chemical modeling of iron (ii)/(iii) solutions in hydrometallurgy

CHEMICAL MODELING OF IRON(II)/(III) SOLUTIONS IN HYDROMETALLURGY USING OLI™

by

Michael Kardono Carlos

A thesis submitted in conformity with the requirements for the degree of Master of Applied Science

Department of Chemical Engineering and Applied Chemistry University of Toronto

© Copyright by Michael Kardono Carlos 2013

ii

CHEMICAL MODELING OF IRON(II)/(III) SOLUTIONS IN HYDROMETALLURGY USING OLI™

Michael Kardono Carlos

Master of Applied Science

Department of Chemical Engineering and Applied Chemistry

University of Toronto

2013

Abstract

Iron is the most common impurity in hydrometallurgy which is usually removed by precipitation

of insoluble iron compounds, such as hematite and jarosite. The knowledge of iron solubility in

multicomponent solutions is important for design and optimization of the iron removal steps.

The OLI Software package is a chemical modeling tool that incorporates the powerful mixed-

solvent electrolyte (MSE) model capable of performing simulations of multicomponent

electrolyte solutions from the freezing point up to the limit of fused salt and near the critical

temperature of the solution. Literature or experimental solubility data was fitted on the OLI

MSE model to improve the performance in simulating multicomponent Fe(II)/Fe(III) solutions.

The particular focus of this work aimed at developing simulation capability for the FeCl3-MgCl2-

HCl-H2O system through experimental solubility measurement and modeling, relevant to

atmospheric processing of saprolites by HCl using MgCl2 brines.

iii

Acknowledgments

Foremost, I would like to acknowledge my supervisor, Prof. Vladimiros G. Papangelakis for his

expertise, enthusiasm, patience, support and guidance over these three years which have made

me possible to finish my master study and achieve invaluable lessons which prepared me to step

into the real world.

I would also like to thank Prof. Charles Jia and Prof. Donald Kirk for spending their precious

time to sit down during my oral defense.

I would like to thank APEC group for their continuous support and creating great lab experience,

in particular, Igor Guzman, for offering his helping hand with my experimental setup; Ilya

Perederiy, Georgiana Moldoveanu, Douglass Duffy, and Anirudha Dani for their constructive

advice for my research. I also thank Srinath Garg, Steven Palmer, Wendy Zhou, Nazanin

Samadifard for their support.

I would like to thank George Kretschmann with the advice and lesson to conduct XRD analysis,

also to Dr. John Graydon for sharing his expertise with the TGA analysis.

I would like to thank Barrick Gold, Vale Canada Ltd., and OLI for their generous financial

support. I would also like to thank Dr. Peiming Wang for the technical support with performing

OLI regression.

I would like to acknowledge my families and friends for the continuous love and support which

has encouraged me to accomplish this milestone.

iv

Table of Contents

Acknowledgments .......................................................................................................................... iii

Table of Contents ........................................................................................................................... iv

List of Tables ............................................................................................................................... viii

List of Figures ................................................................................................................................ xi

Chapter 1 ......................................................................................................................................... 1

1 Introduction ................................................................................................................................ 1

1.1 Iron chemistry and its importance in hydrometallurgy ....................................................... 1

1.2 OLI Analyzer ...................................................................................................................... 4

Chapter 2 ......................................................................................................................................... 6

2 Literature Review and Theoretical Background ........................................................................ 6

2.1 OLI MSE Model ................................................................................................................. 6

2.1.1 Chemical Equilibrium ............................................................................................. 6

2.1.2 The HKFT Model ................................................................................................... 7

2.1.3 Activity Coefficient Model ..................................................................................... 7

2.2 Hydrometallurgical Treatment of Saprolite in Concentrated MgCl2 Brines ...................... 9

2.2.1 Current Technologies to Process Saprolite ............................................................. 9

2.2.2 More concentrated MgCl2 solutions ..................................................................... 10

2.3 Default OLI MSE's Performance in Modeling FeCl3 and MgCl2 Systems ...................... 14

2.3.1 Solubility of FeCl3 in water .................................................................................. 14

2.3.2 Solubility of FeCl3 in HCl solutions ..................................................................... 15

2.3.3 Solubility of Hematite in HCl solutions ................................................................ 15

2.3.4 Solubility of MgCl2 in Water ................................................................................ 16

2.3.5 Solubility of MgCl2 in HCl Solutions ................................................................... 17

v

2.3.6 Solubility of Hematite in MgCl2 and HCl Solutions ............................................ 18

2.4 Objectives ......................................................................................................................... 20

Chapter 3 ....................................................................................................................................... 22

3 Experimental ............................................................................................................................ 22

3.1 Experimental Setup ........................................................................................................... 22

3.2 Reagents ............................................................................................................................ 23

3.3 Experimental Procedures .................................................................................................. 24

3.3.1 FeCl3 Solubility in MgCl2 Solutions ..................................................................... 24

3.3.2 MgCl2 Solubility in FeCl3 Solutions ..................................................................... 24

3.3.3 Hematite Solubility in Acidic MgCl2 Solutions .................................................... 25

Chapter 4 ....................................................................................................................................... 26

4 Experimental Results and Discussion ...................................................................................... 26

4.1 Solubility of FeCl3 in MgCl2 Solutions ............................................................................ 26

4.1.1 Reproducibility Tests for FeCl3 Solubility in Water ............................................. 26

4.1.2 Kinetic Tests for FeCl3 Solubility in MgCl2 solutions.......................................... 26

4.1.3 FeCl3 Solubility Data in MgCl2 Solutions ............................................................ 28

4.2 Solubility of MgCl2 in FeCl3 Solutions ............................................................................ 32

4.2.1 Reproducibility Tests for MgCl2 Solubility in Water ........................................... 32

4.2.2 Kinetic Tests for MgCl2 Solubility in FeCl3 Solutions ......................................... 33

4.2.3 MgCl2 Solubility in FeCl3 Solutions ..................................................................... 34

4.3 Solubility of Hematite in MgCl2 and HCl Solutions ........................................................ 35

4.3.1 Reproducibility of Hematite Solubility in MgCl2 and HCl Solutions .................. 35

4.3.2 Kinetic Tests for Hematite Solubility in MgCl2 and HCl Solutions ..................... 37

4.3.3 Hematite Solubility Data in MgCl2 and HCl Solutions ........................................ 38

Chapter 5 ....................................................................................................................................... 40

5 OLI Modeling Results and Discussion .................................................................................... 40

vi

5.1 Fitting on the FeCl3-MgCl2-H2O System .......................................................................... 40

5.2 Maximum Achievable Solid Loading during Leaching .................................................... 44

5.3 Effect of FeCl3 Addition on MgCl2 Phase Diagram ......................................................... 47

5.4 Limitation of The OLI Model ........................................................................................... 48

5.5 OLI Model Improvement for Other Fe(II) and Fe(III) Systems ....................................... 51

Chapter 6 ....................................................................................................................................... 53

6 Conclusions .............................................................................................................................. 53

Chapter 7 ....................................................................................................................................... 56

7 Recommendation for Future Work .......................................................................................... 56

References ..................................................................................................................................... 57

Appendices .................................................................................................................................... 67

Appendix A: Determination of The Concentration of Fe(III), Mg and Free HCl ........................ 67

Appendix A-1: Determination of the concentration of Fe(III) ................................................. 67

Appendix A-2: Determination of the concentration of Mg ...................................................... 68

Appendix A-3: Free HCl determination ................................................................................... 70

Appendix B: Solubility Data of FeCl3, MgCl2 and Hematite in the Chloride Systems ................ 72

Appendix B-1: Solubility Data of FeCl3 in MgCl2 (and HCl) Solutions ................................. 72

Appendix B-2: Solubility Data of MgCl2 in FeCl3 Solutions .................................................. 78

Appendix B-3: Solubility Data of Hematite in MgCl2 and HCl Solutions .............................. 80

Appendix C: XRD Patterns of The Equilibrating Solid Phases .................................................... 88

Appendix C-1: XRD Patterns for FeCl3 Solubility Experiments ............................................. 88

Appendix C-2: XRD Patterns for MgCl2 Solubility Experiments ......................................... 102

Appendix C-3: XRD Patterns for Hematite Experiments ...................................................... 107

Appendix D: Analysis of the Double Salt 2.5FeCl3.MgCl2.7.5H2O ........................................... 109

Appendix E: Validation Plots and Model Improvement for Fe(II)/Fe(III) Systems .................. 113

Appendix E-1: FeSO4-H2O System ....................................................................................... 113

vii

Appendix E-2: FeSO4-H2SO4-H2O System ........................................................................... 114

Appendix E-3: FeSO4-MgSO4-H2O System .......................................................................... 117

Appendix E-4: FeSO4-MgSO4-H2SO4-H2O System .............................................................. 118

Appendix E-5: FeSO4-ZnSO4-H2SO4-H2O System ............................................................... 120

Appendix E-6: FeCl2-H2O System ......................................................................................... 122

Appendix E-7: FeCl2-HCl-H2O System ................................................................................. 123

Appendix E-8: FeCl2-MgCl2 System ..................................................................................... 125

Appendix E-9: FeCl2-MgCl2-HCl-H2O System ..................................................................... 126

Appendix E-10: Fe2O3-H2SO4-H2O System .......................................................................... 128

Appendix E-11: Fe2O3-MgSO4-H2SO4-H2O System ............................................................. 130

Appendix E-12: NaFe3(SO4)2(OH)6-H2SO4 System .............................................................. 132

Appendix F: Lists of Regressed OLI Parameters ....................................................................... 134

viii

List of Tables

Table 1: Regressed OLI parameters for FeCl3-MgCl2-H2O system ............................................. 40

Table 2: Comparison of the thermodynamic properties of the double salts ................................. 41

Table 3: Saprolite Ore Composition (Dry Basis) [6] .................................................................... 44

Table 4: Theoretical maximum dissolution of Fe, Ni, and Mg based on the solid loading .......... 45

Table 5: Summary of improved OLI model on Fe(II) and Fe(III) systems .................................. 51

Table A-1: The comparison between Fe(III) measurement from ICP and EDTA titration .......... 67

Table A-2: The comparison between Mg measurement from ICP and EDTA titration ............... 69

Table B-1: Solubility of FeCl3 in water from 25 up to 100°C ...................................................... 72

Table B-2: Kinetic data for FeCl3 solubility in water at 25°C ...................................................... 72

Table B-3: Kinetic data for FeCl3 solubility in 6 molal MgCl2 solutions at 25°C ....................... 73

Table B-4: Kinetic data for FeCl3 solubility in water at 40°C ...................................................... 74




Table B-8: Solubility data of FeCl3 in MgCl2 solutions from 25 up to 100°C ............................. 75

Table B-9: Solubility data of FeCl3 in MgCl2 and HCl solutions from 25 up to 100°C ............... 77

ix

Table B-10: Solubility of MgCl2 in water from 25 up to 100°C .................................................. 78

Table B-11: Kinetic data for MgCl2 solubility in water at 25°C .................................................. 78

Table B-12: Kinetic data for MgCl2 solubility in 0.5 molal FeCl3 solutions at 25°C .................. 79

Table B-13: Solubility data of MgCl2 in FeCl3 solutions from 25 up to 100°C ........................... 79

Table B-14: Kinetic data for hematite solubility in 0.1 molal init. HCl at 60°C (measured) ....... 81

Table B-15: Kinetic data for hematite solubility in 0.1 molal init. HCl at 60°C (mass balance) . 81

Table B-16: Kinetic data for hematite solubility in 2.5 molal MgCl2 and 0.1 molal init. HCl at

60°C (measured) ........................................................................................................................... 82


60°C (mass balance) ..................................................................................................................... 82

Table B-18: Kinetic data for hematite solubility in 4 molal MgCl2 and 0.1 molal init. HCl at

60°C (measured) ........................................................................................................................... 83


60°C (mass balance) ..................................................................................................................... 83

Table B-20: Hematite Solubility in MgCl2 and HCl solutions at 60 and 90°C (measured) ......... 84

Table B-21: Hematite Solubility in MgCl2 and HCl solutions at 60 and 90°C (mass balance) ... 86

Table C-1: XRD peak lists for FeCl3 solubility in 1 molal init. MgCl2 solutions at 60°C ........... 94

Table C-2: XRD peak lists for FeCl3 solubility in 6 molal init. MgCl2 solutions at 60°C ........... 96

Table D-1: Double salt 2.5FeCl3.MgCl2.7.5H2O chemical analysis ........................................... 109

x

Table F-1: Default OLI MSE standard state parameters for solid and aqueous species ............. 134

Table F-2: Regressed standard state parameters of solid and aqueous species .......................... 134

Table F-3: Default OLI MSE mid-range binary interaction parameters between ions/aqueous

species ......................................................................................................................................... 135

Table F-4: Regressed mid-range binary interaction parameters between ions/aqueous species 136

xi

List of Figures

Figure 1: Variation of laterite profile with climate [4] ................................................................... 1

Figure 2: Simplified process flowsheet for saprolite processing with MgCl2 brines [6] .............. 10

Figure 3: MgCl2 phase diagram in H2O under MgO saturation [5] .............................................. 13

Figure 4: Default OLI MSE's simulated solubility of FeCl3 in water ........................................... 14

Figure 5: Default OLI MSE's simulated solubility of FeCl3 in HCl solutions.............................. 15

Figure 6: Default OLI MSE's simulated solubility of hematite in HCl solutions at 100°C .......... 16

Figure 7: Default OLI MSE's simulated solubility of MgCl2 in water ......................................... 17

Figure 8: Default OLI MSE's simulated solubility of MgCl2 in HCl solutions ............................ 18

Figure 9: Default OLI MSE's simulated solubility of hematite in MgCl2 solutions with constant

init. HCl at 60 and 90°C. The lines correspond to default OLI MSE's prediction results on the

solubility of hematite .................................................................................................................... 19

Figure 10: Default OLI MSE's simulated solubility of hematite in HCl solutions with constant

init. MgCl2 at 60 and 90°C. The lines correspond to default OLI MSE's prediction on hematite

solubility ....................................................................................................................................... 20

Figure 11: Experimental setup for low temperature solubility experiment .................................. 22

Figure 12: Experimental setup for high temperature solubility experiment ................................. 23

Figure 13: Reproducibility tests for the solubility of FeCl3 in water ............................................ 26

Figure 14: Kinetics for FeCl3 solubility in MgCl2 solutions at 25°C ........................................... 27

Figure 15: Kinetic tests for FeCl3 solubility in MgCl2 solutions at 40°C ..................................... 28

Figure 16: Solubility of FeCl3.6H2O in MgCl2 solutions at 25 and 35°C..................................... 29

xii

Figure 17: Solubility of 2.5FeCl3.MgCl2.7.5H2O in MgCl2 Solutions from 25 up to 100°C.

Symbols at 0 MgCl2 represent the solubility of FeCl3 hydrates in water. .................................... 30

Figure 18: XRD Spectra of the double salt from 3 molal init. MgCl2 experiment at 60°C .......... 32

Figure 19: Reproducibility tests for the solubility of MgCl2 in Water ......................................... 33

Figure 20: Kinetic Tests for Solubility of MgCl2 in FeCl3 Solutions ........................................... 34

Figure 21: Solubility of MgCl2 in FeCl3 Solutions from 25 up to 100°C ..................................... 35

Figure 22: Reproducibility issues in the solubility of hematite in MgCl2 and HCl solutions ...... 36

Figure 23: Vapor-Liquid Equilibria of HCl in MgCl2 Solutions .................................................. 36

Figure 24: Kinetic tests for hematite solubility in 0.1 molal HCl initially and varying MgCl2

concentration ................................................................................................................................. 38

Figure 25: Solubility of hematite in MgCl2 and HCl solutions at 60 and 90°C ........................... 39

Figure 26: OLI fitting on the solubility of FeCl3.6H2O in MgCl2 solutions ................................. 42

Figure 27: OLI fitting on the solubility of 2.5FeCl3.MgCl2.7.5H2O in MgCl2 solution. The

dashed lines represent the solubility of FeCl3 hydrates at low MgCl2 concentration. .................. 42

Figure 28: OLI fitting on the solubility of MgCl2 in FeCl3 solutions ........................................... 43

Figure 29: OLI prediction on the solubility of FeCl3 in MgCl2 and HCl solutions ...................... 43

Figure 30: Maximum solid loading achievable based on OLI Prediction .................................... 45

Figure 31: OLI's Predicted MgCl2-FeCl3-H2O ternary phase diagram ......................................... 48

Figure 32: Predicted solubility of hematite in HCl solutions ....................................................... 49

Figure 33: Predicted solubility of hematite in MgCl2 and HCl solutions ..................................... 50

Figure 34: Predicted hematite solubility in MgCl2 solutions at const. free HCl ........................... 51

xiii

Figure C-1: XRD pattern for FeCl3 solubility in water at 25°C ................................................... 88

Figure C-2: XRD pattern for Solubility of FeCl3 in 6 molal init. MgCl2 solutions at 25°C ......... 89



Figure C-5: XRD pattern for FeCl3 solubility in 1 molal init. MgCl2 solutions at 40°C .............. 92





Figure C-10: XRD pattern for FeCl3 solubility in 3 molal init. MgCl2 solutions at 80°C ............ 99

Figure C-11: XRD pattern for FeCl3 solubility in 1 molal init. MgCl2 at 100°C ....................... 100

Figure C-12: XRD pattern for FeCl3 solubility in 3 molal init. MgCl2 solutions at 100°C ........ 101

Figure C-13: XRD pattern for MgCl2 solubility in 1.5 molal init. FeCl3 solutions at 25°C ....... 102

Figure C-14: XRD pattern for MgCl2 solubility in 0.5 molal FeCl3 solutions at 80°C .............. 103

Figure C-15: XRD pattern for MgCl2 solubility in 2 molal init. FeCl3 solutions at 80°C .......... 104

Figure C-16: XRD pattern for MgCl2 solubility in 0.5 molal FeCl3 solutions at 100°C ............ 105

Figure C-17: XRD pattern for MgCl2 solubility in 1 molal FeCl3 solutions at 100°C ............... 106

