January/February 1998 Table of Contents
NASA Ames Research Center
is a very special planet indeed. It is the only place in the solar
system where you can pummel your little brother (or, more likely,
get pummeled by him) with snowballs.
made the Walker "kneel down" a few times - that is, lowered
the bubble - to warm up the hydraulic fluids. It's hard to remember
that the legs of the Cat are working at temperatures a hundred degrees
below freezing, when you're sitting in a toasty cabin, but it can
be dangerous to forget. While I was doing this I looked out at the
life dome rising in the distance. I could pick out people sledding
down a hill and further away a crowd in a snowball fight. A scramble
like that is more fun on Ganymede than on Earth; somebody a hundred
yards away can pick you off with an accurate shot, because low gravity
extends the range of your throwing arm.
I first read the 1984 book Jupiter Project by Gregory Benford,
one of my favorite science-fiction authors, I liked the image of
the future it presented - people taking their humanity with them
to space along with their technology. Alas, this feeling only lasted
a moment, for I quickly realized that a snowball fight is not possible
on Ganymede. Nor, it turns out, is any kind of snowball fight possible
on any planet or moon in our solar system, with the sole exception
of snowball fights using ordinary H2O snow on Earth. The spoilsports
are the physics of snowball manufacture and certain peculiarities
of the H2O molecule.
trouble stems from how water and other materials behave when squeezed,
heated, or cooled. Figure 1, a "phase diagram," summarizes
this behavior for H2O and carbon dioxide (CO2). On a phase diagram,
the horizontal axis represents temperature, with higher temperatures
to the right. The vertical axis represents pressure, with higher
pressures at the top. The different areas of the plot represent
conditions under which the material is gaseous, liquid, or solid.
The thick solid lines represent the transitions between these phases.
example, suppose you put an ice cube in a frying pan and turn on
the burner. First the ice cube melts and then the resulting puddle
boils away. On the phase diagrams in Figure 1, this is represented
by starting at point 'A' and moving horizontally to the right (higher
temperature, same pressure) to 'B' and then 'C'.
behavior differs somewhat if the ice cube starts at a low enough
pressure. In this case, raising the temperature does not cause it
to melt, but instead evaporate directly, or "sublimate."
For water ice to sublimate, the pressure must be lower than 0.6
percent of standard atmospheric pressure; for CO2, lower than 5.1
atmospheres. Therefore, dry ice (CO2 ice) doesn't melt when it warms
up at atmospheric pressure, but instead gives off gas directly.
On the phase diagrams in Figure 1, ice sublimates if it starts at
point 'D' and warms up at constant pressure to 'E' and 'F'.
the process of sublimation is what makes your ice cubes shrink with
time in your frost-free freezer. Although the pressure is much higher
than 0.006 atmospheres, the dehumidifier in a frost-free freezer
removes water vapor from the air - preventing the ice and vapor
from reaching the equilibrium implied in the phase diagram.
about changes in pressure? Suppose we start with gas and squeeze
it while holding the temperature constant (that is, move vertically
upward on the plot). If the temperature is low enough, the gas will
turn into a solid ('E' to 'A'); at higher temperatures, to a liquid
('F' to 'B').
have probably heard the warning against removing the cap from your
car radiator when the engine overheats. This is because your radiator
contains liquid water that is both very hot and under pressure.
Removing the cap allows the pressure in the radiator to drop precipitously
(from, say, 'B' down to 'F'). As a result, the contents of your
radiator try to change instantly to steam - your radiator boils
over and you are in danger of being badly scalded. People who work
with high-pressure air lines have to deal with these pressure effects,
too. If humid air seeps into high pressure lines, the H2O vapor
in the air can condense into liquid within the lines. This is one
reason high-pressure lines often have problems with corrosion.
1. The low-pressure phase diagrams of water (H2O, left) and
carbon dioxide (CO2, right). The different areas of these diagrams
represent conditions under which these substances are gaseous, liquid,
or solid. At points 'A' and 'D', they are solid; at 'C', 'E', and
'F, gaseous; at 'B', liquid. The thick solid lines represent transitions
between phases. Along these lines, two phases can exist together.
