with hydrogen peroxide are thought to have similar transfers
with the displacement of either Cl" or OH" as initial steps. However, the
exact nature of these initial steps has not been elucidated.
The solvent can play a key role in a redox reaction. The hydration of an
ion may significantly affect its suceptibility to redox reactions. The
hydratlon of an ion may either stabilize or destabilize the intermediates.
Stabilizing the intermediate lowers the energy barrier thereby facilitating
the reaction.
OXYGEN OXIDATION
Theoretically oxygen should be capable of oxidizing trlvalent arsenic (see
Figure 2). However, this reaction does not appear to occur except very
slowly. Latimer [37] proposed that direct oxidation by 02 proceeds through
the formation of the intermediate, hydrogen peroxide. This reaction has a
very low potential. Thus, this step becomes the rate determining step for the
direct oxidation by oxygen via this mechanism. Furthermore, pH would also be
an important variable as it determines arsenic speci-.tion and, therefore, the
half reaction potentials.
-------
SECTION 5
EXPERIMENTS
PRELIMINARY COLUMN STUDIES
The objective of Che preliminary column studies waff to verify the
reportedly efficient removal of pentavalent arsenic and the non-removal of
trivalent arsenic on activated alumina. Since some ground waters in the
United States contain both excess arsenic and fluoride, and activated alumina
is known as a successful adsorbent for fluoride, these column studies were
conducted on water that contained both fluoride and arsenic.
Column Experiment Design
Two column runs were conducted simultaneously. Each column was
constructed from 1/4" ID plexiglass tubing with stainless steel Swagelocli
fittings at each end. Each column was loaded with 5 mL of conditioned Alcoa
F-l activated alumina (U. S. standard mesh size 24 X 48). The conditioning
procedure for activated alumina is given in Appendix 4. Mini-pumps from
Milton Roy Model # NS1-33R were used to control the flowiate at 1.7 mL/min for
an EBCT of 2.9 minutes. These pumps provided a steady flow rate over the
maximum length of the entire column run-- 58 days.
The column feedwater was a synthetic ground water designed to resemble
ground water that had been acidified to pH 6.0 with sulfuric acid. This
pretreatment creates a sulfate enriched, bicarbonate-free water. An arsenic
concentration of 0.100 mg/L was chosen because it is twice the MCL and not
uncommon in arsenic contaminated ground waters (e.g. Hanford, CA and San
Ysidro, NM). The name line of reasoning applies to the choice of 3 og/L as
the fluoride concentration. The composition of the column feedwater is given
in Table 1.
TABLE 1. COMPOSITION OF COLUMN FEEDWATER
SO.
mg/L
71
384
3.0
Cations mg/L meq/L
210 9.16
Ca'T 20 1.0
fl+
.2+
25
-------
The Influent to each column differed only in the valence state of the arsenic;
one column received 0.100 ng/L As(V), while the other received 0.100 mg/L
AB(IH). The feedwater was made from reagent grade sodium fluoride, sodium
chloride, calcium chloride, and sodium, aulfate. The pH was adjusted to pH 6
with eulfurlc acid. For the As(V), column the feed water was made in 40- liter
batches. For the As(lII) coHmn, the feedwater with the exception of arsenic
addition was also made in 'iO-liter batches. In order to avoid arsenic
oxidation during the long run, freshly prepared trivalent arsenic was added to
a 4-liter aliquot each morning for the day's feedstock.
An ISCO Model 0 1850 fraction collector was used to collect 25.5-mL
samples every 15 minutes. Various 25.5-mL samples were analyzed for arsenic
content with a Perkin-Elmer Model 5000 Atomic Adsorption Spectrophotometer
with a graphite furnace, Zemnan background correction, and a HGA-400 Furnace
Temperature Programmer. An electrotieless discharge arsenic lamp was used as
the source lamp. Details of this analysis are given in Appendix 5. Flouride
analyses were conducted using an Orion, fluoride Ion-selective electrode Model
# 94-09-00 with an Orion 50i digital pH/ion analyzer.
OXIDATION STUDIES
Objectives
The oxidation studies were undertaken with the fol lowing objectives in
nindi
(1) to examine the kinetics of arsenic(IIl) oxidation by chlorine.
(2) to examine the effect of pH on the oxidation of arsenic(III) by
chloiine.
(3) to examine the effect of anions and cations upon the oxidation of
arsenic(lll) by chlorine.
(4) to examine the effect of chloramlnes upon the oxidation of
arsenic(III).
(5) to examine the effect of TOC upon the oxidation of arsenic(III) by
chlorine.
(6) to Investigate the oxidation of arsenic(III) by aerobic conditions.
Reagents
All stock arsenic(V) solutions were prepared according to Standard Methods
|62) using reagent grade KKjAsO^. All stock arsenic(III) solutions were
prepared according tc Standard Methods (62] using reagent grade As203. All
arsenic(III) solutions wore prepared immediately preceding their use. All
other solutions with the exception of chlorine oxldant were prepared using the
appropriate reagent grade chemicals. Stock chlorine oxldant was prepared froa
26
-------
a commercial bleach (Chlorox) containing sodium hypochlorite. The titer of
this solution was checked immediately preceding use. Analysis showed no
arsenic to be present in tl-e bleach. The stock chlorine solution was
standardized according to Standard Methods [62].
Procedure Development for Chlorine Oxidation
Ultimately a quenching agent was found to be necessary because the
oxidation reaction could not be monitored on a continuous basis. No method
was available for instantaneous arsenic speciation and quantification.
Methods available for analytical arsenic(IIl) separation from arsenic(V)
include ion exchange, pH controlled arsine generation, and selective
extraction. Each of these methods requires subsequent analysis for total
arsenic in various fractions. The method chosen for arsenic speciation, ion
exchange separation, does not provide instantaneous separation. The method
chosen for speciation is based upon the work of Clifford et al. (15]. Five mL
of pretreated, chloride-form IRA-458 (finer than 35 mesh) in a mini-column at
a flow rate of 10 mL/min was used to sepcrate Ac(IIl) from As(V) in a 100 mL
sample. Under these conditions, the fastest possible speciation time was
greater than 10 minutes. Verification experiments with known arsenic
standards showed good results as seen in Table 2. Arsenic was analyzed by
GFAA as described earlier.
