Cold Water

To the ancients, one of the remarkable things about ice, perhaps the most remarkable, was the very property that had caused my near disaster‑its transparency. To the Greeks, ice was krystallos from kryos, meaning "frost," so the first strong impression seems to have been left by its manner of formation.

Once that was established, however, another property supervened and the word came to be more significant for the connotation of transparency than of cold. After all, anything at all could be cold, but in ancient times few objects were known that were at the same time solid but not opaque.

It followed, then, that when pieces of quartz were discovered, and found to be transparent, they were called krystallos, too, and were considered (at first) to be a form of ice that had been subjected to such intense cold as to have attained permanent solidity and an inability ever to melt again.

Then the word achieved still another change of connotation. One interesting fact about transparent quartz was its surprising regularity in shape. It bad plane faces that met to form clearly defined angles and edges. Consequently, krystallos came to mean any solid with such a regular geometry. From this came our modern word "crystal."

Nevertheless, the older meaning of transparency persists in vestigial fashion. One still hears of the "crystalline spheres" which held the planets in the old Ptolemaic cosmology. This was not because they consisted of solid crystals; heavens, no. It was because they were perfectly transparent so they could not be seen.

And in modem times, the fortuneteller, gazing mystically into a glass sphere, is pretending to see something in her "crystal ball." This is not because the sphere is crystalline in the modern sense, for glass happens to be one of the very few common solids that is not crystalline (and, therefore, not truly solid), but because it is transparent.

And yet, none of that really represents the true wonder of ice. It may seem to have wonders enough. Its mere existence as "hard water" may seem amazing and paradoxical enough to the life‑long inhabitants of tropic climes, and its coldness and transparency may be of interest, but all that is really nothing.

Consider instead something that is often remarked on, to the point, in fact, where it becomes something of a clich6. Have you never heard a statement such as this one: "Like an iceberg, nine tenths of the significance of the remark was hidden"?

Like an iceberg!

Being a non‑traveler, I have never seen a real iceberg, but if I were on a ship and one hove into view (at, I hope, a safe distance) I am sure that the passengers, crowding against the rail to see it, would say to one another, "Just imagine, Mabel (or Harry), nine tenths of that iceberg is under water."

Then I would say, "That's not surprising, ladies and gentlemen. The surprising thing is that one tenth of that iceberg is above water." Naturally, that would mean I would start getting those queer looks that would indicate once again (oh, how many times!) bow much of a nut I appear to be to my beloved fellow‑creatures.

But it's true --

In general the density of any substance increases as temperature goes down. The lower the temperature, the more slowly the atoms or molecules of a gas move, the less forcefully they bounce off each other, and the closer they can crowd to each other. When the kinetic energy of the gas molecules is insufficient to overcome the attractive forces between the molecules (see the previous two chapters), the gas liquefies.

In liquids, the molecules are in virtual contact, but they have enough energy to slip and slide past each other freely. They also vibrate and keep each other at greater distances than would be the case if all were absolutely motionless. As the temperature drops, the vibrations decrease in force and amplitude and the molecules settle a bit closer together. The density continues to increase.

Eventually, the energy of vibration isn't enough to keep the molecules slipping and sliding. They settle into a fixed position and the substance solidifies. The settling is more compact than is possible (usually) in the liquid form but there is still vibration about the fixed position. As the temperature continues to drop, the vibrations continue to die down until they are reduced to a minimum at the temperature of absolute zero (‑273.1o C.). It is then that density is at a maximum.

To summarize ‑ As a general rule, there is an increase in density with decrease in temperature. There is a sudden sharp increase in density when a gas becomes a liquid* and another, but lesser, sharp increase when a liquid becomes a solid. This means that the solid form of a substance, being denser than the liquid form of that same substance, will not float in the liquid form.
* Except at the "critical temperature," something which need not concern us now.

As an example, liquid hydrogen has a density of about 0.071 grams per cubic centimeter, but solid hydrogen has a density of about 0.086 grams per cubic centimeter. If a cubic centimeter of solid hydrogen were completely immersed in liquid hydrogen it would still weigh 0.015 grams and would be pulled downward by gravity. Sinking slowly (against the resistance of the liquid hydrogen) but definitely, it would eventually reach the bottom of the container, or the bottom of the ocean, if there were that much liquid hydrogen.

