Identifying Redox Reactions and Oxidation Numbers

A reaction is a redox reaction if the oxidation number of at least one element changes; the element whose oxidation number increases is oxidized, and the element whose oxidation number decreases is reduced. In brief: in 2Mg + O₂ → 2MgO, Mg goes from 0 to +2 (oxidized) and O from 0 to −2 (reduced); in HCl + NaOH → NaCl + H₂O no oxidation numbers change, so it is not a redox reaction.

Oxidation numbers make it possible to find redox reactions even when electrons are shared, as in covalent compounds. This article gives a short procedure, reviews the rules for oxidation numbers, shows patterns that signal a redox reaction, classifies common reactions in a table, works through examples with oxidation-number changes and electron counts, and gives some history.

What is the procedure?

1. Assign oxidation numbers to every atom on both sides.
2. Compare the numbers for each element.
3. If none changes, the reaction is not redox.
4. If some change, the element that increases is oxidized (its compound is the reducing agent), and the one that decreases is reduced (its compound is the oxidizing agent).

What are the oxidation number rules?

• Free elements: 0. Monatomic ions: their charge.
• Group 1: +1; group 2: +2; fluorine: −1; hydrogen: +1 (−1 in metal hydrides); oxygen: −2 (−1 in peroxides).
• The sum equals the charge of the species (0 for a neutral compound).

Which reactions are always redox?

• Any reaction in which a free element appears or disappears: single replacement, and synthesis or decomposition involving elements.
• Combustion.
• Metal with acid (hydrogen gas forms).

Double replacement (precipitation, neutralization, gas formation without change of oxidation number) is not redox.

Classifying reactions

Worked examples

1. Cl₂ + 2KBr → 2KCl + Br₂. Cl: 0 → −1 (reduced). Br: −1 → 0 (oxidized). Redox.
2. MnO₂ + 4HCl → MnCl₂ + Cl₂ + 2H₂O. Mn: +4 → +2 (reduced). Cl: −1 → 0 (oxidized; only 2 of 4 chlorine atoms change). MnO₂ is the oxidizing agent; HCl is the reducing agent.

3. Permanganate in acid. In MnO₄⁻ → Mn²⁺, manganese goes from +7 (4 × −2 + x = −1) to +2, a gain of 5 electrons per Mn.

4. Fe₂O₃ + 3CO → 2Fe + 3CO₂. Fe: +3 → 0 (reduced, 3 e⁻ each). C: +2 → +4 (oxidized, 2 e⁻ each). Total: 6 electrons transferred for 2 Fe and 3 C.

5. Not redox. BaCl₂ + Na₂SO₄ → BaSO₄ + 2NaCl: Ba +2, Cl −1, Na +1, S +6, O −2 throughout.
6. Hydrogen peroxide. In 2H₂O₂ → 2H₂O + O₂, oxygen is −1 in peroxide. Two oxygen atoms go to −2 (reduced) and two go to 0 (oxidized): H₂O₂ is both oxidant and reductant.

How do you count electrons transferred?

Multiply the change in oxidation number of the oxidized (or reduced) element by the number of atoms changing. In 2Mg + O₂ → 2MgO, 2 Mg × 2 = 4 electrons lost and 2 O × 2 = 4 electrons gained. The total number of electrons lost always equals the total number gained.

How did the method develop?

Oxidation numbers were developed in the early 20th century. Wendell Latimer used oxidation states systematically in his 1938 work on the potentials of the elements, and IUPAC later gave an exact definition.

Which mistakes are common?

Mistake 1: Treating all reactions of oxygen as oxidation of oxygen

Oxygen is reduced, the other element is oxidized.

Mistake 2: Forgetting that only some atoms of an element may change

In MnO₂ + 4HCl, half of the chlorine does not change.

Mistake 3: Calling acid-base reactions redox

No electron transfer occurs in neutralization.

Mistake 4: Using the wrong oxidation number for peroxides

Oxygen is −1 in H₂O₂.