Cyanide Regeneration from Thiocyanate with the Use of Anion Exchange Resins

Cyanide Regeneration from Thiocyanate with the use of Anion Exchange Resins

A thesis submitted for the degree of

Doctor of Philosophy

Kenneth Chung-Keong Lee B.E. Chemical Engineering (UNSW)

School of Chemical Engineering and Industrial Chemistry Faculty of Engineering The University of New South Wales Sydney, Australia October 2005 Abstract

It is known in many gold operations that less than 2% of the cyanide consumed accounts for the gold and silver dissolution. The majority of the cyanide is consumed by minerals contained in the gold ore to produce many different cyanide soluble complexes or used in converting cyanide to other related species such as thiocyanate and cyanate. The high costs associated with cyanide and thiocyanate detoxification and excessive cyanide utilisation encountered when treating ores with high cyanide consumption constitutes a significant proportion of the overall processing costs. This study examines the possibility of recovering free cyanide from thiocyanate using a process based on the Acidification-Volatilisation-Regeneration (AVR) circuit in conjunction with a pre-concentration stage using commercially available ion- exchange resin.

From thermodynamic modelling based on the STABCAL program it was found that it was thermodynamically possible to recover cyanide from thiocyanate if the oxidation of cyanide to cyanate can be stopped. Addition of copper to the system found that the majority of the thiocyanate exists as copper(I) thiocyanate (CuSCN) solid.

Using ion-exchange resins can be an effective way to concentrate thiocyanate from tailing solutions or slurries. Four different models were successfully used to model the equilibria between thiocyanate and chloride on commercial ion-exchange resins. By normalising the equilibria data when applying the Mass action law the equilibria becomes independent of ionic strength within the range of concentration considered. An advantage of this is that only one unique equilibrium constant is used to describe the ion-exchange process.

The electrochemical and kinetic studies showed that the reaction between thiocyanate and hydrogen peroxide is catalysed by hydrogen ions. Secondly under acidic conditions the rate of cyanide recovery by the AVR circuit was faster than at higher pH conditions. The overall reaction of thiocyanate with respect to the concentration of thiocyanate and hydrogen peroxide is an overall third order reaction. The derived

i third order rate expression is first order with respect to thiocyanate concentration and second order with respect to hydrogen peroxide concentration. Previous studies showed that the production of cyanide inhibits the reactions between thiocyanate and hydrogen peroxide, but by removing cyanide from the reaction by air stripping, this was not observed. Addition of copper to the system did not show a catalytic effect on the reaction but it was found that copper (II) ions suppresses competing reactions that occurred without affecting the reaction between thiocyanate and hydrogen peroxide.

ii Table of Contents

Abstract...... i

Table of Contents...... iii

List of Figures...... vii

List of Tables ...... xii

List of Publications ...... xiii

Chapter 1 General Introduction ...... 1

1.1 Introduction...... 1 1.2 Objectives of this study...... 2 1.2.1 Literature review ...... 2 1.2.2 Solution equilibria of the thiocyanate-water and copper-thiocyanate-water systems...... 3 1.2.3 Binary ion-exchange equilibria between thiocyanate and chloride using commercially available Purolite A500u/2788 ...... 3 1.2.4 Recovery of cyanide by air stripping from the oxidation of thiocyanate by hydrogen peroxide ...... 4 1.2.5 Electrochemical oxidation half-cell reaction of thiocyanate ...... 4

Chapter 2 Cyanide Chemistry...... 5

2.1 Introduction...... 5 2.2 Free Cyanide ...... 6 2.3 Complex cyanides...... 8 2.4 Other Cyanide Related Compounds...... 12 2.5 The Cyanide Cycle...... 12 2.6 Degradation of Cyanide ...... 14 2.6.1 Complexation...... 14 2.6.2 Cyanide Complex Precipitation ...... 14 2.6.3 Adsorption...... 14 2.6.4 Oxidation of cyanide to cyanate ...... 15 2.6.5 Volatilization...... 16 2.6.6 Biodegradation...... 17 2.6.7 Formation of Thiocyanate...... 18 2.6.8 Hydrolysis and Saponification of HCN ...... 18

Chapter 3 Cyanide Management Using Chemical Oxidation Processes ...... 23

3.1 Introduction...... 23 3.2 Cyanide Destruction...... 25

iii 3.2.1 The INCO process...... 25 3.2.2 Peroxygens...... 27 3.2.2.1 The Efflox Process...... 28 3.2.2.2 Ozone ...... 31 3.2.3 Chlorination ...... 33 3.3 Cyanide Recovery of free and weak-acid dissociable cyanides ...... 36 3.3.2 MNR (SART) Process ...... 40 3.3.3 Copper Electrowinning Process...... 41 3.4 Pre-concentration Processes...... 44 3.4.1 Resin Technologies...... 44 3.4.1.1 AuGMENT Process ...... 44 3.4.1.2 Vitrokele Process ...... 45 3.4.1.3 Elutech Process ...... 47 3.4.1.4 Cognis AURIX Resin ...... 48 3.4.1.5 Dow-Mintek (Minix) Ion exchange resin ...... 49 3.4.1.6 Superhydrophilic Ion Exchange Polymers...... 49 3.4.1.7 Elution of Thiocyanate from Strong based ion exchange resins...... 51 3.4.1.7.1 Elution of Thiocyanate Using Mineral Acids ...... 51 3.4.1.7.2 Elution of thiocyanate with potassium chloride and potassium hydroxide ...... 52 3.4.1.7.3 Elution of thiocyanate with Ferric ions...... 52 3.4.2 Solvent Extraction...... 55 3.5 Processes for the recovery of Cyanide From Thiocyanate ...... 57 3.5.1 Chemistry of Cyanide Regeneration from Thiocyanate ...... 58 3.5.1.1 Ozone ...... 59 3.5.1.3 Hydrogen Peroxide ...... 64 3.5.1.4 Caro’s Acid ...... 65 3.5.1.5 Chlorination ...... 66 3.5.1.6 Oxygen...... 67 3.5.1.7 Sulphur Dioxide...... 67

Chapter 4 Solution Equilibria ...... 76

4.1 Introduction...... 76 4.2 Chemical Reactions ...... 78 4.3 Theoretical Background on the formation of Eh-pH ...... 79 4.4 Distribution Diagrams...... 85 4.4.1 Theory behind the distribution diagram...... 85 4.5 Developments in computer aided programs to produce Pourbaix diagrams .....91 4.5.1 Point-by-point method ...... 91 4.5.2 Line elimination method...... 92 4.5.3 Convex polygon method...... 93 4.6 Pourbaix Diagrams for the Thiocyanate-Water and Copper-Thiocyanate-Water Systems ...... 95 4.6.1 Thiocyanate-Water System...... 95 4.6.2 Copper-Thiocyanate-Water System...... 98 4.8 Conclusion ...... 114

iv Chapter 5 Ion Exchange Resin...... 119

5.1 Introduction...... 119 5.2 Theory...... 120 5.2.1 Characterisation of ion-exchange equilibrium in a binary mixture ...... 120 5.2.2 Adsorption Isotherm Models ...... 121 5.2.2.1 Langmuir Isotherm...... 121 5.2.2.2 Freundlich Isotherm...... 122 5.2.2.3 Dubinin-Radushkevich Isotherm ...... 123 5.2.2.4 Mass-Action Sorption Model...... 124 5.3 Experimental...... 131 5.3.1 Reagents and Materials...... 131 5.3.2 Equipment and Instrumentation...... 132 5.3.3 Procedure for analysis of thiocyanate...... 133 5.3.4 Experimental Procedures ...... 134 5.4 Results and Discussion ...... 136 5.4.1 Equilibrium Study...... 136 5.4.2 Langmuir Model Fitting...... 138 5.4.3 Freundlich Model Fitting ...... 141 5.4.4 Dubinin-Radushkevich Model Fitting ...... 142 5.4.6 Mass Action Law Model Fitting...... 145 5.5 Conclusions...... 148

Chapter 6 Recovery of Cyanide from Thiocyanate using Hydrogen Peroxide ...... 152

6.1 Introduction...... 152 6.2 Chemistry of thiocyanate oxidation using hydrogen peroxide ...... 153 6.3 Decomposition of hydrogen peroxide...... 158 6.4 Reaction rate expressions...... 159 6.4.1 Derivation of a first order integrated rate expression ...... 159 6.4.2 Derivation of a second order integrated rate expression...... 160 6.4.3 Derivation of a third order integrated rate expression ...... 161 6.5 Experimental...... 165 6.5.1 Reagents and Materials...... 165 6.5.2 Equipment and Instrumentation...... 166 6.5.3 Analytical Procedures to determine the concentration of thiocyanate, free cyanide and hydrogen peroxide ...... 169 6.5.3.1 Procedure for thiocyanate determination...... 169 6.5.3.2 Procedure for free cyanide determination...... 169 6.5.3.2.1 Preparation of 0.01M standard silver nitrate solution...... 169 6.5.3.2.2 Preparation of silver sensitive 5-(4- (dimethylamino)benzylidene)rhodanine indictor...... 169 6.5.3.2.3 Analytical procedure for cyanide analysis...... 169 6.5.3.3 Procedure for hydrogen peroxide determination ...... 170 6.5.3.3.1 Preparation of standard potassium permanganate solution...... 170 6.5.3.3.2 Standardisation of potassium permanganate solution...... 171 6.5.3.3.3 Analytical procedure for hydrogen peroxide analysis ...... 171 6.5.4 Experimental Procedure...... 172 6.6 Results...... 175 6.7 Conclusions...... 189

Chapter 7 Electrochemical Studies of Thiocyanate Oxidation...... 193

7.1 Introduction...... 193 7.2 Voltammetry ...... 194 7.2.1 Rotating Disc Electrode ...... 196 7.3 Experimental...... 199 7.3.1 Reagents and Materials...... 199 7.3.2 Equipment and Instrumentation...... 199 7.3.2.1 Electrodes...... 199 7.3.2.2 Equipment...... 200 7.3.3 Experimental Procedure...... 202 7.4 Results...... 204 7.5 Conclusions...... 216

Chapter 8 Conclusions and Recommendations...... 218 8.1 Conclusions...... 218 8.1.1 Solution Equilibria for thiocyanate-water and copper-thiocyanate-water systems...... 218 8.1.2 Ion-exchange equilibria between thiocyanate and chloride...... 219 6.1.3 Recovery of cyanide by air stripping from the oxidation of thiocyanate by hydrogen peroxide ...... 220 6.1.4 Electrochemical studies of thiocyanate oxidation...... 221 8.2 Recommendations for further work...... 222

Appendix A: Fortran Code for Marquardt algorithm applied to the Mass-action Law ...... 224

Appendix B: Experimental data for ion-exchange equilibria between thiocyanate and chloride ...... 236

Appendix C: Experimental for the recovery of cyanide from thiocyanate using hydrogen peroxide ...... 239

Appendix D: Experimental for the electrochemical study of thiocyanate oxidation 243

Publications…………………………………………………………………………248

vi List of Figures Figure 2.1: Distribution of HCN and CN as a function of pH @25oC ...... 7 Figure 2.2: Distribution of different copper cyanide species as a function of pH. Thermodynamic data for system used from NBS Tables (Wagman et al., 1982)..8 Figure 2.3: The cyanide cycle, (Whitlock and Mudder, 1986)...... 13 Figure 2.4: Eh-pH diagram for Cyanide-Water System at 25oC generated from Stabcal using equilibrium data from NBS Tables (Wagman et al., 1982)...... 16 Figure 3.1: Illustration of the difference between hydrogen peroxide and Caro’s acid treatment of weak-acid dissociable cyanides in real liquor solutions (Nugent et al., 1991, Case study 1)...... 30 Figure 3.2: Illustration of the difference between hydrogen peroxide and Caro’s acid treatment of weak-acid dissociable cyanides in slurries (Nugent et al., 1991, Case study 1)...... 30 Figure 3.3: Oxidation of cyanide to cyanate using ozone as an oxidant. Operating o conditions: T=27 C, Flow rate of gas =4 L/min, Vsg=2.92 cm/s, ozone addition rate=0.11g/min. (Carrillo-Pedroza, 2000)...... 323 Figure 3.4: Flow diagram of the AVR Circuit...... 39 Figure 3.5: May Day Mine’s AVR Circuit (left) and Resin Columns (right)...... 39 Figure 3.6: Pilot scale SART plant flow sheet. (Dreisinger, 2001)...... 41 Figure 3.7: Schematic diagram of the electrolytic cell described by Lu et al. (2002)...... 43 Figure 3.8: Process diagram to regenerate ion exchange resin which was loaded with thiocyanate (Fleming,1986)…………………………………………………….54 Figure 3.9: Oxidation of thiocyanate using ozone as the oxidant. (Layne et al., 1984). Operating conditions: pH=11.2, I=0.23, T=25oC, gas flow=1.60 L/min...... 60 Figure 3.10: The effect of ozone oxidation on free cyanide (Soto et al., 1995).Note: CN (f) is the total concentration of HCN and CN-...... 61 Figure 3.11: Schematic of anodic and cathodic boundary layer formation for reactor with parallel plate bipolar electrodes. (Byerley et al., 1984)...... 63 Figure 3.12: Electrolytic oxidation of thiocyanate. A graph showing SCN/CN vs. time in a bipolar flow reactor. (Process conditions: temperature 30oC (av), current 8A (av), initial pH 9 un-buffered (0.09M Na2SO4), final pH 2.4) (Byerley et al., 1984)...... 64 Figure 4.1: Potential-pH equilibrium diagram for the system iron-water, at 25oC. Equilibrium data for each species are from NBS tables (Wagman et al., 1982). (Fe, Fe(OH)2 and Fe(OH)3 considered as the solid substances) ...... 77 Figure 4.2: Potential-pH equilibrium diagram for thiocyanate-water system at 25oC...... 97 Figure 4.3: Potential-pH equilibrium diagram for thiocyanate-water system at 25oC. Excluding cyanate and hydrogen cyanate from the simulation...... 98 Figure 4.4: Potential-pH equilibrium diagram for the copper-thiocyanate-water system at 25oC, at a copper and thiocyanate concentration of 0.001M...... 100 Figure 4.5: Potential-pH equilibrium diagram for the copper-thiocyanate-water system at 25oC, at a copper concentration of 0.001M and thiocyanate concentration of 0.01M...... 100

vii Figure 4.6: Potential-pH equilibrium diagram for the copper-thiocyanate-water system at 25oC, at a copper concentration of 0.001M and thiocyanate concentration of 0.1M...... 101 Figure 4.7: Potential-pH equilibrium diagram for the copper-thiocyanate-water system at 25oC, at a copper concentration of 0.001M and thiocyanate concentration of 1M...... 101 Figure 4.8: Potential-pH equilibrium diagram for the copper-thiocyanate-water system at 25oC, at a thiocyanate concentration of 0.001M and changing copper concentration from 0.001 to 1M...... 102 Figure 4.9: Potential-pH equilibrium diagram for the copper-thiocyanate-water system at 25oC, the thiocyanate and copper concentrations are at the same of either 0.001 or 0.01M...... 103 Figure 4.10: Potential-pH equilibrium diagram for the copper-thiocyanate-water system at 25oC, the thiocyanate and copper concentrations are at the same of either 0.1 or 1M...... 103 Figure 4.11: Distribution diagram for copper-thiocyanate and water system. The initial concentration of thiocyanate and copper is 0.001M at 25oC with the addition of copper (II)...... 106 Figure 4.12: A comparison of free thiocyanate concentration (SCN-) at different pH and Eh conditions as the ratio of thiocyanate to copper is changed...... 107 Figure 4.13: A comparison of hydrogen thiocyanate concentration (HSCN) at different pH and Eh conditions as the ratio of thiocyanate to copper is changed...... 107 Figure 4.14: A comparison of copper (II) thiocyanate concentration (CuSCN+) at different pH and Eh conditions as the ratio of thiocyanate to copper is changed...... 108 Figure 4.15: A comparison of the amount of copper (I) thiocyanate (CuSCN) at different pH and Eh conditions as the ratio of thiocyanate to copper is changed...... 109 Figure 4.16: A comparison of copper (I) di-thiocyanate concentration (Cu(SCN)2) at different pH and Eh conditions as the ratio of thiocyanate to copper is changed...... 109 3- Figure 4.17: A comparison of copper (I) tetra-thiocyanate concentration (Cu(SCN)4 ) at different pH and Eh conditions as the ratio of thiocyanate to copper is changed...... 110 - Figure 4.18: A comparison of free thiocyanate (SCN ) at different pH and Eh conditions as the ratio of copper(II) to thiocyanate changed...... 111 Figure 4.19: A comparison of hydrogen thiocyanate (HSCN) under different Eh conditions at a pH 0 as the ratio of copper(II) to thiocyanate changed...... 112 Figure 4.20: A comparison of copper (II) thiocyanate (CuSCN+) at different pH and Eh conditions as the ratio of copper(II) to thiocyanate changed...... 113 Figure 4.21: A comparison of copper (I) thiocyanate (CuSCN) at different pH and Eh conditions as the ratio of copper(II) to thiocyanate changed...... 113 Figure 4.22: A comparison of copper (II) di-thiocyanate (Cu(SCN)2) at different pH and Eh conditions as the ratio of copper(II) to thiocyanate changed...... 114 Figure 5.1: Ion exchange with a solution. For example a cation exchanger containing counter ions A is placed in a solution containing counter ions B (left). The counter ions are redistributed by diffusion until equilibrium is attained (right). (Helfferich, 1995)...... 120

viii Figure 5.2: Algorithm to calculate the three parameters: KAB, ΛAB, ΛBA to describe the ion exchange equilibria using the Mass-action law...... 130 Figure 5.3: Bench-Top Orbital Shaker Incubator...... 132 Figure 5.4: Concentration profile of thiocyanate adsorption at an initial thiocyanate concentration for each run of 1200ppm NaSCN, temperature 30oC at atmospheric pressure...... 136 Figure 5.5: Equilibrium adsorption curves for the adsorption of thiocyanate at different thiocyanate concentrations using Purolite A500u/2778 strong base ion exchange resin...... 137 Figure 5.6: The Langmuir Isotherm. A plot of the thiocyanate concentration on the resin versus the ratio of the thiocyanate concentration to equilibrium solution concentration...... 139 Figure 5.7: Parameter b for the Langmuir isotherm at different solution concentration...... 140 Figure 5.8: Comparison of experimental data to predicted values from the Langmuir Isotherm...... 140 Figure 5.9: The Freundlich Isotherm. A plot of log (X) (resin concentration) versus log (C) (solution concentration)...... 141 Figure 5.10: Comparison of experimental data to predicted values from the Freundlich Isotherm...... 142 Figure 5.11: The Dubinin-Radushkevich Isotherm. A plot of the natural log of thiocyanate concentration on the resin versus ε2 which is defined by equation 6...... 143 Figure 5.12: The change in the value of constant K in the Dubinin-Radushkevich isotherm with a change in ionic strength of solution...... 144 Figure 5.13: Comparison of experimental data to predicted values from the Dubinin- Radushkevich Isotherm...... 145 Figure 5.14: The Mass Action Law. A plot of resin fraction versus solution fraction of thiocyanate...... 147 Figure 5.15: Comparison of experimental data to predicted values by the Mass- Action Law...... 147 Figure 6.1: Graphical representation of bonding between copper(I) and the thiocyanate group...... 158 Figure 6.2: Schematic diagram of equipment set-up for oxidation study of thiocyanate using hydrogen peroxide. (1) Bubble Column (2) Caustic Trap (3) Precision Bore Flowrator (4) Cadet Air Pump...... 167 Figure 6.3: Front view of equipment arrangement for oxidation study of thiocyanate using hydrogen peroxide, showing: Bubble reactor, caustic trap, Precision bore flowrator and air pump...... 167 Figure 6.4: Dimensions of the bubble column reactor. (A) Top view, (B) Side View, (C) Front View...... 168 Figure 6.5: Oxidation of thiocyanate using hydrogen peroxide to recovery cyanide. Initial thiocyanate concentration 1000ppm, stoichiometric amount of hydrogen peroxide and 10g/l sulphuric acid at room temperature and pressure...... 176 Figure 6.6: A comparison of thiocyanate concentration against time at different solution pH. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution...... 177

ix Figure 6.7: A comparison of cyanide recovery at different initial pH’s. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution...... 177 Figure 6.8: The pH profile of hydrogen peroxide oxidation of thiocyanate with different initial starting pH. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution...... 179 Figure 6.9: The Eh profile of hydrogen peroxide oxidation of thiocyanate with different initial starting pH. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution...... 179 Figure 6.10: A comparison of thiocyanate concentration against time at different experimental conditions outlined in Table 6.2...... 181 Figure 6.11: A comparison of hydrogen peroxide concentration against time under different experimental conditions outlined in Table 6.2...... 182 Figure 6.12: A comparison of cyanide recovered against time under different experimental conditions outlined in Table 6.2...... 183 Figure 6.13: Applying the third order kinetic model for the reaction between thiocyanate and hydrogen peroxide under different initial pH’s. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution...... 185 Figure 6.14: Applying the third order kinetic model for the reaction between thiocyanate and hydrogen peroxide with a acid concentration of 10g/l sulphuric acid. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution...... 186 Figure 6.15: Applying the third order kinetic model for the reaction between thiocyanate and hydrogen peroxide with an initail pH of 12. The initial concentration of thiocyanate was 1000ppm (17.2mM), the amount of hydrogen peroxide added to the experimentis double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution ...... 186 Figure 6.16: Applying the third order kinetic model for the reaction between thiocyanate and hydrogen peroxide in the presence of copper(II). Experimental conditions outlined in Table 6.2……………………………………………….187 Figure 6.17: Applying the first order kinetic model for the reaction between thiocyanate and hydrogen peroxide with an initial pH of 7. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution …………………...188 Figure 6.18: Applying the second order kinetic model for the reaction between thiocyanate and hydrogen peroxide with an initial pH of 7. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution……………………188 Figure 7.1: Current-potential axes for voltammetric technquies...... 195

x Figure 7.2: Typical three electrode system for electrochemical measurements (Bard & Faulkner, 1980)...... 196 Figure 7.3:Schematic diagram of the fluid flow for a rotating disc system (Miller, 1979)...... 197 Figure 7.4: Schematic of equipment set-up for electrochemical half-cell studies. 1) Rotating disc electrode, 2) Speed controller, 3) Potentiostat, 4) Data Acquisition unit, 5) Counter electrode, 6) Reference electrode, 7) Thermometer, 8) 5 port Pyrex flask, 9) Water bath...... 201 Figure 7.5: Front view of the equipment arrangement for electrochemical studies showing: IBM computer with Quicklog PCTM software, TLC 548/9 Anolog to digital converter, Pine intrument Rotating Disc Electrode rotator with 5 port reaction cell and wath bath, MSR speed controller and CV-27 Potentiostat.....201 Figure 7.6: Cyclic voltammogram of the glassy carbon electrode in 0.1M Na2SO4 at pH 2, rest potential 0V and a scan rate of 40mV/s at a temperature of 25oC. ...204 Figure 7.7: Cyclic voltammogram of a glassy carbon in a solution of 0.1M Na2SO4+ 0.01M NaSCN, rest potential=0.5V (vs. Ag-AgCl electrode), pH 2, temperature 25oC. The scan rate was varied from 10 to 40 mV/s. (A) 10 mV/s, (B) 20 mV/s, (C) 30 mV/s, (D) 40mV/s...... 207 Figure 7.8: Stationary cyclic voltammogram of a glassy carbon in a solution of 0.1M Na2SO4, 0.01M NaSCN, rest potential=0V (vs. Ag-AgCl electrode), pH 12, temperature 25oC. The scan rate was varied from 10 to 40 mV/s. (A) 10 mV/s, (B) 20 mV/s, (C) 30 mV/s, (D) 40mV/s.………………………………………208 Figure 7.9: Peak currents versus square root of scan rate for the first peak in the voltammograms shown in Figure 7.7 (pH 2) and Figure 7.8 (pH 12).………...209 Figure 7.10: Peak currents versus square root of scan rate for the second peak in the voltammograms shown in Figure 7.7 (pH 2) and Figure 7.8 (pH 12).………...209 Figure 7.11: Rotation disc voltammograms of the anodic half-cell study of thiocyanate. The solution was made from 0.1M Na2SO4+0.01M NaSCN, pH 12, temperature 25oC and a rest potential of 0.5V vs Ag-AgCl reference. The rotation frequency was varied from 50rpm to 500 rpm. (A) 50rpm, (B) 100rpm, (C) 200rpm and (D) 500rpm.………………………………………………….211 Figure 7. 12: Rotation disc voltammograms of the anodic half-cell study of thiocyanate. The solution contained 0.1M Na2SO4+0.01M NaSCN, pH 2 and 12, temperature 25oC and a rest potential of 0.5V vs Ag-AgCl reference. The rotation frequency was varied from 50rpm and 100 rpm at each pH. (A) 50rpm and (B) 100rpm.……………………………………………………………….212 Figure 7.13: A plot of limiting current versus the square root electrode rotation frequency for experiments conducted in Figure 7.11.…………………………212 Figure 7.14: A comparison of the stationary cyclic voltammogram of a glassy carbon in a solution of 0.1M Na2SO4, 0.01M NaSCN with copper added to one experiment (concentration of 0.005M CuSO4), rest potential=0.5V (vs. Ag-AgCl electrode), pH 12, temperature 25oC and a voltage scanning rate of 40mV/s...214 Figure 7.15: Rotation disc voltammograms of the anodic half-cell study of thiocyanate. The solution contained 0.1M Na2SO4+0.01M NaSCN with one set o of experiments containing 0.005M CuSO4, pH 2, temperature 25 C and a rest potential of 0.5V vs Ag-AgCl reference. The rotation frequency was varied from 50rpm to 100 rpm. (A) 50rpm, (B) 100rpm………………………………….215

xi List of Tables Table 2.1: Classification of cyanide compounds ...... 6 Table 2.2: Dissociation constants for metal cyanides in decreasing stability in water..9 Table 2.3: Solubilities of Ferrocyanide and Ferricyanide in Water...... 11 Table 3.1: Operating cost associated with thiocyanate formation (Botz et al., 2001) 57 Table 3.2: Half-cell reactions for the oxidation of thiocyanate...... 58 Table 3.3: Stoichiometric reagents demands for thiocyanate oxidation with chlorine (Botz et al., 2001)...... 67 Table 4.1: Gibbs Free Energies for Thiocyanate and related species at 25oC...... 96 Table 4.2: Operating conditions for the simulations to determine the distribution diagrams for the copper-thiocyanate-water system...... 105 Table 5.1: Values for Parameters in the Debye-Huckel law (Appendix 7.1, Robinson et. al., 1965) ...... 128 Table 5.2: Typical Chemical and Physical Characteristics for Purolite A-500 and A- 500u/2788. (Purolite 2002) ...... 131 Table 5.3: Parameters in the model based on the Langmuir isotherm...... 139 Table 5.4: Parameters in the model based on the Freundlich isotherm...... 142 Table 5.5: Parameters in the model based on the Dubinin-Radushkevich isotherm.144 Table 6.1: Heat of decomposition of hydrogen peroxide ...... 159 Table 6.2: Experimental conditions for thiocyanate oxidation experiments with the addition of copper. The initial concentration of thiocyanate for each experiment is 1000 ppm SCN-. Experiments were carried out at room temperature and pressure. Airflow rate 1.2 L/min...... 181 Table 6.3: 3rd order reaction constants for a reaction between thiocyanate and hydrogen peroxide. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution...... 185 Table 7.1: Specifications of the cyclic voltammograph...... 200 Table 7.2: Peak Currents for different scan rates for figure 7.7...... 207 Table 7.3: Peak Currents for different scan rates for figure 7.9...... 208 Table 7.4: Rotating frequency of the working electrode and the related limiting currents for the experiments shown in Figure 7.11...... 213

xii List of Publications

1. Use of ion exchange resin for cyanide management during the processing of copper-gold ores. T. Tran, K. Lee, K. Fernando and S. Rayner In: Proceedings of Minprex 2000 International Congress on Mineral Processing and Extractive Metallurgy, held in Melbourne, Australia, 11-13 September 2000, pp 207-215. 2. Use of ion exchange resin for the treatment of cyanide and thiocyanate during the processing of gold ores. T. Tran, K. Fernando, K. Lee and F. Lucien In: Cyanide: Social, Industrial and Economic Aspects. Proceedings of a symposium held at Annual Meeting of TMS (The Minerals, Metals and Materials Society), held New Orleans, Louisiana. February 12-15, 2001, pp 289-302. 3. Halide as an alternative lixiviant for gold processing-an update. T. Tran, K. Lee and K. Fernando In: Cyanide: Social, Industrial and Economic Aspects. Proceedings of a symposium held at Annual Meeting of TMS (The Minerals, Metals and Materials Society), held New Orleans, Louisiana. February 12-15, 2001, pp 501-508. 4. Patent: Selective recovery of precious metal(s). K. Lee, F. Lucien, R. Rajasingam and T. Tran Patent Number: WO02099144, Publication date: 2002-12-12. ------5. Ion-exchange equilibria for Au(CN)2 /Cl , Au(CN)2 /SCN , and SCN /Cl in aqueous solution at 303K. N. Jayasinghe, K. Lee, F. Lucien and T. Tran In: Journal of Chemical and Engineering Data, American Chemical Society, In print. Web Release Date: July 9, 2004.

xiii Chapter 1 General Introduction 1.1 Introduction

Cyanide has been used to recover gold economically from ore for over the past one hundred years, but only a small amount of the cyanide consumed accounts for leaching the precious metals (gold and silver) from the gold bearing ores. It is known in many gold operations that less than 2% of the cyanide consumed accounts for the gold and silver dissolution. The majority of cyanide is consumed by minerals contained in the gold ore to produce many different cyanide soluble complexes or used in converting cyanide to other related species such as thiocyanate and cyanate. A small amount of cyanide is lost through volatilisation of hydrogen cyanide when cyanide is protonated. In New South Wales, Australia, the Environmental Protection

Authority limits the various metal cyanide species which constitute CNWAD (weak- acid dissociable cyanides) for discharge into tailing dams to less than 30ppm. Currently there are no limits to discharge thiocyanate into tailing dams in Australia but in the future due to consumer pressure upon the government may enforce discharge limits for thiocyanate into tailing dams.

Detoxification of cyanide is commonly used to remove CNWAD from tailing waters. The high costs associated with cyanide and thiocyanate detoxification and excessive cyanide utilisation encountered when treating ores with high cyanide consumption constitutes a significant proportion of the overall processing costs. Cyanide consumption higher than 5kg/tonne ore is generally found to be prohibitive for processing, often preventing the development of these ore resources. This has prompted the development of several alternative technologies that recover base metal cyanides and other species such as thiocyanate to recycle cyanide for re-use in the leaching stage.

1 One of the first cyanide recovery processes is based on the use of acid to regenerate cyanide as hydrogen cyanide gas, which is then adsorbed by a caustic or lime solution to produce sodium cyanide or calcium cyanide, respectively. The Acidification- Volatilisation-Regeneration (AVR) circuit in its original form (such as the Cyanisorb process) is expensive and less effective in treating dilute liquors or slurries containing less than 500ppm cyanide due to the handling of large volumes of liquor. However, if the AVR circuit is operated in conjunction with a pre-concentration stage such as ion- exchange or solvent extraction the AVR technique is more useful for handling low volumes of concentrated cyanide streams.

1.2 Objectives of this study

This study examines the possibility of recovering of free cyanide from thiocyanate using a process based on the AVR circuit in conjunction with a pre-concentration stage using commercially available ion exchange resin. The study will include a review of the chemistry of cyanide in a mining operation and management processes used to control the cyanide levels discharged to tailing ponds. Experimental work and discussions have been focused on the stability of different thiocyanate-water and copper-thiocyanate-water systems, binary ion-exchange equilibria between thiocyanate and chloride using commercially available Purolite A500u/2788, recovery of cyanide by air stripping from the oxidation of thiocyanate by hydrogen peroxide and electrochemical oxidation half cell reaction of thiocyanate.

1.2.1 Literature review

Literature review examines the chemistry of cyanide in a gold operation, dividing the many cyanide species formed and other related species into different classes. As a part of the cyanide cycle the review will cover discussions on the several natural pathways via which cyanide will degrade in the natural environment. Literature on different cyanide management processes used to control cyanide and thiocyanate levels discharged to tailing pond either by destruction of cyanide and thiocyanate or recovery of free cyanide from waste cyanides or thiocyanate was also reviewed.

2 1.2.2 Solution equilibria of the thiocyanate-water and copper-thiocyanate- water systems

Thermodynamic modelling has been successfully applied to a wide range of metallurgical, hydrometallurgical and geochemical processes. The production of Pourbaix and distribution diagrams is required to provide an understanding of the distribution of the predominant species in a thiocyanate-water and copper- thiocyanate-water system. This may give support to the understanding of the electrochemistry and reaction mechanisms for the oxidation of thiocyanate in the presence with or without copper. The thermodynamic modelling was established by using a modelling program (STABCAL) to generate the Pourbaix and distribution diagrams.

1.2.3 Binary ion-exchange equilibria between thiocyanate and chloride using commercially available Purolite A500u/2788

Using commercially available ion exchange resin (Purolite A500/2788) ion-exchange equilibria was examined between thiocyanate and chloride at 303 K. The experimental data was correlated and compared using various adsorption isotherm models to determine the apparent equilibrium constants for the binary system and to look at the effectiveness of using ion exchange resins to concentrate and remove thiocyanate from mine liquors. By examining different binary systems with other species such as base and precious metal cyanides using models to predict the equilibria for ion-exchange, a comparison of the relative affinity of different counter ions can be determined by using the cross product rule. This can be valued in the prediction of the loading characteristics of more complex systems found in the mining industry.

3 1.2.4 Recovery of cyanide by air stripping from the oxidation of thiocyanate by hydrogen peroxide

The recovery of cyanide from thiocyanate was studied under various different conditions using hydrogen peroxide. The experiments conducted involved oxidising thiocyanate to cyanide where an Acidification-Volatilisation-Reneutralisation circuit recovered the cyanide generated. The rate of reaction was observed under different operating conditions by changing the pH and the concentration of copper and hydrogen peroxide of the system. The aim of this study is to understand the overall reaction mechanism between thiocyanate and hydrogen peroxide to recover cyanide that can be re-used in the leach stage of the gold operation.

1.2.5 Electrochemical oxidation half-cell reaction of thiocyanate

Electrochemical studies were used to examine the oxidation half-cell reaction of thiocyanate using a voltametric technique based on the Rotating Disk Electrode (RDE) system. Many researchers have successfully used the RDE system to examine different redox reactions. The technique involves the oxidation half-cell reaction of thiocyanate in the absence of any chemical oxidants. The effects of various factors, namely, solution pH, scanning rates and disc rotation speeds and addition of copper on the current-potential relationship of the thiocyanate half-cell reaction were evaluated.

4 Chapter 2 Cyanide Chemistry 2.1 Introduction

Over the past hundred years cyanide has been used commonly within the gold industry to leach gold and silver from ores because it is an efficient lixiviant (or leachant) to dissolve these precious metals. Dilute cyanide solutions can be used to leach the precious metal and at current market prices for gold it is economical to use cyanide for many different ores. Only a small amount of the cyanide consumed accounts for leaching gold and silver from gold bearing ores. It is known that in many gold operations less than 2% of the cyanide consumed accounts for the of gold and silver dissolution (Marsden and House, 1992). Side reactions with other minerals within the ore consume most of the free cyanide.

The majority of the cyanide that is consumed reacts with many cyanide soluble minerals found commonly in gold bearing ores. For example copper minerals such as azurite (Cu3(CO3)2(OH)2) malachite (Cu2CO3(OH)2), cuprite (CuO2), tenorite (CuO),

chalcocite (Cu2S) and covellite (CuS). Other less soluble minerals are bornite

(Cu5FeS4), chrysocolla (Cu2H2Si2O5(OH)4) and chalcopyrite (CuFeS2) (Tran et al, 2000). Sulphide minerals can also react with free cyanide present in leach solutions to form thiocyanate. The loss of free cyanide is also due to its oxidation to cyanate or hydrogen cyanate and the volatilisation in the form of hydrogen cyanide (HCN).

The chemistry of cyanide is complex and dynamic. The different species of cyanide have been classified into 3 different categories namely free cyanide, weak-acid dissociable and strong acid dissociable cyanides according to the relative stability of each species. This classification can be further subdivided into 5 sub-categories as shown in Table 2.1. Other related species such as thiocyanate and cyanate are species related to the cyanide family but have to be considered under a different category.

5 Table 2.1: Classification of cyanide compounds Classification Compound

Free cyanide 1. Free cyanide CN, HCN

Weak-acid dissociable cyanide 2. Simple compounds

(a) Readily soluble salts NaCN, KCN, Ca(CN)2, Hg(CN)2

(b) Neutral insoluble salts Zn(CN)2, Cd(CN)2, Cu(CN), Ni(CN)2, Ag(CN)

2- - 2- 3. Weak complexes Zn(CN)4 , Cd(CN)3 , Cd(CN)4

Strong-acid dissociable cyanide -1 -2 -2 -1 4. Moderately strong complexes Cu(CN)2 , Cu(CN)3 , Ni(CN)4 , Ag(CN)2

-4 -4 -1 -3 5. Strong complexes Fe(CN)6 , Co(CN)6 , Au(CN)2 , Fe(CN)6

Source: Scott and Ingles (1987)

2.2 Free Cyanide

Free cyanide refers to two different species, the cyanide ion (CN-) and hydrocyanic acid or otherwise known as hydrogen cyanide (HCN). The relationship between the two species depends on the pH of the system. In water the free cyanide is also in equilibrium with HCN according to equation 2.1

C N H HCN H + +⇔ CN - [2.1]

where the pKa for HCN is

2347.2 pK 1.3440 += [2.2] a + 273.16T and T is temperature given in degrees Celsius (Izatt et al., 1962).

6 100

30 [HCN]/([HCN]+[CN]) (%) [HCN]/([HCN]+[CN])

0 6 7 8 9 10 11 12 pH

Figure 2.1: Distribution of HCN and CN as a function of pH @ 250C.

The distribution of CN- and HCN as a function of pH is shown in figure 2.1. Most gold plants operate at a pH of 10.5 to 12 to keep the predominant species of cyanide as CN-. If the pH drops too low as in natural and brackish waters, which have pH below 8.3, cyanide will exist predominantly as HCN. This is important because HCN has a relatively high vapour pressure (100 kPa at 26oC, Huiatt, 1982) and will volatilise readily from the liquid surface under ambient conditions causing a loss of cyanide from the solution. The rate of cyanide loss depends greatly on the formation of HCN, which is dependent on pH.

7 2.3 Complex cyanides.

Simple cyanide complexes are salts of hydrocyanic acid. These salts will dissolve completely in solution to produce a free alkali cation and cyanide anions (such as sodium cyanide).

Complex cyanides include those formed between cyanide and cadmium, nickel, copper, silver and zinc. In some cases, multi-cyanide complexes are formed in a step- wise manner; an example of this is copper cyanide. Depending on the ratio of CN- to - 2- 3- Cu(I) the predominant species are CuCN(s) or Cu(CN)2 , Cu(CN)3 and Cu(CN)4 as illustrated in Figure 2.2.

100%

Cu(CN)3 Cu(CN)4 90%

CuCN 80%

70%

60%

50%

40%

Cu(CN)2 30%

20%

10% Distribution of Copper Cyanide Species

0% 23456789101112 pH

Figure 2.2: Distribution of different copper cyanide species as a function of pH. Thermodynamic data for system used from NBS tables (Wagman et al., 1982).

8 Weak acid dissociable cyanides create serious problems during waste handling especially those which exist in more than one form such as copper cyanide complexes. As free cyanide decomposes with time by ultra-violet light, more cyanide can be liberated from copper cyanide complexes and the equilibrium will shift from 3- 2- - Cu(CN)4 and Cu(CN)3 to complexes of Cu(CN)2 and finally to CuCN(s). This maintains the high level of free cyanide within the tailing ponds. If excess cyanide exist in the tailing waters, copper cyanide can uptake the cyanide to form stable complexes of Cu(I) species and acts as a sink to store the free cyanide. This equilibrium effectively buffers the free cyanide concentration within the tailing waters and complicates the cyanide management of the tailings.

Stability of each cyanide compound is dependant on the metal associated with it, the most stable being iron and the weakest being zinc and cadmium. The dissociation constants of selected cyanide complexes are presented in Table 2.2. The rate of dissociation of the complex or the release of free cyanide depends on a number of factors. These factors include intensity of light, temperature of the water, pH, total dissolved solids and concentration of the cyanide complex.

Table 2.2: Dissociation constants for metal cyanides in decreasing stability in water. Complex Dissociation Constant 1.0 x 10-52 -3 Fe(CN)6 -47 -4 1.0 x 10 Fe(CN)6 -42 -2 4.0 x 10 Hg(CN)4 -28 -2 5.0 x 10 Cu(CN)3 -22 -2 1.0 x 10 Ni(CN)4 -21 -1 1.0 x 10 Ag(CN)2 -17 -2 1.4 x 10 Cd(CN)4 -17 -2 1.3 x 10 Zn(CN)4 Source: Caruso (1975)

As with weak acid dissociable cyanides, strong-acid dissociable cyanides will liberate free cyanide by reflux, but at a much lower pH. These include cyanides of iron, gold

9 and cobalt. These cyanide compounds are very stable and therefore are classified as non-toxic.

From environmental standpoint, the most significant strong-acid dissociable cyanides are iron cyanides. This is because iron cyanide has very high chemical stability, which results in extremely slow rates of dissociation and very low solubilities in water. Secondly, iron cyanide complexes can undergo a variety of chemical reactions, for example they can combine with other metals to form very stable double salts without the loss or exchange of the cyanide ligand as shown in Table 2.3.

Under certain conditions iron cyanide will decompose to release cyanide into the environment. Theses complexes resist natural degradation until the free cyanide and more degradable metal cyanides have dissipated. Iron cyanides are capable of breaking down and releasing cyanide when exposed to intense ultraviolet radiation as illustrated in the following reaction (equation 2.3):

UV Fe(CN)-4+ H O ⎯⎯⎯→ Fe(CN) .H O + CN -1 [2.3] 62←⎯⎯dark⎯ 52

This maybe a problem at mine sites across the world if iron cyanides are present in their tailing dams. It is possible that the free cyanide can be reintroduced into the dam sites after the removal of free and more degradable metal cyanides. It would be unlikely that iron cyanides pose a significant problem concerning for the management of cyanide in tailing waters but its presence introduces another variable for consideration.

10 Table 2.3: Solubilities of Ferrocyanide and Ferricyanide in Water. Name Solubility (g/l) Formula

Ammonium Ferricyanide (NH4)3Fe(CN)6 Very soluble

Ammonium Ferrocyanide (NH4)3Fe(CN)6.3H2O Soluble o Barium Ferrocyanide Ba2Fe(CN)6.6H2O 1.7 (15 C)

Cadmium Ferrocyanide Cd2Fe(CN)6.XH2O Insoluble o Calcium Ferrocyanide Ca2Fe(CN)6.12H2O 868 (25 C)

Cobalt Ferrocyanide Co2Fe(CN)6.XH2O insoluble

Copper (I) Ferricyanide Cu2Fe(CN)6 insoluble

Copper (II) Ferricyanide Cu3(Fe(CN)6)2.14H2O insoluble

Copper (I) Ferrocyanide Cu2Fe(CN)6.XH2O insoluble

Iron (II) Ferricyanide Fe3(Fe(CN)6)2 insoluble

Iron (III) Ferricyanide FeFe(CN)6 -

Iron (II) Ferrocyanide Fe2Fe(CN)6 insoluble

Iron (III) Ferrocyanide Fe4(Fe(CN)6)3 insoluble

Lead Ferricyanide Pb3(Fe(CN)6)2.5H2O slightly soluble

Magnesium Ferrocyanide Mg2Fe(CN)6.12H2O 330

Manganese (II) Ferrocyanide Mn2Fe(CN)6.7H2O insoluble

Nickel Ferrocyanide Ni2Fe(CN)6.XH2O insoluble o Potassium Ferricyanide K3Fe(CN)6 330 (4 C) o Potassium Ferrocyanide K4Fe(CN)6.3H2O 278 (12 C) o Silver Ferricyanide Ag3Fe(CN)6 0.00066 (20 C)

Silver Ferrocyanide Ag4Fe(CN)6.H2O insoluble o Sodium Ferricyanide Na3Fe(CN)6.H2O 189 (0 C) o Sodium Ferrocyanide Na4Fe(CN)6.10H2O 318.5 (20 C)

Strontium Ferrocyanide Sr2Fe(CN)6.15H2O 500 o Thallium Ferrocyanide Tl4Fe(CN)6.2H2O 3.7 (18 C)

Tin (II) Ferrocyanide Sn2Fe(CN)6 insoluble

Tin (IV) Ferrocyanide SnFe(CN)6 insoluble

Zinc Ferrocyanide Zn2Fe(CN)6 insoluble Source: Huiatt et al. (1982)

11 2.4 Other Cyanide Related Compounds

A number of cyanide related compounds are produced during the cyanidation process such as thiocyanate, cyanate and ammonia. Thiocyanate is formed by the reaction of cyanide and liable sulphur during the leaching process. The sulphur may originate from the attack of lime or cyanide on pyrrhotite, or be formed by air oxidation of sulphide minerals. Sulphur ions can be introduced by the dissolution of the readily soluble metal sulphide minerals. Thiocyanate is chemically and biologically degradable to produce ammonium, carbonate and sulphate as final products. These products especially ammonium ions at sufficiently high levels can become toxic to plant and animal life.

