Question

Nitrogen in the atmosphere exists as very stable diatomic molecules. Why does phosphorus form less stable \({{\bf{P}}_4}\)molecules instead of \({{\bf{P}}_2}\)molecules?

Short answer

Phosphorus atoms have an empty d orbital, which makes it a larger atom and unable to form multiple bonds with other phosphorous atoms.

Step-by-step solution

Electronic configuration

The distribution of electrons in an element's orbitals is described by its electron configuration. They are represented by a standard memorandum that arranges all electron-containing subshells in a logical order.

Reasons why phosphorus forms a less stable\({{\bf{P}}_4}\) molecules

• Due to the electronic configuration, phosphorus atom ha an empty d orbital, which makes it a larger atom and unable to form multiple bonds with other phosphorous atoms.

• Nitrogen atom doesn't have a d orbital and as such can makepibonds, which can form multiple bonds with the same element.

• Nitrogen atom has the ability to create \(\pi - \pi \)multiple bonds. Due to the repulsion between the non-bonded electrons in the inner core, phosphorus cannot form such bonds.

• There is no such repulsion in nitrogen atom since it possesses only \(1s\) electron in its inner shell, making the overlap of p orbitals to create pi-bonds simply.