TABLE 2. VERIFICATION OF ARSENIC SEPARATION
Sample Z Expected Recovery
As(III) As(V)
100 ppb As(III) 97 — .
50 ppb As(III)/50 ppb As(V) 102 98
100 ppb As(V) — 98
An indirect method to monitor the oxidation reaction of AsdII}, by
chlorine is to measure the disappearance of chlorine. In addition to the
inherent disadvantage of an indirect method, two other disadvantages arose:
(1) Arsenic(III) can be the only species present that exerts a chlorine
demand.
(2) The colorimetric methods for free chlorine determination proved to
be not sensitive enough.
In experiments with 1.0 mg/L chlorine and 100 ppb As(III) in 0.01 M NaCl, the
DPD colorimetric method [62] showed no significant difference between a blank
of 1.0 mg/L chlorine and the oxidation samples after 1, 2, 5, 10 minutes
27
-------
contact tine. Furthermore, the adsorbance of a given sample changed over a
tine span of several minutes (see Appendix 6). The color change was visible;
the color changed from an initial reddish color to a deep purple color often
within 60 minutes.
As continuous monitoring of the reaction proved unfeasible, oxidation was
conducted in batch reactors each sampled at a different time. The bacch
reactors were 2-L polyethylene beakers filled with 1-L of solution. In
oxidation tests on 100 ppb As(III) in DT water by 1.1 mg/L chlorine, samples
taken at 1, 5, IS, 30 minutes then separated (10 minutes) had no detectable
arsenic(lll) remaining (see Table 3).
TABLE 3. OXIDATION IN PI WATER WITH NO QUENCHING
Time
1 minute
5
IS
30
Therefore, a quenching agent was needed because the reaction time of the
oxidation reaction was faster than the ininimum separation time of 10 minutes.
Several criteria weie required of the quenching agent.
(1) The quenching agent must react with chlorine instantaneously in a
mannner to prevent further oxidation of As(III).
(2) The quenching agent cannot oxidize As(III) nor reduce As(V).
(3) The quenching agent must not deleteriously effect the separation
procedure.
(4) The quenching agent should not substantially interfere with arsenic
analysis by GFAA.
Ammonia Quenching--
Ammonia, at first glance, appeared to be a good candidate for a quenching
agent. It reacts very fast with chlorine to form monochloroamines and was
unlikely to oxidize As(III) or reduce As(V), or interfere with the arsenic
separation or analysis procedures. However, previous success in using the
addition of ammonia and the subsequent formation of chlorocmines as a
quenching agent for the oxidation of Se(IV) [4] and Cr(III) [14] by chlorine
did not follow for the oxidation of As(III). Oxidation of 100 ppb As(III) in
28
-------
DI water occured in the initial presence of 1.0 mg/L chlorine and a 20-molar
excess of NHjCadded as NH^-C1). These results are in Table 4.
TABLE 4. SCREENING OF AMMONIA QUENCHING
Sample
control
control
blank
Time
30 sec
1 min
1 min
As(III) Remaining
43
39
98
Thiosulfate Quenching—
Thiosulfate, a common dechlorinating agent, was also considered for use as
a quenching agent. However, the reaction of chlorine with thiosulfate was
reported to occur in a somewhat slow stepwise fashion [71). In view of the
apparent ease of oxidation of As(III), thiosulfate was abandoned as a
potential quenching agent without performing any experiments. Furthermore,
thiosulfate proved to be unsuccessful in the quenching of Se(IV) oxidation by
chlorine apparently by interfering with the separation by ion exchange [4].
Scdium Sulfite Quenching--
Sodium sulfite is another common dechlorinating agent. It is reported to
react instantaneously with chlorine [71]. However, experiments indicated that
As(V) was reduced in the presence of a three times excess of sodium sulfite
based upon a 1.0 mg/L C12 dosage. The results of these experiments are
summarized in Table 5.
TABLE 5. SODIUM SULFITE RESULTS
Sample As(IIl) Recovered, ppb
100 ppb As(V) 18
100 ppb As(III) with 1.0 mg/L C12 98
100 ppb As(V) with 1.0 mg/L C\2 14
50 ppb As(III)/50 ppb As(V) 61
29
-------
DPD Quenching —
DPD, N,N-diethy 1-p-pheny lene dlamine oxalate. which is used in the
colorimetric t»;st for chlorine, is reported to react instanteously with free
chlorine under the proper conditions. The color change associated with the
DPD reaction with chlorine can be seen immediately upon addition of DPD
reagent under the proper conditions. This reaction is one-to-one, follows
Beer's Law and results in the destruction of free chlorine. Furthermore, DPD
is unlikely to oxidize Ap(III) or reduce As(V). Therefore, DPD was a likely
candidate as a quenching jgent.
A standard DPD solution (62] has a pH of approximately 2 because of the
addition of su If uric coid to promote dissolution. In order to keep the pH in
the range 6.2 to 6.5 a phosphate buffer is used in the procedure for the
determination of chlorine. Initial experiments with this quenching method
were conducted using a standard DPD solution and a standard phosphate buffer
solution [62], 5 mL each per 100 mL of sample. However, these experiments
could not be analyzed successfully for arsenic. The peak height absorbance
was drastically reduced. Suspicion that the phosphate buffer was primarily
responsible was confirmed by addition of 5 mL of phosphate buffer to 100 mL of
arsenic standard. In order to reduce interference, the amount of phosphate
buffer added was reduced 10 fold and the pH adjusted to the correct range with
1 N NaOH (or, in the case of extremely high initial pH, with 1 N HC1). Jar
tests were conducted to determine the amount oi NaOH solution to be added for
each initial pH. Once these amounts were determined (see Table 6), they were
used consistent Ij. This procedure allowed the arsenic analysis by CFAA to be
performed satisfactorily. As shown in Table 7, DPD appeared to quench the
oxidation reaction and neither oxidize As(lII) nor reduce Ae(V).N
TABLE 6. AMOUNTS OF NaOH ADDED FOR
Initial pH IN NaOH, mL
5.5
6.5
7.5
8.5
11.5
12.5
1.5
1.5
1.5
1.0
0.0
1.5
(1 N HC1)
30
-------
TABLE 7. DATA ON DPP QUENCHl.'C
Sample As(III) Remaining, ppb
100 ppb As(III) 97
100 ppb As(III) 98
with C12 and quenching agents
SO ppb As(III)/50 ppb As(V) 51
with C12 and quenching agents
For a dosage of 1.0 mg/L chlorine In an oxidation test, the quenching
dosage of DPD was at 4.3 molar excess. Further details are given in Table 8.
