SOLUTION CHEMISTRY OF SOME DICARBOXYLATE SALTS OF RELEVANCE TO THE BAYER PROCESS

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1 SOLUTION CHEMISTRY OF SOME DICARBOXYLATE SALTS OF RELEVANCE TO THE BAYER PROCESS Andrew John Tromans B.Sc. Hons (Edith Cowan University) This thesis is presented for the degree of Doctor of Philosophy of Murdoch University Western Australia November 2003

2 i I declare that this thesis is my own account of my research and contains as its main content work which has not previously been submitted for a degree at any tertiary education institution. Andrew John Tromans

3 ii Abstract This thesis deals with certain aspects of the solution chemistry of the simple dicarboxylate anions: oxalate, malonate and succinate, up to high concentrations. These ions are either significant impurities in the concentrated alkaline aluminate solutions used in the Bayer process for the purification of alumina, or are useful models for degraded organic matter in industrial Bayer liquors. Such impurities are known to have important effects on the operation of the Bayer process. To develop a better understanding of the speciation of oxalate (the major organic impurity in Bayer liquors) in concentrated electrolyte solutions, the formation constant (Logβ) of the extremely weak ion pair formed between sodium (Na + ) and oxalate (Ox 2 ) ions was determined at 25 o C as a function if ionic strength in TMACl media by titration using a Na + ion selective electrode. Attempts to measure this constant in CsCl media were unsuccessful probably because of competition for Ox 2 by Cs +. Aqueous solutions of sodium malonate (Na 2 Mal) and sodium succinate (Na 2 Suc) were studied up to high (saturation) concentrations at 25 o C by dielectric relaxation spectroscopy (DRS) over the approximate frequency range 0.1 ν/ghz 89. To complement a previous study of Na 2 Ox, formation constants of the Na + -dicarboxylate ion pairs were determined and they were shown to be of the solvent-shared type. Both the Mal 2 and Suc 2 ions, in contrast to Ox 2, were also shown to possess large secondary hydration shells

4 iii Apparent molal volumes (V φ ) and heat capacities at constant pressure (C pφ ) of aqueous solutions of Na 2 Ox, Na 2 Suc, Na 2 Mal and K 2 Ox were determined at 25 o C up to their saturation limits using vibrating tube densitometry and flow calorimetry. These data were fitted using a Pitzer model. The adherence of V φ and C pφ of various Na + and K + salts to Young s rule was examined up to high concentrations using the present and literature data. Young s rule was then used to estimate hypothetical values of C pφ and V φ for the sparingly soluble Na 2 Ox at high ionic strengths, which are required for the thermodynamic modelling of Bayer liquors. The solubility of Na 2 Ox in various concentrated electrolytes was measured, at temperatures from 25 o C to 70 o C in media both with (NaCl, NaClO 4, NaOH) and without a common ion (KCl, CsCl, TMACl). The common ion effect was found to dominate the solubility of Na 2 Ox. The solubility of calcium oxalate monohydrate (CaOx H 2 O) was also determined. The solubilities of both Na 2 Ox and CaOx H 2 O in media without a common ion increased with increasing electrolyte concentration, except in TMACl media, where they decreased. The solubility of Na 2 Ox was modelled using a Pitzer model assuming the Pitzer parameters for Na 2 SO 4 and minimising the free energy of the system. The data were modelled successfully over the full concentration and temperature range of all the electrolytes, including ternary (mixed electrolyte) solutions.

5 iv Publications The following publications have arisen from work completed by the candidate for the present thesis. A. Tromans, E. Königsberger, G. T. Hefter, P. M. May, Solubility of Sodium Oxalate in Concentrated Electrolyte solutions, 10 th International Symposium on Solubility Phenomena, Varna, Bulgaria, (2002).

6 v Acknowledgements I would like to sincerely thank my supervisors, Associate Professors Glenn Hefter and Peter May for their continuous guidance, advice, encouragement and help throughout this entire project. I would also like to express my gratitude to Dr Richard Buchner, Institute für Physikalische und Theoretische Chemie, Universität Regensburg, Germany for his role in modeling and interpreting my dielectric relaxation spectra. I would like to sincerely thank Dr Erich Königsberger for his expertise in developing Pitzer models for my experimental data. His advice throughout this project is also greatly appreciated. I would also like to thank Dr Pal Sipos for his advice and encouragement through the initial, crucial stages of this project. I am further indebted to the following individuals for their help in various respects. Mr Simon Bevis, for his continual assistance throughout this project. Mr Doug Clarke, Mr Tom Osborne and Mr Andrew Foreman of the Chemistry Department technical staff. Mr Kleber Claux, Mr Ernie Etherington and Mr John Snowball of the Murdoch University Mechanical and Electrical Workshops. I am especially thankful for the love and support of my family and friends and in particular my parents Edmund and Helen. Finally, I would like to thank the A. J. Parker CRC for Hydrometallurgy for their financial assistance in the form of a research scholarship.

