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ELECTROWINNING COUPLED TO GOLD LEACHING
BY ELECTROGENERATED CHLORINE
A THESIS SUBMITTED FOR THE DEGREE OF DOCTOR
OF PHILOSOPHY OF THE UNIVERSITY OF LONDON
AND
THE DIPLOMA OF IMPERIAL COLLEGE
BY
MIGUEL ANGEL DIAZ
Department of Mineral Resources Engineering Imperial College of Science and Technology University of LondonLondon SW7 2BP December 19S6
1 .
Dedicated to my family for all that they
have given me.
2 .
ABSTRACT
The kinetics of the conventional gold leaching process by dissolved
oxygen in cyanide solution suffers from slow kinetics limited by the
aqueous solubility of oxygen. The objectives of this Ph.D project
were:
a) to improve the leaching kinetics of gold by using chlorine/hypochlorite in solutions containing chloride ions.as complexants.
b) to recover that gold from very dilute solutions by electrowinning in packed bed electrodes, and
c) to test the feasibility of coupling the two processes, using anodically generated chlorine as the oxidant in an external leach reactor.
The thermodynamics of the AU-CI-H2O and Ag-Cl-HgO systems were summarised
in the form of potential-pH and activity-pH diagrams; silver was included
in these calculations since it often occurs alloyed with gold and may passivate the leach reaction.
Cyclic voltammetry at a (Pt) rotating disc electrode (RDE) was used to
study the kinetics of the various reactions in the system; for some
experiments the disc was gold-plated. Electrochemical reduction of
chlorine and mixed potential leaching of gold, were studied over the pH
range 0 - 7 using the RDE. The higher solubility of chlorine/hypochlorite
compared with oxygen enabled gold leaching rates more than two orders of
magnitude greater than achievable with a conventional alkaline cyanide-
dissolved oxygen system. Gold leaching rates obtained by the mixed
potential and net current methods agreed within 10% and AuCL^ ions
were the predominant product in CI2/HCIO systems.
Batch recycle electrolyses through the cathode were carried out to deplete
the dissolved gold from various initial ‘concentrations as functions of
3 .
flow rate, applied potential and pH. The Au(I)/Au(III) molar ratio was
determined; Au(I) species were depleted more rapidly than Au(III), but
at longer times and lower total gold concentrations, Au(III) reduction
produced Au(I) species which were then dispersed into bulk solution
before being further reduced. A PBE enabled the dissolved gold con
centration to be depleted to below analytically detectable levels
( <0.1 g Au m”3). in coupled leach-electrowinning experiments the
dissolved gold concentration-time relationships at the inlet and outlet
of the leach reactor and the outlet of the cathode compartment were stable at pH = 0, demonstrating the feasibility of the coupled system.
4 .
ACKNOWLEDGEMENTS
I should thank Dr G.H. Kelsall for his guidance and support during
the course of this work. Thanks must also go to Dr C.I. House for
reading the manuscript and useful suggestions and to my colleagues and technicians in the Mineral Technology research section. Thanks to
Janel for typing and assisting in the production of this thesis at very
short notice. Thanks to my wife T.J.V. Sokoloff de Diaz for her tremendous
support and understanding. I would also like to acknowledge the financial
support of the British Council and the Universidad Catolica Madre y
Maestra, Santiago, Dominican Republic. Finally, thanks to the human kind
for the pool of knowledge it has created.
5 .
CONTENTSPAGE
ABSTRACT 2
ACKNOWLEDGEMENT 4
CONTENTS 5
LIST OF FIGURES 7
LIST OF TABLES 13
NOMENCLATURE 14
CHAPTER 1 : INTRODUCTION 17
CHAPTER 2 : ALTERNATIVE LEACHING SYSTEMS TO CHLORINATION 20
2.1 Types of Gold Ores 212.1.1 Free Gold Ores 212.1.2 Gold with Iron Sulphides 222.1.3 Gold with Arsenic and/or Antimony Minerals 222.1.4 Gold Tellurides 242.1.5 Gold with Copper Porphyries 242.1.6 Gold with Lead and Zinc Minerals 252.1.7 Carbonaceous Ores 25
2.2 Gold Leaching in Oxygenated Alkaline Cyanide Solutions 252.2.1 Thermodynamics of the Ag/Au-CN-^O System 262.2.2 Mechanism of Cyanidation 272.2.2.1 Oxygen Reduction on Gold 322.2.2.2 The Anodic Reaction of Gold in Alkaline Media 352.2.2.3 Cyanide Consumption 382.2.3 Kinetics of Cyanidation 402.2.4 Future Research 43
2.3 Acid Thiourea Leaching of Gold 43
CHAPTER 3 : REDUCTION OF DISSOLVED AQUEOUS CHLORINE/HYPOCHLORITE 48
3.1 Platinum 493.2 Graphite 503.3 Titanium-Base Electrodes 51
CHAPTER 4 : THE CHEMISTRY OF GOLD AND SILVER IN ACIDIC SOLUTIONS 52
4.1 Gold 534.1.1 Introduction 534.1.2 Oxidation of Gold in Acidic Media 564.1.3 Gbrrosion of Gold in Chloride Media 674.1.4 The Electrodeposition of Gold from Chloride Media 774.1.5 Corrosion of Gold-Silver Alloys in Chloride Media 794.1.6 Chlorine as an Oxidant in Gold Hydrometallurgy 814.1.7 Recent Attempts at Chlorination 82
6 .PAGE
4.2 Silver 884 .2 .1 Introduct ion 884.2.2 Oxidation of Silver in Acidic Non-Oomplexing Media 894.2.3 The Oxidation of Silver in Acidic Chloride Media 90
CHAPTER 5 : EXPERIMENTAL 93
5.1 Analytical Techniques 945.2 Solution Preparation 945.3 Electrode Construction and Preparation 955.4 Electrochemical Studies 965.4.1 Electrochemical Instrumentation 965.4.2 Cyclic Voltammetry 965.4.3 Cbnstant Potential Electrolysis 985.4.4 Constant Current Electrolysis 985.5 Spectrophotometric Studies 995.6 Mixed Potential Measurements 995.7 Net Current Method 1015.8 Flow Circuit Experiments 101
CHAPTER 6 : RESULTS AND DISCUSSIONS 107
6 .1 Silver 1086.2 Gold 1146.2.1 Thermodynami cs 1146.2.2 Cyclic Voltammetry 1236.2.3 UV Spectrophotometry 1326.2.4 Electrochemical Dissolution of Gold 1356.2.5 Mixed Potential Leaching Rates 1406.2.6 Chlorine Reduction on Gold Surfaces 1476.2.7 Electro-deposition of gold in a Packed Bed Electrode 1556.2.8 Coupled Chlorine Leaching and Electrowinning of Gold 1746.2.9 Process Considerations 179
CHAPTER 7 : CONCLUSIONS 186
REFERENCES 190
APPENDIX I : Eh-pH activity equations for the Ag/^O/Cl/ClOq system 209
APPENDIX II : Eh-pH activity equations for the AU/H2O/CI system 218ERRATA 237
7 .
LIST OF FIGURES
Figure 2.1 Generalized cyanidation flowsheet.
Figure 2.2 Eh-pH diagram for the AU/H2O at 298 K.
Figure 2.3 Eh-pH diagram for the Au/^O-CN system at 298 K.
Figure 2.4 Eh-pH diagram for the Ag/H20-CN system at 298 K.
Figure 2.5 General scheme of oxygen reduction.
Figure.4.1 Potent10static polarization of gold in 0.25 M sulphuric acid as a function of NaCl concentration.
Figure 4.2 Partial current density-potential curves for Au(III) and Au(I) dissolution in 0.25 M sulphuric acid as a function of NaCl concentration.
Figure 4.3 Process proposed by Walker for the recovery of gold from Merrill slimes.
Figure 4.4 Proposed process by Finkelstein for the wet chlorination treatment of Merrill slimes.
Figure 4.5 Gold dissolution as a function of redox potential.
Figure 5.1 Electrochemical cell design for experiments with planar electrodes.
Figure 5.2 Electrochemical cell design for experiments with Pt rotating disc electrode.
Figure 5.3 Apparatus to study the disproportionation of AuCl^.
Figure 5.4 Stop-flow apparatus.
Figure 5.5 Perspex packed bed electrode cell.
Figure 5.6 Reactor flow circuit.
8.
Figure 6.1 Potential - pH diagram for the Ag/^O system at 298 K.
Figure 6.2 Potential - pH diagram of the Ag/^O system at 298 K by Pourbaix.
Figure 6.3 Potential - pH diagram for the Ag/H20-Cl-C104 system at 298 K, with dissolved silver, chloride, and perchlorate activities of 10-4, 10“5 and 10“5 respectively.
Figure 6.4 Potential - pH diagram for the Ag/H20-Cl-CL04 system at 298 K, with dissolved silver, chloride, and perchlorate activities of 10-4, 1 .0 and 1 .0 respectively.
Figure 6.5 Activity - pH diagram for the Ag(I)/H20-Cl system at 298 K( - ) chloride activity of 1.0 , ( --- ) chloride activityof 5.0.
Figure 6 .6 Potential - pH diagram for the AU/H2O system at 298 K, with a dissolved gold activity of 10-4.
Figure 6.7 Potential - pH diagram for the AU/H2O-CI system at 298 K, with dissolved gold and chloride activities of 10*"4 and 10-5 respectively.
Figure 6 .8 Potential - pH diagram for the AU/H2O-CL system at' 298 K, with dissolved gold and chloride activities of 10~4 and 10-5, respectively, considering mixed hydroxide-chloride species.
Figure 6.9 Potential - pH diagram for the AU/H2O-CL system at 298 K,with dissolved gold and chloride activities of 5 x 10~5 and0.5 respectively.
Figure 6.10 Potential - pH diagram for the AU/H2O-CL system at 298 K, with dissolved gold and chloride activities of 2 .5 x 10“5 and 5.0, respectively.
Figure 6.11 Activity - pH diagram for the Au(III)/H20-C1 system at 298 K, chloride activity of 1.0 .
Figure 6.12 The effect of rotation rate on the current-potential behaviour of a gold plated Pt disc electrode in 1 kmol NaCl m“3, pH =5.4, sweep rate = 10 mV s“l, ( - ) stationary ( -- ) 4 Hz at295K.
Figure 6.13 Cyclic voltammogram of a planar Pt electrode in a quiescentelectrolyte containing 162.2 g m”3 total gold, Au(III) = 123.1 g m -3 + Au(I) = 39.1 g m~3 , in 1 kmol HC1 m~3, sweep rate =1 mV s"1 at 295 K.
9.
Figure
Figure
Figure
Figure
Figure
Figure
Figure
Figure
Figure
Figure
Figure
Figure
Figure
6.14 Cyclic voltairmogram of a planar Pt electrode in an electrolyte containing 162.2 g m-3 total gold, Au(III) = 123.1 g m-3 +Au(I) = 39.1 g nr3 , in 1 kmol H d m“3 , moderate stirring, sweep rate = 1 mV s“l at 295 K.
6.15 Cyclic voltammogram of a planar Pt electrode in a quiescent electrolyte containing 51.4 gm“3 total gold, Au(III) = 1.2 g nr3 + Au(I) = 50.2 g m“3 , in 3.9 kmol N a d +0.1 kmol HC1 m“3 , sweep rate = 1 mV s“l at 295 K.
6.16 Cyclic voltammogram of a Pt disc electrode in a quiescent electrolyte containing 64.7 gm-3 total gold, Au(III) = 16.7 gm-3 + Au(I) = 48.0 g m~3 , in 0.99 kmol N a d + 0.01 kmol H d nr3 , sweep rate = 10 mV s-1 at 295 K.
6.17 Steady-state reduction current density vs square root rotation rate at constant potential and theoretical currents in the same electrolyte as Figure 6.16.
6.18 UV absorption spectra of solution containing 1 kmol H d m-3( --- ) total Au = 16.6 g m“3, Au(III) = 12.6 g m“3 + Au(I)= 4.0 g m-3, ( - ) total Au = 14.7 g m“3 , Au(III) = 0.7 g m -3+ Au(I) = 14 g m“3.
6.19 Effect of time on the Au(I)/Au(III) ratio of unstirred solutions under nitrogen atomsphere, (0) PTFE, (A) PLASTIC and (□) GLASS beakers at 298 K.
6.20 Steady-state current vs potential curve for a gold electrode, moderate stirring, (□) 1 kmol H d m-3, (0) 0.9 kmol H d +0.1 kmol N a d m-3.
6.21 Partial currents vs potential curve for a gold electrode, moderate stirring, (□) 1 kmol H d m-3, (0) 0.9 kmol H d +0.1 kmol N a d M“3.
6.22 Effect of d “ concentration on the generation of Au(I)/Au(III) species at a potential = 0.8 V vs SCE and pH = 1.
6.23 Effect of H d concentration on the generation of Au(I)/Au(III) species at a potential = 0.8 V vs SCE.
6.24 Open circuit potential vs normalized chlorine concentration,(0 ) l kmol H d nr3 , (□ ) 0.5 kmol H d nr3 .
6.25 Effect of d ” ion and H d concentration on corrosion potentials and the mixed potential leaching rate of a gold- plated Pt rotating disc electrode, at 4 Hz rotation rate,2 mol 0 - 2 m”3 at 295 K.
6.26 Activity - pH diagram of chlorine/hypochlorite species at 298 K and total concentration of 2 mol m -3 in 1 kmol d - nr3.
10 .
Figure 6.27 Effect of chlorine concentration on the mixed potentialleaching rate of a gold-plated Pt rotating disc electrode at 295 K, 4 Hz rotation rate in 1 kmol HC1 m"3 (0).
Figure 6.28 Effect of rotation rate and pH on the mixed potential leaching rate of a gold-plated rotating disc electrode at 295 K, constant chlorine concentration (2 mol m”3) in 1 kmol Cl”, (0) pH = 0, (A) pH = 2, (t) pH = 4.
_oFigure 6.29 Reduction currents for 2 mol Clg m ° at different rotationrates on a gold-plated Pt rotating disc electrode and oxidation currents of Au in 1 kmol HC1 m”3 at 295 K obtained using the net current method, sweep rate = 10 mV s“l.
Figure 6.30 Reduction currents for total chlorine species of 2 mol m“3at different rotation rates on a gold-plate Pt rotating disc electrode and oxidation current of gold in 1 kmol Cl” + 0.9999 kmol Na+ +0.1 mol H+ m”3 at 295 K obtained using the net current method, sweep rate = 10 mV s”l.
Figure 6.31 Experimental reduction currents for total chlorine species of 2 mol m-3 at a potential of 0.2 V vs SCE (□ ), 0.5 V vs SCE ( 0) .
Figure 6.32 Effect of rotation rate and pH on leaching rates obtained by the net current method, total chlorine species concentration = 2 mol m“3 in 1 kmol Cl” m”3 at 295 K, (0) pH = 0, (A) pH = 2, and (□ ) pH = 4.
Figure 6.33 True kinetic current at infinite mass transport for mixed potential leaching rates data (0 ) and net current method (□ ), total chlorine species concentration of 2 mol m”3 in 1 kmol Cl” m”3 at 295 K.
Figure 6.34 Effect of rotation rate on leaching rates obtained by the net current method, 10 mol HC10 + 1 kmol Cl” + 5 mol C00H.C6H4C00K + 3.4 mol NaOH m"3 at 295 K.
Figure 6.35 Depletion of total dissolved gold by adsorption on flow circuit surfaces (A), and unused carbon bed particles (#) of area 0.082 m2 .
Figure 6.36 Total dissolved gold concentrations (open symbols) and current densities (solid symbols) as functions of time and flow rate. Feeder electrode-membrane potential 0.53 V, bed of 3 mm cylindrically-shaped carbon chips, flow rates ( o ,♦) 1.9 x10”6 m3 s”1, ( 0 , ■ ) 8.4 x 10”3 m3 s”1, (O ,• ) 16.2 x 10”6m3 s“l.
Figure 6.37 Total dissolved gold concentration dependence of the PBEcross-sectional current density. Feeder electrode-membrane potential 0.53 V, bed of 3 mm cylindrically-shaped carbon chips, flow rate 1.9 x 10”3 m3 s”l, (O ), 8.4 x 10”3 m3 s”1 ( □), 16.2 x 10”3 m3 s”1 (• ).
11.
Figure 6.38 Total dissolved gold concentration, [Au(I)]/[Au(III)]molar ratio, and incremental, and cumulative, Faradays per mole of gold deposited in the PBE operating under conditions specified in Figure 6.36.
Figure 6.39 Total dissolved gold concentration, [Au(I)]/[Au (III)]molar ratio, and incremental, and cumulative, Faradays per mole of gold deposited in the PBE operating under conditions specified in Figure 6.36.
Figure 6.40 Total dissolved gold concentration, [Au(I)]/[Au(III)]molar ratio, and incremental, and cumulative, Faradays per mole of gold deposited in the PBE operating under conditions specified in Figure 6.36.
Figure 6.41 Total dissolved gold concentrations and current densities as functions of time and flow rate. Feeder electrode-membrane potential 0.53 V, bed of -200 urn carbon chips, flow-rates of 1.9 x 10"® nr* s“l, 8.4 x 10"® m® s"l, and 16.2 x 10”® m^ s”l.
Figure 6.42 Gold concentration decay rate for bed of 3 mn cylindrically- shaped carbon chips and bed of -200 ym carbon chips, initial total dissolved gold = 100 g m"® feeder electrodemembrane potential 0.53 V, 1 kmol H d m"®.
Figure 6.43 Total dissolved gold concentrations and current densities as functions of time and flow-rrate.Feeder electrode-membrane potential 0.53 V, bed of -200 urn carbon chips, flow-rates of 1.9 x 10"® m3 s"l,8.4 x 10”® m3 s”l, and 16.2 x 10~® m”3 s"l.
Figure 6.44 Gold concentration decay rate for bed of -200 ym carbon chips, initial total dissolved gold concentration = 15 g m"3, feeder electrode-membrane potential 0.53 V, 1 kmol HC1 m”3.
Figure 6.45 Total dissolved gold concentration and current densities as functions of time and initial dissolved gold concentration. Feeder electrode-membrane potential 0.53 V, bed of -200 ym carbon chips, flow rate 16.2 x 10“® m3 s'"l, intial dissolved gold concentrations 100 g m“3, 75 g m”3 15 g m"® in l kmol H d m“3.
Figure 6.46 Gold concentration decay rate as a function of pH, feederelectrode-membrane potential 0.229 V, initial dissolved gold concentration = 100 g m"3, flow-rate = 16.2 x 10"® m3 s"*, bed of -200 ym carbon chips, in 1 kmol Cl" m"3. Dashed line indicates possible solubility problems.
Figure 6.47 Idealized behaviour for couple chlorine leaching and electrowinning of gold at steady-state.
12.
Figure 6.48 Total dissolved gold concentrations as a function time at inlet to leach reactor, outlet of leach reactor, outlet of cathode compartment. Leaching bed of gold-plated 3 mn cylindrically-shaped carbon chips, packed bed of - 200 ym carbon chips, feeder electrode-membrane potential 0.53 V, flow rate = 16.2 x 10”° m° s , 0.1 mol CL9 + 1 kmol H Q nr3.
Figure 6.49 Au(I) : Au(III) molar ratios at inlet leach reactor, outletof leach reactor, outlet of cathode compartment, total Faradays, Fardays per mole of gold deposited in the PBE operating under conditions specified in Figure 6.48.
Figure 6.50 Membrane current density as a function of time, operating under conditions specified in Figure 6.48.
Figure 6.51 Total dissolved chlorine concentration as a function of time at outlet of anode compartment, inlet to leach reactor, outlet to leach reactor, outlet to cathode compartment, at pH = 2.
Figure 6.52 Total dissolved gold concentrations as a function of time at inlet to leach reactor, outlet of leach reactor, outlet cathode compartment, operating at pH = 2.
Figure 6.53 Total dissolved gold concentrations at inlet to leach reactor, outlet of leach reactor, outlet of cathode compartmment.
Figure 6.54 Potential - pH diagram for the Fe-S-H20 system at 298 K, with dissolved iron and sulphur activities of 10-4, considering Fe00H(s) as solid Fe(III) oxide phase.
Figure 6.55 Activity - pH diagram for the Fe(III)/H20 system at 298 K, considering Fe00H(s).
Figure 6.56 Activity - pH diagram for the Fe(III)/H20 system at 298 K, considering Fe203(s).
13.
LIST OF TABLES
Table 2.1
Table 2.2
Table 4.1
Table 4.2
Table 4.3
Table 4.4
Table 4.5
Table 6.1
Table 6.2
Table 6.3
Table 6.4
Stability Constant for some Selected Metal-Cyanide Complexes.
Summary of Acid Thiourea Leaching done on Gold Ores.
Pattern of Electronegativities for Possible Au Ligands.
Overall Stability Constants for a Selection of Complexes of Au(I) and Au(III).
Gold Alloys.
Typical Composition of Gold Slimes used by Walker.
Logarithmic Solubility Constants for AgCl(n~l) Species where n = 1, 2, 3, 4.
Selected Free Energy of Formation Data for Species in the Ag/H20-Cl System at 298 K.
Selected Free Energy of Formation Data for Species in the AU/H2O-CI System at 298 K.
Anodic Tafel Slopes for AU-CI-H2O System at Different pHs.
Anodic Tafel Slopes for Au(I) and Au(III) at Different pHs.
14.
NOMENCLATURE
Subscripts
a anodic+ anodicc cathodic
cathodiceq equilibrium0 pertaining to species 0 in 0 + ne” = RR pertaining to species R in 0 + ne" = R
Roman Symbols
Symbol Meaning Dimensions
A (a) Area m^(b) Electrode area(c) Absorbance
cm^
AUp Total dissolved gold concentration g m °
a Specific surface area of bed electrode nri
cj Concentration of species j kmol m""2 , Mc* Bulk concentration of species j kmol m-2
ci Inlet concentration mol m
co Outlet concentration Qmol m
d j Diffusion coefficient of species j m2 s-1
dP Equivalent diameter particle m£ Bed voidageE Electrode potential, all quoted values
versus SHE V
Ea Activation energy kJ mol-^
Eh. Potential versus SHE VE° Standard potential of an electrode or
half-reaction VF Faraday 98 485 C (mol e")"1
f F/RT v-iG Gibbs free energy kJ mol-l
15
AG° Standard Gibbs free energy change kj mol-1
i Current density A m“2} niA m“2
Anodic component current density A m”2, mA m~2
xc Cathodic component current density A m“2 , mA m-2
n Limiting current A, mA
k Mass transfer rate constant m s ~ l
L Electrode length m
n Number of electrons per species oxidised or reduced
Q Charge passed in electrolysis C
Q Volumetric flow rate m3 s“l
R Gas constant (= 8.314) J mol"1 K"1
Re Reynolds number = udp/v
Sc Schmidt number = v/D
Sh Sherwood number = KL/D
t Time sT Residence time (= V/Q) s
T Absolute temperature KV Kinematic Viscosity m^ s-1
V Reservoir volume m3
z Number of F mol-1
Greek Symbolsoc Transfer coefficient
3 Cumulative equilibrium constant0 Coveragee Molar adsorption coefficient mol-1 m2
n Overpotential, E - Egq VOJ Rotation frequency HzX Wavelength nm
16.
Standard Abbreviations
AAS Atomic absorption spectrophotometry
CV Cyclic voltamnetry
ECE Heterogeneous electron transfer, homogeneous chemicalreaction, and heterogeneous electron transfer, in sequence
RDE Rotating disc electrode
RRDE Rotating ring disc electrode
SCE Saturated calomel electrode
SHE Standard hydrogen electrode
17 .
CHAPTER ONE
INTRODUCTION
18.
CHAPTER 1 - INTRODUCTION
Conventional leaching of gold ores by alkaline cyanide solutions
suffers from rate limitations resulting from the restricted solubility of
dissolved oxygen, the transport controlled reduction of which determines
the maximum leaching rate in the presence of an adequate cyanide concentration^ >2 >3) . The use of chlorine instead of oxygen as the oxidant,
and its precursor, chloride instead of cyanide as the complexant, has
been reported(4) as enabling gold to be leached at rates more than two
orders of magnitude greater than attainable with the now traditional
chemistry, which in fact superseded it as the industrial leaching process
at the end of the nineteenth century(5).
In the envisaged gold leaching and recovery process reported here, chlorine evolved at the anode of a cation-exchange membrane divided cell,
chlorine would be reacted with gold ore (or electronic scrap etc.) in an
external reactor in which gold dissolution would occur at a mixed pot
ential, and the solubilized gold would then be electrowon from dilute
solution using a three dimensional (packed or fluidised bed) cathode
in the same electrochemical cell to achieve high mass transport rates,
cross sectional current densities and space-time yields(®). While chlo
rine/hypochlorite has been used as the oxidant in this work, in principle
any other soluble oxidant which could be electro-(regenerated, and with
an adequately high reversible potential (> EAuCln/Au^ could be used, with chloride as the complexant.
The advantages of such a leaching-recovery system would be:
i) Faster leach kinetics, and therefore decreased capital costs.
19.
ii) Energy efficient oxidant generation and soluble gold complex recovery.
iii) The possibility of decreasing the N a d consumption by recycling the chloride ion.
iv) Sodium chloride is a chemical which is easy to transport, store and is not toxic.
v) Despite chlorine being a toxic and corrosive reagent, the integrated nature of the process would minimize safety risks and allow the isolation of high risk areas so that stringent safety measures may be applied.
vi) The process waste would be more environmentally acceptable than ' the cyanide waste.
vii) Certain gold ores (e.g. carbonaceous ores) require chlorination presently, prior to conventional cyanidation.
The major disadvantages of chlorination are:
i) The dissolution of base metals and sulphides.
ii) The presence of silver in certain gold ores may passivate the dissolution reaction.
The objectives of the presently reported work were:
(a) To determine the kinetics and mechanism of gold dissolution by the reduction of chlorine.
(b) To determine the kinetics and mechanism of deposition of gold-chloro complexes in a three dimensional electrode.
(c) To test the feasibility of coupled chlorine leaching and electrowinning of gold.
20 .
CHAPTER TWO
ALTERNATIVE LEACHING SYSTEMS TO CHLORINATION
21.
CHAPTER 2 - ALTERNATIVE LEACHING SYSTEMS TO CHLORINATION
The aim of this chapter is to provide a framework for comparison
between the envisaged process and present day practice in the mineral
industry for the extraction of gold. The types of gold ores available
as raw material, a generalized flowsheet for gold extraction and a review
of the chemistry of cyanidation and thiourea leaching systems are des
cribed.
2.1 TYPES OF GOLD ORES
A useful classification of gold ore types, listed below, has
been published by Mcquinston and Shoemaker(?):
1. Free gold ores.
2. Gold with iron sulphide.
3. Gold with arsenic and/or antimony minerals.
4. Gold tellurides.
5. Gold with copper porphyries.
6 . Gold with lead and zinc minerals.
7. Carbonaceous ores.
2.1.1 Free Gold Ores
These are ores in which the gold is in elemental state and
not locked in other su]Dhide minerals. Part of the gold may be
sufficiently coarse to allow the application of gravity separation.
Cyanidation is normally used in conjunction with gravity separation
or, if the precious metal values are particularly fine, as the
only process of extraction. The gravity concentrate (coarse gold) is either amalgamated, or intensive cyanidation is used.
22.
Intensive Cyanidation uses pure oxygen rather than air, a very
high concentration of cyanide (30 % mass), and a temperature of
around 3 0 °c (8 ) . Sulphide content, caimonly pyrite, is generally
low, usually less than 2 per cent by mass. Two typical examples
of operating plants treating these type of ores are Dome Mines
Ltd., South Porcupine, Ontario, Canada(^) and Ashanti Goldfields
Corporation (Ghana) Ltd., Ghana(9).
2.1.2 Gold with Iron SulphidesGold occurs external to, and disseminated within, sulphide
minerals. Some sulphides, particularly pyrrhotite, tend to decom
pose in solution, consuming cyanide by forming thiocyanate ions
and inhibiting precious metal extraction. Aeration with lime
prior to cyanidation is often practised. Auriferous pyrite
flotation concentrates are often produced which may be directly
smelted if of sufficiently high grade. However they are normally
reground, then subjected to cyanidation or calcination followed
by cyanidation. Parmour Porcupine Mines Ltd., Ontario, Canada(^)
operate this type of process. The generalized flowsheet (Figure
2.1) shows the different possibilities available.
2.1.3 Gold with Arsenic and/or Antimony MineralsThe presence of As oi 3b minerals usually make the ore
refractory to direct amalgamation or cyanidation. These ores are usually regarded as complex sulphides, in which gold tends
to occur as sub-microscopic inclusions or in solid solution in
23.
Ore
\
Figure 2.1 Generalized cyanidation flowsheet.
24 .
the gold-bearing minerals. Therefore the aim is to expose the
gold to the dilute alkaline cyanide solution. In a conventional
plant this is achieved by roasting followed by cyanidation.
Flotation is usually performed before roasting in order to
minimise the capital investment and sulphur dioxide emissions. An
example of this type of flowsheet is operated at Giant Yellowknife
Mines Ltd., Northwest Territories, Canada^).
2.1.4 Gold Tellurides
These are the only gold minerals, other than metallic native
gold, that are of economic significance. They usually occur with
native gold and with sulphides. The ore or flotation concentrate
normally requires some form of oxidation (by roasting or chemical
oxidation) prior to cyanidation. The most important host mineral
for gold and silver *are the tellurides sylvanite (AuAgTe^.) and hessite (Ag2T e ) ( ^ . An example of this type of flowsheet is oper
ated at Bnperor Gold Mines, Vatukoula, Fiji(7).
2.1.5 Gold with Copper Porphyries
Substantial quantities of gold are recovered in copper
sulphide concentrates, produced from exploitation of porphyry copper
deposits. Any associated gold follows the copper through smelting and is recovered during electrolytic refining.
The Magna Copper Company, San Manuel Division, provides
an interesting case example of this type. As with many copper ores, the San Manuel ore contains a small amount of gold, but in
this particular case much of the gold is associated with the
25 .
molybdenite (M0S2). Separation of the molybdenite from the
bulk sulphide concentrate by flotation produces a molybdenite
concentrate which is leached by the normal alkaline cyanide
procedure.
2.1.6 Gold with Lead and Zinc Minerals
As with the case of gold with copper porphyries, the gold
is usually recovered in a sulphide flotation concentrate. This
concentrate is then treated by the conventional cyanidation
procedure.
2.1.7 Carbonaceous Ores
These ores contain some form of carbon which adsorbs
dissolved gold during leaching, resulting in gold loss to the
tailings. Such ores also often contain sulphides. Treatment
requires oxidation of the ore prior to cyanidation, although in
some instances it is possible to separate the carbonaceous
material by froth flotation and in others to mask the effect of
the carbon with light fuel oil. Recently chlorine has been used to oxidize the carbon at Carlin, Nevada^!»12) # a typical flow
sheet of this type of ore is provided by Kerr Addison Mines Ltd.,
Virginiatown, Ontario, Canada^).
2.2 GOLD LEACHING IN OXYGENATED ALKALINE CYANIDE SOLUTIONIn this section the thermodynamics, mechanisms and kinetics
of the cyanidation process will be reviewed. Several comprehensive reviews^’1^,14,15,16,17) have been published.
26.
The treatment of gold ores is based largely on the very high
specific gravities of native gold (Au) and gold-silver tellurides
(Au - AgTex) their associated gangue and their solubility in oxygenated dilute alkaline cyanide solutions(lO). Since the
introduction of cyanidation at the beginning of the 20th century
by McArthur and the Forrest brothers in South Africa(^), cyanidation
has become the main chemical process for gold extraction from its
ore.
The advantages of cyanidation are:
1. Cyanide is relatively selective of gold.
2. It is an established technology that has been improved quite
dramatically throughout the years in operation, e.g. carbon-
in-pulp, carbon in leach, etc.
3. Silver is readily dissolved.
The disadvantages are:
1. Slow leach kinetics.
2. Cyanide toxicity.
3. The increase in capital expenditure when gold is found with
certain minerals like arsenopyrite, stibinite, tellurides, carbon, etc.
2.2.1 Thermodynamics of the Ag/Au - GN - H^O System
The thermodynamics of the system in the form of Eh-pH diagrams
and activity-pH diagrams has been discussed recently by Osseo-
27.
Asare et al(18,19,20)# Their results are very similar to thoseof Finkelstein(13).
From the Eh-pH diagram for the AU-H2O system (Figure 2.2)
derived by Pourbaix(21) ? it can be seen that in the absence
of complexants elemental gold is thermodynamically stable even at
potentials slightly in excess of the reversible oxygen potential.
At even higher potentials gold passivates by oxide formation,
except at the far extremities of the pH range. A suitable oxidant-
complexant is necessary for a successful leaching system. Gold
forms the most stable complexes with cyanide, the cumulative
stability constant for AuCNTj is 7.7 x 10**®.
