The Canon (24 page)

Yet water's hydrogen bonds are slippery and will slide aside to stir things up when something thicker this way comes. Water has been called the universal solvent, for there are precious few substances that will not dissolve in its embrace. Stir a spoonful of salt into water, and water's mighty mice swiftly interpose themselves between the individual crystals, the positive ears tempting the negative chlorines, the negative jaws jockeying for sodium, until the salt grains have disintegrated into a fine mist. Give polarized water molecules about 6 million years, and they'll squeeze blood red beauty from stone, chipping 6,000 feet deep and 277 miles wide into Arizona's northern plateau, through limestone and sandstone and iron-rich shale, to scoop out a canyon the whole world can call Grand.

Hydrogen bonding, though, is not unique to water. It arises in other cases where hydrogen, the lightest of the elements, enters into a covalent compact with a bulkier element, such as nitrogen, and the shared electrons tip their allegiance toward the nucleus of hydrogen's partner. From that asymmetry, a molecule that is electrically neutral in totality assumes a Mickey-like fuzz of charge around the ears and chin.

Another intermolecular melder is the van der Waals force, christened after the Dutch physicist who discovered and mathematically characterized it in the late nineteenth century. Despite the intimidating length of its name, van der Waals is the weakest of the links, as anybody who has ever left a pottery class streaked in enough clay for a solstice fete can attest; its strength is less than a quarter that of a hydrogen bond. Still, a mild manner has its advantages, and van der Waals is essential to the integrity of many solids and liquids, and to the properties of a wide assortment of substances on which we depend. Whereas in other bonds, including the hydrogen bond, electrons tend to know their place and to accord the resulting molecule or compound a fairly fixed arrangement
of negative and positive charges, the van der Waals force showcases the electron's improvisational skills.

Electrons, of course, don't like the feel of other electrons, and that antipathy explains why we can touch objects constructed of the near nothingness of atoms and not go right through them. At the same time, electrons are drawn toward protons—the positive particles of their own nucleus or of any nucleus in the neighborhood. The same predispositions apply to electrons in the corporate environment of a liquid or solid, when they are part of molecular or ionic teams. Nuclear protons, good; other electrons, bad. The consequence of this fundamental preference is that, when atoms and molecules come into close proximity with one another, their electrons tend to shift themselves to one side of the home cloud or another, avoiding regions of ambient electron glut, seeking out patches of heightened proton pull. The molecules hence become ever so mildly polarized, or asymmetrically charged, and this gentle layering of negative and positive charges helps bind many substances together. It's a frail fraternity. Electrons are not formally shared between atoms, as they are in molecules and ions, nor are they committed to their unbalanced orbits, as they are in the Disney design of a water molecule.

Still, van der Waals is sometimes the only force holding big chunks of matter together. Pottery clay, for example, consists of sheets of diverse atoms—silicon, aluminum, oxygen, hydrogen, calcium, nitrogen, iron, maybe a sprinkling of cobalt, copper, manganese, and zinc. Within each sheet, the atoms are lashed together by persuasive covalent and ionic bonds. But between the sheets, only the van der Waals force can be found. That's why it's so easy to smear a bit of the putty onto your fingertip; all you're doing with the pickup is interrupting the circumstantial attraction between sheets of mildly polarized clay particles. The integrity of the molecular bonds themselves becomes obvious, though, as you struggle to remove the fine clay slick from your fingertip and discover that it is very difficult to disengage or smoosh apart. Hours later, you may still feel a greasy residue clinging to every whorl—the lingering clay molecules that can be removed definitively only by fracturing their covalent bonds with a serious detergent or chemical solvent.

Your ordinary, standardized test–taking pencil offers another instance of the sound of a van der Waals snapping. The "lead" of the pencil (graphite was long thought to be a soft form of lead, and by the time chemists realized otherwise, the term "lead pencil" was already lodged deep in the schoolhouse lexicon) is not lead at all, but graphite—countless sheets of carbon atoms stacked one on top of another, rather
like the tissue-fine layers of caramelized toffee inside a Butterfinger candy bar. Within each graphite sheet, the carbon atoms lock elbows covalently in repetitive, crystalline patterns, but van der Waals alone joins one leaf to those above and below. Press the point of the pencil on your paper to fill in your oval, and you shave a stratum or two of carbon crystals away from the larger deck.

