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| {{redirect|Electronegative|the Nightfall EP|Electronegative (EP)}}
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| [[Image:Electrostatic Potential.jpg|thumb|Electrostatic Potential Map|300px|right|This electrostatic potential map shows how the oxygen atom has a more negative charge (red) than the positive (blue) hydrogen atoms of a water molecule .]]
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| '''Electronegativity''', symbol '''[[Chi (letter)|<span class="nounderlines">χ</span>]]''', is a [[chemical property]] that describes the tendency of an [[atom]] or a [[functional group]] to attract [[electron]]s (or [[electron density]]) towards itself.<ref name="definition">{{GoldBookRef|file=E01990|title=Electronegativity}}</ref> An atom's electronegativity is affected by both its [[atomic number]] and the distance that its [[valence electrons]] reside from the charged nucleus. The higher the associated electronegativity number, the more an element or compound attracts electrons towards it. First proposed by [[Linus Pauling]] in 1932 as a development of [[valence bond theory]],<ref name="paulingJACS">{{cite journal |author= Pauling, L.
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| |authorlink=Linus Pauling
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| |year= 1932
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| |journal= [[Journal of the American Chemical Society]]
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| |volume= 54
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| |issue= 9
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| |pages= 3570–3582
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| |title= The Nature of the Chemical Bond. IV. The Energy of Single Bonds and the Relative Electronegativity of Atoms
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| |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]].
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| 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''', on a relative scale running from around 0.7 to 3.98 ([[hydrogen]] = 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'''.
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| Electronegativity, as it is usually calculated, is not strictly a property of an atom, 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|publisher=Cornell University Press|pages=88–107|isbn=0-8014-0333-2}}</ref> Properties of a free atom include [[ionization energy]] and [[electron affinity]]. It is to be expected that the electronegativity of an element will vary with its chemical environment,<ref>{{cite book|author=Greenwood, N. N.; Earnshaw, A. |year=1984|title=Chemistry of the Elements|publisher=Pergamon|isbn=0-08-022057-6|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.
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| 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 negative electrons) and the number/location of other electrons present 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 orbitals will act to [[Shielding effect|shield]] the valence electrons from the positively charged nucleus).
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| The opposite of electronegativity is '''[[electropositivity]]''': a measure of an element's ability to donate electrons.
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| ==Electronegativities of the elements==
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| {{Periodic table (electronegativity by Pauling scale)}}
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| ==Methods of calculation==
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| ===Pauling electronegativity===
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| [[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–B) is stronger than would be expected by taking the average of the strengths of the A–A and 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 [[canonical form]]s to the bonding.
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| The difference in electronegativity between atoms A and B is given by:
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| ::<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>
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| 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>–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–Br, 3.79 eV; H–H, 4.52 eV; Br–Br 2.00 eV)
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| 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
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| |author= Allred, A. L.
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| |year= 1961
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| |journal= Journal of Inorganic and Nuclear Chemistry
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| |volume= 17
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| |issue= 3–4
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| |pages= 215–221
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| |title= Electronegativity values from thermochemical data
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| |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 is in fact overdetermined, and the signs are unique once a reference point is fixed (usually, for H or F).
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| To calculate 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. 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.
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| The essential point of Pauling electronegativity is that there is an underlying, quite accurate, semi-empirical formula for dissociation energies, namely:
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| ::<math>E_{\rm d}({\rm AB}) =[E_{\rm d}({\rm AA}) + E_{\rm d}({\rm BB})]/2+(\chi_{\rm A} - \chi_{\rm B})^2 eV</math>
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| or sometimes, a more accurate fit
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| ::<math>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 eV</math>
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| This is an approximate equation, but holds with good accuracy. Pauling obtained it 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 approximately, by quantum mechanical calculations, the [[geometric mean]] of the two energies of covalent bonds of the same molecules, and there is an additional energy that comes from ionic factors, i.e. polar character of the bond.
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| The geometric mean is approximately equal to the [[arithmetic mean]] - which is applied in the first formula above - when the energies are of the 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 a 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 this semi-empirical formula for bond energy that underlies Pauling electronegativity concept.
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| The formulas are approximate, but this rough approximation is in fact relatively good and gives the right intuition, with the notion of polarity of the bond and some theoretical grounding in quantum mechanics. The electronegativities are then determined to best fit the data.
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| In more complex compounds, there is 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 energy of formation of a molecule containing only single bonds then can be approximated from an electronegativity table, and depends on the constituents and sum of squares of differences of electronegativities of all pairs of bonded atoms. Such a formula for estimating energy typically has relative error of order of 10%, but can be used to get a rough qualitative idea and understanding of a molecule.
