You never hear someone say, Bartender, another round of C2H5OH for my friends here, or, Do you want coffee with C8H10N4O2 or decaf? The only substance commonly known by its chemical formula is good old H2O. Water is not only ubiquitous, it’s simple: two hydrogens perch on a lone oxygen like ears on an atomic mouse. When schoolteachers search for the perfect molecule to introduce elementary concepts of chemistry, they naturally turn to water.
But first impressions are deceiving. Underneath that surface simplicity, water’s mysteries run deep. The Physics of Simple Liquids, a standard, 500-plus-page guide that sits on many researchers’ bookshelves, mentions water only once--in the preface, explaining that water is too complicated to be discussed in the book. It’s the hardest liquid, says physicist Eugene Stanley of Boston University.
Water’s subtle complexity manifests itself in many ways. Unlike other liquids, it flows faster under high pressure than under low pressure. Its solid form is lighter than its liquid form. As a solvent, it’s gentle enough to serve as the cradle of life yet brutal enough to carve the Grand Canyon. This contradictory character results from an atomic arrangement that makes water almost promiscuously friendly. While other liquids are generally aloof, water’s oxygen and hydrogen atoms like to form bonds with practically anything that crosses their path, even other water molecules.
During the past decade computers have been helping researchers visualize the invisible choreography of H2O molecules that produces water in all its familiar and exotic forms. As in all chemical reactions, the macroscopic effects we see are determined by the arrangements of molecules and the changing configurations of the electric fields they create. Water’s simple structure makes it possible to simulate these features with exacting detail. Researchers put their imitation water through its paces, creating analogues of complex situations, observing how the molecules behave; then they compare the computer results with lab experiments using the real thing. These computer-simulated close-ups of water as ice, liquid, and vapor are drawn from some of the most recent research. Perhaps they’ll inspire you to gaze with somewhat more respect at the mystery dripping from your kitchen faucet.
Judging by its chemical structure, water ought to be stiff and syrupy, more a gel than a liquid. The chemical bonds that hold water molecules together are extremely powerful; and a molecule in liquid water can make four bonds with other molecules--its two hydrogen atoms can latch onto other oxygen atoms, while its own oxygen atom can link with two hydrogen atoms. Just about every molecule in a glass of water is bonded at a given instant to four other molecules.
So why is water so, well, watery? Eugene Stanley and his colleagues have found an explanation in their computer simulations. They filled an imaginary cube with 216 water molecules and had the computer track the changes in energy that each one experienced as it moved around. They then made a movie showing how the molecules rearranged themselves over a millionth of a millionth of a second. It turns out that into each five- molecule cluster a sixth molecule insinuates itself. The interloper does so by forcing one of the molecules in the cluster to share a hydrogen bond with it. A shared bond becomes weakened and unstable, so that the intruding molecule can eventually push out another. Clusters are therefore continuously rearranging themselves in a microscopic game of musical chairs that keeps liquid water flowing.
The image on the preceding pages is a snapshot from the movie created by Stanley’s team. The hydrogen bond between two molecules is represented by a stick connecting the centers of each. Together the sticks create a sort of three-dimensional net. Two clusters of actual molecules are superimposed on the mesh. The one in the lower left-hand corner is made up of five molecules. The one in the upper right-hand corner has been joined by a sixth.
Chemists have discovered that in vapor form, water molecules can arrange themselves into elegant cages. These images, generated by A. Welford Castleman and his colleagues at Penn State, depict clusters of molecules that the chemists created in their lab.
Researchers had long known that water vapor often contains clusters of molecules, as many as 21 in a single clump. But they didn’t know what shape those clumps took. To determine that, they had to know how many unbonded hydrogens there were in such a grouping. To find out, the Penn State team spiked a beam of vapor with a chemical compound that glued itself to any available hydrogen atom. The more unbonded hydrogen atoms on a cluster’s surface, the chemists reasoned, the heavier it would become as more molecules hung on. In fact, the mass of a tagged cluster turned out to be an exact measure of the number of hydrogen atoms facing outward. Using a computer, they discovered that the only possible arrangement ten unbonded hydrogens could take was the soccer-ball shape shown in these pictures (red spheres are unbonded hydrogen atoms).
Such clusters seem to carry an extra proton that snags passing molecules to build the cage. (This stray hydrogen nucleus is probably a remnant of an H3O molecule that broke up as two of its hydrogens and an oxygen--H2O--hooked onto the cage, leaving the lone H behind.) On a 20- molecule cluster (opposite page, left), the molecules pass this proton to one another like a hot potato, so that it races around the surface, binding briefly with the oxygen atoms before moving on. When a twenty-first water molecule joins the cage, however, the structure becomes unstable. Then both the free proton and the new molecule get sucked into the center, where they remain imprisoned (right, page 104). Using beams of other compounds, Castleman and his colleagues have discovered how to trap atoms of metal inside water cages as well. On the preceding page, a water cage traps a cesium atom.
