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An Introduction to Aqueous Electrolyte Solutions is a comprehensive coverage of solution equilibria and properties of aqueous ionic solutions. Acid/base equilibria, ion pairing, complex formation, solubilities, reversible emf’s and experimental conductance studies are all illustrated by many worked examples. Theories of non-ideality leading to expressions for activity coefficients, conductance theories and investigations of solvation are described; great care being taken to provide detailed verbal clarification of the key concepts of these theories. The theoretical development focuses on the physical aspects, with the mathematical development being fully explained. An overview of the thermodynamic background is given.

Each chapter includes intended learning outcomes and worked problems and examples to encourage student understanding of this multidisciplinary subject.

An invaluable text for students taking courses in chemistry and chemical engineering. This book will also be useful for biology, biochemistry and biophysics students who may be required to study electrochemistry as part of their course. 

  • A comprehensive introduction to the behaviour and properties of aqueous ionic solutions, including clear explanation and development of key concepts and theories
  • Clear, student friendly style clarifying complex aspects which students find difficult
  • Key developments in concepts and theory explained in a descriptive manner to encourage student understanding
  • Includes worked problems and examples throughout
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Product Details

  • ISBN-13: 9780470842935
  • Publisher: Wiley
  • Publication date: 6/15/2007
  • Edition number: 1
  • Pages: 602
  • Product dimensions: 6.87 (w) x 9.80 (h) x 1.61 (d)

Table of Contents


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 Species 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 in 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 Δ 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 the passage of an electric current through a conducting medium.

9.2 Electrolysis.

9.3 Electrochemical cells.

9.4 Some examples of electrodes used in electrochemical cells.

9.5 Combination of electrodes to make an electrochemical cell.

9.6 Conventions for writing down the electrochemical cell.

9.7 One very important point: cells corresponding to a ‘net chemical reaction’.

9.8 Liquid junctions in electrochemical cells.

9.9 Experimental determination of the direction of flow of the electrons, and measurement of the potential difference.

9.10 Electrode potentials.

9.11 Standard electrode potentials.

9.12 Potential difference, electrical work done and ΔG for the cell reaction.

9.13 ΔG for the cell process: the Nernst equation.

9.14 Methods of expressing concentration.

9.15 Calculation of standard emfs values for cells and ΔGθ values for reactions.

9.16 Determination of pH.

9.17 Determination of equilibrium constants for reactions where K is either very large or very small.

9.18 Use of concentration cells.

9.19 ‘Concealed’ concentration cells and similar cells.

9.20 Determination of equilibrium constants and pK values for reactions which are not directly that for the cell reaction .

9.21 Use of concentration cells with and without liquid junctions in the determination of transport numbers.

10 Concepts and Theory of Non-ideality.

10.1 Evidence for non-ideality in electrolyte solutions.

10.2 The problem theoretically.

10.3 Features of the simple Debye-Hückel model.

10.4 Aspects of electrostatics which are necessary for an understanding of the procedures used in the Debye-Hückel theory and conductance theory.

10.5 The ionic atmosphere in more detail.

10.6 Derivation of the Debye-Hückel theory from the simple Debye-Hückel model.

10.7 The Debye-Hückel limiting law.

10.8 Shortcomings of the Debye-Hückel model.

10.9 Shortcomings in the mathematical derivation of the theory.

10.10 Modifications and further developments of the theory.

10.11 Evidence for ion association from Debye-Hückel plots.

10.12 The Bjerrum theory of ion association.

10.12.6 Fuoss ion pairs and others.

10.13 Extensions to higher concentrations.

10.14 Modern developments in electrolyte theory.

10.15 Computer simulations.

10.16 Further developments to the Debye-Hückel theory.

10.16.7 Use of these ideas in producing a new treatment.

10.17 Statistical mechanics and distribution functions.

10.18 Application of distribution functions to the determination of activity coefficients due to Kirkwood; Yvon; Born and Green; and Bogolyubov.

10.19 A few examples of results from distribution functions.

10.20 ‘Born-Oppenheimer level’ models.

