Enthalpies of Hydration and Solution

When an ionic compound dissolves in water, two processes occur: the ionic lattice must be broken apart, and the separated ions must become surrounded by water molecules. Understanding the enthalpy changes associated with these processes allows us to predict whether dissolving will be exothermic or endothermic, and even to rationalise why some salts are more soluble than others.

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Enthalpy of Hydration (ΔH°hyd)

The standard enthalpy of hydration is the enthalpy change when one mole of gaseous ions becomes surrounded by water molecules to form an aqueous solution (at infinite dilution).

Na⁺(g) → Na⁺(aq) ΔH°hyd = −406 kJ mol⁻¹

Cl⁻(g) → Cl⁻(aq) ΔH°hyd = −363 kJ mol⁻¹

Key points:

• Hydration is always exothermic (negative ΔH°hyd) because ion-dipole attractions form between the ion and surrounding water molecules, releasing energy.
• Hydration enthalpies are more exothermic for ions with higher charge and smaller radius (higher charge density), because these ions attract water molecules more strongly.
• Cations are typically more strongly hydrated than anions of similar size because cations are generally smaller.

Trends in Hydration Enthalpy

For Group 1 cations: