Asimov's New Guide to Science

by Isaac Asimov

  For some twenty years, German chemists tackled the problem, turning out dozens of compounds, some of which were pretty good. The most successful modification was obtained in 1909, when a compound with the following formula was prepared:

  Compare this with the formula for cocaine, and you will see the similarity and also the important fact that the double ring no longer exists. This simpler molecule—stable, easy to synthesize, with good anesthetic properties and little in the way of side effects—does not exist in nature. It is a synthetic substitute far better than the real thing. It is called procaine, but is better known to the public by the trade-name Novocaine.

  Perhaps the most effective and best-known of the general pain deadeners is morphine. Its very name is from the Greek word for “sleep.” It is a purified derivative of the opium juice or laudanum used for centuries by peoples, both civilized and primitive, to combat the pains and tension of the workaday world. As a gift to the pain-wracked, it is heavenly, but it, too, carries the deadly danger of addiction. An attempt to find a substitute backfired. In 1898, a synthetic derivative, diacetylmorphine—better known as heroin—was introduced in the belief that it would be safer. Instead, it turned out to be the most dangerous drug of all.

  Less dangerous sedatives (sleep inducers) are chloral hydrate and, particularly, the barbiturates. The first example of this latter group was introduced in 1902, and they are now the most common constituents of sleeping pills. Harmless enough when used properly, they can nevertheless induce addiction, and an overdose can cause death. In fact, because death Comes quietly as the end product of a gradually deepening sleep, barbiturate overdosage is a rather popular method of suicide, or attempted suicide.

  The most common sedative, and the longest in use, is, of course, alcohol. Methods of fermenting fruit juice and grain were known in prehistoric times. Distillation to produce stronger liquors than could be formed naturally was introduced in the Middle Ages. The value of light wines in areas where the water supply is nothing but a short cut to typhoid fever and cholera, and the social acceptance of drinking in moderation, make it difficult to treat alcohol as the drug it is, although it induces addiction as surely as morphine and, through sheer quantity of use, does much more harm. Legal prohibition of sale of liquor seems to be unhelpful; certainly the American experiment of Prohibition (1920–33) was a disastrous failure. Nevertheless, alcoholism is increasingly being treated as the disease it is rather than as a moral disgrace. The acute symptoms of alcoholism (delirium tremens) are probably not so much due to the alcohol itself as to the vitamin deficiencies induced in those who eat little while drinking much.

  THE PROTOPORPHYRINS

  Man now has at his disposal all sorts of synthetics of great potential use and misuse: explosives, poison gases, insecticides, weed-killers, antiseptics, disinfectants, detergents, drugs-almost no end of them, really. But synthesis is not merely the handmaiden of consumer needs. It can also be placed at the service of pure chemical research.

  It often happens that a complex compound, produced either by living tissue or by the apparatus of the organic chemist, can be assigned only a tentative molecular structure, after all possible deductions have been drawn from the nature of the reactions it undergoes. In that case, a way out is to synthesize a compound by means of reactions designed to yield a molecular structure like the one that has been deduced. If the properties of the resulting compound are identical with the compound being investigated in the first place, a chemist can place confidence in the assigned structure.

  An impressive case in point involves hemoglobin, the main component of the red blood cells and the pigment that gives the blood its red color. In 1831, the French chemist L. R. LeCanu split hemoglobin into two parts, of which the smaller portion, called heme, made up 4 percent of the mass of hemoglobin. Heme was found to have the empirical formula C34H32O4N4Fe. Since such compounds as heme were known to occur in other vitally important substances, in both the plant and animal kingdoms, the structure of the molecule was a matter of great moment to biochemists. For nearly a century after LeCanu’s isolation of heme, however, al1that could be done was to break it down into smaller molecules. The iron atom (Fe) was easily removed, and what was left then broke up into pieces roughly one-quarter the size of the original molecule. These fragments were found to be pyrroles—molecules built on rings of five atoms, of which four are carbon and one nitrogen. Pyrrole itself has the following structure:

  The pyrroles actually obtained from heme possessed small groups of atoms containing one or two carbon atoms attached to the ring in place of one or more of the hydrogen atoms.

