Lattice Enthalpy

This lesson covers: 

  1. Definitions of enthalpy change and standard conditions
  2. Types of enthalpy change
  3. Types of lattice enthalpy

Defining enthalpy change

Enthalpy change (ΔH) is the heat energy transferred in a reaction at constant pressure. Its units are kilojoules per mole (kJ mol-1).

  • ΔH⦵ indicates that the enthalpy change was measured under standard conditions of 298 K and 100 kPa.
  • Exothermic reactions release heat energy to the surroundings and have a negative ΔH value.
  • Endothermic reactions absorb heat energy from the surroundings and have a positive ΔH value.

Types of enthalpy change

There are many types of enthalpy change, each referring to a specific chemical or physical process:

  1. Enthalpy change of formation (ΔH⦵f) - The enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions. For example:

2C(s) + 3H2(g) + 1⁄2O2(g) ➔ C2H5OH(l)

ΔH⦵f values are usually exothermic as energy is released when bonds form between the elements to make the compound.


  1. Enthalpy change of atomisation (ΔH⦵at) - The enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state under standard conditions. For example:

1⁄2Cl2(g) ➔ Cl(g)

ΔH⦵at values are always endothermic as energy must be supplied to break the bonds holding the atoms in the element together.


  1. First ionisation energy (ΔH⦵IE1) - The enthalpy change when 1 mole of gaseous 1+ ions is formed from 1 mole of gaseous atoms. For example:

Mg(g) ➔ Mg+(g) + e-

ΔH⦵IE1 values are always endothermic as energy is needed to overcome the electrostatic attraction between the nucleus and the electron being removed.


  1. Second ionisation energy (ΔH⦵IE2) - The enthalpy change when 1 mole of gaseous 2+ ions is formed from 1 mole of gaseous 1+ ions. For example:

Mg+(g) ➔ Mg2+(g) + e-

ΔH⦵IE2 values are always endothermic as even more energy is required to remove an electron from a positively charged ion.


  1. First electron affinity (ΔH⦵ea1) - The enthalpy change when 1 mole of gaseous 1- ions is formed from 1 mole of gaseous atoms. For example:

O(g) + e- ➔ O-(g)

ΔH⦵ea1 values are usually exothermic as the attraction between the nucleus and the incoming electron releases energy.


  1. Second electron affinity (ΔH⦵ea2) - The enthalpy change when 1 mole of gaseous 2- ions is formed from 1 mole of gaseous 1- ions. For example:

O-(g) + e- ➔ O2-(g)

ΔH⦵ea2 values are always endothermic as energy must be supplied to overcome the repulsion between the negative ion and the second incoming electron.


  1. Bond enthalpy (ΔH⦵diss) - The enthalpy change when 1 mole of a particular covalent bond in the gaseous state is broken. For example:

Cl2(g) ➔ 2Cl(g)

ΔH⦵diss values are always endothermic as energy is required to overcome the attraction between the bonded atoms and break the bond.

Lattice enthalpy measures ionic bond strength

Lattice enthalpy is a measure of the strength of the electrostatic forces holding ions together in an ionic lattice.


It can be defined in two ways:

  1. Lattice enthalpy of formation (ΔH⦵latt) - The enthalpy change when one mole of a solid ionic compound is formed from its gaseous ions under standard conditions (298 K, 100 kPa). When the term "lattice enthalpy" is used without specifying formation or dissociation, it generally refers to the lattice enthalpy of formation.

For example:  Na+(g) + Cl-(g) ➔ NaCl(s)   ΔH⦵latt = −787 kJ mol-1 (exothermic)

Lattice enthalpy of formation values are always exothermic as energy is released when the oppositely charged ions come together to form the solid lattice.


  1. Lattice enthalpy of dissociation (ΔH⦵latt) - The enthalpy change when 1 mole of a solid ionic compound is completely dissociated into its gaseous ions under standard conditions. This is the reverse of the lattice enthalpy of formation.

For example:  NaCl(s) ➔ Na+(g) + Cl-(g)   ΔH⦵latt = +787 kJ mol-1 (endothermic)

Lattice enthalpy of dissociation values are always endothermic as energy must be supplied to overcome the electrostatic attractions between ions and separate them into the gas phase.


The more negative the lattice enthalpy of formation (or the more positive the lattice enthalpy of dissociation), the stronger the ionic bonding in the compound.

For instance, MgO has a more negative lattice enthalpy of formation (-3,791 kJ mol-1) than NaCl (-787 kJ mol-1), indicating that MgO has stronger ionic bonds.