The Disappearing Spoon: And Other True Tales of Madness, Love, and the History of the World from the Periodic Table of the Elements
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As we move horizontally across the periodic table, each element has one more electron than its neighbor to the left. Sodium, element eleven, normally has eleven electrons; magnesium, element twelve, has twelve electrons; and so on. As elements swell in size, they not only sort electrons into energy levels, they also store those electrons in different-shaped bunks, called shells. But atoms, being unimaginative and conformist, fill shells and energy levels in the same order as we move across the table. Elements on the far left-hand side of the table put the first electron in an s-shell, which is spherical. It’s small and holds only two electrons—which explains the two taller columns on the left side. After those first two electrons, atoms look for something roomier. Jumping across the gap, elements in the columns on the right-hand side begin to pack new electrons one by one into a p-shell, which looks like a misshapen lung. P-shells can hold six electrons, hence the six taller columns on the right. Notice that across each row near the top, the two s-shell electrons plus the six p-shell electrons add up to eight electrons total, the number most atoms want in the outer shell. And except for the self-satisfied noble gases, all these elements’ outer-shell electrons are available to dump onto or react with other atoms. These elements behave in a logical manner: add a new electron, and the atom’s behavior should change, since it has more electrons available to participate in reactions.
Now for the frustrating part. The transition metals appear in columns three through twelve of the fourth through seventh rows, and they start to file electrons into what are called d-shells, which hold ten electrons. (D-shells look like nothing so much as misshapen balloon animals.) Based on what every other previous element has done with its shells, you’d expect the transition metals to put each extra d-shell electron on display in an outer layer and for that extra electron to be available for reactions, too. But no, transition metals squirrel their extra electrons away and prefer to hide them beneath other layers. The decision of the transition metals to violate convention and bury their d-shell electrons seems ungainly and counterintuitive—Plato would not have liked it. It’s also how nature works, and there’s not much we can do about it.
There’s a payoff to understanding this process. Normally as we move horizontally across the table, the addition of one electron to each transition metal would alter its behavior, as happens with elements in other parts of the table. But because the metals bury their d-shell electrons in the equivalent of false-bottomed drawers, those electrons end up shielded. Other atoms trying to react with the metals cannot get at those electrons, and the upshot is that many metals in a row leave the same number of electrons exposed. They therefore act the same way chemically. That’s why, scientifically, many metals look so indistinguishable and act so indistinguishably. They’re all cold, gray lumps because their outer electrons leave them no choice but to conform. (Of course, just to confuse things, sometimes buried electrons do rise up and react. That’s what causes the slight differences between some metals. That’s also why their chemistry is so exasperating.)
F-shell elements are similarly messy. F-shells begin to appear in the first of the two free-floating rows of metals beneath the periodic table, a group called the lanthanides. (They’re also called the rare earths, and according to their atomic numbers, fifty-seven through seventy-one, they really belong in the sixth row. They were relegated to the bottom to make the table skinnier and less unwieldy.) The lanthanides bury new electrons even more deeply than the transition metals, often two energy levels down. This means they are even more alike than the transition metals and can barely be distinguished from one another. Moving along the row is like driving from Nebraska to South Dakota and not realizing you’ve crossed the state line.
It’s impossible to find a pure sample of a lanthanide in nature, since its brothers always contaminate it. In one famous case, a chemist in New Hampshire tried to isolate thulium, element sixty-nine. He started with huge casserole dishes of thulium-rich ore and repeatedly treated the ore with chemicals and boiled it, a process that purified the thulium by a small fraction each time. The dissolving took so long that he could do only one or two cycles per day at first. Yet he repeated this tedious process fifteen thousand times, by hand, and winnowed the hundreds of pounds of ore down to just ounces before the purity satisfied him. Even then, there was still a little cross-contamination from other lanthanides, whose electrons were buried so deep, there just wasn’t enough of a chemical handle to grasp them and pull them out.
Electron behavior drives the periodic table. But to really understand the elements, you can’t ignore the part that makes up more than 99 percent of their mass—the nucleus. And whereas electrons obey the laws of the greatest scientist never to win the Nobel Prize, the nucleus obeys the dictates of probably the most unlikely Nobel laureate ever, a woman whose career was even more nomadic than Lewis’s.
