Editing Electronegativity

{{Short description|Tendency of an atom to attract a shared pair of electrons}}
{{Redirect|Electronegative|the Nightfall EP|Electronegative (EP){{!}}''Electronegative'' (EP)}}
[[Image:Electrostatic Potential.jpg|thumb|alt=A water molecule is put into a see-through egg shape, which is color-coded by electrostatic potential. A concentration of red is near the top of the shape, where the oxygen atom is, and gradually shifts through yellow, green, and then to blue near the lower-right and lower-left corners of the shape where the hydrogen atoms are.|upright=1.5|right|Electrostatic potential map of a water molecule, where the oxygen atom has a more negative charge (red) than the positive (blue) hydrogen atoms]]
'''Electronegativity''', symbolized as <span class="nounderlines">''[[Chi (letter)|χ]]''</span>, is the tendency for an [[atom]] of a given [[chemical element]] to attract shared [[electron]]s (or [[electron density]]) when forming a [[chemical bond]].<ref name="definition">{{GoldBookRef|file=E01990|title=Electronegativity}}</ref> An atom's electronegativity is affected by both its [[atomic number]] and the distance at which its [[valence electrons]] reside from the charged nucleus. The higher the associated electronegativity, the more an atom or a substituent group attracts electrons. Electronegativity serves as a simple way to quantitatively estimate the [[bond energy]], and the sign and magnitude of a bond's [[chemical polarity]], which characterizes a bond along the continuous scale from [[Covalent bonding|covalent]] to [[ionic bonding]]. The loosely defined term '''electropositivity''' is the opposite of electronegativity: it characterizes an element's tendency to donate valence electrons.

On the most basic level, electronegativity is determined by factors like the [[effective nuclear charge|nuclear charge]] (the more [[protons]] an atom has, the more "pull" it will have on electrons) and the number and location of other electrons in the [[Electron shell|atomic shells]] (the more electrons an atom has, the farther from the [[Atomic nucleus|nucleus]] the valence electrons will be, and as a result, the less positive charge they will experience—both because of their increased distance from the nucleus and because the other electrons in the lower energy core [[atomic orbital|orbitals]] will act to [[Shielding effect|shield]] the valence electrons from the positively charged nucleus).

The term "electronegativity" was introduced by [[Jöns Jacob Berzelius]] in 1811,<ref name="Jensen">{{cite journal |author= Jensen, W.B.
 |author-link=William B. Jensen
 |year= 1996
 |journal= [[Journal of Chemical Education]]
 |volume= 73
 |issue= 1
 |pages= 11–20
 |title= Electronegativity from Avogadro to Pauling: Part 1: Origins of the Electronegativity Concept
 |doi= 10.1021/ed073p11|bibcode = 1996JChEd..73...11J }}</ref>
though the concept was known before that and was studied by many chemists including [[Amadeo Avogadro|Avogadro]].<ref name="Jensen"/>
In spite of its long history, an accurate scale of electronegativity was not developed until 1932, when [[Linus Pauling]] proposed an electronegativity scale which depends on bond energies, as a development of [[valence bond theory]].<ref name="paulingJACS">{{cite journal |author= Pauling, L.
 |author-link=Linus Pauling
 |year= 1932
 |journal= [[Journal of the American Chemical Society]]
 |volume= 54
 |issue= 9
 |pages= 3570–3582
 |title= The Nature of the Chemical Bond. IV. The Energy of Single Bonds and the Relative Electronegativity of Atoms
 |doi= 10.1021/ja01348a011
 }}</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 [[periodic trends]] between [[Chemical element|elements]].<ref>{{Cite journal |last=Sproul |first=Gordon D. |date=2020-05-26 |title=Evaluation of Electronegativity Scales |url=https://doi.org/10.1021/acsomega.0c00831 |journal=ACS Omega |volume=5 |issue=20 |pages=11585–11594 |doi=10.1021/acsomega.0c00831|pmid=32478249 |pmc=7254809 }}</ref>

The most commonly used method of calculation is that originally proposed by Linus Pauling. This gives a [[dimensionless quantity]], commonly referred to as the '''Pauling scale''' (''χ''<sub>r</sub>), on a relative scale running from 0.79 to 3.98 ([[hydrogen]]&nbsp;= 2.20). When other methods of calculation are used, it is conventional (although not obligatory) to quote the results on a scale that covers the same range of numerical values: this is known as an electronegativity in ''Pauling units''.

