Or, to rephrase the title of this post, "Isn't all ice created equal?" No.
Depending on its environmental conditions and its solute, water can take on a variety of forms. Below is a great article from Chemical Heritage Magazine. Author credit: Sam Kean.
Wild Ice. By Sam Kean.
A story about ice might seem out of place in the summer issue of
Chemical
Heritage, but today we’re going to embrace that topsy-turviness. This is a
story about bizarro ice—ice that burns, ice that sinks instead of floating, ice
literally out of this world. So if you have a chilled drink at hand, take a
gander at the cubes in your glass and let your imagination run because this is a
side of ice you’ve never seen.
The common ice you find in ice cubes—called ice I
h, or “ice
one-h”—is technically a mineral since it is inorganic and has a regular crystal
structure. Specifically, its molecules arrange themselves into a lattice of tiny
hexagons, a six-fold symmetry that ultimately underlies the shape of snowflakes.
Virtually all ice on Earth is ice I
h, and a good thing too! Its
spacious hexagons make it less dense than liquid water; so it floats on lakes
and estuaries and actually insulates fragile aquatic creatures below, protecting
them from wind and chill. Without ice I
h life as we know it wouldn’t
exist.
But talking about ice and mentioning only ice I
h is like talking
about chocolate and mentioning only Hershey’s. Exotic ices are still made up of
H
2O, of course, but the individual molecules break free from the
hexagon straitjacket and reshuffle. Lots of solids can undergo a similar
rearrangement. If you’ve ever opened an ancient Hershey’s kiss and found a tan
and chalky cone inside the wrapper, you’ve seen chocolate do just that. (During
this “chocolate bloom” the cocoa molecules squeeze together, increasing the
chocolate’s density and pushing fat to the surface.) But few solids can form as
many distinct “phases” as ice.
Scientists create different phases of ice by subjecting a tiny sample to
monstrously high pressures, millions of times higher than atmospheric pressure.
And with pressures so high the ice can stay solid at temperatures of thousands
of degrees—a true freezer burn. If you could somehow plop chunks of these ices
into a glass of liquid water, they’d vaporize it. (Imagine the party tricks.) On
a molecular level the high pressure deforms the hexagonal bonds, forcing
H
2O molecules into rhombuses, tetragons, and other alternative
geometries. High pressure can also force H
2O molecules to squeeze
into the holes at the centers of these shapes, trapping them like bugs in tiny
cages. This action increases the density and makes these ices heavy enough to
sink in water. At super-high pressures some chemists predict that ice transforms
into a metal.
Scientists created the first exotic ices, ice II and ice III, around 1900;
the list now extends up to ice XV, discovered in 2009. Creation of these ices is
more than an academic exercise. Ice molecules are held together by the same
hydrogen bonds that, among other things, hold DNA strands together; so forming
new ices helps probe the nature of that bond. What’s more, while ice
I
h dominates the biosphere, other ices exist naturally. An ice that’s
structurally similar to diamonds, ice I
c, probably exists in the
upper atmosphere. The dense, hot interiors of Neptune and Uranus probably
contain chunks of nonhexagonal ices, as do exoplanets around distant stars, a
potentially important consideration as we search for life beyond our solar
system.
In the universe at large, however, ices I through XV are vastly outpopulated
by so-called amorphous ice, ice whose molecules arrange themselves randomly,
without any crystal structure. Amorphous ice forms in flash freezes in deep
space. Microscopic nuggets of this ice also tend to look amorphous since there
aren’t enough molecules to hold themselves together in a regular crystal.
That transition from amorphous ice to crystal ice has long intrigued
scientists, and some of them have even tried to pin down exactly how many
molecules are needed to form genuine crystal ice. That may seem like an
unanswerable question—like asking at what point a man losing his hairs one by
one becomes bald. But believe it or not, an experiment last fall in Germany
actually determined the answer.
The experiment involved slowly adding H
2O molecules to a nucleus
of sodium atoms and probing what wavelengths of infrared light they absorbed.
Amorphous ice showed an absorption peak at a certain wavelength; crystal ice had
a peak at a slightly longer wavelength. The shift from one to the other occurred
surprisingly quickly. Below 250 molecules the amorphous peak dominated. But at
275 molecules the crystal’s longer-wavelength peak began peeking out as a
rudimentary crystal took shape. By 475 molecules that peak alone dominated. So
depending on where you draw the line, as little as 0.000000000000000000008 grams
of water “counts” as crystal ice.
That’s vastly smaller, of course, than even the rapidly shrinking slivers
still left in your glass. They’ll nevertheless soon cross that threshold and
wink out of existence. Midnight will chime, and all these fantasies of burning
ice, ice that sinks, and metallic ice will evaporate. At least until ice
chemists discover the next new wonderful incarnation.
Sam Kean is the best-selling author of The Violinist’s Thumb.
