Occurrence and Electronic Structure
Sulphur sits in Group VI, alongside oxygen, selenium, tellurium and polonium; electronic configuration 1s²2s²2p⁶3s²3p⁴. It’s extracted by the Frasch process. It has two main allotropes, rhombic and monoclinic sulphur. Sulphur shows a wide range of oxidation states, from −2 to +7, because its available 3d-orbitals let it expand its octet.
Uses: manufacturing bleaching agents; producing dyes and skin ointments; sulphur and its compounds as fungicides.
Hydrogen Sulphide
Prepared from hydrochloric acid and iron(II) sulphide: 2HCl(aq) + FeS(s) → FeCl₂(aq) + H₂S(g). A colourless gas with the characteristic smell of rotten eggs.
Reactions of H₂S — always as a reducing agent:
Reduces brownish-yellow iron(III) chloride to green iron(II) chloride: 2FeCl₃(aq) + H₂S(g) → 2FeCl₂(aq) + 2HCl(aq) + S(s)
Decolourises purple potassium permanganate (Mn reduced from +7 to +2): 2MnO₄⁻(aq) + 6H⁺(aq) + 5H₂S(g) → 2Mn²⁺(aq) + 8H₂O(l) + 5S(s)
Turns orange potassium dichromate green (Cr reduced from +6 to +3): Cr₂O₇²⁻(aq) + 8H⁺(aq) + 3H₂S(g) → 2Cr³⁺(aq) + 7H₂O(l) + 3S(s)
Oxides of Sulphur
Sulphur dioxide forms when sulphur burns in air — a dense, colourless gas with a choking, pungent smell: S(s) + O₂(g) → SO₂(g). Sulphur trioxide forms when sulphur dioxide reacts further with oxygen, in the presence of a vanadium(V) oxide catalyst: 2SO₂(g) + O₂(g) → 2SO₃(g).
Sulphuric Acid — the Contact Process
Sulphur (from the Frasch process) is roasted in air to sulphur dioxide, S(s) + O₂(g) → SO₂(g). The SO₂ is mixed with excess air and purified through a dust precipitator (removing dust and catalyst-poisoning impurities like arsenic compounds), then dried by passing through concentrated sulphuric acid. The dried mixture passes over a vanadium(V) oxide catalyst at about 450°C and 1 atm to form sulphur trioxide: 2SO₂(g) + O₂(g) → 2SO₃(g). The SO₃ is cooled and dissolved in existing concentrated sulphuric acid (not water directly — SO₃ dissolving straight into water forms a difficult-to-control acid mist) to form oleum, or fuming sulphuric acid: SO₃(g) + H₂SO₄(l) → H₂S₂O₇(l). The oleum is then diluted with a calculated amount of water to give 98% sulphuric acid: H₂S₂O₇(l) + H₂O(l) → 2H₂SO₄(l).
Properties of sulphuric acid: a strong acid, ionising in two stages: H₂SO₄(l) + H₂O(l) → HSO₄⁻(aq) + H₃O⁺(aq); HSO₄⁻(aq) + H₂O(l) → SO₄²⁻(aq) + H₃O⁺(aq). Acts as a dehydrating agent — e.g. pulling water out of ethanol to form ethene, C₂H₅OH(l) → C₂H₄(g) + H₂O(l). Acts as a drying agent for acidic gases (not alkaline ones, which it would simply react with). Displaces more volatile acids — hydrochloric, hydrofluoric, nitric, phosphoric — from their salts, since sulphuric acid itself is comparatively non-volatile.
Uses: manufacturing sulphate fertilisers, detergents, and synthetic fibres.
Nitrogen
Occurrence and Oxidation States
Nitrogen is Group V, alongside phosphorus, arsenic, antimony and bismuth — general outer electronic configuration ns²np³. It shows oxidation states from −3 to +5 across its compounds.
