View source for Electronegativity



'''Electronegativity''', symbol &chi;, is a [[chemical property]] which describes the power of an [[atom]] (or, more rarely, a [[functional group]]) to attract electrons towards itself.<ref name="definition">"[http://goldbook.iupac.org/E01990.html Electronegativity.]", ''[[Compendium of Chemical Terminology]]''</ref> First proposed by [[Linus Pauling]] in 1932 as a development of [[valence bond theory]],<ref name="paulingJACS">[[Linus Pauling|Pauling, Linus]] (1932). ''[[Journal of the American Chemical Society|J. Am. Chem. Soc.]] '''54''':3570.</ref> it has been shown to correlate with a number of other chemical properties. Electronegativity cannot be directly measured and must be calculated from other atomic or molecular properties. Several methods of calculation have been proposed and, although there may be small differences in the numerical values of the electronegativity, all methods show the same [[Periodicity|periodic trends]] between [[Chemical element|elements]].

The most commonly used method of calculation is that originally proposed by Pauling. This gives a [[dimension]]less quantity, commonly referred to as the '''Pauling scale''', on a relative scale running from 0.7 to 4.0 ([[hydrogen]]&nbsp;= 2.2). When other methods of calculation are used, it is conventional (although not obligatory) to quote the results on a scale which covers the same range of numerical values: this is known as an electronegativity in '''Pauling units'''.

Electronegativity, as it is usually calculated, is not strictly an [[atomic property]], but rather a property of an atom ''in a molecule'':<ref name="NOTCB">[[Linus Pauling|Pauling, Linus]] (1960). ''[[Nature of the Chemical Bond]]'' (3rd Edn.). [[Ithaca, NY]]: Cornell University Press. pp.&nbsp;88&ndash;107.</ref> the equivalent property of a free atom is its [[electron affinity]]. It is to be expected that the electronegativity of an element will vary with its chemical environment,<ref>Greenwood, N. N.; Earnshaw, A. (1984). ''Chemistry of the Elements''. Oxford: Pergamon. ISBN 0-08-022057-6. p.&nbsp;30.</ref> but it is usually considered to be a [[Transferability (chemistry)|transferable property]], that is to say that similar values will be valid in a variety of situations.

