Benzene Structure

How is benzene structured?

Benzene has a unique structure. There are some important structural aspects you need to know about, such as its formula, bond lengths, and electron arrangement.

Structural and displayed formulae

As we mentioned above, benzene has the molecular formula \(C_6H_6\) . It forms a hexagonal molecule, which we often represent as a hexagon with a circle inside.

Fig. 2 - The conventional representation of a benzene molecule

Bond length

We explored some potential structures for benzene in Aromatic Chemistry, each containing three C=C double bonds. But in actual fact, we know that benzene doesn’t contain any double bonds at all. Instead, all of its carbon-carbon bonds are identical intermediates - halfway between a single and a double bond in length. We’ll explore why in just a second.

Bond angle

Each carbon atom in benzene is bonded to two other carbon atoms and just one hydrogen atom, and the bond angle between each bond is 120°. This makes benzene a trigonal planar molecule.

Fig. 3 - Benzene has a bond angle of 120°

However, we know that carbon has four valence electrons. Only three electrons have formed bonds - what has happened to the last one? To answer that, let’s look at electron orbitals and something called the delocalised pi system.

The delocalised pi system

Carbon’s fourth outer-shell electron is found in a pi orbital, whereas the bonded three are found in sigma orbitals. Sigma orbitals stretch between atoms whilst pi orbitals extend above and below the atom. In benzene, all the pi orbitals of the carbon atoms overlap, producing a connected region that stretches above and below the molecule.

Fig. 4 - Benzene’s six pi orbitals overlap and delocalise

The electrons can move anywhere within this overlapping region. We say that they are delocalised. The overall structure is called the delocalised pi system.

Because of these delocalised electrons, benzene doesn’t need to form any double bonds. As we mentioned above, all of its C-C bonds are instead identical intermediates.

To summarise, benzene has the following structure:

• It has a planar hexagonal shape.
• Each carbon atom is bonded to two other carbon atoms and one hydrogen atom using three of its valence electrons.
• Each of its C-C bonds is an intermediate, halfway between a single and double bond in length.
• The angle between bonds is 120°.
• Each carbon atom’s fourth valence electron is delocalised in a region above and below the molecule.

Kekulé's structure of benzene, and other theories

In 1865, the German organic chemist Friedrich August Kekulé published a paper on the structure of benzene. This was a mystery that had puzzled scientists for years. He claimed that he dreamt about a snake biting its own tail, which led him to conclude benzene’s cyclic nature. Kekulé proposed that benzene contained alternating C-C single and C=C double bonds, as shown below. This molecule is systematically known as cyclohexa-1,3,5-triene.

Fig. 5 - Kekulé’s proposed structure for benzene: cyclohexa-1,3,5-triene

However, there were a number of pieces of evidence that didn’t quite support this structure.

Electrophilic addition reactions

A common test for alkenes is to mix them with bromine water. If a C=C double bond is present, the water will become decolourised as the bromine atoms join on to the hydrocarbon in an electrophilic addition reaction. Kekulé’s structure for benzene contains three C=C double bonds, so we’d expect it to react in this way. However, it doesn’t - when mixed with bromine water, the solution remains red-brown. This suggests that benzene doesn’t have any double bonds.

Enthalpy of hydrogenation

Reactions that add hydrogen to a molecule are known as hydrogenation reactions. Let’s look at cyclohexene, shown below. It has an enthalpy of hydrogenation of $-120\mathrm{kJ}{\mathrm{mol}}^{-1}$, meaning that $120\mathrm{kJ}$ of energy are released when two hydrogen atoms add on to its single double bond. This produces cyclohexane.

Fig. 6 - The enthalpy of hydrogenation of cyclohexene

If we now look at Kekulé’s structure ,we can see that it has three C=C double bonds. We’d therefore expect it to have an enthalpy of hydrogenation three times as great as cyclohexene, which has just one double bond:

$3\mathrm{x}-120=-360\mathrm{kJ}{\mathrm{mol}}^{-1}$

However, experiments show that benzene’s enthalpy of hydrogenation is only $-208\mathrm{kJ}{\mathrm{mol}}^{-1}$. It is $152\mathrm{kJ}{\mathrm{mol}}^{-1}$ more stable than expected. This is known as benzene’s resonance energy. We now know that this stability is due to benzene’s ring of delocalised electrons, which stabilises the molecule by spreading the electrons’ negative charges over a larger area.

