In 1834, Eilhardt Mitscherlich conducted vapor density measurements on benzene. Based on data from these experiments, he determined the molecular formula of benzene to be C 6H 6. This formula suggested that the benzene molecule should possess four modes of unsaturation because the saturated alkane with six carbon atoms would have a formula of C 6H 14. These unsaturations could exist as double bonds, a ring formation, or a combination of both.
In 1866, August Kekulé used the principles of structural theory to postulate a structure for the benzene molecule. Kekulé based his postulation on the following premises:
- The molecular formula for benzene is C 6H 6.
- All the carbons have four bonds as predicted by structural theory.
- All the hydrogens are equivalent, meaning they are indistinguishable from each other.
Based on these assumptions, Kekulé postulated a structure that had six carbons forming a ring structure. The remaining three modes of unsaturation were the result of three double bonds alternating with three single bonds. This arrangement allowed all the carbon atoms to have four bonds as required by structural theory.
Scientists soon realized that if Kekulé's structure were correct, substituting substituent groups for hydrogens on the 1,2 positions would lead to a different compound than substitution on the 1,6 positions.
Because no such isomers could be produced experimentally, Kekulé was forced to modify his proposed structure. Kekulé theorized that two structures existed that differed only in the location of the double bonds. These two structures rapidly interconverted to each other by bond movement.
Although Kekulé's structure accounted for the modes of unsaturation in benzene, it did not account for benzene's reactivity.
Modern instrumental studies confirm earlier experimental data that all the bonds in benzene are of equal length, approximately 1.40 pm. (A picometer equals 1 × 10 −12 meter.) This bond length falls exactly halfway between the length of a carbon‐carbon single bond (1.46 pm) and a carbon‐carbon double bond (1.34 pm). In addition, these studies confirm that all bond angles are equal (120°) and that the benzene molecule has a planar (flat) structure.
Modern descriptions of the benzene structure combine resonance theory with molecular orbital theory.
Resonance theory postulates that when more than one structure can be drawn for the same molecule, none of the drawn structures is the correct structure. The true structure is a hybrid of all the drawn structures and is more stable than any of them. The greater the number of structures that can be drawn for a molecule, the more stable the hybrid structure will be. The difference between the calculated energy for a drawn structure and the actual energy of the hybrid structure is called the resonance energy. The greater the resonance energy of a compound, the more stable the compound.
The two Kekulé structures that can be drawn for the benzene molecule are actually two resonance structures.
The hybrid of these structures would be drawn as
where the circle represents the movement of the electrons throughout the entire molecule. This delocalization of π electrons (electrons found in π molecular orbitals) is also found in conjugated diene systems. Like benzene, the conjugated diene systems show increased stability.
Because of resonance, the benzene molecule is more stable than its 1,3,5‐cyclohexatriene structure suggests. This extra stability (36 kcal/mole) is referred to as its resonance energy.
Because experimental data shows that the benzene molecule is planar, that all carbon atoms bond to three other atoms, and that all bond angles are 120°, the benzene molecule must possess sp 2 hybridization. With sp 2 hybridization, each carbon atom has an unhybridized atomic p orbital associated with it. The overlap of the sp 2 hybrid orbitals would create the σ bonds that hold the ring together, while the side‐to‐side overlap of the atomic p orbitals can occur in both directions, leading to complete delocalization in the π system. This complete delocalization adds great stability to the molecule. Figure 1 illustrates this idea.
Molecular orbital theory predicts that overlapping six atomic p orbitals will lead to the generation of six π molecular orbitals. Three of these π molecular orbitals will be bonding orbitals, while the other three will be antibonding orbitals, as shown in Figure 2.
The three low‐energy orbitals, denoted π 1, π 2, and π 3, are bonding combinations, and the three high‐energy orbitals, denoted π 4 *, π 5 *, and π 6 *, are antibonding orbitals. Two of the bonding orbitals (π 2 and π 3) have the same energy, as do the antibonding orbitals π 4 and π 5. Such orbitals are said to be degenerate.
Because the electrons are all located in bonding orbitals, the molecule is very stable. Additional stability occurs because all the bonding orbitals are filled and all the π electrons have paired spins. Molecules that possess all these characteristics are said to have a closed bond shell of delocalized π electrons. Molecules such as benzene that possess a closed bond shell of delocalized π electrons are extremely stable and show great resonance energies.