Death by Black Hole: And Other Cosmic Quandaries (21 page)

While we know water to be essential for life on Earth, we can only presume it to be a prerequisite for life elsewhere in the galaxy. Among the chemically illiterate, however, water is a deadly substance to be avoided. A now-famous science fair experiment that tested antitechnology sentiments and associated chemical-phobia was conducted in 1997 by Nathan Zohner, a 14-year-old student at Eagle Rock Junior High School in Idaho. He invited people to sign a petition that demanded either strict control of, or a total ban on, dihydrogen monoxide. He listed some of the odious properties of this colorless and odorless substance:

• It is a major component in acid rain
  
• It eventually dissolves almost anything it comes in contact with
  
• It can kill if accidentally inhaled
  
• It can cause severe burns in its gaseous state
  
• It has been found in tumors of terminal cancer patients
  

Forty-three out of 50 people approached by Zohner signed the petition, six were undecided, and one was a great supporter of dihydrogen monoxide and refused to sign. Yes, 86 percent of the passersby voted to ban water (H
₂
O) from the environment.

Maybe that’s what really happened to all the water on Mars.

I
f you ask people where they’re from, they will typically say the name of the city where they were born, or perhaps the place on Earth’s surface where they spent their formative years. Nothing wrong with that. But an astrochemically richer answer might be, “I hail from the explosive jetsam of a multitude of high-mass stars that died more than 5 billion years ago.”

Outer space is the ultimate chemical factory. The big bang started it all, endowing the universe with hydrogen, helium, and a smattering of lithium: the three lightest elements. Stars forged all the rest of the ninety-two naturally occurring elements, including every bit of carbon, calcium, and phosphorus in every living thing on Earth, human or otherwise. How useless this rich assortment of raw materials would be had it stayed locked up in the stars. But when stars die, they return much of their mass to the cosmos, sprinkling nearby gas clouds with a portfolio of atoms that enrich the next generation of stars.

Under the right conditions of temperature and pressure, many of the atoms join to form simple molecules. Then, through routes both intricate and inventive, many molecules grow larger and more complex. Eventually, in what must surely be countless billions of places in the universe, complex molecules assemble themselves into some kind of life. In at least one cosmic corner, the molecules have become so complex that they have achieved consciousness and attained the ability to formulate and communicate the ideas conveyed by the marks on this page.

Yes, not only humans but also every other organism in the cosmos, as well as the planets or moons on which they thrive, would not exist but for the wreckage of spent stars. So you’re made of detritus. Get over it. Or better yet, celebrate it. After all, what nobler thought can one cherish than that the universe lives within us all?

TO COOK UP
some life, you don’t need rare ingredients. Consider the top five constituents of the cosmos, in order of their abundance: hydrogen, helium, oxygen, carbon, and nitrogen. Take away chemically inert helium—which is not fond of making molecules with anybody—and you’ve got the top four constituents of life on Earth. Awaiting their cue within the massive clouds that lurk among a galaxy’s stars, these elements begin making molecules as soon as the temperature drops below a couple thousand degrees Kelvin.

Molecules made of just two atoms form early: carbon monoxide and the hydrogen molecule (hydrogen atoms bound together in pairs). Drop the temperature some more, and you get stable three-or four-atom molecules such as water (H
₂
O), carbon dioxide (CO
₂
), and ammonia (NH
₃
)—simple but top-shelf ingredients in the kitchen of life. Drop the temperature even more, and hordes of five-and six-atom molecules form. And because carbon is both abundant and chemically enterprising, most of the molecules include it; indeed, three-quarters of all molecular “species” sighted in interstellar space have at least one carbon atom.

Sounds promising. But space can be a dangerous place for molecules. If the energy from stellar explosions doesn’t destroy them, ultraviolet light from nearby ultraluminous stars will. The bigger the molecule, the less stable it is against assault. Molecules lucky enough to inhabit uneventful or shielded neighborhoods may endure long enough to be incorporated into grains of cosmic dust, and ultimately into asteroids, comets, planets, and people. Yet even if none of the original molecules survives the stellar violence, plenty of atoms and time remain available to make complex molecules, not only during the formation of a particular planet but also on and within the planet’s nubile surface. Notables on the short list of complex molecules include adenine (one of the nucleotides, or “bases,” that make up DNA), glycine (a protein precursor), and glycoaldehyde (a carbohydrate). Such ingredients, and others of their caliber, are essential for life as we know it and are decidedly not unique to Earth.

