STUDIES ON GOLD DISSOLUTION AND LOSS OF GOLD

DUE TO GANGUE MINERALS IN CHLORIDE SOLUTIONS

A thesis submitted in fulfilment of the requirements for the degree of

Doctor of Philosophy (Ph. D.)

by

You Jin Kim (B.Sc., M.Sc.)

Centre for Minerals Engineering

School of Mines

Faculty of Applied Science

University of New South Wales

AUSTRALIA

July 1995 UNiVtrtSi'i V OF N.S.W. 26 JUN 1996 LIBRARY STATEMENT

I hereby declare that this submission is my own work and that, to the best of my knowledge and belief it contains no material previously published or written by another person nor material which to a substantial extent has been accepted for the award of any other degree or diploma of a university or other institute of higher learning, except where due acknowledgement is made in the text.

I also declare that the intellectual content of this thesis is the product of my own work, even though I may have received assistance from others on style, presentation and language expression. ACKNOWLEDGMENTS

I would like to express my heartfelt gratitude, first and foremost, to Dr. Tam Tran, my supervisor, for directing my studies. His enthusiasm, advice and guidance have contributed immeasurably to my research.

1 also wish to extend my appreciation to Dr. A.C. Partidge, Dr. P.L.M. Wong and Dr. D. Cheng for their encouragement and advice in various aspects of my work. My sincere thanks are due to Dr. A. Davis and Mr. M. Yastreboff for reviewing my thesis.

Also, I wish to thank Mr. J.K. Song and Ms. L. Lau of the Centre for Minerals Engineering, for their helpful advice and assistance in the experimental work.

To Glenn, Bill, Hoan, and Raul, I thank them for their help relating to my thesis, and to the postgraduate students of the Centre for Minerals Engineering, for their friendship and in sharing the experimental equipment.

Last but not least, I deeply appreciate the moral support and encouragement of my family during my study. A special thanks is expressed to my wife, Hea Seon, for her love, encouragement and support of my endeavour, and also, to my two sons, Han and Soo Han, whose forbearance have kept me sensible through the progression of my study. ABSTRACT

In certain applications, the extraction of gold by the chlorine/chloride system could offer a clear advantage over cyanidation. This is because it is possible to electrolytically produce chlorine economical from ground water brines, which are common in the gold fields of Western Australia where natural saline waters are available. The precipitation of gold from chloride solutions and the corrosiveness of acidic chloride solutions are major concerns for the application of such a system.

The present study has been conducted to investigate basic aspects of the chemistry of the gold-chloride system as well as to identify scientific considerations for the leaching of gold-bearing refractory ores and concentrates using the chlorination process.

Firstly, the kinetic study involved measurements of the gold dissolution rates and development of the modified rate equation for gold strips of thin rectilinear shapes in chloride-hypochlorite solutions. Tests to determine the stability of gold chloride solutions in the presence of sulfide minerals (Muir, 1987) such as sphalerite, galena, pyrite and chalcopyrite were also conducted. In all these investigations, the rate of gold dissolution was measured for solutions containing NaCl (50 to 200 g/1), and OC1 (1 to 10 g/1) within the pH range from 2 to 6. The dissolution of gold in these chloride-hypochlorite systems is controlled by a kinetic step. A reaction rate equation at the gold-solution interface derived from the experimental results, was as follows:

= k exp{- 2Q1U1"'oI}[H* ]°3 [ NaCl]0 5 [OCt ] dt RT On the basis of the shrinking core model developed for application to thin rectilinear geometry, the reaction kinetics can be represented as:

-32.9 A;/ / mol 1 - (1 - 7?)1/3 kc exp[ ][H+ ]°3 [NaCl]0'6 [OCl~ ]111 RT

The stability of gold chloride in the presence of sulfide minerals is strongly dependent on pH indicating that the gold chloride complex can be readily produced to metallic gold in the presence of sulfide minerals.

Secondly, electrochemical techniques based on scanning voltammetry were used to investigate the half-cells of the electrochemical reactions in chloride- oxidant solutions. Evans diagrams obtained from the anodic and cathodic half-cell reactions required for gold dissolution in chloride solutions at various pH values showed that the oxidation of gold by hypochlorite is chemically controlled. The reaction rates measured on several stationary sulfide electrodes were also established (for sulfide mineral oxidation in sodium hypochlorite solution) using the Butler-Volmer equation (Miller, 1981).

Finally, the computer modelling program was used to construct the log [M]- pH diagram. In the gold-chloride-hypochlorite system, it was found that with increasing NaCl concentration, the predominant domain of AuC14 was extended toward a higher pH. The reaction equations of gold-sulfides- chloride-hypochlorite systems based on the observed results were proposed, and based on the presence of AuS (Gaspar et al., 1994) as an intermediate species. VI

Page TABLE OF CONTENTS

STATEMENT ii

ACKNOWLEDGEMENTS iii ABSTRACT iv

TABLE OF CONTENTS vi

CHAPTER ONE

INTRODUCTION 1

CHAPTER TWO HYDROMETALLURGY OF PRECIOUS METALS PROCESSING 5

2.1 Background 5 2.1.1 Hydrometallurgical Processing of Precious Metals 5 2.1.1.1 Hydrometallurgical Processing of Minerals 5 2.1.1.2 Hydrometallurgy Processing of Electronic Scraps 7 2.1.1.3 Hydrometallurgy Processing of Metal and Mineral Wastes 10 2.1.2 Electrochemical Processes 11

2.1.3 Thermodynamic Aspects of Precious Metal Dissolution 13

2.1.3.1 Stability of Aqueous Gold Complexes 13 2.1.3.2 Eh (Potential)-pH Relationship 14

2.1.3.2.1 Gold-Water System 15 2.1.3.2.2 Gold-Ligand-Water System 16 vii

2.1.3.3 Modelling of Gold Dissolution Process 19 2.1.4 Kinetics Considerations of Precious Metal Dissolution 20 2.2 Cyanidation 22 2.2.1 Gold Dissolution in Cyanide Solutions 22 2.2.2 Electrochemical Processing of Gold in Cyanide 23 2.2.3 Thermodynamic Considerations 26 2.2.4 Kinetic Aspects 28 2.3 Chlorination 32 2.3.1 Chlorine and Hypochlorite Chemistry 33 2.3.1.1 General 33 2.3.1.2 Aqueous Chemistry 33 2.3.2 Chloride Leaching Process 37 2.3.3 Electrochemical Processing of Gold in Chloride Systems 38 2.3.4 Thermodynamic Considerations 43 2.3.5 Kinetic Aspects 45 2.3.6. Process Selection Considerations 48

CHAPTER THREE LEACHING OF REFRACTORY GOLD ORES

3.1 Refractory Ores 50 3.1.1 General 50 3.1.2 Mineralogical Characteristics 51 3.1.2.1 Siliceous Refractory Ores 51 3.1.2.2 Pre-robbing of Carbonaceous Ores 52 3.1.2.3 Sulfidic Refractory Ores 52 3.2 Leaching of Refractory Ores 53 V111

3.2.1 Direct Gold Leaching 53 3.2.1.1 Alkaline Pressure Cyanidation 53 3.2.1.2 Carbon-in-leach (CIL) and Carbon-in-pulp (CIP) 53 3.2.1.3 Leaching with Non-Conventional Lixiviants 56 3.2.1.3.1 Chloride 57 3.2.1.3.2 Iodide 60 3.2.1.3.3 Bromide 62 3.2.1.3.4 Thiourea 64 3.2.1.3.5 Thiosulphate 66 3.2.1.3.6 Thiocyanate 68 3.2.1.3.7 Polysulfide 69 3.2.2 Pretreatment of Refractory Ores 71 3.2.2.1 Roasting 72 3.2.2.2 Biological Oxidation 74 3.2.2.3 Chemical Oxidation of Refractory Ores 77 3.2.2.3.1 Oxidative Acid Leaching at Low Pressure 77 3.2.2.3.2 Oxidative Alkaline Leaching at High Pressure 79 3.2.2.3.3 Oxidative Acid Leaching at High Pressure 81

CHAPTER FOUR REACTION KINETIC MODELS

4.1 Homogeneous Kinetics 84 4.2 Heterogeneous Kinetics 85 4.2.1 Fundamentals 85 4.2.2 Heterogeneous Reactions 87 4.2.2.1 Shrinking Core Model (SCM) 87 4.2.2.1.1 Product Layer Diffusion Control 88 4.2.2.1.2 Fluid Film Diffusion Control 89 4.2.2.1.3 Surface Chemical Reaction Control 90 4.2.2.2 Grain Model 92 4.2.2.3 Capillary Pore Model 93 4.2.2.4 Random Pore Model 94

CHAPTER FIVE FUNDAMENTALS OF ELECTROCHEMISTRY AT ROTATING DISC ELECTRODES

5.1 Introduction 96 5.2 Rotating Disc Electrode System 96 5.3 Electrochemical Considerations 98 5.4 Electrode Kinetics and Reaction Mechanism 101 5.4.1 Determination of the Rate Constant 101 5.4.2 Determination of the Reaction Order 103 5.5 Interpretation of Current-Potential Curves 103 5.6 Evans Diagram 107 5.7 Butler-Volmer Equation 111

CHAPTER SIX SUMMARY OF LITERATURE SURVEY AND RESTATEMENT OF MAIN AIMS 114 CHAPTER SEVEN STUDY ON KINETICS OF GOLD DISSOLUTION AND THE STABILITY OF GOLD SPECIES IN A CHLORIDE SYSTEM

7.1 Introduction 118 7.2 Experimental 119 7.2.1 Materials 119 7.2.1.1 Chemicals 119 7.2.1.2 Sulfide Minerals 120 7.2.1.3 Gold Chloride Solution 123 7.2.2 Equipment 123 7.2.2.1 Experimental 123 7.2.2.2 Gold Leaf Tests 124 7.2.2.3 Stability of Gold Chloride Solutions 126 7.2.2.4 Analysis of Gold Chloride Solutions by Solvent Extraction 126 7.3 Kinetics of Gold Dissolution 127 7.3.1 Effect of pH 128 7.3.2 Effect of NaCl Concentration 130 7.3.3 Effect of OC1" Concentration 132 7.3.4 Effect of Temperature 134 7.4 Determination of the Reaction Order 136 7.4.1 Reaction Order with respect to Hydrogen Ion Concentration 136 7.4.2 Reaction Order with respect to Sodium Chloride Ions 137 7.4.3 Reaction Order with respect to Hypochlorite Ions 138 7.4.4 Activation Energy with respect to Temperature 140 7.5 Rate Equation for the Dissolution of Gold in the Chloride/ Hypochlorite Solution 141 7.6 Development of the Thin Rectilinear Shrinking Core Model 142 7.7 Reaction Kinetics using the Shrinking Core Model 148 7.7.1 Effect of pH 148 7.7.2 Effect of NaCl 151 7.7.3 Effect of OC1' 153 7.7.4 Effect of Temperature 155 7.8 Stability Test in Chloride System 157 7.8.1 Effect of pH 157 7.8.2 Effect of Different Sulfide Minerals 159

CHAPTER EIGHT ELECTROCHEMICAL STUDY OF GOLD DISSOLUTION IN CHLORIDE SYSTEM

8.1 Introduction 163 8.2 Experimental 164 8.2.1 Materials 164 8.2.1.1 Chemical Reagents 164 8.2.1.2 Gold Electrode 164 8.2.1.3 Glassy Carbon Electrode 165 8.2.1.4 Sulfide Electrodes 165 8.2.2 Equipment 169 8.2.2.1 Electrochemical Equipment 169 8.2.2.2 Three Electrode System 173 8.2.3 Experimental Procedure 174 xii

8.3 Scanning Voltammetry 176 8.3.1 Anodic Dissolution of Gold in Chloride Solution 176

8.3.2 Cathodic Reduction of Hypochlorite 181

8.4 Interpretation of Current-Potential Curves 186 8.5 Evans Diagram 190

8.6 Reaction Rates on Gangue Minerals 195

8.6.1 Zinc Sulfide 195 8.6.2 Lead Sulfide 200

8.6.3 Iron Sulfide 203 8.6.4 Copper Sulfide 207

CHAPTER NINE THERMODYNAMIC MODELLING

9.1 Introduction 211 9.2 Thermodynamic Data and Modelling Procedures 212 9.3 Equilibrium Solution Chemistry 216 9.3.1 The Au - Cl - NaOCl - H20 System 216 9.3.2 The Au - ZnS - Cl - NaOCl - H20 System 221 9.3.3 The Au - PbS - Cl' - NaOCl - H20 System 224 9.3.4 The Au - FeS2 - Cl - NaOCl - H20 System 228 9.3.5 The Au - CuFeS2 - Cl - NaOCl - H20 System 232 9.4 Related Kinetic Considerations on the Dissolution of Gold 235

CHAPTER TEN CONCLUSIONS AND RECOMMENDATIONS X111

10.1 CONCLUSIONS 237 10.2 SUGGESTIONS FOR FURTHER WORK 242

REFERENCE 243

APPENDICES 288 CHAPTER ONE

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INTRODUCTION

Over the last one hundred years cyanidation has been used extensively as a hydrometallurgical process for the extraction of gold from conventional oxidic gold ores. Because of the low cost and simplicity of the process for treating these ores, cyanidation for gold processing has been widely accepted. However, cyanidation is not suitable for ores containing cyanicides (cyanide consumers), such as copper oxides or silicates, as significantly higher cyanide consumptions are experienced in these cases. To overcome the major disadvantage of low extraction, alternative gold recovery processes using more powerful lixiviants have been considered. Furthermore, because of increasing awareness of environmental pollution, non-toxic reagents for gold extraction should also be considered.

The use of lixiviants other than cyanide to extract goid has been examined in an attempt to overcome these problems. Recently, several lixiviants for gold extraction, such as thiourea, thiosulphate, polysulphides and halides, have been investigated by many researchers (Davis and Tran, 1991; Qi and Hiskey, 1993 etc.). Especially, the halides have been applied for the processing of gold-bearing refractory ores and concentrates. Stable complexes with gold are formed as Au(I) complex for iodide and both Au(I) and Au(III) complexes for chloride and bromide depending on the dissolution chemistry. The relative stability of various oxidation states in aqueous 2 solution increases according to the order Cl" < Br < I", where the order of electronegavities is Cf > Br > I".

Historically (Hiskey and Atluri, 1988), gold dissolution using dissolved chlorine, had been adopted before cyanidation was first used on a commercial scale. Recently, chlorination has also been re-introduced as a pretreatment process to oxidise the carbonaceous material of certain gold ores prior to cyanidation and has been successfully employed at Carlin Mines since the early 1970's.

Generally, several advantages of chlorination are claimed. The efficient leaching of gold in acidified chloride systems is achieved for ores containing low concentrations of sulfides (less than 0.5% sulphur) because gold dissolution with chlorine is much faster than with cyanide. Such a process for chlorine pretreatment has been proposed for the heap leaching of low grade gold ores and in-situ gold extraction systems (Ahmadiantehrani et al., 1992; Fagan, 1992). In Australia, processes utilising chlorine pretreatment could have an economic advantage over cyanidation. This is because of the ability to electrolytically produce chlorine economically from ground water brines, which are common in the gold fields of Western Australia when natural saline waters are available. Major problems of reprecipitation of gold from chloride solutions and the corrosiveness of acidic chloride solutions are main concerns for the application of such a system.

Fundamental aspects of gold dissolution using chloride as a complexing lixiviant have been studied by several investigators (Nicol, 1981a; Tran et al., 1992 etc.). Their results showed that AuC14 is a predominant species at a higher pH, but gold can also exist as AuC12 , especially at high chloride concentrations and at a neutral pH. AuC12 appeared as an intermediate 3 species, which could be further oxidised to Au(III) into the bulk solution. However, chlorination has been limited in its application and information on its process chemistry is limited. Additionally, the behaviour of gold chloride complexes in the presence of sulfide minerals is not fully understood.

Electrochemical techniques based on scanning voltammetry are effective tools for studying the half-cells of electrochemical reactions. Under the imposition of a potential, characteristics of the gold oxidation and hypochlorite reduction can be determined in the form of current-potential plots. Evans diagrams, known as the graphical super-position of the anodic and cathodic current- potential curves using the mixed potential theory, can be derived to determine the reaction rate by analysing the relationship between the potential and the current density of the system. More detailed information on the kinetics of gold dissolution and the decomposition of sulfide minerals in the presence of dissolved chlorine in chloride solutions, can also be obtained from electrochemical measurements.

Stability diagrams are useful for interpreting many hydrometallurgical reactions. A detailed scientific evaluation of these reactions enhances the understanding of the stability of gold chloride complexes in the presence of sulfide minerals.

Thus, the present study was undertaken to:

• Determine the reaction order and the rate equation for gold dissolution in acidic chloride solutions, • Develop the equation for the thin rectilinear shrinking core model, • Propose the governing rate equation by using the above developed shrinking core model, 4

• Determine the stability of gold chloride complexes in the presence of sulfide minerals, • Investigate the mixed potential and the exchange current density required for gold dissolution in chloride solutions of different compositions, • Describe the relationship between the exchange current density and the mixed potential by using the Butler-Volmer equation, • Determine the rate of dissolution of different sulfide electrodes, • Construct stability diagrams for gold complexes and chlorine species, • Construct stability diagrams for gold complexes and sulphur species in the presence of sulfide minerals. CHAPTER TWO

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HYDROMETALLURGY OF PRECIOUS METALS PROCESSING

2.1 Background

As the world's demand for precious metals increases and the grade of ores decreases, finding effective and efficient methods for processing the ores available to us, and recycling previously used metals, will assume greater importance. Leaching is the process of dissolving the valuable metals from ores or concentrates and extracting the metals of interest into the solution. The feed for the leaching can also be from secondary sources such as electronic scraps, slags, dusts, residues, and wastes from other processes.

2.1.1 Hydrometallurgical Processing of Precious Metals

2.1.1.1 Hydrometallurgical Processing of Minerals

Processes for gold recovery from ores have continuously improved over the years (Eisele, 1988). Agricola (1950), in his book "De Re Metallica" (1556), suggested the use of a mixture of nitric and hydrochloric acid (aqua regia) to dissolve gold. After that, the Plattner process in 1851 (Hiskey and Atluri, 1988) had been proposed, based on the use of dissolved chlorine to extract gold from oxidised ores. This was an early hydrometallurgical approach. 6

In early 1793, Scheel (Hiskey and Sanchez, 1990) suggested the use of potassium cyanide for the dissolution of gold. John MacArther and Robert and William Forrest patented the use of cyanide as an extractive metallurgy process for gold dissolution in 1887 and 1889 (Habashi, 1970). This cyanide process, researched and patented in 1887 by MacArthur and the Forrest brothers, dramatically changed the way in which gold was processed. Cyanidation has since been recognised as a powerful lixiviant for dissolving precious metals from free-milling ores.

However, over the last few years there has been a decline in the number of high grade deposits. The surviving operations have to treat some form of refractory ore (Vassiliou, 1988; Hausen, 1989; Worsted, 1990) and appropriate extractive metallurgical technologies have been considered and developed (Barnes et al., 1960; Hansen and Snoeberger, 1970; Cornwell and Hisshion, 1976; Loo, 1982; Fleming and Cromberge, 1984; Shanks et al., 1985; Sibrell and Miller, 1986; Sheya and Palmer, 1989; Palmer, 1990; McClincy, 1990; Fraser et al., 1991; Lacovangelo and Zarnoch, 1991; Daocheng et al., 1992; Habashi, 1992; Kordosky et al., 1992; Wang et al., 1993; Shi and Lipkowski, 1994).

During the last ten years, as a consequence of the discovery of numerous large low grade deposits, increasing attention has been given to the extraction of gold from refractory ores. Problems still remain with the extraction of gold from such ores when using conventional cyanidation (Haque, 1987; Udupa et al., 1990). These include:

(i) the inefficiency of conventional direct cyanidation, (ii) the problem of winning precious metals from refractory ores, (iii) the toxicity and environmental safety considerations. 7

Various treatment methods involved in the extraction of gold are summarised in Figure 2.1.

Gold ore

I Free milling ore conv. leaching Refractory ore

Pretreatments and Direct Au Au leaching leaching

I Carbon-ln-leach Non.conv. Alkaline Ilxlvants cyanidatlon (CIL) (e.g. thiourea, under O2 malononitrite) pressure

I j Chemical Roasting Blo-oxidatlon oxidation

Oxidative Alkaline Acid acid leaching pressure pressure leaching leaching

Figure 2.1 Various gold extraction process for free milling and refractory ores (Haque, 1987)

2.1.1.2 Hydrometallurgical Processing of Electronic Scraps

Electronic scraps (Reddy, 1989; Sum, 1991; Hoffmann, 1992) can be defined as the waste generated from the manufacture of electronic hardware and the obsolete electronic products, particularly a variety of computer circuits such as printed circuits, memory circuits, integrated circuits and transistors. 8

Generally, the main components of electronic scrap are plastic (30%), refractory oxides (30%) and metals (40%).

The metallic components of electronic scraps include copper (20%), iron (8%), nickel (2%), tin (4%), lead (2%), aluminium (2%), and zinc (1%), while that of the precious metals are gold (0.1%), silver (0.2%), and palladium (0.005%). The major economic concerns for the recycling of electronic scraps is the recovery of precious metals. Of these metals, most attention is given to the recovery of gold, because considerable quantities of gold have been used in the electronic industry during the last three decades.

Recently, an increase in the use of printed circuits, memory chips, integrated circuits, transistors, and precious-metal edge connectors, particularly with the advent of computer scrap, may greatly exceed the value of the precious metals content. Thus the recovery of precious metals from electronic scrap has attracted the interest of many researchers (Miller, 1969; Beaglehole and Hendrickson, 1969; Ganu and Mahapatra, 1987; Reddy, 1989; Sum, 1991; Hoffmann, 1992; Guan and Han, 1993).

The major disadvantage of hydrometallurgical processing of precious metals scraps is the inability to achieve complete extraction of the precious metal content. Thus the residues from hydrometallurgical processing has been conventionally followed by smelting (in conjunction with secondary copper or nickel smelting) to complete extraction.

Despite this, the hydrometallurgical processing of scrap has been a major area of research since the mid-1980s because it has several advantages over pyrometallurgical processes, such as better environmental protection, easier 9

separation of main scrap components, and lower power consumption and recycling of chemical reagents.

The hydrometallurgical process used to treat precious metals scrap can employ either cyanide leaching or chlorine leaching. One of the most effective and selective methods for removing gold plating from a variety of substrates is cyanide leaching.

The overall reaction of cyanide leaching of gold can be presented by:

4 Au + 02+ SNaCN + 2 H20 = 4NaAu(CN2) + ANaOH (2.1)

Moreover, the dissolution of precious metals in oxidic scraps is sometimes accomplished using chlorine leaching. Regardless of its high cost and rapid reactions with virtually all metals commonly present in scrap, chlorine leaching is favoured for materials where the metallic components are already oxidised.

Many ceramic elements that are plated with gold and other precious metals can be dissolved together with the metallic components. During chlorine leaching, acidity must be kept low to avoid excessive acid consumption by reaction with the oxidic phase. A systematic block diagram for chlorination is presented in Figure 2.2. 10

Chlorine Water To Scrubber

Chlorination

Water Wash

Solids tc^ Filtration Smelter

Recycle DBC Add Scrub

Solvent Extraction Raffinate of Gold Hydrazine

Spent Liquor Gold Reduction

Gold Precipitate

To Sale

Tellurium Collection

nitration Effluent Disposal

Tellurium Tellurous Acid Recovery to Recycle

PGM Concentrate to Refining

Figure 2.2 A systematic block diagram for chlorination of oxidic scrap (Hoffmann, 1992)

2.1.1.3 Hydrometallurgical Processing of Metal and Mineral Wastes

Over die last few years, significant emphasis has been placed on the development of new technologies for the processing of slag and dust wastes (Mikhail and Webster, 1992). Minimization of hazardous waste generation 11 through secondary processing has been the main studies lor many scientists around the world, leading to expanded research interests.

Despite the constraints placed on the mineral industry by increasingly stringent environmental regulations, aqueous processing routes have a tremendous potential for treating metallurgical wastes and residues to comply with new regulations and to recover marketable products and by­ products (Finkelstein et al., 1966; Boateng and Phillips, 1978; Sato and Lawson, 1983; Bhatetal., 1987; Bhappu, 1990; Butwell, 1990; Hagnietal., 1991; Doyle, 1991).

2.1.2 Electrochemical Processes

An electrochemical reaction (Hiskey and Atluri, 1988; Felker and Bautista, 1990; Fang and Muhammed, 1992; Cheng and Iwasaki, 1992) may be described by an anodic or a cathodic reaction involving, respectively, loss or gain of electrons as illustrated in Figure 2.3.

Figure 2.3 A systematic electrochemical model for an electron transfer reaction (Hiskey and Atluri, 1988) 12

In general, a half-reaction can be expressed by

Reduced state = Oxidised state + n Electrons (2.2)

The electrode potential can be measured against a reference electrode. When using the standard hydrogen electrode (SHE) as the reference, the potential for reaction (2.2), Eh, is represented by the Nernst equation:

Vo, r>O, E. =E°+ — Ln< (2.3) ' nF ik

where E°: standard electrode potential,

R : gas constant,

T: absolute temperature, n : number of electrochemical equivalents per mole, F: Faraday constant (96,485 Coulomb/equivalent), activities of the oxidised and reduced species involving in

the electrode reaction, respectively, v , v, : stoichiometric coefficient for the oxidised and reduced °i ’j species, respectively.

A mixed potential related to the leaching rate is generated at the interface between the internally short circuited cathodic and anodic sites. 13

2.1.3 Thermodynamic Aspects of Precious Metal Dissolution

Gold is unreactive in pure water and over a very wide range of pH. Gold’s stability can fortunately be partially controlled in a potential range within the water stability region by using solutions containing suitable complexing agents (Finkelstein and Hancock, 1974; Nicol, 1980a; Hiskey and Atluri, 1988). Many processes for gold dissolution can take place if suitable oxidants such as oxygen, chlorine, hydrogen peroxide and complexants such as cyanide, chloride, and thiourea, are present in the leach solution.

2.1.3.1 Stability of Aqueous Gold Complexes

The stability of gold complexes (Andersen et al., 1964; Cornejo and Spottiswood, 1984; IUPAC, 1995) is of great importance in determining the behaviour of the metal in solution. The stability constants show that some complexing ligands form more stable complexes with Au(l) and others with Au(III). This preferred oxidation state is related to the electron configuration of the donor (or complexing) ligand. The Au(I) and Au(lII) oxidation states are considered B-type metal ions (Finkelstein and Hancock, 1974), preferring to bond to large polarizable ligands because the stability of their complexes tends to decrease with increasing electronegativity of the ligand donor atom. Both Au(I) and Au(III) tend to form stable compounds with halides in the order I" > Br' > Cf > F , while the order of electronegativities is F > Cl > Br > I . A periodic representation of the elements with regard to the properties of the elements which effect their formation of complexes with gold is as follows: 14

C N O F Electronegativity P S Cl increasing As Se Br Sb Te I Stability increasing

This indicates that low valency metal ions, Au(I), preferentially form complexes with the less electronegative, or soft, donor ligands, such as cyanide, iodide, thiourea, thiocyanate and thiosulphate. The higher valency, Au(III), forms more stable complexes with the more electronegative, or hard, donor ligands, such as chloride and bromide.

This is in accordance with the order of stability, Se = C(NH2)2 > S = C(NH2)2 > O = C(NH2)2 or CN > NH3 > H20. The argument can be used to account for the observation that gold tellurides and aurostibnite (AuSb2) are stable in nature. The above general rule is useful in predicting the stability of gold complexes.

2.1.3.2 Eh (potential)-pH Relationship

The most suitable thermodynamic method for studying hydrometallurgical processes of gold dissolution with different complexing ligands is by constructing corresponding Eh - pH diagrams. A number of investigators describe the development and use of these diagrams in hydrometallurgical systems (Garrels and Christ, 1965; Peters, 1976). E,rpH diagrams are useful for interpreting many hydrometallurgical reactions in practice (Latimer, 1952; Sillen and Martell, 1971; Hancock et al., 1974). 15

2.1.3.2.1 Gold-Water System

The pe - pH diagram (Pourbaix, 1966), showing stability relations among the hydroxo species of Au(I) and Au(III), has been constructed at 25°C as shown in Figure 2.4. The diagram shows that the reduction potential for AuJ+/Au is greater than the upper stability limit of water. This means that gold is extremely stable in pure water and over a very wide pH range.

Au|OH)J

Figure 2.4 pe-pH diagram for the Au-H20 system, showing stability relations among the hydroxo species of Au(I) and Au(III); [Au] = 1 pg/1 (ppq) to 1 mg/1 (ppm) (Vlassopoulos and Wood, 1990)

Moreover, aurous and auric ions are thermodynamically unstable in water because the value of the standard reduction potential for reaction (2.6) is lower than those for reactions (2.4) and (2.5). This indicates that oxygen in 16

water is not able to oxidize Au to either Au+ or Au3 + , whereas these ionic species (Auf/Au3t) will be readily reduced to metallic gold and water spontaneously oxidised to oxygen. It can be shown that the auric ion is also more stable in the presence of a strong oxidant than the aurous ion.

An" +e = Au E° = 1.73 V (2.4)

Ati}i + 3 e = An E° = 1.50 V (2.5) + 4 H+ +4e~ = 2H20 E° = 1.23 V (2.6)

However, in the presence of oxidizing agents such as oxygen, hydrogen peroxide and chlorine, gold can be dissolved to form gold complexes with a number of other ligands such as cyanide, chloride and iodide in aqueous solution. The dissolved gold is also stable only if the equilibrium potential for the gold complex in the presence of ligands remains above the required reduction potential of the system.

2.1.3.2.2 Gold-Ligand-Water System

Gold in the presence of a complexing ligand and an oxidizing agent can be readily dissolved in an aqueous solution. In general, the half-cell reaction associated with gold oxidation and complexation with the complexing ligand,

X, may be expressed by:

Au + 2X = AuX2+ -t- e (2.7)

and E = E°„, + 2.303(log[AuA',+]/[X|2) (2.8) 17

E°01, related to its free energy change (AG° = - nFE°), is the standard electrode potential for the oxidation of the couple Au(0)/Au(I). The accompanying cathodic reaction is the reduction of a suitable oxidant:

Nz+ +e~ -> N(Z~l)+ (2.9)

Oxidation of Au(0) to Au(III) in the presence of suitable complexing ligands occurs as follows:

Au + 4X —> AuX43+ + 3e" (2.10)

The standard electrode potential E°03 for the couple Au(0)/Au(III) is related

•i i_ to the stability constant for the formation of the complex (AilT4 ).

It is possible to oxidize Au(I) to Au(III) in the presence of complexing ligands, X, as represented by:

AuA',+ + 2X -> AuX43+ + 2e’ (2.11)

The standard potential, E°0i3, where E°1>3 is the standard potential for reaction (2.11) with complexing ligands is expressed as:

3E°0i3 = 2E°,3 +E°0, (2.12)

It is possible to oxidize Au(0) to either the (Au2f2+) or (Au2f43+) complex in the presence of the appropriate ligand. The most stable gold complex formed will depend on the strength of the oxidizing agent used and on the relative 18 magnitudes of E°u and E°0fl. Standard potentials for E°01, E°03, and E°lt3 witli several ligands are summarised in Table 2.1.

Table 2.1 Standard potentials for gold complexes in aqueous solution at

25°C (Hiskey and Atluri, 1988; Fang and Muhammed, 1992)

Standard potentials, mV vs SHE

c° c° Ligand I j 0,1 ^ 0,3 E°L 1,3 H20 1700 1500 140

Cl 1150 1000 930 Br 960 860 810

SCN’ 670 640 620

nh3 563 325 206

I 560 570 570

Thiourea 380 - -

Selenourea 200 - -

SA2' 150 - -

s2- -460 - -

CN' -610 - -

Equation (2.7) of the half-cell reaction for gold represents the formation of the Au(I) complexes for cyanide and iodide systems and equation (2.9) represents the formation of Au(III) complexes for chloride and bromide systems. Furthermore, the magnitude of the complexation is illustrated more clearly by the Eh-pH diagram (Osseo-Asare et al., 1984; Hiskey and Atluri,

1988; Marsden and House, 1992) for the Au-CN-H20, Au-Cl-H20, Au-I -

H20, Au-Br"-H20, and Au-(NH2)2CS-H20 systems, respectively. These diagrams can predict the feasibility of reactions and may also play important roles in the final selection of oxidant and ligand. 19

2.1.3.3 Modelling of Gold Dissolution Process

The conditions for the extraction of valuable metals depend on the metal’s composition, its dissolution behaviour and the response of the associated minerals to the leach solution (Peters, 1991). The recovery of gold from low-grade (less than 3g Au/metric tonne) ores containing a variety of impurities is typically accomplished by hydrometallurgical processes used for the extraction of gold. The gold-bearing and gangue minerals can be decomposed in a ligand solution. Metal ligand complexes and other dissolved compounds may not only affect the leaching step to produce a gold solution as an intermediate product but also other process solutions for the subsequent recovery of gold.

A thorough analysis of the aqueous chemistry and thermodynamic constraints affecting the speciation of aqueous ligand solutions need to be carried out in order to predict the feasibility of dissolution reactions and to determine the soluble species or compounds. The thermodynamic study of hydro- metallurgical processes involving numerous species and compounds can be very complex and generally require a computer programme for equilibrium thermodynamic calculations containing complex chemical equilibria, process modelling and phase diagram calculations (Xue and Osseo-Asare, 1985; Barbosa and Monhemius, 1988; Li, 1993; Gaspar et al., 1994).

More than 25 software packages were reviewed at the third conference and workshop on Computer Software (Morris et al., 1990) for Chemical and Extractive Metallurgy Calculations held in June, 1989. These packages ranged from powerful multipurpose programs with large databases to simple 20 data manipulation programs, some of which are still under development. One of the most comprehensive of the integrated multipurpose programs is the CSIRO (Division of Mineral Products) Thermodata Package which was developed in Australia by the CSIRO.

2.1.4 Kinetic Considerations of Precious Metal Dissolution

For the commercial application of gold leaching process, a knowledge of the mechanism and rate-controlling factors (Wagner, 1938; Kudryk and Kellogg, 1954; Vermilyea, 1966) in gold dissolution should be understood. In general,

(1) If the rate-determining process is controlled by a slow chemical reaction,

• an increased temperature should have a greatly accelerating effect, • agitation in the solution should have a negligible effect on the dissolution rate, and • increased concentration of reactants should cause acceleration of the dissolution rate.

(2) If the rate-determining process is controlled by the diffusion- controlled reaction of one or the other of the reactants,

• increased agitation should increase the dissolution rate, • increased temperature is expected to slightly increase the dissolution rate, 21

• increased concentration of the reactant which is diffusion limited will increase the rate, and • the concentration of other reactants should have no effect on the rate.

Sequential stages for the dissolution of gold in alkaline cyanide solution are as follows (Cornejo and Spottiswood, 1984):

• Absorption of oxygen into the solution, • Transportation of dissolved oxygen and cyanide from the bulk solution to the gold-solution interface, • Adsorption of reactants onto the gold surface, • Electron transfer, • Desorption of the gold cyanide complexes and other products from the surface, and • Transportation of the desorbed products into the bulk solution.

The rate of the dissolution of gold is controlled by the diffusion rate of either dissolved oxygen or cyanide ion species through the diffusion boundary layer (Kudryk and Kellogg, 1954). It has been suggested that at low cyanide concentrations, the dissolution rate is controlled by the rate of the cyanide diffusion, while, at higher cyanide concentrations (in the case of a molar ratio of cyanide to oxygen greater than 10), the rate is dependent on the rate of diffusion of oxygen. However, at high oxygen levels and strong agitation, the gold surface may become passivated, and its rate of dissolution can approach a constant value, or even decrease with increasing level of oxygen (Cathro, 1963; Cathro and Koch, 1964b). 22

In view of the above, it can be concluded that, with a knowledge of the mechanisms and rate-controlling factors of gold dissolution in the cyanide system, the application of a commercial process for the extraction of gold from ores should be readily achieved.

2.2. Cyanidation

2.2.1 Gold Dissolution in Cyanide Solutions

The cyanide leaching process (Cornejo and Spottiswood, 1984; Espiell et al., 1986) carried out in alkaline cyanide solutions involves agitating the ground ore for periods of time that vary from 12 to 96 hrs. The concentration of cyanide solution commonly required varies from 0.5 to 2.5 g/1 KCN or NaCN. To keep the solution pH between 10.5 and 11.5, lime is added to cyanide solution. The role of lime is to prevent the loss of cyanide by causing the precipitation of undesirable metal ions and aiding in the flocculation and settling of ore slimes. A large excess of lime, however, has been found to increase the cyanide consumption and to decrease the rate of dissolution of gold in comparison to when less lime is used. The rate of dissolution depends on several factors such as particle size, degree of liberation, alkalinity of the solution, cyanide concentration, aeration, and gold content.

Cyanidation (Habashi, 1966) has been widely adopted due to its many advantages. Many investigations (Miller et al., 1990; Kenna et al., 1990; Hiskey and Sanchez, 1990) have been conducted to describe the mechanisms and kinetics of the reaction for cyanidation by using various investigative techniques. According to several recent papers, researchers are still in 23 dispute in several aspects of the intriguing chemistry of the cyanidation reaction.

Eisner (1846) has suggested that the gold dissolution in cyanide solutions was dependant on concentration of CN and partial pressure of 02. The overall reaction with a free energy change of -376.4 kJ/mole is expressed as:

4 Au + 8CN~ +02 +2H20 —» 4 An (CN)2 +4OH~ (2.13)

Furthermore, Bodlander (1896) indicated that the dissolution of gold, with H202 as an intermediate product, may occur in the following two steps:

2Au+4CN~ +02 + 2H20->2Au(CN)~ + 20H~ +H202 (2.14)

and

2 Au +4CN~ +H202 -> 2 Au(CN)~ +2OH' (2.15)

With regard to the commercial extraction of gold, the dissolution of gold in the presence of secondary reactions with soluble sulfides has also been discussed by several researchers (Baum, 1990; Riveros, 1990; Lorenzen and van Deventer, 1992). Aspects related to the effect of sulfides during gold cyanide will be discussed in a later chapter.

2.2.2 Electrochemical Processing of Gold in Cyanide

A considerable research effort has been carried out to determine the mechanisms involved in the anodic dissolution of gold in alkaline cyanide 24 solutions. Many authors (Shutt and Walton, 1932, 1933 and 1934; Cathro and Koch, 1964c; Thacker and Hoare, 1964; Bazan and Arvia, 1965; MacArthur, 1972; McIntyre and Peck, 1976; Nicol, 1980b) agreed that the oxidation of gold in aqueous solution, where no other complex-forming species are present, generally results in passivation by the formation of oxide films on the gold surface. However, little has been published specifically relating to the electrochemistry of gold dissolution in alkaline solutions. The most stable complexes of Au+ and Au3+ are those which are formed with cyanide ions. The cyanide complex of Au+ is more stable than that of Au3 + , indicating by a lower reduction potential by as much as 0.5V.

Kirk et al. (1978 and 1980) have revealed using a detailed coulometric method that the anodic dissolution reaction in cyanide solutions is a single­ electron transfer reaction throughout the potential range -0.65 to 0.55 V:

Au + 2CN~ = Au(CN)2~ +

At high potentials, there is evidence of the formation of Au(III) complexes (Eisenmann, 1978; Pan and Wan, 1979). The current-potential curve for the dissolution of gold in an alkaline cyanide solution occurs with three peaks at potentials of about -0.4 V, 0.3 V, and 0.6 V, with respect to the standard hydrogen electrode (SHE).

At potentials below -0.4 V, it is generally accepted that the dissolution of gold occurs via the mechanism of reactions (2.17) and (2.18) (Cathro and Koch, 1964b; Eisenmann, 1978). This is because the potential does not exceed 0.0 V during the cyanidation of gold. An adsorbed layer of AuCN was considered to cause the passivation of gold, according to the following reactions: 25

An + CN = AuCN ads + e (2.17)

AuCNads + CN~ - Au(CN)2~ (2.18)

The second passivation peak at a potential of 0.3 V is considered to occur via the formation of a surface layer of AuOH corresponding to the reaction (Kirk etal., 1978):

Au + WH- = Au(OH)3 +3e (2.19)

The third passivating film occurs at 0.6 V. It has been generally accepted that the formation of Au203 (Cathro and Koch, 1964c) causes the passivation of gold.

The presence of impurities such as lead, mercury, bismuth, and thallium can result in a significant influence on both the anodic and cathodic characteristics of the Au/Au(CN)2 couple. Several investigators (Cathro, 1964a; McIntyre and Peck, 1976) found that the presence of small amounts of lead, mercury, bismuth, and thallium can reduce the passivity of gold and hence increase the dissolution rate. Cathro (1964a) and Nicol (1980b) confirmed that heavy-metal ions, such as lead and thallium, can reduce the passivation effect of the dissolution of gold in cyanide solution. The works of McIntyre and Peck (1976) and Dinan and Cheh (1992) showed that these effects are due to the strong adsorption of the foreign metal ions. It is further noted that the presence of small amounts of sulfide ions has no significant effect on either the anodic or cathodic current (Stanley, 1987). 26

Cornejo and Spottiswood (1984) summarised that the retarding effects on the dissolution of gold in a cyanide solution are as follows:

(1) Consumption of free cyanide from solution due to: • Formation of complex cyanides by metal ions acting as "cyanicides", for example, Fe2+, Cu+, Zn2+, Ni2+, • Formation of inert thiocyanates by the reaction of sulfur with free cyanide, • Decomposition loss in the absence of protective alkali such as calcium ions, • Adsorption on gangue minerals.

(2) Formation of a passive film due to the adsorption of sulphur anions on the surface of the gold particles.

(3) Consumption of oxygen in the solution due to side reaction with dissolved sulfide ions.

2.2.3 Thermodynamic Considerations

A clear knowledge of the reactions, their products, and the factors that influence the rates at which they proceed is very important from a practical and economic point of view. As stated earlier, in the presence of cyanide as a complexing ligand and 02 as an oxidizing agent, gold is readily dissolved. The formation or the stability constant of the Au(CN)2" complex ion in an aqueous solution is:

Au+ +2CN- = Au(CN)~, J32= 2.8xlO39 (2.20) 27

The Eh-pH diagram (Hiskey and Atluri, 1988) for the Au-CN -H20 system at 25°C and [CN ] = 10 JM is shown in Figure 2.5. From the constructed Eh- pH diagram, in the presence of cyanide, a strong complexing ligand, it can be seen that there are domains of stability of Au(CN)2 within the region of the stability of water.

The stability region (Xue and Osseo-Asare, 1985) of the aurocyanide complex increases with increasing cyanide concentration and decreases with increasing dissolved metal concentration. At a pH value of about 10, metallic gold is readily oxidised to the stable Au(CN)2 complex at a potential of about -0.52 V (SHE), as seen from Figure 2.5, at a cyanide concentration of 10"6 to 1 M.

Au 0

Au(OH) (s)

.(c)-

Eh i o

Au(CN):

(a) 10 M CCN] (b) 10 -O 5 (c) 1.0 [Au] =10' M

pH Figure 2.5 Eh-pH diagram for the Au-CN~-H20 system at 25°C; [Au] = 10'3M, [CN ] = 106 to 1 M (Xue and Osseo-Asare, 1985) 28

The practical application of Eh-pH diagrams may be limited by kinetic considerations because thermodynamic data is concerned only with the change in state between the start and end of a reaction. Therefore, the factors (Thurgood et al., 1981) that influence the rate of a chemical reaction have to be considered.

2.2.4 Kinetic Aspects

The dissolution of gold in an alkaline cyanide solution occurs through a heterogeneous reaction (von Hahn and Ingraham, 1968) that takes place between a solid surface and a fluid. The transport of reactants from the bulk solution to the gold-solution interface may be considered as an important factor to determine the overall kinetics. The rate of reaction depends on the chemical steps or on the diffusion processes of the reactants to the gold surface and the area of the reaction interface, during the reaction.

Generally, the dissolution of gold in cyanide solutions can be controlled by the diffusion of both the dissolved oxygen and the cyanide ion species, through the Nernst boundary layer. Several investigators (Kudryk and Kellogg, 1954; Cathro, 1963; Cathro and Koch, 1964b) have confirmed that the kinetics of the dissolution of gold in cyanide solutions become limited by the mass transport of cyanide ions to the gold surface, as described below:

• The rate of dissolution is determined by the nature of the gold surface • The rate of gold dissolution is controlled by the level of agitation 29

• The elevation of the temperature raises the rate of dissolution, but not as much at higher cyanide concentrations, because of the decreased solubility of oxygen at higher temperature.

On the basis of the studies mentioned above, the anodic and cathodic reaction of gold in cyanide solutions, can be described by Fick's law for the steady- state diffusion of oxygen and cyanide to the gold surface.

^l = £^A,{(P0i)b-(P0i)J (2.21)

d[C'?■ ~ = ^-A2{[CN-]„-[CN-]s} (2.22) at o where d{P0 )ldt, d[CN~]/dt: rates of diffusion of the 02 and CN ions,

respectively, Da,, Dcn. : diffusion coefficients of 02 and CN", respectively, {PQ )b , (PGj ),: partial pressures of 02 in the bulk solution and at the

surface, respectively, [CN~]b , [CN ], : concentrations of tlie CN ion in the bulk and at the

surface, respectively,

Ai, A2i surface areas at which the cathodic and those at which anodic reactions take place, respectively,

8: boundary layer thickness,

t: time.

The stoichiometry of the anodic reaction in cyanide solutions is :

Au + 2CN~ = Au(CN)l +e~ (2.23) 30

For the case of diffusion control, the rate of anodic dissolution of gold , R(l, can be reduced to:

Ra=kDcN_[CN-}b/2 (2.24)

where A2 k: a constant, — o diffusion coefficient of the cyanide ion DCN- • [av-]6: concentration of cyanide factor 2: 2 moles of cyanide ion dissolved for one mole of gold.

The stoichiometry of cathodic reaction can be represented by:

02 + 4/T + 4e~ = 2 H20 (2.25)

The rate of the cathodic reduction of gold, Rc, can be reduced to:

Rc (2.26)

where k: a constant, —Ar o diffusion coefficient of 02 concentration of 02 in the bulk solution

factor 4: 1 mole of oxygen dissolved for four moles of gold. 31

At steady state, the rate of the anodic reaction must equal the rate of the cathodic reaction.

Thus,

(2.27)

For a system of gold in a cyanide solution, it can be concluded that the rate of dissolution of gold is controlled either by the rate of diffusion of oxygen or cyanide to the gold surface. The ratio of cyanide to oxygen is also very important in the diffusion-controlled process determined by the rate of diffusion of dissolved oxygen or cyanide to the gold surface. In the presence of small amounts of heavy-metal impurities, it is generally accepted that the formation of a film takes place, which in turn, passivates the gold surface, or alternatively leads to its activation.

From a practical point of view, although the aspects discussed above are of great importance, there are a lot of additional difficulties that arise directly from industrial design problems (Urban et al., 1973; Wan and Miller, 1986;

Stanley, 1990). One reason is that almost always all laboratory tests have been under well-controlled conditions and the samples used were pure.

Moreover, the complexity of low-grade gold ores and the refractory nature of some ores create significant limitations to the direct cyanidation process.

