Problem 93 - A Complete Resource Book in Chemistry for JEE Main

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 are larger than \(\mathrm{P}\) atoms 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

Step-by-Step Solution

Answer

(b) Phosphorus has low lying 3d orbitals for bonding, unlike Nitrogen.

Step 1: Analyze Atomic Size and Orbitals

Nitrogen (\(\mathrm{N}\)) belongs to the second period of the periodic table, while Phosphorus (\(\mathrm{P}\)) is part of the third period. \(\mathrm{P}\) has additional orbitals available in its electronic configuration, namely the \(3d\) orbitals, which \(\mathrm{N}\) does not have due to its smaller size and fewer energy levels.

Step 2: Formation of PCl₃ and PCl₅

Phosphorus can form \(\mathrm{PCl}_3\) using its \(3p\) orbitals by hybridization. Additionally, it can undergo \(sp^3d\) hybridization to expand its valency and form \(\mathrm{PCl}_5\). The \(3d\) orbitals enable bonding with more atoms than what is possible with \(\mathrm{N}\).

Step 3: Consideration for NCl₃ Formation

Nitrogen, being limited to 2\(s\) and 2\(p\) orbitals, forms \(sp^3\) hybridizations leading to compounds like \(\mathrm{NCl}_3\), but cannot utilize \(\mathrm{d}\) orbitals for expanded bonding and thus cannot form \(\mathrm{NCl}_5\).

Step 4: Analyze Answer Choices and Confirm

Given that \(\mathrm{P}\) uses its \(3d\) orbitals to allow expanded bonding compared to \(\mathrm{N}\), the correct reason is \((b)\). Phosphorus having low lying \(3 \,d\) orbitals but nitrogen not having them in its valence shell explains why \(\mathrm{P}\) can form \(\mathrm{PCl}_5\) and \(\mathrm{N}\) cannot.

Key Concepts

Periodic Table

The periodic table is a fundamental tool in chemistry that organizes elements based on their atomic number and properties. Each element in the table occupies a specific position that indicates a trend in physical and chemical behavior. Elements are arranged in rows named periods and columns called groups. The position of an element in the periodic table helps to predict its bonding capacity and reactivity.

When we look at nitrogen and phosphorus, they both belong to the pnictogen family in Group 15. However, nitrogen is in the second period, while phosphorus is in the third period. This difference in periodic location explains some of their chemical behavior, such as their ability to form different types of bonds. For example, since phosphorus is in a higher period, it contains a larger number of orbitals, including the 3d orbitals, which have an impact on its bonding abilities.

Hybridization

Hybridization is a key concept that describes how atomic orbitals mix to form new hybrid orbitals. These hybrid orbitals are critical for forming bonds in molecular compounds.

In the context of phosphorus, the process of hybridization allows it to utilize additional orbitals available beyond 3p, namely the 3d orbitals. By undergoing sp³d hybridization, phosphorus can expand its bonding capacity to form compounds like PCl₅. Nitrogen, on the other hand, does not have access to 3d orbitals due to its position in the second period, limiting it to sp³ hybridization, which is why it forms NCl₃ and not NCl₅.

Understanding hybridization is important for predicting the geometry of molecules, which directly impacts their chemical properties and reactions.

Valence Orbitals

Valence orbitals are the outermost orbitals of an atom that are involved in chemical bonding. The availability and configuration of these valence orbitals largely determine how an element can bond with others.

For nitrogen, the valence orbitals include just the 2s and 2p orbitals, which combine through hybridization to form sp³ orbitals. Phosphorus, in contrast, has its valence electrons present in the 3s, 3p, and potentially 3d orbitals. This access to 3d orbitals gives phosphorus more flexibility and enables it to form five bonds in structures like PCl₅ by hybridizing its orbitals to include these additional d-type orbitals as needed.

This difference in available orbitals between nitrogen and phosphorus is essential to explaining why they form different types and numbers of chemical compounds with similar elements such as chlorine.

Nitrogen and Phosphorus

Nitrogen and phosphorus, though part of the same group in the periodic table, display distinct chemical behaviors due to their position in different periods. Both elements are nonmetals and engage in various types of chemical bonding, but their ability to form specific compounds varies.

Nitrogen tends to form three covalent bonds due to its limited orbital availability, as seen in compounds like NCl₃. It lacks the d-orbitals needed for expanded valency. On the other hand, phosphorus is more versatile, easily forming both three and five covalent bonds, resulting in compounds like PCl₃ and PCl₅. This is due to its access to a third energy level incorporating 3d orbitals.

This variation highlights the impact of orbital availability and atomic structure on the chemical behaviors of nitrogen and phosphorus.

Molecular Compounds

Molecular compounds are made up of molecules formed from atoms sharing electrons through covalent bonds. Understanding how these compounds form requires knowledge of bonding and atomic structure.

For nitrogen and phosphorus, their molecular compounds with chlorine offer a great example. Nitrogen, due to its electron configuration and limited orbital set, forms simple molecular compounds like NCl₃. In contrast, the extra orbitals available in phosphorus allow it to form not only PCl₃, mirroring nitrogen’s bonding pattern, but also PCl₅ through a different, expanded hybridization process.

The study of molecular compounds considers both the number and type of bonds an atom can form, which is directly influenced by its electronic structure, especially the available valence orbitals. This demonstrates how the unique orbital configuration of each element guides its possible molecular formations.