1.1.1 Amount of substance and amount

2.1.1 Extensive and intensive properties

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

2.3.1 Mass

2.3.2 Volume

2.3.3 Density

2.3.4 Pressure

2.3.5 Temperature

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

2.6.1 Energy and reference frames

2.6.2 Internal energy

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

3.2.1 Reversible processes

3.2.2 Irreversible processes

3.2.3 Purely mechanical processes

3.3.1 Heating and cooling

3.3.2 Spontaneous phase transitions

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

3.5.1 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

3.7.1 Stirring work

3.7.2 The Joule paddle wheel

3.8.1 Electrical work in a circuit

3.8.2 Electrical heating

3.8.3 Electrical work with a galvanic cell

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

4.4.1 The existence of the entropy function

4.4.2 Using reversible processes to define the entropy

4.4.3 Some properties of the entropy

4.5.1 Irreversible adiabatic processes

4.5.2 Irreversible processes in general

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

6.2.1 Third-law molar entropies

6.2.2 Molar entropies from spectroscopic measurements

6.2.3 Residual entropy

6.3.1 Joule–Thomson expansion

6.3.2 Magnetization

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

7.5.1 Tables of partial derivatives

7.5.2 The Joule–Thomson coefficient

7.6.1 Ideal gases

7.6.2 Condensed phases

7.8.1 Gases

7.8.2 Liquids and solids

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

8.2.1 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

8.3.1 Molar transition quantities

8.3.2 Calorimetric measurement of transition enthalpies

8.3.3 Standard molar transition quantities

8.4.1 Chemical potential surfaces

8.4.2 The Clapeyron equation

8.4.3 The Clausius–Clapeyron equation

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

9.2.1 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

9.3.1 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

9.4.1 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

9.5.1 Reference states and standard states

9.5.2 Ideal mixtures

9.5.3 Real mixtures

9.5.4 Nonideal dilute solutions

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

9.7.1 Standard states

9.7.2 Activities and composition

9.7.3 Pressure factors and pressure

9.8.1 Gas mixture in a gravitational field

9.8.2 Liquid solution in a centrifuge cell

10.3.1 Solution of a single electrolyte

10.3.2 Multisolute solution

10.3.3 Incomplete dissociation

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

11.2.1 An example: ammonia synthesis

11.2.2 Molar reaction quantities in general

11.2.3 Standard molar reaction quantities

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

11.4.1 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

11.5.1 The constant-pressure reaction calorimeter

11.5.2 The bomb calorimeter

11.5.3 Other calorimeters

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

11.8.1 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

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

12.2.1 Freezing-point measurements

12.2.2 Osmotic-pressure measurements

12.4.1 Freezing-point depression

12.4.2 Boiling-point elevation

12.4.3 Vapor-pressure lowering

12.4.4 Osmotic pressure

12.5.1 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

12.6.1 Miscibility in binary liquid systems

12.6.2 Solubility of one liquid in another

12.6.3 Solute distribution between two partially-miscible solvents

12.7.1 Osmotic membrane equilibrium

12.7.2 Equilibrium dialysis

12.7.3 Donnan membrane equilibrium

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

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

13.2.1 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

13.3.1 Three liquids

13.3.2 Two solids and a solvent

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

14.2.1 Cell potential

14.2.2 Measuring the equilibrium cell potential

14.2.3 Interfacial potential differences

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