What Redox Means
Oxidation and reduction always happen together — one cannot occur without the other, which is why they’re jointly called redox. The old definitions (oxidation as gaining oxygen or losing hydrogen; reduction as the reverse) only work when oxygen or hydrogen is actually involved. The modern, general definition is framed instead around oxidation number (also called oxidation state): oxidation is a rise in an element’s oxidation number, reduction is a fall.
Rules for Assigning Oxidation Numbers
- An element in its free (uncombined) state has oxidation number zero — e.g. H₂, Br₂, Cl₂, S₈, P₄, or a lone atom of any metal.
- For a simple ion, the oxidation number equals the ion’s charge — e.g. Ca²⁺ is +2, O²⁻ is −2, N³⁻ is −3, Na⁺ is +1.
- Hydrogen is +1 in most compounds, but −1 when bonded to a metal (a hydride) — e.g. +1 in H₂O and HCl, but −1 in NaH and MgH₂.
- Some elements almost always keep the same oxidation number, which makes them useful reference points for working out others: fluorine is always −1; oxygen is −2 (except when bonded to fluorine, or in peroxides); chlorine is −1 (except when bonded to fluorine or oxygen); hydrogen is +1 (except when bonded to a metal).
- In a polyatomic ion, the oxidation numbers of all the atoms in it must add up to the ion’s overall charge.
Worked Examples
a) CO₂: oxygen is fixed at −2, and there are two of them, so (ox. state of C) + (−2 × 2) = 0, giving carbon = +4.
b) KMnO₄: potassium is +1, oxygen is −2 (×4), so (+1) + (ox. state of Mn) + (−2 × 4) = 0, giving manganese = +7.
c) Fe(NO₃)₂: the nitrate ion (NO₃⁻) is a unit with charge −1, and there are two of them, so (ox. state of Fe) + (−1 × 2) = 0, giving iron = +2.
Many elements — especially transition metals — take different oxidation states in different compounds, and Roman numerals in a compound’s name signal which one: iron(II) chloride, FeCl₂, has iron at +2; potassium dichromate(VI), K₂Cr₂O₇, has chromium at +6; manganese(IV) oxide, MnO₂, has manganese at +4.
| Element | Oxidation states seen (with examples) |
|---|---|
| Manganese | 0 (Mn), +2 (MnCl₂), +4 (MnO₂), +7 (KMnO₄) |
| Chromium | +2 (CrCl₂), +3 (CrCl₃), +6 (K₂Cr₂O₇) |
| Iron | +2 (FeCl₂), +3 (FeCl₃) |
| Sulphur | −2 (FeS), 0 (S), +4 (SO₂), +6 (H₂SO₄) |
| Carbon | 0 (C), +2 (CO), +4 (CO₂) |
Oxidation and Reduction, in Terms of Electrons
Oxidation is a reaction in which an element’s atoms lose electrons and its oxidation number rises. The substance being oxidised is called the reducing agent (RA) — it “gives away” electrons to make something else’s reduction possible.
Reduction is a reaction in which an element’s atoms gain electrons and its oxidation number falls. The substance being reduced is called the oxidising agent (OA) — it “takes” electrons, enabling something else’s oxidation.
Worked examples, with oxidation numbers marked and the reducing agent (RA) and oxidising agent (OA) identified:
- 2NH₃ (N: −3) + 3Cl₂ (Cl: 0) → N₂ (N: 0) + 6HCl (Cl: −1) — nitrogen rises from −3 to 0 (oxidised, so NH₃ is the RA); chlorine falls from 0 to −1 (reduced, so Cl₂ is the OA).
- O₂ (O: 0) + 2SO₂ (S: +4) → 2SO₃ (S: +6) — sulphur rises from +4 to +6 (oxidised, so SO₂ is the RA); oxygen falls from 0 to −2 in the product (reduced, so O₂ is the OA).
- 2Cu (0) + O₂ (0) → 2CuO (Cu: +2, O: −2) — copper rises from 0 to +2 (oxidised, so Cu is the RA); oxygen falls from 0 to −2 (reduced, so O₂ is the OA).
Common Oxidising and Reducing Agents
| Oxidising agents | Reducing agents |
|---|---|
| O₂, MnO₄⁻, H₂O₂, acidified KMnO₄, acidified K₂Cr₂O₇, Cl₂, concentrated HNO₃, concentrated H₂SO₄, non-metals generally | H₂, C, CO, H₂S, SO₂, NH₃, KI, LiAlH₄, metals generally |
