Electronegativity

The charge distribution in bonds is commonly discussed in terms of the electronegativity, χ, of the elements involved (there should be little danger of confusing this use of χ with its use to denote an atomic orbital, which is another common convention). The electronegativity is a parameter introduced by Linus Pauling as a measure of the power of an atom to attract electrons to itself when it is part of a compound. Pauling used valence-bond arguments to suggest that an appropriate numerical scale of electronegativities could be defined in terms of bond dissociation energies, D, in electronvolts and proposed that the difference in electronegativities could be expressed as , | χA − χB|=0.102{D(A-B) − 1/2[D(A-A) +D(B-B)]}^1/2 , Electronegativities based on this definition are called Pauling electronegativities. A list of Pauling electronegativities is given in Table 11.4. The most electronegative elements are those close to fluorine; the least are those close to caesium. It is found that the greater the difference in electronegativities, the greater the polar character of the bond. The difference for HF, for instance, is 1.78; a C-H bond, which is com monly regarded as almost nonpolar, has an electronegativity difference of 0.51. The spectroscopist Robert Mulliken proposed an alternative definition of electronegativity. He argued that an element is likely to be highly electronegative if it has a high ionization energy (so it will not release electrons readily) and a high electron affinity (so it is energetically favorable to acquire electrons). The Mulliken electronegativity scale is therefore based on the definition , χ_M= 1/2(I + Eea) , where I is the ionization energy of the element and Eea is its electron affinity (both in electronvolts, Section 10.4e). The Mulliken and Pauling scales are approximately in line with each other. A reasonably reliable conversion between the two is χP =1.35χM ^1/2− 1.37.

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