In a brown glass bottle, a colourless liquid can quietly turn yellow, and then red, as nitrogen dioxide creeps into it. That is nitric acid, HNO3, a highly corrosive mineral acid that most shops sell at about 68% in water, the azeotrope with a boiling point of 120.5 °C at 1 atm. When the strength climbs above 86%, it earns the name fuming nitric acid, and above 95% it is white fuming nitric acid, while red fuming acid holds more dissolved nitrogen dioxide. The chemistry looks simple, but the substance is never still.

Its earliest trail runs back into 13th-century European alchemy, and the usual starting point is pseudo-Geber’s De inventione veritatis, written after about 1300. Yet the record reaches further, into Arabic works tied to Jabir ibn Hayyan and the Fatimid caliph al-Hakim bi-Amr Allah. In the Ṣundūq al-ḥikma, the recipe calls for five parts of pure flowers of nitre, three parts of Cyprus vitriol, and two parts of Yemen alum, all powdered and heated in a flask with a glass receiver. What flowed out was described as an oil like cow’s butter, the first glimpse of a mineral acid drawn from fire and salts.

After 1300, nitric acid appears again in writings falsely attributed to Albert the Great and Ramon Llull, where a mixture of niter and green vitriol is distilled and named aqua fortis, strong water. That old name mattered, because the acid did not merely bite metals, it opened a path to new craft. The same spirit of experiment linked the alchemists to practical distillation, and the substance became part of a language of vitriol, nitre, and glassware long before anyone spoke of formulae or molarities. In those workshops, the acid was not yet a commodity, only a hard-won liquid with a sharp purpose.

In the 17th century, Johann Rudolf Glauber found a workable route by distilling potassium nitrate with sulfuric acid. Then the science sharpened. In 1776 Antoine Lavoisier, citing Joseph Priestley, showed that nitric acid could be made from nitric oxide, which he called nitrous air, combined with about an equal volume of the purest part of common air and a considerable quantity of water. By 1785 Henry Cavendish had fixed its composition and proved it could be synthesised by electric sparks through moist air. In 1806 Humphry Davy pushed further, using a high-voltage battery, gold electrode cones, and damp asbestos to show that nitric acid formed at the anode from dissolved atmospheric nitrogen gas.

Industrial scale arrived in 1905 with the Birkeland–Eyde process, also called the arc process. Air was blasted through a very hot electric arc, reaching about 3,000 °C, where atmospheric nitrogen and oxygen combined to make nitric oxide, but only in modest yield, roughly 4 to 5%. The gas was cooled, oxidised to nitrogen dioxide, and absorbed in water through packed or plate towers. Early towers bubbled the gas through water and quartz fragments, while the final towers used alkali to catch the last 20% of unreacted oxides. It worked, but it drank power greedily, and once cheap ammonia became available it was overtaken.

The decisive rival was the Ostwald process, made practical once the Haber process began supplying cheap ammonia in 1913. Here anhydrous ammonia burns to nitric oxide, then oxidises to nitrogen dioxide, which disproportionates in water to nitric acid and more nitric oxide, ready to be recycled. The chemistry is elegant, but the industrial effect was larger still. From that point, nitric acid production from ammonia became the dominant method and remains so today. Commercial grades now usually sit between 52% and 68% by mass, while further dehydration can reach 98%, and the last plant in the United States that made stronger acid by dissolving extra nitrogen dioxide shut that route down in 2012.

Around 1913, French engineer Albert Nodon proposed another production method, and it was as strange as it sounds. Calcium nitrate, first converted by bacteria from nitrogenous matter in peat bogs, was electrolysed in a pit dug into the peat and lined with tarred timber stakes. A porous earthenware vessel sat inside crushed limestone, filled with coke around a graphite anode. Nitric acid was pumped out through a glass tube almost to the bottom, while fresh water entered from above, and cast iron cathodes sat in the peat around it. Resistance was about 3 ohms per cubic metre, the supply about 10 volts, and a one-hectare bog six and a half feet deep was estimated to yield more than 600 tons a year.

At room temperature, the ordinary acid is a colourless liquid, but the pure substance is difficult to hold without change. Nitric acid and water form a 68% azeotrope, known as concentrated nitric acid, and it boils at 120.5 °C. Two solid hydrates are known, the monohydrate HNO3·H2O, also written as oxonium nitrate [H3O]+[NO3]−, and the trihydrate HNO3·3H2O. The anhydrous acid is much closer to white fuming nitric acid, available at 99.9% assay, with a density around 1.512 to 1.3 g/cm3 and a boiling point of 83 °C. It solidifies at −42 °C into white crystals and must be stored in shatterproof amber glass, with extra head space and monthly venting.

Chemically, nitric acid is a strong acid at ambient temperatures, with a pKa usually reported below −1, though the figure rises to 1 at 250 °C. It can also act as a base towards sulfuric acid, and in that partnership it produces the nitronium ion, [NO2]+, the active agent in aromatic nitration. The acid can even undergo autoprotolysis, much as water does. This dual nature is why it serves so well in synthesis: it gives up protons readily, yet in concentrated mixtures it also helps build the species that strikes aromatic rings and installs the nitro group that so many industries rely upon.