Dinitrogen and dioxygen the main constituents of air do not react with each other in atmosphere to form oxides of nitrogen because

We need to explain why dinitrogen ($$N_2$$) and dioxygen ($$O_2$$), despite being the main constituents of the atmosphere, do not react with each other under normal atmospheric conditions to form oxides of nitrogen.

The reaction between $$N_2$$ and $$O_2$$ to form nitric oxide is:

$$N_2(g) + O_2(g) \rightarrow 2NO(g)$$

This reaction is highly endothermic ($$\Delta H \approx +180 \text{ kJ/mol}$$). The reason is that the $$N \equiv N$$ triple bond is extremely strong (bond dissociation energy $$\approx 941 \text{ kJ/mol}$$), and breaking it requires a very large amount of energy. Similarly, the $$O = O$$ bond energy is about 498 kJ/mol. The energy released in forming the N-O bonds in NO is not sufficient to compensate for the energy required to break both the $$N \equiv N$$ and $$O = O$$ bonds.

Because the reaction is endothermic and has a very high activation energy, it requires extremely high temperatures (above 2000°C) to proceed at a significant rate. Such high temperatures occur in lightning strikes and in internal combustion engines, which is why small amounts of NO are produced under those conditions, but not under normal atmospheric conditions.

Let us examine the other options:

Option A states that $$N_2$$ is unreactive under atmospheric conditions — while $$N_2$$ is relatively inert, the fundamental reason is the thermodynamics of the reaction (endothermic nature), not just the inertness of nitrogen alone.

Option B states that oxides of nitrogen are unstable — this is not correct in general; $$NO_2$$ and other nitrogen oxides do exist stably.

Option C states the reaction can occur with a catalyst — while catalysts can lower activation energy, the key reason they don't react is the high temperature requirement due to the endothermic nature.

Hence, the correct answer is Option D.