Solubility of Solids in Liquids
Saturated solutions, dissolution equilibrium, and heats of solution
When you dissolve KNO₃ in water, the solution feels noticeably cold. When you dissolve NaOH in water, the solution becomes very hot. Using only this thermal observation, predict which salt will become more soluble as temperature increases — and explain your reasoning without any memorisation.
Instant cold packs used in sports medicine contain ammonium nitrate crystals and a sealed water bag. When you snap the pack, the inner bag breaks, water contacts , and the dissolving reaction absorbs so much heat that the pack temperature drops to ~2°C — cold enough to reduce swelling. Instant heat packs work in reverse: calcium chloride dissolving is exothermic, releasing enough heat to warm muscles. Two different dissolution reactions, two opposite thermal effects — both applications of heats of solution.
Solubility and Saturation
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:
Heats of Solution — Why Dissolution is a Two-Step Process
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.
The sign of determines whether the solution cools or warms:
- (endothermic) → solution cools: , ,
- (exothermic) → solution warms: , in water,
AI Generation Prompt
Enthalpy level diagram for KI dissolving in water. Show three energy levels as horizontal lines: bottom level labelled "K⁺(aq) + I⁻(aq) dissolved" (lowest), middle level labelled "KI(s) + H₂O" (starting material), top level labelled "K⁺(g) + I⁻(g) separated ions" (highest). An upward arrow from KI(s) to separated ions labelled "+632 kJ/mol (Lattice energy, endothermic)" in orange. A downward arrow from separated ions to dissolved ions labelled "−619 kJ/mol (Hydration enthalpy, exothermic)" in blue. A small upward net arrow from KI(s) to dissolved ions labelled "ΔH_soln = +13 kJ/mol (net endothermic)" in amber. Dark background, orange accent labels, clean technical illustration style.
Effect of Temperature on Solubility
For solid solutes: Most dissolve more at higher temperatures (endothermic dissolution). A few exceptions — , — 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.
Q1.In a saturated solution of KNO₃ at 25°C, the rate of dissolving and the rate of crystallisation are:
When you dissolve KNO₃ in water, the solution feels noticeably cold. When you dissolve NaOH in water, the solution becomes very hot. Using only this thermal observation, predict which salt will become more soluble as temperature increases — and explain your reasoning without any memorisation.
Instant cold packs used in sports medicine contain ammonium nitrate crystals and a sealed water bag. When you snap the pack, the inner bag breaks, water contacts , and the dissolving reaction absorbs so much heat that the pack temperature drops to ~2°C — cold enough to reduce swelling. Instant heat packs work in reverse: calcium chloride dissolving is exothermic, releasing enough heat to warm muscles. Two different dissolution reactions, two opposite thermal effects — both applications of heats of solution.
Solubility and Saturation
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:
Heats of Solution — Why Dissolution is a Two-Step Process
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.
The sign of determines whether the solution cools or warms:
- (endothermic) → solution cools: , ,
- (exothermic) → solution warms: , in water,
AI Generation Prompt
Enthalpy level diagram for KI dissolving in water. Show three energy levels as horizontal lines: bottom level labelled "K⁺(aq) + I⁻(aq) dissolved" (lowest), middle level labelled "KI(s) + H₂O" (starting material), top level labelled "K⁺(g) + I⁻(g) separated ions" (highest). An upward arrow from KI(s) to separated ions labelled "+632 kJ/mol (Lattice energy, endothermic)" in orange. A downward arrow from separated ions to dissolved ions labelled "−619 kJ/mol (Hydration enthalpy, exothermic)" in blue. A small upward net arrow from KI(s) to dissolved ions labelled "ΔH_soln = +13 kJ/mol (net endothermic)" in amber. Dark background, orange accent labels, clean technical illustration style.
Effect of Temperature on Solubility
For solid solutes: Most dissolve more at higher temperatures (endothermic dissolution). A few exceptions — , — 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.
Q1.In a saturated solution of KNO₃ at 25°C, the rate of dissolving and the rate of crystallisation are: