by Sam Kean
Nonetheless, the Drake Equation is important: it outlines what data astronomers need to collect, and it put astrobiology on a scientific foundation. Perhaps someday we’ll look back on it as we do early attempts to organize a periodic table. And with vast improvements recently in telescopes and other heavenly measuring devices, astrobiologists have tools to provide more than guesses. In fact, the Hubble Space Telescope and others have teased out so much information from so little data that astrobiologists can now go one better than Drake. They don’t have to wait for intelligent alien life to seek us out or even scour deep space for an alien Great Wall of China. They might be able to measure direct evidence of life—even mute life such as exotic plants or festering microbes—by searching for elements such as magnesium.
Obviously, magnesium is less important than oxygen or carbon, but element twelve could be a huge help for primitive creatures, allowing them to transition from organic molecules to real life. Almost all life forms use metallic elements in trace amounts to create, store, or shuttle energetic molecules around inside them. Animals primarily use the iron in hemoglobin, but the earliest and most successful forms of life, especially blue-green algae, used magnesium. Specifically, chlorophyll (probably the most important organic chemical on earth—it drives photosynthesis by converting stellar energy into sugars, the basis of the food chain) is crowned with magnesium ions at its center. Magnesium in animals helps DNA function properly.
Magnesium deposits on planets also imply the presence of liquid H2O, the most probable medium for life to arise in. Magnesium compounds sponge up water, so even on bare, rocky planets like Mars, there’s hope of finding bacteria (or bacterial fossils) among those kinds of deposits. On watery planets (like a great candidate for extraterrestrial life in our solar system, Jupiter’s moon Europa), magnesium helps keep oceans fluid. Europa has an icy outer crust, but huge liquid oceans thrive beneath it, and satellite evidence indicates that those oceans are full of magnesium salts. Like any dissolved substances, magnesium salts depress the freezing point of water so that it stays liquid at lower temperatures. Magnesium salts also stir “brine volcanism” on the rocky floors beneath the oceans. Those salts swell the volume of water they’re dissolved in, and the extra pressure from the extra volume powers volcanoes that spew brackish water and churn the ocean depths. (The pressure also cracks surface ice caps, spilling rich ice into the ocean water—which is good, in case intra-ice bubbles are important in creating life.) Moreover, magnesium compounds (among others) can provide the raw materials to build life by eroding carbon-rich chemicals from the ocean floor. Short of landing a probe or seeing alien vegetation, detecting magnesium salts on a bare, airless planet is a good sign something might be happening there bio-wise.
But let’s say Europa is barren. Even though the hunt for far-flung alien life has grown more technologically sophisticated, it still rests on one huge assumption: that the same science that controls us locally holds true in other galaxies and held true at other times. But if alpha changed over time, the consequences for potential alien life could be enormous. Historically, life perhaps couldn’t exist until alpha “relaxed” enough to allow stable carbon atoms to form—and perhaps then life arose effortlessly, without any need to appeal to a creator. And because Einstein determined that space and time are intertwined, some physicists believe that alpha variations in time could imply alpha variations across space. According to this theory, just as life arose on earth and not the moon because earth has water and an atmosphere, perhaps life arose here, on a seemingly random planet in a seemingly unremarkable pocket of space, because only here do the proper cosmological conditions exist for sturdy atoms and full molecules. This would resolve Fermi’s paradox in a cinch: nobody has come calling because nobody’s there.
At this moment, the evidence leans toward the ordinariness of earth. And based on the gravitational perturbations of far-off stars, astronomers now know of thousands of planets, which makes the odds of finding life somewhere quite good. Still, the great debate of astrobiology will be deciding whether earth, and by extension human beings, have a privileged place in the universe. Hunting for alien life will take every bit of measuring genius we have, possibly with some overlooked box on the periodic table. All we know for sure is that if some astronomer turned a telescope to a far-off star cluster tonight and found incontrovertible evidence of life, even microbial scavengers, it would be the most important discovery ever—proof that human beings are not so special after all. Except that we exist, too, and can understand and make such discoveries.