From this versatile cast of bonds, all the stuff of life and site can be staffed. Ionic bonds characterize much of the glittering landscape on which we and our steel-belted radials tread—the mountains, hills, rocks, sand, the shattered seashells by the seashore, the bleaching coral reefs beneath. Ionic solids tend to be rigid, their ions bound together so stiffly they cannot easily be pushed aside. That rigidity makes ionic solids ideal for load-bearing tasks: what better way to begin a bridge than with a few ionically bound concrete pylons, and what better batter for a sidewalk than an ionic compound like cement? Our skeletons, too, are made in part of ionic solids, tightly interdigitated concatenations of calcium, phosphorus, and other atoms. Through bone, we clawed free of the bog; we carry our steppingstones inside us.

Yet ionic solids will get you only so far, and their strength tends to be brittle. You can push down on them and they'll hold up stoutly, but give them a few good rotations, or maybe one swift blow with a hammer, and their ionic bonds will rupture and the crystal palace crumble. For that reason pylons are buried underground, to stabilize them; an uppity tree root can buckle a sidewalk so badly the panel will crack in half; and the twisting of an ankle may end up fracturing the bone. Luckily for us, our bones are marbled through with a soft mortar of proteins, which gives them far more flexibility and torque resistance than they'd have if they were nothing but ionic columns. And lucky, too, that beneath the most brittle outer sheath of our bones lies a network of regenerative tissue that can give birth to new bone cells, sealing cracks and healing breaks and doing for our vertebrate frames what could not be done for the ionic solid of an eggshell: put us together again.

Most of our body tissue consists of covalently rather than ionically bonded compounds—of molecules, not salts. We're abundantly hydrated, of course, and can blame at least 60 percent of our body weight on water molecules, more if we're a pedestrian in Manhattan too embarrassed to try sneaking into a restaurant's
FOR CUSTOMERS ONLY
bathroom when the maître d' isn't looking. Wring us out—finally!—and the bulk of what's left explains the sci-fi canard about humans as self-replicating carbon units: about two-thirds of our dry weight consists of carbon. Water may be the solvent of the universe, but carbon is
the duct tape of life. Every cell, every component of the cell, is based on carbon. If you're seated somewhere on the tree of life, or are the tree of life,
ipso facto
you contain carbon, and that goes for bacteria, amoebas, lichen, dust mites, pinworms, creationists. Even viruses, considered by many to be less than certifiably alive, nonetheless contain carbon, as part of the genetic backpack they tote around from host to host. Small wonder that half of all chemists work in the field of organic chemistry, which has nothing to do with the pure 'n' natural foods industry but instead is the study of compounds that contain carbon.

We are "just" carbon-based units because carbon makes for a just-right class of molecules. Carbon is strong, resourceful, flexible, sociable. With its outer shell of four electrons and four electron slots for rent, carbon is supremely suited to molecular bondage. It happily collaborates with nearly every actor on the periodic table, save helium, neon, and the four other noble elements,
^*
so-called for their snobby refusal to connect chemically to anything. Moreover, carbon is unparalleled among elements in its ability to join with itself almost indefinitely, forming carbon chains and carbon loops and branching carbon prongs and broad carbon planes and bouncing carbon buckyballs. Whatever shape you need to suit whatever cell part or enzyme you desire, chances are it is best draped on a carbon frame. Moreover again, the bond between two carbon atoms is one of the strongest bonds known, far stronger than that between two atoms of silicon, an element that otherwise has much in common with carbon. The strength of the carbon bond helps explain why it is the basis of life: we need molecular stability now, and we really needed it when life was new and the world was a considerably harsher place than it is today. At the same time, the carbon bond under ordinary conditions can bend, spring, and curl, hence the capacity of carbon molecules to array themselves as rings, cages, and coils. Carbon is as good as Goldilocks for building the spiraling, switchbacking molecule called DNA, and so the sugary spine of the double helix, and the individual chemical letters of which its code is composed, are carbonated through and through.

