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 \(\mathrm{P}\) 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

The non-existence of \(\mathrm{NCl}_{5}\) is due to the absence of vacant d orbital in nitrogen, option (c).

Step 1: Analyze the Compounds

The exercise presents two pairs of phosphorus and nitrogen chlorides: \(\mathrm{PCl}_{3}\) and \(\mathrm{PCl}_{5}\) versus \(\mathrm{NCl}_{3}\) and \(\mathrm{NCl}_{5}\). \(\mathrm{PCl}_{3}\) and \(\mathrm{NCl}_{3}\) both exist as stable compounds, whereas \(\mathrm{PCl}_{5}\) exists while \(\mathrm{NCl}_{5}\) does not.

Step 2: Consider Electron Configuration

Phosphorus has an electron configuration that allows for the expansion of its octet, as its configuration is \([Ne] 3s^2 3p^3\). Without a filled 3p orbital, phosphorus can use the available 3d orbitals to expand its valency further, allowing it to form \(\mathrm{PCl}_{5}\).

Step 3: Examine Nitrogen's Limitations

Nitrogen, however, has an electron configuration \([He] 2s^2 2p^3\) with no available 2d orbitals, as the second energy level contains only s and p orbitals. This means nitrogen can't expand its valency beyond four bonds, preventing it from forming \(\mathrm{NCl}_{5}\).

Step 4: Evaluate the Given Options

The difference in the ability to form \(\mathrm{PCl}_{5}\) and not \(\mathrm{NCl}_{5}\) is mainly due to the availability of vacant d orbital in phosphorus, while nitrogen lacks this feature, corresponding to option (c).

Key Concepts

Electron Configuration

Electron configuration describes how electrons are distributed within different atomic orbitals. For an atom, this configuration determines its chemical properties through the arrangement of electrons in reserved energy levels or shells.
Let's break it down:

• Energy Levels: Electrons are arranged in shells around the nucleus, where each shell has a specific energy level.
• Types of Orbitals: Within these energy levels, there are different types of orbitals: s, p, d, and f. These orbitals can hold a certain maximum number of electrons: 2 for s, 6 for p, 10 for d, and 14 for f.

Phosphorus and nitrogen demonstrate different electron configurations:
Phosphorus has the configuration \[Ne] 3s^2 3p^3\], meaning its electrons are spread across s and p orbitals in the third shell. This allows phosphorus to form compounds like \( \mathrm{PCl}_5 \) when utilizing d orbitals.
Nitrogen has an electron configuration of \[He] 2s^2 2p^3\] with electrons only in the second energy level, which is why it cannot expand beyond three bonds. Understanding this helps you see why \( \mathrm{PCl}_5 \) can exist, but \( \mathrm{NCl}_5 \) cannot.

d Orbitals

The concept of d orbitals is crucial for understanding why some elements can form more compounds than others. d orbitals are sets of orbitals found in the third and higher energy levels, capable of holding additional electrons.

• What they are: d orbitals can contain a maximum of 10 electrons, spread over five different d orbitals: \(d_{xy}, d_{yz}, d_{zx}, d_{x^2-y^2},\) and \(d_{z^2}\).
• Where they appear: Unlike 1st and 2nd energy levels that only have s and p orbitals, the 3rd energy level introduces d orbitals. This additional capacity is a key factor for elements like phosphorus which can expand their valency.

In terms of molecular formation, having available d orbitals allows an atom like phosphorus to accommodate more bonding pairs, leading to higher oxidation states and enabling compounds such as \( \mathrm{PCl}_5 \), where five Cl atoms can bond to a single P atom.

Valency Expansion

Valency expansion refers to the ability of an atom to form more bonds than it would normally:

• Understanding Basic Valency: Typically, an atom forms bonds based on its unpaired electrons in its outer shell, as dictated by its electron configuration.
• How expansion happens: Elements like phosphorus can use vacant d orbitals to expand beyond their typical bonding capacity. This is seen in the formation of \( \mathrm{PCl}_5 \).

While elements in the second period (like nitrogen) lack d orbitals, elements in the third period, such as phosphorus, have them. These vacant orbitals accommodate additional electrons, hence increasing bonding possibilities. Valency expansion is responsible for phosphorus forming both \( \mathrm{PCl}_3 \) and \( \mathrm{PCl}_5 \), while nitrogen is limited to \( \mathrm{NCl}_3 \).

Understanding valency expansion shows why some elements can form complex compounds, while others cannot, based on their position in the periodic table.