La electronegatividad , simbolizada como χ , es la tendencia de un átomo a atraer electrones compartidos (o densidad de electrones ) hacia sí mismo. [1] La electronegatividad de un átomo se ve afectada tanto por su número atómico como por la distancia a la que residen sus electrones de valencia del núcleo cargado. Cuanto mayor es la electronegatividad asociada, más electrones atrae un átomo o un grupo sustituyente. Lo opuesto a la electronegatividad es la electropositividad : una medida de la capacidad de un elemento para donar electrones de valencia.
En el nivel más básico, la electronegatividad está determinada por factores como la carga nuclear ( cuantos más protones tiene un átomo, más "atracción" tendrá sobre los electrones) y el número y ubicación de otros electrones en las capas atómicas ( cuantos más electrones tenga). tiene un átomo, cuanto más lejos del núcleo estén los electrones de valencia y, como resultado, menos carga positiva experimentarán, tanto por su mayor distancia del núcleo como porque los otros electrones en los orbitales del núcleo de menor energía actuarán para proteger los electrones de valencia del núcleo cargado positivamente).
El término "electronegatividad" fue introducido por Jöns Jacob Berzelius en 1811, [2] aunque el concepto se conocía antes y fue estudiado por muchos químicos, incluido Avogadro . [2] A pesar de su larga historia, no se desarrolló una escala precisa de electronegatividad hasta 1932, cuando Linus Pauling propuso una escala de electronegatividad que depende de las energías de enlace, como un desarrollo de la teoría del enlace de valencia . [3] Se ha demostrado que se correlaciona con otras propiedades químicas. La electronegatividad no se puede medir directamente y debe calcularse a partir de otras propiedades atómicas o moleculares. Se han propuesto varios métodos de cálculo, y aunque puede haber pequeñas diferencias en los valores numéricos de la electronegatividad, todos los métodos muestran las mismas tendencias periódicas entre elementos .
El método de cálculo más utilizado es el propuesto originalmente por Linus Pauling. Esto da una cantidad adimensional , comúnmente conocida como escala de Pauling ( χ r ), en una escala relativa que va de 0,79 a 3,98 ( hidrógeno = 2,20). Cuando se utilizan otros métodos de cálculo, es convencional (aunque no obligatorio) citar los resultados en una escala que cubra el mismo rango de valores numéricos: esto se conoce como electronegatividad en unidades de Pauling .
Como se suele calcular, la electronegatividad no es una propiedad de un átomo solo, sino más bien una propiedad de un átomo en una molécula . [4] Las propiedades de un átomo libre incluyen energía de ionización y afinidad electrónica . Es de esperar que la electronegatividad de un elemento varíe con su entorno químico, [5] pero generalmente se considera una propiedad transferible , es decir, que valores similares serán válidos en una variedad de situaciones.
El cesio es el elemento menos electronegativo (0,79); el flúor es el más (3,98).
Métodos de cálculo
Electronegatividad de Pauling
Pauling propuso por primera vez [3] el concepto de electronegatividad en 1932 para explicar por qué el enlace covalente entre dos átomos diferentes (A – B) es más fuerte que el promedio de los enlaces A – A y B – B. Según la teoría del enlace de valencia , de la cual Pauling fue un defensor notable, esta "estabilización adicional" del enlace heteronuclear se debe a la contribución de las formas canónicas iónicas al enlace.
La diferencia de electronegatividad entre los átomos A y B viene dada por:
donde las energías de disociación , E d , de los enlaces A – B, A – A y B – B se expresan en electronvoltios , incluyéndose el factor (eV) - 1 ⁄ 2 para asegurar un resultado adimensional. Por lo tanto, la diferencia en la electronegatividad de Pauling entre el hidrógeno y el bromo es 0,73 (energías de disociación: H – Br, 3,79 eV; H – H, 4,52 eV; Br – Br 2,00 eV)
Como solo se definen las diferencias de electronegatividad, es necesario elegir un punto de referencia arbitrario para construir una escala. Se eligió el hidrógeno como referencia, ya que forma enlaces covalentes con una gran variedad de elementos: su electronegatividad se fijó primero [3] en 2,1, luego se revisó [6] a 2,20. También es necesario decidir cuál de los dos elementos es más electronegativo (equivalente a elegir uno de los dos signos posibles para la raíz cuadrada). Esto generalmente se hace usando "intuición química": en el ejemplo anterior, el bromuro de hidrógeno se disuelve en agua para formar iones H + y Br - , por lo que se puede suponer que el bromo es más electronegativo que el hidrógeno. Sin embargo, en principio, dado que se deben obtener las mismas electronegatividades para dos compuestos de enlace cualesquiera, los datos están de hecho sobredeterminados y los signos son únicos una vez que se fija un punto de referencia (generalmente, para H o F).
