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

<b>Michel Soustelle</b> is a chemical engineer and Emeritus Professor at Ecole des Mines de Saint-Etienne in France. He taught chemical kinetics from postgraduate to Master degree level while also carrying out research in this topic.

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