The Eh-pH diagrams for the Au/Ag -CN- are
shown in Figures 2.3 and 2.4 from which it is clear that:
1. Au (CN)2 ions predominate over a wide potential region at
all pH's, the reversible Au(CN)2/Au potential lying well
below that for the reduction of oxygen at all pH's.
2. Ag(CN)2 ions predominate over a wide potential region at
pH's above 4 explaining the ability of cyanidation to cope with the elemental silver content of gold ores.
2.2.2 Mechanism of Cyanidation
The electrochemical nature of gold leaching in oxygenated
alkaline cyanide solutions, i.e. that O2 reduction on the surface
drives the anodic gold oxidative dissolution has been accepted for more than 30 years(l).
28.
-1,2
-0.2
-0,4
-0,6
-0,8
-I
0,6
-I,~
1,2
0,8
-1,6
Au-~---------------
---------
2 3 "" 5 6 7 8 9 10 II 12 13 I~ 1~ 15o';...~~;"""'''';''''''':;''~:'''''''T--:;:''''--T~T--T--T-----:'':'-T-''';::--TT-r-''''''';;''-, 2,B
AuOZ 2,6
2,4
-I
-1,2
-1,04
-1,6
-1,8-2 -1 0
Figure 2.2 Eh-pH diagram for the AU/H20 at 298 K(21).
29 .
Figure 2.3 Eh-pH diagram for the Au/I^O-CN system at 298
Figure 2.4 Eh-pH diagram for the Ag/I^O-CN system at 298 K^1®)
3 0.
The accepted overall chemical reaction is:
4Au + 8CN” + 02 + 2H20 4Au (CN)2 + 40H" [2.l](Eisner's Eq.)
Bodlander(22) proposed that the dissolution occurs through two
steps:
2Au + 4 Q T + 02 + 2H20 > 2Au(CN) 2 + 20H“ + H202 [2.2]
2Au + 4CN“ + H202 ■> 2Au (CN) 2 + 20H" [2.3]
which add to give Eisner's equation [2.1].
Similar equations may be written for silver(23,24)? which
has been less studied than gold.
However, the presence of other ions in solution can have a
dramatic effect on the mechanism and kinetics of both reactions.
The addition of trace quantities of heavy metal ions such as
Pb(II), Tl(I), Bi(III), and Hg(II) have been found(25) to en
hance gold electrodeposition rates in cyanide baths, and being
bifunctional catalysts, they also accelerate the dissolution of
gold. These catalysts deposit uniformly on gold surfaces to form an adsorbed monolayer at electrode potentials positive of
those at which bulk metal deposition begins, i.e. at "underpotentials". Two mechanisms have been proposed(25) to explain the
effect:
1. The enhancement of gold nucleation rate by strongly adsorbed foreign metal atoms and;
3 1.
2. The occurrence of electrochemical displacement reactions,
made possible by the specific adsorption of heavy metal
depolarizer ions.
The effect of Pb(II) on gold has been reported previously by HadenC^S)^ Kameda(27,28)^ pink and Putnam(29).
Engelsmann et al(30>31) used galvanostatic plus A. C. impedance
measurements to study underpotential deposition of lead on gold,
which they postulated to occur by a homogeneous (i.e.non-nucleative)
Mez+ transfer/surface transport mechanism. In the early stages
of polarization (t < 100 us), Pb(II) adsorption proceeded by pre
ferential, fast transfer at linear discontinuities in the gold sur
face, in addition to slow transfer at homogeneous (two-dimensional)
surface regions. Evolving surface excess gradients are levelled
out by surface diffusion at t » 100 ys. A simple RC model was
followed.
Other ions such as Fe(II), Cu(II), Zn(II), Mn(II), Ca(II)
and Ba(II) have been reported to have a retarding effect on gold
electrodeposition or dissolution. These effects have been
attributed to one or a combination of the following:
1. Cbnsumption of oxygen from solution, depriving gold of
its oxidant.
2. Cbnsumption of free cyanide from solution, generating lack
of suitable complexant.
3. Film formation on the surface of the metal.
32 .
2.2.2.1 Oxygen Reduction on Gold
The reduction of oxygen provides the driving force for gold
dissolution in cyanidation. There has been debate whether the
reduction occurs by a 4 electron reaction either direct or by a
sequential mechanism with H2O2 species as intermediates. A recent review on oxygen reduction was published by Schiffrin(32)# a
diagnostic criterion was developed by Wroblowa et al(33) to
distinguish between the direct and sequential mechanism using a
rotating ring-disc electrode (RRDE). It was found(33) that the
sequential mechanism operates even at high overpotential in 2 M
KOH.
Figure 2.5 shows the general scheme for oxygen reduction
which does not consider all the possible permutations. At lower
potentials, 0.3 - 0.7 V, oxygen was found to be reduced to H2O2 ,
which did not undergo further reactions. These findings agree with the results of other workers(36,34,35)#
Several oxygen reduction mechanisms have been proposed by Damjanovic et al(36);
O2 + e + 1^0 = O2H + OH [2 .4 ]O2H + e O2H [2 .5 ]02H“ + H20 = H202 + 0H“ [2 .6 ]
I <M0II10+<MO [2 .7 ]O2 + 6 + H2O + O2H + OH [2 .8 ]02H" + H20 = H202 + OH- [2 .9 ]
33 .
k-[ (+ 4 e ~ )
k ^ ( + 2 e ; )
b = b u l k
* = v i c i n i t y o f t h e d i s c e l e c t r o d e
a = a d s o r b e d s p e c i e s
(33).Figure 2.5 General scheme of oxygen reduction
34 .
In either reaction path, the second step was rate controlling. Zurilla et al(34) suggested the following mechanism:
O2 + e" O2 (ads) [2 .10]
205 (ads) + H20 HOg + 02 + OH" [2.1l]
The first step [2.10] being rate determining. The further re
duction of HO2 is a chemical step with no or very little potential dependence. Fischer and Heitbauna(37) have proposed a similar
mechanism to Zurilla et al(34).
Oxygen reduction on gold is probably more complex than either
of the mechanisms suggested. Pillai and Bockris^5) tried to model 17 different proposed pathways for oxygen reduction, none of which
fitted their experimental data.
The effect of lead on the oxygen reduction on gold has been studied by several workers. Strand(39), USing DC polarography,
found that the presence of small quantities of Fb(II) catalyzed
the oxygen reduction, though the effect disappeared with increasing
Pb(II) concentration. He proposed that Pb(II) was oxidized to
Pb(IV) and formed an intermediate PbC>2 which was reduced back
to Pb(II). However this latter mechanism is not thermodynamically
possible. Pieterse et al(4^), instead proposed that the activation energy of the reduction of H202 to OH" ions is lowered by an
amount proportional to the surface excess of the Pb(II) species.
Azidic et t using ring-disc and in-situ specular reflection,
35 .
determined that on an Au(OH) surface with Pb ad-atoms, O2 is
reduced to 0H~ and that very few HO^ ions leave the surface
of the Au disc. For Au surfaces the underpotential deposition of
lead was responsible for the catalytic effect.
2.2.2.2 The Anodic Reaction of Gold in Alkaline MediaNicol(^) hag reviewed the anodic behaviour of gold in
alkaline solutions. It was pointed out that any mechanism pro
posed to explain passivation during anodic dissolution should
also explain the inhibition of the corresponding cathodic process,
because gold dissolution and its deposition from aurocyanide
ions are both subject to passivation at potentials close to
the equilibrium potential.
The existence of three peaks in the anodic reactions of gold in alkaline cyanide solutions has been established(42,43,
44,45,46,47):
Peak I : -0.56 to -0.26 V vs SCE
Peak II : -0.26 to +0.39 V vs SCE
Peak III : +0.39 to +0.54 V vs SCE
Kudryk and Kellog(l) did not observe any peak, but this was
probably due to a 0.5% addition of KC1 to increase the ionic
conductivity. However gold chloro complexes could also form. McArthur(47) found only two peaks at 0.0 and at 0.4 V vs SCE
at 61°C.
36.
Several mechanisms have been proposed by different workers:
Kirk et al(44»45»46), using potentiostatic and potentiodynamic
techniques in the temperature region of 0 - 50°C, proposed the
following mechanism:
Au + CN“ = AuCN- (ads) [2.12]
AuCN-(ads) = AuCN(ads) + e [2.13]
Au(CN)(ads) + O T = Au (CN) 2 [2.14]
Coulometric measurements were performed for each peak and
n, the number of electrons involved in the reaction studied, was
equal to 1. The apparent requirement of 1.05 F (mol AuCCN)^)"”1
for peak III was interpreted as due to the onset of oxide formation.
Additionally, they found an apparent activation energy of 93± 8
kj mol-1 for peak I and a lack of dependence on mass transfer; the
second step reaction [2.13] was considered to be rate-determining.
The activation energy associated with the process giving rise to
peak II was 47-55 kj mol-1, and a small dependence on the mass
transfer; again the second step reaction [2.13] was considered
rate-determining. The apparent activation energy for the peak
III process was found to be 16-18 kj mol-1, and the reaction was diffusion controlled, the third step [2.14] being ratecontrolling. Finally, they observed a fourth peak in the region
-0.81 to -0.56 V vs SGE with an apparent activation energy of 62 kJ mol-1 with little dependence on mass transfer. This was
considered to be an unknown adsorption reaction.
37-
Eisenmann^S) proposed the same sequence of steps, but
for peak I and II the second step was rate-determining and limited
by diffusion of CN“ to the surface. For peak III he suggested
a change of mechanism to a 3 electron process and oxide formation;
this was later disproved by Kirk et al(^4).
Cathro and Koch(42) found three peaks within the potential
regions stated previously as peaks I, II and III and proposed the
following mechanisms:
Au + CN“ Au(CN) (ads) + e [2.15]
Au(CN) (ads) + CN~ -* Au(CN>2 [2.14]
They considered step 1 [2.15] to be the r.d.s. and that
a surface film (basic cyanide) produced passivation at -0.6 V and
+ 0.1 V vs SCE. Auric oxide formation was considered to be re
sponsible for the last peak.
McArthur(47) working at higher temperatures, suggested two
mechanisms:
1. At low overvoltage, the same mechanism as that suggestedby Cathro and Koch (39) ̂ but; with the second step as rate
determining [2.14].
2. At larger overvoltages, two possibilities were suggested:
a) direct oxidation to Au(CN)2 and;b) the oxidation of gold to a gold (III) complex.
Finally, Pan and Wan(4^) found four peaks, slightly offset
to more positive potentials compared to those reported by Thurgood
38.
et al(46). The mechanism suggested involved the competitive
adsorption of OH" and CN" ions on surface site, with the adsorption
of OH" producing passivation. In concentrated cyanide solutions
they proposed the same mechanism as Cathro and Koch(4^) with the
second step as rate determining(4^).
2.2.2.3 Cyanide Consumption
Free cyanide concentration is one of the most important
variables in cyanidation. Cyanide could be in the form of
cyanic acid HCN or cyanide ion CN". Their relative concentrations
depend on pH according to the following equation(50):
HCN = H+ + CN" [2.16]
Ka = 4.93 x 10"10 (pKa = 9.3)
The cyanic acid can react with oxygen to generate cyanate by:
HCN + 1/2 02 HCNO [2.17]
The cyanate can be hydrolysed to carbonate and ammonia at a pH
up to 8.5, being greatly accelerated at lower pHs(51).
HCNO + H20 -► NH3 + 002 [2.18]
Cyanide can react with hydrogen peroxide to form cyanate:
CN" + H202 -► CN0~ + H20 [2.19]
Cyanide can form complexes with 28 elements with the possibility of forming some 72 metal-cyanide complexes(52).
Polysulphide and thiosulfate are oxidation products of sulphides which could be abundant in certain gold ores. These
39 .
could react with cyanide producing thiocyanates according to(53):
s|_ + c n ~ = s(x_X) 2“ + s o r [2.20]
s 20§“ + CN“ = S0§~ + SCN- [2.21]
Iron, copper and zinc among others form complexes with cyanide.
Table 2.1 shows some selected stability constants of relevant
metal-cyanide complexes.
TABLE 2.1 STABILITY CONSTANTS OF SOME
SELECTED METAL-CYANIDE COMPLEXES COMPLEX_________________________ Log 3
Zn (CN)f~ aq. 19.96Cu (CN) 2 aq. 14.06
Cu (CN)|“ aq. 25.80Cu (CN)|“ aq. 26.15Cu (SCN)+ aq. 2.34
Ag CN aq. -1.10
Ag (CN) 2 aq. 18.14
Ag (CN)2- aq. 19.01
Ag (CNS) 2 aq. 8.29
Ag (CNS)2- aq. 9.50
Ag (CNS)2- aq. • 9.75Au (CN) 2 aq. 38.89Au (SCN) 2 aq. 13.20Au (SCN)^ aq. 43.65Fe (CN)|“ aq. 43.79Fe (CN)|“ aq. 36.76Fe (SCN)2+ aq. 0.85
Source : Bard, Parsons and J o r d a n (54)
40
The above possible reactions explain partly why in cyanida-
tion practice the amount of cyanide added is about an order of magnitude higher than the optimum cyanide concentration determined
by Habashi (3). Assuming mass transport control some of the
above reactions are kinetically slow and do not consume much
cyanide, other reactions require the ore to be pre-treated as
explained in section 2 .1.
2.2.3 Kinetics of Cyanidation
Cyanidation has several drawbacks, the slow kinetics being
the most serious. If cyanidation is diffusion-controlled, then
for an adequate cyanide concentration the maximum rate of dis
solution is limited by the maximum solubility of oxygen, the solubility of which in aqueous KOH of concentration (C) 0 to 12 mol dm”3 at 25°C and 1 Atm, is given by(55):
iog [02]/mol dm-3 = log 1.26 x 10-3 - 0.1746 C [2.22]
where 0.1746 is the solubility coefficient.
The rate of gold dissolution under well defined hydro-
dynamic conditions has not been extensively reported in the literature. Kakovskii and Kholmanskikh(56)t using a rotating
disc electrode at constant 02 concentration of 1.26 x 10“3 M in
10“3 m KOH found the dissolution rate to increase steeply with
increased rotation rate up to 3 Hz. From 3 to 6 Hz the rate
decreased quite dramatically and above 6 Hz the rate decreased
further at a slower pace. Increased cyanide concentration shifted the curve upwards but basically the same behaviour was observed.
41.
At [KCN] = 1 x 10”2 M the maximum dissolution rate was 8 x 10-5 toqi m-2 S-1 an(̂ 2% [k c n ] =0.3 x 10“^ M, 4 x 10“5 mol
m~2 s~l.
Similar behaviour has been found by Chthro(57)> though
he was using a gold disc with forced convection provided by
a propeller at the bottom of the cell, and by Cheh(58) the
O2 concentration was not measured.
These observations can be explained by either:
1. Assuming that some or all H2O2 escapes from the gold surface,
then as the transport of oxygen increases, the local con
centration of hydogen peroxide increases and this could:
a) accumulate in the solution
b) be reduced on the gold surface and,
c) oxidize the cyanide to cyanate.
Condition (a) is very unlikely to occur and if (b) occurs
then O2 would be regenerated and further reduced, requiring
the transport of more cyanide ions to the interface in
order to complex the gold. If the transport rate of cyanide
ions is inadequate then passivation would occur. If (c) occurs, then obviously cyanide ions will not be available
for complexation of the gold and passivation again might
occur.
2. Assuming that oxygen is reduced to 0H“ . This would require
the transport of 8 CN“ ions to the gold interface and the local pH increase would favour passivation.
42.
By either of mechanisms 1 or 2 above, the surface film produced
should be porous to explain the increased dissolution rate in
creasing with OST concentration.
Habashi(2,3)? assuming diffusion-controlled, has postulated
the following theoretical formula for the rate of dissolution of
gold in oxygenated alkaline cyanide:
2 A DcfT °02 [°2] [2-23]Rate = •— -------------------------6 {Dq j- [ o r ] + 4D02 [02 ]}
where A = the surface area of the metal in contact with the
aqueous phase (cm^).
Dq -̂ and Dq = the diffusion coefficients of cyanide and dis- 2
solved oxygen; 1.83 x 10“5 and 2.76 x 10“5 cm^
s”l, respectively.
[CN“] and [c^] = the concentration (in mol cm"1) of cyanide and
dissolved oxygen, respectively.
6 = the thickness of the boundary layer, which varies
between 2 and 9 x 10”3 cm, depending on the
speed and method of agitation.
The optimum [CN“] / [O2 ] ratio is 6 which gives an optimum concentration of cyanide equal to 7.53 x 10“3 M at 1 atm of
pressure.
4 3 .
2.2.4 Future Research
It seems that the future research areas will be:
1. Pressure oxidation of gold using pure oxygen.
2. Decreased capital expenditure by improving heap-leaching.
3. Recovery of Au(CN)2 by carbon-in-leach or resin technology.
2.3 ACID THIOUREA LEACHING OF GOLD
Acid thiourea leaching exploits the fact that gold forms
a very stable complex with thiourea, the driving force for the gold oxidation being provided by the reduction of ferric ion or
formamidine disulphide. The advantages of this system are(59);
1. Leach kinetics are an order of magnitude faster than
cyanidation.
2. Thiourea is considered non-toxic.
3. It is less affected by elements such as Cu, Zn, As, Sb, C, etc. than cyanidation.
The disadvantages are:
1. High cost of thiourea.
2. High consumption of acid and thiourea.
The reaction describing gold dissolution in acidic thiourea solution is(®9,61,62,63 y
Au (CS(NH2 ) 2 )2 + e" = Au + 2CS(NH2 ) 2 E° = 0.38V [2.24]
Electrochemical studies^l) using a rotating gold disc electrode,
showed that in 0.001 to 0.03 M thiourea and 0.1 M sulphuric acid,
gold dissolution nearly reached the diffusion controlled rate calculated from the Levich equation(64) up to an overpotential
of 0.3 V. The exchange current density was greater than 10“^ A m”^
and the dissolution was 100% current efficient. At higher
overpotentials thiourea was oxidized to formamidine disulphide
(RSSR) and other sulphur compounds, the dissolution becoming
inhibited.
The reduction of gold (I) thiourea complex species [2.24]
was diffusion-controlled at overpotentials between -0.15 to
-0.35 V.
Thiourea is oxidized to formamidine disulphide (RSSR) by(®^)•
2CS(NH2 )2 = NH2C(NH)SSC(NH)NH2 + 2H+ + 2e~ [2.25]EP = 0.420 V
which itself is considered a suitable oxidant for gold dissolution (62,63) though the kinetics should be slower than for
ferric ion due to the relatively little driving force for the
dissolution reaction [2.25].
Formamidine disulphide can degrade further in an irrever
sible reaction to thiourea and unidentified sulfinic compounds,
which in a third irreversible step decompose to cyananide and elemental sulphur (65). Thiourea decomposes at temperatures higher
than 35 °C resulting in decreased gold extraction.
The formation of a sulphur surface film has been con
sidered to be the reason for inhibition of the gold dissolution
reaction(63) ancj increased thiourea consumption. Therefore it
would seem that a key factor in thiourea leaching would be to
choose an appropriate oxidant to convert only a portion of
thiourea to RSSR(66) or the addition of a selective reductant of
RSSR in slight deficiency of the molar requirement to oxidize half of the thiourea to RSSR(65).
Schulze(65) has claimed that this procedure decreases re
agent consumption to 0.5 kg thiourea tonne--*- of ore treated.
Several oxidants have been tried(63): formamidine dis-
sulphide, hydrogen and sodium peroxide, oxygen and ferric ions.
Oxygen, hydrogen and sodium peroxide oxidize thiourea to RSSR
so that the leaching reaction with these oxidants is:
2Au + 2CS(NH2 ) 2 + RSSR + 2H+ = 2Au(CS(NH2 ) 2 ) 2
E° = 0.07 V [2.26]
Iron (III) reacts with thiourea very slowly to form
[Fe(III) S04(CS(NH2)2)2]+
which is further reduced to [Fe(II) CS (NH2>2)] SO4 with time.
The leaching reaction when ferric ions are present is described by equation [2.27].
Au + Fe(III) + 2CS(NH2 ) 2 = Au(CS(NH2) 2 )2 + Fe(II)AG = -37.63 kJ mol-1 [2.27]
The dissolution rate is a function of oxidant and complexant concentration and the pH tends to increase with time(59)#
4 6 .
Table 2.2 summarises the applied research done using thiourea.
TABLE 2.2 - SUMMARY OF ACID THIOUREA LEACHING DONE ON ORES
ORE OXIDANT REAGENTCONSUMPTION
RECOVERY REFERENCE
Pure Gold + South African Ore
Fe(III), 02,h 2o 2 , r s s r
1.4 to 0.4 kg/t 96% 62
CarbonaceousOre
Fe(III) 2 kg/t > 80% 59
Sulphide Cone. + Smelter Flue Dust
Fe(III), 02 — 80% 67,68,69
Carbonaceous Clayey Ore
AIR, Fe(III) 1.5 to 2.2 kg/t 95% 70
Complex Au-Ag Cone. (Pb,Zn)
Fe(III), AIR 15-19 kg/t 85% 71
Pyrite Cinders 02, AIR 7-4 kg/t > 94% 72 .
Pynte Cone Fe(III), 02 , AIR
5 kg/t for 02 15 kg/t for Fe(III)
> 95% 73
Broken Porcelain (10-100 g/t Au)
Fe(III), H202 — > 95% 74
Fran Table 2.2 it is clear that most of the applied research has
been done on ores or materials which do not respond well to conventional
cyanidation.
47 .
Pyper and Hendrix(59) found no practical kinetic advantage
when thiourea leaching was applied to carbonaceous ores of the
Carlin, Nevada type when compared with pre-oxidation by CI2
plus cyanidation. Gabra(?3) observed a gold extraction of first
order with respect to thiourea concentration and zero order with
respect to Fe(III) and H2SO4 . This shows that when Fe(III) is
available in the ore, the addition of an oxidant is unnecessary.
In conclusion, of the two systems reviewed, cyanidation
suffers from slow leach kinetics and thiourea leaching from high
reagent cost. Futhermore, thiourea leaching does not have pra- tical kinetic advantage over CI2 pretreatment plus conventional
cyanidation in the treatment of carbonaceous ores.
4 8 .
CHAPTER THREE
REDUCTION OF DISSOLVED AQUEOUS CHLORINE/HYPOCHLORITE
49.
CHAPTER 3 - REDUCTION OF DISSOLVED AQUEOUS CHLORINE/HYPOCHLORITE
Chlorine reduction has been studied on Pt, Graphite, Ti-base/oxide4
Rhodium and iridium electrodes and reviewed by Mussini and FaitaC7^).
No work on the reduction of chlorine or hypochlorous acid on gold elec
trodes was found in the literature as might be expected since gold dis
solves in the presence of O ^ / d - couple unless cathodically protected.
Chlorine reduction on different substrates has been explained by
one of the following three mechanisms:
i) Inverse Volmer [3.l] - Heyrovsky [3.2],
d-ads + e d [3.1]
d-2 + e = C3. + d-ads 1.3.2]
ii) Inverse Volmer [3.1] - Tafel [3.3],
2 dads * d 2 [3.3]
iii) Inverse Krishtalik Mechanism
dads + e d [3.1]
d ads * d acjs + e [3.4]
a 2 * a ads + a " [3.5]
3.1 PLATINUMFrumkin and TedoradseC76), using a rotating Pt disc elec
trode, proposed an inverse Volmer-Heyrovsky mechanism; step [3.2]
was considered rate-determining and the stoichiometric number was
found to be 2. It has been pointed out that a stoichiometric number of 2 would be obtained if the exchange currents of steps
50 .
[3.l] and [3.2] were only slightly different(76,77)# Tafel slopes
of 85 mV/decade and reaction orders of 1 with respect to chlorine
and zero with respect to chloride were found(^). The inverse Volmer-Heyrovsky mechanism has been confirmed by other workers(^8 >
79,80)# a reaction order of 0.6 to 0.7 with respect to chlo
rine has been reported^-*-), but Dickinson et al(^) found a
reaction order of 1 for low chlorine concentrations and 0 .66 at
high chlorine concentrations, with step [3.1] as .rate-determining.
Yokoyama and Enyo(^) ̂ using short-time galvanostatic pulses,
observed an inverse Volmer-Tafel mechanism. The rate determining
step was [3.3] at overpotentials < -120 mV, and step [3.1] at
overpotentials > -120 mV. A tentative model was suggested(^6) in
which a chlorine adsorption step preceeded charge-transfer to ex
plain the lack of linearity in the plots of log current vs potential
obtained by Chang and Wick(83). However, no supporting evidence
was provided.
3.2 GRAPHITE
The reduction of chlorine on graphite was diffusion-controlled (84,85) a-£ overpotentials of - 70 mV. The reaction orders were 1
with respect to chlorine and zero with respect to chloride ion. Hine and Masuda(^4) postulated an inverse Volmer-Heyrovsky mechanism
with [3.2] as rate-determining.
51 .
3.3 TITANIUM-BASE ELECTRODES
This section covers a wide variety of electrodes with electro-
catalyst coatings applied to titanium substrates. Thick catalyst
coatings behave as diodes, i.e. they and therefore CI2 are not reduced, or the kinetics is influenced by the behaviour of the
Ti/MOx interface.
On platinized Ti electrodes, an inverse Volmer-Tafel mechanism
was observed at overpotentials < -0.5 V and an inverse Volmer-
Heyrovsky mechanism at > -0.5 v ( 78) . Some authors^86) observed
that the activities of Ti/Ir02 and Ti/Ru02 electrodes for cathodic
chlorine reduction was considerably less than for anodic chlorine
evolution. Oxygen would be reduced and this might account for
the decrease in activity. Hepel et al(87>88) studied chlorine
reduction on planes (110) and (101) on Ru02 single crystals. A peak was observed at 0.442 V on both planes which was mass-transport
controlled and it was attributed to the Tafel-Volmer or Heyrovsky-
Volmer mechanism on the basis of the assumption that the rearranged
(reduced) surface oxide is not capable of stabilizing highly
oxidized chloronium species in this range of low electrode potentials.
A second peak at 1.022 V is only observed on the (110) plane.
The availability of O2 sites on this plane allowed the chlorine reduction to follow a Krishtalik(244) mechanism according to:
c i2 -► 0-Cl+ + Cl“ [3.6]
0-Cl+ + e- = 0-C1 [3.7]
0-C1 + e- = 0 + Cl“ [3.8]
where 0-C1 represent the chemisorbed chloronium ion or surface 0-C1groups.
52
CHAPTER FOUR
THE CHEMISTRY OF GOLD AND SILVER IN ACIDIC SOLUTIONS
53.
CHAPTER 4 - THE CHEMISTRY OF GOLD AND SILVER IN ACIDIC SOLUTIONS
4.1 GOLD
4.1.1 Introduction
Gold has an atomic weight of 196.967 and a density at 20°C
of 19.32 g cm“3, the pure metal melts at 1063°C and boils at 2966°C.
In common with other group IB elements, gold crystallizes in a
face-centred cubic lattice and the closest inter-nuclear distance
is 288.4 pm.
Compounds of gold are found in two oxidation states, aurous
(+1) and auric (+3). The auric state is generally but not always,
as in the case of Au(CN)2 , more stable than the aurous. Gaseous
and solid solution gold compounds are invariably bound covalently.
The auric ion forms co-ordination compounds with most inorganic and many organic liquids(5,89). The auric ion can exist free
in aqueous solution only when the solution is pure, strongly acid
(to prevent the formation of hydroxide complexes) and strongly oxidizing (to prevent reduction to the metal). All auric complexes
are strong oxidizing agents. The ability of gold to form complexes
is of the utmost importance in its technological applications and
recovery from gold ores.
The stability of Au(I) and Au(III) complexes tends to de
crease with increasing electronegativity of the liquid donor atom (90,91,92)#
54.
For example in complexes such as Au(SCN)2 and Au(SCN)^ it
is the S atom of each SCN“ ligand that is bound directly. The
order of stability of complexes with the halides is I- > Bi~
> Cl” > F” , while the order of electronegativities is F > Cl” >
Bi“ > I”. Electronegativity increases from left to right, and
from the bottom to the top of the periodic table. For elements
that would possibly form complexes with gold, Table 4.1 can be produced:
TABLE 4.1 - PATTERN OF ELECTRONEGATIVITIES FOR POSSIBLE Au LIGANDS
c N 0 F ,
P S ClAs Se Bi
Sb Te I
IncreasingStability
IncreasingElectronegativity
These trends account for the stability order noted for halogen
complexes above and explain why gold tellurides (AuTex) and the
antimonide, aurostibite (AuSb2>, are stable enough to be found in nature.
The co-ordination numbers for Au(I) and Au(III) are always 2 and 4, respectively. Compounds such as AUCI3 form dimers in the
solid or gaseous state to satisfy the co-ordination number.
55.
c i
90°
Cl
Cations such as Au+ and Au^+ occur in a hydrated state as
complexes containing a number of water molecules appropriate to
the co-ordination requirements of the particular oxidation state:
h2o - Au - o h2 + and h2o CMo
Au
h2o oh2
Since the atom bound to the gold in each case is oxygen, these
complexes are of low stability. The simple ions tend to react
in solution to replace the HgO molecules by stronger ligands.
The overall stability constants for gold ions are shown in Table4.2.
TABLE 4.2 - OVERALL STABILITY CONSTANTS FOR A SELECTION OF COMPLEXESOF Au(I) AND Au(III)
Au( I) Au(III)COMPLEX log 3 9 COMPLEX log 34
Au (CN)2 38.89
Au (I)2 21.07 Aul^ 48.75Au (SCN)2 13.20 Au(SCN)J 43.65AuBr2 14.55 AuBr^ 33.48Au C12 11.37 AuCL^ 26.34
Source: Bard, Parsons, Jordan(54).
56.
4.1.2 Oxidation of Gold in Acidic Media
The oxidation of gold has been studied extensively and reviewed comprehensively(93,94,95,96,97)#
At potentials below 1.7 to 1.8 V in aqueous solution the
oxide is one or two monolayers thick and above 2.0 V bulk phase
oxide is produced, which has been identified as AU2O3 or an
oxide of similar composition. However, only one or two monolayers thick may be considered to be chemisorbed oxygen(94). The structure
of the oxide depends on the potential and time of anodization. A
ratio of 0.7 to 1 for the cathodic/anodic charge has been measured.
Several mechanisms have been proposed to explain the behaviour of
the gold oxide: place-exchange, blister, island reduction, simul
taneous generation and dissolution of oxide films, etc.
The pre-1940 work has been reviewed(98). The anodic polar
ization of gold electrodes in acidic solutions forms an oxide
according to the following overall reaction(99):
Au 203 + 6H+ + 6e” = 2Au + 3H20; e£ = 1.36 V [4 .1]
This oxide or one of its hydrated forms has been identified in gold electrodes by ESCA(190,101,102).
A blister mechanism was defined(193) for oxide films with a weak adhesion force which allowed the film to rise like a blister
so that excess energy could be released. This mechanism has been proposed(194,105,106) ^0 explain gold oxidation in concentrated
solutions of sulphuric acid. However, no supporting evidence has been provided.
57.
At apparent current densities of 10“^ and 1CT4 A cm“2 an efficiency of 1.13% and 0.53% were measured for gold oxide formation(104:,107) #
Potential sweep experiments in 4 M H2SO4 demonstrated that
the onset of AU2O3 formation was at 1.35 V and that the maximum reduction current was at 1.1 v(108). Potentiostatic anodization
(up to 1.75 V) and galvanostatic reduction in 1M HC104 has shown(109)
that:
a) surface oxidation was highly irreversible.
b) the extent of oxidation was determined by the potential and
the electrode history.
c) there was only one form of oxide. The cathodic charge was
less than the anodic charge possible due to the decomposition of an intermediate(109). Reduction and oxidation potentials
shifted to less positive values with an increase of 1 pH
unit (56 mV for the cathodic and 60 mV for the anodic process).
The following mechanism was proposed:
Au + H20 = Au0Ha^s + H+ + e" [4.2]
The chemisorbed OH readily undergoes the following reaction:
Au0Hads = AuO + H+ + e” [4.3.A]
58 .
where AuO represents a chemisorbed oxygen atom layer rather than a
definite oxide of divalent gold. Alternatively, the following
reaction could occur in which oxygen is evolved through electrochemical attack:
Au0Hads + h2° = Au00Hads + 2H+ + 26“ [4.3.B]Au0Hads * Au + 02 + H+ + e [4.3.B]
The clean surface regenerated in reaction [4.3.B] may undergo [4.2]
again.
Further oxidation may take place on the AuO surface:
AuO + H2O -► AuOOHa(js + H+ + e” [4.4]
AuOHads has the same empirical formula as hydrated AU2O3 . The
potential of cathodic arrest supports the formation of AU2O3 .
Also, the electrode potential after open circuiting suggested the
existence of a potential-determining system in the neighbourhood
of 1.36 V, the experimental value of the standard potential of Au (0H)3/Au (99).
