Michel Soustelle 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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• Preface xiNotations and Symbols  xvPart 1. Ionic Equilibria 1Chapter 1. Dissociation of Electrolytes in Solution 31.1. Strong electrolytes – weak electrolytes 31.1.1. Dissolution 31.1.2. Solvolysis 41.1.3. Melting 41.2. Mean concentration and mean activity coefficient of ions 51.3. Dissociation coefficient of a weak electrolyte 61.4. Conduction of electrical current by electrolytes 91.4.1. Transport numbers and electrical conductivity of an electrolyte 91.4.2. Equivalent conductivity and limiting equivalent conductivity of an electrolyte 101.4.3. Ionic mobility 111.4.4. Relation between equivalent conductivity and mobility – Kohlrausch’s law 141.4.5. Apparent dissociation coefficient and equivalent conductivity 161.4.6. Variations of equivalent conductivities with the concentrations 161.5. Determination of the dissociation coefficient 201.5.1. Determination of the dissociation coefficient by the cryometric method 211.5.2. Determination of the dissociation coefficient on the basis of the conductivity values 221.6. Determination of the number of ions produced by dissociation 231.6.1. Use of limiting molar conductivity 231.6.2. Use of cryometry 241.7. Thermodynamic values relative to the ions 271.7.1. The standard molar Gibbs energy of formation of an ion 271.7.2. Standard enthalpy of formation of ions 291.7.3. Absolute standard molar entropy of an ion 291.7.4. Determination of the mean activity of a weak electrolyte on the basis of the dissociation equilibrium 30Chapter 2. Solvents and Solvation 312.1. Solvents 312.2. Solvation and structure of the solvated ion 332.3. Thermodynamics of solvation 352.3.1. Thermodynamic values of solvation 362.3.2. Gibbs energy of salvation – Born’s model 372.4. Transfer of a solute from one solvent to another 442.5. Mean transfer activity coefficient of solvation of an electrolyte 482.6. Experimentally determining the transfer activity coefficient of solvation 492.6.1. Determining the activity coefficient of a molecular solute 502.6.2. Determination of the mean transfer activity coefficient of a strong electrolyte 512.6.3. Evaluation of the individual transfer activity coefficient of an ion 512.7. Relation between the constants of the same equilibrium achieved in two different solvents 552.7.1. General relation of solvent change on an equilibrium constant 552.7.2. Influence of the dielectric constant of the solvent on the equilibrium constant of an ionic reaction 56Chapter 3. Acid/Base Equilibria 613.1. Definition of acids and bases and acid–base reactions 623.2. Ion product of an amphiprotic solvent 633.3. Relative strengths of acids and bases 643.3.1. Definition of the acidity constant of an acid 643.3.2. Protic activity in a solvent 673.4. Direction of acid–base reactions, and domain of predominance 693.5. Leveling effect of a solvent 713.6. Modeling of the strength of an acid 753.6.1. Model of the strength of an acid 753.6.2. Comparison of an acid’s behavior in two solvents 783.6.3. Construction of activity zones for solvents 813.7. Acidity functions and acidity scales 843.8. Applications of the acidity function 883.8.1. Measuring the pKa of an indicator 893.8.2. Measuring the ion products of solvents 893.9. Acidity in non-protic molecular solvents 913.10. Protolysis in ionic solvents (molten salts) 923.11. Other ionic exchanges in solution 933.11.1. Ionoscopy 933.11.2. Acidity in molten salts: definition given by Lux and Flood 943.12. Franklin and Gutmann’s solvo-acidity and solvo-basicity 963.12.1. Definition of solvo-acidity 963.12.2. Solvo-acidity in molecular solvents 963.12.3. Solvo-acidity in molten salts 983.13. Acidity as understood by Lewis 100Chapter 4. Complexations and Redox Equilibria 1014.1. Complexation reactions 1014.1.1. Stability of complexes 1014.1.2. Competition between two ligands on the same acceptor 1064.1.3. Method for studying perfect complexes 1084.1.4. Methods for studying imperfect complexes 1104.1.5. Study of successive complexes 1154.2. Redox reactions 1174.2.1. Electronegativity – electronegativity scale 1174.2.2. Degrees of oxidation 1244.2.3. Definition of redox reactions 1284.2.4. The two families of redox reactions 1284.2.5. Dismutation and antidismutation 1304.2.6. Redox reactions, and calculation of the stoichiometric numbers 1314.2.7. Concept of a redox couple 132Chapter 5. Precipitation Reactions and Equilibria 1355.1. Solubility of electrolytes in water – solubility product 1355.2. Influence of complex formation on the solubility of a salt 1365.3. Application of the solubility product in determining the stability constant of complex ions . 