An Introduction to Aqueous Electrolyte Solutions

An Introduction to Aqueous Electrolyte Solutions is a comprehensive coverage of the subject including the development of key concepts and theory that focus on the physical rather than the mathematical aspects. Important links are made between the study of electrolyte solutions and other branches of chemistry, biology, and biochemistry, making it a useful cross–reference tool for students studying this important area of electrochemistry.
Carefully developed throughout, each chapter includes intended learning outcomes and worked problems and examples to encourage student understanding of this multidisciplinary subject.a comprehensive introduction to aqueous electrolyte solutions including the development of key concepts and theoriesemphasises the connection between observable macroscopic experimental properties and interpretations made at the molecular levelkey developments in concepts and theory explained in a descriptive manner to encourage student understandingincludes worked problems and examples throughout
An invaluable text for students taking courses in chemistry and chemical engineering, this book will also be useful for biology, biochemistry and biophysics students required to study electrochemistry.

Spis treści:
Preface.
Preliminary Chapter Guidance to Student.
List of symbols.
1 Concepts and Ideas: Setting the Stage.
1.1 Electrolyte solutions – what are they?
1.2 Ions – simple charged particles or not?
1.3 The solvent: structureless or not?
1.4 The medium: its structure and the effect of ions on this structure.
1.5 How can these ideas help in understanding what might happen when an ion is put into a solvent?
1.6 Electrostriction.
1.7 Ideal and non–ideal solutions – what are they?
1.8 The ideal electrolyte solution.
1.9 The non–ideal electrolyte solution.
1.10 Macroscopic manifestation of non–ideality.
1.11 S pecies present in solution.
1.12 Formation of ion pairs from free ions.
1.13 Complexes from free ions.
1.14 Complexes from ions and uncharged ligands.
1.15 Chelates from free ions.
1.16 Micelle formation from free ions.
1.17 Measuring the equilibrium constant: general considerations.
1.18 Base–lines for theoretical predictions about the behaviour expected for a solution consisting of free ions only, Debye–Hückel and Fuoss–Onsager theories and the use of Beer’s Law.
1.19 Ultrasonics.
1.20 Possibility that specific experimental methods could distinguish between the various types of associated species.
1.21 Some examples of how chemists could go about inferring the nature of the species present.
2 The Concept of Chemical Equilibrium: An Introduction.
2.1 Irreversible and reversible reactions.
2.2 Composition of equilibrium mixtures, and the approach to equilibrium.
2.3 Meaning of the term ‘position of equilibrium’ and formulation of the equilibrium constant.
2.4 Equilibrium and the direction of reaction.
2.5 A searching problem.
2.6 The position of equilibrium.
2.7 Other generalisations about equilibrium.
2.8 K and pK.
2.9 Qualitative experimental observations on the effect of temperature on the equilibrium constant, K.
2.10 Qualitative experimental observations on the effect of pressure on the equilibrium constant, K.
2.11 Stoichiometric relations.
2.12 A further relation essential to the description of electrolyte solutions – electrical neutrality.
3 Acids and Bases: A First Approach.
3.1 A qualitative description of acid–base equilibria.
3.2 The self ionisation of water.
3.3 Strong and weak acids and bases.
3.4 A more detailed description of acid–base behaviour.
3.5 Ampholytes.
3.6 Other situations where acid/base behaviour appears.
3.7 Formulation of equilibrium constants i n acid–base equilibria.
3.8 Magnitudes of equilibrium constants.
3.9 The self ionisation of water.
3.10 Relations between Ka and Kb: expressions for an acid and its conjugate base and for a base and its conjugate acid.
3.11 Stoichiometric arguments in equilibria calculations.
3.12 Procedure for calculations on equilibria.
4 Equilibrium Calculations for Acids and Bases.
4.1 Calculations on equilibria: weak acids.
4.2 Some worked examples.
4.3 Calculations on equilibria: weak bases.
4.4 Some illustrative problems.
4.5 Fraction ionised and fraction not ionised for a weak acid; fraction protonated and fraction not protonated for a weak base.
4.6 Dependence of the fraction ionised on pKa and pH.
4.7. The effect of dilution on the fraction ionised for weak acids lying roughly in the range: pKa = 4.0 to 10.0.
4.8 Reassessment of the two approximations: a rigorous expression for a weak acid.
4.9 Conjugate acids of weak bases.
4.10 Weak bases.
4.11 Effect of non–ideality.
5 Equilibrium Calculations for Salts and Buffers.
5.1 Aqueous solutions of salts.
5.2 Salts of strong acids/strong bases.
5.3 Salts of weak acids/strong bases.
5.4 Salts of weak bases/strong acids.
5.5 Salts of weak acids/weak bases.
5.6 Buffer solutions.
6 Neutralisation and pH Titration Curves.

6.1 Neutralisation.
6.2 pH titration curves.
6.3 Interpretation of pH titration curves.
6.4 Polybasic acids.
6.5 pH titrations of dibasic acids: the calculations.
6.6 Tribasic acids.
6.7 Ampholytes.
7 Ion Pairing, Complex Formation and Solubilities.
7.1 Ion pair formation.
7.2 Complex formation.
7.3 Solubilities of sparingly soluble salts.
8 Practical Applications of Thermodynamics for Electrolyte Solutions.
8.1 The first law of thermodynamics.
8.2 The enthalpy, H.
8.3 The reversible process.
8.4 The second law of thermodynamics.
8.5 Relations between q, w and thermodynamic quantities.
8.6 Some other definitions of important thermodynamic functions.
8.7 A very important equation which can now be derived.
8.8 Relation of emfs to thermodynamic quantities.
8.9 The thermodynamic criterion of equilibrium.
8.10 Some further definitions: standard states and standard values.
8.11 The chemical potential of a substance.
8.12 Criterion of equilibrium in terms of chemical potentials.
8.13 Chemical potentials for solids, liquids, gases and solutes.
8.14 Use of the thermodynamic criterion of equilibrium in the derivation of the algebraic form of the equilibrium constant.
8.15 The temperature dependence of ΔHθ.
8.16 The dependence of the equilibrium constant, K, on temperature.
8.16.2 Determination of ΔHθ from values of K over a range of temperatures.
8.17 The microscopic statistical interpretation of entropy.
8.18 Dependence of K on pressure.
8.19 Dependence of ΔGθ on temperature.
8.20 Dependence of ΔSθ on temperature.
8.21 The non–ideal case.
8.22 Chemical potentials and mean activity coefficients.
8.23 A generalisation.
8.24 Corrections for non–ideality for experimental equilibrium constants.
8.24.1 Dependence of equilibrium constants on ionic strength.
8.25 Some specific examples of the dependence of the equilibrium constant on ionic strength.
8.25.3 The weak acid where there is extensive ionisation.
8.26 Graphical corrections for non–ideality.
8.27 Comparison of non–graphical and graphical methods of correcting for non–ideality.
8.28 Dependence of fraction ionised and fractiion protonated on ionic strength.
8.29 Thermodynamic quantities and the effect of non–ideality.
9 Electrochemical Cells and EMFs.
9.1 Chemical aspects of