Asimov's New Guide to Science

by Isaac Asimov

  So far, so good, but where did the asymmetry lie? What was there about the two molecules that made them mirror images of each other? Pasteur could not say. And although Biot, who had suggested the existence of molecular asymmetry, lived to be eighty-eight, he did not live long enough to see his intuition vindicated.

  It was in 1874, twelve years after Biot’s death, that the answer was finally presented. Two young chemists, a twenty-two-year-old Dutchman named Jacobus Hendricus Van’t Hoff and a twenty-seven-year-old Frenchman named Joseph Achille Le Bel, independently advanced a new theory of the carbon valence bonds that explained how mirror-image molecules could be constructed. (Later in his career, Van’t Hoff studied the behavior of substances in solution and showed how the laws governing their behavior resembled the laws governing the behavior of gases. For this achievement he was the first man, in 1901, to be awarded the Nobel Prize in chemistry.)

  Kekulé had drawn the four bonds of the carbon atom all in the same plane, not necessarily because this was the way they were actually arranged but because it was the convenient way of drawing them on a flat piece of paper.

  Van’t Hoff and Le Bel now suggested a three-dimensional model in which the bonds were directed in two mutually perpendicular planes, two in one plane and two in the other. A good way to picture this model is to imagine the carbon atom as standing on any three of its bonds as legs, in which case the fourth bond points vertically upward. If you suppose the carbon atom to be at the center of a tetrahedron (a four-sided geometrical figure with triangular sides), then the four bonds point to the four vertexes of the figure. The model is therefore called the tetrahedral carbon atom. (see figure 11.2).

  Figure 11.2. The tetrahedral carbon atom.

  Now let us attach to these four bonds two hydrogen atoms, a chlorine atom, and a bromine atom. Regardless of which atom we attach to which bond, we will always come out with the same arrangement. Try it and see. With four toothpicks stuck into a marshmallow (the carbon atom) at the proper angles, you can represent the four bonds. Now suppose you stick two black olives (the hydrogen atoms), a green olive (chlorine), and a cherry (bromine) on the ends of the toothpicks in any order. Let us say that when you stand this on three legs with a black olive on the fourth pointing upward, the order on the three standing legs in the clockwise direction is black olive, green olive, cherry. You might now switch the green olive and cherry so that the order runs black olive, cherry, green olive. But all you need to do to see the same order as before is to turn the structure over so that the black olive serving as one of the supporting legs sticks up in the air and the one that was in the air rests on the table. Now the order of the standing legs again is black olive, green olive, cherry.

  In other words, when at least two of the four atoms (or groups of atoms) attached to carbon’s four bonds are identical, only one structural arrangement is possible. (Obviously this is true when three or all four of the attachments are identical.)

  But when all four of the attached atoms (or groups of atoms) are different, the situation changes. Now two different structural arrangements are possible—one the mirror image of the other. For instance, suppose you stick a cherry on the upward leg and a black olive, a green olive, and a cocktail onion on the three standing legs. If you then switch the black olive and green olive so that the clockwise order runs green olive, black olive, onion, there is no way you can turn the structure to make the order come out black olive, green olive, onion, as it was before you made the switch. Thus with four different attachments you can always form two different structures, mirror images of each other. Try it and see.

  Van’t Hoff and Le Bel thus solved the mystery of the asymmetry of optically active substances. The mirror-image substances that rotated light in opposite directions were substances containing carbon atoms with four different atoms or groups of atoms attached to the bonds. One of the two possible arrangements of these four attachments rotated polarized light to the right; the other rotated it to the left.

  More and more evidence beautifully supported Van’t Hoff’s and Le Bel’s tetrahedral model of the carbon atom; and, by 1885, their theory (thanks, i~ part, to the enthusiastic support of the respected Wislicenus) was universally accepted.

  The notion of three-dimensional structure also was applied to atoms other than carbon. The German chemist Viktor Meyer applied it successfully to nitrogen, while the English chemist William Jackson Pope applied it to sulfur, selenium, and tin. The German-Swiss chemist Alfred Werner added other elements and, indeed, beginning in the 1890s, worked out a coordination theory in which the structure of complex inorganic substances was explained by careful consideration of the distribution of atoms and atom groupings about some central atom. For this work, Werner was awarded the Nobel Prize in chemistry for 1913.

  The two racemic acids that Pasteur had isolated were named “d-tartaric acid” (for “dextrorotatory”) and “l-tartaric acid” (for “levorotatory”), and mirror-image structural formulas were written for them. But which was which? Which was actually the right-handed and which the left-handed compound? There was no way of telling at the time.

  To provide chemists with a reference, or standard of comparison, for distinguishing right-handed and left-handed substances, the German chemist Emil Fischer chose a simple compound called glyceraldehyde, a relative of the sugars, which were among the most thoroughly studied of the optically active compounds. He arbitrarily assigned left-handedness to one form which he named “L-glyceraldehyde,” and right-handedness to its mirror image, named “D-glyceraldehyde.” His structural formulas for them were:

  Any compound that could be shown by appropriate chemical methods (rather careful ones) to have a structure related to L-glyceraldehyde would be considered in the “L—series” and would have the prefix “L” attached to its name, regardless of whether it was levorotatory or dextrorotatory as far as polarized light was concerned. As it turned out, the levorotatory form of tartaric acid was found to belong to the D-series instead of the L-series.Nowadays, a compound that falls in the D-series structurally but rotates light to the left has its name prefixed by “D(–)”; similarly, we have “D(+),” “L(–),” and “L(+).”

  This preoccupation with the minutiae of optical activity has turned out to be more than a matter of idle curiosity. As it happens, almost all the compounds occurring in living organisms contain asymmetric carbon atoms. And in every such case, the organism makes use of only one of the two mirror-image forms of the compound. Furthermore, similar compounds generally fall in the same series. For instance, virtually all the simple sugars found in living tissue belong to the D-series, while virtually all the amino acids (the building blocks of proteins) belong to the L-series.

  In 1955, a Dutch chemist named Johannes Martin Bijvoet finally determined what structure tends to rotate polarized light to the left, and vice versa. It turned out that Fischer had, by chance, guessed right in naming the levorotatory and dextrorotatory forms.

  THE PARADOX OF THE BENZENE RING

  For some years after the secure establishment of the Kekulé system of structural formulas, one compound with a rather simple molecule resisted formulation. That compound was benzene (discovered in 1825 by Faraday). Chemical evidence showed it to consist of six carbon atoms and six hydrogen atoms. What happened to all the extra carbon bonds? (Six carbon atoms linked to one another by single bonds could hold fourteen hydrogen atoms, and they do in the well-known compound called hexane, C6H14.) Evidently the carbon atoms in benzene were linked by double or triple bonds. Thus, benzene might have a structure such as CH ≡ C – CH = CH – CH = CH2. But the trouble was that the known compounds with that sort of structure had properties quite different from those of benzene. Besides, all the chemical evidence seemed to indicate that the benzene molecule was very symmetrical, and six carbons and six hydrogens could not be arranged in a chain in any reasonably symmetrical fashion.

  In 1865, Kekulé himself came up with the answer. He related some years later that the vision of the
benzene molecule came to him while he was riding on a bus and sunk in a reverie, half asleep. In his dream, chains of carbon atoms seemed to come alive and dance before his eyes, and then suddenly one coiled on itself like a snake. Kekulé awoke from his reverie with a start and could have cried “Eureka!” He had the solution: the benzene molecule is a ring.

