Sec 4 Chemistry Oxidation & Reduction notes

Chemistry borrowed "oxidation" from combustion: the substance that gains oxygen is oxidised. In CuO + H₂ → Cu + H₂O, copper(II) oxide loses oxygen (it is reduced) while hydrogen gains it (it is oxidised). The two always travel together — if one substance is reduced, another must be oxidised, which is why the reactions are called redox.

The hydrogen view handles reactions without oxygen: when ammonia NH₃ is converted to nitrogen N₂ it loses hydrogen, so it is oxidised; when hydrogen chloride gains hydrogen to become H₂ it is reduced. Every definition below picks out the same winners; you choose whichever is easiest to see in the equation in front of you.

When zinc displaces copper from copper(II) sulfate, Zn + CuSO₄ → ZnSO₄ + Cu, the ions swap partners but the real news is electron movement. Zinc atoms shed two electrons each: Zn → Zn²⁺ + 2e⁻ (oxidation, loss). Copper ions collect them: Cu²⁺ + 2e⁻ → Cu (reduction, gain). The electrons cancel when the half-equations are added back together.

Any metal displacement along the reactivity series works this way: the more reactive metal is always oxidised, the less reactive metal ion is always reduced. Displacement of halogens flips it — chlorine takes electrons from iodide: Cl₂ + 2KI → I₂ + 2KCl, where Cl₂ is reduced (Cl₂ + 2e⁻ → 2Cl⁻) and iodide is oxidised (2I⁻ → I₂ + 2e⁻).

An oxidation state is a signed number attached to each atom. The rules, in order of priority: free elements are 0; a single-element ion equals its charge; fluorine is −1; hydrogen is +1 (−1 only in metal hydrides); oxygen is −2 (−1 in peroxides); and the states in a neutral compound sum to zero, or to the ion charge.

Once states are assigned: an increase is oxidation, a decrease is reduction. In 2Mg + O₂ → 2MgO, magnesium climbs 0 → +2 (oxidised) and oxygen falls 0 → −2 (reduced). If no element changes state, the reaction is not redox — neutralisation and precipitation are the classic non-redox traps.

An oxidising agent accepts electrons (and is reduced); a reducing agent supplies electrons (and is oxidised). The laboratory heavyweights: acidified potassium manganate(VII) (purple → colourless when it oxidises something) and acidified potassium dichromate(VI) (orange → green). Common reducing agents: hydrogen, carbon, carbon monoxide, reactive metals, sulfur dioxide, and the sulfite/ion family.

Rusting is iron losing electrons to oxygen in the presence of water: iron is oxidised to hydrated iron(III) oxide, the flaky brown solid. All three ingredients — iron, oxygen and water — must be present; boil water free of air or dry the air free of vapour and iron survives. Salt speeds rusting by improving the electrolyte, which is why winter roads and ship hulls suffer.

Prevention is either barrier (paint, oil, grease, plastic — keep oxygen and water off the metal) or sacrificial (galvanising: coat with zinc, or bolt zinc blocks to a hull — zinc oxidises first because it is more reactive, protecting the iron even if the coating scratches).

The same chemistry browns a sliced apple and spoils cooking oil: oxygen oxidises food molecules. Antioxidants — vitamin C (ascorbic acid) is the textbook one — are reducing agents added deliberately: they are oxidised first, sparing the food. Packaging flushed with nitrogen removes the oxidant entirely.