Lesson 9: Polarity - Simple Cases
In the last module we used diatomic molecules as examples, where it was easy to see whether the molecule was polar or nonpolar. If the two atoms were the same, the molecule had to be nonpolar, and if they were different, the molecule was polar.
In general, a bond is polar if the two atoms involved are not the same. (Two common exceptions to this are C-H bonds and Cl-N bonds. C and H have very similar electronegativities and Cl and N have almost identical electronegativities.)
A molecule that contains polar bonds may or may not be polar, depending on its shape. Molecules that do not contain polar bonds are not polar unless the central atom has one or more nonbonded pairs of electrons.
The general principle is that if a molecule is symmetric, its electrons will be uniformly distributed and it will be nonpolar.
It’s the same idea you used to determine whether a bond was polar. If the two atoms were the same, the bond was symmetric and therefore nonpolar.
A molecule with polar bonds can be nonpolar if, for example, it has two polar bonds that are equal in polarity but exactly opposite in direction, effectively cancelling each other out. There are regular geometric arrangements of three, four, five, and six equally polar bonds in which all of the polar bonds together cancel one another out, making the molecule nonpolar. These regular arrangements are found in highly symmetric molecules.
The symmetry of a molecule is determined by its shape. The next page in this lesson will take you through the process of determining the shapes of molecules. But there are times when you can determine that a molecule is polar (or nonpolar) without taking the trouble to determine its shape.
We will be interested in the shapes of small molecules only. Large molecules rarely have symmetric shapes, so their polarity is determined mainly by whether or not they contain polar bonds and by the number of polar bonds compared to the number of nonpolar bonds they have. In fact, many large molecules have both polar and nonpolar regions and so behave a little like both types of compounds. The most common examples of these molecules are soaps and detergents.
The first case involves molecules that contain an H-N, an H-O, or an H-F interatomic bond. These molecules are always polar, and hydrogen bonding will always occur. Several examples are given in the figure.
HF is the only molecule that contains an H-F bond.
The second case involves molecules in which there is a central atom that has one or more unbonded (lone) pairs of electrons. To determine this, you may have to draw the Lewis structure. The lone pairs guarantee that the molecule will not be symmetric and that it will be polar.
You can be sure that a central nitrogen, oxygen, or fluorine atom has lone pairs of electrons. Molecules with only carbon and hydrogen never have lone pairs.
In molecules that have sulfur, phosphorus, chlorine, bromine, or iodine as the central atom, that atom may or may not have lone pairs. In these cases you must draw the Lewis structure to decide.
There are also two cases in which you can be certain that a molecule is nonpolar. The first is the trivial case in which a diatomic molecule consists of only one kind of atom.
Larger molecules, even if they have only one kind of atom, are sometimes polar. This will occur when the central atom has one or more pairs of nonbonded electrons. One example of this is ozone, O3. The middle oxygen atom has a lone pair of electrons and this lone pair gives the molecule its polarity.
The second is the case in which the molecule consists solely of carbon and hydrogen atoms. In this case, the molecule is always very close to being symmetric, and even when it is not, the polarity is so slight that the molecule behaves as if it were completely nonpolar.
Compounds that contain only carbon and hydrogen are known as hydrocarbons. They include almost all petroleum products. Oil and gas are mixtures of a large number of different hydrocarbons. Natural gas is largely a single hydrocarbon named methane, CH4. Some other hydrocarbons with which you may be familiar include butane, propane, acetylene, paraffin (a mixture), and benzene.
In general, however, to determine whether a molecule is polar, you must first determine its shape, and to determine its shape, you must start with its Lewis structure. That process is the subject of the next section of the lesson.
You may wish to review the part of lesson 8 on drawing Lewis structures before you proceed. The next section includes a few reminders, but if you are not confident you can draw the Lewis structure given the formula of a molecule, it may be worth your while to review the relevant section of Lesson 8.