Editing Dipole (section)

== Molecular dipoles ==
<!-- This section is linked from [[Ammonia]] and redirects from [[Molecular dipole]] -->
{{See also|Chemical polarity|Dipole moments of molecules}}

Many [[molecule]]s have such dipole moments due to non-uniform distributions of positive and negative charges on the various atoms.  Such 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 that should not be confused with a [[magnetic dipole]], which generates a magnetic field.

The physical chemist [[Peter Debye|Peter J. W. Debye]] was the first scientist to study molecular dipoles extensively, and, as a consequence, dipole moments are measured in the non-[[SI]] unit named ''[[debye]]'' in his honor.

For molecules there are three types of dipoles:
; Permanent dipoles: These occur when two atoms in a molecule have substantially different [[electronegativity]] : One atom attracts electrons more than another, becoming more negative, while the other atom becomes more positive. A molecule with a permanent dipole moment is called a ''polar'' molecule. See [[Intermolecular force#Dipole–dipole interactions|dipole–dipole attractions]].
; Instantaneous dipoles : These occur due to chance when [[electron]]s happen to be more concentrated in one place than another in a [[molecule]], creating a temporary dipole. These dipoles are smaller in magnitude than permanent dipoles, but still play a large role in chemistry and biochemistry due to their prevalence. See [[London dispersion force|instantaneous dipole]].
; Induced dipoles : These can occur when one molecule with a permanent dipole repels another molecule's electrons, ''inducing'' a dipole moment in that molecule. A molecule is ''polarized'' when it carries an induced dipole. See [[Intermolecular force#Debye (permanent–induced dipoles) force|induced-dipole attraction]].

More generally, an induced dipole of ''any'' polarizable charge distribution ''ρ'' (remember that a molecule has a charge distribution) is caused by an electric field external to ''ρ''. This field may, for instance, originate from an ion or polar molecule in the vicinity of ''ρ'' or may be macroscopic (e.g., a molecule between the plates of a charged [[capacitor]]). The size of the induced dipole moment is equal to the product of the strength of the external field and the dipole [[polarizability]] of ''ρ''.

Dipole moment values can be obtained from measurement of the [[dielectric constant]]. Some typical gas phase values given with the unit [[debye]] are:<ref>
{{cite book
 | last = Weast | first = Robert C.
 | title=CRC Handbook of Chemistry and Physics
 | edition = 65th
 | publisher=CRC Press
 | year=1984
 | isbn=0-8493-0465-2
}}</ref>
* [[carbon dioxide]]: 0
* [[carbon monoxide]]: 0.112&nbsp;D
* [[ozone]]: 0.53&nbsp;D
* [[phosgene]]: 1.17&nbsp;D
* [[ammonia]]: 1.42&nbsp;D
* [[water vapor]]: 1.85&nbsp;D
* [[hydrogen cyanide]]: 2.98&nbsp;D
* [[cyanamide]]: 4.27&nbsp;D
* [[potassium bromide]]: 10.41&nbsp;D

[[File:Carbon-dioxide-2D-dimensions.svg|thumb|160 px|The linear molecule CO<sub>2</sub> has a zero dipole as the two bond dipoles cancel.]]
Potassium bromide (KBr) has one of the highest dipole moments because it is an [[ionic compound]] that exists as a molecule in the gas phase.

[[File:H2O 2D labelled.svg|thumb|160 px|The bent molecule H<sub>2</sub>O has a net dipole. The two bond dipoles do not cancel.]]
The overall dipole moment of a molecule may be approximated as a [[Euclidean vector#Addition and subtraction|vector sum]] of [[bond dipole moment]]s. As a vector sum it depends on the relative orientation of the bonds, so that from the dipole moment information can be deduced about the [[molecular geometry]].

For example, the zero dipole of CO<sub>2</sub> implies that the two C=O bond dipole moments cancel so that the molecule must be linear. For H<sub>2</sub>O the O−H bond moments do not cancel because the molecule is bent. For ozone (O<sub>3</sub>) which is also a bent molecule, the bond dipole moments are not zero even though the O−O bonds are between similar atoms. This agrees with the Lewis structures for the resonance forms of ozone which show a positive charge on the central oxygen atom.

[[File:Ozone-resonance-Lewis-2D.svg|center|400px|Resonance Lewis structures of the ozone molecule]]
{{multiple image
 | align=right
 | image1=Cis-1,2-dichloroethene.png
 | width1=150
 | caption1=''Cis'' isomer, dipole moment 1.90&nbsp;D
 | image2=Trans-1,2-dichloroethene.png
 | width2=150
 | caption2=''Trans'' isomer, dipole moment zero
}}
An example in organic chemistry of the role of geometry in determining dipole moment is the [[cis–trans isomerism|''cis'' and ''trans'' isomers]] of [[1,2-dichloroethene]]. In the ''cis'' isomer the two polar C−Cl bonds are on the same side of the C=C double bond and the molecular dipole moment is 1.90&nbsp;D. In the ''trans'' isomer, the dipole moment is zero because the two C−Cl bonds are on opposite sides of the C=C and cancel (and the two bond moments for the much less polar C−H bonds also cancel).

Another example of the role of molecular geometry is [[boron trifluoride]], which has three polar bonds with a difference in [[electronegativity]] greater than the traditionally cited threshold of 1.7 for [[ionic bonding]]. However, due to the equilateral triangular distribution of the fluoride ions centered on and in the same plane as the boron cation, the symmetry of the molecule results in its dipole moment being zero.