Question
Why does nitrogen show anomalous behaviour compared to other Group 15 elements? Explain why N₂ is relatively inert while phosphorus is highly reactive.
(NCERT Class 11, Chapter 7 — The p-Block Elements)
Solution — Step by Step
Nitrogen is the smallest element in Group 15 with the highest electronegativity (3.0). This small size leads to:
- Very high charge density
- Strong tendency to form multiple bonds (p-p bonding)
- Absence of d-orbitals in the valence shell (only 2s and 2p available)
Nitrogen forms a triple bond () with a bond dissociation energy of 944 kJ/mol — one of the strongest bonds in chemistry.
This triple bond consists of one sigma and two pi bonds, formed by efficient p-p lateral overlap between the small 2p orbitals.
Breaking this bond requires enormous energy, making N₂ kinetically inert at room temperature.
Phosphorus has larger 3p orbitals. The lateral overlap between 3p orbitals is too weak to form effective pi bonds. So instead of forming , phosphorus forms single bonds and exists as (tetrahedral structure with P-P single bonds).
P-P single bond energy is only about 200 kJ/mol — much easier to break than the N≡N triple bond. This is why phosphorus is reactive and catches fire in air, while N₂ just sits there.
- Maximum covalency = 4 (no d-orbitals to expand octet), while P can show covalency up to 5 (e.g., PCl₅)
- Forms strong hydrogen bonds (N-H…N), which P cannot form effectively
- Does not form d-p bonds with oxygen, while P, As do
- Exists as diatomic gas (), while P exists as solid
Why This Works
The root cause of nitrogen’s anomalous behaviour is its small atomic size and absence of d-orbitals. Small size allows efficient 2p-2p overlap (hence strong triple bond in N₂), but it also means nitrogen cannot accommodate more than 4 bonds (maximum covalency = 4).
The general trend in Group 15: as we go down, the ability to form p-p bonds decreases (larger orbitals overlap poorly), and the ability to expand the octet increases (d-orbitals become available). Nitrogen sits at the extreme — best at pi bonding, worst at expanding its octet.
Alternative Method — Comparison Table
| Property | N | P |
|---|---|---|
| Molecular form | N₂ (gas) | P₄ (solid) |
| Bond type | N≡N (triple) | P-P (single) |
| Bond energy | 944 kJ/mol | ~200 kJ/mol per P-P |
| Reactivity | Inert at RT | Catches fire in air |
| Max covalency | 4 | 5 (e.g., PCl₅) |
| d-orbitals | Absent | Available |
NEET often asks: “Why can’t nitrogen form NCl₅ while phosphorus forms PCl₅?” The answer is simple — nitrogen has no d-orbitals in its valence shell, so it cannot expand beyond an octet. Its maximum covalency is 4 (as in NH₄⁺). This one-liner is enough for the MCQ.
Common Mistake
Students often say “N₂ is inert because it has a triple bond” without explaining why nitrogen forms a triple bond in the first place. The deeper reason is efficient p-p overlap due to the small size of 2p orbitals. If the question asks “why is N₂ inert,” start with the small size → effective pi overlap → strong triple bond → high bond energy → inert. This chain of reasoning is what examiners look for.