Helium acts this way for a reason. All atoms contain negative particles called electrons, which reside in different tiers, or energy levels, inside the atom. The levels are nested concentrically inside each other, and each level needs a certain number of electrons to fill itself and feel satisfied. In the innermost level, that number is two. In other levels, it’s usually eight. Elements normally have equal numbers of negative electrons and positive particles called protons, so they’re electrically neutral. Electrons, however, can be freely traded between atoms, and when atoms lose or gain electrons, they form charged atoms called ions.
What’s important to know is that atoms fill their inner, lower-energy levels as full as possible with their own electrons, then either shed, share, or steal electrons to secure the right number in the outermost level. Some elements share or trade electrons diplomatically, while others act very, very nasty. That’s half of chemistry in one sentence: atoms that don’t have enough electrons in the outer level will fight, barter, beg, make and break alliances, or do whatever they must to get the right number.
Helium, element two, has exactly the number of electrons it needs to fill its only level. This “closed” configuration gives helium tremendous independence, because it doesn’t need to interact with other atoms or share or steal electrons to feel satisfied. Helium has found its erotic complement in itself. What’s more, that same configuration extends down the entire eighteenth column beneath helium—the gases neon, argon, krypton, xenon, and radon. All these elements have closed shells with full complements of electrons, so none of them reacts with anything under normal conditions. That’s why, despite all the fervid activity to identify and label elements in the 1800s—including the development of the periodic table itself—no one isolated a single gas from column eighteen before 1895. That aloofness from everyday experience, so like his ideal spheres and triangles, would have charmed Plato. And it was that sense the scientists who discovered helium and its brethren on earth were trying to evoke with the name “noble gases.” Or to put it in Plato-like words, “He who adores the perfect and unchangeable and scorns the corruptible and ignoble will prefer the noble gases, by far, to all other elements. For they never vary, never waver, never pander to other elements like hoi polloi offering cheap wares in the marketplace. They are incorruptible and ideal.”
The repose of the noble gases is rare, however. One column to the west sits the most energetic and reactive gases on the periodic table, the halogens. And if you think of the table wrapping around like a Mercator map, so that east meets west and column eighteen meets column one, even more violent elements appear on the western edge, the alkali metals. The pacifist noble gases are a demilitarized zone surrounded by unstable neighbors.
Despite being normal metals in some ways, the alkalis, instead of rusting or corroding, can spontaneously combust in air or water. They also form an alliance of interests with the halogen gases. The halogens have seven electrons in the outer layer, one short of the octet they need, while the alkalis have one electron in the outer level and a full octet in the level below. So it’s natural for the latter to dump their extra electron on the former and for the resulting positive and negative ions to form strong links.
This sort of linking happens all the time, and for this reason electrons are the most important part of an atom. They take up virtually all an atom’s space, like clouds swirling around an atom’s compact core, the nucleus. That’s true even though the components of the nucleus, protons and neutrons, are far bigger than individual electrons. If an atom were blown up to the size of a sports stadium, the proton-rich nucleus would be a tennis ball at the fifty-yard line. Electrons would be pinheads flashing around it—but flying so fast and knocking into you so many times per second that you wouldn’t be able to enter the stadium: they’d feel like a solid wall. As a result, whenever atoms touch, the buried nucleus is mute; only the electrons matter.*
One quick caveat: Don’t get too attached to the image of electrons as discrete pinheads flashing about a solid core. Or, in the more usual metaphor, don’t necessarily think of electrons as planets circling a nucleic sun. The planet analogy is useful, but as with any analogy, it’s easy to take too far, as some renowned scientists have found out to their chagrin.
Bonding between ions explains why combinations of halogens and alkalis, such as sodium chloride (table salt), are common. Similarly, elements from columns with two extra electrons, such as calcium, and elements from columns that need two extra electrons, such as oxygen, frequently align themselves. It’s the easiest way to meet everyone’s needs. Elements from nonreciprocal columns also match up according to the same laws. Two ions of sodium (Na+) take on one of oxygen (O−2) to form sodium oxide, Na2O. Calcium chloride combines as CaCl2 for the same reasons. Overall, you can usually tell at a glance how elements will combine by noting their column numbers and figuring out their charges. The pattern all falls out of the table’s pleasing left-right symmetry.
Unfortunately, not all of the periodic table is so clean and neat. But the raggedness of some elements actually makes them interesting places to visit.