Brummer and Makrides(HO) working the same potential region
and electrolyte as Laitinen and Chao^O^) found the following:
a) Most of the oxide i.e. chemisorbed oxygen was reduced at a
fixed potential which depended on the cathodic current density.
b) A Tafel slope of 41 mV was observed, independent of oxide
thickness or potential of formation between 1.45 to 1.85 V.
c) (8 log i/9 pH)^ was - 1.39 independent of oxide thickness or its potential of formation.
59.
d) The oxide became harder to reduce the higher the potential
of formation.
e) In more acid solutions, the charge decreased linearly with
increase in the galvanostatic reduction transition time, at
a rate in excess of that observed for simple dissolution of
the oxide.
The following mechanisms were proposed(HO) assuming that
AuOOH was the oxide found:
AuOOH + H+ + e' = AuO + H2O [4 .5 ]
AuO + H+ + e“ -► AuOH [4.6]
AuOH + H+ + e" + Au + H2O [4.7]
[4.5] and [4.7] were fast and [4.6] slow. However this mechanism
is only congruent with their observations b and c. Brummer(lll)
in a later paper stated that the oxide grows slowly with time and
becomes harder to reduce, apparently according to a semi-logarithmic
relationship:
Qa = a + b log Ta [4 .8 ]
where a and b are potential dependent constants and Qa is charge
passed. This "ageing” effect was greatest at the lower potential
of formation despite the greatest thickness of the oxide at the
higher potential and despite the larger change in the amount of
oxide at the higher potentials.
Makrides(H2) later confirmed the logarithmic relationship
between time and charge [4.8].
60.
Grueneberg(113) t using potentiodyanamic and capacity measure
ments in 0.5M H2SO4 found three peaks between 1.3 V and 1.8 V, at
which oxygen started to evolve. It was found that Qa = Qc and
proposed that the reduction of Au(0H )3 occurred in two steps:
Au (0H) 3 + 2H+ + 2e -► AuOH + 2H20 [4.9.A]
or
AuO(OH) + 2H+ + 2e" + AuOH + H20 [4.9.B]
AuOH + H+ + e“ Au + H20 [4.10]
Rand and Woods using potentiodynamic techniques in 1M
H2SO4 , distinguished between chemisorbed oxygen and bulk phase
oxide, since chemisorption is characterized by an almost linear
increase in coverage with potential. However, when bulk phase
oxide is produced at potentials > 2.0 V, there is a sharp rise in
coverage, which is often irreproducible and accompanied by changes
in surface roughness.
Ogura et al(H^) observed a similar logarithmic charge-time
relationship [4.8] to Brurnrner(m) but a and b were pH dependent
and the amount of chemisorbed oxygen increases with increasing
pH. It was proposed that anodization at low pH caused Au AuO -►
AU2O3 , followed by oxygen evolution. Qa was pH dependent in weak
acid solutions. These results were interpreted by considering OH
radical adsorption followed by complex hydroxide formation and a
stoichiometric layer of Au(0H)3 , as illustrated in the following process:
Au -► AuOH -► Au(0H)x > Au(0H)3
61 .
Au (0H)3 could grow without the accompanying evolution of oxygen.
It was concluded that AU2O3 was a better electronic conductor than
Au (0H>3 and conversely that Au(0H)3 was a better ionic conductor
than AU2O3 . Contrary to the work of Qgura et a l ^ ^ ^ , Vetter and
Brendt(H6 ) found that the amount of charge involved in the forma
tion of a layer of oxide was independent of pH.
Schultze and Vetter^117) measured galvanostatic anodic and
cathodic potential-time relations in 0.5M sulphuric acid. A place
exchange mechanism was proposed for the anodic oxidation of gold
and island reduction mechanism for the cathodic reduction of the
oxide. These mechanisms were later proposed for the oxidation of platinum(H8,119). The properties of the oxide layer were seen
to be dependent on the formation conditions of the layer and at
constant formation conditions the anodic layer growth was described
by a modified Tafel relationship:
log i = A+ + (E - E+)/b+ [4.1l]
This 0 - dependent Tafel equation is valid for 9 = 0 to
1.27 A+ = Log 2 x lCT^ A cm“^ and E+ = 1.15 V.
For the cathodic reduction
log i- = A- - (E - E_)/b_ [4.13]
where A- = log 6 x 1CT14 A cm~2 , E- = 1.55 V
b- = bO (1 + a_ Q) [4.14]
and b° = 38 mV, a- = 0.1
62 .
The double layer capacitance Cp decreased with increasing
oxide thickness according to the following relationship:
1 / CD = (1/Ci) (1 + 3.09) [4.15]
The kinetics of the layer formation was explained by the
following model:
At low fractional coverage (0.01 < eacj„ < 0.1) chemisorbed
oxygen ions were in equilibrium with the electrolyte according to:
H2°aq = °̂ ds + [4.16]
Gold ions are then replaced by oxygen ions in the surface monolayer
according to
°ld + Au3+ = °ox + Auld [4.17]
which depends on the field strength. [4.17] is rate determining
and yields an epitactic surface oxide. The slope of changing
curves was explained by the increasing potential drop in the growing
oxide layer at constant field strength. The cathodic reduction
proceeded only at the oxide edges (island reduction). The chemi
sorbed oxygen was not stable outside the electrolyte; gradual
decomposition took p!«ce in air to leave a clean gold surface(423)#
Several authors(120,121,122,123,124,125) have observed that
the oxygen adsorbed at a gold electrode was present in two forms: a less strongly bonded form stripped at potentials 1.30 - 1.35 V
and a more strongly bonded form whose stripping potential was approximately 400 mV more negative. It was suggested that as
6 3.
the time spent by the oxygen on the surface of the electrode in
creased, its bonding to the surface strengthened, and the reduction
current peak gradually shifted towards less positive potentials.
With an adsorption time greater than 0.1 second only the more
strongly bonded form of adsorbed oxygen was observed. The rate of
strengthening of the oxygen-gold bond increased with an increase
in potential. The following mechanism was proposed to account for the strengthening of the oxygen-gold bond(1^0,121).
Au + H20 -► Au(0H)ads + H+ + e~ [4.17]
leading to the formation of the first form, and
Au(0H)ads + Au(0)adg + H+ + e" [4.18]
leading to the formation of the second form. [4.17] was fast while
[4.18] was slow.
Moslavac et alC^-26)^ using the galvanostatic pulse method
in 0.5M sulphuric acid observed no linear relationship between i vs
1/T and proposed a mechanism based on the simultaneous formation
of a blocking anodic layer and its dissolution. This model assumes
a dependence of the corrosion current on the coverage of the gold
electrode. The following relationship was then developed(126)•
T = 1 In (1 - a) [4.19]a b
where T is the transition time and a and b are defined as:
a = kik S 1 [4.20]b = kis S_1 [4 .2 1]
6 4.
pwhere k is the number of cm covered by one coulomb, and and ig
are respectively the corrosion current and the anodic apparent
current density and S is a constant.
Ferro et al( ̂ 7,128,129) using very fast sweep rates in 1 M
HCLO4 and 0.5, 2.5 and 5.0 M H2SO4 in the temperature range of
-10°C to 70°C found three well-defined cathodic peaks between 0.9
and 1.3 V. These peaks were interpreted as due to three different
electrode reactions. Each reaction took place within a definite
potential range, suggesting thus the presence of species with
different bond strengths to the surface. The mechanism proposed by Ferro et al(^29)> envisaged the initial formation of adsorbed
AuOH species that were susceptible either to further oxidation to
a higher "oxide" or to chemical disproportionation to a more stable
surface oxide with a stoichiometry comparable to that of AU2O3 .
These species give rise to different reduction peaks, the relative
magnitudes of which varied with the positive potential limit and
the time allowed for the chemical disproportionation reaction to
take place.
Hamelin and Sotto^*^ working at neutral pH in 0.05M K^SO^
compared the planes (100), (110) and (111) on single gold crystals,
showing that oxidation occurred at more positive potentials on
(100), and (110) planes than (111).
Dickertmann et al(131) studied single crystal planes, (111)
and (100), and poly crystalline gold electrodes in 1M H2SO4 at 25°C using potentiodynamic techniques to test the hypothesis of epitaxial oxide(H7). Having used a novel cell design, it was
found Qc/Qa = 0.7 to 0.8 and one peak for each plane(131)#
65 .
The anodic peaks of planes (111) and (100) differed in their potential by 200 mV. This work was in qualitative agreement with the work of Sotto(^ 2 ,133,134) ̂ although Sotto showed two peaks
on plane (111), Dickertmann et al (131) seated that Sotto was
probably unable to isolate the planes properly. The 200 mV differ
ence for the anodic peaks of different planes was explained by the
difference in the effective charge of the reactive species or by
the difference in potential distribution on • both planes. It was
further concluded that on polycrystalline gold, planes with lower
packing density, e.g.' (110) were also stable and contributed to
the shape of the curve(131). Also the structure of the oxide
seemed to be independent of the substrate orientation, considering
that the cathodic reduction was similar for both planes and previous polycrystalline gold electrodes.
Rotating gold disc-gold ring electrode (RGRDE) and a rotating
gold disc-platinum ring electrode (RGPRDE) studies in 0.2M f^SO^l*^
found that:
i) At potentials > 1.37 V gold dissolves at a rate of 0.035 x 10~® A cm“2 and dissolution increases with potential. Brummer and
Makr ides (HO) had reported 0.3 x 10-® A cnr^ at 1.37 V in
HCIO4 . Therefore it was considered that the difference in
rate might have been due to the presence of trace chloride
ions.
ii) Gold dissolution observed during cathodic reduction of
oxide films was dependent on rotation rate and the extent ofoxide reduction.
iii) The oxidation state of cathodically produced gold was depen
dent on the gold oxide formation potential. If a gold disc electrode was oxidized at E < 1.7 V, Au(III) were the pre
dominant species collected at the ring. If the disc electrode
was oxidized at potentials >1.7 V then significant Au(I)
species were collected at the ring. Cadle and Bruckenstein(135) proposed the following mechanism:
K(E) K(0)AuOx -------— Au(III) -------- ^Au(O) [4.22]
+solution
where K(E) is potential-dependent and some AuOx might be
reduced directly through a parallel path. This mechanism
[4.22] does not explain the significant amount of Au(I) pro
duced at higher potentials. These findings may have practical
consequences in thin film and intergrated circuit technology(136) .
The role of ion adsorption on gold electrodes has been studied(137,138,139). The presence of cations apparently affect
the state of adsorbed water and decrease the lateral repulsion of
oxide species and the electric field in the double layer. These
effects tend to block the initial stages of oxidation, retarding
the place-exchange mechanism and leading to a stabilization of
oxides.
Ellipsometric( 1^0,141,142,143) anci optical reflectance tech
niques^^, -^5,146) have been applied to in-situ observation of
the formation of oxides film on gold electrodes. Both techniques
have detected changes in the optical properties of gold electrodes
6 7 .
in the potential region where chemisorbed oxygen appeared. However,
at present the techniques have added little new information.
Several authors^147*148) have applied polaromicrotribometry
(i.e. the measurement of coefficient of friction and current- potential curves in situ) to gold electrodes. They(147>148) have
correlated the electrochemical characteristic (chemisorbed oxygen)
with the friction phenomena successfully. However, again, little
new information has been derived.
4.1.3 Corrosion of Gold in Chloride Media
Corrosion and passivation of gold in chloride media has very
important technological implications in the fields of electronic(149)
and gold refining(89). The work on this area can be divided
into the pre-1960 period characterized by instrumental constraints,
and the post-1960 period when potential-dynamic and rotating disc
electrode techniques were available.
Shutt and Walton(l®0il51)^ using constant current chrono-
potentiometry observed that the time to passivation was described by the following equation:
t (i - ix) = k [4.23]
where i = current density
ix = limiting current density, above which passivation occurs
t = time of passivation
k = number of coulombs in excess of the limiting current density required for passivation of 1 cm2 of gold surface
68.
For low stirring rates and current densities:
t0 ,5 (i - ix) = k1 [4.24]
The limiting current density was proportional to chloride ion
concentration and pH had only a minor effect. It was postulated
that the limiting current density (i^) was due to the rate of
chloride adsorption.
In an earlier paper Shutt and Stirrup(152) postulated that
the passivation time was determined by the surface concentration of oxygen on gold metal. Later work(153,154,155) confirmed the
validity of [4.23], but Armstrong and Butler(153) explained the
passivation of gold by depletion of chloride ions at the surface,
the limiting current density (ii) being the rate of transport of
chloride ions through the diffusion layer. Both k and wereproportional to the bulk chloride concentration. Muller and Low(154)
proposed that passivation occurred by a surface layer of oxide,
which could be dissolved in 5M HC1, and that the limiting current
density (i^) was due to limited convective diffusion.
Just and Landsberg(156,157) us-Lng a rotating gold disc elec
trode in the range 0.5 to 4M H G demonstrated the validity of
[4.24] and that convection caused deviation from the Sand equation (156) producing a relationship of the type of [4.23]. The diffusion
coefficient for chloride ions derived from the Sand equation, Dq -
= 1 .2 x 10“9 m^ s~l, agrees well with that calculated from con
ductivity data, Dq - = 1.33 x 10”® n f i S“l.
6 9.
Prepolarization at current, ii < i]_ and subsequent application of current, i > i-̂ confirmed diffusion control(156,157)
(i - i-^ t1/ 2 = 0.570 v1/ 6 D- 1 / 6 W 1/ 2 - i1) [4.25]
where v = kinematic viscosity in m2 s~l
D = diffusion coefficient of chloride ion in m2 s”l
W = rotation rate in s”1
and 0.570 is a dimensionless constant. A slope of 0.327 s^*^
compared with a theoretical value of 0.344 s®*^ was determined (156,157).
Heumann and Panesar(158) investigated the mechanism of the
active dissolution of gold in acidic chloride solution and suggested
formation of both Au(I) and Au(III) species. The polarization
curves were characterized by active gold dissolution, chloride
evolution and oxygen evolution (Figure 4.1). Passivation occurred
at about 1.45 V, independent of the chloride concentration. The
authors used weight loss data and/or the gold content of the
solution to separate the current-potential curves of the gold
dissolution region into the two partial current-potential relation
ships for Au(I) and Au(III) dissolution (Figure 4.2). A current
efficiency of 100% was assumed and from the Tafel slope the anodic
transfer coefficients for Au(I) and Au(III) were 0.71 and 0.36,
respectively. The deviation of the Tafel plot from linearity,
observed at high current densities where the dissolution occurred
primarily through Au(III) formation, was attributed to chloride ion depletion, in accordance with previous workers(153,154,156,157).
70 .
Figure 4.1 Potentiostatic polarization of gold in 0.25 M sulphuric acid and (Q) 0.5 M NaCl, ( A 0.25 M NaCl, ( O ) 0.1 M NaCl, (x) 0.05 M NaCl, (0) 0.02 M NaCl, (Q) 0.005 M NaCl. (Source: ref. 158).
E(V) e (v )
Figure 4.2 Partial current density-potential curves for Au(III) and Au(I) dissolution in 0.25 M sulphuric acid and (0) 0.5 M Nad, (A) 0.25 M NaCl, (O) 0.1 M NaCl, ( x) 0.05 M Nad, (0) 0.02 M Nad. (Source: ref. 158).
A potentiostatic study at 3.5% NaCl concentration by Robinson
and Frost(159) surprisingly showed that the shape of the potential
current curves was similar to that in 20% H2SO4 ; passivation
occurred at 0.3 V at a current density of 160 ya cm“2. Chloride
ions addition at concentrations > 2 x 10“^ M caused the complete
disappearance of the hydrogen oxidation current observed in chloride
free solutions^®®). This was due to the adsorption of chloride
ions at potentials close to and anodic to the point of zero charge (p.z.c.) of gold(161,162,163,164) thus blocking the surface sites
necessary for hydrogen oxidation. The anodic dissolution of gold in chloride media was assumed to occur with formation of a soluble
tetrachloro-gold complex at a standard potential of 1.0 V, although the production of Au(I) complexes had been proved previously(158).
In acidic, chloride-free solutions, a potential of 1.0 V is quoted
for the start of the formation of an oxide and/or adsorbed oxygen
film on gold electrodes^*®9). Presumably the first step in de
passivation is the adsorption of chloride ions in preference to
the oxygenated species forming the passive film. The competition
between adsorbed oxygen and chloride ions was postulated in the
process of passivation(149>160>*67). The dissolution of gold was
charge transfer controlled and occurred only at the oxygen-free surface sites(*®®).
Cadle and BruckensteinC*®®), USing a rotating gold disc -
platinum ring electrode in 0.2 M H2SO4 with 2 x 10 M to 10 M C1-, observed that adsorbed chloride ions exist in two different
72.
states on a gold surface. No Au(I) ions were collected at the
ring in this range of chloride concentrations during the anodic
scan. However two peaks were found at the ring (Er = 0.0 V).
Herrera-Gallego et al(166) j_n a study of the Au - CL
- H2O system using potentiodynamic techniques plus rotating disc
electrodes, described the anodic behaviour of Au as follows.
At E > Ep>z>c chloride ions were specifically adsorbed on the
Au surface and when E > 0.9 V, Au dissolved anodically. In the
range 0.9 V < E < Epassivation 'the chloride ion concentration at
the interface diminishes with potential until the anodia dissolution
became diffusion-controlled. When E = Epass, the chloride con
centration at the electrode surface tended towards zero and at
E > Epass either 0H~ ions or H2O discharge yielding a surface
oxide which passivates the metal.
Epass = 1-568 + RT In Cq - - 2.3 RT pH [4.26]F F
At constant Cl" concentration a change of pH between 0.5 to 7,
produced no appreciable modification of the poteniodynamic E/I
curve in the active dissolution region, but shifted the potential
of passivation, which was independent of rotation rate. The elec
trodissolution of gold showed a first order dependent on Cl" ion
concentration. The temperature dependence of the anodic current
peak fitted an Arrhenius plot with an apparent activation energy
equal to 7.5 kj mol"l. The initial portions of the anodic E/I curve,
corresponding to the activated electrode, were independent of AuClJ addition and yield linear E/log I plots with slope close to 2.3(RT) V/decade.F
73 .
Two mechanisms were postulated(16(5) which involved the highly
unlikely heterogeneous disproportionation of adsorbed Au(I) species
and did not allow for the anodic oxidation of Au(I) to Au(III).
Furthermore, these mechanisms predicted that Au(I) should be the major dissolution product at lew chloride concentrations in con
tradiction to the experimental observations(96).
Nicol et al studied gold dissolution(168,169) and the dts_
proportionation of Au(I) species(169»170) in acid chloride, in
relatively concentrated gold solution typical of the electrolytes
used m electrorefining of gold. Gold dissolution was found always
to produce some fraction (x) of solubilised gold as Au(I) species, m accordance with previous workers(158,149) and a mechanism was
proposed involving:
Au -► Au(I)ads + e [4.27]
x Au(I)ads * x Au(III) + 2xe“ [4.28]
(1 - x) Au(I)ads - (1 - x) AuCDaq. [4.29]
x was calculated from the following equation:
x 6ia^a *" Ic^c
^ic^c [4.30]
where ic
^c
la
*a
constant cathodic current
time of cathodic deposition from a gold(III) solution (usually 0.2M)
constant stripped currenttime required for complete dissolution.
ta was determined by the change in potential of the working electrode
due to the onset of chlorine evolution. This method requires
that the Pt disc electrode strips evenly across the whole surface
for its accuracy. However, the current density can vary between
50 and 100% of the average current density in certain circumstances (171). The current due to [4.27] and [4.28], since [4.29] is a
mass-transport step with no potential dependence, is proportional
to:
i = F Ki exp { (1-0!) FE } + 2FK2 [Au(I)ads] exp { (l-02) FE }
The mass-transport step [4.29] is described by K3 [Au(I)ads] •
Under steady-state and rearranging Nicol arrived at:
By plotting log (1/x - 1) vs potential (E) and assuming 02 = 0.5,
Tafel slopes of 60 to 80 mV per decade were predicted, which agreed
well with the experimental result and were consistent with a rate
determining step involving the transfer of two electrons. A value of 42.2 kj for K2 and 20.9 kJ for K3 was found.
The kinetics of Au(I) disproportionation were followed by
determining Au(I) species as a function of time by oxidation at a Pt disc electrode^ 169, 170). The homogeneous Au(I) disproportionation rate was found to be very slow in agreement with Lingane’s work(1^2) ̂
but enhanced in the presence of elemental gold, which would have provided low energy sites on which Au growth could occur. The
presence of oxygen also increased the rate of Au(I) disproportiona
RT RT
Kg exp { (l-02) FE/RT } [4.32]
75.
tion, though only in the presence of elemental gold. Presumably
this corroded by oxygen reduction to produce Au(I) directly or by
its homogeneous oxidation to Au(III), which was then reduced by
elemental gold to produce Au(I) in solution. The mechanism proposed
for homogeneous reaction was
Au(I) + Au(I) = Au(II) + Au(0) [4.33]
Au(I) + Au(II) = Au(III) + Au(0) [4.34J
Heterogeneous catalysis of the reaction is not unexpected since
there is a significant potential region in which the oxidation of
.Au(I) to Au(III) and the reduction of Au(I) to Au(0) overlap, i.e.
Au(I) = Au(III) + 2e“ [4.35]
2Au(I) + 2e" = 2Au(0) [4.36]
is possible on a gold surface. Nicol(170) obtained 8% collection
efficiency rather than the theoretical 33% because apparently
reaction [4.35] is not mass-transport controlled at a potential of
1.241 V above which chlorine evolution precluded the measurements.
Podesta et al(173) studied the current oscillations found
in the potential region 1.5 to 1.7 V (Figure 4.1) prior to passiva
tion. Following a potential step (Es) the oscillation frequency
(f), increased linearly with time. In the active dissolution
region the current decay fitted a linear I/t“0*5 plot characteristic
of C l - ions diffusion control and in agreement with the results of Franck(174). in agreement with other workers(175,176)
number of electrons per mol of reacting Cl“ ions was found(173)
76.
to be between 0.5 to 0.75, compared with a value of 0.66 derived
from voltammetry. It was concluded that two main electrochemical
reactions (gold dissolution and oxide layer formation) coupled to
two diffusion processes were responsible for the periodic effect.
Frankenthal and Siconolfi(177) used saturated solutions of
NaCl and observed that at sufficiently low potentials (< 0.8V),
gold anodically dissolved as Au(I) with a reaction order in d “
of 1.9. The reaction rate was independent of pH for pH > 1.5, but
increased markedly for pH < 0.
The following mechanism was postulated(177):
Au + 2 d = (Au Cl2)ads + e [4.37]
For pH < 0
(Aud2)ads + H+ -► HAud2(soln)
HAud2(Soin) = H+ + A u d 2 ; (K* * 1 + 10) [4.39]
[4.38]
For pH > 1.5
(Aud2)adg -► Au(̂ -2(soln) [4.40]
where step [4.40] is the rate determining step.
77 .
4.1.4.The Electrodeposition of Gold from Chloride MediaEvans and Lingane(178) showed as expected that gold electro-
deposition from chloride electrolytes was an irreversible process
using chronopotentiometric measurements. At concentrations of 10“^—9 — —to 10 M AuCl^, AUCI2 intermediates could not be detected in
1 M HC1. In mixtures of AuCl^ and AuClJ, two ill-defined waves
were obtained, the first due to AuClg reduction.
Harrison and Thompson(l^), using a rotating disc electrode,
proposed an E.C.E. mechanism for the reduction of AuClJ, however
the chemical reaction was not defined. The electrodeposition was
found to be diffusion-controlled, first order in gold, independent
of chloride concentration and with n=3 (number of electrons involved
m the reaction) for an assumed value of D^u° = 1 0 rrrs accord
ing to the Levich equation [6 .8 ]. At lower overpotential a Tafel
slope of 60 mV was found.
The following mechanism was postulated(1^9):
AuCl^ = AuClg + Cl~ [4.41]AuClg + 3e Au + 3C1“ [4.42]
The 6QnV Tafel slope was interpreted as a fast one-electron first
step followed by a chemical step, although the possibility of an
initial slow two electron transfer was u o 'c excluded.
Herrera-Gallego et al(166) obtained two cathodic current
peaks provided that in the anodic scan, the potential limit was
lower than the anodic peak. The cathodic peak at 1 V was related
78.
to AuCl^ reduction and that at 0.675V for Aud^ reduction. By plotting the potential of the cathodic current peak vs. log sweep
rate a slope of RT/F or 2RT/F was found depending on the experimental
conditions. The peak potential was more cathodic at high chloride
ion concentrations. No appreciable increase in electrode roughness
was found contrary to the results of Evans and Lingane^^).
Schalch et al(169,180) studied cathodic depositon in concen
trated gold solutions. Using the rest potential of the platinum
working electrode as an indication of the concentration of aurous
ions, according to the Nernst equation:
E = Eo + RT In [Au(III)l [4.4312F [Au(I)]
The ratio of Au(III)/Au(I) was changed by evolving chlorine. This
method assumed:
1. The rest potential of Pt electrode was governed by the Au(III)/
Au(I) couple even in the presence of chlorine, and
2. The oxidation of Au(I) to Au(III) is a fast process at any
chlorine concentration.
In agreement with Harrison and Thompson(179) they found a diffusion-
controlled deposition which was first order in Au(III) concentration. The deposition from gold (I) is favoured both thermodynamically and
kinetically over that from gold (III). The following mechanism was postulated(180) for -the reduction of gold (III):
79.
AuCl^ = An Cl g + d ~ [4.4l]A u a 3 + 2e“ -► Au(I) [4.44]
Au(I) + e- Au(0) [4.45]
The slow step was [4.44]; however NicolC^O) was able to detect
Au(I) at the ring, which implies that the rates of reactions [4.44]
and [4.45] were comparable, otherwise Au(I) species would have been
consumed.
4.1.5 Cbrrosion of Gold-Silver Alloys in Chloride Media
Several varieties of native gold are known to contain some
amount of other metals(181) as listed in Table 4.3.
TABLE 4.3 GOLD ALLOYS
Argentian gold (electrum) (Au, Ag)Cuprian gold (cuproauride) (Au, Cu)
Palladian gold (porpezite) (Au, Pd)
Rhodian gold (rhodite) (Au, Rh)
Iridic gold (Au, Ir)Platinum gold (Au, Pt)Bismuthian gold (Au, Bi)
Gold Amalgan (Au 2 , Hg3)Maldonite (Au 2 , Bi)Auricupride (Au, Q 13)
Palladium Cuproauride ((Cu, Pd) 3 Au 2)
Source : Boyle(1^1)
80.
From the above alloys only argentian gold (electrum) could result
in corrosion inhibition in chloride media. Native gold may contain
silver, which usually is a lattice constituent; there is a complete
substitutional series from gold through argentian gold to aurican
silver (kustelite) to native silver. The term electrum has been
applied to gold containing 20 per cent or more of silver.
Andonova and Kamenetskaya^182) studied the effect of Ag, Qi,
Pb, Bi, Sb, S and Sn on the anodic dissolution of gold alloys.
In their experiments they used a current density of 1000 A m“2
and an electrolyte typical of gold refining plants (150-200 g Au dm~3
and 80-100 g HC1 dm”3). it was concluded(182) that when silver and
lead content were not over 13%, the electrolysis proceeds normally.
Arsenic and anitimony, within limits of 0.2-0.4% have no influence in the electrolysis. Gubeidulina and Zyryanov(183) found that
Au dissolution ceased only when silver content reached 30% and the
dissolution rate of gold was halved in the presence of 10% Ag.
In the envisaged process the silver content of certain gold
ores might present a problem. However, the Au/Ag ratio varies
through a wide range and for the South African gold ores the ratio Au/Ag is typically 12.3(181)# South Africa produce between
40-60% of the world production of gold(184). An alternative would
be to increase the chloride ion concentration of the leach solution
as well as the temperature (section 4.2.3).
81.
4.1.6 Chlorine as an Oxidant in Gold Hydrometallurgy
Gubaylovskiy et al(185) studied the kinetics of the gold
in aqueous solutions containing chlorine using a rotating disc
electrode. The dissolution reaction was diffusion-controlled. The
presence of nitric, sulphuric and hydrochloric acid lowered the pH,
chlorine being the only oxidant present, as opposed to hypochlorous
acid, allowing a more efficient dissolution. The rate of dissolution
in the presence of hydrochloric acid is faster because the hydro
chloric acid dissolves the AuCl film formed at the surface of the
gold. At pH < 0 chlorination is a first order reaction with respect
to chlorine. In 0.5 M H d with a chlorine concentration of 2.75
mol CL2 m”2, the dissolution rate was 8 x 10“^ g m “2 s--*-.
Gubeidulina and Zyryanov(183) founcj a dissolution rate of
0.28 g m~2 S-1 in o.l M Hd;but no chlorine concentration was stated.
Kakowskii et al(*®®) studied the dissolution of gold tellurides
by using a synthetic ore of composition equivalent to calaverite
(AuTe2). They found that a decrease in pH from 1 to 0.2 decreased
the dissolution rate of gold by 25% and increased the dissolution
rate of tellurium by seven times. In a comparative experiment they
found that the maximum dissolution rate achieved when agitating the
electrolyte at 600 r.p.m. and a cyanide concentration of 1 % mass,
was 7 x 10"® g m~2 s-l. However, using the same rotation rate
and a chlorine concentration of 0.03 M, the dissolution rate was
8 .2 x IQ"5 g nr2 s"1.
82.
Volkova and Filipot1^?) obtained for the dissolution of gold
in the presence of chlorine, at temperature < 30°C an activation
energy of 14 - 23 kJ mol”1 AuCL^ and 6 - 9 kJ mol”1 AuCl^ for > 30°C. The dissolution product was AuCl^ which was further oxidised
to AudJ.
4.1.7 Recent Attempts at Chlorination
Walker( ^ 8) studied the possibility of using chlorination
to treat gold slimes in the later 1950s. The typical composition
of the slimes used is shown in Table 4.4.
%
Gold 19.3Silver 3.9Zinc 27.6Lead 3.8Copper 2.7Mercury 2.7Filter Aid 5.0Minor Constituents 5.0Moisture 30.0
100.0
TABLE 4.4
Typical composition of gold slimes used by WalkerC1^).
The process proposed by Walker was not very different from
the chlorination process used in the 19th century . in essence
the process (Figure 4.3) involved:
1. Chlorine gas was passed into the pulp of gold slimes with
rapid agitation to convert all metals present into their respective chlorides.
Figure 4.3 Process proposed by Walker for the recovery of gold from Merril slimes(190).
2. Insoluble material, including the chlorides of silver and
lead, were filtered off and well-washed to remove any soluble
salts.
3. A solution was added to the filtrate of sodium sulphite to
precipitate the gold.
4. After filtering and washing, the precipitated gold was dried,
melted under a borax cover and cast.
5. The filtrate containing copper, mercury, etc., was further
treated to recover silver
It was found that the amount of heat evolved was sufficient
to raise the temperature of the pulp so that the greater portion
of the liquid present would boil off. The maximum absorption of chlorine took place at 80°C and the quantity of chlorine used was
almost that theoretically required to form the chlorides of the metals present in the slime(188)#
Finkelstein et at^l^O) ancj zyl ai( 191,192) tried to adapt
the process proposed by Walker(188) to treat South African gold
slimes and gravity concentrates. The following differences were
found with respect to the results reported by Walker:
1. Reaction rates were apparently slower than those reportedby W a l k e r w h i c h probably accounts for the fact that Finkelstein et al(190,191,192) did not need to use cooling
at any stage of the process.
85 .
2. Contrary to the results of Walker, it was found that acid
treatment of the slimes before chlorination did not reduce
the extraction efficiency. The rate of chlorination was also
appreciably unaffected and considerably less heat was pro
duced, as conversion of zinc to the chloride was the major
source of heat during the chlorination of Merill slimes.
The flowsheet shown in Figure 4.4 was proposed( to
treat Merrill slimes. Sodium hypochlorite generated in situ was
studied as an alternative to chlorine. However, although it
was technically feasible, no economic advantage was apparently
found. A pilot plant was built and it was concluded that the
cost of chlorinating the gravity concentrates would be similar
to the cost of applying cyanidation.
More recently, Muir et al(l®3) described the development of
a process to recover gold from a highly antimonial slag using
chlorine. The process consisted of a flotation concentrate pre
pared from the slag and leached with CI2 and HC1. After the
separation of Sb as a hydrated pentoxide formed by hydrolysis,
gold was adsorbed onto activated carbon. Figure 4.5 shows a
plot of the percentage dissolution of Au and Sb as a function of the redox potential.
The rate of the dissolution of Au became rapid only when
the redox potential of the slurry reached a value of 0.85 V vs
SCE, and at a value of over 1.0 V vs SCE it was considered to be
complete. The time that was taken for completion was dependent on the rate of CI2 addition, and usually took from 4 to 6 hours.
Figure 4.4 Proposed process for the wet chlorination treatment of Merril slimes(190).
86.
