Which of the following molecules and ions contain polar bonds? Which of these molecules and ions have dipole moments? (a) ClF \(_{5}\) (b) \(\mathrm{ClO}_{2}^{-}\) (c) \(\mathrm{TeCl}_{4}^{2-}\) (d) \(\mathrm{PCl}_{3}\) (e) \(\operatorname{Se} \mathrm{F}_{4}\) (f) \(\mathrm{PH}_{2}^{-}\) (g) \(\mathrm{XeF}_{2}\)

Electronegativity is a measure of an atom's ability to attract and hold onto electrons within a bond. It plays a pivotal role in determining the polarity of bonds in molecules. Atoms with high electronegativity, like fluorine and oxygen, will pull electrons more closely towards themselves in a covalent bond, creating an uneven distribution of charge.

This disparity in electronegativity between atoms, such as in a chlorine-fluorine bond (Cl-F), can cause the bond to become polar. A polar bond is characterized by a partial negative charge near the more electronegative atom and a partial positive charge near the less electronegative atom. The larger the difference in electronegativity, the more polar the bond.

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. It is a critical factor in determining whether a molecule has an overall dipole moment. When atoms in a molecule share electrons unevenly due to differences in electronegativity, the geometric configuration dictates how these polar bonds align.

For example, molecules like \( \mathrm{PCl}_3 \) have trigonal pyramidal geometry, which leads to a lack of symmetry and results in a net dipole moment. On the other hand, when the molecular shape is symmetrical, as in the case of \( \mathrm{XeF}_2 \), the polar bonds may evenly oppose each other, leading to dipole cancellation.

Valence Shell Electron Pair Repulsion (VSEPR) theory is used to predict the molecular geometry of a molecule based on the repulsion between electron pairs in the valence shell of the central atom. According to VSEPR theory, electron pairs—whether they are in bond pairs or lone pairs—will arrange themselves as far apart as possible to minimize repulsion.

This arrangement determines the molecule’s shape and has a direct impact on whether individual bond polarities in a molecule sum to create an overall dipole moment. VSEPR theory helps explain why \( \mathrm{ClF}_5 \) adopts a square pyramidal shape leading to a net dipole, whereas other molecules might have shapes that cause their bond dipoles to cancel.

Dipole cancellation occurs in molecules where polar bonds are symmetrically arranged so that the individual bond dipoles negate each other's effects, resulting in a molecule with no overall dipole moment. The key to dipole cancellation lies in the molecule's geometry.

Symmetrical molecules, like \( \mathrm{XeF}_2 \), have a linear arrangement of polar \( \mathrm{Xe-F} \) bonds that cancel out due to the molecule’s 180-degree symmetry. In contrast, molecules with asymmetrical shapes, such as \( \mathrm{SeF}_4 \), do not experience dipole cancellation and thereby exhibit dipole moments. These concepts are essential when predicting molecular properties, such as boiling points, solubility, and reactivity.