Resonance
Benzene and its derivatives belong into this group, some of whose members occurring in plants are characterized by a strange, 'aromatic' smell. The benzene molecule has a ring-shaped structure with six C-H-groups that are linked alternately with C-C single and C=C double bonds. The structure should thus be written
The electron theory distinguishes between two types of bonds in such a ring system, the sigma-bonds and the pi-bonds that are arranged vertically to each other and are conjugated so that the double bonds of such a system cannot be localized directly. Such a system is also called a mesomere system. There exist consequently six equal 'aromatic' bonds
Other Explanation :
we'll find that resonance is very important in understanding both the structure and the reactions of aromatic compounds. First, let's take a look at the structural representations which distinguish aromatic compounds from those that aren't aromatic.
The most commonly encountered aromatic compound is benzene. The usual structural representation for benzene is a six carbon ring (represented by a hexagon) which includes three double bonds. Each of the carbons represented by a corner is also bonded to one other atom. In benzene itself, these atoms are hydrogens. The double bonds are separated by single bonds so we recognize the arrangement as involving conjugated double bonds. An alternative symbol uses a circle inside the hexagon to represent the six pi electrons. Each of these symbols has good and bad features. We'll use the three double bond symbol simply because it is also routinely used in the text.
Keep in mind that if the hexagon contains neither the three double bonds nor the circle, the compound is not aromatic. It is simply cyclohexane and there are two hydrogens on each carbon atom. This is easy to mistake when hurrying, so be careful when you are intepreting any structural formulas which include hexagons.
The structure with three double bonds was proposed by Kekule as an attempt to explain how a molecule whose molecular formula was C
6
H
6
could be built out of carbons which make four bonds. The ring and the three double bonds fit the molecular formula, but the structure doesn't explain the chemical behavior of benzene at all well. Each of the double bonds would be expected to show the characteristic behavior of an alkene and undergo addition reactions, but this is not how benzene reacts. In particular, we would expect a carbon-carbon double bond to react quickly with bromine to make a dibromo compound. This is what alkenes do very readily, and in fact it is a useful test for alkenes in the laboratory. Benzene does not react with bromine unless a very bright light or a strong catalyst is used, and then the reaction is not an addition reaction. We conclude that there is something quite unusual about the double bonds in benzene. Kekule (thinking about this problem before bonds were understood as pairs of electrons) suggested that there are two forms of benzene which differ in the locations of the double bonds. His idea was that these were in rapid equilibrium, so rapid that there was never a fixed location for the double bond. One could say that an approaching bromine molecule could not "find" a double bond to react with.
There were several other structures proposed for benzene, but a much more satisfactory approach became possible when we began to understand that covalent bonds consist of pairs of electrons shared between atoms. The difference between the two structures Kekule envisioned (called Kekule structures) is only the difference between the locations of three pairs of electrons. This is exactly the type of situation where resonance must be involved. The hybrid or "average" of the two Kekule structures has one sigma bond and one-half of a pi bond between each two carbon atoms. Thus each carbon is joined to each of its neighbors by a one-and-half bond. Each bond in the benzene ring has the same number of electrons and is the same length. This picture is in complete accord with experiments which show that all carbon-carbon bonds in benzene are the same length, with no hint of shorter (double) or longer (single) bonds. It also helps explain why benzene does not undergo addition reactions: there are no simple pi bonds.
Recall that resonance has another important feature: when resonance is involved, the real structure is more stable than we would expect from any of the structures we write using the one line = two electrons symbolism. This extra lowering of energy, which for benzene is about one-third as much as making a typical covalent bond, is quite important in the reactions of benzene and other aromatic compounds. As we will see, reactions of the benzene ring almost always result in products which in which the benzene ring persists -- an outcome of its stability.
Benzene & Resonance
Since its discovery in 1825 by
Michael Faraday,the properties of
benzene have been studied more than any other organic compound.
August Kekule had derived the structure of many carbon containing compounds but
benzene confused him. With a molecular weight of 78, the formula for benzene should be C6H6. Kekule put six carbon atoms in a row but had trouble placing the six hydrogen atoms. No arrangement really worked.
