Why does ammonia have a different molecular shape than methane despite both having four electron groups?

Ammonia has a lone pair of electrons on the central nitrogen atom, while methane has only bonding pairs. The lone pair occupies more space than bonding pairs because it is bound to only one nucleus rather than shared by two. This lone pair-bonding pair repulsion compresses the H-N-H bond angles below the ideal tetrahedral angle of 109.5°, creating a trigonal pyramidal molecular shape instead of tetrahedral.

How does electron-pair geometry differ from molecular geometry?

Electron-pair geometry describes the arrangement of all electron groups around a central atom, including both bonding pairs and lone pairs. Molecular geometry describes only the positions of atoms, excluding lone pairs. When no lone pairs are present, these geometries are identical. However, when lone pairs exist, electron-pair geometry and molecular geometry differ, as lone pairs occupy space but are not part of the molecular structure.

What is the order of electron-pair repulsions in VSEPR theory?

According to VSEPR theory, electron-pair repulsions follow this order from greatest to least: lone pair-lone pair repulsion is strongest, followed by lone pair-bonding pair repulsion, and bonding pair-bonding pair repulsion is weakest. This hierarchy explains why lone pairs occupy larger regions of space than bonding electrons and why their placement significantly affects molecular geometry and bond angles.

Why do lone pairs prefer equatorial positions in trigonal bipyramidal geometry?

In trigonal bipyramidal electron-pair geometry, axial positions are surrounded by 90° bond angles, while equatorial positions have more space available due to 120° bond angles. Lone pairs prefer equatorial positions because these spacious regions can more easily accommodate the larger lone pairs, minimizing repulsive forces and stabilizing the molecular structure.

How does water's molecular geometry result from its electron groups and lone pairs?

Water has four electron groups around the central oxygen atom: two bonding pairs and two lone pairs. The electron-pair geometry is tetrahedral, but the two lone pairs exert greater repulsion than bonding pairs, further compressing the H-O-H bond angle well below the ideal tetrahedral angle of 109.5°. This results in a bent molecular geometry with significantly reduced bond angles.

What is the molecular geometry of xenon difluoride, and why does it have that shape?

Xenon difluoride has a trigonal bipyramidal electron-pair geometry with three lone pairs and two bonding pairs. All three lone pairs occupy equatorial positions to minimize repulsion, leaving only the two fluorine atoms in axial positions. This arrangement results in a linear molecular geometry with the two Xe-F bonds aligned at 180°.

Why is xenon tetrafluoride square planar rather than octahedral?

Xenon tetrafluoride has an octahedral electron-pair geometry with six electron groups: four bonding pairs and two lone pairs. The two lone pairs occupy opposite sides of the octahedron, positioned 180° apart. This arrangement minimizes lone pair-lone pair repulsion while leaving the four fluorine atoms in a square planar configuration around the central xenon atom.

10.2

VSEPR Theory and the Effect of Lone Pairs

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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.