Some molecules contain only nonpolar bonds - for example, methane, CH4. Such molecules are nonpolar molecules. Other molecules contain polar bonds - that is, bonds between atoms whose electronegativities differ by more than 0.4 units. Whether these latter molecules are polar or nonpolar depends on the arrangement in space of these bonds and the resulting geometry of the molecules. If we picture the geometry of a molecule and show its polar bonds with an arrow () aimed at the more electronegative atom, we can usually obtain a picture of the molecule that indicates whether or not it is polar. Figure 7.11 illustrates this method for several molecules.
|FIGURE 7.11 Polarity in molecules: (a) methane has no polar bonds and is a nonpolar molecule; (b) methyl fluoride has one polar bond (denoted by ) and is therefore a polar molecule; (c) carbon tetrafluoride has four counter-balanced polar bonds and thus is a nonpolar molecule.|
Which of the molecules in Figure 7.11 is polar? Methane, which contains no polar bonds, is clearly nonpolar. Methyl fluoride contains one polar bond between carbon (EN = 2.5) and fluorine (EN = 4.0). Methyl fluoride is a polar molecule; the negative end of the dipole is at the fluoride atom. Now look at carbon tetrafluoride. It contains four polar carbon-fluorine bonds but they counteract one another, so the molecule itself is nonpolar.
The principle is like that of erecting an antenna tower (Figure 7.12). If the guy wires are correctly balanced against one another, the tower stays erect. (Compare the balanced pull of the carbon-fluorine bonds in carbon tetrafluoride.) If the guy wires are not balanced, the tower topples. (Compare the unbalanced pull of the carbon-fluorine bond in methyl fluoride.)
|FIGURE 7.12 Erecting an antenna tower: (a) antenna tower with balanced guy wires (compare with carbon tetrafluoride); (b) antenna tower with unbalanced guy wires (compare with methyl fluoride).|
Figure 7.13 shows three other molecules and their polarities. Notice that both ammonia and water have unshared electrons on the atom at the negative end of the dipole. These unshared electrons enhance the polarity of the bond to make these molecules very polar.
|FIGURE 7.13 Predicting the polarity of molecules: (a) ammonia, NH3, has unbalanced polar bonds and is a polar molecule; (b) carbon dioxide, CO2, has balanced polar bonds and is a nonpolar molecules; (c) water, H2O, has unbalanced polar bonds and is a polar molecule.|
Predict whether the following molecules are polar or nonpolar.
a. CH2O b. SiBr4 c. SO2
The Lewis structure of CH2O is:
The electronegativity difference between carbon and oxygen is 1.0. Thus, the C = O bond will be polar, with a partial positive charge on the carbon and a partial negative charge on the oxygen. Thje C - H bond is not polar. By showing the polar bond as an arrow, the diagram shows clearly that the molecule is polar ( the bond is not balanced ). The unshared electrons on the oxygen should enhance this polarity.
b. The Lewis structure of silicon tetrabromide is:
The electronegativity of silicon is 1.8, that of bromine is 2.8. The silicon - bromine bond is polar. Showing these bonds as arrows in a tetrahedral structure clarifies that silicon tetrabromide is a nonpolar molecule. The polar bonds balance each other as in carbon tetrafluoride and the molecule s nonpolar.
c. The Lewis structure of sulfur dioxide is:
The electronegativity of sulfur is 2.5 and that of oxygen is 3.5; thus the sulfur-oxygen bonds are polar. By drawing these polar bonds as arrows in the bent molecule of sulfur dioxide, we show its polar nature:
The molecule is polar. This molecule is a resonance hybrid, but this fact does not affect its polarity.