Kekule was to say later that he must have dozed off at this point. In his dream the black balls of carbon turned into black imps with forked tails that began racing around the room and would soon be upsetting the apparatus of the laboratory. He was ready to run the rascals out. Then, almost suddenly, the confusion died away as each imp grabbed the tail of the one ahead of him, the six forming a whirling circle. One hand of each imp held a tail, the other a white handkerchief--and they waved to him as the group whirled by. He said that he came awake with a start, realizing that the imps were acting out the formula for benzene. As his hand grabbed the sketching pencil, the imps were back to black balls again and the handkerchiefs had changed to hydrogen atoms. How simple the arrangement turned out to be. "The carbon atoms of benzene form a ring." In 1858, August Kekule proposed a rapid oscillation (I & II) between the 3 C=C and the 3 alternating C-C of hexagonal benzene (C6H6).
However, Kekule's representation of benzene had serious flaws.
Flaws with Kekule Model
1. Heat of Hydrogenation is the heat liberated when hydrogen adds to C=C. Benzene is compared to cyclohexene because both form cyclohexane upon complete hydrogenation. Since benzene supposedly contains 3 C=C, it should release three times the heat of hydrogenation reported for cyclohexene with only 1 C=C. Heats of hydrogenation for cyclohexene and benzene are reported as: 
Since cyclohexane releases 28.6 kcal/mol, benzene is expected to release (3x28.6) or 85.8 kcal/mol.
Why does benzene possess 36 kcal/mol less energy than predicted?
2. Addition versus Substitution: Compounds with C=C (alkenes) typically undergo addition reactions such as bromination. Shown below are reactions of ethene/bromine and benzene/bromine:
Why does benzene with 3 C=C resist addition and prefer substitution?
3. X-Ray studies reveal the following bond lengths: 1.53A for typical C-C, 1.34A for typical C=C, and 1.39A for the 6 carbon-carbon bonds in benzene.
Why does benzene possess 6 bonds intermediate between C=C and C-C?
Resonance
To help explain molecules like benzene, Linus Pauling proposed resonance theory in 1931 (Pauling also gave us hybrid orbitals, electronegativity, and valence bond theory). As a result of Pauling's resonance, benzene is viewed as a hybrid of III& IVand represented byV. 
The six carbons are arranged in a hexagon with one hydrogen atom attached to each carbon. The 12 atoms of benzene are planar with carbon in the sp2 hybrid state and 120° bond angles. Because the six bond lengths are equivalent, Kekule's rapid equilibrium must be ruled out. To describe benzene we need to use contributing structures III& IVor one structure with a circle in the middle (V). Structures III &IV do not exist! Whatever benzene might be, it displays characteristics of its contributing structures. For an analogy, consider crossing a bloodhound with a greyhound: Although the offspring is unique, it displays characteristics (smell & speed) of its parents--it does not flip flop between a greyhound one instant and a bloodhound the next.
Why does benzene possess 36 kcal/mol less energy than predicted?
This missing energy, called resonance energy, is attributed to overlap of p-orbitals. Structures VI and VII show pi bonding for the contributing structures while VIII illustrates delocalization of pi electrons over the 6 carbon atoms. Benzene can be represented as IX using molecular models with p-orbitals. The circle in the middle of V is an abbreviated way to represent the delocalization of the 6 pi electrons. Resonance energy is the difference in energies between III (or IV) and V (V has lower energy). Due to resonance, benzene is 36 kcal/mol more stable than calculations would predict.
Why does benzene resist addition and prefer substitution?
If benzene were to undergo addition,resonance would be disrupted and this would render the structure less stable. Therefore benzene prefers to substitute (Br for H) and maintain resonance.
Why does benzene possess 6 bonds intermediate between C=C and C-C?
According to resonance, the bonds are not C-C or C=C but a hybrid of the two. X-ray studies confirms this with the intermediate bond lengths.
Moderation Team
![[Post New]](/templates/default/images/icon_minipost_new.gif){kind=icon}
![[Up]](/templates/default/images/icon_up.gif "[Up]"){kind=icon}
41
