The Lewis structure is not necessarily a good pictorial representation of the three-dimensional appearance of a molecule. The actual geometric arrangement of the bonds and different atoms is obtained by using the VSEPR theory described below. The shape of a molecule can affect its polarity.
The valence shell electron-pair repulsion (VSEPR) theory uses Lewis structures to predict the molecular geometry of covalently bonded molecules. It states that the three-dimensional arrangement of atoms surrounding a central atom is determined by the repulsions between the bonding and the nonbonding electron pairs in the valence shell of the central atom. These electron pairs arrange themselves as far apart as possible, thereby minimizing repulsion.
The following steps are used to predict the geometrical structure of a molecule using the VSEPR theory.
Valence electron arrangements are summarized in the following table:
| number of valence electrons | example | geometric arrangement of electron pairs around the central atom | shape | angle between electron pairs |
|---|---|---|---|---|
| 2 | BeCl2 |  |
linear | 180° |
| 3 | BH3 |  |
trigonal planar | 120° |
| 4 | CH4 |  |
tetrahedral | 109.5° |
| 5 | PCl5 |  |
trigonal bipyramidal | 90°,120°,180° |
| 6 | SF6 |  |
octahedral | 90°,180° |
While the number of electron pairs dictates their overall arrangement around the central atom, it is only a starting point in arriving at the actual description of the geometry of the molecule. If one of the X’s in the table above is a lone pair of electrons rather than an actual atom or group of atoms, new terms need to be introduced to describe the spatial arrangement of the atoms. The example below illustrates this point.
| Example: | Predict the geometry of NH3. |
| Solution: |
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In describing the shape of a molecule, only the arrangement of atoms (not electrons) is considered. Even though the electron pairs are arranged tetrahedrally, the shape of NH3 is described as trigonal pyramidal. It is not trigonal planar because the lone pair repels the three bonding electron pairs, causing them to move as far away as possible.
The double bond behaves just like a single bond for purposes of predicting molecular shape. This compound has two groups of electrons around the carbon. According to the VSEPR theory, the two sets of electrons will orient themselves 180° apart, on opposite sides of the carbon atom, minimizing electron repulsion. Therefore, the molecular structure of CO2 is linear.
Earlier we talked about the concept of the dipole moment in a polar covalent bond. If a molecule has more than two atoms, there will be more than one bond. Each bond may or may not be a dipole, and in such cases one can talk about the polarity of the molecule as a whole. A molecule is polar if it has polar bonds and if the dipole moments of these bonds do not cancel one another (by pointing in opposite directions, for example). The polarity of a molecule, therefore, depends on the polarity of the constituent bonds and on the shape of the molecule. A molecule with only nonpolar bonds is always nonpolar; a molecule with polar bonds may be polar or nonpolar, depending on the orientation of the bond dipoles. For instance, CCl4 has four polar C–Cl bonds. According to the VSEPR theory, the shape of CCl4 is tetrahedral. The four bond dipoles point to the vertices of the tetrahedron and cancel each other, resulting in a nonpolar molecule.
However, if the orientation of the bond dipoles is such that they do not cancel out, the molecules will have a net dipole moment and therefore be polar. For instance, H2O has two polar O–H bonds. According to the VSEPR model, its shape is angular. The two dipoles add together to give a net dipole moment to the molecule, making the H2O molecule polar.
A molecule of two atoms bound by a polar bond must have a net dipole moment and therefore be polar. The two equal and opposite partial charges are localized at the ends of the molecule on the two atoms.
