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• Cover
• Halftitle
• Title
• Copyright
• Preface
• Preface
• Contents
• Notation
• 1 Introduction
• 1.1 The scope and nature of chemical thermodynamics
• 1.2 Equilibrium in mechanical systems
• 1.3 Reversibility and equilibrium
• 1.4 Why we need thermodynamics
• 1.5 The mole
• 1.6 The perfect gas
• 2 Energy
• 2.1 Work
• 2.2 Heat and temperature
• 2.3 Measurement of temperature
• 2.4 Heat and molecular motion
• 2.5 Conservation of energy
• 2.6 State functions: a digression
• 2.7 Enthalpy
• 2.8 Heat capacity
• Problems
• 3 Entropy and equilibrium
• 3.1 Reversibility and equilibrium: a recapitulation
• 3.2 Condition of equilibrium
• 3.3 Entropy
• 3.4 Entropy as a state function
• 3.5 Entropy of expansion of a gas
• 3.6 Entropy changes accompanying heat flow
• 3.7 Entropy and equilibrium
• 3.8 A cosmological aside
• 3.9 Entropy as a function of pressure and temperature
• 3.10   Molecular basis of entropy
• 3.11   Statistical basis of the Second Law
• 3.12   Magnitudes of entropy changes
• 3.13   Heat engines
• Problems
• 4 Equilibrium in chemical systems
• 4.1 Free energy
• 4.2 Gibbs energy
• 4.3 Pressure-dependence of Gibbs energy
• 4.4 Temperature variation of Gibbs energy
• 4.5 Phase equilibria
• 4.6 Clapeyron equation
• 4.7 Clausius–Clapeyron equation
• 4.8 The vapour pressure of liquids
• 4.9 Chemical potential
• 4.10   Chemical potential and Gibbs energy
• 4.11   Equilibrium between gaseous reactants
• 4.12   Temperature-dependence of equilibrium constants
• 4.13   Effect of pressure on equilibrium constants
• 4.14   Basic results of chemical thermodynamics
• 4.15   Le Chatelier’s Principle
• Problems
• 5 Determination of thermodynamic quantities
• 5.1 Hess’s Law
• 5.2 Standard enthalpies of formation
• 5.3 Average bond energies
• 5.4 Temperature-dependence of enthalpy changes
• 5.5 Standard Gibbs energies of formation
• 5.6 Determination of Gibbs energy changes
• 5.7 Determination of entropies of substances
• 5.8 Example of the determination of thermodynamic quantities
• 5.9 Calculation of thermodynamic quantities at temperatures other than 298 K
• 5.10   Ellingham diagrams
• 5.11   Free-energy functions
• Problems
• 6 Ideal solutions
• 6.1 The ideal solution
• 6.2 Properties of truly ideal solutions
• 6.3 Mixtures of liquids
• 6.4 Ideal solutions of solids in liquids
• 6.5 Ideal dilute solutions
• 6.6 Colligative properties
• 6.7 Freezing-point depression
• 6.8 Elevation of boiling point
• 6.9 Osmotic pressure
• 6.10   Properties of the solute in dilute solutions
• 6.11   Solubility of solids
• Problems
• 7 Non-ideal solutions
• 7.1 The concept of activity
• 7.2 Activity of solids in liquids
• 7.3 Activity in aqueous solutions
• 7.4 Chemical equilibria in solution
• 7.5 Electrochemical cells
• 7.6 Standard electrode potentials
• Problems
• 8 Thermodynamics of gases
• 8.1 Expansion of a perfect gas
• 8.2 Irreversible expansion
• 8.3 Equation of state of gases
• 8.4 The Joule–Thomson experiment
• 8.5 Imperfect gases: fugacity
• 8.6 Calculation of fugacities
• Problems
• 9 The molecular basis of thermodynamics
• 9.1 Energy levels
• 9.2 Microstates.160
• 9.3 The Boltzmann factor
• 9.4 The behaviour of heat capacity
• 9.5 Partition functions
• 9.6 Entropy and the partition function
• 9.7 Calculation of the translational partition function
• 9.8 The rotational partition function
• 9.9 The vibrational partition function
• 9.10   Evaluation of the thermodynamic properties of gaseous nitrogen
• 9.11   Chemical equilibrium
• Problems
• Answers to problems
• Further reading
• Appendix 1: Thermochemical data at 298.15K
• Appendix 2: Thermodynamic data for ions in aqueous solution at 298.15K
• Appendix 3: Units and fundamental constants
• Periodic table of elements
• Index