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"THE SURFACE CHEMISTRY OF CHALCOPYRITE DURING ELECTROCHEMICAL DISSOLUTION  by  Ahmad Ghahremaninezhad M.Sc., Sharif University of Technology, 2007 B.Sc., Sharif University of Technology, 2005  A THESIS SUBMITTED IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREE OF  DOCTOR OF PHILOSOPHY in The Faculty of Graduate Studies (Materials Engineering)  THE UNIVERSITY OF BRITISH COLUMBIA (Vancouver) July 2012  \u00A9 Ahmad Ghahremaninezhad, 2012  i  \u000CAbstract Hydrometallurgy may be an alternative to the currently practiced smelting process for copper  extraction  from  chalcopyrite  (CuFeS2).  However,  the  low  temperature  hydrometallurgical processes for chalcopyrite continue to face challenges, mostly relating to their slow dissolution rates or high sulfuric acid production. The slow dissolution rate of the mineral is strongly linked to the formation of the passive film on its surface. However, despite 40 years of research on this topic, there is still not a complete agreement between researchers about the composition and stability of chalcopyrite\u2019s passive film in sulfuric acid solutions. In this work, the nature of chalcopyrite\u2019s passive film and its stability were studied by application of a variety of electrochemical techniques. Additionally, the electrochemical results of the chalcopyrite study were compared to those obtained for a pyrrhotite electrode (Fe1-xS), as pyrrhotite electrochemistry represents a simplified case of the chalcopyrite system. X-ray photoelectron spectroscopy (XPS) was used to analyze the composition of the product layers formed on the surface. It is shown that the chalcopyrite electrode is passive for potentials up to 0.90 V SHE. Above this potential, transpassive dissolution occurs. Results of XPS studies have suggested that a metal-deficient sulfide film (Cu1-xFe1-yS2-z) is the most plausible copper and iron containing sulfide phase which passivates the surface of chalcopyrite. In addition, an outer layer of iron oxyhydroxide (FeOOH) forms on the passive film. FeOOH forms via oxidation of the passive film\u2019s ferrous sulfide phases. The thickness of the sulfide passive film was calculated to be approximately 6.7 nm. It is demonstrated that the transpassive dissolution of chalcopyrite is significantly linked to oxidation of sulfur (from sulfide in the passive film to elemental sulfur and maybe sulfur species with higher oxidation states, e.g. thiosulfate). No elemental sulfur or polysulfide species were detected on the surface for potentials below 0.90 VSHE.  ii  \u000CPreface A version of chapter 5 has been published: Ghahremaninezhad, A., Asselin, E., Dixon, D.G., 2010, Journal of Electrochemical Society, 157, C248\u2013C257. Chapter 6 has also been published: Ghahremaninezhad, A., Asselin, E., Dixon, D.G., 2010, Electrochimica Acta, 55, 5041\u20135056. I have conducted all the testing and written both of the manuscripts. Chapter 9 has been accepted for publication: Ghahremaninezhad, A., Dixon, D.G., Asselin, E., 2012, Hydrometallurgy, 125\u2013126, 42\u201349. I have conducted all of the experiments for these publications. I also wrote the manuscripts.  iii  \u000CTable of Contents ABSTRACT................................................................................................................................................... ii PREFACE .................................................................................................................................................... iii TABLE OF CONTENTS ............................................................................................................................. iv LIST OF TABLES ...................................................................................................................................... vii LIST OF FIGURES.................................................................................................................................... viii ACKNOWLEDGEMENTS........................................................................................................................ xiv DEDICATION............................................................................................................................................. xv 1  INTRODUCTION ................................................................................................................................ 1  2  LITERATURE REVIEW .................................................................................................................... 5 2.1  GENERAL ASPECTS ............................................................................................................................. 5  2.1.1  Chalcopyrite crystal structure ................................................................................................... 5  2.1.2  Semiconductive properties of chalcopyrite ................................................................................ 6  2.2  DISSOLUTION OF SEMICONDUCTOR MATERIALS ................................................................................... 7  2.2.1  Semiconductor-electrolyte interface ........................................................................................ 11  2.2.2  Mott-Schottky equation ........................................................................................................... 13  2.2.3  Current-potential curves for semiconductors ........................................................................... 15  2.3  CHALCOPYRITE DISSOLUTION AND ELECTROCHEMISTRY .................................................................... 18  2.3.1  Polarization studies ................................................................................................................ 18  2.3.2  Chronoamperometry studies ................................................................................................... 21  2.3.3  Electrochemical impedance spectroscopy (EIS) studies ........................................................... 23  2.3.4  Cyclic voltammetry (CV) studies ............................................................................................. 27  2.4  SURFACE ANALYSIS OF CHALCOPYRITE ............................................................................................. 33  2.5  ELECTROCHEMICAL (AND CHEMICAL) DISSOLUTION REACTIONS ......................................................... 37  2.6  WHAT IS MISSING ? ........................................................................................................................... 40  3  OBJECTIVES .................................................................................................................................... 42  4  EXPERIMENTAL ............................................................................................................................. 43 4.1  MINERALS USED ............................................................................................................................... 43  4.2  ELECTRODE PREPARATION ................................................................................................................ 43  4.3  SOLUTION PREPARATION .................................................................................................................. 44  4.4  ELECTROCHEMICAL EXPERIMENTS .................................................................................................... 44  4.5  X-RAY PHOTOELECTRON EXPERIMENTS ............................................................................................. 46  iv  \u000C4.6 5  SCANNING ELECTRON MICROSCOPE STUDY ........................................................................................ 47  ELECTROCHEMICAL DISSOLUTION OF PYRRHOTITE IN HYDROCHLORIC ACID  SOLUTION ................................................................................................................................................. 48 5.1  INTRODUCTION ................................................................................................................................ 48  5.2  RESULTS AND DISCUSSION ................................................................................................................ 51  5.2.1  Open circuit potential study .................................................................................................... 51  5.2.2  Potentiodynamic study ............................................................................................................ 51  5.2.3  Electrochemical impedance spectroscopy study....................................................................... 52  5.2.4  Mott-Schottky study ................................................................................................................ 56  5.2.5  Model development ................................................................................................................. 61  5.2.6  SEM studies............................................................................................................................ 74  5.3 6  CONCLUSION ....................................................................................................................................... I  ELECTROCHEMICAL EVALUATION OF CHALCOPYRITE\u2019S SURFACE DURING  DISSOLUTION IN 0.5 M H2SO4 SOLUTION ........................................................................................... 77 6.1  INTRODUCTION ................................................................................................................................ 77  6.2  RESULTS AND DISCUSSION ................................................................................................................ 77  6.2.1  Characterization of used chalcopyrite electrodes .................................................................... 77  6.2.2  Open circuit potential study .................................................................................................... 79  6.2.3  Potentiodynamic study ............................................................................................................ 79  6.2.4  Electrochemical impedance spectroscopy study....................................................................... 81  6.2.5  Mott-Schottky study ................................................................................................................ 93  6.2.6  Model development ................................................................................................................. 96  6.3 7  CONCLUSION.................................................................................................................................. 107  X-RAY PHOTOELECTRON SPECTROSCOPY ANALYSIS OF CHALCOPYRITE\u2019S SURFACE  IN SULFURIC ACID SOLUTION ........................................................................................................... 109 7.1  INTRODUCTION .............................................................................................................................. 109  7.2  RESULTS AND DISCUSSION .............................................................................................................. 109  7.2.1  Passivation of chalcopyrite; a brief introduction ................................................................... 109  7.2.2  XPS study of chalcopyrite (survey spectra)............................................................................ 114  7.2.3  Cu2p peaks in chalcopyrite ................................................................................................... 115  7.2.4  Fe2p and Fe3p peaks of chalcopyrite .................................................................................... 117  7.2.5  S2p and S2s peaks of chalcopyrite......................................................................................... 118  7.2.6  XPS study of potentiostatically polarized chalcopyrite ........................................................... 122  7.3  CONCLUSION .................................................................................................................................. 145  v  \u000C8  MULTIPLE CYCLIC VOLTAMMETRY STUDY OF CHALCOPYRITE IN 0.5 M SULFURIC  ACID SOLUTION..................................................................................................................................... 147 8.1  INTRODUCTION .............................................................................................................................. 147  8.2  RESULTS AND DISCUSSION .............................................................................................................. 147  8.2.1  Effect of oxygen content of the solution on the CV experiments .............................................. 147  8.2.2  Cathodic reactions on the surface of chalcopyrite ................................................................. 148  8.2.3  Anodic reactions on the surface of chalcopyrite..................................................................... 157  8.3 9  CONCLUSION .................................................................................................................................. 164  FERRIC/FERROUS REDOX REACTION ON ANODICALLY PASSIVATED CHALCOPYRITE  (CUFES2) ELECTRODE .......................................................................................................................... 166  10  11  9.1  INTRODUCTION .............................................................................................................................. 166  9.2  EXPERIMENTAL .............................................................................................................................. 168  9.3  RESULTS AND DISCUSSION .............................................................................................................. 170  9.4  CONCLUSIONS ................................................................................................................................ 185  A PROPOSED ELECTROCHEMICAL DISSOLUTION MODEL............................................... 186 10.1  INTRODUCTION .......................................................................................................................... 186  10.2  MODEL DEVELOPMENT .............................................................................................................. 186  10.3  CONCLUSION ............................................................................................................................. 193  CONCLUSIONS AND RECOMMENDATIONS............................................................................ 194 11.1  CONCLUSIONS ........................................................................................................................... 194  11.2  RECOMMENDATIONS .................................................................................................................. 196  REFERENCES .......................................................................................................................................... 198  vi  \u000CList of Tables Table 2-1: Materials properties of CuFeS2 at room temperature [Hiskey, 1993]. .............................................. 7 Table 2-2: Various sulfur, copper, and iron species on the surface of the chalcopyrite after oxidative dissolution; between 50 to 80 \u00B0C [Parker et al., 2003]. .................................................................................. 36 Table 2-3: Atomic percentage of different sulfur species on the surface of chalcopyrite before and after leaching [Harmer et al., 2006]. ..................................................................................................................... 37 Table 4-1: Composition of the as received sulfide minerals used in this research. ........................................... 43 Table 5-1: Flatband potentials for pyrrhotite in 1 M HCl electrolyte in dark and under irradiation. ............... 60 Table 5-2: Model parameters for equivalent circuits of Figure 5-6. ................................................................ 66 Table 6-1: Composition of the as received chalcopyrite ore............................................................................ 78 Table 6-2: Charge carrier density and effective Debye length for chalcopyrite and its surface layers at 0.5 M H2SO4 electrolyte. ......................................................................................................................................... 96 Table 6-3: Model parameters for equivalent circuit of Figure 6-12a (presented in next page). ........................ 97 Table 6-4: Model parameters for equivalent circuit of Figure 6-15a. ............................................................ 101 Table 7-1: Binding energy (eV) of Cu, Fe and S elements for chalcopyrite after Ar+ ion etching. .................. 122 Table 7-2: Binding energy and FWHM values for XPS spectra of S2p peaks................................................. 132 Table 8-1: Reactions associated with cathodic peaks of chalcopyrite dissolution. ......................................... 151 Table 8-2: Anodic reactions on chalcopyrite surface. ................................................................................... 159 Table 9-1: Model parameters for equivalent circuit of Fig. 9-7..................................................................... 179  vii  \u000CList of Figures Figure 2-1: Unit-cell crystal structure of CuFeS2 [C\u00F3rdoba et al., 2008].......................................................... 6 Figure 2-2: Schematic energy band diagram for chalcopyrite CuFeS2 [Torma, 1991]. ..................................... 7 Figure 2-3: Band model in solids [Bard and Faulkner, 2001]. ......................................................................... 8 Figure 2-4: Schematic diagram of the energy levels of a semiconductor. .......................................................... 9 Figure 2-5: Schematic diagram of the energy levels of (a) an n-type semiconductor and (b) a p-type semiconductor. .............................................................................................................................................. 10 Figure 2-6: Potential (a) and charge (b) distribution at the semiconductor(left side)/electrolyte(right side) interface. Note that a high ion concentration for solution is assumed hence the diffuse layer on the solution side is neglected [Memming, 2001]. ..................................................................................................................... 11 Figure 2-7: Schematic current\u2013potential curve for n-type and p-type semiconductors. .................................... 16 Figure 2-8: Anodic polarization curves for CuFeS2 from 6 different locations in 1 M H2SO4, 30 mV min-1, 25\u00B0C [Warren et al., 1982]. .................................................................................................................................... 19 Figure 2-9: Potentiostatic anodic polarizations of chalcopyrite, 10 mV hr-1 [Viramontes-Gamboa et al., 2007]. ..................................................................................................................................................................... 20 Figure 2-10: Anodic oxidative behavior of chalcopyrite microparticles at 20, 45 and 60\u00BAC and 90 g/L of H2SO4 [Vasquez et al., 2012].................................................................................................................................... 21 Figure 2-11: Current-time curves for a chalcopyrite electrode at the potentials (vs. SHE) indicated in (a) 1 M H2SO4, (b) 1 M HC1 [Biegler and Swift, 1979]. ............................................................................................. 22 Figure 2-12: Current\u2013time curves representing the re-oxidation of chalcopyrite at 800 mV vs. SHE [Lu et al., 2000]. ........................................................................................................................................................... 23 Figure 2-13: Impedance plots for the chalcopyrite electrode; (a) with or without the ferrous and the cupric ion, (b) coexisting of ferrous and the cupric ion, at 0.65 V vs. SHE [Hiroyoshi et al., 2004]................................... 25 Figure 2-14: Measured and simulated Nyquist plots at; (a) \u2013 0.15 V, (b) 0.2 V, (c) 0.45 V and (d) 0.6 V vs. SCE [Velasquez et al., 2005]. ................................................................................................................................ 27 Figure 2-15: Equivalent circuits for; (a) \u2013 0.15 V, (b) 0.2 V, (c) 0.45 V and 0.6 V vs. SCE [Velasquez et al., 2005]. ........................................................................................................................................................... 27 Figure 2-16: First positive going current\u2013voltage sweep on a freshly polished chalcopyrite surface [Biegler and Swift, 1979]. ........................................................................................................................................... 29 Figure 2-17: Anodic stripping curves for a platinum ring held at \u2013 0.06 V after dissolution of a CuFeS2 disk which was held at potentials of (a) 0.45, (b) 0.49, (c) 0.54, (d) 0.59 and (e) 0.65 V vs. SHE for 1 h at 200 rpm [Lazaro and Nicol, 2006]. ............................................................................................................................. 32  viii  \u000CFigure 5-1: Open circuit potential (OCP) of the three different pyrrhotite electrodes vs. time; two of the curves show the OCP of same samples after polishing with carbide paper................................................................. 51 Figure 5-2: Potentiodynamic curves for the pyrrhotite electrode in 1 M HCl electrolyte. ................................ 52 Figure 5-3: Impedance spectra of a pyrrhotite electrode in 1 M HCl at (a) OCP, (b) 692, (c) 782, (d) 917, (e) 967, (f) 1042, (g) 1092 mV vs. SHE. Parameter is frequency in Hz. ................................................................ 54 Figure 5-4: Mott-Schottky plots for pyrrhotite in 1 M HCl electrolyte in the dark and under irradiation. ........ 57 Figure 5-5: Mott-Schottky plots for pyrrhotite electrode in 1M HCl electrolyte, (a) under irradiation and, (b) in dark. ............................................................................................................................................................. 59 Figure 5-6: Equivalent electrochemical circuit for the Pyrrhotite/electrolyte interface at, (a) low anodic potentials (OC \u2013 692 mV), (b) 782 mV, (c) high anodic potentials (917 \u2013 1042 mV). ....................................... 61 Figure 5-7: Measured and calculated Nyquist plots for pyrrhotite electrodes at (a) OCP, (b) 692 mV, (c) 782 mV, (d) 917 mV, (e) 992 mV, and (f) 1042 mV vs. SHE. .................................................................................. 65 Figure 5-8: Simplified equivalent circuit model for the electrode.................................................................... 66 Figure 5-9: Tafel region polarization of the pyrrhotite electrode in 1 M HCl electrolyte. ................................ 68 Figure 5-10: Anodic current density-time transient in response to potential step from the OC potential to, (a) 540, and (b) 850 mV; inset of figure (a) shows the earlier stage of dissolution of the corresponding condition with higher resolution.................................................................................................................................... 72 Figure 5-11: Tafel region of ferrous to ferric oxidation reaction on different oxidized metals compared to the polarization curve of pyrrhotite electrode in HCl solution. ............................................................................. 74 Figure 5-12: SEM images of the pyrrhotite electrode after 2 hrs dissolution in 1 M HCl solution at (a) 782, (b) 1042 mVSHE. .................................................................................................................................................. 75 Figure 6-1: SEM images of (a) polished electrode surface, (b) siderite phase and EDX mapping for element distribution on the surface of prepared electrode. .......................................................................................... 78 Figure 6-2: Open circuit potential (OCP) of different chalcopyrite electrodes vs. time.................................... 79 Figure 6-3: Potentiodynamic curves for the chalcopyrite electrode in 0.5 M H2SO4 electrolyte at different scan rates.............................................................................................................................................................. 80 Figure 6-4: Electrochemical impedance spectra of chalcopyrite electrode for anodic potentials of OC to 300 mVMSE. .......................................................................................................................................................... 84 Figure 6-5: Bode diagrams of (a) phase angle, and (b) total impedance magnitude of the chalcopyrite electrode at potential range of OC\u2013150 mVMSE (OC to 750 mVSHE). ............................................................................... 85 Figure 6-6: Bode diagrams of total impedance magnitude of the chalcopyrite electrode at (a) 250, (b) 300 mVMSE. .......................................................................................................................................................... 86 Figure 6-7: Electrochemical impedance spectra of chalcopyrite electrode for anodic potentials of 325\u2013390 mVMSE (925\u2013990 mVSHE). ............................................................................................................................... 88  ix  \u000CFigure 6-8: Bode diagrams of (a) total impedance magnitude, and (b) phase angle of the electrode at potential range of 250\u2013390 mV. ................................................................................................................................... 90 Figure 6-9: Electrochemical impedance spectra of chalcopyrite electrode for anodic potentials of 0.50\u20131.10 VMSE (1.10 to 1.7 VSHE). .................................................................................................................................. 91 Figure 6-10: Bode diagram of total impedance magnitude for the chalcopyrite electrode at potential range of 390\u20131100 mVMSE (990 to 1700 mVSHE). .......................................................................................................... 92 Figure 6-11: Mott-Schottky plots for chalcopyrite electrode at different frequencies. ...................................... 94 Figure 6-12: (a) Equivalent electrochemical circuit for the Chalcopyrite/electrolyte interface at low anodic potentials; Measured and calculated Nyquist plots for chalcopyrite electrode at (b) OC potential, (c) -200 mV, (d) -100 mV, (e) 0.0 mV. ................................................................................................................................ 98 Figure 6-13: (a) Equivalent electrochemical circuit for the Chalcopyrite/electrolyte interface at 0.15 VMSE (0.75 VSHE); Measured and calculated (b) Complex-impedance plane representation, and (c) Bode representation of the magnitude of the impedance. .................................................................................................................... 99 Figure 6-14: (a) Equivalent electrochemical circuit for the Chalcopyrite/electrolyte interface at potential range of 0.2 to 0.25 VMSE (0.8 \u2013 0.85 VSHE); Measured and calculated (b) Nyquist representation at 0.2 VMSE, (c) Bode representation of the magnitude of the impedance at 0.2 VMSE, (d) Nyquist representation at 0.25 VMSE, (e) Bode representation of the magnitude of the impedance at 0.25 VMSE. ................................................................... 100 Figure 6-15: (a) Equivalent electrochemical circuit for the Chalcopyrite/electrolyte interface at potential range of 0.3 to 0.4 VMSE (0.9-1.0 VSHE); Measured and calculated Nyquist representations at (b) 325, (c) 350, and (d) 390 mVMSE. .................................................................................................................................................. 102 Figure 6-16: (a) Equivalent electrochemical circuit for the Chalcopyrite/electrolyte interface at potential of 0.5 VMSE (1.1 VSHE), (b) Measured and calculated Nyquist representation at 0.5 VMSE. ........................................ 103 Figure 6-17: (a) Equivalent electrochemical circuit for the Chalcopyrite/electrolyte interface at potential of 0.75 VMSE (1.35 VSHE), (b) Measured and calculated Nyquist representation at 0.75 VMSE. ............................. 105 Figure 6-18: (a) Equivalent electrochemical circuit for the Chalcopyrite/electrolyte interface at potential of 1.1 VMSE (1.7 VSHE), (b) Measured and calculated Nyquist representation at 1.1 VMSE, (c) magnification of Figure 618b.............................................................................................................................................................. 106 Figure 7-1: (a) Multicyclic voltammogram (10 cycles) of a chalcopyrite electrode in the potential window of 0.1 to 1 VSHE, (b) associated charge with the anodic peak of Figure 7-1a at 0.476 V at each cycle, (c) first cycle of Figure 7-1a. ............................................................................................................................................ 110 Figure 7-2: (a) Potentiodynamic (SR = 0.1 mV sec-1) and (b) potentiostatic polarization of chalcopyrite electrode in 0.5 M sulfuric acid. .................................................................................................................. 112 Figure 7-3: Potentiostatic polarization of chalcopyrite at (a) 0.5, 0.55 and 0.6 V for 6 hours, and (b) 800 mVSHE for 60 hours................................................................................................................................................. 114  x  \u000CFigure 7-4: The survey (full range) XPS spectrum of chalcopyrite (a) after polishing and 6 hours of air exposure, (b) after ion etching inside the XPS chamber. ............................................................................... 116 Figure 7-5: (a) XPS spectrum of Cu2p, and (b) copper Auger spectrum excited by X-rays, obtained after ion etching. ....................................................................................................................................................... 117 Figure 7-6: XPS spectra of (a) Fe2p, and (b) Fe3p of chalcopyrite, obtained after ion etching...................... 118 Figure 7-7: XPS spectrum of S 2p peak of chalcopyrite. ............................................................................... 120 Figure 7-8: S2p core level peaks obtained from minerals (a) 2 (pyrite, FeS2), (b) 3 (pyrrhotite, Fe1-xS), (c) 4 (chalcocite, Cu2S), and (d) 5 (covellite, CuS). Composition of each mineral is presented in Table 4-1. .......... 121 Figure 7-9: S2s core level peak of chalcopyrite. ........................................................................................... 121 Figure 7-10: Comparison of Cu2p peaks of chalcopyrite after dissolution in 0.5 M sulfuric acid at different potentials for 6 hours................................................................................................................................... 123 Figure 7-11: Copper auger spectrum (CuL3M45M45) excited with X-rays obtained after potentiostatic polarization of chalcopyrite electrode at 700 mVSHE for 6 hours....................................................................................... 124 Figure 7-12: Comparison of Fe2p peaks of chalcopyrite (a) after ion etching inside XPS instrument, and (b) after dissolution at 700 mVSHE in 0.5 M sulfuric acid solution. ...................................................................... 125 Figure 7-13: XPS spectrum of (a) Fe2p3/2 and (b) Fe3p obtained from a chalcopyrite electrode after 6 hours dissolution at 700 mVSHE.............................................................................................................................. 127 Figure 7-14: Comparison of Fe 2p peaks of chalcopyrite after dissolution at different potentials for 6 hours. 129 Figure 7-15: Comparison of S2p peaks of chalcopyrite after dissolution at different potentials for 6 hours. .. 131 Figure 7-16: S2p XPS spectra for chalcopyrite electrode surface obtained after dissolution at (a) 450, (b) OCP (~550 mVSHE), (c) 700 and (d) 870 mVSHE illustrating the characteristics of the S 3p\u2192Fe 3d interband excitation related peaks. .............................................................................................................................. 133 Figure 7-17: S2p XPS spectra for chalcopyrite electrode surface obtained after dissolution at (a) 900, (b) 950, and (c) 1000 mVSHE. .................................................................................................................................... 137 Figure 7-18: Comparison of S2s peaks of chalcopyrite after dissolution at different potentials for 6 hours. Red dashed-line represents the binding energy of S2s peak of elemental sulfur, 228.4 eV..................................... 138 Figure 7-19: Sulfur species on the surface of chalcopyrite formed after dissolution in 0.5 M sulfuric acid solution at different potentials for 6 hours. ................................................................................................... 140 Figure 7-20: Concentration of O, S, Cu and Fe species on the outer-most surface of chalcopyrite after dissolution at different potentials for 6 hours in 0.5 M sulfuric acid solution. C contamination is excluded from the analysis. ................................................................................................................................................ 142 Figure 7-21: Oxygen species on the surface of chalcopyrite formed after dissolution at different potentials for 6 hours........................................................................................................................................................... 144 Figure 7-22: O1s core level peak obtained from chalcopyrite after 6 hours of dissolution at 650 mVSHE........ 145  xi  \u000CFigure 8-1: Multicyclic voltammogram for stationary chalcopyrite electrode in 0.5 M H2SO4 deaerated (a) and aerated (b) solutions. Scan rate is 50 mV s-1. ................................................................................................ 148 Figure 8-2: Typical multicyclic voltammogram of stationary chalcopyrite in 0.5 M H2SO4 solution (a) and effect of the rotation of the electrode on the evolution of the peaks (b). Scan rate is 50 mV s-1. ............................... 150 Figure 8-3: Cyclic voltammetry of chalcopyrite electrode between 0.75 to 0.1 VSHE with (1) 0, (2) 10, (3) 100 and (4) 1000 seconds of potentiostatic polarization at 0.75 VSHE before cyclic voltammetry experiment. Scan rate is 50 mV s-1. ......................................................................................................................................... 153 Figure 8-4: Cyclic voltammetry of a rotating chalcopyrite disk electrode (500 rpm) between 0.75 and 0.1 V. Scan rate is 50 mV s\u22121. ................................................................................................................................ 155 Figure 8-5: Cyclic voltammogram of a chalcopyrite electrode in the potential window of \u22120.1 to 1 V. Scan rate is 50 mV s\u22121................................................................................................................................................. 156 Figure 8-6: Multicyclic voltammetry of chalcopyrite electrode presenting the first, second, and tenth cycles. Scan rate is 50 mV s\u22121. ................................................................................................................................ 157 Figure 8-7: Multicyclic voltammetry (10 cycles) of a stationary chalcopyrite electrode in the potential window of (a) 0.1 to 0.75 V, and (b) \u22120.1 to 0.75 V. Scan rate is 50 mV s\u22121. .............................................................. 158 Figure 8-8: Multicyclic voltammetry of a stationary chalcopyrite electrode in the potential window of \u22120.5 to 1.2 V. Scan rate is 50 mV s\u22121. ...................................................................................................................... 161 Figure 8-9: OCP measurements of a chalcopyrite electrode after polarizing at different passive potentials of 0.65, 0.75 and 0.85 V for 10 hrs. The OCP behavior of a polished chalcopyrite electrode (without previous polarization) is also presented with a solid black line................................................................................... 161 Figure 8-10: Multicyclic voltammetry of a stationary chalcopyrite electrode in the potential window of \u22120.5 to 1.0 V. Scan rate is 50 mV s\u22121. ...................................................................................................................... 164 Figure 9-1: Potentiostatic polarization of chalcopyrite in 0.5 M sulfuric acid solution.................................. 170 Figure 9-2: i-t behavior of chalcopyrite electrode in 0.5 M sulfuric acid solution at between 0.45 to 1.00 V. . 171 Figure 9-3: Potentiostatic polarization of chalcopyrite in 0.5 M sulfuric acid solution. This figure magnifies the anodic region as shown in Fig. 9-1. ............................................................................................................. 173 Figure 9-4: (a) Tafel region of the Fe3+/Fe2+ couple, and (b) potentiostatic behavior of a chalcopyrite electrode at 0.70 V before and after addition of the Fe3+/Fe2+ couple (the difference in current was used to derive Fig. 94a). ............................................................................................................................................................. 174 Figure 9-5: (a) A typical plot of charge density vs time for potentiostatic polarization at 0.60 V. The solid-red line represents the stage at which Fe-containing solution was added to the original solution. The inset figure shows an expanded view of the first 6 hours of polarization in Fe-free solution (dashed-green line). ............. 175 Figure 9-6: Measured and calculated Nyquist plots for the chalcopyrite electrode at 0.70 V in (a) Fe3+/Fe2+free, (b) Fe3+/Fe2+-containing solution......................................................................................................... 177  xii  \u000CFigure 9-7: Equivalent electrochemical circuit for the chalcopyrite/electrolyte interface. ............................. 177 Figure 9-8: Schematic illustration for the dissolution process of a passivated chalcopyrite mineral particle in ferric-containing sulfuric acid solution. High concentration of ferric (0.1 M) is available on the surface for cathodic reduction, and anodic oxidation of chalcopyrite must provide the required electrons for ferric reduction. However, diffusion of iron and copper atoms through the passive film is slow (rate-controlling step) and so the electron production rate will be slow and consequently so will be the ferric reduction rate. ......... 184 Figure 10-1: Passive current vs. time respond of a chalcopyrite electrode to potentiostatic polarization at 800 *  mV, and related thickness of the passive film ( CuS 2 ) calculated from reaction 10-1. .................................. 188 Figure 10-2: Schematic illustration for the growth of surface films on chalcopyrite surface. ......................... 189 Figure 10-3: XPS spectrum of Fe 3p level obtained from a chalcopyrite electrode after 6 hours dissolution at 700 mV. The take-off angle of the photoelectrons was (a) 45\u00BA, (b) 90\u00BA. .......................................................... 190  xiii  \u000CAcknowledgements This thesis would not have been completed without substantial supports from many individuals at UBC. First, I would like to sincerely thank my research advisors, Professor Edouard Asselin and Professor David Dixon for giving me the opportunity to join their group and for their continual support during my work at UBC. I would also like to thank my thesis committee members, Professor David Dreisinger and Professor Dan Bizzotto, for their great comments and input about my research. I would like to thank Professor Akram Alfantazi who has offered me opportunities to work closely with his group members. Without resources from his laboratory, I would not have been able to perform many of my electrochemical experiments, especially in my first year of work at UBC. Colleagues such as Dr. Farzad Mohammadi, Wei Wang, Dr. Debasis Dhak, Maziar Eghbalnia, Tirdad Nickchi, and Dr. John Skrovan are also acknowledged for their helpful discussions. I would like to express sincere thanks to the people in the UBC Hydrometallurgy and Corrosion labs for their kind help and suggestions. All the help from the department technical staff, especially Ross McLeod, Carl Ng, Rudy Cardeno, Jacob Kabel and Kim Wonsang are gratefully acknowledged. Last, but not least, I would like to thank my wife, Parisa, for her understanding and love during the past few years. Her support and encouragement was in the end what made this dissertation possible. My parents, Jabraeel and Soghra, receive my deepest gratitude and love for their dedication and the many years of support during my undergraduate studies that provided the foundation for this work.  xiv  \u000CDedication This thesis is dedicated to the memory of my late brother Mahmoud Ghahremaninezhad, who passed away on 14 June 2010.  Mahmoud Ghahremaninezhad 1980-2010  I never imagined you'd ever be so far away!  xv  \u000C1 Introduction The major copper producing countries include USA, Canada and Mexico; Chile, Peru and Brazil in South America; Zaire, Zambia, Zimbabwe, South Africa and Namibia in Africa; Papua New Guinea and Australia in Oceania; Russia, Kazakhstan, Japan, Uzbekistan, Philippines, Indonesia, India, Iran and Turkey in Asia; and Poland, Spain, Norway, Sweden and Finland in Europe [Habashi F., 1997]. According to the database of the Australian Bureau of Agricultural and Resource Economics (ABARE), world copper production has shown a progressive increase in the period 1984 to 2008 from 9Mt to 18.5Mt per annum (data for 2009 to 2011 was not available). Above 20% of world copper production is via hydrometallurgy [Watling, 2006]. The market for copper has increased each year while the quantity produced is hardly sufficient to allow all countries to reach developed world levels of usage [Gordon et al., 2006]. There are more than 200 copper-containing minerals whereas only about 20 of them are potentially relevant to copper extraction processes [Habashi F., 1997]. Since copper is a chalcophilic element, i.e. an element that would rather react with sulfur than oxygen, most of its minerals occur as sulfides. Some examples of the most significant copper bearing minerals are chalcopyrite (CuFeS2), bornite (Cu5FeS4) and chalcocite (Cu2S). Chalcopyrite is the world\u2019s most abundant copper bearing mineral [Dutrizac, 1978]. Industry standard practice is to treat this mineral via smelting to extract the copper. However, gas emissions from smelters result in soil and air pollution. As an example, a study by Aznar et al. [Aznar et al., 2008] has shown that metal pollution is 5.8 times higher in the vicinity of a smelter than in the \u2018background\u2019 sites. Using Lichens to evaluate the metal pollution in a forest ecosystem around the Murdochville smelter, Canada, Aznar et al. [Aznar et al., 2008] have shown that the forest contamination was detectable up to 30 km from the smelter. The observed enrichment of pollutants was reported to be 2 times higher for 3 metals (i.e. Cu, Cd and Pb) and the highest measured As, Ba, Cd, Cu, and Pb concentrations in lichens were 4, 112, 1, 23, 50 and 952 mg/kg respectively. Elsewhere, Hedberg et al [Hedberg et al., 2005] have shown that the fine particles originating from smelters (Quillota and Linares, two rural sites in vicinity of two smelters in Central Chile) account for more than 72% of the arsenic levels in the vicinity of the smelters. Due to environmental concerns such as these and the 1  \u000Chigh capital cost of building a smelter, the construction of new smelters in Canada is unlikely. Canada currently operates only a single primary copper smelter (The Horne, Rouyn-Noranda, PQ) which is known to treat mostly custom feed. This represents a strategic problem for a developed country such as Canada that relies heavily on natural resources for economic development.  Since smelters are unlikely to be built in Canada, there is a significant  opportunity for a hydrometallurgical process implementation in this country. Hydrometallurgical copper extraction is more attractive from many perspectives such as its ability to deal with a large spectrum of impurities, its applicability to low grade materials, its capability of low capacity operation and, very importantly, low CO2 and fine particle emissions. Considering the current state of the technology of hydrometallurgical treatment of copper sulfide concentrates, there exist a large number of quite different technologies of which today only very few have commercially been implemented for particular applications. Most of the proposed processes cannot attain sustained commercial production due to facing one or more of the following problems [Dreisinger, 2006]: (1) low copper recovery in the primary leach step; (2) copper co-precipitation with iron hydrolysis products resulting in copper loss; (3) incomplete recovery of precious metals; (4) issues with electrowinning; (5) high acid production due to unsuccessful elemental sulfur recovery; (6) corrosion problems especially in chloride circuits; (7) incomplete fixation of toxic byproducts; (8) high energy demand for mixing and oxygen dispersion and finally (9) low quality copper production which requires further processing (electrorefining). Some of the leading technologies for concentrate leaching involve high temperature and pressure leaching steps which suffer from high cost and excess sulfuric acid production. The main problem associated with atmospheric leaching processes is the slow dissolution rate of chalcopyrite due to its passivation (failure to react) [Dutrizac, 1978]. The term \u201Cpassivation\u201D is widely used for chalcopyrite dissolution in different literature [Viramontes-Gamboa et al., 2007; Vel\u00E1squez et al., 2005; Mikhlin et al., 2004; Abed and Dreisinger, 2003; Warren et al., 1982; Hackl et al., 1995]. The passive state is defined in Shreir\u2019s \u201CCorrosion\u201D book [2000] as: \u201CPassivity is a state of low corrosion rate brought about under a high anodic driving force, or potential, by the presence of an interfacial solid film, usually an oxide. The typical passivating oxide film on many metals is only some  2  \u000C1-10 nm in thickness, and is produced by oxidation of the surface to a depth measured in monolayers of atoms.\u201D Such passive films (ultra\u2013thin films) would normally be directly detectable by the sensitive surface analysis methods such as X-ray photoelectron spectroscopy (XPS/ESCA), Auger electron microscopy (AES), or secondary ion mass spectroscopy (SIMS) [Marcus and Mansfeld, 2006]. Generally, different in-situ electrochemical techniques (i.e. cyclic voltammetry, potentiodynamic and potentiostatic polarization, chronoamperometry) and exsitu spectroscopic methods (i.e. XPS, AES, Raman spectroscopy) have been applied to understand the dissolution mechanism of chalcopyrite and furthermore to characterize their surface layers.  