Asimov's New Guide to Science

by Isaac Asimov

  Thanks to chromatography, purified rare-earth elements can now be prepared by the pound or even by the ton. It turns out that the rare earths are not particularly rare: the rarest of them (excepting promethium) are more common than gold or silver, and the most abundant—lanthanum, cerium, and neodymium—are more plentiful than lead. Together the rare-earth metals make up a larger percentage of the earth’s crust than copper and tin combined. So scientists have pretty well dropped the term rare earths and now call this series of elements the lanthanides, after its lead-off member. To be sure, the individual lanthanides have not been much used in the past, but the ease of separation now has multiplied their uses, and by the 1970s, 25,000,­000 pounds a year were being used. Mischmetal, a mixture that consists chiefly of cerium, lanthanum, and neodymium, makes up three-fourths the weight of cigarette-lighter flints. A mixture of the oxides is used in polishing glass, and different oxides are added to glass to produce certain desirable properties. Certain mixtures of europium and yttrium oxides are used as red-sensitive phosphors in color television, and so on.

  THE ACTINIDES

  Nor are the rewards that arise from the better understanding of the lanthanides confined to their practical uses. The new knowledge also provided a key to the chemistry of the elements at the end of the periodic table, including the synthetic ones.

  The series of heavy elements in question begins with actinium, number 89. In the table it falls under lanthanum. Actinium has two electrons in the Q-shell, just as lanthanum has two electrons in the P-shell. Actinium’s eightyninth and last electron entered the P-shell, just as lanthanum’s fifty-seventh and last entered the O-shell. Now the question is: Do the elements after actinium continue to add electrons to the P-shell and remain ordinary transition elements? Or do they, perchance, follow the pattern of the elements after lanthanum, where the electrons dive down to fill the skipped subshell below? If the latter is true, then actinium may start a new series of rare-earth metals, which would be called actinides after the first member.

  The natural elements in this series of actinides are actinium, thorium, protactinium, and uranium. They were not much studied before 1940. What little was known about their chemistry suggested that they were ordinary transition elements. But when the man-made elements neptunium and plutonium were added to the list and studied intensively, these two showed a strong chemical resemblance to uranium. Glenn Seaborg was therefore prompted to propose that the heavy elements were in fact following the lanthanide pattern and filling the buried unfilled fourth subshell of the O-shell.

  With lawrencium that subshell is filled, and the fifteen actinides exist, in perfect analogy to the fifteen lanthanides. One important confirmation is that ion-exchange chromatography separates the actinides in just the same way it separates lanthanides.

  Elements 104 (rutherfordium) and 105 (hahnium) are transactinides and, chemists are quite sure, come underneath hafnium and tantalum, the two elements that follow the lanthanides.

  Gases

  From the dawn of chemistry, it was recognized that many substances could exist in the form of gas, liquid, or solid, depending on the temperature. Water is the most common example: sufficiently cooled, it becomes solid ice; and sufficiently heated, it becomes gaseous steam. Van Helmont, who first used the word gas, differentiated between substances that are gases at ordinary temperatures, such as carbon dioxide, and those that, like steam, are gases only at elevated temperatures. He called the latter vapors, and we still speak of water vapor rather than water gas.

  The study of gases, or vapors, continued to fascinate chemists, partly because they lent themselves to quantitative studies. The rules governing their behavior were simpler and more easily worked out than those governing the behavior of liquids and solids.

  LIQUEFACTION

  In 1787, the French physicist Jacques Alexandre Cesar Charles discovered that, when a gas is cooled, each degree of cooling causes its volume to contract by about 1/273 of its volume at 0° C; and, conversely, each degree of warming causes it to expand by the same 1/273. The expansion with warmth raised no logical difficulties, but, if shrinkage with cold were to continue according to Charles’s law (as it is called to this day), at—2730 C, a gas should have shrunk to nothing! This paradox did not particularly bother chemists, for they were sure that Charles’s law would not hold all the way down, since the gases would condense to liquids as the temperature dropped, and liquids do not contract as drastically as gases do with falling temperature. Still, chemists did not, at first, have any way of getting to very low temperatures to see what actually happens.