(You might suspect that the solid hydrogen would melt on the way downward, but not if the ocean of liquid hydrogen were at its freezing point‑and we'll suppose it is.)

In the same way, solid iron would sink downward through an ocean of liquid iron, solid mercury through liquid mercury, solid sodium chloride through liquid sodium chloride, and so on. This is so general a situation that if you took a thousand solids at random, you would be very likely to find that in each case the solid form would sink through the liquid form and you would be tempted to make that a universal rule.

‑ But you can't, for there are exceptions. And of these, by far the most important one is water.

At 100 o C. (water's boiling point under‑ordinary conditions), water is as un‑dense as it can be and still remain liquid. Its density then is about 0.958 grams per cubic centimeter. As the temperature drops the density rises: 0.965 at 90 o C., 0.985 at still lower temperature, and so on until at 40 o C., it is 1.000 grams per cubic centimeter.

To put it another way, a single gram of water has a volume of 1.043 cubic centimeters at 100 o  C., but contracts to a volume of 1.000 cubic centimeters at 4 o C.

Judging from what is true of other substances, we would have every right to expect that this increase in density and decrease in volume would continue as the temperature dropped below 4' C. It does notf

The temperature of 4 o C.* represents a point of maximum density for liquid water. As the temperature drops below that, the density starts to decrease again (very slightly, to be sure) and by the time one reaches 00 C., the density is 0.9999 grams per cubic centimeter, so that a gram of water takes up 1.0001 cubic centimeters. The difference in density at 00 C. as compared with that at 40 C. is trifling, but it is in the "wrong" direction, and that makes it crucial.
* 3.98 o C., to be more accurate.

At 0o C. water freezes if further heat is withdrawn, and by everything we learn from other solidifications we would have a right to expect a sharp increase of density. We would be wrong! There is a sharp decrease in density.

Whereas water at 0o C. has, as I said, a density of 0.9999 grams per cubic centimeter, it freezes into ice at 0o C with a density of only about 0.92 grams per cubic centimeter.

If a cubic centimeter of ice is completely immersed in water, with both at a temperature of 0o C., then the weight of the ice is ‑0.08 grams and there is, so to speak, a negative gravitational effect upon it. It therefore rises to the surface of the water. The rise continues till only enough of it is submerged to displace its own weight (as measured in air) of the denser, liquid water. Since a cubic centimeter of ice at 0 o C. weighs 0.92 grams and it takes only 0.92 cubic centimeters of water at 0 o C. to weigh 0.92 grams, it turns out that when the ice is floating, 92 per cent of its substance is below water and 8 per cent is above.

What we would ordinarily expect, judging from almost all other solids immersed in their own liquid form, is that 100 per cent of the ice would be submerged and 0 per cent exposed. It follows, then, as I said earlier, that the surprising thing is not that so much of an iceberg is invisible, but that so much of it (or, indeed, any of it at all) is exposed.

Well, why is that?

Let's begin with ice. In ordinary ice, each water molecule has four other molecules surrounding it with great precision of orientation. The hydrogen atom of each water molecule is pointed in the direction of the oxygen atom of a neighbor and this orientation is maintained through the small electrostatic attraction involved in the hydrogen bond (as described in the previous chapter).

The hydrogen bond is weak and does not suffice to draw the molecules very close together. The molecules rem‑tin unusually far apart therefore, and if a scale model is built of the molecular structure of ice, it is seen that there are enough spaces between the molecules to make up a very finely ordered array of "holes." Nothing visible, you understand, for the holes are only about an atom or so in diameter.

Still, this makes ice less dense than it would be if there were a closer array of molecules.

As the temperature of the ice rises, its molecules vibrate and move still farther apart, so that its density falls, reaching a minimum of the aforementioned 0.92 grams per cubic centimeter at 0o C. At that temperature of 0o C. however, the molecular vibration has reached the point where it just balances the attractive forces between the molecules. If further heat is added, the molecules can break free and can begin to slip and slide past each other. In doing so, however, some of them fall into the holes.