However, there remain uncertainties about the dissolution process. The  trouble with most of the DC electrochemical techniques used so far is that only circumstantial evidence is derivable to help with the elucidation of possible reaction mechanisms. The ex-situ analytical methods are controversial for the case of sulfide minerals due to the possible evaporation of elemental sulfur under experimental conditions [Buckley and Woods, 1994]. Thus, the precise nature of the passive layers is not yet understood. In particular, there are some fundamental questions that should be answered as to the nature of the passive layers and sulfur film that form during leaching. In the present study we will concentrate on the electrochemical dissolution of chalcopyrite in simple sulfuric acid solution. The ultimate goal of any electrochemical study pertaining to the dissolution of chalcopyrite is to overcome its slow dissolution rate. Clearly, the path to this goal runs through obtaining a comprehensive knowledge on the surface chemistry of chalcopyrite in the desired dissolution environment. Thus, the present work seeks to understand the exact nature of the surface layers (mainly the passive layers) which form at low anodic potential (&lt; 1 V vs. SHE) through their careful characterization to obtain their composition. In order to characterize the surface properties of chalcopyrite electrodes at different anodic potentials and understand the growth sequence of surface layers, we will apply different electrochemical and analytical methods. In the following chapter (chapter 2) the literature pertaining to this study will be discussed. Chapter 3 will present the objectives of this study. The experimental procedures for the electrochemical and analytical methods will be provided in chapter 4. Chapter 5 discusses the results of electrochemical tests on a pyrrhotite (Fe1-xS) electrode, which is a 3  \u000Csimplified case of chalcopyrite (CuFeS2) electrochemistry. Chapter 6 presents the results of an electrochemical impedance spectroscopy study on chalcopyrite. The results of XPS and cyclic voltammetry studies on a chalcopyrite electrode will be presented in chapters 7 and 8, respectively. Chapter 9 discusses the kinetics of the ferric/ferrous couple on a passivated chalcopyrite electrode. In chapter 10, a dissolution model for chalcopyrite in 0.5 M sulfuric acid solution is proposed. Finally, chapter 11 presents the conclusions and recommendations for future work in this area.  4  \u000C2 Literature review 2.1 General aspects 2.1.1 Chalcopyrite crystal structure The crystal structure of chalcopyrite was determined by Burdick and Ellis in 1917 [Burdick and Ellis, 1917]. Chalcopyrite\u2019s structure was the first determined crystal structure in USA and in addition, its structure was the first structure of a complex sulfide material to be determined in the World [Donnay et al., 1958]. The crystal structure of chalcopyrite is derived from the zincblende (sphalerite, \u03B2\u2013ZnS) structure, in which the two zinc atoms are replaced by copper and iron [Habashi, 1978; Hiskey, 1993]. Figure 2-1 shows the crystal structure of chalcopyrite. Each metal atom is surrounded by sulfur and each sulfur atom by metal in tetrahedral interstices. In this structure, each sulfur atom has bonds with two copper atoms and two iron atoms. The copper and iron atoms have opposite magnetic moments and hence chalcopyrite shows antiferromagnetic property at room temperature [Donnay et al., 1958]. From neutron diffraction studies, the moments of iron and copper in the chalcopyrite structure were found to be 3.85 and 0\u00B10.2 \u03BCB respectively. The moment of iron atoms, 3.85 \u03BCB, is lower than that of a typical trivalent iron, 5 \u03BCB [Greenwood and Whitfield, 1968]. It is worth mentioning that iron atoms in the pyrite structure (FeS2) have no magnetic moment [Hamajima et al., 1981]. Hence the iron atom in chalcopyrite can be considered neither as Fe3+ nor Fe2+ and the bonding of Fe atoms is undoubtedly complex. Donnay et al., [Donnay et al., 1958] have suggested that the complex behavior of iron atoms can be explained by postulating further participation of the 3d level electrons of trivalent iron atoms in the covalent bonding. By this, one would expect a decrease in the iron moment (i.e. from 5 to 3.85 \u03BCB) and at the same time a stronger bond between sulfur atoms and trivalent iron atom. Considering the above discussion, from the two probable ionic states for chalcopyrite, Cu \uF02B Fe 3\uF02B S22 \uF02D and Cu 2\uF02B Fe 2\uF02BS 22\uF02D , the first is the most appropriate. Indeed, the Mossbauer  spectra study by Greenwood and Whitfield [Greenwood and Whitfield, 1968] has further confirmed that the best chemical formula for chalcopyrite is Cu \uF02B Fe 3\uF02B S22 \uF02D . In fact, most of the recent workers have considered this structure in their studies [Mikhlin et al., 2004; Boekema  5  \u000Cet al., 2004; C\u00F3rdoba et al., 2008]. The average bond distances for S\u2013S, Fe\u2013S and Cu\u2013S are 0.368, 0.225 and 0.235 nm, respectively [Robie et al., 1967; Hall and Stewart, 1973]. Lattice constants for the unit cell are a \uF03D 5.2988 \uF0B1 0.0010 and  c \uF03D 10.4340 \uF0B1 0.0050 angstroms  [Robie et al., 1967].  Figure 2-1: Unit-cell crystal structure of CuFeS2 [C\u00F3rdoba et al., 2008].  2.1.2 Semiconductive properties of chalcopyrite Results from infra\u2013red studies show an energy gap (Eg) of 0.53 eV for natural chalcopyrite specimens [Austin et al., 1956]. For a thin film of evaporated chalcopyrite the band gap energy was determined to be 0.60 eV [Teranishi, 1974]. Due to the antiferromagnetic property of chalcopyrite and its crystal structure, it is very difficult to precisely compute the energy bands for chalcopyrite. However, studies on Fe-doped CuAlS2 and CuGaS2 materials provided evidence that the valence band edge is made up of Cu 3d and S 3p orbits and the conduction band is mainly Fe 3d in character [Shuey, 1975]. Figure 2-2 shows the schematic energy band diagram for an n-type chalcopyrite sample.  6  \u000CFigure 2-2: Schematic energy band diagram for chalcopyrite CuFeS2 [C\u00F3rdoba et al., 2008].  Conduction of semiconductor materials is related to the movement of mobile charge carriers; namely, free electrons of the normally empty conduction band or holes of the normally full valence band. Charge carriers in solid materials may be derived from: (a) non\u2013 stoichiometry; (b) impurity defects; (c) thermal excitation [Rockett, 2008]. The main donor in chalcopyrite is a metal interstitial. The conduction band carrier density for large natural single crystal chalcopyrite specimen at room temperature is calculated to be 1025 m-3 [Teranishi, 1961]. For a polycrystalline sample it is shown that the charge carrier density is around 4 \uF0B410 25 m-3 [Donovan and Reichenbaum, 1958]. Some of the material properties of chalcopyrite, extracted from Hiskey\u2019s work [Hiskey, 1993] are summarized in Table 2-1. Electrical studies show that chalcopyrite clearly has properties of a good semiconductor material and, with only a few exceptions chalcopyrite demonstrates an n\u2013type semiconductor behavior. Note that the general semiconductor type on chalcopyrite strongly depends on the impurities of the mineral. Electrons show less mobility rather than holes in natural chalcopyrite samples [Tell et al., 1972]. Since chalcopyrite is a semiconductor material, it is worth having a brief discussion of the dissolution of semiconductor materials.  2.2 Dissolution of semiconductor materials In order to understand the properties of semiconductor electrodes, and their differences from those of metallic electrodes, one should understand the electronic structures of semiconductors and be familiar with the concept of energy bands in solids. Figure 2-3 shows the generation of a band model in solids from atomic orbitals of isolated atoms [Bard and Faulkner, 2001]. Table 2-1: Materials properties of CuFeS2 at room temperature [Hiskey, 1993]. 7  \u000CStructure  Tetragonal  Ionic structure  Cu \uF02B Fe 3\uF02B S22\uF02D  Density (g cm-3)  4.0878  Hardness, Moh  3.5 \u2013 4  Cleavage surface  (011)  C op (J mol-1 K-1) 298 \u2013 830 K  88 . 71 \uF02B 5 . 19 \uF0B4 10 \uF02D2 T \uF02D 7 .45 \uF0B4 10 5 T \uF02D2  \uF044G of (kJ mol-1)  -173.4  \uF044H of (kJ mol-1)  -173.3  \u03B1-CuFeS2 (tetragonal) \u2192 \u03B2-CuFeS2 (cubic)  550 \u00B0C  Energy gap at room temperature (eV)  0.60  Thermo electric power (\u03BCVK-1)  480  Resistivity (\u2126m)  2 \uF0B4 10 \uF02D4 \uF02D 3 \uF0B4 10 \uF02D1  Carrier density (m-3)  10 25 \uF02D 4 \uF0B4 10 25  Electron mobility (cm2V-1s-1)  18  Hole mobility (cm2V-1s-1)  21  Figure 2-3: Band model in solids [Bard and Faulkner, 2001].  8  \u000CThe highest and lowest energy levels of a band are referred to as the band edges. The energy bands of interest are the highest occupied (called the valence band) and the lowest unoccupied (called the conduction) bands (Figure 2-3). It is the energy gap (the band gap) between these bands (i.e., the difference in energy between the upper edge of the valence band and the lower edge of the conduction band, EB) that determines the properties of the material [Yacobi, 2004]. A solid state material is considered as an electrically conductive material if the electrons occupy partially filled orbitals. In other words, for conductivity, electrons should occupy the conduction band (lowest unoccupied band). For metals, the conduction and valence bands overlap, and so, the conduction band can be readily occupied (Figure 2-3). For insulators, the band gap is sufficiently large that electrons cannot be promoted from the valence band to the conduction band (typically above 4 eV). However, for semiconductors, the band gap is not as large (below 4 eV), and electrons can be moved into the conduction band. The promotion of electrons from valence band to the conduction band leaves positively charged vacancies in the valence band, which is referred to as a hole. This is shown in the schematic of Figure 2-4. In this figure, the dark dots represent electrons while holes are represented by the plus sign. These holes can be moved through space by the transfer of an electron to the vacancy; therefore, holes are considered to be mobile [Yacobi, 2004; Rockett, 2008].  Figure 2-4: Schematic diagram of the energy levels of a semiconductor.  The promotion of electrons from the valence band to the conduction band is possible either thermally or photochemically; however, there is another method for generating charge carriers (i.e., electrons or holes) within a semiconductor, referred to as doping. Doping is the addition of a different element(s) into the semiconductor. The simplest example of this process involves the doping of a group V element (e.g., P) or a group III element (e.g., Al) 9  \u000Cinto a group IV element (e.g., Si), as presented in Figure 2-5. For instance, the doping of phosphorous (P) into Si semiconductors introduces occupied energy levels into the band gap close to the lower edge of the conduction band, thereby allowing facile promotion of electrons from the valance band into the conduction band. Thus the addition of aluminum (Al) into Si introduces vacant energy levels into the band gap close to the upper edge of the valence band, which allows facile promotion of electrons from the valence band by formation of holes in the valence band. Doped semiconductors in which the dominant (or majority) charge carriers are electrons are referred to as n-type semiconductors, whereas those in which holes are the majority charge carriers are referred to as p-type semiconductors. Un-doped semiconductors are referred to as intrinsic semiconductors [Yacobi, 2004; Rockett, 2008].  Figure 2-5: Schematic diagram of the energy levels of (a) an n-type semiconductor and (b) a p-type semiconductor.  An important concept in the study of solid state materials is the Fermi level. The Fermi level is the level of energy which is occupied by an electron with a probability of 0.5. For example, as shown in Figure 2-4, for an intrinsic (un-doped) semiconductor the Fermi level lies at the mid-point of the band gap. Doping changes the distribution of electrons within the solid, and hence changes the Fermi level. For instance, in an n-type semiconductor the Fermi level lies closer to the conduction band (Figure 2-5a), whereas for p-type semiconductors it lies closer to the valence band (Figure 2-5b). It should be mentioned that the Fermi level of a semiconductor electrode changes with applied potential, just as it does in metal electrodes; for example, moving to more negative potentials will raise the Fermi level [Morrison, 1980].  10  \u000C2.2.1 Semiconductor-electrolyte interface Generally, in electrochemical science, the understanding of reactions and phenomena at the interface between electrode and electrolyte is important. In order for the two separate phases to be thermodynamically in equilibrium, their electrochemical potential must be the same. The electrochemical potential of the solution is determined by the redox potential of the electrolyte solution, and the electrochemical potential of the semiconductor is determined by the Fermi level. If the redox potential of the solution and the Fermi level do not lie at the same energy, then charges will be exchanged between the semiconductor and the solution in order to equilibrate the two phases [Memming, 2001]. In contrast to metals, the carrier density in a semiconductor is much smaller. Therefore, in a semiconductor electrode the charge in the solid is distributed over a certain distance below the electrode surface. This region, with a thickness of 100-10,000 \u00C5, is called the \u201Cspace charge region\u201D, and has an associated electrical field. The resulting potential and charge distribution is shown in Figure 2-6 (considering an electrolyte with a high ion concentration so that the diffuse layer on the solution side is neglected). Hence, in the case of semiconductors, there are two double layers to consider: the interfacial (electrode/electrolyte) double layer and the space charge layer [Morrison, 1980; Memming, 2001].  Figure 2-6: Potential (a) and charge (b) distribution at the semiconductor(left side)/electrolyte(right side) interface. Note that a high ion concentration for solution is assumed hence the diffuse layer on the solution side is neglected [Memming, 2001].  11  \u000CConsidering the contact between an n-type semiconductor electrode and an electrolyte at open circuit, the Fermi level (electrochemical potential of the semiconductor) is typically higher than the redox potential of the electrolyte, and hence electrons will be transferred from the semiconductor electrode into the solution. Therefore, the space charge region will have positive charge and the band edge will be bent upward. Since the majority charge carrier of the semiconductor has been removed from the space charge region, this region is also referred to as a \u201Cdepletion layer\u201D. In the case of a p-type semiconductor, the Fermi layer is generally lower than the redox potential. Hence, in order to have equilibrium at the interface, electrons must transfer from the solution to the electrode. This generates a negative charge in the space charge region, which causes a downward bending in the band edges. Since the holes in the space charge region are removed by this process, this region is again a depletion layer [Morrison, 1980; Memming, 2001]. As for metallic electrodes, the Fermi level of a semiconductor is a function of the applied potential. The band edges in the interior of the semiconductor (i.e., away from the depletion region) also vary with the applied potential in the same way as the Fermi level. However, the energies of the band edges at the interface are not affected by changes in the applied potential. Therefore, going from the interior of the semiconductor to the interface, the energies of the band edges varies with the applied potential. There are three different scenarios to be considered [Morrison, 1980; Memming, 2001]: (a) The Fermi energy level of the semiconductor lies at an energy equal to the solution redox potential. In this case, there is no net transfer of charge, and hence there is no band bending. This potential is therefore referred to as the \u201Cflat band potential\u201D, Efb. (b) Depletion regions arise at potentials positive of the flat band potential for an n-type semiconductor and at potentials negative of the flat band potential for a p-type semiconductor. (c) In the case of an n-type semiconductor at potentials negative of the flat band potential there is an excess of the electrons (majority charge carrier) in this space charge region, and so the region is referred to as an \u201Caccumulation\u201D region. For p-type semiconductors an accumulation region arises at potentials more positive than the flat band potential. It is important to note that having an accumulation layer or a depletion layer determines the charge transfer ability of a semiconductor electrode. If there is an 12  \u000Caccumulation layer, the semiconductor electrode will behave similarly to a metallic electrode, since there is an excess of the majority charge carrier available for charge transfer. In contrast, if there is a depletion layer, then there are few available charge carriers, and electron transfer reactions will occur slowly.  2.2.2 Mott-Schottky equation There are a few different experiments for measuring the various parameters of semiconductors. As an example, by applying different radiation intensity to a semiconductor electrode and measuring the photo-potential, it is possible to determine the flat band potential of the semiconductor. The simplest method is to measure the open-circuit potential (photopotential) of the electrochemical cell as a function of irradiation intensity. Another important method for studying semiconductors is measuring the apparent capacitance of the surface of the semiconductor as a function of potential under depletion conditions and this approach is based on the Mott-Schottky relationship [Morrison, 1980; Memming, 2001]. The potential and charge distribution within the space charge region is quantitatively described by the Poisson equation as given by: d 2 \uF044\uF066SC \uF02D 1 \uF03D \uF072 \uF028x \uF029 \uF065\uF065 0 dx 2  2-1  in which the charge density is given by \uF072 \uF028 x \uF029 \uF03D e\uF05BN d \uF02D N a \uF02D n\uF028 x \uF029 \uF02B p \uF028x \uF029\uF05D  2-2  here, x is the distance from the surface and Nd and Na, are the fixed ionized donor and acceptor densities, respectively, which are given by the doping of the semiconductor. The \u03B5 is the dielectric constant of the medium and \u03B5o is the permittivity of free space. The electron and hole densities, n(x) and p(x), vary with the distance x according to:  \uF0E6 e\uF044\uF066SC ( x) \uF0F6 n( x) \uF03D no exp\uF0E7 \uF02D \uF0F7 kT \uF0F8 \uF0E8  2-3  \uF0E6 e\uF044\uF066 SC ( x) \uF0F6 p( x) \uF03D po exp\uF0E7 \uF0F7 \uF0E8 kT \uF0F8  2-4  Considering the charge neutrality in the bulk of the semiconductor material, then we have, 13  \u000C2-5  N d \uF02D N a \uF03D no \uF02D po  Inserting of Eqs. 2-2, 2-3, 2-4 and 2.5 into the Eq. 2-1 then one can obtain: d 2 \uF044\uF066 SC e \uF0E9 \uF0E6 e\uF044\uF066 SC ( x) \uF0F6 \uF0E6 e\uF044\uF066 SC ( x ) \uF0F6\uF0F9 \uF03D no \uF02D po \uF02D no exp\uF0E7 \uF02D \uF0F7 \uF02B po exp\uF0E7 \uF0F7\uF0FA \uF0EA 2 \uF065\uF065 o \uF0EB kT dx \uF0E8 \uF0F8 \uF0E8 kT \uF0F8\uF0FB  2-6  After integration of Eq. 2-6 the electric field at the surface (Es) can be derived as given by  d\uF044\uF066 SC Es \uF03D dx  2  x\uF03D0  \uF0E6 kT \uF0F6 \uF0EC Po \uF0E9 \uF0E6 e\uF044\uF066 SC \uF0F6 \uF0F9 no \uF0F7\uF0F7 \uF0ED \uF0EAexp\uF0E7 \uF03D \uF0E7\uF0E7 \uF0F7 \uF02D 1\uF0FA \uF02B \uF0E8 eLD \uF0F8 \uF0EE ni \uF0EB \uF0E8 kT \uF0F8 \uF0FB ni  \uF0E9 \uF0E6 e\uF044\uF066SC \uF0F6\uF0F9 \uF0FC \uF02D exp \uF02D 1\uF0F7\uF0FA \uF0FD \uF0E7 \uF0EA \uF0E8 kT \uF0F8\uF0FB \uF0FE \uF0EB  2-7  where, ni is the intrinsic carrier density which is determined by the equilibrium equation, nopo = ni2. The term LD is the so-called Debye length as given by, \uF0E6 \uF065\uF065 kT LD \uF03D \uF0E7\uF0E7 o 2 \uF0E8 2 ni e  \uF0F6 \uF0F7\uF0F7 \uF0F8  1  2  2-8  Furthermore, it is possible to determine the required space charge for producing the field Es by Gauss\u2019 law, QSC \uF03D \uF065\uF065 o E S  2-9  The differential capacity of the space charge layer is defined by  CSC \uF03D  dQSC d \uF028\uF044\uF066 SC \uF029  2-10  Inserting Eqs. 2-7 and 2-8 into 2-10 and differentiation, then we will have,  CSC \uF03D  \uF065\uF065 o \uF0E6 e\uF044QSC \uF0F6 cosh\uF0E7 \uF0F7 LD \uF0E8 kT \uF0F8  2-11  This equation shows the strong relationship between the space charge capacity (CSC) and potential across the space charge layer. Rearranging the equation we will have the famous Mott-Schottky relationship [Memming, 2001],  14  \u000C\uF0E6 2 LD.eff 1 \uF0E7\uF0E7 \uF03D 2 C SC \uF0E8 \uF065\uF065 o  2  \uF0F6 \uF0E6 e\uF044\uF066 SC \uF0F6 \uF0F7\uF0F7 \uF0E7 \uF02D 1\uF0F7 \uF0F8 \uF0F8 \uF0E8 kT  2-12  where,  LD.eff  \uF0E6 \uF065\uF065 kT \uF03D \uF0E7\uF0E7 o 2 \uF0E8 2 ni e  \uF0F6 \uF0F7\uF0F7 \uF0F8  1  2  ; Csc is the capacitance of the space charge region, \u03B5 is the dielectric  constant of the semiconductor and \u03B5o is the permittivity of free space. In equation 12-2,  \uF044\uF066SC \uF03D E \uF02D E fb and in here E represents applied potential and Efb presents flat band potential. Making the equation simpler, then we will have [Morrison, 1980],  1 2 \uF0E6 kT \uF0F6 \uF03D \uF0E7 E \uF02D E fb \uF02D \uF0F7 2 CSC e\uF065\uF065 o no \uF0E8 e \uF0F8  2-13  Equation 2-13 predicts a linear relation between 1/C2 and applied potential (E) in the depletion region. A classical example of a Mott-Schottky experiment on an n-type and a ptype silicon electrode has been presented by Schmuki et al. [Schmuki et al., 1995]. The donor density can be calculated from the slope of the Mott-Schottky plots (1/C2 vs. E), and the flat band potential can be determined by extrapolation to C = 0. The capacitance values are calculated from electrochemical impedance spectroscopy (EIS) measurements. In the case of a semiconductor electrode, there are two capacitances to be considered, that of the space charge region and that of the double layer. These two capacitances are in series and as a result, the total capacitance is the sum of their reciprocal. The capacitance for the space charge region is much smaller than that of the double layer (2-3 orders of magnitude) such that the double layer capacitance can be practically neglected in calculations. Therefore, the measured capacitance value is equal to the space charge capacitance. Through EIS experiments the capacitance can then be calculated from the imaginary component of the impedance (Z\") using the relationship Z\" = 1/2\u03C0fC.  2.2.3 Current-potential curves for semiconductors The application of a potential between a semiconductor electrode and a counter electrode in an electrolyte leads to a current flux between the two electrodes. Since the charge transfer across a semiconductor/electrolyte interface can only occur via the 15  \u000Cconduction or valence band, the processes in the two bands have to be treated separately. In general anodic and cathodic currents are given by equations similar to those for metallic electrodes, but with some corrections (using boundary conditions which are specific for semiconductors) in the generation of the relationships [Memming, 2001]. For an n-type semiconductor in the accumulation region, electrons accumulate at the surface and a cathodic current can flow. At high anodic potentials, holes accumulate in the surface region and an anodic current can flow. For a p-type semiconductor in the accumulation region an anodic current is observed. In the depletion region current cannot flow, and in the degeneration region a cathodic current can flow [Plieth, 2008]. This is shown in the schematic current\u2013 potential curve in Figure 2-7.  Figure 2-7: Schematic current\u2013potential curve for n-type and p-type semiconductors.  The charge transfer on a metallic electrode depends on the overlap between occupied and unoccupied energy levels in the metal and electrolyte. This can be found around the Fermi level. However, in the case of a semiconductor electrode in contact with electrolyte, the redox system can overlap with both conduction and valence bands. Depending on the redox system and electrode potential, charge transfer is possible between the conduction band and redox system and between the valence band and redox system [Memming, 2001; Plieth, 2008]. If the density of electrons in the conduction band on the surface of the  16  \u000Csemiconductor electrode is ns and the density of unoccupied energy levels in the electrolyte is Dox, then the partial current density of conduction band reduction is \uF02B\uF0A5  2-14  jc ,c \uF03D k c ,c \uF0F2 nS Dox dE \uF02D\uF0A5  Otherwise, the anodic conduction band current is proportional to the density of occupied states in the electrolyte Dred times the density of states in the conduction band; \uF02B\uF0A5  2-15  jc ,a \uF03D k c ,a \uF0F2 N C Dred dE \uF02D\uF0A5  Similarly, it is possible to formulate the anodic valence band process with ps denoting the surface density of holes in the valence band \uF02B\uF0A5  2-16  jv ,a \uF03D k v ,a \uF0F2 Dred ps dE \uF02D\uF0A5  For the cathodic valence band process also, with Nvb, the density of energy levels in the valence band, we have \uF02B\uF0A5  2-17  jv ,c \uF03D k v ,c \uF0F2 Dox N v dE \uF02D\uF0A5  In a first approximation Dred, Dox, Nc, and Nv are independent of the potential. The surface density of electrons in the conduction band and the surface density of holes in the valence band, ns and ps, depend exponentially on the potential. The current density at the flat band potential can be used as a reference point. In this approximation one obtains [Plieth, 2008],  \uF0E6 F \uF028E \uF02D E fb \uF029\uF0F6\uF0F7, ic ,c \uF03D ic ,o exp\uF0E7 \uF02D \uF0E8 RT \uF0F8  and  ic ,a \uF03D const  2-18  and  iv,c \uF03D const  2-19  and  \uF0E6 F \uF028E \uF02D E fb \uF029\uF0F6\uF0F7, iv ,a \uF03D iv ,o exp\uF0E7 \uF02B \uF0E8 RT \uF0F8  This behavior is schematically represented in Figure 2-7. The slope of the rise of the current in the accumulation region in Figure 2-7, often differs from the 59 mV expected according to Eqs. 2.18 and 2.19. Otherwise, in the accumulation region the semiconductor  17  \u000Celectrodes approach metal-like properties and the classical theory of electron transfer must be applied [Plieth, 2008].  2.3 Chalcopyrite dissolution and electrochemistry Electrochemical investigations can build valuable insight about the characterization of leaching mechanisms and rate\u2013determining steps for chalcopyrite. Generally the anodic dissolution of chalcopyrite reveals two distinct regions, namely, low potential processes (up to 0.8 V vs. SHE), and high potential processes (1.0 to 1.2 V vs. SHE). The low potential region is characterized by \u201Cpassive\u201D behavior which shows progressively thickening surface films. Since this potential region is the typical oxidation potential range employed in hydrometallurgy, the reactions in this potential region are very important. In the high potential region (1.0 to 1.2 V vs. SHE) substantial oxidation of chalcopyrite occurs [Hiskey, 1993]. The Pourbaix diagram for the CuFeS2-H2O system [Garrels and Christ, 1965] suggests that the acidic dissolution (25 \u00BAC and 1 atm of total pressure) of chalcopyrite is thermodynamically possible at pH values lower than 6 and oxidation potentials higher than 0.2 V vs. SHE [Garrels and Christ, 1965]. Additionally, the Pourbaix diagram of chalcopyrite suggests that the dissolution process of chalcopyrite occurs through a solid transformation in different intermediate sulfides (Cu5FeS4, CuS, Cu2S), increasingly richer in copper content [Biswas and Davenport, 1976]. In order to study the electrochemical behavior of chalcopyrite in different solutions, various electrochemical techniques have been used previously. In the next section, we will discuss the electrochemical dissolution of chalcopyrite based on the applied electrochemical technique.  2.3.1 Polarization studies Potentiodynamic polarization studies on natural chalcopyrite samples in 1M H2SO4 indicate passive\u2013like behavior during anodic polarization. Figure 2-8 shows the anodic polarization curves for CuFeS2 from 6 different mines in 1 M H2SO4 [Warren et al., 1982]. As presented in this figure, the corrosion potential, passive current density and transpassive potential of chalcopyrite depends on the origin of the mineral, i.e. the chemical composition of mineral. It has been suggested by these workers that the impurities present in the mineral play a significant role in the dissolution behavior of the electrode. From potentiodynamic  18  \u000Cstudies, Warren et al [Warren et al., 1982] have suggested that during the anodic polarization two distinct intermediate layers form on the surface. Additionally, it was concluded that the second surface film forms by decomposition of the first film. In this study, it was suggested that the first surface layer might be bornite (Cu5FeS4), which forms due to selective dissolution of iron atoms from chalcopyrite structure, and the second layer was suggested to be covellite (CuS). Thermodynamic predictions from the system\u2019s Pourbaix diagram were the origin for these phase choices. [Warren et al., 1982].  Figure 2-8: Anodic polarization curves for CuFeS2 from 6 different locations in 1 M H2SO4, 30 mV min-1, 25\u00B0C [Warren et al., 1982].  Studies by Viramontes-Gamboa et al. [2006] have demonstrated that many of the uncertainties associated with potentiodynamic anodic polarizations of chalcopyrite samples are due to the fast scanning rates of the tests and that these will disappear if the potentiostatic method, effectively a very slow scan rate, is applied. Thus the scan rate should be slow enough to detect the effects of the passive layer [Viramontes-Gamboa et al., 2006]. Potentiostatic anodic polarization studies from the same group relating to the oxidation of bulk chalcopyrite electrodes in H2O\u2013H2SO4\u2013Fe2(SO4)3\u2013FeSO4 media (the acid concentration varies from 2 to 100 gL-1 and temperature from 25 to 80\u00B0C) have demonstrated that chalcopyrite oxidation displays the classic active\u2013passive behavior of passivating metals [Viramontes-Gamboa et al., 2007]. As shown in Figure 2-9, the passivation potential is a function of temperature. These workers have claimed that increasing temperature increases  19  \u000Cthe transpassive potential of chalcopyrite. It has also been stated that the time required for passivation of chalcopyrite ranges roughly from 4 to 24 h, which means that the formation of the passive layer is a very slow process at chalcopyrite electrodes.  Figure 2-9: Potentiostatic anodic polarizations of chalcopyrite, 10 mV hr-1 [Viramontes-Gamboa et al., 2007].  In a very recent study [Vasquez et al., 2012] on the electrochemical response of chalcopyrite microelectrodes in 90 g/l sulfuric acid solution, it has been shown that the transpassive potential of chalcopyrite decreases with increased temperature. This apparent discrepancy with previous findings [Viramontes-Gamboa et al., 2006] has been attributed to the use of micro-particles (ranging between 20 \u2013 100 \u03BCm) instead of bulk chalcopyrite (Figure 2-10).  20  \u000CFigure 2-10: Anodic oxidative behavior of chalcopyrite microparticles at 20, 45 and 60\u00BAC and 90 g/L of H2SO4 [Vasquez et al., 2012].  Generally, all of the polarization work by different researchers has demonstrated active\u2013 passive behavior for chalcopyrite in acidic media during the dissolution process.  2.3.2 Chronoamperometry studies Current vs. time behavior of chalcopyrite at a constant anodic potential has been studied by several researchers. Studies by Biegler and Swift [1979] showed critical potentials of 1.01 V and 1.13 V vs. SHE for the dissolution of chalcopyrite in 1 M H2SO4 and 1 M HCl solutions, respectively, below which potentials the current always decays with time and above which, a minimum current value is registered (Figure 2-11). By increasing the dissolution potential the minimum current value appears after a shorter time. Biegler and Swift considered that the dissolution of the chalcopyrite occurs at localized sites, and the number of sites varies with potential. Also, the potentiostatic current\u2013time curves were interpreted in terms of the kinetics of nucleation of dissolution centers, their 3D growth, and finally overlapping of the dissolution centers. Raising the potential increased the number of corrosion sites dramatically. In the case of the hemispherical growth of pits, the current should increase with t2, which is seen in Figure 2-11b. The dissolution potentiostatic current transient is more complex in Figure 2-11a. The presence of the second minimum is related to the overlapping of separate growth centers in a surface process [Biegler and Swift, 1979]. 21  \u000CFigure 2-11: Current-time curves for a chalcopyrite electrode at the potentials (vs. SHE) indicated in (a) 1 M H2SO4, (b) 1 M HC1 [Biegler and Swift, 1979].  Another study considers the current-time transient behavior of a freshly polished chalcopyrite electrode in 0.3 M HCl and 0.3 M H2SO4 solutions over the temperature range of 25 to 90\u00B0C [McMillan et al., 1982]. Results show an initial capacitance charging current followed by a faradaic reaction for which the current decays with time for an overpotential range of 20 to 600 mV vs. SHE. For the current transient at potentials higher than 600 mV (less than 1 V) there is no longer a decay of current by time. These results are similar to those obtained by Biegler and Swift [1979]. McMillan et al. [1982] have also suggested the formation of an iron deficient copper sulfide species (similar to the surface layer proposed by Warren et al. [1982]) on the surface which acts as an interphase between the electrode and the electrolyte; this layer has characteristics of a solid electrolyte which is electronically insulating but can transfer the ionic species from the chalcopyrite to the electrolyte [McMillan et al., 1982]. In another study, the current decay measurements for a chalcopyrite electrode at constant potential in 1 M H2SO4 solution adhered to the Sato-Cohen (logarithmic) model for solid film formation [Warren et al., 1982]. In the low potential passive region (approximately &lt; 0.79 V vs. SHE) current and mass balance measurements support the formation of two intermediate sulfide phases in the sequence CuFeS2 \u2192 S1 \u2192 S2. At higher potentials (&gt; 1.0 V), in the transpassive region, the observed increase in current is  22  \u000Ccompatible with the decomposition of water to form chemisorbed oxygen which releases copper and forms sulfate ions [Warren et al., 1982]. Previously, results of a study on the anodic oxidation of galena in acidic electrolyte showed that \u201Csurface relaxation\u201D occurs on the metal\u2013deficient products of the surface, with lead diffusing from the bulk of the mineral to the metal\u2013deficient surface [Buckley and Woods, 1996]. The same process has been detected with carbon paste chalcopyrite electrodes in 1 M H2SO4 electrolyte. The chalcopyrite electrode has been oxidized at 800 mV vs. SHE for 120 s and then allowed to rest at its open circuit potential for different time durations, and afterward, the 800 mV reapplied. As shown in Figure 2-12, the resulting current transients reveal the dependence of the current on the rest time of the electrode. This behavior provides further evidence that the passive layer of chalcopyrite is a metal-deficient surface layer and not elemental sulfur. A longer rest time results in more copper and iron diffusion from the bulk electrode to the surface, and thus, the re\u2013oxidation process of chalcopyrite occurs at a much more rapid rate [Lu et al., 2000].  Figure 2-12: Current\u2013time curves representing the re-oxidation of chalcopyrite at 800 mV vs. SHE [Lu et al., 2000].  2.3.3 Electrochemical impedance spectroscopy (EIS) studies The electrochemical impedance spectroscopy (EIS) method provides a kinetic model in the form of an equivalent circuit. One can attribute a clear physical meaning to each of the parameters (circuit elements) of the model. Moreover, EIS is a powerful technique for surface layer studies that can be used to study the layers formed on the surface of a chalcopyrite electrode. EIS has almost been disregarded for chalcopyrite studies despite the 23  \u000Cfact that it presents a powerful way of analyzing surface processes. The number of published papers in reputable journals which have studied the anodic dissolution of chalcopyrite electrodes by EIS is fewer than ten. This is partially due to the complexity of the EIS responses of chalcopyrite electrodes which make the analysis and modeling of the electrode processes difficult. Velasquez et al. [1998] were the first group to have studied chalcopyrite dissolution by EIS. Their studies focused on the dissolution of natural chalcopyrite samples from the \u201CEl Teniente\u201D mine, Chile, in an alkaline solution (pH 9.2) for different oxidation potentials. They concluded that for potentials lower than 0.4 V vs. SCE (0.642 vs. SHE) a Randles equivalent circuit can be used and for potentials higher than 0.4 V vs. SCE (0.642 vs. SHE) a layer model had to be applied. The Randles equivalent circuit is a circuit that consists of an active electrolyte resistance (RS) in series with the parallel combination of the doublelayer capacitance (Cdl) and an impedance of a faradaic reaction. The detected layers in the high potential region, E &gt; 0.642 vs. SHE, are suggested to be an inner layer of CuO and outer layer of Fe2O3 [Velasquez et al., 1998]. Another EIS study has been done by Hiroyoshi et al. [2004] on the effect of cupric and ferrous on the active-passive behavior of natural chalcopyrite in 0.1 M H2SO4 solution (with or without cupric and/or ferrous ions). The results indicated that the coexistence of cupric and ferrous ions promote the anodic dissolution current at low anodic potentials (&lt; 0.65 V vs. SHE). It has also been shown that the coexistence of cupric and ferrous ions causes the formation of a product layer on the electrode surface which inhibits the growth of a high resistance passive layer \u2013 typically a high resistance passive layer forms on the surface in the case of simple sulfuric acid solutions without cupric and ferrous ions. Figure 2-13a shows the impedance plots for a chalcopyrite electrode with or without the ferrous and the cupric ion (not coexistence of cupric and ferrous ions) and Figure 2-13b shows the effect of the coexistence of cupric and ferrous [Hiroyoshi et al., 2004].  24  \u000Ca  b  Figure 2-13: Impedance plots for the chalcopyrite electrode; (a) with or without the ferrous and the cupric ion, (b) coexisting of ferrous and the cupric ion, at 0.65 V vs. SHE [Hiroyoshi et al., 2004].  From the analysis of the EIS results Hiroyoshi et al. have proposed that in the case of the coexistence of cupric and ferrous ions in solution, a chalcocite film (Cu2S) might form on the surface of the electrode which inhibits the formation of a typical passive film on the surface. The model proposed by these workers has the following reaction scheme [Hiroyoshi et al., 2004]: CuFeS 2 \uF02B 2Cu 2\uF02B \uF02B 3Fe 2\uF02B \uF03D 2Cu 2S \uF02B 4Fe 3\uF02B  2-20  Cu 2S \uF02B 4Fe 3\uF02B \uF03D 2Cu 2\uF02B \uF02B S \uF02B 4Fe 2 \uF02B  2-21  Cu 2S \uF02B 4H \uF02B \uF02B 2O 2 \uF03D 2Cu 2\uF02B \uF02B S \uF02B 2H 2 O  2-22  It must be mentioned that the formation of elemental sulfur and its passivating effect via reactions 2-21 and 2-22 have not been verified in this study. Additionally, the most well known dissolution path for chalcocite is through formation of covellite (CuS) by the following reactions [Dutrizac and MacDonald 1974]: Cu 2S \uF03D Cu 2\uF02D x S \uF02B xCu 2\uF02B \uF02B 2xe \uF02D  2-23  CuS \uF03D S \uF02B Cu 2\uF02B \uF02B 2e \uF02D  2-24  Reaction 2-23 is reportedly a fast reaction while the second step of dissolution (reaction 224) is a slow reaction that controls the rate of chalcocite dissolution. In the study by Hiroyoshi et al. [2004], the reason for considering a direct dissolution path for chalcocite to copper ions via reactions 2-21 and 2-22, without formation of covellite intermediate, has not been explained.  25  \u000CAnother published EIS study on the dissolution of chalcopyrite electrodes is by Velasquez et al. [2005] which followed from their previous study [Velasquez et al., 1998] in a pH 9.2 solution. Figure 2-14 shows the experimentally obtained and simulated Nyquist plots at different potentials and the applied electrical circuit models are presented in Figure 2-15. The circuit for the potential of \u2013 0.150 V vs. SCE (0.092 vs. SHE) appears to be related to three relaxation processes (Figure 2-15a): Re is the ohmic resistance; Rct is the charge transfer resistance; Qc corresponds to the capacitance of a surface layer (which forms spontaneously during the polishing of the electrode). Rp is the pore resistance of this film, Qdl is the double-layer capacitance, and Rsl\u2013Qsl refers to the adsorbed species (e.g. iron species). Qc, Qsl, and Qdl were modeled by a CPE; CPE behaves as a capacitance, which varies with the frequency. For the potential of 0.20 V vs. SCE (0.442 vs. SHE), the equivalent circuit is shown in Figure 2-15b. This circuit has been proposed to be related to two Faradaic relaxation processes. The change in the equivalent circuit is associated with a decrease of Qc and Rct values together and with an important increase in the Rp value, which shows that the kinetics of the electrochemical process and the thickness and porosity of the film on the electrode surface are increased. For potentials of 0.45 and 0.60 V vs. SCE (0.692 and 0.842 vs. SHE, respectively), considerable changes were observed in the Nyquist plots, showing two loops: one at high frequency which is associated with a capacitive loop, and one at low frequencies which is a pseudo\u2013inductive loop. The equivalent circuit for both potentials was considered to be composed of two sub\u2013circuits, as shown in Figure 2-15c: one forms from the association parallel to a CPE, with Qdl/Rct pair representing the double-layer capacitance and the resistance of the electron transfer of the electrochemical process, and the other Rsl/L pair, corresponds to the resistance associated to the accumulation of superficial species, and an inductive element. The element L may be due to the transfer of ions across the double layer or an increase of the area of the electrode by the partial dissolution of the surface film. Partial dissolution of the film will result in the formation of islands that partially cover the electrode surface [Velasquez et al., 2005].  26  \u000CFigure 2-14: Measured and simulated Nyquist plots at; (a) \u2013 0.15 V, (b) 0.2 V, (c) 0.45 V and (d) 0.6 V vs. SCE [Velasquez et al., 2005].  Figure 2-15: Equivalent circuits for; (a) \u2013 0.15 V, (b) 0.2 V, (c) 0.45 V and 0.6 V vs. SCE [Velasquez et al., 2005].  2.3.4 Cyclic voltammetry (CV) studies Without a doubt, the most popular electrochemical technique to study chalcopyrite dissolution has been cyclic voltammetry. Since CV is a very powerful electrochemical 27  \u000Ctechnique for identifying reactions, almost all researchers who have studied chalcopyrite electrochemistry have availed themselves of this method. The pioneers in this area are Biegler and Swift [1979], who published a study on the anodic behavior of natural chalcopyrite in 1 M H2SO4 and 1 M HC1 by linear sweep voltammetry. The obtained voltammograms with a fresh electrode surface showed a small prewave, attributed to a surface oxidation process, followed by a region of active dissolution characterized by a steeply rising anodic current (Figure 2-16). The current in Figure 2-16 has been drawn on a logarithmic scale to show features over a broad range. From the OC potential (around 0.4 V vs. SHE) to 0.85 V there was a wide prewave in which the current passed through a peak. If the electrode potential was cycled, this prewave was absent from the second and subsequent scans. Hence, the prewave related to a surface oxidation process which involved a thin layer of chalcopyrite. The oxidation process consumes a charge of around 1 mC cm-2 and it has been suggested that this surface reaction involves the selective dissolution of iron [Biegler and Swift, 1979]. With the same approach, Parker et al. [1981] showed that the surface film in the prewave region (in 1 M H2SO4) is a metal deficient polysulfide, with semiconductive behavior, and forms on chalcopyrite during oxidation. Also, the film is thermally unstable (higher temperatures will decompose the film and faster oxidation of chalcopyrite will be achieved). This latter finding is inconsistent with the results of Viramontes-Gamboa et al. [2006] where they detected higher stability for the passive film as temperature was increased (refer to Figure 2-9). Moreover, Parker et al. [1981] have shown that the metal deficient polysulfide surface film slows transport of ions, slows electron transfer to oxidants and slows the supply of holes to the solid-solution interface. Furthermore, it has been mentioned by Parker et al. [1981] that the sulfur film which forms on corroded chalcopyrite does not account for the electrochemical or kinetic aspects of chalcopyrite leaching. The same results regarding the sulfur film were obtained by McMillan et al. [1982].  28  \u000CFigure 2-16: First positive going current\u2013voltage sweep on a freshly polished chalcopyrite surface [Biegler and Swift, 1979].  Under nitrogen sparged conditions, Biegler and Horne [1985] suggested that the prewave peak is not a consequence of atmospheric oxidation during surface preparation. Also, it has been shown that the accumulation of the products of the initial voltammetry cycle causes diminishing decrease in the size of the prewave current at more cycles [Biegler and Horne 1985]. The passed charge for the initial formation of the film is measured to be 2 mC cm-2, while previously, in the former study of Biegler and Swift [1979], it was reported to be 1 mC cm-2. Also these researchers showed that the surface film is stable at 25 \u00B0C for long periods at open circuit and the film is not stable at high temperatures (&gt; 80 \u00B0C). Through study of the relative charge passed during film formation, the electrochemical reduction of the surface film and the presence of dissolved copper in solution, Biegler and Swift [1979] came to the following suggested stoichiometry for the prewave reaction: CuFeS 2 \uF0AE 0.75CuS \uF02B 0.2Cu 2\uF02B \uF02B Fe 2\uF02B \uF02B 1.25S \uF02B 2.5e \uF02D  2-25  where, the product layer of elemental sulfur and the CuS is about 3 nm thick. Results of rotating ring disk electrode (RRDE) studies showed that the prewave reaction caused a partial dissolution of the chalcopyrite with iron being more readily dissolved in solution than copper. The ratio of the Fe2+ to Cu2+ released by oxidation was roughly 5:1 [Holliday and Richmond, 1990], while the rotating disk electrode (RDE) studies  29  \u000Cpreviously had showed it to be 4:1 [Biegler and Horne, 1985]. The overall oxidation reaction in the prewave region resulting from RRDE studies is suggested to be: CuFeS 2 \uF0AE Cu x S 2 \uF02B yCu 2 \uF02B \uF02B Fe 2\uF02B \uF02B 2(2 \uF02D x)e \uF02D  2-26  where x, the portion of oxidized Cu2+ is 0.8 and y, the solvated copper is 0.2 [Holliday and Richmond, 1990]. The proposed reaction is to some extent similar to that suggested by Biegler and Horne [1985]. Parker et al. [1981] and Biegler and Horne [1985] were of the opinion that the prewave peak was not a consequence of atmospheric oxidation during surface preparation. The prewave peak reaction was considered to be associated with the reaction of a clean chalcopyrite surface rather than reaction with atmospheric oxygen prior to voltammetry. However, the atmospheric oxidation studies of chalcopyrite by Yin et al. [1995] showed the opposite. In their study on a natural chalcopyrite electrode in 0.1 M HCl solution, it was shown that exposing a chalcopyrite electrode to the atmosphere will oxidize the surface and a passive film will form which diminishes the extent of the subsequent electrochemical oxidation in the prewave region. Calculations elucidated that after 64 h of exposure to the atmosphere, a passive film of approximately 4 nm forms on the surface. To form such a film, about seven monolayers of chalcopyrite must be oxidized. Also the effect of immersion of the electrode in hydrochloric acid on the prewave reaction has been studied by the same researchers [Yin et al., 1995]. Immersing the electrode for 18 hr in 1 M HCl solution saturated with oxygen-free nitrogen causes a larger prewave rather than immersing in the same solution for less than 1 min. However, immersing the electrode in HCl solution saturated with air for 43 hr will diminish the prewave current. Therefore, it was concluded that oxygen is the cause for passivation of the chalcopyrite surface, whereas, hydrochloric acid can activate the surface oxidation [Yin et al., 1995]. These same authors also studied the relationship between the solution pH and the prewave. The relationship was reported to be quite complex where, in acidic (pH &lt; 2) and basic (pH &gt; 9) electrolytes the prewave is easily distinguishable from the main anodic peak of higher current (and higher potential), and at pH between 2 to 9 the prewave is difficult to differentiate from the main anodic peak [Yin et al., 1995]. In addition to the previous study, the general surface oxidation behavior of natural chalcopyrite in basic solutions (pH 9.2 and 12.7) has been studied by Yin et al. [2000]. Their 30  \u000Cresults showed that the oxidation process consists of three potential and pH dependent stages. By increasing the electrode potential from \u2013 0.6 to 0.0 V vs. SCE (\u2013 0.358 to 0.242 vs. SHE) the iron in the top layer of the chalcopyrite surface was oxidized. However, the copper and sulfur atoms remained unoxidized and the surface of the electrode was covered by a metastable CuS2* phase. CuS2*, together with the oxidized iron layer formed a passivating layer on the electrode. As the potential increased further, deeper layers were involved in the oxidation, which caused a thickening of the passivating film. At this stage, the oxidation process was under solid-state mass transport control. A further increase of the potential \u2013 &gt; 0.4 V vs. SCE (&gt; 0.642 vs. SHE) \u2013 decomposed the passive film and the dissolution rate was accelerated. It is also mentioned that higher concentrations of OH\u203E enhance the dissolution rate [Yin et al., 2000]. The effect of chloride ion and temperature has been studied in 1 M H2SO4 solution by CV on carbon paste-chalcopyrite electrodes [Lu et al., 2000]. From this work, it has been shown that the presence of the chloride ion and higher temperature significantly enhance the oxidation rate of chalcopyrite. The observed higher dissolution rate of chalcopyrite in the presence of chloride is thought to be related to the aggressive nature of halides in corrosion processes [Lu et al., 2000]. Through a ring\u2013disk electrode study on a natural chalcopyrite electrode in 0.1 M H2SO4 solution at 60 \u00B0C, Lazaro and Nicol [2006] proposed the formation of soluble sulfur species such as thiosulfate. In order to investigate the effect of potential on the dissolution process, different potentials within the non\u2013oxidative and oxidative region (0.45\u20130.65 V vs. SHE) were applied on the chalcopyrite disk for 1 hr and anodic stripping curves for the platinum ring (held at \u20130.06 V) were obtained (Figure 2-17). In addition to the peak for the stripping of copper metal (peak A1 in Figure 2-17), the appearance of peaks (A2 and A3) is noteworthy. Some interesting features of these peaks are that A2 appears at all potentials and the magnitude of A2 does not seem to be dependent on potential during the prior potentiostatic period while the magnitude of peaks A1 and A3 do vary with this potential. These additional peaks (A2 and A3) are observed in a potential region similar to those which are known to be associated with the oxidation of the copper sulfides. Since different copper sulfides can be oxidized in this potential region it is not easy to assign these peaks to a particular copper sulfide (e.g. CuS, Cu2S and CuxS). As a result, if these peaks are 31  \u000Crepresentative of the oxidation of copper sulfides, then in addition to copper ions, soluble sulfur species such as thiosulfate ( S2 O32\uF02D ) and/or hydrogen sulfide (H2S) could be formed in the dissolution reaction at the disk [Lazaro and Nicol, 2006]:  CuFeS2 \uF02B 3H 2O \uF03D Cu 2\uF02B \uF02B Fe 2\uF02B \uF02B S2O32\uF02D \uF02B 6H \uF02B \uF02B 8e\uF02D  2-27  Recently the interaction between water molecules and the (001) plane of chalcopyrite has been studied by means of density functional modeling [Ferreira de Lima et al., 2011]. This work has suggested that hydrogen sulfide gas is evolved at anodic potentials. These authors show that in a highly acidic solution, protons adsorb to the surface of the disulfides ( S22\uF02D species) and this will cause elongation of the S\u2013S bond and formation of H2S. Indeed, it has been acknowledged in the literature that the sulfur dimers ( S22\uF02D species) are the precursor for polysulfide formation and that polysulfide, with time, will transform to elemental sulfur [Acres et al., 2011; Hackl et al., 1995]. In the study by Ferreira de Lima et al., [2011] it has been concluded that adsorption of positively charged protons to the sulfur dimers inhibits the initiation of the polysulfide formation, and in return H2S forms.  Figure 2-17: Anodic stripping curves for a platinum ring held at \u2013 0.06 V after dissolution of a CuFeS2 disk which was held at potentials of (a) 0.45, (b) 0.49, (c) 0.54, (d) 0.59 and (e) 0.65 V vs. SHE for 1 h at 200 rpm [Lazaro and Nicol, 2006].  Nava et al. [2006; 2008] have done an extended electrochemical study (mostly by CV technique) on the characterization of the chemical species that form during the electrochemical anodic and cathodic dissolution of carbon paste chalcopyrite electrodes in 1.7 M H2SO4 solution. They detected dissolved copper ions using a mercury film electrode (MFE) and anodic stripping voltammetry of the carbon disk. Also they analyzed the surface 32  \u000Cof the chalcopyrite electrode before and after the dissolution by cyclic voltammetry. Their results demonstrated different electrode behavior that had not been previously reported. They suggest that there are at least two types of non-stoichiometric sulfides: the first one forms at E \u2264 1.015 V vs. SHE with passive behavior (Cu1\u2212rFe1\u2212sS2\u2212t), whereas the second one forms at 1.015 &lt; E \u2264 1.085 V with an active character (Cu1\u2212xFe1\u2212yS2\u2212z); the formation of only one nonstoichiometric sulfide was recognized previously [Hackl et al., 1995; Parker et al., 1981; Elsherief, 2002]. Consequently, Nava and Gonzalez [2006] suggest that in the higher potential region (1.015 &lt; E \u2264 1.085 V) partial oxidation of chalcopyrite occurs through the dissolution of porous layers of non-stoichiometric polysulfide (Cu1\u2212xFe1\u2212yS2\u2212z).  2.4 Surface analysis of chalcopyrite The term \u201Cpassivation\u201D is widely used for chalcopyrite dissolution in different literature [Viramontes-Gamboa et al., 2007; Vel\u00E1squez et al., 2005; Mikhlin et al., 2004; Abed and Dreisinger, 2003; Warren et al., 1982; Hackl et al., 1995]. The passive state is defined in Shreir\u2019s \u201CCorrosion\u201D book [2000] as a state of low corrosion rate usually due to the presence of an (oxide) film on the surface. The typical thickness for a passive film is about a few nanometers. Such passive films (ultra thin films) would normally be directly detectable by sensitive surface analysis methods such as X-ray photoelectron spectroscopy (XPS/ESCA), Auger electron microscopy (AES), or secondary ion mass spectroscopy (SIMS) [Marcus and Mansfeld, 2006]. Generally, different in-situ electrochemical techniques (e.g. cyclic voltammetry, potentiodynamic and potentiostatic polarization, electrochemical impedance spectroscopy, and chronoamperometry) and ex-situ spectroscopic methods (e.g. X-ray photoelectron spectroscopy, auger electron spectroscopy, Raman spectroscopy) have been applied to understand the dissolution mechanism of chalcopyrite and furthermore to characterize the surface layers on electrodes, but there remains uncertainty about the dissolution process. The trouble with most of the DC electrochemical techniques used so far is that only circumstantial evidence is derivable to help with the elucidation of possible reaction mechanisms. The ex-situ analytical methods are plagued with controversy associated with the possible evaporation of elemental sulfur under experimental conditions [Buckley and Woods, 1994]. Nevertheless, Klauber et al. [Klauber et al., 2001] have pointed out that there is a significant difference between the loss of multiple layers of elemental 33  \u000Csulfur and the loss of a monolayer of elemental sulfur. In other words, the adsorption energy of sulfur atoms interacting with a sulfur mono(or multi) layer is expected to be less that the adsorption energy of adsorbed sulfur to the surface of chalcopyrite or another leached mineral. In order to justify this reasoning, Klauber et al. [Klauber et al., 2001] have compared the situation with that of the adsorbed water to the surface of a material; adsorbed water molecules are easily detectable with XPS (considering the point that water has a much higher vapor pressure than sulfur). In this section we will consider some of the previous analytical work on the surface chemistry of chalcopyrite. Hackl et al. [1995] provided a brief introduction to XPS studies of chalcopyrite from 1977 to 1995. Generally, the studies supported the electrochemical studies with regards to the formation of copper-rich sulfide layer(s) during the dissolution of chalcopyrite. Hackl et al. [1995] provided narrow scan Cu 2p, Fe 2p, and S 2p XPS spectra for the unleached chalcopyrite as well as leached chalcopyrite at 110 \u00B0C and 1.38 MPa oxygen for 3 hr. The results demonstrated that the unleached sample does not have any detectable oxidized copper and the leached sample is free of Cu2+ species. In the case of iron, even exposing the sample for a very short time in air caused formation of significant quantity of iron oxide (possible species are Fe2O3, Fe3O4 and FeOOH) on the surface. On the other hand, the analysis on the surface of the leached sample illustrated that the leaching causes dissolution of the iron oxide species. Analyzing the S 2p spectra, Hackl et al. [1995] detected monosulfide (S2-) referring to unaltered CuFeS2, disulfide ( S22\uF02D ) which was thought to be an iron deficient sulfide with the composition of Cu1-xFe1-yS2 (y\u00BBx and x+y \u22481) and polysulfide ( S2n\uF02D ) which was assumed to be CuSn (n &gt; 2). The precise value of n was not defined in the study [Hackl et al., 1995]. Yin et al. [1995] identified the oxidation products of natural chalcopyrite on a freshly cleaved surface as well as after electrochemical dissolution in the potential region of &gt; 0.5 V vs. SCE in 1 M HClO4 solution. The Cu 2p spectra are almost identical in binding energy. It is proposed by the authors that the CuO, Cu2O and Cu2S cannot be formed by oxidation of chalcopyrite in HClO4 solution, and they suggested the formation of CuS and S or CuS2* [Yin et al., 1995]. Their study also confirms the results of Hackl et al. [1995] on the dissolution of iron species during oxidation. In addition to their study in HClO4 solution, Yin et al. [2000] also studied the formation of species on the surface of a natural chalcopyrite electrode in alkaline solutions of 0.1 M borax (pH 9.2) and 0.05 M NaOH (pH 12.7). In the 34  \u000Ccase of the borax solution, they have observed more oxidized iron at higher anodic potential of 0.4 V vs. SCE (0.642 vs. SHE) rather than at 0.15 V (0.392 vs. SHE). Also the formation of Fe(OH)3 and Fe2O3 species have been suggested. Furthermore, it has been mentioned that argon ion bombardment may cause the formation of FeO species during the XPS test. Yin et al. [2000] have proposed that CuO is the only oxidized form of copper and that the identical peak of Cu2O originates from the reduction of CuO to Cu2O by argon ion bombardment in the XPS chamber. So, According to Yin et al. [2000], the existence of the Cu2O species is related to the reduction of CuO during the XPS analysis and not to the oxidation of copper during the electrochemical dissolution process. Moreover, CuFe1-xS2 and CuS2* were identified and both had a distorted chalcopyrite structure. In other words, these species form by dissolution of iron atoms from the chalcopyrite crystal structure. The induced distortion in the chalcopyrite crystal structure is due to the vacancies introduced in the iron atom positions during the dissolution process. In NaOH solution (pH 12.7), in addition to the abovementioned species, the formation of CuO and CuSO4 was reported at 0.8 V vs. SCE (1.042 vs. SHE). The results of this work by Yin et al. at pH 9.2 are in agreement with the work of others [Vel\u00E1squez et al., 1998; Vel\u00E1squez et al., 2005]. Farquhar et al. [2003] have studied the electrochemical dissolution of natural chalcopyrite samples in 0.1 M NaNO3 solution (pH 4). Their results are very close to those of the previous studies. However, it was mentioned that after dissolution at 0.5 V vs. Ag/AgCl (0.697 vs. SHE) the sulfide or natural sulfur concentrations gradually increase with depth below the surface and at approximately 40 \u00C5 all the sulfur is present as sulfide or elemental sulfur. Parker et al. [2003] have studied species on the surface of chalcopyrite after leaching in four different sulfuric acid solutions (starting pH values of 1.3 to 1.9) in the presence of: (1) ferric, (2) ferrous, (3) ferric plus silver and (4) microbes. Table 2-2 summarizes the XPS results that highlight the different species found in each case. According to Table 2-2, the species formed in the presence of ferric and ferrous are almost the same while in the case of Ag+ addition, quite different surface compositions are produced. In the silver containing solution the only forms of sulfur are sulfide (S2-) and sulfate ( SO 24 \uF02D ). Chalcopyrite demonstrated a very high reactivity toward weak silver concentrations (down to 0.0002 M) whether or not ferric was present. The reaction accelerated at the lower pH of 1.9 and  35  \u000Cproduced a silver dominated surface phase of Ag 2S with several thousand \u00C5\u2019s thickness in less than 30 min [Parker et al., 2003]. Table 2-2: Various sulfur, copper, and iron species on the surface of the chalcopyrite after oxidative dissolution; between 50 to 80 \u00B0C [Parker et al., 2003].  The atomic percentage of sulfur species on the top layer of natural chalcopyrite (extracted from XPS results) as a function of dissolution time in 1 M HClO4 (pH 1) and 85 \u00B0C was studied by Harmer et al. [2006]. From this study (summarized in Table 2-3), the total concentration of sulfur species increases from 25.3% to 39.3% during the first two hours of leaching. The SO 24 \uF02D species which was present on the unleached sample was removed from the surface by leaching. The concentrations of both S22\uF02D and S2n\uF02D species was increased by leaching. The total surface concentration of S was 21.7% after 24 h of leaching and was increased dramatically to 77.3% after 144 h of leaching. Also, within this extended leaching period, a significant increase in leaching rate was reported. The S 2p XPS spectrum of the sample leached for 144 h showed a combination of S2\uF02D , S22\uF02D , S2n\uF02D , and So species on the surface, wherein, the concentration of S2\uF02D species increased from 5.9% at 24 h to 17.7% at 144 h (an almost threefold increase). The S concentration and speciation on the surfaces after 144 and 313 h of leaching were reported to be very similar. Because the concentration of S increased on the surface during the leach and because the leaching rate increased during this same period, it was concluded by Harmer et al. [2006] that neither the formation of So or S2n\uF02D is passivating.  36  \u000CTable 2-3: Atomic percentage of different sulfur species on the surface of chalcopyrite before and after leaching [Harmer et al., 2006].  2.5 Electrochemical (and chemical) dissolution reactions During the dissolution process, increasing the potential of chalcopyrite in the anodic direction will cause the formation of surface layers which hinder the dissolution rate of the mineral [Warren et al., 1982]. As mentioned in previous sections, the precise composition of the layer(s) is still in dispute. Previous studies have concluded that these layers could be sulfur [Dutrizac, 1978; Munoz et al., 1979], precipitated iron compounds [Dutrizac, 1978], a solid electrolyte interphase (SEI) [McMillan, et al., 1982], a metal deficient sulfide [Linge, 1976], or polysulfidic in nature [Parker, et al., 1981; Hackl, et al., 1995]. Studies by Biegler [1977] show nucleation and growth of a thin two-dimensional sulfide phase on the surface of chalcopyrite during the exposure of a polished sample in H2SO4 solution or air. More studies by Biegler and Swift [1979] have shown evidence for nucleation and growth of a 3dimensional phase at low anodic potentials (around +200 mV vs. OCP). The current-time behavior of the dissolution process in this potential area has the characteristics of expansion and overlap of localized dissolution centers. Anodic dissolution (around +200 mV vs. OCP) of the chalcopyrite is considered to be a bulk dissolution process which is a distinct mechanism from the surface oxidation process at lower potentials (anodic potentials close to OCP) [Biegler and Horne, 1985]. It was mentioned earlier that the surface oxidation process is detectable as a \u201Cprewave\u201D in cyclic voltammetry tests. Technically, the prewave represents the initial stage of the dissolution of chalcopyrite and furthermore, determines the surface properties of the mineral during the rest of the dissolution process. Many publications on chalcopyrite have concentrated on the dissolution behavior in the prewave region [Biegler, 1977; Biegler and Swift, 1979; Biegler and Horne, 1985]. Studies by Lazaro and Nicol [2006] on the dissolution of chalcopyrite in dilute sulfuric acid solutions have produced some 37  \u000Cevidence for the formation of Cu2+, Fe2+, and soluble sulfur species at the initial stages of dissolution at 60 \u00B0C. More kinetic studies on the prewave reaction by Holliday and Richmond [1990] have demonstrated a sequential mechanism for the surface oxidation which consists of an initial rate determining step for the production of adsorbed Cu2+ followed by the production of Fe2+ ions. These reactions are responsible for making a passive surface layer on the mineral and hindering the dissolution rate. These results are along the same lines as those presented by Biegler and Horne [Biegler and Horne, 1985]. Calculations by Biegler [1977] have shown that the thickness of the layer formed during the prewave reaction in sulfuric acid solution is around 2 nm and in another reference [Biegler and Horne, 1985] this layer is reported to be a CuS and S compound with thickness of about 3 nm. Studies by Yin et al. [1995] indicate that exposing polished chalcopyrite to air causes the formation of a passive layer with a thickness of 4 nm. Moreover, it has been shown that the anodic prewave disappears if the sample is exposed to air due to the formation of the passive layer on the surface. For a 1 M HCl solution the prewave peak appears at around 0.52 V vs. SHE and at higher potentials (around 0.8 V vs. SHE) the second anodic peak appears. The main anodic reactions proposed by these researchers in 1 M HCl solution are [Yin, et al., 1995]: CuFeS 2 \uF0AC \uF0BE\uF0AE Cu 2\uF02B \uF02B Fe 3\uF02B \uF02B 2S \uF02B 5e \uF02D  2-28  and CuFeS 2 \uF02B 8H 2 O \uF0AC \uF0BE\uF0AE Cu 2\uF02B \uF02B Fe 3\uF02B \uF02B 2SO 24\uF02D \uF02B 16H \uF02B \uF02B 17e \uF02D  2-29  Nava et al. [2006; 2008] have studied the dissolution of a carbon paste/chalcopyrite electrode in sulfuric acid solution. It was indicated that the dissolution of chalcopyrite starts at low anodic potentials (0.615 \u2264 E &lt; 1.015 V vs. SHE) and results in the formation of a nonstoichiometric polysulfide layer of composition Cu1-xFe1-yS2-z according to the following reaction:  CuFeS2 \uF02B 4zH 2O \uF0AC \uF0BE\uF0AE Cu 1\uF02Dx Fe1\uF02D yS2\uF02Dz \uF02B xCu 2\uF02B \uF02B yFe 3\uF02B \uF02B zSO 24\uF02D \uF02B 8zH \uF02B \uF02B (2x \uF02B 3y \uF02B 6z)e \uF02D  2-30  In this potential region the concentration of the dissolved copper ion is very low and the anodic stripping voltammetry (ASV), which is known to have a very good detection limit, was not able to detect the copper ion in solution [Nava and Gonzalez, 2006]. Consequently, these authors have concluded that Cu1-xFe1-yS2-z is a passive film which slows down the 38  \u000Cdissolution rate of chalcopyrite and keeps the concentration of copper ion in solution to a very low value. At higher potentials (1.015 \u2264 E &lt; 1.085 V vs. SHE), a small amount of copper ion was detected in solution which confirmed the partial dissolution of the polysulfide layer. The reaction for chalcopyrite dissolution in this higher potential range was proposed to be:  CuFeS2 \uF02B 4pH2O \uF0AC \uF0BE\uF0AE Cu1\uF02Dm Fe1\uF02DnS2\uF02Dp \uF02B mCu2\uF02B \uF02B nFe3\uF02B \uF02B pSO 24\uF02D \uF02B 8pH \uF02B \uF02B (2m \uF02B 3n \uF02B 6p)e \uF02D  2-31  While reactions 2-30 and 2-31 are similar, it was concluded by Nava and Gonzalez [Nava and Gonzalez, 2006] that the sulfide product in reaction 2-31, i.e. Cu1-mFe1-nS2-p, is active as opposed to the passive layer formed in reaction 2-30, i.e. Cu1-xFe1-yS2-z. The dissolution rate of chalcopyrite was found to be a function of the decomposition rate of the layer obtained through reaction 2-31. Increasing the potential into the range of 1.085 \u2264 E &lt; 1.165 V vs. SHE caused the complete release of iron from the lattice of chalcopyrite into the solution and the formation of a CuS layer on the surface of the electrode according to reaction 2-32:  2CuFeS2 \uF02B 13H2 O \uF0AC \uF0BE\uF0AE 0.75CuS\uF02B 1.25Cu2\uF02B \uF02B Fe 2 (SO4 )3 \uF02B 0.25SO 24\uF02D \uF02B 26H \uF02B \uF02B 28e \uF02D  2-32  The CuS layer is not stable at potentials higher than 1.165 V vs. SHE and so the complete dissolution of chalcopyrite will occur by further increasing the potential, according to reactions 2-33 and 2-34: CuS \uF02B 4H 2 O \uF0AC \uF0BE\uF0AE Cu 2\uF02B \uF02B SO 24\uF02D \uF02B 8H \uF02B \uF02B 8e \uF02D  2-33  2CuFeS2 \uF02B 16H2O \uF0AC \uF0BE\uF0AE 2Cu 2\uF02B \uF02B Fe 2 (SO 4 )3 \uF02B SO 24\uF02D \uF02B 32H\uF02B \uF02B 34e\uF02D  2-34  At high potentials, the formation of jarosite on the surface of chalcopyrite was found to decrease the dissolution rate of the electrode. In fact, the products of reaction 2-34 result in the formation of jarosite on the electrode surface [Nava and Gonzalez, 2006; Nava, et al., 2008]. This dissolution model is similar to that of Hackl et al. [1995]. In alkaline solutions, researchers [Yin, et al., 2000; Vaughan, et al., 1997] have observed three different stages of oxidation. The prewave occurs at low potentials (around 0.26 V vs. SHE) which they considered to be the top atomic layer oxidation of iron to a monolayer of Fe(OH)3 and Fe2O3, while the Cu and S of the chalcopyrite structure remained 39  \u000Cunchanged as CuS2*. The layer consisting of oxidized iron and CuS2* acted as a passive layer which inhibits further oxidation of chalcopyrite. At higher potentials (around 0.64 V vs. SHE) the same reactions were found to occur and iron atoms located deeper within the electrode surface were found to take part in the oxidation reaction. At this potential the passive film was stable and thickened with time. The electrochemical reactions under these conditions were considered to be solid-state mass transport-controlled. For higher potentials (&gt; 0.64 V vs. SHE) the CuS2* layer oxidized to CuO, and SO 24 \uF02D which accelerated the rate of the total electrode dissolution. Studies by Velasquez et al. [2005] confirm these results except that they were not able to detect the formation of elemental sulfur due to possible evaporation of this element under the experimental conditions of XPS. Studies by electrochemical impedance spectroscopy (EIS) have shown the formation of islands of passive layers on the surface of chalcopyrite electrodes [Vel\u00E1squez et al., 2005]. Another EIS study has demonstrated the active-passive behavior of chalcopyrite but did not point out the passive islands on the surface of electrode [Viramontes-Gamboa et al., 2006; Hiroyoshi et al., 2004; Velasquez et al., 1998].  2.6 What is missing? As one can see from the literature, a fully accepted mechanism for the dissolution of chalcopyrite is missing and the passive layer(s) which form on the surface of this mineral result in a very slow dissolution and thus a practical engineering challenge for the hydrometallurgist. Furthermore, from literature, the properties of the chalcopyrite passive film are not clear. It is not known whether mass transport through the passive film (i.e. a diffusion process), dissolution of the passive film itself (i.e. reaction) or electron transfer through the passive film is controlling the chalcopyrite dissolution rate in the leaching solution. In this study, the semiconductor, electrochemical and chemical properties of chalcopyrite\u2019s passive film will be studied. Additionally, the electronic and ionic conductivity of chalcopyrite\u2019s passive film is a crucial missing point. Indeed, to interpret many of the phenomena currently being observed in chalcopyrite leaching studies, it is important to understand the extent to which chalcopyrite\u2019s passive film conducts electrons or ions. For example, during development of leaching processes, the answer to this question 40  \u000Cwill help us to decide if the electron transfer rate of the electrochemical reactions should be increased or the rate of ion transfer through the passive film should be promoted.  41  \u000C3 Objectives The ultimate goal of any electrochemical study pertaining to the dissolution of chalcopyrite is to overcome its slow dissolution rate. Clearly, the path to this goal runs through obtaining comprehensive knowledge of the surface chemistry of chalcopyrite in the desired dissolution environment. Thus, the proposed research seeks the following objectives: \uF0B7  To investigate the electrochemical dissolution mechanism of pyrrhotite (Fe1-xS).  Pyrrhotite electrochemistry may be considered to represent a simplified case of chalcopyrite electrochemistry; as pyrrhotite is a Fe-S binary system (metal to sulfur ratio of c.a. 1) while chalcopyrite is the more complex ternary system of Cu-Fe-S (metal to sulfur ratio of 1). Hence, the results obtained regarding the dissolution mechanism of pyrrhotite may be applicable to chalcopyrite dissolution. Additionally, it must be mentioned that the dissolution mechanism of both pyrrhotite and chalcopyrite minerals are very similar as they both passivate in acidic solutions by formation of a metal-deficient sulfide film (this will be further proven in the next few chapters). However, other Fe-S minerals (e.g. pyrite, FeS2) dissolve through different mechanisms. For instance, pyrite leaching proceeds by simultaneous dissolution of Fe and S and no passive film forms on the surface. Hence, we believe that the electrochemical dissolution of pyrrhotite is similar to that of chalcopyrite. \uF0B7  To understand the nature of the surface layers (mainly the passive layers) which form  at low anodic potential (&lt; 1 V vs. SHE) on the surface of chalcopyrite through careful characterization of said layers to obtain their composition. \uF0B7  To offer a comprehensive dissolution reaction scheme for chalcopyrite in sulfuric  acid solutions which takes into account all of the observed critical characteristics of chalcopyrite dissolution. \uF0B7  To investigate the physical properties of the chalcopyrite passive film. The  semiconducting properties of the passive film and its electrical conductivity will be studied.  42  \u000C4 Experimental 4.1 Minerals used High quality natural polycrystalline pyrrhotite (Fe1-xS) and chalcopyrite (CuFeS2) samples originating from the Creighton Mine (Sudbury, Ontario) were used for electrode fabrication. In some parts of the study, other iron sulfide and copper sulfide minerals (i.e. pyrite, pyrrhotite, chalcocite and covellite) were used for comparison purposes. The composition of the minerals, obtained by Rietveld method and X-ray powder diffraction, and associated reference number to each mineral (which will be used during the course of study), are presented in Table 4-1. Table 4-1: Composition of the as received sulfide minerals used in this research. # of mineral Mineral 1 (chalcopyrite) 2 (pyrite) 3 (pyrrhotite) 4 (chalcocite) 5 (covellite) Chalcopyrite (CuFeS2) 98.1 Pyrite (FeS2) 97.6 Pyrrhotite (Fe1-xS) 93.7 Chalcocite (Cu2S) 73.7 Djurleite (Cu31S16) 26.3 Covellite (CuS) 98.3 Digenite (Cu9S5) 1.7 Magnetite (Fe3O4) 6.3 Siderite (FeCO3) 1.9 Quartz (SiO2) 0.6 Anhydrite (CaSO4) 1.8 -  4.2 Electrode preparation Massive electrodes were prepared by cutting the minerals into approximately cubic shapes with areas of 1 cm2 exposed to the solution. In order to have a uniform current distribution on the samples, a thin layer of copper was electrodeposited on the back of the electrodes (the side to which the electrical connection was attached). Copper was electrodeposited from a copper sulfate solution with low current density of 4 mA cm-2. An electrical contact with copper wires was made in the back of the samples with silver-loaded epoxy resin (Mg Chemicals). The electrode contacts were protected from the solution by 43  \u000Cmounting the whole assembly in epoxy resin (LECO), leaving only one face of the electrode exposed to the solution. Prior to each electrochemical experiment the electrodes were checked to ensure that there was no evidence of crevice corrosion in the electrode. This was possible due to the transparent epoxy resin mounts. Before each test, the exposed face of the electrode was polished with brand new carbide papers, starting from no. 300 paper up to no. 1200, to begin each experiment with a fresh surface. Afterwards, the surface of the electrodes was checked with optical microscope in order to make sure that the morphology of the electrodes was similar. The electrodes were then washed with DI water, sonicated in DI water, rinsed with DI water, and then finally dried in a stream of air.  4.3 Solution preparation The goal of this study was to understand the fundamentals of chalcopyrite and pyrrhotite electrochemical dissolution. Thus, the electrolytes for electrochemical experiments were prepared by addition of only acids into deionized (DI) water. Solutions were prepared by dissolving reagent grade hydrochloric acid (HCl, Fisher Scientific) or sulfuric acid (H2SO4, Fisher Scientific) in DI water. The solution used for all of the experiments pertaining to pyrrhotite was 1 M HCl (pH within 0.2 \u2013 0.3) and for chalcopyrite study it was 0.5 M sulfuric acid (pH ~ 0.3). In all of the experiments the oxygen content of the solution was controlled. In some of the experiments the oxygen in solution was removed to less than 10 ppb by sparging pure argon gas through the solution. In these experiments, the argon gas was sparged for 30 min prior to the experiment and sparging was continued during the test. When the experiment was performed in quiescent conditions the argon sparging was stopped during the experiment. All other experiments were performed in aerated solution by sparging filtered air into the cell. The volume of solution for all of the electrochemical experiments was 800 mL of fresh electrolyte.  4.4 Electrochemical experiments All electrochemical experiments were performed at 25 \u00BAC (this temperature was selected because most of the available data in literature are extracted from experiments at 25 \u00BAC) using a VersaSTAT 4 potentiostat/galvanostat controlled by VersaStudio software (version 1.51, Princeton Applied Research). Some of the electrochemical data were obtained 44  \u000Cusing of a PAR 273 potentiostat-galvanostat connected to a PAR 1025 frequency response detector controlled by PowerCorr, PowerStep, and PowerSine software (Princeton Applied Research). The ZSimpWin 3.20 (2004) software was used for fitting the impedance data. A conventional three-electrode electrolytic cell was used for electrochemical analysis. The test cell was jacketed and a silicon fluid was circulated through the jacket to maintain the desired test temperature of 25 \u00BAC. The working electrode was always the prepared mineral electrode (electrodes are presented in Table 4-1). Opposite to the working electrode, a graphite electrode (99.999%, Alfa Aesar) was used as counter electrode. Three different reference electrodes, based on availability in the laboratory, were used. The first one was a glass body mercury/mercurous sulfate electrode (MSE, 0.600V vs. SHE at 25 \u00BAC) filled with saturated K2SO4 electrolyte (Princeton Applied Research). The second electrode was a glass body Ag/AgCl reference electrode filled with 4 M KCl electrolyte saturated with AgCl (Accumet 13-620-523), 199 mV vs SHE at 25 \u00BAC. Lastly, a glass body saturated calomel reference electrode (Accumet 13-620-525), 242 mV vs SHE at 25 \u00BAC. The reference electrode was connected to a Luggin capillary to minimize IR drop. The electrolyte inside the reference electrodes was refreshed after each 20 hours of use and the electrodes were tested against a lab-standard Ag/AgCl electrode before each experiment. The electrode potential was allowed to stabilize for 6 hours before starting the measurements and, in order to ensure reproducibility, tests were begun only after reaching the same OCP within \u00B15 mV. The potentiodynamic polarization experiments were conducted at potential sweep rates between 0.05 to 50 mV s-1. The potentiostatic polarization experiments were performed by application of a constant potential and the current flow through the sample was registered. In addition to the above DC tests, cyclic voltammetry was applied to study the electrodissolution behavior of chalcopyrite. The sweep rate was 50 mV/s for all CV experiments, as the essential features were evident at these rates. Electrochemical impedance spectroscopy tests were conducted at different potentials and in the frequency range of 10-2\u2013105 Hz with peak-to-peak amplitude of 10 mV. Capacitance measurements were carried out at frequencies of 1, 10, and 50 kHz with potential amplitude of 10 mV peak to peak. Each electrochemical experiment was repeated at least twice to insure the reproducibility of results.  45  \u000C4.5 X-ray photoelectron experiments In order to investigate the chemical composition of the surface films, the chalcopyrite electrodes were polarized at different constant potentials of 450, 650, 700, 870, 900, 950 and 1000 mV (vs. SHE) and then analyzed using X-ray photoelectron spectroscopy. For the sake of comparison, the surface composition of a polished sample of chalcopyrite with no exposure to solution and that of a sample immersed in the electrolyte for 6 hours were also analyzed. In order to minimize surface modifications after polarization, the chalcopyrite electrodes were rinsed with distilled water to remove residual acid and then transferred immediately to an evacuated desiccator and stored before transfer into the XPS vacuum. XPS was conducted with an Axis Ultra DLD spectrometer (Kratos Analytical) with a monochromatic Al excitation (1486.7 eV) operating at 108W. For reference, the binding energy of the Au 4f7/2 line was found to be 83.88 eV with the FWHM of 1.4 eV. Binding energy calibration was based on C 1s at 284.6 eV. In addition to all of the above XPS experiments, in order to remove the passive film, only a few XPS experiments were performed after etching of the mineral with an Ar+ ion beam (1 keV beam energy). The results of these experiments were used as reference for the peak position of different species in chalcopyrite. However, a consequence of the ion beam etching might be alteration of the surface species. In order to eliminate the acquisition of invalid results, the peak positions obtained were compared with those in the literature. One parameter of particular concern for the alteration of surface species is the argon ion beam energy and this parameter should be selected carefully. Ar+ ion etching with energies higher than 1.5 keV at long durations (more than 20 min) can result in species alteration and enrichment of the metals on the surface [Otte et al., 2000; Glatzel et al., 2002]. Therefore, a beam energy of 1 keV and short etching times were applied. Additionally, in order to limit the creation of etch pits, the ion beam was rastered on the surface of chalcopyrite to etch a large surface area in comparison to the area that was used to collect the spectra. XPSPeak 4.1 software was used to fit the XPS peaks.  46  \u000C4.6 Scanning electron microscope study In order to characterize the surface morphology of the minerals and the surface layers which were formed after dissolution, a scanning electron microscope with light element EDX (Hitachi S-3000N) was used.  47  \u000C5 Electrochemical dissolution of pyrrhotite in hydrochloric acid solution1 5.1 Introduction Minerals of pyrrhotite (Fe1\u2013xS and Fe7S8) are among the major minerals found in ores of the Sudbury basin (Ontario, Canada). These minerals are usually separated from the other desired metal sulfides through ore beneficiation and as a result, a large amount of pyrrhotite tails are available in the area as waste material. The paymetal content of the disposed tailings make them an interesting mineral resource for further processing and extraction of nickel [Droppert and Shang, 1995; Ke and Li, 2006; Wang, 2008], and gold [Wang, 2008; Lehmann et al., 2000; Ellisa and Senanayake, 2004]. Therefore, studies aimed at understanding the dissolution behavior of pyrrhotite in acidic electrolytes are essential. However, the most interesting aspect of the electrodissolution study of pyrrhotite is its compositional similarity to chalcopyrite. In fact, pyrrhotite electrochemistry may be considered to represent a simplified case of chalcopyrite electrochemistry; as pyrrhotite is a Fe-S binary system (metal to sulfur ratio of c.a. 1) while chalcopyrite is the more complex ternary system of Cu-Fe-S (metal to sulfur ratio of 1). Hence, the results obtained regarding the dissolution mechanism of pyrrhotite would be applicable to chalcopyrite dissolution. Additionally, it must be mentioned that the dissolution mechanism of both pyrrhotite and chalcopyrite minerals are very similar as they both passivate in acidic solutions by formation of a metal deficient sulfide film (this will be further proven in the next few chapters). However, other Fe-S minerals (e.g. pyrite, FeS2) dissolve through different mechanisms. For instance, pyrite leaching proceeds by simultaneous dissolution of Fe and S and no passive film forms on the surface. Hence, we believe that the electrochemical dissolution of pyrrhotite is most like that of chalcopyrite. Moreover, it should be mentioned that electrochemical impedance spectroscopy (EIS) is the major technique applied in this thesis. Although EIS is a powerful method for such studies, due to the complexity of the layers which form when complex minerals are dissolved and due to the intricacy of EIS, this technique has been disregarded in 1  A version of this chapter has been published: \u201CGhahremaninezhad, A., Asselin, E., Dixon, D.G., 2010, Journal of Electrochemical Society, 157, C248\u2013C257\u201D.  48  \u000Cmineral electrochemistry. Thus application of this technique on a simpler system (i.e. Fe-S) before the more complex chalcopyrite system seemed essential. Thermodynamic data (via E-pH diagrams) indicate that pyrrhotite will dissolve at pH less than 2 [Hamilton and Woods, 1981; Biernat and Robins, 1972]. Several workers have suggested that main issue in the acidic dissolution of pyrrhotite is the formation of an impermeable layer of elemental sulfur around the pyrrhotite particles which hinders their dissolution [Filippou et al., 1997; Beolchini and Veglio, 1999]. In order to increase the leaching kinetics of pyrrhotite, characterization of the layers formed during dissolution is essential and electrochemistry is a powerful method to this end. Cyclic voltammetry studies on pyrrhotite carbon paste electrodes by Almeida and Giannetti [2003] in acetic acid/sodium acetate solution (pH 4.5) show three anodic peaks above the open circuit potential (EOC). These results are similar to those of Hamilton and Woods [1981] and Mikhlin [2000]. The first anodic peak is considered to be the formation of elemental sulfur: Fe1-x S \uF0BE \uF0BE\uF0AE (1 - x) Fe 2\uF02B \uF02B S \uF02B 2 (1 - x) e -  5-1  In this step, ferrous ions dissolve and elemental sulfur accumulates on the surface of the electrode [Hamilton and Woods, 1981]. As more sulfur mounts up on the surface, it becomes more difficult for Fe2+ ions to diffuse and a high concentration of ferrous ions remains between the mineral surface and sulfur layer. Further increase in potential will result in the second peak which represents the oxidization of a portion of the trapped Fe2+ to Fe3+ [Almeida and Giannetti, 2003]. However, according to Paul et al. [1978], the accumulation of Fe2+ can reduce the proton concentration near the electrode which favors the formation of thiosulfate [Almeida and Giannetti, 2003]. While their study focused on the dissolution behavior of galena, their reaction mechanism may be extended to pyrrhotite yielding the following reactions [Almeida and Giannetti, 2003]: 2 S \uF02B 3 H 2O \uF0BE \uF0BE\uF0AE S 2 O 32- \uF02B 6 H \uF02B \uF02B 4 e -  5-2  2FeS \uF02B 3H 2O \uF0BE \uF0BE\uF0AE S2O32- \uF02B 2Fe2\uF02B \uF02B 6H\uF02B \uF02B 8e-  5-3  According to Biernat and Robins [1972] these reactions are thermodynamically feasible; however these reactions may not be possible in very low pH solutions with high proton acitivity such as in the case of the test solution used in the present study. A further increase in potential shows the third peak in a cyclic voltammogram which is possibly  49  \u000Cassociated with the oxidation of reaction-formed intermediates and/or sulfate formation [Almeida and Giannetti, 2003]:  S2 O32- \uF02B 5 H 2O \uF0BE \uF0BE\uF0AE 2 SO2-4 \uF02B 10H\uF02B \uF02B 8e-  5-4  Various studies show that, in addition to the species mentioned, the formation of different sulfur species during the electrochemical dissolution of sulfide minerals is possible [Bevilaqua et al., 2004; Viramontes-Gamboa et al., 2007; Abraitis et al., 2004; Mikhlin et al., 2004; Harmer et al., 2006]. There is not complete agreement between researchers on the reaction mechanisms and reaction products, but the final product of pyrrhotite dissolution is considered to be some ferric compounds, elemental sulfur, and sulfate according to the following overall reaction [Mikhlin 2000]: Fe1-x S \uF02B 4yH 2 O \uF0BE \uF0BE\uF0AE (1 - x)Fe 3\uF02B \uF02B y SO 2-4 \uF02B (1 - y) S \uF02B 8yH \uF02B \uF02B 3(1 - x \uF02B 2y)e -  5-5  The evolution of surface layers during the dissolution of sulfide minerals has been studied with different electrochemical and analytical methods and there is some controversy associated with the application of some ex-situ methods such as X-ray photoelectron spectroscopy [Bevilaqua et al., 2004; Mikhlin et al., 2004; Buckley and Woods, 1994]. Hence, in-situ analytical techniques are among the major methods of interest for studies related to dissolution processes. Most of the sulfide minerals are considered to be semiconductor  materials  and  therefore  applying  proper  analytical  methods  for  semiconductors would be very useful. Mott-Schottky analysis is a powerful in-situ electrochemical technique which can characterize the semiconducting behavior of surface layers formed during the dissolution of minerals. Applying this method makes it possible to understand the semiconducting behavior of the surface layers at different potentials [Morrison 1980; Memming 2001]. In this chapter, which is concerned with the electrochemical dissolution of pyrrhotite, the nature of the surface layers and their semiconducting properties at different potentials will be discussed. EIS is applied to understand the growth sequence of different surface layers. In-situ AC impedance is used to extract Mott-Schottky plots and characterize the electrolyte/semiconductor interface. Using the EIS and chronoamperometry results, a model is generated for the dissolution sequence of pyrrhotite at different potential intervals.  50  \u000C5.2 Results and discussion 5.2.1 Open circuit potential study Figure 5-1 presents five open circuit potential (OCP) vs. time curves for three different pyrrhotite electrodes. Two of the curves are repeat OCP measurements of the same sample after polishing with carbide paper. There is a slight increase during the first 20 minutes of the test before reaching steady state. The maximum potential variation at initial immersion and after 20 min immersion is \u00B1 35 mV and \u00B1 5 mV, respectively. All electrodes used in this study displayed a zero immersion time OC potential of 350 \u00B1 35 mV.  OC Potential/V vs. SHE  0.65  0.625  0.6  0.575  0.55 0  30  60  90  120  time/min  Figure 5-1: Open circuit potential (OCP) of the three different pyrrhotite electrodes vs. time; two of the curves show the OCP of same samples after polishing with carbide paper.  5.2.2 Potentiodynamic study Figure 5-2 shows the potentiodynamic curves for the electrode in the dark and under irradiation at scan rates of 0.5 and 5 mV s\u20131. Potentiodynamic polarization curves (a) and (b) are very similar, and both demonstrate a change in slope around 470 mV, indicating increasing resistance to dissolution. Afterward, a rapid increase in the dissolution rate begins in the potential range of 630 to 700 mV. Curve (c) of Figure 5-2 presents the polarization at a lower scan rate of 0.5 mV s\u20131. Generally, applying a lower scan rate will reveal a closer approximation of steady state electrode behavior. Through comparison of curves (a) and (b) 51  \u000Cin Figure 5-2, it is clear that the increased slope of the E\u2013i curve in the potential range of 470 to 600 mV is not associated with the formation of a true passive film on the surface. This will be discussed further in the next section. On the other hand, the slower scan rate plot lends support to the idea that a passive layer is formed at potentials above 700 mV. Therefore, the potentiodynamic polarization curves at different scan rates reveal the existence of two distinct regions during pyrrhotite polarization: a quasi-passive region at low potentials (470 to 600 mV) and a passive region at potentials higher than 700 mV. The nature of these layers will be discussed in more detail in the next section.  Potential/V vs. SHE  1.3 (a) 5 mV/s (Irradiadtion)  1.1  (b) 5 mV/s (Dark) (c) 0.5 mV/s (Irradiation)  0.9 0.7 0.5 0.3 -8  -7  -6  -5 -4 Log i/mAcm-2  -3  -2  Figure 5-2: Potentiodynamic curves for the pyrrhotite electrode in 1 M HCl electrolyte.  5.2.3 Electrochemical impedance spectroscopy study Figure 5-3 presents impedance spectra for the pyrrhotite electrode at different potentials. Figure 5-3a shows a response at the OC potential which is associated with mass transfer. Asselin et al [2007] have related this kind of behavior to the formation of a thin surface layer on the electrode. Perhaps, this layer was formed during sample preparation at the polishing stage or, alternatively, during the 10 min stabilization time in the electrolyte prior to each electrochemical experiment. Figure 5-3b shows the same behavior of Figure 5-3a, but with smaller total impedance at 692 mVSHE. Furthermore, in Figure 5-3b a deviation from pure 52  \u000Clinear behavior is seen as a partially resolved semicircle is formed at high frequencies. Hence, the stability of the surface layer decreases with an increase of the potential and simultaneously the effect of charge transfer in the system increases [Asselin et al., 2007]. Similar frequency response has been observed, also at low potentials, on chalcopyrite as reported by Vel\u00E1squez et al. [2005]. In Figure 5-3c, at 782 mVSHE, a partially resolved semicircle at high frequencies (normally attributed to charge transfer resistance at the solution/surface layer interface) and an inductive loop at low frequencies are apparent. The latter is characteristic of the dissolution of the passive layer through transpassive dissolution [Betova et al., 2002; Bojinov and Tezvetkoff 2003; Erne and Vanmaekelbergh 1997]. Figure 5-3d presents impedance spectra for the electrode at 917 mVSHE which contains a charge transfer related semicircle at high frequencies, an inductive loop at medium frequencies and a loop consisting of combination of a capacitance impedance and a related resistance at lower frequencies (0.1 \u2013 0.01 Hz). A comparison of Figures 5-3c and d shows that, by increasing the potential from 782 to 917 mVSHE, a secondary surface layer forms as the behavior at the lowest frequencies changes from inductive to capacitive [Bojinov et al., 2002]. Both Figures 5-3e and f demonstrate the growth of this second layer at lower frequencies through the increase in the total impedance of the system. It is worth mentioning that in Figure 5-3f there is an apparently large difference between the impedance of the first (small semicircle) and second (larger partially resolved semicircle) surface layers \u2013 a finding which will be supported with Mott-Schottky studies discussed below. Figure 5-3g presents the behavior of the electrode at 1092 mVSHE which exhibits negative real impedance at this potential and which can be attributed to the transpassive dissolution of the second surface layer. This is a finding similar to that presented by Keddam et al. [1985] for transpassive dissolution of nickel. Other researchers have further attributed these inductive loops to small, localized areas of increased anodic activity and transpassive dissolution [Vanmaekelbergh and Erne 1999; Bojinov et al., 2000].  53  \u000C8  4  a  b  4  10  Z\"/k\u03A9  Z\"/k\u03A9  10-2 5  10-1  10-2  2 105 10-1  0  0  0  4 Z'/k\u03A9  2  0.2  10-1  10  Z\"/k\u03A9  0  Z\"/k\u03A9  0  8  5  -2  d  10-2  10 5  -0.1 0  3  6  0  0.1  Z'/k\u03A9 0.8  e  10-2  0.5 105  0.2  0.2 Z'/k\u03A9  0.3  f  2.5 Z\"/k\u03A9  Z\"/k\u03A9  -1  0 10  -4  4  0.1  10-2  c  2 Z'/k\u03A9  10-2  1.5 105  0.5  10-1  10-1  -0.5  -0.1 0  0.25 0.5 0.75 Z'/k\u03A9  1  0  1  Z'/k\u03A9  2  3  Figure 5-3: Impedance spectra of a pyrrhotite electrode in 1 M HCl at (a) OCP, (b) 692, (c) 782, (d) 917, (e) 967, (f) 1042, (g) 1092 mV vs. SHE. Parameter is frequency in Hz.  54  \u000C1  g  10-1  Z\"/k\u03A9  0  105  -1 -2 10-2  -3 -4  -3  -2  -1  0  1  Z'/k\u03A9 Figure 5-3: continued.  As a brief conclusion to this section, the EIS studies have shown the formation of two distinct surface layers on pyrrhotite as a function of potential. The first quasi-passive layer forms during preparation or upon immersion of the electrode and the dissolution of this layer begins at around 782 mVSHE (cf. Figure 5-2). This layer would be the reason for the change in slope observed through the potential range of 712 to 842 mV SHE. The formation of the first layer can be considered independent of the concurrent process which also occurs at low anodic potentials (about 100 mV above the OCP) and results in the production of ferrous ions (cf. reaction 5-1) and elemental sulfur [Hamilton and Woods 1981]. This sulfur layer product would remain on the surface of the sample and could impede the rate of dissolution of the electrode, thus contributing to the apparent passivation or slowing of the reaction as potential is increased. However, Figure 5-3c shows that this quasi-passive layer, which is dissolving between 712 and 842 mVSHE, is not stable and so the effect of the sulfur layer on the surface could in fact be more dominant than the effect of this passive film. This idea is substantiated by the potentiodynamic study undertaken at different scan rates (Figure 5-2) which demonstrated that the resistance to polarization in the potential range of 712 to 842 mVSHE is not because of the formation or thickening of a true passive film on the surface. Beyond this primary film, and as potential is increased to the range of 842 to 917 mVSHE, the dissolution rate of pyrrhotite increases (Figure 5-2). Studies by Hamilton and Woods [1981] and Almeida and Giannetti [2003] have attributed this to the oxidation of Fe2+. EIS results of the present study show that the dissolution of the quasi-passive layer in this potential region 55  \u000Cis the dominant process. It is possible that this quasi-passive film is an Fe2+ compound which at higher potentials oxidizes to Fe3+. At more anodic potentials (917 mVSHE and higher) the dissolution of the first layer and formation of the second layer occurs simultaneously which is in accordance with the results of potentiodynamic scans. Almeida and Giannetti [2003] have suggested that this second surface layer would be a sulfate layer which forms by reaction 5-4. Increasing the potential of the electrode from 917 to 1042 mVSHE increases the total real impedance of the electrode by about a factor of ten, which means that the second layer is substantially more passivating than the first.  