Oxidants can be used during the leaching process to increase the rate of dissolution of gold, but in turn also increase the oxidation of cyanide to cyanate. Common oxidants used are oxygen, chlorine, ozone and hydrogen peroxide. In acidic conditions cyanate will decompose to ammonia and carbonate via a hydrolysis reaction.

2.5 The Cyanide Cycle

There are several natural pathways via which cyanide will degrade in the natural environment. The degradation is due to complexation of cyanide with metal ions or cyanide complex precipitation to form stable precipitates. Other pathways include the adsorption of cyanides on minerals, the breakdown of cyanide to form other related compounds, such as cyanate and thiocyanate and then there is biodegradation of cyanide under aerobic conditions. The last two pathways that are worth discussion are volatilization and hydrolysis/saponification of hydrogen cyanide. This is illustrated well in Figure 2.3 showing the different pathways that can occur to render cyanide harmless to the environment. The following section will discuss each pathway in more detail.

12 (Whitlock and Mudder, 1986) 1986) Mudder, and (Whitlock Figure 2.3 : The cyanide cycle, cyanide : The 2.3 Figure

13 2.6 Degradation of Cyanide

2.6.1 Complexation

As discussed before cyanide can form many stable complexes with many different metals. The solubilities of these complexes range from very soluble to insoluble. Complexes that are not very stable will decompose to release free cyanide.

The formation of metal complexes reduces the toxicity by the removal of free cyanide in solution. Complexation of cyanide with transitional metals act as an intermediate for the formation of more stable compounds to remove free cyanide from being mobile in the environment; for example the iron cyanide double salts. These complexes can be immobilized from the environment by either precipitation or absorption onto organic or inorganic materials.

2.6.2 Cyanide Complex Precipitation

Formation of cyanide double salts as stable precipitates reduces the mobility of cyanide within the environment. Ferrocyanide and ferricyanide form very stable cyanide insoluble precipitates with a number transitional metals like iron, copper and zinc (seen in Table 2.3). Formation of these precipitates can occur over a wide range of pH from 2 to 11.

2.6.3 Adsorption

Adsorption of cyanide is another mechanism to remove cyanide from a waste stream. Soils that possess high anion exchange capacity would attenuate cyanide freely, for example soils that contain a high content of kaolin clay, chlorite, gibbsite clay, and/or iron and aluminium oxides and soils that contain cation exchange minerals like montmorillonite would have a lesser ability to uptake cyanide. Organic materials within soils and rock will absorb cyanide from solution. It is well known that

14 carbonaceous materials will uptake cyanide. Cyanide bound with copper, zinc or nickel has shown a higher absorption capacity on carbon.

2.6.4 Oxidation of cyanide to cyanate

The Pourbaix diagram representing the cyanide-water system (Figure 2.4) shows that cyanate is the predominant stable species within the cyanide-water system at 25oC. Cyanide converts to cyanate according to the simplified reaction (equation 2.4):

1 HCN + 2 O2 → HCNO [2.4]

Under natural conditions, the oxidation of cyanide to cyanate can be slow. An oxidant such as hydrogen peroxide, ozone or chlorine can be used to accelerate the oxidation. Alternatively the use of ultraviolet light in conjunction with catalysts like titanium dioxide, cadmium sulphide and zinc oxide has been shown to convert cyanide to cyanate (Domenech et al., 1989; Marsman et al., 1995; Dabrowski et al., 2002; Chiang et al., 2002 and 2003). It have been theorized that the ultraviolet light activates the catalyst to convert the dissolved oxygen within solution to ozone which then reacts with the cyanide.

15 2

OCN-

0 Eh (volts)

-1 HCN(a) CN- [CN]=0.001M aat 25 degrees Celcius

-2 0 2 4 6 8 10 12 14 pH

Figure 2.4: Eh-pH diagram for Cyanide-Water System at 25oC generated from Stabcal using equilibrium data from NBS Tables (Wagman et al., 1982).

2.6.5 Volatilization

Hydrogen cyanide is produced by a hydrolysis reaction between cyanide ion and water according to the following reaction (equation 2.5):

− − 2 +→+ OHHCNOHCN [2.5]

- o The pKa value for the CN /HCN equilibrium above is 9.22 at 25 C.

Most mines will operate leach solutions at a pH of 10.5 to 12. The barren solutions after metal recovery will have free cyanide remaining. With time the solution will absorb carbon dioxide, reducing the pH of the barren solutions and form hydrogen cyanide, which will slowly volatilise.

16 2.6.6 Biodegradation

Certain bacteria in the environment have been found to consume cyanide for a nitrogen source to produce ammonia and nitrates. Biodegradation has been successfully used at Homestake Mining Company, South Dakota USA, to detoxify cyanide waste effluents using a strain of fungus, Fusarium Lateritium (Smith and Mudder, 1993).

Under aerobic conditions it was reported that up to 200ppm cyanide could be converted to nitrogen in soils (Fuller, 1984), which provides an ideal source of nutrients for plants. Under anaerobic digestion research has shown that the bacteria will degrade cyanide to a much lower level (2ppm cyanide), but at higher cyanide concentrations (> 200ppm) for found to be toxic to the bacteria (Fuller, 1984).

The use of biodegradation to reduce the concentration of cyanide in the effluent during the primary treatment of tailing waters can encounter many problems. For example, biological organism cannot tolerate low temperatures causing the biological organisms to die and cyanide will be released into the tailing ponds. Then it will take some time to re-establish the environment to culture enough bacteria to consume cyanide again.

17 2.6.7 Formation of Thiocyanate

Many gold ores contain sulphide materials. In neutral or basic solutions, polysulphides and thiosulfate are produced from the oxidation of these sulphides. Free cyanide will react with these various polysulphides (equation 2.6) and thiosulfate (equation 2.7) compounds to form thiocyanate, as seen in the following reactions:

−2 − −2 − x + → −1)(x + SCNSCNS [2.6]

−3 − −2 − 32 3 +→+ SCNSOCNOS [2.7]

The formation of thiocyanate from different minerals that have been associated with gold bearing ores was studied by Smith and Mudder (1993). They concluded that the minerals that contributed to the formation of thiocyanate at a solution pH of 10 were chalcopyrite, chalcocite, pyrrhotite and free sulphur. Iron sulphide minerals form thiocyanate but to a lesser extent. The sulphur contained in sphalerite and pyrites is relatively non-reactive with cyanide at a pH of 10.

2.6.8 Hydrolysis and Saponification of HCN

HCN can undergo two different hydrolysis reactions at low pH to form ammonium formate (equation 2.8) and formic acid (equation 2.9). The extent of each product formed is dependent on the pH of the system. Lower pH favours the formation of formic acid. The two reactions are represented as:

2O2HHCN →+ HCOONH 4 [2.8]

HCN + 2H23 O → NH + HCOOH [2.9]

Cyanide degradation via the above reaction is not rapid. Hoecker and Muir (1987) have conducted high temperature hydrolysis reaction of cyanide in autoclaves. From these experiments they estimated the rate of cyanide hydrolysis to ammonium formate

18 of 4% per month by extrapolating the data from the high temperature experiments performed at room temperature. Similar results were found by DuPont with the rate of cyanide hydrolysis to ammonium formate of 2% per month (Longe and DeVries, 1988).

19 References

Caruso S.C., “The Chemistry of cyanide Compounds and their Behaviour in the Aquatic Environment”, Carnegie Mellon Institute of Research, June 1975.

Chiang K., Amal R., Tran T. “Photocatalytic degradation of cyanide using titanium dioxide modified with copper oxide.” Advances in Environmental Research, (2002), Vol. 6(4), pp 471-485.

Chiang K., Amal R., Tran T. “Photocatalytic oxidation of cyanide: Kinetic and mechanistic studies.” Journal of Molecular Catalysis A:Chemical, (2003), Vol. 193(1-2), pp 285-297.

Dabrowski B., Zaleska A., Janczarek M., Hupka J., Miller J. “Photooxidation of dissolved cyanide using TiO2 catalyst.” Journal of Photochemistry and Photobiology A: Chemistry, (2002), Vol. 151(1-3), pp 201-205.

Domench X., Peral J. “Cyanide photo-oxidation using a titanium(IV) oxide –coated zeolite.” Chemistry and Industry, (1989), Vol. 18, pp 606.

Fuller W. “Cyanides in the Environment with Particular Attention to the soil.” Cyanide and the Environment, Vol. 1, Colorado State University, Fort Collins, Colorado, 19-46, 1984.

Hoecker W. and Muir D., “Degradation of Cyanide.” The AusIMM Adelaide Branch, Research and development in Extractive Metallurgy, May, 1987.

Huiatt J, Kerrigan J., Olson F. and Potter G., Proceedings of a cyanide Workshop, Cyanide from Mineral Processing, U.S. Bureau of Mines, Salt Lake City, Utah, February 2-3, 1982

Izatt R.M. et al. “Thermodynamics of metal-cyanide coordination. I. PK, Ho, and So values as a function of temperature for hydrocyanic acid dissociation in aqueous solution”, Inorganic Chemistry 1: 828, 1962.

Longe G.K. and DeVries F.W., “Some Resent Considerations on the Natural Disappearance of Cyanide.” Economics and Practice of Heap Leaching in Gold Mining. Cairns Queensland, Australia, August 1988.

Marsman E., Appelman J., Tauw Milieu B., Deventer N. “Removal of complexed cyanides by means of UV-irradation and biological mineralization.” Soil and Environment (1995), Vol. 5 (Contaminated Soil 95, vol 2), pp 1295-1296.

Marsden J., House I. “The Chemistry of Gold Extraction”, Ellis Horwood, Chichester, UK. 1992.

Tran T., Lee K., Fernando K. and Rayner S. “Use of Ion exchange Resin for Cyanide Management During the Processing of Copper-Gold Ores”, Proceedings of Minprex 2000, International Congress on Mineral Processing and Extractive Metallurgy, 11- 13 September 2000.

Scott J. and Ingles J. “State-of-the-art Process for the Treatment of Gold Mill Effluents”, Mining, Mineral and Metallurgical Processes Division, Industrial Programs Branch, Environment Canada, Ottawa, Ontario, March 1987.

Smith A., Mudder T. “The Environmental geochemistry and fate of cyanide”, Proceedings of the Society of Economic Geologists Meeting, Denver, Colorado, April, 1993.

Wagman D., William E., Parker V., Schumn R., Halow I., Bailey S., Churney K.and Nuttal R. “The NBS tables of Chemical thermodynamic properties. Selected values for inorganic and C1 and C2 organic substances in SI units.” Journal of Physical and Chemical Reference Data. Vol. 11, 1982, Supplement No. 2. Published by American Chemical Society and the American Institute of Physics for the National Bureau of Standards.

Whitlock J.L. and Mudder T.I. “The Homestake Wastewater Treatment Process: Biological Removal of Toxic Parameters from Cyanidation Wastewaters and

21 Bioassay Effuent Evaluation”, R.W. Lawrence et al, eds., Fundamental and Applied Biohydrometallurgy, 1986, 327-339.

Young C., “Cyanide: Just the Facts.” Proceedings of Cyanide: Social, Industrial and Economic Aspects. 97-113. Annual Meeting of TMS (The Minerals, Metals and Materials Society) New Orleans, Louisiana, Feburary 12-15, 2001.

Chapter 3 Cyanide Management Using Chemical Oxidation Processes 3.1 Introduction

Since the development of the cyanidation process in the late 19th century there has been no other processes to rival its use in the industry. However, a number of accidents involving cyanide-contaminated wastes have affected the local environment. For example the cyanide spill at the Aural gold mine in Baia Mare, Romania became an international incident in early 2000 (DeVries, 2001). In this case, the dam containing cyanide-contaminated waste broke and cyanide was released into the Tisza River (which flows into the Danube) affecting one of the major waterways in Europe. This incident was compounded by the negative attitude toward mining activities worldwide. Tighter regulations to reduce the effect of future accidents have been implemented as a result. Greece, Turkey, and the state of Montana of the United States of America have banned the use of cyanide in new mining operations.

Environmental protection authorities in Australia have imposed strict regulations on the level of cyanide released into tailing dams. Most gold operators have to treat their cyanide waste prior to discharge into their tailing ponds. Within New South Wales, the Environmental Protection Authority specifies that the weak acid dissociable cyanide levels for discharge into tailing dams must be less than 30 ppm (Tran, 2000). On an international stage, the International Cyanide Management Code has established an upper limit of 50 ppm of weak acid dissociable cyanide to open pond areas (Ciminelli, 2002).

The control of cyanide in gold operations can be achieved by different techniques. The technique that can be used depends on the type of cyanide complexes present and the total level of cyanide in solution. There are two types of cyanide control that can be used. Either the waste cyanide is destroyed to products that are not as hazardous (for example cyanate), or recovered to form free cyanide to be recycled to be re-used within the operation.

The detoxification techniques that are commonly used by plant operations are based on the use of an oxidant such as hydrogen peroxide (Interox’s Efflox or Degussa process) or SO2/air mixture (Inco process). Earlier cyanide recovery processes used acid to regenerate free cyanide as HCN gas, which was adsorbed by a caustic or lime solution to produce sodium or calcium cyanide, respectively. An example of such a process is the Acidification-Volatisation-Regeneration (AVR) circuit. Compared with oxidation, AVR is expensive and less efficient when treating dilute tailing waters or slurries having less than 500 ppm cyanide due to the handling of large volumes of liquors or slurries. However, the AVR technique is more suitable for handling low volumes of concentrate cyanide streams produced from solvent extraction (Henkel’s SX-MNR process) or ion-exchange.

Cyanide recovery from thiocyanate has not been commercialised from lab-scale research. This technology would be suitable to cyanidation operations that have a high cyanide consumption that is attributed to thiocyanate formation. In Australia it is not required for a mine to treat the thiocyanate before releasing it into tailing dams. In the reaction to recover cyanide from thiocyanate carbon is reduced from C(IV) to C(II) while sulphur is oxidised from S(-II) to S(VI) (Botz 2001).

This chapter reviews current chemical oxidation processes that are used to control cyanide and cyanide related products. It examines the different management techniques from cyanide destruction to recovery of cyanide from cyanide waste streams with particular attention to cyanide recovery from thiocyanate.

24 3.2 Cyanide Destruction

Detoxification of tailings involves mainly the destruction of weak-acid dissociable cyanide species from solution. This is achieved by a wide range of processes such as oxidation to cyanate (CNO-), using chlorine, hydrogen peroxide or sulphur dioxide/air mixtures; precipitation as copper-iron cyanide complexes and biological treatment. Chlorine when used can form potential toxic chloro-cyanide species and its use to destroy cyanide is unfavoured in mining operations. The commonly used oxidation techniques practised within the industry are based on hydrogen peroxide and sulphur dioxide/air mixture. These oxidation techniques have been used extensively around the world to destroy cyanide-contaminated wastes because of proven process performance, low capital and running costs and relatively simple plant design.

3.2.1 The INCO process

The INCO process (named after the International Nickel Company) uses sulphur dioxide-air mixtures with copper (as a catalyst) to oxidise weak-acid dissociable cyanides to cyanate. Removal of free cyanide and complex metal cyanides species from wastewater streams is very slow with SO2 and air alone. Copper catalyses the removal of free cyanide (equation 3.1), complex heavy metal cyanide (equation 3.2) and thiocyanate (equation 3.3) species from the stream. Normally the SO2/air process is configured as a single stage continuous treatment facility in which the reagents are dispersed into solution using a well agitated vessel. The process can be applied to pulps and clear wastewaters. The process was first patented in 1984 (Borberly et al., 1984) and has been used successfully at many mines over the world.

− − 2 2 2 +→+++ SOHCNOOH(g)O(g)SOCN 42 [3.1]

−xy − y+ M(CN)x + 2 (g)SOx 2 (g)Ox 2OHx →++ CNOx + 42 + MSOHx [3.2]

− − SCN 2 2 2 +→+++ SO5HCNOO5H(g)4O(g)4SO 42 [3.3]

One of the benefits of the INCO process is that the sulphur dioxide used can be provided in many different forms. The least expensive is to generate sulphur dioxide on site from roaster gas or burning elemental sulphur. Otherwise reagents such as liquid SO2 and sodium sulfite or sodium metabisulfite can be brought into the operation (Devuyst 1991). A typical process removes total cyanide to less than 2 mg/l, from feeds containing total cyanide in the range of 50-2000 mg/l (Robbins et al. 2001).

The preferred oxidation occurs between a pH from 5 to 10 with the introduction of lime or caustic to neutralise the acid produced within the redox reaction. If the pH of the system is too low the rate of oxidation of cyanide is reduced. Therefore the theoretical consumptions for sulphur dioxide and lime are 2.5mg SO2/mg CNWAD and

2.2 g CaO/mg CNWAD. Many operations increase the level of reagents several times higher to increase the rate of oxidation of cyanide reducing reaction time to less than an hour.

This process is not effective for the removal of thiocyanate due to increased consumption of SO2. It was found that with an addition of nickel (as a catalyst) with or without copper a higher concentration of SO2 of about 4.5 mg SO2/mg SCN is required to destroy thiocyanate (Borbely et al., 1984). Any iron cyanide present in the wastewaters is removed by precipitation as double salts with copper, zinc or nickel as with other oxidation processes. The presence of iron cyanide affects the level of copper added as catalyst because copper is consumed in the formation of different iron cyanide double salts. Copper is normally added as copper sulphate at 50 mg/l or higher if iron cyanide is present.

26 3.2.2 Peroxygens

Peroxygens are a group of powerful oxidants that have been used to detoxify cyanide waste streams. The most commonly used peroxygens in the mining industry for cyanide detoxification include hydrogen peroxide (Degussa process), Caro’s acid (Efflox process), and ozone.

The Degussa process uses hydrogen peroxide as the oxidant in the destruction of cyanide. Hydrogen peroxide reacts with cyanide to produce cyanate (equation 3.4) and where added in excess the products produced are carbonate and nitrite (equation 3.5). Nitrite can continue to be oxidised to form nitrate (equation 3.6). Otherwise, in acidic conditions, cyanate will hydrolyse to produce ammonium and bicarbonate ions (equation 3.7).

− − 22 +→+ 2OHCNOOHCN [3.4]

− − 2− + 22 2 3 2 +++→+ 2HO2HCONOO3HCNO [3.5]

− − 222 +→+ 23 OHNOOHNO [3.6]

− + − − 2 NHO3HCNO 4 3 ++→+ OHHCO [3.7]

This process has great success in removing most weak acid dissociable cyanides but has no effect on thiocyanate and strong-acid dissociable cyanides. Even with this drawback many plants use this method as the primary destruction process as well as a stand-by process for emergencies (Young et al., 1995).

Other reagents have been used to increase the reaction rate and efficiencies of hydrogen peroxide. The most common additive is cupric ions. The copper acts as a catalyst for the oxidation when hydrogen peroxide is the oxidant and is normally not consumed. However, if iron cyanide is present within the waste, the copper will be

27 consumed due to the formation of the copper-iron cyanide double salt. Degussa technology was used at the Ok Tedi gold mine in Papua New Guinea, using 70% hydrogen peroxide at 2.0 to 4.5 times the stoichiometric requirements and copper sulphate catalyst added at 10-20% of the cyanide concentration at a pH of 9.0 to 9.5. All the weak acid dissociable cyanides were oxidised to cyanate, and the metal ions released during oxidation were precipitated as hydroxides.

Another addition to the process has been developed by Dupont to increase the efficiencies of the oxidation of cyanide with hydrogen peroxide containing the same starting levels of cyanide. This was the addition of formaldehyde/Kastone to the process. Cyanide will react with formaldehyde (HOCH) (equation 3.8) in the presence of Dupont’s Kastone reagent to form glycolnitrile (HOCH2CN). In the presence of hydrogen peroxide glyconitrile hydrolyses to glycolic acid amide

(HOCH2CONH2) (equation 3.9) (Young et al., 1995).

− − CN + HOCH(aq) 2OH →+ HOCH 2CN(aq) + OH [3.8]

HOCH 2 2 →+ HOCHOHCN(aq) 2 2 (aq)CONH [3.9]

3.2.2.1 The Efflox Process

The Efflox process was developed by Interox Chemicals and has been used extensively in Australia at mines like Kanowna Belle in Western Australia and Beaconsfield in Tasmania. The process uses peroxymonosulphuric acid, otherwise known as Caro’s acid. This oxidant is made by the reaction of hydrogen peroxide and sulphuric acid (equation 3.10). An advantage of using this oxidant is that it is stronger than hydrogen peroxide (shown in equations 3.11 and 3.12).

4222 →+ + 252 OHSOHSOHOH [3.10]

−+ o 22 →++ 2O2H2e2HOH =1.77E Volt [3.11]

−+ o 52 +→++ 242 OHSOH2e2HSOH = 81.1E Volt [3.12]

With a benefit of having a higher oxidation potential, Caro’s acid has been reported to have greater stability than hydrogen peroxide, in the presence of transition metals. This leads to higher efficiencies in terms of oxidant consumption for particular processes (Nugent et al., 1991).

Nugent et al. (1991) compared the rate of oxidation of cyanide using hydrogen peroxide and Caro’s acid. The experiments were carried out in solution and slurries containing with the same starting levels of cyanide. Each experiment was conducted batch wise on a 250ml to 1000ml scale. The hydrogen peroxide was added as a 50% w/w solution to a portion of cyanide liquor or to 45% w/w solid slurry. The results are shown in Figure 3.1 and Figure 3.2, which compare the final weak-acid dissociable cyanide concentration when using hydrogen peroxide or Caro’s acid dosed at different levels into either the solution or slurries. For comparison, the Caro’s acid doses are expressed in equivalent units of hydrogen peroxide. It can be seen that for the equivalent amount of hydrogen peroxide used, Caro’s acid is more efficient in destroying weak acid dissociable cyanides. The treatment was also applied to slurry solution (Figure 3.2) and although the consumption of reagents was higher as expected, Caro’s acid is still more effective than hydrogen peroxide alone. Caro’s acid experiments showed that in either solutions or slurries the residual concentrations of weak-acid dissociable cyanides were similar. While when using hydrogen peroxide the concentration of residual weak-acid dissociable cyanides in the slurry experiments were higher than with experiments conducted in solutions.

29 6 Hydrogen peroxide Caro's acid 5

4 ) WAD 3 ln(CN

0 0 5 10 15 20 25

Dose equivalence of oxidant (Mole H2O2/Mole CNWAD)

Figure 3.1: Illustration of the difference between hydrogen peroxide and Caro’s acid treatment of weak-acid dissociable cyanides in real liquor solutions (Nugent et al., 1991, Case study 1).

Hydrogen Peroxide 5 Caro's Acid

4 ) WAD 3 ln(CN

0 024681012

Dose equivalence of oxidant (Mole H2O2/Mole CNWAD) Figure 3.2: Illustration of the difference between hydrogen peroxide and Caro’s acid treatment of weak-acid dissociable cyanides in slurries (Nugent et al., 1991, Case study 1).

30 3.2.2.2 Ozone

Ozone is another commonly used oxidant that can be used for cyanide management (Sondak et al., 1967; Rowley et al., 1980; Zeevalkink et al. 1980 and Carrillo-Pedroza et al., 2000). Ozone is a superior oxidant to oxygen and has two mechanisms to oxidise cyanide to cyanate: ozonation (equations 3.13) and catalytic ozonation (equation 3.14).

− − 3 (aq)OCN +→+ 2 (aq)OCNO [3.13]

− − 3 (aq)O3CN →+ 3CNO [3.14]

Cyanate can undergo further oxidation with ozone to form carbonate and nitrogen gas (equation 3.15 and 3.16) (Young, 1995), and because of this cyanate does not undergo the hydrolysis reaction to produce nitrites or nitrates in the system.

− − 3 2 →++ + 23 (g)N2HCOOH(aq)O2CNO [3.15]

− − 2CNO 3 2 →++ 23 ++ 2 (aq)3O(g)N2HCOOH(aq)3O [3.16]

There is an operational limit when using ozone, as the efficiency of oxidation drops when the pH is greater than 11 (Young, 1995). This is because hydroxide has been found to decompose ozone. Most mine operations will operate leach solutions in alkaline conditions, typically above 11 to prevent the loss of cyanide by volatilisation, as seen in Figure 2.1.

The mechanism that controls the reaction is the rate of mass transfer of ozone (Garrison et al., 1975 and Mathieu, 1977). Work reported by Zeevalkink in 1980 showed that the presence of cyanide increased the mass transfer of ozone into the aqueous cyanide solution by 1.5 times compared to pure water alone. Carrillo- Pedroza (2000) found that as long as cyanide was present in solution the redox potential, pH level of the system and dissolved ozone level remained constant. When

31 cyanide was completely oxidised the pH of the system decreased and the redox potential of the system increased to around 600mV. The change was attributed to other reactions between ozone and other species present such as cyanate and ammonia. Figure 3.3 shows the destruction of cyanide with respect to ozone addition and the continued oxidation to cyanate.

1.2

0.8 (mol) - 0.6 ;CNO - CN 0.4 Cyanide Cyanate 0.2

0 0 0.2 0.4 0.6 0.8 1 1.2 1.4 1.6 1.8 Ozone added (g) Figure 3.3: Oxidation of cyanide to cyanate using ozone as an oxidant. Operating conditions: o T=27 C, Flow rate of gas =4 L/min, Vsg=2.92 cm/s, ozone addition rate=0.11g/min. (Carrillo- Pedroza, 2000).

32 3.2.3 Chlorination

One of the first techniques developed for cyanide remediation was chlorination, which has been used in the gold mining industry since 1889 (Young, 1995). Alkaline chlorination is economical and efficient for the removal of free and weak-acid dissociable cyanides. The disadvantages of using chlorination to control the level of cyanide in an operation are the high reagent cost associated with controlling the pH for optimal oxidation and high discharge concentrations of chlorine or hypochlorite leaving the cyanide management system, which are as toxic to the environment as cyanide. Chlorine is no longer the preferred process due to the potential formation of toxic chloro-cyanide species. It has been replaced by oxidation techniques which employ hydrogen peroxide or sulphur dioxide/air.

The oxidation of cyanide to cyanate using chlorination occurs by the formation of an intermediate species. Firstly, chlorine reacts with cyanide (equation 3.17) to form cyanogen chloride (CNCl) and in alkaline conditions cyanogen chloride is converted to cyanate (equation 3.18). The overall reaction from cyanide to cyanate is described in equation 3.19.

− − 2 (g)ClCN →+ CNCl(aq) + Cl [3.17]

− − − CNCl(aq) 2OH 2 ++→+ ClOHCNO [3.18]

− − - − 2 2OH(g)ClCN 2 ++→++ 2ClOHCNO [3.19]

When chlorine is in excess, cyanate continues to oxidise to produce nitrogen and carbon dioxide (equation 3.20). This reaction consumes large amounts of chlorine and hydroxide causing problems in the release of hydrogen cyanide and cyanogen chloride if the pH of the system drops below pH of 10.

− − − 2CNO 2 2 +→++ 2 ++ 2O2H6Cl(g)2CO(g)N4OH(g)3Cl [3.20]

Chlorine can be used to convert thiocyanate and weak acid metal cyanides to cyanate. Thiocyanate reacts with chlorine with a higher consumption rate compared to cyanide (equation 3.21). The sulphur forms sulphate within the oxidative environment.

− − − 2− − 2 10OH(g)4ClSCN 4 +++→++ 2O5H8ClSOCNO [3.21]

Because of the alkaline conditions for the conversion of cyanogen chloride to cyanate, the weak-acid dissociable metal cyanides form metal hydroxides (equation 3.22). Strong acid dissociable cyanides are not oxidised by chlorination.

−xy − − − M(CN) x 2 +++ y)(2x(g)Clx OH → CNOx Cl2x ++ M(OH) y + 2OHx

where x > y [3.22]

Lancy et al. (1960) studied the rate of reaction for continuous oxidation of cyanide wastes using chlorination. The chlorination was carried out in two reaction vessels. The primary reactor where the majority of the oxidation took place required reaction time of 15-30 minutes with an operational pH of 10-11 and excess active chlorine. The size of the reactor depended on the presence of copper and nickel. Secondary stage of oxidation required a shorter retention period of 15 minute with neutral pH.

Lancy et al. (1960) found that hypochlorite was produced during the chlorination process (equation 3.23). They observed that hypochlorite could oxidise cyanide. The reactions that occur with hypochlorite are very similar to chlorination. Hypochlorite will oxidise free and weak-acid dissociable metal cyanides and thiocyanate. Other cyanide compounds will not react readily with hypochlorite. Research done by Baden (1973) on the use of hypochlorite in the oxidation of iron cyanides demonstrates the reaction kinetics are very slow. Addition of silver nitrate accelerates the oxidation rate but the rate is still slow in comparison to that for weak-acid dissociable or free cyanides.

− − − 2 ++→+ 2OHOClCl2OH(g)Cl [3.23]

The reaction with hypochlorite, as with chlorination, is a two-step reaction in which there is the formation of an intermediate species cyanogen chloride, which forms cyanate under alkaline conditions (equation 3.24 and 3.25).

− − − 2OHOClCN →++ CNCl(aq) + 2OH [3.24]

− −− CNCl(aq) 2OH ++→+ 2OHClCNO [3.25]

When in excess hypochlorite solution, as with chlorination, cyanate will react with hypochlorite to produce nitrogen and carbon dioxide (equation 3.26).

− − − − 2CNO 2 2 +→++ 2 ++ 2OH3Cl(g)2CO(g)NOH3OCl [3.26]

The following reactions show the stoichiometry for the reactions of thiocyanate and weak acid dissociable metal cyanides with hypochlorite (equations 3.27 and 3.28).

− − − − 2− − SCN 4OCl 2OH 4 +++→++ 2OH4ClSOCNO [3.27]

−xy − − − − M(CN) x OClx OHy →++ CNOx Clx ++ M(OH) y where x > y [3.28]

35 3.3 Cyanide Recovery of free and weak-acid dissociable cyanides

Although several technologies have been developed for the recovery of cyanide from tailing solutions, only the Acidification-Volatilisation-Reneutralisation (AVR) circuit has been commercialised successfully. Considerable efforts in recent years have been spent in improving techniques for the recovering of cyanide producing many promising processes to reclaim cyanide for example Cyanisorb, AuMENT, Vitrokele, MNR(SART) and the Elutech processes.

For the long term at the mine site if cyanide were recovered from the beginning of the mines life, rehabilitation of the site becomes fairly simple because there are no significant amounts of cyanide related product present in the mine wastewaters. It can be a benefit for the operation because recovery circuits, and by the sale of metals recovered as by-products, the cost to recover cyanide can become less expensive than cyanide. The availability of such options makes the recovery of cyanide a very attractive proposition.

The recovery of cyanide from thiocyanates can be costly because of the higher reagent consumption required. At present, legislation in Australia does not limit the amount of thiocyanate that can be released into tailing waters. However, in view of the growing practices and the possibility of more stringent legislation, the elimination of thiocyanates from tailing waters poses a potential problem for mine operators.

3.3.1 Acidification-Volatilisation-Reneutralisation (AVR) Circuit

The AVR circuit has been known since 1929. It was first referred to as the Mille Crows Process, but since the 1970’s, after developments by CANMET and McNamara, it has been referred to as the AVR process (Tran et al., 2000). This process takes advantage of the volatile nature of hydrogen cyanide at low pH. Many other cyanide recovery techniques use a variation of the AVR circuit, either using it as the recovery step for cyanide or to concentrate the cyanide to a useable concentration.

36 A considerable effort has been made to improve the performance of AVR since its introduction. It has become accepted by the industry as an option for treating moderate or strong cyanide liquors (over 500 ppm CNWAD). However, copper is precipitated as copper cyanide (CuCN) during the acidification stage and is an unsaleable product which binds up the available cyanide, leading to a recovery of 67% of the maximum amount of cyanide available. Equations 3.29 and 3.30 describe the recovery of cyanide from copper cyanide and then the recovery of cyanide in alkaline solutions.

2− 2− Cu(CN)3 SOH 42 →+ CuCN(s) + 2HCN(g) + SO 4 [3.29]

2HCN(g) + Ca(OH)2 → Ca(CN)2 + 2OH2(aq) [3.30]

The cyanide associated with thiocyanate cannot be recovered by acidification alone. This becomes an issue when thiocyanate levels are too high in leach liquors. For example, it has been reported that levels of thiocyanate in mine liquors have reached levels in excess of 2000ppm (Soto et al., 1995). In such cases, the cyanide that has been consumed in the formation of this by-product becomes economically significant.

The AVR circuit involves acidification of the cyanide liquors or slurries to lower the pH of the liquor or slurry from pH ~ 10 to pH ~ 3 to convert free cyanide and weak acid dissociable cyanides (complexes of Zn, Cd, Ni) to HCN. The HCN is then volatilised by passing air bubbles through the liquor or pulp. The HCN/air stream is scrubbed in a lime slurry or sodium hydroxide solution to convert HCN into free cyanide. The process involves a series of aeration columns as illustrated in Figure 3.4. It consists of two process units, one to volatilise hydrogen cyanide using a carrier gas and the second column is used to absorb the hydrogen cyanide gas into an aqueous alkaline solution. The system is a closed loop circuit with respect to cyanide. A number of carrier gases can be used with the preferred gas being air. Other gases that have been suggested include nitrogen and carbon dioxide from other processes.

The US patent 5,254,153 (Mudder et al., 1993) suggested the stripping tower operated under negative pressure and the air-HCN mixture being positively removed. If

37 packing is used the operating pressure drop is about 15 to 30 mm water per meter of packing material. The flow rate of air needed to carry the cyanide from the waste stream varies from 250 to 1000 cubic meters of air per cubic meter of solution. The patent states a preferred operation flow rate of 300-800 m3 air/m3 solution, which corresponds to a flux of 2.8 to7.4 cubic meters of air per square meter of pH adjusted mixture per minute, based on a period of 3 to 4 hours.

At the May Day Mines, Cobar, N.S.W. Australia, the stripping and adsorption towers were packed with plastic media to improve the recovery of cyanide. The preferred packing materials have a size from 50 to 75 mm and are made from either stainless steel, ceramic materials or plastics. The most commonly used packing materials are Pall rings, Rashig rings, Tellerette rings, saddles and grids. This system at May Day Mines was designed and used to effectively handle 100 tonne of cyanide per hour. Air was used as the carrier gas with a flow rate of 10,000 m3/hour (Figure 3.5).

A similar process based on the AVR circuit has been marketed by Unifield Engineering (USA) (called the Cyanisorb process). A number of US patents covering the process (US 4,994,243, US 5,078,977 and US 5,254,153) have been filed. The process is slightly different from the original AVR circuit in that clear solutions or slurries are processed at near neutral pH.

Air Acid Caustic Solution

Cyanide Waste

Acidification Tank Volatilisation Cyanide Zone Recovery Tower

Recovered Cyanide Solution

Air Air Blower

Figure 3.4: Flow diagram of the AVR Circuit.

Figure 3.5: May Day Mine’s AVR Circuit (left) and Resin Columns (right).

39 3.3.2 MNR (SART) Process

The MNR process was developed by Metallgesellschaft Natural Resources in Germany. It is also known as the SART process which stands for Sulphidisation, Acidification, Recycle and Thickening process. As implied in the name of the process, it uses sulphidisation and acidification (to less than pH 5) of copper/cyanide rich liquors to precipitate copper as synthetic chalcocite (Cu2S) and recover free cyanide.

The acidification and sulphidisation steps are conducted simultaneously as follows:

2− 2− 2Cu(CN)3 242 →++ 2S(s)CuSHSO2H + 6HCN(g) + 2SO 4 [3.31]

The main advantage of using this method over other recovery processes is that all the cyanide from the cyanide complex can be recycled. The copper precipitate is a marketable product which can be processed into to copper metal. The copper sulphide can also be blended with other copper flotation concentrates produced on site as in the Telfer mine in Western Australia (Dreisinger et al., 2001).

The SART process produces a slurry of copper sulphide precipitate and free cyanide. This slurry is fed to a thickener where the copper sulphide is recovered in the underflow and either added to other copper flotation concentrates or refined for sale. The overflow from the thickener is adjusted with an alkaline solution to raise the pH of the solution to convert the hydrogen cyanide to sodium or calcium cyanide. Another option is to acidify the overflow further to form gaseous hydrogen cyanide and recover cyanide via adsorption columns (Davis et al., 1998; Dreisinger, 1998). Figure 3.6 shows a flow sheet of a pilot scale SART plant (Dreisinger, 2001).

Dreisinger (2001) conducted pilot plant trails using the SART process for the Telfer Mine. They found that the reaction is nearly stoichiometric with respect to copper, where the sulphur additions were able to remove 95% of the copper from solution. The sulphuric acid consumption was 1.7-2.4 kg of 100% sulphuric acid/kg weak acid dissociable cyanide to achieve a pH of 4-5. At this pH range the recovery of copper

40 was over 95% and the cyanide recovery was over 94%. Dreisinger (2001) also tested

H2SO3 as an alternative to sulphuric acid. The use of H2SO3 was not as successful, leaving over 5 times more copper in solution compared to sulphuric acid. It was proposed that H2SO3 reacted with the sulphur and formed a copper complex preventing it from precipitating from the solution.

Figure 3.6: Pilot scale SART plant flow sheet. (Dreisinger, 2001).

Free cyanide can react with sulphur to form thiocyanate. This represents a disadvantage of the process because of the loss of recovered cyanide and the formation of copper thiocyanates (soluble and insoluble species), which can contaminate the final product.

3.3.3 Copper Electrowinning Process

This process was developed to recover copper metal and cyanide from dilute copper cyanide wastes. Fleming et al. (1995) have filed a patent (US patent 5,411,575) which describes a process to electrowin copper from solution using a membrane-type electrowinning cell. The cell uses a sodium permeable ion-exchange membrane (Nafion 417) developed by Du Pont. The membrane separates the two cell

41 compartments, which prevents the oxidation of cyanide to cyanate at the anode. The anolyte solution is made from concentrated sodium hydroxide solution. This provides sodium ions to the cell for cation transfer through the membrane from the anolyte to the catholyte solution. Copper is electroplated from the catholyte solution leaving the cyanide in the catholyte. The half-cell reactions for the cell are shown below (equations 3.32 to 3.35):

−+ 4NaOH:Anode 2 2 +++→ 4e4NaO2H(g)O [3.32]

4Na:Membrane + (anolyte) → 4Na + )(catholyte [3.33]

+ − + 4Na4Na:Cathode 2Cu(CN)3 4e 4Cu(s) +→+ 12NaCN [3.34]

4Na:Overall 2Cu(CN)3 + 4NaOH → 2 + 2 + 4Cu(s)O2H(g)O +12NaCN [3.35]

The consumption of sodium hydroxide has been reported to be significantly higher than reaction stoichiometry, which can be attributed to the side reaction of hydrogen formation at the cathode (Lu et al., 2002). To reduce the effect of these side reactions, Lu et al. (2002) proposed a second system using sodium chloride. As a by-product chlorine can be formed at the anode. The half-cell reaction as the anode is given be equation 3.36 and the overall reaction shown equation 3.37:

+ − 2NaCl:Anode 2 ++→ 2eCl2Na [3.36]

4Na:Overall 2Cu(CN)3 4NaCl →+ 4Cu(s) +12NaCN + 2Cl2 [3.37]

Lu et al. (2002) used a flow through electrolytic membrane cell to show the effect of the cyanide to copper ratio on the recovery of copper and cyanide. The electrolytic cell used a graphite-felt anode compressed between painted stainless steel mesh, shown in Figure 3.7. Their main conclusion was that the current efficiency of copper plating on the felt decreases with an increase in the cyanide to copper ratio and

42 current density. The accumulation of copper deposited on the felt increases the specific surface area of the cathode thus increasing the conductivity of the felt, which improves copper deposition. Lu et al. (2002) reported that 60% of the copper was successfully removed from diluted copper cyanide solutions (~1g/l copper). The energy requirement for the cell was 1-2 kWh/kg Cu at a current density of 30-100 A/m2.

Inlet/Outlet Inlet/Outlet Anolyte Anolyte Anode Anode Cathode Thermometer

Catholyte inlet Catholyte outlet

Anode Anode Chamber Chamber

Membrane Membrane Cell Cell

Figure 3.7: Schematic diagram of the electrolytic cell described by Lu et al. (2002).

43 3.4 Pre-concentration Processes

Two types of processes that have been used to concentrate cyanide waste streams from mine waste are solvent extraction and ion-exchange resin based technologies. Solvent extraction is widely used in the processing of copper ores. Ion-exchange resins have been used in the mining industry since the 1980’s to recover gold from gold liquors at the Golden Jubilee Mine in South Africa (Fleming et al, 1984 (A); 1984 (B); 1990) and earlier in the former USSR states (Bolinski and Shirley 1996; Marsden and House 1992).

The processes that were developed focused on the recovery of precious metals such as gold and silver cyanides. Later studies focussed on the recovery of precious metals and cyanide species, especially the copper cyanide species. The early resin-based technologies have used basic eluants such as thiocyanate, chloride and hydroxide to remove gold cyanide from the loaded resin for further processing. However, in the presence of base metal cyanide complexes in the liquor, sulphuric acid is also used to strip these complexes off the resin. Acid elution employed in several ion exchange resin processes destroys the cyanide complexes loaded on the resin, regenerating cyanide for recycling via hydrogen cyanide formation.

3.4.1 Resin Technologies

3.4.1.1 AuGMENT Process

The AuGMENT process (Fleming et al. 1998, Le Vier et al. 1997) uses commercially available strong-base ion-exchange resin to recover and pre-concentrate free cyanide and copper cyanide species. The resin is pre-treated by impregnating the ion- exchange resin with CuCN(s). This has been found to promote the adsorption of copper cyanide species and free cyanide. The affinity of the resin for copper cyanide, -1 in particular, the Cu(CN)2 species, is higher than free cyanide itself. The loading mechanism can be described by equation 3.38:

44 2− 2− 1− 1− − 2− 2 − SO2R 4 (s) + Cu(CN))(CuCN 3 2CN 3R −→+ Cu(CN)2 +−+ 2SOCNR 4 where R represents the resin backbone and functional group. [3.38]

Once resin is loaded, copper cyanide can be eluted from the resin using a copper cyanide/caustic eluant (CNTotal:Cu ratio between 3.5:1 to 4.0:1 and a copper concentration between 10 to 70 g/l Cu and 10 g/l NaOH) (Davis et al., 1998). With short elution times the resin is substantially eluted and the resulting eluant having a

CNTotal:Cu ratio of <4:1. Elution is described by equation 3.39:

1− 2− 1− 2− 2− 2R − Cu(CN)2 + Cu(CN)3 R2CN 2 −→+ Cu(CN)3 + Cu(CN)3 [3.39]

As part of the AugMENT process, copper is produced from copper cyanide using the Du Pont membrane cell (as described earlier). Gold has to be recovered prior to copper electrowinning and cyanide recovery. Cyanide can also be recovered via an AVR circuit, where the copper cyanide is precipitated and re-dissolved in the loaded catholyte ahead of the electrowinning circuit.

3.4.1.2 Vitrokele Process

The Vitrokele process uses a proprietary strong-base ion-exchange resin to recover waste cyanides from either clear solution or slurries. The resin is based on a highly cross-linked polystyrene structure (VitrokeleTM 911 and 912). Ligands bound to the resin structure have not been publicly disclosed but V912 is believed to be a type 1 quaternary amine functionality (Jay W., 2000).

