THE PRECIPITATION OF ARSENIC FROM AQUEOUS SOLUTIONS.
by
Mary Glastras
Submitted as part of the requirements for the degree of Doctor of Philosophy.
School of Mines
The University of New South Wales
Australia (1988). SR PT01 Form 1 WAIVER THE UNIVERSITY OF NEW SOUTH WALES
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ACKNOWLEDGEMENT
I would like to thank my supervisor
Associate Professor R. G. Robins for his encouragement and guidance during the course of this project . iv
ABSTRACT
Arsenic is an impurity in the processing of many minerals, the more important of which are the sulphide ores. Under natural weathering conditions the arsenic component of these minerals is only slightly soluble, but through chemical processing for the recovery of valuable metals arsenic can be transformed into more soluble and less stable arsenic compounds.
With the consumption of arsen1c used for commercial purposes not increasing, and the need for removing the arsenic impurity in the processing of these minerals, increased quantities of arsenical wastes may occur. A safe means of disposing the arsenic has become essential.
A review of the literature provides little useful or reliable information on the stability and solubility of the various arsenic compounds which are used or could be used for disposal purposes . For that reason the stability and solubility of various relevant arsenic compounds was investigated.
A variety of analytical procedures were used to determine the solubility of the arsenic compounds and the stability of the various dissolved and solid species.
These results were developed into standard free energies of formation or complexation constants for these species.
The free energies were then used to compute stability diagrams for the particular systems in order to predict v
the likely equilibrium conditions over a wider field than was covered in the actual experiments. The stability diagrams are presented in the form of log activity versus pH diagrams and these show the stability regions of the various dissolved species, as well as the solubility and stable regions of solid compounds.
The systems investigated were the arsenic(iii) sulphur(-ii) water system, the iron{ii) arsenic{v) water system, and the iron(iii} arsenic(v} water system. These systems were chosen due to their current usage or likely application to finding a suitably stable and insoluble arsenic residual.
For the arsenic(iii} sulphur(-ii) water system a log activity versus pH diagram was obtain~d showing the regions of stability for the various dissolved .species and the solid arsenic sulphide {Asz S3) .
The iron(ii) arsenic(v) water system was studied in perchlorate and sulphate mediums. Log activity versus pH diagrams 1n both these mediums were obtained showing the solubility and stable regions of the two ferrous arsenates {FeHAs04, and Fe3 {As04)z).
The iron(iii) arsenic{v) water system was studied in perchlorate, chloride, nitrate, and sulphate mediums . Log activity versus pH diagrams were obtained showing the solubility and stable region of ferric arsenate {FeAs04) for the perchlorate, chloride, and nitrate mediums. Due to the apparent complexity of the vi sulphate medium no satisfactory model for a log activity versus pH diagram was obtained.
The iron arsenic sulphur water system and the iron arsenic water system were also studied. Potential pH diagrams were obtained followed by a number of unsuccessful attempts at the hydrothermal syntheses of arsenopyrite and/or loellingite.
From these studies it was determined that the particular arsenic compounds investigated are not as insoluble as has been believed from past studies, and some other means of stablising the arsenic for disposal in residues should be investigated. vii
TABLE OF CONTENTS
Page no.
STATEMENT ii
ACKNOWLEGDEMENT iii
ABSTRACT iv
TABLE OF CONTENTS vii
LIST OF FIGURES xi
LIST OF TABLES xxii
CHAPTER 1. INTRODUCTION 1
1.1 Historical 1
1.2 Uses of Arsenic 1
1.3 Occurrence 2
1.4 Chemistry of Arsenic 4
1.5 Economic Aspects 7
1.6 Industrial ·Applications 9
CHAPTER 2. STABILITY DIAGRAMS 17
2.1 Theory 17
2.2 Arsenic Water System 19
2.3 Sulphur Water System 22
2.4 Iron Water System 22
CHAPTER 3. INSOLUBLE ARSENIC COMPOUNDS 30
3.1 Metal Arsenates/Arsenites 30
3.1.1 Magnesium Arsenic Water System 33
3.1 . 2 Calcium Arsenic Water System 35
3.1.3 Barium Arsenic Water System 42
3.1.4 Ferrous Arsenate Water System 46
3.1.5 Ferric Arsenate Water System 48 viii
Page No.
3.2 Other Insoluble Arsenic Compounds 53
3 . 2.1 Arsenic Sulphur Water System 53
3.2.2 Iron Arsenic Sulphur Water
System 59
CHAPTER 4. EXPERIMENTAL 64
4.1 Introduction 64
4 . 2 Instrumental Techniques 65
4.2.1 PH Titrations 65
4 .2.2 Spectrophotometric Measurements 66
4.2.3 Turbidity Measurements 67
4.2.4 PH Measurements 67
4.3 Preparation of Solutions 68
4.3.1 Arsenic Solutions 68
4.3.2 Sulphide Solutions _ 70
4 .3.3 Iron Solutions 70
4.3.3.1 Ferric Solutions at
an Ionic Strength of
1.0. 71
4.3.4 Carbonate Free Hydroxide
Solutions 71
CHAPTER 5. RESULTS AND DISCUSSION 72
5.1 Arsenic(iii) Sulphur(-ii) Water
System 72
5.1.1 PH Titrations 72
5 . 1 . 2 Spectrophotometric Measurements 72 ix
Page No.
5.1.3 Discussion 75
5.2 Iron(ii) Arsenic(v) Water System 84
5.2.1 PH Titrations 84
5.2.1.1 Perchlorate Medium 84
5.2.1.2 Sulphate Medium 86
5.2.2 Discussion 88
5.3 Iron(iii) Arsenic(v) Water System 102
5.3.1 PH Titrations 102
5.3.1.1 Perchlorate Medium 108
5.3.1.2 Sulphate Medium 108
5.3.2 Turbidity Measurements 111
5.3.2.1 Perchlorate Medium 114
5.3.2.2 Chloride Medium 114
5.3.2.3 Nitrate Medium 118
5.3.2.4 Sulphate Medium 118
5.3.3 Influence of Ionic Strength 122
5.3.3.1 Nitrate Medium 127
5.3.4 Spectrophotometric Measurements 127
5.3.4.1 Perchlorate Medium 130
5.3.4.2 Sulphate Medium 130
5.3.5 Discussion 133
5.4 Arsenopyrite (FeAsS) and
Loellingite (FeAs2) 154
CHAPTER 6. CONCLUSION 161
CHAPTER 7 . REFERENCES 165
APPENDIX 1 182 X
Page No.
APPENDIX 2 184
APPENDIX 3 185
APPENDIX 4 186
APPENDIX 5 187 xi
LIST OF FIGURES
Page No. CHAPTER 2
2.1 Potential pH diagram for the
arsenic water system at 25°C. 21
2.2 Log activity pH diagram for
the arsenic(iii) hydroxy
species. 23
2.3 Log activity pH diagram for
the arsenic(v) hydroxy
species. 24
2.4 Potential pH diagram for the
sulphur water system at 25°C. 25
2.5 Potential pH diagram for the
iron water system at 25°C . 28
CHAPTER 3
3.1 Solubilities of some metal
arsenic(v) compounds from
Chukhlansev. 31
3.2 Solubilities of some metal
arsenic(iii) compounds from
Chukhlansev. 32
3.3 The variation of arsenic(v)
and magnesium concentrations
with pH in equilibrium with
Mg3 (As04) 2 at 25° C. 34 xii
Page No.
3.4 Log activity pH diagram for
the magnesium arsenic(v) water
system. 36
3 . 5 The solubility of calcium
arsenate (Ca3 (As04) 2) at
various pH's. 37
3 . 6 The variation of arsenic(v)
and calcium concentrations
with pH in equilibrium with
Ca3 (AS04) 2 at 25° C. 39
3.7 Log activity pH diagram for
the calcium(ii) arsenic(v)
water system. 43
3 . 8 Log activity pH diagram for
the calcium(ii) arsenic(iii)
water system. 44
3.9 Log activity pH diagram for
the barium(ii) arsenic(v)
water system. 47
3 . 10 Log activity pH diagram for
the iron(ii) arsenic(v) water
system. 49
3.11 The variation of arsenic(v)
and iron(iii) concentrations
with pH in equilibrium with
FeAs04 at 25°C. 52 xiii
Page No .
3 . 12 Log activity pH diagram for
the iron(iii) arsenic(v) water
system. 55
3.13 Potential pH diagram for the
arsenic sulphur water system. 58
3.14 Potential pH diagram for the
iron arsenic sulphur water
system. 60
3.15 Potential pH diagram for the
iron arsenic sulphur water
system at 25° c at
concentrations for iron,
arsenic, and sulphur at lQ- 3,
10- 6 , and 10- 4 M respectively. 62
3.16 Potential pH diagram for the
1.ron arsenic sulphur water
system at a c oncentration of
unity for all species. 63
CHAPTER 4
4.1 PH correction versus measured
pH graph. 69
CHAPTER 5
5 . 1 PH titration curves for the
arsenic(iii) sulphur(-ii)
water system. 73 xiv
Page No.
5.2 Absorbance versus pH curves at
400nm for the arsenic(iii)
sulphur(-ii) water system. 74
5.3 Log activity pH diagram for
the arsenic(iii) sulphur(-ii)
water system. 76
5.4 Ionic distribution curves .for
the arsenic(iii) sulphur(-ii)
water system at a
concentration of O.lM with
respect to arsenic(iii). 77
5.5 Ionic distribution curves for
the arsenic(iii) sulphur(-ii)
water system at a
concentration of O.OlM with
respect to arsenic(iii). 78
5.6 Ionic distribution curves for
the arse nic(iii) sulphur(-ii)
water system at a
concentration of O.OOlM with
respect to arsenic(iii). 79
5.7 Ionic distribution curves for
the arsenic(iii) sulphur(-ii)
water system at a
concentration of O.OOOlM with
respect to arsenic(iii). 80 XV
Page No.
5.8 PH titration curves for the
iron(ii) arsenic(v) water
system at a ratio of 1:1 for
iron(ii) and arsenic(v) in the
perchlorate medium. 85
5.9 PH titration curves for the
iron(ii) arsenic(v) water
system at a ratio of 1:1 for
iron(ii) and arsenic(v) in the
sulphate medium . 89
5.10 Log activity pH diagram for
the iron(ii) arsenic(v) water
system in the perchlorate
medium taking into
consideration atmospheric
carbon dioxide. 92
5.11 Ionic distribution curves for
the iron(ii) arsenic(v) water
system in the perchlorate
medium at a concentration of
0.2M with respect to iron(ii) . 93
5.12 Ionic distribution curves for
the iron(ii) arsenic(v) water
system in the perchlorate
medium at a concentration of
0.1M with respect to iron(ii). 94 xvi
Page No .
5.13 Ionic distribution curves for
the iron(ii) arsenic(v) water
system 1n the perchlorate
medium at a concentration of
0.01M with respect to
iron(ii). 95
5.14 Ionic distribution curves for
the iron(ii) arsenic(v) water
system in the perchlorate
medium at a concentration of
0.001M with respect to
iron(ii). 96 5.15 Log activity pH diagram for
the iron(ii) arsenic(v) water
system in the perchlorate
medium. The ratio of iron(ii)
to arsenic(v) is 3:2. 99
5 . 16 Log activity pH diagram for
the iron(ii) arsenic(v) water
system in the perchlorate
medium. The ratio of iron(ii)
to arsenic(v) is 4:3. 100
5.17 Log activity pH diagram for
the iron(ii) arsenic(v) water
system in the sulphate medium. xvii
Page No.
The ratio of iron(ii) to
arsenic(v) is 1:1. 101
5.18 Ionic distribution curves for
the iron(ii) arsenic(v) water
system in the sulphate medium
at a concentration of 0.2M
with respect to iron(ii). 103
5.19 Ionic distribution curves for
the iron(ii) arsenic(v) water
system in the sulphate medium
at a concentration of 0.1M
with respect to iron(ii). 104
5 . 20 Ionic distribution curves for
the iron(ii) arsenic(v) water
system in the sulphate medium
at a concentration of 0.01M
with respect to iron(ii). 105
5.21 Ionic distribution curves for
the iron(ii) arsenic(v) water
system in the sulphate medium
at a concentration of 0.001M
with respect to iron(ii). 106
5.22 PH titration curves for the
iron(iii) arsenic(v) water
system in the perchlorate
medium. 109 xviii
Page No.
5.23 PH titration curves for the
iron(iii) arsenic(v) water
system in the sulphate medium. 110
5.24 Turbidity as a function of pH
for the iron(iii) arsenic(v)
water system in the
perchlorate medium. 115
5.25 Turbidity as a function of pH
for the iron(iii) arsenic(v)
water system in the
chloride medium. 117
5.26 Turbidity as a function of pH
for the iron(iii) arsenic(v)
water system in the
nitrate medium. 120
5.27 Turbidity as a function of pH
for the iron(iii) arsenic(v)
water system in the
sulphate medium. 123
5.28 Turbidity as a function of pH
for the iron(iii) arsenic(v)
water system. Back titrating
with sulphuric acid in the
sulphate medium. 124
5.29 Turbidity as a function of pH
for the iron(iii) arsenic(v) xix
Page No.
water system. Back titrating
with perchloric acid in the
sulphate medium. 125 5.30 Turbidity as a function of pH
for the iron(iii) arsenic(v)
water system in the nitrate
medium with an ionic strength
correction of one. 128 5.31 Spectrophotometric scans of ferric perchlorate and arsenic
acid mixed at a ratio of 1:1. 131
5 . 32 Spectrophotometric scans of
ferric sulphate and arsenic
acid mixed at a ratio of 1:1. 132
5·. 33 Spectrophotometric scans
showing ferric s ulphate
complexes . 134
5.34 Log activity pH diagram for
the iron(iii) arsenic(v) water
system in the perchlorate
medium. 136
5.35 Ionic distribution curves for
t h e iron(iii) arsenic(v) water
system at a concentration of
0.01M with respect to
iron(iii). 1 38 XX
Page No.
5.36 Ionic distribution curves for
the iron(iii} arsenic(v} water
system at a concentration of
0.001M with respect to
iron(iii} . 140
5.37 Log activity pH diagram for
the iron(iii} arsenic(v} water
system in the nitrate medium. 143
5.38 Log activity pH diagram for
the iron(iii} arsenic(v} water
system in the chloride medium. 146
5.39 Log activity pH diagram for
the iron(iii} arsenic(v) water
system in the perchlorate
medium. 150
5.40 Log activity pH diagram for
the iron(iii) arsenic(v} water
system in the perchlorate
medium showing the stability
region of crystalline goethite
(FeOOH}. 153
5.41 Potential pH diagram for the
iron arsenic sulphur water
system at a concentration of
0.0001M for all aqueous
species . 158 xxi
Pa ge No.
5 . 42 Potential pH diagram for the
iron arsenic water system at a
concentration of O.OOOlM for
all aqueous species. 160 xxii
LIST OF TABLES
Page No.
CHAPTER 1
1.1 1981 United States markets for
arsenious oxide (As203} [ 8] . 3
1.2 Some naturally occurring
minerals of arsenic [ 3 ] . 5
1.3 Physical Properties of a rsen~c
[ 5] . 6
1.4 More common arsenic compounds
[ 11] . 8
1.5 Total imports of arsenious
oxide (As2 Oa } to Australia
[ 20] . 10
1.6 Composition of flue dust [ 19] . 13
CHAPTER 2
2.1 Standard free energy of
formation data for the arsenic
species at 25° c. 20
2.2 Standard free energy of
formation data for the sulphur
species at 25° c. 26
2.3 Standard free energy of
formation d ata for the iron
species at 25° c . 29 xxiii
Page No.
CHAPTER 3
3.1 Various calcium arsenites and
arsenates [62] . 40
3.2 Various free energies for
calcium arsenites/arsenates. 41
3. 3 Free energies of Ba3 (As04 ) z
from various sources. 45
3.4 List of X-ray diffraction
patterns for ferric arsenate. 50
3.5 Common X-ray diffraction
patterns. 51
3.6 Free energy of formation data
for ferric arsenate from
various literature sources. 54
3.7 Various free energy of
formation data for orpiment
( Asz S3 ) • 57
CHAPTER 5
5.1 Free energy of formation data
for the arsenic(iii)
sulphur(-ii) water system. 83
5.2 PH at the point of
precipitation under the
different conditions specified
for the iron(ii) arsenic (v) xxiv
Page No.
water system l.n the
perchlorate medium . 87
5.3 PH at the point of
precipitation under the
different conditions specified
for the iron(ii) arsenic(v)
wate+ system in the sulphate
medium. 90
5.4 Set of free energies of
formation for the iron(ii)
arsenic(v) water system. 107 5.5 Corrected pH at the point of
precipitation and the various
concentrations of iron(iii)
and arsenic(v) in the
perchlorate medium. 116
5.6 Corrected pH at the point of
precipitation and the various
concentrations of iron(iii)
and arsenic(v) in the chloride
medium. 119
5.7 Corrected pH at the point of
precipitation and the various
concentrations of iron(iii)
and arsenic(v) in the nitrate
medium. 121 XXV
Page No.
5.8 Corrected pH at the point of
precipitation and the various
concentrations of iron(iii)
and arsenic(v) in the sulphate
medium. 126
5.9 Corrected pH at the point of
precipitation and the var1ous
concentrations of iron(iii)
and arsenic(v) in the nitrate
medium with an ionic strength
of 1. 0. 129
5.10 Set of free energies for the
iron(iii) arsenic(v) water
system in the perchlorate
medium. 142
5.11 Set of free energies for the
iron(iii) arsenic(v) water
system in the nitrate medium. 144
5.12 Set of free energies for the
iron(iii) arsenic(v) water
system in the chloride medium. 147
5.13 Iron(iii) sulphate complexes
from various sources. 149
5.14 Standard free energies of
formation data for the iron xxvi
Page No . arsenic sulphur species at
25° c . 155 1
CHAPTER 1. INTRODUCTION
1.1 Historical
The term arsenic was derived from the Greek name arsenikon meaning bold, valiant, and masculine
[1]. The Greeks had originally given the name to the sulphi de minerals containing arsenic, as a result of their reactivity with other metals. These sulphide mineral s were later identified with other arsenic minerals and the name arsenikon was transferred to arsenic itself when it was isolated [2].
Although it is not clear who actually isolated metallic arsenic this has been accredited to Albert the
·- Great in 1250, Paraclesus in the 16th century, and
Shroeder in 1649 [1] .
1.2 Uses of Arsenic
Arsenic compounds have been used from the earliest times. The naturally occurring pigments drew attention in the Copper and Bronze Age cultures.
Arsenical pigments were also used for decoration in
Egyptian tombs and arsenic was mined as a dye by the
Persians [3]. The Chinese 2000-3000 years ago were aware of its toxicity, and devised means of detecting its use in suspected poisonings. Records have also made it clear 2
that arsenic compounds were used extensively by the professional poisoners of the Middle Ages [4-6] .
