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Electronegativity, symbol χ, is a chemical property that describes the tendency of an atom to attract shared electrons towards itself. At the most basic level, electronegativity is determined by factors such as nuclear charge (the more protons an atom has, the more it will attract electrons) and the number and position of other electrons present in the various atomic orbitals (the more electrons an atom has, the farther away from the nucleus the valence electrons are, which will therefore be subject to less positive charge, either because they are farther from the nucleus or because they are shielded by the other electrons present in the lower energy orbitals).
Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called electronegativity. Electronegativity is a measure of the tendency of an atom to attract electrons (or electron density) towards itself. It determines how the shared electrons are distributed between the two atoms in a bond. The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity. Electrons in a polar covalent bond are shifted toward the more electronegative atom; thus, the more electronegative atom is the one with the partial negative charge. The more significant is the difference in electronegativity, the more polarized are the electron distribution, and the larger is the partial charges of the atoms.
We must be careful not to confuse electronegativity and electron affinity. Electronegativity describes how tightly an atom attracts electrons in a bond. It is a dimensionless quantity that is calculated, not measured. Linus Pauling derived the first electronegativity values by comparing the amounts of energy required to break different types of bonds. He chose an arbitrary relative scale ranging from 0 to 4.
The term electronegativity was introduced by Jöns Jacob Berzelius in 1811, although the concept was known even before and was studied by many chemists including Amedeo Avogadro. Despite its long history, an accurate scale of electronegativity was proposed only in 1932, when Linus Pauling in developing the theory of valence bond proposed a scale of electronegativity based on bond energies. Electronegativity has been shown to correlate with various other chemical properties. Electronegativity cannot be measured directly and must be determined from other atomic or molecular properties. Several methods of calculation have been proposed and all methods give similar results, although with small differences in the numerical values of electronegativity. However, the most commonly used electronegativity values remain those of Pauling.
Electronegativity is not a property of a single atom, but rather a property of an atom in a molecule. It can be expected that the electronegativity of an element may depend in part on its chemical surroundings, oxidation state, and coordination number, but it is usually considered a transferable property, meaning that it will retain similar values even in very different chemical species.
Variations along the periodic table
Regardless of the chosen scale, electronegativity values show a fairly regular trend along the periodic table of elements. In general, electronegativity increases from left to right across a period in the periodic table and decreases down a group. Thus, the nonmetals, which lie in the upper right, tend to have the highest electronegativities, with fluorine the most electronegative element of all (EN = 4.0). Metals tend to be less electronegative elements, and the group 1 metals have the lowest electronegativities. Note that noble gases are excluded from this figure because these atoms usually do not share electrons with other atoms since they have a full valence shell.
As a result, fluorine is the most electronegative element (excluding the noble gases), and cesium is almost always the least electronegative. This often leads to the view that cesium fluoride is the compound where the bond has the most ionic character.
Variations with oxidation number
In inorganic chemistry, it is considered that a single electronegativity value can be applied to a given element, valid in most “normal” situations. This approach has the advantage of simplicity, but it must be remembered that the electronegativity of an element is not an invariant atomic property, and in particular depends on the charge and oxidation state of the element. As can be expected, positive ions are more electronegative than the corresponding neutral atom, while negative ions are less electronegative than the corresponding neutral atom.
Similarly, the electronegativity of an element increases as its oxidation state increases. However, attributing precise numerical values is difficult. Allred used Pauling’s method to calculate the electronegativity for the different oxidation states of some elements (including tin and lead) for which adequate data were available. However, for most elements this calculation is not possible, since there are not enough different covalent compounds whose dissociation energies are known. This is particularly the case for transition elements, for which the reported electronegativity values are inevitably an average over different oxidation states, and consequently the respective electronegativity trends are not very meaningful.
Variations with hybridization
The electronegativity of an atom also changes with the hybridization of the orbitals used in the bond. Electrons in s orbitals are more bonded to the nucleus than electrons in p orbitals. Thus, an atom using a spx hybrid will be increasingly polarizable as the p contribution increases. Consequently, the electronegativity of an atom as a function of the hybridization scheme used will drop in the order χ(sp) > χ(sp2) > χ(sp3).
In general these considerations are valid for any element of the main groups, but they are mostly used for carbon (see next table). In organic chemistry these electronegativity values are invoked to predict or rationalize bond polarity in organic compounds containing double and triple bonds to carbon.
Electronegativity and bond type
The absolute value of the difference in electronegativity (ΔEN) of two bonded atoms provides a rough measure of the polarity to be expected in the bond and, thus, the bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is large, the bond is polar covalent or ionic.
The absolute values of the electronegativity differences between the atoms in the bonds H–H, H–Cl, and Na–Cl are 0 (nonpolar), 0.9 (polar covalent), and 2.1 (ionic), respectively. The degree to which electrons are shared between atoms varies from completely equal (pure covalent bonding) to not at all (ionic bonding).
|between 0.4 and 1.8
This table is just a general guide, however, with many exceptions. For example, the H and F atoms in HF have an electronegativity difference of 1.9, and the N and H atoms in NH3 a difference of 0.9, yet both of these compounds form bonds that are considered polar covalent. Likewise, the Na and Cl atoms in NaCl have an electronegativity difference of 2.1, and the Mn and I atoms in MnI2 have a difference of 1.0, yet both of these substances form ionic compounds.
The best guide to the covalent or ionic character of a bond is to consider the types of atoms involved and their relative positions in the periodic table. Bonds between two nonmetals are generally covalent; bonding between a metal and a nonmetal is often ionic.
Some compounds contain both covalent and ionic bonds. The atoms in polyatomic ions, such as OH–, NO3−, and NH4+, are held together by polar covalent bonds. However, these polyatomic ions form ionic compounds by combining with ions of opposite charge. For example, potassium nitrate, KNO3, contains the K+ cation and the polyatomic NO3− anion. Thus, bonding in potassium nitrate is ionic, resulting from the electrostatic attraction between the ions K+ and NO3−, as well as covalent between the nitrogen and oxygen atoms in NO3−.