Once loaded, base metal cyanides are eluted using a strong cyanide eluant (equation 3.40). The precious metal cyanides and other strongly bound complexes are eluted from the resin using a zinc tetracyanide eluant (equation 3.41). The elution step using the zinc tetracyanide eluant is only required every second or third base metal elution depending on the relative loadings of the precious and other strongly bound species. The other elution step uses sulphuric acid as an eluant, which removes most cyanide complexes. An AVR circuit is applied to recover the cyanide for recycling in the

45 leaching plant or used as an eluant in one of the stripping cycles (equation 3.40 and 3.41).

− 2− R 2 − Cu(CN)3 2CN CN2R +−→+ Cu(CN)3 [3.40]

2− − 2R 2R − Au(CN)2 + Zn(CN)4 R 2 −→ Zn(CN)4 + 2Au(CN)2 [3.41] where R represents the resin backbone and functional group.

This process has been successfully applied at the Connemara Mine in Zimbabwe for processing oxidised ores using heap leaching (Satalic et al., 1996). During the period from July 1997 to June 1998, this process was tested on copper-gold ores at May Day Mines, Cobar, N.S.W. Australia under sponsorship of an Industrial Research & Development syndicate scheme (Paterson et al., 1997).

At the May Day Mine, the Vitrokele technology was used in a vat-leach operation to treat 250 000 tonnes of agglomerated ore. The ore was graded at 2.2 gpt Au, 15.2 gpt Ag and 0.12% Cu. The resultant leach solution contained 2-3ppm Au, 25ppm Ag, 50ppm Zn, 200ppm Cu and <1ppm Fe. As report earlier a strong cyanide (100g/l NaCN, 10 g/l NaOH) eluant was used to elute copper from the resin. This step produces a copper liquor in a strong cyanide background, which is then acidified to precipitate the copper as copper cyanide. The hydrogen cyanide is regenerated using an AVR circuit (as seen in Figure 3.5).

Gold was eluted from the resin and recovered from an electro-elution circuit, similar to the operation in the former Golden Jubilee plant in South Africa. Zinc tetracyanide at 50-60oC was used as the eluant to strip gold from the resin in a closed loop in which gold is recovered by electrowinning immediately after elution (Marsden et al 1992). For efficient elution of gold from the resin, the gold concentration has to be maintained <1ppm Au. Even under these conditions the gold elution has been reported to take as long as 48 hours the Golden Jubilee (Fleming et al., 1984, 1990). Similar elution times were used at the May Day site.

46 Several technical problems developed during the operation of the Vitrokele process that led to the abandonment of the Vitrokele Process at the May Day Mines. One of the major problems was that the base metal elution was not effective to remove most of the copper from the resin. The interaction between residual copper cyanide and sulphuric acid eluant lead to the formation of solid copper cyanide in the resin, which passivated the resin surface and active sites. Other species such as insoluble copper- iron or zinc-iron cyanide double salts have the same effect. These precipitates can form at neutral or acidic conditions (Tan, 2000).

Imperial Mining concluded its Syndicate R&D project at the May Day Mines on 30 June 1998. In their June 1998 Quarterly Report to the Australian Stock Exchange, Imperial Mining stated:

“The R&D program clearly proved that the Vitrokele technology was not able to cope with the extraction of copper cyanides adsorbed onto the Vitrokele 912 resin. This critical failure affected almost all other crucial unit operations from adsorption of metals from pregnant liquor solution onto the resin to elecro-elution of the precious metals. Above all, the removal of copper from the circuit and the recovery and regeneration of sodium cyanide within the plant became impossible to control.”

3.4.1.3 Elutech Process

After the abandonment of the Vitrokele process at May Day Mines, another elution technique was developed to overcome the problem of the elution of copper cyanides from commercially available ion exchange resin (Tran et al., 2000). The improved elution technique was developed based on an acid-oxidation technique. Using hydrogen peroxide as the oxidant in an acidic media, copper is oxidised to cupric ions, thus destroying the copper cyanide complex and liberating hydrogen cyanide (equation 3.42). An AVR circuit recovers the cyanide. Once the copper cyanide is oxidised the copper is no longer attracted to the resin.

1− + 2+ 2Cu(CN)2 22 2Cu6HOH ++→++ 2O2H4HCN [3.42]

47 The advantage of this elution technique was that in every cycle conducted at the May Day operation all the adsorbed copper was eluted from the resin. However, in the presence of iron cyanides, cupric ions will readily form stable copper-iron double salts, which can passivate the ion exchange resin.

After its elution from the resin, copper can be recovered either by precipitation as copper hydroxide or by solvent extraction/electrowinning. For the removal of precious metals from the resin, May Day used conventional using zinc tetracyanide eluant to elute the precious metals followed by electrowinning.

3.4.1.4 Cognis AURIX Resin

Gold selective resins have been developed to process gold liquors. One such resin is Cognis Aurix resin that contains guanidine-based ligands. These resins have an intermediate strength between weak base and strong base ion exchange resins. They exhibit a pKa of approximately 12 and are capable of being protonated at the operating pH of liquors from gold plants (Jay, 2000). The resin is easily deprotonated by 1M sodium hydroxide, thus it can be stripped easily by a change in pH (Mackenzie et al., 1995).

The chemistry of the loading and stripping can be illustrated in the following reactions (equation 3.43 & 3.44):

Loading [3.43]

NH2 NH2 - RN=〈 + H2O ⇔ RN(H)=〈⊕ + OH

NH2 NH2

NH2 NH2 - - - - RN(H)=〈⊕ + OH + [Au(CN)2] ⇔ RN(H)=〈⊕ . [Au(CN)2] + OH

NH2 NH2

48 Elution [3.44]

NH2 NH2 - - - RN(H)=〈⊕ .[Au(CN)2] + OH ⇔ RN(H)=〈 + H2O + [Au(CN)2]

NH2 NH2 where R represents the resin backbone.

3.4.1.5 Dow-Mintek (Minix) Ion exchange resin

This particular resin was developed for the gold industry to be gold selective over other base metal cyanides. To increase the selectivity of gold over copper cyanide, the theoretical ion exchange capacity of the resin was lowered to 0.7-0.9 mmol/g, which is 25% of the capacity of other strong base ion exchange resins (Jay, 2000). This Minix resin does not readily absorb metal cyanides like cobalt and iron. The eluant that is used to elute the gold from the resin is acidic thiourea, because zinc cyanide is poorly adsorbed.

3.4.1.6 Superhydrophilic Ion Exchange Polymers

Another alternative is the use superhydrophilic polymers. These polymers are urethane-based containing distributed urea, biuret and allophanate groups, which will sorb anionic metal groups. It has been found that water-soluble oxide-based polymers bind with metal cyanide complexes. If the charge density of the polymer is high enough, the metal may be displaced from the accompanying cyanide ions. This is the basis of a novel technique for the separation and recovery of copper and other cyanide species. Oretek holds the patent rights for this process (Jay, 1999).

Recovery of free cyanide and metals that are associated with the cyanide complex from slurries or solution is done without adjustment of pH. By avoiding pH adjustment by acidification, generation of hydrogen cyanide is eliminated. The loose macrocycle structure formation of the polymer matrix enables the copper ion to be

49 sequestered at the centre of the macrocycle cavity, thus coordinating with copper as follows (equation 3.45):

R N NR2 H R R N N N H N - N R + Cu(CN)X Cu + X (CN ) R N N N R R NH2

where R represents either H, -CH3, -aromatic group, -alkyl chain, -functional group, -

CH2-CH2-NH2. (Jay 2001) [3.45]

An ultrafiltration unit can separate the loaded polymer where the metal-polychelator complex (retentate) is separated from the cyanide ions (permeate). The free cyanide can be directly returned to the leach circuit. Copper is recovered by conventional electrowinning directly from the polychelated copper polymer solution releasing the polychelator for recycle. The half-cell reactions for the electrowinning cell are given in equations 3.46-3.48:

(P)Cu2eCu(P) :Anode - 2+ +→+ (P)Cu2eCu(P) [3.46]

1 −+ 2e2HOOH 2 2 2 ++→ 2e2HOOH [3.47]

C u2eCu :Cathode 2 −+ →+ Cu2eCu [3.48]

50 3.4.1.7 Elution of Thiocyanate from Strong based ion exchange resins

Thiocyanate has been used as an eluant to elute precious metals like gold cyanide - (Au(CN)2 (Davison, 1962; Fleming,1986). It was found that thiocyanate’s affinity to strong base ion exchange resins, which lead to great difficulty in displacing the thiocyanate anion in order to regenerate the resin. In an early case study where thiocyanate was not eluted from the resin before exposing the resin to gold cyanide solution found that the adsorption of gold was slower than when the resin was in a chloride form (Fleming, 1986).

Three processes were developed to convert strong base ion exchange resin in the form of thiocyanate (R-SCN, where R represents the resin backbone and functional group) back into the resin original form, either chloride, hydroxide or sulphate (R-Cl, R-OH,

R2-SO4), depending which application the ion exchange resin was used.

3.4.1.7.1 Elution of Thiocyanate Using Mineral Acids

In 1962, Davison filed a patent that proposed a process involving the use of thiocyanate eluant to strip gold from ion-exchange resins. Sulphuric acid was used to strip the thiocyanate from the eluted resin and fresh resin is then used to recover the thiocyanate from the acid liquor. Equation 3.49 illustrates the elution of thiocyanate from ion-exchange resin in the form thiocyanate:

+ − 42 42 2HSORSOHSCN2R ++−→+− 2SCN [3.49] where R represents the resin backbone and functional group.

Fleming (1986) repeated Davison’s work and found that when using sulphuric acid to elute thiocyanate from the resin. After treating the resin that was loaded with thiocyanate with ten times its volume with acid solution only 50% of the thiocyanate was eluted from the resin. Fleming (1986) examined other mineral acids such as nitric and hydrochloric acid and found more satisfactory recoveries from the resin but

51 found that in these eluants thiocyanate decompose fairly rapidly to form elemental sulphur.

3.4.1.7.2 Elution of thiocyanate with potassium chloride and potassium hydroxide

Sapjeta et al. (1987) developed a two-stage elution technique to convert resin in the form of thiocyanate to hydroxide. They proposed the following mechanism for regeneration (equations 3.50 and 3.51):

− +−→+− SCNClRClSCNR − [3.50]

− +−→+− ClOHROHClR − [3.51] where R represents the resin backbone and functional group.

From the first stage of elution, it was reported that 75% of the SCN could be removed from the resin in 80 minutes using 2M KCl at 45oC. In the second stage it was found that 100% of the chloride form was converted to hydroxide form using 2M KOH after 1 hour at 45oC. Upon subsequent regeneration there was no significant change in elution rate of thiocyanate and inferred that this process had no negative effect on the resin. The resin was checked if there would be a change in gold capacity, after one regeneration cycle the capacity was 2.9 meqiv Au/g of wet resin. This value did not change after subsequent regeneration cycles (Sepjeta et al., 1987).

3.4.1.7.3 Elution of thiocyanate with Ferric ions

Fleming (1986) developed an elution process to regenerate resin loaded with thiocyanate using ferric ions. Regeneration takes place under conditions where ferric thiocyanate complex cations (Fe(SCN)2+) are formed and then washed from the resin by the eluant solution. Thiocyanate can be recovered by neutralising the regenerating

52 solution to precipitate iron oxide (Fe(OH)3) to leave thiocyanate ions in solution. Figure 3.8 illustrates the flow diagram for this process.

The regeneration of ion exchange resin loaded with thiocyanate is contacted with a eluant containing ferric ions to form Fe(SCN)2+ as shown in equation 3.52. The resin from this stage can re recycled in the adsorption of precious metals from pregnant solutions.

2+ 2− +− 32 SOR(SO)FeSCN2R 42 +−→ 2Fe(SCN) + 2SO4 [3.52]

Thiocyanate is recovered by precipitating iron as ferric oxide; the reaction can be described by equation 3.53. By solid liquid separation, thiocyanate solution can be separated from iron precipitate and recycled for the elution of gold loaded ion exchange resin.

2+ − − Fe(SCN) 3OH →+ Fe(OH)3(s) + SCN [3.53]

Iron can be reused in the process by digesting the ferric oxide using a mineral acid to form ferric ions. This reaction is described by equation 3.54.

3+ 2− 2Fe(OH) + 423 4 ++→ 2O6H3SO2FeSO3H [3.54]

Fleming (1986) found if iron concentration was greater than 40g/l, the elution efficiency dropped. This was due to the formation of anionic ferric thiocyanate species that reabsorb onto the resin. Another way formation of anionic ferric thiocyanate species could occur if there is an uneven flow pattern throughout the resin bed, this could influence localised iron concentrations. The recovery of thiocyanate was reported to be approximately 90% (Sapjeta et al., 1987).

53 Metal source material

Leach Liquor Leaching Residue

Loaded Liquor

Loading of Resin Barren Liquor

SCN- Loaded-Resin

Regenerated Resin Elution Metal Recovery

Metal

SCN- SCN-Resin

Regeneration

2+ Fe(SCN) Ca(OH)2

Precipitation

- FeSO4 Fe(OH)3 + SCN

Liquid Solid

Fe(OH)3

H2SO4 Acid Digestion

CaSO4

Figure 3.8: Process diagram to regenerate ion exchange resin which was loaded with thiocyanate. (Fleming, 1986)

54 3.4.2 Solvent Extraction

Solvent-extraction processes have been developed for the removal of copper cyanide complexes from aqueous copper/gold cyanide solutions. Many different reagents have been developed, for example XI-78 from the Henkel Corporation based on mixtures of quaternary amine and a weak acid (Davis et al., 1998). An alternate solvent extractant is LIX® 79 developed by COGNIS, a guanidine-based extractant (Virnig et al., 1996; Kordosky et al., 1993). These extractants have similar extraction properties but XI-78 is characterised by a higher loading capacity compared to LIX® 79.

Normal operation of solvent extraction process for the recovery of a metal product and cyanide regeneration from cyanide bearing solutions comprises three main steps. Firstly to recover the metal cyanide species from cyanide solutions solvent extraction is used, using one of the developed extractants. This is followed by stripping the metal cyanide species from the loaded organic phase into an aqueous phase. The third step involves recovering the metal by electrowinning or precipitation while recovering cyanide by electrolysis or acid volatilisation and alkali adsorption, as discussed before.

The extractants are very good at adsorbing metal cyanide complexes but stripping the anions from the extractant can be difficult as the metal complexes must be eluted from the loaded organic with a second anionic species (for example chloride and hydroxide). This can be overcome by the addition of a hydrocarbon-soluble weak acid to the organic phase. The extraction and stripping is then controlled by adjusting the pH of the aqueous phase. The pH range of operation can be optimised by the selection of a weak acid having the appropriate pKa.

Davis et al. (1998) describes that the quaternary amine cation (using XI-78) must extract an anion to maintain charge neutrality in the organic phase. Large anionic metal complexes are therefore extracted via an ion-pair mechanism. Copper cyanide is extracted from aqueous solution in the accordance with equation 3.55 and 3.56:

55 −+ (org)ROQ →+ + (org)QHOH + ROH(org) + − (aq)OH [3.55]

+ 2− + 2− (org)2Q + Cu(CN)3 → 2 Cu(CN))(Q(aq) 3 (org) [3.56] where ROH represents nonylphenol and Q+ is the quaternary amine cation.

Once the loaded organic phase is contacted with an aqueous strip solution having a pH >13, the nonylphenol is converted to the highly hydrocarbon-soluble phenoxide anion. This forms an ion pair with the quaternary amine cation stripping the loaded anionic metal species to the aqueous phase (equation 3.57).

+ 2− − 2 Cu(CN))(Q 3 (org) + 2ROH(org) + (aq)2OH → [3.57] −+ 2− 2 ++ Cu(CN)O2H(org)RO2Q 3 (aq)

Other existing technologies can be employed to separate the metal cation and the cyanide. The most suitable technologies are AVR, MNR and the Du Pont processes.

56 3.5 Processes for the recovery of Cyanide From Thiocyanate

Thiocyanate formation from cyanidation of gold bearing ores is becoming a more common problem during gold processing. Depending on the sulphur content of the ore that is contacted with cyanide, the level of thiocyanate that is formed as a by- product from cyanidation may range from a few mg/l to well over 1000 mg/l. In extreme cases it has been reported that the thiocyanate level in solution is in excess of 5000 mg/l at mine sites which use cyanidation circuits following biological sulphur- oxidation processes (Botz et al., 2001). The high level of thiocyanate is caused by the interaction of cyanide with oxidized and partially oxidized sulphur species.

Thiocyanate is not regulated in mine discharge waters, but it is known that thiocyanate is toxic to many aquatic species. If mine waters are required to comply to a toxicity test, thiocyanate may be of concern (for example US E.P.A.’s Whole Effluent Toxicity W.E.T. test). Therefore thiocyanate detoxification may be needed in plant operations to comply with the test. Cyanide detoxification that many mines use may not destroy thiocyanate and other measures may be needed to reduce the level of thiocyanate.

Another problem caused by the side reaction of cyanide with sulphide materials within the ore is the cost of cyanide that is consumed. It has been estimated that given that the cost of cyanide was given at USD 1.30 per kilo of cyanide. The formation of thiocyanate can account for USD 1.10 as reported by Botz et al. (2001). Table 3.1 shows that if the level of thiocyanate is in excess of 1000 ppm, the additional cost to a mine is in excess of USD $1 million on an annual basis.

Table 3.1: Operating cost associated with thiocyanate formation (Botz et al., 2001) Thiocyanate Concentration (ppm) Annual Sodium Cyanide Cost Due to Thiocyanate Formation (USD) 500 0.8 million 1000 1.5 million 5000 7 million Note: Calculated for a solution flow of 150 m3/hr and a NaCN purchase price of USD 1.30/kg.

57 3.5.1 Chemistry of Cyanide Regeneration from Thiocyanate

The regeneration of cyanide from thiocyanate can be described by the half-cell reaction in equation 3.58. There is potentially a continued oxidation of cyanide to cyanate as described in other oxidation techniques for cyanide destruction to cyanate.

− − −+ 2 42 +++→+ 6e6HSOHCNO4HSCN [3.58]

Many different oxidants have been studied for the recovery of cyanide from thiocyanate. Some examples are listed in Table 3.2.

Under acidic conditions several authors (Botz et al., 2001; Soto et al., 1995) have noted that the oxidation of thiocyanate to cyanide does not continue to form cyanate. This may be explained in terms of the hydrolysis of cyanide ions. Under acidic conditions the oxidant oxidises the cyanide ion but not the acid form, hydrogen cyanide. Thus at low pH, when the concentration of cyanide ion is very low, cyanide destruction basically stops (Jara et al., 1996).

Table 3.2: Half-cell reactions for the oxidation of thiocyanate. Oxidant Half-cell Reaction Oxygen −+ 2 →++ 2O2H4e4H(g)O

− − Chlorine 2 →+ 2Cl2e(g)Cl

−+ Hydrogen Peroxide 22 →++ 2O2H2e2HOH

− 2− Caro’s Acid 52 4 +→+ 2OHSO2eSOH

− 2− Sulphur Dioxide 2 2 →++ SO2e(g)O(g)SO 4

Ozone −+ 3 →++ 2O3H6e6HO

58 3.5.1.1 Ozone

Ozone has been widely used in many water purification processes. To date, the majority of experimental work conducted using chemical oxidants to regenerate cyanide from thiocyanate has been performed using ozone. In comparison to cyanide destruction, ozone is typically used in operations with small volume and low concentration of cyanides. Ozone operations have relatively high operation costs compared to other oxidants such as chlorine, hydrogen peroxide and sulphur dioxide (Botz et al., 2001).

The oxidation reaction for thiocyanate can take place via two different reactions (Soto et al., 1995) as shown by equations 3.59 and 3.60. Further oxidation of cyanide to cyanate can also occur (equation 3.61).

− − − 2− 3 2OH(g)3OSCN 4 +++→++ 22 OH2OSOCN [3.59]

− − 23 ++→++ SOHOCNOHOSCN 422 [3.60]

− − 3 +→+ OCNOOCN 2 [3.61]

Both reactions for the oxidation of thiocyanate to cyanide and the continuation of cyanide to cyanate are fast and are probably mass transfer controlled (Layne et al., 1984). Representative results for ozonation of thiocyanate are shown in Figure 3.9 which show the concentration of cyanide and related compounds as a function of time in an ozonation reactor (Layne et al., 1984).

As seen in Figure 3.9 during the first 20 minutes there is a rapid decrease in thiocyanate concentration that corresponds to a rapid increase in cyanide concentration. Beyond 20 minutes the concentration of cyanide decreases due the continued oxidation of cyanide to cyanate. After 40 minutes the concentration of cyanide approaches zero. The highest cyanide concentration occurs approximately after 20 minutes of reaction time and this corresponds to approximately 75% of

59 cyanide being regenerated from the oxidation of thiocyanate. Similar results have been found by Kemker et al. (1980) in which cyanide regeneration was approximately 80 to 85% conversion of thiocyanate to cyanide at a pH of 10.

1.8 Total N 1.6 CNO- 1.4 SCN- 1.2 CN- 1

0.8 Millimoles/litre

0.6

0.4 NO3

0.2

0 0 102030405060 Time (min)

Figure 3.9: Oxidation of thiocyanate using ozone as the oxidant. (Layne et al., 1984). Operating conditions: pH=11.2, I=0.23, T=25oC, gas flow=1.60 L/min.

Soto et al. (1995) studied the effect of pH on the oxidation of cyanide using ozone as the oxidant. Their findings showed that the rate of cyanide destruction to cyanate decreases as the pH decreases with negligible oxidation below a pH of 4, as seen in Figure 3.9. In comparison with the work by Layne et al. (1984) it appears that with an operational pH of 10, the cyanide formed through the oxidation of thiocyanate is subsequently oxidised to cyanate. In Figure 3.10, the small decrease observed at pH 4 and 6 is largely due to volatilisation of HCN and confirmed by the control test with nitrogen instead of ozone (curve A, Figure 3.10).

60 250

200

150 pH 4 pH 6

100 CN (f) (mg/L)

pH 8 50

pH 12 pH 10 0 01234567 Ozone addition (g)

Figure 3.10: The effect of ozone oxidation on free cyanide (Soto et al., 1995).Note: CN (f) is the total concentration of HCN and CN-.

The effect of dissolved copper on the rate of oxidation of cyanide to cyanate using ozone is significant when the copper level is in excess of approximately 10 – 50 ppm at any pH. This catalytic advantage is commonly used to promote the destruction of cyanide and forms the basis of several destruction patents (Mathre, 1971; Neville, 1980 and Borbely et al., 1984).

Pilot tests using ozone for the recovery of cyanide from thiocyanate were conducted for TVX Hellas for their Olympias project in Greece in 1998 and 1999 (Botz, 2001). The ore from Olympias Project is refractory and near complete oxidation of mineral sulphides was required before cyanidation. This was accomplished by using a bacterial leach to oxidise the sulphide ore. Upon cyanidation, significant levels of thiocyanate were formed, ranging from 2.5 g/l to in excess of 12 g/l. This level of thiocyanate accounted for approximately 90% of the total cyanide consumption in the circuit. Regeneration of cyanide was conducted in a bubble column reactor. The consumption of ozone during the tests was approximately 0.93 gram ozone per gram of oxidised thiocyanate or about 1.12 moles of ozone per mole of oxidised thiocyanate giving an average cyanide recovery of 85%.

61 3.5.1.2 Electrolytic Oxidation of Thiocyanate

An alternative to chemical oxidation of thiocyanate is the use of an induced electric potential to provide the oxidising environment to regenerate free cyanide. This technology has been successfully tested on a laboratory scale but to date there has been no full scale testing. An inherent feature of the electro-regeneration of cyanide from thiocyanate is that the regenerated cyanide is always present under strongly oxidising conditions. The half-cell reaction is represented by:

− − − − 2OHCN 2 ++→+ eOHCNO [3.62]

At the anode surface many different electro-chemical reactions occur involving thiocyanate, cyanide, cyanate and various oxysulphur species. While the identification of these reactions is difficult, some of the overall stoichiometries representing electro-decomposition of thiocyanate are described in equations 3.63– 3.65.

− − 2− −+ 2 4 +++→+ 6e8HSOCNO4HSCN [3.63]

− − 2− −+ 2 4 +++→+ 8e10HSOCNOO5HSCN [3.64]

− 1 + 2− − 2 2 2 2 4 ++++→+ 11eSO12HCONO6HSCN [3.65]

In addition to the above reactions, water and hydroxide ions are oxidised at the anode to form oxygen gas according to equations 3.66 and 3.67. At the cathode there is the reduction of water to form hydrogen gas (equation 3.68). If other electro-reducible species are present in the solution for example copper cyanide complexes they will form elemental copper at the cathode.

−+ 2 2 ++→ 4e4H(g)OO2H [3.66]

− − 2 2 ++→ 4eOH2O4OH [3.67]

− − 2 2 +→+ 2OHH2eO2H [3.68]

Byerley et al. (1984) used a bipolar flow reactor to regenerate cyanide from thiocyanate (Figure 3.11). The electrode was made from graphite or coke, depending on the experimental set-up. They presented data showing the electrolytic oxidation of thiocyanate using synthetic solutions at varying pH. It was found that 80% of the cyanide was regenerated at a pH of 2.0 – 2.5. A decrease in cyanide recovery was observed at neutral to alkaline pH levels. This was attributed to the protonation of the cyanide ion, which is less susceptible to oxidation to cyanate.

Bi Polar Electrode ++++++++++++++++++++++

Basic cathodic boundary layer

- - + 2- - SCN + 4H2O = CN + 8H + SO4 + 6e

------

++++++++++++++++++++++ - - 2H2O + 2e = H2 + 2OH Bulk pH 9.5 – 11.5

Acidic anodic boundary layer ------

Figure 3.11: Schematic of anodic and cathodic boundary layer formation for reactor with parallel plate bipolar electrodes. (Byerley et al., 1984)

Representative electrolytic cyanide regeneration data presented by Byerley et al. (1984) are shown in Figure 3.12. The initial pH of the system (pH of 9) was reduced to a final pH of 2.4 after 200 minutes of cell operation since the oxidation of thiocyanate is acid positive (equation 3.65). Cyanide recovery for this experiment was 82%. Other experiments that were buffered in alkaline conditions gave a

63 maximum cyanide regeneration of 13% after 30 minutes of reaction time. After this period, the regeneration efficiency decreased and approached zero at 120 minutes. The power consumption of the cell during the experiment was approximately 0.03 kW.hr/g of thiocyanate oxidised.

6000

5000

4000

3000

SCN,CN (mg/l) CN 2000

1000

SCN 0 0 50 100 150 200 250 Time (min)

Figure 3.12: Electrolytic oxidation of thiocyanate. A graph showing SCN/CN vs. time in a bipolar flow reactor. (Process conditions: temperature 30oC (av), current 8A (av), initial pH 9 un-buffered (0.09M Na2SO4), final pH 2.4) (Byerley et al., 1984).

3.5.1.3 Hydrogen Peroxide

Hydrogen peroxide has been widely used for cyanide destruction. There have also been studies for the study of the production of cyanide from the oxidation of thiocyanate using hydrogen peroxide (Orban, 1986; Wilson et al., 1960(A) and 1960 (B)). The studies conducted by Wilson et al. (1960 (A) and (B)) showed continued oxidation of cyanide to cyanate that affects the regeneration of cyanide from thiocyanate at low pH. The reaction of thiocyanate with hydrogen peroxide to produce cyanide is shown in equation 3.69.

− 2 +− 22 4 +++→+ 2O2HHSOHCNO3HSCN [3.69]

64 Soto et al. (1995) have stated that hydrogen peroxide could not be used for the regeneration of cyanide from thiocyanate because it is not a sufficiently strong oxidant. They suggested using stronger oxidants for example ozone or chlorine. Chlorine can produce toxic derivatives and its use in the mining as well as other industries is questionable. The preferred oxidant that Soto et al. (1995) used for cyanide regeneration from thiocyanate is ozone.

Wilson et al (1960 (A) and (B)) proposed that the reaction between thiocyanate and hydrogen peroxide occurs by two different mechanisms, depending on the pH of the system. At low pH the oxidation reaction forms hydrogen cyanide and sulphate and the rate of thiocyanate destruction is much faster than in alkaline solutions. In alkaline conditions, sulphur di-cyanide (S(CN)2) is formed as an intermediate before decomposing to cyanate (equation 3.70):

− − + 22 SCNOHS(CN) ++→+ 2HCNO [3.70]

It was noted that dissolved copper resulted in rapid catalytic oxidation of cyanide to cyanate, even at low pH, destroying any regenerated cyanide. Therefore, in using hydrogen peroxide as an oxidant it is critical to maintain a sufficiently low copper concentration to minimise the catalytic oxidation of cyanide to cyanate.

3.5.1.4 Caro’s Acid

At low pH, Caro’s acid is capable of regenerating cyanide from thiocyanate (Smith et al., 1966 (A) and (B)). The chemistry is very similar to that for hydrogen peroxide as the oxidant, under alkaline conditions in that sulphur di-cyanide is formed as an intermediate. Equations 3.71-3.73 describe the oxidation of thiocyanate with Caro’s acid to form free cyanide, di-cyanide and cyanate.

− − 2− + SCN 25 4 ++→++ 4H4SOHCNOH3HSO [3.71]

− − 2− SCN HSO5 SOHCN 4 ++→++ 22 OHS(CN) [3.72]

− − 2− + HCN HSO5 4 ++→+ 2HSOCNO [3.73]

From the experiment conducted by Smith et al. (1966 (B)), no cyanate was produced by the direct oxidation of cyanide by Caro’s acid. Cyanate was formed by the decomposition of the intermediate sulphur di-cyanide.

3.5.1.5 Chlorination

Chlorine has been widely used in the mining industry for cyanide destruction. The process is operated under alkaline conditions (above pH 10) and results in the complete oxidation of thiocyanate and cyanide to cyanate (Botz et al., 2001). It has been theorised that at low pH, chlorine partially oxidises thiocyanate to yield hydrogen cyanide (Flynn et al., 1995) as illustrated in equation 3.74.

− 2− + − 2 2 4 +++→++ 6Cl7HSOHCNO4H(g)3ClSCN [3.74]

In addition, chlorine oxidises cyanide to cyanate according to the following reaction:

− − + − 2 2 ++→++ 2Cl2HCNOOH(g)ClCN [3.75]

Botz et al. (2001) calculated the stoichiometric requirements for thiocyanate and cyanide oxidation with chlorine and the lime needed to neutralise the acid generated in these reactions (shown in Table 3.1)

66 Table 3.3: Stoichiometric reagents demands for thiocyanate oxidation with chlorine (Botz et al., 2001). Reaction Lime Addition Chlorine Addition

(g CaO/g SCN) (g Cl2/g SCN) Cyanide Regeneration 3.86 3.67 Complete Thiocyanate 4.83 4.90 Destruction

3.5.1.6 Oxygen

The advantages in using oxygen for cyanide regeneration from thiocyanate waste solutions are that it is relatively inexpensive and readily available. However, a study conducted by Botz et al. (2001) suggested that the regeneration reaction is prohibitively slow for practical application. At present, efforts are being made to use oxygen in the presence of either zeolites or aqueous catalysts consisting of homogeneous metal salts of copper, manganese and cobalt (Vicente et al., 2003).

Vicente et al. (2003) studied the oxidation of thiocyanate using oxygen at a pressure of 100 atm and a temperature of 200oC. The experiments were conducted under alkaline conditions to stabilise any cyanide formed. High temperatures were used to ensure that the oxidation took place at an adequate rate, while a high pressure was used to maintain a high oxygen concentration, thereby further increasing the oxidation rate. The degree of oxidation was a function of temperature, oxygen partial pressure, residence time and oxidizability of the pollutants.

3.5.1.7 Sulphur Dioxide

Sulphur Dioxide based processes have been used for cyanide destruction but not as of yet for cyanide recovery. When sulphur dioxide was used for cyanide destruction it was found that only 10% of the thiocyanate was oxidised (Botz, 2001). The use of sulphur dioxide and oxygen for thiocyanate oxidation to regenerate cyanide has not been thoroughly investigated. The following reaction occurs under acidic conditions (equation 3.76):

− 2− + SCN 2 2 2 4 ++→+++ 7H4SOHCNO4H(g)3O(g)3SO [3.76]

As described in section 3.2.1 sulphur dioxide is used to oxidise cyanide to cyanate in the presence of a copper catalyst and oxygen.

68 References Bolinski L., Shirley J. “Russian Resin-In-Pulp technology, current status and resent developments.” Proceedings from Randol Gold Forum ’96. Randol International 1996, pp 419-423.

Borberly G., Devuyst E., Ettel V., Mosoiu M., Schitka K. “Cyanide removal from Aqueous Streams.” Canadian Patent No. 1,165,474, 10 April 1984; Covered by US Patent 4,537,686 27 August 1985.

Baden H. “Silver catalysed oxidation complex of metal cyanides.” US patent 3,772,194. Filed November 13th 1973.

Botz M., Dimitriadis D., Polglase T., Phillips W. and Jenny R. “Processes for the regeneration of cyanide from Thiocyanate.” Journal of Minerals and Metallurgical Processing Vol. 18, No. 3, August 2001, pp 126-132.

Byerley J., Enns K. “ Electrochemical regeneration of cyanide from waste thiocyanate for cyanidation.” CIM Bulletin, Vol. 77, No. 861, 1984. pp 87-93.

Carrillo-Pedroza F.R, Nava-Alonso F. and Uribe-Salas A. “Cyanide ocidation by ozone in Cyanide Tailings: Reaction Kinetics.” Minerals Engineering, Vol. 3, No. 5, 2000, pp 541-548.

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75 Chapter 4 Solution Equilibria 4.1 Introduction

Potential-pH equilibria shown as Pourbaix diagrams have been used extensively to describe complex chemistry of aqueous redox systems. These diagrams have been used to help evaluate problems in corrosion science and the stability of different metal-water systems. The concept has been expanded to different areas in hydrometallurgy and geochemistry. These diagrams describe regions of stability of various species under different Eh and pH conditions, which are desired or undesired. The species will be formed as predicted by the diagram unless some features of the process kinetics interfere. Kinetic aspects are very important considerations, but it is clear that they can be studied efficiently only after equilibrium aspects treated by Pourbaix diagrams are fully understood (Guy G. et al., 1962).

Pourbaix diagrams originated from the theoretical prediction of oxidation-reduction catalytic processes and the conditions under which these reactions take place in aqueous solutions. The theory originated on work done by Michaelis which was based from experimental work conducted by Travers and Thiesse for chemical systems involving gaseous phase reactions. Pourbaix saw a correlation between aqueous systems and the theory defined by Michaelis was applied to electrochemical systems involving aqueous solution chemistry (Pourbaix, 1974).

The theory has been applied to problems that involve the electrochemistry of metals and non-metals and its application spreads to different branches of electrochemistry and related fields. Pourbaix describes the application of this theory in the study on the corrosion of iron. For example, iron can react to form iron oxides and or iron ions within the aqueous system depending on the aqueous environment where iron is kept, as seen in Figure 4.1. This figure is divided into several different areas that define the most predominant species at a particular pH and Eh and relates how the predominant

76 species change as the pH and Eh is changed. In the case iron-water system as seen in 2+ Figure 4.1, the most predominant species at a pH of 4 and Eh of 0V is Fe . If it is of 2+ interest to produce a product of Fe(OH)3 from the solution of Fe at a pH of 4 and Eh of 0V, it can be produced by either increasing the pH or Eh of the system above 6 or 2+ 0.25V. If the Eh of the Fe solution is reduced to –0.25V at the same pH of 4 and then the pH of the system is increased the predominant species changes from Fe2+ to form Fe(OH)2 before forming the product Fe(OH)3. It maybe disadvantage in forming Fe(OH)2 as an intermediate through the formation of Fe(OH)3 product because a higher pH is required which increases the amount of base needed.

Secondly, contamination of the Fe(OH)3 product with Fe(OH)2 which will not happen if the formation of Fe(OH)3 could be produced without forming Fe(OH)2 as an intermediate product. By considering each possible reaction that can occur within the aqueous system, Pourbaix has shown that the limiting conditions for the occurrence of all known reactions can be represented conveniently in a single diagram.

Fe3+ Fe(OH) 1 3

2+ 0 Fe

Eh (volts) Fe(OH)2 H2 -1 Fe

-2 0 2 4 6 8 10 12 14 pH

Figure 4.1: Potential-pH equilibrium diagram for the system iron-water, at 25oC. Equilibrium data for each species are from NBS tables (Wagman et al., 1982). (Fe, Fe(OH)2 and Fe(OH)3 considered as the solid substances)

77 This chapter examines the theory behind the production of Pourbaix and distribution diagrams and recent developments to automate the production of these diagrams for different chemical systems. This theory is applied to produce Pourbaix diagrams for thiocyanate-water and copper-thiocyanate-water systems using the modelling package STABCAL (version W32-Stabcal). Distribution diagrams were constructed to show how the ratio of copper to thiocyanate changes the equilibrium concentration of the different copper thiocyanate species under different pH and Eh conditions.

4.2 Chemical Reactions

Aqueous reactions can be divided into two categories, namely chemical and electrochemical reactions as defined by Pourbaix (1974). A chemical reaction is a “reaction which only neutral molecules and positively or negatively charged ions take part, with the exclusion of electrons”. The electrochemical (or electrode) reaction shall be defined as “being a reaction involving, besides molecules and ions, negative electrons (e-) arising from a metal or other substance by metallic conduction. Such reactions are oxidations if they proceed in the direction corresponding to the liberation of electrons; or reductions if they proceed in the direction corresponding to the consumption of electrons”.

These two processes can be summarised in terms of half-cell reactions, which are shown in generalised form by equation 4.1, written in the form of a reduction:

+ - aA + cH +ne = bB + mH2O [4.1]

By describing a chemical system using the generalised reaction formula we obtain equilibrium relations that relate the reaction as a function of pH and the electrode potential. pH measures the effect of the H+ ions and the electrode potential measures the effect of the charges (e-).

78 4.3 Theoretical Background on the formation of Eh-pH

Any chemical reaction can be described with the general equation (equation 4.1). By bringing all the terms to one side of the equation the reaction can be written in a more concise form (equation 4.2):

− ∑ γγ =− 0nemv [4.2] where e- is an electron, n is the number of electrons transferred in the reaction, γ an identification or index of the species and vγ is the stoichiometric coeffiecient of species mγ. vγ are define as positive for product species and negative for reactants when the equation is written in reduction form. As an example for equation 4.1:

va = -a, vH+ = -c, vb = +b and vH2O = +m where a,b,c and m are positive integers.

It is useful to compare the electrode process to some reversible reference electrode in which the state of equilibrium of a given reversible electrochemical reaction is obtained. In most cases the standard hydrogen electrode is used for which the reaction is defined by equation 4.3:

−+ ↔+ H2e2H 2 [4.3]

The standard hydrogen electrode behaves ideally, that is, the activity of H+ in solution + and the fugacity of the gas are unitary (aH = pH2 = 1). By applying equation 4.2, the reference electrode and any reference electrode can be described by equation 4.4:

− ∑ refref =− 0nemv [4.4]

79 The overall cell reaction when comparing the reaction studied referenced to a reference electrode would be given by combining equations 4.2 and 4.4 to give (equation 4.5)

γγ −∑∑ refref = 0mvmv [4.5]

In an equilibrium situation the equilibrium potential of the cell can be described from the relationship between Gibbs free energy of the reaction and the potential of the cell:

ΔG ref =−+ 0)EnF(E [4.6] where F is Faraday’s constant (F= 96 485 coulomb/mole) and ΔG is the chemical free energy for the overall reactions and can be calculated from the chemical potentials of each species. Therefore ΔG can be defined as:

G ∑∑v μ γγ −=Δ v μ refref [4.7]

where μγ, μref are the chemical potentials of species mγ and mref respectively.

E – Eref is the electrode potential compared to a reference electrode and is defined as chemical potential (Eh). The most commonly used reference electrode is the standard hydrogen electrode (SHE) and by definition ESHE = 0V.

For the reaction given in equation 4.1, the relationship between Eh and pH can be determined:

ΔG ∑∑v μ γγ −= v μ refref [4.8] n (-aμ cμ + Bμ μ ()m nμ + +−−++−= μ ) A H B 2OH H 2 H2

80 The chemical potential (μγ) for each species concerned can be related to its standard o potential (μγ ) by the activity of the species aγ and temperature:

o μμ γγ += RTlna γ [4.9]

For the SHE reaction, μΗ+ and μH2 are the same as its standard chemical potential. This is because each of these species are at standard state and the activity of H+ and the fugacity of H2 is defined as unitary.

Therefore the Gibbs free energy of the reaction of interest can be written as:

b m o o o o o .aa n B 2OH ΔG a μ n)(c + μμ bμ mμ +++−−−−= RTln [4.10] A H 2 H2 B 2OH a C A .aa H+

Equation 4.10 can be simplified to:

b m o .aa B 2OH ΔG +Δ= RTlnG a C [4.11] .aa + A H where ΔGo is the standard free energy change for the overall reaction. Substitution of equation 4.6 in equation 4.11 yields:

b m ΔG RT B .aa OH E −−= ln 2 [4.12] h nF nF a C A .aa H+

The general relationship for the electrode potential at equilibrium can be described as:

ΔG RT E h −−= ∑ v aln γγ [4.13] nF nF γ

81 The ∑v aln γγ term can be expressed a different way to give the following term: γ

γ vγ vγ alnaln aln v aln γγ ∑∑ γ ∏== alnaln γ [4.14] γ γ

Therefore equation 4.12 can be rewritten to give:

ΔG o RT a b RT RT E −−= ln B − 2.303 c.pH − am.ln [4.15] h a 2OH nF nF a A nF nF

From equation 4.15 it can be seen that Eh varies linearly with respect to pH. There are special cases where there are no electrons transferred or only hydrolysis occurs between the species of interest A and B. In the case that there is only a redox without any hydrolysis reactions, equation 4.15 that defines the relationship between the two species lies horizontal and parallel to the pH ordinate. The other case is when there are only hydrolysis reactions occurring between species A and B, and then the defining lines are vertical and parallel to the Eh ordinate. By comparing a series of related chemical reactions at equilibrium, stable areas will be traced out, and represent stable regions for a particular or dominant species.

The stability of water is of important consideration when constructing Pourbaix diagrams, since most systems are aqueous. Water can be oxidised or reduced if the Eh of the system is too high or low at the respective pH. The stability of water limits the range of Eh for which the system can be operated. If the Eh is too high or low, water would either breakdown to produce oxygen (an oxidative environment, high Eh conditions) or hydrogen (a reducing environment, low Eh conditions).

Decomposition of water into its elements of hydrogen and oxygen follow the following reaction:

2 +↔ O2HO2H 22 [4.16]

82 Equation 4.16 can be described by two half-cell reactions, reduction to form hydrogen and oxidation to form oxygen. For the reduction half-cell reaction (equation 4.17):

−+ →+ H24e4H 2 [4.17]

For the oxidation half-cell reaction (equation 4.18):

− −+ 2 ++→ 4e2HO2OH [4.18]

The overall reaction when combining equation 4.17 and equation 4.18 will give the following equation:

+ − 2 +→+ 2HO2OH2H 2 [4.19]

The dissociation of water can be represented by

+ − 2 +↔ OHHOH [4.20] and when combined with equation 4.19 will give the original equation 4.16, which describes the decomposition of water.

For the oxidation half cell which is represented by equation 4.21:

−+ 2 2 ++→ 4e4HOO2H [4.21] the corresponding potential at 25oC is given by

228.1E −= 0.0591pH − 0.0147log p [4.22] h O2

Similarly for the reduction half-cell the corresponding potential is given by:

0.000E −= 0.0591pH − 0.0295log p [4.23] h H2

For a pressure of one atmosphere for hydrogen and oxygen the logarithmic term within the equation is equal to zero leaving the following relationships.

h 228.1E −= 0.0591pH [4.24]

h 0.000E −= 0.0591pH [4.25]

Between these two lines on a Pourbaix diagram the pressure of hydrogen and oxygen is less than 1 atmosphere. Therefore at one atmosphere and a temperature of 25oC the area between these two lines is where water is thermodynamically stable.

