by Sam Kean
In the end, it’s probably impossible to tease out whether the heads or tails of science, the theory or the experiment, has done more to push science ahead. That’s especially true when you consider that Mendeleev made many wrong predictions. He was lucky, really, that a good scientist like Lecoq de Boisbaudran discovered eka-aluminium first. If someone had poked around for one of his mistakes—Mendeleev predicted there were many elements before hydrogen and swore the sun’s halo contained a unique element called coronium—the Russian might have died in obscurity. But just as people forgave ancient astrologers who spun false, even contradictory, horoscopes and fixated instead on the one brilliant comet they predicted exactly, people tend to remember only Mendeleev’s triumphs. Moreover, when simplifying history it’s tempting to give Mendeleev, as well as Meyer and others, too much credit. They did the important work in building the trellis on which to nail the elements; but by 1869, only two-thirds of all elements had been discovered, and for years some of them sat in the wrong columns and rows on even the best tables.
Loads of work separates a modern textbook from Mendeleev, especially regarding the mess of elements now quarantined at the bottom of the table, the lanthanides. The lanthanides start with lanthanum, element fifty-seven, and their proper home on the table baffled and bedeviled chemists well into the twentieth century. Their buried electrons cause the lanthanides to clump together in frustrating ways; sorting them out was like unknotting kudzu or ivy. Spectroscopy also stumbled with lanthanides, since even if scientists detected dozens of new bands of color, they had no idea how many new elements that translated to. Even Mendeleev, who wasn’t shy about predictions, decided the lanthanides were too vexed to make guesses about. Few elements beyond cerium, the second lanthanide, were known in 1869. But instead of chiseling in more “ekas,” Mendeleev admitted his helplessness. After cerium, he dotted his table with row after row of frustrating blanks. And later, while filling in new lanthanides after cerium, he often bungled their placement, partly because many “new” elements turned out to be combinations of known ones. It’s as if cerium was the edge of the known world to Mendeleev’s circle, just like Gibraltar was to ancient mariners, and after cerium they risked falling into a whirlpool or draining off the edge of the earth.
In truth, Mendeleev could have resolved all his frustrations had he traveled just a few hundred miles west from St. Petersburg. There, in Sweden, near where cerium was first discovered, he would have come across a nondescript porcelain mine in a hamlet with the funny name of Ytterby.
An early (sideways) periodic table produced by Dmitri Mendeleev in 1869. The huge gap after cerium (Ce) shows how little Mendeleev and his contemporaries knew about the convoluted chemistry of the rare earth metals.
In 1701, a braggadocian teenager named Johann Friedrich Böttger, ecstatic at the crowd he’d rallied with a few white lies, pulled out two silver coins for a magic show. After he waved his hands and performed chemical voodoo on them, the silver pieces “disappeared,” and a single gold piece materialized in their place. It was the most convincing display of alchemy the locals had ever seen. Böttger thought his reputation was set, and unfortunately it was.
Rumors about Böttger inevitably reached the king of Poland, Augustus the Strong, who arrested the young alchemist and locked him, Rumpelstiltskin-like, in a castle to spin gold for the king’s realm. Obviously, Böttger couldn’t deliver on this demand, and after a few futile experiments, this harmless liar, still quite young, found himself a candidate for hanging. Desperate to save his neck, Böttger begged the king to spare him. Although he’d failed with alchemy, he claimed he knew how to make porcelain.
At the time, this claim was scarcely more credible. Ever since Marco Polo had returned from China at the end of the thirteenth century, the European gentry had obsessed over white Chinese porcelain, which was hard enough to resist scratching with a nail file yet miraculously translucent like an eggshell. Empires were judged by their tea sets, and wild rumors spread about porcelain’s power. One rumor held that you couldn’t get poisoned while drinking from a porcelain cup. Another claimed the Chinese were so fabulously wealthy in porcelain that they had erected a nine-story tower of it, just to show off. (That one turned out to be true.) For centuries, powerful Europeans, like the Medici in Florence, had sponsored porcelain research but had succeeded in producing only C-minus knockoffs.
