Example Question 2

Question: Iron(III) ions react with iodide ions in aqueous solution:2 Fe³⁺ (aq) + 2 I⁻ (aq) → 2 Fe²⁺ (aq) + I₂ (aq)(a) Use oxidation numbers to state which species is oxidised and which is reduced. (b) Identify the oxidising agent and the reducing agent. (c) Write the two half-equations.(a) Iron goes from +3 in Fe³⁺ to +2 in Fe²⁺. Its oxidation number decreases, so Fe³⁺ is reduced. Iodine goes from −1 in I⁻ to 0 in I₂. Its oxidation number increases, so I⁻ is oxidised.(b) The oxidising agent is the species that is reduced: Fe³⁺. The reducing agent is the species that is oxidised: I⁻. Name the species (the ion), not just the element.(c)Fe³⁺ + e⁻ → Fe²⁺
2 I⁻ → I₂ + 2 e⁻Each Fe³⁺ gains one electron, and two iodide ions lose two electrons in total, so two Fe³⁺ ions are needed for every I₂ formed. That is why the full equation has a 2 in front of Fe³⁺.Check: the total increase in oxidation number (2 × 1 for iodine) equals the total decrease (2 × 1 for iron), and the charges balance: 2(3+) + 2(1−) = +4 on the left and 2(2+) = +4 on the right.