Aromatic Hydrocarbons and Benzene Structure

Aromatic hydrocarbons are cyclic, planar compounds with a ring of delocalized pi electrons that make them unusually stable, and the simplest and most important example is benzene, C₆H₆. In brief: all six carbon–carbon bonds in benzene are equal (139 pm, between a single and a double bond), and the ring is about 150 kJ/mol more stable than a hypothetical cyclohexatriene with three separate double bonds.

Benzene puzzled chemists for decades: it has the formula of a highly unsaturated compound but does not react like an alkene. The explanation, delocalization of electrons, is one of the key concepts of organic chemistry. This article describes the structure of benzene, the evidence for it, the rule that defines aromatic compounds and the typical reactions of aromatic rings.

What is the structure of benzene?

Benzene consists of six carbon atoms in a planar hexagonal ring, each bonded to one hydrogen atom. Every carbon is sp² hybridized, so all C–C–C and H–C–C bond angles are 120°. Each carbon has one p orbital perpendicular to the plane of the ring. The six p orbitals overlap sideways to form a continuous cloud of pi electrons above and below the ring, which holds six electrons that are shared by all six atoms.

This delocalization is shown in the Lewis structures as two equivalent resonance forms (the Kekulé structures), with alternating single and double bonds. The real molecule is a hybrid of the two and is often drawn as a hexagon with a circle inside it.

What is the evidence for delocalization?

• Bond lengths: all C–C bonds in benzene are 139 pm, between a single bond (154 pm) and a double bond (134 pm). A cyclohexatriene with localized bonds would have three long and three short bonds.
• Stability: benzene does not decolorize bromine water or undergo addition reactions easily.
• Isomers: there is only one 1,2-disubstituted benzene. With fixed double bonds there would be two isomers, one with a single bond and one with a double bond between the substituted carbons.
• Hydrogenation enthalpy: benzene releases much less energy on hydrogenation than expected.

Hydrogenation enthalpies

Three times the value for cyclohexene is −360 kJ/mol. The measured value for benzene is −208 kJ/mol, which is 152 kJ/mol less exothermic. This difference is the delocalization (resonance) energy: benzene is that much more stable than it would be if the electrons were localized.

What makes a compound aromatic?

Erich Hückel showed in 1931 that a planar, cyclic molecule with a continuous ring of p orbitals is aromatic if it has 4n + 2 pi electrons, where n is zero or a positive integer. Benzene has 6 pi electrons (n = 1). The rule explains why other compounds are aromatic: naphthalene has 10 pi electrons (n = 2), the cyclopentadienyl anion has 6, and pyridine has 6. A planar ring with 4n pi electrons, such as cyclobutadiene (4 electrons), is antiaromatic and unstable.

1. The molecule must be cyclic.
2. Every atom of the ring must have a p orbital (the ring is conjugated).
3. The molecule must be planar.
4. It must have 4n + 2 pi electrons.

Common aromatic hydrocarbons

Naphthalene is the substance that was used in mothballs. Compounds with several fused rings are called polycyclic aromatic hydrocarbons (PAHs); they are formed by the incomplete combustion of organic matter, and some, such as benzo[a]pyrene in tobacco smoke, are carcinogens.

How do aromatic compounds react?

Benzene prefers substitution to addition, since addition would destroy the stable delocalized ring. Substitution keeps the ring intact. The typical reaction is electrophilic aromatic substitution, in which an electrophile replaces a hydrogen atom:

• Nitration: benzene + HNO₃ (with H₂SO₄ as a catalyst, at about 50 °C) → nitrobenzene + H₂O.
• Halogenation: benzene + Br₂ (with FeBr₃) → bromobenzene + HBr.
• Sulfonation: benzene + H₂SO₄ (hot, concentrated) → benzenesulfonic acid.
• Friedel–Crafts alkylation and acylation: benzene + RCl (with AlCl₃) → alkylbenzene.

Substituents already on the ring direct the position of the next group: –OH and –CH₃ direct to the ortho and para positions, and –NO₂ to the meta position.

How was the structure of benzene found?

Michael Faraday isolated benzene in 1825 from the oily residue of illuminating gas. Its formula, C₆H₆, was established by Eilhard Mitscherlich in 1834. August Kekulé proposed the ring structure with alternating double bonds in 1865. According to a story told by Kekulé himself 25 years later, the idea came to him in a daydream of a snake biting its own tail, although historians question how accurate this account is. Quantum mechanics later explained the stability: Hückel’s work of 1931 and Linus Pauling’s resonance theory.

Is benzene dangerous?

Yes. Benzene is a known human carcinogen linked to leukemia, and its vapors are toxic. For this reason, benzene has been largely replaced as a laboratory solvent by toluene, which is considerably less toxic. The benzene content of gasoline is regulated in many countries.

Which mistakes are common?

• Drawing localized double bonds as if they were real: the Kekulé structures are resonance forms, not rapidly interconverting isomers.
• Expecting benzene to behave like an alkene: it undergoes substitution and not addition.
• Mixing up aromatic with “having an odor”: the name is historical; aromaticity is about electrons.
• Forgetting the 4n + 2 rule: cyclooctatetraene (8 electrons) is not aromatic and not planar.
• Counting the wrong number of pi electrons: only electrons in the conjugated ring count.