5.2.4 Mott-Schottky study In order to understand the electrical properties of the passive layers which are observed here, Mott-Schottky studies were undertaken. Figures 5-4 and 5-5 show the MottSchottky plots for the pyrrhotite electrode which were constructed from capacitance vs. electrode potential measurements at 1, 10 and 50 kHz. Experimentally, the total capacitance value of an electrode (C) can be determined from the measured imaginary part of the impedance (Z\") at sufficiently high ac frequency (f) using Eq. 5-6 [Stansbury and Buchana 2000]: Z \uF0A2\uF0A2 \uF03D  \uF02D1 2\uF070fC  5-6  Generally, in the case of semiconductor electrochemistry, the interface between a semiconductor electrode and an electrolyte can be considered as two capacitors in series: the space charge region of the semiconductor and the Helmholtz layer. Thus the measured capacitance of the film-electrolyte interface can be described by the following relationship:  1 1 1 \uF03D \uF02B C CH CSC  5-7  where Csc and CH are the space charge region and the Helmholtz capacitance, respectively. The capacitance for the space charge region is much smaller than that for the double layer (CSC \u00AB CH) [Bard and Faulkner, 2001; Kareh et al., 2009; Schmuki et al., 1995], such that one can neglect the Helmholtz double layer capacitance in most calculations. Therefore, the measured capacitance value (C) is equal to the space charge capacitance which is presented as the Mott-Schottky relation in the following expression:  56  \u000C1 2 \uF03D 2 CSC e\uF065\uF065 o no  kT \uF0F6 \uF0E6 \uF0E7 E \uF02D E fb \uF02D \uF0F7 e \uF0F8 \uF0E8  5-8  where e represents the charge of an electron, \u03B5 the dielectric constant of the semiconductor, \u03B5o the permittivity of free space, no the donor density, E the applied potential, Efb flat band potential, k the Boltzmann constant, and T the absolute temperature, where the term kT/e is 25 mV at 25\u00B0C [Gryse et al., 1975]. The Nyquist plots presented in the previous section show complex responses with several time constants for each studied potential. This means that the capacitance obtained via Mott-Schottky analysis cannot represent the definitive space charge capacitance. However, a change in slope in the Mott-Schottky plots is necessarily due to a change of the electrical nature of the surface of the electrode. The Mott-Schottky plots of Figure 5-4 show that the surface layers of pyrrhotite (in the potential range of 0.2 to 1.25 V vs. SHE) are n-type semiconductors, based on the increasing slopes of the curves. An interesting feature of the plots of Figure 5-4 is the sharp break at around 942 mVSHE. The break was observed for all the samples around this potential under both dark and irradiated conditions. Not surprisingly this potential is very close to the potential of formation of the second passive layer shown in the potentiodynamic and EIS studies (Figures 5-2 and 5-3). Therefore, it is thought that the change in slope of the Mott-Schottky plots is a consequence of the formation of a separate surface layer on the electrode during the dissolution process.  Figure 5-4: Mott-Schottky plots for pyrrhotite in 1 M HCl electrolyte in the dark and under irradiation.  57  \u000CKennedy and Frese [1978] have suggested that the break in the Mott-Schottky curve can also be explained by assuming the existence of two different kinds of donors in the semiconductor surface layer of pyrrhotite. Hence, the two Mott-Schottky slopes could be:  S1 \uF03D  2 e\uF065\uF065 o n1  (E \uF03C EB )  5-9  2 e\uF065\uF065 o n2  (E \uF03E EB )  5-10  and  S2 \uF03D  where EB is the break potential and n1 and n2 are the densities of donors before and after EB, respectively. This suggests that one kind of donor is very close to the conduction band and the second kind is deeper (i.e., far from the conduction band). Assuming the scenario suggested by Kennedy and Frese [1978], increasing the potential of the semiconductor up to potentials more than that of the break potential will cause the ionization of deep donors in the space charge region. A consequence of this phenomenon is an increase in the charge carrier density in the space charge region and a decrease in the slope of the Mott-Schottky plots as potential is increased. In the present study we have an increase in the slope of the MottSchottky plots which means that the mechanism proposed by Kennedy and Frese [1978] is likely not the case here. As a result, from Figure 5-4, the linear behavior of the MottSchottky plots over the potential interval of 200 to 950 mVSHE is related to the first quasipassive layer which forms during sample preparation and the break in the plots near 950 mVSHE is due to the formation of the second surface layer which is described in the EIS studies. The semiconductive properties of the second layer are to some extent different from those of the first layer. From a practical hydrometallurgical perspective, the dissolution of pyrrhotite at lower potentials (&lt; 950 mV vs. SHE) is more interesting. Thus, more study on the properties of the first surface layer is essential. For this reason we have focused on the properties of the surface layer formed over the potential range of 200 to 950 mVSHE. Figures 5-5a and b present the Mott-Schottky plots for pyrrhotite under irradiation and in the dark, respectively. Figure 5-5a shows good linear behavior for potentials lower than 750 mV SHE while increasing the potentials above 750 mV causes a scattering of the data. It is thought that the dissolution of the first passive layer would be the reason for such scattered behavior and 58  \u000CFigure 5-3c supports this idea by presenting the dissolution of this layer in this potential region. Table 5-1 shows the flat band potentials which were obtained from the intercepts of 1/C2 vs. potential by subtracting the term kT/e = 25 mV from the intercept (refer to Eq. 5-7). The dielectric constant, \u03B5, of pyrrhotite is around 82 [Rosenholtz and Smith 1936] and the donor density in the electrode for the two different potential regions was determined using this number (Table 5-1). Table 5-1 also contains the effective Debye length which was calculated according to the Debye equation [Memming 2001]: LD ,eff  \uF0E6 \uF065\uF065 kT \uF0F6 \uF03D\uF0E7 o 2 \uF0F7 \uF0E8 2ne \uF0F8  1  2  5-11 3  1 kHz 10 kHz 50 kHz  C-2 x 10-10/F-2cm4  2.5 2 1.5 1 0.5  a  0 0  0.2  0.4  0.6  0.8  1  Potential/V vs. SHE  10  10 kHz  7.5  50kHz 5  -2  -10  -2  C x 10 /F cm  4  1kHz  2.5  b 0 0  0.2  0.4 0.6 0.8 Potential/V vs. SHE  1  Figure 5-5: Mott-Schottky plots for pyrrhotite electrode in 1M HCl electrolyte, (a) under irradiation and, (b) in dark.  59  \u000CTable 5-1 clearly shows that the charge carrier density of the quasi-passive surface layer in the dark is lower than that under irradiation. This is the case at higher potentials as well. Table 5-1 also shows that the values of the donor density and effective Debye length for pyrrhotite change dramatically at the break potential. This significant difference between the two surface films was also observed by EIS (cf. Figure 5-3f). Figure 5-6 shows the thickness of space charge region, dSC, as a function of potential, \u2206\u03C6SC, which is obtained by [Azumi et al., 1987]: d SC  \uF0E9 e\uF044\uF066SC \uF0F9 \uF03D 2 LD ,eff \uF0EA \uF02D 1\uF0FA \uF0EB kT \uF0FB  1  2  5-12  Eq. 5-12 considers a semiconductor electrode with simple dissolution behavior which is not the case for the pyrrhotite electrode presented in this study with complex Nyquist responses. For this reason, the thickness of the space charge region presented in Figure 5-6 might not be very accurate but does, nevertheless, present evidence for a significant change in the space charge region\u2019s thickness at the break potential. A comparison of the two distinct surface layers reveals that the second, higher potential, layer has a lower donor density and consequently will have lower capacitance. This lower donor density infers that the second layer is less conductive than the first and, therefore, is likely better at hindering the dissolution of pyrrhotite. These results are similar to the results of Hamilton and Woods [1981], Almeida and Giannetti [2003], and Mikhlin [2000] who suggest the formation of different sulfur species during the electrochemical dissolution of sulfide minerals.  Table 5-1: Flatband potentials for pyrrhotite in 1 M HCl electrolyte in dark and under irradiation. f Efb n1 LD,eff,1 n2 LD,eff,2 \u20133 \u20133 kHz mV vs. SHE cm nm cm nm 1 \u2013423 2.19\u00D71020 0.810 7.15\u00D71018 4.505 Irradiated 10 \u2013227 0.92\u00D71020 1.258 4.37\u00D71018 5.769 20 18 50 \u2013332 0.89\u00D710 1.277 3.98\u00D710 6.041 1 \u2013436 1.51\u00D71020 0.981 5.21\u00D71018 5.275 20 18 Dark 10 \u2013243 0.35\u00D710 2.031 1.64\u00D710 9.420 50 \u2013483 0.27\u00D71020 2.302 1.02\u00D71018 11.914  60  \u000C5.2.5 Model development The EIS data of section 5.2.3 were fitted to three equivalent electrochemical circuits of Figure 5-6. These equivalent circuits represent a preliminary step in the elucidation of the electrochemical processes taking place at the electrolyte/electrode interface. The observed Nyquist impedance plot at high anodic potential (1092 mV SHE; Figure 5-3g) is complex due to the existence of the negative real resistance.  Figure 5-6: Equivalent electrochemical circuit for the Pyrrhotite/electrolyte interface at, (a) low anodic potentials (OC \u2013 692 mV), (b) 782 mV, (c) high anodic potentials (917 \u2013 1042 mV).  The authors have tried to fit the plot by using the circuits previously applied, such as [Epelboin and Keddam 1972; Bojinov et al., 1993]. The proposed circuit by Epelboin and 61  \u000CKeddam [1972] considers the dissolution and passivation of metals (iron and nickel) in sulfuric aqueous media. Also, the paper by Bojinov et al. [1993] studies the impedance response of a tin metal electrode in sulfuric aqueous media. Their work and also the work by Glarum and Marshall [1985] have proposed that the negative real impedance, like that of Figure 5-3g, is associated with a time constant which includes a negative resistance connected with a negative capacitance. However, the results of the fitting process were not completely satisfactory and needed more study. The present author is of the opinion that almost all of the publications to date which have dealt with such complex impedance behavior have been done on simple metallic electrodes such as iron, nickel [Epelboin and Keddam 1972], tin [Bojinov et al., 1993] and copper [Glarum and Marshall 1985]. Hence, it is reasonable to consider that since mineral electrodes have a substantially intricate nature than metallic electrodes, the observed complex behavior of the pyrrhotite electrode, as a mineral, would be more complicated to interpret. However, the proposed mechanisms for the metallic electrodes may still be applicable for the pyrrhotite electrode. Therefore, in the present contribution, instead of proposing an electrochemical circuit model for the observed behavior at 1092 mVSHE, the authors have looked into the physical meaning of such behavior as is consistent with previous efforts in the literature (this interesting feature of the pyrrhotite mineral electrode at high anodic potentials will be studied in a subsequent paper). Keddam and Takenouti [1985] have attributed the negative real impedance for a nickel electrode to transpassive dissolution. Elsewhere, Epelboin and Keddam [1972] have related this behavior to the competitive dissolution and passivation reactions of nickel. Studying the dissolution of a tin electrode, Bojinov et al. [1993] have suggested that the low frequency response is a complicated coupling between the passivation and active dissolution processes at the film/solution interface. In their work, a negative capacitance has been applied in the modeled electrochemical circuit and has been considered to be a measure of the degree of coverage of intermediate species - those which are assisting the dissolution of the passive film. However, studying the dissolution of copper in phosphoric acid solution, Glarum and Marshall [1985] have explained the complex impedance behavior by the dependence of the dissolution process on high mobility hydrogen ions. From their work, by increasing the potential of the electrode, the electrical fields external to the barrier lower the concentration of H+ and so the dissolution rate will be reduced. Considering a kinetic system including two adsorbed 62  \u000Cintermediates and two dissolution paths, Bojinov [1996] has reproduced the general features of the impedance spectra of copper, nickel, tin and antimony in sulfuric acid solutions. However, the proposed model does not take into account the possibility of bilayer film formation on the electrode. The modeled circuit for low anodic potentials (OCP \u2013 692 mVSHE) is presented in Figure 5-6a. Figure 5-6b shows the modeled circuit at 782 mVSHE where transpassive dissolution of the surface layer occurs. The model for high anodic potentials (917 \u2013 1042 mVSHE) is illustrated in Figure 5-6c. In all of the models the Rsol element represents the solution resistance. The Cp/Rp and the CSC/RSC pairs represent the capacitive and resistive behavior of the surface layer and the space charge region of the surface layer, respectively. The constant phase element (CPE) presents the effect of defects in the surface layer [Chen et al., 2008] and RCPE represents the resistance of the CPE in the equivalent circuit. The CPE element, which compensates for the non-homogeneity in the system, is defined by two values, Q and n. The impedance, Z, of CPE is presented by Eq. 13 [Hsu and Mansfeld 2001; Orazem and Tribollet 2008]: \uF02Dn  Z CPE \uF03D Q \uF02D1 \uF028 j\uF077 \uF029  5-13  where Q is the CPE parameter, i = (\u22121)1/2, \u03C9 is the frequency in rad s\u22121, \u03C9 = 2\u03C0f, f is the frequency in Hz, and n is the measure of the non-ideality of the capacitor and has a value in the range 0 &lt; n &lt; 1. If n equals 1, then, ZCPE is identical to that of a capacitor, ZC = (i\u03C9C)\u22121, where C is the capacitance. In fact, a capacitor is a constant phase element with a constant phase angle of -90\u00B0. Generally, a CPE behaves like a capacitance which varies with frequency. The presence of the CPE behavior may explain the frequency dependence of the capacitance measured in Figure 5-5. The element L may be due to the transfer of ions across the double layer or an increase of the area of the electrode by the partial dissolution of the surface layer [Vanmaekelbergh and Erne 1999] and RL represents its resistance. In these circuits, contributions of the double layer capacitance and charge transfer resistance as well as the sample resistance have been neglected. This can be justified by the relatively high values of the double layer capacitance compared to those of the space charge region (CSC) and capacitance of the surface layer Cp. Since these three capacitances are in series the double layer capacitance makes only a minor contribution to the total system capacitance. On the other hand, the electrode resistance is typically smaller than the resistance of the solution, 63  \u000Cthus the sample resistance may be ignored. Otherwise, it may introduce a slight error in the obtained values for Rsol. Figure 5-7 presents the experimentally measured and calculated Nyquist curves obtained with the equivalent circuits of Figure 5-6 and the values of different elements of the equivalent circuits are shown in Table 5-2.  64  \u000C4 Msd.  Msd. Cal.  Z\"/k\u2126  Z\"/k\u2126  8  4  Cal.  2  0  0 0  4 Z'/k\u2126  0  8  2 Z'/k\u2126  4  0.2  2 0  Z\"/k\u2126  Z\"/k\u2126  Msd.  Msd.  -2  Cal.  0.1 0  Cal.  -4  -0.1 0  3 Z'/k\u2126  0  6  0.1  0.2 0.3 Z'/k\u2126  0.4  1.3 0.8  Z\"/k\u2126  Z\"/k\u2126  2.8 Msd. Cal.  Cal.  1.8  0.3  0.8  -0.2  -0.2  0  0.5 1 Z'/k\u2126  1.5  Msd.  0  1  2  3  Z'/k\u2126  Figure 5-7: Measured and calculated Nyquist plots for pyrrhotite electrodes at (a) OCP, (b) 692 mV, (c) 782 mV, (d) 917 mV, (e) 992 mV, and (f) 1042 mV vs. SHE.  65  \u000CTable 5-2: Model parameters for equivalent circuits of Figure 5-6. E (mV)  RSC (\u2126 cm2)  106 CSC (F cm-2)  RP (\u2126 cm2)  103 CP (F cm-2)  RCPE (\u2126 cm2)  105 Q (\u2126 s-n)  n  RL (\u2126 cm2)  L (H)  OCP  952.1  116.80  21710.0  0.39  24510.0  24.03  0.56  -  -  692 782  716.6 5.2  136.00 2.40  5434.0 19140.0  2.95 0.28  975.2 5.1  21.15 0.20  0.56 0.36  24.0  91150.0  917 942  11.0 8.8  0.55 0.41  8.9 4156.0  35.40 23.15  243.3 276.4  1.58 2.45  0.74 0.72  8.9 4156.0  88.0 51.4  967 992  7.0 8.4  0.44 0.43  1137.0 1743.0  12.49 12.61  175.2 19.6  2.55 3.34  0.80 0.66  1137.0 1743.0  951.1 267.3  1042  16.7  0.34  4355.0  4.43  308.8  1.61  0.68  4355.0  945.4  In order to determine the capacitance value of the electrode (C) from the measured impedance spectra, a simple circuit shown in Figure 5-8 was fitted. Here, Rohmic is representing the ohmic resistance of the electrode and the CPE and Rt are describing the capacitance and resistance of the electrode\u2019s surface, respectively.  Figure 5-8: Simplified equivalent circuit model for the electrode.  In the potential region of active dissolution of the electrode (842 \u2013 912 mVSHE), the Rt element of the Figure 5-8 represents the resistance of the surface (which can be the resistance of the space charge region, Helmholtz double layer, or a combination of both). Then, the effective capacitance of the CPE will be calculated by Eq. 5-14 [Hirschorn et al., 2010]:  66  \u000C(1\uF02Dn )  Ceff  \uF0E6 R R \uF0F6 \uF03D Q n \uF0E7\uF0E7 ohmic t \uF0F7\uF0F7 \uF0E8 Rohmic \uF02B R t \uF0F8 1  n  5-14  where n and Q are the CPE parameters, as described before, Rohmic and Rt are representing the ohmic resistances of the system and the surface, respectively. Considering Rt \uF0F1\uF0F1 Rohmic , then the Eq. 5-14 will be [Hirschorn et al., 2010]: 1  (1\uF02Dn )  Ceff \uF03D Q n Rohmicn  5-15  where the Ceff represents an estimation of the capacitance of the electrode (C) as a function of Rohmic. However, in the potential range of OC \u2013 842 mVSHE and also potentials &gt; 912 mVSHE the surface of the electrode is blocked with the surface layers. Therefore, the effective capacitance of the CPE element at these potential ranges will be calculated by Eq. 5-16 [Chen et al., 2008; Hirschorn et al., 2010]: 1  (1\uF02Dn )  Ceff \uF03D Q n Rt  n  5-16  where the Rt shows the resistance of the surface layer (equal to the Rp values in Table 5-2) and other parameters have the same meaning of those in Eq. 5-15. The results from electrochemical studies and modeled electrochemical circuits can be used to propose a reaction sequence for the dissolution of the pyrrhotite electrode. The first stage of the dissolution which takes place at potentials ranging between OC \u2013 692 mVSHE may be related to the formation of a metal\u2013deficient sulfide layer: Fe1\uF02D x S \uF0BE \uF0BE\uF0AE Fe1\uF02D x \uF02D y S \uF02B yFe 2\uF02B \uF02B 2ye \uF02D  5-17  where the Fe1-x-yS layer is a semiconductor layer with donor density of around 1020 cm-3 (ref. Table 5-1). Both of the CSC/RSC and Cp/Rp pairs of Figure 5-6a are referring to this layer and the constant phase element of the model indicates the effect of surface imperfections. If the electrode obeys Tafel kinetics from OC - 692 mVSHE, then the Tafel slope, \uF062 a , will be [Marcus 2002]:  67  \u000C\uF062a \uF03D  2.03RT \uF061 a nF  5-18  where \uF061 a is the anodic transfer coefficient which can usually be approximated by 0.5 [Bard and Faulkner 2002], and n is the number of electrons taking place in the charge transfer limited reaction. For the case of reaction 5-17, Eq. 5-18 becomes:  \uF062a \uF03D  2.03RT \uF061 a (2 y ) F  5-19  and by applying Tafel analysis on Figure 5-9 the \uF062 a value was calculated to be 131 mV/decade. Hence, the 2y value will be 0.796 and it is possible to rewrite the dissolution reaction 5-17 for the case of x \uF03D 0.125 (Fe7S8) with the stoichiometry of: Fe0.875 S \uF0BE \uF0BE\uF0AE Fe 0.477 S \uF02B 0.398Fe 2\uF02B \uF02B 0.796e \uF02D  5-20  Potential/V vs. SHE  0.7 0.6 0.5 0.4 0.3 0.00001  0.0001  0.001  0.01  0.1  1  -2  i/mAcm  Figure 5-9: Tafel region polarization of the pyrrhotite electrode in 1 M HCl electrolyte.  The metal\u2013deficient sulfide layer of Fe0.602-xS (or Fe0.477S when x \uF03D 0.125 ) is stable up to around 742 mVSHE and by increasing the potential to higher values (i.e. 782 mV SHE) this layer will dissolve:  68  \u000CFe1\uF02D x \uF02D yS \uF0BE \uF0BE\uF0AE(1 \uF02D x \uF02D y \uF02D \u03BB)Fe 2\uF02B \uF02B \u03BBFe 3\uF02B \uF02B S \uF02B (2 \uF02D 2x \uF02D 2y \uF02B \u03BB)e \uF02D  5-21  where \u03BB is a constant ( 0\uF0E1 \uF06C \uF0E11 ) at each given potential (the ferric ion formation will be considered in more detail below). After complete dissolution of this first layer, the pyrrhotite might start to dissolve via Eq. 5-22: Fe1\uF02D x S \uF0BE \uF0BE\uF0AE (1 \uF02D x \uF02D \u03BD)Fe 2 \uF02B \uF02B \u03BBFe 3\uF02B \uF02B S \uF02B (2 \uF02D 2x \uF02B \u03BD)e \uF02D  5-22  From the proposed reaction sequence for the dissolution of the pyrrhotite electrode, the first stage of the dissolution which takes place at potentials ranging between OC \u2013 692 mVSHE is related to the formation of a metal\u2013deficient sulfide, by reaction 5-17. The results of the EIS experiments support this passive anodic behavior at low anodic potentials (see Figures 5-2a and b and related discussion). However, at higher potentials (i.e. 782 mV SHE) the EIS experiments provide evidence of transpassive dissolution (see Figure 5-2c and related discussion) which we believe is associated with reaction 5-21. According to this reaction, in the potential range of 712 to 842 mVSHE, the Fe1-x-yS passive film dissolves. However, elemental sulfur is also a product of reaction 5-21. Being a poor conductor, elemental sulfur film will impede the dissolution rate of the electrode. Thus, we have concluded that the decrease in the polarization current observed by lowering the scan rate from 5 to 0.5 mVs-1 (compare curves (a) and (c) of Figure 5-1) is due to the formation of an elemental sulfur reaction product and not because of the passivation process proposed in reactions 5-17 and 5-20. This dissolution sequence of the electrode was also studied by chronoamperometry. Figure 5-10a shows the current (density)\u2013time transient for the electrode after applying a potential step from OC potential to 782 mVSHE. After a peak current (67 \u03BCAcm-2) at initial immersion due to double layer formation, the Faradaic current results in the dissolution of iron atoms at the electrode surface. Consequently, a layer of the semiconductor Fe1-x-yS (or generally metal\u2013deficient sulfide layer) forms on the surface and further thickening of this layer as time progresses slows down the dissolution rate. The dissolution rate reaches a minimum value after 60 sec (see inset of Figure 5-10a). Therefore, in the early stage of dissolution at 782 mVSHE, reaction 5-17 \u2013or more specifically reaction 5-20 \u2013 is believed to be the major dissolution reaction. However, because 782 mV is higher than the reversible potential for the ferrous to ferric ion oxidation reaction (770 mVSHE), the production of ferric ion should also be considered at 782 mV SHE. The sources of the ferric ion could be: (1) 69  \u000Cferrous to ferric oxidation on the surface of the mineral and (2) direct dissolution of the mineral as ferric (and ferrous) ions. Studying the rate of the ferrous to ferric oxidation reaction on the pyrrhotite electrode was not the goal of this work. Determining the rate of the ferrous to ferric oxidation reaction on the pyrrhotite electrode (and generally mineral electrodes) is difficult due to existence of some other co-current oxidation reactions such as that of the pyrrhotite electrode itself, as Fe2+ and/or Fe3+, and the oxidation of impurities which may also be present. On the other hand, it is possible to discuss about the ferrous to ferric oxidation reaction according to the literature. The ferrous to ferric oxidation reaction is a fast reaction on the surface of oxide free metals such as platinum and has been reported to have the exchange current density of 15 mA cm-2 [Makrides 1964]. However, superficially oxidizing the surface of metals will cause a ten thousand fold decrease in the exchange current density of Fe2+/Fe3+ couple [Stansbury and Buchana 2000]. Reported values for superficially oxidized Fe, Ti, Fe-13.2% Cr, and 304 stainless steel electrodes are 10, 26, 1, 0.045 \u03BCA.cm-2, respectively [Makrides 1964; Stern 1957]. Figure 5-11 shows the Tafel region of ferrous to ferric oxidation reaction on different oxidized (passive) metals and also compares those curves with the polarization curve of the pyrrhotite electrode measured in our work. Also it have been demonstrated that the oxidation of ferrous ion on the surface of the chalcopyrite (CuFeS2) mineral, a mineral similar to that used in the present study, in HCl solution is extremely slow and almost negligible [Hirato et al., 1986]. Thus, it can be assumed that the ferrous to ferric oxidation reaction is a slow reaction on the pyrrhotite electrode too. Indeed, from the literatures on the electrochemistry of pyrrhotite, this is a proper assumption. Cyclic voltammetry studies on pyrrhotite electrodes conducted in acidic solutions by different researchers have reported existence of an anodic peak related to the formation of Fe3+. The peaks are reported to be at 0.43 V vs. SHE in 0.5 M CH3COOH solution (pH 4.6) [Hamilton and Woods 1981], 0.45 V vs. sat. Ag|AgCl in 0.5 M CH3COOH solution (pH 4.5) [Almeida and Giannetti 2003], and 0.48 V vs. sat. Ag|AgCl in 1M HCl solution [Mikhlin 2000]. However, it has been observed by Hamilton and Woods [1981] that running the cyclic voltammetry experiment either with a solid stationary electrode or rotating disk electrode does not significantly change the current measured at this peak. Therefore, they have concluded that the oxidation reaction which gives rise to the peak is originated directly from the mineral and is not produced from a 70  \u000Cdiffusion dependent dissolved species. In other words, if the source of the Fe3+ is the dissolved Fe2+ ion in the vicinity of the electrode surface, then by increasing the rate of diffusion of this species away from the mineral surface (through the use of a rotating disk electrode) one would expect to eliminate (or decrease) this peak. Consequently, the peak has been assigned to the direct oxidation of pyrrhotite electrode as ferric ion [Hamilton and Woods 1981]. Hence, the oxidation of the ferrous ions to ferric ions on the surface of the pyrrhotite electrode can be neglected. Therefore, in the first 60 s of the chronoamperometry curve of Figure 5-10a at 782 mVSHE, it should be considered that iron dissolves as ferrous and ferric (both by direct oxidation of the electrode) and thus the reaction for the early stage of dissolution at 782 mVSHE might be: Fe1\uF02D x S \uF0BE \uF0BE\uF0AE Fe1\uF02D x \uF02D m\uF02D n S \uF02B mFe 2\uF02B \uF02B nFe 3\uF02B \uF02B (2m \uF02B 3n)e \uF02D  5-23  After reaching a critical thickness of the metal\u2013deficient sulfide layer (Fe1-x-m-nS) at approximately 60 seconds \u2013see Figure 5-10a\u2013 transpassive dissolution of this layer will be the dominant reaction according to reactions 5-21 and 5-22. Subsequently, a layer of elemental sulfur will form on the surface which hinders the dissolution rate (Figure 5-10a). Application of this proposed dissolution sequence to the electrochemical circuit of Figures 56b, infers that the CSC/RSC and Cp/Rp pairs represent the semiconducting and passivation behavior of the metal\u2013deficient sulfide layer, respectively, while the constant phase element may be representative of the thin elemental sulfur layer at the electrode/solution interface. The inductive element, L, is related to the transpassive dissolution of the semiconducting metal\u2013deficient sulfide layer. The chronoamperometry curve of Figure 5-10a shows a gradual increase in dissolution current after 2000 sec. The current increases up to a steady state value of approximately 40 \u03BCAcm-2. The reason behind this behavior is not clear, however, either oxidation of the elemental sulfur layer to sulfate or further oxidation of ferrous ions to ferric might be the cause.  71  \u000C80 Anodic current/\u03BCA cm  Anodic current/\u03BCA cm  -2  -2  a  60  40  30  20  10 0  60  120  180  240  300  360  time/sec  20  0 0  1800  3600  5400  7200  time/sec 1.2  Anodic current/mA cm  -2  b 0.9  0.6  0.3  0 0  1800  3600  5400  7200  time/sec Figure 5-10: Anodic current density-time transient in response to potential step from the OC potential to, (a) 540, and (b) 850 mV; inset of figure (a) shows the earlier stage of dissolution of the corresponding condition with higher resolution.  The dissolution path for the electrode changes when potential is increased to more anodic regions (&gt; 917 mVSHE). Figure 5-10b shows an almost uniform decrease in the dissolution current density of the electrode which is clearly due to the formation of a surface 72  \u000Clayer on the electrode. Reaction 5-55 is the likely overall reaction at very high anodic potentials which considers the dissolution of iron as ferric ion and formation of both elemental sulfur and sulfate ions. Due to the high dissolution rates which are observed (steady state dissolution current is 0.12 mAcm-2) a thick, relatively un-protective, sulfur layer or sulfate product is believed to form on the electrode surface. The Mott-Schottky plots of Figure 5-3 show semiconducting behavior for the surface layer which means that the layer cannot be a perfect insulator layer of elemental sulfur. For these reasons, the layer likely consists of a sulfate product (or other sulfur species) or contains a considerable amount of sulfate. In the electrochemical circuit of Figure 5-6c for high anodic potentials the CSC/RSC and L/RL pairs are related to the formation of the space charge region and dissolution of the semiconductive pyrrhotite, respectively. Cp/Rp represents the capacitive and resistive properties of the surface layer and constant phase element and its corresponding resistance may be due to the imperfections on the surface of film.  73  \u000CFigure 5-11: Tafel region of ferrous to ferric oxidation reaction on different oxidized metals compared to the polarization curve of pyrrhotite electrode in HCl solution.  5.2.6 SEM studies Considering reactions 5-5 and 5-22, elemental sulfur is the main product of the dissolution process. In order to observe the morphology of the sulfur layer formed on the surface of the electrode, samples were kept at different constant potentials for two hours and then analyzed by scanning electron microscope (SEM). Figures 5-12a and b show the morphology of the sulfur layer formed at 782 and 1042 mVSHE, respectively. The sulfur layer formed at lower potentials (Figure 5-12a) is rough and discontinuous and does not cover the whole surface of the electrode and, as a result, this film likely would not be very successful at impeding dissolution as demonstrated through electrochemical measurement. The layer formed at higher potentials after two hours (Figure 5-12b) is thicker and contains fewer 74  \u000Ccracks and discontinuities. The latter appears to cover the whole surface of the sample and, as a result, this layer is more successful at hindering the dissolution rate of the electrode. This is consistent with the results obtained in the electrochemical study.  a  a  b Figure 5-12: SEM images of the pyrrhotite electrode after 2 hrs dissolution in 1 M HCl solution at (a) 782, (b) 1042 mVSHE.  Indeed, from Figure 5-10, the steady dissolution current at 782 mV and at 1092 mVSHE are 18 \u03BCAcm-2 and 160 \u03BCAcm-2, respectively. Thus, increasing the potential by 310 mV (from 782 to 1092 mV) increases the dissolution current by less than an order of magnitude (around 8.9 times). As a comparison, the anodic Tafel slope for the electrode in the potential region of OCP \u2013 692 mVSHE is 131 mVdecade-1, which is 2.47 times higher than that of 782 \u2013 1092 mVSHE range. This comparison shows that the dissolution kinetics increases very slowly in the higher potential range (782 \u2013 1092 mVSHE) which we believe is a result of the physical properties of the second surface layer (&gt; 912 mVSHE). Viz. the donor 75  \u000Cdensity of the second surface layer (~ 1018 cm-3) is lower than that of the first (~ 1020 cm-3, ref. to Table 5-1), and the total impedance of the electrode increases rapidly through the formation of this second layer (roughly from 0.3 k\u2126 to 3.5 k\u2126, ref. to section 5.2.3 and Figure 5-2).  5.3 Conclusion In this chapter, the electrochemical dissolution behavior of pyrrhotite has been studied in 1M HCl solution with in-situ electrochemical methods. Results show formation of two distinct surface layers on the electrode. The first layer forms during the preparation of the sample or during initial immersion of the pyrrhotite in the test solution. At the initial time, the irona atoms from the pyrrhotite structure partially dissolve into the solution and a metal\u2013 deficient sulfide layer with semiconducting properties (Fe1-x-yS) forms on the surface of the electrode. The donor density of the layer is around 1020 cm-3 and the thickening of the layer decreases the dissolution rate of the mineral. The layer is stable up to potential of 842 mVSHE. The second film forms at higher anodic potentials (&gt; 912 mVSHE) and is substantially more resistant to dissolution than the first layer. Considering reaction 5-5 as the overall dissolution path for this potential range, the ferric ion, elemental sulfur and sulfate ion are the main products of reaction. The results of this chapter relate to those presented below, and pertaining to chalcopyrite dissolution, as the pyrrhotite electrode (a Fe-S system) is a simplified model for the chalcopyrite electrode (Cu-Fe-S system). The electrochemical methods developed in this chapter for the simplified Fe-S system were then applied below to help elucidate the mechanism of chalcopyrite dissolution.  76  \u000C6 Electrochemical evaluation of chalcopyrite\u2019s surface during dissolution in 0.5 M H2SO4 solution2 6.1 Introduction Different in-situ electrochemical techniques (i.e. cyclic voltammetry, potentiodynamic and potentiostatic polarization, and chronoamperometry) as well as ex-situ spectroscopic methods (i.e. X-ray photoelectron spectroscopy, auger electron spectroscopy, Raman spectroscopy) have been applied to understand the dissolution mechanism of chalcopyrite, and furthermore, to characterize the surface layers on electrodes (refer to chapter 2 of the thesis). However, there remains uncertainty about the dissolution procedure. The trouble with most of the DC electrochemical techniques used so far is that only circumstantial evidence is derivable to help with the elucidation of possible reaction mechanisms. The ex-situ analytical methods are plagued with controversy associated with the possible evaporation of elemental sulfur under experimental conditions [Buckley and Woods 1994]. In this chapter, the electrochemical dissolution of chalcopyrite in simple sulfuric acid solution will be studied. In order to characterize the surface properties of the chalcopyrite electrode at different anodic potentials and understand the growth sequence of surface layers, we will apply electrochemical impedance spectroscopy. As chalcopyrite is a semiconducting material, Mott-Schottky plots will be extracted to understand the electrolyte/electrode interface. The lessons learned from the pyrrhotite electrochemistry (EIS and Mott-Schottky) study were significantly useful to the development of the methods and analysis applied in this chapter.  6.2 Results and discussion 6.2.1 Characterization of used chalcopyrite electrodes A high quality natural polycrystalline chalcopyrite sample originating from the Creighton Mine (Sudbury, Ontario) was used for electrode fabrication. The quantitative phase analysis of the chalcopyrite ore was determined using the Rietveld method and X-ray powder diffraction and is shown in Table 6-1. Massive electrodes were prepared by cutting 2  A version of this chapter has been published: \u201CGhahremaninezhad, A., Asselin, E., Dixon, D.G., 2010, Electrochimica Acta, 55, 5041\u20135056\u201D.  77  \u000Cthe chalcopyrite into approximately cubic shapes with areas of 1 cm2 exposed to the solution. Electrical contact with copper wires was made in the back of the sample with silver-loaded epoxy resin. The electrode contacts were protected from the solution by mounting the whole assembly in epoxy resin, leaving only one face of the electrode exposed to the solution. Before each test, the exposed face of the electrode was polished with no. 1200 carbide paper to begin each experiment with a fresh surface. The reference electrode in all of the performed tests was a Mercury/mercury Sulfate electrode which is 0.600 mV vs. SHE. The microstructure of a polished electrode is shown in Figure 6-1a. Also, Figure 6-1b presents an SEM picture and its associated EDX mapping for elemental distribution which shows the presence of the siderite phase in the chalcopyrite electrode structure.  Table 6-1: Composition of the as received chalcopyrite ore. Mineral Chemical Formula Mass Percentage Chalcopyrite  CuFeS2 2+  Siderite  Fe CO3  98.1 1.9  a  b Siderite  Chalcopyrite Figure 6-1: SEM images of (a) polished electrode surface, (b) siderite phase and EDX mapping for element distribution on the surface of prepared electrode.  78  \u000C6.2.2 Open circuit potential study In order to elucidate the surface activity of the electrode after 10 min immersion in the test solution, the open circuit potential curves of different chalcopyrite electrodes as a function of time are shown in Figure 6-2. In all of the curves the potential of the electrode increases for 10 min and afterward almost reaches to a steady state. The rapid increase in the potential during the first 10 min of immersion in the electrolyte is approximately 60 mV which indicates the spontaneous growth of a passive film on the surface of the electrodes. The steady-state OC potential was found to be approximately \u22120.25 VMSE (0.35 VSHE) and was established within roughly 10 min (where the steady-state was defined here as a change of less than 2 mV/min). All of the electrodes used in the experiments shown here displayed an open circuit potential of \u2013250 \u00B1 10 mVMSE (350 \u00B1 10 mVSHE) after 10 min of stabilization time in the electrolyte.  Figure 6-2: Open circuit potential (OCP) of different chalcopyrite electrodes vs. time.  6.2.3 Potentiodynamic study Figure 6-3 shows the potentiodynamic curves for the chalcopyrite electrode at different scan rates of 0.05, 0.5, 5, and 50 mVs\u20131. Generally, applying a lower scan rate will reveal a closer approximation of steady state electrode behavior. In the case of the present study, the polarization current decreases with decreasing scan rate which is similar to the results published by Warren et al. [1982] for CuFeS2 in 1 M sulfuric acid solution. For cathodic potentials, all of the curves show similar E-I profiles which means that the initial sample surfaces have similar properties. In general, three different potential regions are 79  \u000Capparent in Figure 6-3. At low anodic potentials, OC to 0.3 VMSE (0.9 VSHE), the dissolution current is a function of scan rate which is clearly related to the formation of a surface layer. Within the small potential range of 0.30 to 0.42 VMSE (0.90 to 1.02 VSHE), the measured current is the same for all scan rates. In other words, the surface of the electrode is electrochemically active in this region. At higher anodic potentials (&gt; 0.42 VMSE) another film(s) forms on the surface of the electrode and hinders the dissolution rate of the electrode again.  Figure 6-3: Potentiodynamic curves for the chalcopyrite electrode in 0.5 M H2SO4 electrolyte at different scan rates.  A closer look at Figure 6-3 reveals further details about the anodic behavior of chalcopyrite. In the case of the low anodic potentials, OC to 0.3 VMSE (OC to 0.9 VSHE), the measured current at lower scan rates is less than that registered at higher scan rates. This behavior, decrease in current by descending the scan rate, is likely related to the formation of a thin surface layer(s) via prewave reactions [Nava and Gonzalez 2006]: CuFeS 2 \uF02B 4zH 2 O \uF0AC \uF0BE\uF0AE Cu 1\uF02D x Fe1\uF02D yS 2\uF02D z \uF02B xCu 2\uF02B \uF02B yFe 3\uF02B  80  \u000C\uF02B zSO 24\uF02D \uF02B 8zH \uF02B \uF02B (2x \uF02B 3y \uF02B 6z)e \uF02D  6-1  CuFeS2 \uF02B 4pH2 O \uF0AC \uF0BE\uF0AE Cu1\uF02Dm Fe1\uF02Dn S2\uF02Dp \uF02B mCu2\uF02B \uF02B nFe3\uF02B \uF02B pSO 24\uF02D \uF02B 8pH \uF02B \uF02B (2m \uF02B 3n \uF02B 6p)e \uF02D  6-2  For lower scan rates the layer can have more coverage on the surface. Also, the layer would be thicker or would have less porosity (more time to react) than at higher scan rates. As a result, the anodic current would be smaller (as seen in Figure 6-3). Performing the experiments at very low scan rate reveals the existence of two different passivation regions at low anodic potentials. According to the low scan rate polarization curves of Figure 6-3 (0.05 and 0.5 mVs-1), the first passive layer forms in the potential range of OC to 0.125 VMSE (0.725 VSHE) and afterward the second passive layer forms during a small potential window of 0.19 to 0.3 VMSE (0.79 to 0.9 VSHE). The formation and dissolution of these layers will be studied in more detail in the next section. At more anodic potentials of 0.30 to 0.42 VMSE (0.90 to 1.02 VSHE), changing the scan rate does not significantly affect the dissolution current. This is likely due to the electrochemically active dissolution of the passive surface layers. At potentials higher than 0.42 VMSE, the lower scan rate shifts the curve to lower currents which could again be considered as evidence for the formation of another surface layer(s) on the electrode. At an even higher potential of around 0.68 VMSE (1.28 VSHE) there is a further decrease in the measured current which could be considered as resulting from the formation of a second, and different, layer on the surface of the electrode. These findings will be discussed in more detail in the next section.  6.2.4 Electrochemical impedance spectroscopy study EIS studies on the chalcopyrite electrodes were conducted in three different potential regions: at low anodic potentials of OC to 0.3 VMSE (OC \u2013 0.9 VSHE) with a passive electrode surface, in a mid-potential region of 0.3 to 0.42 VMSE (0.9 \u2013 1.02 VSHE) with an active electrode surface and at higher anodic potentials of 0.42 to 1.10 VMSE (1.02 \u2013 1.70 VSHE). The interpretation of the measured Nyquist and Bode plots shown in this section will be supported by model development in section 6.2.6.  81  \u000CThe impedance spectra for the chalcopyrite electrode at low anodic potentials are presented in Figure 6-4. The Nyquist curve obtained at OC potential is similar to that obtained by Asselin et al [2007] where it was shown that such behavior is related to the existence of a thin surface layer on the electrode. The EIS study by Vel\u00E1squez et al. [2005] in alkaline solution also revealed the formation of a coating layer which is initially present on the electrode and spontaneously forms during the polishing of the sample. Generally, this layer forms during the preparation of the sample or during initial immersion in the electrolyte for stabilization purposes as studied by Biegler [1977]. Nava and Gonzalez [2006] and Nava et al., [2008], have proposed the formation of a non stoichiometric passive layer of Cu1-xFe1yS2-z  on the surface of chalcopyrite (reaction 6-1). In alkaline solutions, Yin et al., [2000] and  Vaughan et al., [1997] have suggested that, at low anodic potentials (near OCP), a monolayer of surface iron should dissolve first while the Cu and S atoms remain unchanged as CuS2* and inhibit further dissolution. According to their findings, a small increase in potential will cause further dissolution of iron atoms from deeper sections of the chalcopyrite surface while the CuS2* layer remains stable and, due to its thickening, the dissolution current will decrease quickly. Possibly, in the case of the present study, the following reaction may occur as considered by Hackl et al., [1995]: CuFeS 2 \uF0BE \uF0BE\uF0AE Cu 1\uF02D x Fe1\uF02D y S2 \uF02B xCu 2\uF02B \uF02B yFe 2\uF02B \uF02B 2(x \uF02B y)e \uF02D  y\uF0F1\uF0F1 x  6-3  The Cu1-xFe1-yS2 layer is a passive layer and is responsible for a decrease in dissolution rate. The electrode shows similar behavior for higher anodic potentials of \u20130.1 and 0.0 VMSE (0.5 and 0.6 VSHE, respectively) as presented in Figure 6-4. It could be concluded that by increasing the potential, more atomic layers are involved in reaction 6-3 such that the passive film grows into the electrode surface. Simultaneously, the thickness of the passive layer increases and inhibits dissolution [Yin et al., 1995]. Increasing the potential shows a partially resolved semi-circle at high frequencies which is related to charge transfer process on the electrode. This result is close to that of Velasquez et al., [1998]. The passive surface layer is stable at higher potentials (up to 100 mV) and furthermore the effect of charge transfer increases. The impedance curve at 150 mVMSE (750 mVSHE), presented in Figure 6-4, shows a dramatic change in electrochemical behavior. Three well pronounced time constants are detectable in the curve. The first time constant, at high frequencies, is a capacitive loop 82  \u000Crelated to the combination of the double-layer capacitance impedance combined with a charge transfer resistance (due to the electrochemical dissolution of the electrode) and the second one, at middle frequencies (0.49 \u2013 1.5 Hz) is a capacitive loop related to the combination of a pseudo-capacitance impedance (due to the passive layer) and a resistance. At very low frequencies (0.01 \u2013 0.29 Hz) the Nyquist curve reveals inductive behavior. Most probably, the inductive loop (third time constant) is related to the transpassive dissolution of the passive layer [Betova et al., 2002; Bojinov and Tezvetkoff 2003; Erne and Vanmaekelbergh 1997].  83  \u000CFigure 6-4: Electrochemical impedance spectra of chalcopyrite electrode for anodic potentials of OC to 300 mVMSE.  The phase angle diagrams at the low anodic potentials are presented in Figure 6-5a. In the potential range of OC to 0 mVMSE (OC to 600 mVSHE) the electrode shows a maximum phase angle of around 80 degrees which is considered to be related to a solid state mass transport process (through passive layer) and at 150 mV MSE (750 mVSHE) the electrode has a negative phase angle at high frequencies which supports the inductive behavior shown in Nyquist curve of Figure 6-4 for 150 mVMSE. As a result, at 150 mVMSE the passive layer on the electrode is porous due to the transpassive dissolution of the electrode. 84  \u000CFigure 6-5: Bode diagrams of (a) phase angle, and (b) total impedance magnitude of the chalcopyrite electrode at potential range of OC\u2013150 mVMSE (OC to 750 mVSHE).  Figure 6-5b shows the total impedance magnitude diagram for the low anodic potentials. From Figure 6-5b, by increasing potential from 0 to 150 mVMSE (600 to 750 VSHE) the total impedance of the electrode at low frequencies decreases more than an order of magnitude. This is believed to be related to transpassive dissolution of the passive layer and its concomitant increase in porosity. The porous nature of the passive layer in this potential range has also been described by Vel\u00E1squez et al., [2005] and Nava and Gonzalez [2006]. If the Cu1-xFe1-yS2 compound identified by reaction 6-3 is assumed to be the passive layer  85  \u000Cformed from OCP to 0.1 VMSE then it electrochemically dissolves at higher anodic potentials according to Hackl et al., [1995] as: Cu 1\uF02D x Fe1\uF02D yS 2 \uF0BE \uF0BE\uF0AE Cu 1\uF02D x \uF02D zS2 \uF02B zCu 2\uF02B \uF02B (1 \uF02D y)Fe 2\uF02B \uF02B 2(z \uF02B 1 \uF02D y)e \uF02D  6-4  where the Cu1-x-zS2 layer is also passive. The presented Nyquist plot for the chalcopyrite electrode at 250 mVMSE (850 mVSHE) in Figure 6-4 shows five different time constants; a capacitive loop related to the combination of the double-layer capacitance impedance combined with a charge transfer resistance (100 kHz \u2013 14 Hz), a capacitive loop in the 13 \u2013 4.5 Hz related to the combination of a pseudo-capacitance impedance (due to the passive layer) and a resistance, an inductive loop in middle frequencies (5 \u2013 0.85 Hz), and third capacitive and second inductive loops at very low frequencies (0.5 \u2013 0.01 Hz). The capacitive loop at very low frequencies (third capacitive loop) can be related to the formation of a new surface layer on the electrode which according to reaction 6-4 would be Cu1-x-zS2. Similar behaviors have been reported in the case of transpassive dissolution of highly alloyed stainless steels [Betova et al., 2002; Erne and Vanmaekelbergh 1997; Betova et al., 2004]. Figure 6-6a presents the total impedance magnitude diagram for the chalcopyrite electrode at 250 mVMSE (850 mVSHE).  