The leaching of metal sulfides in ammoniacal carbonate ......It was observed that the highest...

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The leaching of metal sulfides in

ammoniacal carbonate solutions

Anna Nowosielska

This thesis is presented for the Degree of Honours in Mineral Science,

Murdoch University

School of Engineering and Information Technology

June 2017

Supervisor : Dr. Aleks Nikoloski

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Declaration

I declare that this thesis is my own account of my research and

contains, as its main content, work which has not previously been

submitted for a degree or examination at any tertiary educational

institution.

Anna Nowosielska

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Abstract

The leaching of metal sulfides in ammoniacal solutions is of industrial importance for

the production of copper, nickel and cobalt. In some processes, such as the Sherritt

Gordon process, they are introduced as feed; while in others, such as the Caron process,

they form in-situ, as intermediate products. The focus of this study is to understand the

leaching behaviour of the in situ formed copper, nickel and cobalt sulfides, but the

findings are expected to be relevant to the other related systems.

The Caron process involves an oxidative dissolution of pre-reduced iron-based nickel

and cobalt alloy grains in an aqueous solution containing ammonia and ammonium

carbonate. Under the process conditions, these iron-based alloy grains undergo

passivation because of the formation of an iron oxide surface layer, as nickel and cobalt

become trapped within the passive iron matrix, resulting in reduced extractions.

Thiosulfate ions have been reported to facilitate the leaching process. The presence of

metal ammines and thiosulfate ions in the same system, however, has been observed to

result in a formation of a metal sulfide layer on the surface of the dissolving iron-based

alloy grains. This, on the other hand, results in a partial loss of nickel and cobalt from

the solution, while also further increasing the tendency for iron passivation. It has been

postulated, that after iron passivates, nickel and cobalt from the sulfide layer redissolve,

however, there is little evidence of this, or the effect of process parameters on the rate

and extent of such re-dissolution. The present study is aiming to shed some light on these

processes.

The dissolution of copper, nickel and cobalt sulfides was examined using three forms of

these compounds - metal sulfides formed electrochemically (EFMS) under simulated

Caron process leach conditions, synthetic monometallic metal sulfides (SMMS) and real

metal sulfides, naturally found in ore (RMS). In the first instance, electrochemically

formed metal sulfides were studied by techniques which included open circuit potential

and cyclic voltammetry measurements. The dissolution rate and extent of the nickel,

cobalt and copper from the metal sulfide was monitored.

The leaching of the synthetic monometallic metal sulfides was carried out under similar

conditions to those of the Caron process. The effects of key process parameters such as

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concentration of ammonia, concentration of ammonium carbonate and temperature

were investigated. The pH and Eh of the studied systems were monitored and solution

samples were analysed for Ni, Co and Cu. It was observed that the highest extraction of

nickel and copper from their respective metal sulfides was at 3M [NH3]T + 1M [CO2]T at

60 °C, whereas the maximum extraction of cobalt from cobalt sulfide was observed at

5M [NH3]T + 1M [CO2]T at 45 °C.

The kinetic data shows that leaching of nickel follows the ash layer diffusion model, with

the activation energy for the process reported as ‘different’ at the different tested

temperatures (69.6 kJ/mol at 25 °C and 14.7 kJ/mol at 60 °C). An attempt was made to

derive the reaction orders with respect to ammonia (0.18) and ammonium carbonate

(0.09). The available data is possibly not very reliable, however.

The rate of copper leaching shows that the surface reaction model applies best to this

system. The activation energy was 11.5 kJ/mol at 25 °C and 57.6 kJ/mol at 60 °C. The

reaction order with respect to ammonia was 1.9 and for ammonium carbonate it was 4.0.

The leaching kinetic data for cobalt from monometallic cobalt sulfide, showed that the

process follows the surface reaction model similar to copper, with a steady activation

energy of 6.5 kJ/mol. In contrast to nickel, the leaching of cobalt from cobalt sulfide had

a higher order with respect to ammonium carbonate of 1.8 than with respect to ammonia

of 0.9.

The observed leaching behaviour of the electrochemically formed metal sulfides (EFMS)

and the synthetic monometallic metal sulfides (SMMS) was compared to leaching of real

metal sulfides naturally found in ore (RMS) by testing flotation concentrate under the

conditions which gave the highest extractions for the former. Tested were two types of

flotation concentrate, one containing high nickel sulfide and another containing high

copper sulfide. A sample of cobalt sulfide was not available. In these experiments, the

effect of adding 0.25M sodium sulfite (Na2SO3) was also examined. The results have

shown that the addition of sulfite improved both rate and extent of the extraction of

nickel. It significantly hindered the extraction of copper, however. Namely, the rate of

nickel extraction with sodium sulfite present was approximately 1.31 mg/L/min vs. 0.07

mg/L/min without it. The rate of copper extraction on the other hand, decreased from

4.09 mg/L/min without sodium sulfite, to just 0.89 mg/L/min when it was added. On

the basis of this data, and assuming similar behaviour between copper and cobalt

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sulfides, it does not appear that addition of sodium sulfite could reduce the losses of

copper and cobalt from the system.

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Table of Contents

Abstract .................................................................................................................................................. iv

Table of Contents .............................................................................................................................. vii

List of Figures ..................................................................................................................................... ix

List of Tables .....................................................................................................................................xiv

List of Appendices ............................................................................................................................. xv

Abbreviations ....................................................................................................................................xvi

Acknowledgements ......................................................................................................................... xviii

1 INTRODUCTION ............................................................................................................................... 1

2 LITERATURE REVIEW ....................................................................................................................... 2

2.1 ORE DEPOSITS, MINERALOGY AND PROCESSING ................................................................... 2

2.2 AMMONIACAL LEACHING PROCESS ........................................................................................ 5

2.2.1 The Caron Process ........................................................................................................... 7

2.2.2 The Sherritt Gordon Process ........................................................................................... 9

2.3 FACTORS AFFECTING NICKEL AND COBALT EXTRACTION ..................................................... 11

2.4 IRON PASSIVATION LAYER FORMATION ............................................................................... 13

2.5 ELECTROCHEMICAL PRINCIPLES ............................................................................................ 16

2.5.1 Metal Electrolyte Interface - The Electrical Double Layer ............................................ 16

2.5.2 Corrosion Potential ....................................................................................................... 18

2.5.3 Mixed Potential Theory ................................................................................................. 20

2.5.4 Cyclic Voltammetry ....................................................................................................... 22

2.5.5 Passivity of Metals ........................................................................................................ 24

2.5.6 In situ XRD Electrochemical Analysis ............................................................................ 25

3 EXPERIMENTAL MATERIALS AND METHODS ................................................................................ 30

3.1 MATERIALS ............................................................................................................................ 30

3.1.1 Electrochemically formed Metal Sulfides (EFMS) ......................................................... 30

3.1.2 Synthetic Monometallic Metal Sulfides (SMMS) .......................................................... 30

3.1.3 Flotation Concentrate Samples (RMS) .......................................................................... 30

3.1.4 Test Solutions ................................................................................................................ 30

3.2 METHODS USED FOR ELECTROCHEMICAL STUDY ................................................................ 31

3.2.1 Electrochemical Apparatus ........................................................................................... 31

3.2.2 OCP Measurements ...................................................................................................... 32

3.2.3 Rotating Disk Cyclic Voltammetry ................................................................................. 32

3.2.4 SEM/EDX Surface Analysis ............................................................................................ 33

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3.2.5 Dissolution Studies ........................................................................................................ 33

3.3 LEACHING METHODS ............................................................................................................ 33

3.3.1 Leaching of Synthetic Monometallic Metal Sulfides (SMMS) ....................................... 33

3.3.2 Kinetic model for Synthetic Monometallic Metal Sulfides ........................................... 34

3.3.3 Leaching of Flotation Concentrate (RMS) ..................................................................... 35

4 RESULTS AND DISCUSSION ............................................................................................................ 36