7 vi Contents DECLARATION ABSTRACT PUBLICATIONS ACKNOWLEDGEMENTS CONTENTS LIST OF TABLES LIST OF FIGURES ABBREVIATIONS AND SYMBOLS i ii iv v vi xii xiv xvi CHAPTER ONE INTRODUCTION THE BAYER PROCESS Digestion Precipitation Calcination Re-causticisation Smelting 5

8 vii 1.2 BAYER LIQUOR IMPURITIES Sources of impurities Effects of Bayer liquor organic impurities CHEMISTRY OF BAYER LIQUOR IMPURITIES MODELLING THE SOLUTION CHEMISTRY OF 12 DICARBOXYLATE SALTS 1.5 OBJECTIVES OF THE PRESENT RESEARCH 13 CHAPTER TWO POTENTIOMETRY INTRODUCTION Inter-ionic attraction theory The Debye-Huckel theory Extentions to the Debye-Huckel limiting law Ion pairing Ion pairing studies on dicarboxylate ions Potentiometry Potentiometric determination of formation 32 constants

9 viii 2.2 EXPERIMENTAL METHOD Reagents and glassware Electrodes Titration apparatus Titration procedure Data analysis CHARACTERISATION OF THE BEHAVIOUR OF 40 THE NaISE Calibration Determination of sodium impurities SODIUM ION-ASSOCIATION CONSTANTS FROM 44 NaISE TITRATIONS Titrations in CsCl media Titrations in TMACl media Standard value of β(naox ) 46 CHAPTER THREE DIELECTRIC RELAXATION SPECTROSCOPY INTRODUCTION Structure of electrolyte solutions 54

10 ix Dielectric relaxation spectroscopy Information obtained from DRS Application of DRS to di-carboxylic acids EXPERIMENTAL Solution preparation Instrumentation Calibration of VNA Measurement procedure and data analysis Raman spectroscopy RESULTS Conductivity of Na 2 Mal and Na 2 Suc solutions Modeling of relaxation processes General features of the spectra Ion hydration Quantitation of ion pairing Kinetics of ion pairing 86 CHAPTER FOUR PARTIAL MOLAL PROPERTIES INTRODUCTION Partial molal properties of a solution 91

11 x Heat capacity Partial molal volumes EXPERIMENTAL Reagents Densitometry Flow microcalorimetry RESULTS Calorimeter calibration Calorimeter asymmetry Apparent molal volumes Apparent molal heat capacities MODELING PARTIAL MOLAL PROPERTIES Redlich-Rosenfeld-Meyer equation Ion Interaction approach VALUES OF C φ o AND V φ o PARTIAL MOLAL PROPERTIES AT HIGH IONIC 123 STRENGTH

12 xi CHAPTER FIVE: SOLUBILITY INTRODUCTION Thermodynamic principles Practical considerations EXPERIMENTAL Reagents Methods Testing of titrimetric analysis RESULTS AND DISCUSSION Na 2 Ox solubility in concentrated electrolyte media CaOx solubility in concentrated electrolyte media Modeling the solubilities using the Pitzer 163 approach CHAPTER SIX: CONCLUSIONS 169 REFERENCES 172

13 xii List of Tables Table 2.1. Association constants of the sodium oxalate ion pair. 47 Table 2.2. Calculation of β NaOx - from titration data with no allowance 48 for Na + impurities in 1M CsCl at 25 o C Table 2.3. Calculation of β NaOx - from titration data, taking into 49 account Na + impurity in 1M CsCl at 25 o C. Table 2.4. Calculation of β NaOx from titration data, taking into 50 account Na + impurity in 1 M TMACl at 25 o C. Table 3.1. Concentration, c, and average effective conductivity, κ e, 64 of aqueous Na 2 Mal and Na 2 Suc solutions at 25 o C. Table 3.2. Concentration, c, dielectric relaxation parameters of a 4D 68 fit and variance of the fit, σ 2, of aqueous Na 2 Suc at 25 o C. Table 3.3. Concentration, c, dielectric relaxation parameters of a 4D 69 fit and variance of the fit, σ 2, of aqueous Na 2 Mal at 25 o C. Table 3.4. Refractive index of aqueous Na 2 Mal and Na 2 Suc solution 83 at 25 o C. Table 3.5. Dipole moment of ion pair models 83 Table 3.6. Formation constants (β) of Na 2 Suc and Na 2 Mal ion pairs 87 calculated for each ion pair model. Table 4.1. Experimental densities, ρ, heat capacities, Cp, and apparent 104 molar heat capacities, C pφ, of aqueous solutions of NaCl at 25 o C. Table 4.2. Apparent molal volume of electrolytes at 25 o C 107 Table 4.3 Apparent molal heat capacities of electrolytes at 25 o C. 110 Table 4.4. Pitzer model parameters for heat capacities and molal 121 volumes (25 o C) derived from this work. Table 4.5. V o φ for electrolytes at 25 o C. Extrapolated from experimental 122 data using the R-R-M model and Pitzer model. Table 4.6. C o p,φ for electrolytes at 25 o C. Extrapolated from experimental 122 data using the R-R-M model and Pitzer model. Table 5.1. Solubility of Na 2 Ox in Na + -containing electrolyte media 146 at ± 0.02 o C.

