Asimov's New Guide to Science

by Isaac Asimov

  Since something like a billion dollars a year is spent over the world in the not too successful effort to keep iron and steel from corroding, the search for a general rust inhibitor goes on unabated. One interesting recent discovery is that pertechnetates (compounds containing technetium) protect iron against rusting. Of course, this rare, laboratory-made element may never be common enough to be used on any substantial scale, but it offers an invaluable research tool. Its radioactivity allows chemists to follow its fate and to observe what happens to it on the iron surface.

  One of iron’s most useful properties is its strong ferromagnetism. Iron itself is an example of a soft magnet. It is easily magnetized under the influence of an electric or magnetic field—that is, its magnetic domains (see chapter 5) are easily lined up. It is also easily demagnetized when the field is removed, and the domains fall into random orientation again. This ready loss of magnetism can be useful, as in electromagnets, where the iron core is magnetized easily with the current on, but should be as easily demagnetized when the current goes off.

  Since the Second World War, a new class of soft magnets has been developed. These are the ferrites, an example being nickel ferrite (NiFe2O4) and manganese ferrite (MnFe2O4), which are used in computers as elements that must gain or lose magnetism with the utmost ease and rapidity.

  Hard magnets, with domains that are difficult to orient or that, once oriented, to disorient, will, once magnetized, retain the property over long periods. Various steel alloys are the commonest examples, though particularly strong, hard magnets have been found among alloys that contain little or no iron. The best known example is alnico, discovered in 1931, one variety of which is made of aluminum, nickel, and cobalt (the name of the alloy being derived from the first two letters of each of the substances), plus a bit of copper.

  In the 1950s, techniques were developed to use powdered iron as a magnet, the particles being so small as to consist of individual domains. These could be oriented in molten plastic, which would then be allowed to solidify, holding the domains fixed in their orientation. Such plastic magnets are very easy to shape and mold but can be made adequately strong as well.

  NEW METALS

  We have seen in recent decades the emergence of enormously useful new metals—ones that were almost useless and even unknown up to a century or so ago and in some cases up to our own generation. The most striking example is aluminum. Aluminum is the most common of all metals—60 percent more common than iron. But it is also exceedingly difficult to extract from its ores In 1825, Hans Christian Oersted (who had discovered the connection between electricity and magnetism) separated a little aluminum in impure form. Thereafter, many chemists tried unsuccessfully to purify the metal, until in 1854 the French chemist Henri Etienne Sainte-Claire Deville finally devised a method of obtaining pure aluminum in reasonable quantities. Aluminum is so active chemically that he had to use metallic sodium (even more active) to break aluminum’s grip on its neighboring atoms. For a while aluminum sold for a hundred dollars a pound, making it practically a precious metal. Napoleon III indulged himself in aluminum cutlery and had an aluminum rattle fashioned for his infant son; and in the United States, as a mark of the nation’s great esteem for George Washington, the Washington Monument was capped with a slab of solid aluminum in 1885.

  In 1886, Charles Martin Hall, a young student of chemistry at Oberlin College, was so impressed by his professor’s statement that anyone who could discover a cheap method of making aluminum would make a fortune, that he decided to try his hand at it. In a home laboratory in his woodshed, Hall set out to apply Humphry Davy’s early discovery that an electric current sent through a molten metal can separate the metal ions by depositing them on the cathode plate. Looking for a material that could dissolve aluminum, he stumbled across cryolite, a mineral found in reasonable quantity only in Greenland. (Nowadays synthetic cryolite is available.) Hall dissolved aluminum oxide in cryolite, melted the mixture, and passed an electric current through it. Sure enough, pure aluminum collected on the cathode. Hall rushed to his professor with his first few ingots of the metal. (To this day, they are treasured by the Aluminum Company of America.)

  As it happened, a young French chemist named Paul Louis Toussaint Héroult, who was just Hall’s age (twenty-two), discovered the same process in the same year. (To complete the coincidence, Hall and Héroult both died in 1914.)