Figure C-18: XRD pattern for hematite solubility in 0.1 molal HCl and 2.5 molal MgCl2 at 60°C

..................................................................................................................................................... 107

Figure C-19: XRD pattern for hematite solubility in 0.1 molal HCl, sat'd MgCl2 solutions at 60°C

..................................................................................................................................................... 108

xiv

Figure D-1: First TGA analysis of the double salt 2.5FeCl3.MgCl2.7.5H2O .............................. 111

Figure D-2: Second TGA analysis of the double salt 2.5FeCl3.MgCl2.7.5H2O ......................... 112

Figure E-1: Default OLI MSE's simulated solubility of FeSO4 in water .................................... 113

Figure E-2: Improved OLI MSE's simulated FeSO4 solubility in water .................................... 114

Figure E-3: Default OLI MSE's simulated FeSO4 solubility in H2SO4 solutions up to 100°C .. 115

Figure E-4: Improved OLI MSE's simulated FeSO4 solubility in H2SO4 solutions up to 100°C 115

Figure E-5: Default OLI MSE`s simulated FeSO4 solubility in H2SO4 solutions above 100°C 116

Figure E-6: Improved OLI MSE`s simulated FeSO4 solubility in H2SO4 solutions above 100°C

..................................................................................................................................................... 116

Figure E-7: Default OLI MSE's simulated FeSO4 solubility in MgSO4 solutions ..................... 117

Figure E-8: Improved OLI MSE's simulated FeSO4 solubility in MgSO4 solutions .................. 118

Figure E-9: Default OLI MSE's simulated FeSO4 solubility in MgSO4 and H2SO4 solutions ... 119

Figure E-10: Improved OLI MSE's simulated FeSO4 solubility in MgSO4 and H2SO4 solutions

..................................................................................................................................................... 119

Figure E-11: Default OLI MSE's simulated FeSO4 solubility in ZnSO4 and H2SO4 solutions .. 120

Figure E-12: Improved OLI MSE's simulated FeSO4 solubility in ZnSO4 and H2SO4 solutions121

Figure E-13: Default OLI MSE's simulated FeSO4 solubility in ZnSO4 and H2SO4 solutions at

200°C .......................................................................................................................................... 121

Figure E-14: Improved OLI MSE's simulated FeSO4 solubility in ZnSO4 and H2SO4 solutions at

200°C .......................................................................................................................................... 122

xv

Figure E-15: Default OLI MSE's simulated FeCl2 solubility in water ....................................... 123

Figure E-16: Default OLI MSE's simulated FeCl2 solubility in HCl solutions .......................... 124

Figure E-17: Improved OLI MSE's simulated FeCl2 solubility in HCl solutions....................... 124

Figure E-18: Default OLI MSE's simulated FeCl2 solubility in MgCl2 solutions ...................... 125

Figure E-19: Improved OLI MSE's simulated FeCl2 solubility in MgCl2 solutions ................... 126

Figure E-20: Default OLI MSE's simulated FeCl2 solubility in MgCl2 and HCl solutions ........ 127

Figure E-21: Improved OLI MSE's simulated FeCl2 solubility in MgCl2 and HCl solutions .... 127

Figure E-22: Default OLI MSE's simulated Fe2O3 solubility in H2SO4 solutions at 130-170°C 128

Figure E-23: Improved OLI MSE's simulated Fe2O3 solubility in H2SO4 solutions at 130-170°C

..................................................................................................................................................... 129

Figure E-24: Default OLI MSE's simulated Fe2O3 solubility in H2SO4 solutions at 230-270°C 129


..................................................................................................................................................... 130

Figure E-26: Default OLI MSE's simulated Fe2O3 solubility in MgSO4 and H2SO4 solutions .. 131

Figure E-27: Improved OLI MSE's simulated Fe2O3 solubility in MgSO4 and H2SO4 solutions

..................................................................................................................................................... 131

Figure E-28: Default OLI MSE's simulated Na-jarosite solubility in H2SO4 solutions ............. 132

Figure E-29: Improved OLI MSE's simulated Na-jarosite solubility in H2SO4 solutions .......... 133

1

Chapter 1

1 Introduction

1.1 Iron chemistry and its importance in hydrometallurgy

Iron is considered as a major impurity in hydrometallurgical processing circuits due to its

mineralogical abundance. At a certain point of the process, and in order to obtain purified pay

metal products, iron needs to be separated most often by precipitation of various iron

compounds, such as hematite (Fe2O3), goethite (FeOOH) or jarosite (XFe3(SO4)2(OH)6, where X

is usually a monovalent cation, e.g. H3O+, NH4

+, Na

+, K

+, Ag

+, 1/2Pb

2+). Hematite is the most

desired iron compound since it is thermodynamically stable, has high Fe content per unit mass,

and is commercially saleable if pure [1]. On the other hand, jarosite precipitation allows sulfate

and alkali metal removal but it has low Fe content per unit mass and it is unstable under certain

storage conditions [2, 3]. For example, in processing nickel from lateritic ores, the

hydrometallurgical treatment varies depending on the types of laterites, whether it is limonitic

laterite or saprolitic laterite. Figure 1 shows the distribution of lateritic ores with depth and

climate.

Figure 1: Variation of laterite profile with climate [4]

2

Limonitic laterites are oxide ores, which are rich in iron and low in magnesium content and are

commonly processed by using pressure acid leach with sulfuric acid. Iron which mainly exists as

goethite in the ore is dissolved by acid attack and precipitated in-situ as hematite, which

regenerates the stoichiometric equivalent of acid as per reactions given below.

2FeOOH (s) + 3H2SO4 (aq) → Fe2(SO4)3 (aq) + 4H2O (l) (1-1)

Fe2(SO4)3 (aq) + 3H2O (l) → Fe2O3 (s) + 3H2SO4 (aq) (1-2)

As a result, acid is only required to bring the value metals (i.e. Ni and Co) into the solution, as

well as the minor impurities, e.g. Ca, and Al. The presence of monovalent cations in the leach

solution promotes jarosite precipitation instead of hematite. While this process is efficient in

processing limonitic laterites, it suffers from high acid consumption when processing saprolitic

laterites. The reason is that saprolitic laterites are rich in Mg, which will be dissolved by acid

attack, thus greatly increasing the acid requirements. This is partially due to the sulfate-bisulfate

equilibrium, i.e. sulfate addition shifts the equilibrium towards bisulfate formation, thus

decreasing the available protons for leaching.

Alternatively, this type of ore can be treated by chloride hydrometallurgy, in particular in the

presence of MgCl2 brines [5]. This process, which is mainly targeted to limit the Mg dissolution,

also improves the acid activity and allows the recovery of dry HCl(g). Furthermore, the use of

chloride brines would elevate the boiling point of the solutions, thus making it possible to run the

process at temperatures higher than 100°C while maintaining atmospheric pressure, which

simplifies the types of equipment used without the need to use pressurized vessels. Moreover,

chloride salts are generally extremely soluble, which could be utilized to minimize the size of the

equipment, thus reducing the capital costs of running the process.

An HCl/MgCl2 solution used as lixiviant will dissolve the oxides in the ore to metal chlorides.

Recent study indicates that Ni and Co dissolution follows Fe dissolution from the oxide phase [6]

and since chloride ion strongly forms complexes with Fe(III), in-situ hematite precipitation is

hindered. Therefore, the idea behind chloride hydrometallurgy is to dissolve all of the Fe to

completely release the Ni and Co to the solution. In this case, the leaching process is not focused

3

on minimizing acid consumption, as compared to the sulfate hydrometallurgy, but to regenerate

all of the chloride units consumed (from HCl) as HCl(g). Based on these findings, the

knowledge of the solubility of Fe(III) under this condition would be beneficial to evaluate the

maximum solid to liquid ratio (solid loading) that can be processed during the leaching step

without losing the chloride units from precipitation of metal chlorides, in this case FeCl3. The

dissolved Fe(III) content is removed as hematite during the neutralization stage by MgO addition

which further increases the MgCl2 concentration in the solution. The knowledge of the solubility

of hematite under concentrated MgCl2 solutions can be utilized to monitor the efficiency of the

iron removal during this process. A detailed description of this process will be discussed in the

next chapter.

Other examples, such as the Akita Zinc Hematite Process, involves neutral leaching of residue in

the form of zinc ferrite (ZnFe2O4) to remove iron as hematite and recover valuable metals (e.g.

Cu, Ga and In) [7, 8]. First, the residue is leached in hot spent H2SO4 solutions, with the addition

of SO2 to reduce Fe(III) to Fe(II) as follows [7]:

ZnFe2O4 (s) + SO2 (l) + 2H2SO4 (aq) → ZnSO4 (aq) + 2FeSO4 (aq) + 2H2O (l) (1-3)

Following neutralization takes place using limestone (CaCO3) to remove valuable metals and

certain impurities from the solution. This reduction of iron allows it to remain in the solution

without undergoing hydrolysis at pH up to 4 [7]. The solution is then subjected to oxidation-

hydrolysis, namely oxydrolysis, inside an autoclave at 200°C, at O2 overpressure of 100-500

kPa, which regenerates H2SO4 and produces hematite as seen in Eq. 1-4 [7, 8].

2FeSO4 (aq) + 0.5O2 (g) + 2H2O (l) → Fe2O3 (s) + 2H2SO4 (aq) (1-4)

However, the solubility of FeSO4 decreases with increasing temperature above 55°C, even more

drastic at temperatures above 100°C [9, 10, 11], thus the knowledge of solubility of FeSO4 in the

presence of acidic ZnSO4 solutions at elevated temperatures is beneficial in avoiding the

potential precipitation of FeSO4.H2O prior to/during oxydrolysis [12]. For the case of steel

pickling process using HCl to remove the FeO scale [13], as well as the non-oxidative/oxidative

chloride leaching process for sulfide minerals using FeCl2+HCl [14] and FeCl3+HCl [15],

respectively, it is important to understand the solubility of FeCl2 in multicomponent HCl

4

solutions at elevated temperatures for figuring out the maximum holding of Fe(II) in that system

to prevent chloride loss for HCl regeneration.

In order to control and optimize such iron removal processes requires the knowledge of the

solubility of iron (II) and (III) in the multicomponent solutions of interest. Development of such

data takes significant amount of experimental work. However, the use of reliable chemical

models can save the industry a considerable amount of time and money.

1.2 OLI Analyzer

The OLI analyzer software package, which was developed by OLI Systems Inc., is capable of

performing simulations of electrolyte solutions at equilibrium. The software features

simultaneous calculation of phase and speciation equilibria along with thermal and volumetric

properties of the mixture. It runs under two main thermodynamic frameworks, either the older

aqueous (AQ) model or the new mixed-solvent electrolyte (MSE) model. The limits of the MSE

model range from below the freezing point to 90% Tcrit of the solution, from 0 to 1500 bar, and

no limit for the ionic strength [16]. The new MSE model works for systems from infinite

dilutions up to the fused salt limit, as well as systems involving mixed-solvent system of organic-

water [16]. For accurate simulations of hydrometallurgical processing solutions, which are

concentrated and possess high degree of non-ideality, the MSE model was selected as the

framework in this study. In addition, model customization was achieved by utilizing the built-in

regression facility with the purpose of either adding new chemical systems or improving the

performance of the current commercial OLI database.

The breakdown of the chapters in this thesis is as follows:

Chapter 2 discusses the fundamentals behind the OLI MSE model, the chemistry behind

the hydrometallurgical treatment of saprolites using MgCl2 brine, and the project

objectives.

Chapter 3 focuses on the experimental methodology used to perform the solubility study.

This chapter includes the experimental setup, the reagents used and the experimental

procedures.

5

Chapter 4 discusses the experimental results for the solubility of FeCl3 in MgCl2, the

solubility of MgCl2 in FeCl3 and the solubility of hematite in MgCl2 and HCl solutions.

Chapter 5 discusses the OLI MSE model improvement through regression of the new

solubility data of FeCl3, MgCl2, and Fe2O3 in chloride systems, and attempts some

extensions relevant to hydrometallurgical applications.

Chapter 6 and Chapter 7 discuss the conclusion of this work and the recommendation

for future work, respectively.

6

Chapter 2

2 Literature Review and Theoretical Background

2.1 OLI MSE Model

2.1.1 Chemical Equilibrium

The equilibrium expression for the general chemical reaction:

aA + bB + … = cC + dD + … (2-1)

is expressed as:

(

) (2-2)

Where R is the ideal gas constant, T is the temperature, xi is the mol fraction of species i, and γi is

the activity coefficient of species i. Meanwhile, the standard-state Gibbs free energy of the

reaction ΔG°, is given by:

∑

(2-3)

which summates all the standard-state chemical potential μi0 of the species i, participating in the

reaction shown in Eq. 2-1 and multiplied by its stoichiometric coefficient, vi with the value

being positive and negative for the species on the product side and reactant side, respectively. In

order to calculate the standard-state chemical potential of the solid species, basic thermodynamic

relationships are used to derive the following equation [17]:

( ) ∫ ∫ (

)

∫

(2-4)

where Δ and

are respectively, the standard partial molal Gibbs free energy of formation

and standard partial molal entropy of the species at 298.15K and 1 bar. P and T correspond to

the pressure and temperature of the system, respectively, while those with the subscript "r"

correspond to the reference conditions which are 298.15K and 1 bar.

7

For aqueous neutral or ionic species, the model needs to predict standard-state properties up to

high temperatures and pressures. In OLI MSE, this is achieved by employing the Helgeson-

Kirkham-Flower-Tanger (HKFT) equation of state. Furthermore, solving for the mole fractions

in Eq. 2-2 it is understood that the activity coefficients are also functions of the mole fractions.

2.1.2 The HKFT Model

The revised HKFT equation of state was developed by Helgeson and co-workers to estimate the

standard state thermodynamic properties of aqueous species [18, 19, 20] at elevated temperatures

and pressures. The expression of the standard partial molal Gibbs Free energy of formation of

any aqueous species ΔG0

P,T is as follows:

( ) ( ) (

)

[ ( ) (

)] (

) [ (

) ]

[{(

) (

)} (

)

(

( )

( ))] (

)

(

) ( ) (2-5)

The species dependant HKFT parameters are a1, a2, a3, a4, c1, c2, and ω. Ψ and θ are the

solvent dependent parameters having the values of 2600 bar and 228K for water, respectively. ε

is the dielectric constant of water and the Born Function,

(

)

while the value of

= -5.8x10-5 K-1

. The database for the HKFT parameters for various aqueous organic and

inorganic species is available in the series of papers published by Helgeson, Shock, and

coworkers [21, 22, 23, 24, 25] or they can be estimated through regression.

2.1.3 Activity Coefficient Model

The activity coefficient is the parameter that describes the degree of non-ideality of the solution.

It is computed based on the excess Gibbs free energy of the solution, Gex

, as follows:

8

( (

)

)

(2-6)

The activity coefficient model incorporated in OLI MSE takes into account the full concentration

range, from a dilute solution which is dominated by long range electrostatic forces, to

concentrated solutions up to the fused salt limit, which is governed by the mid-range ion-ion/ion-

molecule interactions and the short range ion-ion/ion-molecule/molecule-molecule interactions.

As seen in Eq. 2-7, the total excess Gibbs free energy is defined as the sum of the contributions

from all the three ranges [16].

(2-7)

The long-range interaction is described by Pitzer-Debye Hückel expression [26, 27] due to its

effectiveness in reproducing experimental in electrolyte solutions [28, 29, 30, 31], as well as

mixed-solvent electrolyte system [28, 32, 33]. On the other hand, the short-range contributions

are more relevant to mixed-solvent electrolyte solutions which are modeled by using the

UNIQUAC model. The expressions for the long and short-range interactions are available in the

original paper [16]. The concentrated electrolyte solutions of interests to hydrometallurgical

applications are dominated by ionic strength dependent mid-range interactions having the form

of second virial coefficient-type of expression as follows:

(∑ )∑ ∑ ( ) (2-8)

where Bij(Ix) represents the binary interaction parameter between two charged species or a

charged species with neutral molecule with Bij(Ix)=Bji(Ix) and Bii=Bjj=0. The functionality of the

Bij(Ix) is shown in Eqs. 2-9 to 2-11 with b(0-2)ij and c(0-2)ij as the adjustable temperature

independent parameters which are proven to be effective in reproducing experimental data [34].

( ) ( √ ) (2-9)

bij= b0,ij+ b1,ijT + b2,ij/ T (2-10)

cij= c0,ij+ c1,ijT +c2,ij/ T (2-11)

9

2.2 Hydrometallurgical Treatment of Saprolite in Concentrated MgCl2 Brines

2.2.1 Current Technologies to Process Saprolite

A pyrometallurgical process, rotary-kiln electric furnace (RKEF) has been commercially used to

treat saprolite [35]. This process starts with ore-drying inside a rotary kiln at 100°C which

removes the moisture content of the ore. The dried ore moves to another rotary kiln at 700-

900°C which removes the crystalline water by calcination. In this step, coke is generally added

to reduce the nickel and iron oxide back to the metallic state to form ferronickel which is

separated from the iron oxide and silicate slag inside an electric furnace at 1600°C. While the

nickel recovery is above 90% and it can handle high magnesium content, this process requires

high capital cost, and is highly energy intensive [36]. Attempts have been made to process this

type of ore using hydrometallurgy. The high acid requirement due to Mg dissolution and the

unavailability to recover the acid used to dissolve Mg has rendered saprolite not feasible to be

treated using HPAL.

Alternatively, chloride hydrometallurgy to process saprolites has been attempted by Intec [37],

Jaguar [38, 39] and Anglo Base [40] due to the following benefits: high solubility of most metal

chlorides, aggressiveness of HCl to leach the silicate ores even at atmospheric conditions,

minimized scale formation because no calcium sulfate or alunite forms, and the potential to

regenerate HCl. More so, high concentration of chloride increases the activity of the hydrogen

ion [41, 39, 42, 43] (which results in a more efficient leaching using close to stoichiometric HCl

requirements) as well as increase of the solution boiling point to above 100°C which improves

leaching kinetics while maintaining atmospheric conditions. The use of concentrated MgCl2 as

the lixiviant would serve for this purpose as well as limit the dissolution of Mg from the ores

[44]. However, none of the process to treat saprolite has been commercialized, mainly due to

issues with the recovery of HCl. Solution pyrohydrolysis is a conventional method to recover up

to azeotropic HCl (20%), as conceptually implemented by Jaguar Nickel [39, 45] and performed

by the steel pickling industry [46, 47, 48, 49], but it suffers from high energy requirement from

boiling off the excess water and water balance issues of the recycled HCl solutions. In this

process, the MgCl2 brine which is fed to a fluidized bed/spray roaster is contacted with hot

combustion gas at 560-700°C and underwent a hydrolysis reaction to produce MgO and HCl.

