Open Collections

UBC Theses and Dissertations

UBC Theses Logo

UBC Theses and Dissertations

Fundamental study of the controlled-potential leaching of chalcopyrite Lai, Jerry Cheng-Yen 2003

Your browser doesn't seem to have a PDF viewer, please download the PDF to view this item.

Item Metadata

Download

Media
831-ubc_2003-0429.pdf [ 8.9MB ]
Metadata
JSON: 831-1.0078711.json
JSON-LD: 831-1.0078711-ld.json
RDF/XML (Pretty): 831-1.0078711-rdf.xml
RDF/JSON: 831-1.0078711-rdf.json
Turtle: 831-1.0078711-turtle.txt
N-Triples: 831-1.0078711-rdf-ntriples.txt
Original Record: 831-1.0078711-source.json
Full Text
831-1.0078711-fulltext.txt
Citation
831-1.0078711.ris

Full Text

FUNDAMENTAL STUDY OF THE CONTROLLED-POTENTIAL LEACHING OF CHALCOPYRITE by JERRY CHENG-YEN LAI B.Sc. (Hon.), Chemistry, The University of British Columbia, Vancouver, Canada, 1997 A THESIS SUBMITTED IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF SCIENCE in THE FACULTY OF GRADUATE STUDIES (Department of METALS AND MATERIALS ENGINEERING) We accept this thesis as conforming to the required standard THE UNIVERSITY OF BRITISH COLUMBIA July 2003 Copyright ® Jerry Cheng-Yen Lai, 2003 In presenting this thesis in partial fulfillment of the requirements of an advanced degree at the University of British Columbia, I agree that the Library shall make it freely available for reference and study. I further agree that permission for extensive copying of this thesis for scholarly purposes may be granted by the head of my department or by his or her representatives. It is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission. Department of Metals and Materials Engineering The University of British Columbia Vancouver, B.C., Canada Date A B S T R A C T ABSTRACT There is a well-established trend in copper hydrometallurgy to develop processes for the treatment of chalcopyrite ( CuFeS2 ). Chalcopyrite receives considerable attention not only because of its relative abundance and widespread distribution in almost all sulfide deposits, but also because it is the most recalcitrant copper sulfide to leaching in most practical systems, particularly in sulfate media. The main practical difficulty encountered in the commercial acid ferric sulfate leaching of chalcopyrite (in heaps, dumps or concentrates) is its slow rate of reaction, which renders a long residence time and/or incomplete copper extractions. Chalcopyrite is generally agreed to be "passivated" in some fashion, exhibiting slow leaching kinetics and low copper extraction in sulfate media. This thesis work investigated the oxidation of chalcopyrite through the application of controlled-potential at different potentials (0.400-0.600 V A g / A g C i ) and temperatures (60-78°C). A l l experiments were carried out in a compact bench-scale 3-L glass jacketed bioreactor equipped with a servo motor and speed controller. Since the solution redox potential is dependent on the concentrations of ferric and ferrous ions in solution, a constant redox potential was maintained by controlled addition of potassium permanganate using an automatic titration unit. Chalcopyrite leaching in sulfuric acid is shown to be dependent on solution redox potential determined by the concentration ratio of ferric to ferrous ions and temperature. The most significant finding in this thesis work is the fact that the leaching; rate increases with increasing potentials but decreases above a critical potential under the reaction conditions studied. This critical potential is in the vicinity of 0.500 V A g /Agc i and is within the mid-potential range of 0.45 to 0.55 V A g / A g C i at which the leaching rate is maximal from 60 to 78°C. Leaching rates increase with increasing temperature. The leach residues contain sulfate (5-33%), but primarily elemental sulfur (46-78%). XRD analysis confirmed the presence of sulfate sulfur, jarosite, and elemental sulfur in the leach residues. SEM analysis suggests that neither jarosite nor elemental sulfur is responsible for surface passivation or for preventing, the dissolution of chalcopyrite. TABLE OF CONTENTS / TABLE OF CONTENTS ABSTRACT ii TABLE OF CONTENTS ..iii LIST OF TABLES vi LIST OF FIGURES ix LIST OF SYMBOLS xii LIST OF ABBREVIATIONS AND ACRONYMS xv ACKNOWLEDGEMENTS xvi CHAPTER 1 INTRODUCTION 1 1.1 Outline of the Thesis 3 CHAPTER 2 LITERATURE REVIEW 4 2.1 Properties of Chalcopyrite 4 2.2 Theoretical Considerations of Ferric Sulfate Leaching 7 2.2.1 Thermodynamics of Ferric Sulfate Leaching of Chalcopyrite 7 2.2.2 Importance of Iron Precipitation and Jarosite Formation 10 2.2.3 Speciation Distribution oflron in Ferric Sulfate Solution 15 2.3 An Introduction to the Ferric Sulfate Leaching of Chalcopyrite 17 2.3.1 Chemical (Ferric Sulfate) Leaching of Chalcopyrite 18 2.4 An Introduction to Bacterial Leaching 23 2.5 The Nature of the Passivating Layer...-:..:. 26 2.5.1 Iron Compounds and Jarosite Passivating Layer 26 2.5.2 Elemental Sulfur and Bimetallic Intermediate Sulfide Passivating Layer 27 2.5.3 Silver Catalyzed Ferric Sulfate (Bacterial) Leaching Of Chalcopyrite 29 2.5.4 Electrochemical Studies of Chalcopyrite 32 2.5.5 Surface Studies (AES and XPS) 38 2.5.5.1 Air Oxidation of Chalcopyrite 38 2.5.5.2 Oxidation in Alkaline Solution (Studies Related to Flotation) 39 2.5.5.3 Oxidation in Acid Solutions: Studies Related to (Bio)Leaching 41 2.5.5.4 Passivating Layers and the Limitation of Experimental Techniques 43 2.6 Electrochemistry , 47 2.6. J Mixed-Potential Theory 49 2.6.2 Redox-Controlled Potential Leaching Studies 52 T A B L E OF CONTENTS / 2.6.2.1 Redox-Controlled Leaching of Chalcopyrite (Kametani and Aoki) 55 2.6.2.2 Redox-Controlled Chalcopyrite Bioleaching Patent of Mintek 58 2.7 Objective of the Research 62 CHAPTER 3 EXPERIMENTAL PROCEDURES 63 3.1 Materials 63 3.1.1 Chalcopyrite Samples 63 3.1.2 Reagents 67 3.2 General Experimental Techniques and Instrumentation 68 3.2.1 Leaching Apparatus 68 3.2.2 Leaching Procedures 72 3.2.3 Quantitative Analysis 76 3.2.4 Qualitative Analysis 78 3.2.4.1 Powder X-ray Diffraction (XRD) 78 3.2.4.2 Scanning Electron Microscopy (SEM) 78 3.3 Experimental Considerations and Computations 7° 3.3.1 Permanganate Oxidation of Ferrous Iron in Acidic Solution 79 3.3.2 Determination of Percent Extraction and Sulfur Species 83 3.3.3 Statistical Analysis of Experimental Results 84 CHAPTER 4 RESULTS AND DISCUSSION 86 4.1 Reproducibility and Consistency of Leach Tests 86 4.2 Acid Soluble Copper and the Effect of Initial Ferrous Iron Concentration 88 4.3 Analysis of Leach Residues 89 4.3.1. Analysis of Sulfur Species in the Leach Residues :'. 89 4.3.2. Morphology of Leach Residues 96 4.3.3. Copper Extraction and Other Experimental Results 100 4.4 Leaching Kinetics: Interpretation of Batch Reactor Data 103 4.4.1 Effect of Solution Redox Potentials and Temperature on the Dissolution Rate of Copper 103 4.4.2 Effect of Solution Redox Potentials and Temperature on the Dissolution Rate of Iron 114 TABLE OF CONTENTS/ v CHAPTER 5 CONCLUSIONS 125 CHAPTER 6 RECOMMENDATIONS FOR FUTURE WORK 126 REFERENCES 127 APPENDIX A SAMPLE CALCULATIONS 144 A-l Calculation of Solid Pulp Density (SPD) 145 A-2 Calculation of Copper Extraction Based on Calculated Head 146 A-3 Calculation of Iron Extraction Based on Calculated Head 147 A-4 Calculation of Total Ferrous Ions Concentration [Fe(Il)T ] During Leaching... 148 A-5 Determination of Sulfide Sulfur Conversion [ S^jdaljon - S2$~ - (S04~ )F ] 149 A-6 Calculation of Copper Extraction Based on the Consumption of Potassium Permanganate (KMn04) Solution 151 A-7 Calculation of Evaporation Rate During Leaching 153 A-8 Statistical Significance of Temperature on Experimental Results 154 A-9 Statistical Significance of Solution Redox Potential on Experimental Results 156 A-10 Statistical Analysis of Experimental Results 158 A-l 1 Reference Electrode Potential Correction on a Standard Hydrogen Electrode (SHE) Scale : 160 A-12 The Criss-Cobble Extrapolation Method and E"r Values for the Fei+n*Couple 162 APPENDIX B SELECTED LEACH DATA 164 APPENDIX C ANALYTICAL TECHNIQUES 173 C-l Titrimetric Determination of Ferrous Iron Against Potassium Permanganate (KMnOJ Solution 174 C-2 Standardization Procedure for 0.2 M (or 1 N) KMn04 176 C-3 Method of Multi-Acid ICP Multi-element Analysis 178 C-4 Method of Iron Assay by Titration Analysis 179 C-5 Method of Copper Assay by Titration Analysis 180 LIST OF TABLES LIST OF T A B L E S / Table 2-1 Crystal structure of chalcopyrite (CuFeS2) 6 Table 2-2 Physical and thermodynamic properties of chalcopyrite (CuFeS 2) at 25°C 6 Table 2-3 Chemical and mineral names of common jarosites34 11 Table 2-4 Analytical data for sulfur species in the ferric sulfate oxidation of chalcopyrite 18 Table 2-5 Copper extraction in the ferric sulfate oxidation of chalcopyrite 22 Table 2-6 Summary of the influence of process conditions on the ferric sulfate oxidation kinetics of chalcopyrite 22 Table 2-7 Stoichiometric equations for the direct and the indirect bacterial oxidation of chalcopyrite8'77 24 Table 2-8 Summary of the analysis of initial stages of CuFeS2 dissolution of the passive region (E < 1 VSHE ) 46 Table 2-9 Summary of electronic properties of selected sulfide minerals14 48 Table 2-10 Standard electrical rest potential of selected base metal sulfides 48 Table 2-11 Standard potentials of iron coupling and chalcopyrite half reactions at 25 and 90°C 5 3 51 Table 3-1 Chemical analyses of the floated CuFeS2 concentrate ". 64 Table 3-2 Chemical and mineralogical composition of copper concentrate 64 Table 3-3 Standard operating conditions for all leaching experiments 74 Table 4-1 Summary of XRD results of leach residues from data set #2 92 Table 4-2 Statistical significance of temperature dependency on total sulfide ' oxidations relative to that of 60 and 70°C at various potentials (0.400-0.600 V A g / A g C l ) 92 Table 4-3 Statistical significance of redox potential dependency on total sulfide oxidations relative to that of 0.400 V A g / A g C i at various temperatures (60-78°C) 92 Table 4-4 Statistical significance of temperature dependency on total elemental sulfur formations relative to that of 60 and 70°C at various potentials (0.400-0.600 V A g / A g C l ) 93 Table 4-5 Statistical significance of redox potential dependency on total elemental sulfur formations relative to that of 0.400 V A g / A g C i at various temperatures (60-78°C) 93 LIST OF T A B L E S / vii Table 4-6 Statistical significance of temperature dependency on total sulfate ! formations relative to that of 60 and 70°C at various potentials (0.400-0.600 VAg/AgCl) 93 Table 4-7 Statistical significance of redox potential dependency on total sulfate formations relative to that of 0.400 VAg/Agci at various temperatures (60-78°C) 93 Table 4-8 Statistical significance of temperature dependency on total copper extractions relative to that of 60 and 70°C at various potentials (0.400-0.600 VAg/AgCl) 106 Table 4-9 Statistical significance of redox potential dependency on total copper extractions relative to that of 0.400 VAg/Agci at various temperatures (60-78°C) 106 Table 4-10 Comparison of experimental results with reported critical potentials for chalcopyrite leaching 106 Table 4-11 Copper extractions as a function of potential (0.400-0.600 VAg/Agci) and temperature (60-78°C) 107 Table 4-12 Major iron extraction peaks and iron extractions as a function leaching time ; under potentials of 0.400 to 0.600 VAg/Agci at temperatures from 60 to 78°C j for data set #2 115 i Table 4-13 Statistical significance of temperature dependency on total iron extractions relative to that of 60 and 70°C at various potentials 117 Table 4-14 Statistical significance of redox potential dependency on total iron extractions relative to that of 0.400 VAg/AgCi at various temperatures (60-78°C) 118 Table 4-15 Calculated ferric to ferrous mole ratio at 60, 70 and 78°C using the E° calculated from the Criss and Cobble theory 118 Table 4-16 Iron extractions as a function of potential (0.400-0.600 VAg/Agci) and temperature (60-78°C) 118 dE Table A - l Standard electrode potential correction values [ E°H and — ] of some dT common reference electrodes systems as a function of temperature 161 Table A-2 Standard electrode potential versus SHE ( E" ) in volts (VSHE) for the Ag/ AgCl electrode system with saturated KCl electrolyte as a function of potential (0.400-0.600 VAg/Agci) and temperature (60-78°C) 161 Table A-3 Criss-Cobble heat capacity constants (ar,f3r) for simple cations at 60, 70 and 78°C 163 Table A-4 G°298, S°29% and z values for ferrous and ferric ionic species at 25°C(298K)... 163 LIST OF T A B L E S / viii Table A-5 C°P] , (G°.), AGj and E° ( V s H E , V A g / A g C i ) values for ferrous and ferric ions at 60, 70 and 78°C 163 Table B - l Summary of calculated head grade and extractions for iron and copper 165 Table B-2 Summary of sulfide sulfur conversion [ S0~Mation - S2~„ - {SO]~)F ] 166 Table B-3 Summary of copper extractions (Cu^'"^°") based on the consumption of potassium permanganate (KMn04) solution 167 Table B-4 Statistical analysis of copper extractions based on calculated head tLj 168 Table B-5 Statistical analysis of iron extractions based on calculated head , r? Extraction x j g Q Table B-6 Statistical analysis of total sulfide sulfur oxidation (S20xklatjoll) 170 Table B-7 Statistical analysis of total elemental sulfur formation (S2~,) 171 Table B-8 Statistical analysis of total sulfate formation [(S04~)F] 172 LIST OF FIGURES LIST OF FIGURES / i Figure 2-1 a) Unit-cell model showing the crystal structure of chalcopyrite (CuFeS2)u b) Interatomic bond lengths and angles of each atom in chalcopyrite 5 Figure 2-2 The metastable EH - pH diagram for the Cu - Fe - S - H20 system at 25°C (without FeS2 as a product); all solutes at 1 mol/L activity.2 9 9 Figure 2-3 Stability fields of various iron precipitates as a function of pH and temperature [0.5 mol/L F<? 2(S0 4) 3] 3 5 12 Figure 2-4 System Fe2(SOA)i - H2S04 - H20 a) Polytherm 50 to 200°C, 0 to 40 wt%.S04 3 7 14 Figure 2-5 System Fe2{SOA)3 - H2SOA - H20 b) Polytherm 50 to 200°C, 30 to 70 w t % S 0 4 3 7 14 Figure 2-6 Effect of pH on species distribution of Fe(Il) and Fe(lll) in sulfate solution4 5 (25°C, [ S 0 4 ~ ] w = 1 mol/L, Fe(Il) =Fe(JII) = 0.1 mol/L) 16 Figure 2-7 Leaching rate curve of chalcopyrite with ferric sulfate61 [70°C(343K), [SOA']To,nl =0.2 mol/L, Fe(III) = 1.0 mol/L)] 21 Figure 2-8 Effect of temperature on the polarization curve of chalcopyrite (0.1 MH2SOA, scan rate 300 mV/min) 1 3 1 34 Figure 2-9 Anodic polarization curves for CuFeS2 from six different locations (1 MH2S04, scan rate 30mV/min, temperature 25°C) 8 0 36 Figure 2-10 Current density versus time plot for Transvaal CuFeS2 at various anodic ; potentials (1 MH2SOA, temperature 25°C) 8 0 36 Figure 2-11 Current density potential diagram showing curves for metal sulfide ( M S ) with equilibrium (or rest) potential ( E c i ) and redox-ions { N ( n + 1 ) / n + } with equilibrium (or rest) potential (E C 2), which forms a mixed potential ( E m ) 51 Figure 2-12 Reaction curves (KMnOA added vs. log t) for the oxidation of chalcopyrite i in the potential range 0.30-0.43 V S C E (or 0.600-0.630 V S H E ) (90°C, initial ' iron cone. = 0.1 mol /L) 1 6 3 57 Figure 2-13 Reaction curves (KMnOA added vs. log t) for the oxidation of chalcopyrite in the potential range 0.47-0.65 V S C E (or 0.670-0.850 V S H E ) (90°C, initial iron cone. = 0.1 mol/L). Dotted Curve: reaction curve at 0.43 V S C E (or 0.630 VSHE) in Figure 2-12 for comparison. ' 57 Figure 2-14 Ferric oxidation of fine-milled South-African chalcopyrite at 35°C (initial ; iron cone. = 3 g/L, pH = 1.5)168 60 LIST OF FIGURES / x Figure 2-15 Ferric oxidation of fine-milled South-African chalcopyrite at 55°C (initial iron cone. = 3 g/L, pH = 1.5)'68 60 Figure 2-16 Ferric oxidation of fine-milled South-African chalcopyrite at 80°C (initial iron cone. = 3 g/L, pH = 1.5)168 6\ Figure 2-17 Ferric oxidation of a North American chalcopyrite concentrate at 70°C 1 6 8 61 Figure 3-1 Cumulative particle size distribution of the floated chalcopyrite concentrate as determined by the Elzone 280PC counter 65 Figure 3-2 XRD pattern of the floated chalcopyrite concentrate.! 66 Figure 3-3 Schematic diagram of the controlled-potential leaching system in batch mode 70 Figure 3-4 Latimar diagram for manganese species 80 Figure 3-5 EH - pH diagram for the Mn - S04~ - H20 system at 25°C, 0.1 dissolved manganese and unity sulfate activities 81 Figure 4-1 Comparison of total copper and iron extractions as a function of potential (0.400-0.600 VAg/Agci) and temperature from different data sets (set #1, set #2, set #3): (a) 60°C; (b) 70°C; (c) 78°C 87 Figure 4-2 Total sulfide oxidation ( Solillali0ll ), total elemental sulfur formation (S2~ ) and total elemental sulfate formation [(S0 4 2~) F] as a function of potential (0.400-0.600 VAg/Agci) and temperature (60-78°C) 94 Figure 4-3 XRD pattern of the leached residue performed at 60°C under a potential of 0.500 VA G/Agci 95 Figure 4-4 Backscattered electron micrograph for chalcopyrite leached at 78°C under a potential of 0.450 V A g/ A gci after 20 hours leaching 98 Figure 4-5 Backscattered electron micrograph for chalcopyrite leached at 78°C under a potential of 0.500 VA G/Agci after 17 hours leaching 98 Figure 4-6 Backscattered electron micrograph for chalcopyrite leached at 78°C under a potential of 0.550 V A g/ A gci after 7 hours leaching 99 Figure 4-7 Backscattered electron micrograph for chalcopyrite leached at 78°C under a potential of 0.600 VA G/Agci after 9 hours leaching 99 Figure 4-8 Comparison of different reaction products of chalcopyrite dissolution as a function of potential (0.400-0.600 VAg/AgCi) and temperature: (a) 60°C; (b) , 70°C; (c) 78°C 102 Figure 4-9 Extraction-time plot for copper as a function of potential (0.400-0.600 VAg/Agci) at 60°C in ferric sulfate medium 108 Figure 4-10 Extraction-time plot for copper as a function of potential (0.400-0.600 V~Ag/AgCi) at 70°C in ferric sulfate medium 109 LIST OF FIGURES / Figure 4-11 Extraction-time plot for copper as a function of potential (0.400-0.600 V A g / A g c i ) at 78°C in ferric sulfate medium 110 Figure 4-12 Extraction-time plot for copper at 0.400 V A g / A g c i and different temperatures (60-78°C) in ferric sulfate medium 111 Figure 4-13 Extraction-time plot for copper at 0.450 V A g / A g c i and different temperatures (60-78°C) in ferric sulfate medium 111 Figure 4-14 Extraction-time plot for copper at 0.500 V A g / A g c i ) and different temperatures (60-78°C) in ferric sulfate medium 112 Figure 4-15 Extraction-time plot for copper at 0.550 V A g / A g c i and different temperatures (60-78°C) in ferric sulfate medium 112 Figure 4-16 Extraction-time plot for copper at 0.600 V A g / A g c i ) and different temperatures (60-78°C) in ferric sulfate medium 113 Figure 4-17 Extraction-time plot for iron as a function of potential (0.400-0.600 V A g / A g c i ) at 60°C in ferric sulfate medium 119 Figure 4-18 Extraction-time plot for iron as a function of potential (0.400-0.600 V A g / A g c i ) at 70°C in ferric sulfate medium 120 Figure 4-19 Extraction-time plot for iron as a function of potential (0.400-0.600 V A g / A g c i ) at 78°C in ferric sulfate medium 121 Figure 4-20 Extraction-time plot for iron at 0.400 V A g / A g c i and different temperatures (60-78°C) in ferric sulfate medium 122 Figure 4-21 Extraction-time plot for iron at 0.450 V A g / A g c i and different temperatures. (60-78°C) in ferric sulfate medium 122 Figure 4-22 Extraction-time plot for iron at 0.500 V A g / A g c i and different temperatures (60-78°C) in ferric sulfate medium 123 Figure 4-23 Extraction-time plot for iron at 0.550 V A g / A g c i and different temperatures '! (60-78°C) in ferric sulfate medium 123 Figure 4-24 Extraction-time plot for iron at 0.600 V A g / A g c i and different temperatures : (60-78°C) in ferric sulfate medium 124 LIST OF SYMBOLS / xii LIST OF SYMBOLS Symbol a i Description Charge transfer coefficient Criss Cobble heat capacity constants Current density Exchange current density Electron stoichiometric coefficient Rate for the forward reaction Unit k p 1 80 SPD (Cu2^ S° JM (Cu2^ 2+ Fe (Fe2^ S" JM F R AG° AG 2° 9 S J 2 9 8 C° 298 Rate for the reverse reaction Mesh or grind size that results in 80% passing Solid Pulp Density Copper to sulfur molar ratio based on calculated copper head assay Copper to iron molar ratio based on calculated head assay Iron to sulfur molar ratio based on calculated iron head assay Faraday's constant (96485) Universal gas constant (8.3144) Standard Gibbs energy of reaction at temperature (7) in degree Kelvin (K) Change in Gibbs energy for a reaction at temperature (7) Standard free energy of reaction at 25°C(298K) Standard entropy Standard entropy change of reaction at 25°C(298K) Average heat capacity change at constant pressure for the reaction between 25°C(298K) and temperature (7) in degree Kelvin (K) (C/mol) j (J-mof'lC'1) (kJ/mol) , (kJ/mol) (kJ/mol) (J/mol-K) (J/mol-K) (J/mol/K) LIST OF SYMBOLS / EH E E° la Cu, CalcHead Cu Cu, AvgCalcHead Extraction CalcHead s~* Extraction L'"AvgCalcHead Cu Extraction KMnO, s~* Extraction s rp rp \ ^ U AvgCalcHead V 1 \ 1 2> ^ E x t r a c t i o n (p . C \ ^ u A v g C a l c H e a d ^ ^2 > SOxidation (^ 1 ~ * ^ 2 ) SOxidation (^1 — * ^2 ) Sp(T}^T2) (5042-)[r,^r2] Mixed potential of the system (19 Oxidation-reduction potential with respect to SHE (V) Solution redox potential (V) Standard reduction potential (V) Standard reduction potential at any temperature (J) (V) Anodic overpotential (V) Cathodic overpotential (V) Calculated head assay for copper (%) Average calculated head assay for copper (%) Copper extraction based on calculated head assay (%) Copper extraction based on average calculated head „ f /fry assay Copper extraction based on potassium permanganate , 0 / . solution consumption Effect of temperature on copper extraction between a higher temperature (T2) and lower temperature (Jj) Effect of solution redox potential on copper extraction between a higher redox potential (E2 ) and (%), a lower redox potential ( £ , ) Effect of temperature on total sulfide oxidation between a higher temperature ( T2 ) and lower (%} temperature (Jj) Effect of solution redox potential on total sulfide oxidation between a higher redox potential (E2 ) and (%) a lower redox potential (Ex) Effect of temperature on total elemental sulfur formation between a higher temperature (T2) and (%) lower temperature (Jj) Effect of solution redox potential on total elemental sulfur formation between a higher redox potential (%) (E2) and a lower redox potential (E]) Effect of temperature on total sulfate formation between a higher temperature ( T2 ) and lower (%) temperature (Jj) LIST OF S Y M B O L S / xiv (S024~)[E} ->£2] Cu Extraction AvgCalcHead-KMnO. Fe Fe, Fe, Fe AvgCalcHead Cole Head Ext rod ion CalcHead Extraction AvgCalcHead p Extraction / p . p \ 1 KAvgCalcHead y1 ] ^ 12> p Extraction (p p \ 1 ^AvgCalcHead ^ ^2) S I -O.xidation s2; (so24-)F [Fe2+] [Fe>+] [ F e T o , a l ] [F4"oT] Effect of solution redox potential on total sulfate formation between a higher redox potential (E 2) and a lower redox potential (Ex) Absolute difference between copper extraction based on average calculated head assay and potassium permanganate solution consumption Average calculated head assay for iron Calculated head assay for iron Iron extraction based on calculated head assay Iron extraction based on average calculated head assay Effect of temperature on iron extraction between a higher temperature (T2) and lower temperature (7j) Effect of solution redox potential on iron extraction between a higher redox potential (E 2) and a lower redox potential ( ) Total sulfide sulfur oxidation of the head sample Total sulfide sulfur conversion to elemental sulfur Total sulfide sulfur conversion to sulfate Concentration of ferrous iron Concentration of ferric iron Total iron concentration Total soluble iron concentration (%) (g) •(g) (g) (g) LIST OF ABBREVIATIONS A N D A C R O N Y M S / LIST OF ABBREVIATIONS AND ACRONYMS Abbreviation Description AAS Atomic Absorption Spectrophotometry AES Auger Electron Spectroscopy Ag/AgCl Silver/Silver Chloride Electrode Reference ASTM American Society for Testing and Materials ICP Inductively-coupled Plasma Multi-element Analysis IPL International Plasma Laboratory (Vancouver, Canada) SCE Standard Calomel Electrode Reference SEM Scanning Electron Microscopy SHE Standard Hydrogen Electrode Reference jj-gQ ^ n e UmvCTsity of British Columbia (Vancouver, Canada) (http://www.ubc.ca) VWR® V W R International (a company of the Merck Group) (http://www.vwrcanlab.com) XPS X-ray Photoelectron Spectroscopy XRD Powder X-ray Diffraction Technique A C K N O W L E D G E M E N T S / xvi ACKNOWLEDGEMENTS I wish to express my sincere appreciation for the assistance provided by Dr. Ralph Peter Hackl and Dr. David G. Dixon of The University of British Columbia (UBC), who acted as my supervisors for the duration of this thesis work. Being a student just graduated from the Department of Chemistry at U B C [B.Sc. (Hon.)], the field of hydrometallurgy was a completely different experience beyond imagination. My supervisors revealed to me the significance of scientific investigation through the rearrangement of available data. Different equations and graphs on existing data could sometimes reveal insights not previously conceived, and inspired a different research direction. Throughout the preparation of this master thesis, their guidance, constructive criticism, patience and encouragement, as well as full financial support, contributed significantly to the successful completion of this study. i Sincere gratitude is extended to the members of the (bio)hydrometallurgy groups for their helpful comments, discussions and support. They have provided an enjoyable atmosphere in which working and socializing existed seamlessly. The technical assistance of my colleague1 Mr. Mike Timmins in the setup of the leaching reactor is greatly acknowledged. Special thanks to Mr. Mike Timmins (M.Sc), Mr. Masoud Aftaita and Ms. Sylvie Bouffard (M.A.Sc) for their assistance during the course of several sixty hour experiments, which rendered me the time to sleep, eat and relax, while they acquired experimental samples and data on my behalf. Moreover, 1 thank both Ms. Jane Howard and Mr. Charlie Joseph Loo for their assistance with solution assays on many experimental samples. The vending service of Mr. Glenn Smith is also greatly recognized. A l l solid assays performed externally by IPL (International Plasma Laboratory) are appreciated. Many people provided constructive criticism and technical assistance during the preparation of this thesis. I thank Ms. Sally Finora at the Department of Mining Engineering for determining the continuous size distribution of the chalcopyrite feed mineral. I thank Dr. L .A. Groat, Dr. Mati Raudsepp and Dr. Elisabetta Pani at the Department of Geology for the powder XRD studies of the chalcopyrite feed mineral. Moreover, I thank Dr. Elisabetta Pani for the powder XRD and SEM studies of the leached residual solids resulted from individual experiments. Furthermore, I should thank Dr. Berend Wassink of the group for his kind assistance and patience in developing my experimental techniques, particularly for the analyses A C K N O W L E D G E M E N T S / xvii of dissolved copper and soluble iron by AAS. He provided valuable training and suggestions necessary for the completion of this work. I thank Mr. Matthew Chang (J.D.) for proof reading and editing the draft copy of my thesis. I am very grateful to all of them for taking the time to do so. A l l the remaining errors and misconceptions in this thesis work are regretfully mine alone. During the years 2002-2003 when I was preparing my thesis in Vancouver (Canada), many people provided me with food, which meant a lot to a person such as me who lacked culinary skills. I would like to thank the Leung siblings for inviting me to many dinners with them. I would also like to mention Mr. Jerry Chang and his mother, Mrs. Chang, who treated me as a family member, and invited me to their house during holiday seasons (e.g., Chinese New Years and X'mas). Moreover, the numerous shopping trips to Safeway with Mr. Clark Chang and his delivery of breads and buns from Maxim's bakery were greatly cherished. Furthermore, I would like to thank Ms. Priscilla Leung, who kept me company throughout the year, which included numerous dinners and supermarket shopping throughout Burnaby and Richmond City. The study was made possible through the research and equipment grant provided by the Natural Sciences and Engineering Research Council of Canada (NSERC). Last, but not the least, my deepest gratitude belongs to my parents. This thesis work was temporary interrupted for the duration of around three years before its final submission. My father had lung cancer. At that time, my father's health was deteriorating, and his eventual defeat to death on July 29, 2001 brought the family enormous distress. Without my parents' determination to mediate between the traditional Chinese and the modern Occidental education, I could not have broadened my horizon and envisioned a world beyond the small world of Formosa (Taiwan). Speaking English would have become solely a dream and not a reality. I thank my parents for their unwavering love, encouragement and unlimited support through this time, which has at times been a difficult, but ultimately rewarding, experience. C H A P T E R 1 I N T R O D U C T I O N / 1 C H A P T E R 1 INTRODUCTION Conventional pyrometallurgical smelting of copper sulfide concentrates has been the most direct and economical means of metal recovery, but requires the melting of all materials. The copper smelter requires high capital and operational costs, and must be operated on a large scale in order to be profitable. It also releases environmentally unacceptable emissions (i.e., S02 and As2Oi) to the atmosphere. The situation is made worse by low and unstable copper prices in international markets, low grade ores and the progressive exhaustion of reserves. In recent years, many copper producers have moved away from conventional copper pyrometallurgy towards more benign, but time-consuming, copper hydrometallurgy, which is based on the dissolution of metals in aqueous solutions using appropriate l i x i v i a n t s . 1 ' 2 Hydrometallurgical processing schemes virtually eliminate air pollution, and can be economically operated on a much smaller scale than copper smelters, especially when the ore is too low in grade for concentration. Copper hydrometallurgical processes involve leaching, solvent extraction and electrowinning. Among the most promising lixiviants (chloride 3, ammine 4, nitrate 5, sulfate), sulfate is the cheapest and offers the least corrosive solution. Also, sulfuric acid has been the most successful lixiviant for solvent extraction and electrowinning (SX/EW) of copper. 6 Following purification by solvent extraction, the direct electrowinning of copper from sulfate-based strip liquor produces copper at significantly lower cost than most smelting processes. The quality of the electrowon copper is also fully comparable with electro-refined copper produced by smelting and refining. Since sulfuric acid is also regenerated during electrowinning of copper, an on-going interest in copper hydrometallurgy is to develop cost-effective sulfate-based processes for both environmental and operational reasons. There is a well-established trend in copper hydrometallurgy to develop processes for the treatment of chalcopyrite ( CuFeS2 ). Chalcopyrite receives considerable attention not only because of its relative abundance and widespread distribution in almost all sulfide deposits, but also because it is the most recalcitrant copper sulfide to leaching in most practical systems, particularly in sulfate media. The rapid and complete oxidation of chalcopyrite in sulfate media is possible only through oxygen pressure leaching at temperature exceeding 200°C.7 However, the complete oxidation of sulfide sulfur to sulfuric acid (a waste product) requires large oxygen CHAPTER 1 INTRODUCTION / 2 and limestone consumption, which renders this process economically uncompetitive in comparison with smelting. Alternatively, ferric sulfate leaching is a widely applied sulfate-based hydrometallurgical process, which operates at 20-100°C and low pH. 8 A further advantage of this process is the production of mainly elemental sulfur, rather than sulfuric acid, which avoids all S02 emissions and separates copper production from the ever-fluctuating sulfuric acid market. The ferric sulfate leaching is sometimes assisted by mesophilic bacteria in dump leaching, referred to as "bacterial leaching," and operated at 20-40°C.8 While these low temperature processes are technically and economically attractive, chalcopyrite is generally agreed to be "passivated" in some fashion, exhibiting slow leaching kinetics and low copper extraction in sulfate media.9 From an industrial point of view, ores containing chalcopyrite mineral must be treated by conventional concentration/smelting/refining operations, irrespective of cost. 1 0 Although many alternative processes have been tested, no commercial hydrometallurgical processes are available for the recovery of copper from chalcopyrite. There is considerable incentive to design sulfate-based hydrometallurgical processes, which can achieve high extraction from chalcopyrite under ferric sulfate and biological leaching conditions in a reasonable time frame. The development of such processes requires further understanding of the "passivation" phenomenon of chalcopyrite in sulfate media in order to propose innovative ways to prevent the occurrence of passivation. It is this line of investigation that provides the impetus for the present study. : C H A P T E R 1 I N T R O D U C T I O N / 3 1.1 Outline of the Thesis This thesis is an embodiment of six chapters. Chapter l introduces the significance of the current research interest within the scope of chemical (ferric sulfate) leaching of chalcopyrite. Chapter 2 examines the existing literature concerning the chemical (ferric sulfate) and biological leaching of chalcopyrite. The literature highlights many mechanistic investigations o f chalcopyrite leaching, particularly surface analysis studies of leach residues, which help explain chalcopyrite passivation in ferric sulfate based systems. The most significant advance in the ferric sulfate leaching of chalcopyrite has been the use o f controlled potentials, which considerably enhances the extraction of copper from the mineral. Chapter 3 presents the experimental procedures, experimental considerations, and experimental setups. Chapter 4 reports and discusses the results of the controlled-potential test work with chalcopyrite. Chapter 5 draws conclusions on the role of controlled-potential in chemical (ferric sulfate) leaching systems, and some recommendations for future work are given in Chapter 6. CHAPTER 2 LITERATURE REVIEW / 4 C H A P T E R 2 L I T E R A T U R E REVIEW This review of literature will introduce the reader to several aspects of chemical (ferric sulfate) and bacterial leaching of chalcopyrite. It is useful to discuss the nature of the passivating layer and its inhibiting role during the ferric sulfate leaching of chalcopyrite. In particular, the results of previous testwork related to controlled-potential leaching of chalcopyrite are presented in this chapter. 2.1 Properties of Chalcopyrite Chalcopyrite ( CuFeS2 ) behaves as an «-type semiconductor (having excess metal cations). ' It has many properties of metals, including metallic luster and electrical conductivity. It is a good conductor with relatively small resistivity varying from 9x1 OJ to 2X10" 4 Q.m, and an energy band gap (Es) of 0.6 eV at room temperature.13'14 The resistivity of chalcopyrite increases slowly with temperature, and declines rapidly after reaching 310°C." Chalcopyrite is brittle, with a Mohs' hardness of 3.5-4.0, and exhibits a low solubility in water (2.35xl0" 6mol/Lat 110°C). 1 3 ' 1 5 ' The crystal structure of chalcopyrite is derived from the zincblende structure (ZnS ) in which alternate zinc atoms are replaced by copper and iron atoms along the c axis (Figure 2-1). 1 6 Due to the slight asymmetry caused by alternating copper and iron atoms, the sulfur atoms are displaced slightly from the center of each of the metal tetrahedral towards the iron-iron edge. The tetragonal unit cell of chalcopyrite is twice the cubic unit cell structure of the zincblende structure. The short bonding distances between the cations and anions indicate an essentially covalent bonding nature, but there is no extensive direct covalence among the anions (Table 2-O . M - M ' T h e s u ] f u r a t o m s are arranged in an essentially face-centered cubic (fee) array. Each sulfur atom is tetrahedrally coordinated to four metal atoms (twoFe and two Cu), which are located in interstices in the close-packed sulfur atom lattice. Similarly, each metal atom is tetrahedrally coordinated to four sulfur atoms.17 Metal atoms occupy alternate planes of tetrahedral interstices, and octahedral interstices are unoccupied. CHAPTER 2 LITERATURE REVIEW / 5 Although two bonding models have been suggested for the ionic states of chalcopyrite, a combination of Mossbauer and X-ray photoelectron spectroscopic methods strongly support the ionic bonding model of chalcopyrite as " Cu+Fe}+(S2~)2 " . 1 8 ' 1 9 The other bonding model "Cu2+Fe2+(S2~)2" has also been suggested to co-exist within the same mineral sample.20 The magnetic structure of chalcopyrite has been determined by neutron diffraction.21 The spins oh the two iron atoms tetrahedrally bonded on a common sulfur atom are anti-parallel along the c-axis, resulting in anti-ferromagnetism, or a net zero magnetic moment. Although the anti-ferromagnetic structure observation of chalcopyrite supports the former state, the latter structure cannot be disregarded entirely. Some physical and thermodynamic properties of chalcopyrite are depicted in Table 2-1 and Table 2-2. (a) (b) Figure 2-1 a) Unit-cell model showing the crystal structure of chalcopyrite (CwFe5 2 ) 1 6 b) Interatomic bond lengths and angles of each atom in chalcopyrite17 CHAPTER 2 LITERATURE REVIEW / 6 Table 2-1 Crystal structure of chalcopyrite (CuFeS2) Mineral Name : Chalcopyrite Ref. Ionic Structure : Cu+Fe}+(S2')2 18,19 Crystal Structure : Tetragonal 17 Cell Content: 4[CuFeS2 ] -CuFeS2 per cm : 1.3 x 10 2 2 17 Unit Cell Dimensions (A) : a = 5.2988 ±0.0010 c = 10.434 ± 0.005 (c = 2a = 2b) 22 Cu-S Bond Distance (A): 2.302 22 Fe-S Bond Distance (A) : 2.257 22 S-S Bond Distance (A) : 3.68 22 Table 2-2 Physical and thermodynamic properties of chalcopyrite (CuFeS2 ) at 25°C Mineral Name : Chalcopyrite Ref. Color: Brass-yellow, often tarnished and iridescent 23 .. C ; (J mol"1 K" 1 ) T =298-830 K : 88.71 + 5.19x10"2 T - 7.45x105 T~2 13 AG;. (kJ mol"1) : -173.4 13 AH} (kJ mol"1) : -173.3 13 a - CuFeS2 (tetragonal)—> (3 - CuFeS2 (cubic) : 550°C 13 Most intense XRD Peaks (2-theta angle): 29.371,49.014, 57.847 24 Energy Band Gap at Room Temp (eV) : 0.60 13 Resistivity (C>m) : 9xl0" 3 -2xl0" 4 14 Copper (Cu): Iron (Fe): Sulfur (S) Content: 34.6 : 30.4 : 34.9 (total: 99.9 wt%) -Molecular Weight (g/mol) : 183.52 -Specific Gravity at 25°C (g cm"3) : . 4.1 -4.3 ' 25 Hardness, Mohs a : 3.5-4.0 13 a The Mohs' hardness scale is defined with respect to the hardness of diamond, which is assigned the value of 10. 2.2 CHAPTER 2 LITERATURE REVIEW / 7 Theoretical Considerations of Ferric Sulfate Leaching The EH - pH, precipitation and speciation diagrams are useful guides for the evaluation and development of possible ferric sulfate leaching mechanisms in both chemical and biological oxidation of chalcopyrite. 