The rules for oxidation numbers

The oxidation number is the charge an atom would have if every bond in the species were fully ionic. It is a bookkeeping device — but it is the key to every redox question.

• Any element in its uncombined state = 0 (Cl₂, Na, O₂, S₈).
• A simple ion = the charge on the ion (Mg²⁺ = +2; Cl⁻ = −1).
• Hydrogen = +1 (but −1 in metal hydrides such as NaH).
• Oxygen = −2 (but −1 in peroxides such as H₂O₂, and +2 in OF₂).
• Fluorine = −1 always (it is the most electronegative element).
• The oxidation numbers in a neutral compound add up to 0; in an ion they add up to the charge.

Your turn — sulfur in sulfuric acid

Your turn — chromium in dichromate(VI)

Your turn — nitrogen in the ammonium ion

Oxidation, reduction and the agents

Two equivalent definitions — use whichever is easier:

• Electrons: Oxidation Is Loss, Reduction Is Gain.
• Oxidation number: oxidation = an increase (more positive); reduction = a decrease.

An oxidising agent oxidises something else, so it gains electrons and is itself reduced (e.g. MnO₄⁻, Cl₂, H₂O₂). A reducing agent loses electrons and is itself oxidised (e.g. Fe²⁺, I⁻, Zn, CO).

Naming the agent

Building ionic half-equations

A half-equation shows the electrons explicitly. Build it in this order:

• 1. Balance the atoms of the element being oxidised/reduced.
• 2. Balance oxygen by adding H₂O.
• 3. Balance hydrogen by adding H⁺.
• 4. Balance the charge by adding electrons to the more positive side.

Combining half-equations: scale each one so the electrons cancel, then add. Never leave electrons in the final full ionic equation.

Your turn — counting the electrons

Is it balanced?

Disproportionation

Disproportionation is a reaction in which the same element in a single species is simultaneously oxidised and reduced.

• Cold dilute NaOH gives chlorate(I), ClO⁻ (bleach: NaClO). Hot concentrated NaOH gives chlorate(V): 3Cl₂ + 6NaOH → 5NaCl + NaClO₃ + 3H₂O (Cl: 0 → −1 and 0 → +5).
• Chlorine with water: Cl₂ + H₂O ⇌ HCl + HClO — again 0 → −1 and 0 → +1. HClO kills bacteria, which is why chlorine is used to treat drinking water. The benefit (killing pathogens such as cholera) is judged to outweigh the toxicity risk of chlorine itself.
• Copper(I) disproportionates in acid: 2Cu⁺ → Cu²⁺ + Cu (+1 → +2 and +1 → 0).

Your turn — chlorine in bleach

Spotting disproportionation

Oxidised, reduced, or neither?

For each change, tap the card, then tap what has happened to that element.

From oxidation numbers to formulae — and back

Oxidation numbers let you write formulae and name compounds you have never met:

• Vanadium(V) oxide: V is +5, O is −2 → two V (+10) balance five O (−10) → V₂O₅.
• Iron(III) sulfate: Fe³⁺ with SO₄²⁻ → Fe₂(SO₄)₃.
• Sodium chlorate(V): Cl is +5 in ClO₃⁻ → NaClO₃.
• Potassium manganate(VII): Mn is +7 in MnO₄⁻ → KMnO₄.

Metals in general form positive ions by losing electrons (they are reducing agents); non-metals in general form negative ions by gaining electrons (they are oxidising agents). That is why the strongest oxidising agents (F₂, Cl₂) sit at the top right of the Periodic Table and the strongest reducing agents (Group 1 metals) at the bottom left.

Writing a formula

Half-equation for a reducing agent

Match the species to the oxidation number of the highlighted element

Tap an item on the left, then its partner on the right.

The big ideas to know

You've now covered Topic 3: Redox I of the Edexcel A-level Chemistry (9CH0) specification. Press Finish to see your score.