1375.4. Solution with multiple electrolytes at equilibrium with pure solid phases 1385.4.1. Influence of a salt with non-common ions on the solubility of a salt 1395.4.2. Influence of a salt with a common ion on the solubility of a salt 1415.4.3. Crystallization phase diagram for a mixture of two salts in solution 1415.4.4. Formation of double salts or chemical combinations in the solid state 1425.4.5. Reciprocal quaternary systems – square diagrams 1445.5. Electrolytic aqueous solution and solid solution 1475.5.1. Thermodynamic equilibrium between a liquid ionic solution and a solid solution 1475.5.2. Solubility product of a solid solution 1505.6. Solubility and pH 1555.6.1. Solubility and pH 1555.6.2. Solubility of oxides in molten alkali hydroxides 1565.6.3. Solubility in oxo-acids and oxo-bases (see section 3.12.2) 1575.7. Calculation of equilibria in ionic solutions 158Part 2. Electrochemical Thermodynamics 163Chapter 6. Thermodynamics of the Electrode 1656.1. Electrochemical systems 1656.1.1. The electrochemical system 1666.1.2. Electrochemical functions of state 1676.1.3. Electrochemical potential 1676.1.4. Gibbs–Duhem relation for electrochemical systems 1696.1.5. Chemical system associated with an electrochemical system 1706.1.6. General conditions of an equilibrium of an electrochemical system 1716.2. The electrode 1736.2.1. Definition and reaction of the electrode 1736.2.2. Equilibrium of an insulated metal electrode – electrode absolute voltage 1746.2.3. Voltage relative to a metal electrode – Nernst’s relation 1756.2.4. Chemical and electrochemical Gibbs energy of the electrode reaction 1786.2.5. Influence of pH on the electrode voltage 1796.2.6. Influence of the solvent and of the dissolved species on the electrode voltage 1816.2.7. Influence of temperature on the normal potentials 1836.3. The different types of electrodes 1846.3.1. Redox electrodes 1846.3.2. Metal electrodes 1896.3.3. Gas electrodes 1926.4. Equilibrium of two ionic conductors in contact  1936.4.1. Junction potential with a semi-permeable membrane 1936.4.2. Junction potential of two electrolytes with a permeable membrane 1946.5. Applications of Nernst’s relation to the study of various reactions 1966.5.1. Prediction of redox reactions 1966.5.2. Relations between the redox voltages of different systems of the same element 1976.5.3. Predicting the dismutation and anti-dismutation reactions 2016.5.4. Redox catalysis 2026.6. Redox potential in a non-aqueous solvent 2036.6.1. Scale of redox potential in a non-aqueous medium 2036.6.2. Oxidation and reduction of the solvent 2066.6.3. Influence of solvent on redox systems in a non-aqueous solvent 207Chapter 7. Thermodynamics of Electrochemical Cells 2097.1. Electrochemical chains – batteries and electrolyzer cells 2097.2. Electrical voltage of an electrochemical cell 2107.3. Cell reaction 2127.4. Influence of temperature on the cell voltage; Gibbs–Helmholtz formula 2137.5. Influence of activity on the cell voltage 2147.6. Dissymmetry of cells, chemical cells and concentration cells 2157.7. Applications to the thermodynamics of electrochemical cells 2167.7.1. Determining the standard potentials of cells 2167.7.2. Determination of the dissociation constant of a weak electrolyte on the basis of the potential of a cell 2187.7.3. Measuring the activity of a component in a strong electrolyte 2217.7.4. Influence of complex formation on the redox potential 2247.7.5. Electrochemical methods for studying complexes 2267.7.6. Determining the ion product of a solvent 2347.7.7. Determining a solubility product 