TABLE 8. REAGENTS
ReagentConcentration
meq/L mol/L
0.34 0.34
0.34 0.17
**
C12 2.8 X 10'5 2.8 X 10'5
DPD
1.2 X 10'* 1.2 X 10'*
The DPD was added under conditions of extremely turbulent nixing in order to
achieve rapid dispersion. The color change that DPD underwent proved
beneficial in monitoring the effeciency of rapid mixing.
Mixing tests were conducted to determine if As(III) oxidation occured
during mixing. No detectable oxidation took place after 60 minutes of rapid
mixing. These results are presented in Table 9.
31
-------
TABLE 9. MIXING TEST DATA
Stirring Speed Time of Mtxlng.niin Percent AsCllI) Recovered
low
tied
high
high
high
high
1
1
1
2
5
10
98
98
98
97
97
98
Chlorine Oxtdant Experiments
An artlfical ground veeer, composition given in Table 10, was used as
background water for the oxidation test except where otherwise noted.
TABLE 10. BACKGROUND WATER
Anlons meq/L mg/L Cations meq/L mg/L
soA2'
HC03"
Cl"
2.0
6.0
2.0
96
366
71
Ma* 10.0 230
The pH was adjusted, if needed, by dropwise addition of either 0.1 n HC1 or
1.0 M NaOK solution. All pH measurements were made with an electrode
calibrated on using two buffers bracketing the expected pH value. An initial
arsenic(HI) concentration of 100 ppb was chosen because it is twice the MCL
and is not an uncommon concentration in arsenic contaminated ground waters. A
chlorine dosage of 1.0 mg/L was chosen for the batch tests because there was
low chlorine demand in the artifical ground water and the dosage was
sufficient for reaction completion. The batch chlorine oxidation tests were
conducted in 2-L polyethylene beakers with 1-L sample size. A propeller mixer
provided rapid mixing for the addition of stock chlorine solution and
quenching agent. The shaft and impeller were plastic coated. All tioe
increments were measured with a laboratory timer. A 100-mL pipette was used
to collect the sample which was immediately separated. An initial sample and
the As(III) fraction were stored in a 60-mL and 125-mL respectively
polyethylene bottles for subsequent GFAA analysis. An occasional As(V)
fraction was collected as a spot check of separation recovery; otherwise,
Ae(V; was determined by difference.
32
-------
Oxygen Oxidant Tests
Some oxygen oxidation tests were conducted on various background waters
containing 100 ppb As(III). For al 1 these experiments, a 500 mL sample In a
Pyrex gas vanhlng bottle. Model f 31760 was bubbled with Oj at a flow rate of
300 to 40C nL/minute for 60 minutes. One sample to serve as a blank was
bubbled with nitrogen under these sane conditions.
In another experiment, orsenic(Ill) containing sampled vere allowed to sit
in 125 ml polyethylene bottles with air in the headspace for 61 days. Samples
vere 200 ppb As(lII) in DI water at pH 6.0, 7.0, and 8.0; SO ppb As(V)/ SO ppb
As(III) in DI water at pH 6.0, 7.0, and 8.0; and 50 ppb As(III)/ 50 ppbAs(V)
in artificial groundwater with 1 raeq Ca2+ (composition given in Table 11) «t
pH 7.3 and 8.3.
TABLE 11. BACKGROUND WATER WITH CALCIUM"
An ions neq/L tag/L Cations meq/L. ng/L
S^2" 2.0 96 Na+ 9.0 207
HC03" 6.0 366 Ca2+ 1.0' 20
«' 2.0 71
33
-------
SECTION 6
RESULTS AND DISCUSSION
COLUMN STUDIES
The effluent history from the As(IIl) column is shown in Figure 5, and the
effluent history from the As(V) column is shown in Figure 6. As Figure 5
shows, effluent trivalent arsenic concentration reaches the MCL much sooner
than fluoride (300 and 1600 bed volumes, respectively), indicating that
fluoride is a more preferred species than trivalent arsenic. This result is
expected because, at the influent pH of 6.0, trivalent arsenic is
predominantly non-ionic arsenious acid. As Figure 6 indicates, effluent
arsenic(V) reached the MCL significantly later than fluoride breakthrough,
i.e. 23,400 and 1550 bed volumes, respectively. This behavior was expected
because arsenate is a more preferred species than fluoride and on a
milliequivalent basis the As concentration is much lower than fluoride--
0.00135 meq/L compared to 0.158 meq/L.
As shown in Figure 7, a comparison of the arsenite versus arsenate
breakthrough curves, As(III) reached the MCL almost immediately at 300 bed
volumes(0.6 days), while As(V) did not reach the MCL until 23,400 bed volumes
(48 days). Thus, the presence of pentavalent arsenic results in column runs
nearly 80 times longer than trivalent arsenic. Therefore, pre-oxidation of
Ad(III)-containing waters is essential for efficient treatment using activated
alumina.
Even though trivalent arsenic appeared in the effluent very quickly, some
arsenic (III) was removed. In fact, at 90 percent of equilibrium, a mass
balance showed that 0.344 mg As(III) had been removed yielding an approximate
mass loading of 0.078 mg As(III)/ g alumina. Speciation of the influent and
the effluent showed the arsenic to be total ly As(III). Therfore, it is the
trivalent species that is actually absorbed onto the alumina, assuming no
oxidation of As(III) occured.
The comparatively poor removal of As(III) versus As(V) is not suprising in
light of alumina surface chemistry and of arsenic speciation. Alumina is an
aophoteric ion exchanger, capable of both anionic and cationic exchange (34].