Arsenic has been used in local applications in medicine, for example in the form of ointments in the treatment of ulcers and plague as far back as the second century. At the same time the knowledge gained about the toxic properties of arsenic led to its widespread use as a poison.
In more recent times arsenic and its compounds have been used extensively in pigments, pesticides, herbicides, wood preservatives, and as semi conductors.
Some uses of arsenic have been based on its chemical and physical properties, but over eighty percent of its uses have been based on its toxicity [5]. In agriculture, arsenic compounds were used as insecticides which were mainly calcium arsenate (Ca3 (As04)2) and lead arsenate (Pb3 (As04 )2) [7]. Metallic arsenic was used as an alloy with lead in the making of shot and smaller percentages are often used in copper and brass .
The use of arsenic in some applications has declined due to the imposition of environmental regulations [8]. Table 1.1 lists some of the most common commercial applications of arsenic in 1981 [8].
1.3 Occurrence.
Arsenic, the 20th most abundant element, 3
TABLE 1.1
1981 UNITED STATES MARKETS FOR ARSENIOUS OXIDE {As203)
[ 8] .
USE %
- Herbicides 31
Wood Preservatives 36
Glass 5
Flotation 8
Cotton Dessicants 15
Miscellaneous 5 4
composes an average of 0.0005 percent of the earth's crust [3] . It is known to appear in 245 minerals[9,10] with the most common ores being arsenic sulfides [11] . Some of these sulphide minerals and others are listed in table 1.2 [3]. A review of the geochemistry of arsenic has been presented by Onishi and co-workers [12] .
Under neutral pH conditions arsenic containing minerals are only slightly soluble, but through chemica l processing for the recovery of valuable metals they can be transformed into more soluble arsenic compounds.
1.4 Chemistry of Arsenic
Arsenic is the third member in the nitrogen family of elements in group 15 (1985 IUPAC
Recommendation) of the periodic table. It has four stable oxidation states -3, 0, +3, and +5 [ 13] . The toxicity of arsenic increases markedly when reduced from a +5 to a -3 oxidation state [4] .
Naturally occurring arsenic consists as mainly one stable isotope As74.9216. The metallic form is a steel-gray, crystalline solid with a brillant luster
[14] . It is a good conductor of heat and a poor c onductor of electricity. Some of the physical properties of arsenic are listed in table 1.3 [5].
A large variety of arsenic compounds exist.
Arsenic forms hydrides of which the most common being 5
TABLE 1.2
SOME NATURALLY OCCURRING MINERALS OF ARSENIC (3] .
Mineral Name Formula
Arsenolite As2 Oa
Arsenopyrite FeAsS
Cobaltite CoAsS
Enargite Cua AsS4
Kupfernickel NiAsS
Loellingite FeAs2
Mimetite Pb~Cl(AS04 )a
Mispickel FeSAs
Niccolite NiAs
Olivenite Cu2 OHAS04
Orpiment As2 Sa
Proustite Aga AsS2
Realgar AsS
Scorodite FeAs04 .2H20
Smaltite CoAs2
Te nnantite Cue As2 S7 6
TABLE 1.3
PHYSICAL PROPERTIES OF ARSENIC [5].
Physical Property Value
Atomic Number 33
Atomic Weight 74.91
Melting Point at 36 atm oc 814
Boiling Point °C 604.3(sublimes)
Density at 20°C g/cm 5 .73
Specific Heat cal/gm/°C .0822
Latent heat of sublimation
cal/gm 60
Crystal System Hexagonal (rhombohedral)
Hardness, Mohs Scale 3.5
Hardness, Brinell 147 7
arsine (ASH3 ) . It also forms sulphides, halides, numerous organic compounds and many others [5,11,14-18].
The most common compounds are listed in table 1.4 [11].
On exposure to dry air, arsenic undergoes no reaction, but in the presence of moisture, arsenious oxide (As203) is slowly produced [19] . The most important commercial arsenic compound is arsenious oxide [3]. When arsenious oxide is dissolved in water it forms arsenious acid (H3AS03). Anoth~r oxide of arsenic, which is of commercial importance is arsenic oxide
(As2 0!!) (11] . In water, arsenic oxide forms arsenic acid
(H3 AS04 ) . The salts of arsenic acid and arsenious acid are known as arsenates and arsenites respectively.
Arsenic generally behaves as an anion in the form of arsenites and arsenates. It does not react to form carbonates, bicarbonates, or phosphates.
1.5 Economic Aspects
Arsenic is produced as a by-product from the processing of arsenic containing ores, but due to strict environmental regulations imposed on the storage and disposal of arsenic compounds, a safe means of storing and disposing of these compounds has become essential.
Most arsenic in Australia is imported in the chemical form of arsenious oxide (As203). Table 1. 5 illustrates the amount of arsenious oxide imported to 8
TABLE 1.4
MORE COMMON ARSENIC COMPOUNDS [11] .
Classification Formula
Hydrides AS Arsenides Fe2 As, Fe3 As2 , FeAs, CU2AS, Cu3 As. Halides ASF3 , ASF6 AsC13 AsBr3 As2 I Oxyhalides AsOCl, AsOBr Oxides As2 03 , As2 o~ Acids H3 AS03 , H3 AS04 Sulphides As2 S3 , As2 S2 , As2 S~ . 9 Australia during the period 1958-1982 [ 20] . These figures represent the total Australian consumption of arsenious oxide. As shown in table 1.5 there has been a decline in the consumption of arsenic since the 1970's. With increases in environmental pressures, and the major use of arsenious oxide being in agriculture, the sale of arsenicals for this purpose has declined. At the same time extensive environmental studies were undertaken to determine whether arsenic was safe both in the work place and in the applications which have dispersed it into the environment [8]. During the 1970's the applications to which arsenic was applied have also changed, for example arsenic used in the glass industry has been largely replaced and is no longer a major market. The use of arsenic for metal alloying has shifted to arsenic metal rather than the trioxide. It is however in the field of agricultural chemicals that environmental restrictions have had the greatest effect [8]. A situation has thus developed whereby the arsenic output is in excess of demand and the problem of disposing large tonnages of arsenic safely, has become increasingly important. 1.6 Industrial Applications. The majority of arsenic entering into the 10 TABLE 1.5 TOTAL IMPORTS OF ARSENIOUS OXIDE (As~_ Oa) TO AUSTRALIA [20] . YEAR TONNES 1958-59 2,637 1959-60 2,660 1965-66 1,217 1969-70 1 , 237 1978-79 745 1980-81 1,118 1981-82 1,073. 11 world's waters, is contributed by industrialisation and from arsenic containing herbicides [15]. However it must also be taken into account that an appreciable amount of arsenic enters the environment because of smelting operations. The mineral industry contributes 23,600 tons per year of total arsenic emission to the atmosphere where as 7,800 tonnes per year is from natural sources [21] . Arsenic is produced as a by-product in two main smelting operations: 1. In the processing of gold bearing sulphides [22-25]. 2. In the copper, cobalt, and nickel smelting industry [9,~6 - 37]. Owing to the volatility of arsenic metal and its sulphides, and trioxide, arsenic is found in the flue dust and off-gases from roasting operations [3,5,7,28,37,38]. In the processing of gold bearing sulphides, arsenic is an important source of contamination. The contamination of water by arsenic occurs after the roasting stage where arsenic is transformed into a finely divided powder of arsenious oxide. This is then discharged into the environment in two ways: 1 . It can escape as fumes with the combustion gases, or 2. It is retained, the g o ld b e a r ing 12 particles are trapped 1n an electrostatic precipitator and subsequently leached in the presence of cyanide. It is the latter that is responsible for the solubilisation of most of the arsenic . Arsenious oxide is extremely soluble at basic pH levels and it is in this range that cyanide leaching of gold occurs [39]. In the copper, cobalt, and nickel smelting industry, waste products are slags, flue dusts, and sludges[26]. Most of the arsenic is accumulated in the flue dust and part of it is carried over into the gas phase as the trioxide. A typical composition of flue dust is shown in table 1.6 [19]. Crude flue dust can contain up to 30 per cent of arsenious oxide. In the mining industry the removal and subsequent safe disposal of arsenic is achieved by the direct precipitation of arsenic upon the addition of a metal cation to form an insoluble metal arsenate, or arsenite, or the addition of sulphide to form arsenic sulphide (As2Sa). One such accepted process is the addition of ferric ion to produce ferric arsenate [9,22,24,26,33,35,37]. Initially the slurry of the ore is leached in an acidic solution under oxidising conditions at elevated temperature and pressure. Thereby not only are the acid soluble metals dissolved but the 13 TABLE 1.6 COMPOSITION OF FLUE DUST [19]. ELEMENT % Arsenic 17.5 Copper 8.2 Zinc 5.9 Lead 4.1 Iron 12.7 Cadmium 0.3 Antimony 0.35 Bismuth 0.38 Tin 0.48 Sulphur(total) 9.3 Silver 7.4 oz/t 14 arsenic 1s also dissolved. An equivalent amount of iron is converted by the dissolved arsenic to sparingly soluble ferric arsenate. Any additional dissolved arsenic is precipitated by the addition of a cation which forms an insoluble arsenate. The slurry is then cooled, and the solution separated from the residue [24] . Another accepted method is the addition of lime to stabilise the arsenic in the form of calcium arsenate ( Ca3 ( As04 ) z ) [9,27 ,28,39,40]. The slurry containing dissolved arsenic is saturated with calcium ions in the form of calcium hydroxide (Ca(OH)z) which induces the precipitation of either calcium arsenate or calcium arsenite which are both almost insoluble in water. Due to the low solubility of arsenic sulphide the addition of sulphide to form arsenic sulphide (AszS3) is used extensively. In the copper smelting industry the arsenic in the gas phase is scrubbed using dilute sulphuric acid and is often recovered as arsenic sulphide (AszS3) by the addition of hydrogen bisulphide, or sodium hydrosulphide to the solution [26]. In Japan at the Toyo smelter the low solubility of arsenic sulphide has been accepted and the arsenic is precipitated as arsenic sulphide. The arsenic sulphide is then used to produce arsenious oxide. Kondo [41] has reviewed the recovery and fixation of arsenic from metallurgical intermediates and refers to a number of operating processes where sulphide 1 ~ precipitation of arsenic sulphide (As2S3) is used. Another main area where arsenic occurs as an impurity is in the processing of uranium ores. Since most uranium ores contain arsenic and with an increase in the use of uranium, arsenic again poses an impurity problem. A patent presented by the Key Lake Mining Corporation and Sherritt Gordon Mines [42,43], claims that the addition of lime and/or ferric sulphate precipitates the dissolved arsenic as calcium arsenate and/or ferric arsenate. This then reduces the concentration of dissolved arsenic to levels of 2 mg/1 in the pH range of 11-12. Then, by the addition of a barium salt the remaining arsenic in solution is precipitated as barium arsenate (Ba3 (As04 )2) and thus the concentration of dissolved arsenic can be decreased even further. The final concentration of arsenic can be less than 0.5 mg/1 and at high pH values probably less than 0.2 mg/1. This process is especially useful in the recovery of uranium from uranium ores which contain both arsenic and iron. Industrialisation has altered the natural geochemical cycle of many trace elements into the environment. Thus a need exists to understand what factors control the concentrations of these elements and to obtain comprehensive thermochemical data in order to understand the various reactions of arsenic in aqueous systems. 16 The efficiency of removing arsenic from smelter gases has increased and the use of arsenic in pesticides and herbicides has been inhibited by environmental restrictions . The use of arsenic as wood preservatives however has expanded rapidly around the world. Since arsenic is a major impurity in sulphide ores and the arsenic in the recovered dusts and fumes will excee~ world demand for arsenic chemicals and commodities, a means of safe disposal of arsenical residues has become important. Currently, the main method for disposing arsenic is by stabilising and precipitating it in the form of an insoluble arsenic compound. This is then contained in dumps, tailing ponds, and other forms of underground storage. 17 CHAPTER 2. STABILITY DIAGRAMS 2.1 Theory Thermodynamic stability diagrams provide a useful basis for predicting under different environmental conditions the regions where the various dissolved species and solids exist in a given system. Such thermodynamic stability diagrams have been presented by Pourbaix [44]. These diagrams have become increasingly important in explaining likely reactions in various systems. To generate these diagrams it is essential to have a detailed knowledge of the various species in the system together with the free energy of formation for each species. When considering a solid in equilibrium with dissolved ionic species, the solubility 1s generally dependant on the pH of the solution. The solublity of a solid can be measured by the concentration of the ions in equilibrium with it. Thermodynamic theory allows a relationship between ion activity and pH to be developed and this relationship is linear. Equilibrium lines, showing stability limits, can be drawn between two dissolved species, a solid and dissolved species and two solids and associated dissolved species. The equilibrium relation for an aqueous 18 solution can be repre sented by the general reduction reaction , --(2.1). The standard reaction isotherm for this reaction is: 6 G - 6G0 = RTlnK ------(2.2) By substituting and manipulating the following : where 6G = -zFE (kcals) E = potential (volts) F = 23 . 0609 (kcal/volt) K = ratio of activ ities asb = aH?Oc aAa aH+ x and if R = 1.98717 (ca l / deg .mol) and T = 298 . 15 K then assuming aH 0 = 1 2 and defining· pH = -log aH+ equation 2.2 becomes : zE 6G0 aob = xpH - log (2 . 3} 0 . 05916 1.3642 a A 8 For the reaction in which no ox idation states are altered i.e. z=O, equation 2 . 3 becomes a s b xpH = log ---(2. 4) 1.3642 6G0 for any reaction is constant and can be calculated accordingly: = 6 G0 products - 6 G0 reactants 19 Stability diagrams provide a useful device for illustrating the stability fields for the different species in an aqueous environment. The diagrams used in this study were generated by computer using selected free energy data from the literature as well as derived from experimental results . 2.2 Arsenic Water System Arsenic chemistry in aqueous systems is quite complicated with oxidation-reduction, hydrolysis, complexation, precipitation and adsorption reactions all taking place. Arsenic can be stable in four oxidation states +5, +3, 0, -3) under different Eh conditions, but the +5 and +3 oxidation states are the. most commonly encountered in solution. A number of aqueous species of arsenic exist, the most important for arsenic{v) are As04 3-, HAs04 2-, H2AS04-, and H3AS04 and for arsenic{iii) As03 3-, HAS032- , H2AsOa- and HaAS03. Pourbaix [44] has constructed a potential pH diagram which can be used to explain the oxidation reduction and hydrolysis reactions. More recent free energy of formation data for the various arsenic species at 25°C are listed in table 2.1. From this information a slightly modified potential pH diagram was obtained and is shown in figure 2.1. Log activity versus pH diagrams were also obtained using the same free energy data and these are 20 TABLE 2.1 STANDARD FREE ENERGY OF FORMATION DATA FOR THE ARSENIC SPECIES AT 25°C. Species Free Energy Data Source ( f:..Gt , 298.15K Kcal/mol) H3 Aso.. {aq} -183.08 [ 45] H2 Aso.. - - 180.01 [ 4 5] HAs0,. 2- -170.79 [45] As04 3- -154.97 [ 45 J AsO+ -39.15 [45] H3AS03 (aq) - 152. 92 [ 45) H2 As03- -140. 33 [45) HAs03 2 - -121 . 71 [46) AS03 3 - -102.95 [46) AsH3 {g) +16 . 47 [ 45) As2 03 ( s) -137.72 [ 45] As2 0 :~ {s) - 186.97 [45) As{s) 0 .0 [ 45 J ARSENIC WRTER SISTEM RT 0. 1M 2. 0.------.------.------,------, A H3 AS Oil CAQ) B H2 AS Oil E CACl C H AS OY E2 CRQl AS OY E3 CAGil H3 AS 03 CAGil H2 AS 03 E CACl 1. AS 03 E3 CAGil AS H3 CGl AS ..c w H -2. 0 ~------~-.----~----8~.----~----1~2----~--~ PH Figure 2. 1. Potential pH diagram for the arsenic water system a·t 25°0. 22 shown 1n figures 2.2 and 2.3 for the arsenic(iii) and arsenic(v) systems respectively. 