Stretching liquid to its physical limit.

Stretching liquid to its physical limit

Stretch a confined sample of water enough and it becomes unstable, rupturing violently into vapor. Researchers have now withnessed this bizarre scenario in tiny, water-loaded cavities locked within crystals. Their experiments provide a general route for studying liquids under exotic conditions of "negative pressure" and should yield a fuller understanding of the liquid state of matter, they say.

"In normal liquids, particles are buzzing around banging into each other, exerting forces on each other in a fairly random way," explains chemist C. Austen Angell of the Arizona State University in Tempe. "The average force acting on a particle is randomly directed."

That's why water in a glass doesn't spontaneously lurch into your face.

High pressure, on the other hand, can force a liquid's molecules close enough for short-range repulsive forces to kick into high gear so that the average force acting on a particle points away from the center of the sample. "They're all trying to push out and get to a larger volume," Angell says.

That's why squeezing the trigger of a water pistol yields a stream of liquid.

But Angell and his colleagues focus on what happens to liquids under tension, or negative pressure -- a largely unexplored condition in which the average force on a particle pulls it in toward the center. Attractive rather than repulsive forces govern the stretched liquid's behavior.

I the Aug. 10 SCIENCE, the researchers describe their use of microscopic, liquid-containing cavities in quartz as windows onto stretched liquids. To build tension, they heat and then cool the crystals. When heated, the liquid expands to fill space formerly taken up by vapor bubbles at room temperature. During cooling, some of the liquid clings to the cavity walls as the rest tries to contract, causing a rise in tension. The cooler the temperature gets before bubbles repopulate the imprisoned liquid, the higher the negative pressure gets.

At tensions apparently equivalent to 800 to 1,400 negative atmospheres, Angell and his colleagues observe bubbles of vapor suddenly appearing throughout the liquid. Angell notes that the van der Waals equation -- an enduring 19th-century formula describing gas behavior -- predicts that liquids must break up and form vapor bubbles when stretched beyond a point at which their molecules hover between liquid and vapor states.

For a liquid, Angell says, that point "is like the edge of the world." And when he stretches his sample beyond the edge, "it looks like the whole thing suddenly becomes a froth."