Death by Black Hole: And Other Cosmic Quandaries - Neil deGrasse Tyson (2014)



For nearly all of the first 400 millennia after the birth of the universe, space was a hot stew of fast-moving, naked atomic nuclei with no electrons to call their own. The simplest chemical reactions were still just a distant dream, and the earliest stirrings of life on Earth lay 10 billion years in the future.

Ninety percent of the nuclei brewed by the big bang were hydrogen, most of the rest were helium, and a trifling fraction were lithium: the makings of the simplest elements. Not until the ambient temperature in the expanding universe had cooled from trillions down to about 3,000 degrees Kelvin did the nuclei capture electrons. In so doing, they turned themselves into legal atoms and introduced the possibility of chemistry. As the universe continued to grow bigger and cooler, the atoms gathered into ever larger structures—gas clouds in which the earliest molecules, hydrogen (H2) and lithium hydride (LiH), assembled themselves from the earliest ingredients available in the universe. Those gas clouds spawned the first stars, whose masses were each about a hundred times that of our Sun. And at the core of each star raged a thermonuclear furnace, hell-bent on making chemical elements far heavier than the first and simplest three.

When those titanic first stars exhausted their fuel supplies, they blew themselves to smithereens and scattered their elemental entrails across the cosmos. Powered by the energy of their own explosions, they made yet heavier elements. Atom-rich clouds of gas, capable of ambitious chemistry, now gathered in space.

Fast forward to galaxies, the principal organizers of visible matter in the universe—and within them, gas clouds pre-enriched by the flotsam of the earliest exploding stars. Soon those galaxies would host generation after generation of exploding stars, and generation after generation of chemical enrichment—the wellspring of those cryptic little boxes that make up the periodic table of elements.

Absent this epic drama, life on Earth—or anywhere else—would simply not exist. The chemistry of life, indeed the chemistry of anything at all, requires that elements make molecules. Problem is, molecules don’t get made, and can’t survive, in thermonuclear furnaces or stellar explosions. They need a cooler, calmer environment. So how in the world did the universe get to be the molecule-rich place we now inhabit?

RETURN, FOR A MOMENT, to the element factory deep within a first-generation high-mass star.

As we just saw, there in the core, at temperatures in excess of 10 million degrees, fast-moving hydrogen nuclei (single protons) randomly slam into one another. The event spawns a series of nuclear reactions that, at the end of the day, yield mostly helium and a lot of energy. So long as the star is “on,” the energy released by its nuclear reactions generates enough outward pressure to keep the star’s enormous mass from collapsing under its own weight. Eventually, though, the star simply runs out of hydrogen fuel. What remains is a ball of helium, which just sits there with nothing to do. Poor helium. It demands a tenfold increase in temperature before it will fuse into heavier elements.

Lacking an energy source, the core collapses and, in so doing, heats up. At about 100 million degrees, the particles speed up and the helium nuclei finally fuse, slamming together fast enough to combine into heavier elements. When they fuse, the reaction releases enough energy to halt further collapse—at least for a while. Fused helium nuclei spend a bit of time as intermediate products (beryllium, for instance), but eventually three helium nuclei end up becoming a single carbon nucleus. (Much later, when carbon becomes a complete atom with its complement of electrons in place, it reigns as the most chemically fruitful atom in the periodic table.)

Meanwhile, back inside the star, fusion proceeds apace. Eventually the hot zone runs out of helium, leaving behind a ball of carbon surrounded by a shell of helium that is itself surrounded by the rest of the star. Now the core collapses again. When its temperature rises to about 600 million degrees, the carbon, too, starts slamming into its neighbors—fusing into heavier elements via more and more complex nuclear pathways, all the while giving off enough energy to stave off further collapse. The factory is now in full swing, making nitrogen, oxygen, sodium, magnesium, silicon.