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There’s an old joke about a lab assistant who bursts into a scientist’s office one morning, hysterical with joy despite a night of uninterrupted work. The assistant holds up a fizzing, hissing, corked bottle of green liquid and exclaims he has discovered a universal solvent. His sanguine boss peers at the bottle and asks, “And what is a universal solvent?” The assistant sputters, “An acid that dissolves all substances!”
After considering this thrilling news—not only would this universal acid be a scientific miracle, it would make both men billionaires—the scientist replies, “How are you holding it in a glass bottle?”
It’s a good punch line, and it’s easy to imagine Gilbert Lewis smiling, perhaps poignantly. Electrons drive the periodic table, and no one did more than Lewis to elucidate how electrons behave and form bonds in atoms. His electron work was especially illuminating for acids and bases, so he would have appreciated the assistant’s absurd claim. More personally, the punch line might have reminded Lewis how fickle scientific glory can be.
A wanderer, Lewis grew up in Nebraska, attended college and graduate school in Massachusetts around 1900, and then studied in Germany under chemist Walther Nernst. Life under Nernst proved so miserable, for legitimate and merely perceived reasons, that Lewis returned to Massachusetts for an academic post after a few months. That, too, proved unhappy, so he fled to the newly conquered Philippines to work for the U.S. government, taking with him only one book, Nernst’s Theoretical Chemistry, so he could spend years rooting out and obsessively publishing papers on every quibbling error.*
Eventually, Lewis grew homesick and settled at the University of California at Berkeley, where, over forty years, he built Berkeley’s chemistry department into the world’s best. Though that may sound like a happy ending, it wasn’t. The singular fact about Lewis is that he was probably the best scientist never to win the Nobel Prize, and he knew it. No one ever received more nominations, but his naked ambition and a trail of disputes worldwide poisoned his chances of getting enough votes. He soon began resigning (or was forced to resign) from prestigious posts in protest and became a bitter hermit.
Apart from personal reasons, Lewis never secured the Nobel Prize because his work was broad rather than deep. He never discovered one amazing thing, something you could point to and say, Wow! Instead, he spent his life refining how an atom’s electrons work in many contexts, especially the class of molecules known as acids and bases. In general, whenever atoms swap electrons to break or form new bonds, chemists say they’ve “reacted.” Acid-base reactions offer a stark and often violent example of those swaps, and Lewis’s work on acids and bases did as much as anyone’s to show what exchanging electrons means on a submicroscopic level.
Before about 1890, scientists judged acids and bases by tasting or dunking their fingers in them, not exactly the safest or most reliable methods. W
ithin a few decades, scientists realized that acids were in essence proton donors. Many acids contain hydrogen, a simple element that consists of one electron circling one proton (that’s all hydrogen has for a nucleus). When an acid like hydrochloric acid (HCl) mixes with water, it fissures into H+ and Cl−. Removing the negative electron from the hydrogen leaves just a bare proton, the H+, which swims away on its own. Weak acids like vinegar pop a few protons into solution, while strong acids like sulfuric acid flood solutions with them.
Lewis decided this definition of an acid limited scientists too much, since some substances act like acids without relying on hydrogen. So Lewis shifted the paradigm. Instead of saying that H+ splits off, he emphasized that Cl− absconds with its electron. Instead of a proton donor, then, an acid is an electron thief. In contrast, bases such as bleach or lye, which are the opposites of acids, might be called electron donors. These definitions, in addition to being more general, emphasize the behavior of electrons, which fits better with the electron-dependent chemistry of the periodic table.
Although Lewis laid this theory out in the 1920s and 1930s, scientists are still pushing the edge of how strong they can make acids using his ideas. Acid strength is measured by the pH scale, with lower numbers being stronger, and in 2005 a chemist from New Zealand invented a boron-based acid called a carborane, with a pH of –18. To put that in perspective, water has a pH of 7, and the concentrated HCl in our stomachs has a pH of 1. But according to the pH scale’s unusual accounting methods, dropping one unit (e.g., from 4 to 3) boosts an acid’s strength by ten times. So moving from stomach acid, at 1, to the boron-based acid, at –18, means the latter is ten billion billion times stronger. That’s roughly the number of atoms it would take to stack them to the moon.