14 xiii Table 5.2. Solubility of Na 2 Ox in Na + -containing electrolyte media 148 at ± 0.02 o C. Table 5.3. Solubility of Na 2 Ox in Na + -containing electrolyte media 150 at ± 0.02 o C. Table 5.4 Solubility of Na 2 Ox in non-na + electrolyte media at ± 0.02 o C. Table 5.5 Solubility of Na 2 Ox in CsCl media at and o C 156 (± 0.02 o C) Table 5.6. Solubility of CaOx.H 2 O at ± 0.02 o C. 159 Table 5.7. Solubility of CaOx H 2 O at ± 0.02 o C. 160 Table 5.8 Solubility of CaOx.H 2 O at ± 0.02 o C. 161

15 xiv List of Figures Figure 2.1. The response of the NaISE to Na + concentration in the 41 presence of high concentrations of TMACl background electrolyte Figure 2.2. Calibration curve for NaISE in TMACl media 44 Figure 2.3. Values of logβ NaOx - in TMACl media at 25 o C. 51 Figure 3.1. Average effective conductivity, κ e, determined from the 65 total loss spectrum, η (ν) Figure 3.2. Dielectric dispersion, ε (ν), and loss ε (ν) of Na 2 Suc 70 solution in water at 25 o C. Figure 3.3. Dielectric dispersion, ε (ν), and loss ε (ν) of Na 2 Mal 71 solution in water at 25 o C. Figure 3.4. Relaxation processes for Na 2 Suc at two sample 72 concentrations. Figure 3.5. Relaxation processes for Na 2 Mal. 73 Figure 3.6. Measured limiting permittivities of the ion-pair process, 73 ε and ε 2, and of the slow water relaxation, ε 2 and ε 3, of Na 2 Suc(aq) at 25 C. Figure 3.7. Measured limiting permittivities of the ion-pair process, 74 ε and ε 2, and of the slow water relaxation, ε 2 and ε 3, of Na 2 Mal(aq) at 25 C. Figure 3.8. τ for relaxation processes in Na 2 Suc aqueous solutions. 77 Figure 3.9. τ for relaxation processes in Na 2 Mal aqueous solutions. 78 Figure Fraction of slow water, c 2 /c s, of Na 2 Mal(aq) (a) and 80 Na 2 Suc(aq) (b) at 25 C Figure Effective solvation numbers, Z ib, of Na 2 Mal(aq) (a) and 81 Na 2 Suc(aq) (b) at 25 C assuming slip boundary conditions for the kinetic depolarization effect. Figure 3.12 Refractive index of Na 2 Mal and Na 2 Suc solutions at 25 o C. 84 Figure 3.13 Possible structural isomers of NaMal and NaSuc ion pairs. 85 Figure Association constants (β) for [NaMal] - (aq) ( ) and [NaSuc] - 86 (aq) (, multiplied by 0.1) as a function of the ionic strength at 25 C assuming SIP Figure 3.15 Values of (log β 0 ) 2 for Na 2 Suc obtained from S 1 with 88 Eq. 3.8 for tested ion-pair structures versus the reciprocal ion-pair dipole moment, µ -1 IP.

16 Figure 4.1. Apparent molar heat capacity of NaCl solutions at 25 o C. 105 Figure 4.2. Difference between first leg, C p (A) and second leg, 107 C p (B), heat capacities Figure 4.3 Apparent molal volumes of electrolytes at 25 o C; 109 Figure 4.4 Apparent molal heat capacities of electrolytes at 25 o C 113 Figure 4.5. Differences between apparent molal volumes of Na + and K salts containing various common anions at 25 o C Figure 4.6. Differences between apparent molal heat capacities of Na and K + salts of various common anions at 25 o C Figure 4.7. Estimation of apparent molal heat capacity of Na 2 Ox at 129 high ionic strengths. Figure 4.8. Estimation of apparent molal volume of Na 2 Ox to 130 high ionic strengths. Figure 5.1. Na 2 Ox solubility in Na + containing media at 25 o C 144 Figure 5.2. Na 2 Ox solubility in NaOH at various temperatures. 144 Figure 5.3 Measured and calculated solubility of Na 2 Ox 145 Figure 5.4. Na 2 Ox solubility in non Na + media at 25 o C 152 Figure 5.5. Solubility of Na 2 Ox in CsCl medium at various 153 temperatures. Figure 5.6. CaOx.H 2 O solubility in electrolyte media at 25 o C. 158 Figure 5.7. Solubility of CaOx.H 2 O in CsCl media at various 158 temperatures. Figure 5.8 Pitzer model of Na 2 Ox solubility in NaCl media 167 Figure 5.9 Pitzer model of Na 2 Ox solubility in NaOH media. 167 Figure 5.10 Pitzer model of Na 2 Ox solubility in NaOH / NaCl, 168 I = 5 M media. xv

