Sam Kean
Page 21
In later decades, the region around Hannan’s Find, soon renamed Kalgoorlie, became the world’s largest gold producer. The Golden Mile, they called it, and Kalgoorlie bragged that its engineers outpaced the rest of the world when it came to extracting gold from the ground. Seems like the later generations learned their lesson—not to be throwing aside rocks all willy-nilly—after their fathers’ fool’s fool’s gold rush.
Midas’s zinc and Kalgoorlie’s tellurium were rare cases of unintentional deception: two innocent moments in monetary history surrounded by aeons of deliberate counterfeiting. A century after Midas, the first real money, coins made of a natural silver-gold alloy called electrum, appeared in Lydia, in Asia Minor. Shortly after that, another fabulously wealthy ancient ruler, the Lydian king Croesus, figured out how to separate electrum into silver and gold coins, in the process establishing a real currency system. And within a few years of Croesus’s feat, in 540 BC, King Polycrates, on the Greek isle Samos, began buying off his enemies in Sparta with lead slugs plated with gold. Ever since then, counterfeiters have used elements such as lead, copper, tin, and iron the way cheap barkeeps use water in kegs of beer—to make the real money stretch a little further.
Today counterfeiting is considered a straight case of fraud, but for most of history, a kingdom’s precious-metal currency was so tied up with its economic health that kings considered counterfeiting a high crime—treason. Those convicted of such treason faced hanging, if not worse. Counterfeiting has always attracted people who do not understand opportunity costs—the basic economic law that you can make far more money plying an honest trade than spending hundreds of hours making “free” money. Nevertheless, it has taken some brilliant minds to thwart those criminals and design anything approaching foolproof currency.
For instance, long after Isaac Newton had derived the laws of calculus and his monumental theory of gravity, he became master of the Royal Mint of England in the last few years of the 1600s. Newton, in his early fifties, just wanted a well-paying government post, but to his credit he didn’t approach it as a sinecure. Counterfeiting—especially “clipping” coins by shaving the edges and melting the scraps together to make new coins—was endemic in the seedier parts of London. The great Newton found himself entangled with spies, lowlifes, drunkards, and thieves—an entanglement he thoroughly enjoyed. A pious Christian, Newton prosecuted the wrongdoers he uncovered with the wrath of the Old Testament God, refusing pleas for clemency. He even had one notorious but slippery “coiner,” William Chaloner—who’d goaded Newton for years with accusations of fraud at the mint—hanged and publicly disemboweled.
The counterfeiting of coins dominated Newton’s tenure, but not long after he resigned, the world financial system faced new threats from fake paper currency. A Mongol emperor in China, Kublai Khan, had introduced paper money there in the 1200s. The innovation spread quickly in Asia at first—partly because Kublai Khan executed anyone who refused to use it—but only intermittently in Europe. Still, by the time the Bank of England began issuing paper notes in 1694, the advantages of paper currency had grown obvious. The ores for making coins were expensive, coins themselves were cumbersome, and the wealth based on them depended too much on unevenly distributed mineral resources. Coins also, since knowledge of metalworking was more widespread in centuries past, were easier for most people to counterfeit than paper money. (Nowadays the situation is vice versa. Anyone with a laser printer can make a decent $20 bill. Do you have a single acquaintance who could cast a passable nickel, even if such a thing were worth doing?)
If the alloy-friendly chemistry of metal coins once favored swindlers, in the age of paper money the unique chemistry of metals like europium helps governments combat swindling. It all traces back to the chemistry of europium, especially the movement of electrons within its atoms. So far we’ve discussed only electron bonds, the movement of electrons between atoms. But electrons constantly whirl around their own nuclei, too, movement often compared to planets circling a sun. Although that’s a pretty good analogy, it has a flaw if taken literally. Earth in theory could have ended up on many different orbits around the sun. Electrons cannot take any old path around a nucleus. They move within shells at different energy levels, and because there’s no energy level between one and two, or two and three, and so on, the paths of electrons are highly circumscribed: they orbit only at certain distances from their “sun” and orbit in oblong shapes at funny angles. Also unlike a planet, an electron—if excited by heat or light—can leap from its low-energy shell to an empty, high-energy shell. The electron cannot stay in the high-energy state for long, so it soon crashes back down. But this isn’t a simple back-and-forth motion, because as it crashes, the electron jettisons energy by emitting light.
