POTENTIOMETRIC ph MEASUREMENTS IN THE PRESSURE ACID LEACHING OF NICKEL LATERITES

Transcription

1 Potentiometric ph Measurements in the Pressure Acid Leaching of Nickel Laterites by Zoran Jankovic A thesis submitted in conformity with the requirements for the degree of Doctor of Philosophy Graduate Department of Chemical Engineering and Applied Chemistry University of Toronto Copyright by Zoran Jankovic 2010

2 POTENTIOMETRIC ph MEASUREMENTS IN THE PRESSURE ACID LEACHING OF NICKEL LATERITES Zoran Jankovic Doctor of Philosophy 2010 Graduate Department of Chemical Engineering and Applied Chemistry University of Toronto ABSTRACT An electrochemical cell consisting of a flow-through yttria-stabilized zirconia (YSZ) sensor and a flow-through Ag/AgCl reference electrode has been employed to measure ph of high-temperature acidic sulphate solutions relevant to the pressure acid leaching (PAL) of nickel laterites. In a previous study, this cell was used to measure ph of H 2 SO 4, Al 2 (SO 4 ) 3 -H 2 SO 4 and MgSO 4 -Al 2 (SO 4 ) 3 -H 2 SO 4 solutions at 250 o C. In this work, the solutions range in complexity from the binary MgSO 4 -H 2 SO 4, NiSO 4 -H 2 SO 4, and Al 2 (SO 4 ) 3 -H 2 SO 4, through the ternary MgSO 4 -Al 2 (SO 4 ) 3 -H 2 SO 4 and NiSO 4 -Al 2 (SO 4 ) 3 - H 2 SO 4, to the PAL process solutions, whereas the temperature ranges from 200 o C to 250 o C. The measured and theoretical ph values typically agree within less than 0.1 ph unit and 0.2 ph units in synthetic solutions and PAL solutions, respectively. This is an improvement over the results of the previous study in synthetic solutions, which show differences between theory and experiment as high as 0.4 ph units. The conversion of measured potentials into ph values is based on the new mixed-solvent electrolyte (MSE) ii

3 speciation model of the OLI Systems software calibrated independently based on solubility measurements. Both Henderson s equation and the exact definition of the diffusion potential were employed in treating the obtained experimental data. Experimental ph values calculated using the diffusion potentials evaluated by either approach are essentially the same. This finding suggests that Henderson s equation, which is based on readily available limiting ionic mobilities, can be effectively used. Lithium chloride is found to be a suitable alternative to sodium chloride as the reference electrode solution for the measurement of ph of aluminium-containing solutions, because it did not induce precipitation of aluminium as an alunite-type compound. The experimental results indicate that the high-temperature behaviour of Ni, Co and Mn sulphates can be satisfactorily approximated with that of MgSO 4. The experimental findings also support the postulation that acid should be added to a PAL process so that the solution ph is around 1 at the leach temperature, regardless of the feed composition. The cell can be used for hydrometallurgical process research and development on a laboratory scale with very satisfactory performance, provided that a well-behaved YSZ sensor is available. iii

4 ACKNOWLEDGEMENTS This thesis was conducted under the supervision of Professor V.G. Papangelakis. I would like to thank him for his guidance, advice and support during the course of this thesis. I would like to express my thanks to Professor S.N. Lvov for helpful discussions and advice on ph measurement. I would like to thank my reading committee, Professors D.W. Kirk and R.C. Newman, for their advice, suggestions and comments. I would also like to thank Professor S.J. Thorpe for constructive criticism at my Departmental Defence. I would like to especially thank Professor A. Alfantazi for reviewing my thesis and providing valuable feedback. I would like to thank Dr. M.J. Collins for useful discussions on Pressure Acid Leach solutions. Thanks are also due to the Aqueous Process Engineering and Chemistry Group. I would like to thank Dr. D. Seneviratne, Dr. J. Adams, Mr. G. Singh, Mr. J. Brown, Mr. M. Reid, Dr. I. Bylina, Mr. S. Peters, Mr. R. Saini and Mr. I. Perederiy for their help in the lab. I would also like to thank Dr. H. Liu, Dr. M. Huang, Dr. G. Azimi and Mr. S. Roshdi for their support on OLI. I would finally like to thank my wife and my family for their support and patience. iv

5 TABLE OF CONTENTS ABSTRACT ACKNOWLEDGEMENTS TABLE OF CONTENTS LIST OF TABLES LIST OF FIGURES LIST OF APPENDICES ii iv v ix xii xvii 1 INTRODUCTION Background Objectives 4 2 THEORETICAL SECTION Pressure acid leaching of nickel laterites Low-temperature ph measurement Definition of ph Primary method of measurement and primary ph standards Secondary methods of measurement and secondary ph standards Glass electrode cells and their calibration High-temperature potentiometric ph measurement techniques Reference electrodes for high-temperature measurements High-temperature indicator electrodes Hydrogen electrode Yttria-stabilized zirconia ph sensor 20 v

6 2.4 Principle of operation of the yttria-stabilized zirconia ph sensor Structural properties of yttria-stabilized zirconia Electrical properties of yttria-stabilized zirconia Mechanism of ph response Effects of chemical purity Resistance to chemical attack Thermodynamics of the high-temperature electrochemical cell Thermoelectric potential Diffusion potential Thermal diffusion potential Streaming potential Calculation of ph Speciation models 49 3 EXPERIMENTAL SECTION Measurement of ph Apparatus Flow-through yttria-stabilized zirconia ph electrode Flow-through Ag/AgCl reference electrode Measurement procedure Pressure acid leaching 63 4 RESULTS AND DISCUSSION Effect of nickel sulphate and magnesium sulphate on ph of sulphuric acid solutions at elevated temperatures 66 vi

