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Mechanism and kinetics of chalcopyrite passivation and depassivation during ferric and microbial leaching Tshilombo, Alain Fuamba 2006

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MECHANISM AND KINETICS OF CHALCOPYRITE PASSIVATION AND DEPASSIVATION DURING FERRIC AND MICROBIAL LEACHING By A L A I N F U A M B A T S H I L O M B O B.Eng . , Faculte Polytechnique, University of Lubumbashi , 1994 M.Sc . , The University of Pretoria, 2000 A T H E S I S S U B M I T T E D IN P A R T I A L F U L F I L M E N T O F T H E R E Q U I R E M E N T S F O R T H E D E G R E E O F DOCTOR OF PHILOSOPHY In T H E F A C U L T Y O F G R A D U A T E S S T U D I E S (Department of Materials Engineering) T H E U N I V E R S I T Y O F BRITISH C O L U M B I A December 2004 A B S T R A C T Chalcopyrite is known to be recalcitrant to conventional hydrometallurgical and biohydrometallurgical processes. Formation of passive layers on the chalcopyrite surface results in slow and incomplete leaching. The nature of how these passive layers are formed is the subject of much controversy. T h e most likely explanation is the formation of polysulphide compounds or copper-rich intermediate products on the chalcopyrite surface. The formation of these products depends mainly on temperature and solution potential. Based on these observations, electrochemical techniques were used to study the behaviour of chalcopyrite under a variety of conditions similar to ferric and microbial leaching. Electrochemical techniques have the advantage over other techniques of measuring properties at the solid-liquid interface. Slow chalcopyrite leaching was mainly observed under the following conditions: > low temperature ( 2 5 ° C ) and low potential (0.45 to 0.6 V S C E ) > high temperature ( 6 5 ° C ) and high potential (above 0.6 V S C E ) Leaching was accelerated at high temperature ( 6 5 ° C ) under mildly oxidizing conditions (0.45 to 0.55 V S C E ) . T h e study also indicated that a polarized chalcopyrite surface inhibits ferric reduction and that the presence of pyrite during chalcopyrite leaching can be beneficial. The electrochemical study was validated in leaching tests carried out in a stirred-tank reactor with fine chalcopyrite particles. Leaching was retarded at low temperatures due to the presence of an induction period. The duration of the induction period decreased with increasing temperature. The addition of pyrite significantly increased both the rate and the extent of chalcopyrite leaching. Complete conversion of chalcopyrite was obtained within 16 hours at 6 5 ° C at a pyritexhalcopyrite mass ratio of 2:1. A n electrochemical model that takes into consideration the galvanic interaction with pyrite and the "passivation" of chalcopyrite was proposed. The addition of microorganisms to the leaching system was investigated. Chalcopyrite was leached almost to completion (95%) within 30 days in the presence of thermophilic bacteria at low potentials and high temperatures. The bioleaching rate of chalcopyrite was further increased with the addition of pyrite. Finally, atmospheric leaching of chalcopyrite was carried out at 8 0 ° C under a range of conditions. Complete copper extraction was attained in 4 hours at a pyritexhalcopyrite ratio of 4:1. The present study has shown that chalcopyrite passivation can be prevented at low solution potentials, high temperatures and in the presence of moderate amounts of pyrite. T A B L E O F C O N T E N T S A B S T R A C T ii T A B L E O F C O N T E N T S iv LIST O F F I G U R E S vii LIST O F T A B L E S xii LIST O F S Y M B O L S xiii A C K N O W L E D G M E N T S xv C H A P T E R 1. I N T R O D U C T I O N 1 C H A P T E R 2. L I T E R A T U R E R E V I E W 5 2.1. Properties of chalcopyrite 5 2.2. Hydrometallurgical processes for copper sulphide minerals in sulphate media 6 2.2.1. Mt Gordon process 6 2.2.2. Activox process 8 2.2.3. MIM/Highlands Albion (Nenatech) process 8 2.2.4. A A C / U B C Hydrometallurgy process 9 2.2.5. Dynatec process 10 2.2.6. Total Pressure Oxidation process 10 2.2.7. C E S L process 11 2.2.8. BacTech/Mintek process 11 2.2.9. B ioCop process 12 2.2.10. G e o C o a t process 12 2.2.11. Summary 13 2.3. Thermodynamic considerations 13 2.4. Retarding effect on the dissolution of chalcopyrite 15 2.4.1. Passivation of chalcopyrite during ferric leaching 16 2.4.2. Passivation of chalcopyrite during bioleaching 27 2.4.3. Passivation during the electrochemical dissolution of chalcopyrite 32 2.5. Methods for activating the passive film during the ferric leaching and bioleaching of chalcopyrite 45 2.5.1. Mechanical activation 45 2.5.2. Addition of silver 46 2.5.3. Beneficial effect of galvanic interactions between chalcopyrite and associated minerals 47 2.5.4. Thermophilic bioleaching 49 iv 2.6. Conclusions and focus of the present study 50 C H A P T E R 3. M E T H O D O L O G Y A N D E X P E R I M E N T A L P R O C E D U R E 52 3.1. Electrochemical measurements 53 3.1.1. Working electrodes 53 3.1.2. Apparatus 55 3.1.3. Procedure 57 3.2. Chemical leaching experiments 58 3.2.1. Chalcopyrite samples 58 3.2.2. Reagents 58 3.2.3. Apparatus 58 3.2.4. Procedure 60 3.2.5. Sampling and analysis 62 3.2.6. Oxidation of Fe(ll) to Fe(lll) by hydrogen peroxide 62 3.3. Bioleaching experiments 63 3.3.1. Material 63 3.3.2. Microorganisms 63 3.3.3. Equipment 64-3.3.4. Procedure 64 3.3.5. Cell counting 65 3.4. Column tests 65 3.4.1. Material 66 3.4.2. Apparatus 67 3.4.3. Procedure.. 68 3.5. Atmospheric leaching 69 3.5.1. Chalcopyrite samples 69 3.5.2. Pyrite samples 70 3.5.3. Reagents 70 3.5.4. Apparatus 70 C H A P T E R 4. R E S U L T S A N D D I S C U S S I O N 71 4.1. Electrochemical leaching 71 4.1.1. Anodic behaviour of chalcopyrite in acidic medium 71 4.1.2. Cathodic reduction of ferric ions on chalcopyrite 83 4.1.3. Mixed potential of chalcopyrite as a function of Fe(lll) and Fe(ll) concentrations 87 4.2. Chemical leaching 90 4.2.1. Effect of solution potential on reaction rate 90 v 4.2.2. Effect of temperature at constant Fe(lll):Fe(ll) ratio 101 4.2.3. Effect of pyrite on the ferric leaching of chalcopyrite 102 4.2.4. Development of an electrochemical model for the dissolution of chalcopyrite 114 4.2.5. Development of an electrochemical model for the galvanic interaction between pyrite and chalcopyrite 117 4.2.6. Electrochemical reaction and surface passivation model 121 4.2.7. Validation of the electrochemical-passivation model 122 4.3. Bioleaching of chalcopyrite 131 4.3.1. Mesophiles 131 4.3.2. Moderate thermophiles 135 4.3.3. Extreme thermophiles 136 4.3.4. Comparison of bioleaching with mesophiles and thermophiles 139 4.4. Column leaching 144 4.5. Atmospheric leaching 145 4.5.1. Effect of pyrite addition 146 4.5.2. Effect of initial solution potential 148 4.5.3. Effect of particle size 150 4.5.4. Effect of pyrite source 154 4.5.5. Effect of pulp density 156 4.5.6. Effects of impeller speed, choice of primary oxidant and acidity 156 4.5.7. Yield of elemental sulphur 160 C H A P T E R 5. C O N C L U S I O N S 162 C H A P T E R 6. F U T U R E W O R K A N D R E C O M M E N D A T I O N S 165 References 166 Appendix A . Statistical analysis of experimental results 180 vi L I S T O F F I G U R E S Figure 2.1 Crystal structure of chalcopyrite 5 Figure 2.2 Potential-pH diagram of the C u - F e - S - H 2 0 system at 2 5 ° C : all solutes at 0.1 M activity except C u 2 + at 0.01 M 14 Figure 2.3 Concentrations of Fe(lll)-sulfato and bisulfato complexes in 0.2 M H2SO4 solutions having different Fe(S0 4 ) i .5 concentrations 18 Figure 2.4 Variation of reaction curves with temperature 19 Figure 2.5 Elemental sulphur-water E h - p H diagram with extended sulfur stability 21 Figure 2.6 Leaching curves for the dissolution of various size fractions of natural chalcopyrite in sulfate solutions 24 Figure 2.7 Leaching rate curve of chalcopyrite with ferric sulfate 24 Figure 2.8 Stability of various iron precipitates as a function of pH and temperature 26 Figure 2.9 Scheme visualizing indirect leaching and direct leaching of metal sulphide 28 Figure 2.10 Effect of ferrous ion on the anodic polarization curve of Transvaal C u F e S 2 in 1 M H 2 S 0 4 , 40 m V m i n - 1 , 2 5 ° C 36 Figure 2.11 Anodic polarizarion curves for C u F e S 2 from 6 different locations in 1 M H 2 S 0 4 l 30 m V m i n - 1 , 2 5 ° C 37 Figure 2.12 Schematic representation of galvanic interactions between chalcopyrite and intermediate phases formed during the dissolution of chalcopyrite , 39 Figure 2.13 Mixed potential for Type I leaching 42 Figure 2.14 Mixed potential for Type II leaching 42 Figure 2.15 Mixed potential for Type III leaching 43 Figure 2.16 Mixed potential for Type IV leaching 43 Figure 3.1 Electrochemical apparatus 56 Figure 3.2 Schematic representation of the controlled-potential chemical leaching system 59 Figure 3.3 Solution potential as a function of the Fe(lll):Fe(ll) ratio, pH 1.4, 2 5 ° C 61 Figure 3.4 Schematic drawing of a column 67 Figure 3.5 Column leach apparatus 68 vii Figure 4.1 Effect of temperature on the anodic dissolution of chalcopyrite, pH 1.5, scan rate = 1 mV s~1 73 Figure 4.2 Effect of scan rate on the anodic dissolution of chalcopyrite, pH 1.5, 2 5 ° C , de-aerated solutions 76 Figure 4.3 Effect of scan rate on the anodic dissolution of chalcopyrite, pH 1.5, 4 5 ° C , de-aerated solutions 76 Figure 4.4 Effect of scan rate on the anodic dissolution of chalcopyrite, pH 1.5, 6 5 ° C , de-aerated solutions 77 Figure 4.5 Effect of temperature on the anodic dissolution of chalcopyrite, pH 1.5, scan rate = 0.1 mV s _ 1 77 Figure 4.6 Constant potential experiments at 2 5 ° C on chalcopyrite in acidic solutions, pH 1.5 80 Figure 4.7 Anodic behaviour of fresh and polarized chalcopyrite surfaces in acidic solutions at 2 5 ° C , pH 1.5 80 Figure 4.8 Constant potential experiments at 4 5 ° C on chalcopyrite in acidic solutions, pH 1.5 81 Figure 4.9 Constant potential experiments at 6 5 ° C on chalcopyrite in acidic solutions, pH 1.5 81 Figure 4.10 Reduction of Fe(lll) on chalcopyrite as a function of applied potential at various concentrations of Fe(lll), pH 1.5, 2 5 ° C 84 Figure 4.11 Reduction of Fe(lll) on polarized chalcopyrite, pH 1.5, 0.001 M Fe(lll) 85 Figure 4.12 Reduction of Fe(lll) on chalcopyrite and pyrite, pH 1.5, 0.01 M Fe(lll) 86 Figure 4.13 Mixed potential of fresh and polarized chalcopyrite as a function of the Fe(lll):Fe(ll) ratio, pH 1.5, 2 5 ° C 89 Figure 4.14 Mixed potential of chalcopyrite and pyrite as a function of the concentration of Fe(lll) at constant Fe(ll), pH 1.5, 2 5 ° C 89 Figure 4.15 Leaching rate curves of chalcopyrite in ferric sulphate solution, pH 1.4, 3 5 ° C 92 Figure 4.16 H 2 0 2 added during leaching of chalcopyrite in ferric sulphate solution, pH 1.4, 3 5 ° C 92 Figure 4.17 Leaching rate curves of chalcopyrite in ferric sulphate solution, pH 1.4, 4 5 ° C 96 Figure 4.18 H 2 0 2 added during leaching of chalcopyrite in ferric sulphate solution, pH 1 . 4 , 4 5 ° C 97 Figure 4.19 Leaching rate curves of chalcopyrite in ferric sulphate solution, pH 1.4, 6 5 ° C 98 viii Figure 4.20 H2O2 added during leaching of chalcopyrite in ferric sulphate solution, pH 1.4, 6 5 ° C 98 Figure 4.21 Leaching curves of chalcopyrite residues rinsed and not rinsed in carbon disulphide solutions, p H 1.4, 6 5 ° C (R refers to the Fe(lll):Fe(ll) ratio) 101 Figure 4.22 Leaching rate curves of chalcopyrite in ferric sulphate solution at various temperatures, pH 1.4, Fe(lll):Fe(ll) = 1 102 Figure 4.23 Leaching rate curves of pyrite in ferric sulphate solution, pH 1.4, 6 5 ° C 104 Figure 4.24 Leaching rate curves of chalcopyrite at different F e S 2 : C u F e S 2 ratios, Fe(lll):Fe(ll) = 1, p H 1.4, 6 5 ° C 105 Figure 4.25 Schematic representation of the galvanic interaction between chalcopyrite and pyrite 107 Figure 4.26 Effect of increased F e S 2 : C u F e S 2 ratio on the relative anodic and cathodic areas 109 Figure 4.27 Leaching rate curves of chalcopyrite at different ratios of F e S 2 : C u F e S 2 , Fe(lll):Fe(ll) = 15, pH 1.4, 6 5 ° C 111 Figure 4.28 Leaching rate curves of chalcopyrite from various sources, F e S 2 : C u F e S 2 = 2:1, Fe(lll):Fe(ll) = 15, pH 1.4, 6 5 ° C 113 Figure 4.29 Electrochemical-passivation model for the leaching of chalcopyrite in ferric sulphate solutions, pH 1.4, 6 5 ° C 123 Figure 4.30 Electrochemical-passivation model for the leaching of chalcopyrite at different F e S 2 : C u F e S 2 ratios, Fe(lll):Fe(ll) = 1, pH 1.4, 6 5 ° C 124 Figure 4.31 Predicted values for the effect of increased F e S 2 : C u F e S 2 ratio on the leaching rate of chalcopyrite, Fe(lll):Fe(ll) = 1, pH 1.4, 6 5 ° C 125 Figure 4.32 Predicted values for the effect of particle size on the leaching of chalcopyrite, F e S 2 : C u F e S 2 = 2, Fe(lll):Fe(ll) = 1, pH 1.4, 6 5 ° C 126 Figure 4.33 Predicted values for the effect of increased F e S 2 : C u F e S 2 ratio and particle size on the leaching of chalcopyrite, Fe(lll):Fe(ll) = 1, pH 1.4, 6 5 ° C , ofpy = 25 pm, c / C p = 100 pm 127 Figure 4.34 Fraction of metal bioleached in the presence of mesophiles at 2 8 ° C .... 132 Figure 4.35 Molar ratio C u : F e during bioleaching in the presence of mesophiles at 2 8 ° C 132 Figure 4.36 Solution potential evolution and microbial counts during bioleaching in the presence of mesophiles at 2 8 ° C 134 Figure 4.37 Fraction of metal bioleached in the presence of moderate thermophiles a t 4 5 ° C 135 Figure 4.38 Molar ratio C u : F e during bioleaching in the presence of moderate thermophiles at 4 5 ° C 135 ix Figure 4.39 Solution potential evolution and microbial counts during bioleaching in the presence of moderate thermophiles at 45°C 136 Figure 4.40 Fraction of metal bioleached in the presence of extreme thermophiles at 68°C 137 Figure 4.41 Molar ratio Cu:Fe during bioleaching in the presence of extreme thermophiles at 68°C 138 Figure 4.42 Solution potential evolution and microbial counts during bioleaching in the presence of extreme thermophiles at 68°C 138 Figure 4.43 Fraction of copper bioleached in the presence of extreme thermophiles at 68°C with various initial concentrations of Fe(lll) and Fe(ll), total Fe = 1 g L~1 139 Figure 4.44 Fraction of copper bioleached in the presence of various microorganisms 141 Figure 4.45 Solution potentials during the bioleaching of chalcopyrite with various microorganisms 142 Figure 4.46 Schematic diagram of the bioleaching mechanism of chalcopyrite in the presence of thermophiles 143 Figure 4.47 Copper extraction in thermophilic bioleaching tests conducted in small columns with and without pyrite, 68°C 146 Figure 4.48 Solution potentials in thermophilic bioleaching tests conducted in small columns with and without pyrite, 68°C 147 Figure 4.49 Leaching rate curves of chalcopyrite at different ratios of FeS 2 :CuFeS 2 l Initial Fe(lll):Fe(ll) =1, 80°C 148 Figure 4.50 Solution potentials during leaching of chalcopyrite at different ratios of FeS 2 :CuFeS 2 , initial Fe(lll):Fe(ll) = 1, 80°C 149 Figure 4.51 Effect of initial potential on the leaching rate curves of chalcopyrite at different ratios of FeS 2 :CuFeS 2 , 80°C 150 Figure 4.52 Evolution of solution potentials during leaching of chalcopyrite at different initial potentials, 80°C 151 Figure 4.53 Effect of particle size on the leaching of chalcopyrite 151 Figure 4.54 Effect of particle size on solution potential during leaching of chalcopyrite 152 Figure 4.55 Effect of particle size of both chalcopyrite and pyrite on the leaching of chalcopyrite 153 Figure 4.56 Effect of particle size of both chalcopyrite and pyrite on solution potential during leaching of chalcopyrite 154 Figure 4.57 Effect of pyrite source on the leaching of chalcopyrite, FeS 2 :CuFeS 2 = 4, 80°C 155 x Figure 4.58 Effect of pyrite source on solution potential during leaching of chalcopyrite, F e S 2 : C u F e S 2 = 4, 8 0 ° C 156 Figure 4.59 Effect of pulp density on the leaching rate of chalcopyrite 157 Figure 4.60 Effect of pulp density on solution potential during leaching of chalcopyrite 158 Figure 4.61 Effects of impeller speed, choice of primary oxidant and acidity on the leaching rate of chalcopyrite 158 Figure 4.62 Effects of impeller speed, choice of primary oxidant and acidity on solution potential during leaching of chalcopyrite 159 Figure 4.63 Effect of pyrite addition, acidity and chalcopyrite origin on the leaching rate of chalcopyrite 160 xi L I S T O F T A B L E S Table 2.1 Electronic and structural properties of selected sulphide and oxide minerals 33 Table 2.2 Reactions of chalcopyrite 37 Table 2.3 Reaction rate orders for selected leaching systems 44 Table 2.4 Rest potentials of selected sulphide minerals in acidic solutions 48 Table 3.1 Detailed chemical analysis of the chalcopyrite mineral 53 Table 3.2 Mineralogical composition of chalcopyrite electrode 54 Table 3.3 Detailed chemical analysis of the Huanzala pyrite mineral 54 Table 3.4 Results of quantitative phase analysis (wt %) 66 Table 3.5 Chemical analysis of the Park City pyrite mineral 70 Table 4.1 Analysis of solutions after potentiostatic experiments on chalcopyrite 82 Table 4.2 Extraction data for the ferric leaching of chalcopyrite at 3 5 ° C 91 Table 4.3 Extraction data for the ferric leaching of chalcopyrite at 4 5 ° C 94 Table 4.4 Extraction data for the ferric leaching of chalcopyrite at 6 5 ° C 97 Table 4.5 Results obtained during the leaching of chalcopyrite and chalcopyrite residues after 4 hours of leaching at 6 5 ° C 100 Table 4.6 Leaching of chalcopyrite in the presence of pyrite at low solution potential 105 Table 4.7 Leaching of chalcopyrite in the presence of pyrite at high solution potential 110 Table 4.8 Chemical composition of chalcopyrite from various sources 112 Table 4.9 Input parameters for the electrochemical-passivation model 128 Table 4.10 Input parameters for the simulation of the electrochemical model, effect of pyrite addition 129 Table 4.11 Input parameters for the simulation of the electrochemical model, effect of particle size 130 Table 4.12 Calculated elemental sulphur yields from selected tests 161 Table A.1 Statistical analysis of copper extraction after 24 hours of controlled leaching with peroxide 180 xii L I S T O F S Y M B O L S E 0 Standard reduction potential, V E Oxidation-reduction potential, V P 8o Mesh of grid size that results in 80% passing SX/EW Solvent extraction and electrowinning SHE Standard hydrogen reference electrode SCE Standard calomel reference electrode Ag/AgCI Standard silver/silver chloride reference electrode En Redox potential vs. standard hydrogen electrode, V Ea Activation energy, kJ/mol/°K Ads Adsorbed species AAS Atomic absorption spectrophotometry k Rate constant for the oxidation of ferrous ions by H 2 0 2 , M _ 1 s a, Activity for species /' mi Molal concentration for species /, mol/kg H 2 0 Yi Activity coefficient for species /' /a Anodic current density, A cm - 2 /c Cathodic current density, A cm - 2 /0bs Sum of the anodic and cathodic cyrrent densities, A c m - 2 rpm Revolutions per minute, min - 1 o Sulphate yield /Acp Chalcopyrite surface, cm 2 >Apy Pyrite surface, cm 2 0 C p Fraction of the anodic area on chalcopyrite 0 P y Fraction of the anodic area on pyrite A Total surface during the galvanic interaction, cm 2 Aa,cp Anodic area on chalcopyrite, cm 2 Aa>py Anodic area on pyrite, cm 2 rr?cp Mass of chalcopyrite, g m P y Mass of pyrite, g p C p Density of chalcopyrite, g cm - 3 xiii P P y Density of pyrite, g cm-3 dcp Mean diameter of chalcopyrite particles, pm Mean diameter of pyrite particles, pm ]6 Mass ratio pyrite to chalcopyrite la Current for the anodic half-reaction, A Io Current for the cathodic half-reaction, A k Rate constant for the forward anodic reaction k Rate constant for the forward cathodic reaction k Rate constant for the reverse cathodic reaction Ofa Anodic transfer coefficient da Cathodic transfer coefficient Em Mixed potential rCp Dissolution rate of chalcopyrite, mol s~1 r>cP Number of moles of chalcopyrite ncp.o Initial number of moles of chalcopyrite Xcp Conversion of chalcopyrite r Radius of chalcopyrite particles, pm R Initial radius of chalcopyrite particle, pm kp Rate constant for the growth of the passive layer kcp Rate constant Kcp Rate constant \ xiv A C K N O W L E D G M E N T S I would like to express my sincere gratitude to my advisor, Professor David G . Dixon, for his guidance and support. Throughout the preparation of this Ph .D. thesis, his patience, personality, encouragement and financial support contributed to the successful completion of this program. Helpful suggestions from Professor David B. Dreisinger, Dr. Jochen Petersen and the members of my supervising committee are gratefully acknowledged. I would also like to thank the many people who assisted with the construction and operation of my experiments: Mr. Kodjo Afewu, Mr. Berny Rivera-Vasquez, Mr. Jackson So, and Mr. Kevin Tsai . I also thank my parents, brothers and sisters for their support, love and encouragement. Lastly, my deepest gratitude belongs to my wife Annie and our son Daniel, for their endless love, patience and understanding, in good and bad times, and it is to them that I dedicate this work. xv Chapter 1 INTRODUCTION Chalcopyrite ( C u F e S 2 ) is the most abundant copper bearing mineral in the world and is therefore extremely important to the copper industry [1]. Traditionally, this mineral has been treated by pyrometallurgical methods, i.e., in a copper smelter. The search for a hydrometallurgical process as an alternative to the conventional smelting to process copper sulphide minerals is driven by the desire to avoid SO2 emissions, the H2SO4 marketing problems associated with smelter operations and the selectivity of metal extraction. T h e s e advantages, cited for hydrometallurgical processes, are important factors in decision making as environmental regulations become more strict and the margin of profits in mineral processing decrease due to the lower grade of available ores. However, chalcopyrite is known to be particularly recalcitrant to hydrometallurgical processes [2-6]. Despite the relatively slow kinetics in both chemical and biological leaching reactions, interest in the recovery of copper from chalcopyrite concentrates remains strong because in many concentrates it is the sole source of copper. Most authors agree that the slow leaching kinetics is due to the formation of a film, which builds up on the surface of the mineral. The nature of this passivating film has been the subject of much controversy. One theory is that elemental sulphur produced during ferric leaching forms a compact and protective layer, which is believed to impede the transport of reactants and products to and from the chalcopyrite [3]. However, Buttinelli ef al. removed the sulphur with an organic solvent and failed to observe any effect on the rate [7]. A second theory proposes that two intermediate products, covellite and bomite, act as a "solid electrolyte interphase" between the unreacted mineral and the solution, causing passivation of chalcopyrite [8]. However, bornite and covellite are known to dissolve much more readily than chalcopyrite in sulphate media and are, therefore, unlikely to be the cause of chalcopyrite passivation [5,9]. A third theory suggests that the passivating layer is a copper-rich polysulphide, which forms as a result of solid state changes that occur in the mineral during leaching [10-12]. The exact nature of this polysulphide is complex and is the subject of considerable dispute. A Introduction 1 fourth hypothesis proposes that the passive layer is comprised of iron compounds [6]. However, it is considered doubtful that these compounds alone could be responsible for the slow dissolution kinetics [10,13]. These examples indicate that, although the dissolution of chalcopyrite in sulphuric acid has been investigated for many years, there are still doubts about the conditions under which the passivating layers are formed on the chalcopyrite surface. T he principal oxidation reaction of chalcopyrite in acid ferric sulphate solutions can be expressed by: C u F e S 2 + 2 F e 2 ( S 0 4 ) 3 -> C u S 0 4 + 5 F e S 0 4 + 2 S ° (1.1) where it is understood that some sulphate can be formed at the expense of elemental sulphur. The leaching rate and the mechanism of chalcopyrite passivation in ferric sulphate media can be conveniently studied using electrochemical techniques. The principle underlying this type of study is that, in general, the dissolution of semiconducting sulphide minerals in oxidizing solutions is a corrosion process in which a soluble oxidant is reduced at the mineral surface, which itself behaves as an anode [14]. Chalcopyrite is known to be a particularly good conductor with a resistivity of about 10~3 Q - m at 2 5 ° C [15]. The ferric leaching of chalcopyrite can then be interpreted electrochemically by considering reaction (1.1) to take place in two steps: Anodic: C u F e S 2 -> C u 2 + + F e 2 + + 2 S ° + 4 e" £ ° = 0.47 V ( S H E ) (1.2) Cathodic: 4 F e 3 + + 4 e" -> 4 F e 2 + E ° = 0.77 V ( S H E ) (1.3) Electrochemical studies might therefore hope to establish the mechanism and rate-determining steps of the anodic and cathodic reactions involved, the overall rate-controlling process, the factors affecting the passivation of the mineral, the potential dependence of the reaction mechanisms and product compositions. In order to obtain an adequate description of the kinetics of chalcopyrite leaching and to provide confirmation of the electrochemical results, it is essential to conduct experiments with finely ground minerals in a stirred tank reactor. Early studies Introduction 2 established the stoichiometry and thermodynamics of reaction (1.1). However, few studies have developed the rate expressions describing the leaching of chalcopyrite by an electrochemical mechanism. Furthermore, they have not attempted to propose a model that takes into account the passivation commonly observed during the dissolution of chalcopyrite. In an attempt to overcome the slow leaching rates observed during the aqueous oxidation of chalcopyrite, several solutions have been proposed. A m o n g them are: the use of microorganisms to help oxidize the passive layer [16-20]; an increase of temperature to decompose the passive layer [21]; the addition of A g ions, which changes the electrochemical behaviour of chalcopyrite [22] and the addition of pyrite to favour the galvanic interaction between the two minerals [23]. The use of microorganisms for the treatment of sulphide minerals has attracted a great deal of attention in recent years. Certain microorganisms are capable of catalyzing the oxidation of Fe(ll) to Fe(lll), thereby indirectly enhancing the oxidation of sulphide minerals. Microorganisms used in bioleaching are usually classified according to their temperature ranges for growth, so that the mesophilic microorganisms develop at ambient temperature ( 2 0 - 4 0 ° C ) while thermophilic microorganisms are more active at higher temperatures ( 4 0 - 9 0 ° C ) . Various authors have shown that chalcopyrite does not respond well to mesophilic microorganisms such as acidithiobacillus ferrooxidans and leptospirillum ferrooxidans [16-20]. In contrast, thermophilic microorganisms, such as sulfolobus and acidianus, have been shown to leach chalcopyrite considerably faster in the temperature range 6 5 - 7 5 ° C [16]. However, chalcopyrite still leaches slowly compared to the thermophilic bioleaching of other copper sulphide minerals [24]. Furthermore, copper dissolution from chalcopyrite typically stops after about 90% recovery, probably due to the formation of a passivating film [25]. According to various authors, the decrease in leaching rates is also caused by high solution potentials generated by the metabolism of the microbes. Since decreasing leaching rates were observed at high ferric concentration, it seems apparent that some change must be occurring at the surface of the chalcopyrite at these high potentials. Numerous authors have observed such changes in surface speciation during the ferric leaching of copper Introduction 3 sulphide minerals. For example, Koch and Mclntyre [26] observed the formation of metastable phases during the electrochemical oxidation of chalcocite to covellite. The following sequence was proposed: C U 2 S - » C U 1 . 9 3 S —» Cui . 8 3 S —» Cui.67S -> C U 1 . 4 S -> C u S This sequence progresses from states of lowest rest potential to states of highest rest potential. It is therefore possible that similar intermediate products, formed during the oxidative leaching of chalcopyrite, are also responsible for its slow dissolution kinetics. The interactions between chalcopyrite leaching rate, microbial type, Fe in solution, solution potential, temperature and surface modifications are closely interlinked and complex. The aim of this study is to improve our understanding of chalcopyrite passivation and depassivation during the oxidative dissolution of chalcopyrite. Electrochemical measurements using a chalcopyrite electrode, coupled with careful kinetic studies of finely ground mineral, may improve our understanding of chalcopyrite dissolution under oxidizing conditions in acidic solutions. Electrochemical studies should provide useful information about the conditions under which passivation of the mineral occurs and the potential dependence of the reaction mechanisms. The electrochemical behavior of chalcopyrite may then be validated in a stirred-tank reactor with finely ground mineral in acid ferric sulphate solutions at various solution potentials and temperatures. Leaching experiments may also be carried out to identify the role of microbes during mesophilic and thermophilic bioleaching of chalcopyrite. Finally, a model may be proposed which integrates the electrochemical process of chalcopyrite dissolution with the galvanic interaction with pyrite and the passivation observed during the leaching process. Introduction 4 Chapter 2 L I T E R A T U R E R E V I E W 2.1. Properties of chalcopyrite The crystal structure of chalcopyrite is derived from the zinc blende (ZnS) structure, in which copper and iron replace the two zinc atoms, as illustrated in Figure 2.1. The average bond distances for S - S , F e - S and C u - S are 0.368, 0.225 and 0.230 nm respectively. T h e bonding in chalcopyrite is considered to be essentially covalent, where each atom is bonded to its four nearest neighbors. Lattice constants (in angstroms) for the unit c e l l a r e a = 5.2988 and c =..10.434 [27]. Chalcopyrite is an n-type semiconductor, due to the metal excess, with a band gap of about 0.6 eV. Figure 2.1. Crystal structure of chalcopyrite Various authors have explained the dissolution of sulphide minerals based on the mineral structure and valence bond theories. In semiconductors with a band gap greater than 1 eV, anodic dissolution is believed to occur almost entirely as a result of hole injection into the valence band [15]. In the case of chalcopyrite, which has a band gap of 0.6 eV, it is proposed that both holes and electrons contribute to the anodic dissolution. The following sequence is proposed [15]: © Cu Fe Literature Review 5 C u F e S 2 + 3 h + -> F e 3 + + . C u S 2 and . C u S 2 + 2 h + -> C u 2 + + 2 S (2.1) (2.2) or . C u S 2 -> C u 2 + + 2 S + 2 e" (2.3) where . C u S 2 is an unstable radical intermediate. The first step indicates the breaking of the F e - S bond and the second, the breaking of the C u - S bond. This mechanism implies that copper remains in the lattice in the initial stages of leaching. This is consistent with the observation that the F e / C u ratio in solution is greater than one in the earlier stages of leaching. S o m e aspects of chalcopyrite electrochemistry will be reviewed in details in section 2.4.4. 2.2. Hydrometallurgical processes for copper sulphide minerals in sulphate media Leaching of copper sulphide minerals in sulphate media has received extensive study. The sulphate-based processes have some potential advantages over the others in that the leaching chemistry is generally simpler and the technology to recover high purity copper is well established. Unfortunately, chalcopyrite is known to dissolve extremely slowly in sulphate media. For example, Wadsworth [9] reported that chalcocite dissolves 20 times faster than chalcopyrite in oxygenated acid solutions at 2 5 ° C and pH 1.25. Regardless of the slow leaching rates, there is significant new developmental activity in hydrometallurgical processes for the treatment of copper sulphide minerals. These processes can be grouped into three major categories: (1) ferric sulphate leaching, (2) oxygen pressure leaching, (3) bioleaching. This section provides an overview of these methods. 