Matthew J. Dowd, Graduate Student
Department of Medicinal Chemistry
Virginia Commonwealth University
Richmond, VA 23298-0540 USA

Click on the structures to view a 3-D version of the molecule using CHIME or RASMOL.

Benzene. C6H6. The epitome of aromatic compounds. For anyone who has taken organic chemistry, benzene is an important player. We learn that benzene is unusually stable compared to other unsaturated hydrocarbons. We also learn that benzene and other aromatic molecules can undergo certain reactions called electrophilic aromatic substitutions and nucleophilic aromatic substitutions. Because it is made of just carbons and hydrogens, we tend to think benzene is most similar to other hydrocarbons, such as gasoline (petrol), oil, or paraffin, which are generally nonpolar and hydrophobic (from Greek words meaning "water fearing"). Well, chemists are starting to learn that benzene may not be so afraid of water after all.

Ever since benzene was first isolated, it was the oddball of compounds isolated from compressed oil gas and coal distillates. In 1825, Michael Faraday (1791-1867), more famous for his insights into electricity and magnetism, was the first to identify benzene and determine its composition as containing six carbon atoms and six hydrogen atoms. But how were they connected to each other? At that time, it was known that unsaturated hydrocarbons (hydrocarbons with double or triple bonds between carbons) would react with hydrogen gas to take up a molecule or more of hydrogen. Based on the ratio of carbons to hydrogens in benzene, it was considered unsaturated. However, it didn't easily react with hydrogen. There were also some unanswered questions concerning the nature of certain isomers of substituted benzene derivatives.

The exact structure of benzene remained an enigma for several decades after its discovery. Many structures were proposed, but none withstood the scrutiny of experimental evidence. For example, the Scottish chemist Archibald Scott Couper (1831-1892) and the German chemist Josef Loschmidt (1821-1895) both proposed the structure shown below; however, neither chemist could provide empirical support.

It was not until 1865 when a young German by the name of Friedrich August Kekule (1829-1896) devised the hexagonal structure which we all know and love today. According to Kekule, the idea came to him in a dream. During a stay in Ghent, he was writing in an elegant bachelor's apartment. Soon, Morpheus attacked, and Kekule needed an afternoon nap. In his dream, the atoms were gamboling before his eyes. As he later recalled, there were "long rows, sometimes more closely fitted together, all twining and twisting in snake-like motion. But look! What was that? One of the snakes had seized hold of its own tail, and the form whirled mockingly before my eyes."

Fast forward almost 135 years and you would think that we, as chemists, completely understand benzene and other aromatic molecules. Well, guess again! We still have more to learn about benzene, aromatic compounds, and how they interact with polar and charged compounds.

Above, I stated that benzene is nonpolar and hydrophobic. In a nonpolar molecule, the negatively charged electrons are evenly distributed within the molecule. In a polar molecule, such as water (H2O), there is a greater negative charge on the end with the oxygen atom, because oxygen pulls electrons more strongly than hydrogen atoms. Technically, we say that this molecule has a "net dipole moment."

In benzene, no single carbon atom is more negatively charged than the other carbon atoms. In this sense, benzene is nonpolar; there is no net dipole moment. However, the electrons in benzene are not evenly distributed throughout the molecule. The carbon atoms, which are slightly more electronegative than the hydrogen atoms, pull electrons closer to themselves. And we can see the result of this when we view benzene's electrostatic potential surface (See Figure 3). The blue area, towards the center of the molecule, shows higher negative charge; the red color, located near the hydrogens, indicates more positive charge and lower electron density.

Now, when you look at benzene from the side, or edgewise, we notice that there is yet another uneven distribution of electrons, with more negative charge both above and below the hexagonal plane of carbons.

Figure 3. Electrostatic potential images of benzene (Blue indicates more negative charge; red more postive charge).

So, although benzene does not have a net dipole moment, its distribution of electrons is uneven across the molecule. This situation is termed a quadrupole.

Some chemists are trying to learn more about the quadrupole of aromatic compounds. Some early pioneering work in this field was performed by Dr. Lemont Kier. More recently, Dennis Dougherty, a physical organic chemist at the California Institute of Technology, has lead the investigation of the quadrupolar nature of benzene for the past decade. In particular, he is interested in how the aromatic compounds are attracted to other charged or polar molecules. One area of interest is the so-called pi-cation interaction.

A cation is a positively charged atom or molecule. With the pi-cation interaction, the positive cation is attracted to the negative area of benzene's quadrupole. It's not too hard to see how this attraction comes about. What takes a little more work is to determine the importance of this attraction in relation to other molecular forces, such as traditional electrostatic forces, hydrogen bonding, and van der Waals attraction.

