Anomalous Behaviour of Nitrogen
Why the first member of the Pnictogen family breaks all the rules.
In the p-block of the periodic table, the very first element of every group exhibits behavior that is drastically different from the rest of the heavier elements in its family. Nitrogen, the head of Group 15, is no exception. Its unique chemistry governs everything from the atmosphere we breathe to why certain explosive compounds exist.
1. Why does Nitrogen behave abnormally?
Whenever an examiner asks why the first element of a p-block group behaves differently, you must immediately state the "Big Four" reasons:
- Exceptionally small atomic size.
- Extremely high electronegativity.
- Very high ionization enthalpy.
- Complete absence of d-orbitals in its valence shell.
2. $p\pi - p\pi$ Multiple Bonding
Because Nitrogen is so small, its $p$-orbitals are highly concentrated. This allows them to overlap effectively side-by-side to form strong pi ($\pi$) bonds.
Nitrogen easily forms $p\pi - p\pi$ multiple bonds with itself. It exists as a diatomic molecule with a triple bond ($N \equiv N$). This triple bond is incredibly strong (bond enthalpy $\sim 941.4\text{ kJ mol}^{-1}$), making $N_2$ gas practically inert at room temperature.
Heavier elements like P, As, and Sb are too large. Their $p$-orbitals are diffuse, so they cannot effectively overlap side-by-side. Instead of $\pi$ bonds, they satisfy their valency by forming single bonds with other atoms, resulting in polyatomic solids like $P_4$.
Figure 1: Because N is small, it forms triple bonds. P is too large for pi bonds, so it forms a cage of single bonds.
3. The Weak N-N Single Bond & Catenation
Catenation is the property of an element to form long chains with its own atoms. Phosphorus shows strong catenation, but Nitrogen does not. Why?
Because the Nitrogen atom is so small, when two N atoms form a single bond, the non-bonding electrons (lone pairs) on each atom are forced extremely close to each other. This results in massive interelectronic repulsion, which heavily destabilizes and weakens the single bond.
4. Restricted Covalency (The Absence of d-Orbitals)
Nitrogen belongs to the second period, meaning its valence shell ($n=2$) contains only one $2s$ orbital and three $2p$ orbitals.
Because it only has 4 available orbitals (one $s$ and three $p$), Nitrogen's maximum covalency is restricted to 4 (e.g., in the Ammonium ion, $NH_4^+$). It cannot expand its octet.
Phosphorus has vacant 3d-orbitals. It can excite a 3s electron into the 3d orbital to form 5 bonds ($PCl_5$). Because Nitrogen completely lacks d-orbitals, it cannot expand its covalency to 5. Therefore, $NCl_5$ does not exist.
5. Hydrogen Bonding
Because of its small size and high electronegativity, Nitrogen is the only element in Group 15 that can form strong Hydrogen Bonds.
- Boiling Point of Hydrides: Due to strong intermolecular hydrogen bonding, Ammonia ($NH_3$) has a significantly higher boiling point than Phosphine ($PH_3$).
- Solubility: $NH_3$ is highly soluble in water because it forms hydrogen bonds with water molecules, whereas $PH_3$ does not.
Mastery Check: Anomalous Nitrogen
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