17 xvi Abbreviations and Symbols ƒ Ι field factor Infinite 4D Four-Debye relaxation model a Activity A Debye Huckel constant α, β Relaxation time distribution parameters (Eq 3.4) A C Debye-Huckel constant for heat capacity A φ Ion Interaction Debye-Huckel constant Debye-Huckel constant for enthalpy A H α Ι A V polarisability Molarity based Debye-Huckel constant for volumes b Parameter in Eq 4.14 B, C Parameters in Eq 4.17 c Concentration (Mol L -1 ) C A Concentration of standard solution added (Eq 2.26) CaOx Calcium oxalate (CaC 2 O 4 nh 2 O) Heat capacity at constant pressure C p C p, φ D ε E ε γ (i,j) ε (ν) ε (ν) ε ε 0 E o Apparent molal heat capacity at constant pressure Dielectric constant Static permittivity. Cell potential SIT interaction parameter between aqueous species i and j Dielectric dispersion. Dielectric loss. Infinite frequency permittivity Permittivity of a vacuum Formal cell potential F Faraday s constant f(i) Debye-Huckel term (Eq 4.15) F * j (ν) Complex relaxation function F OR (t) Time dependant autocorrelation function γ Activity coefficient G Gibbs energy

18 xvii g Γ γ ± GHz g i H η (ν) η κ (ν) I J k κ k K κ E K IP κ max K s L φ Grams Relative complex reflection coefficient Mean ionic activity coefficient GigaHertz dipole correlation factor Enthalpy Total dielectric loss Conductivity-related ohmic loss Ionic strength Joule Boltzman s constant Conductivity Density proportionality constant Kelvin Effective conductivity Ion pairing association constant. Maximum conductivity Solubility product constant Relative enthalpy µ Chemical potential m Molality (mol kg -1 ) M Molecular weight m ± Mean molality Mal Malonate anion (C 3 H 2 O 2 4 ) µ χ, α Parameters in Eq 3.8 µ ι dipole moment of species i µ ijk Tertiary ion interaction parameter ν Frequency of electric field change. n General number n Number of moles N Total number of titration points (Eq 2.25). N A Avogadro s number n e Number of parameters to be optimized. (Eq 2.25) ν ι Stoichiometric number of species i n p Total number of electrodes. (Eq 2.25) Ox Oxalate anion (C 2 O 2 4 ) P Power (Chapter Four) P Pressure P(t) Time dependant polarization created by an alternating electric field P α eq P µ (t) Polarization arising from an induced dipole Polarization due to permanent dipole

19 xviii ps ρ R R ρ ο S S σ 2 Pico seconds Charge density (Chapter two) Gas constant Molar refractivity Density of pure solvent Analytical solubility Relaxation process amplitude Variance of fit Suc Succinate anion (C 4 H 4 O 2 4 ) S V Empirical slope for Masson equation τ Period of oscillation τ Relaxation time T Temperature (K) U Objective function V Volume V A Volume of standard solution added (Eq 2.26) V φ Apparent molal volume V o Initial volume of test solution (Eq 2.26). ω Valency factor in Redlich-Rosenfeld-Meyer equation w nq Weighting parameter (Eq 2.25) W w Mass in kilograms of the solvent (water) x Mole fraction X o Property at infinite dilution ψ Potential difference between reference ion and ionic cloud. ψ Ternary short range ionic interactions (Chapter Five) y nq Total concentration of electrode ion q / potential of electrode q at n th point. (Eq 2.25) z Stoichiometric atomic charge Z ib Hydration number β Formation constant of ion pair β o Formation constant at infinite dilution. β β, C β Parameters in Equation 2.11 θ Unsymmetrical mixing term in Pitzer equation λ ij Binary ion interaction parameter µ o Standard chemical potential ρ Solution density Φ Binary short range ionic interactions Partial molal heat capacity at constant pressure C p V Partial molal volume 25 n D Refractive index

THERMODYNAMIC AND RELATED STUDIES OF AQUEOUS COPPER(II) SULFATE SOLUTIONS

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