The color of the emitted light depends on the relative heights of the starting and ending energy levels. A crash between closely spaced levels (such as two and one) releases a pulse of low-energy reddish light, while a crash between more widely spaced levels (say, five and two) releases high-energy purple light. Because the electrons’ options about where to jump are limited to whole-number energy levels, the emitted light is also constrained. Light emitted by electrons in atoms is not like the white light from a lightbulb. Instead, electrons emit light of very specific, very pure colors. Each element’s shells sit at different heights, so each element releases characteristic bands of color—the very bands Robert Bunsen observed with his burner and spectroscope. Later, the realization that electrons jump to whole-number levels and never orbit at fractional levels was a fundamental insight of quantum mechanics. Everything wacky you’ve ever heard about quantum mechanics derives directly or indirectly from these discontinuous leaps.
Europium can emit light as described above, but not very well: it and its brother lanthanides don’t absorb incoming light or heat efficiently (another reason chemists had trouble identifying them for so long). But light is an international currency, redeemable in many forms in the atomic world, and lanthanides can emit light in a way other than simple absorption. It’s called fluorescence,* which is familiar to most people from black lights and psychedelic posters. In general, normal emissions of light involve just electrons, but fluorescence involves whole molecules. And whereas electrons absorb and emit light of the same color (yellow in, yellow out), fluorescent molecules absorb high-energy light (ultraviolet light) but emit that energy as lower-energy, visible light. Depending on the molecule it’s attached to, europium can emit red, green, or blue light.
That versatility is a bugbear for counterfeiters and makes europium a great anticounterfeiting tool. The European Union (EU), in fact, uses its eponymous element in the ink on its paper bills. To prepare the ink, EU treasury chemists lace a fluorescing dye with europium ions, which latch onto one end of the dye molecules. (No one really knows which dyes, since the EU has reportedly outlawed looking into it. Law-abiding chemists can only guess.) Despite the anonymity, chemists know the europium dyes consist of two parts. First is the receiver, or antenna, which forms the bulk of the molecule. The antenna captures incoming light energy, which europium cannot absorb; transforms it into vibrational energy, which europium can absorb; and wriggles that energy down to the molecule’s tip. There the europium stirs up its electrons, which jump to higher energy levels. But just before the electrons jump and crash and emit, a bit of the incoming wave of energy “bounces” back into the antenna. That wouldn’t happen with isolated europium atoms, but here the bulky part of the molecule dampens the energy and dissipates it. Because of that loss, when electrons crash back down, they produce lower-energy light.
So why is that shift useful? The fluorescing dyes are selected so that europium appears dull under visible light, and a counterfeiter might be lulled into thinking he has a perfect replica. Slide a euro note beneath a special laser, though, and the laser will tickle the invisible ink. The paper itself goes black, but small, randomly oriented fibers laced with europium pop out like parti-colored constellations. The charcoal sketc
h of Europe on the bills glows green, as it might look to alien eyes from space. A pastel wreath of stars gains a corona of yellow or red, and monuments and signatures and hidden seals shine royal blue. Officials nab counterfeits simply by looking for bills that don’t show all these signs.
There are really two euros on each banknote, then: the one we see day to day and a second, hidden euro mapped directly onto the first—an embedded code. This effect is extremely hard to fake without professional training, and the europium dyes, in tandem with other security features, make the euro the most sophisticated piece of currency ever devised. Euro banknotes certainly haven’t stopped counterfeiting; that’s probably impossible as long as people like holding cash. But in the periodic table–wide struggle to slow it down, europium has taken a place among the most precious metals.