BUT ORGIES OF
organic molecules are not life, just as flour, water, yeast, and salt are not bread. Although the leap from raw ingredients to living individual remains mysterious, several prerequisites are clear. The environment must encourage molecules to experiment with one another and must shelter them from excessive harm as they do so. Liquids offer a particularly attractive environment, because they enable both close contact and great mobility. The more chemical opportunities an environment affords, the more imaginative its resident experiments can be. Another essential factor, brought to you by the laws of physics, is a generous supply of energy to drive chemical reactions.

Given the wide range of temperatures, pressures, acidity, and radiation flux at which life thrives on Earth, and knowing that one microbe’s cozy nook can be another’s house of torture, scientists cannot at present stipulate additional requirements for life elsewhere. As a demonstration of the limits of this exercise, we find the charming little book
Cosmotheoros
, by the seventeenth-century Dutch astronomer Christiaan Huygens, wherein the author speculates that life-forms on other planets must grow hemp, for how else would they weave ropes to steer their ships and sail the open seas?

Three centuries later, we’re content with just a pile of molecules. Shake ’em and bake ’em, and within a few hundred million years you might have thriving colonies of organisms.

LIFE ON EARTH
is astonishingly fertile, that’s for sure. But what about the rest of the universe? If somewhere there’s another celestial body that bears any resemblance to our own planet, it may have run similar experiments with its similar chemical ingredients, and those experiments would have been choreographed by the physical laws that hold sway throughout the universe.

Consider carbon. Its capacity to bind in multiple ways, both to itself and to other elements, gives it a chemical exuberance unequalled in the periodic table. Carbon makes more kinds of molecules (how does 10 million grab you?) than all other elements combined. A common way for atoms to make molecules is to share one or more of their outermost electrons, creating a mutual grip analogous to the fist-shaped coupler between freight cars. Each carbon atom can bind with one, two, three, or four other atoms in this way, whereas a hydrogen atom binds with only one, oxygen with one or two, and nitrogen with three.

By binding to itself, carbon can generate myriad combinations of long-chain, highly branched, or closed-ring molecules. Such complex organic molecules are ripe for doing things that small molecules can only dream about. They can, for example, perform one kind of task at one end and another kind at the other; they can coil and curl and intertwine with other molecules, creating no end of features and properties. Perhaps the ultimate carbon-based molecule is DNA: a double-stranded chain that encodes the identity of all life as we know it.

What about water? When it comes to fostering life, water has the highly useful property of staying liquid across what most biologists regard as a fairly wide range of temperatures. Trouble is, most biologists look to Earth, where water stays liquid across 100 degrees of the Celsius scale. But on some parts of Mars, atmospheric pressure is so low that water is never liquid: a freshly poured cup of H
₂
O boils and freezes at the same time! Yet in spite of Mars’s current sorry state, its atmosphere once supported liquid water in abundance. If ever the Red Planet harbored life on its surface, it would have been then.

Earth, of course, happens to have a goodly—and occasionally deadly—amount of water on its surface. Where did it come from? As we saw earlier, comets are a logical source: they’re chock full of (frozen) water, the solar system holds countless billions of them, some are quite large, and they would regularly have been slamming into the early Earth back when the solar system was forming. Another source of water could have been volcanic outgassing, a frequent phenomenon on the young Earth. Volcanoes erupt not simply because magma is hot, but because hot, rising magma turns underground water to steam, which then expands explosively. The steam no longer fits in its subterranean chamber, and so the volcano blows its lid, bringing H
₂
O to Earth’s surface from below. All things considered, then, the presence of water on our planet’s surface is hardly surprising.

ALTHOUGH EARTH-LIFE
takes multifarious forms, all of it shares common stretches of DNA. The biologist who has Earth-on-the-brain may revel in life’s diversity, but the astrobiologist dreams of diversity on a grander scale: life based on alien DNA, or on something else entirely. Sadly, our planet is a singular biological sample. Nevertheless, the astrobiologist may glean insights about life-forms that dwell elsewhere in the cosmos by studying organisms that thrive in extreme environments here on Earth.