2.2.1 Thermodynamics of Ferric Sulfate Leaching of Chalcopyrite The EH - pH (or Pourbaix) diagrams focus attention on the thermodynamics of possible reactions as determined by the equilibrium relationships between dissolved and solid species. These diagrams are commonly used as an initial indication of the stability of the sulfide minerals in aqueous solutions under the reducing or oxidizing environments at various pH values. Since thermal equilibrium between species is usually not observed in hydrometallurgical systems, the formation of many non-stoichiometric phases gives rise to apparent discrepancies and 6 26 27 28 ambiguities in the interpretation of results in transient systems. Peters ' ' ' has promoted the application of metastable stability diagrams to systems in which there are kinetic controls, including leaching and electrolytic oxidation of copper sulfide minerals. He has said in relation to hydrometallurgy; "when you are in a hurry, you always observe the faster processes, but you 28 may treat some of the slower processes as i f they were not taking place at all." Figure 2-2 is a metastable Cu-Fe-S-H20 diagram, which is consistent with hydrometallurgical observations under acidic ferric ion leaching conditions.29 Although pyrite is a relatively good catalyst for oxygen reduction, the oxidative leaching of chalcopyrite can be performed under a nitrogenous atmosphere. Pyrite (FeS2) shows little reactivity towards ferric sulfate leaching as in all experiments performed within and it' is excluded from the figure. According to Peters6, pyrite does not seem to leach at all (except at temperatures approaching 200°C) during oxidative ferric sulfate leaching of chalcopyrite, and reports almost quantitatively to the residue. Experimental results also indicate that elemental sulfur is an extremely thermodynamically stable species due to the high over-potential required to form sulfate. The free energy of formation of sulfur is artificially extended by 300 kJ/mol to account for the observed stability of elemental sulfur in the presence of ferric oxidant, which has the standard CHAPTER 2 LITERATURE REVIEW / 8 redox potential of 0.771 VSHE at 25°C. Alternatively, aqueous lixiviants may not be effective towards the hydrophobic elemental sulfur. The shaded areas represent the predominance areas for chalcopyrite, covellite (CuS) and (Cu2+ + S"). Bornite (Cu5FeS4) and chalcocite (Cu2S) appear only as reduction products. Chalcopyrite is thermodynamically unstable at potentials above 0.3 V$HE- Since cuprous ion is not stable in solution without complexing anions (e.g., Cl ) at temperatures less than about 200°C, cuprous ion is rapidly oxidized to cupric ions even i f initially present in the solution. The ferric ion is then expected to act as an oxidant (E°for Fey+I2+ = 0.771 VSHE) for chalcopyrite dissolution in the Cu-Fe-S-H20 system. According to this diagram, iron can be leached out of chalcopyrite as ferrous ion, leaving the mineral covellite (CuS) as the only intermediate sulfide phase. Covellite is thermodynamically stable up to 0.6 VSHE over the pH range 0 to 3.5. Under more strongly oxidizing conditions, the predominance area of Cu2+ and S° species predicts the dissolution products for the pxidation of covellite and represents the operating conditions for the oxidation of chalcopyrite. The dissolution of chalcopyrite in acidic ferric sulfate solution can then be interpreted as an electrochemical reaction, in which chalcopyrite is dissolved anodically with liberation of elemental sulfur. The liberated electrons are consumed for the reduction of ferric iron, according to the equations: CuFeS, -» CuS + S° + Fe2+ + 2e (2-1) CuS Cu2+ +S° +2e (2-2) Reduction of ferric oxidants: 4Fei+ + 4e 4Fe 2+ (2-3) The combined reaction including ferric reduction.is described by the equation: CuFeS, + 4Fei+ -» Cu2+ + 5Fe2+ + 2S° (2-4) CHAPTER 2 LITERATURE REVIEW / 9 Co <2*> Cu <2*> Cu <2*> Fe <2+> Fe <2+> • Fe <2+> (AO) • Fe <3*> (fldl •• Fe2 03 • Cu S Cu Fe S2 (ROI (ROI IRQ) (ROI (ROI Fe2 03 • CU S Fe2 03 • Cu Fe S2 Cu S • Cu F« S2 Fe S * Cu Fe S2 Cu 0 + Fe2 03 Fe <2+> (ROI • Cu2 5 Fe <2+> (RQI * Cu5 Fe SU Fe <2*> (0Q) • Cu Cu2 S + Cu5 F« St Cu S + CuS Fe S<i Fe S + Cu5 Fe SU Fe S + Cu Cu Fe S2 • Cu5 Fe S« Cu5 Fe S4 * Cu Cu + Fe 1 Cu 2 + + S° CuS CuFeS2 Figure 2-2 The metastable EH-pH diagram for the Cu - Fe - S - H20 system at 25°C (without FeS2 as a product); all solutes at 1 mol/L activity 29 CHAPTER 2 LITERATURE REVIEW / 10 2.2.2 Importance of Iron Precipitation and Jarosite Formation In chemical (ferric sulfate) systems, the rate of leaching depends significantly on the concentration of iron in solution. The oxidation of sulfide minerals leads to the release and accumulation of ferric (Fe3+) and sulfate (S04~) ions, and increases the acidity of the leaching system. Ultimately, ferric ion in acid sulfate systems may reach saturation and hydrolyze to form various iron precipitates on the sulfide surfaces.3 0'3 1'3 2 These precipitates may impart a barrier for the diffusive transport of oxidants (ferric ion or dissolved oxygen) to the unreacted material, subordinating the leaching kinetics to its dissolution, and can give poor copper extraction from sulfide minerals.33 The hydrolysis and precipitation of iron is an extremely complicated process and is affected by a large number of variables. Hydrolysis in sulfate systems is further complicated by the fact that the concentration of free sulfate ions depends on the pH through the S04~ IHS04 equilibrium. In solutions not containing sulfate, the precipitates that form are oxide and hydroxides. In the presence of sulfate, additional solid compounds, particularly jarosites, can precipitate. Regardless of the precise process, the course of hydrolysis and precipitation is extremely sensitive to variables such as pH, temperature, and solution compositions (particularly the presence of complexing ions, the total iron content, and other monovalent cations). The overall hydrolysis reactions of ferric sulfate are shown below: Ferric hydroxide [Fe{OH) i, Fe2 (OH)4 ] and ferric hydroxysulfates [ Fe{OH)S04 ]: Fe2{S04)y +6H20 2Fe{OH)i + 3H2S04 (2-5) Fe2(S04), + 2H20 2Fe(OH)SOA +H2S04 (2-6) 2Fe(OH)S04+2H20 + 2H+ +2e~ Fe2(OH)4 + 2H2S04 (2-7) Ferric oxyhydroxides (or goethite) [FeO • OH]: Fe2(S04), + 4H20 2FeO • OH + 3H2S04 (2-8) Jarosite-group precipitates [ AFez {S04 )2 {OH)b ]: Fe2(S04), + 4H20 + 0.67 A + -» 0.67 AFe,(S04)2(OH)b + 1.61H2S04 +0.67/T (2-9) CHAPTER 2 LITERATURE REVIEW / 11 Jarosite is commonly detected on the surface of leached chalcopyrite particles. In equation (2-9), jarosite-group precipitates (basic ferric hydroxysulfates) have the general formula AFe3(S04)2(OH)6, where A isK+, Na+ ,NH+4 ,H30\ Ag+ or 0.5Pb2+ . 3 4 These jarosite-group precipitates have been found as naturally occurring minerals, and can be synthesized under laboratory conditions (Table 2-3). Table 2-3 Chemical and mineral names of common jarosites Formula Chemical Name Mineral Name KFe3(SOA)2(OH)6 Potassium jarosite Jarosite NaFe3{S04)2{OH)b Sodium jarosite Natrojarosite (NH4)Fe3(S04)2(OH)b Ammonium jarosite Ammoniojarosite {H30)Fe3{S04)2(OH)b Hydronium jarosite Hydronium jarosite (or carphosiderite in early literature) AgFe,{S04)2{OH)b Silver jarosite Argentojarosite Pb05Fe3(SO4)2(OH)b Lead jarosite Plumbojarosite Figure 2-3 presents the stability fields for potassium jarosite, hematite, goethite (ferric oxyhydroxides) and ferric hydroxide precipitates.35 The hydrolysis reactions of ferric ions depend largely on the pH reached during the course of leaching. Low pH is necessary to enhance the solubility of ferric ions, and prevent them from hydrolyzing. Appreciable amounts ofFei+ can remain in solution only at pH values below 2. At pH around 2, potassium jarosite and other related jarosite-group precipitates form as secondary minerals in the presence of suitable monovalent cations, which cannot be avoided during ferric sulfate leaching that operates at 20-100°C. When the pH rises above 3, complete ferric ion precipitation occurs and forms highly insoluble iron precipitates (ferric hydroxide, goethite and hematite). Hematite forms at temperature exceeding 100°C, and does not take place at any extent during ferric sulfate leaching (i.e., 20-100°C). When dissolved sulfate levels decline, potassium jarosite and other related jarosite-group precipitates are expected to dissolve and reprecipitate as goethite (a -FeO • OH ), which is usually considered to be the most stable form of Fe(III) oxide. 3 6 CHAPTER 2 LITERATURE REVIEW / 12 ^120 U o . 80 40 0 HEMATITE F e 2 0 3 JAROSITE GOETHITE FeOOH F ^ J O H J T J I i pH Figure 2-3 Stability fields of various iron precipitates as a function of pH and temperature [0.5mol/LFe 2(S<9 4),] 3 5 Figures 2-4 and 2-5 present the precipitation of hydronium jarosite (or carphosiderite), ferric hydroxysulfates and ferric sulfates.37 Sulfate is the most common anion association during ferric sulfate leaching, and it is expected to play a major role in precipitation in the Fe2(<S,04)3 - H2SOA - H20 system. The potential precipitation product's depend on the concentrations of chemical species [ Fei+, SO]~ ] in solution and temperature (50-200°C) of a leaching system. At temperatures below 100°C, goethite appears at less than 5 wt% of dissolved sulfate levels, and hydronium jarosite predominates between 5 and 30 wt%. Ferric hydroxysulfates present themselves between 30 and 43 wt%, and other related ferric sulfates emerge at higher wt% of dissolved sulfate levels. It is fortunate, therefore, that bioleaching bacteria thrive under strongly acid conditions. Since typical bioleaching is carried out in the pH range 1.5-2.5, much of the bacterially oxidized iron may hydrolyze as jarosite-group precipitates. The acid produced by this reaction may CHAPTER 2 LITERATURE REVIEW / 13 ultimately lead to lower pH values, which encourage the presence of aqueous Fei+ species. In general, pH values less than 1.8 seem to limit the extent of Fei+ precipitation effectively, and the 38 solution pH should not exceed 3.0 if jarosite is to remain a stable phase. The relative order of formation of common jarosite-group precipitates is as follows: 3 9 Potassium jarosite > ammonium jarosite > sodium jarosite > hydronium jarosite The jarosite produced abiotically and in the presence of bacteria have shown no chemical or structural differences.40 Clearly, unwanted precipitation of ferric iron must be avoided, as it would coat the minerals and prevent bacterial attachment at the mineral surfaces. However, some researchers believe that Acidi thiobacillus ferrooxidans cells serve as nucleation sites for the formation of jarosite. 4 1 ' 4 2 ' 4 3 ' 4 4 The bacteria promote the precipitation of jarosite, by local reaction of ferric sulfate with potassium or sodium ions acid-dissolved from the clay/feldspar component of the heap, and the combined sulfur-jarosite mixture forms an impermeable layer on the surface of the chalcopryite.23 One researcher argued that bacteria lower the activation energy for the formation of ferric hydroxysulfate polymers, which are the precursors for the precipitation of either amorphous ferric hydroxysulfates or crystalline jarosites.41 CHAPTER 2 LITERATURE REVIEW / 14 Sulplut*. Weight % Figure 2-4 System Fe2 (S04), - H2S04 -H2Oa) Polytherm 50 to 200°C, 0 to 40 wt% S04 JO 40 SO 60 70 S<4pt>aU, weight % Figure 2-5 System Fe2(S04)3 -H2S04- H2Ob) Polytherm 50 to 200°C, 30 to 70 wt%5C 4 3 CHAPTER 2 LITERATURE REVIEW / 15 2.2.3 Speciation Distribution of Iron in Ferric Sulfate Solution Speciation diagrams allow the estimation of quantitative distributions of reaction species in solution as a function of pH and temperature. Many non-predominant iron species are present in solution, and many different kinds of reactions are possible during acidic ferric sulfate leaching.4 5 Therefore, the "oxidizing activity" of ferric sulfate solution must be viewed as a function of all the possible ferric species present in the ferric sulfate solution. Because of the industrial application of bioleaching, the influence of these chemical species on leaching bacteria such as Acidi thiobacillus ferrooxidans is often investigated. The accumulation of some ionic forms of iron species may be toxic to bacteria and inhibit the microbial oxidation of ferrous iron during the course of reaction.4 6'4 7 The speciation equilibrium distribution diagram (Figure 2-6) clearly supplies information regarding the activity of various species of iron as a function of pH at 25°C. 4 5 Sulfate forms complexes with both ferric and ferrous ions. The majority of iron(lll) species exists in the form of Fe(S04)2 complexes and the rest in the form of FeS04 within the pH range of 1 to 2. In this pH range, the FeS04 is probably the more electroactive species in the anodic dissolution of chalcopyrite in acidic ferric sulfate solution, where liberated electrons are consumed by the positively charged ferric monosulfate complex more so than the negatively charged disulfate complex. Under similar conditions, ferrous ions ( F e 2 + ) form considerably weaker complexes with sulfates. About 12% of iron(II) occurs in the form of uncomplexed Fe2+ ions, and the rest in the form of FeS04 molecules. The effect of adding more H2S04 (lower pH) to the system increases the concentration of the more reactive ferric monosulfate complexes, FeS04 and HS04 , followed by the formation of FeHSQ4+ complexes. This would lower theFeS04 ion-pair concentration.48 Fe(SQ4y2 + H+ -» FeSO; +HSQ4 (2-10) Fei+ + HSO; -> FeHSO] 2 + (2-11) Speciation diagrams allow us to provide an alternate mechanism that leads to the formation of jarosite-group precipitates in the presence of acid solutions of ferric sulfate with CHAPTER 2 LITERATURE REVIEW / 16 alkali metal ions. It is postulated that the dissolved ferric sulfato [Fe(S04)2] complexes may slowly convert to a variety of hydroxo ferric sulfate minerals, most notably jarosite-group • • 44 precipitates: Fe3* +2SO]- ->Fe{S04)2 (2-12) Fe(OH)j+Fe(S04)~ +H20 +0.67A + 0.67A[Fei(S04)2(OH)6] + 0.67SO2' + H+ (2-13) Therefore the overall reactions result in: Fe(OH),+\.33S04- +Fei+ +H20+0.67A+ -> 0.67 A[Fei{S04)2(OH)b] + H+ (2-14) Where: A can be K+, Na+ , NH+4,H,0\Ag+ or 0.5Pb2+etc. CHAPTER 2 LITERATURE REVIEW / 17 2.3 An Introduction to the Ferric Sulfate Leaching of Chalcopyrite Acid ferric sulfate leaching has been studied extensively because of its significance in the commercial dump leaching of chalcopyrite, and as a possible leaching reaction for the hydrometallurgical processing of other copper sulfides.9'23 The main practical difficulty encountered in the commercial acid ferric sulfate leaching of chalcopyrite (in heaps, dumps or concentrates) is its slow rate of reaction, which renders a long residence time and/or incomplete copper extractions. Both chemical and electrochemical studies agree that a passivating surface layer exists and limits the oxidation rate, resulting in the observed parabolic kinetics (rate °c / l / 2 ) found in sulfate media. The exact nature of the mechanism whereby passivity occurs is not entirely clear, but the parabolic kinetics of chalcopyrite dissolution is thought to result from the formation of a "passivating layer" on the chalcopyrite surface that causes a dramatic decrease in the rates of copper release. However, there is less unanimity in the interpretation of the reaction kinetics to explain rate control for chalcopyrite dissolution. In this regard, the influence of process conditions on the leaching rate of chalcopyrite and the nature of the passivating surface film must be investigated to establish the optimum conditions for copper recovery. CHAPTER 2 LITERATURE REVIEW / 18 2.3.1 Chemical (Ferric Sulfate) Leaching of Chalcopyrite The chemistry of chalcopyrite has been described extensively in the literature. 49,50,51,52,53,54 p e r r j c S Upp]j e fj j n m e form of ferric sulfateFe 2(S0 4) 3, is a powerful oxidant, which gives more rapid sulfide dissolution compared with molecular oxygen.5 5 At moderately low temperatures (T < 95°C), the sulfide sulfur of chalcopyrite (CuFeS2) can be completely or incompletely oxidized to sulfate or elemental sulfur, respectively, in an acidic solution of ferric sulfate. The overall dissolution reactions may be expressed stoichiometrically by .the following reactions: CuFeS2 + 2Fe2 (S04 )3 -> CuSQ4 + 5FeS04 + 25° (2-15) CuFeS2 + SFe2 (S04 ) 3 + SH20 -» CuS04 +1 IFeSO, + SH2S04 (2-16) Table 2-4 Analytical data for sulfur species in the ferric sulfate oxidation of chalcopyrite Ref. Temperature Leaching Time Medium S° dissolved in CS2 56 T=35°C 42 days 0.45 mol/L Fe2(S04)3 65-75% . 53 T=90°C 38 days 1 mol/L Fe2(S04)3 82% 49 T=95°C 0-70 hours 0-1 mol/L Fe2(SQ4)3 94% Table 2-4 shows that the incomplete chemical oxidation of chalcopyrite is the predominant reaction, as depicted in equation (2-15). Advantageously, the sulfide sulfur of chalcopyrite is oxidized to elemental sulfur, rather than to sulfate, thereby substantially decreasing the requirement for oxidative agents, and also subsequent neutralization requirements. The experimental study at 95°C indicates the generation of about 94% 5 ° , regardless of leaching time, ferric sulfate concentration and chalcopyrite particle size. 4 9 Another study found virtually stoichiometric yields of elemental sulfur and ferrous ion in the temperature range 50 to 94°C. 5 7 Since only orthorhombic (Sa) and monoclinic (Sp) elemental sulfur are soluble in CS2, the percentage of elemental sulfur ( 5 8 ) listed in Table 2-4 should be regarded as minimum to account for the amorphous plastic sulfur, which is not extractable in CS2,53 It was suggested that the physical properties of plastic sulfur may allow it to form the dense tenacious layer found in leaching experiments.50'58 Over a period of days, the plastic sulfur eventually transforms into the CHAPTER 2 LITERATURE REVIEW / 19 most stable form of sulfur (Sa) at room temperature. It is interesting to note that the morphology of precipitated sulfur should be different from the residual sulfur left after ferric leaching. For instance, sulfur on leached chalcopyrite is present in isolated sulfur globules, which is difficult to explain unless the sulfur has been precipitated from solution and deposited on pre-existing sulfur 49 grains. The copper leaching kinetics in acidic sulfate media is unusual during the ferric leaching of chalcopyrite. Studies have shown that the initial copper release is very rapid. 5 7 ' 5 9 As the reaction proceeds, the leaching rate becomes extremely slow at or before 50% copper recovery.9 In tests, regrinding of the chalcopyrite residue restores the fast leaching rate, and achieves adequate copper recovery.9'60 Munoz et al.50 achieved a maximum recovery of about 58% with 4 micron particles in about 20 hours at 90°C, 1200 rpm and 1.0 M sulfuric acid. An extended leaching time of 100 hours was necessary to achieve the same recovery of 58% with 12 microns particle size. Less than 20% copper recovery was achieved even after 160 hours with 47 microns particle size. Consequently, many investigators concluded that the rapid formation of a "passivating" surface film on partially leached chalcopyrite particles is responsible for the low copper extraction, and that the passivating surface film can be removed by grinding. However, regrinding of chalcopyrite is not often possible in commercial leaching operations (dump, heap, and in-situ processes).9 Various fundamental studies of the reaction of chalcopyrite in/ re 2(5(9 4), media have been reported, and a summary of the relevant information is given in Tables 2-5 and 2-6. In the ferric sulfate system, reaction rates are thought to be dependent on temperature and the surface area of the particles. Temperature is an important parameter affecting the rate of hydrometallurgical processing of copper concentrates. The slow kinetics of ferric sulfate leaching for chalcopyrite ores often requires elevated temperatures (>60°C). In most cases, there is a direct correlation between the reaction rate and the surface area (or grain size) of the chalcopyrite. Sufficient agitation must be provided to suspend the particles, but higher stirring rates do not promote the leaching of chalcopyrite.52'54 Most researchers agree upon the negative effect of FeSOA addition on ferric sulfate leaching of chalcopyrite. This effect was ascribed to the slow kinetics of ferrous oxidation with oxygen on the chalcopyrite surface in sulfate media.6 1 Interestingly, recent studies observed that ferrous sulfate solution in the presence of dissolved oxygen is a more effective leaching reagent than ferric sulfate solution under the same leaching CHAPTER 2 LITERATURE REVIEW / 20 conditions [30°C, 168 hours, I M Fe2(S04)3 or FeS04].62,63 In contrast, initial ferric sulfate concentration (adequate to initiate leaching, above 0.005M) 5 7, the concentration of sulfuric acid (sufficient to prevent iron hydrolysis at concentration > 0.1 M ) 5 0 ' 5 4 ' 6 4 and the presence of CuSOA 5 0 ' 5 4 ' 6 4 have little or no effect on the dissolution of chalcopyrite. The reported kinetics for the leaching of chalcopyrite is inconsistent in proposing the rate determining step in ferric sulfate systems. Most researchers concluded that ferric sulfate leaching obeys a parabolic rate relationship between chalcopyrite conversion and time in the temperature range of 50-95°C (Table 2-5). Parabolic leaching kinetics implies the formation of a product layer that controls or inhibits the rate of leaching. Many oxidation products have been observed on the surface of chalcopyrite and insoluble reaction products (mainly precipitated iron compounds65, jarosites66 or S") are thought to hinder chalcopyrite dissolution. Other researchers reported linear kinetics, and they attributed rate control to a surface reaction (chemisorption or electrochemical surface reactions) as the rate-determining processes. 5 3' 6 7 Lowe 6 7 reported the calculated activation energy (E = 75 kJ/mol) that characterized a chemical reaction, and he attributed the rate limiting step to the surface-controlled chemisorption (surface saturation of ferric sulfate in solution). Jones and Peters53 found that leaching was halted even though 18% of the sulfur formed was oxidized to sulfate in 17 days, suggesting the possibility of some limiting factor other than the diffusion-limiting elemental sulfur layer. The overall leaching process was concluded to be under mixed control, where the combination of anodic and cathodic processes taking place on the mineral surface controls the rate of dissolution. A mixed linear-parabolic kinetics is observed by Hirato et al.61 In this study, the parabolic kinetics was attributed to the formation of a dense sulfur layer during the initial stage (-100 hours); thereafter, the leaching rate increases and shows linear kinetics over an extended period as the sulfur layer grows and peels off the surface (Figure 2-7). The present knowledge of the leaching mechanism does not allow us to give a consistent interpretation of experimental observations. The variations in sample preparation, experimental conditions and procedures, particle sizes as well as the sample purity of different investigations make it difficult to establish universally acceptable parameters employable for all systems. The other major problem remains the general transferability of short-term laboratory data to predict long-term events in commercial dump and heap operations. Since most experiments were CHAPTER 2 LITERATURE REVIEW / 21 completed within about 100 hours, the parabolic kinetics that was observed by many investigators may possibly be valid only in the initial stage of commercial leaching practices. If a longer leaching period is employed, the kinetics may well be linear as supported by the work of Jones and Peters.53 In the case of Hirato et , who observed a mixed linear-parabolic kinetics, the overall kinetics can possibly become linear i f the leaching time were further prolonged (Figure 2-7). Time/ h Figure 2-7 Leaching rate curve of chalcopyrite with ferric sulfate61 [70°C(343K), [S024-]Tolal = 0.2 mol/L, Fe{IIl) = 1.0 mol/L)] C H A P T E R 2 L I T E R A T U R E R E V I E W / 22 CD CM O CD 13 X o <+-. o c t3 12 '* o <L> 3 C U O » o 03" I-. C D V* C D CM Q. O U CN C D X) nperature u o i n O N o o o O N U o O O N o C r o O N U c i n O N O c O O N cu H <u Leaching Tin -50 hours -100 hours 15 days -3 hours -50 hours -70 hours cu g P P 3. 1 CN 6 B . i =L Particle -14+10 3. o o m 1 i n o -45+38 -14+10 a o Extracti 0 s -O O N 0 s -00 i n 0 s -o r -o O N I T ) r o O I T ) Cu nee Refere O N o i n r o i n i n m N O ^ J -N O 1— CM O o 13 C D o C D C c _o 13 'S o C D =3 .0) M>M C D X » c o c/3 C o r-, O O c/a cn C D C D O C D C D c -C D C D X <+M o E 00 N O I CN C D X l 03 Inert Gas Atmosphere No No No No Yes No Yes Yes Yes Inert Gas Atmosphere Activation Energy (kJ/mole) i oo i • 71 ± 13 r o oo 38-63 • m c— Temp. O c i n ON O o O ON I o NO U c O ON. U c r o ON U o ON o i n o c o r -U c i n ON o r o O c O ON U c O i n CN r o Leaching Time -50 hours -100 hours Over 55 days -3 hours -60 hours -300 hours -50 hours 65 hours -8 hours o Leaching Kinetics Parabolic Parabolic Linear Parabolic Parabolic Linear-Parabol Parabolic Parabolic Linear Particle Size/ Surface Area i Positive No Positive Positive • Positive Positive X a i No i No No I i No + a i No l No 1 i No l + • No Negative No Negative Negative i Negative i 2 " .ti 7^  a k, NH I i • No Slightly No No Slightly • No Ref. ON o m r o >n i n r — i n MS >n NO NO t — NO CHAPTER 2 LITERATURE REVIEW / 23 2.4 An Introduction to Bacterial Leaching The introduction of bacteria to leaching systems, or bioleaching, has assumed significant importance for the mining industry. Bioleaching is an effective and essentially safe technology to extract metals from low-grade ores (unsuited to concentration by flotation) and refractory mineral concentrates by naturally occurring microorganisms. Since bioleaching operates essentially at ambient temperature and pressure, process equipment is simple and inexpensive. Since metal sulfides are commercially extracted in acid leaching operations, acidophilic bacteria are natural candidates, which catalyze much faster oxidation of metal-bearing sulfide minerals, ferrous ion, and elemental sulfur under milder conditions than conventional chemical leaching.69 It is especially successful for copper ores, because the resulting copper sulfate products are very water soluble, and can be processed by the SX/EW process. Current literature on the application of bioleaching of sulfide minerals is dominated by studies of a mixed culture of mesoacidophilic bacteria, including Acidi thiobacillus ferrooxidans, Acidi thiobacillus thiooxidans and Leptospirillum ferrooxidans.70 The advantage of a mixed culture is found in its ability to break down minerals that no single species could oxidize alone. They are aerobic, mesophilic and acidophilic, i.e., they require oxygen and their optimum temperature and pH fall within the range 25-35°C and 1.5-2.5, respectively. They are able to adapt to pH values much lower than 1.5, given enough time. In practice, however, the pH is not allowed to exceed the range 1.5-1.8 to prevent jarosite precipitation, which would possibly interfere with leaching.33 Recent investigations have focused on moderate fhermoacidophiles such as Sulfobacillus and extremothermoacidophilic bacteria such as Sulfolobus, which grow in the temperature range 40-60°C and 60-80°C, respectively. 3 3 ' 7 1 ' 7 2 Since the oxidation rate increases with increasing temperature, thermophilic bacteria have been considered as a means to improve the leaching kinetics of chalcopyrite. The bioleaching mechanisms of chalcopyrite and the passivation phenomenon are discussed briefly below. Bacterial leaching is defined as the chemical attack of an ore in the presence of bacteria. The presence of bacteria has been proposed to assist the leaching of chalcopyrite by both direct and indirect mechanisms. The distinguishing feature between the two mechanisms is the role played by ferric ions in the dissolution of the mineral as shown by the stoichiometric equations listed in Table 2-7. The direct mechanism affects leaching at the point of attachment to sulfide CHAPTER 2 LITERATURE REVIEW / 24 minerals, without the use of the ferric iron oxidant, but by some biological agents; this mechanism is essentially biological reaction that requires oxygen for the generation and the regeneration of the ferric oxidant.73 Chalcopyrite is completely oxidized to copper sulfate and ferrous sulfate via equation (2-17), and bacterial oxidation of ferrous sulfate also takes place via equation (2-18). In the indirect mechanism, the bacteria are only indirectly involved since they affect leaching through the continuous oxidation of ferrous ions to ferric. The sulfide moiety of chalcopyrite can be completely or incompletely oxidized to sulfate or elemental sulfur, respectively. In the case of an incomplete chemical oxidation, the specific role of bacteria includes the continual oxidation of elemental sulfur to sulfate. The accumulation of rather porous elemental sulfur (S s) can coat unreacted mineral surfaces and hinder the progress of leaching by imposing a diffusion limitation of the reacting species. The overall stoichiometric equation is the same for both mechanisms as shown in equation (2-24). Table 2-7 Stoichiometric equations for the direct and the indirect bacterial oxidation of 8 11 chalcopyrite ' Direct Mechanism CuFeS2 + 40 2 -» CuS04 + FeS04 FeSQ4 + 0.25<?2 +0.5H2SO4 -> 0.5Fe2(SOA)} +0.5H2O (Biological) (Biological) (2-17) (2-18) Incomplete Indirect Mechanisms CuFeS2 + 2Fe2(S04)i -» CuS04 + 5FeSOA + 2S° 5FeS04 +1.2502 +2.5H2S04 -» 2.5Fe2(SOA)3 + 2.5H20 2S° +302 +2H20-^2H2S04 (Chemical) (Biological) (Biological) (2-19) (2-20) (2-21) Indirect Mechanisms CuFeS2 + 8Fe 2(S0 4), +SH20^ CuS04 +1 !FeS04 + W2SOA MFeSQ4 +4.250 2 +S.5H2SQ4 -> S.5Fe2(S04), +S.5H20 (Chemical) (Biological) (2-22) (2-23) The Overall Equation for Both Direct and Indirect Mechanisms CuFeS2 +4.25<92 +0.5H2SO4 -> CuS04 +0.5Fe2(SO4)y +0.5H2O (2-24) Bacterial leaching of chalcopyrite is similar to that of the chemical ferric leach. It can be generalized into two distinct s tages . 7 4 ' 7 5 ' 7 6 ' 7 7 The first stage is controlled by the rate of chalcopyrite dissolution. This stage is characterized by an initial rapid release of copper in solution and a rapid increase in solution potential. A rapid increase in solution potential indicates a high rate of bacterial growth. When a product layer becomes thick enough to become rate CHAPTER 2 LITERATURE REVIEW / 25 controlling by diffusion limitation of reactants and/or products, a leaching process is ready to move into the second stage. The second stage is characterized by a rapid decrease in leaching of copper which eventually forms a plateau at a final value when dissolution is completely passivated by this layer. The passivating layer may be elemental sulfur, jarosite precipitate or a metal depleted sulfide layer. CHAPTER 2. LITERATURE REVIEW / 26 2.5 The Nature of the Passivating Layer The exact nature of the passivating layer is still under considerable debate, and has not been elucidated with certainty. During both bacterially assisted and chemical ferric sulfate leaching of chalcopyrite, the mineral is oxidized by iron oxidants. The success of regrinding (bio)leach residues is probably due to exposure of fresh chalcopyrite surfaces.9 A loosely termed "passivating layer", also known as the reaction zone, is thought to form on the chalcopyrite surface.78 The reaction zone may be a diffusion barrier or a completely inaccessible solid barrier. This layer may hinder greater copper extraction by restricting flow of bacteria, 81 nutrients, oxidants and reaction products to and from the mineral surface. Various theories have been proposed to explain the nature of the passivating layer, including precipitates (e.g., precipitated iron compounds, jarosite, elemental sulfur) and a partly altered non-stoichiometric phase, commonly called a reaction zone (e.g., bimetallic intermediate sulfide layer, iron-deficient polysulfide like covellite (CuS), ferric hydroxyl sulfate). 5 4 ' 7 9 ' 8 0 ' 8 1 The passivating layers will be discussed in terms of leach residues from chemical studies, electrochemical studies of chalcopyrite, and surface studies of chalcopyrite (AES and XPS). Two general schemes of leach residues are considered below: (1) precipitated iron compounds and jarosite passivating layers and (2) elemental sulfur and bimetallic intermediate sulfide passivating layers. The implication of catalytic silver ion in ferric sulfate (bacterial) leaching of chalcopyrite is also discussed in relation to the passivating layer. On the other hand, the evidence from electrochemical and surface analytical studies of leached chalcopyrite surfaces have led to the formulation of some interesting passivation theories for chalcopyrite leaching. 2.5.1 Iron Compounds and Jarosite Passivating Layer The hypothesis that the passivation layer is comprised of iron compounds and jarosites has received less support.82 Both Fe(ll) and Fe(lll) are frequently present in acidic iron sulfate leaching solutions of chalcopyrite. The immediate consequence of ferric precipitation is the reduction of available ferric oxidants in solution. While iron and jarosite compounds are also associated with sulfur layer under some leaching conditions, there is still no direct evidence that CHAPTER 2 LITERATURE REVIEW / 27 8 66 83 84 they alone could be the culprits of the significant decrease in leaching rates. ' ' ' As was discussed in Section 2.2.2, the formation of jarosites is promoted by sulfate ions and high pH values. Other ferric precipitates include ferric hydroxysulfates and ferric sulfates. Although Cooney et al.66 had suggested the role of jarosite as the passivating layer, it was not considered to be a diffusion barrier, but a non-porous layer. Dutrizac and Hirato 6 1 ' 6 5 performed experiments at low pH with 0.3 M and 0.2 M H2SOA, respectively. Although low pH prevented jarosite formation, parabolic kinetics was still observed in these experiments. Consequently, the formation of jarosite as an explanation of the occurrence of parabolic kinetics is not plausible. 2.5.2 Elemental Sulfur and Bimetallic Intermediate Sulfide Passivating Layer Elemental sulfur (Ss) is hardly oxidized at temperatures below its melting point (119°C) during chemical ferric sulfate leaching of chalcopyrite.85 As a result of equation (2-15), a diffusion-limiting elemental sulfur layer is hypothesized by many investigators. They believe that the transport of ions through the porous adherent elemental sulfur product layer on the chalcopyrite surface limits the reaction rate and explains the slow leaching rate of chalcopyrite. 4 9 ' 5 6 ' 5 7 ' 5 9 ' 8 4 As the sulfur layer thickness increases, the rate of dissolution of the chalcopyrite decreases leading to surface passivation. It has been proposed that the diffusion of ferric ion through the elemental sulfur layer is the rate limiting step.54 However, many researchers have noted that initial ferric sulfate concentration (adequate to initiate leaching) has little or no effect on the dissolution of chalcopyrite, suggesting ferric ion diffusion is not limiting as shown in Table 2-6. It is also possible that the rate limitation is due to diffusion of ferrous and cupric ion products rather than by diffusion of the ferric ion reactant.1 The diffusion of electrons through the elemental sulfur layer as the rate-limiting step has also been sugg ested.50 The electron transport limitation theory was tested by leaching chalcopyrite in the presence of finely 86 ground graphite. Since graphite was semi-conductive, it would improve electron transport as it became embedded in the electronic insulating sulfur layer during leaching. The leaching of chalcopyrite was enhanced, lending support to the theory. Nevertheless, the interpretation of the parabolic kinetics as a layer of 5"'apparently cannot explain many experimental facts. A few of the most important ones include: CHAPTER 2 LITERATURE REVIEW / 28 1. A layer of S° is formed on all copper minerals, but only the leaching of chalcopyrite 26 obeys the parabolic law. 2. When the oxidation of one mole of chalcopyrite (42.65 cm ) renders 100% elemental sulfur (31 cm3), the negative volume change indicates a high shrinkage of 27%. 8 7 The sulfur film is then expected to form cracks or a porous morphology so that uniform coatings could not form and interfere with leaching. 3. The removal of S° with carbon tetrachloride ( C C / 4 , a sulfur solvent) in the course of reaction does not improve the rate of leaching. 4. The calculated pore diffusion rates of Fe 3 + ions are four times higher than the measured rate of chalcopyrite dissolution (in an acidic ferric nitrate medium, 25-40°C). 8 9 Although acidified ferric chloride leaching of chalcopyrite in the presence of carbon tetrachloride has shown an increased copper extraction as supported by the third statement90, presumably the presence of complexing chloride anions has affected the leaching rate, and does not suggest that the sulfur deposited during the leaching hindered further dissolution of chalcopyrite. Based on the fourth statement, the passage of ferric oxidants and dissolved metal ions is slow and rate-limiting by the transport processes through the porous sulfur layer; however, the porous layer should not lead to an eventual surface passivation. Linge 8 9 argued that S" cannot control the leaching reaction. He proposed that ferric leaching of chalcopyrite is controlled by the solid-state cation diffusion process; this is a much slower process that also supports the observed parabolic leaching kinetics. This process is further defended by the measured values of activation energies (63-88 kJ/mol), which are too high to be consistent with a pore diffusion process.61'65' According to the model advanced by Burk in 9 1 , and supported by several investigators, the non-stoichiometric diffusion rates of copper and iron in the bulk of chalcopyrite mineral give rise to a compact, bimetallic intermediary sulfide layer, which is able to influence the reaction rate and retard the reaction development. 2 6 ' 9 2 ' 9 3 ' 9 4 jhis intermediate layer grows at the chalcopyrite surface, but is chemically and structurally different from the bulk of CuFeS2. It is proposed that the metal-deficient polysulfide slows the transport of Cu2+ and Fe2+ reaction products. Parker et al.95 has extended the model to include the slow electron transport on oxidants such as Fei+ and Cu2+ as well as electron transfer from reduced species such as Fe2+ on the corroding surface. In CHAPTER 2 LITERATURE REVIEW / 29 this model, the region of oxidation involves both the chalcopyrite surface as well a reaction zone of finite depth. If this model is valid, this reaction zone of finite thickness should be detectable on chalcopyrite leached surfaces. 