Short on time? This page is the recap. For explanations and worked examples, use the full study guide.
Recap
- In a redox reaction, electrons pass from the species that is oxidized (loses e⁻, oxidation number rises) to the species that is reduced (gains e⁻, oxidation number falls).
- A half-reaction shows just one of these processes, with the electrons written in.
- A balanced redox equation is built by balancing each half-reaction, making the electrons equal, and adding.
- The coefficients of the balanced equation are mole ratios for stoichiometry and redox titrations.
Key relationships
| Idea | What it looks like | Use it to |
|---|---|---|
| Oxidation half-reaction | Zn → Zn²⁺ + 2e⁻ (e⁻ on the right) | show electrons lost |
| Reduction half-reaction | Cu²⁺ + 2e⁻ → Cu (e⁻ on the left) | show electrons gained |
| Balancing order (acid) | atoms other than O and H → H₂O for O → H⁺ for H → e⁻ for charge | balance any half-reaction in acid |
| Electron count | electrons = change in oxidation number × number of atoms that change | check step 5 |
| Combining | multiply to the lowest common multiple of the electrons, then add | cancel the electrons |
| Basic solution (extension) | balance as in acid, then add OH⁻ to both sides for each H⁺ | turn H⁺ into H₂O |
Assumptions behind the method
- The reaction is in aqueous solution, so H₂O is always available, plus H⁺ (acidic) or OH⁻ (basic).
- You know (or are told) the main reactants and products; the method balances them but does not predict them.
- Spectator ions are left out, so the result is a net ionic equation.
Mistakes to avoid
- Checking atoms only. Always check total charge on each side too.
- Electrons left over. If e⁻ is in the final equation, the half-reactions were not multiplied correctly.
- Electrons on the wrong side. Confirm with the oxidation-number change.
- Adding electron numbers. A 3e⁻ and a 2e⁻ half-reaction combine with 6e⁻ transferred, not 5.
- Not cancelling H⁺ and H₂O that appear on both sides after adding.
- Charges “made to equal zero”. The two sides must have equal charge, which need not be zero.
- Using O₂ or O²⁻ to balance oxygen. Use H₂O.
Quick self-check
- Balance: Fe(s) + Ag⁺(aq) → Fe²⁺(aq) + Ag(s). (Fe + 2Ag⁺ → Fe²⁺ + 2Ag)
- Write the oxidation half-reaction for magnesium metal forming Mg²⁺. (Mg → Mg²⁺ + 2e⁻)
- Balance: Mg(s) + Au³⁺(aq) → Mg²⁺(aq) + Au(s). How many electrons are transferred? (3Mg + 2Au³⁺ → 3Mg²⁺ + 2Au; 6 electrons)
- Balance the half-reaction NO₃⁻ → NO in acidic solution. (NO₃⁻ + 4H⁺ + 3e⁻ → NO + 2H₂O)
Next: practice questions.