How Do You Know the Electronegativity of an Element

Trend of an cantlet to attract a shared pair of electrons

"Electronegative" redirects hither. For the Nightfall EP, see Electronegative (EP).

Electrostatic potential map of a water molecule, where the oxygen atom has a more than negative charge (red) than the positive (blue) hydrogen atoms

Electronegativity, symbolized as χ , is the tendency for an atom of a given chemical element to attract shared electrons (or electron density) when forming a chemical bond.^[1] An atom's electronegativity is afflicted past 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 mode to quantitatively estimate the bond free energy, and the sign and magnitude of a bail's chemical polarity, which characterizes a bail along the continuous scale from covalent to ionic bonding. The loosely defined term electropositivity is the opposite of electronegativity: information technology characterizes an element's tendency to donate valence electrons.

On the nigh basic level, electronegativity is adamant past factors like the 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 atomic shells (the more electrons an cantlet has, the farther from the nucleus the valence electrons will exist, and as a result, the less positive charge they volition experience—both because of their increased distance from the nucleus and because the other electrons in the lower energy core orbitals volition act to shield the valence electrons from the positively charged nucleus).

The term "electronegativity" was introduced by Jöns Jacob Berzelius in 1811,^[2] though the concept was known earlier that and was studied by many chemists including Avogadro.^[2] In spite of its long history, an accurate scale of electronegativity was non developed until 1932, when Linus Pauling proposed an electronegativity scale which depends on bail energies, equally a development of valence bond theory.^[3] Information technology 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 take been proposed, and although there may be small-scale differences in the numerical values of the electronegativity, all methods testify the same periodic trends betwixt elements.^[4]

The about ordinarily used method of calculation is that originally proposed past Linus Pauling. This gives a dimensionless quantity, unremarkably referred to as the Pauling scale (χ _r), on a relative scale running from 0.79 to 3.98 (hydrogen = 2.xx). When other methods of calculation are used, information technology is conventional (although not obligatory) to quote the results on a scale that covers the same range of numerical values: this is known equally an electronegativity in Pauling units.

Equally it is unremarkably calculated, electronegativity is not a property of an atom alone, but rather a holding of an cantlet in a molecule.^[5] Still, the electronegativity of an cantlet is strongly correlated with the showtime ionization energy, and negatively correlated with the electron affinity. It is to be expected that the electronegativity of an element volition vary with its chemic environs,^[six] just it is ordinarily considered to be a transferable belongings, that is to say that similar values will be valid in a variety of situations.

Caesium is the least electronegative element (0.79); fluorine is the most (3.98).

Methods of calculation [edit]

Pauling electronegativity [edit]

Pauling offset proposed^[3] the concept of electronegativity in 1932 to explain why the covalent bond between 2 dissimilar 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 bail is due to the contribution of ionic canonical forms to the bonding.

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

$${\displaystyle |\chi _{\rm {A}}-\chi _{\rm {B}}|=({\rm {eV}})^{-1/ii}{\sqrt {E_{\rm {d}}({\rm {AB}})-{\frac {E_{\rm {d}}({\rm {AA}})+E_{\rm {d}}({\rm {BB}})}{two}}}}}$$

where the dissociation energies, E _d, of the A–B, A–A and B–B bonds are expressed in electronvolts, the factor (eV)^{− 1⁄ii} being included to ensure a dimensionless result. Hence, the difference in Pauling electronegativity between hydrogen and bromine is 0.73 (dissociation energies: H–Br, three.79 eV; H–H, four.52 eV; Br–Br two.00 eV)

Every bit merely differences in electronegativity are defined, it is necessary to cull an arbitrary reference betoken in society to construct a scale. Hydrogen was called as the reference, equally it forms covalent bonds with a large variety of elements: its electronegativity was stock-still get-go^[iii] at ii.1, later revised^[7] to 2.20. It is besides necessary to decide which of the two elements is the more electronegative (equivalent to choosing 1 of the two possible signs for the square root). This is usually done using "chemical intuition": in the above case, hydrogen bromide dissolves in water to form H⁺ and Br⁻ ions, so it may exist assumed that bromine is more electronegative than hydrogen. However, in principle, since the same electronegativities should be obtained for whatever ii bonding compounds, the data are in fact overdetermined, and the signs are unique once a reference point is stock-still (normally, for H or F).

To calculate Pauling electronegativity for an chemical element, it is necessary to have data on the dissociation energies of at to the lowest degree two types of covalent bonds formed past that chemical element. A. 50. Allred updated Pauling'southward original values in 1961 to take account of the greater availability of thermodynamic data,^[7] and it is these "revised Pauling" values of the electronegativity that are virtually oft used.

