Fundamentals of Partial Pressure Oxidation of Gold Bearing Sulphide Minerals

FUNDAMENTALS OF PARTIAL PRESSURE OXIDATION OF GOLD BEARING SULPHIDE MINERALS

Ambrosia Maria Ivana BE (Chemical Engineering)

A thesis submitted for the degree of Doctor of Philosophy at The University of Queensland in April in 2020 School of Chemical Engineering

Abstract

Pressure oxidation (POX) technology has been used worldwide to treat refractory gold ores. Despite decades of operation, the chemistry inside the autoclave has not been fully established, and in particular the chemistry of iron is of critical importance. The aims of this thesis are to investigate the role of ferric as surrogate oxidant, determine the solubility of hematite, basic ferric sulphate (BFS) and potassium jarosite at 220°C and develop Eh-pH diagrams for the Fe-S-K-H2O and Fe-S-H2O systems which cover the range of potential operating conditions.

The role of ferric as pyrite surrogate oxidant was investigated in an oxygen-deprived environment at 220°C using a microwave digester. The results showed that ferric was capable to act as pyrite surrogate oxidant in the high temperature system despite of the removal of ferric from solution via iron hydrolysis. Pyrite oxidation by ferric was found to be surface limited. The oxidation extent was found to be linearly dependent with ferric concentration when the reaction has not reached the surface limit. Ferrous was found to hinder pyrite oxidation but its negative impact was able to be offset to some extent with higher solid loading (i.e. surface area).

The solubility studies were carried out in a titanium autoclave with high temperature sampling that allows rapid solid-liquid separation and in-line ORP measurement. The results 0 from hematite solubility study indicated that aqueous FeHSO4SO4 species was the 0 predominant species at higher acidity solution and it switched to aqueous Fe2(SO4)3 at lower acidity by considering the reaction stoichiometry. New Gibbs free energy of formation 0 -1 (ΔGf ) data for BFS at 220°C was determined to be -919.4 ± 0.3 kJ mol . Similarly, the Gibbs free energy of formation data for potassium jarosite at 220°C was determined to be -3009.4 ± 0.7 kJ mol-1 which is within the range of values reported in the literature. Subsequently,

Eh-pH diagram for Fe-S-H2O and Fe-S-K-H2O system at 220°C were developed based on this new data and validated.

Using the developed Eh-pH diagram, a combination of potassium jarosite, hematite, BFS were predicted to precipitate in Lihir autoclave depending on the acidity and potassium concentration. Potassium jarosite was identified as the possible main lime consumers. Due to the small particle size (<6 µm), possible surface defect and crystallinity of the potassium jarosite precipitated at Lihir autoclave, its kinetic decomposition with lime could be relatively

ii fast. Basic ferric sulphate was also predicted to form if the sulphuric acid concentration in Lihir autoclave is higher than approximately 35 g/L at autoclave temperature.

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Declaration by author This thesis is composed of my original work, and contains no material previously published or written by another person except where due reference has been made in the text. I have clearly stated the contribution by others to jointly-authored works that I have included in my thesis.

I have clearly stated the contribution of others to my thesis as a whole, including statistical assistance, survey design, data analysis, significant technical procedures, professional editorial advice, financial support and any other original research work used or reported in my thesis. The content of my thesis is the result of work I have carried out since the commencement of my higher degree by research candidature and does not include a substantial part of work that has been submitted to qualify for the award of any other degree or diploma in any university or other tertiary institution. I have clearly stated which parts of my thesis, if any, have been submitted to qualify for another award.

I acknowledge that an electronic copy of my thesis must be lodged with the University Library and, subject to the policy and procedures of The University of Queensland, the thesis be made available for research and study in accordance with the Copyright Act 1968 unless a period of embargo has been approved by the Dean of the Graduate School.

I acknowledge that copyright of all material contained in my thesis resides with the copyright holder(s) of that material. Where appropriate I have obtained copyright permission from the copyright holder to reproduce material in this thesis and have sought permission from co- authors for any jointly authored works included in the thesis.

Submitted manuscripts included in this thesis

No manuscripts submitted for publication

Other publications during candidature

Ivana, A., Vaughan, J. and Hawker, W. 2017. “Ferric sulphate precipitation at gold pressure oxidation conditions”, Proceedings of Conference of Metallurgist (COM) 2017, Vancouver, Canada, 28 - 31 August 2018.

Contributor Statement of contributions

Ambrosia Ivana Designed experiments (%) 100

Wrote the paper (%) 80

Dr. James Vaughan Edited and discussed the paper (%) 10

Dr. William Hawker Edited and discussed the paper (%) 10

Contribution by others to the thesis

Associate Professor James Vaughan and Dr William Hawker provided technical support for project design, experimental design and data analysis. They have reviewed technical outcomes of this thesis and provided invaluable guidance.

Dr. Kathryn Stewart and Dr. Miree Leslie provided valuable advice on QXRD analysis and ore mineralogy.

John O’Callaghan and Luke Vollert provided technical advise and reviewed the thesis.

Statement of parts of the thesis submitted to qualify for the award of another degree

None

Research Involving Human or Animal Subjects

No animal or human subjects we involved in this research

Acknowledgements

Firstly, I would like to express my deepest gratitude and utmost appreciation to my advisors, Associate Professor James Vaughan and Dr. William Hawker, for their continuous support, constant guidance, patience and motivation throughout this Ph.D. Thank you for not giving up on me. Particular thanks go to the rest of my thesis committee members: Dr. Greg Birkett, Dr. Grant Ballantyne and Dr Liguang Wang, for their insightful comments and constructive suggestions.

I would also like to sincerely thank Newcrest Mining, in particular John O’Callaghan and Luke Vollert for their continuous financial support and insightful comments which make this study possible. Also, Dr. Kartherine Stewart and Dr. Miree Leslie for their continuous support in mineralogy side of the study.

My sincere thanks also goes to Professor Edouard Asselin, who gave access to the hydrometallurgy laboratory at the University of British Columbia and supported me throughout my time there together with Dr. Jing Liu and Baseer Abdul.

Special thanks to all my friends at the university: Luisa, Dilini, John, Pritii, Harrison, Stefan, Yufan, Na, George, Toby and all of you who I’ve shared fun moments with in the last four years. Many thanks to James G for his continuous support in the lab especially with the autoclave experiments.

Last but not the least, I would like to thank my family, especially my mom Catharina and my sister Janice for their understanding, continuous prayer and moral support throughout my Ph.D.

Financial support

This research was supported by an Australian Government Research Training Program Scholarship.

Financial support and a titanium autoclave for this project was provided Newcrest Mining Pty Ltd.

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Keywords Pressure oxidation, gold, iron, hydrolysis, basic ferric sulphate, hematite, potassium jarosite, Gibbs energy of formation, solubility, Eh-pH diagram

Australian and New Zealand Standard Research Classifications (ANZSRC) ANZSRC code: 091403, Hydrometallurgy, 80% ANZSRC code: 090499 Chemical Engineering not elsewhere classified, 20%

Fields of Research (FoR) Classification FoR code: 0914, Resources Engineering and Extractive Metallurgy, 80% FoR code: 0904, Chemical Engineering, 20%

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Table of Contents

Chapter 1: Project background and literature review 1 1.1 Gold ...... 1 1.2 Refractory gold ore ...... 2 1.3 Pressure oxidation ...... 6 1.3.1 Pyrite Oxidation Chemistry 7 1.3.2 Ferric hydrolysis 15 1.3.3 Factors affecting iron hydrolysis product 19 1.3.4 Ferrous solubility 25 1.4 Industry context ...... 27 1.4.1 Lihir pressure oxidation process overview 29 1.4.2 Lihir partial pressure oxidation operation 30 1.4.3 Lihir iron hydrolysis product 31 1.4.4 Lihir lime consumption 32

1.5 Thermodynamic model for Fe-SO4-H2O system at high temperature ...... 35 1.6 Summary ...... 38 1.7 Aim and objectives ...... 39 1.8 Outline of the thesis ...... 40 Chapter 2: Common Experimental methodology 41 2.1 Materials ...... 41 2.1.1 Reagents 41 2.1.2 Pyrite crystal 41 2.1.3 Lihir Ore 42 2.2 Equipment ...... 42 2.2.1 Speedwave Four Microwave Digester 42 2.2.2 Autoclave 43 2.2.3 High temperature sampling 44 2.2.4 In-situ high temperature ORP 45 2.3 pH measurement ...... 48 2.4 Ferrous titration ...... 49 2.5 Solid identification ...... 49 2.5.1 Scanning electron microscope (SEM) 49 2.5.2 X-ray Diffraction (XRD) 50 2.5.3 Quantitative X-Ray Diffraction (QXRD) 50 2.6 Solution analysis ...... 51 Chapter 3: The role of ferric as surrogate oxidant during pyrite pressure oxidation 52 ix

3.1 Introduction ...... 52 3.2 Experimental ...... 55 3.2.1 Equipment 55 3.2.2 Methodology 55 3.2.3 95% Confidence interval calculation 56 3.2.4 Ferrous Quantification 57 3.2.5 Pyrite oxidation determination 58 3.2.6 Ferric balance 59 3.2.7 Quantitative X-ray Diffraction method 60 3.3 Result ...... 61 3.3.1 Effect of initial ferric concentration and solid loading on pyrite oxidation 61 3.3.2 Effect of initial ferric concentration and solid loading on ferric deportment 64 3.3.3 Effect of ferrous concentration 67 3.3.4 The combined effect of initial ferric concentration, solid loading and ferrous concentration on pyrite oxidation 69 3.3.5 Solid phase precipitates 71 3.4 Implication of findings...... 75 3.4.1 Understanding confliction results of previous studies 75 3.4.2 Change of ferrous and ferric concentration at Lihir 76 3.5 Summary ...... 77 Chapter 4 Solubility of hematite, basic ferric sulphate and potassium jarosite 79 4.1 Introduction ...... 79 4.2 Experimental design ...... 81 4.2.1 Solubility experiment methodology 81 4.2.2 Residence time – equilibrium determination 83 4.3 Chemical thermodynamic approach ...... 84 4.3.1 Thermodynamic calculation 84 4.3.2 Activity coefficient estimation 85 4.3.3 Solid and Aqueous species selection 87 4.3.4 pH estimation at high temperature 94 4.3.5 95% confidence interval methodology 95 4.3.6 High temperature thermodynamic data review 95 4.4 Assumptions ...... 99 4.5 Solubility of hematite at 220°C ...... 99 4.5.1 Solid identification and morphology 99 4.5.2 Solution assay 102

0 4.5.3 Fitting of the FeHSO4SO4 and Fe2(SO4)3 aqueous species 106 x

4.6 Solubility of basic ferric sulphate at 220°C ...... 108 4.6.1 Solid XRD and morphology 108 4.6.2 Results and discussion 109 4.6.3 Equilibrium between basic ferric sulphate, hematite and aqueous ferric species 113 4.7 Solubility of potassium jarosite at 220°C ...... 114 4.7.1 Solid XRD and morphology 114 4.7.2 Results and discussion 115

4.8 Development of the Eh-pH diagram at 220°C ...... 119 4.8.1 Methodology 119 4.8.2 Reactions 120

4.8.3 Eh-pH diagram of Fe-S-H2O at 220°C 121

4.8.4 Eh-pH diagram of Fe-K-S-H2O at 220°C 132 4.9 Summary ...... 143 Chapter 5: Industrial implication 145 5.1 Introduction ...... 145 5.2 Experimental methodology ...... 149 5.2.1 Experimental plan 149 5.2.2 Oxygen injection methodology 150 5.2.3 ORP data processing 152 5.2.4 QXRD analysis 153 5.3 Results ...... 153 5.3.1 QXRD on solid residues 153 5.3.2 Solution assay 157 5.3.3 Solid assay 161 5.3.4 ORP Data 161 5.2 Discussion ...... 167

5.3 Eh-pH thermodynamic model validation ...... 169

5.4 Application of Eh-pH model using Lihir site data ...... 171 5.7 Industrial implication ...... 175 5.8 Summary ...... 177 Chapter 6: Conclusion and recommendations 179 6.1 Conclusion ...... 179 6.2 Industrial implication ...... 180 6.3 Limitations and recommendations ...... 180 References 182 Appendix A 189

Appendix A.2 XRD patterns from hematite solubility experiments ...... 190 Appendix A.3 XRD patterns from basic ferric sulphate solubility experiments ...... 191 Appendix A.4 XRD patterns from potassium jarosite solubility experiments ...... 192 Appendix B 193 Appendix B.1 Liquid-liquid Junction Potential sample calculation ...... 193 Appendix B.2 Lime consumption calculation ...... 196 Appendix C 199 Appendix C.1 pH Estimation at high temperature ...... 199 Appendix D QXRD Codes for TOPAS 203

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List of Figures

Figure 1 Gold prices (USD/oz) of the day in the last 50 years ...... 1 Figure 2 Horizontal autoclave at Barrick Goldstrike Mine, USA (Thomas & Pearson, 2016)6 Figure 3 General illustration of pyrite oxidation during pressure oxidation process ...... 8 Figure 4 Rate of thiosulphate oxidation by ferric vs. acid decomposition as a function of pH (adapted from Williamson & Rimstidt, 1993) ...... 13 Figure 5 Tetrathionate possible oxidation paths (Druschel et al., 2003) ...... 14 Figure 6 Typical hematite crystals (Mindat, 2019) ...... 16 Figure 7 Schematic illustration of hematite – basic ferric sulphate - jarosite stability. Dashed 2 3 lines represent contours of log a Fe3+ + a SO42- and increase to the left (Stroffregen, 1993) ...... 17

Figure 8 Structure of KFe3(SO4)2(OH)6. (a) Single unit, (b) View along a and b axes, (c) View along c- and a-axis (Das et al., 1996) ...... 17

Figure 9 Projection of the structure of FeOHSO4 along the a-axis (left) and b-axis (right) (Johansson, 1962)...... 18 Figure 10 TEM image of basic ferric sulphate with magnification of 41000X (Gomez et al., 2013) ...... 18 Figure 11 Relationship between sulphur content in hydrolysis product and free sulphuric acid concentration in Fe-S-H2O system at 170°C, 185°C and 200°C (Tozawa & Sasaki, 1986) ...... 19 Figure 12 Jarosite stability field at 200°C and 100 bar with log m ∑S = -0.5 and log m K = - 1.5 (Stroffregen, 1993) ...... 20 3+ + 2- Figure 13 Ferric stability diagram for the Fe -K -SO4 -H2O system as a function of temperature and pH (adapted from Babcan, 1971) ...... 21

Figure 14 The effect of ferric concentration Potassium jarosite stability field in Fe-K-SO4- + 2- H2O system at 95 °C, [K ] = 0.01M and [SO4 ] = 1M (Umetsu et al., 1977) ...... 23 Figure 15 The effect of hematite seed addition on product yield and composition of the product precipitated from 0.5M Fe(SO4)1.5 solution at pH 1.4 at 225°C with 2-hour retention time (Dutrizac & Chen, 2011) ...... 24 Figure 16 Iron precipitation curves for Na-Jarosite precipitation in the presence of K-jarosite seed at 98°C (Dutrizac, 1999) ...... 25 Figure 17 Ferrous sulphate solubility as a function of temperature (adapted from Cheng, 2002 and Hasegawa et al., 1998) ...... 26

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Figure 18 SEM image of FeSO4.H2O crystals with magnification of 1,500X (Elgersma et al.,1993) ...... 27 Figure 19 Secondary Ion Mass Spectrometry result of Lihir ore (Ketcham et al., 1993) ..... 28 Figure 20 Lihir Processing Plant Block Flow Diagram (Newcrest, 2013) ...... 29 Figure 21 Mineral liberation analyser (MLA) for three different particle sizes of Lihir Ore .. 30 Figure 22 The relationship between hematite and jarosite in the Lihir neutralisation feed (Newcrest Mining, 2015) ...... 32 Figure 23 Solubility profiles of several metals hydroxides as a function of pH (Aube & Zinck, 2003) ...... 33 Figure 24 Partially decomposed jarosite particle made of a gel halo, a reaction front and an unreacted core (Patino et al., 2013) ...... 35 +3 -2 Figure 25 Iron stability diagram for the Fe -SO4 -H2O system as a function of temperature and pH (adapted from Fleming, 2009) ...... 36

Figure 26 Eh-pH diagram of Fe-S-H2O at 200°C (Biernat & Robins, 1972) ...... 37

Figure 27 Eh-pH diagram of Fe-S-H2O at 220°C (Huang, 2008) ...... 38 Figure 28 Speedwave Four Microwave Digester and its control box (left), DAK-100 TFM pressure vessel (bottom right) and TFM liner and ceramic pressure jacket (top right) ...... 43 Figure 29 Autoclave setup at UQ Hydrometallurgy ...... 44 Figure 30 High Temperature Autoclave sampling setup ...... 45 Figure 31 In-situ ORP Probe ...... 45 Figure 32 Estimation of conversion from Ag/AgCl reference electrode to standard hydrogen electrode (SHE) as a function of KCl concentration ...... 47

Figure 33 Equilibrium solubility of Fe(OH)3 at 25°C (Kim et al., 2015) ...... 52 Figure 34 Oxygen solubility in pure water as a function of temperature (Tromans, 2000).. 53 Figure 35 Pyrite conversion over time as a function of oxygen partial pressure and temperature at ferric concentration of 0.2 M (left) and 1 M (right) (King & Lewis, 1980) ..... 54 Figure 36 Typical Fe2+ concentration profile as a function of time (Cheng, 2002) ...... 58 Figure 37 Iron mass balance illustration ...... 59 Figure 39 Example of Rietveld refinement graphics from QXRD analysis in TOPAS ...... 61 Figure 40 The effect of initial ferric concentration on pyrite oxidation in microwave digester experiment at 220°C under deoxygenated conditions at 2.5 g-pyrite/L ( ∆ ), 5 g-pyrite/L () and 10 g-pyrite/L (O) solid loading...... 63

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Figure 41 the effect of initial ferric concentration on the amount of ferric used for pyrite oxidation at 220°C under deoxygenated conditions at 2.5 g-pyrite/L ( ∆ ), 5g-pyrite/L () and 10g-pyrite/L (O) solid loading...... 64 Figure 42 the effect of initial ferric concentration on ferric utilisation (%) for pyrite oxidation at 220°C under deoxygenated conditions at 2.5 g-pyrite/L ( ∆ ), 5g-pyrite/L () and 10g- pyrite/L (O) solid loading...... 66 Figure 43 the effect of initial ferric concentration on filtrate pH (left-axis) and ferric concentration in the filtrate (right-axis) for pyrite oxidation at 220°C under deoxygenated conditions at 2.5 g-pyrite/L ( ∆ ), 5g-pyrite/L (  ) and 10g-pyrite/L (O) solid loading...... 67 Figure 44 the effect of ferrous concentration on ferric utilisation and pyrite oxidation at 220°C for 2 hours, 5g/L solid loading, 25g/L [Fe3+] ...... 68 Figure 45 Ferric utilisation at various initial ferric concentration, ferrous concentration and solid loading (  = 5 g-pyrite/L, 25 g/L Fe3+ , ▲ = 5 g-pyrite/L, 50 g/L Fe3+, ◇ = 10 g- pyrite/L, 25 g/L Fe3+, O = 10 g-pyrite/L, 50 g/L Fe3+) ...... 70 Figure 46 SEM images from 2.5 g/L solid loading at (A) 16 g-Fe3+/L, (B) 27 g-Fe3+/L and (C) 35 g-Fe3+/L, 5g/L solid loading at (D) 14 g-Fe3+/L, (E) 27g-Fe3+/L and (F) 35 g-Fe3+/L and 10 g/L solid loading at (G) 14 g-Fe3+/L and (H) 51 g-Fe3+/L...... 74 Figure 47 Monthly average of ferrous concentrations in Lihir autoclave discharge as a function of oxidation extent from plant data in December 2013 to December 2015 ...... 77 Figure 48 Aqueous module in HSC Chemistry v9 ...... 86 - Figure 49 Activity coefficient of HSO4 species as a function of concentration and temperature generated using HSC Chemistry v.9 at [H+] = 0.158 and [OH-] = 6.3E-14 ...... 87

Figure 50 Eh-pH diagram for sulphur at 220°C, total P = 32bar, ∑S = 0.224M and ionic strength = 0.28. Pressure of H2S (g) was set at 0.1 bar...... 88 Figure 51 Ferric aqueous species distribution of at equilibrium with hematite at 200°C (Papangelakis, 1994) ...... 90 Figure 52 Ferric aqueous species distribution at equilibrium with hematite at 250°C (Liu et al., 2003) ...... 91 Figure 53 Potassium aqueous speciation at 220°C, total [S] = 0.2 m and total [K] = 0.05 m ...... 94 Figure 54 Pyrite crystal cluster precipitated from 12 g/L Fe3+ solution with no addition of LiOH (A), and 17 g/L Fe3+ solution with no addition of LiOH (B), 13 g/L Fe3+ solution with addition of 1 g/L LiOH (C and D), 5 g/L LiOH (E) and 10 g/L LiOH (F). The magnification of A is x4000, B is x5500, C is x5,000 and D, E, F are x20000...... 101

0 Figure 55 Equilibrium line fitting for hematite solubility data with aqueous FeHSO4SO4 species at 푎퐻푆푂4 − of 0.2 ...... 104

Figure 56 Equilibrium line fitting for hematite solubility data with aqueous Fe2(SO4)3 species at 푎퐻푆푂4 − of 0.2 ...... 106

Figure 57 Fitting of equilibrium line between hematite, aqueous Fe2(SO4)3 and aqueous

0 FeHSO4SO4 species at 푎퐻푆푂4 − of 0.2 ...... 107 Figure 58 Distribution of aqueous Fe(III) species as a function of pH at 푎퐻푆푂4 − of 0.2. 108 Figure 59 Morphology of basic ferric sulphate precipitated in basic ferric sulphate solubility experiment. Magnification 5,000X, 15kV and 10mm WD...... 109 Figure 60 Equilibrium line fitting for basic ferric sulphate solubility data with aqueous 0 FeHSO4SO4 species at 푎퐻푆푂4 − of 0.2...... 111 Figure 61 Equilibrium pH between basic ferric sulphate and hematite as a function of bisulphate activity ...... 112 Figure 62 Solubility of hematite and basic ferric sulphate as a function of pH and aqueous ferric species activity ...... 114 Figure 63 Potassium Jarosite Morphology. Magnification 5000X, 15kV and 10mm WD .. 115 Figure 64 Equilibrium line fitting for potassium jarosite solubility data with aqueous 0 FeHSO4SO4 species at 푎퐻푆푂4 − of 0.2...... 117

Figure 65 Illustration #1 of Eh-pH diagram for Fe-S-H2O system at 220°C ...... 123

Figure 66 Illustration #2 of Eh-pH diagram for Fe-S-H2O system at 220°C ...... 124

Figure 67 Illustration #3 of Eh-pH diagram for Fe-S-H2O system at 220°C ...... 124

Figure 68 Fitting of Hematite solubility data on Eh-pH diagram for Fe-S-H2O system at 220°C at 푎퐻푆푂4 − of 0.1 and 푎퐹푒퐻푆푂4푆푂40 of 0.01...... 126 Figure 69 Mixture of basic ferric sulphate and hematite precipitate. Magnification 5,000X, 15 kV, WD 10mm...... 127 Figure 70 Fitting of pure basic ferric sulphate and mixed basic ferric sulphate-hematite precipitates on the Eh-pH diagram at 푎HSO4 − of 0.3, 푎FeHSO4SO40 of 0.1 and 푎Fe2 + of 0.0001 ...... 129

Figure 71 Effect of bisulphate ion activity on Eh-pH diagram for Fe-S-H2O system at 220°C with 푎퐹푒퐻푆푂4푆푂40 = 0.1 and 푎퐹푒2+= 1 × 10 − 4 ...... 130

Figure 72 Effect of aqueous ferric complex activity on Eh-pH for Fe-S-H2O system at 220°C with 푎퐻푆푂4 − = 0.1 and 푎퐹푒2+= 1 × 10 − 4 ...... 131

Figure 73 Effect of aqueous ferrous activity on Eh-pH for Fe-S-H2O system at 220°C evaluated at 푎FeHSO4SO40 = 0.1 and 푎HSO4 − = 0.1 ...... 132

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Figure 74 Illustration #1 for Eh-pH diagram of Fe-K-S-H2O system at 220°C ...... 134

Figure 75 Illustration #2 for Eh-pH diagram of Fe-K-S-H2O system at 220°C ...... 135

Figure 76 Illustration #3 for Eh-pH diagram of Fe-K-S-H2O system at 220°C ...... 135

Figure 77 Illustration #4 for Eh-pH diagram of Fe-K-S-H2O system at 220°C ...... 136 Figure 78 Mixed potassium jarosite-hematite precipitate formed with the addition of 4 g/L (left) and 7g/L (right) lithium hydroxide. Magnification 5,000X, 15 kV and 10mm WD...... 137

Figure 79 Plot of pure potassium jarosite and mixed precipitate experiments data on Eh-pH diagram for Fe-S-K-H2O system at 220°C. The Eh-pH diagram was plotted with 푎퐹푒퐻푆푂4푆푂40 = 0.02, 푎퐻푆푂4 − = 0.1, 푎퐾 + = 0.1 and 푎퐹푒2+= 1 × 10 − 4 ...... 139

Figure 80 The effect of potassium activity on Eh-pH diagram for Fe-K-S-H2O system at 220°C with 푎퐹푒퐻푆푂4푆푂40 = 0.1, 푎퐻푆푂4 − = 0.1 and 푎퐹푒2+= 1 × 10 − 4 ...... 140 - Figure 81 The effect of bisulphate activity (a HSO4 ) on Eh-pH diagram for Fe-K-S-H2O system at 220°C with 푎퐹푒퐻푆푂4푆푂40 = 0.1, 푎퐾 += 0.1 and 푎퐹푒2+= 1 × 10 − 4...... 141 Figure 82 The effect bisulphate activity (푎퐻푆푂4 −) on minimum potassium concentration for potassium jarosite solid stability ...... 141 0 Figure 83 The effect of aqueous ferric activity (a FeHSO4SO4 ) on Eh-pH diagram for Fe-

K-S-H2O system at 220°C with 푎퐻푆푂4−= 0.1 , 푎퐾+= 0.1 and 푎퐹푒2+= 1 × 10 − 4 ...... 142 2+ Figure 84 The effect of aqueous ferrous activity (a Fe ) on Eh-pH diagram for Fe-K-S-H2O system at 220°C with 푎퐹푒퐻푆푂4푆푂40 = 0.1, 푎퐻푆푂4−= 0.1 and 푎퐾+= 0.1 ...... 143 Figure 85 Lihir Lime consumption breakdown estimate during high oxidation period based on Lihir plant data from December 2013 to January 2014 ...... 148 Figure 86 Lihir Lime consumption breakdown estimate during low oxidation period (Partial POX) based on Lihir plant data from October 2015 to November 2015 ...... 148 Figure 87 Oxygen injection setup ...... 151 Figure 88 Illustration of beam absorption in high and low absorbing minerals (Madsen, 2017) ...... 156 Figure 89 Illustration of electrode potential as a function of time in POX test ...... 163 Figure 90 Measured electrode potential during POX of Lihir gold-bearing sulphide ore at 220°C at different solid loadings (Test ID: Ore 1 to Ore 3) with 120 minutes retention time ...... 164 Figure 91 Measured electrode potential during POX of Lihir gold-bearing sulphide ore at 220°C at different retention time and 20 wt% solid loading (Test ID: Ore 2, Ore 4, Ore 5) ...... 166

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Figure 92 Measured electrode potential during POX of Lihir gold-bearing sulphide ore at 220°C at different initial ferric and ferrous concentration (Test ID: Ore 2, Ore 6, Ore 7) .. 167

Figure 93 Eh-pH diagram for Lihir ore POX experiment at 220°C ...... 170

Figure 94 Eh-pH diagram for high potassium condition at 220°C on Lihir Ore NTS200 .... 171 3+ 2+ + - Figure 95 Eh-pH diagram at 220°C. [Fe ] = 1 g/L, [Fe ] = 2 g/L, [K ] = 3 g/L, [HSO4 ] = 4.3 g/L ...... 173 3+ 2+ + - Figure 96 Eh-pH diagram at 220°C. [Fe ] = 1 g/L, [Fe ] = 2 g/L, [K ] = 3 g/L, [HSO4 ] = 9 g/L ...... 173 3+ 2+ + - Figure 97 Eh-pH diagram at 220°C. [Fe ] = 1g/L, [Fe ] = 2 g/L, [K ] = 3g/L, [HSO4 ] = 20g/L ...... 174 3+ 2+ + - Figure 98 Eh-pH diagram at 220°C. [Fe ] = 1 g/L, [Fe ] = 2 g/L, [K ] = 0 g/L, [HSO4 ] = 20 g/L ...... 175 Figure 99 Histograms of Lihir autoclave discharge free acid concentration from November to December 2015 ...... 176 Figure 100 XRD of Pyrite used for experiment ...... 189

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List of acronyms and abbreviations ACF – autoclave feed BFS – basic ferric sulphate CCD – counter current decantation DoF – degree of freedom DSTP – deep sea tailing placement DI – deionised EDS – energy dispersive x-ray spectroscopy HOMO – highest occupied molecular orbital ICP-OES – Inductively coupled plasma optical emission spectrometry LUMO – lowest unoccupied molecular orbital MLA – mineral liberation analysis MSE – mixed solvent electrolyte NCAF – neutralisation cyanidation adsorption feed NHE – normal hydrogen electrode ORP – oxidation reduction potential POX – pressure oxidation SEM – scanning electron microscopy SHE – standard hydrogen electrode US – united states XRD – x-ray powder diffraction QXRD – qualitative x-ray powder diffraction

List of shortenings

푎푖 – activity of species 𝑖

훾푖 – aqueous molar activity coefficient of species 𝑖

Eh – half cell electrode potential measured against standard hydrogen electrode E° – standard half cell electrode potential ∆E° – standard electrode potential

∆Grxn° – standard Gibbs energy of reaction

∆Gf° – standard Gibbs energy of formation

∆Hf° – standard enthalpy of formation

Keq – equilibrium constant 푆°– standard entropy

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Chapter 1: Project background and literature review

1.1 Gold Gold has been a coveted metal throughout history. Although it is not the rarest metal on earth, gold is hard to find, and it is normally present in ore at a low concentration. The ore grade in high-quality underground mine typically varies between 8 to 10 ppm, while low- quality underground mine has a gold concentration between 1 to 4 ppm. This drives the high production cost and ultimately become the main reason for its high value (World Gold Council, 2020).

Figure 1 shows the fluctuations of gold price from 1970 to June 2020. The first notable rise in gold price happened in 14 January 1980 when it hit a record high of USD 850 an ounce. This was due to the release of gold price from the United States (US) government control (History, 2010). The price of gold then started to soar high due to the 2008 financial crisis and kept rising until it hit another record high in 2011. A decade after the 2008 financial crisis, gold has once again been recognised as the “safe haven” during the recent Coronavirus (COVID-19) pandemic. In times of economic uncertainty like this, investors tend to retreat to safety and hold gold as a strategic asset due to its enduring value which in turn increases the price of gold. This shows the power of gold in stabilising international economy and world currency markets (Folger, 2020; McKay & Peters, 2017).

Figure 1 Gold prices (USD/oz) of the day in the last 50 years (Adapted from World Gold Council, 2020) 1

Apart from being used as investment, gold has been used as jewellery and decorative ornaments due to its lustrous colour and tarnish resistance property since ancient times. To date, it has also been widely utilised in various industrial fabrications, such as electronics, dental and medical, due to its extreme ductility, malleability and high electrical conductivity and resistance to corrosion (Marsden & House, 2006).

Gold mining is a growing global business. With the starting of new gold operations worldwide, such as Gruyere in Western Australia and Agnico Eagle’s Meliadine mine in Northern Canada, the gold global output is predicted to reach a record of 110 million ounces by the end of 2019 (Hosie, 2019). In 2018, China was the biggest gold producer, followed by Australia, Russia and the United States.

1.2 Refractory gold ore

With the declining grades of gold reserves, gold production from refractory deposits have become a major contributor to global gold production (Habashi, 2016). Ore is classified as refractory when finely ground samples yield less than 30% recovery from direct cyanidation. In general, the unresponsiveness of refractory ores to conventional cyanidation could be caused either by the inaccessibility of the gold within the ore or by the high reactivity of the gangue mineral within the ore. The different mechanisms known to date are summarised in Table 1 (Vaughan, 2004; Fraser et al., 1991; Swash, 1988; Cook and & Chryssoulis 1990; Mason, 1992; King et al., 2011).

Table 1 Causes of ore refractory behaviour

Cause of gold ore Mechanisms Description refractoriness Gold ‘locking’ in the Physical Gold is physically locked up within the mineral host minerals locking matrix (e.g. sulphide minerals, siliceous materials) with no practicable amount of grinding that are capable to expose the fine particles of gold. The most common example is colloidal gold where gold inclusion has a dimeter smaller than 200 Å.

Cause of gold ore Mechanisms Description refractoriness Chemical Presence of insoluble or poorly soluble gold locking minerals in alkaline cyanide solution. Some commonly found gold minerals includes Au-Ag tellurides, aurostibite, maldonite and auricupride. Presence of gold in solid solution. This means gold is atomically distributed in the crystal structure of the ore minerals. Presence of Leach- Gangue mineral consumes excessive amount of reactive gangue robbing ores cyanide and/or oxygen in side reactions. minerals Preg-robbing Gangue mineral re-adsorbs or precipitates the ores dissolved gold cyanide complex from the solution resulting in the loss of gold from leach liquor.

Of all these, physical encapsulation of colloidal gold in the host minerals is believed to be the most common cause for ore refractory behaviour (Chryssoulis & McMullen, 2005;

Yannopoulos, 1990). Pyrite (FeS2) is known as the preferential host mineral for gold in sulphide ores. With the high affinity of gold to arsenic, arsenic-containing pyrites, namely arsenopyrite and As-rich pyrite or more commonly known as arsenian pyrite have been acknowledged as excellent hosts for gold (Cook & Chryssoulis, 1990; Fleet et al. 1993; Fleet & Mumin 1997; Arehart et al., 1993; Mason, 1992). The highest reported concentration of gold carried in natural arsenian pyrite exceeds 1.1 wt% (Pals and Spry, 2003) while up to 1.52 wt% in natural arsenopyrite (Johan et al., 1989).

The recoverability of gold from refractory ore is challenging as ore pre-treatment prior to cyanidation is required (Marsden & House, 2006; La Brooy et al., 1994; Vaughan, 2004; Newcrest, 2013). There are a number of processing technologies developed for different types of refractory ore. For refractory sulphidic ores, the most common technologies are summarised in Table 2 (Lunt & Weeks, 2016; Fraser et al., 1991; Marsden & House, 2006; Adams, 2016; Baron et al., 2016; Miller et al., 2016).

This current project focuses on pressure oxidation, more specially on acidic pressure oxidation, and therefore other technologies will not be further discussed in this literature review. 3

Table 2 Common pre-treatment technologies for processing sulphidic refractory ore Industrial Working Principle Application examples Type Processes Roasting is considered to be the most effective oxidative pre- Giant Yellowknife (Canada) treatment for severe refractory gold ores such as double Kalgoorlie Consolidated—Gidji (Australia) refractory carbonaceous sulphidic ore. The process is carried out New Consort (South Africa) in the range of 450 to 820°C, and most typically below 600°C. Big Springs, Carlin, Cortez, and Jerritt Canyon (Nevada) Oxygen is introduced into the process to oxidise the Minahasa (Indonesia) carbonaceous matter and sulphides in which hematite and Roasting Tongguan Zongjin (China) various offgas are typically formed. These include, but not limited

to, carbon dioxide (CO2), carbon monoxide (CO), sulphur dioxide

(SO2) and nitrogen oxides (NOx). The formation of hematite is Oxidative beneficial for the downstream gold leaching as its porous characteristic allows for cyanide leach solution to penetrate and dissolve the gold.

Pressure Refer to Section 1.3 Refer to Table 3. oxidation

Biological Oxidation involves the action of bacteria as a catalyst Fairview (South Africa) for the conversion of iron sulphides in the ore to soluble ferric iron. Sao Bento (Brazil) Biological Some bacteria examples are Thiobacillus thio-oxidans, Wiluna and Youanmi (Australia) oxidation Sulfobacillus acidophilus and Sulfolobus acidocaldarius. The Ashanti Sansu (Ghana) Tonkin Springs (Nevada) 4

Industrial Working Principle Application examples Type Processes exact mechanism for this process remains debatable and Olympias (Greece) controversial. However, the generalised mechanism shows that, the bacteria play a role in two different ways. The first one involves their attachment to the mineral surface which enables them to oxidise the elemental sulphur that is on the mineral surface to form sulphate. Bacteria is not involved in the formation of this elemental sulphur on the mineral surface as it is the product between ferric and the sulphide mineral reaction. Secondly, the bacteria derive energy by oxidising ferrous to ferric in the bulk solution phase which then is used to create valuable oxidant to assist the action of bacteria mentioned earlier.

Ultrafine grinding has continually improved in terms of equipment Eleonore (Canada) design and efficiency. Some examples of available specialist Kumtor (Kyrgyztan) machines are Xstrata’s IsaMill®, Metso’s Vertimill, Outotec’s High KCGM (Australia) Cowal (Australia) Ultrafine Intensity Grinding (HIG) mill and the Metprotech mill. These Physical grinding machines typically fine grind in the range of 1µm to 10 µm by using a finer media size in the range of 2 to 3 mm. With a higher surface area generated, UFG can be used to pre-treat finely grained refractory ore.

1.3 Pressure oxidation Pressure oxidation, or commonly abbreviated to POX, has been used extensively for the pre-treatment of refractory gold ores over the past three decades with Homestake Mining Company’s McLaughlin being the first facility operated in 1985 (Thomas & Pearson, 2016; Fraser et al., 1992). This technology has been used successfully within the gold industry to process whole ore and sulphide concentrates using acidic route, while alkaline route is typically used exclusively for whole ore only. The process is carried out in a pressure vessel called an autoclave. Figure 2 shows an example of horizontal autoclave design (Thomas & Pearson, 2016)

Figure 2 Horizontal autoclave at Barrick Goldstrike Mine, USA (Thomas & Pearson, 2016)

During pressure oxidation, sulphide minerals which occlude fine gold are oxidised which effectively liberates the encapsulated gold particles. This process is typically carried out at elevated temperatures ranging from 180 to 240°C with oxygen partial pressure of 350 to 700 kPa (Rusanen et al., 2013; Papangelakis & Demopoulos, 1990; Ruonala et al., 2016; King et al., 2011). Table 3 shows some examples of current gold pressure oxidation operations worldwide (Alacer Gold, 2014, Thomas & Pearson, 2016, Habashi, 2014).

Table 3 Current Gold Pressure Oxidation Facilities Temperature Mine Operations Location Company (°C) Amursk Russia Polymetal International 205 Campbell Red Lake Canada Goldcorp 195 Copler Turkey Alacer Gold 220 Pueblo Viejo Dominican Republic Barrick 230 Goldstrike Nevada Barrick 225 Lihir Papua New Guinea Newcrest Mining 215 Kittila Finland Agnico-Eagle 210 Macraes New Zealand Oceana Gold 225 Mercur USA Barrick 240 Porgera Papua New Guinea Barrick 195 Sao Bento Brazil Anglo Ashanti 220 Twin Creeks Nevada Newmont 225

Although the objective of the alkaline and acidic pressure oxidation is identical, the chemistry inside the autoclave is very different which leads to different gold recovery and residence time. Alkaline pressure oxidation typically gives a lower gold recovery (up to 10%) due to the gold entrapment in the hematite which is formed at the oxidising surface of the sulphide mineral. Additionally, residence time in alkaline POX is generally longer due to difference in the chemistry inside the two processes. to the lower capital and operating costs. Despite of these disadvantages, alkaline POX is often preferred for its lower capital and operating cost due to a more conventional materials of construction. The performance of this technology may also surpass that of acidic POX with certain types of ore, such as high carbonaceous containing ore (Thomas & Pearson, 2016). As this thesis focuses on the acidic pressure oxidation, alkaline pressure oxidation will not be further discussed hereafter.

1.3.1 Pyrite Oxidation Chemistry

In acidic pressure oxidation, the three major reactions occurring inside the autoclave are illustrated in Figure 3. Note that this study focuses on pyrite (FeS2) as the main sulphide mineral in the ore and therefore the oxidation of other sulphide minerals will not be discussed in this review. The first reaction involves the aqueous oxidation reaction for pyrite by oxygen (1) and also by ferric ion (2) which will be discussed in Section 1.3.1. The second is the

7 oxidation of ferrous to ferric (3) due to the oxidising condition in the autoclave. The last one is the hydrolysis of ferric to form various forms of iron precipitate (4). Hematite (Fe2O3), basic + ferric sulphate (FeOHSO4) and jarosite compounds (MFe3(OH)6(SO4)2) where M can be K , + + + + + 2+ 2+ Na , H3O , Rb , Ag , Tl , Pb or Hg are the three possible iron hydrolysis products that could form at gold pressure operating temperature (Fleming, 2009; Demopoulos & Papangelakis, 1987; King et al., 2011; Fleuriault, 2016; Long & Dixon, 2004; Ruonala et al. 2016; Chen & Cabri, 1986). Further information on these iron phases will be discussed in 1.3.2 while factors that influence their formations will be discussed in 1.3.3.

Figure 3 General illustration of pyrite oxidation during pressure oxidation process

Despite of many years of study and diverse techniques utilised, aqueous pyrite oxidation mechanisms, with both ferric and oxygen as oxidants, no agreement on a comprehensive reaction pathway has yet been reached. Pyrite oxidation is a complex reaction which is electrochemical in nature involving multiple electron transfers. As there could only be one or at most two electrons transferred at a specific time and location, it has been suggested that this oxidation-reduction process occurs in multiple elementary steps (Luther, 1987; Basolo & Person, 1967; Rimstidt & Vaughan, 2003; Descostes et al., 2004).

There are variations in the reaction pathway proposed in literatures. These are due to the uncertainty in the form of iron released in the solution (Fe2+ or Fe3+), and also the form of sulphur species present at one time as sulphur goes through multiple oxidation states along the oxidation process.

Various intermediate sulphur species are suspected to be involved in the oxidation process. 0 2- These include, but not limited to, elemental sulphur (S ), sulphite (SO3 ), thiosulphate 2- 2- (S2O3 ) and polythionates (SnO6 , n = 4, 5, and 6) (Moses et al., 1987; Rimstidt & Vaughan, 2003; Descostes, 2004; Druschel & Borda, 2005; Druschel et al., 2003).

Using a molecular orbital theory approach i.e. HOMO/LUMO (highest occupied molecular orbital/ lowest unoccupied molecular orbital) arguments, Luther (1987) predicted that pyrite 2- oxidation would proceed through a thiosulphate (S2O3 ) intermediate as the first sulphur species released from the pyrite surface, regardless of the type of oxidant used for oxidation. This prediction is supported by research which investigated pyrite oxidation by ferric ion and oxygen as the oxidants (Goldharber, 1983; Moses et al., 1987; Druschel et al., 2003; Descostes et al., 2004; Descostes et al., 2001; Rimstidt & Vaughan, 2003).

Different possible pathways for sulphur species oxidation are proposed depending on the type of oxidant available, solution pH (Goldharber, 1983; Descostes, 2004; Moses et al., 1987; Smith & Hitchen, 1976) and temperature (Smith & Hitchen, 1976). Nevertheless, regardless of which pathway is undertaken, Moses et al (1987) and Descostes (2004) reported a lack of observed intermediate sulphur species during the experiment, and 2- sulphate (SO4 ) was the only sulphur species detected by ionic chromatography (IC) at the end of the experiment. This observation is attributed to the rapid conversion of thiosulphate to sulphate in acidic solution (pH under 3) either via decomposition or an oxidation reaction.