2.2.1. Mt Gordon process This process is based on the hot acid ferric sulphate dissolution of copper sulphide minerals, mainly chalcocite with accessory covellite and minor chalcopyrite [28]. Generally chalcocite accounts for over 90% of the total copper present. The process Literature Review 6 uses a moderate feed grind of 75-106 um. T h e effect of grind size was investigated and leach recovery was insensitive to grind size below Pso = 120 um: T h e process operates in an autoclave below the boiling point ( ~ 9 5 ° C ) at a pressure of 8 atm with a residence time of 60 minutes. Chalcocite leaches in two steps in ferric sulphate solutions: Chalcocite leaching - Stage 1: C u 2 S + F e 2 ( S 0 4 ) 3 - » C u S 0 4 + C u S + 2 F e S 0 4 (2.4) Chalcocite leaching - Stage 2: C u S + F e 2 ( S 0 4 ) 3 -> C u S 0 4 + S + 2 F e S 0 4 (2.5) Ferric sulphate regeneration: 4 F e S 0 4 + 0 2 + 2 H 2 S 0 4 -> 2 F e 2 ( S 0 4 ) 3 + 2 H 2 0 (2.6) Overall chalcocite leaching reaction: C u 2 S + 0 2 + 2 H 2 S 0 4 -> 2 C u S 0 4 + S + 2 H 2 0 (2.7) The rate of ferric generation was found to be second-order with respect to ferrous sulphate concentration and first order with respect to oxygen partial pressure. The process uses abundant pyrite to resupply acid and iron lost from the leach circuit. Pyrite oxidation: F e S 2 + 3.5 0 2 + H 2 0 -> F e S 0 4 + H 2 S 0 4 (2.8) Typical iron solution concentration is 35 g/L Fe and sulphuric acid concentration is maintained at 10-15 g/L in the autoclave discharge. Leach recovery is about 92-93%. T h e elemental sulphur formed inside the autoclave remains relatively unattacked. Copper recovery is by conventional solvent extraction and electrowinning ( S X / E W ) . The electrolyte sent to electrowinning contains 48 g/L C u . This process was operated Literature Review 7 successfully at Western Metals' Mt Gordon operation in Queensland, producing about 50,000 tpy of high grade cathode copper. 2.2.2. Activox process Activox is a leaching process operating at about 1 1 0 ° C and oxygen pressure of 10 atm for the treatment of sulphide concentrates containing copper, nickel, cobalt, zinc and gold. Activox consists of activation of the mineral surface by ultrafine grinding ( P 8 0 = 7 -10 pm) [28]. Decreasing the initial mineral particle size gives rise to an increase in the surface areas for leaching with a corresponding increase in the copper leached. T h e disintegration of the matrix by high energy grinding is also called "mechanical activation" [29]. Mechanical activation increases internal and surface energies, increases surface area, and decreases the coherence energy of solids. Base metals are extracted into the leach solution, while precious metals remain in the leach residue. If insufficient acid is present during the Activox leach, copper may precipitate as a basic copper sulphate in the leach residue. This copper can be readily recovered by acid washing the leach residue. Sulphur reports largely as elemental sulphur and remains in the leaching residue. Most of the iron precipitates selectively as hematite, goethite or jarosite. Laboratory tests indicated that Activox was suitable to treat chalcopyrite minerals with high recoveries (about 95% conversion). Copper is recovered by conventional S X / E W . Western Minerals Technology in Perth has successfully run pilot scale campaigns. 2.2.3. MIM/Highlands Albion (Nenatech) process T h e Nenatech process uses ferric leaching under atmospheric pressure, with oxygen or air sparging at temperatures of about 8 0 ° C [30]. This process also uses ultrafine grinding to reduce the feed size to about 16-18 pm. Elemental sulphur formed in the process remains in the leaching residue, leading to difficulty in recovering precious metals. Copper recovery is by conventional S X / E W . T h e process has been run successfully at the pilot plant scale. MIM and Highlands Pacific jointly own this process. Literature Review 8 2.2.4. AAC/UBC Hydrometallurgy process Anglo American Corporation and The University of British Columbia have developed a process for chalcopyrite leaching under moderate oxygen pressure (10-12 atm) at 1 5 0 ° C [30]. Leaching of chalcopyrite at 1 5 0 ° C in the presence of elemental sulphur requires process conditions that prevent coating of unreacted chalcopyrite by liquid sulphur and avoidance of passivation due to the presence of a copper sulphide layer [10]. It was observed that the viscosity of the sulphur decreases from 115 to 1 5 9 ° C , but increases dramatically between 159 and 1 9 0 ° C . It is therefore essential to work at temperatures below 1 5 9 ° C . The process uses surfactant additives such as calcium lignosulfonate and ortho-phenylenediamine (OPD) to disperse the molten sulphur. The feed is finely ground (Pso = 5-20 um). Copper and zinc extractions of 95% are achieved at 1 5 0 ° C in two hours of residence time. For coarser particle sizes, the leach ceases at about 80% copper extraction due to passivation. The degree of sulphur oxidation is relatively high. T h e range of elemental sulphur yields is 59-80%. Sulphide sulphur in chalcopyrite may be oxidized to elemental sulphur or through to sulphate sulphur according to the following reactions: Formation of hematite and elemental sulphur: C u F e S 2 + 1.25 0 2 + H 2 S 0 4 -> C u S 0 4 + 0.5 F e 2 0 3 + 2 S + H 2 0 (2.9) Formation of hematite and sulphate: C u F e S 2 + 4.25 0 2 + H 2 0 C u S 0 4 + 0.5 F e 2 0 3 + H 2 S 0 4 (2.10) With oxygen as a reaction product, the oxygen consumption is 1.25 mol/mol C u ; with sulphate as a reaction product, the oxygen consumption is 4.25 mol/mol C u . Therefore, high elemental sulphur yields (80-100%) are desirable to reduce the oxygen consumption. Copper recovery is by conventional S X / E W . Precious metals in the ore can be recovered by cyanidation. A pilot plant program is in progress at the A A R L laboratory in South Africa. Literature Review 9 2.2.5. Dynatec process The Dynatec process was first introduced for the treatment of zinc concentrates and refractory gold ores. Dynatec zinc pressure leaching operates at 1 5 0 ° C using oxygen, with dissolution of about 98% zinc and conversion of most of the sulphur to the elemental form. In the case of refractory gold ores, more aggressive conditions (generally between 190 and 2 2 5 ° C ) are used to dissolve the pyrite and/or arsenopyrite matrix. A Dynatec process for chalcopyrite was recently introduced and involves oxidative leaching at 1 5 0 ° C under a pressure of 10-12 atm [30,31]. T h e process uses low-grade coal as a surfactant additive to disperse molten sulphur. The feed is ground to 30-40 pm. Experimental results indicate that the first leaching step dissolves about 80% copper, and the unreacted sulphide is recovered by flotation. The removal of elemental sulphur is done by melting and filtration. After recycling of the sulphide to the leach, a high extraction of copper (98%) is readily attainable. T h e precious metals from leaching residue are treated by cyanidation. Iron is precipitated as hematite and jarosite. Copper recovery is by conventional S X / E W . 2.2.6. Total Pressure Oxidation process This process is owned by Placer Dome and uses high temperature oxidation conditions ( 2 2 0 ° C ) and high pressure (30-40 atm) to achieve high copper extractions. The process was first used for the treatment of refractory gold ores. Pyrite and arsenopyrite were oxidized with molecular oxygen at temperatures of 1 8 0 - 2 3 5 ° C to liberate sulfide-encapsulated gold. Gold was recovered from the oxidized residue by cyanidation. Copper sulphide minerals are similarly decomposed by high temperature oxidation conditions. This process is characterized by rapid and complete copper leaching, a high degree of iron hydrolysis and impurity fixation. The oxygen consumption is high due to the total conversion of sulphur to sulphate. Iron is precipitated mainly as hematite. Gold and silver are recovered by cyanidation of the washed residue. The total oxidation process also achieves a high extraction of copper (95%) in two hours of residence time. Copper recovery is by conventional S X / E W . A demonstration plant treating 136 Literature Review 10 tonnes/day of concentrate has been constructed in Bagdad, Arizona by Phelps Dodge [28,30,31]. 2.2.7. CESL process T h e C E S L process, owned by Cominco Engineering Services Limited of Vancouver, B C , uses mixed sulphate-chloride solutions for pressure oxidation. It operates at about 1 5 0 ° C with chloride ions playing the role of catalyst. The concentrate is reground to about 40 um and leached under moderate pressure (10-12 atm). Copper is converted to a basic copper sulphate salt, iron to hematite, and sulphur to elemental sulphur. Following liquid-solid separation, basic copper precipitates are re-dissolved in an atmospheric leach step using the acidic raffinate recycled from S X / E W . The resultant residue, comprising elemental sulphur, a small amount of unreacted sulphide and hematite, is floated to recover a sulphur-sulphide fraction for removal of elemental sulphur with an organic solvent. The residue of the organic wash is recombined with the hematite for further leaching at 2 2 0 ° C to maximize the extraction and to produce a residue amenable to precious metals recovery by cyanidation. Copper is recovered by conventional S X / E W from a predominantly sulphate solution [28, 30,31]. 2.2.8. BacTech/Mintek process Mintek and B a c T e c h have developed a bioleaching process for the treatment of various copper concentrates [31]. The feed is moderately ground to a Pso of 75 um. For chalcocite concentrates, mesophilic cultures are used and the leaching and ferrous biooxidation are carried out in the same tank. For chalcopyrite concentrates, moderate to extreme thermophilic cultures are used and the leaching and ferrous biooxidation are done in separate tanks. Regrinding of concentrate to a Pso of 10 um is also required for chalcopyrite bioleaching. The residence time is 5 to 6 days for complete chalcopyrite oxidation. T h e sulphur is converted to elemental sulphur and sulphate, the sulphate is precipitated from the bleed stream as gypsum and soluble iron is precipitated as hydronium jarosite. Gold and silver are recovered from the leach residue by cyanidation. B a c T e c h and Mintek have now jointly developed their tank bioleaching process up to demonstration plant level at the Penoles operation in Monterrey, Mexico. This Literature Review 11 demonstration plant is designed to produce 500 kg/day of cathode copper by conventional S X / E W [31]. 2.2.9. BioCop process B H P Billiton and Codelco have joined forces to develop a new process called BioCop for the bacterial leaching of copper sulphide minerals [32]. In this process, mesophiles are used to leach chalcocite at 4 2 - 4 5 ° C , while thermophiles leach chalcopyrite at 7 0 -8 0 ° C . (Mesophiles give unacceptably low copper extraction unless the chalcopyrite is finely ground prior to bioleaching.) The oxygen consumption is high due to the complete conversion of sulphur to sulphate. The BioCop process is therefore similar to the Total Pressure Oxidation process described above in that leaching is conducted under total oxidation conditions (i.e. to sulphate). B H P Billiton and Codelco have formed a joint venture called Alliance Copper and built a pilot plant capable of producing 20,000 tpy of copper cathode. The pilot plant is located at the Chuquicamata mine in northern Chile. 2.2.10. GeoCoat process The G e o C o a t process involves the coating of concentrates onto sized support rocks -either barren rock or low-grade ore - then heap bioleaching the coated material [33]. The heap is inoculated with suitable cultures and irrigated with acidic solutions containing iron and nutrients, while ambient air is applied at the heap base. Temperatures of 6 0 - 7 0 ° C are attained within the heap to facilitate rapid copper recovery. The process was developed to solve two major problems in heap leaching: obstruction of liquid and air flow by fine particles of ore, and limited exposure of sulphide particles to air, leach solution and microbes. GeoBiotics of Lakewood, Colorado is developing the process. T h e technology was initially developed for gold recovery, and an extensive program is being carried out for copper sulphides. Plans are now underway for the first large-scale field test. Literature Review 12 2.2.11. Summary The challenge of process development for chalcopyrite leaching is to dissolve copper quickly and completely with a high yield of elemental sulphur. T o date, none of the processes mentioned have achieved sustained commercial production above the demonstration plant scale. The main problems hindering commercial application of these processes are poor copper recoveries from chalcopyrite, high reagent costs (e.g. oxygen and/or limestone), and the need for special reagents and/or process conditions (autoclaves, fine grinding, surfactants, microbes) to encourage chalcopyrite dissolution which render many of these processes uneconomical relative to toll smelting. T o overcome the slow and incomplete leaching of chalcopyrite in sulphate media, two major considerations should be taken into account: the extent to which the reaction will proceed, and the rate at which the reaction will proceed. T h e former depends on the thermodynamic tendency of the chemical system, which determines the overall reaction driving force. This may be conveniently followed on potential-pH diagrams. Th e latter, the reaction kinetics, depend on a combination of physical, chemical and mass transfer factors. 2.3. Thermodynamic considerations The thermodynamic stability diagram of compounds in the C u - F e - S - H 2 0 system is shown in Figure 2.2. Since the mixed potentials for most common oxidants for the leaching of chalcopyrite are between 0.45 and 0.85 V (SHE) [34], our attention will be focused within this potential range. Literature Review 13 O 2 4 6 8 IO 12 14 pH Figure 2.2. Potential-pH diagram of the C u - F e - S - H 2 0 system at 2 5 ° C : all solutes at 0.1 M activity except C u 2 + at 0.01 M [35] According to the diagram, as the positive values of potential on the surface of chalcopyrite increase, a number of successive oxidation reactions will occur in the acid region (pH = 1 to 2): Formation of bornite and pyrite: 5 C u F e S 2 + 2 H 2 S + 5 Ox -> C u 5 F e S 4 + 4 F e S 2 + 4 H + + 4 Ox" (2.11) Formation of chalcocite and pyrite: 2 C u F e S 2 + H 2 S + 2 Ox -> C u 2 S + 2 F e S 2 + 2 H + + 2 Ox" (2.12) Formation of covellite and pyrite C u F e S 2 + H 2 S + 2 Ox -+ C u S + F e S 2 + 2 H + + 2 Ox" (2.13) Literature Review 14 Formation of covellite and elemental sulphur: C u F e S 2 + 2 O x -> C u S + F e 2 + + S ° + 2 Ox" (2.14) It follows that the solid compounds which may passivate the surface of chalcopyrite are pyrite, bornite, chalcocite, covellite, and sulphur. From thermodynamic calculations (En-p H diagrams) these products dissolve in acidic medium at potentials higher than 0.5 V (SHE) . Therefore, maintaining the potential in the system at 0.65 V (SHE) would be sufficient to dissolve chalcopyrite at a reasonable rate. Unfortunately, that is not the case in practice. Disagreements between predictions made using the potential-pH diagram of the Cu-Fe -S -H20 system and experimental results may result from the fact that the oxidation reactions (2.11) to (2.13) require H 2 S as a reactant, while F e S 2 is a product. If H 2 S is depleted in solution during the leaching process and not supplied as reactant, these reactions will not take place. Even if H 2 S is available, pyrite does not nucleate and grow at a measurable rate [35]. Another problem inherent with the direct application of potential-pH diagrams to sulphide leaching systems is the possible formation of metastable intermediate phases. These phases may be attributed to slow chemical kinetics associated with the products of reaction. There is no way to predict the presence of metastable phases from thermodynamics considerations alone. It is therefore appropriate to review the kinetic aspects of chalcopyrite leaching in sulfate media in order to have a better understanding of the mechanisms of chalcopyrite passivation. 2.4. Retarding effect on the dissolution of chalcopyrite T he formation of passive layers on the chalcopyrite surface is a phenomenon observed in acidic solutions in which an external current is applied to dissolve the mineral and in sulfate solutions containing oxidizing reactants such as ferric, oxygen under pressure and bacteria. Since the mechanisms for the formation of passive layers depend on the technique used to dissolve the mineral, it is convenient to discuss separately the various cases of chalcopyrite passivation. For the present work, much attention will be given to the ferric, microbial and electrochemical leaching of chalcopyrite. Literature Review 15 2.4.1. Passivation of chalcopyrite during ferric leaching Th e first systematic study of the dissolution of copper sulphides by dilute sulphuric acid in the presence of ferric sulphate can be credited to J . Sullivan during the late 1920's and early 1930's [36]. Fine grinding was considered necessary because of the refractory nature of the mineral used. Only 33% of the copper from a sample ground to - 4 0 pm was leached after 57 days at 3 5 ° C , using a 1% ferric sulphate solution. T h e chalcopyrite leaching reaction was initially rapid but soon slowed! Many explanations have been presented to justify the rapid decrease in copper leaching rate, but there are still serious points of controversy about the role of ferric during the leaching process. Th e reactions by which chalcopyrite dissolves are believed to be: C u F e S 2 + 4 F e 3 + -> C u 2 + + 5 F e 2 + + 2 S ° (2.15) C u F e S 2 + 4 F e 3 + + 2 H 2 0 + 3 0 2 -> C u 2 + + 5 F e 2 + + 2 H 2 S 0 4 (2.16) T h e variables found to influence the ferric leaching of chalcopyrite are: ferric sulphate concentration, formation of elemental sulphur and precipitation of iron compounds. Ferric sulphate concentration The most logical place to begin is to evaluate the effects of reactants since limitations on reactant availability would be the simplest explanation for the rapid decrease in copper leaching rate. It has been reported consistently [4,5] that increasing concentration of ferric does not result in a corresponding increase in the leaching rate. Hirato ef al. [37] observed that the leaching rate of chalcopyrite increased with an increase in the ferric sulphate concentration up to 0.1 M Fe 2 (S04)3. showing a first-order dependency, but declined at a higher concentration. Leaching experiments were conducted at 7 0 ° C with 3 different specimens. Jones and Peters [38] observed similar results. Increasing the concentration from 0.03 to 0.1 M Fe 2 (S04 )3 increased the leaching rate slightly, but a further increase to 1.0 M Fe 2 (S04 )3 yielded a leaching rate lower than that found at 0.03 M . One explanation for the decreasing copper leaching rate would be depletion of ferric in a boundary layer close to the chalcopyrite surface. Literature Review 16 Such an explanation is unlikely, however, since electrochemical measurements using chalcopyrite rotating electrodes have demonstrated no dependence between reaction rate and rotational speed [39], which is an indication of a lack of reaction control by mass transfer. Since decreasing leaching rates were observed above a ferric concentration of 0.1 M , it seems apparent that some factor other than ferric concentration becomes rate controlling after an initial leaching phase. The slow kinetics at high ferric concentration may also be attributed to the distribution of iron species in the solution. The speciation for the aqueous solution system H2SO4-Fe2(SC>4)3 was examined by Sapieszko et al. [40] under acidic leaching conditions. It was found that the principal Fe(lll) species in this system are F e 3 + , F e S 0 4 + and F e H S 0 4 2 + . By using these data, Hirato and co-workers [37] plotted the distribution of iron species as a function of F e 2 ( S 0 4 ) 3 concentrations in 0.2 M H 2 S 0 4 solutions (Figure 2.3). The rapid decline in the concentration of F e 3 + and F e H S 0 4 2 + at higher F e 2 ( S 0 4 ) 3 concentrations is evident in this figure. For example, for a thirty-fold increase in Fe2(S0 4 ) 3 , the ferric activity increases by a factor of 6. By contrast, the F e S 0 4 + concentration increases linearly. They concluded that the F e 3 + and/or F e H S 0 4 2 + ions play a significant role in the ferric leaching of chalcopyrite. Crundwell [41] obtained similar results during the leaching of sphalerite in acid ferric sulphate solutions at 2 5 ° C . He observed that the rate of sphalerite dissolution was proportional to the sum of the concentration of the F e 3 + and F e H S 0 4 2 + species in solution. Literature Review 1 7 10"' 10"' ? E •D O E .- 10"! u " 10"3 10"2 10"1 I C(Fe(S0 4) l 5)/ moldm-3 Figure 2.3. Speciation of Fe(lll)-sulfato and bisulfato complexes in 0.2 M H2SO4 solutions as a function of total Fe(S04)i.5 concentration [37] Another concern, besides speciation, arises from recent reports describing the effects of ferrous on chalcopyrite oxidation with ferric. Kametani and Aoki [42] reported that ferric concentration had little effect on the oxidation of chalcopyrite within the temperature range of 50 to 9 0 ° C . They suggested that copper extraction was mainly controlled by the redox potential of the solution, a function of the ferric:ferrous ratio, and that there was a critical potential at which the leaching rate is maximum. It was observed that the amount of copper released during chalcopyrite leaching was much lower than that of iron, particularly at low potentials - 0.3 to 0.33 V ( S C E ) - indicating the formation of a copper-rich intermediate. X-ray diffraction examination of the residue revealed the formation of C u S on the surface of chalcopyrite. In the potential range 0.37 to 0.43 V , the oxidation of chalcopyrite reached completion and equal amounts of C u and Fe were released into solution. T h e maximum rate of chalcopyrite oxidation was attained only over this narrow range of solution potential. Above 0.45 V there was a sudden decrease in the leaching rate of chalcopyrite. This critical potential corresponded with the start of the oxidation of pyrite present in the concentrate. It is possible that pyrite was providing Literature Review 18 the cathodic sites for the ferric reduction at low potentials. Kametani and Aoki also reported an induction period for the leaching of chalcopyrite. T h e induction period decreased with increasing temperature as shown in Figure 2.4. Figure 2.4. Variation of reaction curves with temperature [42] Fascinatingly, Hiroyoshi and co-workers [4] observed that ferrous sulphate extracted copper more efficiently than ferric sulphate. The amount of extracted copper was larger in the presence of an adequate amount of ferrous than in the absence of ferrous. However, the results indicated that ferrous alone does not leach chalcopyrite and that oxygen is needed to extract copper from chalcopyrite. Hiroyoshi and co-workers believe that the ferrous-promoted chalcopyrite leaching is an electrochemical phenomenon where ferrous acts as a reductant. T h e y proposed a model consisting of two steps: In the first step, chalcopyrite is reduced by ferrous to form CU2S in the presence of cupric as follows: C u F e S 2 + 3 C u 2 + + 3 Fe : ,2+ 2 C u 2 S + 4 Fe 3+ (2.17) Next, C u 2 S is oxidized by dissolved oxygen and/or ferric to form cupric: 2 C u 2 S + 8 H + + 2 0 2 -> 4 C u 2 + + 2 S ° + 4 H 2 0 (2.18) and/or Literature Review 19 2 C u 2 S + 8 F e 3 + -> 4 C u 2 + + 8 F e 2 + + 2 S ° (2.19) Following the proposed model, enhanced copper extraction by ferrous occurs only at lower potentials when enough C u 2 + ions are present in solution. T h e intermediate C u 2 S does not form without cupric or at higher potentials (above 0.7 V (SHE)). This might explain the decrease in the leaching rate above Kametani's critical potential. By contrast, Dutrizac [1] observed that small quantities of F e S 0 4 (0.1 M F e S 0 4 ) , added to a 0.1 M F e 2 ( S 0 4 ) 3 + 0.3 M H 2 S 0 4 solution, could cause the rate of chalcopyrite leaching to fall appreciably (by approximately 30%) at 8 0 ° C . In addition, from a theoretical point of view, an oxidizing agent is required because the non-oxidative dissolution of chalcopyrite cannot be sustained on a thermodynamic basis [5]. This raises some concerns about the role of ferrous during the ferric leaching of chalcopyrite. In view of the above discrepancy, the role of ferrous and ferric in chalcopyrite leaching is worthy of further study. Formation of elemental sulphur Another explanation for the slow kinetics would be the formation of reaction products on the surface of chalcopyrite. T h e product that has received the most attention is elemental sulphur. Dutrizac [3] noted that elemental sulphur is the major sulphidic reaction product of ferric sulphate attack on chalcopyrite, but some sulphate is also formed. He reported the formation of 94% elemental sulfur and less than 6% sulphate, regardless of the leaching time (0-70 h), the ferric concentration (0-2 M F e 3 + ) or the chalcopyrite particle size. Experiments were conducted at 9 5 ° C in acidic ferric sulphate solution. T h e E/ , -pH diagram of the sulphur-water system indicates that increasing acid concentration and increasing sulphate concentration result in increasing the area of stability of elemental sulphur (Figure 2.5). Literature Review 20 1.0 05 5 g 0 I xz L U -05 " 0 4 7 10 14 pH Figure 2.5. Elemental sulphur-water E/,-pH diagram with extended sulfur stability [35] Elemental sulphur enjoys an extraordinary stability under a wide range of conditions and at temperatures at least as high as its melting point [35]. T h e formation of a sulphur layer on the chalcopyrite surface may significantly influence the reaction kinetics by providing a diffusion barrier. Several authors observed parabolic kinetics with respect to ferric concentration and concluded that this was indicative of diffusion control through the elemental sulphur layer. Linge [43] contests this view using as evidence measurements of ferric diffusion through sulphur films that are four times higher than the leaching rate of chalcopyrite. The diffusion of electrons across a sulphur film as the rate-limiting step has also been considered [44]. Leaching chalcopyrite in the presence of finely ground graphite tested this assumption. T h e presence of graphite embedded in the insulating suphur layer would improve electron transport, since graphite is conductive. T h e leaching of chalcopyrite was enhanced, providing support to the electron transport limitation theory. However, some aspects of this interpretation of rate control through elemental sulphur are not satisfactorily explained. T h e most consistent characteristic of chalcopyrite leaching is the high activation energy associated with the reaction, ranging from 70 to 84 kJ/mol [42-45]. T h e s e large activation energies are v HS0,- so4- ^ v Extended Sulphur >>. y Zone so,-(so,-) 1 ., H2S 1 atm. * » » * HS" 1m. is-i i "—1 Literature Review 21 normally not observed for rate control by pore transport processes in other hydrometallurgical processes. Typically, the following may be assumed: > Ea < 20 kJ/mol: mass transfer control > 20 < Ea < 40 kJ/mol: mixed control > E a > 40 kJ/mol: surface reaction control Another concern is the form of the rate equation. A number of publications have reported that the leaching time may influence the form of the rate equation. Dutrizac et al. [45] reported parabolic kinetics for leaching experiments completed within 100 hours (Figure 2.6). Jones and Peters [38], who used a leaching time of over 55 days, found a linear kinetics for the leaching of chalcopyrite in ferric media. Hirato et co-workers [37] proposed a two-stage kinetics for the ferric leaching of chalcopyrite at 7 0 ° C ; they found parabolic kinetics during the initial stage (0-100 h) and linear kinetics over extended leaching periods (Figure 2.7). The first stage is characterized by the diffusion of reactants and products through the elemental sulfur layer. This layer is porous in the initial stage of leaching and the leachant can easily migrate through it. For this stage the rate can be expressed by parabolic kinetics. However, the sulphur layer becomes thicker and more compact with an increase in leaching time. For example, Dutrizac [1,3] observed that the sulphur layer formed in sulphate media is considerably more dense and compact than that formed in chloride media and that the sulphur coverage was complete on the chalcopyrite surface after 3 hours of leaching. There was no evidence of any macroscopic porosity or cracks, which would allow easy solution access to the surface of the remaining core of CuFeS2 . This would have affected the leaching rate of chalcopyrite during the second stage, which is characterized by linear kinetics. However, it is observed in practice that chalcopyrite still leaches even after the formation of a compact elemental sulphur layer on its surface [37]. T h e reason is not clear, but Peters [48,49] believes that if the oxidation of one mole of chalcopyrite (42.65 cm 3 ) produces 100% elemental sulphur (31 cm 3 ) , it will result in a negative volume change of 27%. This difference in molar volume between sulphur and chalcopyrite would cause sulphur films to crack so that uniform coatings could not form. Literature Review 22 Similar conclusions were reached by Hirato and co-workers [37]. T h e y reported that geometrical factors such as roughness of the surface affect the leaching rate. The dense and compact layer of elemental sulphur peels off at random during the second stage of leaching characterized by linear kinetics. A s a result, the effective reacting surface is kept almost constant, and the leaching follows linear kinetics. However, it is important to note that linear kinetics was mainly observed on polished surfaces such as chalcopyrite electrodes cut from a massive piece. Based on this observation, Dutrizac [3] tried to examine the morphology of the sulphur formed on finely ground chalcopyrite (10-14 um) where extensive dissolution occurs in a relatively short time at 9 5 ° C . The individual CuFeS2 particles were extensively coated after 4 hours of leaching by elemental sulphur, and agglomerated into significantly larger masses via sulphur bridges. About 25% of the chalcopyrite was dissolved during that time. After 17 hours of leaching, 60% copper extraction occurred and the leach residue consisted of agglomerated sulphur globules with cores of residual chalcopyrite. After 72 hours of leaching, 90% C u extraction was reported with corresponding amounts of elemental sulphur formation. T h e particles did not develop roughened surfaces during leaching, and this fact coupled with the extensive sulphur coverage is likely responsible for the observed non-linear ("parabolic") leaching kinetics. A n additional concern is published data showing that chalcopyrite still leaches slowly even if elemental sulphur is removed during the leach with an organic solvent [7]. This observation suggests that elemental sulphur is not the only product responsible for the slow kinetics observed during the ferric sulphate leaching of chalcopyrite. The formation of an intermediate sulphide phase beneath the sulphur layer might be a contributing factor for the slow kinetics of chalcopyrite leaching [3]. Several authors tried to identify the passive species beneath the sulphur layer but were not successful. They may have failed because the passive species are metastable intermediates which form extremely thin layers (3 nm) [3]. Literature Review 23 0 2 M F « ( S 0 4 ) | S . O-.S" HZ$°4- ?