THE SAME OLD ELEMENTS John W. Campbell ========== an editorial by John W. Campbell ========== In 1925 the last naturally-occurring stable element in Earth’s mass —rhenium—was finally isolated. It wasn’t that it was difficult to handle chemically—as fluorine definitely wasn't—but that it’s an exceedingly rare element. As of 1928 the price—if you could call it a quotation!—was $10,000 per gram. As of today, the total estimated free-world reserve of rhenium is about one hundred tons—and considering the amount of free-world real estate being calculated on, that’s not very much. Hurt date of discovery, 1925, makes an interesting dividing line in time. It’s forty-four years ago; the chemical engineering students who graduated that year are just about at retirement age now——and it’s worth looking at what's happened in the span of one man’s active engineering career. We’ll stick to chemistry; 1932 is the Year of the Neutron, and we might have considered that—but nuclear physics added to the chemistry developments would be too lengthy a discussion! In two respects, nuclear physics sort of overlapped into chemical engineering; the Hanford piles produced plutonium by transmutation of uranium, and some hot-shot chemical engineering was called for to separate the brand-new elements from the old-fashioned uranium and its lethally radioactive daughters. That called for developing some completely new engineering techniques—which were, naturally, presently applied in hundreds of other chemical engineering areas! Particularly, the techniques of liquid-liquid separations, and ion-exchange reactions on a large commercial scale came in for some heavy development. The spent reactor fuel cans had to be broken down, and uranium separated from the desired product, plutonium, and from the neutron-absorbing, reactor-poisoning fission product elements. And it had to be done untouched by human hands, to a never-before achieved extent. It being a great deal easier to pump liquids around than to pick up and manipulate precipitates and ashed residues where no human being can go—all-liquid techniques were attractive. Of course it wasn’t really necessary to separate the mess of fission product substances from each other—just getting them out of the uranium, so that could be recycled for further fission reaction, was all that had to be done. But no chemist could have just let that unresolved mixture go without even looking into it. Naturally, new, and vastly superior methods of rare-earth element separations be came necessary. The original isolation of the lanthanide “rare earth” elements— which aren’t very rare, save that it was such a horrible job to separate them that it was rarely done—-had been done by fractional crystallization. It was said in the last century, when the process of discovering new rare earths was creeping on, that “generations of Ph.D.s were earned tending hundreds of evaporating dishes in the rare-earth laboratories.” It took years for the slow multiple fractional crystallizations to separate one from another. For the work facing the chemists of the Manhattan Project days, this was impossible; most of the elements they were working with wouldn’t hang around that long. They’d transmute to something else long before the crystallization was more than started. And anyone working with hundreds of those evaporating dishes definitely wouldn’t have generations. At least, not human! So the nuclear physicists did force on the chemists an engineering development of hitherto esoteric laboratory techniques—ion exchange, and liquid-liquid solution partition techniques for purifying and recovering materials. They also, of course, presented them with a terrific challenge; the problem of handling uranium hexafluoride. UF, is one of those things that exists only in equilibrium with a certain amount of free elemental fluorine. The Germans, during WWII, considered the matter of gaseous diffusion to separate U-238 from U-235, took a look at UF. and its properties, shuddered, and spent the war years looking for a more tractable uranium compound. Fluorine is a very, very nasty element. Chlorine is poisonous, and terrifically corrosive; it chews holes in stainless steel with ease. A little chlorine dissolved in water corrodes away platinum, gold, and al-most any metal of ordinary use. But chlorine has an oxidation potential of only about —1.35. Fluorine rates —2.83—it’s really corrosive. It’s also really toxic; one single bubble of fluorine escaping from a piece of apparatus has been known to kill the careless chemist. It is, incidentally, a substance that no human being will ever smell; since it reacts spontaneously and instantly with water, a human being would smell only the reaction products. Trying to handle fluorine was an enormously difficult problem. Because of its vicious habits, the first fluorocarbon compound, CF4, was not isolated until 1926, and another four years passed before a second, C2F6 was isolated. The first isolation, of fluorine, by Moisson, depended on the fact that Moisson was a truly great chemist —and his father-in-law had two important characteristics—faith in his son-in-law, and many millions of francs. Moisson used an apparatus constructed entirely of platinum-iridium alloy—density about 21; raw material costs about three hundred dollars per ounce; melting point so high that construction was exceedingly difficult and