Despite all the counterfeiting, many elements have been used as legitimate currency throughout history. Some, such as antimony, were a bust. Others became money under gruesome circumstances. While working in a prison chemical plant during World War II, the Italian writer and chemist Primo Levi began stealing small sticks of cerium. Cerium sparks when struck, making it an ideal flint for cigarette lighters, and he traded the sticks to civilian workers in exchange for bread and soup. Levi came into the concentration camps fairly late, nearly starved there, and began bartering with cerium only in November 1944. He estimated that it bought him two months’ worth of rations, of life, enough to last until the Soviet army liberated his camp in January 1945. His knowledge of cerium is why we have his post-Holocaust masterpiece The Periodic Table today.
Other proposals for elemental currency were less pragmatic and more eccentric. Glenn Seaborg, caught up in nuclear enthusiasm, once suggested that plutonium would become the new gold in world finance, because it’s so valuable for nuclear applications. Perhaps as a send-up of Seaborg, a science fiction writer once suggested that radioactive waste would be a better currency for global capitalism, since coins stamped from it would certainly circulate quickly. And, of course, during every economic crisis, people bellyache about reverting to a gold or silver standard. Most countries considered paper bills the equivalent of actual gold or silver until the twentieth century, and people could freely trade the paper for the metal. Some literary scholars think that L. Frank Baum’s 1900 book The Wonderful Wizard of Oz—whose Dorothy wore silver, not ruby, slippers and traveled on a gold-colored brick road to a cash-green city—was really an allegory about the relative merits of the silver versus the gold standard.
However antiquated a metals-based economy seems, such people have a point. Although metals are quite illiquid, the metals markets are one of the most stable long-term sources of wealth. It doesn’t even have to be gold or silver. Ounce by ounce, the most valuable element, among the elements you can actually buy, is rhodium. (That’s why, to trump a mere platinum record, the Guinness Book of Records gave former Beatle Paul McCartney a disk made of rhodium in 1979 to celebrate his becoming the bestselling musician of all time.) But no one ever made more money more quickly with an element on the periodic table than the American chemist Charles Hall did with aluminium.
A number of brilliant chemists devoted their careers to aluminium throughout the 1800s, and it’s hard to judge whether the element was better or worse off afterward. A Danish chemist and a German chemist simultaneously extracted this metal from the ancient astringent alum around 1825. (Alum is the powder cartoon characters like Sylvester the cat sometimes swallow that makes their mouths pucker.) Because of its luster, mineralogists immediately classified aluminium as a precious metal, like silver or platinum, worth hundreds of dollars an ounce.
Twenty years later, a Frenchman figured out how to scale up these methods for industry, making aluminium available commercially. For a price. It was still more expensive than even gold. That’s because, despite being the most common metal in the earth’s crust—around 8 percent of it by weight, hundreds of millions of times more common than gold—aluminium never appears in pure, mother lode-al form. It’s always bonded to something, usually oxygen. Pure samples were considered miracles. The French once displayed Fort Knox–like aluminium bars next to their crown jewels, and the minor emperor Napoleon III reserved a prized set of aluminium cutlery for special guests at banquets. (Less favored guests used gold knives and forks.) In the United States, government engineers, to show off their country’s industrial prowess, capped the Washington Monument with a six-pound pyramid of aluminium in 1884. A historian reports that one ounce of shavings from the pyramid would have paid a day’s wages for each of the laborers who erected it.
Dapper engineers refurbish the aluminium cap atop the Washington Monument. The U.S. government crowned the monument with aluminium in 1884 because it was the most expensive (and therefore most impressive) metal in the world, far dearer than gold. (Bettmann/Corbis)
Aluminium’s sixty-year reign as the world’s most precious substance was glorious, but soon an American chemist ruined everything. The metal’s properties—light, strong, attractive—tantalized manufacturers, and its omnipresence in the earth’s crust had the potential to revolutionize metal production. It obsessed people, but no one could figure out an efficient way to separate it from oxygen. At Oberlin College in Ohio, a chemistry professor named Frank Fanning Jewett would regale his students with tales of the aluminium El Dorado that awaited whoever mastered this element. And at least one of his students had the naïveté to take his professor seriously.