Once you look for them, you find these extremophiles practically everywhere: nuclear dump sites, acid-laden geysers, iron-saturated acidic rivers, chemical-belching vents on the ocean floor, submarine volcanoes, permafrost, slag heaps, commercial salt-evaporation ponds, and a host of other places you would not elect to spend your honeymoon but that may be more typical of the rest of the planets and moons out there. Biologists once presumed that life began in “some warm little pond,” to quote Darwin (1959, p. 202); in recent years, though, the weight of evidence has tilted in favor of the view that extremophiles were the earliest earthly life-forms.

As we will see in the next section, for its first half-billion years, the inner solar system resembled a shooting gallery. Earth’s surface was continually pulverized by crater-forming boulders large and small. Any attempt to jump-start life would have been swiftly aborted. By about 4 billion years ago, though, the impact rate slowed and Earth’s surface temperature began to drop, permitting experiments in complex chemistry to survive and thrive. Older textbooks start their clocks at the birth of the solar system and typically declare that life on Earth needed 700 million or 800 million years to form. But that’s not fair: the planet’s chem-lab experiments couldn’t even have begun until the aerial bombardment lightened up. Subtract 600 million years’ worth of impacts right off the top, and you’ve got single-celled organisms emerging from the primordial ooze within a mere 200 million years. Even though scientists continue to be stumped about how life began, nature clearly had no trouble creating the stuff.

IN JUST A FEW
dozen years, astrochemists have gone from knowing nothing of molecules in space to finding a plethora of them practically everywhere. Moreover, in the past decade astrophysicists have confirmed that planets orbit other stars and that every exosolar star system is laden with the same top four ingredients of life as our own cosmic home is. Although no one expects to find life on a star, even a thousand-degree “cool” one, Earth has plenty of life in places that register several hundred degrees. Taken together, these discoveries suggest it’s reasonable to think of the universe as fundamentally familiar rather than as utterly alien.

But how familiar? Are all life-forms likely to be like Earth’s—carbon-based and committed to water as their favorite fluid?

Take silicon, one of the top ten elements in the universe. In the periodic table, silicon sits directly below carbon, indicating that they have an identical configuration of electrons in their outer shells. Like carbon, silicon can bind with one, two, three, or four other atoms. Under the right conditions, it can also make long-chain molecules. Since silicon offers chemical opportunities similar to those of carbon, why couldn’t life be based on silicon?

One problem with silicon—apart from its being a tenth as abundant as carbon—is the strong bonds it creates. When you link silicon and oxygen, for instance, you don’t get the seeds of organic chemistry; you get rocks. On Earth, that’s chemistry with a long shelf life. For chemistry that’s friendly to organisms, you need bonds that are strong enough to survive mild assaults on the local environment but not so strong that they don’t allow further experiments to take place.

And how important is liquid water? Is it the only medium suitable for chemistry experiments—the only medium that can shuttle nutrients from one part of an organism to another? Maybe life just needs a liquid. Ammonia is common. So is ethanol. Both are drawn from the most abundant ingredients in the universe. Ammonia mixed with water has a vastly lower freezing point (around–100 degrees Fahrenheit) than does water by itself (32 degrees), broadening the conditions under which you might find liquid-loving life. Or here’s another possibility: on a world that lacks an internal heat source, orbits far from its host star, and is altogether bone-cold, normally gaseous methane might become the liquid of choice.

IN
2005, the European Space Agency’s
Huygens
probe (named after you-know-who) landed on Saturn’s largest moon, Titan, which hosts lots of organic chemistry and supports an atmosphere ten times thicker than Earth’s. Setting aside the planets Jupiter, Saturn, Uranus, and Neptune, each made entirely of gas and having no rigid surface, only four objects in our solar system have an atmosphere of any significance: Venus, Earth, Mars, and Titan.

Titan was not an accidental target of exploration. Its impressive résumé of molecules includes water, ammonia, methane, and ethane, as well as the multiringed compounds known as polycyclic aromatic hydrocarbons. The water ice is so cold it’s as hard as concrete. But the combination of temperature and air pressure has liquefied the methane, and the first images sent back from
Huygens
seem to show streams, rivers, and lakes of the stuff. In some ways Titan’s surface chemistry resembles that of the young Earth, which accounts for why so many astrobiologists view Titan as a “living” laboratory for studying Earth’s distant past. Indeed, experiments conducted two decades ago show that adding water and a bit of acid to the organic ooze produced by irradiating the gases that make up Titan’s hazy atmosphere yields sixteen amino acids.