Down the periodic table we go, until iron. The buck stops at iron, the final element to be fused in the core of first-generation stars. If you fuse iron, or anything heavier, the reaction absorbs energy instead of emitting it. But stars are in the business of making energy, so it’s a bad day for a star when it finds itself staring at a ball of iron in its core. Without a source of energy to balance the inexorable force of its own gravity, the star’s core swiftly collapses. Within seconds, the collapse and the attendant rapid rise in temperature trigger a monstrous explosion: a supernova. Now there’s plenty of energy to make elements heavier than iron. In the explosion’s aftermath, a vast cloud of all the elements inherited and manufactured by the star scatters into the stellar neighborhood. And consider the cloud’s top ingredients: atoms of hydrogen, helium, oxygen, carbon, and nitrogen. Sound familiar? Except for helium, which is chemically inert, those elements are the main ingredients of life as we know it. Given the stunning variety of molecules those atoms can form, both with themselves and with others, they are also likely to be the ingredients of life as we don’t know it.

The universe is now ready, willing, and able to form the first molecules in space and construct the next generation of stars.

IF GAS CLOUDS are to make enduring molecules, they must hold more than the right ingredients. They must also be cool. In clouds hotter than a few thousand degrees, the particles move too quickly—and so the atomic collisions are too energetic—to stick together and sustain molecules. Even if a couple of atoms manage to come together and make a molecule, another atom will shortly slam into them with enough energy to break them apart. The high temperatures and high-speed impacts that worked so well for fusion now work against chemistry.

Gas clouds can live long, happy lives as long as the turbulent motions of their inner pockets of gas hold them up. Occasionally, though, regions of a cloud slow down enough—and cool down enough—for gravity to win, causing the cloud to collapse. Indeed, the very process that forms molecules also serves to cool the cloud: when two atoms collide and stick, some of the energy that drove them together is captured in their newly formed bonds or emitted as radiation.

Cooling has a remarkable effect on a cloud’s composition. Atoms now collide as if they were slow boats, sticking together and building molecules rather than destroying them. Because carbon readily binds with itself, carbon-based molecules can get large and complex. Some become physically entangled, like the dust that collects into dust bunnies under your bed. When the ingredients favor it, the same thing can happen with silicon-based molecules. In either case, each grain of dust becomes a happening place, studded with hospitable crevices and valleys where atoms can meet at their leisure and build even more molecules. The lower the temperature, the bigger and more complex the molecules can become.

AMONG THE EARLIEST and most common compounds to form—once the temperature drops below a few thousand degrees—are several familiar diatomic (two-atom) and triatomic (three-atom) molecules. Carbon monoxide (CO), for instance, stabilizes long before the carbon condenses into dust, and molecular hydrogen (H2) becomes the prime constituent of cooling gas clouds, now sensibly called molecular clouds. Among the triatomic molecules that form next are water (H2O), carbon dioxide (CO2), hydrogen cyanide (HCN), hydrogen sulfide (H2S), and sulfur dioxide (SO2). There’s also the highly reactive triatomic molecule H3+, which is eager to feed its third proton to hungry neighbors, instigating further chemical trysts.

As the cloud continues to cool, dropping below 100 degrees Kelvin or so, bigger molecules arise, some of which may be lying around in your garage or kitchen: acetylene (C2H2), ammonia (NH3), formaldehyde (H2CO), methane (CH4). In still cooler clouds you can find the chief ingredients of other important concoctions: antifreeze (made from ethylene glycol), liquor (ethyl alcohol), perfume (benzene), and sugar (glycoaldehyde), as well as formic acid, whose structure is similar to that of amino acids, the building blocks of proteins.

The current inventory of molecules drifting between the stars is heading toward 130. The largest and most structurally intricate of them are anthracene (C14H10) and pyrene (C16H10), discovered in 2003 in the Red Rectangle Nebula, about 2,300 light-years from Earth, by Adolf N. Witt of the University of Toledo in Ohio and his colleagues. Formed of interconnected, stable rings of carbon, anthracene and pyrene belong to a family of molecules that syllable-loving chemists call polycyclic aromatic hydrocarbons, or PAHs. And just as the most complex molecules in space are based on carbon, so, of course, are we.

