The Canon (25 page)

The supple power of the molecular bond that gives us edible carbon fare and breathable oxygen pairs is crucial to life, but a covalent commitment can still be too ham-fisted when life demands Nijinsky. Here the secondary bonds come into play, and weakness becomes a source of strength. The backbone of DNA may be held together by carbon bonds, but the double helix is like a zipper, the "teeth" designed to fasten together or come apart as needed. If one of your body cells is about to replicate, for example, the two halves of the DNA molecule must separate to allow a copy, a carbon copy, to be made. If the cell isn't dividing, but simply needs to generate a fresh supply of an essential protein like hemoglobin or insulin, for example, the DNA must still zip apart, just a bit, to expose the spot where the recipe for hemoglobin is written in nucleic script. Open, shut, unwind, tighten up. The molecule of life jiggles viscously. Life was, after all, born in water, and DNA, the sentinel of heredity, doesn't forget its roots. The bond that holds the two halves of the helix together, that joins each tooth, or chemical letter, of one strand to the complementary tooth of the other strand, is the hydrogen bond, the same bond that keeps water clinging in the lugubrious bulge of a teardrop. The attractive tippiness that a molecule gains when a hydrogen atom shares electrons with a bigger, bolder, and more possessive element like oxygen, nitrogen, or carbon is just right for the needs of our genetic code. A hydrogen bond is strong enough to maintain the serpentine shapeliness of DNA during its tranquil hours within the nucleus, but it's an easy bond to interrupt for the sake of making new proteins or a whole new set of chromosomes.

The same goes for the protein molecules themselves. Proteins must have specific shapes to perform their tasks in the cell, but they must also be nimble, sinuous, squashable. Hydrogen bonds help define the most Gumbyish contours of a protein, allowing it to pleat out a bit on one side or buckle inward on another. Through hydrogen bonding, a hemoglobin protein can tangle itself until it looks like a plate of spaghetti and meatballs, each ball a lump of iron for clasping the oxygen we crave. Or an antibody protein of the immune system can array its four floppy chains into a straitjacket to form-fit any microbe encountered.

Sometimes, when you irreparably rupture a hydrogen bond, you can see life harden before your eyes. The clear liquid inside a freshly laid egg is a delicate matrix of some forty proteins designed to cosset a fetal chick through development, and those proteins owe their three-dimensional contours to hydrogen bonds. Fry the egg, and you destroy the bonds, freeing the protein components to rearrange themselves as haphazardly as they please. The buoyant, forward-looking, transparent syrup congeals into an opaque sedentary solid that now deserves the designation of egg white.

And while the hydrogen bond is first among secondaries when talking about the microscopic side of life, we organics mustn't neglect our van der Waals, the lumper and clumper and soft tissue wrangler. The layers of our visceral organs, and the figgy pudding furrows of our brain, are largely held together by van der Waals bonds. Our adipose depots in particular owe their cohesion to this loosest of glues, which is why it is easy to slice through fat with a steak knife or surgeon's scalpel—easier than breaking down the constituent fat molecules with exercise, for they are energy-rich stores of carbon bonds. Plants, too, rely on the van der Waals magnetism of their cellulose walls to survive. The internal surfaces of a plant's roots and stem are slightly charged, attracting water molecules from the soil and prompting them to begin crawling upward, as water will crawl up a paper towel if the corner of the paper is dipped in a puddle. Hydrogen bonds then ensure that more water molecules will follow their leaders along the cellulose road. The name of life is bonds, all bonds, an ecumenical band of bonds, each contributing what talent it can to uphold order, bolster morale, and resist the cosmic drift toward rot, at least for one more day.

Chemistry is about molecules and bonds, and it is also about a match-stick found, brandished, and at last ignited. "Chemistry is the science of change, the study of transformation," said Rick Danheiser. Its roots lie in alchemy, the ancient effort to turn lead into gold, the mundane into the glamorous, the dead into the reborn; the word "alchemy," from the Greek
khemia,
means "black sand," which the Greeks associated with ancient Egypt and its elaborate devotion to ensuring the pharaonic class a good life in the afterlife. The Chinese words for "chemistry" and "change" share a common ideogram, which shows a simple but unmistakable postural transformation, from a person standing to a person sitting.