4.1 THE INVESTIGATION OF NICKEL SULFIDE .............................................................................. 36

4.1.1 Electrochemical Testwork ............................................................................................. 36

4.1.2 Leaching of a Synthetic Sample (SMMS) ....................................................................... 40


4.2 THE INVESTIGATION OF COBALT SULFIDE ............................................................................ 57



4.3 THE INVESTIGATION OF COPPER SULFIDE ............................................................................ 76




4.4 COMPARISON OF REACTION RATES AND EXTENTS OF LEACHING ....................................... 98

5 CONCLUSIONS ............................................................................................................................... 99

6 REFERENCES ................................................................................................................................ 102

7 APPENDICES ................................................................................................................................ 110

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List of Figures

FIGURE 1. Nickel production and resources (World resources on land) left, and (Nickel production in

2009) right, of laterite and sulfide ores (Oxley et al., 2016). .................................................................. 2

FIGURE 2. Past and future predictions of nickel production from laterite and sulfide ores (Oxley et al.,

2016). ...................................................................................................................................................... 2

FIGURE 3. Generalised laterite processing flowsheets (Norgate and Jahanshahi, 2011)....................... 4

FIGURE 4. A flowsheet showing the Sherritt Gordon process. ............................................................. 10

FIGURE 5. The electric field consisting of two layers of charge – The Electrical Double Layer

(faculty.kfupm.edu.sa) .......................................................................................................................... 18

FIGURE 6. Polarization curves of corrosion reactions, where ia is the anodic reaction current, ic is the

cathodic reaction current, icorr is the corrosion current, Ea is the equilibrium potential of the cathodic

reaction, Ec is the equilibrium potential of the cathodic reaction and Ecorr is the corrosion potential

(Sato, 2012). .......................................................................................................................................... 19

FIGURE 7. A schematic representation of the Mixed Potential Theory for two redox couples

(Nikoloski, 2002). .................................................................................................................................. 22

FIGURE 8. Potential versus time profile applied to the working electrode (Compton and Banks, 2007).

.............................................................................................................................................................. 23

FIGURE 9.Current versus potential Cyclic Voltammogram (Compton and Banks, 2007). .................... 23

FIGURE 10. A schematic representation of current-potential relationship as metal undergoes

passivation (Bockris and Reddy, 1970). ................................................................................................ 24

FIGURE 11. Different cell configurations in reflection and transmission geometries. Thick black arrow

shows the path of X-ray beam; dashed arrows show the scattered beam (De Marco and Veder,

2010). .................................................................................................................................................... 26

FIGURE 12. Synchrotron radiation X-ray real time images of nickel morphology deposited at pH 10

(a) -1.2V, (b) -1.4V and (c) -1.6V (Song et al., 2016). ............................................................................ 27

FIGURE 13.SEM micrographs showing nickel electrodeposition obtained at different potentials

(deposition time of 400s) (a) -1.2V, (b) -1.4V and (c) -1.6V (Song et al., 2016). ................................... 27

FIGURE 14. Schematic of the electrochemical flow cell (a) exploded view of the cell with its main

components, (b) partial section view of the assembled cell, highlighting the path of the flowing

electrolyte solution, and (c) top view of the cell (Clancy et al., 2015). ................................................ 28

FIGURE 15. Schematic diagram of an electrochemical cell used for in-situ X-ray analyses. Upper

diagram shows a side profile of the cell; lower diagram shows a top profile (Veder et al., 2011). ..... 29

FIGURE 16. Electrochemical set-up, showing working electrode (WE), counter electrode (CE),

reference electrode (RE), Luggin capillary(LC) (D’Aloya, 2015). ........................................................... 32

FIGURE 17. Schematics of the leaching set up. .................................................................................... 34

FIGURE 18. OCP of iron in a deoxygenated ammoniacal-carbonate solution containing 0.15M nickel

(II) and 0.022M thiosulfate ions. ........................................................................................................... 36

FIGURE 19. CV of iron following different times of immersion in a deoxygenated Ni (II) thiosulfate

ammoniacal-carbonate solution. .......................................................................................................... 37

FIGURE 20. SEM image of iron RDE following 3 hour immersion in a 3M [NH3]T and 1M [CO2]T

containing 0.15M Ni(II) and 0.022M S2O32-. ......................................................................................... 39

FIGURE 21. Nickel dissolution from iron RDE in 5M [NH3]T + 1M [CO2]T following pre-treatment in 3M

[NH3]T + 1M [CO2]T + 0.15M Ni(II) + 0.022M S2O32- solution. ................................................................ 40

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FIGURE 22. SEM/EDX analysis of the nickel sulfide sample used in the leaching tests (A) SEM image of

the nickel sulfide particles, (B) corresponding spectra for image (A), and (C) EDX image showing the

spots that were quantitatively analysed. .............................................................................................. 41

FIGURE 23. The effect of ammonia concentration on the leaching efficiency of nickel from nickel

sulfide. ................................................................................................................................................... 43

FIGURE 24. The effect of ammonium carbonate concentration on the leaching efficiency of nickel

from nickel sulfide. ................................................................................................................................ 45

FIGURE 25. The effect of temperature on the leaching efficiency of nickel from nickel sulfide. ........ 47

FIGURE 26. The effect of sodium sulfite addition on the leaching efficiency of nickel from nickel

sulfide. ................................................................................................................................................... 48

FIGURE 27. Plot of surface reaction model vs time for the leaching of nickel from nickel sulfide at

different temperatures. ........................................................................................................................ 49

FIGURE 28. Plot of ash layer diffusion model vs time for the leaching of nickel from nickel sulfide at


FIGURE 29. Plot of the mixed model vs time for the leaching of nickel from nickel sulfide at different

temperatures. ....................................................................................................................................... 50

FIGURE 30. Arrhenius plot for nickel leaching from nickel sulfide. ...................................................... 52

FIGURE 31. Plot of –log k vs log [NH3] for reaction order estimation................................................... 52

FIGURE 32. Plot of – log k vs log [CO2] for reaction order estimation. ................................................. 53

FIGURE 33. The effect of sodium sulfite addition on the leaching of nickel from nickel flotation

concentrate. .......................................................................................................................................... 54

FIGURE 34. SEM images of (A) nickel feed, (B) Leach Residue A, and (C) Leach Residue B. ................ 56

FIGURE 35. OCP of iron immersed in a deoxygenated ammoniacal-carbonate solution containing

cobalt (II) and thiosulfate ions. ............................................................................................................. 57

FIGURE 36. CV of iron following different immersion times in a deoxygenated Co (II) thiosulfate

ammoniacal solution. ............................................................................................................................ 58

FIGURE 37. SEM image of the iron RDE following 3 hour immersion in 5M[NH3]T and 1M[CO2]T

containing 0.012M Co(II) and 0.022M S2O32- a) image of a larger area, b) close up on spot marked

with an arrow. ....................................................................................................................................... 60

FIGURE 38. Cobalt dissolution from iron RDE in a 5M [NH3]T + 1M [CO2]T following pre-treatment in

5M [NH3]T + 1M [CO2]T + 0.012M Co(II) + 0.022M S2O32- solution. ....................................................... 61

FIGURE 39. SEM/EDX analysis of the cobalt sulfide sample used in the leaching tests (A) SEM image

of the cobalt sulfide particles, (B) corresponding spectra for image (A), and (C) EDX image showing

the spots that were quantitatively analysed. ....................................................................................... 62

FIGURE 40. The effect of ammonia concentration on the leaching efficiency of cobalt from cobalt

sulfide. ................................................................................................................................................... 64