°*C 0 766 TRANSVAAL CHALCOPYRITE TIME IH) Figure 2.6. Leaching curves for the dissolution of various size fractions of natural chalcopyrite in sulfate solutions [1] Time/ h Figure 2.7. Leaching rate curve of chalcopyrite with ferric sulfate [37] Literature Review 24 Precipitation of iron compounds Another passivation theory has it that the impermeable layer is composed of iron compounds [6, 50]. W h e n the sulphuric acid in the leaching system has been consumed due to the continued formation of copper and iron sulphate, the ferric sulphate begins to hydrolyze according to the following reaction: F e 2 ( S 0 4 ) 3 + 6 H 2 0 ^ 2 F e ( O H ) 3 + 3 H 2 S 0 4 (2.20) In addition to the formation of F e ( O H ) 3 , other iron compounds such as ferric hydrosulphate, goethite and jarosite have also been identified in leach residues. The reactions accompanying formation these compounds are as follows: Ferric hydroxysulphates: F e 2 ( S 0 4 ) 3 + 2 H 2 0 -> 2 F e ( O H ) S 0 4 + H 2 S 0 4 (2.21) Goethite: F e 2 ( S 0 4 ) 3 + 4 H 2 0 2 F e O O H + 3 H 2 S 0 4 (2.22) Jarosite: 3 F e 2 ( S 0 4 ) 3 + 1 2 H 2 0 + 2 A + -> 2 A F e 3 ( S 0 4 ) 2 ( O H ) 6 + 5 H 2 S 0 4 + 2 H + (2.23) where A represents K + , N a + , N H 4 \ H 3 0 + , A g + or (Pb 2 + ) 0 . 5 -Figure 2.8 presents the predominant iron precipitates as a function of pH and temperature. Ferric is mainly present in solution at pH values below 2. T h e precipitation of jarosite starts at around pH 2 in the presence of suitable monovalent cations. Jarosite formation cannot be avoided during ferric leaching that operates within the temperature range 20 to 1 0 0 ° C . At p H values above 3, complete ferric precipitation occurs and forms highly insoluble iron precipitates (ferric hydroxide, goethite, hematite). Hematite forms only at temperatures higher than 1 0 0 ° C . Literature Review 25 O L^JL 1 , ?„„; . I 1 , j . £>Hf Figure 2.8. Stability of various iron precipitates as a function of pH and temperature [51] T h e formation of jarosite and other iron-hydroxy precipitates may retard chalcopyrite leaching by restricting the mass transfer of ions to solution and by preventing Fe(lll) access to the mineral surface. Stott and co-workers [6] observed that extensive jarosite formation was detected on the surface of chalcopyrite during ferric and microbial leaching. While it was thought that this layer prevented further leaching, the partial removal (70%) of this layer via the bioreduction of Fe(lll) in the precipitates that coated the surface did not significantly increase the leaching rate of chalcopyrite. Further incubation failed to remove this residual jarosite from the chalcopyrite surface. A thin layer of jarosite (as little as 1 pm) was strongly bound to the chalcopyrite surface and was sufficient to diminish the copper leaching significantly. Parker and co-workers [52], using X-ray photoelectron spectroscopy, detected basic ferric sulphate compounds on chalcopyrite, akin to, but not stoichiometrically identical to, jarosite. Their results indicated that polysulphide products did not play any role in inhibiting chalcopyrite dissolution. T h e slow kinetics was mainly due to the formation of jarosite. In practice, the acidity of the ferric solution is usually maintained between pH 1 and pH 1.8. A s Literature Review 26 shown in Figure 2.8, pH values less than 1.8 limit the extent of F e 3 + precipitation. It is therefore doubtful that the iron compounds alone could be responsible for the significant decrease in chalcopyrite leaching rates. 2.4.2. Passivation of chalcopyrite during bioleaching Microbes have been active in dissolving copper sulphide minerals for the commercial recovery of copper since about 1670 [53]. However, it was not until the late 1940's that a link was established between the bacterium acidithiobacillus ferrooxidans and acid mine drainage from pyrite inclusions in bituminous coal deposits. Since that time the copper industry has made extensive use of heap and dump leaching with more than 25% of copper production estimated to come from that source [54]. Bioleaching is based on the ability of certain microorganisms to catalyze the oxidation of sulphide minerals. The energy needed for their growth is usually obtained from the oxidation of reduced compounds of sulphur and ferrous. The mechanisms of microbial action on sulfide minerals are usually discussed in terms of a direct and an indirect mechanism. In the indirect mechanism, microbes produce oxidising conditions in the system through their ability to oxidise ferrous to ferric. Sutton and Corrick [55] reported an indirect mechanism for the bioleaching of chalcopyrite. They proposed the following sequence: C u F e S 2 + 4 0 2 -> C u S 0 4 + F e S 0 4 (2.24) The ferrous sulphate product is oxidized by oxygen in the presence of microbes: 4 F e S 0 4 + 2 H 2 S 0 4 + 0 2 -> 2 F e 2 ( S 0 4 ) 3 + 2 H 2 0 (2.25) The ferric sulphate product then reacts with chalcopyrite to form more ferrous sulphate: C u F e S 2 + 2 F e 2 ( S 0 4 ) 3 C u S 0 4 + 5 F e S 0 4 + 2 S ° (2.26) The elemental sulphur can also be oxidized to acid: 2 S ° + 3 0 2 + 2 H 2 0 -> 2 H 2 S 0 4 (2.27) Literature Review 27 Fe(lll) species in the bulk solution phase act as electron acceptors. This mechanism is similar to the chemical ferric leaching of chalcopyrite discussed in the previous section. The only difference comes from the fact that the formation of sulphate is favoured in the presence of sulphur oxidizing bacteria (reaction 2.25). This can make a significant difference in the rate of leaching. T h e direct mechanism is the situation where the microbes oxidise the mineral surface directly using enzymes as indicated in Figure 2.9. This means that the microbes maintain a close contact across the extracellular polymeric layer with the sulphide surface in order to dissolve it. Fe2+ leaching leaching Figure 2.9. Scheme visualizing indirect leaching and direct leaching of metal sulphide [57] The electron transfer involves Fe(lll) bound in the cell envelope and the extracellular polymeric layer. This bound iron acts as an electron shuttle between the electron donor and the electron transport system of the cell, which conveys a major portion of electrons to O2 and the rest to CO2 [56]. T h e number of sites on a metal sulphide may be limited for microbial attachment. Therefore, once all sites are occupied, further multiplication of attached cells will result in the displacement of microbes into the bulk phase. This can Literature Review 28 lead to the oxidation of ferrous by an indirect mechanism. This combined mechanism is also called cooperative leaching [57]. T o date, no one has deduced the relative importance of the respective mechanisms. Bioleaching processes are particularly suited for treating copper sulphides such as covellite and chalcocite, but are problematic for chalcopyrite. For example, Dew et al. [24] noted that the bioleaching of various copper sulphide concentrates with mesophiles resulted in 98% dissolution of chalcocite, 94% dissolution of bornite, 90% dissolution of covellite and 43% dissolution of chalcopyrite after 6 days. The slow kinetics of chalcopyrite leaching observed in the presence of microbes has been the subject of much controversy, with different authors concluding that their presence enhances leaching [58,59], has no effect [60] or is detrimental to leaching [4]. According to Third et al. [5], it is not the bacteria that are inhibitory, but rather the ferric and high solution potentials generated by their metabolism. A s a result, if the rate of ferrous biooxidation exceeds of rate the ferric consumption at the mineral surface, then any further microbial activity can be detrimental for the leaching process. A large number of publications have reported that the kinetics of chalcopyrite bioleaching are characterized by three distinct phases, which yield a characteristic sigmoidal leach curve. The first is a lag period marked by little leaching because the microbial population needs time to grow. The duration of the lag phase is dependent on many factors such as the extent of adaptation to the substrate and the presence of inhibiting substances. During this phase the initial rates of leaching are low and only acid soluble minerals are dissolved. Rivera-Santillan et al. [25] observed a lag phase of 10 to 20 days, depending on the temperature. In the second phase, copper leaches rapidly and the redox potential increases indicating the distinctive action of microbes during the dissolution process. The cell population increases exponentially by cellular division. A n d finally, in the third phase, leaching slows abruptly and then ceases entirely. Various authors observed that copper dissolution ceased after about 80% recovery [25,61]. T h e reasons for the abrupt cessation of leaching have not been determined with certainty. This so-called death phase can be attributed to the depletion of the energy source and nutrients. In addition, a toxic effect may be produced by a Literature Review 29 build-up of metabolic products or an increase in the ionic strength of the solution. Norris and Parrot [62] suggested that the incomplete dissolution of copper might result from the formation of a barrier of secondary minerals on chalcopyrite surfaces that will inhibit those microbes capable of oxidizing elemental sulphur. Given that the basic reactions occurring in biological systems are the same as those in purely chemical systems, with the microbes acting as a catalyst, the passive layers in chalcopyrite bioleaching are similar to those formed during ferric leaching as discussed in the previous section. In addition to the formation of passive layers on the chalcopyrite surface, other factors related to the activity of microbes might also contribute to the slow kinetics. These factors are discussed individually below. Metal tolerance T h e tolerance level of microbes for specific ions in solution depends on their physiological state, the chemical form of the metals, and the degree of interaction of the metals with the environment [62]. It has been reported that some bacteria become ineffective when the copper content of the solution reaches 25 to 30 g/L [53]. Tuovinen and Kelly [63] demonstrated that copper in concentrations of 6.5 to 65 g/L inhibits iron oxidation and CO2 fixation. Apparently tolerance levels of microbes with respect to copper can be a problem in bioleaching of rich chalcopyrite concentrates. Aeration The supply of oxygen to a microbial culture is probably the most important factor determining their activity [54]. The microbes also use carbon dioxide as their source of carbon for cell growth. The limitations of oxygen and carbon diffusion may cause a decrease in the rate of biooxidation. It should be noted that the dissolved equilibrium concentrations of oxygen and carbon dioxide decrease with increasing temperature. It has also been observed that the oxygen and carbon dioxide consumption rates of thermophiles are significantly larger than for mesophiles [64]. This implies that the transfer of oxygen and carbon dioxide may easily become rate limiting during Literature Review 30 thermophilic bioleaching of sulphide minerals. It follows that attention needs to be paid to prevent gas-liquid mass transfer limitations during bioleaching of sulphide minerals. Nutrients and additives T o become active, bacteria need oxygen, carbon dioxide, nitrogen, and phosphate. Nitrogen is generally supplied in the form of N H 4 + . The addition of N H 4 + and P 0 4 3 ~ is limited because of cost and the risk of precipitation as ammonium jarosite and potassium phosphate [65]. In addition, other nutrients such as magnesium and calcium are supplied to the medium. Magnesium is necessary for C 0 2 fixation [66]. Types of microbes It is also important to understand the behavior of microbes on the mixed substrates F e 2 + and S ° , which are formed during the bacterial leaching of chalcopyrite. For example, Acidithiobacillus ferrooxidans is capable of oxidizing ferrous and reduced sulphur compounds (e.g., elemental sulphur or thiosulphate) while Leptospirillum ferrooxidans can only oxidize ferrous [67]. This means that, when compared to Leptospirillum ferrooxidans, Acidithiobacillus ferrooxidans can increase the dissolution rate of sulphide minerals by oxidizing the sulphur layer. However, Norris et al. [68] found that Leptospirillum ferrooxidans dissolved pyrite more extensively than Acidithiobacillus ferrooxidans. It is possible that these two bacteria leach sulphide minerals by different mechanisms. Knowledge of the mechanism of microbial attack could improve our understanding of chalcopyrite passivation. This succinct review of ferric leaching and bioleaching processes has shown that the surface of the chalcopyrite is altered in some manner, shielding it from further oxidation. It is apparent that the solution potentials and formation of metastable compounds are certainly two important factors in the study of chalcopyrite passivation. These two factors are conveniently studied using an electrochemical approach. This is discussed in the next section. Literature Review 31 2.4.3. Passivation during the electrochemical dissolution of chalcopyrite A s a result of their conductivity, certain minerals can participate in coupled charge transfer processes analogous to a metal corroding in an electrolyte, and the kinetics of leaching can be related to the potential of the solid in contact with the aqueous electrolyte [69-71]. Resistivities, electronic and structural information for selected sulphides and oxides are shown in Table 2.1. The covalent character of most sulphide minerals provides non-localization of charge, resulting in appreciable intrinsic electronic conductivity. Oxide minerals are in general more ionic in nature and usually have higher resistivities compared to sulphide minerals [71]. The electronic properties shown in Table 2.1 indicate that chalcopyrite is a good conductor with a resistivity of about 1CT3 Q-m. Therefore, the dissolution reaction of chalcopyrite in ferric solutions can be interpreted electrochemically. T h e electrochemical dissolution of chalcopyrite can be written in terms of the half reactions for the oxidation of chalcopyrite and the reduction of ferric: Anodic: C u F e S 2 -> C u 2 + + F e 2 + + 2 S ° + 4 e~ (2.28) Cathodic: 4 F e 3 + + 4 e~ -> 4 F e 2 + (2.29) Reactions (2.28) and (2.29) occur simultaneously on the entire surface of the dissolving chalcopyrite at rates that satisfy the condition that the net production of electrons is zero. T h e potential at which this condition is met is called the mixed potential. Since reactions (2.28) and (2.29) take place independently, they can be studied separately, and rate expressions for the two half-reactions can be established. According to the mixed potential theory, the corrosion potential and the net dissolution rate can be obtained by combining the anodic and cathodic current-potential curves on so-called Evans diagrams. The main interest in electrochemical studies of chalcopyrite arises from the information such studies provide concerning the mechanism and kinetics of chalcopyrite dissolution and passivation in acidic solutions. Literature Review 32 Table 2.1. Electronic and structural properties of selected sulphide and oxide minerals at 2 5 ° C [71]. Formula Name Resistivity (Q-m) Conductor Structure Ionic structure Cu 5 FeS 4 Bornite 10~3 - 10 - 6 P Tetragonal (Cu +) 5Fe 3 +(S 2-) 4 Cu 2 S Chalcocite 4x10" 2 -8x1Cr 5 P Orthorhombic (Cu+ ) 2 S 2 _ CuFeS 2 Chalcopyrite 2x10"4 -9x10 - 3 n Tetragonal Cu + Fe 3 + (S 2 ") 2 CuS Covellite 8x10" 5 -7x10" 7 metallic hexagonal (Cu + ) 2 S 2 2 -PbS Galena i x i o ^ - z x i o - 6 n&p Cubic P b 2 + S 2 " MoS 2 Molybdenite 7 .5 - 8x10"3 n&p Hexagonal Mo 4 +(S 2") 2 FeS 2 Pyrite 3 X 1 C T2 - 1 X 1 0 - 3 n&p Cubic F e 2 + S 2 2 " ZnS Sphalerite 3 X 1 0 " 3 - 1 X 1 0 ^ Cubic Zn 2 + S 2 " Sn0 2 Cassiterite 1 0+ 2 - 1 0 " 2 n Tetragonal Sn 4 + (0 2 ") 2 Cu 2 0 Cuprite 10~1 - 10 P Cubic (Cu+ ) 2 0 2 _ F e 2 0 3 Hematite 2.5x10"1 -4x10" 2 n&p Trigonal (Fe 3 +) 2(0 2-) 3 F e 3 0 4 Magnetite 2x10^-4x10"5 n&p Cubic Fe 3 + [Fe 3 + Fe 2 + ] (0 2 - ) 4 Mn0 2 Pyrolusite 10 " 1 -10 " 3 n Tetragonal Mn4 +(0 2") 2 Ti0 2 Rutile 104 - 10 n&p Tetragonal Ti4 +(0 2") 2 The anodic reaction The anodic polarization curves of chalcopyrite usually exhibit three electro-active regions, namely, the dissolution, passive and transpassive regions [70,72]. The passive region reflects the formation of progressively thickening surface films. Several theories concerning the nature of this passive film have been proposed in the literature. These will be reviewed in the present section. Formation of defect structures According to various authors, the anodic reaction, expressed by reaction (2.28), occurs through a series of intermediate structures [69-71]. Warren et al. [70] proposed the following sequence for the anodic dissolution of chalcopyrite: CuFeS2 -> Cui_ xFe 1_ yS 2-z + x C u 2 + + y F e 2 + + z S° + 2 (x+y) e" (2.30) where y> x initially. Literature Review 33 The intermediate or defect structure Cui_ x Fei -yS2- z releases ferrous much faster than cupric. T h e shorter the time of measurement, the higher is the F e : C u ratio obtained. For example, Baur et al. [73] found a ratio of 6.6 at 30 s, while Linge [43] reported a ratio of about 2 over the first 5 min. These results clearly indicate an induction period in the early stage of chalcopyrite leaching at;low potentials. It is only after a longer time that a congruent release of C u 2 + and F e 2 + can be expected. Undoubtedly, the leaching time is an important variable in studies dealing with the preferential dissolution of either copper or iron from the chalcopyrite mineral. Other factors such as temperature and redox potentials may also affect the selective dissolution of copper. The intermediate Cui_ x Fei_yS2- z decomposes further to form a second (non-stoichiometric) intermediate, CuS( n- S), which forms in the region 700 to 750 m V (SHE) and may be represented by the following reaction: Cui-xFei_yS2-z -> (y-x) CuS(n-s) + (y-x) CuS ( n - S ) + (1-y) C u 2 + +(1-y) F e 2 + + (2+x-y-z) S ° + 4(1-y) e~ (2.31) Cu-i_xFe-i_yS2-z and CuS (n-s) are termed S1 and S2 respectively. The intermediate products S1 and S2 serve as a passivation layer by limiting ionic transport to the mineral surface. Equation 2.31 indicates a congruent release of C u 2 + and F e 2 + during this reaction. Warren et al. [70] found a ratio of dissolved iron to dissolved copper of 1:1 at a constant potential of 1.09 V (SHE) after extended leaching times (17-84 h). Baur et al. [73] obtained similar results for extended leaching times. Stankovic [74] studied the anodic dissolution of chalcopyrite using galvanostatic pulse chronopotentiometric techniques and proposed a two-stage oxidation process. T h e first stage involves the release of C u 2 + and acts as the limiting step of the reaction, whereas the second stage involves the release of F e 2 + , in accordance with the following scheme: n C u F e S 2 -> C u 2 + + C u , ^ F e „ S 2 n + 2 e~ (rate-limiting) (2.32) CUn-iFenS2n -> F e 3 + + C u n _ 1 F e n _ 1 S 2 n + 3 e~ (2.33) Literature Review 34 The overall reaction can be expressed by: n C u F e S 2 -»• C u 2 + + F e 3 + + C u ^ F e ^ S ^ + 5 e" (2.34) According to this scheme, the dissolution of chalcopyrite would lead to the formation of a defect structure, also called polysulphide. Formation of elemental sulphur The anodic oxidation reactions of chalcopyrite in acidic medium (Equations 2.28 and 2.30) indicate that elemental sulphur forms on the chalcopyrite surface. Jones and Peters [38] believe that the layer of elemental sulphur formed during the anodic dissolution does not passivate chalcopyrite. This is illustrated in Figure 2.10 where anodic polarization curves are shown for various ferrous levels. T h e plateau region, indicating diffusional control, moves up with increasing ferrous concentration. Since sulphur is an electronic insulator, the plateau should not have been influenced by the addition of ferrous to the solution. However, the increased current in the plateau region can also be attributed to the oxidation of ferrous to ferric ions. Lu and al. [75] found that the solvent C S 2 does not remove the passive film from the chalcopyrite surface. T h e y conducted potentiostatic experiments after C S 2 washing within the potential range 450-900 m V (SHE) and failed to observe any significant change in the leaching rate. They concluded that chalcopyrite passivation was mainly due to a metal-deficient structure rather than elemental sulphur. Parker et al. [34] also support the concept of a metal deficient copper-rich polysulfide, which has different semiconductor properties from C u F e S 2 and slows not only transport of C u 2 + and F e 2 + but also electron transfer to or from electroactive species in solution. Reddy et al. [11], using E S C A and microscopic analysis, confirmed the formation of a semi-conducting, metal-deficient, polysulphide film on the chalcopyrite surface whose composition varied with temperature, pulp density, particle size and acid concentration. Literature Review 35 ,j_ S A M P L E 3E, T R A N S V A A L 1 m -0.5b I0"7 lO* 6 ICf5 10"* I0' 3 CURRENT DENSITY (A/cm?) Figure 2.10. Effect of ferrous on the anodic polarization curve of Transvaal CuFeS2 in 1 M H 2 S 0 4 ) 40 mV m i n - 1 , 2 5 ° C [70] Formation of covellite and bornite The anodic polarization curves obtained by Warren et al. [70] indicate that the rest potential of chalcopyrite is located between 0.52 V and 0.62 V (SHE) , (0.28-0.38 V (SCE)) at 2 5 ° C in a solution of 1 M H 2 S 0 4 (Figure 2.11). In order to aid the interpretation of current-potential curves, a number of possible reactions were considered within the range 0.4-0.9 V (SHE) because practical leaching of chalcopyrite usually takes place within this potential range [34]. O n the basis of thermodynamics alone, and assuming that there is no substantial activation overpotential for the reaction on the chalcopyrite electrode, all of the reactions given in Table 2.2 would be feasible within the potential range 0.4-0.9 V (SHE). Literature Review 36 c * m o . « » « « i c o — —. suoeuRv.oNTftfiio 0.5--IO"6 IO"5 IO"4 10"3 CURRENT DENSITY ( A / c m z ) Figure 2.11. Anodic polarizarion curves for C u F e S 2 from six different locations in 1 M H 2 S 0 4 , 30 m V m i n - 1 , 2 5 ° C [70] Table 2.2. Reactions of chalcopyrite [69] Reactions £° (V (SHE)) C u F e S 2 F e S 2 + C u2 + + 2e~ (2.35) + 0.50 C u F e S 2 - » F e S + S ° + C u2 + + 2e~ (2.36) + 0.81 C u F e S 2 -> C u S + S ° + F e3 + + 3e~ (2.37) + 0.45 C u F e S 2 -> C u 2 + + F e 2 + + 2 S ° + 4 e " (2.38) + 0.47 2 C u F e S 2 -> C u 2 S + 2 F e2 + + 4 S ° + 4 e" (2.39) + 0.87 C u F e S 2 -> C u 2 + + F e 3 + + 2 S ° + 3e~ (2.40) + 0.51 C u S -> C u 2 + + S ° + 2 e" (2.41) + 0.47 F e 2 + F e 3 + + e" (2.42) + 0.77 Based on the potential-pH diagram for the C u - F e - S - H 2 0 system and the data presented in Table 2.2., Warren et al. proposed that S1 and S2 are bornite ( C u 5 F e S 4 ) and covellite (CuS). McMillan et al. [8] reported results on the electrochemical oxidation of chalcopyrite in both acidic sulphate and acidic chloride media under the conditions Literature Review 37 relevant to chemical leaching; i.e., at temperatures greater than 7 0 ° C and over the potential range 0.45-0.85 V (SHE) . Their results indicated the formation of a surface layer, a "solid electrolyte interphase", which slows the rate of electron transfer. The passive layer had the properties of covellite. Biegler and Home [76] used cyclic voltammetry to propose the following reaction: 4 C u F e S 2 3 C u S + C u 2 + + 4 F e 2 + + 5 S ° + 10e" (2.43) where C u S (covellite) and S ° would be in a ratio of 3:5 and F e 2 + : C u 2 + in solution would be in a ratio of 4:1. O n the other hand, Holliday and Richmond [69] also reported the formation of covellite during the anodic dissolution of chalcopyrite and suggested the following sequence in the low potential region: C u F e S 2 - » C u 2 + + F e S 2 + 2 e" (2.44) C u 2 + x C u 2 + (ads) + (1-x) C u 2 + (aq) (2.45) C u 2 + (ads) + F e S 2 -+ C u S + F e 2 + + S ° (2.46) Jones and Peters [72] reported the formation at 1 7 5 ° C of large and well-defined crystals of covellite, C u S , in the first region of the anodic polarization curve of chalcopyrite. Covellite is a cuprous polysulphide and is reported to behave energetically like free sulphur [77]. Peters [35] notes that by adjusting Eh-pH diagrams to remove the contribution for reduction of sulphate to sulphur, the thermodynamic stability regions are altered so that covellite is thermodynamically stable up to 0.6 V (SHE) over the range pH 0 to pH 14. A s early as 1930, Sullivan [36] postulated the conversion of chalcocite to covellite accompanied by a deceleration of copper leaching. King [78] studied the dissolution of chalcocite in ferric chloride media and observed that covellite was an intermediate product during the leaching process. There was evidence that C u S was non-stoichiometric, and was, therefore, different from the natural covellite mineral used by most other workers. However, Hackl [10,79] rejected the theory of chalcopyrite passivation by bornite and covellite on the grounds that these two minerals are known Literature Review 38 to dissolve much more readily than chalcopyrite in sulphate medium. He proposed that the passive layer was a copper polysulfide, C u S „ , where n > 2. Galvanic interaction between chalcopyrite and intermediate products Numerous authors have observed changes in surface speciation during the ferric leaching of sulphide minerals. For example, Koch and Mclntyre [80] observed the formation of metastable phases during the electrochemical oxidation of chalcocite to covellite. The following sequence was proposed: C u 2 S - > C-U1.93S -> CU1.83S - > C U L B / S - > CU1.4S -> C u S This sequence progresses from states of lowest rest potential to states of highest rest potential. It is therefore possible that similar intermediate products are formed during the oxidative leaching of chalcopyrite. The intermediate phases are believed to be semi-conductors, and because they are in contact with unreacted chalcopyrite, a galvanic couple may be formed between the two phases as illustrated in Figure 2.12. Intermediate compounds Porous S ° Figure 2.12 Schematic representation of galvanic interactions between chalcopyrite and intermediate phases formed during the dissolution of chalcopyrite The galvanic interaction will depend on the rest potentials of the two semi-conductors, the exchange current densities, the Tafel slopes and the relative areas of the minerals involved. Since the nature of the intermediate phases has not yet been elucidated with Literature Review 3 9 certainty, it is difficult to predict the magnitude of the galvanic interaction between the chalcopyrite and the intermediates phases. Additional investigations are required to determine the contribution of this effect on the passivation of chalcopyrite. Natural flotation of chalcopyrite Another possible aspect of passivation is wetting of the chalcopyrite surface by leaching solutions. Several studies on the flotation of sulphide minerals have dealt with the changes occurring at the mineral surface and their effect on flotation [81-83]. It has been reported that, under appropriate conditions, chalcopyrite is subject to collectorless flotation [81-83]. This means that the surface becomes hydrophobic. Electrochemical studies confirmed that the surface of chalcopyrite became increasingly hydrophobic when conditioned at electrochemical potentials of 0.64, 0.84, and 1.04 V (SHE) for 4 h at 2 5 ° C . This characteristic inhibits contact of the chalcopyrite surface with leaching solutions. The cathodic reaction In studying the passivation of chalcopyrite, the cathodic reduction of ferric on chalcopyrite, expressed by equation (2.29), is of great significance since the true leaching rate can be predicted based on the mixed potential theory, i.e., by combining the cathodic and anodic branches of the half-cell reactions. Depending on the position of the cathodic and anodic branches, Nicol [84] has classified the oxidative leaching of sulphide minerals in three categories: Type I Leaching For this type, the oxidant and the dissolving mineral are in their so-called "high-field regions" and the mixed potential intersects their Tafel slopes. This type represents the simplest situation to analyze and the leaching rate is proportional to the square root of the oxidant concentration [85]. In general, Type I leaching is observed when the exchange current densities of the half-cell are similar in magnitude, but the reversible potentials are far apart. The situation is shown schematically in Figure 2.13. Literature Review 40 Type II Leaching O n many sulphide minerals, the reduction of ferric to ferrous takes place more or less reversibly. In this case, the exchange current density of the oxidizing couple is higher than that of the dissolution reaction, and the mixed potential thus corresponds to the reversible potential of the oxidizing couple. The situation is shown schematically in Figure 2.14. For this case, the leaching rate is proportional to the square root of the ratio of the concentrations of the oxidized and reduced forms of the oxidant. Type III Leaching In some cases , the leaching reaction is particularly fast and mass transfer of oxidant to the mineral surface limits the leaching rate. A s shown in Figure 2.15, Type III leaching is observed when the reversible potentials of the half-reactions are relatively far apart. In this situation the leaching reaction will proceed at the limiting current density of the oxidant, i.e., the leaching rate is proportional to the oxidant concentration. Type IV Leaching This situation, shown schematically in Figure 2.16, is encountered when the reversible potentials are close to one another. Luckily, most sulfide leaching reactions of industrial interests are well described by one of the three types outlined above, as shown in Table 2.3. below. Literature Review 41 E E M Type 1 X / ^ / ^ x ' \ \ I \ J 1 \ In j d In j Figure 2.13. Mixed potential for Type I leaching Type II F ei + <-> Fe 2 + / /i In jd In j Figure 2.14. Mixed potential for Type II leaching Literature Review E M . Type III Fe3* <-> Fe 2 + / X • A / In j«, In j Figure 2.15. Mixed potential for Type III leaching Type IV Figure 2.16. Mixed potential for Type IV leaching Literature Review Table 2.3. Reaction rate orders for selected leaching systems [84, 85] Medium Reactants (Order) Type OXIDATIVE U 0 2 + Fe(lll) S 0 4 2 " Fe(lll) (0.5) I F e S + Fe(lll) S 0 4 2 " Fe(lll) (0.5), Fe(ll) (-0.5) II F e S 2 + 0 2 S 0 42 " 0 2 , (0.5) I F e S 2 + Fe(lll) S 0 4 2 " Fe(lll) (0.5), Fe(ll) (-0.5) II F e S 2 + 0 2 O H " 0 2 , (0.31) I Z n S + Fe(lll) S 0 4 2 " Fe(lll) (0.45), Fe(ll) (-0.5) II Z n S + Fe(lll) c r Fe(lll) (0.36), Fe(ll) (<-1) IV C u F e S 2 + 0 2 S 0 4 2 " 0 2 , (0.35) I C u F e S 2 + Fe(lll) c r Fe(lll) (0.5, 0.3-0.4) I C u F e S 2 + Cu(ll) c r Cu(ll) (0.54), Cu(l) (-0.5) II A g 2 S + C N " + 0 2 C N " 0 2 , (0.5) I P b S + 0 2 O H " 0 2 , (0.5) I N i 3 S 2 + Fe(lll) S 0 4 2 " Fe(lll) (0.2), Fe(ll) (-0.2) II C u 2 S + Fe(lll) stage 1 S 0 4 2 _ Fe(lll) (1.0), III C u 2 S + Fe(lll) stage II s o 4 2 _ Fe(lll) (0.4), Fe(ll) (-0.3) II A u + C N " + 0 2 C N " C N " (0 or1) , 0 2 , ( 0 o r 1 ) III REDUCTIVE M n 0 2 + S 0 2 S 0 4 2 " S 0 2 (0.5) I M n 0 2 + Fe(ll) S 0 4 2 " Fe(ll) (0.58) I F e O O H + S 0 2 S 0 4 2 " S 0 2 (0.5-0.76) I F e 3 0 4 + Fe(ll) S 0 4 2 " Fe(ll) (0.5), Fe(lll) (-0.5) II Z n F e 2 0 4 +. Fe(ll) S 0 4 2 " Fe(ll) (0.6), Fe(lll) (-0.5) II It is of interest to investigate the factors affecting the position of the cathodic branches and the magnitude of the cathodic current during the dissolution of chalcopyrite. A s discussed in the previous section, a compact solid oxidation product covers the chalcopyrite surface as leaching proceeds. These oxidation products may slow the transfer of electrons and transport of ferric to the cathodic surface. Jones and Peters [38] reported that there is a significant degree of irreversibility on a chalcopyrite surface during leaching in ferric sulphate solutions. In addition, it has been demonstrated that Literature Review 44 bioleaching of chalcopyrite is a very slow process and high redox potentials can be maintained in the system. These two factors, namely leaching time and high redox potentials, are also known to contribute to the formation of a passive layer on the chalcopyrite surface. It is then possible that the Tafel slope and the exchange current density for the cathodic reaction on the chalcopyrite surface become less uniform as bioleaching proceeds. A more complete understanding of ferric reduction on chalcopyrite surfaces is crucial in studies devoted to chalcopyrite passivation. 