Electronegativity

The electronegativity is the force, the power of an atom to attract electrons towards itself. We must also consider the distribution of electron density around an atom determined compared to other different ones, both in a molecular species and in non-molecular systems or species.

The electronegativity of a given atom is fundamentally affected by two magnitudes: its atomic number and the average distance of the valence electrons from the atomic nucleus. This property has been able to be correlated with other atomic and molecular properties. Linus Pauling was the researcher who first proposed this magnitude in 1932, as a further development of his valence bond theory. Electronegativity cannot be directly measured experimentally, like ionization energy, for example, but it can be determined indirectly by making calculations from other atomic or molecular properties.

Different methods have been proposed for its determination and although there are small differences between the results obtained, all the methods show the same periodic trend between the elements.

The most common calculation procedure is the one initially proposed by Pauling. The result obtained by this procedure is a dimensionless number that is included within the Pauling scale. This scale varies between 0.65 for the least electronegative element (francium) and 4.0 for the most electronegative (fluorine).

It is interesting to note that electronegativity is not strictly an atomic property, as it refers to an atom within a molecule and, therefore, it can vary slightly when the “environment” of the same atom varies in different bonds. different molecules. The equivalent property of electronegativity for an isolated atom would be electron affinity or electroaffinity.

Two atoms with very different electronegativities form an ionic bond. Pairs of atoms with small differences in electronegativity form polar covalent bonds with the negative charge on the higher electronegativity atom.

Electronegativity scales

The different values of electronegativity are classified according to different scales, among them the aforementioned Pauling scale and the Mulliken scale.

In general, the different electronegativity values of atoms determine the type of link that will be formed in the molecule that combines them. Thus, according to the difference between electronegativity (Δ Δ χ χ {displaystyle scriptstyle Delta chi }) of these can be determined (conventionally) if the link will be, according to the Linus Pauling scale:

• Covalent apolar: 0≤ ≤ Δ Δ χ χ ≤ ≤ 0.4{displaystyle 0leq Delta chi leq 0.4}.
• Polar covalent: 0.5≤ ≤ Δ Δ χ χ ≤ ≤ 1.6{displaystyle 0.5leq Delta chi leq 1.6}.
• Ionic:1.7≤ ≤ Δ Δ χ χ ≤ ≤ r2.0{displaystyle 1.7leq Delta chi leq r2.0}.

The smaller the atomic radius, the higher the ionization energy, the higher the electrongativity and vice versa. Electrongativity is the tendency or ability of an atom, in a molecule, to attract the electrons. Neither quantitative definitions nor electronity scales are based on electronic distribution, but on properties that are supposed to reflect electronivity. The electronegativity of an element depends on its oxidation state and therefore is not an invariable atomic property. This means that a single element can present different electrongativities depending on the type of molecule in which it is found, for example, the ability to attract electrons from a hybrid orbital spn{displaystyle sp^{n} in a carbon atom linked with a hydrogen atom, it increases in line with the percentage of character s in the orbital, according to the ethanus series Δ ethylene(eth) ≤ acetylene(ethin). The Pauling scale is based on the difference between the A-B link energy in the compound ABn{displaystyle AB_{n}} and the average of the energies of the A-A and B-B homopolar links.

Fluorine is the most electronegative element on the periodic table, while Francium is the least electronegative element on the periodic table. It is very important to know that electronegativity values go from bottom to top and from left to right.

R. S. Mulliken proposed that the electronegativity of an element can be determined by averaging the ionization energy of its valence electrons and the electron affinity. This approximation agrees with Pauling's original definition and gives orbital electronegativities and not invariant atomic electronegativities.

The Mulliken scale (also called the Mulliken-Jaffe scale) is a scale for the electronegativity of chemical elements, developed by Robert S. Mulliken in 1934. is based on the Mulliken electronegativity (c_M) which averages the electron affinity A.E. (magnitude that can be related to the tendency of an atom to acquire a negative charge) and the ionization potentials of its valence electrons P.I. or E.I. (magnitude associated with the ease, or tendency, of an atom to acquire a positive charge). The units used are the kJ/mol:

χ χ =12(Eea+Ei){displaystyle chi ={frac {1}{2}}}(E_{rm {ea}+}E_{rm {i}})}

The following table shows some values of electronegativity for representative elements on the Mulliken scale:

E. G. Rochow and A. L. Alfred defined electronegativity as the attractive force between a nucleus and an electron of a bonded atom.

Electronegativities of the elements

Measurements on the Pauling scale.

Other systems for measuring electronegativity

Allred and Rochow Electronegativity

Albert L. Allred and Eugene G. Rochow considered that electronegativity must be related to the charge experienced by an electron at the "surface" of an atom: the greater the charge per unit atomic surface area, the greater the tendency of that atom to attract electrons. The effective nuclear charge, z _ef , experienced by an electron of valence s can be estimated using Slater's rules, while the surface area of an atom in a molecule can be can be taken as proportional to the square of the Covalent Radius, R _Cov . When R _Cov is expressed in picometers, alt URL

Sanderson Electronegativity Equalization

Robert Thomas Sanderson also noted the relationship between Mulliken electronegativity and atomic size, and proposed a calculation method based on the reciprocal of atomic volume. With knowledge of bond lengths, Sanderson's model allows the estimation of binding energies for a wide range of compounds. The Sanderson model has also been used to calculate molecular geometry, S -Electron Energy, spin-spin NMR coupling constants, and other parameters for organic compounds. This work underlies the concept of electronegativity equalization, which suggests that electrons are distributed around a molecule to minimize or equal Mulliken electronegativity. This behavior is analogous to chemical potential equalization in macroscopic thermodynamics.

Allen's 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 valence electrons in a free atom,

Electronegative group

In organic chemistry, electronegativity is associated more with different functional groups than with individual atoms. The terms electronegative group and electronegative substituent can be considered synonymous terms. It is quite common to distinguish between inductive and resonance effects, effects that could be described in terms of σ and π electronegativities, respectively. There are also a number of linear relationships with free energy that have been used to quantify these effects, with Hammet's equation being the best known.