In acidic solutions, surface-bonded protons are electrically balanced by
anions adsorbed. In basic solutions, cations Are electrically balanced by
surface-bonded anions. Hydroxide ions are highly preferred as counterions to
the adsorbed protons in neutral and slightly acidic solutions. Other anions
such as fluoride and arsenate must compete with hydroxide for ion-exchange
sites.
34
-------
in
Time,Days
7 8 9 10 11 12 13 14 15 16 17 18
100
90
80
70
£ 60
a
•" 50
<
40
30
20
10
0
Run Number 1
Influent Composition
70.9 mg/L, 2.0 meq/L
384 mg/L, 8.0 meq/L
3.0 mg/L. 0.16 meq/L
100 ppb
211 mg/L, 9.16 meq/L
20.0 mg/L, 1.0 meq/L
6.0
2,9 mln.
9
1000 Bed Volumes
Figure 5. Arsenlte, As(III), and fluoride breakthrough curves from a minicolumn
containing 28 X 48 mesh, F-l activated alumina.
-------
Days
10
20
30
40
91
JQ
a
a
100
90
80
70
60
50
40
30
20
10
0
OJ
E
Run Number 2
Influent Composition
70.9 mg/L, 2.0 meq/L
384 mg/L, 8.0 meq/L
3.0 mg/L, 0.16 meq/L
100 ppb
211 mg/L, 9.16 meq/L
20.0 mg/L, 1.0 meq/L
6
8
10 12 14 16
1000 Bed Volumes
18 20 22 24 26
Figure 6. Arsenate, A$(V), and fluoride breakthrough curves from a minlcolumn containing 28 X 48
mesh, F-l activated alumina.
-------
100
10
—I—
20
Days
30
40
50
—I—
•• As(lll). Arsenlte. Run No. 1
As(v). Araenate. Run No. 2
10
15
20
25
1000 Bed Volumes
Figure 7. Comparison of As(III) and As(V) breakthrough curves
from minicolumns of activated alumina. CT = 100 ppb.
EBC1 = 2.9 min. PH - 6.0
37
-------
In fact, a generally accepted selectivity sequence for anlon adsorption
(or Ion exchange) on alumina exists [6. 67, 68). This sequence is presented
below:
OH" > HP042", HAsO^2" > F~ > Si(OH)30' > CrO^2' > Cl" >
N03" > MnOA~ > C104~ > CH3COO"
At the column influent pH of 6.0, trivalent arsenic exists predominantly
(1000:1) as the non-ionic species--arsenious acid. However, as Figure 8
shows, at the influent conditions (pH = 6.0, 0.100 mg/L As(III)), the hydroxyl
ion concentration is ten tines greater than the arsenite anlon concentration.
An arsenic mass balance on Run 1 showed that 0.344 ng As(III) was adsorbed;
however, only 0.0001 ng As(III) was available as the anion of arsenious acid
to be adsorbed. Therefore, it is unlikely that ion exchange is the only
mechanism of adsorption of As(lll) because more arsenic was adsorbed than
existed as the anion. Furthermore, at pH 6.0 the arsenite anion always exists
In concentrations less than the concentration of the more preferred hydroxyl
ion.
Unlike trivalent arsenic, pentavalent arsenic exists predominantly as the
singly-charged anion, HjAsO^". As Figure 9 shows, at pH 6.0, this species
concentration is 100 times that of the highly preferred hydroxyl ion. Even
the doubly-charged anion, HAs042", exists at over 10 times the concentration
of the preferred hydroxyl ion. Thus, based on concentration, As(V) can
compete favorably for the available adsorption sites.
As indicated in Figure 10, fluoride adsorption was little affected by the
difference In adsorption of As(III) as compared to As(V). The effluent
concentration of fluoride reached the MCL at 1600 bed volumes for the column
with As(III) and at 1500 bed volumes for the column with As(V). Although
As(V) is more preferred than fluoride and As(III) is less preferred, the
difference in arsenic adsorption is unlikely to effect fluoride adsorption.
This situation occurs because the molar ratio of F~/As is 117. Thus, arsenic
vould typically occupy only one site for every 117 sites occupied by fluoride,
if the selectivities were equal.
Studies on arsenic and fluoride removal were conducted in San Ysidro, NM
using the University of Houston/EPA Mobile Drinking Water Treatment Research
Facility [16], The fluoride and total arsenic effluent histories from that
work are presented in Figure 11. Their column was 1" ID filled with Alcoa F-l
activated alumina (mesh 24 X 48). A comparison of results from the field and
laboratory data are presented in Table 12. Although the bed volumes to the
fluoride MCL were greater in the field study, the capacity to the fluoride MCL
was almost exactly the saca as the capacity obtained in laboratory tests.
38
-------
VO
o
O>
O
-10
PK--9.2
- -10
10 12 14
PH
Figure 8. Log concentration vs. pH diagram for arsenlous acid, As(III), system.
Total As(III) concentration is 100 ppb.
-------
PK,
O
O)
o
PH
Figure 9. Log concentration vs. pH diagram for arsenic acid, As(Y)j system.
Total As(V) concentration Is 100 ppb.
-------
D)
E
i
Days
4 6
8
10
= 2 meq/L
= 8 meq/L
3.0 mg/L
1.0 meq/L
pH - 6.0
• As(v) = 100 ug/L
• As(lll)=100 ug/L
1000 Bed Volumes
Figure 10. Comparison of fluoride breakthrough curves from mini-
columns of activated alumina. CT = 3.0 ir,g/L EBCT = 2.9 min
41
-------
TABLE 12. COMPARISON OF LABORATORY AND FIELD DATA
FOR FLUORIDE AND ARSENIC REMOVAL
UH Laboratory Test San Ysidro
As(III) As(V) Field Test
BV to As MCL 300 23400 8760
Arsenic capacity, g/cubic n 18 1610 575
BV to F~ MCL 1600 1550 2520
Fluoride capacity, g/cubic m 4190 4280 4160
Notesi The San Ysidro water contained 48 ppb As(V), 37. ppb As(III), and
2.0 ppm fluoride.