2357.7.8. Determining the enthalpies, entropies and Gibbs energies of reactions 2367.7.9. Determining the standard Gibbs energies of the ions 2377.7.10. Determining the standard entropies of the ions 2387.7.11. Measuring the activity of a component of a non-ionic conductive solution (metal solution) 2387.7.12. Measuring the activity coefficient of transfer of a strong electrolyte 2417.7.13. Evaluating the individual activity coefficient of transport for an ion 242Chapter 8. Potential/Acidity Diagrams 2458.1. Conventions 2458.1.1. Plotting conventions 2458.1.2. Boundary equations 2468.2. Intersections of lines in the diagram 2498.2.1. Relative disposition of the lines in the vicinity of a triple point 2498.2.2. Shape of equi-concentration lines in the vicinity of a triple point 2508.3. Plotting a diagram: example of copper 2568.3.1. Step 1: list of species and thermodynamic data 2568.3.2. Step 2: choice of hydrated forms 2568.3.3. Step 3: study by degrees of oxidation of acid–base reactions; construction of the situation diagram 2578.3.4. Step 4: elimination of unstable species by dismutation 2598.3.5. Step 5: plotting the e/pH diagram 2618.4. Diagram for water superposed on the diagram for an element 2628.5. Immunity, corrosion and passivation 2638.6. Potential/pX (e/pX) diagrams 2648.7. Potential/acidity diagrams in a molten salt 265Appendix 267Bibliography 275Index 279

Michel Soustelle 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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• Preface xiNotations and Symbols  xvPart 1. Ionic Equilibria 1Chapter 1. Dissociation of Electrolytes in Solution 31.1. Strong electrolytes – weak electrolytes 31.1.1. Dissolution 31.1.2. Solvolysis 41.1.3. Melting 41.2. Mean concentration and mean activity coefficient of ions 51.3. Dissociation coefficient of a weak electrolyte 61.4. Conduction of electrical current by electrolytes 91.4.1. Transport numbers and electrical conductivity of an electrolyte 91.4.2. Equivalent conductivity and limiting equivalent conductivity of an electrolyte 101.4.3. Ionic mobility 111.4.4. Relation between equivalent conductivity and mobility – Kohlrausch’s law 141.4.5. Apparent dissociation coefficient and equivalent conductivity 161.4.6. Variations of equivalent conductivities with the concentrations 161.5. Determination of the dissociation coefficient 201.5.1. Determination of the dissociation coefficient by the cryometric method 211.5.2. Determination of the dissociation coefficient on the basis of the conductivity values 221.6. Determination of the number of ions produced by dissociation 231.6.1. Use of limiting molar conductivity 231.6.2. Use of cryometry 241.7. Thermodynamic values relative to the ions 271.7.1. The standard molar Gibbs energy of formation of an ion 271.7.2. Standard enthalpy of formation of ions 291.7.3. Absolute standard molar entropy of an ion 291.7.4. Determination of the mean activity of a weak electrolyte on the basis of the dissociation equilibrium 30Chapter 2. Solvents and Solvation 312.1. Solvents 312.2. Solvation and structure of the solvated ion 332.3. Thermodynamics of solvation 352.3.1. Thermodynamic values of solvation 362.3.2. Gibbs energy of salvation – Born’s model 372.4. Transfer of a solute from one solvent to another 442.5. Mean transfer activity coefficient of solvation of an electrolyte 482.6. Experimentally determining the transfer activity coefficient of solvation 492.6.1. Determining the activity coefficient of a molecular solute 502.6.2. Determination of the mean transfer activity coefficient of a strong electrolyte 512.6.3. Evaluation of the individual transfer activity coefficient of an ion 512.7. Relation between the constants of the same equilibrium achieved in two different solvents 552.7.1. General relation of solvent change on an equilibrium constant 552.7.2. Influence of the dielectric constant of the solvent on the equilibrium constant of an ionic reaction 56Chapter 3. Acid/Base Equilibria 613.1. Definition of acids and bases and acid–base reactions 623.2. Ion product of an amphiprotic solvent 633.3. Relative strengths of acids and bases 643.3.1. Definition of