H2SO4 Properties: Sulphuric Acid Explained — Structure, Reactions and Uses

Sulphuric acid, H₂SO₄, is the most-produced industrial chemical in the world and one of the most-asked in Class 11 and 12 papers. Its exam questions come down to three roles it plays — a strong dibasic acid, a dehydrating agent and an oxidising agent — plus how it is made.

Structure and physical properties

• Sulphur is at the centre, sp³ hybridised, tetrahedral: two S=O double bonds and two S–OH single bonds.
• Molar mass 98 g/mol.
• A colourless, oily, dense liquid (density 1.84 g/cm³ for the 98 % acid), boiling point about 338 °C.
• Miscible with water in all proportions with a large release of heat — hence "add acid to water, never water to acid."
• High boiling point and viscosity come from strong hydrogen bonding between molecules.

Chemical properties

1. A strong dibasic acid

It ionises in two steps:

H₂SO₄ → H⁺ + HSO₄⁻ (complete, Ka₁ very large) HSO₄⁻ ⇌ H⁺ + SO₄²⁻ (Ka₂ ≈ 1.2 × 10⁻²)

So it forms two series of salts: NaHSO₄ (sodium hydrogen sulphate) and Na₂SO₄ (sodium sulphate).

Acid reactions: with metals above hydrogen (Zn + H₂SO₄ → ZnSO₄ + H₂), with bases (2NaOH + H₂SO₄ → Na₂SO₄ + 2H₂O), with carbonates (Na₂CO₃ + H₂SO₄ → Na₂SO₄ + H₂O + CO₂).

Because it is non-volatile (high boiling point), it displaces more volatile acids from their salts: NaCl + H₂SO₄ → NaHSO₄ + HCl; 2NaNO₃ + H₂SO₄ → Na₂SO₄ + 2HNO₃. This is how HCl and HNO₃ are prepared in the laboratory.

2. A dehydrating agent

Concentrated H₂SO₄ has a strong affinity for water and removes the elements of water from compounds:

• Sugar chars: C₁₂H₂₂O₁₁ → 12C + 11H₂O (the black "carbon snake" demonstration).
• Copper sulphate crystals turn white: CuSO₄·5H₂O → CuSO₄ + 5H₂O.
• Formic acid gives carbon monoxide: HCOOH → CO + H₂O.
• Ethanol at 170 °C gives ethene: C₂H₅OH → C₂H₄ + H₂O.

It is used as a drying agent for gases that do not react with it (Cl₂, SO₂, CO₂ — but not NH₃ or H₂S).

3. An oxidising agent (hot, concentrated)

Hot concentrated H₂SO₄ is reduced to SO₂:

• Cu + 2H₂SO₄ → CuSO₄ + SO₂ + 2H₂O (copper is below hydrogen, so no H₂ — this is oxidation, not acid displacement).
• C + 2H₂SO₄ → CO₂ + 2SO₂ + 2H₂O.
• S + 2H₂SO₄ → 3SO₂ + 2H₂O.
• 2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O (so HBr and HI cannot be made from their salts with H₂SO₄ — they are oxidised; HCl can).

Manufacture: the Contact process

1. Burn sulphur (or roast a sulphide ore) to SO₂: S + O₂ → SO₂.
2. Oxidise SO₂ to SO₃ over a V₂O₅ catalyst at about 450 °C and 2 bar: 2SO₂ + O₂ ⇌ 2SO₃, ΔH = −196 kJ/mol. The reaction is exothermic and reduces moles, so by Le Chatelier low temperature and high pressure favour SO₃ — the actual conditions are a compromise between yield and rate; a catalyst supplies the rate.
3. Absorb SO₃ in 98 % H₂SO₄ to give oleum, H₂S₂O₇, then dilute with water: SO₃ + H₂SO₄ → H₂S₂O₇; H₂S₂O₇ + H₂O → 2H₂SO₄.

SO₃ is not dissolved directly in water because the reaction is so violent that it forms a mist of acid that cannot be condensed.

Uses

Fertilisers (superphosphate, ammonium sulphate), petroleum refining, manufacture of dyes, drugs and detergents, lead–acid batteries, pickling of metals, and as a laboratory reagent for HCl and HNO₃.

How it is tested

• Why H₂SO₄ is dibasic, with the two ionisation steps and both salt series.
• The Contact process: catalyst, conditions, why oleum, why not water.
• Distinguish acid displacement (NaCl) from oxidation (HBr, HI) — the classic trap.
• One reaction each as dehydrating and oxidising agent, balanced.
• The safety line: acid into water.