10.2:
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.
10.2:
VSEPR Theory and the Effect of Lone Pairs
Effect of Lone Pairs of Electrons on Molecule Geometry
It is important to note that electron-pair geometry around a central atom is not the same thing as its molecular structure. Molecular structure describes the location of the atoms, not the electrons. The geometry that includes all electron pairs is the electron-pair geometry. The electron-pair geometries describe all regions where electrons are located, bonds as well as lone pairs. The structure that includes only the placement of the atoms in the molecule is called the molecular structure. The electron-pair geometries will be the same as the molecular structures when there are no lone electron pairs around the central atom, but they will be different when there are lone pairs present on the central atom.
For example, the methane molecule, CH4, which is the major component of natural gas, has four bonding pairs of electrons around the central carbon atom; the electron-pair geometry is tetrahedral, as is the molecular structure. On the other hand, the ammonia molecule, NH3, also has four electron pairs associated with the nitrogen atom and thus has a tetrahedral electron-pair geometry. One of these regions, however, is a lone pair, which is not included in the molecular structure, and this lone pair influences the shape of the molecule.
Angle Distortions Based on VSEPR Theory
The small distortions from the ideal angles can result from differences in repulsion between various regions of electron density. VSEPR theory predicts these distortions by establishing an order of repulsions and an order of the amount of space occupied by different kinds of electron pairs. The order of electron-pair repulsions from greatest to least repulsion is:
lone pair-lone pair > lone pair-bonding pair > bonding pair-bonding pair
This order of repulsions determines the amount of space occupied by different regions of electrons. A lone pair of electrons occupies a larger region of space than the electrons in a triple bond; in turn, electrons in a triple bond occupy more space than those in a double bond, and so on. The order of sizes from largest to smallest is:
lone pair > triple bond > double bond > single bond
In the ammonia molecule, the three hydrogen atoms attached to the central nitrogen are not arranged in a flat, trigonal planar molecular structure, but rather in a three-dimensional trigonal pyramid with the nitrogen atom at the apex and the three hydrogen atoms forming the base. The ideal bond angles in a trigonal pyramid are based on the tetrahedral electron pair geometry. Again, there are slight deviations from the ideal because lone pairs occupy larger regions of space than do bonding electrons. The H–N–H bond angles in NH3 are slightly smaller than the 109.5° angle in a regular tetrahedron because the lone pair-bonding pair repulsion is greater than the bonding pair-bonding pair repulsion.
According to VSEPR theory, the terminal atom locations are equivalent within the linear, trigonal planar, and tetrahedral electron-pair geometries. It does not matter which X is replaced with a lone pair because the molecules can be rotated to convert positions. For trigonal bipyramidal electron-pair geometries, however, there are two distinct X positions, an axial position (if we hold a model of a trigonal bipyramid by the two axial positions, we have an axis around which we can rotate the model) and an equatorial position (three positions form an equator around the middle of the molecule). The axial position is surrounded by bond angles of 90°, whereas the equatorial position has more space available because of the 120° bond angles. In a trigonal bipyramidal electron-pair geometry, lone pairs always occupy equatorial positions because these more spacious positions can more easily accommodate the larger lone pairs.
When a central atom has two lone electron pairs and four bonding regions, we have an octahedral electron-pair geometry. The two lone pairs are on opposite sides of the octahedron (180° apart), giving a square planar molecular structure that minimizes lone pair-lone pair repulsions.
This text has been adapted from Openstax, Chemistry 2e, Section 7.6: Molecular Structure and Polarity.