== Electronegativities of the elements ==
<center>
{|  WIDTH="80%"  align="CENTER"
!  colspan=20 | → Atomic radius decreases → [[Ionization energy]] increases → Electronegativity increases →
|-  align="CENTER"
| [[Periodic table group|'''Group''']] (vertical)
| [[Alkali metal|'''1''']]
| [[Alkaline earth|'''2''']]
| [[Group 3 element|'''3''']]
| [[Group 4 element|'''4''']]
| [[Group 5 element|'''5''']]
| [[Group 6 element|'''6''']]
| [[Group 7 element|'''7''']]
| [[Group 8 element|'''8''']]
| [[Group 9 element|'''9''']]
| [[Group 10 element|'''10''']]
| [[Coinage metal|'''11''']]
| [[Group 12 element|'''12''']]
| [[Boron group|'''13''']]
| [[Carbon group|'''14''']]
| [[Pnictogen|'''15''']]
| [[Chalcogen|'''16''']]
| [[Halogen|'''17''']]
| [[Noble gas|'''18''']]
|-  align="CENTER"
| [[Periodic table period|'''Period''']] (horizontal)
|  colspan=19 | 
|-  align="CENTER"
| [[Period 1 element|'''1''']]
|  bgcolor="#ff8a00" | [[Hydrogen|H]]<br/>2.20
|  colspan=16 | 
|  bgcolor="#bbbbbb" | [[Helium|He]]<br/>&nbsp;
|-  align="CENTER"
| [[Period 2 element|'''2''']]
|  bgcolor="#ffe900" | [[Lithium|Li]]<br/>0.98
|  bgcolor="#ffbb00" | [[Beryllium|Be]]<br/>1.57
|  colspan=10 | 
|  bgcolor="#ff9700" | [[Boron|B]]<br/>2.04
|  bgcolor="#ff6f00" | [[Carbon|C]]<br/>2.55
|  bgcolor="#ff4900" | [[Nitrogen|N]]<br/>3.04
|  bgcolor="#ff2a00" | [[Oxygen|O]]<br/>3.44
|  bgcolor="#ff0000" | [[Fluorine|F]]<br/>3.98
|  bgcolor="#bbbbbb" | [[Neon|Ne]]<br/>&nbsp;
|-  align="CENTER"
| [[Period 3 element|'''3''']]
|  bgcolor="#ffed00" | [[Sodium|Na]]<br/>0.93
|  bgcolor="#ffd000" | [[Magnesium|Mg]]<br/>1.31
|  colspan=10 | 
|  bgcolor="#ffb800" | [[Aluminium|Al]]<br/>1.61
|  bgcolor="#ffa200" | [[Silicon|Si]]<br/>1.90
|  bgcolor="#ff8b00" | [[Phosphorus|P]]<br/>2.19
|  bgcolor="#ff6d00" | [[Sulfur|S]]<br/>2.58
|  bgcolor="#ff4000" | [[Chlorine|Cl]]<br/>3.16
|  bgcolor="#bbbbbb" | [[Argon|Ar]]<br/>&nbsp;
|-  align="CENTER"
| [[Period 4 element|'''4''']]
|  bgcolor="#fff600" | [[Potassium|K]]<br/>0.82
|  bgcolor="#ffe800" | [[Calcium|Ca]]<br/>1.00
|  bgcolor="#ffcc00" | [[Scandium|Sc]]<br/>1.36
|  bgcolor="#ffbe00" | [[Titanium|Ti]]<br/>1.54
|  bgcolor="#ffb700" | [[Vanadium|V]]<br/>1.63
|  bgcolor="#ffb400" | [[Chromium|Cr]]<br/>1.66
|  bgcolor="#ffbd00" | [[Manganese|Mn]]<br/>1.55
|  bgcolor="#ffa700" | [[Iron|Fe]]<br/>1.83
|  bgcolor="#ffa300" | [[Cobalt|Co]]<br/>1.88
|  bgcolor="#ffa100" | [[Nickel|Ni]]<br/>1.91
|  bgcolor="#ffa200" | [[Copper|Cu]]<br/>1.90
|  bgcolor="#ffb500" | [[Zinc|Zn]]<br/>1.65
|  bgcolor="#ffa900" | [[Gallium|Ga]]<br/>1.81
|  bgcolor="#ff9900" | [[Germanium|Ge]]<br/>2.01
|  bgcolor="#ff8c00" | [[Arsenic|As]]<br/>2.18
|  bgcolor="#ff6f00" | [[Selenium|Se]]<br/>2.55
|  bgcolor="#ff4f00" | [[Bromine|Br]]<br/>2.96
|  bgcolor="#ff4c00" | [[Krypton|Kr]]<br/>3.00
|-  align="CENTER"
| [[Period 5 element|'''5''']]
|  bgcolor="#fff600" | [[Rubidium|Rb]]<br/>0.82
|  bgcolor="#ffec00" | [[Strontium|Sr]]<br/>0.95
|  bgcolor="#ffd700" | [[Yttrium|Y]]<br/>1.22
|  bgcolor="#ffce00" | [[Zirconium|Zr]]<br/>1.33
|  bgcolor="#ffb900" | [[Niobium|Nb]]<br/>1.6
|  bgcolor="#ff8d00" | [[Molybdenum|Mo]]<br/>2.16
|  bgcolor="#ffa200" | [[Technetium|Tc]]<br/>1.9
|  bgcolor="#ff8a00" | [[Ruthenium|Ru]]<br/>2.2
|  bgcolor="#ff8400" | [[Rhodium|Rh]]<br/>2.28
|  bgcolor="#ff8a00" | [[Palladium|Pd]]<br/>2.20
|  bgcolor="#ff9f00" | [[Silver|Ag]]<br/>1.93
|  bgcolor="#ffb200" | [[Cadmium|Cd]]<br/>1.69
|  bgcolor="#ffab00" | [[Indium|In]]<br/>1.78
|  bgcolor="#ff9d00" | [[Tin|Sn]]<br/>1.96
|  bgcolor="#ff9600" | [[Antimony|Sb]]<br/>2.05
|  bgcolor="#ff9200" | [[Tellurium|Te]]<br/>2.1
