1. Determining if polar molecules are involved.

Steps and hints:

A. Draw the Lewis Dot Structure (section 8.6 - 8.8 in your text)
B. Determine the number of electron domains around the central atom
C. Determine the electron domain geometry and the molecular geometry (section 9.2 in your text)
D. Determine whether the molecule is polar (section 9.3 in your text)

Example:

Is ammonia, NH3, a polar molecule?
A. Draw the Lewis Dot Structure:

B. To determine the number of electron domains around the central atom remember that lone pairs, single bonds, double bonds and triple bonds each count as one domain so we have a total of 4 electron domains around the nitrogen.

C. Based on 4 electron domains, the electron domain geometry is tetrahedral. The geometries that correspond to the various numbers of domains can be found on Table 9.1 in your text book. The molecular geometry is then found by replacing the electron domains with bonds to atoms (single, double or triple bonds) or lone pairs. The final drawing should reflect the 3-D nature of the molecule.

D. For the most part, the molecules we deal with have polar bonds, the question is whether or not the polar bonds (the dipoles) will cancel out to give a non-polar molecule. Usually, if there are any lone pairs on the central atom or if there is more than one type of atom bonded to the central atom, then the dipoles will not cancel and the molecule is polar. If there are no lone-pairs on the central atom and all the atoms bonded to the central atom are identical then the dipoles will all cancel (even though it might not look like it) and the molecule will be non-polar. In our example, the nitrogen has a lone pair on it so, ammonia is a polar molecule.

Try this:

Is methanol, CH3OH polar?
answer: Yes, since the central C atom has a tetrahedral electronic geometry and is surrounded by three H's and one OH, the corresponding dipoles will not cancel out.