Question

Among the following pairs, the one in which the two species are not isostructural is (a) \(\mathrm{SiF}_{4}\) and \(\mathrm{SF}_{4}\) (b) \(\mathrm{IO}_{3}^{-}\)and \(\mathrm{XeO}_{3}\) (c) \(\mathrm{BH}_{4}^{-}\)and \(\mathrm{NH}_{4}^{+}\) (d) \(\mathrm{PF}_{6}^{-}\)and \(\mathrm{SF}_{6}\)

Short answer

The non-isostructural pair is (a) \(\mathrm{SiF}_{4}\) and \(\mathrm{SF}_{4}\).

Step-by-step solution

Understand Isostructurality

Two compounds are isostructural if they have the same structure or geometry, meaning they have the same number of atoms arranged in a similar way.

Analyze the Geometry of \(\mathrm{SiF}_{4}\) and \(\mathrm{SF}_{4}\)

\(\mathrm{SiF}_{4}\) is tetrahedral in shape as silicon is sp3 hybridized. \(\mathrm{SF}_{4}\) has a trigonal bipyramidal structure with one lone pair on sulfur, leading to a seesaw shape.

Analyze the Geometry of \(\mathrm{IO}_{3}^{-}\) and \(\mathrm{XeO}_{3}\)

Both \(\mathrm{IO}_{3}^{-}\) and \(\mathrm{XeO}_{3}\) have a pyramidal shape with a lone pair on iodine in \(\mathrm{IO}_{3}^{-}\). Each has a central atom with sp3 hybridization.

Analyze the Geometry of \(\mathrm{BH}_{4}^{-}\) and \(\mathrm{NH}_{4}^{+}\)

Both \(\mathrm{BH}_{4}^{-}\) and \(\mathrm{NH}_{4}^{+}\) have a tetrahedral structure because both the boron and nitrogen atoms are surrounded by four atoms, leading to sp3 hybridization.

Analyze the Geometry of \(\mathrm{PF}_{6}^{-}\) and \(\mathrm{SF}_{6}\)

Both \(\mathrm{PF}_{6}^{-}\) and \(\mathrm{SF}_{6}\) are octahedral because phosphorus and sulfur are surrounded by six fluorine atoms each, indicating sp3d2 hybridization.

Determine the Non-Isostructural Pair

Comparing all pairs, only \(\mathrm{SiF}_{4}\) and \(\mathrm{SF}_{4}\) have different structures: tetrahedral versus seesaw. Hence, they are not isostructural.

Key concepts

Molecular geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule, and it's a fundamental aspect of understanding chemical structures. The shape of a molecule greatly influences its physical and chemical properties, such as polarities and reactivities. Understanding molecular geometry requires considering each atom's position relative to others and the angles between bonding pairs. For example, consider methane (CH_4"). It exhibits a tetrahedral geometry where the carbon atom is centrally located and bonded to four hydrogen atoms, each 109.5 degrees apart, forming a three-dimensional shape.

Different molecular geometries arise primarily due to the presence of lone pairs and bond pairs around central atoms. For instance, if a molecule includes lone electron pairs, as seen in water (H_2O"), these pairs can influence the molecular geometry by pushing bonded pairs closer together, changing a tetrahedral arrangement to a bent shape. As a result, molecular geometry provides essential insights into predicting the behavior and interaction of molecules in different environments.

Hybridization in chemistry

Hybridization in chemistry explains how atomic orbitals mix to form new, hybrid orbitals that accommodate electron pairs when atoms form covalent bonds. This concept helps us describe and predict the molecular geometry. For example, when carbon in methane (CH_4") forms bonds, its one s and three p orbitals hybridize to form four equivalent sp3 hybrid orbitals. These orbitals arrange themselves to minimize electron pair repulsion, giving methane its tetrahedral geometry.

Different types of hybridization lead to different geometries:

• sp Hybridization: Forms linear geometry as seen in acetylene (C_2H_2").
• sp2 Hybridization: Results in trigonal planar geometry, typical in ethylene (C_2H_4").
• sp3 Hybridization: Creates tetrahedral geometry as observed in methane.
• sp3d and sp3d2 Hybridization: Lead to trigonal bipyramidal and octahedral geometries, respectively, evident in molecules like phosphorus pentafluoride (PF_5") and sulfur hexafluoride (SF_6").

Hybridization provides a comprehensive framework for understanding how orbital mixing influences molecular shape and solves the discrepancy between atomic orbital availability and molecular geometry.

Tetrahedral and octahedral structures

Tetrahedral and octahedral structures are specific types of molecular geometries that describe how atoms are arranged around a central atom. These structures are largely determined by the hybridization of the central atom.

**Tetrahedral Structure:** The tetrahedral shape is common in molecules where the central atom forms four bonds, as in molecules like methane (CH_4") and ammonium ion (NH_4^+"). In these cases, the central atom exhibits sp3 hybridization, forming a structure where the bond angle between any two bonds is approximately 109.5 degrees. This geometry is very stable and symmetrical, reducing the overall energy of the molecule.

**Octahedral Structure:** An octahedral structure arises when a central atom forms six bonds. Common examples include sulfur hexafluoride (SF_6") and the hexafluorophosphate ion (PF_6^-", SF_6"). These atoms utilize sp3d2 hybridization, arranging the bonds such that the angles between them are 90 degrees. This configuration provides exceptional symmetry, contributing to the stability and unique properties of octahedrally shaped molecules.

Understanding these structures is crucial because they not only affect how the molecule appears but also its reactivity, polarity, and interactions with other molecules.