After Module 7, we move into Module 8. Module 8 discusses molecular structure which should have been covered in the very first week of class. This module is important. It is important to understand how atoms bond. Understanding this module is important for all Chemistry.

I do have good news. It isn’t hard. I touched on it a little in my discussion of Module 7. If you have explained the periodic table in detail, your student shouldn’t have any difficulty with this module.

I wrote about nomenclature in my post about on Module 3. You can find some links there to help you write the chemical formulas. Remember that all the atoms want 8 valence electrons in their outer shells. It makes them happy. They want to be like the noble gases. The cool kids on the block. They want to grab electrons. Some more than others. Some are happy to share as in a covalent bond. Some are takers as in an ionic bond.

It is easiest for me to remember that two non-metals form covalent bonds, and a metal and a non-metal form an ionic bond.  If you are still have trouble check out this page to compare the two. Here is a quiz to check your knowledge.

Lewis structures

Drawing Lewis structures is not difficult. You only need to know the number of  valence electrons which you get from the group number of the element. That is all you have to know.

For example, let’s take Oxygen.

Oxygen is in Group 6A. This tells me that Oxygen has 6 valence electrons. See? Easy. (My lovely drawing in Paint)

oxygenI just drew six, happy little dots around an oxygen atom. Now, I know that oxygen is a diatomic atom, meaning that they never go it alone. So, if I want to draw the Lewis structure for the oxygen bond, I just match up the lonely electrons.

oxygen bond
Yes, this may look like your face while you study Chemistry, but it is really just a Lewis structure showing how an O2 is bonded.

oxygen bond 2

You don’t need to draw all the electrons after you know how they are bonded. It is redundant. We know all the other electrons are happily paired together.

Let’s do one more because it is fun. How about PCl3

PClHere are the Lewis structures of Phosphorus and Chlorine.

Phosphorus has three unpaired electrons. No, you can’t pair up two of the singles. There must be an electron at each of the spots on the compass before you start pairing them up. If I were to pair up two electrons, there would be an empty spot. This is how we draw the Lewis structure in order to get the correct bonds.

Chlorine has one unpaired electron. Phosphorus wants three. One chlorine atom can’t give a Phosphorus atom three of its electrons. Remember, it wants to have a full 8 as well. (the octet rule). So, instead two more Chlorine atoms join the party.

Pcl 2


Does that just look like a bunch of red dots? It does. But let’s draw the bond lines.



Remove the rest of your paired electrons.

final pcl3And there you go. There is just a little more that you will have to learn in the next module about bond shapes. This is the foundation of drawing Lewis structures. I hope this helps.