2.5.3 Silver Catalyzed Ferric Sulfate (Bacterial) Leaching Of Chalcopyrite Several catalysts and/or promoters have been proposed to improve chalcopyrite leaching, such as surfactants 9 6 ' 9 7 ' 9 8 ' 9 9 ' 1 0 0 carbon particles 1 0 1 ' 1 0 2 ' 1 0 3 , iron powder' 0 4, hematite105 and silver ions . Silver ion is the most effective catalyst that has been proposed, but the processes using silver ion are costly unless ores contain silver minerals. However, the mechanism of the catalytic effect of silver ions can contribute to the understanding of the passivation phenomenon. The use of silver catalysis during ferric sulfate leaching was proposed by Miller and Portillo (1981)1 0 6 and later reported by many researchers 1 0 7 ' 1 0 8 ' 1 0 9 ' 1 1 0 ' 1". The mechanism of silver catalysis of chalcopyrite remains uncertain. According to Miller and Portillo, ferric sulfate leaching of chalcopyrite is retarded by a dense elemental sulfur layer formed on the mineral surface during leaching, acting as a diffusion barrier. In the presence of silver ions, however, the enhancement of dissolution is due to the formation of conducting silver sulfide film (Ag2S), which is deposited on the chalcopyrite instead of the film of elemental sulfur. The silver sulfide film is then oxidized by ferric ion to generate silver ions and elemental sulfur, which liberate silver ions for further catalysis. Step 1: Silver ions react with chalcopyrite to form a silver sulfide film on the CuFeS2: CuFeS2 + 4Ag+ -> Cu2+ + Fe2+ + 2Ag2S (2-25) Step 2: Silver catalyst regenerated by ferric oxidation of the silver sulfide: 2Ag2S + 4Fe3' -» 4Ag+ + 4Fe2+ + 2S° (2-26) Overall: CuFeS2+4Fe 3+ >Cu2+ +5Fe2+ +2S o (2-27) The continuous formation and dissolution of the silver sulfide film forms a mixed product layer of elemental sulfur and silver sulfide on the chalcopyrite surface, which has a different CHAPTER 2 LITERATURE REVIEW / 30 morphology than the elemental sulfur layer alone. Since this mixed product layer is porous and non-protective, copper extraction with silver ions is suggested to be higher because it allows electron transfer between chalcopyrite and the solution containing ferric ions. They concluded that the rate is controlled by an electrochemical reaction of Ag2S with Fei+ as shown in step 2 (equation 2-26). This explained their finding that the reaction is dependent on the initial Fe3+concentration but neither the initial Cu2+ concentration nor the initial Ag+ concentration. When a small amount of silver(I) ions (up to about 10 M) is added to the ferric sulfate solution, the rate of chalcopyrite dissolution increases dramatically."1-112 It is important to add the silver to the slurry and allow the silver sulfide layer to form, before adding ferric sulfate. If silver was added after the ferric sulfate, the ferric would attack the chalcopyrite resulting in a passivating layer. Larger silver additions did not give them improved overall extraction. In fact, excess silver impeded the reaction. Moreover, it has been proposed that the nucleation of elemental silver improves conductivity1 1 3, and the coexistence of elemental silver (in a much smaller quantity) with the silver sulfide film on the chalcopyrite surface has been shown by Price and Warren. 1 1 1 Gomez et a/. 1 0 7 suggested that different cathodic zones are created in the presence of a silver sulfide film that facilitates the establishment of reversible redox processes on the chalcopyrite surface, leading to more positive corrosion potentials, and favors the oxidation of chalcopyrite. In opposition to Miller's theory, Parker 1 1 4 argued that Ag2S nuclei form uncoated sites in the passivating sulfide barrier. Unlike the uncatalyzed leach, silver catalyzed ferric leaching of chalcopyrite was improved with increasing concentration of ferric ion, and severely impeded if any ferrous ion was in solution similar to the effect of excess silver addition. The rate of reaction was found to be temperature dependent, and had an activation energy of 66 kJ/mole. Peters and Doyle 1 1 5 theorized that the effectiveness of silver as a catalyst has to do with its high solid-state diffusivity in sulfide minerals. In the case of chalcopyrite, Chen and Harvey" 6 reported that the solid state diffusivity of silver in chalcopyrite is greater than copper and iron self-diffusion above 100°C. Extrapolation of the result shows that the relationship holds at room temperature: Ag>Cu> Fe. The high mobility of silver can easily enter into step 1 and form silver sulfide, which can later be oxidized by the ferric oxidant. Similarly, a silver sulfide film is also observed during silver catalyzed bioleaching of chalcopyrite. Scientists at B.C. Research in Vancouver (British Columbia, Canada) discovered that silver catalyzed bioleaching not only increased the copper leach rates and extractions, but CHAPTER 2 LITERATURE REVIEW / 31 also converted sulfide sulfur quantitatively to elemental sulfur rather than sulfate . 1 1 7 ' " 8 ' 1 1 9 However, it was necessary to mix the concentrate with a solution containing silver, thiosulfate ions and cupric ions prior to bioleaching to ensure elemental sulfur production in high yield at ambient temperature and pressure."8'"9 The chemistry of silver catalyzed bioleaching can be represented by the following reaction:"9 CuFeS2 +02+ 2H2SOA -» CuSOA + FeSOA + 2S° + 2H20 (2-28) Typical bacterial leaching of sulfide minerals expects the oxidation of elemental sulfur to sulfuric acid. The reason for the elemental sulfur production in this case is still uncertain. According to SEM and AES studies on chalcopyrite dissolution in the presence of Ag+ , no significant differences between the leached and bioleached chalcopyrite were found. 1 2 0 Another study showed that the microbial formation of elemental sulfur during silver-catalyzed bioleaching of chalcopyrite increased the silver sulfide film thickness.107 A negative aspect of the process is the toxicity of silver to bacteria. Mesophiles are not only more tolerant of silver than thermophiles, but copper extraction in silver assisted bioleaching systems is much more favorable in the presence of mesophiles than thermophiles. 1 2 1 , 1 2 2 Ahonen et a/ . 1 0 9 ' 1 1 0 have studied the role of silver catalysis on bacterial leaching of chalcopyrite. They explain that the role of ferrous-oxidizing bacteria such as Acidi thiobacillus ferrooxidans is to maintain a favorable ratio of Fe(III) to Fe(II) by oxidizing ferrous ion, and to 29 contribute to the ferric dependent oxidation of silver sulfide [equation (2-18)]. Hackl et al. have suggested the potential inhibiting of bacterial sulfide oxidizing mechanism by silver but not the ferrous oxidation mechanism, so that ferric oxidants are continuously regenerated in an essentially catalyzed ferric leach. These authors also suggested that silver causes the bacterial sulfide oxidation process to cease at elemental sulfur. Sukla et al.123 studied the kinetics of silver catalyzed and uncatalyzed bioleaching of chalcopyrite. They reported that uncatalyzed bioleaching exhibits parabolic kinetics with the likely rate limiting step being the diffusion through a smooth jarosite-sulfur reaction product; whereas the reaction product in catalyzed bioleaching was much more porous, and exhibited chemical reaction control. Their conclusions are similar to the conclusions presented by Miller et al.m for the silver catalyzed ferric sulfate leaching of chalcopyrite. These authors also attributed the catalytic effect to the prevention of the formation of jarosite. CHAPTER 2 LITERATURE REVIEW / 32 2.5.4 Electrochemical Studies of Chalcopyrite Since most metal sulfides are good semiconductors, the same electrochemical techniques developed for metallic electrodes are applicable to metal sulfides. 3 ' 1 3 ' 1 2 4 These techniques study the physical and chemical phenomena associated with electron transfer at the interface of the electrode and the solution. Valuable data regarding the mechanisms and kinetics of chalcopyrite dissolution and some details of their transient behaviors under different acidic oxidizing conditions are obtainable. An excellent review of the electrochemical behavior of chalcopyrite has been written by Hiskey. 1 3 The dissolution of chalcopyrite is slow, even in highly acidic solutions. Aside from the formation of elemental sulfur, many electrochemical studies showed the initial formation of a thin sulfide as the product film on the chalcopyrite surface during dissolution, and this film is agreed to be responsible for its low copper extraction. However, there is disagreement as to its nature and little is known about the electrochemistry involved in its formation. For the purpose of discussion, it is useful to consider the low anodic potential passive region (or prewave) (E < 1 VSHE) and the high anodic potential transpassive region (E >1 VSHE) separately. At low anodic potential regions, McMillan et al.125 concluded that the product film on chalcopyrite electrodes in acidic solutions had the properties of covellite (CuS). The film was referred to as a "Solid Electrolyte Interphase" between the unreacted mineral and the solution, i.e., the layer was said to be electronically insulating, which slowed the rate of electron transport, but allowed the transport of ionic species from chalcopyrite to the solution. Gardner and Woods 1 2 6 reported a mixture of covellite (CuS) and elemental sulfur ( 5 8 ) as the surface products in acidic solution. However, Hamilton and Woods 1 2 7 later reported the surface film to be more consistent with the stoichiometry of CuS2 known in nature as "villamaninite". Biegler et a/. 5 8 ' 1 2 8 conducted extensive studies of the anodic behavior of chalcopyrite in deoxygenated acid solutions ( IM H2S04 , HCl , HNO} and HC104 ) at 25°C using cyclic voltammetry. The observed prewave [0.2-0.7 V S C E or 0.444-0.944 VSHE] at the first anodic sweep was relatively insensitive to the type of acid electrolyte and the specimen origin. The ratio of covellite and sulfur (3:5) was determined based on the relative charges associated with formation and reduction of the product layer. Moreover, the ratio of sulfur to sulfate formed was found to CHAPTER 2 LITERATURE REVIEW / 33 be around 6:1. 5 8 Furthermore, the ratio of copper(II) to iron(II) dissolved in solution was between 1:1 to 1:4. Since iron was known to be selectively leached from chalcopyrite, they postulated the formation of a thin covellite-like passivating film (3 nm) on the chalcopyrite surface, using the upper limit of the copper(II) to iron(II) ratio: CuFeS2 -+0.75CwS + 0.25Cu 2 + +Fe2+ +1.255° + 2.5e~ (2-29) This passivating film was observed to be stable for long periods at open circuit at 25°C as opposed to the reported transient passivating behavior at 80°C observed by Parker et al.95'129 Holliday and Richmond 1 3 0 observed an anodic prewave in the range 0.2-0.4 VSCE (or 0.444-0.644 V S H E ) in acidic solutions at 25°C on the first scan in rotating ring-disc electrode experiments. The observed ratio of iron(ll) to copper(II) was determined to be 5:1 to give the following prewave equation: CuFeS2 0.2Cw2+ + Fe2+ + CuQAS2 + 2Ae~ (2-30) This surface rim was a layer of iron-deficient polysulfide such asCw 0 8 5 2 (a non-stoichiometric villamaninite). Parker et a/. 9 5 ' 1 2 9 studied the potentiostatic oxidation of chalcopyrite and observed a strong passivity at potentials up to 0.7 V S C E (or 0.9 V S H E ) at elevated temperatures (80-90°C). This passivating film was a thermally unstable and metal-deficient, copper-rich polysulfide, which formed from chalcopyrite as the reaction proceeded. The film was suggested to transport ions slowly to the solution and transfer electrons slowly to oxidants, thus accounting for the slow kinetic of the leaching of chalcopyrite. They concluded that the product sulfur layer and copper salts were of little kinetic consequence in the leaching of chalcopyrite. Jones and Peters'31 found that chalcopyrite electrochemistry varied significantly with the specimen origin and the method of preparation. Similar to other investigators, they observed decay curves of current density with time (typical of activation polarization) at relatively low constant potentials, and reported the passivity of chalcopyrite in the potential range 0.6-1 VSHE at low temperatures (6-98°C) and 0.1 M H2S04 (Figure 2-8). However, the passive region gradually disappeared at higher temperatures and disappeared entirely at 175°C, suggesting the inherent instability of the passivating layer at high temperatures. They proposed that covellite (CuS) was responsible for passivation during anodic dissolution of chalcopyrite. The anodic dissolution of chalcopyrite in general was found to be a non-reproducible process, which preferentially corroded deeply in the grain boundaries while the electrode surface remained passive. The C H A P T E R 2 L I T E R A T U R E R E V I E W / 34 continued current decay reflected the formation of a progressively thickening passive surface film. The dependence of current on scan rate in the passive region was also reported, and Peters1 4 7 suggested that passivation is probably a superficial phenomenon. At higher potential [greater than IVSHE], the current density was linear with increasing potential in the transpassive region at the lower temperatures, and the transport of cuprous ions through the mineral lattice was thought to be rate controlling. I O " 4 I O " 3 IO" 2 IO"' CURRENT DENSITY ( A / c m 2 ) Figure 2-8 Effect of temperature on the polarization curve of chalcopyrite (0.1 MH2S04, scan rate 3 0 0 mV/min) 1 3 1 Warren et al. ' studied the passive and transpassive anodic behavior of chalcopyrite specimens in acid solutions at low temperature. The anodic polarization curves for specimens from six different sources in acid solutions at low temperature (\MH2S04, 2 5 ° C ) are illustrated in Figure 2-9. The general shape of the polarization curves is classified into three regions, namely, dissolution, passive and transpassive regions. A l l samples exhibited similar passive-like response during anodic polarization. The "passive" region ranged from 0.6-0.9 VSHE, where the CHAPTER 2 LITERATURE REVIEW / 35 current density suddenly decreased to a very low value that remains independent of potential. Finally, at potential values greater than about 0.9-1 VSHE, the current density again increased with potential in the "transpassive" region. The results of the constant potential experiment (Figure 2-10) further confirmed the passivity of chalcopyrite at potentials less than about IVSHE (\MH2S04, 25°C). Based on current and mass-balance measurements, the surface passivating film was considered a mixture of two intermediate phases, [ Cui_xFe]_vS2_z ( St )] and [ CuSit (5 2)], which formed in the sequence CuFeS2 —» S, —» S2 in the passive region: CuFeS2 -> Cw 1_ vFe l_ v5 2_ z(5 l) + xCu2+ + yFe2+ + zS° + 2(x + y)e~ (2-31) Cu\-xFe\-yS2-z^S\ ^ ~* ( > ' " x ) C u S n (S2) + (] " y)Cu2+ + (1 - y)Fe2+ +(2 + nx - ny - z)S° + 4(1 - y)e~ (2-32) In the first step, equation (2-31), chalcopyrite reacts initially to form an intermediate defect structure (S ]). Since ferrous ions are known to release much more rapidly into the aqueous phase than cupric ions in the initial stages of dissolution 5 9 ' 8 9 ' 9 4 ' 1 3 1, y>x and x: + y = 2 to account for the overall charge balance. In the second step, equation (2-32), theS, intermediate decomposes further to form a second intermediate (S2) as x approached y. The release of Cu2+ and F e 2 + i n equal amounts was consistent with the results of Baur et al.59 The S| and S2 were speculated to be iron-depleted sulfides such as bornite (Cu5FeS4) and covellite (CuS), respectively. In the transpassive region (1.0-1.2 VSHE), two overall anodic oxidation reactions of chalcopyrite were reported as shown below: CuFeS2 -> Cu2+ + Fei+ + 2S° + 5e~ (2-33) CuFeS2 +8H20-> Cu2+ + Fei+ + 2S02' +\6H+ +17e" (2-34) According to the amount of cupric ion and the charge passed, 90% of the sulfide sulfur converted to elemental sulfur and 10% to sulfate for equations (2-33) and (2-34), respectively. They also confirmed that the passive film which formed at low potentials was electron-conducting as observed by Ammou-Chokroum et al.m. An addition of ferrous ions drastically increased the current during anodic polarization, most likely due to the oxidation of ferrous at the chalcopyrite surface. Since sulfur is an excellent electronic insulator at room temperature, i f the electrode were passivated by elemental sulfur, it would not be expected to respond to the additions of ferrous sulfate.53 CHAPTER 2 LITERATURE REVIEW / 36 -i 1 r CURRENT OENSITY ( A / c m * ) Figure 2-9 Anodic polarization curves for CuFeS2 from six different locations (1 M H2S04, scan rate 30mV/min, temperature 25°C) 8 0 j i i i— 10 20 30 40 Time (min) Figure 2-10 Current density versus time plot for Transvaal CuFeS2 at various anodic potentials (1 M H2S04, temperature 25°C) 8 0 CHAPTER 2 LITERATURE REVIEW / 37 Although the leaching of chalcopyrite in the passive region involves the dissolution of iron as ferrous ions, dissolution reactions involving ferric ions have also been published in the 1 3 3 literature. The rotating split-cylinder bi-electrode measurements of Kelsall and Page suggested the following reactions for the electro-dissolution of chalcopyrite, which produced mainly ferric ion and little accompanying cupric ion. Step 1 : 5CuFeS2 -» Cu5FeS4 + 4 / V + + S° + 12e" (2-35) Step 2: ' Cu5FeS4 -> 2.5Cu2S + Fe3+ + 1.55° + 3e~ (2-36) Another case that involved the dissolution of ferric ion in the passive region was provided by Stankovic134, who examined the prewave region using galvanostatic pulse chronopotentiometric techniques in solutions of Fe(II) and Cu(Il) . That author proposed a two-stage oxidation process, leading to the formation of an iron polysulfide ( Cun_]Fell_]S2n ). The same non-stoichiometric sulfide was reported by Arce and Gonzalez135, and no covellite was formed in either cases. The first stage involved the liberation of Cu2+, which was rate-limiting: Stage 1: nCuFeS2 —> Cu2+ + Cun_xFenS2n + 2e~ (slow and rate-limiting) (2-37) The second stage involved the formation of Fe 3 + as shown below: Stage 2: Cun_,FenS2n -> Fe3* + CuH_Fe„_xS2ll + 3e~ (2-38) CHAPTER 2 LITERATURE REVIEW / 38 2.5.5 Surface Studies (AES and XPS) Chalcopyrite is a ternary sulfide with a distinctive surface chemistry. It possesses surface geometric structures that exhibit different properties from the bulk solid. 1 3 6 AES (Auger Electron Spectroscopy) and XPS (X-ray Photoelectron Spectroscopy) are the two most popular surface-sensitive spectroscopic techniques for sulfide minerals.1 3 6 AES is used to determine the elemental composition of chalcopyrite surfaces; whereas XPS is used to determine the oxidation state as well as the elemental composition of a surface species. XPS interpretations of the initial, intermediate and final oxidation products of (bio)leaching of chalcopyrite have been used to support leaching rate and yield measurements.79 The term 'surface' actually means the near-surface region, extending from the top monolayer to a typical depth of 1 -10 nm. Surface studies of chalcopyrite in different media and experimental conditions can be used as supplementary tools to reveal insights to the tentative nature of the so-called "passivating" layer, responsible for low yields of copper extraction from chalcopyrite. 2.5.5.1 Air Oxidation of Chalcopyrite AES study of chalcopyrite surfaces was pioneered by Eadington. 1 3 7 Exposing a fresh fracture surface for 30 minutes to air and water resulted in a thin layer of 0.9 nm and a thick layer of 9.5 nm of oxidation products, respectively. He concluded that the oxidation of a chalcopyrite surface when exposed to water was about 10 times greater than when exposed to air. Brion 1 3 8 investigated the oxidation of air-ground chalcopyrite particles by XPS. He reported the initial formation of iron hydroxide/oxyhydroxides on the surface, which subsequently transformed into basic iron sulfates (jarosite). The copper bonded to the sulfur remained unoxidized. The broadening of the S (2p) spectra toward higher binding energy was interpreted as disulfide (S2) group formation and indicated the presence of an iron-depleted copper sulfide lattice. No elemental sulfur was observed in this study. The XPS study of chalcopyrite (CuFeS2) fracture surfaces exposed to air by Buckley and Woods 1 3 9 was consistent with Brion's XPS study1 3 8. They concluded that in air, iron atoms migrated to the surface to give an overlayer of ferric oxyhydroxides, and neither sulfur nor CHAPTER 2 LITERATURE REVIEW / 39 copper atoms are involved in the initial oxidation. The remaining copper bonded to sulfur as Cu(I), resulting in an iron-deficient sulfur-rich sulfide lattice with composition close to CuS2: CuFeS2 + 0.15xO2 +1.5xH20 -> CuFe^S2 + xFe(OH)3 ( 2 " 3 9 ) Where: x ~ 1. Since the values for the S (2p) binding energy does not resemble the values for artificially synthesized CuS2, which has the pyritic S2~ groups, the iron-deficient sulfide was concluded to retain the original chalcopyrite lattice structure. Subsequently, the CuS2 species was oxidized to sulfate after prolonged exposure (>10 days), probably as copper(II) sulfate, and no elemental sulfur formation was observed in this study.43 Ruzakowski et al. employed a combined AES, XPS and optical microreflectometry to study the air oxidation of the polished chalcopyrite surface at temperatures between 23°C and 300°C. 1 4 0 Although exposure times varied from as long as 1.5 years to only 1 to 2 hours at 300°C, three layers were found to form above the chalcopyrite bulk solid in all cases. Based on the reflectance data, these authors reported that the outermost surface layer was most consistent with magnetite (Fe 3 0 4 ) , followed by a layer of chalcocite (Cu2S) and a subsequent layer of bornite (Cu5FeS4) on top of chalcopyrite surface. No elemental sulfur formation was reported in this study. The calculated total film thicknesses were 10 to 35 nm at 23°C, 8 to 51 nm at 200°C and 12 to 85 nm at 300°C. Balaz et al.m reported elemental sulfur formation on chalcopyrite particles that were ground ultra-fine to a mean grain diameter (dso) of 2 microns. These authors attributed the higher proportion of sulfur oxidation to surface oxidation, and reported the existence of sulfur in three different chemical forms based on the analysis of the S(2p) spectrum: sulfide (S2~), elemental sulfur (5° ) and sulfate ( 5 6 + ) . 2.5.5.2 Oxidation in Alkaline Solution (Studies Related to Flotation) Collectorless flotation of chalcopyrite in alkaline solution (pH 10.5) with controlled-potential was studied by Luttrell and Y o o n 1 4 2 using XPS. They detected covellite ( CuS ), chalcopyrite and sulfur oxidation on the surface to either an elemental sulfur or polysulfide CHAPTER 2 LITERATURE REVIEW / 40 species. Since polysulfides can be proposed as intermediates for elemental sulfur formation , Hackl 2 9 has suggested the following reaction after considering the reaction proposed by Gardner and Woods' 2 6: Although S" to CuS ratio was much less than 1, elemental sulfur was suggested to be lost by volatization in the spectrometer vacuum. Buckley and Woods 1 4 4 investigated the oxidation of chalcopyrite (CuFeS2) fracture surfaces in air-saturated dilute ammonia solution, which gave results similar to those for air as mentioned above. The copper remained bonded to sulfur, and an iron oxyhydroxide appeared. However, elemental sulfur was observed along with an altered copper sulfide layer (possibly 1 ^9 CuS). The altered copper sulfide layer was later proposed to be Cu0 SS2. - They considered the layer to be chalcopyrite with a layer of copper and iron deficiencies or a metal-deficient copper sulfide. Vaughan et al.136 used a combination of XPS, XAS (X-ray Absorption Spectroscopy) and electrochemical methods to establish a sequence of oxidation processes for the oxidation of chalcopyrite electrodes in alkaline solution (pH 9.2) at 25°C. These authors proposed that a monolayer of Fe2Oz IFe(OH)iwas initially formed leaving copper and sulfur unoxidized in the original chalcopyrite structure as a metastable phase CuS*2 .Together these phases were then responsible for the passivation of the chalcopyrite surface in alkaline solution. Iron diffused from within the chalcopyrite bulk solid with increasing potential and at a rate controlled by solid-state diffusion. Above a critical (pH-dependent) potential, the passivating CuS2 layer was decomposed as shown below: CuFeS2 + 3H2Q -> Fe{OH)3 + CuS + S° +3H + 3e (2-40) 2CuFeS2 + 3xOH —> 2CuFe}_xS2 + xFe203 + 3xH+ + 6xe (2-41) 2CuFeS2 + 60H~ -» 2CuS'2 + Fe2Oy + 3H20 + 6e CuFeS2 + 30H~ -» CuS*2 + Fe(OH), + 3e~ (2-42) (2-43) CuS, + 20H~ -» CuO + 2S + HnO + 2e (2-44) CuS*2 + 2QH' -> Cu(OH)2 +2S + 2e (2-45) CHAPTER 2 LITERATURE REVIEW / 41 2.5.5.3 Oxidation in Acid Solutions: Studies Related to (Bio)Leaching 8 A new theory in the mechanism of passivation was proposed recently by Hackl et al. in their study of the oxygen pressure leaching of chalcopyrite in the temperature range 110-220°C. Surface analysis of leached residues was performed using AES and XPS. Leaching was carried out at 110, 155 and 200°C. In all cases, the compositions of leached surfaces deviated from stoichiometric CuFeS2, and showed little difference at different leach temperatures. AES analysis showed that the leached surfaces displayed an elevation with respect to copper and sulfur, but a reduction in iron (Cu/Fe atomic ratio 1.6-1.7:1). This indicated that iron leached preferentially to copper to form an iron-deficient sulfide on the chalcopyrite surface. No significant amount of iron hydrolysis products such as ferric hydroxide, hematite or basic ferric sulfate were found on the leached surfaces. The broad S (2P) spectra from XPS were interpreted in terms of five sulfide species: 68% monosulfide ( S 2~ , unaltered chalcopyrite), 18% disulfide (S 2 ' , iron deficient, altered sulfide), 10% polysulfide (S]~ ,CuSn) and 4% elemental sulfur (S°). A complicated reaction sequence was devised to account for the results. It was proposed that an iron deficient sulfide (Cu]_xFe]_vS2) formed initially, and subsequently broke down slowly to form a copper polysulfide ( C u S n ) , which was alternatively expressed as C M , _ v _ 2 5 , in the reaction sequence. The copper polysulfide broke down further to release cuprous ions and form porous elemental sulfur. CuFeS2 -> Cul_xFe]_vS2 + xCu2+ + yFe 2+ + 2(x + y)e~ (fastest) • (2-46) Cux_xFex_vS2 -> C M , _ _ W S 2 + zCu 2+ + (1 - y)Fe 2+ + 2(z + 1 - y)e~ (slow) (2-47) Cul_x_zS2 -> (1 - x - z)Cu 2+ + 2S° + 2(1 - x - z)e~ (slowest) (2-48) Where: y » x,x + y ~ I. A thin film of copper polysulfide ( C u S u , where n>2) (less than 1 p:m thick), which formed as a result of solid-state change, was proposed to passivate chalcopyrite. The slow decomposition of CuSn was treated as the rate-determining step, and the decomposition rate increased with CHAPTER 2 LITERATURE REVIEW / 42 increasing temperature up to 200°C. The leaching kinetics at 110°C can be explained in terms of a mixed diffusion/chemical reaction model, where the diffusion through the thickening passivating layer is parabolic and the eventual decomposition of the passivating layer is linear (or chemical). Diffusion through the elemental sulfur layer was thought to be non-rate-limiting, and the presence of sulfur masked the real passivating layer. The presence of elemental sulfur, which is ubiquitous in chalcopyrite leaching, could then be construed as supporting evidence for the interim presence of the copper polysulfide passivating layer. This is supported by Steudel 1 4 5 who proposed the formation of elemental sulfur via the intermediary polysulfides. These authors have suggested that this model can be equally valid during ferric sulfate leaching and bacterial leaching of chalcopyrite in sulfate media as a result of similar chemistries operative in oxygen pressure leaching. Sulfur speciation on chalcopyrite surfaces during the initial stage of ferrous and feme sulfate leaching in acidic solution (pH 1.3-1.9) at 50°C was examined recently by Klauber et al19 Leaching was carried out in all cases for 120 minutes. Similar surface speciation was observed and examined by XPS for both ferrous and ferric leaching. Neither severely metal-deficient sulfides nor polysulfides were found to dominate the leached surfaces. Polysulfide with a chain length greater than two was evidenced in acidic ferrous leaching, but not in the ferric leach. The primary surface species was elemental sulfur, followed next by disulfide (S2~). Although both acidic ferrous and ferric leaches produced elemental sulfur and disulfide on the surface, a greater amount of elemental sulfur was produced during the ferric leach. The cation association of the S22~was not identified, but the formation of any CuS2 type species was discounted. De Filippo et al. 1 4 6 characterized the surfaces of chalcopyrite flotation concentrate before and after bioleaching by XPS quite simply in terms of sulfide and sulfate. These authors observed a decrease in ratio of copper to iron on the surface of bioleached chalcopyrite, and attributed this copper-deficient layer as a direct proof of a passivating layer with the formulation of Cw,_vFe,_i S2, where y < x. However, the XPS analysis indicated the presence of sulfate on the residue surface, indicating the probability of jarosite precipitation. This decrease in the copper to iron ratio may simply due to the presence of jarosite coating the leached surface. CHAPTER 2 LITERATURE REVIEW / 43 2.5.5.4 Passivating Layers and the Limitation of Experimental Techniques Chalcopyrite can undergo significant distortion in its surface region during ferric sulfate leaching, which imparts the passivating phenomenon. The retardation of leaching by S° , precipitated iron compounds and jarosite formation is less likely. The elemental sulfur layer is readily oxidized in bioleach systems. Even when S° is continuously oxidized to sulfuric acid by the catalytic action of bacteria, the bioleaching of chalcopyrite is equally slow, especially during the indirect leaching mechanism. A general agreement can be found in surface studies of oxidized and leached residues of chalcopyrite as well as electrochemical studies with regard to the formation of a copper-rich sulfide layer, which is chemically and structurally different from the chalcopyrite underneath i t . 8 ' 2 6 ' 9 4 ' 8 9 ' 9 1 Various phases, such as Cu0gS2 and Cu]_xFe{_YS2, have been suggested to be potentially responsible as surface passivating agents during leaching of chalcopyrite. The formation of such metal-deficient sulfide or metal polysulfide films at the chalcopyrite surface can be explained by the process of solid-state diffusion. The passivating layer possibly consists of a mixture of several precipitates or a combination of precipitates and altered phases. Silver catalyzed bacterial leaching and ferric sulfate leaching of chalcopyrite yield much higher copper extraction under acid oxidizing conditions at low to moderate temperatures, i.e., temperatures below the sulfur melting point of 119°C. 2 9 Although the silver catalyzed mechanism is not well understood, the quantitative transformation of sulfide sulfur in chalcopyrite to silver sulfide has important implications for the passivation of chalcopyrite. Since all the sulfide sulfur is tied up by silver ions (by forming silver sulfide, Ag2S ), sulfide sulfur is therefore not available to participate in the formation of a passivating layer, whether the passivating layer is iron-sulfide or copper-sulfide in nature. A major contribution of the electrochemical studies of chalcopyrite has been to provide evidence in support of the hypothesis of a potential-related surface passivation mechanism to explain leaching kinetics. Chalcopyrite dissolution is controlled, at least initially, by surface phenomena. Anodic polarization studies on chalcopyrite electrodes have revealed the existence of a passive region where the current does not increase with increasing potential. This passive region results form the formation of a passivating film, which prevents electrochemical reactions CHAPTER 2 LITERATURE REVIEW / 44 from occurring on the chalcopyrite surface. The thin metal-deficient sulfide film is produced in the low anodic potential passive region (E < 1 VSHE) at sub-boiling temperatures (6-98°C) on the chalcopyrite surface, and it is chemically and structurally different from chalcopyrite.131 The ratio of dissolved copper to dissolved iron is in the range of 1:1 to 1:5, indicating the preferential dissolution of iron from the chalcopyrite lattice. The nature of the passivating film is classified into two general groups, which include either iron-deficient and copper-rich polysulfides or non-stoichiometric sulfides. These non-stoichiometric sulfides support the solid-state cation diffusion control of chalcopyrite dissolution as postulated by many researchers previously mentioned in Section 2.5.2. 8 9 ' 9 3 ' 9 4 However, Warren et a/. 8 0 ' 1 3 2 suggested that the non-stoichiometric intermediate phases (5, and S2) could be bornite (Cu $FeSA) and covellite (CuS). However, bornite and covellite are much more readily oxidizable than chalcopyrite, so their roles as the passivating layers are highly unlikely.8 Several different routes for initial oxidation and product layer formation are summarized in Table 2-8. There are important differences between the dissolved iron that leaves the mineral as either Fe(JJ) or Fe(lII) . Most studies report the dissolved iron asFe(IJ). Nevertheless, Stankovi 1 3 4 and Kelsall and Page 1 3 3 report the dissolved iron as Fe(lll), indicating a different kind of electrochemical mechanism that is associated with chalcopyrite oxidation than those that have dissolved iron as Fe(ll) . The differences in electrochemical mechanisms could be associated to the other bonding model of chalcopyrite " Cu2+Fe2+ (S2~)2 "(see Section 2.1), indicating the reduction of lattice Cu(II) and oxidation of latticeFe(II). Studies of the oxidation of chalcopyrite surfaces in air, acid and alkaline solutions provided by AES and XPS seem to compliment the electrochemical studies with regards to the formation of a copper-rich sulfide surface layer. However, the limitations of AES and XPS techniques have to be realized while considering the validity of these surface studies. Since AES and XPS are ex-situ techniques, their applications are only accurate to a certain degree, and may not reflect the true identity of the passivating surface in the unperturbed leaching conditions. Moreover, application of the AES technique is rather limiting in determining the passivating layer, since Auger intensities only indicate the total element concentration, rather than distinguishing the leached surface from the underlying unaltered mineral.7 9 Besides, there is probably also an altered phased between the leached residue and unleached mineral, which makes any identification rather difficult and without absolute validity. On the other hand, CHAPTER 2 LITERATURE REVIEW / 45 although XPS is a powerful technique for understanding surface speciation, interpretations can often be ambiguous, especially with sulfur speciation such as the overlapping states as in the S (2p) spectrum79, and with similar peak shifts136. Two types of elemental sulfur are capable of forming during leaching. It was well demonstrated that 5° will vaporize unless the sample is cooled to < 200 K . 7 9 At that temperature, the difference between sub-monolayer sulfur (attached to the mineral surface as S [0]), which does not vaporize, and multi-layer sulfur (attached to other sulfur molecules as 58°), which does, cannot be differentiated. The direct comparison between leaching data and electrochemical measurements are difficult. Since the electrochemical experimental techniques investigate the initial phase of surface reactions in oxidation, the associated initial leaching rate in these experiments tends to delineate the initial stages of reactions during oxidation. These stages are often rapid, and the reaction time is often measured in seconds. Unlike electrochemical studies that are often only a few hours, leaching studies can carry on for days and even weeks. The ferric sulfate leaching study by Jones and Peters53 lasted over 55 days. The estimation of initial rate from leaching data is extremely difficult, where the reaction time is often measured in hours, or in the very least, half an hour. However, a potential-relating surface passivation phenomenon of chalcopyrite leaching is apparent from the electrochemical studies, and the existence of a passivating layer on the chalcopyrite surface in the potential passive region (E < 1 VSHE) at low temperature (6-98°C) is definite. This metal-deficient polysulfide layer is not only extremely thin, but also chemically and structurally different from chalcopyrite. CHAPTER 2 LITERATURE REVIEW / 46 tu X IS) > V w c o '5b — C D > 03 & C D -g <+-. O c o 3 C O o (/) bO 03 O m '55 c C D s-o3 3 00 O O I C D X> f— ^ 4 (U Oi pCo a ^ cj a _o o 03 C D ai C D o Pi o o 00 ( N CO in C N + C N + + a O C N O + C O a in r-©• t CO a ( J S C C o o o m + + CJ a CJ C N H + + t o DC co~ o S \ o t o " + + + a -+ O a T i C N O T co a CJ co CJ 83 m CN 'co + H C N + t o + CO a C J I + o C O Nl I + CN + CN + CN s <J CX OO O 0 EX C D I C O K C O s CJ I C O N I CN ?• <? H I CJ c o a cj c o co c = J , tO CO a ro io CN CN CO C N + c CO + + Co C o .CO ~ CJ CO CO C O k , a CJ C N T T CO Co <u. rco k , k , ° c l UT) CJ J : r4 cx C D cx C D OO 0 0 *3 co „ Cj ro 2 a 03 o CO C N C O a CJ CO m co a cj CJ t 1 Co CO ,C0 £ ^_ a -CJ s a CJ cx cx C D O 0 0 0 C O C O k . C O a CJ — 03 a CJ a cj CHAPTER 2 LITERATURE REVIEW / 47 2.6 Electrochemistry Most copper sulfides are semiconductors that are essentially covalent in nature. The distinctive delocalized electron behavior of most copper sulfides should, at the very least, be partly responsible for the electrochemical nature of sulfide leaching.1 4 7 Unlike chemical redox reactions, charge transfer in electrochemical reactions does not occur at the same physical position but rather the half-cell reactions are separated by some finite distance. The extensive covalent bonding nature of metal sulfides provides delocalized electrons that are free to move throughout the lattice of metal sulfides, and results in appreciable intrinsic electronic conductivity; thus, most semiconducting sulfide minerals have sufficient conductivity to permit electrochemical reactions on their surfaces.1 4'1 4 8 K o c h 1 4 9 has reviewed the electrochemistry of sulfide minerals and Shuey 1 5 0 has reviewed the electronic properties of both oxide and sulfide minerals. The resistivities, electronic, and structural information for selected sulfide minerals are listed in Table 2-9. These sulfide minerals can be arranged in an electrochemical series based on their electrical rest potentials (Table 2-10). The electrical rest potential is the open-circuit potential of minerals measured against a suitable reference electrode in solution, which corresponds to the equilibrium (no net anodic or cathodic current) electrode potential.151 The characteristic "rest" potential is a fundamental property that explains the mineral oxidation sulfide minerals. The rest potential of chalcopyrite in acidic solutions has been found to be between 0.45-0.55 VSHE- Since semiconducting sulfide minerals would typically be at a rest potential lower than that of solution, the redox reaction between a soluble cationic oxidant such as ferric ion and a solid, semiconducting, metal sulfide minerals results in electrochemical reaction that would cause mineral oxidation to be coupled to reduction of the ferric sulfate oxidant. CHAPTER 2 LITERATURE REVIEW / 48 Table 2-9 Summary of electronic properties of selected sulfide minerals Mineral Resistivity (Q-m) Semiconduction Type Structure Ionic Structure Pyrite 3xl(T 2-lxl(r 3 n,p Cubic. Fe2+S22~ Chalcopyrite 9xl0~3-2xl0~4 n Tetragonal Cu+Fe"+(S2~)2 Chalcocite 4xl0~2-8xl(r5 P Orthorhombic (Cu+)2S2-Covellite 8xl0" 5-7xl0" 7 Metallic Hexagonal (Cu+)2S2~ Galena Ixl0~5-7xl0~6 n,p Cubic Pb2+S2~ Sphalerite 3xl0~ 3-lxl0~ 4 - Cubic Zn2+S2~ Bornite io- 3 - io - 6 P Tetragonal (Cu+)5Fei+(S2~)4 Table 2-10 Standard electrical rest potential of selected base metal sulfides Metal Sulfide Chemical Structure Rest Potential (VSHE) Measuring Conditions Ref. Pyrite FeS2 0.63 1.0 mo\/L H2S04; 25°C 14 Chalcopyrite CuFeS2 0.52-0.53 1.0 mol/L H2 SO,; 20°C 152 CuFeS2 0.52 1.0 mol/L 7 / 2 50 4 ; 20°C 14 CuFeS2 0.45-0.55 1.0 mol/L H2SOA; 25"C 80 Chalcocite Cu2S 0.44 1.0 mol/LH 2SO 4;20°C 14 Covellite CuS 0.42 " 1.0 mol/L HC104; 25°C 14 Galena PbS 0.28 1.0 mo l /L / / 2 S0 4 ; 20°C 14 Sphalerite ZnS -0.24 1.0 mo\/LH2S04; 20°C 14 Pyrrhotite FeS -0.28 1.0 mo\/LH2S04; 20°C 14 CHAPTER 2 LITERATURE REVIEW / 49 2.6.1 Mixed-Potential Theory Several theories have been investigated to explain the kinetics of metal sulfide dissolution, including the well-known mixed-potential theory, the recently applied semiconductor electrochemistry and the solid-state electronic structure theory. 