Nitric acid

Laboratory preparation: Nitric acid is prepared in the laboratory by gently heating equal amounts of potassium or sodium nitrate with concentrated sulphuric acid: KNO3 + H2SO4 -> KHSO4 + HNO3. The temperature is kept as low as possible to minimise decomposition of the acid as it forms; the product condenses as a fuming liquid, tinted brown by traces of nitrogen dioxide produced from the acid's own thermal decomposition (4HNO3 -> 4NO2 + 2H2O + O2).

Commercial preparation — Ostwald's process: On an industrial scale, nitric acid is manufactured by Ostwald's process. Ammonia from the Haber process is mixed with about ten times its volume of air, preheated, and passed into a catalyst chamber where it contacts a platinum gauze at about 1275 K; the gauze catalyses the rapid oxidation of ammonia to nitric oxide (4NH3 + 5O2 -> 4NO + 6H2O + 120 kJ). The nitric oxide is further oxidised by more air to nitrogen dioxide (2NO + O2 -> 2NO2), which is then passed through a series of absorption towers where it reacts with water to give nitric acid, with unreacted nitric oxide by-product being recycled/bled off by blowing air through (3NO2 + H2O -> 2HNO3 + NO).

Properties: Pure nitric acid is a colourless liquid boiling at 86 °C, completely miscible with water, forming a constant-boiling azeotropic mixture of 98% HNO3 that boils at 120.5 °C. Fuming nitric acid contains dissolved oxides of nitrogen. On standing or exposure to sunlight it decomposes (4HNO3 -> 4NO2 + 2H2O + O2), and the liberated NO2 gives the concentrated acid its characteristic yellow tinge on standing. In most of its reactions nitric acid behaves as a strong oxidising agent, in which its own nitrogen is reduced from the +5 state to some lower oxidation state; as a consequence it essentially never liberates hydrogen gas when reacting with metals (unlike a typical non-oxidising acid). Nitric acid can act simultaneously as an acid, as an oxidising agent, and as a nitrating agent. As an ordinary acid it reacts with basic oxides and bases to give salts and water, e.g. ZnO + 2HNO3 -> Zn(NO3)2 + H2O, and 3FeO + 10HNO3 -> 3Fe(NO3)3 + NO + 5H2O (note this second example shows the acid simultaneously behaving as an oxidant, converting Fe2+ to Fe3+). As an oxidising agent, non-metals such as carbon, sulphur, phosphorus and iodine are oxidised by hot concentrated nitric acid: C + 4HNO3 -> CO2 + 2H2O + 4NO2; S + 2HNO3 -> H2SO4 + 2NO; P4 + 20HNO3 -> 4H3PO4 + 20NO2 + 4H2O; 3I2 + 10HNO3 -> 6HIO3 + 10NO + 2H2O; it can also oxidise hydrogen sulphide to sulphur (3H2S + 2HNO3 -> 3S + 2NO + 4H2O) and react with fluorine (HNO3 + F2 -> HF + NO3F). As a nitrating agent, in the presence of concentrated sulphuric acid nitric acid replaces an aromatic -H with a -NO2 group (nitration), as in the conversion of benzene to nitrobenzene: C6H6 + HNO3 --H2SO4--> C6H5NO2 + H2O; this proceeds through generation of the electrophilic nitronium ion, NO2+ (HNO3 + 2H2SO4 -> NO2+ + H3O+ + 2HSO4-). …