Group 16 (Oxygen Family)

Ozone (O₃) is prepared in the laboratory by passing dry, pure oxygen through a silent electric discharge (ozoniser). The electrical energy breaks O₂ molecules into oxygen atoms (O•), which then combine with intact O₂ molecules to form O₃. The reaction is endothermic and requires continuous energy input. Approximately 10–15% conversion is typically achieved. Industrially, ultraviolet radiation at 185 nm also converts O₂ to O₃ in the upper atmosphere via the same two-step mechanism. The resulting gas mixture (ozonised oxygen) is pale blue when concentrated.

The ozone molecule has a bent geometry with a bond angle of approximately 117°. The two O–O bond lengths are equal at 128 pm, intermediate between a single (148 pm) and double (121 pm) bond, indicating resonance delocalisation. The central oxygen atom is sp²-hybridised. Ozone is a powerful oxidising agent (E° = +2.07 V in acidic solution), stronger than O₂, and can oxidise iodide to iodine (a qualitative test), bleach dyes, and kill bacteria. It decomposes slowly at room temperature and rapidly on contact with alkenes (ozonolysis). Ozone distinguishes itself from O₂ by liberating iodine from acidified KI solution: O₃ + 2KI + H₂SO₄ → K₂SO₄ + I₂ + H₂O + O₂, whereas O₂ alone cannot do so.

The Contact process is the industrial method for manufacturing concentrated sulphuric acid and proceeds in three main stages. Stage 1 — Sulphur dioxide production: Sulphur is burned in air (S + O₂ → SO₂), or sulphide ores such as pyrite are roasted (4FeS₂ + 11O₂ → 2Fe₂O₃ + 8SO₂). Stage 2 — Catalytic oxidation: SO₂ is oxidised to SO₃ over a vanadium pentoxide (V₂O₅) catalyst at 450–500 °C and 1–2 atm pressure. The reaction is reversible and exothermic, so an optimum temperature is chosen to balance reaction rate and equilibrium yield (Le Chatelier's principle). Stage 3 — Absorption: SO₃ cannot be absorbed directly in water because it forms an acid mist. Instead, it is absorbed in concentrated H₂SO₄ (98%) to form oleum (fuming sulphuric acid, H₂S₂O₇), which is then carefully diluted with water to produce the desired concentration of H₂SO₄.

The V₂O₅ catalyst operates via a redox cycle: it oxidises SO₂ to SO₃ and is itself reduced from V⁵⁺ to V⁴⁺, then reoxidised by atmospheric O₂ back to V⁵⁺. The gas mixture must be thoroughly purified before entering the catalyst chamber, as arsenic, dust, and fluoride compounds act as catalyst poisons. Modern double-absorption plants achieve over 99.5% conversion efficiency by removing SO₃ at an intermediate stage before passing unconverted gases over the catalyst a second time. The choice of 450–500 °C reflects the compromise: below this temperature the reaction is too slow; above it the equilibrium shifts back toward SO₂, reducing yield.

Sulphur forms a rich variety of oxoacids due to its ability to adopt multiple oxidation states (+2, +4, +6) and its capacity to form S–S bonds and S–O–S bridges. The most important oxoacids include: (1) Sulphurous acid (H₂SO₃, S oxidation state +4) — formed when SO₂ dissolves in water; it is a weak diprotic acid and a mild reducing agent, though it exists mainly as SO₂·H₂O in solution. (2) Sulphuric acid (H₂SO₄, S = +6) — a strong diprotic acid; concentrated H₂SO₄ is a powerful dehydrating and oxidising agent. (3) Peroxomonosulphuric acid (Caro's acid, H₂SO₅, S = +6) — contains a peroxo –O–O– linkage, strong oxidant. (4) Peroxodisulphuric acid (Marshall's acid, H₂S₂O₈, S = +6) — two sulphate groups bridged by –O–O–, prepared by electrolytic oxidation of H₂SO₄. (5) Pyrosulphuric acid (oleum, H₂S₂O₇, S = +6) — formed when SO₃ is absorbed in H₂SO₄. (6) Thiosulphuric acid (H₂S₂O₃, average S oxidation state = +2) — the sulphate ion in which one oxygen is replaced by sulphur; its sodium salt (Na₂S₂O₃) is widely used as a photographic fixer and in iodometric titrations. All these acids share the sulfonyl (–S(=O)–) core but differ in the number of –OH groups, –O–O– bridges, or additional sulphur atoms.

• \ce{SO2} — Sulphur dioxide (mol): Reactant gas; sulphur in +4 oxidation state; produced by burning sulphur or roasting pyrite
• \ce{O2} — Oxygen gas (mol): Co-reactant; supplied from air after purification
• \ce{V2O5} — Vanadium pentoxide (—): Heterogeneous catalyst; participates via V⁵⁺/V⁴⁺ redox cycle and is regenerated
• \ce{SO3} — Sulphur trioxide (mol): Product; sulphur in +6 oxidation state; absorbed in H₂SO₄ to form oleum
• \Delta H — Enthalpy change (kJ mol⁻¹): Reaction is exothermic (ΔH ≈ −196 kJ mol⁻¹); temperature must be controlled to maintain yield

Problem. A sample of ozonised oxygen contains 15% ozone by volume. When 500 cm³ of this mixture is passed through an excess of acidified potassium iodide (KI) solution, the liberated iodine is titrated against 0.10 mol dm⁻³ sodium thiosulphate (Na₂S₂O₃) solution. Calculate the volume of Na₂S₂O₃ solution required. [Reactions: O₃ + 2KI + H₂SO₄ → K₂SO₄ + I₂ + H₂O + O₂ ; I₂ + 2Na₂S₂O₃ → Na₂S₄O₆ + 2NaI]

Answer. 62.5 cm³ of 0.10 mol dm⁻³ sodium thiosulphate solution is required.

• Group 17 — The Halogens — Adjacent p-block group; comparing oxidising power, oxoacid formation, and interhalogen compounds deepens understanding of periodic trends seen in Group 16.
• Le Chatelier's Principle — Explains the choice of temperature and pressure in the Contact process; understanding equilibrium shifts is essential for optimising SO₃ yield.
• Oxidation States and Redox Reactions — Ozone's oxidising action, the V₂O₅ catalytic cycle, and the varied sulphur oxidation states in oxoacids all require a firm grasp of redox concepts.
• Heterogeneous Catalysis — The V₂O₅ catalyst in the Contact process is a classic example of heterogeneous catalysis; understanding surface reactions and catalyst poisoning is directly relevant.
• Strong and Weak Acids — The oxoacids of sulphur range from weak (H₂SO₃) to strong (H₂SO₄); comparing their dissociation behaviour links to acid-base equilibrium theory.