Much research attention has been focused towards hydrolytic distillation of HCl which mainly

10

involves hydrolysis of the ferric chlorides at atmospheric pressure and temperatures below

200°C to form superazeotropic HCl (8-9M) and hematite [42, 50, 51, 52]. This process was first

implemented in the 1970s by PORI Inc. In the PORI process HCl is regenerated from the spent

steel pickling liquor by oxydrolysis of a FeCl2 solution [53, 54, 55, 56]. Even though higher

grade of HCl can be achieved as compared to that of pyrohydrolysis, the pre-evaporation step to

concentrate the FeCl2 brine and the use of autoclave during oxidation to FeCl3 rendered this

process economically unattractive. A recent modification by McGill University allows HCl

regeneration from MgCl2-FeCl3 brines but the precipitation of MgCl2.xH2O as well as the MgCl2

hydrolysis product, Mg(OH)Cl.xH2O at 200°C hinders the hydrolysis step [51].

2.2.2 More concentrated MgCl2 solutions

A process idea was recently developed in our group, uses concentrated MgCl2 and HCl as the

lixiviants, and is capable of treating saprolite ore as well as recover dry HCl gas by employing

precipitation, dehydration, and decomposition of magnesium hydroxychloride [5].

Figure 2: Simplified process flowsheet for saprolite processing with MgCl2 brines [6]

First, saprolite ore is subjected to leaching with HCl in MgCl2 brine, where most of the metal

contents of the ore are dissolved to form a metal chloride solution (with minimal dissolution of

magnesium due to common ion effect) while the remaining unleached residue is disposed. The

dissolved iron fraction is removed using MgO as either anhydrous or hydrated iron oxide or iron

oxyhydroxychloride (akageneite). The leaching and the iron control steps will be discussed

further in the next sections. Further MgO addition precipitates base metals as hydroxides. The

remaining MgCl2 brine is precipitated as the double salt, magnesium hydroxychloride by excess

MgO based on the following reactions [5]:

11

2MgO (s) + MgCl2 (aq) + 6H2O (aq) 2MgO.MgCl2.6H2O (s) (2-11)

3MgO (s) + MgCl2 (aq) + 11H2O (aq) 3MgO.MgCl2.11H2O (s) (2-12)

The double salts, 2MgO.MgCl2.6H2O and 3MgO.MgCl2.11H2O are referred as “2-form” and “3-

form”, respectively. While the “2-form” forms at temperature above 90°C and 35 wt% MgCl2,

the “3-form” requires temperature as low as 80°C and 30 wt% MgCl2. This way, the chloride

units are locked into the solid phase, which then is subjected to dehydration to remove the

moisture and crystalline water contents at 230°C and 180°C for the “2-form” and “3-form”,

respectively, as follows [5]:

2MgO.MgCl2.6H2O (s) 2MgO.MgCl2.H2O (s) + 5H2O (g) (2-13)

3MgO.MgCl2.11H2O (s) MgO (s) + 2MgO.MgCl2.H2O (s) + 10H2O (g) (2-14)

The final step is the decomposition of the magnesium hydroxychloride at 560°C to produce dry

HCl gas for recycling and saleable/recycleable MgO product [5].

2MgO.MgCl2.H2O (s) 3MgO (s) + 2HCl (g) (2-15)

The advantages of this process are:

1. Only the water that remains in the magnesium hydroxychloride solid phase needs to be

evaporated in the dehydration step, thus minimizing the energy usage as compared to

solution pyrohydrolysis. The energy required for “2-form” and “3-form” are 10.7

GJ/tonne HCl and 15.9 GJ/tonne HCl, respectively, as compared to 15 wt% and 25 wt%

MgCl2 solution pyrohydrolysis which require 16.5 GJ/tonne HCl and 28.6 GJ/tonne HCl,

respectively [5].

2. The temperature of the decomposition step (560°C) is much lower than that of solution

pyrohydrolysis (700-900°C)

3. The production of anhydrous HCl (g) eliminates water balance issues and produces

water-free HCl.

12

2.2.2.1 Leaching Step

During chloride leaching of saprolite, the following general reaction occurs:

(Fex, Niy, Mg(2-1.5x-y))SiO4 (s) + 4HCl (aq) xFeCl3 (aq) + yNiCl2 (aq) + (2-1.5x-y)MgCl2 (aq)

+ SiO2 (s) + 2H2O (l) (2-16)

The leach solutions consist primarily of FeCl3, MgCl2, NiCl2 and free HCl and leaving behind

unleached magnesium silicate and silica. Leaching tests were studied at MgCl2 concentration of

2 and 4.5 molal, temperature of 100°C and solids loading of 10, 15, and 25% [6]. The

hydrolysis of iron is inhibited due to the high chloride content of the solution, which leads to the

formation of ferric chloride complexes, thus shifting the equilibrium from Fe2O3 towards

FeCl3(s). The question arises at what is the maximum solid loading that can be achieved without

the crystallization of FeCl3. There is no published literature data on the phase diagram of FeCl3-

MgCl2-H2O at elevated temperatures, thus it is of interest to study the solubility of FeCl3 in

concentrated MgCl2 concentration (up to 4.5 molal) with free acid concentration up to 1 molal.

This data serves two purposes: to establish a new phase diagram for the FeCl3-MgCl2 system and

more importantly to regress new interaction parameters between Fe3+

species (including the

ferric chloride complexes) and Mg2+

ion. The resulting model can then be used to predict FeCl3

solubility in the quaternary systems of MgCl2-HCl-H2O.

2.2.2.2 Iron Control

The iron removal step was performed through MgO addition, where the free acid is neutralized at

moderate temperature, and Fe(III) is hydrolyzed as akaganeite (Fe(OH)2.7Cl0.3), which may then

be converted to stable hematite [1]. Higher temperatures and the presence of hematite seeding

were found to promote the formation of hematite [1]. There are published data on hematite and

akaganeite solubility under concentrated MgCl2 solutions (up to 4 molal) at 60 and 90°C [57].

However, as MgO is used to neutralize the solution, more MgCl2(aq) is formed, thus increasing

the MgCl2 brine concentration above the levels encountered during the leaching step. Therefore,

the knowledge of solubility of hematite under more concentrated MgCl2 solutions, i.e., up to the

solubility limit of MgCl2 would be beneficial in order to identify the behavior of Fe(III) in this

particular quaternary FeCl3-MgCl2-HCl-H2O system.

13

2.2.2.3 Effect of FeCl3 on the Phase Diagram of MgCl2-H2O

During leaching, MgCl2 concentration increases due to the dissolution of Mg from the saprolite

ore and second, during iron control due to the addition of MgO. During these two processes,

especially prior to and during Fe(III) precipitation, it is important to prevent the precipitation of

MgCl2.6H2O which contributes to the premature chloride loss. Figure 3 describes the phase

diagram of MgCl2-H2O in the presence of saturated magnesium oxide [5]. The conditions for the

formation of magnesium hydroxychloride including the "2-form" and "3-form" are depicted in

the figure. The operating window for magnesium hydroxychloride precipitation occurs above 30

wt% MgCl2 but below the MgCl2.6H2O solubility limit. There is no published phase diagram of

MgCl2 in the presence of FeCl3. Therefore, it would be of interest to perform solubility

experiments of MgCl2 in the presence of FeCl3 to identify whether FeCl3 addition would shift the

the phase diagram of MgCl2.6H2O.

Figure 3: MgCl2 phase diagram in H2O under MgO saturation [5]

14

2.3 Default OLI MSE's Performance in Modeling FeCl3 and MgCl2 Systems

2.3.1 Solubility of FeCl3 in water

The solubility data for FeCl3 in water at elevated temperatures is provided by the literature [9].

Figure 4 describes the solubility of FeCl3 in water at elevated temperatures, which in general

increases with increasing temperatures, except when eutectic behavior occurs for hydrated salts,

such as FeCl3.6H2O, FeCl3.3.5H2O, or FeCl3. 2H2O. In this case, it is possible to have more than

one solubility value at the same temperature, as well as different stable solid phase, which

depends on the starting concentration of the FeCl3. While the lowest solubility value can be

achieved when equilibrium is approached from the dissolution pathway, the higher solubility

values can only be achieved from the precipitation pathway. The OLI MSE simulation shows

consistency with these literature data.

Figure 4: Default OLI MSE's simulated solubility of FeCl3 in water

15

2.3.2 Solubility of FeCl3 in HCl solutions

OLI MSE simulations were carried out to validate the solubility data obtained from literature [9],

as seen in Figure 5. While higher temperature increases the solubility of FeCl3, higher HCl

concentration is a bit complex. While the default model shows the correct solubility trends, it

suffers from lack of accuracy, especially for the 25°C, FeCl3.6H2O and FeCl3.3.5H2O solid

phases. Nevertheless, the default model provides a good estimation of the FeCl3 solubility in

HCl solutions.

Figure 5: Default OLI MSE's simulated solubility of FeCl3 in HCl solutions

2.3.3 Solubility of Hematite in HCl solutions

The solubility data, which was obtained from literature [1] was converted from molarity to

molality using OLI’s predicted density of the same solution compositions at room temperature.

It was then compared with OLI MSE default model, as seen in Figure 6. Solubility of hematite

increases with increasing HCl concentration. The model showed good agreement at low HCl

concentration (up to 0.3 molal) whereas it overestimated the solubility at higher acid level.

16

Figure 6: Default OLI MSE's simulated solubility of hematite in HCl solutions at 100°C

2.3.4 Solubility of MgCl2 in Water

The default OLI MSE model on the solubility of MgCl2 in water was compared with literature

data obtained from Linke and Seidell [9]. OLI is consistent with the literature data for up to

200°C, as seen in Figure 7. Solubility of MgCl2 increases with increasing temperatures. Higher

temperatures favor the dehydration of the equilibrating solid phase from MgCl2.6H2O to

MgCl2.4H2O and finally MgCl2.2H2O.

17

Figure 7: Default OLI MSE's simulated solubility of MgCl2 in water

2.3.5 Solubility of MgCl2 in HCl Solutions

The literature solubility data of MgCl2 in HCl solutions taken from [9] was used to validate the

default OLI MSE model. It can be seen from Figure 8 that the solubility of MgCl2 increases with

increasing temperature but decreases with increasing HCl concentration. The model is in a good

agreement with literature data but OLI overestimates the solubility data at 70°C.

18

Figure 8: Default OLI MSE's simulated solubility of MgCl2 in HCl solutions

2.3.6 Solubility of Hematite in MgCl2 and HCl Solutions

The solubility data was taken from Konigsberger et al. [57]. It was shown in Figure 9 that OLI

default model predicted incorrect solubility trends, i.e. the solubility decreases when MgCl2 is

added as supposed to increase based on the Fe(III) chloride complex formation. The effect of

HCl addition (Figure 10) and temperature (Figure 9 and Figure 10) on the hematite solubility are

captured by the model, but the predicted values are still inaccurate. This identifies a missing

gap in the OLI database.

19

Figure 9: Default OLI MSE's simulated solubility of hematite in MgCl2 solutions with constant

init. HCl at 60 and 90°C. The lines correspond to default OLI MSE's prediction results on the

solubility of hematite

20

Figure 10: Default OLI MSE's simulated solubility of hematite in HCl solutions with constant

init. MgCl2 at 60 and 90°C. The lines correspond to default OLI MSE's prediction on hematite

solubility

2.4 Objectives

The overall objective of this thesis was to optimize and expand the commercial database of iron

(II) and (III) to the conditions of interests to hydrometallurgy, especially from room temperature

up to 100°C and from 0 up to 3 molal of acid (either HCl or H2SO4). The specific objectives

were:

1. Perform solubility experiments on FeCl3-MgCl2-H2O and FeCl3-MgCl2-HCl-H2O systems

from 25°C up to 100°C in order to determine the maximum concentration of Fe(III) during

leaching of saprolitic ore under concentrated magnesium chloride solution

2. Peform solubility experiments on MgCl2-FeCl3-H2O systems from 25 up to 100°C in order to

determine the effect of FeCl3 on the MgCl2 phase diagram, in order to identify an operating

window for magnesium hydroxychloride precipitation.

21

3. Perform solubility experiments of Fe2O3-MgCl2-HCl-H2O systems at 60 and 90°C for

determining the limit of iron removal during iron precipitation step at concentrated MgCl2

solutions through neutralization by MgO addition.

4. Model the new Fe(III) and Mg(II) solubility data on the chloride system through regression

and utilize the developed model to simulate the leaching and iron removal steps.

On the other hand, the secondary objectives of this thesis were:

1. Validate the OLI models on the solubility of Fe(II) and Fe(III) in sulfate or chloride systems

in the presence of acid and/or other metal sulfates/chlorides and compare with literature data.

2. Improve the chemical model by regression of select Fe(II) and Fe(III) systems relevant to

hydrometallurgy

22

Chapter 3

3 Experimental

3.1 Experimental Setup

Low temperature (up to 80°C) FeCl3 solubility in MgCl2 experiments were conducted in

Erlenmeyer flasks immersed in water bath inside an acrylic tank maintained at constant

temperature using Cole Palmer Polystat Digital Immersion Circulator and the solutions were

stirred continuously using magnetic stirrers (See Figure 3-1). The flasks were capped using

rubber stopper equipped with glass thermometer and the vacuum grease was applied to seal the

joints. In addition, parafilm was wrapped around the outer joints of the flasks to enhance the

seal. In order to minimize the evaporation of water from the bath, hollow plastic balls were

added to cover the water surface. For later experiments involving the solubility of MgCl2 in

FeCl3 solutions, as well as the solubility of hematite in acidic MgCl2 solutions, the bath was

modified by replacing the material of the tank with polypropylene, as well as the water bath with

a mixture of polypropylene glycol (PG): water = 70:30 vol% enabling the setup to be used up to

100°C.

Water Bath

60°C

Stirrer

Thermometer

Erlenmeyer

Flask

Magnetic

Stirrer Bar

Thermostat

Rubber

Stopper

Figure 11: Experimental setup for low temperature solubility experiment

23

On the other hand, high temperature (100°C) FeCl3 solubility in MgCl2 experiments were

conducted in 1L jacketed glass reactor heated with oil bath maintained at the reaction

temperature using a VWR temperature-controlled circulator (See Figure 3-2). The stirring was

provided using a motor with the stirrer made of PTFE and the shaft made of glass. A glass

thermometer was attached to the reactor using one of the available top ports to monitor the

temperature. Vacuum grease was applied to seal all of the joints.

Motor

Thermometer

Jacketed

Glass

Reactor

Oil Inlet From

Heater

Oil Outlet to

Heater

Stirrer

Liquid

Sampling Port

Solid Sampling

Port

Figure 12: Experimental setup for high temperature solubility experiment

3.2 Reagents

All the reagents used, i.e. FeCl3.6H2O(s), FeCl3(s), MgCl2.6H2O(s), MgCl2(s), concentrated HCl and

Fe2O3(s) were ACS grade from various chemical suppliers, including Fischer Scientific, Sigma-

Aldrich and Alfa Aesar. The water used for the experiments was Millipore-Q deionized water.

24

3.3 Experimental Procedures

3.3.1 FeCl3 Solubility in MgCl2 Solutions

The solubility experiments were approached from the dissolution pathway. Since the solubility

of FeCl3 increases with temperature, the following procedures were implemented. For

experiments below the melting point of FeCl3.6H2O at 37°C, solutions of various MgCl2

concentrations were prepared using either MgCl2.6H2O or anhydrous MgCl2 and excess

FeCl3.6H2O was added (above its saturation level). On the other hand, for experiments above

37°C, solids containing a predetermined amount of MgCl2.6H2O, FeCl3.6H2O and excess FeCl3

were added to the flasks at ambient temperature and the temperature was raised to the reaction

temperature. This method was selected to control the heat released from the exothermic

dissolution reaction of the anhydrous FeCl3 by melting the FeCl3.6H2O. It prevented the

solutions to boil up which would otherwise cause the experiment to be approached from

supersaturation.

All the equipment used for sampling were kept inside an oven heated slightly above the reaction

temperature to prevent precipitation during sampling. Kinetics studies were performed to

estimate the time required to reach equilibrium. 3 Samples were taken at equilibrium over a

period of 1-2 hours using dip tubes and immediately filtered using a 0.45μm syringe filter. The

supernatant solution was placed in a 1 mL volumetric flask for density measurement at

temperature by measuring its weight using an analytical balance. The resulting solution was

diluted 10-fold using 5% HCl. For experiments involving HCl, the solution was diluted with

deionized water for free acid measurement by an acid-base titration method (See Appendix A-3).

The Fe and Mg concentration were analyzed by using a complexometric titration method (See

Appendix A-1 and A-2, respectively) instead of by ICP-AES due to the error associated with

high dilution from highly concentrated solutions. Solid samples were taken using a dip tube,

filtered using a vacuum filtration apparatus and kept inside vacuum desiccators for phase

analysis by powder X-Ray Diffraction (XRD).

3.3.2 MgCl2 Solubility in FeCl3 Solutions

The solubility experiments were conducted from the dissolution pathways. Since the solubility

of MgCl2 increases with increasing temperature, the experiments were started at 25°C and the

25

temperature was progressively raised up to 100°C. Solutions of various concentration of FeCl3

were prepared and put inside the Erlenmeyer flasks which were kept at 25°C using the water

bath. Following thermal equilibration, excess MgCl2.6H2O was added to the flask marking the

start of the experiment. Initial kinetics was measured to estimate the equilibration time. After

equilibrium was achieved, 2 liquid samples were obtained using a plastic syringe equipped with

a 0.45μm syringe filter and placed inside a 1 mL volumetric flask for density measurement. The

solution was then diluted 10 fold (the first sample in 0.1M HCl solution and the second sample in

DI water to check whether hydrolysis reaction occurs or not). The samples were immediately

titrated for Fe, Mg, and HCl determination using the same method mentioned previously. Solid

samples were obtained as previously mentioned. After sampling, the temperature of the water

bath was raised to the next reaction temperature while ensuring that excess MgCl2.6H2O was

maintained in each flask by visual observation.

