What is oxidation?

The three definitions of oxidation

Chemists' idea of oxidation grew over time, so you will meet three definitions. Each one has a matching definition of reduction, which is the opposite.

The electron definition is the most general, because it covers reactions with no oxygen or hydrogen at all. AQA GCSE Chemistry uses the oxygen definition for everyone and adds the electron one at Higher tier: "Oxidation is the loss of electrons and reduction is the gain of electrons." OpenStax Chemistry, used in many US courses, defines oxidation as loss of electrons, or an increase in oxidation number.

Oxidation and reduction always happen together

Electrons can't just disappear. If one substance loses electrons, another must gain them. So every oxidation is paired with a reduction, and the whole reaction is called a redox reaction (reduction plus oxidation).

Example: magnesium burning

By the oxygen definition, magnesium gains oxygen, so it is oxidised. By the electron definition, each magnesium atom loses 2 electrons to become Mg²⁺, and each oxygen atom gains 2 electrons to become O²⁻. So magnesium is oxidised and oxygen is reduced. Both definitions agree.

Example: zinc in copper sulfate solution

Drop zinc into blue copper sulfate solution and copper metal appears while the blue fades. There is no oxygen gained or lost here, so you need the electron definition. Leaving out the sulfate ions, which don't change, gives the ionic equation:

Zinc atoms lose 2 electrons each (oxidised), and copper ions gain them (reduced). You can split this into two half equations, one for each change:

In an oxidation half equation the electrons are on the right, as a product. In a reduction they are on the left. The electrons lost must equal the electrons gained, which is why both halves here have 2e⁻.

Oxidising agents and reducing agents

An oxidising agent (oxidizing agent in the US) causes something else to be oxidised by taking its electrons, so the oxidising agent itself is reduced. A reducing agent gives electrons away, so it is oxidised. In the zinc reaction, Cu²⁺ is the oxidising agent and zinc is the reducing agent. The names feel backwards at first: they describe what the substance does to the other one.

• Common oxidising agents: oxygen, chlorine, hydrogen peroxide, potassium manganate(VII).
• Common reducing agents: reactive metals, carbon, carbon monoxide, hydrogen.

Iron is made this way. In a blast furnace, carbon monoxide removes oxygen from iron oxide: Fe₂O₃ + 3CO → 2Fe + 3CO₂. The iron oxide loses oxygen (reduced), and the carbon monoxide gains it (oxidised). Check that the equation balances with our Equation Balancer.

Oxidation numbers (oxidation states)

For complicated molecules, chemists track electrons with oxidation numbers, also called oxidation states. An increase in oxidation number means oxidation; a decrease means reduction. These rules, as set out in OpenStax Chemistry, give the number for each atom:

1. An uncombined element is 0, for example Mg, O₂ or Cu.
2. A simple ion's oxidation number is its charge: Zn²⁺ is +2, Cl⁻ is −1.
3. Fluorine is always −1.
4. Oxygen is usually −2 (but −1 in peroxides).
5. Hydrogen is usually +1 with non-metals (but −1 with metals).
6. The numbers add up to the overall charge: zero for a neutral compound, or the ion's charge.

Worked example: chromium in potassium dichromate

Find the oxidation number of chromium in K₂Cr₂O₇. Potassium is +1 (group 1), oxygen is −2, and the compound is neutral. Call chromium x:

So 2 + 2x − 14 = 0, giving 2x = 12 and x = +6. That is why it is also called potassium dichromate(VI): the Roman numeral is the oxidation number.

How to tell if a reaction is a redox reaction

1. Write the balanced equation.
2. Give every atom an oxidation number, on both sides.
3. Look for changes. If any element’s number goes up, that element is oxidised; if any goes down, it is reduced. If nothing changes, it isn’t redox.

When sulfur dioxide reacts with oxygen, 2SO₂ + O₂ → 2SO₃, sulfur goes from +4 to +6 (oxidised) and the oxygen molecule's atoms from 0 to −2 (reduced), so it is redox. In neutralisation, HCl + NaOH → NaCl + H₂O, hydrogen stays +1, chlorine −1, sodium +1 and oxygen −2 throughout. Nothing changes, so neutralisation is not a redox reaction.

Oxidation in organic chemistry

With carbon compounds, the oxygen and hydrogen definitions are usually the easiest to use. Wine left open to the air turns sour because microbes oxidise the ethanol in it to ethanoic acid, the acid in vinegar: the molecule gains an oxygen atom and loses two hydrogen atoms.

Oxidation in electrolysis and batteries

In electrolysis, AQA states that negative ions lose electrons at the positive electrode (the anode), so those reactions are oxidations, while positive ions gain electrons at the negative electrode (the cathode), which are reductions. Remember "an ox, red cat": anode oxidation, reduction cathode. For example, chloride ions are oxidised at the anode: 2Cl⁻ → Cl₂ + 2e⁻. Batteries and fuel cells run redox reactions too, with the oxidation and reduction kept apart so the electrons have to flow through your device.

Examples of oxidation in everyday life

• Rusting: iron reacts with oxygen and water to form hydrated iron(III) oxide. Iron is oxidised from 0 to +3.
• Burning fuels: methane, petrol and wood are oxidised to carbon dioxide and water, releasing energy.
• Respiration: your cells oxidise glucose to release the energy you live on.
• Browning fruit: a cut apple goes brown as compounds in it are oxidised by oxygen in the air, helped by an enzyme.
• Food going off: fats oxidise and taste rancid, which is why foods contain antioxidants.
• Bleach: it works as an oxidising agent, breaking down coloured compounds and killing microbes.

Common mistakes

• Thinking oxidation needs oxygen. Under the electron definition, the zinc and copper ion reaction is oxidation with no oxygen involved.
• Mixing up the agents. The oxidising agent is reduced; the reducing agent is oxidised.
• Putting electrons on the wrong side. Oxidation half equations have electrons as products.
• Forgetting the partner reaction. If something is oxidised, something else must be reduced.
• Saying a positive ion was oxidised when it gained electrons. Gaining electrons is reduction, even for a positive ion.
• Unbalanced electrons. Electrons lost in the oxidation must equal electrons gained in the reduction.

Read next

Chemistry · Reactions and electrolysisHow to find oxidation numbers: rules and worked examplesHow to find oxidation numbers step by step: the rules, worked examples for compounds and ions, the exceptions, and using them to spot redox.