10.21 Lattice calculations for concentrated solutions.

11 Conductance: The Ideal Case.

11.1 Aspects of physics relevant to the experimental study of conductance in solution.

11.2 Experimental measurement of the conductivity of a solution.

11.3 Corrections to the observed conductivity to account for the self ionisation of water.

11.4 Conductivities and molar conductivities: the ideal case.

11.5 The physical significance of the molar conductivity, Λ.

11.6 Dependence of molar conductivity on concentration for a strong electrolyte: the ideal case.

11.7 Dependence of molar conductivity on concentration for a weak electrolyte: the ideal case.

11.8 Determination of Λ0.

11.9 Simultaneous determination of K and Λ0.

11.10 Problems when an acid or base is so weak that it is never 100% ionised, even in very, very dilute solution.

11.11 Contributions to the conductivity of an electrolyte solution from the cation and the anion of the electrolyte.

11.12 Contributions to the molar conductivity from the individual ions.

11.13 Kohlrausch’s law of independent ionic mobilities.

11.14 Analysis of the use of conductance measurements for determination of pKas for very weak acids and pKbs for very weak bases: the basic quantities involved.

11.15 Use of conductance measurements in determining solubility products for sparingly soluble salts.

11.16 Transport numbers.

11.17 Ionic mobilities.

11.18 Abnormal mobility and ionic molar conductivity of H3O+(aq).

11.19 Measurement of transport numbers.

12 Theories of Conductance: The Non-ideal Case for Symmetrical Electrolytes.

12.1 The relaxation effect.

12.2 The electrophoretic effect.

12.3 Conductance equations for strong electrolytes taking non-ideality into consideration: early conductance theory.

12.4 A simple treatment of the derivation of the Debye-Hückel-Onsager equation 1927 for symmetrical electrolytes.

12.5 The Fuoss-Onsager equation 1932.

12.6 Use of the Debye-Hückel-Onsager equation for symmetrical strong electrolytes which are fully dissociated.

12.7 Electrolytes showing ion pairing and weak electrolytes which are not fully dissociated.

12.8 Empirical extensions to the Debye-Hückel-Onsager 1927 equation.

12.9 Modern conductance theories for symmetrical electrolytes – post 1950.

12.10 Fuoss-Onsager 1957: Conductance equation for symmetrical electrolytes.

12.11 A simple illustration of the effects of ion association on experimental conductance curves.

12.12 The Fuoss-Onsager equation for associated electrolytes.

12.13 Range of applicability of Fuoss-Onsager 1957 conductance equation for symmetrical electrolytes.

12.14 Limitations of the treatment given by the 1957 Fuoss-Onsager conductance equation for symmetrical electrolytes.

12.15 Manipulation of the 1957 Fuoss-Onsager equation, and later modifications by Fuoss and other workers.

12.16 Conductance studies over a range of relative permittivities.

12.17 Fuoss et al. 1978 and later.

Appendix 1.

Appendex 2.

13 Solvation.

13.1 Classification of solutes: a resume.

13.2 Classification of solvents.

13.3 Solvent structure.

13.4 The experimental study of the structure of water.

13.5 Diffraction studies.

13.6 The theoretical approach to the radial distribution function for a liquid.

13.7 Aqueous solutions of electrolytes.

13.8 Terms used in describing hydration.

13.9 Traditional methods for measuring solvation numbers.

13.10 Modern techniques for studying hydration: NMR.

13.11. Modern techniques of studying hydration: neutron and X-ray diffraction.

13.12 Modern techniques of studying solvation: AXD diffraction and EXAFS.

13.13 Modern techniques of studying solvation: computer simulations.

13.14 Cautionary remarks on the significance of the numerical values of solvation numbers.

13.15 Sizes of ions.

13.16 A first model of solvation – the three region model for aqueous electrolyte solutions.

13.17 Volume changes on solvation.

13.18 Viscosity data.

13.19 Concluding comment.

13.20 Determination of ΔGθ hydration.

13 21 Determination of ΔHθ hydration.

13.22 Compilation of entropies of hydration from ΔGθ hydration and ΔHθ hydration.

13.23 Thermodynamic transfer functions.

13.24 Solvation of non-polar and apolar molecules – hydrophobic effects.

13.25 Experimental techniques for studying hydrophobic hydration.

13.26 Hydrophobic hydration for large charged ions.

13.27 Hydrophobic interaction.

13.28 Computer simulations of the hydrophobic effect.

Subject Matter of Worked Problems.


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