2.5. Methods for activating the passive film during the ferric leaching and bioleaching of chalcopyrite A s discussed in the previous sections, chalcopyrite appears to passivate under oxidizing conditions in acidic sulphate media. The literature refers to several methods used to depassivate chalcopyrite during ferric and bacterial leaching. These methods will be briefly reviewed in the next sections. 2.5.1. Mechanical activation Mechanical activation consists of comminution of mineral particles that results in changes to a great number of physicochemical properties of a particular system [29]. This disintegration by high energy grinding is accompanied by an increase in surface area, increases in internal and surface energies, a decrease in the coherence energy of solids, an increase in reactivity, and an increase in the number of particles by generation of fresh, previously unexposed surfaces. Dutrizac [1] ground chalcopyrite particles down to 10-14 um and observed rapid dissolution. There was a linear relationship between the initial rate constants and Mr, where r is the mean particle size of chalcopyrite. Beckstead et al. [47] used attrition grinding to reduce the particle size of chalcopyrite to 5 um and reported 90% copper conversion after 3 hours of ferric leaching at 9 3 ° C . They found that the initial rate of leaching increased in direct proportion to the initial surface area. Gerlach et al. [86] ground chalcopyrite in a vibrating impact mill and observed that copper extraction (at 1 0 0 ° C and under 2 M P a oxygen pressure) was improved from 15% for unground concentrate to 98% after grinding for 3 hours. These results indicate that high copper recoveries can be obtained Literature Review 45 by ultrafine grinding. However, these methods are energy intensive and costly, and they complicate solid/liquid separation downstream. 2.5.2. Addition of silver It is well known that small amounts of A g added to solution greatly accelerate the ferric leaching and bioleaching of chalcopyrite [87-94]. Pawlek [87], for example, has found that 30 min was required to extract 51% of the copper at 1 0 0 ° C during the oxygen leaching of chalcopyrite. With the addition of only 0.75% dissolved silver by weight, 95% of the copper was extracted in the same length of time under the same conditions. Snell [95] obtained a similar beneficial effect when small amounts (50-500 mg/L) of silver were added to the ferric leaching solution. The copper extraction increased from 46 to 90% after three hours of leaching at 91 ° C . Different mechanisms have been proposed to explain the accelerating effect of silver on the leaching rate of chalcopyrite. In one mechanism, Miller et al. [88] believe that the accelerating effect of silver is due to the deposition of silver as A g 2 S according to the equation (2.47): C u F e S 2 + 4 A g + -> 2 A g 2 S + C u 2 + + F e 2 + (2.47) The silver ions are regenerated by the reaction between ferric and A g 2 S : Ag(l) thus released will again react with chalcopyrite. T h e y believed that, unlike the uncatalysed ferric sulphate leach in which elemental sulphur forms a tight and tenacious film resistant to transport, the silver catalysed ferric sulfate leach results in a porous sulphur layer which does not prevent electron transport at the mineral surface. In addition, through the use of polarization curves and constant potential experiments, they confirmed the occurrence of the anodic dissolution of A g 2 S electrodes in 1 M H2SC>4: This would indicate that the chemical oxidation of silver sulphide via reaction (2.48) is just as probable as the electrochemical oxidation via reaction (2.49). In another A g 2 S + 2 F e 3 + -> 2 A g + + 2 F e 2 + + S' o (2.48) A g 2 S -> 2 A g + + S ° + 2 e" (2.49) Literature Review 46 mechanism the reactivation effect due to Ag2S can be explained by an electrochemical mechanism based on the formation of the A g 2 S / C u F e S 2 galvanic cell, with A g 2 S acting as cathode and C u F e S 2 acting as anode [90]. T h e beneficial effect of silver was also observed during bioleaching of chalcopyrite. T h e presence of microbes accelerates ferrous oxidation, providing the necessary conditions to dissolve the Ag2S film. In these various mechanisms of silver-catalysed chalcopyrite leaching, the formation of Ag2S is a critical step. The main problem with silver as a catalyst for the chalcopyrite leaching process is that recovery of the silver from the residue is difficult because of the formation of silver jarosite. Recycling the silver is important because of its high cost. Others ions ( H g 2 + , C o 2 + , B i 3 + and C u 2 + ) have been tested with far less dramatic results [91]. 2.5.3. Beneficial effect of galvanic interactions between chalcopyrite and associated minerals In the opinion of Majima and Peters [96], the galvanic effect may be one of the most important electrochemical factors which governs the dissolution rate of sulphide minerals in hydrometallurgical systems. It has been shown that the rate of oxidation of certain natural sulphides is greatly increased by the presence of either marcasite or pyrite [23,97]. A s mentioned previously, in a mixture of two different sulphides, significant oxidation of one sulphide occurs while the other is galvanically protected, due to the difference in potential between the two sulfide minerals. T h e rest potential for chalcopyrite is lower than that of pyrite as shown in Table 2.4. W h e n the two minerals are in contact in an acidic medium, chalcopyrite, acting as an anode, will dissolve faster, while pyrite is galvanically protected. Literature Review 47 Table 2.4. Rest potentials of selected sulphide minerals in acidic solutions Mineral Rest potential (V (SHE)) in 1 M H2SO4 at 20°C References F e S 2 0.63 Hiskey [71] C u F e S 2 0.52 Hiskey [71] A g 2 S 0.51 Warren [98] C u 2 S 0.44 Hiskey [71] C u S 0.42 Hiskey [71] P b S 0.28 Hiskey [71] Few studies have been conducted to understand the mechanism of galvanic interactions when various minerals are in contact in the presence of the microorganisms commonly met with in dump or heap leaching. Mehta and Murr [23] examined the effect that galvanic interactions had on the rate of bacterial leaching of chalcopyrite. It was observed that as the amount of pyrite in contact with chalcopyrite increases the rate of leaching of chalcopyrite increases. A C u F e S 2 : F e S 2 ratio of 5:5 (g/g) and size fraction of - 7 4 um was found to be the most preferable when leaching was carried out in the presence of A. ferrooxidans. The presence of pyrite in the optimized quantities enhanced the bacterial leaching of chalcopyrite by a factor of 15 after 48 days of leaching. The potential in the inoculated experiments increased from 320 m V to 570 m V ( S C E ) . However, these two values of potential are higher than the rest potentials of pyrite and chalcopyrite. This indicates that the pyrite surface will not act totally as cathode. One part of pyrite will act as anode and the other part will act as cathode. This may reduce the available surface for the cathodic reaction and consequently reduce the dissolution rate of chalcopyrite. Since the bacterial leaching of sulphide minerals takes place within a wide potential range, it is of interest to investigate in two steps the galvanic interaction between pyrite and chalcopyrite. T h e first step will concern the galvanic interaction at lower potentials (between the rest potential of chalcopyrite and that of pyrite) and the second step will be associated with higher potentials. It will be discussed in the present work whether or not the redox potential can influence the Literature Review 48 dissolution of chalcopyrite in galvanic contact with pyrite. A better understanding of this mechanism is crucial in the study of the ferric leaching and bioleaching of chalcopyrite. 2.5.4. Thermophilic bioleaching Bioleaching of chalcopyrite has been studied extensively, especially with mesophilic bacteria such as Acidithiobacillus ferrooxidans, Acidithiobacillus thiooxidans and Leptospirillum ferroxidans [55,99,100]. Mesophiles operate optimally between the temperature ranges of 30 to 3 7 ° C . However, chalcopyrite does not respond well to mesophilic bioleaching. Chalcopyrite leaching using moderate and extreme thermophiles such as Acidianus brierleyi, Archaea and Sulfolobus has met with some success and various authors [101-104] have suggested that thermophiles might be superior to mesophiles for bioleaching chalcopyrite. For example, Le Roux and coworkers [101] observed that the average rate of copper dissolution from chalcopyrite in a batch reactor with mesophiles was 2.5 mg L" 1 h"1 with a conversion of 19%. Under comparable conditions, the equivalent rate with thermophiles was 11.5 mg L ~ 1 h~1 with a conversion of 83%. The six-fold increase in the rate of reaction makes the thermophilic bioleaching process economically attractive for bioleaching chalcopyrite. T h e success of thermophiles over mesophiles has not been clearly explained. Leaching rates are considerably increased in the presence of thermophiles. Three significant facts are observed during bioleaching tests with extreme thermophiles: > Iron in solution is mainly found as ferrous [105]. > Thermophilic bioleaching of chalcopyrite takes place in a fairly low potential environment (380-500 m V (Ag/AgCI)) [16]. > Thermophiles appear to attach to the ore and perform a direct attack on chalcopyrite [54]. They appear to float freely and oxidize dissolved ferrous to a much lesser extent than mesophiles. These facts suggest that chalcopyrite leaches faster in the presence of thermophiles because of the low solution potentials. The critical potential for chalcopyrite passivation is believed to be around 600 mV (Ag/AgCI) [10,42]. Literature Review 49 Commercial application of thermophiles in conventional stirred-tank bioleaching reactors may be potentially limited by a number of factors such as a more fragile cellular structure that can be damaged at high pulp density and by the high shear conditions generated in these reactors, and the significantly reduced solubilities of oxygen and carbon dioxide at the operating temperatures required for extreme thermophiles. A number of studies have apparently addressed these challenges, although specific design details have not yet been published. For example, B H P Billiton, B a c T e c h and Mintek have developed their bioleaching processes using thermophiles up to the demonstration plant scale. 2.6. Conclusions and focus of the present study A key factor for the successful application of hydrometallurgical processes for the treatment of chalcopyrite is a better understanding of those factors that prevent rapid chalcopyrite dissolution. The slow kinetics observed during the ferric leaching and bioleaching of chalcopyrite have been studied extensively by a number of researchers but there are still doubts about the exact nature of the passive layers and the conditions under which they are formed. The formation of metastable intermediates of copper-rich polysulphide seems to be the more likely explanation for the slow leaching rates. It has been shown that the formation and stability of these phases depends on temperature, microbial type, solution potential and leaching time. Better control of these parameters may counteract the passivation of chalcopyrite. Since the dissolution of chalcopyrite in acidic medium involves an electrochemical mechanism, the intermediate phases and parameters affecting their presence can be conveniently studied by electrochemical methods. Electrochemical measurements using a chalcopyrite electrode, coupled with careful kinetic studies of finely ground mineral, may improve our understanding of chalcopyrite dissolution under oxidizing conditions in acidic solutions. Electrochemical studies might therefore hope to establish the conditions under which passivation of the mineral occurs, and the potential dependence of the reaction mechanisms, while stirred-tank reactor studies with fine mineral particles Literature Review 50 will provide information about the effects of temperature, solution potential, galvanic interactions, and the mechanism of microbial attack. T h e key questions that are addressed by this research are: > What is the potential range for severe passivation of chalcopyrite? > How can copper be dissolved preferentially from chalcopyrite? > What is the effect of temperature on the extent and rate of chalcopyrite passivation? > What is the behaviour of ferric on a polarized chalcopyrite surface? > What is the effect of pyrite on the ferric dissolution of chalcopyrite? The aim of this research is to provide an adequate description of chalcopyrite passivation and depassivation by studying the electrochemical response of chalcopyrite, and to validate the resulting electrochemical model under a variety of experimental conditions similar to those found during ferric leaching and bioleaching. Literature Review 51 Chapter 3 METHODOLOGY AND EXPERIMENTAL PROCEDURE The literature survey has revealed that the surface of chalcopyrite is altered during the ferric leaching process causing "passivation" of the mineral. This surface modification is mainly dependent on time, potential and temperature. In order to have an adequate description of chalcopyrite leaching in acidic medium, it is therefore appropriate first to study the mechanism controlling the formation of passive layers by electrochemical techniques. Electrochemical techniques have the advantage over other techniques as they measure properties at the solid-liquid interface. In other words, they elucidate the reactions that are taking place on the mineral surface itself. The effects of temperature and potential and the presence of pyrite are investigated. In addition to the electrochemical study, oxidation in suspension (chemical leaching and bioleaching) with fine chalcopyrite grains is studied in a stirred-tank reactor. Various parameters such as temperature, solution potential, presence of pyrite (various pyritexhalcopyrite ratios), and mineral grain size distribution are investigated. Validation of the electrochemical study and the development of a leaching model for chalcopyrite is an integral part of this work. After identification of the best conditions for chalcopyrite leaching, these results have been applied to two leaching processes used in the copper industry, namely heap leaching (column leaching) and atmospheric leaching. Consequently, the experimental program developed for this study has been divided into five phases: > Electrochemical study with a chalcopyrite electrode; > Controlled-potential chemical leaching of chalcopyrite with fine chalcopyrite particles in a stirred-tank reactor; > Bioleaching of fine chalcopyrite particles in shake flasks; > Small column bioleaching tests with fine chalcopyrite coated on low-grade chalcopyrite ore Methodology and Experimental Procedure 52 > Atmospheric leaching of chalcopyrite in a stirred-tank reactor. 3.1. Electrochemical measurements A s discussed in previous sections, it is believed that the dissolution of sulphide minerals in acidic medium is an electrochemical process involving cathodic and anodic half-reactions. The existence of an electrochemical process suggests that the anodic dissolution of chalcopyrite and ferric reduction on the chalcopyrite surface can be studied separately using chalcopyrite working electrodes. 3.1.1. Working electrodes All chalcopyrite samples used were cut from a massive piece obtained from Santa Eulalia Mining district, Mexico. The detailed chemical analysis of the mineral is given in Table 3.1. Table 3.1. Detailed chemical analysis of the chalcopyrite mineral Element Mass, % Element Mass (%) Copper 27.10 Manganese 0.02 Iron 32.51 Mercury 0 Sulfur total 32.22 Molybdenum 0 Aluminium 1.22 Nickel 0.03 Antimony 0.02 Phosphorus 0 Arsenic 0 Potassium 0.27 Barium 0 Silver 0 Bismuth 0.02 Sodium 0.54 Cadmium 0 Strontium 0 Calcium 1.39 Thallium 0 Chromium 0 Titanium 0.06 Cobalt 0.05 Tungsten 0 Lanthanum 0 Vanadium 0 Lead 0.41 Zinc 1.08 Magnesium 0.52 Zirconium 0 Methodology and Experimental Procedure 53 The mineralogical composition of the chalcopyrite electrode is given Table 3.2 below: Table. 3.2. Mineralogical composition of chalcopyrite electrode Mineral Percent Chalcopyrite (CuFeSa) 80.5 Pyrite (FeS 2 ) 8.2 Sphalerite (ZnS) 1.6 Galena (PbS) 0.5 The pyrite used was also of high purity and from Huanzala Mine, Peru. The detailed chemical analysis of the Huanzala pyrite is given in Table 3.3. Table 3.3. Detailed chemical analysis of the Huanzala pyrite mineral Element Mass, % Element Mass (%) Copper 0.31 Manganese 0.02 Iron 45.30 Mercury 0 Sulfur total 51.40 Molybdenum 0 Aluminium 0.39 Nickel 0.06 Antimony 0.26 Phosphorus 0 Arsenic 0.31 Potassium 1.23 Barium 0 Silver 0 Bismuth 0.04 Sodium 1.34 Cadmium 0 Strontium 0 Calcium 0.13 Thallium 0 Chromium 0.02 Titanium 0.01 Cobalt 0 Tungsten 0.02 Lanthanum 0 Vanadium 0 Lead 0.11 Zinc 0.02 Magnesium 0.09 Zirconium 0 All chalcopyrite and pyrite samples were mounted, with one surface exposed, in a cold-setting resin. T h e exposed face was metallographically polished with 3 um alumina, Methodology and Experimental Procedure 54 followed by washing with water and acetone to remove any species that remained on the electrode surface after polishing. 3.1.2. Apparatus The electrochemical polarization was accomplished with a potentiostat. A Schlumberger Solartron 1286 potentiostat was used in the present study. Data were collected by a personal computer via a commercial software package called CorrWare for Windows. The software is designed to provide potentiostat control, data acquisition, and on-screen real time graphics. The companion software, CorrView, was used for post-test data analysis and charting. The electrochemical kinetics is generally investigated using a standard three-electrode cell. The cell consists of a working electrode, an auxiliary electrode, and a reference electrode. The auxiliary electrode supplies the current to the working electrode in order to polarize it. Figure 3.1 illustrates schematically a typical experimental set-up. The cell itself was made of glass and had a capacity of 1 L. The cell was immersed in a water-bath in order to maintain the desired working temperature constant. The cover of the leaching vessel had five openings, which allowed insertion of the counter electrodes, the working electrode, the Luggin capillary connecting the cell with the reference electrode compartment and the tube for nitrogen or oxygen purging. The experimental work was performed with chalcopyrite and pyrite electrodes, which had respective surface areas of 1.45 and 1.5 cm 2. Platinum electrodes were used as counter-electrodes. A saturated calomel electrode (SCE), with a standard potential of 0.24 V relative to the standard hydrogen electrode (SHE) was used as the reference electrode, for all electrochemical tests. Methodology and Experimental Procedure 55 Personal computer purge tube (N2 gas) j saturated Calomel reference electrode working electrode Pt counter electrode magnetic stirrer Electrochemical Cell Figure 3.1. The electrochemical set-up Methodology and Experimental Procedure 56 3.1.3. Procedure The procedure for anodic potentiodynamic experiments consisted of preparing an acidic solution at the required concentrations, placing the solution and the electrode in the cell, and de-aerating the solution by passing nitrogen gas through the solution for at least 30 minutes before each test to eliminate the effect of oxygen. A magnetic stirrer agitated the solution. Polarization curves were obtained by continuously measuring the current as a function of potential between 0 and 0.8 V ( S C E ) starting at the more negative potential. This potential range was selected because practical leaching of chalcopyrite usually takes place in the potential range of 0.2 to 0.6 V ( S C E ) [34]. T h e scan rate was 1 m V s _ 1 unless otherwise noted. The anodic characteristics of chalcopyrite were studied in acid solution (pH 1.5) at 25, 45 and 6 5 ° C . T h e choice of the working temperature was based on the type of microbes generally used in bioleaching, namely mesophiles, moderate thermophiles and extreme thermophiles. It has been observed in practice that mesophiles optimize their metabolism within a fairly narrow temperature range ( 2 5 - 3 5 ° C ) , moderate thermophiles are active between 40 and 5 5 ° C and above 5 5 ° C extreme thermophiles are more active. In order to evaluate accurately the behavior of chalcopyrite at specific potential values and to investigate the parameters affecting the stability of passive layers, potentiostatic experiments were conducted. The chalcopyrite specimen was held at constant potentials for 24 hours in deaerated solutions and soluble species of C u and Fe were analyzed by atomic absorption spectroscopy technique (AAS). Potentiodynamic experiments were also performed on chalcopyrite electrodes polarized for 24 hours in the potential region of severe passivation. The mixed potentials of chalcopyrite were measured for different ferric to ferrous ratios on "fresh" chalcopyrite electrodes and on polarized chalcopyrite electrodes. Mixed potential measurements were made in acidic solutions of various concentrations of ferric and ferrous ions. The mixed potential of the chalcopyrite electrode with respect to the reference electrode was measured by determining the potential difference between the two electrodes. Methodology and Experimental Procedure 57 3.2. Chemical leaching experiments 3.2.1. Chalcopyrite samples The chalcopyrite used for the controlled potential experiments in stirred tanks was obtained from Santa Eulalia Mining District, Mexico. The chemical and mineralogical analyses of the massive sample are presented in Tables 3.1 and 3.2. The mineral received from the mine was crushed, ground and screened to the required size fractions. T h e standard size fraction employed for most tests was 38-75 pm. 3.2.2. Reagents Distilled water and reagent grade chemicals were used to prepare the desired leaching solutions. The reagents grade chemicals, listed below, were used as received, without further purification. > Iron(ll) sulfate heptahydrate [ F e S 0 4 . 7 H 2 0 ] (99% purity) > Iron (III) sulfate pentahydrate [Fe 2 (S0 4 )3 . 5H 2 0] (97% purity) > Hydrogen peroxide [H 2 0 2 ] (30% solution) > Sulfuric acid [ H 2 S 0 4 ] (96% purity) 3.2.3. Apparatus Particulate leaching tests were carried out in a covered 3-L glass jacketed bioreactor equipped with a motor and speed controller manufactured by Applikon Dependable Instruments (ADI) (Schiedam, The Netherlands). Figure 3.2 illustrates the components of the set-up used for the stirred tank leaching tests. T h e ADI Bioreactor is a digital controller capable of monitoring and controlling four parameters throughout the experiments: acidity (pH), solution redox potential, temperature and rate of agitation. T h e desired temperature was controlled within 1 ° C by circulating water through the reactor jacket. T h e reactor was completely sealed in order to reduce evaporation losses during leaching. Particles were suspended at 750 rpm. T h e controller controlled the valves that inject the hydrogen peroxide and the sulfuric acid solutions to the reactor at a pre-defined solution redox potential and p H . The temperature compensator of the Methodology and Experimental Procedure 58 controller enables the measurement of pH and solution redox potential at its actual temperature. A Corning pH electrode (model 477006) was used for pH measurements. The probe was 25 cm long for use in deep containers. T h e temperature range of the probe was 0 to 1 0 0 ° C . Th e 3M KCI filling solution was changed every 1-2 months. Combination redox electrodes with Calomel and Double Junction Ag/AgCI reference half-cells (with a standard potential of 0.243 V relative to the standard hydrogen electrode) were used for redox measurements. T h e 4 M KCI filling solution was changed regularly. Cole-Parmer Masterflex peristaltic pumps were used for H2O2 and H 2 S O 4 solution additions into the reactor. Stirrer Condenser H202 port (tubing to pump not shown) Water Stirrer »?02 jacket controller purnp w ? # f Sparger pH and E h probes (wiring to controller not shown) Figure 3.2. Schematic representation of the controlled-potential chemical leaching system Methodology and Experimental Procedure 59 3.2.4. Procedure Controlled-potential leaching experiments were carried out in duplicate with 30 g of chalcopyrite ore added to 1500 mL of acid solutions. T h e retention time was 24 hours. The redox potential in the leaching solution was maintained to within 2 m V by the addition of 2% hydrogen peroxide, the amount of the addition being monitored by a digital scale. The ferric to ferrous ratios and temperature varied from 0.1 to 20, and 25 to 6 5 ° C respectively. The choice of these ferric to ferrous ratios was made after measuring the solution potentials at various ferric to ferrous ratios. Nernst's Equation (Eq. 3.1) was assumed to show the relationship between the ferric to ferrous ratio and the redox potential of the solution. E = E° + *T\n ' F e3 V a F e 2 + J (3.1) Using molal concentrations instead of activities, E q . (3.1) is rearranged to: £ ° = £ ° ' + — I n ' F e3 + (3.2) V " ' F e 2 V where £ ° = E ° + — I n ' F e3 + V ' F e 2 V (3.3) and YI is the activity coefficient for species /'. The solution potential increased logarithmically with the ratio of Fe(lll) to Fe(ll) with slopes of 0.059 as indicated in Figure 3.3. This value is exactly the same as the theoretical value of the Nernst's equation (59 mV/decade). T h e solution potential at the concentration ratio of 1:1 was about 0.71 V (SHE) . T h e difference between the experimental value of the solution potential (E°' = 0.71 V) and the theoretical value ( £ ° = 0.77 V) at a ratio of 1:1 in Nernst's equation is due to complexation of ferric and ferrous ions with sulphate and bisulphate in acidic solutions. Methodology and Experimental Procedure 60 A low solid to liquid ratio (2% w/v) was chosen so that the addition of hydrogen peroxide during the reaction would not change the total volume of the solution significantly. Direct oxidation of chalcopyrite by hydrogen peroxide was assumed to be negligible due to the low solid-liquid ratio, and the fast oxidation reaction with ferrous ions. The standard conditions used in this study were: > Total iron: 5 g/L > Ratio F e 3 + / F e 2 + = 1 > pH = 1.3 > Temperature = 6 5 ° C > Mass of chalcopyrite: 30 g > Volume of solution: 1.5 L 0.7 0.4 -I 1 1 1 . 1 - 1 0 1 2 3 4 log ([Fe3+]/[Fe2+]) Figure 3.3. Solution potential as a function of the ferric to ferrous ratio, pH 1.4, 2 5 ° C Methodology and Experimental Procedure 61 3.2.5. Sampling and analysis Samples of the solution were taken during the reaction with a 20 mL syringe. Soluble copper and iron were determined by atomic absorption spectrophotometry (AAS), and leaching curves were generated from the analytical results and sampling times. For the calculation of the conversion, correction factors were applied to account for the volume of peroxide solution added and mass losses due to sampling. T h e addition of hydrogen peroxide to the leaching vessel was monitored as a check on the copper concentration in solution. The addition of peroxide was always found to be stochiometrically proportional to the copper concentration. 3.2.6. Oxidation of Fe(ll) to Fe(lll) by hydrogen peroxide Preliminary tests indicated that during the first few minutes of leaching there is a rapid drop in the redox potential of the solution. This is due to the release of ferrous ions and the reduction of ferric ions according to reactions 2.8 and 2.9. Therefore, it is difficult to determine the contribution of chemical leaching by ferric ions to the overall leaching reaction. In order to obtain good kinetic data, the concentration of the reacting species should be kept almost constant during the course of the reaction. In practice, this is usually achieved by ensuring that the reactants are present in stoichiometric excess. O n the other hand, if the ratios ferric to ferrous (redox potential) are controlled at constant values, chemical, electrochemical and bioleaching experiments can be compared. Constant redox potential of the solution was maintained by controlled addition of hydrogen peroxide, continuously oxidizing F e 2 + ions and regenerating F e 3 + in a cyclic manner. Hydrogen peroxide is completely miscible with water. It is a powerful oxidant, as demonstrated by the E° value of 1.76 V (SHE) measured for the half-reaction: W h e n hydrogen peroxide is added to a ferrous solution, the following reaction takes place: H 2 0 2 + 2 H + + 2 e - = 2 H 2 0 (3.4) H O O H + 2 F e 2 + + 2 H + = 2 F e 3 + + H 2 0 (3-5) Methodology and Experimental Procedure 62 The oxidation rate is given by: Rate = k [Fe 2 + ] [ H 2 0 2 ] (3.6) The rate constant is 41.4 M _ 1 s _ 1 at 2 0 . 2 ° C with small variation for acidity and ionic strength [106]. 3.3. Bioleaching experiments 3.3.1. Material The chalcopyrite used in bioleaching tests was essentially the same as that used for the ferric leaching tests. The mineralogical composition and chemical analysis of the bulk sample are given in Tables 3.1 and 3.2 respectively. T h e mineral was pulverized in a ring grinder to 80% passing 35 um with a mean diameter of 22 um. 3.3.2. Microorganisms Mesophiles Mesophilic microorganisms containing strains of Leptospirillum and acidithiobacillus ferrooxidans and thiooxidans, cultured at 3 0 ° C , were used for this study. The nutrient contained the following salts: 2 g/L M g S 0 4 . 7 H 2 0 , 0.8 g/L ( N H 4 ) 2 S 0 4 l 0.3 g/L K H 2 P 0 4 l 1 g/L F e S 0 4 . 7 H 2 0 . Microbial cultures were grown in 250 mL Erlenmeyer flasks with 100 mL of the medium, 2 g of concentrate and 0.1 g of sulphur powder as energy source. The pH was adjusted to 1.8 using 6 M sulphuric acid. The solution was then inoculated with 25 mL of the previous culture. The old flask was well stirred when pouring into the new flask so that fine solids were also transferred. The new flask was covered with a foam stopper and placed in the incubator at 3 0 ° C and 100 rpm. Every week a new preparation was made in the same way in order to obtain enough cultures to inoculate the bioleaching experiments. Methodology and Experimental Procedure 63 Moderate and extreme thermophiles For moderate thermophilic leaching, a mixed culture containing strains of Sulfolobus, Acidianus and Sulfobacillus cultured at 4 8 ° C was used. For extreme thermophilic leaching, a mixed culture of containing strains of Sulfolobus, Acidianus and Metallosphaera cultured at 6 8 ° C was used. The growth medium for moderate/extreme thermophiles contained the following salts: 0.7 g/L M g S 0 4 . 7 H 2 0 , 1.5 g/L (NH 4 )2S0 4 , 0.3 g/L K H 2 P 0 4 , 0.5g/L F e S 0 4 . 7 H 2 0 . The pH was adjusted to 1.8 with sulphuric acid. The stock culture was prepared in the same way as described in the previous paragraph. The temperature inside the incubator was maintained at 4 8 ° C for the moderates and 6 8 ° C for the extreme thermophiles. A stirring speed of 150 rpm was used in the two incubators. 3.3.3. Equipment A n orbital incubator shaker from Lab-line Instrument Inc. (model 3527), was used for mesophilic leaching. The incubator holds 21 flasks and has a controller for the speed (40 to 400 rpm) and the temperature (25 to 6 5 ° C ) . Moderate thermophilic leaching was conducted in an orbital incubator shaker, model Innova 4300, from New Brunswick Scientific. The incubator was designed to handle up to 50 flasks. The temperature inside the incubator could range from ambient to 6 0 ° C . The stirring speed ranged from 25 to 500 rpm. Extreme thermophilic leaching tests were performed in a Giromax 737 orbital incubator from Amerex Instruments Inc. The incubator has a digital display controller for the temperature (25 to 8 0 ° C ) and the stirring speed (20 to 460 rpm). 