87,
(193),Figure 4.5 Gold dissolution as a function of redox potential
88.
The main problem encountered by Muir et al(193) was the deposition
of Au in pipes, pumps and other parts of the plant by the reduction
of gold-chloride complexes. Finally, gold refractory ores con
taining activated aarbon and various types of carbon compounds
have the problem of very low recovery (typically 30-40%) with
cyanide, (Section 2.1.7). At present chlorine is being used at Carlin, Nevada ( H >12) to oxidise the carbonaceous materials to 00
and' 002* The chlorine reacts with the limestone in the ore to
produce calcium hypochlorite which in turn reacts with the car
bonaceous matter to form calcium chloride, 00 and OO2 . Following
the oxidation treatment the ore is subject to conventional cyani-
dation.
4.2 SILVER
4.2.1 Introduction
Silver is a transition metal, belonging to the IB group
elements together with copper and gold. Like the other group
IB elements, silver crystallizes in the face-centred cubic lattice
and the closest inter-nuclear distance is 288.8 pm. Silver has
an atomic weight of 107.87 and a density of 10.5 gcm_3} the pure
metal melts at 419.5°C and boils at 906°C.
Compounds of silver are found in three oxidation states:
+1, +2 and +3. The +1 oxidation is stable and +2 and +3 oxidation
states of silver are powerful oxidizing agents. The Ag(I) ion
forms stable complexes with ligands of strong it -bonding character. Those ligands which form strong complexes tend to form linear structures, L-Ag-L^1^4).
69.
4.2.2 Oxidation of Silver in Acidic Non-Complexing Media
The Ag/Ag(I) couple has been studied widely because of its reversibility. The standard potential for this couple is(^4):
Ag+ + e -* Ag E° = 0.7991 V [4.46]
Several authors(195,196) have found that the Berzins-Delabray(197)
theory concerning the reversible deposition of an insoluble com
pound fits the Ag/Ag+ electrode well. The couple Ag(I)/Ag(II) has a standard potential of 1.98 V. Fleischmann et al(198) studied the
anodic oxidation of Ag(I) in very acid solutions. It was found that
in the concentration range of 10“^ M to 1 M Ag(I) and over the pot
ential range of 1.7 V to 1.95 V, the current was first order with
respect to silver (I) concentration. At potentials higher than
1.9 V a silver (II) oxide occurred. The actual potential where
growth of solid phase began was dependent on the concentration of
Ag(I) and pH. Tafel slopes of 60 mV/decade were obtained and the
oxidation of silver (I) to silver (II) was found to be diffusion-
controlled. In less acidic solutions Ag203 is formed according to:
Ag203 + 6H+ + 4e -► 2Ag+ + 3H20; E° = 1.67 V [4.47]
Ag2()3 has .been shown to exist only when stabilized by oxygen
anions(199). Several comprehensive reviews have been published
concerning the thermodynamics(200,201) and the electrochemistry (199,202,203) Qf the oxides of silver.
90 .
4.2.3 The Oxidation of Silver in Acidic Chloride Media
The Ag/AgCl electrode has been very well studied due to its
applications as a reference electrode of the second kind and its
analytical capabilities in the determination of solubility products, halide salts activities, etc.(203). Several comprehensive reviews of the Ag/AgCl electrode have been published(203,204,205)#
The standard potential for the silver/silver chloride electrode is(54):
AgCl (s) + e” -► Ag(s) + Cl"aq# ; E° = 0.2223 [4.48]
The solubility product, Kgp(AgCl) = 1.77 x 10“^ . The electro
oxidation of silver in aqueous chloride solutions to produce silver chloride occurred via complex formation(206)#
At high chloride concentration (> 4M) and high current
densities, a barrier layer film was formed which exhibited high field conduction(207). The barrier film changed to a porous
condition at potentials higher than 20 V or when the oxidizing
current was either interrupted or briefly reversed. These films
were found to be non-stoichiometric "silver-excess" silver chlo- ride(^58). The ratio of Ag/Cl in the films varied between 1.07
to 1.10 depending on the current density applied. Katan et al(209,
210) observed that in porous silver chloride films in 1M KC1,
the anodic oxidation progressed within the pores by pitting at
dissolution sites of silver, probably at dislocation sites, and by
deposition of AgCl at some distance from the pits at nucleation
sites. It was found that the controlling step in the reduction of
91.
these silver chloride films in 2 M and 4 M KC1 was the surface
diffusion to growth sites on the silver surface.
Giles(211) reported that in concentrated chloride solutions,
silver dissolution proceeded at a rate too fast to measure,
according to equation [4.49 j:
Ag + nCL = Agd^n_1)" + e” [4.49]
— 9 Q_The ratio between the species AgCl^, Agdg , and Agd^ is pre
dicted by the thermodynamic solubility constraints of the individual
species (Table 4.5).
TABLE 4.5
Logarithmic Solubility Constants for Agd^j}-^ Species where n = 1, 2, 3, 4
KS1 Agd (mol 1 *) -6.63 *-7.00 -6.60 -6.60 -6.24
k S2 A g d 2 (mol l- 1 )2 -4.76 -4.70 -4.70 -4.70 -4.35
K S 3A g d 2- (mol l"1 )3 -4.58 -3.85 -4.40 -4.70 -4.12
KS4 A g d 3- (mol 1“^)^ -4.67 -4.52 -3.52 -4.46 -4.52
Source : ref. 212, taken from ref. 211
Kakovskii and Gubailovskii(213) studied the formation of
A g d in the presence of chlorine. It was found that the film
growth was according to a parabolic law at concentrations less
than 0.2 M chlorine and that the rate determining step was the
diffusion of silver through the film. To overcome the problem
of low solubility of silver chloride, the addition of organic solvents has been proposed(206)# Other authors(214,215) have
92.
demonstrated that the solubility of silver is satisfactory in the
presence of high concentrations of chloride ions and ferric chloride. Dinardo and Dutrizac(215) found that the solubility of
silver in 0.3 M HC1 can be described by the following equation:
S = a + bT + cT2 [4.50]
where S = solubility in g m~3 of saturated solution
a,b,c = constants
T = temperature
For instance in 1.5 kmol Fedg + 2 kmol N a d m”^ at 20°C, the
solubility of silver is 600 g m”3. it can be seen that the solubility
of silver increases with temperature.
The leaching of silver ores using brines has been reported(216)
where, for certain Peruvian silver ores; recoveries of 80% have
been achieved as opposed to 40% using cyanidation. The key parameter
has been a pH below 2. It seems that the oxidant used was oxygen, though this was not reported explicitly(216)#
93 .
CHAPTER FIVE
EXPERIMENTAL
94.
CHAPTER 5 - EXPERIMENTAL
5.1 ANALYTICAL TECHNIQUESTotal dissolved gold concentrations were determined by atomic
absorption spectrophotometry (AAS) (Baird Alpha III). Aqueous
solutions of gold chloride complexes are yellow in colour which
allows the use of U.V. spectrophotometry to characterize solution
composition. U.V. spectrophotometry was used as a routine analy
tical method for the determination of Au(III) concentrations
(*max = 312 11111 > nrolar absorptivity, ^ 1 2 = 662,1 11101-1 m2)« AuClg does not absorb in the U.V. region, allowing the determination
of gold(I) concentrates by subtraction from the total gold concen
tration. Two spectrophotometers were used during the project:
i) Perkin-Elmer model '200 double beam UV - Visible Spectrophotometer.
ii) Hewlett-Packard Diode Array Spectrophotometer 8451 A equipped with computer, disc drive and extended memory.
Chlorine was determined by Na2S203 titration following the standard procedure described by Vogel(217) and by U.V. spectrophotometry
using the multicomponent analysis programme available in the Hewlett-
Packard 8451 A.
5.2 SOLUTION PREPARATIONAll electrolytes were produced from analytical grade reagents
(BDH pic) dissolved in deionised, distilled water, which had been
redistilled from acid diohromate then alkaline permanganate to oxidize residual organics. All solutions were deoxygenated by
95.
nitrogen purging for at least 4 hours prior to chemical and
electrochemical experiments. This procedure reduced the dissolved
oxygen concentration to less than 2 x lcH> kmol the detectionlimit of a Philips PW 9600 dissolved oxygen meter(218). Nitrogen
('white spot' B.O.C. cylinder, < 10 ppm O2) bubbling was continued
above the solutions throughout the duration of the experiments.
5.3 ELECTRODE CONSTRUCTION AND PREPARATIONThe platinum rotating disc electrode was constructed by
cementing a disc of platinum, with silver-loaded epoxy resin, to
a suitably shaped brass rod, which was screwed on to the electrically
conducting rotor of an Oxford Electrodes rotating ring-disc electrode
assembly. The electrode was encased in epoxy resin to produce a
smooth cylinder, insulating all components except the flat disc electrode surface. The geometric surface area was 38.5 mm2.
The gold electrodes were prepared by cutting a gold foil
(Goodfellows Metals 99.99% purity), the rectangular shaped (5 inn
x 25 mm) gold foil was soldered to copper wire and then encap
sulated into glass. One side of the electrode was coated in
lacomit (W. Cannings Ltd, Birmingham) to improve the potential
and current density distribution.
The platinum indicator electrode was prepared by soldering
a Pt wire (Goodfellows Metals, 0.46 mm diameter) to a copper wire
and then encapsulated in glass so that only 0.5 cm of Pt was exposed to the solution.
96 .
5.4 ELECTRQCHMICAL STUDIES
5.4.1 Electrochemical Instrumentation
The working electrode potential was controlled with respect
to a reference electrode by a potentiostat (Thomson Ministat
28 V/1.25 A). In addition, a modular bi-potentiostat, manufactured
in the laboratory, was used for experiments involving a second
independent working electrode (mixed potential measurements).
Potential functions were controlled with a Hi-Tek PPR1 waveform
generator, and charge were determined by a Hi-Tek gated digital
integrator/DVM. Data was recorded direct into a JJ Lloyd Instru
ments PL4 xy/t analogue recorder. Sinclair/Thandar DM 450 digital
voltmeters were used. Rotating electrodes were mounted on an Oxford electrodes motor, driven by variable speed controller capable
of rotation frequencies up to 50 Hz.
5.4.2 Cyclic VoltammetryExperiments with planar or rotating disc Pt electrodes were
carried out in three compartment electrochemical cells, shown in
Figures 5.1 and 5.2 respectively.
The counter electrode compartments were isolated from the
working electode by a Nafion 425 cation exchange membrane (Du
Pont Inc.). A Luggin capillary probe provided contact between the
working and reference electrodes compartments. The tip of the Luggin
capillary was located close to the working electrode surface to
minimize any uncompensated ohmic drop due to solution resistance between the reference electrode and working electrode. The Luggin
tip was generally 1-2 mm from the working electrode surface to avoid
97-
Figure 5.1 Electrochemical cell design for experiments with planarelectrodes, indicating Pt working electrode (A), Pt counter electrode (B), Luggin capillary (C), N2 bubbler (D), Pt indicator electrode (E), sampling port (F), drain (G), vent (H), nation membrane (I), clamp (J), magnetic stirrer (K).
Figure 5.2 Electrochemical cell design for experiments with Pt rotating disc electrode. Chlorine evolving electrode (A), Pt counter electrode (B), Luggin capillary (C), N2 bubbler (D), Pt indicator electrode (E), sampling port (F), drain (G), vent (H), nation membrane (I), clamp (J), magnetic stirrer (K).
98.
shielding the electrode, preventing non-uniform current densities,
and scratching the working electrode. The cell in Figure 5.2
was equipped with PTFE lid with the minimum necessary openings.
A commercial saturated calomel electrode (SCE) (KENT/EIL) was
used as a reference electrode. All potentials are quoted relative
to the standard hydrogen electrode (SHE) unless otherwise stated.
The SCE was assumed to have a reversible potential of 0.242 V vs
s h e(205).
5.4.3 Constant Potential Electrolysis
Constant potential electrolysis was performed using gold
electrodes to derive the steady-state current vs potential curves
and as a way to generate solutions of known Au(I) to Au(III) ratio.
This was later checked using U.V. Spectrophotometry. The anolyte
was introduced into a three compartment cell (Figure 5.1) and purged with nitrogen to remove dissolved oxygen. The aatholyte was
introduced just prior to electrolysis to minimize any transport
across the membrane. Enough gold was dissolved in each experiment
to minimize analytical error.
5.4.4 Constant Current Electrolysis
Chlorine was generated at constant current using a power
suoply (Weir 4000T 30V/1A) in a three compartment cell (Figure
5.1). Previously the anolyte was introduced in the cell and purged
with nitrogen to remove dissolved oxygen. The catholyte was intro
duced just a few minutes before electrolysis. Samples were taken every 10 or 15 minutes, after the current had been interrupted and the Pt indicator potential and charge passed noted. Chlorine was
99.
determined by Na2S203 titration and a Pt indicator potential
vs CI2 concentration curve was plotted before every mixed
potential and net current experiment.
5.5 SPECTROPHOTOMETRIC STUDIESThe disproportionation reaction [6.4] was studied by genera
ting a AuClrJ solution (Section 5.4.3) and putting into three beakers
(PTFE, plastic and glass) in a closed container under a vapour
saturated nitrogen atmosphere in thermostated bath at 25°C (Figure
5.3). Samples were withdrawn, using a syringe with a hypodermic
glass needle, every 15 days on average and analyzed for total gold
and Au(III); the Au(I)/Au(III) ratio vs time curve was plotted.
The kinetics of AuG^ oxidation by G 2 was studied using the
stop-flow apparatus (Figure 5.4) and Hewlett-Packard Spectrophoto
meter. The spectrophotometric cell was clamped inside the spec
trophotometer and the cell was flushed 5 times with the electrolyte
for study. The computer was programmed to take spectra for different
experiments at intervals of 10 seconds, 30 seconds, 1 minute and
10 minutes. The spectra were stored in the extended memory of the
computer for later transfer to disc and analysis.
5.6 MIXED POTENTIAL MEASUREMENTSThe gold leaching rates were determined by maintaining a
rotating gold-plated Pt disc electrode in contact with a solution
of constant G 2 concentration. Samples were taken every 10 to 15 minutes depending on the G 2 concentration and analyzed for total
gold content. The slope of plots of gold concentration vs time provided the leaching rates. Chlorine was generated (Section 5.4.4)
100.
A= ^2 cylinder
B = rotameter C = N2 bubblers 0 = conical flask
Figure 5.3 Apparatus to study the disporportionation of AuCl^-
Figure 5.4 Stop-flcw apparatus
101.
and a Pt indicator potential vs CI2 concentration curve plotted
for every experiment. The chlorine was transported through a PTFE
tubing to the cell shown in Figure 5.2, which was previously flushed
with nitrogen gas. A bi-potentiostat was used to control the CI2
evolving electrode and the gold-plated Pt disc electrode. Once
the desired CI2 concentration was achieved, the gold-plated Pt disc electrode was allowed to corrode. The chlorine concentration
was kept within ± 5% of the chosen concentration by evolving CI2
using the bi-potentiostat. The chlorine concentration in solution
was determined by the potential of the Pt indicator electrode.
5.7 NET CURRENT METHOD
The reduction currents of chlorine on gold were determined
by the net current method. A polarization curve of the gold-plated
Pt rotating disc electrode was obtained in the chosen de-oxygenated
electrolyte. Then a second de-oxygenated electrolyte (same com
position as the first) was used and chlorine was electrogenerated
following the same procedure as described for the mixed potential
measurements (Section 5.6). Once the desired concentration of
CI2 was reached, then a second polarization curve was taken. This was repeated for the different rotation rates and chlorine
concentrations. The net reduction current w^s obtained by subtracting the first polarization curve from the second.
5.8 FLOW CIRCUIT EXPERIMENTSFigure 5.5 shows the Perspex packed/fluidised bed electrode
cell, which incorporated a Ti/Ru02 mesh anode (IMI Ltd) and a
Nafion 425 cation exchange membrane (Du Pont Inc.). The cathode
102 .
A
i j i t '
l i \2"Vu-"si!i
■ iji_,
0 cm 81 _i_____ i i i
Figure 5.5
Packed bed electrode (PBE). B - catholyte outlet, C - anolyte outlet, D - reference electrode compartment, E - Ti/Ru02 anode, F - Ti/Pt cathode feeder electrode for packed bed, G - flow distributor, H - Nafion membrane, I - anolyte inlet, J - .catholyte inlet.
103.
feeder electrode was a Ti/Pt mesh contacting the 10 mm thick bed,
which consisted of 2-3 nm cylindrically-shaped carbon chips (esti
mated projected area 0.09 m2 ) or a second bed of carbon particles
100% -200 um (-200 um carbon chips). The cross sectional active
area of the electrode and membrane was 0.145 m x 0.046 m = 6.67 x
10”3 m2. The membrane prevented transport of anionic gold species
to the anode, at which Au(I) species would otherwise have been
oxidized, and more importantly minimized transport of anodically
generated chlorine to the cathode, at which its reduction would
have decreased the current efficiency for gold deposition and
lowered the ahlorine utilisation.
The flow circuit shown in Figure 5.6 was constructed from
uPVC pipework, valves and fittings (G. Fischer Ltd.) and in
corporated 5 dm3 aspirators as reservoirs, Totton Electrics Ltd.
EMP 50/7 magnetically coupled polypropylene pumps and flow meters
with aaid resistant ceramic floats (Fischer Controls Ltd.). The
only corrodible materials in the flow circuit were the gold coated
particles in the leach reactor (Figure 5.6), which could be switched
into the circuit for those experiments involving leaching coupled
to electrowinning.
Gold containing solutions were produced by generating chlorine
in a separate three compartment cell, absorbing it in a solution
containing 1 kmol HC1 m-3 and reacting it with gold powder (99.99+%
- Goodfellow Metals Ltd.). Subsequently the residual ahlorine was desorbed by bubbling nitrogen through the solution for 12 hours.
104.
+ -
L E G E N D
A CENTRIFUGAL P U M P
A N G L E SEAT VALVE
BALL VALVE
o R O T A M E T E R
P A C K E D B E D
L- PORT VALVE
Figure 5.6 Reactor flow circuit
105.
The supporting electrolyte (1 kmol HC1 nr3) was nitrogenated
for 12 hours and pre-electrolysed until the residual current was
< 10 mA. Only then was gold solution added to the system, and
thoroughly mixed to give the required bulk solution concentration,
before being passed through the aell. For the packed bed electro
deposition experiments, the anolyte was 0 .5 kmol m”^ which
avoided the evolution of chlorine required in the proposed coupled
leach-electrowinning process.
A Wenking ST72 potentiostat was used to control the feeder
electrode (Figure 5.5) of the packed bed against a saturated
calomel electrode (SCE), with the Luggin probe tip located at the
bed/membrane interface, at which the maximum potential occurs in packed bed electrodes(6). This enabled that potential to be
constrained so that hydrogen evolution could be prevented. While
most of the bed could have been operated under more positive
potentials (i.e. lower overpotentials), the voltammetry indicated
a potential range of > 0.5 V for mass transported Au(III) deposition
in the absence of hydrogen evolution. Thus the whole bed volume
would have been operated under mass transport control, supporting
evidence whioh was provided by the visually uniform distribution
of the gold deposit, corresponding to the expected uniform aurrent
density distribution. However, no local potential distribution
measurements were made because of the difficulty of probing the
potential in the particulate carbon bed.
A Nioolet Explorer 1 digital oscilloscope and Gould 60000 XY/t recorder were used to acquire current/potential or current/time data.
106.
The apparatus used for coupling electrowinning and leaching
was the same as used in the electrodeposition experiments (Figures
5.5 and 5.6), two changes were made:
1. The by-pass was taken out of the circuit since it would add
dead volume in the apparatus.
2. The angle-seat valves were re-set so that the output from the
anode compartment would go to the cathode tank and the exit
flow from the cathode compartment to the anode tank.
The electrolyte was de-oxygenated overnight with oxygen-free
nitrogen. For this set of experiments it was not possible to
reduce oxygen electrochemically due to the new configuration of the flow circuit, since it would generate chlorine.
Samples were taken using hypodermic syringes and pipette
through sampling parts and analyzed immediately for:
1. CI2 from the anode exit flow.
2. CI2 and Au(III) using UV Spectrophotometry.
Later these samples were analyzed for:
1. Total gold concentration.
2. pH measurement where possible.
The data acquisition was performed in the same manner as in
the electrodeposition experiments in packed bed. Gold-plated carbon
cylinders of 3 mm length x 3 mm diameter were used in the leach reactor, the packed bed in the cathode consisted of -200 pm carbon particles.
107
CHAPTER SIX
RESULTS AND DISCUSSION
108.
CHAPTER 6 - RESULTS AND DISCUSSION
6.1 SILVERPotential-pH diagrams are a useful tool in the development
of new hydrometallurgioal processes. They summarize the thermo
dynamics of aqueous chemical systems, provided reliable thermody
namic data is available. A series of of Eh-pH diagrams were
computer-generated for the systems Ag-H20 and Ag-H20-CL-C104
to study the effect of Cl” ions on the redox reactions of silver.
Thermodynamic data are now available for three solid and four
aqueous silver species (Table 6.1) that were not considered by
Pourbaix(21). The Ag+ and Ag2+ cations hydrolyse according to the
general reactions,
mAgp+ + nH20 = mAg(0H)^p n +̂ + nH+ ;
For Ag+ (P = 1), n = 1 and 2, 3X>1 = -13.98, 32>i = -23.97.For Ag2+ (P = 2), n = 2 and 3, 32 ̂= —1*87, 33 ,i = -14.07.
The Eh-pH diagram for the Ag/H20 system (Figure 6.1) shows
predominance areas for Ag(OH) 2 and Ag(0H)2, and the predominance
of AgO rather than Ag202 contrary to the original Pourbaix diagram (
(Figure 6.2). AgO is relatively stable in practice, when dry and
in alkaline solutions. It has been proposed that AgO exists as
Ag(I).Ag(III)02 The existence of an equilibrium between
Ag(0H) 2 and both Ag20 and AgO in alkaline solutions, as indicated in Figure 6.1, has been well established experimentally(220).
109
Selected Free Energy ofTABLE 6.1
Formation Data System at 298
for Species K
in the Ag/P^O'
SPECIESOXIDATION
STATEAG°
kj mol~l REFERENCE
Ag° 0 0 .0 221
Ag20 (S) +1 -11.22 221
AgO (S) +2 +3.50 54
Ag2°3 (s) +3 +121.39 221
Ag202 (S) +2 +27.63 54
Ag02 (S) +4 -10.99 54Ag(OH) (S) +1 -91.99 222
Ag+ +1 +77.16 221
Ag2+ +2 +269.16 221
AgO+ +3 +225.63 54AgO“ +1 -22.99 222
AgOH +1 -80.21 221
Ag(0H)2 +1 -260.37 221
Ag(0H0)2 +2 -194.50 202
Ag(0H)3 +2 -362.07 202
A g a (s) +1 -109.86 221
AgC102 (S) +1 +94.18 221
AgCIO3 (S) +1 +73.67 54AgC104 (S) +1 +8 .2 203AgCl +1 -54.16 221
AgCLg +1 -215.58 221
Aga|- +1 -345.97 54
1COrfd
+1 -478.46 54AgC102 +1 +94.18 221
AgCl03 +1 +73.67 221
AgC104 +1 +68.52 221
110
Figure 6.1 Potential - pH diagram for the Ag/H20 system at 298 K, with a dissolved silver activity of 10-4.
Figure 6.2 Potential - pH diagram of the Ag/H20 system at 298 K^21^
111-
Eh-pH diagrams were generated to establish the effect of
chloride and perchlorate ion addition in the activity ranges 10”®
to 10“4 dissolved silver and 10”® to 1 perchlorate and chloride.
Inclusion of 10”® chloride^perchlorate (Figure 6.3) results in a
predominance area for AgCl(s), rather than for Ag+ ions. An increase of the activities of chloride-perchlorate ions to 10”®,
and then to 1.0 (Figure 6.4) was found to increase the predominance
area of AgCl(s). Under the latter conditions Ag20 was no longer
predominant and silver perchlorate was stable in preference to
Ag+ ions. However, the formation of AgClO^ is very unlikely due
to kinetic constraints.
The activity-pH diagrams for the Ag(I)-Cl-H20 system for
chloride-perchlorate activities of 1.0 and 5.0 is shown in Figure
6.5. The increase in chloride ion activity to 5.0’ produced an
increase in the solubility of silver-chloride species in acidic conditions. This is congruent with the work of Giles(211), who
found that the dissolution of silver in concentrated chloride
solution (6.09 M) proceeded with extremely fast kinetics according to equation:
Ag + n Cl” = AgCl^11”1)” + e” [4.49]
High chloride activities increase the silver solubility to formQ_AgCl^ (Figures 6.4 and 6.5). This result accounts for the re
ported successful leaching of Peruvian silver ores in brines at pH < 2(216). in conclusion, it would seen that in the envisaged
process, the required pH and chloride ion concentration may be
dependent on the need for concomitant silver dissolution.
112.
Figure 6.3 Potential - pH diagram for the Ag/H^O-Cl-ClC^ system at 298 K, with dissolved silver, chloride, and perchlorate activities of H T 4 , 1(T5 and 1CT 5 respectively.
Figure 6.4 Potential - pH diagram for the Ag/H20-Cl-ClC>4 system at 298 K, with dissolved silver, chloride, and perchlorate activities of 1CT"4 , 1.0 and 1.0 respectively.
113.
pHFigure 6.5 Activity - pH diagram for the Ag(I)/H20-Cl system at 298 K
( - ) chloride activity of 1 .0 ( --- ) chloride activityof 5.0.
Figure 6 .6 Potential - pH diagram for the AU/H2O system at 298 K, with a dissolved gold activity of 10“4 .
6.2 GOLD
114.
6.2.1 Thermodynamics
The effect of chloride ions on the solubility and redox
behaviour of gold was investigated by generating potential-pH
diagrams for the Au/f^O-Cl” system at 298 K. These diagrams represent graphically the thermodynamically favourable chemical and
electrochemical reactions of aqueous gold species and therefore aid
the interpretation of Au corrosion/leaching 9 Au electrodeposition
and Au passivation.
Thermodynamic data are now available for three dissolved species that were not considered either by Pourbaix(^l) or Finkelstein( 13)
(Table 6.2). The Au^+ cation hydrolyses according to the general
reactions,
mAu3+ + nH20 = mAu(0H)3-n + nH+; 3n,m
for n = 1 to 5, 3X>1 = 3.64, 32>1 = -3.21, B3>1 = 2.09,
34 ̂= —9.30, 85^1 — -22.66.
The main difference between the derived Au/H20 system (Figure 6 .6 )
Eh-pH diagram, and that calculated by Pourbaix(21) (Figure 2.2) is
that Au (0H)2+ rather than Au^+ is the predominant species at low
pH and high potentials. Au(s) and Au(0H)3(s) remained predominant
over wide potential and pH regions. Also, AuO^ or Au(0H)g ions
are stable at very high pH and potentials.
115.
TABLE 6.2Selected Free Energy of Formation Data for Species in the
Au/H20 - Cl System at 298 K
SPECIEOXIDATION
STATEAG°
kj mol“l REFERENCE
Au (s) 0 0 .0 21
Au (0H) 3 (s ) +3 -317.0 221
Au02 ( s ) +4 +200.8 21
AuO (s) +2 +27.3 223Au+ +1 +176 54Au2+ +3 +440 54
Au (0H) 3 +3 -283.5 221
1coco1 +3 -51.9 221
AuOH2+ +3 +182 194
Au (0H) 2 +3 -16 194Au (OH)J +3 -455.6 194Au(OH)g +3 -616.5 194
Au(OH) +1 24.0 223AuClg +1 -151.0 221
AuClg +3 -79.3 194A11CI4 +3 -234.6 221
Au(OH)Clg +3 -275.7 194
Au(0H)2Clg +3 -340.9 194Au (0H)3C1" +3 -400.3 194Au (0H)C12 +3 -151.0 194
Au (0H)2C1 +3 -220.0 194
116.
Eh-pH diagrams were generated to study the effect of Cl" ions
on the thermodynamic behaviour of the Au/H^ system. The range of
activities used were 2.5 x 10"5 to 10"2 dissolved gold and 10"^ to
5.0 Cl~ ions. Diagrams were generated for the system with both the
hydroxide-chloride complexes considered and absent. The mixed hydroxide-chloride complexes were proposed by Bjerrun(224) and
Chateau et al(225). The data for these species str'w some discrepancies(194) and other sources(54,221) have not considered
them.
The addition of Cl" ions with an activity of 10”5 to the
Au/HgO system with a dissolved gold activity of 10~4 , generates
a predominance area for AuClg and Au(0H )2 in the acidic pH and
high potential region (Figure 6.7). The solubility of gold increases
at very acidic pH (< 0). Au(s) is still stable within the
boundaries of H2O stability. If the mixed hydroxide-chloride
species are considered (Figure 6.8 ), the predominance areas for
AUCL3 and Au(0H )3 (s ) increase. The Au(0H ) d 2 and Au (0H)2C1 species
predominate at high potentials ( > 1.2 V) and pH 1 to 2.5. At
chloride activities < 10”5, chloride ions have no effect on the
AU/H2O diagram (Figure 6 .6 ).
An increase in chloride activity to 10~3, produces a pre
dominance area for AuCl^ at potentials higher than 1.5 V and
pH’s < 3.5. The equilibrium between AuClg and AuClJ is governed
by:
Au C14 = AuClg + Cl ; log K = -4.24 [6 .1 ]
117.
Figure 6.7 Potential - pH diagram for the AU/H2O-CI system at 298 K, with dissolved gold and chloride activities of 10-4 and 10”5 respectively.
Figure 6 .8 Potential - pH diagram for the Au/HgO-Cl system at 298 K, with dissolved gold and chloride activities of 10”4 and 10-5, respectively, considering mixed hydroxide-chloride species.
118.
If the mixed hydroxide-chloride species are considered, Au(0H)2d
is stable only in the pH range 4 to 4.5 at potentials > 1.0V.
At chloride activities > 10“3, as applicable in the envisaged
process, the mixed hydroxide-chloride species have no area of
stability for the range of gold activities investigated.
Figure 6.9 shows the Eh-pH diagram for activities of 5 x 10”5
Au and 0.5 Cl” ions. Comparison with the corresponding diagram
for the AU/H2O system (Figure 6 .6), demonstrates the powerful
depassivating and solubilising effect of CL” ions, due to reaction
Au (0H) 3 (s) + 3H+ + 4 d ~ = AudJ + 3H20 [6.2]
log(AuClJ) = 18.38 - 3pH + 4 log(Cl”)
Audg and AuCLJ ions are stable at potentials > 0.93 V and pH > 7,
indicating the thermodynamic possibility of leaching gold at pHs
as high as 7. Gold is oxidized to A u d 2 according to reaation,
Au C12 + e = Au(s) + 2 d ” [6.3]
E = 1.152 + 0.059 log(AuCl2) - 0.118 log(Cl")
At gold concentrations greater than those defined by equation
[6 .4 .a] Au(I) species are thermodynamically unstable, dispropor-
tionatmg according to reaction [6 .4 ].
3Aud2 -► AudJ + 2Au + 2C1” [6.4]AG° = -43.713kJ (mol AuCip -1
3 log(AuCl2) = 7.66 + log(AuClJ) + 2 log(Cl”) [6 .4 .a]
119.
Figure 6.9 Potential - pH diagram for the Au/H^O-Cl system at 298 K,with dissolved gold and chloride activities of 5 x 10“^ and 0.5 respectively.
Figure 6.10 Potential - pH diagram for the AU/H2O-CI system at 298 K, with dissolved gold ‘and chloride activities of 2 .5 x 10-^ and 5.0, respectively.
120.
Therefore high chloride and low gold activities improve the AuCl^
predominance (Figure 6.10). In the region of AuCl^ predominance
(Figure 6.10), Au(III) species could be reduced by the reaction:
AuClJ + 2e~= Aud^ + 2 d ” [6.5]
E = 0.925 + 0.029 log(AudJ) - 0.029 log^udg) - 0.059 log(Cl")
[6.5a]
At constant gold activity of 5 x 10“®, an increase in chloride
activity from 0.5 to 5.0 increases the AuClTJ predominance area
as predicted by equation [6 .4.a]. Similarly, at constant chloride
activity of 5.0 an increase in gold activity from 2.5 x 10”® to
10”^ decreases the area of predominance of AuCl^.
At lower chloride activities (~ 10“^ < Cl" < 1.0) gold is
oxidized to AuCl^ by reaction [6 .6 ]
AuClJ + 3e” = Au(s) + 4C1” [6 *6 ]
E = 1.001 + 0.0197 log(AuCl^) - 0.0789 log(Cl_) [6 .6 .a]
The high standard electrode potentials for the AuCl^/Au [6.3] and
AudJ/Au [ 6 .6 ] couples would allow electrodeposition of gold at high positive potentials and hence high current efficiencies.
