Aromaticity

Let's compare the π molecular orbitals of benzene and hexatriene. You will find a similarity, delocalization, and a very important difference: degeneracy. Benzene has two pairs of degenerate orbitals, i.e. orbitals with the same energy. We already know that delocalization means stabilization. Degeneracy is also a synonym of added stability.

YouTube video: the acidity of cyclopentadiene

π molecular orbitals of benzene: three bonding orbitals occupied by six electrons, with two degenerate pairs — π molecular orbitals of benzene. The nodal planes are marked with **dashed lines** across the ring: the lowest-energy orbital has none; the two degenerate bonding orbitals, one; the two degenerate antibonding orbitals, two; and the highest-energy orbital, three.

π molecular orbitals of hexatriene: six non-degenerate levels, with the three bonding ones occupied — π molecular orbitals of hexatriene. The nodes are marked with **black dots** on the chain: the lowest-energy orbital has none and each orbital has one more node than the previous one, up to five in the highest-energy orbital.

Both systems have six π molecular orbitals, formed by combining the six p orbitals, and six π electrons occupying the three bonding orbitals. The energy of each orbital increases with the number of nodes (planes where the sign of the wave function changes).

In hexatriene, an open chain, the six orbitals have different energies (from 0 to 5 nodes). In benzene, being cyclic, the orbitals with 1 and with 2 nodes come in pairs of equal energy: they are degenerate orbitals. As a result, the six π electrons of benzene occupy orbitals of lower overall energy than those of hexatriene, and benzene is more stable.

Hückel set forth rules for aromaticity and antiaromaticity:

A system of π bonds is aromatic if:

fully conjugated
cyclic
planar
contains 4n + 2 π electrons

A set of 4n electrons in the same conditions is antiaromatic and highly unstable.

Hückel rule: with 4n π electrons two unpaired electrons remain in degenerate orbitals (antiaromatic); with 4n + 2 all the electrons are paired (aromatic)

All the molecules below have 4n + 2 π electrons, are cyclic conjugated systems and, as you may check, are planar. All of them are examples of aromatic compounds.



Benzene	Naphthalene	Anthracene	Phenanthrene



[18]-Annulene	Pyridine	Furan	Thiophene



Pyrrole	Indole	Pyrimidine	Purine

Resonance forms of aromatic hydrocarbons

Benzene

Kekulé almost hit the bull's eye! Benzene is explained today as a resonance hybrid between two equivalent structures, not as an equilibrium. Look at the double-headed arrow!

Geometry of benzene: C–C bonds of 1.39 Å, C–H bonds of 1.09 Å and 120° angles, compared with the double (1.34 Å) and single (1.54 Å) bonds

The C–C distance in benzene, 1.39 Å, lies between those of a single bond (1.54 Å) and a double bond (1.34 Å). The identical contribution of the two resonance forms to the hybrid explains why all the C–C bonds in benzene have the same length.

Naphthalene

The additional stability gained by naphthalene as compared to benzene is not double: it is a little lower. This is because the second and third resonance forms do not contain two true benzene rings: they are just like a cyclohexadiene fused to a benzene.

Bond lengths of naphthalene: C(1)–C(2), 1.37 Å; C(2)–C(3), 1.42 Å; C(1)–C(9), 1.40 Å; central bond, 1.39 Å; 121° angles

Unlike in benzene, the C–C distances in naphthalene are not all equal. The C(1)–C(2) bond, 1.37 Å, is the shortest because it has double bond character in two of the three resonance forms. The rest are longer because they have double bond character in only one of the three resonance forms.

Anthracene

Anthracene is not three times as stable as benzene: it is actually only a bit more stable than naphthalene. None of its four resonance forms contains three benzene rings, don't you agree?

Diels-Alder reaction of anthracene with an alkene at the central ring

Anthracene can sacrifice a small part of its aromaticity and act as the diene in a Diels-Alder reaction. The middle ring reacts because this leads to an adduct that keeps the full aromaticity of two benzene rings; no other possibility would retain as much aromaticity.

Phenanthrene

Phenanthrene is somewhat more aromatic than anthracene because its non-linear geometry allows more resonance forms bearing complete benzene rings to be written.

Partial hydrogenation of phenanthrene with H₂, copper chromium oxide and pressure: 9,10-dihydrophenanthrene

Yet, under relatively harsh conditions (H₂, copper chromium oxide, pressure), phenanthrene can be partially hydrogenated to 9,10-dihydrophenanthrene, giving away the minimum possible aromaticity. The 9,10 bond is the one hydrogenated because that is the only way of keeping the full aromaticity of two benzene rings.

The cyclopentadienyl anion

The sp³ carbon of cyclopentadiene is much more acidic than expected (as acidic as an alcohol!). As usual, this can be explained by looking at the conjugate base, the cyclopentadienyl anion: don't you realize that it is aromatic and thus exceptionally stable?

Compared acidity: propane (pKa = 50), propene (pKa = 40) and cyclopentadiene (pKa = 16), whose anion, with 4n + 2 π electrons, is aromatic

Resonance forms of the cyclopentadienyl anion: the negative charge is spread over the five carbons and all the forms are aromatic

Ferrocene: "sandwich" complex of Fe(II) between two aromatic cyclopentadienyl anions

The cycloheptatrienyl (tropylium) cation

1,3,5-Cycloheptatriene with Br₂ and heat loses a hydride (as HBr) to give tropylium bromide

1,3,5-Cycloheptatriene can lose a hydride, leading to a very stable cation. Why is that? Because the tropylium cation is aromatic.

Resonance forms of the tropylium cation: the positive charge is spread over the seven carbons

The tropylium cation is a planar system of seven laterally overlapping p orbitals in which 6 π electrons reside (4n + 2, with n = 1). Its seven equivalent resonance forms indicate great stability.