Animated Solution for Chemistry - s and p-Block Elements: PCl_3 and PCl_5 both exist; NCl_3 exists but NCl_5 does not exist. It is due to

Have you ever wondered why nature allows certain molecules to exist while strictly forbidding others? In the world of chemistry, the rules of molecular architecture are dictated by quantum mechanics. A classic example of this is the tale of two closely related elements: Nitrogen ($N$) and Phosphorus ($P$). Both belong to Group 15 of the periodic table, meaning they share similar chemical properties. Yet, when they react with chlorine, phosphorus can form both $PC{l}_3$ and $PC{l}_5$, whereas nitrogen can only form $NC{l}_3$. The molecule $NC{l}_5$ is completely unheard of. Why does phosphorus get to break the rules while nitrogen is held back? Let's dive into their electronic configurations to uncover the truth.

To understand nitrogen's limitations, we must look at its atomic structure. Nitrogen has an atomic number of 7. Its ground-state electronic configuration is $1s^{2}2s^{2}2p^3$. The outermost shell, or the valence shell, is the second shell ($n=2$).

In the quantum world, the second shell only contains $s$ and $p$ subshells. There is absolutely no such thing as a $2d$ orbital. This is a hard physical constraint. Nitrogen has three unpaired electrons in its $2p$ orbitals, which it happily shares with three chlorine atoms to form $NC{l}_3$.

But what if it wanted to form five bonds? To do that, it would need five unpaired electrons. It could theoretically try to unpair the two electrons in the $2s$ orbital, but where would the excited electron go? Since there are no $2d$ orbitals, the electron would have to jump all the way to the $3s$ orbital. The energy required for such a massive jump is astronomically high, making it chemically impossible under normal conditions. Therefore, nitrogen's maximum covalency is restricted to 4 (by donating its lone pair), and it can never form $NC{l}_5$.

The Rule Breaker

Now, let's shift our focus to phosphorus. Sitting right below nitrogen in Group 15, phosphorus has an atomic number of 15. Its ground-state electronic configuration is $[Ne]3s^{2}3p^3$.

At first glance, it looks very similar to nitrogen—it has three unpaired electrons in the $p$ orbitals, which explains why $PC{l}_3$ forms easily. However, phosphorus has a hidden superpower. Because its valence shell is the third shell ($n=3$), it has access to $d$ orbitals! Even though the $3d$ orbitals are completely empty in the ground state ($3d^0$), they are energetically accessible.

When highly electronegative atoms like chlorine approach phosphorus, they create a strong chemical demand for more bonds. To satisfy this, phosphorus undergoes a process called excitation. It absorbs a small amount of energy, and one of the paired electrons from the $3s$ orbital is promoted into one of the empty $3d$ orbitals.

The new excited-state configuration becomes $[Ne]3s^{1}3p^{3}3d^1$.

Suddenly, the landscape changes. Phosphorus now has exactly five unpaired electrons—one in the $s$ orbital, three in the $p$ orbitals, and one in the $d$ orbital. These five orbitals hybridize to form five equivalent $s{p}^{3}d$ hybrid orbitals, allowing phosphorus to form five strong covalent bonds with five chlorine atoms. This results in the stable formation of $PC{l}_5$.

The existence of $PC{l}_5$ and the non-existence of $NC{l}_5$ perfectly illustrates how the availability of atomic orbitals dictates chemical bonding. Phosphorus can expand its octet and exhibit a covalency of 5 because it has vacant d-orbitals in its valence shell. Nitrogen, lacking these d-orbitals, is strictly bound by the octet rule. Therefore, the correct explanation for this phenomenon is the availability of vacant d-orbitals in phosphorus but not in nitrogen.