Nitrogen compounds

The chemical element nitrogen is one of the most abundant elements in the universe and can form many compounds. It can take several oxidation states; but the most common oxidation states are −3 and +3. Nitrogen can form nitride and nitrate ions. It also forms a part of nitric acid and nitrate salts. Nitrogen compounds also have an important role in organic chemistry, as nitrogen is part of proteins, amino acids and adenosine triphosphate.

Dinitrogen complexes

The first example of a dinitrogen complex to be discovered was ²⁺, and soon many other such complexes were discovered. 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 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.
Dinitrogen is able to coordinate to metals in five different ways. The more well-characterised ways are the end-on M←N≡N and M←N≡N→M, 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 or to just one. The fifth and unique method involves triple-coordination as a bridging ligand, donating all three electron pairs from the triple bond. A few complexes feature multiple N₂ ligands and some feature N₂ bonded in multiple ways. Since N₂ is isoelectronic with carbon monoxide and acetylene, 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. Theoretical studies show that σ donation is a more important factor allowing the formation of the M–N bond than π back-donation, which mostly only weakens the N–N bond, and end-on donation is more readily accomplished than side-on donation.
Today, dinitrogen complexes are known for almost all the transition metals, accounting for several hundred compounds. They are normally prepared by three methods:

Replacing labile ligands such as H₂O, H⁻, or CO directly by nitrogen: these are often reversible reactions that proceed at mild conditions.
Reducing metal complexes in the presence of a suitable coligand in excess under nitrogen gas. A common choice include replacing chloride ligands by dimethylphenylphosphine to make up for the smaller number of nitrogen ligands attached than the original chlorine ligands.
Converting a ligand with N–N bonds, such as hydrazine or azide, directly into a dinitrogen ligand.

Occasionally the N≡N bond may be formed directly within a metal complex, for example by directly reacting coordinated ammonia with nitrous acid, but this is not generally applicable. Most dinitrogen complexes have colours within the range white-yellow-orange-red-brown; a few exceptions are known, such as the blue .

Nitrides, azides, and nitrido complexes

Nitrogen bonds to almost all the elements in the periodic table except the first three noble gases, helium, neon, and argon, and some of the very short-lived elements after bismuth, creating an immense variety of binary compounds with varying properties and applications. Many binary compounds are known: with the exception of the nitrogen hydrides, oxides, and fluorides, these are typically called nitrides. Many stoichiometric phases are usually present for most elements. They may be classified as "salt-like", covalent, "diamond-like", and metallic, although this classification has limitations generally stemming from the continuity of bonding types instead of the discrete and separate types that it implies. They are normally prepared by directly reacting a metal with nitrogen or ammonia, or by thermal decomposition of metal amides:
Many variants on these processes are possible. The most ionic of these nitrides are those of the alkali metals and alkaline earth metals, Li₃N and M₃N₂. 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. However, the alkali metal azides NaN₃ and KN₃, featuring the linear anion, are well-known, as are Sr₂ and Ba₂. Azides of the B-subgroup metals are much less ionic, have more complicated structures, and detonate readily when shocked.
Many covalent binary nitrides are known. Examples include cyanogen, triphosphorus pentanitride, disulfur dinitride, and tetrasulfur tetranitride. The essentially covalent silicon nitride and germanium nitride are also known: silicon nitride in particular would make a promising ceramic if not for the difficulty of working with and sintering it. In particular, the group 13 nitrides, most of which are promising semiconductors, are isoelectronic with graphite, diamond, and silicon carbide and have similar structures: their bonding changes from covalent to partially ionic to metallic as the group is descended. In particular, since the B–N unit is isoelectronic to C–C, and carbon is essentially intermediate in size between boron and nitrogen, much of organic chemistry finds an echo in boron–nitrogen chemistry, such as in borazine. Nevertheless, the analogy is not exact due to the ease of nucleophilic attack at boron due to its deficiency in electrons, which is not possible in a wholly carbon-containing ring.
The largest category of nitrides are the interstitial nitrides of formulae MN, M₂N, and M₄N, where the small nitrogen atoms are positioned in the gaps in a metallic cubic or hexagonal close-packed lattice. They are opaque, very hard, and chemically inert, melting only at very high temperatures. They have a metallic lustre and conduct electricity as do metals. They hydrolyse only very slowly to give ammonia or nitrogen.
The nitride anion is the strongest π donor known amongst ligands. Nitrido complexes are generally made by thermal decomposition of azides or by deprotonating ammonia, and they usually involve a terminal ³⁻ group. The linear azide anion, being isoelectronic with nitrous oxide, carbon dioxide, and cyanate, forms many coordination complexes. Further catenation is rare, although is known.

