Mechanism and Kinetics of Chalcopyrite Passivation and Depassivation During Ferric and Microbial Leaching
MECHANISM AND KINETICS OF CHALCOPYRITE PASSIVATION AND DEPASSIVATION DURING FERRIC AND MICROBIAL LEACHING
By
ALAIN FUAMBA TSHILOMBO
B.Eng., Faculte Polytechnique, University of Lubumbashi, 1994 M.Sc. , The University of Pretoria, 2000
A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF
DOCTOR OF PHILOSOPHY
In
THE FACULTY OF GRADUATES STUDIES (Department of Materials Engineering)
THE UNIVERSITY OF BRITISH COLUMBIA
December 2004 ABSTRACT
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. The 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 (25°C) and low potential (0.45 to 0.6 V SCE)
> high temperature (65°C) and high potential (above 0.6 V SCE)
Leaching was accelerated at high temperature (65°C) under mildly oxidizing conditions
(0.45 to 0.55 V SCE). The 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 65°C at a pyritexhalcopyrite mass ratio of 2:1. An
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 80°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. TABLE OF CONTENTS
ABSTRACT ii TABLE OF CONTENTS iv LIST OF FIGURES vii LIST OF TABLES xii LIST OF SYMBOLS xiii ACKNOWLEDGMENTS xv
CHAPTER 1. INTRODUCTION 1
CHAPTER 2. LITERATURE REVIEW 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. AAC/UBC Hydrometallurgy process 9 2.2.5. Dynatec process 10 2.2.6. Total Pressure Oxidation process 10 2.2.7. CESL process 11 2.2.8. BacTech/Mintek process 11 2.2.9. BioCop process 12 2.2.10. GeoCoat 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
CHAPTER 3. METHODOLOGY AND EXPERIMENTAL PROCEDURE 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
CHAPTER 4. RESULTS AND DISCUSSION 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
CHAPTER 5. CONCLUSIONS 162
CHAPTER 6. FUTURE WORK AND RECOMMENDATIONS 165
References 166
Appendix A. Statistical analysis of experimental results 180
vi LIST OF FIGURES
Figure 2.1 Crystal structure of chalcopyrite 5
Figure 2.2 Potential-pH diagram of the Cu-Fe-S-H20 system at 25°C: all solutes at 0.1 M activity except Cu2+ 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(S04)i.5 concentrations 18
Figure 2.4 Variation of reaction curves with temperature 19
Figure 2.5 Elemental sulphur-water Eh-pH 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
-1 CuFeS2 in 1 M H2S04, 40 mV min , 25°C 36
Figure 2.11 Anodic polarizarion curves for CuFeS2 from 6 different locations in
-1 1 M H2S04l 30 mV min , 25°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,
25°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, 25°C, de-aerated solutions 76
Figure 4.3 Effect of scan rate on the anodic dissolution of chalcopyrite, pH 1.5, 45°C, de-aerated solutions 76
Figure 4.4 Effect of scan rate on the anodic dissolution of chalcopyrite, pH 1.5, 65°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 25°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 25°C, pH 1.5 80
Figure 4.8 Constant potential experiments at 45°C on chalcopyrite in acidic solutions, pH 1.5 81
Figure 4.9 Constant potential experiments at 65°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, 25°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, 25°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, 25°C 89
Figure 4.15 Leaching rate curves of chalcopyrite in ferric sulphate solution, pH 1.4, 35°C 92
Figure 4.16 H202 added during leaching of chalcopyrite in ferric sulphate solution, pH 1.4, 35°C 92
Figure 4.17 Leaching rate curves of chalcopyrite in ferric sulphate solution, pH 1.4, 45°C 96
Figure 4.18 H202 added during leaching of chalcopyrite in ferric sulphate solution, pH 1.4,45°C 97
Figure 4.19 Leaching rate curves of chalcopyrite in ferric sulphate solution, pH 1.4, 65°C 98
viii Figure 4.20 H2O2 added during leaching of chalcopyrite in ferric sulphate solution, pH 1.4, 65°C 98
Figure 4.21 Leaching curves of chalcopyrite residues rinsed and not rinsed in carbon disulphide solutions, pH 1.4, 65°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, 65°C 104
Figure 4.24 Leaching rate curves of chalcopyrite at different FeS2:CuFeS2 ratios, Fe(lll):Fe(ll) = 1, pH 1.4, 65°C 105
Figure 4.25 Schematic representation of the galvanic interaction between chalcopyrite and pyrite 107
Figure 4.26 Effect of increased FeS2:CuFeS2 ratio on the relative anodic and cathodic areas 109
Figure 4.27 Leaching rate curves of chalcopyrite at different ratios of
FeS2:CuFeS2, Fe(lll):Fe(ll) = 15, pH 1.4, 65°C 111
Figure 4.28 Leaching rate curves of chalcopyrite from various sources,
FeS2:CuFeS2 = 2:1, Fe(lll):Fe(ll) = 15, pH 1.4, 65°C 113
Figure 4.29 Electrochemical-passivation model for the leaching of chalcopyrite in ferric sulphate solutions, pH 1.4, 65°C 123
Figure 4.30 Electrochemical-passivation model for the leaching of chalcopyrite
at different FeS2:CuFeS2 ratios, Fe(lll):Fe(ll) = 1, pH 1.4, 65°C 124
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, 65°C 125
Figure 4.32 Predicted values for the effect of particle size on the leaching of
chalcopyrite, FeS2:CuFeS2 = 2, Fe(lll):Fe(ll) = 1, pH 1.4, 65°C 126
Figure 4.33 Predicted values for the effect of increased FeS2:CuFeS2 ratio and particle size on the leaching of chalcopyrite, Fe(lll):Fe(ll) = 1, pH
1.4, 65°C, ofpy = 25 pm, c/Cp = 100 pm 127
Figure 4.34 Fraction of metal bioleached in the presence of mesophiles at 28°C .... 132
Figure 4.35 Molar ratio Cu:Fe during bioleaching in the presence of mesophiles at 28°C 132
Figure 4.36 Solution potential evolution and microbial counts during bioleaching in the presence of mesophiles at 28°C 134
Figure 4.37 Fraction of metal bioleached in the presence of moderate thermophiles at45°C 135
Figure 4.38 Molar ratio Cu:Fe during bioleaching in the presence of moderate thermophiles at 45°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
FeS2:CuFeS2l Initial Fe(lll):Fe(ll) =1, 80°C 148 Figure 4.50 Solution potentials during leaching of chalcopyrite at different ratios
of FeS2:CuFeS2, 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 FeS2:CuFeS2, 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, FeS2:CuFeS2 = 4, 80°C 155
x Figure 4.58 Effect of pyrite source on solution potential during leaching of
chalcopyrite, FeS2:CuFeS2 = 4, 80°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 LIST OF TABLES
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 35°C 91
Table 4.3 Extraction data for the ferric leaching of chalcopyrite at 45°C 94
Table 4.4 Extraction data for the ferric leaching of chalcopyrite at 65°C 97
Table 4.5 Results obtained during the leaching of chalcopyrite and chalcopyrite residues after 4 hours of leaching at 65°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 LIST OF SYMBOLS
E0 Standard reduction potential, V E Oxidation-reduction potential, V
P8o 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
_1 k Rate constant for the oxidation of ferrous ions by H202, M s a, Activity for species /'
mi Molal concentration for species /, mol/kg H20 Yi Activity coefficient for species /'
-2 /a Anodic current density, A cm
-2 /c Cathodic current density, A cm
-2 /0bs Sum of the anodic and cathodic cyrrent densities, A cm rpm Revolutions per minute, min-1 o Sulphate yield
2 /Acp Chalcopyrite surface, cm
2 >Apy Pyrite surface, cm
0Cp Fraction of the anodic area on chalcopyrite
0Py Fraction of the anodic area on pyrite A Total surface during the galvanic interaction, cm2
2 Aa,cp Anodic area on chalcopyrite, cm
2 Aa>py Anodic area on pyrite, cm
rr?cp Mass of chalcopyrite, g
mPy Mass of pyrite, g
-3 pCp Density of chalcopyrite, g cm
xiii -3 PPy Density of pyrite, g cm 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
Of a Anodic transfer coefficient da Cathodic transfer coefficient E m Mixed potential 1 rCp Dissolution rate of chalcopyrite, mol s~ r>c P 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 k p Rate constant for the growth of the passive layer kcp Rate constant Kcp Rate constant \
xiv ACKNOWLEDGMENTS
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 (CuFeS2) 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. These 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.
The principal oxidation reaction of chalcopyrite in acid ferric sulphate solutions can be
expressed by:
CuFeS2 + 2Fe2(S04)3 -> CuS04 + 5 FeS04 + 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 25°C [15]. The ferric leaching of chalcopyrite can then be interpreted
electrochemically by considering reaction (1.1) to take place in two steps:
2+ 2+ Anodic: CuFeS2 -> Cu + Fe + 2 S° + 4 e" £° = 0.47 V(SHE) (1.2)
Cathodic: 4 Fe3+ + 4 e" -> 4 Fe2+ E° = 0.77 V(SHE) (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. Among 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 Ag 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 (20-40°C) while thermophilic microorganisms are more active at
higher temperatures (40-90°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 65-75°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:
CU2S -» CU1.93S —» Cui .83S —» Cui.67S -> CU1.4S -> CuS
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
LITERATURE REVIEW
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, Fe-S and Cu-S are 0.368, 0.225 and 0.230 nm
respectively. The 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 cellare 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.
