Redox Processes: Oxidation Numbers and Electron Transfer

This guide covers subtopic 6.1, Redox processes: electron transfer and changes in oxidation number, from Topic 6, Electrochemistry, of Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. This is AS Level content.

Before studying this

At IGCSE or O Level, you met oxidation and reduction as oxygen or electron transfer, learned the basic oxidation-number rules, and used Roman-numeral naming for compounds like iron(III) oxide. Redox Reactions covers the Core and Supplement treatment of this in full.

AS Level treats oxidation number as a complete, systematic tool: assigning it confidently to any element in any compound or ion, using changes in oxidation number to balance whole equations, and using it to define disproportionation precisely. Nothing here is Core/Extended-tiered — 9701 does not use that distinction — but you’re expected to apply the rules more fluently and to a wider range of species than at IGCSE.

Syllabus coverage

CAMBRIDGE INTERNATIONAL AS & A LEVEL CHEMISTRY 9701 — AS Level, Topic 6

6.1 Redox processes: electron transfer and changes in oxidation number (oxidation state) — calculating oxidation numbers of elements in compounds and ions; using changes in oxidation numbers to help balance chemical equations; explaining and using the terms redox, oxidation, reduction and disproportionation in terms of electron transfer and changes in oxidation number; explaining and using the terms oxidising agent and reducing agent; using a Roman numeral to indicate the magnitude of the oxidation number of an element.

Assigning oxidation numbers

An oxidation number is assigned using a small set of rules, applied in order until the value is fixed:

• An atom in its elemental form has oxidation number 0 (e.g. Cl in Cl₂, Fe in Fe(s)).
• A simple (monatomic) ion has an oxidation number equal to its charge (e.g. Na⁺ is +1, O²⁻ is −2).
• Combined oxygen is −2, except in peroxides (−1) and when bonded to fluorine.
• Combined hydrogen is +1, except in metal hydrides (−1).
• The oxidation numbers in a neutral compound sum to 0; in a polyatomic ion, they sum to the ion’s overall charge.

Worked example. Find the oxidation number of manganese in MnO₄⁻.

Oxygen contributes −2 each, four oxygens = −8. The ion’s overall charge is −1, so Mn’s oxidation number + (−8) = −1, giving Mn = +7. This is why the manganate(VII) ion is named with the Roman numeral (VII) — it states the oxidation number of the manganese directly.

Oxidation, reduction and the vocabulary of redox

In terms of oxidation number: oxidation is an increase in oxidation number; reduction is a decrease. In terms of electrons — the equivalent, older definition — oxidation is loss of electrons, reduction is gain of electrons (“OIL RIG”: oxidation is loss, reduction is gain). The two descriptions always agree: losing an electron makes an atom’s charge (and therefore oxidation number) more positive.

An oxidising agent is the species that oxidises another substance — in doing so, it is itself reduced (it gains the electrons the other substance loses). A reducing agent is the species that reduces another substance, and is itself oxidised.

Disproportionation is a single species being simultaneously oxidised and reduced in the same reaction. A clear example is chlorine reacting with cold dilute sodium hydroxide:

Cl₂ + 2NaOH → NaCl + NaOCl + H₂O

Chlorine starts at oxidation number 0. In NaCl, it ends at −1 (reduced); in NaOCl, it ends at +1 (oxidised) — the same element, from the same starting reagent, going in both directions in the same reaction.

Balancing equations by oxidation-number change

Because electrons lost must equal electrons gained, you can balance a redox equation by tracking oxidation-number changes rather than balancing atoms by inspection alone.

Worked example. Balance: MnO₄⁻ + Fe²⁺ + H⁺ → Mn²⁺ + Fe³⁺ + H₂O

Step 1 — find the oxidation-number changes. Manganese goes from +7 (in MnO₄⁻) to +2 (in Mn²⁺): a decrease of 5, so each Mn gains 5 electrons. Iron goes from +2 to +3: an increase of 1, so each Fe loses 1 electron.

Step 2 — equalise electrons lost and gained. One Mn gains 5 electrons, so 5 Fe²⁺ are needed to supply them:

MnO₄⁻ + 5Fe²⁺ + H⁺ → Mn²⁺ + 5Fe³⁺ + H₂O

Step 3 — balance oxygen and hydrogen. The 4 oxygens in MnO₄⁻ need 4 H₂O; those 4 H₂O need 8 H⁺:

MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O

Step 4 — check. Charge: left = (−1) + 5(+2) + 8(+1) = +17; right = (+2) + 5(+3) + 0 = +17. Atoms: 1 Mn, 5 Fe, 4 O, 8 H on each side. Balanced.

Common mistakes

• Assigning oxygen −2 without checking for a peroxide. In H₂O₂ or Na₂O₂, oxygen is −1, not −2 — always check whether the rule’s exception applies before assuming the default.
• Confusing “oxidised” and “oxidising agent.” The oxidising agent is the species that causes oxidation in something else, and is itself reduced — it is easy to say the reverse by mistake.
• Missing disproportionation because both products “look like normal salts.” Check the oxidation number of the element in question in every product, not just the most obvious one, before concluding a reaction is a simple redox reaction rather than a disproportionation.
• Balancing electrons but forgetting to balance charge overall. Equal electrons lost and gained is necessary but not sufficient — always do a final charge and atom count on the whole equation.

Quick revision checklist

• The oxidation-number rules, applied in order, including the oxygen and hydrogen exceptions
• Oxidation = increase in oxidation number / electron loss; reduction = decrease / electron gain
• Oxidising agent (itself reduced) vs reducing agent (itself oxidised)
• Disproportionation: one species, one reaction, both oxidised and reduced
• Balancing redox equations by equalising electrons lost and gained via oxidation-number change, then balancing O and H

Related resources

• Redox Reactions — the IGCSE/O Level oxidation-number rules and Core/Supplement treatment this resource builds on
• Chemical Equilibria: Kc, Kp and Le Chatelier’s Principle — the Contact process, itself a redox reaction run reversibly
• Electrochemistry: Electrolysis and Standard Electrode Potentials — the A Level topic that quantifies redox as a voltage
• Cambridge AS & A Level Chemistry hub

Written against Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. Always check the current syllabus for your examination year.