Question

\(\mathrm{N}\) forms \(\mathrm{NCl}_{3}\) whereas \(\mathrm{P}\) can form both \(\mathrm{PCl}_{3}\) and \(\mathrm{PCl}_{5}\). Why? (a) \(\mathrm{N}\) atoms is larger than \(\mathrm{P}\) in size (b) \(\mathrm{P}\) has low lying \(3 \mathrm{~d}\) orbitals, which can be used for bonding but \(\mathrm{N}\) does not have \(3 \mathrm{~d}\) orbitals in its valence shell (c) \(\mathrm{P}\) is more reactive towards \(\mathrm{Cl}\) than \(\mathrm{N}\) (d) None of these

Short answer

(b) \(\mathrm{P}\) has low lying \(3 \mathrm{d}\) orbitals for bonding, \(\mathrm{N}\) does not.

Step-by-step solution

Understanding Elemental Differences

Both nitrogen (\(\mathrm{N}\)) and phosphorus (\(\mathrm{P}\)) belong to group 15 of the periodic table, but they are in different periods. \(\mathrm{N}\) is in the second period, while \(\mathrm{P}\) is in the third. This information is crucial as it affects their ability to expand their valence shell through available orbitals.

Valence Electron Configuration

\(\mathrm{N}\), being in the second period, has a valence electron configuration of \(2s^2 2p^3\). \(\mathrm{P}\), in the third period, has \(3s^2 3p^3\), and importantly, access to \(3d\) orbitals, although not occupied, they are available for bond formation.

Capability to Expand Valence

\(\mathrm{P}\) can accommodate more electrons (as in \(\mathrm{PCl}_5\)) because it can potentially use the \(3d\) orbitals in bonding, which provides more space for electrons. \(\mathrm{N}\) cannot form analogous \(\mathrm{NCl}_5\) as it lacks \(3d\) orbitals to facilitate additional bond formation, thus can only form \(\mathrm{NCl}_3\).

Conclusion on Reactions with Chlorine

Given \(\mathrm{P}\)'s ability to use \(3d\) orbitals, option (b) explains why \(\mathrm{P}\) can form \(\mathrm{PCl}_3\) and \(\mathrm{PCl}_5\), whereas \(\mathrm{N}\) can only form \(\mathrm{NCl}_3\). \(\mathrm{N}\) lacks these orbitals, blocking its ability to expand the number of bonds.

Key concepts

Valence Shell Expansion

Valence shell expansion is a fascinating concept in chemistry that explains why certain elements can form more bonds than others. It is all about the ability of an atom to use additional orbitals beyond the typical s and p orbitals found in its valence shell.
In the periodic table, elements like phosphorus (\(\mathrm{P}\)) can access d-orbitals. This is because phosphorus is in the third period where the 3d orbitals are available, even though they aren’t occupied in the ground state. On the other hand, nitrogen (\(\mathrm{N}\)), which is in the second period, lacks these d-orbitals entirely.

The presence of these additional orbitals allows phosphorus to expand its valence shell. This expansion is what makes it possible for phosphorus to form molecules like \(\mathrm{PCl}_5\), where five chlorine atoms bond with phosphorus. In contrast, nitrogen cannot expand beyond three valence electrons for bonding as it can only form \(\mathrm{NCl}_3\). This intrinsic limitation happens because nitrogen has no room for expansion without access to additional orbitals.

d-Orbitals

d-Orbitals play a crucial role in the reactivity and bonding capabilities of elements. These orbitals become available to elements starting from the third period on the periodic table. Although they are not filled in the ground state, they can participate in bonding.
For instance, phosphorus can use its 3d orbitals to form additional bonds, thereby increasing its coordination number. This is why \(\mathrm{P}\) can form both \(\mathrm{PCl}_3\) and \(\mathrm{PCl}_5\).

It is essential to recognize that the 3d orbitals act as a reserve, enabling elements in the third period and beyond to accommodate extra electrons. This capability is what prevents nitrogen, lacking accessible d-orbitals, from forming compounds such as \(\mathrm{NCl}_5\). Hence, understanding the availability and usage of d-orbitals allows us to predict and explain the unique bonding behavior of different elements.

Nitrogen and Phosphorus Reactivity

Nitrogen and phosphorus, while similar in that they both belong to group 15 of the periodic table, exhibit markedly different chemical reactivity due to their positions in different periods. This difference fundamentally arises from their capability to expand their valence shells and the availability of d-orbitals.
Phosphorus shows greater versatility in its reactivity, thanks to its ability to form \(\mathrm{PCl}_5\), whereas nitrogen is limited to forming \(\mathrm{NCl}_3\). The lack of d-orbitals in nitrogen means it cannot accommodate additional bonds needed for higher chlorides.

In general:

• \(\mathrm{N}\), with a limited valence shell, often gets involved in the formation of a limited number of bonds.
• \(\mathrm{P}\), with more flexible bonding capabilities, can increase its coordination number using those "extra" d-orbitals.

This distinct behavior in nitrogen and phosphorus reactivity stems from their differential orbital availability and usage, influencing their chemical reactivity significantly.