Thus, pretreatment of these ores before cyanidation, by such methods as roasting, leaching with strong acids, or oxidation by oxidants in solution or gaseous oxidants under pressure has been suggested. 32

2.3 Chlorination

Historically, chlorination (referred to as "the Plattner process") (Hiskey and Atluri, 1988), an early hydrometallurgical approach, was adopted extensively to extract gold from oxidised ores in Australia and North America during the second half of the nineteenth century. However, since the advent of the commercial success of the cyanide process, chlorination has been rapidly replaced by the cyanide process. The chlorination process has consequently been limited to refining processes and analytical applications, where the speed of dissolution and capacity for simultaneous sulfide destruction are more important than reagent consumption.

In the early 1970's, a chlorination process (Scheiner et al., 1971; Guay and Peterson, 1973) was developed to be used as a pre-treatment stage for carbonaceous gold ores prior to cyanidation. The main advantages (Hiskey and Atluri, 1988) of alkaline pretreatment of refractory carbonaceous gold ores are:

• The elimination of the pre-robbing behaviour of the carbonaceous materials, • The liberation of gold in a combined chemical form with carbonaceous materials, • The decomposition of associated sulfide minerals containing locked finely dispersed gold which might not be otherwise exposed during cyanidation. 33

2.3.1 Chlorine and Hypochlorite Chemistry

2.3.1.1 General

Chlorine is the 11th most abundant element in the earth's lithosphere. Although chlorine is relatively abundant and widely dispersed, a high percentage is present in the form of chlorides of alkali metals. Chlorine is the most industrially important element of the halogens which represent a major part of Group VIIA of the periodic table. Chlorine has certain chemical characteristics which are similar to the other halogen members such as fluorine, bromine, and iodine.

Chlorine is highly reactive and forms compounds with virtually all other elements. Chlorine dissolves in water and hydrolyses to a limited extent when dissolved in aqueous solution, thus the solubility is dependent on the pH of the solution. Moist chlorine reacts readily and corrosively with almost all common metals except platinum, tantalum, and titanium (McKetta, 1993).

Chlorine reacts with hydrated lime or caustic soda to form calcium hypochlorite or sodium hypochlorite used for bleaching and sterilizing, because of their oxidizing action, in laundries, paper and textile mills, swimming pools, and water treatment plants.

2.3.1.2 Aqueous Chemistry

Chlorine both in its elemental and hypochlorite form can be rapidly dissolved in water. Dissolved chlorine produces aqueous chlorine, hydrochloric acid, hypochlorous acid, and hypochlorite ion. The relative proportion of the two 34 forms of available free chlorine (hypochlorous acids, hypochlorite ions or mixture of both to the total amount of dissolved chlorine) is dependent on the solution pH and temperature.

Chlorine evolved during anodic reactions dissolves in water, producing various species, according to the following equation (Bayrakceken et al., 1990),

CL +H,0= HOCl + H+ + Cl~ (2.28)

The equilibrium constant for the above reaction, obtained by Bayrakceken et al. (1990) from the free energy values of the species at 25°C, is:

[HOCl][H+}[Cr] = 4.1x10 4 (2.29)

Aqueous chlorine is a predominant species at a pH of 2 or below, as presented in the distribution diagram for chlorine species (Figure 2.6). At a higher pH value (pH range from 3 to 6), hydrochloric acid (HC1) as a strong acid, which completely dissociates in dilute aqueous solution is generated, together with hypochlorous acid (HOC1) as a weak acid. This means that moist chlorine can corrode metals which are usually unaffected by gas or liquid chlorine.

Moreover, the hypochlorous acid (HOC1) formed in reaction (2.28) dissociates, according to the following reaction (2.30) and, at pH values above 6, hypochlorite ion (OCf) predominates in aqueous solution.

HOCl = H+ + ocr (2.30) 35

PH

Figure 2.6 Distribution diagram for chlorine species at 25°C; [Cl ] = 10 3 M (Marsden and House, 1992)

The equilibrium constant at 25°C is:

r [u'Wocr] (2.31) K>- ihoch 23x10

The half-cell reactions (Pourbaix, 1966; Wu, 1987; Marsden and House, 1992) for hypochlorous acid (HOC1), chlorine, and hypochlorite (OC1) are as follows:

2HOCI + 2H+ + 2e~ = C12+H20 E° =+1.64V (2.32)

Cl2 (aq) + 2e = 2 Cl E° =+1.36V (2.33) 36

0Cr+H20 + 2e~ =Cr+20H~ E° = +1.07 V (2.34)

The OC1 , HOCl, and Cl2(aq) species are strong oxidizing agents, with hypochlorous acid (HOCl) as a preferred species for dissolving carbonaceous materials and sulfides.

Above 50°C, chlorine dissolves to produce the strongly oxidizing chlorate species (C103). This reaction virtually only varies with pH and brine strength, according to the following equation:

3 C/2 + 3 H20 = CIO- +5 Cl~ + 6 H+ (2.35) and 2 CIO' +12 H+ + \0e~ = Cl2 + 6 H20 (2.36)

Reaction (2.35) implies that the lower the pH and the higher the brine strength, the lesser is the amount of chlorate available in the solution. Alternatively, the addition of inorganic hypochlorite salts to a solution results in its decomposition, generating the strongly oxidizing hypochlorite species as follows:

NaOCl = Na+ + OCC (2.37)

Ca{OCl)2 = Ca2+ +2 OCC (2.38)

The hypochlorite species, as stated above, further reacts with water to form available free chlorine, with a corresponding variation in pH. The overall profile for both the cathodic and anodic reactions on the electrode surface is shown in Figure 2.7. 37

Cathodic

OCl~ + H20 = HOCl + OH oc/-

HOCl + e- -» HOCr

Hocr =oh+ci~ cr

OH+ e~ —» OH^ A * OH'

Anodic Aqueous Phase Electrons

OH~ ->OH + e OH

OH + OH" = 0 + H20 + e 20 = a «------►.

Nernst boundary layer

Figure 2.7 Schematic representation for both the cathodic and anodic reactions on the electrode surface (Wu, 1987)

2.3.2 Chloride Leaching Process

The chlorination process (Piret et al., 1978) was replaced by cyanidation for three major reasons, namely the high reactivity of base metal sulphides, the inability to form stable gold complexes, and the highly corrosive nature of acidic chloride solutions. Recently, chlorination (Cui et al., 1990; Mason et al., 1990) has been re-introduced for the treatment of complex metal sulfide ores and concentrates, because it has several advantages over the cyanide process. Firstly, the dissolution rate of gold during chlorination is much faster than with cyanidation. Secondly, carbonaceous ores are more 38 amenable to chlorination. Finally, the reagent cost (1.5 $/t ore) of chlorine for gold leaching is lower than that (2.0 $/t ore) of cyanide (Ximing et al., 1992).

Preoxidation of ores at Carlin Mines, Nevada, with the proposed flash chlorination (Marsden and House, 1992), can be optimised to maximise gold extraction during the chlorination of refractory gold sulfides and carbonaceous gold ores. Furthermore, heap leaching with hypochlorite (Barratt and McElroy, 1990; Ahmadiantehrani et al., 1991 and 1992) and in- situ leaching with chlorine (Fagan, 1991) as economical processes for the recovery of gold from low grade ores have been proposed and the feasibility of employing hypochlorite pretreatment (Ahmadiantehrani et al., 1991) for the conventional cyanide heap leaching process to treat gold ores has been investigated. Gold recoveries of greater than 85% were achieved utilising hypochlorite pretreatment, but ores with more than 0.5% sulfide content are uneconomical to treat in a chloride system due to the high chlorine consumption.

However, in Australia, most gold processing regions are arid and mines arc often forced to use ground water brines for gold processing (La Brooy et al., 1994). Processes using chlorine pretreatment has a cost advantage over the cyanide process. This is because of the higher cost of cyanide in comparison to the electrolytic production of chlorine from ground water brines when natural saline waters are available (Sandberg and Huiatt, 1986).

2.3.3 Electrochemical Processing of Gold in Chloride Systems

Many investigations (Pearson and Buder, 1938; Just and Landsberg, 1964; Heumann and Panesar, 1966; Gaur and Schmid, 1970; Gallego et al., 1975; 39

Schalch and Nicol, 1978a; Nicol, 1980a and 1981b; Schalch et al., 1978b; Frankenthal and Siconolfi, 1982) have been carried out to determine the fundamental electrochemistry of gold dissolution when using chloride as the complexing lixiviant. Gallego et al. (1975) and Nicol (1980a) agree that the anodic dissolution of gold in an aqueous chloride solution produces Au(I) and Au(III) complexes according to the following reactions:

An+ 20 = AuCl2+e E°=1.113V (2.39)

Au + ACr = AuCll E°=0.994V (2.40)

Polarization curves for electrochemical oxidation and reduction have been reported by several investigators, but their results are not generally in agreement. As summarised by Nicol (1980a), the electrochemical dissolution of gold in acidic chloride solutions occurs at potentials lower than 1.2 V for both Au(I) and Au(III) species. At higher potentials in the range of 1.2 to 1.4 V, the reactions of gold species at limiting current is controlled by the diffusion of chloride ions to the gold surface. At a potential of about 1.5V, passivation is generally observed due to the formation of an oxide layer. Both oxygen and chlorine evolution take place at potentials higher than 1.6V.

Using cyclic voltammograms Gaur and Schmid (1970) have extensively studied the effect of chloride ions on the passivity of gold in dilute chloride solutions containing 0.1M HC104. According to their work, a current peak occurs at about 1.2V corresponding to the oxidation of Au to Au(III). Its height was found to be dependent on the chloride concentration and the stirring rate. At approximately 1.2V, passivation was observed due to the formation of "adsorbed oxygen" on the electrode surface (Bonewitz, 1969; Burke et al., 1990 and 1992), which prevented further dissolution of gold as 40

Au(III). The reaction (Takamura et al., 1971) of the adsorption of Cl , Br’, and I anions onto gold, is controlled by the diffusion of halide ions to and from the electrode. The negative potential of adsorption follows the order I" > Br > Cl . This means that the extent of adsorption is greatest with I" ions.

A mechanism (Lingane, 1962; Gallego et al., 1975) describing the electro­ dissolution and electro-deposition of gold in concentrated HC1 solutions, through the use of cyclic voltammograms has been proposed. For the gold electro-dissolution, the proposed reactions are as follows:

Au+cr = Aucrd (2.41)

AuCl~d +cr - (AuCl~)ad + e~ (2.42)

3(AuCl')ad = AuCl4 +2Au + 2Cr (2.43) with reaction (2.42) being the rate determining step.

From their investigation, the rate for the electro-dissolution of gold was found to depend on the disproportionation of adsorbed Au(I). The electro- deposition involving either the AuC12 ion or the AuC14 ion was found to be dependent on the chloride concentration. Contrary to the previous investigators (Nicol, 1981a; Tran et al., 1992), they reported that at low concentration of chloride, Au(I) species occurred as the major dissolution product.

Frankenthal and Siconolfi (1982) concluded that at potentials below 0.8 V, gold in chloride solutions dissolves as Au(I) species, while at higher potentials, above 1.1 V, gold forms Au(III) species. They also explained the results of the steady-state E-/ curves of gold, by five potential regions: active, 41 prepassive, prepassive-to-passive transition, passive and transpassive. In the active region, gold dissolved to Au(I) and the rate-determining step was the activated desorption of (AuCl2’)ads. In the prepassive region, a film of an oxide or hydroxide of gold was observed. In the prepassive-to-passive transition region, an oxide or adsorbed oxygen fdm was formed. In the passive and transpassive region, the Cl2 and 02 evolution took place. In another work on the dissolution of gold in 1 M H2S04 containing dilute chloride solutions reported by Frankelthal and Thompson (1976), two mechanisms of oxygen evolution were suggested: at low potentials oxygen is evolved on the metal surface, while at high potentials the process occurs on the Au(OH)3 film (Ferro et al., 1974a).

Nicol and his co-workers (Schalch and Nicol, 1978a; Nicol, 1980a; Schalch et al., 1978b; Nicol, 1981a) reviewed previous investigations and reported a detailed study of the dissolution and passivation of gold in chloride solutions. Nicol (1980a) noted that the dissolved Au(III) species concentration increased with potential, and decreased with the increase in chloride concentration and the decrease in the rotating speed, at all potentials. The Au(I) species occurred as an intermediate species which could be further oxidised to Au(III) or could diffuse away from the electrode into the bulk solution. Contrary to the previous investigator (Gallego et al., 1975), the author showed that especially at a high chloride concentration (>0.1 M Cl), the dissolution rate for Au(I) increased with the increase in the chloride concentration.

Moreover, electrochemical studies have also been carried out by Previous investigators (Ferro et al., 1974b and 1975; Podesta et al., 1979; Nicol, 1980a). These studies confirm that the general shape of the curves proposed for gold oxidation, reduction were similar to those observed by previous 42 investigators. The following mechanism is proposed as a means of presenting these results for the dissolution of gold in aqueous chloride solutions. The dissolution of gold can be roughly divided into two stages. As the first stage, an intermediate Au(I) chloride forms on the gold surface as follows :

2Au + 2Cr = 2AuCl + 2e (2.44)

As the second stage, AuC12 species, which could be either oxidised further to Au(III) or diffuse into the solution, can be formed as a secondary intermediate with respect to the oxidizing potential of the solution, possibly according to:

AuCl~ +2 Cl~ = AuCl; +2e~ (2.45)

At potentials of above approximately 1.4V vs SHE, passivation was observed due to the formation of an oxide layer on the gold surface. Oxygen and chlorine evolution starts at potentials higher than 1.6V vs SHE.

On the other hand, the reduction (Armstrong and Butler, 1934; Evans and Lingane, 1964; Schalch and Nicol, 1978a) of Au(III) chloride or Au(I) chloride to the metallic gold can be expected as follows:

AuCL +e = Au + 2 Cl (2.46)

AuCl4 +3e = Au + 4C/ (2.47) 43

On the assumption of the formation of Au(I), as an intermediate species of the reduction of Au(III) chloride to gold metal, the following reaction has been derived by Nicol (1981a):

AuCl; + 2e~ = AuCl~ + 2 Cl~ (2.48)

Furthermore, gold in the presence of impurities can be readily precipitated from the resulting solution either by reducing with hydrogen sulphide, sulphur dioxide, ferrous sulphate, or by carbon adsorption.

2.3.4 Thermodynamic Considerations

The dissolution of gold in aqueous chloride solutions produces both the aurous and auric complexes, according to the reactions:

AuCl-+e-=Au + 2Cr E°= 1.113V (2.49)

AuCl; +3e' = Au + 4Cr E° = 0.994V (2.50)

Reactions (2.49 and 2.50) show standard electrode potentials at 25°C. These indicate that the Au(III) chloride complex is more stable than the Au(I) chloride species by approximately 0.12V. However, these gold chloride complexes are not as stable as the Au(I) cyanide complex, nevertheless one advantage of this is that they are more readily reduced to metallic gold. The dissolution of gold for the production of Au(I) and Au(III) chloride complexes will only occur above about 1.2V vs SHE. Figure 2.8 shows a potential-pH equilibrium diagram for the Au-Cl -H20 system at 25°C and [Cl ] = 10 2M. 44

Au(OH) AuCL HAuO

-0.5

Figure 2.8 Eh-pH diagram for the Au-Cl -H20 system at 25°C; [Au] = 10‘5 M and [Cl ] = 10 2 M (Hiskey and Atluri, 1988)

With increasing chloride concentration, the Au/AuC14 boundary falls to a lower potential within the region of stability of water. When the chloride level is extremely high, such as in the case of sea water (0.004 to 0.008 ppb Au), gold can be dissolved at about a neutral pH. Tran et al. (1992) showed that AuC12 species is generally the predominant species under these conditions. Moreover, if a ligand stronger than chloride is present, the stability zone of the oxidised species would be expanded.

Chloride can be readily oxidised by a strong oxidant, such as chlorine (E° = 1.395 V), hypochlorite (E° = 1.715 V), hydrogen peroxide (E° = 1.763 V), chlorate ions or ozone. Of these, chlorine is considered as the most suitable oxidant for the gold-chloride system because it also supplies chloride ions for 45 gold dissolution. Furthermore, chlorine dissolves to form hydrochloric acid and strongly oxidizing hypochlorous species. Nitric acid in aqua regia (a mixture of 33% HN03 and 66% HC1), as a powerful oxidant can also oxidize gold in the chloride system (Marsden and House, 1992).

2.3.5 Kinetic Aspects

The dissolution of gold in chloride solutions can be controlled by the diffusion of both the dissolved chlorine and chloride ion species through the Nernst boundary layer. Several investigators (Nicol, 1980a; Warren and Mounsey, 1983; Wu, 1987) have confirmed that the kinetics of dissolution of gold in chloride solutions is controlled by diffusion of both chloride and hypochlorite ions to the gold surface as described below:

• the rate of dissolution is proportional to the chloride ion concentration, • the mass transport of chloride ions to the gold surface is the rate­ determining step, • chloride ion mass transport rates are enhanced by high hypochlorite - chloride concentrations and by increased temperatures.

On the basis of the above studies, the anodic and cathodic reaction of gold in chloride solutions can be described by Fick's law for the steady-state diffusion of hypochlorite and chloride to the gold surface. The reaction may be written as:

d([ocr\) [ocr AMOCl-])b-{[OCl]),} (2.51) dt 8 46

d[Cl~] D, [cr] A2{[cr]b-[cru (2.52) dt 5 where

d{[OCr])!dt, d[Crydt: rates of diffusion of the OC1 and Cl ions,

respectively D^cr]» 2) : diffusion coefficients of OCf and Cf, respectively

([OCr\)b , ([OC/'])s: concentrations of OCf in the bulk solution and at

the surface, respectively

[Cr]b , [Cr]s : concentrations of the Cl ion in the bulk and at the

surface, respectively

The stoichiometry of the anodic reaction of gold in chloride solutions is:

Au + 4CC = AuCl~ + 3e~ (2.53)

For the case of diffusion-limiting control, the rate of the anodic process of gold, Ra, reduces to:

R‘=kD

concentration of chloride

factor 4: 4 moles of chloride ion dissolved for one mole of gold. 47

Similarly, the stoichiometry of the cathodic reaction is represented by:

0Cr+H20 + 2e = Cl~ +20H~ (2.55)

The rate of the cathodic reaction of gold, Rc, reduces to:

(2.56)

where It:; a constant, —Ai o D .j: diffusion coefficient of OCT ([0C/~]) : concentration of OCf in the bulk solution

factor 2/3: 3 mole of hypochlorite dissolved for 2 moles of gold.

At steady state, the rate of the anodic reaction must be equal to the rate of the cathodic reaction.

Thus, 8 D ['OCX-] {[ocn\ (2.57) icr]

For electrochemical reactions, the reduction potentials of reaction (2.53) at 25°C can be calculated from the Nernst equation, as follows:

„ „ 0.059, [Audi] E-E +------log------7- (2.58) n [cr]

The effect of chloride concentration can be expressed in terms of electrode potentials. The potential decreases when the chloride concentration is 48

increased. This corresponds to a decrease in the adsorption driving force due to the decrease in the difference between the reduction potential of chlorine and the gold deposition reaction.

A comparative study (Marsden and House, 1992) of the rate of dissolution of gold in aqueous solutions by using a gold leaf test showed that the rates in a chloride solution, 0.3 g/(m2.s), is much faster than the rates in an equivalently strong cyanide solution, 0.008 g/(m2.s). Such a preferential dissolution is attributed to the high solubility of chlorine in water compared to that with oxygen. In his investigation, the presence of sodium chloride (3%) in the chlorine solution had a considerable accelerating effect, possibly due to the retarding effect of chloride ions on chlorine dissociation.

2.3.6 Process Selection Considerations

In general, chlorination is more difficult to use directly in industrial applications for two basic reasons:

• the leaching media is of a highly corrosive nature and requires special materials of construction to withstand corrosion in the plant resulting in higher costs • the chlorine gas must be contained to minimize the emission of reagent and avoid health risks associated with its use.

Gold dissolution by chlorination (Solozhenkin et al., 1993; Wang et al., 1993) is achieved in the treatment of a wide range of ores. Unfortunately, chlorination of whole ores was discontinued because of the more favourable economics and process simplicity of cyanide leaching after the advent of the cyanide process. For this reason, chlorination has been limited in its 49 application and little process data is available. Recently, Newmont Mining Corporation (Nevada, USA) has developed the chlorination process (Guay and Peterson, 1973) for the treatment of carbonaceous ore prior to cyanidation. Scheiner et al. (1972) have determined that a sodium hypochlorite pre-leach of carbonaceous refractory gold ores, followed by cyanidation, enables high subsequent gold recovery to be achieved up to 90% gold dissolution. Furthermore, other methods of chlorination-gold recovery process continue to be investigated.

The information presented in this chapter suggests that:

• Gold can be recovered effectively by chlorination during the processing of ores and recycling of metals such as electronic scraps, slags, dusts, residues, and wastes from other processes.

• Many processes for gold dissolution can take place if suitable oxidants such as oxygen, chlorine, hydrogen peroxide and complexants such as cyanide, chloride, and thiourea, are present in the leach solution.

• With a knowledge of mechanism and rate-controlling factors of gold dissolution in the cyanide and chloride systems, the application of a commercial process for the extraction of gold from ores should be readily achieved.

• Chlorination has been re-introduced as a pretreatment process to oxidize carbonaceous materials and sulfide minerals of certain gold ores to liberate the gold prior to cyanidation. CHAPTER THREE

LEACHING OF REFRACTORY GOLD ORES

3.1 Refractory Ores

3.1.1 General

Ores that are typically unresponsive to conventional cyanidation and require pretreatment before cyanidation are termed refractory. Several problems are encountered in the recovery of gold from such ores (La Brooy et al., 1994);

(i) In highly refractory ores, the inefficiency of conventional direct cyanidation owing to the lock up of gold in the mineral matrix.

(ii) In complex ores, the high consumption of leach reagents and the insufficient cyanide and/or oxygen in the pulp, due to the side reactions with reactive minerals in the ore.

(iii) The loss of the dissolved gold cyanide complex from the leach liquor because of the adsorption or precipitation onto components of the ore.

The degree of ore refractoriness related to recovery might be determined and classified as shown in Table 3.1. 51

Table 3.1 Classification of Ore Refractoriness (La Brooy et al., 1994)

Refractoriness degree Recovery(%) Highly refractory < 50 Moderately refractory 50-80 Mildly refractory 80-90 Non-refractory(Free Milling) 90-100

3.1.2 Mineralogical Characteristics

Mineralogical characteristics responsible for the refractoriness in gold ores can be classified as siliceous refractory ores, "preg-robbing" of carbonaceous ores, and sulfidic refractory ores.

3.1.2.1 Siliceous Refractory Ores

Siliceous refractory ores (Hausen, 1989) are referred to as locked fine metallic gold or auriferous sulfides within an impermeable matrix of microcrystalline quartz, chalcedony, chert or amorphous forms of opaline silica. In the category of this type of ore, roasting or chlorination does not effectively increase the recovery of gold. Cyanide solutions cannot attack the finely locked gold because the pore size from micro-meter to submicro-meter in many siliceous host rocks does not allow for sufficient impregnation. Moreover, roasting or chlorination does not significantly increase the recovery of gold because of the poor permeability of ore fragments. However, after metallic gold is increasingly exposed by finer grinding, the degree of gold recovery usually improves. 52

3.1.2.2 "Preg-robbing" of Carbonaceous Ores

The active adsorption of gold from cyanide solutions, by selective mineralogic components, such as organic components and elemental carbon, can be referred to as preg-robbing. This term was proposed to characterize the deleterious effect of carbonaceous materials in gold ores (Rota and Ekburg, 1988) on cyanidation by the Carlin Gold Mining Co. in 1968. It is generally accepted that various forms of carbon were many times more preg- robbing than the noncarbonaceous mineral phases. This preg-robbing behaviour can be reduced by use of pretreatments such as flash chlorination (with Cl2), roasting, and bacteria leaching to deactivate the carbonaceous materials.

3.1.2.3 Sulfidic Refractory Ores

The auriferous sulfides (Wells and Mullens, 1973; Hausen, 1989; Mason, 1990; McDonald et al., 1990), containing iron sulfides averaging between 1.5 and 2% for the total orebody, cannot be ground sufficiently fine to liberate sub-microscopic gold from the sulfide minerals in which gold is encapsulated. These gold ores typically contain sulfides, such as pyrite (FeS2), arsenopyrite (FeS2 FeAs2), galena (PbS), and copper minerals. Furthermore, these sulfide minerals can not only consume oxygen but form cyanicides in the cyanide solutions, thus depleting the level of dissolved oxygen and the cyanide content of the pregnant cyanide solutions.

Pretreatment required to produce high extractions of gold are necessary to deal with sulfidic refractory ores amenable to conventional cyanidation. Several investigators (Kontopoulos and Stefanakis, 1990; Yang and Gao, 53

1993) have reported that various types of pretreatment methods, including roasting, chlorination, hot-water treatment, and pressure oxidation, can be economic means, with which to expose the gold to cyanidation recovery.

3.2 Leaching of Refractory Ores

3.2.1 Direct Gold Leaching

The process of direct gold leaching can be roughly classified as alkaline pressure cyanidation, carbon-in-leach (CIL) and carbon-in-pulp (CIP) process, and non-conventional lixiviants leaching.

3.2.1.1 Alkaline Pressure Cyanidation

Alkaline cyanide leaching under oxygen pressure is very effective in the treatment of free milling ore and refractory ores. As reviewed by Haque (1987), a comparative study showed that a gold ore containing 17.6g Au/t was leached up to 98% gold in only 15mins by alkaline pressure cyanidation, whereas atmospheric pressure leaching yielded 96% of the gold in 24 hours for the same ore. Alkaline pressure cyanidation technology, for the recovery of gold, has only been applied commercially on a very limited scale.

3.2.1.2 Carbon-in-leach (CIL) and Carbon-in-pulp (CIP)

Over the last few decades several investigators (Gilman, 1962; Rhodes, 1976; McDougall and Hancock, 1980; Abotsi and Osseo-Asare, 1986; Hughes and Linge, 1989; Stange et al., 1990; Gallagher et al., 1990; Klauber and Vernon, 1990; Kongolo et al., 1990; McDonald et al., 1990; White, 1990; Afenya, 1991) have extensively studied the carbon adsorption systems for 54 improving the recovery of gold from refractory ores. Several modifications of the carbon adsorption system, with cyanidation, were designed to establish a process, such as carbon-in-leach (CIL), carbon-in-pulp (CIP), and carbon- in-column (CIC), that simultaneously treats the refractory ore. Recently, carbon-in-chlorine-leach (CICL) has been suggested for treating the refractory gold ores by Greaves et al. (1990).

CIL (carbon-in-leach) is a process where the leaching and precipitation of gold, occur simultaneously in a single stage. This process (Haque, 1987) involves the use of active carbon to adsorb the gold cyanide complex

[Au(CN)2 ] from the pregnant cyanide leaching solutions. The mild preg- robbing behaviour of carbonaceous type refractory ore can be reduced by the use of CIL because the gold loading kinetics of activated carbon may be faster than that of the preg-robbers. The CIL process enhances the dissolution of gold, but still some problems remain, such as carbon attrition, fouling, and a lower gold loading than with the CIP process.

CIP (carbon-in-pulp) recovery processes (Udupa et al., 1990) generally involve carbon passing counter-current to the process slurry in adsorption contactors. The CIP process involves the extraction of gold from cyanide leach slurries by the use of granular carbon (6 to 16 mesh). The loaded carbon is eluted prior to the recovery of gold by electro-winning. Generally, a modified Zadra process is used to recover the adsorbed gold.

The CIC (carbon-in-column) for gold recovery from heap leach solution is a con-current technique which involves the direct contact of the gold solution with activated carbon, as shown in Figure 3.1. This process followed by electrowinning was used at Round Mountain (Nevada, USA) to treat approximately 40,000 tpd ore. 55

Ore column

Carbon columns

Barren cyanide solution

Figure 3.1 Laboratory CIC unit for the recovery of gold from a pregnant solution by using activated carbon (Heinen et al., 1978)

CICL (carbon-in-chlorine-leach) is a process by which gold can be simultaneously recovered via leaching and precipitation from the pretreated pulp by the chlorination of refractory carbonaceous and sulfidic ores. Greaves et al. (1990) have extensively studied the recovery of gold by batch chlorination tests. Their work indicated that the use of CICL with highly carbonaceous ores achieves a higher recovery than CICL with primarily sulfidic refractory ores. In the case of cyanide amenability testing, two carbonaceous ores achieved recoveries of only 5.5% and 46%, respectively. 56

Recoveries with CICL treatment were 90% and 92%, respectively, from the same ores.

3.2.1.3 Leaching with Non-Conventional Lixiviants

Although cyanide leaching has been recognized for a long time as a powerful lixiviant of gold, cyanide solutions are toxic and their use is accompanied with concerns of accidental cyanide spills which might cause environmental harm. Moreover, cyanide leaching is not economic for certain classes of gold ores and concentrates containing cyanicides (sulfide minerals, copper oxide or silicates etc.) which are considered refractory to the cyanidation process. Refractory gold ores are those from which a significant percentage of gold cannot be recovered by conventional cyanidation. Over the last few decades, alternative gold recovery processes using more powerful lixiviants other than cyanide have drawn increasing attention due to the limitations and increasingly stringent pollution control regulations associated with cyanidation.

The application of lixiviants, such as chloride, iodide, bromide, thiourea, thiosulfate etc., other than cyanide has been seriously investigated by several researchers because they have certain advantages, such as faster leaching rates and their relative lower-toxicity reagents, over cyanide. As reviewed by several investigators (Benari et al., 1983; Haque, 1987; Hiskey and Atluri, 1988; Demopoulos and Papangelakis, 1989; Udupa et al., 1990; Tran, 1992), although the gold recovery processes using these lixiviants are not widely used at present, few advantages over conventional cyanidation are still being discovered. 57

3.2.1.3.1 Chloride

Gold can be dissolved in the presence of chloride ions and a strong oxidant. In chlorination, dissolved chlorine, hypochlorous acid, and hypochlorite ions, as strong oxidants, are capable of oxidizing gold in a chloride media as described in the previous chapter. Generally, it is accepted that the dissolution rate of gold in chlorination is favoured by low pH, high chloride and chlorine levels, increased temperature, and high ore surface area. The use of chlorine for the pretreatment of refractory material has been examined in an attempt to overcome problems associated with refractory ores. Alkaline chlorination pretreatment has been afforded significant attention for two classes of gold ores as described below.

Carbonaceous gold ores

In the past two decades, the deleterious effect of the carbonaceous material, present in gold ores, on cyanidation has been investigated by several researchers (Nice, 1971; Scheiner et al., 1971; Scheiner, 1986; Abotsi and Osseo-Asare, 1986; Hutchins et al., 1988; Afenya, 1991). This phenomenon can be referred to as "preg robbing".

The primary application of chlorination (Scheiner et al., 1971) is to destroy the deleterious effect of the gold-adsorbing native organic carbon which occurs in certain ores, reducing the aurocyanide complex formed during cyanidation. In the presence of chlorine, the process of deactivation, or passivation is not well understood. The direct oxidation reaction of organic carbon to carbon dioxide can be represented as: 58

2 HOCl + C = C02 + 2 HCl (3.1)

The results of several investigations (Heinen et al., 1978; Avraamides et al., 1985; Abotsi and Osseo-Asare, 1986; Afenya, 1991) have shown that carbonaceous materials consist of two major components, elemental carbon and organic carbon composed mainly of carboxyl groups (COOH). As discussed in the previous chapter, chlorine oxidation between pH 3 and 5 is the most achievable, because in this pH range the hypochlorous species with strong oxidizing properties predominates over other chlorine solution species. Deactivation properties are reduced below pH 2 and above pH 6 due to the reduced hypochlorous activity.

The exact hypochlorous concentration level required for deactivation of carbonaceous materials is limited by the amount of competing species in the ore, i.e. sulphides and carbonates, and the solution conditions. In general, hypochlorous concentrations of about 1 g/1 are used to deactivate carbonaceous materials. The results of a temperature test for oxidation (Marsden and House, 1992) have shown that chlorine gas absorption is favoured at low temperatures, while the diffusion rate of hypochlorous species and overall chemical kinetics are favoured at higher temperatures.

In addition, Table 3.2 shows the effects of temperature on the solubility (Chao, 1968) of hypochlorous species. As shown in this table, the solubility decreases with increasing temperature. In practice, optimum oxidation is obtained at approximately 50°C. 59

Table 3.2 The effect of temperature on chlorine solubility

Temperature(°C) Cl2 solubility(mCl2/m H20)

20 7 x 10'3

40 4.1 x 103

60 2.5 x 10'3

Refractory sulfide gold ores

Aqueous chlorine will readily oxidize all sulfides commonly associated with gold. The overall reaction for sulfide minerals, such as sphalerite, pyrite, chalcopyrite, and chalcocite within the range of pH values can be represented as (Cho, 1987a and b; Marsden and House, 1992):

3 ZnS + 6 HOCl = 3 Zn2+ + 2S + SOf + 2H+ + 6 C/“ + 2 H20 (3.2)

2 FeS2 +15 HOCl + IHfj = 2Fe(OH)3 + 23/T + 45'Of + 15C/" (3.3)

2CuFeS2 +1 l^OC/ = 2Cw2+ + Fe203 + 2SOf +2S + 11 Cl~ +11H+ (3.4)

Cu2S + 2HOCl + 2H+ =2Cu2+ +S + 2H20 + 2Cl- (3.5)

The major reactions involve the total oxidation of the sulfidic sulfur to the sulfate form, in the strongly oxidizing conditions. Other sulfides are similarly oxidized. Arsenic minerals are oxidized to arsenic trichloride states which are precipitated and could be recovered as a stable product. The rate­ controlling and sulfide oxidation factors are similar to those affecting other 60 aqueous sulfide oxidation processes and the rate will also depend on the area of the reaction interface.

Chlorination is not an economical process to treat ores which contain high levels of sulfide, due to excessive chlorine consumption. For this reason the efficient leaching of gold in an acidified chloride system is restricted to the treatment of ores containing low sulfide content, typically less than 0.5%. Gold is extracted from such ores and concentrates with higher sulfide contents by employing one of the alternative processes, such as pre-aeration, pressure oxidation or roasting.

In Australia, such processes utilizing chlorine pretreatment have price advantages over the cost of cyanide. This is because of the ability to electrolytically produce chlorine from ground water brines, which are common in the gold fields of Western Australia (La Brooy et al., 1994).

3.2.1.3.2 Iodide

Iodide is a powerful complexing agent for gold and the gold iodide complexes are the most stable among those of the halide systems. Recently, several researchers (Lai et al., 1951; Hubbard et al., 1966; Falanga and MacDonald, 1982; Murphy, 1988; Davis and Tran, 1991; Qi and Hiskey, 1992) have extensively studied the kinetics and mechanisms of the gold-halide system with regard to the extraction of gold. According to their works, iodide is oxidized preferentially to tri-iodide when dissolved iodine reacts with iodide. These studies agree that the process is controlled by surface electrochemical reaction. The chemistry of gold dissolution in iodine solution has been elucidated by Hiskey and Aduri (1988) and Davis et al. (1993). Potendal-pH diagram for the Au-f-H20 system is presented in Figure 3.2. 61

2.0

1.5

1.0 Eh

0.5

0.0

-0.5 -2 0 2 4 6 8 10 12 14 16 PH

Figure 3.2 Eh-pH diagram for the Au-I-H20 system at 25°C;[Au] = 105 M, [I'] = 10'2M (Hiskey and Atluri, 1988)

This diagram indicates that metallic gold will dissolve to generate the Aul2 complex ion at about 0.51V vs SHE, while the Aul4 becomes stable at approximately 0.69V. The extent of this domain varies directly with the concentration of gold in solution. The domain of stability of the Aul2’ species is greater than that of either AuC14 (Figure 2.8) or AuBr4 (Figure 3.3).

Solid iodine is highly soluble in aqueous iodide solution and dissolves to form the tri-iodide, according to the following reaction:

^2(aq) + ^ _ A (3.6) 62

The electrochemical reactions can be written as follows:

Anodic Au + 2I~ -> Aul~ +e~ (3.7)

Cathodic I~ + 2e~ =37“ (3.8)

Overall 2Au + I~ + I~ -> 2AuI~ (3.9)

From the above reactions, the I3 species can act as an oxidizing agent for the dissolution of gold. Moreover, the Nernst equation for the I3/I couple is:

E = 0.54 - 0.03 log (3.10)

Application of the iodide system for the extraction of gold has certain advantages, such as high stability, low oxidation potential, selectivity in leaching other minerals, and high gold dissolution rates over other leachants. Furthermore, an in-situ leaching process based on the oxidation by tri-iodide is described in a patent by McGrew and Murphy (1985). Until recently, iodide leaching technology for the recovery of gold has not yet been applied commercially on a large scale.

3.2.1.3.3 Bromide

Bromide systems (Tran and Armoo, 1990; Pesic et al., 1992) have been studied for many years for their suitability in extracting gold. Bromide is a powerful complexing agent for gold and gold can be dissolved at very fast rates within this system. 63

The diagram in Figure 3.3 presents the thermodynamics of the Au-Br-H20 system. The concentrations of gold and bromide are 10s M and 10 2 M, respectively. This diagram indicates that metallic gold will dissolve to generate the AuBr2 complex ion at about 0.90V vs SHE, while the AuBr4 is formed at approximately 0.95V. The domain of stability of the AuBr4 species is greater than that of AuC14 (Figure 2.8).

The electrochemical reaction in a bromide solution can be written as follows:

Au + 4Br~ = AuBr~ +3e~ (3.11)

Au(OH).

HAuO

AuBr

Figure 3.3 Eh-pH diagram for the Au-Br -H20 system at 25°C;[Au] = 10 5 M and [Br ] = 10 2M (Hiskey and Atluri, 1988) 64

This system is strongly oxidizing and the achieved dissolution rate of gold is faster than that of cyanidation. Several advantages for bromide leaching are claimed, such as rapid extraction, non-toxicity, and applicability to a wide range of pH values.

Dadger and Howarth (1992) have suggested that the high dissolution rate and low recovery costs of the bromine process have a certain economic advantage over cyanidation. However, despite some potential advantages, the commercial application of bromine solutions for gold leaching is still not economical due to high reagent consumption.

3.2.1.3.4 Thiourea

Thiourea, CS(NH2)2, is capable of dissolving gold in acidic conditions and forms a cationic complex. The thiourea leaching system (Groenewald, 1975; Deschenes and Ghali, 1988) offers a relatively non-toxic and a greater rate of gold dissolution compared to that by cyanidation. A number of publications (Filmer et al., 1984; Huyhua et al., 1989; van Deventer et al., 1990) have suggested that thiourea is as an attractive alternative to cyanide for the treatment of sulfidic, cyanide-consuming ores, and for use in sites where environmental concerns make the use of cyanide difficult.

The gold dissolution reaction in acidic thiourea solutions involves the formation of a stable gold complex species as follows:

Au + 1CS(NH2)2 = Au[CS(NH2)2]+2 +

i?h

Figure 3.4 Eh-pH diagram for the Au-CS(NH2)2-H20 system at 25°C; [Au] = 10"4 and [CS(NH2)2] = 10'2M (Gaspar et al„ 1994)

On the basis of equilibrium reactions and thermodynamic data, the Au- CS(NH2)2-H20 system (Gaspar et al., 1994) is expressed in the following Eh- pH diagram (Figure 3.4). In the absence of oxidants, the thiourea solution achieves inadequate gold extraction, within practical time scales, with an extraction of only 20.5% of the gold in 8 hrs. Gold in the thiourea solution can be readily oxidized by the addition of oxidants, such as formamidine disulfide, hydrogen peroxide, sodium peroxide, dissolved oxygen, the ferric ion, ozone, and manganese dioxide. Of these, it is generally accepted that ferric ion can significantly improve the driving force of the reactions. The overall reaction can be written:

2Au + 4CS(NH2)2 +2 Fe^ =2 Au[CS(NH2)2]+2 +2 Fe2+ (3.13) 66

The dissolution of gold is dependent on the pH, temperature, redox potential, thiourea and oxidant concentration. The rate of gold dissolution in acidic thiourea solutions is usually controlled by the diffusion of reactants to the gold surface. Several metal species form metal-thiourea complexes to varying extents, in acidic thiourea solutions, as illustrated in Table 3.3. Most of the process development work related to thiourea has been focused on the leaching step. Moreover, Schulze (1984) has suggested that activated carbon can be used to recover gold from thiourea solutions.

Table 3.3 Stabilities of metal-thiourea (Tu) complexes

Complexes PK

Au(Tu)2+ 21.75 Ag(Tu)3+ 13.10

Cu(Tu)42+ 15.40 Zn(Tu)22+ 1.77 FeS04(Tu)+ 6.64 Cd(Tu)42+ 3.55 Pb(Tu)42+ 2.04

3.2.1.3.5 Thiosulphate

Gold in an alkaline thiosulphate (S2032) solution can be oxidized by dissolved oxygen to form the Au(I) complex, according to the following equation:

AAii + SSjOl’ +2 H20 + 02 = 4Au(S203)3- +AOH~ (3.15) 67 1.2

i.i -

i.o -

0.9 -

0.8 -

0.7 - Au(S 2 O3)2 _ 0.6 - > 0.5 - -C Uj 0.4 -

6 7 8 9 10 11 12 13

Figure 3.5 Eh-pH diagram for the Au-S2032 -H20 system at 25°C (Marsden and House, 1992)

An Eh-pH diagram (Marsden and House, 1992) of the Au-S2032 -H20 system is shown in Figure 3.5, showing a large stability zone for the gold thiosulphate complex. The dissolution of metals in an alkaline thiosulphate solution (Changlin et al., 1992) depends on many factors, such as temperature, thiosulphate and oxidant concentrations. Moreover, gold dissolution (Hiskey and Atluri, 1988; Langhans et al., 1992) has been observed to be very sensitive to the concentration of thiosulfate and copper. The rate of gold dissolution increases to some extent on the addition of small amounts of copper(II) ions which have a strong catalytic effect. This effect on the rate of the dissolution of gold is presented by the following reactions:

4Cu(S203)]- +\6NH,+02 +2H20 = 4Cu(NH3)24+ +SS2Oj- + 40H~ (3.16) and

Au + \S2Ol~ + Cu(NH3 )]+ = Au(S203)32- +4NH3+Cu(S203y2- (3.17) 68

A process using thiosulphate ions, generated in situ under alkaline oxidation conditions, for the treatment of sulfidic gold ores has been proposed by Kerley (1981). However, despite several potential advantages, there are no commercial processes using the reagent at present.

3.2.1.3.6 Thiocyanate

Gold can be oxidized in the presence of thiocyanate (SCN) to form both the

Au(SCN)2 and Au(SCN)4’ complexes (Stanley, 1987; Marsden and House, 1992):

Au(SCN)2 +e = Au + 2SCN~, E° = 0.662 V (3.18)

Au(SCN)-+ 3e = Au + 4SCN~, E° = 0.623 V (3.19)

Gold in the presence of thiocyanate can be readily dissolved in the presence of oxidizing agents, such as ferric ion, hydrogen peroxide, dissolved oxygen, and Caro's acid (H2S05). The ferric ion (Fe3+/Fe2+, E° = 0.77V) is considered as the most suitable oxidant for this system. This is because the rate of gold dissolution is slow, due to its low solubility in water, particularly under atmospheric conditions. Thiocyanate is oxidized very rapidly by hydrogen peroxide and dissolved oxygen. Caro's acid is a more powerful oxidizing agent than hydrogen peroxide.

Figure 3.6 represents the Eh-pH diagram for the Au-CNS-H20 system at 25°C. The stability (Barbosa and Monhemius, 1988) of Au(I) complexes in the presence of thiocyanate is greater than that of chloride complexes and less than that of cyanide complexes. Thiocyanate complexes with the ferric ion 69

are less stable than those with gold. This means that Au cations would be favoured in the competition for thiocyanate ions in the presence of ferric ion.

Au(OHMs) Au0 Au(OH)

[Au] = 10 [CNS] = 10

Figure 3.6 Eh-pH diagram for the Au-CNS-H20 system at 25°C (Marsden and House, 1992)

The rate of the dissolution of gold increases with increasing thiocyanate, and to a lesser extent, ferric ion, concentration. However, such reactions will incur a high degree of reagent consumption, especially in acidic solutions where the metal ion concentration is high because thiocyanate forms complexes with metal cations (e.g. Fe(SCN)2+). Further investigations are required on the kinetics of both gold dissolution and thiocyanate oxidation.

3.2.1.3.7 Polysulfide

Polysulfides have been proposed as alternative lixiviants for gold (Krauskopf, 1951). The general equation for the overall formation of a polysulfide could be written as: 70

nS2~ S2n~ +(2n-2)e~ (3.20)

A Russian group of scientists have extensively studied the complexation reactions between gold and polysulfides (Hiskey and Atluri, 1988). According to their work, gold dissolves in aqueous solutions of polysulfides to form the Au(I) complex as follows:

Au+ + S2~ = AuS~ (3.21)

In another work, S22 anion was identified as the most suitable oxidant for the reaction, according to the following equation:

2 Au + S2- = 2AuS~ (3.22)

It is worth emphasizing that sulfide/polysulfides may play important roles in the breakdown of refractory gold-bearing sulfides. In addition, the leaching chemistry of As2S3 (orpiment) and Cu3AsS3 (tennantite) dissolution with sulfide/polysulfides has been considered and the general equations representing these reactions are described briefly below.

The general equation for As2S3 dissolution with polysulfides is given as follows:

As2S3+3S2~ ->2AsS33~ (3.23) 71

The general equation for Cu3AsS3 dissolution with sulfide ions is represented by:

2Cu3AsS3 + 3S2~ -> 3Cu2S + 2AsS]~ (3.24)

3.2.2 Pretreatment of Refractory Ores

Over the last decade the treatment of refractory gold ores, as described earlier, has received considerable attention because processes for the extraction of gold have improved significantly. Economic gold extraction processes have been applied to low grade and refractory gold ores and even to old mill tailings. This is due to advanced leaching techniques, such as heap leaching, vat leaching, in-situ leaching, carbon-in-pulp (CIP), carbon-in­ leach (CIL), carbon-in-column (CIC), carbon-in-chlorine-leach (CICL), and pretreatment of refractory gold ores (Fraser et al., 1991).

Recently, the pretreatment of refractory gold ores has been studied extensively. A range of assessment has been carried out into some of the different alternatives available for pretreatment by several investigators (Berezowsky and Weir, 1984; Kontopoulos and Stefanakis, 1990). According to their work, several factors including technical, environmental, and economic characteristics related to the specific claims of different processes have been considered. Pretreatment of refractory ores can include: roasting, biological oxidation, and pressure oxidation. 72

3.2.2.1 Roasting

The major characteristics of roasting are that it is a simple process and a conventional method of pretreating refractory ores or concentrates which contain a high sulphur content. Roasting is generally carried out in a single or two-stage process. The single stage process involves the direct roasting of the material in an oxidizing atmosphere. The two-stage process, employs a first stage which operates under reducing conditions, creating a porous, intermediate product, followed by a second stage roast in an oxidizing atmosphere to complete the oxidation. A systematic block diagram for the two-stage processing of arsenopyrite ores and concentrates is illustrated in Figure 3.7.

