Volume 4

Benzene. August Kekulé's Uroboros Dream • The sp2 Model of Benzene • Isomerism and Nomenclature of Aromatic Compounds • Chemical Properties of Benzene • Production and Applications of Benzene and Other Aromatic Hydrocarbons • Exercises
4.5.1. Benzene. August Kekulé's Uroboros Dream. In the old times, aromatic compounds were called "aromatic" due to their fragrance. This name is traditionally maintained in modern chemistry, despite the fact that some aromatic compounds are odorless and some have an unpleasant smell. The central aromatic compound is benzene, C6H6. I personally would hesitate to identify the smell of benzene as particularly pleasant. Mothballs are made of naphthalene (C10H8), another aromatic compound. I doubt that the odor of mothballs could be described as nice or "aromatic". The structures of benzene and naphthalene are shown in Figure 4-58.
Figure 4-58. Chemical structures of benzene and naphthalene.

Aromatic hydrocarbons, also known as arenes, are ubiquitous, but what is an aromatic compound? One might think that there should be a pretty straightforward answer to this question. As surprising as you may find it, the issue is much more complicated than it seems. In advanced organic chemistry courses, one can find lengthy discussions of aromaticity. For us, it would suffice to define a hydrocarbon as aromatic if it (a) is cyclic, (b) is planar, (c) has double bonds, (d) has a total number of (4n + 2) electrons involved in the formation of the π bonds within the cycle, and (e) is prone to undergo substitution reactions with electron-deficient reagents rather than addition reactions. Is this too much for you? Let me try to put you at ease.

Years ago, I read an article telling the story of an international scientific meeting on aromaticity. At that meeting, a bunch of world experts discussed and argued over how to define an aromatic compound. In the end, they arrived at an interesting conclusion that an aromatic compound is a compound like benzene! Perhaps, this is just an urban legend, but even if it is, the story is still didactic – aromaticity is a nontrivial matter that is hard to define both strictly and broadly. In our course, we will study only benzene and its simple derivatives.

What is it that makes benzene so special in chemistry?

First of all, it is the three C=C double bonds (Figure 4-58) that do not behave as conventional, "normal" C=C double bonds like those found in alkenes. We will see that shortly.

Second, it is the exceptional stability of the benzene ring, its distinct chemical behavior, and the way it influences atoms and groups it is bonded to.

Last but not least, aromatic compounds, especially benzene and its derivatives, play an extremely important role in the production of drugs, dyes, crop protection agents, polymers, resins, synthetic fibers, lubricants, unique materials for electronics, detergents, explosives, coatings… you name it!

Benzene is a colorless liquid that boils at 80 oC at 1 atmosphere and freezes to a white crystalline solid at 5 oC. Benzene is flammable, immiscible with water, and has a density of 0.87 g/cm3 at 23 oC. Although benzene is both toxic and carcinogenic, it is produced on a huge scale of dozens of millions of tons annually due to its industrial importance.
Digression. We rarely think of what it took the true pioneers of scientific research to make their discoveries. We open a chemistry textbook and see the structure of benzene, but how often do we ask ourselves the questions, "Where does this structure come from? How hard was it to establish the cyclic structure of benzene?"

We do take a lot of things for granted. We use our smartphones on a daily basis, but do we ever think about what materials and chemical elements make smartphones possible? Take a look at Figure 4-59 to see that eight out of the total of fifteen different lanthanides are needed to build the device.
Figure 4-59. Lanthanides used in smartphones (source).
Moreover, all these lanthanides and their compounds must be super-pure. Natural sources of lanthanides contain all of them plus many other metals. The lanthanides are like twins exhibiting chemical properties that are so similar that separating them by conventional means represents a big challenge. Unsurprisingly, it took over 110 years of dedicated hard work of many scientists, countless failures, and hundreds of erroneous reports and incorrect claims before all of the lanthanides were discovered.