2.3 Sulphur Water System The most common ores containing arsenic are the sulphides. The study of the chemistry of sulphur species and their removal is extremely important in aqueous systems. Several of the thermodynamically stable species are environmentally sensitive the more noted ones being sulphides and to a lesser degree, sulphur and sulphate. Of the many ionic and molecular sulphur species that exist, only six are thermodynamically stable in aqueous solutions at room temperature and one atmosphere pressure namely HS04-, S04 2 -, S(s), H2S ( aq )~ Hs-, and S2 - [47]. Other species such as thiosulphates, polysulphides and polythionates also appear in the natural environment, but they are thermodynamically unstable [ 48] • A potential pH diagram for the stable sulphur species in aqueous solution is shown in figure 2.4. The free energy of formation data used to produce this diagram is listed in table 2.2. 2.4 Iron Water System As the fourth most abundant element in the earth's crust, iron becomes a very important component in aqueous systems . r:z9a.ls 0 1 H3Rs Cl3 ( RQ l 2 H2As Cl3·- ~------6------~~ 3 HAs Cl3-- 1 .. 4 As Cl3-- .. 5 Rsrl+ ·· 6 Rs2Cl3 (5) 2 pH2Cl:O/O/O .. pH: 0/0/. 25 ·· pAs {I I I l : 0/. 25/0 3 2 3 1 4 --s (/) ~6~------~~~~0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH Figure 2.2. Log activity pH diagram for the arsenic(iii) N hydroxy species. VJ 0 r:298.1S 1 H3Rs(j4 (RQ.l 2 H2Rs(j4- 3 HRs (j4-- 1 4 Rs (j 4·--- 5Rs2(j5(5) pH2(j:O/O/O 2 pH:0/0/.25 pRs (VJ: 0/.25/0 3 1 2 3 4 4 >-s (/) ~6~------~~~--~~~~~~~~0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH Figure 2.3. Log activity pH diagram tor the arsenic(v) hydroxy species. SULPHUR WRTER SYST EM RT 0. 1M 2. 0r---r------.,A H S Oll E CAQl B S Oll E2 CRQI c s D H2 S (RQI E H S E CACl F S E2 tRCI 1. .c. w F PH Figure 2. 4. Potential pH diagram for the sulphur water N system at 25°C. V1 26 TABLE 2.2 STANDARD FREE ENERGY OF FORMATION DATA FOR THE SULPHUR SPECIES AT 25°C. ' ' Species Free Energy Data Source . ( 6Gt , 298.15K Kcal/mol) HS04- -180.67 [45] S04 2- -177 .95 [ 45] S(s) 0.0 [ 45] H2 S ( aq) -6.65 [ 45] HS- +2.89 [45] S2- +21.17 [49] 27 The main aqueous hydroxy species for Fe2 + are FeOW, Fe(OH}z (aq}, Fe(OH}a- and Fe(OH}4 2 - and for Fea+ they are FeOH2 +, Fe(OH)z+, Fe(OH)a (aq), Fe(OH)4 - and the polynuclear ions Fez (OH} z 4 + and Fea (OH) 4 ° + • A potential pH diagram for the iron water system is shown in figure 2.5. The reactions considered have been listed elsewhere [50] . The free energies of formation used in computing this potential pH diagram are listed in table 2.3. IR ON WATER SYSTEM AT 0. 1M 2. 0 .---~------~ A FE E-2 !AQl B FE 0 H E·l CAQl C FE [0 Hl2 D FE E-3 !AQl D E FE 02 H F FE 1. E A £ w F PH Figure 2. 5. Potential pH diagram tor the iron water N co system at 25°C. 29 TABLE 2.3 STANDARD FREE ENERGY OF FORMATION DATA FOR THE IRON SPECIES AT 25°C. Species Free Energy Data Source. ( /::, Gt , 298.15K kcal/rnol) Fe2 + -18 .86 [ 45] FeOW -66.3 [45] Fe(OH)2 (aq) -104.24 [49] Fe (OH) 2 (s) -114 . 7 [ 49] Fe(OH)a- -146.96 [ 45] Fe(OH) 42- -183 . 96 [ 45] Fe3 + -1.12 [ 45] FeOH2 + -54.83 [45] Fe(OH)z+ -104.68 [ 45] Fe (OH) 3 (aq) -157 . 57 [45] Fe(OH)4- -198.39 [49] Fez (OH)z 4 + -111.7 [ 45] Fe3 ( OH) 4 11 + -221.46 [49] FeOOH(s) -109.82 [51] Fe(s) 0.0 [ 45] 30 CHAPTER 3. INSOLUBLE ARSENIC COMPOUNDS 3.1 Metal Arsenates/Arsenites Most of the solubility data for metal arsenates and metal arsenites have been based on the work by Chukhlansev in 1956-57 [52-54] . These results have been used in several compilations of thermodynamic data both in the form of solubility products and free energies of formation [45,55]. From Chukhlansev's work a general linear relationship between log solubility and pH can be observed, for both metal arsenates and arsenites and this relationship can be seen in figures 3.1 and 3.2 respectively [52-54]. This linear relationship has been used by other workers [56] to obtain solubilities at higher pH values by extrapolating these lines and it is this extrapolation that leads to the misconception that some of these metal arsenates and metal arsenites are extremely insoluble under higher pH conditions than in the original measurements. Direct extrapolation of solubility products without considering the particular equilibria involved can lead to large errors when assessing the solubility at various pH values. Even though earlier work by Guerin [57], Masson 31 \ 1·0 > -VI Fe . \ ·01 ~\ y Hg •• 8o \\ ~.\ ·0 01 1 2 3 4 5 6 7 8 9 p~ Figure 3. 1. Solubilities of some metal arsenic(v) compounds from Chukhlansev. 32 5 1 Zn 0.5 Ag Pb Pb \ \ . '· \ \...... \ O'l 0.1 ...... ~ 0.05 <( 0.01 0.005 0.001 L___.l__....J..__ __L.__ _L.__ ___.L..._ ___L_ ___...L_ ____.J 1 2 3 4 5 6 7 8 9 pH Figure 3.2. Solubilities of some metal arsenic(iii) compounds from Chukhlansev. 33 [58 J , and Mas [59], reported solubility data for metal arsenates, it is Chukhlansev's work that has been widely accepted. Recently, Nishimura and Tozawa [60] have reported a variation to this linear relationship for calcium, magnesium and ferric arsenates. At higher pH values than was used by Chukhlansev the linear -relationship does not exist and the solubility of the various metal arsenates could be lower by two orders of magnitude. 3.1.1 Magnesium Arsenic Water System Magnesium arsenate has been formed as a residual using either magnesium oxide or hydroxide, as a safe means for disposing of the arsenic. Work by Nishimura and Tozawa [60] in precipitating magnesium arsenate determined the variation of arsenic and magnesium concentrations with pH at 25°C. At pH values above 8 the linear relationship assumed from Chukhlansev's work was no longer the case and the solubility increased significantly. This is shown in figure 3.3 [60]. Further work by Nishimura and Robins [61] concluded that over a wide pH range there are two stable magnesium arsenates, MgHAs04 and Mga (As04)2 which are both in hydrate form at 25°C, and that no low solubility 34 5 As.. \ ~ ·~ 05 >.., ...... 0.1 0' 0.01 0.005 o.ooiGL -----l7~-.-_ _..a__ _..9 __ ...... 1o__ ...... ll __ '"':"2 pH Figure 3.3. The variation of arsenic(v) and rnagnesium concentrations with pH in equilibrium with Mg3(As04)2 at 25°C. Nishimura ------· Chukhlansev. 35 magnesium arsenite compounds are formed. Mga (AS04 ) 2 in the absence of carbon dioxide is stable but is effected by atmospheric carbon dioxide above pH values of 9.3 where it decomposes to yield the basic magnesium carbonate (4MgCOa .Mg(OH)2 .4H20) as shown by the stability diagram in figure 3.4. Free energies of formation for the various species are listed in Appendix 1 . 3.1. 2 Calcium Arsenic Water System The use of lime to precipitate the arsenic as calcium arsenate has been and is still being used extensively in a variety of industrial applications. The solubility product and free energy of formation found in the literature for calcium arsenate (Caa (AsO" l 2 l originates from the work of Chukhlansev [53]. It is this work that has resulted in a wide acceptance that calcium arsenate ( Caa ( AS04 ) 2 ) is extremely insoluble and therefore ideal to stabilise arsenic in that form. In his work Chukhlansev only measured the calcium concentration, figure 3 . 5 represents the solubility of Ca3 (As04)2 at various pH values [53]. Nishimura and Tozawa [60] investigated the solubility of calcium arsenate over a wide pH range , by measuring both calcium and arsenic concentrations in 0 T:298.1S 1 H3RsCI4 CRQJ I 2 H2Rs Cl4- I 3 HRs Cl4-- 1 I 4 Rs Cl4--- 10 I 5 Mg++ I 6 Mg ClH+ 2 I 7 Mg CC1Hl2(RQJ I 8Mg(CIHl2CSJ I 9MsHR~CI4CSl I 3 10 Ms3 CR~Cl4l 2 (Sl 8 pH2CI : O/O/O ,...... ; , 5 : ' ... pH : 0/0/.25 ...... pRs(VJ:0/.25/0 ::4 pMg (Ill :0/.25/0 C) ~ 0... -s > Cl) ~6 6 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH Figure 3.4. Log activity pH diagram for the rhagnesium(ii) arsenic(v) water system. (The effect of atmospheric carbon dioxide is to cause the basic carbonate to be stable to the upper right of the dashed line, and the neutral dissolved spec;ies to be stable below the dashed line.) - --4MgC03Mg(OH)2 4H20(s) ...... MgC03 (aq) 37 -1 >- . t- 0::: ::E: ::L :::::> -2 L.) __J <..!:) c:::::> __J Figure 3.5. The solubility of calcium arsenate (CaiAs0 4) 2) at various pH's. 38 solution. Their results showed a minimum solubility at about pH 8 which was significantly higher -than previously accepted, this is shown in figure 3.6 [60]. Work by Nishimura, Tozawa and Robins [62] resulted in the identification of two calcium arsenites and five calcium arsenates and their solubility determined as a function of pH. These compounds are listed in table 3.1. Numerous calcium arsenates and arsenites are mentioned within the literature and these vary widely [63-68]. The calcium arsenites are more soluble than the calcium arsenates. Table 3.2 shows the variation in free energies of formation calculated from various sources. Further work by Nishimura and co workers [62,70] suggested that the results were influenced by carbon dioxide in the atmosphere which caused a decrease in the stability of calcium arsenate and its decomposition to calcium carbonate. Robins and Tozawa [71] interpreted these results in the form of stability diagrams. It was determined that at a pH above 8.3 calcium carbonate was more stable than calcium arsenite and calcium arsenate. Further work by Robins and Tozawa [71] showed that as the carbon dioxide pressure or carbonate concentrations were increased the solubility of calcium arsenate was increased significantly. Stability diagrams taking into account 39 5 0.5 ' 0.1 0' -Cl) ~ 0.05 u0 0.01 • 0.00 • • 0.001 '-----'-___..___ _ _.~.__ _..__ _ _.li....-----'---li..- 12 . 13 7 8 9 10 II 6 pH Figure 3.6. The variation of arsenicfv) and calcium concentrations with pH in equilibrium with Ca3(As04) 2 at 25°C. Nishimura -·-··-·- Chukhlansev 40 TABLE 3.1 VARIOUS CALCIUM ARSENITES AND ARSENATES [62]. Calcium Arsenites Ca(AsOz)z Ca(AsOz ) z .Ca(OH)z Calcium Arsenates Caz Hz (As04 ) z Cao Hz (AS04 ) 4 Ca3 (As04 ) z Ca3 (As04 )z .Ca(OHh CaH4 ( AS04 ) z TABLE 3. 2 VARIOUS FREE ENERGI ES FOR CALC IUM ARSENITES/ARSENATE S. t,Go (Kcal/mole) f Compounds Nishi mu r a Guerin Chukhlansev Nishimur a Gorochova et.al & Tozawa e t . al. [62] [57) [53] [60 ] [ 69] 25°C l7°C 20 °C 30 °C 298K Ca(As0 ) -~308. 84 2 2 Ca(As0 ) . Ca(OH) -531.91 2 2 2 CaH (As0 ) -490.91 -491.26 4 4 2 Ca H (As0 ) -615.29 -615.515 2 2 4 2 Ca H (As0 ) -1347.24 -1347.54 I 5 2 4 4 c a (As0 ) - 731.47 -732.16 - 732. 1 - 732 . 27 - 736.55 3 4 2 ca (As0 ) .Ca(OH) - 948 . 42 3 4 2 2 42 atmospheric carbon dioxide and showing stability regions of calcium arsenate and arsenite are shown in figures 3.7 and 3.8 respectively. Free energies of formation for the various species are listed in Appendix 1. From these results the solubility of calcium arsenate i s considerably higher than previously thought, particularly in relation to the treatment of gold processing waste waters and in the copper smelting industry. 3.1.3 Barium Arsenic Water System Barium arsenate has been accepted as one of the most insoluble arsenates. This originated from Chukhlansev 's work in 1956 [53], even though in 1938 another worker Guerin [72] published the results of the solubility of barium arsenate which were five orders of magnitude higher than that reported by Chukhlansev. Guerin's solubility data is to be found l.n the authoritative data source by Seidell and Linke [73] . Two barium arsenates have been reported by Robins [7 4] these are BaHAs04 and Ba3 (As04 ) 2 which exist in the pH ranges 5-9.3 and 9.3-12 respectively. The various free energy of formation data for Ba3 (As04) 2 throughout the literature differs greatly. Some of this data is reported in table 3.3. Robins [ 74 ] showed that if atmospheric carbon dioxide is taken into account at higher pH values the barium arsenate decomposes to barium carbonate. 0 T:298.15 1 H3RsCJ4 (RQJ 2 H2Rs Cl4- 3 HRs eJ4-- 1 4 Rs~4--- 12 5 Co.++ 6 Co.CJH+ 2 13 7 Ce. (CJHJ 2 (RQJ 8Ce.(CJHJ2(SJ \ 9Ce.H4(Rs(j4J2(SJ \ 10Ce.HRs(j4(SJ 3 \ 11 Ce.SH2(Rs(j4J4(SJ 5 \ 12Ce.3(Rs(j4J2(SJ \ 13 Co.3 (Rs(j4) 2. Co. ((jHJ 2 ::4- \ pHZCJ:0/0/0 cS u \ pH:0/0/.25 0... \ 6 p Rs ( V) : 0/. 25/0 ---5 \ . . pCo.(IJl:0/.25/0 > ...... -·------.... Cl) ~6 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH Figure 3. 7 Log activity pH diagram for the calcium(ii) arsenic(v) water system. (The effect of atmospheric carbon dioxide is to cause calcium carbonate to be stable to the upper right of the dashed line.) -----CaC03 (s) 0 T :· 298. 15 1 H3Rs(j3 (RQJ 2 H2Rs (j3- 10 3 HRs (j3-- 1 4 Rs (j3--- ~ 5 Rs (j+ \ 6 Co.++ \ 1 1 2 \ 7 Co.(jH+ \ 8 Co. ((jHJ 2 £RQJ \ 9Ce.£llHJ2(SJ .-.3 \ 10 Ce. £Rsll2l 2 (SJ \ 11 Ce.lRsll2l2 . Ce.£(jHJ2 - 6 \. pH2(j:0/0/0 \ .___.._ pH: 0/0/. 25 ~4 \ pAs (!Ill :0/ . 25/0 a.. \ p Co. l I I J : 0 I . 2 5 / ,0 .-. \ --s \ 7 - '-..--- .- - ·- _..: - ··- · Cl) ~6 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH Figure 3.8. Log activity pH diagram for the calcium(ii) arsenic(iil) water system. ( The effect of atmospheric carbon dioxide is to cause calcium carbonate to be stable to the upper right of the dashed line.) ----CaC03(S) 45 TABLE 3.3 FREE ENERGIES OF Ba3 (As04 )z FROM VARIOUS SOURCES. Source I::.Gt I 298.15K 1 Ba3 (As04 ) z I kcal/mol Chukhlansev [53] - 780.4 Guerin [7 2] -736 Gorochova [69] Method 1 -739.44 Method 2 -768.4 Robins [7 4] -736.62 46 The stability diagram in figure 3.9 represents the areas where the two barium arsenates are stable. When atmospheric carbon dioxide is taken into account barium carbonate becomes more stable. This area is represented in figure 3.9 by the dotted line which intersects the solubility curve of BaHAs04 at pH 8.3. This point represents the pH above which barium arsenate is unstable 1n the presence of atmospheric carbon dioxide. Free energy of formation data used to produce the diagram is listed in Appendix 1. These results are contradictory to the patent presented by the Key Lake Mining Coorporation and Sherritt Gordon Mines Ltd mentioned in section 1.6 . 3.1.4 Ferrous Arsenate Water System Ferrous arsenate was first mentioned in the British Pharmacopedia in 1864 [75]. There is very little information in the literature concerning ferrous arsenate. The removal of arsenic from solution with ferrous iron is usually followed by aeration or oxidation to oxidise ferrous to ferric and hence precipitate FeAS04 .2H20 . There are in fact two ferrous arsenates FeHAS04 and Fea (As04 ) 2 . Goroch ova [69] states that the free energy of formation for Fea (AS04 ) 2 is -422.21 Kcal/mol at 298K. If this data is used to produce a stability diagram the minimum solubility of Fea (As04 )2 0 T:298.15 1 H3Rs Cl4 ( RQ J I 2 H2Rs Cl 4- I 3 HRsCl4-- 1 I 8 I 4 As Cl4·--- I 5 Bo. ++ I 9 6 Bo.ClH+ 2 I 7 Bo. (ClHJ 2. 8H2Cl lS J I 8 Bo.HRso Cl4 ( S J I 9 Ba3 (RsCl 4 l 2 (SJ pHZCl:0/0/0 3 \ \ pH:0/0/ .25 ...... \ pAs (VJ: 0 /. 25/0 5 \ pBo.(IIJ : 0/ . 25/0 ::4 \ ~ (C \ 0... \ . 6 --5 \.._ ------...... > (I) ~6 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH Figure 3.9. Log activity pH diagram for the bariUm(ii) arsenic(v) water systern. (The effect of atmospheric carbon dioxide is to cause barium carbonate to be stable to the upper right of the dashed line.) - - -aaC03(s) 48 lies in the pH range of 6-7. This stability diagram showing the ferrous arsenate water system is shown in figure 3.10. Free energies of formation for the species are listed in Appendix 1. 3.1.5 Ferric Arsenate Water System It is generally agreed that only one ferric arsenate compound (FeAs04 .xH20} exists [76] although there are still references to compounds such as 6FeAS04 . 2Fe (OH} 3 [77], Fe3 AsO? and Fe4 As2 04 (in hydrated form) [78, 79] and basic ferric arsenates [80]. Numerous f erric arsenates have been identified in the literature by X-ray diffraction, a list of these is shown in table 3.4. Upon investigation of the X-ray diffraction patterns of these' compounds it was found that some of the patterns were identical, these are listed 1n table 3 . 5. Nishimura and Tozawa [60] have shown that the solubility of ferric arsenate at high pH values deviates from the linear relationship of log concentration versus pH at about a pH of 2 . Figure 3.11 shows Nishimura and Tozawa's [60] measurements in determining the solubility of ferric arsenate at 25°C . Robins [92] claimed that this was probably due to the instability of ferric arsenate and its decomposition to ferric hydroxide. 0 T:298.15 1 H3Rsll4 (RQJ 2 H2Rs Ll4- 3 HRs Ll4-- 1 4 Rs Ll4--- 5 Fa++ 6FollH+ 2 7 Fe (LJHJ2(RQJ 8 Fe (LJHJ 2 (SJ 9fe(LJHJ3- 8 10fo(CJHJ4-- 3 11 FeHRsll4(SJ ...... 12 Fe3 (Rsll4J 2 (5) pH2ll:O/O/O ::4 pH:0/0/.25 (J) 5 l1.. pRs rVJ :o/.25/0 a.. pfe(lll:0/.25/0 --5 > Cl) ~6 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH Figure 3. 