Question

\(\mathrm{PCl}_{3}\) and \(\mathrm{PCl}_{5}\) both exists; \(\mathrm{NCl}_{3}\) exists but \(\mathrm{NCl}_{5}\) does not exist. It is due to (a) lower electronegativity of \(\mathrm{P}\) than \(\mathrm{N}\) (b) lower tendency of \(\mathrm{N}\) to form covalent bond (c) availability of vacant d orbital in \(\mathrm{P}\) but not in \(\mathrm{N}\) (d) statement is itself incorrect

Short answer

The existence of {PCl}_{5} and not {NCl}_{5} is because of the availability of vacant d orbitals in phosphorus.

Step-by-step solution

Understanding the Existence Criteria

To solve this problem, we must determine under what conditions these compounds exist. The key lies in the ability of an element to expand its octet using vacant d orbitals, which is possible for elements in period 3 and beyond like phosphorus, but not for nitrogen, which is in period 2.

Analyze Nitrogen Compounds

Nitrogen ({N}) has electronic configuration 1s² 2s² 2p³. It lacks vacant d orbitals needed to expand its valency beyond 8 electrons. Hence, it can form {NCl}_{3}, fulfilling the octet rule, but not {NCl}_{5}, as it cannot accommodate more than 8 electrons.

Analyze Phosphorus Compounds

Phosphorus ({P}) has the electronic configuration 1s² 2s² 2p⁶ 3s² 3p³ with vacant 3d orbitals available. This allows it to form {PCl}_{3} with a typical octet, and {PCl}_{5} by expanding its octet using the d orbitals, unlike nitrogen.

Conclusion

Given these analyses, the existence of {PCl}_{5} and non-existence of {NCl}_{5} is primarily due to the presence of vacant d orbitals in phosphorus, which are absent in nitrogen. Thus, option (c) is correct.

Key concepts

Octet Rule

The octet rule is a fundamental concept in chemistry that helps us understand how atoms bond and form stable molecules. This rule suggests that atoms tend to form bonds until they are surrounded by eight electrons in their valence shell, mimicking the electron configuration of noble gases which are inherently stable.
When atoms bond, they either share or transfer electrons to achieve a full outer shell of eight electrons. For example, in the compound \(Cl_\3\), nitrogen follows the octet rule by sharing electrons with three chlorine atoms. This completes its valence shell with a total of eight electrons, leading to a stable and common nitrogen compound.
The octet rule is crucial in predicting the types of bonds and the number of bonds an atom can form but has limitations for elements beyond the second period.

Vacant d Orbitals

Vacant d orbitals play a pivotal role in chemical bonding, especially for elements in the third period and beyond. These orbitals allow certain elements to expand their octet, going beyond the typical eight electrons.
Phosphorus, an element in period 3, possesses vacant 3d orbitals in its electronic configuration. These additional orbitals enable it to accommodate more than eight valence electrons and form compounds like \(PCl_\5\), where phosphorus uses these d orbitals to hold extra electrons beyond the octet.
In contrast, nitrogen, located in period 2, doesn't have d orbitals in its valence shell and therefore cannot expand its octet. This explains why \(NCl_\5\) is nonexistent while \(PCl_\5\) is possible, highlighting the significance of these d orbitals in chemical bonding.

Covalent Bonding

Covalent bonding occurs when two atoms share one or more pairs of electrons, allowing them to achieve a full outer shell and gaining stability.
In molecules like \(NCl_\3\), nitrogen forms covalent bonds with chlorine by sharing its electrons. Each chlorine atom contributes one electron to a pair with nitrogen, satisfying both their valence shells.
Covalent bonds are characterized by the sharing of electrons, as opposed to ionic bonds which involve the transfer of electrons. This type of bonding typically occurs between nonmetals and is essential for forming a variety of organic and inorganic compounds.
The strength and type of covalent bonding can vary, leading to different molecular properties, such as polarity and bond angles across different molecules.

Electronic Configuration

Electronic configuration describes the arrangement of electrons in an atom's orbitals. This configuration determines how an atom can bond and interact with others.
For elements like nitrogen, its electronic configuration \(1s^2 2s^2 2p^3\) defines its bonding capacity, as it can form a maximum of three covalent bonds like in \(NCl_\3\) due to having five valence electrons.
Phosphorus, with the electronic configuration \(1s^2 2s^2 2p^6 3s^2 3p^3\), has five valence electrons but also accessible 3d orbitals. This configuration affords it more versatility in bonding, enabling it to form compounds with expanded octets, such as \(PCl_\5\).
Understanding an element's electronic configuration helps predict and explain its chemical behavior and the types of compounds it can form.