Aromatic Hydrocarbons — Structure of Benzene

What Are Aromatic Hydrocarbons

Aromatic hydrocarbons, also called arenes, form the third major class of hydrocarbons after the saturated alkanes and the unsaturated alkenes and alkynes. The name comes from the Greek aroma, meaning pleasant smelling, because many of the earliest members of this family were isolated from fragrant natural sources. Most aromatic compounds were found to contain the benzene ring as their structural core.

Compounds that contain a benzene ring are classified as benzenoids, while those that show aromatic character without a benzene ring — instead containing some other highly unsaturated ring — are called non-benzenoids. Benzene itself, toluene, naphthalene and biphenyl are common benzenoid examples. The defining chemical feature of the benzene ring is that although it is highly unsaturated, in the great majority of its reactions this unsaturation is retained rather than destroyed.

That a highly unsaturated ring should resist addition is the central puzzle this subtopic resolves. The answer lies entirely in the structure of benzene, which is best understood by following the historical sequence in which chemists arrived at it. For a wider view of where arenes sit among the hydrocarbon families, see the companion note on the classification of hydrocarbons.

The Kekule Structure of Benzene

Benzene was isolated by Michael Faraday in 1825. Its molecular formula, $\ce{C6H6}$, signals a high degree of unsaturation that did not fit the pattern of any alkane, alkene or alkyne. Benzene forms a triozonide, which indicates the presence of three double bonds, yet it produces one and only one monosubstituted derivative — proof that all six carbon atoms and all six hydrogen atoms are equivalent.

On the basis of these observations, August Kekule (1865) proposed a cyclic arrangement of six carbon atoms joined by alternate single and double bonds, with one hydrogen atom on each carbon. This is the Kekule structure, and it captures two essential facts at once: the ring is six-membered and cyclic, and the formula $\ce{C6H6}$ is satisfied by three double bonds.

Limitations of the Kekule Structure

The first difficulty appears when two hydrogen atoms are replaced. Kekule's structure, with fixed alternating double bonds, predicts two different 1,2-disubstituted isomers. In one isomer the two substituents sit on a pair of carbons joined by a double bond; in the other they sit on a pair joined by a single bond. Experimentally, however, benzene gives only one ortho-disubstituted product.

Kekule tried to rescue the model by proposing that the double bonds oscillate — that the two arrangements rapidly interconvert so that, on average, no two ortho positions are distinguishable. Even this modification was not enough. The oscillating-bond picture still could not explain benzene's unusual stability, nor its strong preference for substitution over addition, which is the opposite of what a molecule with three genuine double bonds should do.

The repeated failures of even an oscillating Kekule structure point to a deeper truth: the double bonds in benzene are not localised between specific pairs of carbon atoms at all. Resolving this required the language of resonance, developed within Valence Bond Theory.

Resonance and the Stability of Benzene

According to Valence Bond Theory, the old idea of oscillating double bonds is replaced by resonance. Benzene is not flipping back and forth between two structures; rather, it is a single, fixed structure that is a hybrid of several contributing canonical forms. The two Kekule structures, A and B, are the two main contributing structures, and the real molecule is their resonance hybrid.

The resonance hybrid is represented by drawing a circle, or a dotted circle, inside the hexagon, as in structure C of Figure 1. This circle stands for the six π electrons that are delocalised between the six carbon atoms of the ring. Because the electrons spread evenly over all six carbons, every carbon–carbon link is identical, and the molecule has no distinct single or double bonds.

Resonance is the phenomenon by which a single molecule can be represented by two or more contributing structures, the actual molecule being the resonance hybrid of them all. None of the individual Kekule forms is correct by itself; the true structure is intermediate between them.

Resonance Energy of Benzene

The delocalisation of the six π electrons does more than equalise the bonds — it makes benzene markedly more stable than any hypothetical molecule with three isolated double bonds. This extra stability can be measured directly through heat of hydrogenation data, which provides the cleanest experimental proof of resonance.

