Thermodynamics and Chemistry
Free

Thermodynamics and Chemistry

By Howard DeVoe
Free
Book Description
Table of Contents
  • Cover
  • Title page
  • Copyright page
  • Short Contents
  • Contents
  • Biographical Sketches
  • Preface to the Second Edition
  • From the Preface to the First Edition
  • Epigraphs
  • 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
      • 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
  • 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
    • Problem
  • 7 Pure Substances in Single Phases
    • 7.1 Volume Properties
    • 7.2 Internal Pressure
    • 7.3 Thermal Properties
      • 7.3.1 The relation between C(V,m) and C(p,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 to T, p, and V
      • 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
    • 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
      • 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–Hückel Theory
    • 10.5 Derivation of the Debye–Hückel 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
      • 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 of K
      • 11.8.2 Reaction in a gas phase
      • 11.8.3 Reaction in solution
      • 11.8.4 Evaluation of K
    • 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 of mu(i)/T with temperature
      • 12.1.2 Variation of mu(i)o/T with temperature
      • 12.1.3 Variation of ln K with 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
    • 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 between molar reaction Gibbs energy and equilibrium cell potential
      • 14.3.2 Relation between molar reaction Gibbs energies of the cell and the direct reaction
      • 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
  • A Definitions of the SI Base Units
  • B Physical Constants
  • C Symbols for Physical Quantities
  • D Miscellaneous Abbreviations and Symbols
    • D.1 Physical States
    • D.2 Subscripts for Chemical Processes
    • D.3 Superscripts
  • E Calculus Review
    • E.1 Derivatives
    • E.2 Partial Derivatives
    • E.3 Integrals
    • E.4 Line Integrals
  • F Mathematical Properties of State Functions
    • F.1 Differentials
    • F.2 Total Differential
    • F.3 Integration of a Total Differential
    • F.4 Legendre Transforms
  • 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
  • H Standard Molar Thermodynamic Properties
  • I Answers to Selected Problems
  • Bibliography
  • Index
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