23.2 Enthalpies of solution and hydration

Hydration enthalpy is the enthalpy change when one mole of gaseous ions becomes aqueous. Solution enthalpy is the enthalpy change when one mole of an ionic solid dissolves to form aqueous ions.

Hydration involves ion–water attractions; solution combines lattice separation and hydration. State the ion charge, state symbols and direction before comparing values.

Dissolving NaCl(s) gives Na⁺(aq) and Cl⁻(aq); the solution enthalpy equals lattice dissociation plus the two hydration enthalpies.

Hydration is not the same as dissolving a molecular solute, and solution enthalpy is not automatically exothermic.

A solution cycle connects the direct dissolution route with an indirect route: lattice dissociation of the solid followed by hydration of the gaseous ions. Hess’s law makes the two routes equivalent.

Choose lattice formation or dissociation consistently. The indirect route is ΔHsol = ΔHlatt(dissociation) + ΣΔHhyd, with signs determined by direction.

For MgCl₂, the cycle contains one Mg²⁺ hydration and two Cl⁻ hydration steps. Omitting the coefficient on chloride gives a wrong energy balance.

Do not use one hydration value for the whole salt or mix lattice formation and dissociation arrows.

Write the known enthalpy values on a labelled cycle, choose a sign convention, then use Hess’s law to isolate the unknown. Include every ion with its stoichiometric coefficient.

Check units and direction after rearrangement. A quick magnitude check should match the expected strength of ion–water or lattice attraction.

If ΔHsol, lattice dissociation and one hydration term are known, the missing hydration term is the residual after subtracting the other route contributions.

Do not average the hydration values or change a sign because the numerical answer is negative; the arrow direction determines the sign.

Hydration enthalpy measures attraction between an ion and polar water molecules. Higher charge density gives stronger attraction and a more negative hydration enthalpy.

Compare charge first, then radius among ions of the same charge. Smaller ions place charge closer to water dipoles, but hydration also depends on the ion’s coordination environment.

Mg²⁺ has a more negative hydration enthalpy than Na⁺ because of its higher charge density; Na⁺ is more strongly hydrated than K⁺ because it is smaller.

Do not claim that every larger ion has a more negative value; for a fixed charge, the trend is usually toward less exothermic hydration.