There are even worse acids based on antimony, an element with probably the most colorful history on the periodic table.* Nebuchadnezzar, the king who built the Hanging Gardens of Babylon in the sixth century BC, used a noxious antimony-lead mix to paint his palace walls yellow. Perhaps not coincidentally, he soon went mad, sleeping outdoors in fields and eating grass like an ox. Around that same time, Egyptian women were applying a different form of antimony as mascara, both to decorate their faces and to give themselves witchlike powers to cast the evil eye on enemies. Later, medieval monks—not to mention Isaac Newton—grew obsessed with the sexual properties of antimony and decided this half metal, half insulator, neither one thing nor the other, was a hermaphrodite. Antimony pills also won fame as laxatives. Unlike modern pills, these hard antimony pills didn’t dissolve in the intestines, and the pills were considered so valuable that people rooted through fecal matter to retrieve and reuse them. Some lucky families even passed down laxatives from father to son. Perhaps for this reason, antimony found heavy work as a medicine, although it’s actually toxic. Mozart probably died from taking too much to combat a severe fever.
Scientists eventually got a better handle on antimony. By the 1970s, they realized that its ability to hoard electron-greedy elements around itself made it wonderful for building custom acids. The results were as astounding as the helium superfluids. Mixing antimony pentafluoride, SbF5, with hydrofluoric acid, HF, produces a substance with a pH of –31. This superacid is 100,000 billion billion billion times more potent than stomach acid and will eat through glass, as ruthlessly as water through paper. You couldn’t pick up a bottle of it because after it ate through the bottle, it would dissolve your hand. To answer the professor in the joke, it’s stored in special Teflon-lined containers.
To be honest, though, calling the antimony mix the world’s strongest acid is kind of cheating. By themselves, SbF5 (an electron thief) and HF (a proton donor) are nasty enough. But you have to sort of multiply their complementary powers together, by mixing them, before they attain superacid status. They’re strongest only under contrived circumstances. Really, the strongest solo acid is still the boron-based carborane (HCB11Cl11). And this boron acid has the best punch line so far: It’s simultaneously the world’s strongest and gentlest acid. To wrap your head around that, remember that acids split into positive and negative parts. In carborane’s case, you get H+ and an elaborate cagelike structure formed by everything else (CB11Cl11−). With most acids it’s the negative portion that’s corrosive and caustic and eats through skin. But the boron cage forms one of the most stable molecules ever invented. Its boron atoms share electrons so generously that it practically becomes helium, and it won’t go around ripping electrons from other atoms, the usual cause of acidic carnage.
So what’s carborane good for, if not dissolving glass bottles or eating through bank vaults? It can add an octane kick to gasoline, for one thing, and help make vitamins digestible. More important is its use in chemical “cradling.” Many chemical reactions involving protons aren’t clean, quick swaps. They require multiple steps, and protons get shuttled around in millionths of billionths of seconds—so quickly scientists have no idea what really happened. Carborane, though, because it’s so stable and unreactive, will flood a solution with protons, then freeze the molecules at crucial intermediate points. Carborane holds the intermediate species up on a soft, safe pillow. In contrast, antimony superacids make terrible cradles, because they shred the molecules scientists most want to look at. Lewis would have enjoyed seeing this and other applications of his work with electrons and acids, and it might have brightened the last dark years of his life. Although he did government work during World War I and made valuable contributions to chemistry until he was in his sixties, he was passed over for the Manhattan Project during World War II. This galled him, since many chemists he had recruited to Berkeley played important roles in building the first atomic bomb and became national heroes. In contrast, he puttered around during the war, reminiscing and writing a wistful pulp novel about a soldier. He died alone in his lab in 1946.
There’s general agreement that after smoking twenty-some cigars per day for forty-plus years, Lewis died of a heart attack. But it was hard not to notice that his lab smelled like bitter almonds—a sign of cyanide gas—the afternoon he died. Lewis used cyanide in his research, and it’s possible he dropped a canister of it after going into cardiac arrest. Then again, Lewis had had lunch earlier in the day—a lunch he’d initially refused to attend—with a younger, more charismatic rival chemist who had won the Nobel Prize and served as a special consultant to the Manhattan Project. It’s always been in the back of some people’s minds that the honored colleague might have unhinged Lewis. If that’s true, his facility with chemistry might have been both convenient and unfortunate.
In addition to reactive metals on its west coast and halogens and noble gases up and down its east coast, the periodic table contains a “great plains” that stretches right across its middle—columns three through twelve, the transition metals. To be honest, the transition metals have exasperating chemistry, so it’s hard to say anything about them generally—except be careful. You see, heavier atoms like the transition metals have more flexibility than other atoms in how they store their electrons. Like other atoms, they have different energy levels (designated one, two, three, etc.), with lower energy levels buried beneath higher levels. And they also fight other atoms to secure full outer energy levels with eight electrons. Figuring out what counts as the outer level, however, is trickier.
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 som
ething 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.
Sam Kean Page 2