|  bgcolor="#ff6700" | [[Iodine|I]]<br/>2.66
|  bgcolor="#ff6b00" | [[Xenon|Xe]]<br/>2.6
|-  align="CENTER"
| [[Period 6 element|'''6''']]
|  bgcolor="#fff800" | [[Caesium|Cs]]<br/>0.79
|  bgcolor="#fff000" | [[Barium|Ba]]<br/>0.89
|  *<br/>&nbsp;
|  bgcolor="#ffd000" | [[Hafnium|Hf]]<br/>1.3
|  bgcolor="#ffc100" | [[Tantalum|Ta]]<br/>1.5
|  bgcolor="#ff7e00" | [[Tungsten|W]]<br/>2.36
|  bgcolor="#ffa200" | [[Rhenium|Re]]<br/>1.9
|  bgcolor="#ff8a20" | [[Osmium|Os]]<br/>2.2
|  bgcolor="#ff8a00" | [[Iridium|Ir]]<br/>2.20
|  bgcolor="#ff8400" | [[Platinum|Pt]]<br/>2.28
|  bgcolor="#ff7000" | [[Gold|Au]]<br/>2.54
|  bgcolor="#ff9a00" | [[Mercury (element)|Hg]]<br/>2.00
|  bgcolor="#ffb700" | [[Thallium|Tl]]<br/>1.62
|  bgcolor="#ff8000" | [[Lead|Pb]]<br/>2.33
|  bgcolor="#ff9800" | [[Bismuth|Bi]]<br/>2.02
|  bgcolor="#ff9a00" | [[Polonium|Po]]<br/>2.0
|  bgcolor="#ff8a00" | [[Astatine|At]]<br/>2.2
|  bgcolor="#bbbbbb" | [[Radon|Rn]]<br/>&nbsp;
|-  align=CENTER
| [[Period 7 element|'''7''']]
|  bgcolor="#ffff00" | [[Francium|Fr]]<br/>0.7
|  bgcolor="#ffef00" | [[Radium|Ra]]<br/>0.9
|  **<br/>&nbsp;
|  bgcolor="#bbbbbb" | [[Rutherfordium|Rf]]<br/>&nbsp;
|  bgcolor="#bbbbbb" | [[Dubnium|Db]]<br/>&nbsp;
|  bgcolor="#bbbbbb" | [[Seaborgium|Sg]]<br/>&nbsp;
|  bgcolor="#bbbbbb" | [[Bohrium|Bh]]<br/>&nbsp;
|  bgcolor="#bbbbbb" | [[Hassium|Hs]]<br/>&nbsp;
|  bgcolor="#bbbbbb" | [[Meitnerium|Mt]]<br/>&nbsp;
|  bgcolor="#bbbbbb" | [[Darmstadtium|Ds]]<br/>&nbsp;
|  bgcolor="#bbbbbb" | [[Roentgenium|Rg]]<br/>&nbsp;
|  bgcolor="#bbbbbb" | [[Ununbium|Uub]]<br/>&nbsp;
|  bgcolor="#bbbbbb" | [[Ununtrium|Uut]]<br/>&nbsp;
|  bgcolor="#bbbbbb" | [[Ununquadium|Uuq]]<br/>&nbsp;
|  bgcolor="#bbbbbb" | [[Ununpentium|Uup]]<br/>&nbsp;
|  bgcolor="#bbbbbb" | [[Ununhexium|Uuh]]<br/>&nbsp;
|  bgcolor="#bbbbbb" | [[Ununseptium|Uus]]<br/>&nbsp;
|  bgcolor="#bbbbbb" | [[Ununoctium|Uuo]]<br/>&nbsp;
|-  align=CENTER
|   
|-  align=CENTER
|  [[Lanthanides]]
|  *<br/>&nbsp;
|  bgcolor="#ffe000" | [[Lanthanum|La]]<br/>1.1
|  bgcolor="#ffde00" | [[Cerium|Ce]]<br/>1.12
|  bgcolor="#ffde00" | [[Praseodymium|Pr]]<br/>1.13
|  bgcolor="#ffdd00" | [[Neodymium|Nd]]<br/>1.14
|  bgcolor="#ffde00" | [[Promethium|Pm]]<br/>1.13
|  bgcolor="#ffda00" | [[Samarium|Sm]]<br/>1.17
|  bgcolor="#ffd800" | [[Europium|Eu]]<br/>1.2
|  bgcolor="#ffd800" | [[Gadolinium|Gd]]<br/>1.2
|  bgcolor="#ffe000" | [[Terbium|Tb]]<br/>1.1
|  bgcolor="#ffd700" | [[Dysprosium|Dy]]<br/>1.22
|  bgcolor="#ffd600" | [[Holmium|Ho]]<br/>1.23
|  bgcolor="#ffd500" | [[Erbium|Er]]<br/>1.24
|  bgcolor="#ffd400" | [[Thulium|Tm]]<br/>1.25
|  bgcolor="#ffe000" | [[Ytterbium|Yb]]<br/>1.1
|  bgcolor="#ffd300" | [[Lutetium|Lu]]<br/>1.27
|-  align=CENTER
|  [[Actinides]]
|  **<br/>&nbsp;
|  bgcolor="#ffe000" | [[Actinium|Ac]]<br/>1.1
|  bgcolor="#ffd000" | [[Thorium|Th]]<br/>1.3
|  bgcolor="#ffc100" | [[Protactinium|Pa]]<br/>1.5
|  bgcolor="#ffca00" | [[Uranium|U]]<br/>1.38
|  bgcolor="#ffcc00" | [[Neptunium|Np]]<br/>1.36
|  bgcolor="#ffd200" | [[Plutonium|Pu]]<br/>1.28
|  bgcolor="#ffde00" | [[Americium|Am]]<br/>1.13
|  bgcolor="#ffd200" | [[Curium|Cm]]<br/>1.28
|  bgcolor="#ffd000" | [[Berkelium|Bk]]<br/>1.3
|  bgcolor="#ffd000" | [[Californium|Cf]]<br/>1.3
|  bgcolor="#ffd000" | [[Einsteinium|Es]]<br/>1.3
|  bgcolor="#ffd000" | [[Fermium|Fm]]<br/>1.3
|  bgcolor="#ffd000" | [[Mendelevium|Md]]<br/>1.3
|  bgcolor="#ffd000" | [[Nobelium|No]]<br/>1.3
|  bgcolor="#bbbbbb" | [[Lawrencium|Lr]]<br/>&nbsp;
|-  align="CENTER"
|  colspan=20 | 
|}</center>
<center>[[Periodic table]] of electronegativity using the Pauling scale</center>