Figure 6-6: Bode diagrams of total impedance magnitude of the chalcopyrite electrode at (a) 250, (b) 300 mVMSE.  At the lowest frequencies (Figure 6-6a) the total impedance of the electrode increases which means that the second inductive loop for 250 mVMSE (inductive loop at the lowest frequencies) is not related to the dissolution of the passive layer. This inductive behavior between 0.15 and 0.01 Hz may be related to the relaxation of negative surface charges on the 86  \u000Cpassive layer which form by accumulation of metal vacancies at the passive layer/solution interface [Bojinov et al., 1999]. In the case of the anodic behavior of iron in sulfuric acid, Gabrielli et al. [1996] explains this behavior as the relaxation of the mobile Fe3+ ion. Both of the interpretations of the second inductive loop are analogous from an electrical circuit perspective. The last curve of Figure 6-4 shows the Nyquist plot at 300 mVMSE (900 mVSHE) for the electrode. Two capacitive time constants are visible at high frequencies and one inductive loop at low frequencies. From the low scan rate polarization curve of Figure 6-3, the potential of 300 mVMSE is the very early stage of active dissolution region for the electrode. In other words, 300 mVMSE is a transient potential from passive state to active dissolution state. Also comparing Figures 6-6a and b shows a rapid decrease (around order of magnitude) of total impedance of electrode by increasing the electrode potential from 250 to 300 mVMSE (850 to 900 mVSHE). Consequently, at this potential, the Cu1-xFe1-yS2 passive film is almost dissolved by reaction 6-4, and the Cu1-x-zS2 layer is the only present passive film on the surface (which also is the product of reaction 6-4). Thus, in Nyquist plot of 300 mV MSE (refer to Figure 6-4), the high frequency loop will be due to the combination of a double layer capacitance impedance and a charge transfer resistance, the combination of a capacitance impedance and a resistance for mid frequencies (due to Cu1-x-zS2 passive layer) and the combination of a inductance impedance and a resistance for low frequencies (due to a transpassive dissolution of the film). The effect of low frequency transpassive dissolution of the electrode can be seen in the Bode diagram for the electrode at 300 mVMSE (Figure 66b), which shows a decrease in the total impedance magnitude of the system. The decrease is due to the transpassive dissolution of the passive film and decrease in their thickness. Hackl et al. [1995] have suggested a reaction for the dissolution of the passive layer: Cu 1\uF02D x \uF02D z S 2 \uF0BE \uF0BE\uF0AE (1 \uF02D x \uF02D z)Cu 2 \uF02B \uF02B 2S 0 \uF02B 2(1 \uF02D x \uF02D z)e \uF02D  6-5  In summary, the EIS studies at low anodic potentials of OC \u2013 0.3 VMSE (OC to 0.9 VSHE) show evidence for the formation of two distinct passive layers according to reactions 6-3 and 6-4 which is in accordance with the findings of the potentiodynamic polarization studies. Generally, the first passive layer forms during sample preparation and is stable in the potential range of OC to 0.1 VMSE (Reaction 6-3). Afterward, a second passive layer forms on the surface (0.1 to 0.3 VMSE) and, at same time, the dissolution of the first passive layer is occurring (Reaction 6-4). 87  \u000CFigure 6-7 shows the impedance spectra for the electrode at 325\u2013390 mVMSE (925 to 990 mVSHE). In this potential range the electrode shows a capacitive time constant at high frequencies and an inductive loop at low frequencies. The high frequency impedance loop is the result of charge transfer resistance and double layer capacitance coupling. The inductive impedance loop at low frequencies is characteristic of anodic dissolution of semiconductor electrodes [Gomes and Vanmaekelbergh 1996; Vanmaekelbergh and Searson 1994; Hiskey 1993]. In this potential range, electrochemically active dissolution of the electrode occurs and none of the passive layers formed at lower potentials are stable. Figure 6-8a presents the total impedance magnitude Bode diagram for the dissolution of chalcopyrite in the potential range of 250\u2013390 mVMSE (850\u2013990 mVSHE). Increasing the potential from 275 to 325 mV MSE causes a two order of magnitude decrease in the total impedance of the electrode. Simultaneously, the phase angle diagram of Figure 6-8b shows a decrease in capacitive behavior of the electrode at the same potentials. Hence, at the potential window of 300 to 390 mVMSE the anodic dissolution of the electrode occurs.  Figure 6-7: Electrochemical impedance spectra of chalcopyrite electrode for anodic potentials of 325\u2013390 mVMSE (925\u2013990 mVSHE).  Biegler and Swift [1979] suggest that this dissolution reaction occurs as:  88  \u000CCuFeS 2 \uF02B 8H 2 O \uF0BE \uF0BE\uF0AE Cu 2\uF02B \uF02B Fe 3\uF02B \uF02B 2SO 24\uF02D \uF02B 16H \uF02B \uF02B 17e \uF02D  6-6  A study by Lazaro and Nicol [2006] has indicated that soluble sulfur species are involved in the dissolution of chalcopyrite. Their results show that S 2O32 \uF02D would be the soluble sulfur species and the proposed reaction is: CuFeS 2 \uF02B 3H 2 O \uF0BE \uF0BE\uF0AE Cu 2\uF02B \uF02B Fe 2 \uF02B \uF02B S 2 O 32\uF02D \uF02B 6H \uF02B \uF02B 8e \uF02D  6-7  which considers the formation of thiosulfate as a product of chalcopyrite dissolution. An interesting feature of Nyquist plots at these potentials (Figure 6-7) is a jump in real impedance values around the frequency of 10 Hz. This behavior occurs in the potential range of 300 to 400 mVMSE and by increasing the potential the jump in the real impedance decreases. It is possible to consider that a frequency dependent adsorption of a dissolution product occurs at this point. The sulfur species from reactions 6-6 and 6-7 could be good candidates for this phenomenon. However, more studies are needed here and the present authors are not sure about the nature of this process. Impedance spectra for the chalcopyrite electrode at high anodic potentials of 0.50 \u2013 1.10 VMSE (1.1 to 1.7 VSHE) are presented in Figure 6-9. At 500 mVMSE (Figure 6-9a), three major time constants are detectable; a capacitive loop at high frequencies related to Faradaic resistance and double layer formation on the electrode (105 \u2013 386 Hz); an inductive loop at middle frequencies which is a characteristic of anodic dissolution of semiconductor electrodes (386 \u2013 4.5 Hz); and finally a second capacitive loop (4.5 \u2013 10-2 Hz) which is related to the formation of a new film (surface layer). The potentiodynamic studies presented in the previous section (Figure 6-3) also show evidence of film growth on the surface of the electrode. Moreover, the total impedance magnitude Bode diagram of Figure 6-10 shows an increase at low frequencies which is related to formation of the surface layer. Nava and Gonzalez [2006] suggest the formation of ferric sulfate at this potential according to the reaction: 2CuFeS 2 \uF02B 13H 2 O \uF0BE \uF0BE\uF0AE 0.75CuS \uF02B 1.25Cu 2\uF02B  \uF02B Fe 2 (SO 4 ) 3 \uF0AF \uF02B0.25SO 24\uF02D \uF02B 26H \uF02B \uF02B 28e \uF02D  6-8  The formation of covellite (CuS) at these potentials has also been proposed by Warren et al., [1982] as well.  89  \u000CFigure 6-8: Bode diagrams of (a) total impedance magnitude, and (b) phase angle of the electrode at potential range of 250\u2013390 mV.  CuS might be the surface layer which results in the pseudo-passive behavior of the electrode at 0.50 mVMSE (1.10 VSHE). Increasing the potential by more than 250 mV to a value of 750 mVMSE (1.35 VSHE) adds one more time constant to the Nyquist plot of the electrode (Figure 6-9b). The low frequency capacitive behavior would be related to either the formation of a new surface layer on the electrode or a notable change in the morphology and roughness of the previously formed surface layer. As presented in Figure 6-10, there are two 90  \u000Csteps in the total impedance magnitude of the electrode at 0.75 VMSE (1.35 VSHE). The first step (around the frequency of 10 Hz) is thought to result from the formation of the former surface layer (CuS layer) and the second step, at lower frequencies (around the frequency of 0.1 Hz), is thought to be related to the latter phenomenon.  Figure 6-9: Electrochemical impedance spectra of chalcopyrite electrode for anodic potentials of 0.50\u2013 1.10 VMSE (1.10 to 1.7 VSHE).  Increasing the potential to 1.10 VMSE (1.7 VSHE) result in an inductive loop at very low frequencies of 92 \u2013 10 mHz (Figure 6-9c) and the Bode diagram of total impedance magnitude at 1.10 VMSE (Figure 6-10) shows a drop in total impedance at the same frequency range. It may thus be assumed that the low frequency inductive loop is related to a transpassive dissolution process. Figure 6-9d presents a higher magnification of the inductive loop for the anodic dissolution of the electrode and the capacitive loop associated with the formation of surface layer which occur at 1.10 VMSE. Studies by Nava and Gonzalez [2006] indicate that the covellite layer which forms around 0.50 VMSE (1.10 VSHE) is not stable at higher potentials (&gt;750 mVMSE) and the suggested dissolution reaction is:  91  \u000CCuS \uF02B 4H 2 O \uF0BE \uF0BE\uF0AE Cu 2\uF02B \uF02B SO 24\uF02D \uF02B 8H \uF02B \uF02B 8e \uF02D  6-9  where the covellite decomposes to cupric ions (Cu2+). Furthermore, some authors previously have proposed the formation of jarosite at high anodic potentials [Klauber 2008; Dutrizac 1982; Dutrizac 1989]. X-ray photoelectron spectroscopy studies by Parker et al. [2003] indicate that having ferric sulfate on the surface of the chalcopyrite electrode acts as a nucleation template for jarosite formation which hinders the dissolution of the electrode. Ferric sulfate could be directly produced by reaction 6-8 or by the formation of sulfate in the system. For example, further oxidation of the thiosulfate formed according to reaction 6-7 forms sulfate in the system. The proposed reaction for jarosite formation is [Klauber 2008]: 2Fe 2 (SO 4 ) 3 \uF02B (K, Na, H 3O) 2 SO 4 \uF02B 12H 2 O \uF0BE \uF0BE\uF0AE 2(K, Na, H 3O)Fe 3 (SO 4 ) 2 (OH) 6 \uF0AF \uF02B6H 2SO 4  6-10  The jarosite formation rate increases very quickly by increasing the temperature and the rate is maximum at pH around 2. The presence of a nucleation template enhances the jarosite formation significantly [Klauber 2008, Parker et al., 2003]. Thus, the formation of jarosite on the surface of electrode could explain the very low frequency capacitive behavior of the electrode at 750 mVMSE (1350 mVSHE) of Figure 6-9b.  Figure 6-10: Bode diagram of total impedance magnitude for the chalcopyrite electrode at potential range of 390\u20131100 mVMSE (990 to 1700 mVSHE).  92  \u000C6.2.5 Mott-Schottky study It is useful here to summarize the results of the impedance work provided above: (1) It was observed that for the potential range of OC to 0.3 VMSE (0.9 VSHE) two different passive films formed on the surface of chalcopyrite; (2) in the potential range of 0.3 to 0.42 VMSE (0.9 \u2013 1.02 VSHE) the surface of the chalcopyrite dissolves actively without any passive layer formation; and (3) at higher potentials of 0.42 \u2013 1.1 VMSE (1.02 \u2013 1.7 VSHE) another surface film forms. In this section, Mott-Schottky tests are used to probe the semiconducting properties of the chalcopyrite and its surface films as a function of potential. Figure 6-11 shows the Mott-Schottky plots for chalcopyrite electrodes at different frequencies of 1 and 10 kHz. A polarization curve is also included in Figure 6-11. Generally, as explained in the previous chapter, the imposed potential at a semiconductor/electrolyte interface results in the formation of the electrical equivalent of two capacitors in series. The first capacitance results from the Helmholtz (double) layer while the second is associated with the space charge region. The overall capacitance of the interface can be described as:  1 1 1 \uF03D \uF02B C CH CSC  6-11  where Csc and CH are the space charge and the Helmholtz capacitances, respectively. The capacitance for the space charge region is much smaller than that of the double layer (2 to 3 orders of magnitude), such that one can neglect the Helmholtz double layer capacitance in most calculations. Therefore, the measured capacitance value is equal to the space charge capacitance which is presented as the Mott-Schottky relation in the following expression:  1 2 \uF0E6 kT \uF0F6 \uF03D \uF0E7 E \uF02D E fb \uF02D \uF0F7 2 CSC e\uF065\uF065 o no \uF0E8 e \uF0F8  6-12  where e represents the charge of an electron, \u03B5 the dielectric constant of the semiconductor, \u03B5o the permittivity of free space, no (or po for p-type semiconductors) the donor density, E the applied potential, Efb flat band potential, k the Boltzmann constant, and T the absolute temperature, where the term kT/e is 25 mV at 25\u00B0C [Memming 2001]. Figure 6-11 shows 5 specific potential ranges, each with particular slope. The slopes of the Mott-Schottky plot reveal the electronic behavior of the outer surface layers of the electrode which are near the electrolyte. The Nyquist plots presented in the previous section 93  \u000Cshow complex responses with several time constants for each studied potential. This means that the capacitance obtained through Mott-Schottky analysis cannot represent the definitive space charge capacitance. However, a change in slope in the Mott-Schottky plots is necessarily due to a change of the electrical nature of the surface of the electrode. In the case of Figure 6-11, for potentials less than around 0.3 V the electrode surface has n-type semiconductive properties due to the positive slope of the Mott-Schottky plot. For higher potentials the slope is negative, which means that the electrode surface has changed to p-type behavior. Minor slope changes are noticeable at different potentials of around -0.1, 0.2, 0.3, and 0.4 VMSE (0.5, 0.8, 0.9 and 1.0 VSHE, respectively) for both curves. The potentiodynamic and EIS studies of previous sections have revealed the formation and dissolution of different surface layers at these critical potentials.  Figure 6-11: Mott-Schottky plots for chalcopyrite electrode at different frequencies.  For potentials less than -0.1 VMSE (0.5 VSHE), the surface is n-type and most probably this is associated with the surface layer formed on the electrode during sample preparation. The first break in the Mott Schottky plots occurs at around -0.1 VMSE. Some authors have suggested that a monolayer thick passive film forms \u2013reaction (6-3)\u2013 on the surface of the electrode in this potential region (around -0.1 VMSE) [Yin et al., 1995; Vaughan et al., 1997] 94  \u000Cwhile other studies propose the formation of a layer with thickness around 2 to 4 nm [Biegler 1977; Biegler and Horne 1985]. Thus, the thin passive film explains the break at -0.1 VMSE (0.5 VSHE). This observed increase in slope implies a decrease in the charge carrier density of surface layer compared to the surface at lower potentials. Increasing the anodic potential causes the deeper layers of the chalcopyrite electrode to be involved in reaction 6-3 and the thickness of the Cu1-xFe1-yS2 layer (product of reaction 6-3) increases on the surface. As a result the conductivity and capacitance of the surface layer decreases. The second break in the plot occurs around 0.2 VMSE (0.8 VSHE). This break as seen in the Mott-Schottky plot is very close to the potential at which reaction 6-4 is believed to start. As a result, it is suggested that the formation of the Cu1-x-zS2 layer on the surface \u2013reaction 6-4\u2013 changes the electrical properties of the surface in this potential region. It has already been mentioned that chalcopyrite actively dissolves in the potential range of 0.3 \u2013 0.42 VMSE (potentiodynamic and EIS studies section) and reactions 6-6 and 6-7 are the suggested reactions. Since the dissolution of the chalcopyrite is electrochemically active in this potential region, it is likely that either no surface layer is present on the chalcopyrite surface or that which is present is ineffective. Thus, the observed behavior of the electrode in the Mott Schottky plots of Figure 6-11 in the potential range of 0.3 \u2013 0.42 VMSE (0.9 \u2013 1.02 VSHE) is directly associated with the properties of the chalcopyrite. The Mott Schottky curves of Figure 6-11 show a negative slope for the electrode in this potential region which means that the chalcopyrite electrode of the present study is a p-type semiconductor. At close to 0.4 VMSE (1.1 VSHE) another break in Figure 6-11 appears which is be related to the formation of a CuS layer on the chalcopyrite surface according to reaction 6-8. Table 6-2 shows the charge carrier density and effective Debye length for chalcopyrite and its surface layers in 0.5 M H2SO4 electrolyte for different frequencies. In the development of this table it is assumed that the capacitance measured by the Mott-Schottky method is very close to the space charge region capacitance. Actually, for potentials less than \u20130.1 VMSE (0.5 VSHE) and in the potential range of 0.3 to 0.4 VMSE (0.9 to 0.1.1 VSHE) the Nyquist curves are not very complex. It is therefore acceptable to employ the Mott-Schottky technique for the measurement of the space charge region capacitance and the calculated charge carrier density and effective Debye length likely represent good estimates in these potential ranges. For other potential ranges (for example \u20130.1 to 0.3 VMSE) the Nyquist 95  \u000Ccurves are very complex and calculated numbers in Table 6-2 are likely poor estimates. The dielectric constant, \u03B5, of chalcopyrite at room temperature is close 10 [Salsman et al., 1996]. Assuming that the dielectric constant of all surface layers (Cu1-xFe1-yS2 and Cu1-x-zS2) are close to that of the chalcopyrite, the donor density of surface layers for four different potential regions are determined using the slope of Eq. 6-12 (refer to Table 6-2). Table 6-2 also contains the effective Debye length which was calculated according to the Debye equation [Memming 2001]: LD ,eff  \uF0E6 \uF065\uF065 kT \uF0F6 \uF03D\uF0E7 0 2 \uF0F7 \uF0E8 2ne \uF0F8  1  2  6-13  Table 6-2 shows that the charge carrier density and effective Debye length of the interface decreases and increases by increasing the potential, respectively. Under the assumption that there is no surface layer on the electrode in the 0.3 to 0.4 VMSE (0.9 to 1.1 VSHE) range (active dissolution range), the charge carrier density in this region is that of the electrode (chalcopyrite) and is calculated to be close to 1022 m-3.  Table 6-2: Charge carrier density and effective Debye length for chalcopyrite and its surface layers at 0.5 M H2SO4 electrolyte.  f kHz 1 10  &lt; \u2013 0.1 VMSE n1 LD,eff,1 \u20133 m nm 8.8\u00D71024 1.9\u00D71024  0.9 1.9  \u2013 0.1 to 0.2 VMSE n2 LD,eff,2 \u20133 m nm 6.2\u00D71023 2.5\u00D71023  3.3 5.3  0.2 to 0.3 VMSE n3 LD,eff,3 \u20133 m nm 2.6\u00D71023 0.8\u00D71023  5.2 9.5  0.3 to 0.4 VMSE P4 LD,eff,4 \u20133 M nm 5.3\u00D71022 2.8\u00D71022  11.5 15.8  6.2.6 Model development Considering the surface properties and dissolution processes already elucidated at various potentials, the EIS data were fitted to different equivalent electrochemical circuits. Figure 6-12a shows the modeled circuit for the potential range of OC to 0 VMSE (OC to 0.6 VSHE) and the calculated Nyquist plots according to the model. In this model, the constant phase element (CPE) most probably presents due to defects in the passive layer [Chen et al., 2008]. Another physical explanation for the CPE would be electrode roughness [Mulder et al., 1990]. RCPE represents the resistance of the CPE in the equivalent circuit. The Cp/Rp pair and the CSC/RSC pair represent the capacitive and resistive behavior of the passive film on the 96  \u000Csurface and in the space charge region of the passive film, respectively. In all of the models Rsol is the contribution of the solution resistance to the electrochemical circuit. This model was also previously considered for the Si/SiO2 electrode in HCl electrolyte [Kareh et al., 2009]. Figures 6-12b to e present the measured and modeled (calculated) Nyquist plots for the anodic dissolution potentials of OC to 0 VMSE. The passive layer formed at this potential is believed to be the Cu1-xFe1-yS2 layer which forms through reaction 6-3. The values of the different elements in the equivalent circuit of Figure 6-12a are shown in Table 6-3. Table 6-3: Model parameters for equivalent circuit of Figure 6-12a (presented in next page). Potential (mV)  RSC (\u2126 cm2)  106 CSC (F cm-2)  105 RP (\u2126 cm2)  105 CP (F cm-2)  RCPE (\u2126 cm2)  105 Q (\u2126 s-n)  n  OCP -200 -100 0  14.5 9.0 7959.0 51210.0  5.83 12.70 6.69 12.10  9.5 7.5 7.9 3.4  1.51 2.43 1.21 35.24  13.4 12.6 16.6 5690.0  4.26 5.48 12.75 1.77  0.710 0.685 0.606 0.800  Increasing the potential to 0.15 VMSE (0.75 VSHE) results in transpassive dissolution of the passive layer formed at lower potentials. This transpassive dissolution results in a porous surface layer as shown by Vel\u00E1squez et al. [2005]. Figure 6-13a shows the proposed model for such a surface layer in which the surface roughness is taken into account. In this model (Figure 6-13a) C1/R1 pair refers to the passive layer at sections with maximum thickness and C2/R2 pair refers to the passive layer at sections with minimum thickness. Such a model has previously been used by Mansfeld and Kendig [1988] for modeling the interfacial behavior of Al with a porous alumina passive layer on the surface. The CSC||Ra||(LRb) section of the model refers to the dissolution of the semiconductor at the electrolyte as used by different researchers [Gomes and Vanmaekelbergh 1996; Vanmaekelbergh and Searson 1994]. The parallel connection of the Ra||(LRb) corresponds to the real impedance values during the depletion layer control of the dissolution process while the Helmholtz layer control of the process corresponds to the imaginary impedance values [Vanmaekelbergh and Searson 1994]. Figures 6-13b and c represent the measured and modeled Nyquist plots and total magnitude of the impedance Bode plot for the chalcopyrite electrode at 0.15 VMSE (0.75 VSHE).  97  \u000CCPE Rsc  Q Sol  Rsol  (a)  RCPE  CP  Cp  P1  Csc Rp  Figure 6-12: (a) Equivalent electrochemical circuit for the Chalcopyrite/electrolyte interface at low anodic potentials; Measured and calculated Nyquist plots for chalcopyrite electrode at (b) OC potential, (c) -200 mV, (d) -100 mV, (e) 0.0 mV.  From the EIS and polarization studies presented above, a second passive layer (Cu1-xzS2)  forms on the surface of the electrode through reaction 6-4 in the potential range of 0.2 \u2013  0.3 VMSE (0.8 to 0.9 VSHE). The model for the electrochemical circuit at this potential range is illustrated in Figure 6-14a. In this figure, the Rp/Cp represent the resistance and capacitance of the inner passive film (i.e. directly on the chalcopyrite surface) while CPE and RCPE are used to model the outer surface layer. The parallel connection of the Ra||(LRb) corresponds to 98  \u000Cthe dissolution of the semiconductor electrode and the CSC/RSC pair represent the capacitive and resistive behavior of the space charge region of the passive film. The calculated and measured Nyquist plots and impedance magnitude Bode plots are shown in Figures 6-14e \u2013 b for 0.2 and 0.25 VMSE (0.8 to 0.85 VSHE).  Csc  C1 Sol  Rsol  Ra  CP  R2  P1  (a)  L Rb  R1 C2  Figure 6-13: (a) Equivalent electrochemical circuit for the Chalcopyrite/electrolyte interface at 0.15 VMSE (0.75 VSHE); Measured and calculated (b) Complex-impedance plane representation, and (c) Bode representation of the magnitude of the impedance.  99  \u000CCPE  Rsc  Q RP  Ra  Sol  Csc  (a)  CP  P2  Rsol  P1  L Rb  CP  RCPE  Figure 6-14: (a) Equivalent electrochemical circuit for the Chalcopyrite/electrolyte interface at potential range of 0.2 to 0.25 VMSE (0.8 \u2013 0.85 VSHE); Measured and calculated (b) Nyquist representation at 0.2 VMSE, (c) Bode representation of the magnitude of the impedance at 0.2 VMSE, (d) Nyquist representation at 0.25 VMSE, (e) Bode representation of the magnitude of the impedance at 0.25 VMSE.  In the potential range of 0.3 to 0.4 VMSE (0.9 to 1.0 VSHE) active dissolution of chalcopyrite electrode occurs. The model chosen for this potential range is shown in Figure 6-15a. The Css||RL||(LRSS) section of the model represents the dissolution of semiconductor electrodes in which CSS and RSS contribute to the surface (or interface) states on electrode and L, RL denote the imaginary part of the impedance. A model close to this has been used by Cardon [1972] for the dissolution of a semiconductor-electrolyte interface with electron 100  \u000Ctransfer through surface states. The model considers surface state interactions with both the conduction and valence bands and with the filled and empty states of a redox system. The Rsol and CSC/RSC pair addresses the solution resistance and capacitive and resistive properties of the space charge region, respectively. As discussed before, the Nyquist plots at these potentials show a jump in real impedance values around the frequency of 10 Hz. However, the nature of this phenomenon is not clear to these authors. As a result, the proposed model for the 0.3 \u2013 0.4 VMSE potential range (Figure 6-15a) does not address the jump in EIS results. In order to generate best fit results, the data points at the frequencies less than 10 Hz are disregarded for models fit at 325 and 350 mV MSE. At the potential of 390 mV MSE (990 mVSHE) the 10 Hz jump is very small and considered to be of negligible importance to the modeling procedure. Figures 6-15b \u2013 d present the measured and modeled Nyquist plots for the three potentials tested between 300 and 400 mVMSE. The values of the different elements in the equivalent circuit of Figure 6-15a are shown in Table 6-4.  Table 6-4: Model parameters for equivalent circuit of Figure 6-15a. Potential (mV) 325 350 390  RSC  cm )  107 CSC (F cm-2)  RSS (\u2126 cm2)  105 CSS (F cm-2)  122.2 92.9 24.9  5.55 5.47 8.52  20.2 9.3 7.5  1.92 2.79 1.64  (\u2126 2  RL  (\u2126 cm2)  1087.0 117.2 21.9  L  (H) 15.61 9.77 2.04  101  \u000CCSS Rsc Sol CP  Rsol  RSS  L  Csc  (a)  RL  Figure 6-15: (a) Equivalent electrochemical circuit for the Chalcopyrite/electrolyte interface at potential range of 0.3 to 0.4 VMSE (0.9-1.0 VSHE); Measured and calculated Nyquist representations at (b) 325, (c) 350, and (d) 390 mVMSE.  Figure 6-16a shows the model applied for the 0.5 VMSE (1.1 VSHE) Nyquist plot. In this model the circuit components CPE||RL||L1RCPE represent the formation and associated properties of the surface layer on the electrode. The surface layer is a CuS layer which forms as a product of reaction 6-8. The element L1 represents ion transfer from the double layer or an increase in the surface area of the electrode due to partial dissolution [Vanmaekelbergh 102  \u000Cand Ern\u00E9 1999]. The components labeled C||Rb||L2Ra represent the dissolution of the semiconductor electrode as discussed previously. Figure 6-16b shows the measured and calculated Nyquist plots at 0.5 VMSE. C  CPE  Rsc Rsol  Q Ra  Sol  CP  RL  L2  P3  Csc  (a)  L1 Rb  RCPE  Figure 6-16: (a) Equivalent electrochemical circuit for the Chalcopyrite/electrolyte interface at potential of 0.5 VMSE (1.1 VSHE), (b) Measured and calculated Nyquist representation at 0.5 VMSE.  The modeled electrochemical circuit for the potential of 0.75 VMSE (1.35 VSHE) is presented at Figure 6-17a and the modeled Nyquist plot in Figure 6-17b. At this potential the effect of the negative imaginary part is not very significant and so the CPE||Rp||CpRCPE part of the modeled circuit does not contain any inductive elements. Thus, the CPE||Rp||CpRCPE section of circuit represents the mass transport in solid state for both of the surface layers. As discussed before, the surface layers are considered to be the CuS layer from reaction 6-8 and the gradual formation of jarosite on the surface is expected to occur according to reaction 610. The CSC||RL||LRSC section addresses the dissolution of the semiconductor. Generally, the C||R||LR circuit could be considered as a general model for the dissolution of semiconductors 103  \u000C[Gomes and Vanmaekelbergh 1996] and the physical meaning of the elements depends upon the details of the entire dissolution (or electrochemical reactions) mechanism. So, the C in the equivalent circuit of Figure 6-16a might be related to the surface states or the reaction intermediates at the surface which are separate from the space charge capacitance, CSC. However, at a potential of 0.75 VMSE, the mechanism of dissolution changes such that the CSC||RSC section of the circuit dissolves into the main circuit representing semiconductor dissolution (CSC||RL||LRSC). This could be attributed to the importance of charge transfer in the space charge region at this potential. Clearly, the proposed reactions for high potential anodic dissolution of chalcopyrite are very complex in mechanism. For example, reaction 6-8 has been proposed for chalcopyrite dissolution around 0.5 VMSE and this reaction involves the transfer of 28 electrons. At higher potentials, Nava and Gonzalez [2006]: 2CuFeS 2 \uF02B 16H 2 O \uF0AC \uF0BE\uF0AE 2Cu 2\uF02B \uF02B Fe 2 (SO 4 ) 3 \uF02B SO 24\uF02D \uF02B 32H \uF02B \uF02B 34e \uF02D  6-14  which involves the transfer of 34 electrons. As a consequence, any increase in potential results in significant dissolution mechanism changes which are illustrated in this work by the significant changes in circuit parameterization. The simulated electrochemical circuit for 1.1 VMSE (1.7 VSHE) is presented in Figure 6-18a. In this model, the C4/R4 and C3/R3 represent jarosite as a surface film (P4) and CuS as an inner surface layer (P3), respectively. L1 is most probably present due to the change in surface area from partial surface layer dissolution [Vanmaekelbergh and Ern\u00E9 1999]. Physically this can be explained by transpassive dissolution of the inner or outer surface layers. Another explanation for this inductive behavior could be the relaxation of surface charges as discussed before. The C||Rb||L2Ra section of the circuit refers to the dissolution of the chalcopyrite (as the general reaction occurs at high anodic potential) and the CSC/RSC pair addresses the capacitance and resistance of the space charge region. Figure 6-18b shows the best fit for the Nyquist plot at 1.1 VMSE which is fitted by the proposed model for this potential. Higher magnification of the Nyquist plot for the frequency region related to the formation of the first surface layer (CuS layer) is presented in Figure 6-18c.  104  \u000CCSC  CPE  Q Rsol  RSC  L  P4  Rp  Sol  (a)  RL  CP  P3  Cp RCPE  Figure 6-17: (a) Equivalent electrochemical circuit for the Chalcopyrite/electrolyte interface at potential of 0.75 VMSE (1.35 VSHE), (b) Measured and calculated Nyquist representation at 0.75 VMSE.  105  \u000CC  C4  Rsc  P4  Rsol  Ra  L2  C3  (a)  R4  Sol  Csc  CP  P3  L1 Rb  R3  Figure 6-18: (a) Equivalent electrochemical circuit for the Chalcopyrite/electrolyte interface at potential of 1.1 VMSE (1.7 VSHE), (b) Measured and calculated Nyquist representation at 1.1 VMSE, (c) magnification of Figure 6-18b.  All of the equivalent electrochemical circuit models are generated by considering the sequence of proposed electrochemical reactions of section 6.2.4. Several points are worth mentioning when summarizing the results of this chapter: \uF0B7  At the outset, a thin Cu1-xFe1-yS2 layer is present on the surface of the electrode. This is a passive layer which forms during the sample preparation and by application of an anodic potential its thickness increases via reaction 6-3. The layer is stable up to the potential of 0.1 VMSE (0.7 VSHE).  \uF0B7  In the potential range of the 0.1 \u2013 0.3 VMSE (0.7 \u2013 0.9 VSHE) the Cu1-xFe1-yS2 layer dissolves and a second passive layer of Cu1-x-zS2 forms through reaction 6-4. Generally, in the lower half of the potential range (around 0.1 \u2013 0.15 VMSE) the transpassive dissolution of the Cu1-xFe1-yS2 layer is the major phenomenon on the surface (Nyquist curve of Figure 6-4 at 0.15 VMSE) and by increasing the potential to the upper half of the range (0.15 \u2013 0.3 VMSE) a complete film of Cu1-x-zS2 forms on the outer surface of the porous Cu1-xFe1-yS2 layer (Nyquist curve of Figure 6-4 at 0.25 VMSE). The model of Figure 6-13 addresses the transpassive dissolution of the Cu1-xFe1-yS2 layer while the 106  \u000Cmodel of Figure 6-14 elucidates the effect of the Cu1-x-zS2 film as a second passive layer formed on the surface of the first passive layer of Cu1-xFe1-yS2. At this point, the Cu1-xFe1yS2  \uF0B7  layer is a thin porous layer directly on the surface of chalcopyrite.  In the potential range of 0.3 \u2013 0.42 VMSE (0.9 \u2013 1.02 VSHE) the chalcopyrite electrode dissolves with an electrochemically active surface. Dissolution of the previously formed Cu1-x-zS2 layer occurs through reaction 6-5 and chalcopyrite dissolves via reaction 6-6 and/or 6-7.  \uF0B7  Increasing the potential of the electrode into the range of 0.42 \u2013 0.7 VMSE (1.02 \u2013 1.3 VSHE) causes the formation of a CuS layer on the surface of the electrode via reaction 6-8. This layer impedes the dissolution rate of the electrode (ref. to the low scan rate potentiodynamic curves of Figure 6-3). However, considering the high dissolution current, the layer is not a passive layer.  \uF0B7  For the potentials higher than 0.7 V MSE (1.3 VSHE), the previously formed surface layer of CuS dissolves through reaction 6-9. Moreover, the conditions are favorable for the formation of jarosite on the surface via reaction 6-10. Thermodynamic data (via E-pH diagrams [Garrels and Christ, 1965]) indicate that  chalcopyrite will dissolve at pH less than 6 and an oxidizing potential higher than \u20130.4 VMSE (0.2 VSHE). However, the main issue in the dissolution of chalcopyrite under atmospheric conditions appears to be the passivation of chalcopyrite through the formation of Cu1-xFe1-yS2 and Cu1-x-zS2 passive layers.  6.3 Conclusion The present chapter considered the electrochemical dissolution of chalcopyrite in 0.5 M sulfuric acid solution. Potentiodynamic polarization, electrochemical impedance spectroscopy, and Mott-Schottky techniques were applied. Initially, during the sample preparation a thin layer of Cu1-xFe1-yS2 forms on the surface of the electrode where y\uF0F1\uF0F1 x . During electrochemical dissolution of chalcopyrite this layer is stable and its thickness increases with potential by reaction 6-3 (OC \u2013 0.1 VMSE). At higher potentials (0.1 \u2013 0.3 VMSE) the previously formed surface layer partially dissolves and a second passive layer (Cu1-x-zS2) forms via reaction 6-4. Both of these layers display the characteristics of passive layers at low scan rate (0.05 mVs-1) and for high scan rates they are observed to be pseudo107  \u000Cpassive. In the potential range of 0.3 to 0.4 VMSE, none of the surface layers are stable and electrochemically active dissolution of the electrode occurs by reactions 6-6 and/or 6-7. By increasing the potential to 0.5 VMSE, CuS is believed to form through reaction 6-8 and this only slightly impedes the dissolution of chalcopyrite. In fact, due to the high dissolution current observed at this potential (around 0.1 Acm-2) this layer cannot be considered a passive layer. At higher potentials (around 0.75 VMSE) the conditions are suitable for the formation of jarosite on the surface of electrode. The ferric sulfate phase (which is produced through reaction 6-8) has been shown to act as a nucleation template for jarosite. Jarosite forms according to reaction 6-9 and this reaction product does decrease the dissolution rate. Understanding the surface behavior of chalcopyrite during the electrochemical dissolution was the main goal of this study. The observations of the present study were explained based on the literature and results of previous studies. Since EIS was the main electrochemical technique applied in the present study, different equivalent electrochemical circuits were modeled and compared to experimental results to support the proposed sequence of chalcopyrite dissolution and surface layer formation. The models have been used to describe the experimental data in a satisfactory way. In the next chapter, the XPS method will be applied to explore the composition of chalcopyrite\u2019s passive film. XPS can provide very useful information about the chemistry of the passive film and the information obtained will be used to understand further details of the dissolution mechanism.  108  \u000C7 X-ray photoelectron spectroscopy analysis of chalcopyrite\u2019s surface in sulfuric acid solution 7.1 Introduction In the previous chapter, electrochemical methods were applied to analyze the surface of a chalcopyrite electrode in 0.5 M sulfuric acid. All the discussion and the proposed reactions were extracted from literature. In the present chapter, X-ray photoelectron spectroscopy will be applied to study the surface of chalcopyrite and link the results to the previous electrochemical studies.  7.2 Results and discussion 7.2.1 Passivation of chalcopyrite; a brief introduction Figure 7-1a shows the multi-cyclic voltammogram (10 cycles) of the chalcopyrite electrode in the potential window of 0.1 to 1 VSHE. In this experiment, the potential of the electrode was initially swept from 0.55 to 1 VSHE then 10 cycles, from 1 VSHE to 0.1 VSHE, were performed continuously (scan rate was 50 mV s-1). Two anodic peaks at 0.48 VSHE (peak a1) and 0.84 VSHE (peak a2) and a cathodic peak c1 at 0.29 VSHE are clear in Figure 71a. Both cathodic and anodic current peaks declined as the experiment proceeded due to the passivation of the electrode surface. The decline of the associated charge with the peak a1, as the number of cycles increased, is shown in Figure 7-1b. The anodic peak a2 has been shown to be related to the initial dissolution of chalcopyrite [Yin et al., 1995]. This peak is completely undetectable in all of the following cycles due to passivation of chalcopyrite (this point is very clearly shown in Fig. 7-1c). In other words, during the following cycles, the surface of chalcopyrite appears to be in the passive state and the dissolution rate seems to be controlled by mass transfer through the passive film which substantially reduces the dissolution rate [Yin et al., 1995]. For the same reason, the associated charge with this peak of Fig. 1a decreases significantly to a negligible value and the peak is not detectable in the second cycle. The anodic peak at 0.84 VSHE (peak a2) has been attributed to the following reaction [Yin et al., 1995]:  109  \u000CCuFeS 2 \uF03D CuS *2 \uF02B Fe 2\uF02B \uF02B 2e \uF02D  7-1  Figure 7-1: (a) Multicyclic voltammogram (10 cycles) of a chalcopyrite electrode in the potential window of -0.1 to 1 VSHE, (b) associated charge with the anodic peak of Figure 7-1a at 0.476 V at each cycle, (c) first cycle of Figure 7-1a.  It has been suggested that the initial chalcopyrite dissolution at 0.84 VSHE results in formation of a metastable copper sulfide phase with the stoichiometry of Cu:2S shown as CuS *2 . In addition, it has been proposed that this phase is responsible for the subsequent  passivation of chalcopyrite [Yin et al., 1995]. It is worth mentioning that peaks c1 and a1 of Fig. 7-1c are not related to the direct oxidation of chalcopyrite; however, the observed decrease in the associated charge at peak a1 (and also c1) is indirectly due to chalcopyrite passivation. Fig. 7-2a presents the potentiodynamic polarization of chalcopyrite. After stabilizing the electrode in solution at OCP for 6 hours, the electrode was polarized in the cathodic direction (solid-blue curve in Fig. 7-2a). The anodic part of the curve was separately obtained  110  \u000Cby anodic polarization of a fresh chalcopyrite electrode, also after 6 hours of stabilization at OCP (dashed-red curve in Fig. 7-2a). The applied scan rate in both of the polarization experiments was 0.1 mV s-1 and the solution was gently aerated during the experiment. The separate experiments for the measurement of the anodic and cathodic branches were conducted to avoid the impact of cathodic dissolution of chalcopyrite on the anodic section of the curve. The registered OCP for the electrode after 6 hours, before starting the polarization experiment of Fig. 7-2a, was ca. 0.55 VSHE and the OCP was typically reached after 1 hour of immersion. At potentials between OCP and 0.94 VSHE the electrode shows a slow dissolution rate owing to the passivation of the electrode in this potential window [Viramontes-Gamboa et al. 2007]. This result is in agreement with our previous chapter. It is important to note that a scan rate of 0.1 mV s\u22121 would be sufficiently low to obtain a steady-state polarization response from most metallic electrodes; however, this scan rate is not low enough if one seeks to measure the most stable E-I response of chalcopyrite [Viramontes-Gamboa et al., 2007]. This is clearly shown by comparison of the potentiodynamic polarization diagram of Fig. 7-2a with the potentiostatic polarization diagram of Fig. 7-2b. The passivation of chalcopyrite in sulfuric acid solution is more evident in Fig. 7-2b which shows a decrease in current density as the potential is increased above 0.67 VSHE.  111  \u000CFigure 7-2: (a) Potentiodynamic (SR = 0.1 mV sec-1) and (b) potentiostatic polarization of chalcopyrite electrode in 0.5 M sulfuric acid.  The formation of a stable passive film on the surface of chalcopyrite is the reason for the sharp decrease in current prior to the steady state dissolution current density in the passive potential region being established. This can be confirmed by long term potentiostatic polarization of chalcopyrite at a passive potential, as shown in Fig. 7-3. Fig. 7-3a presents i-t behavior of chalcopyrite at potentials of 0.5, 0.55 and 0.65 VSHE. All the curves of Fig. 7-3a show a sharp decrease in the current density during the first few hours of the experiment, which is due to the formation of a passive film. Fig. 7-3b presents the long-term potentiostatic polarization of chalcopyrite at a passive potential. This figure presents the 112  \u000Cpolarization behavior of chalcopyrite in sulfuric acid solution for about 60 hours at 0.80 V SHE. From Fig. 7-2b, it is obvious that the potential of 0.80 VSHE is within the passive range. Under this polarization condition the electrode reaches the steady state current in about 3 hours and afterwards the current density is almost constant (ca. 1.3 \u03BCA cm\u22122) for the rest of the experiment. This confirms the formation of a stable film on the surface and suggests that the thickness of the passive film most probably remains constant during the 60 hours of experiment. Considering that diffusion is the rate-limiting mechanism in the passive film, which is shown to be the case in chapter 9 of the present thesis, from Fick\u2019s first law of diffusion one would expect a substantial decrease in the current density of Fig. 7-3b if the passive film were thickening. For instance, if the thickness of the passive film were to double (during the 60 h of experiment), then the current density should decrease by half. However the current density of Fig. 7-3b (and also Fig. 7-3a) is almost constant, suggesting a passive film of constant thickness. The formation of a compact elemental sulfur film is also very unlikely at this potential since accumulation of elemental sulfur (as an insulating material) on the surface must lower the current density as a function of time. Additionally, after this experiment (about 60 h), there was no visible product layer on the surface of the electrode. This is in agreement with the findings of other workers [Warren et al. 1982; Holliday and Richmond, 1990; Yin et al., 1995] and will be further verified by our XPS results. The noise observed in the current density of Fig. 7-3b (sudden increases in dissolution current) is most probably related to the formation of pits on the electrode [Biegler and Swift, 1979]. Chemical characterization of the species in the passive film is very complicated or impossible by electrochemical methods. Hence, in the next section, XPS was used to provide detailed information about the oxidation state of different species and the composition of the surface film(s).  113  \u000CFigure 7-3: Potentiostatic polarization of chalcopyrite at (a) 0.5, 0.55 and 0.6 V for 6 hours, and (b) 800 mVSHE for 60 hours.  7.2.2 XPS study of chalcopyrite (survey spectra) XPS is a very successful surface analytical tool for mineral dissolution research. Although XPS requires sample transfer from the electrolyte to ultrahigh vacuum (UHV) with a concomitant loss of contact with the electrolyte and loss of electrode potential control, it can provide reliable data on the chemical situation of the surface at the time of removal. This information is required to obtain a sound basis for the interpretation of the mechanisms and kinetics of dissolution processes. In this section we have studied the general XPS spectra of chalcopyrite and also the high-resolution spectra of copper, iron and sulfur species. In some of the experiments the XPS data were obtained after Ar+ ion etching of the sample. These  114  \u000Cresults will be used as reference for the peak position of copper, iron and sulfur atoms in the chalcopyrite structure. As mentioned above, the ion etching might alter the chemical state of the surface species. In order to eliminate possible errors in the registered results, a low beam energy (1 keV) was applied (for further details refer to the experimental section). Also, all the acquired results were compared with the available results from the literature, in order to ensure the validity of the results. The survey (full range) XPS spectrum of a chalcopyrite sample is provided in Fig. 74. Fig. 7-4a presents the spectrum for polished chalcopyrite after 6 hours of air exposure which shows the carbon, oxygen, copper, iron and sulfur peaks. The binding energy of C1s at 284.6 eV is characteristic of aromatic/aliphatic carbons. The O1s peak may be related to air contamination of the samples and/or surface oxidation of the chalcopyrite. Figure 7-4b presents the spectrum of the same sample after ion etching (in order to remove the passive film) inside the XPS chamber. Etching of the sample removed the oxygen and carbon peaks from the spectrum and only copper, iron and sulfur related peaks remained.  7.2.3 Cu2p peaks in chalcopyrite Fig. 7-5a shows a typical high-resolution spectrum of Cu 2p for chalcopyrite obtained after ion etching. The fitting results of peaks by XPSPeak 4.1 software revealed that the Cu 2p3/2 peak is centered at 932.4 eV (FWHM = 1.19), and the Cu 2p1/2 peak is centered at 952.3 eV (FWHM = 2.43). The suggested binding energies for the Cu 2p3/2 and Cu 2p1/2 peaks of chalcopyrite are 932.4 eV and 952.1 eV, respectively [Nakai et al. 1978], which are in excellent agreement with our data. It is worth mentioning that the presence of Cu2+ usually results in the appearance of a satellite peak around the binding energy of 942 eV [Yin et al. 2005; Chusuei et al., 1995]. In the case of the spectrum presented in Fig. 5a, no evidence for such a satellite peak exists and subsequently the divalent state of copper (Cu2+) is not observed. This agrees with the results of other workers suggesting the presence of only monovalent copper atoms in the chalcopyrite structure [Nakai et al. 1978]. Fig. 7-5b presents the copper auger spectrum (CuL3M45M45) excited with X-rays, which is centered at 568.9 eV. This value is close to that measured by Brion [1980], which was 568.65 eV. The CuL3M45M45 values for the Cu-Cu bond, CuS and Cu2S are 568.1, 568.5 and 569.5 eV, respectively [Ghahremaninezhad et al., 2011]. 115  \u000CFigure 7-4: The survey (full range) XPS spectrum of chalcopyrite (a) after polishing and 6 hours of air exposure, (b) after ion etching inside the XPS chamber.  116  \u000CFigure 7-5: (a) XPS spectrum of Cu2p, and (b) copper Auger spectrum excited by X-rays, obtained after ion etching.  7.2.4 Fe2p and Fe3p peaks of chalcopyrite  A typical high-resolution spectrum of Fe 2p for chalcopyrite obtained after ion etching is presented in Fig. 7-6. The spectrum shows two strong peaks of Fe 2p3/2 and Fe 2p1/2 at 707.8 eV and 721.1 eV, respectively. These are related to fully coordinated Fe atoms in the chalcopyrite structure and are consistent with the results of other workers [Brion, 1980].  117  \u000CFigure 7-6: XPS spectra of (a) Fe2p, and (b) Fe3p of chalcopyrite, obtained after ion etching.  