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Molecules have shapes. There is an abundance of experimental evidence to that effect — from their physical properties to their chemical reactivity. Small molecules — molecules with a single central atom — have shapes that can be easily predicted.

The basic idea in molecular shapes is called valence shell electron pair repulsion (VSEPR). It says that electron pairs, being composed of negatively charged particles, repel each other to get as far away from each other as possible. VSEPR makes a distinction between electron group geometry, which expresses how electron groups (bonding and nonbonding electron pairs) are arranged, and molecular geometry, which expresses how the atoms in a molecule are arranged. However, the two geometries are related.

There are two types of electron groups: any type of bond — single, double, or triple — and lone electron pairs. When applying VSEPR to simple molecules, the first thing to do is to count the number of electron groups around the central atom. Remember that a multiple bond counts as only one electron group.

Any molecule with only two atoms is linear. A molecule whose central atom contains only two electron groups orients those two groups as far apart from each other as possible — 180° apart. When the two electron groups are 180° apart, the atoms attached to those electron groups are also 180° apart, so the overall molecular shape is linear. Examples include BeH₂ and CO₂:

[latex]\begin{array}{lr} \begin{array}{c} \chemfig{H-Be-H} \\ \\ 180^{\circ} \end{array} &\hspace{10mm} \begin{array}{c} \chemfig{\Lewis{2:6:,O}=C=\Lewis{2:6:,O}} \\ \\ 180^{\circ} \end{array} \end{array}[/latex]

A molecule with three electron groups orients the three groups as far apart as possible. They adopt the positions of an equilateral triangle — 120° apart and in a plane. The shape of such molecules is trigonal planar. An example is BF₃:

[latex]\begin{matrix} 120^{\circ}&& \\ &\searrow & \\ &&\chemfig{\Lewis{2:4:6:,F}-[:30]B(-[:-30]\Lewis{0:2:6:,F})-[:90]\Lewis{0:2:4:,F}} \end{matrix}[/latex]

Some substances have a trigonal planar electron group distribution but have atoms bonded to only two of the three electron groups. An example is GeF₂:

[latex]\chemfig{\Lewis{2:4:6:,F}-[:30]\Lewis{2:,Ge}-[:-30]\Lewis{0:2:6:,F}}[/latex]

From an electron-group-geometry perspective, GeF₂ has a trigonal planar shape, but its real shape is dictated by the positions of the atoms. This shape is called bent or angular.

A molecule with four electron groups around the central atom orients the four groups in the direction of a tetrahedron, as shown in Figure 9.4 “Tetrahedral Geometry.” If there are four atoms attached to these electron groups, then the molecular shape is also tetrahedral. Methane (CH₄) is an example.

[latex]\chemfig{H-[:30]C(-[:90]H)(<[:-30]H)(<:[:-5]H)}[/latex]

This diagram of CH₄ illustrates the standard convention of displaying a three-dimensional molecule on a two-dimensional surface. The straight lines are in the plane of the page, the solid wedged line is coming out of the plane toward the reader, and the dashed wedged line is going out of the plane away from the reader.

NH₃ is an example of a molecule whose central atom has four electron groups but only three of them are bonded to surrounding atoms.

[latex]\chemfig{H-[:30]\Lewis{2:,N}(<:[:-5]H)<[:-30]H}[/latex]

Although the electron groups are oriented in the shape of a tetrahedron, from a molecular geometry perspective, the shape of NH₃ is trigonal pyramidal.

H₂O is an example of a molecule with a central atom that has four electron groups but only two of them are bonded to surrounding atoms.

[latex]\chemfig{H-[:30]\Lewis{0:2:,O}<[:-30]H}[/latex]

Although the electron groups are oriented in the shape of a tetrahedron, the shape of the molecule is bent or angular. A molecule with four electron groups around the central atom but only one electron group bonded to another atom is linear because there are only two atoms in the molecule.

Double or triple bonds count as a single electron group. CH₂O has the following Lewis electron dot diagram.

[latex]\begin{array}{c} \Lewis{0:4:6:,O} \\ \ce{H}\Lewis{0:2:4:,C}\ce{H} \\ \end{array}[/latex]

The central C atom has three electron groups around it because the double bond counts as one electron group. The three electron groups repel each other to adopt a trigonal planar shape:

[latex]\chemfig{H-[:30]C(=[:90]O)-[:-30]H}[/latex]

(The lone electron pairs on the O atom are omitted for clarity.) The molecule will not be a perfect equilateral triangle because the C–O double bond is different from the two C–H bonds, but both planar and triangular describe the appropriate approximate shape of this molecule.

Table 9.4 “Summary of Molecular Shapes” summarizes the shapes of molecules based on their number of electron groups and surrounding atoms.

Molecular Polarity

The overall polarity of molecules with more than one bond is determined from both the polarity of the individual bonds and the shape of the molecule. Each bond’s dipole moment can be treated as a vector quantity, having a magnitude and direction. Therefore the molecular polarity is the vector sum of the individual bond dipoles.

How to determine the vector sum

One method to determine the vector sum of dipole arrows is known as the tail-to-head method. Let’s examine this method for a molecule of water.

1. First draw the Lewis electron dot diagram for water and determine its molecular shape. Water has four electron groups, but only two atoms attached to the central atom so it is bent.

   

2. Draw in dipole arrows for all polar covalent bonds, starting the arrow at the more electropositive atom, and ending at the more electronegative atom.

   

3. Connect the dipole arrows tail-to-head.

   

4. Draw a new line connecting the tail of the first vector. This is the net molecular dipole.

   

5. Now superimpose the net molecular dipole arrow onto the molecule.

   

An alternative method for determining the vector sum

An alternative method to determine the vector sum of dipole arrows is known as the vector component method. Let’s examine this method again for a molecule of water. The first two steps remain the same as the tail-to-head method:

1. First draw the Lewis electron dot diagram for water and determine its molecular shape. Water has four electron groups, but only two atoms attached to the central atom so it is bent.

   

2. Draw in dipole arrows for all polar covalent bonds, starting the arrow at the more electropositive atom, and ending at the more electronegative atom.

   

3. For dipole arrows at an angle, separate them into horizontal and vertical vector components.

   

4. Superimpose the vector components onto the molecule.

   

5. Cancel out any vector components that are of equal magnitude and pointing in the opposite direction.

   

6. The remaining vector components both point vertically, showing the net molecular dipole.

   

As a result of the vector quantity nature of bond dipoles, some molecules may contain polar bonds, yet have no net molecular dipole moment. For example, CO₂: