This is Higher Level (HL) content.
Key Definitions
| Term | Definition | Sign |
|---|---|---|
| Enthalpy of solution (ΔHsol) |
Enthalpy change when one mole of a solute dissolves completely in excess solvent to form an infinitely dilute solution | ± either |
| Enthalpy of hydration (ΔHhyd) |
Enthalpy change when one mole of isolated gaseous ions is completely surrounded by water molecules | Always − (exo) |
The Dissolving Energy Cycle
\( \Delta H_{sol} = \Delta H_{latt} + \sum \Delta H_{hyd} \)
Or equivalently: ΔHsol = −ΔHlatt(formation) + ΔHhyd(cation) + ΔHhyd(anion)
Two-Step Process
- Lattice breaking. Separate the ionic solid into gaseous ions (endothermic, + ΔHlatt)
- Hydration. Surround each gaseous ion with water molecules (exothermic, − ΔHhyd)
If hydration releases more energy than lattice breaking absorbs → ΔHsol is negative (exothermic dissolving, e.g. NaOH).
If lattice breaking absorbs more → ΔHsol is positive (endothermic dissolving, e.g. NH₄NO₃).
Hydration enthalpy depends on charge density
Smaller ions and higher charges → stronger ion-dipole attractions with water → more exothermic ΔHhyd. E.g. Mg²⁺ has a far more exothermic hydration enthalpy than K⁺.
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