VSEPR Theory and the Effect of Lone Pairs

Predicting molecular geometry has more steps if the central atom has both bonding pairs and lone pairs of electrons. 

The nitrogen atom in ammonia has four electron groups arranged in a tetrahedral fashion: 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.

Lone pair-lone pair repulsions are greater than lone pair-bonding pair and bonding pair-bonding pair repulsions.

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.

The effect of lone pairs on molecular geometry is evident by looking at methane, ammonia, and water, all with four electron groups. The bond angle becomes smaller as the number of lone pairs increases.

In VSEPR theory, the terminal atom locations are structurally equivalent in each of the linear, trigonal planar, and tetrahedral electron-pair geometries. A lone pair may replace any of these atoms.

However, for trigonal bipyramidal electron-pair geometries, there are two distinct positions, the axial position and the equatorial position, which could be replaced by a lone pair.

The axial position is surrounded by bond angles of 90°, whereas the equatorial position has more space available because of the 120° bond angles. A lone pair prefers an equatorial position.

Consider three examples with trigonal bipyramidal electron pair geometry.

Sulfur tetrafluoride has one lone pair of electrons, and the molecular geometry is seesaw.

Bromine trifluoride has two lone pairs at the equatorial positions, and therefore its molecular geometry is T-shaped.

Xenon difluoride has three lone pairs, all placed equatorially, and the molecule is linear.

When a central atom has six electron groups, the electron-pair geometry is octahedral, as seen in bromine pentafluoride. A lone pair can occupy any position, as they are all equivalent. The molecular geometry is square pyramidal.

When the electron pair geometry is octahedral and a central atom has two lone pairs, for example, in xenon tetrafluoride, the lone pairs occupy opposite sides of the octahedron. The molecular geometry is square planar with minimum lone pair-lone pair repulsions.