1.7:
VSEPR Theory
The valence shell electron-pair repulsion, or VSEPR, theory assumes that electron groups involved in single bonds, multiple bonds, or lone pairs repel each other and try to stay at the maximum possible distance from each other.
The molecular geometry is dictated by the arrangement of various electron groups around the central atom.
There are five basic molecular shapes: linear for two bonding electron groups, trigonal planar for three, tetrahedral for four, trigonal bipyramidal for five, and octahedral for six.
While predicting the molecular geometry, remember that lone pair–lone pair repulsions are greater than lone pair–bonding pair and bonding pair–bonding pair repulsions.
Consider the Lewis structures of methane, ammonia, and water. In each, the central atom is surrounded by four electron groups.
In methane, the four bonding electron pairs are arranged tetrahedrally with an ideal H–C–H angle of 109.5°.
In ammonia, the nitrogen atom has three bonding pairs and one lone pair.
The lone pair of electrons occupies a larger space than the bonding pairs. This is because a lone pair is bound to only one nucleus, whereas a bonding electron group is shared by two nuclei.
The H–N–H bond angles are smaller than the expected tetrahedral angle of 109.5°, as observed in methane. This compression of the bond angle is attributed to the repulsive force exerted by a lone pair on the adjacent bonding electron groups.
The arrangement of electron pairs is called electron-pair geometry. The molecular geometry describes the arrangement of the atoms, and differs from the electron-pair geometry. The electron-pair geometry for ammonia is tetrahedral, whereas the molecular shape is trigonal pyramidal.
A water molecule also has four electron groups around the central atom. The electron pair geometry is also tetrahedral with two bonding electron groups and two lone pairs.
The greater repulsion exerted by two lone pairs further compresses the H–O–H bond angle in water molecules. It is much smaller than the ideal tetrahedral bond angle, and the molecular geometry is bent.
1.7:
VSEPR Theory
Valence shell electron-pair repulsion theory (VSEPR theory) enables us to predict the molecular structure around a central atom from an examination of the number of bonds and lone electron pairs in its Lewis structure. The VSEPR model assumes that electron pairs in the valence shell of a central atom will adopt an arrangement that minimizes repulsions between these electron pairs by maximizing the distance between them. The electrons in the valence shell of a central atom form either bonding pairs, located primarily between bonded atoms, or lone pairs.
Two regions of electron density in a molecule are oriented linearly on opposite sides of the central atom to minimize repulsions. Similarly, three electron groups are arranged in the trigonal planar geometry, four electron groups form a tetrahedron, five electron groups prefer the trigonal bipyramidal geometry, while six such groups are oriented octahedrally.
It is important to note that the electron-pair geometry around a central atom is not always the same as its molecular structure. The electron-pair geometry describes all the regions where electrons are located in a molecule, in bonds as well as lone pairs. The molecular structure describes the location of the atoms in a molecule, not the electrons. Thus, the electron-pair geometry is the same as the molecular structure only when there are no lone electron pairs around the central atom.
A lone pair of electrons occupies a larger space than a bonding pair; this is because a lone pair is bound to only one nucleus, whereas a bonding electron group is shared by two nuclei. Thus, lone pair–lone pair repulsions are greater than lone pair–bonding pair and bonding pair–bonding pair repulsions.
According to VSEPR theory, the terminal atom locations are equivalent in the linear, trigonal planar, tetrahedral, and octahedral electron-pair geometries. Thus, any of the positions can be occupied by a single lone pair. In the trigonal bipyramidal geometry, however, the two axial positions are distinct from the three equatorial positions. Here, the equatorial position has more space available because of the 120° bond angles and is preferred by the larger lone pairs. Similarly, when two lone pairs and four bonding pairs are arranged octahedrally around a central atom, the two lone pairs are 180° apart, resulting in a square planar molecular structure.
This text has been adapted from Openstax, Chemistry 2e, Section 7.6 Molecular Structure and Polarity.