An ionic equation is a balanced equation in which the substances are expressed as dissociated ions in aqueous solution. Always make sure the type of elements and the number of each type of elements are balanced on both sides of the equation. Equally important, but quite often left unchecked by students, is to make sure the net charge of the ions is balanced on both sides of the equation.

You can start off an ionic equation by first writing down the full chemical equation, or if it is a simple and common reaction, you can write down the ionic equation directly. (For simplicity, we assume an ionic equation means a net ionic equation, with the spectator ions omitted)

There are only 3 general rules to follow:

– If the substance is an ionic compound in the aqueous state, split it into its ions.

– If the substance is a strong alkali or a strong acid, split it into its ions.

– If it does not fall in categories 1 or 2, do not change anything. Leave it as it is.

Examples:

Precipitation Reaction:

AgNO₃ (aq) + NaCl (aq) –> AgCl (s) + NaNO₃ (aq)

Splitting into ions: Ag⁺ + NO₃^– + Na⁺ + Cl^– –> AgCl + Na⁺ + NO₃^–

Cancel away the spectator ions: Ag⁺ + NO₃^– + Na⁺ + Cl^– –> AgCl + Na⁺ + NO₃^–

Final ionic equation: Ag⁺ (aq) + Cl^– (aq) –> AgCl (s)

Displacement Reaction:

Ca (s) + Cu(NO₃)₂ (aq) –> Ca(NO₃)₂ (aq) + Cu (s)

Splitting into ions: Ca + Cu²⁺ + 2NO₃^– –> Ca²⁺ + 2NO₃^– + Cu

Cancel away the spectator ions: Ca + Cu²⁺ + 2NO₃^– –> Ca²⁺ + 2NO₃^– + Cu

Final ionic equation: Ca (s) + Cu²⁺ (aq) –> Ca²⁺ (aq) + Cu (s)

Neutralisation Reaction:

2NaOH (aq) + H₂SO₄ (aq) –> Na₂SO₄ (aq) + 2H₂O (l)

Splitting into ions: 2Na⁺ + 2OH^– + 2H⁺ + SO₄^2- –> 2Na⁺ + SO₄^2- + 2H₂O

Cancel away the spectator ions: 2Na⁺ + 2OH^– + 2H⁺ + SO₄^2- –> 2Na⁺ + SO₄^2- + 2H₂O

Final ionic equation, in simplest ratio: OH^– (aq) + H⁺ (aq) –> H₂O (l)