Hydrogenating the single double bond of cyclohexene releases 119.5 kJ per mole:

$\ce{Cyclohexene + H2 ->[\text{catalyst}] Cyclohexane}\quad \Delta H = -119.5~\text{kJ mol}^{-1}$

If the three double bonds of a hypothetical "cyclohexatriene" did not interact, hydrogenating all three should release three times as much, that is $3 \times 119.5 = 358.5~\text{kJ mol}^{-1}$. The actual heat of hydrogenation of benzene, however, is only 208.2 kJ per mole:

$\ce{Benzene + 3H2 ->[\text{catalyst}] Cyclohexane}\quad \Delta H = -208.2~\text{kJ mol}^{-1}$

The shortfall is the resonance energy of benzene:

$E_{\text{resonance}} = 358.5 - 208.2 = 150.3~\text{kJ mol}^{-1}$

Benzene therefore lies far lower in energy than the localised model would predict. The delocalised π electron cloud is attracted more strongly by the carbon nuclei than electrons confined between any two carbons would be, and this is precisely why benzene is so reluctant to undergo the addition reactions typical of alkenes — addition would break up the stabilising delocalised cloud.

Molecular Orbital Picture of Benzene

Orbital overlap gives the most complete picture of benzene's structure. All six carbon atoms are sp² hybridised. Each carbon uses two of its three sp² hybrid orbitals to overlap with the sp² orbitals of its two neighbours, forming six C–C σ bonds that lie in one plane. The third sp² orbital of each carbon overlaps with the 1s orbital of a hydrogen atom, forming six C–H σ bonds, also in the plane.

This leaves one unhybridised p orbital on every carbon, oriented perpendicular to the plane of the ring. These six parallel p orbitals overlap sideways (laterally). There are two equivalent ways to pair them into three localised π bonds — overlap of $\mathrm{C_1\text{–}C_2}$, $\mathrm{C_3\text{–}C_4}$, $\mathrm{C_5\text{–}C_6}$ on one hand, or $\mathrm{C_2\text{–}C_3}$, $\mathrm{C_4\text{–}C_5}$, $\mathrm{C_6\text{–}C_1}$ on the other. These two pairings correspond exactly to the two Kekule structures.

X-ray diffraction shows that the internuclear distance between all the carbon atoms is the same, which means there is equal probability for each p orbital to overlap with both of its neighbours. The result is not two alternative sets of localised π bonds but a single delocalised system. The six π electrons spread into two continuous ring-shaped clouds — pictured as two doughnuts of electron density, one above and one below the hexagonal plane.

Equal Bond Lengths and Planarity

The two structural facts that the resonance and molecular orbital pictures predict are confirmed by X-ray diffraction. First, benzene is a planar molecule — a flat, regular hexagon. This follows directly from sp² hybridisation at every carbon, which fixes all the C–C–C bond angles at exactly 120°. Second, all six C–C bonds are of the same length, 139 pm.

Had either Kekule structure been correct on its own, benzene would have shown two distinct C–C bond lengths — short double bonds and longer single bonds. Instead, the single observed value of 139 pm sits neatly between the two extremes. The absence of any pure double bond is exactly what accounts for benzene's reluctance to undergo addition.

Nomenclature of Benzene Derivatives

Because all six hydrogen atoms in benzene are equivalent, replacing one of them gives one and only one type of monosubstituted product, such as methylbenzene (toluene). When two hydrogen atoms are replaced by two similar or different groups, three position isomers become possible, distinguished by the relative positions of the substituents.

The fact that only one ortho isomer of any 1,2-disubstituted benzene exists — rather than the two that the rigid Kekule model predicts — is itself a consequence of the equal, delocalised bonds. The naming of substituted benzenes therefore rests on the very symmetry that the structure of benzene establishes. The way different substituents direct further substitution to the ortho, meta or para positions is taken up in the note on the directive influence of functional groups, while the substitution reactions themselves are covered under electrophilic aromatic substitution.