The essential point of Pauling electronegativity is that there is an underlying, quite accurate, semi-empirical formula for dissociation energies, namely:

$${\displaystyle E_{\rm {d}}({\rm {AB}})={\frac {E_{\rm {d}}({\rm {AA}})+E_{\rm {d}}({\rm {BB}})}{2}}+(\chi _{\rm {A}}-\chi _{\rm {B}})^{ii}{\rm {eV}}}$$

or sometimes, a more than accurate fit

$${\displaystyle E_{\rm {d}}({\rm {AB}})={\sqrt {E_{\rm {d}}({\rm {AA}})E_{\rm {d}}({\rm {BB}})}}+1.three(\chi _{\rm {A}}-\chi _{\rm {B}})^{ii}{\rm {eV}}}$$

This is an approximate equation merely holds with skillful accurateness. Pauling obtained information technology past noting that a bond tin be approximately represented as a quantum mechanical superposition of a covalent bond and 2 ionic bond-states. The covalent energy of a bail is approximate, by quantum mechanical calculations, the geometric mean of the two energies of covalent bonds of the same molecules, and in that location 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 practical in the first formula above—when the energies are of a similar value, e.yard., except for the highly electropositive elements, where there is a larger difference of ii dissociation energies; the geometric mean is more authentic and nearly always gives positive backlog 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 bail free energy that underlies the concept of Pauling electronegativity.

The formulas are approximate, merely this rough approximation is in fact relatively adept 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 than complex compounds, there is an boosted fault since electronegativity depends on the molecular environment of an atom. Besides, the energy estimate tin be only used for unmarried, not for multiple bonds. The energy of the formation of a molecule containing only unmarried bonds can subsequently exist approximated from an electronegativity tabular array 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 a relative fault of an gild of 10% only can be used to go a rough qualitative idea and understanding of a molecule.

1. ^ The electronegativity of francium was called by Pauling as 0.seven, close to that of caesium (also assessed 0.7 at that indicate). The base value of hydrogen was later increased by 0.x and caesium's electronegativity was afterwards refined to 0.79; nevertheless, no refinements have been fabricated for francium as no experiment has been conducted. All the same, francium is expected and, to a small extent, observed to be more electronegative than caesium. See francium for details.
2. ^ See Brown, Geoffrey (2012). The Inaccessible Globe: An integrated view to its structure and composition. Springer Scientific discipline & Business concern Media. p. 88. ISBN9789401115162.

Mulliken electronegativity [edit]

The correlation between Mulliken electronegativities (10-axis, in kJ/mol) and Pauling electronegativities (y-axis).

Robert S. Mulliken proposed that the arithmetic hateful of the first ionization energy (Due east_i) and the electron affinity (E_ea) should be a measure of the trend of an atom to attract electrons.^{[eight] [nine]} As this definition is not dependent on an arbitrary relative scale, it has also been termed accented electronegativity,^[ten] with the units of kilojoules per mole or electronvolts.

$${\displaystyle \chi ={\frac {E_{\rm {i}}+E_{\rm {ea}}}{2}}}$$

However, it is more usual to use a linear transformation to transform these absolute values into values that resemble the more than familiar Pauling values. For ionization energies and electron affinities in electronvolts,^[xi]

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

and for energies in kilojoules per mole,^[12]

$${\displaystyle \chi =(i.97\times 10^{-iii})(E_{\rm {i}}+E_{\rm {ea}})+0.19.}$$

The Mulliken electronegativity can just exist calculated for an element for which the electron analogousness is known, fifty-vii elements as of 2006. The Mulliken electronegativity of an atom is sometimes said to be the negative of the chemical potential.^[13] By inserting the energetic definitions of the ionization potential and electron affinity into the Mulliken electronegativity, information technology is possible to show that the Mulliken chemical potential is a finite divergence approximation of the electronic energy with respect to the number of electrons., i.due east.,

$${\displaystyle \mu ({\rm {Mulliken)=-\chi ({\rm {Mulliken)={}-{\frac {E_{\rm {i}}+E_{\rm {ea}}}{2}}}}}}}$$

Allred–Rochow electronegativity [edit]

The correlation betwixt Allred–Rochow electronegativities (x-axis, in Å⁻²) and Pauling electronegativities (y-centrality).

