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Thermodynamics and Chemistry

(Kiana) #1

Biographical Sketches



  • Biographical Sketches

  • Preface to the Second Edition

  • From the Preface to the First Edition

  • 1 Introduction

  • 2 Systems and Their Properties

  • 3 The First Law

  • 4 The Second Law

  • 5 Thermodynamic Potentials

  • 6 The Third Law and Cryogenics

  • 7 Pure Substances in Single Phases

  • 8 Phase Transitions and Equilibria of Pure Substances

  • 9 Mixtures

  • 10 Electrolyte Solutions

  • 11 Reactions and Other Chemical Processes

  • 12 Equilibrium Conditions in Multicomponent Systems

  • 13 The Phase Rule and Phase Diagrams

  • 14 Galvanic Cells


  • Appendix A Definitions of the SI Base Units



  • SHORTCONTENTS

  • Appendix B Physical Constants

  • Appendix C Symbols for Physical Quantities

  • Appendix D Miscellaneous Abbreviations and Symbols

  • Appendix E Calculus Review

  • Appendix F Mathematical Properties of State Functions

  • Appendix G Forces, Energy, and Work

  • Appendix H Standard Molar Thermodynamic Properties

  • Appendix I Answers to Selected Problems

  • Bibliography

  • Index

  • Biographical Sketches CONTENTS

  • Preface to the Second Edition

  • From the Preface to the First Edition

  • 1 Introduction

    • 1.1 Units

      • 1.1.1 Amount of substance and amount



    • 1.2 Quantity Calculus

    • 1.3 Dimensional Analysis

    • Problem



  • 2 Systems and Their Properties

    • 2.1 The System, Surroundings, and Boundary

      • 2.1.1 Extensive and intensive properties



    • 2.2 Phases and Physical States of Matter

      • 2.2.1 Physical states of matter

      • 2.2.2 Phase coexistence and phase transitions

      • 2.2.3 Fluids

      • 2.2.4 The equation of state of a fluid

      • 2.2.5 Virial equations of state for pure gases

      • 2.2.6 Solids



    • 2.3 Some Basic Properties and Their Measurement

      • 2.3.1 Mass

      • 2.3.2 Volume

      • 2.3.3 Density

      • 2.3.4 Pressure

      • 2.3.5 Temperature



    • 2.4 The State of the System

      • 2.4.1 State functions and independent variables

      • 2.4.2 An example: state functions of a mixture


      • 2.4.3 More about independent variables







  • CONTENTS

    • 2.4.4 Equilibrium states

    • 2.4.5 Steady states

    • 2.5 Processes and Paths

    • 2.6 The Energy of the System

      • 2.6.1 Energy and reference frames

      • 2.6.2 Internal energy



    • Problems



  • 3 The First Law

    • 3.1 Heat, Work, and the First Law

      • 3.1.1 The concept of thermodynamic work

      • 3.1.2 Work coefficients and work coordinates

      • 3.1.3 Heat and work as path functions

      • 3.1.4 Heat and heating

      • 3.1.5 Heat capacity

      • 3.1.6 Thermal energy



    • 3.2 Spontaneous, Reversible, and Irreversible Processes

      • 3.2.1 Reversible processes

      • 3.2.2 Irreversible processes

      • 3.2.3 Purely mechanical processes



    • 3.3 Heat Transfer

      • 3.3.1 Heating and cooling