7.2.5 S2p and S2s peaks of chalcopyrite Fig. 7-7 shows the S 2p core level peak after ion etching. The S 2p peaks also occur as doublets, S 2p3/2 and S 2p1/2, as a result of spin\u2013orbit splitting. XPSPeak 4.1 software was used to fit the peaks of the S 2p spectra. The S 2p spectra were fitted using a 2:1 peak area ratio for S 2p3/2 to S 2p1/2 and 1.2 eV splitting. The peaks fitted to the spectra use a summed Gaussian-Lorentzian (SGL) function to describe the convoluted Gaussian-Lorentzian line shapes of the peaks. The weighting used was 80% Gaussian and 20% Lorentzian. The background of the spectrum was obtained using the Shirley method [Shirley 1972]. The extracted peak positions from literature were used to assign species to the peaks. The presented S 2p spectrum of Fig. 7-7 consists of two major doublets representing bulk monosulfide and disulfide. The S 2p3/2 peaks of two doublets of Fig. 7-7 are centered at 161.4 eV and 162.5 eV. The two S 2p3/2 peaks are 1.1 eV apart. It has been reported that the S 2p3/2 peak originating from fully coordinated sulfur atoms in chalcopyrite\u2019s crystal structure, i.e. monosulfide (S2\u2212), centers at 161.4 eV [Mielczarski et al., 1996] and elsewhere it has been reported to be at 161.5 eV [McCarron et al., 1990]. These results agree well with those of our experiments (S 2p3/2 peak of Fig. 7-7, centered at 161.4 eV). On the other hand, the position of the S 2p3/2 peak for the disulfide species ( S22\uF02D ) is reported to be 162.48 eV [Klauber et al., 2001], and this again agrees well with the position of the second S 2p3/2 peak 118  \u000Cin our data. In fact, it is well established that sulfur anions in the chalcopyrite structure (S2\u2212) undergo a dimerization process to form S22\uF02D species once they reach the surface of the mineral [Klauber 2003]. Hence, the detected sulfur species on the surface of chalcopyrite are sulfide species (i.e. sulfur species bonded to metallic atoms of copper and iron). In addition to the two S 2p doublets, Fig. 7-7 presents another peak positioned at 163.7 eV. This peak is a broad peak (FWHM = 2.8) and has previously been related to the formation of polysulfide ( S2n\uF02D , n &gt; 2) [Hackl et al., 1995, Klauber et al., 2001]. On the other hand, it has been shown by other researchers [Fujisawa et al., 1994; Goh et al., 2010] that this high-energy feature is a small tail related to the S 3p\u2192Fe 3d interband excitation. In other words, the tail is related to the energy-loss satellite due to the S 3p\u2192Fe 3d excitation [Chernyshova and Andreev 1997; Buckley et al., 1994]. In order to prove that this result is valid in our study, a systematic XPS study was conducted to compare the S 2p core level peaks obtained from Fe-S minerals (i.e. pyrite and pyrrhotite) and Cu-S minerals (i.e. chalcocite and covellite). Table 4-1 shows the compositions of the minerals used in this study; minerals 2 (pyrite rich) and 3 (pyrrhotite rich) were used as Fe-S minerals and minerals 4 (chalcocite rich) and 5 (covellite rich) were used as Cu-S minerals. The results are presented in Fig 7-8a to d. Figs. 7-8a, b, c and d show the high-resolution sulfur spectrum of minerals 2 (pyrite), 3 (pyrrhotite), 4 (chalcocite) and 5 (covellite), respectively. From the comparison of the figures, it is clear that the energy loss feature at high binding energy (~ 163.5 eV) is present in Fe-S minerals (Figs. 7-8a and b), while it is absent in Cu-S minerals (Figs. 7-8c and d). This result further clarifies that the high-energy feature in sulfur species spectra is most likely related to the S 3p\u2192Fe 3d excitation and does not originate from polysulfide. The S 2s core level peak of chalcopyrite after ion etching is presented in Fig. 7-9. The peak is positioned at binding energy of 225.6 eV. The S 2s peaks of the S ions in CuS and Cu2S are reported to be 225.6 and 225.3 eV, respectively; while the S 2s peak of elemental sulfur occurs at a binding energy of 228.4 eV [Chernyshova and Andreev 1997; Buckley et al., 1994]. The energy difference between the S 2s level of elemental sulfur and sulfides is appreciably wide (~3 eV) and this fact can be used to confirm sulfur formation on the surface of leached chalcopyrite samples. The XPS results obtained for a chalcopyrite sample conditioned with Ar+ ion is summarized in Table 7-1. These data can be used as the reference 119  \u000Cvalues for XPS analysis of chalcopyrite. In the next section the chalcopyrite electrodes were polarized at different anodic potentials and afterwards the outermost surface of chalcopyrite was analyzed by XPS. In order to eliminate any alteration of surface species during the XPS experiment, no argon ion etching was performed on the surface of the samples.  Figure 7-7: XPS spectrum of S 2p peak of chalcopyrite.  120  \u000CFigure 7-8: S2p core level peaks obtained from minerals (a) 2 (pyrite, FeS2), (b) 3 (pyrrhotite, Fe1-xS), (c) 4 (chalcocite, Cu2S), and (d) 5 (covellite, CuS). Composition of each mineral is presented in Table 4-1.  Figure 7-9: S2s core level peak of chalcopyrite.  121  \u000CTable 7-1: Binding energy (eV) of Cu, Fe and S elements for chalcopyrite after Ar+ ion etching.  Cu Valence I  Cu 2p 3/2  1/2  CuLMM  932.4  952.3  568.9  Binding energy of core levels (eV) Fe 2p S 2p3/2 Fe 3p 3/2 1/2 monosulfide disulfide 707.8  721.1  53.6  161.4  162.5  S 2s 225.6  7.2.6 XPS study of potentiostatically polarized chalcopyrite In this section, XPS experiments are applied to study the surface of chalcopyrite electrodes after dissolution in 0.5 M sulfuric acid solution by application of potentiostatic potentials for 6 hours. Dissolution of chalcopyrite for 6 hours will provide information regarding the composition and stoichiometry of the most stable surface film(s) on chalcopyrite. These results are useful for long term dissolution of chalcopyrite, especially the dissolution processes such as acidic leaching of chalcopyrite in hydrometallurgical processes, which typically take several hours. Note again that no Ar+ etching was performed on the experiments of this section. Fig. 7-10 presents the Cu 2p peaks of chalcopyrite electrodes obtained after 6 hours of potentiostatic polarization at 450, 650, 700, 900, 950 and 1000 mV. Since the satellite peak is absent in the Cu 2p core level peaks of Fig. 10, cupric is absent on the mineral surface even after dissolution at different potentials [Yin et al. 2005; Chusuei et al., 1995]. In Fig. 7-10, the two peaks centered at about 931.9 and 951.9 eV represent the Cu 2p3/2 and Cu 2p1/2, respectively. Both of the Cu 2p3/2 and Cu 2p1/2 are centered at lower binding energies than that of non-dissolved chalcopyrite, i.e. 932.4 eV and 952.1 eV (Fig. 7-5a). It has been well established that the Cu 2p3/2 peak with high binding energy (about 933.0\u2013933.8 eV) and the presence of a shake-up peak (about 939\u2013944 eV) are two major XPS characteristics of cupric, while a lower Cu 2p3/2 binding energy (about 931.8\u2013933.1 eV) and the absence of the shake-up peak are characteristics of reduced copper species [Avgouropoulos and Ioannides 2003; Liu and Flytzani-Stephanopoulos 1995].  122  \u000CFigure 7-10: Comparison of Cu2p peaks of chalcopyrite after dissolution in 0.5 M sulfuric acid at different potentials for 6 hours.  So again, this means that even after dissolution of chalcopyrite at different potentials between 450 and 1000 mV, surface copper exists only in the monovalent form. Probable cuprous (Cu+) compounds are cuprous oxide (Cu2O) and/or cuprous sulfides (CuS and Cu2S). The Cu 2p3/2 peak of cuprous species bonded to oxygen and sulfur in Cu2O, CuS and Cu2S is reported to be centered at 932.8 [Lefevre et al., 2000], 932.0 [Nakai et al., 1978] and 932.6 [Nakai et al., 1978], respectively. From the compareson of the Cu 2p3/2 peak value of Fig. 710 with these values, one can conclude that the most plausible cuprous specie would be CuS. This conclusion can be further confirmed by using the copper auger spectrum (Cu L3M45M45) 123  \u000Cexcited with X-rays, which is centered at 568.5 eV (Fig. 7-11). The Cu L3M45M45 values for CuS and Cu2S are 568.5 and 569.5 eV, respectively [Ghahremaninezhad et al., 2011]. This result rules out the possibility of the formation of Cu2S on the surface. It is worth mentioning that the CuLMM peak of Cu2O is reported to be at 569.7 eV [Lefevre et al., 2000], which is about 1.25 eV higher than that of Fig. 7-11.  Figure 7-11: Copper auger spectrum (CuL3M45M45) excited with X-rays obtained after potentiostatic polarization of chalcopyrite electrode at 700 mVSHE for 6 hours.  The formation of CuS on the surface of oxidized chalcopyrite has been detected by other researchers as well. In a recent study on electrochemical dissolution of chalcopyrite in 0.1 M sulfuric acid solution by Majuste et al. [2012], the formation of covellite (CuS) has been detected by Raman spectroscopy of chalcopyrite surface after electrochemical dissolution of the electrode at 0.6 and 0.8 V. Covellite (CuS) formation during electrochemical dissolution of chalcopyrite in acid chloride solutions has also been suggested by Ammou-Chokroum et al. [1979]. Elsewhere, Warren et al. [1982] have suggested that the formation of a nonstoichiometric CuS phase on massive chalcopyrite electrode surface during electrochemical dissolution in 1.0 M sulfuric acid is plausible. Yin et al., [1995] and C\u00F3rdoba et al. [2008] are among other researchers who have suggested the formation of CuS on the surface of chalcopyrite during electrochemical oxidation. Fig. 7-12 compares the Fe 2p spectra of the chalcopyrite surface after 6 hours of dissolution at 700 mV with that of Fig. 7-6 (non-dissolved chalcopyrite sample). The 124  \u000Cspectrum of etched chalcopyrite, curve a in Fig. 7-12, shows two characteristic Fe 2p3/2 and Fe 2p1/2 peaks of chalcopyrite centered at 707.8 eV and 721.1 eV, respectively. On the other hand, the Fe 2p spectrum obtained from the surface of a chalcopyrite electrode after dissolution at 700 mV for 6 hours, curve b in Fig. 7-12, has two extra peaks centered around 710 eV and 724 eV. In order to further explore the chemical state of iron species of curve b, the Fe 2p3/2 peak was fitted with three peaks centered at 707.4 (peak 1), 710.3 (peak 2) and 712.6 eV (peak 3), presented in Fig. 7-13a. Peak 1 of Fig. 7-13a at 707.4 eV shows a 0.4 eV left shift versus the Fe 2p3/2 peak of chalcopyrite (curve a of Fig. 7-12). The Fe 2p3/2 peak at 707.4 eV can be assigned to the binding energy of Fe2+ species bonded to sulfur, e.g. FeS which is 707.5 eV [Pratt et al., 1994; Jones et al., 1992] or Fe2+ in FeS2 which is reported to be 707.4 eV [Brion 1980; Pratt et al., 1994] and elsewhere is reported to be 707.5 eV [Ennaoui et al., 1986]. Unfortunately, from this information we were not able to confidently differentiate between FeS and FeS2, especially due to lack of XPS studies on FeS. It is worth mentioning that the Fe 2p3/2 peak of FeO and ferrous hydroxide are around 709.1 and 711.6 eV [Nesbitt and Muir 1994; Descostes et al., 2000], respectively. Both of these values are substantially different from the Fe 2p3/2 peak obtained in this study (i.e. 707.4 eV) which rules out the possibility of the presence of FeO and ferrous hydroxide on the surface.  Figure 7-12: Comparison of Fe2p peaks of chalcopyrite (a) after ion etching inside XPS instrument, and (b) after dissolution at 700 mVSHE in 0.5 M sulfuric acid solution.  125  \u000CThe peak at 710.3 eV can be attributed to the Fe 2p3/2 peak of FeOOH or Fe3O4 which are both centered at 710.2 eV [Nesbitt and Muir 1994; Descostes et al., 2000]. From the Fe 2p peak of Fig. 7-13a it is not possible (or very risky) to distinguish between FeOOH or Fe3O4. For this reason, the Fe 3p spectrum of chalcopyrite under the same condition of Fig. 7-13a is presented in Fig. 7-13b. Two main peaks centered at 53.3 (peak 1) and 55.5 eV (peak 2) have been used for fitting of the main spectrum (Fig. 7-13b). The peak at 53.3 eV perfectly matches with the Fe 3p peak of Fe2+ bonded to sulfur in species such as FeS2, which is 53.3 eV [Descostes et al., 2000] (here it must be mentioned that the Fe 3p value for FeS is not available; hence yet again we are not sure if FeS or FeS2 or a mixture of both species are present on the surface as iron sulfide). The peak at 55.5 eV matches well with the Fe 3p peak of Fe3+ in FeOOH which is 55.6 eV [Descostes et al., 2000]. However, the Fe 3p peak of iron species in Fe3O4 occurs at 53.7 eV [Brion 1980] which is about 1.8 eV below the Fe 3p binding energy of curve 2 in Fig. 7-13b. Thus it is most likely that after dissolution of chalcopyrite at constant potential, surface iron exists as iron sulfide and FeOOH. The reported value for the Fe 3p peak of chalcopyrite by Brion [1980] is 54.2 eV which is 0.9 eV higher than peak 1 of Fig. 7-13b and 1.3 eV lower that peak 2 of Fig. 7-13b. These values again confirm that no chalcopyrite is present on the surface. The formation of FeOOH on chalcopyrite is generally associated with high pH solutions (&gt; 2). Nevertheless, in addition to the results in this study, the formation of ferric oxy-hydroxide (FeOOH) on the surface of chalcopyrite has been previously reported. C\u00F3rdoba et al. [2008] have detected goethite on chalcopyrite by XRD, after ferric assisted leaching of chalcopyrite in sulfuric acid solution (pH = 1.8) at 35\u00BAC. Elsewhere, Kaplun et al. [2011] and Li et al [2010], have detected goethite as a chalcopyrite leaching product. Kaplun et al. [2011] detected goethite by XPS after leaching chalcopyrite at 0.75 V and pH 1 (experiment at 85 \u00BAC). Li et al. [2010], detected goethite, also by XPS, after leaching of chalcopyrite at pH 1. In another study, Pratesi and Cipriani [2000] have used XPS, Auger electron spectroscopy (AES) and Rutherford backscattering spectrometry (RBS) methods to detect the formation of goethite on chalcopyrite surfaces after air oxidation. Based on all of the above literature and our XPS results, we suggest that the formation of FeOOH in the case of our studies is possible. This will be further discussed below.  126  \u000CFigure 7-13: XPS spectrum of (a) Fe2p3/2 and (b) Fe3p obtained from a chalcopyrite electrode after 6 hours dissolution at 700 mVSHE.  The other iron species detected in this study, in addition to FeOOH, are those bonded to sulfur (iron sulfide species). In fact, it is highly plausible that these iron sulfide species may be related to a metal-deficient sulfide film formed during electrochemical oxidation. Several workers have reported the formation of a metal-deficient sulfide film (e.g. Cu1-xFe1yS2)  on the surface of chalcopyrite during dissolution in acidic media [Hackl et al., 1995;  Warren et al., 1982]. This occurs by surface depletion of metallic atoms from the chalcopyrite surface, where the depletion rate of iron atoms is higher than copper atoms due to preferential dissolution of iron (i.e. y &gt; x). Additionally, it has been suggested by several 127  \u000Cresearchers that the metal-deficient layer is the chalcopyrite passive film and is thus responsible for the slow dissolution rate of chalcopyrite [Hackl et al., 1995; Warren et al., 1982; Ghahremaninezhad et al., 2010]. In fact, it seems very probable that the detected ironsulfide species in our XPS study originates from a metal-deficient sulfide phase such as Cu1xFe1-yS2.  Cuprous sulfide species were also detected (refer to Fig. 7-10 and related  discussion). The ratio of iron sulfide to FeOOH (i.e. the ratio of Fe-S species to Fe-O species) can be obtained from the ratio of their peak areas. From the Fe 2p spectrum (Fig. 7-13a) the ratio of ferric oxy-hydroxide to iron sulfide is 87:13. Similarly, using the Fe 3p spectrum (Fig. 713b) the ratio of ferric oxy-hydroxide to iron sulfide is 85:15. The last peak of Fig. 7-13a (peak 3) is centered at 712.5 eV. This peak does not match with any cited peak of iron species [Wagner et al., 1979]. This peak has been assigned to a CuLMM auger peak excited with X-rays which is around 712.5 eV [Harmer et al., 2005; Schon 1972]. Based on this analysis, the latter peak at 712.5 eV (Fig. 7-13a) was excluded from the interpretation of the ferric oxy-hydroxide to iron sulfide ratio. Fig. 7-14 presents the Fe 2p spectra of chalcopyrite electrodes obtained after potentiostatic dissolution at different potentials. In addition to the Fe 2p3/2 peaks centered at 707.4 eV, originating from iron sulfide, the Fe 2p3/2 peak related to FeOOH (centered at 710.3 eV) is clear in all of the samples.  128  \u000CFigure 7-14: Comparison of Fe 2p peaks of chalcopyrite after dissolution at different potentials for 6 hours.  Fig. 7-15 shows a series of S 2p XPS spectra for chalcopyrite electrode surfaces obtained after dissolution at different potentials. The evolution of various sulfur species as a function of applied potential can be seen in this figure. The fitted XPS spectra of S 2p peaks of Fig. 7-15 are presented in Fig. 7-16 and 7-17. For clarity, only the S 2p3/2 peaks of the fitted curves have been displayed in the Figs. 7-16 and 7-17. The binding energies of the first S 2p2/3 core level are all in the range of 161.0 to 161.1 eV (peak 1 of Figs. 7-16 and 7-17), and this value is in agreement with what is generally found for monosulfide sulfur (S2-) species of cuprous sulfides (CuS) [Bhide et al., 1981] and/or iron sulfides (FeS) [Yu et al., 1990]. The second S 2p2/3 core level peak (peak 2 of Figs. 7-16 and 7-17) with the binding energy of about 162.1 eV agrees well with the S 2p3/2 peak position of disulfide ( S22\uF02D ) sulfur species of cuprous sulfide (CuS) [Nakai et al., 1978] and iron disulfide (FeS2) [Descostes et al., 2000]. The formation of cuprous sulfide (CuS) was suggested from the Cu 2p peak 129  \u000Canalysis (Fig. 7-10) and iron sulfides (FeS and/or FeS2) were also possible based on Fig. 7-13. It is well established in the literature that all copper atoms in the CuS phase are in the monovalent copper (Cu+) state [Nakai et al., 1978] which implies that the best formulation for the description of copper and sulfur\u2019s chemical state in covellite (CuS) is Cu 3\uF02BS 22\uF02DS \uF02D [Goh et al., 2006; Folmer and Jellinek 1980]. Additionally, it was mentioned above that the detected copper sulfide and iron sulfide species by XPS may be related to Cu-S and Fe-S bonds in a metal-deficient sulfide phase (Cu1-xFe1-yS2). The XPS spectrum of S 2p peaks at potentials of 450, 650, 700 and 870 mV (Figs. 715 and 7-16) show a tail in the high binding energy region of about 163 eV which we believe is due the S 3p\u2192Fe 3d interband excitation (peak 3 of Figs. 7-16a to d). The characteristic of this peak is its high FWHM value, i.e. larger than 2 [Klauber 2003; Acres et al., 2010]. The tail of the low potential S 2p XPS spectra of chalcopyrite obtained after potentiostatic dissolution at potentials up to 870 mV SHE has the same characteristics as the S 3p\u2192Fe 3d interband excitation, i.e. a broad peak (high FWHM) centered around 163 eV [Acres et al., 2011]. Figs. 7-16a\u2013d illustrate this point by presenting the XPS spectrum of the S 2p peak of chalcopyrite after dissolution at 450, 550 (OCP), 700 and 870 mVSHE. The binding energies and FWHM values for all of the peaks of S 2p3/2 spectra at different potentials are presented in Table 7-2. Peak 1 and peak 2 of table 7-2 are the same as explained already in Fig. 7-7. Peak 3 of Table 7-2 has a high FWHM value for the no dissolution condition (not immersed in solution) and chalcopyrite dissolved at 450, OCP, 650, 700 and 870 mVSHE and refers to S 3p\u2192Fe 3d interband excitation (also presented in Figs. 7-16a to d).  130  \u000CFigure 7-15: Comparison of S2p peaks of chalcopyrite after dissolution at different potentials for 6 hours.  131  \u000CPeak 1 Condition No dissolution 450 mV OCP 650 mV 700 mV 870 mV 900 mV 950 mV 1000 mV  132  Table 7-2: Binding energy and FWHM values for XPS spectra of S2p peaks. Peak 2 Peak 3 Peak 4  Peak 5  B.E. (eV)  FWHM  B.E. (eV)  FWHM  B.E. (eV)  FWHM  B.E. (eV)  FWHM  B.E. (eV)  FWHM  161.00 160.97 161.04 160.96 161.08 161.01 160.97 161.05 161.01  0.74 0.70 0.72 0.70 0.69 0.73 0.69 0.65 0.67  162.15 162.14 162.21 162.12 162.26 162.13 162.18 162.17 162.29  1.01 0.99 0.91 0.95 0.95 0.96 0.90 0.89 1.02  163.17 163.29 163.22 163.32 163.40 163.44 163.36 163.29 163.36  2.46 2.49 2.45 2.46 2.35 2.41 0.98 0.98 0.96  \u2013 \u2013 \u2013 \u2013 \u2013 \u2013 164.54 164.33 164.43  \u2013 \u2013 \u2013 \u2013 \u2013 \u2013 0.87 0.99 1.05  \u2013 \u2013 \u2013 168.23 168.09 168.04 168.43 168.36 168.02  \u2013 \u2013 \u2013 1.54 1.27 1.30 2.01 0.99 1.12  \u000CFigure 7-16: S2p XPS spectra for chalcopyrite electrode surface obtained after dissolution at (a) 450, (b) OCP (~550 mVSHE), (c) 700 and (d) 870 mVSHE illustrating the characteristics of the S 3p\u2192Fe 3d interband excitation related peaks.  133  \u000CIn the case of high applied potentials (i.e. 900, 950 and 1000 mV SHE) the spectra are significantly different from those obtained at low potentials (Fig. 7-15). From the potentiostatic polarization study presented in Fig. 7-2a, it was clear that, for the potential range up to about 900 mVSHE, chalcopyrite is in its passive state. The potentials of 900 mVSHE and above are transpassive dissolution potentials, i.e. the dissolution rate of chalcopyrite increases drastically. Hence, one would expect a substantial change in the sulfur species on the surface of chalcopyrite for potentials above 900 mVSHE. It is clear from Fig. 715 that the two peaks at ~ 161.0 and 162.1 eV, are still present; this fact can be seen in Table 7-2. However, the area under the peaks has decreased. In addition, two extra peaks centered at higher binding energies of 163.3 and 164.4 eV have grown (Fig. 7-15). The fitted XPS spectra of S 2p peaks for chalcopyrite after dissolution at 900, 950 and 1000 mVSHE are presented in Figs. 7-17a\u2013c (For clarity, only S 2p3/2 peaks of fitted curves have been presented). Each XPS spectrum of Fig. 7-17 is fitted with 5 peaks. Peaks 1 and 2 are the same peaks seen in Fig. 7-16 (related to iron sulfide and copper sulfide bonds, which most probably originated from a metal-deficient sulfide film on the surface). The last peak (peak 5) of Figs. 7-17a\u2013c, centered at a very high binding energy of ~168.5 eV, is related to sulfate and most probably originates from sulfate phase formation [Klauber 2008] or electrolyte (sulfuric acid) residue on the surface [Bhide et al., 1981]. This peak was also present on the samples polarized at 650, 700 and 870 mVSHE (Table 7-2 and Figs. 7-16c and d). Peak 3 of Figs. 7-17a\u2013c has similar binding energy to that of the last peaks of Figs. 7-16a\u2013d, however the FWHM is smaller. The binding energy for peak 3 is between 163.3 and 163.4 eV and the FWHM is between 0.96 and 1.04, which is a smaller FWHM value than those obtained at low potentials (~ 2.5). We suggest that this S 2p 3/2 peak (peak 3 in Figs. 7-17a\u2013c) may be related to polysulfide ( S2n\uF02D ) formation. This is in agreement with the literature, which suggests the formation of polysulfide sulfur species around the binding energy 163.3 eV and low FWHM of around 1.0 [Hackl et al., 1995]. The S 2p3/2 peak 4 of Figs. 7-17a\u2013c, centered around the binding energy of 164.4 eV (Table 7-2) is most probably related to the S 2p3/2 of elemental sulfur as the binding energy of this peak is in agreement with the value reported for elemental sulfur (164.4 eV) [Manocha and Park 1977]. The reason for the absence of the S 3p\u2192Fe 3d interband excitation peak in the polarized electrodes at high anodic potentials  134  \u000C(Fig. 7-17) is not clear; however, we suspect that it may be related to elemental sulfur formation on the surface. A better investigation of the formation of elemental sulfur requires further study of the XPS spectra for S, and the S 2s peak provides a convenient way of doing this. The binding energy of chalcopyrite\u2019s S 2s peak is measured to be 225.6 eV (Fig. 7-9). The S 2s peaks of the S atoms in CuS, Cu2S and FeS2 are reported to be 225.6, 225.3 and 226.1 eV, respectively; while the S 2s peak of elemental sulfur occurs at a binding energy of 228.4 eV [Chernyshova and Andreev 1997; Panzuner and Egert 1984; Buckley et al., 1994]. Hence, the broad difference between the S 2s level of sulfides vs. elemental sulfur can be used for detection of elemental sulfur formation. Fig. 7-18 presents XPS spectra of S 2s peaks of chalcopyrite after dissolution at different potentials for 6 hours. At low potentials, the S2s peaks are narrow and, by increasing the potential above 900 mVSHE, the peaks become wider. From Fig. 7-18 it is clear that the binding energy of the maximum intensity point (peak point) is almost the same for all the potentials below 870 mV SHE and by increasing potential to 900 mVSHE (or above), the binding energy of the peak point shifts to higher energies. This is due to the formation of elemental sulfur on the surface, which has higher binding energy for the S 2s level. The above study on the S 2p and S 2s XPS spectra of chalcopyrite (Figs. 7-15 to 718) suggest that for elemental sulfur to be formed as a direct product of electrochemical chalcopyrite dissolution, a fairly high potential of 900 mV SHE or above is required. Warren et al. [1982] have proposed the formation of elemental sulfur as a product of chalcopyrite dissolution at high potentials (E &gt; 0.7 V): CuFeS 2 \uF03D Cu 2 \uF02B \uF02B Fe 3\uF02B \uF02B 2S o \uF02B 5e \uF02D  7-2  Additionally, they have mentioned that this elemental sulfur may be the product of covellite (CuS) dissolution at high potentials (E &gt; 0.9 VSHE), where covellite was considered to be the product of chalcopyrite dissolution at lower potentials. It was suggested by Warren et al. [1982] that during the electrochemical dissolution of chalcopyrite in 1 M sulfuric acid solution at 25\u00BAC, a nonstoichiometric covellite film forms and passivates the chalcopyrite surface. Our results also agree with the results of Yin et al. [1995], as they did not report the formation of elemental sulfur during the dissolution of chalcopyrite electrodes in 0.1 and 1 M HCl solutions at 25\u00BAC. Yin et al. [1995] have shown that, in strongly acidic solutions (pH &lt; 135  \u000C2) at potentials below 0.77 VSHE, dissolution of chalcopyrite proceeds through reaction (1), and for higher potentials (up to 1.04 VSHE) the following reaction is proposed [Yin et al. 1995]: CuFeS 2 \uF02B 8H 2 O \uF03D Cu 2 \uF02B \uF02B Fe 3\uF02B \uF02B 2SO 24 \uF02D \uF02B 16H \uF02B \uF02B 17e \uF02D  7-3  Nevertheless, neither reaction (1) nor (3) suggests the formation of elemental sulfur. A recent synchrotron small-angle X-ray diffraction study by Majuste et al. [2012] has shown that polarization of a chalcopyrite electrode in 0.1 M sulfuric acid solution at 0.8 and 1.0 VSHE yields covellite (CuS) and elemental sulfur (S8) on the chalcopyrite electrode surface. These researchers did not detect elemental sulfur at low anodic potentials of 0.6 and 0.7 VSHE. These results are also consistent with the observations of the present study.  136  \u000CFigure 7-17: S2p XPS spectra for chalcopyrite electrode surface obtained after dissolution at (a) 900, (b) 950, and (c) 1000 mVSHE.  137  \u000CFigure 7-18: Comparison of S2s peaks of chalcopyrite after dissolution at different potentials for 6 hours. Red dashed-line represents the binding energy of S2s peak of elemental sulfur, 228.4 eV.  Fig. 7-19 shows the speciation of sulfur versus potential on the outermost surface of chalcopyrite after 6 hours of dissolution at different constant potentials. To obtain the data of Fig. 7-19, it was assumed that the percentage of each species ( S2\uF02D , S22\uF02D , S2n\uF02D , elemental S and SO 24\uF02D ) is directly proportional to the area under its S 2p peak presented in Figs. 7-15, 7-16  and 7-17. It is worth mentioning that for the experimental data from the 550 mV SHE condition presented in Fig. 7-19, no potential was applied to the electrode; indeed, the electrode was immersed in solution at open circuit and throughout the 6 hours of immersion the potential was fairly stable at ca. 550 mVSHE. From Fig. 7-19, the major species of sulfur at low potentials (450 to 870 mV SHE) are monosulfide ( S2\uF02D ) and disulfide ( S22\uF02D ). Additionally, about 5% of sulfate ( SO 24\uF02D ) species form in this potential interval as well. Interestingly, by increasing the potential from 450 mV to 870 mV SHE, the fraction of monosulfide increases from 43 to 56% of total sulfur while the fraction of disulfide decreases from 57 to 39%. In the passive potential region, 650 &lt; E &lt; 900 VSHE, the ratio of different sulfide species ( S2\uF02D / S22\uF02D ) remains relatively constant (at c.a. 1.4). This point can be justified by considering 138  \u000Cthe passive state of chalcopyrite in this potential region; in this region a passive sulfide film is formed on the surface, which has approximately the same composition throughout all passive potentials. Note from Fig. 7-19 that in the transpassive region, E &gt; 0.90 VSHE, the fraction of monosulfide and disulfide species show a notable decrease (monosulfide decreases from 56% to 19% and disulfide decreases from 39% to 18%) and the fraction of polysulfides and elemental sulfur shows a sudden increase (polysulfide increase from 0 to 35% and elemental sulfur increases from 0 to 21%). From Fig. 7-19, the concentration of polysulfide and elemental sulfur for potentials below 0.90 VSHE is zero and the concentration of sulfate at potentials below OCP is negligible. Considering the fact that transpassive dissolution of chalcopyrite occurs at potentials greater than 0.90 VSHE, these findings clearly demonstrate that monosulfide and disulfide species are responsible for passivation of chalcopyrite and elemental sulfur would have an insignificant effect on passivation under the conditions tested here. As concluded from the XPS analysis above, copper sulfide (CuS) and iron sulfide bonds were detected on the surface of the electrochemically dissolved chalcopyrite electrode (Figs. 7-10 and 7-13). The literature has previously indicated that copper sulfide and iron sulfide bonds likely originate from the metal-deficient sulfide film which forms on the surface [Warren et al., 1982; Hackl et al., 1995; Viramontes-Gomboa et al., 2007]. If true, this phase should be considered to constitute the passive film on the surface. The following reaction may then be suggested for passivation of chalcopyrite where a metal-deficient sulfide film is the passivating film: CuFeS 2 \uF03D Cu 1\uF02D x Fe1\uF02D y S 2 \uF02B xCu 2\uF02B \uF02B yFe 2 \uF02B \uF02B 2(x \uF02B y)e \uF02D  y&gt;x  7-4  Several other researchers have concluded that a metal-deficient sulfide passive film forms via reaction (7-4) or very similar reactions [Warren et al., 1982; Hackl et al., 1995; Viramontes-Gomboa et al., 2007]. In all of these articles it has been acknowledged that y should be larger than x due to preferential dissolution of iron atoms to copper atoms.  139  \u000CFigure 7-19: Sulfur species on the surface of chalcopyrite formed after dissolution in 0.5 M sulfuric acid solution at different potentials for 6 hours.  Nevertheless, reaction (7-4) or similar reactions such as reaction (7-1) may not be the best way to describe the dissolution of chalcopyrite. The reason for this is the accumulation of sulfur (in sulfide form) by reaction (7-4) and the accumulation of both copper and sulfur by reaction (7-1) on the surface of the electrode. During dissolution by reactions (7-4) and (7-1), as the sulfide film grows on the surface (CuS2 in reaction 7-1 and metal-deficient sulfide layer in reaction 7-4), it will result in thickening of the film which, in turn, will cause a decrease of the dissolution current density (during potentiostatic polarization of the electrode at a passive potential, OCP &lt; E &lt; 0.9 VSHE). From Fick\u2019s law of diffusion, one would estimate that if the passive film thickness doubles, then the dissolution current density of chalcopyrite must decrease by half. However, Fig. 7-3b clearly shows that even after 60 hours of potentiostatic dissolution of chalcopyrite at 0.80 VSHE, the current density remains approximately constant.  This means that sulfur species (i.e. monosulfide and disulfide  species) and copper ions cannot accumulate on the surface and must dissolve instead. This process can only be described if all the elements of chalcopyrite (i.e. Fe, Cu and S) are considered to dissolve during the dissolution of chalcopyrite. In fact, through a ring\u2013disk electrode study on a natural chalcopyrite electrode in 0.1 M H2SO4 solution, Lazaro and Nicol [2006] proposed the formation of soluble sulfur species such as thiosulfate. In their study, different potentials (in the passive region) were applied on the chalcopyrite disk for 1 h while the platinum ring was held at a constant cathodic potential 140  \u000Cof \u20130.06 VSHE.  Under this condition, any oxidized species of chalcopyrite dissolution  emanating from the disk were reduced on the platinum ring (except ferrous, as the potential for its reduction to metallic iron is very low). Afterwards, during the anodic stripping of the platinum ring, in addition to the peak for the stripping of copper metal, two other peaks were obtained. The potential of the latter two peaks were in the potential region similar to those that are known to be associated with the oxidation of the copper sulfides. Since different copper sulfides can be oxidized in this potential region, it was not easy to assign these peaks to a particular copper sulfide (e.g. CuS, Cu2S and CuxS). Nevertheless, since these peaks are representative of the oxidation of copper sulfides, then in addition to copper ions, it was suggested that soluble sulfur species such as thiosulfate ( S 2 O 23 \uF02D ) and/or hydrogen sulfide (H2S) could be formed in the dissolution reaction at the disk [Lazaro and Nicol 2006]. Hence, if one assumes that thiosulfate is a soluble product of chalcopyrite dissolution, then a reaction such as the following may be a good description for chalcopyrite dissolution in the passive region: CuFeS 2 \uF03D Cu 1\uF02D x Fe 1\uF02D y S 2-z \uF02B xCu 2 \uF02B \uF02B yFe 2 \uF02B  z \uF02B S 2 O 32 \uF02D \uF02B 3zH \uF02B \uF02B 2(x \uF02B y \uF02B z)e \uF02D 2  y&gt;x  7-5  In reaction (7-5), a portion of the iron, copper and sulfur atoms dissolve as cupric, ferrous and thiosulfate species while the metal-deficient sulfide film ( Cu1\uF02D x Fe1\uF02D y S 2-z ) passivates the electrode and hinders the dissolution rate. In fact, reaction (7-5) is similar to reaction (7-4), but reaction (7-5) also accounts for the proposed dissolution of sulfur. It is worth mentioning here that thiosulfate is known to be a major dissolution product of some other sulfide minerals such as pyrite (FeS2) [Mishra and Osseo-Asare 1988; Rohwerder et al., 2003]. In addition to thiosulfate, one would also consider H2S as a potential soluble sulfur product. The interaction between water molecules and the (001) plane of chalcopyrite has recently been studied by means of density functional modeling [Ferreira de Lima et al., 2011]. This work has suggested that hydrogen sulfide gas may evolve at anodic potentials. These authors show that in a highly acidic solution, protons adsorb to the surface of the disulfides ( S22\uF02D species) and this causes elongation of the S\u2013S bond and formation of H2S  141  \u000Cspecies. Additionally, it has been concluded by Ferreira de Lima et al., [2011] that the adsorption of positively charged protons to sulfur dimers on the chalcopyrite surface inhibits the initiation of the polysulfide formation, and H 2S forms instead. The solubility of H2S in water is very low and has a very strong foul odor (50% of humans can recognize the characteristic odor of hydrogen sulfide at concentrations as low as 0.005 ppm), and we have never noticed this odor during the course of our experiments. Nevertheless, further experimental studies are warranted for analysis of the soluble sulfur species formed during chalcopyrite dissolution. Fig. 7-20 presents the concentration of various elements (oxygen, sulfur, iron and copper) on the surface of chalcopyrite as a function of potential. Each element of Fig. 7-20 reflects its total concentration after 6 hours of dissolution. From Fig. 7-20, the concentration of copper, iron and sulfur remains relatively constant on the surface in the passive potential region (650 to 900 mV SHE). This result would suggest that a passive film with a defined and time-invariant composition (e.g. Cu1\uF02D x Fe1\uF02D y S2-z ) is responsible for hindering the rate of chalcopyrite dissolution in this potential region. This point agrees with our results presented above. However, oxygen concentration increases from ca. 24% to 42% upon increasing the potential from OCP (~ 550 mVSHE) to 870 mVSHE.  Figure 7-20: Concentration of O, S, Cu and Fe species on the outer-most surface of chalcopyrite after dissolution at different potentials for 6 hours in 0.5 M sulfuric acid solution. C contamination is excluded from the analysis.  142  \u000CTo further elucidate the role of oxygen on the passivation of chalcopyrite, its surface speciation is presented in Fig. 7-21 as a function of different potentials. To obtain Fig. 7-21, the XPS spectra of the O 1s peak of oxygen at different potentials were fitted with three peaks; one peak at 529.6 \u00B1 0.1 eV corresponding to O2\uF02D in oxide phases [Tan et al. 1990], one at 531.5 \u00B1 0.1 eV corresponding to OH \uF02D in hydroxide species [Clayton and Lu 1986] and the last one at high binding energy of 532.8 \u00B1 0.1 eV corresponding to oxygen in sulfate or/and water [Shuxian et al., 1986]. Since oxygen in sulfate and water gives rise to peaks at the same binding energy, we were not able to interpret water and sulfate separately and thus we have excluded water from the data presented in Fig 7-21. A sample of the fitting process for XPS spectra of the O1s peak of chalcopyrite obtained after 6 hours of potentiostatic dissolution at 650 mV is presented in Fig. 7-22. From Fig. 7-21, a major part of the oxygen present at all potentials is in hydroxyl form ( OH \uF02D ). Fig. 7-21 shows that the concentration of sulfate and/or adsorbed water on the surface increases on the surface as potential increases (450 &lt; E &lt; 900 mVSHE) and, in contrast, the concentration of O2\uF02D and OH \uF02D species is almost constant. This fact would imply that the observed increase of total oxygen fraction on the surface in Fig. 7-20, is most likely due to an increase in sulfate concentration and adsorbed water. This finding is consistent with that of Fig. 7-19 which shows the formation of a sulfate phase on the surface as potential is increased. In addition, the adsorbed water on the electrode constitutes further evidence for the formation of FeOOH on the surface of the electrode. As noted above, the formation of FeOOH is generally associated with pH higher than that used in our experiments (above pH = 2) [Pourbaix 1966]. It appears that adsorption of water to the iron-sulfide species (Fe-S bonds were detected by XPS, as shown above) of the passive film (highly hydrated passive film) provides an environment of H2O and OH \uF02D (lattice or bound water) around Fe species analogous to a high pH environment where the FeOOH phase can be stable. This suggested path for the formation of FeOOH agrees reasonably well with the anodic dissolution model of pyrite in acidic media suggested by Mishra and Osseo-Asare [1988], which assumes that the first step of pyrite dissolution involves the deprotonation of adsorbed water:  143  \u000CFeS 2 \uF02B H 2 O \uF03D Fe(OH)S 2 \uF02B H \uF02B \uF02B e \uF02D  7-6  This model [Mishra and Osseo-Asare 1988] predicts the formation of thiosulfate as a dissolution product. Hence, we would suggest that the FeOOH species on the surface of the dissolved chalcopyrite (detected by XPS) may originate from the dissolution of the sulfurbonded iron species of the passive film (i.e. Fe-S species). This pathway for formation of FeOOH is very similar to the proposed model for hydration of molybdenum and chromium atoms in the passive film of stainless steel in acidic solution (0.1 M HCl) which suggests the formation of highly hydrated species such as MoO(OH)2 , MoO 24\uF02D , Cr(OH) 3 and CrO 24\uF02D [Clayton and Lu 1986]. The latter four species are thermodynamically stable only in high pH solutions, i.e. higher than that of Clayton and Lu\u2019s study. We conclude it is likely that, during the dissolution of chalcopyrite in acidic solutions, water molecules adsorb to its surface and provide an environment of H2O and OH\u00AF, which facilitate the formation of FeOOH.  Figure 7-21: Oxygen species on the surface of chalcopyrite formed after dissolution at different potentials for 6 hours.  144  \u000CFigure 7-22: O1s core level peak obtained from chalcopyrite after 6 hours of dissolution at 650 mVSHE.  7.3 Conclusion In summary, the surface layers formed during the electrochemical dissolution of chalcopyrite in 0.5 M sulfuric acid solution were studied. It was shown that potentiodynamic polarization of chalcopyrite, even at a very low scan rate, cannot successfully demonstrate the active-passive dissolution of chalcopyrite. On the other hand, the results of potentiostatic polarization have agreed very well with those of XPS studies, which suggest that this technique can accurately demonstrate the active-passive dissolution behavior of chalcopyrite. From potentiostatic polarization it was concluded that chalcopyrite is passive at potentials below 0.90 VSHE, beyond which it actively dissolves. Results of XPS studies have shown a substantial change in the oxidation state of sulfur species on the surface of chalcopyrite at potentials above 0.90 VSHE. XPS experiments revealed that copper and iron sulfide bonds are present in the passive film of chalcopyrite, but their binding energies are different from those of chalcopyrite. These bonds were believed to originate from a metal-deficient sulfide passive film with composition Cu1\uF02D x Fe1\uF02D y S2-z . In addition, FeOOH was detected on the surface and a possible reaction mechanism for its formation was suggested. XPS results showed that the surface of passive chalcopyrite is blocked by a sulfide film and beyond the passive region (above 0.90 VSHE) the percentage of sulfides in the film decreased significantly, i.e. monosulfide decreased from 56% to 19% and disulfide decreased from 145  \u000C39% to 18%. Furthermore, no elemental sulfur or polysulfide was detected on the surface at potentials below 0.90 VSHE. Hence, it is concluded that the transpassive dissolution of chalcopyrite is significantly linked to the oxidation of sulfur (from sulfide in the passive film to elemental sulfur and perhaps to sulfur species with higher oxidation states, e.g. thiosulfate). Additionally, by comparing XPS sulfur spectra of chalcopyrite with the results of sulfur spectra from iron sulfide (pyrite and pyrrhotite) and copper sulfide minerals (chalcocite and covellite) we have concluded that the high energy feature in sulfur species spectra (~ 163.3 eV) is most likely related to the S 3p\u2192Fe 3d excitation and does not originate from polysulfide. In the next chapter the multi-cyclic voltammetry method will be used to further understand the dissolution mechanism of chalcopyrite in 0.5 M sulfuric acid solution. This study will help us to suggest proper electrochemical reactions for chalcopyrite dissolution.  146  \u000C8 Multiple cyclic voltammetry study of chalcopyrite in 0.5 M sulfuric acid solution 8.1 Introduction In the last two chapters, the surface of the chalcopyrite electrode was studied mostly by electrochemical impedance spectroscopy and X-ray photoelectron spectroscopy. EIS studies in chapter 6 revealed the formation of two distinct surface layers on chalcopyrite after dissolution in 0.5 M sulfuric acid solution. Also, EIS results have suggested that the outermost surface layer must be the product of the dissolution of the inner passive film. The formation of a metal deficient sulfide passive film on the chalcopyrite surface was further confirmed by XPS studies in chapter 7. In this chapter, cyclic voltammetry will be used as an additional method to understand the dissolution mechanism of chalcopyrite in 0.5 M sulfuric acid solution.  8.2 Results and discussion 8.2.1 Effect of oxygen content of the solution on the CV experiments Figure 8-1 presents the multicyclic voltammograms for a stationary chalcopyrite electrode in both deaerated (Fig. 8-1a) and aerated (Fig. 8-1b) 0.5 M H2SO4 solutions. In the case of deaerated solution, argon gas was sparged into the solution for 1 hour and was stopped about 2 minutes prior to the experiment. It can be seen from the comparison of these two figures that the position of peak currents are almost identical. All of the anodic and cathodic peaks either increase with time (i.e., the number of cycles) or are invariant in time. None of the peaks vanish. Both experiments show seven anodic peaks at 0.02, 0.10, 0.19, 0.24, 0.35, 0.65 and 0.79 VSHE and four cathodic peaks at 0.37, 0.30, \u22120.10 and \u22120.20 VSHE. Hence throughout the rest of this paper, only the results of aerated solution will be discussed. In addition, each of the current peaks of Fig. 8-1 will be explained in depth.  147  \u000CFigure 8-1: Multicyclic voltammogram for stationary chalcopyrite electrode in 0.5 M H2SO4 deaerated (a) and aerated (b) solutions. Scan rate is 50 mV s-1.  8.2.2 Cathodic reactions on the surface of chalcopyrite In general, during the multicyclic voltammetry experiments, anodic and cathodic peaks may grow or vanish in time (i.e. number of cycles) or be insensitive to the number of cycles. The response of current peaks to the number of cycles is indicative of the depletion or accumulation of electroactive species in the vicinity of the electrode.  148  \u000CFigure 8-2a presents a typical multicyclic voltammogram for stationary chalcopyrite in 0.5 M H2SO4 solution (quiescent condition). In this experiment, the chalcopyrite electrode was polarized at 0.55 V (very close to its OCP) for 20 seconds and afterwards, the potential of the electrode was increased linearly to 0.75 V, followed by 10 cycles from 0.75 V to \u20130.5 V and back at a scan rate of 50 mV s\u22121. The evolution of the current peaks as the number of cycles is increased is clear. The current for all of the cathodic and anodic peaks increases with an increasing number of cycles. In the tenth cycle, anodic peaks occur at 0.01, 0.09, 0.19, 0.24, 0.35, 0.64 and 0.84 V and cathodic peaks occur at 0.36, 0.28, \u22120.11 and \u22120.20 V. In order to understand the origin of the cathodic peaks, the tenth cycle of the experiment was compared to the tenth cycle of a rotating (500 rpm) chalcopyrite disk electrode in Fig. 8-2b. If a peak originated directly from the oxidation or reduction of chalcopyrite or other solid phases on the surface then it remains in the voltammogram for the rotating electrode; however, peaks originating from the oxidation or reduction of dissolved species should be significantly reduced or entirely absent. From Fig. 8-2b, it is clear that anodic peaks a1, a2, a3 and a4 are related to the oxidation of dissolved species (these peaks are discussed in more detail in section 8.2.3). The same explanation can be provided for cathodic peaks c1 and c2, which may be related to the reduction of ferric and cupric dissolved during the anodic scans (the assigned reactions to all cathodic peaks are presented in Table 8-1). In the case of the stationary electrode, iron and copper dissolve and remain near the electrode surface (quiescent condition), becoming involved in reduction reactions at low potential. In Fig. 8-2a, peak c1 is present from the second cycle and grows with each additional cycle; however, peak c2 is not observed in the first two cycles. Peak c1 can be related to ferric reduction: 3\uF02B  \uF02D  Fe \uF02B e \uF0AE Fe  2\uF02B  [Fe 3\uF02B ] E (VSHE ) \uF03D 0.771 \uF02B 0.0591 log [Fe 2\uF02B ]  8-1  Under the conditions of our experiment (dissolution of chalcopyrite up to 0.75 V, Fig. 8-2a), iron dissolves mostly as ferrous. Hence in the first cycle the cathodic peak c1 is not observed. Nevertheless, after the first cycle, the ferrous concentration in the vicinity of the electrode surface increases rapidly (ferrous ions form during both anodic and cathodic dissolution of chalcopyrite). Hence, the equilibrium potential of reaction 8-1 (cf. Table 8-1) will decrease significantly and even in the high regions of the potential window of Fig. 8-2a,  149  \u000Cferric will be stable. Therefore, after the second cycle and for all further cycles, peak c1 appears in Fig. 8-2a.  Figure 8-2: Typical multicyclic voltammogram of stationary chalcopyrite in 0.5 M H2SO4 solution (a) and effect of the rotation of the electrode on the evolution of the peaks (b). Scan rate is 50 mV s-1.  150  \u000CTable 8-1: Reactions associated with cathodic peaks of chalcopyrite dissolution. peak  Reaction number  Reaction  Reference  c1  8-1  Fe3\uF02B \uF02B e\uF02D \uF0AE Fe2 \uF02B  This study  c2  8-4  CuFeS 2 \uF02B 3 Cu 2 \uF02B \uF02B 4 e \uF02D \uF0DE 2 Cu 2 S \uF02B Fe 2 \uF02B  Biegler and Horne 1985  8-5  2 CuFeS2 \uF02B 6 H \uF02B \uF02B 2 e \uF02D \uF0AE Cu 2S \uF02B 3 H 2S(aq) \uF02B 2 Fe 2 \uF02B  Biegler and Swift 1979  8-6  9 CuFeS2 \uF02B 4 H \uF02B \uF02B 2 e \uF02D \uF0AE Cu 9 Fe8S16 \uF02B 2 H 2S(aq) \uF02B Fe 2 \uF02B  Hiroyoshi et al., 2002  8-7  5 CuFeS2 \uF02B 12 H \uF02B \uF02B 4 e\uF02D \uF0AE Cu 5 FeS4 \uF02B 6 H 2S(aq) \uF02B 4 Fe 2 \uF02B  Nava and Gonzalez 2006  c3/c4  151  \u000CPeak c2 of Fig. 8-2a has been assigned by various authors to the reduction of a solid ferric sulfate phase [Nava and Gonzalez 2006], cupric reduction to copper [Holliday and Richmond 1990; Eghbalnia and Dixon 2001], and covellite reduction to chalcocite [Yin et al., 1995]. Nava and Gonzalez have assigned the following reaction for peak c2 of Fig. 8-2a [Nava and Gonzalez 2006]: Fe 2 (SO 4 ) 3(solid) \uF02B 2 e \uF02D \uF0AE 2 Fe 2 \uF02B \uF02B 3 SO 24\uF02D  E ( VSHE ) \uF03D 0.795 \uF02D 0.0591 log[ Fe 2 \uF02B ] \uF02D 0.0887 log[SO 24\uF02D ]  8-2  Argument against the validity of reaction 8-2 is that peak c2 must be related to a reduction reaction which involves at least one soluble reactant, since, from Fig. 8-2b, the peak is absent in the case of a rotating electrode. However, reaction 8-2 consumes no soluble species during reduction. In order to investigate peak c2 further, several cyclic voltammetry experiments were performed on a chalcopyrite electrode after four different conditioning procedures. First, the electrode was polarized at 0.55 V (~OCP) for 20 s. Next, the potential was scanned linearly up to 0.75 V, where it was polarized for 0, 10, 100, or 1000 s, after which cyclic voltammetry was performed from 0.75 to 0.1V and back to 0.75 V. The results of these experiments are presented in Fig. 8-3. From Fig. 8-3 the cathodic reaction seems to be very much related to the anodic reaction which follows at 0.47 V. Yin et al. [1995] and Beigler and Horne [1985] have acknowledged that the associated reactions to these peaks must be the forward and reverse portions of a reversible reaction. Indeed, from the observed behavior of the reactions in Fig. 8-3, we also believe that both anodic and cathodic peaks are related to a reversible reaction with a reversible potential of ca. 0.4 V. Additionally, it should be pointed out that the maximum positive potential in this experiment (Fig. 8-3) is 0.75 V which is well below the transpassive dissolution potential of chalcopyrite (~0.90 V) [Ghahremaninezhad et al., 2010; results of chapter 7]. In this potential region mostly ferrous and ferric will be released to solution rather than copper ions; Nova and Gonzales have shown that the concentration of dissolved copper is low and anodic stripping voltammetry (ASV) was unable to detect copper in solution [Nava and Gonzalez 2006]. At this low copper concentration level, the reversible potential of cupric reduction will be well below the reversible potential of Fig. 8-3 (ca. 0.4 V). Hence, cupric reduction to metallic copper is not a valid reaction for peak c2 of Fig. 8-2a. 152  \u000CFigure 8-3: Cyclic voltammetry of chalcopyrite electrode between 0.75 to 0.1 VSHE with (1) 0, (2) 10, (3) 100 and (4) 1000 seconds of potentiostatic polarization at 0.75 VSHE before cyclic voltammetry experiment. Scan rate is 50 mV s-1.  