Lewis Structures are very similar to electron dot diagrams except for the fact that shared electrons between atoms are shown as lines. Lone pairs of electrons themselves are usually represented by dots around the atoms. These diagrams are helpful because they allow us to predict the shape of the molecule and see the positions of each atom relative to the others.
In 1916, Gilbert Lewis Newton introduced a way to show the bonding of atoms in a molecule though diagrams. Creating Lewis Structures is rather simple and only requires a small series of steps and little bit of math.
Basic Steps for Simple Structures
Lewis Structure are very simple to construct and only take a few short steps. Make sure to have a paper and pencil handy because you will be drawing these molecules, as well as doing some simple math.
For this module, we will use Boron trifluoride as our example.
The final structure can be seen in a diagram below. Because there are three bonded atoms to a central atom with no lone pairs on it, be can also conclude that BF3 has a trigonal planar geometry with sp2 hybridization.
BF3 is only one of the many simple Lewis Structures. There are others that require a little bit more time in creating because they contain a little bit more criteria to complete it. Some of the more complex structures may contain charges as well as double or triple bonds. There are also many molecules that break the octet rule and can form a total of 5 or 6 bonds to the central atom.
When drawing Lewis Structures you need to check the structure's formal charge. Unless a charge is preferred you want the overall charge of the molecule to be 0.
Formal Charge = # of valence electrons - # of unpaired electrons - (1/2) # of paired electrons.
So the formal charges are:
FCBoron = 3-0-(1/2)6 = 0
FCFluorine = 7-6-(1/2)2 = 0
FCTotal = FCBoron + FCFluorine = 0 + 0 = 0
If a formal charge is preferred, then the final answer should be in brackets with the charge outside of the bracket.
Lewis structures of Charged Molecules
In the case of Triiodide (I3-), you notice that there is a negative charge. That means that instead of the 21 valence electrons (7 for each I), you add one more electron to make it 22. You will then create the structure though the steps that were given above. The only additional step required is to put your final structure in brackets with the charge of the molecule outside of the brackets. Triiodide can be seen below.
Multiple Bonds in Lewis Strutures
In the case of CO2, when it comes time to create a lewis structure, it seems as though you don not have enough electrons to complete the octets of all the atoms. When you run into this problem, it usually means that the molecule will contain a double or triple bond.
There are 16 valence electrons in a CO2 molecule. If you go about the steps that were given to you above, you will notice that the octets of the oxygen atoms have been filled, but the carbon still needs 4 more electrons. What you need to do it move one pair of electrons from each oxygen atom and turn it into a double bond on each side of the carbon. The final structure for CO2 can be seen below.
Complete the following problems to to evaluate your understanding of this module:
Draw the Lewis structure for the following molecules:
Find the Formal Charge on all atoms of:
1. 2. 3.
4. FCN = 5-0-(0.5)8 = +1
FCO = 6-4-(0.5)4 = 0
5. FCC =4-0-(0.5)8 = 0
FCH = 1-0-(0.5)2 = 0
6. FCS = 6-0-(0.5)12 = 0
FCF = 7-6-(0.5)2 = 0
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