A colourless oily liquid, H₂SO₄; r.d. 1.84; m.p. 10.36°C; b.p. 338°C. The pure acid is rarely used; it is commonly available as a 96–98% solution (m.p. 3.0°C). The compound also forms a range of hydrates: H₂SO₄.H₂O (m.p. 8.62°C); H₂SO₄.2H₂O (m.p. –38/39°C); H₂SO₄.6H₂O (m.p. –54°C); H₂SO₄.8H₂O (m.p. –62°C). Its full systematic name is tetraoxosulphuric(VI) acid.

Until the 1930s, sulphuric acid was manufactured by the lead-chamber process, but this has now been replaced by the contact process (catalytic oxidation of sulphur dioxide). More sulphuric acid is made in the UK than any other chemical product; production levels (UK) are commonly 12 000 to 13 000 tonnes per day. It is extensively used in industry, the main applications being fertilizers (32%), chemicals (16%), paints and pigments (15%), detergents (11%), and fibres (9%).

In concentrated sulphuric acid there is extensive hydrogen bonding and several competing equilibria, to give species such as H₃O⁺, HSO₄⁻, H₃SO₄⁺, and H₂S₂O₇. Apart from being a powerful protonating agent (it protonates chlorides and nitrates producing hydrogen chloride and nitric acid), the compound is a moderately strong oxidizing agent. Thus, it will dissolve copper:

It is also a powerful dehydrating agent, capable of removing H₂O from many organic compounds (as in the production of acid anhydrides). In dilute solution it is a strong dibasic acid forming two series of salts, the sulphates and the hydrogensulphates.

Sulphur acids