© Cu
Fe
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]:
Literature Review 5 + 3+ CuFeS2 + 3 h -> Fe + .CuS2 (2.1)
+ 2+ and .CuS2 + 2 h -> Cu + 2 S (2.2)
2+ or .CuS2 -> Cu + 2 S + 2 e" (2.3)
where .CuS2 is an unstable radical intermediate.
The first step indicates the breaking of the Fe-S bond and the second, the breaking of the Cu-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 Fe/Cu ratio in solution is greater than one in the earlier stages of leaching. Some 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 25°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. The effect of grind size was investigated and leach recovery was insensitive to grind size below Pso = 120 um: The process operates in an autoclave below the boiling point (~95°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:
Cu2S + Fe2(S04)3 -» CuS04 + CuS + 2 FeS04 (2.4)
Chalcocite leaching - Stage 2:
CuS + Fe2(S04)3 -> CuS04 + S + 2 FeS04 (2.5)
Ferric sulphate regeneration:
4 FeS04 + 02 + 2 H2S04 -> 2 Fe2(S04)3 + 2 H20 (2.6)
Overall chalcocite leaching reaction:
Cu2S + 02 + 2 H2S04 -> 2 CuS04 + S + 2 H20 (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:
FeS2 + 3.5 02 + H20 -> FeS04 + H2S04 (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%.
The elemental sulphur formed inside the autoclave remains relatively unattacked.
Copper recovery is by conventional solvent extraction and electrowinning (SX/EW). The electrolyte sent to electrowinning contains 48 g/L Cu. 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 110°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 (P80 = 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. The 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 SX/EW.
Western Minerals Technology in Perth has successfully run pilot scale campaigns.
2.2.3. MIM/Highlands Albion (Nenatech) process
The Nenatech process uses ferric leaching under atmospheric pressure, with oxygen or
air sparging at temperatures of about 80°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 SX/EW. The 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
150°C [30]. Leaching of chalcopyrite at 150°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 159°C, but increases dramatically between 159 and 190°C. It is therefore essential to work at temperatures below 159°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 150°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. The 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:
CuFeS2 + 1.25 02 + H2S04 -> CuS04 + 0.5 Fe203 + 2 S + H20 (2.9)
Formation of hematite and sulphate:
CuFeS2 + 4.25 02 + H20 CuS04 + 0.5 Fe203 + H2S04 (2.10)
With oxygen as a reaction product, the oxygen consumption is 1.25 mol/mol Cu; with sulphate as a reaction product, the oxygen consumption is 4.25 mol/mol Cu. Therefore, high elemental sulphur yields (80-100%) are desirable to reduce the oxygen consumption. Copper recovery is by conventional SX/EW. Precious metals in the ore can be recovered by cyanidation. A pilot plant program is in progress at the AARL 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 150°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 225°C) are used to dissolve the pyrite and/or arsenopyrite matrix. A Dynatec process for chalcopyrite was recently introduced and involves oxidative leaching at 150°C under a pressure of 10-12 atm [30,31]. The 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. The precious metals from leaching residue are treated by cyanidation. Iron is precipitated as hematite and jarosite.
Copper recovery is by conventional SX/EW.
2.2.6. Total Pressure Oxidation process
This process is owned by Placer Dome and uses high temperature oxidation conditions
(220°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 180-235°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 SX/EW. 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
The CESL process, owned by Cominco Engineering Services Limited of Vancouver,
BC, uses mixed sulphate-chloride solutions for pressure oxidation. It operates at about
150°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 SX/EW. 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 220°C to maximize the extraction and to produce a residue amenable to precious metals recovery by cyanidation. Copper is recovered by conventional SX/EW from a predominantly sulphate solution [28, 30,31].
2.2.8. BacTech/Mintek process
Mintek and BacTech 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. The 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.
BacTech 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 SX/EW [31].
2.2.9. BioCop process
BHP 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 42-45°C, while thermophiles leach chalcopyrite at 70-
80°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). BHP 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 GeoCoat 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 60-70°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. The 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. To 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.
To 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. The 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. The 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 Cu-Fe-S-H20 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 Cu-Fe-S-H20 system at 25°C: all solutes at
0.1 M activity except Cu2+ 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:
+ 5CuFeS2 + 2H2S + 5 Ox -> Cu5FeS4 + 4 FeS2 + 4 H + 4 Ox" (2.11)
Formation of chalcocite and pyrite:
+ 2 CuFeS2 + H2S + 2 Ox -> Cu2S + 2 FeS2 + 2 H + 2 Ox" (2.12)
Formation of covellite and pyrite
+ CuFeS2 + H2S + 2 Ox -+ CuS + FeS2 + 2 H + 2 Ox" (2.13)
Literature Review 14 Formation of covellite and elemental sulphur:
2+ CuFeS2 + 2 Ox -> CuS + Fe + 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-
pH 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 H2S as a reactant, while FeS2 is a
product. If H2S is depleted in solution during the leaching process and not supplied as
reactant, these reactions will not take place. Even if H2S 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
The 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
The 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
-40 pm was leached after 57 days at 35°C, using a 1% ferric sulphate solution. The 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.
The reactions by which chalcopyrite dissolves are believed to be:
3+ 2+ 2+ CuFeS2 + 4Fe -> Cu + 5 Fe + 2 S° (2.15)
3+ 2+ 2+ CuFeS2 + 4Fe + 2 H20 + 3 02 -> Cu + 5 Fe + 2 H2S04 (2.16)
The 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 Fe2(S04)3. showing a first-order
dependency, but declined at a higher concentration. Leaching experiments were
conducted at 70°C with 3 different specimens. Jones and Peters [38] observed similar
results. Increasing the concentration from 0.03 to 0.1 M Fe2(S04)3 increased the
leaching rate slightly, but a further increase to 1.0 M Fe2(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
3+ + was found that the principal Fe(lll) species in this system are Fe , FeS04 and
2+ FeHS04 . By using these data, Hirato and co-workers [37] plotted the distribution of
iron species as a function of Fe2(S04)3 concentrations in 0.2 M H2S04 solutions (Figure
3+ 2+ 2.3). The rapid decline in the concentration of Fe and FeHS04 at higher Fe2(S04)3 concentrations is evident in this figure. For example, for a thirty-fold increase in
+ Fe2(S04)3, the ferric activity increases by a factor of 6. By contrast, the FeS04
3+ 2+ concentration increases linearly. They concluded that the Fe and/or FeHS04 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 25°C.
He observed that the rate of sphalerite dissolution was proportional to the sum of the
3+ 2+ concentration of the Fe and FeHS04 species in solution.
Literature Review 17 10"'
10"' ?E •D O E
! u.- 10"
" 10"3 10"2 10"1 I
C(Fe(S04)l5)/ 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 90°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 (SCE) - indicating the formation of a copper-rich intermediate. X-ray diffraction examination of the residue revealed the formation of CuS 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 Cu and Fe were released into solution. The 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. The 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. They 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:
2+ ,2: + 3+ (2.17) CuFeS2 + 3 Cu + 3 Fe 2 Cu2S + 4 Fe
Next, Cu2S is oxidized by dissolved oxygen and/or ferric to form cupric:
+ 2+ (2.18) 2 Cu2S + 8 H + 2 02 -> 4Cu + 2 S° + 4 H20
and/or
Literature Review 19 3+ 2+ 2+ 2Cu2S + 8Fe -> 4 Cu + 8 Fe + 2 S° (2.19)
Following the proposed model, enhanced copper extraction by ferrous occurs only at
2+ lower potentials when enough Cu ions are present in solution. The intermediate Cu2S 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 FeS04 (0.1 M FeS04), added
to a 0.1 M Fe2(S04)3 + 0.3 M H2S04 solution, could cause the rate of chalcopyrite leaching to fall appreciably (by approximately 30%) at 80°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. The 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 Fe3+) or the
chalcopyrite particle size. Experiments were conducted at 95°C in acidic ferric sulphate
solution. The 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
v HS0,- so4- ^v
05
Extended so,- Sulphur >>. y Zone 5 g 0 (so,-) I xz LU
-05
*»» *
H2S HS" 1 atm. 1m. is- 1 ., i i "—1 " 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]. The 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. The presence of graphite embedded in the insulating suphur layer would improve electron transport, since graphite is
conductive. The 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. The most consistent
characteristic of chalcopyrite leaching is the high activation energy associated with the
reaction, ranging from 70 to 84 kJ/mol [42-45]. These large activation energies are
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
> Ea > 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 70°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]. The reason is not
clear, but Peters [48,49] believes that if the oxidation of one mole of chalcopyrite (42.65
cm3) produces 100% elemental sulphur (31 cm3), 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]. They 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. As 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 95°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% Cu extraction was reported with corresponding amounts of elemental sulphur formation. The 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.
An 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 H C 02M F«(S04)|S. O-.S" Z$°4- ?°* 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]. When 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:
Fe2(S04)3 + 6H20^2 Fe(OH)3 + 3 H2S04 (2.20)
In addition to the formation of Fe(OH)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:
Fe2(S04)3 + 2 H20 -> 2 Fe(OH)S04 + H2S04 (2.21)
Goethite:
+ 4 Fe2(S04)3 H20 2 FeOOH + 3 H2S04 (2.22)
Jarosite:
+ +
3Fe2(S04)3 + 12H20 + 2 A -> 2 AFe3(S04)2(OH)6 + 5 H2S04 + 2 H (2.23)
+ + + + 2+ where A represents K , Na , NH4\ H30 , Ag or (Pb )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. The 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 100°C. At pH values above 3, complete ferric precipitation occurs and forms
highly insoluble iron precipitates (ferric hydroxide, goethite, hematite). Hematite forms
only at temperatures higher than 100°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]
The 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. The 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. As
Literature Review 26 shown in Figure 2.8, pH values less than 1.8 limit the extent of Fe3+ 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:
CuFeS2 + 402 -> CuS04 + FeS04 (2.24)
The ferrous sulphate product is oxidized by oxygen in the presence of microbes:
4 FeS04 + 2 H2S04 + 02 -> 2 Fe2(S04)3 + 2 H20 (2.25)
The ferric sulphate product then reacts with chalcopyrite to form more ferrous sulphate:
CuFeS2 + 2 Fe2(S04)3 CuS04 + 5 FeS04 + 2 S° (2.26)
The elemental sulphur can also be oxidized to acid:
2 S° + 3 02 + 2 H20 -> 2 H2S04 (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.
The 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]. The 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]. To 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. As 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. And 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]. The 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
The 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
To become active, bacteria need oxygen, carbon dioxide, nitrogen, and phosphate.
+ + 3 Nitrogen is generally supplied in the form of NH4 . The addition of NH4 and P04 ~ 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 C02 fixation [66].
Types of microbes
It is also important to understand the behavior of microbes on the mixed substrates Fe2+ 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
As 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. The electrochemical dissolution of chalcopyrite can be written in terms of the half reactions for the oxidation of chalcopyrite and the reduction of ferric:
2+ 2+ Anodic: CuFeS2 -> Cu + Fe + 2 S° + 4 e~ (2.28)
Cathodic: 4 Fe3+ + 4 e~ -> 4 Fe2+ (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. The 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 25°C [71].
Formula Name Resistivity (Q-m) Conductor Structure Ionic structure 3 -6 Tetragonal + 3+ 2 Cu5FeS4 Bornite 10~ - 10 P (Cu )5Fe (S -)4
+ 2_ 2 5 Orthorhombic (Cu ) S Cu2S Chalcocite 4x10" -8x1Cr P 2
4 -3 + 3+ 2
CuFeS2 Chalcopyrite 2x10" -9x10 n Tetragonal Cu Fe (S ")2
5 7 + 2 CuS Covellite 8x10" -7x10" metallic hexagonal (Cu )2S2 - PbS Galena ixio^-zxio-6 n&p Cubic Pb2+S2"
3 4+ 2
MoS2 Molybdenite 7.5- 8x10" n&p Hexagonal Mo (S ")2
2 3 2+ 2 3X1CT -1X10- Fe S2 " FeS2 Pyrite n&p Cubic ZnS Sphalerite 3X10"3-1X10^ Cubic Zn2+S2"
+2 2 4+ 2 Sn02 Cassiterite 10 -10" n Tetragonal Sn (0 ")2 1 (Cu+) 02_ Cu20 Cuprite 10~ - 10 P Cubic 2
1 2 3+ 2 Fe203 Hematite 2.5x10" -4x10" n&p Trigonal (Fe )2(0 -)3
5 3+ 3+ 2+ 2-
Fe304 Magnetite 2x10^-4x10" n&p Cubic Fe [Fe Fe ](0 )4
1 3 4+ 2 Mn02 Pyrolusite 10" -10" n Tetragonal Mn (0 ")2
4 4+ 2 Ti02 Rutile 10 - 10 n&p Tetragonal Ti (0 ")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:
2+ 2+ CuFeS2 -> Cui_xFe1_yS2-z + x Cu + y Fe + z S° + 2 (x+y) e" (2.30)
where y> x initially.
Literature Review 33 The intermediate or defect structure Cui_xFei-yS2-z releases ferrous much faster than cupric. The shorter the time of measurement, the higher is the Fe:Cu 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 Cu2+ and Fe2+ 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_xFei_yS2-z decomposes further to form a second (non-
stoichiometric) intermediate, CuS(n-S), which forms in the region 700 to 750 mV (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) Cu2+ +(1-y) Fe2+ + (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 Cu2+ and Fe2+ 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. The first
stage involves the release of Cu2+ and acts as the limiting step of the reaction, whereas
the second stage involves the release of Fe2+, in accordance with the following scheme:
2+ nCuFeS2 -> Cu + Cu,^Fe„S2n + 2 e~ (rate-limiting) (2.32)
3+ CUn-iFenS2n -> Fe + Cun_1Fen_1S2n + 3 e~ (2.33)
Literature Review 34 The overall reaction can be expressed by:
2+ 3+ nCuFeS2 -»• Cu + Fe + Cu^Fe^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. The 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 CS2 does not remove the passive film from the
chalcopyrite surface. They conducted potentiostatic experiments after CS2 washing
within the potential range 450-900 mV (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
2+ 2+ semiconductor properties from CuFeS2 and slows not only transport of Cu and Fe
but also electron transfer to or from electroactive species in solution. Reddy et al. [11],
using ESCA 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_ SAMPLE 3E, TRANSVAAL
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
-1 M H2S04) 40 mV min , 25°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 25°C in a solution of 1 M H2S04 (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]. On 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*mo.«» ««ico — —. suoeuRv.oNTftfiio
0.5--
IO"6 IO"5 IO"4 10"3 CURRENT DENSITY (A/cmz)
Figure 2.11. Anodic polarizarion curves for CuFeS2 from six different locations in 1 M
-1 H2S04, 30 mV min , 25°C [70]
Table 2.2. Reactions of chalcopyrite [69]
Reactions £° (V (SHE))
2+ (2.35) + 0.50 CuFeS2 FeS2 + Cu + 2e~
2+ (2.36) + 0.81 CuFeS2 -» FeS + S° + Cu + 2e~
3+ + 0.45 CuFeS2 -> CuS + S° + Fe + 3e~ (2.37)
2+ 2+ + 0.47 CuFeS2 -> Cu + Fe + 2 S° + 4e" (2.38)
2+ + 0.87 2 CuFeS 2 -> Cu2S + 2 Fe + 4 S° + 4 e" (2.39)
2+ 3+ + 0.51 CuFeS2 -> Cu + Fe + 2S° + 3e~ (2.40)
CuS -> Cu2+ + S° + 2 e" (2.41) + 0.47
Fe2+ Fe3+ + e" (2.42) + 0.77
Based on the potential-pH diagram for the Cu-Fe-S-H20 system and the data presented
in Table 2.2., Warren et al. proposed that S1 and S2 are bornite (Cu5FeS4) 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 70°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:
2+ 2+ 4CuFeS2 3 CuS + Cu + 4 Fe + 5 S° + 10e" (2.43) where CuS (covellite) and S° would be in a ratio of 3:5 and Fe2+:Cu2+ in solution would be in a ratio of 4:1.