3.3.4. Procedure Bioleaching tests were conducted in duplicate in 250 mL conical flasks containing 150 mL of leach solution. The pulp density used in all test work was 5% (w/v). The progress of the leach was followed by daily measurement of pH levels and redox potentials of the solution. Distilled water was added to the flasks in order to compensate for evaporation Methodology and Experimental Procedure 64 losses. Liquid samples (5 mL) were taken periodically for analysis of C u and Fe by atomic absorption spectrophotometry. O n c e bioleaching tests were finished, solids were removed by filtration and analyzed for the presence of copper, iron, elemental sulphur, sulphide and sulphate. 3.3.5. Cell counting The determination of the number of cells in solution and attached to solids was conducted in the Department of Earth and O c e a n Sciences at the University of British Columbia. T he procedure for liquid samples consisted of taking a liquid sample of a precise volume form the well-mixed sample, and filtering it onto a 0.2-um pore size, 25-mm diameter Anodisc filter. Cells collected on the filter were stained with the nucleic acid specific fluorochrome S Y B R Green 1 (Molecular Probes, O R ) . E a c h filter was incubated with 70 uL of dye (0.3 % S Y B R Green I, made up in 0.02-um filtered Milli-Q water) for 20 min, then mounted onto a glass slide with a small drop of 0.1% p-phenylenediamine and stored at 4 ° C , or - 2 0 ° C , if not handled immediately. The number of cells per field of view was counted directly using an epifluorescence microscope with 1000* magnification. About 30 fields of view and a minimum of 200 cells were counted per slide. The procedure for solid samples consisted of detaching the cells from their solid substrate by first adding 10 mL of extraction buffer, and topping it up with wash buffer to yield a ratio of 1 mL liquid per 1.5 g of ore. Th e slurry was vortexed vigorously for 5 s, followed by 30 min mixing at 225 rpm in a rotary orbit shaker. Th e supernatant was transferred into a sterile plastic tube and set aside. Cells in the supernatant were enumerated using the same protocol described above. 3.4. Column tests Small column tests were conducted at 6 8 ° C with fine particles of high grade chalcopyrite coated on low-grade chalcopyrite mineral. T h e G e o C o a t process, developed by Geobiotics, involves spray-coating a ground mineral concentrate onto relatively coarse inert or low-grade rock and then heap bioleaching [16, 33]. This aims Methodology and Experimental Procedure 65 to utilize the advantages of heap leaching (low capital and operating costs) while avoiding long leaching times and mineral occlusion. The G e o C o a t process may also be advantageous in terms of operating a thermophile chalcopyrite heap leach [16]. T h e galvanic interaction between chalcopyrite and pyrite was also investigated using this process. 3.4.1. Material Ore concentrate The materials used for the small columns tests were the chalcopyrite mineral from Santa Eulalia (Mexico) and the pyrite from Huanzala mines (Peru) as identified in the previous sections. T h e mineralogical composition and chemical analysis of the concentrates are given in Tables 3.1, 3.2, and 3.3. T h e chalcopyrite and pyrite concentrates were pulverized in a ring grinder to 80% passing 35 and 26 um respectively. Support rocks The low-grade chalcopyrite used in this test came from the Escondida mine in Chile. The mineral composition of the Escondida ore is detailed in Table 3.4. The size fraction employed for the supporting rocks (Escondida mineral) was 5 to 10 mm. Table 3.4. Results of quantitative phase analysis (wt %) Quartz (S i0 2 ) 41.5 Chalcopyrite ( C u F e S 2 ) 0.4 Pyrite (FeS 2 ) 3.5 Muscovite (KAI 2(AISi 3Oio)(F, O H ) 2 ) 31.5 Albite (NaAISi 3 0 8 ) 5.9 Kaolinite ( A I 2 S i 2 0 5 ( O H ) 4 ) 12.6 Clinochlore ((Mg,Fe,AI) 6(Si, A I ) 4 O 1 0 ( O H ) 8 ) 3.6 Calcite (CaC03) 1.0 Total 100.1 Methodology and Experimental Procedure 66 3.4.2. Apparatus The column leach experiments were conducted in 4 mini columns, 10 cm in diameter and 50 cm high. The columns were arranged as indicted in Figure 3.4. The ore was resting on a layer of support pebbles atop a sieve plate through which an air supply line was protruding. The sieve plate was covered with filter paper acting as a seal to air flow. The sieve plate rested on a spacer atop the botom plate, which closed the column. Al effluents were colected in the center of the botom plate and drained from there into the receiving container. Each column was fed with the leaching solution from a feed container via a multi-channel peristaltic pump at a rate of 5 L rrf2 h~1. The feed solution was dripped onto the layer of plastic bals covering the ore in each column. The columns were immersed in a plastic tub filed with water and fitted with an electric immersion heater and a circulating pump as shown in Figure 3.5. The temperature of the water bath was maintained at 68°±1°C. Air, enriched with 1% (vol.) C 0 2 , was blown into the base of each column at a rate of approximately 360 mL min-1. concentrate on support rock plastic balls plastic balls gas sparger spacer ring f — g i e v e p| a { e pipe flange _ gas hose gas supply received solution Figure 3.4. Schematic drawing of a column Methodology and Experimental Procedure 67 Figure 3.5. T h e column leach apparatus 3.4.3. Procedure Coating and charging the support rock The coating slurry was made of 50 mL of de-ionised water and 100 g of chalcopyrite concentrate. For the second set of experiments, the effect of pyrite on the bioleaching rate of chalcopyrite was studied with 100 g of chalcopyrite mixed with 200 g of pyrite and 150 mL of water. The thick slurry was poured onto 2.5 kg support rock in a bucket. T h e bucket was rotated several times until the slurry evenly coated all of the support rock. The coated rocks were immediately charged into the column through a funnel and covered with plastic balls. Methodology and Experimental Procedure 68 Column operation The columns were rinsed with an acidic solution at pH 1.3 for the first 2 days, and then inoculated with 100 mL of inoculum. In order to prevent the ore from being flushed off the support rocks, small doses (10 mL) were injected into the top of the columns at ten minutes intervals. Collected solutions were sampled twice weekly. Before sampling, the received solution was combined with the remaining feed solution. Evaporation losses were made up with distilled water to a total circulation volume of 4 L per column. Sulphuric acid was added to maintain the pH at 1.3. This sulphuric acid addition was made to avoid the precipitation of iron in the form of jarosites which may retard the bioleaching process. This combined solution was sampled and then used as the new feed solution. Solution potential was measured with a platinum combination redox electrode with calomel and double-junction Ag/AgCI reference half-cells. A Corning pH electrode (model 477006) was used for pH measurements. T h e copper and iron concentrations in the liquid samples were determined by A A analyses. Column Shut Down At the end of the tests, the feed flow was disconnected and the column was left to dry for several hours. A large bucket was then placed under the column, the base plate unbolted, and the entire column charge was allowed to drop into the bucket. T h e support rock and the columns were rinsed repeatedly with deionised water, and the residual solid was filtered from the slurry. The resulting wash solid was dried, weighed and analysed. T h e final volume was recorded. 3.5. Atmospheric leaching 3.5.1. Chalcopyrite samples High grade chalcopyrite ores from Selwyn (Australia), and Temagami (Canada) were used for the atmospheric leaching tests. Table 4.8 shows the chemical analyses of the two ores in the condition in which they were received. T h e concentrate received from the mine was crushed, ground and screened to the required size fractions. T h e standard size fraction employed for most tests was 38-75 urn. Methodology and Experimental Procedure 69 3.5.2. Pyrite samples The effect of the source of pyrite was investigated with two pyrite ores coming from Huanzala mines (Peru) and Park City (USA). The chemical compositions of the Huanzala pyrite and Park City pyrite are given in Tables 3.3. and 3.5. Table 3.5. Chemical analysis of the Park City pyrite mineral Element Mass(%) C u 0.05 Fe 42.9 S(total) 47.93 S(elemental) 0.01 S(sulphate) 0.12 3.5.3. Reagents Distilled water and reagent grade chemicals were used to prepare the desired leaching solutions. The reagents grade chemicals, listed below, were used as received, without further purification. > Iron(ll) sulfate heptahydrate [ F e S 0 4 . 7 H 2 0 ] (99% purity) > Iron (III) sulfate pentahydrate [ F e 2 ( S 0 4 ) 3 . 5 H 2 0 ] (97% purity) > Sulfuric acid [ H 2 S 0 4 ] (96% purity) 3.5.4. Apparatus Atmospheric leaching tests were carried out in a covered 3-L glass jacketed bioreactor equipped with a motor and speed controller manufactured by Applikon Dependable Instruments (ADI) (Schiedam, The Netherlands). A similar set-up was described for the chemical leaching of chalcopyrite (section 3.2.3). T h e reactor was maintained at 8 0 ° C by circulating water through the reactor jacket. Particles were suspended at 750 rpm. Air or 0 2 were pumped to the reactor through air tubes fitted with spargers. Methodology and Experimental Procedure 70 Chapter 4 RESULTS AND DISCUSSION 4.1. Electrochemical leaching 4.1.1. Anodic behaviour of chalcopyrite in acidic medium Potentiodynamic experiments A convenient way to study the electrochemical reaction kinetics of conductive minerals is to plot the current-potential curves for the respective half-cells involved in the leaching reaction. The applied potential may have a direct effect on the kinetics of the reaction if slow discharge is involved and may also serve to stabilize intermediate product phases. The overall anodic reaction during leaching of chalcopyrite in acidic medium can be expressed by: C u F e S 2 C u 2 + + F e 2 + + 2 S ° + 4 e - (4.1) The anodic characteristics of chalcopyrite were studied in acid solution (pH 1.5) at 25, 45 and 6 5 ° C , as shown in Figure 4.1. The open circuit potential ( O C P ) was about 0.34 V ( S C E ) (0.58 V (SHE)) at 2 5 ° C . This value is very close to the rest potential of chalcopyrite reported in the literature (0.52 V (SHE) in 1 M H 2 S 0 4 ) [70]. This small difference is probably due to differences in crystal orientation, the presence of impurities or perhaps slight deviations in stoichiometry. The open circuit potential was decreasing with increasing temperature. The O C P at 6 5 ° C was approximately 50 m V more negative than that at 2 5 ° C . This lower O C P indicates that the chalcopyrite surface was becoming more active and less noble at higher temperatures. A s the scan proceeded in the positive potentials, the anodic polarization curves of chalcopyrite at various temperatures revealed three electro-active regions. T h e first region, from 320 to 450 mV ( S C E ) , indicates that the current increased exponentially with applied potential, with Tafel slopes ranging from 0.12 to 0.16 V d e c a d e - 1 . T h e plateau region in the potential range 450 to 600 m V ( S C E ) , indicates that the current levels out. T h e third region, at Results and Discussion 71 potentials higher than 600 m V ( S C E ) , reveals that the current increases rapidly with applied potential. The second portion of the curve is referred to as the "passive" region because the current does not increase significantly with increasing potential. T h e presence of this plateau region is attributed to the formation of a product layer, which eventually covers the chalcopyrite surface. This "passive" behaviour is different from that of typical passivating metals such as zinc and aluminium. For these metals, the reduced corrosion rate in the passive state may be as much as 10 6 times lower than the maximum rate in the active state [107]. The test for the anodic dissolution of chalcopyrite at 2 5 ° C was repeated two times under the same reaction conditions to examine the issues of reproducibility and it was observed that the polarization curves were similar as shown in Figure 4.1 The results from other electrochemical studies indicate that new phases are formed during chalcopyrite leaching and the rate-limiting step for the anodic reaction at low potentials is a surface reaction or diffusion through the new phases. This can be expressed by the following equation: C u F e S 2 -> New phase + m C u 2 + + n F e 2 + + p S ° + q e~ (4.2) Ammou-Choukroum [108] proposed that covellite (CuS) was the new phase formed on the chalcopyrite, leading to a decrease of the anodic current. Peters [109] supported this conclusion and proposed the following reaction from thermodynamic considerations: C u F e S 2 + 2 H + = C u S + F e 2 + + H 2 S (4.3) Oxidation of C u F e S 2 and C u S could then occur in parallel: C u F e S 2 = C u 2 + + F e 2 + + 2 S + 4 e" (4.4) C u S = C u 2 + + S ° + 2 e" (4.5) A recent study by Arce et al. [110] has shown that the chalcopyrite oxidation process does not produce covellite as claimed by Ammou-Choukroum and Peters. They Results and Discussion 72 conducted a comparative study of the electrochemical behavior of chalcopyrite, chalcocite and bornite in acidic solution and did not observe a cathodic peak associated with the covellite reduction on the reverse potential scan from voltammetric chalcopyrite oxidation. In addition, Lazaro et al. [111] confirmed the high reactivity of bornite and covellite in acidic solutions, which are therefore unlikely to be formed as intermediates in the oxidation of chalcopyrite under oxidizing conditions. 0.2 -I 1 ' 1 1 0.000001 0.00001 0.0001 0.001 0.01 I (A/cm2) Figure 4.1 Effect of temperature on the anodic dissolution of chalcopyrite, p H 1.5, scan sate = 1 mV s~1. Numerous studies have indicated the preferential release of iron from chalcopyrite and the formation of a defect chalcopyrite structure during the earlier stage of leaching [43, 73]. The defect structure is believed to cause chalcopyrite passivation. Kelsall and Page [112] suggested that the initial release of iron and the formation of a defect structure are non-oxidative dissolution processes. This conclusion was based on the fact that the immersion time prior to the measurements with the chalcopyrite electrode affects the reproducibility of the voltammetric study. They proposed the following reaction: Results and Discussion 73 C u F e S 2 + 2 x H + = CuFe (i_x)S ( 2_X) + x H 2 S + x F e : (4.6) The copper-rich phase may then be oxidized on the subsequent anodic potential sweep to release C u 2 + . Lazaro and Nicol [111], using a rotating ring-disk electrode reported the same observation. They indicated that there was a change in the initial open-circuit potential, with the O C P becoming more positive at longer immersion times. The surface became less reactive causing a decrease of the anodic current. In contrast, Warren et al. [70] suggested an oxidative process for the preferential dissolution of iron from chalcopyrite and also suggested that the passivating reaction is a low-potential process. They proposed the following reaction: C u F e S 2 -> Cu^xFd-yS^z + x C u 2 + + y F e 2 + + z S ° + 2 (x+y) e~ (4.7) where y > x. In order to clear up the confusion concerning the nature of the defect structure, several authors used surface analysis techniques to identify the products of the anodic dissolution of chalcopyrite. Hamilton and W o o d s [113], using X-ray photoelectron spectroscopy (XPS) studies of the chalcopyrite surface, suggested that the initial products of oxidation of chalcopyrite include a copper sulphide with stoichiometry close to C u S 2 . They also observed that copper was present in these species as copper(l). This implies an oxidation state for sulphur of -Vz, the same value found in the polysulphide ion S 4 2 ~ . This metal-deficient sulphide (or polysulphide) had different semi-conductor properties from the underlying chalcopyrite. Thus , his composition approaches C u 2 S 4 in electronic structure. Hackl [10, 79] reported similar conclusions. Compounds involving polysulphides are known to be poor conductors and have a passivating effect. It is evident from Figure 4.1 that increasing temperature had a favorable effect on the dissolution rate of chalcopyrite, particularly in the first and second regions of the anodic polarization curves. In order to find the boundary between the three regions and to identify the potential range of severe passivation, polarization curves at different scan rates were carried out. Three scan rates, namely 0.1, 1 and 10 m V s~1, were used to Results and Discussion 74 generate the polarization curves (Figures 4.2, 4.3 and 4.4). It appears from Figs. 4.2 to 4.4 that lower scan rates shift the anodic polarization curves to lower currents. This is because, when the scan rate is very low, the amount of product layer is more important and this layer produces a decrease in current density. Based on these results, a scan rate of 0.1 m V s - 1 was the best choice for examining the effect of temperature on the extent of chalcopyrite passivation. Typical anodic polarization curves at various temperatures are shown in Figure 4.5. At low temperature ( 2 5 ° C ) , passivation is evident by the plateau region and the low currents in the potential range 0.45 to 0.55 V ( S C E ) . At moderate temperature ( 4 5 ° C ) , the curve does not show a plateau region within the potential range 0.45 to 0.55 V ( S C E ) and the anodic currents are much higher. This indicates that the stability of the passive layer diminished as the temperature increased. At high temperature ( 6 5 ° C ) , the anodic current increases rapidly within the potential range 0.45 to 0.55 V ( S C E ) , with no clear indication of a flat region. This rapid increase in the anodic current indicates that chalcopyrite passivates at a lesser extent at high temperatures and low potentials. However, above 0.6 V there was a sharp increase in current analogous to the pitting corrosion of metals, particularly at 2 5 ° C . It is not yet clear if the high current generated at low temperature and high potentials is due to the breakdown of the passive layer or to other oxidation reactions such as the oxidation of elemental sulphur to sulphate and the oxidation of ferrous to ferric. Braithwaite and Wadsworth [13] have observed that higher temperatures increase the amount of sulphate formed. In addition the ferric oxidation is faster at high temperatures. Thus, it may be expected that higher currents should result if the temperature increases. However, it was also observed that there was inversion between the current densities as the temperature increased in the third region (above 0.65 V (SCE)) . Jones and Peters [72] obtained similar results at high temperatures (above 9 5 ° C ) and high potentials (above 1 V (SHE)) and they postulated that the reaction was "saturated". It is possible that cupric ions accumulate in the product layers formed at high temperatures. A s indicated earlier, the nature of the reaction products is speculative at this time. Results and Discussion 75 0.9 0.2 H 0.1 A 1.00E-06 1.00E-05 1.00E-04 1.00E-03 1.00E-02 I (A/cm2) Figure 4.2 Effect of scan rate on the anodic dissolution of chalcopyrite, pH 1.5, 2 5 ° C , de-aerated solutions Figure 4.3 Effect of scan rate on the anodic dissolution of chalcopyrite, pH 1.5, 4 5 ° C , de-aerated solutions Results and Discussion 76 0.9 0.2 -0.1 -0 1 , = — . . 1 1.00E-06 1.00E-05 1.00E-04 1.00E-03 1.00E-02 l(A/cm2) Figure 4.4 Effect of scan rate on the anodic dissolution of chalcopyrite, pH 1.5, 6 5 ° C , de-aerated solutions Figure 4.5 Effect of temperature on the anodic dissolution of chalcopyrite, pH 1.5, scan rate = 0.1 m V s - 1 Results and Discussion 77 Potentiostatic experiments Conditions during potentiodynamic polarization experiments are non-steady state. Hence, several authors have observed that protective coatings on the mineral surface do form after sufficient reaction has occurred [3, 21]. In order to evaluate accurately the time dependence and the stochiometry of chalcopyrite dissolution, constant potential experiments were conducted. The chalcopyrite specimen was held at constant potential for 24 hours within the potential range 0.45 to 0.65 V ( S C E ) and the current was recorded as shown in Figures 4.6 to 4.8. This potential range is situated in the plateau region of the anodic polarization curves, i.e., in the region where "passivation" occurrs. The time dependence at 2 5 ° C shows a drop of 80% in the current at 0.45 and 0.55 V ( S C E ) (Figure 4.6). At higher potentials, the drop in current was about 50%. These results indicate that a progressively thickening passive film is formed at low temperature and potential and that the nature of the film changes as the potential is increased. T o further delineate the nature of the passive layer at low temperature, anodic polarization experiments were carried out with chalcopyrite electrodes initially polarized at 0.45, 0.50 and 0.55 V ( S C E ) for 24 hours. Pretreatment of the chalcopyrite electrode by anodic polarization increased its open circuit potential with respect to that measured with a fresh electrode. T h e O C P of polarized surfaces was about 50 m V more positive than that of a "clean" surface. The exchange current densities also decreased for polarized surfaces. A s shown in Figure 4.9, the overall shape of the anodic curves for polarized chalcopyrite was the same, with each including dissolution, passive, and transpassive regions. The current densities of polarized chalcopyrite were considerably lower than those of a fresh surface. The polarized surfaces became less reactive causing a decrease of the anodic current. The most pronounced effect was observed when chalcopyrite was initially polarized at 0.55 V ( S C E ) . This indicates that there was a change in surface speciation during the anodic dissolution of chalcopyrite. It is likely that changing potential could produce very different surfaces or intermediates. The time dependence at 4 5 ° C shows a drop of 60% in the current at 0.45 V ( S C E ) (Figure 4.7). However, as the potential increases, current-time curves show an induction Results and Discussion 78 period during which the current decays and a period when the dissolution rate begins to increase. This induction period was mainly observed between 0.45 and 0.6 V ( S C E ) . This anomalous observation can be attributed to the instability of intermediate phases with temperature and potential. At higher potentials (> 0.65 V (SCE)) the current decays with time to a steady-state value, but still remains high. T h e time-current curves at 6 5 ° C show three different behaviors (Figure 4.8): > At low potential (0.45 V (SCE)) , the current decreases with time to a steady-state value. The drop in current was about 30%, which is much less than the drop in current observed at 25 and 4 5 ° C . It seems that the passive layers formed at 6 5 ° C are less stable than those formed at 25 and 4 5 ° C . > At medium potentials (0.5 V (SCE)) , the time dependence shows an induction period followed by an active dissolution of chalcopyrite. Similar curves were observed at 4 5 ° C . > At high potentials (0.6 to 0.65 V (SCE)) , there is a rise in current for the first hour followed by a drop in current as the leaching proceeds. T h e drop was more severe as the potential increased. It seems that the anodic dissolution of chalcopyrite at high temperatures and high potentials takes place in two steps: The initial step is the oxidation of chalcopyrite to produce C u 2 + , F e 3 + and S C v 2 - and the second step is the precipitation of ferric that hydrolyze to form iron-hydroxy compounds [6]. It is possible that, given the higher current, and hence, the higher amount of material forced to dissolve, resulting from an increase in the applied potential, the pH value in the vicinity of the electrode may increase and favor the precipitation of iron-hydroxy compounds (e.g. iron hydroxide, oxyhydroxides and jarosite). Gardner and W o o d s [114], postulated that the anodic dissolution of chalcopyrite was associated with the following reaction: C u F e S 2 + 3 H 2 0 = C u S + F e ( O H ) 3 + S ° + 3 H + + 3 e" (4.8) Results and Discussion 79 0 2 4 6 8 10 12 14 16 18 20 22 24 Time (H) Figure 4.6 Constant potential experiments at 2 5 ° C on chalcopyrite in acidic solutions, pH 1.5 Figure 4.7 Anodic behaviour of fresh and polarized chalcopyrite surfaces in acidic solutions at 2 5 ° C , pH 1.5 Results and Discussion 80 0.0002 0.00018 0.00016 0.00014 _ 0.00012 c~f | 0.0001 ~ 0.00008 0.00006 0.00004 0.00002 0 E = 0.6 V E = 0.55 V E = 0.45 V —I 1 I I 1 1 1 1 8 10 12 14 16 18 20 22 24 Time (H) Figure 4.8 Constant potential experiments at 4 5 ° C on chalcopyrite in acidic solutions, pH = 1.5 0.0004 0.00035 | 0.0002 0.00015 0.0001 0.00005 • • • • • B-B-B E = 0.55 V E = 0.45 V 0 2 4 6 8 10 12 14 16 18 20 22 24 Time (H) Figure 4.9 Constant Potential Experiments at 6 5 ° C on Chalcopyrite in Acidic Solutions, pH = 1.5 Results and Discussion 81 T o further confirm that the anodic currents generated at constant potentials were proportional to the amount of copper and iron released, solution samples were collected at the end of each potentiostatic experiment and then analyzed by atomic absorption spetrophotometry (AAS). Results are summarized in Table 4.1. Table 4.1 Analysis of solutions after potentiostatic experiments on chalcopyrite Time (H) Potential (V (SCE)) T e m p ( ° C ) C u (10~ 6 M) Fe (10"6 M) Ratio C u : F e I (10 - 6 A) 24 0.45 25 23.3 62.0 0.4 11.6 24 0.55 25 25.0 75.1 0.3 19.8 24 0.60 25 59.0 55.5 1.1 36.5 24 0.65 25 80.9 90.8 0.9 55 24 0.45 45 81.4 111.2 0.7 55.8 24 0.55 45 90.5 79.3 1.1 74.8 24 0.60 45 117.0 100.8 1.1 119 24 0.65 45 149.7 180.8 0.8 165 24 0.45 65 276.0 251.1 1.1 132 24 0.55 65 332.0 260.9 1.3 240 24 0.60 65 486.2 376.1 1.3 322 24 0.65 65 893.2 523.3 1.7 399 Table 4.1 confirmed that the anodic currents measured were indeed proportional to the amount of copper dissolved. Analysis of the solution indicated that iron was preferentially dissolved at low potential and low temperature. A ratio of dissolved copper to iron of 1:3 was obtained within the potential range 0.45-0.55 V ( S C E ) . The preferential dissolution of iron in the first region clearly indicates the passivation of chalcopyrite by an iron-deficient sulphide layer. Potentiostatic tests have shown that the iron-deficient sulphide layer dissolves in the potential range 0.6 to 0.65 V ( S C E ) , with the release of equal amounts of copper and iron to the solution. Increasing temperature had also a favorable effect on the dissolution of the iron-deficient sulphide layer. Equal amounts of copper and iron were released at 0.55 V ( S C E ) (Cu:Fe =1:1), at 45 and 6 5 ° C compared with only 1:3 at 2 5 ° C . At high temperature ( 6 5 ° C ) and high potential Results and Discussion 82 (0.65 V), the amount of copper released in solution appears to be much higher than iron. This is probably due to the precipitation of ferric as mentioned earlier. Based on these results, the three electro-active regions can be described by the following reactions: First region (0.3 to 0.45 V ( S C E ) ) : C u F e S 2 -> Cui_xFei_yS 2-z + x C u 2 + + y F e 2 + + z S ° + 2 (x+y) e" (4.8) where y > x. Second region (0.45 to 0.6 V (SCE)): Cui_ x Fei_yS 2 -z--> (2-z)CuS ( n-s) + ( z - x - 1 ) C u 2 + + (1-y) F e 2 + + 2 (z-x-y) e" (4.9) where CuS( n- S) represents a nonstoichiometric intermediate product phase. Third region (above 0.6 V (SCE)): C u S -> C u 2 + + S ° + 2 e " (4.10) It appears that several intermediate products can be formed as the potential increases. These results confirm the suggestion that the properties and stability of the passive layers are temperature and potential dependent. 4.1.2. Cathodic reduction of ferric on chalcopyrite The true leaching rate in electrochemical processes can be predicted based on the mixed potential theory, i.e., by combining the cathodic and anodic branches of the half-cell reactions. Two kinds of cathodic reactions are observed on sulphide minerals: > Cathodic reduction of oxidants, such as F e 3 + and 0 2 > Cathodic reduction of the mineral itself due to the presence of a strong reducing agent. The present study will focus mainly on the cathodic reduction of ferric: Results and Discussion 83 F e 3 + + e -> Fe : (4.11) In order to avoid the contribution of the anodic dissolution of chalcopyrite to the total current, the reduction of ferric was studied at potentials that were lower than the rest potential of chalcopyrite. T h e effect of the applied potential on the rate of reduction of ferric on chalcopyrite is shown in Figure 4.10. The cathodic currents increased with increasing ferric concentrations. The limiting current was not reached within the potential range selected for the tests, i.e., 0 to 0.35 V ( S C E ) . A s mentioned in the previous section, the oxidation of chalcopyrite takes place through a series of intermediate phases, depending on temperature and applied potentials. Since these intermediate products are known to be semi-conductors [11, 26, 78], their presence may also influence the reduction of ferric. In order to test whether the formation of polysulphide layers at various temperatures and applied potentials can influence ferric reduction, and consequently the overall process, the chalcopyrite surface was polarized for 20 hours before each experiment. Figure 4.10 Reduction of Fe(lll) on chalcopyrite as a function of applied potential at various concentrations of Fe(lll), pH 1.5, 25 ° C Results and Discussion 84 0.000001 0.00001 0.0001 0.001 I (A/cm2) Figure 4.11 Reduction of Fe(lll) on polarized chalcopyrite, pH 1.5, 0.001 M Fe(lll) It is evident from Figure 4.11 that the reduction of ferric is much slower when the surface has been polarized. The polysulphide layers formed at high potentials and low temperatures strongly inhibit the rate of ferric reduction. T h e simplest explanation for the slow kinetics would be the formation of a protective layer, which impedes the transport of ferric to the chalcopyrite surface, the transfer of electrons between ferric and chalcopyrite, and the transport of cupric from the chalcopyrite surface. It follows that the slow kinetics for the ferric reduction on polarized chalcopyrite surface is a contributing factor for chalcopyrite passivation. In order to improve the understanding of the ferric reduction, we decided to conduct some experiments with a pyrite electrode. It is well known that the rest potential of chalcopyrite is lower than that of pyrite in acidic solution (Table 2.4). Therefore, when the two minerals are in contact in a ferric acidic medium, chalcopyrite will act as an anode, while the ferric reduction will take place on the pyrite surface. It was thus of Results and Discussion 85 interest to test the response of a pyrite electrode in ferric sulphate solutions by conducting polarization experiments. 0.00001 0.0001 0.001 0.01 l(A/cm2) Figure 4.12 Reduction of Fe(lll) on chalcopyrite and pyrite, pH 1.5, 0.01 M Fe(lll) T h e cathodic reduction of ferric on pyrite and chalcopyrite are shown in Figure 4.12. By comparing the two polarization curves, it is apparent that reduction of ferric on chalcopyrite is characterized by slow kinetics. The ferric reduction on the pyrite surface was much faster. For example, the cathodic current for pyrite at 0.35 V ( S C E ) was 10 times higher than that observed for chalcopyrite at the same potential. Biegler [115] obtained similar results for the oxygen reduction on sulphide minerals. His measurements indicated that pyrite (FeS 2 ) and pentlandite ((Fe,Ni) 9 S 8 ) are the most electroactive sulphide minerals towards the cathodic reduction of oxygen in acid solutions. Furthermore, Peters and Majima [116] showed that pyrite is "passivated" by a protective layer of elemental sulphur when exposed to atmospheric oxygen, i.e., F e S 2 + 1 / 2 0 2 + 2 H + -> 2 S ° + F e 2 + + H 2 0 (4.12) Results and Discussion 86 Their experimental results showed that the normal rest potential of pyrite was 0.4 V higher than would be expected from the E h - p H diagram. This increase in the rest potential can extend the potential range for the cathodic reduction of pyrite and amplify the driving force of the galvanic reaction when pyrite is in galvanic contact with other sulphide minerals. These interesting results clearly demonstrate that, if pyrite can support significant cathodic currents at potentials just below its rest potential, it will enhance the overall leaching rate of chalcopyrite and other sulphide minerals with which it is in galvanic contact. 