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| ===Mulliken electronegativity===
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| [[Image:Pauling and Mullikan electronegativities.png|thumb|right|300px|The correlation between Mulliken electronegativities (''x''-axis, in kJ/mol) and Pauling electronegativities (''y''-axis).]]
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| [[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
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| |author = Mulliken, R. S.
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| |year =1934
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| |journal = [[Journal of Chemical Physics]]
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| |volume = 2
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| |title = A New Electroaffinity Scale; Together with Data on Valence States and on Valence Ionization Potentials and Electron Affinities
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| |doi = 10.1063/1.1749394
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| |pages = 782–793
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| |issue = 11|bibcode = 1934JChPh...2..782M }}</ref><ref>{{cite journal
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| |author = Mulliken, R. S.
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| |year =1935
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| |title = Electronic Structures of Molecules XI. Electroaffinity, Molecular Orbitals and Dipole Moments
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| |journal = ''[[Journal of Chemical Physics|J. Chem. Phys.]]''
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| |volume = 3
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| |doi = 10.1063/1.1749731
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| |pages = 573–585
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| |issue = 9|bibcode = 1935JChPh...3..573M }}</ref> 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.
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| |title = Absolute electronegativity and absolute hardness of Lewis acids and bases
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| |year = 1985
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| |journal = ''[[Journal of the American Chemical Society|J. Am. Chem. Soc.]]''
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| |volume = 107
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| |pages = 6801
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| |doi = 10.1021/ja00310a009|issue = 24}}</ref> with the units of [[Joule per mole|kilojoules per mole]] or [[electronvolt]]s.
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| ::<math>\chi = (E_{\rm i} + E_{\rm ea}) /2 \,</math>
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| 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|author = Huheey, J. E.
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| |year = 1978
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| |title = Inorganic Chemistry (2nd Edn.)
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| |publisher = New York: Harper & Row. p. 167}}</ref>
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| ::<math>\chi = 0.187(E_{\rm i} + E_{\rm ea}) + 0.17 \,</math>
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| 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>
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| ::<math>\chi = (1.97\times 10^{-3})(E_{\rm i} + E_{\rm ea}) + 0.19.</math>
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| The Mulliken electronegativity can only be calculated for an element for which the electron affinity is known, fifty-seven elements as of 2006.
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| The Mulliken electronegativity of an atom is sometimes said to be the negative of the chemical potential. 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.,
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| ::<math>\mu(\rm Mulliken) = -\chi(\rm Mulliken) = -(E_{\rm i} + E_{\rm ea}) /2 \,</math>
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| ===Allred–Rochow electronegativity===
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| [[Image:Pauling and Allred-Rochow electronegativities.png|thumb|right|300px|The correlation between Allred–Rochow electronegativities (''x''-axis, in Å<sup>–2</sup>) and Pauling electronegativities (''y''-axis).]]
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| A. Louis Allred and [[Eugene G. Rochow]] considered<ref>{{cite journal|author=Allred, A. L.; Rochow, E. G.|year=1958|journal=Journal of Inorganic and Nuclear Chemistry|volume=5|page=264|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>Housecroft C.E. and Sharpe A.G. ''Inorganic Chemistry'' (2nd ed., Pearson Prentice-Hall 2005) p.38</ref>
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| ::<math>\chi = 3590{{Z_{\rm eff}}\over{r^2_{\rm cov}}} + 0.744</math>
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| ===Sanderson electronegativity equalization===
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| [[Image:Pauling and Sanderson electronegativities.png|thumb|right|300px|The correlation between Sanderson electronegativities (''x''-axis, arbitrary units) and Pauling electronegativities (''y''-axis).]]
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| 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|page=2259|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=0-12-618080-6}}</ref> Sanderson's model has also been used to calculate molecular geometry, ''s''-electrons energy, [[NMR]] spin-spin constants and other parameters for organic compounds.<ref>{{cite journal|last=Zefirov|first=N. S.|coauthors=M. A. Kirpichenok, F. F. Izmailov, and M. I. Trofimov|journal=[[Doklady Akademii Nauk SSSR]]|year=1987|volume=296|page=883}}</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|last2=Smolenskii|first2=E. A.|issue=9}}</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>
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| {{cite book |title=Reviews in computational chemistry |author=SW Rick &SJ Stuart |chapter=Electronegativity equalization models |editor=Kenny B. Lipkowitz, Donald B. Boyd
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| |url=http://books.google.com/?id=IqWXSLz6QE8C&pg=PA106 |page=106 |isbn=0-471-21576-7 |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 |author= Robert G. Parr, Weitao Yang |url=http://books.google.com/?id=mGOpScSIwU4C&pg=PA91 |page=91 |isbn=0-19-509276-7 |year=1994 |publisher=Oxford University Press}}</ref>
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| ===Allen electronegativity===
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| [[Image:Pauling and Allen electronegativities.png|thumb|right|300px|The correlation between Allen electronegativities (''x''-axis, in kJ/mol) and Pauling electronegativities (''y''-axis).]]