Further details on the different pathways for pyrite oxidation and thiosulphate conversion influenced by either the presence of oxygen and ferric will be discussed in Section 1.3.1.1 and Section 1.3.1.2 respectively.

1.3.1.1 Pyrite Oxidation by oxygen The most common pyrite oxidation mechanism presented is the oxidation of pyrite by molecular oxygen where iron passes into the solution in the ferrous (Fe2+) state while sulphur 2- 0 passes in the form of sulphate ion (SO4 ) and/or elemental sulphur (S ). These reactions are represented by Reactions 1.1 and 1.2 respectively. At typical industrial operating conditions of pressure oxidation (above 180°C), Reaction 1.1 is believed to be the dominant reaction due to the poor sulphur stability at high temperature and high oxygen pressure (Demopoulos & Papangelakis, 1987; Papangelakis & Demopoulos, 1990; Papangelakis & Demopoulos, 1991, Ruonala et al., 2016; Thomas, 2015; Lowson, 1982; Habashi & Bauer,1966; Bailey & Peters, 1976).

FeS 2 (s) + 7/2O2 (g) + H2O (l)→ FeSO4 (a) + H2SO4 (a) (1.1)

0 FeS 2 (s) + O2 (g) → FeSO4 (a) + S (s) (1.2)

These overall reactions are the sums of multiple elementary steps where the sulphur in pyrite is oxidised though two or more intermediate sulphur species until sulphate is formed (Reaction 1.4 to Reaction 1.7).

2- As mentioned in 1.3.1, thiosulphate (S2O3 ) was identified to be the possible first intermediate sulphur species generated from pyrite oxidation. As a result, the first elementary aqueous pyrite oxidation reaction could be expressed in Reaction 1.3 where a transfer of six electrons from anodic to cathodic sites was involved as shown in Reactions 1.3a and 1.3b.

2+ 2- Overall: FeS2 (s) + 1.5O2 (g) → Fe (a) + S2O3 (a) (1.3)

2+ 2- + - Anode: FeS2 (s) + 3H2O (l) → Fe (a) + S2O3 (a) + 6H (a) + 6e (1.3a)

+ - Cathode: 3/2O2 (g) + 6H (a) + 6e → 3H2O (l) (1.3b)

In the absence of ferric, thiosulphate undergoes decomposition due to its instability in highly acidic solution. The decomposition products seem to be dependent on acidity and temperature. In the early studies, Davis (1958) predicted that thiosulphate decomposes to elemental sulphur and sulphite based on disproportionation reaction as shown in Reaction 1.4. This pathway was later adopted by Long & Dixon (2004).

2- 0 2- S2O3 (a) → S (s) + SO3 (a) (1.4)

2- However, Descostes et al. (2004) proposed that instead of sulphite, tetrathionate (S4O6 ) is the direct thiosulphate decomposition product which is expressed according to Reaction 1.5. The elemental sulphur (S0) produced in Reaction 1.4 and Reaction 1.5 will be oxidised possibly through one or more thionate intermediates (Long & Dixon, 2004).

2- + 0 2- 5S2O3 (a) + 6H (a) → 2S (s) + 2S4O6 (a) + 3H2O (l) (1.5)

These two decomposition reactions are in agreement with the data presented by Smith & Hitchen (1976) who detected tetrationate at high acidity conditions while sulphite was detected only at low acidity. Therefore, tetrationate seems to be the favoured species generated at high acidity while the formation of sulphite seems to be favoured at low acidity.

Under oxidising condition, both tetrationate and sulphite will be further oxidised to sulphate 2- (SO4 ). The oxidation of tetrationate to sulphate is likely to occur through sulphite formation as expressed in Reaction 1.6 and Reaction 1.7 consecutively (Smith & Hitchen, 1976).

2- 2- + 2S4O6 (a) + 7O2 (g) + 3H2O (l) → 4SO3 (a) + 6H (a) (1.6)

2- 2- 2SO3 (a) + O2 (g) → 2SO4 (a) (1.7)

The other pyrite oxidation mechanism by molecular oxygen presented in literatures involves the direct release of iron into the solution as ferric (Fe3+) instead of ferrous. Earlier electrochemical studies (Biegler & Swift, 1979; Meyers, 1979; Peters & Majima; 1969) and a more recent pressure oxidation study (Long & Dixon, 2004) suggested ferric ion is directly released into the solution from pyrite. Considering the formation of thiosulphate as the first sulphur immediate, the first pathway for the proposed pyrite oxidation can be expressed according to Reaction 1.8 which involves Reaction 1.8a in the anodic site and oxygen evolution (Reaction 1.8b) in the cathodic site.

+ 3+ 2- Overall: 4FeS2(s) + 7O2(g) + 4H (a) → 4Fe (a) + 4S2O3 (a) + 2H2O (a) (1.8)

3+ 2- + - Anode: FeS2(s) + 3H2O(l) → Fe (a) + S2O3 (a) + 6H (a) + 7e (1.8a) + - Cathode: 3/2O2 (g) + 6H (a) + 6e → 3H2O (l) (1.8b)

This mechanism was proposed due to the absence of ferrous in the electrochemical studies where colometric-phenanthrolene method was used to quantify the concentration of ferric and ferrous. In the pressure oxidation experiment (Long & Dixon, 2004), the direct 3+ generation of ferric was deduced due to the decrease in [Fe /FeT] ratio at the beginning, which then increased with time. The drop was attributed to local build-up of ferrous near pyrite surface which was proposed to be generated from the reduction of ferric by thiosulphate (refer to Reaction 1.10 in Section 1.3.1.2). However, this drop could also be explained by other causes, such as generation of ferrous from pyrite oxidation by ferric ion reduction which will be discussed in Section 1.3.1.2. Therefore, the results from Long & Dixon (2004) study is deemed to be insufficient in confirming the direct generation of ferric.

To date, this dispute about the pyrite oxidation reaction has not been resolved. However, in the most recent study by Santos et al (2016), who used electronic structure calculations approach, proposed that both pathways occur in parallel. They concluded that the ferric- producing pathway (referred as Type I) occurs initially and it involves the consumption of water adsorbed on the pyrite surface which leads to the oxidation of iron sites to form Fe(III)- OH- groups. Following this, after the transfer of a radical hydrogen to the previously generated Fe(III)-OH- groups, ferrous and sulphate species are released into the solution via the formation of S-OH and Fe-OH2 groups on the pyrite surface (referred as Type 2).

Despite this recent prediction, there has not been any definite evidence that can support any pathway proposed. Therefore, both interpretations are still considered plausible, with a possibility that they are occurring in parallel.

1.3.1.2 Pyrite oxidation by ferric ion The importance of ferric iron as an effective pyrite oxidant was not recognised until early investigations by Moses et al. (1987). To date, there have been several studies done to investigate the role of ferric ion as pyrite oxidant at lower temperatures (below 90°C) (Moses et al., 1987; Singer & Stumm, 1970; Garrels & Thompson, 1960; Holmes & Crundwell, 2000; Goldhaber, 1983). These studies concluded that the rate of pyrite dissolution with ferric ion is much faster than that with oxygen. This indicates that ferric is the main direct pyrite oxidant at these low temperature conditions. These results are in agreement with the prediction by Luther (1987) who suggested that the cause of the difference lies in the different binding mechanism where ferric could readily bind to the pyrite surface chemically whereas oxygen could only associate with the pyrite surface by adsorption.

2- The oxidation of pyrite by ferric ion to form thiosulphate (S2O3 ) species as the first sulphur intermediate can be expressed as in Reaction 1.9. This reaction involves the transfer of six electrons between the anodic and cathodic sites which can be expressed in Reactions 1.9a and 9b respectively.

3+ 2+ 2- + Overall: FeS2(s) + 6Fe (a) + 3H2O (l) → 7 Fe (a) + S2O3 (a)+ 6H (a) (1.9)

2+ 2- + - Anode: FeS2(s) + 3H2O(l) → Fe (a) + S2O3 (a) + 6H (a) + 6e (1.9a)

Cathode: 6Fe3+(a) + 6e- → 6Fe2+(a) (1.9b)

In the presence of excess ferric in acidic solution, the decomposition reaction of thiosulphate (Reaction 1.6) is competing with thiosulphate oxidation reaction by ferric. Williamson & Rimstidt (1993) found that under acidic conditions (below pH ~1.8 at 25°C), the rate of thiosulphate oxidation by ferric ion (represented by solid line) is faster than the rate of thiosulphate acidic decomposition (represented by dashed line) as illustrated in Figure 4. The sharp decrease of thiosulphate oxidation by ferric is likely to be correlated to the decrease of ferric solubility as pH increases. This means thiosulphate has a tendency to undergo oxidation in the presence of excess ferric under acidic condition.

Figure 4 Rate of thiosulphate oxidation by ferric vs. acid decomposition as a function of pH (adapted from Williamson & Rimstidt, 1993)

o As ferric is known to be a strong oxidant (E Fe3+/Fe2+ = 0.77 V vs. standard hydrogen electrode (SHE)), thiosulphate would likely be completely oxidised to sulphate according to Reaction 1.10 via the formation of various sulphur intermediate. This is supported by Moses et al. (1987) who did not observe intermediate sulphoxy anions in acidic, ferric saturated, anaerobic experiments. Further, Luther (1987) claimed that this oxidation reaction is rapid enough such that thiosulphate would not be able to co-exist with Fe3+.

2- 3+ 2+ + 2- S2O3 (a) + 8Fe (a) + 5H2O (l) → 8Fe (a) + 10H (a) + 2SO4 (a) (1.10)

Similar to thiosulphate acid decomposition, the first sulphur intermediate species generated 2- from thiosulphate oxidation by ferric is proposed to be tetrathionate (S4O6 ) where its formation reaction could be expressed in Reaction 1.11 (Druschel et al. 2003; Williamson & Rimstidt, 1993; Druschel & Borda, 2005).

2- 3+ 2+ 2- 2S2O3 (a) + 2Fe (a) → 2Fe (a) + S4O6 (a) (1.11)

However, the further tetrathionate oxidation to sulphate appears to be more complicated in the presence of ferric. Figure 5 shows the possible oxidation pathways of tetrationate to sulphate ions (Druschel et al., 2003).

Figure 5 Tetrathionate possible oxidation paths (Druschel et al., 2003)

2- With sulphate (SO4 ) being the final sulphur product, the overall reaction of pyrite oxidation by ferric then can be expressed according to Reaction 1.12.

3+ 2+ 2- + FeS2 (s) + 14Fe (a)+ 8H2O (l) → 15Fe (a) + 2SO4 (a) + 16H (a) (1.12)

1.3.1.3 Pyrite oxidation in gold pressure oxidation During continuous pressure oxidation operation, high purity oxygen gas is sparged into the autoclave at all times while ferric and ferrous are continuously present in the background. This means the oxidation reactions taking place in the autoclave would likely to proceed via a combination of mechanisms discussed in Section 1.3.1.1 and Section 1.3.1.2.

One possible mechanism is that pyrite oxidation reaction largely proceeds via oxidation by ferric although oxygen is the overall oxidant in the system. As discussed in Section 1.3.1.2, several studies (Moses et al., 1987; Singer & Stumm, 1970; Garrels & Thompson, 1960; Holmes & Crundwell, 2000; Goldhaber, 1983), that investigated the efficacy of ferric as pyrite oxidant at oxidant at low temperature system, have concluded that the kinetic of pyrite oxidation by ferric is faster than the oxidation by oxygen. Additionally, the mass transfer of ferric to the pyrite surface is likely to be higher due to its higher concentration (in the form of ferrous and ferric) in the solution compared to oxygen whose solubility is relatively low at higher temperature. With these two reasons, majority of the pyrite oxidation reaction could possibly go through ferric as a surrogate oxidant with oxygen as the overall oxidant for the system. Although the proportion of pyrite oxidation by direct oxygen could be relatively small, the presence of oxygen is deemed to be necessary to regenerate ferrous to ferric following Reaction 1.13 (Demopoulos & Papangelakis, 1987; Long & Dixon 2004; Ruonala et al., 2016; Thomas, 2015; Bayley & Peters, 1976) and to initiate the pyrite oxidation reaction when the initial ferric concentration is relatively low. 14

2+ + 3+ Overall: 4Fe (a) + O2 (g) + 4H (a) → 4Fe (a) + 2H2O (a) (1.13)

Anode: 4Fe2+ (a) → 4Fe3+ (a) + 4e- (1.13a)

+ - Cathode: O2 (g) + 4H (a) + 4e → 2H2O (l) (1.13b)

However, the role of ferric in catalysing the kinetic of pyrite oxidation at pressure oxidation conditions remains debatable (Papangelakis & Demopoulus, 1991) with contradicting result have been reported in literatures (McKay & Halpern, 1958, Gerlach et al., 1966, Warren, 1958, King & Lewis, 1980). This means the role of ferric as pyrite surrogate oxidant at high temperatures maybe insignificant at high temperatures. McKay & Halpern (1958) suggested that the oxidation of ferrous to ferric (Reaction 1.13) is slower than the direct oxidation of pyrite by oxygen although pyrite oxidation by ferric may be relatively fast. This could mean the catalysing effect of ferric on the kinetic of pyrite oxidation at high temperatures largely depends on the initial concentration of ferric. This could possibly explain the contradicting results reported in Gerlach et al. (1966) and King & Lewis (1980). This speculation, however, cannot be confirmed as there is currently no study that has been done to investigate the role of ferric as pyrite oxidation at high temperature where ferric undergoes hydrolysis.

1.3.2 Ferric hydrolysis The hydrolysis of ferric generates a large amount of acid which leads to an acidic environment inside the autoclave. Reactions 1.14, Reaction 1.15 and Reaction 1.16 are the hydrolysis reactions for hematite, jarosites and basic ferric sulphate consecutively. The crystal structure and physical properties of these solids are summarised in Section 1.3.2.1 to 1.3.2.3.

Fe2(SO4)3(a) + 3H2O(l) → Fe2O3(s) + 3H2SO4(a) (1.14)

3Fe2(SO4)3(a) + 12H2O(l) + M2SO4(a) → 2MFe3(SO4)2(OH)6(s) + 6H2SO4 (1.15)

Fe2(SO4)3(a) + 2H2O(l) → 2FeOHSO4(s) + H2SO4(a) (1.16)

1.3.2.1 Hematite

Hematite (α-Fe2O3) is one of the two polymorphs of Fe2O3 found in nature and was originally named “Aematitis lithos” for “blood stone”. The other polymorph of Fe2O3 is maghemite (γ-

Fe2O3) which is always metastable with respect to hematite (Mindat, 2019a; Chen & Cabri, 1986). Hematite has an extremely variable appearance where its luster can range from earthy to submetallic to metallic. Hematite precipitated by iron hydrolysis in hydrometallurgy

15 processes is earthy and therefore its colour is reddish brown (rust-red), ranging from dull to bright. Figure 6 illustrates the typical crystal structure of hematite (Mindat, 2019a).

Figure 6 Typical hematite crystals (Mindat, 2019)

1.3.2.2 Jarosites The jarosite group of minerals is part of the alunite super group and is characterised with a general formula AFe3(SO4)2(OH)6, where A can be substituted by various mono- or divalent + + + + + + 2+ 2+ cations such as K , Na , H3O , Rb , Ag , Tl , Pb and Hg . There are more than one hundred fifty jarosite compounds known to exist but only six are known to occur in nature. Table 4 summarises the six jarosites in the order of stability from high to low (Chen & Cabri, 1986; Das et al., 1996).

Table 4 The most common jarosite compounds Alkali Class General Mineralogy Common Name cations

+ Ag AgFe3(SO4)2(OH)6 Silver Jarosite Argentojarosite + K KFe3(SO4)2(OH)6 Potassium Jarosite Jarosite + NH4 NH4Fe3(SO4)2(OH)6 Ammonium Jarosite Ammoniojarosite

+ Na NaFe3(SO4)2(OH)6 Sodium Jarosite Natrojarosite 2+ Pb Pb1/2Fe3(SO4)2(OH)6 Lead Jarosite Plumbojarosite + H3O H3OFe3(SO4)2(OH)6 Hydronium Jarosite Carphosiderite

Since this research is tailored to focus on the Newcrest Lihir operation, potassium jarosite is the only jarosites’ member of interest and therefore no other members of this group will be discussed hereafter.

The formation of jarosites is favoured in the presence of high concentration of the associated alkali cations (Brown, 1971; Das et al., 1991; Stroffregen, 1993). Figure 7 shows that the stability of potassium jarosite increases at higher activity of K+ i.e. higher potassium concentrations, and the minimum jarosite stability limit corresponds to the saturation value

16 of the potassium sulphate (Stoffregen, 1993). This means there is a stability region for hematite and basic ferric sulphate in the presence of potassium although they are narrow (Stroffregen, 1993; Tozawa & Sasaki,1986).

Figure 7 Schematic illustration of hematite – basic ferric sulphate - jarosite stability. Dashed 2 3 lines represent contours of log a Fe3+ + a SO42- and increase to the left (Stroffregen, 1993)

The colour of potassium jarosite is known to vary from amber-yellow, yellow-brown, to brown or light yellow. Figure 8 shows that the structure of jarosite consists of sheets of hydroxyl- and sulphate bridged Fe3+ distorted octahedra (Das et al., 1996). The Fe-octahedra is 2- surrounded by four hydroxyl-groups and two oxygen atoms belonging to two different SO4 groups. Each octahedra group links with its neighbour by 4 hydroxyl groups and three 2- - 2 oxygen of each SO4 group bonded with three Fe(OH)4 units with one oxygen in the SO4 - being free which is doubly bonded to sulphur. The cation (e.g. K+) is located in octahedron formed by the oxygen and hydroxyl group.

Figure 8 Structure of KFe3(SO4)2(OH)6. (a) Single unit, (b) View along a and b axes, (c) View along c- and a-axis (Das et al., 1996)

1.3.2.3 Basic ferric sulphate

There is relatively limited information available about basic ferric sulphate (BFS, FeOHSO4). Its crystal structure was first determined by Johansson (1962) which was then followed by Gomez et al. (2013). Figure 9 shows that the structure of the basic ferric sulphate consists of Fe-octahedra which is surrounded by two water molecules, two hydroxyl groups and two oxygen atoms belonging to two different tetrahedra sulphate group. The Fe-octahedra is forming an infinite chain along the a-axis by sharing two opposite hydroxyl group with the neighboring octahedra. The two adjacent octahedra are linked by sulphate-tetrahedra which alternates along the chain (Gomez et al., 2013; Johansson, 1962).

Figure 9 Projection of the structure of FeOHSO4 along the a-axis (left) and b-axis (right) (Johansson, 1962)

Using FEG-TEM JEOL** 2100F, Gomez et al. (2013) captured the image of BFS crystal and they observed an elongated directional crystal growth as shown in Figure 10.

Figure 10 TEM image of basic ferric sulphate with magnification of 41000X (Gomez et al., 2013) 18

1.3.3 Factors affecting iron hydrolysis product The predominance of each of these solids depends on the prevailing conditions in the autoclave. The effect of acidity, temperature, initial ferric concentration, presence of divalent metal sulphates, retention time and seeding have been studied (Cheng & Demopoulos, 2004; Tozawa & Sasaki, 1986; Fleuriault, 2016; Umetsu et al., 1977) and will be discussed in the following sub-sections.

1.3.3.1 Effect of acidity Numerous studies have shown that acidity has a great effect on the phase of iron hydrolysis precipitates (Tozawa & Sasaki, 1986; Fleming, 2009; Demopoulos & Papangelakis, 1987; King et al., 2011; Fleuriault, 2016; Umetsu et al., 1977). The precipitation of hematite (Reaction 14) is favoured at low acidity, while the formation of basic ferric sulphate (Reaction 15) and jarosites (Reaction 16) formations are favoured at high acidity.

In the absence of monovalent cations, Tozawa and Sasaki (1986) investigated the relationship between free sulphuric acid concentration and sulphur content in the iron hydrolysis product at three different temperatures. They observed the transition of hydrolysis products from hematite to basic ferric sulphate as acidity was increased during the experiment (refer to Figure 11).

Basic Ferric Sulphate Precipitation

Hematite Precipitation

Figure 11 Relationship between sulphur content in hydrolysis product and free sulphuric acid concentration in Fe-S-H2O system at 170°C, 185°C and 200°C (Tozawa & Sasaki, 1986)

The low-sulphur content region represents hematite precipitation, while the high-sulphur content region represents basic ferric sulphate precipitation. The free sulphuric acid concentration at which this transition occur corresponds to the upper limit of free sulphuric 19 acid at which hematite is stable (Tozawa & Sasaki, 1986; Umetsu et al, 1977; McDonald & Robinson, 2016). Free sulphuric acid concentration was determined from the following stoichiometric relation (Tozawa & Sasaki, 1986; Umetsu et al, 1977; McDonald & Robinson, 2016).

2- 3+ 2+ [H2SO4] = [SO4 ]total – 1.5[Fe ] – [Zn ] Equation 1

In the presence of a monovalent cation, Stroffregen (1993) modelled the stability field of potassium jarosite at 200°C and 100 bar which was constructed with log m∑S of -0.5 and log mK of -0.5 as shown in Figure 12. This shows that the stability field of potassium jarosite extends to pH 1.2 at 200°C, before hematite becomes the predominant phase at lower acidity.

Figure 12 Jarosite stability field at 200°C and 100 bar with log m ∑S = -0.5 and log m K = -1.5 (Stroffregen, 1993)

1.3.3.2 Effect of Temperature Temperature also has a significant effect on the stability of different phases of the iron hydrolysis products. Table 5 shows that the upper limit of free sulphuric acid for hematite precipitation is higher as temperature increases. This implies the stability of hematite increases, while that of basic ferric sulphate decreases, at higher temperature (Rubisov & Papangelakis, 2000; Fleming, 2009; Stroffregen, 1993; Tozawa & Sasaki, 1986, Demopoulos & Papangelakis, 1987; Umetsu et al., 1977).

Table 5 The effect of temperature on free sulphuric acid concentration upper limit for hematite precipitation

Temperature (°C) Upper limit of sulphuric acid Reference Note1 (g-H2SO4/L) 53 1 170 55 3 185 56 1, 2 64 1 200 70 3 1 Tozawa & Sasaki, 1986 2 Umetsu et al., 1977 3 Demopoulos & Papangelakis, 1987

Note1 Based on room temperature measurement

Similar to basic ferric sulphate, the stability field of potassium jarosite also narrows down and shift towards lower pH with the increase of temperature (Stroffregen, 1993; Das et al., 1996; Umetsu et al., 1977; Babcan, 1971) as shown in Figure 13. This means the formation of jarosite takes place in higher acidic region as the temperature increases. For example,

Stroffregen (2006) suggested that jarosite is stable, relative to hematite, below H2SO4 concentration of 0.45 m at 250°C, but the sulphuric acid concentration required for jarosite stability decreases to 0.26 m at 200°C.

3+ + 2- Figure 13 Ferric stability diagram for the Fe -K -SO4 -H2O system as a function of temperature and pH (adapted from Babcan, 1971) 21

Das et al (1986) and Dutrizac & Jambor (2000) also reported that temperature has a great effect on the kinetics of the jarosite precipitation. The period of jarosite precipitation decreases significantly above 100°C (Dutrizac & Jambor, 2000) where it was found to decrease from up to 6 months at 25°C to two hours at 210°C (Das et al., 1986).

1.3.3.3 Effect of Iron Concentration Initial iron concentration also affects the stability of hematite. Umetsu et al. (1977) found that hematite is stable at low initial ferric concentration, while basic ferric sulphate is favoured at higher concentration. 14.9 g/L Fe3+ was found to be the highest concentration before basic ferric sulphate starts precipitating, which is indicated by the increase of sulphur content in the precipitate as shown in Table 6.

Table 6 The effect of initial iron concentration on sulphur content in the precipitates (Umetsu et al., 1977)

Initial concentration of Sulphur in iron (g/L) precipitate (%) 5.5 1.6 12.2 1.4 14.9 1.3 20.7 5.4 24.9 15.2

The effect of initial iron concentration is closely related to the effect of pH. Higher initial iron concentration leads to higher concentration of free sulphuric acid due to its generation from ferric sulphate hydrolysis according to Reaction 14 (Tozawa & Sasaki, 1986).

In the presence of excess alkali ion, jarosite generally precipitates from a solution containing 0.001M to up to 3M of ferric. In this range of ferric concentration, the composition of jarosite is independent of the iron concentration (Das et al., 1996). Umetsu et al. (1977) and Brown (1971) suggested that as the initial iron concentration (iron activity) is higher, the stability field of potassium jarosite expands towards lower pH and Eh as shown in Figure 14.

Figure 14 The effect of ferric concentration Potassium jarosite stability field in Fe-K-SO4- + 2- H2O system at 95 °C, [K ] = 0.01M and [SO4 ] = 1M (Umetsu et al., 1977)

1.3.3.4 Presence of Divalent Metal Sulphates

The presence of divalent metal sulphates, such as ZnSO4, CuSO4, MgSO4 and MnSO4 has been suggested to favour the precipitation of hematite. The addition of these metal sulphates shifts the upper limit of acidity for hematite stability region to a higher value as shown in Table 7 (McDonald & Robinson, 2016; Tozawa & Sasaki, 1986; Umetsu et al., 1977).

Table 7 The effect of divalent metal sulphates presence to free sulphuric acid concentration Metal Metal (M2+) Upper limit of free Reference Sulphates Concentration (g/L) sulphuric acid (g/L) Note 1 14-16 Approx. 82 3 15 76 2 MgSO4 32 100 2 27-33 Approx. 105 3 28 68 2 46 74 2 50 80 1 ZnSO4 68 83 2 80 100 1 101 92 2 *1 Umetsu et al.,1977 *2 Tozawa & Sasaki, 1986 *3 McDonald & Robinson, 2016 Note1 Based on room temperature measurement

Studies speculated that these divalent metal sulphates buffer the excess acid in the solution by either promoting the formation of bisulphate ions (Tozawa & Sasaki, 1986) or forming metal bisulphate complex (Cheng & Demopoulos, 2004; King et al., 2011; McDonald & 23

Robinson, 2016). This buffering action decreases the activity of hydrogen ion (H+) in the solution which increases the “at temperature” pH and consequently favours the formation of hematite.

When potassium jarosite is the favoured precipitate, the presence of divalent metal sulphate could lead to the substitution of these metals into the jarosite structure. Dutrizac (2008) observed incorporations of small amount (< 3%) of copper, zinc and lead into the structure potassium jarosite. The extent of the substitution was found to increase with the increase of divalent metal ions concentrations.

1.3.3.5 Seeding Dutrizac & Chen (2001) investigated the effect of hematite seeding over a wide range of temperature and concentration. At temperatures above 160°C, increasing the amount of hematite seed up to 40 g/L was found to decrease the amount product yield. Little effect was observed with further addition of hematite as shown in Figure 15. However, hematite seeding was found to be a critical factor in favouring the precipitation of hematite over other iron sulphate hydroxide species. For example, hematite was able to be precipitated at higher free sulphuric acid concentration (up to 0.2M H2SO4) and at higher initial ferric concentration (up to 0.5M Fe3+) at 225°C before basic ferric sulphate started to form.

Figure 15 The effect of hematite seed addition on product yield and composition of the product precipitated from 0.5M Fe(SO4)1.5 solution at pH 1.4 at 225°C with 2-hour retention time (Dutrizac & Chen, 2011)

The effect of seeding on jarosite precipitation in sulphate solution at temperatures above 160°C could not be found to date. However, Dutrizac (1999) investigated the effect of jarosite seeding on the precipitation of Na-jarosite at 98°C. This study reported that the presence of jarosite seed would promote jarosite precipitation where an increase in seed addition was found to increase the rate of jarosite precipitation, almost linearly. The effectiveness of four different jarosite seed types, Na, K, Pb and Ag- jarosite, on the rate of Na-jarosite precipitation were found to be similar. Figure 16 shows the iron precipitation curves for Na-jarosite precipitation in the presence of various concentration of potassium jarosite seed at 98°C.

Figure 16 Iron precipitation curves for Na-Jarosite precipitation in the presence of K- jarosite seed at 98°C (Dutrizac, 1999)

1.3.4 Ferrous solubility Similar to ferric sulphate, the solubility of ferrous sulphate also decreases notably at high temperature (Cheng, 2002; Hasegawa et al., 1998; Kobylin et al., 2011). At pressure oxidation operating temperatures, ferrous sulphate would precipitate as ferrous sulphate monohydrate (FeSO4.H2O), or also known as szomolnokite as illustrated in Figure 17. Crystallisation of this solid has been reported in autoclave test from 140 to 200°C (Elgersma et al., 1993; Cheng, 2002; Hasegawa et al., 1998).

Figure 17 Ferrous sulphate solubility as a function of temperature (adapted from Cheng, 2002 and Hasegawa et al., 1998)

Apart from temperature, Hasegawa et al. (1998) also indicated that the co-existence of other sulphate electrolyte in the aqueous solution also affects the solubility of ferrous sulphate. The presence of sulphuric acid was found to increase the solubility of ferrous sulphate linearly in the temperature range of 160 to 220°C. Table 8 summarises the ferrous sulphate solubility at 220°C in pure water and sulphuric acid.

Table 8 Ferrous sulphate solubility as a function of sulphuric acid concentration at 220 °C (Hasegawa et al., 1998)

Sulphuric Acid concentration, g/L Fe2+ solubility, g/L 0 1.25 12.0 6.70 28.2 13.5 50.3 22.1

Although this solid is favoured to form at high temperature, it was found to readily dissolve during cooling where Elgersma et al. (1993) observed rounding of crystal edges and rough surfaces of the crystals as shown in Figure 18.

Figure 18 SEM image of FeSO4.H2O crystals with magnification of 1,500X (Elgersma et al.,1993)

1.4 Industry context This project focuses the pressure oxidation circuit at the Lihir gold processing operation, which is located on Niolam Island in the New Ireland Province of Papua New Guinea (PNG). The plant was acquired by Newcrest Mining for $9.5 billion in August 2010 and is considered as the company’s second most important asset (Ker, P 2015; Fitzgerald, 2016).

The Lihir gold deposit is known to be one of the largest gold deposits in the world with an estimated reserved of 188 Mt grading 2.48 g/t-gold. The gold is hosted in potassium-rich sulphide ore body, which is deposited in volcanic, intrusive and breccias within an extinct volcanic crater known as Luise Caldera. Majority of the sulphides are in the form of pyrite, with minor to trace level of other base metal sulphides and sulfosalts, such as chalcopyrite. The major gangue minerals are potassium feldspar, phyllosilicates such as biotite and illite, and quartz (Newcrest, 2013). Table 9 tabulates the mineralogy of the ore samples received from Lihir analysed using Mineral Liberation Analyser (MLA) technique.

Table 9 Mineralogy of Ore Samples received in 2016 (ALS Mineralogy, 2016)

Mineral Composition, wt% Pyrite 15 Arsenopyrite 0.05 Chalcopyrite 0.04 FeTi Oxide 0.73 Carbonates 1.62

Mineral Composition, wt% Sulphate-Phosphates 0.91 K-Feldspar 64.95 Phyllosilicates 9.10 Other silicates 7.57

This Lihir gold-bearing sulphide ore is classified as highly refractory which is caused by the presence of gold as solid-solution gold within the grains of pyrite. This type of gold is often referred as invisible gold. Invisible gold cannot be detected using conventional microscopy technique (Chryssoulis & McMullen, 2005; Arehart et al., 1993; Yang et al. 1998; Cook & Chryssoulis 1990; Cabri et al., 2000). As an alternative, Lihir mineralogist use arsenic as the pathfinder element for gold because arsenic association with gold has been established using secondary ion mass spectroscopy (SIMS) in the early mineralogy study or Lihir ore as shown in Figure 19 (Ketcham et al. 1993). Therefore, they use these arsenic-containing pyrites to identify the location of majority of the gold at Lihir.

Figure 19 Secondary Ion Mass Spectrometry result of Lihir ore (Ketcham et al., 1993)

Of all the refractory ore pre-treatment technologies available, pressure oxidation was selected based on the result from the oxidation testwork campaigns conducted during the plant feasibility study (Newcrest, 2013). With the inclusion of pressure oxidation into the Lihir process route, the general flowsheet of Lihir processing plant is illustrated in Figure 20. This project focuses on the pressure oxidation operation and the downstream process, up to the neutralisation stage.

Figure 20 Lihir Processing Plant Block Flow Diagram (Newcrest, 2013)

1.4.1 Lihir pressure oxidation process overview There are four autoclaves running in parallel on site, where the first three autoclaves are identical while the fourth one is 2.2 times the capacity of the other three autoclaves. The autoclaves are operating at 215°C with a total pressure of 2650 kPa. During pressure oxidation, sulphide minerals which occluded the invisible gold are oxidised to liberate the encapsulated gold particles with the injection of high purity oxygen (98% purity) which is produced onsite. The residence time in the autoclave is approximately 2 hours but may vary depending on the feed flowrate (Newcrest, 2013).

Upon exiting the autoclave, the oxidised slurry splits into two trains of two stage counter- current decantation (CCD) stage. Where majority of the acid and soluble salts were removed via the CCD overflow. This overflow stream is sent to a common disposal system where it will be combined with other tailings streams and eventually is discharged through a de- aeration tank to the ocean via a pipeline outfall at a depth of 128 meter below sea level. This tailings disposal method is also known as deep sea tailing placement (DSTP) which was chosen as the most sustainable method considering the specific conditions at Lihir Island. Continuous monitoring is exercised regularly to ensure no contamination of any metals concentration in the food chain nor in fish species (Newcrest, 2013).

The CCD underflow is pumped to a neutralisation tank, in which milk-of-lime is added. Lime is added to neutralise the residual acid and raise the slurry pH to 10, which is the alkalinity required for the subsequent cyanidation to avoid the formation of poisonous HCN gas. As the pH increases, some solubilised metals will precipitate out of the slurry. The neutralisation tank overflows to the leach tank for cyanidation process (Newcrest, 2013).

1.4.2 Lihir partial pressure oxidation operation In the past, Lihir run their autoclaves by specifying high degree of sulphide oxidation extent, typically above 95%, which is known as total pressure oxidation. This is common practice to ensure satisfactory chemical liberation for effective gold recovery to be achieved.

In December 2014, however, Lihir changed their autoclave operating philosophy to run on partial pressure oxidation due to the mineralogical distribution of gold throughout the ore. As shown in Figure 21, the MLA result of the Lihir ore indicates that majority of the pyrite present in the ore are the low-to-none arsenic containing pyrite (denoted by Pyrite) - from herein this pyrite will be referred as barren pyrite. More importantly, the gold-containing pyrite, i.e. arsenopyrite and arsenian pyrite (denoted by Pyrite_As), are mainly located on the outer edge of the ore. These arsenic-containing pyrite grains have higher reactivity which makes them available to be preferentially oxidised compared to barren pyrite. Therefore, there is an opportunity for Lihir to lower their oxidation extent and selectively oxidise the gold- containing pyrite.

+106 µm +38 µm ± 6 µm 2

Figure 21 Mineral liberation analyser (MLA) for three different particle sizes of Lihir Ore

Partial POX at Lihir autoclave is achieved by varying the feed rate into the autoclave while keeping the same amount of oxygen injected into the vessel. As autoclave feed is increased, the slurry residence time in the vessel inherently decreases and consequently results in the drop of pyrite oxidation extent which could go down to as low as 30%.

By running the autoclave on partial POX, the condition in the autoclave could be relatively oxygen starved and consequently the extent of ferrous oxidation to ferric ion (denoted by 2 in Figure 3) is expected to decrease. This would shift the ferric and ferrous proportion in the 30 autoclave. Based on Lihir plant data, the Fe3+:Fe2+ ratio was found to decrease at lower oxidation extent as expected. However, eventhough the Fe3+:Fe2+ ratio decreases at lower oxidation extent, the residual ferric concentration in the autoclave discharge was found to exhibit a tendency to increase. This is possibly due to the reduced extent of iron precipitation in the autoclave and causing the residual ferric in the solution to increase.

Additionally, with the increase of ferrous proportion relative to ferric in the autoclave, a local build-up of ferrous iron around the pyrite and/or precipitation of ferrous sulphate monohydrate could possibly occur. The former could lead to a decrease in potential near the pyrite surface and promotes the formation and accumulation of elemental sulphur which would initiate pyrite passivation by sulphur (Long & Dixon, 2004). However, due to the high operating temperature of Lihir autoclave, sulphur passivation is unlikely to occur. Lihir autoclave operating temperature of 220°C is well above the highest temperature limit reported in literatures for sulphur stability, which is 170°C (Papangelakis & Demopoulos, 1991; Bailey & Peters, 1976; Mackiw et al., 1966; Habashi & Bauer, 1966; Long & Dixon, 2004). However, it should be noted that this temperature varies with other autoclave conditions, such as oxygen partial pressure.

1.4.3 Lihir iron hydrolysis product Newcrest completed a mineralogical analysis of ten monthly composite pairs of autoclave feed (ACF) and neutralisation cyanidation adsorption feed (NCAF) during the period between June 2011 and April 2014. Based on the MLA and QXRD (Quantitative X-Ray Diffraction) analysis, hematite and jarosites were found to be the main phases precipitated inside Lihir autoclave. The jarosites are predominantly precipitate in the form of potassium + jarosite (KFe3(OH)6(SO4)2) due to the high potassium (K ) content in Lihir ore, sourced from K-bearing layer silicate such as biotite and illite (Newcrest Mining, 2015).

Figure 22 shows the mineral wt% difference between the ACF and NCAF for hematite and jarosites at Lihir processing plant. From this monthly baseline data, it was concluded that there is a negative correlation between the amount of hematite precipitated and that of jarosite, meaning when more hematite is being formed, less jarosite is being formed and vice versa.

Figure 22 The relationship between hematite and jarosite in the Lihir neutralisation feed (Newcrest Mining, 2015)

The report also concluded that Basic ferric sulphate (BFS) is absent from Lihir autoclave discharge slurry based on the QXRD analysis result. Newcrest believed that either the amount of BFS was below the detection limit which is in the order of 0.5 to 2 wt% or its formation was suppressed by the monovalent cations present in the solution (Newcrest Mining, 2015).

1.4.4 Lihir lime consumption Lime is used as the neutralising agent on site and it has been flagged as a significant contributor to the process plant operating cost (OPEX). Based on plant data from December 2013 to November 2015, the daily lime consumption at Lihir was found to significantly fluctuate from 10 to 22 kg-lime/ t-ore. To optimise the lime consumption, a good understanding of some factors that influence the lime consumption are necessary. The three major lime consumers in the neutralisation stage are listed below: • Neutralisation of free acid in CCD underflow. Lime reacts with any excess acid in the slurry that enters the neutralisation stage until the target pH of 11 is achieved. This free acid originates from the autoclave discharge and any acid generated in the CCD circuit. Although minimal, small ore particles of unreacted sulphides leaving the autoclave are likely to keep undergoing oxidation in the CCD train and therefore releasing additional acid into the slurry. The free acid neutralisation reaction can be written in Reaction 1.17 below. 32

H2SO4 (a) + Ca(OH)2 (a) → CaSO4.2H2O (s) (1.17)

• Precipitation of metal cations, such as Fe2+, Fe3+, Mg2+, Al3+, Cu2+ and Zn2+. These metal cations precipitate at high pH as metal hydroxide as shown in Figure 23. The upswing on the right side of the curve is due to metal anions in solution such as - - Cu(OH)3 and Al(OH)4 .

Figure 23 Solubility profiles of several metals hydroxides as a function of pH (Aube & Zinck, 2003)

The amount of lime consumed for metal hydrolysis is highly dependent on the cations concentration. These cations are introduced into the system from the ore, process water and sea water that enters from the CCD circuit. The general hydrolysis reaction can be expressed in Reaction 1.18.

X+ - (1.18) M (a) + X OH (a) → M(OH)X (s)

• Formation of reactive iron hydrolysis product in the autoclave. The iron hydrolysis products precipitated inside the autoclave react differently with lime in the neutralisation stage. Hematite has been known to be stable while BFS and jarosite, particularly hydronium jarosite, can be chemically instable and therefore would react with lime from a thermodynamic perspective to form ferric hydroxide and gypsum. Their reactions can be expressed in Reaction 1.19 and Reaction 1.20 (Ji et al., 2006).

2FeOHSO4 (s) + Ca(OH)2 (a) + 2H2O → Fe(OH)3 (s) + CaSO4.2H2O (s) (1.19)

(H3O)Fe3(SO4)2(OH)6 (s) + 2Ca(OH)2 (a)+ 2H2O (a) → 3Fe(OH)3 (s) + (1.20) 2CaSO4.2H2O(s) From these three possible lime major consumers, presence of BFS has been commonly associated with high lime consumption in gold pressure. Fleming (2009) stated that 8 kg/t lime is consumed for every 1% sulphate present in the autoclave residue. However, as discussed in Section 1.4.2, Newcrest did not find BFS in Lihir autoclave discharge slurry during the mineralogy study.

However, there is still a possibility that BFS might be present in Lihir autoclave discharge slurry, but it decomposed prematurely before the analysis was performed. This phase is known to be unstable at low temperature (<140°C) and acidic conditions (pH<1) where it is prone to undergo decomposition to form ferric sulphate according to Reaction 1.21. Therefore, depending on the post-sampling treatment of the mineralogy samples, such as delayed filtration and washing, premature decomposition of BFS could occur before the analysis was performed.

2FeOHSO4 (s) + H2SO4 (a) → Fe2(SO4)3 (a) + 2H2O (l) (1.21)

If basic ferric sulphate is formed inside the autoclave, this solid could partially or completely decompose in the CCD train. However, the pH of the slurry in the CCD should be higher than the autoclave discharge slurry due to the washing in the CCD and therefore could potentially stop or slow down the decomposition reaction. Consequently, basic ferric sulphate could reach the neutralisation stage and consume a large amount of lime.

Other than the presence of BFS, there is also another possibility that potassium jarosite or other less stable form of jarosite undergoes decomposition with lime despite of the claimed slow kinetic by Ji et al. (2006). Patino et al (2013) showed that there is an induction period where no changes observed on the jarosite particle surface, which was approximated to be 40 minutes at 30°C after being contacted with lime. This induction period varies depending on the temperature, particle size and OH- concentration. Following the induction period, the decomposition starts by forming a gel halo (Figure 24) in which K and S diffuse from the particle into the solution.

Unreacted Core Gel halo

Figure 24 Partially decomposed jarosite particle made of a gel halo, a reaction front and an unreacted core (Patino et al., 2013)

With the uncertainty of which solid consumes the majority of lime at Lihir, Eh-pH diagrams is needed to provide a prediction and therefore determine which solid phase precipitates under Lihir operating condition.

1.5 Thermodynamic model for Fe-SO4-H2O system at high temperature

For modeling purposes, Lihir system could be simplified and represented by Fe-K-S-H2O system. Hydrometallurgists often refer to the stability diagram shown in Figure 25 as a guide to define the stability of various iron precipitates, in particular iron hydroxyl-sulphates which includes both jarosite and basic ferric sulphates. It is practical and gives a good general indication of the solids’ stability field.

270 (Includes jarosite and BFS)

220

170

120

20 2 4 6 8 10

pH +3 -2 Figure 25 Iron stability diagram for the Fe -SO4 -H2O system as a function of temperature and pH (adapted from Fleming, 2009)

However, this stability field of iron hydroxyl sulphates is based on data for potassium jarosite precipitation from an earlier study by Babcan (1971) shown in Figure 13 in Section 1.3.3.2. As a result, the generalisation of the iron hydroxyl-sulphates stability field is valid but more experimental work is required to accurately define the phase boundaries between basic ferric sulphate and jarosite.

Another form of thermodynamic model which is often used by hydrometallurgist is the Eh- pH diagram. This diagram, which is also known as Pourbaix diagram, is a convenient format to display chemical thermodynamic data to display predominant species and phases in aqueous systems. To predict the solid phases that precipitate in Lihir autoclave, an Eh-pH diagram that includes potassium jarosite and basic ferric sulphate is needed.

