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) + 3H_2(g) + 1⁄2O_2(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: 1⁄2Cl_2(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^- ➔ O^2-_(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.

Lattice energy measures ionic bond strength

Lattice energy (ΔH^⦵_latt), also known as lattice enthalpy, is a measure of the strength of the electrostatic forces holding ions together in an ionic lattice. It is defined as the enthalpy change when one mole of a solid ionic compound is formed from its gaseous ions under standard conditions (298 K, 100 kPa).

For example: Na⁺_(g) + Cl^-_(g) ➔ NaCl_(s) ΔH^⦵_latt = −787 kJ mol^-1

Lattice energy values are always negative (exothermic) because energy is released when the oppositely charged ions come together to form the solid lattice.

The more negative the lattice energy, the stronger the ionic bonding in the compound. For instance, MgO has a more negative lattice energy ($-$3,791 kJ mol^-1) than NaCl ($-$787 kJ mol^-1), indicating that MgO has stronger ionic bonds.

Factors affecting lattice energy

The lattice energy of an ionic compound depends on two key factors.

1. Ionic charge

• Ions with higher charges experience stronger electrostatic attractions than ions with lower charges.
• This leads to more energy being released when the lattice forms, resulting in a more negative lattice energy.
• For example, MgCl₂ has a much more negative lattice energy ($-$2,526 kJ mol^-1) than NaCl ($-$787 kJ mol^-1) because the Mg²⁺ ion has a higher charge than the Na⁺ ion, resulting in stronger electrostatic attractions in the lattice.

2. Ionic radius

• Smaller ions have a higher charge density and can pack more closely together in the lattice.
• This increases the strength of the electrostatic attractions between ions.
• Consequently, compounds with smaller ions tend to have more negative lattice energies.
• For example, LiCl has a more negative lattice energy ($-$853 kJ mol^-1) than NaCl ($-$787 kJ mol^-1) because the Li⁺ ion is smaller than the Na⁺ ion, allowing closer packing in the lattice.

So in general, compounds with small, highly charged ions have the most negative lattice energies.