      • 3.3.2 Spontaneous phase transitions



    • 3.4 Deformation Work

      • 3.4.1 Gas in a cylinder-and-piston device

      • 3.4.2 Expansion work of a gas

      • 3.4.3 Expansion work of an isotropic phase

      • 3.4.4 Generalities



    • 3.5 Applications of Expansion Work

      • 3.5.1 The internal energy of an ideal gas

      • 3.5.2 Reversible isothermal expansion of an ideal gas

      • 3.5.3 Reversible adiabatic expansion of an ideal gas

      • 3.5.4 Indicator diagrams

      • 3.5.5 Spontaneous adiabatic expansion or compression

      • 3.5.6 Free expansion of a gas into a vacuum



    • 3.6 Work in a Gravitational Field

    • 3.7 Shaft Work

      • 3.7.1 Stirring work

      • 3.7.2 The Joule paddle wheel



    • 3.8 Electrical Work

      • 3.8.1 Electrical work in a circuit

      • 3.8.2 Electrical heating

      • 3.8.3 Electrical work with a galvanic cell



    • 3.9 Irreversible Work and Internal Friction

    • 3.10 Reversible and Irreversible Processes: Generalities

    • Problems



  • CONTENTS

  • 4 The Second Law

    • 4.1 Types of Processes

    • 4.2 Statements of the Second Law

    • 4.3 Concepts Developed with Carnot Engines

      • 4.3.1 Carnot engines and Carnot cycles

      • 4.3.2 The equivalence of the Clausius and Kelvin–Planck statements

      • 4.3.3 The efficiency of a Carnot engine

      • 4.3.4 Thermodynamic temperature



    • 4.4 Derivation of the Mathematical Statement of the Second Law

      • 4.4.1 The existence of the entropy function

      • 4.4.2 Using reversible processes to define the entropy

      • 4.4.3 Some properties of the entropy



    • 4.5 Irreversible Processes

      • 4.5.1 Irreversible adiabatic processes

      • 4.5.2 Irreversible processes in general



    • 4.6 Applications

      • 4.6.1 Reversible heating

      • 4.6.2 Reversible expansion of an ideal gas

      • 4.6.3 Spontaneous changes in an isolated system

      • 4.6.4 Internal heat flow in an isolated system

      • 4.6.5 Free expansion of a gas

      • 4.6.6 Adiabatic process with work



    • 4.7 Summary

    • 4.8 The Statistical Interpretation of Entropy

    • Problems



  • 5 Thermodynamic Potentials

    • 5.1 Total Differential of a Dependent Variable

    • 5.2 Total Differential of the Internal Energy

    • 5.3 Enthalpy, Helmholtz Energy, and Gibbs Energy

    • 5.4 Closed Systems

    • 5.5 Open Systems

    • 5.6 Expressions for Heat Capacity

    • 5.7 Surface Work

    • 5.8 Criteria for Spontaneity

    • Problems



  • 6 The Third Law and Cryogenics

    • 6.1 The Zero of Entropy

    • 6.2 Molar Entropies

      • 6.2.1 Third-law molar entropies

      • 6.2.2 Molar entropies from spectroscopic measurements

      • 6.2.3 Residual entropy



    • 6.3 Cryogenics

      • 6.3.1 Joule–Thomson expansion

      • 6.3.2 Magnetization





  • CONTENTS

    • Problem



  • 7 Pure Substances in Single Phases

    • 7.1 Volume Properties

    • 7.2 Internal Pressure

    • 7.3 Thermal Properties

      • 7.3.1 The relation betweenCV;mandCp;m.