Ammonia
Laboratory preparation: heating an ammonium salt with an alkali — 2NH₄Cl(s) + Ca(OH)₂(s) → CaCl₂(s) + 2H₂O(l) + 2NH₃(g) — and drying the gas with quicklime (CaO).
Industrial preparation — the Haber process: nitrogen (from the fractional distillation of liquid air) and hydrogen (from decomposing natural gas) are purified to remove dust and catalyst-poisoning impurities, dried, and mixed in a 1:3 ratio by volume. The mixture passes over a finely divided iron catalyst at roughly 450°C and 250 atm: N₂(g) + 3H₂(g) ⇌ 2NH₃(g). Yield under these conditions is only around 25% per pass — the ammonia formed is liquefied and drawn off, and the large excess of unreacted gas is recycled back over the catalyst.
Properties: dissolves in water to give a weakly alkaline solution (the resulting solution, NH₄OH, is called aqueous ammonia): NH₃(g) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq). Acts as a reducing agent — e.g. reducing heated copper(II) oxide: 3CuO(s) + 2NH₃(g) → 3Cu(s) + N₂(g) + 3H₂O(l). Burns in oxygen: 4NH₃(g) + 3O₂(g) → 2N₂(g) + 6H₂O(l).
Nitric Acid
Laboratory preparation: heating a mixture of concentrated sulphuric acid and potassium nitrate in a retort: H₂SO₄(l) + KNO₃(s) → KHSO₄(s) + HNO₃(g). The vapour condenses to a yellow, fuming liquid — the yellow colour comes from dissolved nitrogen dioxide, itself formed by the acid’s own partial decomposition: 4HNO₃(l) → 4NO₂(g) + O₂(g) + 2H₂O(l).
Industrial preparation — the Ostwald process: ammonia (from the Haber process) is mixed with excess air, purified and dried, then passed over a platinum-rhodium catalyst at around 850°C and 7 atm:
4NH₃(g) + 5O₂(g) → 4NO(g) + 6H₂O(l) (catalytic oxidation)
The hot gases are cooled and mixed with more excess air, oxidising the nitrogen monoxide further: 2NO(g) + O₂(g) → 2NO₂(g)
The nitrogen dioxide is then passed over hot water, again in the presence of excess air, to form nitric acid: 4NO₂(g) + O₂(g) + 2H₂O(l) → 4HNO₃(aq)
The acid produced this way is about 98% concentrated. In terms of pure equilibrium yield, low temperature would favour the (exothermic) forward reaction and low pressure would favour the product side of the first stage specifically (it has slightly more moles of gas than the reactant side) — but the process is actually run hot and under moderate pressure regardless, because that gives a far more useful reaction rate and better selectivity for NO over competing side-products; maximum equilibrium yield and practical industrial throughput aren’t the same target.
Properties: a very strong acid — neutralises bases/alkalis to give salt and water, e.g. NaOH(aq) + HNO₃(aq) → NaNO₃(aq) + H₂O(l). Dilute nitric acid liberates hydrogen with a moderately reactive metal like magnesium: Mg(s) + 2HNO₃(aq) → Mg(NO₃)₂(aq) + H₂(g).
Uses: manufacturing explosives, dyes, and fertilisers (e.g. potassium nitrate).
Fertilisers
Fertilisers are chemical compounds added to soil to supply nutrients essential for plant growth — chiefly nitrogen, phosphorus and potassium.
- Nitrogenous fertilisers: soluble nitrogen-containing compounds — calcium nitrate, ammonium nitrate, urea, ammonium sulphate, potassium nitrate.
- Phosphate fertilisers: soluble phosphorus-containing compounds — ammonium phosphate, calcium superphosphate.
- Potassium fertilisers: soluble potassium-containing compounds — potassium nitrate, potassium chloride, potassium sulphate.
Environmental effects: over-application means the excess dissolves in rainwater and washes into rivers, causing water pollution and, in turn, the death of aquatic organisms (eutrophication).
