1. Learning Objective
By the end of this topic you will be able to:
- State the Cambridge IGCSE core definition of oxidation and reduction (loss or gain of electrons) and recognise the historic oxygen‑centric view as supplementary.
- Assign oxidation numbers, identify oxidation‑number changes and therefore determine which species are oxidised or reduced.
- Identify the oxidising and reducing agents, using the syllabus wording (oxidising agent = species that gains electrons; reducing agent = species that loses electrons).
- Write balanced redox equations and the corresponding oxidation and reduction half‑reactions using the standard 4‑step method (required for the Supplement).
- Apply colour‑change tests to confirm redox activity.
- Explain two industrial redox processes that appear in the syllabus (Contact process and metal extraction).
2. Definitions
- Core (syllabus) definition – electron‑transfer view
- Oxidation: loss of electrons; the oxidation number of the element increases.
- Reduction: gain of electrons; the oxidation number of the element decreases.
- This definition is required for all IGCSE questions (AO1).
- Historic (oxygen‑centric) definition – supplementary
- Oxidation: a substance gains oxygen atoms (or loses hydrogen atoms).
- Reduction: a substance loses oxygen atoms (or gains hydrogen atoms).
- Useful only when oxygen is actually transferred; otherwise rely on the electron‑transfer definition.
3. Oxidation Numbers – Quick Reference
Remember that in the core syllabus the oxidation numbers of hydrogen and oxygen are fixed in most compounds (H +1, O –2). This helps avoid common mistakes.
| Element | Common oxidation number(s) |
|---|
| H | +1 (except in metal hydrides, –1) |
| O | –2 (except in peroxides, –1; in OF₂, +2) |
| Al, Na, K, Mg, Ca, Sr, Ba | +3, +1, +1, +2, +2, +2, +2 respectively |
| Fe | +2, +3 (depends on compound) |
| C | –4 (CH₄), –2 (CO), +2 (CO₂), 0 (elemental C) |
| Cl | –1 (most compounds), +5 (ClO₃⁻), +7 (ClO₄⁻) |
| Mn | +2, +4, +6, +7 (e.g., MnO₂, KMnO₄) |
4. Oxidising and Reducing Agents
| Agent | Role (as defined by the syllabus) | Example reaction |
|---|
| Potassium permanganate, KMnO₄ |
Oxidising agent – it is reduced (its oxidation number decreases) |
KMnO₄ + 5 H₂C₂O₄ + 6 H⁺ → K⁺ + Mn²⁺ + 10 CO₂ + 8 H₂O |
| Hydrogen sulphide, H₂S |
Reducing agent – it is oxidised (its oxidation number increases) |
H₂S + Cl₂ → 2 HCl + S |
| Carbon monoxide, CO |
Reducing agent |
Fe₂O₃ + 3 CO → 2 Fe + 3 CO₂ |
| Oxygen, O₂ |
Oxidising agent |
2 Mg + O₂ → 2 MgO |
5. Writing Redox Half‑Reactions (4‑Step Method – required for the Supplement)
- Separate the overall equation into two halves: one for oxidation, one for reduction.
- Balance all atoms except O and H.
- Balance oxygen atoms by adding H₂O molecules.
- Balance hydrogen atoms by adding H⁺ ions (in acidic solution) or OH⁻ ions (in basic solution).
- Balance the charge by adding electrons (e⁻) to the more positive side.
- Multiply each half‑reaction by the smallest integer that makes the number of electrons equal, then add the halves and cancel species that appear on both sides.
6. Representative Redox Reactions (Core syllabus)
| # | Balanced equation | Oxidation‑number changes (electron‑transfer view) | Oxidising / Reducing agents | Notes (O‑centric view, where useful) |
| 1 | 2 Mg + O₂ → 2 MgO |
Mg: 0 → +2 (oxidised) O: 0 → –2 (reduced) |
Oxidising agent: O₂ Reducing agent: Mg |
Mg “gains” O – matches O‑centric rule. |
| 2 | Zn + CuSO₄ → ZnSO₄ + Cu |
Zn: 0 → +2 (oxidised) Cu: +2 → 0 (reduced) |
Oxidising agent: CuSO₄ (Cu²⁺) Reducing agent: Zn |
O‑centric rule not helpful (no O transferred). |
| 3 | Fe₂O₃ + 3 CO → 2 Fe + 3 CO₂ |
Fe: +3 → 0 (reduced) C: +2 → +4 (oxidised) |
Oxidising agent: Fe₂O₃ Reducing agent: CO |
CO “gains” O – O‑centric rule works. |
| 4 | 2 Al + 3 Cl₂ → 2 AlCl₃ |
Al: 0 → +3 (oxidised) Cl: 0 → –1 (reduced) |
Oxidising agent: Cl₂ Reducing agent: Al |
O‑centric rule not applicable. |
| 5 | KMnO₄ + 5 H₂C₂O₄ + 6 H⁺ → K⁺ + Mn²⁺ + 10 CO₂ + 8 H₂O |
Mn: +7 → +2 (reduced) C: +3 → +4 (oxidised) |
Oxidising agent: KMnO₄ Reducing agent: H₂C₂O₄ |
MnO₄⁻ “gains” O – not a reliable shortcut. |
| 6 | 2 NaCl(l) → 2 Na(l) + Cl₂(g) (electrolysis of molten NaCl) |
Na: +1 → 0 (reduced) Cl: –1 → 0 (oxidised) |
Oxidising agent: Cl⁻ (it is oxidised) Reducing agent: Na⁺ (it is reduced) |
Classic binary halide electrolysis – appears in the syllabus. |
| 7 | Fe₂O₃ + 3 CO → 2 Fe + 3 CO₂ (metal extraction – blast furnace) |
Fe: +3 → 0 (reduced) C: +2 → +4 (oxidised) |
Oxidising agent: Fe₂O₃ Reducing agent: CO |
Half‑reactions are required for the Supplement (see Section 5). |
7. Colour‑Change Tests Frequently Used in IGCSE Exams
- Potassium permanganate, KMnO₄ – deep purple solution becomes colourless (or faint pink) when it oxidises a substance such as oxalic acid.
