by David Toomey
It’s worth noting that the common metaphor for the activity of a living cell—a time-lapse video of a day of city traffic compressed into a few frenetic minutes—hardly does justice to its speed. In an average human cell, 2,000 new proteins are created each and every second. Belgian biochemist Christian de Duve observes that the speed of metabolic processes within the cell is “beyond the powers of our imagination.”4
As complex as chores necessary to maintaining a metabolism are, they are in some ways mere prelude and preparation for the main event: reproduction. Familiar life can reproduce, of course, because cells divide. For cells with nuclei, it all begins inside the nucleus, when proteins don’t pull apart merely a section of the DNA molecule; they unwind and unzip the entire molecule along its entire 3-billion-base-pair length. Other proteins then scan one strand, make a copy, correct and repair proofreading errors, and, from material in the surrounding cytoplasm, fashion a matching strand that winds together with the copy, base locking neatly to base. Then the parent DNA, its own strands zipped up and rewound, is pulled to one side of the nucleus, the child DNA is pulled to the other, and the nucleus itself is squeezed in the middle until it splits into halves. Shortly thereafter the cell does likewise, with each half holding a nucleus. Where there was one cell, now there are two.
If the metabolic and reproductive processes in a cell are precise and fast, they are also delicate. In order to work, they need a barrier that protects and insulates them from the harsh world outside, even as they take energy and nutrients from that world and discharge waste into it. For this reason, life based in chemistry may have another requirement: a semipermeable barrier. In familiar life, that barrier is the cell membrane, a structure composed largely of lipids, the large molecules that water will not dissolve.5
In a much-cited 2001 article, microbiologist Norman Pace made a case that all life—including what we’ve been calling weird—would use the same or similar molecules to metabolize and reproduce, going so far as to argue that even the sugars used by nucleic acids, given their evolutionary advantages, are probably universal.”6 While some astrobiologists think that Pace may overvalue the advantages of those particular nucleic acid sugars and undervalue life’s ingenuity, most expect that life based in chemistry will use chemicals and processes that are roughly analogous to those used by familiar life. They also expect that if weird life doesn’t use proteins and DNA per se, it will use molecules just as large. How large? A molecule of water has a molecular weight of 18. Some protein molecules have molecular weights in the millions. A DNA molecule from a multicellular organism is likely to have a molecular weight in the billions; such a molecule—untangled, uncoiled, and stretched out—would be a meter long.
Biochemists call these gargantuan molecular assemblies “macromolecules.” These are not, we should note, haphazard constructions, and you can’t make them with just any element. Macromolecules need “backbones”—that is, long chains or sequences of the same kind of atom. Carbon, of course, serves as the molecular backbone of all familiar life. Only one other element is capable of making such backbones, and that element is silicon.
SILICON LIFE
Ideas of silicon-based life have a surprisingly long history. Among the first hypothesizers was Herbert George Wells, a figure better known today for his science fiction than his science. But Wells did have scientific training, having studied at the Normal School of Science in London under the tutelage of Darwin’s advocate T. H. Huxley. Wells was captivated by Darwin’s ideas, and he reasoned that natural selection and competition would operate for living beings anywhere. Accordingly, the Martians of his 1897 novella War of the Worlds, having evolved in a world with weaker gravity, needed prosthetics to move about on Earth. Because Wells’s Mars was Earthlike, its inhabitants had evolved with the same needs as Earth’s organisms (water among them) and with the same frailties (a vulnerability to certain bacteria).
However, Wells also imagined organisms far weirder, and he imagined them in the context not of science fiction, but of science. Taking inspiration from an address by an English chemist named Emerson Reynolds, in 1884 Wells penned a short piece making a case for life based on silicon. Momentarily, he allowed his imagination free reign, writing of “visions of silicon-aluminium organisms—why not silicon-aluminium men at once?—wandering through an atmosphere of gaseous sulphur, let us say, by the shores of a sea of liquid iron some thousand degrees or so above the temperature of a blast furnace.”7 He acknowledged that such an idea was fantasy—“merely a dream” were his words—yet he allowed that a biochemistry based in silicon and aluminum, what he called “an analogue to protoplasm,”* was possible.
Since Wells, many science fiction authors have imagined silicon-based organisms. So have a few biologists—but only a few. Until recently, most suspected that the resemblances between carbon and silicon were superficial at best, and that any putative silicon-based biochemistry would meet with at least three showstoppers.
The first was silicon’s relatively discriminating nature. Silicon was long thought to form stable bonds with only a handful of elements. The price paid for such exclusivity was that even silicon’s macromolecules—those built on backbones of silicon atoms being “silanes,” and those built on backbones of silicon atoms alternating with oxygen atoms being “silones”—were little more than endlessly repeating sequences of the same atoms, sequences that chemists, sounding rather like jaded music critics, call “monotonous.” If nature playing on a theme of carbon could compose symphonies, it seemed that with silicon she could manage only hour-long compositions on two or three notes.
The second showstopper was that, unlike carbon, silicon was generally thought unable to form double and triple bonds, and so unable to capture and channel electrical energy.
