Nitrogen compounds

These complexes, in which a nitrogen molecule donates at least one lone pair of electrons to a central metal cation, illustrate how N2 might bind to the metal(s) in nitrogenase and the catalyst for the Haber process: these processes involving dinitrogen activation are vitally important in biology and in the production of fertilisers.

The more well-characterised ways are the end-on M←N≡N (η1) and M←N≡N→M (μ, bis-η1), in which the lone pairs on the nitrogen atoms are donated to the metal cation.

The less well-characterised ways involve dinitrogen donating electron pairs from the triple bond, either as a bridging ligand to two metal cations (μ, bis-η2) or to just one (η2).

The fifth and unique method involves triple-coordination as a bridging ligand, donating all three electron pairs from the triple bond (μ3-N2).

Since N2 is isoelectronic with carbon monoxide (CO) and acetylene (C2H2), the bonding in dinitrogen complexes is closely allied to that in carbonyl compounds, although N2 is a weaker σ-donor and π-acceptor than CO.

[3] Many binary compounds are known: with the exception of the nitrogen hydrides, oxides, and fluorides, these are typically called nitrides.

These can formally be thought of as salts of the N3− anion, although charge separation is not actually complete even for these highly electropositive elements.

Nitrido complexes are generally made by thermal decomposition of azides or by deprotonating ammonia, and they usually involve a terminal {≡N}3− group.

The linear azide anion (N−3), being isoelectronic with nitrous oxide, carbon dioxide, and cyanate, forms many coordination complexes.

The presence of hydrogen bonding has very significant effects on ammonia, conferring on it its high melting (−78 °C) and boiling (−33 °C) points.

However, the hydrogen bonding in NH3 is weaker than that in H2O due to the lower electronegativity of nitrogen compared to oxygen and the presence of only one lone pair in NH3 rather than two in H2O.

The reason for adding gelatin is that it removes metal ions such as Cu2+ that catalyses the destruction of hydrazine by reaction with monochloramine (NH2Cl) to produce ammonium chloride and nitrogen.

Only when heated does it act as a fluorinating agent, and it reacts with copper, arsenic, antimony, and bismuth on contact at high temperatures to give tetrafluorohydrazine (N2F4).

[9] Nitrogen trichloride (NCl3) is a dense, volatile, and explosive liquid whose physical properties are similar to those of carbon tetrachloride, although one difference is that NCl3 is easily hydrolysed by water while CCl4 is not.

Its adduct with ammonia, which was known earlier, is very shock-sensitive: it can be set off by the touch of a feather, shifting air currents, or even alpha particles.

[9][10] For this reason, small amounts of nitrogen triiodide are sometimes synthesised as a demonstration to high school chemistry students or as an act of "chemical magic".

It is rather unreactive (not reacting with the halogens, the alkali metals, or ozone at room temperature, although reactivity increases upon heating) and has the unsymmetrical structure N–N–O (N≡N+O−↔−N=N+=O): above 600 °C it dissociates by breaking the weaker N–O bond.

In mammals, including humans, it is an important cellular signaling molecule involved in many physiological and pathological processes.

The latter two compounds are somewhat difficult to study individually because of the equilibrium between them, although sometimes dinitrogen tetroxide can react by heterolytic fission to nitrosonium and nitrate in a medium with high dielectric constant.

Dinitrogen tetroxide is very useful for the preparation of anhydrous metal nitrates and nitrato complexes, and it became the storable oxidiser of choice for many rockets in both the United States and USSR by the late 1950s.

This is because it is a hypergolic propellant in combination with a hydrazine-based rocket fuel and can be easily stored since it is liquid at room temperature.

Gaseous dinitrogen pentoxide decomposes as follows:[15] Many nitrogen oxoacids are known, though most of them are unstable as pure compounds and are known only as aqueous solution or as salts.

Acidic solutions are quite stable but above pH 4 base-catalysed decomposition occurs via [HONNO]− to nitrous oxide and the hydroxide anion.

Sodium nitrite is mildly toxic in concentrations above 100 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 diazotise primary aromatic amines as follows:[19] Nitrite is also a common ligand that can coordinate in five ways.

In the 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.

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:[19] The thermal stabilities of nitrates (involving the trigonal planar 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.

[19] Finally, although orthonitric acid (H3NO4), which would be analogous to orthophosphoric acid, does not exist, the tetrahedral orthonitrate anion NO3−4 is known in its sodium and potassium salts:[19] These white crystalline salts are very sensitive to water vapour and carbon dioxide in the air:[19] 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.

This may be offset by other factors: for example, amides are not basic because the lone pair is delocalised into a double bond (though they may act as acids at very low pH, being protonated at the oxygen), and pyrrole is not acidic because the lone pair is delocalised as part of an aromatic ring.