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
Thermodynamic Cycle
Enthalpy of Solution Equation
\( \Delta H_{\text{sol}}^\ominus = \Delta H_{\text{latt}}^\ominus + \sum \Delta H_{\text{hyd}}^\ominus \)
Where \(\Delta H_{\text{latt}}^\ominus\) is defined endothermically (lattice dissociation).
Dissolution Mechanism
The Two Steps of Ionic Dissolution
- Lattice breaking (endothermic): Separate the ionic lattice into gaseous ions (+ΔHlatt).
- Hydration (exothermic): Surround gaseous ions with polar water molecules (−ΣΔHhyd).
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).
Periodic Trend
Hydration Enthalpy and Ionic Charge Density
Smaller ionic radius and higher ionic charge (→ higher charge density) produce stronger ion-dipole attractions with water molecules, making ΔHhyd more exothermic (e.g. Mg²⁺ has a much more negative ΔHhyd than Na⁺ or K⁺).
AQA GCSE & IB Chemistry
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