oxoacids of nitrogen (nitric acid)

When you combine nitrogen with oxygen and a proton you get the nitrogen oxoacids — acids built around a central nitrogen atom surrounded by oxygens, with an O-H group that can let a proton go. The two that matter most are nitric acid, HNO3 (nitrogen in the +5 state), and the weaker, more fragile nitrous acid, HNO2 (nitrogen in the +3 state). Nitric acid in particular is one of the great industrial acids and a staple of every chemistry laboratory.

Nitric acid is made industrially by the Ostwald process: ammonia is burned over a platinum catalyst to nitric oxide (NO), which is oxidised in air to nitrogen dioxide (NO2), which then reacts with water to give HNO3. Structurally the nitrate ion, NO3-, is a flat triangle with the negative charge spread evenly over all three oxygens by resonance, which is part of why nitrate salts are so stable. Nitric acid is a strong acid (fully ionised in water) and, just as importantly, a powerful oxidising agent: it attacks many metals not by the simple acid route of releasing hydrogen gas, but by the nitrogen itself being reduced (to NO with dilute acid, to NO2 with concentrated acid). Nitrous acid, by contrast, is a weak acid, exists only in cold dilute solution, and readily disproportionates — meaning the same nitrogen splits into both a higher and a lower oxidation state at once.

These oxoacids are everywhere in real chemistry and life. Nitric acid makes nitrate fertilisers, explosives (TNT, nitroglycerine), and dyes, and concentrated nitric mixed with hydrochloric gives aqua regia, the only common mixture that dissolves gold. Nitrous acid and its nitrite salts are used in curing meats and are central to diazonium chemistry in dye-making. The recurring theme is nitrogen's wide range of oxidation states — from -3 in ammonia up to +5 in nitric acid — which lets it shuttle through a rich and sometimes hazardous redox chemistry.