70. Log activity pH diagram for the iron(ii) arsenic(v) water system. 50 TABLE 3 .4 LIST OF X-RAY DIFFRACTION PATTERNS FOR FERRIC ARSENATES. Chemical Formula Name Source Fe4 As2 Ot 1 Angellite [81] Fe3 AS07 Iron Arsenate [ 82] FeAS04 Iron Arsenate [ 83] FeHAs04 Iron Hydrogen [ 83] (Fet-xH3xAS04) Arsenate Fe2 As4 Os Karibibite [ 8 4] (Fe2 03 . 2AS2 03) Fe2 03 . 2FeAS04 Iron Oxide-2 [ 85] Iron Arsenate Fe As04 .3 . 5H20 Kankite [ 86] FeAS04 .2H20 Scorodite [87] FeAs30s .4H20 Iron Arsenate [88] (Fe203 .3As205 . 8H2 0) Hydrate FeAS04 .2H20 Scorodite [89] Fe(H2AS04 )3H20 Iron Hydrogen [90] Arsenate Hydrate Fe(H2 AS04 ) 2 .5H20 Iron Hydrogen [ 91] Arsenate Hydrate FeAS3 Os . 8H2 0 Iron Arsenate [83] (Fe203 .3As205 .16H20) Hydrate FeAs04 .xH20 Iron(iii) Arsenate [ 85] Hydrate 51 TABLE 3.5 COMMON X-RAY DIFFRACTION PATTERNS. Pattern Formula Source 1. FeAS04 .2Hz0 (Scorodite) [87] FeAS04 .2Hz0 (Scorodite) [ 89] FeAs04 .xHzO (Iron(iii) Arsenate Hydrate) [ 85] 2 Fe (Hz As04) 3 .Hz 0 [90] FeAS3 Og . 4Hz 0 [ 8 8] 3 Fe (Hz As04 ) z . 5Hz 0 [91] FeAS3 Og . 8Hz 0 [88] 4 Fez 03 . 2FeAs04 [ 8 5] Fe4 Asz Ott [ 81] 52 05 Ql QO -...... 0.01 "' -U) QOOI 0000 O.OOOIOI..--..J.---~2--3...... _--'4---"----6" pH Figure 3. 11. The variation of arsenic(v) and iron(iii) concentrations with pH in equilibrium with FeAs04 at 25°C. Nishimura ------Chukh/ansev 53 Work by Dove and Rimstidt [93] reported a solubility product for scorodite (FeAs04 .2H20) which is approximately 1.5 log units lower than that reported by Chukhlansev [54). This work has prompted the comments of Nordstrom and Parks [94) and Robins [95) who dispute the accuracy of Dove and Rimstidt results since the measurements of solubility were taken in the incongruent region, that is in the pH range 5.5-6.3,_ where ferric hydroxide is stable and ferric arsenate unstable. Table 3.6 shows the free energy of formation of ferric arsenate from various literature sources. Tozawa, Utmetsu, and Nishimura [ 96 J have produced a log activity pH diagram for the ferric arsenic{v) water system this is shown in figure 3.12. 3.2 Other Insoluble Arsenic Compounds 3.2.1 Arsenic Sulphur Water System It was assumed that the compound (orpiment) has an extremely low solubility and would be an ideal compound to be precipitated to decrease arsenic levels in solutions. The thermodynamic properties of reported in the literature are incomplete and vary widely. In the most commonly used NBS Table s [45 ) it is listed as having a free energy of formation of -40.3 54 TABLE 3.6 FREE ENERGY OF FORMATION DATA FOR FERRIC ARSENATE FROM VARIOUS LITERATURE SOURCES. Source flGt , 298. 15K, Kcal/mol Gorochova [69] -184 .9 Dove and Rimstidt [93] -301.99 Sillen [55] -183.71 Nishimura [ 60] -185 .57 55 I 'I 1 I' I I I I I I 1. I l I -0 I -0 I ~ I Fe(OH) I 3 3 I ' I Q 0 · ~ :o - ~~ 0 I - Figure 3. 72. Log activity pH diagram for the iron(iii) arsenic(v) water system. 56 Kcal/mol whereas other work lists a value o f -23.0 and -21.68 Kcal/ moll these are listed in table 3.7 . There has also been a diverse variety of thio-complexes of arsenic(iii) in solution reported. Complexes such as AsS2- and AsS(OH)2- have been reported by Babko and Lisetskaya [99]. At 90°C Mironova and Zotov 2 [100] have reported H2 As2 S4 1 As2 S4 - and HAs2S4- in solution. Also complexes of polynuclear and mononuclear based on spectrophotometric and pH measurements have been reported by Vorobe'va and co-workers [101]. Angeli and Souchay [102] have suggested the formation of the polynuclear ion AS3 S6 3- . The work of Angeli and Souchay [102] has provided data for the sulphide complex As3S6 3- which together with data for the complex AsS2- and the hydroxy complexes of arsenic(iii) 1 has been used by Tozawa and Nishimura [103] to construct a potential pH diagram for the arsenic sulphur water system. This is shown in figure 3.13. From this data the solubility of orpime nt (As2S3) is shown to be extremely low and it is this low solubility that has generally been accepted, and has an important bearing on normal practice for the removal of arsenic from solution using sulphide precipitation. 57 TABLE 3.7 VARIOUS FREE ENERGY OF FORMATION DATA FOR ORPIMENT ( As2 S3 ) • Source I:::,Gr , 298.15K, kcal/mol NBS Tables [ 45] -40.3 i Barton [97] -23.0 Johnson [98] -21.68 58 1.0 .... 10 - 0.2 > HAsoi- . tiJ '-...a 0 '-.. AszSJ AsJs~- 11 HAs02 1 ~ .... I I -0.2 37 I I ~ I I A so~- I ' 5 1 25 127 ', I 24 I " I -0.4 ~ I 28 8 A!=- AsJs~- 26 1 '. ~02 29"'- -0.6 1-bS I 30 .... I As I 31 -0.8 I Hs- 19 12 14 0 2 4 6 8 10 pH Figure 3. 13. Potential pH diagram for the arsenic sulphur water system. 59 3.2.2 Iron Arsenic Sulphur Water System Arsenopyrite (FeAsS) is the most common arsenic bearing mineral and is mainly associated with copper and gold ores. Various thermochemical studies have been reported by various authors [97,104,105] but these have been in the high temperature region. It has been suggested that arsenopyrite(FeAsS) and loellingite (FeAs2) may also be appropriate stable compounds to precipitate arsenic for disposal as a residue. Arsenopyrite and loellingite are extremely stable compounds and if they can be synthesised under room temperature conditions, then perhaps they can be used as a more favourable means for the long term disposal of arsenic. In such a disposal situation oxidising conditions would have to be limited. A number of potential pH diagrams for this system have been computed. Firstly Osseo- Asare [106] calculated the potential pH diagram for the iron arsenic sulphur water system, which is shown in figure 3.14. From this diagram it would appear that arsenopyrite is stable under reducing and neutral pH conditions. However, Osseo-Asare has neglected to include the stable arsenic sulphides, realgar and orpiment and the stable iron arsenide, loellingite which all should appear in this diagram. Bur'Yanova [107] has also obtained a 60 2.0 ,----,----y----.-----,r-----, 1.0 ------~ ---FeOOH------Eh o.o FeASS------1.0 Fe(OH'& Fe 2- -4 Fe(OH)4 [Fe]= 10 M - -4 [As] =[ S] =10 M -2.0L___ _l_ ___ _J._ __ _IL__------J 0.0 4.0 8.0 12.0 16.0 pH Figure 3. 14. Potential pH diagram for the iron arsenic sulphur water system. 61 potential pH diagram which is shown in figure 3.15. It can be seen that arsenopyrite is stable under alkaline and reducing conditions. Tahija and Haung [108] produced another diagram for the same system where similar stability regions for the arsenic sulphides and iron arsenic sulphides are shown is presented in figure 3.16. If arsenopyrite or loellingite can be synthesised under room temperature conditions they would appear to be ideal compounds for stabilising arsenic residues. Although their stability may be influenced by oxidising conditions, it is likely that the disposal of such a material could be carried out using earth cover to produce an anaerobic environment. 62 I EhJJ F~J~' 081-----1. z.+ Fe 0.11 r~ts / ,~ F~ ~s#A_, S ~ .1 s.#r~s~. -O.I.t r.• Fe •AsH, -- - - F~•AsH3 -0./3 -'--· 0 l . l..• Figure 3. 15. Potential pH diagram for tho iron arsenic sulphur water systern at 25°C at concentrations for iron, 6 4 arsenic, and sulphur at 10 - ~ 10- , and 10- M respectively. 63 .5 Eh 0 -.5 - I AsH3 []=I - 1.5 -2 0 2 4 6 8 10 12 14 16 pH Figure 3. 16. Potential pH diagram for the iron arsenic sulphur water system at a concentration of unity for all species. 64 CHAPTER 4. EXPERIMENTAL 4.1 Introduction Chukhlansev (53,54] and Nishimura and Tozawa (60], determined the solubility of various solids by mixing the solid and solution in equilibrium and then measuring the concentrations of the various species in ~olution. In this work a series of pH titrations were carried out in order to determine the point of precipitation of the solid in each system under study and to elucidate a ny speciation of aqueous species in the systems. Spectrophotometric a n alyses were also made for the purposes of speciation of the different species in solution and turbidity measurements were made in order to determine the pH at which precipitation or dissolution of the solid had commenced. The arsenic(iii) sulphur(-ii) water system was studied by pH titrations and spectrophotometry to determine the solubility and stability region of arsenic sulphide (AszS3). The iron(ii) arsenic(v) water system was studied in perchlorate and sulphate mediums by pH titrations to determine the solubility and stability regions of the two f errous arsenates. The iron(iii) arsenic(v) water system was studied in perchlorate, chl oride, nitrate and sulphate 65 mediums using pH titrations, spectrophotometric measurements and turbidity measurements to determine the solubility and stability region of ferric arsenate. 4.2 Instrumental Techniques 4.2.1 PH Titrations A series of pH titrations using either aci d or base were carried out in order to determine the pH at which precipitation occurred and any speciation of the aqueous species in the systems. The method employed 1n this study is similar to that described by Angeli and Souch ay [102] . Titration curves were recorded on a Houston recorder Omniscribe B-5000. The pH was varied by the addition of either base or acid, added at a rate of 5 ml/hr unless otherwise specified. All experiments were carried out at a temperature of 25°C for 5-10 hours. For the arsenic(iii) sulphur(-ii) work the sol utions were continuously purged with high purity nitrogen to prevent interferences from carbon dioxide. A series of working solutions of KH2AS03 and KHS were mixed so that a constant ratio of 2:3 existed between the arsenic(iii) and sulphur(-ii) respectively, since the solid aiming to precipitate is 2:3. The pH titration curves for these solutions at different concentrations have been recorded. For the iron(ii) arsenic(v) work the 66 solutions were continuously purged with high purity nitrogen to prevent interferences from carbon dioxide. A series of pH titration curves were recorded, the ratio of iron(ii) to arsenic(v) was maintained constant at 1:1. A number of experiments were performed where the ratio was altered and these are specified. This system was studied in both perchlorate and sulphate mediums. For the iron(iii} arsenic(v} work a series of working solutions were mixed for two of the mediums studied perchlorate and sulphate, so that a constant ratio existed between iron(iii) and arsenic(v} of 1:1. A series of pH titration curves for these solutions at different concentrations have been recorded. 4.2.2 Spectrophotometric Measurements Spectrophotometric analyses were made on a Perkin Elmer Lambda 3 UV-VIS Spectrophotometer. The wavelength was scanned from 750nm to 190nm. For the arsenic(iii) and sulphur(-ii) work , samples of the solution during the titrations were withdrawn periodically and scanned. For the iron(iii} arsenic(v} work, solutions were prepared and scanned for the wavelength range specified. Iron(iii} solutions were used as the reference solution, because the iron(iii) sulphate complexes have an absorption band which is in the same 67 region as the iron(iii) hydroxy ions and so by placing an iron(iii) solution in the reference cell the iron(iii) interferences will be cancelled out and only the sulphate absorption bands will appear. 4.2.3 Turbidity Measurements The turbidity was measured with a HACH Tubidimeter, Model 2100. Turbidity measurements were made 1n the iron(iii) arsenic(v) water system, this system was studied in perchlorate, chloride, nitrate and sulphate mediums. The iron(iii) and arsenic(v} solutions were allowed to reach equilibrium for 24 hours under continuous stirring and the turbidity measured daily. Base or acid was added to these solutions in aliquots of approximately lml and the turbidity was recorded. 4.2.4 PH Measurements. Solution pH was measured on an EIL 7050 pH meter using a combined EIL pH electrode. The electrodes were calibrated using EIL buffer solutions of pH 7.0 and 4.0 at 25°C. In solutions that contained perchlorate ion it was found necessary to use an EIL all purpose pH electrode in conjunction with a calomel reference electrode filled with 3.8M sodium chloride. The readings 68 were found to be erratic when using a r egular potassium chloride calomel reference electrode, due to the precipitation of potassium perchlorate at the electrode interface which in turn interfered with the electrolytic flow causing unsteady pH readings. The behaviour of the sodium chloride filled electrode was different to that of the normal potassium chloride filled electrode, du~ to the imbalance of ionic mobilities between sodium ions and chloride ions. This in turn increases the liquid junction potential and hence decreases the pH readings. In order to standardise pH using this electrode several readings of hydrochloric acid solutions of different concentrations were used as buffer solutions and the pH of the solution was obtained. Figure 4.1 shows the pH correction required with the measured pH. 4.3 Preparation of Solutions 4.3.1 Arsenic Solutions The stock solution of 1.5M solution was prepared by dissolving arsenic triox ide (As203) with the required amount of carbonate free potassium hydroxide solution. Stock solutions of 1.0M arsenic(v) were prepared by dissolving arsenic pentoxide (As20~). This 1 0 c: Q -0 cp ...... Q u :c ~~ ~------~------~0 1 2 3 4 5 6 Hee.surecl pH Figure 4. 1. PH correction versus measured pH graph. 70 solution was stirred and heated for 24 hours and then filtered. The solution was standardised using sodium thiosulphate as outlined in Appendix 2. 4.3.2 Sulphide Solutions A stock solution of 1.5M KHS was obtained ~y passing hydrogen sulphide through a carbonate free potassium hydroxide solution. 4.3.3 Iron Solutions. Stock solutions of l.OM ferric ion were pre pared by dissolving ferric perchlorate, ferric chloride, ferric nitrate, or ferric sulphate for the particular mediums studied. The stock solution was then standardised using EDTA as outlined in Appendix 3. Sample solutions were prepared by diluting the stock solution of l.OM to the required concentration. Stock solutions of l.OM ferrous ion were prepared by dissolving an excess amount of metallic iron powder in either sulphuric acid or perchloric acid. This solution was a llowed to stand overnight, continuously purged with high purity nitrogen. The solution was then filtered and the ferrous ion was standardised using eerie sulphate as outlined in Appendix 4. It was found that after two days the solution had discoloured and the ferrous ion was partly oxidised to ferri c ion. 71 4.3.3 . 1 Ferric Solutions at an Ionic Strength of 1.0 To determine the effect of ionic strength on the solubility of ferric arsenate 1M sodium nitrate was added to ferric nitrate and arsenic acid solutions mixed at a ratio of 1:1. To obtain an ionic strength equal to 1.0 J.n accordance with the equation: I = 112 I c Z1 2 where I = ionic strength c = concentration and Zt = is the charge [109] . The appropriate amount of sodium nitrate was added to solutions of ferric ion concentrations of ·0.1M, 0 .01M, 0.0075M, 0 . 0025M, and 0.001M. 4.3.4 Carbonate Free Hydroxide Solutions Extremely tigh concentrations of 17M of potassium and sodium hydroxide solutions were prepared. At this concentration the potassium or sodium carbonate precipitates, due to the low solubility of the hydroxide ion thus leaving a pure hydroxide solution . The solution was filtered to separate the carbonate precipitate and then purged with high purity nitrogen and sealed. This stock solution was then diluted to the required concentration. 72 CHAPTER 5. RESULTS AND DISCUSSION 5.1 Arsenic(iii) Sulphur(-ii) Water System 5.1.1 PH Titrations A series of hydrochloric acid titrations were performed, where the ratio of arsenic(iii) to sulphur(-ii) was kept constant at a ratio of 2:3 even though the concentration with respect to arsenic(iii) was altered. Titration curves were obtained for 0.3M, 0.15M, 0.015M, 0.0015M, and 0.00015M concentrations of arsenic(iii). With each different concentration of arsenic(iii) the concentration of hydrochloric acid used for the titrations was always ten times that of the arsenic(iii). Titration curves for millilitre of hydrochloric acid added versus pH for each of the particular solutions are shown in figure 5.1. 5.1.2 Spectrophotometric Measurements Aliquots of the solutions titrated were withdrawn periodically and the spectra recorded. The absorbance was measured at 400nm and plotted against pH. The absorbance versus pH for these solutions are shown in figure 5.2. 