However, these two solids have not been presented in any high temperature Eh-pH diagram published for this system at high temperature. For example, Huang (2008) and Biernat &

Robins (1972) presented Eh-pH diagram for the Fe-S-H2O system at 200°C and 220°C respectively as shown in Figure 26 and Figure 27. At acidic gold pressure oxidation condition

(pH < 4), both studies predicted that the stable iron precipitates at 200°C is hematite (Fe2O3) as the stability of basic ferric sulphate was not considered in both studies. This is likely due to the lack of thermodynamic data for basic ferric sulphate at high temperature.

Potassium jarosite, sodium jarosite and hydronium jarosite were considered by Huang (2008) but were found to be less favourable compared to hematite. This could be due to the low concentration of potassium specified in the model which was not stated in the paper. As shown in Figure 7 in Section 1.3.2.2, the stability of potassium jarosite depends on the potassium concentration in the system where there is a minimum concentration required before potassium jarosite becomes stable in the solution. Another reason could be due the discrepancies in the potassium jarosite thermodynamic data at high temperature (Majzlan et al. 2010; Stoffregen, 2000; Stoffregen, 1993; Drouet & Navrotsky, 2003).

Figure 26 Eh-pH diagram of Fe-S-H2O at 200°C (Biernat & Robins, 1972)

Figure 27 Eh-pH diagram of Fe-S-H2O at 220°C (Huang, 2008)

With the lack of relevant chemical thermodynamic data at high temperature for this

Fe-S-K-H2O system, in particular the stability data of BFS and potassium jarosite, a development of Eh-pH diagram to accurately predict the phase boundaries is currently not possible. Therefore, there is a need for a methodology that allows critical assessment of these thermodynamic data.

1.6 Summary Pressure oxidation (POX) is used worldwide to treat refractory sulphide gold-bearing ore. High purity oxygen is injected into the autoclave in order to oxidise the sulphide material, mostly pyrite, so that the occluded gold could be exposed and recovered during cyanidation.

The Lihir processing plant has four autoclaves operating at 220°C. Due to the mineralogy of the ore where the gold-containing pyrites (i.e. arsenian pyrite and arsenopyrite) are located on the outer rim of the ore particles, Lihir changed their autoclave operation philosophy from total oxidation (i.e. high oxidation extent) to partial oxidation. There was an indication that running the autoclave on partial POX affects the Fe3+ and Fe2+ behaviour at high temperature. Based on Lihir plant data from December 2013 to November 2015, Fe2+: Fe3+ ratio was found to increase at lower oxidation extent as expected, but the Fe3+ concentration was also found to increase which is likely due to the reduced iron hydrolysis extent. These

38 changes could affect both pyrite oxidation and iron hydrolysis that are occurring in the autoclave.

During pyrite oxidation, the literature review has indicated that ferric has been proved to be a potent pyrite oxidant at low temperature. The increase of both ferric and ferrous in the solution at partial POX condition could possibly affect the pyrite oxidation. However, the efficacy of ferric as surrogate oxidant at high temperature, where it undergoes iron hydrolysis concurrently, has never been investigated

From the iron hydrolysis perspective, the reduced iron hydrolysis extent means less acid is generated in the autoclave. Literatures has indicated that acidity plays a major role in favouring the stability of different iron phases. Based on the Lihir lime consumption, although not conclusive, there was an indication that lower amount of lime was consumed by reactive solid at partial POX condition. Therefore, running the autoclave on partial POX could possibly affect the phase and/or the amount of reactive solid precipitating in Lihir autoclave. However, the phase of this reactive solid remains unclear as basic ferric sulphate (BFS) was not detected in Lihir autoclave discharge while the reaction between potassium jarosite and lime was reported to be kinetically slow.

Thermodynamic model could help to identify the iron phase precipitating in Lihir autoclave. The model needs to include the three possible iron precipitates which are hematite, BFS and potassium jarosite, to make prediction more accurate. However, there is a lack of thermodynamic data for this Fe-K-S-H2O system at high temperature, more specifically on BFS. There are also discrepancies on potassium jarosite at high temperature reported in literatures. Therefore, the thermodynamic data for these two solids need to be determined to enable the development of this thermodynamic model (i.e. Eh-pH diagram) at 220°C.

1.7 Aim and objectives The aim of this thesis is to determine the stability of hematite, basic ferric sulphate and potassium jarosite at high temperature and pressure and to investigate the behaviour of ferric ion at partial pressure oxidation condition. To achieve this, the major objectives of the current research are to: • Investigate the role of ferric ion as a surrogate oxidant at pressure oxidation conditions; • Determine solubility of hematite, BFS and jarosite as a function of sulphate concentration, acidity and potassium concentration at pressure oxidation condition;

• Develop thermodynamic models for iron stability in solution at 220°C through experimental validation; • Apply learning to Lihir.

1.8 Outline of the thesis

This thesis is divided into six chapters and brief description of each chapter follows.

Chapter 1 is a general introduction to the present work. It provides a detailed review of pyrite oxidation mechanisms by oxygen and ferric ion, brief description of the structure of hematite, jarosite and basic ferric sulphate, and factors that affecting the formation of these aforementioned phases. It introduces Newcrest and its gold operation, general flowsheet of Lihir processing plant, Lihir pressure oxidation operation and high lime consumption in neutralisation stage. The objectives and aims of this study are also listed.

Chapter 2 provides a description of experimental methodology. It includes the microwave digester and autoclave descriptions, autoclave high temperature sampling set-up, in-situ high temperature ORP probe and redox ferrous titration methodology.

Chapter 3 describes the efficacy of ferric iron as pyrite surrogate oxidant at Lihir pressure oxidation operating temperature.

Chapter 4 presents a thermodynamic data review for species at high temperature and a methodology to critically assess existing thermodynamic data and define new thermodynamic data for basic ferric sulphate and potassium jarosite. The methodology was developed by relating solubility experimental data to chemical thermodynamic predictions using a series of assumptions and aqueous species modelling. It also describes the 3+ 2+ + 2- development of Eh-pH diagram of Fe -Fe - K -SO4 -H2O system at 220°C.

Chapter 5 presents the validation of the developed Eh-pH diagram using Lihir ore pressure oxidation experiments and the application of the validated Eh-pH on Lihir system to predict the reactive phase that precipitates in Lihir autoclave.

Chapter 6 presents overall conclusions and scope for future research.

Chapter 2: Common Experimental methodology

All experiments in this study was carried out at the University of Queensland hydrometallurgy laboratory. Speedwave Four microwave digester (Section 2.2.1) was used to investigate the role of ferric as surrogate oxidant while 2L Parr autoclave (Section 2.2.2) was used to investigate the solubility of hematite, basic ferric sulphate and potassium jarosite. For the autoclave experiments, sampling was done at high temperature with Millipore pressure filter (Section 2.2.3) and high temperature ORP measurement was recorded throughout the experiment (Section 2.2.4). The methodology covered in this chapter are the common methodology across the study. The specific methodology are covered in each Thesis Chapter.

2.1 Materials

2.1.1 Reagents All solutions were prepared with reagent grade chemicals and deionised water. The chemicals used in present work for feed solutions were:

97% iron (III) sulphate hydrate (Fe2(SO4)3.xH2O),

99% iron (II) sulphate hepta hydrate (FeSO4.7H2O),

99% potassium sulphate anhydrous (K2SO4),

98% lithium hydroxide monohydrate (LiOH.H2O),

98% concentrated sulphuric acid (H2SO4).

The chemicals used for the ferrous titration were:

99.7% potassium dichromate (K2Cr2O7),

Analytical grade Barium diphenylamine-4-sulfonate ((C12H10NO3S)2Ba),

98% concentrated sulphuric acid (H2SO4),

85% concentrated orthophosphoric acid (H3PO4).

2.1.2 Pyrite crystal High purity pyrite crystal clusters were purchased from GEOdiscoveries in New South Wales, Australia. The samples were crushed, ground and sieved to size range between 212 and 300 µm. Samples were stored in an air tight container in the freezer at all times to avoid premature pyrite oxidation. Samples were not chemically pre-treated or washed prior to use.

The XRD pattern and SEM image of the pyrite are shown in Appendix A.1. 41

2.1.3 Lihir Ore Ore samples (NTS 200), approximately 30-35 mm in size, were obtained from Newcrest’s Lihir operation in Papua New Guinea. Samples were crushed with hammer crusher to a size of approximately 1-1.5 mm with 100% passing. They were then split and ground to a target size p80 of 180 µm using ball mill. Samples were passed through a sample splitter that split the sample into two equal portions and was repeated three times to ensure samples homogeneity was achieved. Samples were stored in air tight container in the freezer after grinding.

The head assay of the ore was determined by: • Mineral liberation analysis (MLA) which was done by ALS Mineralogy Pty Ltd. • Solid digestion + inductively coupled plasma optical emission spectrometry (ICP- OES) which was done by the School of Agriculture and Food Sciences analytical laboratory at The University of Queensland. • Qualitative X-ray diffraction (QXRD). The XRD pattern was obtained using Bruker D8 Advance MKII X-ray diffractometer in the Centre of Microscopy and Microanalysis (CMM) at The University of Queensland. The QXRD analysis was done by the QUT Central Analytical Research Facility.

The MLA and QXRD results are presented in Section 5.3.1 and solid digestion + ICP result is presented in Section 5.3.3.

2.2 Equipment 2.2.1 Speedwave Four Microwave Digester Pyrite oxidation experiments were conducted in an oxygen deprived environment using a Speedwave Four microwave digester. There are eight TFM pressure vessels (DAK-100) inside the digester and each vessel consists of a ceramic pressure jacket and a TFM liner as shown in Figure 28. The maximum temperature that can be achieved by this set up is 300°C and therefore the operating temperature of 220°C for this study is well within the equipment capability. The temperature in the vessel was monitored and controlled using optical sensor technology (Speedwave DIRC thermometer) embedded in the digester.

A variation on the maximum temperature reached during the run was observed across the vessels. This is likely due to the measuring technique and the condition of the vessels. Therefore, the temperature of each vessels was closely monitored during the run and any vessels that deviated more than  3°C were discarded.

Note that the slurry in the microwave digester test was essentially stagnant (i.e. no mixing). 42

Figure 28 Speedwave Four Microwave Digester and its control box (left), DAK-100 TFM pressure vessel (bottom right) and TFM liner and ceramic pressure jacket (top right)

2.2.2 Autoclave

All solubility and Lihir ore pressure oxidation experiments were conducted using a 2-L Parr titanium autoclave. Glass liner was used to contain the experimental solutions to avoid cross contamination and to minimise reactor scaling. The autoclave temperature and agitation were controlled with a PAR 4843 controller system. Agitation was provided by dual 4- pitched-blade impeller and an agitation speed of 500 RPM was set for all experiments. Slurry temperature was measured using a thermocouple probe to an accuracy of about ±2°C and maintained at the target temperature by external heating and by-passing cooling water through the internal cooling coil. Cooling water flow was regulated by a solenoid valve. The setup of the autoclave is shown in Figure 29.

Figure 29 Autoclave setup at UQ Hydrometallurgy

2.2.3 High temperature sampling The online high temperature sampling was carried out with Millipore pressure filter and Supor®0.45µm hydrophilic polyethersulfone filter membrane. The slurry samples were taken using a sampling dip tube which was fitted to the autoclave head. From the dip tube, the samples were directed to the pressure filter via a 150mL 316SS Swagelok sample cylinder as shown in Figure 30. With this setup, both solid and solution samples were able to be retrieved at high temperature.

Once the high temperature sampling was completed, the pressure filter was detached from the sample cylinder by releasing the pressure in the filter via the relief valve. In the fumehood, high pressured compressed air was used to remove most of the entrained solution. Subsequently, 20 mL of DI water was poured into the filter to wash the solid. High pressured compressed air was once again used to separate the wash liquid from solid residue. Due to the thin solid bed on the filter paper (~1mm in thickness), the filtration process took less than 3 seconds to complete and therefore dissolution of solid during washing would be minimal.

Figure 30 High Temperature Autoclave sampling setup

2.2.4 In-situ high temperature ORP High temperature ORP measurement was done with a pressure balanced external reference probe supplied by Corr Instruments shown in Figure 31. This type of probe is more stable than internal probes, considering that the autoclave temperature is above 100°C. It can operate up to 300°C with a maximum pressure rating of 140 bar (g).

Figure 31 In-situ ORP Probe

The electrode design incorporates a Ag/AgCl junction exposed to KCl solution. Two ORP probes with different concentrations of KCl solutions were tested in the preliminary experiments. The probe with the internal solution of 0.5 M KCl solution was found to give a more stable reading compared to the probe with an internal solution of 0.1 M KCl. Therefore, the probe with an internal solution of 0.5 M was used for this study.

The titanium vessel of the autoclave was used as the working electrode for the ORP probe. Therefore, the signals obtained in this study represent the reduction-oxidation reactions that are occurring on the titanium surface.

The ORP cell can then be defined as:

2+ 3+ Ti (s) | H2SO4 solution containing Fe /Fe | 0.5M KCl (a) | AgCl (s) | Ag

All potential values in this study were reported with respect to the normal hydrogen electrode (V vs. NHE). The conversion of the measured potential values obtained in this study (V vs. Ag/AgCl) to NHE was done via standard hydrogen electrode (SHE) by assuming that V vs. SHE is equal to V vs. NHE. Although there is technically a difference of 0.006 V between NHE and SHE (Connelly et al., 2015), the difference is small enough to be within the experimental error.

The conversion from Ag/AgCl to SHE reference electrode depends on the KCl solution concentration as shown in Table 10 (Roberge, 2008; Smith & Stevenson, 2007). The conversion for KCl concentration of 0.5 M used in this study could not be found in literature review. Therefore, it was approximated to be 250 mV as shown in Figure 32.

Table 10 conversion from Ag/AgCl reference electrode to standard hydrogen electrode (SHE) at different KCl concentration

KCl concentration in Potential (V vs. SHE), Ag/AgCl electrode, M mV 0.1 288 1 235 3.5 205

Figure 32 Estimation of conversion from Ag/AgCl reference electrode to standard hydrogen electrode (SHE) as a function of KCl concentration The measured potential data was corrected for thermal junction potential (TJP) which was created from the temperature gradient (ΔT) that exists between the working electrode environment and the reference electrode temperature. For 0.5 M KCl solution, Equation 2 was used in present work to estimate the corrected solution potential with respect to the standard hydrogen electrode (SHE) (Macdonald, 1979).

−8 3 −5 2 −3 ∆퐸푇퐽푃(SHE) = 0.247 − 0.257 × 10 ∆푇 − 0.102 × 10 ∆푇 + 0.858 × 10 ∆푇 Equation 2

Throughout autoclave preliminary testwork, the temperature of the Ag/AgCl section of the probe was measured and found to increase to a maximum value of 33°C, which translates to T of 187°C given that the autoclave temperature was maintained at 220°C. Using Equation 2, the TJP correction was calculated to be 34 mV.

Additionally, liquid-liquid junction potential (LJP) correction was also applied to the measured potential data. This junction potential is generated from the differences in the ionic composition between the reference electrolyte and the solutions. The Henderson equation, described in Equation 3 to Equation 6, was used to calculate the LJP between the two solutions. 47

푅푇 퐵Ι⁄퐵ΙΙ Equation 3 ΦΙ − ΦΙΙ = − × 퐴 × ln 퐹 퐵Ι−퐵ΙΙ

Ι ΙΙ Equation 4 퐴 = ∑ 푧푖푢푖(푐푖 − 푐푖 )

Ι 2 Ι Equation 5 퐵 = ∑ 푧푖 푢푖푐푖

ΙΙ 2 ΙΙ Equation 6 퐵 = ∑ 푧푖 푢푖푐푖

Where R is 8.314 J mol-1K-1, T is the temperature in Kelvin, F is 96500 Coulombs mol-1 and

푧푖 is the ionic charge of species i. The ionic mobility 푢푖 in 퐴 and 퐵 can be replaced by ionic diffusion coefficient 퐷푖 (Newman & Thomas-Alyea, 2004). Table 11 lists the diffusion coefficient of the related ions at infinite dilution in water at 25°C. Sample calculations can be found in Appendix B.1.

Table 11 Values of diffusion coefficients of selected ions at infinite dilution in water at 25°C

ퟓ 2 Ion 풛풊 푫풊 × ퟏퟎ , cm /s H+ 1 9.312 Li+ 1 1.030 K+ 1 1.957 Ca2+ 2 0.792 Mg2+ 2 0.706 Fe2+ 2 0.604 Fe3+ 3 0.719 Al3+ 3 0.559 OH- -1 5.260 - HSO4 -1 1.330 2- SO4 -2 1.070 Cl- -1 2.032

2.3 pH measurement pH measurement was taken with two TPS intermediate junction pH sensors. The probe is using Ag/AgCl reference electrode with intermediate junction reference design. Two-point calibration was done in standard buffer solution of pH 4 and pH 1 prior to every measurement. pH measurement was taken at room temperature (below 30°C) together with a temperature electrode. pH was allowed to stabilise until the value remained the same for approximately 1 minute. 48

2.4 Ferrous titration

Quantification of ferrous iron concentration was done by potentionmetric titration with 0.1N potassium dichromate (K2Cr2O7) solution. The reaction involved in this titration follows Reaction 2.1 below.

-2 2+ + +3 3+ Cr2O7 (a) + 6Fe (a) + 14H (a) → 6Fe (a) + 2Cr (a) + 7H2O(l) (2.1) For each titration, the sample aliquot was mixed with 50mL of de-oxygenated distilled water and 10mL of sulphuric/phosphoric acid (H2SO4/H3PO4) mixture. To minimise the oxidation of ferrous to ferric during the titration, the DI water was previously sparged with nitrogen gas for 10 minutes. The sulphuric/phosphoric acid mixture consisted of 1:1:1 ratio of 98% sulphuric acid: 85% phosphoric acid: DI water. 5 drops of indicator (0.5 g/L barium diphenylamine sulphonate) was also added to trigger a permanent colour change from clear to a deep purple colour at titration endpoint. Equation 7 was used to calculate ferrous concentration based on the volume of titrant used.

퐹푒2+ [푇𝑖푡푟푎푛푡] × 푉 × 푛 × 푀푊 푡푖푡푟푒 퐶푟 푂 2− 퐹푒 Equation 7 [퐹푒2+] = 2 7 푉푠푎푚푝푙푒 Where, [Fe2+] = concentration of ferrous in solution (g/L),

[Titrant] = concentration of K2Cr2O7 solution which is 0.0167M,

Vtitre = volume of K2Cr2O7 titrant solution used (mL)

2+ 2- 2+ 2- n Fe /Cr2O7 = molar ratio of reacted Fe ions to Cr2O7 ions, which is 6

MW Fe = molecular weight of iron which is 55.8 g/mol

Vsample = volume of sample being titrated (mL)

2.5 Solid identification

2.5.1 Scanning electron microscope (SEM)

JEOL JSM-7001F and JEOL JSM-6610 which are available in the Centre of Microscopy and Microanalysis (CMM) at The University of Queensland, were used interchangeably for this study. The JEOL JSM-7001F was used mainly to study the morphology of the crystals while JEOL JSM-6610 was used for elemental analysis as it is equipped with an Oxford 50mm2 X-Max SDD X-ray detector.

Powder particles were mounted on a high purity conductive double sided adhesive carbon tabs. All samples were put inside vacuum oven for at least 8 hours at 70°C to allow complete degassing before they were coated with carbon. Prior and after carbon coating, all samples were cleaned to remove hydrocarbon contamination from surfaces using the Evactron 25 De-Contaminator RF Plasma Cleaner.

An accelerating voltage of 10 - 15kV and working distance of 10mm were typically used. For EDS spectra acquisition, the spot size was varied from 50 to 54mm to allow a high count while maintain a dead time value of between 12 and 16%.

2.5.2 X-ray Diffraction (XRD)

X-ray diffraction patterns of pyrite samples, solubility solid residues and Lihir ore POX residues were conducted using Bruker D8 Advance MKII X-ray diffractometer in the Centre of Microscopy and Microanalysis (CMM) at The University of Queensland. All samples were scanned with a copper (Cu) target source at the test conditions of 40 kV and 150 mA. The scan was run with a scan step of 0.01° of 2θ and a counting time of one second per step with 5 ° to 90° scanning range.

For all raw patterns, the K-alpha-2 radiation was stripped prior to peak identification. Peak identification was carried out by matching the peaks with the mineral data files in the RGB- 2019 database. Due to the detection limit of the X-ray diffraction, any minerals that occupies less than 5% of the mixture will not be detected.

2.5.3 Quantitative X-Ray Diffraction (QXRD)

2.5.3.1 Sample preparation

All samples were ground with Retch XRD-Mill McCrone using agate grinding media. A grinding time of 2 minutes at grinding power of 2 (out of 4 settings) was specified to avoid excessive breakage which can potentially induce phase transformations.

A known amount of internal standard was added into the samples to help increasing the accuracy of the quantification. Corundum (Al2O3) was selected as the internal standard.

2.5.3.2 Analysis TOPAS Academic Software v.4.1 was used in present investigation to conduct the quantitative XRD (QXRD) analysis. This analysis was used to measure phase fractions in all samples that contained more than one phase. QXRD mathematically calculates the XRD pattern and compares the calculated pattern with a known diffraction pattern obtained from 50

PDF-2019 data base in a process known as the fundamental parameters approach (FPA). Parameters such as crystal structure, atom type and ordering were refinement within the software until a “best fit” is achieved between the entire observed diffraction pattern and the predicted pattern.

The codes used for present investigation are shown in Appendix D.1.

2.6 Solution analysis

Composition of elements in solution samples was determined by inductively coupled plasma – optical emission spectrometry (ICP-OES) technique. All samples were stabilised by diluting 10X in 1wt% HCl solution. The analysis was done by the School of Agriculture and Food Sciences analytical laboratory at The University of Queensland.

Chapter 3: The role of ferric as surrogate oxidant during pyrite pressure oxidation

3.1 Introduction

At low temperature (below 100°C), ferric has been shown to be an effective oxidant for pyrite (Moses et al., 1987; Singer & Stumm, 1970; Garrels & Thompson, 1960; Holmes & Crundwell, 2000; Goldhaber, 1983). Not only for pyrite, the capability of ferric as an effective oxidant for other sulphide minerals, such as chalcopyrite (Kinnunen et al., 2006, Hiratori et al., 1987) and Pyrrhotite (Janzen et al., 2000) has also been established.

Moses et al (1987) demonstrated that ferric was able to compete with oxygen as pyrite oxidant in basic environment (between pH 7 and 9) at 25°C even though the solubility of ferric is much lower than the oxygen solubility at that condition. Ferric solubility is approximately 0.01 ppb as shown in Figure 33 (Kim et al., 2015) while oxygen solubility is approximately 40 ppm as shown in Figure 34 (Tromans, 2000).

Figure 33 Equilibrium solubility of Fe(OH)3 at 25°C (Kim et al., 2015)

Figure 34 Oxygen solubility in pure water as a function of temperature (Tromans, 2000)

As discussed in Section 1.3.1.1, pyrite oxidation by ferric may involve the formation of 2- thiosulphate (S2O3 ) as the first sulphur intermediate species (Luther, 1987; Moses et al., 1987; Druschel et al., 2003; Descostes et al., 2004; Descostes et al., 2001; Rimstidt & Vaughan, 2003). However, thiosulphate is unstable in highly acidic solution where it would undergo decomposition. In the presence of excess ferric, thiosulphate would be completely oxidised to sulphate via the formation of various sulphur intermediate. Luther (1987) reported that the kinetic of this reaction is rapid enough such that thiosulphate would not be able to co-exist with ferric in solution. The overall reaction of pyrite oxidation by ferric then can be written according to Reaction 4.3 where fourteen moles of ferric is required to oxidise one mol of pyrite.

3+ 2+ 2- + FeS2 (s) + 14Fe (a)+ 8H2O (l) → 15Fe (a) + 2SO4 (a) + 16H (a) (3.1)

Despite of the importance of ferric ion as pyrite oxidant at low temperatures, the role of ferric ion in catalysing the kinetic of pyrite oxidation at pressure oxidation condition has not been clearly established (Papangelakis & Demopoulus, 1991). Contradicting experimental results were reported. McKay & Halpern (1958), Gerlach et al. (1966) and Warren (1958) reported that the presence of ferric played an insignificant role in pyrite oxidation. McKay & Halpern (1958) suggested that although the oxidation of pyrite by ferric is relatively rapid at 110°C,

53 the re-oxidation of ferrous to ferric is significantly slower than the rate of direct oxidation of pyrite by oxygen. In this case, the catalysing effect of ferric on the kinetic of pyrite oxidation at high temperatures would largely depend on the initial concentration of ferric.

This could possibly explain the contradicting results reported in Gerlach et al. (1966) and King & Lewis (1980) who carried out their experiments at the same conditions. In autoclave experiment at 100°C and 120 psig oxygen pressure, Gerlach et al. (1966) added 0.05 M Fe3+ into the system and found that the rate of pyrite oxidation remained unchanged after 1 hour of test. With the addition of 2 M Fe3+, they observed some slight increase in the overall pyrite oxidation extent but it was within their experimental error. On the other hand, King & Lewis (1980) claimed that adding 0.2 M and 1 M Fe3+ enhanced the overall pyrite oxidation by approximately 1.5 and 2 times respectively after a 1-hour test as shown in Figure 35. From these results, it could mean that the catalytic effect of ferric on pyrite oxidation could only be observed at high initial ferric concentration.

Figure 35 Pyrite conversion over time as a function of oxygen partial pressure and temperature at ferric concentration of 0.2 M (left) and 1 M (right) (King & Lewis, 1980)

At pressure oxidation conditions, which is typically carried out at above 170°C, ferric is unstable and tends to undergo hydrolysis to form various iron precipitates. Depending on the acidity and ferric concentration, hematite, basic ferric sulphate or both phases could be form (refer to Section 1.3.3 for further details on factors affecting the iron phase stability). This precipitation causes the ferric to leave the solution phase and consequently only the 54 residual ferric would be capable to oxidise pyrite. This means that the effective ferric concentration in the autoclave may not be high enough for the catalytic effect of ferric to be observed, depending on the proportion of ferric that precipitates and remains in the solution which has never been investigated.

This study will explore the behaviour of ferric as pyrite oxidant at gold pressure oxidation operating temperature. To do this, the extent of pyrite oxidation and iron deportment needs to be assessed discretely at 220°C under deoxygenated conditions. The relationship of pyrite oxidation extent with initial ferric concentration, solid loading and ferrous concentration will be explored and iron deportment in the system will also be determined.

3.2 Experimental

3.2.1 Equipment For this study, all experiments were carried with microwave digester described in Section 2.2.1. With the use of microwave digester, eight experiments can be done at the same time which means experiment repeats can be carried out with the same feed solution. This is beneficial in removing any variance on ferric and ferrous concentrations in the feed. However, mixing could not be introduced into the system as the vessel are essentially stagnant. This means the oxidation extent reported in this study may not be optimised.

3.2.2 Methodology

Feed solution was prepared by mixing powder ferric sulphate (Fe2(SO4)3) with nitrogen- sparged Deionised (DI) water. The solution was stirred until all solids have evidently dissolved. It was then sampled and stabilised in 2wt% HCl solution for assaying by ICP-OES and pH measurement was taken. Subsequently, pyrite was weighed into the microwave digester liner before adding in 20mL of feed solution.

The solutions required 25 minutes to reach the target temperature of 220°C based on - observation during preliminary experiments. During the heating, bisulphate (HSO4 ) will be - formed as it becomes stable at high temperature. The formation of HSO4 reduces the activity of hydrogen ion (H+) in the solution which then increases the “at temperature” pH. Once the target temperature was reached, the experiment was allowed to run for 2 hours. Pressure vessels were allowed to cool down for 15 minutes inside the digester before unloading. This was done to reduce the risk of spontaneous pressure release during unloading. During the cool down, however, some solids (i.e. basic ferric sulphate) become unstable at lower temperature and therefore dissolution of these solids may occur to a

55 certain extent. Table 12 shows the settings entered into the digester control system for the 2-hour experiment done in this present work.

Table 12 SpeedWave Four Microwave Digester Settings

Temperature Pressure (bar) Ramp-up (min) Time (min) Power 120°C 60 5 5 80 180°C 60 5 5 80 220°C 60 5 55 80 220°C 60 10 55 80 50°C 60 5 10 80

After the vessels were unloaded, they were quench cooled in an ice bath for 30 min. A ferrous titration was performed immediately upon opening the liner’s lid. A specific amount of filtrate was taken out of the liner to be filtered with 0.2 micron PTFE syringe filter. The filtered filtrate, varied between 2mL and 10mL depending on the expected ferrous concentration, was then added to the acid mixture ready for titration.

The remaining of the slurry in the liner was then vacuum filtered with 0.45 micron pore-size PTFE membrane filters. Solids were washed with DI water (1:1 ratio) to displace entrained solution and were oven dried at 60°C, typically for 24 hours. They were then characterised using SEM-EDS and XRD for particle morphology and phase identification. Filtrate was sampled and stabilised in 2wt% HCl solution for assaying by ICP-OES and pH measurement was taken.

3.2.3 95% Confidence interval calculation Each experiment was repeated four times. All data presented in this chapter is an average of the four data points and the error bars that represent the 95% confidence interval is presented. This uncertainty was estimated based on a two-tailed t-test using Equation 8.

s E95% = ±tα/2,DOF × Equation 8 √n

Where 푡훼/2,퐷푂퐹 is critical t value, s is standard deviation and n is the number of samples.

The critical t-value was obtained directly via Excel using the T.INV(/2, DOF) formula for a two-tailed test. The degree of freedom (DOF) is equal to the number of samples minus the number of estimates (i.e. DOF = n-1). For four samples, the t-critical value was calculated to be ± 2.78.

The standard deviation was estimated based on the relevant data, such as pyrite oxidation and ferric utilisation for oxidation, using the STDEV.S command in excel.

3.2.4 Ferrous Quantification Ferrous was measured using redox titration with potassium dichromate which is discussed in Section 2.4. The sample aliquot volume was varied between 5 and 2 mL depending on the expected ferrous concentration in the sample. Since the sample temperature remains hot upon opening the lid, the accuracy of volume pipette is uncertain. Therefore, the mass of sample was measured and subsequently converted to volume using the filtrate density measured once the filtrate has cooled down.

It is understood that the solubility of ferrous sulphate decreases at high temperature. However, Cheng (2002) found that minimal amount (~5 - 20%) of the initial ferrous crystallised out of the solution in the absence of stirring as shown in Figure 36 . Since there is no stirring in microwave digester experiment, the precipitation of ferrous as ferrous sulphate monohydrate (FeSO4.H2O) is still expected to occur but the effect of ferrous investigation on pyrite oxidation should still be valid as the majority of the initial ferrous added should remain in the solution.

When ferrous sulphate monohydrate (FeSO4.H2O) was detected in the solid phase based on the XRD result, the amount of ferrous measured by titration is adjusted by adding an equivalent amount of ferrous measured in the solid phase. Quantification of FeSO4.H2O phase in the solid was done using QXRD method in TOPAS program. The QXRD methodology will be discussed in Section 3.2.7.

Figure 36 Typical Fe2+ concentration profile as a function of time (Cheng, 2002) 3.2.5 Pyrite oxidation determination The extent of pyrite oxidation by ferric ion is determined by measuring the amount of ferrous ion generated during the experiment. This pyrite oxidation determination methodology is based on the following assumptions which were made based on the literature review and preliminary experiments: • Oxidation of ferrous to ferric during the experiment was assumed to be insignificant considering the unlikelihood of oxygen to enter the sealed digester liner and the slow ferrous oxidation at low temperature upon opening the lid. This was confirmed using a blank test with ferrous sulphate solution. • Pyrite oxidation by dissolved oxygen present in the feed solution is assumed to be insignificant. • The thiosulphate oxidation to sulphate is assumed to go to completion i.e. no intermediate sulphoxy anions and/or sulphur (S0) present in the system as Moses et al (1987) did not observe any presence of intermediate sulfoxy anions in their acidic ferric-saturated anaerobic experiments. Therefore, all ferrous generated during the experiment is assumed to be generated from pyrite oxidation and ferric reduction reactions. For every one mole of pyrite oxidised, a total of 15 moles of ferrous is produced where 1 mole is generated from pyrite oxidation and 14 moles are generated from ferric reduction.

Based on these assumptions, the amount of pyrite oxidised can be calculated using Equation 9 below. 1 푚표푙 퐹푒푆 푚표푙 푃푦푟𝑖푡푒 (푚표푙) = 2 × [퐹푒2+] ( ) × 푉 (퐿) Equation 9 푂푥푖푑푖푠푒푑 15 푚표푙 퐹푒2+ 푓푖푙푡푟푎푡푒 퐿 푓푖푙푡푟푎푡푒

The extent of pyrite oxidation was then estimated using Equation 10 below where molecular weight (MW) of pyrite (FeS2) is 119.98 g/mol.

푃푦푟𝑖푡푒푂푥푖푑푖푠푒푑 (푚표푙) 푔 % 표푥𝑖푑푎푡𝑖표푛 푒푥푡푒푛푡 = × 푀푊푃푦푟푖푡푒 ( ) × 100% Equation 10 푃푦푟𝑖푡푒푓푒푒푑 (푔) 푚표푙

3.2.6 Ferric balance To understand how iron hydrolysis affects the role of ferric ion as pyrite surrogate oxidant, a mass balance on iron was carried out around the system, as illustrated in Figure 37, to estimate the deportment of ferric iron. The amount of ferric ion in the feed as ferric sulphates (1) is equal to the amount of ferric used for pyrite oxidation (2), ferric that precipitates (3) and ferric that remains in filtrate (4). The amount of ferric in feed and that in filtrates are taken as the difference between total iron and ferrous concentration. Total iron concentration was measured by ICP-OES, while ferrous concentration was determined by titration with potassium dichromate. Refer to Section 2.2.3 for further details on redox titration with potassium dichromate.

. Figure 37 Iron mass balance illustration

Based on Reaction 3.1, the amount of ferric ion used for pyrite oxidation was estimated using Equation 11.

14 푚표푙 퐹푒3+ 푚표푙 퐹푒3+ = × [퐹푒2+] ( ) Equation 11 표푥. 15 푚표푙 퐹푒2+ 푓푖푙푡푟푎푡푒 퐿

The amount of ferric that remains in the filtrate and precipitates were estimated using Equation 12 to Equation 14.

푚표푙 푚표푙 퐹푒3+ = [퐹푒 ] ( ) − [퐹푒2+] ( ) Equation 12 푟푒푚푎푖푛푖푛푔 푇 푓푖푙푡.(퐼퐶푃) 퐿 푓푖푙푡푎푡푒 퐿 푚표푙 푚표푙 퐹푒3+ = [퐹푒 ] ( ) − [퐹푒2+] ( ) Equation 13 푓푒푒푑 푇 푓푒푒푑 (퐼퐶푃) 퐿 푓푒푒푑 퐿 푚표푙 푚표푙 푚표푙 퐹푒3+ = [퐹푒3+] ( ) − [퐹푒3+] ( ) − [퐹푒3+] ( ) Equation 14 푝푟푒푐푖푝. 푓푒푒푑 퐿 푟푒푚푎푖푛푖푛푔 퐿 표푥. 퐿

During the preliminary studies, the iron deportment of three different initial ferric concentrations were compared in the presence and absence of pyrite as summarised in Table 13. At the same initial ferric concentration, the residual ferric in the solution was found to remain relatively constant within the limit of ICP-OES accuracy regardless of the pyrite presence. In the presence of pyrite, however, a certain proportion of ferric that underwent hydrolysis seems to be used for pyrite oxidation. Consequently, the extent of ferric that precipitates was evidently less in the presence of pyrite compared to that in the absence of pyrite.

Table 13 The effect of initial ferric concentration on iron deportment in the absence and presence of pyrite at 5 g/L solid loading in microwave digester at 220°C

Initial Fe3+ Fe3+ Fe3+ for No. [Fe3+], g/L precipitated (g) remaining (g) oxidation (g) 5.7 Pyrite 0.06 0.0026 0.05 1 6.4 No Pyrite 0.11 0.0076 - 14 Pyrite 0.14 0.016 0.13 2 14 No Pyrite 0.27 0.017 - 66 Pyrite 0.62 0.35 0.35 3 66 No Pyrite 1.01 0.31 -

3.2.7 Quantitative X-ray Diffraction method Sample preparation and analysis procedures can be viewed in Section 2.2.3. Based on the

QXRD analysis, four solid phases were identified in this study; they are hematite (Fe2O3), basic ferric sulphate (FeOHSO4), ferrous sulphate monohydrate (FeSO4.H2O) and pyrite

(FeS2). The codes used for the analysis are shown in Appendix D.1.

Figure 38 shows an example of the Rietveld refinement graphics from the analysis in TOPAS. The hematite (00-033-0664), basic ferric sulphate (04-012-6256) and the corundum (00-005-0712) peaks shown were taken from the PDF-2019 data base. Although the fitting 60 of the refinement to the original pattern is not perfect where there are some peaks that could not be fitted, the overall fitting is deemed to be acceptable

Figure 38 Example of Rietveld refinement graphics from QXRD analysis in TOPAS

3.3 Result

3.3.1 Effect of initial ferric concentration and solid loading on pyrite oxidation Table 14 shows the solution conditions of both the feed and filtrate solutions for 2.5, 5 and 10 g-pyrite/L solid loading at various initial ferric concentrations. Ferrous was detected in all experiments where pyrite was contacted with ferric sulphate solution (Test 1 to Test 20). Note that the data shown are the average of four tests carried out in parallel.

In contrast, the amount of ferrous generated from pyrite oxidation in the acidified DI water (initial pH of 1) was found to be very low with only 0.02 g/L Fe2+. With a total iron concentration of 0.02 g/L, the ferric concentration was assumed to be negligible. This indicates minimal pyrite oxidation had occurred (<0.1%) and therefore pyrite oxidation by dissolved oxygen in present work can be considered as insignificant and no ferrous offset was applied for further solution analysis.

Table 14 Oxidation extent data at different ferric concentrations and solid loadings

Feed Filtrate

Test Test ID Pyrite [Fe3], [Fe2+], Pyrite Fe3+, g/L pH Loading, g/L g/L g/L Oxidation (%) 0 No Ferric 0 5 - 0.02 <0.1% 0.99 1 D2.5_50 8.8 0.16 3.5 19.3 0.9 2.5 2 D2.5_120 16.4 0.76 7.9 43.1 0.6

Feed Filtrate

Test Test ID Pyrite [Fe3], [Fe2+], Pyrite Fe3+, g/L pH Loading, g/L g/L g/L Oxidation (%) 3 D2.5_140 21.2 1.89 10.0 54.4 0.65 4 D2.5_180 27.0 8.54 10.9 59.3 0.59 5 D2.5_200 30.7 12.12 11.0 60.1 0.60 6 D2.5_240 35.1 13.26 11.2 61.1 0.57 7 D2.5_280 40.4 16.80 11.3 63.3 0.56 8 D2.5_360 52.7 20.30 11.7 63.7 0.43 9 D5_10 2.7 0.99 1.1 2.9 1.1 10 D5_20 5.7 0.12 2.7 7.4 0.95 11 D5_50 14.5 0.13 7.4 20.2 0.68 5 12 D5_90 26.5 5.3 14.8 40.4 0.58 13 D5_140 40.4 14 19.7 54.1 0.55 14 D5_180 51.0 20 20.7 57.2 0.44 15 D10_20 13.5 0.089 8.2 11.1 - 16 D10_50 30.2 0.19 18.6 25.3 - 17 D10_60 34.1 11.6 20.3 30.3 0.58 10 18 D10_70 41.6 14.0 24.0 34.6 0.49 19 D10_80 48.7 19.7 24.5 35.3 0.48 20 D10_90 53.1 20.1 24.6 35.6 0.46

The effects of initial ferric concentration on pyrite oxidation extent and ferric utilisation for each solid loading are shown in Figure 39 and Figure 40. In general, oxidation extent was found to increase linearly as a function of initial ferric concentration at the lower end of ferric concentration. In this region, the amount of ferric used for pyrite oxidation for all three solid loading were found to be relatively similar, with slightly higher amount of ferric used at higher solid loading.

However, upon reaching a certain initial ferric concentration, the pyrite oxidation extent seemed to reach a plateau which could indicate that the addition of ferric at higher initial concentration was not used for oxidation. This independence of pyrite oxidation on initial ferric concentration at high ferric concentration is likely to be caused by the higher driving

62 force for ferric to undergo hydrolysis. This causes the ferric to leave the solution phase by precipitation, and consequently only the residual ferric would be capable to oxidise pyrite.

The onset concentration at which the plateau was observed, was different for each solid loading where it seems to increase at higher solid loading. For 2.5 g/L solid loading test, the onset concentration was at ~25 g-Fe3+/L, while ~40 g-Fe3+/L and ~50g-Fe3+/L initial ferric concentration were the onset concentrations for the 5 g/L and 10 g/L solid loading. These values are not conclusive, but they give a good indication that the onset concentration is higher at higher solid loading. This is likely because there are more surface area available for ferric to oxidise pyrite before the hydrolysis is completed at higher solid loading. All these observations suggest that pyrite oxidation is surface area limited.

Figure 39 The effect of initial ferric concentration on pyrite oxidation in microwave digester experiment at 220°C under deoxygenated conditions at 2.5 g-pyrite/L ( ∆ ), 5 g-pyrite/L () and 10 g-pyrite/L (O) solid loading.

Figure 40 the effect of initial ferric concentration on the amount of ferric used for pyrite oxidation at 220°C under deoxygenated conditions at 2.5 g-pyrite/L ( ∆ ), 5g-pyrite/L () and 10g-pyrite/L (O) solid loading.

3.3.2 Effect of initial ferric concentration and solid loading on ferric deportment To understand how the ferric deportment is affected by the initial ferric concentration and solid loading, iron mass balance was performed as described in Section 3.2.5. Table 15 tabulates the proportion of ferric that is used for pyrite oxidation, ferric that undergoes iron hydrolysis and ferric that remains in the solution as a function of initial ferric concentration.

There is a clear negative correlation between the proportion of ferric used for oxidation and the proportion of ferric that undergoes iron hydrolysis. Below the onset concentration where pyrite oxidation become independent of initial ferric concentration, the proportion of ferric ion that is used for pyrite oxidation increased while the proportion for iron hydrolysis decreased as the initial ferric concentration was increased. Above the onset concentration, the independence of ferric used for oxidation is reflected in the drop of ferric proportion used for oxidation. Excess ferric was found to be utilised in iron hydrolysis which caused the proportion of iron hydrolysis to increase at high initial ferric concentration.

Table 15 Ferric deportment at 220°C under deoxygenated conditions at 2.5 g/L, 5 g/L and 10 g/L solid loadings

Fe3+ for Oxidation (%) Fe3+ remaining (%) Fe3+ precipitation 3+ [Fe ]i % g % g % g Pyrite Loading 2.5 g/L 8.8 42% 0.063 8% 4.9E-3 51% 0.076 16.4 43% 0.14 9% 0.15 48% 0.16 21.2 42% 0.18 9% 0.16 50% 0.18 27.0 36% 0.19 30% 0.16 34% 0.18 30.7 32% 0.20 37% 0.23 31% 0.19 35.1 28% 0.20 36% 0.25 36% 0.25 40.4 25% 0.21 40% 0.31 36% 0.29 52.7 20% 0.21 37% 0.39 44% 0.46 Pyrite Loading 5 g/L 2.7 36% 0.019 1% 3.6E-4 63% 0.034 5.7 42% 0.048 2% 2.3E-3 56% 0.064 14.5 46% 0.13 6% 0.017 49% 0.14 26.5 48% 0.26 18% 0.10 34% 0.19 40.4 44% 0.35 33% 0.26 24% 0.19 51.0 37% 0.37 37% 0.37 27% 0.27 Pyrite Loading 10 g/L 13.5 54% 0.14 33% 0.089 13% 0.036 30.2 61% 0.33 25% 0.15 14% 0.084 34.1 59% 0.41 28% 0.19 13% 0.090 41.6 54% 0.45 27% 0.22 19% 0.16 48.7 47% 0.46 36% 0.35 17% 0.17 53.1 44% 0.47 33% 0.35 23% 0.25

Depending on the solid loading, a variation from approximately 20% to 60% of the available ferric is capable to oxidise pyrite despite of the removal of ferric through iron hydrolysis that occurs concurrently. The proportion of ferric utilised for pyrite oxidation as a function of initial ferric concentration for all three pyrite solid loading is graphed in Figure 41. It can be seen that ferric utilisation for pyrite oxidation at higher solid loading is higher for all initial ferric concentration tested.

