Electronegativity

"Electronegativity" is antipodally distinguished from "Electropositivity," which describes an element's ability to donate electrons.

Electronegativity, symbol χ, is a chemical property that describes the power of an elements.

The most commonly used method of calculation is that originally proposed by Pauling. This gives a dimensionless quantity, commonly referred to as the Pauling scale, on a relative scale running from 0.7 to 4.0 (hydrogen = 2.2). 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.

Electronegativity, as it is usually calculated, is not strictly an atomic property, but rather a property of an atom in a molecule:^[3] 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,^[4] but it is usually considered to be a transferable property, that is to say that similar values will be valid in a variety of situations.

Electronegativities of the elements

→ Atomic radius decreases → Ionization energy increases → Electronegativity increases →
Group (vertical)	1	2	3	4	5	6	7	8	9	10	11	12	13	14	15	16	17	18
Period (horizontal)
1	H 2.20																	He 3.89
2	Li 0.98	Be 1.57											B 2.04	C 2.55	N 3.04	O 3.44	F 3.98	Ne 3.67
3	Na 0.93	Mg 1.31											Al 1.61	Si 1.90	P 2.19	S 2.58	Cl 3.16	Ar 3.3
4	K 0.82	Ca 1.00	Sc 1.36	Ti 1.54	V 1.63	Cr 1.66	Mn 1.55	Fe 1.83	Co 1.88	Ni 1.91	Cu 1.90	Zn 1.65	Ga 1.81	Ge 2.01	As 2.18	Se 2.55	Br 2.96	Kr 3.00
5	Rb 0.82	Sr 0.95	Y 1.22	Zr 1.33	Nb 1.6	Mo 2.16	Tc 1.9	Ru 2.2	Rh 2.28	Pd 2.20	Ag 1.93	Cd 1.69	In 1.78	Sn 1.96	Sb 2.05	Te 2.1	I 2.66	Xe 2.67
6	Cs 0.79	Ba 0.89	*	Hf 1.3	Ta 1.5	W 2.36	Re 1.9	Os 2.2	Ir 2.20	Pt 2.28	Au 2.54	Hg 2.00	Tl 1.62	Pb 2.33	Bi 2.02	Po 2.0	At 2.2	Rn 2.2
7	Fr 0.7	Ra 0.9	**	Rf	Db	Sg	Bh	Hs	Mt	Ds	Rg	Uub	Uut	Uuq	Uup	Uuh	Uus	Uuo

Lanthanides	*	La 1.1	Ce 1.12	Pr 1.13	Nd 1.14	Pm 1.13	Sm 1.17	Eu 1.2	Gd 1.2	Tb 1.1	Dy 1.22	Ho 1.23	Er 1.24	Tm 1.25	Yb 1.1	Lu 1.27
Actinides	**	Ac 1.1	Th 1.3	Pa 1.5	U 1.38	Np 1.36	Pu 1.28	Am 1.13	Cm 1.28	Bk 1.3	Cf 1.3	Es 1.3	Fm 1.3	Md 1.3	No 1.3	Lr 1.291

Periodic table of electronegativity using the Pauling scale See also Periodic table

Methods of calculation

Pauling electronegativity

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 forms to the bonding.

The difference in electronegativity between atoms A and B is given by:

$\chi_{\rm A} - \chi_{\rm B} = ({\rm eV})^{-1/2} \sqrt{E_{\rm d}({\rm AB}) - [E_{\rm d}({\rm AA}) + E_{\rm d}({\rm BB})]/2}$

where the bromine is 0.73 (dissociation energies: H–Br, 3.79 eV; H–H, 4.52 eV; Br–Br 2.00 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^[2] at 2.1, later revised^[5] 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⁺ and Br⁻ ions, so it may be assumed that bromine is more electronegative than hydrogen.

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. Allred updated Pauling's original values in 1961 to take account of the greater availability of thermodynamic data,^[5] and it is these "revised Pauling" values of the electronegativity which are most usually used.

Mulliken electronegativity

kilojoules per mole or electronvolts.

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,^[8]

$\chi = 0.187(E_{\rm i} + E_{\rm ea}) + 0.17 \,$

and for energies in kilojoules per mole,^[9]

$\chi = (1.97\times 10^{-3})(E_{\rm i} + E_{\rm ea}) + 0.19$

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

Allred and Rochow considered^[10] 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 covalent radius, r_cov. When r_cov is expressed in ångströms,

$\chi = 0.359{{Z\star}\over{r^2_{\rm cov}}} + 0.744$ .