3.3.3 Hematite Solubility in Acidic MgCl2 Solutions

The solubility experiments were performed from the dissolution pathway. Solutions containing

MgCl2 and HCl of various concentrations were prepared and initial samples were taken to

confirm the initial solution composition. The solutions were then heated to the reaction

temperature. After thermal equilibrium was achieved, excess Fe2O3 was added to the flasks to

start the solubility experiments. Kinetic tests were performed to estimate the equilibration time.

After equilibrium was achieved, 2 liquid samples were obtained by using dip tubes and

immediately filtered using 0.45 μm syringe filters. The non-diluted solutions remained stable due

to increased solubility with decreasing temperature. For experiments involving saturated MgCl2

solutions, since MgCl2 solubility decreases with decreasing temperature, the filtered solution was

transferred to a 2 mL volumetric flask for density measurement at temperature and immediately

diluted 2.5-fold using deionized water to prevent precipitation of MgCl2.6H2O. Solid samples

were obtained by using dip tubes, and filtered immediately using vacuum filtration apparatus and

stored inside vacuum desiccators. Metals and free acid concentrations were analyzed using the

same methods described in the previous section.

26

Chapter 4

4 Experimental Results and Discussion

4.1 Solubility of FeCl3 in MgCl2 Solutions

4.1.1 Reproducibility Tests for FeCl3 Solubility in Water

Initial tests were performed to reproduce literature data for solubility of FeCl3 in water [9]. As

seen in Figure 13, they show close agreement with the literature up to 60°C and they

underestimate the solubility at 80 and 100°C within 10%.

Figure 13: Reproducibility tests for the solubility of FeCl3 in water

4.1.2 Kinetic Tests for FeCl3 Solubility in MgCl2 solutions

Kinetics were measured at 25°C (see Figure 14) and 40°C (see Figure 15) in order to estimate

the time required for the system to reach equilibrium. At 25°C, the stable solid phase is

FeCl3.6H2O and there is no phase transformation that occurs under the conditions studied.

However, at 40°C, the stable solid phase may differ depending on the conditions, i.e. for no

MgCl2, the stable solid phase is FeCl3.2.5H2O, whereas in the presence of MgCl2, the stable solid

27

phase is the double salt 2.5FeCl3.MgCl2.7.5H2O, which will be described in the later section. It

turns out that the kinetics at 40°C is faster than the kinetics of the phase transformation to the

double salt, indicating that 12 hours is sufficient to reach equilibrium.

Figure 14: Kinetics for FeCl3 solubility in MgCl2 solutions at 25°C

28

Figure 15: Kinetic tests for FeCl3 solubility in MgCl2 solutions at 40°C

4.1.3 FeCl3 Solubility Data in MgCl2 Solutions

The solubility data for FeCl3 in MgCl2 solutions is shown in Figure 16 and Figure 17. It can be

seen that at 25°C the solubility first decreases with increasing MgCl2 concentration due to the

common ion effect; the solubility starts to increase at about 2.5 molal MgCl2 which is due to the

effect of complex formation between Fe3+

and Cl- ions. This complexation effect also explains

the increasing solubility trend at 35°C. Under these conditions, when the equilibrium is

approached from the dissolution pathway, FeCl3.6H2O is the stable solid phase. It is also shown

that the solubility increases with increasing temperature up to 100 °C. At 40°C and above, the

solubility decreases with increasing MgCl2 concentration whereas the stable solid phase under all

conditions studied is the double salt, 2.5FeCl3.MgCl2.7.5H2O. One set of the experiments was

conducted by cooling the equilibrated solutions at 40 °C down to 25 °C. It turns out that the

solubility trends behave similarly with those at above 40 °C and that the double salt remains

stable at 25 °C. The two different solubility values at 25 °C for the same MgCl2 concentrations,

one in equilibrium with the hexahydrate while the other with the double salt, can be traced back

to the nature of the solubility of FeCl3 in water, as seen in Figure 13. It was clear that the phase

29

diagram for FeCl3 in water contains multiple eutectic points, which make it possible to have two

distinct solubility values for the same temperature and solid phase. Thus, it is likely that the

higher solubility values are achieved only when the equilibrium is approached from precipitation

resulting from supersaturated solutions upon cooling from 40 °C back to 25 °C.

Figure 16: Solubility of FeCl3.6H2O in MgCl2 solutions at 25 and 35°C

30

Figure 17: Solubility of 2.5FeCl3.MgCl2.7.5H2O in MgCl2 Solutions from 25 up to 100°C.

Symbols at 0 MgCl2 represent the solubility of FeCl3 hydrates in water.

The XRD spectra for the equilibrating solid phase are included in the Appendix C-1. One

particular feature of the double salt is the unique XRD spectrum which does not correspond to

any FeCl3 solid compound, or any MgCl2 solid compound, as seen in Figure 18. As a

consequence, the identity of the unknown double salt could not be determined by just performing

XRD analysis. In order to resolve this issue, chemical analysis was performed by dissolving the

solid sample in 5% HNO3 and then determining the stoichiometry of the solid phase as follows:

- Metal analysis (Mg, Fe): ICP OES, or complexometric titration

- Chloride analysis: Chloride Ion Selective Electrode (ISE), or gravimetric analysis

- Crystalline water analysis: mass balance. In order to find out whether the only anion is

chloride ion or there is any other anion, e.g. O2-

, OH-, it is simply checked by performing

electroneutrality of the metal cations and chloride anion. If it is electroneutral, all the metals

exist in the form of metal chloride, which is indeed the case. The crystalline water content is

finally determined from mass balance. There are several factors that affect the accuracy of the

measurements:

31

1. The nature of the double salt, a chloride salt which is extremely hygroscopic, thus

absorbing moisture during sample preparation for the analysis and affecting the

crystalline water content of the solid

2. Solid was not washed during filtration due to extreme solubility with any solvents tested,

including water, alcohols, and acetone, which could lead to absorption of the residual

solution on the surface of the double salt, thus affecting the Fe:Mg stoichiometry to vary

between 2.4 up to 3

The complete chemical analysis can be found in Appendix D. The best analysis was performed

by using complexometric titration for Fe and Mg determination, coupled with AgCl gravimetric

analysis for Cl determination which gives the lowest error for the electroneutrality, under 1%.

Finally, the proposed final formula of the double salt is 2.5FeCl3.MgCl2.7.5H2O.

32

Figure 18: XRD Spectra of the double salt from 3 molal init. MgCl2 experiment at 60°C

There were additional attempts to confirm the water content by conducting TGA analysis, as

shown in Appendix D, by measuring the weight difference of crystalline water evaporation, but

this measurement was not clear due to the existence of multiple unidentified peaks, as well as the

high amount of absorbed (not crystalline) water due to the hygroscopic nature of the sample

which blurred the calculation of the crystalline water content based on the weight loss.

4.2 Solubility of MgCl2 in FeCl3 Solutions

4.2.1 Reproducibility Tests for MgCl2 Solubility in Water

The reproducibility check with literature data showed that the results of this work are in an

excellent agreement with the literature solubility data, with errors below 2%, as shown in Figure

19. It clearly indicates that the proposed sampling procedures, as well as the complexometric

titration method for Mg determination work with a high degree of accuracy.

33

Figure 19: Reproducibility tests for the solubility of MgCl2 in Water

4.2.2 Kinetic Tests for MgCl2 Solubility in FeCl3 Solutions

The kinetics of MgCl2.6H2O dissolution in water and in 0.5 molal FeCl3 even at 25°C is very fast

with equilibrium attained in less than 4 hours, as seen in Figure 20. However, for each

experiment, even at higher FeCl3 concentrations, up to 2 molal, 12 hours of holding time was

chosen to ensure that the solutions have reached equilibrium.

34

Figure 20: Kinetic Tests for Solubility of MgCl2 in FeCl3 Solutions

4.2.3 MgCl2 Solubility in FeCl3 Solutions

The solubility of MgCl2 decreases with increasing FeCl3 concentration which is due to the

common ion effect and increases with increasing temperature, as seen in Figure 21. However,

the decrease of the solubility diminishes with increasing temperature. At 100°C, the effect of

FeCl3 addition has no significant impact (just under 5%) on the decrease of solubility of MgCl2.

The solid phase was identified by XRD as MgCl2.6H2O; details are included in Appendix C-2.

35

Figure 21: Solubility of MgCl2 in FeCl3 Solutions from 25 up to 100°C

4.3 Solubility of Hematite in MgCl2 and HCl Solutions

4.3.1 Reproducibility of Hematite Solubility in MgCl2 and HCl Solutions

Tests were also performed to check if we could reproduce hematite solubility data in MgCl2 and

HCl as measured by Konigsberger [57]. However, the results were inconsistent due to HCl loss

to the vapor phase. HCl loss affects the mass balance calculation as well as the extent of

hematite dissolution. It becomes worse when the concentration of the chloride brines and the

temperature is high as seen by the difference between the measured vs. calculated lines in Figure

4-9. Vapor-Liquid Equilibria of HCl in MgCl2 solutions, as seen in Figure 23, has been studied

in the literature [58] and it has been shown that chloride addition promotes the formation of

molecular HCl which is in equilibrium with the HCl vapor based on the following reaction:

H+ (aq) + Cl

- (aq) → HCl (aq) → HCl (g) (4-11)

which makes it possible to distill HCl(g) using very concentrated chloride brines.

36

Figure 22: Reproducibility issues in the solubility of hematite in MgCl2 and HCl solutions

Figure 23: Vapor-Liquid Equilibria of HCl in MgCl2 Solutions

37

The main sources of error come from:

1. HCl losses during initial sampling at temperature and prior to adding the hematite

powder,

2. HCl losses during sampling (kinetic or equilibrium), and

3. High solubility of hematite in MgCl2 brines which leaves a very small residual free

HCl at equilibrium.

The above cause the measured initial HCl to be overestimated, the equilibrium HCl to be

underestimated, and the free acid measurement by titration to be less reliable. Konigsberger et

al., in performing their hematite solubility measurements used a recycled plug flow reactor

(PFR) with hematite placed in a fixed bed [59], which in theory prevents supersaturation due to

excessive stirring. They successfully closed the mass balance by using a pH electrode specially

calibrated with respect to the concentration (as opposed to activity) of H+ and at constant MgCl2

concentration [57]. This sophisticated continuous measurement of the free acid cannot be

reproduced inside a batch reactor due to the need to separate the supernatant from the solid to

prevent buildup of hematite powder on the tip of the pH probe.

4.3.2 Kinetic Tests for Hematite Solubility in MgCl2 and HCl Solutions

Figure 24 shows that the kinetics of hematite dissolution reaches equilibrium in 24 hours in the

absence of MgCl2. On the other hand, the kinetics is faster with increasing MgCl2 concentration

due to an increase of the activity of proton in these water-deprived systems.

38

Figure 24: Kinetic tests for hematite solubility in 0.1 molal HCl initially and varying MgCl2

concentration

4.3.3 Hematite Solubility Data in MgCl2 and HCl Solutions

Despite the inability of the current measurements to obtain solubility data at prescribed acid

concentration levels due to HCl losses, several trends can be observed. The new hematite

solubility in MgCl2 and HCl solutions, as seen in Figure 25 shows that MgCl2 addition increases

the solubility of hematite, as previously confirmed by Konigsberger [57] up to 4 molal and then

decreases as the MgCl2 concentration approaches its saturation level. The increase of the

solubility is due to ferric chloride complex formation and the decrease is due to salting out effect,

at which there is no more available water to hydrate the Fe3+

ions in the solution. This salting

out effect seems to have a more pronounced effect at higher HCl concentration. The solubility of

hematite also decreases with increasing temperature, which is also consistent with the

Konigsberger's observation [57]. The XRD spectra for the equilibrating solid phases are

included in Appendix C-3.

39

Figure 25: Solubility of hematite in MgCl2 and HCl solutions at 60 and 90°C

40

Chapter 5

5 OLI Modeling Results and Discussion

5.1 Fitting on the FeCl3-MgCl2-H2O System

The default OLI model has successfully simulated the binary and ternary chloride systems of

interests: FeCl3-H2O, MgCl2-H2O, FeCl3-HCl-H2O, MgCl2-HCl-H2O, but not for the system

FeCl3-MgCl2, as seen previously in Chapter 2.3. The new solubility data of FeCl3 in MgCl2

solutions, as well as the MgCl2 in FeCl3 solutions, was fitted on the OLI MSE model and several

parameters were regressed. First, the basic thermodynamic properties of the new double salt,

2.5FeCl3.MgCl2.7.5H2O was regressed, i.e. ΔG°f, S° and heat capacity of the solid, Cp, since this

solid compound is new and there is no published thermodynamic parameters of the salt. Second,

the mid-range binary interaction parameters between major species in the solution, FeCl2+

- Mg2+

and FeCl2+

- Mg2+

were regressed as well. The regressed parameters and the comparison

between the thermodynamic properties of the new double salt and the 3 form of the magnesium

hydroxychloride are shown in Table 1 and Table 2, respectively. It was clear that the regressed

thermodynamic parameters of the new double salt is in the same order of magnitude with the “3-

form” of the magnesium hydroxychloride.

Table 1: Regressed OLI parameters for FeCl3-MgCl2-H2O system

Species Standard State Properties BMD CMD

2.5FeCl3.MgCl2.

7.5H2O(s)

ΔG°f = -3.379986 MJ/mol

S° = 0.64759 kJ/mol.K

ΔH°f = -4.054028 MJ/mol

Cp = 52.189 J/mol.K

- -

FeCl2+

- Mg2+

- BMD0 = 218.3465 CMD0 = -472.2865

41

BMD1 = -0.545778

BMD2 = 16731.82

CMD1 = 0.9531058

CMD2 = 3236.161

FeCl2+

- Mg2+

- BMD0 = 1333.926

BMD1 = -2.607153

BMD2 = -67653.59

CMD0 = -581.7210

CMD1 = 1.652331

CMD2 = -138093.5

Table 2: Comparison of the thermodynamic properties of the double salts

Solid species ΔG°f (kJ/mol) S° (J/mol.K) Cp (kJ/mol.K)

2.5FeCl3.MgCl2.7.5H2O -3379.86 647.59 52.189

3MgO.MgCl2.11H2O [5] -6077.45 643 -

It could be seen in Figure 26, Figure 27 and Figure 28 that the OLI model is consistent with both

systems FeCl3-MgCl2-H2O and MgCl2-FeCl3-H2O with AARD less than 5%. The fitted model

was then validated using the new solubility of FeCl3 in the presence of MgCl2 and HCl solutions

within around 3% error, as presented in Figure 29.

42

Figure 26: OLI fitting on the solubility of FeCl3.6H2O in MgCl2 solutions

Figure 27: OLI fitting on the solubility of 2.5FeCl3.MgCl2.7.5H2O in MgCl2 solution. The

dashed lines represent the solubility of FeCl3 hydrates at low MgCl2 concentration.

43

Figure 28: OLI fitting on the solubility of MgCl2 in FeCl3 solutions

Figure 29: OLI prediction on the solubility of FeCl3 in MgCl2 and HCl solutions

44

5.2 Maximum Achievable Solid Loading during Leaching

The high solubility of FeCl3 in MgCl2 brines enables processing of saprolite ores with minimum

amount of solution, thus potentially reducing the overall equipment sizes. A previous

experimental study on leaching of saprolite under MgCl2 brines was conducted up to 25% solid

loading by using HCl gas instead of concentrated HCl(aq) [6]. Mined saprolite comes often in

the form of a slurry with about 30% solids. Since Ni dissolution follows Fe dissolution from

saprolite, it is expected that in order to achieve near complete (>99%) extraction of Ni, all Fe

should be dissolved. Due to the high chloride strength of the brine, Fe in-situ hydrolytic

precipitation as hematite is hindered. Therefore, estimation of the maximum achievable solid

loading is based on how much Fe can be dissolved without precipitation of any chloride salt,

mainly as FeCl3. Calculation of solids loading was performed using the ore composition as in

Table 5-2:

Table 3: Saprolite Ore Composition (Dry Basis) [6]

Metal Ni Co Fe Mg Al Mn Cr

Wt % 3.89 0.07 25.47 3.12 0.60 0.44 0.41

The MgCl2 brine strength of 2 and 4.5 molal was chosen as the constituent of the liquid phase

during calculation of the solid loading which represents the brine strength used in the previous

study during leaching step [6]. The maximum Fe dissolution data from Table 4 was used to

construct the dissolved iron curve in Figure 30. This information, combined with the OLI model

prediction of FeCl3 solubility in 2 or 4.5 molal MgCl2 and 0 up to 1 molal HCl solutions at

100°C would serve as the way to estimate the maximum solid loading, i.e. to have the solid

loading curve just below the solubility limit of the double salt. For 2 and 4.5 molal MgCl2, the

maximum solid loading was estimated to be around 75% and 60%, respectively.

