Editing Chemical polarity (section)

== Polarity of bonds ==
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| footer = In a molecule of [[hydrogen fluoride]] (HF), the more [[electronegative]] atom ([[fluorine]]) is shown in yellow. Because the electrons spend more time by the fluorine atom in the H−F bond, the red represents partially negatively charged regions, while blue represents partially positively charged regions.
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Not all atoms attract electrons with the same force. The amount of "pull" an atom exerts on its electrons is called its [[electronegativity]]. Atoms with high electronegativities{{snd}}such as [[fluorine]], [[oxygen]], and [[nitrogen]]{{snd}}exert a greater pull on electrons than atoms with lower electronegativities such as [[alkali metal]]s and [[alkaline earth metal]]s. In a bond, this leads to unequal sharing of electrons between the atoms, as electrons will be drawn closer to the atom with the higher electronegativity.

Because electrons have a negative charge, the unequal sharing of electrons within a bond leads to the formation of an [[electric dipole]]: a separation of positive and negative electric charge. Because the amount of charge separated in such dipoles is usually smaller than a [[fundamental charge]], they are called [[partial charges]], denoted as δ+ ([[Delta (letter)|delta]] plus) and δ− (delta minus). These symbols were introduced by [[Christopher Kelk Ingold|Sir Christopher Ingold]] and [[Hilda Ingold|Edith Hilda (Usherwood) Ingold]] in 1926.<ref>{{cite journal|title=The Origin of the "Delta" Symbol for Fractional Charges|author1-link=William B. Jensen |last=Jensen |first=William B. |journal=J. Chem. Educ. |date=2009 |volume=86 |issue=5 |page=545 |url=http://www.jce.divched.org/Journal/Issues/2009/May/abs545.html |doi=10.1021/ed086p545|bibcode=2009JChEd..86..545J |url-access=subscription }}</ref><ref>{{cite journal |author1=Ingold, C. K. |author2=Ingold, E. H. | title = The Nature of the Alternating Effect in Carbon Chains. Part V. A Discussion of Aromatic Substitution with Special Reference to Respective Roles of Polar and Nonpolar Dissociation; and a Further Study of the Relative Directive Efficiencies of Oxygen and Nitrogen | journal = J. Chem. Soc. | date = 1926 | pages = 1310–1328 | volume=129 | doi = 10.1039/jr9262901310 }}</ref> The bond dipole moment is calculated by multiplying the amount of charge separated and the distance between the charges.

These dipoles within molecules can interact with dipoles in other molecules, creating [[Intermolecular force#Dipole-dipole interactions|dipole-dipole intermolecular forces]].

===Classification===
Bonds can fall between one of two extremes{{snd}}completely nonpolar or completely polar. A completely nonpolar bond occurs when the electronegativities are identical and therefore possess a difference of zero. A completely polar bond is more correctly called an [[ionic bond]], and occurs when the difference between electronegativities is large enough that one atom actually takes an electron from the other. The terms "polar" and "nonpolar" are usually applied to [[covalent bond]]s, that is, bonds where the polarity is not complete. To determine the polarity of a covalent bond using numerical means, the difference between the electronegativity of the atoms is used.

Bond polarity is typically divided into three groups that are loosely based on the difference in electronegativity between the two bonded atoms. According to the [[Electronegativity#Pauling electronegativity|Pauling scale]]:
* ''Nonpolar bonds'' generally occur when the difference in [[electronegativity]] between the two atoms is less than 0.5
* ''Polar bonds'' generally occur when the difference in electronegativity between the two atoms is roughly between 0.5 and 2.0
* ''[[Ionic bonds]]'' generally occur when the difference in electronegativity between the two atoms is greater than 2.0

[[Linus Pauling|Pauling]] based this classification scheme on the ''partial ionic character'' of a bond, which is an approximate function of the difference in electronegativity between the two bonded atoms. He estimated that a difference of 1.7 corresponds to 50% ionic character, so that a greater difference corresponds to a bond which is predominantly ionic.<ref>{{cite book|first=L. |last=Pauling |author-link=Linus Pauling |title=The Nature of the Chemical Bond |url=https://archive.org/details/natureofchemical00paul |url-access=registration |edition=3rd |publisher=Oxford University Press |date=1960 |pages=[https://archive.org/details/natureofchemical00paul/page/98 98–100] |isbn=0801403332}}</ref>

As a [[Introduction to quantum mechanics|quantum-mechanical]] description, Pauling proposed that the [[wave function]] for a polar molecule AB is a [[linear combination]] of wave functions for covalent and ionic molecules: ψ = aψ(A:B) + bψ(A<sup>+</sup>B<sup>−</sup>). The amount of covalent and ionic character depends on the values of the squared coefficients a<sup>2</sup> and b<sup>2</sup>.<ref>{{cite book|first=L. |last=Pauling |author-link=Linus Pauling |title=The Nature of the Chemical Bond |url=https://archive.org/details/natureofchemical00paul |url-access=registration |edition=3rd |publisher=Oxford University Press |date=1960 |page=[https://archive.org/details/natureofchemical00paul/page/66 66] |isbn=0801403332}}</ref>

===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.