Problem 206 - A Complete Resource Book in Chemistry for JEE Main

Question

The molecular shapes of \(\mathrm{SF}_{4}, \mathrm{CF}_{4}\), and \(\mathrm{XeF}_{4}\) are [2005] (a) different with 1,0 and 2 lone pairs of electrons on the central atoms respectively (b) different with 0,1 and 2 lone pairs of electrons on the central atoms respectively (c) the same with 1,1 and 1 lone pair of electrons on the central atoms respectively (d) the same with 2,0 and 1 lone pair of electrons on the central atoms respectively

Step-by-Step Solution

Answer

(a) different with 1, 0, and 2 lone pairs of electrons on the central atoms, respectively.

Step 1: Determine the Shape of SF4

The molecule \( \text{SF}_4 \) consists of sulfur (S), which has six valence electrons. With four fluorine (F) atoms bonded to it, this leaves 2 electrons as a lone pair. The presence of one lone pair gives \( \text{SF}_4 \) a see-saw shape, which is associated with 1 lone pair on the central atom.

Step 2: Determine the Shape of CF4

Carbon (C) in \( \text{CF}_4 \) has four valence electrons and forms four single bonds with fluorine atoms. Therefore, there are no lone pairs of electrons on the carbon atom. The molecular geometry of \( \text{CF}_4 \) is tetrahedral, characterized by 0 lone pairs on the central atom.

Step 3: Determine the Shape of XeF4

Xenon (Xe) in \( \text{XeF}_4 \) has eight valence electrons. It forms four bonds with fluorine atoms and has 4 remaining electrons, which constitute two lone pairs. This leads to a square planar shape, with 2 lone pairs on the central atom.

Step 4: Conclusion on Molecular Shapes

For \( \text{SF}_4 \), \( \text{CF}_4 \), and \( \text{XeF}_4 \), the molecular shapes are different, with 1, 0, and 2 lone pairs of electrons on the central atoms, respectively.

Key Concepts

Lone Pairs of Electrons

Lone pairs of electrons are the electron pairs that do not participate in bonding and remain on the central atom in a molecule. They affect the molecular geometry as they occupy space around the atom, similar to bonding pairs. However, they exert a greater repulsive force due to their negative charge, influencing the shape more significantly.
In the case of sulfur tetrafluoride ( SF_4 ), sulfur has one lone pair. This lone pair is responsible for altering the ideal geometrical arrangement, resulting in a see-saw shape. In contrast, xenon tetrafluoride ( XeF_4 ) has xenon with two lone pairs, contributing to its square planar geometry.
Lone pairs are a crucial component in understanding the actual shape of a molecule, as they provide insight into the molecule's three-dimensional structure by repelling bonding pairs away.

See-Saw Shape

The see-saw shape is a type of molecular geometry that is observed when a molecule has a distorted tetrahedral arrangement due to the presence of a lone pair. This can be seen in molecules like sulfur tetrafluoride ( SF_4 ).
In SF_4 , the central sulfur atom forms bonds with four fluorine atoms and retains one lone pair. The lone pair occupies an equatorial position, pushing the bonded fluorines slightly closer together, creating an asymmetric geometry that resembles a lightweight see-saw.
This unique geometry is recognized in the VSEPR (Valence Shell Electron Pair Repulsion) theory as one of the possible deviations from the classic shapes, contributing to varied molecular properties such as polarity and reactivity.

Tetrahedral Geometry

Tetrahedral geometry is prevalent in molecules where a central atom is symmetrically bonded to four other atoms with no lone pairs affecting the arrangement. A perfect example is carbon tetrafluoride ( CF_4 ).
In CF_4 , the carbon atom is at the center, surrounded equally by four fluorine atoms. With no lone pairs on the central carbon, the bond angles are perfectly equal, approximately 109.5 degrees, leading to a symmetric structure.
This structure is significant as it provides stability and influences chemical properties. Such a configuration minimizes repulsion between electron pairs, resulting in an efficient energy distribution and a nonpolar molecule in isotropic environments.

Square Planar Shape

The square planar shape is another distinct molecular geometry, characterized by a central atom surrounded by four atoms in a single plane, with two lone pairs above and below the plane. This is evident in xenon tetrafluoride ( XeF_4 ).
In XeF_4 , xenon forms bonds with four fluorine atoms and holds two lone pairs. These lone pairs on the opposite sides balance the overall structure, stabilizing the geometry into a square planar shape in line with the electron pair repulsion theory.
This type of geometry often arises in molecules with transition metals or heavy p-block elements and affects the compound's physical and chemical characteristics, including hybridization, reactivity, and magnetic properties.

Valence Electrons

Valence electrons are the electrons located in the outermost electron shell of an atom. They play a crucial role in chemical reactions as they are involved in forming bonds between atoms.
For sulfur in SF_4 , there are six valence electrons. Four of these are used to form bonds with fluorine atoms, and two remain as a lone pair. For xenon in XeF_4 , there are eight valence electrons, with four involved in bonding and four forming two lone pairs.
Understanding valence electrons allows chemists to predict how atoms will bond and interact. The number and arrangement of these electrons are foundational in determining a molecule's three-dimensional geometry, its reactivity, and overall stability.