The main advantages (La Brooy et al., 1994) of the pretreatment of refractory gold ores by roasting are:

• the destruction of the metal sulfide phase, • the exposure of the gold in the sulfide, • the production of a, high surface area/high porosity, hematite (Fe203) calcine, • minimisation of the production of reactive constituents in the calcine, e.g. pyrrhotite, ferrous sulfate etc., which could consume reagents, and • the fixation of arsenic and sulphur in an environmentally acceptable form. 73

Gold ores/concentrates

Superheated Steam To Atmosphere 1st Stage Roasting Gas Sulfuric Acid Cleaning Plant

2nd Stage Roasting Arsenic Trioxide Sulfuric Acid

Calcine to Gold Extraction

Figure 3.7 A systematic block diagram of the two-stage roasting pretreatment of arsenical pyrite gold ores or concentrates (Kontopoulos and Stefanakis, 1990)

Generally oxidative roasting has been used to process pyritic and arsenopyritic ores and concentrates. Pyritic roasting can be performed within a temperature range of 600°C and 700°C as a "dead roast". Under these conditions, diis process converts pyrite, marcasite, and pyrrhotite to magnetite, and then further to hematite (Fe203) which is insoluble in cyanide solution and sufficiently porous to expose the gold particles.

The decomposition of pyrite to hematite can be expected as follows:

4FeS2 +1102-> 2Fe20, + 8S02 (3.25) 74

This reaction shows that the sulphur is volatilised as S02. The operational variables in this process are the partial pressure of oxygen and the temperature in the roasting chamber.

Arsenopyrite roasting can be performed either as a single stage or two-stage processes within a temperature range of 450°C and 550°C. Arsenopyrite is ultimately oxidized to hematite, arsenic tri-oxide (As203), and S02. This is shown by the following overall reaction:

2FeAsS + 502 —» Fe203 + As203 + 2S02 (3.26)

The arsenic, as volatile As203, is cooled and collected as a solid. The roasting process is usually carried out in rotary kilns, stationary fluidised-bed and circulating fluidised-bed reactors.

In Australia, fluidised-bed roasters are currently being used to treat pyritic- gold telluride concentrates at Kalgoorlie in Western Australia. In addition, roasting with a pyritic concentrate containing 5% arsenopyrite is carried out in a single stage, stationary fluidised bed reactor by Western Mining Corp. at Windarra in Western Australia. However, a number of problems, such as the emission of particulates, sulphur dioxide, arsenic, mercury etc., exist. New stringent regulations on the environmental pollution controls limit the application of roasting processes.

3.2.2.2 Biological Oxidation

In the early 1900s biological oxidation was initiated at Rio Tinto (Spain) for heap and dump leaching of low grade copper ores. Many types of naturally occurring bacteria have a catalysing effect on mineral oxidation reactions. 75

Of these bacteria, Thiobacillii (Thiobacillus ferroxidans and Thiobacillus thiooxidans) are the most suitable for the oxidation of gold-bearing sulfide ores and concentrates (Brierley, 1982). The design of a suitable environment for the maintenance of a bacterial culture (i.e., 25°C to 40°C at atmospheric pressure; pH 1-3; 02 and C02 from air) is considered as the most important aspect of bacterial oxidation. A systematic block diagram for bacterial oxidation is illustrated in Figure 3.8.

Gold ores/concentrates

Regrinding

Preconditioning

Nutrients Air + CO Bacterial Lime, Oxidation Limestone

S/L Separation & Washing Neutralization

Solids

To Gold Extraction To Tailings Pond

Figure 3.8 A systematic block diagram for bacterial pretreatment of arsenical pyrite gold ores or concentrates (Kontopoulos and Stefanakis, 1990)

As reviewed by Kontopoulos and Stefanakis(1990), the major reactions occurring during pyrite and arsenopyrite bio-oxidation have been considered 76

as either the direct or indirect oxidation by micro-organisms, according to the reactions: Direct microbial oxidation

2FeS2 +102 + 2H20 bac,ena >2FeSO, +2H2SO, (3.27)

4FeSO, + 02 +2 H2SO, hactena >2Fe2{S04)i+2H10 (3.28)

AFeAsS +1302 + 6H20-jM,n« >4H}As0* +4FeS0i (3.29)

The formation of arsenous (As3+) acid as an intermediate product can be expected. The ferric sulfate generated from the above reactions oxidizes the remaining sulfides alternately, according to the following reactions:

Indirect microbial oxidation

2FeAsS + Fe2 {SO, )3 -> 2 H,AsO, + AFeSO, + H2S04 (3.30)

FeS2 + Fe2 {SO,)3 -> 3FeSO, + 2S° (3.31)

Ferrous sulfate is re-dissolved, while the elemental sulfur generated, is further oxidized by Thiobacili, according to the reaction:

2S° +2H20 + 302 >2H2S04 (3.32)

During pyrite and arsenopyrite bio-oxidation, Fe3+ and H2S04 species generated can promote the dissolution of various mineral species. In Australia (La Brooy et al., 1994), at the Harbour Lights gold mine in Western Australia a 40 tpd plant for the treatment of arsenopyritic 77

concentrates is operated based on Gencor's BIOX technology. The same technology is being used at another bio-oxidation plant designed to treat 115 tpd pyritic/arsenopyritic concentrate at Wiluna. Generally, the bio-oxidation reaction (Brierley, 1990) has some disadvantages, such as slow rate of reaction, careful temperature control, corrosion due to low pH pulps, and possible high cyanide consumption by elemental sulphur. However, much research work of various aspects of bio-oxidation as an available technology for treating refractory ores is currently in progress due to significantly low capital and operating costs and simplicity of the operation (Kleinmann et al., 1981; Lawson etal., 1990; Hackl, 1990).

3.2.2.3 Chemical Oxidation of Refractory Ores

Pretreatment processes by chemical oxidation can be classified into three major categories: • oxidative acid leaching at low pressure, • oxidative alkaline leaching at high pressure, and • oxidative acid leaching at high pressure.

The pressure oxidation of gold-bearing refractory ores and concentrates as a pre-cyanidation stage has been studied (Berezowsky and Weir, 1984; White, 1990; Kontopoulos and Stefanakis, 1990) extensively for different types of oxidants such as ozone, hydrogen peroxide, permanganate, chlorine, Caro's acid, perchlorate, hypochlorite, ferric ion, nitric acid, and oxygen.

3.2.2.3.1 Oxidative Acid Leaching at Low Pressure

Low pressure oxidation using a chemical oxidant can be used to oxidise sulfide minerals. This can be applied as a pretreatment process prior to 78

cyanidation to oxidize and/or deactivate the reactive, reagent-consuming sulfides. This treatment is often only capable of the partial oxidation of sulfides and is usually unsuitable for the treatment of ores where gold is encapsulated in larger unreactive and non-porous sulfide grains. Under this category, the Nitrox process and Caro's acid process as pretreatment stages have been described by several investigators (Haque, 1987; Kontopoulos and Stefanakis, 1990; La Brooy et al, 1994).

Nitrox process

The process (van Weert et al., 1986) uses air at atmospheric pressure and nitric acid as an oxidant to dissolve sulfide minerals such as arsenopyrite and pyrite. The oxidation reactions (Afenya, 1991) for pyrite are represented by:

2FeS2 + 10HNO3 -> Fe2(S04)3 + H2S04 +10NO(g) + AH20 (3.33a)

FeS2 + Fe2(S04)3 -> 3FeS04 + 2S° (3.33b)

2NOM+OaU)-+2NOa(t) (3.34)

3N02(g) + Hfi -> 2HN03 + NO(g) (3.35)

The operating conditions of this process are atmospheric pressure, 90°C with 1-2 hours retention time, and a pH < 1. By such pretreatment, the extraction of gold can be improved from 30 to 90%. About 50% of the sulphur forms elemental sulfur as an intermediate product. The elemental sulphur has to be converted into a soluble compound prior to cyanidation because the sulphur formed during oxidation may occlude the gold or coat unreacted sulfides. 79

Caro's acid process

Caro's acid process is based on the oxidation of gold ores or concentrates using peroxymonosulphuric acid (H2S05) as the oxidant, which is commonly known as Caro's acid. Caro's acid contains the species H2S05 and H202 and is a source of free oxygen in solution, as expressed in the following equations:

H2S05 + H202 -> H2S04 +H20 + 02 (3.36)

2H202 -±H20 + 02 (3.37)

The oxidation reaction is carried out at ambient conditions at pH 1.5, with an oxidation potential of 400 to 450 mV. The attraction of this process is that it enables the treatment of gold associated or occluded in arsenopyrite. The products of the oxidation of the sulfides are sulfates and arsenates and some elemental sulphur. The economics of this process are more favourable for the concentrate than for the total ore.

3.2.2.3.2 Oxidative Alkaline Leaching at High Pressure

Alkaline pressure oxidation is operated in a neutral or slightly alkaline media and the conditions of temperature, pressure, and oxygenation are similar to that required for acidic pressure oxidation. Oxidative pressure leaching in an alkaline system could be suitable for the treatment of refractory ores which contain large amounts of acid-consuming carbonates and have a low sulfide sulphur content. In neutral and alkaline solutions, this process utilizes 80 dissolved oxygen to oxidize sulfide minerals such as pyrite, arsenopyrite, and chalcopyrite, according to the following reactions:

4FeS2 +1502 +14H20 -> 4Fe(OH), + 16H+ + 8SO? (3.38)

IFcAsS + 102 +8H20 -> 2Fe(OH)i + 2HiAs04 +4//+ + 2S042~ (3.39)

4CuFeS2 + \102 + 18H20 -> 4Cu(OH)2 +4Fe(OH), + 16/T +8SOf (3.40)

A retarding effect on gold solubility and the extent of sulfide oxidation is caused by the insoluble metal oxides/hydroxides formed which coat gold surfaces and sulfide minerals. The process is thus hindered as the sulfide content increases, and consequendy this treatment may be only suitable for low sulphur feeds. Carbonates can usually promote non-acidic oxidation not only by its nonreactivity in neutral/alkaline media but also by the neutralization of acid that is generated by sulfide oxidation reactions. For example, this treatment can be applied to high acid-consuming carbonates (>10% C032’) and relatively low sulfide sulphur content (about 2%) ores.

One advantage of alkaline oxidation is that silver jarosite is not produced under these conditions and high silver recoveries can be readily achieved by cyanide leaching after oxidation. Contrary to acidic pressure oxidation, the drawback of this process may be the formation of mercury which remains in the autoclave discharge. Russian scientists have extensively studied the autoclave oxidation of gold-bearing sulfides and arsenopyrite with sodium hydroxide (Haque, 1987). According to their work, this treatment achieved 95 to 99% removal of sulphur and arsenic at temperatures of 100°C to 120°C, and subsequent cyanidation yielded 98% gold extraction, according to the following reaction: 81

2FeAsS + 6NaOH + —02 -> 2Fe{OH)2 + 2Na3As04 + Na2S04 + 3H20 (3.41)

Alkaline pressure oxidation has been applied commercially at Mercur plant in Utah (USA). Oxidation was carried out at high temperature (220°C) and 140 to 180 kPa oxygen partial pressure (total pressure of 3300 kPa) with sodium hydroxide consumption of 2 to 2.5 kg/t.

3.2.2.3.3 Oxidative Acid Leaching at High Pressure

The Arseno process uses pressurised oxygen for high arsenic-bearing concentrates. This process has been used for concentrate leaching under oxygen pressure (200 psi) at temperatures in the 60 to 120°C range with a retention time of only 15mins. The systematic block diagram for high pressure oxidation of arsenical pyrite ores or concentrates is illustrated in Figure 3.9. The major reactions take place, according to the following reactions:

3FeAsS + 3//+ + \4HN03 -> 3Fe}+ +3SOj~ + 3H3As04+\4NO + 4H20 (3.42) 2N02 +H20-> HN02 + HN03 (3.43) 3HN02 -> HN03 + 2NO(g) + H20 (3.44)

The major advantages of the use of nitrate-nitrite ions for sulfide oxidation over oxygen are: • high solubility of nitrogen dioxide in water, • the ability to regenerate the oxidant in the gaseous phase, and • a higher reduction potential resulting in faster oxidation. 82

In spite of these advantages the major drawback of this process is the formation of arsenides and elemental sulphur. The elemental sulfur which is generated is unfavourable for the pressure oxidation because: • occlusion of unreacted sulfides, hindering or preventing complete oxidation, • occlusion of gold particles, hindering or preventing extraction during cyanidation, • increased cyanide consumption during cyanidation, • increased consumption of oxygen during cyanidation.

Gold ores/concentrates

Regrinding

Recycling Feed Pretreatment

Oxygen Oxygen Pressure Plant Oxidation > Flashed Water Vapour Limestone Solids Counter-current Neutralization Washing

To Gold Extraction Tailings Pond

Figure 3.9 A systematic block diagram for oxidative pressure pretreatment of arsenical pyrite gold ores or concentrates (Kontopoulos and Stefanakis, 1990)

Furthermore, recent investigations into using this method on concentrates have shown great potential for this pretreatment process. 83

A review of leaching of refractory gold ores presented in this chapter suggests that:

• Several lixiviants, such as halides, thiourea, thiosulphate, thiocyanate and polysulfide, can be used for gold extraction. Although these lixiviants have been examined mainly in an attempt to treat complex metal sulfide ores and concentrates over the last few years, none has yet been commercialized on a large scale. Among these, the use of halides has assumed greater importance because the rate of gold dissolution with halides is much faster than with other lixiviants.

• The processes of direct gold leaching, such as alkaline pressure cyanidation, carbon-in-leach (CIL) and carbon-in-pulp (CIP), for the treatment of refractory ores have been applied over the past decade. Roasting, biological oxidation and chemical oxidation as pretreatment steps for refractory ores have also been evaluated.

• Chlorination has been successfully employed at Carlin mines since the early 1970's. Such a process of chlorine pretreatment has been proposed for the heap leaching of low grade gold ores and in-situ gold extraction systems.

• In Australia, processes utilising chlorine pretreatment could have a clear advantage over cyanidation. This is because gold plants are able to electrolytically produce chlorine economically from ground water brines, which are common in the gold fields of some of Western Australia where natural saline waters are available. CHAPTER FOUR

REACTION KINETIC MODELS

4.1 Homogeneous Kinetics

In general, the rate of reaction is a function of temperature, pressure, and the concentration of various species in the reaction. Reactions can be kinetically classified as homogeneous or heterogeneous. A homogeneous reaction occurs entirely in one phase, i.e., the reactants and the products are either in the gas phase or in the liquid phase. They can be catalyzed (for example, the hydrolysis of esters in the presence of hydrogen ions as catalyst) or non- catalyzed (for example, the neutralization of an acid with a base).

According to the law of mass action, the rate of a chemical reaction is proportional to the molar concentration (i.e., the number of moles in unit volume) of the reacting species. Consider a chemical equation written in the general form:

A + B-+C + D, (4.1) where A, B : reactants, respectively, participating in the chemical reactions C, D : products, respectively, generated by a bi-molecular reaction

If the rate is directly proportional to the concentration of a reacting species, the reaction is of first order. The order of the reaction is first order with 85 respect to reactant A, when the rate of decrease in the concentration of A at any time is proportional to that of its concentration and we have,

d\A = k[A] (4.2) dt

If the rate depends on two concentration terms; these may both refer to the same reactant or to two different reactants, the reaction is second order in A and B, respectively. The rate is given by,

d\A\ d[B] = k[A][B] (4.3) dt dt where for equations (4.2) and (4.3) [A], [B] : the concentration of A, B species, respectively k : a proportionality constant (velocity constant).

4.2 Heterogeneous Kinetics

4.2.1 Fundamentals

Heterogeneous reactions occurring in more than one phase play an important role in chemistry, metallurgy, geology, and in engineering. In general, a heterogeneous reaction has a rate which depends on the area of the surface that is exposed to the reaction mixture. Surface reactions proceed through the following step,

1. Diffusion of reactants from the bulk phase to the surface, 2. Chemical/electrochemical reaction on the surface, 3. Diffusion of the reaction products from the surface. 86

Any step or a combination of these steps could be rate determining in a heterogeneous reaction (Meng and Han, 1993), even though several steps may be present in the reaction mixture. If the rate of a single step is slower than all other steps, the overall rate of reaction is controlled by the rate of the slowest step, namely the rate determining step. The typical variation between the concentration and the distance from the surface is shown in Figure 4.1.

solid diffusion bulk layer solution

flux to solid (- ve)

flux to solution (+ ve)

Distance, x

Figure 4.1 Typical variation between the concentration and the distance from the surface, with regard to mass transfer 87

4.2.2 Heterogeneous Reactions

Generally, heterogeneous reactions (Smith, 1956; Sahimi and Tsotsis, 1988; Sahimi et al., 1990; Filippou and Demopoulos, 1992) can be classified as fitting either the " shrinking core model", "grain model", "capillary pore model", or the "random pore model".

4.2.2.1 Shrinking Core Model (SCM)

The reaction rate of particles which are of unchanging size may be controlled by one of the following steps: fluid film diffusion, product layer diffusion (converted solids and/or inert material), or the chemical reaction at the surface of the core of unreacted material. In the case of changing particle size, if the unreacted core is not covered by a product layer, there could only be two possible controlling steps, namely, fluid film diffusion and chemical reaction.

Initial Partly Completely unreacted reacted reacted particle particle particle

Final particle is hard, firm, and unchanged in size

Initial unreacted particle

Particle shrinks o with time, finally disappearing.

Flaking ash or gaseous products cause shrinkage in size

Figure 4.2 Systematic models for progressive-conversion and shrinking core reactions (Levenspiel, 1972) 88

The progressive-conversion (unchanging particle size) reaction and shrinking core (changing particle size) reactions are shown schematically in Figure 4.2.

4.2.2.1.1 Product Layer Diffusion Control

When diffusion through the fluid film presents a negligible resistance to the progress of reaction compared with diffusion through the product layer and the rate of chemical reaction is relatively low, the overall rate is controlled by product layer diffusion as shown in Figure 4.3.

fluid film

Figure 4.3 A schematic shrinking core model controlled by product layer diffusion reaction 89

The concentration gradient through the product layer is given by (CA{b) - CA(a))/ (rc -ra) and the distance (rc -ro) is the thickness of the product layer.

Let us consider a reaction of the form:

tfT(y-) + bB(j) —> + dD{s) (4.4)

The rate of reaction in the case of product layer diffusion control may be expressed as:

CA t (4.5) 3 aro Abulk

where R : fraction of component B reacted a,b : stoichiometric ratios V : molar volume of component B De : effective diffusivity of A into the product layer ro : initial radius of the particle CAtuit : bulk concentration of component A

t : time

4.2.2.1.2 Fluid Film Diffusion Control

When the overall rate is controlled by diffusion through the fluid film, the relationship between time and conversion for a reaction (4.4) is given by the following equation: Considering the elementary reaction of the form:

(4.7)

The rate equation for reaction (4.7) can be written as:

(4.8)

where in equations (4.6) and (4.8) km : mass transfer coefficient for component A

4.2.2.1.3 Surface Chemical Reaction Control

When diffusion resistance through the product layer is small due to the formation of a highly porous product layer, the overall rate is controlled by the chemical reaction at the core surface. The relationship between time and conversion in the case of surface reaction control of equations (4.4) and (4.7) can be written in the form:

1-(1 -R) (4.9)

where k : rate constant n : order of reaction with respect to component A 91

An important consideration in the determination of the overall rate of reaction, is the shape of the solid, such as whether the particles are spherical, cubic, or pellets, which may in turn undergo changes during the reaction with a liquid. As the reaction proceeds, the unreacted core diminishes in size. As the surface area of the solid changes, then the rate changes correspondingly. Taking into consideration this change, it is possible to predict the overall rate of the process as a function of changing size.

Kinetic equations (Habashi, 1970) for particles of changing size have been derived for different geometrical shapes as follows:

kcn 1-(1-/?),/3 =—±a-t forspheres (4.10) r„PB

2kC" 1-(1-/?)1/3=-----— f for cubes (4.11) roP B

kc" 1-(1-/^)1/2 = —for pellets (4.12) roPB

where for equations (4.10), (4.11), and (4.12) R : fraction of the solid reacted pB : density of the solid

Since these rate equations are only directly applicable to systems that have mono-size particles, it is unlikely that they can be directly applied to systems that have a range of particle sizes. However, if the size of particles is within a narrow size range, the above rate expression can be applied to 92 quantitatively describe the reaction kinetics for applications in industrial design.

4.2.2.2 Grain Model

The grain model was systematically developed by Szekely and his co-workers (Szekely and Evans, 1970, 1971; Sohn and Szekely, 1972) for certain types of porous and non-porous solids. The grains are usually taken as initially non-porous but the product layer developing could be porous or non-porous. In a porous structure, the space generated between the grains makes up the porous network.

The reaction of non-porous solids can be divided into two steps, namely those controlled by external mass-transfer and those by chemical reaction, with an intermediate stage where both steps affect the overall rate. The characteristic equation for the grain model in the absence of external mass transfer effects is:

bkC" l/Fp (------d"L)( y = i-(\-R) (4.13) Fp Vp where Sp : external surface area of the pellet at a given time

Vp : volume of the pellet at a given time

Fp : shape factor

The shape factor (Szekely et al., 1976), Fp, applies for infinite slabs (Fp =

1), long cylinders (Fp =2), and spheres (Fp = 3). 93

In the case of a porous solid, the reaction rate constant, k, and CnA can be replaced by k' , and C^1)/2, respectively.

k ' is defined as:

(4.14) where Sv : surface area per unit volume De : effective diffusivity of the gaseous reactant in the porous solid

4.2.2.3 Capillary Pore Model

The principles behind the capillary pore model were first proposed by Peterson (1957). One of the major assumptions, is that the porous solid contains uniform and cylindrical pores (capillaries) with random intersections. The rate of reaction for such a system can be written in the form (Szekely et al., 1976):

* A * (4.15) b ’ dt b c/l where Sv : surface area per unit volume of the pellet r : radius of the pores s : porosity ps : molar density of the pore-free solid

b : a stoichiometric coefficient

When the overall rate is controlled by chemical kinetics, i.e. external mass transport and pore diffusion present a negligible resistance to the progress of 94 the reaction, Szekely et al. (1976) have derived equations that predict the conversion in terms of the development of the pore radius. A structural parameter, G, is a function only of ea and is the solution to the cubic equation as follows:

(^OG3-G + 1 = ° (4.16)

The relationship between time and conversion of the solid phase can be written in the form:

Eo 1 ; G-l-(//rc) R= ■; [(i + ) (------^.... )-l] (4.17) \-E„ G - 1 where initial porosity of the solid particle

tc : time constant of the uniform pore model

4.2.2.4 Random Pore Model

Unlike Peterson' formulation, the random pore model (Sotirchos and Burganos, 1988) has attempted to describe the geometry of capillary intersection and overlap and to analyze the evolution of the pore volume distribution and surface area with consideration of pore intersection. The random pore model has been extensively studied by Bhatia and Perlmutter (1980) and independently by Gavalas (1980) who assumed that the solid contains cylindrical capillaries, which are more or less, randomly located and thoroughly cross-linked, porous spaces. With these assumptions if the pore overlap is taken into account the rate expression (Bhatia and Perlmutter, 1981; Sotirchos, 1987; Sahimi, et al., 1990) is given by: 95

dR (1 -/?)[!- i/f\n(l-R)]^2 (4.18) ~dt (1-0

where the initial pore structure parameter, i//, is defined as:

y/ = (4.19) .S’2

Further the relationship (Su and Perlmutter, 1984) between time and conversion can be written in the form:

k/S„t R = 1 - exp -Kr'Sj 1 + y/ (4.20) /_ where for equations (4.18), (4.19), and (4.20) ks : rate constant for the surface reaction So : reaction surface area per unit volume at / = 0

Ci : concentration of A at the reaction surface

s0 : porosity at / = 0 L0 : total pore length of a non-overlapped system at / = 0 r* : reaction rate per unit surface area

In this chapter, reaction kinetic models have been presented and discussed. More specifically, the model which was found to best describe the dissolution pattern of gold in the chloride/hypochlorite system is the shrinking core model for chemical reaction control at the surface of gold. From analysis based on the shrinking core model, this will be discussed in more detail in Chapter Seven. CHAPTER FIVE

FUNDAMENTALS OF ELECTROCHEMISTRY AT ROTATING DISC ELECTRODES

5.1 Introduction

The rotating disc electrode (RDE) is an important electrochemical technique which is often used to study inorganic oxidation-reduction processes. A common application of the RDE system is in the analysis of diffusion coefficients, bulk concentrations and the study of moderately fast electrode reactions. Levich in 1942 was the first to study the kinetics of the rotating disk electrode (Levich, 1962). Since then, considerable work has been carried out to evaluate the theory of this electrode system. Most of these studies accept the general Levich theory used for the study of electrode processes and mass transfer in electrochemical systems.

5.2 Rotating Disc Electrode System

Electrochemical measurements are carried out using a typical three electrode polarographic system. Generally, the three electrodes can be classified as the working electrode (WD), the counter electrode (CE), and the reference electrode (RE). The primary polarographic reaction at the metal-solution interface takes place at the working electrode. Gold, platinum, and glassy carbon are commonly used for the preparation of working electrodes. The 97

term "counter electrode" refers to the electrode which is counterpart to the working electrode. The counter electrode can be characterised by the following:

• an inert electrode, which does not take part in the electrode reaction, dipping directly into the test solution, • a shielded electrode with the compartment separated from the test solution by fine porosity or fritted disk, • a conventional half-cell with a salt bridge probe.

Although the counter electrode serves to complete the electrochemical cell, its potential has no importance to the system. The potential between the working electrode and reference electrode is usually measured. A Luggin capillary, housing the reference electrode, is positioned close to the surface of the working electrode to minimize IR drops in the system. The saturated calomel electrode (SCE) is widely used in experimental and industrial pH measurements, as a reference electrode. The potential of SCE is taken as 0.245 V on the SHE scale. It consists of a pool of mercury covered with a mixture of calomel, Hg2Cl2, and is immersed in a KC1 solution which acts as the electrolyte. This system is expressed as,

Pt/Hg/Hg2Cl2, KC1(C =...... )// (5.1)

The potential of the electrode depends to some extent on the temperature but primarily on the strength of the potassium chloride solution.

A typical three-electrode system for electrochemical measurements is represented in Figure 5.1. Levich (1962) has extensively studied the full 98 details of the hydrodynamic treatment of the RDE system and its application to electrochemical reactions.

Power supply

Working Auxiliary electrode electrode Ewk vs- ref

Reference electrode

Figure 5.1 A typical three electrode system for electrochemical measurements (Bard and Faulkner, 1980)

5.3 Electrochemical Considerations

For a rotating disc electrode, the overall reaction could be controlled by the diffusion of the electroactive species and electron transfer simultaneously takes place at the metal-solution interface. The major advantages of the RDE system include the following:

• a simple relationship between the limiting current density (iL) and the angular velocity of the disc (co). Enabling a quantitative analysis of the i-E curve obtained at a RDE surface, 99

• a simple RDE theory can be assumed for the expression of a more complex pathway without additional data, relying either on published theory or an experimental proof of the existence of an intermediate species.

Assuming a steady-state condition, the expression of the current density, /, can be written as:

i_ JL+1 (5.2)

where ik : kinetically limited current density iL : diffusion controlled current density

The simplified Butler-Volmer expression for the kinetically limited current density, neglecting back reaction kinetics, is:

(5.3)

where i0 : exchange current density n : number of electrons involved ^ : charge transfer coefficient F : Faraday constant, 96,485 coulomb/equivalent E: electrode overpotential R : gas constant = 8.314 Jmol^K'1 T: temperature 100

For a RDE, the limiting current for a reaction controlled only by diffusion can be evaluated by using the "Levich equation":

i, = 0.62 nFACbD2'3 vV2 (5.4)

where i, : limiting current, mA 3 Cb : bulk concentration of the electroactive species, moles/cm 2 A : area of the disc electrode, cm D : diffusion coefficient of the electroactive species, cm /sec co : angular velocity of the disc given by: co — 27iN rad/sec, with N = rps n : equivalents of charge/mole 2 v : kinematic viscosity of the solution, cm /sec

By using the relationship between iL and co , the diffusion coefficient for the electroactive species can be determined. The RDE offers the advantages of determining the diffusion coefficient in a reasonably short time and the associated theory allows for the checking of the experimental limiting current density with theoretical values.

The theory of the RDE applies to a thin plane surface, so large in diameter that the effect of edges can be reduced to a negligible level. In the quantitative study of electrode kinetics by the RDE, the Levich equation assumes that laminar flow exists near the disc surface. The rate of mass transfer to the RDE depends on the RDE rotational speed. At high rotating speeds the fluid flow at the surface of the RDE becomes turbulent and the Levich equation which is based on laminar flow conditions is not applicable. 101

Under laminar flow conditions, Reynolds numbers, vVRc, can be evaluated to be in the range from zero to about between 104 and 105 and are defined by:

r1 co N (5.5) Re v

where r is the radius of disc electrode (cm).

Moreover, the thickness of the diffusion boundary layer, <5, based on the laminar flow at the RDE can be written as,

S = 1.61 Z^W2 (5.6)

5.4 Electrode Kinetics and Reaction Mechanism

5.4.1 Determination of the Rate Constant

According to the general Levich law (Levich, 1962), if a first-order electrode reaction is controlled both by mass transfer and charge transfer, the value of / is expressed as,

nFADCb (5.7) 1.61D,/W,/2 +[Fj

By applying the thickness of the diffusion boundary at the RDE under laminar flow conditions, equations (5.6) and (5.7) may be combined to give the following expression: 102

nFADC (5.8)

where k is the heterogeneous rate constant at the given potential.

If S = 10 3 cm, then for values of k > 10 1 cm/sec, the limiting current of a diffusion-controlled process at the RDE is independent of the ratio D/k.

On the other hand, if 5 = 10'3 cm and k < 10’4 cm/sec, then the limiting current is independent of the thickness of the diffusion boundary layer and equation (5.8) can be written in the form:

/ = tiFAkCh (5.9)

From equation (5.9), the current is only controlled by charge transfer. The rate constant (k) at a given potential can be calculated from equations (5.7) and (5.9).

Moreover, the evaluation of the rate constant for the reaction involves a number of successive approximation techniques and graphical methods. Equation (5.7) can be rewritten in the form,

-i (5.10) knFAC Au' OmnFAD-> v~6C 103

A plot of il vs cSX11 would result in a straight line. Furthermore, it is possible to estimate the rate constant (k) by analysing the slope and the intercept of this line.

5.4.2 Determination of the Reaction Order

The order of the electrochemical reaction can be evaluated from RDE measurements. The order for the charge transfer-controlled reactions will depend on the relationship between the / - a}'1 curves. Alternatively, the order can be obtained by the following relationship:

1-(Vi5i / CbD) /win (5.11) ^-(j2^2/CbD)

where

m : order of reaction

j : 0.62£>2/VvV/2Ct = —s flux nFA

8: 1.61 £vV'\r1/2

A value for the order of reaction can be determined by measuring the current at two different rotational velocities and evaluation of the various parameters

(Levich, 1962).

5.5 Interpretation of Current-Potential Curves

Current-potential curves for a rotating disc electrode can be analyzed in terms of the Nernst-diffusion-layer concept, assuming steady-state conditions.

Generally, a metal corrosion process (Bancroft, 1937; Petrocelli, 1951; 104

Potter, 1956; Bockris et al., 1961; Milazzo, 1963; Marathe and Newman, 1969; Bockris and Reddy, 1970) presents an interesting pattern for the experimental current-potentials plots. Figure 5.2 shows a typical current- potential curve for dissolution and passivation for a rotating disc electrode.

Current density

Corrosion Passivation

Possivolion current Decreasing density dissolution

Flade potential

Increasing dissolution

Passivation I Potential potential Negligible dissolution currents

Figure 5.2 A typical current-potential curve for a rotating disc electrode (Bockris and Reddy, 1970)

The curve shifts to higher anodic potentials with increasing dissolution current as the potential becomes more positive. However, at a certain potential, the current-potential curves change direction and decrease rapidly with increasing potential. At the maximum potential in the current-potential curve (or "passivation potential"), the dissolution of the metal ceases and the metal becomes passive (Cheh and Sard, 1971; Gonzalez-Dominguez, 1994). At this point the surface of metal is stable. A typical current-potential curve (Eisenberg et al., 1964) of the electrochemical reaction for the rotating disc electrode is presented in Figure 5.3. The curve can be divided into three 105 regions, the "plateau region", "Tafel region", and the "mixed control region". In the plateau region, the current is independent of the potential. This means that electron transfer is not the rate determining step. In the Tafel region, the magnitude of the current changes rapidly with a relatively little change in the potential.

Plateau region Mixed control region

Tafel region

Figure 5.3 A typical current-potential curve of an electrochemical reaction

In the Tafel to plateau, transition region, the current is controlled by both the electrode kinetics and the mass transport at a certain potential. This region is referred to as the mixed control region. When the anodic and cathodic reaction occur on a heterogeneous electrode interface, the overall reaction can be analyzed, by reference to the concept (Paunovic and Vitkavage, 1979) of a steady-state mixed potential (Em). The schematic current density- potential diagram of a hypothetical electrochemical reaction is presented in Figure 5.4. As can be seen from Figure 5.4, the anodic and cathodic currents are equal at the mixed potential(or rest potential), E„„. The position 106 and shape of these curves depends on the composition of the electrolyte, sweep rate, electrode material, composition of the solution, and concentration of reactants. Thus, at the steady- state mixed potential, the rate of deposition is equal to the rate of oxidation.

+ z,e

+ z,o — M

Figure 5.4 A schematic current-potential curve by the concept of a steady- state mixed potential (Miller, 1981)

The overpotential, 77, depends on the electrode potential difference between the potential (E) when there is a net current density and the equilibrium potential (E0,m) when the current density is zero. It is given by:

n = E-E0M (5.12)

When the overpotential is less than zero, the cathodic half-cell reaction takes place, whereas when the overpotential is greater than zero, the anodic half­ cell reaction occurs (Barnartt, 1959). The exact position of the mixed potential is determined by the fact that the sum of cathodic current densities 107 is equal to the sum of the anodic current densities for a system at steady state. As the reaction proceeds, any change in the concentration of the reactants will change the current. This correspondingly means that the half-cell current-potential plots will change and the mixed potential will shift.

5.6 Evans Diagram

The current-potential behaviour has been studied extensively to investigate an electrochemical reaction kinetics for the respective half-cells involved in the reaction. The graphical super-position of the anodic and cathodic current- potential curves using the mixed potential theory is known as the Evans diagram. A typical polarization curve for the metal-dissolution is presented in Figure 5.5.

onodlc

Tafel slopes

cathodic M —-M

diffusion

log i

Figure 5.5 A typical polarization curve for metal dissolution in terms of an Evans diagram (Miller, 1981) 108

Bockris and Reddy (1970) demonstrated that corrosion reactions could be represented in an Evans diagram. In another work on the kinetics of cementation reactions by Power and Ritchie (1981), they reported that Evans diagrams applied to the data obtained from anodic and cathodic polarization measurements.

Since then, attempts to explain the mechanism of the short-circuited cementation reaction and other electrochemical reactions have been made by several investigators (Bancroft and Magoffin, 1935; Delahay, 1954; Simpson et al., 1955; Stern, 1955; Schmid and his co-worker, 1962, 1963, 1964 and 1967; Chiu and Genshaw, 1968 and 1969; Takamura et al., 1970; Miller, 1981; Miller and Wan, 1983; Tadayyoni et al., 1986; Stremsdoerfer et al., 1990; Cook et al., 1990; Creus et al., 1992). Tran and Davis (1992) have recently proposed the use of the Evans diagram which additionally combines the dissolution rate of gold. According to their work, the dissolution of gold in an iodide-iodine solution is in the kinetic-controlled region. They also found die potential difference for the gold dissolution to be lower than the typical value of 360 mV for a diffusion-controlled system.

An idealized Evans diagram which represents an electrochemical reaction is illustrated in Figure 5.6. From this diagram, the cathodic branch can be divided into three regions:

• At low current densities, the potential is independent of the current density and equivalent to the reversible electrode potential. • At an intermediate current density, the potential changes linearly with log /. This region of the potential-current curve is referred to as the Tafel region and the slope of this curve is known as the Tafel slope. 109

• At extremely low potentials, a limiting value of the current which is called the limiting current is reached and the current density remains constant.

,36V —Em (,,QnsMlon)

log i/i0

Figure 5.6 An idealized Evans diagram by the concept of the rate-limiting step (Miller, 1981)

In a similar way, the anodic behaviour of the system can be described using a potential-current curve. In the case of the anodic curve, the limiting current (Wilke et al., 1953) is reached only when the metal surface becomes saturated with respect to a salt of the metal. From the linear region of the potential-current curve, the Tafel equation is given by:

b = x±y\0Qi (5.13)

where x and y are constants 110

The current at the point where the polarization curves intersect the mixed potential, determines the rate controlling step. The criterion for the reaction to be mass transfer controlled is that the intersection of the polarization curves occurs in the diffusion-limited region, whereas that for a electrochemically controlled reaction is their intersection in the Tafel region or above it. Study of the intersecting point of the polarization curves results in the following:

• the determination of the rate controlling reaction, • the analysis of the kinetics of the short-circuited cementation reaction, • the theoretical maximum rate for the anodic and cathodic reactions, • the factors that influence the mixed potential.

Moreover, by considering the Faraday Law, the theoretical maximum rate of dissolution of metal related to the current density (Dorin and Woods, 1991) can be expressed as:

RA (5.14)

where 2 Ra : rate of dissolution of the metal, mole/cm s 2 Id : current density, A/cm F : Faraday constant, 96,485 coulomb/equivalent n : equivalents of charge/mole of metal dissolved Ill

5.7 Butler-Volmer Equation

The Butler-Volmer equation by using the basic assumption of absolute reaction-rate theory can be applied to electrolytic as well as electrochemical reactions. The reaction rate with respect to the cathodic (reverse) and anodic (forward) directions of the half-cell reaction (from Figure 5.4) may be given by the Butler-Volmer equation.

The generalized anodic reaction can be represented by:

M, = M;z° + zae~ (5.15)

While, the cathodic reaction can be expressed as:

M+2z‘ +zce~ = M2 (5.16)

If za = zc, the overall reaction is:

mx + m;z‘ = m;2* +m2 (5.17)

The Butler-Volmer equation for reactions (5.15) and (5.16) is given by:

^ My ~= n.Fk.[My\ex.-naFk_a[M^]exp^ (5-18) and -(1 -i>c)zcFEc ]exp( ^^j-ncFk_c[M*!-]exp (5.19) RT J 112

where, respectively /„, ic : current densities of the anodic and cathodic cell reactions, nc : total charge transferred, ka, kc : forward rate constants of the anodic and cathodic reactions, k_a, k_c : backward rate constants of the anodic and cathodic reactions, (j>a, c : transfer coefficients, (j)a = c = 0.5, za, zc : number of electrons transferred, za — zc — \ and Ea, Ec : anodic and cathodic potentials.

The anodic and cathodic reactions, neglecting back reaction kinetics, can be rewritten in the following form:

(5.20) and

= -ncFk_c[M^]e^ (1 • j (5.21)

At the mixed potential, i.e. i(Em) = 0 or -i.c(Em) = ia(Em), the reactions are combined arriving at,

M,} exp FEm) k° (5.22) RT ) RT 113

For symmetrical barriers, the transfer coefficients ( „ and (j)c ) equal 0.5.

Under these conditions the equation reduces to,

exp (5.23) V rtJ

Using the equation (5.14), the rate of the overall reaction in equation (5.17) is given by:

rdn\ (5.24) \

Ac : total area of the electrode of the cathodic cell

IM : current of the anodic cell reaction

By combining equations (5.21) and (5.24) at the mixed potential, the following expression can be derived.

(5.25) with

• k r = a Zo CHAPTER SIX

SUMMARY OF LITERATURE SURVEY

AND

RESTATEMENT OF MAIN AIMS

According to the previous literature review, the major concerns are summarised as follow:

1) Cyanidation has been used extensively in the carbon-in-pulp process for the extraction of gold from a wide range of oxidized ores at a relatively low cost. However, cyanidation is not suitable for ores containing cyanicides, such as copper oxides or silicates, as significantly higher cyanide consumption is experienced during gold processing. Because of the low extraction experienced for these ores and increasing awareness of environmental pollution, alternative gold recovery processes using more powerful lixiviants and non-toxic reagents for gold extraction have been considered.

2) The processes of direct gold leaching, such as alkaline pressure cyanidation, carbon-in-leach (CIL) and carbon-in-pulp (CIP), for the treatment of refractory ores have been applied over the past decade. Roasting, biological oxidation and chemical oxidation as pretreatment steps for refractory ores have also been evaluated. 115

3) Recently, the recycling of previously-used metals has assumed greater importance and the feed for leaching can be obtained from secondary sources such as electronic scraps, slags, dusts, residues, and wastes from other processes. The recovery of precious metals from electronic scraps has attracted the interest of many researchers since the mid- 1980s.

4) Several lixiviants, such as halides, thiourea, thiosulphate, thiocyanate and poly sulfide, can be used for gold extraction. Although these lixiviants have been examined mainly in an attempt to treat complex metal sulfide ores and concentrates over the last few years, none has yet been commercialized on a large scale. Among these, the use of halides has assumed greater importance because the rate of gold dissolution with halides is much faster than with other lixiviants.

5) Chlorination has been re-introduced as a pretreatment process to oxidize the carbonaceous material and sulfide minerals of certain gold ores to liberate the gold prior to cyanidation. The technique has been successfully employed at Carlin Mines since the early 1970's. Such a process of chlorine pretreatment has been proposed for the heap leaching of low grade gold ores and in in-situ gold extraction systems.

6) Generally, chlorination is not an economical process to treat ores which contain high levels of sulfide, due to the excessive consumption of chlorine. For this reason, the efficient leaching of gold in an acidified chloride system is restricted to the treatment of ores containing a low sulfide content, typically less than 5%. Gold is extracted from such ores and concentrates by employing one of several 116

alternative processes, namely pre-aeration, pressure oxidation or roasting.

7) In Australia, processes utilising chlorine pretreatment could have a clear advantage over cyanidation. This is because gold plants are able to electrolytically produce chlorine economically from ground water brines, which are common in some gold fields of Western Australia where natural saline waters are available. Major problems of the reprecipitation of gold from chloride solutions and the corrosiveness of acidic chloride solutions are main concerns for the application of such a system.

8) Basic principles of voltammetry on rotating disc electrodes have been reviewed with the view to using these techniques to study the electrochemical dissolution of gold and reduction of hypochlorite in chloride solutions. Evans diagrams can be derived to determine the initial reaction rate by analysing the relationship between the mixed potential and current density of the system.

In general, major problems of the corrosiveness of acidic chloride solutions and the reprecipitation of gold from chloride solutions are main concerns for the industrial applications of chloride/hypochlorite system. Additionally, relatively little is known about the behaviour of gold chloride complexes in the presence of sulfide minerals. For this reason, chlorination has been limited in its application and information on its process chemistry is limited. Thus, the present study was undertaken to clarify the above problems and to develop a reaction kinetic model to describe quantitatively the dissolution of gold in chloride solutions, with the aim of achieving the following objectives: 117

1) Investigate the effect of solution pH, temperature, chloride and hypochlorite concentrations on gold dissolution in acidic chloride solutions,

2) Determine the reaction kinetics and the rate equation describing the dissolution of gold in chloride solutions,

3) Develop the equation for the thin rectilinear shrinking core model and, by using this model as a basis, to propose the rate equation,

4) Investigate the stability of the gold chloride complex in chloride solutions in the presence of sulfide minerals such as sphalerite, galena, pyrite and chalcopyrite,

5) Determine the effect of pH and chloride concentration on the current- potential relationship for gold dissolution and to establish the mixed potential and current density required for gold dissolution in chloride solutions of different compositions,

6) Develop the relationship between the mixed potential and current density by using the Butler-Volmer equation,

7) Determine the rates of dissolution, from several sulfide electrodes,

8) Use the computer modelling program to identify the predominant species in the gold-chloride system and gold-sulfide systems and hence establish the chemistry of chloride leaching of gold with and without sulfide minerals. CHAPTER SEVEN

STUDY ON THE KINETICS OF GOLD DISSOLUTION AND THE STABILITY OF GOLD SPECIES IN A CHLORIDE SYSTEM

7.1 Introduction

A kinetic study which included measurements of the gold dissolution rates and the development of a modified rate equation for the dissolution of a thin rectilinear shape in chloride-hypochlorite solutions was undertaken. Tests to determine the stability of gold chloride solutions in the presence of gangue minerals such as sphalerite, galena, pyrite and chalcopyrite were also conducted.

The objectives of the gold dissolution and stability studies are:

• to determine the effect of pH, chloride concentration, hypochlorite concentration, and temperature on the dissolution of gold in acidic chloride solutions, • to determine the reaction order and the rate equations describing the dissolution of gold in chloride solutions, • to develop an appropriate equation for the thin rectilinear shrinking core model, • to fit experimental data into a relationship between surface concentration and initial particle size and to propose suitable rate equations by using the developed shrinking core model, 119

• to identify the stability of the gold chloride complex in chloride solutions in the presence of gangue minerals such as sphalerite, galena, pyrite and chalcopyrite.

7.2 Experimental

7.2.1 Materials

7.2.1.1 Chemicals

For the dissolution tests, gold strips of 99.99% high purity grade from Johnson Matthey, Australia were used. The dimensions of all gold strips before and after the experiments were measured by using a vernier calliper and a thickness gauge with an accuracy of ± 5 x 10 4 cm and ±5 x 10 6 cm units, respectively..

The following chemical reagents were used in the experimental investigations: Sodium chloride : (AJAX) Analytical reagent, assay (on dried basis) 99.9% minimum NaCl. Sodium hypochlorite : (AJAX) 5 litres volumetric solution. Available chlorine: 125 g/1 when packed. Potassium iodide : (BDH) Analytical reagent, assay (on dried basis) 99.8% minimum KI. Thiosulfate : (BDH) Analytical reagent, 0.1 N Na2S203. Silica : Technical grade (60G), assay 99.0% Si02, 0.03% Fe203. Sodium hydroxide and sulfuric acid of analytical-grade quality were used to adjust the solution pH and distilled-deionised water was used to prepare the 120 solutions. The concentration of hypochlorite (OC1 ) was determined by the

Kl-thiosulfate titration method (Sherman and Strickland, 1955).

7.2.1.2 Sulfide Minerals

High purity mineral samples of sphalerite, galena, chalcopyrite and pyrite used for the stability tests were supplied by BHP Minerals, Australia. The mineral specimens were initially crushed into particles of several millimetres in diameter. The predominant sulfide mineral fractions were hand-picked to obtain higher purity samples and the selected samples were ground in a ring- disc pulverizer. The ground samples were screened and the size fraction of -

210 +150 pm was used for the experiment. Representative samples from each sulfide mineral were taken, further ground to minus 38 pm and examined by a Simens D 5000 X-Ray Diffractometer (XRD) and its chemical composition also determined by solution Atomic Absorption Spectrometer

(Varian model SpectroAA-20). The XRD patterns for the mineral samples used are represented in Figs 7.1 (a-d).

1000 <1 • Sphalerite 'c 800 13 Almpurities

£ 600 «1

L. 03 400

c 1» Q) 3= 200

• • • A * 1 1 1 L...... , ill A* It 0 1 1 20 40 60 80 100

2 theta

Figure 7.1 (a) X-Ray Diffractometry pattern of the sphalerite sample Figure Intensity (arbitrary unit) ^ Intensity x 1000 (arbitrary unit) re 250 200

150 4 5 3 100 2 0 1 7.1 7.1 20

50 0

20

(c) (b) -

▲ X-Ray X-Ray 1 © u ► 1 > 40

• Diffractometry Diffractometry

40 i • 60 60 • ▲

2

theta 60 2 •

pattern pattern theta 1

of of

the the a •

80

pyrite galena Impurities Pyrite a • 80 100

I

Galena • Impurities

A sample sample

100 121 122

& 600 • Chalcopyrite a Impurities

£ 400

20 40 60 80 100 2 theta

Figure 7.1 (d) X-Ray Diffractometry pattern of the chalcopyrite sample

Based on the information given by X-ray diffractometry and chemical analysis on the size fraction of the sulfide samples used, an approximate mineralogical composition of the samples is determined ( Table 7.1).

