A2-Level Benzene, Aromatic Chemistry
Benzene is a cyclic hydrocarbon made of six carbon atoms and six hydrogen atoms (C H ).
The Kekule structure shows bonding in benzene as alternating carbon double bonds.
Bond length data and enthalpies of hydrogenation for benzene show this to be incorrect
Each carbon atom has an un-bonded electron in a p-orbital. The p-orbitals of all carbon atoms merge to create two rings of electron density - spreading electrons throughout the carbon ring (delocalised electrons).
Due to the delocalised electron rings, benzene molecules are easily attacked by electrophiles.
Compounds that contain benzene are called aromatic compounds.
The study of these compounds is called aromatic chemistry
In organic chemistry, hydrocarbons can form cyclic compounds. A bit like a chain that links back to itself to form a necklace.
Cyclohexane is a common example of a cyclic compound, containing six carbon atoms that are linked to one another.
A specific type of cyclic compound is called benzene. In benzene, there are six carbon atoms arranged in a ring and each one is bonded to a single hydrogen atom. The bonding that takes place between carbon atoms in benzene is what makes it so interesting and unique.
If we look at cyclohexane, it is clear to see that each carbon atom forms two single bonds with other carbon atoms and two single bonds with hydrogen atoms. This makes sense as carbon can form four covalent bonds.
In benzene however, things get a little more complicated. If we draw benzene out it in the same way as cyclohexane, each carbon atom is now only forming three single bonds. Two to carbon atoms and one to a hydrogen atom.
The obvious thing to assume is that each carbon atom forms a double bond with another carbon atom. This gives three double bonds and this representation of benzene is called the Kekule structure.
The Kekule structure does address the number of bonds each carbon atom makes, but does not fully explain the structure of benzene.
Problems with the Kekule Structure
There are two main problems with the Kekule structure – bond lengths and hydrogenation enthalpies.
Single carbon-carbon bonds have a different length to double carbon-carbon bonds.
C-C bond length is 0.154nm
C=C bond length is 0.134nm
Experimental data shows that all the carbon bonds within benzene have a bond length of 0.139nm. This bond length does not make sense if we think of only single and double carbon bonds, therefore the Kekule structure is not an accurate representation of the bonding.
Double carbon bonds can be hydrogenated (reacted with H to break the double bond and add hydrogen atoms to each carbon). As single carbon bonds are more stable than double carbon bonds, when a double bond is hydrogenated, energy is released.
In cyclohexene, there is one double carbon bond. When hydrogenated the amount of energy released per mole of cyclohexene is 120kJ. The enthalpy of hydrogenation is -120kJmol . (Remember energy is released – negative enthalpy change).
In the Kekule structure, there are three double bonds. If one double bond in cyclohexene releases 120kJ per mol when hydrogenated, three times this amount of energy should be released when the double bonds in benzene are hydrogenated.
3 x -120kJmol = -360kJmol .
The predicted enthalpy of hydrogenation is this value, but the actual enthalpy change that occurs when benzene is hydrogenated is -208kJmol .
Clearly there are not three double bonds in benzene.
What is actually going on in benzene?
To really understand the bonding in benzene, we need to look at the electron orbitals around each carbon atom. Remember sigma bonds are formed when orbitals perfectly overlap, giving single covalent bonds.
Each carbon atom has three sigma bonds, leaving one electron left over. The electron left over is in a p-shaped orbital.
Due to orbital hybridisation (not required understanding for A-level chemistry, but very interesting!), the sigma bonds are all in the same plane – this means you could take a slice through the molecule and it would pass through all three atoms. The p-orbital sticks up and below this plane.
The electrons in each p-orbital repel each other. To minimise this repelling, they alternate above and below the plane of the carbon atoms.
There are now three electrons ‘above’ and three electrons ‘below’ the plane of the carbon and hydrogen atoms. These half-filled p-orbitals are not stable and, rather than stay in a p-orbital, the electrons spread out evenly in two regions – one above and one below the plane. Technically, each of the two regions form a pi-bond (merging of p-orbitals to create a bonding orbital). This is why benzene is considered to have a pi-bonding system.
The electrons are no longer held in p-orbitals and become delocalised electrons. Each carbon atom has ‘access’ to a full outer shell of electrons and all bonds are the same length (and energies).
It is not possible to draw the individual carbon bonds as either single or double bonds, so a circle is used in the middle of the ring of carbon atoms to represent delocalised electrons.
Electrons have a negative charge, so the circle represents an area of high electron density that electrophiles are attracted to. This is why benzene reacts well with electrophiles.