Benzene


February 25, 1998

Benzene Lewis Structure
Lewis structure
Benzene Kekule Structure Benzene Hybrid Structure
Kekule structure, left, and Benzene hybrid structure, right

When I worked in a Forensic Chemistry lab in 1996-1997 at Florida International University, I developed an interest in the benzene molecule. I was fascinated that this molecule could be used both in nature and in the lab to create many organic compounds used in every day life, yet by itself, it could be highly dangerous. Its shape is the key in how it can be used as an intermediate building block for compounds and how it can also have such serious health effects. Its hexagonal shape allows it to be used to synthesize massive chains of compounds. Unfortunately, the benzene ring is also a shape ideal for getting between the base pairs in DNA to cause genetic mutations that are usually fatal.

Benzene Computer Generated Drawing, side viewBenzene Computer Generated Drawing, top view
Computer generated drawing, side and top view

Benzene is a clear, colorless flammable liquid at room temperature. It is also aromatic, a term originally referring to fragrance, but now used to refer to benzene and its structural relatives. It does have a distinct fragrance; benzene smells like gasoline fumes. The benzene-related aromatic compounds are part of an enormous family of compounds that includes many solvents, dyes, rubbers, resins, plastics, neurotransmitters, amino acids, DNA bases, drugs, and synthetic fabrics like nylon and polyester. Its use as a solvent has been replaced by the non-carcinogen toluene. Toluene has a methyl group where a hydrogen would normally be attached in benzene; this minute change in shape does not allow toluene to get between DNA base pairs like benzene does.

The Environmental Protection Agency was first required in 1974 to determine safe levels of chemicals in drinking water which do or may cause health problems, thanks to the Safe Drinking Water Act. These non-enforceable levels, based solely on possible health risks and exposure, are called Maximum Contaminant Level Goals (MCLG). Based on the MCLG, the EPA has set an enforceable standard called a Maximum Contaminant Level (MCL), which is set as close to the MCLG as possible and is 5 parts per billion. Any exposure over the MCL can potentially cause severe health problems, depending on amount of exposure. The short-term health risks attributed to benzene exposure include temporary nervous system disorders, immune system depression, and anemia. Long-term risks include chromosome aberrations and forms of cancer that include leukemia.

Benzene's structure is that of six carbons arranged in a hexagonal ring with six hydrogens attached. The ring is made up of alternating single and double bonds, giving benzene a resonance structure that resonates in a circle. This resonance, as well as benzene's large relative mass, accounts for its stability; the resonance sets up a situation in which the electrons are delocalized and thus are able to form very strong London dispersion forces. The atoms in benzene are co-planar because each carbon is sp2 hybridized and has 120 degree bond angles; thus all six carbons are equivalent. All six bond lengths are also equivalent; therefore the circular resonance does not mean that the double and single bonds are rapidly equilibrating, but rather that the structure contains six equal bonds that have bond lengths halfway between that of a single C-C bond and a double C-C bond.

Benzene Ball and Stick Model, side viewBenzene Ball and Stick Model, top view
Ball and stick model, side and top view

The true resonance structure for benzene is the hybrid between the two resonance structures; an analogy for this is one resonance structure of benzene being a dragon and the other a unicorn, with the true structure being the cross between the two, a rhinoceros. While the two structures are frequently drawn in Organic Chemistry books in order to show that resonance occurs, the real structure is the hybrid. This hybridization never changes. Benzene's unusual stability stems from its resonance and the overlap of the delocalized p-orbitals. Benzene's conjugated pi bond system over and under the carbons creates an overlap of p-orbitals that forms an electron cloud in the shape of a doughnut. These clouds aid in the formation of unusually strong London dispersion forces. In addition, benzene should have three double bonds, and the standard test for double bonds is reactivity with Br2. Benzene does not react with Br2 because it is stabilized by its resonance energy.

Benzene Ball and Stick Model with Pi-bond Electron Cloud 'Doughnut', top view
Ball and stick model with pi-bond electron cloud "doughnut"

Benzene is a basic aromatic compound. Its aromaticity is due to the combination of having conjugated double bonds and a six membered ring; however, an obvious drawback to aromatic compounds in general is that they are for the most part carcinogenic. In addition to benzene, there are many other compounds that have fused 6-membered aromatic rings, such as napthalene, anthracene and phenanthrene. Once chemists understood how benzene's structure worked in terms of its aromaticity and its having resonance, they were able to manipulate it to create the compounds we take for granted today, such as polyester, nylon, and plastic. Benzene is a deadly compound if not properly handled; when in the right hands, however, benzene is a tool that can outline basic properties of compounds with similar structures or structures based on benzene, and thus it can have many useful and safe applications.

Electrostatic Surface of Benzene, top viewElectrostatic Surface of Benzene, side view
Electrostatic surface of benzene

EPA Information

https://www.epa.gov/dwreginfo/drinking-water-regulations

General Information

York University Chemistry Department's Benzene Page
Hazard.com's International Safety Card
Medical Chemistry Department of the Albert Szent-Gyorgyi Medical University, Szeged, Hungary
McMurry, John. Organic Chemistry. 1996. Brooks/Cole Publishing Company. Boston.
Lewis Structure image taken from John McMurry's "Organic Chemistry" with permission.

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