7 4.1.1 Effect of MgSO 4 on the ph of H 2 SO 4(aq) solutions at 250 o C with NaCl (aq) as the reference electrode solution Effect of NiSO 4 on the ph of H 2 SO 4(aq) solutions at 250 o C with NaCl (aq) as the reference electrode solution Effect of MgSO 4 on the ph of H 2 SO 4(aq) solutions at 200 o C with NaCl (aq) as the reference electrode solution Effect of MgSO 4 on the ph of H 2 SO 4(aq) solutions at 250 o C and 200 o C with LiCl (aq) as the reference electrode solution Effect of NiSO 4 on the ph of H 2 SO 4(aq) solutions at 250 o C with LiCl (aq) as the reference electrode solution Measurement of ph in aluminium-containing acidic sulphate solutions Measurements at 250 o C Al 2 (SO 4 ) 3 -H 2 SO 4 solutions at 250 o C MgSO 4 -Al 2 (SO 4 ) 3 -H 2 SO 4 solutions at 250 o C NiSO 4 -Al 2 (SO 4 ) 3 -H 2 SO 4 solutions at 250 o C Measurements at 200 o C Al 2 (SO 4 ) 3 -H 2 SO 4 solutions at 200 o C MgSO 4 -Al 2 (SO 4 ) 3 -H 2 SO 4 solutions at 200 o C Measurement of ph in high-temperature nickel laterite pressure acid leach process solutions Pressure acid leach solutions Diluted, acid-adjusted pressure acid leach solutions Synthetic NiSO 4 -H 2 SO 4 solutions 116 vii

8 4.3.4 Process implications Sensitivity analysis Potential variability Comparison of calculations using different speciation models Feasibility of using an internal reference electrode CONCLUSIONS RECOMMENDATIONS NOMENCLATURE REFERENCES 169 APPENDICES 178 Appendix A: Dielectric constants of bisulphate ion and water at temperatures to 250 o C 178 Appendix B: Data for ph measurements with yttria-stabilized zirconia electrodes made from etched tubes 180 Appendix C: Data for ph measurements with yttria-stabilized zirconia electrodes made from as received tubes 187 Appendix D: Resistivities of yttria-stabilized zirconia tubes 209 Appendix E: Data for experiments with an internal reference electrode 211 viii

9 LIST OF TABLES Table No. Title Page 4.1 Calibration data for the MgSO 4 -H 2 SO 4 -H 2 O and NiSO 4 -H 2 SO 4 - H 2 O systems at 250 o C with 0.1 mol kg -1 NaCl as the reference electrode solution Compositions of test solutions and corresponding diffusion potentials for the MgSO 4 -H 2 SO 4 -H 2 O system at 250 o C with 0.1 mol kg -1 NaCl as the reference electrode solution Comparison between experimental ph values (calculated based on diffusion potentials evaluated using Henderson s equation) and theoretical ph values (calculated with OLI) for the MgSO 4 -H 2 SO 4 - H 2 O system at 250 o C with 0.1 mol kg -1 NaCl as the reference electrode solution Comparison between experimental ph values (calculated based on diffusion potentials evaluated using Harper s equation) and theoretical ph values (calculated with OLI) for the MgSO 4 -H 2 SO 4 - H 2 O system at 250 o C with 0.1 mol kg -1 NaCl as the reference electrode solution Compositions of test solutions and corresponding diffusion potentials for the NiSO 4 -H 2 SO 4 -H 2 O system at 250 o C with 0.1 mol kg -1 NaCl as the reference electrode solution Calibration data for the MgSO 4 -H 2 SO 4 -H 2 O system at 200 o C with 0.1 mol kg -1 NaCl as the reference electrode solution Compositions of test solutions and corresponding diffusion potentials for the MgSO 4 -H 2 SO 4 -H 2 O system at 200 o C with 0.1 mol kg -1 NaCl as the reference electrode solution Calibration data for the MgSO 4 -H 2 SO 4 -H 2 O and NiSO 4 -H 2 SO 4 - H 2 O systems at 250 o C with 0.1 mol kg -1 LiCl as the reference electrode solution Calibration data for the MgSO 4 -H 2 SO 4 -H 2 O system at 200 o C with 0.1 mol kg -1 LiCl as the reference electrode solution Compositions of test solutions and corresponding diffusion potentials for the MgSO 4 -H 2 SO 4 -H 2 O system at 250 o C with 0.1 mol kg -1 LiCl as the reference electrode solution. 83 ix

10 4.11 Compositions of test solutions and corresponding diffusion potentials for the MgSO 4 -H 2 SO 4 -H 2 O system at 200 o C with 0.1 mol kg -1 NaCl as the reference electrode solution Compositions of test solutions and corresponding diffusion potentials for the NiSO 4 -H 2 SO 4 -H 2 O system at 250 o C with 0.1 mol kg -1 LiCl as the reference electrode solution Compositions of test solutions and corresponding diffusion potentials for the Al 2 (SO 4 ) 3 -H 2 SO 4, MgSO 4 -Al 2 (SO 4 ) 3 -H 2 SO 4 and NiSO 4 -Al 2 (SO 4 ) 3 -H 2 SO 4 systems at 250 o C Compositions of test solutions and corresponding diffusion potentials for the Al 2 (SO 4 ) 3 -H 2 SO 4 and MgSO 4 -Al 2 (SO 4 ) 3 -H 2 SO 4 systems at 200 o C Equilibrium constants for ion association of divalent metal sulphates at 25 o C Ionic mobilities at infinite dilution and 25 o C Elemental composition (wt%) of laterite feeds Pressure acid leaching conditions under which solutions P#1, P#2 and P#3 were obtained Compositions of leach solutions P#1, P#2 and P#3 in g L Compositions of leach solutions P#1, P#2 and P#3 in mol kg Compositions of synthetic solutions SP#1 and SP# Compositions of diluted, acid-adjusted leach solutions D#1 to D# Compositions of binary NiSO 4 -H 2 SO 4 synthetic solutions Ni#1 to Ni# Values for constants A, B and C in Equation (4.3) depending on the type of equation Calibration data for the MgSO 4 -H 2 SO 4 -H 2 O system at 250 o C Calculated values of the calibration coefficient Composition of test solutions and corresponding diffusion potentials for the MgSO 4 -H 2 SO 4 -H 2 O system at 250 o C. 151 x

11 4.28 Comparison between the experimental and theoretical ph values, obtained using different speciation models, for the MgSO 4 -H 2 SO 4 - H 2 O system at 250 o C List of symbols Values of physical constants List of acronyms 168 xi

12 LIST OF FIGURES Figure No. Title Page 2.1 Schematic representation of the cell with a floating hydrogen electrode Schematic representation of the hydrogen-electrode concentration cell Schematic representation of the ph measuring system involving an yttria-stabilized zirconia electrode with a Ag/AgCl internal reference element Electrochemical diagram of a symmetrical cell consisting of two Ag/AgCl electrodes (Cell I) Electrochemical diagram of Cell II involving a Ag/AgCl reference electrode and a hydrogen electrode Electrochemical diagram of Cell III comprising a Ag/AgCl reference electrode and an yttria-stabilized zirconia ph sensor ZrO 2 -rich part of the ZrO 2 -Y 2 O 3 phase diagram: monoclinic (m), tetragonal (t) and cubic (c) phases Electrochemical diagram of the non-isothermal electrochemical system consisting of a flow-through external pressure-balanced Ag/AgCl reference electrode and a flow-through yttria-stabilized zirconia electrode Schematic representation of the flow-through apparatus used for potentiometric ph measurements Schematic representation of the flow-through yttria-stabilized zirconia ph electrode Schematic representation of the Ag/AgCl flow-through external pressure balanced reference electrode (FTEPBRE) Potential of the flow-through YSZ electrode recorded after replacing 0.01 mol kg -1 H 2 SO 4 with 0.02 mol kg -1 H 2 SO 4 at 250 o C and a flow rate of 0.5 ml min xii

13 4.2 Potential of the flow-through YSZ electrode recorded after replacing 0.3 mol kg -1 H 2 SO mol kg -1 MgSO 4 with mol kg -1 H 2 SO 4 at 250 o C and a flow rate of 0.5 ml min Regression lines for the first calibration of electrode YSZ-E2 at 250 o C Variation of ph as a function of MgSO 4 concentration at different H 2 SO 4 concentrations at 250 o C with 0.1 mol kg -1 NaCl as the reference electrode solution Variation of ph as a function of NiSO 4 concentration at different H 2 SO 4 concentrations at 250 o C with 0.1 mol kg -1 NaCl as the reference electrode solution Variation of ph as a function of MgSO 4 concentration at different H 2 SO 4 concentrations at 200 o C with 0.1 mol kg -1 NaCl as the reference electrode solution Variation of ph as a function of MgSO 4 concentration at different H 2 SO 4 concentrations at 250 o C with 0.1 mol kg -1 LiCl as the reference electrode solution Variation of ph as a function of MgSO 4 concentration at different H 2 SO 4 concentrations at 200 o C with 0.1 mol kg -1 LiCl as the reference electrode solution Variation of ph as a function of NiSO 4 concentration at different H 2 SO 4 concentrations at 250 o C with 0.1 mol kg -1 LiCl as the reference electrode solution Variation of ph as a function of Al 2 (SO 4 ) 3 concentration at different H 2 SO 4 concentrations at 250 o C Effect of Al 2 (SO 4 ) 3 additions on solution speciation at 0.4 mol kg -1 H 2 SO 4 and 250 o C Theoretical (lines) and experimental (symbols) ph values at ( ) 0.1 mol kg -1 H 2 SO 4 (Section 4.1), ( ) 0.3 mol kg -1 H 2 SO 4 (Section 4.1), ( ) 0.2 mol kg -1 H 2 SO mol kg -1 Al 2 (SO 4 ) 3, ( ) 0.4 mol kg -1 H 2 SO mol kg -1 Al 2 (SO 4 ) 3 and ( ) mol kg -1 H 2 SO mol kg -1 Al 2 (SO 4 ) 3 [25] Effect of MgSO 4 additions on solution speciation at 0.4 mol kg -1 H 2 SO 4 (dotted lines) and mol kg -1 Al 2 (SO 4 ) mol kg -1 H 2 SO 4 (solid lines), at 250 o C. 96 xiii

14 4.14 Variation of ph as a function of NiSO 4 concentration, at different H 2 SO 4 concentrations and a fixed Al 2 (SO 4 ) 3 concentration, at 250 o C Variation of ph as a function of Al 2 (SO 4 ) 3 concentration at different H 2 SO 4 concentrations at 200 o C Variation of ph as a function of MgSO 4 concentration, at different H 2 SO 4 concentrations and a fixed Al 2 (SO 4 ) 3 concentration, at 200 o C Variation of ph as a function of H 2 SO 4 concentration at different MgSO 4 concentrations for pressure acid leach solutions P#1 and P# Variation of ph as a function of H 2 SO 4 concentration at different MgSO 4 concentrations for synthetic solutions SP#1 and SP# Variation of ph as a function of H 2 SO 4 concentration at different MgSO 4 concentrations for diluted, acid-adjusted solutions D#1 to D# Variation of ph as a function of H 2 SO 4 concentration at different MgSO 4 concentrations for synthetic solutions SD#1 to SD# Variation of ph as a function of H 2 SO 4 concentration at different NiSO 4 concentrations for binary NiSO 4 -H 2 SO 4 synthetic solutions Ni#1 to Ni# Recalculated experimental and theoretical ph values for binary MgSO 4 *-H 2 SO 4 synthetic solutions Mg*#1 to Mg*# Variation of free H 2 SO 4 as a function of equivalent-magnesium at various temperatures. Symbols refer to experimental data at: (+) 245 o C [122]; ( ) 250 o C [31,121]; ( ) 260 o C (this work); ( ) 260 o C [124]; ( ) 270 o C [31,121]; ( ) 270 o C [123] Effect of variations in ph of the calibrating solutions (ph 1 and ph 2 ) and temperature (T) on the calibration coefficient Effect of variation in measured potentials for the calibrating solutions (E 1 and E 2 ) on the calibration coefficient Effect of variation in the activity of water on the calibration coefficient 127 xiv

