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Dipolar molecules have a permanent separation of positive and negative charge, so one side is partially positive and another partially negative. They remain neutral overall. Their dipoles create directional electric attractions and repulsions with nearby molecules; molecular shape determines whether polar bonds produce a net dipole in the first place.
What makes a molecule dipolar?
A dipole is a separation of positive and negative charge. In a dipolar molecule, the separation gives the molecule a permanent, nonzero electric dipole moment. The ends carry partial charges rather than necessarily being separate ions, and the molecule as a whole is still electrically neutral. IUPAC defines the dipole–dipole interaction in terms of dipoles whose interaction strength depends on their distance and relative orientation.
Polar bonds alone do not guarantee a dipolar molecule. Each bond’s contribution has a direction, and the molecule’s three-dimensional geometry determines whether those contributions reinforce or cancel. Water has a net dipole; in a symmetrical molecule, polar bond contributions can cancel one another.
Dipole moments are often reported in debyes (D). The educational sources cited here give water’s dipole moment as 1.85 D. Georgia Tech also lists HCl at 1.1 D, CH₃Cl at 1.9 D, HCN at 2.9 D, and NH₃ at 1.47 D; those consulted pages do not state publication years. Colorado State University’s key-terms page and Georgia Tech’s molecular-interactions page provide these examples.
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How do two permanent dipoles interact?
Each permanent dipole has an electric field. When two dipolar molecules approach, the force between them is electrostatic: an arrangement with the positive end of one near the negative end of the other tends to attract, while like ends near one another tend to repel. The interaction is not a covalent bond; it is an intermolecular force between molecules.
Orientation matters as much as distance. Molecules move and rotate, so their relative alignment changes; a pair of dipoles can attract in one arrangement and repel in another. The interaction weakens as the molecules move farther apart.
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Why two distance rules are used
For dipoles held in fixed orientations, the interaction energy varies with the inverse cube of separation (1/r³). In a liquid, thermal rotation continually changes their orientations. Averaging over those orientations gives a net attractive interaction with an inverse-sixth-power dependence (1/r⁶). These describe different conditions, not conflicting rules. Georgia Tech’s discussion of molecular interactions describes both scalings.
How dipole interactions relate to other intermolecular forces
Nearby molecules can interact through induced or fluctuating dipoles as well as permanent ones. These mechanisms are related to electrostatic charge distributions, but they are not all dipole–dipole interactions between two permanent molecular dipoles.
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| Interaction | What creates the dipoles? | How to distinguish it |
|---|---|---|
| Permanent dipole–dipole | Two molecules each have a permanent dipole. | Orientation determines whether a particular arrangement attracts or repels; thermal averaging in liquids gives a net attraction. |
| Dipole-induced dipole | A permanent dipole distorts the electron cloud of a nearby polarizable molecule, inducing a dipole. | The treatment described by the National Programme on Technology Enhanced Learning characterizes this interaction as attractive. |
| London dispersion | Fluctuations in electron distributions create instantaneous dipoles. | Dispersion occurs between all molecules, including nonpolar ones; it is not an interaction requiring two permanent dipoles. |
| Hydrogen bonding | A strongly polarized O–H or N–H bond interacts with a neighboring electronegative atom. | A particularly strong, directional case of dipole–dipole attraction. Despite its name, it is an intermolecular force, not a covalent bond. |
OpenStax’s discussion of noncovalent interactions describes hydrogen bonding, while the NPTEL explanation of deviations from ideal gas behaviour covers induced dipoles and dispersion.
Independent reader supportYour contribution helps us test, update, and keep practical guides available for everyone.What dipoles can—and cannot—tell you about boiling points
Intermolecular attractions influence bulk properties such as boiling point, but dipole moment alone does not explain every difference. Dispersion and hydrogen bonding can also contribute, so a comparison should account for the available interactions rather than treating polarity as the only cause.
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For example, OpenStax gives normal boiling points of 188 K for HCl and 85 K for F₂ to illustrate the role of intermolecular attractions. The comparison illustrates a bulk-property outcome; it is not a general formula for predicting boiling points from dipole moments.
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