At the triple point, where the lines intersect, the solid, liquid,
and gas can all co-exist. For H2O, the triple point falls at a temperature
of 0.0099 degrees Celsius and a pressure of about 0.006 times the
pressure of Earth's atmosphere at sea level. For CO2, it falls at
-56.6 degrees and 5.1 atmospheres. The heavy dashed lines represent
sequences of heating (horizontal) or squeezing (vertical):
Start with the ices at 'A'. Keeping the pressure constant, warm
them until they reach their solid-liquid transitions. At this point
the ices melt. Further heating initially results in warmer liquids
('B'). If you warm the liquids enough, they reach their liquid-gas
transitions and evaporate. Further heating results in warmer gases
Start with the ices at 'D'. Raising their temperature does not cause
them to run into their solid-liquid transitions, but rather into
their solid-gas transitions. The ices evaporate directly into the
gas phase in a process called sublimation. Further heating warms
the resulting gases ('E' and 'F').
Start with gases at 'E'. Holding their temperature constant, squeeze
them until they reach their solid-gas transitions. At this point
the gases condense directly into solids ('A'). For CO2 and most
other materials, more pressure will not produce any additional changes
('G' on right diagram). For H2O, however, additional pressure pushes
the solid to the solid-liquid transition, where it melts ('G' on
Start with gases at 'F'. Squeezing them pushes them to their solid-liquid
transitions, where they condense directly into liquids ('B'). For
CO2 and most other materials, further squeezing pushes the material
to the solid-liquid transition, where it solidifies ('H'). For H2O,
squeezing the liquid doesn't do anything (until reaching extremely
high pressures off the top of the diagram). Diagram by Scott A.
2. Making a snowball. First, squeeze the snow hard enough that
it moves vertically up the phase diagram to the solid-liquid transition
('A' to 'B'). Then release the pressure ('B' to 'A'). While at the
solid-liquid transition ('B') some of the edges of the ice crystals
melt. When you release the pressure, they refreeze, cementing the
snow together. If you squeeze too hard, the snow goes all the way
up to 'C' and the snowball melts completely away. This is one reason
your gloves get so wet when you make snowballs on a warm day.
a very cold day, the snow may start at location 'A´' and proceed
up to 'B´' and 'C´'. At 'C´', you are squeezing as
hard as you can, but the pressure is insufficient to reach the solid-liquid
transition. Because no melting occurs, when you release the pressure
you are left with a loose handful of snow rather than a snowball.
Diagram by Scott A. Sandford.
far, the phase transitions discussed have been identical for H2O
and CO2. The two diagrams in Figure 1 indeed look very similar.
But there is a crucial difference: namely, the slope of the nearly
vertical line between the solid and liquid phases. For CO2, this
line slopes upward and to the right - the normal behavior of almost
all materials. However, for H2O, the line slopes upward and to the
left. This seemingly minor change reflects major differences in
the behavior of H2O under pressure.
happens if we take gaseous CO2, squash it at a temperature at which
it will liquefy (above -56.6 degrees Celsius), and squeeze harder
and harder? It will turn into a liquid and eventually into a solid.
In Figure 1, this is represented by the line 'F' to 'B' to 'H'.
If we crank up the pressure still further, nothing new happens;
we move up the diagram, but CO2 remains solid. Below -56.6 degrees
Celsius, squeezing turns CO2 gas directly into a solid without the
liquid go-between ('E' to 'A' to 'G'). Thus, if you squeeze CO2
and most other substances hard enough, they eventually solidify.
behaves very differently. If you start with water vapor hotter than
0.0099 degrees Celsius and squeeze, the gas condenses into liquid,
just as CO2 did. After this point, however, additional squeezing
does not turn the liquid water into solid ice. It stays liquid.
(Under extreme conditions, this is not exactly true; I'll revisit
this issue later.) In Figure 1, this process is represented by the
line 'F' to 'B' to 'H'. Water also behaves differently at lower
temperatures, below 0.0099 degrees Celsius. Applying pressure causes
the gas to condense directly into a solid, as with CO2. But if you
apply more pressure, the H2O ice melts. In Figure 1, this is represented
by the line 'E' to 'A' to 'G'. This behavior is very different from
that of CO2, which stays solid no matter how hard you squeeze. If
you squeeze H2O hard enough, it eventually liquefies.
peculiarity is associated with the fact that solid H2O is less dense
than liquid water. Water ice floats in water - a very unusual behavior.
Dry ice, for example, sinks in liquid CO2. Most solids are essentially
as dense as those substances can possibly be, and further squeezing
cannot make them significantly more compact. Yet when you squeeze
H2O ice, it can ease the squeeze by turning into a liquid, making
itself smaller and denser.
