The dipole moment of a molecule and its overall polarity depends on the magnitude and direction of individual polar bonds and their dipole moments.

Remember, for molecules with one polar bond, the molecular dipole is determined simply based on the dipole moment of that bond:

Now, if there are multiple polar bonds, the molecular dipole moment is determined by the vector sum of those dipole moments.

For example, if we replace one more hydrogen atom in dichloromethane (CH₂Cl₂ shown above), we need to find the vector sum of the two C-Cl dipoles to determine the overall dipole of the molecule:

Lone Pairs and Dipole Moment

It is important to mention that lone pairs of electrons also affect the magnitude and direction of the molecular dipole.

For example, the water molecule has two polar bonds and one lone pair of electrons. Oxygen pulls the bonding electrons shared with the hydrogens since it is more electronegative. The lone pairs, on the other hand, point the dipole moment towards them. Although the resulting dipole vectors are not linearly aligned at 0^o, they all point in the same direction, and when combining them, we see that all the dipoles reinforce each other. This can be proved by the greater value of the water dipole moment (1.85 D) compared to the dipole moment of the individual O-H bonds (1.5 D).

It is very helpful to learn the hybridization and VSEPR theories at his point which suggest that the hydrogens and the lone pairs on the oxygen are in sp³ orbitals at ~109.5^o_.

Therefore, 1.85 D is the vector sum of the two dipoles of 1.5 D at 104.5^o plus the reinforcing effect of the lone pairs.

Similarly, ammonia has a greater dipole moment (1.47 D) than each N-H bond itself (1.3 D). The dipole moment of the lone pair is reinforced by the vector sum of the three N-H bonds:

Geometry and Molecular Dipole

We just discussed examples where the dipole moments were enforced by one another according to their vector sum. So what about molecules when the dipole moments are in the opposite direction?

There are two scenarios here. If the dipole moments have the same magnitude at 180^o, then the molecule is nonpolar sine the diple moments are canceled.

So, to generalize, remember that symmetrical molecules have no dipole moment regardless of how polar the covalent bonds are because the overall dipole moment of the molecule depends on the magnitude and direction of individual dipoles.

For example, carbon dioxide has two polar C=O bonds, however, their dipole moments, being at 180^o, are canceled, and therefore, the molecule has no net dipole and is nonpolar:

If the dipole moments are of different magnitudes but in opposite directions, then the net dipole moment will be the vector sum (subtraction) of these dipole moments. See the practice question below.

What if I Don’t know the Molecular Geometries?

There may be a lot of cases where the molecule is not drawn or the geometry is not given, and you are asked to determine whether the given molecule is polar or not.

To answer this question, you need to first determine the geometry of the molecule following the rules of VSEPR theory. These are summarized in the following table, but are also covered in detail here, so feel free to check them and the Lewis structures before we go over the next example:

 Example:

Determine if the molecule is polar or not by showing the corresponding dipole moment(s).

This molecule is formaldehyde and has a trigonal planar geometry according to the VSEPR rules:

The C-H bonds are nonpolar and therefore, the molecular dipole is mostly defined by the magnitude and direction of the C=O bond. Oxygen, being more electronegative, pulls the electron density of the C=O bond, and thus defining the direction of the molecular dipole:

Dipole Moment of Organic Molecules

For organic molecules containing a hydrocarbon chain with few polar bonds, the molecular dipole is determined by the sum of these dipole moments.

For example, diethyl ether possesses a net dipole moment which is the vector sum of two C-O polar bonds:

For larger molecules with lots of sigma bonds, it is more difficult to determine the net dipole moment because of the free rotation about these single bonds. However, the presence of double bonds that restrict the constant change of the dipole moment direction makes it possible to determine the molecular dipole moment.

This can be illustrated by comparing the boiling points of the cis– and trans-1,2-Dichloroethene:

The dipole moments of the C-Cl bonds in cis-1,2-Dichloroethene reinforce and the molecule exhibits a net dipole. In contrast, the trans isomer has no molecular dipole since the C-Cl bonds are at the opposite direction and their dipole effects are canceled:

As a result, the trans isomer is a non-polar molecule with a lower boiling point (48^o vs 60^o C) than the cis-dichloroethene because of the intermolecular dipole-dipole interactions in the latter.

Interestingly, the trans isomer has a higher melting point! Why is that the case?

We will answer this question in the next post about the boiling and melting points.

To summarize, the molecular dipole moment depends on the magnitude and direction of all the dipole moments in the molecule. The larger the molecular dipole, the more polar the molecule is. Symmetrical molecules with comparable structure are less polar since their individual dipole are canceled and there is no net dipole moment.

As always, below are some practice problems on the molecular dipole moment.