Problem 189 - A Complete Resource Book in Chemistry for JEE Main

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 P than \(\mathrm{N}\) (b) lower tendency of \(\mathrm{N}\) to form covalent bond (c) availability of vacant d orbital in P but not in \(\mathrm{N}\) (d) statement is itself incorrect

Step-by-Step Solution

Answer

(c) availability of vacant d orbital in P but not in N.

Step 1: Analyze the Valence Orbitals

Consider the electronic configuration of phosphorus (P) and nitrogen (N). Phosphorus has an electron configuration ending in 3p8 and has vacant 3d orbitals, while nitrogen ends in 2p3, lacking d orbitals.

Step 2: Evaluate Molecular Formation

9Cl3 can exist for both P and N since three bonds utilize the valence p orbitals. However, extending to five bonds for 9Cl5 is only feasible for elements that can use d orbitals for expansion of the octet.

Step 3: Determine Element Capabilities

Phosphorus can form more than four covalent bonds by using its vacant d orbitals, allowing for 9Cl5 formation. Nitrogen cannot do this as it has no vacant d orbitals.

Step 4: Identify the Correct Explanation

The presence of vacant d orbitals in phosphorus allows it to expand beyond an octet, unlike nitrogen, explaining the discrepancy in existence of 9Cl3 and 9Cl5.

Key Concepts

Valence Orbitals

Valence orbitals are crucial in understanding how atoms form bonds. These orbitals are the outermost electron shells of an atom that participate in chemical bonding. For any given element, the valence orbitals determine how it can interact with other atoms to form molecules. In the case of phosphorus ( 5 ext{P} ) and nitrogen ( 5 ext{N} ), the valence electrons reside in different shells. Phosphorus has an electron configuration that ends in 3p, meaning its valence electrons are in the third energy level, while nitrogen's valence electrons are in the 2p orbital, marking the second energy level. Because phosphorus has electrons in a higher energy level, it has additional orbitals available beyond those occupied in bonding situations, whereas nitrogen does not.

• Phosphorus (P): Ends in 3p, which means it has available 3s, 3p, and 3d orbitals.
• Nitrogen (N): Ends in 2p, with only 2s and 2p orbitals available for bonding.

Due to these differences, phosphorus and nitrogen exhibit varied abilities to form bonds, particularly when the number of bonds requires accessing more orbitals than available in the valence shell.

Vacant d Orbitals

Vacant d orbitals play a critical role when it comes to bond formation, especially for elements beyond the second period of the periodic table. The presence of these orbitals allows atoms to expand their valence shell, making additional space for more electrons. In phosphorus, the 3d orbitals are vacant and can be used once the usual valence orbitals (3s and 3p) are filled. This capacity to utilize 3d orbitals gives it a noteworthy advantage. It allows the atom to accommodate more electrons by forming more bonds, specifically in cases where a molecule desires more than an octet configuration. In contrast, nitrogen lacks d orbitals in its valence shell since it only possesses 2s and 2p orbitals. Thus, nitrogen is restricted to its octet, unlike phosphorus, which can extend beyond this limit using its vacant d orbitals. This is a reason nitrogen can form NCl 3 but not NCl 5, as additional orbitals are not available to support more than three bonds.

Covalent Bond Formation

Covalent bonds involve the sharing of electrons between atoms to fulfill their valence shell, leading to stable molecular structures. The ability of an element to form covalent bonds depends on its electronic configuration and the availability of its orbitals. Phosphorus and nitrogen both form three covalent bonds with chlorine in PCl 3 and NCl 3, respectively. However, phosphorus can form PCl 5 due to its capability to utilize the empty 3d orbitals, accommodating five bonds by expanding its octet. This extended bonding is enabled by the vacant 3d orbitals, which can house the extra electron pairs required for additional covalent bonds.

• Phosphorus Bonding: Utilizes 3s, 3p, and vacant 3d orbitals.
• Nitrogen Bonding: Limited to 2s and 2p orbitals.

In contrast, nitrogen does not possess vacant d orbitals and thus cannot expand beyond its octet. As such, it cannot form more than three covalent bonds, which prevents the formation of NCl 5. This inherently restricts nitrogen to completing its octet configuration without the option of further expansion.