Maria Goeppert was born in Germany in 1906. Even though her father was a sixth-generation professor, Maria had trouble convincing a Ph.D. program to admit a woman, so she bounced from school to school, taking lectures wherever she could. She finally earned her doctorate at the University of Hannover, defending her thesis in front of professors she’d never met. Not surprisingly, with no recommendations or connections, no university would hire her upon her graduation. She could enter science only obliquely, through her husband, Joseph Mayer, an American chemistry professor visiting Germany. She returned to Baltimore with him in 1930, and the newly named Goeppert-Mayer began tagging along with Mayer to work and conferences. Unfortunately, Mayer lost his job several times during the Great Depression, and the family drifted to universities in New York and then Chicago.
Most schools tolerated Goeppert-Mayer’s hanging around to chat science. Some even condescended to give her work, though they refused to pay her, and the topics were stereotypically “feminine,” such as figuring out what causes colors. After the Depression lifted, hundreds of her intellectual peers gathered for the Manhattan Project, perhaps the most vitalizing exchange of scientific ideas ever. Goeppert-Mayer received an invitation to participate, but peripherally, on a useless side project to separate uranium with flashing lights. No doubt she chafed in private, but she craved science enough to continue to work under such conditions. After World War II, the University of Chicago finally took her seriously enough to make her a professor of physics. Although she got her own office, the department still didn’t pay her.
Nevertheless, bolstered by the appointment, she began work in 1948 on the nucleus, the core and essence of an atom. Inside the nucleus, the number of positive protons—the atomic number—determines the atom’s identity. In other words, an atom cannot gain or lose protons without becoming a different element. Atoms do not normally lose neutrons either, but an element’s atoms can have different numbers of neutrons—variations called isotopes. For instance, the isotopes lead-204 and lead-206 have identical atomic numbers (82) but different numbers of neutrons (122 and 124). The atomic number plus the number of neutrons is called the atomic weight. It took scientists many years to figure out the relationship between atomic number and atomic weight, but once they did, periodic table science got a lot clearer.
Goeppert-Mayer knew all this, of course, but her work touched on a mystery that was more difficult to grasp, a deceptively simple problem. The simplest element in the universe, hydrogen, is also the most abundant. The second-simplest element, helium, is the second most abundant. In an aesthetically tidy universe, the third element, lithium, would be the third most abundant, and so on. Our universe isn’t tidy. The third most common element is oxygen, element eight. But why? Scientists might answer that oxygen has a very stable nucleus, so it doesn’t disintegrate, or “decay.” But that only pushed the question back—why do certain elements like oxygen have such stable nuclei?
Unlike most of her contemporaries, Goeppert-Mayer saw a parallel here to the incredible stability of noble gases. She suggested that protons and neutrons in the nucleus sit in shells just like electrons and that filling nuc
lear shells leads to stability. To an outsider, this seems reasonable, a nice analogy. But Nobel Prizes aren’t won on conjectures, especially those by unpaid female professors. What’s more, this idea ruffled nuclear scientists, since chemical and nuclear processes are independent. There’s no reason why dependable, stay-at-home neutrons and protons should behave like tiny, capricious electrons, which abandon their homes for attractive neighbors. And mostly they don’t.
Except Goeppert-Mayer pursued her hunch, and by piecing together a number of unlinked experiments, she proved that nuclei do have shells and do form what she called magic nuclei. For complex mathematical reasons, magic nuclei don’t reappear periodically like elemental properties. The magic happens at atomic numbers two, eight, twenty, twenty-eight, fifty, eighty-two, and so on. Goeppert-Mayer’s work proved how, at those numbers, protons and neutrons marshal themselves into highly stable, highly symmetrical spheres. Notice too that oxygen’s eight protons and eight neutrons make it doubly magic and therefore eternally stable—which explains its seeming overabundance. This model also explains at a stroke why elements such as calcium (twenty) are disproportionately plentiful and, not incidentally, why our bodies employ these readily available minerals.