  In the 1920s, the German chemist Hans Fischer tackled the problem further. Since the pyrroles were one-quarter the size of the original heme, he decided to try to combine four pyrroles and see what he got. What hs finally succeeded in getting was a four-ring compound which he called porphin (from a Greek word meaning “purple,” because of its purple color). Porphin would look like this:

  However, the pyrroles obtained from heme in the first place contained small side chains attached to the ring. These remained in place when the pyrroles were joined to form porphin. The porphin with various side chains attached make up a family of compounds called the porphyrins. Those compounds that possessed the particular side chains found in heme were protoporphyrins. It was obvious to Fischer, upon comparing the properties of heme with those of the porphyrins he had synthesized, that heme (minus its iron atom) was a protoporphyrin. But which one? No fewer than fifteen different protoporphyrins (each with a different arrangement of side chains) could be formed from the various pyrroles obtained from heme, according to Fischer’s reasoning, and anyone of those fifteen might be heme itself.

  A straightforward answer could be obtained by synthesizing all fifteen and testing the properties of each one. Fischer put his students to work preparing, by painstaking chemical reactions that allowed only a particular structure to be built up, each of the fifteen possibilities. As each different protoporphyrin was formed, he compared its properties with those of the natural protoporphyrin of heme.

  In 1928, he discovered that the ninth protoporphyrin in his series was the one he was after. The natural variety of protoporphyrin is therefore called protoporphyrin IX to this day. It was a simple procedure to convert protoporphyrin IX to heme by adding iron. Chemists at last felt confident that they knew the structure of that important compound. Here is the structure of heme, as worked out by Fischer:

  For his achievement Fischer was awarded the Nobel Prize in chemistry in 1930.

  NEW PROCESSES

  All the triumphs of synthetic organic chemistry through the nineteenth century and the first half of the twentieth century, great as they were, were won by means of the same processes used by the alchemists of ancient times—mixing and heating substances. Heat was the one sure way of adding energy to molecules and making them interact, but the interactions were usually random in nature and took place by way of briefly existent, unstable intermediates, whose nature could only be guessed at.

  What chemists needed was a more refined, more direct method for producing energetic molecules—a method that would produce a group of molecules all moving at about the same speed in about the same direction. This method would remove the random nature of interactions, for whatever one molecule would do, all would do. One way would be to accelerate ions in an electric field, much as subatomic particles are accelerated in cyclotrons.

  In 1964, the German-American chemist Richard Leopold Wolfgang accelerated ions and molecules to high energies and, by means of what might be called a chemical accelerator, produced ion speeds that heat would produce only at temperatures of from 10,000° C to 100,000° C. Furthermore, the ions were all traveling in the same direction.

  If the ions so accelerated are provided with a supply of electrons they can snatch up, they will be converted to neutral molecules which will still be traveling at great speeds. Such neutral beams were produced by the American chemist Leonard Wharton in
1969.

  As to the brief intermediate stages of a chemical reaction, computers could help. It was necessary to work out the quantum-mechanical equations governing the state of the electrons in different atom-combinations and to work out the events that would take place on collision. In 1968, for instance, a computer guided by the Italian-American chemist Enrico Clementi collided ammonia and hydrochloric acid on closed-circuit television to make ammonium chloride, with the computer working out the events that. must take place. The computer indicated that the ammonium chloride that was formed could exist as a high-pressure gas at 700° C. This possibility was not previously known but was proved experimentally a few months later.

  In the last decade, chemists have developed brand-new tools, both theoretically and experimentally. Intimate details of reactions not hitherto available will be known, and new products—unattainable before or at least attainable only in small lots—will be formed. We may be at the threshold of unexpected wonders.