Luckily for Böttger, King Augustus had a capable man working on porcelain, Ehrenfried Walter von Tschirnhaus. Tschirnhaus, whose previous job was to sample the Polish soil to figure out where to dig for crown jewels, had just invented a special oven that reached 3,000°F. This allowed him to melt down porcelain to analyze it, and when the king ordered the clever Böttger to become Tschirnhaus’s assistant, the research took off. The duo discovered that the secret ingredients in Chinese porcelain were a white clay called kaolin and a feldspar rock that fuses into glass at high temperatures. Just as crucially, they figured out that, unlike with most crockery, they had to cook the porcelain glaze and clay simultaneously, not in separate steps. It’s this high-heat fusion of glaze and clay that gives true porcelain its lucidity and toughness. After perfecting the process, they returned, relieved, to show their liege. Augustus thanked them profusely, dreaming that porcelain would immediately make him, at least socially, the most influential monarch in Europe. And after such a breakthrough, Böttger reasonably expected his freedom. Unfortunately, the king decided he was now too valuable to release and locked him up under tighter security.
Inevitably, the secret of porcelain leaked, and Böttger and Tschirnhaus’s recipe spread throughout Europe. With the basic chemistry in place, craftsmen tinkered with and improved the process over the next half century. Soon, wherever people found feldspar, they mined it, including in frosty Scandinavia, where porcelain stoves were prized because they reached higher temperatures and held heat longer than iron-belly stoves. To feed the burgeoning industry in Europe, a feldspar mine opened a dozen miles from Stockholm, on the isle of Ytterby, in 1780.
Ytterby, pronounced “itt-er-bee” and meaning “outer village,” looks exactly like you’d hope a coastal village in Sweden would, with red-roofed houses right on the water, big white shutters, and lots of fir trees in roomy yards. People travel around the archipelago in ferries. Streets are named for minerals and elements.*
The Ytterby quarry was scooped from the top of a hill in the southeast corner of the island, and it supplied fine raw ore for porcelain and other purposes. More intriguingly for scientists, its rocks also produced exotic pigments and colored glazes when processed. Nowadays, we know that bright colors are dead giveaways of lanthanides, and the mine in Ytterby was unusually rich in them for a few geological reasons. The earth’s elements were once mixed uniformly in the crust, as if someone had dumped a whole rack of spices into a bowl and stirred it. But metal atoms, especially lanthanides, tend to move in herds, and as the molten earth churned, they clumped together. Pockets of lanthanides happened to end up near—actually beneath—Sweden. And because Scandinavia lies near a fault line, tectonic plate action in the remote past plowed the lanthanide-rich rocks up from deep underground, a process aided by Bunsen’s beloved hydrothermal vents. Finally, during the last Ice Age, extensive Scandinavian glaciers shaved off the surface of the land. This final geological event exposed the lanthanide-rich rock for easy mining near Ytterby.
But if Ytterby had the proper economic conditions to make mining profitable and the proper geology to make it scientifically worthwhile, it still needed the proper social climate. Scandinavia had barely evolved beyond a Viking mentality by the late 1600s, a century during which even its universities held witch hunts (and sorcerer hunts, for male witches) on a scale that would have embarrassed Salem. But in the 1700s, after Sweden conquered the peninsulas politically and the Swedish Enlightenment conquered it culturally, Scandinavians embraced rationalism en masse. Great scientists started popping up all out of proportion to the region’s small population. This included Johan Gad
olin, born in 1760, a chemist in a line of scientific-minded academics. (His father occupied a joint professorship in physics and theology, while his grandfather held the even more unlikely posts of physics professor and bishop.)
After extensive travel in Europe as a young man—including in England, where he befriended and toured the clay mines of porcelain maker Josiah Wedgwood—Gadolin settled down in Turku, in what is now Finland, across the Baltic Sea from Stockholm. There he earned a reputation as a geochemist. Amateur geologists began shipping unusual rocks from Ytterby to him to get his opinion, and little by little, through Gadolin’s publications, the scientific world began to hear about this remarkable little quarry.