In his later years, Professor Jewett bragged to old college chums that “my greatest discovery was the discovery of a man”—Charles Hall. Hall worked with Jewett on separating aluminium throughout his undergraduate years at Oberlin. He failed and failed and failed again, but failed a little more smartly each time. Finally, in 1886, Hall ran an electric current from handmade batteries (power lines didn’t exist) through a liquid with dissolved aluminium compounds. The energy from the current zapped and liberated the pure metal, which collected in minute silver nuggets on the bottom of the tank. The process was cheap and easy, and it would work just as well in huge vats as on the lab bench. This had been the most sought-after chemical prize since the philosopher’s stone, and Hall had found it. The “aluminium boy wonder” was just twenty-three.
Hall’s fortune, however, was not made instantly. Chemist Paul Héroult in France stumbled on more or less the same process at the same time. (Today Hall and Héroult share credit for the discovery that crashed the aluminium market.) An Austrian invented another separation method in 1887, and with the competition bearing down on Hall, he quickly founded what became the Aluminum Company of America, or Alcoa, in Pittsburgh. It turned into one of the most successful business ventures in history.
Aluminium production at Alcoa grew at exponential rates. In its first months in 1888, Alcoa eked out 50 pounds of aluminium per day; two decades later, it had to ship 88,000 pounds per day to meet the demand. And while production soared, prices plummeted. Years before Hall was born, one man’s breakthrough had dropped aluminium from $550 per pound to $18 per pound in seven years. Fifty years later, not even adjusting for inflation, Hall’s company drove down the price to 25 cents per pound. Such growth has been surpassed probably only one other time in American history, during the silicon semiconductor revolution eighty years later,* and like our latter-day computer barons, Hall cleaned up. At his death in 1914, he owned Alcoa shares worth $30 million* (around $650 million today). And thanks to Hall, aluminium became the utterly blasé metal we all know, the basis for pop cans and pinging Little League bats and airplane bodies. (A little anachronistically, it still sits atop the Washington Monument, too.) I suppose it depends on your taste and temperament whether you think aluminium was better off as the world’s most precious or most productive metal.
Incidentally, I use the international spelling “aluminium” instead of the strictly American “aluminum” throughout this book. This spelling disagreement* traces its roots back to the rapid rise of this metal. When chemists in the earl
y 1800s speculated about the existence of element thirteen, they used both spellings but eventually settled on the extra i. That spelling paralleled the recently discovered barium, magnesium, sodium, and strontium. When Charles Hall applied for patents on his electric current process, he used the extra i, too. However, when advertising his shiny metal, Hall was looser with his language. There’s debate about whether cutting the i was intentional or a fortuitous mistake on advertising fliers, but when Hall saw “aluminum,” he thought it a brilliant coinage. He dropped the vowel permanently, and with it a syllable, which aligned his product with classy platinum. His new metal caught on so quickly and grew so economically important that “aluminum” became indelibly stamped on the American psyche. As always in the United States, money talks.
14
Artistic Elements
As science grew more sophisticated throughout its history, it grew correspondingly expensive, and money, big money, began to dictate if, when, and how science got done. Already by 1956, the German-English novelist Sybille Bedford could write* that many generations had passed since “the laws of the universe were something a man might deal with pleasantly in a workshop set up behind the stables.”
Of course, very few people, mostly landed gentlemen, could have afforded a little workshop in which to do their science during the eras Bedford was pining for, the eighteenth and nineteenth centuries. To be sure, it’s no coincidence that people from the upper classes were usually the ones doing things like discovering new elements: no one else had the leisure to sit around and argue about what some obscure rocks were made of.
This mark of aristocracy lingers on the periodic table, an influence you can read without an iota of knowledge about chemistry. Gentlemen throughout Europe received educations heavy in the classics, and many element names—cerium, thorium, promethium—point to ancient myths. The really funny-looking names, too, such as praseodymium, molybdenum, and dysprosium, are amalgams of Latin and Greek. Dysprosium means “little hidden one,” since it’s tricky to separate from its brother elements. Praseodymium means “green twin” for similar reasons (its other half is neodymium, “new twin”). The names of noble gases mostly mean “stranger” or “inactive.” Even proud French gentlemen as late as the 1880s chose not “France” and “Paris” when enshrining new elements but the philologically moribund “Gallia” (gallium) and “Lutetia” (lutetium), respectively, as if sucking up to Julius Caesar.