      • 7.3.2 The measurement of heat capacities

      • 7.3.3 Typical values



    • 7.4 Heating at Constant Volume or Pressure

    • 7.5 Partial Derivatives with Respect toT,p, andV

      • 7.5.1 Tables of partial derivatives

      • 7.5.2 The Joule–Thomson coefficient



    • 7.6 Isothermal Pressure Changes

      • 7.6.1 Ideal gases

      • 7.6.2 Condensed phases



    • 7.7 Standard States of Pure Substances

    • 7.8 Chemical Potential and Fugacity

      • 7.8.1 Gases

      • 7.8.2 Liquids and solids



    • 7.9 Standard Molar Quantities of a Gas

    • Problems



  • 8 Phase Transitions and Equilibria of Pure Substances

    • 8.1 Phase Equilibria

      • 8.1.1 Equilibrium conditions

      • 8.1.2 Equilibrium in a multiphase system

      • 8.1.3 Simple derivation of equilibrium conditions

      • 8.1.4 Tall column of gas in a gravitational field

      • 8.1.5 The pressure in a liquid droplet

      • 8.1.6 The number of independent variables

      • 8.1.7 The Gibbs phase rule for a pure substance



    • 8.2 Phase Diagrams of Pure Substances

      • 8.2.1 Features of phase diagrams

      • 8.2.2 Two-phase equilibrium

      • 8.2.3 The critical point

      • 8.2.4 The lever rule

      • 8.2.5 Volume properties



    • 8.3 Phase Transitions

      • 8.3.1 Molar transition quantities

      • 8.3.2 Calorimetric measurement of transition enthalpies

      • 8.3.3 Standard molar transition quantities



    • 8.4 Coexistence Curves

      • 8.4.1 Chemical potential surfaces

      • 8.4.2 The Clapeyron equation

      • 8.4.3 The Clausius–Clapeyron equation





  • CONTENTS

    • Problems



  • 9 Mixtures

    • 9.1 Composition Variables

      • 9.1.1 Species and substances

      • 9.1.2 Mixtures in general

      • 9.1.3 Solutions

      • 9.1.4 Binary solutions

      • 9.1.5 The composition of a mixture



    • 9.2 Partial Molar Quantities

      • 9.2.1 Partial molar volume

      • 9.2.2 The total differential of the volume in an open system

      • 9.2.3 Evaluation of partial molar volumes in binary mixtures

      • 9.2.4 General relations

      • 9.2.5 Partial specific quantities

      • 9.2.6 The chemical potential of a species in a mixture

      • 9.2.7 Equilibrium conditions in a multiphase, multicomponent system

      • 9.2.8 Relations involving partial molar quantities



    • 9.3 Gas Mixtures

      • 9.3.1 Partial pressure

      • 9.3.2 The ideal gas mixture

      • 9.3.3 Partial molar quantities in an ideal gas mixture

      • 9.3.4 Real gas mixtures



    • 9.4 Liquid and Solid Mixtures of Nonelectrolytes

      • 9.4.1 Raoult’s law

      • 9.4.2 Ideal mixtures

      • 9.4.3 Partial molar quantities in ideal mixtures

      • 9.4.4 Henry’s law

      • 9.4.5 The ideal-dilute solution

      • 9.4.6 Solvent behavior in the ideal-dilute solution

      • 9.4.7 Partial molar quantities in an ideal-dilute solution



    • 9.5 Activity Coefficients in Mixtures of Nonelectrolytes

      • 9.5.1 Reference states and standard states

      • 9.5.2 Ideal mixtures

      • 9.5.3 Real mixtures

      • 9.5.4 Nonideal dilute solutions



    • 9.6 Evaluation of Activity Coefficients

      • 9.6.1 Activity coefficients from gas fugacities

      • 9.6.2 Activity coefficients from the Gibbs–Duhem equation

      • 9.6.3 Activity coefficients from osmotic coefficients

      • 9.6.4 Fugacity measurements



    • 9.7 Activity of an Uncharged Species

      • 9.7.1 Standard states

      • 9.7.2 Activities and composition

      • 9.7.3 Pressure factors and pressure



    • 9.8 Mixtures in Gravitational and Centrifugal Fields



  • CONTENTS

    • 9.8.1 Gas mixture in a gravitational field