84 4.4 Distribution Diagrams

Distribution diagrams represent the speciation of various species in a multicomponent aqueous system. These diagrams have an added benefit over Pourbaix diagrams as they allow an independent variable to be varied in order to see its effect on the distribution of different species present in the system. STABCAL a program developed Huang et al. (1989 (A) and (B)), uses the same principles as those used in the construction of Pourbaix diagrams. The main difference between the Pourbaix and distribution diagrams is that distribution diagrams solve calculations for a mass balances for each component.

The mass balance for each component is defined as the mass of all the related species to be equal to a predetermined total mass. This avoids the system becoming supersaturated and solid species should be included for each component. The other set of equations needed to determine the distribution of the species are defined by chemical equilibria. The program uses thermodynamic data in the form of Gibbs free energies of formation. From this chemical reactions, equilibrium equations and equilibrium constants can be determined.

Two methods are used to solve the mass distribution calculations based on either free energy minimisation and mass action (equilibrium) calculations. STABCAL uses the mass action calculations because it does not involve sophisticated numerical methods such as non-linear programming.

4.4.1 Theory behind the distribution diagram

Consider an aqueous system that consists of I chemical species and has J components. There will be J mass balance equations and I - J chemical reactions, which will be labelled as T reactions.

85 The mass balance equations are defined by equation 4.26:

I ∑ = MmA jiji jfor = to1 J [4.26] =1i

where Mj is the total concentration of jth component, mi is the concentration of the ith species, and Aji is the number of atoms from jth component for ith species.

The equilibrium equations for the distribution diagram are similar to Pourbaix o diagrams by defining from the free energy of reaction ΔG r according to the Nernst equation (equation 4.27). For a system with known pH and Eh, the equilibrium equation depends on the activities of the chemical species involved. If gaseous species are present, this can be accounted according to their reaction coefficients and partial pressures.

− ΔG(nFEh o ) ln(K) = r [4.27] RT

Equation 4.28 describes each reaction within the system (T in total):

I I ∑∑iti = iiti = t )ln(K)mln(rB)ln(aB for t = to1 T [4.28] ==1i 1i

This can be rearranged to give the following equation 4.29.

I I ' ∑∑iti t −= iti = t )ln(K)ln(rB)ln(K)ln(mB [4.29] ==1i 1i

’ where ai, ri are the activity and activity coefficient of the ith species respectively, K t is the tth equilibrium constant that includes also the pH, Eh and activity coefficients and gaseous pressure, and Bti stoichiometric coefficient from ith species to tth equilibrium equation.

86 Stabcal (version W32-Stabcal) uses the Newton-Raphson method to solve the equations for the defined chemical system. Calculations can be solved for one species or all at once, Huang (1989(B)) uses the latter method because the system converges faster.

Usually concentration (mi) or mole fraction (xi) are variables that are used as the iteration variables. Another option to use is to take the logarithmic concentration as the iteration variable. This has a benefit as it expands the concentration to a logarithmic scale and avoids some computational problems, for example when the concentration of mi is small. For example, in a study of the stability of different 4- uranium species (Huang, 1989(B)), the calculated concentration of UO2(CO3)3 was found to be 2.3984 x 10-47 M. With results like this computations can assume that the concentration for this species to be zero. A value of zero for a concentration or equilibrium constant can disturb the equilibrium equations and as a result can cause errors within a simulation. By the introducing a new variable yi = ln(mi) this problem can be avoided. In turn with the introduction of yi, the mass balance equations are no longer linear and the equilibrium equations to become linear equations. The mass balance equations then become (equation 4.30):

I j = ∑ jiji =− 0M)exp(yAf [4.30] =1i

From this the derivative of this equation with respect to y becomes (equation 4.31):

df j = iji )exp(yA = to1jfor J [4.31] dyi

Similarly the equilibrium equations by the introduction of y become (equation 4.32):

I ' t ∑ iti t =−= 0lnKyBg [4.32] =1i

and the derivative of this equation with respect to yi is (equation 4.33):

dg t = Bti t = to1for T [4.33] dyi

From all equations a series of matrices can be formed and solving the equations by repeating a number of iterations till a specified tolerance is met, giving a resultant solution (a vector of y). Haung has described the operations, which are given below:

(0) (0) From an initial value of yi , a new value of yi is calculated by the following matrix operations (equations 4.34 to 3.37).

(0) −= − U)(WV (0)1(0) [4.34] where V is a I vector:

(1) (0) ii −= yyv i [4.35]

W is an I x I matrix:

(0) ⎛ df ⎞ ⎜ j ⎟ w ni = ⎜ ⎟ when n ≤ J. ⎝ dyi ⎠ [4.36] (0) ⎛ dg ⎞ ⎜ t ⎟ w ni = ⎜ ⎟ when n > J. ⎝ dyi ⎠

U is an I vector:

(0) = jn whefu n n ≤ J. [4.37] (0) = tn whegu n n > J.

88 All calculations up to this point have only considered soluble species. When a solid is introduced in the calculations, the calculations can be simple if the solid species are exactly known under the systems operation conditions (Eh and pH). This sort of information is not always available for certain conditions or if the conditions change. The normal procedure is to calculate the distribution for the aqueous species first. Then from the results check to see if any species are super-saturated. If saturation occurs, search for the solid species that will precipitate. Morel and Morgan (1979) used a trial and error method to determine the solid species.

When STABCAL introduces solid species to an aqueous system, two assumptions are made. It assumes that all solids behave as an ideal solid solution and the solid species have unit activity. The solid species equations are basically the same equations that are used for aqueous species the only difference is that for solid species mole fraction are used instead of molality to describe the species.

With all this in mind, assume a system that consists of I chemicals, of this there are Q dissolved species, giving S solids (S = I-Q). The mass balance equations would change to include the solid species and the equilibrium equations for the aqueous species should remain unchanged. The new equations to describe the system are described in equations 4.38 and 4.39 for the mass balance and equilibrium equations, respectively.

I ' ∑ iqi = q )ln(K)ln(mB qfor = to1 Q [4.38] =1i

I ⎛ I ⎞ ' ⎜ mlnB)ln(K)ln(mB ⎟ to1Qsfor I [4.39] ∑∑isi s += ss ⎜ s ⎟ += =+1i ⎝ = 1Qs ⎠

The last term in equation 4.39 has been included to convert mole concentration to mole fraction. The stoichiometric coefficient for the solid, Bss, should be one because if all the master species are ions, there is only one solid species per solid equilibrium equation.

89 Letting the variable ys = ln(ms) and applying the assumption that there is unit activity for solid species, the variable ys will become zero for all solid species. The equilibrium equations will change to give the following equation (equation 4.40):

Q ' ∑ isi = s )ln(K)ln(mB sfor += to1Q I [4.40] =1i

From the solid solution calculations, Huang (1989(B)) has reported that some of the solid species may be rejected because of low mole fractions or a disagreement with the phase rule. Once a solid species is eliminated as one of the possible stable species, iteration can start with the remaining species.

It becomes nearly impossible to reject a solid species because of negative mass since the iteration variable is the logarithm of concentration. The program rejects the solid species during the iteration if it causes the matrix W to become a singular determinant or the mass is extremely small compared to the rest of the other species.

90 4.5 Developments in computer aided programs to produce Pourbaix diagrams

Since the age of the computer, programs have been developed to aid researchers in producing Pourbaix diagrams. The main differences between the different programs available for producing Pourbaix diagrams differ by the approach of removing redundant boundary lines between stable compounds. The main methods of doing this are the point-by-point method, the line elimination method and the convex polygon method.

4.5.1 Point-by-point method

The “point-by-point” method has been named because every point on the plane of the

Eh-pH diagram is considered separately. Duby (1963) at Colombia University developed one of the first programs. One of the limitations of this method is the resolution of the diagram can be limited to the number of calculations required to be performed.

This is a foolproof method in determining the diagram since it calculates the species with the lowest free energy of formation, which is the most stable species at each point on the diagram. This is the basis of all equilibrium calculations in thermodynamics. The problem that is faced by this method of calculation is that it requires a large number of calculations for each point on the diagram and the actual number of calculations is very dependent on the resolution of the diagram.

The output from this method is an array of symbols representing the stable species at each point on the Eh-pH diagram. The boundaries separating the different stable regions can be dependent on the number of points calculated. If the resolution in the diagram is not well defined, the boundaries separating the different stable regions can be confusing; this can be overcome by increasing the resolution of the diagram.

91 4.5.2 Line elimination method

This method focuses establishing equilibrium lines to determine relevant boundaries between stable compounds. Brook (1971) proposed that for any chemical system can be defined by important system characteristics:

1. The valency of stable metallic species increases as the potential increases.

2. With an increase in pH the stable ionic type changes from cationic to non- ionic then anionic.

3. Most intersections of reaction lines are triple points.

Applying important characteristics of a system defined by Brook (1971) to group the common species together and to order all lines generated by decreasing slope (using the theory developed by Pourbaix). The line with the greatest slope for a given species is defined as a phase boundary and continues to be so until it intersects with the line with the next highest slope or plot boundary. Within the plot boundary, the new line intersection becomes the phase boundary until it to reaches an intersection. When the boundaries have been established for the first species (the species with the lowest valency and least anionic), succeeding species are treated similarly except if the line has already been considered previously. The end result is the establishment of the phase boundary lines.

This method overcomes some of the disadvantages that occur when using the point- by-point method by establishing line boundaries. Allowances have to be made to allow inconsistencies in the data to make the program to accept some triple points if the three intersections involved are less than some distance apart.

There are some deficiencies in the proposed algorithm. In some cases the algorithm does not necessarily define a particular species uniquely, for example in the Cu-H2O - - system; the following species can be defined the same, HCuO2 and CuO2 , because they have the same oxidation number of 2+ and are both anionic (for which the ion type parameter is –1). This can be fixed by using an alternative scheme for

92 identifying each species (Linkson, 1979). The main limitation of the method developed by Brook (1971) is that the algorithm will not determine the correct stability diagram if the species involved are not ultimately stable. This limits the generality of the algorithm because before a stable system can be identified, the program would have to be run and rerun until a consistent set of boundaries are found.

4.5.3 Convex polygon method

The convex polygon method as with the point-by-point method developed by Duby (1977) is robust. Compared to the line elimination method the convex polygon method has more calculations to complete increasing the computation time but has the main advantage of eliminating the species that are not stable.

When a predominant area is found it satisfies several inequalities simultaneously and forms a convex polygon. The line elimination method relies on the fact that the boundaries form predominant areas; but the convex polygon method develops procedures to form predominant areas by the formation of polygons.

The first work in defining convex polygons was studied by Froning et al. (1976). They applied convex polygon theory to the formation of Pourbaix diagrams. Their technique involved the use of conventional mathematical theory to determine the convex polygon for each species from the equilibrium lines that are involved. The equilibrium lines are considered in groups and by using a sorting algorithm establishes possible boundaries for a species to form a convex polygon. If the polygon is a stable area for a particular species, the convex polygon satisfies all the equilibrium line equations for each individual species and consequently the predominance area for that species has been determined.

STABCAL (version W32-Stabcal) was developed by Haung and used to generate the stability diagrams for the thiocyanate-water and copper-thiocyanate-water systems. This program uses a linear algebra method similar to the convex polygon method developed by Verink. This model has two basic assumptions based on work done by Garrels and Christ (1965): the total mass of each component is assigned to the

93 predominant species and the total concentration of the ligand component is assigned to the ligand species that are not complexed to the main component. The algorithm defines constraint equations that apply to each species. Then searches for boundary lines and corner points of a predominant area enclosed by 2 constraints imposed to each species, which are boundary lines of each ligand. Equilibrium constraints are determined from equilibrium equations and chemical reactions.

94 4.6 Pourbaix Diagrams for the Thiocyanate-Water and Copper-Thiocyanate-Water Systems

In this study stability diagrams were simulated for the thiocyanate-water system using the Stabcal (Stability Calculation for Aqueous System) version 32-STABCAL SN 8- 19-2001. The species that were considered to determine the stability products for the thiocyanate-water system were products related to thiocyanate and cyanide. These species are summarised in Table 4.1 including the Gibbs free energies of formation that were used within the simulation. A comparison between the American National Bureau of Standards (NBS) of chemical thermodynamic properties (1982) and data that Duby (1977) reported during studies of different copper systems showed that the values for each species used were very similar in each reference. The values used within the simulation were from the NBS Tables.

Other species of different sulphur compounds were included within the simulation to balance the sulphur during the oxidation and reduction of thiocyanate. Many of the sulphur compounds are thermodynamically unstable and tend to decompose at standard room temperature and pressure. Pourbaix (1974) reported that the thermodynamic instability of many sulphur compounds such as thiosulphates, diathionites, sulphites and polythionates shows a false equilibrium between the different sulphur species in aqueous solution. It is added that persulphates are - unstable in water. It was suggested that if the equilibria were obtained, only H2S, HS 2- - 2- , S , S, HSO4 and SO4 would be found in solution. For this reason these 6 sulphur compounds were the chosen species for the simulation.

4.6.1 Thiocyanate-Water System

Figure 4.2 and Figure 4.3 are potential-pH equilibrium diagrams for thiocyanate-water system at 25oC. For Figure 4.3 the cyanate and hydrogen cyanate were removed as stable species to show that if the oxidation of thiocyanate to produce cyanide, either free cyanide or hydrogen cyanide would become the predominant species.

95 Table 4.1: Gibbs Free Energies for Thiocyanate and related species at 25oC. Species Duby (1977) NBS Tables (1982) o o ΔG f (kcal/mol) ΔG f (kcal/mol)

HCN (l) 27.86 29.869 HCN (aq) 28.6 28.609 CN- (aq) 41.2 41.205 CNO- (aq) -23.3 -23.279 HCNO (aq) -28.0 -27.988

SCN- (aq) 22.15 22.158 HSCN (aq) 23.31 23.317

(SCN)2 (l) (79.9) 79.9

H2S (aq) - -6.652 HS- (aq) - 2.887 S2- (aq) - 20.507 S (s) - 0 - HSO4 (aq) - -180.667 2- SO4 (aq) - -177.947

CuSCN (s) 16.7 16.706 3- Cu(SCN)4 (aq) 87.0 87.022 CuSCN+ (aq) 34.7 34.63

Cu(SCN)2 (aq) 54.9 54.947

96 The predominant species formed in solution in the thiocyanate water system is cyanate and its protonated form hydrogen cyanate seen in Figure 4.2. With an Eh less than 1 V over the entire pH range, either cyanide or hydrogen cyanide is the predominant stable species. The division between cyanide and hydrogen cyanide o occurs at the value of pKHCN at 25 C, the same applies for cyanate and hydrogen cyanate.

The concentration of thiocyanate was varied from 0.001 to 1 M thiocyanate to see how the areas of predominant species change with concentration. As seen in Figure 4.2, as the concentration of thiocyanate is increased the stability region for thiocyanate increased becoming the predominant specie more stable over a wider pH range. When cyanate and hydrogen cyanate were removed from the simulation (Figure 4.3) the effect of concentration did not change the predominant region for thiocyanate significantly, with thiocyanate being stable close to the water stability hydrogen line from pH of 0 to 14.

HOCN(a) 1 OCN-

0 - SCN -3

Eh (volts) -2 -1 0 H2 -1 HCN(a) CN-

-2 0 2 4 6 8 10 12 14 pH

Figure 4.2: Potential-pH equilibrium diagram for thiocyanate-water system at 25oC.

1 CN- HCN(a)

0 - -3 SCN Eh (volts) -2 -1 0 H2 -1 HCN(a) CN-

-2 0 2 4 6 8 10 12 14 pH

Figure 4.3: Potential-pH equilibrium diagram for thiocyanate-water system at 25oC. Excluding cyanate and hydrogen cyanate from the simulation.

From Figure 4.2 and Figure 4.3 it can be demonstrated that under oxidising conditions it is thermodynamically possible for the production of free cyanide from thiocyanate if the continued oxidation of free cyanide can be stopped. The effect of increasing the concentration of thiocyanate from 0.001M to 1M does not introduce new predominant species within the thiocyanate-water system.

4.6.2 Copper-Thiocyanate-Water System

In the past, copper (II) has been added as a catalyst in many cyanide destruction processes. For the recovery of cyanide via the oxidation of thiocyanate it was postulated that copper could give the same advantage. Therefore it is of interest to understand the formation of different copper thiocyanate species under different operating conditions.

98 Copper can form stable complexes with thiocyanate in an aqueous system. The four stable complexes that can occur are solid copper (I) thiocyanate (CuSCN) and three 3- other soluble species complexed with either copper (I) (Cu(SCN)4 and Cu(SCN)2) or copper (II) (Cu(SCN)+). These species were included to simulate the copper- thiocyanate-water system.

Figure 4.4 shows the stability diagram for the copper-thiocyanate-water system at 25oC at a copper and thiocyanate concentration of 0.001M. From this diagram the only stable species containing thiocyanate is the solid copper (I) thiocyanate with the other species being copper and different forms of copper oxide. Increasing the thiocyanate concentration to 0.01M with the same copper concentration of 0.001M a new predominant region is established for copper (II) thiocyanate (CuSCN+) at a low pH and in an oxidising environment, as seen in Figure 4.5.

The thiocyanate concentration was increased further to 0.1 and 1M as shown in Figure 4.6 and Figure 4.7, respectively. At a higher concentrations of thiocyanate the predominant stable copper species can exist with a higher proportion of thiocyanate, 3- in particular Cu(SCN)2 and Cu(SCN)4 . At low pH the species that are formed are mainly copper (I) species, either CuSCN(s), and Cu(SCN)2. At higher Eh conditions, i.e. in an oxidising environment copper (I) thiocyanates are oxidised to copper (II) + 3- thiocyanates to form CuSCN . The formation of Cu(SCN)4 occurs when the thiocyanate concentration was 1M, and is stable between the regions of CuSCN and CuO, as seen in Figure 4.7.

As the copper concentration is increased from 0.001M to 1M, at a constant thiocyanate concentration of 0.001M (Figure 4.8), stability regions of Cu2+, CuSCN and CuO change. As the copper concentration is increased, CuSCN region becomes more predominant at a higher Eh, and the CuO region extends more into the acidic environment.

99 2

Cu2+ CuO 1

O2 CuSCN 0 Cu2O Eh (volts)

H2 -1 Cu

-2 0 2 4 6 8 10 12 14 pH

Figure 4.4: Potential-pH equilibrium diagram for the copper-thiocyanate-water system at 25oC, at a copper and thiocyanate concentration of 0.001M.

Cu2+ CuO 1 CuSCN+

O2 CuSCN 0 Cu2O Eh (volts)

H2 -1 Cu

-2 0 2 4 6 8 10 12 14 pH

Figure 4.5: Potential-pH equilibrium diagram for the copper-thiocyanate-water system at 25oC, at a copper concentration of 0.001M and thiocyanate concentration of 0.01M.

100 2

Cu2+ CuO 1 CuSCN+ Cu(SCN)2(a) O2

CuSCN 0 Cu2O Eh (volts)

H2 -1 Cu

-2 0 2 4 6 8 10 12 14 pH

Figure 4.6: Potential-pH equilibrium diagram for the copper-thiocyanate-water system at 25oC, at a copper concentration of 0.001M and thiocyanate concentration of 0.1M.

Cu2+

1 CuSCN+ CuO

Cu(SCN)2(a) Cu(SCN) 3- 4 O2 CuSCN 0 Cu2O Eh (volts)

H2 -1 Cu

-2 0 2 4 6 8 10 12 14 pH

Figure 4.7: Potential-pH equilibrium diagram for the copper-thiocyanate-water system at 25oC, at a copper concentration of 0.001M and thiocyanate concentration of 1M.

101 As the concentration of thiocyanate and copper are changed at the same level from 0.001M to 1M (results can be seen in Figure 4.9 and Figure 4.10), similar trends developed as observed in Figure 4.4 through to Figure 4.8. As the thiocyanate and copper concentration is increased from 0.001 to 0.01M CuSCN+ becomes a predominant species within the diagram above CuSCN under acidic conditions (seen in Figure 4.9). The effect of increasing the copper and thiocyanate concentrations above 0.01M introduces another stable region for Cu(SCN)2 between CuSCN and + 3- CuSCN , shown in Figure 4.10. Cu(SCN)4 is not a stable species within the Eh-pH diagram. In this case the ratio of thiocyanate to copper does not change as with the scenario shown in Figure 4.6. An important factor emerges in that it is important to 3- maintain a high thiocyanate to copper ratio for the stability of Cu(SCN)4 . However, in practice, the concentration of thiocyanate in mine waters does not reach this level 3- to form Cu(SCN)4 in tailing waters.

2 0 -1 Cu2+ -2 CuO 1 -3 0 -1 -2-3 O2 CuSCN 0 Cu2O Eh (volts)

H2 -1 Cu

-2 0 2 4 6 8 10 12 14 pH

Figure 4.8: Potential-pH equilibrium diagram for the copper-thiocyanate-water system at 25oC, at a thiocyanate concentration of 0.001M and changing copper concentration from 0.001 to 1M.

102 2

Cu2+ -2 -3 1 CuO CuSCN+

O2 -3 CuSCN -2 0 -3 -2 Cu2O Eh (volts)

H2 Cu -1

-2 0 2 4 6 8 10 12 14 pH

Figure 4.9: Potential-pH equilibrium diagram for the copper-thiocyanate-water system at 25oC, the thiocyanate and copper concentrations are at the same of either 0.001 or 0.01M.

2+ 0 Cu -1 CuO 1 CuSCN+

Cu(SCN)2(a) O2 CuSCN -1 0 0 -1 Cu2O 0 Eh (volts)

H2 -1 Cu

-2 0 2 4 6 8 10 12 14 pH

Figure 4.10: Potential-pH equilibrium diagram for the copper-thiocyanate-water system at 25oC, the thiocyanate and copper concentrations are at the same of either 0.1 or 1M.

103 4.7 Distribution Diagrams for the Copper-Thiocyanate- Water System

The distribution diagrams can show which species are stable under different copper- thiocyanate systems that are formed during the oxidation of thiocyanate in the presence of hydrogen peroxide. From oxidation experiments using hydrogen peroxide to oxidise thiocyanate (as seen in Chapter 6), copper was added to the system to see if copper had any catalytic effect upon the oxidation of thiocyanate. In these experiments the concentration of copper was constant since it was not consumed in the reaction, while the thiocyanate concentration changed with time as it was consumed, hence changing the ratio of thiocyanate to copper. To emulate the change in thiocyanate to copper ratio, a base case of 0.001M thiocyanate and copper was simulated. Then 0.1 moles of either copper or thiocyanate were added step wise to the base case and the distribution of the different copper-thiocyanate species were calculated.

Species that are of interest and under investigation were free thiocyanate (SCN-) and the protonated form (HSCN) and other thiocyanate species complexed with copper + 3- (CuSCN(s), Cu(SCN)2, CuSCN and Cu(SCN)4 ). To keep the simulation simple and to show the stability between the different copper-thiocyanate species the only other species that were included in the simulation were copper (I) and copper (II) ions.

The solution chemistry of the system was changed by manipulating the pH and Eh of the operating solution. The pH and Eh were changed to three different conditions, i.e. the pH of the system was studied at a pH of 0, 7 and 12 and the Eh of the solution changed from –0.25, 0 and 1 V. The conditions for each simulation are summarised in Table 4.2.

104 Table 4.2: Operating conditions for the simulations to determine the distribution diagrams for the copper-thiocyanate-water system.

Simulation pH Eh (V)

1 7 -0.25 2 7 0.00 3 7 1.00 4 0 -0.25 5 0 0.00 6 0 1.00 7 12 -0.25 8 12 0.00 9 12 1.00

At equilibrium, the majority of the thiocyanate exists as free thiocyanate (SCN-) or species complexed with copper like the solid copper (I) thiocyanate (CuSCN) when the initial concentration of thiocyanate and copper is set at 0.001M with a pH of 7 and an Eh of 0V, seen in Figure 4.11. To this system either copper or thiocyanate were added to change the ratio of thiocyanate to copper under different system conditions

(i.e. pH and Eh) to see how the different ratios change the distribution of species within the system. For example in the case of Figure 4.11 additional copper was added.

Figure 4.12 to Figure 4.17 show the individual components studied for the copper- thiocyanate-water system when additional free thiocyanate is added to the system changing the ratio of thiocyanate to copper ratio. Thiocyanate exists predominantly as either free or hydrogen thiocyanate and when complexed with copper it forms significant amounts of copper (I) and copper (II) complexes. Due to large amounts of 3- excess free thiocyanate, copper will form copper (I) tetra-thiocyanate (Cu(SCN)4 ), but the concentration levels are very low (in the order of magnitude of 10-7 3- moles/litre). Therefore the formation of Cu(SCN)4 is insignificant in the simulations studied.

105

-3 CuSCN CuSCN+

SCN- -4

-5

-6 Cu(SCN)2(a) log(mol/L)

-7

-8 0 0.2 0.4 0.6 0.8 1 Ammount of 0.1M Cu added (L)

Figure 4.11: Distribution diagram for copper-thiocyanate and water system. The initial concentration of thiocyanate and copper is 0.001M at 25oC with the addition of copper (II).

Figure 4.12 and Figure 4.13 show the concentration of free thiocyanate and hydrogen thiocyanate, respectively. Figure 4.12 shows that there is a drop in free thiocyanate at a pH of 0 compared to the simulations that were at the higher pH of 7 and 12. This is because at low pH there is an abundant amount of hydrogen ions to protonate thiocyanate. The effect of changing Eh did not have an effect on the concentration of free and hydrogen thiocyanate (Figure 4.12 and Figure 4.13).

106 0.06

0.05

0.04

0.03

0.02

0.01 Eh = -0.25, 0 or 1 at pH 7

Eh = -0.25,0 or 1 at pH 0 or 12 Thiocyanate concentration (mole/litre) 0 0 102030405060708090100 SCN:Cu Ratio

- Figure 4.12: A comparison of free thiocyanate concentration (SCN ) at different pH and Eh conditions as the ratio of thiocyanate to copper is changed.

0.003 Eh = -0.25,0 or 1 at pH 7 or 12

0.0025 Eh = -0.25,0 or 1 at pH 0

0.002

0.0015

0.001

0.0005 Hydrogen thiocyanate Concentration (mole/litre) 0 020406080100 SCN:Cu Ratio

Figure 4.13: A comparison of hydrogen thiocyanate concentration (HSCN) at different pH and Eh conditions as the ratio of thiocyanate to copper is changed.

107 As the thiocyanate to copper ratio approached a ratio of 15:1, the simulations that were conducted at a pH of 7 and 12 (Figure 4.14) shows that the concentration of CuSCN+ reached a maximum concentration of 0.005M before decreasing gradually as the thiocyanate to copper ratio continued to be increased. As the ratio of thiocyanate to copper ratio increased above 15:1, the concentration of CuSCN(s) was the only species that increased (Figure 4.15), this accounts for the loss of free copper in solution. The concentration of CuSCN+ for simulations conducted at a pH of 0 is lower compared to simulations at a pH of 7 and 12 (Figure 4.14). At each pH level + changing the Eh of the system did not change the concentration of CuSCN .

0.0006

Eh = -0.25, 0 or 1 at pH 0 0.0005 Eh = -0.25, 0 or 1 at pH 7 or 12

0.0004

0.0003 (mole/litre) 0.0002

0.0001 Copper (II) thiocyanate concentration

0 0 20406080100 SCN:Cu Ratio

Figure 4.14: A comparison of copper (II) thiocyanate concentration (CuSCN+) at different pH and Eh conditions as the ratio of thiocyanate to copper is changed.

Figure 5.15 shows that there is no difference between the nine simulations conducted with respect to the amount of CuSCN(s) present as the concentration of thiocyanate was increased. This species is the only species that was not affected when pH and Eh were changed as the ratio of thiocyanate to copper was increased.

108 0.06

0.05

0.04

0.03 (mole/litre) 0.02

0.01

Eh = -0.25, 0 or 1 at pH 0,7 or 12 Copper (I) Thiocyanate concentration

0 0 20406080100 SCN:Cu Ratio

Figure 4.15: A comparison of the amount of copper (I) thiocyanate (CuSCN) at different pH and Eh conditions as the ratio of thiocyanate to copper is changed.

0.00025

0.0002

0.00015

0.0001 (mole/litre)

0.00005 Eh = -0.25, 0 or 1 at pH 7 or 12

Eh = -0.25, 0 or 1 at pH 0 Copper (I) di-thiocyanate Concentration 0 0 20406080100 SCN:Cu Ratio

Figure 4.16: A comparison of copper (I) di-thiocyanate concentration (Cu(SCN)2) at different pH and Eh conditions as the ratio of thiocyanate to copper is changed.

109

Shown by Figure 4.16, there is a lower concentration of copper (II) di-thiocyanate

(Cu(SCN)2) in very acidic conditions (pH = 0) compared to the other simulations carried out at pH of 7 and 12. As the ratio of SCN- to copper increased to 40:1 the concentration of Cu(SCN)2 reached a maximum concentration of 0.0002M for solutions of pH 7 and 12 and 0.00014M at a pH of 0. Figure 4.17 shows the concentration of copper (I) tetra-thiocyanate as the ratio of thiocyanate to copper(II) ratio is increased. The concentration of this species maybe insignificant compared to the other copper-thiocyanate species but there is an interesting observation. When the 3- ratio of SCN:Cu was lower than 15, there was no formation of Cu(SCN)4 . As the 3- SCN:Cu ratio increased past 15:1 the concentration of Cu(SCN)4 increased exponentially until the ratio of SCN:Cu reached a value of 65:1 for simulations carried out at a pH of 0 and a ratio of 90:1 at other pH levels studied. After the 3- SCN:Cu ratio was above 90:1 the concentration of Cu(SCN)4 was the same for all pH levels studied.

7.0E-07

Eh = -0.25, 0 or 1 at pH 0 6.0E-07 Eh = -0.25, 0, 1 at pH 7 or 12

5.0E-07

4.0E-07

3.0E-07 (mole/litre)

2.0E-07

1.0E-07

Copper (I) tetra-thiocyanate concentration 0.0E+00 0 20406080100 SCN:Cu Ratio

3- Figure 4.17: A comparison of copper (I) tetra-thiocyanate concentration (Cu(SCN)4 ) at different pH and Eh conditions as the ratio of thiocyanate to copper is changed.

110 A second set of simulations was carried out by the addition of copper (II) ions instead of free thiocyanate increasing the ratio of copper to thiocyanate. Figure 4.18 to Figure 4.22 show the concentration of the individual species present in the different simulations. For each simulation when copper(II) was added to the system did not show the formation of copper(I) tetra-thiocyanate. This is because the formation of copper(I) tetra-thiocyanate needs a high ratio of thiocyanate to copper which is not achieved when the concentration of copper(II) is increased.

0.001

0.0009 Eh = -0.25, 0, 1 at pH 7 or 12

0.0008 Eh = -0.25, 0, 1 at pH 0

0.0007

0.0006

0.0005

0.0004

0.0003

0.0002

0.0001 Thiocyanate concentration (mole/litre)

0 0 102030405060708090100 Cu:SCN Ratio

- Figure 4.18: A comparison of free thiocyanate (SCN ) at different pH and Eh conditions as the ratio of copper(II) to thiocyanate changed.

Most of the copper added to the simulation remained as free copper (II) ions, but as the level of copper increased the amount of thiocyanate ion (SCN-) within the simulation decreased for all the simulations studied (Figure 4.18) due to the formation of different complexes between copper and thiocyanate. Only in simulations conducted with a pH of 0 showed relatively significants amounts of hydrogen thiocyanate (Figure 4.19). In highly oxidising environments (as in simulation 6 with an Eh of 1V, from Table 4.2) the concentration of hydrogen thiocyanate was ten times the concentration compared to the other simulations (simulations 4 and 5) that were conducted at a pH of 0.

111 0.001 Eh = 0.0 or 1.0V at pH 0 0.0009 Eh=-0.25V at pH 0

0.0008

0.0007

0.0006

0.0005

0.0004 (mole/litre)

0.0003

0.0002

0.0001 Hydrogen Thiocyanate concentration

0 0 20406080100 Cu:SCN Ratio

Figure 4.19: A comparison of hydrogen thiocyanate (HSCN) under different Eh conditions at a pH 0 as the ratio of copper(II) to thiocyanate changed.

Figure 4.20 shows the concentration profile of copper (II) thiocyanate (CuSCN+). The CuSCN+ profile is similar to results when thiocyanate (Figure 4.14) was added to solution instead of copper (II). If the ratio of copper to thiocyanate is either too large or small the concentration of CuSCN+ falls. The maximum CuSCN+ concentration of 0.005M occurs when the copper to thiocyanate ratio was 30:1 (Figure 4.20), a similar result was found when the thiocyanate to copper ratio was 30:1 as seen in Figure 4.14. As the ratio of copper to thiocyanate is increased the amount of CuSCN(s) falls from 0.001 to 0.0006M (Figure 4.21). CuSCN(s) accounts for the majority of thiocyanate within the simulation but as the ratio of copper to thiocyanate changes, thiocyanate forms another copper complex, which is CuSCN+. Figure 4.22 shows the concentration profile for copper (II) di-thiocyanate (Cu(SCN)2). When the pH is 0, the concentration of Cu(SCN)2 is lower than simulations conducted at a pH of 7 and 12. For the simulations where additional copper was added, there was no significant 3- concentrations of copper(I) tetra-thiocyanate (Cu(SCN)4 ). Only when the ratio of 3- thiocyanate to copper was greater 20:1 a small amount of Cu(SCN)4 formed. The 3- maximum concentration of Cu(SCN)4 was found when the SCN:Cu ratio was 60:1 at 6 x 10-7M.

112 0.0006

0.0005

0.0004

0.0003 (mole/litre) 0.0002

0.0001 Eh = -0.25, 0 or 1 at pH 0

Copper (II) thiocyanate Concentration Eh = -0.25, 0 or 1 at pH 7 or 12 0 0 102030405060708090100 Cu:SCN Ratio

+ Figure 4.20: A comparison of copper (II) thiocyanate (CuSCN ) at different pH and Eh conditions as the ratio of copper(II) to thiocyanate changed.

0.0012

0.001

0.0008

0.0006 (mole/litre) 0.0004

0.0002

Copper (I) thiocyanate Concentration Eh = -0.25, 0 or 1 at pH 0, 7 or 12 0 020406080100 Cu:SCN Ratio

Figure 4.21: A comparison of copper (I) thiocyanate (CuSCN) at different pH and Eh conditions as the ratio of copper(II) to thiocyanate changed.

113 0.000004 Eh = -0.25, 0, 1 at pH 7 or 12 0.0000035 Eh = -0.25, 0 or 1 at pH 0 0.000003

0.0000025

0.000002

(mole/litre) 0.0000015

0.000001

0.0000005 Copper (II) di-thiocyanate Concetration 0 0 102030405060708090100 Cu:SCN Ratio

Figure 4.22: A comparison of copper (II) di-thiocyanate (Cu(SCN)2) at different pH and Eh conditions as the ratio of copper(II) to thiocyanate changed.

4.8 Conclusion

Pourbaix diagrams can be used to describe complex chemical systems showing the thermodynamic stable species under different pH and Eh conditions. From this study Pourbaix diagrams have been used to describe the thiocyanate-water and copper- thiocyanate-water systems. As an extension to the Pourbaix diagram, the distribution diagrams were created to show the equilibrium concentrations of species present in the copper-thiocyanate-water system.

In many gold mines, the tailing waters are alkaline and have a chemical potential of 0V. If these solutions contain thiocyanate the Pourbaix diagrams developed in this chapter (such as Figure 4.2) show that the predominant stable species formed under these conditions at equilibrium is cyanate. The pH of the tailing waters will drop due to changing environmental conditions, for example the adsorption of carbon dioxide. Under extreme acidic conditions the Pourbaix diagram show that thiocyanate will become the predominant stable species over cyanate at a neutral chemical potential or

114 a slightly reducing environment. But in the natural environment the pH level will not drop to the level where the predominant species is thiocyanate.

Cyanide can be recovered from thiocyanate if the oxidation of cyanide to cyanate can be stopped. This was confirmed when cyanate and hydrogen cyanate were removed as stable species from the simulation. The resulting Pourbaix diagram (Figure 4.3) shows that the predominant areas of cyanate and hydrogen cyanate were replaced with cyanide and hydrogen cyanide depending on the pH of the system. The addition of copper makes the system more complex. In comparison to the system without copper, at equilibrium thiocyanate will form a stable complex of copper(I) thiocyanate which is a solid under the same tailing water conditions. This reduces the mobility of a toxic substance to the environment trapping thiocyanate within a solid, but this can also inhibit natural processes that can convert thiocyanate to more basic compounds that are non-toxic to the environment. In existing mines there is an opportunity of recovering cyanide from waste species like thiocyanate that is contained in tailing ponds from all the years of operation.

The distribution diagrams can give more information for a particular system considered at a particular pH and Eh. The copper-thiocyanate-water system is a very complex and the concentration of the different copper thiocyanate species is dependent on the concentration of thiocyanate and copper as well as the solution properties like pH and Eh. The species with the highest concentration among the soluble species present in solution was copper (II) thiocyanate (Cu(SCN)+). When the ratio of thiocyanate to copper was increased to 15:1 the concentration of Cu(SCN)+ increased to a maximum concentration of approximately 0.005M, as the thiocyanate to copper ratio continued to increase, the concentration of Cu(SCN)+ decreased to form other copper thiocyanate species that had a higher ratio of thiocyanate to copper, 3- for example, Cu(SCN)2 and Cu(SCN)4 . In the other case where copper concentration was increased compared to thiocyanate, the amount of free thiocyanate and hydrogen thiocyanate decreased, as there was more copper present in the system. The concentration of Cu(SCN)+ approximately reached the same maximum concentration of 0.005M, which corresponds to a copper to thiocyanate ratio of 30:1. A copper to thiocyanate ratio above 30:1 saw all the different soluble copper- thiocyanate concentrations decreased slightly with the production of CuSCN (s).

115

These simulations show equilibrium conditions only and when comparing these conclusions with any kinetic experiments the results may differ since the system has not reached equilibrium. These simulations can provide useful information about the chemistry of the system studied and how different chemical species relate to each other.

116 References

Brook P.A., “A computer method of Calculating Potential-pH Diagrams.” Corrosion Science, Vol. 11, No. 6, pp 389-396, 1971.

Angus J.C and Angus C.T., “Computation of Pourbaix Diagrams Using Virtual Species: Implementation on Personal Computers.” Electrochemical Science and Technology, Vol. 132, No. 5, pp 1014-1019, May 1985.

Duby P., “The Thermodynamic Properties of Aqueous Inorganic Copper Systems.” June 1977, Library of Congress Catalogue Card Number: 77-71709, National Standard Reference Data System, National Bureau of Standards, The United States of America.

Froning M.H., Shanley M.E. and Verink, Jr. E.D., “An improved method for calculation of potential-pH diagrams of metal-ion-water systems by computer.” Corrosion Science, Vol. 16, pp. 371-377, 1976.

Garrels R.M. and Christ C.L., “ Solutions, Minerals and Equilibria” 1965, Harper and Row Publishers New York.

Guy A.G. and Rhines F.N., “Pourbaix Diagrams: A firm basis for understanding corrosion.” Metal treatment and Drop Forging. pp 45-54, February, 1962.

Huang Hsin-Hsiung, Cuentas L., 1989(A), “Construction of Eh-pH and other stability diagrams of Uranium in a muticomponent system with a microcomputer – 1. Domains of Predominance Diagrams.” Canadian Metallurgical Quarterly, Vol. 28, No. 3, pp 225-234, 1989.

Huang Hsin-Hsiung, 1989(B), “Construction of Eh-pH and other stability diagrams of Uranium in a multicomponent system with a microcomputer – II. Distribution Diagrams.” Canadian Metallurgical Quarterly, Vol. 28, No. 3, pp 235-239, 1989.

117 Linkson P.B., Phillips B.D. and Rowles C.D., “Computer methods for the generation of Eh-pH diagrams.” Minerals Sci. Engng, Vol. 11, No. 2., pp 65-79, April 1979.

Morel F. and Morgan J., “A numerical method for computing equilibria in aqueous systems.” Environmental Science and Technology. Vol. 6, pp 58, 1972.

Pourbiax M., “Some applications of electrochemical Thermodynamics.” Corrosion, Vol. 6, pp 395-404, 1950.

Pourbiax M., “Atlas of Electrochemical Equilibria in Aqueous Solutions.” Second Edition 1974, National Association of Corrosion Engineers, Houston Texas USA.

Tran T. and Kim M.J., “Hydrometallurgy: Theory and Practice.” Channam National University Press, ISBN 89-7598-239-4.

Townsend, Jr. H.E., “Potential-pH diagrams at elevated Temperature for the system

Fe-H2O.” Corrosion Science, Vol. 10, pp 343-358, 1970.

Verink, Jr. E.D., “A technical note: Simplified Procedure for Constructing Pourbaix Diagrams.” Corrosion, Vol. 23, pp 371-373, 1967.

118 Chapter 5 Ion Exchange Resin Adsorption of Thiocyanate 5.1 Introduction

Ion-exchange resins have been used in the mining industry to concentrate cyanide or cyanide related wastes and even the precious metals such as gold and silver. In the mining industry, ion-exchange resins have been used in two different roles, either for the explicit recovery of the precious metals from leach liquors or the removal of cyanide species from liquors to meet environmental requirements.

Helfferich (1995) explains the mechanism of ion exchange by considering a system of spherical ion-exchanger beads of uniform size containing a counter ion A are placed in a well stirred solution of an electrolyte BY, where B is another counter ion. As equilibrium is approached, ion A diffuses out of the beads into the solution, and ion B diffuses from the solution into the beads. This interdiffusion of counter ions is known as ion exchange. The rate-controlling step of ion exchange is usually the diffusion of the counter anions rather than the “chemical” exchange reaction at the fixed ionic groups, which means that the ion-exchange process is essentially a diffusion phenomenon. Reaction rate constants maybe defined formally, but in their physical interpretation, have little in common with rate constants of actual chemical reactions. Ion-exchange is inherently a stoichiometric process because each counter ion which leaves the ion exchanger, is replaced with another counter ion to maintain the substrate in an electrically neutral state.

This chapter examines the binary ion-exchange equilibria between thiocyanate and chloride at 303 K using commercially available Purolite A500/2788 as the ion- exchanger. The experimental data was correlated and compared using various

119 adsorption isotherm models to determine the apparent equilibrium constants for the binary system and to look at the effectiveness of using ion exchange resins to concentrate and remove thiocyanate from mine liquors.

5.2 Theory

5.2.1 Characterisation of ion-exchange equilibrium in a binary mixture

Ion exchange equilibrium is obtained when the ion exchanger is placed in an electrolyte solution that contains a counter ion which is different from which is already adsorbed on the ion exchanger. For example, suppose the ion exchanger is in the A-form and the counter ion in the solution is B. Counter-ion exchange occurs, and ion A in the ion exchanger is partially replaced by counter ion B (equation 5.1), illustrated graphically in Figure 5.1.

+↔+ ABBA [5.1]

B A - - - - B B A - B - - - A - - A A - - A - - A A - - - - A - A - A B B B B - A - Initial State Equilibrium

- Matrix with fixed charges A B Counter ions - Co ions Figure 5.1: Ion exchange with a solution. For example a cation exchanger containing counter ions A is placed in a solution containing counter ions B (left). The counter ions are redistributed by diffusion until equilibrium is attained (right). (Helfferich, 1995).

120 5.2.2 Adsorption Isotherm Models

Many authors have used different adsorption models to describe the relationship between equilibrium loading levels and solution concentration of different counter ions under different operation conditions. Three adsorption isotherm models that have been frequently used to describe adsorption processes are the Langmuir, Freundlich and Dubinin-Radushkevich models. More recent studies have employed the law of mass action to define the equilibrium constant for the ion-exchange process.

5.2.2.1 Langmuir Isotherm

The Langmuir isotherm was developed to describe the adsorption of gas molecules onto homogeneous solid surfaces (crystalline materials) that exhibit one type of adsorption site (Langmuir, 1918). This model has been extended to describe the adsorption of ions onto solid adsorbates including heterogeneous solids.

The Langmuir isotherm has three underlying assumptions (Atkins, 1995):

1. Adsorption cannot proceed beyond monolayer coverage.

2. All sites are equivalent and the surface is uniform (that is, the surface is perfectly flat on a microscopic scale).

3. The ability of a molecule to adsorb at a given site is independent of the occupation of neighbouring sites.

At equilibrium the Langmuir isotherm relates the amount of solute adsorbed per unit weight of solid to the equilibrium solution concentration of the adsorbate, shown in equation 5.2:

CbX X = m [5.2] + bC1

121 where X is amount of solute adsorbed per unit weight of solid, b is constant related to the energy of adsorption, Xm is the maximum adsorption capacity of solid and C is the equilibrium solution concentration of the adsorbate.