the acidity constant of an acid 643.3.2. Protic activity in a solvent 673.4. Direction of acid–base reactions, and domain of predominance 693.5. Leveling effect of a solvent 713.6. Modeling of the strength of an acid 753.6.1. Model of the strength of an acid 753.6.2. Comparison of an acid’s behavior in two solvents 783.6.3. Construction of activity zones for solvents 813.7. Acidity functions and acidity scales 843.8. Applications of the acidity function 883.8.1. Measuring the pKa of an indicator 893.8.2. Measuring the ion products of solvents 893.9. Acidity in non-protic molecular solvents 913.10. Protolysis in ionic solvents (molten salts) 923.11. Other ionic exchanges in solution 933.11.1. Ionoscopy 933.11.2. Acidity in molten salts: definition given by Lux and Flood 943.12. Franklin and Gutmann’s solvo-acidity and solvo-basicity 963.12.1. Definition of solvo-acidity 963.12.2. Solvo-acidity in molecular solvents 963.12.3. Solvo-acidity in molten salts 983.13. Acidity as understood by Lewis 100Chapter 4. Complexations and Redox Equilibria 1014.1. Complexation reactions 1014.1.1. Stability of complexes 1014.1.2. Competition between two ligands on the same acceptor 1064.1.3. Method for studying perfect complexes 1084.1.4. Methods for studying imperfect complexes 1104.1.5. Study of successive complexes 1154.2. Redox reactions 1174.2.1. Electronegativity – electronegativity scale 1174.2.2. Degrees of oxidation 1244.2.3. Definition of redox reactions 1284.2.4. The two families of redox reactions 1284.2.5. Dismutation and antidismutation 1304.2.6. Redox reactions, and calculation of the stoichiometric numbers 1314.2.7. Concept of a redox couple 132Chapter 5. Precipitation Reactions and Equilibria 1355.1. Solubility of electrolytes in water – solubility product 1355.2. Influence of complex formation on the solubility of a salt 1365.3. Application of the solubility product in determining the stability constant of complex ions . 1375.4. Solution with multiple electrolytes at equilibrium with pure solid phases 1385.4.1. Influence of a salt with non-common ions on the solubility of a salt 1395.4.2. Influence of a salt with a common ion on the solubility of a salt 1415.4.3. Crystallization phase diagram for a mixture of two salts in solution 1415.4.4. Formation of double salts or chemical combinations in the solid state 1425.4.5. Reciprocal quaternary systems – square diagrams 1445.5. Electrolytic aqueous solution and solid solution 1475.5.1. Thermodynamic equilibrium between a liquid ionic solution and a solid solution 1475.5.2. Solubility product of a solid solution 1505.6. Solubility and pH 1555.6.1. Solubility and pH 1555.6.2. Solubility of oxides in molten alkali hydroxides 1565.6.3. Solubility in oxo-acids and oxo-bases (see section 3.12.2) 1575.7. Calculation of equilibria in ionic solutions 158Part 2. Electrochemical Thermodynamics 163Chapter 6. Thermodynamics of the Electrode 1656.1. Electrochemical systems 1656.1.1. The electrochemical system 1666.1.2. Electrochemical functions of state 1676.1.3. Electrochemical potential 1676.1.4. Gibbs–Duhem relation for electrochemical systems 1696.1.5. Chemical system associated with an electrochemical system 1706.1.6. General conditions of an equilibrium of an electrochemical system 1716.2. The electrode 1736.2.1. Definition and reaction of the electrode 1736.2.2. Equilibrium of an insulated metal electrode – electrode absolute voltage 1746.2.3. Voltage relative to a metal electrode – Nernst’s relation 1756.2.4. Chemical and electrochemical Gibbs energy of the electrode reaction 1786.2.5. Influence of pH on the electrode voltage 1796.2.6. Influence of the solvent and of the dissolved species on the electrode voltage 1816.2.7. Influence of temperature on the normal potentials 1836.3. The different types of electrodes 1846.3.1. Redox electrodes 1846.3.2. Metal electrodes 1896.3.3. Gas electrodes 1926.4. Equilibrium of two ionic conductors in contact  1936.4.1. Junction potential with a semi-permeable membrane 1936.4.2. Junction potential of two electrolytes with a permeable membrane 1946.5. Applications of Nernst’s relation to the study of various