However, this may present a problem as gold-chloride complexes can
therefore be reduced easily. To avoid this problem, an oxidant (CI2 )
should be present. According to Muir et al(^®®), the main problem
encountered in leaching antimonial gold slag with chlorine is the
cementation of gold onto metallic surfaces, e.g. pipes, pumps, etc.
This indicates insufficient chlorine concentration in solution.
1 2 1 .
Peters et al(226) have leached gold using Fe(III) in chloride
media (EFeCL+/Fe2+ = 0.772 + 0.059 log [(Fed2 )/(Fe2+) ] -0.118 2
log(Cl”)) as a final step in the treatment of copper sulphide leach
residues generated by the U.B.C.-OOMINCO process. Gold was reported
to be oxidized to Au CIq , probably according to the reaction
Au + Fed* = Au O -2 + Fe2+ [6.7]
A minimum chloride activity of 12 would be necessary for
gold oxidation to Aud^ = 5 x 10“® according to equation [6.3].
This would decrease the potential to lower than that of the FeClij/
Fe2+ couple, thereby allowing oxidative gold dissolution. The
HC1 activity in mixed Nad/HCL or C a d 2/HCl solution has been studied (227). For instance, at 1.8 M H d and 3.0 M N a d the
activity of H d is about 58(227). However, the use of such
strong salt solutions would incur high reagent consumptions due
to side reactions in practical operations. The use of strong salt
solutions in hydrometallurgical processes has been reviewed com- prehensively(228,229,230).
The aotivity-pH diagram for the Au(III)-d“-H20 system
(Figure 6.11) shows that Aud^ is the main species under acid conditions and that solubility is unlikely to be limited under the
considered practical conditions. AudJ is the most stable species
in acid conditions in the activity range of 0 .1 to 5 .0 and the
solubility is dictated by the concentration of d ” ions.
The thermodynamic calculations have shown the strong de- passivating and solubilizing effect of d ” ions on gold. Therefore
the chloride ion activity would be a very important process
122.
variable. The pH at which gold may be leached increases with
increased chloride concentration. There is a limited AuCl^ pre
dominance area, allowing gold to oxidize by a 1 electron reaction,
thereby requiring lower oxidant consumption. However, the stability of AuCl^ ions is limited and disproportionation to AuCl^ and Au may occur. The high standard potential for the AuClJ/Au couple
allows AuCl^ to be electrodeposited at very high current effici
encies. However, redeposition may occur if the potential conditions
are not closely controlled.
Cj 1 -T4CM U U
pH
Figure 6.11
Activity - pH diagram for the Au( II I)/H20-C1 system at 298 K, chloride activity of 1.0.
123.
6.2.2 Cyclic Voltammetry
Cyclic voltairmetry at rotating diso and planar platinum
electrodes was used to study the effect of Au(I)/Au(III) ratio,
pH, Cl“ concentration on reaction kinetics in the Au-G_-H20
system. Figure 6.12 shows superimposed cyalia voltammograms for
a gold-plated Pt disc in 1M HG. The disappearance of the cathodic
peak at a rotation rate of > 4 Hz, indicates that the anodically
produced gold-chloride species were soluble. The reduction peak
was obtained at 0.575 V vs SCE when the Pt diso was stationary.
Similar voltanmograms were obtained at pH 2 and 4.
Figure 6.13 shows the voltammogram for a solution with a total
gold concentration in 1 M H G of 8.23 x 10“4 M of which 6 .3 5 x 10-"4 M
was Au(III) and 1.98 x 10-4 M was Au(I), as determined by UV
spectrophotometry. From the rest potential of 0.646 V vs SCE, the
first anodic sweep showed no reactions due to gold species. At
potentials above 0.9 V vs SCE there was an indication of Pt oxide
formation. On the cathodic sweep a deposition peak was obtained at
0.46 V vs SCE. The reduction of AuG^ and AuG^ occurs
according to equations [6 .3 ] and [6 .6 ]
AuGg + e“ = Au(s) + 2 G “ [6.3]
E = 1.152 + 0.059 log (AuClg) - 0.118 log (Cl“) [6 .3a]AuClJ + 3e“ = Au(s) + 4 G “ [6 .6 ]
E = 1.001 + 0.0197 log (AuGJ) - 0.0789 log (Cl“) [6.6aJ
Substitution of the Au(I) and Au(III) concentrations into equations
[6 .3a] and [6 .6a] produces equilibrium potentials of 0.692 V vs SCE
for Au(I)/Au(s) and 0.696 V vs SCE for Au( III)/Au(s). These similar
124.
Figure 6.12
The effect of rotation rate on the current-potential behaviour of a gold plated Pt disc electrode in 1 kmol NaCl m~3, pH =5.4, sweep rate = 10 mVs-1, ( - ) stationary ( -- ) 4 Hz at295 K.
Figure 6.13
Cyclic voltammogram of a planar Pt electrode in a quiescent electrolyte containing 162.2 gm-3 total gold, Au(III) = 123.1 gm-3 + Au(I) = 39.1 gm-3, in 1 kmol HC1 m“3 , sweep rate =1 mVs-1 at 295 K.
125.
values may explain the ooourrenoe of only a single deposition peak.
On the second anodia sweep (Figure 6.13) two peaks were observed
at 0.775 V and 0.825 V vs SCE; these were investigated further
using ooulometry (Section 6.2.4) which suggested that the first
peak was due to the dissolution of gold as a gold( I)-chloro complex.
The second peak was gold dissolution as a mixture of gold(I) and
gold(III) chloro complexes.
Under agitated solution conditions (Figure 6.14), the two
anodic gold dissolution peaks merged. This behaviour may be due
to adsorbed AuCl, which could be further complexed with CL” ions
from solution and diffuse away from the electrode and/or oxidized
to a Au(III) chloro complex with an increased potential.
The second anodic peak (Figure 6.13) arose when AuClads
passivated the stationary electrode (Figure 6.13). With increased
potential part of the AuCla^g was further oxidized to Au(III) and
at the same time AuClac[s was further complexed and diffused away
from the electrode.
A similar mechanism has been postulated by Nicol(l88) aoo0rd-
ing to reactions [3.34 - 3.36],
Au = Au(I)ads + e“ [3.34]x Au(I)adg = x Au(III) + 2xe“ [3.35](1 - x) Au(I)ads = (1 - x) Au(I)aq [3.36]
where x is the fraction of Au(I) oxidised to Au(III). However,
NicolC168) was unable to observe two anodic peaks.
126.
POTENTIAL vs SCE/V
Figure 6.14
Cyclic voltammogram of a planar Pt electrode in an electrolyte containing 162.2 gm“3 total gold, Au(III) = 123.1 gm-3 +Au(I) = 39.1 gm~3, in 1 kmol HC1 m-3, moderate stirring, sweep rate = 1 mVs-l at 295 K.
Figure 6.15 Cyclic voltammogram of a planar Pt electrode in a quiescent electrolyte containing 51.4 gm-3 total gold, Au(III) = 1.2 gm~3 + Au(I) = 50.2 gm~3, in 3.9 kmol NaCl + 0.1 kmol HC1 m~3, sweep rate = 1 mV s- ̂at 295 K.
127.
The first anodic sweep of a voltammogram for a predominantly
Au(I) solution (Figure 6.15) showed only the onset of Pt oxide
formation. In the subsequent sweeps two peaks were observed, as in the previous solution (Figure 6.13). Similar results were
obtained in the pH range 0 to 4 and for chloride concentrations
of 1 to 5 M. These results refute the possibility that the oxida
tion of bulk AuCITj to AuClJ was responsible for the second peak.
The two anodic peaks appeared only when gold was pre-deposited on
the cathodic sweep, the solution was quiescent and the sweep rate
was very slow (1 mV s”l). The slow sweep rate required may explain
why other workers have not observed the two peaks previously.
Figure 6.16 shows the voltammogram of a 1 M Cl”, pH = 2
electrolyte, total gold concentration of 3.28 x 10”^ M and a Au(I)/
Au(III) ratio of 2.88. This voltammogram shows two deposition
peaks at 0.595 V and 0.437 V vs SCE. By substituting the gold(I)
and gold(II) concentrations into equations [6.3a] and [6 .6a], the
equilibrium potential for Au(I)/Au(s) is 0.6968 V vs SCE and for
Au(III)/Au(s) is 0.6788 V vs SCE. The reduction of AuCl^ to
AuCl^ according to equation [6.5],
AuClJ + 2e“ = AUCI2 + 2C1~ [6.5]
E = 0.925 + 0.029 log (AuClJ) - 0.029 log (AuClg) - 0.059 log (Cl')
[6.5a]
yields an equilibrium potential of 0.46 V vs SCE, which suggests that at least one of the two peaks is due to the reduction of AuClJ
to AuClp. If the same calculations are applied to the electrolyte used for the voltammogram shown in Figure 6.13, an equilibrium
curr
ent
dens
ity
/A
m”
128.
Figure 6.16 Cyclic voltammogram of a Pt disc electrode in a quiescentelectrolyte containing 64.7 gm-3 total gold, Au(III) = 16.7 gm“3 + Au(I) = 48.0 gnT3 , in 0.99 kmol NaCl + 0.01 kmol HC1 m“3 , sweep rate = 10 mV s-1 at 295 K.
129.
potential of 0.482 V vs SCE is obtained. The deposition peak in
Figure 6.13 was observed at 0.46 V vs SCE; this result supports
the hypothesis that the first deposition peak was due to the re
duction of AuClJ to AuCl^ ions.
Herrara-Gallego et al(l®6 ) observed two cathodic peaks that
were related to the deposition of AuClJ and AuCl^ ions. However,
no analysis for AuCl^ and AuCl^ concentrations were performed.
These two cathodic peaks (Figure 6.16) were investigated further
by steady-state reduction at a Pt rotating disc electrode. The
theoretical and experimental current densities vs square root of
rotation rate relationships have been plotted (Figure 6.17) for
fixed potentials of 0.595 V, 0.437 V and 0.0 V vs SCE. The maximum
possible rate of mass transfer to a rotating disc can be calculatedfrom the Levich equation(64)
I1>c = 1.554 n FA D2/ 3 W1/ 2 V- 1 / 6 C* [6 .8 ]
where n = number of electrons involved in the reaction F = Faraday constant, 96485 C mol“l A = Area, cm^
D = Diffusion coefficient, 1.3 x 10“5 cm2 S-1
C* = Concentration of species, mol cm~3
V = Kinematic viscosity, cm^ s“l
W = Rotation rate, Hz
The theoretical current density vs line [6 .8 ] for the
reduction of Au(III) to Au(I) coincided with the experimental
current densities vs obtained at an applied potential of0.594 V vs SCE (Figure 6.17). This result proves that the cathodic
1 30.
(rotation rate / H z ^
Figure 6.17
Steady-state reduction current density vs square root rotation rate at constant potential and theoretical currents in the same electrolyte as Figure 6.16. (1) Theoretical Au(III)reduce to Au(I), (2) Theoretical Au(I) reduce to Au(s),(3) Theoretical Au(III) reduce to Au(s), (□) 0.594 V vs SCE, (0) 0.437 V vs SCE, and (A) 0.0 V vs SCE.
131 .
peak at 0.594 V vs SCE in Figure 6.16 is due to the mass transport
controlled reduction of Au(III) to Au(I)surf . At an applied
potential of 0.0 V vs SCE, the experimental current density vs
w0.5 line is in excellent agreement with the theoretical line
resulting from the reduction under mass transfer-control of
Au(III) to Au(I) plus Au(I)buik to Au(s). This is congruent with
the detection of Au(I) species at the ring of a rotating ring-disc
electrode, while Au(III) species were being reduced at the disc
electrode(170) • however Au(I) detection by oxidation to Au(III)
species could not be carried out under mass transport control at
potentials lower than those causing co-evolution of chlorine.
At an applied potential of 0.437 V vs SCE, the experimental
line (Figure 6.17) can be explained by two possibilities:
i) The reduction of Au(III) to Au(0) was under mass transport
control and the reduction of Au(I)hulk was partially mass transport controlled. However, the fact that Au(I) was detected at the ring by Nicol(170) makes this mechanism unlikely.
ii) The mass transport reduction of Au(III) to Au(I)surf# and
the partial mass transport of Au(I)hulk "to Au(s).
This latter mechanism implies that Au(I)hulk raass transport controlled only at potentials lower than 0.437 V vs SCE and that the
reduction of Au(I)surf. a sl°w step.
The following mechanism is proposed to explain these results:
Au(III) + 2e + Au(I)surface [6.9]
Au ( I ) s u r f . * [6.10]Au(I)bulk
132.
Au(I)gurf ̂ + e“ + Au(s) [6.11]
Au(I)buik + e" Au(s) [6.12]
[6.9] was under mass transport control at potential < 0.594 V vs
SCE for the electrolyte investigated. [6 .10] is a mass transport
step independent of potential. [6 .1l] is a slow step, whereas
[6.12] is under mass transport control at 0.0V vs SCE.
In conclusion, the cyclic voltarimetry of gold in chloride
media in the pH range from 0 to 4 has shown that gold dissolves
readily in the presence of chloride ions. Two peaks can be observed
on the anodic sweep, provided the sweep rate is very slow (1 mV s~l)
and gold was deposited in the previous cathodic scan. A hypothesis
has been advanced to explain these results (Section 6.2). Two
peaks have been observed on the cathodic scan and a mechanism has
been proposed supported by steady-state measurements.
6.2.3 UV Spectrophotometry
Lingane^*^) has shown that AuCl^ ions do not absorb in theU.V. spectral region. AuCl^ ions absorb at 312 nm and 226 nm, though
the latter peak tends to shift to 228 nm with increasing gold
concentration. The molar absorptivity at 312 nm was found to be—1 2 —662.1 mol m . The lack of absorption by AUCI2 is shown in
Figure 6.18. Although the difference in the total Au concentration
for the two spectra is only about 2 g m“3, the absorption of spectra
of both solutions is proportional to their gold(III) concentrations.
133.
Figure 6.18UV absorption spectra of solution containing 1 kmol HC1 m-3( -- ) total Au = 16.6 gm~3, Au(III) = 12.6 gm“3 + Au(I)= 4.0 gm-3, ( --- ) total Au = 14.7 gm"3, Au(III) = 0.7 gm~3+ Au(I) = 14 gm“3.
Figure 6.19
Effect of time on the Au(I)/Au(III) ratio of unstirred solutions under nitrogen atomsphere, (0) PTFE, (A) PLASTIC and (0 ) GLASS beakers at 298 K.
134.
The disproportionation reaction [6.4]
3 AuClg -► AuClJ + 2Au + 2C1“; AG° = -43.713 kJ (mol AuClJ)
[6.4]
was studied following the procedure described in Section 5.5. Figure 6.19 shows that the Au(III)/Au(I) ratio tends to increase
with time for all three surfaces (PTFE, glass and plastic). Glass
seems to catalyse the disproportionation to a greater extent than
PTFE and plastic. This may be due to the adsorption of Au on
glass(232) providing catalytic sites. A K value, defined as
(AuCl^)(AuClJ) (Cl“)̂ , of 1.15 x 10“® was found for the glass
surface which compares well with a reported value of 1 .0 x 10”®
by Lingane(231). The values of K are 9.72 x 10”® and 2.04 x 10”® on PTFE and plastic, respectively. Gold(I) solutions may require
> 2 month to reach equilibrium. This indicates that for most
practical purposes the disproportionation reaction may be negligible .
The oxidation of AuCl^ to AuClJ by Cl2 was studied using
the stop-flow apparatus shown in Figure 5.3. A gold solution of
1.78 x 10”® M Au(III) and 3.55 x 10”® M Au(I) was mixed with
1.8 x 10-^ M and 9 x 10“^ M CI2 in 1 M HC1. Spectra were taken
for different experiments at intervals of 10, 30, 60 and 600 seconds.
The spectra were deconvoluted using a multicomponent analysis program available for the Hewlett-Packard Diode Array Spectro
photometer, Model no. 8451 A. The reduction of chlorine is governed by:
[6.13]
[6.13a]E = 1.395 + 0.029 log (Cl2) - 0.029 log (Cl-")
and for the oxidation of AuC12 to AuClJ:
AuCl^ + 2e = AuC12 + 2C1“ [6.5]
E = 0.925 + 0.029 log (A11CI4 ) - 0.029 log (AuC12) - 0.059 log (Cl“)[6.5a]
For a CI2 concentration of 9 x 10”^ M, ^q \ 2 / C Y ~ = 1-306 V. For
the gold solution e auC14/AuC12 = 6.887 V, therefore 0.419 V was available as a driving force. However, no oxidation of AuC12 was
_Qobserved. When the Cl2 concentration was increased to 1.8 x 10 M,
the multicomponent analysis programme was unable to deconvolute
the spectra. These results suggest that the Au(III)/Au(I) couple
does not reach equilibrium quickly. Therefore the use of a Pt indicator electrode to measure the Au(III)/Au(I) ratio(180) ±s not
valid.
Cl2 + 2e = 2C1"
The disproportionation reaction [6.4] was studied and found
to be very slow. Also the oxidation of AuC12 to AuCl^ by
9 x 10”^ M Cl2 appears to be slow. However, an independent method
to confirm this latter result would be desirable.
6.2.4 Electrochemical Dissolution of Gold
The aim of these experiments was to determine the relative
yield of gold(I) and gold(III) and the effect of chloride ion and proton concentration on the Au(I)/Au(III) ratio at an applied
potential.
136.
The relationship between applied potential and current density
is described by the Butler-Volmer equation(233):
co(0,t) e““nfn C^O.t) e d - ‘) nfn
Q>* C r *
[6.14]
where i = current density, A nr 2
io = exchange aurrent density, A m~2
Cb(o,t), CR(o,t) = concentration of oxidized and reducedspecies at electrode surface, M
Co*, Q}* = concentration of oxidized and reducedspecies in bulk solution, M
oc = transfer coefficient
n = electrons per molecule oxidized or reducedf = F/RT, V"1
n = overpotential, (E-Eq), V
Eq = equilibrium potential, VE = applied potential, V
In the absence of mass transfer effects the Butler-Volmer equation
may be simplified to the Tafel relationship(233).
[6.15]
If a = 2.3 RT log io, and «nF
b = -2.3 RT «nF
n = RT In io - RT In i «nF anF
then a and b are the Tafel constants in the Tafel equation [6.16].
n = a + b log i [6.16]
A 7 “7\ J ! .
Tafel plots are widely used to evaluate kinetics parameters.
The difficulties of applying Tafel plots to the AU-CI-H2O is that
gold dissolves to a mixture of gold(I) and gold(III) and the Au(I)/
Au(III) molar ratio is dependent on the applied potential.
Average current density vs potential was determined for pH
values of 0, 1, 4 and 6. Figure 6.20 shows the curves for pH’s of
0 and 1; results at other pHs were omitted for clarity. The
deviation from linearity at 1.0 V vs SCE was probably due to Cl" ion depletion and the onset of oxide formation(109, 166). The
Tafel slopes are shown in Table 6.3.
TABLE 6.3 - Tafel Slopes for AU-CI-H2O System at Different pHs
pH = 0 80.5 mV/decadepH = 1 6 6 .0 mV/decade
pH = 4 55.0 mV/decadepH = 6 77.0 mV/decade
Nicol(108) found Tafel slopes of 60 to 80 mV per decade at
pH < 2. Assuming a symmetry factor of 0.5, these slopes are
consistent with a two electron slow transfer for Au oxidation as
the rate-determining step.
The current contributions due to gold oxidation to Au(I) and
Au(III) species can be resolved if gold electrodes are anodised at
100 % current efficiency. It was found that at pH = 1 and an
applied potential of 0.8 V vs SCE in 1 M Cl", a purely gold(I) solution can be generated. This experiment was repeated 10 times
CURR
ENT
DENS
ITY /A
m"
13S.
Figure 6 .2 0
Steady-state current vs potential curve for a gold electrode moderate stirring, (Q) 1 kmol HC1 m~3 , (0) 0.9 kmol HC1 +0.1 kmol NaCl m~3.
Figure 6.21
Partial currents vs potential curve for a gold electrodestirring, (Q) 1 kmol H d m-3, (0) 0.9 kmol HC1 + u.i kmol NaCl
1-0
139.
and an average current efficiency of 100.2 ± 1.45 % was obtained
for Au(I). Consequently, if the same experimental technique is
used at other potentials less than passivation potentials, it is
reasonable to assume a 100 % current efficiency for gold dissolution.
The partial average current densities vs potential for Au(I) and
Au(III) are shown in Figure 6.21 for pHs = 0 and 1. The curves
for pHs = 4 and 6 were omitted for clarity. The Tafel slopes for
Au(I) and Au(III) at different pH’s are shown in Table 6.4.
TABLE 6.4 - Tafel Slopes for Au(I) and Au(III) at Different pHs
pH Au(I), mV/decade Au(III), mV/decade0 90 48
1 89 35
4 100 58
6 103
Anodic transfer coefficients of 0.71 and 0.36 have been
reported(158) for Au(I) and Au(III). The validity of these trans
fer coefficients can be examined by considering the cathodic transfer
coefficients, since the addition anodic and cathodic transfer
coefficients should equal 1. The higher overpotential necessary
to deposit Au from Au(I)-Dujk than to reduce Au(III) -► Au(I)surf
(Section 6.2.2) suggests that the cathodic transfer coefficient
for Au(III) is larger than for Au(I), which is consistent with the anodic transfer coefficients reported previously(158). Using
these transfer coefficients and substituting for b (equation [6.15])
the Tafel slopes should be 83 and 49 mV per decade for Au(I) and Au(III), respectively.
140.
Considering that coulometry value has a relative error of
about 3 %, the Tafel slopes in Table 6.4 are in reasonable agreement with the results of Heumman and Panesar(158) . Figure 6.22 shows the
effect of CL“ ion concentration on the relative current yields due
to gold(I) and gold(III), at pH = 1 and potential = 0.8 V vs SCE.
At this pH and potential, gold(I) is the main dissolution product
at all Cl“ ion concentrations studied. Figure 6.23 shows the
effect of HC1 concentration at a potential = 0.8 V vs SCE on the
relative current yields due to Au(I) and Au(III). The overall
dissolution rates are higher than in Figure 6.22. The main dis
solution product in HCL concentration range studied is gold(I).
The increase in overall current at pH < 0 was not due entirely to
an increased dissolution rate as has been suggested^??), but
was due partially to a higher proportion of gold(III) generated.
Tafel slopes of 55 to 80 mV per decade were found in the over
all gold dissolution current vs potential curve for pH 0 to 4.
Assuming a synsnetry factor (3) of 0.5 this is consistent with
a mechanism having a two electron transfer in the rate of deter
mining step (Section 6.2.2, equations [3.34] to [3.36]). The main
dissolution product, at all pH's investigated and potentials of
practical use, was Au(I). The Tafel slopes changed with pH in
dicating a dependence of the kinetics on the proton activity.
6.2.5 Mixed Potential Leach ng RatesThe dissolution of gold in the presence of chlorine may occur
by one of the two following reactions:
LO
G
CU
RR
EN
T
DE
NS
IT
Y /
Am
-
141 .
Figure 6.22 Effect of Cl“ concentration on the generation of Au(I)/Au(III) species at a potential = 0.8 V vs SCE and pH = 1.
LO
G
CU
RR
EN
T
DE
NS
ITY
/
Am
"
142.
Figure 6.23 Effect of HC1 concentration on the generation of Au(I)/ Au(III) species at a potential = 0.8 V vs SCE.
143.
au + 0 .5 d 2 + a - = auci2 [6 .17]
AE/V = 0.243 + 0.0291 log(Cl2) - 0.0591 log(Aud2) + 0.0891 log(Cl“)
[6.17a]
or
Au + 1.5 a 2 + Cl" = Au C12 [6.18]
AE/V = 0.395 + 0.0291 log(Cl2) - 0.01971 log(Audp + 0.0494 log(Cl")[6.18a]
The leaching rates were determined using the cell shown in Figure 5.2
and the procedure described in section 5.6. The chlorine concen
tration in solution was determined using a Pt indicator electrode.
The potential of the indicator electrode is governed by the Nernst
equation:
E = E° + RT In oxidized species] [6.19]nF [reduced species]
This method can be used provided:
i) The potential-determining couple equilibrium is achieved
rapidly.
ii) Other possible potential-determining couples are absent, or at
mutual equilibrium.
The sensitivity of this method depends on the d 2/Cl“ molar
ratio, as this ratio tends to 1, the indicator electrode potential tends to the standard potential of the d 2/ d ” couple. However,
in the majority of experiments to be reported in this section, the
d " ion concentration was about 500 times higher than d 2, so avoiding this problem.
'144.
Figure 6.24 shows two typical potential vs C^/Cl- molar
ratio curves. As expected from the Nernst equation, the curves are
of exponential form. A potential vs CI2 concentration (determined
by titration) curve was derived before each leaching rate experiment.
The effect of Cl” ion concentration, pH, rotation and chlorine
concentration were studied at a gold-plated rotating platinum
disc. The linear dependence of leaching rates on Cl” ion con
centration at pH = 4, HC1 concentration at constant chlorine con
centration and rotation rate is shown in Figure 6.25. As in the
electrochemical dissolution rates (Section 6.2.4), the proton
concentration increased the leaching rate by a factor of two with
increasing HC1 concentration from 1 M to 5 M at constant oxidant
concentration.
The marked effect of proton concentration on the leaching
rate could be explained partly by the CI2 activity vs pH diagram
(Figure 6.26). At pH < 1, chlorine is the only oxidizing agent in
2 mol m”3 dissolved chlorine, whereas at pH = 4, HC10 species
predominate. The diffusion coefficient of HC10 is less than for chlorine(234) ancj the activity of concentrated HC1 is higher
than concentrated NaCl(229,230)# The corrosion potentials (Figure
6.25), after t = 3600 decreased with increased Cl” ion and HC1 concentration.
There was a linear dependence between leaching rates and chlo
rine concentration (Figure 6.27) in the concentration range 1 -_o3.25 mol CI2 m . The theoretical leaching rates calculated
POT
ENT
IAL
vs
S.H
.E./
V
145.
cviGZD<
cn■c
10
HC1 m~3
0-78<->0—j
0t/i0-73 03
TDO0-68 it>
3
53063 <
in
l/lnm\<
concentration / kmol m-3
Figure 6.25 Effect of Cl~ ion and HC1 concentration on corrosion potentials and the mixed potential leaching rate of a gold- plated Pt rotating disc electrode, at 4 Hz rotation rate, 2 mol CI2 ni at 295 K.
146.
Figure 6.26 Activity - pH diagram of chlorine/hypochlorite species at 298 K and total concentration of 2 mol m~3 in 1 kmol Clm“3.
chlorine concentration / mol m-3
Effect of chlorine concentration on the mixed potential leaching rate of a gold-plated Pt rotating disc electrode at 295 K, 4 Hz rotation rate in 1 kmol HC1 nr3 (Q). Theoretical leaching rates, assuming a 3 electron reaction (g) and a 1 electron reaction (□>. calculated from the Levich equation for mass transport controlled chlorine reduction.
147.
from the Levich equation [6.8] for a 1 and 3 electron reaction
were also plotted, showing that if the dissolution reaction is a 3
electron process then it is mass transport controlled. If the
dissolution reaction is a 1 electron process then it is under
partial kinetic control. The same conclusions can be drawn from the relationship between square root of rotating disc electrode
(RDE) rotation rate and leaching rates at pH’s = 0, 2 and 4 (Figure
6.28). The theoretical lines, derived from the Levich equation [6.8],
have been plotted only for pH = 0 in Figure 6.28.
To summarize, increased proton concentration increased the
leaching rates at constant CI2 concentration. An increase in Cl™
ion concentration also increased the leaching rates, but to a
lesser extent. There was a linear dependence between leaching
rates and both chlorine concentration (Figure 6.27) and square
root of RDE rotation rate (Figure 6.28). This may suggest a 3e“ mass transport controlled dissolution reaction, although some
evidence suggests the possibility of a partially kinetic controlled
le” reaction.
6.2.6 Chlorine Reduction on Gold Surfaces
The reduction of chlorine provides the driving force for the gold oxidation reaction. The net current method was used tu
determine whether chlorine and hypochlorous acid are reduced under
transport control on gold surfaces. Resultant gold dissolution
rates were determined, and compared with those obtained using the
mixed potential method (Section 6.2.5).
148.
(rotation rate / H z^
Figure 6.28
Effect of rotation rate and pH on the mixed potential leaching rate of a gold-plated rotating disc electrode at 295 K, constant chlorine concentration (2 mol m-3) in l kmol Cl- , (0) pH = 0, (A) pH = 2, (t) pH = 4. Theoretical leaching rates assuming a 3 electron reaction (g) and a 1 electron reaction (□). calculated from the Levich equation for mass transport controlled chlorine reduction.
149.
The anodic and oathodio net currents for the oxidation of gold
and the reduction of 2 mol CI2 m“^ at pH = 0, on a gold-plated
platinum disc electrode, are shown in Figure 6.29. A chlorine
reduction transport-controlled plateau was observed at potentials
< 0.4 V vs SCE, for which a linear dependence relationship was
obtained between current density and the square root of RDE rotation
rate. The correlation with the theoretical values derived from
the Levi oh equation [6.8] was good, indicating that an overpotential
of about 0.7 V was necessary to obtain mass transport control
reduction of Cl2 at pH = 0. The corrosion currents were about one sixth of the mass transport reduction current.
Similar results are shown in Figure 6.30 for pH = 4. Two
plateaux were observed, one at potentials <0.3 V vs SCE and the
other in the potential region 0.4 V to 0.5 V vs SCE. Figure
6.31 was generated by plotting the theoretical values obtained
from the Levi oh equation [6.8] for the reduction of 2 mol Cl 2
m“3 at pH = 0 and pH = 4 and the experimental values of the two
plateaux observed. The difference in the two theoretical lines
arises from the pH dependent solution' composition (Figure 6.26) ;
the diffusion coefficients for chlorine and HC10 are 1.7 x 10”®m2 S-1 and 1.29 x 10”® m2 s“l respectively (234) # These results
(Figure 6.31) are contrary to those expected, since the plateau at
potentials < 0.3 V vs SCE seems to correspond to Cl2 reduction
and the other plateau to the reduction of HC10. The process is
complex as there is a chemical reaction [6.20]
cl2 + h2o H0C1 + H+ + o r [6 . 2 0 ]
150.
_r>Reduction currents for 2 mol Cl^ m at different rotation rates on a gold-plated Pt rotating disc electrode and oxidation currents of Au in 1 kmol HC1 m”^ at 295 K obtained using the net current method, sweep rate = 10 mV s~l.
Reduction currents for total chlorine species of 2 mol m~3 at different rotation rates on a gold-plate Pt rotating disc electrode and oxidation current of gold in 1 kmol Cl- + 0.9999 kmol Na+ +0.1 mol H4" m“3 at 295 K obtained using the net current method, sweep rate = 10 mV s~l.
151-
(rotation rate / Hz^
Figure 6.31
Experimental reduction currents for total chlorine species of 2 mol m-3 at a potential of 0.2 V vs SCE (□), 0.5 V vs SCE(0). Theoretical reduction current at pH = 0 (1), and pH = 4 (2) in 1 kmol Cl“ + 0.9999 kmol Na+ + 0.1 mol H+ m“3 at 295 K form the Levich equation assuming mass transport controlled.
(rotation rate / H z ^
141312 O11 C
310 m
Z3
9 CLrt>Z3
8 (/)7 \6
J>5 34 ro32
Figure 6.32Effect of rotation rate and pH on leaching rates obtained by the net current method, total chlorine species concentration = 2 mol m-3 in 1 kmol Cl" m-3 at 295 K, (0) pH = 0, (A) pH = 2, and (0) pH = 4.
152-
which is dependent on the local pH, and coupled to an electrode
reaction [6.21J:
The reduction of HC10 increases the local pH in the diffusion
cathode reaction layer and the anions concentration. Electro
neutrality would require the transport of cations (H+ and Na+)
towards the reaction/diffusion layer, changing the local pH and
favouring CI2 formation by reaction [6.20]. As the potential
was increased to 0.4 V vs SCE, HC10 rather than chlorine was reduced
under mass transport control. However, this hypothesis requires
more experimental evidence to test its validity.
The corrosion currents and corresponding leaching rates for
pH's 0, 2 and 4 are shown in Figure 6.32. The leaching rates were calculated assuming a le“ dissolution reaction(64). These
leaching rates are of the same order (10-3 mol Au m~2 s-*) as
the mixed potential leaching rates (Figure 6.28), although in
general about 10 % higher. Considering the different time scale
of the experiments and that the mixed potential method has no
associated uncompensated iR potential drop, the results are in
acceptable agreeement. This suggests that the dissolution product is mainly a Au(I)-chloro complex.
The true kinetic current of an irreversible reaction can be found by plotting i~l.vs W~0-5 according to the following equation
HC10 + 2e“ = Cl“ + 0H“ [6 .2 1]
(235).
1 1 1.613 v1/ ^ " 2/3 vr0 *51 FK-l c F C
+ [6.22J
153.
where i = current density, A m~2
F = Faraday constant, 96485 C mol“l
Ki = rate constant, cm s"^
c* = concentration of electroactive species in bulk solution, mol cm-2
V = kinematic viscosity, cm2 s~l
D = diffusion coefficients, cm2 s“l
W = rotation rate, Hz
and extrapolating the infinite rotation rate (W“0*5 -► «), the true
kinetic current is equal to nFKqC*. The rate constant Kq can
not be obtained when the net current method is used, because Kq
varies with the potential and the corrosion potential varies with
the rotation for this case. However, the true kinetic currents
are a measure of the inherent electrochemical kinetics in the
absence of mass transport restrictions. Plots of i“l vs W“0*5
were drawn for the data derived from the corrosion currents and
mixed potential leaching rates for the three pHs and the true
kinetic currents were determined. Figure 6.33 shows the true
kinetic currents for both sets of data, a marked decrease of kinetic
current was observed with increasing pH. These results are in
agreement with the conclusion that an increase in proton concen
tration increases the dissolution rate (Section 6.2.4 and 6.2.5).
It would be desirable to operate the proposed process at pH’s
> 4, since this would decrease the rate of some side reactions
(section 6.2.9) and less corrosion-resistant plant materials could be used. Therefore gold passivation under a high pH and high HC10 concentration leaching condition was tested. An experiment was
15'4.
Figure 6.33 True kinetic current at infinite mass transport for mixed potential leaching rates data (0) and net current method ( □ ), total chlorine species concentration of 2 mol m"3 in 1 kmol Cl~ m-3 at 295 K.
Figure 6.34 Effect of rotation rate on leaching rates obtained by the net current method, 10 mol HC10 + 1 kmol Cl“ + 5 mol C00H.C6H4C00K + 3.4 mol NaOH m“3 at 295 K.
155.
performed at pH = 5.4 (buffered(236)) and io mol m”3 HC10.
This pH was chosen because HC10 is the main specie present. The
corrosion currents and respective leaching rates (Figure 6.34)
showed that although the concentration of the oxidant was increased
five-fold, the leaching rates were very similar to those obtained
at pH = 4, indicating some degree of gold passivation.
In summary, there was good agreement between the leaching
rates derived by the mixed potential and net current methods.
Further evidence was provided to suggest that the gold dissolution
reaction is a le process. Corrosion currents were about one sixth
of the mass transport corrosion currents for CI2 and HC10, indicating
a mixed control dissolution reaction. However, the high true
kinetic currents in acidic pH's indicated that by increasing CI2
concentration and agitation, high leaching rates can be achieved.
Gold was passivated to some extent during leaching at pH = 5.4 and 10 mol HC10 m"2.
6.2.7. Electrodeposition of Gold in a Packed Bed Electrode
The electrowinning of metals from dilute solutions cannot be
carried out economically using classical electrolytic cells without
some pre-concentration stage, since the current density for de
position would be extremely low. Three-dimensional particulate electrodes have been investigated extensively, for applications
such as electrowinning, effluent treatment, fuel cells, batteries,
and organic electrosynthesis(®»237)# These electrodes, either in
the form of packed, slurry or fluidized beds, have been considered
156.
because of their high specific surface areas and mass transfer
rates, which makes them attractive systems for a number of electro
chemical processes, especially those with very low operating
current densities, i.e. low concentration of electroactive species.
Figure 6.35 shows the consequence of the instability to
reduction of AuCl^ ions, and possibly of their electrostatic
adsorption on flow circuit surfaces (uPVC, carbon/graphite etc)
which were likely to have been positively charged due to protonation
of surface groups. The propensity of Au(CN)2 anions to adsorb on carbon surfaces is exploited in the carbon-in-pulp process(238) ̂as a means of concentrating and purifying gold leach solutions, prior
to their reduction to elemental gold by electrowinning in Zadra type cells incorporating steel wool cathodes(239)# With dicyanoaurate
ions the adsorption is complete at well below monolayer levels (240) .
Figure 6.35 shows there was very significant depletion of AuCl^
ions from solution by the previously unused carbon packed bed with
no applied potential. The phenomenon has been exploited for the
recovery of gold from chloride solutions(193) ̂ though probably
at specific rates and gold loadings significantly lower, and to
concentrations greater, than potentially achievable by electrowinning in three dimensional electrodes.
Figure 6.36 shows the flow rate dependence of the exponential
decay of currents and dissolved gold concentrations by the re
circulation of the electrolyte through the packed bed electrode (PBE), the potential between the feeder electrode and solution at the membrane being controlled potentiostatically at 0.529 V.
157.
10"4 T IM E / s
Figure 6.35 Depletion of total dissolved gold by adsorption on flowcircuit surfaces (A), and unused carbon bed particles (t) of area 0.082 m^.
158.
Figure 6.36
Total dissolved gold concentrations (open symbols) and current densities (solid symbols) as functions of time and flew rate. Feeder electrode-membrane potential 0.53 V, bed of 3 mm cylindrically-shaped carbon chips, flow rates (O',-^) 1.9 x10“6 m3 s-1, (0 ) 8.4 x 10-6 m3 s"1, ( © , # ) 16.2 x 10’m3 s_1.
v-6
lAu(I)] ♦ lAu(ffl)] / gm
159.
Although the lowest gold concentration shown in Figure 6.36
is 3 g levels below the detection limit (say 0.1 g m~3) of
atomic adsorption spectrophotometry (AAS) were achieved routinely.
Initial decay rates were sensitive to the history of the bed,
reflecting the changing area for reaction on the carbon particles.
The concentration decay has the general form of the steady-state
stirred tank reservoir, plug flow reactor model equation(241);
c(t) = o(o) exp {-t/x (1 - exp (-kAaL/Q))} [6.23]
where c(t)
o ( o )
tT
k
A
a
L
Q
concentration as a function of time, mol m“3
initial concentration, mol m~3
time, s
residence time (V/Q), s
mass transfer rate constant, m s-1
cross-sectional area of reactor, m2
specific wetted surface area of bed electrode, m~l
eleotrode length, m
volumetric flow rate, m3 s~l
though as discussed below the reaction mechanism is more complex
than a simple mass transport controlled process assumed by the
model and the reactor/reservoir and dead volumes were such that
the application of equation [6.23] was not strictly legitimate for
the conditions used.
However, using such a model as a first level of approximation,
then the inlet (o-̂ ) and outlet (c0) concentrations are related by the equation(241);
160.
°o = °i exP (-kAaL/Q) [6.24]
where oQ = outlet concentration, mol m“^
cL = inlet concentration, mol m”^
From determinations of values of cj_ and cQ of samples taken during
depletion experiments, average values of the product (ka) were
determined as 0.035, 0.094 and 0.150 s_l for the three flow rates
specified in Figure 6.36. These are a factor of 80 greater than
calculated from the projected surface area of the graphite chips
and mass transport coefficients (k) derived from the correlation for packed beds(242) assuming a voidage (e) of 0.5:
Sh = {(l-e)0-5 / (l-e)} Re0«5 ScO-33 [6.25]
where Re = u dp [6.32]v
So = v .[6.33]D
This discrepancy may be attributed to an underestimate of the area
of the carbon chips, which appeared macroporous to the naked
eye. Assuming the derived values of k from equation [6.25] to
be correct, then the effective (wetted) area of the packed bed of
particles for a mass transport controlled process, was 7.3 m2.
In spite of the additional complexities of the gold deposition
process discussed below, substitution of ka values derived from
equation [6.24] gave total dissolved gold (Au^) concentration/time data whiah were in reasonable agreement with the experimental data
given in Figure 6.36, though with greater discrepancies at lower flow rates, as expected(241).
161.
The linear current density vs. concentration relationship
shown in Figure 6.37 is supporting evidence that the reduction
process(es) were transport controlled, and as the lines passed
through the origin, this implied high current efficiency even at
very low concentrations. Analyses of solution samples taken during
depletion experiments showed (Figures 6.38 - 6.40) that while the
total dissolved gold concentrations (Aup) decayed exponentially with
time (Figure 6.36) the Au(I)/Au(III) molar ratio also decreased
initially from a value of 0.35 - 0.25, depending on the particular
solution used, the former being the equilibrium values given by
equation [6.4]:
3 AuClg * AuCLJ + 2 Au + 2 CL" [6.4]3 log (AuClg) = -7.66 + log (AuClJ) + 2 log (Cl") [6.4a]
That ratio then increased, passing through a maximum value
which increased with flow rate, before decaying to zero at long
times, though prior to the dissolved gold (Au^) being totally
depleted. This behaviour was particularly pronounced at the highest
flow rate used (Figure 6.40) at which the Au( I)/Au( III) molar
ratio showed a sharp peak after 700 s, when the gold total dis
solved gold had decreased to < 0.1 mol m-3 (19.7 g m~3).
As the solutions contained both Au(III) and Au(I) species, a
figure of merit of Faradays per mole of gold deposited was used,
rather than the more usual current efficiency (%). The data for
F (mol Au)“l corresponding to the depletion results in Figure 6.36,
are given in Figures 6.38 - 6.40. These show an increase from
162.
Total dissolved gold concentration dependence of the PBE cross-sectional current density. Feeder electrode-membrane potential 0.53 V, bed of 3 mm cylindrically-shaped carbon chips, flow rate 1.9-x 10“^ 'nr* s~* (0), 8.4 x 10“^ m3s"1 (Q), 16.2 x K T 6 m3 s'*1 ($).
Total dissolved gold concentration ( g ), [Au(I)]/[Au(III)] molar ratio ( % ) , and incremental (□). and cumulative (O), Faradays per mole of gold deposited in the PBE operating under conditions specified in Figure 6.36 («o).
163.
Figure 6.39
Total dissolved gold concentration (J|), [Au(I)]/[Au(III)] molar ratio ( % ) , and incremental (□), and cumulative (0), Faradays per mole of gold deposited in the PBE operating under conditions specified in Figure 6.36 (0).
164.
>c
>c
to3iCO
Figure 6.40
Total dissolved gold concentration (|), [Au(I)]/[Au(III)] molar ratio (£), and incremental (□). and cumulative (O), Faradays per mole of gold deposited in the PBE operating under conditions specified in Figure 6.36 (O).
165.
a short time (< 200 s) value of about 1 F (mol Au)-1 (deposition
via reaction [6.12], with the cumulative values reaching a plateau
of about 3 (deposition by reaction [6.6]) at long times. The
incremental F (mol Au)”l values reflected (Figures 6.38 - 6.40) the
Au(I)/Au(III) molar ratio data, particularly in the region of the
peak, which presumably arose from the production of Au(I) species
by reduction reaction [6.9], and decayed due to their deposition
by reaction [6.12], which required 1 F (mol Au)~l.
As an applied potential of + 0.529 V between the feeder
electrode and solution at the bed/membrane interface precluded
hydrogen evolution, the extra Faradaic requirement (Figure 6.40)
above the value of 3 F (mol Au)~l for reaction [6.6] was probably due to the reduction of oxygen. The anolyte would have been super
saturated with dissolved oxygen and recently published data(24:3)
showed two grades of Nafion membrane to have high solubilities
and diffusion coefficients for oxygen. Adventitious ingress
through the various joints in the flow circuit would have con-
sistituted a secondary source of dissolved oxygen.
Another set of experiments was performed in the -200ym
carbon bed, which had a surface area of 39.2 m2 determined
by single point nitrogen absorption. The flow rate dependence
of the exponential decay of current and dissolved gold concentration
(Figure 6.41) for the -200vim carbon bed was very similar to that for the 3 mm aylindrically-shaped carbon bed.
166.
Total dissolved gold concentrations (open symbols) and current densities (solid symbols) as functions of time and flow rate. Feeder electrode-membrane potential 0.53 V, bed of -200 ym carbon chips, flow-rates ( ^ , ^ ) 1.9 x 10“^ m3 s“l, ( 0 , # )8.4 x 10“° itH s-^flT]^) 16.2 x K T 6 m3 s"1.
10̂ flowrate / m-^s^Figure 6.42
Gold concentration decay rate for bed of 3 mm cylindrically- shaped carbon chips (Q) and bed of -200 ym carbon chips (A), initial total dissolved gold = 100 gmr3 , feeder electrodemembrane potential 0.53 V, 1 kmol HC1 m“3.
167.
The two beds were compared by plotting the deposition rate
d log C/dt vs flowrate (Figure 6.42). The 3 nm cylindrically-shaped
carbon bed has a higher deposition rate than the -200 ym carbon
bed. This result may be explained by the different physical charac
teristics of the beds. The -200 ym aarbon bed had smaller particle
sizes and hence lower mean Reynolds [6.32] and Sherwood [6.25] num
bers, the latter being proportional to the mass transfer coefficient.
Furthermore, the effective surface area is a more significant
parameter than the total surface area, since area in certain pore
size ranges may not play any part in the electrodeposition process.
Experiments were run at initial concentrations of 15 g m~3
(Figure 6.43) since this value is typical of the envisaged end use.
The results show very similar behaviour to that for high initial
concentrations, but the deposition rate was faster (Figure 6.44).
The effect of initial concentration on the exponential decay of
current and dissolved gold concentration vs time (Figure 6.45) shows
that the three lines are not parallel. This might be due to the
different initial Au(I)/Au(III) ratio of the different solutions,
or changing area with time. The more cathodic the applied potential,
the faster the deposition rate and at -0.2 V vs SCE, hydrogen evolu
tion occurred. The Au(l)/Au(III) ratio was similar to that at
less oathodia potentials.
An increase in pH decreased the deposition rate (Figure 6.46)
indicating some form of kinetic control or at certain pHs the
formation of gold oxides (Figure 6.11). At pH = 11.7 the species
168.
Total dissolved gold concentrations (open symbols) and current densities (solid symbols) as functions of time and flow-rate. Feeder electrode-membrane potential 0.53 V bed of -200 pm carbon chips, flow-rates (o,-^) 1.9 x 10“^ m3 s-1r (O , #)8.4 x 10”^ m 3 s “*> ( D . H ) 16.2 x 10”® m“3 s-l.
10̂ flowrate /
Figure 6.44
Gold concentration decay rate for bed of -200 pm carbon chips, initial total dissolved gold concentration = 15 gnr3, feeder electrode-membrane potential 0.53 V, 1 kmol HC1 m~3.
169.
>c
>c
i O
3iGO
Figure 6.45
Total dissolved gold concentrations(open symbols) and current densities (solid symbols) as functions of time and initial dissolved gold concentration. Feeder electrode-membrane potential 0.53 V, bed of -200 pm carbon chips, flow rate16.2 x 10“^ m3 s“l, intial dissolved gold concentrations
100 g™-3’ 75 S™-3 ( O . # ) 15 gm-3 in 1 kmolHa nr 3.
170.
pH
Figure 6.46
Gold concentration decay rate as a function of pH, feeder electrode-membrane potential 0.229 V, initial dissolved gold concentration = 100 gm"3 , flow-rate = 16.2 x 10"^ m3 s-*, bed of -200 \ im carbon chips, in 1 kmol Cl" m"3. Dashed line indicates possible solubility problems.
171 .
that was reduced was probably Au(OH)J since when the gold was added to the solution, the solution changed colour to purple. The
dashed line in Figure 6.46 indicates solubility problems.
As the AE° values for reactions [6.17] and [6.18] are only
0.243 and 0.395 V respectively, typical operating voltages for the cell were about 1 V, even with a large anode-membrane gap (Figure
5.5). The cumulative specific charge requirements of about 3 F
(mol Au)“l, shown in Figures 6.37-6.40, correspond to a specific
energy requirement of 400 kWh (tonne Au)“l for gold electrowinning
at a cell voltage of 1 V.
The behaviour of the system could be explained by the follow
ing explanation:
Au(III) + 2e Au( I)surface [6.9]
Au(I)surface * Au(I)bulk [6.10]
Au(I)surface + e * Au(s) [6.11]
A u W b u l k + e + Au(s) [6.12]
3Au(I) -► 2Au + Au(III) [6.26]
The rapid initial decay of the Au(I)/Au(III) molar ratio, and the
initial F (mol Au)“ ̂ values of 1-3, imply that while reaction [6.12] was transport controlled, as has been reported(169,170) ancj as shown
(Section 6.2.2) for a solution containing predominantly Au(I) spe
cies, reactions [6.9] and [6.ll] were only partially so, for the
range of potentials applied. Having depleted the Au(I) species
from bulk solution to concentrations well below their equilibrium
value given by equation [6.26], reaction [6.9] became increasingly
significant, causing a rise in the Au(I)/Au(III) molar ratio,
172.which increased with increasing flow rate due to mass transport
dependent dispersion process [6.10]. However, depletion of the
Au(III) species, causing decreasing rates of formation of Au(I)
species by reaction [6.9J, and the transport controlled removal
of Au(I) species by reaction [6.12], would cause the Au(I)/Au(III) molar ratio to pass through a maximum and decrease to zero, prior
to total depletion of the total dissolved gold, as observed.
It is not immediately apparent why the further reduction of
Au(I) species formed from Au(III) species (reaction [6.11 ]) should be comparatively slow, enabling their dispersion to the bulk solu
tion (step [6.10]), while reaction [6.12] is fast. Nicol(170) found a chloride ion reaction order for reduction of AuClJ ions to be -1,
but did not determine a value for the reduction of AuCl^ ions.
If a similar inverse order were operative, the high local chloride
ion concentration resulting from the reduction of AuClJ ions would
inhibit the further reduction of the Au(I) intermediate, whereas the
reduction of (predominant) AuCl^ i°ns from the bulk electrolyte could still be fast. If there was insignificant potential drop
in the particulate graphite phase, so that the potential applied
appeared near the bed/membrane interface, then the remaining bed
volume would have operated as lower overpotentials(®). This would
have had the effect of further favouring the faster of the elementary steps in the overall process. At lower applied potentials than
used in the experiments reported here both Au(I) reduction processes
may be fast.
In summary, the feasibility of electrowinning gold-chloro complexes in a packed bed electrode at high current efficiencies
173.
has been established. An explanation was provided for the behaviour
of the Au(I)/Au(III) molar ratio vs time which is consistent with
the results of Section 6.2.2. The concentration decay has the
general form:
c(t) = c(o) exp {-t/x(l - exp (-kAaL/Q))} [6.23]
though as discussed the reaction mechanism is more complex than a
simple mass transport controlled process assumed by the model.
Although only the packed bed electrode has been tested
experimentally, it is reasonable to expect that other high mass
transport rate cells, for example incorporating a fluidized bed
electrode, would also recover the gold efficiently.
174.
6.2.8 Gbupled Chlorine Leaching and Electrowinning of Gold
The aim of these experiments was to test the feasibility of
coupling electrowinning and leaching and to gain some insight
into the behaviour of the overall process and define control
requirements. Such information could be used for the design of an
optimized system at a later stage.
Tests were performed at pH 0, 2 and 5.5 under potential
control of the cathode at 0.53 V and a flowrate of 16.2 x 10~^ m3
s~l. The ahlorine concentration was selected to be 0.1 mol m~3,
since these levels enable more accurate analytical determination
from the multicomponent analysis programme with the UV Diode Array
Spectrophotometer and as the pH was increased, make the decon
volution of the spectra possible. To compensate for this low
chlorine concentration, a high surface area of gold was used in
the leach reactor.
The idealized behaviour for the process (Figure 6.47), at
steady-state, assuming mass transport controlled reactions and
plug flow behaviour could be described as follows: gold is oxidized
in the leach reactor by the reduction of chlorine, then both are
reduced in the cathode compartment of the electrolytic reaotor and
finally chlorine is regenerated at the anode compartment. Figure
6.48 shows the concentration of gold vs. time at pH = 0 at three
key points, i.e. the inlet to the leach reactor, the outlet of
the leach reaotor which was the inlet concentration to the cathode compartment and the outlet of the cathode compartment.
175.
Figure 6.47
position in process circuit
Idealized behaviour for couple chlorine leaching and electro- winning of gold at steady-state.
Figure 6.48
Total dissolved gold concentrations as a function time at(0) inlet to leach reactor, (^) outlet of leach reactor, ( □ ) outlet of cathode compartment. Leaching bed of gold- plated 3 mm cylindrically-shaped carbon chips, packed bed of - 200 um carbon chips, feeder electrode-membrane potential0.53 V, flow rate = 16.2 x 1Cf6 m3 s“ , 0.1 mol Cl9 + 1 kmol H d ro-3. J
176.
The inlet and outlet gold concentrations at the leach reactor
tended to decrease slightly with time. The concentration at the
outlet to the cathode compartment followed an inverse parabolic
behaviour, which was unexpected, especially when compared with
the Au(I)/Au(III) molar ratio vs. time (Figure 6.49). There was a
peak in the Au(I)/Au(III) molar ratio at the outlet of both leach
reactor and packed bed electrode. The peak observed at the outlet
of the cathode compartment was correlated (Section 6.2.7) to the
dispersion of Au(I)surf.
The overall and specific (i.e. allowing for the change due
to chlorine reduction) Faradaic requirements for Au deposition
are shown in Figure 6.49. As expected the presence of ahlorine
increased the charge requirements. A few experiments were performed
to test whether CL2 was transported across the Nafion membrane.
When the anode compartment was saturated with chlorine, chlorine
was detected at the 10“® M level after four hours in the oatholyte
at pH = 0 and 5.5. This transport across the membrane decreased
current efficiencies. The chlorine concentration at the inlet
of the leach reactor stabilized at around 1.9 x 10“4 M and the
membrane current density vs. time (Figure 6.50) showed cyclical
variations, possibly due to the change in total gold concentration
and in Au(I):Au(III) molar ratios at the inlet of the cathode
compartment.
Figure 6.51 shows the time dependence of the chlorine con
centrations in different parts of the circuit at pH = 2. Contrary to the behaviour at pH = 0, the chlorine concentration did not stabilize at the inlet to the leach reactor., with the consequence
177.
10'3 time/s
Au(I) : Au(III) molar ratios at inlet leach reactor (©)■ outlet of leach reactor (^), outlet of cathode compartment ( □ ), total Faradays (fl), Fardays per mole of gold deposited(0), in the PBE operating under conditions specified in Figure 6.48.
Figure 6.50 Membrane current density as a function of time, operating under conditions specified in Figure 6.48.
178.
Figure 6.51
Total dissolved chlorine concentration as a function of time at ( ^ ) outlet of anode compartment, (O) inlet to leach reactor, (A) outlet to leach reactor, (□) outlet to cathode compartment. Leaching bed of gold-plated 3 mm cylindrically- shaped carbon chips, packed bed of -200 v»m carbon chips, feeder electrode-membrane potential 0.53 V, flow rate =16.2 x 10“® m3 s'"1, electrolyte = 0.99 kmol HC1 + 0.01 kmol Nad
----------------1________________ i________________i________________i________________ i1 2 3 ^ 5
10̂ time/s Figure 6.52Total dissolved gold concentrations as a function of time at(O) inlet to leach reactor, (A) outlet of leach reactor,<□ ) outlet cathode compartment, operating under conditions specified in Figure 6.51.
179.
that the total gold concentration around the circuit tended to
fall (Figure 6.52). The Au(I)/Au(III) molar ratio and the overall
and specific charge requirements at pH = 2, behaved similarly at
pH = 0. At pH = 5.5 it was impossible to deconvolute the spectra
of HC10 and Au(III), and therefore it was impossible to calculate
the Au(I)/Au(III) ratios and specific charge requirements. The
average cumulative charge requirements was 5.69 F(mol Au)“l and it
was found that the concentration of total gold at the inlet of the
leach reactor tended to increase with time (Figure 6.53). This
indicates that gold was not electrowon at an adequate rate at the
cathode, which was expected since it was found that the deposition
rate decreased rapidly with pH (Figure 6.46). The slow deposition
rate in turn, decreased the current and the rate of Cl2 generation.
The current decayed with time, which is consistent with the explana
tion given above. The pH decrease at the cathode exit was particularly noticeable.
The coupling of CI2 leaching to electrowinning of gold oper
ated most efficiently at pH's > 2 at 22 °C, though more experiments
are required. However, an optimized design of the electrolytic
reactor is necessary, with the anode placed away from the membrane
to minimize Cl2 transport across the membrane and the packed bed
should have a larger characteristic length.
6.2.9 Process Considerations
In the previous sections of this chapter, the dissolution, deposition and solution chemistry of gold in acidic chloride media has been discussed. The feasibility of solubilizing gold
180.
Figure 6.53
Total dissolved gold concentrations at ( Q ) inlet to leach reactor, (A) outlet of leach reactor, (□) outlet of cathode compartmment. Leaching bed of gold-plated 3 rim cylindrically- shaped carbon chips, packed bed of -200 pm carbon chips, feeder membrane potential 0.53 V flow rate = 16.2 x 10“^ m^ s”l, electrolyte = 1 kmol Cl" m~3 at pH = 5.5.
181.
by the reduction of CI2 and electrowinning the gold-chloro complexes
on a 3 dimensional electrode has been established. However, gold
occurs associated with other minerals (Section 2.1) of which silver
and pyrite (FeS2) are the most common. The effect of the silver
content on the proposed process has been discussed in Section 6.1.
Pyrite is an electroactive mineral which could increase chlorine consumption according to reaction(218,245) [6.27].
FeS2 + 7/7.5 Cl2 + 8H20 = Fe2+/Fe3+ + 14/15 Cl" + 2So|" + 16H+
[6.27]
The reduction of any Fe(III) species produced would then constitute
a loss of current efficiency for gold electrowinning, although 'Fed^'
has been used previously to leach gold(226)# However, by operating
at pH > 2-3 (Figure 6.54) and using HC10 species:
HC10 + H+ + 2e" = Cl" + H20 [6.28]
E = 1.494 - 0.0295 pH + 0.0295 log (HC10) - 0.0295 log (Cl“)
[6.28a]
rather than chlorine:
Cl2 + 2e~ = 2C1" [6.13]
E = 1.395 + 0.029 log (Cl2) - 0.029 log (Cl") [6.13a]
to drive the gold leaching reactions
Au + 1.5 HC10 + 1.5H+ + 2.5 Cl" = AuCl^ + H20 [6.29]
Au + 0.5 HC10 + 0.5H+ + 1.5C1" = AuClJ + H20 [6.30]
182.
Figure 6.54
Potential - pH diagram for the Fe-s-H20 system at 298 K, with dissolved iron and sulphur activities of lCT'S. considering FeOOH(s) as solid Fe(III) oxide phase (218) #
183.
FeOOH would be formed by an oxidative surface transformation,
(Figure 6.55 and 6.56) and the sulphide mineral oxidation rate
would be decreased. A decrease in proton concentration would
decrease the dissolution rate (Sections 6.24, 6.2.5, 6.2.6) and
the deposition rate (Sections 6.2.7 and 6.2.8) therefore an increase
in temperature would be desirable. The solubility of chlorine
decreases with an increase in temperature, at 20 °C the maximum solubility is 98.7 mol m-3 and 35.2 mol m~3 at 60 °c(246). The
solubility of chlorine at 60 °C would be at least an order of
magnitude higher than oxygen at 25 °C, according to(54):
The presence of oxygen would increase the gold dissolution rate,
increasing kinetic advantages over other leaching systems.
Another benefit of increased temperature would be an increased
solubility of silver. Therefore, there is a need to optimize the
operating temperature with respect to the product of diffusion
coefficient x solubility of chlorine. However, in certain cases
where the gold value is physically locked by pyrite, it might be
advantageous to operate at pH < 2 so that the dissolution of pyrite
would provide the porosity necessary to allow leaching solutions
to attack the gold. Present day practice to deal with this type
of ore is to roast the ore prior to oyanidation; therefore one
unit operation would be obviated if Cl2 leaching at pH < 2 were used, decreasing the capital investment required.
02 (g) = 02 (aq.) ; AG = 16.3 kj mol-1
log (02 aq.) = -2.855 + log Pq2
[6.3l][6.31a]
1 m .
Figure 6.55 Activity - pH diagram for the Fe(III)/H20 system at 298 K, considering FeOOH(s)(218)#
Figure 6.56 Activity - pH diagram for the Fe(III)/H20 system at 298 K, considering Fe2C>3(s)(218) #
185.
The high standard potential of the AuCl”/Au couple could in
theory be used advantageously to selectively deposit the gold,
leaving base metals in solution for subsequent recovery.
More experimental work is necessary to establish the condi
tions under which pyrite would be passivated or at least the con
sumption of chlorine reduced to an acceptable level and to optimize
the product of the diffusion coefficient x maximum solubility of
chlorine. Application for this process may be found in the treat
ment of alluvial gravity concentrates, electronic scrap and many
refractory gold ore types.
186.
CHAPTER SEVEN
CONCLUSIONS
CHAPTER 7 - CONCLUSIONS
Thermodynamic calculations have shown the strong depassivating and
solubilizing effect of Cl" ions on the Au/HpO system. This
effect is appreciable, for practical purposes, at concentrations_q _ _
> 0.5 kmol Cl" m . Aud^ and Audp are stable species at low
pH and high potential, the latter being favoured by high chloride
concentration and low total gold concentration. The concentration
of Cl” ions was found to be the controlling factor in the solubility
of gold in acidic chloride media.
In the cyclic voltammetry of gold electrodes. Two anodic peaks were
observed, provided the sweep rate was very slow (1 mV s”l) and
gold was deposited in the previous cathodic scan. The first was
due to the dissolution of gold as Audp and the second one was
due to both Audp and Aud^ ions being formed.
The oxidative dissolution of gold by chlorine reduction at 22 °C
was found to be under partial kinetic control. However, the
true kinetic currents, at pH = 0 and 2 mol d 2 m were found
to be between 560 and 340 A m“2, depending on the experimentalqmethod used. The maximum solubility of d p is 98.7 mol m at
20 °C indicating that the true kinetic currents at maximum d p
solubility would be at least an order of magnitude higher. The
main dissolution product was Audp and at a given d Q concen
tration, the dissolution rates increased with proton concentration and chloride ion concentration.
Gold electrodeposition in chloride media occurred through a complex mechanism, in which the slowest step was the reduction of:
Au(I)surf. + Au(s).
'188.
5. The feasibility of electrowinning AuCl^ and AuCl^ in a packed bed
electrode at high current efficiencies was demonstrated. Despite
the mechanism being more complex than two parallel mass transport
reduction reactions, a model which assumes mass transport control
may be used to describe the exponential decay of gold concentration
with time. The deposition rate was found to decrease markedly
with increasing pH. Although only the packed bed electrode has
been tested experimentally, it is reasonable to expect that other
high mass transport rates cells, for example incorporating a fluid
ized bed electrode, would also recover the gold efficiently.
6. The concept of coupled chlorine leaching and electrowinning of gold
has been tested. It was found that the process performance improved
with pH < 2 in 1 kmol Cl” m”3 solutions and that at pH = 5.5 the
system passivated. However, the optimization of the product of
diffusion coefficient x maximum solubility of CI2 and an increased
in Cl” ion concentration would allow the operation of the process
at pH's between 2 and 4.
7. Silver, commonly associated with native gold, could succesfully
be treated by decreasing the pH (< 2) and using a high chloride
activity.
8. The oxidation of pyrite by CI2 would incur considerable oxidant
consumption. However, thermodynamic calculations have shown the
possibility that by operating at pH = 2-3, this side reaction may
be decreased considerably.
9. Increased operating temperature would have the following benefical
effects:
189.
i) Increased leaching rates by optimizing the product of diffusion coefficient x maximum solubility of chlorine.
ii) Increased solubility of silver in chloride media.
iii) Operation of the process in the pH range 2-4 at acceptable leaching rates for gold and silver may be possible.
190.
REFERENCES
191.
REFERENCES
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209.
APPENDIX I : Eh-pH ACTIVITY EQUATIONS FOR THE Ag/H20/Cl/Cl04 SYSTEM
0 APPENDIX I : Eh-pH ACTIVITY EQUATIONS FOR THE Ag/H20/Cl/C104 SYSTEMo00
ION INUMBER I IONIC SPECIES
I CL- -I---
I FREE ENERGY I IONIC I I CHARACTERISTIC II OF FORMATION I CHARGE I ATOMS I ATOMS I.j------ x_..—I-------- 1-----1—I -31323.00 I -1 I 1 I 0 I-I.... ........ I....... I............. -I-------- 1—
OXYGEN I HYDROGENATOMS I ATOMS----X-----0 I------1--------
I FREE ENERGY I SPECIES
METALLIC SPECIES I OF FORMATION I TYPE*
-I-
I CHARGE -I....AG (S) I .00 I 1 I 0AG20 (S) I -2682.00 I 1 I 0AGO (S) I 837.00 I 1 I 0AG203 (S) I 29013.00 I 1 I 0AG202 (S) I 6604.00 I 1 I 0AG ♦ I 18442.00 I 0 I 1AG 2 ♦ I 64331.00 I 0 I 2AGO* I 53927.00 I 0 I 1AG(OH)2 - I -62230.00 I 0 I -1AG(OH)3- I -86537.00 I 0 I -1AGCL (S) I -26257.00 I 1 I 0AGCL I -12945.00 I 0 I 0ACCL2- I -51525.00 I 0 I -1AGCL3 2- I -82689.00 I 0 I -2AGCL4 3- I -114355.00 I 0 I -3AGCL04 I 16377.00 I 0 I 0
1-P7 SPECIES; O-DISSOLVED SPECIES
I METAL I OXYGEN I HYDROGEN I NUMBER OF CHARACTERISTIC ATOMS OF EACH IONIC SPECIESMETALLIC SPECIES I ATOMS I ATOMS I ATOMS I ION1 I ION2 I ION3 I ION4 I ION5 I ION6 I ION7 I ION8 I ION9 I IONIO--------- 1---- 1---- x------x--- 1------ 1 1 1 1 1 1-1... i...