expensive. This kind of apparatus—with its ground insulator-stoppers made from calcite crystals (not even free fluorine can attack CaFj)—was not considered suitable for engineering use. Other men before Moisson appeared to have released elemental fluorine— but their apparatus seemed to have developed leaks; they all died. This made studies to develop a more economical form of fluorine generator somewhat unpopular. It was finally solved when it was found that several metals-—copper, nickel and a number- of others— were completely immune to fluorine, even hot fluorine, as aluminum is immune to oxygen. Aluminium on exposure to oxygen, , instantly acquires a dense, completely impervious coating of oxide that protects the metal from further attack. Since ordinary, metals would serve, the next step was to get insulators for the electric current that would also resist hot, nascent fluorine. Jewel-crystals, such as calcite or lithium fluoride, don’t come in large, industrial-size masses, and it was quite some while before somebody discovered that, for ‘some weird reason, not even hot fluorine would attack common Portland cement! Once fluorine became available in laboratories not supported by indulgent, multimillionaires, the pace of fluorine research had picked up—yet it was 1937 (a dozen years after the graduation of our hypothetical Chemical Engineering Class of 1925) before it was recognized that the fluorocarbons represented a whole field of chemistry—an unexplored territory. Tetrafluorethylene, CJFt, was first isolated in 1933—and the concept of fluorocarbon chemistry began to develop. Just in time to provide the fluorine-handling equipment for the Manhattan Project—the vitally necessary stoppers, insulators, flexible tubing, and gaskets that made handling the violently poisonous and viciously corrosive UF. a practicable chemical engineering job. ========== By 1945, our Chemical Engineering graduate was twenty years out of school—and there had been a few other changes in the field of chemistry. The chlorofluorocarbon gases were being used for bug bombs, paint spraying, refrigerants, and their higher-boiling brothers were breaking into the industrial cleaning field. Silicon had moved from glass to silicone rubbers, oils and greases. Meanwhile the Manhattan Project, now opening some of its secret files, was discussing elements that did not exist on Earth, because they had no stable isotopes, and were not produced by any natural radioactive decay*. Promethium, element 61, has no isotope with a half-life greater than 2.7 years; understandably it’s not around, whatever may have been its concentration when Earth first formed. Since it’s a rare-earth element of the lanthanide group, it’s not particularly missed—its chemical properties are practically identical with the rest of the group. Technitium, Tc, element 43, is more interesting; it has a reasonably long-lived isotope, Tc-97 has a half-life of 2.6 million years, and Tc-98, a half-life of 1.5 million. And it has, for chemists, an interesting characteristic. Manganese is, of course, one of the most important elements both in metallurgy, chemistry, and geology—and the odd thing about it was that it was “odd Mn out” so to speak. In the Group VII elements, the A family of Fluorine, Chlorine, Bromine and Iodine made a nearly complete family. (Only Astatine, At, the heavy radioactive halogen, was missing.) But the VII-B family was represented only by manganese; neither technitium nor rhenium was found before 1925, and the quantity of rhenium was so minute that thorough studies were almost impossible. The family characteristics of the Group VH-B elements couldn’t be determined properly… That area has now been cleared up; there’s enough technitium accumulated from fission products to make a very fair supply for chemical research. It’s available at about ninety dollars a gram in metallic form. Rhenium is now produced commercially from various flue dusts; it sells for about two dollars a gram in powder form, or eight hundred to fifteen hundred in strip form. It’s a wee bit hard to convert from powder to solid metal; rhenium is second only to tungsten in melting point—it’s still solid at 3100 ° C; its relationship to the -> platinum metals corresponds to that of manganese to the iron-nickel-cobalt group. It has a density of 21, and makes some highly useful alloys with tungsten. But at its price and natural rarity, it isn’t of wide application. Those changes over the past forty years or so are largely due to special problems, or theoretical advances. However, what’s happened to the periodic table of elements is almost as remarkable as what’s happened to organic chemistry, nuclear physics and astronomy. ========== When our Class of ‘25 chemical engineers left school, there were a lot of squares in the periodic table chart with neatly lettered names, atomic weights and numbers, and perhaps a sample stored in some , University museum somewhere. Things like zirconium and europium and gallium. In 1925, in the Group IV elements, there were three of engineering importance—carbon, tin and lead. Of course silicon, as silicon dioxide and various complex silicate rocks was important, and silicon carbide was an excellent abrasive—but what else could you do with the stuff? The other squares under Group IV contained odd names like zirconium, titanium, hafnium and thorium—and aside from the fact that thorium