  Polymers and Plastics

  When we consider molecules like those of heme and quinine, we are approaching a complexity that even the modern chemist can cope with only with great difficulty. The synthesis of such a compound requires so many steps and such a variety of procedures that we can hardly expect to produce it in quantity without the help of some living organism (other than the chemist). This is nothing to make for an inferiority complex, however. Living tissue itself approaches the limit of its capacity at this level of complexity. Few molecules in nature are more complex than heme and quinine.

  To be sure, there are natural substances composed of hundreds of thousands, even millions, of atoms, but these are not really individual molecules, constructed in one piece, so to speak. Rather, these large molecules are built up of units strung together like beads in a necklace. Living tissue usually synthesizes some small, fairly simple compound and then merely hooks the units together in chains. And that, as we shall see, the chemist also is capable of doing.

  CONDENSATION AND GLUCOSE

  In living tissue, this union of small molecules (condensation) is usually accompanied by the over-all elimination of two hydrogen atoms and an oxygen atom (which combine to form a water molecule) at each point of junction. Invariably, the process can be reversed (both in the body and in the test tube): by the addition of water, the units of the chain can be loosened and separated. This reverse of condensation is called hydrolysis, from Greek words meaning “loosening through water.” In the test tube, the hydrolysis of these long chains can be hastened by a variety of methods, the most common being the addition of a certain amount of acid to the mixture.

  The first investigation of the chemical structure of a large molecule dates back to 1812, when a Russian chemist, Gottlieb Sigismund Kirchhoff, found that boiling starch with acid produced a sugar identical in properties with glucose, the sugar obtained from grapes. In 1819, the French chemist Henri Braconnot also obtained glucose by boiling various plant products such as sawdust, linen, and bark, all of which contain a compound called cellulose. It was easy to guess that both starch and cellulose were built of glucose units, but the details of the molecular structure of starch and cellulose had to await knowledge of the molecular structure of glucose. At first, before the days of structural formulas, all that was known of glucose was its empirical formula, C6H12O6· This proportion suggested that there was one water molecule, H2O, attached to each of the six carbon atoms. Hence, glucose, and compounds similar to it in structure, were called carbohydrates (“watered carbon”).

  The structural formula of glucose was worked out in 1886 by the German chemist Heinrich Kiliani. He showed that its molecule consists of a chain of six carbon atoms, to which hydrogen atoms and oxygen-hydrogen groups are separately attached. There are no intact water combinations anywhere in the molecule.

  Over the next decade or so, the German chemist Emil Fischer studied glucose in detail and worked out the exact arrangement of the oxygen-hydrogen groups around the carbon atoms, four of which were asymmetric. There are sixteen possible arrangements of these groups, and therefore sixteen possible optical isomers, each with its own properties. Chemists have, indeed, made all sixteen, only a few of which actually occur in nature. It was as a result of his work on the optical activity of these sugars that Fischer suggested the establishment of the L-series and D-series of compounds. For putting carbohydrate chemistry on a firm structural foundation, Fischer received the Nobel Prize in chemistry in 1902.

  Here are the structural formulas of glucose and of two other common sugars, fructose and galactose:

  These are the simplest structures that adequately present the asymmetries of the molecule; but in actual fact, the molecules are in the shape of nonplanar rings, each ring made up of five (sometimes four) carbon atoms and an oxygen atom.

  Once chemists knew the structure of the simple sugars, it was relatively easy to work out the manner in which they are built up into more complex compounds. For instance, a glucose molecule and a fructose can be condensed to form sucrose—the sugar we use at the table. Glucose and galactose combine to form lactose, which occurs in nature only in milk.

  There is no reason why such condensations cannot continue indefinitely, and in starch and cellulose they do. Each consists of long chains of glucose units, condensed in a particular pattern.

