Redox Reactions (Class 11): Oxidation Numbers and Balancing Made Simple
Once oxidation numbers are assigned correctly, balancing a redox equation stops being guesswork and becomes a procedure you can finish every time.
The idea in one sentence
A redox reaction is one in which electrons move from one species to another. The species losing electrons is oxidised; the one gaining them is reduced. Because electrons cannot simply vanish, the two always happen together and in exactly matching numbers — and that matching is what balancing exploits.
Rules for assigning oxidation numbers
Apply these in order. Where two rules conflict, the earlier one wins.
| # | Rule | Example |
|---|---|---|
| 1 | A free element has oxidation number 0 | Na, O2, P4, S8 all 0 |
| 2 | A monatomic ion equals its charge | Na+ is +1; Cl− is −1 |
| 3 | Fluorine is always −1 in compounds | Fluorine is the most electronegative element |
| 4 | Group 1 is +1; Group 2 is +2 in compounds | K in KMnO4 is +1 |
| 5 | Hydrogen is +1, except −1 in metal hydrides | +1 in HCl; −1 in NaH |
| 6 | Oxygen is −2, except −1 in peroxides and +2 in OF2 | −1 in H2O2 |
| 7 | The sum equals the overall charge on the species | Zero for a neutral molecule |
Worked assignment: KMnO4
K is +1 (rule 4). Each O is −2, and there are four, giving −8 (rule 6). The molecule is neutral, so by rule 7:
Manganese is therefore in the +7 state, its maximum, which is precisely why permanganate is such a strong oxidising agent: it can only go down.
Balancing by the oxidation-number method
Take the reaction of permanganate with iron(II) in acid.
- Assign oxidation numbers and find what changed. Mn goes +7 → +2, a fall of 5. Fe goes +2 → +3, a rise of 1.
- Balance the electron count. One Mn gains 5 electrons; one Fe loses 1. So 5 Fe are needed per Mn.
- Insert those coefficients, then balance the remaining atoms — oxygen with water, hydrogen with H+ in acidic medium.
- Check the charge on both sides. If the charges do not match, the balancing is wrong even when the atoms do.
Verify: left charge = (−1) + (+10) + (+8) = +17; right charge = (+2) + (+15) = +17. Balanced.
The ion-electron (half-reaction) method
Split the reaction into an oxidation half and a reduction half, balance each completely, then combine so electrons cancel.
In acidic medium
- Balance all atoms except O and H.
- Balance O by adding H2O.
- Balance H by adding H+.
- Balance charge by adding electrons.
- Multiply the halves so the electrons cancel, then add.
In basic medium
Balance exactly as for acid first, then add OH− to both sides equal to the number of H+ present. The H+ and OH− combine to water, and any water appearing on both sides is cancelled. Doing this conversion at the end rather than the start is much less error-prone.
Types of redox reaction worth recognising
| Type | What happens | Example |
|---|---|---|
| Combination | Two species combine with electron transfer | Formation of an oxide from its elements |
| Decomposition | One compound breaks up with electron transfer | Decomposition of a chlorate on heating |
| Displacement | A more reactive element displaces a less reactive one | Zinc displacing copper from its salt solution |
| Disproportionation | The same element is both oxidised and reduced | Decomposition of hydrogen peroxide |
Disproportionation is the one most often missed. It requires the element to have an intermediate oxidation state — one it can move both up and down from. An element already at its highest or lowest state cannot disproportionate, and that reasoning is a standard question.
Frequently asked questions
Is oxidation number the same as valency?
No. Valency is the combining capacity and has no sign. Oxidation number is a signed bookkeeping device that may even be fractional in some species, which valency never is.
Why can an oxidation number be fractional?
Because it is an average across identical atoms in different environments. A fractional value signals that the atoms are not all equivalent, not that anything is wrong with the calculation.
Which method should I use in an exam?
Use the ion-electron method for ionic equations in solution, especially where medium matters, and the oxidation-number method for simpler molecular equations. Both give the same answer; the ion-electron method makes charge balance harder to forget.
How do I spot the oxidising and reducing agent quickly?
The oxidising agent is the species that gets reduced, and the reducing agent is the one that gets oxidised. Students routinely swap these, so it is worth stating the rule to yourself each time until it is automatic.
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