<center>See also [[Periodic table]]</center>

== Methods of calculation ==
=== Pauling electronegativity ===<!-- This section is linked from [[Rhenium diboride]] -->
[[Linus Pauling|Pauling]] first proposed<ref name="paulingJACS"/> the concept of electronegativity in 1932 as an explanation of the fact that the covalent bond between two different atoms (A&ndash;B) is stronger than would be expected by taking the average of the strengths of the A&ndash;A and B&ndash;B bonds. According to [[valence bond theory]], of which Pauling was a notable proponent, this "additional stabilization" of the heteronuclear bond is due to the contribution of ionic [[canonical form]]s to the bonding.

The difference in electronegativity between atoms A and B is given by:
::<math>\chi_{\rm A} - \chi_{\rm B} = ({\rm eV})^{-1/2} \sqrt{E_{\rm d}({\rm AB}) - [E_{\rm d}({\rm AA}) + E_{\rm d}({\rm BB})]/2}</math>
where the [[Bond dissociation energy|dissociation energies]], ''E''<sub>d</sub>, of the A&ndash;B, A&ndash;A and B&ndash;B bonds are expressed in [[electronvolt]]s, the factor (eV)<sup>&minus;½</sup> being included to ensure a dimensionless result. Hence, the difference in Pauling electronegativity between [[hydrogen]] and [[bromine]] is 0.73 (dissociation energies: H&ndash;Br, 3.79&nbsp;eV; H&ndash;H, 4.52&nbsp;eV; Br&ndash;Br 2.00&nbsp;eV)

As only differences in electronegativity are defined, it is necessary to choose an arbitrary reference point in order to construct a scale. Hydrogen was chosen as the reference, as it forms covalent bonds with a large variety of elements: its electronegativity was fixed first<ref name="paulingJACS"/> at 2.1, later revised<ref name="Allred">Allred, A. L. (1961). [[Journal of Inorganic and Nuclear Chemistry|''J. Inorg. Nucl. Chem.'']] '''17''':215.</ref> to 2.20. It is also necessary to decide which of the two elements is the more electronegative (equivalent to choosing one of the two possible signs for the square root). This is done by "chemical intuition": in the above example, [[hydrogen bromide]] dissolves in water to form H<sup>+</sup> and Br<sup>&minus;</sup> ions, so it may be assumed that bromine is more electronegative than hydrogen.

To calculate a Pauling electronegativity for an element, it is necessary to have data on the dissociation energies of at least two types of covalent bond formed by that element. Allred updated Pauling's original values in 1961 to take account of the greater availability of thermodynamic data,<ref name="Allred"/> and it is these "revised Pauling" values of the electronegativity which are most usually used.