oxide was useful in gas mantles, that was about it. Of course the first explosive development in the Group IV elements came about when the transistor was developed—it turned out that germanium and silicon were the most practical sources of usable semiconductor materials. The Atomic Energy Commission has been interested in zirconium for some time; zirconium is not a rare element—it was just an unexploited element heretofore. The HI metal is very similar in density, appearance, and general feel to high-quality stainless steel—but unlike stainless, it doesn’t corrode in chloride solutions. And the zirconium nucleus does not absorb neutrons—which hugely interested the AEC. It makes an ideal material for cladding reactor fuel elements, and for structural members in nuclear reactors. The chemical engineer likes the stuff because of its extremely good corrosion resistance; even at about five dollars a pound it paid off for him. But zirconium has one joker; that other Group IV element, hafnium, that no one had paid any attention to. Rhenium was the last of the stable elements discovered— but it held out only by rarity. Hafnium isn’t rare, but its chemical properties are almost identical to those of zirconium—so much so that it had been masked by its more plentiful twin. And unlike the chemically near-identical zirconium, hafnium soaks up neutrons greedily. The ion-exchange techniques that worked well on rare earth elements win also work in separating Zr from Hf—but they do add to the cost. Reactor grade Zr is about 99.999 % pure Zr, but it is not sold at $5 a pound. The extracted hafnium can be used for reactor control rods, of course. However it appears in another form. In 1925, everyone knew that tungsten remained solid at a higher temperature than any other element—up to 3400°C. a temperature so high that nearly all chemical bonds break down. Since compounds break, and no element has a higher fusion temperature, clearly nothing could exist that had a higher melting point, for all the‘ elements were known… Hafnium carbide has a melting point high enough so that the compound can float around like a sort of iceberg in a sea of molten tungsten. Moreover, it forms a sort of quasi-compound with tantalum carbide that has an even higher melting point, somewhere above 3900.° Exact temperatures are a bit hard to measure when you get up in that region. And the two carbides do not break down at that temperature; alas for the theory about compounds! Titanium, the other member of Group IV that was of no practical use as of 1925 is currently being mined and processed in enormous tonnages. Fortunately there’s no danger of running out of the stuff-1-it’s one of Earth’s most common elements. No, it is not being used in such quantities as a metal—but our chemical engineer of 1925 may well have been working on it as the dioxide. It’s the whitest white pigment known; just a little tricky to manipulate, but so markedly superior that it’s well worth the tricks. The metal itself, of course, is used in supersonic planes because it can stand heat, is tough, strong, and non corroding, and is lighter than stainless steel. The metallurgists have problems with titanium metal, too—it’s got tricks that are anything but helpful. Take a nicely rolled flat sheet of titanium, and slit it into thin strips; the strips cut from the nice straight, flat sheet proceed to roll, curl, twist, and act like a dying snake. Before the tricks of the metal were learned, it ruined tools, drove machinists nuts, and terrorized metallurgists. If you think nitrogen is an inert gas—you just don’t know how it reacts to titanium. Titanium reacts more violently in nitrogen than in pure oxygen! (It also burns •nicely in pure carbon dioxide, water, or all other standard extinguishing materials.) Titanium’s excellent mechanical and chemical properties—it’s extremely resistant to corrosion, and a lot is used by chemical industry for corrosion-proof equipment— remained unknown and unusable until the late ‘30s, when processes for getting high-purity, gas-free metal became available. Zirconium, hafnium, titanium and a number of other active metals, such as vanadium, have a tremendous ability to absorb gases—hydrogen, nitrogen, oxygen, even some of the “inert” gases—into the solid metal. And the metal is thereby rendered brittle as glass. Ductile vanadium was not available until 1950, although ferro-alloys had been used for years. In 1925 there were a lot of other names-in-squares that didn’t mean anything in engineering terms. Things like selenium, tellurium, gallium, indium and niobium. Of course the rare-earth elements were just curiosities, not useful for anything except for “misch metal,” a mixture of the earth metals smelted to the metallic state and becoming popular as cigarette lighter “flints.” Who needs neodymium, europium, and the rest of that bunch? Currently there’s quite a market for europium, as you know—it’s the necessary impurity in the fluorescent crystals of your color TV set that produce a good, strong, bright red. In 1925 TV wasn’t even “just around the corner” yet. And lasers weren’t even a vague hunch in some way-out theoretician’s mind. So the potential of neodymium doped laser crystals couldn’t have been imagined. Nor could the possibilities of things like gallium arsenide solid-state lasers, cadmium selenide photocells, photoresistors, and the various selenide and