4.1.3. Mixed potential of chalcopyrite as a function of Fe(lll) and Fe(ll) concentrations A s mentioned in the previous section, it was difficult to study the reduction of ferric on chalcopyrite in the potential range 0.35 to 0.8 V ( S C E ) due to the contribution of the anodic current of the mineral to the total current. T h e observed currents (/0bs) at potentials higher than the rest potential of the mineral are the sum of positive anodic currents due to mineral oxidation, ia (mineral); the oxidation of ferrous present in the leaching solution, ia (ferrous); and the negative cathodic current for the reduction of ferric, ic (ferric): kbs = ia (mineral) + ia (ferrous) + ic (ferric) (4.13) However, the leaching of chalcopyrite takes place between 0.3 and 0.6 V ( S C E ) [34]. This difficulty of studying ferric reduction above the rest potential of chalcopyrite can be overcome by measuring directly the mixed potential in synthetic solutions containing various concentration ratios of Fe(lll) and Fe(ll). Chemical changes in the condition of the chalcopyrite surface can be followed by studying the effect of concentration ratios of Fe(lll) and Fe(ll) on the mixed potential. The mixed potential of the mineral electrode with respect to the calomel reference electrode was measured by determining the potential difference between the two electrodes. The effect of the concentration Fe(lll):Fe(ll) ratio on the mixed potential of chalcopyrite is shown in Figure 4.13. The mixed potential increased logarithmically with the Results and Discussion 87 Fe(lll):Fe(ll) ratio, with a slope of 0.040 V d e c a d e - 1 . T o further define the effect of ferric and ferrous on the intermediate phases, chalcopyrite was initially polarized at various potentials and the mixed potential was measured on those polarized surfaces. It was observed that the mixed potentials of polarized chalcopyrite were 40 m V lower than that of fresh surface. This result confirms the previous suggestion that the ferric reduction is much slower when the surface is polarized. Polarization of chalcopyrite within the potential range 0.45 to 0.55 V leads to low values of ic (ferric) and /'a (chalcopyrite) and thus slow oxidation of CuFeS2-A similar series of experiments was performed with a pyrite electrode in the presence of ferric to determine the effect of pyrite on the leaching rate of chalcopyrite. The results are shown in Figure 4.14. It was observed that the mixed potential of chalcopyrite was 100 mV lower than that of pyrite. For ferric concentrations in the range of 0.01 to 0.1 M the mixed potential of chalcopyrite was located in the region of severe passivation on the chalcopyrite anodic curve. It is important to note that severe passivation was observed between 0.45 and 0.55 V at 2 5 ° C . This implies that an increase in the ferric concentration from 0.01 M to 0.1 M would not affect significantly the leaching rate of chalcopyrite, as the mixed potential is located in the "passive" region. In contrast, the mixed potential of pyrite was located at much higher potentials. Therefore, the presence of pyrite in a chalcopyrite ore might stimulate ferric reduction and create more oxidizing conditions favorable for chalcopyrite leaching. Results and Discussion 88 Figure 4.13 Mixed potential of fresh and polarized chalcopyrite as a function of the Fe(lll):Fe(ll) ratio, pH 1.5, 2 5 ° C 350 0.001 0.01 0.1 Fe(lll) (M) Figure 4.14 Mixed potential of chalcopyrite and pyrite as a function of the concentration of Fe(lll) at constant Fe(ll), pH 1.5, 2 5 ° C Results and Discussion 89 Electrochemical measurements have demonstrated that chalcopyrite "passivation" is mainly observed under the following conditions: > Low temperature ( 2 5 ° C ) and low potentials (0.45 to 0.6 V) > High temperature ( 6 5 ° C ) and high potentials (above 0.65 V) Leaching of chalcopyrite was accelerated at high temperatures 6 5 ° C and mild oxidizing conditions (0.45 to 0.55 V), and that the presence of pyrite in the chalcopyrite leaching system can be beneficial. Based on these results a series of controlled potential experiments with fine particles was conducted in order to validate the electrochemical study and to determine intrinsic leaching kinetics. 4.2. Chemical leaching 4.2.1. Effect of solution potential on reaction rate Results of leaching chalcopyrite in acidic solutions containing various ratios of Fe(lll) to Fe(ll) concentration are shown in Figures 4.15 to 4.19. T h e s e experiments were conducted at temperatures varying from 35 to 6 5 ° C . Measurements at 35°C Figure 4.15 illustrates the effect of solution potentials on the extraction of copper from chalcopyrite under atmospheric conditions at 3 5 ° C . Leaching experiments were conducted in acid sulphate solutions. Two different Fe(lll):Fe(ll) ratios were tested, namely 1 and 15, corresponding to a solution potential of 478 and 550 mV (Ag/AgCI) respectively (Table 4.2). The potentials were kept constant by the addition of hydrogen peroxide. The amount of peroxide was also monitored as shown in Figure 4.16. Comparison between Figures 4.15 and 4.16 indicates that the addition of peroxide was proportional to the copper extraction. Results and Discussion 90 Table 4.2 Extraction data for the ferric leaching of chalcopyrite at 3 5 ° C Test ID Temp (°C) Leaching time (h) Fe(lll):Fe(ll) ratio Solution potential (mV (Ag/AgCI)) Cu extraction (%) S° formation (%) Final S° stoichiometry (%) 1 35 24 1 478 6.05 6.2 100 2 35 24 15 550 7.1 6.1 86 The results indicate that the copper extraction was not dependent on the Fe(lll):Fe(ll) ratio at 3 5 ° C for leaching test conducted in 24 hours. For a Fe(lll):Fe(ll) ratio of 1, copper extraction reached a value of only 6% in 24 hours. With an increase in the amount of ferric to a ratio of 15, copper extraction increased slightly to a value of 7%. C u leaching commenced with the release of acid soluble copper over the first hour and then slowed down unexpectedly between 2 and 6 hours. It seems that the ferric leaching of chalcopyrite exhibits two-stage kinetics at 3 5 ° C . The first stage shows a quasi-parabolic induction period for approximately 6 hours. During this stage, the leaching rate did not increase with an increase in the Fe(lll):Fe(ll) ratio. The reaction curves were practically identical. This is a typical case of passivation, showing a potential region where the copper extraction is independent from the potential. From this observation, the formation of a passive film seems to have slowed the leaching rate. The formation of this passive layer in the earlier stage of leaching is similar to that observed during the electrochemical study. It is then possible that the dissolution of chalcopyrite in this first stage is controlled by the formation of an unstable metal-deficient polysulphide film on the chalcopyrite surface. T h e second stage indicates linear kinetics at longer reaction time. A n increase in solution potential (Fe(lll):Fe(ll) ratio) resulted in a relatively small increase in the leaching rate of chalcopyrite. Results and Discussion 91 0.1 Time (H) Figure 4.15 Leaching rate curves of chalcopyrite in ferric sulphate solution, pH 1.4, 3 5 ° C 40 -y 0 5 10 15 20 25 Time (H) Figure 4.16 H 2 0 2 added during leaching of chalcopyrite in ferric sulphate solution, pH 1.4, 3 5 ° C Results and Discussion 92 The two-stage kinetics can be expressed by the following reactions: Stage 1: C u F e S 2 + 2 ( x + y ) F e 3 + -> C u ^ x F e ^ ^ - z + x C u 2 + + (2x+3y) F e 2 + + z S ° (4.14) Stage 2: Cih-xFei-ySa-z + (4-2x-2y) F e 3 + -> (1-x) C u 2 + + (5-2x-3y) F e 2 + + (2-z) S ° (4.15) The change from quasi-parabolic to linear kinetics may be attributed either to a change in the composition of the reaction product or to the instability of the passive layers with time. Table 4.2 also shows the portion of sulphide that is oxidized to elemental sulphur and the sulphur yield. Analysis of the residue after leaching indicated a sulphur yield of 88 to 100%. The results showed that the oxidation of chalcopyrite in ferric sulphate solutions generates a near stoichiometric yield of elemental sulphur. T h e elemental sulphur reaction product was not attacked by the ferric under the conditions used. According to Dutrizac [3], elemental sulphur is relatively inert in ferric leaching systems given that the temperature is kept below the melting point of sulphur. He reported 86 to 100% elemental sulphur formation during the ferric leaching of chalcopyrite. Some researchers claimed that the parabolic kinetics in the first stage results from the production of elemental sulphur at the chalcopyrite surface. In contrast, Munoz et al. [44] suggested that the leaching process is not limited by a transport process through the elemental sulphur product because it is not present initially. Kametani and Aoki [42] reported that a stoichiometric yield of elemental sulphur was formed during chalcopyrite oxidation, but the induction period resulted from the formation of an intermediate phase (CuS). A s discussed previously, it is believed that the sulphur layer is not responsible for the parabolic kinetics in the earlier stage of chalcopyrite leaching. Hirato and co-workers [37] also got a two-stage kinetics for the leaching of chalcopyrite in ferric sulphate solutions. They obtained a parabolic kinetics for approximately 100 hours and thereafter, the leaching rate increased and showed linear kinetics. In the present study, "parabolic" kinetics were observed for approximately 6 hours at 3 5 ° C . Results and Discussion 93 The difference between Hirato's results and our results may come from the fact that Hirato used a disk specimen rotating at 300 rpm, while we used fine particle chalcopyrite in suspension at 750 rpm. The kinetics in our system were surely enhanced by the fine particle size and the rapid agitation. Measurements at 45"C The dependence of the leaching rate on the Fe(lll):Fe(ll) ratio at 4 5 ° C is shown in Figure 4.17. Four Fe(lll):Fe(ll) ratios were used, namely, 0.33, 1, 5 and 20. The solution potentials corresponding to these ratios are given in Table 4.3. Comparison of Tables 4.2 and 4.3 shows that, for the same ratio, the solution potentials increased with increasing temperature. Table 4.3 Extraction data for the ferric leaching of chalcopyrite at 4 5 ° C Test ID Temp (°C) Leaching time (h) Fe(lll):Fe(ll) ratio Solution potential (mV (Ag/AgCI)) Cu extraction (%) s ° formation (%) Final S° stoichiometry (%) 3 45 24 0.33 450 10.9 8.7 80 4 45 24 1 488 18.9 14.0 74 5 45 24 5 540 39.8 32.3 81 6 45 24 20 588 27.0 22.8 84 The leaching of chalcopyrite in acid sulphate solutions containing ferric and ferrous at 4 5 ° C was much faster compared to the results obtained at 3 5 ° C . For a Fe(lll):Fe(ll) ratio of 1, copper extraction was about 18% in 24 hours, i.e., 3 times faster than that at 3 5 ° C . The results shown in Figure 4.17 and 4.18 indicated that the amount of peroxide added was proportional to the amount of copper extracted. The leaching curve also revealed two-stage kinetics. The first stage showed a parabolic induction period of about 4 hours. During this stage, increasing the Fe(lll):Fe(ll) ratio moved the copper extraction to slightly higher values. This implies that the presence of more ferric at 4 5 ° C disrupts the formation of the passive layer formed in the earlier stage of leaching. Electrochemical studies have shown that the intermediates formed at 2 5 ° C were more stable than those Results and Discussion 94 formed at 4 5 ° C . It is then possible that these two intermediates have different compositions. For the second stage the leaching rate increased rapidly with a rise in the Fe(lll):Fe(ll) ratio until it reached a maximum at a ratio of 5, after which there was a marked decrease in the rate. T he decrease in the rate of leaching with an increase of solution potential is in good agreement with previously reported work by Kametani and Aoki [42]. They observed that the maximum rate of chalcopyrite leaching was attained only over a very narrow range of solution potential. This potential range was situated between 0.4 and 0.43 V ( S C E ) . They also mentioned that a portion of C u was precipitated as C u S (covellite) at low potentials. Several authors [8, 36, 69, 75] have reported the presence of covellite as an intermediate product during the ferric leaching of chalcopyrite. Surface analyses have shown that the covellite intermediate was non-stiochiometric and was therefore different from the natural covellite mineral [42]. Based on these results, the two-stage kinetics at 4 5 ° C can be expressed by the following reaction: Stage 1: C u F e S 2 + 2 F e 3 + -> C u S + 3 F e 2 + + S ° (4.17) Stage 2: C u S + 2 F e 3 + C u 2 + + 2 F e 2 + + S ° (4.18) T he sudden decrease in the rate above a critical potential (or Fe(lll):Fe(ll) ratio) indicates that chalcopyrite passivation at high temperatures and high solution potentials takes place according to a different mechanism. It is also possible that the formation of iron precipitates, favored at high temperatures and high Fe(lll):Fe(ll) ratios, is responsible for the slow kinetics observed at a Fe(lll):Fe(ll) ratio of 20. Stott et al. (21) suggested that the passivation of chalcopyrite is partly due to the precipitation of iron-hydroxy compounds. Although the nature of this passive layer is not yet known with certainty, it has been observed consistently in this study that high potentials or increased ferric concentrations at higher temperatures promote rapid passivation of chalcopyrite. Significant amounts of elemental sulphur were formed, as shown in Table 4.3. The presence of elemental sulphur did not seem to inhibit the leaching of chalcopyrite. This Results and Discussion 95 confirmed our suggestion that the sulphur layer was not responsible for the induction period in the earlier stage of leaching. 0.45 25 Figure 4.17 Leaching rate curves of chalcopyrite in ferric sulphate solution, pH 1.4, 4 5 ° C Measurements at 65°C T h e leaching of chalcopyrite at 6 5 ° C was performed with 5 different Fe(lll):Fe(ll) ratios, namely, 0.1, 0.33, 1, 5 and 15. The values of the solution potentials corresponding to these ratios are given in Table 4.4. The copper extraction curves as a function of time at 6 5 ° C are shown in Figure 4.19. The amount of peroxide added to control the potential was directly proportional to the amount of copper extracted, except at high Fe(lll):Fe(ll) ratios (Figure 4.20). T h e high consumption of hydrogen peroxide at high Fe(lll):Fe(ll) ratios is possibly due to the oxidation of ferrous to ferric by hydrogen peroxide which is limited above 600 m V (Ag/AgCI) [106]. At low Fe(ll) concentration, the kinetics are very slow and hydrogen peroxide simply decomposes . Results and Discussion 96 Time (H) Figure 4.18 H2O2 added during leaching of chalcopyrite in ferric sulphate solution, pH 1.4, 4 5 ° C A s indicated in Figure 4.20, increased temperature enhanced the copper extraction, particularly at low solution potentials. For a Fe(lll):Fe(ll) ratio of 1, the copper extraction was about 48% in 24 hours; compared to 18% at 4 5 ° C and 6% at 3 5 ° C . Table 4.4 Extraction data for the ferric leaching of chalcopyrite at 6 5 ° C Test ID Temp (°C) Leaching time (h) Fe(lll):Fe(ll) ratio Solution potential (mV (Ag/AgCI)) Cu extraction (%) S° formation (%) Final S° stoichiometry (%) 7 65 24 0.1 444 17.9 15.3 86 8 65 24 0.33 486 44.0 37.0 85 9 65 24 1 520 47.7 42.1 88 10 65 24 5 580 33.8 27.8 82 11 65 24 15 597 32.1 31.0 96 Results and Discussion 97 Figure 4.19 Leaching rate curves of chalcopyrite in ferric sulphate solution, pH 1.4, 6 5 ° C Figure 4.20 H2O2 added during leaching of chalcopyrite in ferric sulphate solution, pH 1.4, 6 5 ° C Results and Discussion 98 The leaching of chalcopyrite did not show an induction period during the initial stage of leaching at 6 5 ° C . Clearly, the presence of an induction period is mainly a temperature dependent phenomenon, which takes place below 6 5 ° C . T he oxidation reaction at 6 5 ° C can be expressed by the following reaction: C u F e S 2 + 4 F e 3 + -> C u 2 + + 5 F e 2 + + S ° (4.19) T he leaching rate increased with an increase in the Fe(lll):Fe(ll) ratio until it reached a maximum at a ratio of 1:1. Above the ratio of 1:1 the copper extraction decreased significantly. Hiroyoshi et al. [117] proposed a reaction model to explain this phenomenon. They believe that at higher solution potentials, cupric are extracted more slowly from chalcopyrite by direct oxidation with ferric, according to reaction (4.19). At lower potentials, chalcopyrite is reduced by ferrous in the presence of cupric to form C u 2 S (Eq.(4.20)) and the intermediate C u 2 S is oxidized by ferric to release cupric (Eq. (4.21)) as: C u F e S 2 + 3 C u 2 + + 3 F e 2 + -> 2 C u 2 S + 4 F e 3 + (4.20) The accelerating effect can only take place if there are enough cupric ions in solution. The oxidation rate of the intermediate C u 2 S is higher than that of chalcopyrite and this causes the rapid copper extraction at lower potentials. A similar scheme was proposed by Wadsworth [9]. He reported that metallic iron additions increased the leaching rate of chalcopyrite. The iron has a lower potential than chalcopyrite, producing, on contact, a negative potential causing the chalcopyrite to react cathodically according to Equation C u 2 S + 4 F e 3 + -> 2 C u 2 + + S ° + 4 Fe : 2+ (4.21) (4-22): 2 C u F e S 2 + 6 H + + 2 e" -> C u 2 S + 3 H 2 S + 2 Fe : 2+ (4.22) T he anodic reaction is: F e ° F e 2 + + 2 e" (4.23) Results and Discussion 99 giving the overall reaction : 2 C u F e S 2 + 6 H + + F e ° -> C u 2 S + 3 H 2 S + 3 Fe : (4.24) However, Dutrizac ef al. [1], and Jones and Peters [38] have shown that ferrous additions suppress chalcopyrite leaching. In addition, in the present study, the leaching curve at very low potential (Fe(lll):Fe(ll) = 0.1) was characterized by a slow kinetics. This raises doubts about the proposed mechanisms of Hiroyoshi and Wadsworth. T h e decrease in the leaching rate above the Fe(lll):Fe(ll) ratio of 1:1 is probably due to precipitation as mentioned in the previous section. T h e yield of sulphur ranged between 82 and 96%. This result confirmed the observation that elemental sulphur is relatively inert in ferric. In order to elucidate the exact role of elemental sulphur in the passivation of chalcopyrite, the residue was rinsed in carbon disulphide ( C S 2 ) solution to remove sulphur. The residue collected after washing with carbon disulphide was leached again in ferric sulphate solutions. T h e results shown in Figure 4.21 indicate that the residues washed in CS 2 dissolve faster than those without a CS 2 treatment. Table 4.5 also indicates that the removal of sulphur does not increase significantly the overall leaching rate of chalcopyrite. T h e average amounts of copper released from the residues treated in C S 2 solutions were much lower than those from fresh mineral. This observation suggests that an intermediate sulphide phase beneath the sulphur layer might be the main element responsible for the slow kinetics of chalcopyrite leaching. Table 4.5 Results obtained during the leaching of chalcopyrite and chalcopyrite residues after 4 hours of leaching at 6 5 ° C . Fe(lll):Fe(ll) ratio C u dissolved from the residue not treated in C S 2 (mg L - 1 h"1) C u dissolved from the residue treated in C S 2 (mg L " 1 h- 1) C u dissolved from fresh mineral (mg L - 1 h"1) 0.33 8 17 264 1 10 20 250 5 12 32 228 Results and Discussion 100 0.1 0.09 0.08 -\ •a 0.07 u !» ~ 0.06 4) O 0.05 I 0.04 H •- R = 5, CS2 I - R= 1, CS2 •- R = 0.33, CS2 »— R = 5, No CS2 I — R = 1, No CS2 >— R=0.33, No CS2 Figure 4.21 Leaching curves of chalcopyrite residues rinsed and not rinsed in carbon disulphide solutions, pH 1.4, 6 5 ° C (R refers to the Fe(lll):Fe(ll) ratio). 4.2.2. Effect of temperature at constant Fe(lll):Fe(ll) ratio Figure 4.22 shows the effect of temperature on the leaching of chalcopyrite for a Fe(lll):Fe(ll) ratio of 1. It is immediately clear that temperature affects the induction period. At 3 5 ° C the induction period lasted about 6 hours, while at 4 5 ° C it only lasted about 4 hours. T h e leaching curves at 55 and 6 5 ° C did not show an induction period. A s discussed in the previous sections, the thermal breakdown of the layer causing this induction period may be linked to the thermal instability of polysulphides in solution. This result, together with the hypothesis that the induction period is caused by a copper-rich polysulphide, suggest that the nature and composition of the passive layers change with leaching time, temperature and solution potential. This is in good agreement with the results obtained from the electrochemical study. Results and Discussion 101 0.5 0 5 10 15 20 25 Time (H) Figure 4.22 Leaching rate curves of chalcopyrite in ferric sulphate solution at various temperatures, pH 1.4, Fe(lll):Fe(ll) = 1 4.2.3. Effect of pyrite on the ferric leaching of chalcopyrite In many ores, chalcopyrite is usually associated with other sulphide minerals, particularly pyrite. These associated minerals may affect the leaching of chalcopyrite in a variety of ways such as scavenging of reagents and galvanic interactions. Pyrite is of particular interest, since the electrochemical study (sections 4.1.2 and 4.1.3) has shown that its presence can enhance the overall leaching rate of chalcopyrite under specific conditions of temperature and potentials. In order to fully understand the combined effect of pyrite additions and solution potentials on the galvanic interactions between chalcopyrite and pyrite, several tests were conducted at 6 5 ° C . The choice of this temperature was based on the results of the previous section, which indicate no induction period above 5 5 ° C . Three series of experiments were conducted for this section: Results and Discussion 102 > Oxidation of pyrite at various potentials in order to find the electrochemical conditions under which pyrite is either inert or active in acid conditions. > Oxidation of chalcopyrite in the presence of pyrite at low potentials (when pyrite is inert) > Oxidation of chalcopyrite in the presence of pyrite at high potentials (when pyrite is oxidized) Effect of solution potential of the leaching rate of pyrite Th e leaching of pyrite in solution containing Fe(lll) and Fe(ll) at 6 5 ° C was conducted at low solution potential (Fe(lll):Fe(ll) = 1) and high solution potential (Fe(lll):Fe(ll) = 15). Th e oxidation of pyrite by ferric can be expressed by the following reaction: F e S 2 + 8 a H 2 0 + (2 +12c) F e 3 + (4.25) (3 + 12cr)Fe 2 + + 2 C T S 0 2 - + ( 2 - 2 a ) S ° + 1 6 C T H + where a is the sulfate yield. This equation indicates that pyrite dissolution generates ferrous, sulphur, sulphate and acid. Figure 4.23 indicates that pyrite is almost inert at low solution potentials. For a Fe(lll):Fe(ll) ratio of 1, iron extraction reached a value of 3% in 24 hours. With an increase in the amount of ferric to a ratio of 15, iron extraction increased to a value of 43%. It seemed that there was no sign of passivation at high Fe(lll):Fe(ll) ratios. The observation that pyrite does not leach significantly at a Fe(lll):Fe(ll) ratio of 1 has a practical implication for the galvanic interaction between chalcopyrite and pyrite. At low solution potentials, pyrite will not dissolve and the leaching rate of chalcopyrite may be enhanced due to the presence of cathodic sites on pyrite. At high solution potentials, the number of cathodic sites on pyrite will decrease and the beneficial effect of the galvanic interaction may stop. T o clarify this phenomenon, several experiments have been conducted with various amounts of pyrite mixed with chalcopyrite. Results and Discussion 103 0.5 Time (h) Figure 4.23 Leaching rate curves of pyrite in ferric sulphate solution, pH 1.4, 6 5 ° C Effect of pyrite on the leaching rate of chalcopyrite at low solution potential (Fe(lll):Fe(ll) = 1) T h e galvanic interaction between chalcopyrite and pyrite at low solution potential was conducted at 6 5 ° C with a Fe(lll):Fe(ll) ratio of 1. Pyrite of various amounts were added to chalcopyrite to produce F e S 2 : C u F e S 2 ratios of 0:3, 1:2, and 2:1 as shown in Table 4.6. Results and Discussion 104 Table 4.6 Leaching of chalcopyrite in the presence of pyrite at low solution potential Test ID 12 13 14 Chalcopyrite source Santa Eulalia Santa Eulalia Santa Eulalia Pyrite source Huanzala Huanzala Huanzala Mass of chalcopyrite (g) 30 20 10 Mass of pyrite (g) 0 10 20 Mean diameter of chalco particles (u.m) 47.6 47.6 47.6 Mean diameter of pyrite particles (u.m) 45.5 45.5 45.5 Density of chalcopyrite (g c m - 3 ) 4.2 4.2 4.2 Density of pyrite (g c m - 3 ) 5.0 5.0 5.0 Temperature ( ° C ) 65 65 65 Fe(lll):Fe(ll) ratio 1 1 1 Solution potential (mV(Ag/AgCI)) 520 514 511 Copper extraction after 24 hours (%) 47.7 57.8 81.1 0.9 0.8 0) • Chalco =30, Pyr = 0 Chalco = 20, Pyr =10 - Chalco = 10, Pyr = 20 1 0 Time(H) 1 5 20 25 Figure 4.24 Leaching rate curves of chalcopyrite at different F e S 2 : C u F e S 2 ratios, Fe(lll):Fe(ll)= 1, pH 1.4, 6 5 ° C Results and Discussion 105 Figure 4.24 shows copper extraction versus time at different ratios of FeS2 :CuFeS 2 . A s the ratios increase, the rate of copper extraction increases. For a ratio of 2:1, the dissolution rate of copper was increased by a factor of 3 during the first 6 hours of leaching. In the absence of pyrite, the anodic dissolution of chalcopyrite is coupled to the reduction of ferric, which must occur exclusively on the chalcopyrite surface. W h e n the chalcopyrite is in contact with the pyrite, the reduction of ferric is taking place on pyrite and on a portion of chalcopyrite. The dissolution rate of chalcopyrite increases because the rate of reduction of ferric is greater on the pyrite surface than on the chalcopyrite surface, as shown in the electrochemical study. In addition, the galvanic effect is more pronounced at high F e S 2 : C u F e S 2 ratios because pyrite provides more cathodic area for the reduction reaction, and the anodic dissolution current (rate) of chalcopyrite must increase to compensate. A schematic representation of the galvanic interaction between pyrite and chalcopyrite is shown in Figure 4.25. The chalcopyrite surface ( A C p ) c a n be divided into an anodic zone and a cathodic zone, the areas of which may be expressed, respectively, as follows: Aa,cP=0CpACp (4.26) A:,cp=(1-0cp)A;p (4-27) where 0Cp represents the fraction of the anodic area on chalcopyrite. Similarly, the pyrite surface (/\P y) can be divided into anodic and cathodic zones, respectively, thus: Vy=0pyA>y (4-28) A:,Py=(1-epy)A>y (4-29) where 0Py represents the fraction of the anodic area on pyrite. T h e total surface during the galvanic interaction is given by: A = ACp+APy (4-30) Results and Discussion 106 Fe 3+ Fe 2+ Fe 3+ Fe 2+ Chalcopyrite Anodic area CuFeS2=Cu2++Fe2+ + 2S +4e Chalcopyrite Cathodic area Fe 3 + + e-=Fe2+ Pyrite Cathodic area Fe 3 + + e" = Fe 2 + Pyrite Anodic Fes2=Fe2+ + 2S042" +16H++ 14 e" Figure 4.25 Schematic representation of the galvanic interaction between chalcopyrite and pyrite T h e total anodic and cathodic areas for the leaching of chalcopyrite in galvanic contact with pyrite can therefore be expressed by: \ - Ai.Cp + Ai,Py ~~ ^Cp^Cp + ^Py A>y (4.31) A, = Ac,Cp + 4,Py = (1 " #Cp )A* + (1 ~ #Py )4>y (4.32) The ratio of the anodic and cathodic areas during galvanic interaction can then be expressed by: Aa _ eCp^ cp+QpyA>y K (i-eCp)A;p+(i-epy)A>y (4.33) T h e F e S 2 : C u F e S 2 mass ratio can be expressed by: rTJpy Ppy^Py^Py m c P PcpA^Cp (4.34) Results and Discussion 107 where p P y is the density of pyrite, p C p is the density of chalcopyrite, d P y is the mean diameter of pyrite particles and d C p is the mean diameter of chalcopyrite particles. The pyrite surface can then be expressed by the following expression: Pcp^Cp m P y m C p Ppydpy (4.35) Substitution of Eqn . (4.35) into (4.33) gives: A , _ 0 C P + £ 0 p y Ac ( 1 - 0 C p ) + )8(1-0p y) (4.36) T h e effect of the F e S 2 : C u F e S 2 mass ratio on the galvanic interaction can be quantified using Equation (4.36). T h e results of the previous section (Figure 4.23) have shown that pyrite was inert at low solution potential (ratio Fe(lll):Fe(ll)=1). Based on this result, we can assume that the pyrite surface is mainly cathodic, i.e., 6Py=0. Using this value, E q . (4.36) gives: T he correlation between the surface areas and the amount of pyrite added during galvanic interaction, represented by E q . (4.37), can be conveniently followed by plotting AaIAc vs mpy/mcp. T h e values of p P y , p C p , d P y and d C p are given in Table 4.6. It is evident from Figure 4.26 that there is a notable effect on the anodic to cathodic area the surface ratio for a F e S 2 : C u F e S 2 ratio of 2:1 and 0Cp > 0.5. A small ratio of anode to cathode surface area is beneficial because the ferric reduction takes place favorably on a larger catalytic surface (pyrite surface). This is in good agreement with our results. Ac ( 1 - 0 C P ) + J3 (4.37) ratio with increasing the amount of pyrite. The curve shows a drop of more than 75% on Results and Discussion 108 Figure 4.26 Effect of increased FeS2:CuFeS2 ratio on the relative anodic and cathodic areas. Effect of pyrite on the leaching rate of chalcopyrite at high solution potential (Fe(lll):Fe(ll) = 15) The effect that galvanic interaction had on the leaching rate chalcopyrite at high solution potential was investigated at 6 5 ° C with a Fe(lll):Fe(ll) ratio of 15. Various amounts of pyrite were added to chalcopyrite to produce FeS2:CuFeS2 ratios of 0:3, 1:2, and 2:1 as shown in Table 4.7. Extraction data at high solution potentials showed that the addition of pyrite was detrimental to the leaching of chalcopyrite during galvanic interaction. A s shown in Figure 4.27, it was also observed that this effect was more pronounced at high FeS2:CuFeS2 ratios. The slow kinetics at high FeS2:CuFeS2 ratios may be attributed to a decrease in the cathodic surface area at high solutions potentials. Results and Discussion 109 Table 4.7 Leaching of chalcopyrite in the presence of pyrite at high solution potential Test ID 15 16 17 Chalcopyrite source Santa Eulalia Santa Eulalia Santa Eulalia Pyrite source Huanzala Huanzala Huanzala Mass of chalcopyrite (g) 30 20 10 Mass of pyrite (g) 0 10 20 Mean diameter of chalco particles (u,m) 47.6 47.6 47.6 Mean diameter of pyrite particles (urn) 45.5 45.5 45.5 Density of chalcopyrite (g c m - 3 ) 4.2 4.2 4.2 Density of pyrite (g c m - 3 ) 5.0 5.0 5.0 Temperature ( ° C ) 65 65 65 Fe(lll):Fe(ll) ratio 15 15 15 Solution potential (mV (Ag/AgCI)) 597 594 590 Copper extraction after 24 hours(%) 32.1 30.5 26.4 Chalcopyrite and pyrite both dissolve at high solution potentials (Fe(lll):Fe(ll) = 15). T h e number of cathodic sites of the system is therefore affected because of the presence of an anodic surface area on pyrite. This unusual behaviour can be explained in terms of current densities. The galvanic system will have a larger anode surface and the total anodic current density due to the galvanic interaction will be reduced. In addition, the total current in this case (high solution potentials) will be the sum of the current for the oxidation of chalcopyrite and the current for the oxidation of pyrite, while in the first case (low solution potential) the total anodic current was only due to the oxidation of chalcopyrite. It is therefore difficult to quantify the galvanic interaction of chalcopyrite and pyrite at high solution potentials, i.e., when both minerals dissolve. Results and Discussion 110 0.35 Figure 4.27 Leaching rate curves of chalcopyrite at different ratios of FeS2:CuFeS2, Fe(lll):Fe(ll)= 15, pH 1.4, 6 5 ° C Effect of chalcopyrite origin on leaching rate in the presence of pyrite at low solution potential (Fe(lll):Fe(ll) = 1) Warren et al. [70] and others have noticed that chalcopyrite from various sources exhibit different behaviour in acidic solutions. This could be due to differences in crystal orientation, the presence of impurities or perhaps slight deviations in stoichiometry. Based on this observation, and to further.confirm the beneficial effect of pyrite on the leaching rate of chalcopyrite at low solution potential, a series of experiments were conducted with high grade chalcopyrite ore from different origins, namely Santa Eulalia (Mexico), Temagami (Canada) amd Selwyn (Australia). Table 4.8 provides the chemical composition for the various chalcopyrite samples. All samples were ground to the following fraction: - 7 5 + 38 pm. Leaching experiments were carried out at 6 5 ° C with a Fe(lll):Fe(ll) ratio of 1. A FeS 2 :CuFeS2 ratio of 2:1 was used for the tests. T h e pyrite used for these experiments was obtained from Huanzala mines in Peru. The leaching response of chalcopyrite from various origins is presented in Figure 4.28. Results and Discussion 111 Table 4.8 Chemical composition of chalcopyrite from various sources Origin Cu (%) Fe(%) S (%) Santa Eulalia, Mexico 27.10 32.51 32.22 Temagami , C a n a d a 24.05 32.40 31.20 Selwyn, Australia 23.60 28.10 28.70 A s predicted by Warren ef al. [70], the leaching of chalcopyrite from various sources revealed significant differences. The Temagami sample exhibited the lowest leaching rate, with just 13% copper recovery in 24 hours. The Santa Eulalia sample showed slow kinetics in the absence of pyrite, but not to the same dramatic extent as observed for the Selwyn and Temagani samples. The addition of pyrite had a beneficial effect on the leaching rate of all chalcopyrite samples. The copper extraction of the Temagani sample was 4 times higher in the presence of pyrite. T h e Selwyn sample achieved total conversion in less than 24 hours in the presence of pyrite. T h e s e interesting results indicate that, whatever the origin of the chalcopyrite mineral, the ferric reduction on the chalcopyrite surface is the limiting step at high temperature. T h e addition of pyrite is a key factor in the "depassivation" of chalcopyrite because it can act as cathode and catalyze the ferric reduction reaction. Results and Discussion 112 Figure 4.28 Leaching rate curves of chalcopyrite at different ratios of FeS2 :CuFeS 2 , Fe(lll):Fe(ll) = 15, pH1.4, 6 5 ° C T he results of the chemical leaching tests have confirmed the following conclusions from the electrochemical studies: > T h e leaching of chalcopyrite is accelerated at high temperatures and low solution potentials. > Leaching in ferric sulphate solutions shows two-stage kinetics: an induction period followed by an active dissolution period. > The induction period depends mainly on temperature and not on potential. Increasing temperature had a favorable effect on the induction period. > At higher solution potentials, the ferric participate in reactions responsible for passivation. Results and Discussion 113 > The ferric reduction on chalcopyrite is a contributing factor for its passivation. > The addition of moderate amounts of pyrite accelerates leaching dramatically. Based on these observations, it was of interest to predict the behaviour of chalcopyrite in ferric sulphate using an electrochemical model. This is discussed in the next section. 