The third showstopper was that unlike carbon-based compounds, silicon-based compounds are highly reactive with many naturally occurring chemicals on Earth, including water and oxygen. Drop some carbon tetrachloride into a beaker of water, and not much happens. It will sit there and stay stable, quite literally, for years. Do the same with its cousin silicon tetrachloride, and the compound will dissolve in seconds. Or, take the simple compound of carbon and hydrogen known as methane gas, expose it to air, and you may expect them to coexist peacefully. But substitute methane’s silicon-based counterpart silane, and you’d better put on your lab goggles first, because you are guaranteed the spectacle of spontaneous combustion.
Many biologists long believed there was a fourth reason to doubt the utility of silicon for life, and it was more or less self-evident. After oxygen, silicon is the second most abundant element on Earth—a fact made evident by a glance at a map of northern Africa, western China, or the western United States. Much of Earth’s silicon, combined with oxygen, is silica—or sand. Clearly, life has all the silicon it could ask for. But with the exception of diatoms that use it in their hard cell walls and a few plants that use it in various supporting structures, familiar life has opted for carbon instead.
For decades, the case for carbon as the only viable basis for all living organisms seemed strong. Even Carl Sagan, who regularly inveighed against parochial thinking in the search for extraterrestrial life, had difficulty imagining alternatives. “When I try to think of other elements as a basis for life,” he said, “I always wind up being what I call a carbon chauvinist.”8 Exactly how carbon became a cause for chauvinism is an interesting question. If we expect to answer it, we’ll need to back up a bit.
THE STUDY OF CARBON
The periodic table, its 118 elements named and numbered, arranged by period and group, seems the very model of order—a cabinet exquisitely crafted and custom-made to contain all of nature’s essences. But any suggestion of real containment is belied by the possibilities inherent in combinations of those essences. It is those combinations, and the temperatures and pressures at which those combinations occur, that make possible ships and shoes and sealing wax, and at a somewhat higher level of complexity, cabbages and kin
gs—as well as every material thing there is, was, will be, and can be.
Exactly how do chemists know what a given combination will produce? The surprising truth is they don’t—at least not always. Frank DiSalvo, a professor of chemistry at Cornell University, admits, “Much of what we come up with we happen on by trial and error, and we can’t predict what we’ll get ahead of time.”9 Even in an age of computer-generated 3D simulations of tumbling, pirouetting macromolecules, chemistry defiantly remains a science of measuring, mixing, and waiting around to see what happens—in other words, a science of experimentation.
Naturally, there has been more experimentation with some elements than others, and so much with carbon that it is accorded its very own specialization: the field of organic chemistry. That there are more organic chemists than any other kind is certainly owed to the field’s commercial profitability. Organic chemists working in research divisions of pharmaceutical, polymer, and petrochemical firms have derived millions of compounds—and generated revenues that rival the GNP of many nations. But the number of organic chemists is also owed to chemists’ quite natural interest in the chemicals that chemists themselves are made of, and to the endlessly fascinating nature of carbon.
And so comes a question. Is carbon the spoiled child of the periodic table? Might its astonishing chemical virtuosity, while much deserving of attention, have distracted scientists from seeing or seeking the innate talents of other elements—especially, considering our interests here, those conducive to biochemistry? Many biologists think not. They suspect carbon is essential to all life that is based in chemistry. Norman Pace, in the spirit of scientific humility, says, “I would never say never,” but he nonetheless admits, “I’m not at all optimistic about finding non-carbon life.”10
There are, though, a few dissenters. One is British biochemist William Bains, currently affiliated with the University of Cambridge Institute of Biotechnology. Bains looks younger than his fifty-odd years. With his pale complexion and rimless eyeglasses, you might guess him once upon a time to have captained his school’s chess team, and to have been the sort of student who openly challenged his teachers. In fact, Bains is something of a provocateur. And he can make a case for silicon.
THE CASE FOR SILICON
Bains can answer the arguments against silicon point by point. He admits that some silanes are indeed “monotonous” as charged, but he notes that the right kind of monotony can work wonders of complexity and variation. After all, the 3 billion base pairs of the DNA molecule itself, that very template of all genetic diversity, are built from only four nucleic acids. Moreover, Bains says, the claim that silicon can bind with only a few elements amounts to “a myth.”11 It is held as a truth by many astrobiologists only because they are unaware of recent work in chemical engineering, work showing that under certain circumstances—most notably very cold temperatures—silicon can bond stably with many elements.12
The stability of any given chemical, of course, is directly tied to its temperature. For a living organism to continue to be living, its biochemistry must be stable enough that its cellular structures hold together, but not so stable that nothing moves—the latter condition being a fair description of death. A working biochemistry is a bit like a juggler. A juggler can’t be too stable; he has to juggle, after all. Neither can he be unstable; that is, he can’t drop a bowling pin or let one go flying off in the direction of the audience. To keep his audience interested, he must always seem about to miss a catch, without ever actually missing a catch. He must be continually approaching instability, without ever actually getting there. It is likewise with an organism.