Although he didn’t have the chemical tools (or chemical theory) to tweeze out all fourteen lanthanides, Gadolin made significant progress in isolating clusters of them. He made element hunting a pastime, even an avocation, and when, in Mendeleev’s old age, chemists with better tools revisited Gadolin’s work on the Ytterby rocks, new elements started to fall out like loose change. Gadolin had started a trend by naming one supposed element yttria, and in homage to all the elements’ common origin, chemists began to immortalize Ytterby on the periodic table. More elements (seven) trace their lineage back to Ytterby than to any other person, place, or thing. It was the inspiration for ytterbium, yttrium, terbium, and erbium. For the other three unnamed elements, before running out of letters (“rbium” doesn’t quite look right), chemists adopted holmium, after Stockholm; thulium, after the mythic name for Scandinavia; and, at Lecoq de Boisbaudran’s insistence, Gadolin’s namesake, gadolinium.
Overall, of the seven elements discovered in Ytterby, six were Mendeleev’s missing lanthanides. History might have been very different—Mendeleev reworked his table incessantly and might have filled in the entire lower realm of the table after cerium by himself—if only he’d made the trip west, across the Gulf of Finland and the Baltic Sea, to this Galápagos Island of the periodic table.
Part II
MAKING ATOMS, BREAKING ATOMS
4
Where Atoms Come From: “We Are All Star Stuff”
Where do elements come from? The commonsense view that dominated science for centuries was that they don’t come from anywhere. There was a lot of metaphysical jousting over who (or Who) might have created the cosmos and why, but the consensus was that the lifetime of every element coincides with the lifetime of the universe. They’re neither created nor destroyed: elements just are. Later theories, such as the 1930s big bang theory, folded this view into their fabric. Since the pinprick that existed back then, fourteen billion years ago, contained all the matter in the universe, everything around us must have been ejected from that speck. Not shaped like diamond tiaras and tin cans and aluminium foil quite yet, but the same basic stuff. (One scientist calculated that it took the big bang ten minutes to create all known matter, then quipped, “The elements were cooked in less time than it takes to cook a dish of duck and roast potatoes.”) Again, it’s a commonsense view—a stable astrohistory of the elements.
That theory began to fray over the next few decades. German and American scientists had proved by 1939* that the sun and other stars heated themselves by fusing hydrogen together to form helium, a process that releases an outsized amount of energy compared to the atoms’ tiny size. Some scientists said, Fine, the population of hydrogen and helium may change, but only slightly, and there’s no evidence the populations of other elements change at all. But as telescopes kept improving, more head-scratchers emerged. In theory, the big bang should have ejected elements uniformly in all directions. Yet data proved that most young stars contain only hydrogen and helium, while older stars stew with dozens of elements. Plus, extremely unstable elements such as technetium, which doesn’t exist on earth, do exist in certain classes of “chemically peculiar stars.”* Something must be forging those elements anew every day.
In the mid-1950s, a handful of perceptive astronomers realized that stars themselves are heavenly Vulcans. Though not alone, Geoffrey Burbidge, Margaret Burbidge, William Fowler, and Fred Hoyle did the most to explain the theory of stellar nucleosynthesis in a famous 1957 paper known simply, to the cognoscenti, as B2FH. Oddly for a scholarly paper, B2FH opens with two portentous and contradictory quotes from Shakespeare about whether stars govern the fate of mankind.* It goes on to argue they do. It first suggests the universe was once a primordial slurry of hydrogen, with a smattering of helium and lithium. Eventually, hydrogen clumped together into stars, and the extreme gravitational pressure inside stars began fusing hydrogen into helium, a process that fires every star in the sky. But however important cosmologically, the process is dull scientifically, since all stars do is churn out helium for billions of years. Only when the hydrogen burns up, B2FH suggests—and here is its real contribution—do things start shaking. Stars that sit bovinely for aeons, chewing hydrogen cud, are transformed more profoundly than any alchemist would have dared dream.