4.2.4. Development of an electrochemical model for the dissolution of chalcopyrite Electrochemical models are generally based on the Butler-Volmer equation, which quantitatively describes reactions limited by charge transfer at an interface. The overall leaching reaction for the ferric or bioleaching of chalcopyrite may be expressed in terms of its respective half-cell reactions: Anodic: C u F e S 2 -> C u 2 + + F e 2 + + 2 S ° + 4e" (4.1) Cathodic: 4 F e 3 + + 4 e" = 4 Fe 2+ (4.11) Assuming that the back reaction of equation (4.1) is negligible, the currents for the anodic and cathodic half reactions may be represented by the equations: L.cp - 4/~A»,cp Kcp exp RT (4.38) 'c,Cp _ '"A.Cp / c c C p [Fe2 + ] exp RT - / c c , C p [Fe 3 + ] exp ( Vcp ( 1 - g e C p ) F E RT (4-39) where / a C p , / c C p , AaCp and AcCp are the anodic current, the cathodic current, and the anodic and cathodic surface area of chalcopyrite, respectively. The sum of anodic oxidation currents must equal the sum of cathodic reduction currents at the mixed potential, i.e., 'a,Cp c^,Cp (4.40) Results and Discussion 114 Substitution of Eqns . (4.38) and (4.39) into Eqn . (4.40) gives the following expression for the mixed potential: 4 F A 'a,Cp ka,CP e X P RT -FA, 'c,Cp / c c , D [Fe2 + ] exp V,Cp RT - / c c . C D [Fe 3 + ]exp V,Cp ( 1 - o r e , C p ) F E RT (4.41) It should be noted that, in the region of the mixed potential ( E m ) , the oxidation of minerals is highly irreversible, i.e., the rates of the reverse processes can be ignored. This is not valid for the oxidant ( F e 3 + in this case). Assuming transfer coefficients c r a C p and a c C p equal to 0.5, then: (4A,.cP*a.cp + A:,cpWF e l ) e x P [2RT) = A,cP^c,cp[Fe3+]exp FEm 2RT (4.42) This expression can be simplified to obtain the following expression for the mixed potential: E m = In £c,cA,cpFe3+] F 4 / \ a C p / c a C p + / c c C p / \ c C p [ F e ] (4.43) Substitution of Eqn . (4.43) into Eqn . (4.38) gives the following expression for the anodic current: / -AFk A I Aj.cpftc.cp[Fe3+] a C p a ' C p a C P l , 4^ C p ^ C p + / \ c .cp^,cp[Fe 2 + ] (4.44) The dissolution rate of chalcopyrite, - r C p , may be obtained by noting that the rate and the current density are related by stoichiometry, expressed by Faraday's law: Results and Discussion 115 where n C p is the number of moles of chalcopyrite. The effect of surface area during the dissolution of chalcopyrite can be conveniently followed if the rate of leaching is expressed in terms of the conversion of the mineral. The dissolution rate of the mineral may be related to the conversion by: dX^ = _^_dn^ dt n C p 0 dt where X C p and n C p 0 are respectively the conversion and the initial moles of chalcopyrite. • t ' Leaching is a heterogeneous reaction, which involves a diminishing surface area as the reaction proceeds. T h e change in the number of moles of chalcopyrite is related to the change in size of the chalcopyrite particles. Assuming a shrinking-particle model for spherical chalcopyrite particles with radius r, the conversion is related to surface area by: * c P = 1 - R) ( 4 - 4 7 ) r2 = R 2 ( 1 - X C p ) 2 ' 3 (4.48) / l C p = 4 7 r r 2 = 4 r r R 2 ( 1 - X C p ) 2 / 3 (4.49) where R is the initial particle size. Substitution of Eqns . (4.45), (4.49), (4.26) and (4.27) in E q n . (4.46) gives the following equation: Results and Discussion 116 ^T"^7 e c B , cl4eC p/<.,C p +(i-eC p)^ p[Fe-] ( 1 x " ] ( 4 5 0 ) or f ^ i H g l t y J 0 - ^ > ' F e " l , ( 1 - X c p ) - (4.51, eft n c p 0 C p 0 p p C p K C p + (1 -e 0 p ) [Fe 2 * ] 1 c>' where k^k^j-f*- and K C p = - ^ ^ c , C p ^ c . C p Integration of equation (4.51) will give the following expression for the conversion of chalcopyrite as a function of time: 4 T T R 2 (1-e c p )[Fe 3 + ] 3nCp,o C P C l V < c p + a - e c p ) [ F e 2 + ] (1_x )1/3=1_^e , y ^ ' V ' 2 t (4-52) 4.2.5. Development of an electrochemical model for the galvanic interaction between pyrite and chalcopyrite For this case the half-reactions are the anodic dissolution of chalcopyrite and pyrite, CuFeS 2 -» C u 2 + + Fe 2 + + 2 S° + 4e _ (4.1) FeS 2 + 8a H20 -> Fe 2 + + 2a SO 2 " + (2 - 2a) S° +16tr H + + (2 +12a) e" (4.53) and the cathodic reduction of ferric on the surface of both minerals: Fe 3 + + e~ -> Fe 2 + (on chalcopyrite) (4.11a) F e 3 + + e~ -> F e 2 + (on pyrite) (4.11b) The anodic current due to the oxidation of chalcopyrite (/a,cP) is the same as that given in the previous section (Eqn 4.38). Holmes and Crundwell [118] found that the anodic dissolution of pyrite was dependent on the pH and proposed the following expression for the anodic current of pyrite (slightly altered to reflect the stoichiometry of Eqn. 4.53): Results and Discussion 117 / a P y = ( 2 + 1 2 a ) F ^ P y / c a P y [ H + ] - 1 / 2 e x p X P y F E ^ RT (4.54) The currents due to the oxidation and reduction of dissolved iron on chalcopyrite and pyrite, given by Eqns (1.2a) and (1.2b), are described by: 'c,Cp - '"Aj.Cp kc.cD[Fe2+]exp fa^FE' *c,Cp 'c.Cp - WFe 3 + ]exp ( (1-ae,Cp)FE RT (4.39) ^c,Py "~ FAC Py / c c P y [ F e2 + ] e x p - S g — - / c c P y [ F e J + ] e x p (1-g c P y)FE RT (4.55) At the mixed potential: L.Cp + 'a.Py - Cc.Cp + c^,Py ) (4.56) Substitution of Eqns (4.38), (4.39), (4.54) and (4.55) into E q n . (4.56) gives the following expression for the system: 4 F A a , C p / c a i C p e x p <xa,cpFEm RT + (2 + 12a )F^ , P y / c a , P y [H + ] - 1 / 2 exp RT V:,Cp WFe 2 + ]exp + A 'c,Py / c c P y [Fe ' + ]exp RT fa^FE, ^ / c c C p [Fe 3 + ] exp ' c C p ^ ' ' - m RT *c ,Py ' * - m RT / c c P y [Fe 3 + ] exp ( (1-ac,Cp)FE, ^ f (1-ac,Py)FE, ^ RT (4-57) If the approximation is made that araCp = c r a P y = o r c C p = a c P y =0 .5 , this expression can be simplified to obtain the following expression: E RT^ ( ^ , C p A c , c P + ^ . P y 4 , P y ) F e J + ] m F n 4 / \ a , C p / c a C p + (2 + 1 2 a ) A , , P y / c a , P y [ H T ° 5 + ( W W + W W ) [ F e 2 + ] (4-58) Results and Discussion 118 Substitution of Eqn. (4.58) into Eqn. (4.38) gives the following expression for the anodic current: / =4Fk A 1 ( ^ c ,Cp^ c ,C P + ^ c , P y A , P y ) [ F e +3 a,cP ° ^ l 4 A a C p h : a C p + ( 2 + 12a)/\a,Py/ca,Py[H+]-05 +(/cc,CpA,cP +^,Py4.Py)Fe 2 +] (4.59) The dissolution rate of chalcopyrite, -rCp, may be expressed by: dnCp /a,Cp c p dt 4F _ h, A I (^c .Cp^c .Cp+^cPyA.Py)^ ] (4.60) a,CP ° ^ \ 4 A a C p k a C p + ( 2 + 120-)\P y/C a,P y[H+]-°- 5 +(/Cc,c p4,Cp +^,PyA,Py)Fe 2 +] Substitution of Eqns (4.60) and (4.26) through (4.29) in Eqn. (4.46) gives the following expression for the leaching rate of chalcopyrite: * k - J L e k A - c'cp'Vcp^cp U L "Cp,0 [(1-0Cp)>4Cp/ceCp +(1-9P y)AP y/c e P y][Fe 3 +] Cp^a.Cp + (2 +12a)0P yA3y/ca,Py [H +]-° 5 + [(1 - 6Cp )ACpkc,Cp + (1 - 0 P y )APykcPy][Fe2+] (4.61) Substitution df Eqn. (4.35) in Eqn. (4.61) gives the following expression: dXcP _ J L 0 n A ^Cp^a.Cp^Cp d t ncP,o [(1 - Sep )*c,cp + (1-6Py)/3/cc,Py][Fe3+] (4.62) 4GCpkaCp +(2 + ^2o)9PvBkaPv[H+]-05^ Substitution of (4.49) in Eqn. (4.62) gives the following equation: Results and Discussion 119 dXCp _ ATTR2 r 2 / 3 —IT - — y C p K a , C p ( 1 - * C p ) U L "Cp , 0 [(1-0Cp)^,Cp+(1-0py)^c,Py][Fe3+] (4.63) l49CpkaCp + ( 2 + 12a)0 P y J 8/c a , P y [H + r 0 5 + [(1 - 0 C p )^ .cP +(1-0p y ) /3/c c , P y ] [Fe 2 + ] It is known that the rest potential of chalcopyrite is lower than that of pyrite. When contact is made, the folowing cases may then occur: > Em< Ecp < Epy (chalcopyrite and pyrite are inert) > Ecp < Em < Epy (chalcopyrite is oxidized and pyrite is inert) > Ecp < Epy < Em (chalcopyrite and pyrite are oxidized) where Ecp is the rest potential of chalcopyrite and EPy is the rest potential of pyrite. As shown previously, the beneficial effect of the galvanic interaction was only observed within the potential range where chalcopyrite was oxidized and pyrite was inert (Ecp < Em < Epy). Based on this observation, it was of practical interest to investigate only this second case. We can assume for this case that the total pyrite surface is cathodic (0Py = 0), in which case Eqn. (4.63) is simplified as folows: ,3+1 _ \2 /3 dt "cP,o 4 0 c p / c a , c p + [(1 -6 C p )k , C p + /3/cc,Py][Fe2+] dXCp 4TTR2q r I [(1-QcP)^ ,cp+^^ ,Py][Fe3+] 2 —77- = y C D K a . C D J r 77.—77~r T r 7771 oITV1 _ Ac P/ (4.64) dXCp 4nR2 a or —— = 0, dt nCp,o ,• I [(1-eCp) + ^ c][Fe3-] C p C l0c p K C p + [(1-a Cp) + ^ c][Fe2+)(1 X c p ) ( 4' 6 5 ) k k where 5>. =-^H. and <pc = K c , C p ftc,Cp Integration of Eqn. (4.65) gives the folowing expression for the conversion of chalcopyrite as a function of time during galvanic interaction with pyrite: Results and Discussion 120 ( 1 - * C p ) 1 / 3 = 1 -477-R2 3«Cp,0 [ (1-0 C p ) + i 8^] [Fe 3 + ] (4.66) Fcp c^P+[(1-0cp) + ^ c ] [ F e 2 + ] 4.2.6. Electrochemical reaction and surface passivation model A s mentioned previously, the slow kinetics for the chalcopyrite leaching is attributed to the formation of passive layers on the chalcopyrite surface. It is then necessary to account for this in the electrochemical model. This model incorporates the following assumptions: > The passivating films grow at a rate proportional to the amount of surface area that is not covered by the film [119]. > T h e passivating films are formed on the anodic area of chalcopyrite only. > Ferric reduction does not take place on the passivating films. In the absence of pyrite If A represents the fraction of the anodic area covered by passivating films, then Eqn . (4.26) may be rewritten as follows: Based on the the first assumption above, the growth rate of the passive layer can be expressed by: 4,CP=( 1-^CA P (4.67) (4.68) or / = 1 - e x p ( - y ) (4.69) and AaCf> = exp(-/cpi)ec py4C p (4.70) from which Eqn (4.51) may be rewritten as follows: Results and Discussion 121 c /X C p _ AnR^g k , k , (1-eCp)[Fe"] _ 2 , 3 ( 4 7 1 . Integration of Eqn . (4.71) gives the following expression for the conversion of chalcopyrite as a function of time when passivation is taken into account: 4 r r R 2 2kc d " X C p ) 1 ' 3 = 1 - T K T K V(1-0cP)[Fe3+] J"cP.o K p A c P (4.72) x (V^cp^cp + (1 - 0 C P )Fe2+] - VeCpKcp exp(-V) + (1 - QCp ) [Fe 2 + ]) /n fA7e presence of pyrite at potentials lower than the rest potential of pyrite In the presence of pyrite, assuming that pyrite does not dissolve and that passivation takes place only on the anodic area of chalcopyrite, then E q n . (4.70) may be applied in a similar fashion to E q n . (4.65) with the following result: * " "cp, w C ° P < "'"\|exp(-fcp()eI:pKCp + [ ( 1 - e c p ) + W J [ F e - ] ( 1 - X c ' ) ( 4 J 3 ) Integration of Eqn . (4.73) gives the following expression for the conversion of chalcopyrite as a function of time during galvanic interaction with pyrite, when passivation is taken into account: ( 1 _ ^ r = u f R 2 2 k * V[(1-eCp) + ^ e ] [ F e 3 - ] ^"Cp.O K p ^ C p *(VecpKcP + [ ( i ^ (4-74) 4.2.7. Validation of the electrochemical passivation model T h e model was validated at 6 5 ° C because the extraction curves did not show an induction period in the early stage of leaching. T h e initial leaching rates were determined from the curves by considering that passivation did not occur during the first Results and Discussion 122 3 hours for Fe(lll):Fe(ll) ratios of 0.1, 0.33 and 1. This was consistent with the plot representing the copper conversion (1 - (1 -Xc P ) 1 / 3 ) as a function of time. Electrochemical and passivation parameters used in the model are presented in Tables 4.9 to 4.11. The model fits well the copper extraction data in the absence of pyrite and during galvanic interaction with pyrite as shown in Figures 4.29 and 4.30. The model indicates that little passivation (low value of k p) occurred at a Fe(lll):Fe(ll) ratio of 0.33. Severe passivation was observed at high Fe(lll):Fe(ll) ratios. T h e model can be used to predict the copper extraction when pyrite is added as catalyst for the ferric reduction. Typical copper extraction (predicted by the model) at various pyrite to chalcopyrite ratios are shown in Figure 4.31. A s can be seen, more than 90% copper extraction can be achieved in 13 hours for FeS2:CuFeS2 mass ratio of 4. Complete conversion is obtained in 15 hours for FeS2:CuFeS2 mass ratio of 8. • Fe(lll)/Fe(ll) = 0.1 • Fe(lll)/Fe(ll) = 0.33 Fe(lll)/Fe(ll) = 1 • Fe(lll)/Fe(ll) = 5 o Fe(lll)/Fe(ll) = 15 model-Fe(lll)/Fe(ll) =0.1 -model-Fe(lll)/Fe(ll) = 0.33 -model-Fe(lll)/Fe(ll) = 1 -model-Fe(lll)/Fe(ll) = 5 model-Fe(lll)/Fe(ll) = 15 10 15 Time (H) 25 Figure 4.29 Electrochemical-passivation model for the leaching of chalcopyrite in ferric sulphate solutions, pH 1.4, 6 5 ° C Results and Discussion 123 T h e model can also predict the effect of particle size during the galvanic interaction. Decreasing the particle size of pyrite has a much more significant effect than that of the chalcopyrite. A s shown in Figure 4.32, complete conversion is obtained in 15 hours while using finer pyrite particles (25 um) and relatively coarser chalcopyrite particle (100 um). T h e combine effect of fine pyrite and amount of pyrite added is shown in Figure 4.33. Complete conversion is obtained in 7 hours for FeS2:CuFeS2 mass ratio of 8. It is therefore essential to have sufficient catalytic cathodic sites to support the ferric reduction in order to counteract chalcopyrite 'passivation" and achieve complete copper extraction. 0.9 -i - - -0.8 -| . „ - A " A A "S 0 7 -O Chalco=30, Pyr=0 • Chalco=20, Pyr=10 A Chalco=10, Pyr=20 Model-Chalco=30, Pyr=0 Model-Chalco=20, Pyr=10 - - - Model-Chalco=10, Pyr=20 Time (H) Figure 4.30 Electrochemical-passivation model for the leaching of chalcopyrite at different F e S 2 : C u F e S 2 ratios, Fe(lll):Fe(ll) = 1, pH 1.4, 6 5 ° C Results and Discussion 124 H 3 - P y / C p = 0 Py/Cp = 0.5 -A -Py /Cp = 1 Py/Cp = 2 Py/Cp = 4 - A - Py/Cp = 6 - • - Py/Cp = 8 0 * 1 ! i 1 1 0 5 10 15 20 25 Time (H) Figure 4.31 Predicted values for the effect of increased FeS2:CuFeS2 ratio on the leaching rate of chalcopyrite, Fe(lll):Fe(ll) = 1, pH 1.4, 6 5 ° C Results and Discussion 125 Time (H) Figure 4.32 Predicted values for the effect of particle size on the leaching of chalcopyrite, F e S 2 : C u F e S 2 = 2, Fe(lll):Fe(ll) = 1, pH 1.4, 6 5 ° C Results and Discussion 126 Figure 4.33 Predicted values for the effect of increased F e S 2 : C u F e S 2 ratio and particle size on the leaching of chalcopyrite, Fe(lll):Fe(ll) = 1, pH 1.4, 6 5 ° C , dPy = 25 pm, d c p = 100 pm Results and Discussion 127 Table 4.9. Input parameters for the electrochemical-passivation model Fe(lll):Fe(ll) ratio 0.1 0 . 3 3 1 5 15 1 1 Physical Parameters mCP ( g ) 3 0 3 0 3 0 3 0 3 0 2 0 1 0 mPy (g) 0 0 0 0 0 1 0 2 0 PcP ( 9 cm -3) 4 . 2 4 . 2 4 . 2 4 . 2 4 . 2 4 . 2 4 . 2 Ppy (g cm -3) 5 . 0 5 . 0 5 . 0 5 . 0 5 . 0 5 . 0 5 . 0 dcP (um) 4 7 . 6 4 7 . 6 4 7 . 6 4 7 . 6 4 7 . 6 4 7 . 6 4 7 . 6 dpy(um) 4 5 . 5 4 5 . 5 4 5 . 5 4 5 . 5 4 5 . 5 4 5 . 5 4 5 . 5 )8 1 . 8 3 . 5 7 . 1 0 0 0 . 4 1 . 8 Electrochemical Parameters [Fe(lll)] (M) 0 . 0 0 8 0 . 0 2 2 g 0 . 0 4 5 0 . 0 7 5 0 . 0 8 4 0 . 0 4 5 0 . 0 4 5 [Fe(ll)] (M) 0 . 0 8 1 0 . 0 6 7 0 . 0 4 5 0 . 0 1 5 0 . 0 0 6 0 . 0 4 5 0 . 0 4 5 [ H + ] ( M ) 0 . 1 8 0 0 . 1 8 0 0 . 1 8 0 0 . 1 8 0 0 . 1 8 0 0 . 1 8 0 0 . 1 8 0 . kcp 1 . 1 2 * 1 0 " 5 1 . 1 2 x 1 0 " 5 1 . 1 2 x 1 0 " 5 1 . 1 2 x 1 0 " 5 1 . 1 2 * 1 0 " 5 1 . 1 2 * 1 0 ~ 5 1 . 1 2 x 1 0 " 5 Kcp 0 . 0 0 1 0 . 0 0 1 0 . 0 0 1 0 . 0 0 1 0 . 0 0 1 0 . 0 0 1 0 . 0 0 1 <Pc 1 . 4 5 2 1 . 4 5 2 1 . 4 5 2 1 . 4 5 2 1 . 4 5 2 1 . 4 5 2 1 . 4 5 2 <Pc 1 . 0 0 0 1 . 0 0 0 1 . 0 0 0 1 . 0 0 0 1 . 0 0 0 1 . 0 0 0 1 . 0 0 0 Passivation Parameters 0 . 3 2 8 0 . 1 6 7 0 . 2 4 8 0 . 5 7 7 0 . 6 0 5 0 . 2 4 8 0 . 2 4 8 Results and Discussion 128 Table 4.10. Input parameters for the simulation of the electrochemical model, effect of pyrite addition Fe(lll):Fe(ll) ratio 1 1 1 1 1 1 1 Physical Parameters mCp (g) 30 20 15 10 6 4.3 3.3 mpy (9) 0 10 15 20 24 25.7 26.7 Pep (gem - 3) 4.2 4.2 4.2 4.2 4.2 4.2 4.2 p P y (g cm"3) 5.0 5.0 5.0 5.0 5.0 5.0 5.0 dcp (um) 47.6 47.6 47.6 47.6 47.6 47.6 47.6 dpy (Mm) 45.5 45.5 45.5 45.5 45.5 45.5 45.5 P 0.0 0.4 0.9 1.8 3.5 5.3 7.1 Electrochemical Parameters [Fe(lll)] (M) 0.045 0.045 0.045 0.045 0.045 0.045 0.045 [Fe(ll)] (M) 0.045 0.045 0.045 0.045 0.045 0.045 0.045 [H+] (M) 0.180 0.180 0.180 0.180 0.180 0.180 0.180 kcP 1.12*10"5 1.12*10"5 1.12x10"5 1.12x10 - 5 1.12*10~5 1.12x10 - 5 1.12x10"5 Kcp 0.001 0.001 0.001 0.001 0.001 0.001 0.001 <Pc 1.452 1.452 1.452 1.452 1.452 1.452 1.452 <Pc 1.000 1.000 1.000 1.000 1.000 1.000 1.000 Passivation Parameters kp 0.248 0.248 0.248 0.248 0.248 0.248 0.248 Results and Discussion 129 Table 4.11. Input parameters for the simulation of the electrochemical model, effect of particle size Fe(lll):Fe(ll) ratio 1 1 1 Physical Parameters mcp (g) 10 10 10 m P y (g) 20 20 20 PcP (9 cm"3) 4.2 4.2 4.2 Ppy (g cm - 3) 5.0 5.0 5.0 dcP (Mm) 47.6 25 100 dpy (Mm) 45.5 100 24 P 1.8 3.5 7.1 Electrochemical Parameters [Fe(lll)] (M) 0.045 0.045 0.045 [Fe(ll)] (M) 0.045 0.045 0.045 [H +] (M) 0.180 0.180 0.180 kCp 1.12x10" 5 1.12*10" 5 1.12x10" 5 Kcp 0.001 0.001 0.001 Vc 1.452 1.452 1.452 <Pc 1.000 1.000 1.000 Passivation Parameters kp 0.328 0.167 0.248 Results and Discussion 130 4.3. Bioleaching of chalcopyrite Chemical leaching tests conducted in the previous section have shown that chalcopyrite leaching is strongly dependent on the concentrations of ferrous and ferric (solution potential) and temperature. Hydrogen peroxide was used to regenerate the ferric during the leaching process. However, the use of chemical reagents for ferrous oxidation is impractical for industrial application. The addition of microorganisms to regenerate ferric and overcome chalcopyrite passivation has attracted a lot of attention in recent years. The response of Santa Eulalia chalcopyrite ore during bioleaching was tested with three types of microorganisms, namely mesophiles, moderate thermophiles and extreme thermophiles. 4.3.1. Mesophiles Bioleaching experiments with mesophiles were carried out at 2 8 ° C . After 30 days of leaching, the copper extraction reached 35% (Figure 4.34). The leaching curve showed 3 distinct phases; a lag phase, an active phase and a "death" phase. From day 0 to day 7, the copper extraction was relatively low (Figure 4.34). Copper was released in this phase at a rate of 0.22 g L ~ 1 h" 1. A s shown in Figure 4.35, iron seemed to leach more rapidly than copper. A C u : F e molar ratio of about 1:3 was obtained in the earlier stage of leaching. Similar results were obtained during the electrochemical leaching of chalcopyrite (section 4.1.1.2). This result confirmed the conclusion suggesting the preferential dissolution of Fe at low temperature and the formation of an iron-deficient layer on the chalcopyrite surface. The solution potential did not increase and the number of bacteria was about the same during this phase (Figure 4.36). The concentration of the cells in the liquid phase was considered as indicative of microbial growth. Results and Discussion 131 10 15 20 Time (Days) Figure 4.34 Fraction of metal leached with mesophiles at 2 8 ° C 0.7 3 o O « o S 10 15 20 Time (Days) Figure 4.35 Molar ratio C u : F e during bioleaching with mesophiles at 2 8 ° C . Results and Discussion 132 700 350 300 -solution potential, mesophiles -Bacteria counts, mesophiles 10 15 20 Time (days) 25 30 2.5 2 72 E « o X 1.5 c 3 1 O 1 o TO O 0.5 m 35 Figure 4.36 Solution potential evolution and microbial counts during bioleaching with mesophiles at 2 8 ° C . Many authors have observed similar behavior and believe that microbes need time to grow and adapt to the substrate. The leaching rate was relatively fast from day 7 to day 15. Copper was released during this phase at a rate of 0.32 g L - 1 h~ 1. The solution potential increased from 450 to 625 m V (Ag/AgCI). The increase in the redox potential is due to the bacterial oxidation of ferrous to ferric. The molar ratio C u : F e also increased from 1:3 to 2:3, which was an indication that a different leaching mechanism was occurring on the chalcopyrite surface. The leaching rate of chalcopyrite gradually declined from day 15 to day 30. Copper was extracted at a rate of about 0.08 g L" 1 h~ 1. T h e solution potential and the number of bacteria in the inoculated flasks leveled off. This was an indication that the mineral was either passivated or the microorganisms were not active on the chalcopyrite surface. The formation of passive films is the most likely explanation for the decrease in the leaching rate at low temperature and high solution potentials. It appears that mesophiles Results and Discussion 133 turned to the use of available ferrous in solution when the chalcopyrite became passivated, and the energy required for chalcopyrite oxidation was no longer enough to sustain a large number of bacteria. Analysis of the residue indicated that the sulphur yield was about 20%, i.e., the oxidation of sulphur yielded almost entirely sulphate. This result renders the role of an elemental sulphur layer in the passivation of chalcopyrite unlikely. 4.3.2. Moderate thermophiles Bioleaching tests with moderate thermophiles were performed at 4 5 ° C . A final extraction of 70% was achieved in 30 days (Figure 4.37). Copper was released at a rate of 0 .25-0.45 g L - 1 h - 1 . The shake-flask tests confirmed that moderate thermophiles were more active than the mesophiles for the leaching of chalcopyrite. The leaching curve was almost linear, with no sign of an induction period in the earlier stage of leaching and no sign of passivation after extended leaching times. The C u : F e molar ratio was about 1:2 in the earlier stages of leaching and rose slowly until it reached a level of 8:9 after 30 days (Figure 4.38). It is unclear at this stage what causes this low C u : F e initially, as chalcopyrite leach kinetics did not seem to be affected. Solution potentials increased from 410 to 520 m V (Ag/AgCI) from the beginning of the test to day 15 and remained at 520 m V until the end of the test (Figure 4.39). It is interesting to note that chemical leaching tests at 4 5 ° C indicated that chalcopyrite leaching was much faster at a redox potential of about 540 m V (Fe(lll):Fe(ll) ratio of 5). T h e relatively high dissolution rate in the presence of thermophiles is likely due to their poor iron-oxidizing capabilities. Analysis of the solid residue indicated a sulphur yield of 48%. The presence of elemental sulphur did not seem to stop the leaching of chalcopyrite. There was good agreement between the chemical and bioleaching tests. Results and Discussion 134 0.8 Time (Days) Figure 4.37 Fraction of metal leached with moderate thermophiles at 4 5 ° C . o S 0.3 -0.2 -0.1 -0 J 1 . 1 1 I 1 1 0 5 10 15 20 25 30 35 Time (days) Figure 4.38 Molar ratio C u : F e during bioleaching with moderate thermophiles at 4 5 ° C . Results and Discussion 135 Figure 4.39 Solution potential evolution and microbial counts during bioleaching with moderate thermophiles at 4 5 ° C . 4.3.3. Extreme thermophiles Bioleaching tests with extreme thermophilic microorganisms were performed at 6 8 ° C . Chalcopyrite leached faster with extreme thermophiles with a copper extraction of 91% after 21 days compared with only 32% in the presence of mesophiles (Figure 4.40). The curves showed three stages of kinetics. The leaching rate was relatively slow during the first 5 days and was followed by a rapid dissolution of the mineral from day 5 to day 21. The leaching rate then declined after 21 days of leaching. T h e C u : F e molar ratio was about 0.9-1.2 during the first 21 days of leaching, indicating that almost equal amounts of copper and iron were released into solution (Figure 4.41). Similar results were obtained during the electrochemical leaching of chalcopyrite at 6 5 ° C . Solutions potentials were very low (350 m V (Ag/AgCI)) during the first leaching stage and increased up to 450 m V in the second and third stage (Figure 4.42). Bacterial counts showed similar trends. The number of microorganisms slightly increased during the first 5 days of leaching and rapidly increased over the following 15 days, but decreased very Results and Discussion 136 gradually after 21 days of leaching. The slow kinetics observed initially is probably due to the absence of enough ferric in solution for the oxidation of chalcopyrite. T o further confirm our conclusions, two more tests were conducted with relatively high concentrations of ferric at the beginning of the tests. A s shown in Figure 4.43, the "induction" period disappeared when more ferric were present in solution. The same slope and the linear leaching rate observed after 5 days of leaching in all three cases at 6 5 ° C indicate good repeatability of the tests. This suggests that a certain concentration of ferric in solution is necessary to initiate chalcopyrite oxidation. Analysis of the residue after leaching revealed a sulphur yield of 70%. The plateau observed at about 92% conversion is probably related to the formation of a thick layer of elemental sulphur, restricting the supply of oxygen and ferric to the chalcopyrite surface. 0 5 10 15 20 25 30 35 Time (Days) Figure 4.40 Fraction of metal leached with extreme thermophiles at 6 8 ° C . Results and Discussion 137 u. "3 o (0 CC JS o 1.4 1.2 1 0.8 0.6 0.4 0.2 0 Extreme thermophiles, 68°C 10 15 20 Time (Days) 25 30 35 Figure 4.41 Molar ratio C u : F e during bioleaching with extreme thermophiles at 6 8 ° C . 600 ° 350 Solution potential, extreme thermophiles Bacter ia counts, extreme thermophiles 10 15 20 Time (days) 30 + 1.5 2.5 E a> o w ra u 0.5 S 35 Figure 4.42 Solution potential evolution and microbial counts during bioleaching with extreme thermophiles at 6 8 ° C . Results and Discussion 138 Time (h) Figure 4.43 Fraction of copper leached with extreme thermophiles at 6 8 ° C with various initial concentrations of Fe(lll) and Fe(ll), total Fe = 1g L ~ 1 . 4.3.4. Comparison of bioleaching with mesophiles and thermophiles Bioleaching tests have shown that thermophile cultures are able to attain complete dissolution of chalcopyrite (Figure 4.44). T h e use of mesophiles for chalcopyrite leaching was not successful due to the tendency of this mineral to 'passivate" at low temperature and high solution potentials. The metabolism of thermophiles generates low solution potentials (360 to 480 mV (Ag/AgCI)) favorable to chalcopyrite leaching, while mesophiles generate high solutions potentials (Figure 4.45). T h e success of thermophiles over mesophiles has not yet been clearly elucidated. It is possible that mesophiles act on the chalcopyrite surface through a indirect mechanism. The main characteristic of this mechanism is the hypothesis that ferric and/or protons are the only chemical agents dissolving the mineral. The bacteria's sole function is to regenerate ferric and to concentrate them at the mineral-solution interface. The chemical attack of the mineral takes place in a tiny exopolymer layer, the glycocalyx, with a thickness in the nanometer range, surrounding the cells [120]. Two Results and Discussion 139 reactions are in competition during the indirect attack of sulphide minerals: ferrous production from the chemical ferric leach reaction: M S + 2 F e 3 + - » M e 2 + + 2 F e 2 + + S ° (4.98) and ferrous consumption to generate ferric and in some cases the oxidation of elemental sulphur: Due to passivation of the mineral, the rate of ferrous production from the leaching step can be much lower than that of the ferrous consumption by the bacteria. In addition, the ferric reduction is retarded when the surface is polarized at low temperatures and high solution potentials. These combined effects are responsible for the low kinetics observed during the bacterial oxidation of chalcopyrite with mesophiles. In contrast, it is possible that thermophilic leaching of chalcopyrite takes place through a direct mechanism. According to the direct mechanism, the microbes maintain close contact across the exopolymer layer and obtain energy from direct oxidation of the reduced sulfur substrate. This mechanism is similar to localized corrosion (pitting corrosion) giving rise to high dissolution rates. The direct mechanism is predominant when the microbes enhance the rate of oxidation above what can be achieved by chemical reaction with ferric sulphate under the same solution conditions. It is believed that the enhancement in the rate of dissolution is caused by an enzymatic oxidant secreted by the microbes. The attachment of microbes is a necessary condition for the direct mechanism, but not a sufficient condition [121]. 