Lewis structures

Drawing Lewis structures for molecules or ions is fairly easy once you practice on a few examples. Simply follow the steps below.

Count the total number of valence electrons in the molecule or ion
Draw a skeleton structure for the molecule, joining the atoms with single bonds. This can be a problem if you are not certain which atoms are joined, and is one of the reasons that condensed structural formulas are used. For molecules that consist of a central atom and some outlying atoms, the central atom is written first. (I.e., in ammonia, NH3, the N is the central atom.)
Subtract the number of electrons used in step 2 (Two electrons/bond) from the total number of valence electrons from step 1. If the remainder is zero, you're done. If not, go to step 4.
Count the number of electrons it will take to create an octet around each atom. (Or two electrons for H atoms.) If this is the same as the number of extra valence electrons from step 3, then simply add lone pairs to all the atoms that don't have octets. If it is more than the number from step 3, you will have to place multiple bonds. If you are two electrons short, create 1 double bond, if you are four short create a triple bond or two double bond, and so on.
This is best illustrated with an example: let's consider the Lewis structure for formaldehyde, CH2O.

First, count the valence electrons. Hydrogen has 1, and there are two of them, so total 2. Carbon has 4, and oxygen has 6. Add all these up and you'll get 12 total valence electrons.
Now draw a skeleton structure. Everything in the formaldehyde molecule is attached to the central carbon, so draw single bonds from the carbon to each hydrogen and the oxygen





There are three single bonds in the above structure. Each uses two electrons, so we have a total of 12 - 6 = 6 valence electrons left.
Each hydrogen has two electrons around it, so they have the correct complement of electrons. The carbon has only 6 electrons, but wants 8. The oxygen has only 2 and wants 8, so we need a total of 8 electrons to fill out octets on the carbon and oxygen. We only have 6 available though: thus, we must put in a double bond. The only place we can do this is between the carbon and oxygen, since the hydrogens already have the maximum number of electrons. This uses two of the six electrons we have left.




The carbon now has eight electrons around it, so its octet is filled. The oxygen only has four electrons and needs eight, but we have four electrons left, so place two lone pairs on the oxygen.



This structure used all electrons and has an octet around all atoms, so it's correct.

Example 1: What is the Lewis stucture for ethylene, CH2CH2?

Solution 1:

Each hydrogen has 1 valence electron, each carbon has four, so we have a total of 12.
The condensed structural formula for ethylene implies that the molecule has the two carbons joined and two hydrogens bonded to each carbon, like the skeleton below:



This uses up 10 total electrons
We have 12 electrons and used 10 in the skeleton above, so we need to place 2 more
The hydrogens all have 2 electrons, so they are happy. The carbons each have six, but need eight, so we need a total of 4 electrons. We only have two, so we must place a double bond.



This structure has an octet for all atoms, so we're done.
Example 2: What is the Lewis structure for HCN, hydrogen cyanide?

Solution:

The hydrogen has 1 valence electron, the carbon 4 and the nitrogen 5 for a total of 10
The molecule is arranged exactly as it is spelled: H-C-N.



The two bonds use up four electrons
We had 10 valence electrons, used 4, so we have 6 left
The hydrogen has 2 electrons, so it is happy. The carbon has four, but needs 8, the nitrogen has two but needs 6, so we need a total of 10 electrons. We only have 6 from step 3, so we need to place two double bonds or one triple. Since the hydrogen already has the correct complement of electrons, we place a triple bond between C and N



This uses 4 electrons, leaving 2. The carbon now has 8 electrons, the nitrogen 6, so the nitrogen needs 2 more: place the last two electrons in a lone pair on nitrogen

reference website:

http://www.upei.ca/~chem342/Resources/Reviews/Molecular_Orbital_Tutorial.pdf
http://www.cerlabs.com/experiments/10534977707.pdf
http://www.kentchemistry.com/links/bonding/lewisdotstruct.htm
http://lewisdotstructure.music2.fat.philipsproducts.co.cc/lewisdotstructureforh2/