Physical Chemistry Third Edition

862 20 The Electronic States of Diatomic Molecules

Electronegativity

Relative energies of atomic orbitals on different atoms are not always immediately

available when we want to construct an approximate description of a molecule with

unequal sharing of electrons. Theelectronegativityis an empirical parameter that can

be used to estimate the degree of inequality of electron sharing in a bond between atoms

of two elements. It was introduced by Pauling, who observed that polar covalent bonds

generally have larger bond dissociation energies than nonpolar covalent bonds. If the

electronegativity of element A is denoted byXAand that of element B is denoted by

XB, Pauling defined

|XA−XB|(0.102 mol^1 /^2 kJ−^1 /^2 )(∆EAB)^1 /^2 (20.4-15)

where∆EABis the difference between the average bond energy of an A–B bond and

the mean of the average bond energies of A–A and B–B bonds:

∆EABEAB−

1

2

[EAA+EBB] (20.4-16)

Linus Pauling, 1901–1994, was a
prominent American chemist who won
the 1954 Nobel Prize in chemistry for
his work on molecular structure and the
1963 Nobel Peace Prize for his work on
nuclear disarmament.

Because Eq. (20.4-15) defines only electronegativity differences, Pauling arbitrarily

chose the value of 4.0 for fluorine, which makes the electronegativities range in value

from 0.7 to 4.0. Table A.21 in Appendix A contains electronegativity values for several

elements. Fluorine is the most electronegative element, followed by oxygen, nitrogen,

and chlorine. The alkali metals are the least electronegative.

In any row of the periodic table the electronegativity increases from left to right, and

in any column it decreases from top to bottom. We can understand these trends on the

basis of effective nuclear charges. As one moves from left to right across a row of the

periodic table, the effective nuclear charge increases, corresponding to lower orbital

energies and to the observed increase in the electronegativity. As one moves from top

to bottom in a column of the periodic chart, there are opposing tendencies. The nuclear

charge increases, which would seem to increase the electronegativity. However, each

row corresponds to a larger value of the principal quantum number than the shell above

it, so that the average distance of the electron from the nucleus is larger and there is

an additional shell of electrons to shield the electron from the nuclear charge. This

tendency dominates, corresponding to a less negative orbital energy for electrons in

the valence shell and a decrease in the electronegativity.

EXAMPLE20.14

Using average bond energies from Table A.9 of the appendix, calculate the difference in

electronegativity between H and F.

Solution

We omit the units on the factors:

|XH−XLi|(0.102)[568−(1/2)(436+158)]^1 /^2  1 .7 (Table value 1 .9)

Exercise 20.20

Using average bond energies from Table A.9 of the appendix, calculate the differences in elec-

tronegativity between (a) C and O, and (b) C and Cl. Compare your results with the values in

Table A.21 of the appendix.