Simplifying equation 5.2 by the substitution of B = 1/b leads to:

CX X = m [5.3] + CB

Taking the reciprocal of equation 5.3 and multiplying both sides by X.Xm gives a linear relationship relating X to X/C (equation 5.4). By plotting X versus X/C, one can determine the value for –B from the slape of the line of best-fit and the value of

Xm from the intercept.

⎛ X ⎞ −= BX ⎜ ⎟ + X m [5.4] ⎝ C ⎠

5.2.2.2 Freundlich Isotherm

The Freundlich isotherm is an empirical model that relates the solute adsorbed on the solid to the equilibrium solute solution concentration by a logarithmic change. Freundlich (1926) defined this relationship shown in equation 5.5.

= KCX N [5.5] where X (mmol/ml Resin) is the amount of solute adsorbed per unit weight of solid, C (mmol/L) is the equilibrium solute solution concentration and K and N are constants.

Equation 5.5 can be transformed to a linear relationship by taking the logarithms of both sides of the equation, shown below:

122 C log NK logX log logX += NK log C [5.6]

In a plot of log X versus log C, the line of best-fit has a slope of N, and log K is its intercept. When N = 1, the Freundlich isotherm that is represented by equation 5.5 reduces to a linear relationship, and because X/C is the ratio of the amount of solute adsorbed to the equilibrium solution concentration (the definition of Kd), the constant

K is equivalent to the value of Kd.

At low solute concentrations the adsorption isotherms are often linear, either the Freundlich isotherm with N = 1 or the Langmuir isotherm with bC much greater than 1 fits the adsorption data. In many cases the value of N differs from 1 showing non- linear isotherms and the Freundlich isotherm provides a superior correlation of the data than the Langmuir isotherm.

5.2.2.3 Dubinin-Radushkevich Isotherm

This model has been used to describe the adsorption of trace components. Should the adsorbent surface become saturated or the solute exceeds its solubility product, the model is inappropriate (Serne, 1992). Compared to the Langmuir isotherm, Dubinin- Radushkevich isotherm model is more general because it does not require either homogeneous adsorption sites or constant adsorption potential.

The model is described by equation 5.7 shown below (Dubinin et al., 1947):

−Kε 2 = meXX [5.7]

where X is the observed amount of solute adsorbed per unit weight, Xm is the sorption capacity of adsorbent per unit weight, K is a constant and ε is defined as

⎛ 1 ⎞ ε ⎜1RTln += ⎟ [5.8] ⎝ C ⎠

123 where R is the gas constant, T is temperature (as Kelvin) and C is the solute equilibrium solution concentration (in mM)

Equation 5.7 can be rearranged to show a linear relationship for the Dubinin- Radushkevich isotherm (equation 5.8):

2 K Xln Xln Xln Xln m −= K ε [5.9]

2 A plot of ln X versus ε allow the estimation of the ln Xm as the intercept and –K as the slope of the straight-line.

5.2.2.4 Mass-Action Sorption Model

Mass action laws and equilibrium constants have been used to described ion exchange processes. For high concentrations of counter ions loaded on the resin and in solution, the mass action law approach to model equilibrium is more realistic (Gomes et al., 2001). de Lucas et al. have successfully used mass-action models to describe different equilibrium cation ion exchange processes (de Lucas et al., 1993,1994,1998 and 2001; Valverde et al. 1999 and 2001).

Suppose that an ion exchanger (r) is initially in the form of counter ion A, and that the counter ion in the solution is B. Counter ion exchange occurs, and the A ions in the ion exchanger are partially replaced by B ions (equation 5.10).

α− β− α− β− βA r αBs βAs +⇔+ αBr [5.10]

In this reversible equilibrium, both the ion exchanger and the solution contain both competing counter ion species, A and B.

The simplest form of the mass action law, which is without activity coefficients is defined in equation 5.11:

124 mm AB K AB == Constant [5.11] mm BA

where mi is the concentration of ion i in the solution and mi is the concentration of ion i loaded on the ion exchanger. The selectivity coefficient (KAB) is a constant and is independent of the relative amounts of the competing species A and B.

True thermodynamic mass-action conceptual models assume that only free ion species are adsorbed; the ion exchange capacity is constant and independent of pH and the solution composition or total ionic strength; solutes are a true solution, which means the species’ activities are not affected by the presence of a solid and exchange of the counter ions is reversible and there is no hysteresis between adsorption and desorption. Most deviations away from the mass-action law consider only one type of adsorption site with a fixed adsorption energy (Serne, 1992).

From this basic form of the mass-action law de Lucas (de Lucas et al., 1993,1994,1998 and 2001) expands equation 5.11 to include individual activity coefficients for each ion on the ion exchanger and solution. Secondly de Lucas introduces ionic fraction normalising the ion exchange equilibria data. The ionic fraction can be defined by equation 5.12 for solution species A and B and equation 5.13 for ions adsorbed by the ion exchanger.

αC βC x = A x; = B [5.12] A N B N

αq A βq B y A = ; y B = [5.13] q o q o

th where Ci is the concentration of the i ionic species in solution (mmol/L), N is the total ionic concentration in the solution phase or the solution normality (mmol/L), qi th is the solute concentration of the i species in the solid phase (mmol/L resin) and qo is the useful capacity of the resin in the system studied (mmol/L resin)

125 de Lucas uses the mass action law to define the equilibrium constant for the ion reaction (equation 5.10) giving a relationship to explain the exchange of two counter ions A and B (equation 5.14). Activity coefficients for each ion (γA and γΒ) have been included to take into account real behaviour of the ions in solution and on the ion- exchanger. CT is defined as total concentration of all ions in solution (mmol/L).

(y γ α − )Cx[(1) γ ]β (T)K = BB ATB [5.14] AB β α − )y[(1 γ AB (x] γ TBB )C

Past work described by de Lucas et al. (1994) explains that there is a relative similarity between the phenomena of adsorption and ion exchange. The activity coefficients for the resin phase can be correlated using analogous equations proposed by Wilson (1964) for liquid-vapour equilibria.

Wilson (1964) proposed a relationship to describe liquid-vapour equilibria using Gibbs Excess Energy of Mixing (equation 5.15).

G E ∑∑i −−= xln(1x Λ jij ) [5.15] RT ij

where xi is the mole fraction of component i, Λji are the Wilson interaction parameters

(Λii = 0, Λij ≠ Λji).

The following equation results for a binary system (equation 5.15):

G E −= x x(xln Λ −+ x) + x(xln Λ ) [5.16] RT A ABBA B BAAB

From equation 5.15, Wilson (1964) defined the activity coefficient for each component to be defined as below:

126 j (1x − Λij ) x(1ln ln γi −−= ∑∑x(1ln Λ jii 1) −+ [5.17] jj− ∑ x1 Λ kjk k

For a binary system A and B, the activity coefficients for the resin phase can be defined by equation 5.18 and 5.19:

y A y Λ BAB lnγ A 1 y(yln Λ ABBA ) −+−= − [5.18] + yy Λ ABBA + yy Λ BAAB

y B y Λ ABA lnγ B 1 y(yln Λ BAAB ) −+−= − [5.19] + yy Λ BAAB + yy Λ ABBA

The activity coefficients in the liquid phase were determined from the extended Debye-Huckel limiting law (Robinson et al., 1965) shown by equation 5.20.

21 IzzA log γi10 −= [5.20] + IBa1

where γi is the activity coefficient for counter ion i, A and B are Debye-Huckel parameters involving absolute temperature and the dielectric constants of the solvent, I is the ionic strength in the bulk phase and a is the ionic diameter of the hydrated ions i.

The extended Debye-Huckel limiting law often holds up to an ionic strength of about I = 0.1 molar, when the ions are separated on the average by no more than about 20 angstroms (Robinson et al., 1965). Since the counter ion concentration is very dilute and the resulting ionic strength of the solution is less than 0.1 molar, the values of A and B for the Debye-Huckel parameters for the ion exchange system can be estimated by A and B values for water. Table 5.1 shows values for A and B for water at various temperatures.

127 The ionic diameters of hydrated ions were calculated from the ionic mobilities of the respective ion by Kielland (1937). For the counter ions of chloride and thiocyanate, -8 -8 Kielland estimated them as aCl = 3.0 x 10 cm and aSCN = 3.5 x 10 cm, respectively.

Table 5.1: Values for Parameters in the Debye-Huckel law (Appendix 7.1, Robinson et. al., 1965) Temperature (oC) A [(mol/L)-1/2] 10-8B [(mol/L)-1/2cm-1] 0 0.4918 0.3248 5 0.4952 0.3256 10 0.4989 0.3264 15 0.5028 0.3273 18 0.5053 0.3278 20 0.5070 0.3282 25 0.5115 0.3291 30 0.5161 0.3301 35 0.5211 0.3312 38 0.5242 0.3318 40 0.5262 0.3323 45 0.5317 0.3334 50 0.5373 0.3346 55 0.5432 0.3358 60 0.5494 0.3371 65 0.5558 0.3384 70 0.5625 0.3397 75 0.5695 0.3411 80 0.5767 0.3426 85 0.5842 0.3440 90 0.5920 0.3456 95 0.6001 0.3471 100 0.6086 0.3488

128 Ionic strength of the solution phase in molar units can be calculated by equation 5.20:

1 2 = 2 ∑ i czI i [5.21] i

where zi is the charge number of an ion i (positive for cations and negative for anions) and ci is the molar concentration of ion i.

The mass-action equilibrium constant (KAB) and Wilson binary interaction parameters

(ΛAB and ΛBA) were determined for the system comprising counter ions of chloride and thiocyanate using a non-linear regression method based on the Marquardt algorithm shown in Figure 5.2. The value of these constants (KAB, ΛAB and ΛBA) should be viewed as fitted parameters and do not represent the true thermodynamic equilibrium constants for the ion-exchange processes of interest. The determined values for ΛAB and ΛBA generally range from 0 to 2, which is similar for values reported for cation-exchange systems (de Lucas et al., 1993; Valverde et al., 2001). Note when using equation 5.14 to describe the equilibrium constant assumes that the swelling-pressure effects are negligible (Helfferich, 1995) and the equilibrium constant (KAB) is independent of the total concentration of ions in solution while the binary interaction parameters (ΛAB and ΛBA) depend on the total concentration of ions in the resin phase.

129

Start

Data

X, Y, I

Initial

KAB, ΛAB, ΛBA

Calculate

γ , Y i calc

Calculate

2 ε = Σ(Ycalc – Yexp) Yexp

ε Print Minimum? Finish

KAB, ΛAB, ΛBA

Marquardt Subroutine

New: KAB, ΛAB, ΛBA

Figure 5.2: Algorithm to calculate the three parameters: KAB, ΛAB, ΛBA to describe the ion exchange equilibria using the Mass-action law.

130 5.3 Experimental

5.3.1 Reagents and Materials

A commercially available anion exchange resin was used as the ion exchanger. Purolite has developed an ion exchange resin that is suited to the gold industry namely, Purolite A-500u/2788. This resin is based on a resin developed for water purification Purolite A-500. The difference between A-500 and A-500u/2788 is that A-500u/2788 has been sized to give a bead size of approximately 1mm in diameter. Table 5.2 shows the typical chemical and physical characteristics for Purolite A-500 and A-500u/2788 anion exchange resin. Prior to usage, the resin was soaked in distilled water for 24 hours to stabilize resin swelling.

Analytical grade sodium thiocyanate (98%) was supplied by BDH Chemical Ltd. Hydrated ferric nitrate (98%) and nitric acid (70% wt/wt), which was used in the determination of thiocyanate was supplied by Asia Pacific Speciality Chemicals Limited. All solutions were prepared with distilled water.

Table 5.2: Typical Chemical and Physical Characteristics for Purolite A-500 and A-500u/2788. (Purolite 2002) Typical Chemical and Physical Characteristics Polymer Structure……………………….Polystyrene crosslinked with divinylbenzene + Functional Groups…………………………………………………………..R(CH3)3N Physical Appearance………………………………………….Opaque Spherical Beads Ionic form (as shipped)……………………………………………………...…Chloride Uniformity Coefficient……………………………………………………1.7 (average) Effective size (A-500)……………………………………………….0.50mm (average) (A-500u/2788)…...…………………………………...1.00mm (average) Water Retention……………………………………………………...53-58% (Cl form) Swelling (Cl to OH)….…………………………………………………………….10% pH Limitations…………………………………………………………………….None Temperature Limitations (in OH from Max.)…………………………….140oF (60oC) (in salt form Max.)……..……………………..212oC (100oC) Chemical Resistance……… ...Unaffected by dilute acids, alkalies and most solvents Shipping Weight…………… …………………………………….…43 lbs/ft3 (690 g/l) Total Capacity………………………………………………..1.15 meq./ml Volumetric 3.9 meq./g min. Weight Whole beads…...…………… ………………………………………….95 % minimum Mechanical Strength…………………………………………..300 g minimum average

131 5.3.2 Equipment and Instrumentation

The experiment consisted of several 500 ml conical flasks where ion-exchange resin was contacted with solutions of thiocyanate. The flasks were sealed with a rubber stopper to minimise evaporation of the solution and placed in a temperature controlled bench top shaker incubator made by Paton Scientific (model number 01-2116) as seen in figure 5.3. The incubator temperature was controlled by an air heater in conjunction with a water recirculatory loop to maintain the temperature within the incubator to ±0.5oC of the desired temperature. A rotating platform was fixed to the bottom of the incubator in which the conical flasks were held in place with metal clamps. The orbital rotation speed platform can be controlled from 0 to 400 orbits per minute.

Figure 5.3: Bench-Top Orbital Shaker Incubator.

For the analysis of thiocyanate, after addition of the ferric ions, the solutions absorbance was measured at 460nm using a Perkin Elmer Lambda 3 ultra violet/visual spectrophotometer.

132 5.3.3 Procedure for analysis of thiocyanate

The method used to analyse the concentration of thiocyanate was adapted from ASTM D 4193-95, a standard test method for thiocyanate in water. Thiocyanate reacts with ferric ions at a pH of < 2 to form a coloured complex, which is determined colourmetrically at 460 nm and adheres to Beer’s Law.

The main influences that affect the effectiveness of this method are interferences due to the formation of ferric ion with thiocyanate and other coloured or interfering organic compounds that influences the colour of the solution at 460 nm. Hexavalent chromium interferes with the analysis of thiocyanate, which can be removed by adjusting the pH to 2 with concentrated nitric acid and adding ferrous sulphate. Raising the pH to 8.5 to 9 with sodium hydroxide precipitates Fe(III) and Cr(III) as the hydroxides, which are removed by filtration. Reducing agents that can reduce Fe(III) to Fe(II) prevent the formation of ferric thiocyanate. Addition of an oxidant such as hydrogen peroxide can remove reducing agents from solution. High concentrations of cyanide in proportion to the concentration of thiocyanate will react with the iron to form coloured complexes. If this problem occurs in analysis, thiocyanate is stable in both the acid and alkaline pH range. Addition of acid to reduce the pH and any cyanide will form hydrogen cyanide and can be stripped from solution by air stripping.

Coloured or interfering organic compounds must be removed by adsorption on macroreticular adsorption resin prior to analysis. Examples of interfering compounds are fluoride, phosphate, oxalate, tartrate and borate, which all form complexes with iron (Newman 1975). Other organic substances that form red colour solutions with ferric ions are complexes with phenols, enols, oximes and acetates (Shriner et al. 1948). Testing method (ASTM D 4193-95) explains a procedure to remove these interferents, but would not be needed within the analysis for these experiments because pure systems are used that do not contain these contaminants.

The ferric nitrate solution used for the reaction between Fe3+ and SCN- to form the red complex was made by dissolving 404 g of ferric nitrate (Fe(NO)3.9H2O) in

133 approximately 800 ml of distilled water. To this solution 80 ml of concentrated nitric acid was added, mixed well and diluted to 1 L with distilled water.

The ASTM states that for every 50ml of thiocyanate sample 2.5ml of ferric nitrate solution was added and mixed well. Within 5 minutes the absorbance of the solution at 460nm in a 5.0-cm cell was determined using water as a reference. The absorbance of known thiocyanate concentration was used to calculate a calibration graph to determine the concentration of the thiocyanate samples. The standard solutions were made fresh for each use to ensure the quality of the standard. Thiocyanate is a biodegradable material; some samples may contain bacteria that can consume thiocyanate. These samples can be preserved at a pH < 2 by the addition of a mineral acid and refrigerated.

5.3.4 Experimental Procedures

Ion-exchange equilibria for the SCN-/Cl- binary system was determined by contacting the Cl- form of the resin with a solution containing a fixed concentration of SCN-. A 10ml sample of Purolite A-500u/2788 that has been soaked in distilled water for 24 hours was measured and placed into 500ml conical flasks followed by the addition of 250ml of thiocyanate solution. The conical flask was sealed with a rubber stopper to minimise the evaporation of the solution during the experiment and placed in a temperature controlled bench top orbital shaker incubator made by Paton Scientific (model number 01-2116) for which the temperature was set at 30± 0.5oC. The speed of the rotating platform was set at 200 orbits per minute in all experiments.

Solution samples were taken from the conical flask at the following times during the experiment: 0, 5, 10, 15, 30, 60, 120, 240, 360 and 480 minutes and the concentration of thiocyanate was measured in accordance with ASTM D 4193-95 (as described in section 5.3.3) to monitor the adsorption of thiocyanate on the resin. Duration of 480 minutes was found to be sufficient to ensure that equilibrium has been achieved between the resin and solution. After equilibrium was established, the solution phase was separated from the resin phase. From the knowledge of the change in concentration of thiocyanate in the solution phase, it was possible to calculate the composition of thiocyanate on the resin phase at equilibrium. A mass balance was

134 used to determine the equilibrium compositions of the resin and solution phases with respect to chloride.

The partially loaded resin was then contacted with a fresh 250ml of thiocyanate solution and allowed to reach equilibrium. The process of draining the solution phase at equilibrium, followed by the addition of a fresh quantity of thiocyanate solution to the same sample of resin was repeated until no further loading of thiocyanate occurred, which was confirmed when the change in the concentration of thiocyanate in the solution was found to be negligible. Experimental data reported represent the average of duplicate run.

The effect of the total concentration on the ion-exchange isotherms was determined by contacting the resin with solutions containing different concentrations of thiocyanate. In the case of the SCN-/Cl- binary system, the SCN- concentrations of 300, 600, 900, 1200 and 1500 ppm NaSCN.

135 5.4 Results and Discussion

5.4.1 Equilibrium Study

Purolite A500u/2788 resin was contacted with five different thiocyanate solution concentrations. To ensure the resin loading reached equilibrium, regular samples were taken and the thiocyanate concentration measured. Figure 5.4 shows the adsorption profile of thiocyanate onto the resin at a starting solution concentration of 1200 ppm NaSCN. After an eight-hour period it can be seen that the concentration in solution has reached equilibrium with no change in the concentration of thiocyanate in solution. The partially loaded resin was then contacted with a fresh 250ml of thiocyanate solution and allowed to reach equilibrium. The process of draining the solution phase at equilibrium, followed by the addition of a fresh quantity of thiocyanate solution to the same sample of resin was repeated until no further loading of thiocyanate occurred, which was confirmed when the change in the concentration of thiocyanate in the solution was found to be negligible.

1400

1200

1000 Run 1 Run 2 Run 3 800 Run 4 Run 5 600 Run 6 Run 7 Run 8 400 Concentration (ppm NaSCN) 200

0 0 100 200 300 400 Time (min)

Figure 5.4: Concentration profile of thiocyanate adsorption at an initial thiocyanate concentration for each run of 1200ppm NaSCN, temperature 30oC at atmospheric pressure.

136

The average loading capacity of thiocyanate over the five different concentrations was 96.26 mg NaSCN/ml resin (1.19x10-3 mol NaSCN/ml resin) with a relative standard deviation of 2.6%. Figure 5.5 is a plot of thiocyanate loading on the resin and concentration of thiocyanate in solution. It can be seen that there is a significant difference in the adsorption performance of the resin at different solution concentrations. As the ionic strength is increased the equilibrium solution concentration is higher at a given thiocyanate resin loading.

100

50 Resin) 40

30 300ppm 600ppm 20 900ppm

Resin Concentration (mg NaSCN/ml 10 1200ppm 1500ppm 0 0 200 400 600 800 1000 1200 1400 1600 Solution Concentration (mg/l NaSCN)

Figure 5.5: Equilibrium adsorption curves for the adsorption of thiocyanate at different thiocyanate concentrations using Purolite A500u/2778 strong base ion exchange resin.

Four different models were used to define the relationship between the loading capacity and the solution concentration. Three of the models, based on the analogies with the phenomenon of adsorption are the Langmuir, Freundlich and Dubinin- Radushkevich isotherms. The other model was developed for a true ion exchange system using the Mass Action Law.

137 5.4.2 Langmuir Model Fitting

The Langmuir model was developed to understand phenomenon of adsorption and has been frequently used to describe the adsorption using ion exchange resins. Figure 5.6 is a plot of the thiocyanate concentration on the resin versus the ratio of the thiocyanate concentration to equilibrium solution concentration. This describes the Langmuir adsorption relationship of thiocyanate between the resin phase and solution phase seen by equation 5.2.

Some outlying points can be seen for the initial thiocyanate concentration of 300 ppm NaSCN. These points occur at the first initial adsorption runs that were carried out at 300 ppm NaSCN. For these first few runs the equilibrium thiocyanate concentrations after 8 hours of adsorption were very low, with concentrations of less than 1 ppm NaSCN reaching the limit of detection of the analytical technique. These points were removed when calculating the two constants for the Langmuir isotherm.

The slope of the straight line gives the constant b, described in equation 5.4, which relates to the energy of adsorption. At each different thiocyanate solution concentration there was a different value for the constant b, implying that with increasing solution concentration there is an increase in the energy of adsorption. A plot of initial concentration of thiocyanate in solution versus the constant b (Figure 5.7) was found to give a linear relationship.

The second constant that helps describe the ion exchange process is Xm, the maximum adsorption capacity of the solid in this case the ion exchange resin. For the different solution concentration carried out gave very similar values for Xm with an average value of 1.21x10-3 mol NaSCN/ ml resin with a relative standard deviation 3.3%. This was very similar to the experimental value of 1.19x10-3 mol NaSCN/ml resin. A summary of the different parameters that were determined based on the Langmuir isotherm is given in Table 5.3.

138 Table 5.3: Parameters in the model based on the Langmuir isotherm. Initial NaSCN Concentration Constants (ppm NaSCN) (mmol/l NaSCN) b (mmol/l) Xm (mmol/ml resin) 300 3.70 0.18 1.25 600 7.40 0.20 1.14 900 11.10 0.44 1.21 1200 14.80 0.51 1.20 1500 18.50 0.64 1.25

1.4

300ppm 1.2 600ppm 900ppm 1200ppm 1 1500ppm

0.8

0.6

0.4 X (mmol SCN/ml resin)

0.2

0 02468101214 (X/C) (l solution/ml resin)

Figure 5.6: The Langmuir Isotherm. A plot of the thiocyanate concentration on the resin versus the ratio of the thiocyanate concentration to equilibrium solution concentration.

139 0.7

0.6

0.5

0.4

0.3

0.2

0.1 Constant b, for the Langmuir Isotherm

0 0 200 400 600 800 1000 1200 1400 Initial Thiocyanate Concentration (ppm NaSCN)

Figure 5.7: Parameter b for the Langmuir isotherm at different solution concentration.

1.6

1.4

1.2

0.8

0.6

0.4

0.2 X Experimental (mmol NaSCN/ml resin) 0 0 0.2 0.4 0.6 0.8 1 1.2 1.4 X Model (mmol NaSCN/ml resin)

Figure 5.8: Comparison of experimental data to predicted values from the Langmuir Isotherm.

A comparison of the actual experimental results to the predicted values using the Langmuir isotherm is shown in Figure 5.8. The line within Figure 5.8 shows when the experimental and the predicted values are equal. It can be seen that the Langmuir isotherm shows good correlation (R2 = 0.985) between the experimental results and predicted values of the adsorption of thiocyanate onto Purolite A500u/2788 ion

140 exchange resin. There was greater agreement between the experimental and predicted results at higher thiocyanate concentration on the resin.

5.4.3 Freundlich Model Fitting

The second model that was used to describe the adsorption of thiocyanate onto ion exchange resin is the Freundlich Isotherm. This isotherm is an empirical model relating the thiocyanate resin concentration and solution concentration by a log-log relationship.

0.2

0 -2.5 -2 -1.5 -1 -0.5 0 0.5 1 1.5

-0.2

-0.4

log (X) -0.6

300ppm -0.8 600ppm 900ppm 1200ppm -1 1500ppm

-1.2 log (C) Figure 5.9: The Freundlich Isotherm. A plot of log (X) (resin concentration) versus log (C) (solution concentration).

The Freundlich isotherm does not describe the adsorption isotherm as well as the Langmuir isotherm. It can be seen from Figure 5.9 that the plotted results do not give a true linear relationship (R2 = 0.766). A summary of the parameters for the Freundlich isotherm is shown in Table 5.4. Comparing the experimental results to predicted values as shown in Figure 5.10, it can be seen that the data exhibit a high degree of scattering. With exception of the results for the 300 ppm NaSCN

141 adsorption experiments the constant N measured at the different concentrations of NaSCN were very similar with values ranging from 0.230 to 0.271.

Table 5.4: Parameters in the model based on the Freundlich isotherm. Initial NaSCN Concentration Constants (ppm NaSCN) (mmol/l NaSCN) N K (l/ml resin) 300 3.70 0.370 0.909 600 7.40 0.230 0.787 900 11.10 0.271 0.649 1200 14.80 0.250 0.956 1500 18.50 0.239 0.641

1.6

1.4

1.2

0.8

0.6

0.4

0.2 X Experimental (mmol NaSCN/ml resin) 0 0 0.2 0.4 0.6 0.8 1 1.2 1.4 X Model (mmol NaSCN/ml resin)

Figure 5.10: Comparison of experimental data to predicted values from the Freundlich Isotherm.

5.4.4 Dubinin-Radushkevich Model Fitting

The Dubinin-Radushkevich adsorption isotherm is a correlation between the resin loading and an inverse relationship of the equilibrium concentration. A comparison of different equilibrium relationships between the thiocyanate in solution and adsorbed onto the resin at the different initial thiocyanate concentration is shown in Figure 5.11. One of the constants that are determined by the Dubinin-Radushkevich

142 isotherm is the maximum adsorption capacity for the solid Xm. As with the Langmuir isotherm at the five different initial thiocyanate concentrations the value for Xm converge to the same value. The average value for Xm found by this isotherm was 1.11 mmol NaSCN/ml resin, with a relative standard deviation 2.4%. This value is very close to the calculated value of the total capacity of the resin for thiocyanate of 1.19 mmol NaSCN/ml resin. A summary of the parameters found is tabulated in Table 5.5. As seen in Figure 5.12, for this model the value for K is dependent on the ionic strength of the solution with K changing linearly from 2x10-8 at a concentration of 300 ppm NaSCN to 8x10-8 at 1500 ppm NaSCN. This relationship was similar to constant b in the Langmuir isotherm.

0.5 0.0E+00 2.0E+07 4.0E+07 6.0E+07 8.0E+07 1.0E+08 1.2E+08 1.4E+08 1.6E+08

300ppm -0.5 600ppm 900ppm 1200ppm -1 1500ppm ln X

-1.5

-2

-2.5 ε2

Figure 5.11: The Dubinin-Radushkevich Isotherm. A plot of the natural log of thiocyanate concentration on the resin versus ε2 which is defined by equation 5.8.

143 9.00E-08

8.00E-08

7.00E-08

6.00E-08

5.00E-08

4.00E-08 Isotherm

3.00E-08

2.00E-08

1.00E-08

Constant K for the Dubinin-Radushkevich 0.00E+00 0 200 400 600 800 1000 1200 1400 1600 Initial Thiocyanate concentration (ppm NaSCN)

Figure 5.12: The change in the value of constant K in the Dubinin-Radushkevich isotherm with a change in ionic strength of solution.

Figure 5.13 compares experimental results and the predicted values that have been 2 determined from the Dubinin-Radushkevich isotherm (R = 0.945). For a given Xexp,

Xcalc is typically higher than Xexp.

Table 5.5: Parameters in the model based on the Dubinin-Radushkevich isotherm. Initial Solution Concentration (ppm NaSCN) Xm (mmol/ml resin) K (1/(J2.mol2) 300 1.082 2.00E-08 600 1.094 3.00E-08 900 1.124 5.00E-08 1200 1.109 6.00E-08 1500 1.161 8.00E-08

144 1.6

1.4

1.2

0.8

0.6

0.4

0.2 X Experimental (mmol NaSCN/ml resin) 0 0.000 0.200 0.400 0.600 0.800 1.000 1.200 1.400 X Model (mmol NaSCN/ml resin)

Figure 5.13: Comparison of experimental data to predicted values from the Dubinin- Radushkevich Isotherm.

Ion-exchange equilibria are presented in Figure 5.14 where the data are plotted in terms of equivalent ionic fractions defined by equations 5.12 and 5.13. It was interesting when plotting the data in the form of ionic fractions that the data at the different ionic strengths converged to give the same relationship. This means that by modelling the fraction of the counter ion of interest (in this case thiocyanate) in solution and on the ion exchange resin the model becomes independent of ionic strength of the solution within the range of concentration considered.

5.4.6 Mass Action Law Model Fitting

There are three parameters that were determined from the Mass-Action Law, KAB and the two Wilson parameters to determine the activity coefficients for the solid phase. An advantage of the Mass-Action Law is that the relationship is independent of concentration. Therefore there is only one unique value of KAB to describe the ion exchange process while for the other isotherms the relationships show the dependence on the concentration of the solution.

145

To determine the three parameters for the Mass-Action Law, a Fortran program was developed (appendix A). The data that was supplied to the program was the experimental values for the fraction of thiocyanate in solution and the corresponding fraction of thiocyanate on the ion exchange resin. At each corresponding fraction the ionic strength was determined for the solution used. Then an initial guess was given to the three parameters to be determined, which were KAB = 1, ΛAB = 1 and ΛBA = 1.

Each time the system looped to recalculate the three parameters, the use of the

Marquardt Subroutine was used to determine new values for KAB and the two Wilson parameters until the error between the calculated and experimental data was minimised. KAB was determined to be 22.20. The Wilson parameters were found to be ΛAB = 1.805 and ΛBA = 0.265.

146 1

0.9

0.8

0.7

0.6

0.5

0.4

300ppm 0.3 600ppm 900ppm 0.2 Thiocyanate Fraction in Resin 1200ppm 0.1 1500ppm

0 0 0.10.20.30.40.50.60.70.80.91 Thiocyanate Fraction in Solution

Figure 5.14: The Mass Action Law. A plot of resin fraction versus solution fraction of thiocyanate.

0.9

0.8

0.7

0.6

0.5

0.4

0.3

0.2 Resin Fraction (Experimental) 0.1

0 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1 Resin Fraction (Model)

Figure 5.15: Comparison of experimental data to predicted values by the Mass-Action Law.

147

It can be seen that there is good agreement between the model and experimental data as seen in Figure 5.15. The average absolute relative deviation (AARD) was 7.20%. From Figure 5.15, the line represents the model compared to the experimental data, the model tends to slightly over estimates the resin fraction but is still good agreement between the model and the experimental data.

5.5 Conclusions

From experiments using Purolite A500u/2788 the effective capacity of the ion exchange resin for thiocyanate is 96.26 mg NaSCN/ml resin (1.19 mmol NaSCN/ml resin). This capacity is compatible to the loading capacity for chloride of 1.15 meq/ml resin quoted by the manufacturer. Since the charge of thiocyanate is unitary the unit is equivalent to mmol NaSCN/ml resin.

Four different relationships were used to describe the equilibrium between solution concentration and resin loading of thiocyanate. Three isotherms have been based on adsorption phenomenon of species onto solids, which are Langmuir, Freundlich and Dubinin-Radushkevich isotherms. The other model has been specifically developed for the ion exchange process, which is based on the Mass-Action Law.

From the different models based on the adsorption phenomenon the Langmuir isotherm gave the closest approximation of the loading characteristics compared to the experimental results. The Dubinin-Radushkevich isotherm showed good correlation as well whereas the Freundlich isotherm showed poor correlation between the calculated values and the experimental results. Normally the Freundlich isotherm is used to describe adsorption equilibria with low solute concentrations and low resin loading capacity, this usually corresponds to a linear relationship, i.e. N = 1 for this isotherm. For this study the conditions are not ideal for the Freundlich isotherm, the concentration of thiocyanate above trace levels and the loading capacity of the resin has been reached, resulting with a range of N values between 0.230 – 0.370. As the thiocyanate concentration increased the resulting N value deviated further away from

148 the value of 1, therefore showing a poor non-linear relationship for the Freundlich isotherm.

For the Langmuir and Dubinin-Radushkevich isotherms, one of the constants determined is the maximum effective loading capacity of the solid, in this case the ion exchange resin. These values are 1.21 and 1.11 mmol NaSCN/ml resin respectively. The other constant determined was b for the Langmuir isotherm and k for the Dubinin-Radushkevich isotherm. As the solution concentration was increased these constants increased linearly with solution concentration. Therefore these isotherms are solution concentration dependent.

The Mass-Action law was developed to study the ion exchange process. The constant

KAB quantifies to the relationship between the counter ion loaded on the ion exchange material and the solution concentration. In the model used to describe the adsorption of thiocyanate on a strong base anion exchange resin assumptions were made to account for solution and ion-exchange resin non-ideality by using activity coefficients for the liquid phase and Wilson parameters for the resin phase respectively. The Wilson parameters are fitting parameters for this model.

The main advantage of the Mass-Action Law over the other isotherms is that the model is independent on the concentration of thiocyanate in solution. This means that there is only one parameter that is needed to describe the thiocyanate adsorption displacing chloride ions from the ion exchange resin. The parameters determined for o the Mass-Action Law at a temperature of 30 C is KAB = 22.20 and the two Wilson parameters of ΛAB = 1.805 and ΛBA = 0.265.

By examining different binary systems with other species such as base and precious metal cyanides using the Mass-Action Law, a comparison of the relative affinity of different counter ions can be determined by using the cross product rule. This can be valued in the prediction of the loading characteristics of more complex systems found in the mining industry.

149 References

American Standard Testing Method “Standard Test Method for Thiocyanate in Water.” Designation: D 4193-95.

Atkins P. “Physical Chemistry: 5th Edition” 1995. Published by Oxford University Press, Oxford. ISBN 0-19-855730-2. de Lucas Martinez A., Canizares P. Diaz J. “Binary ion exchange equilibrium for Ca2+, Mg2+, K+, Na+ and H+ ions on Amberite IR-120.” Chemical Engineering Technology, 1993, Vol. 16, pp 35-39. de Lucas Martinez A., Zarca Diaz J., Canizares P. “Ion exchange equilibrium in a binary mixture. Models for its characterisation.” International Chemical Engineering, 1994, Vol. 34, No. 4, pp 486-497. de Lucas A., Canizares P., Garcia M., Gomez J., Rodriguez J. “Recovery of Nicotine from aqueous extracts of Tobacco wastes by an H+ form strong acid ion exchanger.” Ind. Eng. Chem. Res., 1998, Vol. 37, pp 4783-4791. de Lucas A., Valverde J., Romero M., Gomez J., Rodrigues J. “Ion exchange equilibria in nonaqueous and mixed solvents on the cationic exchanger amberlite IR- 120.” Chemical Engineering Data, 2001, Vol. 46, pp 73-78.

Dubinin M., Radushkevich L. “Equation of the Characteristic Curve of Activated Charcoal.” Proc. Acad. Sci. Phys.Chem. Sec. USSR, Vol. 55, 1947, pp 331-333.

Freundlich H. “Colloid and Capillary Chemistry.” 1926. Methuen, London, England.

Gomes C., Almeida M., Loureiro J. “Gold recovery with ion exchange used resins.” Separation and Purification Technology, Vol. 24, 2001, pp 35-57.

Helfferich F. “Ion Exchange” 1995. Published by Dover Publications Incorporated, New York. ISBN 0-486-68784-8.

150

Kielland J. “Individual Activity Coefficients of Ions in Aqueous Solutions.” Journal of the American Chemical Society. Vol. 59, No. 9, 1937, pp 1675-1678.

Langmuir I. “The Adsorption of Gases on Plane Surfaces of Glass, Mica, and Platinum.” Journal of the American Chemical Society, 1918, 40:1361-1403.

Newman A. “Chemistry and Biochemistry of Thiocyanic Acid and Its Derivatives.” Publisher: Academics Press, New York, NY, 1975.

Purolite “A-500 and A500P Anion exchange resins, Type 1 Strong-Base Macroporus Resin. Technical Data.” Purolite Product Specifications. 2002.

Robinson R., Stokes R. “Electrolyte Solutions. Second Edition (Revised)” Publishers: Butterworths, London, 1965.

Shriner R., Fuson R. “Identification of Organic Compounds.” Publisher: John Wiley & Sons Inc., New York, NK, 1948.

Serne R. “Conceptual Adsorption models and open issues pertaining to performance assessment.” Radionuclide Sorption Saf. Eval. Perspect. Proc. NEA Workshop (1992), Meeting 1991. Publisher: OECD, Paris, France.

Valverde J., de Lucus A., Rodriguez J. “Comparison between heterogeneous and homogeneous MASS action models in the prediction of ternary ion exchange equilibria.” Ind. Eng. Chem. Res., 1999, Vol. 38, pp 251-259.

Valverde J., de Lucas A., Gonzalez M., Rodriguez J. “Ion exchange equilibria of Cu2+, Cd2+, Zn2+, and Na+ ions on the cationic exchanger Amberlite IR-120.” Journal of Chemical and Engineering Data, 2001, Vol. 46, No. 6, pp 1404-1409.

Wilson G. “Vapor liquid equilibrium XI. A new expression for the excess free energy of Mixing.” Journal of the American Chemical Society. Vol. 86, 1964, pp 127-130.

151 Chapter 6

Recovery of Cyanide from Thiocyanate using Hydrogen Peroxide

6.1 Introduction

Hydrogen peroxide has been used successfully in the mining industry to control the cyanide being released to tailing ponds from leach liquors. The advantage of using hydrogen peroxide over other strong chemical oxidants is that it will naturally decompose into water and oxygen and does not cause secondary contamination if used in excess. Hydrogen peroxide can also assist the rate of gold dissolution in leaching operations by elevating the dissolved oxygen levels during its decomposition.

The study of thiocyanate oxidation by hydrogen peroxide was first studied by Kastle and Smith (Kastle et al., 1904) and it was determined that the reaction between hydrogen peroxide and thiocyanate can be approximated by a third order reaction for mixtures containing 10-2M thiocyanic acid and 3x10-2M hydrogen peroxide. Since then further studies have found that the mechanism pathway between hydrogen peroxide and thiocyanate is not simplistic. There has been strong debate on the mechanism pathway with studies showing contradictory information (Figlar et al., 2000).

152 This chapter examines the recovery of cyanide by air stripping from the oxidation of thiocyanate by hydrogen peroxide under different conditions, where hydrogen cyanide is recovered via a caustic trap. The rate of reaction was observed under different operating conditions by changing the pH and the concentration of copper and hydrogen peroxide of the system.

6.2 Chemistry of thiocyanate oxidation using hydrogen peroxide

Wilson et al. (1960 (A) and (B)) proposed that the reaction mechanism between hydrogen peroxide and thiocyanate occurs differently depending on the pH of the system. In acidic media (pH<2) the overall reaction is:

+ ][H − − 22 SCNO3H HCN 4 ++→+ 2O2HHSO [6.1]

Deviations from ideal stoichiometry shown by equation 6.1 can occur through the formation of intermediate products like S(CN)2 and products that come from the hydrolysis of S(CN)2. S(CN)2 is formed when HCN reacts with another intermediate product HOSCN. Wilson et al. (1960 (A)) proposed the following mechanism of reaction between hydrogen peroxide and thiocyanate. The initial steps involve the formation of the intermediates HOSCN and HOOSCN. The rate-limiting step was thought to be the formation of HOSCN. The formation of HOSCN proceeds via two different ways shown by equation 6.2 at pH>4 or equation 6.3 at pH<2.

− − 22 SCNOH →+ HOSCN + OH > 4)(pH [6.2]

− + 22 3OHSCNOH →++ HOSCN + 2O2H < 2)(pH [6.3]

153 HOSCN reacts with either H2O2 (equation 6.4) or HCN (equation 6.5) to form either

HOOSCN and S(CN)2 respectively:

HOSCN OH 22 →+ HOOSCN + 2OH [6.4]

HOSCN HCN →+ + 22 OHS(CN) [6.5]

Figlar et al (2000) stated that the reaction between HOSCN and HCN (equation 6.5) could reduce the rate of consumption of H2O2 by as much as a factor of 2. When the pH > 4, HOOSCN continues to react with hydrogen peroxide to produce hydrogen cyanate (equation 6.6), otherwise when the pH < 2 HOOSCN reacts with hydrogen peroxide to form hydrogen cyanide (equation 6.7).

HOOSCN + 22 → SOHOH 32 + HCNO > 4)(pH [6.6]

+ + HOOSCN 3 32 ++→+ HHCNSOHOH < 2)(pH [6.7]

Hydrogen cyanate will undergo hydrolysis under acidic conditions to produce ammonium and carbonates (equation 6.8). In an oxidative environment H2SO3 will oxidise to H2SO4 (equation 6.9).

− + HCNO 2 HCOO2H 3 +→+ NH 4 [6.8]

2232 →+ + 242 OHSOHOHSOH [6.9]

As stated previously free cyanide (CN- or HCN) in neutral or basic solutions will react with hydrogen peroxide to form cyanate (equation 6.10). Under very acidic solutions this reaction is very slow (Wilson et al., 1960 (B)) with little or no oxidation occurring.

154 OHHCN 22 +→+ 2OHHCNO [6.10]

The reason that HOSCN was used by Wilson et al. (1960(B)) as a proposed intermediate product between the reaction of hydrogen peroxide and thiocyanate (seen in equations 6.2 and 6.3) is due to the known chemical similarity between thiocyanates and halide ions and the fact that HOX is the first product in peroxide/halide reactions (Fortnum et al., 1960).

The production of S(CN)2 from thiocyanate with hydrogen peroxide was explained by the reaction between cyanide with a reaction intermediate most likely HOSCN (Wilson et al. 1960(A) and Filar et al., 2000). If the amount of hydrogen peroxide is limited to produce relatively low concentrations of hydrogen cyanide compared to thiocyanate, the complicating reaction between HOSCN and HCN could be avoided (Filar et al., 2000). Therefore in acidic conditions where cyanide exists predominantly as hydrogen cyanide, it will compete with hydrogen peroxide in - different reactions for the intermediate (HOSCN) to produce either S(CN)2 and CN , respectively. While at higher pH cyanide predominantly exists as CN-, which does not compete for the intermediate HOSCN, but in the presence of hydrogen peroxide cyanide would more likely continue to be oxidised to form cyanate.

S(CN)2 has been found to decompose to yield thiocyanate and either hydrogen cyanide or cyanate depending on the pH of the system (Wilson et al., 1960 (B); Smith et al., 1966). Under acidic conditions (pH<2) S(CN)2 will undergo acid hydrolysis according to equation 6.11. Otherwise (pH>4) S(CN)2 will undergo the following hydrolysis reaction (equation 6.12). This reaction is known to be catalysed by nucleophiles and would be significant above a pH of 4 with the hydroxyl ion being the main catalyst (Smith et al., 1966).

5 5 − 2 1 1 + 1 − 1 + S(CN)2 3 2 →+ 6 SCNOH 3 HCN 2 2 2 NHCO 4 6 4 +++++ 2 HHSO [6.11]

− − + 22 SCNOHS(CN) ++→+ 2HCNO [6.12]

155 - Figlar et al (2000) studied the reaction between ClO2 and SCN in highly acidic conditions and found that there exists an equilibrium HOSCN and (SCN)2 (shown in equation 6.13). This equilibrium was expected since thiocyanate is considered as a pseudohalide. In the presence of high concentrations of acid and thiocyanate Filar et al. (2000) believed that (SCN)2 acts as an intermediate species to eventually form products like cyanide and cyanate.