1 3 ' 1 4 7 ' 1 5 3 ' 1 5 4 ' 1 5 5 Among them, the mixed-potential theory of metallic corrosion has been successfully applied to explain dissolution kinetics of a number of hydrometallurgical systems, including the ferric sulfate leaching of chalcopyrite. 5 0 ' 5 3 ' 8 0 ' 1 3 2 ' 1 5 6 Both chemically and biologically catalyzed oxidation reactions can be considered to be corrosion systems in which the anodic oxidation of metal sulfides and the cathodic reduction of the oxidant (e.g., ferric ion or oxygen) occur on the mineral surface. 1 5 7 ' 1 5 8 The concept of mixed potential can be illustrated schematically (Figure 2-11). The dissolution of a metal sulfide (MS) is taking place in the presence of a cationic oxidant N"+ (such as ferric ions in acid solutions), which has a more positive equilibrium potential. The ferric leaching of chalcopyrite can be represented by two half-reactions as shown below: CuFeS2 <=> Cu2+ + Fe2+ + 2S° + 4e~ (2-49) Fe'++ e~ <^ Fe2+ (2-50) The dependence of the equilibrium (or rest) potential on the concentration of the reactant and product species in solution as predicted by the Nernst equation are shown below, where E" and E are the values of the standard reduction potential and formal potential, respectively. 4T a y, 4T [re(W)\ F -F° +L-L]n-l£--F" • ^ [FejIJl)] A l l measurements are referred relatively to the standard hydrogen electrical (SHE) potential, which is set to zero at all temperatures. The value of £°for Fe3+'2+ is 0.771 V S H E (or 0.572 V A g / A g C i ) a n d the value of the formal potential (E ) in 0.1 kmol/m3 sulfuric acid solutions is about 0.67 V S H E (or 0.471 V A g / A g c i ) at 25°C. l 5 9 At temperatures other than 25°C, the standard CHAPTER 2 LITERATURE REVIEW / 50 reduction potential ( E° ) for the ferric/ferrous couple is approximated via the "entropy correspondence principle" developed by Criss and Cobble (Appendix A-12) . ' 6 0 ' 1 6 1 Using the Criss and Cobble theory, the E° values for ferrous and ferric ions at 60, 70 and 78°C are shown in Table A-5. The mixed potential (Em ) lies somewhere between the two equilibrium half reaction potentials. The exact position of the mixed potential is a function of the magnitude of the two individual overpotentials: anodic overpotential (?]a) and cathodic overpotential (rjc). The anodic overpotential (rjn) is positive, and is defined as (Em -Ee]). The cathodic overpotential (r/c) is negative and is defined as (Ein — Ee2). Overpotential (rj) defines the necessary potential relative to the equilibrium electrode potential necessary to drive a particular reaction at a given rate, producing zero net current in the process. The rate of the reactions is exponentially dependent on the potential across the mineral-solution interface. The charge transfer process in a simple reversible reaction may be expressed quantitatively by the Butler-Volmer equation, which relates the current density at the mineral-electrolyte interface to the established overpotential (for a single electron transfer process):147 / = ; 0jexp| where /„ -pry RT exp RT (2-53) is identified as the "exchange current density," i.e., the implied current density in each direction at zero current. (3 is the fraction of polarization potential rj that affects the activation energy in the cathodic direction. (1 - /?) is the fraction that affects the anodic direction. The standard potential of equation (2-52) of iron coupling and equation (2-51) of chalcopyrite at 25°C and 90°C is shown below in Table 2-11 as presented by Jones and Peters.53 It can be seen that the typical mixed potentials of the ferric sulfate leaching of chalcopyrite lie approximately midway between these standard potentials, and that the equilibrium potential ( Ee2) of the Fe(IIl)- Fe(ll) couple is greater than the mixed potential ( Em ). This explains why hydrometallurgical solutions have measured potentials in the order of ~ 0.6 - 0.8 VSHE-CHAPTER 2 LITERATURE REVIEW / 51 Table 2-11 Standard potentials of iron coupling and chalcopyrite half reactions at 25 and 90°C 5 3 Reaction 4c(VSHB) £ 9 ° 0 . C ( V S H E ) (2-51) 0.44 0.50 (2-52) 0.77 0.85 t N , n + 1 ) + = N n + + e- | E m = Mixed Potential Figure 2-11 Current density potential diagram showing curves for metal sulfide (MS) with equilibrium (or rest) potential ( E e i ) and redox-ions { N ( n + l ) / n + } with equilibrium (or rest) potential ( E C 2 ) , which forms a mixed potential ( E m ) 1 6 2 CHAPTER 2 LITERATURE REVIEW / 52 2.6.2 Redox-Controlled Potential Leaching Studies A number of recent publications have focused on the use of redox potentials as a monitor of the rate of reduction of Fe(III) and as a means of controlling the oxidation rate of minerals by controlling the redox potential.1 iMi9,i<a,i64,i65,i66,i67,i68 M a n y m o d e ] s h a v e b e e n p r 0 posed for the enhanced oxidation of sulfide minerals under controlled-potential. The simplest model assumes a linear relationship between the redox potential of the solution and the mineral rest I f>9 73 77 potential. Other postulated models include those based on the Monod equation and electrochemical theory 1 6 4 ' 1 6 6. Boon 7 3 ' 7 7 measures the rate of bacterial growth and pyrite oxidation in batch cultures of Leptospirillum-like bacteria, using off-gas analysis of carbon dioxide and oxygen. A Monod-type equation has been suggested to describe ferric leaching of pyrite: max / e 5 ' (2-54) [FeS2] j [Fe»] The quantities vFeS and B are kinetic constants, where vFeS (mol FeS2lmo\ FeS2/h) is defined as the pyrite specific rate. Thus, the ferrous-ferric iron ratio is the rate-determining variable, and can be expressed with the Nernst equation: [Fe3+] TT = e x P [Fe2+] F - F RT nF (2-55) Electrochemical (corrosion) phenomena have been suggested for ferric leaching of pyrite 1 6 6, arsenopyrite167 and sphalerite164. The rate expressions were derived by applying the Butler-Volmer equation to the anodic and cathodic processes. The Butler-Volmer expression describes the relationship between the rate of an electrochemical reaction and its thermodynamic and kinetic parameters. The leach rate of pyrite as predicted by the Butler-Volmer equation is given as: 1 7 0 - 'res, = >;,(eapu:-n - e < ( 2 - 5 6 ) The ferric leach kinetics of arsenopyrite can also be described by a rate equation of this form. 167 Verbaan and Crudwell 1 6 4 derived a model in which the corrosion (or mixed) potential can be approximated by the redox potential for the ferrous-ferric iron ratio. The charge-transfer CHAPTER 2 LITERATURE REVIEW / 53 coefficient of 0.5 in their rate equation implies a half-order dependence on the concentration of ferric ions at constant ferrous ions. The rate of reaction ( r ) is given by r = *exp(0.5—) (2-57) RT Where: E = E+(—)ln^^ (2-58) 0 F [Fe2+] 73 77 166 167 These models suggest that the ferric leaching kinetics of pyrite ' ' , arsenopyrite , chalcopyrite 1 6 3 ' 1 6 8 and sphalerite164 may be strongly dependent on the ferric-ferrous iron ratio (i.e., redox potential), and not a function of the total iron or ferric ion concentrations. The redox potential of chalcopyrite leaching has been controlled either chemically with the addition of reagents (e.g., potassium permanganate solution) or biologically (a separate two-stage process) in a set range of values to enhance copper extraction. Natarajan171 examines the effect of DC potential on the dissolution of sulfide mineral in the presence of Acidi thiobacillus ferrooxidans. The maximum copper and iron dissolution rates from chalcopyrite occur at 0.600 and 0.650 V S C E , respectively. Other scientists discovered that copper leach rates and extractions are enhanced by controlling the redox potential of the slurry between 0.340-0.460 V A g / A g C i via the addition of thiosulfate and cupric ions prior to bioleaching."8 Their process has been applied to silver catalyzed bioleaching of chalcopyrite as a means of oxidizing sulfide sulfur stoichiometrically to elemental sulfur at ambient temperature and pressure."8'1 1 9 Ahonen and Tuovinen 1 7 2 studied the mixed mesophile bioleaching of pyrite-rich chalcopyrite ores, by controlling the oxygen supply to monitor redox potential. These authors reported an increase of 5-10% in copper extraction at lower redox potentials (0.500-0.550 VSCE) than at higher redox potentials (600-650 V S C E ) - Although a galvanic coupling effect between pyrite and chalcopyrite was proposed to be the major influencing factor on the improved chalcopyrite dissolution at lower redox potential, improved leaching at low redox potential simply due to a decrease in the ferric-ferrous ratio might also explain the result. Therefore, it is reasonable to speculate that the rate of chalcopyrite oxidation with ferric ion's is higher in the presence of an adequate amount of ferrous ions than in the absence of ferrous ions as suggested by studies discussed above. Commercial bioleaching processes such as the silver catalyzed IBES process " and Mintek's Indirect Bioleaching Process 1 7 4 are essentially two-stage bioleaching processes, which are particularly amenable to redox control. In these processes, the continuous regeneration of ferric ions by bacterial ferrous oxidation is separated from the oxidation of the sulfide mineral. Redox-CHAPTER 2 LITERATURE REVIEW / 54 controlled bioleaching of chalcopyrite is an integral part of a recent patent by Mintek. The Mintek bioleaching patent168 of chalcopyrite and the work of Kametani and A o k i 1 6 3 are discussed in detail below. These studies reveal the ferric-ferrous ratio dependency of ferric sulfate leaching of chalcopyrite under constant redox potential. CHAPTER 2 LITERATURE REVIEW / 55 2.6.2.1 Redox-Controlled Leaching of Chalcopyrite (Kametani and Aoki) Kametani and A o k i 1 6 3 investigated the effect of solution redox potential on chalcopyrite leaching in the presence of ferrous and ferric ions at 90°C and 1 MH2S04 . Leaching was maintained in the potential range 0.30-0.65 VSCE (0.500-0.850 VSHE) by controlled addition of potassium permanganate solution. The leaching rate was monitored and followed by the addition of potassium permanganate solution, rather than copper extraction. The potential range 0.30-0.43 VSCE (or 0.500-0.630 V S H E ) and 0.47-0.64 V S C E (or 0.670-0.840 V S H E ) are referred to as the first and second stage, respectively. Figure 2-12 depicts the leaching curves of the first stage reactions. It shows clearly that at potentials below 0.37 VSCE (or 0.570 VSHE) the leaching rate increases rapidly with increasing potentials but decreases above a certain critical potential. This critical potential is in the potential range of 0.40-0.43 VSCE (or 0.600-0.630 VSHE) at which the leaching rate is maximum, and the reaction curves are practically identical. Below 0.33 VSCE (or 0.530 VSHE), covellite is detected in the leach residue, which is suggested to form directly from chalcopyrite: CuFeS2 <=> CuS + Fe2+ + S° + 2e (2-59) or by the reaction between cuprous ions and elemental sulfur. Leaching in the potential range above 0.33 V S C E (or 0.530 V S H E ) and below 0.43 V S C E (or 0.630 VSHE) is shown below, depending on the oxidation state of iron dissolution (Fe y + or Fe2+ ). CuFeS2 <=> Cu2+ + (Fe3+ to Fe2+) + 2S° + (5 to 4)e (2-60) Leaching curves of the second stage reactions are depicted in Figure 2-13 and the reaction curve of 0.430 VSCE (or 0.630 V S H E ) is shown as a dotted curve. It is apparent that the leaching rate declines around 0.450 VSCE (or 0.650 VSHE), and achieves relatively slow increases until reaching 0.600 VSCE (or 0.800 VSHE)- Above this potential the rate becomes independent of suspension potential. Leaching is incomplete even after 40 hours at 0.470 VSCE (or 0.670 VSHE), whereas leaching approaches completion after 50 hours at 0.60-0.65 VSCE (or 0.800-0.850 V S H E ) -Bisulfate ion (HS0 4 ) is believed to be the stable form of sulfur in the second stage: CuFeS2+8H20 <=> Cuu +(Fe3+ to Fe2+) + 2HSO~ + 14H+ + (17 to 16)e~ (2-61) The leaching rate is nearly independent of ferric ion concentration, which was investigated by varying the total iron concentration (0.01-0.5 mol/L) at a set solution potential of CHAPTER 2 LITERATURE REVIEW / 56 0.40 VSCE (or 0.600 V S H E ) , which indicates the absence of diffusion control. It also follows that the variation of oxidation rate as discussed above is due mainly to the effect of suspension potential in the presence of both ferrous and ferric ions. The temperature effect on leaching rate and the induction period is investigated by varying the leaching temperature (50-90°C) at a set solution potential of 0.40 VSCE (or 0.600 V S H E ) - Leaching rate increases with increasing temperature, where leaching rate is higher at 90°C than 50°C. The induction period is when leaching proceeds at a slow but measurable rate, and is inversely proportional to the leaching rate. The dependence of the induction period on potential is not addressed in this paper, except that a much longer induction period is observed in a fresh solution than in a solution from a previous run. The apparent activation energy at 0.40 V S C E (0.600 V S H E ) is 47 kJ/mol. In summary, the maximum rate of leaching is attained only over a very narrow range of suspension potentials. They attribute the critical potential centering at 0.45 VSCE (or 0.650 VSHE) to either a change in the nature of chalcopyrite or the chemical reaction during leaching at that potential. The relatively low optimum redox potential proposed indicates that there is an optimum concentration of ferrous ions that results in increased leaching. Experimental results have revealed that the rate constant (k ) is closely related to the redox potential (Em) rather than the ferric iron concentration at 90°C. In the potential range: 0.33-0.37 V S C E (or 0.530-0.570 VSHE), the leaching of chalcopyrite is reported as: Em =0.375 + 0.030 log k (2-62) CHAPTER 2 LITERATURE REVIEW / 57 Suspension cn it •a o c ic o c D • O £ < 10 V 0-30 V A 0-33 A 0-35 o 0.37 o 0-40 • 0.43 ure 2-12 Time Reaction curves (KMn04 added vs. log t) for the oxidation of chalcopyrite in the potential range 0.30-0.43 V S C E (or 0.600-0.630 V S H E ) (90°C, initial iron cone. = 0.1 mol/L) 163 10 .. 100 T i m e , min 1000 ure 2-13 Reaction curves (KMn04 added vs. log /) for the oxidation of chalcopyrite in the potential range 0.47-0.65 V S C E (or 0.670-0.850 V S H E ) (90°C, initial iron cone. = 0.1 mol/L). Dotted Curve: reaction curve at 0.43 VSCE (or 0.630 VSHE) in Figure 2-12 for comparison.163 CHAPTER 2 LITERATURE REVIEW / 58 2.6.2.2 Redox-Controlled Chalcopyrite Bioleaching Patent of Mintek The influence of redox potential on the leaching of chalcopyrite had not gained much attention until a South African patent was granted to Mintek for the controlled-potential leaching of chalcopyrite. This patent (ZA 9701307)1 6 8 describes the leaching of chalcopyrite using ferric sulfate in a fairly narrow range 0.350-0.450 V S C E , which is consistent with the results of Kametani and A o k i 1 6 3 . The potential range is somewhat ore specific, and varies for different chalcopyrite samples, depending on the crystal structure, impurities and associated minerals present in the ore. Some experimental results are portrayed from Figures 2-14 to 2-16, where the slurry redox potential is controlled between 0.370 and 0.400 V S C E with a particle size of 20 p:m at 35, 55 and 80°C. The desired ferric-ferrous ratio is controlled with the addition of potassium permanganate solution, the pH is maintained at 1.5, and the initial iron concentration is 3 g/L in these tests. The temperature dependence of leaching is once again apparent here where 90% copper extraction can be obtained in less than 4 hours at 80°C, but 2.7 days (64.8 hours) is required at 35°C. The thermal instability of the passivating film at high temperatures such as 80°C (Figure 2-16) together with the preferential dissolution of chalcopyrite within the controlled-potential range may give rise to the facile chalcopyrite dissolution. The leaching curve at 35°C (Figure 2-14) has a striking similarity to the bacterial leaching curve with the characteristic lag phase of the incubation, follow by the growth and a slowly stagnant death phase of bacteria; this observation is similar to the induction period reported by Kametani and A o k i . 1 6 3 Figure 2-17 shows a series of tests carried out at 70°C for 48 hours using concentrates of two particle sizes: 90% passing 103.8 p;m and 90% passing 6.9 \im. Six tests were carried out on each particle size at slurry redox potentials of 0.394, 0.404, 0.414, 0.424, 0.434 and 0.444 V S C E -Copper extractions in excess of 90% were obtained within 48 hours for the fine milled concentrate (90% passing 6.9 |im) at a redox potential window of 0.414 to 0.434 VSCE (or 0.628-0.648 VSHE)- Above or below this range the copper leach kinetics was less favorable, and the copper extraction declined appreciably. However, a much narrower window of redox around 0.414 VSCE (or 0.628 VSHE) had to be maintained to obtain high copper extraction for the more coarse concentrate (90% passing 103.8 jam). In other words, high copper extraction can be CHAPTER 2 LITERATURE REVIEW / 59 achieved by maintaining the redox potential window of around 0.414 VSCE (or 0.628 VSHE) regardless of particle size (fine or coarse). The process is preferably controlled by selection of the surface potential, leach temperature, pH of the leach solution, and the particle size of the chalcopyrite ore. This process is effective in both bioleaching and ferric sulfate leaching processes. The patent describes a process in which the potential control is achieved by controlling the ferric-ferrous ratio, and this ratio is maintained by controlling the oxygen supply to the ferrous ion oxidation process. The oxidation of ferrous to ferric ions is carried out by bacterial oxidation, preferably conducted in a separate reaction vessel so that the residence time of circulating leach solution in the bacterial oxidation vessel may be actively controlled. The chalcopyrite is finely milled prior to leaching in order to increase the available surface area and promote crystallographic dislocations. The Mintek redox-controlled process represents a breakthrough in chalcopyrite leaching. The ferric sulfate leaching of chalcopyrite is enhanced by controlling the redox potential in a pre-defined range. Despite the excitement, however, heap application is bound to be difficult for the controlling of such a narrow range of potential. Hackl 1 7 5 suggested that the heaps should perhaps be made anoxic (oxygen-depleted) to keep the potential low enough, and leach chemically with ferric sulfate. CHAPTER 2 LITERATURE REVIEW / 100 H i 1 i i 1 1 1 1 1 1 0 0.3 0.6 0.9 1.2 1.5 1.8 2.1. 2.4 2.7 3 Time [Days] Figure 2-14 Ferric oxidation of fine-milled South-African chalcopyrite at 35°C 168 (initial iron cone. = 3 g/L, pH = 1.5) CHAPTER 2 LITERATURE REVIEW / 61 O l i : i ~ 1 i i i 0 4 8 12 16 20 24 T i m e [Hrs] Figure 2-16 Ferric oxidation of fine-milled South-African chalcopyrite at 80°C (initial iron cone. = 3 g/L, pH = 1.5)168 CHAPTER 2 LITERATURE REVIEW / 62 2.7 Objective of the Research The present investigation was aimed at an improved understanding of processes in which copper is leached from chalcopyrite (CuFeS7) under oxidizing conditions in acidic ferric sulfate solution. Various researchers have shown that the dissolution of chalcopyrite has little dependence on ferric ions (adequate to initiate leaching, above 0.005M 5 7), and that increasing the concentration of ferric does not result in a corresponding increase in the leaching rate of chalcopyrite. 6 1 ' 6 2 ' 6 3 ' 6 4 On the other hand, most researchers agree upon the negative effect of the addition of ferrous sulfate on acidic ferric sulfate leaching of chalcopyrite, and that its concentration in solution should be minimized to improve leaching rates. The role of ferrous ions has been taken merely as a source of ferric ions in both biological and chemical leaching. However, a review of the literature reveals the suppression of leaching by high concentrations of ferric ions together with the apparent enhancing effect of low redox potential, which suggests that redox potential (EH ) is thus the key parameter in the dissolution of chalcopyrite. The leaching rate of chalcopyrite can therefore be controlled by controlling the mixed potential. Instead of applying external potentials to control the mixed potential and oxidizing conditions at the metal surface, constant redox potential of the leaching solution is maintained by controlling the feed rate of potassium permanganate solution. The experiments are chosen to operate in sulfate-based medium and the temperature range of extremothermoacidophilic bioleaching for future application of current research. Media other than sulfate-based medium will kill the bacteria. Since these bacteria do not enter as reactants in an oxidation reaction sequence but only as mediators, biosystems can carry out only those reactions which are fhermodynamically feasible, the presence and absence of bacteria should produce the same kind of leaching behaviors. The aim of the present research is to gain a better understanding of the passivation behavior of chalcopyrite as a function of redox potential (0.400-0.600 V A g /A g c i ) under ferric sulfate leaching conditions at modestly elevated temperatures between 60 and 78°C. CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 63 C H A P T E R 3 E X P E R I M E N T A L P R O C E D U R E S A general description . of the materials,. instrumentation, analytical techniques and experimental considerations employed during the course of the thesis work are presented in this chapter. 3.1 Materials 3.1.1 Chalcopyrite Samples Leach tests were carried out on chalcopyrite concentrates obtained from Santa Eulalia Mining District (State of Chihuahua, Mexico), and obtained through Mineralogical Research Co., San Jose (California, USA). The as-received concentrates were finely ground slurry samples. Chemical and X-ray diffraction analyses of the as-received concentrate revealed the presence of substantial amounts of pyrite, quartz and other iron-containing gangues (including magnetite, FeFe204) as well as chalcopyrite. To reject these impurities, the as-received concentrate was sent to PRA (Process Research Associates Ltd., Vancouver, Canada) for flotation. The flotation was quite successful, yielding about 0.9 kg of a relatively pure, finely ground CuFeS2 concentrates (93.1% passing 37 um or 400 mesh as determined by sieving). The chalcopyrite concentrates were allowed to air dry at room temperature. Although no protective measures were taken to guard against surface oxidation of the mineral during storage, chemical analyses of the floated concentrate analyzed on five different occasions (at least two months apart) before and during experimentation revealed little variation as shown in Table 3-1. The average composition of these chalcopyrite samples was taken as the composition of the mineral samples in all calculations. The methods of chemical analyses for the metal and sulfur species in the floated concentrates are the same as those procedures employed for the leach residues as described in Section 3.2.3. These methods include copper and iron assays by titration, sulfur determinations by gravimetric analysis and ICP scans for a range of available metals. Although ICP analysis confirmed the presence of approximately 2.70% zinc, the powder XRD technique determined that chalcopyrite was the major constituent and the pyrite gangue was present in trace amounts in the floated chalcopyrite concentrate as shown in Figure 3-2. On the basis of 30.40% elemental CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 64 copper, the chalcopyrite content of the sample was estimated to be 87.79%. Assuming that zinc was preset as sphalerite (ZnS ) in the sample, the sphalerite and pyrite contents of the sample were estimated to be 4.02% and 2.96%, respectively. Although sphalerite is less noble than chalcopyrite, the oxidation of sphalerite will consume an insignificant amount of potassium permanganate solution. In total, 96.67% of the mineral composition could be accounted for. The rest was difficult to predict based on the available information. The chemical and mineralogical compositions are shown in Table 3-2. Table 3-1 Chemical analyses of the floated CuFeS2 concentrate Sample CuT Fer S j s2- Zn Insol.1 No. (%) (%) (%) (%) (%) (%) 2 30.13 29.34 34.67 33.71 3.20 2.66 12 30.46 29.93 34.43 33.69 2.80 2:38 28 30.54 29.04 33.82 33.17 2.90 3.70 29 30.49 29.60 34.70 34.10 2.40 2.81 45 30.40 29.94 33.26 33.23 2.20 4.20 Average 30.40 29.57 34.18 33.58 2.70 3.15 Acid insoluble content, indicative of siliceous gangue. Table 3-2 Chemical and mineralogical composition of copper concentrate C H E M I C A L A N A L Y S I S M I N E R A L O G I C A L COMPOSITION Element Percent Mineral Percent Cu 30.40 CuFeS2 (Chalcopyrite) 87.79 Fe 29.57 FeS2 (Pyrite) 2.96 Zn 2.70 ZnS (Sphalerite) 4.02 Sr (Total Sulfur) 34.18 Fe -containing Gangue Minerals 0.53 S2- (Sulfide Sulfur) 33.58 Finely ground concentrates were necessary due to the refractory nature of chalcopyrite. The size distribution of this concentrate was determined using an Elzone 280PC particle size counter (Particle Data Inc., Elmhurst, Illinois). This instrument determines a continuous size distribution by measuring particle volumes and reports the equivalent spherical diameter. It is capable of measuring subsieve-size particles, but disregards particle sizes below 1 pm. The cumulative particle size distributions detennined by the Elzone 280PC counter are plotted in Figure 3-1. The (arithmetic) mean and median particle sizes (or diameters) were 17.52 pm and 14.87 pm, respectively. CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 65 0 5 10 15 20 25 30 35 40 45 50 55 60 Ffertide Diameter (r^m) Figure 3-1 Cumulative particle size distribution of the floated chalcopyrite concentrate as determined by the Elzone 280PC counter CHAPTER 3 EXPERIMENTAL PROCEDURES / 66 CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 67 3.7.2 Reagents A l l stock solutions were prepared from reagent grade chemicals and distilled and deionized water. Unless otherwise indicated, commercially available chemicals were reagent grade and were used as received, without further purification. Reagent grade chemicals employed for the leaching experiments were obtained from commercial suppliers, either Sigma Aldrich or Fisher Chemical, including those listed below. • Iron(II) Sulfate Heptahydrate [FeS04 • 1H20 ] (Certified ACS, 99% purity) • Iron(III) Sulfate Pentahydrate [Fe2(S04)i -5H20] (High Grade, 97% purity) • Potassium Permanganate (KMn04 crystals) (Certified ACS , 99.6% purity) • Sulfuric Acid (H 2 S0 4 ) (18.0098 M or 36N) (Certified ACS , 96% purity) Nitrogen gas (7V2) (99.5%), obtained from Praxair Technology, Inc, was used without further purification. CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 68 3.2 General Experimental Techniques and Instrumentation 3.2.1 Leaching Apparatus , A l l batch reactor experiments described herein were carried out in a compact bench-scale 3-L glass jacketed bioreactor equipped with a servo motor and speed controller manufactured by Applikon Dependable Instruments (ADI) (Schiedam, The Netherlands). Figure 3-3 illustrates the major components of the system used in the batch operation. A l l parts in contact with the lixiviant were made of material resistant to corrosion by the process solutions, mainly acidic ferric sulfate solution. The bioreactor (with a dished bottom) has a working volume of 2.7 L, a minimum working volume of 0.47 L, and a tank height to diameter ratio (H/D) of 1.5. The weight of the glass reactor with empty water jacket is 3.19 kg. The working volume of the jacket is 1.3 L. The desired reaction temperature was controlled within 1°C by circulating water through the reactor jacket. Agitation was achieved by a variable speed stirrer assembly and the three vertical baffles. This stirrer assembly consists of two 6-bladed pitch blade impellers. These impellers are 6-bladed, 45 degree (vortex) pitch down type with a diameter of 45 mm. The dished bottom also assures good mixing characteristics. The three vertical baffles in the reactor have an insertion length of 220 mm and a width of 6 mm. Preliminary runs were performed to determine a suitable speed for stirring the process liquid under the test conditions. It was found that at 400 rpm the particles were not suspended in solution. Particles were suspended at 600 rpm. At this stirring rate, the diffusion layer encircling individual particles was assumed to be minimized and homogeneous concentration of solutes was maintained throughout the liquid volume. Other mechanisms such as diffusion and advection were assumed to be negligible. Stirring speeds of 600, 800 and 1000 rpm gave virtually the same copper extractions with slight deviations. However, stirring at 1000 rpm was overly vigorous and uniform stirring was deemed impractical at this stirring speed. Therefore, a stirring speed of 800 rpm was used in all experiments reported in this thesis work. The reactor has a tight-fitting head plate with six holes to allow for the insertion of two electrodes, a sampling port, an inlet tube for nitrogen gas, a triplet inlet for the addition of reagents, and a condenser. An air-outlet stainless steel condenser designed for optimal heat exchange between the exiting gas and the cooling medium of the condenser was used to prevent evaporative loss. The nitrogen gas was sparged through a 1/8 inch I.D. stainless steel tube with a CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 69 porous nozzle and an insertion length of 236 mm. The tube was inserted through the reactor head plate with the bottom tip positioned approximately 10 mm from the impeller shaft underneath the bottom impeller. The nitrogen gas flow was controlled with a Matheson (Montgomeryville, PA, USA) mass-flow controller (model 7351H). The pressure generated during the reaction was monitored with a pressure gauge. The nitrogen gas was saturated with water before being passed into the reactor in order to maintain constant solution volume. The flow of nitrogen gas was continued throughout the experiments to prevent diffusion of atmospheric oxygen to the solution surface. A nitrogenous atmosphere prevents undesired redox reactions, particularly pyrite, sulfur or iron oxidation by atmospheric oxygen. The dissolved oxygen and associated temperature was measured with YSI Model 58 Dissolved Oxygen Meter with YSI Model 5739 Dissolved Oxygen Probe. The polarographic sensor in the probe was covered by an oxygen permeable membrane, which prevented impurities from entering the electrode, and interfered with the measurement. The polarographic sensor was equipped with built-in thermistors for temperature measurement and compensation. Although ferrous oxidation by oxygen is thermodynamically spontaneous (AG° = -88 kJ/mole Fe), the reaction is kinetically hindered in acid. 1 7 6 The system was thoroughly agitated and degassed with high purity nitrogen gas for about 30 minutes before the addition of ferrous sulfate salt at the initial stage of the experiment (pH 1.5, 20°C). The dissolved oxygen content was measured to be less than 2 ppm, and was determined to be adequate in preventing ferrous oxidation by atmospheric oxygen. The system is controlled by an ADI 1030 Biocontroller, which is a flexible, digital controller capable of monitoring and controlling four parameters throughout the experiments: acidity (pH), temperature, solution redox potential and the rate of agitation. This controller controlled the valves that lead the liquid addition of the potassium permanganate and sulfuric acid solutions to the analyzing apparatus at a pre-defined pH and solution redox potential. The temperature compensator of the controller enables the measurement of pH and solution redox potential at its actual temperature. The data were monitored on-line through a desktop computer operated using a data acquisition program BIOWATCH (version 2.30).1 7 7 A schematic of the bioreactor, along with the control components, is shown in Figure 3-3. Temperature control Biowatch Program Data Acquisition CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 70 N2/Stirring control E H control pH control N 2(g) 800 rpm K M n 0 4 H 2 S 0 4 Sensors : E H Temperature pH Figure 3-3 Schematic diagram of the controlled-potential leaching system in batch mode A VWR® pH combination epoxy electrode was used to measure pH in the bioreactor. The probe combines a pH electrode and an Ag/AgCl reference electrode filled with a reference solution containing 4M KCl saturated with Ag/AgCl. A glass membrane responds to the hydrogen ion activity by developing an electrical potential across the membrane. The electrical potential follows the Nernst principle, and varies linearly with the pH of the solution being measured. The probe was consistently calibrated in two buffer solutions at pH 2.0 and pH 1.0 before every test. The two-point buffer calibration method was carried out using the pH monitoring function of the bio-controller. The bio-controller automatically generates the efficiency as determined by the slope of the correlation between EH and pH. When the calibration could not be achieved or if the efficiency of the pH probe was less than 90%, the probe was cleaned or replaced. A VWR platinum combination redox epoxy electrode, which combines a platinum electrode and the Ag/AgCl refillable junction reference electrode solution containing 4M KCl saturated with Ag/AgCl, was used for redox measurements. The probe responds to oxidation-reduction potentials in solution. A silver/silver chloride electrode (Ag/AgCl), rather than a saturated calomel electrode (SCE), was used because it could function at temperatures greater CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 71 than 60°C. 1 7 8 The electrical potential again follows the Nernst principle. To ensure good readings, calibration checks and probe maintenance were employed. Between each set of experiments, the electrolyte solution of the redox electrode was replaced with potassium chloride solution (fill solution = 4M KCl w/ AgCl) and the electrode was recalibrated using standard solutions of known redox potential. The A S T M recommended standard reference solution for oxidation-reduction (ORP) measurement, also known as Light's solution, is prepared as follows: 1 7 9 ' 1 8 0 dissolve 39.21 g of ferrous ammonium sulfate [Fe(NH4)2(S04)2 • 6H20 ], 48.22 g of ferric ammonium sulfate [FeNH4{S04)2 \2H20], and 56.2 mL of concentrated sulfuric acid (18M) in 1 L distilled and deionized water. Unlike other EH reference solutions, Light's solution does not precipitate over time. The standard reduction potential of the Ag/AgCl reference electrode in saturated 4M KCl solution in the above ferric-ferrous reference solution at 25°C is 0.475 ± 0.010 VSHE- After every test, the pH and the redox probes were disassembled entirely to facilitate cleaning and dispose of internal precipitate build-up. To prevent build up of an iron-precipitate layer on the probe surface, fine grain sand paper (5 urn grain size) was used to sand off any visible film developed on the probe head. The pH probe was stored with pH 2 reference solution on a daily basis, and the redox probe was stored in Light's solution. Cole-Parmer Master Flex (Model 7014-20) peristaltic pumps were used for liquid addition. The potassium permanganate (KMnOA) and sulfuric acid (H2SOA) solutions were both pumped at a flow rate of 0.005 L/min (or 0.3 L/hour). A l l tests were maintained at a pH below 1.5 by the addition of 6M H2S04 and at the desired solution redox potential by the addition of 0.2 M KMn04 solution. The addition of acid was necessary throughout the leach to minimize iron precipitates, especially iron jarosites. The acid addition pump was connected to a 50 mL burette of 6 M / / 2 S 0 4 , and the oxidant addition pump was connected to a 250 mL burette of 0.2 M KMn04 . Both burettes were covered with parafilms to minimize contamination and evaporation. Since sunlight decomposes permanganate solutions, the apparatus was kept out of direct contact with sunlight. CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 72 3.2.2 Leaching Procedures A standardized test was employed in order to compare the effect of potential at various temperatures. Leach tests were carried out at three temperatures: 60, 70 and 78°C. The upper temperature limit of 78°C was imposed by the thermal instability of the VWR® Redox Probe above 80°C. Although the redox potential difference at the interface of the suspended particles and solution cannot be measured or controlled in leaching experiments, it has been assumed in this study that the redox potential measured in the bulk of the leach solution approximates the mixed potential ( Em ) of the oxidizing mineral surface. Since the redox potential ( E ) corresponds to the oxidizing capability of the bulk solutions, it is termed "solution redox potential" in this study. The solution potential was controlled at selected intervals in the range of 0.400 to 0.600 V A g / A g c i - Preliminary tests showed that potentials above 0.600 V A g / A g C ! consumed too much potassium permanganate solution, which exceeded the maximum allowable working volume of the bioreactor. At high potentials, little ferrous ions exist, and the kinetics between potassium permanganate and ferrous ions become very slow because of the "law of mass action". Thus, the slow kinetics effectively sets an upper limit on the controller of the solution redox potential. If the controller set point exceeds this limit (above 0.600 V A g / A g c i ) , the controller ends up "dumping" potassium permanganate solution uncontrollably into the bioreactor. The leaching time depended on leaching conditions (see Table B- l ) . Duration of the tests was determined based on the addition rate of potassium permanganate solution, and ranged from as long as 60 hours at 60°C to as short as 8 hours at 78°C. Since flotation agents on particle surfaces would degrade at process conditions (pH below 1.5, temperatures between 60 and 78°C), the samples were not pre-washed with methanol prior to leaching. The standard FeS04 -H2S04 leaching medium for all tests was prepared in-situ, and had a redox potential of about 0.315 V A g / A g c i at 25°C (0.514 VSHE)- After the introduction of chalcopyrite, the standard leaching medium in all tests was an aqueous solution of ferrous sulfate containing 2.95-3.13 g/L Fe2+ , 8.8 g/L H2S04 and 5 g/L CuFeS2. The solid pulp density (SPD) was 0.49%. The proportion of ferrous ion used was one-fourth of the stoichiometric requirements to dissolve the copper content of the concentrate (when in ferric form). On one hand, permanganate oxidation of ferrous to ferric ions occurred instantly. Ferric ions were recycled incessantly and would be constantly available for the oxidation of the chalcopyrite concentrate. On the other hand, the deposition of ferric iron CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 73 precipitates (such as jarosites) on mineral surfaces was minimized by the presence of low levels of total soluble iron. These unwanted precipitates might impart a diffusion barrier to mineral surfaces, impeding oxidation, and rendering intermediate compounds that could be mis-identified as the so-called "passivating layers" on chalcopyrite. A l l tests were performed according to the following procedure: 30 mL sulfuric acid (6M H2S04) was first added to 2 liters of distilled and deionized water to produce an acidic solution suitable for the dissolution of ferrous sulfate and chalcopyrite. Upon agitation, this initial acidic solution was purged with purified nitrogen gas at a flow rate of 445 mL/min for about 30 minutes to remove oxygen. Ferrous sulfate powder (30 g) was then introduced into the reactor. Since most investigators agree that ferrous oxidation of chalcopyrite is not possible without an oxidant such as diatomic oxygen or ferric (see Section 2.3.1), sample mineral was introduced into ferrous sulfate leaching medium at start-up. After 10 minutes of stirring, 10 grams of chalcopyrite sample was added into the vessel. Acid soluble copper and iron in the sample material would dissolve into the starting medium during the warm-up time. The warm-up time varied from 30-120 minutes, depending on the desired temperature. Once adequate mixing was achieved (roughly ten minutes) for the added reagents, the heating unit was then engaged. During heat up, only nitrogen gas was fed to the reactor. Time zero was taken as the time at which the desired operating temperature and solution redox potential were achieved. From time zero, the control systems automatically maintained the pH and redox potential levels, activating the peristaltic pumps associated with each. The addition of permanganate (0.2 M KMn04 ) solution was required to oxidize the ferrous iron to ferric iron until the desired solution redox potential was reached. Although sulfuric acid did not directly take part in the reaction, the effect of pH was still very important. There is a limit in the pH value above which ferric ion will be hydrolyzed. The response of the logging system was rapid and the readings were found to be stable over long periods of time. The solution redox potential was maintained within 0.02 V A g / A g c i , and the pH was maintained within 0.1 pH units at pH below 1.5 in all cases by the continuous addition of 6M sulfuric acid. A l l solution redox potentials were taken at the reaction temperature versus the Ag/AgCl electrode, and were latter converted to standard conditions against SHE (see Appendix A - l 1). A summary of the standard operating conditions is given in Table 3-3. CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 74 Liquid samples of leach solutions were withdrawn through the sampling port at timed intervals during the course of the reaction. Samples were taken frequently during the initial stages of leaching when intense oxidation occurred as determined by the rapid addition of potassium permanganate. Later, the frequency of sampling was decreased. Ten millimeters of slurry was removed by a sterile, previously unused syringe with attached PVC tubing. The samples were immediately emptied into plastic centrifuge tubs and sealed with Parafilm laboratory film to minimize the occurrence of premature atmospheric oxidation of ferrous ion to ferric ion. A l l samples were cooled immediately in a room temperature water bath. After cooling, sample solutions were centrifuged for 15 minutes and the supernatants were diluted as required for the analyses of the dissolved copper and iron. Centrifuging was necessary to ensure that no entrained fine particles were present in the supernatants, which might clog the atomic absorption atomizer. Aliquots withdrawn from the leaching solution were analyzed for ferrous iron content in leach solutions as determined by permanganate titration (see Appendix C- l ) . Distilled and deionized water was added to maintain a constant electrolyte volume. When liquid samples were removed from the leaching vessels by sampling, equal volumes of distilled and deionized water were used to rinse the tubing as well as the syringe before being added to the reaction vessel. Table 3-3 Standard operating conditions for all leaching experiments CONTROLLED PARAMETERS PARAMETER V A L U E S Temperature setpoint (+ 1°C) : 60°C/ 70°C/ 78°C Solution redox potential setpoint ( ± 0.02 V A g / A g C i ) : 0.400/0.450/0.500/0.550/0.600 V A g /A E c i Target p H ( ± 0 . 