By Natalie Angier, The New York
Times
Humans are really all wet. As
fetuses, we gestate in bags of water. As adults, we are bags of water: roughly
60 percent of our body weight comes from water, the fluidic equivalent of 45
quarts. Our cells need water to operate, and because we lose traces of our
internal stores with every sweat we break, every breath and excretion we
out-take, we must constantly consume more water, or we will die in three days.
Thirstiness is a universal
hallmark of life. Sure, camels can forgo drinking water for five or six months
and desert tortoises for that many years, and some bacterial and plant spores
seem able to survive for centuries in a state of dehydrated, suspended
animation. Yet sooner or later, if an organism plans to move, eat or multiply,
it must find a solution of the aqueous kind.
In the view of Geraldine
Richmond, a chemistry professor at the University of Oregon who often talks to
the public on the wonders of water, Mark Twain put it neatest: "Whiskey is
for drinking; water is for fighting over."
Behind water's peerless punch,
and the reason it rather than alcohol or any other lubricant serves as the
elixir of life, is the three-headed character whose chemical name we all know:
H2O. Scientists observe that when two atoms of hydrogen conjoin with one of
oxygen, the resulting molecule displays a spectacular range of powers, gaining
the mightiness of a molecular giant while retaining the speed and convenience
of a molecular mite.
"Water behaves very
differently from other small molecules," said Jill Granger, a professor of
chemistry at Sweet Briar College in Virginia. "If you want something else
with similar properties, you'd end up with something much bigger and more
complex, and then you'd lose the advantages that water has in being
small."
Because of water's atomic
architecture, the tendency of its comparatively forceful oxygen centerpiece to
cling greedily to electrons as it consorts with its two meeker hydrogen mates,
the entire molecule ends up polarized, with slight electromagnetic charges on
its foreside and aft. Those mild charges in turn allow water molecules to engage
in mild mass communion, linking up with one another and with other molecules,
too, through an essential connection called a hydrogen bond. From hydrogen
bonds many of water's major and minor properties flow.
With their hydrogen bonds, water
molecules become sticky, cohering as a liquid into droplets and rivulets and
following each other around like a jiggling conga line. Such stickiness means
that water is drawn to the inner plumbing of plants and will crawl up the
fibrous conduits to hydrate even the crowns of redwood trees towering hundreds
of feet above ground.
Pulled together by hydrogen
bonds, water molecules become mature and stable, able to absorb huge amounts of
energy before pulling a radical phase shift and changing from ice to liquid or
liquid to gas. As a result, water has surprisingly high boiling and freezing
points, and a strikingly generous gap between the two. For a substance with
only three atoms, and two of them tiny little hydrogens, Dr. Richmond said,
you'd expect water to vaporize into a gas at something like minus 90 degrees
Fahrenheit, to freeze a mere 40 degrees below its boiling point, and to show
scant inclination to linger in a liquid phase.
That's what happens to hydrogen
sulfide, a similarly sized molecule but with its two hydrogen atoms fastened to
sulfur rather than to oxygen; on our temperate world, hydrogen sulfide has long
since reached its boiling point and exists as a foul-smelling gas. Same for the
tidy troika of carbon dioxide: low freezing point, low boiling point, and, poof,
it's up in the air. But given its vivid power of hydrogen bonding, water proves
less flighty and fickle, with a boiling point at sea level of 212 degrees
Fahrenheit, and a full 180 degrees lying between the tempest of a teapot and
the tinkling of an ice cube at 32 degrees. A vast temperature span over which
water molecules can pool and cling as the liquid assets we love best.
We rely in myriad ways on
water's fluid forbearance, its willingness to take the heat without blinking.
Earth's oceans and lakes soak up huge quantities of solar radiation and help
moderate the climate. As sweat evaporates from our skin, it wicks away large
amounts of excess heat.
Water also serves as a nearly
universal solvent, able to dissolve more substances than any other liquid. It can act as an
acid, it can act as a base, with a pinch of salt it is the solution in which
the cell's thousands of chemical reactions take place.
At
the same time, water's gregariousness, its hydrogen-bonded viscosity, helps
lend the cell a sense of community.
"Water
acts as the contact between biological molecules, not just separating them, but
imparting information among them," said Martin Chaplin, a professor of
applied science who studies the structure of water at London South Bank
University. "In an aqueous environment, all the molecules are able to feel
the structure of all the other molecules that are present, so they can work as
whole rather than as individuals."
There's
no end to water's chemical kinkiness, including the way it freezes from the top
down and becomes buoyant as it chills. Most substances shrink and get denser
and heavier as they cool, and expand and lighten as they melt. Water bucks the
norm, and is lighter and airier as ice than when liquid, and so in winter
marine life can find liquid haven beneath the floating blanket of ice, and so
in summer ice cubes bob and clink in your glass of lemonade. Bottoms up.
