Problem 44 - 43 Years JEE ADVANCED (1978-2020) + JEE MAIN Chapterwise & Topicwise Solved Papers Chemistry

Question

Interpret the non-linear shape of \(\mathrm{H}_{2} \mathrm{~S}\) molecule and non-planar shape of \(\mathrm{PCl}_{3}\) using valence shell electron pair repulsion (VSEPR) theory. (Atomic numbers : \(\mathrm{H}=1, \mathrm{P}=15, \mathrm{~S}=16, \mathrm{Cl}=17 .\) )

Step-by-Step Solution

Answer

\(\mathrm{H}_2\mathrm{S}\) is bent due to its two lone pairs and two bonds. \(\mathrm{PCl}_3\) is trigonal pyramidal due to its three bonds and one lone pair.

Step 1: Determine the Number of Valence Electrons

First, we need to calculate the number of valence electrons for each central atom in the molecules. For \(\mathrm{H}_2\mathrm{S}\): \(\mathrm{S}\) has 6 valence electrons. For \(\mathrm{PCl}_3\): \(\mathrm{P}\) has 5 valence electrons.

Step 2: Count the Electron Pairs Around the Central Atom

Based on the atoms bonded and lone pairs, determine electron pairs around the central atom. In \(\mathrm{H}_2\mathrm{S}\), \(\mathrm{S}\) forms two bonds with hydrogen and has two lone pairs. In \(\mathrm{PCl}_3\), \(\mathrm{P}\) forms three bonds and has one lone pair.

Step 3: Apply VSEPR Theory

VSEPR theory suggests that electron pairs will arrange themselves to minimize repulsion. For \(\mathrm{H}_2\mathrm{S}\), the two bonds with hydrogen and two lone pairs on sulfur create a bent (non-linear) shape. For \(\mathrm{PCl}_3\), the three chlorine-phosphorus bonds and one lone pair create a trigonal pyramidal shape, which is non-planar.

Key Concepts

Molecular Geometry

Molecular geometry is all about the three-dimensional arrangement of atoms in a molecule. It's like figuring out how atoms sit in space relative to each other. Imagine that each molecule is a little structure built with points hanging in various directions. The shape of a molecule greatly affects its physical and chemical properties because it determines how molecules react with each other.
In the case of \( \mathrm{H}_2\mathrm{S} \), its molecular geometry is non-linear, resembling something like the shape of a bent stick. This is similar to a famous molecule, water (\( \mathrm{H}_2\mathrm{O} \)), where two hydrogen atoms form bonds at an angle, resulting in a bent shape.
For \( \mathrm{PCl}_3 \), the molecular geometry is described as trigonal pyramidal – picture a pyramid with a triangular base. This shape is non-planar, meaning it doesn’t perfectly align in a flat plane. Such geometries stem from the specific arrangements around a central atom, shaped by bonds and lone pairs.

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a crucial role in defining an atom’s chemical bonding behavior. They are like the communication tools atoms use to interact and connect with each other.
Each element has a different number of valence electrons, depending on its position in the periodic table. For instance, sulfur \( (\mathrm{S}) \) in \( \mathrm{H}_2\mathrm{S} \) comes equipped with six valence electrons. Phosphorus \( (\mathrm{P}) \) in \( \mathrm{PCl}_3 \), on the other hand, brings along five valence electrons.
These valence electrons are key in forming bonds with other atoms. The bonds in \( \mathrm{H}_2\mathrm{S} \) and \( \mathrm{PCl}_3 \) involve these valence electrons, which ultimately impact the shape or geometry of the molecules. Understanding valence electrons helps in predicting how atoms will bond and what kind of structures they might form.

Lone Pairs

Lone pairs are pairs of valence electrons that are not shared with another atom and are not involved in bonding. Think of them as the electrons that prefer to stay at home. They don’t engage in bonding but still play a vital role in determining a molecule’s shape.
In a molecule like \( \mathrm{H}_2\mathrm{S} \), sulfur has two lone pairs in addition to the two hydrogen atoms it is bonded with. This configuration results in a bent shape due to the presence of these unshared electrons. Lone pairs repel the bonded pairs of electrons more strongly than the bonded pairs repel each other.
Similarly, \( \mathrm{PCl}_3 \) features a lone pair on the phosphorus atom. This lone pair nudges the three \( \mathrm{P-Cl} \) bonds into a trigonal pyramidal shape rather than letting them spread out in a flat triangle. Understanding the impact of lone pairs helps in explaining why the molecular geometry isn’t always perfectly aligned.

Electron Pair Repulsion

The Valence Shell Electron Pair Repulsion (VSEPR) theory is based on a simple, yet powerful principle: electron pairs around a central atom will repel each other. They organize themselves in a way that minimizes the repulsion, helping to define the shape of the molecule.
In \( \mathrm{H}_2\mathrm{S} \), we see a clear example of VSEPR theory at work. The presence of two bonded pairs (\( \mathrm{S-H} \) bonds) and two lone pairs around sulfur cause the electrons to push away from each other. This results in a bent, non-linear structure as they seek to reduce repulsion.
With \( \mathrm{PCl}_3 \), the lone pair on phosphorus exerts a repulsive force on the \( \mathrm{P-Cl} \) bonds, creating a trigonal pyramidal shape. This theory helps us understand how both lone pairs and bonded pairs interact, ultimately shaping the molecular geometry. VSEPR theory provides a simple model to predict and explain the molecular shape based on electron pair repulsions.