Thermochemistry

The study of heat energy associated with chemical reactions and physical transformations.

Basic Concepts

System and Surroundings

  • System: Part of universe under study
  • Surroundings: Everything else
  • Universe: System + Surroundings

Types of Systems

Type Energy Exchange Matter Exchange
Open Yes Yes
Closed Yes No
Isolated No No

State Functions

Properties depending only on initial and final states:

  • Enthalpy (H)
  • Internal energy (U)
  • Entropy (S)
  • Gibbs free energy (G)

First Law of Thermodynamics

$$ΔU = q + w$$

Where:

  • ΔU = change in internal energy
  • q = heat absorbed by system
  • w = work done on system

Enthalpy (H)

Heat content at constant pressure: H = U + PV

Types of Enthalpy Changes

Symbol Name Definition
ΔH°f Standard enthalpy of formation One mole from elements in standard states
ΔH°c Standard enthalpy of combustion One mole burns in O₂
ΔH°neut Enthalpy of neutralization Acid + base → salt + water
ΔH°sol Enthalpy of solution One mole dissolves
ΔH°vap Enthalpy of vaporization Liquid → gas
ΔH°fus Enthalpy of fusion Solid → liquid
ΔH°sub Enthalpy of sublimation Solid → gas
ΔH°at Enthalpy of atomization One mole to gaseous atoms

Standard Conditions

  • 1 bar pressure
  • 298 K (25°C) temperature
  • 1 M concentration for solutions
  • Specified physical state

Calorimetry

Coffee-Cup Calorimeter (Constant Pressure)

$$q = mcΔT$$

Where:

  • m = mass (g)
  • c = specific heat capacity (J/g·K)
  • ΔT = temperature change

At constant pressure: q = ΔH

Bomb Calorimeter (Constant Volume)

  • Measures ΔU directly
  • qv = Cv × ΔT

Hess's Law

The total enthalpy change for a reaction is the same, regardless of the pathway taken.

Applications

Using Formation Enthalpies

$$ΔH°_{rxn} = ΣΔH°_f(products) - ΣΔH°_f(reactants)$$

Using Bond Enthalpies

$$ΔH°_{rxn} = Σ(Bonds\ broken) - Σ(Bonds\ formed)$$

  • Breaking bonds: Endothermic (+)
  • Forming bonds: Exothermic (-)

Born-Haber Cycle

Hess's law application for ionic compounds:

  • Atomization of metal
  • Ionization of metal
  • Atomization of non-metal
  • Electron affinity of non-metal
  • Lattice energy
  • Formation enthalpy

Enthalpy of Hydration (ΔH°hyd)

Energy released when one mole of gaseous ions is surrounded by water molecules:

  • Cations: ΔH°hyd is negative (exothermic); smaller ions and higher charges → more negative
  • Anions: ΔH°hyd is negative; smaller ions → more negative
  • Overall: ΔH°hyd(total) = ΔH°hyd(cation) + ΔH°hyd(anion)

Relationship: Enthalpy of Solution

$ΔH°{sol} = ΔH°{lattice} + ΔH°_{hyd}$

  • If |ΔH°hyd| > |ΔH°lattice|: ΔH°sol is negative (exothermic dissolution)
  • If |ΔH°hyd| < |ΔH°lattice|: ΔH°sol is positive (endothermic dissolution)

Lattice Energy

Energy required to separate one mole of solid ionic compound into gaseous ions.

Factors affecting lattice energy:

  1. Ionic charge: Higher charge → Higher LE
  2. Ionic radius: Smaller ions → Higher LE

Spontaneity and Gibbs Free Energy

$$ΔG = ΔH - TΔS$$

ΔG Spontaneity
< 0 Spontaneous
= 0 Equilibrium
> 0 Non-spontaneous

Temperature Dependence

ΔH ΔS Spontaneity
- + Always spontaneous
- - Spontaneous at low T
+ + Spontaneous at high T
+ - Never spontaneous

Related Topics

Sources