Hydrides

Industrially, ammonia is the most important compound of nitrogen and is prepared in larger amounts than any other compound, because it contributes significantly to the nutritional needs of terrestrial organisms by serving as a precursor to food and fertilisers. It is a colourless alkaline gas with a characteristic pungent smell. The presence of hydrogen bonding has very significant effects on ammonia, conferring on it its high melting and boiling points. As a liquid, it is a very good solvent with a high heat of vaporisation, that also has a low viscosity and electrical conductivity and high dielectric constant, and is less dense than water. 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. It is a weak base in aqueous solution ; its conjugate acid is ammonium,. It can also act as an extremely weak acid, losing a proton to produce the amide anion,. It thus undergoes self-dissociation, similar to water, to produce ammonium and amide. Ammonia burns in air or oxygen, though not readily, to produce nitrogen gas; it burns in fluorine with a greenish-yellow flame to give nitrogen trifluoride. Reactions with the other nonmetals are very complex and tend to lead to a mixture of products. Ammonia reacts on heating with metals to give nitrides.
Many other binary nitrogen hydrides are known, but the most important are hydrazine and hydrogen azide. Although it is not a nitrogen hydride, hydroxylamine is similar in properties and structure to ammonia and hydrazine as well. Hydrazine is a fuming, colourless liquid that smells similarly to ammonia. Its physical properties are very similar to those of water. Despite it being an endothermic compound, it is kinetically stable. It burns quickly and completely in air very exothermically to give nitrogen and water vapour. It is a very useful and versatile reducing agent and is a weaker base than ammonia. It is also commonly used as a rocket fuel.
Hydrazine is generally made by reaction of ammonia with alkaline sodium hypochlorite in the presence of gelatin or glue:
The reason for adding gelatin is that it removes metal ions such as Cu²⁺ that catalyses the destruction of hydrazine by reaction with monochloramine to produce ammonium chloride and nitrogen.
Hydrogen azide was first produced in 1890 by the oxidation of aqueous hydrazine by nitrous acid. It is very explosive and even dilute solutions can be dangerous. It has a disagreeable and irritating smell and is a potentially lethal poison. It may be considered the conjugate acid of the azide anion, and is similarly analogous to the hydrohalic acids.

Halides and oxohalides

All four simple nitrogen trihalides are known. A few mixed halides and hydrohalides are known, but are mostly unstable; examples include NClF₂, NCl₂F, NBrF₂, NF₂H, NFH₂, NCl₂H, and NClH₂.
Five nitrogen fluorides are known. Nitrogen trifluoride is a colourless and odourless gas that is thermodynamically stable, and most readily produced by the electrolysis of molten ammonium fluoride dissolved in anhydrous hydrogen fluoride. Like carbon tetrafluoride, it is not at all reactive and is stable in water or dilute aqueous acids or alkalis. 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. The cations and are also known, as is ONF₃, which has aroused interest due to the short N–O distance implying partial double bonding and the highly polar and long N–F bond. Tetrafluorohydrazine, unlike hydrazine itself, can dissociate at room temperature and above to give the radical NF₂•. Fluorine azide is very explosive and thermally unstable. Dinitrogen difluoride exists as thermally interconvertible cis and trans isomers, and was first found as a product of the thermal decomposition of FN₃.
Nitrogen trichloride 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. It was first synthesised in 1811 by Pierre Louis Dulong, who lost three fingers and an eye to its explosive tendencies. As a dilute gas it is less dangerous and is thus used industrially to bleach and sterilise flour. Nitrogen tribromide, first prepared in 1975, is a deep red, temperature-sensitive, volatile solid that is explosive even at −100 °C. Nitrogen triiodide is still more unstable and was only prepared in 1990. 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. 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". Chlorine azide and bromine azide are extremely sensitive and explosive.
Two series of nitrogen oxohalides are known: the nitrosyl halides and the nitryl halides. The first are very reactive gases that can be made by directly halogenating nitrous oxide. Nitrosyl fluoride is colourless and a vigorous fluorinating agent. Nitrosyl fluoride can continue reacts with fluorine to form nitrogen oxide trifluoride, which is also a strong fluorinating agent. Nitrosyl chloride behaves in much the same way and has often been used as an ionising solvent. Nitrosyl bromide is red. The reactions of the nitryl halides are mostly similar: nitryl fluoride and nitryl chloride are likewise reactive gases and vigorous halogenating agents.