45

Table 4: Theoretical maximum dissolution of Fe, Ni, and Mg based on the solid loading

MgCl2 (molal) Solid Loading

(%)

Maximum Fe

(molal)

Maximum Ni

(molal)

Maximum Mg

(molal)

2 35 2.92 0.42 0.82

2 50 5.43 0.79 1.53

2 75 16.29 2.37 4.58

4.5 35 3.51 0.51 0.99

4.5 50 6.51 0.95 1.83

4.5 60 9.77 1.42 2.75

Figure 30: Maximum solid loading achievable based on OLI Prediction

46

However, in practice, the selection of maximum solid loading is much lower than the theoretical

maximum solid loading because the process is bound to the MgCl2 phase diagram. As MgO is

used as neutralizing agent, more MgCl2 is introduced to the solution, thus further augmenting the

brine concentration. The limit of how much MgCl2 can be dissolved without losing any chloride

unit depends on the knowledge of the phase diagram of MgCl2. In this case, there are several

factors that influence the increase of the MgCl2 concentration, some of which are controllable

and the rest are not, as follows:

- Selection of the starting MgCl2 brine concentration prior to leaching step (controllable)

- Mg dissolution from the saprolite ores (uncontrollable)

- Free acid neutralization reaction: MgO (s) + 2HCl (aq) = MgCl2 (aq) + H2O (l) (controllable)

- Iron precipitation reaction: 3MgO (s) + 2FeCl3 (aq) = 3MgCl2 (aq) + Fe2O3 (s) (uncontrollable)

- Nickel precipitation reaction: MgO (s) + NiCl2 (aq) = MgCl2 (aq) + NiO (s) (uncontrollable)

The biggest factor that influences the increase of MgCl2 concentration in the solution is the iron

precipitation reaction which has 3:2 stoichiometry with respect to MgCl2:FeCl3. Since the

solution is dominated by iron, according to the stoichiometry, precipitating all iron from the

solution significantly increases the MgCl2 concentration. Starting the leaching step with 4.5

molal MgCl2 brine would make it difficult to operate with high solid loading because of the

solubility limit of MgCl2 in water, around 7.7 molal at 100°C. As a result, it would be more

reasonable to operate at lower starting MgCl2 at around 2 molal (or possibly lower) and then to

neutralize the solution to precipitate the iron and nickel before reaching the solubility limit of

MgCl2. With all of these factors taken into consideration, the maximum operable solid loading

decreases to around 35% when leaching is performed at 2 molal MgCl2 brine at 100°C. Below is

the summary of the MgCl2 increase from leaching up to the end of Ni precipitation:

- Starting MgCl2 brine concentration = 2 molal

- Mg dissolution (assuming 50% extraction) = 0.41 molal

- Free acid neutralization (assuming 0.5 molal free acid) = 0.25 molal

47

- Iron removal through MgO neutralization = 4.38 molal

- Ni removal through MgO neutralization = 0.42 molal

- Total = 7.46 molal

The resulting MgCl2 concentration is still just below the solubility limit of MgCl2 in water at

100°C and meets the requirement of at least 30 wt% MgCl2 concentration (4.5 molal) for

magnesium hydroxychloride precipitation. This process could be further improved by lowering

the initial MgCl2 brine concentration below 2 molal to accomodate for possible higher Mg

extraction from the ore as well as lower free acid concentration. Nevertheless, solid loading of

35% is possible to be processed under the proposed saprolite processing circuit.

5.3 Effect of FeCl3 Addition on MgCl2 Phase Diagram

Addition of FeCl3 decreases the solubility of MgCl2 due to common ion effect and it can be seen

from Figure 31. The decrease in the MgCl2 solubility is smaller at higher temperature. However,

the boiling point of the solution keeps rising with additional FeCl3. This could potentially affect

the operating window for magnesium hydroxychloride precipitation. These two contradicting

effects seem to have a neutral effect on the process operating window. In the proposed process,

prior to Fe removal, the system is expected to have high concentration of dissolved FeCl3 but as

more FeCl3 is precipitated through MgO addition, the phase diagram would eventually shift back

to the MgCl2-H2O phase diagram. At any point of time during Fe removal, the MgCl2

concentration should never fall above the solubility limit of MgCl2 in FeCl3 solution and this

could be prevented by good selection of process conditions as outlined in the previous section.

48

Figure 31: OLI's Predicted MgCl2-FeCl3-H2O ternary phase diagram

5.4 Limitation of The OLI Model

While the developed OLI model for the system FeCl3-MgCl2-HCl-H2O performs well, it suffers

from a lack of accuracy when used in predicting the solubility of hematite in HCl, as well in the

presence of MgCl2. As seen in Figure 32, the solubility of hematite increases with increasing

free HCl concentration and decreases with increasing temperature. When compared with the

solubility data obtained from Dutrizac and Riveros, their solubility data seems to be higher than

that of obtained from this work at 60 and 90°C.

49

Figure 32: Predicted solubility of hematite in HCl solutions

Their overestimated data is mainly due to their equilibrium approach from hydrolytic

precipitation of FeCl3 instead of hematite dissolution. The developed OLI model predicted the

correct trends when compared with the results of this work for the system Fe2O3-HCl. When the

model is used to predict hematite solubility in MgCl2 and HCl solutions, it underestimated the

solubility data by around 1 log unit, as seen in Figure 33. Nevertheless, the model still predicted

the correct trends for MgCl2 addition to the hematite solubility, i.e. the solubility increases with

MgCl2 addition due to ferric chloride complex formation, as clearly seen in Figure 33. This

limitation may be due to the absence of higher order complexes, FeCl3(aq) and FeCl4- in the

default OLI database. The default FeCl3-H2O and FeCl3-HCl-H2O was built using only the first

two ferric chloride complexes, i.e. FeCl2+ and FeCl2

+ and it manages to performs well in the

system FeCl3-MgCl2-H2O as well as FeCl3-MgCl2-HCl-H2O. However, when the higher order

complexes are introduced to the OLI database, it would disrupt the performance of the all the

developed FeCl3 models and as a result, refitting of the whole FeCl3 system starting from the

binary FeCl3-H2O may be required. Furthermore, regressing more interaction parameters requires

more data points to prevent over-parameterization. This approach, which may further improve

the overall performance of the model, may be the subject for a future study.

50

Figure 33: Predicted solubility of hematite in MgCl2 and HCl solutions

51

Figure 34: Predicted hematite solubility in MgCl2 solutions at const. free HCl

5.5 OLI Model Improvement for Other Fe(II) and Fe(III) Systems

OLI model improvement was performed on select Fe(II) and Fe(III) systems, relevant in

hydrometallurgy which is included in Appendix E. A summary of all the improvements is as

follows:

Table 5: Summary of improved OLI model on Fe(II) and Fe(III) systems

Type of System Temperature Range (°C) AARD (%)

FeSO4-H2O 25-220 9.16

FeSO4-H2SO4-H2O 25-100 6.58

FeSO4-H2SO4-H2O 160-220 4.45

FeSO4-MgSO4-H2O 25-90 9.74

52

FeSO4-MgSO4-H2SO4-H2O 200 12.76

FeSO4-ZnSO4-H2SO4-H2O 140-180 -

FeSO4-ZnSO4-H2SO4-H2O 200 7.33

FeCl2-HCl-H2O 25-100 1.81

FeCl2-MgCl2-H2O 25-90 5.64

FeCl2-MgCl2-HCl-H2O 50 14.66

Fe2O3-H2SO4-H2O 130-170 20.2

Fe2O3-H2SO4-H2O 230-270 23.8

Fe2O3-MgSO4-H2SO4-H2O 250 11.5

NaFe3(SO4)2(OH)6 70-110 12.9

53

Chapter 6

6 Conclusions

Chemical modeling of Fe(II)/(III) in hydrometallurgy using OLI was investigated in this work

with the focus on modeling of FeCl3-MgCl2-HCl-H2O system relevant in hydrometallurgical

treatment of saprolitic laterites using MgCl2 brines. The conclusions of this work are as follows:

1. The solubility of FeCl3 in MgCl2 solutions was investigated in this work from 25°C up to

100°C and up to 4.5 molal MgCl2. It was found out that the solubility increases with

increasing temperature but it decreases with increasing MgCl2 concentration for the

temperature range of 40 to 100°C due to common ion effect. However, the solubility

below 40°C, especially at 25°C shows two opposite trends with the addition of MgCl2.

The solubility first decreases due to common ion effect and then it increases at higher

concentration of MgCl2 due to Fe3+

-Cl- complex formation. In the presence of MgCl2,

the stable solid phase below 37°C is FeCl3.6H2O while a new type of double salt,

2.5FeCl3.MgCl2.7.5H2O was identified at temperature as low as 25°C and up to 100°C.

2. The solubility of MgCl2 in FeCl3 solutions was experimentally investigated in this work

from 25°C up to 100°C and up to 2 molal FeCl3. It was observed that the solubility

increases with increasing temperature but it decreases with increasing FeCl3

concentration due to common ion effect. Furthermore, the decrease of the solubility with

FeCl3 addition diminishes with increasing temperature. Under all conditions studied, the

stable solid phase is always MgCl2.6H2O.

3. The solubility of hematite in MgCl2 and HCl solutions was experimentally determined in

this work at 60 °C and 90 °C and it further extends the concentration range of literature

data previously produced by Konigsberger et al. [57] up to the saturation limit of MgCl2.

Despite the experimental difficulties associated with HCl losses to the vapour, which

prevents from closing the mass balance, several meaningful trends were observed when

the equilibrium HCl was determined from mass balance, not from measurements. It was

observed that the solubility of hematite decreases with increasing temperature but

drastically increases with MgCl2 concentration due to the formation of Fe3+

-Cl-

54

complexes. However, at MgCl2 concentration above 4 m and up to saturation, the

solubility of hematite drops suggesting that there is a common-ion effect at higher MgCl2

concentration. Further investigations are necessary for understanding the solubility

trends of hematite at more concentrated MgCl2 solutions.

4. The default OLI MSE database reproduces literature data on binary chloride systems

(FeCl3-H2O and MgCl2-H2O) and ternary chloride systems (FeCl3-HCl-H2O and MgCl2-

HCl-H2O). The missing binary FeCl3-MgCl2 system was required for good prediction in

the quaternary FeCl3-MgCl2-HCl-H2O. The experimental data on the solubility in the

FeCl3-MgCl2 that was obtained from this work was fitted on the OLI MSE model in order

to obtain new interaction parameters between Fe3+

and Mg2+

species, as well as the

standard state properties of the double salt, 2.5FeCl3.MgCl2.7.5H2O. The model was

found to be valid when compared with experimental solubility data on the quaternary

system obtained from this work not used in the parameterisation.

5. Due to the high solubility of most metal chlorides, it was possible to further increase the

solid loading during leaching with the purpose of reducing the processing equipment

sizes. Since Fe2O3 hydrolytic precipitation during leaching is inhibited by concentrared

MgCl2 brines and the fact that Ni dissolution follows Fe dissolution from the ore, the

process throughput is limited by FeCl3 solubility in the MgCl2 brine solutions under a

given free HCl concentration. Calculations were performed using the developed model to

predict the maximum solids loading that can be achieved without losing the chloride units

pertinent to HCl recovery step. It was estimated that considering the leaching step only,

up to 75% solid loading can be processed in 2 molal MgCl2 brines. However, since MgO

is used to neutralize the free HCl and precipitate Fe, Ni, Co, the MgCl2 concentration

would keep increasing throughout the process and as a result the ultimate limiting factor

will be the MgCl2 solubility. Thus, the operable solid loading is around 35%, which is

far below the predicted FeCl3 solubility.

6. Since FeCl3 addition decreases the solubility of MgCl2, but it further elevates the boiling

point of the solution, the operating window for magnesium hydroxychloride remains

unchanged. However, as Fe is removed from the solution by MgO addition, it is

expected that at the end of Fe removal, the system will approach the solubility of MgCl2

55

in water. Maintaining the MgCl2 level below the solubility limit at any given time during

MgO addition is necessary to prevent premature chloride unit loss.

7. The predicted OLI MSE model on hematite solubility in MgCl2 and HCl solutions is still

far below the literature values. This could be due to the fact that only the first two Fe3+

-

Cl- complexes (FeCl

2+ and FeCl2

+) were considered during the development of the default

OLI MSE model. Addition of the higher order complexes may improve the overall

performance of the model but reworking of the whole FeCl3 system is required.

56

Chapter 7

7 Recommendation for Future Work

Further improvements are required in some aspects of the present work as follows:

1. Further analysis of the double salt, 2.5FeCl3.MgCl2.7.5H2O is required to obtain the exact

stoichiometry since this salt is extremely soluble in any common solvent used in solid

washing, i.e. H2O, alcohols, and acetone. The fact that solid washing was not performed

during this study could result in metal absorption to the surface of the solid, thus affecting

its overall stoichiometry. Furthermore, this compound is very hygroscopic and thus may

further increase the water content during chemical and TGA analysis. Analysing

conditions in the absence of moisture may be required.

2. Hematite solubility in the presence of MgCl2 and HCl needs to be performed in

pressurized vessels in order to suppress the volatility of HCl in the MgCl2 brines. This

could improve the measurement of free HCl and close the mass balance with high degree

of accuracy.

3. OLI Modeling of hematite in MgCl2 and HCl solutions may be further improved by

reworking the whole FeCl3 system, starting from the binary system by adding the higher

order Fe3+

-Cl- complexes not available yet in the OLI MSE default database.

57

References

[1] J. E. Dutrizac and P. A. Riveros, “The precipitation of hematite from ferric chloride media

at atmospheric pressure,” Metallurgical and Materials Transactions B, vol. 30B, pp. 993-

1001, 1999.

[2] J. E. Dutrizac and J. L. Jambor, “Jarosites and their application in hydrometallurgy,” in

Reviews in Mineralogy and Geochemistry, vol. 40, C. N. Alpers, J. L. Jambor and D. K.

Nordstrom, Eds., Washington D.C., Mineralogical Society of America, 2000, pp. 405-451.

[3] D. Paktunc and J. E. Dutrizac, “Characterization of arsenate-for-sulphate substitution in

synthetic jarosite using x-ray diffraction and x-ray absorption spectroscopy,” The Canadian

Mineralogist, vol. 41, pp. 905-919, 2003.

[4] A. D. Dalvi, W. G. Bacon and R. C. Osborne, “The past and the future of nickel laterites,”

in PDAC 2004 International Convention, Toronto, 2004.

[5] J. S. Bakker, “The recovery of magnesium oxide and hydrogen chloride from magnesium

chloride brines and molten salt hydrates,” Kingston, 2011.

[6] D. Duffy, V. G. Papangelakis and B. Davis, “Saprolite leaching and iron precipitation in

magnesium chloride brine,” in COM 2012, Towards Clean Metallurgical Processing for

Profit, Social and Environmental Stewardship, Niagara Falls, ON, 2012.

[7] A. Onozaki, K. Sato and S. Kuramochi, “Effect of some impurities on iron precipitation at

the Iijima Zinc Refinery,” in Iron Control in Hydrometallurgy, J. E. Dutrizac and A. J.

Monhemius, Eds., Chichester, Ellis Horwood, 1986, pp. 742-752.

[8] T. Yamada, H. Nagata, A. Hosoi, M. Kato and R. Togashi, “Recent improvement of

electrolytic zinc production at Iijima Refinery,” in Zinc & Lead '95, T. Akazami, N.

Masuko, J. E. Dutrizac and E. Ozberk, Eds., Tokyo, The Mining & Materials Processing

Institute of Japan, 1995, pp. 589-598.

58

[9] W. F. Linke and A. Seidell, Solubilities, inorganic and metal-organic compounds vol. I & II,

New Jersey: Van Nostrand Company, 1958 & 1965.

[10] V. G. Bruhn, J. Gerlach and F. Pawlek, “Untersuchungen uber die loslichkeiten von salzen

und gasen in wasser und wabrigen losungen bei temperaturen oberhalb 100 C,” Zeitschrift

fur anorganische und allgemeine chemie, vol. 337, pp. 68-79, 1965.

[11] F. Hasegawa, K. Tozawa and T. Nishimura, “Solubility of ferrous sulfate in aqueous

solutions at high temperatures,” Shigen to Sozai, vol. 112, no. 12, pp. 879-884, 1998.

[12] T. C.-M. Cheng, “Production of hematite in acidic zinc sulphate media,” Montreal, 2002.

[13] M. Regel-Rosocka, “A review of methods of regeneration of spent pickling solutions from

steel processing,” Journal of Hazardous Materials, vol. 177, pp. 57-69, 2010.

[14] B. Harris, C. White and B. Ballantyne, “Chloride leaching of polymetallic massive

sulphide,” in ALTA Copper 10, Perth, 2006.

[15] G. B. Harris, C. W. White, G. P. Demopoulos and B. Ballantyne, “Recovery of copper from

a massive polymetallic sulphide by high concentration chloride leaching,” Canadian

Metallurgical Quarterly, vol. 47, no. 3, pp. 347-356, 2008.

[16] P. Wang, A. Anderko and R. D. Young, “A speciation-based model for mixed-solvent

electrolyte systems,” Fluid Phase Equilibria, vol. 203, pp. 141-176, 2002.

[17] Honeywell, “Unisim Design OLI Interface,” 2010.

[18] H. C. Helgeson and D. H. Kirkham, “Theroretical prediction of the thermodynamic

properties of aqueous electrolytes at high pressures and temperatures. I & II.,” American

Journal of Science, vol. 274, pp. 1089-1198 & 1199-1261, 1974.

[19] H. C. Helgeson and D. H. Kirkham, “Theoretical prediction of thermodynamic properties of

aqueous electrolytes at high pressures and temperatures. III. Equation of state for aqueous

species at infinite dilution.,” American Journal of Science, vol. 276, pp. 97-240, 1976.

59

[20] H. C. Helgeson, D. H. Kirkham and G. C. Flowers, “Theoretical prediction of the

thermodynamic behavior of aqueous electrolytes at high pressures and temperatures. IV.

Calculation of activity coefficient, and apparent molal and standard and relative partial

molal properties to 600 degrees C and 5kb,” American Journal of Science, vol. 281, pp.

1249-1516.

[21] E. L. Shock and H. C. Helgeson, “Calculation of the thermodynamic and transport

properties of aqueous species at high pressures and temperatures: correlation algorithms for

ionic species and equation of state predictions to 5kb and 1000 degrees C,” Geochimica et

Cosmochimica Acta, vol. 52, pp. 2009-2036, 1988.

[22] E. L. Shock and H. C. Helgeson, “Calculation of the thermodynamic and transport

properties of aqueous species at high pressures and temperatures: standard partial molal

properties of organic species,” Geochimica et Cosmochimica Acta, vol. 54, pp. 915-945,

1990.

