Solubility of Group 2 compounds
Solubility (Group 2)
- When an ionic solid dissolves, two energy changes compete.
- Their balance is the enthalpy of solution.
- This explains why hydroxides and sulfates trend oppositely.
The energy balance
$$\Delta H^{\ominus}_{\text{sol}} = -\Delta H_{\text{latt}} + \Delta H_{\text{hyd}}$$
- first pull the lattice apart (costs the lattice energy), then hydrate the ions (releases hydration energy).
- If hydration roughly matches or beats the lattice energy, the solid dissolves easily.
Practice
The enthalpy of solution is given by:
You reverse the lattice energy (to break the lattice) and add the hydration enthalpy.
Practice
A solid dissolves easily when the hydration energy:
If hydration releases about as much (or more) energy than is needed to break the lattice, dissolving is favourable.
Opposite trends
Down the group both energies get smaller, but at different rates:
- hydroxides: the small $\text{OH}^-$ means the lattice energy falls a lot → solubility increases.
- sulfates: the large $\text{SO}_4^{2-}$ means the lattice energy barely changes, so the fall in hydration energy dominates → solubility decreases.
Practice
Group 2 hydroxides become MORE soluble down the group because:
With the small OH⁻ ion, the lattice energy drops sharply down the group, outweighing the hydration fall.
Practice
Group 2 sulfates become LESS soluble down the group because:
The big sulfate ion makes the lattice energy nearly constant, so the decreasing hydration energy controls the trend.
You've got it
Key idea
- $\Delta H^{\ominus}_{\text{sol}} = -\Delta H_{\text{latt}} + \Delta H_{\text{hyd}}$ (break the lattice, then hydrate)
- hydroxides: lattice energy falls a lot down the group → more soluble
- sulfates: lattice energy ≈ constant, hydration fall dominates → less soluble