40 As(lll) concentration 1•0.3M 11•0.15M JII•0.015M IV•0.0015M V•0.0001.5M 30 20 '0 ,Q) '0 '0= . ~ 10 . -0 -e 2 4 e 8 10 12 14 pH Figure 5.1. PH titration curves for the arsenic(iii) su/phur(-ii) water system. 5 As(lll) concentration s· 1•0.3M 4 2•0.15M 3•0.015.M 4•0.0015M 4 5•0.00015M W3 u 3 z a: cc 0::: o · (1")2 cc a: 1 2 0 1 0 1 2 8 9 10 11 12 13 1-4 Figure 5.2. Absorbance versus pH curves at 400nm for the arsenic(lii) su/phur(-11) water system. 75 5.1.3 Discussion The interpretation of figure 5.1 was made in conjunction with the spectrophotometric results shown in figure 5.2. A computer mass balance modelling technique was used to convert the titration curve data into a series of stability and distribution diagrams figures 5.3-5.7. Upon mixing the two solutions of arsenic(iii) and sulphur(-ii) with a ratio of 2:3 respectively, the overall reactions are: 8H2.AS03- + 12H5- = 6As52- + 2H3AS03 + 4H20 + 140H and then 6As52- + 2H3 As03 = 4As2 53 ( s) + 60H- . However when the arsenic(iii) concentration is less than 0.01M the As52- forms the polynuclear ion HAs254- prior to precipitating As253. Thus 8H2AS03- + 12H5- = 3HAs254- + 2H3AS03 + H20 + 170H and then 3HAS2 54- + 2H3 AS03 = 4AS2 53 ( s) + 3H2 0 + 30H finally the overall reaction is 8H2AS03- + 12H5- = 4As253 (s) + 4H20 + 200H- As an example of the cqnversion of the titration data into the stability diagrams, take plot III in figure 5.1. This curve represents the titration of a 0 T: 298.15 . 1 H3RslJ3 (RQJ . 2 H2Rs LJ3- . 3 HRs LJ3-- . 1 .. 4 Rs Ll 3 ·- -- ·· · 5 Rs (J+ . .... 6 RsS2- 2 .. 7 HRs 254- .. 8 H2Rs2S4 (RQJ .. 9 Rs2S3 (SJ 3 ·· 10Rs2LJ3(S J :3 I1 ~ ~ =( R Q J - •' ·· 13 - •' s-- ~4 z pH2LJ:O / O/ O a.. 8 pH:0 / 0/ . 25 : .. pRs (JJIJ :o/.25/0 --s .· .. pS(JIJ : ·. 176/ . 25/0 ~6 ...... 1 .· 0. . 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH Figure 5.3. Log activity pH diagram fo r the arsenic(iii) sulphur( -ii) water system. 4 1 H3Rsl!3 ( RQJ 2 H2Rs l!3- HRsl!3-- 4 Rs (j 3--- 5 Rs ll+ 6 Rs S2- 7 HRs 2S4- 8 H2R s 2S'4 ( RQ l 9H2S(RQJ 10 HS- 11 s- p H2LJ:o;o pH: 0/ . 1 pRs ( l1 I l : 1/0 pS( I IJ: . 824/ 0 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 pH Figure 5.4. Ionic distribution curves for the arsenic(iii) sulphur(-//) water system at a concentration of 0.1M with respect to arsenlc(iii). T:298.15- I " 100 4 1 H3Rs(j3 (RQJ 2 H2Rs (j3- 90 3 HRs CJ3-- 4 As (j3--- 80 5 Rs (j+. 8 6 Rs 52- 70 7 HRs 254- 8H2Rs2S4(RQJ 60 9H2S(RQJ 10 HS- 50 11 s ·-- pH2(j:O/O 40 pH:0/.1 p Rs (I I I J.; 2/0 30 pS(IIJ; 1.824/0 ~20 ..._, ::1 0 5 ">Q a: :---.: 0 1 2 3 4 5 6 7 8 9 10 1 1 12 13 14 15 16 pH Figure 5.5. Ionic distribution curves for the arsenic(iii) sulphur(-//) water system at. a concentration of 0.01M with respect to arsenic(iii). ,> T:298.15 100 4 1 H3Rs03(RQJ 2 H2Rs 03- 90 3 HRs 03-- 4 As 03--- 80 5 Rs 0+ 8 6 Rs 52- 7 HR s 2S4- 70 8H2Rs2S4(RQJ 60 9H2S(RQJ 1 0 HS·- 1 1 s ·-·- 50 pH20:0/0 pH:0/.1 40 pRs(IIIJ:3/0 pS(!!):2.824/0 30 1 ,.....20 ...... ::10 5 (1) 0 a: :---.: 0 1 2 3 4 s 6 7 8 9 10 11 12 13 14 15 16 pH Figure 5.6. Ionic distribution curves for the arsenic(iii) sulphur(-ii) water system at a concentration of 0.007M with respect to arsenic(iii). T:298.15 100 4 lH3AslJ3(AQJ 2 2 H2As·{j3- 90 3 HAs lJ3-- 4 As lJ3--- 80 5AslJ+. 8 6 As 52·- 70 7 HAs 25 4·- 8H2As2S4CAQJ 60 9H2S(AQJ 10 HS·- so 11 S- pH2lJ:O/O 40 pH:0/.1 pAs (IIIJ :4/0 30 pS(IIJ :3.824/0 ,.... 2 0 .._ ::1 0 ,5 rno CI ~ 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 pH Figure 5. 7. Ionic distribution curves for the arsenic(iil) sulphur(-//) water system at a concentration fo O.OOOJM with respect to arsenlc(lil}. 81 solution of [As] = 0.015M, [S]=0.0225M with 0.1M hydrochloric acid. Initially the solution is at pH 12 where the ion AsS2- is predominant as is shown in figure 5 . 3 . At pH 9 . 3 in figure 5.1 at the point marked A there is an indication of excess H2AS03- to- H3AS03 as can be seen in figure 5.5 which corresponds approximately to those concentrations. At pH 7.5 at the point marked B 1n figure 5.1, the next equivalence point is for AsS2- to at which point there is little precipitation of The precipitation of As2S3 however has commenced at pH 7 . 4 and continues till the completion of AsS2- to HAsS2- and then onto point C at pH 6 . At pH 5 . 25 at At concentrations less than 0 . 00015M there exists a direct change from AsS2- to HAs2s4- without a predominant polynuclear intermediate. Thus: and then and finally the overall reaction is In figure 5.1 curve V represents the titration for the solution of [As] = 0.00015M, [S] = 0.000225M with 0.001M hydrochloric acid . The "overshoot" which is shown at pH 5 indicates the direct change AsS2- to HAsS2 without a predominant polynuclear intermediate. 82 The stoichiometry of the various reactions which are proposed from the trends of the t itration curves are consistent (almost precisely) with the quantity of acid added . The proposition that the species HAs2S4- is the only polynuclear ion is deduced from stoichiometry and the intensity of the absorption spectra at 400nm. The absorption spectra in figure 5.2 indicate an absorptivity which is about ten times that for AsS2- in the pH range 7-13 which is consistent with the formation of a binuclear complex. It is probable that arsenic(iii) and chloride ion complexes occurred since hydrochloric acid was used in the titrations . However , only three main arsenic(iii) chloride complexes have been listed in the literature, [110-113] As(OH)Cl2, As(OH)2Cl , and AsCla and the free energies of formation for these compounds are available from the NBS Tables [45], so that they can be incorporated into the mass balance model. When this is done it was shown that they do not effect the proposed arsenic sulphide complex species. These results lead to the consistent set of free energies which are shown in table 5 . 1. The solubility of As2Sa can be determined directly from figure 5 . 3 . The minimum solubility lie s in the pH range 0-5 where [As] = 28 mg / 1. This value is considerably higher than is commonly accepted and has an important 83 TABLE 5.1 FREE ENERGY OF FORMATION DATA FOR THE ARSENIC(III) SULPHUR(-II) WATER SYSTEM. Species Free Energy Data Source (llGt I 298.15K Kcal/Mol) . H3AS03 (aq) -152.92 [45] H2 AS03- -140.33 [ 45] HAS03 2- -121.71 [46] As03 3 - -102.95 [46] Aso• -39.15 [ 45] As2 03 ( s) -137.72 [ 45] *AsS2- -5. 6 *HAs2 S4- -24.3 *H2AS2 S4 (aq) -31 . 5 *As2 S3 ( s) -21.68 - *Determined in this work . 84 bearing on normal practise for the removal of arsenic from process solutions using sulphide precipitation methods. 5.2 Iron(ii) Arsenic(v) Water System / 5.2.1 PH Titrations A series of titrations were performed to investigate the iron(ii) arsenic(v) system in perchlorate and sulphate mediums . For each iron(ii) and arsenic(v) concentration used, the sodium hydroxide concentration was ten times that of the iron(ii) concentration. 5.2.1.1. Perchlorate Medium A series of sodium hydroxide titration curves were obtained for 0.2M, 0.1M, 0.01M, 0.001M and 0.0001M concentrations of iron(ii) and arsenic(v). Even though the concentrations of iron(ii) and arsenic(v) were changed the ratio of iron(ii) to arsenic(v) was maintained at 1:1. These titration curves are shown in figure 5.8. Additional experiments were performed where the ratio of iron(ii) to arsenic(v) was changed to 3:2 and 4:3. These titration curves were also recorded, but did not produce anything more conclusive than solutions where the ratio of iron(ii) t9 arsenic(v) was 1:1. 30 20 Fe(ll) oonoentratlon 1•0.2M 11•0.1M 111•0.01M IV•0.001M V•0.0001M "C 10 G) "C "C ~ J: 0 z~ -0 E 0 0 2 4 6 8 10 12 14 pH Figure 5.8. PH titration curves for the iron(ii) arsenic(v) During these titrations which were run for approximately seven hours a colouration of deep blue green was observed. Initially it was believed that even though high purity nitrogen was continuously purged through the solution, oxidation of iron(ii) to iron(iii) might still be occurring. As the experiment had been run for such a lengthy time an identical experiment was run with a freshly prepared iron(ii) solution and the sodium hydroxide added at five times the initial rate. Once again the same intensity of colouration occurred thus eliminating the possibility that oxidation of Fe2 + to Fe3 + was the cause of the colour. Instead it was possibly due to the different complexes in the solution in the high pH region of 10-12 . Table 5.2 represents a summary of the various concentrations used, the altered ratios and the actual pH at the point of precipitation. No point of precipitation was recorded for curve V in figure 5 . 8 concentration 0.0001M because at this concentration it was too difficult to visually observe the exact point where precipitation commenced. 5.2.1.2 Sulphate Medium A series of sodium hydroxide titration curves were obtained for 0.2M, 0.1M , 0 . 01M and 0.001M concentrations of iron(ii) and arsenic(v) . Even 87 TABLE 5 .2 ~H AT THE POINT OF PRECIPITATION UNDER THE DIFFERENT CONDITIONS SPECIFIED FOR THE IRON(II) ARSENIC(V) WATER SYSTEM IN THE PERCHLORATE MEDIUM. I' Fe(ii) As(v) Fe:As Initial pH at ' [ ] [ ] pH precipitation 0 . 2 0.13 3:2 1. 22 1.31 0.1 0.07 3:2 1. 38 1.56 0.2 0.2 1:1 1.3 1.325 0.1 0 .1 1:1 1.435 1 .55 0.01 0.01 1 :1 2.2 2.66 0.001 0.001 1:1 3.135 3.6 0.5 0.375 4:3 0 .255 0.5 0.2 0 . 15 4:3 0.76 1. 325 - 88 though the concentration of iron{ii) and arsenic{v) were changed the ratio of iron{ii) to arsenic{v) was maintained at 1:1. These titration curves are shown in figure 5.9. Table 5.3 indicates the actual pH at the point of precipitation at these various concentrations. 5.2.2 Discussion A computer mass balance technique was used to convert the titration data into stability diagrams for the iron{ii) arsenic{v) water syst~m in perchlorate and sulphate mediums . At the commencement of titration, with the two solutions of iron(ii) and arsenic{v) mixed at a ratio of 1:1 and at concentrations above 0.005M in the perchlorate medium, the initial reaction can be represented by: and then as the titration with sodium hydroxide proceeds, FeH2AS04+ = FeHAs04 {s} + H+ and at a pH of 3 -: FeHAS04 (s) + 2Fe2+ + 2H3AS04 = Fe3 (AS04)2 (s) + H2As04- + 5H+ so that the overall reaction becomes 3Fe2 + + 3H3 AS04 = Fe3 {AS04) 2 (s) + H2 AS04 - + 7H+ At concentrations less than 0 . 005M the FeH2AS04+ complex is no longer a predominant species, 30 Ft(ll) concentration 20 1•0;2M 11•0,1M 111•0.01M IV•0,001M , ,IV 10 co . J: 0 zco -0 E 0 2 4 6 8 10 12 14 pH Figure 5.9. PH titration curves for the iron(ii) arsenic(v) water system at a ratio of 1:1 for iron'(it') and arsenic(v) in the sulphate medium. 90 TABLE 5.3 PH AT THE POINT OF PRECIPITATION UNDER THE DIFFERENT CONDITIONS SPECIFIED FOR THE IRON(II) ARSENIC(V) WATER SYSTEM IN THE SULPHATE MEDIUM. .,_Fe(ii) As(v) Initial pH at [ ] [ ] pH Precipitation 0.2 0.2 1.44 1. 48 0.1 0.1 2.15 2.175 ' i : 0.01 I 0.01 2.415 2.855 I I 0.001 i 0.001 2.75 3.44 I 91 the direct precipitation of Fe3 (As04)z occurred and thus the initial reaction at these concentrations is-: 3Fe2 + + 3H3 As04 = Fe3 (As04 ) 2 ( s) + H2 As04- + 7H+ Interpretation of these titration curves consisted of firstly examining area A in figure 5.8 and in conjunction with the stability diagram in figure 5.10, it can be seen that area A is the region where FeHzAS04+ is the predominant ion. The pH at which this region ends i s the pH at which precipitation of ferrous arsenate (FeHAs04) occurs. Curve IV in figure 5.8 represents the titration of a solution of equal concentrations of iron(ii) and arsenic(v) at 0 . 001M, at this concentration is no longer the predominant ion. It can also be observed from figure 5.10 the direct precipitation of Fe3 (As04 )z exists at this concentration. Figures 5.11, 5.12 and 5.13 are distribution diagrams for concentrations of iron(ii) and arsenic(v) at 0.2M, 0.1M, and 0.01M respectively. In the pH range between 1 and 6 the predominant ion is FeHzAs04+. In figure 5.14, which is equivalent to iron(ii) and arsenic(v) concentrations of 0.001M, FeHzAs04+ is no longer the predominant ion in the same pH range. Area B in figure 5~8 represents the region where the conversion of FeHAs04 to Fe3 (AS04)z occurs. This can also be observed by the slope of the 1 H3Rs(j4 CRQl + Represents experlmentaJ points. 2 H2Rs (j4- .. 3 HRs (j4-- 2 .. 4 Rs (j4--- 12 5 Fa++ 3 6Fa(jH+ 7Fa(eJHl2 4 14 BFeCeJHl3- 9 Fe CeJHl 4-- 5 10 Fe CeJHl2CSl 11 FeHRseJ4 (51 5 12 Fa3(Rse1412(5) ::s 13 FeH2RseJ4+ - 14FeCeJ3CSl 15 C(j3-- 6 8 16 HCeJ3- 17 H2CeJ3 ::9 p H2(j: 0/0/0 ~ pH:OIOI.25 ~1 0 p As ( V l : 0 I . 2 5 I 0 ~----~--~------~~----~0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH Figure 5. 70. Log activity pH diagram for the iron(ii) arsenic(v) water system in the perchlorate medium taking into consideration atmospheric carbon dioxide. T:298.1S 100 6 l H3RsLJ4(RQJ 1 0 8 2 H2Rs LJ4- 90 3 HRs LJ4-- 4 Rs LJ4--- 80 S Fe++ 6FeClH+ 70 7 Fe (CJHJ 2 8 Fe (LJHJ 3 - 60 9Fe(LJHJ4-- 10 FeH2RsLJ4+ pH2CJ:O/O so pH:0/.1 pRs (V J;. 69897 40 pfe(!!J:.69897 30 9 20 :::1 0 J 0 1 2 3 4 5 6 7 8 9 1.0 1 1 1 2 1 3 1 4 pH Figure 5.12. Ionic distribution curves for the iron(ii) ar senic(v) water sys.tem in the perchlorate medium at a concentration of 0.1M with respect to iron(ii). T:298.15 100 6 1 H3R~(j4(RQJ 2 H2Rs (j4- 90 3 HRs (j4-- 4 As (j4--- 80 5 Fe++ 6fe(jH+ 10 70 7 Fe ( (jH J 2 8fe((jHJ3- 60 9 Fe ((jHJ 4-- 10 FeH2Rs(j4+ 50 pH2(j:Q/O pH:0/.1 pRsCVJ:2/0 40 pfe(!!l:2/0 30 9 20 ::1 0 J 1 2 · 3 4 5 6 7 8 9 10 11 12 13 14 Figure 5.13. Ionic distribution curves for the Iron(//) arsenic(v) water system In the perchlorate medium at a concentration of 0.01M with respect to Iron(//). T:298.15 100 5 6 1 H3Rs(j4 (RQl 2 H2Rs (j4- 90 3 HRs (j 4-- 4 Rs (j4--- 80 5 Fe +4- 6fe(jH+ 7Fe ((jHJ 2 (RQl 70 8Fe((jHJ3- 60 9 Fe ((jHJ 4-- 10 FcH2Rs(j 4 + pH2(j:O/O 50 pH : 0/ . 1 pRs (VJ : 3 / 0 40 10 pFe(I I l:3/ 0 30 9 20 ::10 q)o J Ll... X 0 1 2 3 4 5 6 7 B 9 10 11 12 13 14 pH Figure 5. 74. Ionic distribution curves for the iron(//) arsenic(v) water system in the perchlorate medium at a concentration of 0.001M with respect' to lron(ii). 97 precipitation points in the log activity pH diagram shown in figure 5.10. If all the precipitation points represent the same solubility line then this line should pass the pH axis at a concentration of one at the same pH. The precipitation points of 0.2M, 0.1M and 0.01M concentrations all pass through the same pH, but the precipitation point at a concentration of 0.001M passes through a different pH at a concentration of one. This can be observed from the slope of the solubility line in figure 5.10. Thus, the conversion of FeHAso. to Fea (As04)2 occurs in the concentration range of 0.01M to 0.001M and in the pH range of 2.5-3.5. Area C in figure 5.8 represents the equivalence points for H2Aso.- to HAso. 2- and for Fe2+ to FeOH+ which can be observed in curves I-IV. Area D represents the equivalence point for FeH2AS04+ to FeOH+. On curves I and II the equivalence point for FeH2AS04+ to FeQH+ is well defined, whereas on curves III and IV the equivalence points for areas C and D become less defined. This is due to the fact that the difference in pH between these equivalence points is minimal at these concentrations. This can also be observed from the distribution diagrams in figures 5.11-5.14. By altering the iron(ii) to arsenic(v) ratio it can be observed from table 5.2 that the pH at which precipitation occurred did not alter. Therefore by obtaining log activity versus pH diagrams at these ratios 98 no changes to the solubility curve could be observed. This is shown in figures 5.15 and 5.16 where the ratios of iron(ii) to arsenic(v) are 3:2 and 4:3 respectively. In the sulphate medium the same interpretation and conversion of data was used. From the pH at the point of precipitation obtained from the titration curves in the sulphate medium and the free energies calculated for the solid species determined in the perchlorate medium, a log activity versus pH diagram was obtained for the iron(ii) arsenic(v) system in the sulphate medium, which is shown is figure 5.17. The complex postulated in the perchlorate medium FeHzAso.+ still exists, but the neutral species Feso. (aq) is more predominant in this medium. Area A in figure 5.9 represents the region where the conversion of FeHAso. to Fe3 (AsO.)z occurs as in area B shown in figure 5.8 representing the perchlorate medium. Area B in figure 5.9 represents the equivalence area C represents the equivalence point for FeHzAso.