  Kekulé suggested that the six carbon atoms of the molecule are arranged as follows:

  Here at last was the required symmetry. It explained, among other things, why the substitution of another atom for one of benzene’s hydrogen atoms always yielded just one unvarying product. Since all the carbons in the ring are indistinguishable from one another in structural terms, no matter where you make the substitution for a hydrogen atom on the ring you will get the same product. Second, the ring structure showed that there are just three ways in which you can replace two hydrogen atoms on the ring: you can make the substitutions on two adjacent carbon atoms in the ring, on two separated by a single skip, or on two separated by a double skip. Sure enough, it was found that just three doubly substituted benzene isomers can be made.

  Kekulé’s blueprint of the benzene molecule, however, presented an awkward question. Generally, compounds with double bonds are more reactive, which is to say more unstable, than those with only single bonds. It is as if the extra bond were ready and more than willing to desert the attachment to the carbon atom and form a new attachment. Double-bonded compounds readily add on hydrogen or other atoms and can even be broken down without much difficulty. But the benzene ring is extraordinarily stable—more stable than carbon chains with only single bonds. (In fact, it is so stable and common in organic matter that molecules containing benzene rings make up an entire class of organic compounds, called aromatic, all the rest being lumped together as the aliphatic compounds.) The benzene molecule resists taking on more hydrogen atoms and is hard to break down.

  The nineteenth-century organic chemists could find no explanation for this queer stability of the double bonds in the benzene molecule, and it disturbed them. The point may seem a small one, but the whole Kekulé system of structural formulas was endangered by the recalcitrance of the benzene molecule. The failure to explain this one conspicuous paradox made all the rest uncertain.

  The closest approach to a solution prior to the twentieth century was that of the German chemist Johannes Thiele. In 1899, he suggested that when double bonds and single bonds alternate, the nearer ends of a pair of double bonds somehow neutralize each other and cancel each other’s reactive nature. Consider, as an example, the compound butadiene, which contains, in simplest form, the case of two double bonds separated by a single bond (conjugated double bonds). Now if two atoms are added to the compound, they add onto the end carbons, as shown in the following formula. Such a view explained the nonreactivity of benzene, since the three double bonds of the benzene rings, being arranged in a ring, neutralize each other completely.

  Some forty years later, a better answer was found by way of the new theory of chemical bonds that pictured atoms as linked by sharing electrons.

  The chemical bond, which Kekulé had drawn as a dash between atoms, came to be looked upon as representing a shared pair of electrons (see chapter 6). Each atom that forms a combination with a partner shares one of its electrons with the partner, and the partner reciprocates by donating one of its electrons to the bond. Carbon, with four electrons in its outer shell, could form four attachments; hydrogen could donate its one electron to a bond with one other atom; and so on.

  Now the question arose: How are the electrons shared? Obviously, two carbon atoms share the pair of electrons between them equally, because each atom has an equal hold on electrons. On the other hand, in a combination such as H2O, the oxygen atom, which has a stronger hold on electrons than a hydrogen atom, takes possession of the greater share of the pair of electrons it has in common with each hydrogen atom. Hence, the oxygen atom, by virtue of its excessive portion of electrons, has a slight excess of negative charge. By the same token, the hydrogen atom, suffering from an electron deficiency, has a slight excess of positive charge. A molecule containing an oxygen-hydrogen pair, such as water or ethyl alcohol, possesses a small concentration of negative charge in one part of the molecule and a smaJl concentration of positive charge in another. It possesses two poles of charge, so to speak, and is called a polar molecule.

  This view of molecular structure was first proposed in 1912 by Peter Debye (who later suggested the magnetic method of attaining very low temperatures; see chapter 6). He used an electric field to measure the amount of separation of poles of electric charge in a molecule. In such a field, polar molecules line themselves up with the negative ends pointing toward the positive pole and the positive ends toward the negative pole, and the ease with which this is done is the measure of the dipole moment of the molecule. By the early 1930s, measurements of dipole moments had become routine; and in 1936, for this and other work, Debye was awarded the Nobel Prize in chemistry.