A. Louis Allred and Eugene Yard. Rochow considered^[14] that electronegativity should be related to the accuse 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 cantlet to attract electrons. The effective nuclear charge, Z _eff, experienced by valence electrons tin can be estimated using Slater'south 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 _cov. When r _cov is expressed in picometres,^[15]

$${\displaystyle \chi =3590{{Z_{\rm {eff}}} \over {r_{\rm {cov}}^{2}}}+0.744}$$

Sanderson electronegativity equalization [edit]

The correlation between Sanderson electronegativities (ten-axis, arbitrary units) and Pauling electronegativities (y-axis).

R.T. Sanderson has also noted the human relationship between Mulliken electronegativity and atomic size, and has proposed a method of calculation based on the reciprocal of the atomic volume.^[16] With a knowledge of bond lengths, Sanderson'south model allows the estimation of bond energies in a wide range of compounds.^[17] Sanderson's model has as well been used to summate molecular geometry, s-electrons energy, NMR spin-spin constants and other parameters for organic compounds.^{[xviii] [19]} 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.^[20] This beliefs is analogous to the equalization of chemic potential in macroscopic thermodynamics.^[21]

Allen electronegativity [edit]

The correlation between Allen electronegativities (x-axis, in kJ/mol) and Pauling electronegativities (y-axis).

Perchance the simplest definition of electronegativity is that of Leland C. Allen, who has proposed that information technology is related to the average energy of the valence electrons in a free cantlet,^{[22] [23] [24]}

$${\displaystyle \chi ={n_{\rm {s}}\varepsilon _{\rm {southward}}+n_{\rm {p}}\varepsilon _{\rm {p}} \over n_{\rm {south}}+n_{\rm {p}}}}$$

where ε _s,p are the ane-electron energies of due south- and p-electrons in the free atom and n _{due south,p} are the number of south- and p-electrons in the valence vanquish. It is usual to apply a scaling gene, ane.75×x^−three 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.

The one-electron energies tin can be adamant directly from spectroscopic data, and so electronegativities calculated by this method are sometimes referred to as spectroscopic electronegativities. The necessary data are available for nearly all elements, and this method allows the interpretation of electronegativities for elements that cannot exist treated by the other methods, e.1000. francium, which has an Allen electronegativity of 0.67.^[25] However, it is not clear what should be considered to be valence electrons for the d- and f-cake elements, which leads to an ambiguity for their electronegativities calculated by the Allen method.

In this scale neon has the highest electronegativity of all elements, followed past fluorine, helium, and oxygen.

v t e Electronegativity using the Allen scale
Group →	ane	two	3	iv	5	6	7	8	9	10	xi	12	13	14	fifteen	xvi	17	18
↓ Menstruation
1	H ii.300																	He 4.160
2	Li 0.912	Be 1.576											B 2.051	C 2.544	Due north 3.066	O 3.610	F 4.193	Ne four.787
three	Na 0.869	Mg i.293											Al i.613	Si 1.916	P 2.253	S two.589	Cl 2.869	Ar 3.242
4	M 0.734	Ca 1.034	Sc 1.xix	Ti 1.38	V 1.53	Cr i.65	Mn one.75	Atomic number 26 1.80	Co i.84	Ni 1.88	Cu 1.85	Zn ane.588	Ga 1.756	Ge 1.994	As 2.211	Se 2.424	Br 2.685	Kr 2.966
v	Rb 0.706	Sr 0.963	Y i.12	Zr 1.32	Nb 1.41	Mo 1.47	Tc 1.51	Ru 1.54	Rh 1.56	Pd one.58	Ag 1.87	Cd one.521	In 1.656	Sn 1.824	Sb one.984	Te two.158	I 2.359	Xe 2.582
6	Cs 0.659	Ba 0.881	Lu 1.09	Hf one.16	Ta one.34	Due west 1.47	Re 1.60	Os one.65	Ir 1.68	Pt i.72	Au 1.92	Hg ane.765	Tl 1.789	Lead 1.854	Bi two.01	Po 2.19	At 2.39	Rn 2.60
7	Fr 0.67	Ra 0.89
See also: Electronegativities of the elements (data page)

Correlation of electronegativity with other backdrop [edit]

The variation of the isomer shift (y-axis, in mm/s) of [SnX_six]²⁻ anions, equally measured by ¹¹⁹Sn Mössbauer spectroscopy, against the sum of the Pauling electronegativities of the halide substituents (ten-axis).