Desperate to maintain high temperatures, stars lacking hydrogen begin to burn and fuse helium in their cores. Sometimes helium atoms stick together completely and form even-numbered elements, and sometimes protons and neutrons spall off to make odd-numbered elements. Pretty soon appreciable amounts of lithium, boron, beryllium, and especially carbon accumulate inside stars (and only inside—the cool outer layer remains mostly hydrogen for a star’s lifetime). Unfortunately, burning helium releases less energy than burning hydrogen, so stars run through their helium in, at most, a few hundred million years. Some small stars even “die” at this point, creating molten masses of carbon known as white dwarfs. Heavier stars (eight times or so more massive than the sun) fight on, crushing carbon into six more elements, up to magnesium, which buys them a few hundred years. A few more stars perish then, but the biggest, hottest stars (whose interiors reach five billion degrees) burn those elements, too, over a few million years. B2FH traces these various fusion reactions and explains the recipe for producing everything up to iron: it’s nothing less than evolution for elements. As a result of B2FH, astronomers today can indiscriminately lump every element between lithium and iron together as stellar “metals,” and once they’ve found iron in a star, they don’t bother looking for anything smaller—once iron is spotted, it’s safe to assume the rest of the periodic table up to that point is represented.
Common sense suggests that iron atoms soon fuse in the biggest stars, and the resulting atoms fuse, forming every element down to the depths of the periodic table. But again, common sense fails. When you do the math and examine how much energy is produced per atomic union, you find that fusing anything to iron’s twenty-six protons costs energy. That means post-ferric fusion* does an energy-hungry star no good. Iron is the final peal of a star’s natural life.
So where do the heaviest elements, twenty-seven through ninety-two, cobalt through uranium, come from? Ironically, says B2FH, they emerge ready-made from mini–big bangs. After prodigally burning through elements such as magnesium and silicon, extremely massive stars (twelve times the size of the sun) burn down to iron cores in about one earth day. But before perishing, there’s an apocalyptic death rattle. Suddenly lacking the energy to, like a hot gas, keep their full volume, burned-out stars implode under their own immense gravity, collapsing thousands of miles in just seconds. In their cores, they even crush protons and electrons together into neutrons, until little but neutrons remains there. Then, rebounding from this collapse, they explode outward. And by explode, I mean explode. For one glorious month, a supernova stretches millions of miles and shines brighter than a billion stars. And during a supernova, so many gazillions of particles with so much momentum collide so many times per second that they high-jump over the normal energy barriers and fuse onto iron. Many iron nuclei end up coated in neutrons, some of which decay back into protons and thereby create new elements. Every natural combination of element and isotope spews forth from this particle blizzard.
Hundreds of millions of supernovae have gone through this
reincarnation and cataclysmic death cycle in our galaxy alone. One such explosion precipitated our solar system. About 4.6 billion years ago, a supernova sent a sonic boom through a flat cloud of space dust about fifteen billion miles wide, the remains of at least two previous stars. The dust particles commingled with the spume from the supernova, and the whole mess began to swirl in pools and eddies, like the bombarded surface of an immense pond. The dense center of the cloud boiled up into the sun (making it a cannibalized remnant of the earlier stars), and planetary bodies began to aggregate and clump together. The most impressive planets, the gas giants, formed when a stellar wind—a stream of ejecta from the sun—blew lighter elements outward toward the fringes. Among those giants, the gassiest is Jupiter, which for various reasons is a fantasy camp for elements, where they can live in forms never imagined on earth.
Since ancient times, legends about brilliant Venus, ringed Saturn, and Martian-laden Mars have pinged the human imagination. Heavenly bodies provided fodder for the naming of many elements as well. Uranus was discovered in 1781 and so excited the scientific community that, despite the fact that it contains basically zero grams of the element, a scientist named uranium after the new planet in 1789. Neptunium and plutonium sprang from this tradition as well. But of all the planets, Jupiter has had the most spectacular run in recent decades. In 1994, the Shoemaker-Levy 9 comet collided with it, the first intergalactic collision humans ever witnessed. It didn’t disappoint: twenty-one comet fragments struck home, and fireballs jumped two thousand miles high. This drama aroused the public, too, and NASA scientists were soon fending off some startling questions during open Q & A sessions online. One man asked if the core of Jupiter might be a diamond larger than the entire earth. Someone else asked what on earth Jupiter’s giant red spot had to do with “the hyper-dimensional physics [he’d] been hearing about,” the kind of physics that would make time travel possible. A few years after Shoemaker-Levy, when Jupiter’s gravity bent the spectacular Hale-Bopp comet toward earth, thirty-nine Nike-clad cultists in San Diego committed suicide because they believed that Jupiter had divinely deflected it and that it concealed a UFO that would beam them to a higher spiritual plane.