reactions 1966.5.1. Prediction of redox reactions 1966.5.2. Relations between the redox voltages of different systems of the same element 1976.5.3. Predicting the dismutation and anti-dismutation reactions 2016.5.4. Redox catalysis 2026.6. Redox potential in a non-aqueous solvent 2036.6.1. Scale of redox potential in a non-aqueous medium 2036.6.2. Oxidation and reduction of the solvent 2066.6.3. Influence of solvent on redox systems in a non-aqueous solvent 207Chapter 7. Thermodynamics of Electrochemical Cells 2097.1. Electrochemical chains – batteries and electrolyzer cells 2097.2. Electrical voltage of an electrochemical cell 2107.3. Cell reaction 2127.4. Influence of temperature on the cell voltage; Gibbs–Helmholtz formula 2137.5. Influence of activity on the cell voltage 2147.6. Dissymmetry of cells, chemical cells and concentration cells 2157.7. Applications to the thermodynamics of electrochemical cells 2167.7.1. Determining the standard potentials of cells 2167.7.2. Determination of the dissociation constant of a weak electrolyte on the basis of the potential of a cell 2187.7.3. Measuring the activity of a component in a strong electrolyte 2217.7.4. Influence of complex formation on the redox potential 2247.7.5. Electrochemical methods for studying complexes 2267.7.6. Determining the ion product of a solvent 2347.7.7. Determining a solubility product 2357.7.8. Determining the enthalpies, entropies and Gibbs energies of reactions 2367.7.9. Determining the standard Gibbs energies of the ions 2377.7.10. Determining the standard entropies of the ions 2387.7.11. Measuring the activity of a component of a non-ionic conductive solution (metal solution) 2387.7.12. Measuring the activity coefficient of transfer of a strong electrolyte 2417.7.13. Evaluating the individual activity coefficient of transport for an ion 242Chapter 8. Potential/Acidity Diagrams 2458.1. Conventions 2458.1.1. Plotting conventions 2458.1.2. Boundary equations 2468.2. Intersections of lines in the diagram 2498.2.1. Relative disposition of the lines in the vicinity of a triple point 2498.2.2. Shape of equi-concentration lines in the vicinity of a triple point 2508.3. Plotting a diagram: example of copper 2568.3.1. Step 1: list of species and thermodynamic data 2568.3.2. Step 2: choice of hydrated forms 2568.3.3. Step 3: study by degrees of oxidation of acid–base reactions; construction of the situation diagram 2578.3.4. Step 4: elimination of unstable species by dismutation 2598.3.5. Step 5: plotting the e/pH diagram 2618.4. Diagram for water superposed on the diagram for an element 2628.5. Immunity, corrosion and passivation 2638.6. Potential/pX (e/pX) diagrams 2648.7. Potential/acidity diagrams in a molten salt 265Appendix 267Bibliography 275Index 279

Michel Soustelle 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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• Preface xiNotations and Symbols  xvPart 1. Ionic Equilibria 1Chapter 1. Dissociation of Electrolytes in Solution 31.1. Strong electrolytes – weak electrolytes 31.1.1. Dissolution 31.1.2. Solvolysis 41.1.3. Melting 41.2. Mean concentration and mean activity coefficient of ions 51.3. Dissociation coefficient of a weak electrolyte 61.4. Conduction of electrical current by electrolytes 91.4.1. Transport numbers and electrical conductivity of an electrolyte 91.4.2. Equivalent conductivity and limiting equivalent conductivity of an electrolyte 101.4.3. Ionic mobility 111.4.4. Relation between equivalent conductivity and mobility – Kohlrausch’s law 141.4.5. Apparent dissociation coefficient and equivalent conductivity 161.4.6. Variations of equivalent conductivities with the concentrations 161.5. Determination of the dissociation coefficient 201.5.1. Determination of the dissociation coefficient by the cryometric method 211.5.2. Determination of the dissociation coefficient on the basis of the conductivity values 221.6. Determination of the number of ions produced by dissociation 231.6.1. Use of limiting molar conductivity 231.6.2. Use of cryometry 241.7. Thermodynamic values relative to the ions 271.7.1. The standard molar Gibbs energy of formation