And while it may be mere coincidence, there's something gratifying about which gem we carbon vessels seize on prior to making a few carbon copies of our own: the diamond. Perhaps nothing underscores carbon's chemical genius better than the breadth of its packaging options, from the dark, slippery, shavable format of graphite on one extreme, to fossilized starlight on the other—translucent, mesmeric, intransigent
diamond, the hardest substance known, save for a human heart grown cold.

What spells the difference here, between carbon as ductile lubricant, a material you can spritz into balky locks, and carbon by De Beers? In graphite, each carbon atom is covalently bonded to three other carbon atoms, all of them lying in the same two-dimensional plane; there is no upstairs-downstairs blending of electrons but only the wan charms of van der Waals holding one floor to the next, so they slip-slide away.

In a diamond, by contrast, the bonds are fully fleshed out in every direction. Now, each carbon is strapped covalently to four of its kind, the maximum possible, and across three-dimensional space. To the left, right, crownward, groundward; wherever a carbon looks, there's a carbon bound to it. They're packed together so tightly and with such crystalline homogeneity that light finds very little impediment to its passage, very few imperfections to bounce off of and muddy the view, and the diamond gleams translucently. And because anywhere one might want to slice, one encounters thickets of jealous carbon-carbon bonds, a diamond feels like forever; to cut a diamond, a professional diamond cutter uses another diamond.

This painstaking compaction and positioning of carbon atoms is extremely difficult to accomplish. Getting every atom just where it needs to be to bond in a sororal three-dimensional mosaic, millions upon millions of flawlessly arrayed rings of four-faceted tetrahedrons, takes time and tremendous force. Until recently, the only place diamond factories could be found was hundreds of kilometers underground, in the Earth's mantle, where carbon stores subjected to great heat and pressure over millions or billions of years finally locked together in fixed constructs. Every so often, a volcanic eruption would spew a geyser of these diamonds to the surface, and another monarch might have his diadem, or Marilyn a pear-shaped friend. Industry also came to rely on diamonds for their unequaled ability to abrade metal machine parts into shape, and semiconductor manufacturers sought diamond bits to install in their microchips, to help prevent the embedded circuits from overheating. Diamonds happen to be excellent heat sinks, which is why even a room-temperature gem will feel cool to the touch. Put your fingertips or puckered lips against a diamond, and the jewel drains warmth from you to it, a heat transfer that your brain interprets as a brush with something cold; in fact, their high thermal conductivity, rather than their crystal clarity, earned diamonds the alias "ice."

Whatever the argot, diamonds clearly were too useful to leave to chance delivery through a magma pipeline. In the mid-twentieth century, scientists figured out how to mimic conditions in the bowels of the earth and began fabricating industrial-grade synthetic diamonds. More recently, researchers have managed to gin up gem-quality diamonds as well, although the process is so expensive that the resulting stone may cost a Tiffany customer more money than would a natural diamond from a Namibian mine.

Carbon bonds less zealous than those in diamonds help tape us together, and carbon bonds keep us alive. Most of the food that we eat, the carbohydrates, fats, proteins, and fiber, are carbon-centered compounds, which works out to an average daily intake of three hundred grams of pure carbon—about the weight of a pair of kidneys—per belly per day. Some of that ingested carbon is put to use directly, to repair damaged cells or assemble hormones, but more often the body simply cracks open the carbon bonds to extract the energy stored therein, and then tosses out the carbon atoms in the form of exhaled carbon dioxide. One species' waste is another's preferred taste, however. Plants blend the carbon dioxide together with water to synthesize sugars—and generate as a lucky byproduct the oxygen we need. The carbon cycle is just one of the many mandalas of life on which we rely, often mindlessly, and which we monkey with slapdashedly. Carbon bonds are dense packets of energy, and we can't get enough of them. The bulk of the energy driving the engines of our economy and our vehicles comes from breaking and remaking carbon bonds in coal, natural gas, and oil. Our cars, like our bodies, take the bond energy and dump the carbon, as carbon dioxide. Humans burn about 7 billion tons of fossil fuels a year, and so carbon deposits that might otherwise have hibernated underground for millennia instead combust into the atmosphere, harrying a carbon cycle that is already spinning as fast as it can.