11. Group 17

The halogens exist as diatomic molecules. Down Group 17, increasing electron count strengthens van der Waals forces, so melting and boiling points rise and volatility falls.

The visible colour also becomes deeper as electron clouds become more polarisable and electronic energy gaps change. Physical-state observations should be tied to the molecular trend.

Chlorine is a yellow-green gas, bromine a red-brown liquid and iodine a dark solid at room conditions. The change reflects stronger intermolecular attraction down the group.

The colour trend is not evidence of stronger covalent bonds within X₂ alone. Volatility and boiling point mainly reflect intermolecular forces between molecules.

Halogen molecules are diatomic, X₂, with a covalent bond between the two atoms. Down Group 17, atomic size and electron count increase, but the trend in X–X bond strength is not explained by molecular mass alone.

The attraction between the two nuclei and the shared pair competes with greater internuclear distance and electron–electron repulsion in larger molecules. Use the measured trend or supplied data rather than assuming every bond becomes stronger down the group.

When comparing Cl₂, Br₂ and I₂, distinguish the covalent bond inside each molecule from the van der Waals forces between molecules. The latter clearly become stronger as the electron cloud becomes more polarizable.

Higher boiling point does not prove a stronger X–X covalent bond; it mainly reflects intermolecular attraction.

Volatility is the tendency of a substance to enter the gas phase. For the halogens, volatility decreases down Group 17 as London dispersion forces between X₂ molecules become stronger.

More electrons make the electron cloud more polarizable. Temporary dipoles therefore induce stronger attractions, raising melting and boiling temperatures and changing the physical state at room temperature.

Chlorine is a gas, bromine a liquid and iodine a solid under ordinary conditions. This sequence is a physical-property trend, not a change from molecular to ionic bonding.

Do not explain the trend as stronger covalent bonds between halogen atoms. The key change is attraction between molecules.

A halogen acts as an oxidising agent when X₂ gains electrons to form halide ions: X₂ + 2e⁻ → 2X⁻. Oxidising power decreases from fluorine to iodine.

Down the group, the incoming electron is attracted less strongly because the atom is larger and more shielded. This makes reduction of X₂ less favourable, even though the molecules have more electrons.

Chlorine can oxidise bromide ions: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂. Bromine cannot similarly displace chloride under the same conditions.

Oxidising strength is not the same as volatility or colour. Decide it from electron-transfer evidence or displacement reactions.

Hydrogen reacts with a halogen to form a hydrogen halide: H₂ + X₂ → 2HX. The HX molecule is covalent in the gas phase, while its aqueous behaviour depends on how readily the H–X bond ionises.

Hydrogen halides dissolve in water to form acidic solutions. Down the group the H–X bond becomes longer and generally weaker, so bond cleavage and ion formation become easier; hydrogen fluoride is the important weak-acid exception in water.

Hydrogen chloride is a strong acid in aqueous solution, whereas hydrogen fluoride is weak and can form hydrogen bonds. Keep gas-phase molecules and aqueous acid strength as separate descriptions.

“Covalent” does not mean “non-acidic”; acid strength is about ionisation in the solvent.

Thermal stability describes resistance to decomposition on heating. Hydrogen halides become less thermally stable down Group 17 as the H–X bond lengthens and weakens.

A larger halogen atom gives poorer orbital overlap with hydrogen, so less energy is needed to break the bond. This bond-strength explanation is distinct from the aqueous acid-strength question, even though both involve H–X cleavage.

HI decomposes more readily than HCl when heated. A comparison should mention bond length/overlap, not merely say iodine is “more reactive”.

Thermal stability is not the same as boiling point, volatility or acid strength; identify the process being measured.

A halide ion is a reducing agent when it donates an electron and is oxidised to its halogen: 2X⁻ → X₂ + 2e⁻. Reducing power increases from fluoride to iodide.

The larger, more shielded ions hold their outer electrons less strongly. Iodide therefore loses an electron more readily, while fluoride is difficult to oxidise in aqueous solution.

Bromide reduces chlorine water, producing bromine: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂. The reverse displacement is not spontaneous under the same conditions.

This is the opposite trend to halogen oxidising power. Always state whether the species is gaining or losing electrons.

Aqueous halide ions form precipitates with silver ions: Ag⁺ + X⁻ → AgX. The precipitate colour and its response to aqueous ammonia provide a structured identification test.

Chloride, bromide and iodide give different AgX observations. Concentrated sulfuric acid gives a second line of evidence: halide ions reduce sulfuric acid to different extents, so balanced equations and oxidation states matter.

Add nitric acid then silver nitrate to an unknown. A pale precipitate that dissolves in dilute ammonia is consistent with AgCl; use the whole observation sequence rather than colour alone.

Ammonia is not just a washing step, and precipitate colour alone is not a complete identification. Keep test conditions and ionic equations explicit.

In a disproportionation reaction, the same element is both oxidised and reduced. Chlorine reacts with cold dilute sodium hydroxide to form chloride and chlorate(I), while hot concentrated alkali forms chloride and chlorate(V).

Use oxidation numbers to show the two pathways. In cold solution chlorine goes from 0 to −1 and +1; in hot solution it goes from 0 to −1 and +5. Balance atoms and charge before naming the reaction.

Cold: Cl₂ + 2OH⁻ → Cl⁻ + ClO⁻ + H₂O. Hot: 3Cl₂ + 6OH⁻ → 5Cl⁻ + ClO₃⁻ + 3H₂O. The conditions determine the product.

Do not use one equation for every sodium-hydroxide condition, and do not call the reaction simple oxidation or reduction.

When chlorine dissolves in water it establishes an equilibrium that produces hypochlorous acid: Cl₂ + H₂O ⇌ HCl + HOCl. HOCl can partially dissociate to H⁺ and OCl⁻.

HOCl and OCl⁻ are the active chlorine species that damage microorganisms. Their relative amounts depend on pH, so purification is a chemical equilibrium and a treatment decision, not simply “chlorine kills bacteria”.

A controlled chlorine dose leaves a measurable residual disinfectant after contact time. Lower pH shifts more active chlorine toward HOCl, but operational safety and local standards still constrain the choice.

Do not treat chloride ions as the disinfectant, or assume any chlorine dose is safe. Distinguish formation of active species from the public-health application.