A particularly intriguing story in the history of lanthanides is that of praseodymium (Pr) and neodymium (Nd), two neighbors in the periodic table. Now we know that neodymium salts are pink in color and praseodymium salts are light green (Figure 4-60). As these two colors are complementary, solutions containing both Pr and Nd salts are often colorless. The deceiving lack of color and the extreme similarity of chemical and physical properties of neodymium and praseodymium compounds made many scientists believe that it was just one element. For nearly 45 years in the 19th century, there was a new "element", which was named didymium and given the symbol Di. Finally, in 1885 the Austrian scientist Carl Auer von Welsbach (1858-1929) separated, after thousands of recrystallizations, the two different lanthanides that constituted "didymium" and that since then have been known as praseodymium and neodymium (Figure 4-60).
Figure 4-60. Samples of Nd(NO3)36H2O (left) and Pr(NO3)36H2O (right); source.
Why are we talking about lanthanides if our current topic is aromatic hydrocarbons and benzene? This is because the path to our current understanding of benzene was as thorny as the long and winding road to the separation and discovery of the lanthanides. It took over 100 years from the original discovery of benzene by Michael Faraday (1791-1867) in 1825 to Linus Pauling's (1901-1994) 1931 work describing the currently used electronic structure of benzene. The cyclic structure of benzene was first proposed in 1865 by August Kekulé (1829-1896) who came up with this idea while daydreaming about Uroboros, an ancient symbol of a snake biting its own tail (Figure 4-61). Seven years later, in 1872, Kekulé proposed that the double and single bonds of the benzene ring rapidly interchanged their positions, as shown in Figure 4-62. As we know now, however, there are no localized single and double bonds in benzene. We will learn that in the next subsection.
Figure 4-61. Left: The famous drawing of Uroboros by Theodoros Pelecanos of Corfu (15th century; source). Right: Kekulé's cyclic structure of benzene, inspired by Uroboros (source).
Figure 4-62. The original drawing in Kekulé's 1872 publication, illustrating his hypothesis of oscillating C=C bonds in benzene (source).

4.5.2. The sp2 Model of Benzene. There have been many experimental observations indicating that there is something wrong with Kekulé's benzene structure shown in Figures 4-58, 4-61, and 4-62.

First, the chemical behavior of benzene is inconsistent with the Kekulé structure. As we already know, such molecules as Br2, HBr, and water readily add to the C=C double bond of alkenes. While cyclohexene is no exception, no such reactions occur with benzene (Figure 4-63). The lack of reactivity of benzene toward reagents that easily add to alkenes suggests that the C=C bonds of benzene are different in nature from those of alkenes.
Figure 4-63. Cyclohexene undergoes addition of Br2 and HBr (left), whereas benzene does not (right).

Second, the Kekulé structure with alternating single and double bonds is inconsistent with the experimental data. We know (Table 4) that the C=C double bond is about 1.33 Å long and the C-C single bond is much longer, 1.54 Å. According to these numbers, the structure of benzene would look like the one shown in Figure 4-64. All experimental structure determination methods, however, show that the molecule of benzene is an ideal hexagon. All six carbon-carbon bonds in benzene are identical and equal, being 1.40 Å in length, longer than the C=C double bond (1.33 Å) and shorter than the C-C single bond (1.54 Å).
Figure 4-64. An incorrect structure of benzene with alternating localized single and double carbon-carbon bonds.

Third, experiments and calculations indicate that benzene is 36 kcal/mol more stable than the Kekulé structure featuring the alternating double and single bonds. Without going into details, we just emphasize that the 36 kcal/mol difference is huge.

The experimental data is reconciled with theory nicely if we build a molecule of benzene from six sp2-hybridized carbon atoms (Figure 4-65 A) and six hydrogen atoms (Figure 4-65 B). (If needed, refresh your memory on sp2 hybridization, see subsection 4.3.2.) In Figure 4-65, each of the six carbon atoms has three sp2 hybrids (drawn in red) lying in one plane and forming an equilateral triangle. Orthogonal to this plane is the non-hybridized 2p orbital (drawn in blue).
Figure 4-65. Building a molecule of benzene from six sp2-hybridized carbon atoms (A) and six hydrogen atoms (B) showing the formation of the C-H and C-C σ-bonds (C) and the positions of the six non-hybridized 2p orbitals (D).

First, we build the σ-bonds, six C-C and six C-H, using six sp2-hybridized carbon atoms and six hydrogen atoms. Let us forget about the existence of the 2p non-hybrids for a moment. We position the six C atoms such that they form a hexagon, with all of their sp2-hybridized orbitals being coplanar. Two of the three sp2 orbitals of each carbon atom overlap with sp2 orbitals of the two neighboring carbon atoms. The third sp2 orbital on each C atom overlaps with the 1s orbital of a hydrogen atom, drawn in purple. In this way, the hexagonal C6H6 skeleton is produced (Figure 4-65 C).

Now that our hexagon of the C-C and C-H σ-bonds has been constructed, it is time to recall that each of the carbon atoms has a 2p non-hybrid orbital (drawn in blue) perpendicular to the plane (Figure 4-65 D). These 2p orbitals are shown in a separate drawing (D) because adding them to image C would produce an overcrowded and confusing picture. Superimpose the dotted line hexagons of D and C in your mind's eye to see the whole pattern.