+ to FeOH+ and area D the equivalence point for FeS04 (aq) to FeOH+ . On curves III and IV in figure 5 . 9 an "overshoot" exists in the pH range 5.5-6.5 which indicates the equivalence points for HzAso.- to HAso.z-, Fe2 + to FeOH+, FeHzAso.+ to FeOH+ , and FeS04 (aq) to FeOH+ which all occur in this pH range. These equivalence points can also be observed 0 ~~~~~------~r-----11 H3As~4CAQl 1 + Represents ex perlmental points. 2 H2As Cl4- 3 HAs ~4-- 2 4 As ~4--- 12 5 Fe •+ 3 6 Fe ~H+ 7 Fe C~Hl 2 CAQl 4 14 8fe(~Hl3- 9fe(~Hl4-- 5 10 Fe £~Hl 2 (Sl 11 Fe HAs~ 4 ( S l 5 1 2 Fe 3 (As~ 41 2 ( S l ::s...... 13FeH2As~4+ 14 FeC~3 (51 -;7 15 C~3-- LL. 6 16 HC~3- ~ 8 - 17 H2C~3 CRQl :::9 pH2~:0/0/0 Cl) pH:0/0/.25 pAs (Vl:. 1761/.25/0 ~10 L------2--~3~~4~~5~~6~~7~~8--~9--~1~0~1~1~1~2~1~3~140 1 pH Figure 5. 15. Log activity pH diagram for the iron(ii) arsenic(v) water system in the perchlorate medium. The ratio of iron(ii) to arsenic(v) is 3:2. 0 T: -298.15 1 H3Rs t'J4 ( RO J .. 11 1 + Represents experimental points. 2 H2As el4- .. 3 HRs eJ 4-- 2 ·. 4 Rs el4.--- 12 5 Fe++ 3 6FeCJH+ . 7fe(eJHJ2CRQJ 14 8fe(eJHJ3- 4 9FaCCJHl4-- 5 10 Fa CeJH1 2 (51 11 FaHRst'J4 (51 5 12 Fe3 (RseJ4l 2 (5) ::s 13 FeH2RseJ4+ :/- 14FeCt'J3CSJ LL. 15 Ct'J3-- a...a 6 a 16 HCCJ3- ~ 17 H2CeJ3 ( RQ 1 > -9 pH2CJ:O/O/O tl) pH:0/0/.25 pAs (VJ:. 12/ . 25/0 ~10 ------~------~0 1 2 3 4 5 6 7 B 9 10 11 12 13 14 · pH Figure 5. 16. Log activity pH diagram for the iron(ii) ...... arsenic(v) water system in the perchlorate medium. The 0 ratio of iron(ii) to arsenic(v) is 4:3. 0 Represents experimental points. 15 _4 ...... ::s CJ) LL-6 a.. . 7 ~7 (1)8 8 a: a.. oo::t-9 0 ~10 0 1 pH Figure 5. 17. Log activity pH diagram for the iron(ii) arsenic(v) water system in the sulphate medium. The ratio of iron(ii) to arsenic(v) is 1:1. 102 in the distribution diagrams for the equivalent concentrations in figures 5 . 18-5.21. These results for the iron(ii) arsenic(v) water system in both the perchlorate and sulphate mediums have led to the set of free energies listed in table 5.4. The solubility of FeHAs04 and Fe3 (As04 )2 can be determined directly from figures 5.10 and 5.17 in their respective mediums. The minimum solubility for Fe3 (As04 )2 is reached between pH 6 and 7 in both the sulphate and perchlorate mediums. By changing the ratio of iron(ii) to arsenic(v) from 1:1 to 3:2 and 4:3 it was found that this had no effect on the solubility of Fe3 (As04 )2 in the pH range 6-7. From the log activity pH dia grams shown in figures 5.10 and 5.17 it can be concluded, that by controlling the pH between 6 and 7 and under anaerobic conditions, Fe3 (As04 )2 might be a suitable compound for stabilising arsenic waste s. 5.3 Iron(iii) Arsenic(v) Water System 5.3.1 PH Titrations A s e ries of pH titrations were performed for the iron(iii) arse nic(v) water system which was studied in sulphate a nd perchlorat e mediums. The ratio of iron(iii) to arsenic(v) was maintained a t 1 : 1 for all expe rime nts in t hese two medi ums. With different T:298.15 100~~~~------.1 H3R~~4£RQJ 8 2 H2Rs ~4- 90 3 HRs ~4-- 4 Rs Cl4--- 5 fe..J-4- 80 6FeCJH+ _ 7 Fe (CJHJ 2 (RQl 70 8fa(CJHl3- 1 1 60 9Fa£CJHl4-- 10 FaH2R~CJ4+ . 11 FQSCJ4 (RQl so 12 HSCJ4- 10 13 SCJ4- 40 pH2CJ:O/O pH:0/.1 30 pAs (VJ ·:. 69897 9 pfe(!Il:.69897 20 pSCJ4:.69897/0 ::10 Q)o J u.. X 0 2 4 6 a ro . 12 14 pH Figure 5. 78. Ionic distribution· curves for the iron(! I) arsenic(v) water system in the sulphate medium at a concentration of 0.2M with r,espect to iron(//). T:298.15 100 1 H3Rs (j 4 £ R Q J · 6 8 2 H2Rs Cl4- 90 3 HAs Cl4-- 4 As CJ4---· 80 5 Fe++ 6FeCJH+ 70 7 Fe (eJHJ 2 £RQl 8fe(ClHl3- 1 1 60 9fo(ClHl4-- 10 Fe H2Rs: CJ4+ 50 11 FeSCl4 (RQJ 12 HSd4- 40 10 13 SCl4-- pH2CJ:O/O 30 pH:0/.1 pAs (Vl: 1/0 9 pfe(!Il:l/0 20 pSCJ4:1/0 ::10 J 0 2 4 6 8 10 12 14 pH Figure 5.21. Ionic distribution cur ves for the iron(ii) arsenic(v) water system in the sulphate medium at a concentration of 0.001M with respect to iron(ii). 107 TABLE 5.4 SET OF FREE ENERGIES OF FORMATION FOR THE IRON(II) ARSENIC(V) WATER SYSTEM. Species Free Energy Data Source (~Gf , - 298.15K KCal/mol). H3AS04 (aq) -183.08 [ 45] H2 As04- -180.01 [45] ' HAs04 2- -179.79 [45] AS04 3- -154.97 [45] HS04- -180.67 [45] ' S04 2 - -177.95 [45] C02 (g) -92.254 [45] ' C03 2- -126.17 [ 45] : HC03- -140.26 [45] H2 C03 ( aq) -148.92 ' [45] FeC03 (s) -159.34 [45] Fe2 + -18.86 [ 45] FeOH+ -66.3 [45] Fe (OH) 2 (aq) -104.24 [49] Fe(OH)3- -146.96 [45] Fe(OH)4 2- -183.96 [ 45] Fe (OH)2 (s) -114.7 [49] *FeHAs04 (s) -202.33 *Fe3 (AS04 )2 (s) -419.00 *FeH2AS04+ -203.00 *FeS04 ( aq) - 201.23 ' *Determine d in this work 108 concentrations of iron(iii) and arsenic(v) the concentration of sodium hydroxide was also changed so that it was always fifteen times that of the iron(iii). 5.3.1.1 Perchlorate Medium Solutions of various concentrations of iron(iii) perchlorate and arsenic(v) were prepared and mixed so that the ratio of iron(iii) to arsenic(v) was 1:1. Titration curves of 0.2M, 0.1M, 0.01M and 0.001M concentrations for iron(iii) and arsenic(v) were recorded and these curves are shown in figure 5.22 . 5.3.1.2 Sulphate Medium Upon mixing the iron(iii) sulphate and arsenic(v) solutions at concentrations of 0.2M and 0 . 1M with respect to iron(iii) and arsenic(v) no precipitation occurred. To these solutions sodium hydroxide was then added. These titration curves were recorded and are shown in figure 5 . 23 . However, at the lower concentrations of 0.01M, 0.001M and 0.0001M upon mixing the iron(iii) sulphate and arsenic(v) solutions precipitation was immediate . These solutions at the lower concentrations were also titrated with sodium hydroxide and these curves are shown in figure 5 . 23. These solutions were prepared again and to determine the dissolution point, back titrated with sulphuric acid. These titration curves 30 Fe(lll) concentration 20 1•0.2M 11•0.1M 111•0.01M JV•0.001M 10 "0 4) "0 "0 ell :I: 0 zCll 0 - IV E 0 0 2 4 6 8 10 12 14 pH Figure 5.22. PH titration curves for the iron(iii) arsenic(v) water system in the perchlorate medium. Fe(lll) oonoentratlon 1•0.:2M 11•0.1M 20 111•0.01M IV•0.001M V•0.0001M 0 - Ill IV Figure 5.23. PH titration curves for the iron(iii) arsenic(v) water system in the sulphate medium.- 111 were ·also recorded but did n o t provide any evidence of any reactions occurring in the low pH range. 5.3.2 Turbidity Measurements Turbidity measurements were made on iron(iii) aresenic(v) sol utions in perchlorate, nitrate, chloride, and sulphate mediums . Changes in the turbidity were recorded with each addition of base or acid and plotted against the measured pH of the solution. Daily measurements of the turbidity were made for all of the different mediums. The solution was allowed to reach equilibrium after each addition of base or acid. These solutions were continuously stirred and were mixed so that the iron(iii) to arsenic(v) ratio was 1 : 1 . A linear relationship exists between tu ~bidity and pH at low turbidity values and it was this linear relationship t hat allowed the extrapolation at low turbidity values to determine the pH at the point of precipitation. As the turbidity of the solution was increased with each addition of base a deviation from this linear relationship was exhibited. This deviation was related to the particle size of the solid in solution. When the par ticle size was less than the wavelength of light , it was found that a linear relationship existed between turbidity and pH. This can be e xplained by the fact that the scattering of 112 light · in the forward and backward direction is equal, thus maki~g the angular dependance negligible . This is known as Rayleigh scattering [114] . As the particle size of the solids is increased and is equal or larger than the wavelength of light, the angle at which the light is scattered has to be taken into account . This is because the incident beam hits the surface and is scattered in all different directions, since the particles are large enough to have an effect on the scattering of light. These reflected beams interfere with one another influencing the measurement of turbidity. This scattering is known as Mie scattering [114]. Therefore at larger particle sizes the turbidity is not directly proportional to the amount of light scattered and the calculations become more complex. The relationship between light scattering and particle size has been well documented [114,115] and relationships between particle size and scattering angle in more detail have also been defined [114,115]. At the initial stages of precipitation the particle size is small and a linear relationship was observed between turbidity and measured pH. It is therefore reasonable to extrapolate to pH values at the onset of precipitation. Another factor which affected these turbidity measurements was that as the particle size increased, the settling of the particles had more effect on obtaining an 113 accurate turbidity measurement. In order to determine the influence of settling, a solution of O. OlM with respect to iron(iii) and arsenic(v) in perchlorate medium was prepared. To this solution sodium hydroxide was added in small aliquots and in this experiment the turbidity was measured over · a period of 30 minutes at time intervals of 5 minutes. It was found that while the particle size was negligible the turbidity over the 30 minute period did not change. However, as the particle size increased settling of the particles began to interfere with the turbidity measurements . Within the first 5 minutes settling had a negligible effect on the turbidity but within 30 minutes the turbidity had increased twofold. Since all of the previous measurements were made within 5 minutes of pouring the suspension into the turbidity cell the influence of settling on the turbidity measurements can be neglected. Turbidity measurements taken during the precipitation of ferric arsenate from perchlorate, nitrate, chloride, and sulphate mediums were recorded. Graphs of turbidity versus measured pH were plotted and the linear sections of the plots were extrapolated to zero turbidity and the measured pH recorded. This measured pH was then corrected using the pH correction versus measured pH graph shown in figure 4.1. The corrected pH was then used as the point of onset of precipitation commencement. 11 4 5.3.2.1 Perchlorate Medium Solutions of various concentrations of iron(iii) perchlorate and arsenic(v) were prepared and mixed so that the ratio of iron(iii) to arsenic(v) was 1:1. Sodium hydroxide was then added in small aliquots to the solution and these were allowed to reach equilibrium. The turbidity and measured pH was recorded. The concentration of sodium hydroxide was always ten times more than the concentration of iron(iii) and arsenic(v) in solution. Figure 5 . 24 represents the change in turbidity and measured pH and table 5.5 represents the change in corrected pH at the point of precipitation with concentration. 5 . 3 . 2.2 Chloride Medium Solutions of various concentrations of iron(iii) chloride and arsenic(v) were prepared and mixed so that the ratio of iron(iii) to arsenic(v) was 1 : 1. Sodium hydroxide was then added in small aliquots and these were allowed to reach equilibrium. The turbidity and measured pH was recorded. The concentration of sodium hydroxide was always ten times more than the concentration of iron(iii} and arsenic(v} in solution. Fi gure 5 .25 represents the change in turbidity with measured pH . By e x trapolating back the straight line relationship to 2ero turbidity the measured 115 500 0.2M -·- 0 .1M -(!J- • 0 .01M 0.0076M -·--·- 400 0.0026M -x- 0.001M -+- 300 • I C!J 200 •I (!JI I :z: X ~ 100 I . I t- :H- . + :z: (!J >- t- ~~~ -Q -CQ a: I(!J :;:) ci li + t-0 • 1 0 1 2 4 pH Figure 5.24. Turbidity as a function of pH for the iron(iii) arsenic(v) wa ter system in the perchlorate medium: 116 TABLE 5.5 CORRECTED PH AT THE POINT OF PRECIPITATION AND THE VARIOUS CONCENTRATIONS OF IRON{III) AND ARSENIC{V) IN THE PERCHLORATE MEDIUM. Fe{iii) & As{v) Corrected pH at [ ] precipitation 0.2 0.53 0.1 0.8 0.01 1 . 39 0.0075 1. 625 0.0025 2 . 065 0.001 2.41 117 400 0 .2M -·- O.tM -[!]- O.OtM -·- 320 0.0075M _ .... _ 0 .0025M -X- 0.001M (!] -+- (!]I 240 160 ·ao ;::). 1!1 t-. :z >- t- I -0 -Q:l a:: ;::) t-0 0 l 2 3 4 pH Figure 5.25. Turbidity as a function of pH for the iron(iii} arsenic(v} water system in the chloride medium. 118 pH at· which precipitation of ferric arsenate occurred was determined. Table 5 . 6 represents the change in corrected pH at the onset of precipitation with concentration. 5.3 . 2 . 3 Nitrate Medium Solutions at various concentrations of iron(iii) nitrate and arsenic(v) were prepared and mixed so that the ratio of iron(iii) to arsenic(v) was 1 : 1 . These solutions were then allowed to reach equilibrium and the turbidity and measured pH was recorded. Small aliquots of sodium hydroxide were added after equilibrium was attained. The turbidity and measured pH was then recorded and plotted, as shown in figure 5 . 26 . The straight lines sections of these plots were extrapolated to zero turbidity and the measured pH at which the onset of precipitation of ferric arsenate occurred was recorded and corrected with the pH correction versus measured pH graph shown in figure 4.1 . Table 5 . 7 illustrates the change in corrected pH at the point of precipi tation and the various concentrations of iron(iii) and arsenic(v ) . 5 . 3 . 2 . 4 Sulphate Medium Solutions of various concentrations of iron(iii) sulphate and arsenic(v) were prepared and 119 TABLE 5.6 CORRECTED PH AT THE POINT OF PRECIPITATION AND THE VARIOUS CONCENTRATIONS OF IRON(III) AND ARSENIC(V) IN THE CHLORIDE MEDIUM. ' Fe(iii) & As(v) Corrected pH at [ J precipitation 0.2 0.52 0.1 0 . 765 0.01 1. 485 0.0075 1. 64 0 . 0025 1. 986 ' 0.001 2.175 120 •I 0 .2M -·- 0 .1M 560 -(!I- 0.01M - 4 - 0.0075M 0.0025M -·- -x- 480 0 . 001M -+- 400 320 • C!l 240 C!l 160 •I . t- 1 C!l 0 1 2 ~ 4 pH Figure 5.26. Turbidity as a function of pH for the iron(iii) arsenic(v) water system in the nitrate medium. 121 TABLE 5.7 CORRECTED PH AT THE POINT OF PRECIPITATION AND THE VARIOUS CONCENTRATIONS OF IRON(III) AND ARSENIC(V) IN THE NITRATE MEDIUM. Fe(iii) & As(v) Corrected pH at [ J precipitation 0.2 0.57 0.1 0.78 0.01 1.52 0.0075 1.62 0.0025 2.2 ' 0.001 2.355 122 mixed so that the ratio of iron(iii) to arsenic(v) was maintained at 1 :1. When the two solutions at 0.2M and 0 . 1M concentrations were mixed no precipitation was observed. These solutions were allowed to reach equilibrium and sodium hydrox ide was added in small aliquots. The turbidity and measure d pH was recorded and the plotted results are shown in figure 5.27. On mixing the solutions having concentrations of 0 . 01M, 0 . 0075M, 0.0025M, and 0.001M, with respect to iron(iii) and arsenic(v) at a ratio of 1 : 1 precipitation occurred immediately . In these experiments sulphuric acid was added to determine the dissolution point of the ferric arsenate and the decreased turbidity and measured pH recorded. Figure 5.28 repre~ents the turbidity and measured pH change with the addition of sulphuric acid. These s olutions were again prepared and the experiments were repeated using perchloric acid instead of sulphuric acid . This was decided due to the fact that perchloric acid is unlikely to have as high a tendency to complex as readily as sulphuri c acid. Figure 5.29 represents the change in turbidity with measured pH using perchloric acid. The corrected pH at which precipitation commences for the various c oncentrations considered are shown in table 5.8. 5 . 3.3 Influence of Ionic Strength Turbidity measurements were also made on 123 0 .2M -·- 560 0 .1M - (!J - 480 • 400 320 . 240 160 . :::;:). t-. ~BO )- t- ~ - 1!1 -Q:l a:: :::;:) t-0 I 0 1 z pH Figure 5.27. Turbidi ty as a func tion of pH for the iron(iii) arsenic(v) water system in the sulphate medium. 124 4 r 0 .01M -·- 560 · 0.0075M 0.0025M -·- ---¥- 0 .001M -+- ~ 480 400 320 240 160 . • ::J lt • + t- • ~80 >- t- -~ ell -Q:: ::J • t-0 rj 0 J pH Figure 5.28. Turbidity as a function of pH for the iron(iii) arsenic(v) water system. Back titrating with sulphuric acid in the sulphate medium. 125 1000 0.01M • ~·- 0.0075M -~- I 0.0026M -z- •I • 0 .001M -+- I ~ .I • 800 I • 600 • f. ::z: X ...... 400 I z • ll 0 1 z . pH Figure 5. 29. Turbidity as a function o f pH for the iron(iii) arsenic(v) water system. Back titrating with p erchloric acid in the sulphate medium. 