1 ) : 1.5 Standard leaching medium composition : 2.95-3.13 g /LFe 2 + ; 8.8 g/L H2S04; 5 g/L CuFeS2 Solid pulp density (SPD) : 0.49% Acid i ty control by : 6 M H2S04 Redox potential control by : 0.2M KMn04 N2 gas flow rate : 445 mL/min Stirring speed : 800 rpm CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 75 The rate of oxidation during leaching was followed and determined by the rate in which potassium permanganate was added to the leaching solution to maintain a constant solution potential. The overall reaction rate was studied as a function of total copper concentration. Quantitative copper analysis at sample time was performed with AAS. The extent of leaching at any time was taken as the fraction of copper released into solution. Each data point in an extraction-time plot represents the copper extraction achieved at a given point in time when a sample was taken, making it possible to measure changes in the solution component concentrations as leaching progressed. In other words, each data point is a calculation that attempts to account for all copper solubilized during the leach test, up to and including the sample. When the test was completed, if was terminated by cooling the reactor rapidly to about 40°C to avoid any undesired side reactions. The reactor contents were allowed to settle for approximately thirty minutes before weighing and filtration. Any material adhering to the vessel was removed and also filtered. The reaction mixtures were suction filtered through 2.7 pm Whatman (hardened, low ash grade) filter paper and the resulting solid residues were washed thoroughly with dilute 6M H2S04 to ensure that all soluble salts were removed. The leach residues were dried in an oven maintained at 50°C, to prevent oxidation of sulfide sulfur. The filtered leach residues from each experiment were sent to 1PL (Vancouver, B.C., Canada) for total copper, total iron and sulfur suite [total sulfur (ST), sulfide sulfur (S2~), elemental sulfur (5°) , sulfate sulfur (5 2.)] analyses. The filtrates were always noticed to have a greenish or yellowish color in most experiments, and the volume of leaching solution at completion was between 2100 mL and 2600 mL, depending on the leach temperatures and solution redox potentials. A sample of the filtrate was also taken and analyzed as described in Sections 3.2.3 and 3.2.4. The leach residues were analyzed by XRD and SEM analyses. CHAPTER 3 EXPERIMENTAL PROCEDURES / 76 3.2.3 Quantitative Analysis Solution Samples from each test were analyzed for dissolved copper and soluble iron (ferric plus ferrous) by AAS to a precision of 1 ppm. The accuracy was 3 to 5% depending on the dilution factor and number of serial transfers performed during dilution. Samples and the standard were all prepared in diluted hydrochloric acid solution (2% w/w). Aliquots of 2 mL of the supernatants were diluted to a sufficient level for analysis by AAS using a Unicam 929 A A Spectrometer (Unicam Limited, Oxford, UK) . Preliminary tests and references both indicated that the effect of sulfate, iron and manganese ions on copper were minimal, and could be safely ignored without matrix matching under the test conditions performed in this thesis work. 1 8 1 However, iron is more sensitive than copper to matrix conditions. A few drops of hydrogen peroxide (30% H202) were added to the diluted samples to oxidize ferrous to ferric ions. Sulfate species were matched with sulfuric acid and manganese species were matched with manganese sulfate (MnSOA ) in both the diluted samples and standards. Serial transfers during dilutions were performed with the aid of an Eppendorf micropipette. Periodically, the micropipette was cleaned and calibrated to ensure reliable apd accurate performance. Solution samples from the experiments of data set #2 were assayed for ferrous ions (Fe 2 + ) 182 using a standard titration with potassium permanganate. Three 1 mL aliquots of solution were subsequently titrated using dilute potassium permanganate solution of known molarity. Upon ferrous oxidation, the intensely colored (purple) potassium permanganate solution turns colorless. The endpoint is now marked by the transition from colorless to a lasting deep pink color with just a single drop of unreduced permanganate solution, where both aqueous ions (Mn2* and Fei+ ) are virtually colorless. The ferric concentration was subsequently determined according to equation (3-1) by subtracting the ferrous ion concentration from the concentration of total soluble iron. The procedure for titrimetric determination of ferrous ions against potassium permanganate solution (KMn0 4 ) is shown in Appendix C - l . [Fei+] = [Fesr°0;:>k]-[Fe2+] (3-1) The filtered leach residues were assayed for copper, iron and sulfur as well as an ICP scan for a range of available metals by personnel at a local Vancouver analytical laboratory, International Plasma Laboratory Ltd. (IPL). The copper and iron contents were determined by CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 77 titrations after multiacid digestion, which is usually considered more precise than the AAS method. This was the same methodology for determining the elemental composition of the feed material. The determination of sulfur species ( ST , S° , S2~ , -Ssg2 ) was accomplished by gravimetric analysis, which involved the dissolution of respective sulfur species into solution, followed by the selective precipitation of individual sulfur species as barium sulfate (BaS04). A sample of the leach residues was washed in boiling perchloroethylene ( C 2 C / 4 ) (b.p. 121°C) to remove elemental sulfur (S°). The resulting solution containing the elemental sulfur content was taken to dryness, and reoxidized in KBr - Br21HNO-^ solution, prior to gravimetric precipitation of barium sulfate (BaS04). The sulfate (<-« 2 .) was then removed from the perchloroethylene leach residue by selective dissolution in 5-10% Na2C03 solution. Again the sulfur is precipitated as BaS04. Finally, the Na2C03 leach residue was digested in a KBr - Br21 HNOy solution and assayed for sulfide sulfur (S2~) in the same method. Total sulfur (ST) determination was accomplished by the same method, except that the steps to remove elemental sulfur (S") and sulfate (S 2„) were not taken. The procedures performed at IPL for copper and iron assay by titration as well as ICP multi-element analysis are shown in Appendix C. CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 78 3.2.4 Qualitative Analysis 3.2.4.1 Powder X-ray Diffraction (XRD) Selected samples from the leach residues and the initial starting chalcopyrite concentrates were studied by step-scan X-ray power diffraction (XRD) in an attempt to determine their chemical and mineralogical composition. The XRD study was carried out with a Phillips 1830 X -ray generator (featuring CuKa^ radiation at a wavelength of 1.5406A) on a standard Siemens D5000 Bragg-Brentano diffractometer located in the Department of Earth and Ocean Sciences at UBC. This diffractometer is equipped with a diffracted-beam graphite monochromator crystal, 2 mm (1°) divergence and antiscatter slits, 0.6 mm receiving slit and incident beam Soller slit. The long sample holder used (43 mm) ensured that the area irradiated by the X-ray beam under these conditions was completely contained within the sample. The Cu X-ray tube was operated at 40 kV and 50 mA. The X-ray powder diffraction data were collected over a range of 3-70°2? with a step size of 0.04°2? and a counting time of 2 s/step with no rotation. The diffraction data were processed using the E V A software as part of the DIFFRAC Plus suite of programs (5 t h ed.), which determines peak positions with intensities and searches the International Centre for Diffraction Data (LCD.D.) for the mineral with the closest match. 3.2.4.2 Scanning Electron Microscopy (SEM) Selected samples from the leach residues were studied by Scanning Electron Microscopy (SEM) in an attempt to examine their surface conditions. The SEM study was carried out on a fully automated PHILIPS XL30 electron microscope (The Netherlands) with Princeton Gamma-Tech (PGT) IMIX-PC energy dispersion and digital image analysis (IAS) systems. The starting chalcopyrite concentrates and leached resides were powder samples that were carbon coated on aluminum stubs using the EDWARDS Coating System E306A. In addition to quick surface morphological examination at both low and high magnifications, digital images of various signals (secondary electrons, backscattered electrons and characteristic X-rays) were taken at selected representative areas on all samples. The obtained X-ray spectra were evaluated qualitatively by the IMIX-PC energy dispersive system to produce element maps. CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 79 3.3 Experimental Considerations and Computations 3.3.1 Permanganate Oxidation of Ferrous Iron in Acidic Solution Preliminary examination showed that the presence of constant proportions of ferric and ferrous sulfate as a means to maintain constant solution redox potential was not possible. Ferric sulfate not only had low solubility, which made the preparation of ferric-ferrous sulfate solution difficult, but the precipitation of ferric iron, particularly when concentrated (i.e., high potential), during the course of experiments constituted a major problem. Alternatively, constant solution redox potential was maintained by controlled addition of potassium permanganate, continuously oxidizing theFe(II) and regenerating theFe(III) in a cyclic manner. Potassium permanganate is highly soluble in aqueous solution (6:34 g KMn04 /100 g H 2 0 at 20°C, or 0.4012 M ) . 1 8 3 Since the standard reduction potential (E°) of potassium permanganate in acid solution is 1.51 volts, it is a powerful oxidizing agent suitable for the oxidation of ferrous iron in sulfuric acid. 1 8 2 Additionally, a dilute titrant solution is necessary to obtain high precision because one mole of permanganate ions (Mn0 4 ) titrates 5 moles of ferrous ions (Fe 2 + ). Permanganate is inherently unstable in the presence of manganese(II) ions (Mn2+), and it 184 leads to the side reaction known as the "Guyard" reaction: 2MnO;(aq) + 3Mn2+(aq) + 2H20 5Mn02(s) + 4H +(aq) (3-2) Complete reduction of permanganate ion is therefore essential to prevent the precipitation of manganese dioxide, which might coat the mineral and impede oxidation. The presence of 185 manganese dioxide (Mn02) also interferes with the role of permanganate oxidation. It could act as an oxidizing agent and oxidizes ferrous ion in the solution to ferric ions: Mn02(s) + 2Fe2+(aq) + 4N+(aq) Mn2+(aq) + 2Fe'+(aq) + 2H20 (3-3) Latimar and EH - pH diagrams provided compact portrayal of the oxidation states of manganese. Intermediate states during permanganate oxidation are shown in the Latimar Diagram (Figure 3-4). A Latimar diagram depicts the species which disproportionate spontaneously. The potential of an element, not complicated by disproportionation, changes steadily from left to right, in a decreasing trend (more positive to more negative). Good reducing agents are on the right, good oxidizing agents are on the left, and stable species toward the CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 80 middle. A species has a thermodynamic tendency to disproportionate into its neighbors i f the potential on the right of the species is higher than the potential on the left. Such trend is broken at MnO\~ and Mrr+ species in acidic solution. Thus, MnO\~ and Mni+ species are unstable in acidic solution and disproportionate spontaneously. The disproportionation of Mn02A~ species gives Mn02 and Mn04 ; whereas the disproportionation of Mni+ gives Mn02 andMn2+ . 3Mn024'(ag) + 4H+(aq)^ Mn02(s) + 2MnO;(aq) + 2H20 £" = + 1.70 V S H E (3-4) 2Mni+(aq) + 2H20-> Mn2+(aq) +Mn02{s) + 4H+(aq) E° = + 0.54 V S H E (3-5) Acidic solution {[H+] = 1 M} : +7 +6 +4 +3 +2 0 0.56 V 7 2.26 V 0.95 V 1.49 V 7 + -1.18 V M n 0 4 M n 0 4 2 M n 0 2 M n 3 + — — M n 2 + Mn (purple) (green) (black) (red) (pale pink) | 1.70 V | | 1.23 V | 186 Figure 3-4 Latimar diagram for manganese species Although the temperature and concentrations of species involved in a standard EH — pH diagram differs from experimental conditions, it is commonly used as an initial indication of the identity of the predominant species under the influence of the pH and solution redox potentials. Hence, the prediction of possible reaction products at leaching conditions is possible. The EH - pH diagram of the Mn-SO\~ -H20 system at 25°C reveals that Mn2+ ions are the major species over a wide area of pH and potential combinations (Figure 3-5). The domain of stability of the manganese species (Mn04~, Mn04 , M« 3 + and Mn02) covers the whole of the upper part of the diagram. The Mn+ species exists as MnHS024+ species in sulfate medium. In other words, Guyard and the disproportionation reactions represented by equations (3-2) to (3-5) are not thermodynamically favored in acidic solution. Since manganese dioxide is stable only in very basic solution, these reactions are extremely slow under the test conditions where pH was consistently lower than 1.5. Since manganese(II) is the only stable manganese CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 81 species in strongly acidic conditions, complete reduction of permanganate to manganese(II) without the formation of manganese dioxide is expected at test conditions of low pH and potential. Similarly, the lack of permanganate ions in leaching solution confirmed the direct oxidation of chalcopyrite by ferric ions, which were readily available as shown by solution assay for soluble iron. • -2 0 . 2 4 6 8 . 10 12 14 pH Figure 3-5 EH - pH diagram for the Mn - S02~ - H20 system at 25°C, 0.1 187 dissolved manganese and unity sulfate activities. The absence of significant amounts of permanganate ions during leaching can be further demonstrated by Nernst equation. On the basis of Nernst equation, the solution redox potential is correlated with the logarithmic ratio of the concentration of permanganate and manganese(II) ions ([Mn04]: [Mn2+]) as shown in equation (3-6). EMIMa =1-3110- 0.0947/?// + 0.0118(log (At 25°C) (3-6) " " [Mn J A calculation at the maximum solution redox potential employed in this thesis work (versus silver/silver chloride reference electrode) E = 0.600 V A g / A g c i and pH = 1.5 yield \og([Mn04]/[Mn2+]) = -48. This low value is in accordance with the EH - pH diagram and CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 82 consistent with the conclusion discussed above. The kinetics of the complete reduction of permanganate to manganese(II) ions is rapid at temperatures higher than 25°C. It is therefore reasonable to conclude that the role of permanganate ions merely served to regenerate ferric ions by the oxidation of ferrous ions at the given potential. The EH - pH diagram also demonstrates that potassium permanganate solutions are not perfectly stable at 25°C. Permanganate, with a standard reduction potential (E° ) above 1.3 volts, oxidizes water to oxygen. A freshly prepared solution of potassium permanganate is deep purple but turns brown on long standing, because permanganate ion tends to oxidize water to oxygen 182 with the evolution of oxygen and separation of Mn02 that is brown: AMnO; {ag) + 2H20(l) -» 4Mn02 (s) + 302 (g) + 40H~(ag) (3-7) purple brown Equation (3-7) is catalyzed by Mn02(s) and is thus autocatalytic. This reaction is fhermodynamically possible, but proceeds extremely slowly. The preparation of potassium permanganate solution by dissolving reagent grade potassium permanganate is therefore possible. A properly stored solution of 0.02 M KMn04 decomposes = 0.1% in several months.182 The oxidation of water by Mn04 is very slow in the absence of such agents as Mn02, Mn2+, heat, light, acids and bases. The standardization procedure for potassium permanganate solution with oxalic acid (H2C204) in strongly acidic solution (pH = 1) is attached in Appendix C-2. The reproducibility of this method has been established experimentally. CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 83 3.3.2 Determination of Percent Extraction and Sulfur Species Solution samples and leach residues from each test were analyzed to determine copper extraction, iron extraction and sulfide oxidation. Copper (CufJ™a'ic°Head ) a n a " ^ r o n extractions (Fe^cnicHemi) were determined using the average of the calculated head values from all tests, where copper and iron in the final residue and final leachate were included in the calculations. Calculated head grades and extractions for every test are listed in Table B - l . The average calculated head copper grade (Cu 4va0,,cHead) of the starting material is 29.4 + 1.3%, which is close to the initial copper grade measured at 30.40%. The average calculated head iron grade (FeAvgCalcHead)'s 32.8 ± 2.3%, which is a bit off from the iron grade measured at 29.57%. The prediction of copper extractions based on the consumption of potassium permanganate solution is slightly higher than the calculated copper extraction in the range of 0.29-6.25% as presented in Table B-3. The total soluble iron concentration was calculated by subtracting the ferrous concentration of the initial standard FeS04 -H2S04 leaching medium (2.95 - 3.13 g/L Fe2+ ). The volume of the leaching medium at the sample time was adjusted for the rate of evaporation and the amount of solution removed and added during sampling was adjusted accordingly in all relevant calculations. The average rate of evaporation (RFvt,pomlioll) was calculated based on the differences between the initial and final solution volume divided by the leach time. Although evaporation was more rapid at 78°C than 70°C, a shorter leach time at 78°C rendered a lower average rate of evaporation. The average rates of evaporation at 60, 70 and 80°C were 0.04, 0.09 and 0.06 mL/min, respectively. A l l analyses were done at room temperature (20°C) before and after the leach tests. Since the density of water exhibits little variation between 20-30°C (1.0018-1.0043 g/cm3), the density of water at 20°C is taken as 1.00 g/cm 3in all relevant calculations.188 Total sulfide oxidation ( So~llllllj0n ), total elemental sulfur formation ( S2~ ) and total sulfate formation (S04')F were also calculated for every test as shown in Table B-2. Sample calculations of all types involved in this thesis work can be found in APPENDIX A. CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 84 3.3.3 Statistical Analysis of Experimental Results Many uncontrolled (and even unknown) variables can give rise to both systematic and random errors in this kind of work. Sources of random error include things such as weighing, material transfers and instrument readings, whose values fluctuate without bias. Systematic error refers to the use of equipment, techniques or the ambient environment, which renders values consistently, biased either high or low. Although the repetition of experiments does not necessarily curtail systematic and random errors reflected in these experimental data, each leach test was repeated three times under the same operating conditions to ensure reliability of experimental results, reduce experimental uncertainties and minimize variations among variables. An average value of the three sets of repeated experiments is used throughout this thesis work. The average of repeated results permits the assessment of the statistical significance of the measurements, including standard deviation (s) and standard errors (se). The experimental standard deviation (s) provides the scatter of individual data points around the experimental mean with a measurable level of confidence for the prediction of experimental means in future experiments.189 The range of values from (mean - s) to (mean + s) for one standard deviation represents a confidence level of 95% that the next measurement will fall in that range. Since only three data points are involved in the calculations of individual copper and iron extractions, the predictive value of the reported error can be improved by dividing the standard deviation with the square root of the number of data points to provide the standard error of the experimental mean (se). On the other hand, each datum on an extraction-time plot represents the copper and iron extraction achieved at a given point in time when a sample was taken. Many different factors, including instrumental errors of AAS such as flame fluctuation, sampling frequency and uncertainty, can contribute to the errors of the datum. Therefore, the reported error that was calculated based on the final extraction values of three repeated tests is taken to better represent the error bars of all data points on the extraction-time plots of the repeated tests. A larger error value accounts for a larger range of uncertainties. The calculated standard deviation (s) is 1.73 times (or the square root of 3 due to 3 data points) larger than the value of standard error (se). Therefore, the larger value associated with the standard deviation (s) is deemed more appropriate to represent the error bars on the graphs of copper and iron extractions than the standard error (se). CHAPTER 3 E X P E R I M E N T A L PROCEDURES / 85 The calculations for copper (Cu*™cT„m d) and iron (Fe^^°"lend ) extractions are based on the average of the calculated head values from all 45 tests. Calculations dependent on calculated head could be subjected to the compound effect of various analytical or procedural errors. Since calculated head determination involves a copper (or iron) mass balance, any large errors are usually noticed, unless the errors are biased in one direction. The standard deviation (s) of the average calculated head copper grade (CuAvoCalcHead) of the starting material is +1.3%, which means 95% of the measurements fall between 28.1% and 30.7%. The standard deviation (s) of the average calculated head iron grade (FeAvaCalcHend) is +2.3%, which means that 95% of the measurements falls between 30.5% and 35.1%. The standard errors of the copper and iron extractions are in the range +0.29-5.28% and ±0.48-4.06%, respectively. The experimental average (X), variance (s2), standard deviation (s) and the standard error (se) of copper and iron extractions are shown in Tables B-4 and B-5 of Appendix B, respectively. Sample calculations of statistical analyses can be found in Appendix A-10. The analyses of sulfur species are performed on leach residues from all 45 tests. The experimental average (X ) , variance (s2), standard deviation (s) and the standard error (se) of the total sulfide sulfur oxidation (S2Jxjdnljm), the total elemental sulfur formation (S2„) and the total sulfate formation [(SOA~)F] are shown in Tables B-6 to B-8, respectively. The standard errors of the determination of sulfur species are less than ±6% in all cases. Sample calculations of statistical analyses can be found in Appendix A-10. The influence of solution redox potential at each temperature is estimated from the differences between extraction values at a higher solution redox potential and a lower solution redox potential (0.400 V A g / A g c i in this case). Similarly, the effect of temperature on solution redox potential can be estimated in terms of extraction values relative to a lower temperature (60 and 70°C in this case). Statistical interpretations of the effect of temperature and solution redox potentials on copper dissolution can be found in Tables 4-8 and 4-9 of Chapter 4. Similar interpretations can be found in Tables 4-2 and 4-3 for total sulfide oxidation, Tables 4-4 and 4.5 for total elemental sulfur formation, and Tables 4-6 and 4-7 for total sulfate formation. Sample calculations of the statistical interpretations of temperature and controlled redox potential are shown in Appendices A-8 and A-9, respectively. CHAPTER 4 RESULTS A N D DISCUSSION / 86 C H A P T E R 4 R E S U L T S A N D DISCUSSION The experimental results are presented and discussed in this chapter. A number of experiments were conducted under a nitrogen atmosphere in order to examine the relationship between temperature (60-78°C) and solution redox potential (0.400-0.600 V A g / A g c i ) on the leaching kinetics of chalcopyrite in acidic ferric sulfate solution. 4.1 Reproducibility and Consistency of Leach Tests The anodic dissolution of chalcopyrite is notorious for its non-reproducibility. The comparison of the effect of temperature and solution redox potential on copper dissolution between different leach tests is possible only i f similar, i f not reproducible, leaching curves are attainable. Each leach test was repeated three times under the same reaction conditions to examine the issues of reproducibility and consistency. A summary of copper and iron extractions for the three sets of repeated leach data (set #l, set #2 and set #3) can be found in Table B - l . The results of overall copper and iron extractions for leach tests that were repeated at 60, 70 and 78°C are plotted in Figure 4-1. The standard deviations (s) of these leach tests as represented by the error bars in these graphs can be found in Tables B-4 and B-5. The standard deviation of the copper and iron extractions is in the range of ±0.50-9.15%. The standard deviation of the copper and iron extractions is in the range of ±0.83-7.03%. Discrepancies among repeated tests could be quite large in some cases, but these extraction values are within the acceptable range of standard deviations. A consistent trend of the copper and iron dissolution rate as a function of potential at each temperature is clearly visible. Therefore, average values of copper and iron extractions from repeated tests can reasonably represent typical leaching behaviors of chalcopyrite as a function of temperature and potential. CHAPTER 4 RESULTS A N D DISCUSSION / 87 100% 90% 80% 70% e JO 60% u 1. 50% 40% 30% 20% e o 5^  Iron 100% 90% 80% h 70% 60% 50% 40% 30% 20% L 100% 90% 80% 70% 60% -50% -40% -30% -20% (b) 70°C (c) 78° C Iron 0.400 V 0.450 V 0.500 V 0.550 V Solution Redox Potential (V A g M g C i ) 0.600 V -X—Cu.Sel#] —±-Cu,Set#2 -+—Cu,Set#3 -*r-Fe,Set#l —±-Fe,Set#2 —*—Fe,Sel#3 Figure 4-1 Comparison of total copper and iron extractions as a function of potential (0.400-0.600 V A g / A g c i ) and temperature from different data sets (set #1, set #2, set #3): (a) 60°C; (b) 70°C; (c) 78°C. CHAPTER 4 RESULTS A N D DISCUSSION / 88 4.2 Acid Soluble Copper and the Effect of Initial Ferrous Iron Concentration The sample chalcopyrite mineral was introduced into the initial standard ferrous sulfate [ Fe(S04) -H2S04 ] leaching medium at start-up. After the introduction of chalcopyrite, the standard leaching medium in all tests was an aqueous solution of ferrous sulfate containing 2.95-3.13 g/LFe2+, 8.8 g/L H2SOA and 5 g/LCuFeS2. Reaction time zero corresponded to the time when the temperature was stabilized before the initiation of the potassium permanganate pump to set the setpoint of the desired solution redox potential. The copper at reaction time zero had to result from acid soluble copper present in the chalcopyrite samples, which dissolved into the starting leaching medium during the warm-up period. The low level of copper at reaction time zero confirmed that ferrous oxidation of chalcopyrite was not possible as anticipated, which agreed with preliminary examinations. Preliminary runs showed that the amount of ferrous sulfate in the initial leaching medium had little effect on the dissolution of chalcopyrite, rendering less than 5% copper extraction at temperature after 3 hours of leaching. An oxidant such as ferric iron (see Section 2.3.1) is required for the leaching of chalcopyrite to take place. Therefore, the introduction of sample mineral with ferrous sulfate leaching medium at start-up was a safe practice implemented for all the leach tests in this thesis work. At 60°C, acid soluble copper at reaction time zero ranged from 15.1 to 49.5 mg/L (or 0.48 to 1.58 mmol Cu). At 70°C, solution copper ranged from 30.8 to 42.3 mg/L (or 0.98 to 1.35 mmol Cu). At 78°C, solution copper ranged from 10.3 to 38.1 mg/L (or 0.33 to 1.22 mmol Cu). Since the maximum leachable copper from 5 g/L of the chalcopyrite used in this work is about 1.50 g/L, the maximum acid soluble copper represents 1 -4% of total leachable copper. CHAPTER 4 RESULTS A N D DISCUSSION / 89 4.3 Analysis of Leach Residues Quantitative (chemical analysis) and qualitative (XRD and SEM analyses) examination of leach residues are presented and discussed in this section. Attention is given to the analysis of sulfur species in the leach resides, followed by the inspection of relationships between copper extraction and other experimental results. The statistical analyses of copper and iron residues are presented in Sections 4.4.1 and 4.4.2, respectively, with the discussions on the rate of copper and iron dissolution. 4.3.1. Analysis of Sulfur Species in the Leach Residues Sulfur assay on the leach residues is a good measure of the degree of oxidation of chalcopyrite, and the results for all tests can be found in Table B-2. Leaching of chalcopyrite with ferric sulfate in acid solutions produces elemental sulfur according to equation (2-15). As the leaching temperature increases, the degree of oxidation increases, as indicated by the enhancement of sulfide sulfur oxidation and elemental sulfur/sulfate formation with temperature (Figure 4-2). The leach residues contain sulfate (5-33%), but primarily elemental sulfur (46-78%). The presence of orthorhombic elemental sulfur ( S a ) is supported by XRD. Since only orthorhombic sulfur ( S a ) is soluble in perchloroethylene ( C 2 C / 4 ) (see Section 3.2.3), the values of elemental sulfur (5 8 ) extraction should be regarded as minimum to account for the amorphous sulfur which can exist in leach solutions. The remnants of elemental sulfur reaction product in the leach residues implies that it is not significantly attacked by the ferric sulfate leaching medium under the reaction conditions, which is supported by fhei i w - pH diagram presented in Figure 2-2. The figure illustrates the stability of the elemental sulfur in the presence of copper and iron species from mildly oxidizing (low pH) to moderately reducing (high pH) conditions. The trace amount of sulfate formed during leaching cannot be determined directly as a result of the high background sulfate concentrations. Alternatively, the production of sulfate anions is estimated indirectly by the differences between total sulfide oxidation and total elemental sulfur formation. The majority of the sulfate sulfur is present in the form of jarosite in the leach CHAPTER 4 RESULTS A N D DISCUSSION / 90 residues due to iron hydrolysis as confirmed by XRD. Elemental sulfur is hardly oxidized, i f at all, at temperatures below its melting point (119°C). 8 5 Therefore, sulfate anion is expected to form directly from sulfide sulfur during anodic breakdown of chalcopyrite. The XRD qualitative diffraction patterns were used to examine the leach residues of data set #2, and the results are summarized in Table 4-1. XRD patterns of partly leached particles display the same significant peaks as were identified in the original chalcopyrite (Figure 3-2). A representative set of diffraction patterns for the leaching reaction at 60°C under a potential of 0.500 V A G / A g c i is shown in Figure 4-3. XRD analysis reveals that leach residues are formed mainly of unleached chalcopyrite, pyrite, orthorhombic sulfur ( Sa) and jarosites. Residual chalcopyrite from incomplete leaching was always observed in the residues, except at 78°C under a potential of 0.500 V A G / A g c i , indicating near complete copper extraction. There was no evidence for the formation of pyrite-type S2" groups aside from the 5% pyrite already present in the starting material. Pyrite (FeS2) dissolution increased with temperature and solution redox potential. Since the rest potential of pyrite in acidic solutions is 0.630 V S H E at 25°C (Table 2-10), dissolution at potentials above 0.500 V A g / A g C i (0.645-0.664 V S H E at 60-78°C, see Appendix A - l 1) is expected. Orthorhombic sulfur (Sa) was always observed in the leach residues in accordance with equation (2-15). Jarosite becomes the dominant phase at high potentials and temperatures. At 60°C, jarosite precipitates do not form until reaching 0.500 V A G / A g c i , but at 70-78°C, the initial presence of jarosites in leach residues is at 0.450 V A G / A g c i - This result is in accordance with the observations of the orange-brown color of the leaching solution at potentials above 0.500 V A G / A g c i - Alkali jarosites have the general formula of AFei(S04)2(OH)b, where A = K+ for potassium [ KFe3(S04)2(OH)b ] and A = H30+ for hydronium [ {H\0)Fe,(S04)2(OH)b ] jarosites, respectively. Both types of jarosite were discovered in the leach residues. The addition of potassium permanganate (KMn04) solution, which introduced potassium impurities into the solution during leaching, supports the presence of potassium jarosite. Moreover, the formation of potassium jarosite in sulfate based leach solution should precede the formation of hydronium jarosite as already discussed in Section 2.2.2. Dutrizac 1 9 0 has noted the incorporation of cupric ion (Cw 2 + ) by substitution of the ferric ion (FeM) in the lattice structure of jarosite. While this incorporation is usually minor, the entrainment of copper in the jarosite portion of the leach residue is a distinct possibility. CHAPTER 4 RESULTS A N D DISCUSSION / 91 Statistical interpretations of the total sulfide oxidation, total elemental sulfur formation and total sulfate formation as a function of increasing temperature are shown in Tables 4-2, 4-4 and 4-6, respectively. A minimum gain of 10% in oxidation efficiency, 6% in elemental sulfur formation, and about 2% in sulfate formation is seen as a function of increasing temperature from 60 to 78°C. Temperature increment has its largest impact on both the sulfide oxidation and elemental sulfur formations at 0.500 VAg/Agci- 1° the case of sulfate formation, the major impact of temperature increment occurred at 0.500 VAg/Agci from an increase of 60 to 70°C, and shifted to 0.600 VAg/Agci as temperature increases from 70 to 78°C. The values of oxidation efficiency or elemental sulfur/sulfate formation are not always proportional to temperature increment. For example, a 10°C change from 60 to 70°C yields an enhancement of 16.99% oxidation efficiency, whereas an 8°C change from 70 to 78°C yields an increase of only 2.05% both at 0.500 V A g/Agci -Similar scenarios are also available for elemental sulfur/sulfate formation. This suggests that the leaching of chalcopyrite is a more complicated phenomenon than a simple function of temperature, and associates immensely with solution redox potentials. Statistical interpretations of the total sulfide oxidation, total elemental sulfur formation and total sulfate formation as a function of increasing solution redox potential are shown in Tables 4-3, 4-5 and 4-7, respectively. The optimal sulfide oxidation efficiency relative to 0.400 VAg/Agci occurs at 0.550 VAg/Agci for. leaching between 60 and 78°C. Maximum elemental sulfur formation occurs at 0.550 V A g /A gci for the 60°C test and at 0.500 V A g /A gci for both the 70 and 78°C tests. In the case of sulfate anions, the maximum formation occurs at 0.600 V A g /A g c i for leaching between 60 and 78°C. The low values of oxidation efficiency and elemental sulfur formation along with the high values of sulfate formation at 0.600 V A g /A gci suggest the possible contribution of high solution redox potential towards the passivity of chalcopyrite. Consequently, leaching at 0.600 VAg/Agci should be avoided at all temperatures. On the other hand, the effect of increasing solution redox potential from 0.400 to 0.600 VAg/Agci yields a negative value of -0.06% at 70°C, which becomes an anomaly when oxidation efficiencies of 10% are gained for both 60 and 78°C tests under similar conditions. An explanation cannot be provided for this anomaly. CHAPTER 4 RESULTS A N D DISCUSSION / 92 Table 4-1 Summary of XRD results of leach residues from data set #2 E (VAg/AgCl) 60°C 70°C '78°C 0.400 Cp,Py, So Cp,Py, So Cp,Py, So 0.450 Cp,Py,So Cp , Py, So, Jar Cp ,So, Jar 0.500 Cp ,Py ,So, Jar Cp , Py, So, Jar So, Jar 0.550 Cp ,So, Jar Cp ,So, Jar Cp ,So, Jar 0.600 Cp ,So, Jar Cp ,So, Jar Cp ,So, Jar Cp: Chalcopyrite (Tetragonal) Py: Pyrite (Cubic) So : Elemental Sulfur (Orthorhombic) Jar : Hydronium or Potassium larosite (Rhombohedral) Table 4-2 Statistical significance of temperature dependency on total sulfide oxidations relative to that of 60 and 70°C at various potentials (0.400-0.600 VAg/Agci) E (VAg/AgCl) 5 ^ ( 6 0 ^ 70° C) 5 ^ ( 6 0 - * 78° C) C , a , , ( 7 0 ^ 7 8 ° C ) 0.400 9.26% 20.16% 10.90% 0.450 10.62% 11.13% 0.51% 0.500 16.99% 19.03% 2.05% 0.550 5.72% 10.56% 4.84% 0.600 -1.12% 20.33% 21.45% Table 4-3 Statistical significance of redox potential dependency on total sulfide oxidations relative to that of 0.400 VAg/Agci at various temperatures (60-78°C) E SaMon{0A00VAglAgO^E2) (VAg/AgCl) 60°C 70°C 78°C 0.450 19.75% 21.11% 10.72% 0.500 13.75% 21.48% 12.63% 0.550 28.20% 24.67% 18.61% 0.600 10.32% -0.06% 10.50% CHAPTER 4 RESULTS A N D DISCUSSION / 93 Table 4-4 Statistical significance of temperature dependency on total elemental sulfur formations relative to that of 60 and 70°C at various potentials (0.400-0.600 VAg/Agci) E (VAg/AgCl) 5 j ( 6 0 - > 7 0 ° C ) 5 ^ ( 6 0 - » 78" C) S £ ( 7 0 - > 7 8 ° C ) 0.400 5.00% 11.93% 6.93% 0.450 5.93% 9.22% 3.28% 0.500 9.89% 12.88% 2.99% 0.550 1.78% 6.23% 4.46% 0.600 1.24% 9.78% 8.54% Table 4-5 Statistical significance of redox potential dependency on total elemental sulfur formations relative to that of 0.400 VAg/Agci at various temperatures (60-78°C) E (VAg/AgCl) S2-(0A0QVAglAgCl^E2) 60°C 70°C 78°C 0.450 0.500 0.550 0.600 14.84% 15.78% 12.13% 11.78% 16.67% 12.73% 17.81% 14.59% 12.11% -7.13% -10.88% -9.27% Table 4-6 Statistical significance of temperature dependency on total sulfate formations relative to that of 60 and 70°C at various, potentials (0.400-0.600 VAg/Agci) E (VAg/AgCl) (5O 4 2 - ) , . [60^70°C] ( ,SO 4 2 -) F [60^78 o C] (SO24-)F[70^18°C] 0.400 4.26% 8.23% 3.97% 0.450 4.68% 1.91% -2.77% 0.500 7.09% 6.15% -0.94% 0.550 3.95% 4.33% 0.38% 0.600 -2.36% 10.55% 12.91% Table 4-7 Statistical significance of redox potential dependency on total sulfate formations relative to that of 0.400 VAg/Agci at various temperatures (60-78°C) E « ) [ 0 . 