Being more diffuse (bulkier) than shown in Figure 4-65 D, the six 2p non-hybrid orbitals are ideally positioned for a side-to-side overlap, just as in the π-bond formation in alkenes. In benzene, however, each of the 2p orbitals overlaps with both of its neighbors, which results in a delocalized system of π-bonds (Figure 4-66). Consequently, there are neither single nor double bonds in benzene, but six identical carbon-carbon bonds with the bond order of 1.5. Bond order is the number of chemical bonds between two atoms. In ethane, ethylene, and acetylene, the bond order between the two carbon atoms is 1, 2, and 3, respectively. In benzene, it is fractional, 1.5. Fractional bond orders are known for other molecules. For example, the bond order between oxygen atoms in ozone, O3, is also 1.5.
Figure 4-66. Formation of the delocalized π system of benzene via side-to-side overlap of all six 2p non-hybridized orbitals.

Although there are no localized single and double bonds in benzene, chemists still often use the Kekulé structure, while keeping in mind that the π-bonds are delocalized. Alternatively, the three double bonds in the Kekulé structure are replaced with a circle to depict the delocalization of the π electron cloud over the entire ring (Figure 67). Note that chemists conventionally omit H atoms when drawing structures of benzene and other aromatic compounds. We will do so in this course, too, albeit not always.
Figure 4-67. Conventional ways to draw structures of aromatic compounds.

4.5.3. Isomerism and Nomenclature of Aromatic Compounds. Numerous aromatic compounds are known and many of them have common names. These common names are vastly preferred over the "official" long and clumsy systematic names (Table 5). You do not have to memorize any of these names for this course, except for just "benzene" and "naphthalene".

Table 5. Common and systematic names for selected simple aromatic hydrocarbons.
Hydrogen atoms of benzene can be substituted with other atoms such as halogens, or groups of atoms. Benzene derivatives with only one substituted H atom do not have isomers and are named by placing the name of a substituent as a prefix to "benzene" (Figure 4-68). Some mono-substituted benzenes have common names, such as toluene for methylbenzene, aniline for aminobenzene, and anisole for methoxybenzene.
Figure 4-68. Examples of mono-substituted benzenes with names.

For benzene molecules bearing two substituents, three positional isomers are possible (Figure 4-69). If the two substituents are next to each other, this substitution pattern is called ortho. Isomers bearing two substituents at carbon atoms separated by one carbon atom of the ring are called meta. If the substituents occupy opposite positions on the ring, such isomers are called para. Alternatively, the carbon atoms of a disubstituted benzene are numbered starting with one of the two bearing the substituents. The second substituent should have the lowest possible number.
Figure 4-69. Three isomers of dimethylbenzene (xylene).

For trisubstituted benzenes, it is the numbering system that is used much more frequently (Figure 4-70).
Figure 4-70. Structures and names of all three isomeric trimethylmenzenes.

One might wonder how to name a benzene derivative bearing two different substituents on the ring, such as the molecule below.
Should we start the numbering at the carbon bearing the methyl group, or the one that is bonded to the ethyl? The numbering starts at that carbon atom of the benzene ring that is connected to a substituent of the highest priority. In our particular case, ethyl has a higher priority than methyl. Therefore, the correct name is 1-ethyl-4-methylbenzene, not 1-methyl-4-ethylbenzene. (It is also appropriate to name this molecule para-ethyltoluene.) The substituent priority rules are beyond the scope of this course, but if you would like to learn them, click here.

Like the CH3 group (methane without one H atom) that has a special name, methyl, the C6H5 group (benzene without one H atom) also has a widely used name: phenyl.

4.5.4. Chemical Properties of Benzene. The addition reactions that are most commonly observed for alkenes are not characteristic of benzene because the π-bonds of benzene are delocalized. While ethylene and other alkenes add Br2 quickly, mixing bromine with benzene does not result in any reaction. However, in the presence of a catalyst such as AlBr3, benzene readily reacts with bromine. The outcome of this reaction is substitution of one of the six hydrogen atoms with a bromine atom (Figure 4-71). Note that the H atoms on the ring are not shown in Figure 4-71, but they are all there, all six on the starting molecule of benzene and five on the product, bromobenzene.
Figure 4-71. Bromination of benzene in the presence of AlBr3.

