Lewis Structures 2

Introduction

In the previous note pack, we learned how to write Lewis structures fir simple compounds. We can use the four simple rules we discussed in most situations to get a good structure. However, there are some things to watch for when drawing Lewis structures. One is the phenomenon of resonance (caused by electrons being shared by more than two atoms). We will discuss resonance in this note pack.

Delocalization and bonding

When you think of a covalent bond (involving the sharing of electrons), you most likely think of a picture like this:

	The shaded area represents the likely location of the shared pair of electrons. The bonded electrons are located in between the atoms being bonded.
Illustration 1: Cartoon diagram of a covalent bond in H₂.

With many bonds, this is indeed a correct picture of how the electrons are distributed between atoms. In some cases, though, electrons are shared by more than two atoms. The shared electrons aren't necessarily located between two of the atoms all the time - they are spread out over several other atoms. These spread-out electrons are said to be delocalized. You can picture delocalized electrons this way:

	The shaded area represents the likely location of the shared pair of electrons. The bonded electrons are not located simply between two atoms - they're shared by a group. These electrons are delocalized - they are shared by more than two atoms.
Illustration 2: Cartoon diagram of delocalized electrons in O₃.

An extreme case of delocalization is the metallic bond, where electrons are shared by all atoms in a solid piece of metal.

Resonance

How can we tell when delocalization occurs in a compound? How do we show this delocalization in a Lewis structure? We'll do an example to illustrate.

Let's examine the sulfur trioxide (SO₃) molecule. First, we draw a Lewis structure the way we learned previously.

1) Count the number of valence electrons:

Atom	Number of valence electrons	Number of atoms in molecule	Total valence electrons
Sulfur (S)	6	1	6
Oxygen (O)	6	3	18
Total			24

So, SO₃ has a total of 24 valence electrons.

2) Draw the skeletal structure. Sulfur goes in the middle. Sulfur is less electronegative than oxygen.

3) Distribute electrons on the atoms - outside first, inside last. Once each oxygen is assigned three lone pairs of electrons, there are no more electrons left. Sulfur is assigned no extra electrons.

	Connect the oxygen atoms to the central sulfur atom using single bonds. We run out of electrons after satisfying all the oxygen atoms.
Illustration 3: Skeletal structure of SO₃

4) Check the structure and use multiple bonds if atoms do not have an octet of electrons (or two electrons for hydrogen) . Look in the structure in Illustration 3. The sulfur atom has a share in only six electrons. So we use a lone pair on one of the oxygen atoms to make a double bond to the sulfur atom:

	The structure checks - all atoms are satisfied. One of the oxygen atoms is bonded differently than the others. This should be an observable difference!
Illustration 4: A possible Lewis structure of SO₃

According to our rules, we're done and Illustration 4 is the Lewis structure of SO ₃ . It's odd that one oxygen would have a double bond while the two others have a single bond. There's no other difference between the oxygens - they're all bonded to the same atom and no others.

In experiments, we should be able to tell that one oxygen atom is bonded differently from the other two. However, experiments on SO₃ have shown that all three oxygen atoms in the molecule are bonded the same way. There is a degree of delocalization in the bonds. Some bonding electrons are shared by all the atoms in the molecule.

How do we represent the fact that all the sulfur-oxygen bonds are the same and still draw a "correct" Lewis structure showing all atoms with filled valence shells? We draw alternate structures with the double bond to the other two oxygen atoms. In essence, we draw three Lewis structures for SO₃ and say that the real structure is a combination of all three: These structures are called resonance structures

Illustration 5: The resonance structures of SO₃

Some notes on the structures above:

The proper way to draw the Lewis structure of SO₃ is to draw all three of the above structures separated by double-headed arrows. Simply drawing one isn't adequate!
The double-headed arrows are used to show that all three structures together are a representation of SO₃'s true structure.
Don't visualize the SO₃ molecule as having a double bond that flips around. Visualize it as having a pair of electrons shared by all the atoms.

How do we tell if a molecule has delocalized bonds in the first place - without experiment? We can tell from the Lewis structure. First, look at any double bonds in the structure you've drawn. If there's an identical atom you could have made that double bond to in your molecule, you are dealing with a molecule that probably has delocalized bonds. In the SO₃ example above, we expect delocalized bonds because the double bond could have been assigned to any of three identical oxygen atoms around the sulfur atom.

For a compound with delocalized bonding, it is not technically correct to draw only one Lewis structure. There's not actually a double bond between sulfur and one of the oxygen atoms. In reality, that pair we draw in the double bond is shared by all the atoms in the structure. Because of this, the single structure we drew in Illustration 4 isn't enough to give us the true picture of the SO₃ molecule, We need to draw all three possibilities, as we did in Illustration 5. These possibilities are called resonance structures, and the correct way to draw the structure of SO₃ is to draw all three of the possible resonance structures and connect them with double-headed arrows. This indicates that the three structures are resonance structure and that the true SO₃ molecule is really something in between all three.

Summary

We've dealt with molecules that exhibit delocalized bonding, where electrons are shared by more than one atom. We discovered that you can tell when delocalization occurs by trying to write a Lewis structure for a molecule and looking at the double or triple bonds. Since a single Lewis structure is not able to show this delocalization, we drew multiple resonance structures and wrote them with double-headed arrows between them to indicate that the real structure of the molecule was somewhere in between. Next, we will discuss the concept of formal charge and how you can use it to choose between several possible structures for a molecule.