Goeppert-Mayer’s theory echoes Plato’s notion that beautiful shapes are more perfect, and her model of magic, orb-shaped nuclei became the ideal form against which all nuclei are judged. Conversely, elements stranded far between two magic numbers are less abundant because they form ugly, oblong nuclei. Scientists have even discovered neutron-starved forms of holmium (element sixty-seven) that give birth to a deformed, wobbly “football nucleus.” As you might guess from Goeppert-Mayer’s model (or from ever having watched somebody fumble during a football game), the holmium footballs aren’t very steady. And unlike atoms with misbalanced electron shells, atoms with distorted nuclei can’t poach neutrons and protons from other atoms to balance themselves. So atoms with misshapen nuclei, like that form of holmium, hardly ever form and immediately disintegrate if they do.
The nuclear shell model is brilliant physics. That’s why it no doubt dismayed Goeppert-Mayer, given her precarious status among scientists, to discover that it had been duplicated by male physicists in her homeland. She risked losing credit for everything. However, both sides had produced the idea independently, and when the Germans graciously acknowledged her work and asked her to collaborate, Goeppert-Mayer’s career took off. She won her own accolades, and she and her husband moved a final time in 1959, to San Diego, where she began a real, paying job at the new University of California campus there. Still, she never quite shook the stigma of being a dilettante. When the Swedish Academy announced in 1963 that she had won her profession’s highest honor, the San Diego newspaper greeted her big day with the headline “S.D. Mother Wins Nobel Prize.”
But maybe it’s all a matter of perspective. Newspapers could have run a similarly demeaning headline about Gilbert Lewis, and he probably would have been thrilled.
Reading the periodic table across each row reveals a lot about the elements, but that’s only part of the story, and not even the best part. Elements in the same column, latitudinal neighbors, are actually far more intimately related than horizontal neighbors. People are used to reading from left to right (or right to left) in virtually every human language, but reading the periodic table up and down, column by column, as in some forms of Japanese, is actually more significant. Doing so reveals a rich subtext of relationships among elements, including unexpected rivalries and antagonisms. The periodic table has its own grammar, and reading between its lines reveals whole new stories.
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Near Twins and Black Sheep: The Genealogy of Elements
Shakespeare had a go at it with “honorificabilitudinitatibus”—which, depending on whom you ask, either means “the state of being loaded with honors” or is an anagram proclaiming that Francis Bacon, not the Bard, really wrote Shakespeare’s plays.* But that word, a mere twenty-seven letters, doesn’t stretch nearly long enough to count as the longest word in the English language.
Of course, determining the longest word is like trying to wade into a riptide. You’re likely to lose control quickly, since language is fluid and constantly changing direction. What even qualifies as English differs in different contexts. Shakespeare’s word, spoken by a clown in Love’s Labor’s Lost, obviously comes from Latin. But perhaps foreign words, even in English sentences, shouldn’t count. Plus, if you count words that do little but stack suffixes and prefixes together (“antidisestablishmentarianism,” twenty-eight letters) or nonsense words (“supercalifragilisticexpialidocious,” thirty-four letters), writers can string readers along pretty much until their hands cramp up.
But if we adopt a sensible definition—the longest word to appear in an English-language document whose purpose was not to set the record for the longest word ever—then the word we’re after appeared in 1964 in Chemical Abstracts, a dictionary-like reference source for chemists. The word describes an important protein on what historians generally count as the first virus ever discovered, in 1892—the tobacco mosaic virus. Take a breath.
