5.1.4 Comparing Lattice Enthalpies

Lattice Enthalpy

Lattice enthalpy represents the strength of the ionic bonds within an ionic compound and can be defined in two ways:

1. Enthalpy of Lattice Formation: The enthalpy change when 1 mole of an ionic compound is formed from its gaseous ions.
2. Enthalpy of Lattice Dissociation: The enthalpy change required to separate 1 mole of an ionic compound into its gaseous ions. A high lattice enthalpy indicates strong ionic bonding, often seen with small, highly charged ions.

Factors Affecting Lattice Enthalpy

1. Ionic Radius: Smaller ions allow closer packing, increasing the attraction between oppositely charged ions, which increases lattice enthalpy.
2. Ionic Charge: Higher charges on ions lead to stronger electrostatic attraction and a greater lattice enthalpy.

The Perfect Ionic Model

The perfect ionic model assumes that:

• The bonding in the compound is entirely ionic, with no covalent character.
• Ions are perfect spheres with a uniform charge distribution. However, this model is idealized. In reality, some ionic compounds display covalent character due to the distortion of ions. When a small, highly charged cation distorts a larger anion, the electron cloud of the anion shifts, creating partial covalent bonding.

Comparing Experimental and Theoretical Lattice Enthalpies

1. Theoretical Lattice Enthalpy: Calculated based on the perfect ionic model, which assumes ideal ionic bonding and spherical ions.
2. Experimental Lattice Enthalpy: Obtained from the Born-Haber cycle, which relies on actual experimental values for enthalpy changes such as atomisation, ionisation energy, and electron affinity.

Evidence for Covalent Character in Ionic Compounds

If an ionic compound has covalent character, the experimental lattice enthalpy (from the Born-Haber cycle) is often higher than the theoretical lattice enthalpy predicted by the perfect ionic model.

• The larger the difference between these two values, the greater the degree of covalent character in the ionic compound.
• This deviation highlights the presence of covalent character, as the actual forces of attraction are stronger than what the perfect ionic model predicts.