Editing Chemical polarity (section)

===Bond dipole moments===
[[Image:Polarity boron trifluoride.png|thumb|250px|A diagram showing the bond dipole moments of [[boron trifluoride]]. δ− shows an increase in negative charge and δ+ shows an increase in positive charge. Note that the dipole moments drawn in this diagram represent the shift of the valence electrons as the origin of the charge, which is opposite the direction of the actual electric dipole moment.]]

The '''bond dipole moment'''<ref>{{cite web |url=https://chem.libretexts.org/Textbook_Maps/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Physical_Properties_of_Matter/Atomic_and_Molecular_Properties/Dipole_Moments |title=Dipole_Moments |last=Blaber |first=Mike |date=2018 |website=Libre Texts |publisher=California State University |access-date= |quote=}}</ref> 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. 

The bond dipole [[μ]] is given by:

:<math>\mu = \delta \, d</math>.

The bond dipole is modeled as δ<sup>+</sup>&nbsp;&mdash;&nbsp;δ<sup>–</sup> with a distance ''d'' between the [[partial charges]] δ<sup>+</sup> and δ<sup>–</sup>. It is a vector,  parallel to the bond axis, pointing from minus to plus,<ref>{{GoldBookRef|title=electric dipole moment, ''p''|file=E01929}}</ref> as is conventional for [[electric dipole moment]] vectors.

Chemists often draw the vector pointing from plus to minus.<ref>{{cite journal |title= Misconceptions in Sign Conventions: Flipping the Electric Dipole Moment |first1= James W. |last1= Hovick |first2= J. C. |last2= Poler |journal= J. Chem. Educ. |year= 2005 |volume= 82 |issue= 6 |page= 889 |doi= 10.1021/ed082p889 |bibcode= 2005JChEd..82..889H }}</ref> This vector can be physically interpreted as the movement undergone by electrons when the two atoms are placed a distance ''d'' apart and allowed to interact, the electrons will move from their free state positions to be localised more around the more [[Electronegativity|electronegative]] atom.

The [[SI unit]] for electric dipole moment is the coulomb–meter. This is too large to be practical on the molecular scale.  
Bond dipole moments are commonly measured in [[debye]]s, represented by the symbol D, which is obtained by measuring the charge <math>\delta</math> in units of 10<sup>−10</sup> [[statcoulomb]] and the distance ''d'' in [[Angstrom]]s.  Based on the [[conversion factor]] of
10<sup>−10</sup> statcoulomb being 0.208 units of elementary charge, so 1.0 debye results from an electron and a proton separated by 0.208&nbsp;Å. A useful conversion factor is 1&nbsp;D&nbsp;=&nbsp;3.335&nbsp;64{{e|-30}}&nbsp;C&nbsp;m.<ref>{{cite book |last1=Atkins |first1=Peter |last2=de Paula |first2=Julio |date=2006 |title=Physical Chemistry |edition=8th |page=[https://archive.org/details/atkinsphysicalch00pwat/page/620 620 (and inside front cover)] |publisher=W.H. Freeman |isbn=0-7167-8759-8 |url-access=registration |url=https://archive.org/details/atkinsphysicalch00pwat/page/620 }}</ref> 

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&nbsp;D.  At one extreme, a symmetrical molecule such as [[bromine]], {{chem|Br|2}}, has zero dipole moment, while near the other extreme, gas phase [[potassium bromide]], KBr, which is highly ionic, has a dipole moment of 10.41&nbsp;D.<ref> ''Physical chemistry'' 2nd Edition (1966) G.M. Barrow McGraw Hill</ref>{{page needed|date=October 2019}}<ref>{{cite journal |title= Dipole Moments of KF and KBr Measured by the Molecular-Beam Electric-Resonance Method |journal= J. Chem. Phys. |volume= 47 |issue= 7 |pages= 2256 |year= 1967 |doi= 10.1063/1.1703301 |first1= R. |last1= Van Wachem |first2= F. H. |last2= De Leeuw |first3= A. |last3= Dymanus |bibcode= 1967JChPh..47.2256V }}</ref>{{verify source|date=October 2021}}

For polyatomic molecules, there is more than one bond. The total [[Dipole#Molecular dipoles|molecular dipole moment]] may be approximated as the [[vector sum]] of the individual bond dipole moments. Often bond dipoles are obtained by the reverse process: a known total dipole of a molecule can be decomposed into bond dipoles. This is done to transfer bond dipole moments to molecules that have the same bonds, but for which the total dipole moment is not yet known. The vector sum of the transferred bond dipoles gives an estimate for the total (unknown) dipole of the molecule.