15 4.27 Effect of variation in diffusion potential on the calibration coefficient Effect of variation in measured potentials for the calibrating solution (E 2 ) and test solution (E 1 ) on the ph of the test solution (ph 1 ) Effect of variations in temperature (T), calibration coefficient (α) and ph of the calibrating solution (ph 2 ) on the ph of the test solution (ph 1 ) Effect of variation in the activity of water on the ph of the test solution (ph 1 ) Effect of variation in diffusion potential on the ph of the test solution (ph 1 ) Effect of variation in temperature and ph on the potential of an ideal YSZ electrode Effect of variation in the activity of water on the potential of an ideal YSZ electrode Effect of temperature on the potential of the reference electrode Measured potentials at 250 o C (0.1 mol kg -1 H 2 SO 4 ) versus measured tube resistivities at 90 o C (0.1 mol kg -1 KCl) for untreated tubes Measured potentials at 200 o C (0.1 mol kg -1 H 2 SO 4 ) versus measured tube resistivities at 90 o C (0.1 mol kg -1 KCl) for untreated tubes Measured potentials at 250 o C (0.1 mol kg -1 H 2 SO 4 ) versus measured tube resistivities at 90 o C (0.1 mol kg -1 KCl) for etched tubes Measured potentials at 200 o C (0.1 mol kg -1 H 2 SO 4 ) versus measured tube resistivities at 90 o C (0.1 mol kg -1 KCl) for etched tubes Calculated ph in H 2 SO 4 solutions at 250 o C using the Alpha_2, OLI MSE and OLI Aqueous speciation models Calculated species concentrations in H 2 SO 4 solutions at 250 o C using the OLI MSE speciation model 149 xv

16 4.41 Calculated species concentrations in H 2 SO 4 solutions at 250 o C using the OLI Aqueous speciation model Effect of MgSO 4 additions on solution speciation at 0.3 mol kg -1 H 2 SO 4 and 250 o C as calculated using the OLI MSE speciation model Effect of MgSO 4 additions on solution speciation at 0.3 mol kg -1 H 2 SO 4 and 250 o C as calculated using the OLI Aqueous speciation model Potential difference between an internal Ag/AgCl electrode and an external Ag/AgCl electrode ( T=225 o C) in 0.1 mol kg -1 KCl Potential difference between an internal Ag/AgCl electrode and an external Ag/AgCl electrode ( T=225 o C) in 0.1 mol kg -1 LiCl Potential difference between an internal Ag/AgCl electrode and an external Ag/AgCl electrode ( T=225 o C) in 0.1 mol kg -1 LiCl mol kg -1 H 2 SO mg L -1 Fe 2 (SO 4 ) 3 5H 2 O 158 xvi

17 LIST OF APPENDICES Appendix No. Title Page Appendix A Appendix B Appendix C Dissociation constants of bisulphate ion and water at temperatures to 250 o C 178 Data for ph measurements with yttria-stabilized zirconia electrodes made from etched tubes 180 Data for ph measurements with yttria-stabilized zirconia electrodes made from as received tubes 187 Appendix D Resistivities of yttria-stabilized zirconia tubes 209 Appendix E Data for experiments with an internal reference electrode 211 xvii

18 1 INTRODUCTION 1.1 Background The measurement of ph at high temperatures is highly desirable for the monitoring and control of industrial processes. A typical example is pressure acid leaching (PAL) of laterite ores, which is conducted in autoclaves at a temperature around 250 o C, to extract nickel and cobalt [1,2]. PAL solutions consist of excess sulphuric acid and a number of metal sulphates. The concentration of H 2 SO 4 ranges from 0.20 to 0.60 mol kg -1, the Ni, Mg, Mn and Al concentrations are between 0.05 and 0.5 mol kg -1, the Co concentration is about 0.01 mol kg -1, and the concentration of other metals (Fe, Cr, Cu, Zn) is below mol kg -1 [1,2]. The sulphuric acid consumption for pressure acid leaching of limonitic laterites is relatively low ( kg per tonne of dry ore). However, limonite/saprolite blends require much higher acid additions ( kg per tonne of dry ore). The extra acid requirement is considerably higher than the stoichiometric requirement to dissolve extra magnesium and other soluble constituents in saprolitic laterites. Although magnesium consumes part of the acid, a significant amount of the acid remains in solution as free H 2 SO 4. Free acid measured by titration at room temperature, however, does not reflect the true concentration of hydrogen ion (or ph) at elevated temperatures, which in fact controls the kinetics of the PAL process [3]. Changes in solution ph result from two primary causes. The first cause is due to changes in equilibrium constants with variation in temperature (Appendix A). Precipitation or dissolution of phases with consumption or generation of hydrogen ions is the second factor responsible for ph change. Acid cost is a primary factor in the economics of the PAL process. Accordingly, optimizing acid 1

19 additions to the PAL process confers savings in acid and downstream neutralization costs, while at the same time ensuring high productivity and minimizing corrosion problems [4]. In the absence of commercial sensors capable of measuring ph under autoclave conditions, acid addition requirements have been calculated using speciation models where PAL process solutions were simplified to ternary MgSO 4 -Al 2 (SO 4 ) 3 -H 2 SO 4 solutions by combining all the divalent metal sulphates into one with the chemical and thermodynamic properties of MgSO 4. These calculations suggest that high nickel and cobalt extractions are achieved when sufficient acid is added to maintain the solution ph equal to 1 at temperature, regardless of feed composition [4]. Initial precise potentiometric ph measurements in high-temperature aqueous solutions were carried out using platinum electrodes as both the test and reference electrodes in hydrogen concentration cells in the 1970s [5,6]. Cells of this type have since been used for potentiometric measurements at temperatures to 300 o C at Oak Ridge National Laboratory [7,8]. This method, however, is limited to the systems that are stable against reduction by hydrogen and therefore they cannot be used for aqueous solutions containing oxidizing species. With the invention of the yttria-stabilized zirconia (YSZ) ph sensor in 1980 [9,10], it became possible to measure ph of high-temperature aqueous solutions in the presence of oxidizing species. In 1984, the YSZ sensor was demonstrated to behave in a Nernstian manner at temperatures to 275 o C in solutions having ph values from about 2 to 9 at 275 o C [11]. The precision of these measurements was estimated to be about ±10 mv corresponding to about ±0.1 ph unit at the highest temperature. Since then, a number of researchers have investigated various designs of the YSZ sensor [12-2