FIGURE 41. The effect of ammonium carbonate concentration on the leaching efficiency of cobalt

from cobalt sulfide. ............................................................................................................................... 67

FIGURE 42. The effect of temperature on the leaching efficiency of cobalt from cobalt sulfide........ 68

FIGURE 43. The effect of sodium sulfite addition on the leaching efficiency of cobalt from cobalt

sulfide. ................................................................................................................................................... 70

FIGURE 44. Plot of surface reaction model vs time for the leaching of cobalt from cobalt sulfide at


FIGURE 45. Plot of ash layer diffusion model vs time for the leaching of cobalt from cobalt sulfide at


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FIGURE 46. Plot of the mixed model vs time for the leaching of cobalt from cobalt sulfide at


FIGURE 47. Arrhenius plot for cobalt leaching from cobalt sulfide. ..................................................... 74


FIGURE 49. Plot of –log k vs log [CO2] for reaction order estimation. .................................................. 75

FIGURE 50. OCP of iron immersed in a deoxygenated ammoniacal-carbonate solution containing

copper (II) and thiosulfate ions. ............................................................................................................ 76

FIGURE 51. CV of iron following different immersion times in a deoxygenated Cu (II) thiosulfate

ammoniacal solution. ............................................................................................................................ 77

FIGURE 52. SEM images of the different areas on the surface of the iron RDE following 3 hour

immersion in 5M [NH3]T and 1M [CO2]T containing 0.008M Cu (II) and 0.022M S2O32-. ...................... 79

FIGURE 53. Copper dissolution from iron RDE in 5M[NH3]T + 1M[CO2]T following pre-treatment in

5M[NH3]T + 1M[CO2]T + 0.008M Cu(II) + 0.022M S2O32- solution. ........................................................ 80

FIGURE 54. SEM/EDX analysis of the copper sulfide sample used in the leaching tests. (A) SEM image

of the copper sulfide particles, (B) corresponding spectra for image (A), and (C) EDX image showing

spots that were quantitatively analysed. .............................................................................................. 81

FIGURE 55. The effect of ammonia concentration on the leaching efficiency of copper from copper

sulfide. ................................................................................................................................................... 84

FIGURE 56. Stability limits of Cu2+ in ammoniacal solution (Ek et al., 1982). ...................................... 85

FIGURE 57. The effect of ammonium carbonate concentration on the leaching efficiency of copper

from copper sulfide. .............................................................................................................................. 86

FIGURE 58. The effect of temperature on the leaching efficiency of copper from copper sulfide. ..... 88

FIGURE 59. The effect of sodium sulfite addition on the leaching efficiency of copper from copper

sulfide. ................................................................................................................................................... 89

FIGURE 60. Plot of surface reaction model vs time for leaching of copper from copper sulfide at


FIGURE 61. Plot of ash layer diffusion model vs time leaching of copper from copper sulfide at


FIGURE 62. Plot of the mixed model vs time for leaching of copper from copper sulfide at different

temperatures. ....................................................................................................................................... 91

FIGURE 63. Arrhenius plot for copper leaching from copper sulfide. .................................................. 93


FIGURE 65. Plot of –log k vs log [CO2] for reaction order estimation. .................................................. 94

FIGURE 66. The effect of sodium sulfite addition on the leaching efficiency of copper from

chalcopyrite flotation concentrate. ...................................................................................................... 95

FIGURE 67. SEM image of (A) copper feed, (B) leach residue A and (C) leach residue B. .................... 97

FIGURE 68. Electrochemical set-up. ................................................................................................... 110

FIGURE 69. Electrochemical cell used in this project ......................................................................... 111

FIGURE 70. Iron RDE used in this project............................................................................................ 112

FIGURE 71. Leaching set up. ............................................................................................................... 113

FIGURE 72. Leaching data 1. ............................................................................................................... 114





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FIGURE 81. Leaching data 11. ............................................................................................................. 123









FIGURE 90. Dissolution of nickel sulfide. ........................................................................................... 132

FIGURE 91. Dissolution of cobalt sulfide. ........................................................................................... 132

FIGURE 92. Dissolution of copper sulfide. .......................................................................................... 132

FIGURE 93. XRD analysis of nickel sulfide. .......................................................................................... 133

FIGURE 94. XRD analysis of cobalt sulfide. ......................................................................................... 133

FIGURE 95. XRD analysis of copper sulfide. ........................................................................................ 133

FIGURE 96. Plot of surface reaction model vs time for leaching nickel from nickel sulfide at different

ammonia concentrations. ................................................................................................................... 134

FIGURE 97. Plot of ash layer diffusion model vs time for leaching nickel from nickel sulfide at

different ammonia concentrations. .................................................................................................... 134

FIGURE 98. Plot of the mixed model vs time for leaching nickel from nickel sulfide at different


FIGURE 99. Plot of surface reaction model vs time for leaching nickel from nickel sulfide at different

ammonium carbonate concentrations. .............................................................................................. 135

FIGURE 100. Plot of ash layer diffusion model vs time for leaching nickel from nickel sulfide at

different ammonium carbonate concentrations. ............................................................................... 136

FIGURE 101. Plot of the mixed model vs time for leaching nickel from nickel sulfide at different


FIGURE 102. Plot of surface reaction model vs time for leaching cobalt from cobalt sulfide at


FIGURE 103. Plot of ash layer diffusion model vs time for leaching cobalt from cobalt sulfide at


FIGURE 104. Plot of the mixed model vs time for leaching cobalt from cobalt sulfide at different


FIGURE 105. Plot of surface reaction model vs time for leaching cobalt from cobalt sulfide at


FIGURE 106. Plot of ash layer diffusion model vs time for leaching cobalt from cobalt sulfide at


FIGURE 107. Plot of the mixed model vs time for leaching cobalt from cobalt sulfide at different


FIGURE 108. Plot of surface reaction model vs time for leaching copper from copper sulfide at


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FIGURE 109. Plot of ash layer diffusion model vs time for leaching copper from copper sulfide at


FIGURE 110. Plot of the mixed model vs time for leaching copper from copper sulfide at different


FIGURE 111 . Plot of surface reaction model vs time for leaching copper from copper sulfide at


FIGURE 112. Plot of ash layer diffusion vs time for leaching copper from copper sulfide at different


FIGURE 113. Plot of the mixed model vs time for leaching copper from copper sulfide at different


FIGURE 114. Reagent amounts calculations. ...................................................................................... 143

FIGURE 115. XRD analysis of Nickel Sulfide flotation concentrate feed. ............................................ 144

FIGURE 116. XRD analysis of Nickel Leach Residue A. ........................................................................ 144

FIGURE 117. XRD analysis of Nickel Leach Residue B. ........................................................................ 145

FIGURE 118. XRD analysis of Copper Sulfide flotation concentrate feed. .......................................... 145

FIGURE 119. XRD analysis of Copper Leach Residue A. ...................................................................... 146

FIGURE 120. XRD analysis of Copper Leach Residue B. ...................................................................... 146

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List of Tables

TABLE 1. Solution compositions used in the electrochemical portion of this project. ........................ 31

TABLE 2. Elemental composition of the nickel sulfide feed (Figure 22 (A)) and the elemental

composition of the feed at certain spots from Figure 22 (C). ............................................................... 42

TABLE 3. Nickel leaching summary. ..................................................................................................... 43

TABLE 4. Correlation coefficients for the different models at various leaching conditions of nickel

sulfide. ................................................................................................................................................... 51

TABLE 5. Elemental composition of the Nickel Sulfide flotation feed and both leach residues. ......... 55

TABLE 6. Mass (%) composition of cobalt sulfide feed sample used and mineral composition at

certain spots from Figure 39 (C). .......................................................................................................... 62