Para calcular la electronegatividad de Pauling para un elemento, es necesario tener datos sobre las energías de disociación de al menos dos tipos de enlaces covalentes formados por ese elemento. AL Allred actualizó los valores originales de Pauling en 1961 para tener en cuenta la mayor disponibilidad de datos termodinámicos, [6] y son estos valores "revisados de Pauling" de la electronegatividad los que se utilizan con mayor frecuencia.
El punto esencial de la electronegatividad de Pauling es que existe una fórmula semi-empírica subyacente, bastante precisa, para las energías de disociación, a saber:
oa veces, un ajuste más preciso
Esta es una ecuación aproximada pero se mantiene con buena precisión. Pauling lo obtuvo al señalar que un enlace se puede representar aproximadamente como una superposición mecánica cuántica de un enlace covalente y dos estados de enlace iónico. La energía covalente de un enlace es aproximada, por cálculos de mecánica cuántica, la media geométrica de las dos energías de enlaces covalentes de las mismas moléculas, y hay energía adicional que proviene de factores iónicos, es decir, el carácter polar del enlace.
La media geométrica es aproximadamente igual a la media aritmética - que se aplica en la primera fórmula anterior - cuando las energías son de un valor similar, por ejemplo, excepto para los elementos altamente electropositivos, donde hay una diferencia mayor de dos energías de disociación; la media geométrica es más precisa y casi siempre da un exceso de energía positiva, debido al enlace iónico. La raíz cuadrada de este exceso de energía, señala Pauling, es aproximadamente aditiva y, por lo tanto, se puede introducir la electronegatividad. Por lo tanto, es esta fórmula semi-empírica para la energía de enlace la que subyace al concepto de electronegatividad de Pauling.
Las fórmulas son aproximadas, pero esta aproximación aproximada es de hecho relativamente buena y da la intuición correcta, con la noción de la polaridad del enlace y alguna base teórica en mecánica cuántica. A continuación, se determinan las electronegatividades para que se ajusten mejor a los datos.
En compuestos más complejos, existe un error adicional ya que la electronegatividad depende del entorno molecular de un átomo. Además, la estimación de energía solo se puede usar para enlaces simples, no múltiples. La energía de la formación de una molécula que contiene solo enlaces simples se puede aproximar posteriormente a partir de una tabla de electronegatividad y depende de los constituyentes y la suma de cuadrados de las diferencias de electronegatividades de todos los pares de átomos enlazados. Una fórmula de este tipo para estimar la energía suele tener un error relativo del orden del 10%, pero se puede utilizar para obtener una idea cualitativa aproximada y una comprensión de una molécula.
1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 | 11 | 12 | 13 | 14 | 15 | dieciséis | 17 | 18 | ||
---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
Grupo → | |||||||||||||||||||
↓ Periodo | |||||||||||||||||||
1 | H 2,20 | Él | |||||||||||||||||
2 | Li 0,98 | Ser 1,57 | B 2.04 | C 2.55 | N 3.04 | O 3,44 | F 3,98 | Nordeste | |||||||||||
3 | Na 0,93 | Mg 1,31 | Al 1,61 | Si 1,90 | P 2.19 | S 2.58 | Cl 3.16 | Arkansas | |||||||||||
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 | Como 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 | En 1,78 | Sn 1,96 | Sb 2.05 | Te 2.1 | Yo 2,66 | Xe 2.60 | |
6 | Cs 0,79 | Ba 0,89 | Lu 1.27 | Hf 1.3 | Ta 1,5 | Ancho 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 >0.79[en 1] | Ra 0.9 | Lr 1.3[en 2] | Rf | Db | Sg | Bh | Hs | Mt | Ds | Rg | Cn | Nh | Fl | Mc | Lv | Ts | Og | |
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 | ||||||
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 |
See also: Electronegativities of the elements (data page)
- ^ The electronegativity of francium was chosen by Pauling as 0.7, close to that of caesium (also assessed 0.7 at that point). The base value of hydrogen was later increased by 0.10 and caesium's electronegativity was later refined to 0.79; however, no refinements have been made for francium as no experiment has been conducted. However, francium is expected and, to a small extent, observed to be more electronegative than caesium. See francium for details.