4 0 0 ^ / 4 ( O - » £ 2 ] (VAg/AgCl) 60°C 70°C 78°C 0.450 4.91% 5.33% -1.41% 0.500 1.97% 4.81% -0.10% 0.550 10.39% 10.08% 6.49% 0.600 17.45% 10.83% 19.77% CHAPTER 4 RESULTS A N D DISCUSSION / 94 s C « •o c B 0 3 100% 90% a 80% o cs 70% B -Q 60% U. In 50% <2 3 40% "3 30% s 20% £ 10% s 0% Figure 4-2 0.400 V 0.450 V 0.500 V 0.550 V 0.600 V 0.400 V 0.450 V 0.500 V 0.550 V 0.600 V 0.400 V 0.450 V 0.500 V 0.550 V Solution Redox Potential (V A g / A g C I) • 60°C 0.600 V • 70°C • 78°C Total sulfide oxidation (S 2 0 x j i l a l j o i i), total elemental sulfur formation (S2~) and total elemental sulfate formation [(S02')F] as a function of potential (0.400-0.600 VAg/Agci) and temperature (60-78°C) CHAPTER 4 RESULTS AND DISCUSSION / 95 . 3 X -4'. -I c O ^ •§ P j? o C M t . <N O CO s a :I5 • 0) 5? J= > 5 ° N a a tn o " , co CO ' r*- o ^ O CN co C N -cj _1 - LL C ™ . « g a I ^ 'S s  -g f i - c | •2 8 & S B <D j? ro O ^ tt: o -c ro 3 o j= ™ O -s oo c 1 s a HEEB u 00 < < > O o IT) o t+M o o CM cS IM <U -o c U o O NO CM T3 o CD 03 CM m i tM (sjunoo) un CHAPTER 4 RESULTS A N D DISCUSSION / 96 4.3.2. Morphology of Leach Residues The morphology of the leach residues performed at temperatures ranging from 60 to 78°C from data set #2 was studied by SEM. The leaching residue samples were thoroughly acid-washed and air-dried at room temperature before inspection. Quick surface morphological examination at both high and low magnifications shows that the leach residues consisted of different particle sizes. The diameter of the particles leached at 78°C is generally smaller than those leached at 60°C. In addition to copper, iron and sulfur, the spectra show a peak corresponding to zinc, which occurred as an impurity in the original chalcopyrite. At the lowest potential of 0.400 V A g / A g c i at 60-78°C, the leached residue show no significant changes of the chalcopyrite surface aside from the appearance of sulfur globules at numerous sites on the surface. Only small amounts of elemental sulfur and jarosites are present, and this is a direct reflection of the low rate of dissolution of chalcopyrite samples. There is no indication of surface pitting on leached surfaces at this potential. However, extensive pitting and a great abundance of jarosite precipitates on leached surfaces are observed at potentials above 0.400 V A g / A g c i . Selected backscattered electron micrographs of the leach residues of tests performed at 78°C are shown from Figures 4-4 to 4-7 over the potential range of 0.450 to 0.600 V A g / A g c i . An appreciable decrease in the exposed surface area of chalcopyrite as a result of increasing jarosite precipitation at increasing potential from 0.450 to 0.550 V A g / A g C ] is visible from these figures. The chalcopyrite core is almost entirely enveloped in a compact film of jarosite at 0.550 V A g / A g c i (Figure 4-6). The formation of jarosite on the mineral surface can hinder and even prevent both the transport of ferric ions to the chalcopyrite and the diffusion of the reaction products to the solution. At 0.450 V A g / A g c i , pitting is highly selective, leaving the greatest part unchanged in appearance (Figure 4-4). These areas have yet to be identified but it seems most likely that they are in certain crystal lattice defects, cracks or grain boundaries. In this micrograph, the presence of sulfur near the pitted area is confirmed by the IMIX-PC energy dispersive system on the SEM, and is shown to be elemental sulfur globules as supported by XRD. Relatively large amounts of tiny sulfur globules are clearly visible in Figure 4-5, and they are non-protective. Thick, compact sulfur layers that may protect the chalcopyrite surface from the attack of oxidants are not observed on the leached surface at any potential. The presence of more ferric ions at 0.600 V A g / A g c i than other potentials may have facilitated the formation of extensive pitting on the CHAPTER 4 RESULTS A N D DISCUSSION / 97 mineral surface. It appears at 0.600 VAg/Agci ferric attacks the entire surface of chalcopyrite, whereas at 0.450 VAg/Agci ferric attacks only in selected areas. The high concentration of ferric ions in the solution should favor jarosite formation, whose presence is confirmed by XRD diffraction. Unlike at other potentials, at 0.600 VAg/Agci jarosites are agglomerated into larger masses than the surrounding chalcopyrite particles. If the rate of leaching in ferric sulfate depends on the total surface area exposed to the solution, leaching at 0.600 VAg/Agci should render the highest copper extractions. Nevertheless, low copper extractions are observed for this potential even though there are no jarosite precipitates or elemental sulfur to block the access of ferric to the unreacted mineral surface. Clearly, neither jarosite nor elemental sulfur is responsible for surface passivation, or for preventing the dissolution of chalcopyrite. CHAPTER 4 RESULTS A N D DISCUSSION / 98 Figure 4-4 < *20 um Figure 4-5 Backscattered electron micrograph for chalcopyrite leached at 78°C under a potential of 0.500 V A g A g c i after 17 hours leaching CHAPTER 4 RESULTS A N D DISCUSSION / 99 < K20 um Figure 4-6 Backscattered electron micrograph for chalcopyrite leached at 78°C under a potential of 0.550 V A g / A g c i after 7 hours leaching * *-20 um Figure 4-7 Backscattered electron micrograph for chalcopyrite leached at 78°C under a potential of 0.600 V A g / A g C i after 9 hours leaching CHAPTER 4 RESULTS A N D DISCUSSION / 100 4.3.3. Copper Extraction and Other Experimental Results The relationships between sulfide oxidation, elemental sulfur/sulfate formation, and copper/iron extraction at 60 to 78°C and 0.400 to 0.600 V A G / A g c i are shown in Figure 4-8. Higher rates of copper/iron extraction and sulfide oxidation are observed at the higher temperatures. At 0.400 V A g / A g c i (0.545 to 0.564 V S H E at 60 to 78°C, see Appendix A - l l ) , low copper extraction and low sulfide oxidation indicate that this potential is not sufficient to leach the chalcopyrite. A higher potential than 0.400 V A g / A g c i is necessary to leach chalcopyrite, which has the rest potential between 0.450-0.550 V S H E at 25°C (Table 2-10) in acidic solutions. It is immediately clear that the oxidation rate is far more rapid at 0.500 V A g / A g c i than at other potentials between 60 and 78°C. In addition, iron extractions are found to be consistently lower than the sulfide oxidation. There is less than 10% increase in copper extraction as potential increases from 0.500 to 0.550 V A g / A g c i . Furthermore, an abrupt decline in the values of copper and iron extraction and sulfide oxidation are observed as solution redox potential increases from 0.550 to 0.600 V A g / A g c i -This reveals the existence of a potential window outside of which sulfide oxidation and copper extraction do not increase with increasing potential. The results support the hypothesis that passivation is related to potential. The apparent lag in both iron extraction and sulfide oxidation, which peaks at 0.550 V A g / A g C i rather than 0.500 V A g / A g C i , could be due to a passivating mechanism that oxidizes the sulfide sulfur to sulfate rather than elemental sulfur. Above 0.500 V A g / A g c i , this passivating mechanism has a bigger influence on copper dissolution, which gradually has a further impact on both iron dissolution and sulfide oxidation at potentials above 0.550 V A g / A g c i . In general, a low level of sulfate formation is shown for the leaching of chalcopyrite in acidic ferric sulfate medium. The formation of sulfate is proportional to increasing potential and temperature. Approximately 5% or less of the sulfur reports as sulfate as potential increases from 0.400 to 0.500 V A g / A g C i . Sulfate in the residue of the 0.600 V A g / A g C i test was at least twice that found in the residue of the 0.400 V A g / A g c i test (Figure 4-2). However, comparison of the 0.400 and 0.600 V A g / A g c i tests shows that corresponding increases in extraction values of copper and iron, sulfide oxidation and elemental sulfur formation are not observed with the additional gain in sulfate formation as solution redox potential increases from 0.400 to 0.600 V A g / A g c i . For example, at 60°C, 66% copper extraction is associated with 5% sulfate formation at 0.400 V A g / A g c i , whereas 56% copper extraction is associated with 23% CHAPTER 4 RESULTS A N D DISCUSSION / 101 sulfate formation at 0.600 V A g /Agci- Similar scenarios are also observed at both 70 and 78°C. Low copper extractions at higher potentials (0.500-0.600 V A g /Agci) indicate that the formation of a passivating layer leads to a decrease in the dissolution rate. A different dissolution mechanism of chalcopyrite may be operative that leads to the eventual formation of the passivating layer, but not at lower potentials. This mechanism converts the sulfide moiety of chalcopyrite to sulfate rather than elemental sulfur as supported by the increase in sulfate formation. In summary, the passivating layer is much less protective at 0.500 V A g / A g c i than other potentials between 60 and 78°C. CHAPTER 4 RESULTS A N D DISCUSSION / 103 4.4 Leaching Kinetics: Interpretation of Batch Reactor Data The effect of solution redox potentials and temperature on the dissolution rates of copper and iron are examined in this section. Among three repeated sets of experiments, results from data set #2 are used to illustrate the general behavior of copper and iron dissolution as a function of time. 4.4.1 Effect of Solution Redox Potentials and Temperature on the Dissolution Rate of Copper The leaching curves showing the release of copper as a function of time at individual temperatures and solution redox potentials are presented in Figures 4-9 to 4-11. The graphs clearly indicate the fact that the rate of copper dissolution of chalcopyrite in acidic sulfate media increases with temperature from 60 to 78°C. Shorter leach times indicate faster rates of copper dissolution. Leach times of individual experiments are listed in Table B - l . The minimum leach time of 60°C tests was at least 50 hours before copper dissolution ceased, whereas the leach time required for 70 and 78°C tests was between 16 and 24 hours. Hence, different leach times confirm the temperature dependency of copper dissolution from chalcopyrite. A better perception of the influence of temperature at different potentials is provided from Figures 4-12 to 4-16. The effect of temperature on copper extraction at different potentials as depicted by Table 4-11 validates once again the increase of copper extraction with increasing temperature. Statistical interpretation of the final copper extraction as a function of temperature on different solution redox potentials is shown in Table 4-8. The temperature dependency of copper extraction is clear, particularly at 0.600 V A g / A g C i > which raises the copper yield 9% at 70°C and 18% at 78°C relative to that of60°C. Comparison of Figures 4-9 and 4-10 shows that reaction curves in these figures have similar shapes, and resemble bacterial leaching curves with their characteristic lag phase during incubation, followed by the growth and stagnant death phases. At 78°C, the resemblance of leaching behavior to a bacterial leaching curve is not seen in Figure 4-11 with dissolution proceeding more rapidly initially. Clearly, the exhibition of an induction period is a temperature CHAPTER 4 RESULTS A N D DISCUSSION / 104 dependent phenomenon, which takes place at 60 and 70°C, but not 78°C. The induction period is when leaching proceeds at a slow but measurable rate. Although at different temperatures, it has been reported by Kametani and A o k i 1 6 3 (see Figure 2-12) and Van de Merwe et a/. 1 6 8 (see Figure 2-14) as already discussed in Sections 2.6.2.1 and 2.6.2.2, respectively. The dependency of the induction period on potential still remains debatable. Induction could play an important role in removing the passivating layers on the chalcopyrite surface. Since this induction period is similar to the early stage of bacterial leaching behavior, where enough bacterial species have to build up before leaching is possible, it is possible that enough ferric have to cover the chalcopyrite surface in this stage to initiate leaching. Although the induction period is not present at 78°C, leaching mechanisms at 60, 70 and 80°C should be similar due to their apparent potential dependence. Attempts were made to deduct the initial rate constants by drawing tangent lines to the copper extraction curves during the early stages of leaching. However, the induction period represents a very slow initial leaching rate, and a very low initial extraction. Although tangent lines were drawn to the rising portion of copper extraction curves and the induction period were ignored where appropriate, the results were disappointing. The initial rate constants obtained for each solution redox potential are random, non-uniform and inconclusive. Statistical interpretation of the final copper extraction as a function of increasing solution redox potentials relative to 0.400 V A g /AgCi at different temperatures is shown in Table 4-9. It can be seen that the copper dissolution of chalcopyrite in acidic ferric sulfate medium is dependent on solution redox potential. Highest copper extractions are observed at 0.500 V A g / A g c i relative to those at 0.400 V A g / A g c i at each temperature, which is shown in Tables 4-9 and 4-11. The effect of increasing solution redox potential is particularly rewarding for the experiments performed at 60°C. Leaching under a potential of 0.500 VA g /A g c i increases the copper extraction by 26.06% for the 60°C test, but by only 21.72% for the 78°C test. The optimal leaching potential of 0.500 V A g / A g c i for maximum copper extraction in the temperature range of 60 to 78°C shows agreement within the critical redox potential range as reported by Kametani and A o k i 1 6 3 and Van de Merwe et al.168 (Table 4-10). The critical redox potential obtained from this work is in the vicinity of 0.500 V A g / A g c i (or 0.645-0.664 V S H E at 60-78°C) and is within the mid-potential range of 0.45 to 0.55 V A G / A g c i - Leaching at 0.500 V A g / A g C i yields a copper extraction of 70% after 13 hours at 60°C, after 9 hours at 70°C, and 6 hours at 78°C (Figure 4-14). Leaching at 0.600 V A g / A g C i is clearly undesirable and should be avoided. Indeed, as shown by Table 4-11, leaching at 0.400 V A g / A g a renders a higher copper extraction than leaching at 0.600 V A g / A g c i . CHAPTER 4 RESULTS A N D DISCUSSION / 105 At 60 and 70°C, the 0.450 and 0.500 V A G / A g c i curves are practically identical and exhibit parallel leaching kinetics and hence similar reaction rates. There is an abrupt change in the rate of leaching in the vicinity of 0.550 V A G / A g c i as temperature increases from 60 to 70°C. At 60°C, the rate of leaching at 0.550 V A G / A g c i is slow; however, the 0.550 V A G / A g c i curve is practically identical to 0.450 and 0.500 V A G / A g c i curves at 70°C as shown in Figures 4-9 and 4-10. Similar leaching behaviors are once again observed for 0.450, 0.500 and 0.550 V A G / A g c i curves at 78°C. The 0.600 V A G / A g c i curve at 78°C seems to share the same behavior in the early stage (the first 6 hours) of leaching, which reaches an eventual plateau with a slow rate of copper dissolution (Figure 4-11). A slight decline in copper extraction at 9 hours for the 0.600 V A G / A g c i curve is thought to be due to random experimental error associated with the analytical technique (Figure 4-16). Although increasing temperature has an enhancing effect on copper extraction under potentials of 0.550 and 0.600 V A G / A g c i , a 10°C increase from 60 to 70°C has a greater impact on dissolution kinetics than a corresponding 8°C increase from 70 to 78°C (Figures 4-15 and 4-16). Passivation of chalcopyrite is apparent at 60 and 70°C for the 0.600 V A G / A g c i curves, as indicated by low copper extraction (Figures 4-9 and 4-10). At 0.400 V A G / A g c i (0.545-0.564 V S H E at 60-78°C, see Appendix A - l 1), low copper extraction indicates that this potential is simply not sufficient to leach the chalcopyrite, and not a manifestation of passivity. However, leaching is incomplete even after 58 hours at 60°C and 22.hours at 70°C.with less than 70% of copper extraction for both 0.400 and 0.600 V A G / A G C | . Passivation behavior is less obvious at 78°C. Copper extraction of 76.31% is obtained after about 23 hours at 0.600 V A g / A g c i . Interestingly enough the copper extraction of 0.400 V A g / A g c i is consistently higher than that of 0.600 V A g / A g c i at 60°C at nearly every extraction point, but the reverse is shown at both 70 and 78°C. CHAPTER 4 RESULTS A N D DISCUSSION / 106 Table 4-8 Statistical significance of temperature dependency on total copper extractions relative to that of 60 and 70°C at various potentials (0.400-0.600 VAg/Agci) E (VAg/AgCl) C ^ S w ( 6 0 ^ 7 0 ° C ) C < c v ™ ( 6 0 ^ 7 8 ° C ) C ^ X / L / ( 7 0 - > 7 8 ° C ) 0.400 4.63% 10.66% 6.03% 0.450 6.31% 9.16% 2.85% 0.500 2.52% 6.33% 3.81% 0.550 1.43% 6.27% 4.84% 0.600 8.55% 18.39% 9.84% Table 4-9 Statistical significance of redox potential dependency on total copper extractions relative to that of 0.400 VAg/Agci at various temperatures (60-78°C) E2 Cu^L.(0.4O0VAglAgCl^E2) (VAg/AgCl) 60°C 70°C 78°C 0.450 18.89% 20.57% 17.39% 0.500 26.06% 23.94% 21.72% 0.550 21.08% 17.87% 16.69% 0.600 -9.91% -5.99% -2.18% Table 4-10 Comparison of experimental results with reported critical potentials for chalcopyrite leaching References Reported Critical Potentials Conversion to SHE scale Mintek-South African patent (ZA 9701307)168 0.414-0.434 VSCE at 70°C (0.474-0.494 V A g / A g c i at 70°C) 0.628-0.648 V S H E a t 70°C Kametani and A o k i 1 6 3 0.40-0.43 V S C E at 90°C (0.467-0.497 V A g / A g C i at 90°C) 0.600-0.630 VSHE at 90°C Experimental Results 0.45-0.55 V A g / A g c i at 60°C 0.45-0.55 V A g / A g C i at 70°C 0.45-0.55 V A g / A g C i at 78°C 0.614-0.714 VSHE at 60°C 0.604-0.704 VSHE at 70°C 0.595-0.695 V S H E a t 78°C CHAPTER 4 RESULTS AND DISCUSSION / 100%-90%-80% 70% 60% 50% 40% 30% 20% 10% 0% 0.400 V 0.450 V 85% 91% 94% 0.500 V 92% 95% 99% 0.550 V 87% 89% 93% 0.600 V 56% 65% 75% Solution Redox Potential (V A g/A gci) Table 4-11 Copper extractions as a function of potential (0.400-0.600 VAg/AgCi) and temperature (60-78°C) CHAPTER 4 RESULTS AND DISCUSSION / CHAPTER 4 RESULTS AND DISCUSSION / 109 CHAPTER 4 RESULTS AND DISCUSSION / 110 CHAPTER 4 RESULTS A N D DISCUSSION / 111 0 2 4 6 8 1012141618 2022 24 26 28 3032 34 36 384042 444648 50 52 54 56 586062 646668 70 Leaching Time (hr) Figure 4-12 Extraction-time plot for copper at 0.400 VAg/Agci and different temperatures (60-78°C) in ferric sulfate medium 0 2 4 6 8 10 12 14 16 18 202224 26 28 30 32 34 36384042 444648 5052 54 56 58 6062 64 6668 70 Leaching Time (hr) Figure 4-13 Extraction-time plot for copper at 0.450 V A g / A gci and different temperatures (60-78°C) in ferric sulfate medium CHAPTER 4 RESULTS A N D DISCUSSION / 112 0 2 4 6 8 10 12 14 16 18 20 22 24 26 28 30 32 34 36 38 40 42 44 46 48 50 52 54 56 58 60 62 64 66 68 70 Leaching Time (hr) Figure 4-14 Extraction-time plot for copper at 0.500 V A g / A g c i ) and different temperatures (60-78°C) in ferric sulfate medium 0 2 4 6 8 10 12 141618202224 26 28 30 3234 36384042 44464850 5254 56586062 64666870 Leaching Time (hr) Figure 4-15 Extraction-time plot for copper at 0.550 V A g / A g c i and different temperatures (60-78°C) in ferric sulfate medium CHAPTER 4 RESULTS A N D DISCUSSION / 113 0 2 4 6 8 10 12 14 16 18 20 22 24 26 28 30 32 34 36 38 40 42 44 46 48 50 52 54 56 58 60 62 64 66 68 70 Leaching Time (hr) Figure 4-16 Extraction-time plot for copper at 0.600 V A g / A g c i ) and different temperatures (60-78°C) in ferric sulfate medium CHAPTER 4 RESULTS A N D DISCUSSION / 114 4.4.2 Effect of Solution Redox Potentials and Temperature on the Dissolution Rate of Iron Leaching curves showing the release of iron as a function of time at various temperatures and solution redox potentials are presented in Figures 4-17 to 4-19. The graphs clearly indicate that iron extractions increased with temperature from 60 to 78°C as the runs progressed. The dissolved iron increased to some maximum before the occurrence of ferric iron precipitation. Ferric iron precipitation results mainly in jarosites as confirmed by XRD and SEM. larosite precipitation is a major concern in leaching processes, which is inevitable and is a characteristic of iron extraction under ferric sulfate leaching at temperatures exceeding 50°C. A better perception of the influence of temperature at various potentials is provided in Figures 4-20 to 4-24. The maximum iron concentration is reached faster at a higher temperature than at a lower temperature in all cases: 78°C > 70°C >60°C. The iron extraction curves fluctuate up and down during leaching. Different peaks on the curves can be designated for the ease of discussion and comparison. Table 4-12 shows the major iron extraction peaks at different temperatures and solution redox potentials, where peak#l precedes peak#2 in leach time. In situations where only one peak is available, the peak is assigned arbitrarily for the purpose of comparison. The maximum solution iron concentration can occur at peak#l, peak#2 or at the end of leaching as represented by the final iron extraction from the leach tests. Comparison of peak#2 shows a general trend that increasing temperature yields an overall higher rate of iron extraction. The transition from peak#l to peak#2 in Figure 4-17 can be the result of either the reductive decomposition of the ferric sulfates (e.g., jarosites) leached into solution, or the further dissolution of iron from the sample. At 60°C, the maximum total iron dissolution is in the range of 49-55%. At 0.400 and 0.550 V A g / A g c i , the maximum iron dissolution represents final iron extraction after 59 hours and 52 hours, respectively. Other curves decrease to a lower iron extraction after reaching a peak due to iron precipitation. Approximately 10% iron has been precipitated at these potentials. The maximum iron dissolution is reached at different potentials in this order of time: 0.450 V > 0.600 V > 0.500 V > 0.550 V > 0.400 V A g / A g a . However, this result is not expected since it is logical to infer that the higher the redox potential, the faster the mineral will oxidize under the prevailing conditions unless surface passivation occurs at high potentials. At 70°C, the maximum total solution iron is in the range of 41-60%. The maximum CHAPTER 4 RESULTS A N D DISCUSSION / 115 iron dissolution of the 0.450 V A g / A g c i curve occurred at the end of the leach test. Approximately 3-8% iron was precipitated at the other potentials at this temperature. The maximum iron dissolution is reached at different potentials in this order of time: 0.600 V > 0.550 V > 0.500 V > 0.400 V > 0.450 V A g / A g c i - This result is anticipated where the highest potential (0.600 V A g / A g c i ) oxidizes the mineral at the fastest rate. At 78°C, the maximum total iron dissolution is in the range of 44-61%. Approximately 2-10% iron has been precipitated at the other potentials at this temperature. The maximum iron dissolution is reached at different potentials in this order of time: 0.550 V > 0.500 V , 0.400 V > 0.450 V > 0.600 V A g / A g a . Similar to the result of 60°C, the oxidation rate is out of order where the oxidation of other potentials predominates over the highest leaching potential (0.600 V A g / A g c i ) . The low iron extraction at 0.600 V A g / A g C i and 78°C is a sign of surface passivation of chalcopyrite, which indicates that the leaching of chalcopyrite is more favorable at lower solution redox potential. Table 4-12 Major iron extraction peaks and iron extractions as a function leaching time under potentials of 0.400 to 0.600 VAg/Agci at temperatures from 60 to 78°C for data set #2 Temp. (°C) E ( V A g / A g C l ) Major Iron Extraction Peak#l ( % ) Leaching Time (hnmin) 60 0.400 23.81 8:30 0.450 54.65* 7:10 0.500 49.93, 10:02 0.550 25.71 3:00 0.600 16.32 6:00 70 0.400 22.35 6:30 0.450 30.13 4:15 0.500 47.27 6:35 0.550 49.74 3:00 0.600 78 0.400 53.80* 6:00 0.450 38.59 2:00 0.500 22.31 1:00 0.550 55.11 2:00 0.600 Major Iron Extraction Peak#2 (%) Leaching Time (hnmin) 32.60 15:30 37.99 26:00 53.35* 19:00 28.98 10:00 53.98* 17:30 49.59* 19:00 48.47* 15:30 60.38* 8:30 40.96* 6:30 53.09 8:00 60.50* 10:00 58.88* 6:00 56.74* 3:30 43.71 2.00 *The maximum solution iron concentration during t Final Iron Extraction (%) Leaching Time '. (hnmin) 49.26* 59:30 44.09 60:00 43.55 66:46 51.26* 52:00 39.51 60:30 46.49 22:00 . 56.63* 21:00 45.40 18:30 52.32 20:00 36.27 21:00 51.04 23:00 54.66 20:00 52.51 17:00 54.43 16:30 54.32* 22:30 be entire leaching process CHAPTER 4 RESULTS A N D DISCUSSION / 116 The influence of temperature on iron dissolution is less obvious in Table 4-16 as a result of ferric iron precipitation. It is clear that the calculations for iron extractions consistently underestimate the amount of iron evolved during the leach tests. The 0.600 V A g / A g c i test exhibits substantially lower extraction than the rest of the potentials, clearly indicative of ferric iron precipitation. However, a general trend that increasing temperature yields increasing iron extractions can be inferred from this table. A temperature-pH diagram constructed for the stability fields of various iron precipitates is shown in Figure 2-3. Although a low level of total soluble iron (2.95-3.13 g/LF<?2+) was used as a precursor for the generation of ferric oxidants and the pH was kept below 1.5 in accordance with Figure 2-3 by the continuous addition of sulfuric acid to prevent precipitation, the oxidation of chalcopyrite leads to the formation and accumulation of ferric (Fe~+). Ultimately, ferric ions in acid sulfate systems hydrolyze to form various iron precipitates on sulfide surfaces. Figures 2-4 and 2-5 offer thermodynamic support for the assertion that the jarosites are formed during leaching. Since the desired solution redox potential is controlled by the addition of potassium permanganate, potassium jarosite formation is the major precipitate. However, these diagrams only present the crystalline jarosite, and the amorphous form that can form during leach tests is not included.1 9 1 The inter-conversion of amorphous jarosite and ferric ions is much more rapid and could be happening during the transition from peak#l to peak#2. The process of inter-conversion can occur at mineral surfaces or in the bulk solution. At the mineral surface, these precipitates may impart a barrier for the diffusion of ferric oxidants to the unreacted mineral surface and subordinate leaching kinetics to its dissolution. The reverse of equation (2-9) can be the mechanism by which the decomposition of jarosites might occur: AFe,(S04)2(OH)b + 2.5H2S04+H + -» +A+ + \.5Fe2(S04), + 6H20 (4-1) Where:A \sK+,Hp+. Statistical interpretation of the final iron extraction as a function of temperature at different redox potentials is shown in Table 4-13. A 13.77% increase in iron extraction as temperature increases from 60 to 78°C implies the reduction of iron precipitation and an increase in the solubility of iron with temperature. The effect of increasing temperature relative to 60 and 70°C on overall iron extraction is proportional to the amount of ferric ions in solution. Upon an applied potential between the semiconductive chalcopyrite and the solution, the concentrations of the species adjust themselves so as to conform to the Nernst equation. The ferric and ferrous ions CHAPTER 4 RESULTS A N D DISCUSSION / 117 continuously undergo consumption and regeneration in a cyclic manner and their concentrations are automatically adjusted at different levels so as to maintain the desired redox potential set point by the addition of potassium permanganate to the leaching solution. Using the Criss and Cobble theory, the E° values of 60, 70 and 78°C are calculated and shown in Table A-5. On the basis of Nernst equation, a calculation of the ferric to ferrous mole ratio involved at each solution redox potential and temperature employed in this thesis work is shown in Table 4-15. At a given potential, the ferric to ferrous mole ratio decreases with increasing temperature. Hence, i f the total iron concentration is constant at a given potential, ferric ions decrease with increasing temperature. In other words, the low level of ferric ions at 78°C than at 60°C for a given potential results in a low level of jarosite formation and an increase in overall iron extraction. At 0.450 V A g / A g c i , the iron extraction increases by 7.45% at 70°C and 8.96% at 78°C relative to that of 60°C. At other potentials, negative extraction values suggest iron precipitation, which consumed iron during the tests. Statistical interpretation of the final iron extraction as a function of increasing redox potentials relative to 0.400 V A g / A g c i at different temperatures is shown in Table 4-14. It can be seen that iron dissolution from chalcopyrite in acidic ferric sulfate medium is dependent on solution redox potential. Leaching at 0.600 V A g / A g c i is clearly undesirable at 60 and 70°C, which suggests that the high redox potential of solution contributes to passivation. Once again, an increase in iron extraction relative to 0.400 V A g / A g C i at 78°C favors the effect of higher solution redox potential on the leaching of iron from chalcopyrite. Table 4-13 Statistical significance of temperature dependency on total iron extractions relative to that of 60 and 70°C at various potentials E ( V A g / A g C l ) Fe^LA™^ 70° C) Fe^TJLAiQ^irc) 0.400 -2.36% -0.65% 1.71% 0.450 7.45% 8.96% 1.51% 0.500 -1.13% 6.89% 8.03% 0.550 -1.75% 5.70% 7.45% 0.600 -0.44% 13.77% 14.21% CHAPTER 4 RESULTS A N D DISCUSSION / 11 Table 4-14 Statistical significance of redox potential dependency on total iron extractions relative to that of 0.400 VAg/AgCi at various temperatures (60-78°C) E2 (VAg/AgCl) EeZ:clZaA0A00VAg/AgCl^E2) 60°C 70°C 78°C 0.450 0.500 0.550 0.600 -1.27% 8.54% 8.34% -0.07% 1.15% 7.47% 5.38% 5.99% 11.72% -8.30% -6.38% 6.12% Table 4-15 Calculated ferric to ferrous mole ratio at 60, 70 and 78°C using the E" calculated from the Criss and Cobble theory 0.400 VAg/Agci 0.450 VAg/Agci 0.500 VAg/Agci 0.550 VAg/Agci 0.600 VAg/AgC! 60°C 0.000077 0.000439 0.002510 0.014335 0.081869 70°C 0.000041 0.000222 0.001207 0.006553 0.035573 78°C 0.000026 0.000134 0.000700 0.003658 0.019106 0.400 V 0.450 V 0.500 V 0.550 V 0.600 • 60 C 47% 46% 47% 52% 39% E170<C 45% 53% 46% 51% 38% 0 7 8 ^ 46% 55% 54% 58% 52% S o l u t i o n R e d o x P o t e n t i a l ( V A g / A g C | ) Table 4-16 Iron extractions as a function of potential (0.400-0.600 VAg/Agci) and temperature (60-78°C) CHAPTER 4 RESULTS AND DISCUSSION / 11 CHAPTER 4 RESULTS AND DISCUSSION / 120 CHAPTER 4 RESULTS AND DISCUSSION / 1 CHAPTER 4 RESULTS A N D DISCUSSION / 1 100 r 95 -90 -85 -80 -75 -0 2 4 6 8 10 12 14 16 18 2022 24 26 2830 32 34 36 384042 444648 505254 56 586062 646668 70 Leaching Time (hr) Figure 4-20 Extraction-time plot for iron at 0.400 V A g / A g c i and different temperatures (60-78°C) in ferric sulfate medium 100 r 95 -90 -85 -0 2 4 6 8 10 12 14 16 182022 24 26 28 30 32 34 36 384042 4446485052 54 56586062646668 70 Leaching Time (hr) Figure 4-21 Extraction-time plot for iron at 0.450 V A g / A g C i and different temperatures (60-78°C) in ferric sulfate medium CHAPTER 4 RESULTS A N D DISCUSSION / 1 •0.500 V,78°C •0.500 V,70C -0.500 V,60>C Figure 4-22 2 4 6 8 10 12 14 16 18 2022 242628 303234 3638404244464850525456586062646668 70 Leaching Time (hr) Extraction-time plot for iron at 0.500 V A g / A g c i and different temperatures (60-78°C) in ferric sulfate medium •0.550 V,78°C -0.550 V,70°C •0.550 V,60°C Figure 4-23 2 4 6 8 10 12 14 16 18 2022 24 2628 30 32 34 36384042 4446485052 54 5658 6062 64 6668 70 Leaching Time (hr) Extraction-time plot for iron at 0.550 VAg/Agci and different temperatures (60-78°C) in ferric sulfate medium CHAPTER 4 RESULTS A N D DISCUSSION / 124 -0.600 V,78°C -0.600 V,70>C -0.600 V,60°C 0 2 4 6 8 10 12 14 16 18 2022 24 26 28 3032 34 36 384042 444648 50 52 54 56 58 6062 646668 70 Leaching Time (hr) Figure 4-24 Extraction-time plot for iron at 0.600 V A g/A gci and different temperatures (60-78°C) in ferric sulfate medium CHAPTER 5 CONCLUSIONS/ 125 C H A P T E R 5 C O N C L U S I O N S Chalcopyrite ( CuFeS2) is the world's most important copper-bearing mineral. The leaching of chalcopyrite is a complex phenomenon, which makes it difficult to design a cost-effective and efficient sulfate based hydrometallurgical process for chalcopyrite. It is widely acknowledged that poor copper extraction can be attributed to the formation of a passivating layer on the chalcopyrite surface in acidic ferric sulfate medium. Considerable differences in opinion exist as to the structure and composition of the passivating layer. Currently, there is no commercial hydrometallurgical process operating for the recovery of copper from chalcopyrite. However, in view of its potential as an environmentally more benign process than pyrometallurgy, hydrometallurgy promises to replace many i f not all pyrometallurgical ore extraction methods in the not too distant future. The present study investigated the dependency of solution redox potential (0.400-0.600 V A g / A g c i ) on chalcopyrite leaching by batch leaching experiments with sulfuric acid solutions at 60, 70 and 78°C. On the basis of the experimental results, the following conclusions are drawn: 1. Both the elemental sulfur and sulfuric acid forming reactions take place and thus suggest that the oxidation mechanism is quite complex. 2. Chalcopyrite leaching in sulfuric acid is shown to be dependent on solution redox potential determined by the concentration ratio of ferric to ferrous ions. 3. The most significant finding in this thesis work is the fact that the leaching rate increases with increasing potentials but decreases above a critical potential. This critical potential is in the vicinity of 0.500 V A g / A g c i and is within the mid-potential range of 0.45 to 0.55 V A g / A g C i at which the leaching rate is maximal from 60 to 78°C. 4. SEM analysis suggests that neither jarosite nor elemental sulfur is responsible for surface passivation or for preventing the dissolution of chalcopyrite. CHAPTER 6 RECOMMENDATIONS FOR FUTURE W O R K / 126 CHAPTER 6 RECOMMENDATIONS FOR FUTURE WORK This work investigated the effect of potential on the leaching of chalcopyrite in acidic ferric sulfate medium. The following suggestions are studies that should be pursued in the future: 1. The use of chalcopyrite concentrates of varying purity is needed to establish whether the critical redox potential, where optimal chalcopyrite leaching occurs, varies with different types of concentrate. 2. Since pronounced differences can exist between results obtained in bench-top experiments and those obtained on a large scale, a major task is to examine the controlled leaching behavior of chalcopyrite at a larger scale (or pilot plant study). Moreover, the most significant problem is the general transferability of short-term laboratory data to predict long-term event as in dump and heap operations. 3. The effect of the potential leaching should be explored with (extreme) thermophiles. 4. A detailed kinetic study should be conducted into the potential leaching of chalcopyrite. 5. The use of electrodes rather than chemical oxidants such as potassium permanganate (KMn04) in potential leaching systems can be beneficial since the post-reaction separation of products is much easier. 6. The influence of different sources of oxidizing potential obtained from appropriate chemical reagent or a potential source (potentiostat) on the leaching kinetics of chalcopyrite can be very interesting. 7. The replacement of potassium permanganate solutions with the widely available manganese oxide such as pyrolusite (Mn02) is a cheap alternative for potential leaching. R E F E R E N C E S / 127 R E F E R E N C E S 1. Ferron, C.J.; Fleming, C A . "Hydrometallurgical Processing of Copper Concentrates," In Electrorefining and Hydrometallurgy of Copper of the Proc. of Copper 95 Intl. Conf. (Vol. III). W.C. Copper, D.B. Dreisinger, J.E. Dutrizac, H. Hein and G. Ugarte (Eds.), Chile, Santiago, 1995, pp. 535-548. 2. Arbiter, N . ; Fletcher, A.W. "Chapter 34. Copper Hydrometallurgy-Evaluation and Milestones," In Hydrometallurgy: Fundamentals, Technology and Innovations of the Proc. of the Milton E. Wadsworth (IV) Intl. Symp. on Hydrometallurgy. J.B. Hiskey and G.W. Warren (Eds.), SME: Littleton, Colorado, 1993, pp. 549-565. 3. Dutrizac, J.E. "The Leaching of Sulphide Minerals in Chloride Media," Hydrometall, 29, 1-45(1992). 4. Bell, Shawn L.; Welch, Glenn D.; Bennett, Paul G. "Development of Ammoniacal Lixiviants for the In-situ Leaching of Chalcopyrite," Hydrometall, 39, 11-23(1995). 5. Brennicke, H .M. ; Bergmann, O.; Ellefson, R.R.; Davies, D.S.; Lueders, R.E.; Spitz, R.A. "Nitric-sulfuric Leach Process for the Recovery of Copper from Concentrate," Miner. Eng., 33.1259-1266(1981). . . . ^ - • . 6. Peters, E. "Hydrometallurgical Process Innovation," Hydrometall, 29, 431-459(1992). 7. King, J.A.; Dreisinger, D.B. "Autoclaving of Copper Concentrates," In Electrorefining and Hydrometallurgy of Copper of the Proc. of Copper 95 Intl. Conf. (Vol. III). W.C. Copper, D.B. Dreisinger, J.E. Dutrizac, H. Hein and G. Ugarte (Eds.), Chile, Santiago, 1995, pp. 511-533. 