Lesson 9: Molecular Polarity
Once we know its shape, we can determine whether a molecule is polar or nonpolar. A polar molecule is one in which one side, or end, of the molecule has a slight positive charge and the other side, or end, has a slight negative charge. This will occur whenever the molecule is not completely symmetric. (Unless, of course, the molecule is a hydrocarbon.)
The other simple case – a molecule that consists of two atoms of the same element – is nonpolar precisely because it is completely symmetric.
A nonpolar molecule is one which is completely symmetric. In the last page of the lesson, we pointed out the symmetric shapes: linear, trigonal planar, and tetrahedral. In order for a molecule to be nonpolar, it must have one of these shapes. But that’s not quite enough.
(Side note: There are a few other equally symmetric shapes that arise when there are five or six groups of electrons around the central atom. We will not deal with them here (they all involved central atoms that have more than an octet, which is called an "expanded" octet). To satisfy your curiosity, they are called a trigonal bipyramid and an octahedron. Interestingly, when there are six groups of electrons around the central atom, it is possible to have a central atom with two lone pairs that still has its outer atoms arranged in a completely symmetric manner. This shape is called square planar.)
In addition to having a symmetric shape, for the molecule to be completely symmetric, all of the atoms that are attached to the central atom must be the same. For this reason, SO3 is completely symmetric but SOCl2 is not, even though they are both trigonal planar.
Symmetry, therefore, has two components: the geometric arrangement of the outer atoms and whether or not they are all the same.
Ultimately, it is the uniform, or non-uniform, distribution of electrons that determines whether a molecule is polar or nonpolar, but this distribution is, in turn, determined by the distribution and identity of the atoms in a molecule.
Let’s illustrate with two examples, PCl3 and CCl4. To determine whether a molecule is polar or nonpolar, unless the molecule is a hydrocarbon, or has only two atoms of the same element, first complete its Lewis structure.
Next, use the Lewis structure to determine the shape of the molecule. There are only five possible shapes we will deal with.
The symmetric shapes are linear, trigonal planar, and tetrahedral.
The unsymmetric shapes are bent and trigonal pyramidal.
Then look to see whether the shape is symmetric. If not, the molecule is polar.
Even if all the outer atoms are the same, a molecule with an unsymmetric shape will be (at least slightly) polar. The lone pair or pairs of electrons on the central atom guarantee a nonuniform distribution of electrons.
The polarity of the bonds – if they are polar – also contributes to the polarity of the molecule.
If the shape is symmetric, look to see whether all of the atoms attached to the central atom are the same. If so, the molecule is nonpolar. If there is more than one kind of atom attached to the central atom, the molecule is polar.
If all the outer atoms are the same (and the molecule has a symmetric shape), the electrons will be distributed uniformly, even if the bonds are very polar.
The symmetric shape and the fact that the polarities of the bonds are exactly the same means that the polarities of the bonds cancel each other out, leaving the molecule as a whole nonpolar.
Many molecules are nonpolar, but have polar bonds. A bond is polar if the two atoms on either end are different. The one general exception to this rule is the bond between C and H. C-H bonds are only very slightly polar, and so behave as though they were nonpolar.
The other, much less common, exception to this rule is the N-Cl bond. Nitrogen and chlorine have almost identical electronegativities, and so the N-Cl bond is essentially nonpolar. Even so, the one common compound that contains N-Cl bonds, NCl3, is still slightly polar because of the lone pair of electrons on the nitrogen atom.
Take a few minutes to try filling in the first two columns in the table in Exercise 7 in your workbook. Indicate in the column labeled “Molecular Polarity” whether the molecule is polar or nonpolar. In the column labeled “Bond Polarity” indicate whether any of the bonds in the molecule are polar or if they are all nonpolar. (Leave the third column blank for now. We will deal with it in a later section of the lesson.)
Here are the answers:
- CH4: nonpolar bonds (or very slighty polar), nonpolar molecule
- NH3: polar bonds, polar molecule
- H2O: polar bonds, polar molecule
- SO3: polar bonds, nonpolar molecule
- SO2: polar bonds, polar molecule
- CH3Cl: polar bond (C-Cl), polar molecule
- HCN: polar bond (C-N), polar molecule
- CH2O: polar bond (C-O), polar molecule
- O3: nonpolar bond, polar molecule (lone pair on the central O)
- BH3: polar bonds, nonpolar molecule
- PCl3: polar bonds, polar molecule
If you had trouble with any of these, be sure to get help from an instructor before moving on.
Except for hydrocarbons, larger and more complex molecules are almost always at least slightly polar. In these cases, the relative number of polar and nonpolar bonds determine the behavior of the molecule. Example 8 in your workbook shows some of these more complex molecules.
The first three are polar because of the C-O and O-H bonds they contain. The fourth and fifth are hydrocarbons and therefore are nonpolar.
When a large molecule, the size of octane or larger, contains only one or two polar bonds, even though the molecule overall is polar, it behaves in many ways like a nonpolar molecule. The very large nonpolar region overshadows the small polar section in determining miscibility, for example. Nonetheless, such a molecule will generally have a higher boiling and melting point than a molecule of equal size with no polar bonds.