Table 7.1. Estimated mineralogical composition of the sulfide minerals

Ore Major minerals Impurities Sphalerite ZnS Pyrrhotites Galena PbS Pyrite Pyrite FeS2 Chalcopyrite, Quartz Chalcopyrite CuFeS2 Pyrrhotite, Quartz 123

7.2.1.3 Gold Chloride Solution

A gold strip of 99.9% purity was weighed on a micro-balance (Mettler E) with a detection limit of 0.001 mg. The gold strip was immersed in 200 g/1 NaCl solution for 5 hrs to prepare the gold chloride stock solution. To accelerate the dissolution of the gold strip in the chloride solution, an anodic electric current was supplied to the Au strip using a laboratory DC power supply (Voltage: 4V; Current: 0.2A). A graphite electrode was utilised as a cathode to complete the electrolytic cell. The difference between the initial and the final weights of the Au strip was compared with the analysis of the solution by AAS to finally determine the assay of the stock solution. The stock solution was diluted to the desired concentration in subsequent experiments.

7.2.2 Equipment

7.2.2.1 Experimental

Two different reactors were used in this study. Dissolution of the gold leaf was carried out in a 500 ml conventional glass reactor placed in a glass water bath set at the test temperature (Figure 7.2). The gold strip was immersed in NaCl-NaOCl solutions. The system was air sealed and samples of the solution were withdrawn regularly for analysis.

To test the effect of sulfide minerals, ore particles were placed in a 500 ml conventional glass reactor, equipped with a mechanical agitator, a pH probe and a condenser in a temperature-controlled water bath (Figure 7.3). All kinetic experiments were conducted at constant temperature (20°) and at an agitation speed of 300 rpm. 124

7.2.2.2 Gold Leaf Tests

For studying the kinetics of gold dissolution, gold strips of known surface areas, which were pre-weighed on a micro-balance (Mettler E) having a detection limit of 0.001 mg, were immersed in NaCl/NaOCl solutions of known compositions for 60 minutes. For each run, 250 ml of liquor was first introduced into the reactor. Samples of approximately 2 mis were withdrawn at predetermined intervals over 60 mins and the Au concentration was analysed by using a Varian SpectroAA 20 Atomic Absorption Spectrometer (AAS). The difference between the initial and the final weights of the Au strip allowed for a comparison with the gold analysis of the solution by AAS.

A. Magnetic stirrer with E. pH probe temperature control F. Glass water bath B. Gold strip G. Water pump C. 500ml glass reactor H. Water bath for D. Thermometer temperature control

Figure 7.2 Schematic representation of gold dissolution experiment using gold strip. 125

water out

water in

A. Mechanical agitator with variable speed B. Condenser C. Water bath with temperature control D. 500ml glass reactor E. pH probe

Figure 7.3 Schematic representation of the experiment testing the stability of gold chloride complex in the presence of gangue minerals. 126

7.2.2.3 Stability of Gold Chloride Solutions

Weighed samples of sphalerite, galena, chalcopyrite and pyrite were reacted separately with a solution of gold chloride to determine the reactivity of the sulfide minerals. For each run, lg of sulfide mineral, 250 ml of liquor and 50 g of silica (unless otherwise stated) were introduced into the reactor. Samples of approximately 3 mis were withdrawn at pre-determined time intervals over 1 hour and analysed for Au by AAS using a modified solvent extraction method (Lockyer and Hames, 1959; Gatehouse and Willis, 1961; Khalifa et al., 1965; Strelow et al., 1966; Groenewald, 1968; Zlatkis et al., 1969; Hildon and Sully, 1971; Mooiman and Miller, 1986; Miller et al., 1987; Kordosky et al., 1992 and 1993).

7.2.2.4 Analysis of Gold Chloride Solutions by Solvent Extraction

Difficulties during analysis (Guy et al., 1983) were encountered due to the presence of excess sodium chloride ion (200g/l NaCl). To prevent the interference of NaCl in the determination of the gold concentration, the following procedures were carried out : 1 ml of auric chloride in 200 g/1 NaCl solution was diluted with 9 mis of distilled-deionised water and then 0.5 ml of distilled-deionised water, 0.5 ml of 40% w/v HC1, and 4 ml of 1 % w/v Aliquot 336 in di-isobutyl ketone (DIBK) were added. The whole sample was shaken for 1 minute. After shaking, the solution was left to stand for 1 day in order for the two phases to settle. The organic layer was separated from the aqueous phase wherein the gold concentration was determined directly by AAS. 127

Table 7.2 Determination of gold in 200g/l NaCl liquors by solvent extraction.

Au, ppm Test No. Micro-balance Solvent extraction Ratio (A) (B) (B/A) 1 1.79 1.76 0.98 2 1.29 1.38 1.06 3 1.57 1.65 1.05

The comparative analysis of the solution is shown in Table 7.2, indicating that the two techniques yield results within ± 5 % of each other.

7.3 Kinetics of Gold Dissolution

A series of experiments was carried out to determine the effect of various parameters on the kinetics of gold dissolution in a chloride system. The following parameters were investigated:

i) pH ( 2, 3, 4, 5 and 6 ), ii) sodium chloride concentration ( 50, 100, and 200 g/1), iii) hypochlorite concentration ( 1, 2, 5 and 10 g/1 ), and iv) temperature (22, 26, and 35°C ).

To get the slope of the linear relationship, the regression line known as the "line of best fit" was drawn for each set of results, so as to fit this line as closely as possible to the data. The width of confidence limits is dependent on the correlation coefficient (r). The correlation coefficient (refer to Appendices 7.1 and 7.2) includes measurement error and/or changes in other 128 variables which were not controlled. The values of the dissolution rate, reaction order, activation energy, etc, were calculated from the slopes of each straight line by using linear regression.

7.3.1 Effect of pH

The dissolution of gold by chloride-hypochlorite solutions was dependent on the solution pH (Cho, 1987a and b). The effect of pH on gold dissolution using solution of 100g/l NaCl and 10 g/1 OCT was examined in the pH range from 2 to 6 as shown in Figure 7.4. The results show that the two characteristic pH regions present during gold dissolution with hypochlorite in the pH = 2-6 range. Gold dissolution rate decreased slightly with pH decrease in the range 2 to 3. This was probably the result of limited solubility of Cl2. However, above these values, the gold dissolution rate decreased dramatically with pH.

Time, mins

Figure 7.4 Effect of pH on gold dissolution in solutions of 100g/l NaCl and ocr iog/i. 129

pH

Figure 7.5 Gold dissolution rate as a function of pH for several concentrations of OC1 in 100g/l NaCl .

The dissolution rate of gold as a function of pH for several concentrations of OCT is presented in Figure 7.5. Two characteristic regions were seen where the system behaved differently, at lower pH (2 and 3) compared with higher pH (4, 5 and 6) for a solution of 10 g/1 00 concentration in this system. The rate of the dissolution of gold increased with a pH increase in the range 2-3. The change in the dissolution rate of gold at pH 4 suggests a change in the dissolution mechanism as shown in Figure 7.6. At a higher pH in this system, the gold dissolution rate decreased. This is due to a decrease in the dissolved hypochlorous acid concentration as the pH increases. Furthermore, with increasing NaCl concentrations, the maximum gold dissolution rate shifts toward a higher pH. The shift of maximum rate to increase the NaCl concentration is in very good agreement with the shift of predominant domain of AuC14’, as shown in Figure 9.2 (a)-(c), obtained from thermodynamic study. 130

ONaCl 50 g/1 -A-NaCl 100 g/1 -&NaCl 200 g/1

pH

Figure 7.6 Gold dissolution rate as a function of pH for several concentrations of NaCl in 10g/l OCF.

In all these investigations, the dissolution rate of gold was measured for solutions of NaCl (50 g/1 to 200 g/1) and OCT (1 g/1 to 10 g/1) in the pH range from 2 to 6. These results are summarised in Appendix 7.1. At low NaCl concentration for gold chloride system, the gold dissolution rate is in good agreement with the observations of Nicol (1980a).

7.3.2 Effect of NaCl Concentration

The effect of NaCl concentration on gold dissolution in solutions containing 10 g/1 OCf at pH 4 was examined for NaCl additions ranging from 50 g/1 to 200 g/1, as presented in Figure 7.7. 131

e-NaCl 50 g/1 -±-NaCl 100 g/1 ■B-NaCl 200 g/1

Time, mins

Figure 7.7 Effect of NaCl concentration on gold dissolution in solutions of 10g/l OCT at pH 4.

It was found that at the initial pH 4, a NaCl concentration increase had a positive effect on the dissolution of gold. The dissolution rate as a function of NaCl concentration in the presence of 10 g/1 00 at several pH ranges is depicted in Figure 7.8. The data in Figure 7.8 indicate an increasing rate with decreasing NaCl concentration at pH 2 and 3. The slopes of these lines were -0.14 for pH 2 and -0.16 for pH 3, respectively. This is due to the decreased diffusion coefficients of gold chloro-complex by increasing the NaCl concentration (Kim et al., 1989; Puvvada and Tran, 1995). For pH 4- 6, the rate of the dissolution of gold increases with increasing NaCl concentrations. The slopes of theses lines were 0.22 for pH 4, 0.05 for pH 5, and 0.03 for pH 6. 132

At high pH and low NaCl concentrations, the decrease in the dissolution rate of gold can be explained in terms of the passivation of gold due to hydro­

oxide ion adsorption. This is because OH is more adsorbed than Cl for gold

(Bode et aJ., 1967).

OpH 2 ■ pH 3 A pH 4 □ pH 5 + pH 6

Figure 7.8 Dissolution rate of gold as a function of sodium chloride concentration at several pH’s for solutions containing 10 g/1 OCT.

7.3.3 Effect of OC1" Concentration

The effect of hypochlorite concentration on gold dissolution in solutions containing 100 g/1 NaCl is shown in Figure 7.9. These data confirm the first order kinetics indicating that an increase in OC1 concentration has a positive effect on the rate of gold dissolution. 133

e-oci- l g/i «-OCl- 2 g/1 -*-OCl- 5 g/1 -&OC1- 10 g/1

Time, mins

Figure 7.9 Effect of hypochlorite concentrations on gold dissolution in 100 g/1 NaCl at pH 4.

The rates of dissolution of gold versus 00 concentrations at different pHs for a solution of 200 g/1 NaCl are shown in Figure 7.10. The results in Figure 7.10 indicate that the addition of hypochlorite to the solution increased the rate of gold dissolution. At low pHs (2 and 3), the reaction rate increased for concentrations of up to 5 g/1 OC1" and did not increase further for the higher concentrations. The observed insensitivity of the reaction rate at higher OCF concentration and lower pH (2 and 3), is in good agreement with the result obtained for the rate of gold dissolution using chlorine by Sun and Yen (1992). This is due to the limited solubility of dissolved chlorine (to form OC1) at this low pH range. 134

O pH 2 H pH 3 A pH 4 □ pH 5 + pH 6

Figure 7.10 Dissolution rate of gold as a function of hypochlorite concentrations at several pH for a solution 200 g/1 NaCl.

7.3.4 Effect of Temperature

The effect of temperature on the dissolution of gold was investigated at three different temperatures using solutions of 200 g/1 NaCl and 10 g/1 OCf at pH

4 as presented in Figure 7.11. The experiments were not performed at temperatures higher than 35°C because of the high volatility and thermal instability of hypochlorite (Cho, 1987b; Qi and Hiskey, 1991). Increasing the temperature results in a faster rate of dissolution, as can be seen in Figure

7.11, within the temperature range of 22 to 35°C. The rate of the dissolution of gold as a function of temperature is presented in Figure 7.12. A linear relationship with a negative slope of 207.3 mg K/(cm2hr) was obtained. This study is consistent with the observations of Bayrakceken and co-workers (1990) in relation to the chlorination of pyrite. 135

Time, mins

Figure 7.11 Effect of temperatures on gold dissolution in 200 g/1 NaCl and 10g/l OCT at pH 4.

51 ioo

Slope = - 207.3

Figure 7.12 Dissolution rate of gold as a function of temperature for a solution of 200 g/1 NaCl and 10 g/1 OC1’ at pH 4. 136

7.4 Determination of The Reaction Order

7.4.1 Reaction Order with respect to Hydrogen Ion Concentration

Considering the hydrogen ion as an aqueous reactant that participates in the gold dissolution reaction, the rate will depend on its concentration, as shown:

roc [H+]n (7.1) where r is the rate of the dissolution of gold, and n the reaction order with respect to H+. The reaction order with respect to H+ for different OCf concentration levels is presented in Figure 7.13.

o OC1- 1 g/1 ■ OC1- 2 g/1 - A OC1- 5 g/1 □ OC1- 10 g/1

pH

Figure 7.13 The reaction order with respect to [H+] for a solution of 200 g/1 NaCl. 137

Two regions were seen to behave differently, pH 2 and 3, as compared to higher pH regions (4, 5 and 6) in this system. In the pH range 2 to 3, the reaction order with respect to the hydrogen ion concentration is independent of the OC1' concentration below 10 g/1. This can be explained in terms of the result of limited solubility of OCf. As shown in Figure 7.13, the values of the reaction order with respect to the hydrogen ion concentration in the pH range from 4 to 6 are 0.35 for 1 g/1, 0.30 for 2 g/1, 0.30 for 5 g/1, and 0.35 for 10 g/1 OCf concentration, respectively. The range of the reaction order with respect to the hydrogen ion concentration is between 0.30 and 0.35. This agrees with the half order dependency observed by Pesic and Sergent (1991) for the bromine leaching system.

7.4.2 Reaction Order with respect to Sodium Chloride Ions

The results for the reaction order with respect to sodium chloride concentrations are presented in Figure 7.14. Generally, the rate increased linearly with an increasing sodium chloride concentration in the higher pH range 4 to 6, according to the following relationship:

roc [NaCiy (7.2) where n is the reaction order with respect to NaCl.

The reaction order for gold dissolution had two characteristic regions at lower pH (2 and 3) and with higher pH (4, 5 and 6) in this system with respect to NaCl concentration changes. The reaction order declinedlinearly with increasing NaCl concentrations in the range between pH 2 and 3. This situation could be attributed to the decreased diffusion coefficients of gold 138 chloro-complex at high NaCl concentrations. Similar trend has recently been observed by Puvvada and Tran (1995) for silver chloro-complexes at an initial pH of 2.0 investigated by using a rotating disc electrode system.

Log NaCl, M

Figure 7.14 The reaction order with respect to NaCl in 10 g/1 OC1 and pH range 2 to 6.

The values of the reaction order obtained from the slopes of the lines in Figure 7.14 are -0.28 for pH 2, -0.30 for pH 3, 0.64 for pH 4, and 0.48 for pH 5 and 0.52 for pH 6. The reaction order in the pH range 4 to 6 was determined to be of almost a half dependency on the NaCl concentration. The values of the reaction order in the pH range 4 to 6 are in very good agreement with the half order dependency in iodide systems as found by Qi and Hiskey (1991).

7.4.3 Reaction Order with respect to Hypochlorite Ions

The reaction order with respect to OC1 (1-10 g/1) was studied in 200 g/1 NaCl concentration in the pH range of 2 to 6. The results are presented in 139

Figure 7.15. The relationship between reaction rate and hypochlorite concentration can be written as followed: roc [OCrX (7.3) where n is the reaction order with respect to OCT.

O pH 2 ■ pH 3 A pH 4 □ pH 5 + pH 6

Log [OC1-], M

Figure 7.15 The reaction order with respect to OCT in 200 g/1 NaCl for the pH range 2 to 6.

The values of reaction order of gold dissolution by OC1 concentration are 0.88 for pH 2, 0.93 for pH 3, 1.06 for pH 4, and 0.96 for pH 5 and 1.09 for pH 6. At pH 2 and 3, the reaction order was almost first order with respect to OCF at lower (1-5 g/1) concentrations and zero at higher (10 g/1) concentrations, all in 200 g/1 NaCl concentration. The reaction of first order in chloride system in this study is comparable to the values reported by Pesic and Sergent (1991) for bromine and Qi and Hiskey (1991) for iodine. As previously stated, however, the situation at higher OC1 (10 g/1) concentration and lower pH (2 and 3) should be attributed to the limited solubility of dissolved chlorine. 140

7.4.4 Activation Energy with respect to Temperature

Figure 7.16 shows a plot of the activation energy with respect to temperature (22 to 35°C) for a solution of 200 g/1 NaCl and 10 g/1 OCT at pH 4. By using an "Arrhenius-type" plot (Majima et al., 1985; Pesic and Sergent, 1990 and 1991; Qi and Hiskey, 1991; Filippou and Demopoulos, 1992), an apparent activation energy for temperature increments was determined to be 28.7 kJ/mole for a gold chloride/ hypochlorite system.

Ea = 28.7 kJ/mol

1/T, K1 x 103

Figure 7.16 Activation energy with respect to temperature in 200 g/1 NaCl and 10 g/1 OCP at pH 4

Such a value of the activation energy indicates a dissolution process would be expected to be controlled by a kinetic step (Wadsworth and Sohn, 1979). This reaction control is consistent with the results studied by Bayrakceken et al. (1990) for the oxidative dissolution of pyrite by chlorine within the temperatures range from 13 to 40°C. 141

7.5 Rate Equation for the Dissolution of Gold in the Chloride/Hypochlorite System

The mechanism of the dissolution of gold in chloride/hypochlorite solutions is dependent on solution of pH, NaCl and hypochlorite. The anodic half-cell can be presented as follows:

An +4CT —> AuCl^ +3e~ (7.4)

The cathodic half-cell is:

OCr +2H+ + 2e~ ->Cr +H20 (7.5)

The overall reaction of gold with the chloride/hypochlorite system in an acidic medium is governed by the equation:

2An + 30Cl~ + 6H+ + 5Cl~ -> 2AuCl4 + 3H20 (7.6)

The effect of important process parameters such as hydrogen ions, sodium chloride, hypochlorite concentrations, and temperature was studied. From the analysis based on an "Arrhenius-type" relationship, the activation energy was found to be 28.7 kJ/ mole for a solution of 200 g/1 NaCl and 10 g/1 OCT at pH 4. Furthermore, the reaction order with respect to the activity of H+ for a solution of 200 g/1 NaCl is almost 0.3 at the pH range from 4 to 6 and zero at pH 2 to 3. The reaction order with respect to NaCl for a solution of 10 g/1 OC1 is half, at the pH range from 4 to 6 and -0.3 at pH 2 and 3. The reaction order with respect to 00 is almost first order at pH 4-6 and zero 142 order in the pH 2-3 at higher OC1 concentration (10 g/1). Therefore, a generalised rate equation obtained from the results indicates that the dissolution process, as presented by Eq (7.6), is controlled by a kinetic step.

By taking [H+] = 10 4-106, [NaCl] = 50-200 g/1, [OC1] = 1-10 g/l, and E„ = 28.7 kJ/mol, the generalised empirical rate equation is as follows:

d[Au(lIl)]ldt = k expj- 287^ mo' \[H* ]° 3[M7C/]05[OCr ] (7.7)

7.6 Development of the Thin Rectilinear Shrinking Core Model

Consider a reaction of the form:

aA(j) +bB{f) + cCU) +dD{s) -> eEy) + JFs) (7.8)

If the shrinking core model with no product layer at the interface is adopted, the following rate equation can be derived (Levenspiel, 1972; Wadsworth and Sohn, 1979):

cJN, (7.9) dt where N A : moles of component A,

S : surface area of solid component D, [A]core,[B]co,e,[C]co;e : concentration of soluble components A, B and C at the interface, respectively,

p, q, r : order for soluble reactant A, B and C, k, : rate constant, 143

k o : concentration at reactive sites of solid species D, m : order of reaction with respect to the reactive sites.

when t - 0 when t = /

Figure 7.17 The rectilinear gold strip illustrating its dimensions specified for tlie derivation of the shrinking core model.

When a rectilinear gold strip as component D is used and its dimensions is specified in Figure 7.17, its surface area can be given as:

S = 2(- + - + — )X2 (7.10) where X: length of the rectilinear gold strip, X width of the rectilinear gold strip, X X thickness of the rectilinear gold strip, y x, y : constants, respectively. 144

By combining Eq (7.9) with Eq (7.10),

cJN, P -2(- + - + — )X2[A] core i (7.11) dt x y xy

Moreover, the fraction reacted, R, can be represented as (Chris, 1993):

R=l- (7.12)

where XQ : initial length of gold strip when t = 0 X : length of gold strip at time /.

The following relationship between the fraction reacted (R) and time (/) from Eq (7.12) can be expressed as follows:

dR_ 3X2 dX (7.13) dt X] dt

Also, if Nd are moles of solid component D, then:

i7j4> where, V is molar volume of component D.

By differentiation with respect to time t with Eq (7.14):

dND _ 3X2 dX (7.15) dt xyV dt 145

From stoichiometry of Eq (7.8)

By combining Eq (7.11) and Eq (7.16), we get:

dNp -2(-H-+-+— )X2[A]p (7.17) dt a x y xy core P1L.ICL.*;* ,

Substituting Eq (7.13) and Eq (7.15) into Eq (7.17) yields:

? = (^r)d-)V[A}l lB]qcore[C]rcoreKk,(x + y +1) (7.18) dt a X\

Using the value of X2/X] = (1 - R)2/3 from Eq (7.12) and rearranging, we can write:

h ‘"L, =1 (~)(~)V[A]L,[fl]L.[CL.(x+y+1 )dt (7.19) (i-y<) a0 ci

The boundary conditions required to solve Eq (7.19) are as follows:

R = 0 when t = 0 R = R when / = t 146

This equation can be easily integrated with boundary conditions:

with k = k"'kx.

To confirm Eq (7.20), this equation can be inserted with the values of x = 1 and y — 1 as applicable to cubes. Eq (7.20) is in good agreement with the proposed equation for cubes by Habashi (1970). This means that the proposed equation can be applied to thin rectilinear strips or cubes.

Furthermore, the following relationship between fraction reacted (R) and time (/) from Eq (7.20) can be written as:

1-(1-/01/3 =kj (7.21)

The apparent rate constant k„ (min *), which involves many factors, can be expressed as follows:

k„ = (-niAIL, IB] L, IQL, (x + y + 1) (7.22) 32f a

The thin rectilinear shrinking core model which has been developed by the author was used to quantitatively describe the dissolution of gold in this study. The dissolution of gold in chloride/hypochlorite solutions can be expressed as an electrochemical process. The addition of sodium hypochlorite to a solution results in its decomposition, generating the strongly oxidising hypochlorite species, as follows: 147

NaOCl -> Na+ + OCl~ (7.23)

The cathodic half-cell related to the reduction of hypochlorite ion to generate chloride ion is:

ocr + 2H+ +2e~ -> cr +H20 (7.24)

The overall reaction between Eqs (7.4) and (7.24) can be rewritten:

3ocr + 6H+ + 5Cl~ + 2An -» 2AuCl4 + 3H20 (7.25)

From Eq (7.20), since k[X\Pcore[B]qcore[C]rcore is exponentially dependent on temperature, the rate equation can be rewritten as:

1 - d - «),/3 = •'”'(*+y+(7-26) oXon a

If as in Eq (7.26), diffusion steps are not rate-determining, the concentration of A at the bulk is equal to that at the interface, i.e. [A] = [A\ = [A]

(same for B and C). The kinetic equation which incorporates the function of time t for the dissolution of gold is as follows:

1 - (1 - 7?)V3 =kc[A\PlB]q[C]re-{EJRT)t (7.27) with^K3t (x+y+^ 148

Using the overall reaction of gold with the chloride/hypochlorite system from Eq (7.25), the rate equation in Eq (7.27) can be expressed as follows:

1-(1-/01/3 = kc[OCr]p[H+]q[NaCl]r e~{E °,RT)t (7.28)

7.7 Reaction Kinetics using the Shrinking Core Model

7.7.1 Effect of pH

The effect of pH on the dissolution of gold in 200 g/1 NaCl and 10 g/1 00 solutions is shown in Figure 7.18. The fraction of gold dissolved at any time was seen to increase with an increase in pH.

O pH 2 ■ pH 3 A pH 4 □ pH 5 4- pH 6 ____

Time, mins

Figure 7.18 Plot of the fraction reacted as a function of time at different pHs in 200 g/1 NaCl and 10 g/1 OCf. 149

From the data in Figure 7.18, the plots of 1-(1-R)1/J versus time t are linear, as shown in Figure 7.19. The slopes of each of the lines in Figure 7.19 are apparent rate constants, ka, obtained from Eq (7.21). In all these investigations, the rate constant for the dissolution of gold was measured for

NaCl (50 g/1 to 200 g/1) and OC1 (1 g/1 to 10 g/1) in the pH range from 2 to 6, as summarized in Appendix 7.2. The apparent rate constant as a function of pH is presented in Figure 7.20. This plot shows the same trends as the results observed by Jin et al. (1984).

Time, mins

Figure 7.19 Plot of 1-(1-R)1/3 as a function of time in 200 g/1 NaCl and 10 g/1 ocr.

From the data in Figure 7.20, the reaction order with respect to H+ based on Eq (7.22) is presented in Figure 7.21. The value obtained from the relationship between ka and H+ is 0.3. This value is in good agreement with the reaction order previously determined from the kinetic study. 150

Figure 7.20 Plot of the rate constant (ka) as a function of pH in 200 g/1 NaCl and 10 g/1 OC1'.

Slope = 0.3

Log [IT]

Figure 7.21 The order of reaction with respect to H+ for a solution of 200 g/1 NaCl and 10 g/1 OCP. 151

7.7.2 Effect of NaCl

Figure 7.22 presents fraction reacted versus time plots for different NaCl concentrations (10 g/1 OCf at pH 4). As can be seen from this figure, in the

NaCl concentration range of 50-100 g/1, the increasing rate of gold dissolution is consistent with the observations of Nicol (1981a).

O NaCl 50 g/1 + NaCl 100 g/1 A NaCl 200 g/1

Time, mins

Figure 7.22 Plot of the fraction reacted as a function of time for several NaCl concentrations (10 g/1 OCf at pH 4).

Plots of 1-(1-R)1/3 against time t are shown in Figure 7.23. The slopes of the straight lines in Figure 7.23 yield apparent rate constants, ka, of which log values are plotted against log NaCl concentration, as shown in Figure 7.24. The value of the reaction order obtained from the slope of the line in Figure 7.24 is 0.6. The value of the reaction order using Eq (7.22) is similar to a 0.5 dependency observed from the previous kinetic study. Figure

7.24 Log /<„, min' ^ X-(l-l

The Plot g/1 OC1'

OCF

order

of at

pH 1-(1-R) at

of pH 4.

reaction

4. 1/3 Slope

as

Log

a Time,

=

with function

0.6 NaCl,

mins respect

M

of

time to

NaCl

for

a for

solution

solutions

of

10 of

152 g/1 10

153

7.7.3 Effect of OCf

The effect of OC1 concentration (1-10 g/1) on the dissolution of gold was examined by measuring the reaction kinetics in solutions of 200 g/1 NaCl at an initial pH 4, as shown in Figure 7.25. As expected, the dissolution rate of gold increased with increasing OC1 concentrations. The plots of l-(l-R) versus time showed linear relationships, as given in Figure 7.26. The apparent rate constants, ka, were obtained from the slopes of the straight lines in Figure 7.26.

o OC1- 1 g/1 + OC1- 2 g/1 AOC1- 5 g/1 □ OC1- 10 g/1

Time, mins

Figure 7.25 Plot of the fraction reacted as a function of time for several OC1 concentrations in NaCl 200 g/1 at pH 4.

These values were used to plot the linear relationship between the log (rate constant) and log (OCF) concentration shown in Figure 7.27. From the data in Figure 7.27, the value of the reaction order obtained from the slope of this line is 1.1 with respect to OCF. In other words, the reaction order was evaluated to be of almost first order dependency to OC1 . 154

Time, mins

Figure 7.26 Plot of 1 -(1 -R)I/3 as a function of time for a solution of 200 g/1 NaCl at pH 4.

Slope = 1.1

Log [OC1-], M

Figure 7.27 The order of reaction with respect to OC1 in 200 g/1 NaCl at pH

4. 155

7.7.4 Effect of Temperature

Experiments were performed to determine the effect of temperature on the dissolution of gold in the temperature range 22 - 35°C. Figure 7.28 shows the fraction of gold dissolved versus time. As can be seen from Figure 7.28, an increase in temperature caused an increase in the dissolution rate of gold. The characteristic function of the shrinking core model for surface reaction control ( l-(l-R) ) showed an excellent straight line when plotted versus time, as presented in Figure 7.29.

Time, mins

Figure 7.28 Plot of the fraction reacted as a function of time at several temperatures in 200 g/1 NaCl and 10 g/1 OC1 at pH 4.

1 The plot of log (apparent rate constants, ka) as a function of — gives straight line, as shown in Figure 7.30. The apparent rate constants were used to determine an activation energy. The activation energy obtained from Figure 7.30 for temperature dependence was calculated as 32.9 kJ/mol. Therefore, 156 at such a high activation energy, the reaction would be expected to be controlled by a kinetic step.

2, 0.2

Time, mins

Figure 7.29 Plot of l-(l-R)13 as a function of time for a solution of 200 g/1 NaCl and 10 g/1 OCT at pH 4.

Ea = 32.9 kJ/mol

Figure 7.30 The activation energy for the dissolution of gold from the slope of Arrhenius plots in 200 g/1 NaCl and 10 g/1 OC1 at pH 4. 157

From the results of the effect of parameters, such as pH, NaCl, OCF and temperature, the generalised rate equation by using the developed shrinking core model (see Eq (7.28)) can be represented as:

1-(1 -R)V3 = kc exp[ 32-9^L"'°^][//*]°3[AfaC/]06[OCl~J11/ (7.29)

7.8 Stability Test in Chloride System

7.8.1 Effect of pH

Figure 7.31 shows the effect of pH on Au loss with time in a solution containing 200 g/1 NaCl, approximately 10 mg/1 AuC14’, 4 g/1 sphalerite and 200 g/1 silica as an inert material. From the data in Appendix 7.3, the plot presented in Figure 7.32 shows the reproducibility (Topping, 1971; Caulcutt, 1992) of the effect of pH on Au loss in solutions determined by the /-test.

The 95% confidence limits for the mean yield of Au loss in solutions is obtained from the following formula:

~X ±t- SD / 4n (7.30) where X : the mean values, SD : the standard deviation, n : the number of test, / : the calculated value from 95 % confidence limits from the /-test. 158

In acidic pH solutions, the stability of Au on sphalerite showed a strong pH- dependence. This result indicates that gold chloride can be rapidly moved to a passive state and reduced to metallic gold by sphalerite. Typically, the almost complete loss of gold from solutions can be achieved in 60 mins. The chemical equation involving the decomposition of zinc sulfide, as sphalerite, can be represented as (Peters, 1976 and 1992; Sudderth et al., 1978; Jan et al., 1976; Cho, 1987a; Wang et al., 1989; Kahanda and Tomkiewicz, 1989):

ZnS + 4Cr -> ZnCl\~ +S°+ 2e" (7.31)

S° +4H20->S02a + 8/T +6

100 + 4 ® "A + * ------f 80 . A D * 5 • _ ■ ■ co 60 " ■ ■ pH 2 oC/3 + pH 3 3 40 □ < A pH 4 □ pH 5 'i n OpH 6

01 1------1------1------1 0 20 40 60 Reaction Time, mins

Figure 7.31 Effect of pH on Au loss with time for a 200 g/1 NaCl,

approximately 10 mg/1 AuC14 solution.

Gold complexes can be precipitated to metallic gold unless sphalerite is completely dissolved and oxidising conditions are maintained throughout the 159 leach. The precipitation of metallic gold from gold complexes can be represented by combining Eq (7.4) and Eqs (7.31) and (7.32), according to the following generalised equation:

And. + — ZnS + xH.O —> An + -Znd2~ + —CT + -SO2,~ +2xH' (7.33) 4 4 3 4

oN wT C/3o

3

Reaction Time, mins

Figure 7.32 The reproducibility of the effect of pH on Au loss from the data of Appendix 7.3.

7.8.2 Effect of Different Sulfide Minerals

The effect of sulfide minerals on Au loss in solutions containing 200 g/1 NaCl, approximately 10 mg/1 AuCl4\ 4 g/1 sulfide minerals and 200 g/1 silica as an inert material at pH 4, was examined for different sulfide minerals, as presented in Figure 7.33. The stability of gold complexes is also strongly affected by sulfide minerals. Gold complexes in solution decreased dramatically in the presence of sulfide minerals to produce metallic gold. 160

The variation of gold losses in solutions with sulfide minerals has the same tendency as those in Figure 7.31. A plot using the data in Appendix 7.4 shows the reproducibility of the effect of sulfide minerals on Au loss in solutions by using the /-test, as shown in Figure 7.34. As expected, in acidic solutions, the loss of gold indicates that gold chloride can be reduced rapidly to metallic gold by sulfide minerals. The reaction with sulfide minerals, such as galena, pyrite, and chalcopyrite, can be postulated with regard to the stability of gold on sulfide minerals in the chloride system. The reaction for the decomposition of lead sulfide, as galena, can be represented as (Peters, 1976 and 1992; Guy et al., 1983; Morin et al., 1985; Byerley and Scharer, 1992; Subrahmanyam and Forssberg, 1993; Ahlberg and Asbjornsson, 1993):

PbS + 3CP +4H20-> PbCl3 + SO]~ + QH+ + 8

Following the same procedure as used for sphalerite, the precipitation of gold complexes to metallic gold can be associated with the following generalised equation:

x x 23x x AuCl: + — PbS + xH20 —» An + — PbClZ +-----CP +-SO^ +2xH+ (7.35) 4 4 2 4 3 12 4 4 '

In a similar way, the stoichiometry of the decomposition of iron sulfide, as pyrite, can be represented by (McKay and Halpern, 1958; Kelly, 1965; Filmer et al., 1984; Zheng et al., 1986; Colak et al., 1987; Murthy, 1990; Bayrakceken et al., 1990; Ramprakash et al., 1991; Ximing et al., 1992; Cheng and Iwasaki, 1992; Robins, 1993; Corrans et al., 1993; Zhu et al., 1993; Chander et al., 1993): FeS2 +8 H2O^Fe3+ + 2SO^ + 16/T + 15

The overall reaction for the precipitation of metallic gold from gold complexes can be represented as:

AuClZ + — FeS2 + xH20 —> An + —Fe3+ + —C/' +-S02~ + 2x.H+ (7.37) 4 8 2 2 8 2 4 4

100 ITT D i " “ ■ 80 . A

60 on o ■ Sphalerite 3 40 < □ Galena A Pyrite 20 O Chalcopyrite

0 20 40 Reaction Time, mins

Figure 7.33 Effect of sulfide minerals on Au loss at pH 4 for a 200 g/1

NaCl, approximately 10 mg/1 AuC14 solution.

Studies by other researchers (Tragert and Robertson, 1955; Groves and Smith, 1973; Biegler and Swift, 1976; Biegler, 1977; Murthy, 1990; Pang and Chander, 1992; Li and Iwasaki, 1992; Das et al., 1992; Barcia et al., 1993; Jana et al., 1993) on the decomposition of copper sulfide, as chalcopyrite, can be represented by the reaction:

2CuFeS2 + \2H20 -» Cu2S + 2Fe2* +3SOf +24H* +22e“ (7.38) 162

The generalised overall reaction for the precipitation of metallic gold from gold complexes can be written as:

And; +^CuFeS2 + xH2() —> Au +-— Cu2S + 4CT +~^2+ +^SO? +2xH+

(7.39)

. ©

Reaction Time, mins

Figure 7.34 The reproducibility of the effect of sulfide minerals on Au loss from the data of Appendix 7.4.

It should be apparent from this discussion that the mechanisms of the dissolution and precipitation of gold in chloride solutions are still not adequately resolved and that much remains to be explained in these areas. Other factors, such as effect of silica and temperature, may also influence the stability of gold complex in the chloride system. These areas describing the electrochemical dissolution of sulfide minerals by two half-cell reactions using either a gold or sulfide minerals electrode together with a reference electrode in Chapter Eight and the reactions involving sulfur species in the chloride/ hypochlorite system, will be discussed in Chapter Nine in further detail. CHAPTER EIGHT

ELECTROCHEMICAL STUDY OF GOLD DISSOLUTION IN CHLORIDE SYSTEM

8.1 Introduction

The dissolution of gold and sulfide minerals in chloride solutions was studied as half-cell reactions using a rotating disc electrode (RDE). The effects of gangue minerals, solution pH, sodium chloride concentration and hypochlorite concentration on the electrochemical behaviour of gold were investigated at ambient temperature (20°C).

The electrochemical study was undertaken to determine:

• The effect of pH and chloride concentration on the current-potential relationship for gold dissolution,

• The mixed potential and current density required for the dissolution of gold in chloride solutions of different compositions,

• The effect of hypochlorite on the current - potential curves for gold reduction,

• The relationship between the current density and the mixed potential for gold dissolution in a chloride system by using the Butler-Volmer equation, 164

• The reaction rates for several sulfide electrodes in a chloride system.

8.2 Experimental

8.2.1 Materials

8.2.1.1 Chemical Reagents

The following chemical reagents were used in the experimental investigations:

Sodium chloride : (M & B) Analytical reagent, assay (on dried basis) 99.9 % minimum NaCl.

Sodium hypochlorite : (AJAX) 5 litres volumetric solution. Available chlorine: 125 g/1 when packed.

8.2.1.2 Gold Electrode

The gold electrodes which were used as the working electrodes (WE) for the electrochemical investigation were made of high purity gold obtained from Johnson Matthey, Australia. A gold disc with a diameter of 0.5 cm was cut from a 0.5 cm diameter gold rod (99.99 % high purity grade) and was then mounted into a teflon holder which could be screwed on to the rotator shaft arrangement (Model MSR of Pine Instrument Co.) so that only the bottom surface of the disc was exposed to the solution. 165

8.2.1.3 Glassy Carbon Electrode

A Pine Instrument Co., Glassy Carbon Rotating Disk Electrode (GCRDE) mounted in epoxy resin (which was similar to that used for the gold electrode) was carefully polished and used as a working electrode for the electrochemical experiments.

8.2.1.4 Sulfide Electrodes

The working electrode used was made from a combination of graphite and sulfide minerals (e.g. sphalerite, galena, pyrite, chalcopyrite) composite.

The preparation of the sulfide electrodes involved the following steps: a) Graphite-sulfide mixture composition b) Pelletizing c) Waxing d) Moulding e) Machining f) Polishing

a) Graphite-sulfide mixture composition

The sulfide electrodes (Figure 8.1) comprised a graphite-sulfide powder mixture (e.g. sphalerite, galena, pyrite and chalcopyrite). The graphite- sulfide mixture consisted of 50 % by volume of sulfide minerals powder (with particle size of less than 45 micron) and 50 % by volume of graphite powder with average particle size of 14.5 micron. 166

Insulated shaft

Stainless Steel Shaft

Interconnecting Piston

Epoxy Casing Graphite + Sulfide Mixture

Figure 8.1 Schematic diagram of the graphite-sulfide mixture electrode. 167

b) Pelletizing

Pellets were formed by compressing a mixture of graphite and sulfide in a die of the desired size. A stainless steel interconnecting piston was used to compress the mixture in the die to form the pellet. A schematic arrangement of the equipment used for the fabrication of the graphite-sulfide pellets is shown in Figure 8.2. A hydraulic press to compress graphite-sulfide mixture was employed to form the pellets suitable for making the working electrode. The working pressure to form the pellets was around 20,670 KPa. The graphite-sulfide pellet was fastened onto a stainless steel rod to produce a tight joint of very low contact resistance.

mm Ifsllll \ Plunger Top Plate

Pellet Die

Interconnecting Piston

Bottom Plate Graphite + Sulfide Mixture

Figure 8.2 A schematic arrangement of equipment used for the fabrication of the graphite-sulfide pellets. 168

c) Waxing

After connecting the stainless steel rod with the pellet, the electrode was immersed in hot wax bath for around 2 minutes, and the temperature was adjusted to approximately 95°C until air bubbles no longer emerged from the graphite-sulfide pellet. The hot wax can penetrate the graphite-sulfide pellet filling in empty pore spaces and subsequently make it impermeable to the chloride solution. The electrode was allowed to cool after removal from the wax bath in order to set in epoxy resin.

d) Moulding

The vessel used to mould the resin was first covered with a small amount of grease to act as a releasing agent. This enabled the epoxy resin to be easily separated from the vessel after setting. The resin covered the steel-graphite sulfide pellet to provide adequate strength for bonding. Two notches were also filed into the steel shaft below the epoxy boundary. The epoxy was set in a vacuum oven at a pressure of -80 kPa relative to atmospheric pressure and at ambient temperature for the first half hour in order to eliminate air bubbles. The epoxy resin was allowed to set for 24 hours. The electrode was embedded in the epoxy resin so that only its lower surface was exposed to the solution.

e) Machining

The moulded electrode was turned on an Emco Compact 5 lathe until an outer diameter of 16 mm was obtained. The graphite sulfide electrode was 169 assembled from a 0.8 cm graphite-sulfide mixture disc, an interconnecting piston, and a 316 stainless steel rod.

f) Polishing

The exposed surface of the electrode was first wet polished on a 1200-grit, silicon-carbide wet and dry paper. After the polishing the electrode was rinsed with distilled water and allowed to dry. The surface was then cleaned with a cotton wool moistened with distilled deionised water. The electrode was again rinsed with distilled deionised water and allowed to dry. The resistance of each graphite sulfide electrode was approximately 1.4 ohm. Prior to each electrochemical experiment the electrodes were subjected to this preparation and only the first cycle was recorded unless stated otherwise.

8.2.2 Equipment

8.2.2.1 Electrochemical Equipment

The equipment used in the electrochemical experiments consisted of a BAS- CV 27 cyclic voltammograph, a MSR rotator and a speed controller, a YEW type 3086 X-Y recorder, an ORION pH meter and a P7080B multimeter (Figure 8.3). All data were acquired via a data acquisition board to transfer the output signals from a BAS-CV 27 cyclic voltammograph to an IBM compatible computer and a YEW type 3086 X-Y recorder. The electrochemical experiments were carried out in a glass reactor placed in constant-temperature water bath. 170

Figure 8.3 The equipment and apparatus set up for electrochemical half-cell studies.

Cyclic Volt ammograph

A BAS-CV 27 cyclic voltammograph with a built-in scanner was used for the scanning voltammetry experiments. The equipment has the following specifications:

Sweep range : 5.00 V Scan rate range : 0.1 mV/s to 10.0 V/s Compliance voltage : 10V 171

Maximum available current : 120 mA typical 100 mA minimum Current to voltage output range : 2 A/V to 10 mA/V

Rotator

A Pine rotator and a model MSR speed controller with a speed range of 50 to 9,999 rpm were used. Its accuracy was better than 1 % of the dial setting for the rotation speed of the electrode.

Recorder

A YEW type 3086 X-Y recorder was used to record the scanning voltammograms. Its features comprise an accuracy of 0.25 % over an effective recording area of 180 x 250 mm with maximum sensitivity of 5 V/cm and a maximum pen speed of 500 mm/s for both X and Y axes. This system detects fast moving input signals with excellent accuracy.

Multimeter

A Model P7080B high impedance multimeter (a maximum sensitivity of 0.1 mV) was used to measure the rest potential of the half-cell and to confirm the voltammograph potential readings. pH meter

An ORION Research, model 520A with a ± 0.01 accuracy for measuring pH was used to measure the pH of the investigated solutions throughout these experiments. pH was measured using a calibrated glass combination 172

(Ag/AgCl) electrode with an ORION pH meter with an accuracy of 0.01 pH units. By choosing buffer solutions of proper pH: potassium biphthalate (pH =4.01), sodium phosphate and potassium phosphate (pH=7.00), and sodium carbonate and sodium bicarbonate (pH =10.01) either hypochlorous or hypochlorite will be the predominant species.

Data acquisition system

Computer data acquisition software was used to record cyclic voltammograms in order to obtain the digital current-potential data (Figure 8.4). All data was acquired via a plug-in acquisition board to an IBM 386 compatible computer operating Quicklog pc™ data acquisition software version 1.04 (Strawberry Tree Incorporated, 1990).

Figure 8.4 The data acquisition system set up for electrochemical half-cell studies. 173

Constant temperature water bath

A temperature controller, manufactured by LABEC, Inc., was used to control the system temperature.

8.2.2.2 Three Electrode System

Electrochemical measurements were obtained using a typical three electrode system. The electrodes consisted of a gold or graphite-sulfide composite working electrode (WE), a standard calomel reference electrode (SCE), placed in a Luggin capillary tube positioned close to the surface of the working electrode and a graphite rod as a counter electrode (CE).

Working Electrode

Gold, glassy carbon, and graphite-sulfide electrodes were used as working electrodes. The gold and glassy carbon electrodes had diameters of 5 mm and geometric surface area of 0.196 cm2 each while the graphite-sulfide electrode was 8 mm diameter with a geometric surface area of 0.503 cm2.

Reference Electrode

A saturated calomel electrode (SCE) was used as the reference electrode for

c. measuring redox potentials. The cell arrangement used may be represented as:

Au\ KCl(al),Hg2Cl2W\Hgm (8.1) Red ) 174

The measured potential readings (Esce) may be related to the potential of hydrogen (Eh) at 25°C by:

Et = Emosura, + 0.245volt at 25°C (8.2)

Counter Electrode

A graphite rod as the counter electrode (CE) was employed for electrochemical measurement during this experiment. This electrode, 5 mm diameter and 140 mm long, was shielded in a Pine Instrument fritted glass compartment and was immersed in electrolytic solutions of the same composition as the test solutions.

8.2.3 Experimental Procedure

A BAS-CV27 cyclic voltammograph with a built-in scanner was used to determine current density-potential curves for the half-cell reactions. The working electrodes used were in the form of rotating disks with only their bottom surface immersed in the solution. A five-port spherical Pyrex cell was used in the experiment. The cell top had entry ports for the working electrode stirrer and a teflon blade. The teflon blade provided enough clearance for the working electrode to rotate smoothly. The side of this cell had entry ports for a counter electrode, a Luggin capillary housing for a reference electrode, a gas bubbler and a thermometer. The counter electrode compartment was separated from the test cell by a fritted glass disk. The Luggin capillary was positioned close to the surface of the working electrode to minimize IR drops in the system. The electrochemical equipment set up is shown in Figure 8.5. 175

Figure 8.5 The electrochemical equipment set up for half-cell studies.

All chemical reagents used in the tests were of analytical grade. Distilled deionised water was used to prepare the required chloride-oxidant solution concentrations for the tests. Approximately 180 ml, with a known NaCl concentration and also with varying quantities of Au(III) and NaOCl, was used as test solution for each electrochemical study. The solution pH was adjusted by adding either NaOH or HC1. Before immersion of the test electrode, the electrolyte was deoxygenated with high purity (oxygen-free) nitrogen for 30 minutes using an appropriate glass nozzle for good gas 176 dispersion with stirring by a magnetic stirrer. The solution was then transferred to the reaction cell which was placed in a water bath equipped with an automatic temperature controller to maintain constant temperature during each experiment. All experiments were conducted at 20°C ± 0.5°C unless stated otherwise.

