4 Ionic equations, spectator ions and half-reactions

Consider the redox reaction 2NaI(aq) + Cl₂(g) → 2NaCl(aq) + I₂(s).

• The overall ionic equation: 2Na⁺ + 2I⁻ + Cl₂ → 2Na⁺ + 2Cl⁻ + I₂.
• Na⁺ is a spectator ion — its oxidation number does not change and it does not take part actively. Removing it gives the net ionic equation: 2I⁻ + Cl₂ → 2Cl⁻ + I₂.
• Iodide I⁻ (−1) → I₂ (0): oxidation. Chlorine Cl₂ (0) → Cl⁻ (−1): reduction.
• Writing each separately and adding electrons to keep the charge balanced:
  • Oxidation half-reaction: 2I⁻ → I₂ + 2e⁻
  • Reduction half-reaction: Cl₂ + 2e⁻ → 2Cl⁻

  Worked example 6.6 — metal with copper sulphate. Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s). Overall ionic: Zn + Cu²⁺ + SO₄²⁻ → Zn²⁺ + SO₄²⁻ + Cu. Spectator: SO₄²⁻. Net: Zn + Cu²⁺ → Zn²⁺ + Cu. Half-reactions: Zn → Zn²⁺ + 2e⁻ (oxidation); Cu²⁺ + 2e⁻ → Cu (reduction).

  Worked example 6.7 — sodium with water. 2Na + 2H₂O → 2NaOH + H₂. Ionic: 2Na + 2H₂O → 2Na⁺ + 2OH⁻ + H₂. Oxidation: Na → Na⁺ + e⁻ (×2). Reduction: 2H₂O + 2e⁻ → H₂ + 2OH⁻. Na is the reducing agent; water (H +1) is the oxidizing agent. OH⁻ is not changed in oxidation number (it appears as product only).

  Worked example 6.8 — calcium with water. Ca + 2H₂O → Ca(OH)₂ + H₂. Net ionic: Ca + 2H₂O → Ca²⁺ + 2OH⁻ + H₂. Oxidation: Ca → Ca²⁺ + 2e⁻. Reduction: 2H₂O + 2e⁻ → H₂ + 2OH⁻.