Although the influent arsenic concentration differed between the field
study and the laboratory study, the arsenic removed per m of alumina in the
field was qualitatively in the range expected from tlie laboratory studies
(i.e. greater than 18 g/m3 and less than 1610 g/nr). However, the shape of
the arsenic breakthrough curve in the field study (see Figure 11} was
suprisingly sharp because a much earlier breakthrough of trivalent arsenic
vas expected. By way of explanation, some oxidation of As(IIl) to As(V) nay
have occured in the field column, and the trivalent arsenic concentration of
the influent in the field study was only 0.032 mg/L--one third the
concentration in the laboratory study. Furthermore, the pentavalent arsenic
breakthrough curve in Figure 6 indicated early breakthrough of A»(V). This
effect was presumably because of the shorter EBCT of the lab column, its
shallow bed depth, and the fact that the adsorption zone was a large fraction
of the bed depth.
CHLORINE OXIDANT STUDIES
Kinetics
The results of the kinetic experiments on the oxidation of arsenic(III) by
chlorine in artificial ground water with sodium as the cation are presented in
Figure 12. The Figure 12 indicates that the reaction has reached completion
by the 5-second data point. However, five seconds is the fastest possible
reproducible quenching time using this method. Under the experimental
conditions, the data imply that -d[As(IIT)]/dt £ 2.66 X 10~7 mol/L sec.
Johnson and Bruckenstein [35), in their study on the oxidation of As(III) by
electrogenerated iodine at pH 8.2 to 9.2, reported a maximum overall rate -
d(As(III)]/dt of approximately 7 mol/L sec. Chlorine oxidant Is considered a
more powerful oxidizing agent than iodine since it is capable of oxidizing
iodine. Assuming the oxidation of As(III) by chlorine follows a rate
42
-------
90
80
70
60
.o
a 50
40
30
20
10
0
3.6
3.4
3.2
3.0
2.8
2.6
2.4
2.2
2.0
1.8
1.6
1.4
1.2
1.0
0.8
0.6
0.4
0.2
Time, Hours
100 200 300 400 500 600 700 600
10
Time. Days
15 20 25
30
35
San Yaldfo. New Mexico
AA, 26x48. Run No. 2
April 12-Uay 8. T984
Bod Volume • 400 ml EBCT • 9 mln.
Flow Rata -80 ml/tnin pH,,,d - 8.0 - 0.1
Total As Adsorbed • 210 mg(or 230 mg to 50 pob Aa)
Total F" Adsorbed • 2025 mg(or 1663 mg to 1.4 pom F~)
.Co • 80 ppb
[40% As(lll)]
10
9
8
7
6
pH
4567
1000 Bed Volumes
8
Figure 11. Total arsenic and fluoride breakthrough from 1
column in San Ysidro. KM.
inch
43
-------
1
O)
£
!
-------
expression of the form
- d[As(III» / dt - k [AidlD]" [Cl2]n (1)
where [As(III)] - [H,AsO3l + lH2As03~], and [C12J - [HOClJ + [OCl'J, the only
variables available for manipulation are (As(III)] and IC12). In American
water treatment, chlorine dosages less than 1.0 mg/L are not often practiced.
Increasing the arsenic concentration 10A times to 1000 mg/L, an unrealistic
ground water concentration, would only allow the determination of
-d(As(III)]/dt < 2.66 X 10"3 mol/L sec. For these reasons, further
experiments to determine the kinetics were abandoned.
From equilibrium calculations, the oxidation of As(III) to As(V) by
chlorine should go tc completion; however, as shown in Figure 12, the reaction
only goes to approximately 95 percent completion. Although at pH 8.0
equilibrium calcula.'->ns predict the As(V)/As(III) to be 2.34 X 10" at these
experimental conditions, the reaction appears to reach completion at an
As(V)/As(lII) ratio of approximately 19. The leakage of 1 ppb As(V) during
the analytical separation procedure would automatically give a false
As(V)/As(III) ratio of 100; the measured As(III) concentrations shown in
Figure 12 may be attributable to leakage of As(V). The apparent 95 percent
complete reaction is surely sufficient for water treatment practice. However,
other factors may effect the extent of reaction including the pH, and the
presence of other ions.
jpjl Effect
As the pH increases, the speciatlon of both arsenic and chlorine changes,
and the free energy of reaction decreases. The experimentally observed effect
of pH on the oxidation of As(III) by chlorine is shown in Figure 13. In the
neutral pH range of 6.5 to 9.5, the pH does not significantly effect the
extent of reaction. It is not until pH values greater than 10.5 that the
oxidation reaction is affected adversely. By pH 11.5 the G of reaction has
been approximately halved from the AC of reaction at pH 6.5. The relative
insensitivity of the oxidation reaction of As(III) to As(V) by chlorine in
the neutral pH range may be attributed to the fact that As(III) exists
predominantly as arsenious acid in this pH range and that the changes in pH do
not significantly affect the activation of the species. Furthermore, from the
data it appears that for pH values less than 10.5 that the AC of reaction is
sufficient with adequate overpotential for the reaction to occur. At pH 11.5,
the overpotential necessary for complete reaction may not be available.
The slight decrease in reaction extent at pH 5.5 can be explained by
examining the oxidation equation for that pH range:
HOC1 -
+ 2 H+ + Cl" (2)
Even though the AC of reaction (2) increases with decreasing pH, this
oxidation reaction is acid producing in the 2.2 to 7.0 pH range. As the pH
decreases, the H* concentration increases representing an accumulation of
products. Another contributing factor may be the increase in Cl" as a result
45
-------
70
60
J3
Ct
a 50
co
c
•5 40
E
c 30
01
< 20
10
Initial Conditions
CI2 Dosage "1.0 mg/L
As(lll) = 100 ug/L
Cl~ • 2.0 meq/L (7 t mg/L)
SO|" • 2.0 meq/L (96 mg/L)
HCO:
• 6.0 meq/L (366 mg/L)
- 10.0 meq/L (230 mg/L)
6
8 9 10 11 12 13
PH
Figure 13. The effect of pH on As(III) oxidation by chlorine
with DPO quenching after 30 sec.
46
-------
of acidification from pH 8.3 with KC1. The effect of background Cl~ is
theoretically Important and will now be discussed.
