We previously discussed whether a covalent bond is polar or nonpolar based on the electronegativity difference of the two atoms involved in the bond. Entire molecules can also be polar or nonpolar. If a molecule has one polar bond, it is a polar molecule. There will be a net dipole moment in the molecule so that the centers of positive and negative charge are separated. This is the case for HCl and CH3F. If a molecule has two or more identical polar bonds, it may or may not be polar depending on the three-dimensional orientation of the polar bonds. For example, carbon dioxide has two polar carbon-oxygen double bonds. However, because carbon dioxide has a linear geometry, these two identical polar bonds cancel each other out, and carbon dioxide is a nonpolar molecule. Water also has two polar bonds, but water is a polar molecule. Since water has a bent molecular geometry, the two identical polar bonds cannot cancel each other out. There are specific cases when identical polar bonds can cancel each other out. If a molecule has all identical polar bonds and there are no lone pairs of electrons on the central atom, then those polar bonds will cancel each other out, and the molecule overall will be nonpolar. This rule only applies when all of the polar bonds are identical. If the central atom has lone pairs of electrons, then even if a molecule has all identical polar bonds, they cannot cancel each other out; the molecule will still be polar. This is the case for bent, trigonal pyramidal, seesaw, T-shaped, and square pyramidal molecular geometries. Even identical polar bonds do not cancel each other out in these geometries. However, there are two exceptions. If we have two bonding groups and three lone pairs, we'll have the linear molecular geometry. In this case, the identical polar bonds can cancel each other out, and the lone pairs of electrons can cancel each other out separately. So with identical polar bonds, the linear molecular geometry is nonpolar. The same is true if we have four bonding groups and two lone pairs, which is square planar molecular geometry. The four identical polar bonds cancel each other out, and the two lone pairs cancel each other out separately, so that the square planar molecular geometry is nonpolar. Let's practice determining whether molecules are polar or nonpolar. We'll start with carbon tetrafluoride. First, we want to draw the Lewis structure. Then we can assess whether or not our four polar carbon-fluorine single bonds cancel each other out. This molecule has a tetrahedral molecular geometry. Since all bonds are identical polar bonds and there are no lone pairs of electrons on our central atom, the polar bonds will cancel each other out, and carbon tetrafluoride is a nonpolar molecule. In our next example, we'll determine if CH2F2 is polar or nonpolar. We'll start by drawing the Lewis structure. We have two polar carbon-fluorine bonds. Although from this Lewis structure it looks like the two polar bonds could cancel each other out, we need to remember that the Lewis model is a two-dimensional model. It only shows what atoms are bonded and where our valence electrons are located. To determine if the two polar bonds cancel, we need to consider the three-dimensional molecular geometry. Our molecular geometry is tetrahedral, so these two bonds are at an angle of about 109.5°. They cannot cancel each other out, so our molecule overall is polar. In our last example, we'll determine if sulfur tetrafluoride is polar or nonpolar. We'll start again with our Lewis structure. We have four bonding groups and one lone pair of electrons on sulfur, so our molecular geometry is seesaw. With a seesaw molecular geometry, even all identical polar bonds cannot cancel each other out, so our molecule is polar.