19
Above (and Beyond) the Periodic Table
There’s a conundrum near the fringes of the periodic table. Highly radioactive elements are always scarce, so you’d think, intuitively, the element that falls apart the most easily would also be the most scarce. And the element that is effaced most quickly and thoroughly whenever it appears in the earth’s crust, ultra-fragile francium, is indeed rare. Francium winks out of existence on a timescale quicker than any other natural atom—yet one element is even rarer than francium. It’s a paradox, and resolving the paradox actually requires leaving behind the comfortable confines of the periodic table. It requires setting out for what nuclear physicists consider their New World, their America to conquer—the “island of stability”—which is their best and perhaps only hope for extending the table beyond its current limitations.
As we know, 90 percent of particles in the universe are hydrogen, and the other 10 percent are helium. Everything else, including six million billion billion kilos of earth, is a cosmic rounding error. And in that six million billion billion kilos, the total amount of astatine, the scarcest natural element, is one stupid ounce. To put that into some sort of (barely) comprehensible scope, imagine that you lost your Buick Astatine in an immense parking garage and you have zero idea where it is. Imagine the tedium of walking down every row on every level past every space, looking for your vehicle. To mimic hunting for astatine atoms inside the earth, that parking garage would have to be about 100 million spaces wide, have 100 million rows, and be 100 million stories high. And there would have to be 160 identical garages just as big—and in all those buildings, there’d be just one Astatine. You’d be better off walking home.
If astatine is so rare, it’s natural to ask how scientists ever took a census of it. The answer is, they cheated a little. Any astatine present in the early earth has long since disintegrated radioactively, but other radioactive elements sometimes decay into astatine after they spit out alpha or beta particles. By knowing the total amount of the parent elements (usually elements near uranium) and calculating the odds that each of those will decay into astatine, scientists can ring up some plausible numbers for how many astatine atoms exist. This works for other elements, too. For instance, at least twenty to thirty ounces of astatine’s near neighbor on the periodic table, francium, exist at any moment.
Funnily enough, astatine is at the same time far more robust than francium. If you had a million atoms of the longest-lived type of astatine, half of them would disintegrate in four hundred minutes. A similar sample of francium would hang on for just twenty minutes. Francium is so fragile it’s basically useless, and even though there’s (barely) enough of it in the earth for chemists to detect it directly, no one will ever herd enough atoms of it together to make a visible sample. If they did, it would be so intensely radioactive it would murder them immediately. (The current flash-mob record for francium is ten thousand atoms.)
No one will likely ever produce a visible sample of astatine either, but at least it’s good for something—as a quick-acting radioisotope in medicine. In fact, after scientists—led by our old friend Emilio Segrè—identified astatine in 1939, they injected a sample into a guinea pig to study it. Because astatine sits below iodine on the periodic table, it acts like iodine in the body and so was selectively filtered and concentrated by the rodent’s thyroid gland. Astatine remains the only element whose discovery was confirmed by a nonprimate.
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br /> The odd reciprocity between astatine and francium begins in their nuclei. There, as in all atoms, two forces struggle for dominance: the strong nuclear force (which is always attractive) and the electrostatic force (which can repel particles). Though the most powerful of nature’s four fundamental forces, the strong nuclear force has ridiculously short arms. Think Tyrannosaurus rex. If particles stray more than a few trillionths of an inch apart, the strong force is impotent. For that reason, it rarely comes into play outside nuclei and black holes. Yet within its range, it’s a hundred times more muscular than the electrostatic force. That’s good, because it keeps protons and neutrons bound together instead of letting the electrostatic force wrench nuclei apart.