[23] E. L. Shock, H. C. Helgeson and D. A. Sverjensky, “Calculation of the thermodynamic and

transport properties of aqueous species at high pressures and temperatures: standard partial

molal properties of inorganic neutral species,” Geochimica et Cosmochimica Acta, vol. 53,

pp. 2157-2183, 1989.

[24] E. L. Shock, D. C. Sasani, M. Willis and D. A. Sverjensky, “Inorganic species in geologic

fluids: correlations among standard molal thermodynamic properties of aqueous ions and

hydroxide complexes,” Geochimica et Cosmochimica Acta, vol. 61, pp. 907-950, 1997.

[25] D. A. Sverjensky, E. L. Shock and H. C. Helgeson, “Prediction of the thermodynamic

properties of aqueous metal complexes to 1000 degrees C and 5kb,” Geochimica et

Cosmochimica Acta, vol. 61, pp. 1359-1412, 1997.

[26] K. S. Pitzer, “Thermodynamics of electrolytes. I. Theoretical basis and general equations,”

Journal of Physical Chemistry, vol. 77, pp. 268-277, 1973.

[27] K. S. Pitzer, “Electrolytes. From dilute solutions to fused salts,” Journal of American

60

Chemical Society, vol. 102, pp. 2902-2906, 1980.

[28] C. Achard, C. -G. Dussap and J. -B. Gros, “Prediction of pH in complex aqueous mixtures

using a group-contribution method,” American Institute of Chemical Engineers Joudnal,

vol. 40, pp. 1210-1222, 1994.

[29] C. -C. Chen, H. I. Britt, J. F. Boston and L. B. Evans, “Local composition model for excess

Gibbs energy of electrolyte systems. Part I: single solvent, single completely dissociated

electrolyte systems,” American Institute of Chemical Engineers Journal, vol. 28, pp. 588-

596, 1982.

[30] C. -C. Chen and L. B. Evans, “A local composition model for the excess Gibbs energy of

aqueous electrolyte systems,” American Institute of Chemical Engineers Journal, vol. 32,

pp. 444-454, 1986.

[31] K. S. Pitzer and J. M. Simonson, “Thermodynamics of multicomponent, mixcible, ionic

systems: theory and equations,” Journal of Physical Chemistry, vol. 90, pp. 3005-3009,

1986.

[32] Y. Liu and S. Watanasiri, “Representation of liquid-liquid equilibrium of mixed-solvent

electrolyte systems using the extended electrolyte NRTL model,” Fluid Phase Equilibria,

vol. 116, pp. 193-200, 1996.

[33] S. P. Pinho and E. A. Macedo, “Representation of salt solubility in mixed solvents: a

comparison of thermodynamic models,” Fluid Phase Equilibria, vol. 116, pp. 209-216,

1996.

[34] P. Wang, R. D. Springer, A. Anderko and R. D. Young, “Modeling phase equilibria and

speciation in mixed-solvent electrolyte systems,” Fluid Phase Equilibria, Vols. 222-223, pp.

11-17, 2002.

[35] F. Stober, C. Walker, N. Voermann, M. Solar and B. Wasmund, “Evolution and future of

Rotary Kiln-Electric Furnace (RKEF) plants for smelting of nickel laterites,” in ALTA

Nickel/Cobalt Conference, Perth, 2008.

61

[36] J. Kyle, “Nickel laterite processing technologies - where to next?,” in ALTA 2010

Nickel/Cobalt/Copper Conference, Perth, 2010.

[37] A. J. Moyes, “The Intec Nickel Laterite Process,” in ALTA 2005 Ni/Co 10, Perth, Australia,

2005.

[38] B. Harris and J. Magee, “Atmospheric chloride leaching:the way forward for nickel

laterites,” in Hydrometallurgy 2003, Vancouver, 2003.

[39] G. B. Harris, T. J. Magee, V. I. Lakshmanan and R. Sridhar, “The Jaguar Nickel Inc. Sechol

Laterite Project atmospheric chloride leach process,” in Proceedings of International

Laterite Nickel Symposium 2004, W. P. Imrie and D. M. Lane, Eds., Charlotte, 2004, p. 219.

[40] J. T. Smit and J. D. T. Steyl, “Leaching process in the presence of hydrochloric acid for the

recovery of a value metal from an ore”. Patent WO2006043158, 27 April 2006.

[41] J. J. Jansz, “Estimation of ionic activities in chloride systems at ambient and elevated

temperatures,” Hydrometallurgy, vol. 11, pp. 13-31, 1983.

[42] G. P. Demopoulos, Z. Li, L. Becze, G. Moldoveanu, T. C. Cheng and B. Harris, “New

technologies for HCl regeneration in chloride hydrometallurgy,” World of Metallurgy -

ERZMETALL, vol. 61, no. 2, pp. 89-98, 2008.

[43] R. G. McDonald and B. I. Whittington, “Atmospheric acid leaching of nickel laterites

review. Part II. Chloride and bio-technologies,” Hydrometallurgy, vol. 91, pp. 56-69, 2008.

[44] B. Harris, C. White, M. Jansen and D. Pursell, “A new approach to the high concentration

chloride leaching of nickel laterite,” in ALTA Ni/Co 11, Perth, WA, 2006.

[45] B. Harris, “Continuing development of the Jaguar Nickel Sechol laterite Project in

Guatemala,” in Second New Caledonian Nickel Conference, Noumea, 2004.

[46] R. T. Poy, “Regeneration of hydrochloric acid from ferrous chloride using the

Keramchemie/Lurgi Fluidized Bed Reactor System,” in Iron Control and Disposal, Second

62

International Symposium on Iron Control in Hydrometallurgy, J. E. Dutrizac and G. B.

Harris, Eds., Ottawa, 1996, p. 471.

[47] E. M. Peek, O. F. Goedhart and G. V. Weert, “Process evaluation of steel pickle liquor

pyrohydrolysis in a commercial Keramchemie Fluid Bed Reactor,” in Iron Control and

Disposal, Second International Symposium on Iron Control in Hydrometallurgy, J. E.

Dutrizac and G. B. Harris, Eds., Ottawa, 1996, p. 483.

[48] F. Baerhold, A. Lebl and J. Statrcevic, “Recycling of spent acids and iron via

pyrohydrolysis,” in Iron Control Technologies, Proceedings of the Third International

Symposium on Iron Control in Hydrometallurgy, 36th Annual CIM Hydrometallurgical

Meeting, J. E. Dutrizac and P. A. Riveros, Eds., Montreal, 2006, p. 789.

[49] K. Adham, C. Lee and D. Small, “Energy consumption for iron chloride pyrohydrolysis: a

comparison between fluid beds and spray roasters,” in Iron Control Technologies,

Proceedings of the Third International Symposium on Iron Control in Hydrometallurgy,

36th Annual CIM Hydrometallurgical Meeting, J. E. Dutrizac and P. A. Riveros, Eds.,

Montreal, 2006, p. 815.

[50] B. Harris, C. White, G. Demopoulos and L. Becze, “Recovery of metal oxides and

associated acid from process solutions using water as reagent,” in ALTA Copper 2008,

Perth, 2008.

[51] L. Becze, S. J. Hock and G. P. Demopoulos, “Precipitation of hematite and recovery of

hydrochloric acid from concentrated iron chloride solutions by a novel hydrolytic

decomposition process,” in EPD Congress 2011, 2011.

[52] B. Harris and C. White, “Recent developments in the chloride processing of nickel

laterites,” in ALTA 2011, Perth, 2011.

[53] J. W. Burtch, “PORI hydrochloric acid regeneration process,” Iron and Steel Engineer, vol.

50, no. 4, p. 40, 1973.

[54] J. W. Burtch, “PORI hydrochloric acid regeneration process,” SEAISI Quarterly, vol. 2, no.

63

4, p. 52, 1973.

[55] J. W. Burtch, “Hydrochloric acid from industrial waste streams - the PORI Process,” CIM

Bulletin, vol. 68, p. 96, January 1975.

[56] J. W. Burtch, “Regeneration of hydrochloric acid from spent pickle liquor,” Wire Journal,

vol. 9, no. 2, p. 57, 1976.

[57] E. Konigsberger, L. -C. Konigsberger, P. May and B. Harris, “Properties of electrolyte

solutions relevant to high concentration chloride leaching. III. Solubility of pertinent solids

and iron(III)/iron(II) redox potential measured in concentrated magnesium chloride

solutions,” Hydrometallurgy, vol. 90, pp. 192-200, 2008.

[58] Y. Zhang and R. Zhou, “Vapor-Liquid Equilibria for Water+Hydrochloric Acid+magnesium

Chloride and Water+Hydrochloric Acid+Calcium Chloride Systems at Atmospheric

Pressure,” Chinese Journal of Chemical Engineering, vol. 14, no. 2, pp. 276-280, 2006.

[59] H. Gamsjager and F. Reiterer, “Investigation of Equilibria Involving Carbon Dioxide,

Carbonates and Water,” Environmental International, vol. 2, pp. 419-424, 1979.

[60] J. Mendham, R. C. Denney and J. D. Barnes, Vogel's Quantitative Chemical Analysis,

Essex: Pearson Education Ltd., 2000.

[61] A. Kosak and H. Ballozo, “Stufenweise chelatometrische titration von calcium und

magnesium mit visueller endpunktsanzeige neben hoheren eisen(III)-konzentrationen,”

Zeitschrift fur Analytische Chemie, vol. 253, pp. 262-266, 1971.

[62] V. J. Zatka, “The alkalimetric determination of free acid in leach liquors,” Toronto, 1975.

[63] G. N. Kononova and B. A. Redzhepov, “Solubility in the FeSO4-MgSO4-H2O and

Fe2(SO4)3-MgSO4-H2O systems at 25 and 90 degrees C,” Zhurnal Neorganicheskoi

Khimii, vol. 41, no. 7, pp. 1225-1229, 1996.

[64] F. Elgersma, G. J. Witkamp and G. M. Rosmalen, “Incorporation of zinc in ferrous sulfate

64

monohydrate,” Hydrometallurgy, vol. 33, pp. 301-311, 1993.

[65] R. Kammel, F. Pawlek and M. Simon, “Loslichkeitsuntersuchungen von FeSO4 in ZnSO4-

haltigen, schwefelsauren Losungen,” Chemie Ingenieur Technik, vol. 58, no. 4, pp. 323-325,

1986.

[66] I. -M. Chou and L. D. Phan, “Solubility relations in the system sodium chloride-ferrous

chloride-water between 25 and 70 degrees C at 1 atm,” Journal of Chemical Engineering

Data, vol. 30, pp. 216-218, 1985.

[67] A. Atbir, L. Aneflous, A. Marrouche, M. E. Hadek, R. Cohen-Adad and M. -T. Cohen-

Adad, “Diagramme polythermique due systeme ternaire KCl-FeCl2-H2O entre 0 et 70

degrees C,” Journal of Thermal Analysis and Calorimetry, vol. 59, pp. 893-900, 2000.

[68] A. Atbir, S. M. Billah, A. Marrouche, M. E. Hadek, R. Cohen-Adad and M. -T. Cohen-

Adad, “Diagramme polythermique du systeme ternaire NaCl-FeCl2-H2O entre 0 et 70

degrees C,” Journal of Thermal Analysis and Calorimetry, vol. 61, pp. 157-164, 2000.

[69] I. A. Belov, G. I. Novikov, P. K. Rud'ko, S. A. Belov and B. A. Butylin, “The FeCl3-FeCl2-

HCl-H2O system at 298K,” Russian Journal of Inorganic Chemistry, vol. 30, no. 1, pp. 135-

137, 1985.

[70] M. V. Ionin and A. A. Kozhakova, “Phase diagram and physicochemical properties of the

system FeCl2-HCl-H2O,” Zhurnal Prikladnoi Khimii, vol. 46, no. 4, pp. 816-819, 1973.

[71] K. Balarev and D. Spasov, “Formation of double salts from chlorides of bivalent metals,”

Russian Journal of Inorganic Chemistry, vol. 25, no. 10, pp. 1551-1556, 1980.

[72] A. P. Shchedrina, L. I. Krasnova and M. I. Ozerova, “The FeCl2-MgCl2-H2O system at 40

degrees C,” Russian Journal of Inorganic Chemistry, vol. 14, no. 1, pp. 138-139, 1969.

[73] A. P. Shchedrina and L. M. Mel'nichenko, “The FeCl2-MgCl2-H2O system at 70 degrees

C,” Russian Journal of Inorganic Chemistry, vol. 19, no. 6, pp. 897-899, 1974.

65

[74] E. C. Chen, G. McGuire and H. Y. Lee, “Solubility isotherm of the FeCl2-MgCl2-HCl-H2O

system,” Journal of Chemical and Engineering Data, vol. 15, no. 3, pp. 448-449, 1970.

[75] Y. Umetsu, K. Tozawa and K. Sasaki, “The hydrolysis of ferric sulphate solutions at

elevated temperatures,” Canadian Metallurgical Quarterly, vol. 16, pp. 111-117, 1977.

[76] K. Sasaki, K. Ootsuka and K. Tozawa, “Equilibrium diagram in the system Fe2O3-SO3-

H2O at elevated temperatures - hydrometallurgical studies on hydrolysis of ferric sulphate

solutions at elevated temperatures,” Shigen to Sozai, vol. 109, pp. 871-877, 1993.

[77] H. Liu, V. G. Papangelakis, B. C. Blakey, M. S. Alam and G. Singh, “Solubility of hematite

in H2SO4 solution at 230-270 degrees C,” Canadian Metallurgical Quarterly, vol. 42, pp.

199-208, 2003.

[78] M. Reid and V. G. Papangelakis, “New data on hematite solubility in sulphuric acid

solutions from 130 to 270 degrees C,” in Iron Control Technology, 36th Annual

Hydrometallurgy Meeting of CIM, J. E. Dutrizac and P. A. Riveros, Eds., Montreal, 2006,

pp. 673-686.

[79] K. Sasaki, K. Ootsuka, M. Watanabe and K. Tozawa, “Solubility of sodium jarosite in

sulphuric acid solution at temperature range of 70 to 110 degrees C,” Shigen to Sozai, vol.

114, pp. 111-120, 1998.

67

Appendices

Appendix A: Determination of The Concentration of Fe(III), Mg and Free HCl

Appendix A-1: Determination of the concentration of Fe(III)

The Fe(III) content of the solutions containing MgCl2 and HCl was determined by both ICP-OES

and complexometric titration method but ICP-OES measurement suffers from high dilution error

to bring down the concentration of the solution to the detection limit of the instrument. On the

other hand, complexometric titration method works well under concentration nature of the

solution, which makes it the preferable method of quantifying Fe(III). The complexometric

titration works by 1:1 chelation of Fe(III) with ethylenediaminetetraacetic acid (EDTA) at pH

range of 2-3 and temperature of 35-40°C and the end point is indicated by color change from

violet blue to clear yellow using a redox indicator, Variamine blue [60].

Table A-1: The comparison between Fe(III) measurement from ICP and EDTA titration

Theoretical

Conc. of Fe (M)

Fe (M) ICP Fe (M) EDTA Error % ICP Error % EDTA

0.1 0.106 0.100 6 0

0.5 0.498 0.493 0.4 1.4

1.0 0.926 0.992 7.4 0.8

Procedures:

1. Transfer a known volume of the test solution into a beaker filled with DI water (about 50

mL) and continuously stir the solution with magnetic stirrer

2. Add pH 2 buffer and adjust until pH fall between range of 2-3

3. Add 3 drops of the variamine blue indicator solution (0.5% in water), which will change

the color of the solution to purple

68

4. Heat up the solution to 40 °C and titrate with 0.05M standard EDTA solution

5. The end point is indicated by a sharp change from purple to clear yellow

Sample calculations:

1 mL of the diluted test solution (x10 dilution) is titrated with 5mL standard EDTA solution

(0.05M)

Molarity of Fe in the diluted solution = M EDTA * Vol EDTA / Vol test solution

= 0.05 M * 5 mL / 1mL

= 0.25 M

Molarity of Fe in the undiluted solution = M Fe * dilution factor

= 0.25 M * 10

= 2.5 M

Appendix A-2: Determination of the concentration of Mg

Determination of Mg concentration in FeCl3 and HCl solutions cannot be normally measured by

EDTA titration using well-known ErioT indicator at pH 10-10.5 due to interference from Fe, i.e.

precipitation of Fe(OH)3. This precipitation which blurs the color change of the solution is also

responsible for a drop in Mg concentration in the solution, mainly due to coprecipitation with

Fe(OH)3. In order to solve this issue, Mg concentration was measured by first complexing the

Fe(III) using 30% v/v triethanolamine (TEA) to prevent the precipitation of Fe at high pH,

according to an adapted method from [61]. Prior to TEA addition, the solution was acidified

using conc. HCl to prevent precipitation of Fe(OH)3 which could potentially inhibit the kinetics

of Fe(III)-TEA complexation due to the basic nature of TEA solution. The resulting solution

was then neutralized to pH 6-7 and buffer pH 10 (NH3/NH4Cl) was added. This buffer not only

maintain the desired pH but it also makes a weak complex with Mg, which further prevent

Mg(OH)2 precipitation. After addition of the buffer, the resulting solution color turned to clear

from orange yellow. Metallochromic indicator of methylthymol blue (25-35 mg) was added

69

which turns the color of the solution to blue and following this step is the EDTA titration. This

metallochromic indicator works by different color of the indicator, i.e. metal-free color and

metal-indicator complex color. During the beginning of the titration, the color of the solution is

that of metal-indicator complex color. When all of the metal has been chelated with EDTA, the

indicator exists as their metal-free form, thus indicating the end point of the titration. The end

point is indicated by a change from blue to either grey (under low Fe concentration) or rose

(under high Fe concentration). Benchmarking was also performed as a comparison between the

measurement Mg from ICP-OES and this complexometric titration. It was confirmed that the

error of this titration method is below 5% as compared to variation of error up to 20% in the

case of measurement using ICP-OES.

