When an ionic compound dissolves in water, two processes occur: the lattice must be broken (endothermic) and the ions must be hydrated (exothermic). The overall enthalpy of solution depends on which process releases more energy.
Lattice Dissociation
NaCl(s) → Na⁺(g) + Cl⁻(g)
Endothermic (+)
Hydration
Na⁺(g) → Na⁺(aq)
Cl⁻(g) → Cl⁻(aq)
Exothermic (−)
Solution (Overall)
NaCl(s) → Na⁺(aq) + Cl⁻(aq)
Can be + or −
Key Definitions
Enthalpy of solution (ΔHsol) — enthalpy change when 1 mole of a solute dissolves in enough water to form an infinitely dilute solution.
Enthalpy of hydration (ΔHhyd) — enthalpy change when 1 mole of gaseous ions is surrounded by water molecules. Always exothermic.
Dissolution Energy Cycle
\( \Delta H_{sol} = -\Delta H_{lat} + \sum \Delta H_{hyd} \)
Factors Affecting Hydration Enthalpy
- Smaller ions → more exothermic ΔHhyd (stronger attraction to water)
- Higher charge → more exothermic ΔHhyd
- This is why Mg²⁺ has a much more exothermic hydration enthalpy than Na⁺
Think About It
NaCl dissolves endothermically (ΔHsol = +3.9 kJ mol⁻¹). If it absorbs energy, why does it still dissolve?
Because dissolving also increases entropy (disorder). The ions spread out in solution, increasing the number of possible arrangements. Even though ΔH is slightly positive, the entropy increase makes the overall process spontaneous (ΔG < 0).