The other reaction suggested in the literature is reduction of covellite [Yin et al. 1995]: CuS \uF02B Cu 2 \uF02B \uF02B 2 e \uF02D \uF0AE Cu 2 S  E ( VSHE ) \uF03D 0.493 \uF02B 0.0295 log[ Cu 2\uF02B ]  8-3  Yin et al. [1995] have proposed that CuS and cupric, which are products of chalcopyrite oxidation [Yin et al. 1995], form in the positive-going cycle and that these are involved in the cathodic direction of reaction 6 during the negative-going cycle. If reaction 6 were valid, then one would expect to observe a cathodic peak in Fig. 8-3 with increasing magnitude as the polarization time at 0.75 V was increased from 0 to 1000 s (due to the formation of more CuS and Cu2+ by further oxidation of chalcopyrite at 0.75 V). However, the results of Fig. 8-3 present exactly the opposite scenario; by increasing the polarization time at 0.75 V from 0 to 1000 s the peak is almost eliminated. This evidence calls the validity of reaction 6 into question. Lastly, Biegler and Horne [1985] proposed the formation of chalcocite by chalcopyrite reduction according to the following reaction: CuFeS 2 \uF02B 3 Cu 2 \uF02B \uF02B 4 e \uF02D \uF0AE 2 Cu 2 S \uF02B Fe 2\uF02B  E ( VSHE ) \uF03D 0.700 \uF02B 0.0443 log[Cu 2 \uF02B ] \uF02D 0.0148 log[ Fe 2\uF02B ]  8-4  153  \u000CIn order to further evaluate this reaction, Biegler and Horne [1985] investigated the influence of 0.5 mM of ferric, ferrous and cupric addition on the magnitude of the cathodic peak of Fig. 8-3. Addition of ferric and ferrous had no effect; however, cupric addition increased the magnitude of the peak significantly. Now, it remains only to justify why the potentiostatic polarization of chalcopyrite at 0.75 V before a cyclic voltammetry experiment eliminates both the anodic and cathodic peaks (refer to Fig. 8-3) of reaction 8-4. The potential of 0.75 V is a passive potential for chalcopyrite [Ghahremaninezhad et al. 2010]; hence, during polarization of the electrode at this potential, a passive film forms on the mineral surface. Anodic polarization of chalcopyrite at 0.75 V extends the area of the passive surface of chalcopyrite, decreasing the exposed surface of chalcopyrite to solution as polarization time is increased. Hence, the extent of reaction 8-4 will decrease as a function of anodic polarization time at 0.75 V. This means that less chalcocite (Cu2S) will be produced by reaction 8-4 during a negative-going cycle, and therefore less chalcocite will be available for oxidation during the following positive-going cycle; i.e., a decrease in the magnitude of the cathodic peak will lower the magnitude of the corresponding anodic peak. As a final validation of reaction 8-4, a rotating chalcopyrite disk electrode (500 rpm) was polarized linearly from 0.55 V (ca. OCP) to 0.75 V, and two successive potential scan cycles were then applied between 0.75 V and 0.1 V (the same potential window as in Fig. 83). The result of this experiment is presented in Fig. 8-4, which clearly shows that both the cathodic and anodic peaks are eliminated by rotation of the electrode. This is due to diffusion of both cupric and ferrous away from the electrode which eliminates the possibility of reaction 8-4.  154  \u000CFigure 8-4: Cyclic voltammetry of a rotating chalcopyrite disk electrode (500 rpm) between 0.75 and 0.1 V. Scan rate is 50 mV s\u22121.  Fig. 8-5 presents the cyclic voltammogram of a chalcopyrite electrode in the potential window from \u22120.1 to 1 V. Increasing the anodic vertex potential during cyclic voltammetry of chalcopyrite electrodes to potentials above 0.9 V shifts the cathodic peak of Fig. 8-3 to more negative potentials, i.e. from 0.33 V to 0.29 V. In addition, increasing the vertex potential from 0.75 V (Fig. 8-3) to 1.0 V (Fig. 8-5) clearly increases the magnitude of the cathodic peak from about 125 \u03BCA cm\u22122 (Fig. 8-3) to about 470 \u03BCA cm\u22122 (Fig. 8-5). This increase in the magnitude of the peak may be related to the increase in the concentration of cupric (reactant of reaction 8-4 \u2212 cf. Table 2) in the vicinity of the electrode, provided by increasing the anodic dissolution potential up to 1.0 V.  155  \u000CFigure 8-5: Cyclic voltammogram of a chalcopyrite electrode in the potential window of \u22120.1 to 1 V. Scan rate is 50 mV s\u22121.  Fig. 8-6 presents the first, second and tenth CV cycles for a stationary chalcopyrite electrode. In this experiment, the electrode was polarized at 0.55 V (ca. OCP) for 20 seconds and then scanned up to 0.75 V and finally 10 cycles were performed between 0.75 and \u22120.5 V. It is clear that both peaks c3 and c4 are absent in the first cycle, appearing only upon (and after) the second cycle. The current of peak c3 remains almost constant regardless of the number of cycles while peak c4 grows as the cycling is continued (Fig. 8-6). Additionally, from Fig. 8-2b, it is clear that rotation of the electrode does not eliminate peaks c3 and c4. The origin of peaks c3 and c4 is most likely reduction of chalcopyrite itself [Li et al. 2011]. From the Pourbaix diagram for chalcopyrite [Garrels and Christ 1965], at ca. pH 0.3 chalcopyrite is not stable at potentials below \u221250 mV, i.e. the Pourbaix diagram predicts formation of chalcocite at this condition. Furthermore, again from the same Pourbaix diagram, chalcocite (Cu2S) is the stable phase at potentials below \u22120.25 V, and may thus form on the surface. However, partial cathodic dissolution of iron from the chalcopyrite structure may result in the formation of copper-iron-sulfur phases such as Cu9Fe8S16 or Cu5FeS4. A list of possible reactions associated with peaks c3 and c4 of Fig. 8-6 are presented in Table 8-1 (reactions 8-5 to 8-7).  156  \u000CFigure 8-6: Multicyclic voltammetry of chalcopyrite electrode presenting the first, second, and tenth cycles. Scan rate is 50 mV s\u22121.  8.2.3 Anodic reactions on the surface of chalcopyrite Fig. 8-1a reveals seven distinct anodic peaks for a typical stationary chalcopyrite electrode. Anodic peaks a1 to a7 are centered at 0.02, 0.10, 0.19, 0.24, 0.35, 0.65 and 0.79 V, respectively. In order to understand the origin of these peaks (the assigned reactions to all anodic peaks are summarized in Table 8-1), the anodic vertex potential was held at 0.75 V (same as Fig. 8-2a) while the cathodic vertex potential was increased from \u22120.5 V to +0.1 V (Fig. 8-7a) and \u22120.1 V (Fig. 8-7b). In both experiments, first the electrode was held at 0.55 V (ca. OCP) for 20 seconds and afterwards polarized linearly up to 0.75 V (50 mV s\u22121), and then 10 cycles were applied (50 mV s\u22121). From Fig. 8-7a, it is clear that anodic peaks a1, a2, a3 and a4 are related to the oxidation of products of cathodic peaks c3 and c4 (cf. Table 8-1); otherwise, they would have appeared in Fig. 8-7a. Moreover, from Fig. 8-2b one can say that the reactions associated with these anodic peaks all involve soluble reactants, such that rotation of the electrode eliminates them. By lowering the cathodic vertex potential from 0.1 V (Fig. 8-7a) to \u22120.1 V (Fig. 8-7b), cathodic peak c3 (and, to some extent, c4 \u2013 cf. Table 8-1) occurs, and subsequently, anodic peaks a2, a3 and a4 also appear. Hence, these anodic peaks are related to the oxidation of soluble reaction products associated with peaks c3 and c4 (ferrous and H2S, based on the reactions given in Table 8-1). Formation of copper sulfides 157  \u000C(with different stoichiometries) through oxidation of other copper sulfides (products of reactions 8\u20135 to 8\u20137 of Table 8-1) are possible reactions [Biegler and Horne 1985; Vinkevicius et al. 1998]. An example of these reactions can be reaction 8-9, presented in Table 8-2. The required H2S for the above reaction originates from the reactions associated with peaks c3 and c4 of Fig. 8-2a, i.e. reactions 8\u20135 to 8\u20137 of Table 8-1. By rotation of the electrode (Fig. 8-2b), any aqueous H2S product diffuses away from the electrode surface and reaction 8-9 (cf. Table 8-2) is prevented.  Figure 8-7: Multicyclic voltammetry (10 cycles) of a stationary chalcopyrite electrode in the potential window of (a) 0.1 to 0.75 V, and (b) \u22120.1 to 0.75 V. Scan rate is 50 mV s\u22121.  158  \u000CTable 8-2: Anodic reactions on chalcopyrite surface.  159  peak  Reaction number  Reaction  Reference  a1 \u2013 a4  8-8  Cu 2S \uF02B H 2S(aq) \uF0AE 2 CuS \uF02B 2 H \uF02B \uF02B 2 e \uF02D  Biegler and Horne 1995  a5  8-4  CuFeS 2 \uF02B 3 Cu 2 \uF02B \uF02B 4 e \uF02D \uF0DC 2 Cu 2 S \uF02B Fe 2 \uF02B  Biegler and Horne 1995  a6  8-9  a7  8-10  Cu 1\uF02D x Fe1\uF02D y S 2 \uF0AE \uF0281 \uF02D x \uF029 Cu 2\uF02B \uF02B \uF0281 \uF02D y \uF029 Fe 3\uF02B \uF02B 2 S \uF02B \uF0285 \uF02D 2 x \uF02D 3 y \uF029 e \uF02D  Hackl et al. 1995  a8  8-12 8-13  CuFeS 2 \uF02B 8 H 2 O \uF0AE Cu 2 \uF02B \uF02B Fe 3 \uF02B \uF02B 2 SO 24 \uF02D \uF02B 16 H \uF02B \uF02B 17 e \uF02D  Yin et al. 1995  CuFeS 2 \uF0AE Cu 2\uF02B \uF02B Fe 3\uF02B \uF02B 2 S \uF02B 5e \uF02D  Yin et al. 1995  CuFeS 2 \uF03D Cu 1\uF02D x Fe1- y S 2 \uF02B xCu 2 \uF02B \uF02B yFe 2\uF02B \uF02B 2( x \uF02B y )e \uF02D  y\u00BBx  Nava and Gonzalez 2006  \u000CFig. 8-8 presents the multicyclic voltammogram of a stationary chalcopyrite electrode over a wide potential range (\u20130.5 to 1.2 V). In this figure, anodic peaks a1 to a4 are the same peaks as those shown in Figs. 8-6 and 8-7. Peak a5 is the reverse of reaction 8-4 (cf. Tables 8-1 and 8-2) as shown in Figs. 8- 3 to 8-5. Peak a6 of Fig. 8-8 represents the anodic oxidation of chalcopyrite. At potentials below 0.9 V (transpassive dissolution potential of chalcopyrite [Ghahremaninezhad et al. 2010]), iron atoms dissolve preferentially from the surface of chalcopyrite, leaving behind a copper-rich film [Ghahremaninezhad et al. 2010; Viramontes-Gamboa et al., 2007]. In fact, it has been claimed by many researchers that a metal-deficient sulfide film is the passive film responsible for slow leaching of chalcopyrite [Ghahremaninezhad et al. 2010; ViramontesGamboa et al. 2007; Biegler and Swift 1979; Yin et al. 1995; Nava and Gonzalez 2006; Hackl et al. 1995]. Hence, the reaction associated with peak a6 of Fig. 8-8 is [Ghahremaninezhad et al. 2010]: CuFeS 2 \uF03D Cu 1\uF02D x Fe1\uF02D y S 2 \uF02B xCu 2\uF02B \uF02B yFe 2\uF02B \uF02B 2( x \uF02B y )e \uF02D  y&gt;x  8-10  The formation of a metal-deficient sulfide film can be investigated by OCP measurements. If a passivated chalcopyrite electrode is immersed in the electrolyte (at OCP), then the metal-deficient sulfide film will relax as iron and copper atoms diffuse from the bulk of the electrode to the surface (i.e. diffusion will cause the x and y values of reaction 8-9 to approach 0). In order to confirm the formation of such a film, chalcopyrite electrodes were polarized at three different passive potentials of 0.65, 0.75 and 0.85 V for 10 hrs in 0.5 M sulfuric acid solution and afterwards allowed to relax to their OCPs. Fig. 8-9 presents the OCP of each electrode after polarization. In addition to the above three OCP curves, for the sake of comparison, the OCP behavior of an unpolarized chalcopyrite electrode is also presented in Fig. 8-9 (solid black curve). The OCP measurements of Fig. 8-9 were conducted for about 9 hrs. All of the OCP values approach the original OCP value of chalcopyrite (that of an unpolarized sample \u2013 ca. 0.54 V) after about 4 hrs. This figure provides further evidence that the passive layer of chalcopyrite is a metal-deficient film. Hence reaction 8-9 (Table 8-2) is the reaction associated with the anodic dissolution of chalcopyrite (peak a6 of Fig. 8-8).  160  \u000CFigure 8-8: Multicyclic voltammetry of a stationary chalcopyrite electrode in the potential window of \u22120.5 to 1.2 V. Scan rate is 50 mV s\u22121.  Figure 8-9: OCP measurements of a chalcopyrite electrode after polarizing at different passive potentials of 0.65, 0.75 and 0.85 V for 10 hrs. The OCP behavior of a polished chalcopyrite electrode (without previous polarization) is also presented with a solid black line.  161  \u000CIn addition to peaks a1 to a6, Fig. 8-8 presents a new anodic peak observable at 0.90 V (peak a7) and a rapid rise in current in the positive-going direction at potentials above 1.1 V (peak a8). The magnitude of peaks a5 and a6 increased with the number of cycles while the magnitude of peak a7 was almost constant. Interestingly, the potential of peak a7, 0.90 V, is the same as the transpassive dissolution of chalcopyrite [Ghahremaninezhad et al., 2010]. Hence, peak a7 can be assigned to dissolution (oxidation) of the passivating film of chalcopyrite (i.e., the metal-deficient sulfide film). The near-constant magnitude of peak a7 can be justified by the fact that the passive film is only a few nanometers thick [Ghahremaninezhad et al., 2010]. Hence, the charge associated with dissolution of this film (and consequently the peak current magnitude) is always about the same. Elemental sulfur (and polysulfide) is one of the oxidation products of chalcopyrite at (transpassive) potentials above 0.90 V [Ghahremaninezhad et al., 2010]. If peak a7 is in fact associated with the oxidation of the passive film, and if elemental sulfur is one of the oxidation products, then one possible reaction for peak a7 would be: Cu 1\uF02D x Fe 1\uF02D y S 2 \uF0AE \uF0281 \uF02D x \uF029 Cu 2\uF02B \uF02B \uF0281 \uF02D y \uF029 Fe 3\uF02B \uF02B 2 S \uF02B \uF0285 \uF02D 2 x \uF02D 3 y \uF029 e \uF02D  8-11  where y &gt; x [Hackl et al. 1995]. In reaction 8-11 we have considered the formation of ferric due to the thermodynamic instability of ferrous at 0.90 V and above. In order to calculate the thickness of the passive film, we make the assumption that all of the charge associated with peak a7 is related to CuS dissolution (since y in reaction 8-11 is much larger than x, so one can assume that the iron content of the metal-deficient phase is almost negligible) [Warren et al. 1982]. In other words, we assume that the main passivating film of chalcopyrite is CuS. This assumption is also consistent with the results of some other workers who have concluded that the copper sulfides formed at passive potentials are responsible for passivation of chalcopyrite [Biegler and Horne 1985; Yin et al. 1995; Lazaro and Nicol 2006; Nava and Gonzalez 2006; Holliday and Richmond 1990]. Under this assumption, it is possible to relate the mass of the dissolved copper sulfide (CuS) through peak a7 to the charge associated with the peak. Integrating the current under peak a7 in Fig. 8-8 gives a charge of 6.2 \u00B1 0.4 mC cm\u22122. Considering the calculated density of 4.6 g cm\u22123 for CuS [Evans and Konnert 1976], and assuming a surface roughness factor of unity, the thickness of the passivating CuS film on the surface of chalcopyrite can be estimated to be approximately 6.7 nm. This result agrees fairly well with that of Yin et al. [1995] who calculated a thickness 162  \u000Cof about 4 nm for the passive film. Elsewhere, a passive film consisting of CuS and elemental sulfur with a thickness of 2.9 nm is proposed [Biegler and Horne 1985]. From Fig. 8 it is very clear that above 1.1 V the anodic current density sharply increases. This is due to the dissolution of chalcopyrite via one or both of the following reactions [Yin et al. 1995]: CuFeS 2 \uF02B 8 H 2 O \uF0AE Cu 2 \uF02B \uF02B Fe 3\uF02B \uF02B 2 SO 24\uF02D \uF02B 16 H \uF02B \uF02B 17 e \uF02D  8-12  CuFeS 2 \uF0AE Cu 2\uF02B \uF02B Fe 3\uF02B \uF02B 2 S \uF02B 5e \uF02D  8-13  The cathodic peak of Fig. 8-8 at \u22120.10 V (peak c*3, whose potential increases with cycle number up to \u22120.18 V) has a different shape and position compared to cathodic peaks c3 and c4 of Fig. 8-2a. Peak positions for c3 and c4 of Fig. 8-2a in the tenth cycle are \u22120.08 and \u22120.20 V respectively, while the potential of peak c*3 in Fig. 8-8 is \u22120.17 V. This dissimilarity may suggest a change in the path of the reactions associated with these peaks. Further evidence for such a change can be obtained by comparing the magnitude of peaks c3 and c4 in Fig. 8-2a with that in Fig. 8-8. In the first cycle shown in Fig. 8-2a peaks c3 and c4 are absent while in Fig. 8-8 the cathodic peak at \u22120.10 V has an appreciable magnitude during the first cycle, ca. 18 mA, and grows to ca. 32 mA after several cycles. In order to determine the origin of this difference, multicyclic voltammetry in the potential window of \u22120.5 to 1.0 V was conducted, as shown in Fig. 8-10. The lower vertex potential of the experiment was kept at \u22120.5 V but, in order to eliminate the oxidation reaction which begins at 1.1 V, the upper vertex was reduced to 1.0 V. In both of the experiments of Figs. 8-8 and 8-10, the electrodes were held at 0.55 V (ca. OCP) for 20 seconds, then the potential was scanned up to the anodic vertex (1.2 V for Fig. 8-8 and 1.0 V for Fig. 8-10), and then 10 cycles were performed. The shape of cathodic peaks c1 and c2 in Fig. 8-10 are the same as those of Fig. 8-2a and different from that of Fig. 8-8 (peak c*3). By comparing Figs. 8-8 and 8-10, one may conclude that the origin of the different magnitudes of peaks c3 and c4 of the first cycle is related to the reaction associated with peak a8 of Fig. 8-8. If peak a8 is associated with dissolution of chalcopyrite via reactions 8-12 and 8-13 (cf. Table 8-2), then at potentials above 1.1 V the surface of the electrode dissolves actively and the exposed surface of the electrode will consist mainly of chalcopyrite and elemental sulfur [1, 22], and all of the other copper sulfide intermediates will have been dissolved. Hence, we conclude that peak c*3 of Fig. 8-8 is related to the reduction of elemental sulfur. By exposure of the 163  \u000Celectrode to the very high anodic potential of 1.2 V (Fig. 8-8), a significant amount of elemental sulfur can form on the surface of chalcopyrite (via reaction 8-13). In fact, exposure of the electrode to this potential causes the formation of a thin black film on the surface of the electrode which can be wiped away with a soft cotton swab. Thus, peak c*3 of Fig. 8-8 can be related to the following reaction: S \uF02B 2 H \uF02B \uF02B 2 e \uF02D \uF03D H 2S (aq)  E ( VSHE ) \uF03D 0.148 \uF02D 0.0591 pH  8-14  Due to the high concentration of sulfur on the surface, elemental sulfur reduction will be the main cathodic reaction in comparison to cathodic reactions 8\u20135 to 8\u20137 of Table 8-1.  Figure 8-10: Multicyclic voltammetry of a stationary chalcopyrite electrode in the potential window of \u22120.5 to 1.0 V. Scan rate is 50 mV s\u22121.  8.3 Conclusion In summary, electrochemical reactions on the surface of a massive chalcopyrite electrode were investigated in 0.5 M sulfuric acid solution (pH 0.3) at 25\u00BAC by multicyclic voltammetry and open circuit potential measurements. Variation of the scan limits under different experimental conditions (quiescent vs. rotating electrodes) was used for analysis of each electrochemical current peak.  Oxidation of chalcopyrite proceeds at high anodic  potentials (E &gt; 0.6 VSHE). It was shown iron atoms dissolve preferentially from the surface of chalcopyrite during anodic dissolution, leaving behind a copper-rich metal-deficient 164  \u000Csulfide passive film (i.e. Cu1\u2212xFe1\u2212yS2 y &gt; x). By integrating the current under the chalcopyrite de-passivation peak (peak a7 of Fig. 8-10), the thickness of the passive film was calculated to be ca. 6.7 nm. Additionally, experiments showed that elemental sulfur is a product of chalcopyrite oxidation at anodic potentials higher than 1.0 VSHE. Cathodic dissolution of chalcopyrite proceeds at potentials below 0 V SHE. The cathodic dissolution process involves the release of ferrous and H2S(aq) to solution. The cathodic reactions which occur in the potential region of 0.1 to 0.6 VSHE and the anodic reactions which occur at potentials lower that 0.5 VSHE are all related to the reduction and oxidation of soluble species. The soluble species involved in those cathodic reactions are ferric and cupric and the soluble species involved in the anodic reactions is H2S(aq).  165  \u000C9 Ferric/ferrous redox reaction on anodically passivated chalcopyrite (CuFeS2) electrode3 9.1 Introduction Typical hydrometallurgical copper extraction processes for chalcopyrite concentrates proceed by reduction of an oxidant (e.g. ferric ion, oxygen) and oxidation of chalcopyrite to cupric and ferrous. The oxidation reaction for chalcopyrite in the presence of ferric can be written as [Dutrizac and MacDonald, 1974]: CuFeS 2 \uF02B 4Fe 3\uF02B \uF03D Cu 2 \uF02B \uF02B 5Fe 2 \uF02B \uF02B 2S o  9-1  Oxygen can re-oxidize the ferrous back to ferric via the following reaction: 4Fe 2\uF02B \uF02B O 2 \uF028aq \uF029 \uF02B 4H \uF02B \uF03D 4Fe 3\uF02B \uF02B 2H 2 O  \uF044G25o o C \uF03D \uF02D193.3 kJ/mole  9-2  The ferric ion represents one of the most important oxidants in hydrometallurgy [Tshilombo et al., 2002]. The ultimate effects of ferric in leaching solutions are to accept electrons at the mineral surface with rapid kinetics and to maintain a relatively high redox potential, at which the electrochemical dissolution of chalcopyrite (generally the mineral of interest) is thermodynamically favored. From the Pourbaix diagram of chalcopyrite [Garrels and Christ 1965], the acidic oxidation of chalcopyrite is thermodynamically possible at pH values lower than 6 and oxidation potentials higher than 0.2 VSHE. Since the ferric-ferrous couple has a relatively high potential, Eo = 0.771 VSHE, the desired solution potential during leaching can be obtained by the addition of ferric. The other reason for the use of ferric as an oxidant is its relatively rapid reduction kinetics on the surfaces of some minerals such as pyrite. In the case of chalcopyrite leaching in sulfuric acid solution, it has been claimed that the kinetics of the ferric-ferrous redox reaction are very slow [Beigler and Horne 1985; Parker et al., 1981]. Elsewhere, McMillan et al. [1982] have shown that the rate of the ferricferrous redox reaction is very slow in comparison with Cu2+/Cu+ and I 3\uF02D /I \uF02D redox couples. Nevertheless, the kinetic parameters of the ferric-ferrous couple on the surface of chalcopyrite under leaching conditions have not yet been quantified. These kinetic parameters are important for hydrometallurgical process optimization and development. The 3  A version of this chapter has been published: \u201CGhahremaninezhad, D.G. Dixon, A. Asselin, 2012, Hydrometallurgy, 125\u2013126, 42\u201349.\u201D  166  \u000Clack of quantitative information on the kinetic parameters of the ferric-ferrous couple on the surface of chalcopyrite (and almost all other minerals) is due to the interference from chalcopyrite oxidation. In other words, during the dissolution of chalcopyrite, the oxidation reaction of chalcopyrite will be coupled with ferrous oxidation; and in the case of reducing conditions, oxygen and chalcopyrite reduction will be coupled with ferric reduction. The interference from unwanted reactions makes the quantification of ferric reduction kinetics by application of typical electrochemical methods almost impossible. Another issue with quantifying the kinetics of redox couples is passivation of the substrate. The kinetics of the ferric-ferrous couple on the surface of oxide-free metals [Samec and Weber 1977] and superficially oxidized metals [Makrides 1964a] have been studied previously. The determination of these kinetics on the surface of oxide-free metals is straightforward. On the other hand, the charge transfer rate of the redox reactions at superficially oxidized metals depends on the thickness of the semiconductor oxide film. This can be described with the quantum mechanical theories of electron transfer [Marcus 1964]. According to these theories, the tunneling probability decreases with the thickness of the passive film. However, in the case of semiconductor electrodes, e.g. chalcopyrite and most of the sulfide minerals, the electrochemical reactions proceed by participation in the electron transfer reaction of electronic states inside the semiconductor electrode. Hence, chalcopyrite itself, as a semiconductor material, and its passive film together make the kinetics of the ferric-ferrous couple more complicated (even than that of superficially oxidized metals). Under anodic potentials of about 0.65 to 0.90 VSHE in sulfuric acid solution, chalcopyrite is in the passive state and several metal-deficient-sulfide passive films, e.g. Cu1\u2212xFe1\u2212yS2 (y &gt; x), Cu1\u2212zS2 and CuS, have been suggested to form at potentials between 0 and 1.0 VSHE [Ghahremaninezhad et al., 2010a; Cordoba et al., 2008; Viramontes-Gamboa et al., 2007; Nava and Gonzalez 2006]. In addition, it has been shown that the passive film of chalcopyrite has semiconductive properties [Ghahremaninezhad et al., 2010a; Hiskey 1993]. Hence, in order to analyze the kinetics of the ferric-ferrous couple on chalcopyrite under leaching conditions, one must include the influence of the semiconducting properties of the passive film in the redox reaction kinetics. In this chapter, the kinetics of the ferric-ferrous couple on the surface of chalcopyrite in acidic sulfate media will be discussed. Despite attempts to measure the kinetics on 167  \u000Cchalcopyrite surfaces [Cordoba et al., 2008; Hiroyoshi et al., 2001], to the best of our knowledge, there is no report of the exchange current density values or reaction orders for the ferric-ferrous couple on the surface of chalcopyrite.  9.2 Experimental A high quality natural polycrystalline chalcopyrite sample (98.1% chalcopyrite, 1.9% siderite FeCO3) originating from the Creighton Mine (Sudbury, Ontario) was used for electrode fabrication. Further characterization of the mineral and morphological distribution of the siderite are given elsewhere in chapter 6 [and also Ghahremaninezhad et al., 2010a]. Massive electrodes were prepared by cutting the mineral into approximately cubic shapes with areas of 1 cm2 exposed to the solution. In order to have a uniform current distribution on the samples, a thin layer of metallic copper was electrodeposited on the back of the electrodes (the side to which the wire was attached). Copper was electrodeposited from a copper sulfate solution at low current density (4 mA cm\u22122). An electrical contact with copper wires was made to the back of the samples with silver-loaded epoxy resin (Mg Chemicals). The electrode contacts were protected from the solution by mounting the whole assembly in epoxy resin (LECO), leaving only one face of the electrode exposed to the solution. Prior to each electrochemical experiment the electrodes were checked to ensure that there was no evidence of crevice corrosion in the electrode. This was possible due to the transparent epoxy resin mounts. Before each test, the exposed face of the electrode was polished with new carbide papers, starting from no. 300 and progressing to no. 1200 paper. After this polishing procedure, the surfaces of the electrodes were inspected with an optical microscope to ensure that the morphology of the electrodes was similar. The electrodes were then washed, sonicated and rinsed in deionized (DI) water and dried in a stream of air. All electrochemical experiments were performed at 25\u00BAC using a VersaSTAT 4 potentiostat/galvanostat controlled by VersaStudio software (version 1.51, Princeton Applied Research). A conventional three-electrode electrolytic cell was used for electrochemical analysis. The test cell was jacketed and water was circulated through the jacket to maintain the desired test temperature of 25\u00BAC. The working electrode was always the prepared chalcopyrite electrode. A graphite counter electrode (99.999%, Alfa Aesar) and a glass-body Ag/AgCl reference electrode filled with 4 M KCl electrolyte saturated with AgCl (Accumet 168  \u000C13-620-523, 199 mV vs SHE at 25\u00BAC) were used. All electrode potentials in this study are reported with respect to the SHE. The reference electrode was connected to a saturated K2SO4 solution-filled Luggin capillary to minimize IR drop. The electrolyte inside the reference electrodes was refreshed after every 20 h of use and the electrodes were tested against a lab-standard Ag/AgCl electrode before each experiment. The electrode potential was allowed to stabilize for 6 h before starting the measurements and, in order to ensure reproducibility, tests were begun only after reaching the same OCP to within \u00B15 mV. The initial solution was 500 mL of 0.5 M H2SO4 (pH 0.3), prepared by dissolving reagent grade sulfuric acid (Fisher Scientific) in DI water. In all experiments the oxygen content of the solution was controlled by sparging air through the solution. The potentiostatic polarization experiments were performed by application of a constant potential and the current flow through the sample was registered. The experiments for the kinetic measurements of ferric-ferrous redox reactions on the chalcopyrite surface were performed in the following order. Before each experiment, the electrode potential was allowed to stabilize for 6 h in 500 mL of 0.5 M H2SO4 solution. The solution was stirred during the experiments (250 rpm). The electrode was then passivated potentiostatically at the desired potential for 6 h (the passive potential region of chalcopyrite was obtained from potentiostatic polarization data). The current density during the first two hours of passivation dropped to a steady-state current of less than 10 \u03BCA cm\u22122. Passivation for more than 24 h showed negligible effect on the current, meaning the thickness of the passive film was almost constant after 6 h. 200 mL of solution containing 0.5 M H2SO4 and the desired amounts of Fe2(SO4)3 and FeSO4 (both from Fisher Scientific) were then added to the solution, in order to change the composition of the original electrolyte to 0.5 M H2SO4, 0.1 M Fe2(SO4)3 and 0.05 M FeSO4. The applied potential was held constant during the entire process and the current was registered for 20 h after addition of the Fe salts. The current jump in the potentiostatic experiment was related to the ferric-ferrous redox reaction. EIS tests were conducted at all potentials one hour prior to and one hour after the addition of Fe salts. The frequency range was 10\u22122 to 105 Hz with a peak-to-peak amplitude of 5 mV.  ZSimpWin 3.20 software (Princeton Applied Research) was used to fit the  impedance data.  169  \u000C9.3 Results and discussion Fig. 9-1 shows a typical potentiostatic polarization curve for the chalcopyrite electrode in 0.5 M H2SO4 solution. The curve was obtained by potentiostatically polarizing the chalcopyrite electrode for 6 h at each potential in order to reach the steady-state current (Iss). The i-t (current density-time) behavior of the chalcopyrite electrode at selected potentials is presented in Fig. 9-2.  Figure 9-1: Potentiostatic polarization of chalcopyrite in 0.5 M sulfuric acid solution.  170  \u000CFigure 9-2: i-t behavior of chalcopyrite electrode in 0.5 M sulfuric acid solution at between 0.45 to 1.00 V.  From Fig. 9-1 and Figs. 9-2a to c, all potentials between 0.50 to 0.90 V can be considered to lie within the passive region. It is clearly stated in the literature that even at potentials below the open circuit potential (OCP) of chalcopyrite (ca. 0.54 V from Fig. 9-1) the surface of chalcopyrite is covered by a passive film which forms during exposure to air (prior to exposing the mineral to sulfuric acid solution) or during the 6 h stabilization of the electrode in solution prior to polarization [Ghahremaninezhad et al., 2010a; Velasquez et al., 2005]. The passive film on the surface of chalcopyrite is a stable film which causes steady state dissolution of chalcopyrite. This can be confirmed by long term potentiostatic polarization of chalcopyrite at a passive potential, as shown in Figs. 9-2a to c. Indeed, formation of a stable passive film on the surface of chalcopyrite is the reason for the sharp decrease in current prior to reaching the steady state dissolution current density in the potential region of 0.60 to 0.90 V. All of the curves in Figs. 9-2a and 9-2b show a sharp decrease in the current density during the first few hours of the experiment. However, in Fig. 9-2c, only a 50 mV increase in the potential of the electrode (from 0.90 to 0.95 V) causes 171  \u000Cbreakdown of the passive film and active dissolution of chalcopyrite. At the high anodic potential of 0.95 V, the dissolution rate increases as a function of time and, even after 6 h of polarization, the current density has not yet stabilized. Dissolution of chalcopyrite at 1.0 V (Fig. 9-2d) also shows a very high dissolution current density. The dissolution current density at this potential is above 3 mA cm\u22122 and the behavior is very different from the lower potentials. In Fig. 9-2d, the current reaches a peak after only a few minutes due to double layer formation. The decrease in current density observed after 1.5 h at 1.0 V (Fig. 9-2d) is probably due to the precipitation of Fe species on the surface of the electrode. Lu et al., [2001] have suggested that at low pH (&lt; 1, as in our experiments), ferric containing species such as jarosite may precipitate on the surface of chalcopyrite. Nava and Gonzalez [2006] and Klauber [2008] have also suggested the formation of ferric sulfate precipitates in this potential region. Nevertheless, even after 10 h of electrodissolution at 1.0 V (Fig. 9-2d) the current density is very high (ca. 1.5 mA cm\u22122) which suggests that, regardless of a decrease in the current density, none of the layers formed on the surface act as passive films [Ghahremaninezhad et al., 2010a]. Further analysis of the i-t behavior of chalcopyrite electrodes at 1.0 V is not relevant to this study. The potentiostatic polarization diagram of chalcopyrite (Fig. 9-1) shows a negative slope in the anodic passive region (0.65 to 0.85 V). This is presented in Fig. 9-3, which shows the anodic section of the potentiostatic polarization of Fig. 9-1. According to the point defect model of passive films, the negative slope of the anodic passive region is related to the presence of an n-type semiconductor passive film with cation interstitial defects or anionic vacancies [Macdonald et al., 1998]. This is in agreement with our previous findings [Ghahremaninezhad et al., 2010a] which have shown that the passive film formed on the same chalcopyrite sample used in this study in the same solution has characteristics of an ntype semiconductor material. Fig. 9-4a shows the points of the measured differences between the current density of the potentiostatic experiments (0.56 \u2264 E \u2264 0.95 V) before and after addition of Fe salts to the electrolyte. This process is illustrated in Fig. 9-4b, which shows the current-time behavior of the electrode after addition of Fe salts at 0.70 V. The \u0394i value, shown in the inset of Fig. 9-4b, is the current density used to plot Fig. 9-4a.  172  \u000CFigure 9-3: Potentiostatic polarization of chalcopyrite in 0.5 M sulfuric acid solution. This figure magnifies the anodic region as shown in Fig. 9-1.  In order to ensure that the data points of Fig. 9-4a are reliably presenting the Tafel region for the ferric-ferrous couple, two conditions must be met. First, the passive film must be stable on the surface of the chalcopyrite. The electrode was potentiostatically polarized at a constant potential (0.56 \u2264 E \u2264 0.95 V) for 6 h and for all of the potentials the current\u2013time curve had the characteristics of the passive state, i.e. decline of current until nearly constant, e.g. Fig 9-4b. Second, excessive overpotential must be avoided. High overpotential can change the charge-transfer-controlled kinetics to diffusion-controlled kinetics. An estimate of the limiting current density for the ferric-ferrous couple was calculated. The solution used in these experiments contains 0.2 M Fe3+ and 0.05 M Fe2+ and, considering the concentration of ferrous, the limiting current density is about 10 mA cm\u22122, from:  idFe  3\uF02B  /Fe 2\uF02B  \uF03D  FDc \uF064  9-3  where D is the diffusion coefficient and \u03B4 is the thickness of the Nernst layer. In the above calculation we have assumed that D = 10\u22125 cm2 s\u22121 and \u03B4 = 5\u00D710\u22123 cm [Makrides 1964b]. In this study, the registered highest current density, before addition of the iron salts, was less than 10\u22125 A cm\u22122 (Fig. 9-3), and the current density for the ferric-ferrous couple was always less than 3\u00D710\u22124 A cm\u22122 (Fig. 9-4a). In addition to the above argument, the data points of Fig. 9-4a show two reasonably straight lines corresponding to the Tafel region. This means that any possible contribution from diffusion current can be neglected and that the major 173  \u000Ccontribution comes from charge transfer. Hence, it is safe to conclude that Fig. 9-4a represents the Tafel behavior of the ferric-ferrous redox couple on the surface of the passivated chalcopyrite electrode.  Figure 9-4: (a) Tafel region of the Fe3+/Fe2+ couple, and (b) potentiostatic behavior of a chalcopyrite electrode at 0.70 V before and after addition of the Fe3+/Fe2+ couple (the difference in current was used to derive Fig. 9-4a).  A typical plot of charge density vs time for potentiostatic polarization at 0.60 V is presented in Fig. 9-5. The curve for the Fe-containing part of the experiment (solid-red line) presents a constant slope which is due to the ferric reduction reaction. The inset of Fig. 9-5 presents an expanded view of the first 6 h of polarization in Fe-free solution (dashed green line). After two hours the slope of the curve is almost constant (inset of Fig. 9-5). However, 174  \u000Cthe reaction rate is higher in the Fe-containing segment due to the rapid ferric reduction reaction.  Figure 9-5: (a) A typical plot of charge density vs time for potentiostatic polarization at 0.60 V. The solidred line represents the stage at which Fe-containing solution was added to the original solution. The inset figure shows an expanded view of the first 6 hours of polarization in Fe-free solution (dashed-green line).  From Fig. 9-4a, the equilibrium potential of the ferric-ferrous couple is calculated to be 0.77 V. From the Nernst equation, this value should be: \uF0E6 Fe 3\uF02B \uF0F6 E \uF028VSHE \uF029 \uF03D 0.770 \uF02B 0.059 log\uF0E7\uF0E7 2 \uF02B \uF0F7\uF0F7 \uF03D 0.805 \uF0E8 Fe \uF0F8  9-4  which is higher than the measured potential of this study. This is justified for two reasons: (1) the formation of species such as Fe(OH)2+ and possibly Fe(OH)+ [Makrides 1964a] and/or other ferric sulfate complexes [Hirato et al., 1987] can reduce the free ferric and ferrous concentrations and cause the reversible potential of the ferric-ferrous couple to be different from the calculated one; (2) experimental measurements may contain an appreciable liquid junction potential which is, however, constant during the experiment. The exchange current density, io, and the anodic and cathodic transfer coefficients, \u03B1a and \u03B1c, were calculated to be 2.44 \u03BCA cm\u22122, 0.37 and 0.50, respectively. Reported io values for superficially oxidized Fe, Ti, Fe-13.2% Cr, and 304 stainless steel electrodes are 10, 26, 1, and 0.045 \u03BCA cm\u22122, respectively [Makrides 1964a, 1964b; Stern 1957; Ghahremaninezhad et 175  \u000Cal., 2010b], which are close to the above results. The reaction order obtained from the summation of the anodic and cathodic transfer coefficients is 0.87. Theoretically, for the ferric-ferrous couple on the surface of bare metal this number should be 1.00. However, blocking of the anodic reactions is a typical property of n-type semiconductors [Kapusta and Hackerman 1981], which lowers the anodic transfer coefficient, \u03B1a, to less than 0.4 while the cathodic transfer coefficients, \u03B1c, remain around 0.5. In fact, this result matches with that of our previous work [Ghahremaninezhad et al., 2010a] which suggested that at high anodic potentials (e.g. 0.80 V or above) the surface of the electrode is negatively charged due to the accumulation of metal vacancies in the passive film/solution interface. This negative charge on the surface is in fact responsible for blockage of oxidation reactions at the electrode surface. Additionally, it should be noted that the formation of n-type passive films on the surface of chalcopyrite was suggested by the potentiostatic experiments (Fig. 9-3) presented above and the Mott-Schottky experiments discussed in a previous paper [Ghahremaninezhad et al., 2010a]. Electrochemical impedance spectroscopy (EIS) has also been used to further investigate the ferric-ferrous couple on a passivated chalcopyrite electrode. The EIS method provides a kinetic model in the form of an equivalent circuit, and one can typically attribute a clear physical meaning to each of the parameters (circuit elements) of the model. EIS is a powerful technique for the study of charge transfer kinetics, and is used here to study the charge transfer rate on the surface of chalcopyrite in the presence and absence of iron salts in solution. Fig. 9-6 shows the measured and calculated Nyquist impedance spectra for iron-free (Fig. 9-6a) and iron-containing (Fig. 9-6b) solutions, measured at 0.70 V. In both figures, two time constants are clearly evident. The first time constant is a well-defined capacitive loop related to the combination of the double-layer capacitive impedance with a charge transfer resistance. The second time constant is one part of a capacitive loop, which is related to the influence of diffusion processes occurring at the electrode surface [MacDonald 1990; Ghahremaninezhad et al., 2010a]. During the application of an anodic potential of 0.70 V to the chalcopyrite electrode in iron-free solution (Fig. 9-6a), iron atoms dissolve preferentially from the surface of chalcopyrite, leaving behind a copper-rich passive film [Hackl et al., 1995; Ghahremaninezhad et al., 2010a]. Hence, the second time constant may be due to the diffusion of iron atoms through the copper rich passive film. Alternatively, the second time 176  \u000Cconstant may be related to ferric reduction on the electrode surface. At the applied potential of 0.70 V, ferric reduction contributes significantly to the overall electrochemical processes occurring on the chalcopyrite electrode. This is clearly demonstrated in Fig. 9-4b, where the addition of iron salts to the solution caused the current density to increase from ca. 3 \u03BCA cm\u22122 (anodic) to ca. \u201317 \u03BCA cm\u22122 (cathodic). In addition, the second time constant of the Nyquist plot of Fig. 9-6b is a straight line with a slope of ca. 1 (characteristics of Warburg impedance), which suggests a diffusion process in solution [Bard and Faulkner 2001]. Hence, the second time constant in iron-containing solution (Fig. 9-6b) may be due to the diffusion of ferric towards the electrode surface.  Figure 9-6: Measured and calculated Nyquist plots for the chalcopyrite electrode at 0.70 V in (a) Fe3+/Fe2+-free, (b) Fe3+/Fe2+-containing solution.  CPE Rdl Rsol  Q RCPE  CP  Cdl RP Figure 9-7: Equivalent electrochemical circuit for the chalcopyrite/electrolyte interface.  177  \u000CThe equivalent electrical circuit chosen to model the measured impedance spectra is shown in Fig. 9-7, and its electrical parameters are given in Table 1. The low chi-squared (\u03C72) values given in Table 9-1 indicate a good fit to the experimental data. Rsol represents the solution resistance and, not surprisingly, this resistance is decreased by the addition of iron salts. Rdl and Cdl represent the resistance and capacitance of the double-layer, respectively. Note the significant decrease in the double-layer resistance (Rdl) upon the addition of iron salts. The double-layer resistance can be defined as follows [Bard and Faulkner 2001]:  Rdl \uF03D  RT nFio  9-5  where R, T , n and F represent the ideal gas constant, temperature in Kelvin, number of transferred electrons and Faraday constant, respectively, and io is the exchange current density of the electrochemical reaction. In iron-free solution, the slow anodic dissolution of chalcopyrite is the electrochemical reaction on the chalcopyrite electrode and its exchange current density value must be applied in Eq. 9-5. In contrast, by addition of iron salts to the solution, ferric reduction will be the predominant reaction on the surface and in this case the exchange current density of ferric reduction should be used in Eq. 9-5. The exchange current density of ferric reduction is higher than that of chalcopyrite oxidation. Hence, one would expect a decrease in the Rdl value by addition of iron salts to the solution. This agrees well with the data of Table 9-1, which shows the Rdl value decreases from 1714 to 470 \u03A9 cm2 upon the addition of iron salts to the solution.  178  \u000CTable 9-1: Model parameters for equivalent circuit of Fig. 9-7.  