20 21]. However, the level of precision noted above was not improved until Lvov and coworkers from the Pennsylvania State University developed an electrochemical cell consisting of a flow-through YSZ ph sensor and a flow-through Ag/AgCl reference electrode in 2001 [22-24]. The precision of this cell was established to be ±5 mv or about ±0.05 ph units up to 350 o C in HCl solutions with ph values between about 2 and 3 at 350 o C. This electrochemical cell was employed by Seneviratne et al. from the University of Toronto in 2003 to measure ph of H 2 SO 4 as well as binary Al 2 (SO 4 ) 3 -H 2 SO 4 and ternary MgSO 4 -Al 2 (SO 4 ) 3 -H 2 SO 4 aqueous solutions at 250 o C [25]. The solutions ranged in ph from about 0.6 to 2.1 at 250 o C. An average difference of ±0.15 ph units was observed between theory and experiment. However, the conversion of measured potentials into ph values was based on a speciation model incorporating the B-dot equation (equation (2.86)), whereas theoretical ph values were obtained from the Aqueous model of the OLI Systems software based on the Bromley-Zemaitis activity equation (equation (2.81)). An exact integration formula was employed for calculating diffusion potentials between acidic solutions and the NaCl reference electrode solution (equation (2.58)). Scale formation at the outlet of the electrochemical cell was observed during subsequent measurements in pressure acid leach solutions in 2004 [26], even though aluminium concentrations were not high enough to cause precipitation of hydronium alunite, H 3 OAl 3 (SO 4 ) 2 (OH) 6, according to equation (2.4). This scale was determined to be sodium alunite, NaAl 3 (SO 4 ) 2 (OH) 6, which was produced as a result of the mixing of the aluminium-containing PAL solutions and the NaCl reference electrode solution according to equation (2.5). This observation was consistent with the fact that 3

21 sodium alunite is less soluble than hydronium alunite. As a result, the Al 2 (SO 4 ) 3 effect on solution ph was not unequivocally clear because theoretical and experimental ph values were based on the two different speciation models and the concentration of aluminium was not controlled. 1.2 Objectives The ultimate objective of this thesis was to measure ph of high-temperature pressure acid leaching (PAL) process solutions using the electrochemical cell consisting of a flow-through yttria-stabilized zirconia (YSZ) ph sensor and a flow-through Ag/AgCl reference electrode in order to check the speciation models that have been developed in the Aqueous Process Engineering and Chemistry group at the University of Toronto based on the solubility data. Specific objectives included: (a) to investigate the effects of NiSO 4, MgSO 4 and Al 2 (SO 4 ) 3 on the ph of H 2 SO 4 solutions, (b) to examine the effect of NiSO 4 and MgSO 4 on the ph of Al 2 (SO 4 ) 3 -H 2 SO 4 solutions, (c) to use the new mixedsolvent electrolyte (MSE) speciation model of the OLI Systems software to convert the measured potentials into ph values, (d) to investigate whether Henderson s equation (equation (2.57)), which is based on readily available limiting ionic mobilities, provides an acceptable accuracy when used instead of the exact integration for calculating the diffusion potential, (e) to test whether the NaCl in the reference electrode solution can be replaced with another chloride electrolyte that produces no alunite-type scale, (f) to check whether the approximation of the high-temperature behaviour of divalent metal sulphates with that of MgSO 4 is acceptable and (g) to test whether the previous conclusion (based on speciation calculations), that acid should be added to the PAL process so that the 4

22 solution ph at temperature is around 1 regardless of feed composition, is justified by direct measurements. 5

23 2 THEORETICAL SECTION 2.1 Pressure acid leaching of nickel laterites While about 70% of the world s terrestrial nickel resources are contained in laterite deposits, they accounted for only about 40% of the world s nickel production in However, the share of laterite source nickel is expected to rise to about 50% by Most of this expansion in nickel production will come from pressure acid leaching (PAL) of laterites [29]. Nickel laterites refer to oxide or silicate ores developed by weathering from the peridotite bedrock. The type of weathering that produces laterites is most frequent in tropical or subtropical regions. Peridotite is composed mostly of olivine or serpentine (hydrothermally altered olivine). These minerals are attacked by ground water and decompose. Nickel and other metals dissolve and then reprecipitate at different depths according to the solubility of their hydroxides. Thus, iron and aluminium are present in the upper zones, while magnesium is concentrated at the base of the ore body. Nickel and cobalt are found in the middle zones [2,30]. Limonitic laterites are iron-rich oxide ores with iron content higher than 40 wt%, magnesium content is up to 2 wt%, while nickel and cobalt contents are about 1.4 wt% and 0.15 wt%, respectively. These ores are commonly processed by pressure acid leaching. Saprolitic laterites are magnesium-rich ores with magnesium content in the range of wt%, iron content is between 10 and 25 wt%, while nickel and cobalt contents are about 2.4 wt% and 0.05 wt%, respectively. Although saprolites are richer in nickel, the high magnesium content results in higher sulphuric acid consumption, which renders the PAL process less economical. However, an optimal blend of limonites and 6

24 saprolites may form a high-grade feed that at the same time has an acceptable acid consumption [31,32]. Pressure acid leaching is performed in acid resistant autoclaves operating at o C. During this process, goethite and gibbsite dissolution is immediately followed by iron and aluminium precipitation mainly in the form of hematite and alunite, respectively, while nickel and cobalt dissolve and remain in the aqueous phase as sulphates. As a result, it is possible to attain a high [Ni+Co]/[Fe+Al] ratio in the resulting leaching solution. The process chemistry can be illustrated by the following reactions [3,32-34]. Goethite dissolution can be described by: FeOOH (s) + 3H + = Fe H 2 O (2.1) Trivalent iron hydrolyzes and precipitates as hematite: 2Fe H 2 O = Fe 2 O 3(s) + 6H + (2.2) During ore preheating, gibbsite, Al(OH) 3, transforms to boehmite, AlOOH, which dissolves according to the following reaction: AlOOH (s) + 3H + = Al H 2 O (2.3) Aluminium precipitates as hydronium alunite: 3Al 3+ +2SO H 2 O = H 3 OAl 3 (SO 4 ) 2 (OH) 6(s) + 5H + (2.4) In the presence of saline water, iron dissolution is not considerably affected. On the other hand, aluminium dissolution is much lower in these conditions because of the precipitation of sodium alunite according to the following equation: 3Al 3+ +2SO H 2 O + Na + = NaAl 3 (SO 4 ) 2 (OH) 6(s) + 6H + (2.5) 7

