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Jump to one of the following Lewis structures' parts:   Choose One Introduction to Lewis Structures Part 1- Basics Part 2- Multiple Bonding and Polyatomic Ions Part 3- Resonance Part 4- Hypervalency Part 5- Formal Charge I Part 6- Formal Charge II Part 7- Electron Deficiency Part 8- Oxoacids Part 9- Quizzes

We're going to shift gears slightly for this section. When we learned about electron configurations, we learned that it could be thought of as a bookkeeping system for electrons. We also learned how to write valence configurations for atoms and ions. If an ion has more electrons than its corresponding neutral atom, it has a negative charge (anion). In a similar manner, if an ion has less electrons than its corresponding neutral atom, it has a positive charge (cation).

We'll do something similar for the atoms in a compound and we'll call this system "formal charge".

Formal charge is the apparent charge an atom has in a compound. To calculate it-

Where we have certain definitions:

• FC- formal charge
• # of valence e^- in the atom- the number of valence electrons in the free atom just as you've been calculating right along.
• # of bonds- the number of bonds to the atom in question where a single bond counts as 1 bond, a double bond counts as 2 bonds, and a triple bond counts as 3 bonds.
• # of unshared electrons- the total number of unshared electrons surrounding the atom in question (note that this is not the number of lone pairs).

Since the formal charge is the apparent charge each atom has in the compound, the sum of all the formal charges must be equal to the total charge on the compound. In other words, the sum of the formal charges on carbon and oxygen in CO must sum up to 0 since CO is a neutral molecule. The sum of the formal charges on the nitrogen and oxygen atoms in NO₃^-1 must sum up to -1. Likewise, the sum of the formal charges on nitrogen and oxygen in NO⁺¹ must equal +1.

Something to remember is that you must include the sign even if it's positive.

So, how do we put all this to work? For right now we'll concentrate simply on calculating formal charge and we'll put this to work in the next section. Let's take a look at a few structures for which you've already learned to draw Lewis structures.

Pull up the Periodic Table if you need one.

Let's look at O₂.

In case you've forgotten how to draw the Lewis structure for O₂.

Now let's calculate the formal charge on each oxygen atom. The last check is to make sure the sum of the formal charges add up to 0. Why 0? What's the overall charge on the O₂? Since there's no indication of a positive or negative charge, it must be 0.

oxygen (left):

• FC = 6 - [2 + 4] = 0
• # valence electrons in the atom = 6
• # bonds to the atom = 2
• # unshared electrons = 4

oxygen (right):

• FC = 6 - [2 + 4] = 0
• # valence electrons in the atom = 6
• # bonds to the atom = 2
• # unshared electrons = 4

sum of the formal charges = 0 (oxygen, left) + 0 (oxygen, right) = 0

A bit excessive, you say? Perhaps. But, if you pay close attention to the process on problems that are relatively easy like O₂ it will help you on the more difficult problems.

Now let's do CH₄. The format will be somewhat similar to the previous format for drawing Lewis structures.

Here's a current list of the rules.

And here's the formula for formal charge if you wish to have it available.

Your turn- enter your answer for the formal charges and hit "Verify"
Determine the Lewis structure for CH4. (Just in case you've forgotten!)
Calculate the formal charge for each of the hydrogens. Notice that we can get away with calculating it once since it's the same for each hydrogen.
Calculate the formal charge for carbon.
Sum up the formal charges.

When summing up the formal charges, you have to do it for each atom. That's why there's a "4*0" above. There are 4 hydrogens and the formal charge on each one is 0.

You've hopefully noticed that the formal charges for everything up until this point have been 0. If it were only that easy all the time! Remember, the sum of the formal charges must be equal to the charge on the species. If we're talking about an ion, then at least one of the atoms must have a nonzero formal charge.

Now let's do NO¹⁺.

Your turn- enter your answer for the formal charges and hit "Verify"
Determine the Lewis structure for NO1+. (Just in case you've forgotten!)
Calculate the formal charge for the oxygen.
Calculate the formal charge for nitrogen.
Sum up the formal charges.

It should be apparent that an ion must have a nonzero formal charge for at least one of its atoms. But, that doesn't mean that's the only time an atom in a compound has a nonzero formal charge. To demonstrate, let's look at a neutral species which has atoms with nonzero formal charges in it.

Now let's do O₃.

Your turn- enter your answer for the formal charges and hit "Verify"
Determine the Lewis structure for O3. (Just in case you've forgotten!)
Calculate the formal charge for oxygen #1 (bottom).
Calculate the formal charge for oxygen #2 (middle).
Calculate the formal charge for oxygen #3 (right).
Sum up the formal charges.

And now you might be able to see a couple of difficulties with identifying formal charges. You need to be very clear how you are assigning them. For O₃, you can't just say something like "the formal charge on oxygen is 0" because there is more than one oxygen atom in the molecule. I chose to label the atoms descriptively. Another way is to give the atoms labels like a, b, c, etc., or 1, 2, 3, etc. Another reason you need to be clear is because there are generally multiple ways to draw Lewis structures. And, in the case of O₃, there are resonance structures to deal with. If you need to refresh your memory about resonance structures, take a gander and then come back. Regardless of how you choose to do it, and if your instructor gives you a choice, be clear about it.