TABLE 7. Cobalt leaching summary. ..................................................................................................... 63

TABLE 8. Correlation coefficients for the different models at various leaching conditions of cobalt

sulfide. ................................................................................................................................................... 73

TABLE 9. Mass (%) composition of the copper sulfide feed sample used and the mineral composition

at certain spots from Figure 54 (C). ...................................................................................................... 81

TABLE 10. Copper leaching summary. .................................................................................................. 83

TABLE 11. Correlation coefficients for the different models at various leaching conditions of copper

sulfide. ................................................................................................................................................... 92

TABLE 12. Elemental composition of the copper sulfide flotation concentrate feed and both leach

residues. ................................................................................................................................................ 96

TABLE 13. Comparison of the reaction rates and efficiencies for the different metal sulfides analysed

in this project. ....................................................................................................................................... 98

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List of Appendices

APPENDIX A. Equipment used in the project ..................................................................................... 110

APPENDIX B. SMMS leaching data and calculations .......................................................................... 114

APPENDIX C. EFMS dissolution data ................................................................................................... 132

APPENDIX D. SMMS XRD data ............................................................................................................ 133

APPENDIX E. SMMS kinetic model plots ............................................................................................ 134

APPENDIX F. Amount of reagents calculations .................................................................................. 143

APPENDIX G. RMS XRD data ............................................................................................................... 144

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Abbreviations

- Total Concentration of Carbonate species

- Total Concentration of Ammonia species

- Concentration of Ammonia added as

- Activation Energy (kJ/mol)

- Equilibrium Potential of Anodic Reaction

- Equilibrium Potential of Cathodic Reaction

- Corrosion Potential

- Partial Reaction Equilibrium Potential

- Steady State Mixed Potential

- Passivation Potential

- Anodic Peak Potential

- Cathodic Peak Potential

- Anodic Half-reaction

- Cathodic Half-reaction

- Exchange Current Density

- Cathodic Current Density

- Anodic Current Density

- Anodic Reaction Current

- Cathodic Reaction Current

- Corrosion Current

- Apparent rate constants

- Atomic Absorption Spectrophotometry

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- Counter Electrode

- Electrode Potential

- Electrical Double Layer

- Electrochemically Formed Metal Sulfides

- Faraday’s Constant

- Grazing Incidence X-ray Diffraction

- Luggin Capillary

- Mixed Hydroxide Precipitate

- Metal Oxide

- Nitrogen Line

- Universal Molar Gas Constant

- Rotating Disk Electrode

- Reference Electrode

- Real Metal Sulfides

- Saturated Calomel Electrode

- Scanning Electrode Microscopy

- Synthetic Monometallic Metal Sulfides

- Solvent Extraction and Electrowinning

- Soft X-ray Microscopy

- Working Electrode

- X-ray Diffraction

- Fraction Extracted

- Transfer Coefficient

- Electrode Overpotential

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Acknowledgements

First and foremost, I would like to thank my supervisor, Dr. Aleks

Nikoloski for providing me with endless guidance, support and advice.

I am forever grateful for all his time and dedication spent in

mentoring me throughout this challenging journey.

I would like to express my gratitude to Dr. Drew Parsons for always

being willing to help when I needed a question answered.

Special thanks to Rorie Gilligan for his time and assistance in the

laboratory.

I wish to place my sincere thanks to Dr. David Hough for his

invaluable input.

I would also like to take this opportunity to acknowledge the other

faculty in the School of Engineering and Information Technology at

Murdoch University. Their teachings have guided my academic

progress and have changed the ways in which I engage with the

world.

Finally, I would like to thank my friends and family for their

unconditional love, support, advice and encouragement. Without

them, none of this would indeed be possible.

I would like to dedicate this work to the memory of my father –

Marek.

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1 INTRODUCTION

In an ammoniacal carbonate leaching nickel, cobalt and copper sulfides are introduced

as either feed material (Sherritt Gordon process) or as formed intermediate products in

the Caron process. The Caron process is a combined pyro-metallurgical and hydro-

metallurgical process of extracting nickel and cobalt from nickel laterite ores. First, the

ore is heated up to ~700°C in a reducing atmosphere, allowing the nickel and cobalt to

be reduced to their metallic states. There are no free nickel and cobalt minerals in the

nickel laterite ore, the reduced metals will, therefore, be associated with iron (Valix and

Cheung, 2002). After the roasting stage, the calcine is cooled down to stop the nickel and

cobalt from oxidizing. An ammonia-ammonium carbonate solution is used to selectively

leach out these alloys, as both nickel and cobalt form stable ammine complexes (Valix

and Cheung, 2002). Boiling will remove the ammonia and allow the basic nickel

carbonate to precipitate out, resulting in nickel oxide being produced at 1200°C. Under

the atmosphere of hydrogen, this nickel oxide will be converted to nickel metal in

reduction furnaces (Norgate and Jahanshahi, 2011). The main advantages of this process

are its great selectivity towards the metal values, and the fact that ammonia can be

recycled in the leaching circuit.

Due to a low metal recovery of ~80% for nickel and ~60% for cobalt, however, the

Caron process has lacked popularity. Over the years the causes of this poor metal

recovery have been studied, but the observed behaviour is still unclear. A deeper

understanding of these limiting factors could help in optimising the current conditions,

leading to a higher metal recovery.

The main objective of this study is to provide a valuable insight into the leaching of

metal sulfides in a solution containing ammonia, ammonium carbonate and thiosulfate

ions. By studying the effects of the different operating conditions, the limiting factors in

the leaching mechanisms can be determined. This study aims to gain a better

understanding of how the leaching mechanisms respond to various conditions, and how

these conditions can improve the effectiveness of the Caron process.

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2 LITERATURE REVIEW

2.1 ORE DEPOSITS, MINERALOGY AND PROCESSING

Nickel (Ni) is an important metal used in the modern infrastructure and technology

(Wood, 2013). Major uses of nickel include stainless steel (~65%), metal alloys (~20%),

electroplating (~9%) and rechargeable batteries (~5%) (Mudd, 2009 & 2010). From a

metallurgical perspective nickel has a number of benefits, such as a high melting point of

1453°C, ferromagnetic properties as well as catalytic behaviour (Moskalyk and

Alfantazi, 2002).

World’s resources of nickel are present as two ore types - sulfide and laterite ores.

Figure 1 shows the nickel resources and production from laterite and sulfide ores, and

Figure 2 shows predictions for the future.

FIGURE 1. Nickel production and resources (World resources on land) left, and (Nickel production

in 2009) right, of laterite and sulfide ores (Oxley et al., 2016).

FIGURE 2. Past and future predictions of nickel production from laterite and sulfide ores (Oxley et al., 2016).

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Nickel sulfide ores derived from hydrothermal processes, often involving copper, cobalt

and sometimes platinum group metals (Oxley et al., 2016; Mudd, 2009). The production

of nickel from sulfide ores requires open cut or underground mining, a concentration

technique (flotation), smelting to produce nickel matte, followed by the refining of matte

to produce pure metal (Mudd, 2009).

Nickel laterite ores were formed as a result of extensive weathering over time, in the

tropical climates around the equator (Mudd, 2009). The main areas of nickel laterite

resources in the world are located in New Caledonia, Australia, Indonesia, South

America and the Philippines (Farrokhpay and Filippov, 2016). It has been noted that

because of the complexity of nickel laterite ores, the deposits located at the same regions

can be quite variable in nature (McDonald and Whittington, 2008).

Nickel laterite deposits are usually divided into three different zones: limonite (oxide),

nontronite (clay) and saprolite (silicate) zones (Forrokhpay and Filippov, 2016).

