Editing Nitrogen (section)

===Oxoacids, oxoanions, and oxoacid salts===
Many nitrogen [[oxoacid]]s are known, though most of them are unstable as pure compounds and are known only as aqueous solutions or as salts. [[Hyponitrous acid]] (H<sub>2</sub>N<sub>2</sub>O<sub>2</sub>) is a weak diprotic acid with the structure HON=NOH (p''K''<sub>a1</sub> 6.9, p''K''<sub>a2</sub> 11.6). Acidic solutions are quite stable but above pH 4 base-catalysed decomposition occurs via [HONNO]<sup>−</sup> to nitrous oxide and the hydroxide anion. [[Hyponitrite]]s (involving the {{chem|N|2|O|2|2-}} anion) are stable to reducing agents and more commonly act as reducing agents themselves. They are an intermediate step in the oxidation of ammonia to nitrite, which occurs in the [[nitrogen cycle]]. Hyponitrite can act as a bridging or chelating bidentate ligand.<ref name="Greenwood459">Greenwood and Earnshaw, pp. 459–72</ref>

[[Nitrous acid]] (HNO<sub>2</sub>) is not known as a pure compound, but is a common component in gaseous equilibria and is an important aqueous reagent: its aqueous solutions may be made from acidifying cool aqueous [[nitrite]] ({{chem|NO|2|-}}, bent) solutions, although already at room temperature disproportionation to [[nitrate]] and nitric oxide is significant. It is a weak acid with p''K''<sub>''a''</sub> 3.35 at 18&nbsp;°C. They may be [[titration|titrimetrically]] analysed by their oxidation to nitrate by [[permanganate]]. They are readily reduced to nitrous oxide and nitric oxide by [[sulfur dioxide]], to hyponitrous acid with [[tin]](II), and to ammonia with [[hydrogen sulfide]]. Salts of [[hydrazinium]] {{chem|N|2|H|5|+}} react with nitrous acid to produce azides which further react to give nitrous oxide and nitrogen. [[Sodium nitrite]] is mildly toxic in concentrations above 100&nbsp;mg/kg, but small amounts are often used to cure meat and as a preservative to avoid bacterial spoilage. It is also used to synthesise hydroxylamine and to [[Diazonium compound#Preparation|diazotise]] primary aromatic amines as follows:<ref name="Greenwood459" />
:ArNH<sub>2</sub> + HNO<sub>2</sub> → [ArNN]Cl + 2 H<sub>2</sub>O

Nitrite is also a common ligand that can coordinate in five ways. The most common are nitro (bonded from the nitrogen) and nitrito (bonded from an oxygen). Nitro-nitrito isomerism is common, where the nitrito form is usually less stable.<ref name="Greenwood459" />

[[File:Fuming nitric acid 40ml.jpg|thumb|right|Fuming nitric acid contaminated with yellow nitrogen dioxide]]
[[Nitric acid]] (HNO<sub>3</sub>) is by far the most important and the most stable of the nitrogen oxoacids. It is one of the three most used acids (the other two being [[sulfuric acid]] and [[hydrochloric acid]]) and was first discovered by alchemists in the 13th century. It is made by the catalytic oxidation of ammonia to nitric oxide, which is oxidised to nitrogen dioxide, and then dissolved in water to give concentrated nitric acid. In the [[United States|United States of America]], over seven million tonnes of nitric acid are produced every year, most of which is used for nitrate production for fertilisers and explosives, among other uses. Anhydrous nitric acid may be made by distilling concentrated nitric acid with phosphorus pentoxide at low pressure in glass apparatus in the dark. It can only be made in the solid state, because upon melting it spontaneously decomposes to nitrogen dioxide, and liquid nitric acid undergoes [[Molecular autoionization|self-ionisation]] to a larger extent than any other covalent liquid as follows:<ref name="Greenwood459" />
:2 HNO<sub>3</sub> {{eqm}} {{chem|H|2|NO|3|+}} + {{chem|NO|3|-}} {{eqm}} H<sub>2</sub>O + [NO<sub>2</sub>]<sup>+</sup> + [NO<sub>3</sub>]<sup>−</sup>
Two hydrates, HNO<sub>3</sub>·H<sub>2</sub>O and HNO<sub>3</sub>·3H<sub>2</sub>O, are known that can be crystallised. It is a strong acid and concentrated solutions are strong oxidising agents, though [[gold]], [[platinum]], [[rhodium]], and [[iridium]] are immune to attack. A 3:1 mixture of concentrated hydrochloric acid and nitric acid, called ''[[aqua regia]]'', is still stronger and successfully dissolves gold and platinum, because free chlorine and nitrosyl chloride are formed and chloride anions can form strong complexes. In concentrated sulfuric acid, nitric acid is protonated to form [[nitronium]], which can act as an electrophile for aromatic nitration:<ref name="Greenwood459" />
:HNO<sub>3</sub> + 2 H<sub>2</sub>SO<sub>4</sub> {{eqm}} {{chem|NO|2|+}} + H<sub>3</sub>O<sup>+</sup> + 2 {{chem|HSO|4|-}}
The thermal stabilities of [[nitrate]]s (involving the trigonal planar {{chem|NO|3|-}} anion) depends on the basicity of the metal, and so do the products of decomposition (thermolysis), which can vary between the nitrite (for example, sodium), the oxide (potassium and [[lead]]), or even the metal itself ([[silver]]) depending on their relative stabilities. Nitrate is also a common ligand with many modes of coordination.<ref name="Greenwood459" />

Finally, although orthonitric acid (H<sub>3</sub>NO<sub>4</sub>), which would be analogous to [[orthophosphoric acid]], does not exist, the tetrahedral [[orthonitrate]] anion {{chem|NO|4|3-}} is known in its sodium and potassium salts:<ref name="Greenwood459" />
:<chem>NaNO3{} + Na2O ->[\ce{Ag~crucible}][\ce{300^\circ C~for~7 days}] Na3NO4</chem>
These white crystalline salts are very sensitive to water vapour and carbon dioxide in the air:<ref name="Greenwood459" />
:Na<sub>3</sub>NO<sub>4</sub> + H<sub>2</sub>O + CO<sub>2</sub> → NaNO<sub>3</sub> + NaOH + NaHCO<sub>3</sub>
Despite its limited chemistry, the orthonitrate anion is interesting from a structural point of view due to its regular tetrahedral shape and the short N–O bond lengths, implying significant polar character to the bonding.<ref name="Greenwood459" />