Preface xi
Notations and Symbols xv
Part 1. Ionic Equilibria 1
Chapter 1. Dissociation of Electrolytes in Solution 3
1.1. Strong electrolytes – weak electrolytes 3
1.1.1. Dissolution 3
1.1.2. Solvolysis 4
1.1.3. Melting 4
1.2. Mean concentration and mean activity coefficient of ions 5
1.3. Dissociation coefficient of a weak electrolyte 6
1.4. Conduction of electrical current by electrolytes 9
1.4.1. Transport numbers and electrical conductivity of an electrolyte 9
1.4.2. Equivalent conductivity and limiting equivalent conductivity of an electrolyte 10
1.4.3. Ionic mobility 11
1.4.4. Relation between equivalent conductivity and mobility – Kohlrausch’s law 14
1.4.5. Apparent dissociation coefficient and equivalent conductivity 16
1.4.6. Variations of equivalent conductivities with the concentrations 16
1.5. Determination of the dissociation coefficient 20
1.5.1. Determination of the dissociation coefficient by the cryometric method 21
1.5.2. Determination of the dissociation coefficient on the basis of the conductivity values 22
1.6. Determination of the number of ions produced by dissociation 23
1.6.1. Use of limiting molar conductivity 23
1.6.2. Use of cryometry 24
1.7. Thermodynamic values relative to the ions 27
1.7.1. The standard molar Gibbs energy of formation of an ion 27
1.7.2. Standard enthalpy of formation of ions 29
1.7.3. Absolute standard molar entropy of an ion 29
1.7.4. Determination of the mean activity of a weak electrolyte on the basis of the dissociation equilibrium 30
Chapter 2. Solvents and Solvation 31
2.1. Solvents 31
2.2. Solvation and structure of the solvated ion 33
2.3. Thermodynamics of solvation 35
2.3.1. Thermodynamic values of solvation 36
2.3.2. Gibbs energy of salvation – Born’s model 37
2.4. Transfer of a solute from one solvent to another 44
2.5. Mean transfer activity coefficient of solvation of an electrolyte 48
2.6. Experimentally determining the transfer activity coefficient of solvation 49
2.6.1. Determining the activity coefficient of a molecular solute 50
2.6.2. Determination of the mean transfer activity coefficient of a strong electrolyte 51
2.6.3. Evaluation of the individual transfer activity coefficient of an ion 51
2.7. Relation between the constants of the same equilibrium achieved in two different solvents 55
2.7.1. General relation of solvent change on an equilibrium constant 55
2.7.2. Influence of the dielectric constant of the solvent on the equilibrium constant of an ionic reaction 56
Chapter 3. Acid/Base Equilibria 61
3.1. Definition of acids and bases and acid–base reactions 62
3.2. Ion product of an amphiprotic solvent 63
3.3. Relative strengths of acids and bases 64
3.3.1. Definition of the acidity constant of an acid 64
3.3.2. Protic activity in a solvent 67
3.4. Direction of acid–base reactions, and domain of predominance 69
3.5. Leveling effect of a solvent 71
3.6. Modeling of the strength of an acid 75
3.6.1. Model of the strength of an acid 75
3.6.2. Comparison of an acid’s behavior in two solvents 78
3.6.3. Construction of activity zones for solvents 81
3.7. Acidity functions and acidity scales 84
3.8. Applications of the acidity function 88
3.8.1. Measuring the pKa of an indicator 89
3.8.2. Measuring the ion products of solvents 89
3.9. Acidity in non-protic molecular solvents 91
3.10. Protolysis in ionic solvents (molten salts) 92
3.11. Other ionic exchanges in solution 93
3.11.1. Ionoscopy 93
3.11.2. Acidity in molten salts: definition given by Lux and Flood 94
3.12. Franklin and Gutmann’s solvo-acidity and solvo-basicity 96
3.12.1. Definition of solvo-acidity 96
3.12.2. Solvo-acidity in molecular solvents 96
3.12.3. Solvo-acidity in molten salts 98
3.13. Acidity as understood by Lewis 100
Chapter 4. Complexations and Redox Equilibria 101
4.1. Complexation reactions 101
4.1.1. Stability of complexes 101
4.1.2. Competition between two ligands on the same acceptor 106
4.1.3. Method for studying perfect complexes 108
4.1.4. Methods for studying imperfect complexes 110
4.1.5. Study of successive complexes 115
4.2. Redox reactions 117
4.2.1. Electronegativity – electronegativity scale 117
4.2.2. Degrees of oxidation 124
4.2.3. Definition of redox reactions 128
4.2.4. The two families of redox reactions 128
4.2.5. Dismutation and antidismutation 130
4.2.6. Redox reactions, and calculation of the stoichiometric numbers 131
4.2.7. Concept of a redox couple 132
Chapter 5. Precipitation Reactions and Equilibria 135
5.1. Solubility of electrolytes in water – solubility product 135
5.2. Influence of complex formation on the solubility of a salt 136
5.3. Application of the solubility product in determining the stability constant of complex ions . 137
5.4. Solution with multiple electrolytes at equilibrium with pure solid phases 138
5.4.1. Influence of a salt with non-common ions on the solubility of a salt 139
