Question

Nitrogen forms \(\mathrm{N}_{2}\), but phosphorous is converted into \(\mathrm{P}_{4}\) from \(\mathrm{P}\), the reason is (a) triple bond is present between phosphorous atom (b) \(\mathrm{p} \pi-\mathrm{p} \pi\) bonding is strong (c) p \(\pi-\mathrm{p} \pi\) bonding is weak (d) multiple bond is formed easily

Short answer

The reason is that p π-p π bonding is weak (option c).

Step-by-step solution

Understanding Nitrogen's Structure

Nitrogen molecules ( N_{2}) possess a triple bond between the nitrogen atoms. This triple bond is very strong due to significant  p - p  overlapping, which results in a stable diatomic molecule. Nitrogen makes effective use of  p - p  multiple bonding.

Analyzing Phosphorus Structure

Phosphorus forms  P_{4}  tetrahedra instead of   P_2  because the   p - p  bonding in phosphorus is not as effective at forming stable multiple bonds as observed in nitrogen. The propensity for   p - p   bonding is considerably weaker due to larger atomic size and less effective orbital overlap.

Final Evaluation of Options

Based on the analysis, phosphorus forms   P_4   because the   p - p  multiple bonding is weak. As a result, simple covalent multiple bonding does not stabilize phosphorus, unlike in nitrogen. Therefore, the correct answer aligns with option (c).

Key concepts

Triple Bond in Nitrogen

When understanding why nitrogen forms \( \mathrm{N}_{2} \) molecules, it's crucial to explore the nature of the triple bond present between two nitrogen atoms. The triple bond in nitrogen is exceptionally robust due to the presence of one sigma (σ) bond and two pi (π) bonds. This combination makes the bond not only strong but also very stable.

This stability is a direct result of effective overlapping. Nitrogen atoms participate in \( p \text{-} p \) overlapping, which leads to the formation of both a sigma bond and two pi bonds. These pi bonds are strong because the orbitals overlap significantly, ensuring that electrons are shared efficiently.

The lower energy and high stability of this triple bond give nitrogen its diatomic nature. Nitrogen molecules become less reactive due to this bonding, which is why \( \mathrm{N}_{2} \) is a very inert gas under standard conditions.

p-p Overlapping

The concept of \( p \text{-} p \) overlapping is vital in understanding molecular bonding, particularly in non-metals like nitrogen and phosphorus. Overlapping occurs when atomic orbitals from different atoms come into close proximity, allowing for shared electrons.

In nitrogen, as mentioned, \( p \text{-} p \) overlapping is particularly effective. This efficiency is due to the smaller atomic size and less shielding effect, which allows the p orbitals to align properly and overlap significantly.

However, in phosphorus, the \( p \text{-} p \) overlapping is not as pronounced. Phosphorus atoms are larger and their orbitals are more diffused compared to nitrogen. This diffusion prevents effective overlapping, resulting in weaker bonding when it comes to forming multiple bonds. This is why phosphorus prefers other types of molecular structures.

Molecular Structure of Phosphorus

In contrast to nitrogen, phosphorus shows a distinct molecular structure by forming \( \mathrm{P}_{4} \) tetrahedra rather than a diatomic \( \mathrm{P}_{2} \). This difference lies in the nature and strength of phosphorus's molecular bonding. The inability to form strong \( p \text{-} p \) pi bonds in phosphorus means that alternative structures are more stable.

The formation of \( \mathrm{P}_{4} \) involves tetrahedral bonding where each phosphorus atom is bonded to three other phosphorus atoms. This creates a more stable molecular structure due to the ample use of sigma bonds, which are essential for stability in the absence of pi bonding.

The tetrahedral \( \mathrm{P}_{4} \) has lower energy compared to any potential diatomic structure, making it the preferred form of phosphorus at room temperature. This arrangement is a direct consequence of weaker \( p \text{-} p \) interactions and larger atomic sizes, dictating a preference for different bonding schemes.