The organic chemistry of familiar life is stable, but approaching instability within the temperature range at which water is liquid, and with extremophile modifications, somewhat beyond that range. The reason silicon has difficulty forming large molecules is that within that same range, its bonds are weak. But at much colder temperatures, where carbon-based chemical activity all but ceases, silicon chemistry is still active, and silicon’s weaker bonds are strong enough to form long molecular chains. (Wells, by the way, had probably erred in placing his silicon-based organisms in warm environments.) Just as impressive, silicon can bond with those elements in ways that produce a variety of molecular structures, including rings and cages.13
Bains notes that because chemists have not made certain compounds with silicon doesn’t mean they can’t make them, and doesn’t mean that, somewhere, nature hasn’t made them already. “There is,” Bains says, “no reason why a high-molecular-weight silane or siloxane [a compound of silicon, oxygen, and hydrogen or hydrocarbons] should not have a highly diverse, highly structured set of side chains, analogous to proteins, nucleic acids or carbohydrates.”14
As to the charge that silicon cannot capture and distribute electricity, Bains summons a robust counterexample in certain polysilanes—that is, polymers with silicon backbones—that are semiconductors, and argues that other silicon-based compounds might channel energy as effectively. Bains contends that if biochemists and biophysicists suspect that silicon chemistry cannot channel energy, it is only because they assume it would use the same chemical pathways that organic chemistry uses. Silicon chemistry, he maintains, would have to use other pathways, and those pathways are yet to be discovered.
That no silicon life is evident on Earth is a point Bains does not dispute. In our oxygen-rich atmosphere, he notes, silicon biochemistry hasn’t a chance. The problem—if it is that—is that silicon bonds with oxygen easily and readily, and the result is silicon dioxide, or silica. The bond between silicon and oxygen is strong and difficult to break. It is the reason that most of the silicon on (and in) Earth is locked into rocks and, like some imprisoned fairy-tale princess, made unavailable for other chemical and biochemical activities. Bains’s point is that the apparent absence of silicon life on Earth is at best a reason to admit that it is unlikely here. It is not a reason to think it unlikely elsewhere.*
WHAT LIFE NEEDS (REDUX)
AND WHERE TO LOOK FOR IT
Let me summarize. Probably all biologists agree that life based in chemistry needs a source of energy. Most believe that it needs a semipermeable barrier that does the work a cell membrane does for familiar life. Most also believe it needs macromolecules with backbones of carbon. A few biologists—Bains among them—believe that those macromolecules may also have backbones of silicon.
If we use this near consensus to guide a wider search for life, we can begin to check a few places off—but, as it happens, only a few. Most places in the known universe have available to them a source of energy, and with a few exceptions—like the inert interiors of certain planets and moons—most places are in thermodynamic disequilibrium. Complex chemistry, the kind necessary to making semipermeable barriers and macromolecules, seems present at many places in the Solar System, and there is no obvious reason not to expect it elsewhere.
As discussed earlier, most biologists believe that chemically based life needs a liquid medium. It is here that we meet real limitations that allow us to narrow the search criteria. The only places we are likely to find liquids of any kind—water, ammonia, methane, or something else—are on the surfaces of planets and moons having atmospheres and orbiting in thin shell-shaped regions around stars, or in the interiors of planets, moons, and smaller bodies heated by some other means. While some liquids, like liquid hydrogen, are for various reasons unlikely, others, like ammonia, are known to exist in abundance.
Bains has graphed the distances from the Sun at which elements and compounds that could support complex chemistry are liquid. In so doing, he has shown us how to keep the search selective; at the same time, he has defined many habitable zones for weird life. It is easy enough for us to overlay these zones onto the orbits of the Solar System’s planets. Recall the placement of the habitable zone for familiar life, in the traditional version—from just inside Earth’s orbit to just beyond Mars’s. And recall the new version, which includes the traditional but also a
dds smaller zones beneath the surfaces of several moons of the outer planets. To astrobiologists grown used to these, the effect of the Bains overlay is agreeably unsettling. It is a bit like seeing a world map inverted or a chessboard turned 180 degrees halfway through a game.
Imagine the orbits of the planets of the inner Solar System—Mercury, Venus, Earth, and Mars—as four concentric circles on a dinner plate, with the Sun at the plate’s center and Mars’s orbit just inside the plate’s rim. A wide ring representing the traditional habitable zone for familiar life covers the orbits of Earth and Mars. It goes dark, and at the orbit of Mars appears another ring—this representing the range at which, on a planet or moon with a substantial atmosphere, the Sun’s warmth is sufficient to melt frozen hydrogen peroxide, but not so great as to boil it. A second ring, overlapping yet extending outside this ring, represents the range at which hydrogen peroxide could exist as a liquid beneath the surface of a planet or moon.
Zoom out for a field of view wide enough to accommodate the vastly larger scale of the outer Solar System, and the inner Solar System is shrunk to the size of a dime, now sitting in the center of the plate. Between the edge of the dime and the plate’s rim appear double rings representing analogous ranges for other liquids. From innermost to outermost are double rings for ammonia, methane, ethane, and, near the rim, liquid nitrogen.*