- Iodine solution, I₂ – brown colour disappears when reduced to colourless I⁻ (e.g., by thiosulphate).
- Blue‑copper sulphate solution – deep blue turns colourless when a reducing sugar (e.g., glucose) reduces Cu²⁺ to Cu⁺/Cu⁰ (Benedict’s test).
- Tincture of iodine – brown solution turns blue‑black on addition of starch; the colour disappears when the iodine is reduced (e.g., by a strong reducing agent).
8. Industrial Redox Processes (Core syllabus)
8.1 Contact Process (SO₂ → SO₃)
- Oxidation step: SO₂ + ½ O₂ → SO₃ (oxidation number of S: +4 → +6)
- Catalyst cycle (V₂O₅ ⇌ VO₂):
- V₂O₅ + SO₂ → VO₂ + SO₃ (V: +5 → +4, reduced)
- VO₂ + ½ O₂ → V₂O₅ (V: +4 → +5, oxidised)
- Overall, oxygen is transferred from O₂ to SO₂, and the vanadium catalyst shuttles electrons, illustrating a redox cycle.
8.2 Metal Extraction – Blast Furnace
Reaction shown in Section 6 (Fe₂O₃ + 3 CO → 2 Fe + 3 CO₂). The iron ore is reduced (oxidation number decreases) by carbon monoxide, which is oxidised.
9. Core vs. Supplement – What Must Be Covered for the IGCSE
- Core (mandatory for all candidates)
- Electron‑transfer definitions of oxidation/reduction.
- Assigning oxidation numbers (including fixed H and O values).
- Identifying oxidising and reducing agents (using the syllabus wording).
- Balancing redox equations and recognising the five core examples listed in Section 6.
- Colour‑change tests listed in Section 7.
- Industrial processes in Section 8.
- Supplement (required for the optional Supplement paper)
- Full 4‑step half‑reaction method (Section 5) – write oxidation and reduction half‑equations for any redox reaction.
- Electrolysis of molten binary halides (e.g., NaCl) – include half‑reactions.
- Detailed catalyst redox cycle for the Contact Process (V₂O₅ ⇌ VO₂).
- Collision theory for rate of reaction (linked to redox kinetics – optional box).
10. Practice Questions (Core + Supplement)
- In the reaction
2Al + 3Cl₂ → 2AlCl₃, identify the substance that is oxidised and the one that is reduced. Explain using both the O‑centric and electron‑transfer definitions.
- Balance the reaction between iron(III) oxide and carbon monoxide. State the oxidising and reducing agents and show the oxidation‑number changes.
- Write the oxidation and reduction half‑reactions for
CuO + H₂ → Cu + H₂O using the gain/loss of oxygen concept, then rewrite them in electron‑transfer form (apply the 4‑step method).
- Predict the colour change when potassium permanganate reacts with oxalic acid in acidic solution. Write the overall redox equation and indicate which species is the oxidising agent.
- For the reaction
Zn + CuSO₄ → ZnSO₄ + Cu:
- Assign oxidation numbers.
- State which element is oxidised and which is reduced.
- Name the oxidising and reducing agents.
- Electrolysis of molten sodium chloride: write the oxidation half‑reaction, the reduction half‑reaction and the overall balanced equation.
- Show the two half‑reactions that constitute the catalyst cycle in the Contact Process (V₂O₅ ⇌ VO₂).
11. Summary
- Oxidation = loss of electrons (oxidation number ↑) – also described historically as gain of oxygen or loss of hydrogen.
- Reduction = gain of electrons (oxidation number ↓) – also described historically as loss of oxygen or gain of hydrogen.
- The O‑centric rule is a useful shortcut **only when oxygen atoms are actually transferred**; otherwise rely on oxidation‑number changes.
- Oxidising agent = species that gains electrons (its oxidation number falls).
Reducing agent = species that loses electrons (its oxidation number rises).
- Use the 4‑step half‑reaction method for any redox equation – essential for the Supplement.
- Colour‑change tests give rapid visual confirmation of redox activity.
- Industrial examples (Contact process and metal extraction) illustrate the importance of redox chemistry in real‑world applications.