Hydration Enthalpy (ΔhydH°) is the enthalpy change when one mole of gaseous ions dissolves in a large excess of water to form an infinitely dilute solution. It measures how strongly water molecules are attracted to and stabilize an ion.
The Li⁺ ion has the highest hydration enthalpy among Group 1 alkali metal ions because it has the smallest ionic radius and therefore the highest charge density, attracting water dipoles most strongly.
Always Negative (Exothermic): Hydration enthalpy is always negative because when water molecules (which are polar) surround an ion, they form ion-dipole attractions, releasing energy in the process.
Factors affecting magnitude:
Whether an ionic compound dissolves in water depends on the competition between:
If |Hydration Enthalpy| > Lattice Enthalpy → Compound dissolves
This explains why NaCl dissolves readily (its hydration enthalpy is large enough to overcome its lattice enthalpy) but some ionic compounds like AgCl do not dissolve (lattice enthalpy is too large).
**Hydration enthalpy** refers only to the ion–water interaction. **Enthalpy of dissolution** (or solution) includes BOTH the energy to break the crystal lattice (lattice enthalpy) AND the hydration enthalpy of all ions. Dissolution enthalpy = -(Lattice enthalpy) + Hydration enthalpy.
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