  Although the Hall-Héroult process made aluminum an inexpensive metal, it was never to be as cheap as steel, because useful aluminum ore is less common than useful iron ore, and because electricity (the key to aluminum) is more expensive than coal (the key to steel). Nevertheless, aluminum has two great advantages over steel. First, it is light—only one-third the weight of steel. Second, in aluminum, corrosion merely takes the form of a thin, transparent film over its surface, which protects deeper layers from corrosion without affecting the metal’s appearance.

  Pure aluminum is rather soft, but alloying can modify that. In 1906, the German metallurgist Alfred Wilm made a tough alloy by adding a bit of copper and a smaller bit of magnesium to the aluminum. He sold his patent rights to the Durener Metal Works in Germany, and they gave the alloy the name Duralumin.

  Engineers quickly realized the value of a light but strong metal for aircraft. After the Germans introduced Duralumin in zeppelins during the First World War, and the British learned its composition by analyzing the alloy in a crashed zeppelin, use of this new metal spread over the world. Because Duralumin was not quite as corrosion-resistant as aluminum itself, metallurgists covered it with thin sheets of pure aluminum, forming the product called Alclad.

  Today there are aluminum alloys that, weight for weight, are stronger than some steels. Aluminum has tended to replace steel wherever lightness and corrosion resistance are more important than brute strength. It has become, as everyone knows, almost a universal metal, used in airplanes, rockets, railway trains, automobiles, doors, screens, house siding, paint, kitchen utensils, foil wrapping, and so on.

  And now we have magnesium, a metal even lighter than aluminum. Its main use is in airplanes, as you might expect; as early as 1910, Germany was making use of magnesium-zinc alloys for that purpose. After the First World War, magnesium-aluminum alloys came into increasing use.

  Only about one-fourth as abundant as aluminum and more active chemically, magnesium is harder to obtain from ores. But fortunately there is a rich source in the ocean. Magnesium, unlike aluminum or iron, is present in sea water in quantity. The ocean carries dissolved matter to the amount of 3.5 percent of its mass. Of this dissolved material, 3.7 percent is magnesium ion. The ocean as a whole, therefore, contains about 2 quadrillion (2,000,­000,­000,­000,­000) tons of magnesium, or all we could use for the indefinite future.

  The problem was to get it out. The method chosen was to pump sea water into large tanks and add calcium oxide (also obtained from the sea, from oyster shells). The calcium oxide reacts with the water and the magnesium ion to form magnesium hydroxide, which is insoluble and therefore precipitates out of solution. The magnesium hydroxide is converted to magnesium chloride by treatment with hydrochloric acid, and the magnesium metal is then separated from the chlorine by means of an electric current.

  In January 1941, the Dow Chemical Company produced the first ingots of magnesium from sea water, and the stage was laid for a tenfold increase in magnesium production during the war years.

  As a matter of fact, any element that can be extracted profitably from sea water may be considered in virtually limitless supply since, after use, it eventually returns to the sea. It has been estimated that if 100 million tons of magnesium were extracted from sea water each year for a million years, the magnesium content of the ocean would drop from its present figure of 0.13 to 0.12 percent.

  If steel was the “wonder metal” of the mid-nineteenth century, aluminum of the early twentieth century, and magnesium of the mid-twentieth century, what will the next new wonder metal be? The possibilities are limited. The
re are only seven really common metals in the earth’s crust. Besides iron, aluminum, and magnesium, they are sodium, potassium, calcium, and titanium.

  Sodium, potassium, and calcium are far too active chemically to be used as construction metals. (For instance, they react violently with water.) That leaves titanium, which is about one-eighth as abundant as iron.

  Titanium has an extraordinary combination of good qualities. It is only a little more than half as heavy as steel; it is stronger, weight for weight, than aluminum or steel; it is resistant to corrosion and able to withstand high temperatures. For all these reasons, titanium is now being used in aircraft, ships, and guided missiles wherever these properties can be put to good use.