  The details of the pattern are important, because although both compounds are built up of the same unit, they are profoundly different. Starch in one form or another forms the major portion of humanity’s diet, while cellulose is completely inedible to human beings. The difference in the pattern of condensation, as painstakingly worked out by chemists, is analogous to the following: Suppose a glucose molecule is viewed as either right side up (when it may be symbolized as u) or upside down (symbolized as n), The starch molecule can then be viewed as consisting of a string of glucose molecules after this fashion “… uuuuuuuuu…,” while cellulose consists of “… ununununun…”. The body’s digestive juices possess the ability to hydrolyze the “uu” linkage of starch, breaking it up to glucose, which we can then absorb to obtain energy. Those same juices are helpless to touch the “un” or “nu” linkages of cellulose, and any cellulose we ingest travels through the alimentary canal and out.

  Certain microorganisms can digest cellulose, though none of the higher animals can. Some of these microorganisms live in the intestinal tracts of ruminants and termites, for instance. It is thanks to these small helpers that cows, to our advantage, can live on grass, and that termites, often to our discomfiture, can live on wood. The microorganisms form glucose from cellulose in quantity, use what they need, and the host uses the overflow. The microorganisms supply the processed food, while the host supplies the raw material and the living quarters. This form of cooperation between two forms of life for mutual benefit is called symbiosis, from Greek words meaning “life together.”

  CRYSTALLINE AND AMORPHOUS POLYMERS

  Christopher Columbus discovered South American natives playing with balls of a hardened plant juice. Columbus and the other explorers who visited South America over the next two centuries were fascinated by these bouncy balls (obtained from the sap of trees in Brazil). Samples were brought back to Europe eventually as a curiosity. About 1770, Joseph Priestley (soon to discover oxygen) found that a lump of this bouncy material would rub out pencil marks, so he invented the uninspired name of rubber, still the English word for the substance. The British called it India rubber, because it came from the “Indies” (the original name of Columbus’s new world).

  People eventually found other uses for rubber. In 1823, a Scotsman named Charles Macintosh patented garments made of a layer of rubber between two layers of cloth for use in rainy weather, and raincoats are still sometimes called mackintoshes (with an added k).

  The trouble with rubber used in this way, however, was that in warm weather it became gummy and sticky, while in cold weather it was leathery and hard. A number of individuals tried to discover ways of treating rubber so as to remov
e these undesirable characteristics. Among them was an American named Charles Goodyear, who was innocent of chemistry but worked stubbornly along by trial and error. One day in 1839, he accidentally spilled a mixture of rubber and sulfur on a hot stove. He scraped it off as quickly as he could and found, to his amazement, that the heated rubber-sulfur mixture was dry even while it was still warm. He heated it and cooled it and found that he had a sample of rubber that did not turn gummy with heat or leathery with cold but remained soft and springy throughout.

  This process of adding sulfur to rubber is now called vulcanization (after Vulcan, the Roman god of fire). Goodyear’s discovery founded the rubber industry. It is sad to have to report that Goodyear himself never reaped a reward despite this multimillion-dollar discovery. He spent his life fighting for patent rights and died deeply in debt.

  Knowledge of the molecular structure of rubber dates back to 1879, when a French chemist, Gustave Bouchardat, heated rubber in the absence of air and obtained a liquid called isoprene. Its molecule is composed of five carbon atoms and eight hydrogen atoms, arranged as follows:

  A second type of plant juice (latex), obtained from certain trees in southeast Asia, yields a substance called gutta percha. This lacks the elasticity of rubber; but, when heated in the absence of air, it, too, yields isoprene.

  Both rubber and gutta percha are made up of thousands of isoprene units. As in the case of starch and cellulose, the difference between them lies in the pattern of linkage. In rubber, the isoprene units are joined in the ”… uuuuu…” fashion and in such a way that they form coils, which can straighten out when pulled, thus allowing stretching. In gutta percha, the units join in the ”… ununununun…” fashion, and these form chains that are straighter to begin with and therefore much less stretchable (figure 11.3).