4 F e 2 + + 4 H + + 0 2 + bacteria 4 F e 3 + + H 2 0 (4.99) 2 S ° + 3 0 2 + 2 H 2 0 -> 2 H 2 S 0 4 (4.100) Results and Discussion 140 1 Time (Days) Figure 4 .44 Fraction of copper bioleached in the presence of various microorganisms. The sufficient condition is that an oxidizing agent (not ferric or oxygen) secreted by the microbes is involved in mineral leaching. Thus , the main difference between the two mechanisms is the role played by ferric in the dissolution of the mineral. In the "indirect mechanism" the mineral is leached only by ferric and/or oxygen, whereas in the "direct mechanism", the mineral is not leached by ferric, but by an oxidant secreted by the microbes [ 1 2 1 ] . This conclusion was supported by the fact that iron is mainly present as ferrous during thermophilic leaching of chalcopyrite. According to Marchbank et al. [54], thermophiles are less mobile than mesophiles. They do not float freely and oxidize dissolved iron as do the mesophiles. Results and Discussion 141 700 T 300 -\ 1 1 1 1 1 1 1 0 5 10 15 20 25 30 35 Time (Days) Figure 4.45 Solution potentials during the bioleaching of chalcopyrite with various microorganisms. Several authors have tried to identify the biological leaching agent secreted by the microbes. Rojas-Chapana and Tributsch [122] identified cysteine as the chemical carrier in A. ferrooxidans. Addition of bacteria to the system cysteine-pyrite increased the bioleaching of pyrite by a factor of about 2.5. However these authors did not show any data in the absence of bacteria under similar conditions of leaching. It was therefore difficult to prove a link between cysteine and the chemical leaching of the mineral. T h e work presented by Rojas-Chapana and Tributsch [122] did not provide satisfactory answers about the exact role played by the biological leaching agent. In addition, if the leaching of chalcopyrite was only due to a biological oxidant, an increase in the ferric concentration should not have affected the dissolution rate. However, it was observed that increasing the ferric concentration had a favorable effect in the earlier stage of the thermophilic leaching of chalcopyrite. It seems that both the anodic and cathodic reactions on chalcopyrite were fast in the presence of thermophiles. It is therefore Results and Discussion 142 possible that thermophiles play the same role as pyrite; i.e., bacteria attach to chalcopyrite and their surfaces support the entire cathodic reaction on behalf of chalcopyrite. This means that the oxidation of ferrous is only due to the presence of dissolved oxygen in solution. This can explain why the leaching rates are high and the solution potentials are low in the presence of thermophiles. T h e mechanism that explains our results is shown in Figure 4.46. C u 2 + Fe Chalcopyrite Chalcopyrite Anodic area Cathodic area j 2 + + F e 2 + + 2 S ° + 4 e~ F e 3 + + e" = F e 2 + Figure 4.46 Schematic diagram of the bioleaching mechanism of chalcopyrite in the presence of thermophiles. This section has shown that thermophiles were more effective than mesophiles. Mesophiles oxidize chalcopyrite by an indirect mechanism similar to chemical leaching with ferric. T h e slow leach rates observed in this case are due to the formation of passive layers at high potentials and low temperature and to the slow kinetics of ferric reduction on chalcopyrite. At higher temperatures, ferric reduction appears to take place on the surface of attached thermophiles. Complete conversion was obtained with thermophiles because of the direct galvanic interaction of the microorganisms with Results and Discussion 143 chalcopyrite. W e suggest that thermophiles accelerate ferric reduction on chalcopyrite by providing alternative cathodic sites. T o further confirm our conclusions and improve the galvanic couples, it was decided to increase the number of cathodic sites by adding pyrite to chalcopyrite during bioleaching with thermophilic cultures. This is discussed in the next section. 4.4. Column leaching The primary objective of this section was to demonstrate that the addition of pyrite could be beneficial to heap bioleaching of chalcopyrite. T h e study was conducted in 4 small columns, all placed in a common heated water bath at 6 8 ° C . In order to improve the kinetics of the process, fine chalcopyrite and pyrite particles were coated onto a low-grade chalcopyrite ore. Details of the experimental set-up are given in section 3.4. The curves show two stages of kinetics (Figure 4.47). T h e leaching rate was relatively slow during the first 10 days and was followed by rapid dissolution from day 10 to day 40. A similar trend was obtained in shake flask experiments. T h e induction period was possibly due to the lack of a significant amount of ferric in solution for the oxidation of chalcopyrite. The addition of pyrite (FeS2:CuFeS2 ratio of 2:1) did not affect this induction period. Solution potentials were below 400 m V in both cases. After 8 days, leaching proceeded to an extraction of 56% in the absence of pyrite and an extraction of 74% in the presence of pyrite by day 40. This increase in the leaching rate by a factor of 1.3 confirmed our previous conclusions suggesting that the presence of pyrite was accelerating the leaching of chalcopyrite. The dissolution rate of copper was increased by a factor of two during controlled potential experiments. Solution potential increased rapidly to levels around 530 m V in both columns (Figure 4.48). This is due to the biooxidation of ferrous. There was not a marked difference between the solution potentials inside the two columns. At these high potentials, it is reasonable to presume that pyrite oxidation had started. A s noted previously, the beneficial effect of pyrite decreases once it starts to dissolve. Results and Discussion 144 In order to achieve high copper recoveries, it is therefore essential to work at high temperature, to maintain low solution potentials, a sufficient amount of ferric in solution, and pyrite in the mineral. However, as mentioned in the literature survey, heap leaching is more suitable for the treatment of low-grade ores. T h e oxidation of ferrous by oxygen might be the limiting step at high temperature in a heap containing high grade chalcopyrite ores. Several studies have been performed in a bioleaching agitated-tank reactor, but they have been largely restricted to the gold sector [54]. Conventional bio-reactors used in mesophilic leaching have counted on high-power input to improve the oxygen mass transfer necessary for rapid bacterial growth and regeneration of ferric. However, as seen previously, chalcopyrite does not respond well to mesophiles. Thermophiles are known to have a much more fragile cell structure that can be damaged by the high shear conditions generated in these reactors [54]. They can only be used with low pulp densities. A n alternative method for the treatment of high grade chalcopyrite ores could be the use of oxygen or air under atmospheric pressure for ferric regeneration in an agitated tank reactor. This is discussed in the next section. 4.5. Atmospheric leaching The previous sections have shown the importance of redox (solution) potential and the addition of pyrite in chalcopyrite dissolution. The methods described in preceding sections used electrical, chemical and biological means of controlling the electrode and solution potentials within a set range of values. For economic reasons, these methods cannot be applied at a larger scale, particularly while treating high grade chalcopyrite minerals. O n the other hand, thermophiles can only be used at low pulp densities in agitated tank reactors. It was therefore of interest to investigate the use of air or oxygen as the oxidant to control the slurry redox potential at elevated temperature. Atmospheric leaching tests were conducted at 8 0 ° C with various amount of pyrite added into the reactor. Results and Discussion 145 0.8 Time (h) Figure 4.47 Copper extraction in thermophilic leaching tests conducted in small columns with and without addition of pyrite, 6 8 ° C 4.5.1. Effect of pyrite addition Leaching tests were conducted in acid sulphate solutions containing Fe(lll) and Fe(ll) at a ratio of 1:1, which corresponded to a solution potential of about 540 m V (Ag/AgCI). Air was supplied as oxidizing agent. T h e leaching curves showed slow kinetics in the absence of pyrite (Figure 4.49). A copper extraction of 22% was recorded after 24 hours of leaching. T h e addition of pyrite increased the7 dissolution of chalcopyrite. The leaching rate increased by a factor of 1.5 when 20 g of pyrite were mixed with 10 g of chalcopyrite. For a ratio FeS2:CuFeS2 of 4:1, the dissolution rate of copper was increased by a factor of 3 during the first 6 hours of leaching. T h e copper extraction rapidly dropped after 7 hours of leaching, possibly due to chalcopyrite passivation. This can be related to the change in solution potential with time. Results and Discussion 146 550 300 H 1 1 - i 1 1 1 . i 0 5 10 15 20 25 30 35 40 Time (H) Figure 4.48 Solution potentials in thermophilic leaching tests conducted in small columns with and without addition of pyrite, 6 8 ° C The potential development showed some interesting trends (Figure 4.50). In the absence of pyrite, the solution potential decreased in the first 6 hours, probably due to the consumption of ferric by the mineral. After 7 hours of leaching, the potential increased rapidly to high levels around 520 mV and decreased slightly towards the end of the test. This could possibly be related to the reduced ferric reduction once the passive films are formed. In the presence of pyrite, for a ratio F e S 2 : C u F e S 2 of 4:1, the potential dropped rapidly to 415 mV for the first 2 hours and increased to 470 m V after 7 hours of leaching, and reached 520 mV towards the end of the experiment. The sudden decrease in the leaching rate after 7 hours of leaching possibly corresponds to the beginning of pyrite oxidation. A s mentioned in the previous sections, the beneficial effect of pyrite stops once it starts to dissolve. For a ratio F e S 2 : C u F e S 2 of 2:1, solution potentials were above 460 mV. This critical potential corresponds to the simultaneous dissolution of chalcopyrite and pyrite. This can explain the slow kinetics observed for this case. In order to confirm these suggestions, it was decided to conduct similar Results and Discussion 147 experiments, starting with low solution potentials (low Fe(lll):Fe(ll) ratios) and possibly maintaining low potentials in the leaching system by adjusting the air flowrate. Figure 4.49 Leaching rate curves of chalcopyrite at different ratios of FeS2:CuFeS2, Initial Fe(lll):Fe(ll) =1, 8 0 ° C 4.5.2. Effect of initial solution potential T h e effect of solution potential was investigated by conducting leaching tests at potentials much lower than 460 m V in order to avoid chalcopyrite passivation and pyrite dissolution. A solution potential of about 410 mV was selected for the tests because maximum copper extraction was observed around this value in the presence of pyrite (Figure 4.49). In the absence of pyrite, decreasing initial solution potential had a favorable effect on the leaching rate of chalcopyrite. The copper extraction increased by a factor of 2.5 (Figure 4.51). T h e leaching rate was initially rapid, and then started to decrease after 3 hours of leaching. C u leaching ceased almost entirely after 16 hours of leaching. Results and Discussion 148 560 540 O I5 520 "5> < 500 (A > > 480 | 460 f 440 (A LU 420 400 - o - Pure chalco -»-Chalco=10-Pyr=20 - o - Chalco=10-Pyr=40 ~ T \ -* 0 10 15 Time (h) 20 25 Figure 4.50 Solution potentials during Leaching of chalcopyrite at different ratios of F e S 2 : C u F e S 2 , Initial Fe(lll):Fe(ll) = 1, 8 0 ° C Solution potentials increased rapidly with time in the absence of pyrite as shown in Figure 4.52. This indicates that a small amount of ferrous was released into solution, possibly due to chalcopyrite passivation. With the introduction of pyrite, complete copper extraction was achieved in 21 hours when the solution potential was initially low. Comparison between the curves at high and low initial redox potentials shows that the leaching rates were almost the same for the first 6 hours and there was a marked difference between the two curves only after 7 hours of leaching. The chalcopyrite oxidation clearly continued to completion in the case where the solution potential was initially low. T h e solution potential remained below the "critical" potential of 460 mV. This means that chalcopyrite was releasing a significant amount of ferrous in solution, and consequently maintaining the solution potential at relatively low values. Based on these results, it was of interest to investigate the effect of other parameter which may enhance the leaching rate of chalcopyrite and the galvanic contact between chalcopyrite and pyrite. T h e s e parameters include particle size, pulp density, impeller speed, the type of primary oxidant (air or oxygen) and acidity of the solution. Results and Discussion 149 > c o u k_ 01 a. a. o o 100 90 80 70 60 50 40 30 20 10 0 — -Cl-Chalco=10, High Redox -•-Chalco-10, Low Redox -O-Chalco=10, Pyr=40, High Redox -•-Chalco=10, Pyr=40, Low Redox m — • -10 15 Time (H) 20 25 Figure 4.51 Effect of initial potential on the leaching rate curves of chalcopyrite at different ratios of F e S 2 : C u F e S 2 , 8 0 ° C 4.5.3. Effect of particle size T h e effect of particle size of both chalcopyrite and pyrite on the leaching rate of chalcopyrite was studied separately and together. It was observed that reduction of particle size below - 7 5 pm did not improve the rate of leaching or the final extraction of chalcopyrite in the absence of pyrite (Figure 4.53). Jones and Peters observed similar results when chalcopyrite was dissolved in ferric sulphate medium at 9 0 ° C [38]. The rate increased substantially with finer particles only in ferric chloride medium. A s shown in Figure 4.54, solution potentials were about the same, which indicated that the copper extraction was mainly dependent on the solution potentials. Results and Discussion 150 600 Figure 4.52 Evolution of solution potentials during leaching of chalcopyrite at different initial potentials, 8 0 ° C 60 50 c 40 o 'I co £ 30 o u i_ 0) §: 20 o o 10 - • - <38 microns -o- 38_75 microns =Q 10 15 Time (h) 20 25 Figure 4.53 Effect of particle size on the leaching of chalcopyrite Results and Discussion 151 Figure 4.54 Effect of particle size on solution potential during leaching of chalcopyrite A s shown in Figures 4.55 and 4.56, when pyrite was added to the leaching system, various behaviors were observed. T h e first thing to notice from these tests is that the particle size of the pyrite had a much more significant effect than that of the chalcopyrite. In every case, copper extraction decreased with increasing pyrite particle size. Complete conversion was obtained with decreasing the chalcopyrite and pyrite particle size. However, the leaching curves also revealed that decreasing the chalcopyrite particle size without also decreasing the pyrite particle size has a detrimental effect on copper conversion. A copper recovery of 58% was obtained after 24 hours of leaching when coarse pyrite and fine chalcopyrite were used. In contrast, the copper recovery reached 88% when fine pyrite and coarse chalcopyrite were used. This result is completely at odds with conventional wisdom regarding the effect of chalcopyrite particle size on the leaching efficacy (where very fine, or even ultrafine, grinding is considered necessary). This mysterious behaviour confirmed that pyrite is the key element for chalcopyrite depassivation. It is essential to have a sufficient amount of pyrite surface area available to support the entire cathodic reaction on behalf of chalcopyrite. Decreasing pyrite Results and Discussion 152 particle size will result in increasing the surface area on which the ferric reduction will take place. With insufficient pyrite surface area, the chalcopyrite must support at least a portion of the cathodic process in order to provide a large enough electron sink for its anodic breakdown reaction. This also explains the detrimental effect of rising potential above 460 m V in some tests. A s long as the entire cathodic process is carried out on the pyrite surface, it doesn't really matter what the solution potential is. However, when the pyrite surface is inadequate, then there will be a certain potential above which the chalcopyrite becomes (at least partly) cathodic. O n c e this "critical" potential is reached, the chalcopyrite begins to exhibit "passive" behavior. A s mentioned in the previous section, this "critical" potential was about 460 m V (Ag/AgCI). Time (h) Figure 4.55 Effect of particle size of both chalcopyrite and pyrite on the leaching of chalcopyrite Results and Discussion 153 480 Time (h) Figure 4.56 Effect of particle size of both chalcopyrite and pyrite on solution potential during leaching of chalcopyrite 4.5.4. Effect of pyrite source Several authors have indicated that minerals from various origins may behave differently in acidic solutions. It was therefore appropriate to verify the beneficial effect of the galvanic leaching of chalcopyrite with pyrite minerals from different sources. Two pyrite minerals, namely Huanzala (Peru) and Park City (USA), were tested (Figure 4.57). A FeS2:CuFeS2 ratio of 4:1 was used for the tests. T h e tests using Huanzala pyrite are shown with open symbols, and those using Park City pyrite are shown with filled symbols. A s indicated in Figure 4.57, copper extraction in the presence of Park City pyrite reached 86% in 22 hours when the test started at relatively low potential (405 mV). Comparing the tests using Huanzala pyrite and Park city pyrite, it seems that Huanzala pyrite seems to have outperformed the Park City pyrite. Indeed, the Park City pyrite exhibited a lag period at the beginning of the test, during which the solution potential dipped to very low levels (almost down to 300 mV) as shown in Figure 4.58. However, Results and Discussion 154 after this initial period, its performance was very similar to the other tests. T h e lag period was likely due to the lack of ferric for chalcopyrite oxidation. It is unclear why the potential decreased so sharply during this period. It may be that there was a certain proportion of marcasite or pyrrhotite in the Park City sample which was subject to oxidation initially. Lorenzen [123] mentioned that marcasite and pyrrhotite dissolve easily in an oxidizing acid leach at high temperatures. T h e problem was eliminated by starting the tests using Park City pyrite at a slightly higher potential (around 470 mV instead of 405 mV). Under these conditions, the two pyrite samples performed identically. These results confirmed previous observations indicating that maintaining the solution potentials between 410 and 450 mV leads to high copper recoveries. Figure 4.57 Effect of pyrite source on the leaching of chalcopyrite, F e S 2 : C u F e S 2 = 4, 8 0 ° C Results and Discussion 155 500 j 480 -300 A 1 1 1 1 1 0 5 10 15 20 25 Time (H) Figure 4.58 Effect of pyrite source on solution potential during leaching of chalcopyrite, F e S 2 : C u F e S 2 = 4, 8 0 ° C 4.5.5. Effect of pulp density The effect of pulp density was investigated because increasing pulp density may enhance the contact between chalcopyrite and pyrite particles. Increases in pulp density were achieved by simply adding more chalcopyrite and pyrite to the reactor. A s shown in Figure 4.59, the test run with more pyrite and chalcopyrite achieved nearly 80% copper recovery in 24 hours. Th e same test, conducted with half the amount of pyrite and chalcopyrite, achieved 100 % copper recovery in 16 hours. Solution potentials remained below 460 mV for the tests (Figure 4.60). There was no sign of passivation in both cases. T h e s e results indicated that the rate of gas-liquid mixing was inadequate for tests involving large amounts of chalcopyrite. 4.5.6. Effects of impeller speed, choice of primary oxidant and acidity In order to improve the gas-liquid mixing in the reactor, it was decided to increase the impeller speed from 750 rpm to 1200 rpm. Each of the tests used 30 g of chalcopyrite (38-75 pm) and 60 g of pyrite (<38 pm). T h e test run at 750 rpm achieved only about Results and Discussion 156 62% copper recovery after 24 hours (Figure 4.61). Increasing the speed to 1200 rpm increased the copper recovery to nearly 78% after 24 hours. However, it was suspected at this point that the rate of gas-liquid mixing was also limited by the low oxygen partial pressure of air as well as stirring speed. Switching to oxygen was extremely beneficial, allowing the leach to achieve its full kinetic potential. The copper extraction increased by a factor of 2 during the first 4 hours of leaching. However, the copper recovery reached a plateau at about 75% after 12 hours of leaching. Solutions potentials remained below 460 mV, which indicated that pyrite was not yet leached (Figure 4.62). The leaching process was then limited by another factor. Doubling the level of acid in solution (from 10 to 20 g L ~ 1 acid) remedied this situation. Finally, increasing the pulp density by decreasing the amount of solution brought about a slight additional improvement, as initially expected. Figure 4.59 Effect of pulp density on the leaching rate of chalcopyrite Results and Discussion 157 Figure 4.60 Effect of pulp density on solution potential during leaching rate of chalcopyrite 100 Time (h) Figure 4.61 Effects of impeller speed, choice of primary oxidant and acidity on the leaching rate of chalcopyrite Results and Discussion 158 u D) < > > E 500 480 460 440 420 400 380 360 340 320 300 - • - C 3 0 - P 6 0 , 750 rpm, air - • - C 3 0 - P 6 0 , 1200 rpm, air - A - C 3 0 - P 6 0 , 1200 rpm, 0 2 -»-C30-P60,1200 rpm,02, High acidity - * - C 3 0 - P 6 0 , 1 2 0 0 rpm, 02,High Pulp, High acidity 10 15 20 Time (h) 25 30 35 Figure 4.62 Effects of impeller speed, choice of primary oxidant and acidity on solution potential during leaching of chalcopyrite However, it bears noting that, even under the best conditions, the system shown in Figure 4.59 did not leach to completion. The reason for this may be that the pyrite to chalcopyrite ratio (2:1) is insufficient under the highly oxygenated conditions of these tests. T o remedy to this situation, more pyrite was added and the test was started with 30 g L" 1 of acid (Figure 4.63). This had the effect of dramatically increasing the leaching rate such that the leach was complete within 4 hours. This test is compared with the only test using Temagani chalcopyrite, run at a pulp density 11 times lower with the same FeS2:CuFeS2 mass ratio of 4:1, and at an initial acid concentration of 10 g L ~ 1 . The results are virtually identical. This confirms that the pulp density can take any value desired without adversely affecting the efficacy of the leach. This is encouraging, since the capital cost of a leaching plant decreases proportionally with increasing pulp density. W h y simply increasing the initial concentration of acid should have had such a large effect on leaching kinetics remains a mystery. However, one possible explanation could be that the higher electrical conductivity of the solution with higher acid concentration Results and Discussion 159 somehow facilitates the transfer of electrons between the pyrite and the chalcopyrite particles; in other words, the acid may have a catalytic effect on the galvanic couple. Another possibility is that the higher acid concentration prevented the precipitation of basic ferric salts on the mineral surfaces, which may have impeded leaching kinetics. A third possibility is that the non-oxidative mechanism suggested by Lazaro and Nicol [111] may play an important role in this galvanically accelerated leaching tests. 4.5.7. Yield of elemental sulphur The final consideration is the yield of elemental sulphur. Results from the more successful tests are shown in Table 4.12. The elemental sulphur yields are consistently above 80%, and near quantitative yields are possible when oxygen gas is used as the primary oxidant. Time (h) Figure 4.63 Effect of pyrite addition, acidity and chalcopyrite origin on the leaching rate of chalcopyrite Results and Discussion 160 Table 4.12 Calculated elemental sulphur yields from selected tests Cp source Selwyn Selwyn Selwyn Selwyn Selwyn Selwyn Selwyn Selwyn Selwyn Selwyn Temagani mCp (g) 10 10 10 10 30 30 30 30 30 30 10 dcP (pm) -38 -38 +38 -75 -38 +38 -75 +38 -75 +38 -75 +38 -75 +38 -75 +38 -75 +38 -75 Py source Huanzala Huanzala Park City Park City Park City Park City Park City Park City Park City Park City Park City "ipy (g) 40 40 40 40 60 60 60 60 120 120 40 Opy (pm) -38 -38 -38 -38 -38 -38 -38 -38 -38 -38 -38 Initial E (mV (Ag/AgCI)) 538 405 470 405 450 438 462 480 468 480 468 [Fe] (g L"1) 5 5 5 5 5 5 5 5 5 5 5 [H 2S0 4] (g L"1) 10 10 10 10 10 10 20 20 20 30 10 Initial volume (mL) 1500 1500 1500 1500 1500 1500 1500 1000 1000 1000 1500 Impeller speed 750 750 750 750 1200 1200 1200 1200 1200 1200 1200 Oxidant Air Air Air Air Air Oxygen Oxygen Oxygen Oxygen Oxygen Oxygen Pulp density (%) 0.0333 0.0333 0.0333 0.0333 0.06 0.06 0.06 0.09 0.15 0.15 0.0333 Temperature (°C) 80 80 80 80 80 80 80 80 80 80 80 Final S° formation 81% 85% 88% 67% 71% 78% 89% 92% 89% 92% 90% Final S° stoichiometry 81% 85% 88% 77% 91% 104% 96% 97% 96% 97% 90% Max Cu recovery (%) 99% 98% 78% 86% 78% 75% 92% 95% 98% 99% 98% Duration (h) 15 21 11 22 24 21 19 21 12 8 6 Results and Discussion 161 Chapter 5 CONCLUSIONS The leaching of chalcopyrite in sulphate medium has been studied by various authors and it is widely recognized that poor copper extraction can be attributed to the formation of a passivating layer on the chalcopyrite surface. There are substantial differences of opinion concerning the mechanism and conditions under which the passivating layers are formed on chalcopyrite. In view of these discrepancies, the present study used electrochemical techniques to investigate surface changes during chalcopyrite leaching. Anodic and cathodic reactions during the ferric leaching of chalcopyrite were studied separately at various temperatures, applied potentials and leaching times. The following conclusions are drawn from the electrochemical results: > A progressively thickening passive layer is formed at low temperature in the potential range 0.45-0.6 V (Ag/AgCI) > T h e properties and stability of the passive layers are temperature and potential dependent > The formation of passive layers strongly inhibit ferric reduction at the chalcopyrite surface > The slow kinetics for ferric reduction on polarized chalcopyrite surfaces is a major factor contributing to chalcopyrite passivation > Ferric reduction is much faster on pyrite surfaces Based on these results a series of controlled potential experiments with fine particles were conducted at various temperatures and solution potentials in order to validate the electrochemical study and to determine intrinsic leaching kinetics. The galvanic interaction between pyrite and chalcopyrite was also investigated at various Fe(lll):Fe(ll) ratios. O n the basis of the experimental results, it was observed that the ferric leaching of chalcopyrite shows an induction period followed by an active dissolution period. The Conclusions 162 induction period decreases with increasing temperature. T h e breakdown of the layer causing this induction period may be linked to the thermal instability of the passive layer. Copper extraction increased rapidly with increasing solution potential up to a critical value beyond which copper extraction started to decrease. This critical value decreased with increasing temperature; i.e, the leaching of chalcopyrite is accelerated at high temperatures and low solution potentials. Copper extraction increased in the presence of a significant amount of pyrite. The galvanic effect was more pronounced at high FeS2 :CuFeS 2 ratios (>2) because pyrite provides more cathodic area for the reduction reaction, and the anodic dissolution rate of chalcopyrite must increase to compensate. T h e beneficial effect of pyrite stops when the two minerals both dissolve at high solution potentials. In this case, pyrite does not provide enough cathodic sites for ferric reduction. A n electrochemical model which takes into account the passivation of chalcopyrite was proposed. This model was also validated during galvanic interaction with pyrite. T he addition of microorganisms was beneficial at high temperature because extreme thermophiles generate a fairly low potential environment (350-500 m V (Ag/AgCI)). Copper extraction was measurably higher in the presence of pyrite. Bacterial leaching of chalcopyrite was retarded in the presence of mesophiles due to the formation of passivating films at high solution potentials (> 550 m V (Ag/AgCI)). T h e s e observations together with previous results show the importance of the solution potential, temperature and the addition of pyrite to chalcopyrite dissolution. These three main factors were tested during the atmospheric leaching of chalcopyrite. Leaching experiments were carried out at 8 0 ° C in order to avoid the induction period and accelerate the leaching process at low solution potentials. A certain quantity of pyrite (typically between 2 and 4 times the mass of chalcopyrite) was also added to the reactor. Complete extraction was achieved in as little as 4 hours with elemental sulphur yields above 80%. This is perhaps the shortest time possible for complete chalcopyrite conversion under atmospheric conditions in sulphate media without the aid of ultrafine grinding. Conclusions 163 The present study has shown that the combined effects of low solution potential, high temperature, pyrite addition and high acidity are desirable for the complete conversion of chalcopyrite in a relatively short time. This work could lead to a novel atmospheric process for leaching copper from chalcopyrite concentrates by admixing a certain amount of pyrite. Such a process would be especially attractive to those operations where the clean separation of pyrite and chalcopyrite by froth flotation is difficult. Conclusions 164 Chapter 6 FUTURE WORK AND RECOMMENDATIONS Based on the results of the present study, several suggestions for future work can be offered: 1. The addition of a certain amount of pyrite to the chalcopyrite leaching system has been shown to prevent chalcopyrite passivation. This effect was investigated with high grade chalcopyrite mixed with high grade pyrite ores. However, typical chalcopyrite ores contain a significant amount of associated pyrite. The galvanic interaction should be enhanced in this case because the two minerals are in intimate contact. T h e potential benefit of using low-grade chalcopyrite containing various amounts of pyrite should be investigated further. 2. T h e leaching of chalcopyrite in this study is surprisingly enhanced in the presence of higher acid concentration. 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References 178 122. Rojas-Chapana, J.A., Tributsch, H., Biochemistry of sulfur extraction in bio-corrosion of pyrite by Thiobacillus ferrooxidans, Hydrometallurgy (59), 2001: 291-300. 123. Lorenzen, L, An electrochemical study of the effect of potential on the selective dissolution of base metal sulphides in sulphuric acid, Minerals Engineering, Vol. 5(3-5), 1992: 535-545. References 1 7 9 Appendix A STATISTICAL ANALYSIS OF EXPERIMENTAL RESULTS Each leaching test was conducted in duplicate under the same conditions to examine the effect of reproducibility. A summary of copper extractions after 24 hours of leaching is presented in Table A . 1 . Mean: * = Z x / / n Variance: s 2 = £ ( X , - X ) 2 n-\ ;=1 / Std deviation: s = Vs2" Table A . 1. Statistical analysis of copper extraction after 24 hours of controlled leaching with peroxide T e m p ( ° C ) Fe(lll):Fe(ll) C u (%) X s 2 s 35 1 6.1 6.5 0.41 0.64 35 1 7 45 1 18.9 18.15 1.45 1.20 45 1 17.2 45 5 39.8 37.7 8.41 2.90 45 5 35.7 55 1 28.7 29.40 0.98 0.99 55 1 30.1 65 0.1 17.9 16.75 2.65 1.63 65 0.1 15.6 65 1 47.7 48.95 3.13 1.77 65 1 50.2 65 15 32.1 33.90 6.48 2.55 65 15 35.7 References 180 

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