A dipole exists when electrons are unevenly shared between two bonded atoms. This creates partial positive and negative charges at either end of the two atoms. Conceptually, it is useful to think of a molecule having a cloud of electrons that get pulled one way or another by the nuclei of the atoms, depending on their electronegativity. Dipole moments increase with ionic bond character, and decrease with covalent bond character.
Bond dipole moment
The bond dipole moment uses the idea of electric dipole moment to measure the polarity of a chemical bond within a molecule. It occurs whenever there is a separation of positive and negative charges. Many molecules have dipole moments due to non-uniform distributions of positive and negative charges among various atoms. This is the case with polar compounds like hydrogen fluoride (HF), where electron density is shared unequally between atoms. Therefore, a molecule's dipole is an electric dipole with an inherent electric field—not to be confused with a magnetic dipole, which generates a magnetic field.
The physical chemist Peter J. W. Debye was the first scientist to study molecular dipoles extensively. Bond dipole moments are commonly measured in debyes, represented by the symbol D, which is obtained by measuring the charge in units of 10-10 statcoulomb and the distance d in Angstroms. Note that 10-10 statcoulomb is 0.48 units of elementary charge. Another useful conversion factor is 1 C m = 2.9979×1029 D.
For diatomic molecules there is only one (single or multiple) bond so the bond dipole moment is the molecular dipole moment, with typical values in the range of 0 to 11 D. At one extreme, a symmetrical molecule such as chlorine, Cl2, has zero dipole moment. At the other extreme, gas phase potassium bromide, KBr, which is highly ionic, has a dipole moment of 10.5 D.
Molecular dipole moment
When a molecule consists of more than two atoms, there is more than one bond holding the molecule together. The dipole for the entire molecule can be calculated by adding up all the individual dipoles of the individual bonds, as their vector. Dipole moment values can be experimentally obtained from measurement of the dielectric constant. Some typical gas phase values in debye units are:
- carbon dioxide: 0 (despite having two polar C=O bonds, the two are pointed in geometrically opposite directions, cancelling each other out and giving a molecule with no net dipole moment) (Figure 1)
- carbon monoxide: 0.112 D
- ozone: 0.53 D
- phosgene: 1.17 D
- water vapor: 1.85 D
- hydrogen cyanide: 2.98 D
- cyanamide: 4.27 D
- potassium bromide: 10.41 D
KBr has one of the highest dipole moments because it is a very ionic molecule (which only exists as a molecule in the gas phase).
The bent molecule H2O has a net dipole. The two bond dipoles do not cancel.
While the above figure in instructive for visualizing the bond angles that lead to a net dipole in water, determining the net dipole requires knowledge about directionality of the polarization in each bond. Another way to describe this phenomenon is to use arrows such as those shown in the figure below, where the arrowhead indicates the polarization of electron density, leading to a partial negative charge. (Figure 3)