  The development of the atomic theory, picturing gases as collections of molecules, presented the situation in new terms. The volume was now seen to depend on the velocity of the molecules. The higher the temperature, the faster they move, the more “elbow room” they require, and the greater the volume. Conversely, the lower the temperature, the more slowly they move, the less room they require, and the smaller the volume. In the 1860s, the British physicist William Thomson, who had just been raised to the peerage as Lord Kelvin, suggested that it was the molecules’ average energy content that declined by 1/273 for every degree of cooling. Whereas volume could not be expected to disappear completely, energy could. Thomson maintained that, at −273° C, the energy of molecules would sink to zero. Therefore −273° C must represent the lowest possible temperature. So this temperature (now put at −273.16° C according to refined modern measurements) would be absolute zero, or, as it is often stated, zero Kelvin. On this absolute scale, the melting point of ice is 273° K. (See figure 6.4 for the Fahrenheit, Celsius, and Kelvin scales.)

  Figure 6.4. A comparison of the Fahrenheit, Celsius (or centigrade), and Kelvin thermometric scales.

  This view made it even more certain that gases would all liquefy as absolute zero approached. With ever less energy available, the gas molecules would require so little elbow room that they would collapse upon each other and be in contact. In other words, they would become liquids, for the properties of liquids can be explained by supposing that they consist of molecules in contact, but that the molecules still contain enough energy to slip and slide freely over, under, and past each other. For that reason, liquids can pour and can easily change their shape to suit a particular container.

  As energy continues to decrease with drop in temperature, the molecules eventually possess too little to make their way past each other but come to occupy some fixed position about which they can vibrate but from which they cannot move bodily. In other words, the liquid has frozen to a solid. It seemed clear then to Kelvin that, as one approached absolute zero, all gases would not only liquefy but freeze.

  Naturally, among chemists there was a desire to demonstrate the accuracy of Kelvin’s suggestion by lowering the temperature to the point where all the gases would first liquefy, then freeze, on the way to actually attaining absolute zero. (There is something about any distant horizon that calls for conquest.)

  Scientists had been exploring extremes of coldness even before Kelvin had defined the ultimate goal. Michael Faraday had found that, even at ordinary temperatures, some gases could be liquefied under pressure. He used a strong glass tube bent into boomerang shape. In the closed bottom, he placed a substance that would yield the gas he was after. He then sealed the open end. The end with the solid material he placed into hot water, thus liberating the gas in increasingly greater quantity; and since the gas was confined within the tube, it developed increasingly greater pressure. The other end of the tube Faraday kept in a beaker filled with crushed ice. At that end the gas would be subjected to both high pressure and low temperature and would liquefy. In 1823, Faraday liquefied the gas chlorine in this manner. Chlorine’s normal liquefaction point is −34.5° C (238.7° K).

  In 1835, a French chemist, C. S. A. Thilorier, used the Faraday method to form liquid carbon dioxide under pressure, using metal cylinders, which would bear greater pressures than glass tubes. He prepared liquid carbon dioxide in considerabl
e quantity and then allowed it to escape from the tube through a narrow nozzle.

  Naturally, under these conditions, the liquid carbon dioxide, exposed to normal temperatures would evaporate quickly. When a liquid evaporates, its molecules are pulling away from those by which it is surrounded and become single entities moving freely about. The molecules of a liquid have a force of attraction among themselves, and to pull free against that attraction requires energy. If the evaporation is rapid, there is no time for sufficient energy (in the form of heat) to enter the system, and the only remaining source of energy to feed the evaporation is the liquid itself. When a liquid evaporates quickly, therefore, the temperature of the residue of the liquid drops.

  (This phenomenon is experienced by us, for the human body always perspires gently, and the evaporation of the thin layer of water on our skin withdraws heat from the skin and keeps us cool. The warmer it is, the more we must perspire; and if the air is humid so that evaporation cannot take place, the perspiration collects on our body and we become uncomfortable indeed. Exercise, by multiplying the heat-producing reactions within our body, also increases perspiration, and we are then also uncomfortable under humid conditions.)

