Question

Antimony pentafluoride, \(\mathrm{SbF}_{5}\), reacts with \(\mathrm{XeF}_{4}\) and \(\mathrm{XeF}_{6}\) to form ionic compounds, \(\mathrm{XeF}_{3}^{+} \mathrm{SbF}_{6}^{-}\) and \(\mathrm{XeF}_{5}^{+} \mathrm{SbF}_{6}^{-}\). Describe the geometries of the cations and anion in these two compounds.

Short answer

The \(\mathrm{XeF}_{3}^{+}\) cation has a trigonal pyramidal geometry, the \(\mathrm{XeF}_{5}^{+}\) cation has a square pyramidal geometry, and the \(\mathrm{SbF}_{6}^{-}\) anion has an octahedral geometry.

Step-by-step solution

Determine the Electron Domains

In the first step, find the number of electron domains around the central atom for the given molecules. For \(\mathrm{XeF}_{3}^{+}\), the Xe atom is surrounded by 3 bond pairs and one lone pair. Meanwhile, the \(\mathrm{XeF}_{5}^{+}\) cation has 5 bond pairs and no lone pairs around the central atom. In the anion \(\mathrm{SbF}_{6}^{-}\), Sb is surrounded by 6 bond pairs and no lone pairs.

Identify Molecular Geometries

Now, knowing the number of electron domains, you can identify the geometries through the VSEPR model. The \(\mathrm{XeF}_{3}^{+}\) cation, with 3 bond pairs and one lone pair, should have a trigonal pyramidal shape. The \(\mathrm{XeF}_{5}^{+}\) cation, with 5 bond pairs and no lone pair, should have a square pyramidal geometry. The \(\mathrm{SbF}_{6}^{-}\) anion, having 6 bond pairs and no lone pair, should have an octahedral geometry.

Double Check

Finally check your understanding and re-interpretation of the problem to ensure that it fits with your understanding of the concepts of molecular geometry and the VSEPR model.

Key concepts

Molecular Geometry

Molecular geometry is a vital concept in understanding the shape and structure of molecules. It refers to the three-dimensional arrangement of atoms within a molecule. The molecular geometry is crucial because it can influence the physical and chemical properties of the substance. It is primarily determined by the number of electron domains, which includes both bonding and non-bonding electron pairs surrounding the central atom.

In the compound, \[\text{XeF}_{3}^{+}\]cation, the central atom xenon is surrounded by three bonding pairs and one lone pair of electrons. The presence of this lone pair results in a trigonal pyramidal shape due to the repulsion between the lone pair and the bond pairs. On the other hand, \[\text{XeF}_{5}^{+}\]exhibits a square pyramidal structure because it has five bonding pairs and no lone pairs.

For the anionic compound \[\text{SbF}_{6}^{-}\], the antimony has six bonding pairs, which leads to an octahedral geometry, a very symmetrical and common formation when six domains are present around a central atom.

Understanding the shapes of these ions with the VSEPR theory helps predict their interactions and behavior during reactions.

Electron Domains

Electron domains refer to the different regions around an atom where electrons are likely to be found. These can include single bonds, double bonds, triple bonds, and lone pairs of electrons. The count of these domains is central in determining the molecular geometry of a compound.

In a typical molecule, you consider each of these domains to exert repulsion forces, pushing each other as far apart as possible.

• For \[\text{XeF}_{3}^{+}\], there are four domains around xenon: three from the bonds and one from the lone pair.
• \[\text{XeF}_{5}^{+}\]has five domains, all originating from bond pairs with no lone pair interference.
• In\[\text{SbF}_{6}^{-}\], all six domains are created by bond pairs.

This framework of quantifying electron domains provides a systematic approach to anticipating the molecular geometry using VSEPR, which predicts how such electron domains influence molecular shape by minimizing repulsive interactions.

Cations and Anions

Cations and anions are charged species formed from atoms or molecules by gaining or losing electrons. Cations are positively charged because they have fewer electrons than protons, while anions bear a negative charge due to an excess of electrons.

In the compounds formed by \[\text{SbF}_{5}\]and xenon fluorides, we see cations \[\text{XeF}_{3}^{+}\]and \[\text{XeF}_{5}^{+}\]specifically acquiring their positive charges by donating electrons, leaving behind a deficit. These cations contribute to the total charge balance in the compound alongside the negatively charged anion \[\text{SbF}_{6}^{-}\], which has gained an electron making it negatively charged.

The interaction between these ions forms ionic compounds, where the overall charge is balanced, creating stable structures. The knowledge of these charges and their distribution helps in understanding the formation and stability of different ionic compounds. These characteristics also impact the resultant geometric confirmations and compound properties.