Lesson Explainer: VSEPR Chemistry

In this explainer, we will learn how to identify the shape of molecules using VSEPR.

The valence shell electron pair repulsion (VSEPR) model can be used with Lewis-structure diagrams to predict the shapes of different molecules.

Lewis-structure diagrams are used to identify valence electrons in polyatomic molecules, and the VSEPR model is used to understand how these valence electrons are arranged in three-dimensional space. The VSEPR model assumes that there will always be electrostatic repulsion between pairs of valence electrons. The valence electrons will be forced to adopt specific three-dimensional shapes that minimize high-energy repulsive electrostatic interactions.

The following figure shows the Lewis-structures of some simple polyatomic molecules and ions. The red color is used to represent the valence electrons on the central atom and the black color is used to represent the valence electrons on the other atoms.

The method is used with Lewis-structure diagrams to categorize polyatomic molecules and ions into different groups. Each group has a preferred three-dimensional geometry that can include, for example, the linear or trigonal planar structures. The correct category can be determined by choosing appropriate , , and terms. The term represents the central atom and the and terms represent the bonded and nonbonded electron pairs. The and values are always integers, and they represent the number of bonding and nonbonding electron pairs.

The following figure shows Lewis-structure diagrams and VSEPR geometries for some of the simplest polyatomic molecular categories.

The figure shows that the number of valence electrons determines how they are arranged around the central atom (). Molecules adopt rather simple linear configurations when there are just two bonded pairs of valence electrons and more complex configurations when there are three or four pairs of bonded or unbonded valence electrons.

Valence electrons always adopt specific three-dimensional structures that minimize the strength of repulsive electrostatic interactions between bonded and nonbonded pairs of electrons. This insight can be used to predict the three-dimensional structures of more complex molecules. The following figure shows the shape of higher-order category molecules.

Less common higher-order molecules are formed as the number of pairs of bonded or unbonded valence electrons increases.

It is useful to move away from generic molecular structures and focus on real molecules instead. Boron trifluoride () is one of the simplest molecules that we can study to better understand the VSEPR model. The central atom () is a single boron atom and it has three bonded valence electron pairs and zero nonbonded electron pairs.

Boron trifluoride molecules preferentially adopt a trigonal planar structure because they are group molecules. The fluorine atoms minimize the repulsive electrostatic energy in molecules by moving to the vertices of a triangle that is centered on the boron atom. There is an angle of between each one of the three covalent – bonds.

Methane molecules () are slightly more complex and are categorized as group molecules. Methane molecules contain one central carbon atom and four pairs of bonded valence electrons. The hydrogen atoms minimize the repulsive electrostatic energy in methane molecules by moving to four vertices of a three-dimensional tetrahedral structure that is centered on the carbon atom. There is an angle of between each one of the four covalent – bonds.

Methane and ammonium ions () both belong to the group and they adopt similar tetrahedral-shaped structures. The molecules have four equivalent bonded pairs of valence electrons because dative bonds are fundamentally identical to “normal” two-electron covalent bonds.

Ammonia molecules are more interesting because they have three equivalent bonded pairs of valence electrons and one nonbonded lone pair of electrons. The nonbonded lone pair takes up more space in the valence shell of the nitrogen atom and this pushes the three bonded pairs (–) closer together. The ammonia molecule ends up having a skewed trigonal pyramidal structure, and there is a relatively small angle of only between the three covalent – bonds.

Lone pairs of electrons take up the most space in the valence shell of the central atom, and they generate the strongest repulsive electrostatic interactions in the VSEPR model. The strongest repulsive electrostatic interactions are between two lone pairs of electrons, and the weakest repulsive electrostatic interactions are between two bonded pairs of electrons. The combination of one lone pair and one bonded pair has a strength somewhere between these two extremes.

Chemists have several rules that help them understand bond angles. They state, for example, that the –– angle will be approximately smaller for every extra pair of nonbonded electrons that are counted on any one quaternary bonded central atom. molecules have –– bond angles of and and molecules have –– bond angles of and .

Water molecules are able to dissolve polar alcohol and carboxylic acid solutes because they have a nonlinear or “bent” structure. Water molecules have an asymmetrical distribution of electron density and they have a permanent dipole moment that runs from the plane of the hydrogen atoms through to the oxygen atom lone pairs. Lewis-structure diagrams and the VSEPR model can be used to understand the shape of water molecules.

The Lewis-structure diagram for water molecules shows that the central oxygen atom has six valence shell electrons. Two of these valence electrons are used to create two covalent (–) bonds and the other four electrons are grouped as two lone pairs.

Water has the group classification, and this means that the valence electrons will adopt a skewed tetrahedron structure. The two lone pairs take up more space in the oxygen valence shell than the bonded pairs, and the –– angle will always end up being smaller than . The –– bond angle usually has a value of because there are two electron lone pairs and each electron lone pair decreases the –– bond angle by approximately from .

The oxygen atom has two electron-rich lone pairs, and this partly explains why it can form hydrogen bonds with other molecules. The other reason is that oxygen atoms are highly electronegative. The oxygen atom attracts a disproportionate amount of electron density from the two – covalent bonds and the water molecule becomes highly polarized. The hydrogen atoms end up being depleted of electron density () and the oxygen atoms become enriched with electron density (). The VSEPR model shows that the water molecule will always adopt a skewed tetrahedron structure. There will always be an electric dipole moment that runs from the plane of the hydrogen atoms through to the oxygen atom lone pairs. Water molecules will also have a permanent dipole moment and they will always be able to form hydrogen bonds with other molecules.

Let us summarize what we have learned in this explainer.