inertness of dinitrogen

About 78% of the air you are breathing right now is nitrogen gas, N2, and yet it does almost nothing. You can stand in a room full of it forever and it will not react with you, with the oxygen beside it, or with most things you could throw at it. This is strange: nitrogen as an element is essential to all life (it is in every protein and every strand of DNA), and yet the form in which it is most abundant — the air itself — is one of the least reactive substances we know.

The reason is the bond holding the two nitrogen atoms together. In N2 the atoms are joined by a triple bond, N≡N, made of one sigma and two pi bonds, and it is one of the strongest bonds in all of chemistry, with a bond dissociation energy of about 945 kilojoules per mole. To make N2 react, that triple bond usually has to be broken, and the first step of breaking it (going from a triple bond to a double bond) is especially costly. There is no easy partial reaction; the molecule sits in a deep energy valley with very high walls. On top of that N2 is nonpolar, has no easily attacked lone-pair geometry, and its filled molecular orbitals leave nothing obvious for a reagent to grab. So the inertness is kinetic (reactions are slow because the barrier is huge) as much as it is about overall thermodynamics.

This inertness is the whole problem of nitrogen chemistry. Plants cannot use N2 directly — they need nitrogen in a reactive form such as ammonia or nitrate. Breaking the triple bond at scale is what the Haber-Bosch process does industrially (with high temperature, high pressure, and an iron catalyst) and what the enzyme nitrogenase does in certain bacteria at ordinary temperature, which chemists still find astonishing. The flip side of inertness is useful: N2 is cheap, abundant, and so unreactive that it is used as a protective blanket gas to keep oxygen and moisture away from sensitive materials and reactions.