+− 22 OH(SCN) ⇔+ HOSCN ++ HSCN [6.13]

Orban (1986) reported that the addition of copper (II) ions to thiocyanate in the presence of hydrogen peroxide leads to oscillation reactions and bistability in alkaline solutions (Orban, 1986; Luo et al., 1989 and Orban et al., 2000). The reactions that make the system oscillate can be described by similar chemistry seen by the autocatalytic decomposition of hydrogen peroxide catalysed with Cu(OH)2. The cycle can be written as follows:

2+ − + 22 2OHOH2Cu 2 ++→++ 2O2HO2Cu [6.14]

+ 2+ − 22 +→+ OH22CuOH2Cu [6.15]

. . The chain carriers in equations 6.14 and 6.15 are HO2 , HO , and copper(I) according to the following reactions:

2+ − + • Cu HOO +→+ HOCu 2 [6.16]

2+ • ++ 2 ++→+ OHCuHOCu 2 [6.17]

+ 2+ − • 22 ++→+ OHOHCuOHCu [6.18]

+ • 2+ +→+ OHCuOHCu − [6.19]

156 Equations 6.14 and 6.15 yields equation 6.20, the copper catalysed decomposition of hydrogen peroxide.

22 2 +→ 2O2HOO2H [6.20]

Orban et al (2000) suggested that the role of thiocyanate is to transiently stabilise copper (I) as a complex resulting in a time delay in its reoxidation to Cu2+. A summary of the essential features of the model under alkaline conditions and its essential features by reference to the key steps are shown by equation 6.21 to 6.27:

− − 22 SCNO2H →+ OS(O)CN + 2O2H [6.21]

2OS(O)CN − → OOS(O)CN − + OSCN − [6.22]

− − • − 2OH + OS(O)CN + OOS(CN) → 2OS(O)CN + 2OH [6.23]

• + − − 2+ − OS(O)CN + }{SCNCu n → OS(O)CN Cu ++ nSCN [6.24]

2+ − 22 OHCuOH →++ HOOCu(I) + 2OH [6.25]

− + − • HOOCu(I) + nSCN → n + HO}{SCNCu 2 [6.26]

− • 2− OS(O)CN 2 SOHO 3 +→+ HCNO [6.27]

The formation and consumption of the yellow intermediate species HOOCu(I) is the species responsible for colour oscillations that occur within the solution (Orban et al., 2000). The system oscillates due to the formation of OS(O)CN- which is generated autocatalytically through the formation and reduction of the radical OS(O)CN. shown . by equations 6.21 to 6.24. As the concentration of HO2 radicals and copper

157 + - thiocyanate complexes Cu {SCN }n accumulate by reactions shown by equations 6.25 to 6.26 it reinitiates the cycle to form OS(O)CN-.

It was found that in the structure of Cu(I) thiocyanates there are strong π-bond interactions between the Cu – S bonds made possible by the overlap of occupied 3dx2- y2 and 3dz2 orbitals of copper(I) atoms with similar vacant 3dπ orbitals of the sulphur atoms of the thiocyanate group (Kabesova et al., 1978; Kabesova et al., 1980). It also was shown that π interactions predominate over the Cu – N bonds. This leads to a shift of electron density from the internuclear space C-S to space C-N, therefore in theory weakening the C-S bond. This suggests that the formation of complexes between copper (I) and thiocyanate weakens the S-C bond in thiocyanate. A graphical representation is shown in Figure 6.1.

π Cu

π π Cu N C S Cu π Cu

Figure 6.1: Graphical representation of bonding between copper(I) and the thiocyanate group.

6.3 Decomposition of hydrogen peroxide

Hydrogen peroxide decomposes with disproportionation in the following equation 6.28:

1 22 2 +→ 2 OOHOH 2 [6.28]

The reaction is highly exothermic and takes place in the presence of small amounts of catalyst even in aqueous solution shown in Table 6.1. In the absence of a catalyst, it occurs only in the gas phase at high temperature.

158 Table 6.1: Heat of decomposition of hydrogen peroxide

1 22 → 2 + 2 2 (g)OO(g)H(g)OH -105.8 kJ/mol

1 -98.3 kJ/mol 22 2 +→ 2 2 (g)OO(l)H(l)OH Source: Ullmann’s Encyclopedia of Industrial Chemistry, 5th Edition.

Decomposition can be catalysed both homogeneously by dissolved ions especially the heavy metals iron of copper, manganese and chromium and heterogeneously by suspended oxides and hydroxides such as manganese, iron, copper, palladium and magnesium) and by metals such as platinum, osmium and silver. The mechanism of decomposition in the presence of copper has been described previously in section 6.2. Choudhary et al. (2003) found that the decomposition of hydrogen peroxide in the presence of a metal catalyst namely palladium decreased as the concentration of sulphuric acid was increased.

6.4 Reaction rate expressions

When studying different chemical reactions reaction rate expressions are useful tools used to evaluate the mechanism or interaction of different components within a reaction. Reaction rates are usually defined as the rate of change of concentration of one or more components involved in the reaction. At a fixed temperature the reaction rate is a function of the concentrations of some or all of the various components of the system, and is usually expressed as a function of reactants only. In general it is not possible to predict the rate expression for a given reaction by just knowing the stoichiometric equation. By examining different rate expressions and experimental data a rate expression can be determined for a particular chemical reaction, in this case the reaction between hydrogen peroxide and thiocyanate.

6.4.1 Derivation of a first order integrated rate expression

From the two half-cell reactions the overall redox reaction for the oxidation for thiocyanate by hydrogen peroxide is given by equation 6.1. Let A correspond to the

159 concentration of either thiocyanate or hydrogen peroxide. The following model shown by equation 6.29, assumes that the pH of the system remains unchanged giving an overall rate of reaction pseudo 1st order. That is that the rate of reaction is dependent on the concentration of one reactant within the reaction.

dA =− k.A [6.29] dt where k is the reaction rate constant.

Rearranging equation 6.29 gives:

dA =− k.dt [6.30] A

Integrating equation 6.30 yields a rate expression comparing the concentration of A with respect to time. The limits of the integration are taken as A = Ao at t = 0 and A = A at t = t, shown by equation 6.31. If the reaction shows first order kinetics with respect to one reactant a plot of ln(A/Ao) against time will yield a linear relationship.

A ln = k.t [6.31] Ao

6.4.2 Derivation of a second order integrated rate expression

In the same manner as a first order rate reaction, let A correspond to the concentration of either thiocyanate or hydrogen peroxide. The following model shown by equation 6.32, assumes that the pH of the system remains unchanged giving an overall rate of reaction pseudo 2nd order with respect to one of the reactants.

dA =− k.A 2 [6.32] dt

160 Rearranging equation 6.32 gives:

dA =− k.dt [6.33] A 2 where k is the reaction rate constant.

Integrating equation 6.33 yields a rate expression comparing the concentration of A with respect to time. The limits of the integration are taken as A = Ao at t = 0 and A = A at t = t, shown by equation 6.34. If the reaction shows second order kinetics with respect to one reactant a plot of (1/A-1/Ao) against time will yield a linear relationship.

1 1 =− k.t [6.34] A Ao

6.4.3 Derivation of a third order integrated rate expression

From the two half-cell reactions the overall redox reaction for the oxidation of thiocyanate by hydrogen peroxide is given by equation 6.1. Let A and B correspond to the concentrations of hydrogen peroxide and thiocyanate respectively. The following model assumes that the pH of the system remains unchanged giving an overall rate of reaction pseudo-3rd order, first order with respect to the concentration of thiocyanate and second order with respect with the concentration of hydrogen peroxide (equation 6.35).

dA =− 2 BkA [6.35] dt where k is the reaction rate constant.

From the redox reaction between thiocyanate and hydrogen peroxide (equation 6.1), for every mole of thiocyanate consumed three moles of hydrogen peroxide is consumed. Therefore the consumption of hydrogen peroxide compared to thiocyanate can be defined as (equation 6.36):

161

= 3dBdA [6.36]

Integrating equation 6.36 between t = 0 and t = t gives:

o =− o − B)3(BAA [6.36] and rearranging equation 6.36 gives:

A A B o +−= B [6.37] 3 3 o

where Ao and Bo are the initial concentrations.

Substituting B from equation 6.37 into equation 6.35 to give a an equation in the two variables A and t, therefore:

dA 2 ⎛ A o A ⎞ ⎜BkA o +−=− ⎟ [6.38] dt ⎝ 3 3 ⎠ and rearranging equation 6.38 gives:

dA − = k.dt [6.39] 2 ⎛ A o A ⎞ ⎜BA o +− ⎟ ⎝ 3 3 ⎠

Now by partial fractions rewrite the left hand side as a sum of 3 simpler terms. Let

1 p q r

2 ++≡ [6.40] 2 ⎛ A o A ⎞ A A ⎛ A o A ⎞ ⎜BA o +− ⎟ ⎜Bo +− ⎟ ⎝ 3 3 ⎠ ⎝ 3 3 ⎠

162 where p, q and r are constants and the identity symbol means equality of the two sides of equation 6.40 for all values of A. The values of p, q and r can be evaluated by using a common denominator and equating coefficients of like powers of A in the numerators. Equating the numerators gives (equation 6.41):

⎛ A o A ⎞ ⎛ A o A ⎞ 2 ⎜Bp1 o +−≡ ⎟ ⎜BqA o +−+ ⎟ + rA [6.41] ⎝ 3 3 ⎠ ⎝ 3 3 ⎠

Equating constant terms when A = 0, gives (equation 6.42)

⎛ A o ⎞ ⎜Bp1 o −= ⎟ [6.42] ⎝ 3 ⎠

Therefore solving for constant p (equation 6.43),

3 p = [6.43] − A3B oo

and equating coefficients of A by setting A equal to Ao-3Bo and 1; constants q and r can be determined as shown in equations 6.44 and 6.45, respectively.

1 B3A oo q −= 2 − 2 [6.44] ()− 3BA oo ()oo ()−+− 3BA1A3B oo

r = 2 [6.45] ()− 3BA oo

Therefore equation 6.39 can be rewritten in the following:

163 ⎛ ⎞ ⎜ ⎟ p q r ⎜ ⎟ dA 2 ++− = k.dt [6.46] ⎜ A A A o A ⎟ ⎜ Bo +− ⎟ ⎝ 3 3 ⎠

Integrating equation 6.46 yields a rate expression comparing the concentration of A with respect to time. The limits of the integration are taken as A = Ao at t = 0 and A = A at t = t. After integration and substituting p, q and r gives equation 6.47.

⎡ ⎤ 3 ⎡ 1 1 ⎤ ⎡ 1 B3A oo ⎤ A ⎢ ⎢ − ⎥ + ⎢ + ⎥ln ⎥ − A3B A A 2 2 A ⎢ oo ⎣ o ⎦ ⎣⎢()− 3BA oo ()oo ()−+− 3BA1A3B oo ⎦⎥ 0 ⎥ − ⎢ ⎥ = kt ⎡ ⎤ ⎢ 1 Bo ⎥ + ⎢ ⎥ln ⎢ 2 B ⎥ ⎣ ⎣⎢()− 3BA oo ⎦⎥ ⎦

[6.47]

If the oxidation of thiocyanate with hydrogen peroxide is a third order reaction as described by equation 6.47, data from the experiment would yield a linear plot if the left hand side of equation 6.47 is plotted against t.

164 6.5 Experimental

6.5.1 Reagents and Materials

Analytical grade hydrogen peroxide (30 wt%) which was a certified ACS reagent was supplied by Sigma-Aldrich. Two stabilising agents were added to the hydrogen peroxide stock solution, 0.5ppm stannate and 1ppm phosphorus. This solution was kept in the original polyethylene bottle and refrigerated to avoid decomposition. To ensure that the concentration of hydrogen peroxide did not show signs of decomposition it was periodically titrated against standard potassium permanganate.

Analytical grade sodium thiocyanate (98%) and hydrated copper sulphate (98%)

(CuSO4.5H2O) were used for these experiments and were supplied by BDH Chemical Ltd and Asia Pascific Speciality Chemicals Limited. The pH adjustment was made by the addition of either analytical grade sulphuric acid (98 wt%) or a solution made from analytical grade sodium hydroxide (97%) and supplied by Asia Pacific Speciality Chemicals Limited.

The following chemicals were used to determine the concentration of hydrogen peroxide, free cyanide and thiocyanate. Analytical grade potassium permanganate (99%) supplied by Asia pacific Speciality Chemicals Limited for the determination of hydrogen peroxide. The analysis of free cyanide concentration used analytical grade silver nitrate (99%) and 5-(4-(Dimethylamino)benzylidene)rhodanine (97%) and supplied by Prolabo Ltd. and Sigma-Aldrich Ltd., respectively. Determination of thiocyanate concentration used hydrated ferric nitrate (98%) and nitric acid (70 wt%), which was supplied by Asia Pacific Speciality Chemicals Limited. All solutions were prepared with distilled water.

165 6.5.2 Equipment and Instrumentation

The experimental set-up to recover cyanide from thiocyanate consisted of a bubble column, a caustic trap, a small air pump, a precision bore flowrator and instruments to measure pH and Eh. The schematic diagram of the equipment set-up is shown in Figure 6.2, while its front view arrangement is shown in Figure 6.3. The bubble column reactor was custom made from Perspex, dimensions of the reactor given in Figure 6.4. It was made from three sections clamped together by three equally spaced tie rods placed 5mm away from the edge of the head and base of the reactor. The middle of the base was counter sunk to conceal the glass frit and a 2mm O-ring, providing even bubble size and good aeration over the cross sectional area of the reactor. When clamped down the O-ring provides a watertight seal. The head of the reactor was also counter sunk to align the middle Perspex section; the rim of the top of the cylinder was lined with vacuum grease to keep the reactor airtight.

The caustic trap consisted of a 500ml measuring cylinder fitted with a ground glass Quickfit fitting with two arms to allow air into and out of the caustic trap. The air was drawn through the caustic solution by extending the inlet from the Quickfit fitting to near the base of the measuring cylinder with 6mm silicon tubing. At the end of the tubing a stone sparger was fitted to provide small air bubbles for better mass transfer. A small diaphragm Air Cadet air pump (model: 7530-50) from Cole-Parmer Instrument Company was used to draw air through the reactor and the caustic trap. The air flow rate was controlled by a needle valve precision bore flowrator made by Fisher and Porter (Workington, London). All tubing used in the apparatus was 6mm silicon tubing supplied made by Nalgen.

A Perkin Elmer Lambda 3 ultra violet/visual spectrophotometer was used in the analysis of thiocyanate. After addition of the ferric ions to the thiocyanate solution, the absorbance was measured at 460nm.

166

1 2 3 4

Figure 6.2: Schematic diagram of equipment set-up for oxidation study of thiocyanate using hydrogen peroxide. (1) Bubble Column (2) Caustic Trap (3) Precision Bore Flowrator (4) Cadet Air Pump.

Figure 6.3: Front view of equipment arrangement for oxidation study of thiocyanate using hydrogen peroxide, showing: Bubble reactor, caustic trap, Precision bore flowrator and air pump.

167 8mm ∅6mm

∅6mm

∅6mm (A)

∅6mm ∅6mm

12mm 15mm 15mm

28mm 5mm

(B) (C) 350mm

6mm 28mm

4mm 14mm ∅6mm 12mm Figure 6.4: Dimensions of the bubble column reactor. (A) Top view, (B) Side View, (C) Front View.

168 6.5.3 Analytical Procedures to determine the concentration of thiocyanate, free cyanide and hydrogen peroxide

6.5.3.1 Procedure for thiocyanate determination

The procedure that was followed to determine the concentration of thiocyanate is described in section 5.3.3.

6.5.3.2 Procedure for free cyanide determination

6.5.3.2.1 Preparation of 0.01M standard silver nitrate solution.

o Silver nitrate (AgNO3) was dried to a constant weight in an oven at 40 C. It was allowed to cool to room temperature in a dessicator. To make the 0.01M silver nitrate solution, 1.699g dry silver nitrate was dissolved in 1000ml of distilled water. This stock solution was stored in a dark-coloured bottle.

6.5.3.2.2 Preparation of silver sensitive 5-(4- (dimethylamino)benzylidene)rhodanine indictor

Dissolve 20mg of 5-(4-(dimethylamino)benzylidene)rhodanine in 100ml of acetone.

6.5.3.2.3 Analytical procedure for cyanide analysis

The titration procedure uses silver nitrate with the rhodanine indicator to measure free cyanide concentrations greater than 1mg/l. The detection limit for this procedure is 0.3 mg/l. The main factors that can affect the effectiveness of this method to analyse free cyanide concentration are interference that actually change the free cyanide concentrations in solution. Sulphides can adversely affect the titration procedure. Samples that contain hydrogen sulphide, metal sulphides or other compounds that may produce hydrogen sulphide (this interference can be removed by distillation). Positive errors may occur with samples that contain nitrate and/or nitrite. Nitrate and

169 nitrite can form nitrous acid which can reacts with some organic compounds to form oximes. These compounds can decompose to generate HCN. The interference of nitrate and nitrite can be eliminated by pre-treatment with sulfamic acid. Weak acid dissociable cyanides present in the sample can change the concentration of free cyanide available in solution. An example of this is copper cyanide. As the level of free cyanide in solution drops there is an equilibrium shift to increase the level of free cyanide and reduce the amount of cyanide associated with copper cyanide.

A 5 ml cyanide sample was added to a conical flask. To this solution four drops of benzal-rhodanine indicator was added. If cyanide is present, the solution turns a yellowish-brown colour. Then this solution was titrated against standard silver nitrate solution until the first change in colour from yellowish-brown to pink.

The free cyanide concentration [CN-] expressed in mol/l, is given by the following formula:

][AgNO x ml)(titre, x 2 x ml)(titre, x ][AgNO - ][CN = standard3 [6.43] (sample volume, ml)

- where [CN ] is the cyanide sample concentration (M), [AgNO3]standard is the standardised concentration (M) of silver nitrate and titre (in mls) is the amount of silver nitrate used in the titration.

6.5.3.3 Procedure for hydrogen peroxide determination

6.5.3.3.1 Preparation of standard potassium permanganate solution

Standard potassium permanganate (KMnO4) solution was made by dissolving 1.86 g

KMnO4 in 800mls of distilled water. This solution was heated to boiling, then the heat was lowered to keep the solution just below boiling point for 1 hour. To remove any undissolved material from the KMnO4 solution it was passed through either glass wool or a fritted glass funnel. Additional distilled water was added to make the final

170 volume of 1000ml. The stock solution was standardised against sodium oxalate and stored in a dark-coloured bottle.

6.5.3.3.2 Standardisation of potassium permanganate solution

o Analytical grade sodium oxalate (Na2C2O4) was dried at 105-110 C for 1 hour and allowed to cool to room temperature in a dessicator. 0.30+0.01g of dry sodium oxalate was measured and quantitatively transferred to a 500ml flask. To the sodium oxalate 250ml of 5wt% H2SO4 (which has been boiled for 10-15 minutes and allowed to cool room temperature) was added to dissolve the oxalate. This solution was titrated with the prepared potassium permanganate solution until the appearance of a pink colour. Once the solution turns pink it was heated to a temperature of 60oC, upon heating the pink colour will disappear if the solution is not over-titrated. The titration was completed when additional potassium permanganate was added until the pink colour persists for about 30s. The normality of potassium permanganate is given by:

g),OC(Na N,KMnO = 422 [6.44] 4 (0.04903 x titrati ml),on

6.5.3.3.3 Analytical procedure for hydrogen peroxide analysis

This method describes the determination of hydrogen peroxide concentration in an aqueous solution. The reaction of potassium permanganate with hydrogen peroxide in acidic medium according to the following redox reaction:

− + 2+ 2MnO 4 22 2 ++→++ 2O8H2Mn5O6HO5H [6.45]

The concentration of hydrogen peroxide was determined by a titration using standardised potassium permanganate. A known volume of H2O2 sample was added to a 250 ml flask, ranging from 1 to 5mls. To this sample 10mls of 5N sulphuric acid

(H2SO4, 2.5mol/l) was added to keep the solution acidic. The hydrogen peroxide

171 sample was titrated against standardised 1.86g/l potassium permanganate solution until the solution turned pale pink colour. The titration end point has been reached when the pale pink colour persists for at least 30s. The standardised potassium permanganate solutions used in the titration must be standardised regularly.

The concentration [H2O2] of hydrogen peroxide solution expressed in mol/l, is given by the following formula:

][KMnO x ml)(titre, x 5 x ml)(titre, x ][KMnO ]O[H = standard4 [6.46] 22 2 x (sample volume, ml)

- where [H2O2] is the hydrogen peroxide sample concentration, [MnO4 ]standard is the standardised concentration of potassium permanganate and titre (in mls) is the amount of potassium permanganate used in the titration.

6.5.4 Experimental Procedure

The rate of reaction between thiocyanate and hydrogen peroxide to recover cyanide was examined by contacting a solution of thiocyanate and a solution of hydrogen peroxide in a bubble column. Experimental set-up is given in detail in section 6.5.2. Various conditions were changed within this set of experiments to evaluate the effect of initial pH of the solution, addition of copper (II) ions as a catalyst and the concentration of hydrogen peroxide to the rate of reaction between thiocyanate and hydrogen peroxide and the recovery of cyanide from the reaction. All experiments were conducted at room temperature and at a constant airflow.

As shown in Figure 6.2 and 6.3, the bubble column was connected to the caustic trap followed by a needle valve and air pump. Silicon tubing was used to join each piece of apparatus. Air was drawn through the bubble column and caustic trap keeping a slight negative pressure throughout the bubble column and caustic trap at all times. If any leaks did occur the air would be drawn through the caustic trap at all times, therefore eliminating the risk of hydrogen cyanide being released from the experiment.

172

Two 250 ml solutions were freshly made for each experiment. The first solution contained the desired amount of thiocyanate and copper sulphate and the second solution contained the desired amount of hydrogen peroxide. The initial pH of the experiment was adjusted before each experiment. Each solution was adjusted with either 0.5M sodium hydroxide solution or concentrated sulphuric acid. Care was taken to add the sulphuric acid slowly to keep the solution cool to avoid hydrogen peroxide decomposition occurring before the experiment.

400ml of 0.5M sodium hydroxide was poured to the caustic trap. An initial sample was taken from the caustic trap for a blank cyanide titration. All seals and tubing were checked before the air pump was turned on to provide negative pressure within the bubble column. The thiocyanate/copper solution was drawn into the column. After the addition of the thiocyanate solution, the hydrogen peroxide solution was drawn into the bubble column. Mixing the two solutions was achieved by continuous air bubbling (this took less than 2 seconds). A sample from the bubble column was taken after the two solutions were mixed and the experiment started. The amount of air drawn through the bubble column and caustic trap was controlled by a needle valve. Solution samples were taken from the bubble column and the caustic trap at the following times during the experiment: 0, 20, 40, 60, 120, 240, 360, 480 and 1440 minutes. The bubble column samples were analysed for thiocyanate concentration, hydrogen peroxide concentration, pH and Eh, the samples from the caustic trap were analysed for free cyanide concentration. The analytical procedures followed to determine the concentration of thiocyanate, hydrogen peroxide and cyanide are described in section 6.5.3 (Analytical procedures) to determine the concentration of thiocyanate, free cyanide and hydrogen peroxide.

Within this set of experiments the pH was varied from 100g/l sulphuric acid to pH 12. The starting concentration of thiocyanate was 1000 ppm SCN- (17.22 mM). It can be seen from equation 6.1 that for every mole of thiocyanate destroyed 3 moles of hydrogen peroxide is consumed. Based on this relationship either stoichiometric or twice the stoichiometric amounts of hydrogen peroxide were used this equates to a concentration of 1.8g/l or 3.6g/l H2O2, respectively. Addition of copper was added in

173 a ratio of 1:1 or 2:1 copper to thiocyanate on a mole basis, resulting in a copper concentration of 0.6g/l to 1.2g/l CuSO4, respectively.

174 6.6 Results

The concentration of thiocyanate in tailing waters found at gold operations could be in excess of 5000ppm SCN- (Botz et al, 2001). These concentrations maybe considered high at mines where thiocyanate is formed as a by-product in the cyanidation process. As a compromise the initial thiocyanate level for the oxidation experiments were set at 1000ppm SCN-. The effect of acid concentration, hydrogen peroxide concentration and the addition of copper (II) ions were examined in this study. The concentration levels of each parameter were related to the stoichiometric relationship of the initial amount of thiocyanate present in each experiment.

Figure 6.5 is an example of typical results from the oxidation of thiocyanate with hydrogen peroxide under acidic conditions. Thiocyanate is oxidised by hydrogen peroxide to form cyanide ion under acidic conditions. In acidic conditions cyanide will protonate to form hydrogen cyanide, which can be stripped from solution using air. This air is then bubbled through the caustic trap where cyanide is effectively removed and trapped in the caustic solution. The concentration of cyanide is determined from samples taken from the bubble column. The advantage of using an Acidification-Volatilisation-Reneutralisation (AVR) circuit to recover the cyanide from solution is it removes cyanide from an oxidative environment stopping further cyanide oxidation occurring to other species like cyanate. It can be seen from other previous thiocyanate oxidation studies using different oxidants such as ozone, chlorine and electrolytic oxidation (Botz et al. 2001, Byerley et al. 1984 and Soto et al. 1995) that the cyanide concentration increases till a point in the experiment where the rate of further oxidation of cyanide is greater than the oxidation of thiocyanate to cyanide. By not stopping further cyanide oxidation it lowers cyanide recovery and increases oxidant consumption, therefore increasing the chemical costs. From Figure 6.5 it can be seen that the chemical reaction between thiocyanate and hydrogen peroxide follows very closely to stoichiometry in acidic conditions. For every mole of thiocyanate consumed, 3.8 moles of hydrogen peroxide was used. Hydrogen peroxide consumption was higher than the stoichiometric relationship of 1:3 (mole basis) which can be due to consumption by other side reactions, for example cyanide oxidation to cyanate or decomposition of hydrogen peroxide.

175 20 90

18 80

16 70

14 60 12 Thiocyanate 50 10 Cyanide Hydrogen peroxide 40 8

30 concentration (mM) 6 2 O 2

20 H 4 SCN or NaCN concentration (mM) 2 10

0 0 0 200 400 600 800 1000 1200 1400 1600 Time (min)

Figure 6.5: Oxidation of thiocyanate using hydrogen peroxide to recovery cyanide. Initial thiocyanate concentration 1000ppm, stoichiometric amount of hydrogen peroxide and 10g/l sulphuric acid at room temperature and pressure.

The most significant influence in the recovery of cyanide from the oxidation of thiocyanate using hydrogen peroxide is the pH of the solution. This has been confirmed already by previous studies (Wilson et al,. 1960(A) and 1960(B)). The pH of the system has a dual effect in the recovery of cyanide from thiocyanate. The reaction between thiocyanate and hydrogen peroxide to produce cyanide is catalysed by hydrogen ions, therefore a low pH is required for faster kinetics. In Figure 6.6 it shows results of a set of experiments where the pH was varied under the following process conditions at an initial thiocyanate concentration of 1000ppm (17.2mM SCN- ), the amount of hydrogen peroxide was double the stoichiometric amount needed to oxidise all the thiocyanate present with an airflow rate of 1.2 L/min. It can be seen that thiocyanate is destroyed at a faster rate at lower pH. In acidic conditions cyanide ions protonate to form hydrogen cyanide, which is not readily oxidised to cyanate in the presence of an oxidant compared to the cyanide ion (CN-) (Jara et al., 1996). Therefore in very acidic media the hydrogen ions inhibit the oxidation of cyanide through the forming hydrogen cyanide, maximising cyanide recovery. This

176 20 10 g/l Sulphuric Acid initial pH 2 Initial pH 7 Initial pH 12 15

5 SCN Concentration (mM)

0 0 200 400 600 800 1000 1200 1400 1600 Time (min)

Figure 6.6: A comparison of thiocyanate concentration against time at different solution pH. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution.

20 10 g/l Sulphuric Acid pH 2 pH 7 pH 12 15

5 NaCN Concentration (mM)

0 0 200 400 600 800 1000 1200 1400 1600 Time (min)

Figure 6.7: A comparison of cyanide recovery at different initial pH’s. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution.

177 can be seen in Figure 6.7 that the lower the pH of the system the higher the cyanide recovery and it was easier to remove hydrogen cyanide from solution.

Under conditions of excess hydrogen peroxide and low pH’s, Wilson and Harris (1960(A)) noticed that the reaction showed mild inhibition by hydrogen peroxide. This may be due to a competing reaction between hydrogen cyanide and thiocyanate in the formation of side products of sulphur di-cyanide (S(CN)2) and its hydrolysis products. But by reducing the concentration of hydrogen cyanide from the system by air stripping, concentrations of hydrogen cyanide can be low enough that the complicating reaction between hydrogen cyanide and thiocyanate could be avoided. Therefore the hydrogen cyanide inhibition can be neglected when studying the reaction kinetics.

Since the reaction between thiocyanate and hydrogen peroxide produces acid the pH of the solution decreases with time. As seen in Figure 6.8 for the experiments with an initial pH of 12 and 7 the pH of the solution decreases over the first 200 minutes of the experiment reaching a constant pH for the remainder of the experiment. When comparing the cyanide recovery (Figure 6.7) and the pH of the solution (Figure 6.8) for the experimental run with an initial pH of 7 it was only after the pH of the system stabilised was when the rate of cyanide recovery increased. In the case with an initial pH of 12, there may not be enough acid produced to protonate the cyanide produced to protect it from further oxidation resulting in a very low cyanide recovery. It is evident that the pH 2 and 10g/l sulphuric acid that the pH affects the rate at which hydrogen cyanide is removed from the bubble column.

As the pH of the solution decreases with reaction time, the Eh of the system increases, which is due to the additional acid formed during the reaction, as seen in Figure 6.9. For experiments where the pH was less than 2, the amount of acid produced did not change the pH of the system and it remained relatively constant and inturn the Eh of the system did not change significantly. When the initial pH of the system started at

12 the initial Eh 0.071 V (vs. SCE) increasing to 0.382 V (vs. SCE) after 200 minutes.

178 12

10 g/l Sulphuric acid

10 Initial pH 2 Initial pH 7 Initial pH 12 8

6 pH

0 0 200 400 600 800 1000 1200 1400 1600 Time (min)

Figure 6.8: The pH profile of hydrogen peroxide oxidation of thiocyanate with different initial starting pH. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution.

0.6

0.5

0.4

0.3 (V vs SCE) h

E 0.2 10 g/l sulphuric acid initial pH 2 0.1 Initial pH 7 Initial pH 12

0 0 200 400 600 800 1000 1200 1400 1600 Time (min)

Figure 6.9: The Eh profile of hydrogen peroxide oxidation of thiocyanate with different initial starting pH. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution.

179 Copper(II) ions was added to the system to see if the reaction between hydrogen peroxide and thiocyanate would be affected by the copper. In alkaline systems it was found that addition of copper(II) ions leads to oscillation reactions and bistability (Luo et al., 1989; Orban, 1986 and Orban et al., 2000). These oscillation reactions were only stable in alkaline conditions (pH >12), since the overall reaction between thiocyanate and hydrogen peroxide produces acid, to maintain a pH > 12 addition of an alkali was needed. In acidic conditions these oscillation reactions are not stable but copper may act in a similar way outlined by Orban (1986). Copper enhances the autocatalytic decomposition of hydrogen peroxide. In one of the steps of decomposition of hydrogen peroxide copper (II) ions are reduced to form copper(I) ions, which then complex with thiocyanate present in solution. From Pourbaix diagrams (Figures 4.4-4.7) it was found that the stability zone of copper (I) thiocyanate complexes was larger than for copper (II) thiocyanate complexes. These are the predominant species formed in the copper-thiocyanate-water systems under the operating conditions carried out in the experiments. By removing free thiocyanate from the system side reactions that consume hydrogen cyanide can be avoided. Oxidation may be aided by the addition of copper to the system because when copper (I) complexes with thiocyanate the bond between the sulphur atom and carbon atom weakens as electron density is rearranged and inturn strengthening the triple bond between the carbon and nitrogen.

From the initial oxidation experiments it was found that a low pH is required for high cyanide recovery. A set of eight experiments was conducted to understand the oxidation of thiocyanate and hydrogen peroxide in the presence of copper. The conditions for each experiment are tabulated in Table 6.2.

Figure 6.10 is a comparison of thiocyanate concentration against time under different experimental conditions outlined in Table 6.2. These experiments showed faster kinetics than experiments shown in Figure 6.6 that did not contain any copper. The lowest residual thiocyanate concentration was 0.3ppm SCN-.

180 Table 6.2: Experimental conditions for thiocyanate oxidation experiments with the addition of copper. The initial concentration of thiocyanate for each experiment is 1000 ppm SCN-. Experiments were carried out at room temperature and pressure. Airflow rate 1.2 L/min. Experiment Copper Hydrogen peroxide Sulphuric acid concentration concentration concentration (g/l) (g/l) (g/l) 1 0.6 1.8 10 2 1.2 1.8 10 3 1.2 3.6 10 4 0.6 3.6 10 5 0.6 3.6 10 6 0.6 3.6 100 7 0.6 1.8 100 8 1.2 3.6 100

20 Experiment 1 Experiment 2 Experiment 3 Experiment 4 15 Experiment 5 Experiment 6 Experiment 7 Experiment 8 10

5 SCN concentration (mM)

0 0 200 400 600 800 1000 1200 1400 1600 Time (min)

Figure 6.10: A comparison of thiocyanate concentration against time at different experimental conditions outlined in Table 6.2.

181 Figure 6.11 is a plot of residual hydrogen peroxide concentration against time for each experiment that contained copper. Four experiments contained double the stoichiometric requirement of hydrogen peroxide (experiments 3,5,6 and 8), Figure 6.11 shows that the consumption of hydrogen peroxide was greater than half the amount of hydrogen peroxide added to the experiments. This is due to the continued decomposition of hydrogen peroxide in the presence of a copper catalyst. After all the thiocyanate was consumed in the reaction the concentration of hydrogen peroxide slowly decreased even further with time which could be due to further oxidation of cyanide to cyanate and then even leading to the oxidation of cyanate to nitrate and carbonates.

180 Experiment 1 160 Experiment 2 Experiment 3 140 Experiment 4 Experiment 5 120 Experiment 6 Experiment 7 100 Experiment 8

20 Hydrogen peroxide concentration (mM) 0 0 200 400 600 800 1000 1200 1400 1600 Time (min)

Figure 6.11: A comparison of hydrogen peroxide concentration against time under different experimental conditions outlined in Table 6.2.

With the addition of copper to the oxidation of thiocyanate, cyanide recovery continued to be high, as shown in Figure 6.12. Experiment 6 recovered 100% of the theoretical cyanide produced if all the thiocyanate present is converted to cyanide, the concentration of copper, hydrogen peroxide and sulphuric acid in the experiment were 2+ [Cu ] = 0.6 g/l, [H2O2] = 3.6 g/l and [H2SO4] = 100 g/l. For all the other experiments cyanide recovery was greater than 80%. Since the air flowrate was set at 1.2 L/min for each experiment the only factor that changes the overall cyanide recovery is

182 related to the operating conditions of each experiment. It can be seen in Figure 6.10 that for experiment 6 that the thiocyanate was all oxidised within the first 40 minutes, therefore increasing the initial concentration of cyanide available to be stripped from solution from an early stage in the experiment. It was found it took over 24 hours to be able to strip the cyanide formed during the oxidation of thiocyanate. Experiment 6 shows that even if the cyanide was left in an oxidative acidic environment (where the sulphuric acid concentration was 100g/l) that cyanide would not continue to oxidise to cyanate. The recovery of cyanide using the AVR circuit occurred over a 24 hour period because of slow mass transfer from solution to the stripping air. In this time hydrogen cyanide was left in contact with residual thiocyanate and oxidants, which can react with hydrogen peroxide to produce cyanate, carbonates and nitrates or their hydrolysis products. This can lead to low in cyanide recovery for the other experiments. At the end of these experiments cyanate was determined ranging in concentration from 10-15 ppm in solution. Due to the hydrolysis of cyanate in acidic condition, the concentration of cyanate will change rapidly therefore it can not be used to close the mass balance.

120

100

60 Experiment 1 Experiment 2 Experiment 3 40 Experiment 4 Experiment 5 Cyanide recovered (%) Experiment 6 20 Experiment 7 Experiment 8

0 0 200 400 600 800 1000 1200 1400 1600 Time (min)

Figure 6.12: A comparison of cyanide recovered against time under different experimental conditions outlined in Table 6.2.

183 The 3rd order reaction kinetics based on the derived rate expression defined in section 6.4.3 was used to describe the rate of reaction between hydrogen peroxide. The kinetic model assumes that the rate of reaction is 1st order with respect to the thiocyanate concentration and 2nd order with respect to the hydrogen peroxide concentration. For the experiments that did not contain any copper the validation of the mechanism shown by equation 6.42 was determined by plotting the left hand side of the equation 6.42 against time showing a linear relationship (shown in Figure 6.13 and Figure 6.14). As the acid concentration was increased the rate reaction increased resulting in a larger reaction constant k, shown in table 6.3. The effect of increasing the acid concentration from a pH of 7 to 10 g/l sulphuric acid the reaction constant k increased an order of magnitude. Under these reaction conditions the reaction mechanism did not show the mild inhibition by hydrogen cyanide as observed by other authors reported by Wilson and Harris (1960(A)). This is a direct result of removing hydrogen cyanide from the bubble column.

The experiment shown in Figure 6.15 (with an initial pH of 12) shows that under these conditions the reaction does not follow the third order reaction as with the other experiments carried out at lower pH. This was the only experiment were the pH level changed significantly over the duration of the experiment. Therefore the pH level has to remain constant for the pseudo third order expression described by equation 6.42 to hold true.

When applying the 3rd order reaction kinetic model defined by equation 6.42 to the experiments that contained copper it was found that these experiments did not show a linear relationship seen in Figure 6.16. This may be due to the reactions between copper (II) ions and hydrogen peroxide to produce free radicals which will react with thiocyanate changing the rate limiting step within the reaction, therefore affecting the overall 3rd reaction mechanism described by equation 4.47.

184

Table 6.3: 3rd order reaction constants for a reaction between thiocyanate and hydrogen peroxide. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution. Acid concentration k (min-1.mol-2) (10 g/l Sulphuric acid) 15.8 pH 2 0.32 pH 7 0.19

300

250

200

150

100 Order Kinetic Model rd 3

50 Initial pH 2 Initial pH 7

0 0 100 200 300 400 500 600 Time (min)

Figure 6.13: Applying the third order kinetic model for the reaction between thiocyanate and hydrogen peroxide with an initial pH of 2 or 7. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution.

185 12000

10000

8000

6000

4000 order kinetic model rd 3

2000

0 0 100 200 300 400 500 600 Time (min)

Figure 6.14: Applying the third order kinetic model for the reaction between thiocyanate and hydrogen peroxide with a acid concentration of 10g/l sulphuric acid. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution.

250

200

150

100 order kinetic model rd 3 50

0 0 100 200 300 400 500 600 Time (min)

Figure 6.15: Applying the third order kinetic model for the reaction between thiocyanate and hydrogen peroxide with an initial pH of 12. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution.

186 1200

1000

800

600

400 3rd oder kinetic model Experiment 5 200 Experiment 6 Experiment 8

0 0 100 200 300 400 500 600 time (min)

Figure 6.16: Applying the third order kinetic model for the reaction between thiocyanate and hydrogen peroxide in the presence of copper(II). Experimental conditions outlined in Table 6.2.

The same data shown in Figures 6.13 and 6.14 was used in examination if the reaction rate was either first or second order with respect of either hydrogen peroxide or thiocyanate concentration. Figure 6.17 shows the plot to determine if the reaction rate is first order described by equation 6.31 at initial pH of 7. From this figure the plot is not linear with respect to either hydrogen peroxide or thiocyanate concentration and therefore is not a first order reaction. A similar result is seen for second order reaction described by equation 6.34 and shown by Figure 6.18. It can be seen that the plot is not linear with respect to either hydrogen peroxide or thiocyanate concentration and therefore the reaction is not a second order reaction with respect to either reactant.

187 0.6 0.3 Thiocyanate

Hydrogen Peroxide 0.5 0.25 )

) 0.4 0.2 o o ] 2 O 2

0.3 0.15 ]/[H 2 O 2 ln([SCN]/[SCN] 0.2 0.1 ln([H

0.1 0.05

0 0 0 100 200 300 400 500 Time (min)

Figure 6.17: Applying the first order kinetic model for the reaction between thiocyanate and hydrogen peroxide with an initial pH of 7. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution.

45 2.5 Thiocyanate 40 Hydrogen Peroxide 2 35 o o

30 ] 2 O

1.5 2 25 ] - 1/[H 20 2 O 1 2 1/[SCN] - 15 1/[H

10 0.5

0 0 0 100 200 300 400 500 600 Time (min) Figure 6.18: Applying the second order kinetic model for the reaction between thiocyanate and hydrogen peroxide with an initial pH of 7. The initial concentration of thiocyanate is 1000 ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution.

188 6.7 Conclusions

The oxidation of thiocyanate using hydrogen peroxide with or without the presence of copper under different acid concentrations was studied to see if cyanide could be recovered. By using an AVR circuit cyanide was stripped from solution in a bubble column by air at a constant airflow (1.2 L/min). The air that was used to strip the cyanide from solution was bubbled through a 0.5M caustic solution to trap the cyanide in solution. An advantage of using the AVR circuit to recover cyanide from solution is that it removes cyanide from solution in a form that can be added easily back into a gold mine cyanidation circuit. Process conditions can be varied to concentrate the cyanide recovered from the process.

Typical cyanide recovery from the oxidation of thiocyanate was greater than 80% over a 24 hour period of the theoretical cyanide produced if all the thiocyanate present is converted to cyanide. Figure 6.12 shows that experiment 6 gave the conditions for 100% cyanide recovery from thiocyanate. The conditions of this experiment with an initial thiocyanate concentration of 1000 ppm SCN- were 0.6 g/l copper (II), 3.6 g/L initial hydrogen peroxide, an acid concentration of 100 g/l sulphuric acid and an air flowrate of 1.2 L/min. The mass transfer to remove hydrogen cyanide from the liquid was very slow, taking over 24 hours to recover the cyanide produced in the reaction. By ensuring that high acid concentration further oxidation of cyanide can be minimised. The mass transfer could be increased if stripping towers were used to increase the liquid/air contact area as with commercial projects. Increasing the airflow rate in the bubble column caused slugging within the column reducing the total air surface area reducing the mass transfer between the two phases.

The most significant factor for increasing the rate of reaction and recovery of cyanide is acid concentration. As the concentration of acid was increased, the rate of oxidation increased, by changing the pH from 7 to an acid concentration of 10g/l sulphuric acid the reaction rate constant from 0.19 to 15.8 min-1.mol-2, respectively. If the pH of the reaction is carried out under conditions that the acid concentration does not change significantly the reaction rate order is independent of acid concentration. Other than increasing the rate of reaction, the acid concentration plays an important

189 role in protonating cyanide to form hydrogen cyanide, which aids in stripping the cyanide from solution and it has been found that hydrogen cyanide is not easily oxidised. Therefore high acid concentrations can minimises further oxidation of the cyanide until it is recovered from the bubble column.

The reaction between thiocyanate and hydrogen peroxide is an overall third order reaction, 1st order with respect to thiocyanate concentration and 2nd order with respect to hydrogen peroxide. Addition of copper to the system did not follow the same reaction kinetic mechanism as with experiments that were carried out without copper. This could be due to the formation of free radicals during the reaction between copper and hydrogen peroxide can oxidize thiocyanate. The rate limiting step described by the third order reaction described by equation 6.42 may not be the rate limiting step therefore deviating from the third order reaction that describes the reaction between thiocyanate and hydrogen peroxide without copper. Also addition of copper to the experiment complicates the reaction mechanism kinetics. In the presence of hydrogen peroxide, copper (II) ions will reduce to copper (I) as a part of the decomposition of hydrogen peroxide. Stable complexes can form between copper (I) ions and thiocyanate. Kabesova et al (1978) noticed the formation of copper(I) thiocyanate complexes, weakens the sulphur-carbon bond and strengthens the triple bond between the carbon-nitrogen bond making oxidation easier. As a result for the experiment in the presence of copper the residual thiocyanate concentration after 24 hours is a lot lower than experiments without copper. The lowest residual thiocyanate concentration was 0.3 ppm. Increasing the copper concentration from 0.6 g/l to 1.2g/l did not significantly change the rate of thiocyanate being consumed.

