Learn first: Non-metals and their compounds
The halogens (Greek for "salt formers") are the Group VIIA elements: fluorine F, chlorine Cl, bromine Br, iodine I and radioactive astatine At. Their atoms have 7 outer electrons, one short of an octet, so the halogens are the most reactive non-metals and strong oxidising agents. With metals they form salts directly (NaCl, KBr), hence their name. Chlorine is widely used to disinfect drinking water and to make PVC, medicines and bleaches; fluorine is used in toothpaste and PTFE; iodine in medicine (tincture of iodine). Exam questions often cover the reactivity order of the halogens, the reactions of chlorine with alkalis and calculations with silver halides.
Electron configuration ns² np⁵; typical oxidation state in compounds −1
Down the group: radius increases, electronegativity and oxidising power decrease: F₂ > Cl₂ > Br₂ > I₂
The elements consist of diatomic molecules (F₂, Cl₂, Br₂, I₂) with non-polar covalent bonds. Their physical state changes with molar mass: F₂ is a pale yellow gas, Cl₂ a yellow-green gas, Br₂ a red-brown liquid and I₂ black-violet crystals that sublime on heating (turn to vapour without melting). All halogens are toxic and have sharp smells.
Fluorine is the most electronegative element, so in compounds it is always −1. Chlorine, bromine and iodine also show +1, +3, +5 and +7 in compounds with oxygen: HClO (+1), HClO₂ (+3), HClO₃ (+5), HClO₄ (+7). In the chlorine oxoacids, acid strength increases and oxidising power decreases as chlorine’s oxidation state rises.
A more reactive halogen displaces a less reactive one from a solution of its salt:
Cl₂ + 2KBr → 2KCl + Br₂; Br₂ + 2KI → 2KBr + I₂; Cl₂ + 2KI → 2KCl + I₂
I₂ + KBr — no reaction; Br₂ + NaCl — no reaction
Fluorine reacts violently with water, releasing oxygen: 2F₂ + 2H₂O → 4HF + O₂, so it cannot be used for displacement in salt solutions. Reactions with hydrogen follow the same order: fluorine combines explosively even in the dark and cold, chlorine explosively in light, bromine on heating, and iodine reversibly.
Excess chlorine was bubbled through a solution containing 0.2 mol of potassium bromide. What mass of bromine is released?
Preparation:
Laboratory: MnO₂ + 4HCl → MnCl₂ + Cl₂ + 2H₂O; 2KMnO₄ + 16HCl → 2KCl + 2MnCl₂ + 5Cl₂ + 8H₂O
Industry: electrolysis of NaCl solution 2NaCl + 2H₂O → H₂ + Cl₂ + 2NaOH
Properties:
15.8 g of potassium manganate(VII) reacted with excess concentrated hydrochloric acid. What volume of chlorine (STP) is released?
HCl is a colourless gas with a sharp smell, extremely soluble in water; its solution, hydrochloric acid, is a strong acid. Laboratory preparation: NaCl + H₂SO₄(conc.) → NaHSO₄ + HCl (heat). The strength of the hydrohalic acids rises in the order HF < HCl < HBr < HI because the H–Hal bond gets longer and weaker; their reducing power rises in the same order. HF is a weak acid but etches glass: SiO₂ + 4HF → SiF₄ + 2H₂O; it is made from fluorite: CaF₂ + H₂SO₄ → CaSO₄ + 2HF.
Oxygen compounds: potassium chlorate KClO₃ releases oxygen on heating (with an MnO₂ catalyst): 2KClO₃ → 2KCl + 3O₂; bleaching powder CaOCl₂ and sodium chlorate(I) are bleaches and disinfectants; HClO₄ is one of the strongest inorganic acids.
Halide ions are detected with silver nitrate:
Cl⁻ → AgCl↓ white; Br⁻ → AgBr↓ cream; I⁻ → AgI↓ yellow; F⁻ — no precipitate (AgF is soluble)
AgCl dissolves in ammonia solution, AgI does not. Iodine gives a deep blue colour with starch — the test for iodine (and for starch). Chlorine water plus starch distinguishes iodide from bromide.
Biological importance: chloride ions are part of the hydrochloric acid in gastric juice and of blood plasma; iodine is needed for thyroid hormones, and its deficiency causes goitre — so table salt is iodised; fluoride ions strengthen tooth enamel.
840 mL of 0.5 M AgNO₃ was added to a solution of 25 g of a mixture of NaCl and KCl, and the precipitate was filtered off. A 100 g copper plate was placed in the filtrate; its mass became 101.52 g. Find the mass fraction of NaCl in the mixture.