=== Mulliken electronegativity ===
[[Image:Pauling and Mullikan electronegativities.png|thumb|right|300px|The correlation between Mulliken electronegativities (''x''-axis, in kJ/mol) and Pauling electronegativities (''y''-axis).]]
[[Robert S. Mulliken|Mulliken]] proposed that the [[arithmetic mean]] of the first [[ionization energy]] and the [[electron affinity]] should be a measure of the tendency of an atom to attract electrons.<ref>Mulliken, R. S. (1934). [[Journal of Chemical Physics|''J. Chem. Phys.'']] '''2''':782. Mulliken, R. S. (1935). [[Journal of Chemical Physics|''J. Chem. Phys.'']] '''3''':573.</ref> As this definition is not dependent on an arbitrary relative scale, it has also been termed '''absolute electronegativity''',<ref>Pearson, R. G. (1985). [[Journal of the American Chemical Society|''J. Am. Chem. Soc.'']] '''107''':6801.</ref> with the units of [[Joule per mole|kilojoules per mole]] or [[electronvolt]]s.

However, it is more usual to make use of a linear transformation to transform these absolute values into values which resemble the more familiar Pauling values. For ionization energies and electron affinities in electronvolts,<ref>Huheey, J. E. (1978). ''Inorganic Chemistry'' (2nd Edn.). New York: Harper & Row. p.&nbsp;167.</ref>
::<math>\chi = 0.187(E_{\rm i} + E_{\rm ea}) + 0.17</math>
and for energies in kilojoules per mole,<ref>This second relation has been recalculated using the best values of the first ionization energies and electron affinities available in 2006.</ref>
::<math>\chi = (1.97\times 10^{-3})(E_{\rm i} + E_{\rm ea}) + 0.19</math>

The Mulliken electronegativity can only be calculated for an element for which the electron affinity is known, fifty-seven elements as of 2006.

=== Allred-Rochow electronegativity ===
[[Image:Pauling and Allred-Rochow electronegativities.png|thumb|right|300px|The correlation between Allred-Rochow electronegativities (''x''-axis, in Å<sup>&minus;2</sup>) and Pauling electronegativities (''y''-axis).]]
Allred and Rochow considered<ref>Allred, A. L.; Rochow, E. G. (1958). [[Journal of Inorganic and Nuclear Chemistry|''J. Inorg. Nucl. Chem.'']] '''5''':264.</ref> that electronegativity should be related to the charge experienced by an electron on the "surface" of an atom: the higher the charge per unit area of atomic surface, the greater the tendency of that atom to attract electrons. The effective nuclear charge, ''Z''* experienced by [[valence electron]]s can be estimated using [[Slater's rules]], while the surface area of an atom in a molecule can be taken to be proportional to the square of the [[covalent radius]], ''r''<sub>cov</sub>. When ''r''<sub>cov</sub> is expressed in [[ångström]]s,
::<math>\chi = 0.359{{Z\star}\over{r^2_{\rm cov}}} + 0.744</math>.

=== Sanderson electronegativity ===
[[Image:Pauling and Sanderson electronegativities.png|thumb|right|300px|The correlation between Sanderson electronegativities (''x''-axis, arbitrary units) and Pauling electronegativities (''y''-axis).]]
Sanderson has also noted the relationship between electronegativity and atomic size, and has proposed a method of calculation based on the reciprocal of the atomic volume.<ref>Sanderson, R. T. (1983). "[http://pubs.acs.org/cgi-bin/archive.cgi/jacsat/1983/105/i08/pdf/ja00346a026.pdf Electronegativity and bond energy.]" [[Journal of the American Chemical Society|''J. Am. Chem. Soc.'']] '''105''':2259.</ref> With a knowledge of bond lengths, Sanderson electronegativities allow the estimation of bond energies in a wide range of compounds.<ref>Sanderson, R. T. (1983). ''Polar Covalence''. New York: Academic Press.</ref>

