Core Concepts
In this tutorial, you will learn all about dipole moments. First, we begin with a basic introduction to the presence of dipole moments in chemical bonds and molecules. Then, we discuss the formula used to calculate dipole moments and explain their classification as a vector (rather than scalar) quantity. Lastly, we provide several examples that demonstrate the importance of molecular geometry in determining net dipole moments.
Topics Covered in Other Articles
- Electron Orbitals & Orbital Shapes
- What Is Electron Affinity?
- Sigma and Pi Bonds
- What Is Formal Charge?
- Cations and Anions
Introduction to Dipole Moments
Dipole moments occur in any system in which there is a separation of positive and negative electrical charges; therefore, they can be found in both ionic and covalent bonds. A dipole moment is a result of unequal electron distribution, which occurs due to differences in electronegativity between chemically-bonded atoms (since electrons tend to move towards more electronegative atoms). Greater differences in electronegativity will consequently lead to larger dipole moments.
Since they have both a magnitude and direction, dipole moments are considered to be a vector quantity. Shown below is an illustration of the dipole moment present in a molecule of hydrofluoric acid. The arrow represents the shift of electron density in the molecule; it points away from the less electronegative (and more positively-charged) hydrogen atom and towards the more electronegative (and more negatively-charged) fluorine atom.
It is important to note the distinction between bond dipole moments and molecular dipole moments. A bond dipole moment is a measure of the polarity of a chemical bond between two atoms in a molecule. On the other hand, the molecular dipole moment (also referred to as the net dipole moment) is a measure of the polarity of the molecule as a whole. Therefore, in polyatomic molecules (i.e., molecules made up of three or more atoms), the net dipole moment is the vector sum of all bond dipole moments present in the molecule. In addition, the net dipole moment is affected by the geometry of the molecule, which can be predicted using VSEPR theory.
Dipole Moment Formula
To calculate the dipole moment of a chemical bond, the following formula is used: 
(where 
is the bond dipole moment, 
is the magnitude of the separated charge, and 
is the distance of separation between the charges). The SI unit for dipole moment is the coulomb-meter (
); however, it is most commonly measured in debyes (
). 1 debye is equal to 
.
As mentioned above, dipole moments are a vector quantity because they have both a magnitude (determined by the electronegativity difference) and a direction (based on the movement of electrons toward the more electronegative atom). Therefore, it is possible for the net dipole moment of a molecule to be zero when a symmetric molecule has individual bond dipole moments (equal in magnitude but opposite in direction) that cancel each other out. In this case, the molecule would be classified as non-polar, despite the fact that it contains multiple polar bonds.
Dipole Moment Examples
CO2 Dipole Moment
In a molecule of carbon dioxide (CO2), there are two polar covalent bonds; each involves the central carbon atom and an attached oxygen atom. Since oxygen is more electronegative than carbon, both oxygen atoms have a partially-negative charge, while the carbon atom has a partially-positive charge. The two arrows representing the molecule’s bond dipole moments point away from the carbon atom and towards the oxygen atoms as a result. In addition, because the central atom bonds with two other atoms and lacks any lone pairs of electrons, CO2 has a linear structure (with a bond angle of 180°). Therefore, the two bond dipole moments run in opposite directions and cancel each other out. This means that CO2 has a net dipole moment of zero and is a non-polar molecule.
H2O Dipole Moment
In a molecule of water (H2O), there are two polar covalent bonds; each involves the central oxygen atom and an attached hydrogen atom. The two arrows representing the molecule’s bond dipole moments point away from the less electronegative hydrogen atoms and towards the more electronegative oxygen atom. In addition, the oxygen atom has a partially-negative charge, while the two hydrogen atoms have partially-positive charges. Because the central oxygen atom bonds with two other atoms and contains two lone pairs of electrons, the structure of H2O is bent (with a bond angle of 104.5°). Unlike in a molecule of carbon dioxide, the individual bond dipole moments in a molecule of water do not cancel out. This means that H2O has a net dipole moment and is a polar molecule.
NH3 Dipole Moment
In a molecule of ammonia (NH3), there are three polar covalent bonds; each involves the central nitrogen atom and an attached hydrogen atom. Electrons tend to move towards the more electronegative nitrogen atom and away from the less electronegative hydrogen atoms (as indicated by the arrows for the three individual bond dipole moments). The nitrogen atom has a partially-negative charge, while the hydrogen atoms have partially-positive charges. Since the central nitrogen atom bonds with three other atoms and has one lone pair of electrons, a molecule of ammonia has a trigonal-pyramidal shape (with a bond angle of 107°). The individual bond dipole moments in a molecule of ammonia do not cancel out; this means that NH3 has a net dipole moment and is a polar molecule.
