Nitric acid | Properties, Formula, Uses, & Facts (2024)

nitric acid, (HNO₃), colourless, fuming, and highly corrosive liquid (freezing point −42 °C [−44 °F], boiling point 83 °C [181 °F]) that is a common laboratory reagent and an important industrial chemical for the manufacture of fertilizers and explosives. It is toxic and can cause severe burns.

The preparation and use of nitric acid were known to the early alchemists. A common laboratory process used for many years, ascribed to a German chemist, Johann Rudolf Glauber (1648), consisted of heating potassium nitrate with concentrated sulfuric acid. In 1776 Antoine-Laurent Lavoisier showed that it contained oxygen, and in 1816 Joseph-Louis Gay-Lussac and Claude-Louis Berthollet established its chemical composition.

The principal method of manufacture of nitric acid is the catalytic oxidation of ammonia. In the method developed by the German chemist Wilhelm Ostwald in 1901, ammonia gas is successively oxidized to nitric oxide and nitrogen dioxide by air or oxygen in the presence of a platinum gauze catalyst. The nitrogen dioxide is absorbed in water to form nitric acid. The resulting acid-in-water solution (about 50–70 percent by weight acid) can be dehydrated by distillation with sulfuric acid.

Nitric acid decomposes into water, nitrogen dioxide, and oxygen, forming a brownish yellow solution. It is a strong acid, completely ionized into hydronium (H₃O⁺) and nitrate (NO₃⁻) ions in aqueous solution, and a powerful oxidizing agent (one that acts as electron acceptor in oxidation-reduction reactions). Among the many important reactions of nitric acid are: neutralization with ammonia to form ammonium nitrate, a major component of fertilizers; nitration of glycerol and toluene, forming the explosives nitroglycerin and trinitrotoluene (TNT), respectively; preparation of nitrocellulose; and oxidation of metals to the corresponding oxides or nitrates.