=== Allen electronegativity ===
[[Image:Pauling and Allen electronegativities.png|thumb|right|300px|The correlation between Allen electronegativities (''x''-axis, in kJ/mol) and Pauling electronegativities (''y''-axis).]]
Perhaps the simplest definition of electronegativity is that of Allen, who has proposed that it is related to the average energy of the [[valence electron]]s in a free atom,<ref>Allen, L. C. (1989). [[Journal of the American Chemical Society|''J. Am. Chem. Soc.'']] '''111''':9003.</ref>
::<math>\chi = {n_{\rm s}\epsilon_{\rm s} + n_{\rm p}\epsilon_{\rm p} \over n_{\rm s} + n_{\rm p}}</math>
where &epsilon;<sub>s,p</sub> are the one-electron energies of s- and p-electrons in the free atom and ''n''<sub>s,p</sub> are the number of s- and p-electrons in the valence shell. It is usual to apply a scaling factor, 1.75&times;10<sup>&minus;3</sup> for energies expressed in kilojoules per mole or 0.169 for energies measured in electronvolts, to give values which are numerically similar to Pauling electronegativities.

The one-electron energies can be determined directly from [[Spectroscopy|spectroscopic data]], and so electronegativities calculated by this method are sometimes referred to as '''spectroscopic electronegativities'''. The necessary data are available for almost all elements, and this method allows the estimation of electronegativities for elements which cannot be treated by the other methods, e.g. [[francium]], which has an Allen electronegativity of 0.67.<ref name="Fr">The widely quoted Pauling electronegativity of 0.7 for francium is an extrapolated value of uncertain provenance. The Allen electronegativity of [[caesium]] is 0.66.</ref> However, it is not clear what should be considered to be valence electrons for the d- and f-block elements, which leads to an ambiguity for their electronegativities calculated by the Allen method.

== Correlation of electronegativity with other properties ==
[[Image:Sn-119 isomer shifts in hexahalostannates.png|thumb|right|300px|The variation of the isomer shift (''y''-axis, in mm/s) of [SnX<sub>6</sub>]<sup>2&minus;</sup> anions, as measured by <sup>119</sup>Sn [[Mössbauer spectroscopy]], against the sum of the Pauling electronegativities of the halide substituents (''x''-axis).]]
The wide variety of methods of calculation of electronegativities, which all give results which correlate well with one another, is one indication of the number of chemical properties which might be affected by electronegativity. The most obvious application of electronegativities is in the discussion of [[bond polarity]], for which the concept was introduced by Pauling. In general, the greater the difference in electronegativity between two atoms, the more polar the bond that will be formed between them, with the atom having the higher electronegativity being at the negative end of the dipole. Pauling proposed an equation to relate "ionic character" of a bond to the difference in electronegativity of the two atoms,<ref name="NOTCB"/> although this has fallen somewhat into disuse.

Several correlations have been shown between [[Infrared spectroscopy|infrared stretching frequencies]] of certain bonds and the electronegativities of the atoms involved:<ref>See, e.g., Bellamy, L. J. (1958). ''The Infra-Red Spectra of Complex Molecules'' (2nd Edn.). New York: Wiley. p.&nbsp;392.</ref> however, this is not surprising as such strectching frquencies depend in part on bond strength, which enters into the calculation of Pauling electronegativities. More convincing are the correlations between electronegativity and chemical shifts in [[NMR spectroscopy]]<ref>Spieseke, H.; Schneider, W. G. (1961). [[Journal of Chemical Physics|''J. Chem. Phys.'']] '''35''':722.</ref> or isomer shifts in [[Mössbauer spectroscopy]]<ref>Clasen, C. A.; Good, M. L. (1970). [[Inorganic Chemistry (journal)|''Inorg. Chem.'']] '''9''':817.</ref> (see figure). Both these measurements depend on the s-electron density at the nucleus, and so are a good indication that the different measures of electronegativity really are describing "the ability of an atom in a molecule to attract electrons to itself".<ref name="definition"/><ref name="NOTCB"/>

== Trends in electronegativity ==
=== Periodic trends ===
[[Image:Periodic variation of Pauling electronegativities.png|thumb|right|300px|The variation of Pauling electronegativity (''y''-axis) as one descends the main groups of the Periodic table from the second period to the sixth period.]]
In general, electronegativity increases on passing from left to right along a period, and decreases on descending a group. Hence, [[fluorine]] is undoubtedly the most electronegative of the elements while [[caesium]] is the least electronegative, at least of those elements for which substantial data are available.<ref name="Fr"/>

There are some exceptions to this general rule. [[Gallium]] and [[germanium]] have higher electronegativities than [[aluminium]] and [[silicon]] respectively because of the [[d-block contraction]]. Elements of the [[fourth period]] immediately after the first row of the transition metals have unusually small atomic radii because the 3d-electrons are not effective at shielding the increased nuclear charge, and smaller atomic size correlates with higher electronegativity (see [[#Allred-Rochow electronegativity|Allred-Rochow electronegativity]], [[#Sanderson electronegativity|Sanderson electronegativity]] above). The anomalously high electronegativity of [[lead]], particularly when compared to [[thallium]] and [[bismuth]], appears to be an artifact of data selection (and data availability)&mdash;methods of calculation other than the Pauling method show the normal periodic trends for these elements.