    • 9.8.2 Liquid solution in a centrifuge cell

    • Problems



  • 10 Electrolyte Solutions

    • 10.1 Single-ion Quantities

    • 10.2 Solution of a Symmetrical Electrolyte

    • 10.3 Electrolytes in General

      • 10.3.1 Solution of a single electrolyte

      • 10.3.2 Multisolute solution

      • 10.3.3 Incomplete dissociation



    • 10.4 The Debye–Huckel Theory ̈

    • 10.5 Derivation of the Debye–Huckel Equation ̈

    • 10.6 Mean Ionic Activity Coefficients from Osmotic Coefficients

    • Problems



  • 11 Reactions and Other Chemical Processes

    • 11.1 Mixing Processes

      • 11.1.1 Mixtures in general

      • 11.1.2 Ideal mixtures

      • 11.1.3 Excess quantities

      • 11.1.4 The entropy change to form an ideal gas mixture

      • 11.1.5 Molecular model of a liquid mixture

      • 11.1.6 Phase separation of a liquid mixture



    • 11.2 The Advancement and Molar Reaction Quantities

      • 11.2.1 An example: ammonia synthesis

      • 11.2.2 Molar reaction quantities in general

      • 11.2.3 Standard molar reaction quantities



    • 11.3 Molar Reaction Enthalpy

      • 11.3.1 Molar reaction enthalpy and heat

      • 11.3.2 Standard molar enthalpies of reaction and formation

      • 11.3.3 Molar reaction heat capacity

      • 11.3.4 Effect of temperature on reaction enthalpy



    • 11.4 Enthalpies of Solution and Dilution

      • 11.4.1 Molar enthalpy of solution

      • 11.4.2 Enthalpy of dilution

      • 11.4.3 Molar enthalpies of solute formation

      • 11.4.4 Evaluation of relative partial molar enthalpies



    • 11.5 Reaction Calorimetry

      • 11.5.1 The constant-pressure reaction calorimeter

      • 11.5.2 The bomb calorimeter

      • 11.5.3 Other calorimeters



    • 11.6 Adiabatic Flame Temperature

    • 11.7 Gibbs Energy and Reaction Equilibrium

      • 11.7.1 The molar reaction Gibbs energy

      • 11.7.2 Spontaneity and reaction equilibrium





  • CONTENTS

    • 11.7.3 General derivation

    • 11.7.4 Pure phases

    • 11.7.5 Reactions involving mixtures

    • 11.7.6 Reaction in an ideal gas mixture

    • 11.8 The Thermodynamic Equilibrium Constant

      • 11.8.1 Activities and the definition ofK

      • 11.8.2 Reaction in a gas phase

      • 11.8.3 Reaction in solution

      • 11.8.4 Evaluation ofK



    • 11.9 Effects of Temperature and Pressure on Equilibrium Position

    • Problems



  • 12 Equilibrium Conditions in Multicomponent Systems

    • 12.1 Effects of Temperature

      • 12.1.1 Variation ofi=T with temperature

      • 12.1.2 Variation ofi=Twith temperature

      • 12.1.3 Variation of lnKwith temperature



    • 12.2 Solvent Chemical Potentials from Phase Equilibria

      • 12.2.1 Freezing-point measurements

      • 12.2.2 Osmotic-pressure measurements



    • 12.3 Binary Mixture in Equilibrium with a Pure Phase

    • 12.4 Colligative Properties of a Dilute Solution

      • 12.4.1 Freezing-point depression

      • 12.4.2 Boiling-point elevation

      • 12.4.3 Vapor-pressure lowering

      • 12.4.4 Osmotic pressure



    • 12.5 Solid–Liquid Equilibria

      • 12.5.1 Freezing points of ideal binary liquid mixtures

      • 12.5.2 Solubility of a solid nonelectrolyte

      • 12.5.3 Ideal solubility of a solid

      • 12.5.4 Solid compound of mixture components

      • 12.5.5 Solubility of a solid electrolyte



    • 12.6 Liquid–Liquid Equilibria

      • 12.6.1 Miscibility in binary liquid systems

      • 12.6.2 Solubility of one liquid in another

      • 12.6.3 Solute distribution between two partially-miscible solvents



    • 12.7 Membrane Equilibria

      • 12.7.1 Osmotic membrane equilibrium

      • 12.7.2 Equilibrium dialysis

      • 12.7.3 Donnan membrane equilibrium