telluride infrared detectors. As far back as the ‘20s, metallic sodium had been used for conducting very heavy currents over relatively short spans by running the molten metal into cast-iron pipes, and sealing the piping against air and water vapor. But the demand for immense quantities of electric power hadn’t really started then—and the copper shortage wasn’t more than hinted at. We’re now headed into an era when the conductors for power lines may not be familiar copper and aluminium —or even the newer sodium-filled polyvinyl plastic pipes. They may turn out to be related to the present superconductor magnet windings—niobium-tin intermetallic compounds. As the demand for power rises, and the supply of copper diminishes, the superconductive cables look more and more practicable, even if liquid-hydrogen temperatures are required. I’ve left out the changes in organic chemistry—the organic chemical engineering that’s exploded in the direction of plastics, of biochemicals such as antibiotics that are produced by fermentation, the immense growth of petro chemistry. You can’t cover everything in one editorial! There are still a number of unused “squares in the table”; some time we -may find that scandium is the element needed for that job. Indium, for example, doesn’t yet have any mass application—but it’s magnificent for soldering at very low temperatures. Indium itself melts at about 150°C., various alloys melt even lower, and it’s the “wettest” of all metals—i.e., it will wet almost all other metals. Mercury wets gold, silver, copper and zinc, a wide variety of metals— but it won’t wet iron, for example. Indium wets germanium and silicon; it makes a great transistor solder! Indium also has a terrific neutron absorption ability—which makes indium foils important in nuclear work. It’s a cognate of aluminum in Group III-A, along with gallium and thallium. Thallium is still one of those “squares in the table”; the only present use is as a poison for killing rats and ants. Wonder what they’ll be using it for when the Chemical Engineering Class of ‘70 is ready to retire about 2015 AD? Incidentally, the abundance of these elements is not easily determined. For a long time, it was thought that rubidium was a relatively rare element—but it’s now known to be considerably more common than chromium, zinc, or copper. It’s just so much like potassium in its behavior, and so easily overlooked, that its plenitude isn’t realized. Selenium seemed relatively rare —until people started looking, when a peculiar pattern manifested itself. It is rare in most of the world—in fact practically all the world supply of the stuff seems to be in Norm America! Here, cattle and other grazing animals are sometimes poisoned by eating plants that have grown in sulfur-poor, selenium-rich soil, producing a form of “loco weed” that’s deadly. Sulfur is necessary for growth of living cells in animals; selenium is a cognate of sulfur in Group VI-A, and metabolism tries to substitute the S. for S. It almost works. That’s why it produces monsters. Yet in Australia and New Zealand, they’ve discovered recently, “milky white disease” of sheep—a serious economic problem-—is caused by lack of selenium in the soil! The animal’s muscles degenerate, and the muscle tissue turns milky white—hence the name. The condition is reversible and preventable by simply spraying selenium compounds on the pasture land. When astrophysicists try to estimate element abundance in the Universe, they’re somewhat like a geologist, confined to studying one small island, trying to judge the composition of the planet. Evidently the relative abundance of elements varies wildly in different areas of Earth—-and gradually astrophysicists are learning that different stars—the furnaces where atoms are “seasoned to taste and cooked until done”—have different recipes for atom-cooking. There is no technitium in the Sun so far as we can see, and none in Earth—even the 2.6 mega-year isotope has run down since Earth was formed. But there are many stars which show a strong spectrum of technitium; evidently those stars are busily cooking it up in quantity… while Sol does not. Other stars are known to show, in their spectra, strong lines of platinum and the platinum metals; they may have very interesting planets! Other stars have wildly different brews of atoms, with relatively far higher proportions of elements that are barely detectable or apparently absent in Sol, and rare on Earth. One thing we can say for sure; we do not know what the relative abundance of the elements in the Universe is. It isn’t easy to account for the fact that iron, nickel and cobalt, all very near the least-energy dip of the packing fraction curve, and, therefore, representing the most stable possible nuclei, should have such wildly different abundance. Iron’s everywhere in Earth; cobalt’s relatively rare, and nickel isn’t too common. Our theories of nucleogenesis leave something to be desired— and from the startling variations in what the different “peculiar” class stars seem to be cooking up, there seems to be a large element of pure whimsy in the chemical table! There are so many, and such variable, types of stars with high abundances of unexpected elements like rhenium or platinum or technitium that a class name was developed—Peculiar is their official designation. One thing appears certain about Out There. We’ll be surprised.