Answer the question on basis of passage given below:

Nitrogen differs from the rest of the members of its group due to its small size, high electronegativity, high ionisation enthalpy and non-availability of d-orbitals. Nitrogen has unique ability to form $pπ-pπ$ multiple bonds with itself and with other elements having small size and high electronegativity (eg, C, O) Heavier elements of this group do not form $pπ-pπ$ bonds as their atomic orbitals are so large and diffused that they cannot have effective overlapping.

Single bond is not formed by two atoms of:

$N - N$

The correct answer is Option 2. $N - N$.

Group 15 of the periodic table includes nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi). These elements have distinct properties due to their differences in size, electronegativity, and the availability of d-orbitals for bonding.

Key Characteristics of Nitrogen

Small Atomic Size: Nitrogen is the smallest element in Group 15. Its smaller size allows for effective overlap of atomic orbitals when forming bonds, particularly p-orbitals.

High Electronegativity: Nitrogen has a high electronegativity (3.04 on the Pauling scale), making it very effective at attracting electrons when forming covalent bonds.

High Ionization Enthalpy: The energy required to remove an electron from nitrogen is high, which makes it less likely to form cations. Instead, nitrogen tends to form covalent bonds.

\(p\pi −p\pi \) Bonding: Nitrogen can form multiple bonds (double and triple bonds) through \( p\pi−p\pi \) overlap with other small and electronegative atoms, such as itself (N−N) and oxygen (N−O).

N−N Bonding

Diatomic Nitrogen (\( \text{N}_2 \)): In its most stable form, nitrogen exists as \( \text{N}_2 \), which has a triple bond (one sigma bond and two pi bonds). The triple bond is extremely strong and stable due to effective overlap of the \( p \) orbitals.

Single Bond Formation: Nitrogen can form single bonds in certain compounds, such as hydrazine (\( \text{N}_2\text{H}_4 \)). However, the bond length and strength of the N−N single bond are weaker compared to the triple bond.

Comparison with Other Group 15 Elements

P−P Bonding:

Phosphorus can easily form single bonds, particularly in \( \text{P}_4 \) (white phosphorus), where each phosphorus atom is bonded to three others. The \( p \) orbitals in phosphorus can overlap effectively enough to form stable single bonds.

As−As Bonding:

Arsenic can also form single bonds in \( \text{As}_4 \). However, these bonds are generally weaker than P−P bonds due to the larger atomic size and more diffuse orbitals of arsenic.

Sb−Sb Bonding:

Antimony does form single bonds, but they are quite weak and less stable compared to the other elements in the group. The larger size of Sb leads to poor overlap of orbitals, making single bonds less favorable.

Why N−N is Unique

Preference for Multiple Bonds: Nitrogen’s ability to form strong triple bonds means that it typically favors multiple bonding over single bonding. In many chemical environments, nitrogen prefers to be part of stable structures like \( \text{N}_2 \) rather than forming single N−N bonds.

Bonding Limitations: While nitrogen can form a single bond, its preference for triple bonding reflects its unique position in the periodic table. The size and electronegativity of nitrogen allow it to engage in \( p\pi−p\pi \) bonding, which is not effectively replicated by heavier elements.

Conclusion

In summary, nitrogen's ability to form strong multiple bonds, combined with its small size and high electronegativity, means that while it can technically form N−N single bonds in certain contexts, it is more commonly recognized for its triple bond in molecular nitrogen (\( \text{N}_2 \)).

Therefore, when considering the question about the formation of single bonds, N−N is unique because it predominantly forms strong multiple bonds instead. This understanding highlights the distinctiveness of nitrogen among its group members.