During each run, oxygen-free high purity nitrogen gas was passed slowly over the solution surface to maintain a low dissolved oxygen content. For each experiment only the first scan was recorded and scanned at 20 mV/s unless stated otherwise. Potentiostatic polarization was performed using a BAS-CV 27 cyclic voltammograph. The electrode’s resistance and rest potential were measured with a P7080B multimeter. A YEW type 3086 X-Y recorder was used for the potential sweep measurements with a scan rate of 20 mV/sec. A quicklog data acquisition system was used to record the current - potential digital data.

8.3 Scanning Voltammetry

8.3.1 Anodic Dissolution of Gold in Chloride Solution

Anodic polarization curves for gold in deoxygenated 0.855 M NaCl solution at various pH values are shown in Figure 8.6. For pH values between 2 and 6, an increase in the pH shifted the anodic polarization curves towards the more positive direction. From Figure 8.6, for potentials of 0.8, 1.0 and 1.2 V vs SCE at which gold dissolution occurs (Latimer, 1952; Gaur and Schmid, 1970; Frankenthal and Siconolfi, 1982), their corresponding current densities are 10, 30 and 60 mA/cm2, respectively. Using the data obtained from Figure 8.6, the plots of potential versus pH is presented in Figure 8.7. Figure Figure

8.7 8.6

deoxygenated Potential, E vs SCE Anodic deoxygenated 3x10 Potential

2 polarization ,

6x10 as

0.855M a

Log

2 0.855M

function A/cm

Current of

NaCl

2 gold

NaCl.

from of

Density, at as pH 45

various

a

the mV/decade

for function with

extrapolated mA/sq.cm current

pH respect

values. of

to density current

[H

+ Tafel ]

of density

line 1x10'

177

in in

2 ,

178

This plot shows a slope of 45 mV/pH. The observed value is slightly smaller than the values of 47 to 115 mV per decade with respect to [Hf] as reported by Hoarse (1966a and b) on bright gold in acid solutions. The reaction order with respect to the hydroxyl ion concentration for gold dissolution was determined from the anodic polarization curves by plotting log current density versus pH, at a constant potential, in the Tafel region, as shown in Figure 8.8.

Figure 8.8 Anodic current density as a function of pH at constant potential (E = 0.8, 1.0, 1.2 V vs SCE, respectively) from the extrapolated Tafel line in deoxygenated 0.855M NaCl.

The reaction order with respect to the hydroxyl ion concentration for gold dissolution was deduced from the slope of the straight line plots in Figure 8.8 as:

( 3log/<1N) = 0.4 (8.3) dpH ) [cr] 179

1.4 -HNTaCl 0.855M ONaCl 1.709M

0.6

-2.0 -1.0 0.0 1.0 2.0 3.0 Log Current Density, mA/sq.cm

Figure 8.9 Anodic polarization of gold as a function of current density in several deoxygenated NaCl concentrations at pH 6.

Anodic polarization curves for gold in deoxygenated chloride solutions at pH 6 are shown in Figure 8.9. For a concentration of CL between 0.855 and 3.419 M, increasing the Cl concentration shifted the polarization curves of gold in the more positive direction. This means that the rate of the anodic reaction increases with an increase in the Cl concentration. For varying CL concentrations, a slope of 118 mV/decade was obtained, as shown in Figure 8.10. The reaction order with respect to the chloride ion concentration for gold dissolution was determined from the slope of the log current density versus log NaCl concentration plot, as shown in Figure 8.11, and is expressed as:

(8.4) 180

118 mV/decade

Log NaCl, M

Figure 8.10 Potential as a function of NaCl for current densities of 1.0, 10, and 60 mA/cm2 from the extrapolated Tafel lines at pH 6.

Slope =1.9

-0.2 0.0 0.2 0.4 0.6 Log NaCl, M

Figure 8.11 Anodic current density as a function of NaCl concentration at constant potential (E = 0.8, 0.9, and 1.0 V vs SCE, respectively) from the extrapolated Tafel line at pH 6. 181

This value is in good agreement with the value of 1.88 reported by Heumann and Panesar (1966) and 1.9 by Frankenthal and Siconolfi (1982) for gold dissolution.

8.3.2 Cathodic Reduction of Hypochlorite

Cathodic polarization curves for gold in deoxygenated 1.709 M NaCl and 10 mM OC1 solutions are shown in Figure 8.12 at different pH values. The rate of hydrogen evolution at a given potential decreases with increasing pH. In Figure 8.13, current densities of 1, 5 and 10 mA/cm2 within the Tafel region as reported by Wu (1987) show a slope of approximately 64 mV/decade This result is close to the value of 86 mV/decade as observed by Wu (1987) for (transfer coefficient) =0.6.

Log Current Density, mA/sq.cm

Figure 8.12 Cathodic polarization as a function of current density at various pH values in deoxygenated 1.709M NaCl and 10 mM OC1 . 182

The reduction of hypochlorite ion to chloride ion (Harrison and Khan, 1971; Hine et al., 1971 and 1974; Wu, 1987; Ahmadiantehrani et al., 1991) can be written as:

ocr +2 H* + 2e~ = cr + h2o (8.5)

The Nernst equation related to Eq. (8.5) is expressed as:

230W (8.6) 2 F [Cl-]

Eq. (8.6) can be rewritten as:

„ 2202RT. [OCr] 2202RT. rr/+1 E = E +———log-—— +------log[//+] (8.7) 2 F [Cl] F

64 mV/decade |^>^with respect to [H ]

Figure 8.13 Potential as a function of pH for a current density of 1, 5 and 10 mA/cm2 from the extrapolated Tafel lines in deoxygenated 1.709M NaCl and 10 mM OCT. 183

I RT\ The value calculated by Eq. (8.7) is 59 mV/decade, i.e. 2.303— , with F ) respect to [H+]. The experimental value of 64 mV/decade obtained from this study is in between the theoretical value of 59 mV/decade and the reported value of 86 mV/decade by Wu (1987). The reaction order with respect to the hydrogen ions can be derived from a plot of log ic versus pH for potentials 0, 0.2 and 0.3 V vs SCE corresponding to current densities of 1, 5, 10 mA/cm2, respectively in the cathodic Tafel region. This plot is shown in Figure 8.14 and shows that

'diogC = -0.2 (8.8) 8pH J [cry

Slope = - 0.2

Figure 8.14 Cathodic current density as a function of pH at constant potential (E = 0, 0.2 and 0.3 V vs SCE) from the extrapolated Tafel line at pH 6 in deoxygenated 1.709M. 184

The cathodic polarization curves for gold at several deoxygenated OCf concentrations and 1.709 M NaCl is shown in Figure 8.15 at pH 8. Increasing the OCf concentrations shifted the potential in the positive direction. A cathodic Tafel slope of approximately 115 mV/decade, ca.

2.303— , was obtained from the Tafel behaviour at cathodic current F densities of 1, 5 and 10 mA/cm2 as shown in Figure 8.16. This value is similar to the value of 103 mV/decade reported by Wu (1987).

ocr (M)

2 0.004 3 0.010 4 0.019 5 0.039

2 3 4 5

Log Current Density, mA/sq.cm

Figure 8.15 Cathodic polarization as a function of current density in several deoxygenated OCf concentrations and 1.709M NaCl at pH 8.

Moreover, the reaction order with respect to the hypochlorite ions was determined from a plot of log /c versus OCF at constant potential within the cathodic Tafel region. The reaction order obtained from Figure 8.17 is expressed as:

feiogO =Q3 (8.9) (Slog[OC/ pH 185

115 mV/decade

Log OC1-, M

Figure 8.16 Potential as a function of OC1 concentrations for current densities (1, 5 and 10 mA/cm2, respectively) from the extrapolated Tafel lines at pH 8 in deoxygenated 1.709M NaCl solution.

Slope = 0.3

Log OC1-, M

Figure 8.17 Cathodic current density as a function of OCF concentrations at constant potential (E = 0, 0.2 and 0.3 V vs SCE, respectively) from the extrapolated Tafel lines at pH 8 in deoxygenated 1.709M NaCl solution. 186

8.4 Interpretation of Current-Potential Curves

Anodic and cathodic polarization data obtained in 3.419 M NaCl and 19 mM OCf solution, as shown in Figure 8.18, were investigated using cyclic voltammetry with the RDE at 500 rpm. The extrapolated anodic and cathodic Tafel lines (dotted) intersect at the observed potential, Ecorrosion, and fix the magnitude of the corrosion current density, Jcorrosion- At different pH values, a Tafel slope of 40 mV/decade, i.e., ^2.303^-J , was obtained in the vicinity of the corrosion potential (Figure 8.19):

= 40/;? V / decade (8.10) dpH J

Log Current Density, mA/sq.cm

Figure 8.18 Anodic and cathodic polarization in deoxygenated 3.419M NaCl and 19mM OCF concentration at various pH values. 187

Tafel slopes ranging from 37 to 61 mV/decade in 2N NaCl were reported by Podesta and Arvia (1965), 52 mV/decade by Kaesche (1959) in perchlorate solution. The result obtained in this study is similar to the values obtained by previous investigators.

40 mV/decade

Figure 8.19 Effect of pH on the corrosion potential from the extrapolated Tafel lines in deoxygenated 3.419M NaCl.

The reaction order with respect to the hydroxyl ion concentration for the dissolution of gold can be determined from the extrapolated polarization curves by a plot of log current densities versus pH in the Tafel region. This plot is shown in Figure 8.20 with the reaction order given by the slope:

d|0g'C,„r^| = -0.3 (8.11) dpH J[cr]

The reaction order in this study is slightly smaller than -0.4 as previously reported by Chin and Nobe (1972) in acidic chloride media. Anodic and cathodic polarization of gold in several deoxygenated NaCl concentrations 188

and 19 mM 00 is shown in Figure 8.21 at pH 8. Increasing the Cl

concentrations shifted the anodic polarization curves of gold towards the more

active direction.

S* -0.5

Slope = - 0.3

Figure 8.20 Effect of pH on the corrosion current density from the

extrapolated Tafel lines in deoxygenated 3.419M NaCl.

-B-NaCl 0.855M -4-NaCl 1.709M ■©■NaCl 3.419M

Log Current Density, mA/sq.cm

Figure 8.21 Anodic and cathodic polarization as a function of current density

in several deoxygenated solutions NaCl and 19mM OCF

concentration at pH 8. 189

A corrosion potential of 80 mV/decade, i.e., 2.303 as shown in F

Figure 8.22, shows that

f dEc = QOmV / decade (8.12) \d logic/"];

The result obtained is in good agreement with the values reported by Nicol (1981) which vary between 60 and 80 mV/decade.

80 mV/decade

Log NaCl, M

Figure 8.22 Effect of NaCl concentration on the corrosion potential from the extrapolated Tafel lines in deoxygenated solutions 19mM 00 at pH 6.

Moreover, the reaction order with respect to chloride ions can be determined from a plot of log zc versus Cf in the Tafel region, as shown in Figure 8.23. 190

dlog jcorr = 0.44 (8.13) Ulog[C/“]J PH

Reaction orders ranging from 0.3 to 1.2 were observed by Frankenthal and Siconolfi (1982) on the dissolution of gold in concentrated chloride solutions.

Slope = 0.44

Log NaCl, M

Figure 8.23 Effect of NaCl concentration on the corrosion current density from the extrapolated Tafel lines in deoxygenated solutions 19mM OCF at pH 6.

8.5 Evans Diagram

An Evans diagram was constructed for the dissolution of gold in a chloride- hypochlorite solution and the cathodic reduction of hypochlorite ions, as shown in Figure 8.24. Figure 8.25 shows the mixed potential and the rate of gold dissolution from the intersection points of anodic and cathodic polarization. 191

Log Current Density(mA/sq.cm)

Figure 8.24 Anodic and cathodic polarisation curves in deoxygenated

3.419M NaCl and 50 pM Au at various pH values.

The most likely electrochemical anodic half-cell reactions involved in the dissolution of gold in a chloride solution are given by the following reactions (Nicol, 1981a):

Au + 2Cl —> AuCl2 + e (8.14)

AuCl2 + 2Cl~ -» AuCIa + 2e~ (8.15)

Au + 4C/_ -> AuCIa + 3e~ (8.16)

The cathodic half-cell reaction (Wu, 1987) is:

CIO' + 2H+ + 2e~ -> Cl' + H20 (8.17)

The overall reaction between Eq. (8.16) and (8.17) is represented as:

2Au + 5Cl~ +3OCr + 6/T -> 2AuCl4 +3H20 (8.18) 192

If the AuC12 ion participates in an equilibrium, the Butler-Volmer equations for reactions (8.15) and (8.17) are:

d) z FE -(1 -*>a)zaFEa 'a =naFka[AuCl2 ][C/ ]exp t a a a -naFk_a[AuCl4]ex p V RT J

(8.19) and

<1hzaFEc -(1-4k)zcfec i_c =ncFkc[Cr]ex p ncFk_c [OCl~ ][//+ ] exp RT 7

(8.20)

Equations (8.19) and (8.20) can be rewritten assuming that the back reaction does not contribute significantly to the kinetics:

ia=2Fka[A“Cl2][Cr]exp (8.21) V RT J and -(1 -

At the mixed potential, E„„ i(Em) = 0 or -i.c(Em) = ia(Em), which yields,

k'lAvCigcn f (4)c-(|K)EEm' (8.23) k_c[OCl~ ][H+ ] RT J

It is generally accepted to assume a = c = 0.5 for symmetry of the activation barrier for the anodic and cathodic charge-transfer reactions (Wadsworth and Sohn, 1979). 193

Eq. (8.23) can be reexpressed as Eq. (8.24) as follows:

FE. kg [AuCll ][CI~ ] exp (8.24) \ RT) k-c[OCr][H+]

At a high chloride ion concentration, the following reaction is proposed by

Frankenthal and Siconolfi (1982) for gold dissolution at potential ranges between 1.1 and 1.6 volts versus SCE before oxygen evolution starts.

Au203 + 8Cl~ + 6H+ -> 2AuCl; + 3H20 (8.25)

Also, the following reaction (8.26) is proposed to account for the dissolution of gold in deoxygenated chloride solution (Gallego et al., 1975; Nicol,

1981a):

3 AuCll = AuCl; + 2 Au + 2 Cl~ (8.26)

Combining Eq. (8.25) and Eq. (8.26), the equilibrium that exists in chloride solutions is represented by:

— Au + — Au203 +2Cl'+H* = AuC12 +-h2o (8.27) 3 6 2

K [AuCl2 ] (8.28) ‘7 [cr]2[H+]

Applying Eq. (8.28) to Eq. (8.24),

kaKa[cr? exi k.c[ocr] (8.29) 194

At the mixed potential, the reaction can be rewritten as:

2.303RT *a,[<7 r E. log (8.30) ~F k.AOCT})

From Eq. (8.22), the mixed current density can be expressed as:

FF i m -2Fk_c[OCl~][H+]exp(-^ (8.31)

120 mV/decade.

Log Mixed CD, mA/sq.cm

Figure 8.25 Plot of mixed potential as a function of mixed current density from the intersection of anodic and cathodic polarization in deoxygenated 3.419M NaCl and 50 pM Au.

The relationship between the mixed current density and the mixed potential

2.303 x 2RT 2.303 x 2RT E_. =- \od2Fkc[OCl-][H*}) + l°9('m) (8.32) F F 195

The slope of Eq. (8.31), —------, is consistent with the observed value F of 120 mV/decade, as shown in Figure 8.25.

8.6 Reaction Rates on Gangue Minerals

8.6.1 Zinc sulfide

Using cyclic voltammetry with a stationary sphalerite electrode, data was collected. A plot of the current density versus potential in 1.709 M NaCl at pH 6, at different concentrations of OCT is shown in Figure 8.26. Based on both leaching experiments and thermodynamic studies, the electrochemical dissolution of sphalerite in chloride and hypochlorite solutions can be expressed as two half-cell reactions.

Sphalerite

B C D E

-2.0 -1.5 -1.0 -0.5 0.0 Log Current Density, mA/sq.cm

Figure 8.26 Anodic and cathodic polarisation curves of the sphalerite electrode in several deoxygenated solutions OC1" ( A = 0.0 M; B = 0.004 M; C = 0.01 M; D = 0.019 M; E = 0.039 M) and 1.709 M NaCl concentrations at pH 6. 196

The anodic reactions (Sanchez and Hiskey, 1991; Peters, 1992) may be represented as:

ZnS +4Cr -> ZnCl]~ + S° + 2e~ (8.33)

S° + 4H20 -> SO]~ + SH+ + 6e~ (8.34)

The overall reaction for the anodic side is:

ZnS + 4Cl~ + 4H2O^ZnCl]- + SO]~ + 8H+ + Qe~ E0 = 0.342 V (vs SHE) k-a

(8.35)

The cathodic reaction (Wu, 1987; Ahmadiantehrani et al., 1991) is:

OCr + H20 + 2e~ <^CJ~ +2OH' E0 = 0.890 V (vs SHE) (8.36) k-C

The overall reaction combining the anodic and cathodic reactions can be written as:

ZnS + 4OCI~ -> ZnCl\- + SO* (8.37)

The Butler-Volmer equation for reactions (8.35) and (8.36) gives:

■(1 -*.)z.FE'

(8.38) 197 and -(1 -ifc)zcFEc <-c =>icFkc[Cl~][OH~]ex p 'I'*.™. - ",Fk [OCr ]exp RT

(8.39) where ia,ic : anodic and cathodic current density, respectively, na,tic : total charge transferred in anodic and cathodic reactions, respectively, (|)t7,(|)c : charge transfer coefficients for reactions (8.35) and (8.36), respectively, za,zc : number of electrons transferred in rate-determining steps in anodic and cathodic reactions, respectively, Fa,Ec : voltage for anodic and cathodic reactions, respectively, ka,k_a : specific rate constants for forward and reverse direction for the anodic reaction, respectively, F: Faraday constant, R : gas constant, T: temperature, °K.

The potential difference is large enough that the back reactions (8.38) and (8.39) are insignificant (Wu, 1987). In this case, na = 8, nc = 2. Commonly, za — zc = 1 for single electron transfer processes are favoured.

Then, the anodic and cathodic reaction can be rewritten as:

4>aFEa 8Fka[Cl ]exp (8.40) RT

i_c = -2Fk_c[OCr]exp\ ~(1 (8.41) RT J 198

At the mixed potential, Em, i(E,„) = 0 or -i.c(Em) = i(l(Em), which yields,

exp raU*.-*.-1)' 4A-Ijg ] (8.42) RT gc[r;a ]

The mixed potential can be written as:

2.303RT , ( 4k\Cr]) ------log ——----- (8.43) \k_c[OCl-]j

Thus, from Eq. (8.41), the mixed current density can be expressed as:

-(1 -*e)FEr /„, =-2F*_c[OC/-]exp (8.44) RT

936 mV/decade

Log Mixed CD, mA/sq.cm

Figure 8.27 Plot of mixed potential as a function of mixed current density from the data of Appendix 8.1 (at different OC1 concentrations). 199

The relationship between the mixed current density and the mixed potential can be rewritten as: 2.303 x RT / 2.303 x RT , x =------log[2Fk_c[OC1 ])+ ^ k x log(/m) (8.45) FV-be) F( 1-

The mixed potential versus mixed exchange current density changes by 936 mV/decade, as shown in Figure 8.27. By determining the slope of the straight line in Figure 8.27, the slope from Eq (8.45) can be written as follows: 2.303 x RT = 936 mV (8.46) H 1-

The transfer coefficient, fl and (J)c, for reaction (8.35) and (8.36) can be derived from Eq. (8.46) and the value of------is calculated to be 16. 1-4>C

Furthermore, the rate of the overall reaction (8.37) can be written in the following form (Levenspiel, 1972):

4tJ'hn. ocr K'-c Ac-L, (8.47a) iiF nF

By combining Eq. (8.44) and Eq. (8.47), we get:

~^dT = ~ W°cr lex p(- ~(1 ~^FEm) (8.47b)

Applying Eq. (8.46) to Eq. (8.47b), the rate of the overall reaction can be expressed as:

FEnA -Ack_c[OCl ]exp| - (8.48) dt 16 RTJ 200

8.6.2 Lead sulfide

The current density versus potential plot in 1.709 M NaCl, at pH 6 at varying concentrations of OC1 with a stationary galena electrode is shown in Figure 8.28. A well-defined anodic peak was observed at 420 mV. According to Ahlberg and Asbjornsson (1993), this anodic peak corresponds to the dissolution of PbS according to the following reaction:

PbS = Pb2+ +S°+ 2e~ (8.49)

In an aqueous chloride solution, the electrochemical dissolution of galena can be described by two half-cell reactions and the formula may be represented by the reactions:

PbS + 3Cl~ -> PbCl~ + S° + 2e~ (8.50a)

S° + AH20 -> SO^ + 8H+ + 6e~ (8.50b)

The overall reaction for the anodic reactions is:

PbS + 3CP + AH20

OCl-+H20 + 2e~ 4cr +2OH- E° = 0.890 V (vs SHE) (8.52) k-c

The overall reaction for the anodic and cathodic sides can be represented as:

Pbs + Aocr -> Pbci~ + sol + cr (8.53) 201

Galena

B C D E

-2.0 -1.5 -1.0 -0.5 0.0 0.5 1.0 1.5 Log Current Density, mA/sq.cm

Figure 8.28 Anodic and cathodic polarisation of the galena electrode in several deoxygenated solutions OCF (A = 0.0 M; B = 0.004 M; C = 0.01 M; D = 0.019 M; E =0.039 M) and 1.709 M NaCl concentrations at pH 6.

The Butler-Volmer equation for reactions (8.51) and (8.52) yields:

-(i-

(8.54) and -(1-'h.FE. i_=ncFk\Cl-][OH]exp - ncFk_ [OCr ] exp RT J

(8.55) 202

In this case, na = 8, nc =2. Single electron transfer processes are zn = zc = 1. If back reaction is considered negligible, the mixed potential can be reduced to:

2.303RT f E (8.56) m ^{+.-+.-1) °\k_c\ocry

The mixed current density can be described by:

-2Fk_c[OCl]exp\ (8.57) i m RT

521 mV/decade

Log Mixed CD, mA/sq.cm

Figure 8.29 Plot of mixed potential as a function of mixed current density from the data of Appendix 8.2. 203

The relationship between the mixed potential and the mixed current density can be written as: 2.303 x RT E_. = \opFkJOCr])+2™3_x«l log(,.) (8.58) T(1-ct>.)

The mixed potential versus mixed current density, as can be seen in Figure 8.29, changes by 521 mV/decade which in turn shows that 2.303 x RT = 521 mV (8.59)

Following the same procedure with the sphalerite electrode, the rate of the overall reaction (8.53) can be written as follows:

AdUphs (1 ~<\>c)FEn —dr = ~Ack_c[OCl~]exp' (8.60a) dt ii F

Applying Eq. (8.59) to Eq. (8.60a), the rate of the overall reaction with 1 ------= 9 can also be written as:

***>

8.6.3 Iron sulfide

Variation of the current density with respect to potential by using a stationary pyrite electrode in 1.709 M NaCl concentration containing different concentrations of OCf at pH 6 is shown in Figure 8.30. The main reactions during the electrochemical dissolution of pyrite in an aqueous chloride solution, in the light of the literature (Bailey and Peters, 1976; Lawson et al., 1992; Li and Iwasaki, 1992), can be expressed as two half-cell reactions. 204

B C D E

-2.0 -1.5 -1.0 -0.5 0.0 0.5 1.0 1.5 Log Current Density, mA/sq.cm

Figure 8.30 Anodic and cathodic polarisation of the pyrite electrode in several deoxygenated solutions OC1’ (A = 0.0 M; B = 0.004 M; C = 0.01 M; D = 0.019 M; E = 0.039 M) and 1.709 M NaCl at pH 6.

The anodic reactions are:

FeS2 -> Fe2* + 2S° + 2

Fe2* -> Fe3t +e“ (8.62) 2S° + 8H./J -> 2SO\- +16H* +12t> (8.63)

The overall reaction for the anodic reactions is:

FeS7+QH70<^Fe3+ +2SOl~ +16H+ + 15e~ E0 = 0.394 V (vs SHE) jt

(8.64) 205

The cathodic reaction is:

OCl~ +H20 + 2e~ +>cr +2OH' E0 = 0.890 V (vs SHE) (8.65) k-c

The overall reaction for both the anodic and cathodic sides is represented as:

2 FeS2 +15 OCr +H20-+ 2 Fe2+ + 4SO*~ +15 Cl'+ 2H+ (8.66)

The Butler-Volmer equation for reactions (8.64) and (8.65) gives:

(8.67) and

Lc=ncFkc[Cr][OH~]ex p| czcFE, -iiFk\OCr]ex p| -(1-4kKfec RT )

(8.68)

Here, na = 15, nc =2. As stated earlier, the mixed potential for Eq. (8.66) can be expressed as:

E = ZZ0ZRT logf 15*° ) (8.69) " JW.-+.-D +2k_c[ocr\)

The mixed current density can be expressed as:

-(1 -*.)FEm L = -2Fk_c[OCF]exp\ (8.70) 206

450 mV/decade

Log Mixed CD, mA/sq.cm

Figure 8.31 Plot of mixed potential as a function of mixed current density from the data of Appendix 8.3.

The relationship between the mixed potential and the mixed current density can be rewritten as:

2.303 x RT E_ = log(2«.JOC/-]) + ^^log(,„,) (8.71)

The slope of Eq (8.71) from Figure 8.31 can be written as:

2.303 x RT = 450 mV (8.72)

The rate of the overall reaction Eq. (8.66) can be written as follows:

15dnt (1-4>.)FE, 4,-'„ = -Ack_c[OCr]ex p (8.73a) 2 cU n.F V RT 7 207

Applying Eq. (8.72) to Eq. (8.73a), the rate of the overall reaction with 1 ------= 8 can be derived as: 1-+C

tin FeS2 FE. Ack_c[OCr]exp - (8.73b) dt 15 QRTJ

8.6.4 Copper sulfide

The potential plot with respect to current density by using a stationary chalcopyrite electrode in 1.709 M NaCl containing different concentrations of OCf at pH 6 is presented in Figure 8.32.

Chalcopyrite

B C D E

-2.0 -1.5 -1.0 -0.5 Log Current Density, mA/sq.cm

Figure 8.32 Anodic and cathodic polarisation of the chalcopyrite electrode in several deoxygenated solutions OC1 (A = 0.0 M; B = 0.004 M; C = 0.01 M; D = 0.019 M; E = 0.039 M) and 1.709 M NaCl at pH 6. 208

The electrochemical dissolution of chalcopyrite in an aqueous chloride solution can be expressed as two half-cell reactions. The anodic reactions (Li and Iwasaki, 1992; Pang and Chander, 1992) are:

CuFeS2 + Cu° + 2H+ -> Cu2S + Fe2+ + H2S (8.74)

CuFeS2 -> CuS + Fe2+ + S° + 2e~ (8.75)

CuS + 2H20 -> Cu(OH)2 + S° + 2H+ + 2e~ (8.76)

Cu(OH)2 + H2S + 2H20 -> Cu° + SO2- + 8H+ + 6e (8.77) 2S° + 8H20 -> 2^042' + 16/T -h 12c-- (8.78)

The overall anodic reaction is:

2CuFeS2 +1 2H20

The cathodic reaction is:

ocr +H20 + 2e~ ++cr +2OH- E0 = 0.890 V (vs SHE) (8.80) k-e

The overall reaction is represented as:

2CuFeS2 +11OCT + #20 -> Cu2S+ 2Fe2* + 3.YC+ +11C/“ + 2//' (8.81)

The Butler-Volmer equation for reactions (8.79) and (8.80) gives:

-(1 -*.)z.FE.' 4kzqFE, -»aFk_a[Fe2'][SOl][H*]ex p 'a = >hFka eXP RT RT (8.82) 209 and

-(1-4>c)zcFEt i-c=",Fkc[Cr)[OH-]ex p| <1KZcFE, - ncFk [OCr ] exp RT RT

(8.83)

In this case, n„ = 22, nc =2. za = zc =1 for single electron transfer processes are favoured. The mixed potential can be derived as:

2.303RT , f 11A„ ) ------log------— (8.84)

The mixed current density can be expressed as:

-0-4>e)FEm im=-2Fk_c[OCr]exp (8.85) RT

879 mV/decade

Log Mixed CD, mA/sq.cm

Figure 8.33 Plot of mixed potential as a function of mixed current density from the data of Appendix 8.4. 210

The relationship between the mixed potential and the mixed current density can be rewritten as:

2.303 x RT 2.303 x RT Eni \og(2Fk_e[OCn) + (8.86) F -

The slope of Eq. (8.86) from Figure 8.33 can be expressed as follows:

2.303 x RT = 879 mV (8.87) F(1-4>,)

The rate of the overall reaction Eq. (8.81) can be written as follows:

1"'“. -XJIOCnexpl" (8.88a) 2 dt nF

Thus, the rate of the overall reaction with = 15 can be expressed as: i-+,

= — A k c[OCl~]exp (8.88b) dt M ‘ ~c K^ 15 RT.

When the backward reaction is considered negligible, the electrochemical dissolution of sulfide minerals in chloride and hypochlorite solutions is proportional to the hypochlorite ion concentration. Although the backward reaction is not specifically mentioned, the process presumably involves the backward reaction by way of a mixed potential mechanism. It should be apparent that the mechanisms of the dissolution of sulfide minerals in chloride solution are still not adequately resolved and much remains to be explained in this area CHAPTER NINE

THERMODYNAMIC MODELLING

9.1 Introduction

Comprehensive thermodynamic modelling is one of the viable methods for predicting and interpreting the solution chemistry. In order to understand the solution chemistry of the gold chloride system, thermodynamic modelling offers a powerful method for examining the speciation of aqueous chloride solutions in the presence of gold and different sulfide minerals at equilibrium. As the solution chemistry of the chloride-oxidant system is complex, the speciation at equilibrium is generated with aid of a CHEMIX computer program (Turnbull and Wadsley, 1984).

The present study is undertaken to investigate the effect of oxidants and sulfide minerals on the solution chemistry of the aqueous gold chloride system. On the basis of the predominant species concentration vs pH diagrams the equilibria of the chloride-oxidant, gold-oxidant, sulfide minerals-oxidant systems were constructed. Such diagrams were used to describe the thermodynamic aspects of the dissolution of gold in different aqueous chloride-oxidant systems. Many of the concentrations of sulfur and chlorine species calculated from thermodynamic data appear to be so low as to be irrelevant, but the concentrations in the 1 x 10 10 to 1 x 10 40 molar range were considered to explain the effect of oxidants and sulfide minerals on the solution chemistry of the aqueous gold chloride system. 212

9.2 Thermodynamic Data and Modelling Procedures

The predominance areas of the different species and the solid/solution equilibria in chloride media have been investigated with the aid of thermodynamic data (Wagman et al., 1982; CSIRO Division of Mineral Products; Pourbaix, 1966) and the CHEMIX computer program (developed by CSIRO Division of Mineral Products, Port Melboune, Australia). The CHEMIX computer program was used for the construction of the solution chemistry of gold-chloride media. This program consists of a data bank compiled of basic thermodynamic data for chemical substances and a variety of compatible programs to perform thermodynamic calculations with these data.

On the basis of a data bank compiled by the U.K. National Physical Laboratory in extended Fortran IV for the CSIRO CYBER 76 computer operating on CSIRONET, the program system was developed in 1975. With continuous extension of the data bank and application programs, the fifth Version which culminated in the present program was written in Fortran 77 for the VAX and IBM PC computers.

CHEMIX is based on a modular method of construction which consists of a shared subroutine library and a generalised file structure. This structure has allowed for an extension of this application to new properties in a simple manner. The system consists of eight major application programs with functional access to the directive language and to the data bank. 213

The specifications of these programs are:

FILER : Fitting and storage of basic data REACT : Properties of species and reactions SYSTEM : Predominance area diagrams CHEMIX : Equilibrium calculations on systems EXERGY : Analysis of efficiency of processes VAPOUR : Vapour pressure and critical data ESTIMA : Estimation of thermodynamic data REFER : Storage and retrieval of references

The purpose of this study is to investigate foundation behind the electrochemical behaviour of sulfide minerals in chloride solutions. All possible species concerned with the required system from a THERMODATA package and the literature were considered and are summarised in Appendices 9.1-3. On the basis of the free energy minimisation method (Eriksson, 1979), all possible species, phases, and their quantities using the minimum total Gibbs energy for the system at equilibrium and desired conditions were calculated. In this method, the variables ni for constant temperature and pressure can be minimised in quantity as illustrated in equation (9.1):

(9.1)

where G : total Gibbs free energy of the system, R : gas constant, 214

T : thermodynamic temperature, K, /?, : amount of the /th substance, g° : standard chemical potential and ai : the activity of component /.

The enthalpy balances between the initial and equilibrium states and the equilibrium concentrations of the predominant species can be calculated with the aid of thermodynamic data for the aqueous phase which explicitly takes into account the formation of aqueous complexes (Garrels and Christ, 1965; Kimura et al., 1984; Ali and Christie, 1984; Forssberg et al., 1984; Albano et al., 1990; Eriksson and Hack, 1990; Escudero et al., 1993; Gaspar et al., 1994). The standard state for minerals by using the CHEMIX program adopted in this study refers to the pure minerals at 25°C and at atmospheric pressure.

The input statement of CHEMIX for this study can be summarized as follows:

• SYSTEM : All the possible phases and component species at equilibrium states in terms of GAS (G), AQUEOUS (AQ) and SINGLE (refer to solid phases). • UNITS : Gibbs energy units for input and output. • GIBBS : Gibbs energy for component defined by the system at 25°C and atmospheric pressure. • FILE: Data files for thermodynamic data sets. CPD as the first three characters, indicate the file type in use as thermodynamic data for species. The second three characters refer to the source of the data bank in the file, as follows: 215

- NBS : National Bureau of Standards, Washington - BPL : U.K. National Physical Laboratory - MRL : C.S.I.R.O Mineral Research Laboratories.

• INSPECIES : Phases and species for the known input. • TEMPERATURE Absolute temperatures for input and output. • PRESSURE : Pressures for input and output. • INMOLE : Moles of input species defined by INSPECIES. • TRY : Trial equilibrium moles of all possible species in the system. • EQUILIBRATE : Calculation of required series at equilibrium states • STEP: Calculation of required series of equilibria. • OUTMOLE: Calculation of moles of required species in the system for a number of steps of another property. • VARY: Variation of property to satisfy the predicted constraint. • ACTIVITY : Calculation of log (base 10) activity of a solution component species in the system.

In this study, all the programs using the CSIRO Thermochemistry Program software were run on an IBM 486 personal computer. The typical thermodynamic data taken from the literature (Pourbaix, 1949; Garrels and Christ, 1965; Wagman et al., 1982; Escudero et al., 1993; Gaspar et al., 1994) and the input statements of CHEMIX computer program (CSIRO, 1988) in the present study are shown in Appendices 9.13-24. 216

9.3 Equilibrium Solution Chemistry

9.3.1 The Au - Cl'- NaOCl - H20 System

As shown in Chapter Two (section 2.3.1.2), the equilibrium involved in the sodium chloride-hypochlorite solution can be predicted by establishing two reactions involving final species such as C1,C12, HOC1, OC1, H+. These reactions yield the following equations: [Hociprjcr] [c/2] = k, (9.2) k7[hoci] [ocr] = (9.3) Iff 1 Equations(9.2)-(9.3) relate the concentrations of these species and the hydrogen ion to their corresponding equilibrium constants. The distribution diagram for chlorine species (Latimer, 1952; Pourbaix, 1966; Marsden and House, 1992) showed aqueous chlorine, Cl2, as the predominant species at a pH of below 3 (Figure 2.6). At a higher pH value (pH range from 3 to 7.5), hypochlorous acid, HOC1, is generated. At pH values above 7.5, hypochlorite ion, OC1, predominates. The equilibrium concentrations of predominant species with respect to pH for the Au - Cl - NaOCl - HzO system at 25°C are presented in Figures 9.1 (a)-(b). The concentrations of initial input were [Au] = 1 mM, [NaCl] = 1.709 M and [OC1 ] = 0.194 M. All possible species were considered in these calculations. Although A number of experiments were conducted the Au - Cl - NaOCl - H20 system at 25°C under these conditions, the results in the difference between the sum of dissolved gold complexes and metal gold are somewhat less certain. A possible reason for this can be attributed to a lack of relevant thermodynamic data for the gold-complexes, such as Au(HS)2, Au2(HS)2S2, Au(HS)(OH) 217 and Au(HS)S2 , containing sulfide (or polysulfide) ions (Renders and Seward, 1989) other than chloride due to the presence of H2S04 in this system.

AuCi; ■OOO QQ- QQ QOO Q

auci ;

Figure 9.1 (a) The equilibrium concentration profiles of the predominant Au species as a function of pH; [Au] = ImM, [NaCl] = 1.709M, and [OCl ] = 0.194M.

HCIO

GOQOOOQ Qg OO QQQOOQ

Figure 9.1 (b) The equilibrium concentration profiles of the predominant chlorine species as a function of pH at the same conditions of Figure 9.1 (a). 218

The major species in this system were AuC14", AuC12 , Cl2, HCIO and CIO with solid Au. The CHEMIX program output shows AuC14 as a predominant species, generated in the pH range of l to 5.5. The AuC14" concentration is dependent on the solution pH. In Figure 9.1 (a), when the AuC14" concentration drops at about pH 5.5, Au precipitation starts. This precipitation point corresponds to the intersection point (indicated by the arrow) of dissolved chlorine and hypochlorite ion, as can be seen in Figure

9.1 (b). This means that the generation of the AuC14 species below pH 5.5 is influenced by the concentration of Cl2. Above pH 5.5, the gold chloride complex, AuC14 , becomes unstable and is reduced to metallic gold. In electrochemical terms, the electrode potentials of the two half-cell reactions, involved in the gold dissolution are equal under these conditions, as follows:

Anodic Au + ACl = AuCl4 +3e (9.4) and

Cathodic 2 OCl~ +4 H+ +2

The Nernst equations related to these half-cells are expressed as:

(9.6) and RT [OCI-]2[H*]4 7T=rln------(9.7) 219

E°, the standard reduction potentials, calculated with standard free energy data (in Appendix 9.1), are 1.001 and -0.417 V versus SHE for E°9A and E°9 5, respectively. Under equilibrium conditions, values from Figure 9.1 (a) at pH 5.5 should be substituted into Equations (9.6) and (9.7) to calculate E°9 4 and E°9 5

At high initial NaCl concentration (1.709 M) (Figure 9.1 (b)), the distribution for chlorine species showed that the aqueous chlorine is predominant at a pH of 4 or below, but the predominance of hypochlorous acid, HOC1, and the hypochlorite ion, OC1", were found to be similar to the observation of previous investigators (Groves and Smith, 1973; Marsden and House, 1992). It is clear that the concentration of dissolved Cl2 is directly related to the formation of AuC14 species.

GDQ 'O Q(MD

Figure 9.2 (a) The equilibrium concentration profiles of the predominant Au species as a function of pH; [Au] = ImM, [NaCl] = 0.855M, and [OC1] = 0.019M. 220

Figure 9.2 (b) The equilibrium concentration profiles of the predominant Au species as a function of pH; [Au] = ImM, [NaCl] = 1.709M, and [OCf] = 0.019M.

GDQCQSDCMSSOO

Figure 9.2 (c) The equilibrium concentration profiles of the predominant Au species as a function of pH; [Au] = ImM, [NaCl] = 3.419M, and [OCf] = 0.019M. 221

The equilibrium concentration profiles of the predominant Au species with respect to pH are presented in Figures 9.2 (a)-(c) for NaCl concentrations of 0.855 M, 1.709 M, and 3.419 M NaCl, respectively, in [Au] = ImM and [OCl] = 0.019 M solutions. With increasing NaCl concentration, the predominant domain of AuC14 is shifted toward higher pH. The CHEMIX program output data for this system are shown in Appendix 9.4.

9.3.2 The Au - ZnS - CF - NaOCl - H20 System

The equilibrium concentrations of predominant species as a function of pH for the Au - ZnS - Cl - NaOCl - H20 system at 25°C are shown in Figures

9.3 (a)-(b). The initial inputs are [AuC14] = 50 pM, [ZnS] = 41.05 mM, [NaCl] = 1.709 M, and [OCl] = 0.194 M. At high OCf (0.194 M) concentrations, sulfur species are not significantly present in the solution, as can be seen in Appendix 9.16. It indicates that sulfur species do not dissolve and directly participate in this system. This system shows that the concentration profiles of the oxidizing species are similar to those of the Au - Cl - NaOCl - H20 system of the same OCl concentrations. However, at lower OCl concentrations (0.019 M), the situation significantly differs from that at higher OCl concentration, as depicted in Figure 9.4 (a)-(b). In Figure 9.4 (a)-(b), chlorine species do not appear in the solution. The species considered were AuS , AuC12 , AuC14 , S042, H2S, HS , and S2 together with solid Au. There has been several modelling studies considering Au-S-Cl-H20 system (Krauskopf, 1951; Garrels and Christ, 1965; Vlassopoulos and Wood, 1990) although no experimental work has yet been done to confirm the existence or non-existence of the AuS*y~ species. We therefore use the thermodynamic data reported by Gaspar et al. (1994) for this modelling. In this system, AuC14 species dramatically disappeared and AuS appeared. AuS’ species with a high sulfur content are formed over a wide range of pH. 222

AuCI' QO GO O..@6D

auci;

Figure 9.3 (a) The equilibrium concentration profiles of the predominant

Au species as a function of pH; [ZnS] = 41.05mM, [AuC14 ]

= 50pM, [NaCl] = 1.709M, and [OCT] = 0.194M.

HCIO

QOOOQOQD€> o ooo ooee

Figure 9.3 (b) The equilibrium concentration profiles of the predominant chlorine species as a function of pH at the same conditions as Figure 9.3 (a). 223

From the results of Figure 9.4 (a)-(b), it is believed that AuC14 species may be altered to Au and AuS by the dissolved sulfur species. S042 as a sulfur species is predominant over a wide range of pFI. At pH values of lower than

7, the dominant sulfur species are S042 and H2S(aq). The major sulfur species in this system are S042 and HS , at pH values higher than 7. The turning point for sulfur species in the gold-sphalerite system is pH 10.

CD QOO OCQGCQ GDOCOGD G>

AuS'

Figure 9.4 (a) The equilibrium concentration profiles of the predominant

Au species as a function of pH; [ZnS] = 41.05mM, [AuC14 ] = 50pM, [NaCl] = 1.709M, and [OCF] = 0.019M.

Based on this data, the reactions with S042, H2S and HS as sulfur species dissolved may be represented as follows:

AuC14 + ZnS + 2e -AuS + ZnCl] (9. 8)

AuS +2//+4"C = Au + H2S (9. 9)

h2s = hs+h+ (9.10) 224

HS'+OH- =S°+H20 + 2e~ (9.11) .S'0 +4 H20 = SO? + 8/T +6e~ (9.12)

Thus, the overall reaction equation should be:

3 AuCl; + 3 ZnS + OH~ + 3 H20 =

2 An + AuS~ + 3ZnCl? + SO? + H2S + 5H+ (9.13)

The output data for this system are shown in Appendices 9.5 and 9.6, respectively.

c£>-0oo ooeeoo gd qqqgd q

Figure 9.4 (b) The equilibrium concentration profiles of the predominant sulphur species as a function of pH at the same conditions as Figure 9.4 (a).

9.3.3 The Au - PbS - Cf - NaOCl - H20 System

Figure 9.5 (a) and (b) show the equilibrium concentrations of predominant species, as a function of pH, for the Au - PbS - Cl - NaOCl - H20 system at 225

25°C. The concentrations of the initial input are [AuC14] = 50 pM, [PbS] = 16.72 mM, [NaCl] = 1.709 M and [OC1 ] = 0.194 M. The variation of

AuC14 in solution with respect to pH for this system seemed to have a similar tendency as the Au - ZnS - Cl - NaOCl - H20 system. The intersection point of the dissolved chlorine and hypochlorite ions occurs at about pH = 6, as shown in Figure 9.5 (b). In Figure 9.6 (a)-(b), the species considered were

AuS , AuC12 , AuC14 , S042, H2S, HS , and S2 as the aqueous components and Au as the solid. The turning point of sulfur species is at pH 7.5 and is shifted toward a lower pH than in the gold-sphalerite system. In this system, the predominant equilibrium components of the sulfur species are S042, H2S, and HS . Except where S042 over a wide range of pH becomes predominant. The concentration profile of AuS’ species is influenced by the concentration of H2S which dominates in the solution below pH 7. Above pH 7, HS' becomes the predominant species.

AuCi; ee—O OOQQQCOQD

Figure 9.5 (a) The equilibrium concentration profiles of the predominant

Au species as a function of pH; [PbS] = 16.72mM, [AuC14 ] = 50pM, [NaCl] = 1.709M, and [OCf] = 0.194M. 226

HCIO

OQ QQQQOCT QQQQQQOQD

Figure 9.5 (b) The equilibrium concentration profiles of the predominant

chlorine species as a function of pH at the same conditions

as Figure 9.5 (a).

Following the same procedure for the gold-sphalerite system, the reactions involving gold chloride complex and lead sulfide in the presence of S042,

H2S and HS as sulfur species dissolved may be represented as:

AuCl- + PbS + 2

AuS-+2H++e~ = Au + H2S (9.15)

H2S = HS+H+ (9.16)

HS-+OH- =S°+H20 + 2e~ (9.17)

S° + AH20 = SOX + 8//+ + 6e~ (9.18)

Therefore, the overall reaction can be expressed as:

3 AuCli + 3 PbS + OH~ + 3 H20 =

2 An + AuS- + 3 PbCl; + SOX + H2S + 3 Cl~ + 5H+ (9.19) 227

OQOQ CQQOO GO 0O OOGD

pH

Figure 9.6 (a) The equilibrium concentration profiles of the predominant

Au species as a function of pH; [PbS] = 16.72mM, [AuC14 ] = 50pM, [NaCl] - 1.709M, and [OCT] = 0.019M.

eeoo oocqo—e-e GO OO -OO

A A AA-AA

■ ■■ ■■■■'P

pH

Figure 9.6 (b) The equilibrium concentration profiles of the predominant sulphur species as a function of pH at the same conditions as Figure 9.6 (a). 228

The output data for this system are shown in Appendices 9.7 and 9.8, respectively.

9.3.4 The Au - FeS2 - Cl' - NaOCl - H20 System

The equilibrium concentration profdes of predominant species versus pH for the Au - FeS2 - Cl - NaOCl - H20 system at 25°C are presented in Figure 9.7

(a) to (c). The concentrations of initial input are [AuC14 ] = 50 pM, [FeS2] - 33.40 mM, [NaCl] - 1.709 M, and [OCF] = 0.194 M. To construct these Figures, the species considered were AuS , AuC12 , AuC14 , S042, H2S, HS , S2 , CIO , HCIO, and Cl2 as the aqueous phase and Au as solid. The turning point of the sulfur species is pH 4 and is shifted toward a lower pH than gold-sphalerite and gold-galena system.

QDOGDOeO OOOOQQ Q OOOOOQQO QQO

pH

Figure 9.7 (a) The equilibrium concentration profiles of the predominant

Au species as a function of pH;[FeS2] = 33.40mM, [AuC14 ] = 50pM, [NaCl] = 1.709M, and [OC1] = 0.194M. Figure

9.7

(c) Log Concentration, Molar @ Log Concentration, Molar The Figure The sulphur as chlorine

Figure

equilibrium equilibrium

9.7

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species species

9.7

(a).

