Oxygen compounds

The most usual encountered oxides, oxoanions and oxocations, are displayed in Fig. 2. All these species contain some multiple bonding, the single N-N and N-O bonds being relatively weak. Nitrous oxide N₂O can be created by heating ammonium nitrate. It is isoelectronic with CO₂ and rather unreactive, and is employed as an anaesthetic ('laughing gas') and like a propellant for aerosols. Nitrogen dioxide NO₂ and Nitric oxide NO are the normal results of reaction of nitrogen and oxygen at high temperatures, or of the oxidation of ammonia. They are both oddelectron molecules. NO₂ dimerizes reversibly at low temperatures to create N₂O₄, but NO has extremely little tendency to dimerize in the gas phase, possibly since the odd electron is delocalized in a π antibonding orbital. NO reacts with oxygen to provide NO₂. It can act like a ligand in transition metal complexes. Another oxide of nitrogen is less stable: N₂O₃ is displayed in Fig. 2. N₂O₅ is generally found as [NO₂]⁺[NO₃]^-; and NO₃ is an unstable radical that (such as NO and NO₂) plays a role within atmospheric chemistry.

NO and (isoelectronic with CO and CO₂, correspondingly) can be formed through the action of strong oxidizing agents on NO or NO₂ in acid solvents like H₂SO₄, and are termed as solid salts (example NO⁺[AsF₆]^-). The nitrite and nitrate ions NO₂^- and NO₃^- are created correspondingly from nitrous acid HNO₂ and nitric acid HNO₃. As supposed from Pauling's rules, HNO₂ is a weak acid in water and HNO₃ a strong acid. Metal nitrates and nitrites are strong oxidizing agents, usually extremely soluble in water. Another less stable oxoacids are known, mainly containing N -N bonds. Even though the free acid subsequent to phosphoric acid H₃PO₄ is unknown, it is feasible to make orthonitrates consisting of the tetrahedral NH^3-₄ ion. Nitric acid is a main industrial chemical created from ammonia by catalytic oxidation to NO₂, followed through reaction with water and more oxygen:

It is employed to make NH₄NO₃ fertilizer, and in several industrial processes.

Fig. 2. Structures of some oxides, oxocations and oxoanions of nitrogen.

The nitrogen compounds' redox chemistry in aqueous solution is demonstrated in the Frost picture in Fig. 1. All oxoacids and oxides are strong oxidizing agents and all oxidation states apart from -3, 0 and +5 are susceptible to disproportionation. The comprehensive reactions are, though, mostly controlled by kinetic before thermodynamic considerations. In conjunction with oxidizable groups, like in ammonium nitrate NH₄NO₃ or in organic nitro compounds, N-O compounds might be powerful explosives.