25 Sodium alunite is reported to be much less soluble than hydronium alunite and is also more prone to scaling the reactor [2,35]. The leaching of nickel and cobalt from goethite can be described as: NiO (s) + 2H + = Ni 2+ + H 2 O (2.6) CoO (s) + 2H + = Co 2+ + H 2 O (2.7) Nickel and cobalt may precipitate at low acidity as monohydrate sulphate salts by: Ni 2+ + SO H 2 O = NiSO 4 H 2 O (s) (2.8) Co 2+ + SO H 2 O = CoSO 4 H 2 O (s) (2.9) Magnesium dissolution from serpentine occurs according to the following reaction: Mg 3 Si 2 O 5 (OH) 4(s) + 6H + = 3Mg SiO 2(s) + 5H 2 O (2.10) Magnesium may also precipitate if its solubility is exceeded by: Mg 2+ + SO H 2 O = MgSO 4 H 2 O (s) (2.11) 2.2 Low-temperature ph measurement Definition of ph The International Union of Pure and Applied Chemistry (IUPAC) formally defines ph in terms of the activity of hydrogen ions in solution [36,37]: ph = log a = log( m ) (2.12) + + γ + H H H 8

26 where a + represents the activity and γ + H H designates the activity coefficient of the hydrogen ion at the molality m + H. However, since this definition includes a single ion quantity, the activity of the hydrogen ion, which cannot be measured by any thermodynamically valid method, Equation (2.12) can be only a notional definition Primary method of measurement and primary ph standards The primary method for measuring ph makes use of a cell without transference (i.e. without liquid junction) known as the Harned cell. The procedure used to assign ph values to primary standards is based on a unit of acidity referred to as the acidity function, p( a γ ) = log( m γ ). In contrast to ph, this quantity is physically γ H Cl H H Cl defined [36]. There are four steps in the assignment of the standard values [34]. First, the acidity function is determined for at least three molalities of added chloride by measuring the electromotive force, E, of the Harned cell: Pt H 2 buffer S, Cl - AgCl Ag (2.13) The cell reaction is given by: ½ H 2 (g) + AgCl(s) = Ag(s) + H + (aq) + Cl - (aq) (2.14) By applying the Nernst equation under p H2 = 1 bar: o E = E 2.303RT / F)log( a m ) ) (2.15) ( + γ H Cl Cl the acidity function is obtained as: p( a log m (2.16) H o + γ ) = log( a + γ ) = ( E E ) /(2.303RT / F) + Cl H Cl Cl where E o is the standard potential of the cell and therefore that of the Ag/AgCl electrode. 9

27 only HCl: Second, the standard potential E o is determined from a Harned cell containing Pt H 2 HCl AgCl Ag (2.17) By applying the Nernst equation to the above cell reaction, E o is obtained as: o E = E + (2.303RT / F)log( m ) (2.18) 2 HClγ ± HCl where γ ± HCl is the mean activity coefficient of HCl. The value of the mean activity coefficient of HCl is either taken from the literature or determined by extrapolation based on the Debye-Hückel theory. Third, log( a is evaluated by linear extrapolation of Equation (2.16) to o + γ ) H Cl zero chloride molality, where the superscript "o" denotes zero chloride molality. Finally, the ph of a primary standard, ph(ps), is calculated as: ph(ps) = log( a (2.19) o + γ ) + log γ H Cl Cl where o γ is the activity coefficient of chloride ion at zero chloride, which is calculated Cl using the Bates-Guggenheim convention: o 2 1/ 2 log γ 1/ = AI /( I ) (2.20) Cl where A is the Debye-Hückel parameter. In total, six buffer solutions are accepted by IUPAC as primary ph standards in the temperature range 0-50 o C [36] Secondary methods of measurement and secondary ph standards Secondary methods for measuring ph involve cells with transference that are more convenient than the Harned cell, such as cells with salt bridges, concentration cells 10

28 and glass combination electrodes. Cells with transference, however, contain liquid junctions and thus have significantly higher uncertainties associated with the results [36]. Secondary ph standards are the standards whose ph value is assigned by comparison with primary ph standards using secondary methods of measurement as well as the standards whose ph value is determined in the Harned cell but do not meet all the criteria of primary ph standards, such as reproducibility, stability, buffer capacity and ease of preparation [36] Glass electrode cells and their calibration Practical low temperature ph measurements are routinely performed using cells comprising a glass electrode and a Ag/AgCl reference electrode [36]: Ag AgCl KCl ( 3.5 mol dm -3 ) solution [ph(s) or ph(x)] glass electrode (2.21) where S denotes a standard solution, and X a test solution (to be measured). Glass electrode cells may yield a sub-nernstian slope and therefore require calibration. In practice, glass electrode cells are usually calibrated by two-point calibration using two standard solutions. The ph values of the two standard solutions, ph(s 2 ) and ph(s 1 ), should be chosen so that they bracket the ph value of the test solution. The practical slope, k, is obtained from: k = [E(S 2 ) E(S 1 )]/[ph(s 2 ) ph(s 1 )] (2.22) where E(S 2 ) and E(S 1 ) are actually the potentials of cell (2.10) measured for standards 1 and 2, respectively. If the potential of cell (2.10) measured for the test solution is E(X), the ph of the test solution, ph(x), is evaluated by: ph(x) = ph(s 1 ) - [E(X) E(S 1 )]/k (2.23) 11