As it is usually calculated, electronegativity is not a property of an atom alone, but rather a property of an atom in a [[molecule]].<ref name="NOTCB">{{cite book|author=Pauling, Linus|year=1960|title=Nature of the Chemical Bond|url=https://archive.org/details/natureofchemical00paul|url-access=registration|publisher=Cornell University Press|pages=[https://archive.org/details/natureofchemical00paul/page/88 88–107]|isbn=978-0-8014-0333-0}}</ref> Even so, the electronegativity of an atom is strongly correlated with the [[Ionization energy|first ionization energy]]. The electronegativity is slightly negatively correlated (for smaller electronegativity values) and rather strongly positively correlated (for most and larger electronegativity values) with the [[electron affinity]].<ref>"[https://web.archive.org/web/20190621000348/https://pubchem.ncbi.nlm.nih.gov/periodic-table/#view=table&property=Electronegativity PubChem ElectroNegativity]," Downloaded, line-graphed, and correlated 'Electronegativity' with 'ElectronAffinity', showing a rather strong positive correlation of '0.712925965'. (Accessed linked site on 2023-09-16.)</ref> It is to be expected that the electronegativity of an element will vary with its chemical environment,<ref>{{cite book|author1=Greenwood, N. N. |author2=Earnshaw, A. |year=1984|title=Chemistry of the Elements|publisher=Pergamon|isbn=978-0-08-022057-4|page=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.

[[Caesium]] is the least electronegative element (0.79);<!--NOT FRANCIUM SEE BELOW--> [[fluorine]] is the most (3.98).

==Methods of calculation==

===Pauling electronegativity===
{{Anchor|Pauling electronegativity}}

[[Linus Pauling|Pauling]] first proposed<ref name="paulingJACS"/> the concept of electronegativity in 1932 to explain why the [[covalent bond]] between two different atoms (A–B) is stronger than the average of the A–A and the B–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 bond|ionic]]  [[Resonance (chemistry)|canonical form]]s to the bonding.

The difference in electronegativity between atoms A and B is given by:
<math display="block">|\chi_{\rm A} - \chi_{\rm B}| = ({\rm eV})^{-1/2} \sqrt{E_{\rm d}({\rm AB}) - \frac{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–B, A–A and B–B bonds are expressed in [[electronvolt]]s, the factor (eV)<sup>−{{frac|1|2}}</sup> being included to ensure a dimensionless result. Hence, the difference in Pauling electronegativity between hydrogen and [[bromine]] is 0.73 (dissociation energies: H–Br, 3.79&nbsp;eV; H–H, 4.52&nbsp;eV; Br–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">{{cite journal
 |author= Allred, A. L.
 |year= 1961
 |journal= Journal of Inorganic and Nuclear Chemistry
 |volume= 17
 |issue= 3–4
 |pages= 215–221
 |title= Electronegativity values from thermochemical data
 |doi= 10.1016/0022-1902(61)80142-5}}</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 usually done using "chemical intuition": in the above example, [[hydrogen bromide]] dissolves in water to form H<sup>+</sup> and Br<sup>−</sup> ions, so it may be assumed that bromine is more electronegative than hydrogen. However, in principle, since the same electronegativities should be obtained for any two bonding compounds, the data are in fact overdetermined, and the signs are unique once a reference point has been fixed (usually, for H or F).

To calculate Pauling electronegativity for an element, it is necessary to have data on the dissociation energies of at least two types of covalent bonds formed by that element. A. L. 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 that are most often used.