126 TABLE 5.8 CORRECTED PH AT THE POINT OF PRECIPITATION AND THE VARIOUS CONCENTRATIONS OF IRON(III) AND ARSENIC(V) IN THE SULPHATE MEDIUM. Fe(iii) & As(v) Corrected pH Corrected pH at at precipitation dissolution [ ] addition of NaOH H2 S04 HC104 0.2 0.885(3: 2 ) 0.1 0.985(3:2) 0.01 1. 37 (1:7) 1 .425(3:2} 0.0075 1.47(1:6) 1.53(3:2) 0.0025 1.97(1:8) 1.615(3:2) 0.001 2. 1(1:18) 1.84(3:2} *Figures in brackets represent iron(iii} to sulphate(-ii} ratios. 127 solutions in which the ionic strength was corrected to 1.0 with sodium nitrate. In solutions where the ionic strength was increased the solubility of ferric arsenate in nitrate medium decreased. Identical behaviour was observed by Sylva and Davidson [116] in the precipitation of uranates. 5 . 3.3 . 1 Nitrate Medium Turbidity me?surements were made on solutions containing both iron(iii) and arsenic(v) mixed at a ratio of 1:1. 1M sodium nitrate was then added to increase the ionic strength to 1 .0. On mixing the solutions having concentrations of 0.01M, 0.0075M, 0.0025M and 0.001M, with respect to iron(iii) and arsenic(v), at a ratio of 1:1 precipitation immediately occurred. In these experiments nitric acid was added to determine the dissolution point of the precipitated ferric arsenate and the turbidity and measured pH was recorded. The turbidity and measured pH curves were recorded and are shown in figure 5.30. Table 5.9 illustrates the corrected pH at the point of precipitation for solutions of ionic strength of 1 .0. 5.3.4 Spectrophotometric Measurements Solutions at various concentrations of iron(iii) and arsenic(v) in both perchlorate and sulphate 128 300 0.1M -(!1- I _._ • 0.01M 0.0075M -4- 1 0.0025M -z- O.OOU.4 -+- I! •I I l!J 4 200 I 4 • 100 4 . t- z • . I 0 1 2 pH Figure 5.30. Turbidity as a function of pH for the iron(iii) arsenic(v) water system in the .nitrate medium with an ionic s trength correction of one. 129 TABLE 5.9 CORRECTED PH AT THE POINT OF PRECIPITATION AND THE VARIOUS CONCENTRATIONS OF IRON(III) AND ARSENIC(V) IN THE NITRATE MEDIUM WITH AN IONIC STRENGTH OF 1.0 . Fe(iii) & As(v) Corrected pH Corrected pH at precipitation at dissolution [ ] addition of NaOH HNOa 0.1 0.86 0.01 0.9 0.0075 1. 01 0 . 0025 1.19 0.001 1. 265 ; 130 mediums were scanned in order to access the possibility of complexing over the wavelength range of 750-190 nm. 5.3.4.1 Perchlorate Medium Iron(iii) perchlorate and arsenic(v) solutions were mixed at various concentrations and the spectra scanned. These spectra were recorded and are shown in figure 5.31 where there is an obvious absorbance band at a wavelength of 425nm. Solutions at 0.1M concentration of iron(iii) perchlorate and arsenic(v) were also scanned separately but these did not produce absorbance bands in the same wavelength region. 5.3.4.2 Sulphate Medium Iron(iii) sulphate and arsenic(v) solutions were mixed so that the final concentration of 0.2M, 0.1M, 0.01M, and 0.001M with respect to iron(iii) and arsenic(v). The spectra of these solutions were scanned over the wavelength range of 750-190 nm. At 0.2M and O.lM concentrations there was no initial precipitation and the recorded spectra are shown in figure 5.32. At concentrat ions of O.OlM and O.OOlM precipitation immediately occurred on mi xing of the iron(iii) and arsenic(v) solutions. The se solutions at the lower c oncentrations we re then filtere d and the filtered solution scanne d. The filtered solution did not produce any absorbance b a nds . Solut i ons of i r on( i ii) 4.0~------~ 3.0 lron(lll) concentration 1•0.2M pH•0.5 11•0. 1M pH•O. 75 111•0.01M pH•1.82 IV•0.001M pH•2.8 2.0 tO cu (.) c a:s .Q.... 0 fiJ .Q < O IV 300 400 500 600 700 800 Wavelength(nm) Figure 5.31. Spectrophotometric scans of ferric perchlorate and arsenic acid mixed at a ratio of 1:1. 4.0~------~~------~ 3.0 Fe( lll) concentration 1•0.2M pH•0.84 11•0.1M pH•1.12 2.0 1.0 4> () c <0 .c.... 0 en .c < 0~------~----~~~~~~~~~==~===T======~~ 300 400 500 600 700 800 Wavelength(nm) Figure 5.32. Spectrophotometric scans of ferric sulphate and arsenic ;acid mixed at a ratio of 1:1. ' ...... VJ N 133 sulp~ate and arsenic(v) were scanned individually and did not produce any absorbance bands in the same wavelength region. In order to identify any sulphate complexes, solutions at various concentrations of iron(iii) sulphate were prepared. Iron(iii) perchlorate solutions were also prepared at the same concentrations. Since iron(iii) hydroxy ions and iron{iii) sulphate complexes have absorbance bands in the same wavelength range the iron(iii) perchlorate solutions were used as the reference solution in order to eliminate any iron{iii) hydroxy ions interferences, so that only absorbance bands of iron{iii) sulphate complexes could be observed. Figure 5.33 illustrates the absorbance bands due to sulphate complexes . Further experiments were performed in order to identify any sulphate complexes . Iron(iii) perchlorate solutions were prepared and small aliquots of sodium sulphate were added. These were then scanned after each addition of sodium sulphate. However these solutions did not produce any absorbance bands . 5.3.5 Discussion Taking into consideration all of the results obtained from the titration, turbidity, and spectrophotometric measurements and by using a computer 1.0~------~ Fe (lll) oono entratlon 1•0.1M pH•1.62 ~ (.) c co ..Q... 0 CT.) ~ 01------~------~--~L-----~~======~==~=====d200 300 400 500 800 700 Wavelength(nm) Figure 5.33. Spectrophotometric scans showing ferric sulphate complexes. 135 mass balance modelling program for producing log activity versus pH · diagrams a revised set of consistent free energy data was obtained for the iron(iii) arsenic(v) water system in the different mediums studied. The results obtained in the perchlorate medium were considered first, since this ligand is less likely to form complexes. The corrected pH, at which precipitation occurred, was taken from the turbidity measurements as the pH at which the onset of precipitation of iron(iii) arsenate had commenced. The titration and spectrophotometric measurements were used to identify any complexes within the system. At the high concentrations of 0.2M and O.lM and at pH values below 0.5 it can be observed from figure 5.34 that upon mixing the iron(iii) perchlorate and arsenic(v) solutions the initial reaction is: Fe3 + + H3AS04 = FeHzAS04 2 + + H+ and then on increasing the pH the reaction is: FeHz As04 2 + = FeAS04 ( s) + 2W making the overall reaction up to the point of precipitation: Fe3 + + H3 As04 = FeAS04 ( s) + 3H+ At concentrations of O.OlM and O.OOlM with respect to both iron(iii) and arsenic(v) the initial pH of the mixed solutions are in the pH region o f 1.5- 2 . 5. From observing the stability diagram in figure 5.34 in -· 1 Fe+++ 2 Fe ~H++ 1 + Represents experimental turbid lty measurements. 3FeC~Hl2+ 4fe((jHJ3CAQJ 9 2 \ Sfe((jHJ4- s ·Fe2 ((jHJ 2++++ 7 Fe3 (eJHJ 4+++++ 3 .. 8FeCICJH(Sl 9FeRsCJ4(5) . a 4 . 10 FoHAsel4+ ,...... 11 Fe H2Rs CJ4++ > 12H3As(j4(AQl 13 H2Rs -sCl) 1 Tl4·- · a: 14 HAs (j4-- c... ,...... 6 1 5 A s CJ 4.- -- ...... pH2(j:O/O/O ...... 3 pH:0/0/_.25 ·2 pfe (III) :0/.25 ::7 4 (J) pRs (VJ :0/.25/0 ~8 0 1 2 3 4 5 6 7 8 9 10 11- 12 13 14 pH Figure 5.34. Log .activity pH .diagram for the iron(iii) arsenic(v) water system in the perchlorate medium...... \...U 0'1 137 this pH region the predominant ion is FeHAs04+ and the initial reaction is- : Fe3+ + H3AS04 = FeHAs04+ + 2H+ and then on increasing the pH the reaction is: FeHAs04+ = FeAs04 (s) + H+ making the overall reaction to the point of precipitation: Fe3+ + H3AS04 = FeAs04 (s) + 3H+ From the titration curves in figure 5.22 it can be observed that curves I,II,III, and IV all exhibit an equivalence point at a pH of 4. When observing the stability diagram in figure 5.34, at a pH of 4 there exists the conversion of FeHAs04+ to Fe (OH) 2 + • From the spectrophotometric measurements an absorbance band exists at the wavelength of 425nm at concentrations of 0.2M and O.lM with respect to iron(iii) and arsenic(v) which can be observed in figure 5.31. This absorbance band is not observed in any of the spect~a obtained for the same concentrations of iron(iii) perchlorate and arsenic(v) solutions scanned separately. Therefore the peak is attributed to the complex FeH2AS04 2+. Figure 5.35 is the distribution diagram of this system at a concentration of 0 . 01M with respect to iron(iii) and arsenic(v) . It can be observed that the complex FeHAs04+ is the predominant ion in the pH range of 1.5-7 but also at a pH of 1.82 the complex FeH2AS04 2 + T:298.15 100 4 5 1 Fe+++ 2 fG ~H++ 90 3fG(~H)2+ 4Fe(~HJ3(RQJ 80 5Fe(l'JHJ4- 6Fe2(l'JHJ2++++ 70 7 Fe3 (l'JHJ 4+++++ 8FeHRs~4+ 60 9 Fe H2Rs l'J4++ 10H3Rs~4(RQJ so 11 H2Rs l'J4- 12 HRsl'J4-- · 40 13 As l'J4--- pH2l'J:O / O 30 pH:0/.1 pfe(IIIJ:2/0 pRs(VJ:2/0 -20 ::1- 0 Q)o I.J._ :--: 0 1 2 3 4 5 s 7 8 9 10 11 12 13 14 pH Figure 5.35. Ionic distribution curves for the iron(iii) arsenic(v) water system at a concentration of 0.01M with respect to iron(iii). 139 still has some effect on the system. Therefore from figure 5.31 on curve III (which represents a solution at pH 1 . 82) a small peak still exists and this is attributed to the small effect of FeH2AS04 2+. Also by observing curve IV 1n figure 5.31 an extremely small absorbance band appears at the same wavelength. This is attributed to the same complex but at this concentration it has an even weaker effect, this can also be observed from the distribution diagram in figure 5.36. In order to observe the conversion of at around a pH of one(1) additional spectrophotometric and titration measurements were obtained. A solution at a concentration of 0.01M was back titrated with perchloric acid. The initial pH of this solution is 1.5, by back titrating,taking aliquots of the solution during titration and scanning for any peaks, the conversion of FeH2AS04 2+ to FeHAs04+ may be observed. No change however was observed in the titration curves, colour of the solution, or spectra during the addition of perchloric acid. Also at a concentration of 0.001M an identical experiment was run to observe possibly the conversion of the complex FeHAs04+ to Fe3 +. Once again no conclusive evidence was obtained . Interpretation of the turbidity measurements and titration results led to the stability diagram shown in figure 5.34 and in turn to the free 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH Figure 5.36. Ionic distribution curves for the iron(iii) arsenic(v) water system at a concentration of 0.001M with respect to iron(iii). 141 energies of formation for the particular species involved. These are listed in table 5.10. The free energies of the hydroxy species of iron(iii) were calculated from the hydrolysis constants obtained from work done by Khoe and co-workers [117]. Iron(iii) is known to complex with perchlorate ion to form FeCl04 2 + [118,119] and the complexation constant for this complex is readily available from th€ NBS tables [45]. This complex was added to the stability diagram in order to observe whether there was any effect. A modelling procedure was used to show that an influence of this complex in this system is negligible. The complex FeC104 2 + was too weak to have any effect at all on the solubility of iron(iii) arsenate. The nitrate medium behaved in a similar manner as the perchlorate medium. The precipitation points obtained from the turbidity measurements were used to produce the stability diagram shown in figure 5.37. The free energies of the solid species were the same as in the perchlorate medium. A list of the free energies used to obtain the stability diagram is shown in table 5.11. Nitrate ion is also known to complex with the ferric ion to form complexes such as FeN03 2 +. The free energy for this species listed in the NBS tables [45] was used and the complex then included in the 142 TABLE 5.10 SET OF FREE ENERGIES FOR THE IRON(III) ARSENIC(V) WATER SYSTEM IN THE PERCHLORATE MEDIUM. Species t:.Gr, 298.15k Source (Kcal/Mol) Fe3 + -4.06 [51) FeOH+ -57.0 [117] Fe(OH)z+ -109.75 [117) Fe (OH) a (aq) -157.7 [117] Fe(OH)4- -198.39 [ 49] Fez (OH) z 4 + -117.1 [117] Fea (OH)4°+ -229 . 34 [117] I ' FeOOH(s) -109.82 r [51] I Ha AS04 ( aq) -183.08 ' [ 45) Hz As04- -180.01 [45] ! ' HAS04z- -170.79 ' [ 45] i AS04 3 - -154.97 [45] ' ' I *FeAS04 ( s) -189.25 *FeHAs04+ -188.5 : ' ! *FeHzAS04Z+ -189.7 : ; *Determined in this work. +Represents experimental 1 Fe+++ turbi.dit~ measurements wi~hout 2 FF.e CJ( ~+H+l 1 an 10n1c strength correcteon. 3 e u 2 + z With an tonic strength 4 Fe ( CJH l 3 ( RQ l 2 ' corrected to one. 5 Fe ( CJ H l 4 ·- l!l With an ionic strength 6 Fe2(CJHl2++++ corrected to [1191. 7 Fe 3 ( CJH l 4+++++ 3 8FGCJCJH(Sl 9FeRsCJ4(5) 4 8 10 FGHRsC14+ 11 Fe H2Rs CJ 4++ > 12 H3RsC14 (RQl ~5 13 H2Rs C14- (J) 1 a: 14 HRsC14-- a.. _s ' . . . .. 15 Rs C14-- 5 pH2~:0/0/0 - 3 pH:0/0/.25 pFe (!Ill :0/.25/0 -::7 2 4 (I) pAs (Vl :0/.25/0 ~8 0 1 2 . 3 4 5 6 9 10 11 12 13 14 pH Figure 5. 3 7. Log activity pH diagram for the iron(iii) arsenic(v) water system in the nitrate medium. 144 TABLE 5 . 11 SET OF FREE ENERGIES FOR THE IRON(III) ARSENIC(V) WATER SYSTEM IN NITRATE MEDIUM. Species 6Gt ,298.15K Source (Kcal/mol). Fe3 + -4.06 . [51] I FeOH2 + -56.95 [117] ! Fe(OH)z+ -108.38 [117] Fe (OH) 3 (aq) -157.7 [117] Fe(OH)4- -198.39 [49] Fez(OH)z4+ -117 . 07 [117] Fe3 (OH) 4 11 + -229 . 34 [117] FeOOH(s) -109.82 [51] H3 AS04 ( aq) -183.08 [ 45] Hz As04- -180.01 [ 45] HAS04 2 - -170.79 [ 45] AS04 3- -154 . 97 [ 45] *FeAS04 (s) -189 . 25 *FeHAs04 + -188 . 5 *FeHzAS04 2 + -189.7 *Determined in this work 145 stability diagram calculations. This had no significant effect on the speciation model. Further experiments in this medium were performed in solutions in which the ionic strength was corrected to one by the addition of sodium nitrate. The pH at which precipitation occurred in these experiments was lower by about 0.5 pH units at concentrations lower than 0.01M. It appears, that with an increase in ionic strength, the pH at which precipitation occurs decreases. These experimental points are plotted in figure 5.37. Three sets of experimental points were plotted, firstly the experimental points obtained without any ionic strength correction, secondly with the ionic strength corrected to one, and thirdly an ionic strength correction greater than one [120]. The iron(iii) arsenic(v) water system in the chloride medium behaved in a similar manner to the nitrate and perchlorate mediums. The precipitation points obtained from the turbidity measurements and the free energies for the solid species used in the perchlorate and nitrate mediums were used to produce the stability diagram for the chloride medium shown in figure 5 . 38. A list of free energies used to obtain the stability diagram is shown in table 5.12 . Chloride ion also complexes with ferric ion. A wide variety of the complexes have been reported . 1 Fe+++ + Represents experimental 2FeelH++ .1 3 F.e ( el H l 2 + turbidity m~aaurementa. 4fe(elHl3(RQJ 2 ' 5fe(elHJ4- .. 6 Fe2 CelHl 2++++ 7 Fe 3 ( eJH J 4+++++ 3 8FeeJelHCSJ 9FeR~eJ4CSl 8 4 10 FeHR~el4+ 11 Fe H2R~ el4++ > 12 H3Rsel4 CRQJ -sCl) 1 13 H2Rsel4- a: 14 HRsel4-- _60.. 15 As el4-- pH2el:O/O/O .. pH:0/0/.25 ·2 4 pfe (!!Il :0/.25/0 C1) pRs (Vl: 0/ . 25/0 ~BL------~--~~~~--~~~~--~~77~~~~ 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH F'i gure 5.38. Log activity pH diagram for the iron(iii) arsenic(v) water system in the chloride medium. 147 TABLE 5.12 SET OF FREE ENERGIES FOR THE IRON(III) ARSENIC(V) WATER SYSTEM IN THE CHLORIDE MEDIUM. Species 6Gt I 298.15 K Source ( Kcal/rnol) . Fea+ -4.06 [51] FeOH2 + -57.0 [117] Fe(OH)2+ -109.75 [117] Fe (OH)a (aq) -157.7 [117] Fe(OH)4- -198.39 [49] Fe2 (OH) 2 4 + -117.1 [117] Fea (OH)4!!+ -229.34 [117] FeOOH(s) -109.82 [51] HaAS04 (aq) -183.08 [ 45] H2 As04- -180.01 [45] HAS04 2- -170.79 [ 45] AS04 3 -:- -154 . 97 [45] *FeAS04 (s) -189.25 *FeHAS04+ -188.4 *FeH2 As04 2 + -189.6 *De termined in this work. 148 Wendt and Strehlow [l21] have d etermined formation constants .f or complexes such as Fe H20Cl 2• and FeC12 + , whe reas Wendt[122] postulates the compl e xes FeC12 + and FeClOHFe4+. Naumov [46] gives the free energy data for the complexes FeC12+, FeCl2+, and the neutral species FeCb(aq) . Naumov ' s data was included in the stability diagram calculations for this medium in order to observe any possible effect of chloride complexing . It was shown that these complexes were to weak to have any significant effect on the speciation model. The model for the iron(iii) arsenic(v) water system in the sulphate medium is suggested as being the most complex. A large variety of iron(iii) sulphate complexes are reported in the lite r a ture and these are l i sted in table 5.13. The experimental points obtained from the turbidity measurements in the sulphate medium are plotted on figure 5.39 which is the log activity pH diagram obtained for the perchlorate medium. The experimental dissolution points at the lower concentrations of O.OlM to O. OOlM were obtained by back titration since precipitation was immediate on the preparation of the solutions. This log activity pH diagram represents a ratio of iron(iii) to sulphate(-ii) of 3:2 for the precipitation point s ob taine d for 0. 