=== Variation of electronegativity with oxidation number ===
In inorganic chemistry it is common to consider a single value of the electronegativity to be valid for most "normal" situations. While this approach has the advantage of simplicity, it is clear that the electronegativity of an element is ''not'' an invariable atomic property and, in particular, increases with the [[oxidation state]] of the element.

Allred used the Pauling method to calculate separate electronegativities for different oxidation states of the handful of elements (including tin and lead) for which sufficient data was available.<ref name="Allred"/> However, for most elements, there are not enough different covalent compounds for which bond dissociation energies are known to make this approach feasible. This is particularly true of the transition elements, where quoted electronegativity values are usually, of necessity, averages over several different oxidation states and where trends in electronegativity are harder to see as a result.

{| class="wikitable" align=left
|-
! Acid
! Formula
! Chlorine<br/>oxidation<br/>state
! p''K''<sub>a</sub>
|-
| [[Hypochlorous acid]]
| HClO
| align=center | +1
| align=right | +7.5
|-
| [[Chlorous acid]]
| HClO<sub>2</sub>
| align=center | +3
| align=right | +2.0
|-
| [[Chloric acid]]
| HClO<sub>3</sub>
| align=center | +5
| align=right | &minus;1.0
|-
| [[Perchloric acid]]
| HClO<sub>4</sub>
| align=center | +7
| align=right | &minus;10&nbsp;
|-
|}

The chemical effects of this increase in electronegativity can be seen both in the structures of oxides and halides and in the acidity of oxides and oxoacids. Hence [[Chromium trioxide|CrO<sub>3</sub>]] and [[Dimanganese heptoxide|Mn<sub>2</sub>O<sub>7</sub>]] are [[acidic oxide]]s with low [[melting point]]s, while [[Chromium(III) oxide|Cr<sub>2</sub>O<sub>3</sub>]] is [[amphoteric oxide|amphoteric]] and [[Manganese(III) oxide|Mn<sub>2</sub>O<sub>3</sub>]] is a completely [[basic oxide]]. 

The effect can also be clearly seen in the [[Acid dissociation constant|dissociation constants]] of the [[oxoacid]]s of [[chlorine]]. The effect is much larger than could be explained by the negative charge being shared among a larger number of oxygen atoms, which would lead to a difference in p''K''<sub>a</sub> of log<sub>10</sub>(¼)&nbsp;= &minus;0.6 between [[hypochlorous acid]] and [[perchloric acid]]. As the oxidation state of the central chlorine atom increases, more electron density is drawn from the oxygen atoms onto the chlorine, reducing the partial negative charge on the oxygen atoms and increasing the acidity.

== Group electronegativity ==
{{main|Electronic effect of substituents}}
In organic chemistry, electronegativity is associated more with different functional groups than with individual atoms. The terms '''group electronegativity''' and '''substituent electronegativity''' are used synonymously. However, it is common to distinguish between the [[inductive effect]] and the [[resonance effect]], which might be described as &sigma;- and &pi-electronegativities respectively. There are a number of [[linear free energy relationship]]s which have been used to quantify these effects, of which the [[Hammett equation]] is the best known. [[Kabachnik parameter]]s are group electronegativities for use in [[organophosphorus chemistry]].

== See also ==
*[[Electronegativities of the elements (data page)]]

== References ==
*Jolly, William L. (1991). ''Modern Inorganic Chemistry'' (2nd Edn.). New York: [[McGraw-Hill]]. ISBN 0-07-112651-1. pp.&nbsp;71&ndash;76.
*Mullay, J. (1987). "Estimation of atomic and group electronegativities." [[Structure and Bonding|''Struct. Bond.'']] '''66''':1&ndash;25.
<references />

== External links ==
{{commonscat|Electronegativity}}
*[http://www.webelements.com/ WebElements], lists values of electronegativities by a number of different methods of calculation

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[[Category:Chemical properties]]

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