When you get to nuclei the size of astatine and francium, the limited reach really catches up with the strong force, and it has trouble binding all the protons and neutrons together. Francium has eighty-seven protons, none of which want to touch. Its 130-odd neutrons buffer the positive charges well but also add so much bulk that the strong force cannot reach all the way across a nucleus to quell civil strife. This makes francium (and astatine, for similar reasons) highly unstable. And it stands to reason that adding more protons would increase electric repulsion, making atoms heavier than francium even weaker.
That’s only sort of correct, though. Remember that Maria Goeppert-Mayer (“S.D. Mother Wins Nobel Prize”) developed a theory about long-lived “magic” elements—atoms with two, eight, twenty, twenty-eight, etc., protons or neutrons that were extra-stable. Other numbers of protons or neutrons, such as ninety-two, also form compact and fairly stable nuclei, where the short-leashed strong force can tighten its grip on protons. That’s why uranium is more stable than either astatine or francium, despite being heavier. As you move down the periodic table element by element, then, the struggle between the strong nuclear and electrostatic forces resembles a plummeting stock market ticker, with an overall downward trend in stability, but with many wiggles and fluctuations as one force gains the upper hand, then the other.*
Based on this prevailing pattern, scientists assumed that the elements beyond uranium would asymptotically approach a life span of 0.0. But as they groped forward with the ultraheavy elements in the 1950s and 1960s, something unexpected happened. In theory, magic numbers extend until infinity, and it turned out that there was a quasi-stable nucleus after uranium, at element 114. And instead of it being fractionally more stable, scientists at (where else?) the University of California at Berkeley calculated that 114 might survive orders of magnitude longer than the ten or so heavy elements preceding it. Given the dismally short life span of heavy elements (microseconds at best), this was a wild, counterintuitive idea. Packing neutrons and protons onto most man-made elements is like packing on explosives, since you’re putting more stress on the nucleus. Yet with element 114, packing on more TNT seemed to steady the bomb. Just as strangely, elements such as 112 and 116 seemed (on paper at least) to get horseshoes-and-kisses benefits from having close to 114 protons. Even being around that quasi-magic number calmed them. Scientists began calling this cluster of elements the island of stability.
A whimsical map of the fabled “island of stability,” a clump of superheavy elements that scientists hope will allow them to extend the periodic table far past its present bounds. Notice the stable lead (Pb) continent of the main-body periodic table, the watery trench of unstable elements, and the small, semi-stable peaks at thorium and uranium before the sea opens up. (Yuri Oganessian, Joint Institute for Nuclear Research, Dubna, Russia)
Charmed by their own metaphor, and flattering themselves as brave explorers, scientists began preparing to conquer the island. They spoke of finding an elemental “Atlantis,” and some, like old-time sailors, even produced sepia “charts” of unknown nucleic seas. (You’d half expect to see krakens drawn in the waters.) And for decades now, attempts to reach that oasis of superheavy elements have made up one of the most exciting fields of physics. Scientists haven’t reached land yet (to get truly stable, doubly magic elements, they need to figure out ways to add more neutrons to their targets), but they’re in the island’s shallows, paddling around for a harbor.
Of course, an island of stability implies a stretch of submerged stability—a stretch centered on francium. Element eighty-seven is stranded between a magic nucleus at eighty-two and a quasi-stable nucleus at ninety-two, and it’s all too tempting for its neutrons and protons to abandon ship and swim. In fact, because of the poor structural foundation of its nucleus, francium is not only the least stable natural element, it’s less stable than every synthetic element up to 104, the ungainly rutherfordium. If there’s a “trench of instability,” francium is gargling bubbles at the bottom of the Mariana.
Still, it’s more abundant than astatine. Why? Because many radioactive elements around uranium happen to decay into francium as they disintegrate. But francium, instead of doing the normal alpha decay and thereby converting itself (through the loss of two protons) into astatine, decides more than 99.9 percent of the time to relieve the pressure in its nucleus by undergoing beta decay and becoming radium. Radium then undergoes a cascade of alpha decays that leap over astatine. In other words, the path of many decaying atoms leads to a short layover on francium—hence the twenty to thirty ounces of it. At the same time, francium shuttles atoms away from astatine, causing astatine to remain rare. Conundrum solved.