  Why was the value of titanium so slow to be discovered? The reason is much the same as for aluminum and magnesium: titanium reacts too readily with other substances and, in its impure forms—combined with oxygen or nitrogen—is an unprepossessing metal, brittle and seemingly useless. Its strength and other fine qualities emerge only when it is isolated in really pure form (in a vacuum or under an inert gas). The effort of metallurgists has succeeded to the point where a pound of titanium that would have cost $3,000 in 1947 cost $2 in 1969.

  The search need not, however, be for new wonder metals. The older metals (and some nonmetals, too) can be made far more “wonderful” than they are now.

  In Oliver Wendell Holmes’s poem “The Deacon’s Masterpiece,” the story is told of a “one-hoss shay” which was carefully made in such a way as to have no weakest point. In the end, the one-horse buggy went all at once—decomposing into a powder. But it had lasted a hundred years.

  The atomic structure of crystalline solids, both metal and nonmetal, is rather like the “one-hoss shay” situation. A metal’s crystals are riddled with submicroscopic clefts and scratches. Under pressure, a fracture will start at one of these weak points and spread through the crystal. If, like the deacon’s wonderful “one-hoss shay,” a crystal could be built with no weak points, it would have great strength.

  Such no-weak-point crystals do form as tiny fibers called whiskers on the surface of crystals. Tensile strengths of carbon whiskers have been found to run as high as 1,400 tons per square inch—or, from 15 to 70 times the tensile strength of steel. If methods could be designed for manufacturing defect-free metal in quantity, we would find ourselves with materials of astonishing strength. In 1968, for instance, Soviet scientists produced a tiny defect-free crystal of tungsten that would sustain a load of 1,635 tons per square inch, as compared with 213 tons per square inch for the best steel. And even if defect-free substances were not available in bulk, the addition of defect-free fibers to ordinary metals would reinforce and strengthen them.

  Then, too, as late as 1968, an interesting new method was found for combining metals. The two methods of historic interest were alloying, where two or more metals are melted together and form a more-or-less-homogeneous mixture, and plating, where one metal is bound firmly to another (a thin layer of expensive metal is usually bound to the surface of a bulky volume of cheaper metal, so that the surface is, for instance, as beautiful and corrosion-resistant as gold but the whole nearly as cheap as copper).

  The American metallurgist Newell C. Cook and his associates were attempting to plate a silicon layer on a platinum surface, using molten alkali fluoride as the liquid in which the platinum was immersed. The expected plating did not occur. What happened, apparently, was that the molten fluoride removed the very thin film of bound oxygen ordinarily present on even the most resistant metals, and presented the platinum surface “naked” to the silicon atoms. Instead of binding themselves to the surface on the other side of the oxygen atoms, they worked their way into the surface. The result was that a thin outer layer of the platinum became an alloy.

  Cook followed this new direction and found that many substances can be combined in this way to form a “plating” of alloy on pure metal (or on another alloy). Cook called the process metalliding and quickly showed its usefulness.

  Thus, copper to which 2 percent to 4 percent of beryllium is added in the form of an ordinary alloy, becomes extraordinarily strong. The same result can be achieved if copper is beryllided at the cost of much less of the relatively rare beryllium. Again, steel metallided with boron (boriding) is hardened. The addition of silicon, cobalt, and titanium, also produces useful properties.

  Wonder metals, in other words, if not found in nature can be created by human ingenuity.

  Chapter 7

  * * *

  The Particles

  The Nuclear Atom

  As I pointed out in the preceding chapter, it was known by 1900 that the atom was not a simple, indivisible particle but contained at least one subatomic particle—the electron, identified by J. J. Thomson. Thomson suggested that electrons were stuck like raisins in the positively charged main body of the atom.