  The new picture explained a number of things that earlier views of molecular structure could not explain—for instance, some anomalies of the boiling points of substances. In general, the greater the molecular weight, the higher the boiling point. But this rule is commonly broken. Water, with a molecular weight of only 18, boils at 100° C, whereas propane, with more than twice this molecular weight (44), boils at the much lower temperature of −42° C. Why the difference? The answer is that water is a polar molecule with a high dipole moment, while propane is nonpolar—it has no poles of charge. Polar molecules tend to orient themselves with the negative pole of one molecule adjacent to the positive pole of its neighbor. The resulting electrostatic attraction between neighboring molecules makes it harder to tear the molecules apart, and substances have relatively high boiling points. Hence, ethyl alcohol has a much higher boiling point (78° C) than its isomer dimethyl ether, which boils at −24° C, although both substances have the same molecular weight (46). Ethyl alcohol has a large dipole moment, and dimethyl ether only a small one. Water has a dipole moment even larger than that of ethyl alcohol.

  When de Broglie and Schrodinger formulated the new view of electrons not as sharply defined particles but as packets of waves (see chapter 8), the idea of the chemical bond underwent a further change. In 1939, the American chemist Linus Pauling presented a quantum-mechanical concept of molecular bonds in a book entitled The Nature of the Chemical Bond. His theory finally explained, among other things, the paradox of the stability of the benzene molecule.

  Pauling pictured the electrons that form a bond as resonating between the atoms they join. He showed that under certain conditions it is necessary to view an electron as occupying anyone of a number of positions (with varying probability). The electron, with its wavelike properties, might then best be presented as being spread out into a kind of blur, representing the weighted average of the individual probabilities of position. The more evenly the electron is spread out, the more stable the compound. Such resonance stabilization was most likely to occur when the molecule possesses conjugated bonds in one plane and when the existence of symmetry allows a number of alternative positions for the electron (viewed as a particle). The benzene ring is planar and symmetrical, and Pauling showed that the bonds of the ring were not really double and single in alternation, but that the electrons were smeared out, so to speak, into an equal distribution which results in all the bonds being alike and in all being stronger and less reactive than ordinary single bonds.

  The resonance structures, though they explain chemical behavior satisfactorily, are difficult to present in simple symbolism on paper. Therefore the old Kekulé structures, although now understood to represent only approximations of the actual electronic situation, are still universally used and will undoubtedly continue to be used through the foreseeable future.

  Organic Synthesis

  After Kolbe had produced acetic acid, there came in the 1850s a chemist who went systematically and methodically about the business of synt
hesizing organic substances in the laboratory. He was the Frenchman Pierre Eugene Marcelin Berthelot. He prepared a number of simple organic compounds from still simpler inorganic compounds such as carbon monoxide. Berthelot built his simple organic compounds up through increasing complexity until he finally had ethyl alcohol, among other things. It was synthetic ethyl alcohol, to be sure, but absolutely indistinguishable from the “real thing,” because it was the real thing.

  Ethyl alcohol is an organic compound familiar to all and highly valued by most. No doubt the thought that the chemist could make ethyl alcohol from coal, air, and water (coal to supply the carbon, air the oxygen, and water the hydrogen), without the necessity of fruits or grain as a starting point, must have created enticing visions and endowed the chemist with a new kind of reputation as a miracle worker. At any rate, it put organic synthesis on the map.

  For chemists, however, Berthelot did something even more significant. He began to form products that did not exist in nature. He took glycerol, a compound discovered by Scheele in 1778 and obtained from the breakdown of the fats of living organisms, and combined it with acids not known to occur naturally in fats (although they occur naturally elsewhere). In this way he obtained fatty substances that were not quite like those that occur in organisms.

  Thus Berthelot laid the groundwork for a new kind of organic chemistry—the synthesis of molecules that nature cannot supply. This meant the possible formation not only of a kind of synthetic which might be a substitute—perhaps an inferior substitute—for some natural compound that is hard or impossible to get in the needed quantity, but are also of synthetics which are improvements on anything in nature.