(a).

as

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function PH

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pH profiles profiles

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at

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the the same

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conditions 229

as

230

The variation of the AuC14~ concentration with pH in this system significantly differs from the gold-sphalerite and gold-galena systems. Despite a higher OC1 (0.194 M) concentration, sulfur species in solution dissolve, and chlorine species (but the concentration is too small) remain in solution, as demonstrated by Figure 9.7 (b)-(c). These results indicate that pyrite can be readily dissolved, unlike other sulfide minerals such as sphalerite, galena, and chalcopyrite, at a high potential and several intermediate species may participate in the mechanism. In Figures 9.8 (a)-(b), the behaviour of the concentration profiles of the oxidizing species at a lower OCf concentration (0.019 M) is similar to the system at a higher OCf concentration (0.194 M). Chlorine species are however not observed in this system. The output data for this system are shown in Appendices 9.9 and 9.10, respectively.

In the gold chloride-pyrite system, the reactions with S042 , H2S, HS and S2 as sulfur species dissolved may be represented as:

AuCl; + FeS2 + 3e~ = AuS~ + Fe>+ + 4C7' + S2~ (9.20)

AuS~+2H++e- =Au + H2S (9.21) H2S = HS~+H+ (9.22) HS~+ OH~ =S°+H20 + 2e~ (9.23) S° +4H20 = S024- +-8/T +6e- (9.24) 2S° + 302 + 2H20 = 2^0^' + 4/T (9.25)

The overall reaction can be written as:

8 AuCIa + 8FeS2 + 2 OH~ +12 H20 =

AAu + AAuS- +8 Fe3+ +2 SO]~ +2H2S + 8S2~ +32CF +30, +22H+ (9.26) Figure

9.8

(b)

Log Concentration, Molar e Lo9 Concentration, Molar The Au The Figure sulphur =

50pM, species

equilibrium equilibrium

9.8

species @

[NaCl] (a).

as oo-ee

a

as function

concentration concentration

= a e

ee function 1.709M,

eooo of

pH;[FeS PH

PH

of and

o

pH profiles profiles

o

[OC1]

at 2 ]

the o =

-

of of o 33.40mM,

= same

oooo 0.019M. the the

conditions predominant predominant

[A u C 231

1 4 as

]

232

9.3.5 The Au - CuFeS2 - Cl' - NaOCl - H20 System

The equilibrium concentrations of predominant species versus pH for the Au -

CuFeS2 - Cl - NaOCl - H20 system at 25°C are shown in Figure 9.9 (a) to

(b). The concentrations of the initial inputs are [AuC14 ] = 50 pM, [CuFeSJ

= 21.80 mM, [NaCl] = 1.709 M, and [OCF] = 0.194 M. The variation of

AuC14 in the solution with pH in this system is similar to that of the gold- sphalerite and gold-galena systems. The intersection point between dissolved chlorine and hypochlorite is at almost pH = 6, as shown in Figure 9.9 (b).

At a lower OC1" concentration (0.019 M), the species considered were AuS ,

AuC12’, AuC14 , S042, H2S, HS , and S2’ as aqueous and Au as solid, as seen from Figures 9.10 (a) to (b). The turning point of sulfur species is at pH 4 and is shifted toward lower pH levels than the gold-sphalerite and gold-galena systems. The output data for this system are shown in Appendices 9.11 and 9.12, respectively.

AuCi; eeeee 00 00

pH

Figure 9.9 (a) The equilibrium concentration profiles of the predominant

Au species as a function of pH; [CuFeS2] = 21.80mM,

[AuC14 ] = 50pM, [NaCl] = 1.709M, and [OCf] = 0.194M. 233

0

05 O

HCIO

v_A—/ v_y v^y C_/ v_y OQQO COCO

Figure 9.9 (b) The equilibrium concentration profiles of the predominant chlorine species as a function of pH at the same conditions as Figure 9.9 (a).

In Figures 9.10 (a) and (b), the predominant equilibrium species of sulfur are S042 *, 4H2S, and HS . The concentration profile of these sulfur species is similar to that other sulfide minerals. The reactions with S042, H2S, HS and S2 as sulfur species dissolved may be represented as:

2 AuCl; + 2CuFeS2 +6e~ =2 AuS~ +2 Fe2+ +Cu2S + SCl~ +S2~ (9.27)

AuS~+2H++e~ =Au + H2S (9.28) H2S = HS~+H+ (9.29) HS~+OH~ =S°+H20 + 2e~ (9.30) S° + 4H20 = SO2" + 8//+ + 6e~ (9.31) 2S° +302 + 2H20 = 2S0l~ + 4 H+ (9.32) Therefore, the overall reaction can be presented as follows: 8 AuCl; + 8 CuFeS2 + 2 OH~ +12 H20 = 4 Au + AAuS' + 4 Cu2S + 8 Fe2' + 2 SO]~ +2 H2S + AS2~ + 32 CF + 30, + 22 H+ (9.33) 234

AuS'

auci;

Figure 9.10 (a) The equilibrium concentration profiles of the predominant Au species as a function of pH; [CuFeS2] = 21.80mM,

[AuC14 ] = 50pM, [NaCl] - 1.709M, and [OCF] = 0.019M.

ooo GGoo o o oeo oeo

pH

Figure 9.10 (b) The equilibrium concentration profiles of the predominant sulphur species as a function of pH at the same conditions as Figure 9.10 (a). 235

9.4 Related Kinetic Considerations of the Dissolution of Gold

Kinetic studies on gold dissolution in the chloride-hypochlorite system has shown that the rate of gold dissolution is dependent on the solution pH and the concentration of chloride and hypochlorite. In the equilibrium state of this system, the predominant species considered were AuC14", AuC12 , Cl2, HCIO, and CIO as the aqueous components and Au as the solid. The equilibrium concentration profile of the AuC14 species has been observed to be strongly influenced by the relationship between dissolved chlorine and the hypochlorite ion in all systems which were tested. Also, the point where the maximum rate of gold dissolution is observed (in Figure 7.6) can be explained in terms of this intersection point. With increasing NaCl concentration, the predominant domain of AuC14 shifted towards a higher pH. This means that industrial design problems, such as corrosiveness of acidic chloride solution and emission of OC1 in highly acidic (pH = 1 to 3) conditions, should be minimised by the use of a high NaCl (3.419 M) concentration.

At a higher OC1 (0.194 M) concentration, gold-sulfide minerals in chloride systems (with the exception of gold-pyrite system), show that the equilibrium concentrations of predominant species with respect to pH at 25°C are similar to those of the Au - Cl" - NaOCl -H20 system. In the gold-pyrite system, FeS2 is readily dissolved in the presence of OCF as oxidant. With increasing OC1 concentration, this situation was not observed to change greatly. At a lower OC1 concentration (0.019 M), the variation of AuC14 with respect to pH follows almost the same trend in all gold-sulfide systems. The species 236 considered were AuS", AuC14’, AuC12 , S042 ,and H2S as the aqueous components and Au as the solid. These results indicate that as the concentration of OC1 decreases the gold chloride complexes in the presence of sulfide minerals can be reduced significantly to metallic gold. The results of the precipitation metallic gold from gold complexes, as shown in section 7.8 in Chapter Seven, can be attributed by both the decrease of OC1 concentration and the presence of sulfide minerals. In this Chapter, the results of thermodynamic study revealed that with increasing NaCl concentration, the predominant domain of AuC14 shifted towards a higher pH. This means that industrial design problems due to the use of the chloride-hypochlorite system can be minimised by the use of higher NaCl (3.419 M) concentration. From analysis based on the gold-sulfide system, as the concentration of OC1 decreases the gold chloride complexes in the presence of sulfide minerals should be reduced significantly to metallic gold. It should be apparent that Au loss in solutions can be attributed by the decrease of OCl concentration as well as the presence of sulfide minerals. CHAPTER TEN

CONCLUSIONS AND RECOMMENDATIONS

10.1 CONCLUSIONS

1) The results of an experimental study on the kinetics of dissolution and stability of gold in a chloride-hypochlorite system have been presented and discussed. The kinetic study included measurements of the gold dissolution rates. A number of experiments were carried out to determine the effect of various parameters such as pH, temperature, chloride and hypochlorite concentrations on the kinetics of the dissolution of gold in a chloride system. In all these investigations, the rate of gold dissolution was measured for solutions containing NaCl (50 to 200 g/1), and OCT (1 to 10 g/1) within the pH range from 2 to 6. The dissolution of gold in these chloride-hypochlorite systems is controlled by a kinetic step. An activation energy of 28.7 kJ/mole was determined using a solution of 200 g/1 NaCl and 10 g/1 OC1 at pH 4. The reaction order with respect to the activity of [H+] for a solution of 200 g/1 NaCl was found to be 0.3-0.35 in the pH range from 4 to 6 and zero at pH 2 and 3. The reaction order with respect to NaCl concentration for a solution of 10 g/1 OC1 is 0.5 in the pH range from 4 to 6 and -0.3 at pH 2 and 3. At a pH 2-3, as the concentration of [H+] ion increases, the dissolution rate decreases significandy. One possible cause might be mass-transfer limitations. The reaction order with respect to OC1" is almost 1 at pH 4-6 and zero at pH 2-3. By 238 taking [H+] = 10'4-10'\ [NaCl] = 50-200 g/1, [OC1] = 1-10 g/1, and Ea = 28.7 kJ/mol, the reaction rate equation at the gold-solution interface was derived from the experimental results, as follows:

If a shrinking core model with no product layer at the interface is adopted, the following rate equation illustrating the relationship between the fraction reacted (R) and time (/) can be expressed as:

1 - (1 - R)v = ^(-mA]UB]L,[C]'coJx + y +1)/ 32f a

On the basis of the above shrinking core model developed for application to thin rectilinear geometry, the reaction kinetics can be represented as:

1 -(1 -R)V3 = k,exp]°3[NaCl]0S[OCl~]111

In acidic solutions, tests to determine the stability of gold chloride solutions in the presence of sulfide minerals such as sphalerite, galena, pyrite and chalcopyrite were also conducted. The stability of gold in the presence of sulfide minerals is strongly dependent on pH and the gold chloride complex could be readily produced to metallic gold in the presence of sulfide minerals. The overall reactions for the precipitation of gold-chloride complexes to metallic gold in the presence of sulfide minerals is proposed as follows: 239

For sphalerite,

AuCl. +-ZnS + xH?0-> AU + -Z11CI2- +—cr +-SO4- +2xH' 4 4 2 4 4 3 4 4 1 For galena,

y v ?3x x /l«a, + -PbS + xH,0 -> A,i + -PbCI: +----CT‘ + -,S,02~ + 2x//+ 4 4 2 4 3 12 4 4 For pyrite,

And; +-FeS,+xH,0-> Au+-Fe2* +—CF +-S02 +2xH+ 4 8 2 2 8 2 4 4 For chalcopyrite,

AuCl: +-CuFeS' +xH,0-> Au + —Cu,S + 4Cr +~Fe2* +-SOj~ +2xH+ 4 6 2 2 12 2 6 4 4

2) Electrochemical techniques based on scanning voltammetry were used to investigate the half-cells of the electrochemical reactions in chloride-oxidant solutions. It was shown that for pH values between 2 and 6, an increase in the pH shifted the anodic polarization curves towards the more positive direction. The rate of the anodic

reaction increased with an increase of the Cl concentration in the

range of 0.855 to 3.419M at pH 6. The study of the cathodic

polarization of gold in deoxygenated 1.709 M NaCl and 10 mM OC1’

solutions, showed that the rate of hydrogen evolution at a given potential decreased at an increasing pH between 4 and 8. Increasing

OCF concentrations shifted the potential in the positive direction.

Evans diagrams obtained from the anodic and cathodic half-cell reactions required for the dissolution of gold in chloride solutions at various pH values indicated that the oxidation of gold by hypochlorite

is controlled by a kinetic step. The reaction rates measured on several stationary sulfide electrodes can be represented (for sulfide mineral oxidation in sodium hypochlorite solutions) by applying the Butler- Volmer equation, as follows: 240

For sphalerite,

FE. = -^4i_e[OC/_]exp at 4 167*77

For galena,

Fr^ = ~Ak\OCr]exp\ - 9 RTJ

For pyrite

1r--V5A‘k-‘locrH-mj

For chalcopyrite

df1cuFes2 _ 2 [(9C/~1expf dt 11 c ~c J Pl 15RTJ

3) The speciation at equilibrium using the CHEMIX computer modelling

program was used to construct the log [M]-pH diagrams. It was found

that AuC14’ is the predominant species in the pH range of 1 to 5.5 for a

solution of [Au] = 1 mM, [NaCl] = 1.709 M and [OCl ] = 0.194 M.

Above pH 5.5, the gold chloride complex, AuC14 , becomes unstable

and is produced to metallic gold. With increasing NaCl concentration,

the predominant domain of AuC14 is expanded toward higher pH

regions. The Au - ZnS - Cf - NaOCl - H20 system showed that the

concentration profiles of oxidizing species at a higher OCF

concentration (0.194 M) was similar to those of the Au - Cl - NaOCl -

H20 system at the same OCF concentration. However, at a lower

OCF concentration (0.019 M), AuC14 species disappeared and AuS

appeared. AuS species in the solutions with decreasing OC1

concentration are formed over a wide range of pH. The variation of

AuC14 concentration in solution with respect to pH in Au - PbS - Cl -

NaOCl - H20 system and Au - CuFeS2 - CF - NaOCl - H20 system 241

was observed to follow almost the same tendency as the gold-sphalerite system. However, in the gold-pyrite system, FeS2 readily dissolved in

the presence of OCF. The AuC14 species may be altered to AuS . With increasing 00 concentration, this situation has not changed

greatly. Based on the observed results, the reactions involving AuC14 species in the presence of AuS as an intermediate species, can be proposed, as follows: For sphalerite, 3 AuCl4 + 3 ZnS + OH~ + 3 H20 =

2 An + AuS' + 3 ZnCl]- + SO]- + H2S + 5 H+ For galena, 3 AuCl4 + 3 PbS + OH~ + 3H20 =

2 An + AuS- + 3PbCl3 + SOI + H2S + + 5H+ For pyrite,

8 AuCIa + 8 FeS2 + 2 OH~ +12 H20 =

4 An + 4 AuS' + 8 Fe3+ + 2SO\~ + 2 H2S + 8.St2- + 32 Cl~ + 302 + 22H+ For chalcopyrite, 8 Audi + 8 CuFeS2 + 2 OH~ +12 H20 =

4 Au + A AuS' + 4 Cu2S + 8F^2+ + 2SO]' +2 H2S + 4S2~ + 32 CH + 302 + 22H+

4) The results of this study revealed that the predominant domain of

AuC14' shifted towards a higher pH as NaCl concentration increases. This behaviour can be applied to minimise industrial design problems due to the use of chloride-hypochlorite system. The kinetics of gold oxidation and precipitation by chloride-hypochlorite system was dependent on the concentration of OCF as well as the presence of sulfide minerals. 242

10.2 SUGGESTIONS FOR FURTHER WORK

The studies conducted in this work aimed at providing data for a detailed scientific evaluation of the dissolution of gold in chloride solutions. Despite the detailed investigation of the dissolution and stability of gold, a number of questions, such as effect of silica and sulfide (or polysulfide) ion, remain unanswered. Therefore, the following areas of work are recommended for further investigation:

1) An economical process to treat high sulfide ores requires some form of preoxidation processes prior to chlorination because the consumption of chlorine is excessive. Efficient pretreatment processes such as roasting and chemical oxidation, for the treatment of gold ores which have high sulfide contents, should be evaluated experimentally. A study of carbon-in-chlorine-leach (C1CL) technology, by which gold can be simultaneously recovered via leaching and precipitation of the pretreated pulp by the chlorination of refractory sulfidic ores, is recommended to determine an economical process.

2) The corrosiveness of acidic chloride solutions, is the main concern in industrial applications. At neutral pH, hypochlorite solutions undergo a slow degradation to the chloride ion. Further studies on hypochlorite treatment prior to cyanidation with respect to heap leaching at pH levels above neutral as well as solvent extraction and carbon adsorption techniques to recover gold from chloride solutions are recommended. 243

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APPENDICES 289

Appendix 7.1 Dissolution rate of gold conducted at pH 2 to 6 under various conditions.

Gold Dissolution Rate (RA), mg/(cm2 hr) NaCl 50 g/1 NaCl 100 g/1 NaCl 200 g/1 pH ocr R, r2 K r2 R, r2 2 1 g/1 10.44 0.970 12.36 0.919 10.74 0.959 2 g/1 24.36 0.956 25.50 0.944 22.38 0.969 5 g/1 49.98 0.984 52.68 0.980 44.94 0.966 10 g/1 65.76 0.985 56.04 0.982 44.82 0.962 3 1 g/1 9.84 0.982 11.28 0.922 9.78 0.984 2 g/1 16.92 0.986 22.92 0.978 19.56 0.982 5 g/1 52.86 0.994 48.66 0.992 44.10 0.977 10 g/1 69.96 0.993 63.66 0.984 46.50 0.983 4 1 g/1 4.32 0.993 5.82 0.995 4.98 0.996 2 g/1 7.08 0.996 9.66 0.996 10.50 0.994 5 g/1 18.84 0.998 20.10 0.995 27.96 0.996 10 g/1 23.22 0.999 32.88 0.999 56.46 0.974

5 lg/I 0.72 0.985 1.56 0.992 1.86 0.985 2 g/1 1.74 0.980 2.94 0.993 4.44 0.984 5 g/1 4.80 0.999 5.82 0.996 9.12 0.998 10 g/1 9.36 0.998 15.36 0.999 18.12 0.980

6 1 g/1 0.36 0.939 0.66 0.989 0.90 0.992 2 g/1 0.78 0.993 1.44 0.998 2.58 0.992

5 g/1 3.18 0.999 4.26 0.999 6.48 0.995 10 g/1 5.52 0.999 11.88 0.996 11.34 0.998 r is the correlation coefficient as obtained from each linear regression line 290

Appendix 7.2 Kinetic data for the dissolution of gold in the chloride- hypochlorite system at pH 2 to 6.

Rate Constants (ka), min NaCl 50 g/1 NaCl 100 g/1 NaCl 200 g/1 -Ae ■A* o © X X Lr\ X pH ocr r2 o r2 r2

2 i g/i 77.2 0.975 142.9 0.929 84.6 0.965 2 g/1 193.4 0.968 349.9 0.966 176.0 0.979

5 g/1 380.5 0.997 538.3 0.998 547.4 0.991 10 g/1 505.0 0.999 562.7 0.997 688.5 0.994

3 1 g/1 84.2 0.986 114.2 0.932 85.4 0.987 2 g/1 228.3 0.994 215.0 0.986 235.0 0.991 5 g/1 343.8 0.999 526.5 0.999 643.2 0.998 10 g/1 551.6 0.992 538.0 0.999 602.4 0.997

4 1 g/1 31.3 0.993 65.1 0.995 71.2 0.996 2 g/1 46.9 0.997 173.7 0.997 113.1 0.996 5 g/1 115.0 0.999 463.7 0.999 213.2 0.999 10 g/1 224.3 0.998 248.9 0.999 549.6 0.993

5 1 g/1 5.1 0.986 16.2 0.992 28.5 0.985 2 g/1 11.5 0.981 68.0 0.994 65.6 0.982 5 g/1 77.8 0.999 96.3 0.998 105.0 0.997 10 g/1 56.7 0.998 126.3 0.998 179.1 0.988

6 1 g/1 2.3 0.940 6.4 0.989 11.9 0.992 2 g/1 11.7 0.993 29.4 0.998 51.0 0.992 5 g/1 44.2 0.999 79.0 0.999 95.4 0.993 10 g/1 84.5 0.999 153.9 0.998 105.6 0.998 r is correlation coefficient obtained for each regression line 291

Appendix 7.3 The reproducibility for the effect ofpH on An loss; [NaClJ = 200 g/l, fAuClf] = approximately 10 mg/l, [ZnS] = 4 g/l and [silica] = 200 g/l (refer to Figure 7.32)

Time X SD n t t ■ sd/ 47i 95% upper 95% lower (min) confidence limits confidence limits

5 71.32 23.17 5 2.78 28.81 100.12 42.51

10 88.51 11.60 5 2.78 14.42 102.93 74.09

15 89.04 10.92 5 2.78 13.58 102.62 75.46

20 93.23 5.36 5 2.78 6.67 99.90 86.57

30 90.97 6.24 5 2.78 7.76 98.73 83.21

45 94.03 7.76 5 2.78 9.65 103.68 84.38

60 96.68 2.80 5 2.78 3.49 100.16 93.19 292

Appendix 7.4 The reproducibility for the effect of sulfide minerals on Au loss; [NaCIJ = 200 g/l, [AuClT] = approximately 10 mg/l, [sulfide minerals] = 4 g/l, respectively, [silica] = 200 g/l at pH 4 (refer to Figure 7.34)

Time X SD n t t ■ SD/-Jit 95% upper 95% lower (min) confidence limits confidence limits

5 81.28 8.72 4 3.18 13.86 95.13 67.42

10 88.16 5.53 4 3.18 8.79 96.96 79.37

15 93.00 6.92 4 3.18 11.00 104.00 82.00

20 96.74 2.18 4 3.18 3.46 100.20 93.28

30 99.54 0.62 4 3.18 0.99 100.52 98.55

45 100.00 0.00 4 3.18 0.00 100.00 100.00

60 100.00 0.00 4 3.18 0.00 100.00 100.00 293

Appendix 8.1 Mixed current densities and mixed potentials measured with

the stationary sphalerite electrode for some chloride-oxidant

solutions at pH 6 (refer to Figure 8.26).

Solution Mixed potential Mixed current density

concentration Em vs SCE mA/cm2

(V)

1.709M NaCl - 0.047 0.269

4mM OCf 0.194 0.537

10mM OCf 0.227 0.585

19mM OC1' 0.274 0.659

39mM OCf 0.389 0.868 294

Appendix 8.2 Mixed current densities and mixed potentials measured with

the stationary galena electrode for some chloride-oxidant

solutions at pH 6 (refer to Figure 8.28).

Solution Mixed potential Mixed current density

concentration Em vs SCE mA/cm2

(V)

1.709M NaCl -0.188 0.314

4mM OCT 0.350 0.327

lOmM OCT 0.381 0.368

19mM OCf 0.460 0.486

39mM OC1' 0.545 0.774 295

Appendix 8.3 Mixed current densities and mixed potentials measured with

the stationary pyrite electrode for some chloride-oxidant

solutions at pH 6 (refer to Figure 8.30).

Solution Mixed potential Mixed current density

concentration Em vs SCE mA/cm2

(V)

1.709M NaCl - 0.205 0.326

4mM OCf 0.322 0.440

lOniM OCf 0.346 0.490

19mM OCf 0.405 0.700

39mM OCf 0.490 1.019 296

Appendix 8.4 Mixed current densities and mixed potentials measured with

the stationary chalcopyrite electrode for some chloride-

oxidant solutions at pH 6 (refer to Figure 8.32).

Solution Mixed potential Mixed current density

concentration Em vs SCE in A/cm2

(V)

1.709M NaCl -0.168 0.301

4mM OCf 0.314 0.382

lOmM OC1' 0.346 0.418

19mM OCf 0.430 0.526

39mM OCf 0.550 0.709 297

Appendix 9.1 Thermodynamic data of components on the Au-CT-OCT -H20 system at 25 XI (GarreIs and Christ, 1965; Wagman et al., 1982)

Formula State G° kJ/mol Formula State G° kJ/mol

o2 g 0 NaCIO ai -298.700 n2 g 0 C102" aq 17.200 H 2 g 0 CIO' aq -36.800

Cl2 g 0 cio2 aq 120.100 HC1 g -95.299 CKV aq -7.950

H20 g -228.572 C104' aq -8.520 h o aq -237.129 2 Na+ aq -261.905 OH' aq -157.245 HC1 ai -131.228 0 H+ aq HAuC14 ai -235.140

AuC12 aq -151.120 S042' aq -744.495

AuC14 aq -235.140 hso4' aq -755.866

Au033 aq -51.800 h2so4 aq -744.495

HAu032 aq -142.200 Au s 0

H2Au03 aq -218.300 Au(OH)3 s -316.920

Au(OH)3 aq -283.370 AuCl s -15.035

Cl' aq -131.228 AuC13 s -45.327

Cl2 aq 6.940 NaCl s -384.138 ci; aq -120.400 HC10 aq -79.900 298

Appendix 9.2 Thermodynamic data of components on the An-ZnS-CT-OCr~ H20 and the Au-PbS-Cr~OCr-H20 system at 25 °C (GarreIs

and Christ, 1965; Wagman et al.t 1982; Escudero et al., 1993; Caspar et al., 1994)

Formula State G° kJ/mol Formula State G° kJ/mol

S2 g 0 Pb(OH)3" aq -575.600 h2so4 g 0 S2' aq 85.800 H S' aq 12.080 S22' aq 79.500 H2S aq -27.830 s32- aq 73.700 Na2S ai -438.100 so2 aq -300.676 Zn2+ aq -147.060 S2032' aq -522.500 -403.700 ZnCl2 aq s2042' aq -600.300 ZnCl+ aq -275.300 s2082- aq -1114.900 aq -540.500 ZnCl3 s4062‘ aq -1040.400 ZnCl42 aq -666.000 Au203 s 0 ZnC104+ aq -163.100 Au2S s 29.037 ZnCl208 ai -164.100 Zn s 0 Zn022 aq -384.240 ZnS s -201.290 HZn02 aq -457.080 ZnS04 s -871.500 ZnOH+ aq -330.100 Pb s 0 Zn(OH)3 aq -694.220 PbS s -98.700 Zn(OH)42" aq -858.520 PbS04 s -813.140 AuS' aq 46.024 PbCl2 s -314.100 ZnS04+ aq -904.900 Pb02 s -217.330 ZnS04 ai -891.590 Pb(OH)2 s -452.200 Pb2+ aq -24.430 PbO s -187.890 PbCl2 aq -297.160 Pb304 s -601.200 PbCl+ aq -164.810 S s 0 PbCl3’ aq -426.300 Na2S s -349.800 HPbOj aq -338.420 NaOH s -379.494 PbOH+ aq -226.300 299

Appendix 9.3 Thermodynamic data of components on the Au-FeS2-Cr~OCr- H20 and the Au-CuFeS2-Cr~0Cr-H20 system at 25°C (Garrels and Christ, 1965; Wagman et al., 1982; Escudero et al., 1993; Gaspar et al., 1994)

Formula State G° kJ/mol Formula State G° kJ/mol

Fe3+ aq -4.700 Fe(OH)42 aq -769.700 Fe2+ aq -78.900 Fe s 0 FeCl2 aq -279.100 FeCl2 s -302.300 FeCl3 aq -404.500 FeCl3 s -334.000 FeCl2+ aq -143.900 Fe203 s -742.200 Fe022 aq -295.300 Fe304 s 0 HFe02‘ aq -377.700 FeS s 29.037 FeS04+ aq -772.700 FeS2 s -167.000 FeS04 ai -823.430 FeS04 s -201.290 s -1270.160 FeS208 aq -1524.500 Na2S04 Cu s 0 NaS04 aq -1010.610 -98.700 Cu2+ aq 65.490 CuCl s -175.700 CuCl2 aq -197.900 CuCl2 s -129.700 CuCl+ aq -68.200 CuO s Cu20 s -146.000 CuCl2’ aq -240.100 Cu2S s -86.200 CuCl32' aq -376.000 CuS s -53.600 Cu022 aq -183.600 CuS04 s -661.800 Cu(OH)2 aq -249.010 CuFeS2 s -227.390 CuS04 aq -692.180 FeO s -245.120 -481.100 CuS03 aq Fe(OH)3 s -696.500 Fe(OH)3 aq -614.900 300

Appendix 9.4 The equilibrium concentration profiles of the predominant Au species as a function of pH; [Au] = 1 mM, [NaCl] = 1.709 M and [OCT] = 0.194 M (refer to Figure 9.1a and b).

pH log[AuCI2] log[AuCI4] log[CIO] log[CI2(aq)] log[HCIO] log[CI03‘] log[Au] 12.63 -11.3515 -26.7811 -17.2722 -31.3653 -22.3466 -24.2610 -3.0000 12.02 -10.7461 -24.9649 -17.2722 -30.1542 -21.7409 -24.2610 -3.0000 11.53 -10.2529 -23.4852 -17.2722 -29.1677 -21.2476 -24.2610 -3.0000 10.92 -9.6426 -21.6542 -17.2722 -27.9470 -20.6372 -24.2610 -3.0000 10.50 -9.2276 -20.4093 -17.2722 -27.1171 -20.2223 -24.2610 -3.0000 10.02 -8.7505 -18.9779 -17.2722 -26.1628 -19.7451 -24.2610 -3.0000 9.17 -7.8971 -16.4178 -17.2722 -24.4561 -18.8917 -24.2610 -3.0000 9.06 -7.7847 -16.0807 -17.2722 -24.2313 -18.7794 -24.2610 -3.0000 8.67 -7.3988 -14.9230 -17.2722 -23.4595 -18.3935 -24.2610 -3.0000 8.16 -6.8850 -13.3815 -17.2722 -22.4319 -17.8797 -24.2610 -3.0001 7.70 -6.4240 -11.9985 -17.2722 -21.5099 -17.4187 -24.2610 -3.0002 7.05 -5.7743 -10.0492 -17.2722 -20.2104 -16.7689 -24.2610 -3.0007 6.65 -5.3789 -8.8633 -17.2722 -19.4197 -16.3736 -24.2611 -3.0018 6.45 -5.1743 -8.2494 -17.2723 -19.0105 -16.1690 -24.2611 -3.0029 5.52 -4.2446 -5.4602 -17.2727 -17.1510 -15.2395 -24.2625 -3.0299 5.01 -3.7440 -3.9575 -17.2788 -16.1490 -14.7413 -24.2798 -3.2880 4.48 -4.3001 -3.4633 -17.2899 -15.0986 -14.2213 -24.3120 - 3.95 -5.3476 -3.4430 -17.2905 -14.0307 -13.6877 -24.3137 - 3.51 -6.2184 -3.4413 -17.2906 -13.1583 -13.2515 -24.3139 - 3.00 -7.2347 -3.4412 -17.2909 -12.1419 -12.7433 -24.3143 - 2.46 -8.3153 -3.4415 -17.2920 -11.0616 -12.2032 -24.3153 - 2.13 -8.9733 -3.4421 -17.2938 -10.4042 -11.8745 -24.3171 - 1.95 -9.3370 -3.4427 -17.2954 -10.0411 -11.6931 -24.3188 - 1.52 -10.1860 -3.4457 -17.3037 -9.1951 -11.2703 -24.3270 - 1.01 -11.1523 -3.4560 -17.3314 -8.2391 -10.7931 -24.3546 - 301

Appendix 9.5 The equilibrium concentration profiles of the predominant Au

species as a function of pH; [ZnSj = 41.05 mM, [AuCl4f = 50

pM, [NaClJ = 1.709 M and [OCT] = 0.194 M (refer to Figure

9.3a and b).

pH log[AuCI2 ] log[AuCI4] log[ZnCI421 log[S042] log[CIO ] log[Cl2(aa)] logfHClOl log[Au] 12.79 -11.4000 -26.9957 -11.4770 -1.6928 -17.2306 -31.5471 -22.3812 -4.3010 12.51 -11.1343 -26.1987 -10.6261 -1.4553 -17.2715 -31.0153 -22.1354 -4.3010 12.13 -10.8454 -25.3319 -10.5697 -1.2518 -17.4511 -30.4366 -21.9355 -4.3010 11.98 -10.7034 -24.9059 -10.7088 -1.2385 -17.4668 -30.1525 -21.8014 -4.3010 10.98 -9.7117 -21.9309 -11.7122 -1.2218 -17.4722 -28.1691 -20.8123 -4.3010 10.43 -9.1555 -20.2622 -12.2696 -1.2207 -17.4722 -27.0567 -20.2560 -4.3010 10.11 -8.8367 -19.3060 -12.5886 -1.2204 -17.4722 -26.4192 -19.9373 -4.3010 9.50 -8.2348 -17.5001 -13.1907 -1.2203 -17.4722 -25.2153 -19.3353 -4.3011 9.03 -7.7587 -16.0717 -13.6669 -1.2202 -17.4722 -24.2630 -18.8592 -4.3012 8.52 -7.2466 -14.5356 -14.1790 -1.2202 -17.4722 -23.2389 -18.3471 -4.3015 8.11 -6.8412 -13.3195 -14.5843 -1.2202 -17.4722 -22.4282 -17.9418 -4.3023 7.73 -6.4631 -12.1850 -14.9625 -1.2202 -17.4722 -21.6719 -17.5636 -4.3040 7.35 -6.0849 -11.0506 -15.3406 -1.2202 -17.4722 -20.9156 -17.1855 -4.3082 7.11 -5.8420 -10.3217 -15.5836 -1.2202 -17.4722 -20.4297 -16.9426 -4.3137 6.95 -5.6798 -9.8351 -15.7458 -1.2202 -17.4722 -20.1052 -16.7803 -4.3196 6.53 -5.2651 -8.5910 -16.1605 -1.2202 -17.4722 -19.2759 -16.3657 -4.3510 6.14 -4.8750 -7.4209 -16.5506 -1.2202 -17.4723 -18.4958 -15.9756 -4.4371 5.23 -4.5056 -5.2158 -17.4687 -1.2201 -17.4732 -16.6601 -15.0582 - 4.80 -5.0197 -4.8828 -17.8923 -1.2201 -17.4737 -15.8130 -14.6349 - 4.54 -5.4881 -4.8198 -18.1580 -1.2201 -17.4738 -15.2817 -14.3693 - 4.06 -6.4182 -4.7940 -18.6358 -1.2204 -17.4739 -14.3257 -13.8914 - 3.51 -7.5119 -4.7909 -19.1834 -1.2216 -17.4740 -13.2290 -13.3430 - 3.00 -8.5260 -4.7908 -19.6884 -1.2252 -17.4743 -12.2147 -12.8359 - 2.47 -9.5850 -4.7911 -20.2109 -1.2372 -17.4753 -11.1561 -12.3066 - 2.02 -10.4890 -4.7920 -20.6486 -1.2645 -17.4780 -10.2529 -11.8551 - 1.51 -11.4857 -4.7950 -21.1256 -1.3293 -17.4870 -9.2592 -11.3587 - 1.47 -11.5656 -4.7954 -21.1647 -1.3364 -17.4883 -9.1797 -11.3189 - 302

Appendix 9.6 The equilibrium concentration profiles of the predominant An species as a function of pH; [ZnS] = 41.05 mM, [Audi] = 50 pM, [Nad] = 1. 709 M and ]Od'] = 0.019 M (refer to Figure 9.4a and b).

PH log[AuCI2] logfAuCU-] log[AuS] log[H2S(aa)] logfHS'l lofl[S2l log[S042] log[Au] 13.29 -28.7648 -79.0066 -13.0398 -17.0375 -10.6639 -10.3039 -2.6127 -4.3010 13.03 -28.2782 -77.5469 -13.4806 -17.4636 -11.3410 -11.2318 -1.6490 -4.3010 12.55 -27.5487 -75.3581 -14.4541 -18.1953 -12.5587 -12.9355 -1.3986 -4.3010 12.18 -27.0197 -73.7711 -15.1413 -18.6701 -13.4042 -14.1518 -1.3474 -4.3010 11.71 -26.3636 -71.8028 -15.9554 -19.2015 -14.4051 -15.6221 -1.3239 -4.3010 9.92 -23.9268 -64.4909 -18.7127 -20.8248 -17.8146 -20.8179 -1.3016 -4.3010 9.24 -23.1781 -62.2410 -18.4980 -19.9968 -17.6695 -21.3558 -1.2821 -4.3010 9.01 -22.9506 -61.5581 -18.2901 -19.5662 -17.4643 -21.3757 -1.2814 -4.3010 8.55 -22.4839 -60.1578 -17.8380 -18.6510 -17.0140 -21.3905 -1.2810 -4.3010 8.13 -22.0622 -58.8927 -17.4202 -17.8125 -16.5967 -21.3944 -1.2809 -4.3010 7.63 -21.5639 -57.3979 -16.9220 -16.8161 -16.0986 -21.3945 -1.2809 -4.3010 7.04 -20.9802 -55.6468 -16.3320 -15.6408 -15.5077 -21.3881 -1.2811 -4.3010 6.83 -20.7645 -54.9998 -16.1104 -15.2021 -15.2854 -21.3822 -1.2812 -4.3010 5.96 -19.9027 -52.4156 -15.1639 -13.3730 -14.3280 -21.2964 -1.2831 -4.3010 5.57 -19.5321 -51.3046 -14.6884 -12.5009 -13.8390 -21.1905 -1.2850 -4.3010 5.26 -19.2317 -50.4043 -14.2518 -11.7301 -13.3850 -21.0534 -1.2869 -4.3010 4.93 -18.9247 -49.4845 -13.7548 -10.8789 -12.8639 -20.8623 -1.2887 -4.3010 4.53 -18.5596 -48.3901 -13.1072 -9.7957 -12.1806 -20.5789 -1.2904 -4.3010 4.10 -18.1743 -47.2345 -12.3797 -8.5974 -11.4102 -20.2363 -1.2916 -4.3010 3.46 -17.6080 -45.5356 -11.2717 -6.7874 -10.2343 -19.6946 -1.2935 -4.3010 3.04 -17.2261 -44.3896 -10.5153 -5.5545 -9.4308 -19.3205 -1.2963 -4.3010 2.58 -16.7137 -42.8507 -10.3952 -5.0339 -9.3677 -19.7150 -1.2359 -4.3010 2.01 -15.9626 -40.5914 -11.1524 -5.4046 -10.3103 -21.2294 -1.2685 -4.3010 1.74 -15.6193 -39.5543 -11.5028 -5.5723 -10.7445 -21.9303 -1.2999 -4.3010 303

Appendix 9.7 The equilibrium concentration profiles of the predominant Au species as a function of pH; [PbS] = 16.72 mM, [AuCU] = 50 pM, [NaClJ = 1.709 M and [OCl'] = 0.194 M (refer to Figure 9.5a and b).

PH log[AuCI2'l logfAuCI4] log[PbCI3'l log[S042 l log[CIO ] log[CI2(aa,l log[HCIO] log[Aul 13.25 -11.7073 -27.9177 -16.0358 -1.7768 -16.9291 -32.1621 -22.5382 -4.3010 13.01 -11.4701 -27.2061 -15.5617 -1.2863 -16.9291 -31.6871 -22.3003 -4.3010 12.59 -11.0471 -25.9370 -14.7159 -1.0877 -16.9291 -30.8406 -21.8767 -4.3010 12.24 -10.7010 -24.8986 -14.0237 -1.0375 -16.9291 -30.1481 -21.5303 -4.3010 11.83 -10.2927 -23.6738 -13.2072 -1.0145 -16.9291 -29.3315 -21.1220 -4.3010 11.41 -9.8678 -22.3992 -12.3575 -1.0056 -16.9291 -28.4818 -20.6971 -4.3010 11.01 -9.4726 -21.2135 -11.5670 -1.0024 -16.9291 -27.6913 -20.3018 -4.3010 10.49 -8.9482 -19.6404 -10.5183 -1.0010 -16.9291 -26.6425 -19.7774 -4.3010 9.98 -8.4401 -18.1161 -9.5021 -1.0005 -16.9291 -25.6264 -19.2694 -4.3011 9.04 -7.4971 -15.2872 -7.6162 -1.0003 -16.9291 -23.7404 -18.3264 -4.3013 7.87 -6.3241 -11.7682 -5.2702 -1.0003 -16.9291 -21.3944 -17.1534 -4.3052 7.48 -5.9367 -10.6056 -4.4952 -1.0000 -16.9291 -20.6191 -16.7656 -4.3112 6.02 -4.4836 -6.2465 -4.4953 -0.9999 -16.9292 -17.7130 -15.3125 -4.8122 5.60 -4.4422 -5.3476 -4.4954 -0.9998 -16.9292 -16.8555 -14.8838 - 5.10 -4.9783 -4.8933 -4.4954 -0.9998 -16.9293 -15.8651 -14.3886 - 4.52 -6.0389 -4.7988 -4.4952 -1.0001 -16.9293 -14.7100 -13.8111 - 4.02 -7.0423 -4.7916 -4.4943 -1.0011 -16.9293 -13.6994 -13.3058 - 3.50 -8.0733 -4.7909 -4.4913 -1.0043 -16.9294 -12.6676 -12.7899 - 3.10 -8.8841 -4.7909 -4.4849 -1.0114 -16.9297 -11.8569 -12.3846 - 2.60 -9.8726 -4.7911 -4.4640 -1.0345 -16.9304 -10.8686 -11.8905 - 2.02 -11.0266 -4.7921 -4.3975 -1.1101 -16.9334 -9.7155 -11.3141 - 1.51 -12.0272 -4.7950 -4.2939 -1.2407 -16.9424 -8.7179 -10.8156 - 1.45 -12.1439 -4.7957 -4.2814 -1.2589 -16.9443 -8.6019 -10.7577 - 304

Appendix 9.8 The equilibrium concentration profiles of the predominant Au species as a function ofpH; [PbSJ = 16.72 mM, [AuCU] =50 pM, [NaClj = 1.709 M and [00'] = 0.019 M (refer to Figure 9.6a and b). o > PH log[AuCI2] log[AuCI4] logfAuS'l log[H2S(aa)] logfHSl iog[S 1 log[S042'l CO 13.31 -28.8283 -79.1970 -12.4557 -16.4302 -10.0365 -9.6562 -2.3117 -4.3010 13.03 -28.3536 -77.7730 -12.5381 -16.4288 -10.3147 -10.2141 -1.3021 -4.3010 12.50 -27.6721 -75.7285 -12.9005 -16.4276 -10.8363 -11.2585 -1.0712 -4.3010 12.03 -27.0729 -73.9309 -13.2502 -16.4273 -11.3107 -12.2076 -1.0228 -4.3010 11.04 -25.8380 -70.2260 -13.9871 -16.4272 -12.2966 -14.1795 -1.0026 -4.3010 10.56 -25.2335 -68.4126 -14.3495 -16.4272 -12.7800 -15.1463 -1.0011 -4.3010 9.99 -24.5229 -66.2807 -14.7758 -16.4272 -13.3485 -16.2832 -1.0005 -4.3010 9.65 -24.1010 -65.0151 -15.0289 -16.4272 -13.6860 -16.9582 -1.0004 -4.3010 8.10 -22.1594 -59.1903 -16.1938 -16.4272 -15.2392 -20.0647 -1.0003 -4.3010 7.47 -21.3785 -56.8477 -16.6623 -16.4271 -15.8639 -21.3142 -1.0000 -4.3010 6.09 -19.9959 -52.6997 -15.2796 -13.6618 -14.4813 -21.3142 -1.0000 -4.3010 5.50 -19.3999 -50.9117 -14.6837 -12.4698 -13.8853 -21.3143 -1.0001 -4.3010 5.10 -19.0022 -49.7187 -14.2861 -11.6745 -13.4877 -21.3144 -1.0001 -4.3010 4.53 -18.4338 -48.0135 -13.7180 -10.5381 -12.9197 -21.3147 -1.0005 -4.3010 4.12 -18.0284 -46.7972 -13.3132 -9.7278 -12.5149 -21.3154 -1.0012 -4.3010 3.08 -16.9869 -43.6720 -12.2823 -7.6548 -11.4840 -21.3266 -1.0124 -4.3010 2.59 -16.4968 -42.2002 -11.8149 -6.6958 -11.0166 -21.3508 -1.0366 -4.3010 2.01 -15.9213 -40.4675 -11.3127 -5.6122 -10.5145 -21.4303 -1.1161 -4.3010 1.52 -15.3287 -38.6724 -11.8036 -5.7364 -11.1271 -22.5313 -1.2468 -4.3010 1.50 -15.3020 -38.5908 -11.8317 -5.7487 -11.1614 -22.5875 -1.2537 -4.3010 305

Appendix 9.9 The equilibrium concentration profiles of the predominant Au species as a function of pH; [FeS2] = 33.40 mM, [AuCl4~] = 50 juM, [NaClJ = 1.709 M and [OCT] = 0.194 M (refer to Figure 9.7a and b).

pH log[AuCI2'] log[AuCI4] log[AuS] log[Fe3*] log[H2S(aq)] log[HS~1 logfS2'] logrSQ42] logfAul 13.00 -29.0914 -80.0701 -6.8208 -41.0619 -9.8324 -3.7470 -3.6750 -1.5761 -4.3023 12.56 -28.5024 -78.3030 -7.2137 -39.6853 -9.9298 -4.2868 -4.6573 -1.3817 -4.3016 12.17 -28.0005 -76.7973 -7.5651 -38.4799 -10.0048 -4.7511 -5.5109 -1.3324 -4.3013 11.51 -27.1647 -74.2899 -8.1444 -36.4851 -10.1080 -5.5103 -6.9260 -1.3080 -4.3011 11.00 -26.5185 -72.3514 -8.5858 -34.9561 -10.1763 -6.0882 -8.0136 -1.3034 -4.3011 10.52 -25.9082 -70.5203 -9.0032 -33.5109 -10.2406 -6.6343 -9.0414 -1.3020 -4.3010 10.03 -25.2857 -68.6530 -9.4290 -32.0369 -10.3061 -7.1911 -10.0896 -1.3016 -4.3010 9.50 -24.6234 -66.6661 -9.8821 -30.4683 -10.3758 -7.7837 -11.2050 -1.3014 -4.3010 9.03 -24.0194 -64.8539 -10.2954 -29.0377 -10.4394 -8.3242 -12.2224 -1.3014 -4.3010 8.51 -23.3609 -62.8786 -10.7459 -27.4782 -10.5087 -8.9133 -13.3313 -1.3014 -4.3010 8.06 -22.7889 -61.1625 -11.1373 -26.1234 -10.5690 -9.4251 -14.2948 -1.3014 -4.3010 7.47 -22.0437 -58.9270 -11.6472 -24.3585 -10.6474 -10.0919 -15.5498 -1.3014 -4.3010 6.90 -21.3267 -56.7759 -12.1378 -22.6603 -10.7229 -10.7334 -16.7574 -1.3014 -4.3010 6.49 -20.8035 -55.2062 -12.4957 -21.4211 -10.7779 -11.2016 -17.6386 -1.3014 -4.3010 6.01 -20.1949 -53.3805 -12.9121 -19.9797 -10.8420 -11.7461 -18.6636 -1.3014 -4.3010 5.51 -19.5598 -51.4753 -13.3467 -18.4756 -10.9089 -12.3143 -19.7332 -1.3013 -4.3010 5.02 -18.9384 -49.6111 -13.7718 -17.0040 -10.9743 -12.8703 -20.7797 -1.3010 -4.3010 4.51 -18.2923 -47.6727 -14.2134 -15.4745 -11.0424 -13.4483 -21.8675 -1.2983 -4.3010 3.94 -17.5670 -45.4966 -14.7057 -13.7644 -11.1202 -14.0960 -23.0852 -1.2722 -4.3010 3.64 -17.1834 -44.3459 -14.9589 -12.8739 -11.1638 -14.4363 -23.7222 -1.2137 -4.3010 3.20 -16.6754 -42.8214 -14.9892 -12.3049 -10.8228 -14.5351 -24.2609 -1.2041 -4.3010 2.76 -16.1703 -41.3052 -14.9894 -11.7984 -10.4421 -14.5979 -24.7671 -1.2101 -4.3010 2.54 -15.9230 -40.5623 -14.9895 -11.5498 -10.2546 -14.6284 -25.0156 -1.2160 -4.3010 2.05 -15.3742 -38.9106 -14.9903 -10.9942 -9.8334 -14.6952 -25.5704 -1.2438 -4.3010 1.51 -14.7842 -37.1220 -14.9921 -10.3818 -9.3621 -14.7648 -26.1809 -1.3148 -4.3010 1.32 -14.5857 -36.5123 -14.9929 -10.1675 -9.1950 -14.7880 -26,3945 -1,3509 -4.3010 306

Appendix 9.10 The equilibrium concentration profiles of the predominant An species as a function of pH; [FeS2] = 33.40 mMy [AuCl4'j = 50 pM, [NaClJ = 1.709 M and [OCT] = 0.019 M (refer to Figure 9.8a and b).

pH log[AuCI2~] logfAuCU] log[AuS] log[Fe3*] log[H2S(aq)] log[HS~] log[S2] log[S042'] logfAu] 13.30 -29.6910 -81.7850 -6.5933 -42.0343 -9.6864 -3.3021 -2.9312 -2.5626 -4.3033 13.02 -29.2074 -80.3345 -6.8009 -41.1328 -9.8233 -3.7165 -3.6230 -1.5989 -4.3024 12.49 -28.5026 -78.2199 -7.2715 -39.4844 -9.9396 -4.3626 -4.7991 -1.3699 -4.3015 11.99 -27.8585 -76.2875 -7.7247 -37.9333 -10.0348 -4.9590 -5.8966 -1.3218 -4.3012 11.45 -27.1694 -74.2203 -8.1982 -36.2970 -10.1135 -5.5796 -7.0592 -1.3071 -4.3011 11.04 -26.6572 -72.6837 -8.5481 -35.0850 -10.1675 -6.0377 -7.9213 -1.3036 -4.3011 10.41 -25.8539 -70.2738 -9.0975 -33.1830 -10.2522 -6.7563 -9.2739 -1.3019 -4.3010 9.65 -24.8905 -67.3836 -9.7566 -30.9014 -10.3536 -7.6183 -10.8964 -1.3014 -4.3010 8.05 -22.8703 -61.3232 -11.1388 -26.1168 -10.5663 -9.4258 -14.2988 -1.3014 -4.3010 7.46 -22.1216 -59.0771 -11.6510 -24.3436 -10.6451 -10.0957 -15.5598 -1.3014 -4.3010 6.89 -21.4020 -56.9182 -12.1434 -22.6392 -10.7208 -10.7396 -16.7718 -1.3014 -4.3010 6.48 -20.8796 -55.3510 -12.5008 -21.4020 -10.7758 -11.2070 -17.6516 -1.3014 -4.3010 6.00 -20.2730 -53.5311 -12.9159 -19.9652 -10.8397 -11.7498 -18.6733 -1.3014 -4.3010 5.58 -19.7430 -51.9410 -13.2785 -18.7099 -10.8955 -12.2240 -19.5660 -1.3013 -4.3010 4.91 -18.8914 -49.3862 -13.8611 -16.6931 -10.9851 -12.9859 -21.0001 -1.3008 -4.3010 4.51 -18.3804 -47.8533 -14.2104 -15.4836 -11.0391 -13.4430 -21.8604 -1.2984 -4.3010 4.03 -17.7923 -46.0888 -14.4920 -14.3321 -10.9519 -13.8344 -22.7302 -1.2905 -4.3010 3.65 -17.3620 -44.7980 -14.4920 -13.9017 -10.6291 -13.8881 -23.1606 -1.2911 -4.3010 3.13 -16.7657 -43.0085 -14.4921 -13.3048 -10.1809 -13.9625 -23.7575 -1.2935 -4.3010 2.47 -16.0147 -40.7534 -14.4924 -12.5510 -9.6135 -14.0555 -24.5110 -1.3055 -4.3010 2.05 -15.5404 -39.3255 -14.4929 -12.0708 -9.2494 -14.1133 -24.9907 -1.3276 -4.3010 1.51 -14.9561 -37.5542 -14.4942 -11.4655 -8.7847 -14.1830 -25.5947 -1.3872 -4.3010 1.34 -14.7718 -36.9885 -14.4947 -11.2673 -8.6311 -14.2050 -25.7923 -1.4150 -4,3010 307

Appendix 9.11 The equilibrium concentration profiles of the predominant Au species as a function of pH; [CuFeS2] = 21.80 mM, [AuCU] = 50 pM, [NaCl] = 1.709 M and [OCT] = 0.194 M (refer to Figure 9.9a and b).