29 2.3 High-temperature potentiometric ph measurement techniques Reference electrodes for high-temperature measurements A reference electrode must have a stable and reproducible potential during the course of measurement. This is a sufficient requirement for non-thermodynamic reference electrodes [38-41]. A great deal of effort has been placed on developing a reliable reference electrode for use in high-temperature aqueous solutions. In general, the reference electrodes that have been used in high-temperature studies can be divided into two categories: (a) internal reference electrodes working within the high-temperature environment and (b) external reference electrodes operating at room temperature, but linked to the hightemperature environment via a non-isothermal electrolyte bridge. Each approach has its own drawbacks and advantages [38-41]. Internal reference electrodes have two general problems. The first problem is the tendency for the solubility of the electroactive element to increase rapidly with increasing temperature. The second problem is hydrolysis of the reference element at elevated temperatures [38-41]. A number of electrochemical couples, such as Ag/AgCl, Ag/AgBr, Ag/Ag 2 SO 4, Hg/Hg 2 Cl 2, Hg/Hg 2 SO 4, Hg/HgO and Pb/PbSO 4, were tested as an internal high-temperature reference electrode [38,40,41]. Only the Ag/AgCl electrode, however, was found to be sufficiently stable for short periods of time at temperatures up to 400 o C [21,42,43]. The choice of a reference system is less important with external electrodes as the temperature is maintained at 25 o C. However, both pressure gradients (if the electrode is not pressure balanced) and temperature gradients develop across the electrolyte bridge. 12

30 These gradients give rise to streaming and thermal diffusion potentials (the Soret effect), both of which affect the measured potential difference. Since these effects are irreversible, external reference electrodes cannot be used as thermodynamic standards. However, they do produce a stable and reproducible potential, and may be calibrated against a suitable standard at various temperatures back to the SHE scale [38,39]. Macdonald and coworkers introduced the external pressure balanced reference electrode (EPBRE) in 1979 [44]. In this configuration, the internal solution (KCl) and the electroactive element (Ag/AgCl) are contained within a flexible PTFE tube so that volume changes on pressurization can be accommodated without significant flow through the isothermal liquid junction. Furthermore, the flexible PTFE tube is also intended to transmit pressure pulses into the internal solution, thereby preventing thermal diffusion and retaining the Soret initial state. However, this electrode is very difficult to maintain in the Soret initial state over the long term. A reproducibility of ±10-30 mv is achievable at temperatures up to 300 o C with an EPBRE [23,45]. Lvov and coworkers developed a flow-through external pressure balanced reference electrode (FTEPBRE) in 1998 [45]. In this design, the reference solution is pumped through the electrolyte bridge at a constant rate, which suppresses thermal diffusion and hence maintains the Soret initial state. Therefore, the thermal diffusion potential remains constant and can be either calculated or experimentally determined. In this way, the electrode potential can be evaluated with respect to the SHE scale. However, this flow-through technique gives rise to a streaming potential that should also be taken into account. The FTEPBRE has been tested in potentiometric measurements at temperatures up to about 400 o C along with a flow-through hydrogen indicator electrode 13

31 and/or a flow-through yttria-stabilized zirconia indicator electrode. The potential of the FTEPBRE was found to be stable within 1-3 mv High-temperature indicator electrodes The indicator electrode must have a stable and reproducible potential during measurement and should respond in a Nernstian manner to changes in ph of the hightemperature aqueous solution [40,41]. A number of electrode systems have been developed for measuring ph at high temperatures: hydrogen electrodes (including palladium hydride electrodes), metal-metal oxide electrodes, the glass electrode, and the yttria-stabilized zirconia (YSZ) electrode. However, only the Pt/H 2 electrode and the YSZ electrode containing a Hg/HgO internal reference element were found to behave in a Nernstian manner at elevated temperatures up to 400 o C [40,41] Hydrogen electrode The hydrogen electrode is adopted as the ultimate standard for the determination of ph values, with which all other electrodes are compared [40,41,46,47]. The hydrogen electrode consists of a platinum wire or strip exposed to a current of hydrogen and partly immersed in an acid solution. The platinum surface is covered with a layer of fine-grained platinum (platinum black) that catalyzes the reaction: H + (aq) + e - = ½ H 2(g) (2.24) and hence permits establishment of equilibrium between the hydrogen gas adsorbed by the metal and the hydrogen ions in solution. 14

32 The potential of the hydrogen electrode is given by the Nernst equation in the form: o 2.303RT 2.303RT E = E + ph log f H /H H (2.25) 2 2 F 2F where f H2 represents the fugacity of hydrogen and E o + H /H 2 is the standard potential of the hydrogen electrode which is by convention equal to zero at all temperatures. The electrode thus responds to both ph and f H. 2 As seen, the primary method of ph measurement is based on a cell without transference consisting of a hydrogen electrode and a silver-silver chloride reference electrode [36]: Pt H 2 H +, Cl - AgCl Ag (2.26) Rouchoudhury and Bonilla were among the first to test the hydrogen electrode above 100 o C in 1956 [48]. They measured the electromotive force of cell (2.26) in a dilute HCl solution at temperatures to 250 o C. The experimental and theoretical values were in excellent agreement up to 150 o C, but differed markedly at higher temperatures. One explanation was that sites existed on the Ag/AgCl electrode which acted as local hydrogen electrodes. Another explanation was reduction of part of AgCl by hydrogen, which would alter the acid concentration [38,39,49]: AgCl (s) + ½ H 2(g) = Ag (s) + H + (aq) + Cl - (aq) (2.27) Lietzke and coworkers employed the hydrogen electrode in cells of type (2.26) at temperatures to 275 o C [49-52]. A schematic representation of this cell is given in Figure 2.1. The Ag/AgCl electrode was placed in a silica tube to isolate it from hydrogen and electrical contact with the body of the solution was made via a silica frit. The hydrogen 15