THE EXISTENCE OF MOLECULES in free space, something now taken for granted, was largely unknown to astrophysicists before 1963—remarkably late, considering the state of other sciences. The DNA molecule had already been described. The atom bomb, the hydrogen bomb, and ballistic missiles had all been “perfected.” The Apollo program to land men on the Moon was in progress. Eleven elements heavier than uranium had been created in the laboratory.

This astrophysical shortfall came about because an entire window of the electromagnetic spectrum—microwaves—hadn’t yet been opened. Turns out, as we saw in Section 3, the light absorbed and emitted by molecules typically falls in the microwave part of the spectrum, and so not until microwave telescopes came online in the 1960s was the molecular complexity of the universe revealed in all its splendor. Soon the murky regions of the Milky Way were shown to be churning chemical factories. Hydroxyl (OH) was detected in 1963, ammonia in 1968, water in 1969, carbon monoxide in 1970, ethyl alcohol in 1975—all mixed together in a gaseous cocktail in interstellar space. By the mid-1970s, the microwave signatures of nearly forty molecules had been found.

Molecules have a definite structure, but the electron bonds that hold the atoms together are not rigid: they jiggle and wiggle and twist and stretch. As it happens, microwaves have just the right range of energies to stimulate this activity. (That’s why microwave ovens work: a bath of microwaves, at just the right energy, vibrates the water molecules in your food. Friction among those dancing particles generates heat, cooking the food rapidly from within.)

Just as with atoms, every species of molecule in space identifies itself by the unique pattern of features in its spectrum. That pattern can readily be compared with patterns catalogued in laboratories here on Earth; without the lab data, often supplemented by theoretical calculations, we wouldn’t know what we were looking at. The bigger the molecule, the more bonds have been deputized to keep it together, and the more ways its bonds can jiggle and wiggle. Each kind of jiggling and wiggling has a characteristic spectral wavelength, or “color”; some molecules usurp hundreds or even thousands of “colors” across the microwave spectrum, wavelengths at which they either absorb or emit light when their electrons take a stretch. And extracting one molecule’s signature from the rest of the signatures is hard work, sort of like picking out the sound of your toddler’s voice in a roomful of screaming children during playtime. It’s hard, but you can do it. All you need is an acute awareness of the kinds of sounds your kid makes. Therein is your laboratory template.

ONCE FORMED, a molecule does not necessarily lead a stable life. In regions where ferociously hot stars are born, the starlight includes copious amounts of UV, ultraviolet light. UV is bad for molecules because its high energy breaks the bonds between a molecule’s constituent atoms. That’s why UV is bad for you, too: it’s always best to avoid things that decompose the molecules of your flesh. So forget that a gigantic gas cloud may be cool enough for molecules to form within it; if the neighborhood is bathed in UV, the molecules in the cloud are toast. And the bigger the molecule, the less it can withstand such an assault.

Some interstellar clouds are so big and dense, though, that their outer layers can shield their inner layers. UV gets stopped at the edge of town by molecules that give their lives to protect their brethren deep within, thereby retaining the complex chemistry that cold clouds enjoy.

But eventually the molecular Mardi Gras comes to an end. As soon as the center of the gas cloud—or any other pocket of gas—gets dense enough and cool enough, the average energy of the moving gas particles gets too weak to keep the structure from collapsing under its own weight. That spontaneous gravitational shrinkage pumps the temperature back up, turning the erstwhile gas cloud into a locus of blazing heat as thermonuclear fusion gets underway.

Yet another star is born.

INEVITABLY, INESCAPABLY, one might even say tragically, the chemical bonds—including all the organic molecules the cloud so diligently made en route to stardom—now break apart in the searing heat. The more diffuse regions of the gas cloud, however, escape this fate. Then there’s the gas close enough to the star to be affected by its growing force of gravity, but not so close as to be pulled into the star itself. Within that cocoon of dusty gas, thick disks of condensing material enter a safe orbit around the star. And within those disks, old molecules can survive and new ones can form with abandon.

What we have now is a solar system in the making, soon to comprise molecule-rich planets and molecule-rich comets. Once there’s some solid material, the sky’s the limit. Molecules can get as fat as they like. Set carbon loose under those conditions, and you might even get the most complex chemistry we know. How complex? It goes by another name: biology.