What it is
Oxidation is a substance losing electrons and reduction is a substance gaining electrons. Electrons do not disappear, so when one substance is oxidised another is reduced: together that is a redox reaction. The substance that is oxidised is the reducing agent (it supplies the electrons) and the one that is reduced is the oxidising agent.
To follow what happens, give each atom an oxidation number, the charge it would have if the bonding electrons were assigned to the more electronegative atom. An element has oxidation number 0, a simple ion its charge, and O is usually −2. When an atom is oxidised its oxidation number goes up; when it is reduced it goes down.
Examples of redox
- Zinc and copper ions. Zn + Cu²⁺ → Zn²⁺ + Cu. Zn goes from 0 to +2 (loses 2 e⁻, oxidised) and Cu goes from +2 to 0 (gains 2 e⁻, reduced). The reaction splits into two half-reactions. From the standard reduction potentials, E° = 0.3419 V for Cu²⁺/Cu and −0.7618 V for Zn²⁺/Zn, so the cell gives 0.3419 − (−0.7618) = 1.1037 V.
- Magnesium burning. 2 Mg + O₂ → 2 MgO. Each Mg goes from 0 to +2 (4 e⁻ lost in all) and O₂ gains the same number as O goes from 0 to −2.
- Rust. 4 Fe + 3 O₂ → 2 Fe₂O₃. Fe goes from 0 to +3 (12 e⁻ lost in all) and the oxygen gains 12 e⁻.
- Burning methane. CH₄ + 2 O₂ → CO₂ + 2 H₂O. C goes from −4 to +4 (loses 8 e⁻), O from 0 to −2.
Not redox
- Neutralisation. HCl + NaOH → NaCl + H₂O. H is +1, Cl −1, Na +1 and O −2 before and after: no atom changes oxidation number.
- Precipitation. Ag⁺ + Cl⁻ → AgCl. Ag is +1 and Cl −1 before and after.
- Decomposing limestone. CaCO₃ → CaO + CO₂. Ca is +2, C +4 and O −2 throughout.
Why it matters
Redox is the chemistry behind batteries, corrosion, combustion, photosynthesis and cellular respiration. It leads into electrochemistry: every galvanic cell is a redox reaction in which the electrons are made to pass through an external circuit.
- The cell potential is worked out from the standard reduction potentials, and the Nernst equation corrects it for concentrations.
- Faraday’s law links charge to amount of substance at an electrode, and the decomposition voltage says how much voltage an electrolysis needs.
- The chemical equation balancer balances the atoms in an equation such as Zn + CuSO4. It does not handle ions and charge yet, so half-reactions are balanced by hand.
- Ions: ions.
Common misconceptions
- “Oxidation means gaining oxygen.” That is where the word came from, but today a substance is oxidised when it loses electrons, with or without oxygen. Mg reacting with Cl₂ is oxidised too.
- “The oxidising agent is oxidised.” It is the other way round: the oxidising agent is reduced, because it is the one that takes the electrons from the other substance.
- “Reduction means something gets smaller.” It means electrons are gained, so the oxidation number goes down.
- “There can be oxidation without reduction.” The electrons have to go somewhere, so the two belong together.
Common questions
How do I tell which substance is oxidised?
Write the oxidation numbers of the atoms before and after the reaction. The atom whose number goes up is
oxidised, the one whose number goes down is reduced. In Zn + Cu²⁺ → Zn²⁺ + Cu, Zn goes up from 0 to +2.
Are all reactions redox reactions?
No. Acid–base reactions, precipitation reactions and the decomposition of limestone change no oxidation
numbers and are not redox.
Why split a redox reaction into half-reactions?
Then you can see which electrons are lost and gained and balance them so that they cancel in the overall
reaction. It is also how a galvanic cell is built: one half-reaction at each electrode.