  IDENTIFYING THE PARTICLES

  But very shortly it developed that there were other particles within the atom. When Becquerel discovered radioactivity, he identified some of the radiation emitted by radioactive substances as consisting of electrons, but other emissions were discovered as well. The Curies in France and Ernest Rutherford in England found one that was less penetrating than the electron stream. Rutherford called this radiation alpha rays and gave the electron emission the name beta rays. The flying electrons making up the latter radiation are, individually, beta particles. The alpha rays were also found to be made up of particles and these were called alpha particles. Alpha and beta are the first two letters of the Greek alphabet.

  Meanwhile the French chemist Paul Ulrich Villard discovered a third form of radioactive emission, which was named gamma rays after the third letter of the Greek alphabet. The gamma rays were quickly identified as radiation resembling X rays, but with shorter wavelengths.

  Rutherford learned by experiment that a magnetic field deflected alpha particles much less than it did beta particles. Furthermore, they were deflected in the opposite direction; hence, the alpha particle had a positive charge, as opposed to the electron’s negative one. From the amount of deflection, it could be calculated that the alpha particle must have at least twice the mass of the hydrogen ion, which possessed the smallest known positive charge. The amount of deflection would be affected both by the particle’s mass and by its charge. If the alpha particle’s positive charge was equal to that of the hydrogen ion, its mass would be two times that of the hydrogen ion; if its charge was double that, it would be four times as massive as the hydrogen ion; and so on (figure 7.1).

  Figure 7.1. DeAection of particles by a magnetic field.

  Rutherford settled the matter in 1909 by isolating alpha particles. He put some radioactive material in a thin-walled glass tube surrounded by a thick-walled glass tube, with a vacuum between. The alpha particles could penetrate the thin inner wall but not the thick outer one. They bounced back from the outer wall, so to speak, and, in so doing, lost energy and therefore were no longer able to penetrate the thin walls either. Thus they were trapped between. Now Rutherford excited the alpha particles by means of an electric discharge so that they glowed. They then showed the spectral lines of helium. (It has become evident that alpha particles produced by radioactive substances in the soil are the source of the helium in natural-gas wells.) If the alpha particle is helium, its mass must be four times that of hydrogen. Hence, its positive charge amounts to two units, taking the hydrogen ion’s charge as the unit.

  Rutherford later identified another positive particle in the atom. This one had actually been· detected, but not recognized, many years before. In 1886, the German physicist Eugen Goldstein, using a cathode-ray tube with a perforated cathode, had discovered a new radiation that streamed through the holes of the cathode in the direction opposite to the cathode rays themselves. He called it Kanalstrahlen (“channel rays”). In 1902, this radiation served as the first occasion when the Doppler-Fizeau effect (see chapter 2) was detected in any earthly source of light. The German ph
ysicist Johannes Stark placed a spectroscope in such a fashion that the rays raced toward it and demonstrated the violet shift. For this research, he was awarded the Nobel Prize for physics in 1919.

  Since channel rays move in a direction opposite to the negatively charged cathode rays, Thomson suggested that this radiation be called positive rays. It turned out that the particles of the positive rays could easily pass through matter. They were therefore judged to be much smaller in volume than ordinary ions or atoms. The amount of their deflection by a magnetic field indicated that the smallest of these particles had the same charge and mass as a hydrogen ion, assuming that this ion carried the smallest possible unit of positive charge. The positive-ray particle was therefore deduced to be the fundamental positive particle—the opposite number of the electron. Rutherford named it proton (from the Greek word for “first”).

  The proton and the electron do indeed carry equal, though opposite, electric charges, although the proton is 1,836 times as massive as the electron. It seemed likely, then, that an atom was composed of protons and electrons, mutually balancing their charges. It also appeared that the protons were in the interior of the atom, for whereas electrons can easily be peeled off, protons cannot. But now the big question was: what sort of structure do these particles of the atom form?

  THE ATOMIC NUCLEUS