Coaching Centre
Dissolving ionic compounds
How water dissolves ionic solids, dissociation equations, the dynamic equilibrium in a saturated solution, the enthalpy and entropy changes of dissolving, solubility curves and the solubility rules.
A sodium chloride crystal holds itself together with some of the strongest forces in chemistry: you would need about 800 °C just to melt it. Yet drop it into a glass of cold water and it falls apart in seconds without any heating at all. Something about water must be prising those ions apart, and paying the energy bill.
How water dissolves an ionic solid
Water molecules are polar: the oxygen end carries a partial negative charge (δ−) and the hydrogen ends carry partial positive charges (δ+). At the surface of an ionic crystal:
- Water molecules are attracted to the exposed ions: δ− oxygen ends point towards cations and δ+ hydrogen ends towards anions. These are ion–dipole attractions.
- Collisions with moving water molecules, together with these attractions, pull ions out of the lattice, starting from corners and edges where ions are held least strongly.
- Each released ion is surrounded by a shell of oriented water molecules. This is called hydration (or, for any solvent, solvation). The hydrated ions spread through the solution.
The process is called dissociation: the ions were already present in the solid and are simply separated. Write it with state symbols:
$$\text{NaCl(s)} \xrightarrow{\text{H}_2\text{O}} \text{Na}^+\text{(aq)} + \text{Cl}^-\text{(aq)}$$
$$\text{Ca(NO}_3)_2\text{(s)} \xrightarrow{\text{H}_2\text{O}} \text{Ca}^{2+}\text{(aq)} + 2\text{NO}_3^-\text{(aq)}$$
$$\text{Al}_2(\text{SO}_4)_3\text{(s)} \xrightarrow{\text{H}_2\text{O}} 2\text{Al}^{3+}\text{(aq)} + 3\text{SO}_4^{2-}\text{(aq)}$$
Polyatomic ions such as nitrate and sulfate stay in one piece. Count the ions carefully: 1 mol of Al₂(SO₄)₃ gives 5 mol of ions.
Saturated solutions are dynamic equilibria
As more solid is added, the concentration of ions rises, and dissolved ions collide with the crystal surface more often and stick back on: precipitation (or crystallisation). Eventually ions leave the solid as fast as they return. The solution is saturated, and a heterogeneous dynamic equilibrium exists:
$$\text{NaCl(s)} \rightleftharpoons \text{Na}^+\text{(aq)} + \text{Cl}^-\text{(aq)}$$
How do we know it's dynamic and not just "stopped"? Two pieces of evidence:
- Radioactive tracers. Add solid lead(II) iodide containing some radioactive iodine-131 to a saturated PbI₂ solution made with ordinary iodine. The amount of solid and the concentration of the solution don't change, yet after a while the solution becomes radioactive and the solid less so. Labelled ions must have left the solid while unlabelled ions joined it.
- Crystal shape. A small, irregular crystal left in its saturated solution for weeks keeps the same mass but changes shape, growing smooth faces as ions dissolve from some places and deposit on others.
A solution holding less solute than the saturated amount is unsaturated. Carefully cooling a hot saturated solution with no dust or scratches can give a supersaturated solution, holding more than the equilibrium amount. It's unstable: drop in a seed crystal and the excess crystallises immediately. "Hot ice" hand warmers (sodium acetate) work this way.
The energy of dissolving
We can think of dissolving as two steps:
- Break the lattice into separate gaseous ions. This needs energy equal to the lattice energy: endothermic, and large for small, highly charged ions.
- Hydrate the ions. Forming ion–dipole attractions releases energy, the hydration energy: exothermic, and also large for small, highly charged ions. Some water–water hydrogen bonds are broken to make room, which is included in this figure.
$$\Delta H_{\text{solution}} = \Delta H_{\text{lattice}} + \Delta H_{\text{hydration}}$$
The two terms are both huge and nearly cancel, so ΔH(solution) is small and can have either sign. If hydration releases more energy than the lattice requires, dissolving is exothermic and the water warms up (CaCl₂, NaOH). If not, dissolving is endothermic and the water cools (NH₄NO₃, the basis of instant cold packs).
Why does an endothermic process happen spontaneously? Use ΔG = ΔH − TΔS from lesson 5B. Releasing ions from an ordered lattice into solution usually increases entropy a lot, so −TΔS outweighs a modest positive ΔH and ΔG is negative. For some salts with small, highly charged ions (Ca²⁺, Mg²⁺, Al³⁺) the entropy change is actually negative, because the water molecules become tightly organised around the ions, and these dissolve thanks to a strongly exothermic ΔH instead. When both terms are unfavourable, as for silver chloride, the salt is barely soluble.
- Dissolve 10 g of NH₄NO₃. By how much does the temperature change? What drives the dissolving?
- Compare 10 g of CaCl₂. Which salt would you put in an instant hot pack?
- Find a salt whose entropy change on dissolving is negative. What makes it dissolve anyway?