In the real world, such water cages are common in the very cold, sparse regions of the atmosphere 50 miles up. There they probably act as seeds for noctilucent clouds, the eerie, glowing formations that you can sometimes see far to the north.
Ice from Space
Before there was Earth, there was ice. It floated in the cloud of primordial matter that formed the sun and the planets. Made sluggish by the extreme cold of space, some of these superchilled H2O grains bumped into each other, stuck, and eventually grew into the cosmic snowballs we know as comets. Many researchers believe that icy comets colliding with Earth deposited much of the water contained in the early oceans.
At -420 degrees, such space ice is not the everyday stuff you find in your freezer. On Earth, when ice forms, water molecules try to organize themselves into the most stable arrangement possible: each hydrogen atom binds to an oxygen atom to form an orderly, hexagon-shaped lattice. Of course, molecules randomly bumping into each other tend to be highly disorganized, and it takes energy for them to twist and turn to fit into the lattice. However, though an ice cube in a freezer is cold, it is still warm enough for the water molecules to jostle themselves into a preferred position--like a sleeper wriggling to get comfortable under the covers.
But an ice grain numbed by the extreme cold of space is more like a sleeper completely knocked out by sleeping pills. It can’t muster enough energy to even get properly settled. When stray water molecules hit one of these grains, they freeze in that position for good. As a result, the grains grow long tentacles that reach out and grab more passing molecules. The process is so complex that researchers need to simulate it on a computer to get an idea of what’s happening. Chemist Victoria Buch of the University of Illinois at Chicago generated these images by letting the grains aggregate, molecule by molecule, under conditions of deep space.
In the images above, the craggy bit (1) was formed on the computer at a typical -420 degrees. To illustrate how temperature determines the shape of ice, Buch then had the computer warm the ice, giving it the energy to start rearranging itself into a much more compact form (2). A compact grain has less surface area and fewer unbonded atoms to lure passing molecules; therefore, it doesn’t form the tentacles typical of an ice grain in space.
Much of an ice grain’s surface is open territory for stray matter. The blue oxygen atoms and yellow hydrogen atoms (3) are not bonded to water molecules. As a result, other molecules can latch onto the ice (4). The nonwater molecules on this dog-shaped grain are shown as the blue spheres. Once attached to the ice, the chemicals can react with each other to create new molecules. Their union releases excess energy that jettisons the molecules off the grain. These serendipitous recombinations are responsible for all sorts of unexpected compounds that are found in interstellar gas clouds, such as cyanide and ammonia--chemicals that also show up in comets.
The actual docking of an atom or a molecule on an ice grain is no simple matter. Here a hydrogen atom (5, blue sphere) approaches the ice. It dances over the surface of the ice grain, too excited with kinetic energy to land. Gradually the ice absorbs the excess energy.
Exhausted, the atom comes to rest.
Drops and Films
Spill water on paper and it spreads out into a thin film, but spill it on rubber and it beads up into little drops. The reason becomes clear if we look at the molecular level. Water molecules in a drop hold on tightly, pulling each other into the most compact shape possible--a sphere. But if other, surrounding molecules exert an even stronger force than neighboring water molecules, these strangers can tear the drop apart. A surface like paper is hydrophilic (water loving), while rubber is hydrophobic (water fearing). The paper pulls water molecules away from their fellows, dispersing the drop; the rubber leaves them intact. Hydrophilia and hydrophobia are more than annoyances in mixing oil-and- vinegar dressing; these forces are central to life. Cell membranes are layered with hydrophilic and hydrophobic surfaces that precisely control the flow of chemicals in and out.
A typical drop of water contains trillions of molecules. But what happens if you spill just a few? Until recently, researchers couldn’t say; they expected the familiar patterns to break down. Now a computer simulation by Michael Klein and Joseph Hautman at the University of Pennsylvania suggests that a minuscule number of molecules behave just like the more familiar multitudes: films as tenuous as one molecule thick form on a hydrophilic surface (right, top); on a hydrophobic surface, the water still beads into drops, even though as many as half the molecules may reside on the surface (right, bottom).
For their hydrophobic surface Klein and Hautman used hydrocarbon molecules fringed with methyl groups, providing a neutrally charged frosting that couldn’t attract water. Their hydrophilic surface consisted of a layer of OH groups (one oxygen atom bonded to one hydrogen atom), which exert a strong pull on water. The researchers spread 90 molecules of water on each surface. When water was spread on the OH-studded hydrophilic surface, the molecules slowly made bonds with the OH groups. On the methane-covered hydrophobic surface, however, the water bunched itself up from a thin film to the droplet shown here in less than a billionth of a second. This microdrop reflects the form of its macroscopic counterpart down to some of the smallest details. The angle of the curve where the drop meets the hydrophobic surface, for instance, is the same for both macro- and microdrop.