On 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:
2+ CuFeS2 -» Cu + FeS2 + 2 e" (2.44)
Cu2+ x Cu2+ (ads) + (1-x) Cu2+ (aq) (2.45)
2+ 2+ Cu (ads) + FeS2 -+ CuS + Fe + S° (2.46)
Jones and Peters [72] reported the formation at 175°C of large and well-defined crystals of covellite, CuS, 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. As 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 CuS 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, CuS„, 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:
-> -> CULB/S -> Cu2S C-U1.93S -> CU1.83S CU1.4S -> CuS
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 39 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 25°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
On 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. As 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 Type 1 E
X /
^ /
EM
^x ' \ \ I \J 1 \
In jd In j
Figure 2.13. Mixed potential for Type I leaching
Type II
i+ 2+ Fe <-> Fe /
/i
In jd In j
Figure 2.14. Mixed potential for Type II leaching
Literature Review . Type III
Fe3* <-> Fe2+
/ X • A /
EM
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
2 U02 + Fe(lll) S04 " Fe(lll) (0.5) I
2 FeS + Fe(lll) S04 " Fe(lll) (0.5), Fe(ll) (-0.5) II
2
FeS2 + 02 S04 " 02, (0.5) I
2 FeS2 + Fe(lll) S04 " Fe(lll) (0.5), Fe(ll) (-0.5) II
FeS2 + 02 OH" 02, (0.31) I
2 ZnS + Fe(lll) S04 " Fe(lll) (0.45), Fe(ll) (-0.5) II ZnS + Fe(lll) cr Fe(lll) (0.36), Fe(ll) (<-1) IV
2 CuFeS2 + 02 S04 " 02, (0.35) I
CuFeS2 + Fe(lll) cr Fe(lll) (0.5, 0.3-0.4) I
CuFeS2 + Cu(ll) cr Cu(ll) (0.54), Cu(l) (-0.5) II
Ag2S + CN" + 02 CN" 02, (0.5) I
PbS + 02 OH" 02, (0.5) I
2 Ni3S2 + Fe(lll) S04 " Fe(lll) (0.2), Fe(ll) (-0.2) II
2_ Cu2S + Fe(lll) stage 1 S04 Fe(lll) (1.0), III
2_ Cu2S + Fe(lll) stage II so4 Fe(lll) (0.4), Fe(ll) (-0.3) II
Au + CN" + 02 CN" CN" (0 or1), 02, (0or1) III REDUCTIVE
2 Mn02 + S02 S04 " S02 (0.5) I
2 Mn02 + Fe(ll) S04 " Fe(ll) (0.58) I
2 FeOOH + S02 S04 " S02 (0.5-0.76) I
2 II Fe304 + Fe(ll) S04 " Fe(ll) (0.5), Fe(lll) (-0.5)
2 II ZnFe204 +. Fe(ll) S04 " Fe(ll) (0.6), Fe(lll) (-0.5)
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. As
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
As 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 93°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 100°C and under 2 MPa 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 Ag 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 100°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 Ag2S according to the equation (2.47):
+ 2+ 2+ CuFeS2 + 4Ag -> 2 Ag2S + Cu + Fe (2.47)
The silver ions are regenerated by the reaction between ferric and Ag2S :
3+ + 2+ o Ag2S + 2 Fe -> 2 Ag + 2 Fe + S' (2.48)
Ag(l) thus released will again react with chalcopyrite. They 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 Ag2S electrodes in 1 M H2SC>4:
+ Ag2S -> 2Ag + S° + 2 e" (2.49)
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
Literature Review 46 mechanism the reactivation effect due to Ag2S can be explained by an electrochemical
mechanism based on the formation of the Ag2S/CuFeS2 galvanic cell, with Ag2S acting as cathode and CuFeS2 acting as anode [90].
The beneficial effect of silver was also observed during bioleaching of chalcopyrite. The 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 (Hg2+, Co2+, Bi3+ and Cu2+) 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]. As 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. The rest potential for chalcopyrite is lower than that of pyrite as shown in Table 2.4. When 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
Rest potential (V (SHE)) in Mineral References 1 M H2SO4 at 20°C
FeS2 0.63 Hiskey [71] Hiskey [71] CuFeS2 0.52
Ag2S 0.51 Warren [98]
Cu2S 0.44 Hiskey [71] CuS 0.42 Hiskey [71]
PbS 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 CuFeS2:FeS2 ratio of 5:5 (g/g) and size fraction of
-74 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 mV to 570 mV
(SCE). 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. The 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 37°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. The 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 mV (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, BHP Billiton, BacTech 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.
The 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
As 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 (FeS2) 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. The 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 cm2. 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 purge tube (N2 gas) computer 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 (SCE) 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 (SCE) [34]. The scan rate was
1 mV s_1 unless otherwise noted. The anodic characteristics of chalcopyrite were
studied in acid solution (pH 1.5) at 25, 45 and 65°C. The 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 (25-35°C), moderate thermophiles are active between 40 and 55°C and above
55°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 Cu 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. The 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 [FeS04.7H20] (99% purity)
> Iron (III) sulfate pentahydrate [Fe2(S04)3.5H20] (97% purity)
> Hydrogen peroxide [H202] (30% solution)
> Sulfuric acid [H2S04] (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. The 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.
The desired temperature was controlled within 1°C by circulating water through the reactor jacket. The reactor was completely sealed in order to reduce evaporation losses during leaching. Particles were suspended at 750 rpm. The 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 pH. 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. The temperature range of the probe was 0 to 100°C. The 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. The 4 M KCI filling solution was changed regularly. Cole-Parmer Masterflex peristaltic pumps were used for H2O2 and
H2SO4 solution additions into the reactor.
Stirrer
Condenser
H202 port (tubing to pump not shown)
Water Stirrer »?02 jacket controller purnp
? # f w
Sparger pH and Eh 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. The retention time was 24 hours.
The redox potential in the leaching solution was maintained to within 2 mV 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 65°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.
3 E = E° + *T\n 'Fe (3.1)
VaFe2+ J
Using molal concentrations instead of activities, Eq. (3.1) is rearranged to:
3+ £°=£°'+—In 'Fe (3.2)
V"'Fe2V
3+ where £° =E°+—In 'Fe (3.3)
V'Fe2V
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). The solution potential at the
concentration ratio of 1:1 was about 0.71 V (SHE). The 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 Fe3+/Fe2+ = 1
> pH = 1.3
> Temperature = 65°C
> Mass of chalcopyrite: 30 g
> Volume of solution: 1.5 L
0.7
0.4 -I 1 1 1 . 1 -10 12 3 4
log ([Fe3+]/[Fe2+])
Figure 3.3. Solution potential as a function of the ferric to ferrous ratio, pH 1.4, 25°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. The 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. On 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 Fe2+ ions and regenerating Fe3+ 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:
+ - H202 + 2 H + 2 e = 2 H20 (3.4)
When hydrogen peroxide is added to a ferrous solution, the following reaction takes place:
2+ + 3+ HOOH + 2 Fe + 2 H = 2 Fe + H20 (3-5)
Methodology and Experimental Procedure 62 The oxidation rate is given by:
2+ Rate = k [Fe ] [H202] (3.6)
The rate constant is 41.4 M_1 s_1 at 20.2°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. The 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 30°C, were used for this study. The nutrient
contained the following salts: 2 g/L MgS04.7H20, 0.8 g/L (NH4)2S04l 0.3 g/L KH2P04l 1
g/L FeS04.7H20.
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 30°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 48°C was used. For extreme thermophilic
leaching, a mixed culture of containing strains of Sulfolobus, Acidianus and
Metallosphaera cultured at 68°C was used. The growth medium for moderate/extreme
thermophiles contained the following salts: 0.7 g/L MgS04.7H20, 1.5 g/L (NH4)2S04, 0.3
g/L KH2P04, 0.5g/L FeS04.7H20. 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 48°C for the moderates and
68°C for the extreme thermophiles. A stirring speed of 150 rpm was used in the two
incubators.
3.3.3. Equipment
An 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 65°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 60°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 80°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 Cu and Fe by atomic absorption spectrophotometry. Once 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 Ocean Sciences at the University of British
Columbia. The 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 SYBR Green 1 (Molecular Probes, OR). Each filter was incubated with 70 uL of dye (0.3 % SYBR 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 -20°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. The slurry was vortexed vigorously for 5 s, followed by 30 min mixing at 225 rpm in a rotary orbit shaker. The 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 68°C with fine particles of high grade chalcopyrite coated on low-grade chalcopyrite mineral. The GeoCoat 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 GeoCoat process may also be
advantageous in terms of operating a thermophile chalcopyrite heap leach [16]. The
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. The mineralogical composition and chemical analysis of the
concentrates are given in Tables 3.1, 3.2, and 3.3. The 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 (Si02) 41.5
Chalcopyrite (CuFeS2) 0.4
Pyrite (FeS2) 3.5
Muscovite (KAI2(AISi3Oio)(F, OH)2) 31.5
Albite (NaAISi308) 5.9
Kaolinite (AI2Si205(OH)4) 12.6
Clinochlore ((Mg,Fe,AI)6(Si, AI)4O10(OH)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 bottom plate, which closed the column. All effluents were collected in the center of the bottom 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 balls covering the ore in each column. The columns were immersed in a plastic tub filled 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.) C02, was blown into the base of each column at a rate of approximately 360 mL min-1.
plastic balls
concentrate on support rock
plastic balls gas sparger
spacer ring f —gieve p|a{e
gas hose pipe flange _
gas supply
received solution
Figure 3.4. Schematic drawing of a column
Methodology and Experimental Procedure 67 Figure 3.5. The 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. The 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. The copper and iron concentrations in the liquid samples were determined by AA 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. The 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. The 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. The concentrate received from
the mine was crushed, ground and screened to the required size fractions. The
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(%) Cu 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 [FeS04.7H20] (99% purity)
> Iron (III) sulfate pentahydrate [Fe2(S04)3.5H20] (97% purity)
> Sulfuric acid [H2S04] (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). The reactor was maintained at 80°C
by circulating water through the reactor jacket. Particles were suspended at 750 rpm.