Figure 41 the effect of initial ferric concentration on ferric utilisation (%) for pyrite oxidation at 220°C under deoxygenated conditions at 2.5 g-pyrite/L ( ∆ ), 5g-pyrite/L () and 10g- pyrite/L (O) solid loading.

The proportion of ferric that remains in the solution was found to constantly increase as a function of initial ferric concentration. At higher initial ferric concentration, more ferric undergoes hydrolysis which generates acid. Reaction 3.2 and Reaction 3.3 show the iron hydrolysis reaction to precipitate hematite and basic ferric sulphate respectively.

Fe2(SO4)3(a) + 3H2O(l) → Fe2O3(s) + 3H2SO4(a) (3.2)

Fe2(SO4)3(a) + 2H2O(l) → 2FeOHSO4(s) + H2SO4(a) (3.3)

If more acid is generated in the solution, there will be a higher residual ferric concentration in solution at the end of the reaction as shown in Figure 42.

Figure 42 the effect of initial ferric concentration on filtrate pH (left-axis) and ferric concentration in the filtrate (right-axis) for pyrite oxidation at 220°C under deoxygenated conditions at 2.5 g-pyrite/L ( ∆ ), 5g-pyrite/L (  ) and 10g-pyrite/L (O) solid loading.

3.3.3 Effect of ferrous concentration Figure 43 shows the effect of ferrous concentration on pyrite oxidation extent and Table 16 summarises the solution assay from the experiments. The results indicate that pyrite oxidation extent by ferric decreased as ferrous concentration was increased. This finding is in agreement with finding from previous low-temperature studies where at ferrous concentration above 10-3 mol Fe2+/L, the presence of ferrous is suggested to be negatively correlated with pyrite oxidation (McKibben & Barnes, 1986; Williamson & Rimstidt, 1994; Holmes & Crundwell, 2000).

The pyrite oxidation extent was found to decrease linearly with ferrous concentration as ferrous concentration was increased from 0 g/L to 10 g/L Fe2+. With an initial addition of 2.8 g/L ferrous, the pyrite oxidation extent dropped by approximately 3% from 42% to 39% while at 10 g/L initial ferrous addition, the oxidation extent was found to further decrease to 34%. Although not conclusive, the decrease in pyrite oxidation extent seems to occur at a decreasing rate from 10 g/L to 20 g/L ferrous addition.

Figure 43 the effect of ferrous concentration on ferric utilisation and pyrite oxidation at 220°C for 2 hours, 5g/L solid loading, 25g/L [Fe3+]

Table 16 Solution assay for the effect of ferrous on pyrite oxidation experiment

Feed Pyrite Feed Filtrates Pyrite Test Fe3+ Loading Fe2+ Oxidation # [Fe3+],g/L [Fe2+],g/L pH (g/L) (g/L) (g/L) (%) 21 26.4 5 2.80 5.0 14.5 0.55 39 22 25.9 5 5.13 4.9 13.9 0.50 38 23 23.7 5 10.2 4.5 12.5 0.47 34 24 25.4 5 20.1 4.6 10.8 0.38 29

The mechanism of pyrite oxidation rate retardation by ferrous is yet to be understood as there was not enough conclusive evidence found in this study. However, the following are the possible reasons for this retardation of pyrite oxidation by ferrous:

• From a molecular mechanism, competition of ferrous and ferric for adsorption on the reactive site on pyrite surface has been proposed in several reports (Garrels & Thompson, 1960; Zheng et al., 1986; Moses et al, 1991). Garrels & Thompson (1960) suggested that the adsorption of ferric and ferrous on pyrite surface is relative to their concentration in the solution. This study disagrees with this statement as ferric has a higher charge density than ferrous and therefore the adsorption of ferrous should be less preferable. However, Papangelakis et al. (1994) stated that the formation of neutral complex ferric species is favoured at high temperature, such as aqueous

FeHSO4SO4 being the predominant ferric species in acidic solution at 200°C. If ferric 68

is present as this large complexed ferric ion species which have lower positive charge density than ferrous, ferrous could be preferentially adsorbed on the pyrite surface (Moses et al, 1991; Zheng et al.,1986) • From electrochemical perspective, the presence of ferrous was noted to decrease the solution mixed potential and therefore slow down the rate of pyrite oxidation. (Holmes & Crundwell, 2000). • From an equilibrium perspective, ferrous is the product of the pyrite oxidation reaction. In the presence of ferrous, less pyrite would be oxidised before the system reaches equilibrium.

• Passivation of pyrite by FeSO4.H2O could be happening in this study as the system was unstirred. However, previous studies at low temperature would not have encountered this passivation problem and they still observed the retardation of pyrite oxidation in the presence of ferrous. Therefore, this passivation could contribute to the lowered pyrite oxidation but would not be the sole reason.

3.3.4 The combined effect of initial ferric concentration, solid loading and ferrous concentration on pyrite oxidation

While ferrous concentration was observed to adversely affect pyrite oxidation and ferric utilisation for oxidation, ferric concentration and/ or solid loading (i.e. surface area) were previously shown to have positive effects. Their combined effect on pyrite oxidation and ferric utilisation are summarised in Table 17, and presented in Figure 44.

Table 17 Solution assay for the combined effect of initial ferric concentration, solid loading and ferrous concentration

Fe3+ utilised 3+ 2+ Pyrite Pyrite Test [Fe ]init [Fe ]init for FeS2 Description Loading Oxidation ID oxidation g/L g/L g/L % % g Initial ferrous concentration = 0 g/L

Base 26.4 0 5 41% 22% 0.11 High ferric (Fe3+) 51.5 0 5 57% 16% 0.16

High surface area (SA) 25.5 0 10 22% 24% 0.12 69

Fe3+ utilised 3+ 2+ Pyrite Pyrite Test [Fe ]init [Fe ]init for FeS2 Description Loading Oxidation ID oxidation g/L g/L g/L % % g 3+ High Fe high SA 53.1 0 10 36% 19% 0.20 Initial ferrous concentration = 10.2 g/L 23 Base 23.7 10.2 5 34% 18% 0.095 25 High ferric (Fe3+) 51.5 10.2 5 41% 11% 0.11 26 High surface area (SA) 25.5 10.2 10 22% 22% 0.12 27 High Fe3+ high SA 50.9 10.2 10 38% 21% 0.21 Initial ferrous concentration = 20.1 g/L 24 Base 25.4 20.1 5 29% 14% 0.082 28 High ferric (Fe3+) 50.6 20.1 5 37% 10% 0.10 29 High surface area (SA) 27.9 20.1 10 20% 21% 0.11 30 High Fe3+ high SA 50.5 20.1 10 30% 15% 0.17

Figure 44 Ferric utilisation at various initial ferric concentration, ferrous concentration and solid loading (  = 5 g-pyrite/L, 25 g/L Fe3+ , ▲ = 5 g-pyrite/L, 50 g/L Fe3+, ◇ = 10 g- pyrite/L, 25 g/L Fe3+, O = 10 g-pyrite/L, 50 g/L Fe3+)

Increasing the ferric concentration was found to increase pyrite oxidation at both 10g/L and 20g/L ferrous concentration. At 10g/L Fe2+, it increased by 20% from 0.095 grams, while it went up by 26% from 0.082 grams at 20 g/L Fe2+. This translates to an increase of pyrite oxidation by 7% and 8% respectively. Although the increase in the oxidation extent of both ferrous concentration was pretty substantial, they were still lower than in the absence of ferrous.

Similarly, increasing the pyrite solid loading, which is equivalent to increasing surface area, was also found to increase ferric utilisation for oxidation at both ferrous concentration. In the presence of ferrous, the increased amount of ferric used for oxidation was found to be similar to that of increasing ferric concentration. At 10 g/L Fe2+, 0.12 grams of ferric were utilised for pyrite oxidation while 0.11 grams were used at 20 g/L Fe2+. This translates to an increase of ferric utilisation proportion to 22% and 21% at 10 and 20 g/L Fe2+ respectively. These values are only slightly less than that in the absence of ferrous.

By increasing the ferric concentration and surface area at the same time, it is expected that amount of ferric used for oxidation to increase significantly in the presence and absence of ferrous ion.

From these observations, it can be concluded that the effect of ferric concentration is significantly stronger than surface area in the absence of ferrous. However, in the presence of ferrous, higher surface area negated the effect of ferrous while the effect of ferric concentration was reduced at higher ferrous concentration.

3.3.5 Solid phase precipitates

From all experiments, the three major phases were precipitated in the solid phase - Hematite

(Fe2O3), basic ferric sulphate (FeOHSO4) and ferrous sulphate monohydrate (FeSO4.H2O). Table 18 summarises all phases identified in the solid residue and their approximated weight percent (excluding leftover pyrite).

Hematite was the main phase precipitated at lower initial ferric concentration. As the initial ferric concentration was increased i.e. increased in solution acidity, basic ferric sulphate started to co-precipitate with hematite. The formation of basic ferric sulphate seems to start at initial ferric concentration of approximately 20 g/L Fe3+. As the initial ferric concentration was increased, hematite became more unstable until it finally disappeared completely leaving only basic ferric sulphate as the major phase in the solid. Hematite seems to disappear completely at ferric concentration higher than approximately 34 g- Fe3+/L. These

71 concentration would be slightly different in the absence of pyrite as pyrite oxidation generates acid which would make the solution acidity higher (pH lower).

Table 18 Phase identification from the effect of initial ferric concentration and solid loading experiments Feed Test Pyrite Major phase and their mass fraction Fe3+, g/L Loading, g/L 0 0 5 - 1 8.8 Hematite (100%) 2 16.4 Hematite (100%) 3 21.2 Hematite (89%), BFS (11%) 4 27.0 Hematite (63%), BFS (27%) 2.5 5 30.7 Hematite (15%), BFS (85%) 6 35.1 BFS (100%) 7 40.4 BFS (100%) 8 52.7 BFS (100%) 9 2.7 Hematite (100%) 10 5.7 Hematite (100%) 11 14.5 Hematite (100%) 5 12 26.5 Hematite (68%), BFS (25%), Fe2SO4.H2O (1%)

13 40.4 BFS (98%), Fe2SO4.H2O (2%)

14 51.0 BFS (98%), Fe2SO4.H2O (2%) 15 13.5 10 Hematite (100%)

Feed Test Pyrite Major phase and their mass fraction Fe3+, g/L Loading, g/L

16 30.2 BFS (59%), Hematite (6%), Fe2SO4.H2O (35%)

17 34.1 BFS (66%), Fe2SO4.H2O (34%)

18 41.6 BFS (78%), Fe2SO4.H2O (22%)

19 48.7 BFS (86%), Fe2SO4.H2O (14%)

20 53.1 BFS (87%), Fe2SO4.H2O (13%)

A B

C D

E F

G H

Figure 45 SEM images from 2.5 g/L solid loading at (A) 16 g-Fe3+/L, (B) 27 g-Fe3+/L and (C) 35 g-Fe3+/L, 5g/L solid loading at (D) 14 g-Fe3+/L, (E) 27g-Fe3+/L and (F) 35 g-Fe3+/L and 10 g/L solid loading at (G) 14 g-Fe3+/L and (H) 51 g-Fe3+/L.

Figure 45 (A) to (C) show the SEM images from the 2.5 g/L solid loading experiment at three different initial ferric concentration. At 16.4 g/L initial ferric concentration (A), all solid were identified as hematite. The primary particles of the hematite seem to have a tendency to aggregate forming a secondary grape-like hemispherical masses. The size of the primary structure varied greatly from approximately 0.1 to 1 µm, while the secondary masses seemed to grow to approximately 10 µm. As the initial ferric concentration was increased to 27 g/L (B), a mixture of hematite and basic ferric sulphate were formed. Basic ferric sulphate (BFS) appears to be in a needle-like structure which looks the same as the one identified by Gomez et al., 2013 shown in Figure 10. The length and the thickness of the BFS particles varies from approximately 1 to 5 µm and 0.1 to 1 µm respectively. The aggregation of primary particles of hematite seems greatly reduced with a maximum size of 1 µm for the secondary masses. In the 35 g/L initial ferric concentration test (C), BFS was the only solid identified. The size and morphology of BFS particle seem to be similar to that in Figure B. Similar trend was observed in the 5g/L and 10 g/L solid loading experiments which are shown in (D) to (F) and (G) and (H) respectively.

Hematite, basic ferric sulphate and ferrous sulphate monohydrate solid phases were also identified in the effect of ferrous experiments. The proportion of FeSO4.H2O crystallised out of the solution was found to increase at higher initial ferrous concentration which was expected. The increase of FeSO4.H2O content in the solid was subtle with the addition of 2.8 g/L, 5.13 g/L to 10.2 g/L ferrous. Based on the QXRD analysis of the solid residues, it was found to be approximately 1%, 3% and 5% respectively. However, the proportion of

FeSO4.H2O was found to significantly increase at 20 g/L ferrous addition with ~28 wt%.

3.4 Implication of findings

3.4.1 Understanding confliction results of previous studies An attempt to understand the conflicting conclusion between the previous studies (King & Lewis, 1980; Gerlach et al., 1968) was done using the current study finding despite of the difference in experiment conditions such as temperature, stirring and residence time. Table 19 summarises the different conditions between the two previous studies. Table 19 Experiment conditions from previous studies on pyrite oxidation by ferric

3+ Pyrite [Fe ]addition, Temperature Residence No. Reference Loading (g/L) g/L (°C) time (hour) 1 Gerlach et al., (1966) 48 ~3 -11 100 1 2 King & Lewis (1980) 20 11 - 56 100 1

The difference in pyrite loading between the two studies was not considered in the analysis as the effect of ferric concentration was found to be stronger than the effect of surface area in the absence of ferrous in this study (Section 3.3.3). Additionally, given the higher pyrite loading in these two studies compared to current study, the amount of pyrite oxidised is expected to be linearly increase with initial ferric concentration.

As shown in Table 19, the range of ferric concentration added in these studies are vastly different where the ferric addition in the study by King & Lewis (1980) was up to 20 times higher than those tested by Gerlach et al. (1966). However, the results of the two studies at 2 M Fe3+ contradicted each other. This is likely to be caused by the difference in pyrite loading as discussed below. Note that 1L was used as the basis of the calculation.

King and Lewis (1980) recorded approximately 20% pyrite conversion in the absence of ferric and 28% in the presence of 0.2 M Fe3+, which is equivalent to 1.6 g additional pyrite oxidised. This value translates to an additional of 3.3% to the overall pyrite conversion in Gerlach et al. (1966). This difference is considered very small and within the experimental error and therefore the catalysing effect of ferric in Gerlach et al. (1966) was concluded to be insignificant.

3.4.2 Change of ferrous and ferric concentration at Lihir The solid loading at Lihir is approximated to be around 100 to 130 g/L of pyrite by assuming 50 wt% solid in the slurry, 15 wt% pyrite content and 1500 g/L slurry density. This solid loading much higher than those used in this study. Therefore, the pyrite oxidation is expected to be in the region where the oxidation extent is linearly dependent on ferric concentration as shown in Figure 39. This means any addition or increase of ferric concentration in the solution is expected to enhance the pyrite oxidation by ferric.

As discussed in Section 1.4.2, with the implementation of partial pressure oxidation at Lihir, there is a tendency for ferric concentration in the autoclave to increase by 3 g/L at maximum with the decrease of pyrite oxidation extent from 80% to 60%. This number was estimated based on the monthly average of ferric in Lihir plant data from December 2013 to December 2015. Additionally, there is also a possibility to increase the ferric concentration by recycling a portion of the autoclave discharge into the feed to accelerates the oxidation process as mentioned in Weir & Berezowsky (1986). The total concentration of iron in that stream at Lihir is likely to be approximately 5 g/L on average based on the same data set. Due to the high solid loading at Lihir solid loading and the low concentration of ferric in both streams, 76 the impact of this additional ferric on the overall pyrite oxidation would likely to be insignificant.

Similarly, the ferrous concentration in the autoclave discharge was reported to increase to approximately 5 g/L at maximum, as shown in Figure 46, due to the implementation of partial pressure oxidation. According to the finding of this study in Section 3.3.3, the increase of ferrous should not affect the pyrite oxidation by ferric due to the high solid loading at Lihir.

Figure 46 Monthly average of ferrous concentrations in Lihir autoclave discharge as a function of oxidation extent from plant data in December 2013 to December 2015

3.5 Summary

In this study, the potency of ferric as pyrite oxidant has been demonstrated despite of the removal of ferric from the solution via iron hydrolysis at high temperature. At the lower region of initial ferric concentration, the oxidation extent was found to increase linearly as a function of ferric concentration. As the ferric concentration was increased, the additional ferric has a tendency to precipitate due to the higher driving force for ferric to undergo hydrolysis. This means at higher ferric concentration region, majority of the ferric would tend to precipitate and only the residual ferric would be capable to oxidise the pyrite. Consequently, a plateau was observed where increasing ferric concentration barely increased the pyrite oxidation. Increasing the pyrite solid loading (i.e. surface area) was shown to delay the onset concentration at which the plateau was observed. This suggests that pyrite oxidation by ferric is surface limited.

In contrast, the presence of ferrous was found to be negatively correlated with pyrite oxidation. The mechanism for this ferrous effect has yet to be determined due to inconclusive evident found in this study. The negative effect of ferrous on pyrite oxidation was found to be negated by increasing the pyrite solid loading (i.e. surface area).

Therefore, the increase of ferrous concentration in Lihir autoclave due to the implementation of partial pressure oxidation should not affect the pyrite oxidation due to the high solid loading on site. The increase of ferric in the autoclave should increase the proportion of pyrite oxidised by ferric. However, due to the low concentration of ferric in Lihir autoclave, even with the addition of ferric into the autoclave feed, the overall oxidation in Lihir autoclave would likely still be governed by oxygen mass transfer in the autoclave.

Chapter 4 Solubility of hematite, basic ferric sulphate and potassium jarosite

4.1 Introduction

The development of accurate Eh-pH diagrams that include basic ferric sulphate and potassium jarosite stability fields is currently not possible. Reliable thermodynamic data at high temperature are, to a great extent, missing for both aqueous species and solid phases. These include the standard Gibbs free energy of formation used to calculate the reaction equilibrium constants.

Although basic ferric sulphate was identified and characterised decades ago (Johansson, 1962), the research into the detailed properties of this solid has only emerged recently. Consequently, not many studies have been done to investigate its thermodynamic properties. To date, there has only been one study by Majzlan et al. (2017) who made the first attempt to determine thermodynamic data of basic ferric sulphate as summarised in Table 20. Based on this thermodynamic data, the Gibbs energy of formation for basic ferric sulphate was estimated to be –1,047.2 kJ/mol at 220°C. This value, however, may not be accurate as the data is reported to be accurate within 1% only between 280 – 300 K. This means the Gibbs energy of formation of basic ferric sulphate at high temperature still cannot be accurately calculated for temperatures above 300 K.

Table 20 Basic ferric sulphate thermodynamic data (Majzlan et al. 2017) Parameter Value Unit

-1 ∆퐻°푓,25°퐶 -1160.2  2.3 kJ  mol

-1 -1 푆° 145.9  1.2 J  mol K

-1 ∆푆°푓,25°퐶 - 491.4  1.2 J  mol

-1 ∆퐺°푓,25°퐶 -1013.7  2.4 kJ  mol

-0.5 -2 -3 -1 -1 Cp = A + BT + CT * + DT , J  mol K A 575.9 B - 460.8 C -7,356,385 D 1,235,146,047

For potassium jarosite, its thermodynamic properties at 25°C were studied extensively in 1990s. However, there are significant discrepancies in the stability data of potassium jarosite data reported at high temperature from different sources listed in Table 21 (Stoffregen, 1993;

Stoffregen, 2000; Majzlan et al., 2010). Drouet & Navrotsky (2003) also reported potassium jarosite thermodynamic data. However, its high temperature thermodynamic properties could not be calculated as heat capacity data was not reported and therefore the data from this study will not be discussed herein.

Table 21. Gibbs energy of formation data for KFe3(OH)6(SO4)2 at elevated temperatures. All values are reported in kJ/mol Reference Stoffregen, 1993 Stoffregen, 2000 Majzlan et al., 2010 0 ∆G f, 150°C -3,388.3 -3,016.4 -3,146.6 0 ∆G f, 200°C -3,416.3 -2,933.8 -3,065.2 0 ∆G f, 220°C -3,426.9 -2,900.8 -3,032.6 0 ∆G f, 250°C -3,449.0 -2,851.5 -2,983.8

Note that the Gibbs energy of formation from Stroffregen (2000) was calculated using HSC Chemistry v7.1, while those from Mazlan et al. (2010) were calculated in HSC Chemistry v.9. The thermodynamic data are summarised in Table 22.

Table 22 Thermodynamic properties for potassium jarosite in HSC Chemistry

Reference Parameter Stoffregen, 2000 Majzlan et al., 2010

-1 ∆퐻°푓,25°퐶 (kJ  mol ) -3715.1 -3829.6

-1 -1 푆°25°퐶 (J  mol K ) 388.9 427.4 -1 ∆퐺°푓,25°퐶 (kJ  mol ) -3309.8 -3349.18 −3 -2 −5 2 −6 -1 -1 Cp = A + BT ∗ 10 + CT ∗ 10 + DT ∗ 10 , J mol K A 616.89 18.724 B 98.74 1769.145 C -199.6 -1.128 D - -1489.465

Tmin-Tmax (K) 298 - 300 100 - 350

Identical thermodynamic properties were used in Stoffregen, 1993 and HSC chemistry v7.1. The discrepancy between these two references arises from the methodology to calculate the standard molal Gibbs free energy at high temperatures. Majzlan et al. (2010) used calorimetry experiments to estimate the thermodynamic properties of potassium jarosite at 25°C while Stoffregen (2000) carried out solubility experiments at high temperature using hematite as the starting material. Although Stoffregen (2000) did the experiments at high

80 temperatures, the solution sampling was done at room temperature after quench cooling. In this way, true concentration of the solution at equilibrium would not be captured as equilibrium changes with temperature.

Additionally, the lack of information available on the equilibrium behaviour of aqueous system is attributed to the difficulties in assessing the solution environment directly at high temperature, especially for the ferric speciation. In-situ measurement at high pressure and high temperature is required to accurately measure aqueous speciation. However, no 3+ 2- measurement has been done to investigate Fe - SO4 - H2O system and therefore chemical speciation modelling has been the only tool used to predict the ferric aqueous speciation at elevated temperature. As a result, there are some uncertainties in the predominant form of ferric species at high temperature and their thermodynamic properties.

0 To date, Fe2(SO4)3 and FeHSO4SO4 are the two proposed ferric species in high acidity solution at temperatures between 200 and 250°C (Papangelakis et al., 1994; Liu et al., 2014). In this study, both ferric species are considered in the determination of their apparent thermodynamic data based on hematite solubility studies at 220°C. The experimental data were related back to chemical thermodynamic predictions using a series of assumptions. This approach permits a critical assessment of the existing thermodynamic data. Subsequently, the present study also investigated the solubility of basic ferric sulphate and potassium jarosite at 220°C to determine their thermodynamic data at 220°C using the same thermodynamic approach. With the newly determined thermodynamic data, the Eh-pH 3+ 2- 3+ + 2- diagram for Fe -SO4 -H2O and Fe -K -SO4 -H2O systems that include the three main iron precipitate phases at 220°C were presented.

Due to the acidic environment in the Lihir autoclave, a pH range from 0 to 4 was selected.

For the oxidation-reduction potential (Eh), a range of -0.5 to 2 was selected as this study 0 considered solid pyrite (FeS2), but not solid pyrrhotite (FeS) and iron (Fe ). Therefore, the very low Eh values are considered irrelevant for this system

4.2 Experimental design

4.2.1 Solubility experiment methodology Two different types of experiments were carried out to investigate the solubility of hematite, basic ferric sulphate and potassium jarosite. The first experiment was to precipitate pure phases of hematite, basic ferric sulphate and potassium jarosite without any adjustment of the solution pH. XRD analysis was used to confirm that only one crystalline phase detected

81 in the solid residues. All feed solutions were prepared in a volumetric flask before being transferred to the autoclave glass liner. Details on reagent purities are listed in Section 2.1.1.

For hematite precipitation, various concentration of ferric sulphate (added as

Fe2(SO4)3.xH2O) were mixed with DI water and stirred until all solids evidently dissolved. Note that no sulphuric acid was added into the feed solution.

For basic ferric sulphate precipitation, various concentrations of ferric sulphate (added as

Fe2(SO4)3.xH2O) and sulphuric acid (added as 98% sulphuric acid) were mixed with DI water and stirred until all solids evidently dissolved. High concentrations of acid, ranging between

30 and 50 g/L of H2SO4 were added to promote basic ferric sulphate precipitation.

For potassium jarosite, various concentrations of ferric sulphate (added as Fe2(SO4)3.xH2O), potassium sulphate (added as K2SO4) and sulphuric acid (added as 98% sulphuric acid) were mixed with DI water and stirred until all solids dissolved.

The second experiment involved regulating the acidity of initial solutions i.e. increasing the solution pH by adding lithium hydroxide (LiOH.H2O) into the feed solution. Lithium hydroxide was selected as lithium does not form an end-member jarosite-type compound and it theoretically should not affect hematite solution as it does not belong to the divalent cation group (Dutrizac & Jambor, 2000; Dutrizac, 2008). For hematite precipitation, no phase transition was expected with the addition of lithium hydroxide. For basic ferric sulphate and potassium jarosite, however, the concentration of lithium hydroxide added was increased until a phase transition from basic ferric sulphate or potassium jarosite to hematite was observed.

Feed solutions were prepared by mixing the required chemicals discussed above with nitrogen-sparged Deionised (DI) water in a volumetric flask. The solution was stirred until all solids have evidently dissolved. It was then sampled and stabilised in 2wt% HCl solution for assaying by ICP-OES and pH measurement was taken. The solution was then transferred into the autoclave liner ready for assembling. The solution was heated up to the target - temperature of 220°C. Similar to the microwave digester experiments, bisulphate (HSO4 ) will be formed as it becomes stable at high temperature during the heating period. Consequently, the “at temperature” pH increases due to the reduced activity of hydrogen ion (H+) in the solution.

Once the target temperature was reached, the experiment was allowed to run for 2 hours. At the end of the experiments, high temperature sampling was done to collect both solid and

82 solutions samples at high temperature. High temperature solid-liquid separation is required to obtain true concentration of the solution at equilibrium. This technique is still not commonly used due to the complexity in setting up the equipment and is only necessary when interested in the chemistry at temperature. Studies which used high temperature sampling technique typically withdraw solution sample via a porous frit filter which is installed in the dip tube to prevent solid leaving the aqueous system.

For the present investigation, a different technique was used where both solid and solution samples were collected at high temperature. This is because solid phase identification at equilibrium is crucial, especially for basic ferric sulphate which is prone to dissolve during cooling due to its instability at low temperature and acidic conditions. The high temperature sampling set-up which involves ex-situ pressure filtration is described in Section 2.2.3. The high temperature pH was estimated based on the low temperature pH measurement using the method described in Section 4.3.4.

Once the sampling finished. The autoclave heating was turned off and the autoclave setpoint was changed to 25°C to trigger the cooling water to run. Once the autoclave had reached approximately 40°C, the autoclave was disassembled and the remaining of the slurry in the liner was then vacuum filtered with 0.45 micron pore-size PTFE membrane filters. Solids were washed with DI water (1:1 ratio) to displace entrained solution and were oven dried at 60°C, typically for 24 hours.

All solids were then characterised using SEM-EDS and XRD for particle morphology and phase identification. Filtrate was sampled and stabilised in 2wt% HCl solution for assaying by ICP-OES and pH measurement was taken.

4.2.2 Residence time – equilibrium determination

Eh-pH diagrams are equilibrium diagrams and therefore it is imperative set a residence time where the system has reached equilibrium. Residence time was counted once the system reached the target temperature of 220°C. It generally took approximately 90 minutes to heat the solution up to the target temperature. To determine the appropriate residence time, a kinetic study on hematite solubility was done with ferric sulphate feed solution containing 13 g/L Fe3+. The sampling times tested were 2, 4, 6 and 24 hours.

Each residence time was tested in a different experimental run as high temperature sampling can only be taken one time for each experiment due to the large amount of samples drawn. Consequently, a variation of the ferric concentration in the feed solution for 83 each experiment was present but very minimal. The ferric concentrations in the feed solution and final filtrate were analysed using ICP-OES, summarised in Table 23.

Table 23 Kinetic test results of 13 g/L Fe3+ feed solution at 220°C

[Fe3+] in feed solution, g/L Residence time, hours [Fe3+] in final filtrate, g/L 12.80 2 1.17 12.60 4 1.15 12.75 6 1.12 12.75 24 1.11

The kinetic of ferric hydrolysis at 220°C were found to be relatively fast where the ferric concentration in solution dropped from approximately 13 g/L to 1.2 g/L in within 2 hours residence time. There was a downward trend in final ferric concentration from 2 to 24 hours residence time, but they were relatively similar with approximately 5% difference which is within the error associated with sampling, dilution and ICP-OES assay.

The residence time at Lihir autoclave varies from 1 to 2 hours depending on the autoclave feed rate. Umetsu et al. (1977) observed a retardation of ferric hydrolysis for the 13 g/L initial ferric concentration at 200°C in one hour residence time. However, this may not be observed at 220°C as Induction time is shortened by the increase of temperature. For example, the induction time of a ferric sulphate solution containing 25 g/L Fe3+ was reduced from 14 hours to less than 3 hours when the temperature was increased from 185 to 200°C (Umetsu et al., 1977). Therefore, the Lihir autoclave should approach equilibrium in one hour and an equilibrium model is therefore relevant. For this study two hours residence time was selected to be residence time for all solubility experiments to ensure that equilibrium has been reached.

4.3 Chemical thermodynamic approach

4.3.1 Thermodynamic calculation The equilibrium constant (K) of each reaction was calculated using Equation 15.

∆Grxn = ∆Grxn° + RTln(K) Equation 15

At equilibrium, Gibbs energy of reaction (∆Grxn) is equal to 0 and therefore the standard

Gibbs energy of reaction (∆G°rxn) can be calculated using Equation 16.

∆Grxn° = −RTln(Keq) Equation 16

The standard Gibbs energy of reaction can be calculated based on the standard Gibb energy of formation of each species involved in the reaction according to Equation 17.

Equation 17 ∆퐺푟푥푛° = ∑ ∆퐺f,product° − ∑ ∆퐺f,reactant°

For general Reaction 4.1 the Gibbs energy of reaction can be calculated using Equation 18.

푎퐴(푎푞) + 푏퐵(푎푞) = 푐퐶(푠) + 푑퐷(푎푞) Reaction 4.1

∆푮풓풙풏° = ∆푮풇,푪° + ∆푮풇,푫° − ∆푮풇,푨° − ∆푮풇,푩° Equation 18

The equilibrium constant of each reaction can be related to the activities of each species using Equation 19.

푎 푏 푎퐴 ∗ 푎퐵 Equation 19 Keq = 푐 푑 푎퐶 ∗ 푎퐷

The activity of species i (푎푖) can be estimated from the species concentration (푚푖) using the aqueous molar activity coefficient (훾푖) as shown in Equation 7. The concentration was defined as molar concentration (mol/kg-H2O)

풂풊 = 후퐢. 퐦퐢 Equation 20

To get the species distribution in the solution, the activity of each species has to be defined. However, the activity coefficient of these species are unknown and have to be determined using activity coefficient model built-in HSC Chemistry v9 which will be further discussed in - Section 4.3.2. The activity coefficient of relevant ionic aqueous species such as HSO4 (a), K+(a), H+(a) were determined using this method. For solid phases and neutral complex species, the activity coefficients were assumed to be unity.

4.3.2 Activity coefficient estimation Activity coefficients are used to account for the deviation of each species from ideal behaviour in the solution. There are many different activity coefficient models available for ionic species, such as the Davies, B-dot, Bromley-Zemaitis, Pitzer, Guggenheim, Bromley, Extended UNIQUAC, Mixed Solvent Electrolyte (MSE) equations.

The current investigation uses HSC Chemistry V.9 which uses a Davies equation to estimate for the activity coefficient values based on the solution concentration. Figure 47 shows the input sheet in aqua module imbedded in HSC Chemistry v.9. All solute species, either ionic or molecular formula, are entered in the ‘water species data’ column, along with their

85 respective concentration obtained from ICP-OES. The concentration entered is based on 1 kg of water which is equivalent to 55.5 kmol of H2O.

Solution temperature

List of aqueous species Species concentration

Figure 47 Aqueous module in HSC Chemistry v9

Pitzer equations have been recognised to provide a versatile theoretically and realistic prediction for aqueous solution modelling when these models are available for the correct system. Pitzer models are available in the HSC aqua module as semi-empirical Pitzer model (with binary interaction only) and Harvie’s modification of the Pitzer model (with binary and ternary parameters). However, the information on binary ion pairs parameters, cation-cation, anion-anion and ternary ion interaction parameters for ferric ion are not currently available yet. In the absence of these interaction parameters, the prediction in HSC aqua module would be based on the extended Debye-Huckel which is shown in Equation 21.

√퐼 2 ln 훾푖 = −퐴ϕ ( + ln(1 + 푏√퐼) Equation 21 1 + 푏√퐼 푏

Where, 퐴휙 and 푏 are the temperature dependent constants and 퐼 is the ionic strength of the solution.

- 2- Despite these limitations, the H2SO4 - HSO4 - SO4 system is well defined, even at high - + temperature and therefore the activity coefficient for HSO4 and H determined using HSC aqua module should still be reliable. Figure 48 shows the trend for activity coefficient for - HSO4 species in solution at pH 0.8 for different temperatures ranging from 50°C to 220°C. - The general trend remains constant where the activity coefficient of HSO4 is lower at higher temperature and more concentration solution.

- Figure 48 Activity coefficient of HSO4 species as a function of concentration and temperature generated using HSC Chemistry v.9 at [H+] = 0.158 and [OH-] = 6.3E-14

4.3.3 Solid and Aqueous species selection

In the Fe-S-K-H2O system, iron (ferric and ferrous), sulphur and potassium could be distributed either as soluble species or as precipitates. The soluble species could be in the form of simple cations or complexes which could be charged or neutral. Depending on the solutions conditions, some species forms are more dominant than the others. In present study, all possible species were collated from published journal papers and HSC Chemistry program. Each of these species were critically analysed and only the predominant species were selected to simplify the construction of the Eh-pH diagram in this study.

4.3.3.1 Solid phases

Hematite (Fe2O3), basic ferric sulphate (FeOHSO4) and potassium jarosite

(KFe3(SO4)2(OH)6) were included in present study as the possible iron precipitates in autoclave. Ferrous sulphate monohydrate (FeSO4.H2O) or also known as szomonolkite was considered as this solid will form in the presence of high concentration of ferrous at the autoclave operating temperature. Lastly, pyrite (FeS2) was also included as it is the main sulphide mineral present in Lihir ore.

0 Other solids such as goethite (FeOOH), pyrrhotite (FeS), iron (Fe ) and magnetite (Fe3O4) were not considered as they are not likely to form under the Eh and pH range of interest in present study.

4.3.3.2 Sulphur species - The five sulphur species that may be present at high temperature system are HSO4 (a), 2- 0 - SO4 (a), S (s), H2S (g) and HS (a). The equilibrium reactions for the S-H2O system are tabulated in Table 24.

Table 24 Equilibrium reactions for S-H2O system Equilibrium Reactions 2- + - SO4 (a) + H (a) ↔ HSO4 (a) 2- + - SO4 (a) + 10H (a) + 8e ↔ H2S (a) + 4H2O 2- + - - SO4 (a) + 9H (a) + 8e ↔ HS (a) + 4H2O 2- + - 0 SO4 (a) + 8H (a) + 6e ↔ S (s) + 4H2O - + - HSO4 (a) + 9H (a) + 8e ↔ H2S (a) + 4H2O - + - - HSO4 (a) + 8H (a) + 8e ↔ HS (a) + 4H2O - + - 0 HSO4 (a) + 7H (a) + 6e ↔ S (s) + 4H2O - + H2S (a) ↔ HS (a) + H (a) 0 + - H2S (a) ↔ S (s) + 2H (a) + 2e - 0 + - HS (a) ↔ S (s) + H (a) + 2e

Figure 49 shows the Eh-pH diagram of the S-H2O system at 220°C under autoclave - operating condition. Within the limit of Eh and pH of interest (highlighted in blue), HSO4 (a) is the predominant sulphur species. The stability of sulphur solid (S0) is almost non-existent and therefore is not considered in present study.

Area of interest

Figure 49 Eh-pH diagram for sulphur at 220°C, total P = 32bar, ∑S = 0.224M and ionic strength = 0.28. Pressure of H2S (g) was set at 0.1 bar. 88

4.3.3.3 Ferric species 3+ 2- Ferric speciation in Fe -SO4 -H2O system has been reported over a wide range of temperatures, however, most of the reports are limited to lower temperature systems (25 – 150°C) (Casas et al., 2005; Cifuentes et al., 2006; Filippou et al., 1995; Sapieszkoe et al., 1977; Stipp, 1990; Yue et al., 2014). Only two studies considered the higher temperature range relevant to pressure oxidation (Liu et al., 2003; Papangelakis et al., 1994). Table 25 summarises the aqueous ferric species reported in literature from 25°C to 250°C.

2+ 3+ 2- Table 25. Aqueous speciation for Fe -Fe - SO4 -H2O system at various temperatures Temperature Ferric Species Reference + FeSO4 1, 2, 3, 4, 7 - Fe(SO4)2 1, 2, 3, 7 Fe3+ 1, 2, 4, 7 2+ FeHSO4 1, 2, 4 25°C 4+ Fe2(OH)2 1, 4 FeOH2+ 3, 4 + Fe(OH)2 3 0 FeH(SO4)2 6, 7 + FeSO4 1, 2 - Fe(SO4)2 1 100°C Fe3+ 1 2+ FeHSO4 1, 2 4+ Fe2(OH)2 1 + FeSO4 2, 8 - Fe(SO4)2 8 2+ FeHSO4 2, 8 0 150°C FeSO4HSO4 8 Fe3+ 8 FeOH2+ 8 + Fe(OH)2 8 + FeSO4 8 2+ FeHSO4 8 0 FeSO4HSO4 8 180°C Fe3+ 8 FeOH2+ 8 + Fe(OH)2 8 + FeSO4 8 200°C 2+ FeHSO4 8

Temperature Ferric Species Reference - Fe(OH)2SO4 8 FeOH2+ 8 + Fe(OH)2 8 0 Fe(OH)3 8 + FeSO4 5 2+ FeHSO4 5 - Fe(SO4)2 5 250°C Fe3+ 5 - Fe(OH)2SO4 5 0 Fe2(SO4)3 5 1Filippou et al., 1995 2 Yue et al., 2014 3 Stipp, 1990 4Sapieszkoe et al., 1977 5Liu et al., 2003 6Casas et al., 2005 7Cifuentes et al. 2006 8 Papangelakis et al., 1994

Based on a high temperature hematite solubility study, Papangelakis et al (1994) proposed 2+ 0 that FeHSO4 and FeSO4HSO4 are the most dominant species at 200°C with the latter become more dominant with increasing acidity and the former is more favoured with increasing temperature as shown in Figure 50.

Figure 50 Ferric aqueous species distribution of at equilibrium with hematite at 200°C (Papangelakis, 1994)

However, in a similar study at higher temperature from 230 to 270°C, Liu et al (2003) found 0 that their thermodynamic prediction failed even with the inclusion of FeSO4HSO4 and 90

+ - 0 Fe(HSO4)2 species. They suggested that Fe(OH)2SO4 and Fe2(SO4)3 are the two predominant species in the range of 230 – 270°C with the latter being the more dominant species at high acidity as shown in Figure 51.

Figure 51 Ferric aqueous species distribution at equilibrium with hematite at 250°C (Liu et al., 2003)

These two studies used different methodologies and thermodynamic data in their modelling work which could contribute to the different outcomes. Despite the differences, both studies suggested a formation of neutral complex species, which is rather controversial, in the high acidity region. Papangelakis (1994) explained that the formation neutral complex species is favoured at high temperature due to the reduced ability of water to stabilise highly charge species via solvation as the water dielectric constant decreases. As a result, these neutrally charged ferric species that are not detectable at room temperature may very well form at an elevated temperature.

This demonstrates the complexity of this system at high temperature and the limited insight that has been gained on ferric speciation at high temperature to date. Direct measurement of ferric speciation at high temperature is currently not possible and therefore the true distribution of ferric species in aqueous system at high temperature remains unknown. The 0 present study considered both ferric species, FeSO4HSO4 and Fe2(SO4)3 to be the dominant ferric species in present study due to the acidic nature of the system investigated.

4.3.3.4 Ferrous species

Ferrous speciation has also been studied extensively over a wide range of temperatures but all studies were limited to lower temperature systems (25 – 150°C) as summarised in Table 26 (Casas et al., 2005; Cifuentes et al., 2006; Filippou et al., 1995; Stipp, 1990; Yue et al., 2014). 2+ 2- Table 26 Ferrous speciation in Fe -SO4 -H2O system at 25, 100 and 150°C Ferrous Temperature Reference Species Fe2+ 1, 2, 4, 5 0 25°C FeSO4 1, 2, 3, 4, 5 + FeHSO4 1, 5 Fe2+ 1, 2 0 FeSO4 2 100°C + FeHSO4 1, 2 FeOH+ 1 Fe2+ 2 0 150°C FeSO4 2 + FeHSO4 2 1Filippou et al., 1995 2 Yue et al., 2014 3 Stipp, 1990 4Casas et al., 2005 5Cifuentes et al. 2006

Ferrous speciation at temperature above 150°C was not found in the literature reviewed. At 150°C, however, approximately 70% of ferrous was reported to exist in the form of free 2+ + ferrous (Fe ) while the remaining was in the form of FeHSO4 (Yue 2005). This trend is assumed to remain unchanged at 220°C and thereby Fe2+ was the only ferrous species considered in present study.

4.3.3.5 Potassium species

+ - 0 Four potassium aqueous species, K (a), KSO4 (a), KHSO4 (a) and KOH (a), were apparent in the HSC Chemistry software database. The aqueous distributions of these species were estimated using the methodology explained in Section 4.3.1 by considering the speciation 2- - of sulphur species which are SO4 , HSO4 and H2SO4. The activity coefficient of all species were assumed to be unity. Table 27 lists the thermodynamic data collected from HSC Chemistry v.9 for all the species involved. The reactions and the calculated thermodynamic data of each reaction at 220°C are listed in Table 28.