- ^ See Brown, Geoffrey (2012). The Inaccessible Earth: An integrated view to its structure and composition. Springer Science & Business Media. p. 88. ISBN 9789401115162.
Mulliken electronegativity
Robert S. Mulliken proposed that the arithmetic mean of the first ionization energy (Ei) and the electron affinity (Eea) should be a measure of the tendency of an atom to attract electrons.[7][8] As this definition is not dependent on an arbitrary relative scale, it has also been termed absolute electronegativity,[9] with the units of kilojoules per mole or electronvolts.
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,[10]
and for energies in kilojoules per mole,[11]
The Mulliken electronegativity can only be calculated for an element for which the electron affinity is known, fifty-seven elements as of 2006. The Mulliken electronegativity of an atom is sometimes said to be the negative of the chemical potential.[12] 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.,
Allred–Rochow electronegativity
A. Louis Allred and Eugene G. Rochow considered[13] 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, Zeff, experienced by valence electrons 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, rcov. When rcov is expressed in picometres,[14]
Sanderson electronegativity equalization
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.[15] With a knowledge of bond lengths, Sanderson's model allows the estimation of bond energies in a wide range of compounds.[16] Sanderson's model has also been used to calculate molecular geometry, s-electrons energy, NMR spin-spin constants and other parameters for organic compounds.[17][18] 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.[19] This behavior is analogous to the equalization of chemical potential in macroscopic thermodynamics.[20]
Allen electronegativity
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 electrons in a free atom,[21][22][23]
where εs,p are the one-electron energies of s- and p-electrons in the free atom and ns,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 that are numerically similar to Pauling electronegativities.
The one-electron energies can be determined directly from 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.[24] 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.
In this scale neon has the highest electronegativity of all elements, followed by fluorine, helium, and oxygen.
Group → | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 |
---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
↓ Period | ||||||||||||||||||
1 | H 2.300 | He 4.160 | ||||||||||||||||
2 | Li 0.912 | Be 1.576 | B 2.051 | C 2.544 | N 3.066 | O 3.610 | F 4.193 | Ne 4.787 | ||||||||||
3 | Na 0.869 | Mg 1.293 | Al 1.613 | Si 1.916 | P 2.253 | S 2.589 | Cl 2.869 | Ar 3.242 | ||||||||||
4 | K 0.734 | Ca 1.034 | Sc 1.19 | Ti 1.38 | V 1.53 | Cr 1.65 | Mn 1.75 | Fe 1.80 | Co 1.84 | Ni 1.88 | Cu 1.85 | Zn 1.588 | Ga 1.756 | Ge 1.994 | As 2.211 | Se 2.424 | Br 2.685 | Kr 2.966 |
5 | Rb 0.706 | Sr 0.963 | Y 1.12 | Zr 1.32 | Nb 1.41 | Mo 1.47 | Tc 1.51 | Ru 1.54 | Rh 1.56 | Pd 1.58 | Ag 1.87 | Cd 1.521 | In 1.656 | Sn 1.824 | Sb 1.984 | Te 2.158 | I 2.359 | Xe 2.582 |
6 | Cs 0.659 | Ba 0.881 | Lu 1.09 | Hf 1.16 | Ta 1.34 | W 1.47 | Re 1.60 | Os 1.65 | Ir 1.68 | Pt 1.72 | Au 1.92 | Hg 1.765 | Tl 1.789 | Pb 1.854 | Bi 2.01 | Po 2.19 | At 2.39 | Rn 2.60 |
7 | Fr 0.67 | Ra 0.89 | ||||||||||||||||
See also: Electronegativities of the elements (data page) |
Correlación de la electronegatividad con otras propiedades.
