Problem 99 - A Complete Resource Book in Chemistry for JEE Main

Question

Among the following species the ones having square planar geometry for central atom are 1\. \(\mathrm{XeF}_{4}\) 2\. \(\mathrm{SF}_{4}\) 3\. \(\left[\mathrm{NiCl}_{4}\right]^{2-}\) 4\. \(\left[\mathrm{PdCl}_{4}\right]^{2-}\) (a) 1 and 4 (b) 1 and 2 (c) 2 and 3 (d) 3 and 4

Step-by-Step Solution

Answer

(a) 1 and 4

Step 1: Identify Lone Pair and Bonded Pair

For each species, determine the electronic configuration of the central atom and use the VSEPR theory to identify the number of lone pairs and bonded pairs.1. **For** \( \mathrm{XeF}_4 \): Xenon (Xe) has 8 valence electrons. It forms 4 bonds with fluorine, so it uses 4 electrons for bonding. This leaves 2 lone pairs.2. **For** \( \mathrm{SF}_4 \): Sulfur (S) has 6 valence electrons. It forms 4 bonds with fluorine, using 4 electrons and leaving 1 lone pair.3. **For** \(\left[\mathrm{NiCl}_4\right]^{2-}\): Nickel (Ni) is typically in a +2 oxidation state in this complex, with a full \(3d^8\) configuration, leading to tetrahedral geometry.4. **For** \(\left[\mathrm{PdCl}_4\right]^{2-}\): Palladium (Pd) has similar electronic configuration to nickel but with strong field ligands leading to square planar geometry.

Step 2: VSEPR Theory Application

Using VSEPR theory, deduce the molecular geometry for each compound.- **\( \mathrm{XeF}_4 \)**: With 4 bonded pairs and 2 lone pairs, the electron geometry is octahedral, but the molecular geometry becomes square planar.- **\( \mathrm{SF}_4 \)**: With 4 bonded pairs and 1 lone pair, the geometry is a see-saw.- **\(\left[\mathrm{NiCl}_4\right]^{2-}\)**: Typically tetrahedral due to \(d^8\) configuration.- **\(\left[\mathrm{PdCl}_4\right]^{2-}\)**: Square planar due to strong field ligands and splitting of \(d\)-orbitals.

Step 3: Determine the Square Planar Species

Compare the geometries identified. **\( \mathrm{XeF}_{4} \)** and **\(\left[\mathrm{PdCl}_{4}\right]^{2-}\)** are square planar species.

Key Concepts

VSEPR Theory

VSEPR stands for Valence Shell Electron Pair Repulsion theory. This fundamental concept in chemistry helps us predict the shapes of molecules based on the idea that electron pairs around a central atom will repel each other. This repulsion causes them to arrange themselves as far apart as possible in three-dimensional space.

• Lone pairs: Electrons not involved in bonding
• Bonded pairs: Electrons shared between atoms to form a bond

For example, in a compound like \( \mathrm{XeF}_4 \), there are four bonded pairs and two lone pairs around the xenon atom. According to VSEPR theory, the molecule will adopt an arrangement where these electron pairs are positioned to minimize repulsion, leading to a specific geometry. By knowing the number of lone and bonded pairs, we can predict the shape of many molecules.

Square Planar Geometry

Square planar geometry is a type of molecular shape that can be predicted using the VSEPR model when the central atom is surrounded by four bonded atoms and some lone pairs. The central atom and four surrounding atoms form a square shape in a plane.
This configuration often occurs in transition metal complexes such as \( \mathrm{PdCl}_4^{2-} \). In this compound, palladium is at the center of the square formed by four chlorine atoms. Here, strong field ligands like chloride ions influence the d-orbitals of the metal, stabilizing the square planar shape.
In general, square planar geometries are favored in molecules where octahedral electron geometry is altered due to symmetrically placed lone pairs or specific ligand interactions that reduce electronic repulsions.

Lone Pairs and Bonded Pairs

Understanding the difference between lone pairs and bonded pairs is crucial in predicting molecular geometry. Lone pairs refer to pairs of electrons that are not shared with another atom in a bond. Bonded pairs, on the other hand, are shared between two atoms.
In \( \mathrm{XeF}_4 \), the xenon atom has two lone pairs and forms four bonds with fluorine atoms. These lone pairs are placed opposite each other, reducing repulsion and bringing the bonded pairs into a planar arrangement.
The number of lone and bonded pairs directly affects the shape of the molecule. Lone pairs tend to take up more space than bonded pairs, often causing deviations from ideal geometric arrangements. By assessing the lone and bonded pairs, one can apply these concepts to predict whether a molecule like \( \mathrm{XeF}_4 \) or \( \mathrm{PdCl}_4^{2-} \) will exhibit square planar geometry.