Effect of Chloride Concentration
Figure 14 shove the effect of Cl~ on the extent of reaction with either
sodium or calcium as the counterion for chloride. The data point at 0.1 M
NaCl is falsely high because the high Cl" concentration drove some As(V) off
the ion-exchange resin during separation, and caused a falsely high As(III)
value. The curve was drawn to an estimated actual concentration. This
estimation was made on the basis of blank separation of 50 ppb As(IIl)/50 ppb
As(V) in 0.1 M NaCl.
As Figure 14 shows, increasing Cl" concentrations in the bacVground water
decrease the extent of oxidation reaction. Although some breakthrough of Ae(V)
nay have occured during separation, a real As(III) concentration remains at
the 0.02 M data point. The open square represents As(V) leakage from a blank
separation of 100 ppb As(V) in 0.02 M Na Cl. Since chloride represents a
product of the oxidation, this behavior is qualitatively expected. In DI
water, no As(III) is detected. However, at 0.01 M Cl", As(IIl) Is detected
after oxidation. At pH 6.0 and {Cl "] » 0.01 M, equilibrium calculations
predict the As(V)/As(III) ratio to be 2.22 X 1085. Three possibilities exist
to explain this behavior. The first explanation is experimental error in the
separation. As(V) leakage during separation nay account for the traces of
arsenlc(III) found after oxidation at salt concentrations up Co 0.01 molar.
The second explanation is that the published thermodynamlc data are in error.
Some error in the published thermodynanic data is likely. The calculated pK
values using Gj from. La timer (37] for arsenic acid are 3.6, 7.26, and 12.47
while the measured values are reported as 2.2, 7.0, and 12.5. And, Latimer Is
ultimately the basic source of the thermodynamic data published for arsenic.
And although errors in the published thermodynamic literature aie possible, it
is unlikely to account for the complete difference between the predicted and
measured result, especially since the errors in the thermodynamic data seem to
be predominantly for the reactions at lower pH values. Chloride ion appears
to exert substantial influence on the extent of oxidation at high chloride
concentration, although this influence is of Insignificant practical value at
the chloride concentrations encountered in drinking water treatccnt.
Furthermore, these data indicate that the increase in As(III) remaining at
lower pH values nay be in part (approximately 5 ppb) because of the added Cl"
during pre-acidiflcation with HC1. However, the trend of decreasing extent of
reaction for pH decreasing below pH 6.5 still appears to hold.
The effect of Cl~ concentration upon arsenic(III) oxidation Is very
different than its effect upon Cr(III) oxidation. The oxidation of Cr(IlI) by
chlorine is much slower than the oxidation of As(lll); the Cr(IIl) oxidation
occurs over several hours at higher chlorine dosages [14]. Furthermore,
increasing chloride concentration appears to enhance the reaction, probably
because the chloro-Cr(HI) complex is more susceptible to oxidation than the
Insoluble alternative—Cr(OH)3.
47
-------
a.
a
di
c
'c
'3
a>
CC
Initial Conditions
As(lll) • 100 ug/L
CI2 Dosage = 1.0 mg/L
pH ° 6.0
A Na * Counter Ion to Cl~
Ca Counter Ion to Cl
100
90
80
70
60
50
40
30
20
10
0 0.001 O.C1 0.1
Cl~ Concentration, molar
Figure 14. The effect of chloride concentration on As(III) remaining after
oxidation using 1.0 mg/L chlorine dosage.
-------
Effect of Counterion
Figure 14 also shows a comparison between the extent of reaction when
sodium is the only cation or when calcium is the only cation to chloride. At
chloride concentrations encountered in ground water (0.0001 M to 0.01 M), no
effect can be attributed to the cation. No real ligand effect can be seen.
Figure 15 shows the kinetics of reaction are not measurably changed in the
presence of 1.0 meq/L Ca2"*". A comparison of the rate of oxidation in
artificial ground water that contains 10 meq/L sodium versus one that contains
9 meq/L sodium and 1 meq/1 calcium shows no measurable difference. If the
presence of calcium effects the rate of As(III) oxidation to As(V), it is
probably of no practical significance for treatment processes.
Effect of Chloramtnes
Figure 16 shows the kinetics of oxidation of 0.100 mg/L As(III) in
artificial ground water containing a dosage of 1.0 mg/L chlorine with a 10
molar excess of ammonium chloride at pH 8.3. These conditions produce
monochloramine. This figure has two salient points. First, the reaction
appears to reach completion at an As(V)/As(III) ratio of approximately two
thirds. And, secondly, the reaction appears to have slower kinetics than the
oxidation of As(IIl) by free chlorine. As expected, analysis of the
chloramine species, showed only the presence of NH2C1. The following reaction
can be written:
NH2C1+ U3Ae03 + H20 = NH4+ + HAsO^" + Cl" + 2 H+ (3)
Using the free energies of formation given in Appendix 7, this reaction has a
standard free energy of -82.4 kcal/mol and an equilibrium constant of 2.57 X
1060. Thus, as written, the above reaction is possible. linger the
experimental conditions, thermodynamic calculations predict an As(V)/As(IlI)
ratio of 5.66 X 10 7® at equilibrium. The observed data does not match *he
predicted value at all. The thermodynamic data available may be somewhat in
error as discussed previously. However, these errors are not likely to
account for the total difference between observed and predicted values.
One possible explanation is that, although the above reaction is possible,
it does not actually occur because the energy barrier of activation is too
high. In this case, oxidation is accomplished by analytically undetectable
amounts (less than 0.1 mg/L) of free chlorine remaining in solution. The
formation of monochloroamine is a reversible reaction; however, the
equilibrium occurs in the direction of monochloramine formation. Given the
apparent ease of oxidation of As(III) by free chlorine oxidant, very small
amounts of free chlorine may be capable of partially oxidizing As(III), as
will be quantified later. The hydrolysis reaction of monochloramine to
produce free chlorine has only recently been studied [29]. Kinetic data show
this reaction to be slow. Current kinetic data indicates that, at 60 minutes,
7 percent of the monochloramine should be hydrolyzed. However, this result
was not observed experimentally.
49
-------
O)
c
'E
CO
E
or
Q
o
CO
100«
90
80
70
• %^
60
50
40
30
20
10
i
1
Initial Conditions
100 ppb As(lll)
1.0 mg/L Chlorine
Cl~ s 2.0 meq/L
SO2" = 6.0 meq/L
HCOg = 2.0 meq/L
pH = 8.3
A Ca2+= 1 meq/L
Na+ = 9 meq/L
Dosage
(70 mg/L)
(288 mg/L)
(122 mg/L)
® Na+ = 10 meq/L
•
»
•
I i 1 1
A
I 1 .