The most unmistakable chemical transformation is that of a matter's state—a solid liquefies, a liquid evaporates, a vapor condenses into rain. For most of the furnishings of our everyday life, we associate a particular substance with only one of those three states. Wood, steel, and stone—solid. Oxygen and helium—gas. Alcoholic beverages—liquid (you can keep a bottle of Bombay Sapphire in the freezer, and somehow it remains an ever pourable starter to a gin and tonic). Water again bucks convention and seems almost equally at home in all three forms, as ice, steam, and liquid. In fact, Earth is exceptional in its possession of tristate water. Mars has a lot of water, but it's frozen away underground. Jupiter and Saturn have traces of water, too, but as orbiting ice crystals or as a gas among miasmic gases. Only on Earth are there ocean flows and Arctic floes and sputtering Yellowstone fumaroles; only the Goldilocks planet has water to suit every bear.

What accounts for the differences among a solid, a liquid, and a gas? And why are certain solids so reluctant to melt, while others begin to ooze out of the bag if you even think about taking them on a picnic? One obvious parameter you can fiddle with to induce a phase change in your sample is heat. Fry an ice cube, and the ice cube melts. By adding heat, you are amplifying molecular anxiety. Granted, molecules are fidgety from the start. Every bit of matter, no matter how sober its appearance, is constantly quivering at the base of itself; protons must spin, electrons gotta fly. In a solid, however—that is, a material with a fairly fixed shape and volume—the constituent molecules can only move so much, the strength of their individual motions counterbalanced by the firmness of their ties. As long as temperatures (and pressures) remain reasonably constant, the particles content themselves with Jack LaLanne isometrics and jogging in place.

Add heat to the solid, though, and the molecular oscillations quicken. The excited particles tug and strain against their bonds, and snap at their personal trainer, until a scattering of tiny rips appears in the three-dimensional array. Now the particles have some room to start sliding over one another. More sliding here means more gapping there, and more opportunity for the oscillating participants to shift out of their frame. When the last of the hindrances to intermolecular glide has disassociated, you're left with a liquid, a flowing substance that has a measurable volume but no fixed shape. If the liquid is heated more, the particles may gain enough kinetic dash to overcome any of the attractive forces that keep molecules clinging together and begin springing free of the surface, as a gas. The components of a gas still have their molecular integrity; the individual water molecules streaming out of your screaming teapot maintain their covalently bonded H
₂
O formulation; but they've shed all volume control and will diffuse out through whatever space you give them.

As a rule, ionic solids of rocks and bones are extremely resistant to melting and boiling. The rigid bonds holding ion to ion defy being loosened and pushed aside, the first step of liquefaction. Many a detective story converges on the telltale hearth, where the victim's skeletal remains refused to be stewed into silence. An ordinary wood fire burns at about 650° Fahrenheit and will barely make a dimple in teeth or bones; even the infernal 1800° of a professional crematorium needs two or three hours to dissolve the bulk of a decedent's skeleton, and still there can be bone fragments lingering in the ash. Metals, too, often have the mettle of Mephistopheles, and melt only at very high temperatures. Not only is the metallic bond that results from a sharing of electrons among multiple atoms quite strong, but the bartering system encourages metal atoms to pack themselves as densely as possible in three dimensions. The degree of solidity and resistance to melting varies considerably from metal to metal, though. Iron atoms have up to three electrons to share with their peers, and they stack together so closely that each atom touches twelve of its neighbors; iron won't melt until 2800° Fahrenheit, or 1538° Celsius. Herring-soft sodium, on the other hand, can share only
one electron with its mates, and so a sodium-sodium federation is comparatively lax and will melt at 208°F. Silver, copper, and gold possess similar orbital architecture and all thaw down at somewhat less than 2000°F.

And then there is mercury, arguably the barmiest of all the elements. Mercury is liquid at room temperature, and it conducts heat and electricity so poorly that it barely merits inclusion in metaldom. Behind mercury's unusual behavior is its massive nucleus and the strong pull of its eighty protons. The positive packet at mercury's core keeps such a powerful lock on all the surrounding electrons that, even though the element theoretically has two negative particles to share in an electron sea, those electrons prefer staying close to their nuclear family, leaving the metallic bonds linking one mercury atom to another weak and easily disrupted.