Table 27 Thermodynamic data at 220°C for potassium aqueous species modelling

Species ΔG°f at 220°C (kJ/mol)

H2O (l) -207.02 H+ (a) 0

H2SO4 (a) -609.32 - HSO4 (a) -662.78 2- SO4 (a) -616.33 K+ (a) -301.039 - KSO4 (a) -937.77

KHSO4 (a) -951.80 KOH (a) -434.78

Table 28 K-SO4-H2O reactions and their respective Gibbs free energy at 220°C

Reactions ∆푮°푹,ퟐퟐퟎ°푪 Log K + -2 - K (a) + SO4 (a) ↔ KSO4 (a) -20.4 2.16 + 0 K (a) + HSO4(-a) ↔ KHSO4 (a) 12.8 -1.27 + + K (a) + H2O(l) ↔ KOH (a) + H (a) 73.3 -7.76 -2 + - SO4 (a) + H (a) ↔ HSO4 (a) -46.5 4.92 - + H2SO4 (a) ↔ HSO4 (a) + H (a) -49.2 5.27

Figure 52 shows the potassium aqueous species distribution at total [S] of 0.2 m and total + [K] of 0.05 m. K is shown to be the absolute dominant species at higher acidity. The pH at + - which the cross over between K and KSO4 occurs varies depending on the total sulphur - + concentration. KHSO4 is almost non-existence in all pH range investigated. Therefore, K is the only potassium species considered.

Figure 52 Potassium aqueous speciation at 220°C, total [S] = 0.2 m and total [K] = 0.05 m

4.3.4 pH estimation at high temperature pH measurement was done at low temperature, approximately 30°C. This section outlines the methodology to estimate the high temperature pH based on the low temperature 2- - measurement. For simplification, only the equilibrium reaction between SO4 and HSO4 is considered and the activity coefficient of all species were assumed to be unity. Table 29 summarises the equilibrium constant value for this reaction at 30°C and 220°C. These values were taken from HSC Chemistry v.9.

2- - Table 29 Equilibrium constant for SO4 -HSO4 reaction at 30 and 220°C

Reaction considered Keq at 30°C Keq at 220°C -2 + - SO4 (a) + H (a) ↔ HSO4 (a) 96.92 83320

The total S concentration obtained from ICP-OES is independent of temperature. For estimation purposes, sulphur complexation with other elements is assumed to be insignificant and therefore the total sulphur in the system is defined as the sum of bisulphate - 2- (HSO4 ) and sulphate (SO4 ) ions according to Equation 22.

푚 = 푚 − + 푚 2− Equation 22 푡표푡푎푙 푆 퐻푆푂4 푆푂4

− At 30°C, 푚퐻푆푂4 can be estimated using equilibrium constant expression in Equation 23 by substituting Equation 22. 푚퐻+ was calculated based on pH measurement value.

− 푚퐻푆푂4 퐾푒푞 = Equation 23 푚 2− 푚 + 푆푂4 퐻

To estimate the pH at 220°c, 푚퐻+was calculated using Equation 23 by substituting Equation 22 and Equation 24. Equation 24 defines the proton balance in the system assuming that total proton is equal to the sum of free H+ and H+ associated with bisulphate.

+ − + 푚푡표푡푎푙 퐻 = 푚퐻푆푂4 + 푚푓푟푒푒 퐻 Equation 24

The equation becomes a quadratic equation shown in Equation 24.

푚푡표푡푎푙 퐻+ − 푚푓푟푒푒 퐻+,220°퐶 퐾푒푞,220°퐶 = Equation 25 (푚푇표푡푎푙 푆 − 푚푡표푡푎푙 퐻+ + 푚푓푟푒푒 퐻+,220°퐶) 푚푓푟푒푒 퐻+,220°퐶

푚푓푟푒푒 퐻+,220°퐶 was solved in Excel to give a value of 83320 which is the 퐾푒푞at 220°C. Sample calculations can be found in Appendix C.1.

4.3.5 95% confidence interval methodology The 95% confidence interval in the solubility experiments (Section 4.5, Section 4.6 and Section 4.7) was estimated based on a two-tailed t-test using Equation 8. s E95% = ±tα/2,DOF × Equation 26 √n

Where 푡훼/2,퐷푂퐹 is critical t value, s is standard deviation and n is the number of sample.

The critical t-value was obtained directly via Excel using the T.INV(/2, DOF) formula for a two-tailed test. The degree of freedom (DOF) is equal to the number of samples minus the number of estimates (i.e. DOF = n-1). For 10 samples, the t-critical value was calculated to be ± 2.23.

The standard deviation was calculated based on the standard deviation of the equilibrium constant and Gibbs energy of formation from all experiments. The calculation was done directly via excel using the STDEV.S command. The 95% confidence values ( E95% ) was then added to the mean equilibrium constant to plot the 95% confidence interval line on the graph using Equation 19, and also to estimate the uncertainty in the solid Gibbs free energy of formation using Equation 16 and Equation 17. Sample calculation can be found in Appendix C.2.

4.3.6 High temperature thermodynamic data review Most of the necessary thermodynamic data were collected from HSC Chemistry v.9 which utilises the Criss-Coble entropy correspondence principle for extrapolation, while a number 95 of the required thermodynamic data were estimated from some scientific publications. Table 30 summarises thermodynamic data for each species from various different sources. Some of the Gibbs free energy of formation values were obtained by interpolating and extrapolating from the data given in the literatures. For example, the Gibbs free energy of formation for 0 Fe2(SO4)3 species obtained from Liu et al. (2003) was interpolated from data given at 230°C, 250°C and 270°C.

The Gibbs free energy of formation at 220°C (G°f 220°C) was estimated for those species for which heat capacity data were available using Equation 27 below.

푻 푻 ∆푪풑 Equation 27 ∆푮°풇,푻 = ∆푮°풇,ퟐퟗퟖ푲 + ∫ ∆푪°풑풅푻 − 푻 ∫ 풅푻 − (푻 − ퟐퟗퟖ)∆푺°ퟐퟗퟖ푲 ퟐퟗퟖ ퟐퟗퟖ 푻 It is important to note that there are variations in the thermodynamic properties for each species, not only at elevated temperatures but also at 25°C, as shown in Table 30. Therefore, the present investigation uses thermodynamic data from HSC Chemistry v.9 to maintain consistency throughout the study. Table 31 summarises the thermodynamic data set at 220°C used in present investigation on solubility.

Table 30 Summary of Thermodynamic properties from various references

° G°f 25°C H°f, 25°C S°f,25°C Cp 25°C Cp (J/K.mol) G°f 220°C Species Ref. kJ/mol kJ/mol J/mol J/K/mol A B X 10-3 C X 105 D X 10-6 kJ/mol

Fe2O3 (s) -744.27 -825.40 -272.11 - 98.3 77.82 -14.85 - 1 -697.3

Fe2O3 (s) -742.83 -824.78 -274.88 104.04 143.6 -36.32 -31.43 71.79 4 -690.1

Fe2O3 (s) -742.20 -824.20 -275.03 103.85 - - - - 2 -

Fe2O3 (s) -742.20 -824.20 -275.03 - 143.6 -36.323 -31.34 71.79 3 -694.6

HSO4 (-a) -756.01 -887.01 -439.38 - -547.3 1342.1 266.78 1 -678.3

HSO4 (-a) -755.31 -886.90 -441.35 22.586 -1734.8 7506.4 305.08 -9183.63 4 -667.9

HSO4 (-a) -756.01 ------3 -668.5

HSO4 (-a) -755.76 -889.1 -447.22 22.18 - - - - 2 -

KFe3(SO4)2(OH)6 (s) -3349.2 -3829.6 -1611.3 412.52 18.724 1769.14 -1.128 -1489.47 4 -3032.6

KFe3(SO4)2(OH)6 (s) -3309.8 -3715.1 -1359.3 616.89 98.74 -199.6 0 6 -3070.7

H2O (a) -237.18 -285.85 -163.24 - 75.44 - - - 1 -209.4

H2O (a) -237.18 -285.85 -163.24 75.35 - - - - 2 -

H2O (a) -237.14 -285.83 -163.30 75.229 186.88 -464.24 -19.565 548.631 4 -207.0

H2O (a) -237.18 - - - 186.88 -464.24 -19.565 548.631 3 -241.2 Fe (+2a) -91.2 ------3 - Fe (+2a) -91.2 -92.5 -4.4 - - - - - 7 - Fe (+2a) -91.8 -92.7 -3.0 - - - - - 7 - Fe (+2a) -91.2 ------7 - Fe (+2a) -91.5 -92.4 -3.0 - - - - - 7 - Fe (+2a) -88.9 ------7 - Fe (+2a) -78.9 -89.1 -34.3 -32.3 -1943.8 8530.47 293.75 -10685.0 4 -70.05

° G°f 25°C H°f, 25°C S°f,25°C Cp 25°C Cp (J/K.mol) G°f 220°C Species Ref. kJ/mol kJ/mol J/mol J/K/mol A B X 10-3 C X 105 D X 10-6 kJ/mol Fe (+3a) -4.6 -48.5 -147.24 - 79.09 -219.35 15.4 1 23.4 Fe (+3a) -17.2 -49.58 -108.66 -76.642 -3668.3 15752.4 580.73 -19551.2 4 8.6 Fe (+3a) -16.7 ------3 - Fe (+3a) -17.2 -49.58 -108.47 -142.67 - - - - 2 46.0

SO4 (-2a) -744.63 ------3 -

SO4 (-2a) -744.63 -909.18 -551.90 - 874.6 -1759.7 - - 1 -650.0

SO4 (-2a) -744.46 -909.6 -553.88 -269.37 - - - - 2 -735.4

SO4 (-2a) -744.36 -909.60 -554.22 -264.96 -3610.8 16366.3 364.89 -21432.5 4 -616.3

Fe2(SO4)3 (a) -2243 -2825 -1952.0 -2243 - - - - 1 -

Fe2(SO4)3 (a) ------2 -2341.3

Fe2(SO4)3 (a) -2242.9 2825.33 -1953.2 - - - - - 4 -

FeSO4.H2O (s) - -1245.7 - - 55.293 0.2798 0 0 7 -

FeSO4.H2O (s) - -1241.7 ------9 -

FeSO4.H2O (s) - -1244.3 ------11 -

FeSO4.H2O (s) - -1243.7 ------12 -

FeSO4.H2O (s) - -1243.5 ------8 -

FeSO4.H2O (s) -1079.3 -1241.8 -545.17 139.519 27.238 376.61 -0.013 0.096 4 -972.7 1Papangelakis et al 1994 2Liu et al., 2003 3Yue et al., 2014 4HSC Chemistry v.9 5Majzlan et al., 2010 6Stoffregen et al., 2000 7 Barin et al., 1973 8Kobilyn et al., 20119Kobilyn et al., 2007 10Sippola, 1992 11DeKock, 1982 12Hemingwat et al. 2002

Table 31 Thermodynamic data set at 220°C used sourced from HSC Chemistry v.9

Species ΔG°f at 220°C (kJ/mol)

H2O (l) -207.02 H+ (a) 0 - HSO4 (a) -662.78 Fe2+(a) -70.05 K+(a) -301.039

Fe2O3 (s) -690.89

FeS2 (s) -150.73

FeSO4.H2O (s) -972.66

4.4 Assumptions

This section outlines the assumption made throughout the current investigation to simplify the methodology. The assumptions made were: • Equilibrium was achieved after 2 hours (Section 4.2.2);

• The activity of water (푎퐻2푂) was assumed to be unity; • The activity of all solids, including hematite, potassium jarosite and basic ferric sulphate, were assumed to be unity;

• All ferric ions were assumed to be in the form of either aqueous neutral Fe2(SO4)3 or

FeHSO4SO4 at high temperature and their activity coefficient were assumed to be unity;

− • The molality of bisulphate ion (푚퐻푆푂4 ) was estimated using the total sulphur measured by ICP-OES. The total sulphur in the system was assumed to be consisted of free bisulphate ion and sulphate and/or bisulphate associated with ferric ion. • Solution molarity (mol L-solution-1) was assumed to be similar to the solution molality (mol kg-water-1). This was confirmed during the data analysis using solution density and total filtrate volume.

4.5 Solubility of hematite at 220°C

4.5.1 Solid identification and morphology Hematite was precipitated from various initial feed conditions as described in Section 4.2.1. Based on XRD result, hematite was established to be the only phase precipitated during the experiment. The addition of lithium hydroxide to adjust the solution pH did not instigate the

99 precipitation of lithium jarosite. XRD results from hematite solubility study are summarised in Appendix A.2.

The presence of lithium ion, however, seems to influence the morphology of hematite particle. It is unlikely for lithium ion to be incorporated into the crystal lattice of hematite due to the different charge density between ferric (Fe3+) and lithium (Li+) ions.

Figure 53 depicts morphologies of hematite precipitated as lithium concentration increases. The addition of lithium ions has a great influence, not only on the structure, but also on the size of the hematite particles. When lithium was not involved in the ferric hydrolysis reaction, the primary particles tend to aggregate forming a secondary grape-like hemispherical masses. The structure of the primary particles looks like an elongated sphere with a size of approximately 1 -2 µm in length. As 1 g/L of lithium was introduced, there was no noticeable difference observed but as the lithium concentration was increased to 5 g/L, the tendency for the particles to aggregate seem to become much less. The primary particles appear to be more uniform and smooth, but the size seems to decrease to approximately 0.3-0.5 µm. With further addition of lithium at 10 g/L, the particle size abruptly decreased to approximately 0.1 µm and the primary crystal seems to become sphere-like structure.

100

A B

C D

E F

Figure 53 Pyrite crystal cluster precipitated from 12 g/L Fe3+ solution with no addition of LiOH (A), and 17 g/L Fe3+ solution with no addition of LiOH (B), 13 g/L Fe3+ solution with addition of 1 g/L LiOH (C and D), 5 g/L LiOH (E) and 10 g/L LiOH (F). The magnification of A is x4000, B is x5500, C is x5,000 and D, E, F are x20000.

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4.5.2 Solution assay

0 Two possible ferric species, aqueous FeHSO4SO4 and aqueous Fe2(SO4)3 were considered in the hematite solubility investigation. For each ferric complex species, the aqueous species was assumed to be the predominant ferric species and was in equilibrium with hematite (Fe2O3). The Gibbs energy of formation for both aqueous species were estimated by fitting the equilibrium line to the measured final aqueous iron concentration and the calculated pH at 220°C tabulated in Table 32. The ferric, total sulphur and lithium concentration in feed were obtained from ICP-OES. The ferric and total sulphur concentration in filtrate was calculated based on the concentration obtained from ICP-OES, filtrate mass and filtrate density. The activity coefficients were estimated using HSC Chemistry as mentioned in Section 4.3.2. High temperature pH was estimated according to the methodology in Section 4.3.4.

Table 32 Hematite solubility experiment result

Feed, g/L Filtrate, g/kg-H2O pH,220°C + − pH,25°C 훾퐻 훾퐻푆푂4 Fe3+ S Li+ Fe3+ S Li+ estimated measured 1 6.43 5.9 - 0.23 5.80 - 0.79 0.81 0.58 0.58 2 9.22 8.3 - 0.50 8.18 - 0.70 0.72 0.54 0.55 3 11.0 9.9 - 1.07 9.95 - 0.59 0.61 0.52 0.52 4 12.1 11.8 - 1.16 10.9 - 0.58 0.60 0.50 0.51 5 17.2 15.4 - 2.08 15.4 - 0.55 0.57 0.45 0.48 6 12.9 11.8 0.98 0.36 10.1 0.94 0.90 0.98 0.52 0.57 7 13.0 12.0 9.82 0.12 11.7 9.74 1.02 1.17 0.54 0.61 8 11.1 9.9 4.89 0.14 9.4 4.85 0.96 1.04 0.55 0.60 9 13.0 12.1 13.7 0.03 11.8 13.8 1.10 1.32 0.53 0.61 10 13.0 12.2 0.96 0.41 11.6 0.96 0.92 1.00 0.62 0.61

0 4.5.2.1 Fitting of FeHSO4SO4 aqueous species

The equilibrium reaction between aqueous neutral FeHSO4SO4 and hematite is described in Reaction 4.2.

0 - + 2FeHSO4SO4 (a) + 3H2O(l) ↔ Fe2O3(s) + 4HSO4 (a) + 4H (a) (4.2)

For each of the experiment data, the equilibrium constant 퐾푒푞was calculated using Equation 28. All ferric species were assumed to be in the form of aqueous neutral complex

FeHSO4SO4 and therefore its molality can be estimated using the ferric concentration 102

3+ measured by ICP-OES where 푚퐹푒 ,25℃ ≈ 푚퐹푒퐻푆푂4푆푂4. The activity coefficient of aqueous

- − FeHSO4SO4 was assumed to unity. The concentration of HSO4 (푚퐻푆푂4 ) was calculated by the difference in Total sulphur and sulphur and/or bisulphate associated by ferric. For

FeHSO4SO4 species, the sulphur concentration incorporated in the aqueous complex is

- − + + twice the concentration of ferric ion. The activity coefficient of HSO4 (훾퐻푆푂4 ) and H (훾퐻 ) tabulated in Table 32 were used.

4 4 푎퐻푆푂4− 푎퐻+ 퐾퐹푒2푂3 = 2 Equation 28 푎퐹푒퐻푆푂4푆푂40

The 퐾푒푞 results for each experiment data are summarised in Table 33. Based on 95% confidence interval, the 퐾푒푞 was estimated to be -3.2  0.2.

Table 33 Activity and Keq values for hematite solubility

log(퐾푒푞) Log 푎퐹푒퐻푆푂4푆푂4 Adjusted pH 1 -3.41 -2.18 1.34 2 -3.37 -2.66 1.17 3 -3.46 -2.60 1.02 4 -3.36 -3.21 0.98 5 -3.34 -2.13 0.85 6 -3.43 -2.38 1.25 7 -2.98 -2.04 1.38 8 -3.08 -1.72 1.37 9 -2.55 -1.43 1.54 10 -3.43 -1.68 1.18

- − There were variations in the concentration of HSO4 (푚퐻푆푂4 ) in for each experimental data - point. To assess the fitting of the experiment data to the equilibrium line, the activity of HSO4

− (푎퐻푆푂4 ) has to be set to the same value for all data point. While keeping the Keq constant, an adjusted value of pH was determined using Equation 28. Figure 54 shows the plot of all

− experiment data against the equilibrium line at 푎퐻푆푂4 of 0.2. The 95% confidence was calculated using a t-test on the equilibrium constant (퐾푒푞) values. The confidence interval methodology can be found in Section 4.3.5 and the sample calculation can be found in Appendix C.1.

103

0 Figure 54 Equilibrium line fitting for hematite solubility data with aqueous FeHSO4SO4 − species at 푎퐻푆푂4 of 0.2

0 Based on the estimated 퐾푒푞, the Gibbs energy of formation for FeHSO4SO4 species at 220°C was calculated to be -1,375.6  1.4 kJ mol-1. This value is higher than the one reported by Papangelakis et al (1994) at 25°C which is -1514.4 kJ mol-1. The difference is - 138.8 kJ mol-1 which is approximately 9% higher.

4.5.2.2 Fitting of the Fe2(SO4)3 aqueous species 0 Similarly, the equilibrium reaction between aqueous neutral Fe2(SO4)3 and hematite can be described in Reaction 4.3

0 - + Fe2(SO4)3 (a) + 3H2O(l)↔ Fe2O3 (s) + 3HSO4 (a) + 3H (a) (4.3) Using the same set of experiment data as summarised in Table 32, the equilibrium constant

퐾푒푞 was calculated using Equation 29. All ferric species were assumed to be in the form of aqueous neutral complex Fe2(SO4)3. Since there are two ferric ions in the aqueous ferric complex species, its molality can be estimated using the ferric concentration measured by 1 ICP-OES where 푚 3+ ≈ 푚 . The activity coefficient of aqueous Fe2(SO4)3 was 2 퐹푒 ,25℃ 퐹푒2(푆푂4)3

- − assumed to unity. The concentration of HSO4 (푚퐻푆푂4 ) was calculated by the difference in

Total sulphur and sulphur and/or bisulphate associated by ferric. For Fe2(SO4)3 species, the sulphur species concentration incorporated in the aqueous complex is 3 of ferric 2

104

- − + + concentration. The activity coefficient of HSO4 (훾퐻푆푂4 ) and H (훾퐻 ) tabulated in Table 32 were used.

3 3 푎퐻푆푂4− 푎퐻+ 퐾퐹푒2푂3 = Equation 29 푎퐹푒2(푆푂4)30

The 퐾푒푞results for are summarised in Table 34. Based on 95% confidence interval, the 퐾푒푞 was estimated to be -3.250.08.

Table 34 Equilibrium constant values assuming hematite is in equilibrium with Fe2(SO4)3 aqueous species

log(퐾푒푞) Log 푎퐹푒2(푆푂4)3 Adjusted pH 1 -3.48 -2.65 1.35 2 -3.25 -2.35 1.17 3 -3.15 -2.02 1.92 4 -3.10 -1.73 0.91 5 -3.12 -1.98 1.00 6 -3.35 -2.48 1.25 7 -3.26 -2.96 1.27 8 -3.31 -2.90 1.37 9 -3.22 -3.51 1.55 10 -3.29 -2.44 1.18

To assess the fitting of the experimental data and equilibrium line, the activity values of - HSO4 of all experiment data were set to the same value. Adjusted value of pH was calculated using Equation 29. Figure 55 shows the plot of all experiment data against the

− equilibrium line between hematite and aqueous Fe2(SO4)3 at 푎퐻푆푂4 of 0.2. The 95% confidence was calculated using a t-test on the equilibrium constant (퐾푒푞) values. The confidence interval methodology can be found in Section 4.3.5 and the sample calculation can be found in Appendix C.1.

105

Figure 55 Equilibrium line fitting for hematite solubility data with aqueous Fe2(SO4)3 − species at 푎퐻푆푂4 of 0.2

0 Based on the estimated 퐾푒푞, the Gibbs energy of formation for aqueous Fe2(SO4)3 species at 220°C was calculated to be -2,088.6  0.8 kJ/mol. This value is higher than the one extrapolated from Liu et al. (2003) at 220°C which is -2,341 kJ mol-1. The difference is -138.8 kJ mol-1 which is approximately 9% higher.

0 4.5.3 Fitting of the FeHSO4SO4 and Fe2(SO4)3 aqueous species Based on Figure 54, it seems that the equilibrium line between hematite and aqueous

FeHSO4SO4 fits the experimental data at the low pH region. In contrast, equilibrium line between hematite and aqueous Fe2(SO4)3 seems to fit the experimental data at higher pH region as shown in Figure 55. This suggests that aqueous FeHSO4SO4 is likely to be the predominant ferric species at lower pH region, while at higher pH aqueous Fe2(SO4)3 is predominant ferric species as shown in Figure 56.

106

Figure 56 Fitting of equilibrium line between hematite, aqueous Fe2(SO4)3 and aqueous

0 − FeHSO4SO4 species at 푎퐻푆푂4 of 0.2

The cross over between the two aqueous ferric complex species can be predicted using the equilibrium reaction between aqueous FeHSO4SO4 and Fe2(SO4)3 in solution which is expressed in Reaction 4.4.

0 0 - + 2FeHSO4SO4 (a) ↔ Fe2(SO4)3 (a)+ HSO4 (a) + H (a) (4.4)

By only considering aqueous FeHSO4SO4 and Fe2(SO4)3 species and assuming activity coefficient of unity, the distribution of these two species in solution as a function of pH (at

− 푎퐻푆푂4 of 0.2) is illustrated in Figure 57. The Gibbs energy of formation of both aqueous ferric complex determined in Section 4.5.2 were used.

It can be seen that the cross over seems to occur at approximately pH 1 where at higher acidity aqueous FeHSO4SO4 is the predominant ferric species. It is important to note that this cross over point is not definitive and can only be confirmed through in-situ speciation study at high temperature which could not be done in this study.

107

− Figure 57 Distribution of aqueous Fe(III) species as a function of pH at 푎퐻푆푂4 of 0.2. Around this cross over point both aqueous ferric complex species are present in the aqueous solution in reality. However, as this study is focusing at predominant diagram, only the predominant species will be considered at that certain condition. For example, aqueous

FeHSO4SO4 is assumed to be the predominant ferric species at pH below 1 at bisulphate activity of 0.2. It is important to note that the cross over pH is dependent on the activity of bisulphate.

4.6 Solubility of basic ferric sulphate at 220°C

4.6.1 Solid XRD and morphology To investigate the solubility of basic ferric sulphate, pure solid phase was precipitated from various feed solutions as discussed in Section 4.2.1. Lithium jarosite did not seem to precipitate as no peaks associated with this phase were detected by XRD. The XRD pattern from basic ferric sulphate solubility experiments can be found in Appendix A.3. Figure 58 shows the needle-like structure of basic ferric sulphate precipitated in the experiment. The size of the particles is not uniform, varying from approximately 1 to 6 µm.

108

Figure 58 Morphology of basic ferric sulphate precipitated in basic ferric sulphate solubility experiment. Magnification 5,000X, 15kV and 10mm WD.

4.6.2 Results and discussion Table 35 summarises aqueous concentration for iron, total sulphur and lithium in the feed and filtrate solution, the measured and estimated pH at 220°C and the activity coefficient of + - proton (H ) and bisulphate (HSO4 ) ions. The aqueous concentration of feed solution were obtained from ICP-OES. The ferric and total sulphur concentration in filtrate was calculated based on the concentration obtained from ICP-OES, filtrate mass and filtrate density. The activity coefficients were estimated using HSC Chemistry as mentioned in Section 4.3.2. High temperature pH was estimated according to the methodology in Section 4.3.4.

Table 35 Solution data for basic ferric sulphate solubility

Feed, g/L Filtrate, g/kg-H2O pH,220°C + − pH,30°C 훾퐻 훾퐻푆푂4 Fe3+ S Li+ Fe3+ S Li+ estimated measured 1 8.7 17.4 - 2.3 15.8 - 0.40 0.41 0.51 0.42 2 17.7 31.6 - 3.4 24.8 - 0.40 0.42 0.51 0.43 3 17.1 31.7 - 3.3 23.8 - 0.40 0.42 0.51 0.44 4 26.5 40.3 - 4.0 28.8 - 0.38 0.40 0.51 0.41 5 15.6 36.5 - 3.1 23.7 - 0.48 0.50 0.53 0.43 6 26.7 31.3 - 4.5 28.0 - 0.34 0.35 0.51 0.44 7 13.7 30.1 0.13 3.8 25.5 0.13 0.36 0.37 0.47 0.46 8 13.2 30.0 0.09 4.1 26.1 0.09 0.35 0.36 0.50 0.45 9 13.0 28.2 0.04 4.5 26.6 0.04 0.34 0.35 0.50 0.44 10 12.8 28.5 0.13 4.2 25.8 0.13 0.36 0.37 0.50 0.45 109

To determine the Gibbs energy of formation for basic ferric sulphate (∆퐺°푓,퐹푒푂퐻푆푂4), aqueous

FeHSO4SO4 species was assumed to be in equilibrium with FeOHSO4 solid. This aqueous ferric complex species was selected as it was found to be the better prediction for predominant aqueous ferric species in acidic condition (pH < 1) at 220°C compared to aqueous Fe2(SO4)3 species (refer to Section 4.5.3). The equilibrium reaction can be expressed according to Reaction 4.5.

0 - + FeHSO4SO4 (a) + H2O(l) ↔ FeOHSO4(s) + HSO4 (a) + H (a) (4.5)

For each experimental data listed in Table 35, the equilibrium constant was calculated using

Equation 30. The activity coefficient of aqueous FeHSO4SO4 was assumed to be unity and therefore its activity is equivalent to its molal concentration. The concentration can be estimated using the ferric concentration measured by ICP-OES, where 푚퐹푒3+,25℃ ≈

- − 푚퐹푒퐻푆푂4푆푂4. The concentration of HSO4 (푚퐻푆푂4 ) was calculated by the difference in total sulphur and sulphur and/or bisulphate associated by ferric in the aqueous ferric complex species. For FeHSO4SO4 species, the number of sulphur moles incorporated in the aqueous

2 - + complex is of total ferric moles. The activity coefficient of HSO4 (훾 − ) and H (훾 +) 3 퐻푆푂4 퐻 used are tabulated in Table 35.

푎퐻푆푂4− 푎퐻+ 퐾퐹푒푂퐻푆푂4 = Equation 30 푎 0 퐹푒퐻푆푂4푆푂4

The 퐾푒푞 results for are summarised in Table 36. Based on 95% confidence interval, the

퐾푒푞 was estimated to be -0.065  0.010. To assess the fitting of the experimental data and

- equilibrium line, the activity values of HSO4 of all experiment data were set to the same value. By keeping 퐾 and 푎 0 constant, adjusted value of pH was calculated using 푒푞 퐹푒퐻푆푂4푆푂4 Equation 30.

Table 36 Equilibrium constant for basic ferric sulphate solubility

log(퐾푒푞) Log 푎퐹푒퐻푆푂4푆푂4 Adjusted pH

1 -0.107 -1.67 1.07 2 -0.053 -1.51 0.86 3 -0.043 -1.53 0.87 4 -0.042 -1.45 0.80 5 -0.092 -1.26 0.65

110

log(퐾푒푞) Log 푎퐹푒퐻푆푂4푆푂4 Adjusted pH

6 -0.065 -1.39 0.75 7 -0.052 -1.17 0.52 8 -0.054 -1.13 0.49 9 -0.086 -1.10 0.48 10 -0.083 -1.12 0.51

Figure 59 shows the plot of all experiment data against the equilibrium line between basic

− ferric sulphate solid and aqueous FeHSO4SO4 at 푎퐻푆푂4 of 0.2. The 95% confidence was calculated using a t-test on the equilibrium constant (퐾푒푞) values. The confidence interval methodology can be found in Section 4.3.5 and the sample calculation can be found in Appendix C.2.

Figure 59 Equilibrium line fitting for basic ferric sulphate solubility data with aqueous 0 − FeHSO4SO4 species at 푎퐻푆푂4 of 0.2.

Based on this study, equilibrium constant of -0.07 ± 0.03 was determined. Using this 퐾푒푞 value, the Gibbs energy of formation for basic ferric sulphate (∆퐺°푓,퐹푒푂퐻푆푂4 ) was calculated to be -919.4 ± 0.3. Comparing this value of ∆퐺°푓,퐹푒푂퐻푆푂4 to the one calculated from Majzlan et al. (2017) which is -1047.2 kJ mol-1 at 220°C, there is a difference of approximately 12% where the one calculated in this study is less negative meaning that basic ferric sulphate is

111 less stable. This difference may arise from the different technique used where Majzlan et al. (2017) determined the thermodynamic properties at 25°C using calorimetry approach.

Using both ∆퐺°푓,퐹푒푂퐻푆푂4 values, the equilibrium pH where the transition between basic ferric sulphate and hematite occurs was assessed. Figure 60 illustrates the equilibrium pH as a function of bisulphate activities. Below the equilibrium pH, basic ferric sulphate is the predominant solid while hematite is favoured above the equilibrium pH. The equilibrium pH was shown to shift towards a higher pH value as bisulphate activity was increased. This shows how basic ferric sulphate is more stable at higher bisulphate activity.

Figure 60 Equilibrium pH between basic ferric sulphate and hematite as a function of bisulphate activity

The equilibrium reaction between basic ferric sulphate and hematite can be expressed according to Reaction 4.6.

- + Fe2O3 (s) + 2HSO4 (a) + 2H (a) ↔ 2FeOHSO4 (s) + H2O(l) (4.6)

Assuming activity of unity for all solids and H2O, this equilibrium pH is affected only by the activity of bisulphate ion as shown in Equation 31.

1 퐾푒푞 = 2 2 Equation 31 푎퐻푆푂4− 푎퐻+

The pH was calculated using Equation 16, Equation 18 and Equation 30. Additional Gibbs free of formation values listed in Table 31 were used to calculated the Gibbs energy of

− reaction (∆퐺°푅 ) for Reaction 4.5. Additionally, equilibrium pH was evaluated at 푎퐻푆푂4 of 1, 112

− 0.3 and 0.1 which were the range of 푎퐻푆푂4 at which basic ferric sulphate and hematite were precipitated in the current experimental campaign.

Table 37 summarises the equilibrium pH values calculated using the two different values of

∆퐺°푓,퐹푒푂퐻푆푂4 . It was found that equilibrium pH values calculated using ∆퐺°푓,퐹푒푂퐻푆푂4 calculated from Majzlan et al. (2017) which is -1047.2 kJ mol-1, were not realistic.

Table 37 Equilibrium pH between basic ferric sulphate and hematite at two different activity of bisulphate for different Gibbs energy of formation of basic ferric sulphate

∆퐺° , Equilibrium pH hematite-BFS No. 푓,퐹푒푂퐻푆푂4 -1 − − − kJ mol 푎퐻푆푂4 = 1 푎퐻푆푂4 = 0.3 푎퐻푆푂4 = 0.1 1 -919.4 1.56 1.05 0.57 2 -1047.2 15.1 14.58 14.11

It is unlikely for basic ferric sulphate to be stable at high pH, especially at pH between 14 and 15. This suggests that the ∆퐺°푓,퐹푒푂퐻푆푂4 values calculated from Majzlan et al. (2017) were too negative which makes the basic ferric sulphate too stable. The equilibrium pH calculated

-1 using ∆퐺°푓,퐹푒푂퐻푆푂4 of -919.4 kJ mol which was determined in this study gave a more realistic estimate which aligns with the stability diagram presented by Fleming (2009) shown in

Figure 25. Therefore, both ∆퐺°푓,퐹푒퐻푆푂4푆푂4 and ∆퐺°푓,퐹푒푂퐻푆푂4 values that were determined in current investigation, were selected for the development of Eh-pH diagram in Section 4.9.

4.6.3 Equilibrium between basic ferric sulphate, hematite and aqueous ferric species

Figure 61 shows the solid-solid and solid-liquid equilibrium as a function of aqueous ferric species activity and pH between the following species: • Hematite and basic ferric sulphate solids;

0 • Basic ferric sulphate solid and aqueous FeHSO4SO4 species;

0 • Hematite solid and aqueous FeHSO4SO4 species;

0 • Hematite solid and aqueous Fe2(SO4)3 species;

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Figure 61 Solubility of hematite and basic ferric sulphate as a function of pH and aqueous ferric species activity From this diagram, it can be seen that the transition of solid phase from basic ferric sulphate to hematite as the solution pH increases, does not seem to be triggered by the change in 0 0 the predominant aqueous ferric species from aqueous FeHSO4SO4 to aqueous Fe2(SO4)3 species. The switch in the predominant aqueous ferric species is shown to occur within the hematite stability region. This means that there could possibly be a non-predominant aqueous ferric species present in the solution that triggers this solid phase transition.

4.7 Solubility of potassium jarosite at 220°C

4.7.1 Solid XRD and morphology To investigate the solubility of potassium jarosite, pure solid phase was precipitated from various feed solutions as discussed in Section 4.2.1. Lithium jarosite did not seem to precipitate as no peaks associated with this phase were detected by XRD. The XRD pattern from potassium jarosite solubility experiments can be found in Appendix A.4. Figure 62 shows the structure of potassium jarosite precipitated in the experiment. The size of the particles are pretty uniform at approximately 2 µm with some small particles of approximately 0.5 µm are seen to be present. The addition of lithium did not seem to affect the morphology of potassium jarosite.

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Figure 62 Potassium Jarosite Morphology. Magnification 5000X, 15kV and 10mm WD

4.7.2 Results and discussion Table 38 summarises aqueous concentration for iron, total sulphur, potassium and lithium in the feed and filtrate solution, the measured and estimated pH at 220°C and the activity + - coefficient of proton (H ) and bisulphate (HSO4 ) ions. The aqueous concentration of feed solution were obtained from ICP-OES. The ferric, potassium and total sulphur concentration in filtrate was calculated based on the concentration obtained from ICP-OES, filtrate mass and filtrate density. The activity coefficients were estimated using HSC Chemistry as mentioned in Section 4.3.2. High temperature pH was estimated according to the methodology in Section 4.3.4.

Table 38 Solution data for potassium jarosite solubility

Feed, g/L Filtrate, g/kg-H2O 훾푖 pH,220°C 3+ + + 3+ + + pH,30°C + - + Fe S K Li Fe S K Li estimated H HSO4 K measured 1 13.2 18.3 14.1 - 0.18 13.3 11.4 - 0.83 0.89 0.49 0.47 0.45 2 12.3 21.0 14.1 - 0.24 16.3 11.3 - 0.81 0.88 0.49 0.45 0.45 3 12.5 19.8 20.5 - 0.16 14.6 17.8 - 0.87 0.95 0.49 0.45 0.42 4 15.6 21.3 15.7 - 0.20 15.2 12.3 - 0.86 0.94 0.49 0.46 0.45 5 12.9 19.8 15.8 - 0.23 14.5 12.7 - 0.82 0.89 0.48 0.46 0.44 6 12.9 23.5 15.9 - 0.33 17.9 13.0 - 0.78 0.85 0.48 0.44 0.43 7 13.4 15.0 5.1 0.61 0.32 12.0 2.1 0.62 0.74 0.78 0.52 0.50 0.48 8 13.3 14.4 5.0 0.35 0.36 8.8 2.0 0.37 0.75 0.80 0.53 0.53 0.50 9 13.3 13.9 5.0 0.36 0.34 8.6 2.1 0.37 0.78 0.81 0.54 0.50 0.51 115

To determine the Gibbs energy of formation of potassium jarosite (∆퐺°푓,퐾퐹푒3 (푆푂4)2(푂퐻)6), aqueous FeHSO4SO4 species was assumed to be in equilibrium with FeOHSO4 solid. This aqueous ferric complex species was selected as it was found to be the better prediction for predominant aqueous ferric species in acidic condition (pH < 1) at 220°C compared to aqueous Fe2(SO4)3 species (refer to Section 4.5.3). The equilibrium reaction can be expressed according to Reaction 4.7.

0 + - + 3FeHSO4SO4 (a) + K (a) + 6H2O(l) ↔ KFe3(SO4)2(OH)6(s) + 4HSO4 (a) + 5H (a) (4.7)

For each experimental data listed in Table 38, the equilibrium constant was calculated using

Equation 32. The activity coefficient of aqueous FeHSO4SO4 was assumed to be unity and therefore its activity is equivalent to its molal concentration. The concentration can be estimated using the ferric concentration measured by ICP-OES, where 푚퐹푒3+,25℃ ≈

2 - + complex is of total ferric moles. The activity coefficient of HSO4 (훾 − ) and H (훾 +) 3 퐻푆푂4 퐻 used are tabulated in Table 38.

4 5 푎퐻푆푂4− 푎퐻+ 퐾퐾퐹푒3 (푆푂4)2(푂퐻)6 = 3 Equation 32 푎 0 푎 + 퐹푒퐻푆푂4푆푂4 퐾

The 퐾푒푞 results for are summarised in Table 36. Based on 95% confidence interval, the 퐾푒푞 was estimated to be -0.065  0.010. To assess the fitting of the experimental data and

- + equilibrium line, the activity values of HSO4 and K of all experiment data were set to the same value. By keeping 퐾 and 푎 0 constant, adjusted value of pH was calculated 푒푞 퐹푒퐻푆푂4푆푂4 using Equation 32.

Table 39 Equilibrium constant for potassium jarosite solubility

log(퐾푒푞) Log 푎퐹푒퐻푆푂4푆푂4 Adjusted pH

1 -0.54 -2.49 1.10 2 -0.57 -2.37 1.03 3 -0.78 -2.53 1.18 4 -0.77 -2.44 1.12 5 -0.67 -2.38 1.09

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log(퐾푒푞) Log 푎퐹푒퐻푆푂4푆푂4 Adjusted pH

6 -0.78 -2.23 0.99 7 -0.54 -2.24 0.94 8 -0.62 -2.19 0.92 9 -0.73 -2.22 0.97

Figure 63 shows the plot of all experiment data against the equilibrium line between

− + potassium jarosite and aqueous FeHSO4SO4 at 푎퐻푆푂4 of 0.2 and 푎퐾 of 0.5. The experimental data seems to fit the equilibrium line pretty well with all the data lie in within the 95% confidence interval. Using this 퐾푒푞 determined, the Gibbs energy of formation for -1 potassium jarosite (∆퐺°푓,퐾퐹푒3 (푆푂4)2(푂퐻)6), was calculated to be -3009.4 ± 0.7 kJ mol . The

95% confidence was calculated using a t-test on the equilibrium constant (퐾푒푞) values. The confidence interval methodology can be found in Section 4.3.5 and the sample calculation can be found in Appendix C.2.

Figure 63 Equilibrium line fitting for potassium jarosite solubility data with aqueous 0 − FeHSO4SO4 species at 푎퐻푆푂4 of 0.2.

As mentioned in Section 4.1, several different Gibbs energy of formation for potassium jarosite values at 220°C have been reported to date; they are -2,900.8 kJ mol-1 (Stoffregen, 2000), -3032.6 kJ mol-1 (Majzlan et al., 2010) and -3,426.9 kJ mol-1 (Stoffregen, 1993). The

117 value calculated in current investigation are within the range of these reference values. To assess these ∆퐺°푓,퐾퐹푒3 (푆푂4)2(푂퐻)6 values, the equilibrium reaction between potassium jarosite and hematite which is shown in Reaction 4.8 was considered.

+ 2- + 3Fe2O3 (s) + 2K (a) + 4HSO4 (a) + 3H2O(l) + 2H (a) ↔ 2KFe3(SO4)2(OH)6 (s) (4.8)

Based on this reaction, at higher acidity (below the equilibrium pH) potassium jarosite is favoured while hematite is favoured above the equilibrium pH. This equilibrium pH at which the hematite and potassium jarosite transition takes place was calculated at different bisulphate and potassium activities as summarised in Table 40.

Table 40 Equilibrium pH at different bisulphate and potassium activities for different values of potassium jarosite Gibbs energy of formation

pH hematite - K jarosite ∆퐺° No. 푓,퐾퐹푒3 (푆푂4)2(푂퐻)6 -1 − − − − kJ mol 푎퐻푆푂4 = 0.3, 푎퐻푆푂4 = 0.1, 푎퐻푆푂4 = 0.1, 푎퐻푆푂4 = 0.3, 푎퐾+ = 0.1 푎퐾+ = 0.1 푎퐾+ = 0.3 푎퐾+ = 0.3 1 -2900.8*1 -9.69 -10.6 -10.2 -9.21 2 -3009.4*2 1.81 0.86 1.33 2.29 3 -3032.6*3 4.27 3.31 3.79 4.75 4 -3426.9*4 46.0 45.1 45.6 46.5 *1 HSC Chemistry v.7.1 (Stoffregen, 2000) *2 This study *3 HSC Chemistry v.9 (Majzlan et al., 2010) *4 Stoffregen, 1993

− + At higher 푎퐻푆푂4 and 푎퐾 , the equilibrium pH shifts to a higher value meaning potassium jarosite is more stable at higher acidity and potassium concentration. Using -1 ∆퐺°푓,퐾퐹푒3 (푆푂4)2(푂퐻)6 of - 2900.8 and - 3426.9 kJ mol , the calculated equilibrium pH values were not realistic where they were either below pH -9 or over pH 44. However, when -1 ∆퐺°푓,퐾퐹푒3 (푆푂4)2(푂퐻)6 of - 3009.4 and - 3032.6 kJ mol were used, the average equilibrium pH were pH 1.5 and 4 respectively.

Figure 13 in Section 1.3.3.2 shows the stability field of potassium jarosite developed by Babcan (1971) as a function of pH and temperature. Based on this, the equilibrium pH between potassium jarosite and hematite at 220°C is close to pH 1. However, the equilibrium pH may vary from this pH value due to the following: • Potassium and free acid concentration were not given and therefore direct comparison with current study data is not possible.

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• pH measurement was taken at room temperature and no adjustment applied to high temperature. This means the pH stated in Babcan (1971) is likely to be overestimated at high temperature experiments.

However, the equilibrium pH should not vary by more than 2 pH unit. Therefore,

-1 ∆퐺°푓,퐾퐹푒3 (푆푂4)2(푂퐻)6 values of - 3009.4 kJ and - 3032.6 kJ mol seem to be appropriate with the earlier value gives a closer estimate of equilibrium pH between hematite and potassium jarosite to that in Babcan (1971) shown in Figure 13. Therefore, -3009.4 kJ mol-1 was selected as the Gibbs energy of formation for potassium jarosite in development of Eh-pH diagram in Section 4.8.4.