Solution  RSol (\u2126 cm2)  Rdl (\u2126 cm2)  Fe3+/Fe2+ free  3.50  1714.0  2.07\u00D710-4  Fe3+/Fe2+ containing  2.65  470.0  2.39\u00D710-6  179  Cdl  (F cm-2)  RP  (\u2126  cm-2)  RCPE (\u2126 cm2)  5655.1  8.53\u00D710-6  2335.4  5.77\u00D710-6  cm2)  CP  (F  Q  (\u2126 s-n)  n  Chi-squared  1765.4  6.57\u00D710-4  0.84  6.34\u00D710-4  592.6  1.45\u00D710-3  0.51  9.59\u00D710-4  \u000CDouble-layer resistance is a characteristic quantity for the charge-transfer step of an electrochemical reaction which indicates its inherent speed (the higher the charge transfer resistance, the lower the reaction rate). From the Rdl values given in Table 9-1, the ferric reduction reaction (Rdl = 470 \u03A9 cm2) proceeds faster than chalcopyrite oxidation (Rdl = 1714 \u03A9 cm2). This means that the passivated chalcopyrite electrode is able to transfer charge at higher rates than its passive state would suggest. This finding agrees with our chronoamperometry results (Fig. 9-4b); and is also made clear in Fig. 9-5, which shows a much higher reaction rate for the iron-containing part of the curve (solid line) than for ironfree solution (dashed line). In general, a passive film (a copper sulfide film such as CuS 2, in the case of chalcopyrite [Yin et al., 1995]) on the surface of a material impedes the dissolution rate owing either to its low electronic or low ionic conduction. It is worth mentioning here that the electrochemical dissolution of the passive film at the film/solution interface cannot be the rate-controlling process [Macdonald 1999]. In fact, if this process were rate controlling then increasing the electrode potential would result in an exponentially increasing current as per the Butler-Volmer equation. In other words, if the electrochemical dissolution of the passive film governed the kinetics then the E-log i diagram of Fig. 9-1 would not show declining current densities at between ca. 0.65 and 0.85 V. Low electronic conduction of a passive film imposes a very high double-layer resistance on electrochemical reactions, while low ionic conduction of a passive film significantly reduces the diffusion rate of atoms through the passive film to the solid-solution interface. In the case of a passivated chalcopyrite electrode, our findings show that the passive film can transfer charge at high rates, indicating that the electronic conduction of the passive film is relatively high. Hence, the passive film of chalcopyrite impedes chalcopyrite dissolution due to its low ionic conduction. In other words, the passive dissolution of chalcopyrite is mass-transport-controlled by the passive film. The Cdl values (double-layer capacitance) of Table 9-1 also show a decrease from 207.00 to 2.39 \u03BCF cm\u22122 by addition of iron salts to the solution. The double-layer capacitance of an electrode can be estimated as follows [Bard and Faulkner 2001]:  180  \u000CCdl \uF03D  \uF065\uF065 o A d  9-6  where \u03B5o, A and d represent the permittivity of free space (8.85\u00D710\u221214 F cm-2), the exposed surface area of the electrode, and the thickness of the double-layer, respectively, and \u03B5 is the dielectric constant of the medium (i.e., solution in the double layer). The dielectric constant of sulfuric acid solutions ca. 25\u00BAC is about 90 and addition of ferrous to the solution lowers the value to below 10 [Greenwood and Earnshaw 1997]. Assuming these values can be applied for our condition, then Eq. 9-6 would give double-layer capacitances of about 100 and 1 \u03BCF cm\u22122 for iron-free and iron-containing solutions, respectively (note that the thickness of the double-layer was assumed to be 1 nm for these estimates). These estimated values agree very well with data given in Table 9-1. RP and CP represent the resistance and capacitance of the passive film, respectively. Table 9-1 shows similar values in both iron-free and iron-containing solution. In general, R P and CP values mostly depend on the thickness of a passive film, and typically this thickness depends on the applied potential. During the related experiment of Figs. 9-6a and b, the applied electrode potential was constant (0.70 V), thus one would not expect a considerable change in RP and CP values, as suggested by data given in Table 9-1. The thickness of a passive film can be estimated simply by rearranging Eq. 6 thus:  dP \uF03D  \uF065\uF065 o A CP  9-7  where dP represents the thickness of the passive film. The capacitance values of the passive film are presented in Table 9-1. If one assumes that the relative dielectric constant \u03B5 for the passive film of chalcopyrite is about 10 [Ghahremaninezhad et al., 2010a], then one can obtain a thickness of ca. 1 to 1.5 nm from Eq. 9-7. Calculations by Biegler [1977] have shown that the thickness of the passive film of chalcopyrite formed in sulfuric acid solution is ca. 2 nm and in another paper [Biegler and Horne 1985] this layer is reported to be ca. 3 nm thick. Both results agree fairly well with our estimates, indicating that the equivalent electrochemical circuit model of Fig. 9-7 and the calculated passive film capacitances given in Table 9-1 are likely valid. The constant phase element (CPE) and its associated resistance (RCPE) are the final elements of Fig. 9-7 to consider. In general, the CPE element behaves as a capacitance which 181  \u000Cvaries with frequency and RCPE represents its resistance in the equivalent circuit. The impedance of a CPE is presented by [Brug et al., 1984]: Z CPE \uF03D Q \uF02D1 ( j\uF077 ) \uF02D n  9-8  where Q is the CPE parameter; j = (\u22121)1/2; \u03C9 is the frequency in rad s\u22121, \u03C9 = 2\u03C0f; f is the frequency in Hz; and n is the measure of the non-ideality of the capacitor and has a value in the range 0 &lt; n &lt; 1. If n = 1, then ZCPE is identical to that of a capacitor, ZC = (j\u03C9C)\u22121, where C is the capacitance. Indeed, a capacitor is a constant phase element with a phase angle of 90\u00BA. Usually, CPE elements represent the effect of crystal defects in the passive film [Hirschorna et al., 2010]. Interestingly, in the case of chalcopyrite dissolution in sulfuric acid solution, the suggested passive film is a metal-deficient sulfide layer which forms by preferential dissolution of iron atoms from the mineral surface and, therefore, contains a high density of point defects (vacancies at iron positions within the chalcopyrite structure). Hence, in the case of iron-free solution, the CPE element and its associated resistance may be related to the point defects in the structure of the chalcopyrite passive film. The suggested n value for such a defective passive film is usually between 0.8 and 0.9 [Jorcin et al, 2006], which agrees well with the results given in Table 9-1 (n = 0.84). In contrast, the addition of iron salts to the solution decreased the value of n to 0.51. A value of n close to 0.5 (together with slope of the second time constant of the Nyquist plot in Fig. 6b which is ca. 1.0) are two independent indicators of a significant diffusion process in solution [Bard and Faulkner 2001]. Thus, in the presence of iron in solution, the CPE element of Fig. 7 most probably represents the diffusion of ferric towards the electrode surface. The potentiostatic polarization diagram of Fig. 9-1 and the i-t curves of Figs. 9-2a to c have shown that a steady-state passive film forms on the surface of chalcopyrite. Additionally, the chronoamperometry results shown in Fig. 9-4b, the results shown in Fig. 95 and a comparison of Rdl values given in Table 9-1 prove that the passive film of chalcopyrite is a better electronic conductor than an ionic conductor. Subsequently, one may safely conclude that the chalcopyrite passive film impedes the dissolution rate of chalcopyrite due to its low ionic conductivity (rather than its low electronic conductivity). In other words, the passive dissolution of chalcopyrite is mass transport-controlled by the passive film. If one assumes that iron dissolves preferentially to copper during chalcopyrite 182  \u000Cleaching and a passive film of metal-deficient sulfide layer forms on the surface [Hackl et al., 1995], then the diffusion of atoms (of mostly Fe) from the bulk of chalcopyrite to the passive film/solution interface is essential for chalcopyrite to be oxidized. In fact, copper and iron atoms must diffuse through the passive film to reach the passive film/electrolyte interface in order to be oxidized. However, as we have shown above, this diffusion step is slow and therefore controls the dissolution rate of chalcopyrite. During the ferric-assisted leaching of chalcopyrite concentrates, ferric reduction represents the major contribution to the total cathodic current. On the other hand, the diffusion of atoms through the passive film is slow, and therefore the number of electrons available for ferric reduction on the surface will be controlled by the diffusion of atoms through the passive film. This is illustrated in the schematic of Fig. 9-8, which shows that the rate of ferric reduction is completely dependent on the diffusion of iron and copper atoms within the passive film and their subsequent oxidation at the passive film/solution interface. Even high concentrations of ferric in solution cannot increase the rate of chalcopyrite dissolution because the electrons for the ferric reduction reaction must be provided by diffusion of atoms through the passive film lattice. Therefore, considering the fact that the main anodic reaction in chalcopyrite leaching is chalcopyrite oxidation, the slow rate of chalcopyrite oxidation (due to the slow diffusion process) will control the rate of ferric reduction on the surface of chalcopyrite. In other words, sluggish chalcopyrite oxidation (dissolution) is the reason for slow ferric reduction on the surface of chalcopyrite. These findings are different from the proposed idea that pyrite-assisted galvanic leaching of chalcopyrite in ferric-containing solution results in a substantially higher dissolution rate [Berry et al., 1978; Koleini et al., 2011]. In fact, the results of a recent study by Eghbalnia and Dixon [2011] have shown that copper extraction from pyrite-assisted leaching in iron sulfate media is only slightly faster than in the absence of pyrite when the pyrite has little or no silver content [Eghbalnia and Dixon 2011; Nazari et al., 2011]. Generally, it is true that the addition of pyrite to a chalcopyrite concentrate leach solution will speed up the ferric reduction rate, as it has already been suggested that the ferric-ferrous couple is rapid on pyrite surfaces [Mishra and Osseo-Asare 1988]. However it cannot be stated conclusively that the rate of chalcopyrite oxidation will also be increased. Indeed, a recent study by Nazari et al. [2011] demonstrated that not all pyrite minerals increase the 183  \u000Cdissolution rate of chalcopyrite leaching significantly, and that silver content of the pyrite is a key factor in determining the leaching rate of chalcopyrite concentrates. Furthermore, Nazari\u2019s study has shown that the co-presence of pyrite and silver is essential for rapid leaching of chalcopyrite, and that the addition of silver or pyrite alone do not result in a substantial increase in the chalcopyrite dissolution rate. It should be mentioned that Nazari et al. [2011] have suggested that chalcopyrite leaching is difficult due to the slow kinetics of ferric reduction on its surface, whereas, the results of this work present a different scenario. Findings of the present study support the idea that the presence of silver in chalcopyrite leach solutions may alter the properties of the chalcopyrite passive film, and this idea will be revisited in a future paper. Nevertheless, a detailed and alternative mechanism of chalcopyrite leaching in the co-presence of pyrite and silver is suggested and discussed elsewhere [Nazari et al., 2012].  4Fe3+  4Fe2+  4e-  CuFeS2  Cu2+ Fe2+  Chalcopyrite Passive film Figure 9-8: Schematic illustration for the dissolution process of a passivated chalcopyrite mineral particle in ferric-containing sulfuric acid solution. High concentration of ferric (0.1 M) is available on the surface for cathodic reduction, and anodic oxidation of chalcopyrite must provide the required electrons for ferric reduction. However, diffusion of iron and copper atoms through the passive film is slow (ratecontrolling step) and so the electron production rate will be slow and consequently so will be the ferric reduction rate.  184  \u000C9.4 Conclusions In summary, the kinetics of the ferric-ferrous couple on an anodically passivated chalcopyrite surface in 0.5 M sulfuric acid solution at 25\u00BAC was investigated for the first time. For this redox reaction, the exchange current density of 2.44 \u03BCA cm\u22122 and the anodic and cathodic transfer coefficients of 0.37 and 0.50 were obtained, respectively. In addition, it was shown that the ferric reduction reaction on the surface of chalcopyrite is relatively fast (i.e. faster than oxidation of chalcopyrite itself). It was shown that the slow kinetics of ferric reduction on chalcopyrite surfaces during the leaching process is due to slow diffusion of Fe atoms through a passive film.  185  \u000C10 A proposed electrochemical dissolution model 10.1 Introduction In this chapter, a possible anodic dissolution model for chalcopyrite in 0.5 M sulfuric acid solution based on the findings of the previous chapters will be proposed.  10.2 Model development The general dissolution behavior of chalcopyrite in 0.5 M sulfuric acid solution can be considered to occur in two different modes: (1) passive dissolution, and (2) transpassive dissolution [Ghahremaninezhad et al. 2010] (note that the anodically active dissolution of chalcopyrite is not an option here since the chalcopyrite surface at potentials even slightly above its OCP (or OCP itself) is covered with a thin layer of passive film [Ghahremaninezhad et al. 2010]). In the passive state, a passive film blocks the surface of the electrode and hinders the dissolution rate, while in the transpassive state, the passive film is not stable, and dissolution of chalcopyrite proceeds without any serious barrier (formation of non-passive surface films is possible at elevated potentials). From our XPS results (chapter 7), the ultimate passivating agent for chalcopyrite may be copper sulfide (CuS) and metal-deficient sulfide (e.g. Cu1-xFe1-yS2-z). There is a distinct difference between the transpassive dissolution of sulfide minerals (e.g. chalcopyrite) and that of most metals and alloys (e.g. stainless steel). In the case of passive metals, a metal oxide film forms on the surface and passivates the metal (e.g. chromium oxide film on the surface of stainless steel). In the transpassive region, metal atoms of the passive film structure on metals oxidize to higher oxidation states and dissolve in solution. In the case of sulfide minerals, e.g. chalcopyrite, the transpassive dissolution proceeds after oxidation of sulfide species to higher oxidation states such as thiosulfate, sulfate or elemental sulfur. Hence, in the case of this study, sulfide ions in the passive film must oxidize to higher oxidation states in order for chalcopyrite to dissolve. We have shown that copper sulfide (CuS) and metal-deficient sulfide (Cu1-xFe1-yS2-z) are the most plausible phases present on the surface of chalcopyrite in the passive state (cf. chapter 7). In addition, iron oxyhydroxide (FeOOH) was also detected on the surface (cf. chapter 7). It was proposed 186  \u000Cthat the low potential anodic dissolution of chalcopyrite proceeds through a reaction such as reaction 7-5 (cf. chapter 7). Reaction 7-5 is an overall anodic reaction which takes into account the formation of a metal-deficient sulfide film (as a passive film phase). In addition, this reaction includes the formation of soluble ions of iron, copper and sulfur as ferrous, cupric and thiosulfate ions, respectively. Indeed, from the characteristics of chalcopyrite dissolution, the formation of the latter soluble ions (or other soluble species of Fe, Cu and S) seems to be a mandatory requirement. One of the most important characteristics of chalcopyrite dissolution, which also was discussed in a previous study (cf. chapter 7), is presented in Fig. 10-1, i.e. potentiostatic dissolution of chalcopyrite at 0.80 V for about 60 hours results in a sharp decrease in the current density value and afterwards a very low steady state current density. In other words, if one considers a reaction such as [Yin et al. 1995]: CuFeS 2 \uF03D CuS *2 \uF02B Fe 2\uF02B \uF02B 2e \uF02D  10-1  as the general passivation reaction of chalcopyrite, the charge associated with the experiment of Fig. 10-1, i.e. total q = 297 mC cm-2, will result in a relatively thick passive film after some hours of dissolution. An estimate of the thickness of the CuS *2 film (product of reaction 10-1, suggested by Yin et al., [1995]) is presented in Fig. 10-2. This figure shows a thickness of the passive film of ca. 200 nm after 20 hours and the thickness increases to ca. 400 nm after 60 h. One would expect a significant decrease in the dissolution current density as a function of time (from Fick\u2019s first diffusion law). For instance, if the thickness of the passive film doubles (during the 60 h of experiment), then the current density (current flux) should decrease to half. However the current density of Fig.10-2 is almost steady state (i.e. not decreasing).  187  \u000CFigure 10-1: Passive current vs. time respond of a chalcopyrite electrode to potentiostatic polarization at 800 mV, and related thickness of the passive film ( CuS *2 ) calculated from reaction 10-1.  On the other hand, the reaction suggested in this study, i.e. reaction 7-5, considers both passivation of chalcopyrite, by formation of metal-deficient sulfide phase, and dissolution of all three elements present in the structure of chalcopyrite. In addition, a fraction of the iron-sulfide species (Fe-S bonds in the passivating film) may react with water to form FeOOH (formation of iron oxyhydroxide (FeOOH) was suggested by our XPS studies in chapter 7). It has been proven that Fe-S species (species with iron-sulfide bond) are good catalysts for oxygen reduction [Susac et al. 2007]. In fact, it has been suggested that pyrite electrodes (Fe-S bonds) can successfully coordinate OH\uF02D and water molecules on their surface and promote the associated reactions [Mishra and Osseo-Asare 1988]. If this is true, then one would expect that the Fe-S bonds of the chalcopyrite passive film coordinate water molecules and form Fe-O species (FeOOH). Then it would be expected that the passive surface of chalcopyrite could be covered with a FeOOH product layer. This schematic is presented in Fig. 10-2. In order to confirm the order of surface films (formation of a FeOOH product layer on top of sulfide passive film), an angle-resolved XPS experiment was performed on the surface of a chalcopyrite electrode after 6 hours of dissolution at 0.70 V. In this experiment (presented in Fig. 10-3) the take-off angle of the photoelectrons for Fig. 103a was 45\u00BA and in the experiment of Fig. 10-3b the take-off angle of photoelectrons was 188  \u000Cincreased to 90\u00BA. The penetration depth for the photoelectrons with the high angle (90\u00BA) is about 10 nm while at low angles, e.g. 45\u00BA, a lower penetration depth is expected. Under this condition, if the position of the FeOOH film is on top of the sulfide passive film, i.e. the Cu1xFe1-yS2-z  film as presented in schematic of Fig. 10-2, then the ratio of Fe-O species to Fe-S  species must show an increase. The fitting procedure of the Fe 3p level spectrum has been explained elsewhere (rf. Chapter 7) and the ratio of Fe-O to Fe-S can be acquired from the area ratio under the related peaks for iron oxyhydroxide and iron sulfide species. The Fe 3p spectra of Figs. 10-3a and b show two main peaks centered at 53.3 (peak 1) and 55.5 eV (peak 2). The peak at 53.3 eV agrees well with the Fe 3p peak of Fe3+ in Fe-S species and the peak at 55.5 eV matches well with the Fe 3p peak of Fe3+ in FeOOH (rf. Chapter 7). The ratio of Fe-O to Fe-S is c.a. 91:9 and 85:2 for Figs. 10-3a and 10-3b, respectively. The reason for this slight increase in the Fe-O to Fe-S ratio, obtained by decreasing the take-off angle of the photoelectrons, is that the FeOOH layer covers the sulfide passive film as shown in Fig. 10-2. Note that the partial dissolution of the passive film, to form FeOOH, might result in flaws on its outer layer and consequently the FeOOH film may be porous. In addition, the thickness of the FeOOH film must be less than 10 nm, otherwise we wouldn\u2019t be able to detect the Fe-S and Cu-S species by XPS.  FeOOH  CuFeS2  Metal- deficient sulfide film  Figure 10-2: Schematic illustration for the growth of surface films on chalcopyrite surface.  189  \u000CFigure 10-3: XPS spectrum of Fe 3p level obtained from a chalcopyrite electrode after 6 hours dissolution at 700 mV. The take-off angle of the photoelectrons was (a) 45\u00BA, (b) 90\u00BA.  The results presented here are very similar to those of chapter 6 which employed EIS to elucidate the nature of the passive film. However, in our previous study, since we didn\u2019t perform any analytical methods on the surface of the electrode for compositional characterization of the surface films, reactions proposed by other workers were selected to explain the dissolution/passivation behavior that we observed. In this previous work the  190  \u000Cdissolution reaction of chalcopyrite at low anodic potentials, which results in passivation of chalcopyrite, was suggested to be: CuFeS 2 \uF0BE \uF0BE\uF0AE Cu 1\uF02D x Fe 1\uF02D y S 2 \uF02B xCu 2\uF02B \uF02B yFe 2 \uF02B \uF02B 2( x \uF02B y )e \uF02D  y\uF0F1\uF0F1 x  10-2  and it was suggested that the layer on top of the passive film forms by the following reaction: Cu 1\uF02D x Fe1\uF02D y S 2 \uF0BE \uF0BE\uF0AE Cu 1\uF02D x\uF02D z S 2 \uF02B zCu 2\uF02B \uF02B (1 \uF02D y )Fe 2 \uF02B \uF02B 2( z \uF02B 1 \uF02D y)e \uF02D  10-3  This conclusion was made based on the measured and modeled impedance spectra, which revealed that in the passive state, there must be two surface layers on the surface of chalcopyrite - the topmost layer forming by partial dissolution of the passive film (reaction 10-3). This conclusion is in good agreement with the results of chapter 7 which suggests the formation of a second layer (FeOOH film) by partial dissolution of the passive film (metaldeficient sulfide film of Cu1-xFe1-yS2-z). In addition, the main product of reaction 10-2 is very similar to the product of reaction 7-5, as they both produce a passive metal-deficient sulfide film. For complete dissolution of chalcopyrite (transpassive dissolution), we would suggest that at high enough potentials (above 0.90 V), the passive film of chalcopyrite dissolves by reaction 8-10. Complete dissolution of chalcopyrite starts after this stage and proceeds according to reactions such as 8-11 to 8-12, but not exclusive to these reactions. One final and very important point which must be discussed is the elemental sulfur yield in the low temperature acidic (around 80\u00BAC) leaching processes of chalcopyrite reported by different workers [Klauber et al. 2001; Koleini et al. 2011; Nazari et al. 2011]. This elemental sulfur yield has led to the conclusion by several authors that passivation is caused by a sulfur film [Klauber et al. 2001]. In contrast, results of electrochemical studies on chalcopyrite dissolution tend to support the passivation of chalcopyrite by a sulfide passive film [Viramontes-Gamboa et al. 2007; Parker et al. 1981; Linge 1976]. We believe that the thiosulfate present in solution, originated from chalcopyrite dissolution by reactions such as 7-5, can be the precursor for elemental sulfur formation. Indeed, the formation of elemental sulfur from thiosulfate is a reduction reaction that may be presented as [Marcus and Protopopoff 1977]: S 2 O 32 \uF02D \uF02B 2O ads \uF02B 10H \uF02B \uF02B 8e \uF02D \uF03D 2S \uF02B 5H 2 O  10-4  Reaction 10-4 is consistent with many experimental leaching process results, as it has been observed that an increase in the oxygen pressure during the chalcopyrite leaching 191  \u000Cincreases the elemental sulfur yield [Dixon and Tshilombo 2010]. Furthermore, it has been shown that higher acidity of chalcopyrite leach solution will result in higher elemental sulfur yield [Marsden and Wilmot 2007] which is consistent again with reaction 10-4. Further evidence for the validity of reaction 10-4 may be gleaned from pyrite leaching studies in acidic ferric sulfate solutions [Bouffard et al. 2006]. The main product of pyrite leaching is known to be thiosulfate [Schippers and Sand 1999]. Studies by Bouffard et al. [2006] has shown that increasing the redox potential in pyrite leaching, lowers the elemental sulfur yield (i.e. increases the sulfate formation, instead). In fact, this can be explained by reaction 10-4, as by increasing of the redox potential the rate of thiosulfate reduction decreases, i.e. increase of redox potential will impose a more oxidizing condition in the leach solution and reduction reactions such as reaction 10-4 will be less favorable). Alternatively, it has been shown by Johnston and McAmish [1973] that in a dilute acidic thiosulfate solution (HCl concentration between 1.95 \u2013 8.16 mM), colloidal sulfur can form with a rapid third-order reaction. The suggested overall reaction is: H \uF02B \uF02B S 2 O 32\uF02D \uF03D S \uF02B HSO 3\uF02D  d [S] \uF03D k[H \uF02B ][S2 O 32\uF02D ] 2 dt  10-5  and the reaction is shown to be a rapid reaction with the k value of: k \uF03D (1 .64 \uF0B1 0.13) \uF0B4 1011 exp(  \uF02D 15930 \uF0B1 160 l2 ) RT mole 2 sec1  10-6  Reaction 10-5 is the disproportionation reaction of thiosulfate which simultaneously reduces and oxidizes the thiosulfate reactant to elemental sulfur and bisulfate products, respectively. If the above pathways for formation of elemental sulfur in leaching of chalcopyrite are true, then reaction 10-4 and 10-5 suggest that the sulfur yield in the concentrate leaching experiments originates from thiosulfate, which itself is a product of chalcopyrite dissolution. Thus So does not form directly from chalcopyrite dissolution. Such a reaction scheme, i.e. the formation of elemental sulfur from thiosulfate in acidic solutions, has also been suggested for the leaching of pyrite [Chandra and Gerson 2010]. Nonetheless, reaction 10-5 can form at most 50% elemental sulfur from its reactant thiosulfate while the reported elemental sulfur yield for chalcopyrite is around 90% or above [Dutrizac 1989]. This means that reaction 10-4 is likely the main sulfur producing reaction, in comparison to  192  \u000Creaction 10-5. If this is true, then there is need for further studies to evaluate the validity of reaction 10-4 in the case of chalcopyrite leaching.  10.3 Conclusion In summary, thiosulfate is proposed to be the soluble sulfur species which forms during the dissolution of chalcopyrite. The absence of elemental sulfur during electrochemical dissolution of chalcopyrite and its presence in leaching processes were discussed. Elemental sulfur is believed to be the product of thiosulfate reduction. However, this needs further study to be confirmed as we do not have any direct evidence for this type of reaction.  193  \u000C11 Conclusions and recommendations The results of the electrochemical and surface analysis experiments of this study on the electrochemistry of pyrrhotite and chalcopyrite minerals have generated the following conclusions and recommendations.  11.1 Conclusions \uF0B7  The electrochemical dissolution behavior of pyrrhotite has been studied in 1M HCl  solution with in-situ electrochemical methods (chapter 5). Results showed the formation of two distinct surface layers on the electrode. At low anodic dissolution potentials, the iron atoms from the pyrrhotite structure partially dissolve and a metal\u2013deficient sulfide layer with semiconducting properties (Fe1-x-yS) forms on the surface of the electrode. This layer is stable up to potential of 0.84 VSHE. The second film forms at higher anodic potentials (&gt; 0.91 VSHE) and is substantially more resistant to dissolution than the first. The results of this study are directly related to those of chalcopyrite electrodissolution. As pyrrhotite represents a simplified case of chalcopyrite dissolution, we would expect to observe relatively similar morphology for the passive film of chalcopyrite. More specifically, in the case of chalcopyrite dissolution, a copper sulfide surface film forms on its surface which is analogous to that in pyrrhotite electrodissolution, i.e. both films are semiconductor films and both form by partial dissolution of atoms from the mineral structure. \uF0B7  From EIS studies on the chalcopyrite electrode (chapter 6), it was shown that at low  potentials (below 0.7 VSHE) a surface layer forms on the surface of chalcopyrite. Above this potential (0.7 \u2013 0.9 VSHE) the previously formed surface layer partially dissolves and a second passive layer forms. Both of these layers display the characteristics of passive layers at low scan rate (0.05 mV s-1) and for high scan rates they are observed to be pseudo-passive. In the potential range of 0.9 to 1.02 VSHE, none of the surface layers are stable and electrochemical active dissolution of the electrode occurs. \uF0B7  From the electrochemical experiments of chapter 7 it was concluded that  potentiodynamic polarization of chalcopyrite, even at a very low scan rate, cannot reliably and successfully demonstrate active-passive dissolution of chalcopyrite. On the other hand, 194  \u000Cthe results of potentiostatic polarization have agreed very well with those of XPS studies, which suggest that this technique can accurately demonstrate the active-passive dissolution behavior of chalcopyrite. From potentiostatic polarization it was concluded that chalcopyrite is passive at potentials below 0.90 VSHE, beyond which it actively dissolves. \uF0B7  Results of XPS studies (chapter 7) have shown a substantial change in the oxidation  state of sulfur species on the surface of chalcopyrite at potentials above 0.90 VSHE. XPS experiments revealed that copper and iron sulfide bonds are present in the passive film of chalcopyrite, but their binding energies are different from those of chalcopyrite. These bonds were  believed  to  originate  from  a  metal-deficient  sulfide  passive  film  with  composition Cu 1\uF02D x Fe1\uF02D y S 2 -z . In addition, FeOOH was detected on the surface and a possible reaction mechanism for its formation was suggested. \uF0B7  XPS results of chapter 7 also showed that the surface of passive chalcopyrite is  blocked by a sulfide film and beyond the passive region (above 0.90 VSHE) the percentage of sulfides in the film decreased significantly, i.e. monosulfide decreased from 56% to 19% and disulfide decreased from 39% to 18%. Furthermore, no elemental sulfur or polysulfide was detected on the surface at potentials below 0.90 VSHE. Hence, it is concluded that the transpassive dissolution of chalcopyrite is significantly linked to the oxidation of sulfur (from sulfide in the passive film to elemental sulfur and perhaps to sulfur species with higher oxidation states, e.g. thiosulfate). Additionally, by comparing XPS sulfur spectra of chalcopyrite with the results of sulfur spectra from iron sulfide (pyrite and pyrrhotite) and copper sulfide minerals (chalcocite and covellite) we have concluded that the high energy feature in sulfur species spectra (~ 163.3 eV) is most likely related to the S 3p\u2192Fe 3d excitation and does not originate from polysulfide. \uF0B7  From the cyclic voltammetry experiments of chapter 8, it was shown that iron atoms  dissolve preferentially from the surface of chalcopyrite during anodic dissolution, leaving behind a copper-rich metal-deficient sulfide passive film (i.e. Cu1\u2212xFe1\u2212yS2 y &gt; x). By integrating the current under the chalcopyrite de-passivation peak, the thickness of the passive film was calculated to be ca. 6.7 nm. Additionally, experiments showed that elemental sulfur is a product of chalcopyrite oxidation at anodic potentials higher than 1.0 VSHE.  195  \u000C\uF0B7  Results of chapter 8 also showed that cathodic dissolution of chalcopyrite proceeds at  potentials below 0 VSHE. The cathodic dissolution process involves the release of ferrous and H2S(aq) to solution. The cathodic reactions which occur in the potential region of 0.1 to 0.6 VSHE and the anodic reactions which occur at potentials lower that 0.5 VSHE are all related to the reduction and oxidation of soluble species. The soluble species involved in those cathodic reactions are ferric and cupric and the soluble species involved in the anodic reactions is H2S(aq). \uF0B7  Chapter 9, for the first time, presents the kinetics of the ferric-ferrous couple on an  anodically passivated chalcopyrite surface in 0.5 M sulfuric acid solution at 25\u00BAC was investigated. For this redox reaction, the exchange current density of 2.44 \u03BCA cm\u22122 and the anodic and cathodic transfer coefficients of 0.37 and 0.50 were obtained, respectively. In addition, it was shown that the ferric reduction reaction on the surface of chalcopyrite is relatively fast (i.e. faster than oxidation of chalcopyrite itself). It was shown that the slow kinetics of ferric reduction on chalcopyrite surfaces during leaching processes is due to slow diffusion of Fe atoms through the passive film. \uF0B7  Coupling the results of chapters 5-9, a reaction model for the anodic chalcopyrite  dissolution was proposed. Thiosulfate is proposed to be the soluble sulfur species which forms during the dissolution of chalcopyrite. The absence of elemental sulfur during electrochemical dissolution of chalcopyrite and its presence in leaching processes were discussed. Elemental sulfur is believed to be the product of thiosulfate reduction. However, this needs further studies to be confirmed.  11.2 Recommendations The recommended future work is the following: \uF0B7  Dissolution of chalcopyrite through the formation of soluble sulfur species should be  further investigated. In this regard, one might put a platinum plate in front of a chalcopyrite electrode. The chalcopyrite electrode should be dissolved at anodic potentials and the platinum plate must be kept at a cathodic potentials. Note that for this experiment two potentiostats are required. Then the platinum electrode can be analyzed for sulfur species by means of different techniques such as XPS, SIMS etc. 196  \u000C\uF0B7  The effect of the temperature increase on the passivity of chalcopyrite was not  considered in this study. Since typical leaching processes of chalcopyrite proceed at elevated temperatures, this effect should be studied. \uF0B7  The aim of this study was to produce very fundamental knowledge about the  dissolution of chalcopyrite in a simple solution. However, the effect of the different acids and acid concentrations on the passivation of chalcopyrite needs to be further elucidated. \uF0B7  The findings of this study, i.e. composition of passive film and its chemical and  semiconductor properties, are helpful regarding the selection of reagents for leaching. Leaching of chalcopyrite at low anodic potentials (below 0.75 VSHE) requires an aggressive reagent (e.g. Cl\u2013) which can break down chalcopyrite\u2019s passive film by pit formation. The selection of proper reagents for further study requires comprehensive knowledge of the physio-chemical properties of chalcopyrite\u2019s passive films, which are provided in this thesis.  197  \u000CReferences Abraitis, P.K., Pattrick, R.A.D., Kelsall, G.H., Vaughan, D.J., 2004, Mineralogical Magazine, 68(2), p. 343. Acres, R.C., Harmer, S.L., Beattie, D.A., 2010, Minerals Engineering 23, p. 928. Acres, R.C., Harmer, S.L., Shui, H.W., Chen, C.H., Beattie, D.A., 2011, J. Synchrotron Radiation, 18, p. 649. Almeida, C.M.V.B., Giannetti, B.F., 2003, J. Electroanal. Chem., 553, p. 27. Asselin, E., Alfantazi, A., Rogak, S., 2007, J. Electrochem. Soc. 154, p. C215. Austin, I.G., Goodman, C.H.L., Pengelly, A.E., 1956, J. Electrochem. Soc. 103, p. 609. Avgouropoulos, G., Ioannides, T., 2003, Appl. Catal. A, 244, p. 155. Aznar, J.C., Richer-Lafleche, M., Cluis D., 2008, Environmental Pollution, 156, p. 76. Azumi, K., Ohtsuka, T., Sato, N., 1987, J. Electrochem. Soc. 134, p. 1352. Bard, A.J., Faulkner, L.R., 2001, Electrochemical methods, 2nd Ed., John Wiley &amp; Sons, NY, p. 745. Beolchini, F., Veglio, F., 1999, Ind. Eng. Chem. Res., 38, p. 3296. Berry, V.K., Murr, L.E., Hiskey, J.B., 1978, Hydrometallurgy, 3, p. 309. Betova, I., Bojinov, M., Laitinen, T., Makela, K., Pohjanne, P., Saario, T., 2002, Corr. Sci. 44, p. 2675. Betova, I., Bojinov, M., Tzvetkoff, T., 2004, Electrochim. Acta 49, p.2295. Bevilaqua, D., Di\u00E9z-Perez, I., Fugivara, C.S., Sanz, F., Benedetti, A.V., Garcia, O., 2004, Bioelectrochemistry, 64, p. 79. Biegler, T., 1977, J. Electrochem. Soc. 85, p. 101. Biegler, T., Horne, M.D., 1985, J. Electrochemical Soc. 132, p. 1363. Biegler, T., Swift, D.A., 1979, J. Appl. Electrochem. 9, p. 545. Biernat, R.J, Robins, R.G., 1972, Electrochimica Acta 17, p. 1261. Biswas, A.K., Davenport, W.G., 1976, Extractive Metallurgy of Copper, Pergamon Press, Oxford. Boekema, C., Krupski, A.M., Varaste, M., Parvin, K.,VanTil, F.,VanDerWoude, F., Sawatzky, G.A., 2004, J. Magn. Magn. Mater. 272\u2013276, p. 559. Bojinov, M, 1996, J. Electroanal. Chem. 405, p. 15. Bojinov, M., Betova, I., Fabricius, G., Laitinen, T., Raicheff, R., Saario, T., 1999, Corr. Sci. 41, p. 1557. AND, Bojinov, M., Fabricius, G., Laitinen, T., Saario, T., 1999, Electrochim. Acta 44, p.4331. Bojinov, M., Fabricius, G., Kinnunen, P., Laitinen, T., Makela, K., Saario, T., Sundholm, G., 2000, Electrochimica Acta 45, p. 2791. 198  \u000CBojinov, M., Fabricius, G., Kinnunen, P., Laitinen, Y., M\u00E4kel\u00E4, K., Saario, T., Sundholm, G., Yliniemi, K., 2002, Electrochim. Acta 47 p. 1697. Bojinov, M., Salmi, K., Sundholm, G., 1993, J. Electroanal. Chem. 347, p. 207. Bojinov, M., Tezvetkoff, T., 2003, J. Phys. Chem. B 107, p. 5101. Brion, D., 1980, Appl. Surf. Sci., 5, p. 133. Buckley, A.N., Kravets, I.M., Shchukarev, A.V., Woods, R., 1994, J. Appl. Electrochem., 24, p. 513. Buckley, A.N., Woods, R., 1994, J. Electroanal. Chem. 370, p. 295. Buckley, A.N., Woods, R., 1996, The galena surface revisited, in: R. Woods, P. Richardson, F. Doyle (Eds.), Electrochemistry in Minerals and Metals Processing IV, The electrochemical Society, Pennington, NJ, p. 1. Burdick, C.L., Ellis, J.H., 1917, J. Am. Chem. Soc., 39, p. 2518. Cardon, F., 1972, Physica 57, p. 390. Chandra, A.P., Gerson, A.R., 2010, Surf. Sci. Reports, 65, p. 293. Chen, L., Zhou, Y., Krause, S., Munoz, A.G., Kunze, J., Schmuki, P., 2008, Electrochim. Acta 53, p. 3395. Chernyshova, I.V., Andreev, S.I., 1997, Appl. Surf. Sci. 108, p. 225. Chusuei, C.C., Brookshier, M.A., Goodman, D.W., 1995, Langmuir, 15, p. 2806. Clayton, C.R., Lu, Y.C., 1986, J. Electrochemical Soc. 133, p. 2465. C\u00F3rdoba, E.M., Mu\u00F1oz, J.A., Bl\u00E1zquez, M.L., Gonz\u00E1lez, F., Ballester, A., 2008, Hydrometallurgy 93, p. 81. Descostes, M., Mercier, F., Thromat, N., Beaucaire, C., Gautier-Soyer, M., 2000, Appl. Sur. Sci. 165, p. 288. Dixon, D.G., Mayne, D.D., Baxter, K.G., 2008, Can. Metall. Q. 47, p. 327. Donovan, B., Reichenbaum, G., 1958, J. Appl. Phys. 9, p. 474. Donnay, G., Coliss, L.M., Donnay, J.D.H., Elliott, N., Hastings, J.M., 1958, Phys. Rev. 112, p. 1917. Dreisinger, D., 2006, Hydrometallurgy 83, p. 10. Dreisinger, D., Abed, N., 2002, Hydrometallurgy 66 p. 37. Droppert, D.J., Shang, Y., 1995, Hydrometallurgy 39, p. 169. Dutrizac, J.E., 1978, Metall. Trans. B 9B, p. 431. Dutrizac, J.E., 1982, Jarosite-type compounds and their application in the metallurgical industry, In: Osseo-Asare, K., Miller, J.D. (Eds.), Hydrometall. Res., Dev. Plant Pract., Proc. Int. Symp., 3rd, Metall. Soc. AIME, Warrendale PA, pp. 531\u2013551. Dutrizac, J.E., 1989, Can. Metall. Q. 28, p. 337. Dutrizac, J.E., 1996, Hydrometallurgy 42, p. 293. 199  \u000CDutrizac, J.E., MacDonald, R.J.C., 1974, Miner. Sci. Eng. 6, p. 59. Eghbalnia, M., Dixon, D.G., 2011, Hydrometallurgy 110, p. 1. Ellisa, S., Senanayake, G., 2004, Hydrometallurgy 72, p. 39. Elsherief, A.E., 2002, Minerals Eng. 15, p. 215. Ennaoui, A., Fiechter, S., Jaegermann, W., Tributsch, H., 1986, J. Electrochem. Soc. 133, p. 97. Epelboin, I., Keddam, M., 1972, Electrochimica Acta 17, p. 177. Erne, B.H., Vanmaekelbergh, D., 1997, J. Electrochem. Soc. 144, p. 3385. Evans, H.T., Konnert, J.A., 1976, Am. Mineral., 61, p. 996. Farquhar, M.L., Wincott, P.L., Wogelius, R.A., Vaughan D.J., 2003, Appl. Surf. Sci. 218, p. 34. Ferreira de Lima, G. Oliveira, C., Avelino de Abreu, H., Duarte, H.A., 2011, J. Phys. Chem. C, 115, p. 10709. Filippou, D., Konduru, R., Demopoulos, G.P., 1997, Hydrometallurgy, 47, p. 1. Folmer, J.C.W., Jellinek, F., 1980, J. Less Common Metals 76, p. 153. Fujisawa, M., Suga, S., Mizokawa, T., Fujimori, A., Sato, K., 1994, Phys. Review B, 49, p. 7155. Gabrielli, C., Keddam, M., Minouflet, F., Perrot, H., 1996, Electrochim. Acta 41, p. 1217. Garrels, R.M., Christ, C.L., 1965, Solution, Minerals, and Equilibria, Harper &amp; Row, NY, p. 213. Ghahremaninezhad, A., Asselin, E., Dixon, D.G., 2011, J. Phys. Chem. C, 115, p. 9320. Ghahremaninezhad, A., Asselin, E., Dixon, D.G., 2010, Electrochim. Acta, 55, p. 5041a. Ghahremaninezhad, A., Asselin, E., Dixon, D.G., 2010, J. Electrochem. Soc. 157, p. C248b. Glarum, S.H., Marshall, J.H., 1985, J. Electrochem. Soc. 132, p. 2878. Glatzel, T., Fuertes-Marr\u00F3n, D., Schedel-Niedrig, T., Sadewasser, S., Lux-Steiner, M.C., 2002, Appl. Phys. Lett., 81, p. 2017. Goh, S.W., Buckley, A.N., Skinner, W.M., Fan, L.L., 2010, Phys. Chem. Minerals, 37, p. 389. Goh, S.W., Buckley, A.N., Lamb, R.N., 2006, Minerals Engineering 19, p. 204. Gomes, W.P., Vanmaekelbergh, D., 1996, Electrochim. Acta 41, p. 967. Gordon, R.B., Bertram, M., Graedel, T.E., 2006, PNAS 103, p. 1209. Greenwood, N.N., Whitfield, H.J., 1968, J. Chem. Soc. A. London 7, p. 1697. Gryse, R.D., Gomes, W.P., Cardon, F., Vennik, J., 1975, J. Electrochem. Soc. 122, p. 711.  200  \u000CHabashi F., 1997, Handbook of Extractive Metallurgy, Vol. 2, Wiley-VCH, Weinheim, Germany, p 491. Habashi F., 1978, Chalcopyrite, its chemistry and metallurgy, McGraw-Hill, New York. Hackl, R.P., Dreisinger, D.B., Peters, E., King, J.A., 1995, Hydrometallurgy 39, p. 25. Hall, S.R., Stewart, J.M., 1973, Acta Cryst. B29, p. 579. Hamajima, T., Kambara, T., Gondaira, K.I., 1981, Phys. Rev. B, 24, p. 3349. Hamilton, I.C., Woods, R., 1981, J. Electroanal. Chem, 118, p. 327. Harmer, S.L., Pratt, A.R., Nesbitt, H.W., Fleet, M.E., 2005, Canadian Mineralogist 43, p. 1619. Harmer, S.L., Thomas, J.E., Fornasiero, D., Gerson, A.R., 2006, Geochim. et Cosmochim. Acta 70 p. 4392. Hedberg, E., Gidhagen, L., Johansson, C., 2005, Atmospheric Environment 39, p. 549. Hiroyoshi, N., Arai, M., Miki, H., Tsunekawa, M., Hirajima, T., 2002, Hydrometallurgy 63, p. 257. Hiroyoshi, N., Kuroiwa, S., Miki, H., Tsunekawa, M., Hirajima, T., 2004, Hydrometallurgy 74, p. 103. Hiroyoshi, N., Miki, H., Hirajima. T., Tsunekawa, M., 2001, Hydrometallurgy, 60, p. 185. Hirschorn, B., Orazem, M.E., Tribollet, B., Vivier, V., Frateur, I., Musiani, M., 2010, Electrochimica Acta 55, p. 6218. Hiskey, J.B., 1993, chalcopyrite semiconductor electrochemistry and dissolution, In: Reddy, R.G.,Weizenbach, R.N. (Eds.), The Paul E. Queneau International Symposium, Extractive Metallurgy of Copper, Nickel and Cobalt, Volume I: Fundamental Aspects, The Minerals, Metals &amp; Materials Society, p. 949. Holliday, R.I., Richmond, W.R., 1990, J. Electroanal. Chem. 288, p. 83. Hsu, C.H., Mansfeld, F., 2001, Corrosion 57, p. 747. Hu, J., Lu, Q., Deng, B., Tang, K., Qian, Y., Li, Y., Zhou, G., Liu, X., 1999, Inorg. Chem. Commun. 2 p. 569. Johnston F., McAmish, L., 1973, J. Colloids Interface Sci. 42, p. 112. Jones, C.F., LeCount, S., Smart, R.S., White, T.J., 1992, Appl. Surf. Sci. 55, p. 65. Kaplun, K., Li, J., Kawashima, N., Gerson, A.R., 2011, Geochimi. et Cosmochimi. Acta, 75, p. 5865. Kapusta, S., Hackerman, N., 1981, J. Electrochem. Soc. 128, p. 327. Kareh, K., Ghahremaninezhad, A., Asselin, E., 2009, Electrochim. Acta 54, p. 6548. Karthe, S., Szargan, R., Suoninen, E., 1993, Appl. Surf. Sci. 72, p. 157. Ke, J., Li, H., 2006, Hydrometallurgy 82, p. 172. Keddam, M., Takenouti, H., Yu, N., 1985, J. Electrochem. Soc. 132, p. 2561. 201  \u000CKennedy, J.H., Frese Jr., K.W., 1978, J. Electrochem. Soc. 125, p. 723. Klauber, C., 2008, Int. J. Miner. Process. 86, p. 1. Klauber, C., 2003, Surf. Inter. Anal., 35, p. 415. Klauber, C., Parker, A., Bronswijk, W.V., Watling, H., 2001, Int. J. Miner. Process., 62, p. 65. Koleini, S.M.J., Aghazadeh, V., Sandstr\u00F6m, \u00C5., 2011, Minerals Engineering, 24, p. 381. Lazaro, I., Nicol, M.J., 2006, J. Appl. Electrochem. 36, p. 425. Lefevre, G., Walcarius, A., Ehrhardt, J.J., Bessiere, J., 2000, Langmuir, 16, p. 4519. Lehmann, M.N., Kaur, P., Pennifold, R.M., Dunn, J.G., 2000, Hydrometallurgy 55, p. 255. Li, J., Kawashima, N., Kaplun, K., Absolon, V.J., Gerson, A.R., 2010, Geochimi. et Cosmochimi. Acta, 74, p. 2881. Li, L.C., Lan, X.J., Yi, Y., Yuan, N.Z., Juan, Z.X., Lei, Z., Yan, M.C., Dong, Z.Y., 2011, Hydrometallurgy, 107, p. 13. Linge, H.G., 1976, Hydrometallurgy 2, p. 51. Liu, W., Flytzani-Stephanopoulos, M., 1995, J. Catal. 153, p. 317. Lu, Z.Y., Jeffrey, M.I., Lawson, F., 2000, Hydrometallurgy 56, p. 145. Macdonald, D.D., Ismail, K.M., Sikora, E., 1998, J. Electrochem. Soc., 145, p. 3141. Majuste, D., Ciminelli, V.S.T., Osseo-Asare, K., Dantas, M.S.S., Magalh\u00E3es-Paniago, R., 2011, Hydromettalurgy, doi:10.1016/j.hydromet.2011.11.003. Makrides, A.C., 1964, J. Electrochem. Soc., 111, p. 392.; Makrides, A.C., 1964, J. Electrochem. Soc., 111, p. 400. Manocha, A.S., Park, R.L., 1977, Appl. Surf. Sci. 1, p. 129. Mansfeld, F., Kendig, M.W., 1988, J. Electrochem. Soc. 135, p. 828. Marcus, P., 2002, Corrosion Mechanisms in Theory and Practice, 2nd ed., Chap. 1, Marcel Dekker, New York. Marcus, P., Mansfeld, F. (Eds.), 2006, Analytical methods in corrosion science and engineering, Taylor &amp; Francis Group, CRC Press, Boca Raton, FL, pp. 1-102. Marcus, R.A., 1964, Ann. Rev. Phys. Chem., 15, p. 155. McCarron, J.J., Walker, G.W., Buckley, A.N., 1990, Intr. J. Mineral Process., 30, p. 1. McMillan, R.S., MacKinnon, D.J., Dutrizac, J.E., 1982, J. Appl. Electrochem. 12, p. 743. Memming, R., 2001, Semiconductor electrochemistry, Wiley-VCH, Weinheim, chapters 1-7. Mielczarski, J.A., Cases, J.M., Alnot, M., Ehrhardt, J.J., 1996, Langmuir, 12, p. 2519. Mikhlin, Y., 2000, Phys. Chem. Chem. Phys. 2, p. 5672.  202  \u000CMikhlin, Y.L., Tomashevich, Y.V., Asanov, I.P., Okotrub, A.V., Varnek, V.A., Vyalikh, D.V., 2004, Appl. Surf. Sci. 225, p. 395. Mishra, K.K., Osseo-Asare, K., 1988, J. Electrochem. Soc. 135, p. 1898. Mishra, K.K., Osseo-Asare, K., 1992, J. Electrochem. Soc. 139, p. 749. Morrison, S.R., 1980, Electrochemistry at semiconductor and oxidized metals electrodes, Plenum Press, NY, chapters 1-6. Mulder, W.H., Sluyters, J.H., Pajkossy, T., Nyikos, L., 1990, J. Electroanal. Chem. 285, p. 103. Munoz, J.A., Gomez, C., Ballestrer, A., Blazquez, M.L., Gonzalez, F., Figueroa, M., 1998, J. Appl. Electrochem. 28, p. 49. Munoz, P.B., Miller, J.D., Wadsworth, J.D., 1979, Metall. Trans. B 10B, p. 149. Nakai, I., Sugitani, Y., Nagashima K., Niwa, Y., 1978, J. Inorg. Nucl. Chem. 40, p. 789. Nava, D., Gonzalez, I., 2006, Electrochim. Acta 51, p. 5295. Nava, D., Gonzalez, I., Leinen, D., Ramos-Barrado, J.R., 2008, Electrochim. Acta 53, p.4889. Nazari, G., Dixon, D.G., Dreisinger, D.B., 2011, Hydrometallurgy 105, p. 251. Nazari, G., Dixon, D.G., Dreisinger, D.B., 2012, Hydrometallurgy doi:10.1016/j.hydromet.2011.12.011. Nesbitt, H.W., Muir, I.J., 1994, Geochim. et Cosmochim. Acta 58, p. 4667. Orazem, M.E., Tribollet, B., 2008, Electrochemical Impedance Spectroscopy, John Wiley &amp; Sons, NJ, p. 233. Otte, K., Lippold, G., Hirsch, D., Schindler, A., Bigl, F., 2000, Thin Solid Films, 498, p. 361. Panzuner, G., Egert, B., 1984, Surf. Sci. 144, p. 651. Parker, A.J., Paul, R.L., Power, G.P., 1981, J. Electroanal. Chem. 118, p. 305. Parker, A., Klauber, C., Kougianos, A., Watling, H.R., Bronswijk, W., 2003, Hydrometallurgy 71, p. 265. Paul, R.L., Nicol, M.J., Dcggle, J.W., Saunders, A.P., 1978, Electrochimica Acta, 23, p. 625. Plieth, W., 2008, Electrochemistry for Materials Science, Elsevier, Oxford, UK, p. 263. Pratesi G., Cipriani, C., 2000, Eur. J. Mineral., 12, p. 397. Pratt, A.R., Muir, I.J., Nesbitt, H.W., 1994, Geochim. et Cosmoch. Acta, 58, p. 827. Price D.W., Warren, G.W., 1986, Hydrometallurgy 15, p. 303. Robie, R.A., Bethke, P.M., Beardsley, K.M., 1967, U.S. geological Survey Bulletin 1248, p. 8. Rockett, A., 2008, The Materials Science of Semiconductors, Springer, NY, p. 21. 203  \u000CRosenholtz, J.L., Smith, D.T., 1936, J. American Mineralogist 21, p. 115. Pourbaix, M., 1966, \u201CAtlas of Electrochemical Equilibrium Aqueous Solutions,\u201D Pergamon Press, Ltd., Oxford. Salsman, J.B., Williamson, R.L., Tolley, W.K., Rice, D.A., 1996, Minerals Eng. 9, p. 43. Samec, Z., Weber, J., 1977, J. Electroanal. Chem., 77, p. 163. Schmuki, P., Bohni, H., Bardwell, J.A., 1995, J. Electrochem. Soc. 142, p. 1705. Schon, G., 1972, J. Electron Spectrosc., 73, p. 377. Shirley, D.A., 1972, Physical Review B, 5, p. 4709. Shreir, L.L., Jarman, R.A., Burstein, G.T., 2000, Corrosion, 3rd Rev., Metal/Invironment Reactions, Vol. 1, Butterworth-Heinemann, Oxford, p. 143. Shuey, R.T., 1975, Semiconducting ore minerals, Elsevier Scientific Publishers Co., NY, p. 242. Shuxian, Z., Hall, W.K., Ertl, G., Konzinger, H., 1986, J. Catalysis 100, p. 167. Stansbury, E.E., Buchana, R.A., 2000, Fundamentals of Electrochemical Corrosion, ASM International, Materials Park, OH, p. 260. Stern, M., 1957, J. Electrochem. Soc., 104, p. 559. Subramanian, K.N., Jennings, P.H., 1972, Can. Metall. Q. 11, p. 387. Susac, D., Zhu, L., Teo, M., Sode, A., Wong, K.C., Wong, P.C., Parsons, R.R., Bizzotto, D., Mitchell, K.A.R., Campbell, S.A., 2007, J. Phys. Chem. C, 111, p. 18715. Tan, B.J., Klabunde, K.J., Sherwood, P.M.A., 1990, Chem. Mater. 2, p. 186. Tell, B., Shay, J.L., kasper, H.M., 1972, J. Appl. Phys. 43, p. 2469. Teranishi, T., 1961, J. Phys. Soc. Jap. 16, p. 1881. Teranishi, T., Sato, K., Kondo, K., 1974, J. Phys. Soc. Jap. 36, p. 1618. Torma, A.E., 1991, Electrochemistry of a semiconductor chalcopyrite concentrate leaching by Thiobacillus ferrooxidans. In: Cooper, W.C., Kemp, D.O., Lagos, G.E., Tan, K.G. (Eds.), Hydrometallurgy and Electrometallurgy of Copper. Pergamon Press, NewYork, p. 73. Tshilombo, A.F., Petersen, J., Dixon, D.G., 2002, Minerals Eng., 15, p. 809. Vanmaekelbergh, D., Ern\u00E9, B.H., 1999, J. Electrochem. Soc. 146, p. 2488. Vanmaekelbergh, D., Searson, P.C., 1994, J. Electrochem. Soc. 141, p. 697. Vaughan, D.J., Becker, U., Wright, K., 1997, Int. J. Miner. Process. 51, p. 1. Vasquez, B.R., Viramontes-Gamboa, G., Dixon D.G., 2012, J. Electrochem. Soc. 159, p. C8. Vel\u00E1squez, P., G\u00F3mez, H., Leinen, D. Ramos-Barrado, J.R., 1998, Colloids and Surfaces A 140, p. 177.  204  \u000CVel\u00E1squez, P., G\u00F3mez, H., Ramos-Barrado, J.R., Leinen, D., 1998, Colloids and Surfaces A 140, p. 369. Vel\u00E1squez, P., Leinen, D., Pascual, J., Ramos-Barrado, J.R., Grez, P., G\u00F3mez, H., Schrebler, R., R\u00EDo, R.D., C\u00F3rdova, R., 2005, J. Phys. Chem. B 109, p. 4977. Vinkevicius, J., Mozginskiene, I., Jasulaitiene, V., 1998, J. Electroanal. Chem., 442, p. 73. Viramontes-Gamboa, G., Rivera-Vasquez, B.F., Dixon, D.G., 2006, ECS Transactions 2, p. 165. Viramontes-Gamboa, G., Rivera-Vasquez, B.F., Dixon, D.G., 2007, J. Electrochem. Soc. 154, p. C299. Wagner, C.D., Riggs, W.M., Davis, L.E., Moulder, J.F., Muilenberg, G.E., 1979, PerkinElmer Corporation/Physical Electronics Division, Eden Prairie, Minnesota, pp. 76-77. Wang, H., 2008, Mineral Processing &amp; Extractive Metall. Rev. 29, p. 1. Wang, S., 2005, JOM 57, p. 48. Warren, G.W., Wadsworth, M.E., El-Raghy, S.M., 1982, Metall. Trans. B 13B, p. 571. Watling, H.R., 2006, Hydrometallurgy 84, p. 81. Wei, D., Osseo-Asare, K., 1997, J. Electrochem. Soc. 144, p. 546. Wei, D., Osseo-Asare, K., 1996, J. Electrochem. Soc., 143, p. 3192. Winkel, L., Alxneit, I., Sturzenegger, M., 2007, Minerals Eng. 20, p. 1179. Yacobi, B.G., 2004, Semiconductor materials, an introduction to basic principles, Kluwer Academic Publishers, NY, chapters 1-4. Yin, Q., Kelsall, G.H., Vaughan, D.J., England, K.E.R., 1995, Geochim. et Cosmochim. Acta 59, p. 1091. Yin, Q., Vaughan, D.J., England, K.E.R., Kelsall, G.H., Brandon, N.P., 2000, J. Electrochem. Soc. 147, p. 2945. Yin, M., Wu,C.K., Lou, Y., Burda, C., Koberstein, J.T., Zhu, Y., O\u2019Brien, S., 2005, J. Am. Chem. Soc., 127, p. 9506. Yu, X.R., Liu, F., Wang, Z.Y., Chen, Y., 1990, J. Electron Spectroscopy and Related Phenomena, 50, p. 159.  205  "@en
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