Rules for creating Lewis Structures

So Far, I've created Lewis Structures for you pretty much 'out of the air'. In some of the simple cases like H2O it's pretty obvious where all the electrons go. In the more complicated cases, multiple bonds, polyatomic ions, etc. It may not be quite so easy to decide the best structure.

Follow these few rules and you should have no trouble creating even quite complicated Lewis structures. I'll enumerate the rules on the left hand side and prepare the structure of CO32– on the right-hand side as an example.

Rule Description Example CO32-
1 Add up all the electrons in the valence shell, including those due to the charge. 
 C ==>  4
3O ==> 18
2- ==>  2
total  24
2 Draw a skeletal structure (single bonds only) paying attention as much as possible to the valence of the atoms.
3 Subtract from the total electron count, the number you used to make the single bonds 
24-6=18
or 9 pairs
4 Add remaining electrons to the structure, starting with the most electronegative atoms first
5 Calculate the formal charge of the atoms (if they're all zero, this is your last step) 
O: 6-6-1= -1
C: 4-0-3= +1
6 Move electron pairs to create double bonds and thereby reducing formal charges as much as possible. Recheck the formal charges If there's no further way you can reduce them you're done.

Note that over all of this, we must be aware of and keep the octet rule.

Let's try another example. The common polyatomic anion, SO42–.

Count the electrons:

    S:  6
   4O: 24
Charge  2
total  32

Now draw the skeletal structure:

8 electrons used in the bonds. 32–8=24 electrons left to place (12 pairs). Calculate Formal Charge.

Now we reduce the formal charges on two of the oxygens and the sulfur by creating double bonds between the O and S using one lone pair from each of the two oxygens.
NOTE: This time we can violate the octet rule because sulfur's valence electrons are in the n=3 shell which has 3-d orbitals where we can place electrons. Only elements in row 3,4,... of the periodic table can do this.

Note that there are five other possible combinations of oxygens on which I could have placed the double bonds and I still would have drawn an identical Lewis Structure. This leads us to the next topic of discussion. In reality, the oxygens on the sulfate ion are identical chemically. That means that any Lewis Structure which seems to make a distinction between the oxygens with and without double bonds is not completely correct.