Air or 02 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:
2+ 2+ - CuFeS2 Cu + Fe + 2 S° + 4 e (4.1)
The anodic characteristics of chalcopyrite were studied in acid solution (pH 1.5) at 25,
45 and 65°C, as shown in Figure 4.1. The open circuit potential (OCP) was about 0.34 V
(SCE) (0.58 V (SHE)) at 25°C. This value is very close to the rest potential of
chalcopyrite reported in the literature (0.52 V (SHE) in 1 M H2S04) [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 OCP at 65°C was approximately 50 mV more negative
than that at 25°C. This lower OCP indicates that the chalcopyrite surface was becoming
more active and less noble at higher temperatures. As the scan proceeded in the
positive potentials, the anodic polarization curves of chalcopyrite at various
temperatures revealed three electro-active regions. The first region, from 320 to 450 mV
(SCE), indicates that the current increased exponentially with applied potential, with
Tafel slopes ranging from 0.12 to 0.16 V decade-1. The plateau region in the potential
range 450 to 600 mV (SCE), indicates that the current levels out. The third region, at
Results and Discussion 71 potentials higher than 600 mV (SCE), 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. The
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 106 times lower than the
maximum rate in the active state [107]. The test for the anodic dissolution of chalcopyrite at 25°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:
2+ 2+ CuFeS2 -> New phase + m Cu + n Fe + 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:
+ 2+ CuFeS2 + 2 H = CuS + Fe + H2S (4.3)
Oxidation of CuFeS2 and CuS could then occur in parallel:
2+ 2+ CuFeS2 = Cu + Fe + 2 S + 4 e" (4.4)
CuS = Cu2+ + 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, pH 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 + : CuFeS2 + 2xH = CuFe(i_x)S(2_X) + xH2S + xFe (4.6)
The copper-rich phase may then be oxidized on the subsequent anodic potential sweep to release Cu2+. 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 OCP 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:
2+ 2+ CuFeS2 -> Cu^xFd-yS^z + x Cu + y Fe + 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 Woods [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 CuS2. 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 S42~. This metal-deficient sulphide (or polysulphide) had different semi• conductor properties from the underlying chalcopyrite. Thus, his composition
approaches Cu2S4 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 mV 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 mV 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 (25°C), passivation is evident
by the plateau region and the low currents in the potential range 0.45 to 0.55 V (SCE).
At moderate temperature (45°C), the curve does not show a plateau region within the
potential range 0.45 to 0.55 V (SCE) and the anodic currents are much higher. This
indicates that the stability of the passive layer diminished as the temperature increased.
At high temperature (65°C), the anodic current increases rapidly within the potential
range 0.45 to 0.55 V (SCE), 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 25°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 95°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.
As 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, 25°C,
de-aerated solutions
Figure 4.3 Effect of scan rate on the anodic dissolution of chalcopyrite, pH 1.5, 45°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, 65°C,
de-aerated solutions
Figure 4.5 Effect of temperature on the anodic dissolution of chalcopyrite, pH 1.5,
scan rate = 0.1 mV 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 (SCE) 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 25°C shows a drop of 80% in the current at 0.45 and 0.55 V
(SCE) (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. To 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 (SCE) 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. The OCP of polarized surfaces was about 50 mV more positive than that of a "clean" surface. The exchange current densities also decreased for polarized
surfaces. As 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 (SCE). 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 45°C shows a drop of 60% in the current at 0.45 V (SCE)
(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 (SCE).
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.
The time-current curves at 65°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 45°C. It seems that the passive layers formed at 65°C
are less stable than those formed at 25 and 45°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
45°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. The 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 Cu2+, Fe3+ and SCv2- 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 Woods [114], postulated that the anodic
dissolution of chalcopyrite was associated with the following reaction:
+ CuFeS2 + 3 H20 = CuS + Fe(OH)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 25°C on chalcopyrite in acidic solutions,
pH 1.5
Figure 4.7 Anodic behaviour of fresh and polarized chalcopyrite surfaces in acidic
solutions at 25°C, pH 1.5
Results and Discussion 80 0.0002 0.00018
0.00016
0.00014 _ 0.00012 c~f
| 0.0001 E = 0.6 V ~ 0.00008 E = 0.55 V 0.00006 0.00004 E = 0.45 V 0.00002
0 —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 45°C on chalcopyrite in acidic solutions,
pH = 1.5
0.0004
0.00035
| 0.0002 • • •• • B-B-B E = 0.55 V 0.00015 E = 0.45 V 0.0001
0.00005
0 2 4 6 8 10 12 14 16 18 20 22 24 Time (H)
Figure 4.9 Constant Potential Experiments at 65°C on Chalcopyrite in Acidic
Solutions, pH = 1.5
Results and Discussion 81 To 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 Potential Temp Cu Fe Ratio I (H) (V (SCE)) (°C) (10~6M) (10"6 M) Cu:Fe (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 (SCE). 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 (SCE), 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 (SCE) (Cu:Fe =1:1), at 45 and
65°C compared with only 1:3 at 25°C. At high temperature (65°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 (SCE)):
2+ 2+ CuFeS2 -> Cui_xFei_yS2-z + x Cu + y Fe + z S° + 2 (x+y) e" (4.8) where y > x.
Second region (0.45 to 0.6 V (SCE)):
2+ 2+ Cui_xFei_yS2-z--> (2-z)CuS(n-s) + (z-x-1)Cu + (1-y) Fe + 2 (z-x-y) e" (4.9)
where CuS(n-S) represents a nonstoichiometric intermediate product phase.
Third region (above 0.6 V (SCE)):
CuS -> Cu2+ + S° + 2e" (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:
3+ > Cathodic reduction of oxidants, such as Fe and 02
> 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 Fe3+ + 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. The 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 (SCE).
As 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. The 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)
The 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 (SCE) 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 (FeS2) and pentlandite ((Fe,Ni)9S8) 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.,
1 + 2+ FeS2 + /2 02 + 2 H -> 2 S° + Fe + H20 (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 Eh-pH 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
As 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 (SCE) due to the contribution of the
anodic current of the mineral to the total current. The 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 (SCE) [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 decade-1. To 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 mV 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 25°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, 25°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, 25°C
Results and Discussion 89 Electrochemical measurements have demonstrated that chalcopyrite "passivation" is
mainly observed under the following conditions:
> Low temperature (25°C) and low potentials (0.45 to 0.6 V)
> High temperature (65°C) and high potentials (above 0.65 V)
Leaching of chalcopyrite was accelerated at high temperatures 65°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. These experiments were
conducted at temperatures varying from 35 to 65°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 35°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 35°C
Solution Cu S° Final S° Temp Leaching Fe(lll):Fe(ll) potential Test ID extraction formation stoichiometry (°C) time (h) ratio (mV (Ag/AgCI)) (%) (%) (%) 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 35°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%.
Cu 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 35°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. The second stage indicates
linear kinetics at longer reaction time. An 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,
35°C
40 -y
0 5 10 15 20 25 Time (H)
Figure 4.16 H202 added during leaching of chalcopyrite in ferric sulphate solution, pH
1.4, 35°C
Results and Discussion 92 The two-stage kinetics can be expressed by the following reactions:
Stage 1:
3+ 2+ 2+ CuFeS2 + 2(x+y)Fe -> Cu^xFe^^-z + xCu + (2x+3y) Fe + zS° (4.14)
Stage 2:
Cih-xFei-ySa-z + (4-2x-2y) Fe3+ -> (1-x) Cu2++ (5-2x-3y) Fe2+ + (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. The 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). As 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 35°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 45°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 45°C
Solution Cu s° Final S° Temp Leaching Fe(lll):Fe(ll) potential Test ID extraction formation stoichiometry (°C) time (h) ratio (mV (%) (Ag/AgCI)) (%) (%) 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
45°C was much faster compared to the results obtained at 35°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 35°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 45°C disrupts the formation of the passive layer formed in the earlier stage of leaching. Electrochemical studies have shown that the intermediates formed at 25°C were more stable than those
Results and Discussion 94 formed at 45°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.
The 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
(SCE). They also mentioned that a portion of Cu was precipitated as CuS (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 45°C can be expressed by the following reaction:
3+ 2+ Stage 1: CuFeS2 + 2 Fe -> CuS + 3 Fe + S° (4.17)
Stage 2: CuS + 2 Fe3+ Cu2+ + 2 Fe2+ + S° (4.18)
The 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,
45°C
Measurements at 65°C
The leaching of chalcopyrite at 65°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
65°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). The 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 mV (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, 45°C
As 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 45°C and 6% at 35°C.
Table 4.4 Extraction data for the ferric leaching of chalcopyrite at 65°C
Solution Cu S° Final S° Temp Leaching Fe(lll):Fe(ll) potential Test ID stoichiometry (°C) time (h) ratio (mV extraction formation (Ag/AgCI)) (%) (%) (%) 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,
65°C
Figure 4.20 H2O2 added during leaching of chalcopyrite in ferric sulphate solution, pH
1.4, 65°C
Results and Discussion 98 The leaching of chalcopyrite did not show an induction period during the initial stage of leaching at 65°C. Clearly, the presence of an induction period is mainly a temperature dependent phenomenon, which takes place below 65°C.
The oxidation reaction at 65°C can be expressed by the following reaction:
3+ 2+ 2+ CuFeS2 + 4 Fe -> Cu + 5 Fe + S° (4.19)
The 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
Cu2S (Eq.(4.20)) and the intermediate Cu2S is oxidized by ferric to release cupric (Eq.
(4.21)) as:
2+ 2+ 3+ CuFeS2 + 3 Cu + 3 Fe -> 2 Cu2S + 4 Fe (4.20)
3+ 2+ :2 + Cu2S + 4 Fe -> 2 Cu + S° + 4 Fe (4.21)
The accelerating effect can only take place if there are enough cupric ions in solution.