33 electrode was made to float at the surface of the solution in order to see both the solution phase and the gas phase. The cells performed well to 200 o C. However, at temperatures of 250 o C and above, it was difficult to keep hydrogen away from the Ag/AgCl electrode, which resulted in more scatter in the measured potentials. Because of the absence of a reliable reference electrode above about 200 o C, Mesmer and coworkers developed a hydrogen-electrode concentration cell (HECC) in 1970 [5]. The HECC employs hydrogen electrodes as both the test and reference electrodes. Cells of this type have since been used for potentiometric measurements at temperatures to 300 o C at Oak Ridge National Laboratory [7,8]. A schematic representation of the hydrogen-electrode concentration cell is given in Figure 2.2. The HECC has two concentric internal compartments made from Teflon. The small inner compartment contains the reference solution, usually a strong acid or base, while the large compartment contains the test solution in which a titration can be performed. The liquid junction is formed through a small plug of porous Teflon in the bottom of the reference compartment. An important feature of the HECC is that, since the inner and outer compartments are connected via a small hole, the fugacity of hydrogen in both compartments is the same. The cell potential is therefore independent of the fugacity of hydrogen in the system, in contrast to the single hydrogen electrode systems. Thus, for the study of a weak acid, HA, the HECC can be represented as: Pt H 2 HCl, NaCl HA, NaA, NaCl H 2 Pt (2.28) reference test Both solutions are maintained at the same ionic strength by addition of a supporting electrolyte such as NaCl, so that m NaCl /m HCl ~ m NaC l/m NaA ~ 100. The supporting 16

34 electrolyte is necessary for the assumption that the activity coefficients of the minor components are the same in both compartments, while the value of the liquid junction potential, E LJ, is minimized and can be calculated using the Henderson equation. The cell potential is given by: RT m + H E,ref = ln + E LJ (2.29) F m + H Lvov and coworkers employed a flow-through hydrogen electrode to measure ph in subcritical and supercritical aqueous systems at temperatures from 25 to 400 o C in 1999 [53]. It was found that the theoretically calculated ph values agree with the experimentally derived ph data to within ±0.03 logarithmic units. The main disadvantage of the hydrogen electrode is that its reversibility can be inhibited by certain poisons. These poisons are substances that, in traces, interfere with the proper operation of the electrode and may be divided into three categories [46,47,54]: (a) Oxidizing agents, which are reducible at the electrode surface and thus deplete the concentration of molecular hydrogen in solution close to the electrode. These agents include gaseous oxygen, oxygen-containing anions such as chromate/dichromate, permanganate, nitrate and the higher valence ions of transition metals such as Fe 3+. (b) Substances that can be reduced and deposited at the electrode surface. These substances involve cations of metals more noble than hydrogen such as Ag, Hg and Cu. (c) Substances that are preferentially adsorbed onto the active centres and paralyse catalytic reaction, such as H 2 S, CN - and arsenic compounds. 17

35 Figure 2.1. Schematic representation of the cell with a floating hydrogen electrode (after [49]). 18

36 Figure 2.2. Schematic representation of the hydrogen-electrode concentration cell (after [8]). 19

37 Yttria-stabilized zirconia ph sensor Niedrach was the first to discover that yttria-stabilized zirconia can function as a ph sensor in 1980 [9,10]. Initial YSZ ph sensors were similar to the glass electrode. That is, a buffered saline solution (e.g., 0.1 mol L -1 NaCl, ph 7) in contact with a AgCl-coated Ag wire served as the internal element inside a YSZ tube. The YSZ sensor was tested over a ph range 3.3 ( mol kg -1 H 2 SO 4 ) to 8.2 (0.001 mol kg -1 NaOH) at 285 o C. The response of the YSZ sensor was compared to that of an oxygen electrode formed at a platinized-pt wire in solutions saturated with air. A very good linear correlation was found, with the response of the YSZ electrode being around 95% of that of the oxygen electrode. The responses of YSZ electrodes were also compared with the theoretical value calculated for the transition between mol kg -1 H 2 SO 4 and mol kg -1 NaOH, and an approximate response of 94% of theoretical was found. The YSZ electrode was also shown to be insensitive to changes in redox potential. Tsuruta and Macdonald also investigated a YSZ electrode containing an internal Ag/AgCl reference element in 1982 [55]. However, the internal element was located at ambient temperature in order to avoid thermal degradation. The YSZ electrode was coupled to an external pressure balanced Ag/AgCl reference electrode. The schematic representation of the ph measuring system is given in Figure 2.3. Since both electrodes were filled with identical solutions (0.1 mol kg -1 KCl mol kg -1 B(OH) mol kg -1 KOH), the thermal diffusion potentials and the thermoelectric potentials on both sides cancelled. As a result, the measured potential needed to be only corrected for the liquid junction potential between the test solution and the reference solution. The liquid junction potential was calculated using the Henderson equation. Experiments were 20

38 performed in the temperature range o C in solutions with calculated ph values ranging from about 3 to 9 at 275 o C. The potential of the YSZ sensor and the potential of an in-situ hydrogen electrode were simultaneously measured against the reference electrode, so that the ph response of the YSZ electrode could be compared to that of the hydrogen electrode. The YSZ electrode was found to yield a near-nernstain ph response above 200 o C, with some deviation being observed in the ph range 3 to 5. Subsequent designs of the yttria-stabilized zirconia sensor featured dry internal reference elements, that is, a mix of a metal and its oxide or two oxides of the same metal in different valence states, such as Cu/Cu 2 O [12,13,15,16,56,57], Ag/Ag 2 O [15] and Hg/HgO [11,15,17-21]. Such solid-state designs offered several advantages over the earlier aqueous system, most notably, simplified fabrication and better stability and reproducibility. Niedrach tested the performance of YSZ sensors containing the Cu/Cu 2 O internal element at 285 o C in 1982 [56]. The potential of the YSZ sensor was first measured against the hydrogen electrode in water that was equilibrated with nitrogen gas containing 10% hydrogen before entering the autoclave. The average difference between the measured potential and theoretical value was as low as 2 mv. Two tests about 40 days long were also carried out to demonstrate long-term performance. In this case, the potential of the YSZ sensor was measured versus an internal Ag/AgCl electrode over a ph range from 3.3 ( mol kg -1 H 2 SO 4 ) to 8.2 (0.001 mol kg -1 NaOH). The potential of the YSZ sensor was found to vary linearly with ph. 21