Both theoretical and empirical evidence indicates that this kind of interaction is relatively strong. Many studies have looked at the interaction of benzene and cations in the gas phase, meaning that there are no solvent molecules to interfere. For example, when a potassium ion (K+) and benzene molecule are present, they are attracted to each other, with an association energy of about 19 kilocalories. Average hydrogen bonds, the all-important intermolecular interaction, are thought to be weaker than this. In fact, in the gas phase, K+ is more strongly attracted to benzene that to water, a very polar molecule. Another surprise is that water can bind to benzene. Apparently, the slightly positive hydrogens of water are attracted to the negatively (blue) charged face of benzene.

What other types of cations are attracted to benzene? So far, much of the research has focused on positively charged nitrogen containing molecules. Examples of these are protonated amines and quaternary amines, as shown in Figure 4.


One reason scientists have focused on nitrogen cations is their relationship to some very important biological molecules. Several neurotransmitters have a positively charged nitrogen portion in their chemical structure. Serotonin, a neurotransmitter involved in processes such as migraine, depression, and anxiety, contains a methylammonium portion. Other molecules include dopamine, epinephrine, and acetylcholine, the last of which contains a quaternary nitrogen.

If we're interested in biological examples of pi-cation attraction, we are not going to look at benzene, an extremely carcinogenic material that is not present in our cells. Certain amino acids, however, contain aromatic rings, including phenylalanine, tyrosine, tryptophan, and histidine. Several research groups have shown that charged nitrogen compounds are attracted to the aromatic portion of these amino acids.

Returning back to the neurotransmitters, these small molecules produce most of their biological effects by binding to the active site of protein receptors. We can examine the active site, looking for the presence of aromatic amino acids. For example, acetylcholine binds to a group of receptors called acetylcholine receptors. To prevent overstimulation of these receptors, there is an enzyme named acetylcholinesterase which degrades the molecule. Located at the binding sites on both proteins are a number of tyrosines and tryptophans. It is believed that the electron rich aromatic rings help stabilize the binding by attracting acetylcholine's charged nitrogen through pi-cation interactions. On their respective receptors, the binding sites for serotonin, dopamine, and epinephrine are also thought to have several aromatic amino acids. In these cases, the pi-cation interaction may augment other intermolecular attractive forces.

From a theoretical perspective, it is quite interesting to investigate the pi-cation interaction. You can even design molecules that tightly bind cations, such as the one reported by Dougherty's group and shown in Figure 7. This molecule, called a cyclophane, is capable of binding small organic cations, such as tetramethylammonium (See Figure 4). The nitrogen cation is the correct size to fit snugly inside the cavity. With the six benzene rings present in the cyclophane, the positively charged molecule is well stabilized - so much so that the cation prefers to remain in the host than in water.

One last molecule of interest is a derivative of benzene called hexafluorobenzene. In this molecule, all six hydrogens are replaced by fluorine atoms. If you examine the electrostatic potential surface of hexafluorobenzene, and compare them to the surface for benzene above, you may notice one interesting thing: the color scheme between to two molecules is reversed.

Figure 8. Electrostatic potential images of hexafluorobenzene.

That is, in hexafluorobenzene, the center of the molecule is more positive (red) and the edge is more negative (blue). Because fluorine is so strongly electronegative, it pulls more negative charge to the edge of the molecule. The result is the quadrupole of hexafluorobenzene is opposite that of benzene. Above, we noted that the polar molecule water is attracted to benzene, with the positive end (the hydrogens) of its molecules pointing towards benzene. Prof. Dougherty has recently shown, through molecular modeling, that water binds to hexafluorobenzene in an opposite fashion, as shown in Figure 9.

When learning about new ideas and concepts, the details can be very interesting, but sometimes overwhelming. Our case, the pi-cation interaction, is no different. But what is extremely important is the big picture - intermolecular forces. Understanding these microscopic forces is crucial in almost every aspect life: designing new drugs and medicines, making non-stick coatings for your cookware, creating the better motor oil. Probably the least studied and understood of these forces is the pi-cation interaction. However, recent work has and continues to shed light on this phenomenon. And, given that benzene entered the scientific forum with questions surrounding it last century, it seems only appropriate that the same molecule is involved in an exciting area of current investigation at the end of this century (and millennium).

Date posted: 11/06/99

Selected References:

1) Roberts, R.M. Serendipity. Accidental Discoveries in Science. New York: John Wiley and Sons, Inc., 1989, p. 75-82.

2) Brock, W.H. The Norton History of Chemistry. New York: W.W. Norton and Co., 1992, p. 263-269.

3) Gallivan, J.P.; Dougherty, D.A. Can Lone Pairs Bind to a Pi System? The Water...Hexafluorobenzene Interaction. Org. Lett. 1999, 1, 103-105.

4)Ma, J.C.; Dougherty, D.A. The Cation-pi Interaction. Chem. Rev. 1997, 97, 1303-1324.

5)Dougherty, D.A. Cation-pi Interactions in Chemistry and Biology: A New View of Benzene, Phe, Tyr, and Trp. Science 1996, 271, 163-168.

6) Kier, L.B.; Aldrich, H.S. A Theoretical Study of Receptor Site Models for Trimethylammonium Group Interaction. J. theor. Biol. 1974, 46, 529-541.