(almost) unique property is in part due to the water molecule's
ability to form strong hydrogen bonds. In hydrogen bonding, the
H atoms in H2O form attachments to adjacent molecules. Such bonding
is weaker than the covalent or ionic bonding that holds together
molecules, but stronger than the normal intermolecular "van
der Waals" forces. Hydrogen bonds are like the sticky portions
of Post-It notes - stickier than just pushing two pieces of ordinary
paper together, but less sticky than glue. It is hydrogen bonding
between H2O molecules that makes the folds in your wet shower curtain
stick together. This ability to grab things firmly but not tightly
is one of the reasons that H2O is so very important for life. Because
of it, liquid water can mediate much of the complex chemistry that
occurs in living organisms.
Can Be Welded
on this understanding of H2O and its phase diagram, let's consider
the physics of snowball manufacture. To make a snowball, you scoop
up some snow and compress it between your cupped hands, after which
you hopefully have a ball that hangs together. What makes the loose
snow stick together?
people assume that the heat from their hands melts enough of the
snow to make it cohere. However, if heat were responsible, a snowball
would be made from the outside in, and you'd have to hold it long
enough for sufficient heat to enter. But if you cut a snowball in
half, it looks pretty much the same throughout; there's nothing
special about its outer layers to indicate they were heated first.
Moreover, snowballs can be compacted quickly.
heat mattered, it would be easier to make a snowball with your bare
hands than with gloves on. This is clearly not the case (thank goodness).
If you need further proof, try this: Drop two spoons in the snow
and let them chill to snow temperature. Then pick the spoons up
with gloved hands, so that you don't warm them, and use them to
pack a snowball. You can make a perfectly good snowball this way,
although the delay may cost you the snowball fight.
some people theorize that snowballs hang together because the intricate
ice crystals in the snowflakes get tangled. If so, fluffy snow would
make better snowballs than denser snow. In fact, the opposite is
often the case.
real explanation is revealed by the H2O phase diagram (see Figure
2). Suppose you are strong enough that you can squeeze snow hard
enough to raise its pressure from 'A' to 'C' or, equivalently, 'A´'
to 'C´'. If the snow is warm enough, you can squeeze it sufficiently
hard that it reaches the solid-liquid transition line. At this point,
some of the ice crystals melt in order to take up less space. If
you then release the pressure, the snow drops away from the solid-liquid
transition line, and the ice that melted refreezes. In Figure 2,
this amounts to moving from 'A' to 'B' and back to 'A'. The refreezing
liquid bonds the ice crystals together. Voilà! A snowball.
suppose you want to have a snowball fight on a very cold day. As
you squeeze the snow it proceeds vertically up the phase diagram,
but because of the low temperature, you can't squeeze hard enough
to reach the solid-liquid transition. You're squeezing as hard as
you can ('A´' to 'B´' to 'C´'), but you can't get
any melting. When you release the pressure, you find that you are
holding a loose handful of snow, rather than a snowball.
how cold is too cold? It depends on how hard you can squeeze. Some
quick-and-dirty measurements with scales in my laboratory indicate
that my hand can apply a pressure of about 0.07 atmospheres (1 pound
per square inch, or 7 kilopascals) without straining too much. I
can get as high as 0.2 atmospheres if I use my arms to squeeze instead
of my hands. Thus, squeezing a snowball only adds a minor pressure
to the pressure that the air is already exerting on the snow (1
atmosphere). Consequently, you can't expect to make a snowball with
your hands if the snow is more than a few degrees below freezing.
lower temperatures, you would need herculean pressures. At -25 degrees
Celsius (-13 degrees Fahrenheit), a summer's day at the South Pole,
it would take 2,000 atmospheres (30,000 pounds per square inch)
to make a snowball. This would crush a submarine, let alone a handful
of snow. Maybe that's why I've never participated in a snowball
fight during any of my trips to Antarctica.
are even worse on deep-freeze worlds such as Ganymede, Pluto, and
Mars. After Earth, Mars is the next warmest place in the solar system
where there is any substantial amount of water ice. But with typical
temperatures between -120 and -40 degrees Celsius (-180 and -40
degrees Fahrenheit), the snows of Mars are way too far to the left
of the phase diagram to mold with our hands. Even the warmest temperature
recorded by Mars Pathfinder, -9 degrees Celsius, would preclude
wait, you say, what if I were willing to drag around a hydraulic
press? Could I make snowballs on Mars then? Sorry. So far, we have
only considered the low-pressure part of the phase diagram. At very
high pressures, H2O gets weirder still. Figure 3 shows that H2O
can actually exist in a number of different solid forms depending
on the temperature and pressure. The entire phase diagram shown
in Figure 2 corresponds to a tiny horizontal strip near the bottom
of Figure 3.