Saprolite, the lowest of these layers reflects the early stages of the weathering process,

and usually contains only around 1.5-3% nickel. The formula for the major mineral is

(Ni,Mg)6Si4O10(OH)8 (Pickles, 2003). The middle layer, nontronite zone, contains mostly

clay and quartz. Limonite, the top zone, consists of goethite and amorphous ferric

hydroxide, and is a result of further weathering (Farrokhpay and Filippav, 2016). The

chemical formula of limonitic type ore is (Fe,Ni)O(OH).n H2O (Pickles, 2003).

Production of nickel from laterite ore is more complex than from sulfide ores (Mudd,

2009; Rice, 2016; Farrokhpay and Filippov, 2016; Norgate and Jahanshahi, 2011). L

The laterite ore mines are all open-cut, due to a large area and shallow nature of this ore.

Figure 3 shows the typical processing routes for the different laterite ores.

4 | P a g e

FIGURE 3. Generalised laterite processing flowsheets (Norgate and Jahanshahi, 2011).

One of the main reasons for the difficulty associated with the processing of nickel from

laterite ores is that these types of ores cannot be significantly upgraded or concentrated

prior to processing (Norgate and Jahanshahi, 2011). Nearly every tonne of the ore mined

needs to be put through the entire process, resulting in very high operating costs. The

complex structure of the laterite ores has prevented the nickel grade from being pre-

concentrated by means of physical beneficiation methods. Some techniques which

include sizing, gravity and magnetic separation either individually or in combinations

have been employed to upgrade/beneficiate the laterite ores. In most cases, however,

these techniques were not able to dramatically improve the quality of the feed (Quast et

al., 2015). Quast et al. (2015), reported that the most commonly used pre-concentration

process was the removal of coarse fraction from the feed, which has been noted to have a

much lower nickel content than the finer material. Lu et al. (2013), reported that nickel in

nickel laterite ore is finely disseminated in the molecular lattice of goethite or hydroxyl

silicate. In this study, Lu et al. (2013), have documented that pyro-metallurgical

techniques are well suited to treating saprolite ore, where the garnierite content is high

and iron content is low.

Another problem arising from the present treatment of laterite ores is the fact that nickel

oxide ores may contain around 45% of water, both as free moisture and also as water

5 | P a g e

chemically combined in hydrated minerals, which means that thermal reduction of the

laterite ore is required prior to processing (Moskalyk and Alfantazi, 2002).

2.2 AMMONIACAL LEACHING PROCESS

Leaching has been described as a primary extractive operation in hydrometallurgy. The

process involves the dissolution of a valuable mineral as the ore comes in contact with an

active chemical solution, known as the leach liquor. This dissolution process is selective,

affecting only the chosen component, with the unwanted (gangue) components

remaining in the solid state. The valuable metal can then be extracted, as the solids are

separated from the leach liquor at the end of the process (Gupta, 2003).

For the efficient leaching of a metal from its ore, the appropriate chemical reactions must

take place, and these reactions are governed by the thermodynamic characteristics of the

particular reaction system, as well as the solubility of the species in the aqueous solution

(Mathobi, 2012).

According to Woollacott and Eric (1994), the following mechanisms will be involved in

a single stage leaching process:

1) There must be a diffusion of the reagents to the mineral surface, as well as

adsorption of the reagents on the surface, while an active chemical reaction is

taking place on the surface, and

2) There must be desorption and diffusion of the reaction products from the surface

into the bulk of the solution.

The rate of the extraction will be controlled by the slowest of these mechanisms. There

are a number of parameters affecting the efficiency of the metal extraction. The size of

the particles, (in other words, the degree of exposure), will affect the rate of leaching,

generally increasing the rate as the particle size decreases. The diffusion rate of reactants

and products will also influence the rate of leaching. This can be improved by an

increase in the degree of agitation of the leach solution. When the rate of chemical

reaction influences the rate of leaching, this can be improved by increasing the degree of

exposure of the valuable metal, or increasing the temperature or pressure of the leaching

system. The rate of leaching will increase with the increase in the concentration of the

leaching agent, and with decreasing pulp density. The rate will also be influenced by a

6 | P a g e

possible insoluble reaction product formed during leaching, with the rate of leaching

dependent on the nature of this product (Woollacott and Eric, 1994).

Metal sulfides are insoluble in water. In the presence of oxygen, however, they will be

solubilised as sulfates. According to Habashi (1970), the desired residue product from the

oxidation of sulfides will be not only sulfates, but elemental sulfur as well. When

leaching sulfides, the temperature and the pH at which the leaching will take place need

to considered. Depending on these two factors, there is a possibility of obtaining both

metal and sulfur in soluble form, with sulfur recovered as ammonium sulfate (Habashi,

1970).

The electrochemical mechanism of leaching sulfides can be represented by the following

reactions:

Cathodic Reaction: [Eq.1]

Anodic Reaction: [Eq.2]

An ammonia-ammonium carbonate system was first used to recover non-ferrous metals

such as copper from oxide ores (Radmehr et al., 2014), but gradually this method was

developed to extract nickel and cobalt as well as other traditional elements. Aqueous

ammonia has been an effective lixiviant, forming stable metal ion ammine complexes

with base metal cations, while rejecting iron (Moyo and Petersen, 2016). Ammonia in

the complex ligands binds to the metal ions through the reactive nitrogen-containing

group, which in turn leads to a higher solubility (Nabizadeh and Aghazadeh, 2015).

According to Meng and Han (1995), there are three main methods of ammonia leaching:

neutral, oxide and the reduction method. During the neutral method, metals will be

leached without the presence of any oxidizing or reducing agents. Most copper and zinc

oxides are leached under this method. The oxide method, is a method which requires the

use of an oxidizer in order to oxidize the solids, such as sulfide minerals. The last type of

method is the reduction method, where a reducing agent is used for dissolving metals

from highly oxidized ores.

Ammonia leaching has a number of advantages over other leaching methods (Radmehr

et al., 2014; Meng and Han, 1995; Wei Wei, 2016; Wetham et al., 2015 and Katsiapi et

al., 2010), which include low toxicity, low cost and the ease of regeneration.

7 | P a g e

The chemistry of the oxidative ammonia leaching is quite complex, and involves a

number of steps. The sulfide component is oxidized to soluble species, mostly sulfate

(SO42-) together with the thiosulfate (S2O3

2-) and sulfamate (NH4SO3NH2). The iron

component will be oxidized to the +3 state, and later precipitated as hydrated ferric

oxide (Fe2O3.H2O). The metal ion will be released in its +2 oxidation state, and

stabilized by the formation of ammine complexes with NH3 ligands (Nabizadeh and

Aghazadeh, 2015).

Katsiapi et al. 2010 presented the dissolution of divalent metal oxide (MeO) in an

ammonia solution by following:

[Eq.3]

There are two main types of ammoniacal leaching processes currently used around the

world, the Caron process and the Sherrit Gordon process.

2.2.1 The Caron Process

The Caron process (Caron, 1955) was first used in the 1950s for the extraction of nickel

and cobalt in the ammoniacal solution (Zuniga et al., 2010). Since then, this process has

been applied in Cuba (Nicaro), Australia (Townsville), the Philippines (Mavinduque)

and China (Qinghai) as well as in Brazil (Niquelandia), where nickel has mainly been

associated with copper rather than cobalt (Ma et al., 2013).

The Caron process has been successfully used for processing nickel laterite ores. It

provides a combined pyro/hydro-metallurgical approach to nickel and cobalt recovery

(Lu et al., 2013; Mudd, 2009; Rice, 2016; Valix and Cheung, 2002; Norgate and

Jahanshahi, 2011). This process consists of drying and grinding laterite ore, reduction

roasting, leaching with ammoniacal-ammonium carbonate solution to dissolve nickel

and cobalt as ammine complexes, and the recovery of base metal from solution to

produce a nickel oxide product (Moskalyk and Alfantazi, 2002).