A very important reaction of benzene is nitration with nitric acid in the presence of sulfuric acid (Figure 4-72). The reaction produces highly industrially important nitrobenzene, a pale-yellow oily liquid with a characteristic almond smell. As in the previous equation (Figure 4-71), all H atoms on the benzene rings in Figure 4-72 are omitted for clarity.
Figure 4-72. Nitration of benzene with HNO3 in the presence of H2SO4.

The mechanism of aromatic substitution reactions such as the bromination (Figure 4-71) and nitration (Figure 4-72) is well-established. These reactions start with the generation of a reactive species that can serve as an electron pair acceptor. Organic chemists call such "hungry" for electrons molecules or ions electrophiles, from electron and philos (φίλος), the Ancient Greek for "that which is loved". In the nitration of benzene, the reactive species (electrophile) is the nitronium cation, NO2+, which is generated from HNO3 and H2SO4 (Figure 4-73), as discussed above in Volume 3, subsection 3.7.3.
Figure 4-73. The formation of NO2+ from HNO3 and H2SO4.

Being hungry for electrons, the nitronium cation is attracted to the π electron cloud of benzene (Figure 4-74). The addition of NO2+ disrupts the delocalized π-system and changes the hybridization of the carbon atom attacked from sp2 to sp3. The positive charge brought to the benzene ring by the NO2+ is delocalized over the remaining π bonds, as depicted by the dotted line in Figure 4-74. Finally, the departure of the proton from the carbon atom bonded to the NO2 group restores the π-system, thereby leading to the formation of nitrobenzene.
Figure 4-74. Simplified mechanism of nitration of benzene.

The mechanism of the bromination of benzene is similar. First, the Br+ electrophile is produced via abstraction of Br- from Br2 by AlBr3 (Figure 4-75). In this process, the AlBr3 acts as a Lewis acid. Up to this point we have considered only Brønsted-Lowry acids and bases. The Brønsted-Lowry classification of acids and bases is based on transfer of a proton. A Brønsted-Lowry acid is a molecule or ion that can serve as a source of a proton (H+). A Brønsted-Lowry base is a molecule or ion that can serve as an acceptor of a proton (H+). The most common Brønsted-Lowry base in simple inorganic reactions such as neutralization is the hydroxide anion, OH-.

Gilbert N. Lewis (1875-1946), the same outstanding American chemist who first formulated the idea of what is now known as the covalent bond and developed the dot diagrams, came up with a different classification of acids and bases. A Lewis acid is a molecule or ion that can accept a lone electron pair. A Lewis base is a molecule or ion that can donate a lone electron pair.

In AlBr3, a typical Lewis acid, the aluminum atom has only six valence electrons and, consequently, is eager to accept an electron pair to expand its valence shell occupancy to a stable octet. The AlBr3 abstracts a Br- from Br2 to generate the Br+ that attacks the benzene ring in the same manner as NO2+ does (Figures 4-74 and 4-75).
Figure 4-75. Simplified mechanism of bromination of benzene.

In fact, it is not the Br+ per se that is produced in the reaction, but rather its chemical equivalent, Brδ+---Brδ----AlBr3. For illustrative purposes, however, we simplify the mechanism.

The displacement of one H atom on the benzene ring with a different atom (e.g., Br) or a group (e.g., NO2) changes the ability of the remaining five C-H bonds to undergo substitution reactions. For example, bromobenzene and nitrobenzene are less reactive than benzene itself. On the contrary, a methyl group on the benzene ring makes the C-H bonds of the ring more reactive. For this reason, trinitrotoluene (TNT), the well-known explosive, is produced by nitration of toluene (Figure 4-76) much more easily than trinitrobenzene from benzene under similar conditions.
Figure 4-76. Nitration of toluene to trinitrotoluene.

Not only does the CH3 group on the benzene ring make it more reactive, but also the benzene ring enhances the reactivity of the methyl group. For example, methane does not react with KMnO4, whereas toluene is quite easily oxidized by potassium permanganate to benzoic acid (Figure 4-77).
Figure 4-77. Lack of reaction of methane with KMnO4 and oxidation of toluene with KMnO4.

