The pair of species having identical shapes for molecules of both species is: (a) \(\mathrm{PF}_{5}, \mathrm{IF}_{5}\) (b) \(\mathrm{XeF}_{2}, \mathrm{BeCl}_{2}\) (c) \(\mathrm{CCl}_{4}, \mathrm{SiCl}_{4}\) (d) \(\mathrm{Bx}_{3}, \mathrm{PF}_{3}\)

When trying to understand molecular shapes, the Valence Shell Electron Pair Repulsion (VSEPR) theory serves as the guiding principle. The theory posits that electron pairs in the valence shell of an atom will arrange themselves as far apart as possible to minimize repulsion between negatively charged particles.

This means that the geometrical shape of a molecule largely depends on the repulsion between lone pair - lone pair, lone pair - bonding pair, and bonding pair - bonding pair of electrons around the central atom. The VSEPR theory can help predict the three-dimensional structure of molecules, which is critical for understanding their chemical behavior and properties.

Electronic geometry refers to the spatial arrangement of all electron pairs (bonding and non-bonding) around the central atom. Visualizing it is key to understanding a molecule's structure. It provides the skeleton that defines the molecule's ultimate shape. Think of it as the frame of a tent before the canvas is draped over it.

The common electronic geometries are linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral shapes. Each geometric setup relates to the number of 'regions of high electron density' or electron pairs around the central atom, which are based on the ideal shapes that minimize electron pair repulsion.

To grasp molecular geometry, it's crucial to distinguish between bonding pairs of electrons, which are shared by atoms to make bonds, and lone pairs, which belong to a single atom without contributing to bonding.

• Bonding Pairs: They are sensed by multiple nuclei and thus hold atoms together, affecting molecular shape by pulling connected atoms toward each other.
• Lone Pairs: Being only under the influence of one nucleus, they tend to take up more space than bonding pairs, often distorting the geometry of a molecule away from that predicted by considering bonding pairs only.

Understanding their influence is crucial because lone pairs can dramatically alter the molecular shape from its basic electronic geometry.

The trigonal bipyramidal shape arises in molecules with five regions of electron density, where three are arranged in a plane (equatorial positions) at 120-degree angles, and two occupy positions above and below this plane (axial positions) at 90-degree angles to the plane.

This structure is interesting because it demonstrates the concept of electron pair repulsions well: the equatorial positions are farther apart, so they're commonly occupied by the larger lone pairs, if present, while the smaller bonding pairs generally occupy axial positions. PF₅ is a canonical example of a molecule that adopts this geometry due to its five bonding pairs and absence of lone pairs.

The tetrahedral shape is a common molecular geometry for molecules with a central atom surrounded by four regions of electron density — the bonding pairs. Envision it as a pyramid with a triangular base; each corner represents a region of electron density. The angles between each electron density region is approximately 109.5 degrees, which is the value that minimizes repulsion between these regions.

Since all positions are equivalent, there's no distinction between 'equatorial' and 'axial' as there is in the trigonal bipyramidal geometry. Molecules like CCl₄ and SiCl₄ exhibit the tetrahedral shape, demonstrating the theory when only bonding pairs are involved and no lone pairs are present to complicate the arrangement.