Pie chart of U.S. nitrous

What is the formula of nitrogen oxide?

The chemistry of nitrogen is dominated by the ease with which nitrogen atoms form double and triple bonds. A neutral nitrogen atom contains five valence electrons: 2s2 2p3. A nitrogen atom can therefore achieve an octet of valence electrons by sharing three pairs of electrons with another nitrogen atom.

Because the covalent radius of a nitrogen atom is relatively small (only 0.070 nm), nitrogen atoms come close enough together to form very strong bonds. The bond-dissociation enthalpy for the nitrogen-nitrogen triple bond is 946 kJ/mol, almost twice as large as that for an O=O double bond.

The strength of the nitrogen-nitrogen triple bond makes the N2 molecule very unreactive. N2 is so inert that lithium is one of the few elements with which it reacts at room temperature.

6 Li(s) N2(g) 2 Li3N(s)

In spite of the fact that the N2 molecule is unreactive, compounds containing nitrogen exist for virtually every element in the periodic table except those in Group VIIIA (He, Ne, Ar, and so on). This can be explained in two ways. First, N2 becomes significantly more reactive as the temperature increases. At high temperatures, nitrogen reacts with hydrogen to form ammonia and with oxygen to form nitrogen oxide.

3 H2(g) 2 NH3(g)
O2(g) 2 NO(g)

Second, a number of catalysts found in nature overcome the inertness of N2 at low temperature.

It is difficult to imagine a living system that does not contain nitrogen, which is an essential component of the proteins, nucleic acids, vitamins, and hormones that make life possible. Animals pick up the nitrogen they need from the plants or other animals in their diet. Plants have to pick up their nitrogen from the soil, or absorb it as N2 from the atmosphere. The concentration of nitrogen in the soil is fairly small, so the process by which plants reduce N2 to NH3or "fix" N2 is extremely important.

Although 200 million tons of NH3 are produced by nitrogen fixation each year, plants, by themselves, cannot reduce N2 to NH3. This reaction is carried out by blue-green algae and bacteria that are associated with certain plants. The best-understood example of nitrogen fixation involves the rhizobium bacteria found in the root nodules of legumes such as clover, peas and beans. These bacteria contain a nitrogenase enzyme, which is capable of the remarkable feat of reducing N2 from the atmosphere to NH3 at room temperature.

Ammonia is made on an industrial scale by a process first developed between 1909 and 1913 by Fritz Haber. In the Haber process, a mixture of N2 and H2 gas at 200 to 300 atm and 400 to 600oC is passed over a catalyst of finely divided iron metal.

N2(g) 3 H2(g) 2 NH3(g)

Almost 20 million tons of NH3 are produced in the United States each year by this process. About 80% of it, worth more than $2 billion, is used to make fertilizers for plants that can't fix nitrogen from the atmosphere. On the basis of weight, ammonia is the second most important industrial chemical in the United States. (Only sulfuric acid is produced in larger quantities.)

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