Benzoic acid is a white crystalline solid, which is broadly used as a food preservative. Oxidation of both methyl groups of para-dimethylbenzene (1,4-dimethylbenzene, para-xylene) gives terephthalic acid. Terephthalic acid is produced on a billion kilogram scale annually for making polyester, one of the most important polymers used in the manufacturing of fabrics, fibers, and plastic bottles. The industrial oxidation of para-xylene is not done with KMnO4 because that would be way too expensive. Instead, catalytic oxidation by oxygen of the air is used (Figure 4-78).
Figure 4-78. Oxidation of para-xylene to terephthalic acid.
Digression. Toluene is cheaper than para-xylene by only a few cents a pound. However, if you discovered an economical way to terephthalic acid from toluene rather than from para-xylene, you would become a multimillionaire due to the huge production scale of polyester.
Although, as discussed above, addition reactions are uncharacteristic of benzene, such reactions may occur under forcing conditions, such as in the presence of an exceedingly active catalyst (addition of H2; Figure 4-79) or under intense light (addition of Cl2; Figure 4-80).
Figure 4-79. Catalytic hydrogenation of benzene to cyclohexane.
Figure 4-80. Intense light-induced addition of chlorine to benzene.

Hexachlorocyclohexane, the product of the chlorination reaction shown in Figure 4-80, was broadly used as an insecticide under the trade name Lindane, but not anymore. Lindane shampoo, however, is still available to treat scabies and lice.

4.5.5. Production and Applications of Benzene and Other Aromatic Hydrocarbons. Modern methods to produce aromatic hydrocarbons are based on various thermal and catalytic processes in petrochemical refineries and coke plants. All these processes produce mixtures of aromatic compounds, which are separated by a variety of techniques. Simple aromatic compounds, including benzene, toluene, naphthalene, and isomers of xylenes as well as many of their derivatives are commercially available.

Aromatic compounds have an exceptionally broad spectrum of applications. Some of these applications have been around for a long time, such as in materials for construction, car, aircraft, and spacecraft building, manufacturing of dyes, fabrics, fibers, explosives, lubricants, coatings, adhesives, etc. Active ingredients of numerous pharmaceuticals and agrochemicals are aromatic derivatives, including Aspirin, Tylenol, Ibuprofen, and Naproxen. The second most widely used pesticide in the U.S.A. (after Glyphosate) is Atrazine, an aromatic compound.

The most modern application of aromatic compounds is in electronics. Without materials based on aromatic compounds, there would have been no LCD and OLED displays, no smartphones, no CDs and DVDs, no modern computers, no modern audio and video equipment. Quite amazingly, aromatic materials originally developed for a particular application, sometimes find use in a totally different area. For instance, Kevlar® (subsection 4.9.5 below), the aromatic polymer originally used to make bullet-proof vests, has more recently found use in top of the line loudspeakers made by the U.K. company Bowers & Wilkins (B&W) and preferred by the renowned Abbey Road Studios in London, England.

4.5.6. Exercises.

1. Aromatic compounds are called aromatic because they all have a pleasant aroma. True or false? Answer

2. Benzene is (a) a liquid that boils below the boiling point of water and solidifies above the melting point of ice; (b) widely used in industry in spite of being toxic and carcinogenic; (c) a nonhazardous compound; (d) highly reactive toward Br2 that easily adds to all three C=C double bonds of benzene in the absence of a catalyst and light to give hexabromocyclohexane; (e) less reactive toward nitration and bromination than toluene. Answer

3. All six carbon-carbon bonds in benzene are equal in length and strength because the alternating double C=C bonds and single C-C bonds constituting the ring quickly interchange their positions. True or false? Answer

4. Draw chemical structures of (a) toluene; (b) naphthalene; (c) para-dichlorobenzene; (d) 1,4-dinitrobenzene; (e) 1,3,5-tribromobenzene; (f) ortho-difluorobenzene; (g) 1,2,3-trimethylbenzene; (h) para-xylene.

5. What is wrong with the following chemical names: (a) 1,3,6-triethylbenzene; (b) 4,5-dinitrotoluene; (c) para-dibromotoluene. Answer

6. Using the hybridization concept, describe the electronic structure of benzene and explain the idea of π electron density delocalization. [Answer: See 4.5.2]

7. To prepare bromobenzene, C6H5Br, from benzene C6H6 and bromine Br2, the amount of bromine used per each mole of benzene should be (a) 0.5 mol; (b) 1 mol; (c) 2 mol. Answer

8. Write the mechanism of nitration of benzene with HNO3 in the presence of H2SO4. Answer

9. Like benzene, benzoic acid is toxic and carcinogenic. True or false? Answer

10. Draw the structures of benzoic and terephthalic acids. [Answer: See Figures 4-77 and 4-78]

11. In industry, terephthalic acid is produced by oxidation of para-xylene (1,4-dimethylbenzene) with KMnO4. True or false? Answer

12. Define a Lewis acid, a Lewis base, a Brønsted-Lowry acid, and a Brønsted-Lowry base. Answer

13. Are toluene and xylenes homologues of benzene? Answer