VSEPR theory can indeed predict the shapes of ions, including polyatomic ions like sulfate (SO₄²⁻). The approach is similar to that used for neutral molecules: considering the repulsion between electron pairs (both bonding and lone pairs) around the central atom. In the case of sulfate, the central sulfur atom is surrounded by four oxygen atoms with double bonds and no lone pairs. According to VSEPR theory, these four regions of electron density (the double bonds) will repel each other equally and arrange themselves as far apart as possible. This arrangement leads to a tetrahedral geometry for the sulfate ion. Each S-O bond is equivalent, and the bond angles are approximately 109.5°, typical of a tetrahedral structure. This example demonstrates how VSEPR theory extends beyond simple molecules to complex ions, using the same principles of electron pair repulsion to predict molecular geometry.

VSEPR theory accounts for multiple bonds (double or triple bonds) in a molecule by considering them as a single electron pair region. Despite being treated as one electron pair for shape prediction, multiple bonds exert a greater repulsive force than single bonds. This increased repulsion can affect the molecule's shape, especially when combined with single bonds or lone pairs. For instance, in carbon dioxide (CO₂), the two double bonds on the central carbon atom are treated as two regions of electron density. These double bonds repel each other more strongly than single bonds would, maintaining a linear shape with a bond angle of 180°. Similarly, in acetylene (C₂H₂), the triple bond between the two carbon atoms is considered as one electron pair region, but its higher electron density ensures a linear configuration. Thus, while multiple bonds are treated as single electron pair regions in VSEPR theory, their higher electron density and repulsion significantly influence the overall shape and bond angles in the molecule.

Hybridization is a concept that integrates well with VSEPR theory to explain molecular shapes. It involves the mixing of atomic orbitals to form new hybrid orbitals that can accommodate bonding and lone pairs of electrons. In ethene (C₂H₄), the carbon atoms undergo sp² hybridization. Each carbon atom in C₂H₄ uses three sp² hybrid orbitals for forming sigma bonds with two hydrogen atoms and one carbon atom. The unhybridized p orbital on each carbon overlaps to form a pi bond, contributing to the double bond between the carbon atoms. According to VSEPR theory, the three regions of electron density (two C-H single bonds and one C-C double bond) around each carbon atom repel each other, arranging in a trigonal planar geometry. This leads to a bond angle of approximately 120° around each carbon atom. Hybridization, therefore, complements VSEPR theory by explaining the formation of specific molecular shapes through the rearrangement and merging of atomic orbitals, as evidenced in the planar structure of ethene.

The bond angle in phosphine (PH₃) is less than that in ammonia (NH₃), a phenomenon that VSEPR theory helps to explain. In NH₃, the bond angle is approximately 107°, whereas in PH₃, it's around 93.5°. This difference is attributed to the size of the central atom and the influence of lone pairs. Nitrogen in NH₃ is smaller than phosphorus in PH₃. The smaller size of nitrogen allows the lone pair to exert a stronger repulsion on the bonding pairs, causing them to spread out more, which leads to a larger bond angle. Conversely, the larger central atom in PH₃ means that the lone pair's influence is less pronounced, leading to a smaller bond angle. Additionally, the bonding pairs in PH₃ are less tightly held due to phosphorus’s lower electronegativity compared to nitrogen, resulting in less repulsion among the bonding pairs. Therefore, the geometry of PH₃ is more compressed than NH₃, illustrating how VSEPR theory considers factors like the size of the central atom and electronegativity in determining molecular shapes and bond angles.

VSEPR theory explains the bond angles in water (H₂O) and oxygen difluoride (OF₂) through the concept of electron pair repulsion. Both molecules have a bent shape due to the two lone pairs on the oxygen atom. In H₂O, the bond angle is approximately 104.5°, while in OF₂, it's slightly larger, around 103°. This difference arises from the nature of the atoms bonded to oxygen. In water, the hydrogen atoms are less electronegative compared to fluorine in OF₂. Less electronegative atoms like hydrogen exert less repulsion on the bonding pairs, leading to a smaller bond angle. In contrast, fluorine, being highly electronegative, exerts a stronger repulsion on the bonding pairs, slightly increasing the bond angle. Furthermore, the larger size of fluorine atoms compared to hydrogen contributes to the difference in bond angles. The VSEPR theory thus accounts for variations in molecular geometries based on the types of atoms involved and their respective electronegativities, as well as the size and number of lone and bonding electron pairs.