Oxidation and reduction always happen together in the same reaction — you cannot have one without the other. The mnemonic OIL RIG (Oxidation Is Loss, Reduction Is Gain — of electrons) is the foundation: when one substance loses electrons, those electrons must go somewhere, so another substance gains them simultaneously.
What are oxidation and reduction?
At GCSE, oxidation and reduction are defined in terms of electrons:
- Oxidation: loss of electrons (OIL)
- Reduction: gain of electrons (RIG)
A reaction in which oxidation and reduction occur simultaneously is called a redox reaction. Because electrons cannot appear or disappear — they are transferred from one substance to another — every oxidation event is accompanied by an equal and opposite reduction event.
Earlier definitions (still valid and tested at GCSE):
| Earlier definition | Still correct? | Electron-transfer equivalent |
|---|---|---|
| Oxidation = gain of oxygen | Yes | Gaining O²⁻ ions means the other element loses electrons to oxygen |
| Reduction = loss of oxygen | Yes | O²⁻ ions are removed, returning electrons to the reduced element |
| Oxidation = loss of hydrogen | Yes (in organic/acid-base contexts) | Hydrogen carries electrons away |
| Reduction = gain of hydrogen | Yes | Hydrogen donates electrons |
The electron-transfer definition is more general and more precise — it applies to all redox reactions, including those where no oxygen is involved.
What are oxidising and reducing agents?
- Oxidising agent: causes oxidation in another substance by accepting electrons from it. The oxidising agent is itself reduced (gains electrons).
- Reducing agent: causes reduction in another substance by donating electrons to it. The reducing agent is itself oxidised (loses electrons).
A key exam point: the oxidising agent and reducing agent always refer to the substance doing the causing — and they always undergo the opposite change themselves.
| In this reaction | What is oxidised | What is reduced | Oxidising agent | Reducing agent |
|---|---|---|---|---|
| Mg + CuSO₄ → MgSO₄ + Cu | Mg (loses 2e⁻) | Cu²⁺ (gains 2e⁻) | CuSO₄ | Mg |
| Fe₂O₃ + 3CO → 2Fe + 3CO₂ | CO (gains O) | Fe₂O₃ (loses O) | Fe₂O₃ | CO |
| 2Cl₂ + 2H₂O → 4HCl + O₂ | H₂O (loses e⁻) | Cl₂ (gains e⁻) | Cl₂ | H₂O |
How do you identify oxidation and reduction in a reaction?
Worked example 1 — metal displacement:
Zinc is added to copper sulfate solution: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
Step 1 — write as the net ionic equation: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
Step 2 — identify electron transfer:
- Zn → Zn²⁺: zinc loses 2 electrons → oxidised
- Cu²⁺ → Cu: copper ion gains 2 electrons → reduced
Step 3 — identify agents:
- Zn is the reducing agent (donates electrons)
- Cu²⁺ (CuSO₄) is the oxidising agent (accepts electrons)
Worked example 2 — blast furnace:
In the blast furnace, iron oxide is reduced by carbon monoxide: Fe₂O₃ + 3CO → 2Fe + 3CO₂
- Fe₂O₃ loses oxygen → iron is reduced (Fe³⁺ gains electrons → Fe⁰)
- CO gains oxygen → carbon is oxidised (C in CO becomes CO₂; C gains oxygen = oxidised)
- Fe₂O₃ is the oxidising agent; CO is the reducing agent
How do you write half-equations for redox reactions?
A half-equation shows one half of a redox reaction — either the oxidation or the reduction — including the electrons transferred:
Example: zinc and copper(II) ions
Oxidation half-equation: Zn → Zn²⁺ + 2e⁻
Reduction half-equation: Cu²⁺ + 2e⁻ → Cu
Add the two half-equations together (the 2e⁻ cancel): Zn + Cu²⁺ → Zn²⁺ + Cu
Rules for writing half-equations:
- Balance atoms of the element being oxidised/reduced.
- Add electrons (e⁻) to balance the charge — electrons go on the right for oxidation, on the left for reduction.
- If needed, add H₂O to balance oxygen atoms and H⁺ ions to balance hydrogen atoms (higher-tier questions only).
Common oxidising and reducing agents to know
| Substance | Typical role | Example |
|---|---|---|
| Oxygen (O₂) | Oxidising agent | Burning metals: Mg + O₂ → MgO |
| Chlorine (Cl₂) | Oxidising agent | Displacement: Cl₂ + 2KBr → 2KCl + Br₂ |
| Carbon (C) | Reducing agent | Blast furnace: FeO + C → Fe + CO |
| Carbon monoxide (CO) | Reducing agent | Blast furnace: Fe₂O₃ + 3CO → 2Fe + 3CO₂ |
| Hydrogen (H₂) | Reducing agent | Metal oxide reduction: CuO + H₂ → Cu + H₂O |
| Metals (Na, Mg, Zn…) | Reducing agents | Displacement: Mg + CuSO₄ → MgSO₄ + Cu |
Frequently asked questions
Why is the oxidising agent the substance that is reduced?
The oxidising agent is the substance that causes oxidation in something else — it does this by accepting the electrons that the other substance loses. Because it accepts electrons, the oxidising agent itself gains electrons, and gaining electrons is reduction. So the oxidising agent is always reduced. This sounds paradoxical at first, but it follows directly from the definition: the agent causes something to happen to its partner, while undergoing the opposite itself.
How can you tell if a reaction is a redox reaction just by looking at it?
Look for changes in the charge (oxidation state) of the elements involved. At GCSE level, a practical shortcut: any reaction involving a pure element as a reactant or product is almost certainly redox, because when an element is formed from a compound (or vice versa), electron transfer must have occurred. Neutralisation (acid + alkali → salt + water) and dissolving are generally not redox reactions because no element changes charge state.
What is the connection between redox reactions and the reactivity series?
The reactivity series ranks metals by their tendency to lose electrons (be oxidised). A more reactive metal loses electrons more readily, making it a better reducing agent. When a reactive metal is added to a solution containing ions of a less reactive metal, the more reactive metal is oxidised and the less reactive metal ion is reduced — this is displacement. Zinc displaces copper because zinc is a better reducing agent (more reactive); copper does not displace zinc ions because copper is a poorer reducing agent.
Is combustion a redox reaction?
Yes. In combustion, the fuel (carbon or hydrogen compounds) is oxidised — it loses electrons to oxygen. Oxygen is the oxidising agent and is itself reduced. The word "oxidation" originally meant "reaction with oxygen", and combustion is the classic example. In the modern electron-transfer definition, combustion is redox because carbon and hydrogen atoms in the fuel go from a neutral state (in organic molecules) to a positive state in CO₂ and H₂O — they have lost electron density to the more electronegative oxygen atoms.
For GCSE chemistry that builds every reaction from particle-level electron transfers — explore Professor Curie at aitutors.me.