The oxidation rate of the intermediate Cu2S 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
(4-22):
+ 2: + 2 CuFeS2 + 6 H + 2 e" -> Cu2S + 3 H2S + 2 Fe (4.22)
The anodic reaction is:
Fe° Fe2+ + 2 e" (4.23)
Results and Discussion 99 giving the overall reaction :
+ : 2 CuFeS2 + 6 H + Fe° -> Cu2S + 3 H2S + 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.
The 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.
The 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 (CS2) solution to remove sulphur. The residue collected after washing with
carbon disulphide was leached again in ferric sulphate solutions. The results shown in
Figure 4.21 indicate that the residues washed in CS2 dissolve faster than those without
a CS2 treatment. Table 4.5 also indicates that the removal of sulphur does not increase
significantly the overall leaching rate of chalcopyrite. The average amounts of copper
released from the residues treated in CS2 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 65°C.
Cu dissolved from Cu dissolved from Cu dissolved from the residue not the residue Fe(lll):Fe(ll) ratio fresh mineral (mg treated in CS (mg treated in CS2 (mg 2 L-1 h"1) L-1 h"1) 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 •- R = 5, CS2
I- R= 1, CS2 0.08 -\ •- R = 0.33, CS2 »— R = 5, No CS2 •a 0.07 I—R = 1, No CS2 u >— R=0.33, No CS2 !» ~ 0.06 4) O 0.05
I 0.04 H
Figure 4.21 Leaching curves of chalcopyrite residues rinsed and not rinsed in carbon
disulphide solutions, pH 1.4, 65°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 35°C the induction period lasted about 6 hours, while at 45°C it only lasted about 4 hours. The leaching curves at 55 and 65°C did not show an induction period.
As 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 65°C. The choice of this temperature was based on the results of the previous section, which indicate no induction period above 55°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
The leaching of pyrite in solution containing Fe(lll) and Fe(ll) at 65°C was conducted at low solution potential (Fe(lll):Fe(ll) = 1) and high solution potential (Fe(lll):Fe(ll) = 15).
The oxidation of pyrite by ferric can be expressed by the following reaction:
3+ FeS2 + 8a H20 + (2 +12c) Fe (4.25) 2+ (3 + 12cr)Fe +2CTS02- +(2-2a)S° +16CTH+ 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. To 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, 65°C
Effect of pyrite on the leaching rate of chalcopyrite at low solution potential (Fe(lll):Fe(ll)
= 1)
The galvanic interaction between chalcopyrite and pyrite at low solution potential was conducted at 65°C with a Fe(lll):Fe(ll) ratio of 1. Pyrite of various amounts were added
to chalcopyrite to produce FeS2:CuFeS2 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 Santa Santa Santa Chalcopyrite source Eulalia Eulalia 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 cm-3) 4.2 4.2 4.2
Density of pyrite (g cm-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 • Chalco =30, Pyr = 0 0.8 Chalco = 20, Pyr =10 - Chalco = 10, Pyr = 20
0)
10 Time(H) 15 20 25
Figure 4.24 Leaching rate curves of chalcopyrite at different FeS2:CuFeS2 ratios,
Fe(lll):Fe(ll)= 1, pH 1.4, 65°C
Results and Discussion 105 Figure 4.24 shows copper extraction versus time at different ratios of FeS2:CuFeS2. As 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. When 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 FeS2:CuFeS2 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 (ACp)can 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 (/\Py) 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. The total surface during the galvanic interaction is given by:
A = ACp+APy (4-30)
Results and Discussion 106 Fe3 + Fe2 + Fe3 + Fe2 +
Chalcopyrite Chalcopyrite Pyrite Pyrite Anodic area Cathodic area Cathodic area Anodic
3+ 2+ 2+ 2 2+ 2+ 3+ 2+ Fes =Fe + 2S0 " CuFeS2=Cu +Fe + 2S +4e Fe + e-=Fe Fe + e" = Fe 2 4 +16H++ 14 e"
Figure 4.25 Schematic representation of the galvanic interaction between chalcopyrite
and pyrite
The total anodic and cathodic areas for the leaching of chalcopyrite in galvanic contact with pyrite can therefore be expressed by:
+ + (4.31) \ - Ai.Cp Ai,Py ~~ ^Cp^Cp ^Py A>y
A (4.32) A, = c,Cp + 4,Py = (1 " #Cp )A* + (1 ~ #Py )4>y
The ratio of the anodic and cathodic areas during galvanic interaction can then be expressed by:
Aa _ eCp^cp+QpyA>y (4.33)
K (i-eCp)A;p+(i-epy)A>y
The FeS2:CuFeS2 mass ratio can be expressed by:
rTJpy Ppy^Py^Py (4.34)
m cP PcpA^Cp
Results and Discussion 107 where pPy is the density of pyrite, pCp is the density of chalcopyrite, dPy is the mean
diameter of pyrite particles and dCp is the mean diameter of chalcopyrite particles. The pyrite surface can then be expressed by the following expression:
Pcp^Cp mPy (4.35) mCp Ppydpy
Substitution of Eqn. (4.35) into (4.33) gives:
A, _ 0 + £0p CP y (4.36)
Ac (1-0Cp) + )8(1-0py)
The effect of the FeS2:CuFeS2 mass ratio on the galvanic interaction can be quantified using Equation (4.36). The 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, Eq.
(4.36) gives:
(4.37)
Ac (1-0CP) + J3
The correlation between the surface areas and the amount of pyrite added during galvanic interaction, represented by Eq. (4.37), can be conveniently followed by plotting
AaIAc vs mpy/mcp. The values of pPy, pCp, dPy and dCp 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 ratio with increasing the amount of pyrite. The curve shows a drop of more than 75% on
the surface ratio for a FeS2:CuFeS2 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.
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 65°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. As 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 Santa Santa Santa Chalcopyrite source Eulalia Eulalia 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 cm-3) 4.2 4.2 4.2
Density of pyrite (g cm-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). The
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, 65°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: -75 + 38 pm. Leaching experiments were carried out at 65°C with a
Fe(lll):Fe(ll) ratio of 1. A FeS2:CuFeS2 ratio of 2:1 was used for the tests. The 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, Canada 24.05 32.40 31.20
Selwyn, Australia 23.60 28.10 28.70
As 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. The Selwyn sample achieved total
conversion in less than 24 hours in the presence of pyrite. These 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. The 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:CuFeS2,
Fe(lll):Fe(ll) = 15, pH1.4, 65°C
The results of the chemical leaching tests have confirmed the following conclusions from the electrochemical studies:
> The 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:
2+ 2+ Anodic: CuFeS2 -> Cu + Fe + 2 S° + 4e" (4.1)
2+ Cathodic: 4 Fe3+ + 4 e" = 4 Fe (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 (4.38) RT
( (1-geCp)FE 2+ 3+ _ /c [Fe ]exp -/cc,Cp[Fe ]exp (4-39) 'c,Cp '"A.Cp cCp RT Vcp RT
where /aCp, /cCp, 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.,
(4.40) 'a,Cp ^c,Cp
Results and Discussion 114 Substitution of Eqns. (4.38) and (4.39) into Eqn. (4.40) gives the following expression for the mixed potential:
4FA k eX 'a,Cp a,CP P RT (4.41)
(1-ore,Cp)FE 2+ 3+ /c , [Fe ]exp -/cc.CD[Fe ]exp -FA, 'c,Cp V,Cc pD V,Cp RT RT
It should be noted that, in the region of the mixed potential (Em), 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 (Fe3+ in this case).
Assuming transfer coefficients craCp and acCp equal to 0.5, then:
3+ FEm Fe ex = A,c ^c,cp[Fe ]exp (4.42) l) P P (4A,.cP*a.cp + A:,cpW [2RT) 2RT
This expression can be simplified to obtain the following expression for the mixed
potential:
£ ,cA,cpFe3+] c (4.43) Em = In
F 4/\aCp/caCp+/ccCp/\cCp[Fe ]
Substitution of Eqn. (4.43) into Eqn. (4.38) gives the following expression for the anodic
current:
Fe3+ / -AFk A I Aj.cpftc.cp[ ] (4.44) aCp a Cp aCPl, 2+ ' 4^Cp^Cp+/\c.cp^,cp[Fe ]
The dissolution rate of chalcopyrite, - rCp, 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 nCp 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 nCp0 dt
where XCp and nCp0 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. The 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 = 1- (4 47) P R) -
2 2 2 3 r = R (1-XCp) ' (4.48)
2 2 2/3 /lCp=47rr =4rrR (1-XCp) (4.49)
where R is the initial particle size.