190 References

Byerley J., Enns K. “ Electrochemical regeneration of cyanide from waste thiocyanate for cyanidation.” CIM Bulletin, Vol. 77, No. 861, 1984. pp 87-93.

Choudhary V., Gaikwad, A. “Kinetics of hydrogen peroxide decomposition in aqueous sulfuric acid over palladium/carbon: effect of acid concentration.” Reaction Kinetics and Catalysis Letters. (2003), Vol. 80(1), pp27-32.

Figlar J. and Stanbury D. “Thiocyanogen as an intermediate in the oxidation of thiocyanate by hydrogen peroxide in acidic solution.” Inorganic Chemistry, (2000), Vol. 39, pp 5089-5094.

Fortnum D., Battagila C., Cohn S. and Edwards J. “Kinetics of the oxidation of halide ions by monosubstituted peroxides.” Journal of the American Chemical Society, (1960), Vol. 82, pp778.

Jara J., Soto H., Nava F. “Regeneration of cyanide by oxidation of thiocyanate.” United States Patent 5,482,694, January 9, 1996.

Kabesova M., Kohout J., Gazo J. “Effect of the bridging mode of the thiocyanate group in compounds on infrared spectra.” Inorganica Chimica Acta. (1978) L435.

Kabesova M., Gazo J. “Structure and classification of thiocyanates and the mutual influence of their ligands.” Chemicke Zvesti. (1980), Vol. 34, No. 6, pp 800-841.

Kastle J., Smith C. “On the oxidation of Sulphocanic acid and its salts by hydrogen peroxide.” Journal of American Chemistry. (1904), Vol. 32, pp 376-385.

191 Luo Y., Orban M., Kustin K. and Epstein I. “Mechanistic study of oscillations and bistability in the Cu(II) catalyzed reaction between H2O2 and KSCN. Journal of American Chemical Society. (1989), Vol. 111, No. 13, pp 4541-4548.

Orban M. “Oscillations and bistability in the Cu(II)-catalyzed reaction between H2O2 and KSCN. Journal of the American Chemical Society. (1986), Vol. 108, pp 6893- 6898.

Orban M., Kurin-Csorgei K., Rabai G and Epstein I. “Mechanistic studies of oscillating copper (II) catalysed oxidation reactions of sulfur compounds.” (2000), Vol 55, pp 267-273.

Soto H., Nava F., Leal J., Jara J. “Regeneration of cyanide by ozone oxidation of thiocyanate in cyanidation tailings.” Minerals Engineering, Vol. 8, No. 3, 1995, pp 273-281.

Ullmann’s Encyclopedia of Industrial Chemistry, 5th Edition. (1989) Publisher: VCH Verlagsgesellschaft mbH.

Wilson I., Harris G., (1960 (A)) “The Oxidation of Thiocyanate Ion by Hydrogen Peroxide. I. The pH-Independent Reaction.” Journal of the American Chemical Society, Vol. 82, 1960, pp 4515-4517.

Wilson I., Harris G. (1960(B)) “The oxidation of thiocyanate by Hydrogen Peroxide, II. The acid catalysed reaction.” Journal of the American Chemical Society, Vol. 83, 1960, pp 286-289.

192 Chapter 7

Electrochemical Studies of Thiocyanate Oxidation

7.1 Introduction

Voltammetry is a versatile technique for studying the mechanism of different reduction-oxidation systems. It enables half-cell studies of the rate of electron transfer against the applied potential. The main concerns of each half-cell reaction are the processes and factors that are associated with the charge transfer across the interfaces between chemical phases. There are two commonly used electrochemical techniques, cyclic stationary electrode voltammetry where the working electrode is stationary and hydrodynamic voltammetry where the working electrode is in motion with respect to the solution.

This chapter examines electrochemical aspects of thiocyanate oxidation using a voltametric technique based on the Rotating Disk Electrode (RDE) system. The technique involves the oxidation half-cell reaction of thiocyanate in the absence of any chemical oxidants. The effects of various factors, namely, solution pH, scanning rates and disc rotation speeds and addition of copper on the current-potential relationship of the thiocyanate half-cell reaction were evaluated.

193 7.2 Voltammetry

A simple electrode reaction maybe represented as:

⇔+ RneO [7.1] where O and R are the oxidised and reduced forms of a species, respectively.

The applied potential controls the surface concentration of the two forms of a redox couple (equation 7.1) as described by the Nernst equation (equation 7.2) for a reversible couple.

RT ⎛ C S ⎞ EE o' += ln⎜ O ⎟ − E [7.2] ,RO ⎜ S ⎟ ref nF ⎝ CR ⎠ where E is the potential applied to the cell, volts (V), Eo’ is the formal reduction potential of O,R couple, V vs. SHE, n is the number of electrons transferred per mole, S eq/mol, C O,R is the surface concentration of O or R, mol/l, Eref is the half cell potential of the reference electrode, V vs. SHE.

Thus the oxidised species (O) can be reduced to the reduced species (R) by making E S S sufficiently negative (relative to EO,R – Eref) to cause the ratio of C O/C R to be very small. Reversely, the reduced species (R) can be oxidised to the oxidised species (O) S S by making E sufficiently positive to cause the ratio of C O/C R to be very large. In other words, the reducing or oxidising strength of the working electrode is controlled by the applied potential, E.

For equation 7.1, the following cases may be considered. When the electron transfer rate is sufficiently fast that the ratio of the surface concentrations of the oxidised and reduced species follow the Nernst equation throughout the entire potential sweep, the system us then said to be reversible. If the reaction is very slow compared to the potential sweep rate, v, so that the surface concentrations of O and R are no longer

194 maintained at Nernstian values, the system is then said to be irreversible. When the situation lies between a reversible and irreversible system, the system is termed quasi- reversible.

As shown in Figure 7.1, scanning the potential in the negative direction makes the electrode a stronger reductant (causing O to be reduced to R), whereas scanning the potential in the positive direction makes it a better oxidant (causing R to be oxidised to O). The conversion of O to R by reduction at the electrode surface results in cathodic current, Ic. Oxidation of R to O gives an anodic current, Ia.

ic, cathodic current

REDUCTION

As the potential is increased in a The current plotted above the line negative direction, the electrode indicates a negative charge flow becomes a stronger “reducing FROM the electrode. agent”

E(+), V E(-), V

OXIDATION As the potential is increased in a positive direction, the electrode The current plotted above the line becomes a stronger “oxidising indicates a negative charge flow agent” TO the electrode.

ia, anodic current Figure 7.1: Current-potential axes for voltammetric technquies.

The current-potential information of an electrochemical half-cell reaction under certain condition can be obtained using a conventional three-electrode polarographic system. This system consists of three separate electrodes, namely a reference electrode (RE), a counter electrode (CE) and a working electrode (WE). Figure 7.2 illustrates the typical three electrode system used for electrochemical measurements. The working electrode is where the primary electrochemical reaction occurs and commonly made from platinum, gold, copper, zinc or glassy carbon. The opposite side of the working electrode is referred to as the counter electrode. The electrochemical properties of a counter electrode do not affect the behaviour of the electrode of interest and may be any inert electrode such as a carbon rod dipping

195 directly into the test solution but normally in a separate compartment from the working electrode (Bard and Faulkner 1980). The role of the counter electrode is to complete the electrochemical cell and its potential is rarely of interest. The potential of the working electrode is normally measured with respect to a reference electrode. The standard hydrogen electrode (SHE) is almost universally accepted as a reference electrode, however the saturated calomel electrode (SCE) or the silver-silver chloride electrodes are usually used in experiments and practical work.

Power Supply

Working Electrode Auxiliary Electrode

V Reference Electrode

Figure 7.2: Typical three electrode system for electrochemical measurements (Bard & Faulkner, 1980).

7.2.1 Rotating Disc Electrode

One of the few electrode systems for which the hydrodynamic and convective- diffusion equations have been solved rigorously for the steady state is the rotating disk electrode (RDE). A schematic diagram is represented in Figure 7.3. Theoretical treatment of the rotating disc system was given by Riddiford (1966), for a smooth disc, the fluid remains laminar when the Reynolds number (refined by equation 7.3) remains under 2x105. As the disc rotates, adjacent solution is pulled along by viscous

196 drag and is thrown away from the axis of rotation in a radial direction by the centrifugal force. The expelled solution is replaced by flow normal (axial flow) to the disc surface. Thus, the rotating disc acts as a pump which moves solution up from below and then away from the electrode.

r 2ω Re = [7.3] v where Re is the dimensionless parameter, r is the radial position on the disc (cm), ω is the angular velocity (rad/s) and v is the kinematic viscosity (cm2/s).

Figure 7.3:Schematic diagram of the fluid flow for a rotating disc system (Miller, 1979).

Levich (1962) showed that for a rotating disc, under laminar flow condition the diffusion layer thickness under the disk is given by:

1 −1 1 3 2 6 o =∂ 1.61Do ω v [7.4]

197 where δο is the “diffusion” layer thickness at an electrode fed by convective transfer 2 (cm), Do is the diffusion coefficient of the reactive species (cm /s), ω is the angular frequency of rotation (1/s) and v is the kinematic viscosity of solution (cm2/s).

For a reaction controlled by a mass-transfer limited condition at the RDE, the limiting current can be described by the Levich equation:

2 1 −1 3 2 6 * i l = 0.620nFADo ω Cv o [7.5]

where n is the number of moles of electrons oxidised or reduced, F is Faraday’s 2 constant (96 486 coulomb/mole), A is the surface area of the electrode (cm ), Do is the diffusion coefficient of the reactive species (cm2/s), ω is the angular frequency of 2 * rotation (1/s) and v is the kinematic viscosity of solution (cm /s) and Co is the bulk concentration of the oxidising or reducing ion (mol/l).

* 1/2 The Levich equation predicts that ic is proportional to Co and ω . One can define 1/2 * the Levich constant, ic/ω Co , which is the RDE analog of the diffusion current constant or current function in voltammetry or the transition time constant in chronopotentiometry (Bard and Faulkner 1980).

The rate of reaction at the electrode surface during the potential sweep can be related to the diffusion coefficient of the reactive species, kinematic viscosity of solution and angular frequency of rotation of the RDE (Bard and Faulkner 1980), which is shown below:

2 −1 1 3 6 2 = o v0.620Dk ω [7.6]

Equating equations 7.5 and 7.6 the reaction rate constant, k, can be expressed as:

i k = l [7.7] nFAC

198 7.3 Experimental

7.3.1 Reagents and Materials

All chemicals reagents used in the studies in this chapter were the same as described in Chapter 6. In addition to these chemical reagents, analytical grade anhydrous sodium sulphate (100%) was used in these experiments and supplied by Asia Pacific Speciality Chemicals Limited.

7.3.2 Equipment and Instrumentation

7.3.2.1 Electrodes

The electrochemical studies were conducted using a standard three-electrode system (Figure 7.2), which consists of a working electrode (WE), a reference electrode (RE) and counter electrode.

Working electrode

A glassy carbon circular polished disc was used in the study as the working electrode. The electrode consisted of a 7mm glassy carbon disk embedded into a high-density polyethylene holder with an outer diameter of 12mm. The working surface area of the electrode is 0.385cm2.

Reference electrode

A standard Silver-Silver Chloride electrode was used in this study as the reference electrode. Its cell arrangement can be represented as follows:

ClAgClAg − [7.10]

199 The electrode was filled with saturated potassium chloride solution giving a potential of 0.199V against a standard hydrogen electrode at 25oC. The measured potential (versus the silver-silver chloride electrode) can be converted to the potential of hydrogen at 25oC by the following equation:

h = measured + 199.0EE V [7.11]

Counter Electrode

The counter electrode was made of a graphite rod of 14 cm in length and 6mm in diameter.

7.3.2.2 Equipment

The equipment used for the electrochemical half-cell studies consist of a rotator, speed controller unit, potentiostat and a data acquisition unit. The schematic diagram of the equipment/apparatus set-up for the half cell study is shown in Figure 7.7.4, while its front view arrangement is shown in Figure 7.7.5.

Potentiostat

A cyclic voltammograph (model:CV-27) made by Bioanalytical Systems Incorporated was used for the voltammetry study. This unit has a built-in scanner. The specifications of the unit are as follows:

Table 7.1: Specifications of the cyclic voltammograph Sweep range +5.0 V Scan rate range 0.1 mV/s to 10 V/s Compliance voltage + 10.0 V Maximum available current 120 mA Current converter 2μA/V to 10.0 mA/V

200

1 2 3

5 6 7

Figure 7.4: Schematic of equipment set-up for electrochemical half-cell studies. 1) Rotating disc electrode, 2) Speed controller, 3) Potentiostat, 4) Data Acquisition unit, 5) Counter electrode, 6) Reference electrode, 7) Thermometer, 8) 5 port Pyrex flask, 9) Water bath.

Figure 7.5: Front view of the equipment arrangement for electrochemical studies showing: IBM computer with Quicklog PCTM software, TLC 548/9 Anolog to digital converter, Pine intrument Rotating Disc Electrode rotator with 5 port reaction cell and wath bath, MSR speed controller and CV-27 Potentiostat.

201 Rotator and speed controller

The rotator and speed controller (model: MSR) was made by Pine Instruments. The speed of rotation ranged from 0 to 10 000 rpm. Its accuracy is better than 1% of the dial setting.

Data acquisition unit

To record the digital current-potential data, a data acquisition unit was made, which consisted of a TLC584/9 LinC MOS 8-bit analog to digital converter linked to a IBM compatible computer and Quicklog PCTM data acquisition software version 1.04 (Strawberry Tree Incorporated, 1990).

The pH of the solution was measured by Orion 520A digital pH meter. Temperature of the solution was kept at constant temperature by immersing the reaction cell in a water bath which was controlled by a Haake water chiller.

7.3.3 Experimental Procedure

The working solutions for the half-cell studies were made in a background solution

0.1M sodium sulphate (Na2SO4) to increase the conductivity of the solution. To the background solution the desired amounts of thiocyanate and copper(II) as copper sulphate were added to make the experimental concentrations. Then the solution was adjusted to the desired pH by adding either 1M sulphuric acid or 1M sodium hydroxide solution. To the 250ml five port spherical Pyrex flask 200ml of the working solution was transferred. The central port of the flask contained the rotating working electrode while the other ports contained the reference electrode, counter electrode and thermometer to measure the temperature of the solution.

The Pyrex flask was set into the water bath and was allowed time to equilibrate to the required temperature. The working electrode was thoroughly polished flat and perpendicular to the electrode’s shaft using a series of polishing pads until a mirror finish. The working electrode is secured to the shaft of the rotator and immersed

202 approximately 5 mm below the surface. Care was taken to ensure that the only electrical contact the working electrode made with the solution was the polished glassy carbon disc. Before each scan, the voltammograph was set at the required potential range, appropriate scan rate and conversion factor for the anolog output for data acquisition. The speed of the rotator was set to the required rotation rate and the voltage was scanned over the potential range. During the scan the current-potential relationship of the half-cell reaction was recorded using the data acquisition unit.

203 7.4 Results

The effects of various factors, namely, solution pH, scanning rates, disc rotation speeds and addition of copper on the current-potential relationship of the thiocyanate half-cell reaction were evaluated. All experiments were conducted at 25oC and at a scan rate of 40mV/s unless stated otherwise.

The stationary voltametric response of the glassy carbon working electrode was determined with respect to the Ag-AgCl reference electrode in 0.1M Na2SO4. The cyclic voltametric response obtained for the first cycle at a scan rate of 40mV/s (Figure 7.6), was scanned from 0 to 1.6V with respect to the Ag-AgCl electrode. From this scan, breakdown of water became apparent above 1.0 V (vs. Ag-AgCl electrode). The increase of current above 1.0 V (vs. Ag-AgCl electrode) was due to - the oxidation of OH ions in solution to form oxygen gas (O2 (g)).

A) 40 μ

20 Current i ( Direction of scan → 10

-10 0 0.2 0.4 0.6 0.8 1 1.2 1.4 1.6 1.8 E (V vs. Ag-AgCl reference electrode)

Figure 7.6: Cyclic voltammogram of the glassy carbon electrode in 0.1M Na2SO4 at pH 2, rest potential 0V and a scan rate of 40mV/s at a temperature of 25oC.

204 In aqueous solution the overall electrolysis reaction of thiocyanate predominantly follows equation 7.10 and is believed to proceed through the formation of thiocyanogen (S(CN)2) which is subsequently hydrolysed to sulphate and cyanide and the electrode process is irreversible (Nicholson, 1959).

− 2− −+ 2 SOO4HSCN 4 +→+ HCN(aq) ++ 6e7H [7.10]

Wilson et al. (1960) proposed that the oxidation of thiocyanate occurred via two different mechanisms depending on the pH of the system to form HOSCN, shown previously by equation 6.2 when the pH >4 or equation 6.3 when the pH <2. HOSCN will further oxidise to form S(CN)2 as described by Nicholson.

Figure 7.7 shows a stationary cyclic voltammogram of a glassy carbon electrode in a solution of 0,1M Na2SO4 + 0.01m NaSCN, pH 2 at a different scan rates, the peak current occurred at E = 1.35V (vs. Ag-AgCl electrode). From the reverse scan from 1.5V back to 0V there are no current peaks produced showing that the electrode process is irreversible. This is in agreement with previous work by Nicholson (1959) where voltammetry of thiocyanate ion in pyridine-pyridinium solutions were studied. The scan rate was increased from 10 mV/s to 40 mV/s shown in Figure 7.7. During the forward scan from 0V to 1.5V, two peaks that be seen. The first peak occur when the potential reached 1.18V (vs Ag-AgCl reference electrode), with the second peak occurring at 1.36V (vs. Ag-AgCl reference electrode) as shown in Figure 7.7. The first peak corresponds to the oxidation of thiocyanate to cyanide; the second peak corresponds to the further oxidation of cyanide generated at the electrode surface when the thiocyanate was oxidised. Since Equation 7.10 may represent the sum of perhaps seven or eight steps, some of them are slow, and deviations from stoichiometry are easily visualised, the oxidation with species such as (SCN)2 or SCN+ can produce broad peaks seen in Figure 7.7 instead of clearly defined peaks that can occur with a simple oxidation, for example the oxidation of ferrous to ferric ions.

The pH of the solution was increased to 12 and the cyclic voltammogram of a glassy carbon electrode in a solution of 0,1M Na2SO4 + 0.01m NaSCN, at a different scan rates were repeated. The voltammograms are shown in Figure 7.8. The shape and

205 position did not change with a change in pH. The peak currents versus scan rate for this set of experiments are tabulated in Table 7.3. At lower scan rates it can be seen that the rate the current drop is not as pronounced in the reverse scan at higher voltage scanning rates. This may suggest that at low scan rates the reaction is autocatalytic, but this effect may be due to the oxidation of cyanate produced at the electrode surface in the forward scan and not from species present in the bulk solution.

Theory predicts direct proportionality between the maximum current and the square root of the voltage scanning rate for reversible and relatively simple irreversible processes (Delahay, 1953), meaning that the reaction occurring at the surface of the electrode is limited to the diffusion of the reactant to the surface of the electrode. In the other case were the maximum current is not directly proportional to the square root of the voltage scanning rate, the reaction rate at the electrode surface is not limited by diffusion. Figure 7.9 shows a plot of maximum current versus the square root of the voltage scanning rate for the first peak in the cyclic voltammogram shown in figures 7.7 and 7.8. From the linear relationships seen in Figure 7.9 the oxidation process at either pH is diffusion controlled, we can see that the diffusion rate constant is higher at lower pH because the gradient of the scanning rate versus maximum current is greater at pH 2 then when the pH was increased to 12. A simular plot was generated for the second peak seen in Figure 7.7, shown by Figure 7.10, it shows that the oxidation process is diffusion controlled and the diffusion rate constant is same when the experiment is carried out at either pH 2 or 12.

206

1400 Peak 1 Peak 2

1200 D C 1000

B 800 A) μ

i ( 600

A 400 Forward scan

200 Reverse scan

0 0 0.2 0.4 0.6 0.8 1 1.2 1.4 1.6 E (V) vs Ag-AgCl

Figure 7.7: Cyclic voltammogram of a glassy carbon in a solution of 0.1M Na2SO4+ 0.01M NaSCN, rest potential=0V (vs. Ag-AgCl electrode), pH 2, temperature 25oC. The scan rate was varied from 10 to 40 mV/s. (A) 10 mV/s, (B) 20 mV/s, (C) 30 mV/s, (D) 40mV/s.

Table 7.2: Peak Currents for different scan rates for figure 7.7. Scan rate (mV/s) Peak 1 Current (μA) Peak 2 Current (μA) 10 349 378 20 749 965 30 732 1165 40 965 1224

207 1600

1400 Peak 1 Peak 2

1200 D C 1000 B A)

μ 800 A i (

600

400 Forward scan

200 Reverse scan

0 00.511.52 E (V) vs. Ag-AgCl

Figure 7.8: Stationary cyclic voltammogram of a glassy carbon in a solution of 0.1M Na2SO4, 0.01M NaSCN, rest potential=0V (vs. Ag-AgCl electrode), pH 12, temperature 25oC. The scan rate was varied from 10 to 40 mV/s. (A) 10 mV/s, (B) 20 mV/s, (C) 30 mV/s, (D) 40mV/s.

Table 7.3: Peak Currents for different scan rates for figure 7.8. Scan rate (mV/s) Peak 1 Current (μA) Peak 2 Current (μA) 10 - 667 20 936 880 30 1174 1000 40 1324 1267

208 1600

1400 pH 2 pH12

1200 A)

μ 1000

800

600 Peak current ( 400

200

0 01234567 Square root of scan rate ((mV0.5/s0.5)

Figure 7.9: Peak currents versus square root of scan rate for the first peak in the voltammograms shown in Figure 7.7 (pH 2) and Figure 7.8 (pH 12).

1400

pH 2 1200 pH 12

1000 A) μ 800

600

Peak current ( 400

200

0 01234567 Square root of scan rate ((mV0.5/s0.5)

Figure 7.10: Peak currents versus square root of scan rate for the second peak in the voltammograms shown in Figure 7.7 (pH 2) and Figure 7.8 (pH 12).

209

The rotating disc electrode was used to determine the oxidation kinetics of thiocyanate over the surface of the working electrode according to the Levich equation. For a given set of mass transfer conditions in each experiment, such as rotation speed of the disc electrode and solution variables (i.e. temperature, concentration of interested species and viscosity of solution) the rate reaction of thiocyanate is limited by how fast the thiocyanate ion can be bought to the electrode surface. Therefore the current flowing through the circuit becomes limited. Figures 7.11 and 7.12 shows the rotating disc voltammograms of a glassy carbon working electrode which was rotated in a solution of 0.1M Na2SO4 + 0.01M NaSCN with a solution pH of 12 or 2, respectively. The rotation speed of the working electrode was varied from 50 rpm to 500 rpm.

From Figure 7.12 it can be seen that by operating in different pH levels, that is, pH 12 and pH 2 that the mechanism of oxidation of thiocyanate changes with a change in pH, the wave formation changes when the chemical potential reaches 1.1V (versus Ag-AgCl standard electrode). This is in agreement that was proposed by Wilson et al. (1960) that the mechanism of oxidation changes with the pH of the system. Therefore when scanning at low pH < 2 will have a different wave formation when comparing it to scans at high pH > 4. When looking at the overall reaction scheme, thiocyanate will oxidise through intermediates HOSCN and HOOSCN (Wilson et al., 1960) to form either cyanide or cyanate. Also the formation of cyanide will oxidise to produce cyanate, therefore the oxidation process of thiocyanate will form intermediate species such as HOSCN, HOOSCN, CN- which then will oxidise form CNO-. In both low and high pH scans the limiting current occurred when the chemical potential reached 1.5V (vs Ag-AgCl electrode) showing that the formation of cyanate is the final oxidation product from thiocyanate.

From Figure 7.11 the limiting current was determined for each rotating speed for the overall reaction of thiocyanate to cyanate, which is shown in Table 7.4. From theory defined by the Levich equation, the limiting current (iL) is proportional to the square root of the rotation frequency of the working electrode (ω). A plot of the limiting current versus the square root of the electrode rotation frequency shown in Figure

210 7.13 exists a linear relationship. Therefore the Levich equation is obeyed for this system. Using equation 7.7, the rate of reaction was calculated, shown in Table 7.4. The rate ranged from k = 0.70 to 2.54 cm/s when the electrode rotation frequency was increased from 50 to 500 rpm, respectively.

7000

D 6000

5000

C 4000 A) μ

i ( 3000 B Direction of scan →

2000 A

1000

0 0.50.70.91.11.31.51.71.9 E (V) vs Ag-AgCl

Figure 7.11: Rotation disc voltammograms of the anodic half-cell study of thiocyanate. The o solution was made from 0.1M Na2SO4+0.01M NaSCN, pH 12, temperature 25 C and a rest potential of 0.5V vs Ag-AgCl reference. The rotation frequency was varied from 50rpm to 500 rpm. (A) 50rpm, (B) 100rpm, (C) 200rpm and (D) 500rpm.

211 3500

pH 2 (B) 3000

pH 12 (B) 2500 pH 2 (A) A)

μ 2000

pH 12 (A) 1500 Current (

1000 Direction of scan → 500

0 00.511.522.5 E (V) vs. Ag-AgCl

Figure 7. 12: Rotation disc voltammograms of the anodic half-cell study of thiocyanate. The o solution contained 0.1M Na2SO4+0.01M NaSCN, pH 2 and 12, temperature 25 C and a rest potential of 0.5V vs Ag-AgCl reference. The rotation frequency was varied from 50rpm and 100 rpm at each pH. (A) 50rpm and (B) 100rpm.

7000

6000

5000 A) μ

4000

3000

limiting current ( limiting 2000

1000

0 012345678 ω0.5 (s-0.5)

Figure 7.13: A plot of limiting current versus the square root electrode rotation frequency for experiments conducted in Figure 7.11.

212 When the solution of the RDE experiment was repeated at pH of 2 with a solution temperature was 25oC, shown by Figure 7.12, it can be seen that by reducing the pH of 12 to 2, the limiting current is slightly higher at the lower pH. At a pH of 2 the limiting current was 1942.86 μA and 2828.57 μA with an electrode rotation speed of 50 and 100 rpm, respectively. This corresponds to a k of 0.87 and 1.27 cm/s. Comparing the rate constant k to the experiments that were conducted at a pH of 12, there was a 20% increase at a rotation speed of 50 rpm, and a 6% increase when the rotation speed was 100 rpm.

Addition of copper to the stationary voltammogram of a glassy carbon electrode in a o base solution of 0.1M Na2SO4 and 0.01M NaSCN, pH 2, temperature 25 C and a voltage scan rate of 40mV/s is shown in Figure 7.14. It can be seen that by the addition of copper the reaction that occurred peak 1 (at E = 1.18 V vs. Ag-AgCl electrode) is suppressed while the oxidation products produced at peak 1 appears in its usual position. This may be due to copper stabilising thiocyanate within solution, as discussed in Chapter 6, but the addition copper does not change the overall rate of reaction of thiocyanate to cyanate because the peak current is unchanged with or without the presence of copper.

Table 7.4: Rotating frequency of the working electrode and the related limiting currents for the experiments shown in Figure 7.11.

Rotation speed ω Limiting current (iL) K (rpm) (s-1) (μA) (cm/s) 50 5.24 1556 0.70 100 10.47 2667 1.20 200 20.94 3722 1.67 500 52.36 5667 2.54

213 1400 Peak 1 Peak 2

1200

1000

800 A) μ

i ( 600

400 Direction of scan →

200

0 0 0.2 0.4 0.6 0.8 1 1.2 1.4 1.6 E (V) vs. Ag-AgCl

Figure 7.14: A comparison of the stationary cyclic voltammogram of a glassy carbon in a solution of 0.1M Na2SO4, 0.01M NaSCN with copper added to one experiment (concentration of o 0.005M CuSO4), rest potential=0.5V (vs. Ag-AgCl electrode), pH 12, temperature 25 C and a voltage scanning rate of 40mV/s.

Figure 7.15 shows a comparison of rotating disc voltammograms between systems that contain with or without copper. As with the stationary cyclic voltammograms (Figure 7.14) when copper was added to the system that the addition of copper to the system reduces the oxidation of thiocyanate as there is a reduction in limiting currents of the RDE experiment at both rotation speeds of 50 and 100rpm in solution with a pH of 2. Therefore addition of copper (II) ions to the system stabilizes thiocyanate ions by the formation of stable copper thiocyanate species as discussed in previous chapters 4 and 6, as a result reducing overall oxidation of thiocyanate in the experiments that contained copper.

214 3500

3000 No Copper B

2500

A 2000 A) μ

i ( 1500 Copper Direction of scan → 1000

500

0 00.511.522.5 E (V) vs. Ag-AgCl electrode

Figure 7. 15: Rotation disc voltammograms of the anodic half-cell study of thiocyanate. The solution contained 0.1M Na2SO4+0.01M NaSCN with one set of experiments containing 0.005M o CuSO4, pH 2, temperature 25 C and a rest potential of 0.5V vs Ag-AgCl reference. The rotation frequency was varied from 50rpm to 100 rpm. (A) 50rpm, (B) 100rpm

215 7.5 Conclusions

The electrochemical studies of thiocyanate under various conditions showed its electrochemical behaviour. Thiocyanate oxidation is not simple and the overall reaction to convert thiocyanate to cyanide may occur over several steps, and if some of the individual steps as slow, deviations from this stoichiometry can be easily visualised.

From peak current and voltage scan rate relationship of the thiocyanate oxidation process suggests that the reaction is diffusion controlled when the pH of the solution was either 2 or 12. Under acidic conditions the oxidation of thiocyanate is not a relatively simple irreversible process.

By rotating disk electrode studies, the rate of reaction at the surface of the working electrode can be determined by applying the Levich equation. The Levich equation relates the limiting current to the square root of the frequency of rotation of the working electrode. For the different thiocyanate systems studied the Levich equation was obeyed. The kinetic rate constants were determined for the overall oxidation of thiocyanate to cyanate in the presence of 0.1M sodium sulphate at pH 12 ranged from 0.70 to 2.54 cm/s.

Addition of copper sulphate to the system did not effect the oxidation of thiocyanate at E = 1.35V (vs. Ag-AgCl reference electrode). But it did suppress competing reactions at a lower potential. Therefore copper(II) ions do not accelerate the oxidation of thiocyanate. Within the thiocyanate oxidation experiments using hydrogen peroxide, copper(II) ions can accelerate the decomposition of hydrogen peroxide. As one of the decomposition intermediate steps copper(II) ions are reduced to copper(I). Copper(I) ions readily complex with free thiocyanate ions which can reduce the activation energy required for the oxidation experiment with hydrogen peroxide. Since this step was missing in the RDE experiment, this phenomenon could not be observed.

216 References

Bard, A. and Faulkner, L.R., “Electrochemical Methods fundamentals and Applications.” John Wiley & Sons, New York. 1980

Delahay P., “New Instrumental Methods in Electrochemistry.” Interscience, New York. 1954.

Kabesova M., Kohout J., Gazo J. “Effect of the bridging mode of the thiocyanate group in compounds on infrared spectra.” Inorganica Chimica Acta. (1978) L435.

Kabesova M., Gazo J. “Structure and classification of thiocyanates and the mutual influence of their ligands.” Chemicke Zvesti. (1980), Vol. 34, No. 6, pp 800-841.

Levich V.G., “Physicochemical Hydrodynamics.” Prentice Hall, Englewood Cliffd, New Jersey. 1962.

Miller J.D., “An analysis of concentration and temperature effects in cementation reactions.” Mineral Science Engineering, 1979, vol. 5, no. 3, pp 242-254.

Nicholson M.M., “Voltammetry of thiocyanate ion at the stationary platinum electrode.” Analytical Chemistry, 1959, vol. 31, no. 1, pp 128-132.

Riddiford A.C., “The rotating disk system.” Advances in electrochemistry and electrochemical engineering, 1966, vol. 4, pp 47-116.

Wilson I. and Harris G., “The oxidation of Thiocyanate Ion by Hydrogen Peroxide. I. The pH-Independent Reaction.” Journal of the American Chemical Society, Vol. 82, 1960, pp 4515-4517.

217 Chapter 8

Conclusions and Recommendations

8.1 Conclusions

The first seven chapters of this thesis deal with an analysis and discussion of the results from experiments conducted on the stability of different thiocyanate-water and copper-thiocyanate-water systems, binary ion-exchange equilibria between thiocyanate and chloride using commercially available Purolite A500u/2788, recovery of cyanide by air stripping from the oxidation of thiocyanate by hydrogen peroxide, electrochemical oxidation half cell reaction of thiocyanate and a critical review and comparisons of the related literature. A summary of the conclusions that may be drawn from the present work is as follows:

8.1.1 Solution Equilibria for thiocyanate-water and copper-thiocyanate-water systems

Pourbaix and distribution diagrams of thiocyanate-water and copper-thiocyanate- water systems were constructed using STABCAL (version 32-STABCAL SN 8-19- 2001). From the Pourbaix diagrams developed for the thiocyanate-water system that the predominant stable species formed in solution at equilibrium is cyanate or hydrogen cyanate depending on the pH of the solution. Under extreme acidic conditions thiocyanate is the predominant stable species over hydrogen cyanate at neutral chemical potential or a slightly reducing environment. When cyanate and hydrogen cyanate were removed as stable species from the simulation, the resulting Pourbaix diagram showed that the predominant areas of cyanate and hydrogen

218 cyanate were replaced with cyanide and hydrogen cyanide depending on the pH of the system. Therefore cyanide can be recovered from thiocyanate if the oxidation of cyanide to cyanate can be stopped.

Addition of copper to the thiocyanate-water system makes the system more complex. The predominant stable species in a copper-thiocyanate-water system is copper (I) thiocyanate which is a solid therefore reducing the mobility of thiocyanate in the environment. As the ratio of thiocyanate to copper was increased the soluble copper thiocyanate complexes which contained a higher ratio of thiocyanate became + 3- predominant species at low pH as either CuSCN , Cu(SCN)2 and Cu(SCN)4 .

The copper-thiocyanate-water distribution diagrams showed that the concentration of the different copper thiocyanate species is dependent on the concentration of thiocyanate and copper, as well as pH and Eh of the solution. The majority of the copper and thiocyanate existed as copper(I) thiocyanate which is a solid. As the concentration of thiocyanate increased compared to copper the concentration of copper thiocyanate complexes that had a higher ratio of thiocyanate to copper 3- increased, i.e. Cu(SCN)2 and Cu(SCN)4 . When more copper was present in the system by the addition of copper (II) ions, the concentration of free thiocyanate and hydrogen thiocyanate decreased.

8.1.2 Ion-exchange equilibria between thiocyanate and chloride

Using ion exchange resins can be an effective way to concentrate thiocyanate from tailing solutions or slurries. The effective capacity of thiocyanate on the ion- exchange resin Purolite A500u/2788 is 1.19 mmol NaSCN/ml resin at 303 K. Four different models were used to describe the equilibria between solution concentration and resin loading of thiocyanate, which were the Langmuir isotherm, the Freundlich isotherm, the Dubnin-Radushkevich isotherm and the Mass-action Law model. The first three models showed that the equilibria between solution concentration and resin loading are solution concentration dependent. Under the Mass-action law model the ion-exchange equilibria are represented as equivalent ionic fraction of ions in solution and on the ion exchanger. By normalising the ion-exchange equilibria it was found

219 that the fraction of ions in solution and on the ion exchanger becomes independent of ionic strength (i.e. solution concentration) within the range of concentration considered. An advantage of this is that only one unique equilibrium constant to describe the ion-exchange process is used while other isotherms the relationships show dependence on the concentration of the solution.

6.1.3 Recovery of cyanide by air stripping from the oxidation of thiocyanate by hydrogen peroxide

The oxidation of thiocyanate by hydrogen peroxide to recover cyanide is an effective way of converting waste thiocyanate to free cyanide which can be recycled in a gold operation. This can be an important consideration for mine rehabilitation as stricter environmental regulations are imposed to limit the discharge of thiocyanate to tailing dams. This could be used as a cost effective way to minimise and recycle cyanide in a mining operation. The pH of the solution considerably affected the rate of oxidation of thiocyanate, showing faster kinetics at lower pH’s levels. This confirmed observations of earlier studies. It is also evident that the pH also affects the rate at which hydrogen cyanide can be removed from the bubble column, taking over 24 hours to remove the hydrogen cyanide. With prolonged time taken to recover cyanide from the bubble column continued oxidation of cyanide to cyanate is very slow resulting in high recoveries in many experiments of over 80% in many cases.

The overall rate of reaction of thiocyanate with respect to the concentration of thiocyanate and hydrogen peroxide is an overall third order reaction. The derived third order rate expression is first order with respect to the concentration of thiocyanate and second order with respect to the concentration of hydrogen peroxide, shown below:

H ]Od[ − 22 = 2 − ][SCN]Ok[H [8.12] dt 22

- where t is time, [SCN ] is the thiocyanate concentration and [H2O2] is the hydrogen peroxide concentration.

220 Compared to previous studies it was found that the production of cyanide inhibits the reaction between thiocyanate and hydrogen peroxide, which was taken into account when defining the rate expression. Since cyanide was removed from the reaction by air stripping, this study did not show the same effect that cyanide had in the previous study.

It was proposed initially that addition of copper (II) ions would accelerate the oxidation of thiocyanate. Previous studies found that copper(I) thiocyanate complexes weaken the sulphur-carbon bond and strengthen the triple bond between the carbon-nitrogen bond making oxidation easier. This was not the case as the addition of copper(II) ions did not dramatically increase the rate of oxidation of thiocyanate. However the residual thiocyanate concentration after a 24 hour period was lower than experiments that were conducted without copper.

6.1.4 Electrochemical studies of thiocyanate oxidation

The anodic voltammetry of thiocyanate ion in aqueous solutions at varying pH has been studied at the stationary and rotating glassy-carbon electrode with linear variation of the applied potential. The oxidation of thiocyanate is not simple and the overall reaction to convert thiocyanate to cyanide may occur over several steps resulting in a voltammogram being a result of all of the various reactions occurring. The reaction at the electrode surface was found to be an irreversible one. From peak current and voltage scan rate relationship of the thiocyanate oxidation process suggest that the reaction is diffusion controlled when the pH was 12. As the pH was reduced to 2, the thiocyanate current deviates from this relationship, showing that under acidic conditions the oxidation of thiocyanate is not a relatively simple irreversible process. With the addition of copper(II) ions as copper sulphate it was found that copper (II) ions suppress competing reactions that occurred at a lower potential without effecting the primary oxidation of thiocyanate to cyanide. It was theorised that the formation of by-products other than cyanide was caused by oxidation of these competing reactions. If copper can suppress these reactions occurring without reducing the rate of reaction higher rates of cyanide recovery can be found. It was found that the Levich equation was obeyed for the rotating disc electrode experiments conducted. The kinetic rate

221 constants were determined for the oxidation of thiocyanate in the presence of 0.1M sodium sulphate at pH 12 and 2 ranged from 0.70 to 2.54 cm/s, respectively.

8.2 Recommendations for further work

Due to the limited scope of this project only the overall reaction rate between thiocyanate and hydrogen peroxide has been studied. The formation of intermediate products formed during the reaction between thiocyanate and hydrogen peroxide and how these products relate to the overall reaction have not been covered and studied. Therefore, the following problems could be the objectives for further study and clarification:

1. It has been theorised that the overall oxidation of thiocyanate to cyanide occurs through the formation of various intermediate products. However, detailed studies on this particular problem have not been established.

2. The effect of various additives such as iron and including copper on the rate and extent of thiocyanate oxidation. These additives may suppress the competing reactions that lead to products other than cyanide and the mechanism on how these additives effect the overall reaction between thiocyanate and hydrogen peroxide can be studied.

3. Using more effective methods to strip cyanide from solution by using AVR stripping towers that are used commercially to remove cyanide from acidic solutions rather than a bubble column used in this study.

4. By examining different binary ion-exchange systems with other species such as base and precious metal cyanides using the Mass-Action Law, a comparison of the relative affinity of different counter ions can be determined by using the cross product rule. This can be valued in the prediction of the loading characteristics of more complex systems found in the mining industry.

222 5. Elution studies for the elution of thiocyanate that has been adsorbed onto anion exchange resins. Direct oxidation on ion exchange resins can affect the ion exchange structure making the resin brittle, that is reduce the effective lifespan of the resin. More benign elution process would overcome this problem. Elution techniques in conjunction with the mass action model would complete the recovery of cyanide from thiocyanate circuit within a gold operation.

223 Appendix A: Fortran Code for Marquardt algorithm applied to the Mass-action Law

PROGRAM IEX IMPLICIT DOUBLE PRECISION (A-H,O-Z) EXTERNAL RESID DIMENSION PARM(4),X(3),F(51),XJAC(51,3),XJTJ(6),WORK(123) DIMENSION XB(51),YB(51),SI(51),YBCALC(51)

! Program for fitting KAB and Wilson coefficients from ion exchange ! equilibrium data. Ion 'A' is initially loaded on the resin ! while ion 'B' is the counter ion.

! The program reads the experimental values of XB(I), YB(I) and SI(I) ! from the file specified in the OPEN statement (unit=2).

! The fitted values of KAB and the Wilson coefficients are obtained ! by minimising the SSRD with respect to YB(I).

! The predicted values of the loading of ion B (YBCALC(I)) ! are sent to the file specified in the OPEN statement (unit=3).

! XB refers to the ionic fraction of ion B in the solution. ! YB refers to the ionic fraction of ion B on the resin. ! SI refers to the total ionic strength of the solution (cations + anions)

OPEN(2,FILE='ascn1.dat',FORM='FORMATTED',STATUS='OLD') OPEN(3,FILE='resultk.dat',FORM='FORMATTED',STATUS='NEW')

! Enter the number of data points (NDP) for each variable. NDP = 51 ANDP = NDP

! Read experimental data from .dat file. DO 5 I=1,NDP READ(2,*) XB(I) 5 CONTINUE

DO 10 I=1,NDP READ(2,*) YB(I) 10 CONTINUE

DO 12 I=1,NDP READ(2,*) SI(I) 12 CONTINUE

! Initialise parameters for ZXSSQ ! Adjustable parameters: X(1)=WCAB, X(2)=WCBA, X(3)=KAB M = NDP N = 3 NSIG = 5 EPS = 0.0 DELTA = 0.0 MAXFN = 500 IOPT = 1 IXJAC = NDP X(1) = 1.0D0

224 X(2) = 1.0D0 X(3) = 1.0D0

CALL ZXSSQ (RESID,M,N,NSIG,EPS,DELTA,MAXFN,IOPT,PARM, & X,SSQ,F,XJAC,IXJAC,XJTJ,WORK,INFER,IER)

! If ZXSSQ converges print YBCALC(I) values to .dat file IF (INFER.GT.0) THEN PRINT*, "Parameters regressed successfully" PRINT*, "WCAB= ",X(1)," WCBA= ",X(2)," KAB= ",X(3)

CALL RESID (X,M,N,F)

DO 15 I=1,M YBCALC(I) = YB(I)*(F(I) + 1.0D0) WRITE (3,50) "YBCALC(",I,") , ",YBCALC(I) 15 CONTINUE

SUMARD = 0.0D0 DO 20 I=1,M SUMARD = SUMARD + ABS(F(I)) 20 CONTINUE AARD = 100.0D0*SUMARD/ANDP

WRITE (3,100) "WCAB , ",X(1) WRITE (3,100) "WCBA , ",X(2) WRITE (3,100) "KAB , ",X(3) WRITE (3,100) "AARD(%) , ",AARD ELSE PRINT *, '------' PRINT *, 'INFER = 0, IER= ',IER PRINT *, 'Convergence failed, reduce NSIG' END IF

END FILE (UNIT=2) END FILE (UNIT=3) CLOSE (UNIT=2) CLOSE (UNIT=3)

50 FORMAT (1X, A, I2, A, E12.5) 100 FORMAT (1X, A, E12.5) STOP END

! ------SUBROUTINE RESID (X,M,N,F) IMPLICIT DOUBLE PRECISION (A-H,O-Z) DIMENSION X(N),F(M) DIMENSION XB(M),YB(M),SI(M),PHI(M),YBCALC(M), & GAMAL(M),GAMBL(M),GAMAR(M),GAMBR(M),DENOM(M)

! A,B = debye-huckel parameters for solvent ! DA = average ionic diameter for ion A ! DB = average ionic diameter for ion B A = 0.5161D0 B = 3.3D7 DA = 3.0D-8 DB = 3.5D-8

225

! Open file containing experimental data

OPEN(4,FILE='ascn1.dat',FORM='FORMATTED',STATUS='OLD')

DO 5 I=1,M READ(4,*) XB(I) 5 CONTINUE DO 10 I=1,M READ(4,*) YB(I) 10 CONTINUE DO 15 I=1,M READ(4,*) SI(I) 15 CONTINUE CLOSE (UNIT=4)

! Calculate activity coefficients for the liquid phase ! and resin phase DO 20 I=1,M GAMAL(I) = GAML(A,B,DA,SI(I)) GAMBL(I) = GAML(A,B,DB,SI(I)) GAMAR(I) = GAMR(1,YB(I),X(1),X(2)) GAMBR(I) = GAMR(2,YB(I),X(1),X(2)) 20 CONTINUE

! Calculate values of YBCALC(I) DO 25 I=1,M DENOM(I) = GAMAL(I)*GAMBR(I)*(1.0D0-XB(I)) + & X(3)*GAMAR(I)*GAMBL(I)*XB(I) YBCALC(I) = X(3)*GAMAR(I)*GAMBL(I)*XB(I)/DENOM(I) F(I) = (YBCALC(I) - YB(I))/YB(I) 25 CONTINUE

RETURN END

! ------DOUBLE PRECISION FUNCTION GAML (A,B,D,S) IMPLICIT DOUBLE PRECISION (A-H,O-Z)

! User defined function for calculating the activity coefficient of a ! component in the liquid phase.