8. Hackl, R.P.; Dreisinger, D.B.; Peters, E.; King, J.A. "Passivation of Chalcopyrite During Oxidative Leaching in Sulfate Media," Hydrometall, 39, 25-48(1995). 9. Rossi, G. Biohydrometallurgy, McGraw-Hill Book Company GmbH: Hamburg, Germany, 1990. rd 10. Biswas, A . K . ; Davenport, W.G. Extractive Metallurgy of Copper (3 ed.). Pergamon, 1994, p. 471. 11. Vaughan, D.J. and Craig, J.R. Mineral Chemistry of Metal Sulphides. Cambridge University Press, New York, (1978). R E F E R E N C E S / 128 12. Pridmore, D.F.; Shuey, R.T. "The Electrical Resistivity of Galena, Pyrite and Chalcopyrite,"/lw. Mineral., 61, 248(1976). 13. Hiskey, J.B. "Chalcopyrite Semiconductor Electrochemistry and Dissolution," In Extractive Metallurgy of Copper, Nickel and Cobalt (the Paul E. Queneau Intl. Symp.), Volume I: Fundamental Aspects. R.G. Reddy and R.N. Weizenbach (Eds.), TMS-AIME, Warrendale, (PA, USA), 1993, pp. 949-969. 14. Hiskey, J.B.; Wadsworth, M.E. "Electrochemical Processes in the Leaching of Metal Sulfides and Oxides," In Process and Fundamental Considerations of Selected Hydrometallurgical Systems. Martin C.K. (Ed.), SME-AIME, New York, 1981, pp. 303-325. 15. Vukotic, S. "Solubility of Galena, Sphalerite, and Chalcopyrite in Water and in the Presence of H 2 S," Bull. Bur. Reck Geol. Minieres, 3, 11-27(1961). C A 57, 2919f. 16. Putnis, A. Introduction to Mineral Sciences. Cambridge University Press, 1992, p. 130. 17. Hall, S.R.; Stewart, J .M. "The Crystal Structure Refinement of Chalcopyrite, CuFeS 2," Acta Crystallographica, B29, 579-585(1973). 18. Greenwood, N.N. ; Whitfield, H-J- "Mossbauer Effect Studies on Cubanite (CuFe 2S 3) and Related Iron Sulfides," J. Chem. Soc. A. London, 7, 1697-1699(1968). 19. Nakai, 1.; Sugitani, Y . ; Nagashima, K. "X-ray Photoelectron Spectroscopic Study of Copper Minerals," J. Inorg. Nucl. Chem., 40, 789-791 (1978). 20. Bierhaus, E .M. ; Perez, J.; Torma, A.E. ; Rossi, G. " A Comparison of Bacterial Leachability of Chalcopyrite Concentrates from Different Origins," In Progress in Biohydrometallurgy, Cagliari, Italy. G. Rossi and A.E. Torma (Eds.), 1983, pp. 127-150. ' 21. Donnay, G.; Corliss, L . M . ; Donnay, J.D.H.; Elliot, N . ; Hastings, J .M. "Symmetry of Magnetic Structures: Magnetic Structure of Chalcopyrite," Phys. Rev., 112, 1917-1923(1958). 22. Robie, R.A.; Bethke, P.M.; Beardsley, K . M . Selected X-ray Crystallographic Data: Molar Volumes, and Densities of Minerals and Related Substances. U.S. Geological Survey Bulletin 1248, 1967, pp. 8-9. 23. Scott, J.D. "The Mineralogy of Copper leaching: Concentrates and Heaps," paper presented in Copper Hydrometallurgy short course, CIM, Ottawa (Ontario, Canada), 1991. R E F E R E N C E S / 129 24. ICDD-PDF-2 Database (Intl. Center for Diffraction Data-Power Diffraction File), Version 2.12, JCPDS-ICDD Copyright, U.S.A, 1991, Volume[CD]: 291.60. 25. George, D.B. "Copper," In Kirk-Othmer: Encyclopaedia of Chemical Technology (4 t h ed.), J.I. Kroschwitz (Ed.), John Wiley and Sons Inc.: New York, USA, 1992, Vol . 7, pp. 381 -428. 26. Peters, E. "Ch.10. The Physical Chemistry of Hydrometallurgical Processes," In Intl. Symp. on Hydrometallurgy. D.J.I. Evans and R.S. Shoemaker (Eds.), A1ME: New York, USA, 1973, pp. 205-228. 27. Peters, E. "Direct Leaching of Sulfides: Chemistry and Applications," Metall. Trans. B, IB, 505-517(1976). 28. Peters, E. "Lecture 3. Metastable E H -pH Diagrams," In Hydrometallurgy II class notes ( M M A T 458), Department of Metals and Materials Engineering, University of British Columbia, (Vancouver, B.C., Canada), 1996. 29 Hackl, R.P. The Leaching and Passivation of Chalcopyrite in Acid Sulfate Media. Ph.D. Thesis, Department of Metals and Materials Engineering, University of British Columbia, (Vancouver, B.C., Canada), 1995. 30. Hackl, R.P." "Elements of Microbiological Systems," In An Introduction to Biohydrometallurgy: A Short Course. (R.P. Hackl, E. Peters and D.B. Dreisinger, contributors). Department of Metals and Materials Engineering, University of British Columbia, (Vancouver, B.C., Canada), 1993, Chapters 1 and 2. 31. Van Breeman, N . "Genesis, Morphology, and Classification of Acid Sulfate Soils in Coastal Plains," In Acid Sulfate Weathering. J.A. Kittrick, D.S. Fanning and L.R. Hossner (Eds.), Soil Science Society of America Special Publication No. 10, Madison, Wisconson, 1982. 32. Lacey, D.T.; Lawson, F. "Kinetics of the Liquid-phase Oxidation of Acid Ferrous Sulphate by the Bacterium Thiobacillus ferrooxidans" Biotechnology and Bioengineering, 12, 29-50(1970). 33. Hackl, R.P. "Ch. 1. An Introduction to Bioleaching," In Biohydrometallurgy class notes ( M M A T 592F), Department of Metals and Materials Engineering, University of British Columbia, (Vancouver, B.C., Canada), 1997. 34. Dutrizac, J.E.; Kaiman, S. "Synthesis and Properties of Jarosite-Type Compounds," Can. Mineralogist, 14, 151-158(1976). R E F E R E N C E S / 130 35. Babcan, Jr. "Synthesis of Jarosite, KFe 3 (S0 4 ) 2 (OH) 6 ," Geol. Zb., 22(2), 299-304(1971). 36. Bigham, J.M.; Schwertmann, U . ; Carlson, L. "Mineralogy of Precipitates Formed by Geochemical Oxidation of Fe(II) in Mine Drainage," In Biomineralizing Process of Iron and Manganese. H.C.W. Skinner and R.W. Fitzpatrick (Eds.), Catena Verlag, Cremlingen, Germany, 1992, pp. 219-232. 37. Posnjak, E.; Merwin, H.E. "The Symbol Fe 2 03 -S03-H 2 0," J. Amer. Chem. Soc., 44, 1965-1994(1922). 38. Smith, J.R.; Luthy, G.R.; Middleton, A.C. "Microbial Ferrous Iron Oxidation in Acidic Solution,"./ Wat. Pollut. Control. Fed., 60(4), 518-530(1988). 39. Lazaroff, N . ; Melanson, L.; Lewis, E.; Santoro, N . ; Pueschel, C. "Scanning Electron Microscopy and Infrared Spectroscopy of Iron Sediments Formed by Thiobacillus ferrooxidans," Geomicrobiol. J., 4, 231-268(1985). 40. Tuovinen, O.H.; Carlson, L. "Jarosite in Cultures of Iron-oxidizing Thiobacilli," Geomicrobiol. J.,1, 205-210(1979). 41. Lazaroff, N . ; Sigal, W.; Wasserman, A. "Iron Oxidation and Precipitation of Ferric Hydroxysulfates by Resting Thiobacillus ferrooxidans Cells." Appl. Environ. Microbiol.,-43,924-938(1982). " ' ' 42. Verbaan, B.; Huberts, R. "An Electrochemical Study of the Bacterial Leaching of Synthetic N i 3 S 2 , " Int. J. Min. Proc, 24, 185-202(1988). 43. Pesic, B; Kim, I. "Electrochemistry of T. ferrooxidans Interactions with Pyrite," In Proc. of Mineral Bioprocessing. Ross W. Smith and Manoranjan Misra (Eds.), TMS-AIME, 1991, pp. 413-431. 44. Toro, L. ; Paponetti, B.; Cantalini, C. "Precipitate Formation in the Oxidation of Ferrous Ions in the Presence of Thiobacillus ferrooxidans,''' Hydrometall., 20, 1 -9(1988). 45. Senanayake, G.; Muir, D . M . "Speciation and Reduction Potentials of Metal Ions in Concentrated Chloride and Sulfate Solutions Relevant to Processing Base Metal Sulfides," Metallall. Trans. B., 19B, February, 37-45(1988). 46. Pivovarova, T.A.; Golovacheva, R.S. "Microorganisms Important for Hydrometallurgy: Cytology, Physiology, and Biochemistry," In Biogeotechnology of Metals, Proc. of Intl. Seminar and Intl. Training Course. G. l . Karavaiko and S.N. Groudev (Eds.), Centre of Intl. Projects GKNT Moscow, 1985, pp. 27-55. R E F E R E N C E S / 131 47. Kelly, D.P.; Jones, C.A.; Green, J.S. "Factors Affecting Metabolism and Ferrous Iron Oxidation in Suspensions and Batch Cultures of Thiobacillus ferrooxidans" In Metallurgical Applications of Bacterial Leaching and Related Microbiolofiical Phenomena. L.E. Murr, A .E . Torma and J.A. Brierley (Eds.), Academic Press: New York, 1978, pp. 19-43. 48. Dreisinger, D.B.; Peters, E. "The Oxidation of Ferrous Sulphate by Molecular Oxygen Under Zinc Pressure-Leach Conditions," Hydrometall, 22, 101-119(1989). 49. Dutrizac, J.E. "Elemental Sulphur Formation During the Ferric Sulphate Leaching of Chalcopyrite," Can. Metallur. Quart., 28(4), 337-344(1989). 50. Munoz, P.B.; Miller, J.D.; Wadsworth, M.E . "Reaction Mechanism for the Acid Ferric Sulfate Leaching of Chalcopyrite," Metall. Trans. B., 10B, 149-158(June,l 979). 51. Wadsworth, M.E . "Advances in the Leaching of Sulfide Minerals," Min. Sci. Eng., 4(4), 36-47(1972). 52. Dutrizac, J.E.; MacDonald, R.J.C. "Ferric Ion as a Leaching Medium," Min. Sci. Eng., 6(2), 59-100(1974). 53. Jones, D.L.; Peters, E. "The Leaching of Chalcopyrite with Ferric Sulfate and Ferric Chloride," In Extractive Metallurgy of Copper, Vol . II. J.C. Yannopoulos and J.C. Agarwal (Eds.), TMS-AIME: New York, 1976, pp. 633-653. 54. Beckstead, L.W.; Munoz, P.B.; Sepulveda, J.L.; Herbst, J.A.; Miller, J.D.; Olson, F.A.; Wadsworth, M.E . "Acid Ferric Sulfate Leaching of Attritor-Ground Chalcopyrite Concentrates," In Extractive Metallurgy of Copper, Vol . II. J.C. Yannopoulos and J.C. Agarwal (Eds.), TMS-AIME: New York, 1976, pp. 611 -632. 55. Osseo-Asare, K. "Ch. 11. Oxidants and Catalysts in the Aqueous Dissolution of Metal Sulfides: Influence of Electronic Structure on Reactivity," In Hydrometallurgy: Fundamentals, Technology and Innovations of the Proc. of the Milton E. Wadsworth (IV) Intl. Symp. on Hydrometallurgy. J.B. Hiskey and G.W. Warren (Eds.), SME: Littleton, Colorado, 1993, pp. 173-187. 56. Sullivan, J. "Chemical and Physical Features of Copper Leaching," Trans. A.I.M.E., 106, 515-546(1933). 57. Dutrizac, J.E.; MacDonald, R.J .C; Ingraham, T.R. "The Kinetics of Dissolution of Synthetic Chalcopyrite in Aqueous Acidic Ferric Sulfate Solutions," Trans. TMS-AIME, 245, May, 955-959(1969). R E F E R E N C E S / 132 58. Biegler, T.; Swift, D.A. "Anodic Electrochemistry of Chalcopyrite," J. Appl. Electrochem., 9, 545-554(1979). 59. Baur, J.P.; Gibbs, H.L.; Wadsworth, M.E . "Initial Stage Sulphuric Acid Leaching Kinetics of Chalcopyrite Using Radiochemical Techniques," U.S. Bureau of Mines Report of Investigations (USBM RI) 7823. 1974. 60. McElroy, R.O.; Bruynesteyn, A . "Continuous Biological Leaching of Chalcopyrite Concentrates: Demonstration and Economic Analysis," In Metallurgical Applications of Bacterial Leaching and Related Microbiological Phenomena. L.E. Murr, A .E . Torma and J.A. Brierley (Eds.), Academic Press: New York, 1978, pp. 441-462. 61. Hirato, T.; Majima, PL; Awakura, Y . "The Leaching of Chalcopyrite with Ferric Sulfate," Metall. Trans. B, 18B, 489-496(1987). 62. Hiroyoshi, N . ; Hirota, M . ; Hirajima, T.; Tsunekawa, M . " A Case of Ferrous Sulfate Addition Enhancing Chalcopyrite Leaching," HydrometalL, 47, 37-45(1997). 63. Hiroyoshi, N . ; Maeda, H.; Mik i , H. ; Hirajima, T.; Tsunekawa, M . "Ferrous Promoted Chalcopyrite Leaching. Ferric Formation and its Effects on the Leaching," Shigen to Sozai (J. Min. Mater. Process. Inst. Jpn.), 114(11), 795-800(1998). 64. Dutrizac,"J.E. "The Dissolution of Chalcopyrite in Ferric Sulfate and Ferric Chloride Media," Metall Trans. B, 12B, June, 371-378(1981). 65. Dutrizac, J.E. "The Kinetics of Dissolution of Chalcopyrite in Ferric Ion Media," Metall. Trans. B, 9B, 431-439(1978). 66. Cooney, C.L.; Kplov, H .M. ; Haeggstroem, M . "Transient Phenomena in Continuous Culture," In Continuous Cultures of Cells. P.H. Calcott (Ed.), CRC Press Inc.: Boca Raton (Florida), 1, 1981, pp. 143-168. 67. Lowe, Donald F. The Kinetics of the Dissolution Reaction of Copper and Copper-Iron-Sulfide Minerals Using Ferric Sulfate Solutions. Ph.D. Thesis, University of Arizona, U.S.A., 1970. 68. Brierley, J.A.; Brierley, C.L. "Chapter 40. Reflections On and Considerations for Biotechnology in the Metals Extraction Industry," In Hydrometallurgy: Fundamentals, Technology and Innovations. Proc. of the Milton E. Wadsworth (IV) Intl. Symp. on Hydrometallurgy. J.B. Hiskey and G.W. Warren (Eds.), SME: Littleton, Colorado, 1993, pp. 647-660. R E F E R E N C E S / 133 69. Jordan, M.A. ; McGinness, S.; Phillips, C.V. "Acidophilic Bacteria - Their Potential Mining and Environmental Application," Miner. Eng., 9(2), 169-181(1996). 70. Brierley, C L . "Bacterial Oxidation; Master Key to Unlock Refractory Gold Ores?" Eng. Min. J., May, 42-44(1995). 71. Brierley, J.A.; Brierley, C L . "Ch. 12 Microbial Mining Using Thermophilic Microorganisms" In Thermophiles: General, Molecular, and Applied Microbiology. T.D. Brock (Ed.), John Wiley & Sons, Inc.: New York, 1986, pp. 279-303. 72. Rawlings, D.E. "Heavy Metal Mining Using Microbes," Annu. Rev. Microbiol., 56, 65-91(2002). 73. Boon, M . ; Brasser, H.J.; Handsford, G.S.; Heijnen, J.J. "Comparison of the Oxidation Kinetics of Different Pyrites in the Presence of Thiobacillus ferrooxidans or Leptospirillum ferrooxidans," Hydrometall., 53, 57-72(1999). 74. De Filippo, D., Rossi, A. , Rossi, G.& Trois, P. "Chalcopyrite Bioleaching: An Investigation on Copper Recovery," In Proc. of the 8th Intl. Biotechnology Symp. G. Durand, L. Bobichon and J. Florent (Eds.), 1988, pp. 1131-1145. 75. Barriga, M.F.; Pereda, M . J.; Palencia, P.I. "Bacterial Leaching of a Bulk Flotation Concentrate of Chalcopyrite - Sphalerite," Biorecovery, 2, 195-218 (1993). 76. Paponetti, B.A. ; Veglio, F.; Toro, L. "Bioleaching of a Chalcopyrite Concentrate By Thiobacillus ferrooxidans," In Mineral Bioprocessing, TMS, 1991. R.W. Smith and M . Misra (Eds.), 1993, pp. 143-152. 77. Boon, M . Theoretical and Experimental Methods in the Modeling of Bio-Oxidation Kinetics of Sulphide Minerals. Ph.D. Thesis, Department of Biochemical Engineering, Kluyver Laboratory for Biotechnology, Delft University of Technology, (Delft, The Netherlands), 1996. 78. Mcintosh, J.M.; Silver, M . ; Groat, L .A . "Ch. 4 Bacteria and the Breakdown of Sulfide Minerals" In Short Course Series: Biological-Mineralogical Interactions. J .M. Mcintosh and L .A. Groat (Eds.), Mineralogical Association of Canada: Ottawa, 25, 1997, pp. 63-92. 79. Klauber, W.; Parker, A. : Van Bronswijk, W; Watling, H. "Sulphur Speciation of Leached Chalcopyrite Surfaces Determined by X-ray Photoelectron Spectroscopy," Int. J. Miner. Process., 62, 65-94(2001). 80. Warren, G.W.; Wadsworth, M.E. ; El-Raghy, S.M. "Passive and Transpassive Anodic Behaviour of Chalcopyrite in Acid Solutions," Metall. Trans. B., 13B, 571-579(1982). REFERENCES / 134 81. Boon, M . ; Heijnen, J.J. "Mechanisms and Rate Limiting Steps in Bioleaching of Sphalerite, Chalcopyrite, Pyrite with Thiobacillus ferrooxidans" In Biohydrometallurgical Technologies - Bioleaching Processes (Vo. 1), A .E . Torma, J.E. Wey and V.I. Lakshmanan (Eds.), TMS: Pennsylvania, 1993, pp. 217-235. 82. Pinches, A ; Al-Jaid, F.O.; Williams, D.J.A. "Leaching of Chalcopyrite Concentrates with Thiobacillus ferrooxidans in Batch culture," Hydrometall, 2, 87-103(1976). 83. Braithwaite, J.W.; Wadsworth, M.E . "Oxidation of Chalcopyrite Under Simulated Conditions of Deep Solution Mining," In Extractive Metallurgy of Copper, Vol . II. J.C. Yannopoulos and J.C. Agarwal (Eds.), TMS-AIME: New York, 1976, pp. 752-775. 84. Tuovinen, O.H. "Biological Fundamentals of Mineral Leaching Processes," In Microbial Mineral Recovery. H.L. Ehrlich and C.L. Brierley (Eds.), McGraw-Hill Publishing Company: New York, 1990, pp. 55-77. 85. Lotens, J.P.; Wesker, E. "The Behaviour of Sulphur in the Oxidative Leaching of Sulphidic Minerals,"Hydrometall, 18, 39-54(1987). 86. Miller, J.D.; Simkovich, G. "Enhanced Leaching of Minerals Which Form Product Layers," US Patent 4,343,773. August 10, 1982. 87. Peters, E. "Leaching Of Sulphides," In Advances In Mineral Processing, SME, 1986, pp. 445-462. 88. Buttinelli, D.; Lavecchia, R.; Pochetti, F.; Geveci, A. ; Guresin, N . ; Topkaya, Y . "Leaching by Ferric Sulfate of Raw and Concentrated Copper-Zinc Complex Sulphide Ores," Int. J. Min. Proc, 36, 245-257(1992). 89. Linge, H.G. " A Study of Chalcopyrite Dissolution in Acid Ferric Nitrate by Potentiometric Titration,"Hydrometall., 2(1), 51-64(1976). 90. Havlik, T.; Kammel, R. "Leaching of Chalcopyrite with Acidified Ferric Chloride and Carbon Tetratchloride Addition," Miner. Eng., 8, 1125-1134(1995). 91. Burkin, A.R. "Solid-state Transformations During Leaching," Min. Sci. Eng., 1(4), 4-14(1969). 92. King, J.A. Solid State Changes in the Leaching of Copper Sulphides. Ph.D. Thesis, University of London, U.K., 1966. 93. Ferreira, R.C.H. Leaching of Chalcopyrite. Ph.D. Thesis, University of London, U.K. , 1972. REFERENCES/ 135 94. Ammou-Chokroum, M.; Sen, P.K.; Forques, F. "Electrooxidation of Chalcopyrite in Acid Chloride Medium; Kinetics, Stoichiometry and Reaction Mechanism," In Proc. of 13th Intl. Miner. Process. Cong., Part A. J. Laskowski(Ed.), Elsevier/Polish Scientific Publishers: Amsterdam, 1979, pp. 759-809. 95. Parker, A.J.; Paul, R.L.; Power, G.P. "Electrochemistry of the Oxidative Leaching of Copper from Chalcopyrite," J. Electroanalyt. Chem., 118, 305-316(1981). 96. Duncan, D.W.; Trussel, P.C., Can. Metali, Q. 3, 43-55(1964). 97. Duncan, D.W; Trussel, P.C.; Walden, C C , Appl. Microbiol. 12, 122-126(1964). 98. Sandoval, S.P.; Pool, D.L.; Schultze, L.E., Rep. Invest. USBM, 9311, 1-12(1990). 99. Hiroyoshi, N.; Nakamura, T.; Tsunekawa, M.; Hirajima, T.; Ito, M., Shigen to Sozai (J. Min. Mater. Process. Inst. Jpn.), I l l , 943-948 (1995). 100. Hiroyoshi, N.; Nishida, S.; Tsunekawa, M.; Hirajima, T. In 20th Int. Miner. Process. Congr.. Proc. 4, 1997, pp. 547-556. 101. Bryner, I.C; Walker, R.B.; Palmer, R., Trans. SME-AIME, 238, 52-56(1967). 102. Wan, R.Y; Miller, J.D.; Foley, J.; Pons, S. In Proc. of the Intl. Svmp. on Electrochemistry in Mineral and Metal Processing. P.E. Richardson, S. Srinivasan and R. Woods (Eds.), The Electrochemical Society, Inc.: Pennington, (New Jersey, USA.), 1984, pp. 391-416. 103. Nakazawa, H.; Fujisawa, H.; Sato, H. "Effect of Activated Carbon on the Bioleaching of Chalcopyrite Concentrate," Int. J. Miner. Process, 55, 87-94(1998). 104. Sanchez, E .C; Umetsu, Y; Saito, F., J. Chem. Eng. Jpn., 29, 720-721(1996). 105 Sanchez, E .C; Umetsu, Y.; Saito, F., Shigen to Sozai (J. Min. Mater. Process. Inst. Jpn.), 113, 631-633(1997). 106. Miller, J.D.; McDonough, P.J.; Portillo, H.Q. "Electrochemistry in Silver Catalyzed Ferric Sulfate Leaching of Chalcopyrite,'7« Process and Fundamental Considerations of Selected Hydrometallurgical Systems. Martin C. Kuhn (Ed.), SME-AIME: New York, 1981, pp.327-338. 107. Gomez, C ; Figueroa, M.; Munoz, J.; Ballester, A.; Blazquez, M.L. "A Study of Bioleached Chalcopyrite Surfaces in the Presence of Ag(I) by Voltammetric Methods," Miner. Eng., 10(1), 111-116(1997). REFERENCES/ 136 108. Munoz, J.A.; Gomez, C ; Ballester, A.; Blazquez, M.L.; Gonzalez, F, Figueroa, M. "Electrochemical Behavior of Chalcopyrite in the Presence of Silver and Sulfolobus Bacteria," J. Appl. Electrochem., 28, 49-56(1998). 109. Ahonen, L.; Touvinen, O.H. "Silver Catalysis of the Bacterial Leaching of Chalcopyrite Containing Ore Material in Column Reactors," Miner. Eng., 3(5), 437-445(1990). 110. Ahonen, L.; Touvinen, O.H. "Catalytic Effects of Silver in the Microbiological Leaching of Finely Ground Chalcopyrite-containing ore materials in Shake Flasks," Hydrometall., 24,219(1990). 111. Price, D.W.; Warren, G.W. "The Influence of Silver Ion on the Electrochemical Response of Chalcopyrite and Other Mineral Sulfide Electrodes in Sulfuric Acid," Hydrometall, 15, 303-324(1986). 112. Ballester, A.; Gonzalez, F.; Blazquez, M.L.; Mier, J.L. "The Influence of Various Ions in the Bioleaching of Metal Ions," Hydrometall, 23, 221-235(1990). 113. Kuzeci, E.; Kammel, R., Erzmetall, 41, 327-338(1981). 114. Mateos, F.B.; Perez, I.P.; Mora, F.C. "The Passivation of Chalcopyrite Subjected to Ferric Sulfate Leaching and its Reactivation with Metal Sulfides," Hydrometall, 19, 159-. 167(1987),. " -115. Peters, E; Doyle, F.M. "Leaching and Decomposition of Sulfide Minerals," In Challenges in Mineral Processing. K.V. Sastry and M.C. Fuerstenau (Eds.), SME-AIME: Littleton, CO, USA, 1988, pp. 509-526. 116. Chen, J.H.; Harvey, W.W. "Cation Self-Diffusion in Chalcopyrite and Pyrite," Metall. Trans. B, 6B, 331-339(1975). 117. McElroy, R.O.; Duncan, D.W. "Copper Extraction By a Rapid Bacteriological Method," US Patent 3,856,913. December, 1974.. 118. Bruynesteyn, A.; Hackl, R.P.; Lawrence, R.W.; Vizsolyi, A.I. "Biological-acid Leach Process," US Patent 4,571,387. February 18, 1986. 119. Lawrence, R.W.; Vizsolyi, A.; Vos, R.J. "The Silver Catalyzed Bioleach Process for Copper Concentrate," In Microbiological Effects on Metallurgical Processes. J.A. Clum and L.A. Haas (Eds.), TMS-AIME: Warrendale, PA, USA, 1985, pp. 65-82. R E F E R E N C E S / 137 120. Gomez, C. Biolixiviacion Con Cultivos Mesofilos de Concentrados Globales de Sulfuros Complejos En Presencia de Iones Catalizadores. Ph.D. Thesis, University of Complutense, Madrid, 1993. 121. Mier, J.L.; Ballester, A. ; Blazquez, M.L . ; Gonzalez, F.; Munoz, J.A. "Influence of Metallic Ions in the Bioleaching of Chalcopyrite by Sulfolobus BC: Experiment Using Pneumatically Stirred Reactors and Massive Samples," Miner. Eng., 8, 949-965(1995). 122. Ballester, A . "Chapter 9" In Separation Process in Hydrometallurgy. G.A. Davies (Ed.), Ellis Horwood Publishers, London, 1987, pp. 99-110. ' 123. Sukla, L.B. ; Chaudhury, G.R.; Das, R.P. "Effect of Silver Ion on Kinetics of Biochemical Leaching of Chalcopyrite Concentrate," Trans. Instn. Min. Metall, C, 9 9 , C43-C46(1990). 124. Koch, D.F.A. Modern Aspects of Electrochemistry. Plenum Press: New York, 1975, pp. 211-237. 125. McMillan, R.S.; MacKinnon, D.J.; Dutrizac, J.E. "Anodic Dissolution of n-type and p-type Chalcopyrite," / Appl. Electrochem., 12(6), 743-757(1982). 126. Gardner, J.R.; Woods, R. "An Electrochemical Investigation of the Natural Floatability of Chalcopyrite," Int. J. Min. Proc, 6, 1-16(1979). : . . . . 127. Hamilton, I.C.; Woods, R. " A Voltammetric Study of the Surface Oxidation of Sulphide Minerals," In Proc. of the Intl. Symp. on Electrochemistry in Mineral and Metal Processing. P.E. Richardson, S. Srinivasan and R. Woods (Eds.), The Electrochemical Society, Inc.: Pennington, (New Jersey, USA.), 1984, pp. 259-285. 128. Biegler, T.; Home, M.D. "The Electrochemistry of Surface Oxidation of Chalcopyrite," In Proc. of the Intl. Symp. on Electrochemistry in Mineral and Metal Processing. P.E. Richardson, S. Srinivasan and R. Woods (Eds.), The Electrochemical Society, Inc.: Pennington, (New Jersey, USA.), 1984, pp. 321-339. 129. Parker, A.J . ; Paul, R.L.; Power, G.P. "Electrochemical Aspects of Leaching Copper from Chalcopyrite in Ferric and Copper Salt Solutions," Aust. J. Chem., 34, 13-34(1981). 130. Holliday, R.I.; Richmond, W.R. "An Electrochemical Study of the Oxidation of Chalcopyrite in Acidic Solution," J. Electroanal. Chem., 288, 83-98(1990). 131. Jones, D.L. The Leaching of Chalcopyrite. Ph.D. Thesis, Department of Metals and Materials Engineering, University of British Columbia, (Vancouver, B.C., Canada), May, 1974. R E F E R E N C E S / 138 132. Warren, G.W.; Wadsworth, M.E.;,E1-Raghy, S.M. "Anodic Behaviour of Chalcopyrite in Sulphuric Acid," In Hydrometallurgy, Research, Development and Plant Practice. K. Osseo-Asare and J.D. Miller (Eds.), TMS-AIME, Warrendale, (PA, USA), 1983, pp. 261-275. 133. Kelsall, G.H.; Page, P.W. "Aspects of Chalcopyrite (CuFeS2) Electrochemistry," In Proc. of the Intl. Symp. on Electrochemistry in Mineral and Metal Processing. P.E. Richardson, S. Srinivasan and R. Woods (Eds.), The Electrochemical Society, Inc.: Pennington, (New Jersey, USA.), 1984, pp. 303-320. 134. Stankovic, Z.D. "The Anodic Dissolution Reaction of Chalcopyrite," Erzmetall., 39, 623-628(1986). 135. Arce, E .M. ; Gonzalez, I. " A Comparative Study of Electrochemical Behavior of Chalcopyrite, Chalcocite and Bornite in Sulfuric Acid Solution," Int. J. Min. Proc, 67, 17-28(2002). 136. Vaughan, D.J.; Becker, U . ; Wright, K. "Sulphide Mineral Surfaces: Theory and Experiment," Int. J. Min. Proc, 51, 1-14(1997). 137. Eadington, P. "Study of Oxidation Layers on Surfaces of Chalcopyrite by use of Auger Electron Spectroscopy," Trans. Inst. Min. Metall, 86, CI 86-C189(1977). -138. Brion, D. "Etude par Spectroscopic de Photoelectrons de la Degradation Superficielle de FeS2, CuFeS 2, ZnS et PbS A L ' A i r et dans UEau," Appl. Surf. Sci., 5, 133-152(1980). 139. Buckley, A . N . ; Woods, R. "An X-ray Photoelectron Spectroscopic Study of the Oxidation ofChalcopyrite,'Muj/ra/.7. Chem., 37, 2403-2413(1984). 140. Ruzakowski, P.H.; Holloway, P.H.; Remond, G. "Complementary Surface Characterization of Chalcopyrite by Electron Microscopy, Electron Spectroscopy, and Optical Reflectance," Scanning Microscopy, 3(1), 71-82(1989). 141. Balaz, P.; Kupka, D.; Bastl, Z.; Achimovicova, M . "Combined Chemical and Bacterial Leaching of Ultrafine Ground Chalcopyrite," Hydrometall., 42, 237-244(1996). 142. Luttrell, G.H.; Yoon, R.H. "Surface Studies of the Collectorless Flotation of Chalcopyrite," Colloids and Surfaces, 12, 239-254(1984) 143. Fullston, D.; Fornasiero, D.; Ralston, J. "Oxidation of Synthetic and Natural Samples of Enargite and Tennantite: 2. X-Ray Photoelectron Spectroscopic Study." Langmuir, 15, 4530-4536(1999). R E F E R E N C E S / 139 144. Buckley, A . N . ; Woods, R. "Investigation of the Surface Oxidation of Sulfide Minerals via ESCA and Electrochemical Techniques," In Interfacial Phenomena in Mineral Processing. B. Yarar and D.J. Spottiswood (Eds.), Engineering Foundation: New York, 1982, pp. 3-17. 145. Steudel, R. "Mechanism for the Formation of Elemental Sulfur from Aqueous Sulfide in Chemical and Microbiological Desulfurization Processes," Ind. Eng. Chem. Res., 35, 1417-1423(1996). 146. De Filippo, D.; Rossi, A. ; Rossi, G.; Trois, P. "Surface Modifications in Copper Sulphide Minerals After Bioleaching," In Proc. of Intl. Biohydrometallurgy Symp. Warwick IBS-87. P.R. Norris and D.P. Kelly (Eds.), Science and Technology Letters, Kew, Surrey, U.K. , 1988, pp. 245-258. 147. Peters, E. "The Electrochemistry of Sulphide Minerals," In Trends in Electrochemistry. J .O 'M. Bockris, D.A.J. Rand and B.J. Welch (Eds.), Plenum Publishing Corp.: New York, U.S.A., 1977, pp. 267-290. 148. Hiskey, J.B. "Chalcopyrite Semiconductor Electrochemistry and Dissolution," In Extractive Metallurgy of Copper, Nickel and Cobalt (the Paul E. Queneau Intl. Symp.), Volume I: Fundamental Aspects. R.G. Reddy and R.N. Weizenbach (Eds.), TMS-AIME, Warrendale, (PA, USA), 1993, pp. 949-969. 149. Koch, G.H. Modern Aspects of the Electrochemistry, Vol . 10. Plenum: New York, 1975, pp.211-237. 150. Shuey, R.T. Semiconducting Ore Minerals. Elsevier: New York, 1975. 151. Wadsworth, M.E . "Electrochemical Reactions In Hydrometallurgy," In Metallurgical Treatises. J.K. Tien and J.F. Elliott (Eds.). AIME, New York, 1981, pp. 1-22. 152. Mehta, A.P.; Murr, L.E. "Fundamental Studies of the Contribution of Galvanic Interaction to Acid-Bacterial Leaching of Mixed Metal Sulfides," Hydrometall, 24, 919-940(1983). 153. L i , J.; Zhong, T.K.; Wadsworth, M.E. "Application of Mixed Potential Theory in Hydrometallurgy," In Hydrometallurgy: Theory and Practice. Proc. of the Ernest Peters Intl. Symp. (Part A). W.C. Copper and D.B. Dreisinger (Eds.). Hydrometall., 29, 47-60(1992). 154. Crundwell, F.K. "The Influence of the Electronic Structure of Solids on the Anodic Dissolution and Leaching of Semiconducting Sulphide Minerals," Hydrometall., 21, 155-190(1988). R E F E R E N C E S / 140 155. Osseo-Asare, K. "Semiconductor Electrochemistry and Hydrometallurgical Dissolution Processes,"Hydrometall, 29, 61-90(1992). 156. Hiskey, J .B.; Wadsworth, M.E . "Galvanic Conversion of Chalcopyrite," Metall Trans. B, 6B, 183-190(1975). 157. Natarajan, K . A . "Electrochemical Aspects of Bioleaching Multisulfide Minerals," Min. Metallurg. Proc., May, 61-65(1988). 158. Torma, A.E. ; Apel, W.A. "Bioleaching of Chalcopyrite Samples," In EPD Cong. 1992. J.P. Hager (Ed.), TMS, 1991, pp. 3-15. 159. Hiroyoshi, N . ; Arai, M . ; Miki , H. ; Hirajima, T.; Tsunekawa, M . " A New Reaction Model for the Catalytic Effect of Silver Ions on Chalcopyrite Leaching in Sulfuric Acid Solutions," Hydrometall, 63, 257-267(2002). 160. Criss, C M . ; Cobble, J.W. "The Thermodynamic Properties of High Temperature Aqueous Solutions. IV: Entropies of the Ions up to 200° and the Correspondence Principle," J. Amer. Chem. Soc, 86(24), 5385-5390(1964). 161. Criss, C M . ; Cobble, J.W. "The Thermodynamic Properties of High Temperature Aqueous Solutions. V: The Calculation of Ionic Heat Capacities up to Entropies of the Ions up to '••200°, Entropies and Heat Capacities above 200°,":J.Amer. Chem. Soc, 86(24), 5390-5393(1964). 162. Wadsworth, M.E., "Heterogeneous Rate Processes in the Leaching of Base Metal Sulfides," In Hydrometallurgical Process Fundamentals. R.G. Bautista (Ed.), Plenum Press, New York, USA, 1984, pp. 41-76. 163. Kametani, H.; , Aoki, A. "Effect of Suspension Potential on the Oxidation Rate of Copper Concentrate in a Sulfuric Acid Solution," Metall. Trans. B, 16B(4), 695-705(1985). 164. Verbaan, B.; Crundwell, F.K. "An Electrochemical Model for the Leaching of a Sphalerite Concentrate," Hydrometall, 16, 345-359(1986). 165. Ralph, D.E.; May, N . ; Handsford, G.S. "An Electrochemical Step in Bioleaching Biomine," In Conf. Proc. of IBS-Biomine. A M F , Sydney, Australia, pp. QP3.1-QP3.10. 166. May, N . ; Ralph, D.E.; Handsford, G.S. "Dynamic Redox Potential Measurements for Determining the Ferric Leach Kinetics of Pyrite," Miner. Eng., 10(11), 1279-1290(1997). 167. Ruitenberg, R.; Handsford, G.S.; Reuter, M.A. ; Breed, A.W. "The Ferric Leaching Kinetics of Arsenopyrite," Hydrometall, 52, 37-53(1999). R E F E R E N C E S / 141 168. Van der Merwe, C ; Myburgh, P.J.; Pinches, A . " A Process for the Leaching of Chalcopyrite," South Africa Patent 971307, February 17, 1997. 169. Nagpal, S.; Dahlstrom, D.; Oolman, T. " A Mathematical Model for the Bacterial Oxidation of a Sulfide Ore Concentrate," Biotechnology and Bioengineering, 43(5), 357-364(1994). 170. Hansford, G.S.; Vargas, T. "Chemical and Electrochemical Basis of Bioleaching Processes," Hydrometall., 59, 135-145(2001). 171. Natarajan, K . A . "Bioleaching of Sulphides Under Applied Potentials," Hydrometall., 29, 161-172(1992). 172. Ahonen, L.; Tuovinen O.H. "Redox Potential-Controlled Bacterial Leaching of Chalcopyrite Ores," In Biohydrometallurgical Technologies - Bioleaching Processes (Vol. 1), A .E . Torma, J.E. Wey and V.I. Lakshmanan (Eds.), TMS: Pennsylvania, 1993, pp. 571-578. 173. Romero, R.; Palencia, I.; Carranza, F. "Silver Catalyzed IBES Process: Application to a Spanish Copper-Zinc Sulphide Concentrate Part 3. Selection of the Operational Parameters for a Continuous Pilot Plant," Hydrometall, 49, 75-86(1998). 174. Vancas, M.F.; Cornejo, R. "MINBAC™ Technology Bio-oxidation Process for Copper," In First Annual Copper Hydromet Roundtable (Vancouver, British Columbia,-Canada), 1995, pp.11-14. 175. Hackl, R.P. "Chapters 4: Prospects for Heap Leaching Chalcopyrite," In Biohydrometallurgy Short Course: Heap Leaching of Copper Ores - Practice, Problems and Possibilities. (R.P. Hackl and David G. Dixon, contributors). Department of Metals and Materials Engineering, University of British Columbia, (Vancouver, B.C., Canada), 1998. 176. Lowson, R.T. "Aqueous Oxidation of Pyrite By Molecular Oxygen," Chem. Rev., 82, 461-497(1982). 177. Manual BIOWATCH (v. 2.30), 1992, Applikon Dependable Instruments, Schiedam, The Netherlands. 178. Skoog, D.A.; Leary, J.J. Principles of Instrumental Analysis (4th ed.). Saunders College Publishing: Fort Worth, 1992, pp. 490-491. 179. Light, Truman S. "Standard Solution for Redox Potential Measurements," Analytical Chemistry, 44(6), May, 1038-1039(1972). REFERENCES / 142 180. Standard Practice for Oxidation-Reduction Potential of Water, Philadelphia, Pennsylvania: Annual Book of American Society for Testing Materials (ASTM) Standards, 1980, Part 31, D1498-76, pp. 215-220. 181. G.R. Kirkbright; M . Sargent. Atomic Absorption and Fluorescence Spectroscopy, Academic Press, 1974. 182. Jeffery, G.H.; Bassett, J.; Mendham, J.; Denney, R.C. Vogel's Textbook of Quantitative Chemical Analysis (5 t h ed.) Longman Scientific & Technical: New York, 1989, pp. 368-372. 183. "Table 5.2: Solubilities of Inorganic Compounds and Metal Salts of Organic Acids in Water at Various Temperatures," In Lange's Handbook of Chemistry (14 t h ed.). J.A. Dean (Ed.). McGraw-Hill, Inc.: New York, 1992, p. 5.18. 184. Turney, T.A. Oxidation Mechanisms. Butterworth & Co (Publishers) Ltd.: London. 185. Hayes, Peter C. Process Principles in Minerals and Materials Production. Hayes Publishing Corporation: Brisbane, 1993, pp. 181-185. 186. Petrucci, R. H. "Ch.24-6, Manganese," from General Chemistry: Principles and Modern Applications (5 t h ed.). Macmillan Publishing Company: New York, 1989, p. 883. 187. Unknown Source. 188. Keenan, Joseph H. ; Keyes, Frederick G.; Hil l , Philip G.; Moore, Joan G. Steam Tables. A Wiley-Interscience Publication: New York, 1978. 189. Streiner, D.L. "Maintaining Standards: Difference between the Standard Deviation and Standard Error, and When to Use Each," The Canadian Journal of Psychiatry, 41(8), 498-502(1996). 190. Dutrizac, J.E. "The Behavior of Impurities During Jarosite Precipitation," In Hydrometallurgical Process Fundamentals. R.G. Bautista (Ed.), Plenum Press, New York, USA, 1984, pp. 125-169. 191. Liddell, K .C. ; Bautista R.G.. " A Partial Equilibrium Model to Characterize the Precipitation of Ferric Ion during the Leaching of Chalcopyrite with Ferric Sulfate," Metall. Trans. B, I4B, 5-14(1983). 192. Walsh, Frank C. A First Course In Electrochemical Engineering. The Electrochemical Consultancy (Romsey) Ltd., Romsey, England, 1993, p.30. R E F E R E N C E S / 143 193. Dixon, D.G. "The Eh-pH Diagram at Elevated Temperatures," In Hydrometallurgy II class notes ( M M A T 458), Department of Metals and Materials Engineering, University of British Columbia, (Vancouver, B.C., Canada), 1998. 194. Bard, J.A.; Parsons, R.; Jordan, J. Standard Potentials in Aqueous Solutions. Mercel Dekker, Inc.: New York, 1985. A P P E N D I X A S A M P L E C A L C U L A T I O N S APPENDIX A S A M P L E CALCULATIONS / 145 A-1 Calculation of Solid Pulp Density (SPD) The standard leaching medium in all tests has similar composition containing approximately 10 g chalcopyrite. This sample calculation is based on the results for the experiment (Test No. 38 of data set #2) performed at the following parameters: Leaching Conditions: Solution Redox Potential Setpoint = 0.500 V A g / A g C i Temperature Setpoint = 60°C Data: mi,emi = m a s s of head sample added (g) = 10.02 p 2 ^ r = density of water at 20°C (g/mL) =1.00 PZH2SO< (96%purity) = density of 96% purity 6M sulfuric acid at 20°C =0.59 V20"c = volume of water added at 20°C (mL) = 2000 VbMH2so4 (96%purity) = volume of 96% purity 6M sulfuric acid added at 20°C = 30 The solid pulp density (SPD ) of the initial standard CuFeS2 - Fe(S04) - H2S04 leaching medium for all tests is defined by the mass of head sample added to the medium: SPD = ( ) • 100% (A-l) mhend + m solution 0.49% 5j P'solution^'solution ^) [pf2o • V%oC ] + IPlZso, (96%purity) • (96%purity)] 2017.7 g Where: SPD = solid pulp density (%) msohmon = m a s s °f solution added (g) Psoiutton = density of solution (g/mL) ^solution = volume of solution added (mL) APPENDIX A S A M P L E C A L C U L A T I O N S / 146 A-2 Calculation of Copper Extraction Based on Calculated Head A summary of calculated head grade and copper extractions can be found at Table B - l in Appendix B. This sample calculation is based on the results for the experiment (Test No. 38 of data set #2) performed at the following parameters: Leaching Conditions: Solution Redox Potential Setpoint = 0.500 V A g / A g c i Temperature Setpoint = 60°C Data: 171 head = m a s s of head sample added (g) = 10.02 m™' = mass of copper solubilized in the leach media (g) . = 2.6404 ^,^usample^sample = s u m 0I"copper mass in samples taken out for assay (g) = 0.1001 m™ = mass of copper in the leach residue (g) =0.1782 Cu'-AveCaicHead = average calculated head assay for copper (%) = 29.37 (from Table B- l ) msro1 + y Cu . V " . + mr" CuCnlcHend = ( c " ^ sa"""e s"'"p,e ^ ) 1 0 0 % (A-3) mhead = 29.13%-29% sol + V p y (Cu'-CalcHead / 1 00) - (#W' w ) = 93.89% - 94% Similarly, Cu^~ead = 93.11% « 93% Where: CuCalcHead = calculated head assay for copper (%) Cuc™Head = copper extraction based on calculated head assay (%) CuAvZaicHead = copper extraction based on average calculated head assay (%) APPENDIX A S A M P L E C A L C U L A T I O N S / 147 A-3 Calculation of Iron Extraction Based on Calculated Head A summary of calculated head grade and iron extractions can be found at Table B - l in Appendix B. This sample calculation is based on the results for the experiment (Test No. 38 of data set #2) performed at the following parameters: Leaching Conditions: Solution Redox Potential Setpoint = 0.500 V A g / A g c i Temperature Setpoint = 60°C Data: = mass of head sample added (g) = mass of iron solubilized in the leach media (g) m m head sol Fe FeSO 1H o — mass of iron from the ferrous sulfate salt added (g) = sum of iron mass in samples taken out for assay (g) m'Fe = mass of iron in the leach residue (g) m V sample sample Fe AvgCalcHead average calculated head assay for iron (%) (from Table B- l ) 10.02 6.9307 6.0282 0.5287 1.8243 32.82 Fe, CalcHead = ( sol , mFe + 5 > sample^sample + m Fe 1 7 1 FeSO, 111,0 ) • 100% m head 32.50% - 33% (A-5) Fe Extraction CalcHead sol , mFe + iFe v sample sample m / 1 0 0 ) - ( m w ) = 43.98% = 44% Similarly, Fe^~end =43.55% = 44% (A-6) Where: Fe CalcHead Fe CalcHead Extraction * eAvgCalcHead calculated head assay for iron (%) iron extraction based on calculated head assay (%) iron extraction based on average calculated head assay (%) APPENDIX A S A M P L E C A L C U L A T I O N S / 148 A-4 Calculation of Total Ferrous Ions Concentration [Fe(II)T ] During Leaching Solution samples were assayed for ferrous ions (Fe 2 + ) using a standard titration with potassium permanganate. The total ferrous concentration [Fe(II)T] in the leaching solution at any given time during leaching is the product of ferrous ions in the titrant and the inverse of the dilution factor. This sample calculation is based on the results for the experiment (Test No. 38 of data set #2) performed at the following parameters: Leaching Conditions: Solution Redox Potential Setpoint = 0.500 V A g / A g C i Temperature Setpoint = 60°C Data: C/cMno, = standardized potassium permanganate solution = 0.002704 concentration (mol/L) VjX"p = volume of leaching solution after the adjustment = 2030 of evaporation rate (mL) Vj's = average titrant volume of potassium = 4.20 permanganate solution (mL) VTesl = test solution volume (mL) • =1.00 ( i \ yDFj [Fe(lI)T] = Fe{ll)rim = \C • Vvs • 51 = 115 mmol (A-7) Where: DF Fe(H)rilrnil [Fe(JJ)T] dilution factor (mmol) total ferrous ion concentration in the titrant(mmol)' total ferrous ion concentration (mmol) APPENDIX A S A M P L E C A L C U L A T I O N S / 149 A-5 Determination of Sulfide Sulfur Conversion [ S20xidaljm - S2S„ - (SO2' )F ] A summary of sulfide sulfur conversion can be found at Table B-2 in Appendix B. This sample calculation is based on the results for the experiment (Test No. 38 of data set #2) performed at the following parameters: Leaching Conditions: Solution Redox Potential Setpoint = 0.500 V A g / A g c i Temperature Setpoint = 60°C Data: = mass of sulfide sulfur in the head sample (g) =3.3647 m'su"fi,r = m a s s °f elemental sulfur in the head sample (g) = 0.0050 m'suin<ie = m a s s of sulfide sulfur in the leach residue (g) = 0.9882 m7utiur = m a s s of elemental sulfur in the leach residue (g) = 2.2337 Total Sulfide Oxidation head res ill — ni S L d a l ! o „ = ( ^ " w ^ " ) • 100% (A-8) 1 7 1 sulfide = 70.63%-71% Where: ^ovidauon = t °* a ' sulfide sulfur oxidation of the head sample (%) Total Elemental Sulfur Formation , „ „ r e s .„„ iii — iii S2S: = ( s"""r heatl CT//I")-100% (A-9) m sulfide = 66.24% = 66% Where: S2~ = total sulfide sulfur conversion to elemental sulfur (%) APPENDIX A S A M P L E C A L C U L A T I O N S / 150 Total Sulfate Formation As a result of the high background sulfate concentrations of the sulfate, the trace amount of sulfate formed during leaching could not be determined directly. Consequently, the amount of sulfate produced was estimated indirectly by the differences between total sulfide oxidation and total elemental sulfur formation. (SOi~)F= s( , 2 -1 Oxidation a- (A-10) [0 - m sulfide res + m sulfur res — m r head sulfur )100%] m t head sulfide = 4 . 