The essential point of Pauling electronegativity is that there is an underlying, quite accurate, semi-empirical formula for dissociation energies, namely:
<math display="block">E_{\rm d}({\rm AB}) = \frac{E_{\rm d}({\rm AA}) + E_{\rm d}({\rm BB})} 2 + (\chi_{\rm A} - \chi_{\rm B})^2 {\rm eV}</math>
or sometimes, a more accurate fit
<math display="block">E_{\rm d}({\rm AB}) =\sqrt{E_{\rm d}({\rm AA}) E_{\rm d}({\rm BB})}+1.3(\chi_{\rm A} - \chi_{\rm B})^2 {\rm eV}</math>

These are approximate equations but they hold with good accuracy. Pauling obtained the first equation by noting that a bond can be approximately represented as a quantum mechanical superposition of a covalent bond and two ionic bond-states. The covalent energy of a bond is approximate, by quantum mechanical calculations, the [[geometric mean]] of the two energies of covalent bonds of the same molecules, and there is additional energy that comes from ionic factors, i.e. polar character of the bond.

The geometric mean is approximately equal to the [[arithmetic mean]]—which is applied in the first formula above—when the energies are of a similar value, e.g., except for the highly electropositive elements, where there is a larger difference of two dissociation energies; the geometric mean is more accurate and almost always gives positive excess energy, due to ionic bonding.  The square root of this excess energy, Pauling notes, is approximately additive, and hence one can introduce the electronegativity. Thus, it is these semi-empirical formulas for bond energy that underlie the concept of Pauling electronegativity.

The formulas are approximate, but this rough approximation is in fact relatively good and gives the right intuition, with the notion of the polarity of the bond and some theoretical grounding in quantum mechanics. The electronegativities are then determined to best fit the data.

In more complex compounds, there is an additional error since electronegativity depends on the molecular environment of an atom. Also, the energy estimate can be only used for single, not for multiple bonds. The [[standard enthalpy of formation|enthalpy of formation]] of a molecule containing only single bonds can subsequently be estimated based on an electronegativity table, and it depends on the constituents and the sum of squares of differences of electronegativities of all pairs of bonded atoms. Such a formula for estimating energy typically has a relative error on the order of 10% but can be used to get a rough qualitative idea and understanding of a molecule.

{{Periodic table (electronegativities)}}

===Mulliken electronegativity===
{{Anchor|Mulliken electronegativity}}

[[Image:Pauling and Mullikan electronegativities.png|thumb|upright=1.35|The correlation between Mulliken electronegativities (''x''-axis, in kJ/mol) and Pauling electronegativities (''y''-axis).]]
[[Robert S. Mulliken]] proposed that the [[arithmetic mean]] of the first [[ionization energy]] (E<sub>i</sub>) and the [[electron affinity]] (E<sub>ea</sub>) should be a measure of the tendency of an atom to attract electrons:<ref>{{cite journal |author = Mulliken, R. S.|year =1934 |journal = [[Journal of Chemical Physics]] |volume = 2 |title = A New Electroaffinity Scale; Together with Data on Valence States and on Valence Ionization Potentials and Electron Affinities |doi = 10.1063/1.1749394 |pages = 782–793 |issue = 11|bibcode = 1934JChPh...2..782M }}</ref><ref>{{cite journal |author= Mulliken, R. S. |year=1935 |title = Electronic Structures of Molecules XI. Electroaffinity, Molecular Orbitals and Dipole Moments |journal = [[Journal of Chemical Physics|J. Chem. Phys.]] |volume = 3 |doi = 10.1063/1.1749731 |pages = 573–585
|issue = 9|bibcode = 1935JChPh...3..573M }}</ref>
<math display="block">\chi = \frac{E_{\rm i} + E_{\rm ea}} 2 </math>