Table A-2: The comparison between Mg measurement from ICP and EDTA titration

Theoretical conc.

of Mg (M)

Mg (M) ICP Mg (M) EDTA Error % ICP Error % EDTA

0.1 0.122 0.103 22 3

0.5 0.544 0.509 8.8 1.8

1 0.990 1.016 1 1.6

Procedures:

1. Transfer a known volume of the test solution into a beaker filled with DI water (about 30

mL) and continuously stir the solution with magnetic stirrer

2. Acidify the solution with 2 mL conc. HCl solution

3. Add 2 mL tartaric acid solution (15g/L) followed by excess 30% TEA solution

determined stoichiometrically from the Fe determination

4. Neutralize the sample to pH 6 using 10N NaOH solution

5. Add 10 mL of the pH 10 buffer solution (NH3/NH4Cl)

6. Ensure the pH of the solution is between 10-10.5 by adjusting with 10N NaOH

70

7. Add 25-35mg of methylthymol blue indicator (1g/100g NaCl) which will change the

color of the solution to blue

8. Titrate the solution with standard EDTA solution (0.05M) until the color change to either

grey or rose

Sample Calculations:

1 mL of the diluted test solution (x10 dilution) is titrated with 2 mL standard EDTA solution

(0.05M)

Molarity of Mg in the diluted solution = M EDTA * Vol EDTA / Vol test solution

= 0.05 M * 2 mL / 1mL

= 0.10 M

Molarity of Mg in the undiluted solution = M Mg * dilution factor

= 0.10 M * 10

= 1 M

Appendix A-3: Free HCl determination

Free HCl concentration is determined by using acid base titration method using 0.1M NaOH

after chelation of Fe(III) using 0.25M Ca-CDTA solution (10-20% excess) [62]. This

complexation prevents hydrolytic precipitation of Fe(III) as Fe(OH)3 during the titration period

which would further increase the HCl concentration as compared to the theoretical value. The

titration was performed using Titroline Easy Autotitrator and the end point was determined

potentiometrically by the use of pH electrode.

Procedures:

1. Transfer a known volume of the test solution into a beaker filled with DI water (60 mL)

71

2. Add 10-20% excess 0.25M Ca-CDTA solution calculated based on the total metal

content (Mg+Fe) from Mg and Fe determinations

3. Immerse the tip of the pH probe into the solution under continuous stirring with magnetic

stirrer

4. Titrate the solution with 0.1N standard NaOH solution using the autotitrator

5. The end point is automatically determined from the 1st derivative of the titration curve

Sample Calculations:

1 mL of the diluted test solution (x10 dilution) is titrated with 0.5 mL standard NaOH solution

(0.1M)

Molarity of HCl in the diluted solution = M NaOH * Vol NaOH / Vol test solution

= 0.10 M * 0.5 mL / 1mL

= 0.05 M

Molarity of HCl in the undiluted solution = M HCl * dilution factor

= 0.05 M * 10

= 0.5 M

72

Appendix B: Solubility Data of FeCl3, MgCl2 and Hematite in the Chloride Systems

Appendix B-1: Solubility Data of FeCl3 in MgCl2 (and HCl) Solutions

Table B-1: Solubility of FeCl3 in water from 25 up to 100°C

T (°C) Density (g/mL) FeCl3 (molal) Stdev from 3

Samples (Fe)

Solid Phase

25 1.5321 6.04 0.09 FeCl3.6H2O

35 1.5830 7.45 0.02 FeCl3.6H2O

40 1.7664 18.19 0.57 FeCl3.2.5H2O

60 1.8160 23.19 0.52 FeCl3.2H2O

80 1.8674 30.26 0.56 FeCl3

100 1.8443 30.92 0.74 FeCl3

Table B-2: Kinetic data for FeCl3 solubility in water at 25°C

Time (h) Density (g/mL) FeCl3 (molal)

0 1.0000 0

1 1.5280 5.64

2 1.5246 5.91

3 1.5395 6.05

4 1.5345 6.11

73

5 1.5380 5.98

6 1.5255 6.08

8 1.5287 5.84

10 1.5282 5.82

12 1.5331 5.95

24 1.5313 6.14

Table B-3: Kinetic data for FeCl3 solubility in 6 molal MgCl2 solutions at 25°C

Time (h) Density (g/mL) MgCl2 (molal) FeCl3 (molal)

0 - 6 0

2 1.5384 3.71 6.44

3 1.5464 4.08 7.24

4 1.5455 3.95 7.24

5 1.5443 4.14 7.56

6 1.5392 3.94 7.43

8 1.5600 3.78 7.95

10 1.5678 3.50 7.62

12 1.5602 3.58 7.93

21 1.5666 3.55 7.87

74

24 1.5606 3.51 7.88

Table B-4: Kinetic data for FeCl3 solubility in water at 40°C

Time (h) Density (g/mL) FeCl3 (molal)

0 - 0

1.5 1.767 16.78

2.5 1.7442 17.32

6 1.7692 18.02

19.5 1.769 17.60

24 1.7617 18.62



0 - 1 0

1 1.6830 1.08 13.84

4 1.7453 0.62 17.34

8 1.7585 0.66 17.26

12 1.7642 0.60 17.98

25 1.7675 0.53 17.75

75



0 - 3 0

1 1.6508 2.57 13.80

4 1.7378 0.88 15.47

8 1.7417 0.84 16.56

12 1.7478 0.65 16.65

25 1.7559 0.55 17.32



0 - 5 0

1.5 1.6234 4.52 12.81

2.5 1.6490 4.60 12.90

6 1.6351 4.53 12.97

19.5 1.6241 4.52 13.07

24 1.6440 4.42 13.02

Table B-8: Solubility data of FeCl3 in MgCl2 solutions from 25 up to 100°C

T (°C) Density MgCl2 FeCl3 Stdev from Stdev from Solid Phase

76

(g/mL) (molal) (molal) 3 Samples

(Mg)

3 Samples

(Fe)

25 1.5229 1.13 5.22 0.026 0.041 FeCl3.6H2O

25 1.5336 2.46 4.89 0.033 0.072 FeCl3.6H2O

25 1.5527 2.92 5.59 0.043 0.153 FeCl3.6H2O

25 1.5603 3.00 5.63 0.067 0.116 FeCl3.6H2O

25 1.5586 3.29 7.10 0.079 0.150 FeCl3.6H2O

25 1.5625 3.54 7.89 0.035 0.033 FeCl3.6H2O

25 1.7269 0.89 13.70 0.048 0.341 DS

25 1.6896 1.93 12.59 0.017 0.141 DS

25 1.6539 3.19 11.52 0.030 0.212 DS

25 1.6412 3.83 11.05 0.081 0.280 DS

35 1.5805 0.29 7.74 0.010 0.079 FeCl3.6H2O

35 1.5999 0.96 8.66 0.022 0.111 FeCl3.6H2O

35 1.5904 1.08 8.74 0.004 0.046 FeCl3.6H2O

40 1.7606 0.56 17.43 0.034 0.772 DS

40 1.7492 0.60 17.50 0.053 0.963 DS

40 1.7099 1.53 13.81 0.106 0.377 DS

40 1.6617 2.60 12.94 0.112 0.458 DS

40 1.6338 4.44 12.97 0.113 0.063 DS

77

60 1.8022 0.74 21.8 0.038 0.309 DS

60 1.7457 1.10 16.94 0.052 1.160 DS

60 1.6928 1.88 15.20 0.020 0.668 DS

60 1.6575 3.54 14.42 0.023 0.121 DS

80 1.8719 0.93 28.41 0.117 2.930 DS

80 1.7675 1.45 20.03 0.070 0.924 DS

80 1.6962 2.80 16.69 0.073 0.242 DS

80 1.6913 3.19 16.69 0.071 0.145 DS

100 1.8636 1.05 33.13 0.070 2.363 DS

100 1.7997 2.48 22.81 0.012 0.264 DS

100 1.7415 2.61 21.03 0.086 0.818 DS

100 1.6852 4.37 15.7 0.066 0.558 DS

Table B-9: Solubility data of FeCl3 in MgCl2 and HCl solutions from 25 up to 100°C

T (°C) Density

(g/mL)

HCl

(molal)

MgCl2

(molal)

FeCl3

(molal)

Stdev (3

samples)

HCl

Stdev (3

samples)

MgCl2

Stdev (3

samples)

FeCl3

Solid Phase

25 1.5527 0.63 2.65 5.94 0.014 0.032 0.051 FeCl3.6H2O

25 1.5481 1.05 2.35 5.99 0.013 0.047 0.063 FeCl3.6H2O

40 1.6709 0.86 2.33 12.87 0.053 0.045 0.544 DS

78

40 1.6398 1.22 3.56 11.93 0.039 0.153 0.508 DS

60 1.6763 0.79 3.38 14.25 0.032 0.086 0.461 DS

60 1.6507 1.30 3.63 13.85 0.0125 0.0274 0.0703 DS

100 1.6852 0.15 4.37 15.70 0.0107 0.0662 0.5579 DS

100 1.6963 0.43 3.99 16.24 0.0376 0.0536 0.1456 DS

Appendix B-2: Solubility Data of MgCl2 in FeCl3 Solutions

Table B-10: Solubility of MgCl2 in water from 25 up to 100°C

T (°C) Density (g/mL) MgCl2 (molal) Solid Phase

25 1.3031 5.79 MgCl2.6H2O

40 1.3048 6.05 MgCl2.6H2O

60 1.3200 6.32 MgCl2.6H2O

80 1.3319 6.80 MgCl2.6H2O

100 1.3463 7.61 MgCl2.6H2O

Table B-11: Kinetic data for MgCl2 solubility in water at 25°C

Time (h) Density (g/mL) MgCl2 (molal)

0 1.0000 0

1 1.3027 5.76

79

2 1.2991 5.81

3.5 1.3070 5.77

Table B-12: Kinetic data for MgCl2 solubility in 0.5 molal FeCl3 solutions at 25°C

Time (h) Density (g/mL) FeCl3 (molal) MgCl2 (molal)

0 - 1.384 0

0.5 1.3455 0.54 5.48

4 1.3319 0.55 5.55

6 1.3410 0.54 5.58

12 1.3439 0.54 5.66

25 1.3374 0.53 5.55

27.5 1.3486 0.53 5.48

Table B-13: Solubility data of MgCl2 in FeCl3 solutions from 25 up to 100°C

T (°C) Density (g/mL) FeCl3 (molal) MgCl2 (molal) Solid Phase

25 1.3430 0.53 5.52 MgCl2.6H2O

25 1.3732 1.09 5.22 MgCl2.6H2O

25 1.4016 1.71 5.03 MgCl2.6H2O

25 1.4402 2.38 4.8 MgCl2.6H2O

80

40 1.3381 0.49 5.81 MgCl2.6H2O

40 1.3657 0.98 5.53 MgCl2.6H2O

40 1.3994 1.51 5.36 MgCl2.6H2O

40 1.4228 2.10 5.32 MgCl2.6H2O

60 1.3414 0.43 6.23 MgCl2.6H2O

60 1.3665 0.85 6.09 MgCl2.6H2O

60 1.3757 1.29 5.96 MgCl2.6H2O

60 1.4013 1.77 5.94 MgCl2.6H2O

80 1.3503 0.36 6.66 MgCl2.6H2O

80 1.3677 0.7 6.65 MgCl2.6H2O

80 1.3680 1.06 6.59 MgCl2.6H2O

80 1.3865 1.46 6.54 MgCl2.6H2O

100 1.3661 0.25 7.36 MgCl2.6H2O

100 1.3714 0.49 7.4 MgCl2.6H2O

100 1.3650 0.74 7.37 MgCl2.6H2O

100 1.3836 1.01 7.35 MgCl2.6H2O

Appendix B-3: Solubility Data of Hematite in MgCl2 and HCl Solutions

81

Table B-14: Kinetic data for hematite solubility in 0.1 molal init. HCl at 60°C (measured)

Time (h) Density (g/mL) HCl Measured

(molal)

FeCl3 (molal)

0 0.9828 0.1001 0

0.33 0.9758 0.0946 0.00082

2 0.9724 0.0929 0.00157

5 0.9715 0.0909 0.00202

9.5 0.9817 0.0889 0.00247

22.5 0.9757 0.0874 0.00300

24 0.9758 0.0864 0.00304

26 0.9778 0.0862 0.00306

Table B-15: Kinetic data for hematite solubility in 0.1 molal init. HCl at 60°C (mass balance)

Time (h) Density (g/mL) HCl from Mass

Balance (molal)

FeCl3 (molal)

0 0.9828 0.1001 0

0.33 0.9758 0.0983 0.00082

2 0.9724 0.0965 0.00157

5 0.9715 0.0952 0.00203

9.5 0.9817 0.0928 0.00247

22.5 0.9757 0.0918 0.00300

82

24 0.9758 0.0917 0.00305

26 0.9778 0.0914 0.00306


60°C (measured)

Time (h) Density (g/mL) MgCl2 (molal) HCl Measured

(molal)

FeCl3 (molal)

0 1.138 2.46 0.1023 0

0.5 1.1449 2.49 0.0726 0.00678

2 1.1393 2.50 0.0403 0.01635

6 1.1361 2.47 0.0251 0.02155

12.5 1.1439 2.45 0.0173 0.02326

24.5 1.1384 2.46 0.0131 0.02394


60°C (mass balance)

Time (h) Density (g/mL) MgCl2 (molal) HCl from Mass

Balance (molal)

FeCl3 (molal)

0 1.138 2.46 0.1023 0

0.5 1.1449 2.49 0.0816 0.00678

2 1.1393 2.50 0.0534 0.01636

83

6 1.1361 2.47 0.0379 0.02156

12.5 1.1439 2.45 0.0319 0.02327

24.5 1.1384 2.46 0.0305 0.02396


60°C (measured)

Time (h) Density (g/mL) MgCl2 (molal) HCl Measured

(molal)

FeCl3 (molal)

0 1.2096 3.86 0.1020 0

0.5 1.2077 3.87 0.0068 0.02472

2 1.217 3.86 0.0034 0.02593

6 1.2219 3.84 0.0028 0.02692

12.5 1.2119 3.89 0.0017 0.02665

24.5 1.2199 3.85 0.0011 0.02641


60°C (mass balance)

Time (h) Density (g/mL) MgCl2 (molal) HCl from Mass

Balance (molal)

FeCl3 (molal)

0 1.2096 3.86 0.1020 0

84

0.5 1.2077 3.87 0.0281 0.02474

2 1.217 3.87 0.0237 0.02595

6 1.2219 3.85 0.0202 0.02693

12.5 1.2119 3.89 0.0221 0.02667

24.5 1.2199 3.85 0.0219 0.02643

Table B-20: Hematite Solubility in MgCl2 and HCl solutions at 60 and 90°C (measured)

T (°C) Init.

Density

(g/mL)

Init.

MgCl2

(molal)

Init. HCl

(molal)

Final

Density

(g/mL)

MgCl2

(molal)

HCl

Measured

(molal)

FeCl3

(molal)

60 0.9865 0 0.0997 0.9758 0 0.0864 0.0030

60 0.9912 0 0.2494 0.9873 0 0.1795 0.0198

60 0.9853 0 1.0320 0.9989 0 0.3328 0.2257

60 1.1444 2.49 0.1008 1.1428 2.47 0.0108 0.0247

60 1.1332 2.48 0.2560 1.1449 2.49 0.0142 0.0732

60 1.1484 2.50 1.0146 1.1735 2.48 0.0187 0.3174

60 1.2330 4.04 0.1014 1.2181 3.98 0.0040 0.0312

60 1.2298 4.08 0.2500 1.2301 4.08 0.0057 0.0797

60 1.2273 4.04 1.0077 1.2445 4.00 0.0063 0.3296

60 1.3010 6.34

(sat'd)

0.0942 1.3057 6.38

(sat'd)

0.0054 0.0270

85

60 1.3048 6.21

(sat'd)

0.2468 1.3148 6.33

(sat'd)

0.0038 0.0744

60 1.3080 6.02

(sat'd)

1.0184 1.3331 6.34

(sat'd)

0.0015 0.3039

90 0.9765 0 0.1007 0.9711 0 0.0878 0.0012

90 0.9711 0 0.2515 0.9784 0 0.1723 0.0179

90 0.9901 0 1.0138 1.0088 0 0.3442 0.2061

90 1.1470 2.47 0.1004 1.1440 2.50 0.0005 0.0132

90 1.1339 2.47 0.2568 1.1416 2.50 0.0044 0.0604

90 1.1472 2.48 1.0143 1.1696 2.50 0.0215 0.3070

90 1.2198 4.00 0.1022 1.2290 3.93 0 0.0171

90 1.2107 3.95 0.2564 1.2239 3.96 0 0.0657

90 1.2190 3.99 1.0156 1.2459 4.04 0.0029 0.3150

90 1.3324 7.07

(sat'd)

0.0613 1.3363 6.88

(sat'd)

0 0.0065

90 1.3383 6.93

(sat'd)

0.2150 1.3317 6.87

(sat'd)

0.0016 0.0517

90 1.3316 6.66

(sat'd)

0.6873 1.3452 6.65

(sat'd)

0.0031 0.2015

86

Table B-21: Hematite Solubility in MgCl2 and HCl solutions at 60 and 90°C (mass balance)

T (°C) Init.

Density

(g/mL)

Init.

