Physical Chemistry

2.

2.1.

Fundamental Concepts

2.1.1.

System Definitions

2.1.1.1.

2.1.1.2.

2.1.1.3.

2.1.2.

Types of Systems

2.1.2.1.

2.1.2.2.

2.1.2.3.

Isolated Systems

2.1.3.

Properties of Systems

2.1.3.1.

Extensive Properties

2.1.3.2.

Intensive Properties

2.1.3.3.

State Functions

2.1.3.4.

2.1.4.

Thermodynamic Equilibrium

2.1.4.1.

Mechanical Equilibrium

2.1.4.2.

Thermal Equilibrium

2.1.4.3.

Chemical Equilibrium

2.1.4.4.

Phase Equilibrium

2.1.5.

Thermodynamic Processes

2.1.5.1.

Reversible Processes

2.1.5.2.

Irreversible Processes

2.1.5.3.

Isothermal Processes

2.1.5.4.

Isobaric Processes

2.1.5.5.

Isochoric Processes

2.1.5.6.

Adiabatic Processes

2.1.5.7.

Cyclic Processes

2.2.

The Zeroth Law of Thermodynamics

2.2.1.

Concept of Temperature

2.2.1.1.

Thermal Contact

2.2.1.2.

Temperature as Intensive Property

2.2.2.

Thermal Equilibrium

2.2.2.1.

Transitive Property

2.2.3.

Temperature Scales

2.2.3.1.

2.2.3.2.

2.2.3.3.

Fahrenheit Scale

2.2.3.4.

2.3.

The First Law of Thermodynamics

2.3.1.

Internal Energy

2.3.1.1.

Definition and Physical Meaning

2.3.1.2.

State Function Nature

2.3.1.3.

Molecular Interpretation

2.3.2.

2.3.2.1.

Definitions and Sign Conventions

2.3.2.2.

Pressure-Volume Work

2.3.2.3.

2.3.2.4.

Other Types of Work

2.3.3.

Mathematical Statement of the First Law

2.3.3.1.

ΔU = q + w for Closed Systems

2.3.3.2.

Differential Form: dU = δq + δw

2.3.3.3.

Cyclic Processes

2.3.4.

2.3.4.1.

Definition: H = U + PV

2.3.4.2.

Physical Meaning

2.3.4.3.

Enthalpy Changes at Constant Pressure

2.3.4.4.

Relationship to Heat Capacity

2.3.5.

2.3.5.1.

Definition and Measurement

2.3.5.2.

Heat Capacity at Constant Volume

2.3.5.3.

Heat Capacity at Constant Pressure

2.3.5.4.

Relationship between Cp and Cv

2.3.5.5.

Temperature Dependence

2.3.6.

Thermochemistry

2.3.6.1.

Standard States and Conditions

2.3.6.2.

Standard Enthalpy of Formation

2.3.6.3.

Standard Enthalpy of Combustion

2.3.6.4.

Standard Enthalpy of Reaction

2.3.6.5.

2.3.6.6.

Kirchhoff's Law

2.3.6.7.

2.3.6.7.1.

Bomb Calorimetry

2.3.6.7.2.

Solution Calorimetry

2.4.

The Second Law of Thermodynamics

2.4.1.

Spontaneous Processes

2.4.1.1.

Characteristics of Spontaneous Processes

2.4.1.2.

Criteria for Spontaneity

2.4.1.3.

Directionality of Natural Processes

2.4.2.

2.4.2.1.

Definition and Physical Meaning

2.4.2.2.

Entropy as a State Function

2.4.2.3.

Molecular Interpretation

2.4.3.

Mathematical Formulations

2.4.3.1.

The Clausius Inequality

2.4.3.2.

Entropy Change for Reversible Processes

2.4.3.3.

Entropy Change for Irreversible Processes

2.4.4.

Entropy Changes in Specific Processes

2.4.4.1.

Isothermal Expansion of Ideal Gas

2.4.4.2.

Phase Transitions

2.4.4.3.

Mixing of Ideal Gases

2.4.4.4.

Temperature Changes

2.4.5.

Heat Engines and Refrigerators

2.4.5.1.

The Carnot Cycle

2.4.5.2.

Carnot Engine Efficiency

2.4.5.3.

Carnot Refrigerator

2.4.5.4.

Real Heat Engines

2.4.6.

Statistical Interpretation of Entropy

2.4.6.1.

Microstates and Macrostates

2.4.6.2.

Boltzmann's Formula: S = k ln W

2.4.6.3.

Probability and Disorder

2.5.

The Third Law of Thermodynamics

2.5.1.

Statement of the Third Law

2.5.1.1.

Perfect Crystal at Absolute Zero

2.5.1.2.

Zero Entropy at 0 K

2.5.2.

Absolute Entropy

2.5.2.1.

Calculation from Heat Capacity Data

2.5.2.2.

Standard Molar Entropies

2.5.2.3.

Tabulated Values

2.5.3.

