Electronegativity . It is a measure of the attractive force exerted by one atom on the electrons of another in a covalent bond. The different electronegativity values are classified according to different scales, among them the Pauling scale and the Mulliken scale.
Summary
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- 1 Definition
- 2 Periodic variation
- 3 Types of links
- 4 Pauling scale
- 1 Calculation of electronegativity
- 5 Source
- 6 See also
Definition
Electronegativity is a chemical property that measures the ability of an atom (or less frequently a functional group) to attract electrons , or electronic density, to it when it forms a covalent bond in a molecule . We must also consider the electron density distribution around a given atom in front of others, both in a molecular species and in a non-molecular compound.
Periodic variation
The electronegativities of the representative elements increase from left to right throughout the periods and from bottom to top within each group. The variations in electronegativities of the transition elements are not as regular. In general, the ionization energies and electronegativities are lower for the elements in the lower left zone of the periodic table than for those in the upper right zone.
The concept of electronegativity is very useful to know the type of bond that two atoms will originate in their union:
The bond between atoms of the same class and of the same electronegativity is apolar. The greater the electronegativity differences between two atoms, the greater the electron density of the molecular orbital in the vicinity of the more electronegative atom. A polar bond originates.
When the difference in electronegativities is high enough, a complete transfer of electrons occurs, leading to the formation of ionic species . Metals are low electronegativity so they tend to give away electrons to form positive ions, and nonmetals are less electronegative, so they form negative ions. Knowing the electronegativity of the elements, it is easy to predict the bonding model they will form (ionic, covalent, coordinated, polar and nonpolar); The stronger the bond, the greater the difference in electronegativities.
Example: potassium is an atom that has very little affinity for electrons. It has an electronegativity of only 0.8 while silicon (Si), whose affinity for electrons is moderate, has an electronegativity of 1.8
The elements, when combined with each other, exchange or sedate electrons, in such a way that they reach the electronic configuration similar to that of a noble gas, thus forming what we call chemical bonding .
Types of links
- In the ionic bonding model, the atom that gives up one electron or more becomes a cation and the one that receives it, an anion. The bond persists because the ions attract each other.
- In the covalent bonding model the valence electrons are shared by the two bonding atoms. The bond persists because the electrons are located between both nuclei and this prevents repulsion between them.
- In the metallic model the electrons move freely between ions (+). Precisely this property causes metals to conduct electric current .
Pauling scale
Pauling defined it as the ability of an atom in a molecule to attract electrons to it. Its values, based on thermochemical data, have been determined on an arbitrary scale, called the Pauling scale, whose maximum value is 4, which is the value assigned to fluorine , the most electronegative element. The least electronegative element, cesium , has an electronegativity of 0.7.
The electronegativity of a given atom is fundamentally affected by two quantities, its atomic mass and the average distance of the valence electrons with respect to the atomic nucleus. This property has been 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 measured experimentally directly, such as ionization energy, but can be indirectly determined by calculating 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 among 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. Scale that varies between 0.7 for the least electronegative element and 4.0 for the largest.
It is interesting to note that electronegativity is not strictly an atomic property, since it refers to an atom within a molecule3 and, therefore, it can vary slightly when the “environment” 4 of the same atom varies in different bonds of different molecules. The equivalent property of electronegativity for an isolated atom would be electronic affinity or electro-affinity.
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 atom with the highest electronegativity.
Calculation of electronegativity
Knowing the energies of the AA and BB bonds, parameter D can be calculated as follows:
- Δ = EAB – ½ (EAA + EBB)
This parameter gives an idea about the deviation that the AB bond may have from a pure covalent bond situation. The greater the D, the greater the ionic contribution to that bond, and in that sense the difference in electronegativity between both elements A and B increases.
Pauling’s definition of electronegativity is given by the following expression:
- | χA – χB | = 0.102 x √Δ Δ (kJ / mol)
The figure below shows the variation of Pauling’s electronegativity with respect to the atomic number, showing a similar evolution to that already observed in ionic radius or ionization energy.
The Pauling scale has some limitations, since for example they depend on the oxidation number of the element (the figure corresponds to the values of the maximum oxidation state of each element). However, this scale is very suitable for calculating bond energies between elements of different electronegativity as well as for a qualitative display of bond polarity.
Another definition of this atomic property was proposed by Robert Mulliken when observing that when an element has a high ionization energy, I, and a high electronic affinity, Ae, then it has a great tendency to acquire electrons rather than lose them. Therefore, when this element is part of a chemical compound, it must be quite electronegative. On the contrary, when both the ionization energy, I, and the electronic affinity, Ae, of an element have low values, this element will have a marked tendency to lose electrons when it is part of a compound and, consequently, it will be classified as an electropositive element. These observations led Mulliken to come up with his own definition of electronegativity known as Mulliken’s electronegativity, cM,
- χM = ½ (I + Ae)
If both I and Ae are high then cM will also have a high value; on the contrary if I and Ae are low then cM will have a small value. The Pauling and Mulliken electronegativity scales can be related to each other using the following expression:
- χP = 1.35 (√χM) -1.37
A third electronegativity scale was proposed by Allred and Rochow, on the basis that this property is determined by the electric field on the surface of the atom. The electric field for such a system must be proportional to Zef / r2, which explains the form of this definition:
- χAR = 0.744 + (0.3590 Zef / r2)
where r is expressed in Å. The numerical constants have been chosen to obtain electronegativity values of the same range as the other commented scales.