Observation 3: Molecular Structures with more than one Bond- Bond-ing Structure

Earlier in this concept study, we learned that dierent molecules can be formed from the same set of atoms.

Isomers are molecules with the same molecular formula but dierent molecular structures, and are therefore distinctly dierent compounds. Rearranging the atoms in the molecule creates a new compound with new physical and chemical properties.

What if, instead of rearranging the atoms in the molecule, we rearrange the electrons? This would possibly change where double or triple bonds are located and whether there are lone pairs of electrons or not. Does this also give rise to new compounds and therefore new isomers?

We rst look at benzene, a compound with the molecular formula C₆H₆. For six carbon atoms, there are not very many hydrogen atoms. Compare benzene to hexane, which has the molecular formula C₆H₁₄. This means that there must be several double or triple carbon-carbon bonds in benzene. Experiments reveal to us two facts about the molecular structure of benzene. First, the six carbon atoms are arranged in a ring, not a chain, and each carbon atom is bonded to a single hydrogen atom. Second, the bonds between the carbon atoms in the ring all have the same length as one another. The second observation tells us that, somehow, all of the bonds in benzene are identical to each other.

The correct molecular structure which explained these observations was a puzzle for chemists. A structure in which the six carbon atoms are arranged in a ring would be:

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78 CHAPTER 8. MOLECULAR STRUCTURES

Figure 8.7

This structure cannot be the right structure though. Although it satises the octet rule and the valences of the carbon atoms, it does not correctly explain why the bonds are all the same length. This structure would predict three shorter bonds and three longer bonds.

A clue to the right structure is found from the value of the bond length, 139 pm. This length is between the typical length of a single bond, 153 pm, and the typical length of a double bond, 134 pm. This is a confusing clue: it suggests that the bonds in benzene are neither single bonds nor double bonds. We clearly need to expand on our model of Lewis structures.

If we look at the molecular structure proposed above, we can see that we made an arbitrary choice of where to put the double bonds. To satisfy the valences of the carbon atoms, we need to alternate the double bonds and single bonds, but we could have chosen the other three C-C bonds to be double:

Figure 8.8

79 The dierence between these two structures is not the arrangement of the atoms, but rather the arrange-ment of the electrons. Are these isomers? If they are, then there are two benzene compounds. Experiarrange-mentally, we nd only one. So these two structures must not be isomers. Rearranging the electrons in a molecular structure does not produce new compounds.

What then are we to do with the two structures drawn above? There is no reason why one would be preferred over the other, so both must be equally correct for benzene. But neither of them alone is correct because each of them predicts bond lengths which do not match experimental observations.

One possible answer is that there is a single molecular structure for benzene which combines both struc-tures together. This means that the double bonds in benzene are not xed in one set of locations or the other. Rather, the bond lengths tell us that the double bonds are spread out around the six carbon ring uniformly. The language that chemists use to describe this phenomenon is that the correct structure of benzene is a hybrid of the two structures drawn above. The word hybrid refers to something that contains properties of more than one element. Benzene has a single molecular structure that combines the properties of both of the above structures at the same time.

How does this explain the experimental bond lengths? If we look at each C-C bond and combine the properties of the two structures, each bond has the properties of a single bond and of a double bond. This means that the bond length would be somewhere between a single bond and a double bond, and this is just what is found experimentally.

A Lewis structure which represents this hybrid is:

Figure 8.9

The drawing on the left uses dotted lines to represent the double bonds which are neither here nor there but are rather delocalized around the six carbon ring. The drawing on the right is another way to represent this idea, but the solid ring represented the double bonds is somewhat easier to draw and therefore more commonly used by chemists. Chemists refer to the delocalization of the electrons as resonance, and the structure above is often called a resonance hybrid.

This concept applies to a number of molecules. A good example is ozone, O₃. Experiments show that the two O-O bond lengths are equal, 128 pm. We can compare this to the double bond length in O₂, which is 121 pm, and to the single bond length in hydrogen peroxide, H-O-O-H, which is 147 pm. From our model, we might conclude that the O-O bonds in O3 are partially double and partially single, just like in benzene.

How would our model account for this?

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80 CHAPTER 8. MOLECULAR STRUCTURES We can draw two equivalent Lewis structures for ozone:

Figure 8.10

Based on our observations and our model, we can conclude that the correct molecular structure of ozone is a resonance hybrid of these two structures in which the double bond is delocalized over both O-O bonds.