3. The high-pressure phase diagram of H2O. On this pressure
scale, the entire area represented by Figure 1 or 2 - including
the entire liquid-gas transition line - is an infinitesimal slice
along the lower axis. At very high pressures, H2O can exist in exotic
solid forms: Ice II, Ice III, Ice IV, and so on. Everyday ice is
Note that the solid-liquid transition line does not extend up and
to the left forever. Instead it meets Ice III and then turns to
the right, as for normal materials. For this reason, if you squeeze
liquid H2O really hard, it ultimately turns into one of the exotic
solid phases. There is no region of stability for liquid water below
-25 degrees Celsius (-13 degrees Fahrenheit). Crushing snow below
this temperature does not cause any melting. Diagram by Scott A.
Sandford, adapted from General Chemistry by Gordon Barrow, p. 311.
everyday ice in our sodas is called Ice I. At extremely high pressures,
H2O ice can take on exotic crystalline forms: Ice II, Ice III, and
so on. (Fortunately there isn't a form of ice called Ice IX which
has the apocalyptic properties Kurt Vonnegut described in "Cat's
Cradle". The solid-liquid transition line does not extend up
and to the left forever. Instead it collides with the area of stability
of Ice III and then turns and heads up and to the right, as for
most other materials.
the way, water also acts strangely at very low temperatures, off
the left edge of Figure 3. When it is very cold - as in comets and
interstellar molecular clouds - non-crystalline ice can form (see
Winter Olympics on Mars
high-pressure ice phases and the reversal of solid-liquid transition
line have several implications for the subject at hand. I stated
earlier that if you started with H2O gas or solid and squeezed,
it would eventually turn into a liquid, and stay a liquid. That's
not quite right. If you squeeze liquid water really hard (more than
4,000 atmospheres) it will ultimately transform into Ice V, VI,
or VII, depending on the temperature.
important, below -25 degrees Celsius, squashing snow (Ice I) doesn't
result in any melting; it just turns Ice I into ices II, III, and
so on. No melting, no snowball. If you were desperate for a snowball
fight on Mars on any but the very warmest days, you'd need a heated
hydraulic press. Of course, you could always throw solid chunks
of ice cut out of the ground, but we all know that our mothers wouldn't
let us do that.
afraid this will also make it difficult to ice-skate on Mars. When
you skate, you are gliding on a thin layer of water that forms between
your blade and the ice. The weight of your body on the skate blades
puts pressure on the ice, which helps to form the water layer. To
be sure, the formation of the water layer is more complicated than
a simple pressure effect. (See Samuel Colbeck's article in the October
1995 issue of the American Journal of Physics.)
snowball fights in our solar system are out, except on Earth. How
about "snowball" fights using something other than frozen
H2O, say, methane or carbon monoxide on Pluto or carbon dioxide
on Mars? Alas, those other compounds have unfavorable phase diagrams.
No amount of compression will get them to melt and pack into snowballs.
Benford's scene just isn't possible. I find this more than a little
sad. Still, it makes you appreciate all the more how special Earth
if you're like me, you are already wondering whether there aren't
any other materials that behave like water - with solid-liquid transition
lines that slope up and to the left. There are, but only under pressure
and temperature conditions in which humans couldn't survive. For
example, plutonium can exist in a form that melts when squeezed.
If you had plutonium dust at 640 degrees Celsius (1184 degrees Fahrenheit)
you might be able to squeeze it hard enough between your hands to
get it to melt. Thus, if you were capable of surviving in a high-pressure
blast furnace and had a strong grip, you could make a plutonium
snowball out of plutonium snow. Of course, you would have to be
extremely careful about it. If you made it too big, above critical
mass, the resulting explosion would end the snowball fight for good.
there you have it, the reason you can't have a snowball fight on
Mars. At least on present-day Mars. A little terraforming might
change that conclusion, but that's another article.
on thick ice. On a 1988-89 expedition to Antarctica, geologist Ralph
Harvey took a break from meteorite-hunting to try on his skates.
Skating on the world's largest ice sheet is even tougher than usual,
because the cold temperatures make it difficult for the blades to
generate a layer of water on which they can glide. This challenge
is compounded by the rough surface of the ice and by the wind. It
takes a few minutes to skate a mile downwind and the better part
of an hour to return upwind. Photo by Roberta Score (perhaps best
known as the discoverer of the martian meteorite Allan Hills 84001).
shape. These electron-microscope pictures (left column) and electron
diffraction patterns (right) reveal the structure - or structurelessness
- of ices on scales of a millionth of a meter (µm). Amorphous
ices, formed when water vapor and other gases are deposited at low
temperatures and pressures, exhibit no physical structure, as evidenced
by their fairly featureless electron diffraction pattern (top row).