2M and O.lM and for the dissolution points obtained from pe rchloric acid addition at the lower concentrations. However, for the 149 TABLE 5.13 IRON(III) SULPHATE COMPLEXES FROM VARIOUS SOURCES. Complex Source FeS04+ [121-137] Fe ( S04 )2 2 - [121-124,126-135,137,138] FeHS04 2 + [121,124,125,129,131-133,135,137] FeS04 .HS04 (aq) [121,129,132] Fe2 (OH)2S04 2 + [123] Fe2 (OH) 2 (S04) 2 [123,138] Fe ( S04 ) 3 3- [126,128] FeOHS04 (aq) [136] ' 1 Fe+++ + Represents experimental 2FeelH++ 1 turbidity measurements in the 3Fe£elHl2+ sulphate medium. 4 Fe (elHJ3(RQJ 2 ' Sfe(elHl4- 6 Fe2 (elHl 2++++ 7 Fe3 (elHl 4+++++ 3 8FeelelH(SJ 9 FeRsel4 (5) a 10 FeHRsel44- ,..... 4 11 Fe H2As el4++ > 12H3Asel4(AQl ~5 Cl) 13 H2As el4- a: 14 HRs el4-- a.. 15 Rsel4--- ~6 5 pH2el:O/O/O -...... pH:0/0/.25 ·2 pfe (!!!) :0/.25/0 ::7 4 Q) pRs (Vl: 0/ . 25/0 ~ 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH Figure 5.39. Log activity pH diagram for the iron(iii) ..... Vl arsenic(v) water system in the p erchlorate medium. 0 151 dissolution points obtained by the addition of sulphuric acid the ~atio of iron(iii) to sulphate{-ii) should be changed for each point on the diagram, however a satisfactory model which could describe these experimental points could not be obtained. The spectrophotometric measurements indicated that the iron(iii) arsenate complexes however still exist in the sulphate medium as is shown in figure 5.32. The absorbance bands on curves I and II at a wavelength of 425nm are identical to the absorbance bands in the perchlorate medium shown in figure 5 . 31. These absorbance bands therefore represent the same complex for both mediums that is the iron{iii) arsenate complex Solutions of iron(iii) sulphate(-ii) and arsenic{v) were individually scanned but did not exhibit any peaks in the 425nm wavelength region. With the fact that in the sulphate medium the iron{iii) arsenate complex exists and together with the precipitation points and dissolution points obtained from the turbidity measurements these results did not allow a satisfactory model to be obtained. Since no model could be obtained with the species considered to this point, the question of whether an iron(iii) arsenate sulphate solid was precipitated instead of iron{iii) arsenate was considered and analysis of the precipitates for sulphate were conducted. 152 In order to do this, solutions at a concentration of O.OlM were prepared, one with an iron{iii) to sulphate{-ii) ratio 1:1.5 and the other with a ratio of 1:5. Upon mixing the solutions at this concentration, precipitation was immediate and this precipitate was separated and scanned for iron, arsenic, and sulphur using EDEX measurements on the electron microscope . No sulphur was detected in either solid, only iron and arsenic. Therefore from these results precipitation of iron(iii) arsenate was confirmed. From the log activity pH diagrams for this system, in perchlorate, nitrate, and chloride mediums, the minimum solublity of iron{iii) arsenate lies at pH 5 at a concentration of O.OOOOlM for arsenic{v) and iron(iii). At this pH, in the sulphate medium, the solubility line of iron(iii) arsenate would not differ substantially. In all the log activity pH diagrams postulated the free energy value for amorphous goethite {FeOOH) was used. This value is consistent for colloidal particles of about 5 nm diameter. This is the size of particles that was precipitated in all the experiments. However, with time, the particle size of goethite increases. If in the stability diagram the free energy of goethite associated with larger particle size is included the log activity pH diagram shown in figure 5 .40 0 r:29B.15 ------~1 Fa+++ 2 FaCJH++ 1 9 3·fa{elHl2+ 4 Fa CelHl 3 {RQJ 5FaCCJHJ4- 2 ' .. ..' 6 Fa2CelHl2++++ 7 Fe 3 CelH l 4+++ ++ 3 a Fe Cl CJH ( S l ( -1 1 4 . 0 l 9 FeRsCJ4 CSJ 10FoHRsCI4+ ,...... 4 8 11 FeH2RsCl4++ > 12 H3AsCI4 CRQl -s(I) 1 13 H2Rs Cl4- a: 1 4 HAs Cl4-- ~ 15 AsCI4-- ,...... 6 pH2Cl:O/O/O ...... pH : 0/0/.25 ·2 pfe {!Ill :o/. 25/0 cP . pAs CVJ: 0/ . 25/0 l.J... ~B · L-----~--~~~~--~~~~~~~~~~~~~~ 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH Figure 5.40. Log activity pH diagram for the iron(iii) arsenic(v) water system in the perchlorate medium showing the stability re gion of crystalline goethite (FeOOI-1}. 154 is obtained. The free energy of crystalline goethite was given the value of -116.87 kcal/mol [51] . From this diagram the stability region of ferric arsenate decreases to a pH of 2. This is consistent with the early work of Nishimura and Tozawa [60]. Therefore it is suggested that at a pH of 2 this is the incongruent point in this system. 5.4 Arsenopyrite (FeAsS) and Loellingite (FeAS2) Several authors [107,108] have produced potential pH diagrams for the iron arsenic sulphur water system in order to show the region of stability of arsenopyrite. A potential pH diagram for this system was generated by computer using the selected thermodynamic data shown in table 5.14 . This potential pH diagram is shown in figure 5.41. This diagram is at a concentration of O. OOOlM for arsenic, sulphur and iron. From this diagram arsenopyrite is shown to exist only as mixtures with other solids and at very low potentials (below the hydrogen potential). 155 TABLE 5.14 STANDARD FREE ENERGY OF FORMATION DATA FOR THE IRON ARSENIC SULPHUR SPECIES AT 25° C . Species f:..Gt ,298.15K Source Kcal/mol. H3AS04(aq) -183.08 [45] H2 AS04- -180.01 [ 45] HAS04 2- -170.79 [45] AS04 3 - -154.97 [ 45] As2 0!5 ( s) -186.97 [ 45] H3 AS03 (aq) -152.92 [ 45] H2 As03- -140.33 [45] HAS03 2 - -121.71 [46] AS03 3- -102.95 [46] Aso+ -39.15 [45] As2 03 ( s) -137.72 [ 45] ASH3 (g) +16.47 [ 45] As(s) 0.0 [45] HS04- -180.67 [ 45] S04 2- -177.95 [ 45] H2 S ( aq) -6.65 [45] Hs- +2.89 [ 45] S2- +21.17 [49] s ( s) 0.0 [ 45] Fe2 + -18.86 [ 45] 156 TABLE 5.14 CONT., Species f).Gt I 298.15K Source Kcal/mol. FeOH+ -66.3 [ 45] Fe(OH)2 (aq) -104.24 [49] Fe(OH)2(s) -114 .7 [49] Fe(OH)3 - -146.96 [ 45] Fe(OH)4 2- -183.96 [ 45] Fe3 + -4.06 [51] FeOH2 + -54.83 [45] Fe(OH)2+ -104 .68 [ 45] Fe(OH)3 (aq) -157.57 [ 45] Fe(OH)4- -198.39 [49] Fe2 ( OH) 2 4 + -111.7 [45] Fe3 ( OH) 4 5 + -221.46 [ 49] FeOOH(s) -109.82 [51] Fe(s) 0.0 [ 45] FeAsS(s) -26.2 [97] FeS2 (s) -39.89 [ 45] FeS(s) -23.99 [ 45] FeAs2 (s) -12.5 [ 97] As4 So~ ( s) -31.43 [ 98] *AsS2- -5.6 *HAs2 S4- -24.3 *H2As2S4(aq) -31.5 157 TABLE 5.14 CONT., Species b.Gt , 298.15K Source Kcal/mol *As2 S3 (s) -21.68 *FeH2As04+ -203.0 *FeS04 (aq) -201.23 *FeHAs04 (s) -202.33 *Fe3 (AS04 )2 (s) -419.0 *FeHAs04+ -188.5 *FeH2AS04 2 + -189.7 *FeAs04 (s) -189.25 *Determined in this work FE-AS-S-H20 SISTEM AT 0.0001M 2.0~----~------~------, AF AC AA -1.0 -1.5 A -2 . ~. 2. 1.!. 6. 8·. 10 . 12 16. PH Figure 5.4 7. Potential pH diagram for the iron arsenic sulphur water system at a concentration of 0.0001M for all aqueous species Refer to Appendix 5 for legend. 159 A potent ial pH diagr am for t h e i r on arse nic water system was also generate d by c o mp u ter u s ing selected thermodynamic data from table 5. 1 4. This potential pH diagram is shown in figure 5 . 42. The diagram is a t a concentr ation of O. OOOlM for bo th arsenic and iron. From this diagram the area marke d K represents where loellingite (FeAs2 ) may be formed . Howe ver , arsenic metal must also be taken into account which is not shown on the diagram, so that only the area shaded in black in figure 5.42 is where loellingite may be formed. In the pH region above 14 under reducing c onditions with hydrogen, it may be possible for the hydrothermal synthesis of loellingite. Attempts to synthesize arsenopyrite and loellingite at 25°C in the laboratory have been unsuccessful. However, there has been reference to the formation of arsenopyrite at temperatures above 300° C in the presence of NH4Cl[139] . Therefore it seems that arsenopyrite and loellingite are probably only stable in high temperature hydrothermal systems, so that it is unlikely that arsenopyrite and loellingite could be produced as stable residues at 25°C. However, there is evidence that arsenopyrite and other sulphides can be formed biologically to b e used as likely end products 1n the chemical treatment of waste solutions. IRON-ARSENIC-WATER AT 0.0001 2. 0..------...... ---r------r-----, A FE E-2 l~l B FE HAS 00 E-1 IAQI C FE H2 AS Qlt E-2 (RQ) D FE02H 8 E FE AS Ott F FE3 RSi 00 1. E C FE [0 Hl3 E IADl H FE (0 Hl~ E2 (RQJ I FE [0 Hl~ E IAQl D J FE [0 Hl2 I t{f'EAS2 l FE r. w l -2.0~-----~~-----~------~------~------~------~-----~------~ 1!. 8. pH· Figure 5. 42. Potential pH diagram for the iron arsenic water system at a concentration of 0.0001M for all aqueous species. 161 CHAPTER 6 CONCLUSION Log activity pH diagrams have been obtained for the arsenic (iii) sulphur(-ii) water system, the iron(ii) arsenic(v) water system in sulphate and perchlorate mediums and the iron (iii ) arsenic(v) water system in perchlorate, nitrate and chloride mediums. R.emov ing arsenic from process solutions in the form of arsenic sulphide (orpiment) is a commonly occurring process. This study resulted in the identific ation of three complexes AsS2-, HAs2S4-, and in the arsenic(iii) sulphur(-ii) water system. This in t urn provided more reliable free energy data for the above mentioned speci es and the solid arsenic sulphide (As2S3 ), than can be found i n the literature. From the log activity pH diagram (figure 5.3) produced from this work the solubility of ars enic sulphide was c alculated. Arsenic s ulphide is least s ~~uble in the pH range of 0-5 and the concent rat ion of arsenic in this range is considerably higher than previously considered. Studying the iron(ii) arsenic(v) water system resulted in the identification of the complex FeH2As04+ in the perchlorate and sulphate me diums, and the neutral species FeS04 (aq) in t he sulphate medium. With the use of free energy data calculated from pH titrati ons 1n 162 this· system log activity pH diagrams were obtained. From these diagrams in both the perchlorate (figure 5.10) and sulphate (figure 5.17) mediums it appears that ferrous arsenate (Fe3 (As04 )2) has a mini mum solubility in the pH range 6-7 and under reducing conditions could possibly be a potential way of stabilising arsenic. The iron(iii) arsenic(v) water system was studied in the perchlorate, nitrate and chloride mediums. Behaviour in these mediums was very similar. The complexes FeHAs04+ and FeH2AS04 2+ were deduced and free energy data calculated. From the log activity pH diagrams (figures 5.34, 5.37, and 5.38) in the three mediums the minimum solubility of ferric arsenate (FeAs04) lies within a pH of 5. This system was also studied in the sulphate medium. In this medium complexing proved to be far more complicated. Further titration and spectrophotmetric studies were undertaken in order to determine complexing in this medium but without obtaining a satifactory model. At a pH of 5 where ferric arsenate is least soluble, from the log activity pH diagrams (figures 5.34, 5.37, and 5.38) in this region it is concluded that in the sulphate medium this solubility would not differ substantially from the other three mediums studied. 163 Ferric arsenate at a pH of 5 is shown be relatively stable , relative t o amorphous ferric hydroxide. Amorphous ferric hydroxide is less stable than goethite ( a -FeOOH) , and in the long term the amorphous ferric hydroxide would be expected to convert to goethite. Under these conditions the stability region for ferric arsenate is decreased to a pH of 2 and a concentration of O.OlM for both iron(iii) and arsenic(v) (figure 5 . 40) . Arsenopyrite (FeAsS} and loellingite (FeAs2) appear to be extremely stable compounds under reducing conditions . From this work loellingite under reducing and high pH regions could possibly be synthesised, however, the hydrothermal synthesis of arsenopyrite may not be possible. In the potential pH diagram (figure 5.41} it can be seen that no region exists where arsenopyrite uniquely occurs. There needs to be further investigations which relate the possibility of forming arsenopyrite and loellingite as a means of stabilising arsenic . One such area could be a study of the formation of these compounds by biological reactions . This work has provided a better understanding of the speciation in the systems studied. 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H3AS04 (aq) -183.08 [45] H2 AS04- -180.01 [ 45] HAS04 2- -170.79 [ 45] AS04 3- -154.97 [45] H3 As03 (aq) -152.92 [ 45] H2 As03- - 140.33 [45] HAS03 2 - - 121.71 [46] AS03 3- -102.95 .. - [46] Aso+ -39.15 [ 45] Mg2+ -108.7 [ 45] MgOH+ -149 .8 [45] Mg (OH)2 ( aq) -183.9 [ 45] Mg (OH)2 (s) -199.21 [ 45] [61] Mg3 (AS04 )2 ( s) -663.3 MgHAs04 (s) -283.8 [61] Ca2 + -132.3 [ 45] caoH+ -171.7 [ 45] Ca (OH) 2 (aq) -207.49 [ 45) Ca(OH)2 (s) -214.76 [ 45 J Ca ( As02 ) 2 ( s) -308.84 [62] 183 APPENDIX 1 CONT., Species Free Energy Data Source ( 6.Gt , 298.15K Kcal/rnol.) Ca ( AsOz ) z • Ca ( OH) z ( s) -531.91 [62] CaH4 (AS04 )z (s) -490 . 91 [62] CaHAs04 (s) -307.65 [62] Cao Hz (As04 ) 4 ( s) -1347.24 [ 62] Caa (As04 )z (s) -73L47 [ 62] Caa (As04 )z .Ca (OH)z (s) -948.42 [62] Baz + -134.0 [45] BaOH+ - 174.6 [45] Ba(OH)z .8Hz0(s) -667.5 [ 45] BaHAs04 (s) -312.38 [7 4] Baa (As04 )z (s) -736.62 [7 4] Fez+ -18.86 [ 45] Fe OW -66.3 [45] Fe(OH)z (aq) -104.24 [49] Fe(OH)z (s) -114.7 [49] Fe(OH)a- -146.96 [ 45] Fe(OH)4Z- -183.96 [ 45] FeHAS04 (s) -204 . 36 [69] Fea (AS04 }z ( s) -422.21 [69] COz (g) -94.254 [ 45] COaz- -126.17 [ 45] HCOa- -140.26 [45] Hz COa ( aq) -148.92 [ 45] 184 APPENDIX 2 Standardisation of Arsenic(v) Solutions. Reagents: Potassium Iodide Sodium Hydrogen Carbonate 0.1N Sodium Thiosulphate Concentrated Hydrochloric Acid Procedure: Place 20.0 ml of 0.025M(0.1N) arsenic pentoxide (As20~) in a conical flask and add 11 ml of concentrated hydrochloric acid to make the solution in the flask 4N. Displace the air by adding two 0.4g portions of sodium hydrogen carbonate, add 1g of pure potassium iodide, replace stopper, mix the solution and allow to stand for 15 minutes. Then titrate the solution against standardised 0.1N sodium thiosulphate with continuous stirring. 185 APPENDIX 3 Standardisation of Iron(iii) using EDTA. Reagents: 0.05M Ethylene Diarnine Tetra Acetic Acid (EDTA) Variarnine Blue Procedure: Pipette 25.0 rnl of 0.05M iron(iii) solution into a conical flask and dilute to 100 rnl with distilled water. Add 5 drops of indicator solution (variarnine blue) and warm the contents of the flask to 40°C, titrate with standardised (0.05M) EDTA solution until the initial blue colour of the solution turns grey just before the endpoint, and with the final drop of reagent changes to yellow. 186 APPENDIX 4 Standardisation of Iron(ii) Solution using eerie Sulphate . Reagents : O.lM eerie Sulphate N-Phenylanthranilic Acid Procedure : Place 25.0 ml of O. lM iron(ii) solution 1n a stoppered flask, add 0.5 ml of N-phenylanthranilic acid and titrate agai nst standardised O. lM eerie sulphate. This standardisation was done continuously under high purity nitrogen. 187 APPENDIX 5 Legend for figure 5.41. Area Species A AsH3 (g) + Fe(s) B AsH3 (g) + Fe2 + c FeAs2 ( s) + AsH3 (g) D FeAs2 (s) + As(s) E AsH3 (g) + FeS(s) F AsH3 (g) + FeAs2 (s) + FeS(s} +Fe(s} G FeAs2 (s} + FeS(s) + FeAsS(s} + As(s} H AsH3 (g} + FeAs2 (s} + Fe(s} I FeAs2 (s) + Fe(s) J FeAs2 (s) + Fe(s} + As(s) K FeAsS(s) + FeAs2 ( s) + FeS2 (s) L FeAs2 (s} + Fe (OH)2 (s} M Fe(OH)2(s) N FeS(s} + FeAs2 (s) 0 FeS(s) + FeAs2 (s} + Fe(s) p FeAsS(s} + FeS2 ( s) + FeAs2 (s) + As(s} Q FeS2(s) + As(s) R As(s) s H2As2S4 (aq) T FeS2(s) u FeS2 ( s) + Fe3 ( As04 ) 2 ( s) v Fe3 (As04 )2 ( s) 188 APPENDIX 5 CONT., Area· Species w Fe3 (As04 )2 (s) + FeOOH(s) X Fe2 + y FeS(s) z FeAs2 ( s) + Fe (OH)2 (s) + As(s) AA FeOOH(s) AB Fe3 ( As04 ) 2 ( s) + Fe (OH)2 (s) ' AC FeAs04 (s) + FeOOH(s) AD FeOOH(s) + Fe3 (As04 )2 (s) + FeAS04 (s) AE FeAs04 ( s) AF FeH2 As04 2 + + Fe3 + :