As this definition is not dependent on an arbitrary relative scale, it has also been termed '''absolute electronegativity''',<ref>{{cite journal |author=Pearson, R. G. |year=1985 |title=Absolute electronegativity and absolute hardness of Lewis acids and bases |journal=[[Journal of the American Chemical Society|J. Am. Chem. Soc.]] |volume=107 |issue=24 |pages=6801–6806 |doi=10.1021/ja00310a009}}</ref> with the units of [[Joule per mole|kilojoules per mole]] or [[electronvolt]]s. However, it is more usual to use a linear transformation to transform these absolute values into values that resemble the more familiar Pauling values. For ionization energies and electron affinities in electronvolts,<ref>{{cite book |last1=Huheey |first1=J.E. |last2=Keiter |first2=E.A. |last3=Keiter |first3=R.L. |date=December 1, 2008 |orig-year=1978 |chapter=17 |editor1-last=Kauffman |editor1-first=G.B. |title=Inorganic Chemistry: Principles of Structure and Reactivity |url=https://www.pdfdrive.com/inorganic-chemistry-principles-of-structure-and-reactivity-e175855674.html |type=digitalized |language=en |edition=3rd |location=New York |publication-date=1900 |page=167 |doi=10.1021/ed050pA379.1 |isbn=9780060429874 |oclc=770736023 |id= inorganicchemist00huhe_0 |archive-url=https://web.archive.org/web/20190908204147/https://www.pdfdrive.com/inorganic-chemistry-principles-of-structure-and-reactivity-e175855674.html |archive-date=September 8, 2019 |access-date=December 15, 2020 |via=Oxford University Press}} [https://archive.org/details/inorganicchemist00huhe_0 Alt URL]</ref>
<math display="block">\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 display="block">\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 whose electron affinity is known. [[Electron affinity (data page)|Measured values are available]] for 72 elements, while approximate values have been [[Electron affinity (data page)|estimated or calculated]] for the remaining elements.

The Mulliken electronegativity of an atom is sometimes said to be the negative of the [[chemical potential]].<ref>{{cite journal |last1=Franco-Pérez |first1=Marco |last2=Gázquez |first2=José L. |title=Electronegativities of Pauling and Mulliken in Density Functional Theory |journal=Journal of Physical Chemistry A |date=31 October 2019 |volume=123 |issue=46 |pages=10065–10071 |doi=10.1021/acs.jpca.9b07468 |pmid=31670960 |bibcode=2019JPCA..12310065F |s2cid=207814569 }}</ref>  By inserting the energetic definitions of the ionization potential and electron affinity into the Mulliken electronegativity, it is possible to show that the Mulliken chemical potential is a finite difference approximation of the electronic energy with respect to the number of electrons., i.e.,
<math display="block">\mu(\rm Mulliken) = -\chi(\rm Mulliken) = {}-\frac{E_{\rm i} + E_{\rm ea}} 2 </math>

===Allred–Rochow electronegativity===
{{Anchor|Allred–Rochow electronegativity}}

[[Image:Pauling and Allred-Rochow electronegativities.png|thumb|upright=1.35|The correlation between Allred–Rochow electronegativities (''x''-axis, in Å<sup>−2</sup>) and Pauling electronegativities (''y''-axis).]]
[[Albert L. Allred|A. Louis Allred]] and [[Eugene G. Rochow]] considered<ref>{{cite journal|author1=Allred, A. L. |author2=Rochow, E. G. |year=1958| journal=Journal of Inorganic and Nuclear Chemistry|volume=5|pages=264–268|doi=10.1016/0022-1902(58)80003-2|title=A scale of electronegativity based on electrostatic force|issue=4}}</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''<sub>eff</sub>, 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 [[picometre]]s,<ref>{{cite book |last1=Housecroft |first1=C.E. |last2=Sharpe |first2=A.G. |title=Inorganic Chemistry |url=https://www.pearson.com/us/higher-education/program/Housecroft-Inorganic-Chemistry-5th-Edition/PGM2178749.html |type=eBook |language=en |volume=3 |edition=15th |location=Switzerland |publisher=Pearson Prentice-Hall |publication-date=November 1, 1993 |page=38 |doi=10.1021/ed070pA304.1|isbn=9780273742753 |via=University of Basel}} [https://archive.org/details/Inorganic_Chemistry_4th_edition_by_Catherine_Housecroft_Alan_G._Sharpe Alt URL]</ref>
<math display="block">\chi = 3590{{Z_{\rm eff}}\over{r^2_{\rm cov}}} + 0.744</math>