3 Q 0 / 0 « 4 o / 0 Where: {SO] )F = total sulfide sulfur conversion to sulfate (%) APPENDIX A S A M P L E C A L C U L A T I O N S / 151 A-6 Calculation of Copper Extraction Based on the Consumption of Potassium Permanganate (KMnOA) Solution The permanganate oxidation of ferrous to ferric ions takes place (almost) stoichiometrically and instantaneously as shown by the following equations: MnO-(aq) + m + (aq) + 5e ^ Mn2+(aq) + 4H20 £ ° = + 1 . 5 1 V S H E ( A - l l ) Fe3+(aq) + e~ <^Fe2+{aq) E" = + 0.770 V S H E (A-12) Overall [Combining equations (A- l 1) and (A-12)]: Mn04(aq) + SH+(aq) + 5Fe2+(aq) -> Mn2+(ag) + 4H20 + 5Fe3+(aq) (A-l3) Since one mole of permanganate ion oxidizes five moles of ferrous ions to ferric ions as given by equation (A-l3), and four moles of ferric ions is related to one mole of copper ions generation during chalcopyrite dissolution as given by equation (2-15), the permanganate ion is therefore proportional to the generation of copper by a factor of 1.25 (or the fraction of five over four). A summary of copper extraction based on the consumption of potassium permanganate solution can be found at Table B-3 in Appendix B. This sample calculation is based on the results for the experiment (Test No. 38 of data set #2) performed at the following parameters: Leaching Conditions: Solution Redox Potential Setpoint = 0.500 V A g / A g C i Temperature Setpoint = 60°C Data: mass of head sample added (g) = 10.02 average calculated head assay for copper (%) = 29.37 copper extraction based on average calculated = 93.11 head assay (%) (from Appendix A-2) standardized potassium permanganate solution - 0.2163 concentration (mol/L) volume of potassium permanganate solution = 0.160 consumed (L) molecular weight of copper (g/mol) = 63.546 CuAvgCalcHead ' *~i Extraction ^ U AvgCalcHead ' r v y KMnO, CUMoim APPENDIX A SAMPLE CALCULATIONS/ 152 Cu Extraction KMnO, _ [CuMitO, ' ^KMnO, ' CuMo]m ) ^  ^ j 25) • I 00% m head • (CuAv AvgCalcHead /100)-(A-14) 93 . 41% -93% \Cu Extraction AvgCalcHead - KMnOA Cu Extraction AvgCalcHead Cu Extraction KMn04 = 0 . 2 9 % - 0 . 3 % (A-15) Where: Cu Extraction KMnO, \Cu Extraction AvgCalcHead-KMnO, = copper extraction based on potassium permanganate solution consumption (%) = absolute difference between copper extraction based on average calculated head assay and potassium permanganate solution consumption (%) APPENDIX A S A M P L E C A L C U L A T I O N S / 153 A-7 Calculation of Evaporation Rate During Leaching Leaching medium is adjusted for the rate of evaporation accordingly in all relevant calculations. This sample calculation is based on the results for the experiment (Test No. 38 of data set #2) performed at the following parameters: Leaching Conditions: Solution Redox Potential Setpoint = 0.500 V A g / A g c i Temperature Setpoint = 60°C Data: mass of the leaching solution at completion (g) = 2000 mass of the leach residue (g) = 2.82 density of leaching solution at completion (g/mL) = 1.006 volume of leaching solution at completion = 2155.4 based on volume addition (mL) total duration of leaching experiment (min.) = 4037 solution res m mass final rsolution 1 total soluiion T Leaching m final res — TV) ytviut / _ solution mass i solution V V final R, ,. = —— ^wiss. (A-16) Evaporation rp . \ / Leaching = 0.042 mL/min - 0.04 mL/min Where: REn = evaporation rate during the course of leaching experiment APPENDIX A S A M P L E C A L C U L A T I O N S / 154 A-8 Statistical Significance of Temperature on Experimental Results The assessment of statistical significance of temperature on experimental results reveals important insights about the temperature dependency of leaching mechanisms. The effect of temperature on copper extraction [ Cu^™" /^"ef l r f(7 , I -» T2) ] is assessed through the differences between repeated tests at different temperatures [a higher temperature ( T2 ) and a lower temperature (71,)]. Similar calculations can be performed on other experimental results, including iron extractions and sulfur formations. This sample calculation is based on the results of copper extractions for experiments (Test No. 19 and 28 of data set #1, Test No. 38 and 42 of data set #2, Test No. 53 and 58 of data set #3) performed at the following parameters: Leaching Conditions: Solution Redox Potential Setpoint = 0.500 V A g / A g C i Temperature Setpoint = 60 and 70°C Data: Cu[^~mtliTesttm) Cu^:Lrl{Test#42) Extraction 1 AvgCalcHead copper extraction based on average calculated = 91.86 , head assay (%)(Test No, 19, data set #1, 60°C) (from Table B- l ) copper extraction based on average calculated =93.11 head assay (%)(Test No. 38, data set #2, 60°C) (from Table B - l ) copper extraction based on average calculated = 91.55 head assay (%) (Test No. 53, data set #3, 60°C) (from Table B- l ) copper extraction based on average calculated = 95.61 head assay (%) (Test No. 28, data set #1, 70°C) (from Table B- l ) : copper extraction based on average calculated = 93.05 head assay (%) (Test No. 42, data set #2, 70°C) (from Table B- l ) (Test#58) = copper extraction based on average calculated = 95.41 head assay (%) (Test No. 58, data set #3, 70°C) (from Table B- l ) APPENDIX A S A M P L E C A L C U L A T I O N S / 155 ry^r Extraction / rp rp \ - i Effect of Temperature on Copper Extraction V~UA vgCalcHead \ J \ ~~^  1 2 > \ Calculation X 1 [ • - T Extraction /rp \ s-i Extraction ('T \ \ 2^ AvgCalcHead^12 ) ~ ^UAvgCalcHead \ J 1 ) \ Cu^SSLi (T^T2)= (A- l 7) \ 1 Extraction tnc\o s-~*\ s~i Extraction / / A O / ~ * \ Zu VCu AvgCalcHead (70 C ) ~ C u AvgCalcHead (60 C)\ Cu^Gdc°Ld (Test#28 + TestMl + Test#5S) - Cu^~ead (Testffi 9 + Test#3S + Test#53) = 2.5167% -2.52% Where n = the number of repeated experiments at a given temperature or solution redox potential APPENDIX A S A M P L E C A L C U L A T I O N S / 156 A-9 Statistical Significance of Solution Redox Potential on Experimental Results The assessment of statistical significance of solution redox potential on experimental results reveals important insights about the redox potential dependency of leaching mechanisms. The effect of redox potential on copper extraction [ Cu^""l°^eMl(Ei —» E2) ] is assessed through the differences between repeated tests at different temperatures [a higher redox potential (E 2) and a lower redox potential (£",)]. Similar calculations can be performed on other experimental results, including iron extractions and sulfur formations. This sample calculation is based on the results of copper extractions for experiments (Test No. 19 and 28 of data set #1, Test No. 38 and 42 of data set #2, Test No. 53 and 58 of data set #3) performed at the following parameters: Leaching Conditions: Solution Redox Potential Setpoint = 0.400 V A g/ A gci and 0.500 V A g/ A gci Temperature Setpoint = 60°C Data: Cu Extraction AvgCalcHead Cu Extraction AvgCalcHead (Test#\9) (Test#38) /-» Extraction / n-> ,_i±c 1 \ CUAvgCalcHead ( T e S t # 5 3 ) Z Cu Extraction AvgCalcHead Cu Extraction AvgCalcHead (Test#2\)-(Test#36) Cu^eod(Test#5l)-copper extraction based on average calculated = 91.86 head assay (%)(Test No. 19, data set #1, 0.500 V A g / A gci) (from Table B- l ) copper extraction based on average calculated = 93.11 head assay (%)(Test No. 38, data set #2, 0.500 VAg/Agci) (from Table B- l ) copper extraction based on average calculated = 91.55 head assay (%) (Test No. 53, data set #3, 0.500 VAg/Agci) (from Table B- l ) copper extraction based on average calculated = 69.26 head assay (%)(Test No. 21, data set #1, 0.400 VAg/Agci) (from Table B- l ) copper extraction based on average calculated = 65.02 head assay (%)(Test No. 36, data set #2, 0.400 VAg / A g c i ) (from Table B- l ) copper extraction based on average calculated = 64.08 head assay (%)(Test No. 51, data set #3, 0.400 VAg/Agci) (from Table B- l ) APPENDIX A S A M P L E C A L C U L A T I O N S / 157 Effect of Solution Redox Potential on Copper Extraction i^uAr^ca"°,'lead {Ex —> E2) ] Calculation X ' \s~< Extraction / -p \ Extraction / 77 \ I / , AvgCalcHead ^ 2 ) AvgCalcHead V £ I )l Cu^SSLiW-> E2) = ^ L _ : (A-18) n I \P^~ead {0.500VAglAgCl) - Cu^~ead (0A00VAg!AgCI)] _ _/z! 3 _ Cu*™ead (Testm 9 + TestmB + Test#53) - Cu*™™md (Test#21 + Test#36 + Test#51) 3 = 26.0533% -26.05% Where n = the number of repeated experiments at a given temperature or solution redox potential APPENDIX A S A M P L E CALCULATIONS / 158 A-10 Statistical Analysis of Experimental Results Statistical analysis of experimental results from repeated experiments is assessed through weighted average of the variances (or squares of the standard deviations). This sample calculation is based on the results of copper extraction for experiments (Test No. 19 of data set #1, Test No. 38 of data set #2, and Test No. 53 of data set #3) performed at the following parameters: Leaching Conditions: Solution Redox Potential Setpoint = 0.500 V A g / A g c i Temperature Setpoint = 60°C Data: CuAZcaicHead (Test#\ 9) = copper extraction based on average calculated =91.86 head assay (%)(Test No. 19, data set #1) (from Table B- l ) Cufv7ca!cHead (Test#3S) = copper extraction based on average calculated = 93.11 head assay (%)(Test No. 38, data set #2) (from Table B - l ) CufvgCaicHemi (Test #53) = copper extraction based on average calculated = 91.55 head assay (%) (Test No. 53, data set"#3) ' (from Table B- l ) Average (X ) Calculation X (A-19) Cu*™eml (Testm 9) + Cu'~etiil (Test#3S) + C u ^ e a d {Test#53) 92.17% -92%. A P P E N D I X A S A M P L E C A L C U L A T I O N S / 159 Variance (s ) Calculation s2 =-!=!-I(^,-^)2 7 1 - 1 = 0.6852 = 0.69 Where: X = 92.17% from above calculation and n = 3. Standard Deviation (s) Calculation (A-20) s =ys = 0.8278% « 0.83% Where: 5 2 = 0.6852 from above calculation. (A-21) Standard Error of the Mean (se) Calculation se * 100% = 0.4779% = 0.48% Where: s = 0.8278% from above calculation and n = 3. (A-22) APPENDIX A S A M P L E C A L C U L A T I O N S / 160 A-11 Reference Electrode Potential Correction on a Standard Hydrogen Electrode (SHE) Scale Calomel and silver/silver chloride electrodes are the most commonly employed reference electrode systems in practices. They can be converted to the standard hydrogen electrode (SHE) scale, ETH , at various temperature by adding the electrode potential correction value, AE(T), as dE provided by equations (A-23) and (A-24). The E°H and —values of various electrode systems dT are given in Table A - l . Relevant standard electrode potential values versus SHE (E°) for the Ag I AgCl reference potential values with saturated KCl electrolyte employed in this thesis work is shown in Table A-2. dE AE(T) = E°H (T-29S) (A-23) dT El, = AE(T)+Ellem (A-24) Where: AE(T) = Electrode potential correction value of various reference electrode systems as a function of temperature ( V ) ETH = Standard Electrode potential at respective temperature of various reference electrode systems versus the standard hydrogen electrode scale ( V S H E ) ELsie,,, = Electrode potential at respective temperature of various reference electrode systems as a function of temperature ( V ) E"H = Standard electrode potential at 298 K of various reference electrode systems ( V ) c\E_ dT Change in electrode potential as a function of temperature of various electrode systems (V/K) T = Temperature in degree Kelvin (K) APPENDIX A SAMPLE CALCULATIONS / 161 Table A-l Standard electrode potential correction values [E°H and —] of some common dT reference electrodes systems as a function of temperature192 Reference Electrode System Equilibrium E° ( V ) — Electrolyte " ' dT (298 K) ( V / K ) Calomel (Hg/Hg2Cl2) (SCE) Hg2Cl2(s) + 2e~ <^2Hg(l) + 2CC(aq) 0.1 M KCl 0.336 0.00008 1 MKC1 0.283 0.00029 Saturated KCl 0.244 0.00067 Silver/Silver Chloride (AglAgCl) AgCl(s) + e~ <=> Ag(s) + Cl~ (aq) 0.1 M KCl 0.288 0.00050 3.5 M KCl 0.205 0.00073 Saturated KCl 0.199 0.00101 Table A-2 Standard electrode potential versus SHE (E°) in volts (VSHE) for the Ag I AgCl electrode system with saturated KCl electrolyte as a function of potential (0.400-0.600 V A g / A g c i ) and temperature (60-78°C) V (vs. Ag/AgCl, Sat.KCl) 0.400 V A g / A g c i 0.450 V A g / A g c i 0.500 V A g ' A g a -0.550 V A g / A g c i " 0.600 V A g / A g c i 60°C(333 K) 70°C(343 K) 78°C (351 K) 0.564 VSHE 0.554 VSHE 0.545 VSHE 0.614 VSHE 0.604 VSHE 0.595 V S H E 0.664 VSHE . 0.654 V S H E 0.645 VSHE 0.714 VSHE 0.704 VSHE 0.695 VSHE 0.764 VSHE 0.754 V S H E 0.745 VSHE Sample Calculation This sample calculation is for silver/silver electrode system for experiments performed at the following parameters: Leaching Conditions: EAZ/ASCI = EMLCI = Solution Redox Potential = 0.500 V (vs. Ag I AgCl, Sat. KCl) T= Temperature Setpoint = 60°C = 273 + 60 = 333 K Substitute data into equations (A-23) and (A-24) give: AE(333K) = 0.199 - 0.00101(333 - 298) = 0.16365 V Therefore, EH = A£(333iQ + £ ™ g f / =0.16365 + 0.500 = 0.66365 = 0.664 V S H E APPENDIX A S A M P L E C A L C U L A T I O N S / 162 A-12 The Criss-Cobble Extrapolation Method and E° Values for the FeM2+ Couple The standard reduction potential (E°T) at any temperature (T) can be calculated from the standard Gibbs energy of formation (AG°-) for a reaction according to equation (A-25). AG r 0 = J V p A G " - ] > XA G r = -nFE°T (A-25) P R The standard Gibbs energy {G°T) at any temperature (T) can be calculated from the following equation as shown by Dixon 1 9 3 , taking 25°C(298K) as the reference temperature: G°T = G°9S + C°P] \g0- (T - 298) S2°98 (A-26) ( T ^ Where: 0 = 7 , - 2 9 8 - 7 , l n , 2 9 8 , T = Temperature in degree Kelvin (K) Very little thermodynamic data are published for temperatures other than 25°C. Average values of heat capacity (C°J 2 9 8 ) under constant pressure for ionic species in the temperature range of interest have been estimated using the "entropy correspondence principle" developed by Criss and Cobble 1 6 0 ' 1 6 1 . The heat capacity constants (aT,/3T) are available for 60, 100, 150, 200, 250 and 300°C. These constants are specific for five difference classes of ions: protons, simple cations, simple anions and OH", oxyanions (XO~m), acid oxyanions {XOn{OH)~'"). Using this method, it is possible to estimate thermodynamic data at any temperature by extrapolation referencing to 25°C(298K) as shown in equation (A-27), where z is the ionic charge of aqueous species. C°P] 'm = aT + pr[S°m - (z)(20.92)] (A-27) By substituting equation (A-27) into (A-26), the general form of the relationship at any temperature, taking 25°C(298K) as the reference temperature, can be shown by equation (A-28): A G ; = A G ; 9 8 + A C ; ] 2r98(r - 298) - T A C ; ] rm i n ( ^ ) - As; 9 8(r - 298) (A-28) APPENDIX A S A M P L E C A L C U L A T I O N S / 163 Determination of (E°T) for ferric-ferrous (Fe3* I Fe2*) couple At temperatures other than 25°C, the standard reduction potential ( E° ) for the ferric/ferrous couple is approximated via the "entropy correspondence principle." The values of aT and (5Tat 60, 70 and 78°C are listed in Table A-3. The thermodynamics data of G°298, S°29% and z values for ferrous and ferric ionic species are listed in Table A-4. Equations (A-25), (A-26) and (A-27) give the values of heat capacity (C^] ), Gibbs energy (G°) and Gibbs energy change ( A G T ) of ferric and ferrous ions at different temperatures (Table A-5). Rearrangement of equation (A-25) gives the values of standard reduction potential ( E ° T ) values for the ferric-ferrous couple, which are in the range of 0.84-0.87 V S H E from 60 to 78°C (Table A-5). Table A-3 Criss-Cobble heat capacity constants (cx r, [5T ) for simple cations at 60, 70 and 78°C Temp. (T) PT 60°C 147 -0.405 70°C* 158.25 -0.442 78°C* 167.25 -0.4716 *Values at 70 and 78°C are interpolatec from values at 60°C and 100°C listed in reference (160). ; - : . fable A-4 G°298, S°2w and z values for ferrous and ferric ionic species at 25°C(298K) Species G°298 (J/mol)* 5°298 (J/mol-K)* z Fe2+(aq) -91,400 -107 2 Fe3*(aq) -16,700 -280.3 3 *A11 G ° , S° data from reference (194). Table A-5 C°P] ^ ,(G°R), AG^ and E° (VsHE,V A g /AgCi) values for ferrous and ferric ions at 60, 70and78°C Temp. (T) c°]T P i 2 9 8 (J/mol/K) (J/mol) A G ; (J/mol) (VSHE) (VAg/AgCl) Fe2+(aq) Fe3+(aq) Fe2+(aq) Fe3+(aq) Fe2+(aq) Fe3+(aq) 60°C 207 286 -88,065 -7,455 -80,610 80,610 0.835 0.671 70°C 224 310 -87,311 -5,090 -82,220 82,220 0.852 0.698 78°C 237 ' 329 -86,787 -3,310 -83,477 83,477 0.865 0.720 A P P E N D I X B S E L E C T E D L E A C H DATA APPENDIX B SELECTED LEACH DATA / 165 Table B- Summary of calculated head grade and extractions for iron and copper Test No. Temp. (°C) E ( V A g / A j j C l ) CU CalcHead ( % ) 1 ^CalcHead ( % ) Extraction ^"AvgCalcHead ( % ) Extraction " e AvgCalcHead ( % ) Leaching Time (hr:min) 19 60 0.500 28.91 31.76 91.86 46.58 68:00 20 0.550 28.74 31.25 89.12 52.96 60:00 21 0.400 28.82 31.74 69.26 46.08 60:00 22 0.450 29.21 31.59 87.72 46.34 60:00 23 0.600 31.28 32.20 55.31 38.48 60:00 24 70 0.400 31.15 31.91 71.16 44.11 22:00 25 0.450 29.47 32.31 89.30 51.05 22:45 27 0.550 29.92 32.06 91.23 51.30 18:00 28 0.500 29.96 32.46 95.61 46.72 20:00 29 0.600 28.76 34.13 74.32 45.97 23:30 30 78 0.400 30.42 33.82 75.80 44.05 20:00 31 0.450 29.16 31.29 90.61 53.62 21:00 32 0.500 31.35 30.70 99.69 54.72 19:00 34 0.550 31.45 34.73 93.20 59.55 9:00 35 0.600 29.01 39.92 69.44 49.21 10:30 36 60 0.400 29.09 32.23 65.02 49.26 59:30 37 0.450 28.54 31.34 86.07 44.09 60:00 38 0.500 29.13 32.50 93.11 43.55 66:46 39 78 0.400 29.37 31.43 82.67 51.04 23:00 40 70 0.400 28.25 30.69 71.81 46.49 22:00 41 0.450 31.81 31.12 93.50 56.63 21:00 42 0.500 29.30 32.04 93.05 45.40 18:30 43 0.550 28.94 34.12 90.18 52.32 20:00 44 0.600 29.18 33.89 63.89 36.27 21:00 45 60 '• 0.550 29.28 31.68 88.96 51.26 • 60:00 46 0.600 26.07 31.94 58.59 39.51 60:30 47 78 0.450 30.39 32.09 94.37 54.66 20:00 48 0.500 29.26 31.34 95.67 52.51 17:00 49 0.550 30.25 34.44 94.04 54.43 7:00 50 0.600 27.01 38.99 76.31 54.32 9:00 51 60 0.400 27.98 31.54 64.08 45.42 60:00 52 0.450 27.03 33.40 81.24 46.51 60:00 53 0.500 27.85 29.64 91.55 50.41 62:00 54 0.550 27.35 34.62 83.51 52.67 60:00 55 0.600 28.31 35.02 54.72 37.87 60:00 56 70 0.400 30.92 33.41 69.28 43.09 15:00 57 0.450 29.77 29.91 91.17 51.61 16:00 58 0.500 30.34 28.13 95.41 45.03 17:00 59 0.550 30.29 33.63 84.46 .48.02 18:00 60 0.600 27.13 33.94 56.08 32.30 19:00 61 78 0.400 29.89 30.30 71.87 43.72 20:00 62 0.450 31.02 33.12 97.52 55.55 20:00 63 0.500 30.38 34.21 100.15 53.99 19:00 64 0.550 31.19 35.81 93.17 60.01 10:00 65 0.600 28.94 38.37 78.04 53.63 10:00 Average : Standard Deviation (.v) : 29.37 32.82 1.32 2.31 = t f c SZ3 APPENDIX B SELECTED LEACH DATA / 166 Table B-2 Summary of sulfide sulfur conversion [ S^] - S2 Oxidation 5° (so24-)F] Test No. Temp. (°C) E ( V A g / A g C l ) S2~ Oxidation ( % ) s2; ( % ) (S02A~)F ( % ) 19 60 0.500 70.00 64.33 5.67 20 0.550 . 87.73 71.37 16.36 21 0.400 60.00 55.78 4.22 22 0.450 79.65 67.79 11.86 23 0.600 62.32 44.31 18.01 24 70 0.400 68.80 61.13 7.67 25 0.450 88.25 74.81 13.44 27 0.550 93.33 76.27 17.06 28 0.500 96.27 75.49 20.78 29 0.600 73.45 54.82 18.63 30 78 0.400 73.62 61.84 11.78 31 0.450 88.62 69.41 19.21 32 0.500 94.25 77.16 17.09 34 0.550 98.65 80.32 18.33 35 0.600 85.30 54.41 30.89 36 60 0.400 57.26 51.72 5.54 37 0.450 79.63 71.16 8.47 38 0.500 70.63 66.24 4.39 39 78 0.400 89.36 75.46 13.90 40 70 0.400 66.09 55.50 10.59 41 0.450 84.50 74.05 10.45 42 0.500 88.22 73.55 14.68 43 0.550 96.32 71.97 24.35 44 0.600 62.21 43.11 19.10 45 60 0.550 89.62 72.10 17.52 46 0.600 75.29 48.42 26.87 47 78 0.450 93.62 78.40 15.22 48 0.500 85.78 75.98 9.80 49 0.550 96.02 77.70 18.32 50 0.600 88.71 58.70 30.01 51 60 0.400 60.03 53.30 6.73 52 0.450 77.26 66.36 10.90 53 0.500 77.92 65.56 12.36 54 0.550 84.54 70.74 13.80 55 0.600 70.65 46.69 23.96 56 70 0.400 70.17 59.16 11.01 57 0.450 95.64 74.25 21.39 58 0.500 85.01 76.77 8.24 59 0.550 89.41 71.30 18.11 60 0.600 69.23 45.21 24.02 61 78 0.400 74.78 59.28 15.50 62 0.450 87.69 85.15 2.54 63 0.500 95.62 81.65 13.97 64 0.550 98.91 74.89 24.02 65 0.600 95.24 55.66 39.58 1/1 =tfc CD APPENDIX B SELECTED L E A C H D A T A / 167 Table B-3 Summary of copper extractions ( Cu'™"^"" ) based on the consumption of potassium permanganate (KMn0 4) solution Test No. Temp. (°C) E (VAg/AgCl) [KMnOA] (mol/L) [KMn04] Consumed (L) Extraction ^UKMn04 (%) \s~< Exirai lion Ar£Cah:Hvud-K\1nOA j (%) 19 60 0.500 0.2102 0.167 94.74 2.88 20 0.550 0.1964 0.172 91.07 1.95 21 0.400 0.1964 0.133 70.42 1.16 22 0.450 0.1964 0.169 89.39 1.67 23 0.600 0.1964 0.109 57.50 2.19 24 70 0.400 0.2059 0.131 72.78 1.62 25 0.450 0.2059 0.167 92.78 3.49 27 0.550 0.2031 0.175 95.53 4.30 28 0.500 0.2031 0.179 98.19 2.58 29 0.600 0.2031 0.142 77.90 3.58 30 78 0.400 0.2031 0.143 78.37 2.57 31 0.450 0.2031 0.172 94.17 3.56 32 0.500 0.2059 0.188 104.34 4.66 34 0.550 . 0.2059 0.179 99.35 6.15 35 0.600 0.2059 0.132 73.41 3.97 36 60 0.400 0.2163 0.119 69.40 4.38 37 0.450 0.2163 0.151 88.24 2.17 38 0.500 0.2163 0.160 93.40 0.29 39 78 0.400 0.2163 0.147 85.73 3.06 40 70 0.400 0.2163 0.125 73.04 1.24 41 0.450 0.2163 0.165 96.32 2.82 42 0.500 0.2163 0.164 95.83 2.79 43 0.550 0.2163 0.162 94.66 4.49 44 0.600 0.2163 0.116 67.36 3.47 45 60 0.550 0.1951 0.174 91.52 2.57 46 0.600 0.1951 0.118 61.81 3.21 47 78 0.450 0.1951 0.185 97.41 3.04 48 0.500 0.1951 0.188 98.89 3.22 49 0.550 0.1951 0.190 99.94 5.90 50 0.600 0.1951 0.156 81.79 5.48 51 60 0.400 0.1977 0.123 65.56 1.48 52 0.450 0.1977 0.156 82.96 1.72 53 0.500 0.1977 0.176 93.99 2.44 54 0.550 0.1977 0.160 85.45 1.94 55 0.600 0.1977 0.108 57.41 2.69 56 70 0.400 0.1977 0.134 71.15 1.87 57 0.450 0.2045 0.169 93.16 1.99 58 0.500 0.2045 0.179 98.59 3.18 59 0.550 0.2045 0.159 87.73 3.27 60 0.600 0.2045 0.107 58.98 2.90 61 78 0.400 0.2045 0.135 74.42 2.55 62 0.450 0.2045 0.183 100.97 3.45 63 0.500 0.2045 0.188 103.84 3.69 64 0.550 0.2045 0.180 99.42 6.25 65 0.600 0.2045 0.151 83.40 5.36 CZ5 CZ5 C D APPENDIX B SELECTED LEACH DATA / 168 Table B-4 Statistical analysis of copper extractions based on calculated head / s ' Extraction \ V ^"AvgCalcHead ) Test No. Temp. CQ E (VAg/AgCl) Data Source y^-i Extraction ^ U AvgCalcHead (%) Average X (%) Variance S2 Standard Deviation 5 (%) Standard Error se (%) 21 Set#l 69.26 36 0.400 Set #2 65.02 66.12 7.60 2.76 1.59 51 Set #3 64.08 22 Set#l 87.72 37 0.450 Set #2 86.07 85.01 11.33 3.37 1.94 52 Set #3 81.24 19 Set#l 91.86 38 60 0.500 Set #2 93.11 92.17 0.69 0.83 0.48 53 Set #3 91.55 20 Set#l 89.12 45 0.550 Set #2 88.96 87.20 10.21 3.20 1.84 54 Set #3 83.51 23 Set#l 55.31 46 0.600 Set #2 58.59 56.21 4.35 2.08 1.20 55 Set #3 54.72 24 Set#l 71.16 40 0.400 Set #2 71.81 70.75 1.73 1.31 0.76 56 Set #3 69.28 25 Set#l 89.30 41 0.450 Set #2 93.50 91.32 4.43 2.11 1.22 57 Set #3 91.17 28 Set#l 95.61 42 70 0.500 Set #2 93.05 " 94.69 2.04 1.43 0.82 58 Set #3 95.41 27 Set#l 91.23 43 0.550 Set #2 90.18 88.62 13.28 3.64 2.10 59 Set #3 84.46 29 Set#l 74.32 44 0.600 Set #2 63.89 64.76 83.78 9.15 5.28 60 Set #3 56.08 30 Set #1 75.80 39 0.400 Set #2 82.67 76.78 29.90 5.47 3.16 61 Set #3 71.87 31 Set #1 90.61 47 0.450 Set #2 94.37 94.17 11.98 3.46 2.00 62 Set #3 97.52 32 Set#l 99.69 48 78 0.500 Set #2 95.67 98.50 6.08 2.47 1.42 63 Set #3 100.15 34 Set #1 93.20 49 0.550 Set #2 94.04 93.47 0.25 0.50 0.29 64 Set #3 93.17 35 • Set #1 69.44 50 0.600 Set #2 76.31 74.60 20.70 4.55 2.63 65 Set #3 78.04 APPENDIX B SELECTED L E A C H D A T A / 169 Table B-5 Statistical analysis of iron extractions based on calculated head ( T? Extraction \ P eAvgCalcHead ) Test No. Temp. CQ E ( V A g / A g C l ) Data Source Extraction " e AvgCalcHead ( % ) Average X ( % ) Variance s2 Standard Deviation S ( % ) Standard Error se ( % ) 21 Set #1 46.08 36 0.400 Set #2 49.26 46.92 4.22 2.06 1.19 51 Set #3 45.42 22 Set#l 46.34 37 0.450 Set #2 44.09 45.65 1.82 1.35 0.78 52 Set #3 46.51 19 Set#l 46.58 38 60 0.500 Set #2 43.55 46.85 11.83 3.44 1.99 53 Set #3 50.41 20 Set#l 52.96 45 0.550 Set #2 51.26 52.30 0.83 0.91 0.53 54 Set #3 52.67 23 Set#l 38.48 46 0.600 Set #2 39.51 38.62 0.69 0.83 0.48 55 Set #3 37.87 24 Set #1 44.11 40 0.400 Set #2 46.49 44.56 3.04 1.74 1.01 56 Set #3 43.09 25 Set #1 51.05 41 0.450 Set #2 56.63 53.10 9.46 3.08 1.78 57 Set #3 51.61 28 Set#l . 46.72 42 70 ' 0.500 • Set #2 45.40 45.71 0.79 0:89 0.51 " 58 Set #3 45.03 27 Set#l 51.30 43 0.550 Set #2 52.32 50.55 5.06 2.25 1.30 59 Set #3 48.02 29 Set #1 45.97 44 0.600 Set #2 36.27 38.18 49.42 7.03 4.06 60 Set #3 32.30 30 Set #1 44.05 39 0.400 Set #2 51.04 46.27 17.09 4.13 2.39 61 Set #3 43.72 31 Set#l 53.62 47 0.450 Set #2 54.66 54.61 0.93 0.96 0.56 62 Set #3 55.55 32 Set #1 54.72 48 78 0.500 Set #2 52.51 53.74 1.27 1.13 0.65 63 Set #3 53.99 34 Set#l 59.55 49 0.550 Set #2 54.43 57.99 9.59 3.10 1.79 64 Set #3 60.01 35 Set#l 49.21 50 0.600 Set #2 54.32 52.39 7.70 2.77 1.60 65 Set #3 53.63 APPENDIX B SELECTED LEACH DATA / 170 Table B-6 Statistical analysis of total sulfide sulfur oxidation (S2 0 s i [ i a l j m ) Test No. Temp. (°C) E (VAg/AgCl) Data Source s2-Oxidation (%) Average X (%) Variance s2 Standard Deviation 5 (%) Standard Error se (%) 21 Set #1 60.00 36 0.400 Set #2 57.26 59.10 2.53 1.59 0.92 51 Set #3 60.03 22 Set #1 79.65 37 0.450 Set #2 79.63 78.85 1.89 1.37 0.79 52 Set #3 77.26 19 Set#l 70.00 38 60 0.500 Set #2 70.63 72.85 19.38 4.40 2.54 53 Set #3 77.92 20 Set #1 87.73 45 0.550 Set #2 89.62 87.30 6.59 2.57 1.48 54 Set #3 84.54 23 Set#l 62.32 46 0.600 Set #2 75.29 69.42 43.19 6.57 3.79 55 Set #3 70.65 24 Set#l 68.80 40 0.400 Set #2 66.09 68.35 4.31 2.08 1.20 56 Set #3 70.17 25 Set #1 88.25 41 0.450 Set #2 84.50 89.46 32.13 5.67 3.27 57 Set #3 95.64 28 Set #1 96.27 42 . 70 • 0.500 Set #2 88.22 , 89.84 33.67 - 5.80 3.35 58 Set #3 85.01 27 Set #1 93.33 43 0.550 Set #2 96.32 93.02 12.01 3.47 2.00 59 Set #3 89.41 29 Set #1 73.45 44 0.600 Set #2 62.21 68.30 32.24 5.68 3.28 60 Set #3 69.23 30 Set #1 73.62 39 0.400 Set #2 89.36 79.25 76.94 8.77 5.06 61 Set #3 74.78 31 Set #1 88.62 47 0.450 Set #2 93.62 89.98 10.17 3.19 1.84 62 Set #3 87.69 32 Set #1 94.25 48 78 0.500 Set #2 85.78 91.88 28.41 5.33 3.08 63 Set #3 95.62 34 Set #1 98.65 49 0.550 Set #2 96.02 97.86 2.56 1.60 0.92 64 Set #3 98.91 35 Set #1 85.30 50 0.600 Set #2 88.71 89.75 25.51 5.05 2.92 65 Set #3 95.24 APPENDIX B SELECTED LEACH DATA / 171 Table B-7 Statistical analysis of total elemental sulfur formation (S2tl) Test No. Temp. (°C) E ( V A g / A g C l ) Data Source <-,2-•V ( % ) Average X ( % ) Variance s2 Standard Deviation S ( % ) Standard Error se ( % ) 21 Set#l 55.78 36 0.400 Set #2 51.72 53.60 4.19 2.05 1.18 51 Set #3 53.30 22 Set#l 67.79 37 0.450 Set #2 71.16 68.44 6.08 2.46 1.42 52 Set #3 66.36 19 Set#l 64.33 38 60 0.500 Set #2 66.24 65.38 0.95 0.97 0.56 53 Set #3 65.56 20 Set#l 71.37 45 0.550 Set #2 72.10 71.40 0.47 0.68 0.40 54 Set #3 70.74 23 Set#l 44.31 46 0.600 Set #2 48.42 46.47 4.27 2.07 1.19 55 Set #3 46.69 24 Set#l 61.13 40 0.400 Set #2 55.50 58.59 8.17 2.86 1.65 56 Set #3 59.16 25 Set#l 74.81 41 0.450 Set #2 74.05 74.37 0.15 0.39 0.23 57 Set #3 74.25 28 Set#l 75.49 42 70 - 0.500 -. Set #2 73.55 • • 75.27- - .. :2.63 1.62 0.94 58 Set #3 76.77 27 Set#l 76.27 43 0.550 Set #2 71.97 73.18 7.26 2.70 1.56 59 Set #3 71.30 29 Set #1 54.82 44 0.600 Set #2 43.11 47.71 39.05 6.25 3.61 60 Set #3 45.21 30 Set #1 61.84 39 0.400 Set #2 75.46 65.52 75.73 8.70 5.02 61 Set #3 59.28 31 Set #1 69.41 47 0.450 Set #2 78.40 77.66 62.35 7.90 4.56 62 Set #3 85.15 32 Set#l 77.16 48 78 0.500 Set #2 75.98 78.26 8.95 2.99 1.73 63 Set #3 81.65 34 Set#l 80.32 49 0.550 Set #2 77.70 77.64 7.37 2.71 1.57 64 Set #3 74.89 35 Set #1 54.41 50 0.600 Set #2 58.70 56.26 4.88 2.21 1.27 65 Set #3 55.66 APPENDIX B SELECTED L E A C H D A T A / 172 Table B-8 Statistical analysis of total sulfate formation [ (SO, )F ] Test No. Temp. (°C) E (VAg/AgCl) Data Source (so2,-)F ( % ) Average X ( % ) Variance s2 Standard Deviation 5 ( % ) Standard Error se ( % ) 21 Set#l 4.22 36 0.400 Set #2 5.54 5.50 1.58 1.26 0.72 51 Set #3 6.73 22 Set#l 11.86 37 0.450 Set #2 8.47 10.41 3.06 1.75 1.01 52 Set #3 10.90 19 Set #1 5.67 38 60 0.500 Set #2 4.39 7.47 18.31 4.28 2.47 53 Set #3 12.36 20 Set#l 16.36 45 0.550 Set #2 17.52 15.89 3.61 1.90 1.10 54 Set #3 13.80 23 Set#l 18.01 46 0.600 Set #2 26.87 22.95 20.36 4.51 2.61 55 Set #3 23.96 24 Set#l 7.67 40 0.400 Set #2 10.59 9.76 3.31 •1.82 1.05 56 Set #3 11.01 25 Set#l 13.44 41 0.450 Set #2 10.45 15.09 32.00 5.66 3.27 57 Set #3 21.39 28 Set#l 20.78 42 70 0.500 Set #2 14.68 14.57 39.34 6.27 3.62 . 58 Set #3 8.24 -27 Set#l 17.06 43 0.550 Set #2 24.35 19.84 15.53 3.94 2.28 59 Set #3 18.11 29 Set #1 18.63 44 0.600 Set #2 19.10 20.58 8.93 2.99 1.73 60 Set #3 24.02 30 Set #1 11.78 39 0.400 Set #2 13.90 13.73 3.48 1.87 1.08 61 Set #3 15.50 31 Set#l 19.21 47 0.450 Set #2 15.22 12.32 75.79 8.71 5.03 62 Set #3 2.54 32 Set#l 17.09 48 78 0.500 Set #2 9.80 13.62 13.37 3.66 2.11 63 Set #3 13.97 34 Set #1 18.33 49 0.550 Set #2 18.32 20.22 10.80 3.29 1.90 ' 64 Set #3 24.02 35 Set#l 30.89 50 0.600 Set #2 30.01 33.49 27.95 5.29 3.05 65 Set #3 39.58 A P P E N D I X C A N A L Y T I C A L T E C H N I Q U E S APPENDIX C A N A L Y T I C A L TECHNIQUES / 174 C-1 Titrimetric Determination of Ferrous Iron Against Potassium Permanganate (KMnOJ Solution In strongly acidic solution (pH = 1), ferrous titration by potassium permanganate solution is represented by equation (A-13), where the permanganate ion is reduced to colorless Mn2+. Solutions of iron(II) salts have a pale greenish-blue color caused by the Fe(H20)2b+ complex ion. Before hydrolysis, solutions that contain ferric ions have a light purple color. In water, the iron(III) salts tend to be somewhat acidic and often appear slightly yellow because of the hydrolysis of the Fe(H20)l+ ion, which has Ka = 9x10" 4. Fe(H20)l+ (ag) + H20 <^ Fe(H20)5OH2+ (ag) + H^O+ (ag) (C-1) yellow The generated Fe(III) iron hydrolyzes forming a range of oxides with various colors (orange, red, yellow....). The color of these hydrolyzed iron compounds and the permanganate solution blend together to generate a continuum of pink, red to purple color, which makes the endpoint of titration hard to spot. A mixture of sulfuric and phosphoric acid, known as the "Spekker Acid" solution,, is added to the sample solution before titration. Phosphoric acid (HyPO,) complexes Fe(lll) ions, rendering colorless ferric complexes and prevent hydrolysis. Since protons (H+) are needed in the reduction of permanganate ions, sulfuric acid is also added to the solution. The endpoint is now marked by the transition from colorless to the faint pink color of unreduced permanganate solution, where both aqueous ions (Mn2+ and Fey+) are virtually colorless. Reagents: 1. Standardized Permanganate Solution (KMnO, solution) 2. Spekker Acid a. Measure out 1600 mL distilled-deionized H20 in 5 L beaker and agitate to wet the glass surface. b. Slowly and carefully add 600 mL concentrated (98%) sulfuric acid and then 600 mL concentrated (85%) phosphoric acid. c. Allow the solution mixture to cool before transferring to a glass storage bottle. Analytical Procedure: APPENDIX C A N A L Y T I C A L TECHNIQUES / 175 1. Add ~50 mL distilled-deionized H20 and ~ 2 mL Spekker acid (a few drops will suffice) into a 500 mL Erlenmeyer flask. 2. Pipette 1 mL of the test solution into and stir thoroughly. 3. Add more "Spekker acid" until the yellow solution turns to colorless. 4. Titrate with standardized KMn04 solution until the color changes from colorless to faint pink. If the faint pink persists for 30 seconds upon the addition of more H2S04 solution, the system has reached its endpoint. N.B.: 1. Ferrous titration with KMn04 is exothermic, generating slight heat during titration. 2. Phosphoric acid ( H i P 0 4 ) gives sever burn and offers less time before treatment is possible than by sulfuric acid ( H 2 S 0 4 ) . Lab coat and eye protection are definitely necessary while working with phosphoric acid. Sample Calculation SmolFe2* , n ,„ ' CKM„OA x • ^ X VT = Mil) (C-2) mol - MnO 4 Where: Fe(Il) = ferrous ion concentration (mol) VT — titration volume (L) CKMti04 - standardized potassium permanganate solution concentration (mol/L) Example using 0.12M KMn04 solution and 12 mL titrant: { Q n mol-MnO 4 y 5mol - Fe2* ^ ^ Q_3 £ = ^ Q_3 ^ ^ L mol • Mn04 References: Jeffery, G.H.; Bassett, J.; Mendham, J . ; Denney, R.C. Vogel's Textbook of Quantitative Chemical Analysis (5 t h ed.). Longman Scientific & Technical: New York, 1989, pp. 368-372. APPENDIX C A N A L Y T I C A L TECHNIQUES / 176 C-2 Standardization Procedure for 0.2 M (or 1 N) KMnOA Potassium permanganate salt (KMn04: 158.04 g/mol; solubility 1 8 3: 6.34 g KMn04 /TOO g H 2 0 at 20°C, or 0.4012 mol/L) is a powerful oxidizing agent of an intense violet color, which is easily reduced by the organic impurities present in the distilled and deionized water to a brown solid, manganese dioxide (Mn0 2 ) . Consequently, potassium permanganate is not a primary standard and must always be standardized, i.e., titrated against a primary standard reducing agent such as sodium oxalate (Nct2C204: 134.00 g/mol, 99.99% purity). In strongly acidic solution (pH = 1), sodium oxalate will become oxalic acid (a weak acid), which is oxidized by the permanganate ion (Mn04) to carbon dioxide gas (C02). The permanganate ion is reduced to colorless manganous ion (Mn2+). As long as any oxalic acid remains in the flash, each added drop of potassium permanganate loses its color. When all the oxalic acid had been consumed, the first excess drop of potassium permanganate will retain its deep color. In effect, potassium permanganate acts as its own indicator. IMnO;(ag) + 5H2C204(ag) + 6H + (ag) -» 2Mn2+ (ag) + l0CO2 (g) + SH20(l) (C-3) Since prolonged elevation of temperature promotes the oxidation of oxalic acid by atmospheric oxygen, 90-95% of sodium oxalate in sulfuric acid should be added rapidly at room temperature. The solution is then warmed to 55-60°C to promote the slow kinetics and titrated until completion. Since oxalate solutions attack glass, it should not be stored more than a few days. H2C204(ag) + 0.5O2 (g) -» H20(l) + 2C02 (g) (C-4) Reagents: 1. Reagent Grade Sodium Oxalate (Na 2 C 2 0 4 ) 2. 0.2 M (or IN) Permanganate Solution (KMn04 solution) a. Dissolve about 32.5 grams potassium permanganate and dilute to 1000 mL distilled and deionized water in a 1500 mL beaker. b. Heat the solution to boiling, and boil gently for 15-30 minutes. c. Allow the solution to cool to the laboratory temperature. d. Filter the solution through a glass funnel containing a plug of purified glass wool into a 1000 mL volumetric flask. e. Dilute to mark with distilled and deionized water and store in an amber bottle. APPENDIX C ANALYTICAL TECHNIQUES / 177 Analytical Procedure: 1. Add -100 mL distilled-deionized H20 and ~10 mL 6M H2S04 into a 500 mL Erlenmeyer flask. 2. Dissolve exactly 1.00g Na2C204 (7.463 mmol) and stir thoroughly. 3. Add 90% required amount of permanganate solution. 4. Heat the solution up to 55-60°C and titrate with KMnO, solution until the color changes from colorless to faint pink. If the faint pink persists for 30 seconds upon the addition of more H2S04 solution, the system has reached its endpoint. Sample Ca lcu la t ion mol.Na&O 2MnQ4 = 134g SNa2C204 ' KM"04 K ' Where: W = weight of sodium oxalate (Na2 C 20 4) used (g) Vr = titration volume (L) • : : . CKMll0, = potassium permanganate solution concentration (mol/L) Example using 14.2 mL KMn04 with l.OOg Na2C204: (1.00g Na2C204 x ^ ' ^ ^ x  2 M n ° 4 ) 4-14.2x10"3L = 0.2102M • KMnO, 134g 5Na2C20, References: Jeffery, G.H.; Bassett, J.; Mendham, J.; Denney, R.C. Vogel's Textbook of Quantitative Chemical Analysis (5th ed.). Longman Scientific & Technical: New York, 1989, pp. 368-372. APPENDIX C A N A L Y T I C A L TECHNIQUES / 178 C-3 Method of Multi-Acid ICP Multi-element Analysis [as provided by International Plasma Laboratory LTD (IPL), Vancouver, Canada, 1999] (a) 0.50 grams of sample was digested with a combination of HNO}, HC104 and HF acid until dryness was achieved. The vessel was then cooled and re-dissolved in a diluted Aqua-Regia solution by heating in a hot water bath. It was then cooled, bulked up to a fixed volume with demineralized water, and thoroughly mixed. (b) The specific elements were determined using an Inductively Coupled Argon Plasma spectrophotometer. A l l elements were corrected for inter-element interference. A l l data were subsequently stored onto computer diskette. Q U A L I T Y CONTROL The machine was First calibrated using six known standards and a blank. The test samples were then run in batches. A sample batch consisted of 38 or less samples. Two tubes were placed before a set. These were an in house standard and an acid blank, which were both digested with the samples. A known standard with characteristics best matching the samples was chosen and placed after every fifteenth sample. At every 38th sample (not including standards), two samples, chosen random, were reweighed and analyzed. At the end of a batch, the standard and blank used at the beginning was rerun. The readings for these known were compared with the pre-rack known to detect any calibration drift. APPENDIX C A N A L Y T I C A L TECHNIQUES / 179 C-4 Method of Iron Assay by Titration Analysis [as provided by International Plasma Laboratory LTD (IPL), Vancouver, Canada, 1999] (a) 0.25 to 2.00 grams of sample was digested in a multi-acid solution of concentrated HCl, HN03, H2S04 and HF until dryness was achieved. (b) Iron was separated from other interfering elements by ammonium hydroxide precipitation and filtration. The ferric ion was then redissolved in HCl acid, reduced to ferrous ion and complexes. The iron concentration in solution was determined by titrating against potassium dichromate using diphenylamine-sulphonic acid sodium salt as an indicator. (c) The result, in percentage, was calculated by standardizing the normality of potassium dichromate solution with a known iron standard. APPENDIX C A N A L Y T I C A L TECHNIQUES / 180 C-5 Method of Copper Assay by Titration Analysis [as provided by International Plasma Laboratory LTD (IPL), Vancouver, Canada, 1999] (a) 0.25 to 2.00 grams of sample was digested with multiple acids ( H C l , HNOj, H2S04 and HF) until dried. The sample was removed, cooled, and then boiled with 100 mL of Bromine water to dissolve any soluble matter. (b) Ammonium acetate, sodium fluoride and potassium iodide were added to the sample solution as complexing reagents. The concentration of copper was determined by titrating against sodium thiosulfate solution, using starch solution as indicator. Color range should be from dark brown to pale yellow before indicator was added. After indicator was added, color change should be from black to white. (c) The result, in percentage, was calculated by standardizing the normality of the sodium thiosulfate solution with a known copper standard. 

Cite

Citation Scheme:

        

Citations by CSL (citeproc-js)

Usage Statistics

Share

Embed

Customize your widget with the following options, then copy and paste the code below into the HTML of your page to embed this item in your website.
                        
                            <div id="ubcOpenCollectionsWidgetDisplay">
                            <script id="ubcOpenCollectionsWidget"
                            src="{[{embed.src}]}"
                            data-item="{[{embed.item}]}"
                            data-collection="{[{embed.collection}]}"
                            data-metadata="{[{embed.showMetadata}]}"
                            data-width="{[{embed.width}]}"
                            async >
                            </script>
                            </div>
                        
                    
IIIF logo Our image viewer uses the IIIF 2.0 standard. To load this item in other compatible viewers, use this url:
http://iiif.library.ubc.ca/presentation/dsp.831.1-0078711/manifest

Comment

Related Items