Consequences of the Third Law

2.5.3.1.

Unattainability of Absolute Zero

2.5.3.2.

Heat Capacity Behavior at Low Temperatures

2.6.

Thermodynamic Potentials

2.6.1.

Helmholtz Energy

2.6.1.1.

Definition: A = U - TS

2.6.1.2.

Physical Significance

2.6.1.3.

Applications at Constant T and V

2.6.2.

2.6.2.1.

Definition: G = H - TS

2.6.2.2.

Physical Significance

2.6.2.3.

Applications at Constant T and P

2.6.3.

Criteria for Spontaneity and Equilibrium

2.6.3.1.

Isolated Systems (ΔS ≥ 0)

2.6.3.2.

Constant T and V (ΔA ≤ 0)

2.6.3.3.

Constant T and P (ΔG ≤ 0)

2.6.4.

Maxwell Relations

2.6.4.1.

Derivation from Exact Differentials

2.6.4.2.

Applications in Thermodynamics

2.6.5.

The Gibbs-Helmholtz Equation

2.6.5.1.

Temperature Dependence of G

2.6.5.2.

Applications to Chemical Reactions

2.7.

Phase Equilibria of Pure Substances

2.7.1.

Chemical Potential

2.7.1.1.

Definition and Physical Meaning

2.7.1.2.

Relationship to Gibbs Energy

2.7.2.

Phase Equilibrium Conditions

2.7.2.1.

Equality of Chemical Potentials

2.7.2.2.

Phase Boundaries

2.7.3.

The Clapeyron Equation

2.7.3.1.

2.7.3.2.

Applications to Phase Transitions

2.7.4.

The Clausius-Clapeyron Equation

2.7.4.1.

Vapor Pressure-Temperature Relationship

2.7.4.2.

Applications to Vaporization and Sublimation

2.7.5.

2.7.5.1.

Construction and Interpretation

2.7.5.2.

One-Component Systems

2.7.5.3.

Water Phase Diagram

2.7.5.4.

Carbon Dioxide Phase Diagram

2.7.5.5.

2.7.5.6.

2.7.5.7.

Supercritical Fluids

2.8.

Properties of Mixtures and Solutions

2.8.1.

Partial Molar Quantities

2.8.1.1.

Definition and Physical Meaning

2.8.1.2.

Partial Molar Volume

2.8.1.3.

Partial Molar Enthalpy

2.8.1.4.

Partial Molar Entropy

2.8.2.

The Gibbs-Duhem Equation

2.8.2.1.

2.8.2.2.

Physical Significance

2.8.3.

Ideal Solutions

2.8.3.1.

Definition and Characteristics

2.8.3.2.

Raoult's Law for Ideal Solutions

2.8.3.3.

Thermodynamic Properties of Mixing

2.8.4.

2.8.4.1.

Deviations from Ideality

2.8.4.2.

Positive and Negative Deviations

2.8.5.

2.8.5.1.

Statement and Applications

2.8.5.2.

Sparingly Soluble Gases

2.8.5.3.

Temperature Dependence

2.8.6.

Colligative Properties

2.8.6.1.

Vapor Pressure Lowering

2.8.6.2.

Boiling Point Elevation

2.8.6.3.

Freezing Point Depression

2.8.6.4.

Osmotic Pressure

2.8.6.5.

Applications and Calculations

2.8.7.

Activity and Activity Coefficients

2.8.7.1.

Definition of Activity

2.8.7.2.

Activity Coefficients

2.8.7.3.

Standard States

2.8.7.4.

Debye-Hückel Theory for Ionic Solutions

2.9.

Chemical Equilibrium

2.9.1.

Equilibrium Condition

2.9.1.1.

Minimum Gibbs Energy

2.9.1.2.

Equality of Chemical Potentials

2.9.2.

The Reaction Gibbs Energy

2.9.2.1.

Definition: ΔrG

2.9.2.2.

Relationship to Spontaneity

2.9.2.3.

Standard Reaction Gibbs Energy

2.9.3.

The Equilibrium Constant

2.9.3.1.

Definition and Types

2.9.3.2.

Concentration Equilibrium Constant (Kc)

2.9.3.3.

Pressure Equilibrium Constant (Kp)

2.9.3.4.

Activity Equilibrium Constant (Ka)

2.9.4.

Relationship between ΔrG° and K

2.9.4.1.

ΔrG° = -RT ln K

2.9.4.2.

Standard State Considerations

2.9.5.

Temperature Dependence of Equilibrium Constants

2.9.5.1.

van 't Hoff Equation

2.9.5.2.

Integrated Form

2.9.6.

Pressure Dependence of Equilibrium Constants

2.9.6.1.

Effect of Pressure on Gas-Phase Equilibria

2.9.7.

Le Châtelier's Principle

2.9.7.1.

Response to Concentration Changes

2.9.7.2.

Response to Pressure Changes

2.9.7.3.

Response to Temperature Changes

2.9.7.4.

Catalysts and Equilibrium

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