The wide multifariousness of methods of calculation of electronegativities, which all give results that correlate well with i some other, is one indication of the number of chemical backdrop that might exist affected by electronegativity. The almost obvious application of electronegativities is in the discussion of bond polarity, for which the concept was introduced by Pauling. In general, the greater the departure in electronegativity between two atoms the more polar the bond that will be formed betwixt them, with the atom having the college electronegativity beingness at the negative terminate of the dipole. Pauling proposed an equation to chronicle the "ionic character" of a bond to the difference in electronegativity of the two atoms,^[5] although this has fallen somewhat into disuse.

Several correlations have been shown between infrared stretching frequencies of certain bonds and the electronegativities of the atoms involved:^[26] even so, this is non surprising equally such stretching frequencies depend in part on bond strength, which enters into the adding of Pauling electronegativities. More than disarming are the correlations between electronegativity and chemical shifts in NMR spectroscopy^[27] or isomer shifts in Mössbauer spectroscopy^[28] (see figure). Both these measurements depend on the s-electron density at the nucleus, then are a good indication that the unlike measures of electronegativity really are describing "the ability of an atom in a molecule to attract electrons to itself".^{[one] [v]}

Trends in electronegativity [edit]

Periodic trends [edit]

The variation of Pauling electronegativity (y-centrality) every bit one descends the main groups of the periodic table from the 2nd menstruation to the 6th period

In general, electronegativity increases on passing from left to right along a menstruation and decreases on descending a grouping. Hence, fluorine is the most electronegative of the elements (non counting noble gases), whereas caesium is the least electronegative, at least of those elements for which substantial data is available.^[25] This would lead one to believe that caesium fluoride is the compound whose bonding features the virtually ionic character.

There are some exceptions to this general rule. Gallium and germanium accept higher electronegativities than aluminium and silicon, respectively, because of the d-block contraction. Elements of the 4th menses immediately after the offset row of the transition metals have unusually pocket-sized atomic radii because the 3d-electrons are not constructive at shielding the increased nuclear charge, and smaller diminutive size correlates with higher electronegativity (see Allred-Rochow electronegativity, Sanderson electronegativity above). The anomalously high electronegativity of lead, in particular when compared to thallium and bismuth

Variation of electronegativity with oxidation number [edit]

In inorganic chemical science, it is common to consider a unmarried value of electronegativity to be valid for most "normal" situations. While this arroyo has the reward of simplicity, information technology is clear that the electronegativity of an chemical element is non an changeless atomic property and, in particular, increases with the oxidation country 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 were available.^[7] However, for well-nigh elements, at that place are not plenty 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 run across every bit a upshot.

Acrid	Formula	Chlorine oxidation country	pG a
Hypochlorous acid	HClO	+i	+seven.5
Chlorous acid	HClO2	+3	+2.0
Chloric acid	HClO3	+5	–1.0
Perchloric acid	HClO4	+7	–10

The chemic furnishings 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 CrO₃ and Mn₂O₇ are acidic oxides with low melting points, while Cr₂O₃ is amphoteric and Mn₂O₃ is a completely basic oxide.

The effect can also be clearly seen in the dissociation constants of the oxoacids of chlorine. The consequence is much larger than could exist explained by the negative charge being shared among a larger number of oxygen atoms, which would pb to a divergence in pM _a of log₁₀( 1⁄iv ) = –0.half dozen betwixt hypochlorous acid and perchloric acid. As the oxidation state of the primal chlorine cantlet increases, more electron density is fatigued from the oxygen atoms onto the chlorine, diminishing the partial negative charge of private oxygen atoms. At the same fourth dimension, the positive fractional accuse on the hydrogen increases with a college oxidation state. This explains the observed increased acidity with an increasing oxidation state in the oxoacids of chlorine.

Electronegativity and hybridization scheme [edit]

The electronegativity of an cantlet 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 ¹⁰ hybrid orbital for bonding will be more than 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 society χ(sp³) < χ(sp^two) < χ(sp) holds (the trend should apply to non-integer hybridization indices also). While this holds truthful in principle for any main-group element, values for the hybridization-specific electronegativity are most oft cited for carbon. In organic chemistry, these electronegativities are frequently invoked to predict or rationalize bond polarities in organic compounds containing double and triple bonds to carbon.

Hybridization	χ (Pauling)[29]
C(spthree)	2.3
C(sp2)	two.6
C(sp)	three.ane
'generic' C	ii.5

Group electronegativity [edit]

In organic chemistry, electronegativity is associated more than with different functional groups than with private atoms. The terms grouping electronegativity and substituent electronegativity are used synonymously. Nevertheless, it is common to distinguish between the inductive issue and the resonance outcome, which might be described as σ- and π-electronegativities, respectively. There are a number of linear free-energy relationships that have been used to quantify these effects, of which the Hammett equation is the best known. Kabachnik parameters are grouping electronegativities for use in organophosphorus chemistry.

