Formal Charges: The Not-So-Mathematical Way
Saturday, January 29, 2011 at 12:46PM Formal Charges...how do they work?
A formal charge is a charge assigned to an electron in a molecule. This is the sum of the electrons and protons associated with the atom. If the atom has more electrons than protons, it will have a negative charge. When the opposite is true, the atom will have a positive charge. Formal charges help keep track of electrons.
That being said, formal charges aren't the actual charge of the atom; they are used as a guide to help draw the different resonance structures of the molecule which will ultimately lead to the resonance hybrid structure. The resonance hybrid structure is the best 'real' representation of a molecule. For now I'll just talk about an easy way to determine those formal charges.
My textbook uses a formal for calculating hybrid structures. It looks something like this:
Formal Charge=Group Number - nonbonding electrons - 1/2(shared electrons)
Group Number: The group number the atom belongs to on the periodic table. For example, carbon is in group four (fourth column from the right...don't count the columns between element 21 and 30). Oxygen is in group six.
Nonbonding Electrons: The number of electrons that are not shared with other atoms in the moledule.
Shared Electrons: The number of electrons that are shared with other atoms in the molecule.
This works, but it takes time to count up all of the electrons and figure out what's being shared and what's not. It would be much easier to just look at an atom and know what the formal charge is. Here's how! I'm going to assume that you know how to draw Lewis structures.
Lets use carbon as our example. Carbon is a very important element.
In this photo, carbon has four single bonds. It's not important where they go.
1. As I've said previously, carbon is in group four. It has four valence electrons in its outer shell. In this drawing, it's sharing all of them. For the purpose of determining formal charge quickly, just know that it's in group four, so keep the number four in mind.
2. We must now determine how many electrons are in carbon's OWN POSESSION. This 'own possession' thing is important. You might be tempted to say eight, but this is not true. It's sharing electrons; we can pretend that for one of single bond, carbon has one electron and the other atom has the other one. Since we have four single bonds on this carbon, we'll say that it has four electrons in its own posession.
3. Very easy math. Take the number of electrons in carbon's own poession (four) and subtract it from the group number (four). That equals zero. This carbon has no charge.
Easy...but what if carbon doesn't feel like sharing. See the drawing below.
This carbon only has three single bonds. How does that work? Lets figure it out.
1. Carbon is still and always will be in group four. Make note of that number.
2. In this case, carbon has three single bonds and a pair of unshared electrons. We only care about the electrons in carbon's own possession. In this case, we'll count one electron for each single bond another two electrons for the ones that are unshared. The unshared pair of electrons belong to carbon and no one else.
3. Do the math. The group number is 4 and the number of electrons in carbon's own possession is 5. 4-5 = -1. This carbon atom has a negative charge!
Lets do it quickly one more time with another atom. Lets use nitrogen.
1. Nitrogen is in group five.
2. Nitrogen has two single bonds and two pairs of unshared electrons. That equals to six electrons in its own possession.
3. 5-6 = -1 This nitrogen atom has a negative charge.
Here is a photo from my textbook (which I have referenced in this post) of some important atoms in organic chemistry. It shows common bonding patterns and their charges. You'll start to recognize the patterns and know right away what charges the atoms.
Wade, Leroy G. "Chapter 1-8." Organic Chemistry. Upper Saddle River [etc.: Prentice Hall, 2010. Print.
Thanks goes out to my O-Chem professor/advisor and to Dmitri Mendeleev.
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