I’m confused as to whether both terms “dipole” and “dipole moment” are the same or different, does the moment have something to do with the molecular geometry? I know the vectors of the charges cancel out however carbon dioxide has no dipole moment because of 180 degrees, this makes it non polar, but it still has partial charges right? So isn’t it still polar?

  • $\begingroup$ There is no difference. CO2 is pretty polar, despite having zero dipole moment. $\endgroup$ Jun 8 '19 at 13:53
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    $\begingroup$ Well, being polar is not the same as being a dipole. A dipole is a phenomena of charge displacement and dipole moment is its quantification. To be polar, a molecule must be a dipole, either as the whole as HCl, or its parts, like 2 polar C=O bonds in CO2 are dipoles, that cancel each other. For polarity in sense of relative permitivity, both dipole moments and polarizability count. $\endgroup$
    – Poutnik
    Jun 8 '19 at 14:26
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    $\begingroup$ As @Poutnik writes the electric dipole is the separation of equal and opposite charges $q$ in a molecule and the dipole moment is this difference multiplied by their separation or $ q \times d$. As a vector, conventionally it points from negative to positive charge. The dipole experiences no net force in a uniform electric field but does experience a torque which rotates it to align with the field. $\endgroup$
    – porphyrin
    Jun 8 '19 at 15:32
  • $\begingroup$ Note that (permanent and oriented) dipole-dipole electrostatic interaction is proportional to 1/r^4. If one of dipoles is induced due molecule polarizability, the distance dependence is even stronger. $\endgroup$
    – Poutnik
    Jun 8 '19 at 15:40

The (electric) dipole moment is a mathematically clearly defined quantity, the product of charge difference and distance

A dipole is anything (e.g. a molecule, or a part of a molecule), that has a non-zero dipole moment.

And a polar molecule is one that has at least local dipoles. CO2 is polar if you get close (i.e. in a condensed phase). From a slightly larger distance, it's unpolar, because the centers of the positive and negative charge distributions in it are identical, the two local C=O dipoles cancel each other out.

  • $\begingroup$ The word polarity is also discouraged by IUPAC because polarity has no specific meaning. It is an umbrella term for several types of interactions. The proper term is dipolar or dipolarity (in a strict sense). $\endgroup$
    – M. Farooq
    Jun 9 '19 at 16:50
  • $\begingroup$ @M.Farooq Yea, I heard about that. I don't think that there can be a context in which "polar" or "polarity" run any immediate danger of being misunderstood. Think IUPAC is slowly running out of things that need a proper new definition. ;-) $\endgroup$
    – Karl
    Jun 9 '19 at 19:27
  • $\begingroup$ Actually, I am in favor of IUPAC's suggestion. Come to the field of analytical separations, the usage of solvent polarity will only disappoint you. I am talking about advanced level/ research papers. Yes organic chemists are free to define whatever solvent polarity is, but to a physical chemist, it has no meaning whatsoever. $\endgroup$
    – M. Farooq
    Jun 9 '19 at 19:57
  • $\begingroup$ @M.Farooq You mean it's not a quantitative measure? Right. But there are ways to quantify solvent polarity. Changing the name to dipolarity is just a new label. $\endgroup$
    – Karl
    Jun 9 '19 at 20:26
  • $\begingroup$ Sort of, but polarity is an umbrella term for so many interactions. One can certainly use solvatochromism to assess "polarity." For example see "dipolar aprotic solvent: A solvent with a comparatively high relative permittivity (or dielectric constant), greater than ca. 15, and a sizable permanent dipole moment, that cannot donate suitably labile hydrogen atoms to form strong hydrogen bonds, e.g. dimethyl sulfoxide. The term (and its alternative 'polar aprotic solvent') is a misnomer and is therefore discouraged." $\endgroup$
    – M. Farooq
    Jun 9 '19 at 21:44

I’m confused as to whether both terms “dipole” and “dipole moment” are the same or different, does the moment have something to do with the molecular geometry?