4.8 Development of the Eh-pH diagram at 220°C

4.8.1 Methodology In Pourbaix diagram, there are two types of reactions considered namely chemical and electrochemical reactions. Chemical reaction is defined as reaction with no unbalanced electrons and therefore its equilibrium is independent on electrode potential. The boundary line is represented by a vertical line with a particular value of pH which can be calculated using its Gibbs energy of reaction ( ∆퐺°푅) and equilibrium constant (퐾푒푞) as described in Equation 16 and Equation 19.

The electrochemical reactions are defined as reaction that involves the transfer i.e. generation or consumption, of electron. When the electrochemical reaction does not involve H+ or OH-, such as in the oxidation of ferrous to ferric ion, it is independent of pH. The boundary line is represented by a horizontal line with a particular Eh value which is equivalent to the ∆퐸° at a specified activity of species which is one at standard condition. Based on the aqueous species considered (Section 4.3.3), there is no pure redox reactions present in current investigation. All electrochemical reactions in current investigation involves H+ which means the equilibrium line is affected by both electrode potential and pH.

The Nernst equation is used to link Gibbs free energy, half-cell potential and species activities as shown in Equation 33 and Equation 34.

푅푇 ∆퐸 = ∆퐸° − ln 퐾 Equation 33 푛퐹 ∆퐺 ∆퐸 = − Equation 34 푛퐹

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Where ∆퐸° is the standard cell potential, R is the universal gas constant 8.314 J K-1 mol-1, n is the number of electrons transferred in the cell reaction, F is the Faraday constant, 9.648 × 104 C mol-1, ∆퐺° is the standard Gibbs free energy and K is the equilibrium constant.

4.8.2 Reactions

Current investigation develops Eh-pH diagram for Fe-S-H2O and Fe-S-K-H2O systems at

220°C. The reactions considered in present work and their respective either 퐸푒푞 (for electrochemical reaction) or ∆퐺rxn (for chemical reaction) are summarised in Table 41. Note that 퐸 and ∆퐺rxn were calculated at 220°C with an activity coefficient of one for all species. These reactions take account reactions between aqueous-aqueous, aqueous -solid and 2+ 0 + - solid-solid species. The aqueous species were Fe , FeHSO4SO4 , K , HSO4 and H2O, whereas the solid phases were Fe2O3, FeHSO4, KFe3(SO4)2(OH)6 and FeSO4.H2O (see Section 4.3.3).

Table 41 Summary of reactions considered for the development of Eh-pH diagram at -1 220°C. Eeq is in V and ΔGrxn is in kJ mol .

Electrochemical reaction 푬풆풒

0 + - 2+ - FeHSO4SO4 (a) + H (a) + e ↔ Fe (a) + 2HSO4 (a) - 0.275

2+ - + - Fe (a) + 2HSO4 (a) + 14H (a) + 14 e (a) ↔ FeS2(s) + 8H2O(l) 0.304

+ - 2+ Fe2O3 (s) + 6H (a) + 2 e ↔ 2Fe (a) + 3H2O (l) - 0.366

+2 - + - Fe (a) + H2O (l) + HSO4 (a) ↔ FeOHSO4 (s) + 2H (a) + e ↔ - 0.213

+ - +2 + - KFe3(SO4)2(OH)6 (s) + 8H (a) + 3e ↔ 3Fe (a) + K (a) + 2HSO4 (a) + 6H2O (l) - 0.240

0 + - FeHSO4SO4 (a) + 15H (a) + 15e ↔ FeS2(s) +8H2O(l) 0.302

0 - - FeHSO4SO4 (a) + H2O(a) +e ↔ FeSO4.H2O (s) + HSO4 (a) 0.615

- + - Fe2O3 (s) + 4HSO4 (a) + 34H (a) + 30e ↔ 2FeS2(s) + 19H2O(l) - 0.309

- + - FeOHSO4 (s) + HSO4 (a) + 16H (a) + 15e ↔ FeS2(s) + 9H2O(l) - 0.298

- + - FeSO4.H2O (s) + HSO4 (a) + 15H (a) + 14e ↔ FeS2(s) + 9H2O(l) - 0.706

2- + - + KFe3(SO4)2(OH)6 (s) + 4HSO4 (a) + 50H (a) + 45e ↔ 3FeS2(s) + K (a) + 30H2O(l) - 0.300

- + - Fe2O3 (s) + 2HSO4 (a) + 4H (a) + 2e ↔ 2FeSO4.H2O (s) + H2O(l) 0.706

+ - FeOHSO4 (s) + H (a) + e ↔ FeSO4.H2O (s) 0.553

- + - + KFe3(SO4)2(OH)6 (s) + HSO4 (a) + 5H (a) + 3e ↔ 3FeSO4.H2O (s) + K (a) + - 0.580 3H2O(l)

Chemical reactions ΔGrxn

+2 - + Fe (a) + HSO4 (a) + H2O (l) ↔ FeSO4.H2O (s)+ H (a) 32.8

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Chemical reactions ΔGrxn

- + 2FeHSO4SO4 (a) + 2H2O(l)↔ Fe2O3 (s) + 4HSO4 (a) + 4H (a) -189.5

- + FeHSO4SO4 (a) + H2O(l) ↔FeOHSO4 (s) + HSO4 (a) + H (a) -6.0

+ - + 3FeHSO4SO4 (a) + K (a) + 6H2O(l) ↔ KFe3(SO4)2(OH)6 (s) + 4HSO4 (a) + 5H (a) -10.2

- + Fe2O3 (s) + 2HSO4 (a) + 2H (a) ↔ 2FeOHSO4 (s) + H2O(l) -29.5

+ 2- + 3Fe2O3 (s) + 2K (a) + 4HSO4 (a) + 3H2O(l) + 2H (a) ↔ 2KFe3(SO4)2(OH)6 (s) -72.8

+ - + 3FeOHSO4 (s) + K (a) + 3H2O(l) ↔KFe3(SO4)2(OH)6 (s) + HSO4 (a) + 2H (a) 7.9

4.8.3 Eh-pH diagram of Fe-S-H2O at 220°C

4.8.3.1 Eh-pH diagram development

After removing redundant lines, the Eh-pH diagram for Fe-S-H2O system can be defined with 0 a total of twelve equilibrium lines showing the stability fields of aqueous FeHSO4SO4 , 2+ aqueous Fe , basic ferric sulphate (FeOHSO4), hematite (Fe2O3), pyrite (FeS2) and szomolnikite (FeSO4.H2O). The equilibrium reactions are summarised in Table 42.

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Table 42 Equilibrium reactions and equilibrium constant expression required for the development of Eh-pH diagram for Fe-S-H2O system at 220°C No Reactions Equilibrium Constant 2 푎 2+ . 푎 − 0 + - 2+ - 퐹푒 퐻푆푂4 R1 FeHSO4SO4 (a) + H (a) + e ↔ Fe (a) + 2HSO4 (a) 퐾1 = 푎 0 푎 + 퐹푒퐻푆푂4푆푂4 . 퐻 푎 2+. 푎퐻푆푂 − + - +2 - 퐹푒 4 R2 FeOHSO4 (s) + 2H (a) + e ↔ Fe (a) + H2O (l) + HSO4 (a) 퐾2 = 2 푎퐻+ 2 2+ + - 2+ 푎퐹푒 R3 Fe2O3 (s) + 6H (a) + 2e ↔ 2Fe (a) + 3H2O (l) 퐾3 = 6 푎퐻+ 1 2+ - + - R4 Fe (a) + 2HSO4 (a) + 14H (a) + 14e ↔ FeS2 (s) + 8H2O(l) 퐾4 = 2+ −2 +14 푎퐹푒 . 푎퐻푆푂4 . 푎퐻 푎퐻푆푂 − . 푎퐻+ 0 - + 4 R5 FeHSO4SO4 (a) + H2O(l) ↔ FeOHSO4(s) + HSO4 (a) + H (a) 퐾5 = 푎 0 퐹푒퐻푆푂4푆푂4 1 - + R6 Fe2O3 (s) + 2HSO4 (a) + 2H (a) ↔ 2FeOHSO4 (s) + H2O(l) 퐾6 = − 2 +2 푎퐻푆푂4 . 푎퐻 푎퐻+ +2 - + R7 Fe (a) + HSO4 (a) + H2O (l) ↔ FeSO4.H2O (s) + H (a) 퐾7 = 2+ − 푎퐹푒 . 푎퐻푆푂4 1 - + - R8 Fe2O3 (s) + 2HSO4 (a) + 4H (a) + 2e ↔ 2FeSO4.H2O (s) + H2O(l) 퐾8 = − 2 +4 푎퐻푆푂4 . 푎퐻 1 - + - R9 FeSO4.H2O (s) + HSO4 (a) + 15H (a) + 14e ↔ FeS2 (s) + 9H2O (l) 퐾9 = − +15 푎퐻푆푂4 . 푎퐻 1 - + - R10 Fe2O3 (s) + 4HSO4 (a) + 34H (a) + 30e ↔ 2FeS2 (s) + 19H2O (l) 퐾10 = −4 +34 푎퐻푆푂4 . 푎퐻 1 + - R11 FeOHSO4 (s) + H (a) + e ↔FeSO4.H2O (s) 퐾11 = 푎퐻+ 푎퐻푆푂 − 0 - - 4 R12 FeHSO4SO4 (a) + H2O(a) + e ↔ FeSO4.H2O (s) + HSO4 (a) 퐾5 = 푎 0 퐹푒퐻푆푂4푆푂4 122

Depending on the activities of the dissolved species, the location of the equilibrium lines will change and therefore affecting the appearance and disappearance of some equilibrium lines. In Figure 64, the equilibrium line between Fe(II) and szomolnokite falls under the hematite stability field and therefore equilibrium line R1 to R10 are shown. When that equilibrium line falls under basic ferric sulphate stability field which is shown in Figure 65, R3 disappears and R11 appears. Again, when this equilibrium line falls under the stability 0 field of aqueous FeHSO4SO4 shown in Figure 66, R2 disappears and R12 appears.

1.4

1.2

1.0

0.8

0. 6

0.4

0.2

0.0

-0.2 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

Figure 64 Illustration #1 of Eh-pH diagram for Fe-S-H2O system at 220°C

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1.2

1.0

0.8

0. 6

0.4

0.2

0.0

-0.2 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

pH Figure 65 Illustration #2 of Eh-pH diagram for Fe-S-H2O system at 220°C

1.2

1.0

0.8

0. 6

0.4

0.2

0.0

-0.2 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

pH Figure 66 Illustration #3 of Eh-pH diagram for Fe-S-H2O system at 220°C

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4.8.3.2 Data validation To validate diagram, experimental data from hematite solubility (Section 4.5) and basic ferric sulphate solubility (Section 4.6) were plotted on the Eh-pH diagram. The pH of all experimental data were re-adjusted at the same 푎 0 and 푎 − as the ones inputted 퐹푒퐻푆푂4푆푂4 퐻푆푂4 in the Eh-pH diagram, by keeping the same equilibrium constant (퐾푒푞). The potential data were measured during the experiment and were adjusted to take account for temperature junction potential (TJP) and liquid junction potential (LJP) as described in Section 2.2.2.2.

− The hematite precipitation conditions in current investigation were approximated at 푎퐻푆푂4 of 0.1 and 푎 0 of 0.01. Based on these activity values, the calculated pH for all 퐹푒퐻푆푂4푆푂4 hematite solubility experiment data are shown in Table 43.

Table 43 Adjusted Eh (V vs. NHE) and pH data for hematite solubility experiments

Measured E Adjusted Eh log(퐾 ) Adjusted pH 푒푞 (V vs. Ag/AgCl) (V vs. NHE) 1 -3.41 0.71 0.92 0.86 2 -3.37 0.71 0.92 0.74 3 -3.46 0.65 0.87 0.77 4 -3.36 0.66 0.87 0.64 5 -3.34 0.67 0.88 0.82 6 -3.43 0.57 0.79 0.63 7 -2.98 0.59 0.81 0.62 8 -3.08 0.41 0.62 0.80 9 -2.55 0.50 0.72 0.79 10 -3.43 0.55 0.77 0.82

− These data were then plotted on the Eh-pH diagram. For the diagram, same values of 푎퐻푆푂4 and 푎 0 were selected. Although ferrous ion was not added into the solubility feed 퐹푒퐻푆푂4푆푂4 solution, its activity (푎퐹푒2+ ) was set to 0.0001 instead of 0 to display the stability field of ferrous on the diagram. This value was selected to represent a very low concentration of ferrous. Figure 67 shows that the hematite solubility data fit the diagram pretty well where all data points were in the hematite stability field.

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1.4

1.2

1.0

0.8

0. 6

0.4

0.2

0.0

-0.2 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

pH Figure 67 Fitting of Hematite solubility data on Eh-pH diagram for Fe-S-H2O system at 220°C at 푎 − of 0.1 and 푎 0 of 0.01. 퐻푆푂4 퐹푒퐻푆푂4푆푂4

For basic ferric sulphate data validation, the precipitation conditions at which pure basic ferric sulphate formed were approximated at 푎 − of 0.3 and 푎 0 of 0.1. Table 44 퐻푆푂4 퐹푒퐻푆푂4푆푂4 summarises the adjusted pH and Eh (V vs. NHE) for all basic ferric sulphate solubility experiment data.

Table 44 Adjusted Eh (V vs. NHE) and pH data for basic ferric sulphate solubility experiments

Measured E Adjusted Eh log(퐾 ) Adjusted pH 푒푞 (V vs. Ag/AgCl) (V vs. NHE) 1 -0.107 0.71 0.93 0.58 2 -0.053 0.74 0.95 0.53 3 -0.043 0.75 0.97 0.52 4 -0.042 0.74 0.96 0.57 5 -0.092 0.72 0.94 0.54 6 -0.065 0.72 0.94 0.53 7 -0.052 0.76 0.98 0.53 8 -0.054 0.74 0.95 0.56 9 -0.086 0.70 0.92 0.56 126

When approximately 1 g/L of lithium hydroxide was introduced into the feed solution 3+ (approximately 13 g/L Fe and 50 g/L H2SO4), hematite was observed in the solid precipitate as shown in Figure 68. The filtrate solution pH was also found to increase to 0.67 where previously it was seemed to be buffered at around 0.36 when 0.04 to 0.13 g/L Li+ was added and BFS was the only precipitate (refer to Table 35).

Figure 68 Mixture of basic ferric sulphate and hematite precipitate. Magnification 5,000X, 15 kV, WD 10mm.

The solution assay of this mixed basic ferric sulphate - hematite experiment data is summarised in Table 45. The filtrate pH at 220°C was estimated according to the + - methodology in Section 4.3.4. The activity coefficient of proton (H ) and bisulphate (HSO4 ) were estimated using HSC Chemistry as described in Section 4.3.2.

Table 45 Solution assay on basic ferric sulphate-hematite mixed precipitate.

Assay (ICP-OES) Feed (g/L) Filtrate (g/kg-H2O) Fe3+ 13.1 4.19 Total S 29.4 28.1 Li+ 0.30 0.31 Filtrate parameters

pH,30°C measured 0.67 1 pH,220°C estimated 0.73

훾퐻+ 0.47

− 훾퐻푆푂4 0.45

127

Since the solid contained both hematite and basic ferric sulphate, the equilibrium constant

(퐾푒푞) was calculated using both methodology as described in Section 4.5.2.2 and Section

4.6.2 for comparison. The pH was adjusted at 푎 − of 0.3 and 푎 0 of 0.1. Note that 퐻푆푂4 퐹푒퐻푆푂4푆푂4 0 both methods assume solid to be in equilibrium with FeHSO4SO4 (a) species.

The 퐾퐹푒2푂3 was estimated to be -3.9 with an adjusted pH of 0.96. While 퐾퐹푒푂퐻푆푂4 was estimated to be -0.42 with and adjusted pH of 0.90. Since the mixture is approximated to contain 50% hematite and 50% basic ferric based on the QXRD result, the experiment pH data point was taken to be 0.93. The Eh of the solution at 220°C was measured to be 0.73 V (vs. Ag/AgCl). After applying thermal and liquid-liquid junction potential correction, the Eh was approximated to be 0.95 V (vs. NHE).

Figure 69 shows the plot of pure basic ferric sulphate data summarised in Table 44 and the hematite/basic ferric sulphate mixture to the Eh-pH diagram at 푎 − of 0.3, 푎 0 of 퐻푆푂4 퐹푒퐻푆푂4푆푂4

0.1 and 푎퐹푒2+ of 0.0001. All pure basic ferric sulphate data are in the stability field of basic ferric sulphate. In theory, when basic ferric sulphate – hematite - solution are in equilibrium i.e. at three point equilibrium shown in Figure 61, the data point should be on the equilibrium line between basic ferric sulphate and hematite on the Eh-pH diagram. The mixed data, however, is also in the basic ferric sulphate stability field although it is close to the equilibrium line with less than 0.1 pH unit difference.

128

1.4

1.2

1.0

0.8

0. 6

0.4

0.2

0.0

-0.2 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 pH

Figure 69 Fitting of pure basic ferric sulphate and mixed basic ferric sulphate-hematite precipitates on the Eh-pH diagram at 푎 − of 0.3, 푎 0 of 0.1 and 푎 2+ of 0.0001 HSO4 FeHSO4SO4 Fe

This difference is still within the accuracy of the pH probe which could be the source of the error. However, there could be another reason for this. Although the precipitation of basic ferric sulphate and hematite may have reached equilibrium in two hours, the solid-solid reaction i.e. transformation from hematite to basic ferric sulphate and vice versa, may need a longer time to reach equilibrium. This factor has not been taken into consideration in current investigation. Also, in real system there is no clear cut between one phase to another, there is a window where both phases may present in the precipitate near the equilibrium line.

4.8.3.3 Effect of bisulphate, aqueous ferric complex and ferrous activities

The three main parameters affecting the Eh-pH diagram for Fe-S-H2O system are activity of bisulphate ion (푎 −), activity of aqueous ferric complex species (푎 0) and activity 퐻푆푂4 퐹푒퐻푆푂4푆푂4 of ferrous ion (푎퐹푒2+ ). Figure 70 shows the effect of bisulphate activity on the Eh-pH diagram

− at 푎퐻푆푂4 of 0.05, 0.1 and 0.2. In general, the activity of bisulphate species has a positive correlation with the stability field of aqueous FeHSO4SO4 species, basic ferric sulphate solid

129 and ferrous sulphate monohydrate solid. Increasing the activity of bisulphate seems to increase the stability field of aqueous FeHSO4SO4 species towards a higher pH and more reducing conditions. For basic ferric sulphate, its stability field shifts towards a higher pH and expands towards a less oxidising condition as bisulphate activity is increased. Similarly, the stability field of ferrous sulphate monohydrate also increases considerably at higher bisulphate activity. Consequently, the stability fields of aqueous ferrous and hematite solid shrink as bisulphate activity increases. The stability field of pyrite solid is hardly affected by bisulphate activity.

1.2

1.0

0.8

0.6

0.4

0.2

0.0

-0.2 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 pH

Figure 70 Effect of bisulphate ion activity on Eh-pH diagram for Fe-S-H2O system at 220°C −4 with 푎 0 = 0.1 and 푎 2+ = 1 × 10 퐹푒퐻푆푂4푆푂4 퐹푒

Figure 71 shows the impact of aqeuous ferric complex activity on the Eh-pH diagram at

푎 0 values of 0.05, 0.1 and 0.2. Activity of aqueous FeHSO4SO4 only has an effect 퐹푒퐻푆푂4푆푂4 the stability field of the adjacent species which are basic ferric sulphate (FeHSO4) solid and 2+ ferrous (Fe ). At higher activity value of aqueous FeHSO4SO4, the stability field of basic ferric sulphate expands towards a more acidic region while that of aqueous ferrous expands toward a more oxidising region. Other species stability fields remain unaffected.

130

1.2

1.0

0.8

0. 6

0.4

0.2

0.0

-0.2 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 pH

Figure 71 Effect of aqueous ferric complex activity on Eh-pH for Fe-S-H2O system at 220°C −4 − 2+ with 푎퐻푆푂4 = 0.1 and 푎퐹푒 = 1 × 10

Figure 72 shows the effect of aqueous ferrous activity on the the Eh-pH diagram at 푎퐹푒2+ values of 1 × 10−4, 1 × 10−3 and 1 × 10−2. It significantly affects the stability field of ferrous sulphate monohydrate (FeSO4.H2O) solid where it expands towards a more acidic region at higher ferrous activity. The stability field of aqueous ferric complex (FeHSO4SO4) and basic ferric sulphate (FeOHSO4) solid also expand towards a more reducing conditions. The effect of ferrous activity on the stability hematite only occurs when the equilibrium pH between 2+ Fe (a) and FeSO4.H2O (s) is at a more acidic condition than equilibrium pH between

FeOHSO4 (s) and Fe2O3 (s). Stability field of pyrite is only slightly affected. Varying the activity of ferrous is equivalent to varying Fe2+:Fe3+ ratio.

131

1.2

1.0

0.8

0. 6

0.4

0.2

0.0

-0.2 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

Figure 72 Effect of aqueous ferrous activity on Eh-pH for Fe-S-H2O system at 220°C evaluated at 푎 0 = 0.1 and 푎 − = 0.1 FeHSO4SO4 HSO4

4.8.4 Eh-pH diagram of Fe-K-S-H2O at 220°C

4.8.4.1 Eh-pH diagram development

For the Fe-K-S-H2O system, in addition to the equilibrium reactions listed in Table 42, five additional equilibrium lines listed in Table 46 are required to define stability field of solid potassium jarosite (KFe3(SO4)2(OH)6) in the Eh-pH diagram.

132

Table 46 Additional equilibrium reactions and equilibrium constant expression required for the development of Eh-pH diagram for Fe-S-K-H2O system at 220°C

No Reactions Equilibrium Constant 2 − + + - + 푎퐻푆푂4 . 푎퐻 R13 3FeOHSO4 (s) + K (a) + 3H2O(l) ↔KFe3(SO4)2(OH)6 (s) + HSO4 (a) + 2H (a) 퐾13 = 푎퐾+ 1 + - + R14 3Fe2O3 (s) + 2K (a) + 4HSO4 (a) + 3H2O(l) + 2H (a) ↔ 2KFe3(SO4)2(OH)6 (s) 퐾14 = +2 − 4 + 2 푎퐻 . 푎퐻푆푂4 . 푎퐾 8 푎 + +2 + - + - 퐻 R15 3Fe (a) + K (a) + 2HSO4 (a) + 6H2O (l) ↔ KFe3(SO4)2(OH)6 (s) + 8H (a) + 3e 퐾15 = 2+3 −2 푎퐹푒 . 푎퐻푆푂4 5 − + + - + - 푎퐻푆푂4 . 푎퐻 R16 3 FeSO4.H2O (s) + K (a) + 3H2O(l) ↔ KFe3(SO4)2(OH)6 (s) + HSO4 (a) + 5H (a) + 3e 퐾16 = 푎퐾+ 4 5 푎 − . 푎 + 0 + - + 퐻푆푂4 퐻 R17 3FeHSO4SO4 (a) + K (a) + 6H2O(l) ↔ KFe3(SO4)2(OH)6 (s) + 4HSO4 (a) + 5H (a) 퐾17 = 3 푎 0 . 푎 + 퐹푒퐻푆푂4푆푂4 퐾

133

Similar to the Eh-pH diagram for Fe-S-H2O system, there are a few scenarios which prompt the changes in equilibrium lines shown on the diagram to define the species stability field.

Figure 73 shows the Eh-pH diagram for Fe-K-S-H2O system at 220°C when the equilibrium line R7 between aqueous Fe(II) and szomonolkite (FeSO4.H2O) falls beneath the stability field of hematite (Fe2O3) solid. R13, R14 and R15 are the additional equilibrium lines required to define the stability field of solid potassium jarosite (KFe3(SO4)2(OH)6). These lines are equilibrium lines between potassium jarosite and basic ferric sulphate (FeOHSO4) solid, hematite (Fe2O3) solid and aqueous Fe(II) respectively.

1.9 1.7 1.5 1.3 1.1 0.9

0.7 0.5 0.3 0.1 -0.1 -0.3 -0.5 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

Figure 73 Illustration #1 for Eh-pH diagram of Fe-K-S-H2O system at 220°C

When R7 falls beneath the stability field of potassium jarosite (KFe3(SO4)2(OH)6) solid e.g. at higher activity of aqueous Fe(II) or higher activity of aqueous potassium (K+) and it, R16 which is the equilibrium line between potassium jarosite and szomonolkite (FeSO4.H2O) appears as illustrated in Figure 74. when R7 falls beneath the stability field of basic ferric sulphate (FeOHSO4) solid, the equilibrium line between potassium jarosite

(KFe3(SO4)2(OH)6) solid and aqueous Fe(II) species disappears (R15) as shown in Figure 75.

134

1.9 1.7 1.5 1.3 1.1 0.9

0.7 0.5 0.3 0.1 -0.1 -0.3 -0.5 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

Figure 74 Illustration #2 for Eh-pH diagram of Fe-K-S-H2O system at 220°C

1.9 1.7 1.5 1.3 1.1 0.9

0.7 0.5 0.3 0.1 -0.1 -0.3 -0.5 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

Figure 75 Illustration #3 for Eh-pH diagram of Fe-K-S-H2O system at 220°C

135 | P a g e

Furthermore, the stability field of basic ferric sulphate (FeOHSO4) solid could disappear at high activity of aqueous potassium (K+) species when equilibrium lines R5 and R13 overlap. Under this condition, the equilibrium lines that define the stability field of basic ferric sulphate solid (R2, R11 and R13) disappear while the equilibrium lines between aqueous ferric 0 complex (FeHSO4SO4 ) species and potassium jarosite solid (R17), and aqueous ferric 0 complex (FeHSO4SO4 ) species and szomonolkite solid (R12) appear as shown in Figure 76. The effect of aqueous potassium activity will be further discussed in Section 4.8.4.3.

1.9 1.7 1.5 1.3 1.1 0.9

0.7 0.5 0.3 0.1 -0.1 -0.3 -0.5 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

Figure 76 Illustration #4 for Eh-pH diagram of Fe-K-S-H2O system at 220°C

4.8.4.2 Potassium jarosite data validation

To validate the Eh-pH diagram for Fe-K-S-H2O system, experimental data from potassium jarosite solubility (Section 4.7) were plotted on the diagram. The pH of all experimental data were re-adjusted at the same 푎 0, 푎 −and 푎 + as the ones inputted in the Eh-pH 퐹푒퐻푆푂4푆푂4 퐻푆푂4 퐾 diagram, by keeping the same equilibrium constant 퐾푒푞. The potential data were measured during the experiment and were adjusted to take account for temperature junction potential (TJP) and liquid junction potential (LJP) as described in Section 2.2.2.2.

The potassium jarosite precipitation conditions in current investigation were approximated at 푎 + of 0.1, 푎 − of 0.1 and 푎 0 of 0.02. Based on these activity values, the 퐾 퐻푆푂4 퐹푒퐻푆푂4푆푂4 adjusted pH for potassium jarosite solubility experiment data are shown in Table 47. 136 | P a g e

Table 47 Adjusted Eh (V vs. NHE) and pH data for potassium jarosite solubility experiments

Measured E Adjusted Eh Adjusted log(퐾 ) 푒푞 (V vs. Ag/AgCl) (V vs. NHE) pH 1 -0.54 0.65 0.86 0.53 2 -0.57 0.64 0.86 0.54 3 -0.78 0.61 0.83 0.56 4 -0.77 0.63 0.84 0.58 5 -0.67 0.61 0.82 0.55 6 -0.78 0.59 0.81 0.54 7 -0.54 0.59 0.81 0.58 8 -0.62 0.60 0.80 0.53 9 -0.73 0.58 0.85 0.53

When approximately 3.5 to 6.5 g/L of lithium hydroxide was added into the feed solution, hematite was observed in the solid precipitate as shown in Figure 77. With the addition of 3.5 g/L LiOH, traces of hematite was started to form in the precipitate and the colour of the precipitate was deep orange (left image). At 6.5 g/L LiOH, the colour of the solid precipitate was red showing that the majority of the solids were hematite with traces of potassium jarosite Figure 77. The solutions assay of these mixed precipitate data are summarised in Table 48.

Potassium Jarosite Hematite Hematite

Potassium Jarosite

Figure 77 Mixed potassium jarosite-hematite precipitate formed with the addition of 4 g/L (left) and 7g/L (right) lithium hydroxide. Magnification 5,000X, 15 kV and 10mm WD.

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Table 48 Solution assay for mixed potassium jarosite-hematite data

Feed, g/L Filtrate, g/kg-H2O 훾푖 pH,220°C 3+ + + 3+ + + pH,30°C + - + Fe S K Li Fe S K Li estimated H HSO4 K measured 1 13.5 15.0 5.1 1.9 0.33 11.4 4.8 1.9 0.88 0.95 0.46 0.44 0.41 2 13.3 14.7 5.0 1.6 0.31 11.0 4.6 1.6 0.87 0.93 0.46 0.44 0.41 3 13.5 14.7 5.0 1.0 0.31 9.2 2.3 1.0 0.84 0.88 0.46 0.45 0.48

In mixed precipitate #1 and #2 where hematite was the major phase in the precipitate, the potassium concentration in the final filtrate is considerably higher than that in mixed precipitate #3 whose major phase was potassium jarosite. To plot these mixed precipitate

+ − data together with the pure potassium jarosite data, the adjusted pH at 푎퐾 of 0.1, 푎퐻푆푂4 of

0.1 and 푎 0 of 0.02 were calculated and summarised in Table 49 together with their 퐹푒퐻푆푂4푆푂4 potential data. The equilibrium constant for each data was calculated differently depending on the major phase present in the precipitate. For hematite majority solid, the equilibrium constant was estimated using Equation 28 in Section 4.5.2.1, while for potassium jarosite majority solid, the equilibrium constant was estimated using Equation 32 in Section 4.7.2.

Table 49 Adjusted Eh (V vs. NHE) and pH data for mixed potassium jarosite-hematite solubility experiments

Measured Eh Adjusted Eh log(퐾 ) Adjusted pH Major Phase 푒푞 (V vs. Ag/AgCl) (V vs. NHE) 1 -3.90 0.57 0.79 0.89 Hematite 2 -3.80 0.59 0.81 0.87 Hematite 3 -1.41 0.61 0.83 0.70 Potassium jarosite

Data tabulated in Table 47 and Table 49 were then plotted on the same Eh-pH diagram. For the diagram, same values of 푎 + , 푎 − and 푎 0 were selected. Although ferrous 퐾 퐻푆푂4 퐹푒퐻푆푂4푆푂4 ion was not added into the solubility feed solution, its activity (푎퐹푒2+ ) was set to 0.0001 instead of 0 to display the stability field of ferrous on the diagram. This value was selected to represent a very low concentration of ferrous.

Figure 78 shows the fitting of the experimental data for both pure potassium jarosite and mixed potassium jarosite-hematite precipitate. All pure potassium jarosite experimental data fall in the potassium jarosite stability field. For the hematite majority precipitate, both data are in the hematite stability field, near (pH±0.05) the equilibrium line between potassium jarosite and hematite. For potassium jarosite majority precipitate, it is in the potassium 138 | P a g e jarosite stability and closer to the potassium jarosite-hematite equilibrium compared to the pure potassium jarosite data. These experiment data have effectively validated the Eh-pH diagram.

1.9 1.7 1.5 1.3 1.1 0.9

0.7 0.5 0.3 0.1 -0.1 -0.3

-0.5 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

Figure 78 Plot of pure potassium jarosite and mixed precipitate experiments data on Eh-pH

diagram for Fe-S-K-H2O system at 220°C. The Eh-pH diagram was plotted with −4 푎 0 = 0.02, 푎 − = 0.1, 푎 + = 0.1 and 푎 2+ = 1 × 10 퐹푒퐻푆푂4푆푂4 퐻푆푂4 퐾 퐹푒

4.8.4.3 Effect of bisulphate, ferrous, potassium and ferric complex activities

The Eh-pH diagram for Fe-K-SO4-H2O system developed in current investigation is affected

− 2+ + by the activities of bisulphate (푎퐻푆푂4 ), ferrous(푎퐹푒 ), potassium (푎퐾 ). and ferric complex

(푎 0) species. 퐹푒퐻푆푂4푆푂4

Figure 79 shows the effect of potassium ion activity on the stability field of potassium jarosite

(KFe3(SO4)3(OH)6) solid as it the only species that is dependent on the potassium activity. At lower potassium activity, the stability field of potassium jarosite shrinks on both sides of the boundaries where potassium jarosite is in equilibrium with hematite and basic ferric sulphate. There is a minimum potassium concentration where below this concentration potassium jarosite will no longer be stable. As a result of the shrinking of potassium jarosite stability field, the stability field of hematite and basic ferric sulphate increase at lower potassium activity. 139 | P a g e

1.9 1.7 1.5 1.3 1.1 0.9

0.7 0.5 0.3 0.1 -0.1 -0.3 -0.5 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

pH Figure 79 The effect of potassium activity on Eh-pH diagram for Fe-K-S-H2O system at −4 220°C with 푎 0 = 0.1, 푎 − = 0.1 and 푎 2+ = 1 × 10 퐹푒퐻푆푂4푆푂4 퐻푆푂4 퐹푒

Figure 80 shows the effect of bisulphate activity on the stability field of aqueous ferric 0 complex (FeHSO4SO4 ), basic ferric sulphate (FeOHSO4) solid, potassium jarosite

(KFe3(SO4)2(OH)6) solid and szomolnokite (FeSO4.H2O) solid. At lower bisulphate activity, the stability fields of aqueous ferric complex, basic ferric sulphate solid and potassium jarosites solid shift to the left towards a more acidic condition and slightly shrinks toward a more oxidising conditions. Since the shift of equilibrium line between aqueous ferric complex and basic ferric sulphate is more prominent than that of equilibrium between basic ferric sulphate solid and potassium jarosite solid, the overall stability field of basic ferric sulphate seems to slightly expand. In contrast, the stability field of potassium jarosite solid diminishes quite significantly at lower bisulphate activity. The minimum potassium concentration for potassium jarosite solid, however, increases at lower bisulphate activity as shown in Figure 81. Similarly, the stability field of szomonolkite solid also decreases at lower bisulphate activity. Consequently, the stability field of hematite solid and aqueous ferrous increase at lower bisulphate activity.

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1.9 1.7 1.5 1.3 1.1 0.9

0.7 0.5 0.3 0.1 -0.1 -0.3

-0.5 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

pH - Figure 80 The effect of bisulphate activity (a HSO4 ) on Eh-pH diagram for Fe-K-S-H2O −4 system at 220°C with 푎 0 = 0.1, 푎 += 0.1 and 푎 2+ = 1 × 10 퐹푒퐻푆푂4푆푂4 퐾 퐹푒

Figure 81 The effect bisulphate activity ( −) on minimum potassium concentration for 푎퐻푆푂4 potassium jarosite solid stability

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0 Figure 82 shows the effect of aqueous ferric complex (FeHSO4SO4 ) activity on the stability field of its adjacent species which are the basic ferric sulphate (FeOHSO4) solid and aqueous ferrous (Fe2+). At higher aqueous ferric complex activity, the stability field of basic ferric sulphate expands towards a more acidic condition while the stability field of ferrous slight expands towards a more oxidising condition.

1.9 1.7 1.5 1.3 1.1 0.9

0.7 0.5 0.3 0.1 -0.1 -0.3

-0.5 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

pH 0 Figure 82 The effect of aqueous ferric activity (a FeHSO4SO4 ) on Eh-pH diagram for −4 Fe-K-S-H2O system at 220°C with − + and 2+ 푎퐻푆푂4 = 0.1 , 푎퐾 = 0.1 푎퐹푒 = 1 × 10

Figure 83 shows the effect of aqueous ferrous activity on the stability field of its adjacent species. At higher ferrous activity i.e. lower Fe3+:Fe2+ ratio, the stability field of szomonolkite

(FeSO4.H2O) solid significantly expands towards a more acidic condition while the stability 0 field of aqueous ferric complex (FeHSO4SO4 ), basic ferric sulphate (FeOHSO4) solid and potassium jarosite (KFe3(SO4)2(OH)6) solid slightly expands towards a less oxidising region.

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1.9 1.7 1.5 1.3 1.1 0.9

0.7 0.5 0.3 0.1 -0.1 -0.3 -0.5 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

pH 2+ Figure 83 The effect of aqueous ferrous activity (a Fe ) on Eh-pH diagram for Fe-K-S-H2O system at 220°C with 푎 0 = 0.1, 푎 − = 0.1 and 푎 + = 0.1 퐹푒퐻푆푂4푆푂4 퐻푆푂4 퐾

4.9 Summary

Through hematite solubility experiments at 220°C, current study proposed that aqueous 0 FeHSO4SO4 is the predominant ferric species at higher acidity before changes to aqueous 0 Fe2(SO4)3 as solution pH keeps increasing. The Gibbs energy of formation of aqueous 0 0 -1 FeHSO4SO4 and aqueous Fe2(SO4)3 were estimated to be -1375.7 and -2088.6 kJ mol respectively. The cross over between the two species were estimated to occur at approximately pH 1 at bisulphate activity of 0.2. Ferric speciation can only be confirmed with direct in-line speciation measurement at high temperature which is not available in current investigation.

The Gibbs free energy of formation for basic ferric sulphate solid at 220°C was estimated to be -919.4  0.1 kJ mol-1. This value is higher than the value calculated from thermodynamic properties reported in Majzlan et al. (2017) by 9%. By assessing the equilibrium pH between hematite and basic ferric sulphate, it was found that the Gibbs free energy determined in this study gave a closer estimate to the stability diagram reported in Fleming (2009).

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Similarly, the Gibbs free energy of formation for potassium jarosite solid at 220°C was estimated to be -3009.4  0.3 kJ mol-1. This value falls within the reported values in literatures (Stoffregen, 1993, Stoffregen, 2000 via HSC Chemistry v.7.1, Majzlan et al, 2010 via HSC Chemistry 9). By assessing the equilibrium pH between hematite and potassium jarosite, it was found that the Gibbs free energy determined in this study gave a closer estimate to the stability diagram reported in Babcan (1971).

Table 50 summarises the experimentally calibrated thermodynamic data which includes the Gibbs free energy of formation for basic ferric sulphate and potassium jarosite at 220°C for the development of Eh-pH diagram for Fe-S-H2O and Fe-S-K-H2O system at 220°C. These thermodynamic data have to be used as a set and cannot be exchanged with other values for the same species.

Table 50 Final set of calibrated thermodynamic data for the development of Fe-S-K-H2O system at 220°C

Species ΔG°f at 220°C (kJ/mol)

H2O (l) -207.02 H+ (a) 0 - HSO4 (a) -662.78 Fe2+(a) -70.05 K+(a) -301.04

Fe2O3 (s) -690.89

FeS2 (s) -150.73

FeSO4.H2O (s) -972.66 0 FeHSO4SO4 (a) -1375.7

Fe2(SO4)3 (a) -2088.6 FeOHSO4 (s) -919.4

KFe3(SO4)2(OH)6 (s) -3009.4

Using these thermodynamic data, Eh-pH diagram for Fe-S-H2O and Fe-S-K-H2O systems at 220°C were developed in current study. Stability field of basic ferric sulphate, potassium jarosite, hematite, pyrite and szomonolkite solids were included.

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Chapter 5: Industrial implication

5.1 Introduction

Lihir is using pressure oxidation (POX) to liberate the finely disseminated gold occluded in their sulphide minerals. It was found that the gold is concentrated in the arsenic containing pyrite i.e. arsenopyrite and arsenian pyrite and tends to be located on the outer rim of the ground particles. Due to this mineralogy, the operation has shifted from a high oxidation extent in the autoclave (total POX), to partial POX. The partial POX condition is achieved by increasing the feed rate into the autoclaves while keeping the oxygen flowrate constant. By increasing the feed rate, the autoclave residence time is reduced and so does the extent of POX. The overall plant productivity for gold, however, is increased.

Large amount of acid is being generated in the Lihir autoclave through a series of reactions where the major reactions can be represented in Reaction 5.1 to Reaction 5.6. Reaction 5.1 and 5.3 represent pyrite oxidation by oxygen and ferric respectively. Reaction 5.2 represents the ferrous oxidation to ferric due to the oxidising conditions in the autoclave. Reaction 5.4 to Reaction 5.6 represent the iron hydrolysis reactions to form hematite, basic ferric sulphate and jarosites respectively, where M can be substituted by various mono-or divalent cations + + + + + + 2+ 2+ such as K , Na , H3O , Rb , Ag , Tl , Pb and Hg . Refer to Section 1.3.1 for further information on pyrite oxidation and Section 1.3.2 for ferric hydrolysis phases.

2+ 2- + FeS 2 (s) + 7/2O2 (g) + H2O (l)→ Fe + 2SO4 (a) + 2H (a) (5.1)

2+ 3+ 4Fe (a) + O2 (g) → 4Fe (a) + 2H2O (a) (5.2)

3+ 2+ 2- + FeS2 (s) + 14Fe (a)+ 8H2O (l) → 15Fe (a) + 2SO4 (a) + 16H (a) (5.3)

Fe2(SO4)3 (a) + 3H2O (l) → Fe2O3 (s) + 3H2SO4 (a) (5.4)

Fe2(SO4)3 (a) + 2H2O (l) → 2FeOHSO4 (s) + H2SO4 (a) (5.5)

3Fe2(SO4)3 (a) + 12H2O (l) + M2SO4(a) → 2MFe3(SO4)2(OH)6 (s) + 6H2SO4 (a) (5.6)

Consequently, the environment in the autoclave becomes very acidic with pH typically below 4. Leaving the autoclave, the majority of the free acid is removed via the CCD circuit overflow by washing the slurry with a mixture of process water and seawater. The free acid concentration in the CCD underflow is controlled and set to a certain setpoint. This means the concentration of free acid entering the neutralisation stage should ideally remain constant in a perfectly controlled system and therefore reducing the oxidation extent in the autoclave should not directly affect the lime consumption in this neutralisation stage.

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Due to the high lime requirement in the neutralisation stage, its consumption has been flagged as the major contributor to the Lihir plant operating cost. The daily lime consumption at Lihir was found to significantly fluctuate from 10 to 22 kg-lime/ t-ore. Based on the plant data available from December 2013 to November 2015, the lime consumption was broken down to three categories; they are (1) lime consumed for free acid neutralisation, (2) lime consumed for the hydroxide precipitation of metal cation and (3) lime consumed by reactive solid (refer to Section 1.4.3 for further detail). Calculations and assumptions are available in Appendix B.2.

To investigate whether running the autoclave on partial Pressure Oxidation has any impact on the lime consumption, the data was split into ‘high oxidation extent’ period and ‘low oxidation extent’ period. The former was calculated based on the plant data between December 2013 and January 2014 where the averaged oxidation extent was 83%. Similarly, the latter was calculated based on the plant data between October and November 2015 where the averaged oxidation extent was 61%. The calculation results are tabulated in Table 51 and illustrated in Figure 84 and Figure 85 respectively.

Table 51 Lime consumption breakdown estimation based on Lihir plant data from December 2013 to November 2015

Averaged value High oxidation extent Low oxidation extent (kg CaO/ t-ore) (Dec’13 – Jan’14) (Oct’15 – Nov’15) Oxidation extent (%) 83 61 Total lime consumption 14.6 15.2 Acid neutralisation 0.9 2.0 Metal cation precipitation 0.3 1.3 Reactive solid 13.4 11.9

In general, the total lime consumptions over the two-month period were found to be similar regardless the difference in the oxidation extent and the variation in ore composition.