  When Thilorier (to get back to him) allowed liquid carbon dioxide to evaporate, the temperature of the liquid dropped as evaporation proceeded, until the carbon dioxide froze. For the first time, solid carbon dioxide was formed.

  Liquid carbon dioxide is stable only under pressure. Solid carbon dioxide exposed to ordinary pressures will sublime—that is, evaporate directly to gas without melting. The sublimation point of solid carbon dioxide is −78.5° C (194.7° K).

  Solid carbon dioxide has the appearance of cloudy ice (though it is much colder); and since it does not form a liquid, it is called dry ice. Some 400,000 tons of it are produced each year, and most of it is used in preserving food through refrigeration.

  Cooling by evaporation revolutionized human life. Prior to the nineteenth century, ice, when obtainable, could be used for preserving food. Ice might be stored away in the winter and preserved, under insulation, through the summer; or it might be brought down from the mountains. At best, it was a tedious and difficult process, and most people had to make do with summer heat (or year-round heat, for that matter).

  As early as 1755, the Scottish chemist, William Cullen, had produced ice by forming a vacuum over quantities of water, enforcing rapid evaporation which cooled the water to the freezing point. This could not compete with natural ice, however. Nor could the process be used indirectly simply to cool food because ice would form and clog the pipes.

  Nowadays, an appropriate gas is liquefied by a compressor and is allowed to come to room temperature. It is then circulated in coiled pipes around a chamber in which food is contained. As it evaporates, it withdraws heat from the chamber. The gas that emerges is again liquefied by a compressor, allowed to cool, and recirculated. The process is continuous, and heat is pumped out of the enclosed chamber into the outside atmosphere. The result is a refrigerator, replacing the older icebox.

  In 1834, an American inventor, Jacob Perkins, patented (in Great Britain) the use of ether as a refrigerant. Other gases such as ammonia and sulfur dioxide also came into use. All these refrigerants had the disadvantage of being poisonous or flammable. In 1930, however, the American chemist Thomas Midgley discovered dichlorodifluoromethane (CF2Cl2), better known under the trade-name of Freon. This is nontoxic (as Midgley demonstrated by filling his lungs with it in public) and nonflammable and suits the purpose perfectly. With Freon, home refrigeration became widespread and commonplace.

  (Although Freon and other fluorocarbons have always proven totally harmless to human beings, doubts did arise, in the 1970s, about their effect on the ozonosphere, as described in the previous chapter.)

  Refrigeration applied, in moderation, to large volumes is air conditioning, so called because the air is also conditioned—that is, filtered and dehumidified. The first practical air-conditioning unit was designed in 1902 by the American inventor Willis Haviland Carrier; since the Second World War air conditioning has become nearly universal in major American cities.

  To get back to Thilorier once again, he added solid carbon dioxide to a liquid called diethyl ether (best known today as an anesthetic; see chapter 11). Diethyl ether is low-boiling and evaporates quickly. Between it and the low temperature of the solid carbon dioxide, which was subliming, a temperature of −110° C (163.2° K) was attained.

  In 1845, Faraday returned to the task of liquefying gases under the combined effect of low temperature and high pressure, making use of solid carbon dioxide and diethyl ether as his cooling mixture. Despite this mixture and his use of higher pressures than before, there were six gases he could not liquefy. They were hydrogen, oxygen, nitrogen, carbon monoxide, nitric oxide, and methane; and he named these permanent gases. To the list, we might add five more gases that Faraday did not know about. One of them was fluorine, and the other four are the noble gases: helium, neon, argon and krypton.

  In 1869, however, the Irish physicist Thomas Andrews deduced from his experiments that every gas has a critical temperature above which it cannot be liquefied even under pressure. This assumption was later put on a firm theoretical basis by the Dutch physicist, Johannes Diderik Van der Waals, who, as a result, earned the 1910 Nobel Prize for physics.

