Lewis Structures
The ability to draw Lewis structures is paramount to success in organic chemistry. Organic chemistry is about drawing structures, so you need to make sure that you can rationalize through a wide variety of compounds. In fact, many questions in organic chemistry place an emphasis on you being able recognize and draw out the structure correctly. If you can draw the right structure, you have succeeded in half the battle. Herein, I give you a fool proof method to draw Lewis structures and also, some tips on some common pitfalls.
Drawing Lewis structures:
Step 1: count up all valence electrons. This brings to question; how do you count the valence electrons? The periodic table is our friend here. We can simply use the periodic table to easily figure out the number of valence electrons. The periodic table has groups that are listed within it and these group numbers tell us the valence electrons. For instance, Li is in group 1A. This means that Li has 1 valence electron. You simply move from left to right to increase the group number. So, after lithium, you have beryllium which has 2, then boron which has 3 and so forth. Not too difficult right?
Practice:
How many valence electrons are in the following?
C:
O:
N:
F:
Once you have figured out the valence electrons, there is one other aspect that you should consider; the number of bonds that are typically present. You need to recognize the norms to be able to quickly assess how many different compounds are assembled. Use the information below to help you, especially in organic chemistry.
B: typically has 3 bonds
C: typically has 4 bonds
N: typically has 3 bonds and one lone pair
O: typically has 2 bonds and 2 lone pairs
F: typically has 1 bond and 3 lone pairs
Keep this information in mind when constructing Lewis structures. Also, remember that many of these traits are the same for those in the same column. For instance, all halogens typically have 1 bond and 3 lone pairs. Also note how these bond characteristics allow for the formation of an octet. Alright, let’s move to step 2.
Step 2: Once you have the values of valence electrons, you will need to analyze the list of what was given and find the central atom. Often, this is easy in that the one that is present in the least amount is the central atom.
Step 3: Once you determine the central atom, put all of them together by drawing a line to connect them. Remember, each line is 2 electrons so we need to subtract this from our total electron count (our total valence electron count that is).
Step 4: After you have everything connected, you need to make a judgement call on how many electrons surround an atom and if this is correct. At this point, formal charge can help you. Wait, what the heck is formal charge?? Well, formal charge allows us to determine whether or not an atom should have a charge. Remember when we were talking about how many bonds elements typically have? This is part of formal charge. For instance, we said oxygen should have 2 bonds and 2 lone pairs. What happens if oxygen has 3 bonds and 1 lone pair? At this point, the oxygen will have a formal charge of 1. So, how did I figure this out? There is an easy formula to simply plug in and get a direct answer. However, you will eventually see these enough that you can do this in your head.
So, for oxygen with 3 bonds how can we calculate formal charge? Well, oxygen is in group 6 so it has 6 valence electrons. We can use this for the formula. Formal charge = valence electrons (group number)-(number of electrons in lone pair +1/2 the number of electrons in bonds). So, let’s look at that example of oxygen with 3 bonds. FC = 6-(2+3) = 1. So, using this formula, you find that oxygen with 3 bonds should have a formal charge of 1. I would recommend that you commit this formula to memory until you can do it in your head. This way, you will never get formal charge wrong.
Back to where we were. If you connect several atoms and find that there should be a formal charge, you must make sure that the molecule has such charges. For instance, if your connections lead to a molecule that has an overall charge of +1, then you should see this in your formula. Note: molecules can have charges but also have a net neutral charge and this is perfectly fine. For instance, if one part of the molecule has a minus charge and another has a positive charge, the overall is neutral and that may be just fine. We will see some of this in the examples section.
Step 5: If the molecule does not work as you have it drawn with all lone pairs, convert some lone pairs to multiple bonds. So for instance, if you have an oxygen with only one bond, you will likely note that this will be a negative formal charge. Take one of the surrounding electron pairs and convert that to a new bond. Often this helps to fix several problems. Note: an important concept is that several atoms can have expanded octets! Atoms such as S and P are known to expand their octets on many occasions. To help, remember that phosphorous typically has either 3 or 5 bonds and sulfur will have 2, 4, and 6.
I can write about this until my fingers fall off, but examples are worth their weight in gold, or palladium catalysts if you will.
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