The Caron process has proved successful for application to processing of low-grade

nickel laterite oxides (~1.5% Ni, ~0.1% Co) (Valix and Cheung, 2002), in particular has

8 | P a g e

been successfully applied to ores with high goethite (FeOOH) or limonite ores, and ores

with low silicate content (Norgate and Jahanshahi, 2011). Norgate and Jahanshahi

(2011) noted that with the increase of saprolite, the recovery of nickel and cobalt will be

reduced, because nickel and cobalt are locked in a silicate matrix and become difficult to

reduce at 700°C.

Rice (2016) reported that limonite ores with a lower magnesia content responded better

to the hydro-metallurgical techniques of the Caron process, and it has also been noted

that this process can tolerate higher amounts of magnesium than any other process

currently used. When the Caron process was studied in the extraction of nickel from

high silicate ores, it was documented that the control of the reduction conditions was

critical for the recovery of nickel from these ores as well as for better iron separation

(Rice, 2016).

Results from Valix and Cheung (2002) study showed that the temperature of the

reduction had a significant effect on the recovery of both nickel and cobalt. The authors

of this report presented the optimum nickel and cobalt recovery conditions from

limonite, with increasing the temperature up to 600°C as recommended, while the

weathered saprolite reduction was favoured at ~800°C. It was noted that at this

temperature, an association with a complete dehydration is achieved, but if the

temperature is being increased further, the recovery will in fact be reduced (Valix and

Cheung, 2002). The phases which form at 800°C will not be reversed upon cooling the

laterite ore (Valix and Cheung, 2002).

Ma et al. (2013), described the Caron process as particularly effective in the treatment of

high goethite containing limonitic laterite ores. This is because silicate-rich minerals tend

to undergo phase transformation during the reduction roasting stage, which impedes the

extraction of nickel (Ma et al., 2013).

There have been a number of problems associated with the current Caron process,

however, reported in the literature (Norgate and Jaharshahi, 2011; Senanayake et al.,

2010; Zuniga et al., 2010; Dyer et al., 2012). The processes involved in the pyro-

metallurgical stage of the Caron process (drying, calcining and reduction) are all energy

intensive and costly (Norgate and Jahanshahi, 2011; Senanayake et al., 2010). The

leaching kinetics of the process are still slow and limited to ~80-86% Ni and ~50-60% Co

9 | P a g e

extractions (Dyer et al., 2012). These require various reagents, high temperature,

ammonia and metallic iron concentrations (Zuniga et al., 2010), and there is still a high

nickel and cobalt content waste in the tailings (Senanayake et al., 2010).

2.2.2 The Sherritt Gordon Process

The developments of ammonia leaching processes gained popularity from their initial

applications to the developments of the base metals recovery. In 1947, Sherritt Gordon

invented a process of ammonia leaching of nickel, cobalt and copper sulfides, conducted

at 105°C with air pressure of 0.8 MPa (Radmehr et al., 2014). Under these conditions

copper, nickel and cobalt could be dissolved with iron precipitating in its hydroxide

form. The pregnant solution was heated to evaporate the excess ammonia, later allowed

to react with air to oxidize its thiosulfates. Reduction continued till the precipitation of

nickel with little effect on the cobalt ions in solution (Radmehr et al. 2014). The Sherritt

Gordon process was the first successful commercial ammonia leaching of nickel sulfide

concentrates (Wei, 2016). It was later developed for nickel matte from smelting and

nickel-containing alloy scraps. Figure 4 shows the Sherritt Gordon process flowsheet

which is currently being used by the Kwinana Nickel Refinery in WA.

10 | P a g e

FIGURE 4. A flowsheet showing the Sherritt Gordon process.

MATTE GRIND

• Contains: Nickel, copper, cobalt, sulfur and iron

LEACH

• Nickel, copper and cobalt dissolved; sulfide oxidised

to sulfate; iron converted to insoluble iron oxide

SOLID / LIQUID

• Iron residue

COPPER BOIL

• Copper seperates as copper sulfide

OXYDROLYSIS

• Unwanted sulfur compounds changed to

sulfate

NICKEL REDUCTION

• Nickel powder precipitated

METAL STRIPPING

• Mixed nickel / cobalt sulfide precipitated

CRYSTALLISATION

• Ammonium sulfate precipitated

11 | P a g e

2.3 FACTORS AFFECTING NICKEL AND COBALT EXTRACTION

There are a number of studies in which higher nickel extraction was attributed to pH

and temperature increases (Kumbasar and Kasap, 2009; Muharok and Lieberto, 2013;

Zuniga et al., 2010; Sarma and Nathsarona, 1996; Hu et al., 2012; Li et al., 2010). Hu et

al. (2012), described various techniques for nickel recovery from ammoniacal leach

solutions. These included precipitation and solvent extraction. In their study Hu et al.

(2012), tested the β-diketones as an alternative for nickel extraction. They presented β-

diketones as a versatile class chelating agents, which had been recommended for nickel

extraction. In this study the authors found that the extraction efficiency of nickel

increased with increased pH, ammonia concentration and temperature. The maximum

extraction was achieved at pH~8.5, and decreased at higher pH.

Hu et al. (2012), reported that the inert electrolyte (Na2SO4) did not affect the distribution

coefficient of nickel (II). The increase of the total ammonia concentration, however, was

seen to depress the extraction of nickel (Hu et al., 2012). The authors attributed the

decrease in the extraction efficiency mainly to the formation of nickel ammine

complexes, because the replacement of H2O in the [Ni(H2O)6]2+ by NH3 ligand increases

in their stability in the aqueous phase (Hu et al., 2012). It was also noted nickel

extraction will increase in temperature at pH~7.5, but will be reduced with the increase

in temperature at pH~8.5 (Hu et al., 2012).

A similar study was presented by Sarma and Nathsarona (1996), who reported that the

influence of pH on the extraction efficiency was found to be more pronounced in the

range 9.5-10.0, compared to 8.6-9.5. They suggested that the ammonium carbonate

concentration had very little influence on the nickel extraction, at the concentration

range of 45-75 kg/m3. It was also noted that extraction of nickel was reduced by 20% at a

salt concentration in solution of 150 kg/m3 (Sarma and Nathsarona, 1996).

Zuniga et al. (2010), reported the kinetics of nickel extraction as slow, and requiring high

temperatures, ammonia and metallic iron concentrations. They suggested that cobalt

concentrations decrease due to precipitation with or on Fe3+ oxide/hydroxides. The

losses of cobalt are affected by increasing ammonia concentrations and temperature,

because these will enhance the formation of Fe3+ oxides/hydroxides. It was found,

however, that if the ammonium sulfate concentration is increased, this will in fact

decrease cobalt losses (Zuniga et al., 2010).

12 | P a g e

In another study led by Kumbasar and Kasap (2009), the increases in nickel extraction

were attributed to ammonia concentrations increases to 6M at a fixed pH.

Senanayake et al. (2010), described the rate of iron dissolution as slow, with low

extraction of

13 | P a g e

dissolved (DeGraaf, 1979). It was noticed that smaller particle size increased the

recovery of nickel but had no effect on the dissolution rate. For the reduction of silicate

ore with CO, fine grinding was necessary for high nickel extraction, regardless of

reducing conditions (DeGraaf, 1979).

In 1980 DeGraaf presented his second study of the leaching reactions of the Caron

process. This study suggested the addition of sulfur as pyrite during the reduction stage

as beneficial to the nickel extraction and beneficial in decreasing the sensitivity to the

reduction conditions. It was also reported that finer grinding and repeated reduction of

leach residue resulted in some increases in the extraction (DeGraaf, 1980).

