Why does white phosphorus consist of tetrahedral \({P_4}\)molecules while nitrogen consists of diatomic \({N_2}\) molecules?

The “transfer/sharing” of electrons gives rise to bond formation. The bond is considered as strong if the electronegativity difference between two elements increases.

The electron configurations of phosphorus and nitrogen are as follows:

\(\begin{array}{l}{\rm{P = }}\left[ {{\rm{Ne}}} \right]{\rm{ 3s^{2}3p^{3}}}{\rm{.}}\\{\rm{N = }}\left[ {{\rm{He}}} \right]{\rm{ 2}}{{\rm{s}}^{\rm{2}}}{\rm{ 2}}{{\rm{p}}^{\rm{3}}}.\end{array}\)

• Due to their electronic configuration, phosphorus atom has an empty d orbital, which makes it a larger atom and unable to form multiple bonds with other phosphorous atoms.
• Nitrogen atom does not have a dorbital and as such can make\(\pi \)bonds, which can form multiple bonds with the same element.
• Due to the repulsion between non-bonded electrons in the inner core, phosphorus cannot form such\(\pi \) bonds.
• There is no such repulsion in nitrogen since it possesses only \(1s\) electron in its inner shell, making the overlap of p orbitals to create \(\pi \) bonds.