When a soluble ionic solid dissolves, water molecules surround separated ions. Oxygen’s partial negative end points towards a cation; hydrogen’s partial positive ends point towards an anion. These ion–dipole attractions stabilise hydrated ions. A dissolution equation preserves atom counts and charge: CaCl2(s)→Ca2+(aq)+2Cl−(aq). Aqueous means dissolved in water, not simply liquid.
A precipitation reaction forms an insoluble solid when appropriate aqueous ions meet. Use supplied solubility rules to predict the solid, then write a net ionic equation containing only particles that change. Spectator ions remain aqueous on both sides and cancel. Silver chloride formation is Ag+(aq)+Cl−(aq)→ AgCl(s). Keep insoluble solids intact when writing ionic equations, and check total charge as well as atom conservation.
Example 1
Mix AgNO3(aq) and NaCl(aq), assuming all nitrates and sodium salts soluble and AgCl insoluble. Give the net ionic equation.
Ag+(aq)+Cl−(aq)→ AgCl(s). Na+ and NO3− remain aqueous spectators.
Example 2
Write dissolution of sodium sulfate.
Na2SO4(s)→2Na+(aq)+SO42−(aq); total charge on the right is 2−2=0.
Watch out:Solubility varies; use solubility data or stated rules to decide whether a precipitate forms.
Ionic dissolution and precipitation practice | Year 11 Chemistry | Ascent Arcade