I I I I l I I I I I I
AG (S) I 1 I 0 I 0 I 0AG20 (S) I 2 I 1 I 0 I 0AGO (S) I 1 I 1 I 0 I 0AG203 (S) I 2 I 3 I 0 I 0AG202 (S) I 2 1 2 I 0 I 0AG ♦ I 1 I 0 I 0 I 0AG 2* I I I 0 I 0 I 0AGO♦ I 1 I 1 I 0 I 0AG (OH) 2 - I 1 I 2 I 2 I 0AG (Oil) 3- I 1 I 3 I 3 I 0AGCL (S) I 1 I 0 I 0 I 1AGCL I 1 I 0 I 0 I 1AGCL2- I 1 I 0 I 0 I 2AGCL3 2- I 1 I 0 I 0 I 3
ACCL4 3- I 1 I 0 I 0 1 4AGCL04 I 1 I 4 I 0 I 1
I
210.
LINE INUMBER I BALANCED CHEMICAL EQUATIONS AND GENERALIZED LINE EQUATIONSI------1------------------------------------I
1 IIIII
2*AG * H20 - AG20 + 2*H+ + 2*E- E - 1.171 + -.0592 PH + .0000 LOG (M2/M1) + .0000 LOG(11)
2 AG + H20 - AGO + 2*H+ + 2*E- E - 1.247 + -.0592 PH + .0000 LOG (M2/M1) + .0000 LOG(11)
3 II
2*AG +■ 3*H20 - AG203 + 6*H+ + 6 E - 1.439 + -.0592 PH +
*E-.0000 LOG (M2/M1) + .0000
LOG (II) I4 I 2*AG + 2*H20 - AG202 + 4i * H * + 4 *E-
I E - 1.301 + -.0592 PH + .0000 LOG(M2/M1) + .0000 LOG(Il)
5 I AG -AG + E-I E - .800 + .0000 PH + .0592 LOG (M2/M1) .0000 LOG(Il)
6II AG - AG + 2*E-I E - 1.395 + .0000 PH * .0296 LOG (M2/M1) + .0000 LOG(11)
7 I AG +■ H20 - AGO* + 2*H + + 3*E-I E - 1.599 * -.0394 PH * .0197 LOG (M2/M1) + .0000 LOG(11)
8 I AG + 2*H20 - AG (OH) 2 * 21 *H+ + E-I E - 2.218 + -.1183 PH + .0592 LOG (M2/M1) .0000 LOG(11)
9 I AG + 3*H20 - AG(OH)3- * 3*11* + 2*E-
I E - 1.811 + -.0887 PH * .0296 LOG(M2/M1) + .0000 LOG(11)
10 I AG + CL- - AGCL + E-I E - .220 + .OOCO PH + .0000 L0G(M2/M1) + -.0592 LOG(11)
11 I AG + CL- - AGCL t E-I E - .797 + .0000 PH + .0592 LOG (M2/M1) + -.0592 LOG(I1)
12 I AG + 2*CL- - AGCL2- * E-I E - .482 + .0000 PH + .0592 LOG (M2/M1) + -.1183 LOG(I1)
13 I AG * 3*CL- - AGCL3 > E-1 E - .489 + .0000 PH * .0592 LOG(M2/M1) + -.1775 LOG(11)
14 I AG t 4 *CL- - AGCL4 + E-II
E - .474 + .0000 PH + .0592 LOG (M2/M1) -.2367 LOG (11)
15 I AG * 4*H20 + CL- - AGCIj04 t 8*11* * 9*E-IT
E - 1.322 + -.0526 PH + .0066 LOG (M2/M1) + -.0066 LOG(11)
16i.I AG2Ci ♦ 1120 - 2*7£0 + 2*IU t 2*E-I E - 1.324 * -.0592 PH * .0000 I/OG (M2/M1) .0000 LOG (11)
17 I 2*AG20 t 4 * H20 - 2*AG20J t 8*li* * 8*1E-1 E - 1.573 * -.0592 PH ♦ .0000 LOG(M2/M1) ♦ .0000 LOG (ID
18 I 2 *AG20 * 2*1120 - 2*AG202 * 4 *111 * 4*!E-I E - 1.430 * -.059? PH * .0000 LOG (M2/M1) + .0000 LOG (11)I19 r 2 *AG H20 AG20 f 2*H*
211 .
I LOG(M2/M1) - -12.55 i 2.CJ PH + .0000 LOG(11)
20 I AG20 - H20 - 2*AG - 2*Hf t- 2*E-II
E - 1.619 + .0592 PH + .0296 LOGIM2/M1) + .0000 LOG(Il)
21 I AG20 + H20 - 2*AG0+ + 2*Ht t 4*E-III
E - 1.813 + -.0296 PH + .0148 LOG (M2/M1) + .0000 LOG(Il)
22 AG20 + 3*H20 - 2*AG (OH)2 + 2*HtIII
LOG(M2/M1) - -35.39 f 2.00 PH + .0000 LOG(Il)
23 AG20 + 5*H20 - 2*AG(OH)3- + 4 *H+ f 2*E-III
E - 2.451 + -.1183 PH t .0296 LOG (M2/M1) + .0000 LOG(11)
24 2*AGCL + H20 - 2*CL- - AG20 ♦ 2*H+III
LOG (M2/M1) - -32.16 + 2.00 PH + -2.0000 LOG(11)
25 2*AGCL + H20 - 2*CL- =■ AG20 + 2*H+I LOG(M2/M1) - -12.64 + 2.00 PH + -2.0000 LOG(Il)
ILINE I
NUMBER I BALANCED CHEMICAL EQUATIONS AND GENERALIZED LINE EQUATIONSI------1------------------------------------I
26 I 2-AGCL2- + H20 - 4*CL- - AG20 + 2*H+I LOG(M2/M1) - -23.28 *• 2.00 PH + -4.0000 LOG(Il)I
27 I 2*AGCL3 + H20 - 6»CI- - AG20 + 2*H+I LOG (M2/M1) - -23.05 + 2.00 PH + -6.0000 LOG(Il)I
28 I 2*AGCL4 + H20 - 8*CL- - AG20 + 2*H+I LOG(M2/M1) - -23.55 * 2.00 PH + -8.0000 LOG(Il)I
29 I AG20 + 7*H20 + 2*CL- - 2*AGCL04 + 14*H+ + 16*E-I E - 1.341 + -.0518 PHI
30 I 2‘AGO + H20 - AG203 t 2 ‘‘litI E - 1.822 f -.0592 PH
31II 2*AGO - AG202I NO LINE GENERATED
32II AG i H20 - AGO i 2"IU i E-
I E - 1.695 t -.118 3 PH
33II AG » 1120 - AGO ♦ 2 *ii ♦I IDG(M2/M1) - 4.99
14 I AGO - AGOt t E-I E - 2.302 * .0000 Pll
35 I AG (OH)2 - H20 - AGO > EIT
E - .277 i .0000 PH
36lI AGO + 2*H20 - AG(OH)3- f H tI L0G(M2/M1) - -19.06I
.0037 LOG(M2/M1) + -.0074 LOG(Il)
2*E-.0000 IX3G (M2/M1) + .0000 LOG(Il)
.0592 LOG(M2/MI) ♦ .0000 LOG(11)
00 PH f- .0000 LOG(II)
.0592 IOG(M2/M1) ♦ .0000 LOG(11)
.0592 LOG(M2/M1) t .0000 LOG(II)
1.00 PH .0000 LOG(II)
212.
LOG(II)
1
37 I AGCL + H20 - CL- - AGO + 2*H + + E-I E - 2.275 + -.1183 PH * .0000 LOG (M2/M1) + .0592 LOG(11)
38 I AGCL + H20 - CL- - AGO + 2*H+ + E-I E - 1.698 + -.1183 PH + .0592 LOG(M2/M1) + .0592
39II AGCL2- + H20 - 2*CL- - AGO + 2*H+ + E-III
E - 2.012 + -.1183 PH + .0592 LOG(M2/M1) + .1183 LOGU1)
40 AGCL3 + H20 - 3*CL- - AGO + 2*H+ + EIT E - 2.005 + -.1183 PH + .0592 LOG(M2/M1) + .1775 LOG(11)
411I AGCL4 + H20 - 4 *CL- - AGO + 2*H+ + EI E - 2.020 + -.1183 PH + .0592 LOG (M2/M1) + .2367 LOG (11)
42II AGO + 3*H20 + CL- - AGCL04 + 6*H+ + 7*E-III
E - 1.344 + -.0507 PH + .0085 LOG (M2/M1) + -.0085 LOG(11)
43 2*AG202 + 2*H2C) - 2*AG203 + 4*H+ + 4*E-III
E - 1.715 + -.0592 PH + .0000 LOG (M2/M1) + .0000 LOG(11)
44 2* AG + 3*H20 -■ AG203 *■ 6*H* *■ 4 *E-
III
E - 1.758 + -.0887 PH + .0148 LOG(M2/M1) + .0000 LOG(Il)
45 2* AG + 3*H20 - AG203 + 6*H+ + 2*E-III
E - 1.527 + -.1775 PH + .0296 LOG(M2/M1) + .0000 LOG(11)
46 2 *AGO+ + H20 -■ AG203 + 2*HtI LOG (M2/M1) - 16.24 + 2.00 PH + .0000 LOG(11)
47II 2*AG(OH)2 - H20 - AG203 + 2*H+ + 4*E-III
E - 1.049 + -.0296 PH + .0148 LOG (M2/M1) + .0000 LOG(11)
48 2*AG (OH) 3- - 3*rH20 - AG203 t 2*E-I E .694 + .0000 PH + .0296 LOG(M2/M1) + .0000 LOG(11)
49 I 2*AGCL + 3*H20 _ 2*CL- - AG203 + 6*H-t + 4‘E-III
E - 2.048 -.0887 PH + .0000 LOG (M2/M1) + .0296 LOG(11)
50 2*AGCL + 3*H20 - 2 *CL- - AG203 t 6*IH■ + 4*E-
I E - 1.760 + -.088/ PH t .0148 LOG(M2/M1) ■f .0296 LOG(11)
ILINE I
NUMBER I BALANCED CHEMICAL EQUATIONS AND GENERALIZED LINE EQUATIONSI---------------- r ------------------------------------------------------------------------------------------------------I51 I 2*AGCL2- + 3*H20 - 4*CL- - AG203 t 6*H + + 4*E-
I E - 1.917 + -.088 / PH + .0148 LOG (M2/M1) + .0592 LOG(11)
52II 2 *AGCL3 + 3*H20 - 6 *CL- - AG203 t 6Mil t 4 *E-I E • 1.914 t -.088 / PH t .0148 LOG(M2/M1) t .0887 LOG(11)
53II 2*AGCL4 < 3*H20 - 8 *CL- • AG20) <■ 6*Ht t 4*E-I E - I
1.921 f -.088 / PH > .0148 IXX1 (M2/M1) + .1103 LOG(11)
54 I AG203 i 5*H20 t 2*CL- - 2 * AGC L04 + 10*H* * 12 *E-
213.
I E - 1.264 + -.0493 PH + .0049 LOG (M2/M1) + -.0099 LOG(11)
55III
2* AG + 2*H20 - AG202 + 4*H+ + 2 E - 1.802 + -.1183 PH +
*E-.0296 LOG (M2/M1) + .0000 LOG(Il)
56III
2*AG + 2*H20 - AG202 + 4*H+LOG(M2/M1) - 6.36 + 4.00 PH + .0000 LOG(11)
57IIIIII
AG202 - 2*AGO+ + 2*E-E - 2.195 + .0000 PH + .0296 LOG (M2/M1) + .0000 LOG (11)
58 2*AG(OH)2 - 2*H20 - AG202 + 2*E- E - .384 + .0000 PH + .0296 LOG (M2/M1) + .0000 LOG(Il)
59IIIII
AG202 + 4*H20 - 2*AG(OH)3- + 2*11 +LOG(M2/M1) - -34.50 + 2.00 PH + .0000 LOG (11)
60 2*AGCL + 2*H20 - 2*CL- - AG202 + 4*11 + + 2*E-
IIIIIIIIII
E - 2.382 + -.1183 PH + .0 0 0 0 L0G(M2/M1) + .0592 LOG(11)
61 2*AGCL + 2*1120 - 2*CL- - AG202 E - 1.804 + -.1183 PH +
+ 4*11 + .0296
+ 2*E- LOG (M2/M1) + .0592 LOG(Il)
62 2*AGCL2- + 2*H20 - 4*CL- - AG202 E - 2.119 + -.1183 PH +
+ 4*.0296
H+ + 2‘E- LOGIM2/M1) + .1183 LOG(Il)
63 2*AGCL3 + 2*1120 - 6*CL- - AG202 E - 2.112 + -.1183 PH +
+ 4 *H+ + 2*E- .0296 LOG (H2/M1) + .1775 LOG(11)
64IIIIIII
2*AGCL4 + 2*1120 - 8*CL- - AG202 E - 2.127 + -.1183 PH +
+ 4*H+ + 2*E- .0296 LOG(M2/M1) + .2367 LOG(11)
65 AG202 + 6*H20 + 2*CL- - 2*AGCL04 E - 1.329 + -.0507 PH +
+ 12 .0042
!*H+ + 14* LOG (M2/M1)
E-+ -.0085 LOG(11)
66 1IIIriii
AG -AG + E-E - 1.990 + .0000 PH + .0592 LOG (M2/M1) + .0 0 0 0 LOG (ID
67 AG + H20 - AGO+ + 2*H + * 2*E- E - 1.999 + -.0592 P1I * .0296 lOG (M2/M1) .0 0 0 0 LOG(11)
68 AG + 2*H20 - AG(OH)2 + 2*11 +LOG (M2/M1) - -23.97 + 2.00 PH + .0000 LOG(11)
69 ii
AG t 3*H20 - AG (OH) 3- f 1*11* * E - 2.822 * -.17 /5 PH ♦
E-.0592 lOG {M2/Ml) •* .0 0 0 0 LOG( I1)
70iii
AG + CL- - AGCL NO LINE GENERATED
71 ii
AG * CL- - AGCL NO LINE GENERATED
72 ii
AG + 2*CL- - AGC 1.2- NO LINE GENERATED
7 3iii
AG * 3*CL- - AGC 1.3 NO LINE GENERATED
74 i AC t 4 *CL- - AGC 1.4 •
214.
I NO LINE GENERATEDI75 I AG + 4*H20 + CL- - AGCL04 + 8*H + + 8*E-
I E - 1.388 + -.0592 PH + .0074 L0G(M2/M1) + -.0074 LOG(Il)
ILINE I
NUMBER I BALANCED CHEMICAL EQUATIONS AND GENERALIZED LINE EQUATIONS I------1------------------------------------I76 I AG + H20 - AGO+ + 2*H+ + E-
II
E - 2.007 + -.1183 PH + .0592 LOG(M2/Ml) + .0000 LOG(11)
77 I AG(OH)2 - 2*H20 -AG - 2*H + + E-III
E - .572 + .1183 PH + .0592 LOG(M2/M1) + .0000 LOG(Il)
?8 AG + 3*H20 - AG (OH) 3- + 3*H+I LOG (M2/M1) - -14.07 * 3.00 PH + .0000 LOG(11)
79 I AGCL - CL- - AG + E-III
E - 2.570 + .0000 PH * .0592 LOG (M2/M1) + .0592 LOG(11)
80 ACCL - CL- - AG + E-III
E - 1.993 + .0000 PH * .0592 LOG (M2/M1) + .0592 LOG(Il)
81 AGCL2- - 2*CL- - AG + E-III
E - 2.307 + .0000 PH * .0592 LOG (M2/M1) + .1183 LOG(11)
82 AGCL3 - 3*CL- - AG + E-II
E - 2.301 + .0000 PH + .0592 LOG(M2/M1) 4 .1775 LOG(11)
83 I AGCL4 - 4*CL- - AG + E-III
E - 2.315 + .0000 PH * .0592 LOG (M2/M1) + .2367 LOG(11)
84 AG + 4*H20 + CL-- - AGCL04 + 8*H* * 7*E-I E - 1.302 + -.0676 PH + .0085 LOG (M2/M1) + -.0085 LOG (11)
85 I AG(OH)2 - H20 - AGO* + 2*E-III
E - 1.289 + .0000 PH * .0296 LOG(M2/M1) 4 .0000 LOG (11)
86 AG(OH)3- - 2*H20 - AGO* - HI* * E-I E - 1.175 + .0592 PH * .0592 LOG (M2/M1) 4 .0000 LOG(11)
87II AGCL + H20 - CL - - AGO* t 2*H+ + 2*1E-I E - 2.289 + -.0592 PH * .0296 IOG (M2/M1) 4 .0296 LOG(11)
88 I AGCL + H20 - CL- - AGO* + 2*H* * 2*iE-I E - 2.000 + -.0592 PH * .0296 IOG (M2/M1) 4 .0296 LOG(11)
89II AGCL2- + H20 - i2*CL- - AGO* * 2*11 + * 2*E-I E - 2.157 t -.0592 PH * .0296 IOG (M2/M1) 4 .0592 LOG(I1)
90 I AGCL3 * H20 - 3*CL- - AGO t * 2*11* + 2*E-I E * 2.154 t -.0592 PH * .0296 10GIM2/M1) f .0887 LOG(11)
91 I ACCL4 * H20 - 4*CL- - AGO* * 2*11+ * 2‘E-I E - 2.161 * -.059? PH * .0296 IOG(M2/Ml) 4 .1183 LOG(11)I
' <LI
Z
I AGO+ + 3*H20 + CL- - AGCL04 + 6*H+ + 6*E-I E - 1.184 + -.0592 PH + .0099 LOG(M2/M1) +
93 I AG(OH)2 + H20 - AG(OH)3- + H+ + E-III
E - 1.404 + -.0592 PH + .0592 LOG (M2/M1) +
94 AGCL + 2*H20 - CL- - AG (OH) 2 + 2*H+I IOG (M2/M1) - -33.77 + 2.00 PH + -1.0000 LOG(11)
95II AGCL + 2*H20 -■ CL- - AG(OH)2 + 2*H+III
LOG (M2/M1) - -24.02 + 2.00 PH + -1.0000 LOG(11)
96 AGCL2- + 2*H20 - 2*CL- - AG(OH)21 + 2*H+III
IOG (M2/M1) - -29.33 + 2.00 PH + -2.0000 LOG(11)
97 AGCL3 + 2*1(20 - 3*CL- - AG (OH) 2 + 2*H +III
LOG (M2/M1) - -29.2 i + 2.00 PH + -3.0000 LOG(11)
98 AGCL4 + 2*H20 - 4*CL- - AG (OH) 2 + 2*H +I LOG (M2/M1) - -29.47 + 2.00 PH + -4.0000 LOG(11)
99 I AG(OH)2 + 2*H20 + CL- - AGCLO4 + 6*H+ + 8*E-IIIE - 1.210 + -.0444 PH + .0074 LOG(M2/M1) +
100 AGCL + 3*H20 -• CL- - AG(OH)3- t 3*H+ + E-I E - 3.402 + -.1775 PH + .0592 LOG(M2/Ml) +
LINE
II
-.0099 LOG(II)
.0000 LOG(II)
-.0074 LOG(II)
.0592 LOG(II)
BALANCED CHEMICAL EQUATIONS AND GENERALIZED LINE EQUATIONS
101 I AGCL + :3*1120 -I E - 2.825I
102 I AGCL2- 3v 3*H20I E - 3.140I
103 I AGCL 3 + 3*H20I E - 3.133
104 I AGCL 4 + 3*H20[ E - 3.148
105 I AG(OH)3- + H20
CL- - AG(OH)3- + 3*H + ̂E-+ -.1775 PH t .0592 L0G(M2/M1) +
.0592 LOG (M2/M1) ♦
AG(OH)3- + 3*H + + E-PH + .0592 LOG(M2/M1) +
AG (OH) 3-
CL- “ AGCLO :
1.183 + -.042 3 PH t
+ 3*H+ + E-.0592 LOG(M2/M1) 3
• 5 * 1(3 t 7*E-
.0085 IOG(M2/M1) +
.0592 LOG(II)
.1183 LOG(II)
.1775 LOG(II)
.2367 LOG(II)
-.0085 LOG (ID
106 I AGCL -AGCLI NO LINE GENERATED
107 I AGCL CL- - AGCL2-I NO LINE GENERATED
108 I AGCL 3 2*CL-- - AGCL3I NO LINE GENERATED
109 I AGCL 3*CL-- - AGCI.4I NO LINE GENERATED
216.
no III
AGCL + 4*H20 - AGCL04 + E - 1.460 + -.0592
8*H+ + PH + 8*E-
.0074 LOG(M2/M1)+ .0000LOG(11)m I
1II
AGCL + CL- - AGCL2- NO LINE GENERATED
112 AGCL + 2*CL- - AGCL3
IIII
NO LINE GENERATED
113 AGCL + 3»CL- - AGCL4 NO LINE GENERATED
114 IIIIIAGCL + 4*H20 - AGCL04 +
E - 1.388 + -.05928*H+ + PH + 8*E-
.0074 LOG(M2/M1) + .0000LOG(Il)
115 AGCL2- + CL- - AGCL3 NO LINE GENERATED
116 IIIII
AGCL2- ♦ 2*CL- - AGCL4 NO LINE GENERATED
111 AGCL2- + 4*H20 - CL- - AGCL04 E - 1.427 + -.0592 PH +■ 8*H +
.0074+ 8*E-
LOGIM2/M1) + .0074 LOG(11)
118 II AGCL3 + CL- - AGCL4 NO LINE GENERATED
119 I AGCL3 +■ 4*H20 - 2*CL- -AGCL04 + 8*H+ + 8*E-I E - 1.427 + -.0592 PH t .0074 LOG (M2/M1) + .0148 LOG (ID
120 II AGCL4 + 4*H20 - 3*CL- - E - 1.428 + -.0592
AGCL04 PH +
t 8*H + .0074
+ 8*E- LOG(M2/M1) + .0222 LOG(Il)
***************************'........ ....... .......... .........DISSOLVED METALLIC SPECIES ACTIVITIES .0001000
ACTIVITY OF CL- I 1.0000000
I NO. I EQUATION I NO. I EQUATION I NO. I EQUATION I NO. I EQUATIONI-
I 1 ][ E- 1.1/10 - •0592PH I 2 I E- 1.2473 - .0592PH I 3 I E- 1.4388 - .0592PH I 4 I E- 1.3007 - .0592PH
I 5 1[ E- .5631 I 6 I E- 1.2765 I 7 I E- 1.5200 - .0394PH I 8 I E- 1.9812 - . 1183PH1 9 1l E- 1.6927 - .0887PH I 10 I E- .219/ 1 11 I E- .5603 I 12 I E- .2456I 13 1[ E- .2525 I 14 I E- .2376 I 15 I E- 1.2961 - .0526PH I 16 1 E- 1.3236 - .0592PHI 17 ][ E- 1.572/ - •0592PH I 18 I E- i .4 305 - •0592PH I 19 I PH-10.2744 I 20 I E- 1.3821 t . 0592PHI 21 ]i E- 1.6946 - •0296PH I 22 IPH-13.6945 I 23 I E- 2.2144 - .1183PH I 24 I PH-16.0/83I 25 J[ PH- 10.3213 1 26 I PH-15.6403 1 27 I Pll-15.5238 I 28 I PH-15./752I 29 ][ E- 1.311 / - •0518PH I 30 I E- 1.8219 - .0592PH I 32 I E- 1.9315 - •1183PH I 33 I PH- -.4946
217
218.
APPENDIX II : Eh-pH ACTIVITY EQUATIONS FOR THE Au /H20/C1 SYSTEM
0 APPENDIX II : Eh-pH ACTIVITY EQUATIONS FOR THE Ag/H20/Cl/C104 SYSTEM
000ION I I FREE ENERGY I IONIC I
NUMBER I IONIC SPECIES I OF FORMATION I CHARGE I-----j------------1------------1------11 I CL- I -31323.00 I -1 I-----X------------1------------1------1X--------------------------------------
I FREE ENERGY I SPECIES I
I CHARACTERISTIC I OXYGEN I HYDROGENATOMS I ATOMS I ATOMS I ATOMS
1 I 0 I 0 I— I-------— I — -I —
METALLIC SPECIES I OF FORMATION I TYPE* I CHARGEAU (S) I .00 I 1 I 0AU(OH)3(S) I -75765.00 I 1 I 0AU02 (S) I 47992.00 I 1 I 0AUO (S) I 6525.00 I 1 I 0AU+ I 42065.00 I 0 I 1AU 3+ I 105162.00 I 0 I 3AU (OH)’3 I -67758.00 I 0 I 0AU03 3- I -12404.00 I 0 I -3AUOH 2+ I 43499.00 I 0 I 2AU(OH)2 1+ I -3824.00 I 0 I 1AU(OH)4 1- I -108890.00 I 0 I -1AU(OH)5 2- I -147347.00 I 0 I -2AUOH(AQ) I 5736.00 I 0 I 0AUCL2- I -36090.00 I 0 I -1AUCL3 I -18953.00 I 0 I 0AUCL4- I -56070.00 I 0 I -1
-I- -I- — I-* l=SOL SPECIES; 0=DISSOLVED SPECIES
00-------------------------------------------------------------------------------------------I METAL I OXYGEN I HYDROGEN I NUMBER OF CHARACTERISTIC ATOMS OF EACH IONIC SPECIES
METALLIC SPECIES I ATOMS I ATOMS I ATOMS I ION1 I ION2 I ION3 I ION4 I ION5 I ION6 I ION7 I ION8 I ION9 I IONIO----------------------------1-------------1------------- 1-----------------1---------- 1-------------------1 x x 1 x x-x-----------x---------
AU (S) I 1 I 0 I 0 I 0 IAU (OH) 3 (S) I 1 I 3 I 3 I 0 IAU02 (S) I 1 I 2 I 0 I 0 IAUO (S) I 1 I 1 I 0 I 0 IAU+ I 1 I 0 I 0 I 0 I
219.
AU 3+ I 1 I 0 I 0 I 0AU (OH) 3 I 1 I 3 I 3 I 0AU03 3- I 1 I 3 I 0 I 0AUOH 2+ I 1 I 1 I 1 I 0AU(OH) 2 1+ I 1 I 2 I 2 I 0AU(OH) 4 1- I 1 I 4 I 4 I 0AU(OH) 5 2- I 1 I 5 I 5 I 0AUOH(AQ) I 1 I 1 I 1 I 0AUCL2- I 1 I 0 I 0 I 2AUCL3 I 1 I 0 I 0 I 3AUCL4- I 1 I 0 I 0 I 4
IIIIIIIIIII-I---- 1---- 1---- 1---- 1---- 1---- 11
LINENUMBER
1
2
3LOG(II)
4
5
6
7
8
9
III BALANCED CHEMICAL EQUATIONS AND GENERALIZED LINE EQUATIONSI■III AU (S) + 3*H20 =AU(OH)3(S) + 3*H+ + 3*E-I E= 1 .36 3 + - .0 5 9 2 PH + .0000 LOG(M2/Ml) + .0000 LOG(11)II AU (S) + 2*H20 =AU02 + 4*H+ + 4*E-IT E= 1 .7 4 9 + - .0 5 9 2 PH + .0000 LOG(M2/Ml) + .0000 LOG(1 1)1I AU (S) +H20 =AUO + 2*H+ + :2*E-I E= 1 .371 + - .0 5 9 2 PH + .0000 LOG(M2/Ml) + .0000II AU (S) _AU+ +E-IT E= 1 .824 + .0000 PH + .0592 LOG (M2/Ml) + .0000 LOG(11)XI AU (S) =AU + :3*1S-iT E= 1 .52 0 + .0000 PH + .0197 LOG(M2/Ml) + .0000 LOG(11)XI AU (S) + 3*H20 =AU(OH) 3 + 3*H+ + 3*E-IT E= 1 .4 7 9 + - .0 5 9 2 PH + .0197 LOG (M2/Ml) + .0000 LOG(11)XI AU (S) + 3*H20 : =AU03 + 6*H+ + 3*E-IT E= 2 .2 7 9 + - .1 1 8 3 PH + .0197 LOG (M2/Ml) + .0000 LOG(11)Xi AU (S) +H20 =AUOH + H+ + 3 *E-I E= 1 .448 + - .0 1 9 7 PH + .0197 LOG(M2/Ml) + .0000 LOG(11)II
I
ro r j o
AU (S) + 2*H20 = AU(OH) 2 + 2*H+ + 3*E-
I E = 1.584 + -.0 3 9 4 PH + .0197 LOG(M2 /M l) +10
II AU(S) + 4*H20 = AU(OH) 4 + 4*H+ + 3*E-IT
E = 1.704 + - .0 7 8 9 PH + .0197 LOG(M2 /M l) +11
XI AU(S) + 5*H20 AU(OH) 5 + 5*H+ + 3*E-IT
E = 1.967 + - .0 9 8 6 PH + .0197 LOG (M2 /M l) +12
XI AU(S) + H20 = AUOH(AQ) + H+ + E -IT
E = 2.707 + - .0 5 9 2 PH + .0592 LOG (M2 /M l) +13
XI AU(S) + 2*C L- = AUCL2- + E -I E = 1.152 + .0000 PH + .0592 LOG (M2 /M l) +
14II AU(S) + 3*C L- AUCL3 + 3* E -II
E = 1.084 + .0000 PH + .0197 LOG(M2 /M l) +
15 I AU(S) + 4*C L- _ AUCL4- + 3* E -IT
E = 1.001 + .0000 PH + .0197 LOG(M2 /M l) +16
XI AU(OH) 3 (S) - H20 = AU02 + H+ + E -iT
E = 2.908 + -.0592 PH + . .0000 LOG (M2 /M l) +17
X1 AUO + 2*H20 = AU(OH)3 (S) + H+ + E -IT E = 1.348 + - .0 5 9 2 PH + .0000 LOG (M2 /M l) +
18XI AU+ + 3*H20 = AU (OH) 3 (S) + 3*H+ + 2*1IT
E = 1.133 + - .0 8 8 7 PH + .0296 LOG (M2 /M l) +19
Xi AU + 3*H20 = AU(OH) 3 (S) + 3*H+IT LOG(M2 /M l) = 7 .9 6 + 3 .00 PH + .0000 LOG(11 )
20XI AU(OH)3 (S) = AU(OH) 3 >IT NO LINE GENERATED
21XI AU(OH) 3 (S) = AU03 + 3*H+IT LOG(M2 /M l) = - 46.44 + 3 .00 PH + .0000 LOG(11 )
22XI AUOH + 2*H20 == AU(OH)3 (S) + 2*H+IT LOG(M2 /M l) = 4 .32 + 2 .0 0 PH + .0000 LOG(11 )
23XI AU(OH) 2 + H20 = AU(OH)3 (S) + H+IT LOG (M2 /M l) = 11.18 + 1 .00 PH + .0000 LOG( 11 )
24XI AU(OH)3 (S) + H20 = AU(OH) 4 + H+
0000 L O G (II)
0000 L O G (II)
0000 L O G (II)
0000 L O G (II)
1183 L O G (II)
0592 L O G ( I I )
0789 L O G (II)
0000 L O G (II)
0000 L O G (II)
0000 L O G ( I I )
221 .