Now that we’ve plumbed the trenches, what about that island of stability? It’s doubtful that chemists will ever synthesize all the elements up to very high magic numbers. But perhaps they can synthesize a stable element 114, then 126, then go from there. Some scientists believe, too, that adding electrons to extra-heavy atoms can stabilize their nuclei—the electrons might act as springs and shocks to absorb the energy that atoms normally dedicate to tearing themselves apart. If that’s so, maybe elements in the 140s, 160s, and 180s are possible. The island of stability would become a chain of islands. These stable islands would get farther apart, but perhaps, like Polynesian canoers, scientists can cross some wild distances on the new periodic archipelago.
The thrilling part is that those new elements, instead of being just heavier versions of what we already know, could have novel properties (remember how lead emerges from a lineage of carbon and silicon). According to some calculations, if electrons can tame superheavy nuclei and make them more stable, those nuclei can manipulate electrons, too—in which case, electrons might fill the atoms’ shells and orbitals in a different order. Elements whose address on the table should make them normal heavy metals might fill in their octets early and act like metallic noble gases instead.
Not to tempt the gods of hubris, but scientists already have names for those hypothetical elements. You may have noticed that the extra-heavy elements along the bottom of the table get three letters instead of two and that all of them start with u. Once again, it’s the lingering influence of Latin and Greek. As yet undiscovered element 119, Uue, is un·un·ennium; element 122, Ubb, is un·bi·bium;* and so on. Those elements will receive “real” names if they’re ever made, but for now scientists can jot them down—and mark off other elements of interest, such as magic number 184, un·oct·quadium—with Latin substitutes. (And thank goodness for them. With the impending death of the binomial species system in biology—the system that gave us Felis catus for the house cat is gradually being replaced with chromosomal DNA “bar codes,” so good-bye Homo sapiens, the knowing ape, hello TCATCGGTCATTGG…—the u elements remain about the only holdouts of once-dominant Latin in science.*)
So how far can this island-hopping extend? Can we watch little volcanoes rise beneath the periodic table forever, watch it expand and stretch down to the fittingly wide Eee, enn·enn·ennium, element 999, or even beyond? Sigh, no. Even if scientists figure out how to glue extra-heavy elements together, and even if they land smack on the farther-off islands of stability, they’ll almost certainly skid right off into the messy seas.
The reason traces back to Albert Eins
tein and the biggest failure of his career. Despite the earnest belief of most of his fans, Einstein did not win his Nobel Prize for the theory of relativity, special or general. He won for explaining a strange effect in quantum mechanics, the photoelectric effect. His solution provided the first real evidence that quantum mechanics wasn’t a crude stopgap for justifying anomalous experiments, but actually corresponds to reality. And the fact that Einstein came up with it is ironic for two reasons. One, as he got older and crustier, Einstein came to distrust quantum mechanics. Its statistical and deeply probabilistic nature sounded too much like gambling to him, and it prompted him to object that “God does not play dice with the universe.” He was wrong, and it’s too bad that most people have never heard the rejoinder by Niels Bohr: “Einstein! Stop telling God what to do.”
Second, although Einstein spent his career trying to unify quantum mechanics and relativity into a coherent and svelte “theory of everything,” he failed. Not completely, however. Sometimes when the two theories touch, they complement each other brilliantly: relativistic corrections of the speed of electrons help explain why mercury (the element I’m always looking out for) is a liquid and not the expected solid at room temperature. And no one could have created his namesake element, number ninety-nine, einsteinium, without knowledge of both theories. But overall, Einstein’s ideas on gravity, the speed of light, and relativity don’t quite fit with quantum mechanics. In some cases where the two theories come into contact, such as inside black holes, all the fancy equations break down.