Enthalpies of solution and hydration — common mistakes

Examiners love to test the sign of the lattice enthalpy in this cycle. Standard lattice enthalpy ($ΔH_{\text{latt}}^{\ominus}$) is defined for the formation of 1 mole of the lattice from gaseous ions, which is exothermic. The cycle for solution involves breaking the lattice, which is the reverse process and therefore endothermic. Always use the value $-ΔH_{\text{latt}}^{\ominus}$ in your calculation.

When comparing solubilities, always link the calculated $\Delta H_{\text{sol}}^{\ominus}$ values back to the physical reality. A large positive value is a significant energy barrier to dissolving, making the substance less soluble. A negative or small positive value means the energy barrier is small or non-existent, favouring dissolution.

Both processes involve the formation of attractive forces, which releases energy and is therefore exothermic. Lattice enthalpy involves forming strong electrostatic attractions between oppositely charged gaseous ions to make a solid lattice. Hydration enthalpy involves forming ion-dipole attractions between gaseous ions and polar water molecules.

Yes. The spontaneity of a process is determined by the Gibbs free energy change, $\Delta G = \Delta H - T\Delta S$. When an ionic solid dissolves, there is a large increase in disorder, so the entropy change ($\Delta S$) is large and positive. If this positive entropy term is large enough, it can make $\Delta G$ negative even if $\Delta H$ is positive (endothermic), allowing the salt to dissolve spontaneously.

The Born-Haber cycle is used to determine the lattice enthalpy ($ΔH_{\text{latt}}^{\ominus}$) from other experimental values (like ionisation energy and electron affinity). The enthalpy cycle for solution uses the lattice enthalpy value, which may have been found via a Born-Haber cycle, to calculate the enthalpy of solution. They are two different, but related, applications of Hess's Law.