GAML = 10**(-1.0*A*SQRT(S)/(1.0+B*D*SQRT(S)))

RETURN END

! ------DOUBLE PRECISION FUNCTION GAMR (N,Y,PAB,PBA) IMPLICIT DOUBLE PRECISION (A-H,O-Z)

! User defined function for calculating the activity coefficient of a ! component in the resin phase. ! For N=1, calculate GAMR for component A.

IF (N.EQ.1) THEN A = 1.0 - LOG((1.0 - Y) + Y * PAB)

226 B = (1.0 – Y)/((1.0 – Y) + Y * PAB) C = Y * PBA/(Y +(1.0 – Y)*PBA) GAMR = EXP(A – B - C) ELSE A = 1.0 - LOG(Y + (1.0 - Y) * PBA) B = Y/(Y + (1.0 - Y) * PBA) C = (1.0 - Y)*PAB/((1.0 - Y)+ Y * PAB) GAMR = EXP(A – B - C) END IF

RETURN END

! ------SUBROUTINE ZXSSQ (FUNC,M,N,NSIG,EPS,DELTA,MAXFN,IOPT, & PARM,X,SSQ,F,XJAC,IXJAC,XJTJ,WORK,INFER,IER) IMPLICIT DOUBLE PRECISION (A-H,O-Z) DIMENSION X(45),F(51),PARM(4),XJAC(51),XJTJ(6),WORK(123) INTEGER U1,U5,U4 COMMON/UNITA/U1,U5,U4 DATA SIG/9.3D0/ DATA AX/0.1D0/ DATA I0,I1,I2,I3,I4,I5/0,1,2,3,4,5/ DATA P01,TENTH,HALF,ZERO,ONE,ONEP5,TWO, & TEN,HUNTW,ONEP10/0.01D0,0.1D0,0.5D0,0.0D0, & 1.D0,1.5D0,2.D0,10.0D0,1.2D2,1.D10/ ! ERROR CHECKS IER = I0 IF (M.LE.0.OR.M.GT.IXJAC.OR.N.LE.0.OR.IOPT.LT.0.OR.IOPT.GT.2) & GOTO 305 IMJC=IXJAC-M IF (IOPT.NE.2) GOTO 5 IF (PARM(I2).LE.ONE.OR.PARM(I1).LE.ZERO) GOTO 305 ! MACHINE DEPENDENT CONSTANTS 5 PREC = TEN**(-SIG-ONE) REL = TEN**(-SIG*HALF) RELCON = TEN**(-NSIG) ! WORK VECTOR IS ONCATENATION OF ! SCALED HESSIAN,GRADIENT,DELX,SCALE, ! XNEW,XBAD,F(X+DEL),F(X-DEL) IGRAD1 = ((N+1)*N)/2 IGRADL = IGRAD1+1 IGRADU = IGRAD1+N IDELX1 = IGRADU IDELXL = IDELX1+1 IDELXU = IDELX1+N ISCAL1 = IDELXU ISCALL = ISCAL1+1 ISCALU = ISCAL1+N IXNEW1 = ISCALU IXNEWL = IXNEW1+1 IXBAD1 = IXNEW1+N IFPL1 = IXBAD1+N IFPL = IFPL1+1 IFPU = IFPL1+M IFML1 = IFPU

227 IFML = IFML1+1 IMJC = IXJAC - M ! INITIALIZE VARIABLES AL = ONE CONS2 = TENTH IF (IOPT.EQ.0) GOTO 20 IF (IOPT.EQ.1) GOTO 10 AL = PARM(I1) F0 = PARM(I2) UP = PARM(I3) CONS2 = PARM(I4)*HALF GOTO 15 10 AL = P01 F0 = TWO UP = HUNTW 15 ONESF0 = ONE/F0 F0SQ = F0*F0 F0SQS4 = F0SQ**4 20 IEVAL = 0 DELTA2 = DELTA*HALF ERL2 = ONEP10 IBAD = -99 ISW = 1 ITER = -1 INFER = 0 IER = 0 DO 25 J=IDELXL,IDELXU WORK(J) = ZERO 25 CONTINUE GOTO 165 ! MAIN LOOP 30 SSQOLD = SSQ ! CALCULATE JACOBIAN IF (INFER.GT.0.OR.IJAC.GE.N.OR.IOPT.EQ.0.OR.ICOUNT.GT.0) GOTO 55 ! RANK ONE UPDATE TO JACOBIAN IJAC = IJAC+1 DSQ = ZERO DO 35 J=IDELXL,IDELXU DSQ = DSQ+WORK(J)*WORK(J) 35 CONTINUE IF (DSQ.LE.ZERO) GOTO 55 DO 50 I=1,M II=IFML1+I G = F(I)-WORK(II) K = I DO 40 J=IDELXL,IDELXU G = G+XJAC(K)*WORK(J) K = K+IXJAC 40 CONTINUE G = G/DSQ K = I DO 45 J=IDELXL,IDELXU XJAC(K) = XJAC(K)-G*WORK(J) K = K+IXJAC 45 CONTINUE 50 CONTINUE GOTO 80 ! JACOBIAN BY INCREMENTING X 55 IJAC = 0 K = -IMJC

228 DO 75 J=1,N K = K+IMJC XDABS =DABS(X(J)) HH =REL*(DMAX1(XDABS,AX)) XHOLD = X(J) X(J) = X(J)+HH ! print*, "about to call func - 1", ",",n CALL FUNC (X,M,N,WORK(IFPL)) ! print*, "call to func - 1 completed", ",",n IEVAL = IEVAL+1 X(J) = XHOLD ! print*, "isw=",isw

IF (ISW.EQ.1) GOTO 65 ! CENTRAL DIFFERENCES X(J) = XHOLD-HH ! print*, "about to call func - 2", n CALL FUNC (X,M,N,WORK(IFML)) IEVAL = IEVAL+1 X(J) = XHOLD RHH = HALF/HH DO 60 I=IFPL,IFPU II=I+M K = K+1 XJAC(K) = (WORK(I)-WORK(II))*RHH 60 CONTINUE GOTO 75 ! FORWARD DIFFERENCES 65 RHH = ONE/HH DO 70 I=1,M K = K+1 II=IFPL1+I XJAC(K) = (WORK(II)-F(I))*RHH ! print*, "xjac(",k,")=", xjac(k) 70 CONTINUE 75 CONTINUE ! CALCULATE GRADIENT 80 ERL2X = ERL2 ERL2 = ZERO K = -IMJC DO 90 J=IGRADL,IGRADU K = K+IMJC SUM = ZERO DO 85 I=1,M K = K+1 SUM = SUM+XJAC(K)*F(I) 85 CONTINUE ! print*, "loop 85 completed" WORK(J) = SUM ERL2 = ERL2+SUM*SUM 90 CONTINUE ! print*, "loop 90 completed" ERL2 = SQRT(ERL2) ! CONVERGENCE TEST FOR NORM OF GRADIENT ! print*, "ijac=",ijac IF(IJAC.GT.0) GOTO 95 IF (ERL2.LE.DELTA2) INFER = INFER+4 IF (ERL2.LE.CONS2) ISW = 2 ! CALCULATE THE LOWER SUPER TRIANGE OF

229 ! JACOBIAN (TRANSPOSED) * JACOBIAN 95 L = 0 IS = -IXJAC DO 110 I=1,N IS = IS+IXJAC JS = -IXJAC DO 105 J=1,I JS = JS+IXJAC L = L+1 SUM = ZERO DO 100 K=1,M LI = IS+K LJ = JS+K SUM = SUM+XJAC(LI)*XJAC(LJ) 100 CONTINUE XJTJ(L) = SUM 105 CONTINUE 110 CONTINUE ! print*, "loop 110 completed" ! ONVERGENCE CHECKs ! print*, "infer=", infer IF (INFER.GT.0) GOTO 315 ! print*, "ieval=", ieval IF (IEVAL.GE.MAXFN) GOTO 290 ! COMPUTE SCALING VECTOR ! print*, "iopt=",iopt IF (IOPT.EQ.0) GOTO 120 K = 0 DO 115 J=1,N K = K+J II=ISCAL1+J WORK(II) = XJTJ(K) 115 CONTINUE ! print*, "loop 115 completed" GOTO 135 ! COMPUTE SCALING VECTOR AND NORM 120 DNORM = ZERO K = 0 DO 125 J=1,N K = K+J II=ISCAL1+J WORK(II) = SQRT(XJTJ(K)) DNORM = DNORM+XJTJ(K)*XJTJ(K) 125 CONTINUE DNORM = ONE/SQRT(DNORM) ! NORMALIZE SCALING VECTOR DO 130 J=ISCALL,ISCALU WORK(J) = WORK(J)*DNORM*ERL2 130 CONTINUE ! ADD L-M FACTOR TO DIAGONAL 135 ICOUNT = 0 140 K = 0 DO 150 I=1,N DO 145 J=1,I K = K+1 WORK(K) = XJTJ(K) 145 CONTINUE II=ISCAL1+I IJ=IGRAD1+I WORK(K) = WORK(K)+WORK(II)*AL II=IDELX1+I

230 WORK(II) = WORK(IJ) 150 CONTINUE ! print*, "loop 150 completed" ! CHOLESKY DECOMPOSITION ! print*, "about to call leq in subr zxssq" 155 CALL LEQ (WORK,1,N,WORK(IDELXL),N,0,G,XHOLD,IER) ! print*, "call to leq done" IF (IER.EQ.0) GOTO 160 IER = 0 IF (IJAC.GT.0) GOTO 55 IF (IBAD.LE.0) GOTO 240 IF (IBAD.GE.2) GOTO 310 GOTO 190 160 IF (IBAD.NE.-99) IBAD = 0 ! CALCULATE SUM OF SQUARES 165 DO 170 J=1,N II=IXNEW1+J IJ=IDELX1+J WORK(II) = X(J)-WORK(IJ) 170 CONTINUE ! print*, "about to call func - 3" CALL FUNC (WORK(IXNEWL),M,N,WORK(IFPL)) IEVAL = IEVAL+1 SSQ = ZERO DO 175 I=IFPL,IFPU SSQ = SSQ+WORK(I)*WORK(I) 175 CONTINUE IF (ITER.GE.0) GOTO 185 ! SSQ FOR INITIAL ESTIMATES OF X ITER = 0 SSQOLD = SSQ DO 180 I=1,M II=IFPL1+I F(I) = WORK(II) 180 CONTINUE GOTO 55 185 IF (IOPT.EQ.0) GOTO 215 ! CHECK DESCENT PROPERTY IF (SSQ.LE.SSQOLD) GOTO 205 ! INREASE PAAMETER AND TRY AGAIN 190 ICOUNT = ICOUNT+1 AL = AL*F0SQ IF (IJAC.EQ.0) GOTO 195 IF (ICOUNT.GE.4.OR.AL.GT.UP) GOTO 200 195 IF (AL.LE.UP) GOTO 140 IF (IBAD.EQ.1) GOTO 310 GOTO 300 200 AL = AL/F0SQS4 GOTO 55 ! ADJUST MARQUARDT PARAMETER 205 IF (ICOUNT.EQ.0) AL = AL/F0 IF (ERL2X.LE.ZERO) GOTO 210 G = ERL2/ERL2X IF (ERL2.LT.ERL2X) AL = AL*DMAX1(ONESF0,G) IF (ERL2.GT.ERL2X) AL = AL*DMIN1(F0,G) 210 AL = DMAX1(AL,PREC) ! ONE ITERATION CYCLE COMPLETED 215 ITER = ITER+1 DO 220 J=1,N IJ=IXNEW1+J X(J) = WORK(IJ)

231 220 CONTINUE DO 225 I=1,M II=IFML1+I WORK(II) = F(I) II =IFPL1+I F(I) = WORK(II) 225 CONTINUE ! RELATIVE CONVERGENCE TEST FOR X IF (AL .GT. 5.0D0) GOTO 30 DO 230 J=1,N IJ=IDELX1+J XDIF =DABS(WORK(IJ))/DMAX1(DABS(X(J)),AX) IF (XDIF.GT.RELCON) GOTO 235 230 CONTINUE INFER = 1 ! RELATIVE CONVERGENCE TEST FOR SSQ 235 SQDIF = DABS(SSQ-SSQOLD)/DMAX1(SSQOLD,AX) IF (SQDIF.LE.EPS) INFER = INFER+2 GOTO 30 ! SINGULAR DECOMPOSITION 240 IF (IBAD) 255,245,265 ! CHECK TO SEE IF CURRENT ! ITERATE HAS CYCLED BACK TO ! THE LAST SINGULAR POINT 245 DO 250 J=1,N IJ=IXBAD1+J XHOLD = WORK(IJ) IF (DABS(X(J)-XHOLD).GT.RELCON*DMAX1(AX,DABS(XHOLD))) GOTO 255 250 CONTINUE GOTO 295 ! UPDATE THE BAD X VALUES 255 DO 260 J=1,N IJ=IXBAD1+J WORK(IJ) = X(J) 260 CONTINUE IBAD = 1 ! INCREASE DIAGONAL OF HESSIAN 265 IF (IOPT.NE.0) GOTO 280 K = 0 DO 275 I=1,N DO 270 J=1,I K = K+1 WORK(K) = XJTJ(K) 270 CONTINUE II=ISCAL1+I WORK(K) = ONEP5*(XJTJ(K)+AL*ERL2*WORK(II))+REL 275 CONTINUE IBAD = 2 GOTO 155 ! REPLACE ZEROES ON HESSIAN DIAGONAL 280 IZERO = 0 DO 285 J=ISCALL,ISCALU IF (WORK(J).GT.ZERO) GOTO 285 IZERO = IZERO+1 WORK(J) = ONE 285 CONTINUE IF (IZERO.LT.N) GOTO 140 IER = 38 GOTO 315 ! TEMINAL ERROR

232 290 IER = IER+1 295 IER = IER+1 300 IER = IER+1 305 IER = IER+1 310 IER = IER+129 IF (IER.EQ.130) GOTO 9000 ! OUTPUT ERL2,IEVAL,NSIG,AL, AND ITER 315 G = SIG DO 320 J=1,N IJ=IDELX1+J XHOLD = DABS(WORK(IJ)) IF (XHOLD.LE.ZERO) GOTO 320 G = DMIN1(G,-DLOG10(XHOLD/DMAX1(AX,DABS(X(J))))) 320 CONTINUE IF(N.GT.2) GOTO 330 DO 325 J = 1,N II=IGRAD1+J 325 WORK(J+5) = WORK(II) 330 WORK(I1) = ERL2+ERL2 WORK(I2) = IEVAL SSQ = SSQOLD WORK(I3) = G WORK(I4) = AL WORK(I5) = ITER IS = G IF(IS .GE. NSIG) IER=0 IF (IER.EQ.0) GOTO 9005 9000 CONTINUE WRITE(U1,9001) IER IF(IER.EQ.129.OR.IER.EQ.132) WRITE(U1,9129) IF(IER.EQ.130) WRITE(U1,9130) M,N,IOPT,(PARM(I),I=1,2) IF(IER.EQ.131) WRITE(U1,9131) WORK(I4),PARM(I3) IF(IER.EQ.133) WRITE(U1,9133) MAXFN IF(IER.EQ.38) WRITE(U1,9038) WRITE(U1,9050) WORK(I1),IS,WORK(I4),ITER,IEVAL 9005 CONTINUE ! ! FORMAT STATEMENTS ! 9001 FORMAT(1H ,' ZXSSQ ERROR NO:',I3) 9038 FORMAT(1H ,'TRIVIAL SOLUTION') 9050 FORMAT(2X,'NORM.GRAD=',E10.4,2X,'NO.OF SIGNIFICANT DIGITS:', & I5,/,3X,'MARQUARDT PARAMETER:',E10.4,/ & 2X,I4,' ITERATIONS;',2X,I4,' FUNCTION EVALUATIONS.') 9129 FORMAT(1H ,'SINGULAR JACOBIAN') 9130 FORMAT(1H ,'INCORRECT: M=',I5,' N=',I5,' IOPT=',I5, & ' PARM(1)=',E10.4,' PARM(2)=',E10.4) 9131 FORMAT(1H ,'MARQUARDT PARAMETER ',E10.4, & ' EXCEED PARM(3)',E10.4) 9133 FORMAT(1H ,'FUNCTION EVALUATIONS EXCEED MAXFN',I5) RETURN END

SUBROUTINE LEQ (A,M,N,B,IB,IDGT,D1,D2,IER) implicit double precision (a-h,o-z) DIMENSION A(1),B(IB,1) DOUBLE PRECISION A,D1,D2

233 ! print*, "about to run first exec lines in leq" ! print*, "idgt=",idgt ! idgt=idgt ! print*, "past exec line 1" ! INITIALIZE IER IER = 0 ! DECOMPOSE A ! print*, "about to call dec in subr leq" CALL DEC (A,A,N,D1,D2,IER) IF (IER.NE.0) GOTO 9000 ! PERFORM ELIMINATiON DO 5 I = 1,M CALL ELM (A,B(1,I),N,B(1,I)) 5 CONTINUE GOTO 9005 9000 CONTINUE WRITE(0,9001) 9001 FORMAT(1H ,' DECOMPOSITION ERROR') 9005 RETURN END SUBROUTINE DEC(A,UL,N,D1,D2,IER) IMPLICIT DOUBLE PRECISION (A-H,O-Z) DIMENSION A(1),UL(1) DATA ZERO,ONE,FOUR,SIXTN,SIXTH/0.0D0,1.D0,4.D0,16.D0,.0625D0/ D1=ONE D2=ZERO RN = ONE/(FLOAT(N)*SIXTN) IP = 1 IER=0 DO 45 I = 1,N IQ =IP IR = 1 DO 40 J=1,I X = A(IP) IF (J .EQ. 1) GOTO 10 DO 5 K=IQ,IP1 X= X - UL(K) * UL(IR) IR = IR+1 5 CONTINUE 10 IF (I.NE.J) GOTO 30 D1 = D1*X IF (A(IP) + X*RN .LE. A(IP)) GOTO 50 15 IF (DABS(D1).LE.ONE) GOTO 20 D1 = D1 * SIXTH D2 = D2 + FOUR GOTO 15 20 IF (DABS(D1) .GE. SIXTH) GOTO 25 D1 = D1 * SIXTN D2 = D2 - FOUR GOTO 20 25 UL(IP) = ONE/DSQRT(X) GOTO 35 30 UL(IP) = X * UL(IR) 35 IP1 = IP IP = IP+1 IR = IR+1 40 CONTINUE 45 CONTINUE GOTO 9005 50 IER = 129

234 9005 RETURN END SUBROUTINE ELM (A,B,N,X) IMPLICIT DOUBLE PRECISION (A-H,O-Z) DIMENSION A(1),B(1),X(1) DATA ZERO/0.0d0/ ! SOLUTION OF LY = B IP=1 IW = 0 DO 15 I=1,N T=B(I) IM1 = I-1 IF (IW .EQ. 0) GOTO 9 IP=IP+IW-1 DO 5 K=IW,IM1 T = T-A(IP)*X(K) IP=IP+1 5 CONTINUE GOTO 10 9 IF (T .NE. ZERO) IW = I IP = IP+IM1 10 X(I)=T*A(IP) IP=IP+1 15 CONTINUE ! SOLUTION OF UX = Y N1 = N+1 DO 30 I = 1,N II = N1-I IP=IP-1 IS=IP IQ=II+1 T=X(II) IF (N.LT.IQ) GOTO 25 KK = N DO 20 K=IQ,N T = T - A(IS) * X(KK) KK = KK-1 IS = IS-KK 20 CONTINUE 25 X(II)=T*A(IS) 30 CONTINUE RETURN END

235 Appendix B: Experimental data for ion-exchange equilibria between thiocyanate and chloride Appendix B reports the experimental results for ion-exchange equilibria between thiocyanate and chloride shown in Chapter 5. All the experiments were repeated in duplicate and the reported data is the average value between the two experiments.

Table B1: Concentration profile of thiocyanate adsorption using Purolite A500u/2778 ion- exchange resin with an initial thiocyanate concentration of 1200ppm NaSCN. Temperature = 30oC, Pressure = 1atm. Ref: Figure 5.4. The amount of NaSCN in solution is reported in ppm NaSCN. Time (min) Run 1 Run 2 Run 3 Run 4 Run 5 Run 6 Run 7 Run 8 0 1208.5 1200.2 1140.0 1158.5 1153.1 1178.3 1179.2 1151.5 5 686.2 542.7 628.8 848.7 1074.9 1157.6 1173.1 1151.5 10 432.2 286.7 421.9 760.7 1063.2 1152.2 1174.1 1155.6 15 279.4 142.2 318.5 731.1 1060.0 1155.4 1172.1 1156.6 30 64.1 89.9 243.4 709.6 1063.2 1156.5 1173.1 1153.5 60 21.1 70.8 240.4 706.5 1062.1 1153.3 1164.0 1150.4 120 18.0 70.9 242.4 709.6 1069.6 1159.8 1181.3 1158.7 240 17.9 71.1 235.3 708.6 1060.0 1153.3 1185.3 1158.7 360 17.5 71.1 239.4 707.6 1050.3 1148.9 1179.2 1148.3 480 17.5 71.6 241.4 708.6 1052.5 1135.9 1176.2 1154.6 Table B2: Equilibrium adsorption data for the adsorption of thiocyanate with an initial thiocyanate concentration of 300ppm NaSCN (3.70 m.mol/l NaSCN) using Purolite A500u/2778 ion-exchange resin. Temperature = 30oC, Pressure = 1atm. Solution Phase Resin Phase Run [SCN] (m.mol/l) [Cl] (m.mol/l) Fraction [SCN] (m.mol/l) [Cl] (m.mol/l) Fraction 1 0.01 3.78 0.00 0.09 1.13 0.08 2 0.05 3.51 0.01 0.18 1.04 0.15 3 0.03 3.65 0.01 0.27 0.95 0.22 4 0.06 3.56 0.02 0.36 0.86 0.30 5 0.13 3.57 0.03 0.45 0.77 0.37 6 0.15 3.50 0.04 0.54 0.69 0.44 7 0.20 3.68 0.05 0.63 0.59 0.52 8 0.20 3.61 0.05 0.72 0.50 0.59 9 0.32 3.53 0.08 0.81 0.41 0.66 10 0.46 3.29 0.12 0.89 0.33 0.73 11 0.66 3.16 0.17 0.97 0.25 0.79 12 0.90 2.66 0.25 1.04 0.19 0.85 13 1.30 2.40 0.35 1.10 0.13 0.90 14 1.93 1.85 0.51 1.14 0.08 0.93 15 2.60 1.08 0.71 1.17 0.05 0.96 16 3.12 0.62 0.83 1.19 0.04 0.97 17 3.11 0.50 0.86 1.20 0.03 0.98 18 3.35 0.37 0.90 1.21 0.02 0.99 19 3.56 0.16 0.96 1.21 0.01 0.99 20 3.42 0.19 0.95 1.22 0.01 0.99 21 3.46 0.30 0.92 1.22 0.00 1.00

236 Table B3: Equilibrium adsorption data for the adsorption of thiocyanate with an initial thiocyanate concentration of 600ppm NaSCN (7.40 m.mol/l NaSCN) using Purolite A500u/2778 ion-exchange resin. Temperature = 30oC, Pressure = 1atm.

Solution Phase Resin Phase Run [SCN] (m.mol/l) [Cl] (m.mol/l) Fraction [SCN] (m.mol/l) [Cl] (m.mol/l) Fraction 1 0.11 7.58 0.01 0.19 0.95 0.17 2 0.09 7.12 0.01 0.37 0.78 0.32 3 0.17 6.59 0.03 0.53 0.61 0.47 4 0.41 6.73 0.06 0.70 0.44 0.61 5 0.75 6.02 0.11 0.85 0.29 0.74 6 1.69 5.02 0.25 0.98 0.17 0.85 7 2.39 3.96 0.38 1.08 0.07 0.94 8 5.50 1.84 0.75 1.12 0.02 0.98 9 6.38 0.68 0.90 1.14 0.01 0.99 10 7.12 0.15 0.98 1.14 0.00 1.00 11 7.27 0.08 0.99 1.14 0.00 1.00

Table B4: Equilibrium adsorption data for the adsorption of thiocyanate with an initial thiocyanate concentration of 900ppm NaSCN (11.10 m.mol/l NaSCN) using Purolite A500u/2778 ion-exchange resin. Temperature = 30oC, Pressure = 1atm. Solution Phase Resin Phase Run [SCN] (m.mol/l) [Cl] (m.mol/l) Fraction [SCN] (m.mol/l) [Cl] (m.mol/l) Fraction 1 0.14 10.96 0.01 0.27 0.90 0.23 2 0.32 10.94 0.03 0.55 0.62 0.47 3 0.81 9.30 0.08 0.78 0.39 0.67 4 2.39 8.27 0.22 0.99 0.18 0.84 5 5.42 5.00 0.52 1.11 0.06 0.95 6 8.71 1.75 0.83 1.16 0.01 0.99 7 10.25 0.11 0.99 1.16 0.01 0.99 8 10.93 0.23 0.98 1.16 0.01 0.99 9 10.84 0.31 0.97 1.17 0.00 1.00 10 10.99 0.01 1.00 1.17 0.00 1.00 11 11.03 -0.08 1.01 1.17 0.00 1.00

Table B5: Equilibrium adsorption data for the adsorption of thiocyanate with an initial thiocyanate concentration of 1200ppm NaSCN (14.80 m.mol/l NaSCN) using Purolite A500u/2778 ion-exchange resin. Temperature = 30oC, Pressure = 1atm.

Solution Phase Resin Phase Run [SCN] (m.mol/l) [Cl] (m.mol/l) Fraction [SCN] (m.mol/l) [Cl] (m.mol/l) Fraction 1 0.22 14.69 0.01 0.37 0.81 0.31 2 0.88 13.92 0.06 0.72 0.46 0.61 3 2.98 11.08 0.21 0.99 0.18 0.84 4 8.74 5.55 0.61 1.13 0.04 0.96 5 12.98 1.24 0.91 1.16 0.01 0.99 6 14.01 0.52 0.96 1.18 0.00 1.00 7 14.51 0.04 1.00 1.18 0.00 1.00 8 14.24 -0.04 1.00 1.18 0.00 1.00

237 Table B6: Equilibrium adsorption data for the adsorption of thiocyanate with an initial thiocyanate concentration of 1500ppm NaSCN (18.50 m.mol/l NaSCN) using Purolite A500u/2778 ion-exchange resin. Temperature = 30oC, Pressure = 1atm. Solution Phase Resin Phase Run [SCN] (m.mol/l) [Cl] (m.mol/l) Fraction [SCN] (m.mol/l) [Cl] (m.mol/l) Fraction 1 0.35 18.03 0.02 0.45 0.60 0.43 2 1.93 16.54 0.10 0.69 0.18 0.79 3 17.89 10.49 0.63 0.95 -0.08 1.09 4 13.67 3.56 0.79 1.04 -0.17 1.19 5 17.60 0.20 0.99 1.05 -0.17 1.20 6 18.29 0.07 1.00 1.05 -0.17 1.20 7 18.15 0.03 1.00 1.05 -0.17 1.20

238 Appendix C: Experimental for the recovery of cyanide from thiocyanate using hydrogen peroxide

Appendix C reports the experimental results for the recovery of cyanide from thiocyanate from thiocyanate using hydrogen peroxide shown in Chapter 6.

Table C1: Data for the oxidation of thiocyanate using hydrogen peroxide to recover cyanide. Initial thiocyanate concentration 1000ppm, stoichiometric amount of hydrogen peroxide and 10g/l sulphuric acid at room temperature and pressure. Ref: Figure 6.5.

- - Time (min) [SCN ] (mM) [H2O2] (mM) [CN ] (mM) 0 17.21 79.46 0.00 20 12.56 67.69 0.48 40 9.39 61.80 1.22 60 7.49 55.92 2.31 120 5.22 38.26 5.81 240 3.44 26.49 10.73 360 2.19 20.60 13.44 480 1.45 17.66 14.94 1440 0.31 8.83 17.39

Table C2: Data for the comparison of thiocyanate concentration against time at different solution pH. The initial concentration of thiocyanate is 100ppm (17.2mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present at room temperature and pressure. No copper was added to solution. Ref: Figure 6.6.

Time 10 g/l pH 2 pH 7 pH 12 (min) [SCN-] (mM) 0 17.21 17.21 17.21 17.21 20 12.56 14.34 16.05 14.63 40 9.39 13.77 15.01 14.20 60 7.49 12.65 14.71 14.14 120 5.22 10.67 14.11 13.05 240 3.44 8.81 12.89 11.09 360 2.19 7.08 11.69 10.00 480 1.45 5.60 10.19 8.77 1440 0.31 1.40 4.70 3.77

239 Table C3: Data for the comparison of cyanide recovery against time at different solution pH. The initial concentration of thiocyanate is 100ppm (17.2mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present at room temperature and pressure. No copper was added to solution. Ref: Figure 6.7.

Time 10 g/l pH 2 pH 7 pH 12 (min) [CN-] (mM) 0 0.00 0.00 0.00 0.00 20 0.48 0.29 0.23 0.16 40 1.22 0.70 0.36 0.16 60 2.31 0.80 0.54 0.17 120 5.81 1.28 0.79 0.29 240 10.73 3.19 1.06 0.50 360 13.44 5.02 1.74 0.55 480 14.94 6.57 2.80 0.61 1440 17.39 14.39 8.72 0.74

Table C4: Data for the pH profile of hydrogen peroxide oxidation of thiocyanate with different initial starting pH. The initial concentration of thiocyanate is 1000ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution. Ref: Figure 6.8.

Time PH (min) 10 g/l 2 7 12 0 1.01 1.77 7.2 12.04 20 1.03 1.91 3.47 9.62 40 1.08 1.91 3.47 7.4 60 0.93 1.9 3.32 7.04 120 1.07 1.88 3.19 5.95 240 1.08 1.87 3.04 4.29 360 1.15 1.87 2.97 4.24 480 1.03 1.86 2.95 4.25 1440 1.02 2.47 2.63 3.98

Table C5: Data for the Eh profile of hydrogen peroxide oxidation of thiocyanate with different initial starting pH. The initial concentration of thiocyanate is 1000ppm (17.2 mM), the amount of hydrogen peroxide added to the experiment is double the stoichiometry required to oxidise the thiocyanate present. No copper was added to solution. Ref: Figure 6.9.

Time Eh (V) (min) 10 g/l 2 7 12 0 0.49 0.464 0.396 0.071 20 0.513 0.47 0.39 0.11 40 0.489 0.472 0.379 0.201 60 0.488 0.433 0.418 0.25 120 0.47 0.421 0.432 0.316 240 0.502 0.42 0.439 0.383 360 0.486 0.43 0.424 0.396 480 0.451 0.414 0.438 0.398 1440 0.464 0.509 0.472 0.386

240 Table C6: Data for the comparison of thiocyanate concentration against time at different experimental conditions outlined in table 6.2. Ref: Figure 6.10.

Time [SCN-] (mM) (min) Experiment 1 Experiment 2 Experiment 3 Experiment 4 Experiment 5 Experiment 6 Experiment 7 Experiment 8 0 17.21 17.21 17.21 17.21 17.21 17.21 17.21 17.21 20 11.44 4.52 8.50 2.76 3.53 1.40 3.53 1.41 40 8.98 3.58 6.97 1.35 2.05 0.15 1.83 0.29 60 7.15 3.04 5.13 0.94 1.36 0.16 1.11 0.09 120 4.82 2.50 2.19 0.54 0.53 0.05 0.67 0.02 240 2.84 1.83 0.36 0.19 0.18 0.14 0.64 0.02 360 2.04 0.65 0.32 0.14 0.16 0.16 0.56 0.05 480 1.59 0.34 0.21 0.09 0.14 0.04 0.43 0.06 1440 0.50 0.21 0.02 0.07 0.01 0.01 0.21 0.01

Table C7: Data for the comparison of hydrogen peroxide concentration against time at different experimental conditions outlined in table 6.2. Ref: Figure 6.11.

Time [H2O2] (mM) (min) Experiment 1 Experiment 2 Experiment 3 Experiment 4 Experiment 5 Experiment 6 Experiment 7 Experiment 8 0 97.12 102.89 164.63 82.31 161.69 144.05 70.56 155.81 20 76.52 82.31 117.59 26.46 99.95 70.56 32.34 73.49 40 61.80 70.56 105.83 17.64 88.19 64.68 23.52 67.62 60 52.97 58.80 91.13 16.17 79.37 61.74 17.64 64.68 120 38.26 44.10 82.31 11.76 70.56 64.68 11.76 64.68 240 26.49 32.34 67.62 8.82 70.56 64.68 10.29 64.68 360 26.49 26.46 64.68 7.35 67.62 61.74 7.35 64.68 480 23.54 22.05 64.68 5.88 70.56 61.74 8.82 64.68 1440 11.77 11.76 55.86 2.94 61.74 55.86 7.06 61.74

241 Table C8: Data for the comparison of cyanide recovered against time at different experimental conditions outlined in table 6.2. Ref: Figure 6.12.

Time Cyanide Recovered (%) (min) Experiment 1 Experiment 2 Experiment 3 Experiment 4 Experiment 5 Experiment 6 Experiment 7 Experiment 8 0 2 2 2 2 2 2 2 2 20 7 6 8 12 10 16 13 11 40 10 10 14 22 19 30 22 22 60 16 12 21 30 27 41 30 31 120 30 21 41 45 48 60 47 50 240 48 37 65 63 66 80 69 69 360 60 47 78 71 74 90 79 81 480 69 55 84 77 77 94 84 86 1440 88 84 88 89 82 101 94 93

242 Appendix D: Experimental for the electrochemical study of thiocyanate oxidation

Appendix D reports the electrochemical study of thiocyanate oxidation shown in Chapter 7.

Table D1: Data for cyclic voltammogram of the glassy carbon electrode in 0.1M Na2SO4 at pH 2, rest potential 0V and a scan rate of 40mV/s at a temperature of 25oC. Ref: Figure 7.6. Potential vs Ag/AgCl Current (V) (μA) 0.1 2.02 0.2 0.91 0.3 0.60 0.4 1.67 0.5 1.38 0.6 2.26 0.7 1.13 0.8 0.28 0.9 1.12 1 2.90 1.1 3.49 1.2 4.91 1.3 8.00 1.4 14.47 1.5 24.28 1.6 42.70 1.5 6.87 1.4 2.96 1.3 0.77 1.2 -1.37 1.1 -1.41 1 -1.70 0.9 -2.18 0.8 -1.80 0.7 -1.75 0.6 -2.34 0.5 -2.12 0.4 -2.26 0.3 -1.97 0.2 -3.60 0.1 -3.10

243

Table D2: Data for the cyclic voltammogram of a glassy carbon in a solution of 0.1M Na2SO4 + 0.01M NaSCN, rest potential=0.5V (vs. Ag-AgCl electrode), pH 2, temperature 25oC. The scan rate was varied from 20 to 40 mV/s. Ref: Figure 7.7. Potential vs Ag/AgCl Current (μA) (V) 10 mV/s 20 mV/s 30 mV/s 40 mV/s 0.5 0.40 1.58 -0.39 2.29 0.6 1.00 3.98 -0.09 3.92 0.7 3.30 17.04 5.44 7.97 0.8 12.80 53.62 22.03 42.88 0.9 53.00 171.29 88.43 215.22 1 163.10 465.40 311.32 657.80 1.1 291.90 779.74 655.71 1099.06 1.2 352.60 907.77 928.65 1247.65 1.3 390.80 944.53 1132.69 1192.76 1.4 395.00 887.54 1218.09 1207.88 1.5 370.30 788.07 972.84 1029.57 1.4 324.40 645.40 708.19 761.24 1.3 223.30 456.77 445.12 468.15 1.2 127.10 295.51 258.75 278.16 1.1 78.30 193.12 151.65 164.39 1 36.20 92.32 72.57 79.61 0.9 10.60 27.17 21.11 19.37 0.8 2.50 5.37 4.96 2.65 0.7 0.80 0.76 -0.25 -1.35 0.6 0.00 -0.68 -0.10 -2.35 0.5 0.00 -0.22 -0.39 -3.04

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Table D3: Data for the stationary cyclic voltammogram of a glassy carbon in a solution of 0.1M o Na2SO4, 0.01M NaSCN, rest potential=0.5V (vs. Ag-AgCl electrode), pH 12, temperature 25 C. The scan rate was varied from 10 to 40 mV/s. Ref: Figure 7.9. Potential vs Ag/AgCl Current (μA) (V) 10 mV/s 20 mV/s 30 mV/s 40 mV/s 0.5 0.23 4.53 3.04 0.60 0.6 1.14 7.85 -0.19 3.69 0.7 0.99 16.41 14.51 8.83 0.8 3.21 55.56 98.78 26.42 0.9 9.07 154.36 264.67 100.48 1 21.53 442.22 675.74 369.72 1.1 50.33 770.84 1071.40 844.56 1.2 98.61 930.46 1172.13 1257.62 1.3 194.72 920.17 1070.00 1351.18 1.4 454.41 905.61 1010.57 1293.90 1.5 696.85 796.04 886.97 1091.47 1.6 640.97 727.97 793.72 937.15 1.7 567.65 709.66 757.03 870.01 1.8 532.66 741.63 779.16 877.23 1.9 493.35 852.04 886.54 992.91 2 847.55 1078.45 1136.15 1209.73 1.9 598.72 803.44 765.97 793.40 1.8 520.01 698.72 628.82 644.74 1.7 489.20 660.41 580.85 574.24 1.6 466.17 643.91 564.31 536.51 1.5 426.05 622.44 557.47 499.57 1.4 298.62 487.22 474.54 395.62 1.3 117.78 276.54 318.15 207.25 1.2 60.77 181.59 244.66 124.33 1.1 49.14 118.42 171.02 74.94 1 31.63 67.90 104.22 43.13 0.9 16.33 29.06 46.62 16.44 0.8 5.83 7.32 14.85 4.66 0.7 1.31 2.11 -0.29 -0.58 0.6 1.12 -0.70 -3.80 -2.18 0.5 0.04 -0.54 -0.29 -1.00

245 Table D4: Data for the Evan’s diagram. The Rotating disc electrode made from glassy carbon in o a solution of 0.1M Na2SO4+0.01M NaSCN, pH 12, temperature 25 Cn and a rest potential of 0.5V vs Ag-AgCl reference. The rotation frequency was varied from 50rpm to 500 rpm. Ref: Figure 7.11. Potential vs Ag/AgCl Current (μA) (V) 50 rpm 100 rpm 200 rpm 500 rpm 0.5 0.59 0.40 1.23 0.83 0.6 1.24 5.14 5.56 4.64 0.7 3.19 10.04 17.89 16.56 0.8 9.76 27.09 120.24 45.61 0.9 29.60 82.99 142.30 144.77 1 72.54 251.24 444.17 376.07 1.1 165.60 488.85 713.76 824.20 1.2 286.76 912.35 1333.15 1319.06 1.3 462.15 1323.38 1853.07 2089.08 1.4 892.50 1740.32 2302.44 2755.48 1.5 1334.66 2285.01 2916.69 3505.50 1.6 1517.11 2587.14 3474.48 4215.03 1.7 1478.69 2628.10 3728.21 5024.66 1.8 1483.22 2652.04 3832.59 5542.02 1.9 1508.98 2676.30 3825.72 5685.02 2 1538.26 2677.99 3876.60 5914.75

Table D5: Data for the Evan’s diagram. Rotating disc electrode made from glassy carbon in a o solution of 0.1M Na2SO4+0.01M NaSCN, pH 2 and 12, temperature 25 Cn and a rest potential of 0.5V vs Ag-AgCl reference. The rotation frequency was varied from 50rpm and 100 rpm at each pH. Ref: Figure 7.13.

Current (μA) Potential vs Ag/AgCl pH 2 pH 12 (V) 50 rpm 100 rpm 50 rpm 100 rpm 0.5 0.91 0.90 0.59 0.40 0.6 0.48 15.13 1.24 5.14 0.7 12.18 87.53 3.19 10.04 0.8 58.21 296.29 9.76 27.09 0.9 196.33 740.97 29.60 82.99 1 468.85 1342.41 72.54 251.24 1.1 801.94 1836.11 165.60 488.85 1.2 1083.16 2040.99 286.76 912.35 1.3 1380.30 2184.00 462.15 1323.38 1.4 1733.13 2404.08 892.50 1740.32 1.5 1956.43 2715.39 1334.66 2285.01 1.6 1928.95 2820.59 1517.11 2587.14 1.7 1946.75 2791.71 1478.69 2628.10 1.8 1978.09 2796.19 1483.22 2652.04 1.9 2085.53 2899.22 1508.98 2676.30 2 2305.16 3034.97 1538.26 2677.99

246 Table D6: Data for the comparison of the stationary cyclic voltammogram of a glassy carbon in a solution of 0.1M Na2SO4, 0.01M NaSCN with copper added to one experiment (concentration of o 0.005M CuSO4), rest potential=0.5V (vs. Ag-AgCl electrode), pH 12, temperature 25 C and a voltage scanning rate of 40mV/s. Ref: Figure 7.14. Potential vs Ag/AgCl Current (μΑ)

(V) CuSO4 No CuSO4 0.5 21.68 0.60 0.6 9.41 3.69 0.7 21.20 8.83 0.8 90.62 26.42 0.9 272.31 100.48 1 556.23 369.72 1.1 746.74 844.56 1.2 850.59 1257.62 1.3 982.00 1351.18 1.4 1155.81 1293.90 1.5 986.29 1091.47

Table D7: Data for the comparison between Evan’s diagram for Rotating disc electrode made from glassy carbon in a solution of 0.1M Na2SO4+0.01M NaSCN with one set of experiments o containing 0.005M CuSO4, pH 2, temperature 25 C and a rest potential of 0.5V vs Ag-AgCl reference. Ref: Figure 7.15. Current (μA) Potential vs Ag/AgCl No Copper Copper (V) 50 rpm 100 rpm 50 100 0.5 0.91 0.90 26.83 29.88 0.6 0.48 15.13 12.12 25.25 0.7 12.18 87.53 30.70 51.31 0.8 58.21 296.29 108.71 186.14 0.9 196.33 740.97 325.10 597.22 1 468.85 1342.41 679.61 1123.48 1.1 801.94 1836.11 914.91 1378.46 1.2 1083.16 2040.99 1041.74 1562.31 1.3 1380.30 2184.00 1286.23 1886.83 1.4 1733.13 2404.08 1677.63 2306.30 1.5 1956.43 2715.39 1726.17 2481.65 1.6 1928.95 2820.59 1714.18 2529.42 1.7 1946.75 2791.71 1746.61 2557.55 1.8 1978.09 2796.19 1807.55 2591.24 1.9 2085.53 2899.22 1929.65 2732.43 2 2305.16 3034.97 2208.16 2974.20

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