Ionic and Electrochemical Equilibria

<p>Preface xi</p>
<p>Notations and Symbols&nbsp; xv</p>
<p>Part 1. Ionic Equilibria 1</p>
<p>Chapter 1. Dissociation of Electrolytes in Solution 3</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>Chapter 2. Solvents and Solvation 31</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>Chapter 3. Acid/Base Equilibria 61</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>Chapter 4. Complexations and Redox Equilibria 101</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>Chapter 5. Precipitation Reactions and Equilibria 135</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>Part 2. Electrochemical Thermodynamics 163</p>
<p>Chapter 6. Thermodynamics of the Electrode 165</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&nbsp; 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>Chapter 7. Thermodynamics of Electrochemical Cells 209</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>Chapter 8. Potential/Acidity Diagrams 245</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>