8
10
Time, Minutes
Figure 15. A comparison of the kinetics of oxidation of 100 ppb
As(III) by 1.0 mg/L chlorine dosage in the presence of
different counterions.
50
-------
.0
a
a
en
"c
'15
E
a
oc
mz
-------
Studies on Se(IV) (4] and Cr(III) [14] showed that they could not be
oxidized by chloranines. The following reaction analogous to reaction (3),
can be written:
NH2C1 + Se032' + H20 = NH4+ + Cl' + SeO^2' (4)
This reaction has a standard free energy of reaction of -25.74 kcal/mol with
an equilibrium constant of 7.43 X 10 . It is potentially less energetic than
the arsenic reaction. No noticible selenium oxidation occurred in the
presence of monochloramine. However, these factors do not necessarily imply
that nonochl'oramine is the oxidizing species. The oxidation kinetics of
selenium(IV) are much slower than the kinetics of arsenic(IIl) oxidation,
implying that the activation energy of arsenic(III) is much lower than that
for Se(IV). As a result the trace amounts of free chlorine may not be capable
of detectable Se(IV) oxidation.
Effect of pH on Chloramlne Oxidation
As Figure 17 shows, the solution pH does not appear to substantially
effect the extent of reaction in the pH range 6.5 to 10.5. This pH
sensitivity of the chloramine reaction is very similar to that obtained for
oxidation of As(III) by free chlorine (See Figure 13) where a somewhat
decreased rate is observed at pH 5.5. The monochloramine formation reaction
should be dependent upon pU because of the different opecies of chlorine and
ammonia present at different pH values. The reaction of monochloramine to
dichloramine is not considered here because the reaction only occurs at pH
less than 6.0. And this reaction occurs only slowly. Calculations- from
kinetic data show that after 1 hour only 7 percent of the monochloramine has
been converted and at the 1 minute quenching time used in this experiment less
than 0.1 percent of the monochloramine should be converted to dichloramine at
the lowest pH examined. The following reactions are possible!
a. pH < 7.5
HOC1 + NH4+ - NH2C1 + H20 + H+ (5)
AC - -2.62 kcal/mol
K - 83.3
b. 7.5 < pH < 9.3
OC1" + NH4* = KH2C1 + H20
- -12.92 kcal/mol
K - 2.97 X 109
52
-------
.0
Q.
a
•
c
'ca
V
QC
=
(0
100
90
80
70
60
50
40
30
20
10
Initial Conditions
As(lll) = 100 ug/L
Cl~ =71 mg/L, 2 meq/L
SOj" 3 96 mg/L, 2 meq/L
HCOT » 366 mg/L, 6 meq/L
A , *
X Na =230 mg/L. 10 meq/L
\
^ A A A
-
-
•
-
i . i . i . i . i . i . i
8
9
10 11
PH
Figure 17. The effect of pH on the oxidation of 100 ppb As(III) by
1.0 mg/L monochloramine and excess aranonla.
53
-------
c. 9.3 < pH
OCI" + NH3 - NH2C1 + OH" (7)
AC - -6.44 kcal/mol
K - 5.27 X 10*
Reaction (6) implies that In the pH range 7.5 to 9.3 the ratio of
monochloramine to chlorine IB constant. At experimental conditions, this
ratio is equal to 7.5 X 105 for the pH range 7.5 to 9.3 using the given
thermodynamic data. At the level of total chlorine used in the experiment,
the free chlorine concentration using the given thermodynanic data would be
3.76 X 10"11 molar. This concentration is not nearly enough to oxidize the
amount of arsenic observed to be oxidized experimentally if the oxidation
occurs on a stolchiometrlc basis. However, thes* thermodynanic data bay be in
error; the values for K and consequently the monochloramine to chlorine ratio
nay be in error by 2 orders of magnitude, at least. Reactions (5) and (7) are
explicitly dependent upon pH. The monochloramine to chlorine ratio depends
upon pH for pH values less than 7.5 or greater than 9.3. At pH 6.0, the
monochloramine to chlorine ratio is 2.1 X 10*; and at pH 10.3 it is 6.7 X 10*.
Outside the pH range of 7.5 to 9.3, the monochloramine to chlorine ratio
decreases implying that the chlorine concentration increases. But the data do
not reflect an increased extent of reaction. However, as mentioned earlier,
the extent of oxidation by chlorine decreased outside the pH range 6.5 to 9.5.
It is, therefore, still possible for the chlorine in solution to be the actual
oxidizing agent.
Effect of TOC
The kinetics of As(III) oxidation in aged Houston, TX tap water with 1.0
mg/L chlorine dosage are shown in Figure IB. Although tie kinetics were
appreciably slowed in this water, the reaction reached 95 percent completion
as did the oxidation of 100 ppb As(III) in artificial ground water by 1.0 mg/L
chlorine. However, the reaction in the aged tap water toc-k 60 minutes to
reach completion while the oxidation reaction in artificial ground water
reached completion within 5 sec. The oxidation of As(III) in artificial
ground water was at least 720 times faster than the reaction in tap water.
Presumably the chlorine demand exerted by the 5 mg/L TOC in the aged tap water
was responsible for slowing the kinetics of reaction. Hypochlorous acid may
undergo the following reactions with organic constituents in water: oxidation
reactions, substitution to form either N-chlorlnated organics or C-chlorinated
organics, or addition reactions. In general, at least part of the chlorine
demand of a water is exerted Immediately, with other Blower reactions occuring
at longer chlorine contact times [29]. In fact, the oxidation may be parallel
reactions with oxidation by free chlorine occuring immediately, and oxidation
by combined chlorine--either chloramlnes or organic chloramines--occuring at a
slower rate for the second stage.
54
-------
O>
CO
E
0)
CC
CO
<
A Aged Tap Water
TOG * 5 mg/L
pH =7.5 mg/L
9 Synthetic Ground Water
pH - 8.0
5 10 15 20 25 30 35 40 45 50 55 60
Time. Minutes
Figure 18. Comparison of the kinetics of oxidation of 100 ppb
As(III) by 1.0 mg/L chlorine dosage in aged tap water
and artificial ground wa-.er.