I LOG (M2/Ml) = -17.27 + 1.00 PH + .0000 L O G (11)
25 I AU(OH)3 (S) + 2*H20 = AU(OH)5 + 2*H+I LOG (M2/Ml) = -30.63 + 2.00 PH + .0000 L O G (11)
LINEII
NUMBER II
BALANCED CHEMICAL EQUATIONS AND GENERALIZED LINE EQUATIONS
26II AUOH(AQ) + 2*H20 = AU(OH)3 (S) + 2*H+ + 2*E-
I E = .691 + -.0592 PH + .0296 L O G (M2/Ml) + .0000 L O G (11)
27 I AUCL2- + 3*H20 - 2*CL- = AU(OH)3 (S) 4- 3*H+ + 2*E-I E = 1.469 + -.0887 PH + .0296 LOG (M2/Ml) + .0592 L O G (11)
28 I AUCL3 + 3*H20 - 3*CL- = AU(OH)3 (S) + 3*H+I L O G (M2/Ml) = -14.13 + 3.00 PH + -■3.0000 L O G (11)
29 I AUCL4- + 3*H20 - 4*CL- = AU(OH)3 (S) 4■ 3*H+I L O G (M2/Ml) = -18.38 + 3.00 PH + --4.0000 L O G (11)
30 I AUO + H20 = AU02 + 2*H+ + 2*E-I E = 2.128 + -.0592 PH 4 .0000 LOG (M2/Ml) + .0000 L O G (11)
31 I AU+ + 2*H20 =■ AU02 + 4*H+ + 3*E-I E = 1.725 + -.0789 PH + .0197 L O G (M2/Ml) -f .0000 L O G (11)
32 I AU + 2*H20 = AU02 + 4*H+ + E-I E = 2.437 + -.2367 PH + .0592 LOG (M2/Ml) + .0000
L O G (11)
33 I AU(OH)3 - H20 = AU02 + H+ + E-I E = 2.561 + -.0592 PH + .0592 L O G (M2/Ml) + .0000 L O G (11)
34 I AU03 - H20 = AU02 - 2*H+ + E-I E = .161 + .1183 PH + .0592 LOG (M2/Ml) + .0000 L O G (11)
35 I AUOH + H20 - AU02 + 3*H+ + E-I E = 2.653 + -.1775 PH + .0592 L O G (M2/Ml) + .0000 L O G (11)
36 I AU(OH)2 = AU02 + 2*H+ + E-I E = 2.247 + -.1183 PH + .0592 LOG (M2/Ml) + .0000 L O G (11)
37 I AU(OH)4 - 2*H20 = AU02 + E-
222.
Ix E = 1.887 + .0000 PH + .0592 LOG(M2/Ml) + • .0000 LOG(11)38
1I AU(OH)5i - 3*H20 = AU02 -- H+ f E-IT E = 1.096 + .0592 PH + .0592 LOG (M2/Ml) + .0000 LOG(11)
39XI AUOH(AQ) + H20 ̂AU02 + 3*H+ + 3*E-IT E = 1.430 + -.0592 PH + .0197 LOG(M2/Ml) + .0000 LOG(11)
40XI AUCL2 + 2*H20 2*CL- == AU02 + 4*H+ + 3*E-IT E = 1.949 + -.0789 PH + .0197 LOG(M2/M1)+ .0394 LOG(11)
41XI AUCL3 + 2*H20 _ 3*CL- = AU02 + 4*H+ + E-IT E = 3.745 + -.2367 PH + .0592 LOG(M2/Ml) + .1775 LOG(11)
42XI AUCL4l- + 2*H20 ■ 4*CL- == AU02 + 4*H+ + E-IT E = 3.996 + -.2367 PH + .0592 LOG(M2/Ml) + .2367 LOG(11)
43XI AU+ + H20 = AUO + 2*H+ + E-II
E = .917 + -.1183 PH + .0592 LOG(M2/Ml) + .0000 LOG(11)
44 I AUO _ H20 = AU - 2*H+ -l- E-IT E = 1.819 + .1183 PH + .0592 LOG(M2/Ml) + .0000 LOG(11)
45XI AUO + 2*H20 = AU(OH)3 + H+ + E-IT E = 1.695 + -.0592 PH + .0592 LOG(M2/Ml) + .0000 LOG(11)
46XI AUO + 2*H20 = AU03 + 4*H+ + E-IT E = 4.096 + -.2367 PH + .0592 LOG(M2/Ml) + .0000 LOG(11)
47XI AUO = AUOH - H+ + E-IT E = 1.603 + .0592 PH + .0592 LOG(M2/Ml) + .0000 LOG(11)
48XI AUO + H20 = AU (OH) 2 + E-IT E = 2.009 + .0000 PH + .0592 LOG(M2/Ml) + .0000 LOG(11)
49XI AUO + 3*H20 = AU(OH)4 + 2*H+ + E-I E = 2.370 + -.1183 PH + .0592 LOG(M2/Ml) + .0000 LOG(11)
50II AUO + 4*H20 =AU(OH)5 + 3*H+ + E-I E = 3.160 + -.1775 PH + .0592 LOG(M2/Ml) + .0000 LOG(11)
LINENUMBER
III BALANCED CHEMICAL EQUATIONS AND GENERALIZED LINE EQUATIONS
223.
I
51II AUOH(AQ) = AUO + H+ + E-I E = .034 + -.0592 PH + .0592 LOG(M2/Ml) + .0000 LOG(11)
52 I AUCL2- + H20 - 2*CL- = AUO + 2*H+ + E-I E = 1.590 + -.1183 PH + .0592 LOG(M2/Ml) + .1183 LOG(11)
53 I AUO - H20 + 3*CL- = AUCL3 - 2*H+ + £I E = .512 + .1183 PH + .0592 LOG(M2/Ml) + -.1775
54II AUO H20 + 4*CL- = AUCL4- - 2*H+ + E-I E = .261 + .1183 PH + .0592 LOG(M2/Ml) + -.2367 LOG(11)
55 I AU+ — AU + 2*E-I E = 1.368 + .0000 PH + .0296 LOG(M2/Ml) + .0000 LOG(11)
56 I AU+ + 3*H20 = AU(OH)3 + 3*H+ + 2*E-I E = 1.306 + -.0887 PH + .0296 LOG(M2/Ml) + .0000 LOG(11)
57 I AU+ + 3*H20 = AUO3 + 6*H+ + 2*EI E = 2.506 + -.1775 PH + .0296 LOG (M2/Ml) + .0000 LOG(11)
58 I AU+ + H20 = AUOH + H+ + 2:*E-I E = 1.260 + -.0296 PH + .0296 LOG(M2/Ml) + .0000 LOG(11)
59 I AU+ + 2*H20 = AU(OH)2 + 2*H+ + 2*E-I E = 1.463 + -.0592 PH + .0296 LOG (M2/Ml) + .0000 LOG(11)
60 I AU+ + 4*H20 = AU (OH) 4 + 4*18+ + 2*E-I E = 1.643 + -.1183 PH + .0296 LOG(M2/Ml) + .0000 LOG(I1)
61 I AU+ + 5*H20 = AU (OH) 5 + 5*1H+ + 2*E-I E = 2.039 + -.1479 PH + .0296 LOG (M2/Ml) + .0000 LOG(11)
62 I AU+ + H20 = AUOH(AQ) + H+I LOG(M2/M1) = -14.92 + 1.00 PH + .0000 LOG(11)
63 I AU+ + 2*CL- = AUCL2-III
NO LINE GENERATED64 AU+ + 3*CL- = AUCL3 + 2*E-
II
E = .714 + .0000 PH + .0296 LOG(M2/Ml) + -.0887 LOG(11)
224.
65 I AU+ + 4*CL— = AUCL4- + 2*E-I E = .589 + .0000 PH + .0296 LOG (M2/Ml) + -.1183
66II AU + 3*H20 = AU(OH)3 + 3*H+IT LOG(M2/Ml) = 2.10 + 3.00 PH + .0000 LOG(11)
671I AU + 3*H20 = AU03 + 6*H+IT LOG(M2/Ml) = -38.48 + 6.00 PH + .0000 LOG(11)
681I AU + H20 = AUOH + H+IT LOG(M2/Ml) = 3.65 + 1.00 PH + .0000 LOG(11)
691I AU + 2*H20 = AU(OH)2 + 2*H+IT LOG(M2/M1) = -3.22 + 2.00 PH + .0000 LOG(11)
70J.I AU + 4*H20 = AU(OH)4 + 4*H+IT LOG(M2/Ml) = -9.31 + 4.00 PH + .0000 LOG(11)
71J.I AU + 5*H20 = AU(OH)5 + 5*H+IT LOG(M2/Ml) = -22.67 + 5.00 PH + .0000 LOG(11)
72i.I AUOH(AQ) - H20 = AU - H+ + 2*E-I E = .927 + .0296 PH + .0296 LOG (M2/Ml) + .0000 LOG(11)
73II AUCL2- - 2*CL- = AU + 2*E-IT E = 1.704 + .0000 PH + .0296 LOG(M2/M1) + .0592 LOG(11)
741I AU + 3*CL- = AUCL3I NO LINE GENERATED
75II AU + 4*CL- = AUCL4-I NO LINE GENERATED
LINE INUMBER I BALANCED CHEMICAL EQUATIONS AND GENERALIZED LINE EQUATIONSI----------j -------------------------------------------------------------
I76 I
ITAU (OH)3 = AU03
LOG(M2/Ml)+ 3*H+= -40.57 + 3.00 PH + .0000 LOG(11)
771I AUOH + 2*H20 == AU(OH)3 + 2*H+I LOG (M2/Ml) = -1.55 + 2.00 PH + .0000 LOG(11)
225.
78II AU(OH)2 + H20 = AU(OH)3 + H+IT LOG (M2/Ml) = 5.31 + 1.00 PH + .0000 LOG(11)
791I AU(OH)3 + H20 = AU (OH) 4 + H+IT
LOG(M2/Ml) = -11.40 + 1.00 PH + .0000 LOG(11)80
1I AU(OH)3 + 2*H20 = AU'OH)5 + 2*H+IT LOG(M2/Ml) = -24 76 + 2.00 PH + .0000 LOG(11)
81XI AUOH(AQ) + 2*H20 = AU(OH)3 + 2*H+ + 2*E-IT
E = .865 + -.0592 PH + .0296 LOG(M2/Ml) +82
1I AUCL2- + 3*H20 - 2*CL- = AU(OH)3 + 3*H+ + 2*E-IT
E = 1.642 + -.0887 PH + .0296 LOG(M2/Ml) +83
1I AUCL3 + 3*H20 _ 3*CL- = AU(OH)3 + 3*H+IT
LOG(M2/Ml) -20.00 + 3.00 PH + -3.0000 LOG(11)84
XI AUCL4- + 3*H20 - 4*CL- = AU(OH)3i + 3*H+IT
LOG(M2/Ml) = -24.25 + 3.00 PH + -4.0000 LOG(11)85
XI AUOH + 2*H20 == AU03 + 5*H+I LOG(M2/Ml) = -42.12 + 5.00 PH + .0000 LOG(11)
86II AU(OH)2 + H20 = AU03 + 4*H+I LOG(M2/Ml) = -35.26 + 4.00 PH + .0000 LOG(11)
87II AU(OH)4 - H20 = AU03 + 2*H+I LOG (M2/Ml) = -29.17 + 2.00 PH + .0000 LOG (11)
88II AU(OH)5 - 2*H20 = AU03 + H+IT
LOG(M2/Ml) = -15.81 + 1.00 PH + .0000 LOG(11)89
XI AUOH(AQ) + 2*H20 = AU03 + 5*H+ + 2*E-I E = 2.065 + -.1479 PH + .0296 LOG(M2/Ml) +
90I1 AUCL2- + 3*H20 - 2*CL- = AU03 * 6*Hf + 2*E-I E = 2.843 + -.1775 PH + .0296 LOG(M2/Ml) +
91II AUCL3 + 3*H20 - 3*CL- = AU03 + 6*H+IT LOG(M2/Ml) -60.57 + 6.00 PH + -3.0000 LOG(11)
92XI AUCL4- + 3*H20 - 4*CL- = AU03 H• 6*H+I LOG(M2/Ml) = -64.82 + 6.00 PH + -4.0000 LOG(11)
.0000 LOG(II)
.0592 LOG(II)
N)CT\
.0000 LOG(II)
.0592 LOG(II)
1
I93 I AUOH + H20 = AU(OH)2 + H+
IT LOG(M2/M1) = -6.86 + 1.00 PH + .0000 LOG(11)94
1I AUOH + 3*H20 = AU(OH)4 + 3*H+IT LOG (M2/Ml) = -12.95 + 3.00 PH + .0000 LOG(11)
951I AUOH + 4*H20 = AU(OH)5 + 4*H+IT LOG(M2/Ml) = -26.31 + 4.00 PH + .0000 LOG(11)
961I AUOH(AQ) = AUOH + 2*E-IT E = .819 + .0000 PH + .0296 LOG(M2/Ml) +
971I AUCL2- + H20 - 2*CL- = AUOH + 1H+ + :2*E-IT E = 1.596 + -.0296 PH + .0296 LOG(M2/Ml) +
981I AUCL3 + H20 - 3*CL- = AUOH + H+
.0000 LOG(II)
.0592 LOG(II)
I LOG(M2/M1) = -18.45 + 1.00 PH + -3.0000 LOG(Il)I99 I AUCL4- + H20 - 4*CL- = AUOH + H+
I LOG(M2/M1) = -22.70 + 1.00 PH + -4.0000 LOG(Il)I100 I AU(OH)2 + 2*H20 = AU(OH)4 + 2*H+
I LOG(M2/M1) = -6.09 + 2.00 PH + .0000 LOG(II)I
LINE INUMBER I BALANCED CHEMICAL EQUATIONS AND GENERALIZED LINE EQUATIONS
I---J.------1-----------------------------------------------------------------------I
101 I AU(OH)2 + 3*H20 = AU(OH)5 + 3*H+I LOG(M2/Ml) = -19.45 + 3.00 PH + .0000 LOG(Il)I
102 I AUOH (AQ) + H20 = AU(OH)2 + H+ + 2*E-I E = 1.022 + -.0296 PH + .0296 LOG(M2/Ml) + .0000 LOG(II)I
103
104
105
I AUCL2- + 2*H20 I E = 1.800II AUCL3 + 2*H20 I LOG(M2/Ml)II AUCL4- + 2*H20 I LOG (M2/Ml)
2*CL- = AU(OH)2 + 2*H+ + 2*E--.0592 PH + .0296 LOG(M2/Ml) + .0592 LOG(II)
3*CL- = AU(OH)2 + 2*H+-25.31 + 2.00 PH + -3.0000 LOG(II)4*CL- = AU(OH)2 + 2*H+-29.56 + 2.00 PH + -4.0000 LOG(Il)
hoho
I106 I AU(OH)4 + H20 = AU(OH)5 + H+
IT LOG (M2/Ml) = -13.36 + 1.00 PH + .0000 LOG(11)107
XI AUOH(AQ) + 3*H20 = AU(OH)4 + 3*H+ + 2*E-IT E = 1.202 + -.0887 PH + .0296 LOG(M2/Ml) +
108XI AUCL2- + 4*H20 - 2*CL- = AU(OH)4 + 4*H+ + 2*E-IT E = 1.980 + -.1183 PH + .0296 LOG(M2/Ml) +
109XI AUCL3 + 4*H20 - 3*CL- = AU(OH)4 + 4*H+I LOG(M2/Ml) = -31.40 + 4.00 PH + -3.0000 LOG(II)I
110 I AUCL4- + 4*H20 - 4*CL- = AU(OH)4 + 4*H+I LOG(M2/M1) = -35.65 + 4.00 PH + -4.0000 LOG(II)I
111 I AUOH(AQ) + 4*H20 = AU(OH)5 + 4*H+ + 2*E-IT E = 1.597 + -.1183 PH + .0296 LOG(M2/Ml) +
112XI AUCL2- + 5*H20 - 2*CL- = AU(OH)5 + 5*H+ + 2*E-IT E = 2.375 + -.1479 PH + .0296 LOG(M2/M1) +
113XI AUCL3 + 5*H20 - 3*CL- = AU(OH)5 + 5*H+IT LOG (M2/Ml) = -44.76 + 5.00 PH + -3.0000 LOG(11)
114XI AUCL4- + 5*H20 - 4*CL- = AU(OH)5 + 5*H+IT LOG (M2/Ml) = -49.01 + 5.00 PH + -4.0000 LOG(11)
115XI AUCL2- + H20 -- 2*CL- = AUOH(AQ) + H+IT LOG (M2/Ml) = -26.29 + 1.00 PH + -2.0000 LOG(11)
116XI AUOH(AQ) - H20 + 3*CL- = AUCL3 - H+ + 2*E-I E = .273 + .0296 PH + .0296 LOG(M2/Ml) +
LOG(11) T117
XI AUOH(AQ) - H20 + 4*CL- = AUCL4- - H+ + 2*E-IT E = .147 + .0296 PH + .0296 LOG(M2/M1) +
118XI AUCL2- + CL- == AUCL3 + 2*E-IT E = 1.051 + .0000 PH + .0296 LOG(M2/M1) +
119XI AUCL2- + 2*0- = AUCL4- + 2*E-IT E = .925 + .0000 PH + .0296 LOG(M2/Ml) +
120XI AUCL3 + CL- = AUCL4-I NO LINE GENERATED
0000 LOG(II)
0592 LOG(II)
0000 LOG(II)
0592 LOG(II)
.0887
1183 LOG(II)
.0296 LOG(II)
0592 LOG(II)
228.
1* DISSOLVED METALLIC SPECIES ACTIVITIES .0001000 ** ACTIVITY OF CL- I 1.0000000 ******************************************************************************
00I NO. I EQUATION I NO. I EQUATION I NO. I EQUATION I NO. I EQUATION
------------- j------------------------------------------------------------------------------------------------
I-I 1 I E= 1.3631 - .0592PH I 2 I E= 1.7494 - .0592PH I 3 I E= 1.3706 - . 0592PH I 4 I E= 1.5875I 5 I E= 1.4412 I 6 I E= 1.3999 - .0592PH I 7 I E= 2.2001 - . 1183PH I 8 I E= 1.3693 - . 0197PHI 9 I E= 1.5047 - .0394PH I 10 - E= 1.6248 - .0789PH I 11 I E= 1.8883 - .0986PH I 12 I E= 2.4703 - . 0592PHI 13 I E= .9149 I 14 I E= 1.0055 I 15 I E= .9217 I 16 I E= 2.9085 - .0592PHI 17 I E= 1.3480 - .0592PH I 18 I E= 1.2508 - . 0887PH I 19 I PH=-1.3214 I 21 I PH=14.1468I 22 I PH= -.1585 I 23 I PH=--7.1804 I 24 I PH=13.2697
I 25 I PH==13.3157I 26 I E= .8094 - . 0592PH I 27 I E= 1.5871 - .0887PH I 28 I PH= 6.0437 I 29 I PH= 7.4593I 30 I E= 2.1282 - . 0592PH I 31 I E= 1.8034 - .0789PH I 32 I E= 2.6740 - .2367PH I 33 I E= 2.7979 - .0592PHI 34 I E= .3975 + . 1183PH I 35 I E= 2.8897 - .1775PH I 36 I E= 2.4837 - . 1183PH I 37 I E= 2.1234I 38 I E= 1.3328 + .0592PH I 39 I E= 1.5091 - .0592PH I 40 I E= 2.0276 - .0789PH I 41 I E= 3.9812 - . 2367PHI 42 I E= 4.2325 - . 2367PH I 43 I E= 1.1537 - . 1183PH I 44 I E= 1.5825 + .1183PH I 45 I E= 1.4585 - . 0592PHI 46 I E= 3.8590 - . 2367PH I 47 I E= 1.3667 + .0592PH I 48 I E= 1.7728 I 49 I E= 2.1331 - . 1183PHI 50 I E= 2.9236 - . 1775PH I 51 I E= .2709 - .0592PH I 52 I E= 1.8263 - . 1183PH I 53 I E= .2752 + . 1183PH
I 54 I E= .0240 + . 1183PH I 55 I E= 1.3681 I 56 I E= 1.3061 - .0887PH I 57 I E== 2.5063 - . 1775PHI 58 I E= 1.2602 - .0296PH I 59 I E= 1.4632 - .0592PH I 60 I E= 1.6434 - .1183PH I 61 I E= 2.0387 - . 1479PHI 62 I PH=14.9214 I 64 I E= .7145 I 65 I E= .5888 I 66 I PH= -.6985I 67 I PH= 6.4127 I 68 I PH=-3.6472 I 69 I PH= 1.6081 I 70 I PH= 2.3264
229.
o o
I 71 I PH= 4.5334 I 72 I E= .9267 + . 0296PH I 73 I E= 1.7044 I 76 I PH=13.5239I 77 I PH= .7758 I 78 I PH=--5.3117 I 79 I PH=11.4010 I 80 I PH=12.3813I 81 I E= .8647 - .0592PH I 82 I E= 1.6424 - . 0887PH I 83 I PH= 6.6667 I 84 I PH= 8.0822I 85 I PH= 8.4247 I 86 I PH= 8.8150 I 87 I PH=14.5853 I 88 I PH=15.8090
I 89 I E= 2.0649 - .1479PH I 90 I E= 2.8426 - . 1775PH I 91 I PH=10.0953 I 92 I PH=10.8030I 93 I PH= 6.8633 I 94 I PH= 4.3176 I 95 I PH= 6.5786 I 96 I E= .8188I 97 I E= 1.5965 - .0296PH I 98 I PH=18.4483 I 99 I PH=22.6950 I 100 I PH= 3.0447I 101 I PH= 6.4837 I 102 I E= 1.0218 - . 0296PH I 103 I E= 1.7995 - .0592PH I 104 I PH=12.6558I 105 I PH=14.7792 I 106 I PH=13.3616 I 107 I E= 1.2020 - .0887PH I 108 I E= 1.9797 -I 109 I PH= 7.8502 I 110 I PH= 8.9119 I 111 I E= 1.5972 - . 1183PH I 112 I E= 2.3749 -I 113 I PH= 8.9525 I 114 I PH= 9.8019 I 115 I PH=26.2887 I 116 I E= .2730 +
0* * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * ** AREA OF PREDOMINANCE FOR% AU(S) *★ ★ ★ ★ ★ ★ ★ ★ ★ ★ ★ ★ ★ ★ ★ ★ A*******************************************************************
NO LOWER LIMITI POTENTIAL I I PH 1 I POTENTIAL 2 I PH 2 I LINE NO. II---I .915 I -2.00 I .915 I 7.57 I 13 II .915 I 7.57 I .578 I 13.27 I 1 II .578 I 13.27 I .571 I 13.36 I 10 II .571 I 13.36 I .329 I 15.81 I 11 II .329 I 15.81 I .307 I 16.00 I 7 IMINIMUM PH = -2.00MAXMIUM PH = 16.00
************************************************************************************* AREA OF PREDOMINANCE FOR% AU(OH)3(E) *********************************A***************************************************o -----------------------------------------------------
. 1183PH
. 1479PH
. 0296PH 230
I POTENTIAL I I PH 1 I POTENTIAL 2 I PH 2 I LINE NO. II— — I- -I-- — I- -I— — II 2.467 I 7.46 I 2.123 I 13.27 I 16 II .925 I 7.46 I .915 I 7.57 I 27 II .915 I 7.57 I .578 I 13.27 I 1 II 2.467 I 7.46 I .925 I 7.46 I 29 II 2.123 I 13.27 I .57.8 I 13.27 I 24 I
* AREA OF PREDOMINANCE FOR% AU02 (S)
0
NO UPPER LIMITI POTENTIAL 1 I PH 1 I POTENTIAL 2 I PH 2 I LINE NO. I
I 4.706 I -2.00 I 2.467 I 7.46 I 42 II 2.467 I 7.46 I 2.123 I 13.27 I 16 II 2.123 I 13.27 I 2.123 I 13.36 I 37 II 2.123 I 13.36 I 2.268 I 15.81 I 38 II 2.268 I 15.81 I 2.291 I 16.00 I 34 IMINIMUM PH MAXMIUM PH
- 2.0016.00
* AREA OF PREDOMINANCE FOR% AUO (S)NO AREAS OF PREDOMINANCE
* AREA OF PREDOMINANCE FOR% AU+NO AREAS OF PREDOMINANCE
* AREA OF PREDOMINANCE FOR% AU 3+) NO AREAS OF PREDOMINANCE)* AREA OF PREDOMINANCE FOR% AU(OH)3
I POTENTIAL I I PH 1 I POTENTIAL 2 1 PH 2 I LINE NO. II ---------------------1----------- 1--------------------- 1----------- 1---------------- 1
' 1Cz
I 2.320 I 8.08 I 2.123 I 11.40 I 33 II .980 I 8.08 I .784 I 11.40 I 45 II 2.320 I 8.08 I .980 I 8.08 I 84 II 2.123 I 11.40 I .784 I 11.40 I 79 I
* AREA OF PREDOMINANCE FOR% AU03 3~
I POTENTIAL I I PH 1 I POTENTIAL 2 I PH 2 I LINE NO. II-- — I- -I-- --I- -I— — II 2.268 I 15.81 I 2.291 I 16.00 I 34 II .329 I 15.81 I .307 I 16.00 I 7 II 2.268 I 15.81 I .329 I 15.81 I 88 II 2.291 I 16.00 I .307 I 16.00 I I
* AREA OF PREDOMINANCE FOR% AUOH 2+NO AREAS OF PREDOMINANCE
★ ★ ★ ★ ★* AREA OF PREDOMINANCE FOR% AU(OH)2 1+
NO AREAS OF PREDOMINANCE
* AREA OF PREDOMINANCE FOR% AU(OH)4 1-
I POTENTIAL 1 I PH 1 I POTENTIAL 2 I PH 2 I LINE NO. II 2.123 I 13.27 I 2.123 I 13.36 I 37 II .578 I 13.27 I .571 I 13.36 I 10 II 2.123 I 13.27 I .578 I 13.27 I 24 II 2.123 I 13.36 I .571 I 13.36 I 106 I
★ ★ ★ ★ ★* AREA OF PREDOMINANCE FOR% AU(OH)5 2-
I POTENTIAL I I PH 1 I POTENTIAL 2 I PH 2 I LINE NO. IX----------------- x--------- 1------------------1----------1------------- x.
0
roVairo
0
I 2.123 I 13.36 I 2.268 I 15.81 I 38 II .571 I 13.36 I .329 I 15.81 I 11 II 2.123 I 13.36 I .571 I 13.36 I 106 II 2.268 I 15.81 I .329 I 15.81 I 88 I
* AREA OF PREDOMINANCE FOR% AUOH(AQ)0 NO AREAS OF PREDOMINANCE0
* AREA OF PREDOMINANCE FOR% AUCL2-
I POTENTIAL I I PH 1 I POTENTIAL 2 I PH 2 I LINE NO. II--- — I- -I- -I- -I— — II .925 I -2.00 I .925 I 7.46 I 119 II .925 I 7.46 I .915 I 7.57 I 27 II .915 I -2.00 I .915 I 7.57 I 13 II .925 I -2.00 I .915 I -2.00 I I
* AREA OF PREDOMINANCE FOR% AUCL3
I POTENTIAL I I PH 1 I POTENTIAL 2 I PH 2 I LINE NO. II— — I -I— — I- -I— — II 4.455 I -2.00 I 2.551 I 6.04 I 41 II 1.051 I -2.00 I 1.051 I 6.04 I 118 II 4.455 I -2.00 I 1.051 I -2.00 I II 2.551 I 6.04 I 1.051 I 6.04 I 28 I
* AREA OF PREDOMINANCE FOR% AUCL4-
I POTENTIAL I I PH 1 I POTENTIAL 2 I PH 2 I LINE NO. II— --I- -I-- — I- -I— — II 4.706 I -2.00 I 2.467 I 7.46 I 42 II .925 I -2.00 I .925 I 7.46 I 119 II 4.706 I -2.00 I .925 I -2.00 I II 2.467 I 7.46 I .925 I 7.46 I 29 I00
roVjsl
1
* AREA OF PREDOMINANCE FOR% AU+ * DISSOLVED SPECIES DIAGRAM0 NO LOWER LIMIT0 -------------------------'----------------------------
I POTENTIAL I I PH 1 I POTENTIAL 2 I PH 2 I LINE NO. II---I .589 I -2.00 I .589 I 8.08 I 65 II .589 I 8.08 I .294 I 11.40 I 56 II .294 I 11.40 I .062 I 13.36 I 60 II .062 I 13.36 I -.168 I 14.92 I 61 IMINIMUM PH = -2.00MAXIMUM PH = 14.92 LINE NO. 62
★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★■a**** AREA OF PREDOMINANCE FOR% AU 3+ * DISSOLVED SPECIES DIAGRAM*★★*★★***★★******★★***★********★***★*★*★*★★★*★****★*★★★*★★★★★★★★******★★★*★*■*★*★ ★ ★ ★ ★0 NO AREAS OF PREDOMINANCE0* AREA OF PREDOMINANCE FOR% AU(OH)3 * DISSOLVED SPECIES DIAGRAM
0 NO UPPER LIMIT0 -------------------------------------------------------------------I POTENTIAL I I PH 1 I POTENTIAL 2 I PH 2 I LINE NO. II------------ I------ j-------------x-------1----------x-------------I .925 I 8.08 I 631 I 11.40 I 82 IMINIMUM PH = 8.08 LINE NO. 84MAXIMUM PH = 11.40 LINE NO . 790
* AREA OF PREDOMINANCE FOR% AU03 3- * DISSOLVED SPECIES DIAGRAM★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★a******
0 NO UPPER LIMIT0 ------------------------------------------------------I POTENTIAL I I PH 1 I POTENTIAL 2 I PH 2 I LINE NO. IX---------------- 1-----------1------------------ 1----------x--------------x------------------I .037 I 15.81 I .003 I 16.00 I 90 IMINIMUM PH = 15.81 LINE NO. 88MAXMIUM PH = 16.00
★ ★ ★ ★ ★ ★ ★ ★ ★ ★ ★ A * ★ ★ ★ ★ ★ ★ ★ ★ * * * * ★ * * * ★ * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * *
234.
* AREA OF PREDOMINANCE FOR% AUOH 2 + * DISSOLVED SPECIES DIAGRAM************************************************************************************0 NO AREAS OF PREDOMINANCE1 0★ ************************************************************************************ AREA OF PREDOMINANCE FOR% AU(OH)2 1+ * DISSOLVED SPECIES DIAGRAM★ ★★★★★★★★★★★★★★★★★★★★★★★★★★★★★it: ★★**★★★★*★★★★★★★★*★*****★★*★★★**★★**★★*******★★★★***
0 NO AREAS OF PREDOMINANCE0************************************************************************************* AREA OF PREDOMINANCE FOR% AU(OH)4 1- * DISSOLVED SPECIES DIAGRAM★ ***********************************************************************************
0 NO UPPER LIMIT0 -------------------------------------------------------------------I POTENTIAL I I PH 1 I POTENTIAL 2 I PH 2 I LINE NO. II------------- 1------1------------ 1------ 1---------- 1------------I .631 I 11.40 I .399 I 13.36 I 108 IMINIMUM PH = 11.40 LINE NO. 79MAXIMUM PH = 13.36 LINE NO. 106
0************************************************************************************* AREA OF PREDOMINANCE FOR% AU(OH)5 2- * DISSOLVED SPECIES DIAGRAM**********************************************»*************************************
0 NO UPPER LIMIT0 -------------------------------------------------------------------
I POTENTIAL I I PH 1 I POTENTIAL 2 I PH 2 I LINE NO. II------------- I----- I------------ 1------ 1---------- 1------------I .399 I 13.36 I .037 I 15.81 I 112 IMINIMUM PH = 13.36 LINE NO. 106MAXIMUM PH = 15.81 LINE NO. 88
0 .* AREA OF PREDOMINANCE FOR% AUOH(AQ) * DISSOLVED SPECIES DIAGRAM
0 NO AREAS OF PREDOMINANCE0★★★★★■ft******************************************************************************* AREA OF PREDOMINANCE FOR% AUCL2- * DISSOLVED SPECIES DIAGRAM0 NO LOWER LIMIT0 --------------------------------------------------------------
I POTENTIAL I I PH 1 I POTENTIAL 2 1 PH 2 I LINE NO. II --------------------1--------------1-----------------------1------------ 1----------------- 1
N3OT
I .925 I -2.00 1I .925 I 8.08 II .631 I 11.40 II .399 I 13.36 II .037 I 15.81 IMINIMUM PH = -2.00MAXMIUM PH = 16.00
925 I 8.08 I 119 I631 I 11.40 I 82 I399 I 13.36 I 108 I037 I 15.81 I 112 I003 I 16.00 I 90 I
10************************************************************************************* AREA OF PREDOMINANCE FOR% AUCL3 * DISSOLVED SPECIES DIAGRAM★★****★★★★*★*★★★*★*★★**★★★★**★*★*★**★***★★*★*★★★★★*★★★***★★*★★****★★**★*★*★**★★**★★*
NO UPPER LIMITI POTENTIAL I I PH 1 I POTENTIAL 2 1 PH 2 I LINE NO. II -------------------------1 ------------- 1 --------------------------1 --------------1 --------------------1 -
I 1.051 I -2.00 I 1.051 MINIMUM PH = -2.00MAXIMUM PH = 6.67 LINE NO. 83
6.67 I 118
★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★★Ik***** AREA OF PREDOMINANCE FOR% AUCL4- * DISSOLVED SPECIES DIAGRAM************************************************************************************
0 NO UPPER LIMITI POTENTIAL I I PH 1 I POTENTIAL 2 I PH 2 I LINE NO. II .925 I -2.00 I .925 I 8.08 I 119 IMINIMUM PH = -2.00MAXIMUM PH = 8.08 LINE NO. 84
236.
237.
< throughout should read
> throughout should read
gm-3 throughout should read
ERRATA
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>
g nr3mVs~l throughout should read mV s-1