Electronegativity

Bond Polarity

An important variation in covalent bonds is in the attraction exerted on the electrons by the two atoms that are bonded together. If there's an equal attraction from both atoms, then we have a nonpolar bond. If one atom exerts a stronger pull on the electrons than the other, then we have a polar bond. Of course, there is a wide range in the degree of polarity.

Lewis diagrams showing polar and nonpolar bonds.

Electronegativity

We have a name for the amount of pull that one atom exerts on the electrons that it is sharing with other atoms. It is called electronegativity.

Electronegativity is useful for all elements on the periodic table but it is most useful for the nonmetals, which are shown here to the right.

Electronegativity is more of a concept than a property. As such, values for it are estimated or calculated rather than measured. Over the years, chemists have come up with a variety of ways of calculating values for electronegativity, and also a variety of values. Here is one particular set of values.

Look at how the electronegativity changes as you go across and up and down the periodic table (for nonmetals,right, and all elements, below). No values are given for the inert gases because they do not bond readily to other atoms. We cannot talk about how strongly they attract electrons when they bond, if they don't bond. Actually, a few do form bonds with fluorine and oxygen.

Periodic table (right half) with electronegativity values.

Periodic table with electronegativity values.

The electronegativity generally increases as you go from left to right across the periodic table. It decreases as you go down the periodic table.

I think you can see that the reason for this is going to depend on those same factors that we used to explain the trends in atomic size, ionization energy, and electron affinity.

First, the horizontal comparison. As you go across a period from left to right, the atoms of each element all have the same number of energy levels and the same number of shielding electrons. Thus the factor that predominates is the increased nuclear charge. When the nuclear charge increases, so will the attraction that the atom has for electrons in its outermost energy level and that means the electronegativity will increase.

Now what about the vertical comparison? As we established previously, when you go from one atom to another down a group, you are adding one more energy level of electrons for each period. The increased shielding nearly balances the increased nuclear charge and the predominant factor is the number of energy levels that are used by the electrons. So as you go from fluorine to chlorine to bromine and so on down the periodic table, the electrons are further away from the nucleus and better shielded from the nuclear charge and thus not as attracted to the nucleus. For that reason the electronegativity decreases as you go down the periodic table.

You might have noticed that there are some high values in the middle of the transition group. Those elements have fairly high electronegativities for metals. The reason for that ties in with the arrangement of electrons and the fact that the atoms are using d orbitals. The way that the d orbitals are shielded is different than the way that s and p orbitals are shielded so there is some variation in the transition metals that is not as easily explained as the general trend that I have just been talking about.

Practice

Take a moment to do parts b, c, d, and e of exercise 26 in your workbook.

Electronegativity - Continued

So far, our comparison of electronegativities has been limited to left and right comparisons within a period and vertical comparisons within a group. Sometimes you have to make comparisons of elements that are not both in the same period or both in the same group. One way is simply to look up values for electronegativities of the elements you are comparing. Another way of remembering how the electronegativities of nonmetals compare is by using the word FONClBrISCHP. It has a spelling that corresponds with the symbols of many of the nonmetals in decreasing order of electronegativity. F for fluorine, O for oxygen, N for nitrogen. Those are the 3 most electronegative elements. Chlorine is very close behind nitrogen, perhaps they are tied. Then comes bromine, iodine, followed by sulfur, carbon, hydrogen, and phosphorus. So FONClBrISCHP is a very useful word in helping you remember the order of the electronegativities of the nonmetals.

We can use electronegativity to predict and explain the polarity of bonds between pairs of atoms.

For example, the bond between hydrogen and chlorine is a polar covalent bond because chlorine is significantly more electronegative than hydrogen so chlorine has a stronger pull on the electrons than does hydrogen. The bond between carbon and oxygen is also a polar covalent bond because oxygen is more electronegative than carbon. The bond between two hydrogen atoms is a nonpolar covalent bond because each atom has the same electronegativity. Because the electronegativities of chlorine and bromine are only slightly different, the bond between them is slightly polar.

H-Cl	polar
C-O	polar
H-H	nonpolar
Cl-Br	slightly polar

Practice

Take some time now to determine the polarity of the bonds shown in exercise 27 in your workbook. The answers follow.

Answers

P-F is a polar bond. F-F is a nonpolar bond. H-O is a polar bond. H-C is a slightly polar bond. There is not much difference between the electronegativities of hydrogen and carbon. C-C is a nonpolar covalent bond. If you had trouble with any of those be sure to check with an instructor to figure out why.

Comparing Covalent and Ionic Bonding

Let's push some of these ideas a bit further. By looking at electronegativity we can talk about gradations in metallic and nonmetallic character. Although there are many inconsistencies, we can generalize that metals have low electronegativities (generally below 2) and nonmetals have high electronegativities (generally above 2). We can also generalize about ionic and covalent bonding in this way. Covalent bonding results when there is a small difference in the electronegativities of the two elements. Ionic bonding results when there is a very large difference in electronegativities between the two elements. Some chemists set the dividing line between a small difference and a large difference at about 1.7 to 1.9.

If we select pairs of elements, such as those shown here (and also in example 28 in your workbook), and compare how different their electronegativities are, you get a wide range of differences. Consequently, you get a gradation of bond types. Not just covalent and ionic, but nonpolar covalent, slightly and very polar covalent, slightly ionic and very ionic. Covalent and ionic bonding can be viewed as extremes on a continuum rather than just different types of bonds. As differences between electronegativities become larger, the bonds become more ionic. As the differences become smaller, the bonds become more covalent.

covalent	F-F	nonpolar covalent
	Cl-F	slightly polar covalent
	H-F	very polar covalent
ionic	Ag-F	slightly ionic
ionic	K-F	very ionic

This approach works well if one or both of the elements is a nonmetal. In the next portion of this lesson we will look at what happens when the bonding involves just metal atoms.

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