Substitution of Eqns. (4.45), (4.49), (4.26) and (4.27) in Eqn. (4.46) gives the following
equation:
Results and Discussion 116 ecB,c (1 x ] (450) ^T"^7 l4eCp/<.,Cp+(i-eCp)^p[Fe-] "
Fe or f^iHgltyJ 0-^>' "l , (1-Xcp)- (4.51, 2 1 c Cp 0p eft ncp0 pCpKCp+(1-e0p)[Fe *] >'
where k^k^j-f*- and KCp=-^ ^c,Cp ^c.Cp
Integration of equation (4.51) will give the following expression for the conversion of chalcopyrite as a function of time:
2 3+ 4TTR (1-ecp)[Fe ] (1_x )1/3=1_^e , y ^'V'2t (4-52) CP C 2+ 3nCp,o lV 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, 2+ 2+ _ CuFeS2 -» Cu + Fe + 2 S° + 4e (4.1) 2+ 2 + FeS2 + 8a H20 -> Fe + 2a SO " + (2 - 2a) S° +16tr H + (2 +12a) e" (4.53) and the cathodic reduction of ferric on the surface of both minerals: Fe3+ + e~ -> Fe2+ (on chalcopyrite) (4.11a) Fe3+ + e~ -> Fe2+ (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 + 1/2 XPyFE^ /aPy=(2 + 12a)F^Py/caPy[H ]- exp (4.54) RT 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: fa^FE' 2+ 'c.Cp ( (1-ae,Cp)FE - 3+ (4.39) 'c,Cp '"Aj.Cp kc.cD[Fe ]exp - WFe ]exp *c,Cp RT (1-gcPy)FE 2+ J+ (4.55) ^c,Py "~ FA /ccPy[Fe ]exp -Sg— -/ccPy[Fe ]exp C Py RT 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 Eqn. (4.56) gives the following expression for the system: ( (1-ac,Cp)FE, ^ 3+ 'cCp^' '-m 2+ /c [Fe ]exp V:,Cp WFe ]exp cCp (4-57) RT RT fa^FE,*c,Py' *-m ^ f (1-ac,Py)FE, ^ + /c [Fe3+]exp + A /ccPy[Fe' ]exp cPy 'c,Py RT RT If the approximation is made that araCp = craPy = orcCp = acPy =0.5, this expression can be simplified to obtain the following expression: RT^ (^,CpA ,c +^.Py4,Py)FeJ+] E c P (4-58) m n 5 2+ F 4/\a,Cp/caCp + (2 +12a)A,,Py/ca,Py[HT° + (WW + WW)[Fe ] Results and Discussion 118 Substitution of Eqn. (4.58) into Eqn. (4.38) gives the following expression for the anodic current: + 1 + Fe / =4Fk A (^c,Cp^c,CP ^c,PyA,Py)[ 3 + 05 2+ a,cP °^l4AaCph:aCp +(2 + 12a)/\a,Py/ca,Py[H ]- +(/cc,CpA,cP +^,Py4.Py)Fe ] (4.59) The dissolution rate of chalcopyrite, -rCp, may be expressed by: dnCp /a,Cp cp dt 4F (4.60) _ h, A I (^c.Cp^c.Cp+^cPyA.Py)^ ] + 5 2+ a,CP °^\4AaCpkaCp +(2 + 120-)\Py/Ca,Py[H ]-°- +(/Cc,cp4,Cp +^,PyA,Py)Fe ] 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-JLek A - c'cp'Vcp^cp UL "Cp,0 3+ [(1-0Cp)>4Cp/ceCp +(1-9Py)APy/cePy][Fe ] Cp^a.Cp + 5 2+ + (2 +12a)0PyA3y/ca,Py [H ]-° + [(1 - 6Cp )ACpkc,Cp + (1 - 0Py )APykcPy][Fe ] (4.61) Substitution df Eqn. (4.35) in Eqn. (4.61) gives the following expression: dX cP _ JL0 n A dt n ^Cp^a.Cp^Cp cP,o (4.62) 3+ [(1 - Sep )*c,cp + (1-6Py)/3/cc,Py][Fe ] + 05 4GCpkaCp +(2+ ^2o)9PvBkaPv[H ]- ^ Substitution of (4.49) in Eqn. (4.62) gives the following equation: Results and Discussion 119 2 dXCp _ ATTR r 2/3 —IT - — yCpKa,Cp( 1 - *Cp ) UL "Cp,0 (4.63) [(1-0Cp)^,Cp+(1-0py)^c,Py][Fe3+] l + 05 2+ 49CpkaCp +(2 + 12a)0PyJ8/ca,Py[H r + [(1 - 0Cp )^.cP +(1-0py)/3/cc,Py][Fe ] It is known that the rest potential of chalcopyrite is lower than that of pyrite. When contact is made, the following 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 follows: 2 ,3+3+1 dXCp 4TTR q r [(1-Qc )^,cp+^^,Py][Fe ] I P _ \2/2 3 (4.64) 1_ A y K —77- = CD a.CDJ r 77.—77~r Tr 7771 oITV2+ cP/ dt "cP,o 40cp/ca,cp + [(1 -6Cp)k,Cp + /3/cc,Py][Fe ] 3 2 dXCp 4nR a ,• I [(1-eCp)+ ^c][Fe-] or —— = 0, 2+ (1 Xcp) (4 65) CpC l0cpKCp+[(1-aCp) ^][Fe ) ' dt nCp,o + c k k where 5>. =-^H. and Integration of Eqn. (4.65) gives the following expression for the conversion of chalcopyrite as a function of time during galvanic interaction with pyrite: Results and Discussion 120 2 3+ 477-R [(1-0Cp) + i8^][Fe ] 1/3 (4.66) (1-*Cp) =1- 2+ 3«Cp,0 Fcp^cP+[(1-0cp) + ^c][Fe ] 4.2.6. Electrochemical reaction and surface passivation model As 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]. > The 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: 1 (4.67) 4,CP=( -^CAP Based on the the first assumption above, the growth rate of the passive layer can be expressed by: (4.68) or / = 1-exp(-y) (4.69) (4.70) and AaCf> = exp(-/cpi)ecpy4Cp from which Eqn (4.51) may be rewritten as follows: Results and Discussion 121 c/XCp _ AnR^g k ,k , (1-eCp)[Fe"] _ 2,3 (4 71. 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: 2 4rrRK 2kc 1 3 1 0 3+ d"XCp) ' = 1 -T TK V( - cP)[Fe ] J K A "cP.o p cP (4.72) 2+ 2+ x (V^cp^cp + (1 - 0CP )Fe ] - VeCpKcp exp(-V) + (1 - QCp )[Fe ]) /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 Eqn. (4.70) may be applied in a similar fashion to Eqn. (4.65) with the following result: w C P< (1 Xc ) (4J3) * " "cp, ° "'"\|exp(-fcp()eI:pKCp + [(1-ecp) + WJ[Fe-] - ' 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: R22k 3 (1 _^r=u f * V[(1-eCp) + ^e][Fe -] ^"Cp.O Kp^Cp *(VecpKcP + [(i^ (4-74) 4.2.7. Validation of the electrochemical passivation model The model was validated at 65°C because the extraction curves did not show an induction period in the early stage of leaching. The 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 1/3 representing the copper conversion (1-(1-XcP) ) 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 kp) occurred at a Fe(lll):Fe(ll) ratio of 0.33. Severe passivation was observed at high Fe(lll):Fe(ll) ratios. The 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. As 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 25 Time (H) Figure 4.29 Electrochemical-passivation model for the leaching of chalcopyrite in ferric sulphate solutions, pH 1.4, 65°C Results and Discussion 123 The 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. As 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). The 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 07 - 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 FeS2:CuFeS2 ratios, Fe(lll):Fe(ll) = 1, pH 1.4, 65°C Results and Discussion 124 H3-Py/Cp = 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, 65°C Results and Discussion 125 Time (H) Figure 4.32 Predicted values for the effect of particle size on the leaching of chalcopyrite, FeS2:CuFeS2 = 2, Fe(lll):Fe(ll) = 1, pH 1.4, 65°C Results and Discussion 126 Figure 4.33 Predicted values for the effect of increased FeS2:CuFeS2 ratio and particle size on the leaching of chalcopyrite, Fe(lll):Fe(ll) = 1, pH 1.4, 65°C, dPy = 25 pm, dcp = 100 pm Results and Discussion 127 Table 4.9. Input parameters for the electrochemical-passivation model Fe(lll):Fe(ll) ratio 0.1 0.33 1 5 15 1 1 Physical Parameters 30 30 30 30 30 20 10 mCP (g) 0 0 0 0 10 20 mPy (g) 0 -3 4.2 4.2 4.2 4.2 4.2 4.2 4.2 PcP (9 cm ) -3 5.0 5.0 5.0 5.0 5.0 5.0 5.0 Ppy (g cm ) 47.6 47.6 47.6 47.6 47.6 47.6 47.6 dcP (um) dpy(um) 45.5 45.5 45.5 45.5 45.5 45.5 45.5 )8 1.8 3.5 7.1 0 0 0.4 1.8 Electrochemical Parameters [Fe(lll)] (M) 0.008 0.022 g 0.045 0.075 0.084 0.045 0.045 [Fe(ll)] (M) 0.081 0.067 0.045 0.015 0.006 0.045 0.045 [H+] (M) 0.180 0.180 0.180 0.180 0.180 0.180 0.180. 5 kcp 1.12*10"5 1.12x10"5 1.12x10"5 1.12x10"5 1.12*10"5 1.12*10~5 1.12x10" Kcp 0.001 0.001 0.001 0.001 0.001 0.001 0.001 0.328 0.167 0.248 0.577 0.605 0.248 0.248 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 30 20 15 10 6 4.3 3.3 mCp (g) 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 3 pPy (g cm" ) 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 45.5 45.5 45.5 45.5 45.5 dpy (Mm) 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 5 5 1.12x10"5 -5 1.12*10~5 1.12x10-5 1.12x10"5 kcP 1.12*10" 1.12*10" 1.12x10 Kcp 0.001 0.001 0.001 0.001 0.001 0.001 0.001 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 mPy (g) 20 20 20 3 4.2 PcP (9 cm" ) 4.2 4.2 -3 Ppy (g cm ) 5.0 5.0 5.0 47.6 25 100 dcP (Mm) 24 dpy (Mm) 45.5 100 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 1.12x10"5 5 1.12x10"5 kCp 1.12*10" Kcp 0.001 0.001 0.001 Vc 1.452 1.452 1.452 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 28°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. As shown in Figure 4.35, iron seemed to leach more rapidly than copper. A Cu:Fe 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 28°C 0.7 3 o O « o S 10 15 20 Time (Days) Figure 4.35 Molar ratio Cu:Fe during bioleaching with mesophiles at 28°C. Results and Discussion 132 700 2.5 -solution potential, mesophiles -Bacteria counts, mesophiles 2 72 E « o 1.5 X c 3 1 O 1 o TO O 0.5 m 350 300 10 15 20 25 30 35 Time (days) Figure 4.36 Solution potential evolution and microbial counts during bioleaching with mesophiles at 28°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 mV (Ag/AgCI). The increase in the redox potential is due to the bacterial oxidation of ferrous to ferric. The molar ratio Cu:Fe 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. The 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 45°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 Cu:Fe 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 Cu:Fe initially, as chalcopyrite leach kinetics did not seem to be affected. Solution potentials increased from 410 to 520 mV (Ag/AgCI) from the beginning of the test to day 15 and remained at 520 mV until the end of the test (Figure 4.39). It is interesting to note that chemical leaching tests at 45°C indicated that chalcopyrite leaching was much faster at a redox potential of about 540 mV (Fe(lll):Fe(ll) ratio of 5). The 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 45°C. o S 0.3 - 0.2 - 0.1 - 0 1 . 