Yet even as mercury's natal spirit is feckless and mercurial, the element readily forms soft amalgams with other metals, including silver and gold. The miners of ancient Egypt and Greece used mercury to extract gold from ore, and alchemists were convinced that if anything could transform lead into gold, it was the bobbling, quasi-animated metal they called chaotic water, or quicksilver. The magnificent Sir Isaac Newton, a passionate if episodic alchemist, considered mercury less a distinct element than a fundamental principle, the essence of all metals, and he sought it in its noblest and most "philosophical" form. Working in his Cambridge laboratory, Newton handled and sampled mercury droplets and inhaled their volatilized fumes, until he became as mad as a hatter or as flaky as a furrier—tradesmen that famously cured their fabrics in mercury and infamously suffered from the metal's neurotoxic effects. Preserved locks of Newton's hair reveal high concentrations of mercury, and, according to contemporary accounts, he grew increasingly hostile and choleric over time. Toward the end of his long life, the man who earlier had discovered the universal laws of gravity, motion, and optics and invented calculus, and whom James Gleick called the "chief architect of the modern world," expressed little interest in anything but that most fantastical of Gospels, the book of Revelation.

In contrast to ionic solids and the less mercurial metals, molecular solids are often disturbingly easy to melt and boil. This is especially true of solids that contain a mixture of different but closely related molecules, as do the soft organs of the body. Such solids are likely to lean heavily for their gross morphology on van der Waals, the promise most easily broken. A stick of butter, for example, which is about 80 percent fat and 20 percent protein, milk sugars, and other dairy components,
melts at just about the temperature of the mouth—a concordance that in no small way explains butter's rich "mouth feel" and its inclusion in so many dishes we judge delicious.

Not every heated substance passes in orderly goosestep from solid to liquid to gas. Take frozen carbon dioxide, or dry ice, the basis of so many memorable children's birthday parties and forgettable stagings of
Macbeth.
On being exposed to room temperature, a block of dry ice bypasses the liquid stage altogether and evaporates directly into billowing white boas of smoke, an act of phase-change denial called sublimation. Dry ice owes its plumosity to both the relative frailty of the bonds binding carbon dioxide molecules together and the paucity of carbon dioxide in the lower atmosphere. At shirtsleeve conditions, the intermolecular links in dry ice quickly begin to dissolve, and the surrounding air essentially sucks the loosened rarities up wholesale and begs for more. Regular H
₂
O ice can also sublimate directly into vapor without pausing at the aqueous phase, though it does so much less dramatically. This is why ice cubes in a freezer tray gradually shrink despite the persistence of ambient frigidity. The circulating air skims off occasional water molecules from the top of the ice and eventually redeposits them as a scrim of frost on the sides of the freezer—or, if the cubes are loose in a container, as a kind of glue welding everything together into a Gaudíesque hoodoo of ice.

Melting, freezing, boiling, condensing, all represent physical changes in matter's state, but not in its composition. The molecular modules may get anarchic or they may get military, but they maintain their molecular identity. A rose petal is a rose petal, whether velvety-limp on the wedding room floor, or Popsicle-stiff in a liquid nitrogen bath. If you want something truly novel, you must change the substance chemically. You must break the extant molecules apart and reshuffle the subunits into new molecular configurations. If you want your loaves to leaven or your juice to ferment, neither boiling nor freezing nor squeezing will do. You need the pith of that allegoric black
khemia
on which the science of change is built. You need a chemical reaction. And what better way to summon the spirit of change than by raising a toast to the toadstool?

Fermentation may well be the oldest chemistry experiment in human history. Nobody knows how or when the first alcoholic beverage was made and sampled and declared "Satiny, supple, and exuberant, with notes of black fruit compote, sassafras, cocoa, cinnamon, meat, mineral, forest floor, Tigris, Euphrates, T'ang, and Tang.® Best if drunk before construction of the first ziggurats." Very likely the event was a total accident, the result of a few yeast spores blowing into a pot of mash that a careless child or a clinically depressed slave had forgotten to clear from the table. Whatever its origins, the vintner's art was domesticated soon after the advent of the agricultural revolution. Chemical traces on pottery shards from nine thousand years ago suggest that the citizens of Jiahu, a village in the Henan province of northern China, brewed a wine made from rice, grapes, and honey, a varietal that may explain why the best thing to drink with Chinese food is beer. And while alcohol has its desolate side, and has killed or made killers of millions, it also has kept millions alive. Through the many millennia before the advent of public sanitation measures, when water was notoriously nonpotable, people of all ages, at least in the West, often quenched their thirst with alcohol instead; given its mild antiseptic properties and its acidity, liquor was far less likely than water to carry parasites. The populace may have been slightly intoxicated much of the time, but better tipsy than typhoid.