of an ion 271.7.2. Standard enthalpy of formation of ions 291.7.3. Absolute standard molar entropy of an ion 291.7.4. Determination of the mean activity of a weak electrolyte on the basis of the dissociation equilibrium 30Chapter 2. Solvents and Solvation 312.1. Solvents 312.2. Solvation and structure of the solvated ion 332.3. Thermodynamics of solvation 352.3.1. Thermodynamic values of solvation 362.3.2. Gibbs energy of salvation – Born’s model 372.4. Transfer of a solute from one solvent to another 442.5. Mean transfer activity coefficient of solvation of an electrolyte 482.6. Experimentally determining the transfer activity coefficient of solvation 492.6.1. Determining the activity coefficient of a molecular solute 502.6.2. Determination of the mean transfer activity coefficient of a strong electrolyte 512.6.3. Evaluation of the individual transfer activity coefficient of an ion 512.7. Relation between the constants of the same equilibrium achieved in two different solvents 552.7.1. General relation of solvent change on an equilibrium constant 552.7.2. Influence of the dielectric constant of the solvent on the equilibrium constant of an ionic reaction 56Chapter 3. Acid/Base Equilibria 613.1. Definition of acids and bases and acid–base reactions 623.2. Ion product of an amphiprotic solvent 633.3. Relative strengths of acids and bases 643.3.1. Definition of the acidity constant of an acid 643.3.2. Protic activity in a solvent 673.4. Direction of acid–base reactions, and domain of predominance 693.5. Leveling effect of a solvent 713.6. Modeling of the strength of an acid 753.6.1. Model of the strength of an acid 753.6.2. Comparison of an acid’s behavior in two solvents 783.6.3. Construction of activity zones for solvents 813.7. Acidity functions and acidity scales 843.8. Applications of the acidity function 883.8.1. Measuring the pKa of an indicator 893.8.2. Measuring the ion products of solvents 893.9. Acidity in non-protic molecular solvents 913.10. Protolysis in ionic solvents (molten salts) 923.11. Other ionic exchanges in solution 933.11.1. Ionoscopy 933.11.2. Acidity in molten salts: definition given by Lux and Flood 943.12. Franklin and Gutmann’s solvo-acidity and solvo-basicity 963.12.1. Definition of solvo-acidity 963.12.2. Solvo-acidity in molecular solvents 963.12.3. Solvo-acidity in molten salts 983.13. Acidity as understood by Lewis 100Chapter 4. Complexations and Redox Equilibria 1014.1. Complexation reactions 1014.1.1. Stability of complexes 1014.1.2. Competition between two ligands on the same acceptor 1064.1.3. Method for studying perfect complexes 1084.1.4. Methods for studying imperfect complexes 1104.1.5. Study of successive complexes 1154.2. Redox reactions 1174.2.1. Electronegativity – electronegativity scale 1174.2.2. Degrees of oxidation 1244.2.3. Definition of redox reactions 1284.2.4. The two families of redox reactions 1284.2.5. Dismutation and antidismutation 1304.2.6. Redox reactions, and calculation of the stoichiometric numbers 1314.2.7. Concept of a redox couple 132Chapter 5. Precipitation Reactions and Equilibria 1355.1. Solubility of electrolytes in water – solubility product 1355.2. Influence of complex formation on the solubility of a salt 1365.3. Application of the solubility product in determining the stability constant of complex ions . 1375.4. Solution with multiple electrolytes at equilibrium with pure solid phases 1385.4.1. Influence of a salt with non-common ions on the solubility of a salt 1395.4.2. Influence of a salt with a common ion on the solubility of a salt 1415.4.3. Crystallization phase diagram for a mixture of two salts in solution 1415.4.4. Formation of double salts or chemical combinations in the solid state 1425.4.5. Reciprocal quaternary systems – square diagrams 1445.5. Electrolytic aqueous solution and solid solution 1475.5.1. Thermodynamic equilibrium between a liquid ionic solution and a solid solution 1475.5.2. Solubility product of a solid solution 1505.6. Solubility and pH 1555.6.1. Solubility and pH 1555.6.2. Solubility of oxides in molten alkali hydroxides 1565.6.3. Solubility in oxo-acids and oxo-bases (see