Electropositivity [edit]

Electropositivity is a measure of an element's ability to donate electrons, and therefore form positive ions; thus, information technology is antipode to electronegativity.

Mainly, this is an attribute of metals, meaning that, in general, the greater the metallic character of an element the greater the electropositivity. Therefore, the brine metals are the nearly electropositive of all. This is because they have a single electron in their outer trounce and, as this is relatively far from the nucleus of the atom, information technology is easily lost; in other words, these metals have depression ionization energies.^[30]

While electronegativity increases along periods in the periodic table, and decreases down groups, electropositivity decreases along periods (from left to correct) and increases downwards groups. This means that elements in the upper correct of the periodic table of elements (oxygen, sulfur, chlorine, etc.) will accept the greatest electronegativity, and those in the lower-left (rubidium, caesium, and francium) the greatest electropositivity.

See besides [edit]

• Chemical polarity
• Electron affinity
• Electronegativities of the elements (data page)
• Ionization energy
• Metallic bonding
• Miedema's model
• Orbital hybridization
• Oxidation land
• Periodic table

References [edit]

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2. ^ ^{a b} Jensen, W.B. (1996). "Electronegativity from Avogadro to Pauling: Part one: Origins of the Electronegativity Concept". Journal of Chemical Instruction. 73 (i): 11–xx. Bibcode:1996JChEd..73...11J. doi:ten.1021/ed073p11.
3. ^ ^{a b c} Pauling, L. (1932). "The Nature of the Chemic Bond. Iv. The Free energy of Unmarried Bonds and the Relative Electronegativity of Atoms". Journal of the American Chemical Order. 54 (nine): 3570–3582. doi:10.1021/ja01348a011.
4. ^ Sproul, Gordon D. (2020-05-26). "Evaluation of Electronegativity Scales". ACS Omega. five (twenty): 11585–11594. doi:10.1021/acsomega.0c00831.
5. ^ ^{a b c} Pauling, Linus (1960). Nature of the Chemic Bond . Cornell Academy Press. pp. 88–107. ISBN978-0-8014-0333-0.
6. ^ Greenwood, N. N.; Earnshaw, A. (1984). Chemistry of the Elements. Pergamon. p. 30. ISBN978-0-08-022057-4.
7. ^ ^{a b c} Allred, A. L. (1961). "Electronegativity values from thermochemical data". Journal of Inorganic and Nuclear Chemistry. 17 (3–4): 215–221. doi:10.1016/0022-1902(61)80142-5.
8. ^ Mulliken, R. S. (1934). "A New Electroaffinity Scale; Together with Data on Valence States and on Valence Ionization Potentials and Electron Affinities". Journal of Chemical Physics. 2 (xi): 782–793. Bibcode:1934JChPh...2..782M. doi:10.1063/ane.1749394.
9. ^ Mulliken, R. South. (1935). "Electronic Structures of Molecules XI. Electroaffinity, Molecular Orbitals and Dipole Moments". J. Chem. Phys. 3 (9): 573–585. Bibcode:1935JChPh...three..573M. doi:10.1063/1.1749731.
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11. ^ Huheey, J.E.; Keiter, E.A.; Keiter, R.L. (December 1, 2008) [1978]. "17". In Kauffman, G.B. (ed.). Inorganic Chemistry: Principles of Construction and Reactivity (digitalized). Inorganic Chemistry: Principles of Construction and Reactivity (tertiary ed.). New York (published 1900). p. 167. doi:ten.1021/ed050pA379.1. ISBN9780060429874. OCLC 770736023. inorganicchemist00huhe_0. Archived from the original on January 1, 2014. Retrieved Dec xv, 2020 – via Oxford University Press.
12. ^ This second relation has been recalculated using the best values of the first ionization energies and electron affinities available in 2006.
13. ^ Franco-Pérez, Marco; Gázquez, José L. (31 October 2019). "Electronegativities of Pauling and Mulliken in Density Functional Theory". Periodical of Physical Chemistry A. 123 (46): 10065–10071. Bibcode:2019JPCA..12310065F. doi:10.1021/acs.jpca.9b07468. PMID 31670960. S2CID 207814569.