The proportion of lime used for acid neutralisation (highlighted in orange) is comparatively smaller than the overall lime consumption for both periods. However, the amount of lime used in the low oxidation extent period is approximately doubled the one in the high oxidation extent period. The difference in these values is not affected by the oxidation extent and ore composition but is a reflection of the CCD performance. Higher free acid concentration seems to present in the CCD underflow at higher feed rate i.e. lower oxidation extent, which is likely due to insufficient washing.

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Similarly, the proportion of lime used for metal cation precipitation (highlighted in yellow) is comparatively smaller to the overall lime consumption. Again, the amount of lime used in the low oxidation extent period is significantly higher than the one in the high oxidation extent period. The difference in these values, however, is affected by the oxidation extent, CCD performance and ore composition. The lime consumption for metal cation precipitation was estimated based on the iron, aluminium and magnesium concentration in the neutralisation feed. Since magnesium and aluminium are not currently monitored at Lihir, their concentrations were assumed to be half of iron concentration.

The remaining of the lime was assigned to reactive solid and it dominates the lime consumption for both periods. This indicates that reactive solid (highlighted in blue) seems to be the main factor responsible for the lime consumption at Lihir processing plant. During low oxidation extent, the estimated amount of lime consumed by reactive solid seems to be lower by an average of 1.5 kg/t. Although not conclusive due to the constant variation in the mineralogy of ore, this suggests that running the autoclave on partial POX could potentially affect the chemistry inside the autoclave. In Section 1.4.1, partial POX is suspected to cause a decrease in Fe3+: Fe2+ ratio but an increase in ferric concentration which is likely due to reduced ferric precipitation in the autoclave. With less iron being hydrolysis, less acid is being generated in the autoclave. These changes in solution acidity and iron concentration could lead to precipitation of different solid phase in the autoclave.

To investigate the effect of partial pressure oxidation on solid phase precipitated in the autoclave, this study explored the effect of oxidation extent, retention time and initial ferric/ferrous concentration individually in a series of pressure oxidation experiment using a high grade Lihir ore (NTS 200) at 220°C. Online high temperature oxidation-reduction potential (ORP) was recorded throughout the experiment to get a better understanding of the reactions undertaken in the autoclave. The results of the study were used to validate the developed Eh-pH thermodynamic model described in Chapter 3. Additionally, using the available site data and the validated Eh-pH thermodynamic model, an attempt was made to predict the solid phase precipitating in Lihir autoclave.

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Figure 84 Lihir Lime consumption breakdown estimate during high oxidation period based on Lihir plant data from December 2013 to January 2014

Figure 85 Lihir Lime consumption breakdown estimate during low oxidation period (Partial POX) based on Lihir plant data from October 2015 to November 2015

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5.2 Experimental methodology

5.2.1 Experimental plan The amount of oxygen injected into the autoclave is kept constant for all experiments. The effect of oxidation extent was investigated by varying the batch autoclave solid loading to achieve different oxidation extents. 10%, 20% and 30% solid loading were selected. At the start of the experiment, the amount of Lihir ore required was weighed before they were mixed with some DI water in autoclave liner. The overhead stirrer was turned on and was adjusted to an acceptable height to ensure that all solids were suspended in the solution. A calibrated pH probe was set in the liner to monitor the slurry pH. Diluted sulphuric acid (100 g/L) was then added into the slurry using pipettes dropper to bring the slurry pH down to pH 2. This pre-acidification step is not done at Lihir as the carbonate content in the ore is considered relatively low and the autoclave has venting capability. However, pre- acidification is required for this study to avoid carbonates build up which was found to hinder the oxygen injection as the autoclave does not have venting capability.

Once the slurry has stabilised at pH 2, the autoclave was assembled and heated up to 220°C. Once the slurry has reached the target temperature, the oxygen injection was started as discussed further in Section 5.2.2. When the oxygen injection had finished, the experiment was let to run at various residence time. The residence time at Lihir varies greatly from 1 to 2 hours, therefore three different residence time were selected for this study. They were 60, 120 and 180 minutes.

At the end of each experiment, high temperature sampling was done to collect both high temperature solid and solution samples. The high temperature sampling set-up which involves ex-situ pressure filtration is described in Section 2.2.3. The high temperature pH was estimated based on the low temperature pH using the method described in Section 4.3.4. The autoclave was then cooled down before disassembling when the slurry has reached approximately 40°C for safety reason. The remaining of the slurry in the liner was then vacuum filtered with 0.45 micron pore-size PTFE membrane filters. Solids were washed with DI water (1:1 ratio) to displace entrained solution and were oven dried at 60°C, typically for 24 hours.

All solids were characterised using SEM-EDS and XRD for particle morphology and phase identification. Filtrate was sampled and stabilised in 2wt% HCl solution for assaying by ICP- OES and pH measurement was taken.

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This study does not only investigate the effect of oxidation extent and residence time. The addition of low concentration of iron in the form of ferric and ferrous was also explored to investigate its effect on oxidation extent and solid precipitate. Table 52 summarised all experiments carried out in the current study.

Table 52 Pressure oxidation with Lihir ore experimental plan at 200°C with NTS 200 sample Solid loading, Retention Ferric/Ferrous Test description wt% time, minute addition, g/L 10 wt% solid 10 120 - 20 wt% solid 20 120 - 30 wt% solid 30 120 - 180 min RT 20 180 - 60 min RT 20 60 - 3 g/L Ferric 20 120 3 g/L Fe3+ 3 g/L Ferrous 20 120 3 g/L Fe2+

5.2.2 Oxygen injection methodology High purity (99.5%) oxygen was injected into the autoclave for Lihir ore pressure oxidation experiments. The oxygen was introduced once the autoclave had reached the target temperature, which was 220°C. The setup of the oxygen regulator and flow control are shown in Figure 86. Pressure gauge 1 indicates the pressure in the line between the cylinder and the regulator and pressure gauge 2 indicates the pressure in the autoclave. A non-return valve was installed to avoid a back flow from the autoclave to the oxygen tank. The metering valve is used to inject the oxygen slowly into the autoclave to allow time for bubbles to dissolve and react.

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Pressure Pressure gauge 1 gauge 2 Cylinder Shutoff valve valve

Pressure Metering regulator valve

Figure 86 Oxygen injection setup

When the autoclave contents reached 220°C, the oxygen line was initially pressurised by opening and closing the oxygen cylinder valve while keeping the injection valve on the autoclave closed (not shown in Figure 86). Regulated pressure was adjusted to 1 bar above the autoclave pressure before opening the injection valve.

Quantifying the amount of oxygen injected into the autoclave was found to be a challenge in this study. Due to the high operating pressure and small quantity required, a suitable oxygen flowrate meter that was within the project budget could not be found at the time of the project. Additionally, accurately weighing the oxygen consumption was not possible due to the small scale of the experiment (i.e. 2-L autoclave) and scale insensitivity.

Since the amount of oxygen injected into the autoclave was kept constant for all experiment to mimic the POX operation on site, measuring the oxygen flow rate quantitatively was deemed unnecessary. Instead, the amount of oxygen injected into the autoclave was measured indirectly by repeatedly filling and emptying the short section of line between the cylinder and regulator, and monitoring the drop on pressure gauge 1. When the pressure lowered to the regulated pressure the cylinder valve was briefly opened to refill the line, and the new full pressure was noted. A combination of adjusting the metering valve and

151 | P a g e increasing regulated pressure was used to maintain a steady low flow during each experiment.

The amount of oxygen required for all experiments was determined based on the amount of oxygen needed for experiment 1 where high oxidation extent was specified (i.e. 10 wt% solids). The oxygen was injected until no apparent pressure drop was observed in pressure gauge 1. The total pressure drop across each refill was recorded and kept constant for all other experiments.

When the oxygen injection is completed, the injection valve on the autoclave was closed to stop the oxygen flow and all other valves were also closed. The remaining oxygen in the line was released at the end of the experiments when the autoclave reached below 40°C and was ready for disassembling.

5.2.3 ORP data processing

Details on the high temperature ORP probe, reference electrode conversion, liquid junction potential (LJP) and thermal junction potential (TJP) can be viewed in Section 2.2.4.

In summary, the measured potential data (V vs. Ag/AgCl) was converted to normal hydrogen electrode (NHE) scale by adding 250 mV to the measured value as the KCl concentration in the ORP probe was 0.5 M. Thermal junction potential of +34 mV was also added to the measured value for TJP correction between working electrode which was at 220°C and reference electrode which was at a maximum temperature of 33°C.

For the liquid-liquid junction potential, the correction values depend on the concentration in the working solution. Table 53 summarises the calculated TJP based on the feed and final filtrate (25°C) concentration measured by ICP-OES shown in Table 58. The difference between the initial and final TJP was deemed to be insignificant with a maximum value of 17 mV vs. NHE.

As there was not enough information to track the changes of the TJP throughout the experiments, the increase of TJP (towards a negative value) was assumed to increase linearly from the start to the end of the experiment.

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Table 53 Thermal junction potential (mV vs. NHE) calculated from solution assay

TJP (mV vs. NHE) Test description @220°C Initial -3.34 10 wt% solid Final -15.2 Initial -3.33 20 wt% solid Final -12.87 Initial -3.32 30 wt% solid Final -10.37 Initial -2.93 180 min RT Final -13.13 Initial -2.62 60 min RT Final -11.34 Initial 2.34 3 g/L Ferric Final -15.03 Initial 1.96 3 g/L Ferrous Final -13.71

5.2.4 QXRD analysis All samples preparations were described in Section 2.2.3.1. The analysis was carried out by QUT Central Analytical Research Facility to maintain consistency with Newcrest solid interpretation. The analysis was done using JADE (V2010, Materials Data Inc.), EVA (V5, Bruker) and X’Pert Highscore Plus (V4, PANalytical) with various reference databases (PDF4+, AMCSD, COD) for phase identification. Rietveld refinement was performed using TOPAS (V6, Bruker) and the fundamental parameters approach for modelling peak shapes.

5.3 Results 5.3.1 QXRD on solid residues QXRD analysis was done on the following samples: • Ore feed • All solid residues collected at the end of each experiments (25°C samples) • ‘High oxidation’ and ‘low oxidation’ solid residues collected after oxygen injection (220°C samples) Table 54 summarises the feed ore and 25°C solid residues data for all six experiments. Table 55 compares the solid phase before and after the precipitation zone for 10 wt% solid loading and 30 wt% solid loading experiments. It is important to that note that diaoyudaoite

(NaAl11O17) was introduced into the samples as an impurity in the corundum internal standard. 153 | P a g e

Table 54 QXRD Results of ore feed and solid residues collected at 25°C as a function of the variable tested

Residues wt % Ore Feed 10 wt% 20 wt% 30 wt% 180 min*3 60 min*3 Ferric *2, 3 Ferrous*2,3 solid*2 solid*2 solid*2 Pyrite 4.9 0.9 3.1 4.2 3.7 3.6 3.7 2.6 Hematite - 4.1 2.0 1.5 2.9 2.0 3.6 4.3 K-Feldspar 57.3 52.9 52.3 62.4 56.5 56.5 54.2 54.7 Quartz 4.1 4.1 2.9 4.6 4.9 4.4 4.8 3.4 Diaoyudaoite*1 0.3 0.4 0.5 0.5 0.4 0.4 0.5 0.4 Anatase trace 0.6 0.8 0.3 0.3 0.3 0.8 0.6 Calcite 0.5 0.6 0.3 0.1 - 0.3 - - Alunite 3.2 2.5 1.2 3.1 1.9 3.4 3.4 Anhydrite 0.4 3.3 4.7 5.0 5.6 5.6 5.5 3.9 Plagioclase 1.8 3.2 2.1 2.5 2.7 3.8 2.7 3.7 Illite/mica (TOTAL) 7.3 10.5 7.8 10.7 8.8 9.0 9.3 7.8 Amorphous 19.3 16.3 21.2 6.9 11.3 12.3 11.6 15.3 Illite - 3.0 ------Biotite 1.6 ------Muscovite 1.7 1.7 ------Phengite 4.0 5.8 7.8 10.7 8.8 9.0 9.3 7.8 Illite/mica (TOTAL) 7.3 10.5 7.8 10.7 8.8 9.0 9.3 7.8

*1 Diayudaoite introduced with internal standard *2 Test were carries out with 120 minutes residence time *3 Tests were carried out with 20 wt% solid loading

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Table 55 QXRD on solid residues comparing before and after precipitation zone 10% wt solid (High oxidation) 30 wt% solid (Low oxidation) Wt% Ore Feed 220°C, t=30min 25°C, t=2h 220°C, t=10min 25°C, t=2h Pyrite 4.9 1.0 0.9 3.7 4.3 Hematite - 2.0 4.1 0.1 1.5 K-Feldspar 57.3 59.6 52.9 61.3 62.4 Quartz 4.1 4.8 4.1 5.4 5.6 Diaoyudaoite 0.3 0.4 0.4 0.4 0.5 Anatase trace 0.6 0.6 0.5 0.3 Calcite 0.5 - 0.6 - 0.1 Alunite - - 3.2 - 1.2 Anhydrite 0.4 2.9 3.3 4.5 5.0 Plagioclase 1.8 2.0 3.2 2.0 2.5 Illite/mica (TOTAL) 7.3 10.0 10.5 10.0 10.7 Smectite - - - 0.2 - Amorphous 19.3 16.8 16.3 12.1 5.9 Illite - - 3.0 - - Biotite 1.6 - - - - Muscovite 1.7 3.7 1.7 - - Phengite 4 6.3 5.8 10.0 10.7 Illite/mica (TOTAL) 7.3 10.0 10.5 10.0 10.7

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The estimations for pyrite and hematite in QXRD are likely to be underestimated. This can be seen from the big discrepancy in the pyrite mass fraction of ore feed between QXRD and MLA results. The MLA results shows that Lihir ore comprised of ~65 wt% K-feldspar, ~15 wt% pyrite, ~5.8 wt% illite and ~5.5 wt% quartz. Traces of arsenopyrite (~0.05 wt%) and chalcopyrite (~0.04 wt%) were found. Other non-sulphide gangue present in minor to trace quantities included iron oxide (~0.3wt%), rutile (~0.4 Wt%), calcite and dolomite (~0.8 Wt% each), anhydrite (~0.3 Wt%), apatite (~0.6 Wt%), albite (~0.6 Wt%), biotite (~1.6 Wt%), diopside (~0.5 Wt%), muscovite (~1.0 Wt%), plagioclase (~0.9 Wt%), and vermiculite (~0.6 Wt%).

There is a discrepancy of approximately 10wt% in pyrite between QXRD and MLA results. The following reasons could be the main contributor to this discrepancy: • Microabsorption contrast due to a mix of low and high absorbing phases In general, higher elements absorb more beam in the surface of the grain and thus only a fraction of the grain diffracting. Pyrite and hematite are considered as higher absorbing minerals in the Lihir solid residue. In contrast, beam penetrates further into grain in the low absorbing minerals and as a result there is a greater likelihood of ‘volume diffraction’ occurring. Some examples of low absorbing minerals in the residues are corundum (internal standard), K-feldspar and illite/mica. The absorptions for pyrite and hematite are approximately ten times higher than those for low absorber minerals. Figure 87 illustrates the difference in the beam absorption for high and low absorbing minerals (Madsen, 2017).

Figure 87 Illustration of beam absorption in high and low absorbing minerals (Madsen, 2017)

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Due to the large contrast in the microabsorption, the intensities of the pyrite and hematite tend to be under-estimated and consequently low phase quantification was reported. • Inaccuracy in MLA technique associated with amorphous phase The MLA was done without internal standard. This means the proportion of the amorphous solid was distributed among the other solid phases. If the amorphous contains Fe and S, it would be assumed to be pyrite and thus the pyrite content was over-estimated.

Despite this discrepancy, the QXRD value can still be used to estimate the relative difference of pyrite content between each test conditions.

To calculate the mass of each of the minerals, the total mass of the solid in the feed and residues are summarised in Table 56. Note that some solids were lost in the high temperature sampling line and during the removal of autoclave head.

Table 56 Mass of solid in feed, residues 220°C and residues 25°C

Test description Solid feed (g) Solid Residue Solid Residue 220°C (g) 25°C (g) 10 wt% solid 130 15 108 20 wt% solid 260 36 207 30 wt% solid 390 64 289 180 min RT 260 42 204 60 min RT 260 30 201 3 g/L Ferric 260 41 204 3 g/L Ferrous 260 9 230

5.3.2 Solution assay Solution assay was obtained using ICP-OES on three different samples: • Feed solution which was sampled after pre-acidification; • Filtrate at 220°C which was sampled after oxygen injection – sampling time is shown in Figure 89 to Figure 91; • Filtrate at 25°C which was sampled at the end of the experiments.

The results of redox titration used to determine ferrous concentration are summarised in Table 57. Note that the redox titration was duplicated for each sample and the volume of

K2Cr2O7 solution reported is the averaged values of the two readings. In most samples, the

157 | P a g e two readings were found to be consistent while for some samples a variation of 0.05 to

0.1mL in the volume of K2Cr2O7 used was observed.

Table 57 Redox titration results for ferrous concentration calculation

Volume Sample Sample Sample [Fe2+] Test condition K2Cr2O7, mL mass, g Density, g/mL Volume, mL g/L 10 wt% solid 220°C 0.25 10.48 1.01 10.4 0.13 25°C 2.0 10.10 1.01 10.0 1.12 20 wt% solid 220°C 2.0 10.38 1.01 10.3 1.09 25°C 1.33 10.21 1.02 10.1 0.74 30 wt% solid 220°C 2.45 10.16 1.01 10.1 1.36 25°C 1.1 10.19 1.02 10.0 0.61 180 min RT 220°C 2.45 10.29 1.01 10.2 1.34 25°C 1.68 10.34 1.01 10.3 0.91 60 min RT 220°C 2.0 10.06 1.01 10.0 1.12 25°C 1.25 10.14 1.01 10.0 0.70 3 g/L Ferric 220°C 1.95 10.14 1.01 10.0 1.09 25°C 2.45 10.20 1.02 10.0 1.36 3 g/L Ferrous 220°C 1.70 10.22 1.01 10.1 0.94 25°C 2.34 10.22 1.01 10.1 1.31

Table 58 summarises the major elements present in these samples which are iron (Fe), sulphur (S), potassium (K), aluminium (Al), magnesium (Mg), calcium (Ca) and silica (Si).

Ferrous concentration was estimated by redox titration using a K2Cr2O7 solution as described in Section 2.4. Ferric concentration was calculated as the difference between total Fe values obtained from ICP-OES analysis and the ferrous concentration. Due to the low concentration of ferric in the solution, the accuracy of [Fe2+]/[Fe3+] ratio is highly affected by the accuracy of ICP-OES assay and ferrous titration.

Silver (Ag), Cadmium (Cd), Cobalt (Co), Chromium (Cr), Copper (Cu), Manganese (Mn), Molybdenum (Mo), Sodium (Na), Nickel (Ni), Lead (Pb), Strontium (Sr), Titanium (Ti) Vanadium (V), and Zinc (Zn) were also present but all concentration were below 0.1 g/L.

Based on the solution assay of the feed samples in Table 58, Calcium (Ca) was released into the solution during the ore pre-acidification step. This is likely from the calcium carbonate reaction with sulphuric acid according to reaction 5.7 below as small gas bubbles were observed to be formed during this step.

CaCO3(s) + H2SO4(a) → CaSO4 (a) + CO2(g) + H2O (l) (5.7)

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Additionally, a small portion of pyrite also seemed to be leached which was indicated by a concentration reading in iron (Fe) and sulphur (S). A minimum amount of mica dissolution also seemed to occur which was indicated by a low reading of aluminium (Al), silica (Si), magnesium (Mg) and potassium (K).

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Table 58 ICP-OES of feed solution (after pre-acidification), filtrate at 220°C (after oxygen injection) and filtrate at 25°C

3+ 2+ Eh (mV vs. NHE) FeT Fe Fe S K Al Mg Si Ca Test ID pH @ 220°C Major Element Concentration, g/L Feed 0.11 - -* 0.70 0.04 0.01 0.08 0.03 0.88 10 wt% Ore 1 220°C 1.25 0.14 3E-3 0.13 3.05 0.22 0.13 0.47 0.91 0.28 solid 25°C 1.21 620 1.13 1E-2 1.12 4.9 0.18 0.07 0.64 0.72 0.56 Feed 0.11 - -* 0.62 0.06 0.03 0.17 0.02 0.56 20 wt% Ore 2 220°C 1.37 1.12 4E-2 1.11 4.03 0.24 0.10 0.64 0.96 0.23 solid 25°C 1.28 490 0.84 2E-3 0.83 5.49 0.24 0.10 1.2 0.72 0.61 Feed 0.17 - -* 0.74 0.10 0.04 0.27 0.02 0.53 30 wt% Ore 3 220°C 1.39 1.36 8E-3 1.35 5.02 0.32 0.11 0.85 0.98 0.22 solid 25°C 1.39 420 0.62 7E-3 0.61 5.78 0.35 0.11 1.74 0.72 0.59 Feed 0.11 - -* 0.78 0.07 0.03 0.17 0.02 0.59 180 min Ore 4 220°C 1.39 1.37 3E-2 1.36 4.19 0.25 0.09 0.65 0.97 0.24 RT 25°C 1.29 440 0.92 2E-3 0.91 5.19 0.28 0.07 1.26 0.73 0.56 Feed 0.13 - -* 0.85 0.06 0.03 0.17 0.02 0.56 60 min Ore 5 220°C 1.39 1.14 3E-2 1.13 3.53 0.27 0.10 0.65 0.95 0.24 RT 25°C 1.37 510 0.71 1E-2 0.70 4.89 0.27 0.12 1.08 0.84 0.59 3 g/L Feed 2.4 - -* 2.9 0.07 0.06 0.17 0.02 0.55 Ore 6 Ferric 220°C 1.25 1.13 4E-2 1.12 5.74 0.27 0.08 1.1 1.05 0.30 addition 25°C 1.16 460 1.37 1E-2 1.36 6.68 0.27 0.05 1.3 0.65 0.61 3 g/L Feed 2.57 - -* 2.6 0.09 0.04 0.16 0.02 0.47 Ore 7 Ferrous 220°C 1.26 0.96 1E-2 0.95 5.98 0.31 0.11 0.89 0.99 0.29 addition 25°C 1.16 460 1.32 5E-3 1.31 7.68 0.32 0.06 1.3 0.69 0.58

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5.3.3 Solid assay Solid analysis was done on the following samples: • Ore feed • All solid residues collected at the end of each experiments (25°C samples) • All solid residues collected after oxygen injection (220°C samples)

All solid samples were digested by microwave digester using a 5:3:2 volume ratio of 70%

HNO3, 37% HCl and 48% HF acids. The digested samples were then analysed by ICP-OES and the results are summarised in Table 59.

Table 59 Solid assay from pressure oxidation ore test experiments

Fe K Al Si S Ca Na Ti Sample Description wt% Feed 36.8 0.28 1.25 3.1 51.6 0.38 n/a n/a 10 wt% solid 220°C 5.17 9.35 4.99 27.01 2.33 0.70 0.44 0.48 25°C 6.61 8.31 3.20 23.50 3.13 0.42 0.37 0.67 220°C 5.95 8.35 3.36 23.79 4.67 1.05 0.37 0.49 20 wt% solid 25°C 8.41 8.10 3.34 23.10 6.17 0.98 0.35 0.57 220°C 6.70 8.26 3.23 23.16 5.45 0.74 0.35 0.50 30 wt% solid 25°C 8.65 8.15 3.05 23.54 7.25 0.91 0.36 0.50 180 min 220°C 7.46 8.02 3.47 23.44 5.81 1.14 0.34 0.55 Residence time 25°C 8.41 8.00 5.18 22.38 6.47 1.50 0.33 0.68 60 min 220°C 7.16 7.72 3.16 22.44 5.64 1.17 0.34 0.50 Residence time 25°C 8.12 7.29 3.38 21.35 6.58 1.31 0.34 0.62 3 g/L Ferric 220°C 7.72 8.22 3.64 23.05 5.72 1.07 0.36 0.52 addition 25°C 9.39 7.81 4.62 21.88 6.91 1.44 0.31 0.61 3 g/L Ferrous 220°C 6.29 9.16 3.75 25.46 3.56 0.38 0.42 0.44 addition 25°C 8.10 8.60 3.74 24.04 4.99 0.82 0.37 0.51 Magnesium (Mg), silver (Ag), cadmium (Cd), cobalt (Co), copper (Cu), manganese (Mn), molybdenum (Mo), nickel (Ni), lead (Pb), strontium (Sr), vanadium (V) and zinc (Zn) were present in traces amount.

5.3.4 ORP Data The oxidation-reduction potential (ORP) data for each test was recorded using the high temperature ORP probe described in Section 2.2.4. The obtained values were adjusted for temperature junction potential (TJP) and liquid junction potential (LJP) as described in 161 | P a g e

Section 2.2.4. The raw data for all oxidation-reduction potential measurement and sample calculations for LJP can be viewed in Appendix B.1.

In this system, the ORP is mainly governed by the potential of Fe3+/Fe2+. Therefore the behaviour of the electrode potential can be explained in terms of ferric and ferrous concentrations in the solution.

Despite of the complex nature of pyrite oxidation, pyrite oxidation by oxygen (Reaction 5.1), pyrite oxidation by ferric (Reaction 5.3), ferrous oxidation to ferric (Reaction 5.2), and iron hydrolysis (Reaction 5.4 to 5.6) were considered to be the main reactions affecting the potential of the solution in the autoclave. These four reactions are occurring concurrently and some are more dominant than other

Based on the oxidation-reduction potential data measured through the experiments, the shape of the potential profile as a function of time is similar for all cases. Figure 88 shows the general illustration of the electrode potential measured as a function of time in the ore pressure oxidation experiment. In general, the potential profile can be divided into three different regions: oxygen injection, oxygen depletion and stabilisation.

The oxygen injection region (Region 1) starts as the oxygen is injected into the autoclave which causes a significant increase in the ORP instantly. This increase in oxidation- reduction potential indicates that there is an increase of [Fe3+]/[Fe2+] in the solution from pyrite oxidation. Despite being generated as the direct product of pyrite oxidation (Reaction 5.1 and Reaction 5.2), ferrous is continuously oxidised to ferric according to Reaction 5.3. However, at the same time ferric is either used as pyrite surrogate oxidant (Reaction 5.2) or removed from the solution via iron hydrolysis to form iron precipitate (Reaction 5.4). The latter could result in a drop in ferric concentration and hence a decrease in the potential. However, since the overall potential increases, the terminal Fe3+/Fe2 ratio is be higher than prior to oxidation.

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Figure 88 Illustration of electrode potential as a function of time in POX test

When the oxygen supply was ceased (Region 2), a prompt decrease in the oxidation- reduction potential was observed due to the drop in dissolved oxygen concentration and reduction of ferric to ferrous in solution. Dissolved oxygen was quickly depleting as it was consumed by Reaction 5.1 and Reaction 5.3. As less oxygen was available to oxidise ferrous, the concentration of ferrous kept increasing while the ferric concentration is going down due to precipitation and pyrite oxidation. Reaction 5.2 and Reaction 5.4 seemed to be the governing reactions as oxygen availability was diminishing.

Once all of the available oxygen has been used up, no ferrous can be further oxidised to ferric. This means no more ferric is being generated in the system and therefore its concentration is continually decreasing through iron precipitation (Reaction 5.4) and pyrite oxidation (Reaction 5.2). Reaction 5.4 stops once the ferric concentration approaches its solubility concentration with respect to the most stable precipitate. On the other hand, Reaction 5.5 should still be occurring while ferric is still present in the solution. However, due to the low concentration of ferric at this stage of the experiment, the extent of pyrite oxidation by pyrite is expected to be minimal based on the finding in Chapter 2. The change in ferric and ferrous concentrations should be minimal and therefore the oxidation-reduction potential should stay relatively constant (Region 3).

Figure 89 shows the oxidation-reduction potential measured during POX of Lihir ore at different solid loadings. The test IDs evaluated are Ore 1, Ore 2 and Ore 3. The potential profile shape for Ore 1 - 10 wt% solid loading is slightly different from the general shape 163 | P a g e illustrated in Figure 88 where the significant drop of oxidation-reduction potential in Region 2 was not observed. This is because the oxygen was continually injected into the autoclave until no apparent pressure drop in the oxygen supply was observed to ensure high pyrite oxidation extent. When the oxygen injection was stopped, the oxidation-reduction potential slightly increased which indicates pyrite oxidation (Reaction 5.1) had stopped momentarily and ferrous oxidation (Reaction 5.3) was dominating which results in the increase of Fe3+/Fe2+ ratio. This is suggested by the low ferrous concentration in the 220°C residues for 10 wt% solid loading in Table 58. After about 10 minutes, however, the oxidation-reduction potential started to decrease gradually until the end of the experiment. This was likely due to the removal of ferric from the solution indicated by the increase of hematite in the residues from the 220°C to 25°C samples (Table 55) and the generation of ferrous from the oxidation of ‘slow reacting’ pyrite.

Despite of the slightly different shape of the Ore 1 - 10 wt% solid loading experiment, it is clear that the terminal potential was higher at higher solid loading given that the same amount of oxygen was injected into the autoclave. This is because the Ore 1 - 10 wt% solid loading test had the most amount of oxygen relative to the pyrite.

Figure 89 Measured electrode potential during POX of Lihir gold-bearing sulphide ore at 220°C at different solid loadings (Test ID: Ore 1 to Ore 3) with 120 minutes retention time

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Figure 90 shows the electrode potential measured during POX of Lihir ore at three different retention time. Test IDs evaluated are Ore 5 - 60 minutes, Ore 2 - 120 minute and Ore 4 - 180 minutes. All three experiments were carried out with 20 wt% solid loading. Since the amount of oxygen injected and the solid loading were the same for all three experiments, the shape of the oxidation-reduction potential were similar. However, the terminal potential was slightly different for all three tests because it was still continuously dropping after 60 minutes, which indicates continuous increase of the Fe2+/Fe3+ ratio in the solution.

As discussed earlier, this marginal increase in Fe2+/Fe3+ ratio could be due to the increase of ferrous or the removal of ferric by pyrite oxidation and iron hydrolysis. Based on the ICP- OES and redox titration of the final filtrate (Table 57), ferrous concentration was found to be higher at longer residence time which suggest continuous ferrous generation from pyrite oxidation. Reaction 5.2 and Reaction 5.4 are both acid generating reactions, but the latter produces more acid than the earlier reaction. From acidity perspective, the 60-minute test had a slightly higher pH compared to the other two tests at pH 1.37 where the pH of Ore 2- 120-minute and Ore 4 - 180-minute tests were very similar at pH 1.28 and 1.29. The pH difference is quite small and is still within the probe accuracy of 0.1 pH unit. Therefore, it is not conclusive but suggests that iron hydrolysis seems to have ceased after 60 minutes.

The terminal potential for the Ore 5 - 60-minute test was 510 mV vs. NHE, while those for Ore 2 - 120-minute and Ore 4 - 180-minute tests were 490 and 440 mV vs. NHE respectively. The oxidation-reduction potential for the first two tests were pretty similar with only 20 mV difference, but there was a substantial difference between the Ore 2 - 120-minute and the Ore 4 - 180-minute tests. This was likely to be caused by a significant drop in the potential at approximately minute 44 in the Ore 4 - 180-minute test. The reason for this drop remains unclear.

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Figure 90 Measured electrode potential during POX of Lihir gold-bearing sulphide ore at 220°C at different retention time and 20 wt% solid loading (Test ID: Ore 2, Ore 4, Ore 5)

Figure 91 shows the electrode potential measured during POX of Lihir ore at different initial ferric and ferrous concentration. The test IDs evaluated are Ore 2 – no ferrous/ferric addition, Ore 6 – 3g/l ferric addition and Ore 7 – 3g/l ferrous addition. All three experiments were carried out with 20 wt% solid loading. The electrode potential in Ore 7 experiment, where 3 g/L ferrous was added to the slurry, was lower at all times in Region 1 and Region 2 compared to the tests without any addition of ferric/ferrous (Ore 2) and with initial addition of 3 g/L (Ore 6). This is due to the higher Fe2+/ Fe3+ in solution in Ore 7 test. In contrast, the addition of 3 g/L ferric seems to increase the oxidation-reduction potential due to the higher Fe3+/ Fe2+ in solution in Ore 6 test.

Despite of the difference in oxidation-reduction potential in Region 1 and Region 3, all three experiments seem to converge to the same potential after about an hour before the potential suddenly drop in Ore 6 and Ore 7 tests. This drop is likely due to further hydrolysis of ferric as hematite proportion in solid residues of Ore 6 and Ore 7 tests were almost doubled than that in solid residue of Ore 2 test (Table 54). Therefore the terminal potential in the absence of initial ferric/ ferrous (Ore 2) was slightly higher (Eh = 490 mV vs. NHE) than the other two experiments which had terminal potentials of approximately 460 mV vs. NHE.

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Figure 91 Measured electrode potential during POX of Lihir gold-bearing sulphide ore at 220°C at different initial ferric and ferrous concentration (Test ID: Ore 2, Ore 6, Ore 7)

5.2 Discussion

From the QXRD of the solid residues in Table 54, Quartz, K-Feldspar and plagioclase seems to remain inert during the pressure oxidation process. In contrast, pyrite is readily oxidised when oxygen is injected into the autoclave. This was indicated by the prompt increase of solution temperature which triggered the autoclave cooling water system. A maximum of 223°C was observed in some experiments but went back down to 220°C due to the cooling water. Since the same amount of oxygen was injected into the system for all experiments, the total amount of pyrite oxidised should theoretically be similar. Using the pyrite mass fraction (Table 54) and solid mass of feed and residues (Table 56), a total of approximately 5 grams of pyrite were oxidised. This translates to 83%, 41% and 26% pyrite oxidation for 10 wt%, 20 wt% and 30 wt% solid loading tests respectively. As The addition of ferric and ferrous as surrogate oxidant did not seem to boost the pyrite oxidation due to its low concentration.

The consumption of oxygen for pyrite oxidation and ferrous oxidation seemed to occur relatively fast. In current investigation, oxygen was consumed within approximately 5 minutes after the injection was stopped. This is confirmed by the similar pyrite wt% in solid residues between 220°C and 25°C summarised in Table 55. This means pyrite oxidation 167 | P a g e reaction is the dominated reaction in the first few chambers of a continuous pressure oxidation operation.

In the pressure oxidation experiments, hematite, alunite and anhydrite (Ca sulphate) were the three main precipitated solid phase by QXRD. Of these three solids, hematite was the only iron phase formed in the autoclave. From the QXRD of solid residues sampled at 220°C and 25°C summarised in Table 55, the amount of hematite was found to increase significantly. This suggests that in the last few chambers of the autoclave, iron precipitation and ferrous oxidation reactions are the dominant reactions. This is further confirmed by the operator observation on site where turning down the oxygen injection in the back end of the autoclave changes the discharge slurry colour from red to green. The red colour comes mainly from hematite which can be mixed with other solid phases, while the green colour comes from ferrous. When oxygen injection is reduced, ferrous oxidation to ferric is hindered and consequently increasing the ferrous concentration while reducing the amount of ferric that undergoes hydrolysis. It is understood that ferrous has low solubility at high temperature where it precipitates as white ferrous sulphate monohydrate. However, this solid readily dissolves as temperature decreases which gives off the green colour observed by operator.

From the solid loading experiment, the amount of hematite, alunite and anhydrite precipitated out of the solution remain relatively constant regardless of different oxidation extent. The solution pH, however, seems to be higher at higher solid loading. From the residence time experiments, the kinetic of hematite precipitation seems to occur relatively fast and little to no change occurred after 60 minutes as discussed in Section 4.2.4. However, it is not the same case for alunite precipitation. Based on the alunite wt% in the 25°C solid residues in Table 54, the concentration of alunite increases at longer residence time. This is likely to be the cause of the higher acidity (i.e. lower pH) in the longer residence time experiment as the amount of pyrite oxidised and hematite precipitation were relatively similar. From the iron addition experiment, the addition of ferric and ferrous seem to increase the amount of hematite and alunite precipitated during the pressure oxidation experiment. These precipitation reactions seem to cause the low pH of the final solution.

Potassium jarosite which is typically found in the discharge slurry of Lihir autoclave (Section 1.4.2) was not found in the solid residues of current investigation pressure oxidation experiment. The absence of potassium jarosite was likely due to the low potassium content in the ore. Biotite has been identified as the main source of potassium contained in Lihir ore. This mineral is readily dissolved during pressure oxidation, in part or totally (Kathryn, 2019, personal conversation). The typical biotite concentration at Lihir feed ore on site is 168 | P a g e approximately 8 wt% while in NTS200 Lihir ore which was used in current investigation has low biotite concentration with only approximately 1 wt%. Consequently, the concentration of potassium released into the solution was also low with a maximum concentration of 0.35 g/L K+ (refer to Table 58) and therefore potassium jarosite formation is not favoured.

5.3 Eh-pH thermodynamic model validation

Based on the measurements, using the final concentrations of filtrate 25°C for all six experiments summarised in the in Table 58, the Eh-pH diagram developed in Chapter 3 will be validated by predicting the possible stable solids. The following assumptions were made to define the conditions that need to be inputted into the model:

• All activity coefficients were assumed to be unity;

0 • All ferric are in the form of aqueous FeHSO4SO4 at high temperature; • All soluble sulphur are either associated with ferric complex species or in the form of bisulphate at high temperature; • All potassium are present as K+ at high temperature.

The concentrations of ferric complex, ferrous, bisulphate and potassium selected for the model inputs were the averaged concentration from all six experiment based on the filtrate 25°C assay in Table 58. The selected concentrations are summarised in

Table 60 Selected concentrations for Eh-pH diagram validation

Selected Concentrations (g/L) Fe3+ (a) 0.01 Fe2+ (a) 0.8 - HSO4 (a) 6 K+ (a) 0.3

The six experiments data were also plotted on the Eh-pH diagram using the oxidation- reduction potential and pH summarised in Table 58. The pH at 220°C is assumed to be similar to the measured pH at 25°C. Figure 92 shows the Eh-pH diagram for the six pressure oxidation experiments carried out at 220°C using the NTS200 Lihir ore. Based on this Eh- pH diagram, potassium jarosite is indeed predicted to be unstable and all six POX experiments are in the stability field of hematite which is in agreement with the observed solid phase in the autoclave discharged residues. This means the developed Eh-pH thermodynamic data can correctly predict the solid phase precipitated in the autoclave in these instances.

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Figure 92 Eh-pH diagram for Lihir ore POX experiment at 220°C

If the potassium source in the NTS200 ore was higher and as a result more potassium were released into the solution, potassium jarosite formation would likely precipitate in the current investigation POX conditions. Using the same ferric, ferrous and bisulphate concentration tabulated in Table 60 but with ten times higher potassium concentration, another Eh-pH diagram was generated and shown in Figure 93. At higher potassium concentration of 3 g/L K+, potassium jarosite is predicted to be the most stable from approximately pH 0.5 to 1.5 where hematite is the predominant solid at higher pH and basic ferric sulphate is the predominant solid at lower pH.

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Figure 93 Eh-pH diagram for high potassium condition at 220°C on Lihir Ore NTS200

If the Eh and pH for all experimental data remain constant, some of them are in the potassium jarosite stability field while the others are in the hematite stability field. Despite of this, potassium jarosite would likely be detected in the solid residues for all experiments because they are close to the potassium jarosite-hematite equilibrium line. This means some potassium jarosite would still be precipitated although hematite is the predominant solid.

5.4 Application of Eh-pH model using Lihir site data

To further examine the performance of the developed Eh-pH diagram, some Lihir site data are plotted on the diagram to predict the predominant solid phase precipitating in Lihir autoclave. Table 61 summarises the autoclave discharge site data obtained from all four autoclaves, both from the night shift and day shift on 31st January 2017. Note that potassium concentration is not reported on site and therefore potassium concentration is unknown. It is also understood that the measured pH data were collected on a filtered and cooled sample. However, there is insufficient data on the solution to estimate the pH at high temperature and therefore all measured pH data are assumed to be relatively similar to pH at high temperature.

To generate the Eh-pH diagram, the potassium concentration was selected to be 3 g/L which is ten times the potassium concentration measured in the pressure oxidation carried out in current investigation. For the ferric and ferrous concentration, 1 g/L Fe3+ and 2 g/L Fe2+ were 171 | P a g e selected. Due to the vast variations in the free acid concentration in the autoclave discharge, three different diagram were generated at 4.3, 9 and 20 g/L H2SO4 as shown in Figure 94, Figure 95 and Figure 96 respectively. The same assumptions as the one listed in Section 5.3.2.2 are applied, except that the bisulphate concentration is assumed to be equal to free acid concentration (H2SO4).

Table 61 Lihir autoclave discharge slurry condition on 31st January 2017

Filtrate, g/L pH Oxidation ORP (mV) 2+ 3+ measured Extent (%) Fe Fe H2SO4 AC1 1.96 0.08 5.4 1.87 376 44 AC2 2.26 0.35 4.3 1.95 364 48 Night shift AC3 2.26 0.66 7.4 1.67 395 63 AC4 1.89 0.96 6.6 1.52 410 64 AC1 2.36 0.77 8.9 1.47 412 67 AC2 2.26 0.60 9.0 1.44 406 68 Day shift AC3 1.84 0.62 17.3 1.16 418 75 AC4 1.66 0.79 21.0 1.1 408 78

Figure 94 shows the Eh-pH diagram at 220°C for the Lihir autoclave discharge conditions at

4.3 g/L H2SO4. At this low free acid concentration, potassium jarosite is unstable at all pH range given that 3/L K+ was specified. Based on the autoclave discharge pH and oxidation- reduction potential shown in Table 61, hematite is predicted to be the predominant solid precipitated in the autoclave. It is located far off from the basic ferric sulphate stability field and therefore it is unlikely for basic ferric sulphate to form.

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1.5

1.3

1.1

0.9

0.7

0.5

0.3

0.1

-0.1

-0.3 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

pH 3+ 2+ + Figure 94 Eh-pH diagram at 220°C. [Fe ] = 1 g/L, [Fe ] = 2 g/L, [K ] = 3 g/L, - [HSO4 ] = 4.3 g/L

As the free acid concentration is increased to 9 g/L H2SO4, potassium jarosite stability starts to appear as shown in Figure 95. The autoclave discharge pH is shifting towards a lower value as the acidity increases. At this condition, hematite is again predicted to be the predominant solid precipitated in the autoclave.

1.5

1.3

1.1

0.9

0.7

0.5

0.3

0.1

-0.1

-0.3 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

pH 3+ 2+ + Figure 95 Eh-pH diagram at 220°C. [Fe ] = 1 g/L, [Fe ] = 2 g/L, [K ] = 3 g/L, - [HSO4 ] = 9 g/L 173 | P a g e

When the free acid concentration reaches 20g/L H2SO4, potassium jarosite is stable from + pH approximately 0.7 to 1.4 at 3 g/L K as shown in Figure 96. At this free acid concentration, the autoclave discharge pH falls to approximately pH 1.1 and therefore potassium jarosite is predicted to be the predominant solid precipitating in Lihir autoclave.

1.5

1.3

1.1

0.9

0.7

0.5

0.3

0.1

-0.1

-0.3 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

pH 3+ 2+ + Figure 96 Eh-pH diagram at 220°C. [Fe ] = 1g/L, [Fe ] = 2 g/L, [K ] = 3g/L, - [HSO4 ] = 20g/L

If there is no potassium in the ore i.e. [K+] = 0 g/L, hematite remains to be the predominant solid predicted as shown in Figure 97. The autoclave discharge condition is relatively close to the basic ferric sulphate stability field and therefore there is a possibility for this solid phase to form at this condition especially at free acid concentration above 35 g/L.