5.4.2. Influence of a salt with a common ion on the solubility of a salt 141
5.4.3. Crystallization phase diagram for a mixture of two salts in solution 141
5.4.4. Formation of double salts or chemical combinations in the solid state 142
5.4.5. Reciprocal quaternary systems – square diagrams 144
5.5. Electrolytic aqueous solution and solid solution 147
5.5.1. Thermodynamic equilibrium between a liquid ionic solution and a solid solution 147
5.5.2. Solubility product of a solid solution 150
5.6. Solubility and pH 155
5.6.1. Solubility and pH 155
5.6.2. Solubility of oxides in molten alkali hydroxides 156
5.6.3. Solubility in oxo-acids and oxo-bases (see section 3.12.2) 157
5.7. Calculation of equilibria in ionic solutions 158
Part 2. Electrochemical Thermodynamics 163
Chapter 6. Thermodynamics of the Electrode 165
6.1. Electrochemical systems 165
6.1.1. The electrochemical system 166
6.1.2. Electrochemical functions of state 167
6.1.3. Electrochemical potential 167
6.1.4. Gibbs–Duhem relation for electrochemical systems 169
6.1.5. Chemical system associated with an electrochemical system 170
6.1.6. General conditions of an equilibrium of an electrochemical system 171
6.2. The electrode 173
6.2.1. Definition and reaction of the electrode 173
6.2.2. Equilibrium of an insulated metal electrode – electrode absolute voltage 174
6.2.3. Voltage relative to a metal electrode – Nernst’s relation 175
6.2.4. Chemical and electrochemical Gibbs energy of the electrode reaction 178
6.2.5. Influence of pH on the electrode voltage 179
6.2.6. Influence of the solvent and of the dissolved species on the electrode voltage 181
6.2.7. Influence of temperature on the normal potentials 183
6.3. The different types of electrodes 184
6.3.1. Redox electrodes 184
6.3.2. Metal electrodes 189
6.3.3. Gas electrodes 192
6.4. Equilibrium of two ionic conductors in contact 193
6.4.1. Junction potential with a semi-permeable membrane 193
6.4.2. Junction potential of two electrolytes with a permeable membrane 194
6.5. Applications of Nernst’s relation to the study of various reactions 196
6.5.1. Prediction of redox reactions 196
6.5.2. Relations between the redox voltages of different systems of the same element 197
6.5.3. Predicting the dismutation and anti-dismutation reactions 201
6.5.4. Redox catalysis 202
6.6. Redox potential in a non-aqueous solvent 203
6.6.1. Scale of redox potential in a non-aqueous medium 203
6.6.2. Oxidation and reduction of the solvent 206
6.6.3. Influence of solvent on redox systems in a non-aqueous solvent 207
Chapter 7. Thermodynamics of Electrochemical Cells 209
7.1. Electrochemical chains – batteries and electrolyzer cells 209
7.2. Electrical voltage of an electrochemical cell 210
7.3. Cell reaction 212
7.4. Influence of temperature on the cell voltage; Gibbs–Helmholtz formula 213
7.5. Influence of activity on the cell voltage 214
7.6. Dissymmetry of cells, chemical cells and concentration cells 215
7.7. Applications to the thermodynamics of electrochemical cells 216
7.7.1. Determining the standard potentials of cells 216
7.7.2. Determination of the dissociation constant of a weak electrolyte on the basis of the potential of a cell 218
7.7.3. Measuring the activity of a component in a strong electrolyte 221
7.7.4. Influence of complex formation on the redox potential 224
7.7.5. Electrochemical methods for studying complexes 226
7.7.6. Determining the ion product of a solvent 234
7.7.7. Determining a solubility product 235
7.7.8. Determining the enthalpies, entropies and Gibbs energies of reactions 236
7.7.9. Determining the standard Gibbs energies of the ions 237
7.7.10. Determining the standard entropies of the ions 238
7.7.11. Measuring the activity of a component of a non-ionic conductive solution (metal solution) 238
7.7.12. Measuring the activity coefficient of transfer of a strong electrolyte 241
7.7.13. Evaluating the individual activity coefficient of transport for an ion 242
Chapter 8. Potential/Acidity Diagrams 245
8.1. Conventions 245
8.1.1. Plotting conventions 245
8.1.2. Boundary equations 246
8.2. Intersections of lines in the diagram 249
8.2.1. Relative disposition of the lines in the vicinity of a triple point 249
8.2.2. Shape of equi-concentration lines in the vicinity of a triple point 250
8.3. Plotting a diagram: example of copper 256
8.3.1. Step 1: list of species and thermodynamic data 256
8.3.2. Step 2: choice of hydrated forms 256
8.3.3. Step 3: study by degrees of oxidation of acid–base reactions; construction of the situation diagram 257
8.3.4. Step 4: elimination of unstable species by dismutation 259
8.3.5. Step 5: plotting the e/pH diagram 261
8.4. Diagram for water superposed on the diagram for an element 262
8.5. Immunity, corrosion and passivation 263
8.6. Potential/pX (e/pX) diagrams 264
8.7. Potential/acidity diagrams in a molten salt 265
Appendix 267
Bibliography 275
Index 279
