Resonance
The Lewis structure notation is useful because it conveys the essential qualitative information about properties of chemical compounds. The main features of the chemical properties of the groups that make up organic molecules,
and so forth, are to a first approximation constant from molecule to molecule, and one can therefore tell immediately from the Lewis structure of a substance that one has never encountered before roughly what the chemical properties will be.
There is a class of structures, however, for which the properties are not those expected from the Lewis structure. A familiar example is benzene, for which the heat of hydrogenation (Equation 1.1) is less exothermic by about 37 kcal/mole than one would have expected from Lewis structure 1 on the basis of the measured
heat of hydrogenation of ethylene. The thermochemical properties of various
types of bonds are in most instances transferable with good accuracy from molecule to molecule; a discrepancy of this magnitude therefore requires a fundamental modification of the bonding model. The difficulty with model 1 for benzene is that there is another Lewis structure, 2, which is identical to 1 except for the placement of the double bonds.
Whenever there are 2 alternative structures for a single compound, and any one of the strucutre becomes an inaccurate representation for the molecular structure. The actual structure of the molecule is actually a hybrid of these 2 strucutres. It is like a “superposition” of these 2 strucutres. The superposition of two more more Lewis structures into a composite picture of the compound is called resonance.
This terminology is well established, but unfortunate, because the term resonance when applied to a pair of pictures tends to convey the idea of a changing back and forth with time. It is therefore difficult to avoid the pitfall of thinking of the benzene molecule as a structure with three conventional double bonds, of the ethylene type, jumping rapidly back and forth from one location to another. This idea is incorrect. The electrons in the molecule move in a field of force created by the six carbon and six hydrogen nuclei arranged around a regular hexagon.
Each of the six sides of the hexagon is entirely equivalent to the otehr side, which is why electrons should, even momentarily, seek out three sides and make them different from the other three, as the two alternative pictures 3 seem to imply that they do. The symmetry of the ring of nuclei (4) is called a sixfold symmetry because rotating the picture by one-sixth of a circle will give the identical picture again. This sixfold symmetry must be reflected in the electron distribution.
A less misleading picture would be the one above, in which the circle in the middle of the ring implies a distribution of the six double bond electrons of the same symmetry as the arrangement of nuclei. We shall nevertheless usually continue to use the notation 3, as it has certain advantages for thinking about reactions.
The most important features of the structures for which the resonance is needed are, first, that the molecule is more stable (of lower energy) than one would expect from looking at one of the individual strucutres. and second that the actual distribution of the electrons in the molecule is different form what one would expect from either of the resonance strucutres. Since the composite picture shows that certain electrons are free to move to a larger area, of the molecule than a single one of the structure implies, resonance is often referred to as delocalization. We shall have more to say about delocalization later in connection with molecular orbitals.
While the benezene ring is the most fimiliar example of the necessity for modifying the Lewis strucutre language by addition of resonance concept, there are many others. The carboxylic acids, for example are much stronger acids than the alcohols; this difference must be largely due to greater stability of the carboxylate ion over the alkoxide iiion. It is the possiblity of writing two equivalent Lewis structures for the carboxylate ions that alertts us to this difference.
Another example is the allylic system. The ally1 cation (8), anion (9), and
radical (10), are all more stable than their saturated counterparts. Again, there is for each an alternativestructure :
In all the examples we have considered so far, the alternative structures have been equivalent. This will not always be the case, as the following examples illustrate :
Whenever there are nonequivalent strucutres, each will contribute to the composite picture to a different extent. The structure would represent the most stable (lowest-energy) molecule, where such a molecule actually to exist, contributes the most to composite, and the others successively less as they represent the higher energy molecules.
It is because the lowest-energy structures are most important that we specifiedin the rules for writing Lewis structures that the number of bonds should be maximum and the valence-shell occupancy not less than 8 whenever possible. Structures that violate these stipulations, such as 11 and 12, represent high-energy forms and hence do not contribute significantly to the structural pictures, which
The followinp; rules are useful in using resonance notatinn:
- All nuclei must be in the same location in every structure. Structures with nuclei in different locations, for example 15 and 16, are chemically distinct substances, and interconversions between them are actual chemical changes, always designated by
.
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Structures with fewer bonds or with greater seperation of formal charge are less stable than those with more bonds or less charge seperation. thus 11 and 12 are higer energy respectively than 12 and 14.
- When 2 structures with formal charge have the same number of bonds and approximately same charge seperation, the structure with charge on the more electronegative atom will usually be somewhat in the lower energy state, but the difference will ordinarily be small enough that both structures can be included in the composite picture.
- All four groups attached to a pair of atoms joined by a double bond in any structure must lie in the same plane. For example, structure 18b cannot contribute because the bridged ring prevents the carbon 6 and 7 from lying in the same place as carbon 3, and the hydrogen on carbon 2. The impossiblity of strucures with double bonds at bridgeheads of small bridged rings is known as Bredt’s rule. Double bonds can occur at bridgehead if the rings are sufficiently large.