section 3.12.2) 1575.7. Calculation of equilibria in ionic solutions 158Part 2. Electrochemical Thermodynamics 163Chapter 6. Thermodynamics of the Electrode 1656.1. Electrochemical systems 1656.1.1. The electrochemical system 1666.1.2. Electrochemical functions of state 1676.1.3. Electrochemical potential 1676.1.4. Gibbs–Duhem relation for electrochemical systems 1696.1.5. Chemical system associated with an electrochemical system 1706.1.6. General conditions of an equilibrium of an electrochemical system 1716.2. The electrode 1736.2.1. Definition and reaction of the electrode 1736.2.2. Equilibrium of an insulated metal electrode – electrode absolute voltage 1746.2.3. Voltage relative to a metal electrode – Nernst’s relation 1756.2.4. Chemical and electrochemical Gibbs energy of the electrode reaction 1786.2.5. Influence of pH on the electrode voltage 1796.2.6. Influence of the solvent and of the dissolved species on the electrode voltage 1816.2.7. Influence of temperature on the normal potentials 1836.3. The different types of electrodes 1846.3.1. Redox electrodes 1846.3.2. Metal electrodes 1896.3.3. Gas electrodes 1926.4. Equilibrium of two ionic conductors in contact  1936.4.1. Junction potential with a semi-permeable membrane 1936.4.2. Junction potential of two electrolytes with a permeable membrane 1946.5. Applications of Nernst’s relation to the study of various reactions 1966.5.1. Prediction of redox reactions 1966.5.2. Relations between the redox voltages of different systems of the same element 1976.5.3. Predicting the dismutation and anti-dismutation reactions 2016.5.4. Redox catalysis 2026.6. Redox potential in a non-aqueous solvent 2036.6.1. Scale of redox potential in a non-aqueous medium 2036.6.2. Oxidation and reduction of the solvent 2066.6.3. Influence of solvent on redox systems in a non-aqueous solvent 207Chapter 7. Thermodynamics of Electrochemical Cells 2097.1. Electrochemical chains – batteries and electrolyzer cells 2097.2. Electrical voltage of an electrochemical cell 2107.3. Cell reaction 2127.4. Influence of temperature on the cell voltage; Gibbs–Helmholtz formula 2137.5. Influence of activity on the cell voltage 2147.6. Dissymmetry of cells, chemical cells and concentration cells 2157.7. Applications to the thermodynamics of electrochemical cells 2167.7.1. Determining the standard potentials of cells 2167.7.2. Determination of the dissociation constant of a weak electrolyte on the basis of the potential of a cell 2187.7.3. Measuring the activity of a component in a strong electrolyte 2217.7.4. Influence of complex formation on the redox potential 2247.7.5. Electrochemical methods for studying complexes 2267.7.6. Determining the ion product of a solvent 2347.7.7. Determining a solubility product 2357.7.8. Determining the enthalpies, entropies and Gibbs energies of reactions 2367.7.9. Determining the standard Gibbs energies of the ions 2377.7.10. Determining the standard entropies of the ions 2387.7.11. Measuring the activity of a component of a non-ionic conductive solution (metal solution) 2387.7.12. Measuring the activity coefficient of transfer of a strong electrolyte 2417.7.13. Evaluating the individual activity coefficient of transport for an ion 242Chapter 8. Potential/Acidity Diagrams 2458.1. Conventions 2458.1.1. Plotting conventions 2458.1.2. Boundary equations 2468.2. Intersections of lines in the diagram 2498.2.1. Relative disposition of the lines in the vicinity of a triple point 2498.2.2. Shape of equi-concentration lines in the vicinity of a triple point 2508.3. Plotting a diagram: example of copper 2568.3.1. Step 1: list of species and thermodynamic data 2568.3.2. Step 2: choice of hydrated forms 2568.3.3. Step 3: study by degrees of oxidation of acid–base reactions; construction of the situation diagram 2578.3.4. Step 4: elimination of unstable species by dismutation 2598.3.5. Step 5: plotting the e/pH diagram 2618.4. Diagram for water superposed on the diagram for an element 2628.5. Immunity, corrosion and passivation 2638.6. Potential/pX (e/pX) diagrams 2648.7. Potential/acidity diagrams in a molten salt 265Appendix 267Bibliography 275Index 279