2.4 IRON PASSIVATION LAYER FORMATION

The passivation of iron has been identified as a major factor limiting the efficiency of the

Caron process and the recovery of nickel and cobalt from laterite ores (D’Aloya and

Nikoloski, 2012). D’Aloya and Nikoloski (2012) studied the passivation of iron in

ammoniacal solutions containing Cu2+ ions, and reported that the presence of these ions

in solutions will in fact promote the passivation of iron. As the concentrations of copper

are increased, and the concentration of ammonia-ammonium bicarbonate is lowered, the

passivation of iron will occur more readily (D’Aloya and Nikoloski, 2012). An open

circuit measurement was performed, which was found to increase from -0.7 V to -0.6 V

after 3 hours in aqueous ammonia and ammonium bicarbonate-ammonium carbamate

double salt and 0.2mM copper sulfate. The experiment was repeated at a 12mM

concentration of Cu2+, with the OCP increasing to around -0.5 V, before a further

increase to a high potential region of 0.1-0.2 V, where it stayed for the remainder of the

test.

The rotating disk cyclic voltammetry studies were also carried out on solution containing

ammonia-ammonium bicarbonate with 6mM Cu2+ ions. The RDE was immersed in a

solution under the OCP conditions for 2 minutes, with the potential scanned in the

positive direction at a rate of 10 mV/s from the OCP to +0.24 V, then reversed back to

-0.56 V. Peak (A) was observed close to -0.43 V during the anodic sweep, which has

been attributed to the anodic dissolution of metallic iron, with a second peak (B) noticed

at around -0.13 V, assigned to the re-dissolution of cemented copper. During the

cathodic scan, the anodic dissolution of iron was noticed at around -0.4 V, peaking at

-0.47 V (D’Aloya and Nikoloski, 2012).

14 | P a g e

The mechanism of passivation, as described by D’Aloya and Nikoloski (2012), involved

cementation of copper onto the actively dissolving iron surface. During this process of

copper cementation, the insoluble oxide compound will become the preferred product of

iron oxidation. A thin, passive layer will form as a result of the re-dissolution of the

cemented copper, when the copper is no longer cathodically protected by metallic iron

(D’Aloya and Nikoloski, 2012).

Dyer et al. (2012), have studied cobalt loss due to iron precipitation in ammoniacal

carbonate solutions. The slow and limited leaching kinetics and poor recovery were

linked to the possible fuel oil containing sulfur, and the formation of nickel sulfide during

the reduction stage, which is less reactive than elemental nickel or alloys, a formation of

iron as oxide/hydroxide and a formation of iron passive layer by precipitated iron

species (Dyer et al., 2012). It was found that because of the sensitivity of iron oxide

formation under conditions such as pH, temperature, residence time or the solution

potential, it is difficult to predict the exact phase produced. The authors of this study

reported that the initial iron concentration was unlikely to influence the adsorption of

cobalt. It was reported that higher iron levels produced greater mass of precipitate,

however, with greater surface area and, therefore, more binding sites for cobalt (Dyer et

al., 2012). It was found that cobalt loss peaks at pH 7, lower at pH 6, and all values

above pH 7, with a sharp decrease observed at pH 10. The presence of cobalt during iron

precipitation resulted in greater cobalt losses at a number of different iron

concentrations. It was noted that the mechanism of cobalt loss during iron precipitation

is adsorption, not the precipitation of a secondary phase (Dyer et al., 2012). The

increases in ammonia concentrations were seen as beneficial, limiting the proportion of

cobalt being removed. No indications that the temperature between 25-45°C has any

influence over cobalt adsorption was reported.

The majority of iron dissolution is thought to take place under oxygen-free conditions,

resulting in the dissolved iron precipitating as ferric hydroxide. The rejection of iron at

the leaching stage is desirable. The passivation of the metallic iron has had a negative

effect on the extraction efficiency, however (D’Aloya and Nikoloski, 2013). The

formation of CoSx and its effects on the anodic dissolution of iron in ammoniacal-

carbonate solutions were studied by D’Aloya and Nikoloski (2013). This study showed

that during the dissolution of iron in ammoniacal-carbonate solutions in the presence of

15 | P a g e

Co2+ and thiosulfate ions, the CoSx species are formed on the iron surface. It has been

reported that the formation of these species is influenced by the passivation of iron, and

that these species produce negative effect in the leaching stage of the Caron process,

causing the loss of the dissolved cobalt from the leach solution and reducing the

efficiency of the extraction (D’Aloya and Nikoloski, 2013). The authors noted that this

passivation layer can be prevented if the concentration of ammonia is maintained higher

than 4M.

In another study D’Aloya and Nikoloski (2014) studied the anodic dissolution of iron in

ammoniacal-carbonate-thiosulfate-copper solutions. It has been shown that the

behaviour of iron strongly influences the dissolution of alloys and the recovery of metals.

The solutions containing both the dissolved cobalt and thiosulfate ions show a tendency

to significantly suppress the passivation and the dissolution of iron (D’Aloya and

Nikoloski, 2014). It was also reported that if both Cu2+ and thiosulfate ions are present in

the leach solutions, they will also suppress the passivation of iron.

D’Aloya and Nikoloski (2014) observed that in solutions which contained cobalt and

thiosulfate ions, even small additions of copper ions were seen to delay the onset of

passivation. The passivation of the RDE at a concentration of copper ions as low as

0.1mM in an ammoniacal-carbonate solution containing 12mM Co2+ and 24mM S2O32-

took place after 2 hours of immersion, compared to 30-40 minutes with no copper

addition.

The rotating disk cyclic voltammetry measurements were performed for 15 minute

immersions in 6M[NH3]T , 2M[CO2]T solutions at different Cu2+ and S2O3

2-

concentrations. The potential was scanned from the OCP, resulting in an anodic peak

(A) around -0.6 to -0.2 V, followed by a second anodic peak (B) at around -0.1 to 0 V

and a third peak (C) above 0.1 V. The increase in Cu2+ concentration from 1mM to

4mM at a constant S2O32- resulted in a 7-fold increase in peak (C) (D’Aloya and

Nikoloski, 2014). It was found, however, that the increase of S2O32- concentration at a

constant Cu2+ results in a 10% decrease in the current density peak (C). The authors

attributed this peak (C) to the oxidation of copper-containing species forming during the

anodic dissolution of iron, and possibly as a result of copper (II) reduction to copper (I)

on the surface of the iron electrode. During the reverse scan, the anodic dissolution of

16 | P a g e

metallic iron resumed around -0.4 V with reported higher current densities reached at

lower Cu2+ and higher S2O32- (D’Aloya and Nikoloski, 2014).

D’Aloya and Nikoloski (2014) observed a black solid layer forming on the iron surface

during the dissolution, a layer which contained Cu2S and dendritic copper. It was

suggested that this layer formation is possibly due to the effect of copper cementation. It

was noted that this layer remained in electrical contact with the iron RDE surface

(D’Aloya and Nikoloski, 2014). The authors commented that the formation of the

deposit on the iron surface did not suppress the dissolution of iron, as was the case with

(D’Aloya and Nikoloski, 2013), nor was it found to promote passivation (D’Aloya

and Nikoloski, 2014).

In another study this iron passivation was linked to the oxide layer being in the presence

of high concentrations of dissolved oxygen and/or other oxidants (Nikoloski and Nicol,

2010). Nikoloski and Nicol (2010) studied the cathodic processes involved in the

leaching reactions in the Caron process. It was found that Co3+ ions were the main

oxidizing agents involved in the dissolution of nickel, cobalt and iron in ammonia-

ammonium carbonate solutions. The authors reported that a sulfide layer will be formed

from a reaction involving the reduction of thiosulfate in the presence of nickel and cobalt

metal ions, and that this additional cathodic reaction could enhance the overall rate of

dissolution of iron alloys in an ammoniacal-carbonate solutions (Nikoloski and Nicol,

2010).