===Sanderson electronegativity equalization===
[[Image:Pauling and Sanderson electronegativities.png|thumb|upright=1.35|The correlation between Sanderson electronegativities (''x''-axis, arbitrary units) and Pauling electronegativities (''y''-axis).]]
[[Robert Thomas Sanderson|R.T. Sanderson]] has also noted the relationship between Mulliken electronegativity and atomic size, and has proposed a method of calculation based on the reciprocal of the atomic volume.<ref>{{cite journal|author=Sanderson, R. T. |year=1983 |title=Electronegativity and bond energy| journal=Journal of the American Chemical Society|volume=105|pages=2259–2261|doi=10.1021/ja00346a026 |issue=8}}</ref> With a knowledge of bond lengths, Sanderson's model allows the estimation of bond energies in a wide range of compounds.<ref>{{cite book|author=Sanderson, R. T.|year=1983|title=Polar Covalence|location=New York|publisher=Academic Press|isbn=978-0-12-618080-0|url-access=registration|url=https://archive.org/details/polarcovalence0000sand}}</ref> Sanderson's model has also been used to calculate molecular geometry, ''s''-electron energy, [[NMR]] spin-spin coupling constants and other parameters for organic compounds.<ref>{{cite journal |last1=Zefirov |first1=N. S. |first2=M. A. |last2=Kirpichenok |first3=F. F. |last3=Izmailov |first4=M. I. |last4=Trofimov |title=Calculation schemes for atomic electronegativities in molecular graphs within the framework of Sanderson principle |journal=[[Doklady Akademii Nauk SSSR]] |year=1987 |volume=296 |pages=883–887}}</ref><ref>{{cite journal |doi=10.1007/s11172-006-0105-6|title=Application of the electronegativity indices of organic molecules to tasks of chemical informatics |year=2005|author=Trofimov, M. I.|journal=Russian Chemical Bulletin|volume=54|pages=2235–2246|last2=Smolenskii|first2=E. A.|issue=9|s2cid=98716956 }}</ref> This work underlies the concept of '''electronegativity equalization''', which suggests that electrons distribute themselves around a molecule to minimize or to equalize the Mulliken electronegativity.<ref name= Lipkowitz>

{{cite book |title=Reviews in computational chemistry |author1=SW Rick |author2=SJ Stuart |chapter=Electronegativity equalization models |editor1=Kenny B. Lipkowitz |editor2=Donald B. Boyd |chapter-url=https://books.google.com/books?id=IqWXSLz6QE8C&pg=PA106 |page=106 |isbn=978-0-471-21576-9 |year=2002 |publisher=Wiley}}</ref> This behavior is analogous to the equalization of chemical potential in macroscopic thermodynamics.<ref name=Parr>{{cite book |title=Density-functional theory of atoms and molecules |author1=Robert G. Parr |author2=Weitao Yang |url=https://books.google.com/books?id=mGOpScSIwU4C&pg=PA91 |page=91 |isbn=978-0-19-509276-9 |year=1994 |publisher=Oxford University Press}}</ref>

===Allen electronegativity===
[[Image:Pauling and Allen electronegativities.png|thumb|upright=1.35|The correlation between Allen electronegativities (''x''-axis, in kJ/mol) and Pauling electronegativities (''y''-axis).]]
Perhaps the simplest definition of electronegativity is that of Leland C. Allen, who has proposed that it is related to the average energy of the [[valence electron]]s in a free atom,<ref>{{cite journal |doi=10.1021/ja00207a003 |title=Electronegativity is the average one-electron energy of the valence-shell electrons in ground-state free atoms|year=1989|author=Allen, Leland C.|journal=Journal of the American Chemical Society |volume=111|pages=9003–9014 |issue=25}}</ref><ref>{{cite journal|doi=10.1021/ja992866e|title=Configuration Energies of the Main Group Elements|year=2000|author=Mann, Joseph B. |author2=Meek, Terry L. |author3=Allen, Leland C. |journal=Journal of the American Chemical Society |volume=122 |pages=2780–2783|issue=12}}</ref><ref>{{cite journal|doi=10.1021/ja9928677|title=Configuration energies of the d-block elements |year=2000|author=Mann, Joseph B. |author2=Meek, Terry L. |author3=Knight, Eugene T. |author4=Capitani, Joseph F. |author5=Allen, Leland C. |journal=Journal of the American Chemical Society|volume=122|pages=5132–5137|issue=21}}</ref>

<math display="block">\chi = {n_{\rm s}\varepsilon_{\rm s} + n_{\rm p}\varepsilon_{\rm p} \over n_{\rm s} + n_{\rm p}}</math>

where ''ε''<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.

