Ionic bonding and lattice enthalpy

An ionic bond forms most easily between:

• a metal with a low ionisation enthalpy (it loses electrons easily, forming a cation); and
• a non-metal with a large negative electron gain enthalpy (it gains electrons readily, forming an anion).

But making the ions alone does not explain why the compound forms. Take sodium chloride. In the gas phase, removing an electron from Na costs 495.8 kJ mol⁻¹, and adding it to Cl gives back only 348.7 kJ mol⁻¹. So making separate Na⁺ and Cl⁻ ions costs 147.1 kJ mol⁻¹ overall.

What makes it worthwhile is what happens next: the oppositely charged ions pack together into a crystal lattice, and a great deal of energy is released. The lattice enthalpy is the energy needed to separate one mole of an ionic solid completely into its gaseous ions. For NaCl it is 788 kJ mol⁻¹, far more than the 147.1 kJ needed to make the ions.

So the stability of an ionic compound comes mainly from its lattice enthalpy, not simply from its ions reaching noble-gas octets. The larger the lattice enthalpy, the more stable the solid. Small ions with high charges attract each other strongly, giving large lattice enthalpies: magnesium oxide (Mg²⁺ and O²⁻) has a far higher lattice enthalpy than sodium chloride.