14. ^ Allred, A. L.; Rochow, E. G. (1958). "A scale of electronegativity based on electrostatic force". Journal of Inorganic and Nuclear Chemistry. five (iv): 264–268. doi:10.1016/0022-1902(58)80003-2.
15. ^ Housecroft, C.E.; Sharpe, A.G. (November 1, 1993). Inorganic Chemistry (eBook). Inorganic Chemical science. Vol. three (15th ed.). Switzerland: Pearson Prentice-Hall. p. 38. doi:10.1021/ed070pA304.1. ISBN9780273742753. Archived from the original on December 16, 2015. Retrieved December 14, 2020 – via Academy of Basel.
16. ^ Sanderson, R. T. (1983). "Electronegativity and bail energy". Periodical of the American Chemic Society. 105 (eight): 2259–2261. doi:ten.1021/ja00346a026.
17. ^ Sanderson, R. T. (1983). Polar Covalence . New York: Academic Press. ISBN978-0-12-618080-0.
18. ^ Zefirov, N. S.; Kirpichenok, M. A.; Izmailov, F. F.; Trofimov, G. I. (1987). "Adding schemes for atomic electronegativities in molecular graphs inside the framework of Sanderson principle". Doklady Akademii Nauk SSSR. 296: 883–887.
19. ^ Trofimov, Chiliad. I.; Smolenskii, Eastward. A. (2005). "Application of the electronegativity indices of organic molecules to tasks of chemic information science". Russian Chemical Message. 54 (ix): 2235–2246. doi:10.1007/s11172-006-0105-6. S2CID 98716956.
20. ^ SW Rick; SJ Stuart (2002). "Electronegativity equalization models". In Kenny B. Lipkowitz; Donald B. Boyd (eds.). Reviews in computational chemical science. Wiley. p. 106. ISBN978-0-471-21576-9.
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23. ^ Isle of mann, Joseph B.; Meek, Terry Fifty.; Allen, Leland C. (2000). "Configuration Energies of the Main Group Elements". Journal of the American Chemical Society. 122 (12): 2780–2783. doi:10.1021/ja992866e.
24. ^ Isle of man, Joseph B.; Meek, Terry 50.; Knight, Eugene T.; Capitani, Joseph F.; Allen, Leland C. (2000). "Configuration energies of the d-block elements". Periodical of the American Chemical Society. 122 (21): 5132–5137. doi:x.1021/ja9928677.
25. ^ ^{a b} The widely quoted Pauling electronegativity of 0.seven for francium is an extrapolated value of uncertain provenance. The Allen electronegativity of caesium is 0.66.
26. ^ Encounter, e.g., Bellamy, L. J. (1958). The Infra-Scarlet Spectra of Circuitous Molecules. New York: Wiley. p. 392. ISBN978-0-412-13850-eight.
27. ^ Spieseke, H.; Schneider, W. K. (1961). "Consequence of Electronegativity and Magnetic Anisotropy of Substituents on C13 and H1 Chemical Shifts in CH3X and CH3CH2X Compounds". Journal of Chemical Physics. 35 (2): 722. Bibcode:1961JChPh..35..722S. doi:10.1063/1.1731992.
28. ^ Clasen, C. A.; Good, M. L. (1970). "Estimation of the Moessbauer spectra of mixed-hexahalo complexes of tin can(IV)". Inorganic Chemistry. ix (4): 817–820. doi:10.1021/ic50086a025.
29. ^ Fleming, Ian (2009). Molecular orbitals and organic chemical reactions (Student ed.). Chichester, West Sussex, U.K.: Wiley. ISBN978-0-4707-4660-8. OCLC 424555669.
30. ^ "Electropositivity," Microsoft Encarta Online Encyclopedia 2009. (Archived 2009-10-31).

Bibliography [edit]

• Jolly, William 50. (1991). Modern Inorganic Chemistry (2nd ed.). New York: McGraw-Loma. pp. 71–76. ISBN978-0-07-112651-9.
• Mullay, J. (1987). "Interpretation of atomic and grouping electronegativities". Electronegativity. Structure and Bonding. Vol. 66. pp. ane–25. doi:10.1007/BFb0029834. ISBN978-three-540-17740-one.

External links [edit]

• Media related to Electronegativity at Wikimedia Eatables
• WebElements, lists values of electronegativities past a number of dissimilar methods of calculation
• Video explaining electronegativity
• Electronegativity Nautical chart, a summary listing of the electronegativity of each element along with an interactive periodic table

kauffmannhistras.blogspot.com

Source: https://en.wikipedia.org/wiki/Electronegativity

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