Benzene (C6H6) is a colorless, flammable, aromatic hydrocarbon, that is a known carcinogen. It boils at 80.1°C and solidifies at 5.5°C. Produced by hydrogen reduction of some allotropes of carbon, or from petroleum, it is used in the creation of drugs, plastics, synthetic rubber, and dyes.
Other aromatic compounds created by the replacement of hydrogen atoms with methyl (CH3) groups are called the benzene series. If one hydrogen is replaced, the new chemical is called toluene, (C6H5CH3), from which trinitrotoluene, or TNT is derived. If two hydrogens are replaced it becomes xylene, (C6H4(CH3)2).
Replacement of the hydrogen atoms with other functional groups produces additional derivaties. A hydroxyl group (OH) produces phenol (C6H5OH), and additional nitridation produces picric acid, or trinitrophenol. Replacement with an amino group (NH3) produces aniline (C6H5(NH3))
Two or more rings may be joined together, as in naphthalene, anthracene, and phenanthrene. Other atoms, such as nitrogen, may be substituted for carbon atoms in the ring, as in pyridine (C5H5N) and pyrimidine (C4H4N2).
The formula of benzene (C6H6), caused a mystery for some time after it's discovery, as no proposed structure could take account of all the bonds (Carbon usually forms four single bonds and hydrogen one). Legend has it that the chemist Kekule dreamt of a snake eating it's own tail, and he discovered the ring structure of benzene. Benzene presents a problem, as to account for all the bonds, there must be alternating double carbon bonds:
However, all of the carbon-carbon bonds in benzene are of the same length, and it is known that a single bond is longer than a double bond. In addition, the bond length (the distance between the two bonded atoms) in benzene is greater than a double bond, but shorter than a single bond. There seems in effect to be a bond and a half between each carbon.
This is explained by electron delocalisation. In order to picture this, we must consider the position of electrons in the bonds of benzene. The single bonds are formed with electrons orbiting in paths in line with this page. The double bonds consist of a single bond and another bond. This second bond has electrons orbiting in paths above and below the plain of this page at each bonded carbon atom. In the diagram below we take a side-view of this occurring:
The '-' denotes the single bond, and the '@'s denote the orbitals of the electrons forming the double bonds.
Being out of the plane of the atoms, these electrons can interact with each other freely, and become delocalised. This means that instead of being tied to one atom of carbon, they are shared by all six in the ring. Thus there are not enough to form double bonds on all the carbon atoms, but the atoms do strengthen all of the bonds on the ring equally.
What in effect is happening is that the structure exists as a superposition of the above forms, rather than either form individually. This type of structure is called a resonance hybrid.
To reflect the delocalised nature of the bonding, benzene is usually depicted as a circle inside a hexagon in chemical structure diagrams:
Benzene occurs sufficiently often as a component of organic molecules that there is a Unicode symbol to represent it; code 232C, ⌬