Sanderson electronegativity

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.^[11] With a knowledge of bond lengths, Sanderson electronegativities allow the estimation of bond energies in a wide range of compounds.^[12] Also Sanderson electronegativities were used for different investigations in organic chemistry.^[13]^[14]

Allen electronegativity

Perhaps the simplest definition of electronegativity is that of Allen, who has proposed that it is related to the average energy of the valence electrons in a free atom,^[15]

$\chi = {n_{\rm s}\varepsilon_{\rm s} + n_{\rm p}\varepsilon_{\rm p} \over n_{\rm s} + n_{\rm p}}$

where ε_s,p are the one-electron energies of s- and p-electrons in the free atom and n_s,p are the number of s- and p-electrons in the valence shell. It is usual to apply a scaling factor, 1.75×10⁻³ 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 francium, which has an Allen electronegativity of 0.67.^[16] 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

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,^[3] although this has fallen somewhat into disuse.

Several correlations have been shown between Mössbauer spectroscopy^[19] (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".^[1]^[3]

Trends in electronegativity

Periodic trends

In general, electronegativity increases on passing from left to right along a period, and decreases on descending a group. Hence, caesium is the least electronegative, at least of those elements for which substantial data are available.^[16]

There are some exceptions to this general rule. 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.

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.^[5] 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.

Acid	Formula	Chlorine oxidation state	pK_a
Hypochlorous acid	HClO	+1	+7.5
Chlorous acid	HClO₂	+3	+2.0
Chloric acid	HClO₃	+5	−1.0
Perchloric acid	HClO₄	+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 basic oxide.

The effect can also be clearly seen in the 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 article: 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 organophosphorus chemistry.

Notes

^ ^a ^b "Electronegativity.", IUPAC Compendium of Chemical Terminology
^ ^a ^b ^c J. Am. Chem. Soc. 54 (9): 3570–3582. doi:10.1021/ja01348a011.
^ ^a ^b ^c Pauling, Linus (1960). Nature of the Chemical Bond (3rd Edn.). Ithaca, NY: Cornell University Press. pp. 88–107.
^ Greenwood, N. N.; Earnshaw, A. (1984). Chemistry of the Elements. Oxford: Pergamon. ISBN 0-08-022057-6. p. 30.
^ ^a ^b ^c Allred, A. L. (1961). "Electronegativity values from thermochemical data". J. Inorg. Nucl. Chem. 17 (3–4): 215–221. doi:10.1016/0022-1902(61)80142-5.
^ Mulliken, R. S. (1934). J. Chem. Phys. 3:573.
^ Pearson, R. G. (1985). J. Am. Chem. Soc. 107:6801.
^ Huheey, J. E. (1978). Inorganic Chemistry (2nd Edn.). New York: Harper & Row. p. 167.
^ This second relation has been recalculated using the best values of the first ionization energies and electron affinities available in 2006.
^ Allred, A. L.; Rochow, E. G. (1958). J. Inorg. Nucl. Chem. 5:264.
^ Sanderson, R. T. (1983). "Electronegativity and bond energy." J. Am. Chem. Soc. 105:2259.
^ Sanderson, R. T. (1983). Polar Covalence. New York: Academic Press.
^ N. S. Zefirov, M. A. Kirpichenok, F. F. Izmailov, and M. I.Trofimov, Dokl. Akad. Nauk SSSR, 1987, 296: 883 [Dokl.Chem., 1987 (Engl. Transl.)].
^ M.I.Trofimov, E.A.Smolenskii, Russian Chemical Bulletin, 2005, 54(9): 2235.(http://dx.doi.org/10.1007/s11172-006-0105-6).
^ Allen, L. C. (1989). J. Am. Chem. Soc. 111:9003.
^ ^a ^b 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.
^ See, e.g., Bellamy, L. J. (1958). The Infra-Red Spectra of Complex Molecules (2nd Edn.). New York: Wiley. p. 392.
^ Spieseke, H.; Schneider, W. G. (1961). J. Chem. Phys. 35:722.
^ Clasen, C. A.; Good, M. L. (1970). Inorg. Chem. 9:817.

References

Jolly, William L. (1991). Modern Inorganic Chemistry (2nd Edn.). New York: McGraw-Hill. ISBN 0-07-112651-1. pp. 71–76.
Mullay, J. (1987). "Estimation of atomic and group electronegativities." Struct. Bond. 66:1–25.

Chemical properties

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