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 that 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.

On this scale, [[neon]] has the highest electronegativity of all elements, followed by [[fluorine]], [[helium]], and [[oxygen]].

{{periodic table (electronegativity by Allen scale)}}

==Correlation of electronegativity with other properties==
[[Image:Sn-119 isomer shifts in hexahalostannates.png|thumb|upright=1.35|The variation of the isomer shift (''y''-axis, in mm/s) of [SnX<sub>6</sub>]<sup>2−</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 that correlate well with one another, is one indication of the number of chemical properties that 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 the "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., {{cite book|author=Bellamy, L. J.|year=1958|title=The Infra-Red Spectra of Complex Molecules|location=New York|publisher=Wiley|page=[https://archive.org/details/infraredspectrao0000bell_w1a8/page/392 392]|isbn=978-0-412-13850-8|url=https://archive.org/details/infraredspectrao0000bell_w1a8/page/392}}</ref> however, this is not surprising as such stretching frequencies 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>{{cite journal|author1=Spieseke, H. |author2=Schneider, W. G. |year=1961|journal=Journal of Chemical Physics|volume=35|page=722|doi=10.1063/1.1731992|title=Effect of Electronegativity and Magnetic Anisotropy of Substituents on C13 and H1 Chemical Shifts in CH3X and CH3CH2X Compounds |issue=2 |bibcode = 1961JChPh..35..722S }}</ref> or isomer shifts in [[Mössbauer spectroscopy]]<ref>{{cite journal|author1=Clasen, C. A. |author2=Good, M. L. |year=1970|journal=Inorganic Chemistry|volume=9|pages=817–820 |doi=10.1021/ic50086a025 |title=Interpretation of the Moessbauer spectra of mixed-hexahalo complexes of tin(IV) |issue=4}}</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===
[[File:Periodic variation of Pauling electronegativities.svg|thumb|upright=1.5|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 the most electronegative of the elements (not counting [[noble gas]]es), whereas [[caesium]]<!-- not francium; please don't change unless you supply a citation for published experimental results --> is the least electronegative, at least of those elements for which substantial data is 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 [[Period 4 element|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]] and [[#Sanderson electronegativity equalization|Sanderson electronegativity]] above). The anomalously high electronegativity of [[lead]], in particular when compared to [[thallium]] and [[bismuth]], is an artifact of electronegativity varying with oxidation state: its electronegativity conforms better to trends if it is quoted for the +2 state with a Pauling value of 1.87 instead of the +4 state.

===Variation of electronegativity with oxidation number===
In inorganic chemistry, it is common to consider a single value of 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.<ref>{{Cite journal |last1=Li |first1=Keyan |last2=Xue |first2=Dongfeng |date=2006-10-01 |title=Estimation of Electronegativity Values of Elements in Different Valence States |url=https://pubs.acs.org/doi/10.1021/jp062886k |journal=The Journal of Physical Chemistry A |language=en |volume=110 |issue=39 |pages=11332–11337 |doi=10.1021/jp062886k |pmid=17004743 |bibcode=2006JPCA..11011332L |issn=1089-5639|url-access=subscription }}</ref>

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 were 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.

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

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]] p''K''<sub>a</sub> 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>({{frac|1|4}})&nbsp;= −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, diminishing the partial negative charge of individual oxygen atoms. At the same time, the positive partial charge on the hydrogen increases with a higher oxidation state. This explains the observed increased acidity with an increasing oxidation state in the oxoacids of chlorine.

