Problem 140 - A Complete Resource Book in Chemistry for JEE Main

Question

Consider the following molecules or ions. (i) \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\) (ii) \(\mathrm{NH}_{4}^{+}\) (iii) \(\mathrm{SO}_{4}^{2}-\) (iv) \(\mathrm{ClO}_{4}^{-}\) (v) \(\mathrm{NH}_{3}\) sp \(^{3}\) hybridization is involved in the formation of (a) (i), (ii), (v) only (b) (i), (ii) only (c) (i), (ii), (iii), (iv) (d) (i), (ii), (iii), (iv), (v)

Step-by-Step Solution

Answer

(a) (i), (ii), (v) only

Step 1: Understanding sp3 Hybridization

For a molecule to show sp³ hybridization, the central atom must form 4 equivalent orbitals to bond with other atoms. This generally happens with a tetrahedral arrangement where the central atom uses one s orbital and three p orbitals to create four sp³ hybrid orbitals.

Step 2: Analyzing CH2Cl2

In \(\mathrm{CH}_{2}\mathrm{Cl}_{2}\), the central atom is carbon. Carbon forms 4 sigma bonds (2 with hydrogen and 2 with chlorine), indicating sp³ hybridization with a tetrahedral shape.

Step 3: Analyzing NH4⁺

In the ammonium ion \(\mathrm{NH}_{4}^{+}\), nitrogen forms 4 sigma bonds with four hydrogen atoms, utilizing sp³ hybridization to maintain a tetrahedral geometry.

Step 4: Analyzing SO4²⁻

In \(\mathrm{SO}_{4}^{2-}\), sulfur is the central atom. Although it has four oxygen atoms linked to it, sulfur typically involves p-d hybridization, not just sp³, due to the ability to expand its octet.

Step 5: Analyzing ClO4⁻

In the perchlorate ion \(\mathrm{ClO}_{4}^{-}\), chlorine uses sp³ hybridization to form sigma bonds with oxygen, organized in a tetrahedral geometry.

Step 6: Analyzing NH3

For \(\mathrm{NH}_{3}\), nitrogen forms 3 sigma bonds with hydrogen and has one lone pair. The nitrogen atom undergoes sp³ hybridization, using 4 equivalent orbitals to accomplish its bonding, arranged in a trigonal pyramidal shape.

Step 7: Conclusion from Analysis

Based on the analysis, sp³ hybridization occurs in \(\mathrm{CH}_{2}\mathrm{Cl}_{2}\), \(\mathrm{NH}_{4}^{+}\), and \(\mathrm{NH}_{3}\) only. Thus, the correct option is (a) (i), (ii), (v) only.

Key Concepts

Tetrahedral Geometry

Tetrahedral geometry is a fundamental concept in chemistry, especially when discussing molecular shapes and hybridization. This structure arises when a central atom forms four covalent bonds positioned as far apart as possible, resulting in a three-dimensional shape. The angle between these bonds is approximately 109.5 degrees.

This shape is created because tetrahedral geometry minimizes the repulsion between electron pairs around the central atom, leading to a stable and balanced arrangement. For instance, in the case of \(\mathrm{CH}_{2}\mathrm{Cl}_{2}\) or \(\mathrm{NH}_{4}^{+}\), the central carbon and nitrogen atoms use their sp³ hybrid orbitals to form these four bond directions. This optimal shape allows molecules with sp³ hybridization to evenly distribute the electron clouds, helping in the stability and reactivity of compounds.

Tetrahedral geometry is often considered whenever sp³ hybridization is involved, and it serves as a key tool in predicting and explaining molecular behavior and interactions in chemistry.

Sigma Bond

A sigma bond is the strongest type of covalent chemical bond and is central to understanding chemical bonding. Formed by the direct overlap of orbitals, sigma bonds are characterized by their high bond strength and symmetry along the bond axis. In general, these bonds form between s orbitals, p orbitals, or hybrid orbitals like sp³.

A classic example of sigma bonds can be seen in the methane molecule, \(\mathrm{CH}_{4}\), where the central carbon atom forms four sigma bonds with hydrogen atoms. This arrangement incorporates the sp³ hybrid orbitals, allowing strong, stable connections.

Because sigma bonds involve the head-on overlap of orbitals, they form the primary framework for molecules, providing flexibility and rotation freedom around the bond axis. This characteristic is part of why sigma bonds are ubiquitous in organic and inorganic chemistry, forming the robust backbones of countless molecules. Understanding sigma bonds is crucial when dissecting the molecular structure and bonding patterns.

Orbitals

Orbitals are the regions around an atom's nucleus where electrons are likely to be found. The concept of orbitals is essential in discussing atomic and molecular structures, and they come in different shapes and energy levels: s, p, d, and f.

Each type of orbital has a distinct shape and capacity:

• s orbitals are spherical and can hold up to two electrons.
• p orbitals have a dumbbell shape and consist of three orientations, each holding two electrons for a total of six.
• d and f orbitals are more complex and are crucial for atoms with more electrons and higher energy levels.

Hybridization alters the typical shapes of atomic orbitals to become suitable for forming bonds. For instance, in sp³ hybridization, one s orbital and three p orbitals blend to form four equivalent hybrid orbitals, enabling the formation of a tetrahedral structure. This transformation allows the central atom to optimally share its electrons with other atoms, creating strong covalent bonds. Understanding orbitals and their hybridization is vital for predicting molecular geometry and reactivity.

Chemical Bonding

Chemical bonding is the force that holds atoms together within molecules, dictating the structure and properties of substances. The two main types of chemical bonds are ionic and covalent, both involving the attraction between electrons and nuclei.

In covalent bonding, atoms share electron pairs, which typically form due to the overlap of atomic orbitals. The shared electron pair results in a stable configuration for the bonded atoms, creating compounds with specific arrangements and characteristics.

Ionic bonds, on the other hand, form through the transfer of electrons, leading to a full positive or negative charge, as seen in salts. Covalent bonds often entail hybridized orbitals, like sp³, to optimize the interaction between atoms.

• Covalent bonds can be single, double, or triple, depending on the number of shared electron pairs.
• Single bonds are usually sigma bonds providing the primary link between atoms.
• Double and triple bonds also incorporate pi bonds, which result from the lateral overlap of p orbitals.

Comprehending the nature of chemical bonding provides insight into molecular durability, electron distribution, and complex chemical interactions, forming a foundation for further study in chemistry.