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1.5

1.3

1.1

0.9

0.7

0.5

0.3

0.1

-0.1

-0.3 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0

pH 3+ 2+ + Figure 97 Eh-pH diagram at 220°C. [Fe ] = 1 g/L, [Fe ] = 2 g/L, [K ] = 0 g/L, - [HSO4 ] = 20 g/L

5.7 Industrial implication

From these examples, it can be seen that the predominant solid phase precipitating inside the autoclave depends highly on the free acid concentration in the autoclave. Figure 98 shows the spread of Lihir autoclave discharge free acid concentration for all four autoclaves from November to December 2015. There were a few times when the autoclave discharge free acid concentration went above 35 g/L H2SO4 and even reached 50 g/L H2SO4. At these conditions, the precipitation of basic ferric sulphate, which is known to cause high lime consumption, is not definite but possible depending on the potassium concentration in the autoclave.

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Figure 98 Histograms of Lihir autoclave discharge free acid concentration from November to December 2015 Because the potassium concentration at Lihir autoclave discharge is unknown, these predictions presented in Section 4.3.5 may not be accurate but it can give a good indication of which solid can precipitate in Lihir autoclave. In general, hematite seems to the one of the main solid phases precipitated in Lihir autoclave. However, it is well established that this solid phase will not react with lime due to its stability. Therefore, the following solids are the possible lime consumers at Lihir: • Potassium jarosite Potassium jarosite is suspected to be the main lime consumers at Lihir. Apart from hematite, potassium jarosite are likely to be the second most abundant solid phase precipitated in the autoclave. Ji et al. (2006) claimed that the kinetic of alkaline decomposition of potassium jarosite with lime was slow. This is perhaps valid for perfectly crystalline synthetic potassium jarosite with a relatively large aggregates diameter (>10 µm). The presence of induction period which is indicated by the halo formation around the particle may contribute to the slow kinetic of potassium jarosite dissolution (Patino et al., 2013; Reyes et al., 2016). However, the potassium jarosite precipitate in Lihir autoclave may be vastly different than those tested in previous studies due to one or more of the following reasons: 1. Majority of the potassium jarosite in the Lihir autoclave discharge are present in the < 6 micron size fractions (Newcrest Mineralogy report, 2016). With a smaller crystallite size of potassium jarosite, the dissolution reaction rate is

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faster. Reyes et al. (2016) found that decreasing the particle does not affect the induction period but if the initial particle size is halved, the reaction rate constant is approximately doubled. 2. The potassium jarosite formed in Lihir autoclave is possibly not completely crystalline or have defects on the crystal surface. Patino et al. (2013) found that the presence of defects on the crystal diminishes the induction period while the reaction rate constant remains constant. This, however, cannot be confirmed in current investigation as potassium jarosite was not formed in the pressure oxidation experiment. • Basic ferric sulphate Considering the occasional high free acidity concentration in the autoclave discharge, the formation of basic ferric sulphate is possible. However, its presence was not detected in the Lihir autoclave discharge during Newcrest mineralogy study in 2016. This could be due to the sample preparation procedure undertook on site. Basic ferric sulphate is known to be unstable at low temperature (< 140°C) and acidic conditions (pH< 1) where it is prone to undergo decomposition to form ferric sulphate according to Reaction 5.8 below.

2FeOHSO4(s) + H2SO4(a) → Fe2(SO4)3(a) + 2H2O(l) (5.8) If the slurry sample was not filtered immediately to separate the solid residue from the acidic solution, basic ferric is readily dissolving.

• Alunite Another possible lime consumer is alunite which was found to precipitate in current investigation pressure oxidation experiments. When the aluminium content in the ore is higher, alunite precipitation is becoming more favoured.

5.8 Summary

Pressure oxidation on Lihir NTS200 was carried out at six different conditions investigating the effect of solid loading, retention time and addition of ferric/ferrous. Hematite was the only solid precipitated in all experiment due to the low biotite concentration which is the main potassium source in the ore. Using these experimental data, Eh-pH diagram has been validated.

An attempt was made to use the Eh-pH diagram to predict the solid precipitating in Lihir autoclave using the available site data. Since potassium concentration is not currently being 177 | P a g e monitored onsite, an accurate and conclusive prediction cannot be made in current investigation. However, at the selected potassium concentration of 3 g/L K+ with free acid concentration ranging from 4 g/L to 20 g/L H2SO4, hematite and potassium jarosite are the two main solid phases likely to be precipitated in Lihir autoclave. This result is in agreement with the observed solid phase on site. Therefore potassium jarosite is suspected to be the main lime consumer at Lihir due to its small crystallite size and possibly lack of crystallinity and surface defect.

The formation of basic ferric sulphate in Lihir autoclave could also occur when the free sulphuric acid in the autoclave exceeds 35 g/L. If BFS presents in Lihir residue, 1 wt% of this solid phase could cause an additional lime consumption of 1.8 kg lime per tonne of slurry assuming 55wt% solid in the autoclave discharge. Therefore, BFS could contribute to the high lime consumption at Lihir.

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Chapter 6: Conclusion and recommendations

6.1 Conclusion • Results from the role of ferric as pyrite surrogate investigation show that ferric was able to oxidise pyrite at gold pressure oxidation temperature eventhough some of the ferric in the solution was removed through iron precipitation. Pyrite oxidation by ferric was evident even at low ferric concentration where the lowest initial ferric concentration tested was 2.7 g/L Fe3+. • Pyrite oxidation by ferric was found to be surface limited. When the reaction is not yet limited by the available surface are, the pyrite oxidation extent was found to increase linearly as a function of ferric concentration. When the reaction is surface limited, additional ferric introduced into the system has a higher tendency to precipitate (i.e. higher proportion of ferric that undergoes hydrolysis). Consequently, increasing initial ferric concentration was barely improved the pyrite oxidation extent. • Presence of ferrous was found to be hinder the pyrite oxidation by ferric. There are a few possible mechanisms proposed on how ferrous is affecting the pyrite oxidation, this includes surface adsorption, pyrite passivation and equilibrium effect. However, the negative effect of ferrous on pyrite oxidation was found to be eased by increasing the pyrite solid loading (i.e. surface area). • From the hematite solubility experiments at 220°C, this study found that aqueous 0 FeHSO4SO4 species fit the solubility data well at higher acidity while aqueous 0 Fe2(SO4)3 species fit the solubility data at lower acidity. At bisulphate activity of 0.2, the cross over between the two species was estimated to occur at pH ~1. The Gibbs 0 0 energy of formation of aqueous FeHSO4SO4 and aqueous Fe2(SO4)3 at 220°C were estimated to be -1375.7 and -2088.6 kJ mol-1 respectively. • A new thermodynamic data for basic ferric sulphate at 220°C was determined in this study. From basic ferric sulphate solubility experiments, the Gibbs free energy of formation for basic ferric sulphate solid at 220°C was estimated to be -919.4  0.1 kJ mol-1. This value is higher than the value calculated from thermodynamic properties reported in Majzlan et al. (2017) by 9%. However, the Gibbs free energy determined in this study gave a closer estimate of equilibrium pH between hematite and basic ferric sulphate compared to the stability diagram reported in Fleming (2009). • Similarly, the Gibbs free energy of formation for potassium jarosite solid at 220°C was estimated to be -3009.4  0.3 kJ mol-1. This value falls within the reported values in

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literatures (Stoffregen, 1993, Stofferegen, 2000 via HSC Chemistry v.7.1, Majzlan et al, 2010 via HSC Chemistry 9). By assessing the equilibrium pH between hematite and potassium jarosite, it was found that the Gibbs free energy determined in this study gave a closer estimate to the stability diagram reported in Babcan (1971). • An experimentally calibrated thermodynamic data set is proposed in this study for the

development of Eh-pH diagram for Fe-S-H2O and Fe-S-K-H2O system at 220°C.

6.2 Industrial implication • The increase of ferrous concentration in Lihir autoclave due to the implementation of partial pressure oxidation should not affect the pyrite oxidation due to the high solid loading on site. • The increase of ferric concentration in Lihir autoclave should increase the proportion of pyrite oxidised by ferric. However, the overall oxidation in Lihir autoclave would still be governed by oxygen mass transfer in the autoclave due to the low concentration of ferric in Lihir autoclave even with the addition of ferric into the autoclave feed. • Potassium jarosite is suspected to be the main lime consumer at Lihir. The kinetic of the potassium jarosite dissolution reaction with lime is suspected to be relatively fast due to the small particle size of potassium jarosite precipitated in the autoclave and also due the possibility of crystal surface defect and lack of crystallinity of the particles. • Basic ferric sulphate was also predicted to precipitate in Lihir autoclave when the

sulphuric acid concentration in the autoclave is above approximately 35 g/L H2SO4. Basic ferric sulphate has been known to cause high lime consumption in gold pressure oxidation operations. • The other phase that could consume lime is alunite. Alunite was found to form during the Lihir ore POX experiments. Alunite formation is favoured in the presence of aluminium.

6.3 Limitations and recommendations • The solution pH data used in this study was estimated based on a room temperature pH measurement. Some errors are introduced in the adjustment calculation due to the assumptions made to simplify the calculation. Therefore, In-situ pH measurement is recommended in the future to improve model accuracy, more specifically the data positioning in the diagram

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• The heat of formation of BFS and potassium jarosite are of great interest from a practical

heat balance point of view. Therefore, it is recommended for these ΔHf data to be generated in future work. • High temperature ferric speciation should be investigated using direct measurement. • The predictions made on the Lihir system was based on an estimated potassium concentration as potassium concentration is not currently monitored on site. When the potassium data is available in the future, a better prediction could be made. • The high temperature sampling setup in this study allows both solid and solution sampling at high temperature. However, due to the presence of sampling bomb, a slight change in the solution composition could occur.

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Appendix A

Appendix A.1 XRD pattern of feed pyrite

Figure 99 XRD of Pyrite used for experiment

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Appendix A.2 XRD patterns from hematite solubility experiments

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Appendix A.3 XRD patterns from basic ferric sulphate solubility experiments

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Appendix A.4 XRD patterns from potassium jarosite solubility experiments

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Appendix B Appendix B.1 Liquid-liquid Junction Potential sample calculation

The liquid-liquid junction potential in this work was estimated using Henderson formula as described in Newman & Thomas-Alyea (2004). To calculate this, the final filtrate test solution and the reference solution first need to be defined. For this calculation, one of the potassium jarosite solubility experiments was used as the test solution. The pH of the final filtrate was 0.81 and this translates to 0.155 mol/L H+. The solution assays from ICP-OES analysis are as follows.

Solution Assay g/L mol/L Fe3+ 0.34 0.0061 Total S 8.6 0.27 K+ 2.1 0.054 Li+ 0.36 0.052

3+ All Fe was assumed to be in the form of FeHSO4SO4 and all sulphur was assumed to be - in the form HSO4 at high temperature. A portion of that is associated with the complex ferric 0 - - species (FeHSO4SO4 ) while the rest is assumed to be in form of free HSO4 . The free HSO4 concentration in solution at high temperature was estimated as follows.

2 푚표푙 표푓 푆 푚표푙 퐹푒3+ 푆푢푙푝ℎ푢푟 푎푠푠표푐𝑖푎푡푒푑 푤𝑖푡ℎ 퐹푒퐻푆푂 푆푂 = × 0.0061 4 4 1 푚표푙 표푓 퐹푒3+ 퐿

푚표푙 푆 = 0.012 퐿

푚표푙 푆 푚표푙 푆 푆푢푙푝ℎ푢푟 푎푠 푓푟푒푒 퐻푆푂 − = 0.27 − 0.012 4 퐿 퐿

푚표푙 푆 1 푚표푙 퐻푆푂 − = 0.26 × 4 퐿 1 푚표푙 푆

푚표푙 퐻푆푂 − = 0.26 4 퐿

The ORP probe refence solution is 0.5M KCl solution. Therefore, there are 0.5 mol/L K+ and 0.5 mol/L Cl- in the reference solution.

Subsequently, the concentration difference for each species (𝑖) was calculated.

퐼 퐼퐼 푐푖 − 푐푖 = 푐푖,푡푒푠푡 푠표푙푢푡푖표푛 − 푐푖,푟푒푓푒푟푒푛푐푒 푠표푙푢푡푖표푛 193 | P a g e

For this calculation, the concentration difference for each species are summarised below.

Aqueous species Test Solution, Reference Solution, Concentration 퐼 퐼퐼 푐푖 (mol/L) 푐푖 (mol/L) difference (mol/L) H+ 0.155 0 0.155 0 FeHSO4SO4 0.0061 0 0.0061 - HSO4 0.2 0 0.2 K+ 0.054 0.5 -0.45 Li+ 0.052 0 0.052 Cl- 0 0.5 -0.5

The liquid-liquid junction potential (ΦΙ − ΦΙΙ) was estimated using the following equation: 푅푇 퐵Ι⁄퐵ΙΙ ΦΙ − ΦΙΙ = − × 퐴 × ln 퐹 퐵Ι−퐵ΙΙ

The A and B was calculated using the following formulas. The ionic mobility 푢푖 in 퐴 and 퐵 can be replaced by ionic diffusion coefficient 퐷푖.

Ι ΙΙ 퐴 = ∑ 푧푖퐷푖(푐푖 − 푐푖 )

Ι 2 Ι 퐵 = ∑ 푧푖 퐷푖푐푖

ΙΙ 2 ΙΙ 퐵 = ∑ 푧푖 퐷푖푐푖

Ι Using the ionic charge 푧푖 and ionic diffusion coefficient 퐷푖 tabulated in Table 10, the A, 퐵 ΙI 0 and 퐵 are summarised below. For FeHSO4SO4 species, the ionic charge is zero and therefore the A and B value would also be zero.

Aqueous species 퐴 × 109, 퐵ΙI × 109, 퐵ΙI × 109, 푚표푙.푚2 푚표푙.푚2 푚표푙.푚2 ( ) ( ) ( ) 퐿.푠 퐿.푠 퐿.푠 H+ 1.44 1.44 0 OH- -3.4 × 10−13 3.4 × 10−13 0 0 FeHSO4SO4 0 0 0 - HSO4 -0.35 0.35 0 K+ -0.87 0.11 0.98 Li+ 0.053 0.053 0 Cl- 1.02 0 1.01 Sum Total: 1.29 1.95 1.99

Sample calculation for K+ species are shown below. Ι ΙΙ 퐴퐾+ = 푧푖퐷푖(푐푖 − 푐푖 ) 푚2 푚표푙 = 1 × 1.957 × 10−9 × −0.45 푠 퐿 194 | P a g e

푚표푙. 푚2 = −0.87 × 10−9 퐿. 푠

퐼 2 Ι 퐵 퐾+ = 푧푖 퐷푖푐푖 푚2 푚표푙 = 12 × 1.957 × 10−9 × 0.054 푠 퐿 푚표푙. 푚2 = 0.11 × 10−9 퐿. 푠

퐼퐼 2 ΙI 퐵 퐾+ = 푧푖 퐷푖푐푖 푚2 푚표푙 = 12 × 1.957 × 10−9 × 0.5 푠 퐿 푚표푙. 푚2 = 0.98 × 10−9 퐿. 푠 The liquid-liquid junction potential (ΦΙ − ΦΙΙ) equation was split into two terms as follows. 푅푇 퐵Ι⁄퐵ΙΙ Ι ΙΙ − × 퐴 × ln Φ − Φ = 퐹 퐵Ι−퐵ΙΙ

st 푅푇 1 term = − × 퐴 퐹 퐽 8.314 × (220 + 273.15) 퐾 푚표푙. 푚2 = − 푚표푙. 퐾 × 1.29 × 10−9 ( ) 퐶 96500 퐿. 푠 푚표푙 = -0.042 V 퐵Ι⁄퐵ΙΙ 2nd term = 퐴 × ln 퐵Ι−퐵ΙΙ −9 −9 푚표푙.푚2 1.95 ×10 ⁄1.99 ×10 퐿.푠 = 1.29 × 10−9 ( ) × ln ( ) 퐿.푠 1.95 ×10−9− 1.99 ×10−9 푚표푙.푚2 = 0.66 ΦΙ − ΦΙΙ = -0.042 V × 0.66 = -0.0279 V = -27.9 mV (vs. NHE)

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Appendix B.2 Lime consumption calculation The lime consumption at Lihir is expressed in kg of CaO per tonne of ore fed into the autoclave. Therefore, the solid (ore) mass in the feed was calculated using the autoclave feed mass times by the solid wt% which is approximately 55% in average. The autoclave feed (slurry) mass was taken as the sum of all four autoclaves feed mass.

푡 − 푠표푙𝑖푑 푂푟푒 푚푎푠푠 (푡) = 퐴푢푡표푐푙푎푣푒 푓푒푒푑 (푠푙푢푟푟푦) 푚푎푠푠 (푡) × 0.55 퐴퐶 퐹푒푒푑 푡 − 푠푙푢푟푟푦

In this sample calculation, the data is based on plant data on 27th December 2013.

푡 − 푠표푙𝑖푑 푂푟푒 푚푎푠푠 (푡) = 13,763 (푡) × 0.55 퐴퐶 퐹푒푒푑 푡 − 푠푙푢푟푟푦

= 7,569 tonne of ore

The total lime consumptions are divided into: • Lime consumed for acid neutralisation; • Lime consumed for metal cation precipitation; • Lime consumed by reactive solid.

The lime consumption for acid neutralisation takes account the amount of lime used to remove the free acid and raise the solution pH to 10. The calculation was based on the

H2SO4 concentration (g/L) in the neutralisation feed recorded onsite. There are two H2SO4 concentrations reported coming from the two CCD circuit. The maximum of the two values was taken. This is because an averaged value would not be suitable in this case as some th data from either CCD were occasionally missing. On 27 December, the H2SO4 concentration from CCD was 1.71 g/L.

The H2SO4 concentration was converted to mol/L using its molecular weight of 98.08 g/mol.

The initial H2SO4 concentration translates to 0.018 mol/L. At pH 10 which is the target final pH, the amount of free H+ in the solution is 1 × 10−10 mol/L. Therefore, the amount of H+ that needs to be removed are as follow.

푚표푙 푚표푙 푚표푙 퐻+푟푒푚표푣푒푑 ( ) = 𝑖푛𝑖푡𝑖푎푙 [퐻+] ( ) − 푓𝑖푛푎푙 [퐻+] ( ) 퐿 퐿 퐿

푚표푙 = ( 2 × 0.018) − 1 × 10−10 퐿

푚표푙 = 0.035 퐿

196 | P a g e

For each mol of CaO, two moles of hydroxide ions (OH-) are released from slaked lime

(Ca(OH)2) according to the lime slaking reaction below. This means 0.018 mol/L of CaO is required to neutralise the free acid coming into the neutralisation stage.

CaO (s) + H2O (l) → Ca(OH)2 (a)

Given that CaO molecular weight is 56.08 g/mol, the amount of CaO required is 0.98 g/L. To convert this value to the desired unit of kg-CaO/ t-ore. The volume of solution coming into the neutralisation stage required to be calculated.

The amount of solid in the CCD overflow was assumed to be 90% of the solid in the autoclave feed. The lost of solid was attributed to the oxidation in the autoclave and solids reporting to CCD overflow. By assuming that solid loading in the CCD underflow is 45 wt%, the solution proportion in the CCD underflow can be calculated as follow.

0.55 푡 − 푠표푙푢푡𝑖표푛 푆표푙푢푡𝑖표푛 푚푎푠푠 푖푛 퐶퐶퐷 푈/퐹 = 푂푟푒 푚푎푠푠 (푡) × 0.85 ÷ 퐴퐶 푓푒푒푑 0.45 푡 − 푠표푙𝑖푑

= 7,569 × 0.9 × 1 푡표푛푛푒 표푓 푠표푙푢푡𝑖표푛

= 8,326 푡표푛푛푒 표푓 푠표푙푢푡𝑖표푛

By assuming that the solution density is 1,030 g/L, the volume of solution in the CCD underflow was estimated to be 8.1 ML of solution. The lime consumption for free acid neutralisation was the calculated as follow.

푔퐶푎푂 푉푠표푙,퐶퐶퐷 푈/퐹 (퐿) 퐿𝑖푚푒 푐표푛푠푢푚푝푡𝑖표푛 푓표푟 퐴푐𝑖푑 푛푒푢푡푟푎푙𝑖푠푎푡𝑖표푛 = 0.98 × 퐿 푚푎푠푠표푟푒,퐴퐶 퐹푒푒푑 (푡)

푔 8.1 × 106 (퐿) = 0.98 퐶푎푂 × 퐿 8,326 (푡)

푘푔퐶푎푂 = 1.1 푡표푟푒 퐴퐶 푓푒푒푑

Lime consumption for metal cation precipitation is mainly attributed to the precipitation of ferric, ferrous, magnesium and aluminium. However, only iron concentration (both ferric and ferrous) is currently monitored. Therefore the total concentration of other cations were assumed to be half of the ferric concentration and majority of them were in the form of magnesium.

The concentration of these species in the neutralisation feed, however, was not available in the plant data. Therefore, it had to be estimated from their concentration in the autoclave discharged and applying a dilution factor across the CCD circuit. The dilution factor was 197 | P a g e estimated based on the H2SO4 concentration in the neutralisation feed and autoclave discharge.

For example, the maximum concentration of Fe3+ and Fe2+ were 1.26 g/L and 0.28 g/L respectively. The concentration of Mg2+ was then assumed to be 0.63 g/L. Based on the

H2SO4 concentration, the dilution factor was estimated to be 19.

For Mg2+ and Fe2+, each mol of this species consumes 2 moles of hydroxide ions which means 1 mol of CaO is required. For each mol of Fe3+, however, 3 moles of hydroxide ions is consumed which means 1.5 moles of CaO is required. The lime consumption was then calculated using the following equation.

1.5 푚표푙 퐶푎푂 (∑[푀3+] × 푔 푛푒푢푡 1 푚표푙 푀3+ 퐿𝑖푚푒 푐표푛푠푢푚푝푡𝑖표푛 = 푐푎푡푖표푛 푝푟푒푐푖푝. 퐿 1 푚표푙 퐶푎푂 + [푀2+] × ) × 푀푊 푛푒푢푡 1 푚표푙 푀2+ 퐶푎푂

For this sample calculations, lime consumption was estimated to be 0.19 g/L CaO. By applying the same method as the previous one to convert the unit, the lime consumption for metal cation precipitation was estimated to 0.21 kg/t.

The lime consumed by reactive solid was calculated by the difference between the total lime consumption and the lime consumed for free acid neutralisation and cation precipitation

퐿𝑖푚푒 푠표푙푖푑. = 퐿𝑖푚푒 푇표푡푎푙 − 퐿𝑖푚푒 푎푐푖푑 푛푒푢푡.. − 퐿𝑖푚푒 푐푎푡푖표푛 푝푟푒푐푖푝.

The total lime consumption was reported to be 16.0 kg/t. Given that the lime consumption for acid neutralisation and cation precipitation were calculated to be 1.1 and 0.2 kg/t respectively, the amount of lime consumed by reactive solid was estimated to be 14.7 kg/t.

198 | P a g e

Appendix C

Appendix C.1 pH Estimation at high temperature The solution pH at high temperature in this study was estimated based on the solution pH recorded at approximately 30°C. The following assumptions were made to simplify the calculations:

2- - • Only equilibrium reaction between SO4 and HSO4 is considered

-2 + - SO4 (a) + H (a) ↔ HSO4 (a) • Activity coefficient for all species were assumed to be unity • Solution molality (푚) ≈ solution molarity (M) • Sulphur complexation with other elements is assumed to be negligible and therefore total sulphur in the solution is attributed to only bisulphate and sulphate ions.

The solution assay for total sulphur was done on filtrate at 220°C. The total sulphur concentration should not change with temperatures and therefore 푇표푡푎푙 푆@220°퐶 ≈

푇표푡푎푙 푆@30°퐶. Based on the assumption, the sulphur balance then can be expressed as follows.

푚 = 푚 − + 푚 2− 푡표푡푎푙 푆 퐻푆푂4 푆푂4

푚 2− = 푚 − 푚 − 푆푂4 푡표푡푎푙 푆 퐻푆푂4

− At 30°C, the 푚퐻푆푂4 can be estimated using the equilibrium constant as shown below.

푚퐻푆푂 − 퐾푒푞 = 4 푚 2−푚 + 푆푂4 퐻 − 푚퐻푆푂4 퐾푒푞 = − + (푚푡표푡푎푙 푆 − 푚퐻푆푂4 ) × 푚퐻

퐾푒푞 × 푚푡표푡푎푙 푆 × 푚퐻+ − 푚퐻푆푂4 = 1 + (퐾푒푞 × 푚퐻+)

For this sample calculation, the data from potassium jarosite solubility experiment #9 was

- 2- used. The 퐾푒푞 for the HSO4 and SO4 equilibrium reaction at 30°C is 133.7 (HSC Chemistry v.9). The 푚퐻+ was calculated based on the pH measurement at 30°C. For pH reading of

0.78, the 푚퐻+was calculated to be 0.166 m. The total sulphur concentration from ICP-OES was estimated to be 0.27 m.

퐾푒푞 × 푚푡표푡푎푙 푆 × 푚퐻+ − 푚퐻푆푂4 = 1 + (퐾푒푞 × 푚퐻+) 133.7 × 0.27 × 0.17 푚 − = 퐻푆푂4 1 + (133.7 × 0.17)

199 | P a g e

− 푚퐻푆푂4 = 0.26 m

By considering the hydrogen species balance in the system, the total H+ in the system can be calculated as follows.

+ − + 푚푡표푡푎푙 퐻 = 푚퐻푆푂4 + 푚푓푟푒푒 퐻 = 0.26 푚 + 0.17 푚 = 0.43 푚

− + By rearranging the equation, the 푚퐻푆푂4 can be related to the 푚푓푟푒푒 퐻 which is equivalent to solution pH.

− + + 푚퐻푆푂4 = 푚푡표푡푎푙 퐻 − 푚푓푟푒푒 퐻

The equilibrium constant expression at 220°C then can be estimate as follows by substituting the hydrogen and sulphur species balances.

푚퐻푆푂 − 퐾푒푞 = 4 푚 2−푚 + 푆푂4 퐻 − 푚퐻푆푂4 퐾푒푞 = − + (푚푡표푡푎푙 푆 − 푚퐻푆푂4 ) × 푚퐻

(푚푡표푡푎푙 퐻+ − 푚푓푟푒푒 퐻+,220°퐶 ) 퐾푒푞 = (푚푇표푡푎푙 푆 − (푚푡표푡푎푙 퐻+ − 푚푓푟푒푒 퐻+,220°퐶)) 푚푓푟푒푒 퐻+,220°퐶

- 2- The equilibrium constant for the HSO4 and SO4 reaction at 220°C is 83,320 (HSC

Chemistry v.9). Based on the ICP-OES analysis, the 푚푇표푡푎푙 푆 is 0.27m. The 푚푡표푡푎푙 퐻+ at 220°C should be similar to that at 30°C which was 0.43 m.

The equation was solved using Solver in Excel and the 푚푓푟푒푒 퐻+,220°퐶 was estimated to be

0.15m which is equivalent to pH@220°C of 0.81.

200 | P a g e

Appendix C.2 95% confidence interval sample calculation for solubility experiment

The potassium jarosite solubility experiments data was used to show sample calculation on 95% confidence interval calculation. The Gibbs free energy of formation for potassium jarosite from all nine experiments was calculated from the log (Keq) values. Example calculations are shown below.

∆Grxn° = −RTln(Keq) 퐽 = − 8.314 × 493.14 퐾 × ln(10−0.54) 푚표푙. 퐾 퐽 = 5,128.5 푚표푙 푘퐽 = 5.1 푚표푙

The potassium jarosite reaction precipitation reaction considered is as follows. 0 + - + 3FeHSO4SO4 (a) + K (a) + 6H2O(a) = KFe3(SO4)2(OH)6 (s) + 4 HSO4 + 5H (a)

With this reaction, the Gibbs energy of formation for potassium jarosite can be calculated. ° ° ° ° ° − + + ∆퐺 푟푥푛 − 4∆퐺 푓,퐻푆푂4 − 5∆퐺 푓,퐻 + 3∆퐺 푓,퐹푒퐻푆푂4푆푂4 + ∆퐺 푓,퐾

∆퐺°푓,퐾퐹푒3 (푆푂4)2(푂퐻)6, = ° + 6∆퐺 푓,퐻2푂 푘퐽 푘퐽 푘퐽 푘퐽 5.1 − (4 × −662.8 ) − (0 ) + (3 × −1,374.5 ) 푚표푙 푚표푙 푚표푙 푚표푙 = 푘퐽 푘퐽 + (−301.0 ) + (6 × −207.0 ) 푚표푙 푚표푙 푘퐽 = −3010.5 푚표푙

These calculations were repeated for all nine experiments and the results are tabulated below.

∆퐺° ∆퐺° log(퐾 ) 푟푥푛,220°퐶 푓,220°퐶 푒푞 (kJ mol-1) (kJ mol-1) 1 -0.54 5.1 -3010.5 2 -0.57 5.3 -3010.3 3 -0.78 7.3 -3008.3 4 -0.77 7.3 -3008.4 5 -0.67 6.3 -3009.4 6 -0.78 7.4 -3008.3 7 -0.54 5.1 -3010.5

201 | P a g e

∆퐺° ∆퐺° log(퐾 ) 푟푥푛,220°퐶 푓,220°퐶 푒푞 (kJ mol-1) (kJ mol-1) 8 -0.62 5.8 -3009.8 9 -0.73 6.9 -3008.8

The 95% confidence interval for the Gibbs free energy of formation was determined using t- test according to the equation below. s E95% = ±tα/2,DOF × √n

The critical t-value (푡훼/2,퐷푂퐹) was calculated using Excel function T.INV(/2, DOF). The degree of freedom (DOF) was determined by the number of samples where DOF = n-1. Since there were 9 samples, the DOF was 8. To calculate the critical t-value for 95% confidence interval (i.e. α = 0.05), T.INV(0.025, 8) was entered in excel and the t-value was calculated to be 2.31. The standard deviation (s) was calculated using Excel formula STDEV.S and it was

-1 calculated to be 1.0 kJ mol . The confidence interval (E95% ) was then estimated to be 0.7 kJ mol-1. The mean Gibbs free energy of formation from all nine experiments was calculated to be -3009.4 kJ mol-1. Therefore, the Gibbs free energy of formation for potassium jarosite was estimated to be -3009.4 ± 0.7 kJ mol-1

202 | P a g e

Appendix D QXRD Codes for TOPAS r_exp 12.5280013 r_exp_dash 9.07950819 r_wp 20.4109782 r_wp_dash 14.7925946 r_p 15.3767917 r_p_dash 12.2398843 weighted_Durbin_Watson 1.51924095 gof 1.62922862 continue_after_convergence iters 1000 do_errors

'change this to the name of your XRD raw file xdd "C:\Users\hydrogroup\Desktop\Ivana\Pyrite_30_Corr_K2A_Stripped.raw"

'do_errors r_wp 20.4109782 r_exp 12.5280013 r_p 15.3767917 r_wp_dash 14.7925946 r_p_dash 12.2398843 r_exp_dash 9.07950819 weighted_Durbin_Watson 1.51924095 gof 1.62922862 lam ymin_on_ymax 0.0001

Lam_recs

{

@ 0.00036`_0.07334 1.4764 3.6854

@ 0.00382`_0.78420 1.540596 0.4370

@ 0.00019`_0.03878 1.3922 0.6000

} bkg @ -27.8330982`_13.9351253 56.74662`_16.9988087 -33.0895274`_10.3582304 17.8187099`_6.28107736 -9.0932512`_3.85328396 4.50374657`_2.37109751 - 1.24308628`_1.49288727 0.295092692`_0.867407068 -0.638721362`_0.53490117 ' add/remove coefficients as required

One_on_X(@, 1339.25435`_293.84581)

Radius(280)

Full_Axial_Model(12, 15, 17.3, 2.5, 2.5) 'synchontron do finger model

Out_Yobs_Ycalc_and_Difference(pyrite30.txt)

' lpsd_th2_angular_range_degrees 2.944

203 | P a g e

' lpsd_equitorial_divergence_degrees 0.12

' lpsd_equitorial_sample_length_mm 25

Zero_Error(@, 0.00976`_0.00205)

LP_Factor(0)

Absorption(@, 500.00000`_1299.70923)

str

' Generated by PDF-4+ 2016 software 4.16.0.4

space_group R-3c

phase_name "Corundum, syn"

' Formula: Al2 O3

a a_010821399 4.760246`_0.000365

b =a_010821399;

c c_010821399 12.993256`_0.001528

al 90.0

be 90.0

ga 120.0

CS_L(@, 123.04606`_10.32906)

spiked_phase_measured_weight_percent 10

' Lattice: Rhombohedral

' Atomic parameters are cross-referenced from PDF entry 04-004-2852

site Al_1 x =0; y =0; z !z_Al_1_010821399 0.35216 occ Al !occ_Al_1_010821399 1.0 beq !beq_Al_1_010821399 0.22424

site O_1 x !x_O_1_010821399 0.30624 y =0; z =1/4; occ O !occ_O_1_010821399 1.0 beq !beq_O_1_010821399 0.27082

site Al_1 x =0; y =0; z !z_Al_1_010821399 0.35216 occ Al !occ_Al_1_010821399 1.0 ADPs { Bani11 0.22029 Bani22 0.22029 Bani33 0.23371 Bani12 0.11054 Bani13 0.0 Bani23 0.0 } 204 | P a g e

site O_1 x !x_O_1_010821399 0.30624 y =0; z =1/4; occ O !occ_O_1_010821399 1.0 ADPs { Bani11 0.25819 Bani22 0.26924 Bani33 0.28819 Bani12 0.13502 Bani13 0.03711 Bani23 0.07422 }

' Unit cell volume [A^3]: 254.42

' Density (calculated) [g/cm^3]: 3.993

scale sc_010821399 0.00818760009`_1.682

Primary Reference: Calculated from ICSD using POWD-12++ (1997).

Crystal Structure: Crystal Structure Source: LPF.

Structure: "Neutron diffraction measurements of the residual stresses in Al2 O3 - Zr O2 (Ce O2) ceramic composites". Wang, X.-L., Hubbard, C.R., Alexander, K.B., Becher, P.F. J. Am. Ceram. Soc. 77, 1569 (1994).

' Database entry no.: 010821399

'------

str

' Generated by PDF-4+ 2019 software 4.19.0.2

space_group C12/c1

phase_name "Szomolnokite, syn"

' Formula: Fe H2 O5 S

a a_000451365 7.084820`_0.000501

b b_000451365 7.549348`_0.000599

c c_000451365 7.777922`_0.000568

al 90

be 118.65

ga 90

205 | P a g e

CS_L(@, 207.35140`_15.41812)

' Lattice: Monoclinic

' Atomic parameters are cross-referenced from PDF entry 04-014-9807

site Fe_1 x =0; y =1/2; z =0; occ Fe !occ_Fe_1_000451365 1.0 beq !beq_Fe_1_000451365 1.30674

site S_1 x =0; y !y_S_1_000451365 0.15307 z =1/4; occ S !occ_S_1_000451365 1.0 beq !beq_S_1_000451365 1.20804

site O_1 x !x_O_1_000451365 0.1697 y !y_O_1_000451365 0.0429 z !z_O_1_000451365 0.3985 occ O !occ_O_1_000451365 1.0 beq !beq_O_1_000451365 1.77732

site O_2 x !x_O_2_000451365 0.0956 y !y_O_2_000451365 0.2683 z !z_O_2_000451365 0.156 occ O !occ_O_2_000451365 1.0 beq !beq_O_2_000451365 1.65572

site O_3 x =0; y !y_O_3_000451365 0.6444 z =1/4; occ O !occ_O_3_000451365 1.0 beq !beq_O_3_000451365 1.63914

site H_1 x !x_H_1_000451365 0.108 y !y_H_1_000451365 0.709 z !z_H_1_000451365 0.315 occ H !occ_H_1_000451365 1.0 beq !beq_H_1_000451365 6.9561

site Fe_1 x =0; y =1/2; z =0; occ Fe !occ_Fe_1_000451365 1.0 ADPs { Bani11 0.95538 Bani22 1.29489 Bani33 1.61862 Bani12 0.0079 Bani13 0.57638 Bani23 0.08685 }

site S_1 x =0; y !y_S_1_000451365 0.15307 z =1/4; occ S !occ_S_1_000451365 1.0 ADPs { Bani11 0.79746 Bani22 1.14487 Bani33 1.61862 Bani12 0.0 Bani13 0.52901 Bani23 0.0 }

site O_1 x !x_O_1_000451365 0.1697 y !y_O_1_000451365 0.0429 z !z_O_1_000451365 0.3985 occ O !occ_O_1_000451365 1.0 ADPs { Bani11 1.08171 Bani22 1.8318 Bani33 2.23448 Bani12 0.33162 Bani13 0.64745 Bani23 0.62376 }

site O_2 x !x_O_2_000451365 0.0956 y !y_O_2_000451365 0.2683 z !z_O_2_000451365 0.156 occ O !occ_O_2_000451365 1.0 ADPs { Bani11 1.42122 Bani22 1.57914 Bani33 2.195 Bani12 0.22897 Bani13 1.05013 Bani23 0.49743 }

site O_3 x =0; y !y_O_3_000451365 0.6444 z =1/4; occ O !occ_O_3_000451365 1.0 ADPs { Bani11 1.26331 Bani22 1.62651 Bani33 2.0055 Bani12 0.0 Bani13 0.76588 Bani23 0.0 }

site H_1 x !x_H_1_000451365 0.108 y !y_H_1_000451365 0.709 z !z_H_1_000451365 0.315 occ H !occ_H_1_000451365 1.0

206 | P a g e

' Unit cell volume [A^3]: 364.86

' Density (calculated) [g/cm^3]: 3.093

scale sc_000451365 0.00781102198`_1.605

prm wt_FeSO4monohydrate= Get(corrected_weight_percent);

Primary Reference: Rosenstingl, J., Hoffmann, C., Irran, E., Lengauer, C., Tillmanns, E., Univ. Vienna, Austria. ICDD Grant-in-Aid 1992.

Crystal Structure: Crystal Structure Source: LPF.

Structure: Wildner, M., Giester, G. Neues Jahrb. Mineral., Monatsh. 1991, 1991, 296.

' Database entry no.: 000451365

str

' Generated by PDF-4+ 2019 software 4.19.0.2

space_group Pa-3

phase_name "Pyrite, syn"

' Formula: Fe S2

a a_040044841 5.391126`_0.017372

b =a_040044841;

c =a_040044841;

al 90

be 90

ga 90

CS_L(@, 12.63988`_25.70037)

Strain_L(@, 5.00000`_4.17603)

' Lattice: Cubic

207 | P a g e

site Fe_1 x =0; y =0; z =0; occ Fe !occ_Fe_1_040044841 1.0

site S_1 x !x_S_1_040044841 0.384 y =x_S_1_040044841; z =x_S_1_040044841; occ S !occ_S_1_040044841 1.0

' Unit cell volume [A^3]: 157.81

' Density (calculated) [g/cm^3]: 5.049

scale sc_040044841 0.0109004353`_2.239

prm wt_pyrite= Get(corrected_weight_percent);

Primary Reference: Calculated from LPF using POWD-12++.

Structure: Elliott N. "Interatomic Distances in FeS2, CoS2 and NiS2". J. Chem. Phys. 1960, 33, 903.

' Database entry no.: 040044841

str

' Generated by PDF-4+ 2019 software 4.19.0.2

space_group Pnma

phase_name "Iron Sulfate Hydroxide"

' Formula: Fe H O5 S

a a_040126256 7.424816`_0.099227

b b_040126256 7.108806`_0.073393

c c_040126256 6.564126`_0.110465

al 90

be 90

ga 90

CS_L(@, 5.57843`_3.99807)

208 | P a g e

Strain_L(@, 0.00010`_3.77181)

' Lattice: Orthorhombic

site Fe_1 x !x_Fe_1_040126256 0.125 y =1/4; z !z_Fe_1_040126256 0.25 occ Fe !occ_Fe_1_040126256 1.0

site S_1 x !x_S_1_040126256 0.875 y =1/4; z !z_S_1_040126256 0.866 occ S !occ_S_1_040126256 1.0

site O_1 x !x_O_1_040126256 0.375 y !y_O_1_040126256 0.067 z !z_O_1_040126256 0.75 occ O !occ_O_1_040126256 1.0

site O_2 x !x_O_2_040126256 0.209 y =1/4; z !z_O_2_040126256 0.517 occ O !occ_O_2_040126256 1.0

site O_3 x !x_O_3_040126256 0.959 y =3/4; z !z_O_3_040126256 0.017 occ O !occ_O_3_040126256 1.0

site O_4 x !x_O_4_040126256 0.375 y =1/4; z !z_O_4_040126256 0.155 occ O !occ_O_4_040126256 1.0

' Unit cell volume [A^3]: 336.13

' Density (calculated) [g/cm^3]: 3.338

scale sc_040126256 0.00202433228`_0.416

prm wt_BFS= Get(corrected_weight_percent);

Primary Reference: Calculated from LPF using POWD-12++.

Structure: Johansson G. "On the Crystal Structures of FeOHSO4 and InOHSO4". Acta Chem. Scand. 1962, 16, 1234-1244.

' Database entry no.: 040126256

str

' Generated by PDF-4+ 2019 software 4.19.0.2

space_group R-3cH

phase_name "Hematite" 209 | P a g e

' Formula: Fe2 O3

a a_000011053 5.029929`_0.000355

b =a_000011053;

c c_000011053 13.803438`_0.001346

al 90

be 90

ga 120

CS_L(@, 53.02649`_3.29681)

Strain_L(@, 0.15006`_0.02705)

' Lattice: Rhombohedral

' Atomic parameters are cross-referenced from PDF entry 04-011-9586

site Fe_1 x =0; y =0; z z_Fe_1_000011053 0.14631`_0.00023 occ Fe !occ_Fe_1_000011053 1.0 beq !beq_Fe_1_000011053 0.1 '0.14476

site O_1 x x_O_1_000011053 0.34610`_0.00217 y =1/3; z =7/12; occ O !occ_O_1_000011053 1.0 beq !beq_O_1_000011053 0.15 '0.36097

' Unit cell volume [A^3]: 300.60

' Density (calculated) [g/cm^3]: 5.293

scale sc_000011053 0.0155104319`_3.186

prm wt_Fe2O3= Get(corrected_weight_percent);

Primary Reference: Hanawalt, J., et al. Anal. Chem. 1938, 10, 475.

Crystal Structure: Crystal Structure Source: LPF.

Unit Cell: Dana's System of Mineralogy, 7th Ed.

' Database entry no.: 000011053

str

phase_name "BFS"

210 | P a g e

a 7.331

b 6.419

c 7.143

al 90

be 90

ga 90

volume 336.133

CS_L(@, 806.56403`_18465.83410)

Strain_L(@, 0.00010`_0.87911)

space_group "Pnma"

site O_1 x= 0.125; y =0.067; z= 0.25; occ O 1

site S_1 x =0.125; y =1/4; z= 0.134; occ S 1

site O_4 x= 0.125; y =1/4; z= 0.655; occ O 1

site O_2 x= 0.291; y =1/4; z= 0.017; occ O 1

site Fe_1 x= 0.375; y =1/4; z= 0.75; occ Fe 1

site O_3 x= 0.459; y =1/4; z= 0.483; occ O 1

scale sc_1 3.26868081e-005`_0.006678

prm wt_BFS2= Get(corrected_weight_percent); prm wtp_Fe2O3 = wt_Fe2O3;: 35.19161`_10208.56478 prm wtp_BFS = wt_BFS;: 3.68824`_1070.14298 prm wtp_BFS2 = wt_BFS2;: 0.05778`_16.71451 prm wtp_pyrite= wt_pyrite;: 6.41774`_1861.66112 prm wtp_FeSO4monohydrate = wt_FeSO4monohydrate;: 15.17575`_4402.19697

'prm wtp_butlerite = wt_butlerite;: 0.49539_11.41132 prm wtp_amorphous = Get(weight_percent_amorphous); : 29.46888`_14764.75794

211 | P a g e