Yes, in colloquial usage you will hear dipole, dipole moment used interchangeably. I think you will feel better if you use the full name: electric dipole moment. This is what the mother of all chemistry related terminologies, IUPAC, tells us to do so.

Dipole moment is a general term because in advanced physics classes you will hear magnetic dipole moment. There you would have to differentiate between electric vs. magnetic dipole moments.

In order to understand O=C=O, you can say that this molecules has two bond dipole moments, since these vectors are 180 degrees apart, there is no net electric dipole moment.

Regarding molecular geometry and dipole: Yes the molecular geometry will let you decide whether there is a net electric dipole moment or not. Nobody in the real world determines the molecular structure on the basis of bond electric dipole moments or the presence of so-called lone pairs. This is good for first year textbooks only. Both molecular geometry and electric dipole moments are experimentally determinable parameters. Electric dipole moments will help to eliminate certain possible molecular geometries. For example, one can be confident that O=C=O is not bent. If it were, it would have a net electric dipole moment.


Dipole moment refers specifically to the second moment in a multipole expansion. An electric dipole moment appears in the expression for the second moment in the multipole expansion of the electric field (or potential) produced by a distribution of electrical charges. "Dipole" is generally a casual abbreviated way to refer to the same thing$\dagger$.

In the case of the distribution of electrical charge in a molecule, the field is divided into a sum of contributions from various multipoles. At large distances from the molecule the sum over the contributions of a small number of terms often suffices to describe the molecule's electric field. The multipole expansion may be applied to the field generated by charges in a local region involving only a pair of bonded atoms, in which case the dipole moment can fairly be referred to as a bond dipole moment. It will differ from the dipole moment for the entire molecule.

$\dagger$ For instance, in the wikipedia entry for electric dipole moment it is stated that

Theoretically, an electric dipole is defined by the first-order term of the multipole expansion;

In this example, "dipole" is used where "dipole moment" is meant. An electric dipole in strict usage consists of two charges separated by a finite distance. The dipole moment is a property of any distribution, including a real dipole, obtained when attempting to approximate the field generated at large distances from the charge distribution by performing a multipole expansion. The higher order multipoles of a true dipole are nonzero, that is, such a dipole will have a dipole moment and other non-zero moments. This is because a dipole moment is a theoretical abstraction used to help represent the field generated by a real distribution.

The wikipedia is a little more careful in making this distinction in the description that follows the above statement:

it [the dipole] consists of two equal and opposite charges that are infinitesimally close together. This is unrealistic, as real dipoles have separated charge. However, because the charge separation is very small compared to everyday lengths, the error introduced by treating real dipoles like they are theoretically perfect is usually negligible. The dipole's direction usually points from the negative charge towards the positive charge.

  • $\begingroup$ I challenge you to make a sensible statement in proper English, casual or not, in which you can leave out the "moment" and it's still correct. $\endgroup$
    – Karl
    Jun 9 '19 at 11:42
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    $\begingroup$ @Karl challenge accepted. The wikipedia was sufficient to provide an example. $\endgroup$
    – Buck Thorn
    Jun 9 '19 at 17:19
  • $\begingroup$ Care to be specific? ;-) I looked in there, and didn't find one. No surprise, you can't replace something that describes an abstract property with a name for a real object and hope to still make sense. $\endgroup$
    – Karl
    Jun 9 '19 at 19:20
  • $\begingroup$ If you mean the sentence you cited above: "Theoretically, the electric dipole moment is defined as the first-order term of the multipole expansion; " is quite a different statement. $\endgroup$
    – Karl
    Jun 9 '19 at 19:31
  • $\begingroup$ @Karl Yes, I meant exactly that example. My point is that people will often use words without literal intent. Someone might write "Theoretically, an electric dipole is" and mean "Theoretically, an electric dipole moment is". It is incorrect usage, therefore my explanation that this is "casual" usage, not strict. $\endgroup$
    – Buck Thorn
    Jun 10 '19 at 8:52

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