2.5 ELECTROCHEMICAL PRINCIPLES

2.5.1 Metal Electrolyte Interface - The Electrical Double Layer

As the metal is immersed in an electrolyte, the electrons from the metal ions will

separate and remain in the metal. Surrounded by the water molecules, these metal ions

will start to diffuse away from the metal. The excess number of electrons on the metal

surface will attract the positively charged metal ions, so instead of diffusing into the bulk

electrolyte, these metal ions will in fact remain near the metal surface. The water layer

around the metal ions, however, will prevent them from making direct contact with the

electrons on the surface, so stopping them from being reduced to metal atoms. This

electrolyte layer adjacent to the electrode, containing water molecules and ions from

both metal and bulk electrolyte, will possess its own unique chemical composition,

17 | P a g e

significantly different from the bulk electrolyte. An electrical double layer (EDL) is a

term used to describe the combination of this negatively charged surface of the metal and

the adjacent electrolyte layer.

The first concept of the ‘electrical double layer’ was developed by Helmholtz in the 19th

century (Helmholtz, 1853). The model proposed a theory of electron transfer occurring

at the electrode with the solution composed only by electrolyte. A good basis for the

understanding of the solid-liquid interface was provided, but, this model failed to

account for the diffusion and mixing in solution, as well as the possible surface

adsorption.

Guoy (1910) and later Chapman (1913) proposed a more advanced theory of a ‘diffused

electrical double layer’. This model described the electrolyte in terms of a number of

oppositely charged ions, whose concentration decreases as the distance from the surface

is increased. This theory was not entirely accurate, however, as the anions and cations

already existed in the solution. There was the probability that with the increased distance

from the surface, the ions of the same charge as the surface will be found within this

double layer would also increase. Stern (1924) worked on developing this ‘diffused

electrical double layer’ theory further and combined the Helmholtz layer with the Guoy-

Chapman theory. His theory described the ‘electrical double layer’ in terms of two

separate layers- the inner Stern layer consisting of adsorbed and immobile ions right on

the solid, and the diffuse outer layer of mobile charge carriers, passing into the bulk

solution. The nature of the interactions between the electrode surface and the ions in

solution was assumed to be electrostatic, caused by either the excess or a deficiency of

electrons at the electrode surface (Stern, 1924). For the interface to remain neutral, ions

close to the electrode surface must have been constantly redistributed. The attracted ions

approaching the surface would need to form a layer, which balances the electrode

charge. This would result in two layers of charge - hence the double layer - and a

potential drop when these two are separated (Wang and Pilon, 2011). Figure 5 shows the

schematics of the double layer theory.

18 | P a g e

FIGURE 5. The electric field consisting of two layers of charge – The Electrical Double Layer (faculty.kfupm.edu.sa)

2.5.2 Corrosion Potential

Corrosion potential is described as the potential of an electrode of metal corroding in an

aqueous solution (Sato, 2012). It represents the point where the anodic oxidation current

of metal dissolution will be equal to the cathodic reduction current of the oxidant. Figure

6 schematically describes the electrode potential versus reaction current curves of anodic

oxidation and cathodic reduction.

19 | P a g e

FIGURE 6. Polarization curves of corrosion reactions, where ia is the anodic reaction current, ic is the cathodic reaction current, icorr is the corrosion current, Ea is the equilibrium potential of the

cathodic reaction, Ec is the equilibrium potential of the cathodic reaction and Ecorr is the corrosion potential (Sato, 2012).

It is possible to control the rate of corrosion of metals by controlling either the anodic or

the cathodic reactions.

The anodic metal dissolution can be represented as an exponential function of the

electrode potential (E), as follows:

[Eq.4]

where and are parameters.

The cathodic current (ic) of the oxidation reaction is also an exponential function of

electrode potential (E), as follows:

[Eq.5]

The cathodic current will increase exponentially with increasing cathodic electrode

potential in the more negative direction (Sato, 2012).

20 | P a g e

2.5.3 Mixed Potential Theory

When a catalystic surface is introduced to an electrolyte containing Mn+ ions and a

reducing agent, the partial oxidation and reduction reactions will occur simultaneously.

A steady-state mixed potential (Emix) will be established as these two partial reactions

strive to establish their own equilibriums (Eeq). It has been noted, that the value of the

mixed potential is not a determination of the two individual thermodynamic equilibrium

potentials, but in fact is determined by their reaction kinetics being closer to the

equilibrium potentials of the faster reacting species (D’Aloya, 2015). As the redox couple

raises anodically from Eeq.Red, the potential of the metal electrode M/Mn+ is reduced

cathodically from the value Eeq.M, down to the value of the mixed potential (Emix).

According to Paunovic and Schlesinger (2006), there are four characteristic aspects of the

mixed potential theory:

1) The characteristic equilibrium potentials of both redox systems will be shifted by

the amount of over-potential , as shown in Eq.6 and Eq.7,

[Eq.6]

[Eq.7]

2) An electrochemical reaction will take place at each redox system, as the mixed

potential shifts their equilibrium potential.

3) During the reduction of Mn+, the cathodic current density (iM), and the anodic

current density (iRed) will be equal, according to Eq. 8, since an isolated system

has no net current,

[Eq.8]

4) The free energy change is not zero, since the system is not in equilibrium.

The Butler-Volmer equation describes the current density at an electrode in terms of the

over-potential and has been used to calculate the rate of electrochemical reaction,

according to the formula:

[Eq.9]

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where:

= current density as a result of imposed potential (A/cm2)

= exchange current density (A/cm2)

= transfer coefficient

= electrode over-potential

F = Faraday’s constant (C/mol)

n = number of electrons in the half-reaction

R = universal molar gas constant (kJ/mol)

T = temperature (°K)

The Butler-Volmer equation represents the sum of the total anodic half-reaction (IA) and

the total cathodic half-reaction (IC), according to these two equations:

[Eq.10]

[Eq.11]

A higher exchange current will results when the anodic half-reaction is paired with

cathodic half-reaction, according to the following:

[Eq.12]

Figure 7 shows the mixed potential (Emix) as the point in the middle of two redox

coupled equilibrium potentials.

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FIGURE 7. A schematic representation of the Mixed Potential Theory for two redox couples (Nikoloski, 2002).

2.5.4 Cyclic Voltammetry

Cyclic Voltammetry is a technique used for analysing electrochemical reactions. It helps

in analysing the thermodynamics of redox processes as well as understanding the

kinetics of the electron-transfer reactions. By applying a voltage to an electrode

immersed in an electrolyte solution, the system’s responses can be analysed. Cyclic

Voltammetry uses the three electrode system, which consists of a working electrode, a

reference electrode and a counter electrode. By varying the magnitude of the applied

potential to the working electrode, the electrode itself will become a stronger oxidant or

reductant, depending on whether the potential is being increased or decreased with time

(Kissinger and Heineman, 1984). The potential is measured between the working and

the reference electrodes, and the current is measured between the working and the

counter electrodes. The role of the counter electrode is to prevent any current running

through the reference electrode, as this would have a negative impact on the potential of

this reference electrode.

During Cyclic Voltammetry, the potential of the working electrode is scanned rapidly

over a wide potential range, returning back to its initial value with an applied potential

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signal, which varies linearly with time (Compton and Banks, 2007). The working

electrode is subjected to a triangular potential sweep, as the potential rises at a constant

rate from the start value (Epa) to the final value (Epc), then returns back to the starting

potential. The current measured during this process is referred to as current density, and

will depend on the surface area of the electrode (Compton and Banks, 2007). A cyclic

voltammogram will represent the current density plotted against the applied potential. A

typical potential versus time profile applied to the working electrode

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