Number of lone pairs of electrons on Xe atoms in \(\mathrm{XeF}_{2}, \mathrm{XeF}_{4}\) and \(\mathrm{XeF}_{6}\) molecules are respectively (a) 2, 3, 1 (b) \(3,2,1\) (c) \(3,2,0\) (d) \(4,3,2\)

Understanding electron pair geometry is vital in predicting the arrangement of electron pairs around a central atom, according to the Valence Shell Electron Pair Repulsion (VSEPR) theory. This theory posits that electron pairs, both bonding and non-bonding (lone pairs), will arrange themselves as far apart as possible within the valence shell to minimize repulsion.

Imagine the central atom as a sphere with electron pairs spreading out in different directions. The geometry emerges from the need to maximize distance between these negatively charged regions. For example, if there are two electron pairs, they will be on opposite sides, forming a linear geometry. With three pairs, they form a trigonal planar shape, spreading out to 120-degree angles. Four pairs morph into a tetrahedral arrangement, suffering the least repulsion at 109.5-degree angles.

Now, let's apply this to Xenon fluorides. With different numbers of bonding and lone pairs, they adopt different electron pair geometries: linear for XeF2 with 3 lone and 2 bonding pairs, square planar for XeF4 with 2 lone and 4 bonding pairs, and pentagonal bipyramidal for XeF6 with 1 lone and 6 bonding pairs.

Molecular shape, often also referred to as molecular geometry, is derived by considering the spatial arrangements of atoms bonded to the central atom. With the VSEPR theory as our guide, we predict these shapes by factoring in both bonding pairs and lone pairs. However, only bonding pairs dictate the actual shape of the molecule.

For instance, though xenon difluoride (XeF2) adopts a linear electron pair geometry due to five pairs (three lone and two bonding), its molecular shape is described merely as linear because only the bonding pairs shape it. Similarly, xenon tetrafluoride (XeF4) has a square planar molecular shape determined by its four bonding pairs, ignoring the two lone pairs for shape considerations.

In more complex examples like xenon hexafluoride (XeF6), the single lone pair affects the distribution of the six bonding pairs but the molecular shape still focuses on how these bonding pairs are arranged, resulting in a distorted octahedral shape. This understanding helps chemists predict and explain molecular properties and reactions.

Valence electrons are the outermost electrons of an atom and play a pivotal role in chemical bonding and reactivity. These are the electrons involved when atoms interact to form chemical bonds. In noble gases like xenon (Xe), the valence electrons initially fill the shell completely, making the atom stable and chemically inert.

However, under certain conditions, xenon can form compounds by using these usually inactive valence electrons. For example, in forming XeF2, XeF4, and XeF6, xenon uses its valence electrons to form bonds with fluorine atoms. Each bond formed uses a pair of valence electrons: one from xenon and one from fluorine. Despite xenon's reluctance to partake in bonding due to its filled valence shell, it can do so by accommodating more electrons in its outer shell, which expands its capacity and exhibits its unique chemistry.

In our exercise, by understanding the count of valence electrons and how they are distributed as bonding and non-bonding pairs, we correctly identified the number of lone pairs on xenon in its fluoride compounds, demonstrating the balance between filled and expanded valence shells in xenon chemistry.