55
-------
OXYGEN OXIDANT STUDIES
Sparging Tests
The results of the oxygen sparging tests are presented in Table 13.
A sample of 100 ppb As(III) in synthetic ground water, pH 8.3, was bubbled
with N2 at a flow rate of 360 mL/min. Even though glass surfaces reportedly
adsorb arsenic 133], ninety seven percent of the arsenic(HI) was recovered
after the test. In the oxidation tests in artificial ground water, very
little, if any, oxidation occurred; whereas, in the DI water, somewhat greater
oxidation occurred. Because Cl~ inhibited As(III) oxidation by chlorine, it
may also stabilize As(III) with respect to oxygen oxidation; however, effects
of the other ions in solution cannot be Ignored. The presence of 1.0 mg/L
Fe(III) appeared to enhance the oxidation reaction by oxygen slightly.
TABLE 13. The Results of 0; Oxidation Batch Tests8
Arsenic(III) Oxidized
Conditions Synthetic Croundwaterb DI Water
pH 6.0 3~Z14 Z
pB 7.5 5 1
pH 8.3 5 I
1.0 mg/L Fe(II) 16 1
1.0 mg/L Fe(TII) 8 Z (pH 7.3) 28 Z (pH 6)
a) All tests were of 60 minutes duration with 100 ppb As(III)
initially.
b) Composition: 2 meq/1 Cl~, 6 meq/L HC03~, 2 meq/L SO^"2;
1 meq/L Ca, 9 meq/L Na
c) 02 flow rate ° 325-400 mL/min
Despite highly favorable thermodynamic conditions, the oxidation of As(III)
appears to occur only very slowly. It is possible, therefore, that the actual
oxidizing species is not 02, but H202. The mechanism was postulated to
account for the very slow kinetics of many oxidation reactions by oxygen. The
formation of hydrogen peroxide Intermediate Is the rate limiting step in this
mechanism [37].
As(III) She If-Life Experiments
Various standard arsenic solutions were stored in polyethylene bottles
and left on a shelf at room temperature and away from direct sunlight. After
61 days, the As(III) fraction was separated at a flow rate of 5 mL/min. The
56
-------
results of the shelf-life experiments are given in Table 14.
TABLE U. As(IIl) REMAINING AFTER 61 DAYS
ppb Aa(IIl) Remaining after 61 days
Initial PI Water Synthetic Groundwater
pH 50 ppb As
-------
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14. Clifford. 'J. and J. Chau. Fate of Cr(III) in Chlorinated Water.
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62
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pH = 6.98
APPENDIX 1
Equations used for construction of Eh-pH diagram for arsenic and water
system:
Dissociation reaction boundaries
[H2As04~]
pH = 2.2
lH3As04l
lHAsQ42"l
[H2As04'J
[As043-J
pH » 11.50
[HAs042'J
[As02-]
pH = 9.22
[HAs02]
Oxidation-reduction reaction boundaries
pE = 9.45 - pH
[HAs02I
I
pE « 16.55 - 2 pH
[H3As04J
pE = 9.31 - 1.5 pH
[AsO+J
63
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'HAsOj]
[HAs02]
pE = 11.27 - 1.5 pH
pE «• 14.91 - 2 pH
pE = 10.30 - 1.5 pH
64
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APPENDIX 2
The equations used for the construction of the Eh-pll diagram for chlorine
and water are as follows.
Dissociation reaction boundary
loci']
pH = 7.5
[HOC1]
Oxidation-reduction reaction boundaries
[HOC1]
pE » 26.9 - pH
pE = 23.6
IC12]'5
IC12]'5
-
[Cl ']
[HOC1]
- pE » 25.25 - 0.5 pH
[CD
loci"]
- pE - 28.9 - pH
65
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APPENDIX 3
Chlorine Oxidant Decay Reactions
Breakpoint reaction
2 NHC12 + H20 = N2 + 3 H* -r 3 Cl" + HOC1
4 HOC1 + NH3 " N03" 4 4 Cl" + 5 H+ + H20
Oxidation of Carbon
CH2OH(CHOH)ACHO + HOC1 - CH2OH(CHOH)4COOC 4 H+ 4 Cl"
R-CmiH2-COOH + HOC1 " R-CHO + NH3 +C02 + H+ + Cl'
Organic Substitution
C6H5OH + 3 HOC1 •= C6H2C13OH + 3 HjO
Disproportionation
3 HOC1 - 3 IT* + 2 Cl" + C103"
Decomposition to Oxygen
2 HOC1 = 2 H+ + 2 Cl" + 02
Inorganic Oxidation
N02" + HOC1 - N03" + H* + Cl"
Mn2+ ••• HOC1 + H20 = Hn02(s) + 3 H+ + Cl"
66
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APPENDIX 4
Conditioning Procedure for Activated Alumina
Bed Volume » 50 ml of Alcoa F-l type alumina, mesh 28 X 48
Rinse Bed Volumes
1 I NaOH 10
DI Water 10
2 * H2S04 10
DI Water 10
1 Z NaOH 10
DI Water 10
2 I H2COA 10
DI Water 100 pH «• 4.5
Air dried overnight, then dried 24 hr at 105 C
67
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APPENDIX 5
Arsenic Analysis
Step
Dry
Char
Atomize
Temperature
110 C
250 C
2700 C
Ramp Time
5 sec
5 sec
1 sec
Hold Time
25 sec
20 sec
10 sec
Wavelength: 197.7 nm
Gas Flow: Interrupt Argon
Low Slit Width: 0.7 nm
Sample Volume: 20x'-L
Matrix Modifier: 20 s>L 1000 mg/L NiN03 as Ni in 2 % HNO3
68
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Sample
Blank
1.0 mg/L Cl, dosage
100 ppb As(III) with
1.0 ng/L C12
APPENDIX 6
Change In Abserbance
Initial Absorbance
0
0.35
0.30
0.28
Absorbance after
2 minutes
0
0.4!>
0.38
0.40
Background water 0.01 M MaCl
69
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