acetylseryltyrosylserylisoleucylthreonylserylprolylseryl-
glutaminylphenylalanylvalylphenylalanylleucylserylseryl-
valyltryptophylalanylaspartylprolylisoleucylglutamyl-
leucylleucylasparaginylvalylcysteinylthreonylserylseryl-
leucylglycylasparaginylglutaminylphenylalanylglutami-
nylthreonylglutaminylglutaminylalanylarginylthreo-
nylthreonylglutaminylvalylglutaminylglutaminylpheny-
lalanylserylglutaminylvalyltryptophyllysylprolylphenyla-
lanylprolylglutaminylserylthreonylvalylarginylphenylala-
nylprolylglycylaspartylvalyltyrosyllysylvalyltyrosylargin-
yltyrosylasparaginylalanylvalylleucylaspartylprolylleucyli-
soleucylthreonylalanylleucylleucylglycylthreonylphenyla-
lanylaspartylthreonylarginylasparaginylarginylisoleucyli-
soleucylglutamylvalylglutamylasparaginylglu-
taminylglutaminylserylprolylthreonylthreonylalanylglutamylthreo-
nylleucylaspartylalanylthreonylarginylarginylvalylaspar-
tylaspartylalanylthreonylvalylalanylisoleucylarginylsery-
lalanylasparaginylisoleucylasparaginylleucylvalylasparagi-
nylglutamylleucylvalylarginylglycylthreonylglycylleucyl-
tyrosylasparaginylglutaminylasparaginylthreonylphenyla-
lanylglutamylserylmethionylserylglycylleucylvalyltrypto-
phylthreonylserylalanylprolylalanylserine
That anaconda runs 1,185 letters.*
Now, since none of you probably did more than run your eyes across “acetyl… serine,” go back and take a second look. You’ll notice something funny about the distribution of letters. The most common letter in English, e, appears 65 times; the uncommon letter y occurs 183 times. One letter, l, accounts for 22 percent of the word (255 instances). And the y and l don’t appear randomly but often next to each other—166 pairs, every seventh letter or so. That’s no coincidence. This long word describes a protein, and proteins are built up from the sixth (and most versatile) element on the periodic table, carbon.
Specifically, carbon forms the backbone of amino acids, which string together like beads to form proteins. (The tobacco mosaic virus protein consists of 159 amino acids.) Biochemists, because they often have so many amino acids to count, catalog them with a simple linguistic rule. They truncate the ine in amino acids such as “serine” or “isoleucine” and alter it to yl, making it fit a regular meter: “seryl” or “isoleucyl.” Taken in order, these linked yl words describe a protein’s structure precisely. Just as laypeople can see the compound word “matchbox” and grasp its meaning, biochemists in the 1950s and early 1960s gave molecules official names like “acetyl… serine” so they could reconstruct the whole molecule from the name alone. The system was exact, if exhausting. Historically, the tendency to amalgamate words reflects the strong
influence that Germany and the compound-crazy German language had on chemistry.
But why do amino acids bunch together in the first place? Because of carbon’s place on the periodic table and its need to fill its outer energy level with eight electrons—a rule of thumb called the octet rule. On the continuum of how aggressively atoms and molecules go after one another, amino acids shade toward the more civilized end. Each amino acid contains oxygen atoms on one end, a nitrogen on the other, and a trunk of two carbon atoms in the middle. (They also contain hydrogen and a branch off the main trunk that can be twenty different molecules, but those don’t concern us.) Carbon, nitrogen, and oxygen all want to get eight electrons in the outer level, but it’s easier for one of these elements than for the other. Oxygen, as element eight, has eight total electrons. Two belong to the lowest energy tier, which fills first. That leaves six left over in the outer level, so oxygen is always scouting for two additional electrons. Two electrons aren’t so hard to find, and aggressive oxygen can dictate its own terms and bully other atoms. But the same arithmetic shows that poor carbon, element six, has four electrons left over after filling its first shell and therefore needs four more to make eight. That’s harder to do, and the upshot is that carbon has really low standards for forming bonds. It latches onto virtually anything.
That promiscuity is carbon’s virtue. Unlike oxygen, carbon must form bonds with other atoms in whatever direction it can. In fact, carbon shares its electrons with up to four other atoms at once. This allows carbon to build complex chains, or even three-dimensional webs of molecules. And because it shares and cannot steal electrons, the bonds it forms are steady and stable. Nitrogen also must form multiple bonds to keep itself happy, though not to the same degree as carbon. Proteins like that anaconda described earlier simply take advantage of these elemental facts. One carbon atom in the trunk of an amino acid shares an electron with a nitrogen at the butt of another, and proteins arise when these connectible carbons and nitrogens are strung along pretty much ad infinitum, like letters in a very, very long word.