Animated Solution for Chemistry - s and p-Block Elements: The metal that forms nitride by reacting directly with ${\text{N}}_2$ of air, is

When we think of exposing highly reactive alkali metals to the open air, we immediately picture rapid oxidation, tarnishing, and sometimes even spontaneous combustion. Air is a mixture of gases, primarily composed of about $78\%$ nitrogen (${\text{N}}_2$) and $21\%$ oxygen (${\text{O}}_2$).

Given that nitrogen is the most abundant gas, one might assume it would be the primary reactant. However, nitrogen gas is notoriously unreactive under standard conditions. The two nitrogen atoms are bound together by a formidable triple bond ($\text{N}\equiv \text{N}$), which boasts a staggering bond dissociation energy of approximately $941\text{ kJ/mol}$. Breaking this fortress of a bond requires an immense thermodynamic driving force.

If we take a piece of sodium ($\text{Na}$), potassium ($\text{K}$), rubidium ($\text{Rb}$), or cesium ($\text{Cs}$) and expose it to air, it will eagerly react with oxygen to form various oxides, peroxides, or superoxides. But they will completely ignore the nitrogen.

Why? Because forming a nitride requires breaking that stubborn $\text{N}\equiv \text{N}$ bond. To make the overall reaction thermodynamically favorable (exothermic), the energy released when the metal ions and nitride ions come together to form a solid crystal lattice—known as the Lattice Energy—must be massive enough to compensate for the energy spent breaking the nitrogen bond. The larger alkali metals simply cannot generate enough lattice energy to pay this thermodynamic toll.

The Tiny Titan

There is one glaring exception in Group 1: Lithium ($\text{Li}$).

When lithium is heated in air, it doesn't just form an oxide; it reacts directly with the nitrogen to form Lithium Nitride (${\text{Li}}_3\text{N}$), a striking ruby-red solid. The chemical equation for this unique reaction is:

But what makes lithium so special? The answer lies in its atomic dimensions.

Fajans' Rules

Lithium is the smallest of the alkali metals. When it loses its single valence electron to become a ${\text{Li}}^{+}$ ion, it becomes exceptionally tiny. This small size concentrates its positive charge over a very small volume, giving it an incredibly high charge density.

On the other side of the reaction, we have the nitride ion, ${\text{N}}^{3-}$. This ion is large and carries a high negative charge, meaning its outermost electrons are loosely held and highly polarizable.

According to Fajans' Rules, when a small, highly charged cation approaches a large, polarizable anion, the cation will strongly attract the anion's electron cloud. The ${\text{Li}}^{+}$ ion effectively pulls and distorts the electron cloud of the ${\text{N}}^{3-}$ ion towards itself. This phenomenon is known as polarization.

Lattice Energy

This intense polarization introduces a significant degree of covalent character into the ionic bond, pulling the ions closer together. The result is an exceptionally tight, stable, and tightly packed crystal lattice.

The lattice energy released during the formation of the ${\text{Li}}_3\text{N}$ crystal is enormous. It is this massive release of energy that successfully offsets the high energy cost of breaking the $\text{N}\equiv \text{N}$ triple bond and the energy required to add three electrons to a nitrogen atom (which is highly endothermic due to electron-electron repulsion).

Because sodium, potassium, rubidium, and cesium are much larger, their cations have lower charge densities. They cannot polarize the nitride ion effectively, resulting in a lattice energy that is too weak to drive the reaction forward.

Therefore, lithium stands alone among the alkali metals in its ability to form a stable nitride directly from the air. This makes option (d) the correct answer.

A Broader Perspective: If high charge density is the key, what about Group 2 elements (the alkaline earth metals)? Elements like beryllium, magnesium, and calcium form divalent cations (${\text{M}}^{2+}$). Because of their higher $+2$ charge and relatively small sizes, their charge densities are even greater than that of lithium. Consequently, all Group 2 metals readily react with nitrogen to form nitrides (e.g., ${\text{Mg}}_3{\text{N}}_2$). Understanding this trend of charge density and lattice energy is a powerful tool for predicting chemical reactivity!