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| 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|issue=25}}</ref>
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| :<math>\chi = {n_{\rm s}\varepsilon_{\rm s} + n_{\rm p}\varepsilon_{\rm p} \over n_{\rm s} + n_{\rm p}}</math>
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| 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. It is usual to apply a scaling factor, 1.75×10<sup>−3</sup> for energies expressed in kilojoules per mole or 0.169 for energies measured in electronvolts, to give values that are numerically similar to Pauling electronegativities.
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| 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.
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| In this scale [[neon]] has the highest electronegativity of all elements, followed by [[fluorine]], [[helium]], and [[oxygen]].
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| {{periodic table (electronegativity by Allen scale)}}
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| ==Correlation of electronegativity with other properties==
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| [[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–</sup> anions, as measured by <sup>119</sup>Sn [[Mössbauer spectroscopy]], against the sum of the Pauling electronegativities of the halide substituents (''x''-axis).]]
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| 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 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.
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| 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=392|isbn=0-412-13850-6}}</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|author=Spieseke, H.; 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|author=Clasen, C. A.; Good, M. L. |year=1970|journal=Inorganic Chemistry|volume=9|page=817|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"/>
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| ==Trends in electronegativity==
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| ===Periodic trends===
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| [[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]]
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| 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"/>
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| 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]], [[#Sanderson electronegativity|Sanderson electronegativity]] above). The anomalously high electronegativity of [[lead]], in particular when compared to [[thallium]] and [[bismuth]], appears to be an artifact of data selection (and data availability)—methods of calculation other than the Pauling method show the normal periodic trends for these elements.
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| ===Variation of electronegativity with oxidation number===
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| 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.
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| 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.
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| {| class="wikitable" style="text-align:center"
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| |-
| |
| ! Acid
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| ! Formula
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| ! Chlorine<br/>oxidation<br/>state
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| ! p''K''<sub>a</sub>
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| |-
| |
| | [[Hypochlorous acid]]
| |
| | HClO
| |
| | +1
| |
| | +7.5
| |
| |-
| |
| | [[Chlorous acid]]
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| | HClO<sub>2</sub>
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| | +3
| |
| | +2.0
| |
| |-
| |
| | [[Chloric acid]]
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| | HClO<sub>3</sub>
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| | +5
| |
| | –1.0
| |
| |-
| |
| | [[Perchloric acid]]
| |
| | HClO<sub>4</sub>
| |
| | +7
| |
| | –10
| |
| |-
| |
| |}
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| 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]].
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| 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) = –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.
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| ==Group electronegativity==
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| 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]].
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| ==Electropositivity==
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| '''Electropositivity''' is a measure of an element's ability to donate [[electrons]], and therefore form positive [[ions]]; thus, it is opposed 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 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>"[http://www.webcitation.org/5kwpIKRbL Electropositivity]," [[Microsoft Encarta]] Online Encyclopedia 2009. (Archived 2009-10-31).</ref>
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| 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.
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| [[Electropositive shark repellent]] utilizes electropositive metals as [[shark repellent]]s, since they generate measurable voltages in a seawater electrolyte relative to a shark.
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| ==See also==
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| *[[Electronegativities of the elements (data page)]]
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| *[[Chemical polarity]]
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| ==References==
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| {{reflist|2}}
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| ==Bibliography==
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| *{{cite book|last=Jolly|first= William L. |year=1991|title=Modern Inorganic Chemistry|edition=2nd| location=New York|publisher=[[McGraw-Hill]]|isbn= 0-07-112651-1|pages=71–76}}
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| *{{cite journal| last=Mullay|first=J. |year=1987|title=Estimation of atomic and group electronegativities|journal=Structure and Bonding|volume=66|pages=1–25| doi=10.1007/BFb0029834| series=Structure and Bonding| isbn=3-540-17740-X}}
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| ==External links==
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| {{commons category|Electronegativity}}
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| *[http://www.webelements.com/ WebElements], lists values of electronegativities by a number of different methods of calculation
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| *[http://sciencehack.com/videos/view/6952235798166539784 Video explaining electronegativity]
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| {{PeriodicTablesFooter}}
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| [[Category:Chemical properties]]
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| [[Category:Chemical bonding]]
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| {{Link FA|sl}}
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