  To liquefy any gas one had to be certain, therefore, that one was working at a temperature below the critical value, or it was labor thrown out. Efforts were made to reach still lower temperatures to conquer the stubborn gases. A cascade method—lowering temperatures by steps—turned the trick. First, liquefied sulfur dioxide, cooling through evaporation, was used to liquefy carbon dioxide; then the liquid carbon dioxide was used to liquefy a more resistant gas; and so on. In 1877, the Swiss physicist Raoul Pictet finally managed to liquefy oxygen, at a temperature of −140° C (133° K) and under a pressure of 500 atmospheres (7,500 pounds per square inch). The French physicist Louis Paul Cailletet, at about the same time, liquefied not only oxygen but also nitrogen and carbon monoxide. Naturally these liquids made it possible to go on at once to still lower temperatures. The liquefaction point of oxygen at ordinary air pressure was eventually found to be −183° C (90° K); that of carbon monoxide, −190° C (83° K); and that of nitrogen, −195° C (78° K). In 1895, the English chemical engineer William Hampson and the German physicist Karl von Linde independently devised a way of liquefying air on a large scale. The air was first compressed and cooled to ordinary temperatures. It was then allowed to expand and, in the process, to become quite cold. This cold air was used to bathe a container of compressed air until it was quite cold. The compressed air was then allowed to expand so that it became much colder. This process was repeated, air getting colder and colder, until it liquefied.

  Liquid air, in quantity and cheap, was easily separated into liquid oxygen and liquid nitrogen. The oxygen could be used in blowtorches and for medicinal purposes; the nitrogen, under conditions where its inertness was useful. Thus, incandescent lightbulbs filled with nitrogen allowed the filaments to remain at white-hot temperature for longer periods before slow metal evaporation broke them, than if those same filaments were burning in evacuated bulbs. Liquid air could also be used as a source for minor components such as argon and the other noble gases.

  Hydrogen resisted all efforts at liquefaction until 1900. The Scottish chemist James Dewar then accomplished the feat by bringing a new stratagem into play. Lord Kelvin (William Thomson) and the English physicist James Prescott Joule had shown that, even in the gaseous state, a gas can be cooled simply by letting it expand and preventing heat from leaking into the gas from outside, provided the temperature is low enough to begin with. Dewar therefore cooled compressed hydrogen to a temperature of −200° C in a vessel surrounded by liquid nitrogen, let this superfrigid hydrogen expand and cool further, and repeated the cycle again and again by conducting the ever-cooling hydrogen back through pipes. The compressed hydrogen, subjected to
this Joule—Thomson effect, finally became liquid at a temperature of about −240° C (33° K). At still lower temperatures, Dewar managed to obtain solid hydrogen.

  To preserve his superfrigid liquids, he devised special silver-coated glass flasks. These were double-walled with a vacuum between. Heat could be lost (or gained) through a vacuum only by the comparatively slow process of radiation, and the silver coating reflected the incoming (or, for that matter, outgoing) radiation. Such Dewar flasks are the direct ancestor of the household Thermos bottle.

  ROCKET FUEL

  With the coming of rocketry, liquefied gases suddenly rose to new heights of glamour. Rockets require an extremely rapid chemical reaction, yielding large quantities of energy. The most convenient type of fuel is a combination of a liquid combustible, such as alcohol or kerosene, and liquid oxygen. Oxygen, or some alternate oxidizing agent, must be carried by the rocket in any case, because it runs out of any natural supply of oxygen when it leaves the atmosphere. And the oxygen must be in liquid form, since liquids are denser than gases and more oxygen can be squeezed into the fuel tanks in liquid form than in gaseous. Consequently, liquid oxygen has come into high demand in rocketry.

  The efficiency of a mixture of fuel and oxidizer is measured by a quantity known as the specific impulse. This represents the number of pounds of thrust produced by the combustion of 1 pound of the fuel-oxidizer mixture in 1 second. For a mixture of kerosene and oxygen, the specific impulse is equal to 242. Since the payload a rocket can carry depends on the specific impulse, there has been an avid search for more efficient combinations. The best liquid fuel, from this point of view, is liquid hydrogen. Combined with liquid oxygen, it can yield a specific impulse equal to 350 or so. If liquid ozone or liquid fluorine could be used in place of oxygen, the specific impulse could be raised to something like 370.