- Look at AgCl. Use the lattice and hydration bars and the ΔG° value to explain why it hardly dissolves.
The temperature change ignores heat absorbed by the container and the dissolved salt itself, so real values are a little smaller.
Solubility and temperature
The solubility of a substance is the maximum amount that dissolves in a given amount of solvent at a stated temperature, often given in g per 100 g of water. Because a saturated solution is an equilibrium, Le Chatelier's principle predicts the effect of temperature:
- If dissolving is endothermic (most salts, such as KNO₃), heating shifts the equilibrium to the right and solubility increases.
- If dissolving is exothermic (such as Ce₂(SO₄)₃, Li₂CO₃ and calcium hydroxide), solubility decreases as temperature rises.
- If ΔH(solution) is close to zero (NaCl), solubility hardly changes.
A saturated solution of KNO₃ in 100 g of water is made at 50 °C and then cooled to 20 °C. What mass of crystals forms? (Solubility: 85.5 g per 100 g at 50 °C; 31.6 g per 100 g at 20 °C.)
Mass of crystals = 85.5 − 31.6 = 53.9 g
If the solution had been made with 250 g of water, every quantity scales by 2.5: 2.5 × 53.9 = 135 g.
This is how chemists purify solids by recrystallisation: impurities present in small amounts stay below their own saturation point and remain in solution.
The solubility rules
"Soluble" generally means more than about 1 g dissolves in 100 mL of water; "insoluble" means less than about 0.1 g. In between, a compound is "slightly soluble". You can find the pattern yourself by mixing pairs of solutions in a spotting tile and recording which combinations form precipitates (see lesson 5I).
| Ions | Generally | Exceptions |
|---|---|---|
| Na⁺, K⁺ (group 1), NH₄⁺ | Soluble | None you need |
| NO₃⁻, CH₃COO⁻ (ethanoate) | Soluble | None you need (silver ethanoate is only moderately soluble) |
| Cl⁻, Br⁻, I⁻ | Soluble | Ag⁺ and Pb²⁺ (PbCl₂ and PbBr₂ are slightly soluble, more so in hot water) |
| SO₄²⁻ | Soluble | Ba²⁺, Pb²⁺, Sr²⁺; Ca²⁺ and Ag⁺ are slightly soluble |
| OH⁻ | Insoluble | Group 1, NH₄⁺ (as NH₃(aq)), Ba²⁺; Ca²⁺ and Sr²⁺ are slightly soluble |
| CO₃²⁻, PO₄³⁻ | Insoluble | Group 1 and NH₄⁺ |
| S²⁻ | Insoluble | Group 1, group 2 and NH₄⁺ |
| O²⁻ | Insoluble | Group 1 and some group 2 oxides react with water to form hydroxides |
Some useful precipitate colours: AgCl white, AgBr cream, AgI pale yellow, PbI₂ bright yellow, Cu(OH)₂ pale blue, Fe(OH)₂ green, Fe(OH)₃ red-brown, BaSO₄ and CaCO₃ white.
Test yourself
How many moles of ions are produced when 0.20 mol of Fe₂(SO₄)₃ dissolves completely?
Each formula unit gives 2 Fe³⁺ + 3 SO₄²⁻ = 5 ions.
Dissolving a salt in water causes the temperature to drop. Which statement is correct?
ΔH is positive, so a positive ΔS is needed for ΔG < 0. Solubility increases with temperature.
Which compound is insoluble in water?
Carbonates are insoluble except group 1 and ammonium.
Solid PbI₂ containing radioactive iodine is added to a saturated PbI₂ solution. After a day, radioactivity is detected in the solution, although the mass of solid is unchanged. This shows that
Ions continue to move between solid and solution at equal rates.
Label the enthalpy cycle for dissolving.
Soluble or insoluble?
When an ionic solid dissolves, polar water molecules form attractions with the ions, a process called . Breaking the lattice is . A solution that cannot dissolve any more solute at that temperature is . For KNO₃, solubility as temperature rises.
Instant cold packs contain solid ammonium nitrate and a pouch of water. Explain, with reference to the energy and entropy changes involved, why ammonium nitrate dissolves spontaneously and why the pack becomes cold.
Dissolving NH₄NO₃ involves breaking the ionic lattice, which is endothermic, and hydrating the NH₄⁺ and NO₃⁻ ions with water molecules, which is exothermic. For ammonium nitrate slightly more energy is needed to break the lattice than is released by hydration, so ΔH(solution) is positive (about +26 kJ mol⁻¹). This energy is absorbed from the surrounding water, so the temperature of the pack falls. Dissolving still happens because the ions move from an ordered lattice into solution, a large increase in entropy (ΔS ≈ +109 J K⁻¹ mol⁻¹). At room temperature TΔS (about 32 kJ mol⁻¹) is larger than ΔH, so ΔG = ΔH − TΔS is negative and the process is spontaneous.