PH log[AuCI2] log[AuCI4'] log[CuCI+] log[Fe2+] log[S042-] log[OCI ] log[CI2] log[HCIO] log[Au] 13.08 -11.8157 -28.2429 -17.7582 -35.3332 -2.4424 -17.4853 -32.3785 -22.9241 -4.3010 13.00 -11.7400 -28.0159 -17.6066 -35.1814 -2.3592 -17.4853 -32.2270 -22.8482 -4.3010 12.57 -11.3034 -26.7060 -16.7324 -34.3067 -2.1574 -17.4853 -31.3533 -22.4109 -4.3010 12.05 -10.7910 -25.1686 -15.7071 -33.2813 -2.0962 -17.4853 -30.3282 -21.8981 -4.3010 11.69 -10.4310 -24.0889 -14.9872 -32.5613 -2.0821 -17.4853 -29.6083 -21.5382 -4.3010 11.00 -9.7413 -22.0197 -13.6076 -31.1817 -2.0735 -17.4853 -28.2288 -20.8484 -4.3010 10.52 -9.2549 -20.5606 -12.6349 -30.2090 -2.0721 -17.4853 -27.2561 -20.3620 -4.3010 10.07 -8.8050 -19.2107 -11.7350 -29.3090 -2.0716 -17.4853 -26.3561 -19.9120 -4.3010 9.50 -8.2408 -17.5181 -10.6066 -28.1807 -2.0714 -17.4853 -25.2278 -19.3479 -4.3011 9.03 -7.7645 -16.0893 -9.6540 -27.2281 -2.0714 -17.4853 -24.2752 -18.8716 -4.3012 7.09 -5.8313 -10.2896 -5.7875 -23.3616 -2.0714 -17.4853 -20.4087 -16.9384 -4.3140 6.51 -5.2467 -8.5357 -4.6182 -22.1922 -2.0711 -17.4854 -19.2395 -16.3538 -4.3534 5.99 -4.7310 -6.9885 -3.5863 -21.1600 -2.0687 -17.4857 -18.2079 -15.8380 -4.5070 5.49 -4.3669 -5.6291 -2.5889 -20.1607 -2.0461 -17.4879 -17.2126 -15.3406 5.24 -4.4832 -5.2408 -2.0798 -19.6474 -1.9945 -17.4920 -16.7079 -15.0883 4.63 -5.2770 -4.8240 -2.0794 -18.4358 -1.9943 -17.4927 -15.4972 -14.4833 4.21 -6.0878 -4.7826 -2.0795 -17.5834 -1.9941 -17.4928 -14.6451 -14.0573 3.99 -6.5223 -4.7781 -2.0795 -17.1444 -1.9940 -17.4929 -14.2061 -13.8378 3.51 -7.4725 -4.7758 -2.0796 -16.1917 -1.9932 -17.4930 -13.2536 -13.3616 3.03 -8.4251 -4.7755 -2.0799 -15.2383 -1.9911 -17.4932 -12.3008 -12.8852 2.53 -9.4369 -4.7756 -2.0809 -14.2248 -1.9841 -17.4941 -11.2890 -12.3793 2.04 -10.3980 -4.7758 -2.0837 -13.2588 -1.9648 -17.4967 -10.3280 -11.8989 1.51 -11.4504 -4.7768 -2.0929 -12.1890 -1.9085 -17.5058 -9.2766 -11.3733 1.18 -12.0900 -4.7793 -2.1049 -11.5226 -1.8462 -17.5205 -8.6396 -11.0551 308

Appendix 9.12 The equilibrium concentration profiles of the predominant Au species as a function of pH; [CuFeS2] = 21.80 mM, [AuCli] = 50 pM, [Nad] = 7.709 M and [OCl'J = 0.019 M (refer to Figure 9.10a and b).

PH log[AuCI2‘] log[AuCI<] log[AuS'] log[Cu2S] log[Fe2t] log[H2S(aq)] log[HS] i°g[s2] log[SO<2] log[Au] 13.31 -29.1514 -80.1662 -11.2310 -2.7766 -18.7617 -14.8811 -8.4881 -8.1086 -3.3538 -4.3010 13.03 -28.6774 -78.7443 -11.3155 -2.7888 -18.3960 -14.8816 -8.7679 -8.6677 -2.3474 -4.3010 12.56 -28.0639 -76.9040 -11.6380 -2.7888 -17.6039 -14.8806 -9.2357 -9.6041 -2.1201 -4.3010 12.03 -27.3953 -74.8981 -12.0269 -2.7888 -16.6857 -14.8802 -9.7643 -10.6618 -2.0588 -4.3010 11.64 -26.9035 -73.4227 -12.3191 -2.7888 -16.0013 -14.8802 -10.1563 -11.4459 -2.0447 -4.3010 9.88 -24.7128 -66.8504 -13.6317 -2.7888 -12.9369 -14.8801 -11.9080 -14.9493 -2.0355 -4.3010 9.54 -24.2772 -65.5437 -13.8930 -2.7888 -12.3272 -14.8801 -12.2564 -15.6462 -2.0354 -4.3010 9.12 -23.7586 -63.9879 -14.2042 -2.7888 -11.6011 -14.8801 -12.6713 -16.4760 -2.0353 -4.3010 8.42 -22.8883 -61.3771 -14.7264 -2.7888 -10.3827 -14.8801 -13.3675 -17.8684 -2.0353 -4.3010 7.63 -21.8972 -58.4038 -15.3210 -2.7888 -8.9952 -14.8801 -14.1604 -19.4541 -2.0353 -4.3010 7.06 -21.1786 -56.2478 -15.7522 -2.7888 -7.9891 -14.8801 -14.7353 -20.6040 -2.0353 -4.3010 6.51 -20.4973 -54.2040 -16.1610 -2.7889 -7.0353 -14.8801 -15.2804 -21.6941 -2.0353 -4.3010 6.00 -19.8588 -52.2886 -16.5441 -2.7890 -6.1414 -14.8801 -15.7911 -22.7156 -2.0353 -4.3010 5.60 -19.3594 -50.7904 -16.8437 -2.7891 -5.4423 -14.8801 -16.1906 -23.5146 -2.0353 -4.3010 5.02 -18.6372 -48.6238 -17.2770 -2.7898 -4.4312 -14.8801 -16.7684 -24.6702 -2.0351 -4.3010 4.58 -18.0827 -46.9603 -17.6095 -2.7901 -3.6552 -14.8801 -17.2119 -25.5572 -2.0341 -4.3010 4.22 -17.6367 -45.6222 -17.8764 -2.7869 -3.0317 -14.8801 -17.5684 -26.2702 -2.0307 -4.3010 3.35 -16.6314 -42.6056 -17.9215 -2.7787 -2.4612 -14.1852 -17.7464 -27.3211 -2.0177 -4.3010 2.69 -15.9075 -40.4325 -17.6761 -2.8000 -2.4666 -13.3360 -17.5618 -27.8012 -2.0114 -4.3010 2.07 -15.2417 -38.4294 -17.4467 -2.8341 -2.4749 -12.5514 -17.3906 -28.2433 -1.9886 -4.3010 1.50 -14.6302 -36.5748 -17.2285 -2.8830 -2.4864 -11.8214 -17.2323 -28.6567 -1.9245 -4.3010 1.21 -14.3356 -35.6663 -17.1179 -2.9102 -2.4927 -11.4579 -17.1549 -28.8653 -1.8699 -4.3010 309

Appendix 9.13 CHEMIX programme input statement for the system Au- Cl -Na0CI-H20 at 25°C: [Au] = ImM, [NaCI] = 1.709M and IOCI'1 = 0.194M 310

TITLE CHEMIX CALCULATION OF GOLD GIBBS SOLUBILITY CL2 (G) IN THE Au-NaCl-Na0Cl-H20 SYSTEM AT 0.0 NaCl 100g/l, OC1- 10g/l H CL (G) TO STUDY THE EFFECT OF pH. -95299. CL E (AQ) SYSTEM -131228. GAS CL2 (AQ) 02 (G) 6940. N2 (G) CL3 E (AQ) H2 (G) -120400. C12 (g) CL 02 E (AQ) H Cl (g) 17200. H2 O (g) H CL O (AQ) -79900. AQUEOUS CL O E (AQ) H20 -36800. O H E (AQ) CL 02 (AQ) H E-l (AQ) 120100. AU CL2 E (AQ) CL 03 E (AQ) AU CL4 E (AQ) -7950. AU 03 E3 (AQ) CL 04 E (AQ) H AU 03 E2 (AQ) -8520. H2 AU 03 E (AQ) AU AU [O H]3 (AQ) 0.0 CL E (AQ) AU 03 H3 CL2 (AQ) -316920. CL3 E (AQ) AU 03 E3 (AQ) H CL O (aq) -51800. Na CL O (ai) H AU 03 E2 (AQ) Cl 02 <-> (aq) -142200. Cl O <-> (aq) H2 AU 03 E (AQ) Cl 02 (aQ) -218300. CL 03 <-> (AQ) AU 03 H3 (AQ) Cl 04 <-> (aq) -283370. Na < + > (aq) AU CL2 E (AQ) H Cl (ai) -151120. Au C14 H (ai) AU CL4 E (AQ) Au C14 Na (ai) -235140. S 04 <2-> (aq) H2S04 <0+> (aq) H S 04 <-> (aq) -744495. H2 S 04 <0+> (aq) FILE SINGLE CPDNBSDAT Au CPDNPLDAT Au [O H]3 CPDMRLDAT Au Cl Au C13 INSPECIES Na Cl GAS N2 (G)

UNITS AQUEOUS JOULE H20 311

Na Cl O (ai) H < + > (AQ) H2 S 04 <0+ > (aq) Na < + > (aq) VARY OH <-> (aq) INMOLE STEPSIZE .00005 .01 .335 AQUEOUS SINGLE H2 s o4 < 0 + > (AQ) Au Na Cl equilibrate

step TEMPERATURE outmole 298.15 single Au PRESSURE 1. VARY INMOLE STEPSIZE .00005 .01 .335 INMOLE AQUEOUS .8 H2 s o4 <0 + > (AQ) 55.52 .1944 equilibrate .00005 .04 STEP .04 OUTMOLE .001 AQUEOUS 1.709 Au CL2 E (AQ) Au CL4 E (AQ) try CL O E (AQ) .2E-1 .869 .19E-36 .15E-24 .26E-16 .25E-1 CL2 (AQ) 55.5 .54E-1 .24E-10 .12E-10 .40E-24 H CL O (AQ) .22E-15 .26E-14 .19E-10 CL 03 E (AQ) . 12E-12 2.22.15E-25 .68E-45 .62E-22 . 17E-15 .48E-22 .76E-15 VARY .24E-32 .34E-20 .97E-22 2.32.16E-10 INMOLE STEPSIZE .01505 . 00024 . 02505 . 16E-32 .92E-22 .22E-2 AQUEOUS . IE-12 .95E-25 0.0 0.0 0.0 0.0 H2so4 <0+> (AQ)

EQUILIBRATE EQUILIBRATE

STEP step OUTMOLE activity AQUEOUS aqueous Au CL2 E (AQ) H < + > (AQ) Au CL4 E (AQ) CL O E (AQ) VARY CL2 (AQ) INMOLE STEPSIZE .01505 .00024 .02505 H CL O (AQ) AQUEOUS CL 03 E (AQ) H2so4 <0+ > (AQ)

VARY equilibrate INMOLE STEPSIZE .00005 .01 .335 AQUEOUS step H2so4 <0+> (AQ) outmole single EQUILIBRATE Au step VARY activity INMOLE STEPSIZE .01505 .00024 .02505 aqueous AQUEOUS 312

H2 s o4 <0 + > (AQ) AQUEOUS H2 s o4 <0+ > (AQ) equilibrate EQUILIBRATE STEP OUTMOLE step AQUEOUS activity Au CL2 E (AQ) aqueous Au CL4 E (AQ) H < + > (AQ) CL O E (AQ) CL2 (AQ) VARY H CL O (AQ) INMOLE STEPSIZE .019993 .0000016 CL 03 E (AQ) .020005 AQUEOUS VARY H2 s o4 < 0 + > (AQ) INMOLE STEPSIZE .05005 .00072 .0554 AQUEOUS equilibrate H2 s o4 <0 + > (AQ) step EQUILIBRATE outmole single step Au activity aqueous VARY H < + > (AQ) INMOLE STEPSIZE .019993 .0000016 .020005 VARY AQUEOUS INMOLE STEPSIZE .05005 .00072 .0554 H2so4 <0+> (AQ) AQUEOUS H2 s o4 <0 + > (AQ) equilibrate equilibrate STEP OUTMOLE step AQUEOUS outmole Au CL2 E (AQ) single Au CL4 E (AQ) Au CL O E (AQ) CL2 (AQ) VARY H CL O (AQ) INMOLE STEPSIZE .05005 .00072 .0554 CL 03 E (AQ) AQUEOUS H2 s o4 <0 + > (AQ) VARY INMOLE STEPSIZE .0199 .00005 .0201 equilibrate AQUEOUS H2 s o4 <0 + > (AQ) STEP OUTMOLE EQUILIBRATE AQUEOUS Au CL2 E (AQ) step Au CL4 E (AQ) activity CL O E (AQ) aqueous CL2 (AQ) H < + > (AQ) H CL O (AQ) CL 03 E (AQ) VARY INMOLE STEPSIZE .0199 .00005 .0201 VARY AQUEOUS INMOLE STEPSIZE .019993 .0000016 H2so4 <0+ > (AQ) .020005 313

equilibrate VARY INMOLE STEPSIZE .0199 .00005 .0201 step AQUEOUS outmole H2 s o4 <0 + > (AQ) single Au equilibrate 314

Appendix 9.14 CHEMIX programme input statement for the system Au- C1 -Na0Cl-H20 at 25°C: [Au] = ImM, [NaCl] = 0.855M and [OCT] = 0.019M 315

TITLE CHEMIX CALCULATION OF GOLD GIBBS SOLUBILITY CL2(G) IN THE Au-NaCl-NaOCl-H20 SYSTEM AT 0.0 NaCl 50g/l, OC1- lg/1 H CL (G) TO STUDY THE EFFECT OF pH. -95299. CL E (AQ) SYSTEM -131228. GAS CL2 (AQ) 02 (G) 6940. N2 (G) CL3 E (AQ) H2 (G) -120400. C12 (g) CL 02 E (AQ) H Cl (g) 17200. H2 O (g) H CL O (AQ) -79900. AQUEOUS CL O E (AQ) H20 -36800. O H E (AQ) CL 02 (AQ) H E-l (AQ) 120100. AU CL2 E (AQ) CL 03 E (AQ) AU CL4 E (AQ) -7950. AU 03 E3 (AQ) CL 04 E (AQ) H AU 03 E2 (AQ) -8520. H2 AU 03 E (AQ) AU AU [O H]3 (AQ) 0.0 CL E (AQ) AU 03 H3 CL2 (AQ) -316920. CL3 E (AQ) AU 03 E3 (AQ) H CL O (aq) -51800. Na CL O (ai) H AU 03 E2 (AQ) Cl 02 <-> (aq) -142200. Cl O < - > (aq) H2 AU 03 E (AQ) Cl 02 (aQ) -218300. CL 03 <-> (AQ) AU 03 H3 (AQ) Cl 04 <-> (aq) -283370. Na < + > (aq) AU CL2 E (AQ) H Cl (ai) -151120. Au C14 H (ai) AU CL4 E (AQ) Au C14 Na (ai) -235140. S 04 <2-> (aq) H2 S 04 <0+> (aq) H S 04 <-> (aq) -744495. H2S 04 <0+ > (aq) FILE SINGLE CPDNBSDAT Au CPDNPLDAT Au [O H]3 CPDMRLDAT Au Cl Au C13 INSPECIES Na Cl GAS N2 (G)

UNITS AQUEOUS JOULE H20 316

Na Cl O (ai) H < + > (AQ) H2 S 04 <0 + > (aq) Na < + > (aq) VARY O H <-> (aq) INMOLE STEPSIZE .00005 .01 .335 AQUEOUS SINGLE H2 s o4 <0+ > (AQ) Au Na Cl equilibrate

step TEMPERATURE outmole 298.15 single Au PRESSURE 1. VARY INMOLE STEPSIZE .00005 .01 .335 INMOLE AQUEOUS .8 H2 s o4 <0+ > (AQ) 55.52 .0194 equilibrate .00005 .04 STEP .04 OUTMOLE .001 AQUEOUS 0.8556 Au CL2 E (AQ) Au CL4 E (AQ) try CL O E (AQ) .2E-1 .869 .19E-36 .15E--24 .26E-16 .25E-1 CL2 (AQ) 55.5 .54E-1 .24E-10.12E-10 .40E-24 H CL O (AQ) .22E-15 .26E-14 .19E-10 CL 03 E (AQ) . 12E-12 2.22 .15E-25 .68E-45 .62E-22 . 17E-15 .48E-22 .76E-15 VARY .24E-32 .34E-20 .97E-22 2.32 .16E-10 INMOLE STEPSIZE .01505 .00024 .02505 . 16E-32 .92E-22 .22E-2 AQUEOUS . IE-12 .95E-25 0.0 0.0 0.0 0.0 H2 s o4 <0+ > (AQ)

EQUILIBRATE EQUILIBRATE

STEP step OUTMOLE activity AQUEOUS aqueous Au CL2 E (AQ) H < + > (AQ) Au CL4 E (AQ) CL O E (AQ) VARY CL2 (AQ) INMOLE STEPSIZE .01505 .00024 .02505 H CL O (AQ) AQUEOUS CL 03 E (AQ) H2 s o4 < 0 + > (AQ)

VARY equilibrate INMOLE STEPSIZE .00005 .01 .335 AQUEOUS step H2so4 <0+> (AQ) outmole single EQUILIBRATE Au step VARY activity INMOLE STEPSIZE .01505 .00024 .02505 aqueous AQUEOUS 317

H2 s o4 <0 + > (AQ) AQUEOUS H2so4 <0+ > (AQ) equilibrate EQUILIBRATE STEP OUTMOLE step AQUEOUS activity Au CL2 E (AQ) aqueous Au CL4 E (AQ) H < + > (AQ) CL O E (AQ) CL2 (AQ) VARY H CL O (AQ) INMOLE STEPSIZE .019993 .0000016 CL 03 E (AQ) .020005 AQUEOUS VARY H2 s o4 <0 + > (AQ) INMOLE STEPSIZE .05005 .00072 .0554 AQUEOUS equilibrate H2 s o4 <0 + > (AQ) > step EQUILIBRATE outmole single step Au activity aqueous VARY H < + > (AQ) INMOLE STEPSIZE .019993 .0000016 .020005 VARY AQUEOUS INMOLE STEPSIZE .05005 .00072 .0554 H2 s o4 <0 + > (AQ) AQUEOUS H2 s o4 <0 + > (AQ) equilibrate equilibrate STEP OUTMOLE step AQUEOUS outmole Au CL2 E (AQ) single Au CL4 E (AQ) Au CL O E (AQ) CL2 (AQ) VARY H CL O (AQ) INMOLE STEPSIZE .05005 .00072 .0554 CL 03 E (AQ) AQUEOUS H2so4 <0+ > (AQ) VARY INMOLE STEPSIZE .0199 .00005 .0201 equilibrate AQUEOUS H2so4 <0+> (AQ) STEP OUTMOLE EQUILIBRATE AQUEOUS Au CL2 E (AQ) step Au CL4 E (AQ) activity CL O E (AQ) aqueous CL2 (AQ) H < + > (AQ) H CL O (AQ) CL 03 E (AQ) VARY IN MOLE STEPSIZE .0199 .00005 .0201 VARY AQUEOUS INMOLE STEPSIZE .019993 .0000016 H2so4 <0+ > (AQ) .020005 318 equilibrate INMOLE STEPSIZE .0199 .00005 .0201 AQUEOUS step H2so4 <0+ > (AQ) outmole single equilibrate Au

VARY 319

Appendix 9.15 CHEMIX programme input statement for the system Au- CI -Na0Cl-H20 at 25°C: [Au] = ImM, [NaCI] = 1.709M and [OCX] = 0.019M 320

TITLE GIBBS CHEMIX CALCULATION OF GOLD CL2 (G) SOLUBILITY 0.0 IN THE Au-NaCl-Na0Cl-H20 SYSTEM AT H CL (G) NaCl 100g/l, OC1- lg/1 -95299. TO STUDY THE EFFECT OF pH. CL E (AQ) -131228. SYSTEM CL2 (AQ) GAS 6940. 02 (G) CL3 E (AQ) N2 (G) -120400. H2 (G) CL 02 E (AQ) C12 (g) 17200. H Cl (g) H CL O (AQ) H2 O (g) -79900. CL O E (AQ) AQUEOUS -36800. H20 CL 02 (AQ) O H E (AQ) 120100. H E-l (AQ) CL 03 E (AQ) AU CL2 E (AQ) -7950. AU CL4 E (AQ) CL 04 E (AQ) AU 03 E3 (AQ) -8520. H AU 03 E2 (AQ) AU H2 AU 03 E (AQ) 0.0 AU [O H]3 (AQ) AU 03 H3 CL E (AQ) -316920. CL2 (AQ) AU 03 E3 (AQ) CL3 E (AQ) -51800. H CL O (aq) H AU 03 E2 (AQ) Na CL O (ai) -142200. Cl 02 < - > (aq) H2 AU 03 E (AQ) Cl O < - > (aq) -218300. Cl 02 (aQ) AU 03 H3 (AQ) CL 03 <-> (AQ) -283370. Cl 04 <-> (aq) AU CL2 E (AQ) Na < + > (aq) -151120. H Cl (ai) AU CL4 E (AQ) Au C14 H (ai) -235140. Au C14 Na (ai) H2 S 04 <0+ > (aq) S 04 <2-> (aq) -744495. H S 04 <-> (aq) H2S 04 <0+> (aq) FILE CPDNBSDAT SINGLE CPDNPLDAT Au CPDMRLDAT Au [O H]3 Au Cl INSPECIES Au C13 GAS Na Cl N2 (G)

AQUEOUS UNITS H20 JOULE Na Cl O (ai) H2 S 04 <0+ > (aq) 321

Na < + > (aq) VARY O H <-> (aq) INMOLE STEPSIZE .00005 .01 .335 AQUEOUS SINGLE H2 s o4 <0+ > (AQ) Au Na Cl equilibrate

step TEMPERATURE outmole 298.15 single Au PRESSURE 1. VARY INMOLE STEPSIZE .00005 .01 .335 INMOLE AQUEOUS .8 H2 s o4 <0 + > (AQ) 55.52 .0194 equilibrate .00005 .04 STEP .04 OUTMOLE .001 AQUEOUS 1.709 Au CL2 E (AQ) Au CL4 E (AQ) try CL O E (AQ) .2E-1 .869.19E-36 .15E-24 .26E-16 .25E-1 CL2 (AQ) 55.5 .54E-1 .24E-10 .12E-10 .40E-24 .22E- H CL O (AQ) 15 .26E-14.19E-10 CL 03 E (AQ) . 12E-12 2.22.15E-25 .68E-45 .62E-22 . 17E-15 .48E-22 .76E-15 VARY .24E-32 .34E-20 .97E-22 2.32.16E-10 INMOLE STEPSIZE .01505 .00024 .02505 . 16E-32 .92E-22 .22E-2 AQUEOUS .IE-12 .95E-25 0.0 0.0 0.0 0.0 H2so4 <0+> (AQ)

EQUILIBRATE EQUILIBRATE

STEP step OUTMOLE activity AQUEOUS aqueous Au CL2 E (AQ) H < + > (AQ) Au CL4 E (AQ) CL O E (AQ) VARY CL2 (AQ) INMOLE STEPSIZE .01505 .00024 .02505 H CL O (AQ) AQUEOUS CL 03 E (AQ) H2 s o4 <0 + > (AQ)

VARY equilibrate INMOLE STEPSIZE .00005 .01 .335 AQUEOUS step H2so4 <0+> (AQ) outmole single EQUILIBRATE Au step VARY activity IN MOLE STEPSIZE .01505 .00024 .02505 aqueous AQUEOUS H < + > (AQ) H2so4 <0+ > (AQ) 322 equilibrate EQUILIBRATE STEP OUTMOLE step AQUEOUS activity Au CL2 E (AQ) aqueous Au CL4 E (AQ) H < + > (AQ) CL O E (AQ) CL2 (AQ) VARY H CL O (AQ) INMOLE STEPSIZE .019993 .0000016 CL 03 E (AQ) .020005 AQUEOUS VARY H2so4 <0+ > (AQ) INMOLE STEPSIZE .05005 .00072 .0554 AQUEOUS equilibrate H2 so4 <0+ > (AQ) step EQUILIBRATE outmole single step Au activity aqueous VARY H < + > (AQ) INMOLE STEPSIZE .019993 .0000016 .020005 VARY AQUEOUS INMOLE STEPSIZE .05005 .00072 .0554 H2 s o4 <0 + > (AQ) AQUEOUS H2 s o4 <0+ > (AQ) equilibrate equilibrate STEP OUTMOLE step AQUEOUS outmole Au CL2 E (AQ) single Au CL4 E (AQ) Au CL O E (AQ) CL2 (AQ) VARY H CL O (AQ) INMOLE STEPSIZE .05005 .00072 .0554 CL 03 E (AQ) AQUEOUS H2so4 <0+ > (AQ) VARY INMOLE STEPSIZE .0199 .00005 .0201 equilibrate AQUEOUS H2so4 <0+ > (AQ) STEP OUTMOLE EQUILIBRATE AQUEOUS Au CL2 E (AQ) step Au CL4 E (AQ) activity CL O E (AQ) aqueous CL2 (AQ) H < + > (AQ) H CL O (AQ) CL 03 E (AQ) VARY INMOLE STEPSIZE .0199 .00005 .0201 VARY AQUEOUS INMOLE STEPSIZE .019993 .0000016 H2so4 <0+ > (AQ) .020005 AQUEOUS equilibrate H2so4 <0+> (AQ) 323

step INMOLE STEPSIZE .0199 .00005 .0201 outmole AQUEOUS single H2 s o4 <0+ > (AQ) Au equilibrate VARY 324

Appendix 9.16 CHEMIX programme input statement for the system Au- CI'-Na0Cl-H20 at 25°C: [Au] =lmM, [NaCI] = 3.422M and [OCT] = 0.019M 325

TITLE GIBBS CHEMIX CALCULATION OF GOLD CL2 (G) SOLUBILITY 0.0 IN THE Au-NaCl-Na0CI-H20 SYSTEM AT H CL (G) NaCI 200g/l, OC1- lg/1 -95299. TO STUDY THE EFFECT OF pH. CL E (AQ) -131228. SYSTEM CL2 (AQ) GAS 6940. 02 (G) CL3 E (AQ) N2 (G) -120400. H2 (G) CL 02 E (AQ) C12 (g) 17200. H Cl (g) H CL O (AQ) H2 O (g) -79900. CL O E (AQ) AQUEOUS -36800. H20 CL 02 (AQ) O H E (AQ) 120100. H E-l (AQ) CL 03 E (AQ) AU CL2 E (AQ) -7950. AU CL4 E (AQ) CL 04 E (AQ) AU 03 E3 (AQ) -8520. H AU 03 E2 (AQ) AU H2 AU 03 E (AQ) 0.0 AU [O H]3 (AQ) AU 03 H3 CL E (AQ) -316920. CL2 (AQ) AU 03 E3 (AQ) CL3 E (AQ) -51800. H CL O (aq) H AU 03 E2 (AQ) Na CL O (ai) -142200. Cl 02 <-> (aq) H2 AU 03 E (AQ) Cl O <-> (aq) -218300. Cl 02 (aQ) AU 03 H3 (AQ) CL 03 <-> (AQ) -283370. Cl 04 <-> (aq) AU CL2 E (AQ) Na < + > (aq) -151120. H Cl (ai) AU CL4 E (AQ) Au C14 H (ai) -235140. Au C14 Na (ai) H2 S 04 <0+> (aq) S 04 <2-> (aq) -744495. H S 04 <-> (aq) H2 S 04 <0+ > (aq) FILE CPDNBSDAT SINGLE CPDNPLDAT Au CPDMRLDAT Au [O H]3 Au Cl INSPECIES Au C13 GAS Na Cl N2 (G)

AQUEOUS UNITS H20 JOULE Na Cl O (ai) H2 S 04 <0+> (aq) 326

Na < + > (aq) VARY O H <-> (aq) INMOLE STEPSIZE .00005 .01 .335 AQUEOUS SINGLE H2 s o4 <0 + > (AQ) Au Na Cl equilibrate

step TEMPERATURE outmole 298.15 single Au PRESSURE 1. VARY INMOLE STEPSIZE .00005 .01 .335 INMOLE AQUEOUS .8 H2so4 <0+ > (AQ) 55.52 .0194 equilibrate .00005 .04 STEP .04 OUTMOLE .001 AQUEOUS 3.4223 Au CL2 E (AQ) Au CL4 E (AQ) try CL O E (AQ) .2E-1 .869 .19E-36 .15E-24 .26E-16 .25E-1 CL2 (AQ) 55.5 .54E-1 .24E-10 .12E-10 .40E-24 .22E- H CL O (AQ) 15 .26E-14 .19E-10 CL 03 E (AQ) . 12E-12 2.22 .15E-25 .68E-45 .62E-22 .17E-15 .48E-22 .76E-15 VARY .24E-32 .34E-20 .97E-22 2.32 .16E-10 INMOLE STEPSIZE .01505 .00024 .02505 . 16E-32 .92E-22 .22E-2 AQUEOUS . IE-12 .95E-25 0.0 0.0 0.0 0.0 H2 s o4 <0 + > (AQ)

EQUILIBRATE EQUILIBRATE

STEP step OUTMOLE activity AQUEOUS aqueous Au CL2 E (AQ) H < + > (AQ) Au CL4 E (AQ) CL O E (AQ) VARY CL2 (AQ) INMOLE STEPSIZE .01505 .00024 .02505 H CL O (AQ) AQUEOUS CL 03 E (AQ) H2 so4 <0+> (AQ)

VARY equilibrate INMOLE STEPSIZE .00005 .01 .335 AQUEOUS step H2 s o4 < 0 + > (AQ) outmole single EQUILIBRATE Au step VARY activity INMOLE STEPSIZE .01505 .00024 .02505 aqueous AQUEOUS H < + > (AQ) H2so4 <0+ > (AQ) 327 equilibrate EQUILIBRATE STEP OUTMOLE step AQUEOUS activity Au CL2 E (AQ) aqueous Au CL4 E (AQ) H < + > (AQ) CL O E (AQ) CL2 (AQ) VARY H CL O (AQ) INMOLE STEPSIZE .019993 .0000016 CL 03 E (AQ) .020005 AQUEOUS VARY H2so4 <0+ > (AQ) INMOLE STEPSIZE .05005 .00072 .0554 AQUEOUS equilibrate H2 s o4 <0+> (AQ) step EQUILIBRATE outmole single step Au activity aqueous VARY H < + > (AQ) INMOLE STEPSIZE .019993 .0000016 .020005 VARY AQUEOUS INMOLE STEPSIZE .05005 .00072 .0554 H2 so4 <0+ > (AQ) AQUEOUS H2so4 <0+ > (AQ) equilibrate equilibrate STEP OUTMOLE step AQUEOUS outmole Au CL2 E (AQ) single Au CL4 E (AQ) Au CL O E (AQ) CL2 (AQ) VARY H CL O (AQ) INMOLE STEPSIZE .05005 .00072 .0554 CL 03 E (AQ) AQUEOUS H2so4 <0+> (AQ) VARY INMOLE STEPSIZE .0199 .00005 .0201 equilibrate AQUEOUS H2so4 <0+ > (AQ) STEP OUTMOLE EQUILIBRATE AQUEOUS Au CL2 E (AQ) step Au CL4 E (AQ) activity CL O E (AQ) aqueous CL2 (AQ) H < + > (AQ) H CL O (AQ) CL 03 E (AQ) VARY INMOLE STEPSIZE .0199 .00005 .0201 VARY AQUEOUS INMOLE STEPSIZE .019993 .0000016 H2 s o4 <0+> (AQ) .020005 AQUEOUS equilibrate H2 s o4 <0+ > (AQ) 328

step INMOLE STEPSIZE .0199 .00005 .0201 outmole AQUEOUS single H2 s o4 <0 + > (AQ) Au equilibrate VARY 329

Appendix 9.17 CHEMIX programme input statement for the system Au-

ZnS-CI -Na0Cl-H20 at 25°C: [ZnS] = 41.05mM, [AuCI4 ]

= 50|iM, [NaCl] = 1.709M and [OCT] = 0.194M 330

Title Au [O H]3 (aq) CHEM1X CALCULATION OF GOLD Au C14 H (ai) SOLUBILITY Au C14 Na (ai) IN THE Au-ZnS-NaCl-Na0Cl-H20 Au S <-> (aq) SYSTEM Zn S 04 < + > (aq) TO STUDY THE EFFECT OF SOLUTION Zn S 04 (ai) pH vs CONCENTRATION S <2-> (aq) S2 <2-> (aq) SYSTEM S 02 (aq) GAS S 04 <2-> (aq) H2 O (g) S2 03 <2-> (aq) H Cl (g) S2 04 <2-> (aq) 02 (g) S2 08 <2-> (aq) N2 (g) S4 06 <2-> (aq) S2 (g) C12 (g) SINGLE H2 (g) Au H2 S 04 (g) Au [O H]3 Au2 03 AQUEOUS Au Cl H2 O Au C13 H < + > (aq) Au2 S O H < - > (aq) Zn Cl <-> (aq) Zn S C12 (aq) ZnS 04 C13 <-> (aq) S H Cl (ai) Na2 S H Cl O (aq) Na Cl Na Cl O (ai) NaOH Cl 02 <-> (aq) Cl O <-> (aq) Cl 02 (aq) units Cl 03 <-> (aq) joules Cl 04 <-> (aq) S 04 <2-> (aq) Gibbs H S <-> (aq) Zn 02 < 2- > (aq) H2 S (aq) -384240. H S 04 <-> (aq) H Zn 02 < - > (aq) H2 S 04 (ai) -457080. Na < + > (aq) Zn O H < + > (aq) Na2 S (ai) -330100. Zn < 2 + > (aq) Zn 03 H3 <-> (aq) Zn C12 (aq) -694220. Zn Cl < + > (aq) Zn S 04 (ai) Zn C13 < - > (aq) -891590. Zn C14 < 2- > (aq) Zn S 04 < + > (aq) Zn Cl 04 < + > (aq) -904900. Zn C12 08 (ai) Au Zn 02 < 2- > (aq) 0.0 H Zn 02 < - > (aq) Au 03 <3-> (aq) Zn O H < + > (aq) -51800. Zn 03 H3 <-> (aq) Au C12 < - > (aq) Zn 04 H4 <2-> (aq) -151120. Au C12 < - > (aq) Au C14 < - > (aq) Au C14 <-> (aq) -235140. Au 03 <3-> (aq) Au 03 H3 H Au 03 < 2- > (aq) -316920. H2 Au 03 <-> (aq) Au 03 H3 (aq) 331

-283370. 55.52 H Au 03 <2-> (aq) 0.1944 -142200. 0.0000000005 H2 Au 03 < - > (aq) 0.00005 -218300. 0.2 Au S < - > (aq) 0.2 46024. 0.04105 Au2 S 1.709 29037. C13 <-> (aq) TRY -120400. 0.023 .24E-16 .36E-2 .8 .62E-80 .28E-26 Zn C12 (aq) .42E-38 .25E-40 -403700. 55.5 .23E-10 .26E-1 3.617 .24E-26 .46E-28 Zn Cl < + > (aq) .22E-10 .68E-20 .42E-16 .19E-25 .62E-16 -275300. .45E-36 .27E-22 .24E-22 .26E-2 .42E-7 Zn C13 <-> (aq) .3IE-10 .26E-10 .48E-24 3.657 .34E-84 -540500. .28E-10 .82E-10 .44E-9 .35E-8 .41E-8 Zn C14 <2-> (aq) . 14E-32 .22E-58 .42E-2 .68E-2 .35E-7 .72E -666000. 2 Zn Cl 04 < -1- > (aq) .22E-2 .25E-10 .69E-22 .72E-16 .46E-15 -163100. .5E-12 .5E-14 .34E-34 .34E-22 .52E-5 .26E Zn S 1 .25E-11 .46E-87 .35E-88 .78E-58 .75E-2 -201290. .36E-78 .12E-92 .28E-52 .83E-94 .57E-4 Zn 04 H4 <2-> (aq) 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 -858520. EQUILIBRATE FILE CPDNBSDAT STEP CPDNPLDAT outmole CPDMRLDAT AQUEOUS CPDJANDAT Au C12 <-> (aq) CPDSGTDAT Au C14 <-> (aq) Au S <-> (aq) INSPECIES Zn C14 <2-> (aq) GAS H2 S (aq) N2 (g) H S <-> (aq) S <2-> (aq) AQUEOUS S 04 <2-> (aq) H2 0 Cl O <-> (aq) Na Cl 0 (ai) C12 (aq) H2 S 04 (ai) H Cl O (aq) Au C14 < - > (aq) Cl 03 <-> (aq) Na < + > (aq) 0 H <-> (aq) VARY Inmole stepsize .0000000005 .005 .250 SINGLE aqueous Zn S H2 S 04 (ai) Na Cl EQUILIBRATE

TEMPERATURE STEP 298.15 activity AQUEOUS PRESSURE H < + > (aq)

1. VARY INMOLE Inmole stepsize .0000000005 .005 .250 0.8 aqueous 332

H2 S 04 (ai) single Au EQUILIBRATE Au2 S Zn STEP Zn S outmole ZnS 04 single S Au Au2 S VARY Zn Inmole stepsize .12 .000065 .125 ZnS aqueous Zn S 04 H2 S 04 (ai) S EQUILIBRATE VARY Inmole stepsize .0000000005 .005 .250 aqueous STEP H2 S 04 (ai) outmole AQUEOUS EQUILIBRATE Au C12 < - > (aq) Au C14 < - > (aq) STEP Au S < - > (aq) outmole Zn C14 <2-> (aq) AQUEOUS H2 S (aq) Au C12 < - > (aq) H S <-> (aq) Au C14 <-> (aq) S <2-> (aq) Au S < - > (aq) S 04 <2-> (aq) Zn C14 <2-> (aq) Cl O <-> (aq) H2 S (aq) C12 (aq) H S <-> (aq) H Cl O (aq) S <2-> (aq) Cl 03 <-> (aq) S 04 <2-> (aq) Cl O <-> (aq) VARY C12 (aq) Inmole stepsize .12043 .0000005 .12046 H Cl O (aq) aqueous Cl 03 <-> (aq) H2 S 04 (ai)

VARY EQUILIBRATE Inmole stepsize .12 .000065 .125 aqueous STEP H2 S 04 (ai) activity AQUEOUS EQUILIBRATE H < + > (aq)

STEP VARY activity Inmole stepsize .12043 .0000005 .12046 AQUEOUS aqueous H < + > (aq) H2 S 04 (ai)

VARY EQUILIBRATE Inmole stepsize .12 .000065 .125 aqueous STEP H2 S 04 (ai) outmole single EQUILIBRATE Au Au2 S STEP Zn outmole ZnS 333

Zn S 04 S

VARY Irunole stepsize .12043 .0000005 .12046 aqueous H2 S 04 (ai)

EQUILIBRATE 334

Appendix 9.18 CHEMIX programme input statement for the system Au- ZnS-Cl -NaOCl-H20 at 25°C: [ZnS] = 41.05mM, [AuCI4 ] = 50|iM, [NaCl] = 1.709M and [OCX] = 0.019M 335

Title Au C14 <-> (aq) CHEMIX CALCULATION OF GOLD Au 03 <3-> (aq) SOLUBILITY H Au 03 <2-> (aq) IN THE Au-ZnS-NaCl-Na0Cl-H20 H2 Au 03 <-> (aq) SYSTEM Au [O H]3 (aq) TO STUDY THE EFFECT OF SOLUTION Au C14 H (ai) pH vs CONCENTRATION Au C14 Na (ai) Au S <-> (aq) SYSTEM Zn S 04 < + > (aq) GAS Zn S 04 (ai) H2 O (g) S <2-> (aq) H Cl (g) S2 <2-> (aq) 02 (g) S 02 (aq) N2 (g) S 04 <2-> (aq) S2 (g) S2 03 <2-> (aq) C12 (g) S2 04 <2-> (aq) H2 (g) S2 08 <2-> (aq) H2 S 04 (g) S4 06 <2-> (aq)

AQUEOUS SINGLE H2 0 Au H < + > (aq) Au [O H]3 O H <-> (aq) Au2 03 Cl <-> (aq) Au Cl C12 (aq) Au C13 C13 <-> (aq) Au2 S H Cl (ai) Zn H Cl O (aq) Zn S Na Cl O (ai) ZnS 04 Cl 02 <-> (aq) S Cl O <-> (aq) Na2 S Cl 02 (aq) Na Cl Cl 03 <-> (aq) NaOH Cl 04 <-> (aq) S 04 <2-> (aq) H S <-> (aq) units H2 S (aq) joules H S 04 <-> (aq) H2 S 04 (ai) Gibbs Na < + > (aq) Zn 02 < 2- > (aq) Na2 S (ai) -384240. Zn < 2 + > (aq) H Zn 02 < - > (aq) Zn C12 (aq) -457080. Zn Cl < + > (aq) Zn 0 H < + > (aq) Zn C13 < - > (aq) -330100. Zn C14 <2-> (aq) Zn 03 FI3 <-> (aq) Zn Cl 04 < + > (aq) -694220. Zn C12 08 (ai) Zn S 04 (ai) Zn 02 < 2- > (aq) -891590. H Zn 02 < - > (aq) Zn S 04 < + > (aq) Zn O H < + > (aq) -904900. Zn 03 H3 < - > (aq) Au Zn 04 H4 <2-> (aq) 0.0 Au C12 < - > (aq) Au 03 <3-> (aq) 336

-51800. TEMPERATURE Au C12 <-> (aq) 298.15 -151120. Au C14 <-> (aq) PRESSURE -235140. 1. Au 03 H3 -316920. INMOLE Au 03 H3 (aq) 0.8 -283370. 55.52 H Au 03 <2-> (aq) 0.01944 -142200. 0.0000000005 H2 Au 03 < - > (aq) 0.00005 -218300. 0.2 Au S <-> (aq) 0.2 46024. 0.04105 Au2 S 1.709 29037. C13 <-> (aq) TRY -120400. 0.023 .24E-16 .36E-2 .8 .62E-80 .28E-26 Zn C12 (aq) .42E-38 .25E-40 -403700. 55.5 .23E-10 .26E-1 3.617 .24E-26 .46E-28 Zn Cl < 4- > (aq) .22E-10 .68E-20 -275300. .42E-16 .19E-25 .62E-16 .45E-36 .27E-22 Zn C13 <-> (aq) .24E-22 .26E-2 .42E-7 -540500. .31E-10 .26E-10 .48E-24 3.657 .34E-84 Zn C14 <2-> (aq) .28E-10 .82E-10 .44E-9 -666000. .35E-8 .41E-8 .14E-32 .22E-58 .42E-2 Zn Cl 04 < -1- > (aq) .68E-2 .35E-7 .72E-2 -163100. .22E-2 .25E-10 .69E-22 .72E-16 .46E-15 Zn S .5E-12 .5E-14 .34E-34 -201290. .34E-22 .52E-5 .26E-1 .25E-11 .46E-87 Zn 04 H4 < 2- > (aq) .35E-88 .78E-58 .75E-2 -858520. .36E-78 .12E-92 .28E-52 .83E-94 .57E-4 0.0 0.0 0.0 0.0 0.0 0.0 FILE 0.0 0.0 0.0 0.0 0.0 CPDNBSDAT CPDNPLDAT EQUILIBRATE CPDMRLDAT CPDJANDAT STEP CPDSGTDAT outmole AQUEOUS INSPECIES Au C12 < - > (aq) GAS Au C14 < - > (aq) N2 (g) Au S < - > (aq) Zn C14 <2-> (aq) AQUEOUS H2 S (aq) H2 0 H S <-> (aq) Na Cl 0 (ai) S <2-> (aq) H2 S 04 (ai) S 04 <2-> (aq) Au C14 <-> (aq) Cl O <-> (aq) Na < + > (aq) C12 (aq) O H <-> (aq) H Cl O (aq) Cl 03 <-> (aq) SINGLE Zn S VARY Na Cl Inmole stepsize .0000000005 .005 .250 aqueous H2 S 04 (ai) 337

H < + > (aq) EQUILIBRATE VARY STEP Inmole stepsize .0996 .0000095 .1 activity aqueous AQUEOUS H2 S 04 (ai) H < + > (aq) EQUILIBRATE VARY Inmole stepsize .0000000005 .005 .250 STEP aqueous outmole H2 S 04 (ai) single Au EQUILIBRATE Au2 S Zn STEP ZnS outmole ZnS 04 single S Au Au2 S VARY Zn Inmole stepsize .0996 .0000095 .1 ZnS aqueous ZnS 04 H2 S 04 (ai) S EQUILIBRATE VARY Inmole stepsize .0000000005 .005 .250 aqueous STEP H2 S 04 (ai) outmole AQUEOUS EQUILIBRATE Au C12 <-> (aq) Au C14 < - > (aq) STEP Au S < - > (aq) outmole Zn C14 < 2- > (aq) AQUEOUS H2 S (aq) Au C12 < - > (aq) H S <-> (aq) Au C14 < - > (aq) S <2-> (aq) Au S < - > (aq) S 04 <2-> (aq) Zn C14 < 2- > (aq) Cl O < - > (aq) H2 S (aq) C12 (aq) H S <-> (aq) H Cl O (aq) S <2-> (aq) Cl 03 <-> (aq) S 04 <2-> (aq) CIO <-> (aq) VARY C12 (aq) Inmole stepsize .1 .000375 .115 H Cl O (aq) aqueous Cl 03 <-> (aq) H2 S 04 (ai)

VARY EQUILIBRATE Inmole stepsize .0996 .0000095 .1 aqueous STEP H2 S 04 (ai) activity AQUEOUS EQUILIBRATE H < + > (aq)

STEP VARY activity Inmole stepsize .1 .000375 .115 AQUEOUS aqueous 338

H2 S 04 (ai) Zn S 04 S EQUILIBRATE VARY STEP Inmole stepsize .1 .000375 .115 outmole aqueous single H2 S 04 (ai) Au Au2 S EQUILIBRATE Zn ZnS 339

Appendix 9.19 CHEMIX programme input statement for the system Au-

PbS-CI -Na0CI-H20 at 25°C: [PbS] = 16.72mM, [AuC14 ] = 50pM, [NaCl] = 1.709M and [OCT] = 0.194M 340

Title 53 <2-> (aq) CHEMIX CALCULATION OF GOLD H2 S (aq) SOLUBILITY H S <-> (aq) IN THE Au-PbS-NaCl-Na0Cl-H20 SYSTEM S 02 (aq) TO STUDY THE EFFECT OF SOLUTION S2 03 <2-> (aq) pH vs CONCENTRATION S2 04 <2-> (aq) S2 08 <2-> (aq) SYSTEM 54 06 < 2- > (aq) GAS H2 O (g) SINGLE H Cl (g) Au 02 (g) Au [O H]3 N2 (g) Au2 03 S2 (g) Au Cl C12 (g) Au C13 H2 (g) Au2 S H2 S 04 (g) Pb Pb S AQUEOUS Pb S 04 H2 O Pb C12 H < + > (aq) Pb 02 O H <-> (aq) Pb 02 H2 Cl <-> (aq) Pb O C12 (aq) Pb3 04 C13 <-> (aq) S H Cl (ai) Na2 S H Cl O (aq) Na Cl Na Cl O (ai) NaOH Cl 02 < - > (aq) Cl O <-> (aq) Cl 02 (aq) units Cl 03 <-> (aq) joules Cl 04 <-> (aq) S 04 <2-> (aq) Gibbs H S 04 <-> (aq) H Pb 02 <-> (aq) H2 S 04 (ai) -338420. Na < + > (aq) Pb O H < + > (aq) Pb <2+ > (aq) -226300. Pb C12 (aq) Pb 03 H3 <-> (aq) Pb Cl < + > (aq) -575600. Pb C13 <-> (aq) Au H Pb 02 < - > (aq) 0.0 Pb O H < + > (aq) Au 03 <3-> (aq) Pb 03 H3 <-> (aq) -51800. Au C12 <-> (aq) Au C12 < - > (aq) Au C14 < - > (aq) -151120. Au 03 <3-> (aq) Au C14 < - > (aq) H Au 03 <2-> (aq) -235140. H2 Au 03 < - > (aq) Au 03 H3 Au [O H]3 (aq) -316920. Au C14 H (ai) Au 03 H3 (aq) Au C14 Na (ai) -283370. Au S < - > (aq) H Au 03 <2-> (aq) S <2-> (aq) -142200. S2 <2-> (aq) H2 Au 03 < - > (aq) 341

-218300. 0.2 Au S <-> (aq) 0.01672 46024. 1.709 Au2 S 29037. TRY Ph S .25E-1 .24E-18 .0638 .8 .78E-78 .62E-28 -98700. .41E-40 .25E-43 Pb 02 H2 55.5 .23E-13 .166 3.617 .24E-30 .46E-30 -452200. .29E-11 .68E-20 C13 <-> (aq) .42E-15 .19E-25 .62E-14 .45E-35 .27E-20 -120400. .24E-21 .24E-2 .26E-12 Pb Cl < + > (aq) .48E-26 3.817 .35E-16 .34E-14 .49E-15 -164810. . 12E-15 .22E-5 .42E-10 Pb C12 (aq) .68E-6 .35E-10 .72E-24 .22E-12 .25E-12 -297160. .69E-10 .72E-12 .46E-38 Pb C13 <-> (aq) .56E-25 .15E-15 0.0 0.0 0.0 0.0 0.0 .16E-63 -426300. 0.0 0.0 .64E-52 0.0 Pb S 04 .5E-4 0.0 0.0 0.0 0.0 0.0 0.0 0.0 -813140. 0.0 0.0 0.0 .12E-1 0.0 0.0 0.0 0.0 0.0 0.0

FILE CPDNBSDAT EQUILIBRATE CPDNPLDAT CPDMRLDAT STEP CPDJANDAT out mole CPDSGTDAT AQUEOUS Au C12 < - > (aq) INSPECIES Au C14 <-> (aq) GAS Au S <-> (aq) N2 (g) Pb C13 <-> (aq) H2 S (aq) AQUEOUS H S <-> (aq) H2 O S <2-> (aq) Na Cl O (ai) S 04 <2-> (aq) H2 S 04 (ai) Cl O <-> (aq) Au C14 <-> (aq) C12 (aq) Na < + > (aq) H Cl O (aq) OH <-> (aq) Cl 03 <-> (aq)

SINGLE VARY Pb S Inmole stepsize .0000000005 .005 .250 Na Cl aqueous H2 S 04 (ai)

TEMPERATURE EQUILIBRATE 298.15 STEP PRESSURE activity

1. AQUEOUS H < + > (aq) INMOLE 0.8 VARY 55.52 Inmole stepsize .0000000005 .005 .250 0.1944 aqueous 0.0000000005 H2 S 04 (ai) 0.00005 0.2 EQUILIBRATE 342

Pb STEP Pb S outmole Pb S 04 single S Au Au2 S VARY Pb Inmole stepsize .082 .00002 .0834 Pb S aqueous Pb S 04 H2 S 04 (ai) S EQUILIBRATE VARY Inmole stepsize .0000000005 .005 .250 STEP aqueous outmole H2 S 04 (ai) AQUEOUS Au C12 <-> (aq) EQUILIBRATE Au C14 <-> (aq) Au S < - > (aq) STEP Pb C13 <-> (aq) outmole H2 S (aq) AQUEOUS H S < - > (aq) Au C12 < - > (aq) S <2-> (aq) Au C14 < - > (aq) S 04 <2-> (aq) Au S < - > (aq) Cl O <-> (aq) Pb C13 <-> (aq) C12 (aq) H2 S (aq) H Cl O (aq) H S <-> (aq) Cl 03 <-> (aq) S <2-> (aq) S 04 <2-> (aq) VARY Cl O <-> (aq) Inmole stepsize .0998 .00003 .1022 C12 (aq) aqueous H Cl O (aq) H2 S 04 (ai) Cl 03 <-> (aq) EQUILIBRATE VARY Inmole stepsize .082 .00002 .0834 STEP aqueous activity H2 S 04 (ai) AQUEOUS H < + > (aq) EQUILIBRATE VARY STEP Inmole stepsize .0998 .00003 .1022 activity aqueous AQUEOUS H2 S 04 (ai) H < + > (aq) EQUILIBRATE VARY Inmole stepsize .082 .00002 .0834 STEP aqueous outmole H2 S 04 (ai) single Au EQUILIBRATE Au2 S Pb STEP Pb S outmole Pb S 04 single S Au Au2 S VARY 343

Inmole stepsize .0998 .00003 .1022 EQUILIBRATE aqueous H2 S 04 (ai) 344

Appendix 9.20 CHEMIX programme input statement for the system Au-

PbS-Cl -Na0CI-H20 at 25°C: [PbS] = 16.72mM, [AuC14 ]

= 50jiM, [NaCl] = 1.709M and [OCT] = 0.019M 345

Title 52 <2-> (aq) CHEMIX CALCULATION OF GOLD 53 <2-> (aq) SOLUBILITY H2 S (aq) IN THE Au-PbS-NaCl-Na0Cl-H20 SYSTEM H S < - > (aq) TO STUDY THE EFFECT OF SOLUTION S 02 (aq) pH vs CONCENTRATION S2 03 <2-> (aq) S2 04 <2-> (aq) SYSTEM S2 08 <2-> (aq) GAS 54 06 <2-> (aq) H2 O (g) H Cl (g) SINGLE 02 (g) Au N2 (g) Au [O H]3 S2 (g) Au2 03 C12 (g) Au Cl H2 (g) Au C13 H2 S 04 (g) Au2 S Pb AQUEOUS Pb S H2 O Pb S 04 H < + > (aq) Pb C12 O H < - > (aq) Pb 02 Cl <-> (aq) Pb 02 H2 C12 (aq) Pb O C13 <-> (aq) Pb3 04 H Cl (ai) S H Cl O (aq) Na2 S Na Cl O (ai) Na Cl Cl 02 <-> (aq) NaO H Cl O <-> (aq) Cl 02 (aq) Cl 03 <-> (aq) units Cl 04 <-> (aq) joules S 04 <2-> (aq) H S 04 <-> (aq) Gibbs H2 S 04 (ai) H Pb 02 <-> (aq) Na < + > (aq) -338420. Pb < 2 + > (aq) Pb O H < + > (aq) Pb C12 (aq) -226300. Pb Cl < + > (aq) Pb 03 H3 <-> (aq) Pb C13 <-> (aq) -575600. H Pb 02 <-> (aq) Au Pb O H < + > (aq) 0.0 Pb 03 H3 <-> (aq) Au 03 <3-> (aq) Au C12 < - > (aq) -51800. Au C14 < - > (aq) Au C12 < - > (aq) Au 03 <3-> (aq) -151120. H Au 03 < 2- > (aq) Au C14 < - > (aq) H2 Au 03 <-> (aq) -235140. Au [O H]3 (aq) Au 03 H3 Au C14 H (ai) -316920. Au C14 Na (ai) Au 03 H3 (aq) Au S < - > (aq) -283370. S <2-> (aq) H Au 03 <2-> (aq) 346

-142200. 0.00005 H2 Au 03 <-> (aq) 0.2 -218300. 0.2 Au S <-> (aq) 0.01672 46024. 1.709 Au2 S 29037. TRY Pb S .25E-1 .24E-18 .0638 .8 .78E-78 .62E-28 -98700. .41E-40 .25E-43 Pb 02 H2 55.5 .23E-13 .166 3.617 .24E-30 .46E-30 -452200. .29E-11 .68E-20 C13 <-> (aq) .42E-15 .19E-25 .62E-14 .45E-35 .27E-20 -120400. .24E-21 .24E-2 .26E-12 Pb Cl < + > (aq) .48E-26 3.817 .35E-16 .34E-14 .49E-15 -164810. . 12E-15 .22E-5 .42E-10 Pb C12 (aq) .68E-6 .35E-10 .72E-24 .22E-12 .25E-12 -297160. .69E-10 .72E-12 .46E-38 Pb C13 <-> (aq) .56E-25 .15E-15 0.0 0.0 0.0 0.0 0.0.16E-63 -426300. 0.0 0.0 .64E-52 0.0 Pb S 04 .5E-4 0.0 0.0 0.0 0.0 0.0 0.0 0.0 -813140. 0.0 0.0 0.0 .12E-1 0.0 0.0 0.0 0.0 0.0 0.0

FILE CPDNBSDAT EQUILIBRATE CPDNPLDAT CPDMRLDAT STEP CPDJANDAT outmole CPDSGTDAT AQUEOUS Au C12 < - > (aq) INSPECIES Au C14 <-> (aq) GAS Au S < - > (aq) N2 (g) Pb C13 <-> (aq) H2 S (aq) AQUEOUS H S <-> (aq) H2 O S <2-> (aq) Na Cl O (ai) S 04 <2-> (aq) H2 S 04 (ai) Cl O <-> (aq) Au C14 <-> (aq) C12 (aq) Na < + > (aq) H Cl O (aq) O H <-> (aq) Cl 03 <-> (aq)

SINGLE VARY Pb S Inmole stepsize .0000000005 .005 .250 Na Cl aqueous H2 S 04 (ai)

TEMPERATURE EQUILIBRATE 298.15 STEP PRESSURE activity

1. AQUEOUS H < + > (aq) INMOLE 0.8 VARY 55.52 Inmole stepsize .0000000005 .005 .250 0.01944 aqueous 0.0000000005 H2 S 04 (ai) 347

Au EQUILIBRATE Au2 S Pb STEP Pb S outmole Pb S 04 single S Au Au2 S VARY Pb Inmole stepsize .0945 .000025 .09525 Pb S aqueous Pb S 04 H2 S 04 (ai) S EQUILIBRATE VARY Inmole stepsize .0000000005 .005 .250 STEP aqueous outmole H2 S 04 (ai) AQUEOUS Au C12 < - > (aq) EQUILIBRATE Au C14 < - > (aq) Au S < - > (aq) STEP Pb C13 < - > (aq) outmole H2 S (aq) AQUEOUS H S < - > (aq) Au C12 < - > (aq) S <2-> (aq) Au C14 < - > (aq) S 04 <2-> (aq) Au S < - > (aq) Cl O <-> (aq) Pb C13 <-> (aq) C12 (aq) H2 S (aq) H Cl O (aq) H S <-> (aq) Cl 03 <-> (aq) S <2-> (aq) S 04 <2-> (aq) VARY Cl O <-> (aq) Inmole stepsize .099915 .000015 .1004 C12 (aq) aqueous H Cl O (aq) H2 S 04 (ai) Cl 03 <-> (aq) EQUILIBRATE VARY Inmole stepsize .0945 .000025 .09525 STEP aqueous activity H2 S 04 (ai) AQUEOUS H < + > (aq) EQUILIBRATE VARY STEP Inmole stepsize .099915 .000015 .1004 activity aqueous AQUEOUS H2 S 04 (ai) H < + > (aq) EQUILIBRATE VARY Inmole stepsize .0945 .000025 .09525 STEP aqueous outmole H2 S 04 (ai) single Au EQUILIBRATE Au2S Pb STEP Pb S outmole Pb S 04 single S 348

VARY EQUILIBRATE Inmole stepsize .099915 .000015 .1004 aqueous H2 S 04 (ai) 349

Appendix 9.21 CHEMIX programme input statement for the system Au-

FeSrCl -NaOCI-H20 at 25°C: [FeS2] = 33.40mM, [AuC14 ]

= 50jiM, [NaCI] = 1.709M and [OCX] = 0.194M 350

Title Fe S 04 (ai) CHEMIX CALCULATION OF GOLD Fe S2 08 <-> (aq) SOLUBILITY S <2-> (aq) IN THE Au-FeS2-NaCl-Na0Cl-H20 52 <2-> (aq) SYSTEM 53 <2-> (aq) TO STUDY THE EFFECT OF SOLUTION H S <-> (aq) pH vs CONCENTRATION H2 S (aq) S 02 (aq) SYSTEM S 04 <2-> (aq) GAS S2 03 <2-> (aq) H2 O (g) S2 08 <2-> (aq) H Cl (g) 02 (g) SINGLE N2 (g) Au S2 (g) Au [O H]3 C12 (g) Au2 03 H2 (g) Au Cl H2 S 04 (g) Au C13 Au2 S AQUEOUS Fe H2 O Fe C12 H < + > (aq) Fe C13 O H <-> (aq) Fe2 03 Cl <-> (aq) Fe3 04 C12 (aq) FeS C13 <-> (aq) Fe S2 H Cl (ai) Fe S 04 H Cl O (aq) S Na Cl O (ai) Na2 S Cl 02 < - > (aq) Na2 S 04 Cl O <-> (aq) Na Cl Cl 02 (aq) NaOH Cl 03 <-> (aq) Cl 04 <-> (aq) S 04 <2-> (aq) units H S 04 <-> (aq) joules H2 S 04 (ai) Na < + > (aq) Gibbs Fe < 3 + > (aq) Fe 02 < 2- > (aq) Fe <2 + > (aq) -295300. Fe C12 (aq) H Fe 02 <-> (aq) Fe C13 (aq) -377700. Fe Cl <2 + > (aq) Fe S 04 (ai) Fe 02 <2-> (aq) -823430. H Fe 02 <-> (aq) Fe S 04 < + > (aq) Au C12 < - > (aq) -772700. Au C14 < - > (aq) Fe S2 08 < - > (aq) Au 03 <3-> (aq) -1524500. H Au 03 <2-> (aq) Fe C12 (aq) H2 Au 03 <-> (aq) -279100. Au [O H]3 (aq) Fe C13 (aq) Au C14 H (ai) -404500. Au C14 Na (ai) Fe Cl <2 + > (aq) Au S < - > (aq) -143900. Fe S 04 < + > (aq) Au 351

0.0 0.8 Au 03 <3-> (aq) 55.52 -51800. 0.1944 Au C12 <-> (aq) 0.0000000005 -151120. 0.00005 Au C14 <-> (aq) 0.2 -235140. 0.2 Au 03 H3 0.0334 -316920. 1.709 Au 03 H3 (aq) -283370. TRY H Au 03 <2-> (aq) .22E-1 .24E-16 .36E-70 .8 .62E-30 .24E-64 -142200. .24E-5 .69E-42 H2 Au 03 <-> (aq) 55.5 .23E-12 .26E-1 3.617 .24E-65 .46E-65 -218300. .2E-11 .68E-56 Au S < - > (aq) .42E-50 .19E-95 .62E-50 .45E-98 .27E-98 46024. .24E-98 .26E-1 .48E-11 Au2 S .39E-26 3.817 .34E-40 .28E-16 .82E-26 29037. .44E-36 .35E-38 .48E-10 Fe S2 .46E-8 .48E-26 .96E-32 .22E-64 .42E-62 -167000. .68E-62 .35E-58 .72E-75 S .22E-76 .52E-5 .25E-35 .69E-18 .2E-34 0.0 .44E-3 .12E-16 .57E-18 C13 <-> (aq) . 19E-3 .16E-9 .19E-25 .56E-12.16E-64 -120400. .5E-4 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 .72E-2 0.0 .66E-2 0.0 0.0 0.0 0.0 0.0 0.0 FILE CPDNBSDAT CPDNPLDAT EQUILIBRATE CPDMRLDAT CPDJANDAT STEP CPDSGTDAT outmole AQUEOUS INSPECIES Au C12 <-> (aq) GAS Au C14 <-> (aq) N2 (g) Au S <-> (aq) Fe < 3 + > (aq) AQUEOUS Fe <2 + > (aq) H2 O H2 S (aq) Na Cl 0 (ai) H S <-> (aq) H2 S 04 (ai) S <2-> (aq) Au C14 <-> (aq) S 04 <2-> (aq) Na < + > (aq) Cl O <-> (aq) O H <-> (aq) C12 (aq) H Cl O (aq) SINGLE Cl 03 <-> (aq) Fe S2 Na Cl VARY Inmole stepsize .0000000005 .005 .250 aqueous TEMPERATURE H2 S 04 (ai) 298.15 EQUILIBRATE PRESSURE L STEP activity INMOLE AQUEOUS 352

H < + > (aq) aqueous H2 S 04 (ai) VARY Iiimole stepsize .0000000005 .005 .250 EQUILIBRATE aqueous H2 S 04 (ai) STEP outmole EQUILIBRATE single Au STEP Au2 S outmole Fe single Fe2 03 Au Fe3 04 Au2 S Fe S Fe Fe S2 Fe2 03 Fe3 04 VARY FeS Inmole stepsize .0463 .000035 .0489 Fe S2 aqueous H2 S 04 (ai) VARY Inmole stepsize .0000000005 .005 .250 EQUILIBRATE aqueous H2 S 04 (ai) STEP outmole EQUILIBRATE AQUEOUS Au C12 <-> (aq) STEP Au C14 <-> (aq) outmole Au S < - > (aq) AQUEOUS Fe < 3 + > (aq) Au C12 <-> (aq) Fe <2 + > (aq) Au C14 <-> (aq) H2 S (aq) Au S < - > (aq) H S <-> (aq) Fe < 3 + > (aq) S <2-> (aq) Fe <2+> (aq) S 04 <2-> (aq) H2 S (aq) Cl O <-> (aq) H S <-> (aq) C12 (aq) S <2-> (aq) H Cl O (aq) S 04 <2-> (aq) Cl 03 <-> (aq) Cl O <-> (aq) C12 (aq) VARY H Cl O (aq) Inmole stepsize .04805 .0000005 .04809 Cl 03 <-> (aq) aqueous H2 S 04 (ai) VARY Inmole stepsize .0463 .000035 .0489 EQUILIBRATE aqueous H2 S 04 (ai) STEP activity EQUILIBRATE AQUEOUS H < + > (aq) STEP activity VARY AQUEOUS Inmole stepsize .04805 .0000005 .04809 H < + > (aq) aqueous H2 S 04 (ai) VARY Inmole stepsize .0463 .000035 .0489 EQUILIBRATE 353

Fe S2 STEP outmole VARY single Inmole stepsize .04805 .0000005 .04809 Au aqueous Au2 S H2 S 04 (ai) Fe Fe2 03 EQUILIBRATE Fe3 04 Fe S 354

Appendix 9.22 CHEMIX programme input statement for the system Au-

FeS2-Cr-Na0CI-H20 at 25°C: [FeS2] = 33.40mM, [AuCI4 ] = 50p.M, [NaCI] = 1.709M and [OCT] = 0.019M 355

Title Fe S 04 (ai) CHEMIX CALCULATION OF GOLD Fe S2 08 <-> (aq) SOLUBILITY S <2-> (aq) IN THE Au-FeS2-NaCl-Na0Cl-H20 52 <2-> (aq) SYSTEM 53 <2-> (aq) TO STUDY THE EFFECT OF SOLUTION H S <-> (aq) pH vs CONCENTRATION H2 S (aq) S 02 (aq) SYSTEM S 04 <2-> (aq) GAS S2 03 <2-> (aq) H2 O (g) S2 08 <2-> (aq) H Cl (g) 02 (g) SINGLE N2 (g) Au S2 (g) Au [O H]3 C12 (g) Au2 03 H2 (g) Au Cl H2 S 04 (g) Au C13 Au2 S AQUEOUS Fe H2 O Fe C12 H < + > (aq) Fe C13 O H < - > (aq) Fe2 03 Cl <-> (aq) Fe3 04 C12 (aq) FeS C13 <-> (aq) Fe S2 H Cl (ai) FeS 04 H Cl O (aq) S Na Cl O (ai) Na2 S Cl 02 <-> (aq) Na2 S 04 Cl O <-> (aq) Na Cl Cl 02 (aq) Na O H Cl 03 <-> (aq) Cl 04 <-> (aq) S 04 <2-> (aq) units H S 04 <-> (aq) joules H2 S 04 (ai) Na < + > (aq) Gibbs Fe < 3 + > (aq) Fe 02 < 2- > (aq) Fe < 2 + > (aq) -295300. Fe C12 (aq) H Fe 02 < - > (aq) Fe C13 (aq) -377700. Fe Cl <2+ > (aq) Fe S 04 (ai) Fe 02 < 2- > (aq) -823430. H Fe 02 <-> (aq) Fe S 04 < + > (aq) Au C12 < - > (aq) -772700. Au C14 < - > (aq) Fe S2 08 <-> (aq) Au 03 <3-> (aq) -1524500. H Au 03 < 2- > (aq) Fe C12 (aq) H2 Au 03 <-> (aq) -279100. Au [O H]3 (aq) Fe C13 (aq) Au C14 H (ai) -404500. Au C14 Na (ai) Fe Cl <2+ > (aq) Au S < - > (aq) -143900. Fe S 04 < + > (aq) Au 356

0.0 0.8 Au 03 <3-> (aq) 55.52 -51800. 0.01944 Au C12 <-> (aq) 0.0000000005 -151120. 0.00005 Au C14 <-> (aq) 0.2 -235140. 0.2 Au 03 H3 0.0334 -316920. 1.709 Au 03 H3 (aq) -283370. TRY H Au 03 <2-> (aq) .22E-1 .24E-16 .36E-70 .8 .62E-30 .24E-64 -142200. .24E-5 .69E-42 H2 Au 03 <-> (aq) 55.5 .23E-12 .26E-1 3.617 .24E-65 .46E-65 -218300. .2E-11 .68E-56 Au S <-> (aq) .42E-50.19E-95 .62E-50 .45E-98 .27E-98 46024. .24E-98 .26E-1 .48E-11 Au2 S .39E-26 3.817 .34E-40 .28E-16 .82E-26 29037. .44E-36 .35E-38 .48E-10 Fe S2 .46E-8 .48E-26 .96E-32 .22E-64 .42E-62 -167000. .68E-62 .35E-58 .72E-75 S .22E-76 .52E-5 .25E-35 .69E-18 .2E-34 0.0 .44E-3 .12E-16 .57E-18 Cl3 <-> (aq) . 19E-3 .16E-9 .19E-25 .56E-12.16E-64 -120400. .5E-4 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 .72E-2 0.0 .66E-2 0.0 0.0 0.0 0.0 0.0 0.0 FILE CPDNBSDAT CPDNPLDAT EQUILIBRATE CPDMRLDAT CPDJANDAT STEP CPDSGTDAT outmole AQUEOUS INSPECIES Au C12 < - > (aq) GAS Au C14 < - > (aq) N2 (g) Au S < - > (aq) Fe <3 4- > (aq) AQUEOUS Fe <2+ > (aq) H2 0 H2 S (aq) Na Cl O (ai) H S <-> (aq) H2 S 04 (ai) S <2-> (aq) Au C14 <-> (aq) S 04 <2-> (aq) Na < + > (aq) Cl O <-> (aq) 0 H <-> (aq) C12 (aq) H Cl O (aq) SINGLE Cl 03 <-> (aq) Fe S2 Na Cl VARY Inmole stepsize .0000000005 .005 .250 aqueous TEMPERATURE H2 S 04 (ai) 298.15 EQUILIBRATE PRESSURE 1. STEP activity INMOLE AQUEOUS 357

H < + > (aq) aqueous H2 S 04 (ai) VARY Inmole stepsize .0000000005 .005 .250 EQUILIBRATE aqueous H2 S 04 (ai) STEP outmole EQUILIBRATE single Au STEP Au2 S outmole Fe single Fe2 03 Au Fe3 04 Au2 S Fe S Fe Fe S2 Fe2 03 Fe3 04 VARY FeS Inmole stepsize .0933 .0001 .098 Fe S2 aqueous H2 S 04 (ai) VARY Inmole stepsize .0000000005 .005 .250 EQUILIBRATE aqueous H2 S 04 (ai) STEP outmole EQUILIBRATE AQUEOUS Au C12 < - > (aq) STEP Au C14 < - > (aq) outmole Au S < - > (aq) AQUEOUS Fe < 3 + > (aq) Au C12 <-> (aq) Fe <2+ > (aq) Au C14 <-> (aq) H2 S (aq) Au S < - > (aq) H S <-> (aq) Fe < 3 + > (aq) S <2-> (aq) Fe <2 + > (aq) S 04 <2-> (aq) H2 S (aq) Cl O <-> (aq) H S <-> (aq) C12 (aq) S <2-> (aq) H Cl O (aq) S 04 <2-> (aq) Cl 03 <-> (aq) Cl O <-> (aq) C12 (aq) VARY H Cl O (aq) Inmole stepsize .09472 .0000005 .09474 Cl 03 <-> (aq) aqueous H2 S 04 (ai) VARY Inmole stepsize .0933 .0001 .098 EQUILIBRATE aqueous H2 S 04 (ai) STEP activity EQUILIBRATE AQUEOUS H < + > (aq) STEP activity VARY AQUEOUS Inmole stepsize .09472 .0000005 .09474 H < + > (aq) aqueous H2 S 04 (ai) VARY Inmole stepsize .0933 .0001 .098 EQUILIBRATE 358

Fe S2 STEP outmole VARY single Inmole stepsize .09472 .0000005 .09474 Au aqueous Au2 S H2 S 04 (ai) Fe Fe2 03 EQUILIBRATE Fe3 04 Fe S 359

Appendix 9.23 CHEMIX programme input statement for the system Au- CuFeS2-Cl -Na0Cl-H20 at 25°C: [CuFeS2] = 21.80mM,

[AuCI4 ] = 50p.M, [NaCl] = 1.709M and [OCX] = 0.194M 360

Title Fe Cl <2 + > (aq) CHEMIX CALCULATION OF GOLD Fe 02 <2-> (aq) SOLUBILITY H Fe 02 <-> (aq) IN THE Au-CuFeS2-NaCl-Na0Cl-H20 Fe 03 H3 <-> (aq) SYSTEM Fe 04 H4 <2-> (aq) TO STUDY THE EFFECT OF SOLUTION Fe S 04 < + > (aq) pH vs CONCENTRATION Fe S 04 (ai) Fe S2 08 < - > (aq) SYSTEM Au C12 <-> (aq) GAS Au C14 <-> (aq) H2 O (g) Au 03 <3-> (aq) H Cl (g) H Au 03 <2-> (aq) 02 (g) H2 Au 03 <-> (aq) N2 (g) Au [O H]3 (aq) S2 (g) Au C14 H (ai) C12 (g) Au C14 Na (ai) H2 (g) Au S < - > (aq) H2 S 04 (g) S <2-> (aq) S2 <2-> (aq) AQUEOUS S 02 (aq) H2 O S2 03 <2-> (aq) H < + > (aq) S2 08 < 2- > (aq) O H < - > (aq) Cl <-> (aq) SINGLE C12 (aq) Cu C13 <-> (aq) Cu Cl H Cl (ai) Cu C12 H Cl O (aq) Cu O Na Cl O (ai) Cu2 O Cl 02 <-> (aq) Cu2 S Cl O <-> (aq) Cu S Cl 02 (aq) Cu S 04 Cl 03 <-> (aq) Cu Fe S2 Cl 04 <-> (aq) Fe S 04 <2-> (aq) FeO H S 04 <-> (aq) Fe2 03 H2 S 04 (ai) Fe3 04 H S <-> (aq) Fe 03 H3 H2 S (aq) Fe C12 Na < + > (aq) FeS Na2 S (ai) Fe S2 Na S 04 <-> (aq) FeS 04 Cu <2 + > (aq) S Cu C12 (aq) Au Cu Cl < + > (aq) Au 03 H3 Cu C12 <-> (aq) Au2 03 Cu C13 <2-> (aq) Au Cl Cu 02 <2-> (aq) Au C13 Cu 02 H2 (aq) Au2 S Cu S 04 (aq) Na2 S Cu S 03 < - > (aq) Na Cl Fe < 3 + > (aq) NaOH Fe <2 + > (aq) Fe C12 (aq) Fe C13 (aq) units 361 joules -349800. C13 <-> (aq) Gibbs -120400. Fe 02 <2-> (aq) -295300. H Fe 02 <-> (aq) FILE -377700. CPDNBSDAT Fe 03 H3 <-> (aq) CPDNPLDAT -614900. CPDMRLDAT Fe 04 H4 <2-> (aq) CPDJANDAT -769700. CPDSGTDAT Fe C12 (aq) -279100. INSPECIES Fe C13 (aq) GAS -404500. N2 (g) Fe S2 08 <-> (aq) -1524500. AQUEOUS Cu C12 (aq) H2 0 -197900. Na Cl O (ai) Cu Cl < + > (aq) H2 S 04 (ai) -68200. Au C14 < - > (aq) Cu C12 <-> (aq) Na < + > (aq) -240100. O H < - > (aq) CuC13 <2-> (aq) -376000. SINGLE Cu S 04 (aq) Cu Fe S2 -692180. Na Cl Cu S 03 <-> (aq) -481100. Cu 02 <2-> (aq) TEMPERATURE -183600. 298.15 Cu 02 H2 (aq) -249010. PRESSURE Cu Fe S2 1. -227390. Au INMOLE 0.0 0.8 Au C12 <-> (aq) 55.52 -151120. 0.1944 Au C14 <-> (aq) 0.0000000005 -235140. 0.00005 Au 03 <3-> (aq) 0.2 -51800. 0.2 H Au 03 < 2- > (aq) 0.0218 -142200. 1.709 H2 Au 03 < - > (aq) -218300. TRY Au 03 H3 (aq) .23E-1 .24E-17 .36E-2 .8 .62E-88 .36E-28 -283370. .28E-40 .78E-42 Au 03 H3 55.5 .23E-12 .113 3.617 .24E-30 .46E-30 -316920. .22E-12 .68E-20 Au S < - > (aq) .42E-15 .26E-26 .39E-15 .19E-36 .52E-22 46024. .26E-24 .48E-2 .39E-12 Au2 S .24E-27 .32E-86 .32E-88 3.775 .54E-98 29037. .24E-1 .62E-16 .36E-15 Na2 S (ai) .45E-16 .32E-15 .48E-14 .44E-5 .48E-18 -438100. . 18E-16 .22E-56 .72E-38 Na2 S 362

.22E-34 .25E-44 .32E-37 .58E-38 .39E-26 outmole .24E-25 .14E-25 .82E-27 single . 14E-38 .74E-36 .63E-40 .52E-10 .65E-14 Au .69E-14.16E-12 .82E-12 Au2 S .5IE-13 .25E-39 .26E-25 .54E-5 .43E-93 Au 03 H3 .41E-97 .16E-60 .56E-93 Cu .22E-52 Cu C12

0.0 0.0 0.0 .21E-1 0.0 0.0 0.0 0.0 Cu2 S 0.0 0.0 0.0 . IE-1 0.0 0.0 0.0 0.0 Cu S 0.0 0.0 0.0 .5E-4 0.0 0.0 0.0 0.0 Cu Fe S2 0.00.00.00.0 Fe Fe C12 Fe2 03 EQUILIBRATE Fe3 04 Fe S2 FeS STEP outmole VARY AQUEOUS Inmole stepsize .0000000005 .003 .22 Au C12 <-> (aq) aqueous Au C14 <-> (aq) H2 S 04 (ai) Au S < - > (aq) H2 S (aq) EQUILIBRATE H S <-> (aq) S <2-> (aq) S 04 <2-> (aq) STEP Cl O <-> (aq) outmole C12 (aq) AQUEOUS H Cl O (aq) Au C12 < - > (aq) Cl 03 <-> (aq) Au C14 < - > (aq) Fe < 3 + > (aq) Au S <-> (aq) Fe <2 + > (aq) H2 S (aq) Fe Cl <2 + > (aq) H S < - > (aq) Fe S 04 < + > (aq) S <2-> (aq) Cu <2 + > (aq) S 04 <2-> (aq) Cu Cl < -I- > (aq) Cl O <-> (aq) Cu C12 <-> (aq) C12 (aq) H Cl O (aq) VARY Cl 03 <-> (aq) Inmole stepsize .0000000005 .003 .22 Fe < 3 + > (aq) aqueous Fe < 2 + > (aq) H2 S 04 (ai) Fe Cl <2+> (aq) Fe S 04 < + > (aq) EQUILIBRATE Cu <2+> (aq) Cu Cl < + > (aq) STEP Cu C12 < - > (aq) activity AQUEOUS VARY H < + > (aq) Inmole stepsize .0557 .00001 .0565 aqueous VARY H2 S 04 (ai) Inmole stepsize .0000000005 .003 .22 aqueous EQUILIBRATE H2 S 04 (ai) STEP EQUILIBRATE activity AQUEOUS STEP H < + > (aq) 363

Fe S 04 < + > (aq) VARY Cu <2 + > (aq) Inmole stepsize .0557 .00001 .0565 Cu Cl < + > (aq) aqueous Cu C12 <-> (aq) H2 S 04 (ai) VARY EQUILIBRATE Inmole stepsize .078 .00008 .080 aqueous STEP H2 S 04 (ai) outmole single Equilibrate Au Au2 S STEP Au 03 H3 activity Cu AQUEOUS Cu C12 H < + > (aq) Cu2 S Cu S VARY Cu Fe S2 Inmole stepsize .078 .00008 .080 Fe aqueous Fe C12 H2 S 04 (ai) Fe2 03 Fe3 04 EQUILIBRATE Fe S2 Fe S STEP outmole VARY single Inmole stepsize .0557 .00001 .0565 Au aqueous Au2 S H2 S 04 (ai) Au 03 H3 Cu EQUILIBRATE Cu C12 Cu2 S CuS STEP Cu Fe S2 outmole Fe AQUEOUS Fe C12 Au C12 < - > (aq) Fe2 03 Au C14 <-> (aq) Fe3 04 Au S < - > (aq) Fe S2 H2 S (aq) FeS H S <-> (aq) S <2-> (aq) VARY S 04 <2-> (aq) Inmole stepsize .078 .00008 .080 Cl O <-> (aq) aqueous C12 (aq) H2 S 04 (ai) H Cl O (aq) Cl 03 <-> (aq) EQUILIBRATE Fe < 3 + > (aq) Fe < 2 + > (aq) Fe Cl <2+ > (aq) 364

Appendix 9.24 CHEMIX programme input statement for the system Au-

CuFeS2-Cl-NaOCI-H20 at 25°C: [CuFeS2] = 21.80mM,

[AuC14 ] = 50jiM, [NaCI] = 1.709M and [OCT] = 0.019M 365

Title Fe Cl < 2 + > (aq) CHEMIX CALCULATION OF GOLD Fe 02 < 2- > (aq) SOLUBILITY H Fe 02 < - > (aq) IN THE Au-CuFeS2-NaCl-Na0Cl-H20 Fe 03 H3 <-> (aq) SYSTEM Fe 04 H4 <2-> (aq) TO STUDY THE EFFECT OF SOLUTION Fe S 04 < + > (aq) pH vs CONCENTRATION Fe S 04 (ai) Fe S2 08 < - > (aq) SYSTEM Au C12 < - > (aq) GAS Au C14 < - > (aq) H2 O (g) Au 03 <3-> (aq) H Cl (g) H Au 03 <2-> (aq) 02 (g) H2 Au 03 <-> (aq) N2 (g) Au [O H]3 (aq) S2 (g) Au C14 H (ai) 012 (g) Au C14 Na (ai) H2 (g) Au S < - > (aq) H2 S 04 (g) S <2-> (aq) S2 <2-> (aq) AQUEOUS S 02 (aq) H2 O S2 03 <2-> (aq) H < + > (aq) S2 08 <2-> (aq) O H < - > (aq) Cl <-> (aq) SINGLE C12 (aq) Cu C13 <-> (aq) Cu Cl H Cl (ai) Cu C12 H Cl O (aq) Cu O Na Cl O (ai) Cu2 O Cl 02 <-> (aq) Cu2 S Cl O <-> (aq) Cu S Cl 02 (aq) Cu S 04 Cl 03 <-> (aq) Cu Fe S2 Cl 04 <-> (aq) Fe S 04 <2-> (aq) FeO H S 04 <-> (aq) Fe2 03 H2 S 04 (ai) Fe3 04 H S <-> (aq) Fe 03 H3 H2 S (aq) Fe C12 Na < + > (aq) FeS Na2 S (ai) Fe S2 Na S 04 <-> (aq) FeS 04 Cu < 2 + > (aq) S Cu C12 (aq) Au Cu Cl < + > (aq) Au 03 H3 Cu C12 < - > (aq) Au2 03 Cu C13 <2-> (aq) Au Cl Cu 02 < 2- > (aq) Au C13 Cu 02 H2 (aq) Au2 S Cu S 04 (aq) Na2 S Cu S 03 < - > (aq) Na Cl Fe < 3 + > (aq) NaOH Fe < 2 + > (aq) Fe C12 (aq) Fe C13 (aq) units 366 joules -349800. C13 <-> (aq) Gibbs -120400. Fe 02 <2-> (aq) -295300. H Fe 02 < - > (aq) FILE -377700. CPDNBSDAT Fe 03 H3 < - > (aq) CPDNPLDAT -614900. CPDMRLDAT Fe 04 H4 <2-> (aq) CPDJANDAT -769700. CPDSGTDAT Fe C12 (aq) -279100. INSPECIES Fe C13 (aq) GAS -404500. N2 (g) Fe S2 08 < - > (aq) -1524500. AQUEOUS Cu C12 (aq) H2 O -197900. Na Cl O (ai) Cu Cl < + > (aq) H2 S 04 (ai) -68200. Au C14 <-> (aq) Cu C12 <-> (aq) Na < + > (aq) -240100. O H < - > (aq) Cu C13 <2-> (aq) -376000. SINGLE Cu S 04 (aq) Cu Fe S2 -692180. Na Cl Cu S 03 < - > (aq) -481100. Cu 02 <2-> (aq) TEMPERATURE -183600. 298.15 Cu 02 H2 (aq) -249010. PRESSURE Cu Fe S2 1. -227390. Au INMOLE 0.0 0.8 Au C12 <-> (aq) 55.52 -151120. 0.01944 Au C14 <-> (aq) 0.0000000005 -235140. 0.00005 Au 03 <3-> (aq) 0.2 -51800. 0.2 H Au 03 <2-> (aq) 0.0218 -142200. 1.709 H2 Au 03 < - > (aq) -218300. TRY Au 03 H3 (aq) .23E-1 .24E-17 .36E-2 .8 .62E-88 .36E-28 -283370. .28E-40 .78E-42 Au 03 H3 55.5 .23E-12 .113 3.617 .24E-30 .46E-30 -316920. .22E-12 .68E-20 Au S <-> (aq) .42E-15 .26E-26 .39E-15 .19E-36 .52E-22 46024. .26E-24 .48E-2 .39E-12 Au2 S .24E-27 .32E-86 .32E-88 3.775 .54E-98 29037. .24E-1 .62E-16 .36E-15 Na2 S (ai) .45E-16 .32E-15 .48E-14 .44E-5 .48E-18 -438100. . 18E-16 .22E-56 .72E-38 Na2 S 367

.22E-34 .25E-44 .32E-37 .58E-38 .39E-26 outmole .24E-25 .14E-25 .82E-27 single . 14E-38 .74E-36 .63E-40 .52E-10 .65E-14 Au .69E-14.16E-12 .82E-12 Au2 S .51E-13 .25E-39 .26E-25 .54E-5 .43E-93 Au 03 H3 .41E-97 .16E-60 .56E-93 Cu .22E-52 Cu C12 0.0 0.0 0.0 .21E-1 0.0 0.0 0.0 0.0 Cu2 S 0.0 0.0 0.0 . IE-1 0.0 0.0 0.0 0.0 Cu S 0.0 0.0 0.0 .5E-4 0.0 0.0 0.0 0.0 Cu Fe S2 0.0 0.0 0.0 0.0 Fe Fe C12 Fe2 03 EQUILIBRATE Fe3 04 Fe S2 FeS STEP outmole VARY AQUEOUS Inmole stepsize .0000000005 .003 .22 Au C12 < - > (aq) aqueous Au C14 <-> (aq) H2 S 04 (ai) Au S <-> (aq) H2 S (aq) EQUILIBRATE H S <-> (aq) S <2-> (aq) S 04 <2-> (aq) STEP Cl O <-> (aq) outmole C12 (aq) AQUEOUS H Cl O (aq) Au C12 <-> (aq) Cl 03 <-> (aq) Au C14 <-> (aq) Fe < 3 + > (aq) Au S < - > (aq) Fe <2 + > (aq) H2 S (aq) Fe Cl <2 + > (aq) H S < - > (aq) Fe S 04 < + > (aq) S <2-> (aq) Cu <2+> (aq) S 04 <2-> (aq) Cu Cl < + > (aq) Cl O < - > (aq) Cu C12 <-> (aq) C12 (aq) H Cl O (aq) VARY Cl 03 <-> (aq) Inmole stepsize .0000000005 .003 .22 Fe < 3 + > (aq) aqueous Fe < 2 + > (aq) H2 S 04 (ai) Fe Cl <2 + > (aq) Fe S 04 < + > (aq) EQUILIBRATE Cu < 2 + > (aq) Cu Cl < + > (aq) STEP Cu C12 < - > (aq) activity AQUEOUS VARY H < + > (aq) Inmole stepsize .09500 .000005 .09530 aqueous VARY H2 S 04 (ai) Inmole stepsize .0000000005 .003 .22 aqueous EQUILIBRATE H2 S 04 (ai) STEP EQUILIBRATE activity AQUEOUS STEP H < + > (aq) 368

Fe Cl <2 + > (aq) VARY Fe S 04 < + > (aq) Inmole stepsize .09500 .000005 .09530 Cu <2 + > (aq) aqueous Cu Cl < + > (aq) H2 S 04 (ai) Cu C12 <-> (aq)

EQUILIBRATE VARY Inmole stepsize .0950460 .0000001 .095050 STEP aqueous outmole H2 S 04 (ai) single Au Equilibrate Au2 S Au 03 H3 STEP Cu activity Cu C12 AQUEOUS Cu2 S H < + > (aq) Cu S Cu Fe S2 VARY Fe Inmole stepsize .0950460 .0000001 .095050 Fe C12 aqueous Fe2 03 H2 S 04 (ai) Fe3 04 Fe S2 EQUILIBRATE Fe S STEP VARY outmole Inmole stepsize .09500 .000005 .09530 single aqueous Au H2 S 04 (ai) Au2 S Au 03 H3 EQUILIBRATE Cu Cu C12 Cu2 S STEP Cu S outmole Cu Fe S2 AQUEOUS Fe Au C12 < - > (aq) Fe C12 Au C14 < - > (aq) Fe2 03 Au S < - > (aq) Fe3 04 H2 S (aq) Fe S2 H S <-> (aq) FeS S <2-> (aq) S 04 <2-> (aq) VARY Cl O <-> (aq) Inmole stepsize .0950460 .0000001 .095050 C12 (aq) aqueous H Cl O (aq) H2 S 04 (ai) Cl 03 <-> (aq) Fe < 3 + > (aq) EQUILIBRATE Fe < 2 + > (aq)