MgCl2

(molal)

Init. HCl

(molal)

Final

Density

(g/mL)

MgCl2

(molal)

HCl from

Mass

Balance

(molal)

FeCl3

(molal)

60 0.9865 0 0.0997 0.9758 0 0.0919 0.0030

60 0.9912 0 0.2494 0.9873 0 0.1912 0.0198

60 0.9853 0 1.0320 0.9989 0 0.3520 0.2259

60 1.1444 2.49 0.1008 1.1428 2.48 0.0265 0.0247

60 1.1332 2.48 0.2560 1.1449 2.49 0.0348 0.0732

60 1.1484 2.50 1.0146 1.1735 2.49 0.0517 0.3178

60 1.2330 4.04 0.1014 1.2181 3.98 0.0085 0.0312

60 1.2298 4.08 0.2500 1.2301 4.00 0.0114 0.0798

60 1.2273 4.04 1.0077 1.2445 4.08 0.0147 0.3297

60 1.3010 6.34

(sat'd)

0.0942 1.3057 6.33

(sat'd)

0.0131 0.0270

60 1.3048 6.21

(sat'd)

0.2468 1.3148 6.36

(sat'd)

0.0240 0.0744

60 1.3080 6.02

(sat'd)

1.0184 1.3331 6.38

(sat'd)

0.1147 0.3052

90 0.9765 0 0.1007 0.9711 0 0.0977 0.0012

90 0.9711 0 0.2515 0.9784 0 0.1960 0.0179

87

90 0.9901 0 1.0138 1.0088 0 0.3859 0.2064

90 1.1470 2.47 0.1004 1.1440 2.50 0..0612 0.0133

90 1.1339 2.47 0.2568 1.1416 2.50 0.0748 0.0606

90 1.1472 2.48 1.0143 1.1696 2.51 0.0860 0.3077

90 1.2198 4.00 0.1022 1.2290 3.94 0.0496 0.0171

90 1.2107 3.95 0.2564 1.2239 3.97 0.0572 0.0658

90 1.2190 3.99 1.0156 1.2459 4.05 0.0625 0.3157

90 1.3324 7.07

(sat'd)

0.0613 1.3363 6.67

(sat'd)

0.0410 0.0065

90 1.3383 6.93

(sat'd)

0.2150 1.3317 6.89

(sat'd)

0.0605 0.0518

90 1.3316 6.66

(sat'd)

0.6873 1.3452 6.89

(sat'd)

0.0798 0.2021

88

Appendix C: XRD Patterns of The Equilibrating Solid Phases

Appendix C-1: XRD Patterns for FeCl3 Solubility Experiments

Figure C-1: XRD pattern for FeCl3 solubility in water at 25°C

89

Figure C-2: XRD pattern for Solubility of FeCl3 in 6 molal init. MgCl2 solutions at 25°C

90


91


92

Figure C-5: XRD pattern for FeCl3 solubility in 1 molal init. MgCl2 solutions at 40°C

93


94


Table C-1: XRD peak lists for FeCl3 solubility in 1 molal init. MgCl2 solutions at 60°C

Pos. [°2Th.] Height [cts] FWHM [°2Th.] d-spacing [Å] Rel. Int. [%]

12.7063 33.37 0.5510 6.96694 12.93

14.5881 258.17 0.1574 6.07221 100.00

14.7306 218.34 0.1181 6.01377 84.57

20.5652 210.53 0.0984 4.31890 81.55

24.2222 187.21 0.2755 3.67449 72.51

25.2573 55.50 0.1181 3.52620 21.50

95

31.7198 102.40 0.1181 2.82099 39.66

32.7218 95.37 0.2755 2.73686 36.94

35.8005 44.31 0.1181 2.50825 17.16

38.2206 38.26 0.1968 2.35481 14.82

41.8729 45.75 0.1574 2.15748 17.72

44.2576 29.34 0.2362 2.04661 11.36

47.0894 33.69 0.1968 1.92993 13.05

49.3051 13.82 0.6298 1.84826 5.35

53.2849 23.62 0.2362 1.71922 9.15

54.0951 27.57 0.1181 1.69537 10.68

55.9376 41.63 0.2362 1.64382 16.12

57.6879 6.83 0.5760 1.59672 2.65

96


Table C-2: XRD peak lists for FeCl3 solubility in 6 molal init. MgCl2 solutions at 60°C

Pos. [°2Th.] Height [cts] FWHM [°2Th.] d-spacing [Å] Rel. Int. [%]

12.9480 26.41 0.2362 6.83743 7.70

14.9053 342.83 0.1181 5.94367 100.00

20.9541 109.72 0.1574 4.23961 32.00

24.5530 76.45 0.2755 3.62573 22.30

25.6360 28.90 0.3936 3.47497 8.43

32.1054 99.89 0.1968 2.78799 29.14

33.1280 48.80 0.3149 2.70423 14.23

97

36.1366 20.76 0.2362 2.48568 6.05

38.5709 14.42 0.4723 2.33423 4.21

42.1446 21.37 0.2362 2.14420 6.23

44.6918 9.24 0.6298 2.02773 2.70

47.2670 28.81 0.1968 1.92309 8.40

49.3947 7.04 0.9446 1.84512 2.05

53.9082 10.78 1.1021 1.70081 3.14

56.3959 11.55 0.6298 1.63155 3.37

58.0460 8.69 0.7680 1.58772 2.53

98


99


100

Figure C-11: XRD pattern for FeCl3 solubility in 1 molal init. MgCl2 at 100°C

101


102

Appendix C-2: XRD Patterns for MgCl2 Solubility Experiments

Figure C-13: XRD pattern for MgCl2 solubility in 1.5 molal init. FeCl3 solutions at 25°C

103

Figure C-14: XRD pattern for MgCl2 solubility in 0.5 molal FeCl3 solutions at 80°C

104

Figure C-15: XRD pattern for MgCl2 solubility in 2 molal init. FeCl3 solutions at 80°C

105

Figure C-16: XRD pattern for MgCl2 solubility in 0.5 molal FeCl3 solutions at 100°C

106

Figure C-17: XRD pattern for MgCl2 solubility in 1 molal FeCl3 solutions at 100°C

107

Appendix C-3: XRD Patterns for Hematite Experiments

Figure C-18: XRD pattern for hematite solubility in 0.1 molal HCl and 2.5 molal MgCl2 at 60°C

108

Figure C-19: XRD pattern for hematite solubility in 0.1 molal HCl, sat'd MgCl2 solutions at

60°C

109

Appendix D: Analysis of the Double Salt 2.5FeCl3.MgCl2.7.5H2O

Table D-1: Double salt 2.5FeCl3.MgCl2.7.5H2O chemical analysis

Condition Fe : Mg : Cl : H2O % Cl error

40°C, 1 molal

MgCl2 (1)

3.10 1 10.74 8.59 4.86

40°C, 1 molal

MgCl2 (2)

3.15 1 10.55 9.79 7.92

40°C, 3 molal

MgCl2 (1)

3.03 1 10.21 9.38 7.94

40°C, 3 molal

MgCl2 (2)

2.82 1 9.56 8.87 8.54

60°C, 1 molal

MgCl2 (1)

2.61 1 9.14 8.70 7.03

60°C, 3 molal

MgCl2 (1)

2.67 1 9.15 8.68 8.53

60°C, 3 molal

MgCl2 (2)

2.68 1 8.70 8.40 13.27

80°C, 3 molal

MgCl2 (1)

2.60 1 9.08 8.51 7.34

80°C, 3 molal

MgCl2 (2)

2.63 1 9.04 8.64 8.65

100°C, 3 2.67 1 9.20 9.16 8.22

110

molal MgCl2

(1)

100°C, 3

molal MgCl2

(2)

2.65 1 9.19 8.52 7.59

40°C, DS (1) 2.46 (EDTA) 1 (EDTA) 9.38 (MB) 7.59 0

40°C, DS (2) 2.54 (EDTA) 1 (EDTA) 9.50 (AgCl) 7.40 0.94

Note: Unless otherwise stated, metals analysis (Fe, Mg) were performed using ICP-OES and Cl

by chloride ISE. Crystalline water content was analyzed by mass balance.

EDTA: complexometric titration

MB: mass balance assuming all of metals (Fe and Mg) are in their chloride form

AgCl: Silver chloride precipitation (gravimetric method) for quantifying chloride content

111

Figure D-1: First TGA analysis of the double salt 2.5FeCl3.MgCl2.7.5H2O

112

Figure D-2: Second TGA analysis of the double salt 2.5FeCl3.MgCl2.7.5H2O

113

Appendix E: Validation Plots and Model Improvement for Fe(II)/Fe(III) Systems

Appendix E-1: FeSO4-H2O System

The solubility data for FeSO4 in water was taken from [9] for data up to 100°C and [10, 11] for

data above 100°C. OLI MSE default database predicts the solubility accurately from room

temperature up to 100°C, but above that temperature, OLI overestimates the solubility data.

Figure E-1: Default OLI MSE's simulated solubility of FeSO4 in water

114

Figure E-2: Improved OLI MSE's simulated FeSO4 solubility in water

Appendix E-2: FeSO4-H2SO4-H2O System

The literature data for the solubility of FeSO4 in H2SO4 solutions at elevated temperature was

taken from [9, 11, 12]. The default OLI MSE model complies with the literature data for up to

100°C. Above that temperature, due to inaccuracy of the model in predicting the solubility of

FeSO4 in water, OLI failed to represent the literature solubility data.

115

Figure E-3: Default OLI MSE's simulated FeSO4 solubility in H2SO4 solutions up to 100°C

Figure E-4: Improved OLI MSE's simulated FeSO4 solubility in H2SO4 solutions up to 100°C

116

Figure E-5: Default OLI MSE`s simulated FeSO4 solubility in H2SO4 solutions above 100°C

Figure E-6: Improved OLI MSE`s simulated FeSO4 solubility in H2SO4 solutions above 100°C

117

Appendix E-3: FeSO4-MgSO4-H2O System

The solubility data for FeSO4 in MgSO4 solutions was taken from [63]. The default OLI model

failed to accurately show the effect of MgSO4 addition on the solubility of FeSO4 at any

temperature studied.

Figure E-7: Default OLI MSE's simulated FeSO4 solubility in MgSO4 solutions

118

Figure E-8: Improved OLI MSE's simulated FeSO4 solubility in MgSO4 solutions

Appendix E-4: FeSO4-MgSO4-H2SO4-H2O System

The solubility data for this system was taken from [11]. The default OLI model overestimated

the literature data due to inaccuracy of the solubility prediction in binary FeSO4-H2O above

100°C.

119

Figure E-9: Default OLI MSE's simulated FeSO4 solubility in MgSO4 and H2SO4 solutions

Figure E-10: Improved OLI MSE's simulated FeSO4 solubility in MgSO4 and H2SO4 solutions

120

Appendix E-5: FeSO4-ZnSO4-H2SO4-H2O System

Solubility of FeSO4 in saturated ZnSO4 solutions at high temperatures was studied by [64, 65,

11]. It is obvious that the carry-over error from the solubility of binary FeSO4-H2O causes

overestimation of the solubility values in this system.

Figure E-11: Default OLI MSE's simulated FeSO4 solubility in ZnSO4 and H2SO4 solutions

121

Figure E-12: Improved OLI MSE's simulated FeSO4 solubility in ZnSO4 and H2SO4 solutions

Figure E-13: Default OLI MSE's simulated FeSO4 solubility in ZnSO4 and H2SO4 solutions at

200°C

122

Figure E-14: Improved OLI MSE's simulated FeSO4 solubility in ZnSO4 and H2SO4 solutions at

200°C

Appendix E-6: FeCl2-H2O System

The default OLI model reproduces the solubility data for FeCl2-H2O system up to 120°C taken

from [9, 66, 67, 68]

123

Figure E-15: Default OLI MSE's simulated FeCl2 solubility in water

Appendix E-7: FeCl2-HCl-H2O System

The solubility data for FeCl2 in HCl solutions was taken from [9, 69, 70]. The default OLI MSE

model is in a good agreement with the literature data up to 100°C. However, due to improve the

performance of FeSO4-H2SO4-H2O, minor changes is needed for shared interaction parameter

Fe2+

- H3O+. This change improves the performance of the model for this system, especially at

100°C.

124

Figure E-16: Default OLI MSE's simulated FeCl2 solubility in HCl solutions

Figure E-17: Improved OLI MSE's simulated FeCl2 solubility in HCl solutions

125

Appendix E-8: FeCl2-MgCl2 System

The solubility of FeCl2 in MgCl2 solutions was studied by [71, 72, 73, 57]. OLI MSE prediction

is consistent with the literature data up to 3 molal MgCl2 solutions and above that, OLI

underestimates the literature solubility data. Furthermore, at 70°C, OLI predicted the presence

of FeCl2.2H2O solid phase which was not seen in the literature data. The double salt

FeCl2.MgCl2.8H2O which was present up to 70°C has not been implemented in the default OLI

MSE database.

Figure E-18: Default OLI MSE's simulated FeCl2 solubility in MgCl2 solutions

126

Figure E-19: Improved OLI MSE's simulated FeCl2 solubility in MgCl2 solutions

Appendix E-9: FeCl2-MgCl2-HCl-H2O System

The solubility data was taken from [74]. It is apparent that OLI underestimated the solubility at

elevated MgCl2 concentration due to the same error seen in the binary FeCl2-MgCl2 system.

However, the effect of acid addition which reduces the solubility was correctly shown in the

default model.

127

Figure E-20: Default OLI MSE's simulated FeCl2 solubility in MgCl2 and HCl solutions

Figure E-21: Improved OLI MSE's simulated FeCl2 solubility in MgCl2 and HCl solutions

128

Appendix E-10: Fe2O3-H2SO4-H2O System

Solubility of hematite in sulfuric acid solutions above 100°C was previously studied by [75, 76,

77] but their results were inaccurate due to sampling errors during solubility measurement. More

recently, Reid and Papangelakis re-performed the solubility experiment by addressing all the

sampling errors from previous studies [78]. For this study, only the literature data provided by

Reid and Papangelakis was used. The default OLI MSE database underestimates the solubility

data at any temperature studied.

Figure E-22: Default OLI MSE's simulated Fe2O3 solubility in H2SO4 solutions at 130-170°C

129


Figure E-24: Default OLI MSE's simulated Fe2O3 solubility in H2SO4 solutions at 230-270°C

130


Appendix E-11: Fe2O3-MgSO4-H2SO4-H2O System

The solubility data provided by [78] was used to validate the OLI MSE default model. Due to

underestimation of the solubility of hematite in Fe2O3-H2SO4-H2O system, the errors are carried

over to this quaternary system.

131

Figure E-26: Default OLI MSE's simulated Fe2O3 solubility in MgSO4 and H2SO4 solutions

Figure E-27: Improved OLI MSE's simulated Fe2O3 solubility in MgSO4 and H2SO4 solutions

132

Appendix E-12: NaFe3(SO4)2(OH)6-H2SO4 System

The solubility data was provided by [79] but the default OLI MSE database does not have

sodium jarosite as the default species, thus no OLI simulation could be performed.

Figure E-28: Default OLI MSE's simulated Na-jarosite solubility in H2SO4 solutions

133

Figure E-29: Improved OLI MSE's simulated Na-jarosite solubility in H2SO4 solutions

134

Appendix F: Lists of Regressed OLI Parameters

Table F-1: Default OLI MSE standard state parameters for solid and aqueous species

Species ΔGof (cal/mol) S

o (cal/mol.K) ΔH

of (cal/mol)

FeSO4.7H2O (s) -598064.1 80.16869 -723891

FeSO4.H2O (s) -256326.5 8.65124 -303812.1

FeSO4 (aq) -199444.1 -5.70945 -234599.9

Fe2O3 (s) -175367.6 15.46522 -196566

Table F-2: Regressed standard state parameters of solid and aqueous species

Species ΔGof

(cal/mol)

So

(cal/mol.K)

ΔHof

(cal/mol)

Other Parameters

FeSO4.7H2O (s) -598439.5 123.1721 -711444.1

FeSO4.H2O (s) -257112.1 38.52326 -295691

FeSO4 (aq) -197086.7 54.0136 -214436

Fe2O3 (s) -178267.4 6.183049 -202234

FeSO4+ (aq) -184680 -31.07075 -222705.5 HA1 = 7.6754

HA2 = -24503

HA3 = -696.62

HA4 = 2.6326E+06

135

HC1 = -0.28386

HC2 = 288120

HW = 385280

Fe(SO4)2- (aq) -369152.9 -15.55063 -443403 HC1 = 69.82967

Fe2(SO4)3 (aq) -537220.4 28.03979 -627275.3 HC1 = 170.9113

NaFe3(SO4)2(OH)6 (s) -779067.7 147.3791 -879385 Cp = 616.39 + 0.09121T

- 20376000/T2

Table F-3: Default OLI MSE mid-range binary interaction parameters between ions/aqueous

species

Binary Interaction BMD CMD

Fe2+

- SO42-

BMD0 = 5.720929

BMD1 = -0.02804112

BMD2 = 1179.33

Fe2+

- HSO4- BMD0 = -33.56567

Fe2+

- H3O+ BMD0 = -14.41686

BMD1 = 0.04981539

Mg2+

- HSO4- BMD0 = 19.9975

BMD2 = -6542.72

CMD2 = -9262.47

Fe3+

- H3O+ BMD2 = -665.362

136

FeOH2+

- HSO4- BMD2 = -6000

Table F-4: Regressed mid-range binary interaction parameters between ions/aqueous species

Binary Interaction BMD CMD

Fe2+

- SO42-

BMD0 = 4967.292

BMD1 = -7.854654

BMD2 = -807851.6

CMD0 = -5326.328

CMD1 = 8.973077

CMD2 = 833187.2

Fe2+

- HSO4- BMD0 = -91.83042

BMD1 = 0.2012855

Fe2+

- H3O+ BMD0 = 118.347

BMD1 = -0.181243

BMD2 = -18886.5

Fe2+

- Mg2+

BMD0 = 43.17417

BMD1 = -0.06165801

BMD2 = -4991.604

Fe2+

- Zn2+

BMD0 = -638.8605

BMD1 = 1.36925

Mg2+

- HSO4- BMD0 = 6.953257

BMD1 = -0.001238579

Fe3+

- H3O+ BMD0 = -39.92635

137

BMD1 = -0.1768866

BMD2 = -10920

FeOH2+

- HSO4- BMD0 = 226.0031

BMD1 = -5.6340103E-2

BMD2 =-61722.12

FeOH2+

- H3O+ BMD0 = 210.5307

BMD1 = -0.3667933

BMD2 = 15460.15

Fe(SO4)2- - HSO4

- BMD0 = -2.598352

BMD1 = -0.4128314

BMD2 = 119172.1

Fe(SO4)2- - H3O

+ BMD0 = 18.71168

BMD1 = -0.4223134

BMD2 = 58631.09

Fe3+

- Na+ BMD0 = 240.6021

BMD1 = -0.9782702

BMD2 = 271251.9

Fe(SO4)2- - Na

+ BMD0 = 116.9141

BMD1 = 0.3929999

BMD2 = 1932.883

chemical modeling of iron (ii)/(iii) solutions in hydrometallurgy

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