1 1 I 1 1 J0 5 10 15 20 25 30 35 Time (days) Figure 4.38 Molar ratio Cu:Fe during bioleaching with moderate thermophiles at 45°C. Results and Discussion 135 Figure 4.39 Solution potential evolution and microbial counts during bioleaching with moderate thermophiles at 45°C. 4.3.3. Extreme thermophiles Bioleaching tests with extreme thermophilic microorganisms were performed at 68°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. The Cu:Fe 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 65°C. Solutions potentials were very low (350 mV (Ag/AgCI)) during the first leaching stage and increased up to 450 mV 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. To further confirm our conclusions, two more tests were conducted with relatively high concentrations of ferric at the beginning of the tests. As 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 65°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 68°C. Results and Discussion 137 1.4 Extreme thermophiles, 68°C 1.2 1 u. "3 o 0.8 (0 CC 0.6 JS o 0.4 0.2 0 10 15 20 25 30 35 Time (Days) Figure 4.41 Molar ratio Cu:Fe during bioleaching with extreme thermophiles at 68°C. 600 2.5 Solution potential, extreme thermophiles Bacteria counts, extreme thermophiles E a> o + 1.5 w ra u ° 350 0.5 S 10 15 20 30 35 Time (days) Figure 4.42 Solution potential evolution and microbial counts during bioleaching with extreme thermophiles at 68°C. Results and Discussion 138 Time (h) Figure 4.43 Fraction of copper leached with extreme thermophiles at 68°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). The 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). The 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: MS + 2 Fe3+ -» Me2+ + 2 Fe2+ + S° (4.98) and ferrous consumption to generate ferric and in some cases the oxidation of elemental sulphur: 2+ + 3+ 4 Fe + 4 H + 02 + bacteria 4Fe + H20 (4.99) 2S° + 3 02 + 2 H20 -> 2 H2S04 (4.100) 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]. 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 [121]. 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. The 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. The mechanism that explains our results is shown in Figure 4.46. Cu2+ Fe Chalcopyrite Chalcopyrite Anodic area Cathodic area j2+ + Fe2+ + 2 S° + 4 e~ Fe3+ + e" = Fe2+ 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. The 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. We suggest that thermophiles accelerate ferric reduction on chalcopyrite by providing alternative cathodic sites. To 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. The study was conducted in 4 small columns, all placed in a common heated water bath at 68°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). The 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. The 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 mV 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 mV 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. As 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. The 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. An 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. On 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 80°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, 68°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 mV (Ag/AgCI). Air was supplied as oxidizing agent. The 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. The 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. The 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, 68°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 FeS2:CuFeS2 of 4:1, the potential dropped rapidly to 415 mV for the first 2 hours and increased to 470 mV 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. As mentioned in the previous sections, the beneficial effect of pyrite stops once it starts to dissolve. For a ratio FeS2:CuFeS2 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, 80°C 4.5.2. Effect of initial solution potential The effect of solution potential was investigated by conducting leaching tests at potentials much lower than 460 mV 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). The leaching rate was initially rapid, and then started to decrease after 3 hours of leaching. Cu leaching ceased almost entirely after 16 hours of leaching. Results and Discussion 148 560 -o- Pure chalco 540 -»-Chalco=10-Pyr=20 O -o- Chalco=10-Pyr=40 5 520 "5I> < 500 ~T\ (A > -* 0 > 480 | 460 f 440 (A LU 420 400 10 15 20 25 Time (h) Figure 4.50 Solution potentials during Leaching of chalcopyrite at different ratios of FeS2:CuFeS2, Initial Fe(lll):Fe(ll) = 1, 80°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. The 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. These parameters include particle size, pulp density, impeller speed, the type of primary oxidant (air or oxygen) and acidity of the solution. Results and Discussion 149 100 -Cl-Chalco=10, High Redox 90 -•-Chalco-10, Low Redox -O-Chalco=10, Pyr=40, 80 — High Redox -•-Chalco=10, Pyr=40, Low 70 Redox 60 > m c 50 o u k_ 40 01 a. a. o 30 o 20 —• 10 - 0 10 15 20 25 Time (H) Figure 4.51 Effect of initial potential on the leaching rate curves of chalcopyrite at different ratios of FeS2:CuFeS2, 80°C 4.5.3. Effect of particle size The 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 -75 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 90°C [38]. The rate increased substantially with finer particles only in ferric chloride medium. As 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, 80°C 60 -•- <38 microns =Q 50 -o- 38_75 microns c 40 o 'I co £ 30 o u i_ 0) §: 20 o o 10 10 15 20 25 Time (h) 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 As shown in Figures 4.55 and 4.56, when pyrite was added to the leaching system, various behaviors were observed. The 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 mV in some tests. As 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. Once this "critical" potential is reached, the chalcopyrite begins to exhibit "passive" behavior. As mentioned in the previous section, this "critical" potential was about 460 mV (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. The tests using Huanzala pyrite are shown with open symbols, and those using Park City pyrite are shown with filled symbols. As 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. The 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. The 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, FeS2:CuFeS2 = 4, 80°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, FeS2:CuFeS2 = 4, 80°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. As shown in Figure 4.59, the test run with more pyrite and chalcopyrite achieved nearly 80% copper recovery in 24 hours. The 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. These 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). The 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 158 500 480 460 u 440 D) < 420 > > 400 E 380 -•-C30-P60, 750 rpm, air -•-C30-P60, 1200 rpm, air 360 -A-C30-P60, 1200 rpm, 02 340 -»-C30-P60,1200 rpm,02, High acidity 320 -*-C30-P60,1200 rpm, 02,High Pulp, High acidity 300 10 15 20 25 30 35 Time (h) 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. To 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. Why 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 10 30 30 30 30 30 30 10 mCp (g) 10 10 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 1 [H2S04] (g L" ) 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) > The 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. On 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. The 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:CuFeS2 ratios (>2) because pyrite provides more cathodic area for the reduction reaction, and the anodic dissolution rate of chalcopyrite must increase to compensate. The 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. An electrochemical model which takes into account the passivation of chalcopyrite was proposed. This model was also validated during galvanic interaction with pyrite. The addition of microorganisms was beneficial at high temperature because extreme thermophiles generate a fairly low potential environment (350-500 mV (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 mV (Ag/AgCI)). These 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 80°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. The potential benefit of using low-grade chalcopyrite containing various amounts of pyrite should be investigated further. 2. The leaching of chalcopyrite in this study is surprisingly enhanced in the presence of higher acid concentration. 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A summary of copper extractions after 24 hours of leaching is presented in Table A.1. Mean: * = Zx//n Variance: s2 =£(X, -X)2 n-\ ;=1 / Std deviation: s = Vs"2 Table A. 1. Statistical analysis of copper extraction after 24 hours of controlled leaching with peroxide Temp (°C) Fe(lll):Fe(ll) Cu (%) X s2 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