Bond lengths

X-ray diffraction is a type of technique using X-rays to work out the structure of molecules. Scientists used it in 1981 to get an image of benzene. In Aromatic Chemistry, we learnt that C-C single bonds are longer than C=C double bonds. This would give Kekulé’s cyclohexa-1,3,5-ene a distorted shape.

Fig. 7 - The distorted shape of cyclohexa-1,3,5-triene

However, the image showed that benzene was in fact a regular hexagon. This meant that all of its bonds were equal length. Furthermore, scientists measured the length of these bonds and found them to be halfway between a single and a double bond in length - suggesting that they were neither one nor the other, but something different instead.

Fig. 8 - A table showing the lengths of different carbon-carbon bonds

Isomeric products

Let’s look at one final piece of evidence against Kekulé’s predicted structure of benzene. Take two of benzene’s adjacent carbon atoms. Imagine swapping the attached hydrogen atoms for bromine, for example. If benzene really was cyclohexa-1,3,5-triene, we’d expect it to form two different isomers: one with a double bond between the two affected carbons, and one with a single bond between them. We can see this below.

Fig. 9 - Two isomeric products of benzene

However, scientists only ever observed one isomer. This meant that benzene couldn’t have Kekulé’s predicted structure. It had to have identical bonds.

Poor Kekulé - he really thought he had cracked the mystery of benzene, but all the evidence was against him! He proposed one final idea - benzene consisted of two structures in equilibrium, rapidly shifting between both. This would result in a hybrid molecule that was neither one nor the other. He called this the resonance model. However, there was no evidence to support this. We instead now believe in the delocalised model described earlier. The delocalisation accounts for benzene’s resonance energy and the identical intermediate C-C bonds explain why benzene is a regular shape.

Fig. 10 - The resonance model. Kekulé believed that benzene was actually two structures that rapidly alternated between each other. Notice how the positions of the double bonds differ between the two molecules

The properties of benzene

Because of its unique structure and ring of delocalisation, benzene has some unique properties. Let’s explore them below.

Combustion

You should know that cyclic alkanes have a general formula ${\mathrm{C}}_{\mathrm{n}}{\mathrm{H}}_{2\mathrm{n}}$. A cyclic hydrocarbon with six carbon atoms would therefore have twelve hydrogen atoms. However, benzene has six carbon atoms but only six hydrogen atoms. This higher ratio of carbon to hydrogen means benzene burns with a characteristically sooty flame.

Melting and boiling point

Because benzene is nonpolar, the only forces it experiences between molecules are weak van der Waals forces, also known as London forces. However, benzene is a planar molecule. This means that in a solid state, it can pack together closely in neat layers. In contrast, cyclohexane is based on tetrahedral arrangements of atoms, meaning it has different hydrogen atoms sticking off in all directions! This means the molecules can’t fit together as neatly as a solid.

Closely-packed molecules experience stronger intermolecular forces than spaced-out molecules, so benzene has a higher melting point than cyclohexane. However, in liquid form this neat arrangement is destroyed. Both molecules, therefore, have similar boiling points.

Fig. 11 - The structure and melting and boiling points of benzene and cyclohexane

Solubility

Like other nonpolar hydrocarbons, benzene is insoluble in water but soluble in other organic solvents.

Reactivity of benzene

As we’ve explored earlier, benzene doesn’t like taking part in addition reactions. This would involve disrupting the strong ring of delocalisation, which is very stable because it spreads the electrons’ negative charges over a larger area.

However, benzene does take part in substitution reactions. These involve swapping one atom or group for another. In this case, we can swap hydrogen atoms for other species like halogens or hydroxyl groups.

Benzene’s ring of delocalisation is an area of electron density. This makes it very attractive to electrophiles.

What are electrophiles?

Electrophiles are electron pair acceptors. As the term -phile comes from the Latin philos, meaning love, we can say that they really just love electrons! Electrophiles have a positive or partial positive charge and a vacant orbital. Some common examples are ${\mathrm{H}}^{+}$ and ${\mathrm{NO}}_2$.

Electrophilic substitution reactions

We now know that benzene is susceptible to attack by electrophiles, and that it commonly reacts in substitution reactions. We can therefore conclude that most of the reactions involving benzene are electrophilic substitution reactions. We’ll cover these in Reactions of Benzene. These include:

• Nitration reactions, swapping a hydrogen atom for the $-{\mathrm{NO}}_2$ group. This produces nitrobenzene which is used in dyes and pharmaceuticals.
• Friedel-Crafts acylation reactions, where benzene reacts with an acid derivative in the presence of an aluminium chloride catalyst. The product is used for plastics and detergents.