Redox chemistry and Groups 1, 2 and 7
Redox and oxidation numbers
Oxidation number rules: elements 0; monatomic ions = charge; O usually −2 (−1 in peroxides); H +1 (−1 in metal hydrides); F always −1; sum = overall charge. Oxidation = increase in oxidation number (loss of electrons); reduction = decrease. Half-equations in acid: balance atoms other than O and H, add H₂O for O, H⁺ for H, then electrons for charge. Combine by matching electrons. Disproportionation: one element is both oxidised and reduced. Redox titrations: (self-indicating, purple → colourless) and iodine–thiosulfate (starch indicator).
≈ 45 minGroup 2 reactions and trends
Down group 2, atomic radius increases, first ionisation energy decreases, so reactivity increases. The metals react with water: Mg slowly (fast with steam to MgO + H₂), Ca–Ba readily to give hydroxide + H₂, e.g. . Hydroxides become more soluble down the group (Mg(OH)₂ sparingly, Ba(OH)₂ soluble), so solutions become more alkaline. Sulfates become less soluble (BaSO₄ insoluble: used to test for sulfate and in barium meals). Uses: Ca(OH)₂ and CaCO₃ neutralise acid soils; Mg(OH)₂ in antacids.
≈ 45 minHalogens and halides
Down group 17, halogens become darker and less volatile (stronger London forces), electronegativity and oxidising power decrease, while halide ions become stronger reducing agents (larger ions lose electrons more easily). A more reactive halogen displaces a less reactive halide. Silver nitrate test (acidified with HNO₃): Cl⁻ white (dissolves in dilute NH₃), Br⁻ cream (concentrated NH₃), I⁻ yellow (insoluble). With concentrated H₂SO₄, NaCl gives only HCl; NaBr gives Br₂ and SO₂; NaI gives I₂, S and H₂S. Chlorine disproportionates in water and in cold or hot NaOH.
≈ 50 min