We now have a better sense of the shapes that molecules adopt and why they do so. We will spend the rest of this chapter looking more closely at the ways in which electrons are shared to form the bonds between atoms in molecules. We will begin by returning to a topic that we first discussed in Section 8.5, namely bond polarity and dipole moments. Recall that bond polarity is a measure of how equally the electrons in a bond are shared between the two atoms of the bond: As the difference in electronegativity between the two atoms of a bond increases, so does the bond polarity. We saw that the dipole moment of a diatomic molecule is a quantitative measure of the amount of charge separation in the molecule.
How can we extend the notion of bond polarity and dipole moments to the collection of bonds in a polyatomic molecule? For a molecule with more than two atoms, the dipole moment depends on both the polarities of the individual bonds and the geometry of the molecule. For each bond in the molecule, we can consider the bond dipole, which is the dipole moment due only to the two atoms in that bond. For example, consider the CO2 molecule, which is linear. As shown in Figure 9.9(a), each CO bond is polar, and, because the CO bonds are identical, the bond dipoles are equal in magnitude.
Figure 9.9 The overall dipole moment of a molecule is the sum of its bond dipoles. (a) In CO2 the bond dipoles are equal in magnitude but exactly oppose each other. The overall dipole moment is zero. (b) In H2O the bond dipoles are also equal in magnitude but do not exactly oppose each other. The molecule has a nonzero overall dipole moment.
Does the fact that both CO bonds are polar mean that the CO2 molecule as a whole is polar? Not necessarily. Bond dipoles and dipole moments are vector quantities; that is, they have both a magnitude and a direction. The overall dipole moment of a polyatomic molecule is the sum of its bond dipoles. Both the magnitudes and the directions of the bond dipoles must be considered in this sum of vectors. The two bond dipoles in CO2, although equal in magnitude, are exactly opposite in direction. Adding them together is the same as adding two numbers that are equal in magnitude but opposite in sign, such as 100 + (-100): The bond dipoles, like the numbers, "cancel" each other. Therefore, the overall dipole moment of CO2 is zero; CO2 is a nonpolar molecule. Note that the oxygen atoms in CO2 carry a partial negative charge and that the carbon atom carries a partial positive charge, as we expect for polar bonds. Even though the individual bonds are polar, the geometry of the molecule dictates that the overall dipole moment be zero.
Now let's consider H2O, which is a bent molecule with two polar bonds [Figure 9.9(b)]. Again, both the bonds are identical, so the bond dipoles are equal in magnitude. Because the molecule is bent, however, the bond dipoles do not directly oppose each other and therefore do not cancel each other. Hence, the H2O molecule has an overall nonzero dipole moment ( = 1.85 D). Because H2O has a nonzero dipole moment, it is a polar molecule. The oxygen atom carries a partial negative charge, and the hydrogen atoms each have a partial positive charge.
Figure 9.10 shows examples of polar and nonpolar molecules, all of which have polar bonds. The molecules in which the central atom is symmetrically surrounded by identical atoms (BF3 and CCl4) are nonpolar. For ABn molecules in which all the B atoms are the same, certain symmetrical geometries—linear (AB2), trigonal planar (AB3), tetrahedral and square planar (AB4), trigonal bipyramidal (AB5), and octahedral (AB6)—must lead to nonpolar molecules even though the individual bonds might be polar.
Figure 9.10 Examples of molecules with polar bonds. Two of these molecules have a zero dipole moment because their bond dipoles cancel one another.
Predict whether the following molecules are polar or nonpolar: (a) BrCl; (b) SO2; (c) SF6.
SOLUTION (a) Chlorine is more electronegative than bromine. Consequently, BrCl will be polar with chlorine carrying the partial negative charge:
Experimentally, the dipole moment of the molecule is 0.57 D. All diatomic molecules with polar bonds are polar molecules.
(b) Because oxygen is more electronegative than sulfur, the molecule has polar bonds. Several resonance forms for SO2 can be written:
For each of these, the VSEPR model predicts a bent geometry. Because the molecule is bent, the bond dipoles do not cancel and the molecule is polar ( = 1.63 D):
(c) Fluorine is more electronegative than sulfur. The bond dipoles therefore point toward fluorine. The six SF bonds are arranged in an octahedral fashion around the central sulfur:
The symmetrical octahedral geometry of the molecule leads to cancellation of the bond dipoles, and the molecule is nonpolar ( = 0).
Are the following molecules polar or nonpolar: (a) NF3; (b) BCl3? Answers: (a) polar; (b) nonpolar