Question

The compound(s) with TWO lone pairs of electrons on the central atom is(are) (A) \(\mathrm{BrF}_{5}\) (B) \(\mathrm{ClF}_{3}\) (C) \(\mathrm{XeF}_{4}\) (D) \(\mathrm{SF}_{4}\)

Short answer

\(\text{XeF}_4\) (Option C) is the compound with two lone pairs of electrons on the central atom.

Step-by-step solution

Understanding VSEPR theory

To determine the number of lone pairs on the central atom of a molecule, we use the Valence Shell Electron Pair Repulsion (VSEPR) theory. This theory states that electron pairs around a central atom will arrange themselves as far apart as possible to minimize repulsion. We determine the molecular geometry after predicting the number of bonding and nonbonding electron pairs.

Analyzing each molecule

For each molecule, calculate the total number of valence electrons available and use the VSEPR theory to predict the geometry. For molecules with expanded octets, remember that some elements in the third period and beyond can accommodate more than eight electrons.

Determining lone pairs on central atom in \(\mathrm{BrF}_{5}\)

\(\text{BrF}_5\) has 7 valence electrons from bromine and 5 * 7 from each of the five fluorine atoms, totaling 42 electrons or 21 electron pairs. With five bonds, bromine uses 10 electrons, leaving 11 electron pairs. After using five for bonds, there are 6 pairs left, but only one lone pair on the central atom, as the rest are used for bromine's expanded octet.

Determining lone pairs on central atom in \(\mathrm{ClF}_{3}\)

\(\text{ClF}_3\) has 7 valence electrons from chlorine and 3 * 7 from each fluorine atom, totaling 28 electrons or 14 electron pairs. With three bonds, chlorine uses 6 electrons, leaving 8 electron pairs. After using three for bonds, there should be 5 pairs left, but only 2 of these are lone pairs on the central atom, while the rest participate in an expanded octet.

Determining lone pairs on central atom in \(\mathrm{XeF}_{4}\)

\(\text{XeF}_4\) has 8 valence electrons from xenon and 4 * 7 from each fluorine, totaling 36 electrons or 18 electron pairs. With four bonds, xenon uses 8 electrons, leaving 10 electrons pairs. After using four for bonds, there are 6 pairs left, and 2 of these are lone pairs on the central atom, while four are used in bonding.

Determining lone pairs on central atom in \(\mathrm{SF}_{4}\)

\(\text{SF}_4\) has 6 valence electrons from sulfur and 4 * 7 from each fluorine, totaling 34 electrons or 17 electron pairs. With four bonds, sulfur uses 8 electrons, leaving 9 electron pairs. After using four for bonds, there are 5 pairs left, but only one of these is a lone pair on the central atom, with the rest as an expanded octet.

Selection of the correct compound

From the analysis above, it's evident that the compound with exactly two lone pairs on the central atom is \(\text{XeF}_4\) (Option C).

Key concepts

Molecular Geometry

Understanding the concept of molecular geometry is crucial for predicting how molecules will interact with each other and with other substances. Molecular geometry refers to the three-dimensional arrangement of the atoms that constitute a molecule. It is determined by the spatial distribution of both bonding electron pairs, which form the bonds between atoms, and nonbonding or lone pairs of electrons, which do not participate in bonding but still influence the shape of the molecule.

According to VSEPR theory, the geometrical shape of a molecule is designed to minimize the electron pair repulsion between the bonding and nonbonding pairs in the valence shell of the central atom. These shapes can include linear, bent, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral, and others, depending upon the number of electron pairs repelling each other. An exercise to ascertain a molecule's geometry would be, for example, to determine which compounds have two lone pairs of electrons on the central atom, impacting the molecule's shape and chemical properties.

Lone Pairs of Electrons

The concept of lone pairs of electrons is an important aspect of molecular geometry and VSEPR theory. Lone pairs refer to pairs of valence electrons that are not shared with another atom and do not participate in bonding. These nonbonding electrons can significantly influence the shape and angles within a molecule due to their presence.

Lone pairs occupy more space around the central atom than bonding pairs because they only need to accommodate the space around one nucleus. Consequently, they tend to push bonded atoms closer together, altering the ideal angles and leading to distinctive geometries for different molecules. For instance, in the exercise, recognizing which compounds have two lone pairs on their central atom helps us predict their molecular geometry. The presence of lone pairs necessitates examining the total valence electrons for each atom and applying VSEPR theory to predict how these pairs will affect the molecule's shape.

Electron Pair Repulsion

The principle of electron pair repulsion is fundamental to understanding the VSEPR theory and, by extension, molecular geometry. Since electrons have similar charges, there is a natural repulsion between them. This theory posits that electron pairs located within the valence shell of an atom, whether they are involved in bonding or not, will arrange themselves as far apart as possible to minimize repulsive forces.

This repulsion affects the position of atoms within the molecule, leading to specific geometries that maximize the distance between valence electron pairs. For example, in a molecule with four pairs of electrons, such as methane (CH4), the electron pairs are positioned at the corners of a tetrahedron, with bond angles of approximately 109.5 degrees. VSEPR theory allows us to predict geometries ranging from simple (linear) to complex (trigonal bipyramidal or octahedral), depending on the number of bonding and lone electron pairs. Exercises that involve electron pair repulsion require students to apply this concept to predict and explain the geometry of various molecules.