=== Electronegativity and hybridization scheme ===
The electronegativity of an atom changes depending on the hybridization of the orbital employed in bonding.  Electrons in s orbitals are held more tightly than electrons in p orbitals. Hence, a bond to an atom that employs an sp''<sup>x</sup>'' hybrid orbital for bonding will be more heavily polarized to that atom when the hybrid orbital has more s character.  That is, when electronegativities are compared for different hybridization schemes of a given element, the order {{math|χ(sp<sup>3</sup>) < χ(sp<sup>2</sup>) < χ(sp)}} holds (the trend should apply to [[Isovalent hybridization|non-integer hybridization indices]] as well). 
{| class="wikitable" style="text-align:center"
|-
! Hybridization
! {{math|χ}} (Pauling)<ref>{{Cite book |title=Molecular orbitals and organic chemical reactions |last=Fleming |first=Ian |date=2009 |publisher=Wiley |isbn=978-0-4707-4660-8 |edition=Student |location=Chichester, West Sussex, U.K. |oclc=424555669}}</ref>
|-
| C(sp<sup>3</sup>)
| 2.3
|-
| C(sp<sup>2</sup>)
| 2.6
|-
| C(sp)
| 3.1
|-
|'generic' C
| 2.5
|-
|}

==Group electronegativity==
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 σ- and π-electronegativities, respectively. There are a number of [[linear free-energy relationship]]s that 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]].

==Electropositivity==
'''Electropositivity''' is a measure of an element's ability to donate [[electrons]], and therefore form [[cations|positive]] [[ions]]; thus, it is antipode to electronegativity.

Mainly, this is an attribute of [[metals]], meaning that, in general, the greater the metallic character of an [[chemical element|element]] the greater the electropositivity. Therefore, the [[alkali metals]] are the most electropositive of all. This is because they have a single electron in their outer shell and, as this is relatively far from the nucleus of the atom, it is easily lost; in other words, these metals have low [[ionization energy|ionization energies]].<ref>"[https://archive.today/20091101004334/http://au.encarta.msn.com/encyclopedia_781538810/Electropositivity.html Electropositivity]," [[Microsoft Encarta]] Online Encyclopedia 2009. (Archived 2009-10-31).</ref>

While electronegativity increases along [[Period (periodic table)|periods]] in the [[periodic table]], and decreases down [[Periodic table group|groups]], electropositivity ''decreases'' along periods (from left to right) and ''increases'' down groups. This means that elements in the upper right of the periodic table of elements (oxygen, sulfur, chlorine, etc.) will have the greatest electronegativity, and those in the lower-left (rubidium, caesium, and francium) the greatest electropositivity.

==See also==
* [[Chemical polarity]]
* [[Electron affinity]]
* [[Electronegativities of the elements (data page)]]
* [[Ionization energy]]
* [[Metallic bonding]]
* [[Miedema's model]]
* [[Orbital hybridization]]
* [[Oxidation state]]
* [[Periodic table]]

==References==
{{Reflist}}

==Bibliography==
* {{cite book|last=Jolly|first= William L. |year=1991|title=Modern Inorganic Chemistry|edition=2nd| location=New York|publisher=[[McGraw-Hill]]|isbn= 978-0-07-112651-9|pages=71–76}}
* {{Cite book| last=Mullay|first=J. |title=Electronegativity |year=1987|chapter=Estimation of atomic and group electronegativities|volume=66|pages=1–25| doi=10.1007/BFb0029834| series=Structure and Bonding| isbn=978-3-540-17740-1}}

==External links==
* {{Commons category-inline}}
* [http://www.webelements.com/ WebElements], lists values of electronegativities by a number of different methods of calculation
* [https://web.archive.org/web/20120516075254/http://sciencehack.com/videos/view/6952235798166539784 Video explaining electronegativity]
* [http://electronegativitychart.com Electronegativity Chart], a summary listing of the electronegativity of each element along with an interactive periodic table

{{Navbox periodic table|state=expanded}}
{{Linus Pauling}}
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[[Category:Chemical properties]]
[[Category:Chemical bonding]]
[[Category:Dimensionless numbers of chemistry]]