Redox Reactions: Oxidation, Reduction, and Everything You Need to Know
Why Your Phone Battery, Rusty Iron, and a Cut Apple All Share One Idea
Think about three everyday things: a freshly cut apple slowly browning on your kitchen counter, an iron railing turning orange-red after weeks of rain, and a smartphone battery quietly powering your screen without any fire or explosion.
These look completely unrelated. But they are all driven by exactly the same chemical process — the movement of electrons from one substance to another.
That process is called a redox reaction. And once you understand it, you will recognise it everywhere — in batteries, medicines, metabolism, metal extraction, water treatment, and combustion.
This article explains redox reactions from the ground up: what they are, how to identify them, how to find the oxidising and reducing agents, and what disproportionation means. By the end, you will be able to analyse any redox reaction with confidence.
Quick Answers
What is redox reaction?
A redox reaction is a chemical reaction in which electrons are transferred from one substance to another.
- Oxidation = loss of electrons → oxidation number increases
- Reduction = gain of electrons → oxidation number decreases
- Oxidation and reduction always happen at the same time — you cannot have one without the other.
- The oxidising agent accepts electrons and is itself reduced.
- The reducing agent donates electrons and is itself oxidised.
Memory Aid —
OIL: Oxidation Is Loss,
RIG: Reduction Is Gain.
What Is an Oxidation Number?
Before you can identify oxidation and reduction in a reaction, you need a way to track what happens to electrons. That tracking tool is the oxidation number.
The oxidation number (also called the oxidation state) of an atom is the charge that atom would carry if all the bonds in its compound were completely ionic — that is, if the more electronegative atom in every bond were awarded full ownership of the shared electrons.
This is a bookkeeping device, not a physical measurement. In covalent compounds such as water (H₂O) or carbon dioxide (CO₂), electrons are shared rather than transferred, but we still assign oxidation numbers to keep track of electron ownership.
Remember: Oxidation numbers are written as signed integers with the sign before the number: +2, −1, 0, and so on. Do not confuse this with ionic charge, where the sign comes after the number , e.g. Ca2+ (2+ for a magnesium ion, not +2).
The Seven Rules for Assigning Oxidation Numbers
Apply these rules in order. Earlier rules take priority over later ones.
| Rule | Rule Name | Statement | Example |
| 1 | Pure elements | Every atom in a pure element has oxidation number = 0 | Zn(s), Cl₂(g), S₈(s): each atom = 0 |
| 2 | Monatomic ions | Oxidation number = ionic charge | Na⁺ = +1; Fe³⁺ = +3; Cl⁻ = −1 |
| 3 | Fluorine | Always −1 in all compounds | HF: F = −1; OF₂: F = −1 |
| 4 | Oxygen | −2 in almost all compounds. Exceptions: peroxides (−1); compounds with F (+2) | H₂O: O = −2; H₂O₂: O = −1; OF₂: O = +2 |
| 5 | Hydrogen | + 1 bonded to non-metals; −1 bonded to metals (metal hydrides) | HCl: H = +1; NaH: H = −1 |
| 6 | Neutral molecules | Sum of all oxidation numbers = 0 | H₂O: 2(+1) + (−2) = 0 ✓ |
| 7 | Polyatomic ions | Sum of all oxidation numbers = ion charge | SO₄²⁻: oxidation numbers sum to −2 |