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,[4] 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:[25] 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[26] or isomer shifts in Mössbauer spectroscopy[27] (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][4]
Tendencias en electronegatividad
Periodic trends
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 gases), whereas caesium is the least electronegative, at least of those elements for which substantial data is available.[24] This would lead one to believe that caesium fluoride is the compound whose bonding features the most ionic character.
There are some exceptions to this general rule. Gallium and germanium have higher electronegativities than aluminium and silicon, respectively, because of the d-block contraction. Elements of the fourth period immediately after the first row of the transition metals have unusually small atomic radii because the 3d-electrons are not effective at shielding the increased nuclear charge, and smaller atomic size correlates with higher electronegativity (see Allred-Rochow electronegativity, 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.
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.[6] 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 | pKa |
---|---|---|---|
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 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 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 pKa of log10( 1⁄4) = –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 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 spx 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 χ(sp3) < χ(sp2) < χ(sp) holds (the trend should apply to non-integer hybridization indices as well). While this holds true in principle for any main-group element, values for the hybridization-specific electronegativity are most frequently 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)[28] |
---|---|
C(sp3) | 2.3 |
C(sp2) | 2.6 |
C(sp) | 3.1 |
'generic' C | 2.5 |
Electronegatividad de grupo
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 relationships that have been used to quantify these effects, of which the Hammett equation is the best known. Kabachnik parameters are group electronegativities for use in organophosphorus chemistry.
Electropositividad
Electropositivity is a measure of an element's ability to donate electrons, and therefore form 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 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 energies.[29]
While electronegativity increases along periods in the periodic table, and decreases down 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.
Ver también
- Electronegativities of the elements (data page)
- Chemical polarity
- Metallic bonding
- Orbital hybridization
- Oxidation state
- Periodic table
- Ionization energy
- Electron affinity
Referencias
- ^ a b IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "Electronegativity". doi:10.1351/goldbook.E01990
- ^ a b Jensen, W.B. (1996). "Electronegativity from Avogadro to Pauling: Part 1: Origins of the Electronegativity Concept". Journal of Chemical Education. 73 (1): 11–20. Bibcode:1996JChEd..73...11J. doi:10.1021/ed073p11.
- ^ a b c Pauling, L. (1932). "The Nature of the Chemical Bond. IV. The Energy of Single Bonds and the Relative Electronegativity of Atoms". Journal of the American Chemical Society. 54 (9): 3570–3582. doi:10.1021/ja01348a011.
- ^ a b c Pauling, Linus (1960). Nature of the Chemical Bond. Cornell University Press. pp. 88–107. ISBN 978-0-8014-0333-0.
- ^ Greenwood, N. N.; Earnshaw, A. (1984). Chemistry of the Elements. Pergamon. p. 30. ISBN 978-0-08-022057-4.
- ^ 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.
- ^ 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 (11): 782–793. Bibcode:1934JChPh...2..782M. doi:10.1063/1.1749394.
- ^ Mulliken, R. S. (1935). "Electronic Structures of Molecules XI. Electroaffinity, Molecular Orbitals and Dipole Moments". J. Chem. Phys. 3 (9): 573–585. Bibcode:1935JChPh...3..573M. doi:10.1063/1.1749731.
- ^ Pearson, R. G. (1985). "Absolute electronegativity and absolute hardness of Lewis acids and bases". J. Am. Chem. Soc. 107 (24): 6801–6806. doi:10.1021/ja00310a009.
- ^ Huheey, J.E.; Keiter, E.A.; Keiter, R.L. (December 1, 2008) [1978]. "17". In Kauffman, G.B. (ed.). Inorganic Chemistry: Principles of Structure and Reactivity (digitalized). Inorganic Chemistry: Principles of Structure and Reactivity (3rd ed.). New York (published 1900). p. 167. doi:10.1021/ed050pA379.1. ISBN 9780060429874. OCLC 770736023. inorganicchemist00huhe_0. Archived from the original on January 1, 2014. Retrieved December 15, 2020 – via Oxford University Press.
- ^ This second relation has been recalculated using the best values of the first ionization energies and electron affinities available in 2006.
