3.2.3 Group 7 (Halogens)

When you go down Group 7, keep the two trends separate: electronegativity decreases while boiling point increases. They have different causes, so one explanation will not do both jobs.

Electronegativity falls because bonding electrons are farther from the nucleus and more shielded. Boiling point rises because the larger halogen molecules have more electrons, are more polarisable and experience stronger London forces.

Talk about the bonding pair when you explain the fall in electronegativity down Group 7. It is farther from the nucleus and more shielded by inner shells, so the attraction for that bonding pair becomes weaker.

Phrase the explanation in terms of attraction for a bonding pair of electrons. That keeps the explanation tied to electronegativity. Ionisation energy is a different property.

For the rise in halogen boiling point, use London forces. From F₂ to I₂ the electron clouds get larger and more polarisable, so the temporary-dipole attractions become stronger and more energy is needed to separate the molecules.

That is why boiling point increases. Do not invoke permanent dipoles or hydrogen bonding; London forces are sufficient for the trend.

Halogens become weaker oxidising agents down the group. A halogen higher in the group can oxidise halide ions below it: for example, Cl₂ oxidises Br⁻ and I⁻, while Br₂ oxidises I⁻.

For a displacement equation, follow electron transfer. Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂ works because chlorine gains electrons and bromide loses them.

Use the concentrated H₂SO₄ reactions to compare halide reducing power: follow the electrons down the group. I⁻ loses electrons most readily, Br⁻ is intermediate and Cl⁻ is much weaker as a reducing agent.

With concentrated H₂SO₄, chloride gives mainly acid–base chemistry; bromide can reduce sulfuric acid to SO₂; iodide is a stronger reducing agent and can reduce sulfur to lower oxidation states. Follow the sulfur product and the halide oxidation product separately.

Acidify the sample with dilute nitric acid, then add AgNO₃. Cl⁻ gives white AgCl, Br⁻ gives cream AgBr, and I⁻ gives yellow AgI.

Nitric acid is used before silver nitrate so interfering basic anions are removed without introducing extra halide ions. If identification needs strengthening, follow with the ammonia-solubility test.

AgCl dissolves in dilute NH₃, AgBr dissolves only in concentrated NH₃, and AgI remains insoluble even in concentrated NH₃. Use this after the precipitate-colour test when you need to distinguish the silver halides more securely.

Keep colour and ammonia behaviour as two observations: white/cream/yellow first, then the appropriate solubility response.

When you explain the halide test, give each reagent its own job. AgNO₃ supplies Ag⁺ to form the silver-halide precipitate, dilute HNO₃ removes interfering ions without adding halide, and NH₃ helps you distinguish AgCl, AgBr and AgI by solubility.

If the question asks why each reagent is present, give each one its own chemical job. A list of colours does not explain the test design.

Chlorine disproportionates in water: Cl₂ + H₂O ⇌ HCl + HClO. Chlorine starts at oxidation state 0 and ends at −1 in chloride and +1 in chlorate(I), so the same element is both reduced and oxidised.

If you are asked to justify disproportionation, show those oxidation-state changes explicitly. Naming both products without the electron-change argument is only a description of the reaction.

If the conditions involve chlorine water in sunlight, do not stop automatically at HClO. You may need the overall change 2Cl₂ + 2H₂O → 4HCl + O₂, which gives chloride ions and oxygen.

Do not assume every chlorine-water question stops at HClO. Check the stated conditions and products, then write the equation that matches them.

For chlorination, weigh the public-health benefit of killing disease-causing microorganisms against the hazards from chlorine and possible chlorinated by-products. Make the comparison at the concentrations and exposures actually involved.

A strong discussion compares the controlled treatment risk with the public-health consequences of inadequately disinfected water. Avoid treating “toxic” as an automatic argument either for or against chlorination.

Chlorine is added to water because it forms oxidising chlorine-containing species, including HClO, that kill microorganisms. The disinfectant chemistry works at low controlled concentrations in the water supply.

Connect the use to oxidation of biological material in microorganisms. You do not need to re-explain the whole Group 7 trend to answer this line.

When a question asks why chlorination can still be justified, do not stop at ‘chlorine is toxic’. Toxicity depends on dose, and controlled chlorination greatly reduces the risk of waterborne infection, so you need to compare the managed chemical risk with that public-health benefit.

For an evaluation, state both sides and make the comparison. Simply writing “chlorine is dangerous” misses why a toxic chemical can still be useful in a carefully controlled process.

Cold dilute NaOH disproportionates chlorine to chloride and chlorate(I): Cl₂ + 2OH⁻ → Cl⁻ + ClO⁻ + H₂O. Hypochlorite/chlorate(I)-containing solutions are used as bleaches and disinfectants.

Learn the cold, dilute condition with these products. The oxidation states are −1 and +1, which gives you a quick way to check that disproportionation has occurred.