But as these ices are warmed, the molecules within them shift into
more organized arrangements. The resulting structures manifest themselves
as visible crystals, which scatter electrons into well-defined directions
(center and bottom). Photos courtesy of David Blake, NASA Ames Research
most materials, water ice comes in both crystalline and amorphous
forms. In crystalline form, the atoms and molecules are arranged
in a well-defined, repeating structure. This regular ordering minimizes
the structure's energy and gives rise to favored directions for
strength, cleavage, light refraction, and other properties. Examples
of crystals include diamonds, quartz, table salt, and sugar.
contrast, the atoms and molecules in amorphous materials are jumbled
together without any long-range order. These materials do not have
favored directions. Examples include glass and soot. Amorphous materials
are "metastable" - that is, their atoms and molecules
would actually prefer to rearrange into crystals if they could.
At low temperatures, for example, atom and molecules lack the energy
to move into a regular structure.
made from H2O can exist in a large number of crystalline forms depending
on the temperature and pressure. The ice in your refrigerator (temperature
slightly below freezing, pressure about 1 atmosphere) is referred
to as hexagonal ice because the molecules stack in a form that has
a six-sided symmetry. At lower temperatures (110 to 220 kelvins,
about -160 to -50 degrees Celsius), the ice prefers to stack with
a four-sided symmetry and is called cubic ice. But if you form an
ice by condensing water vapor at still-lower temperatures, the result
the past decade, scientists largely ignored amorphous ice, in part
because it is harder to make, in part because nobody expected random
jumbles of H2O molecules to be very interesting. But this conventional
wisdom has now been overturned. Just as there are many forms of
crystalline ice, there are multiple forms of amorphous ice. The
different crystalline forms occur when the molecules are stacked
in different ways. You might think that amorphous ices, which have
no stacking order, are all be the same. Jumbled is jumbled, right?
not. At low pressures, amorphous ice comes in at least three forms:
a high-density form (stable between 10 and 38 kelvins); a lower-density
form (38 to 131 kelvins); and a third form (131 to 148 kelvins).
In addition, this summer Peter Jenniskens and his colleagues at
NASA Ames Research Center found that small chunks of amorphous ice
may persist even after an ice crystallizes.
amorphous form has its own special kind of disorder. It is analogous
to a pile of wooden blocks. Suppose you dump a box of blocks onto
the floor in a big pile. The blocks in the pile have no particular
ordering, just like high-density, low-temperature amorphous ice.
Now scatter the blocks so that no block is on top of another. There
is still no long-range order, but the disorder is now very different.
This represents the lower-density, higher-temperature amorphous
astrophysical sites experience low pressures, low temperatures,
or both - ideal conditions for forming amorphous ice. Comets are
one example. If comets are dominated by amorphous ice, it could
have significant implications for cometary activity, such as the
jets that comets squirt into space. When amorphous ice warms up
and crystallizes, the molecules rearrange themselves into more-stable,
less-energetic orientations. As a result, energy is released. Thus,
as a comet approaches the Sun, its ice may transform from amorphous
to crystalline, generating energy. This energy further warms the
ice and accelerates cometary activity.
ice may also be important in the interstellar medium where it coats
dust grains in dense molecular clouds, the birthplaces of stars.
These amorphous ices, like those in comets, aren't just made up
of H2O. They also contain methanol, methane, carbon monoxide, carbon
dioxide, and ammonia. In dense molecular clouds, the ices are irradiated
by ultraviolet photons and charged particles. As seen in laboratory
experiments, such radiation creates complex organic compounds in
the ice. These compounds, under certain conditions, are known to
transform into amino acids as well as other chemicals of biological
interest. Once produced in the interstellar medium, some of these
materials will wind up in new planetary systems.
year, Louis Allamandola and his colleagues, also at Ames, discovered
that the organic residues made in this manner have a tantalizing
property: When dropped in liquid water, they spontaneously organize
into droplets that contain internal structure, are surrounded by
membrane-like walls, and are capable of processing light energy.
In short, they exhibit some of the properties of proto-cells. Maybe
amorphous ice has played a significant role in the origin of life
- here and throughout the Galaxy.
is a researcher at the NASA Ames Research Center in Mountain View,
Calif. Among other things, he is working on the Stardust comet sample-return
mission, due for launch in 1999 and return to Earth in 2006. Sandford
was on the team that found Allan Hills 84001, the martian meteorite
that may contain signs of past life on Mars [see Editorial, September/October
1996, p. 2]. His email address is email@example.com.