    • 12.8 Liquid–Gas Equilibria

      • 12.8.1 Effect of liquid pressure on gas fugacity

      • 12.8.2 Effect of liquid composition on gas fugacities

      • 12.8.3 The Duhem–Margules equation

      • 12.8.4 Gas solubility

      • 12.8.5 Effect of temperature and pressure on Henry’s law constants





  • CONTENTS

    • 12.9 Reaction Equilibria

    • 12.10 Evaluation of Standard Molar Quantities

    • Problems



  • 13 The Phase Rule and Phase Diagrams

    • 13.1 The Gibbs Phase Rule for Multicomponent Systems

      • 13.1.1 Degrees of freedom

      • 13.1.2 Species approach to the phase rule

      • 13.1.3 Components approach to the phase rule

      • 13.1.4 Examples



    • 13.2 Phase Diagrams: Binary Systems

      • 13.2.1 Generalities

      • 13.2.2 Solid–liquid systems

      • 13.2.3 Partially-miscible liquids

      • 13.2.4 Liquid–gas systems with ideal liquid mixtures

      • 13.2.5 Liquid–gas systems with nonideal liquid mixtures

      • 13.2.6 Solid–gas systems

      • 13.2.7 Systems at high pressure



    • 13.3 Phase Diagrams: Ternary Systems

      • 13.3.1 Three liquids

      • 13.3.2 Two solids and a solvent



    • Problems



  • 14 Galvanic Cells

    • 14.1 Cell Diagrams and Cell Reactions

      • 14.1.1 Elements of a galvanic cell

      • 14.1.2 Cell diagrams

      • 14.1.3 Electrode reactions and the cell reaction

      • 14.1.4 Advancement and charge



    • 14.2 Electric Potentials in the Cell

      • 14.2.1 Cell potential

      • 14.2.2 Measuring the equilibrium cell potential

      • 14.2.3 Interfacial potential differences



    • 14.3 Molar Reaction Quantities of the Cell Reaction

      • 14.3.1 Relation betweenÅrGcellandEcell, eq

      • 14.3.2 Relation betweenÅrGcellandÅrG

      • 14.3.3 Standard molar reaction quantities



    • 14.4 The Nernst Equation

    • 14.5 Evaluation of the Standard Cell Potential

    • 14.6 Standard Electrode Potentials

    • Problems



  • Appendix A Definitions of the SI Base Units

  • Appendix B Physical Constants

  • CONTENTS

  • Appendix C Symbols for Physical Quantities

  • Appendix D Miscellaneous Abbreviations and Symbols

    • D.1 Physical States

    • D.2 Subscripts for Chemical Processes

    • D.3 Superscripts



  • Appendix E Calculus Review

    • E.1 Derivatives

    • E.2 Partial Derivatives

    • E.3 Integrals

    • E.4 Line Integrals



  • Appendix F Mathematical Properties of State Functions

    • F.1 Differentials

    • F.2 Total Differential

    • F.3 Integration of a Total Differential

    • F.4 Legendre Transforms



  • Appendix G Forces, Energy, and Work

    • G.1 Forces between Particles

    • G.2 The System and Surroundings

    • G.3 System Energy Change

    • G.4 Macroscopic Work

    • G.5 The Work Done on the System and Surroundings

    • G.6 The Local Frame and Internal Energy

    • G.7 Nonrotating Local Frame

    • G.8 Center-of-mass Local Frame

    • G.9 Rotating Local Frame

    • G.10 Earth-Fixed Reference Frame



  • Appendix H Standard Molar Thermodynamic Properties

  • Appendix I Answers to Selected Problems

  • Bibliography

  • Index

  • Benjamin Thompson, Count of Rumford BIOGRAPHICAL SKETCHES

  • James Prescott Joule

  • Sadi Carnot

  • Rudolf Julius Emmanuel Clausius

  • William Thomson, Lord Kelvin

  • Max Karl Ernst Ludwig Planck

  • Josiah Willard Gibbs

  • Walther Hermann Nernst

  • William Francis Giauque

  • Benoit PaulEmile Clapeyron ́

  • William Henry

  • Gilbert Newton Lewis

  • Peter Josephus Wilhelmus Debye

  • Germain Henri Hess

  • Franc ̧ois-Marie Raoult


  • Jacobus Henricus van’t Hoff



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