Part 4: Halides and hydrogen halides

Follow the reversal from halogen oxidising power to halide reducing power, distinguish acid–base and redox stages, and identify halides with controlled precipitation tests.

Larger halide ions lose electrons more readily

A halide ion is a potential reducing agent because it can lose an electron. Down the group, its outer electrons are farther from the nucleus and more shielded, so electron loss becomes easier. Reducing ability increases from chloride to bromide to iodide; this is the opposite direction from halogen oxidising power.

With a solid Group 1 halide and concentrated sulfuric acid, first consider proton transfer to form a hydrogen halide. Bromide and iodide can then reduce sulfuric acid, so additional products appear. The acid’s role changes from proton donor in the initial stage to oxidising agent in the redox stage.

Chloride gives an acid–base reaction; bromide also gives redox

For sodium chloride, NaCl(s) + H₂SO₄(l) → NaHSO₄(s) + HCl(g). Steamy acidic fumes appear as HCl encounters moisture. Chloride is not a strong enough reducing agent to reduce sulfuric acid under these conditions, so this is an acid–base reaction without oxidation-number changes. Fluoride gives the analogous initial acid–base pattern, but HF requires specialist hazard controls and is not a routine student preparation.

Sodium bromide initially gives HBr. Some HBr/bromide then reduces sulfur in H₂SO₄ from +6 to +4 in SO₂ while bromide goes from −1 to 0 in Br₂. Orange-brown bromine vapour appears as well as colourless sulfur dioxide; steamy HBr fumes may also be present. Bromide does not normally reduce the sulfur as far as elemental sulfur in this standard comparison.

Use the two stages to organise an answer. The molecular redox equation below uses HBr, while a valid acidic ionic version can use sulfate, H⁺ and Br⁻. Do not mix half of each representation without checking atoms and charge.

Iodide can reduce sulfur to three different products

Iodide is the stronger reducing agent. After HI forms in the initial acid–base step, sulfuric acid may be reduced to SO₂, sulfur or H₂S, where sulfur has oxidation numbers +4, 0 and −2 respectively. Iodide is oxidised to iodine, seen as a dark solid and purple vapour. Sulfur can appear as a yellow solid; SO₂ and H₂S are colourless gases.

From sulfur’s initial +6 state, the electron gains are two to form SO₂, six to form S and eight to form H₂S. Each iodide loses one electron, so the corresponding equations consume 2, 6 or 8 HI. The equations describe possible product channels, not a claim that every experiment gives fixed equal amounts of all products.

Hydrogen sulfide is highly toxic and sulfur dioxide is harmful. Although characteristic odours are sometimes described in data, never use deliberate smelling as the identification method. These concentrated-acid reactions require a fume cupboard and a prescribed microscale demonstration with suitable protection.

Hydrogen halides form acids in water and ammonium salts with ammonia

Hydrogen chloride, bromide and iodide are molecular substances as dry gases. They react with water to form aqueous hydrogen/hydroxonium ions and halide ions, giving acidic conducting solutions. For HCl, HCl(g) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq). Dry HCl is not already a bottle of aqueous H⁺ ions.

HF is highly soluble but only partially ionises in dilute water, so do not infer that high solubility means complete ionisation or that every hydrogen halide is equally strong as an aqueous acid. Aqueous acid strength is distinct from the reducing power of the halide ion discussed above.

A hydrogen halide reacts with ammonia to give fine white solid particles of the ammonium halide. HCl + NH₃ → NH₄Cl; analogous equations give NH₄Br and NH₄I. Describe a white smoke/fume of solid salt, rather than a white gas. These are acid–base reactions, not changes in halogen oxidation number.

If a preparation of HBr or HI is required, concentrated sulfuric acid is unsuitable when its oxidising action introduces halogen and sulfur products. A non-oxidising acid such as phosphoric acid can supply H⁺ without that competing redox chemistry. Any practical method must account for the gases’ corrosiveness and very high solubility in water.

Acidify with nitric acid before adding silver nitrate

For an aqueous unknown, add dilute nitric acid, then aqueous silver nitrate. Acid removes carbonate or hydroxide interference that could otherwise produce silver precipitates. Nitric acid is used because nitrate does not introduce a competing insoluble silver salt. Hydrochloric acid would add Cl⁻ and create AgCl, confusing the test.

A precipitate’s colour suggests the halide, but similar colours are better distinguished by its solubility in aqueous ammonia. Test the precipitate with dilute ammonia first, then concentrated ammonia if needed using the approved procedure. Silver nitrate can damage/stain tissue and ammonia fumes are irritating; use small quantities and appropriate eye protection and ventilation.

Pearson 9CH0/01 June 2023 Q3(b)(ii–iii) required identifying the interference and why an unsuitable acid would form another precipitate. The report explains why ‘removes impurities’ or ‘prevents a false positive’ alone is too vague in that context. State the chemical species and unwanted observation.

Ammonia changes the silver-ion equilibrium

Ammonia binds to Ag⁺ to form the soluble complex [Ag(NH₃)₂]⁺. This lowers the concentration of free Ag⁺ and can draw more silver halide into solution. AgCl dissolves readily; AgBr needs more concentrated ammonia; AgI is too sparingly soluble to dissolve appreciably under the test conditions.

At AS, learn the observed distinctions and correct ionic precipitation equations. The complex equation provides a useful link to later transition-metal/equilibrium chemistry: AgCl(s) + 2NH₃(aq) ⇌ [Ag(NH₃)₂]⁺(aq) + Cl⁻(aq). Ammonia is not reducing Ag⁺ to silver metal in this test.

Quick checks

Original Finesse questions. Reveal the indicative worked solutions after attempting each question; these are not official Edexcel mark allocations.

Sources