Solubility of Solids in Liquids

Solubility is the maximum amount of solute that can dissolve in a given quantity of solvent at a specified temperature, when excess solute is in contact with the solution.

Three concentration states:

• Unsaturated solution — less solute than the maximum; more can dissolve
• Saturated solution — exactly at maximum; at dynamic equilibrium with undissolved solute
• Supersaturated solution — contains more dissolved solute than the equilibrium value; unstable — crystallisation triggered by a seed crystal, scratch, or vibration

In a saturated solution, a dynamic equilibrium exists:
$\text{solute}_{\text{undissolved}} + \text{energy} \rightleftharpoons \text{solute}_{\text{dissolved}}$

Dissolving a solid requires two competing energy steps:

Step 1 — Lattice energy (endothermic, always positive):
The ionic lattice must be broken — ions separated from each other. This always requires energy.

Step 2 — Hydration enthalpy (exothermic, always negative):
Water molecules surround and stabilise the separated ions. This releases energy.

$\Delta H_{\text{soln}} = \Delta H_{\text{lattice}} + \Delta H_{\text{hydration}}$

The sign of $\Delta H_{\text{soln}}$ determines whether the solution cools or warms:

• $\Delta H_{\text{soln}} > 0$ (endothermic) → solution cools: $\ce{KI}$, $\ce{KNO3}$, $\ce{NH4NO3}$
• $\Delta H_{\text{soln}} < 0$ (exothermic) → solution warms: $\ce{NaOH}$, $\ce{H2SO4}$ in water, $\ce{CaCl2}$

For solid solutes: Most dissolve more at higher temperatures (endothermic dissolution). A few exceptions — $\ce{Ce2(SO4)3}$, $\ce{Li2SO4}$ — show decreased solubility with temperature.

Practical consequence: Dissolve in hot solvent → cool slowly → crystallisation occurs as solubility falls. This is the principle of recrystallisation purification.

For gases in liquids: Solubility always decreases with temperature. Warm water holds less dissolved oxygen — this is why fish crowd in cooler deep water in summer.