- ^ Franco-Pérez, Marco; Gázquez, José L. (31 October 2019). "Electronegativities of Pauling and Mulliken in Density Functional Theory". Journal of Physical Chemistry A. 123 (46): 10065–10071. doi:10.1021/acs.jpca.9b07468. PMID 31670960.
- ^ Allred, A. L.; Rochow, E. G. (1958). "A scale of electronegativity based on electrostatic force". Journal of Inorganic and Nuclear Chemistry. 5 (4): 264–268. doi:10.1016/0022-1902(58)80003-2.
- ^ Housecroft, C.E.; Sharpe, A.G. (November 1, 1993). Inorganic Chemistry (eBook). Inorganic Chemistry. 3 (15th ed.). Switzerland: Pearson Prentice-Hall. p. 38. doi:10.1021/ed070pA304.1. ISBN 9780273742753. Archived from the original on December 16, 2015. Retrieved December 14, 2020 – via University of Basel.
- ^ Sanderson, R. T. (1983). "Electronegativity and bond energy". Journal of the American Chemical Society. 105 (8): 2259–2261. doi:10.1021/ja00346a026.
- ^ Sanderson, R. T. (1983). Polar Covalence. New York: Academic Press. ISBN 978-0-12-618080-0.
- ^ Zefirov, N. S.; Kirpichenok, M. A.; Izmailov, F. F.; Trofimov, M. I. (1987). "Calculation schemes for atomic electronegativities in molecular graphs within the framework of Sanderson principle". Doklady Akademii Nauk SSSR. 296: 883–887.
- ^ Trofimov, M. I.; Smolenskii, E. A. (2005). "Application of the electronegativity indices of organic molecules to tasks of chemical informatics". Russian Chemical Bulletin. 54 (9): 2235–2246. doi:10.1007/s11172-006-0105-6. S2CID 98716956.
- ^ SW Rick; SJ Stuart (2002). "Electronegativity equalization models". In Kenny B. Lipkowitz; Donald B. Boyd (eds.). Reviews in computational chemistry. Wiley. p. 106. ISBN 978-0-471-21576-9.
- ^ Robert G. Parr; Weitao Yang (1994). Density-functional theory of atoms and molecules. Oxford University Press. p. 91. ISBN 978-0-19-509276-9.
- ^ Allen, Leland C. (1989). "Electronegativity is the average one-electron energy of the valence-shell electrons in ground-state free atoms". Journal of the American Chemical Society. 111 (25): 9003–9014. doi:10.1021/ja00207a003.
- ^ Mann, Joseph B.; Meek, Terry L.; 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.
- ^ Mann, Joseph B.; Meek, Terry L.; Knight, Eugene T.; Capitani, Joseph F.; Allen, Leland C. (2000). "Configuration energies of the d-block elements". Journal of the American Chemical Society. 122 (21): 5132–5137. doi:10.1021/ja9928677.
- ^ 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. New York: Wiley. p. 392. ISBN 978-0-412-13850-8.
- ^ Spieseke, H.; Schneider, W. G. (1961). "Effect 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.
- ^ Clasen, C. A.; Good, M. L. (1970). "Interpretation of the Moessbauer spectra of mixed-hexahalo complexes of tin(IV)". Inorganic Chemistry. 9 (4): 817–820. doi:10.1021/ic50086a025.
- ^ Fleming, Ian (2009). Molecular orbitals and organic chemical reactions (Student ed.). Chichester, West Sussex, U.K.: Wiley. ISBN 978-0-4707-4660-8. OCLC 424555669.
- ^ "Electropositivity," Microsoft Encarta Online Encyclopedia 2009. (Archived 2009-10-31).
Bibliografía
- Jolly, William L. (1991). Modern Inorganic Chemistry (2nd ed.). New York: McGraw-Hill. pp. 71–76. ISBN 978-0-07-112651-9.
- Mullay, J. (1987). "Estimation of atomic and group electronegativities". Electronegativity. Structure and Bonding. 66. pp. 1–25. doi:10.1007/BFb0029834. ISBN 978-3-540-17740-1.
enlaces externos
- Media related to Electronegativity at Wikimedia Commons
- WebElements, lists values of electronegativities by a number of different methods of calculation
- Video explaining electronegativity
- Electronegativity Chart, a summary listing of the electronegativity of each element along with an interactive periodic table