As ice melts, then, the tendency to decrease the density through increased vibrational energy is countered by the disappearance of the holes, and more than countered. For that reason, liquid water is 8 per cent denser at 0o C. than solid water is.

Even in water at 0o C., however, the loose molecular arrangement in ice hasn't utterly vanished. As the temperature rises still higher, there is still a slow disappearance of the last few lingering holes and it is not till a temperature of 4o C. is reached that so few of them are left that they can no longer exert a dominating effect on the density change. At temperatures higher than 4o C., the energy of molecular vibration increases and density decreases steadily as it "ought" to do.

The importance of this density anomaly in water simply can't be exaggerated. Consider what happens to a moderately sized lake, for instance, during a cold winter.

The temperature of the water gradually drops from its mild warmth of the summer. Naturally, it is the water at the surface that cools first, becomes denser, and sinks, forcing up the warmer water at the bottom, so that it can, in turn, cool and sink. In this way the entire body of water cools, and would cool all the way to 0o C. if the density continued to increase steadily as temperature dropped.

As it is, though, when a temperature of 4 o C. is reached, a further cooling of the surface water makes it slightly less dense! It does not sink, but floats on the warmer water below. The surface water drops in temperature all the way to 0 o C., but heat leaves the lower depths only slowly and those depths remain somewhat warmer than 0 o C.

It is the water at the surface, then, that experiences freezing, and the ice, being less dense than water, remains floating. If the cold weather continues long enough, the entire layer of surface water freezes and forms a solid coating of ice that may become thick and strong indeed (to the satisfaction of ice skaters).

But ice is a good insulator of heat, and the thicker it is, the more effective an insulator it is. As it thickens, the deeper layers of water (still liquid) lose heat through the ice to the air above more and more slowly; and more and more slowly does the ice layer thicken further. In short, in any winter that is likely to occur on Earth, a sizeable lake will never freeze solidly all the way to the bottom. This means that life‑forms in it can survive through the winter.

What's more, when the warm weather returns, it is the surface ice that receives the brunt of the Sun's heat. It melts and the liquid water beneath is at once exposed, so that the lake quickly becomes liquid throughout once more.

What would happen, though, if water were like other substances. In cooling, there would be a continual sinking of cooler water all the way down to 0 o C., so that the entire body of the lake would be at that temperature eventually. It would have a tendency to freeze at every point, and any ice that formed near the surface of the lake would sink at once if there were still liquid below it. A winter that under present circumstances would only suffice to form a thick scum of ice on a lake would be enough to freeze that same lake solid, top to bottom, if water were like other substances.

Then, when warm weather came, the surface of the frozen lake would melt, but the water that formed would insulate the deeper layers of ice from the Sun's beat. The thicker the layer of liquid water, the more slowly the Sun's heat would penetrate to the ice below and the more slowly would the deeper ice melt. Through an ordinary summer such as we experience on Earth, a solidly frozen lake would never melt all through. Most of it would remain permanently frozen.

The same would hold true for rivers and for the polar oceans. Indeed, if water were suddenly to change its density characteristics, each winter would see further ice form and sink to the ocean abyss to remain permanently frozen thereafter. Eventually, all of Earth would be a mass of ice‑bound land, with a thin layer of water on the surface of the tropic ocean.

Even though such an Earth would be at the distance from the Sun it now is, and would receive the amount of solar energy it now does, it would be a frigid world and life as we know it would not have formed. It follows then that life depends 'on the hydrogen bond not merely for the reasons I outlined in the last chapter, but because of the loose structure it gives ice.

There's another way to break down the holes in ice besides raising the temperature. Why not simply squeeze the ice together under pressure? To be sure, it takes enormous pressures to squeeze out the holes to the point where ice is as dense as water. (When water is allowed to fill a sealed container tightly and then made to freeze, it exerts an outward pressure equal to the pressure it would take to compress ice to the density of water‑and the container breaks.)

Still, high pressures can be produced in the laboratory. About 1900, a German physicist, Gustav Tammann, began to make use of such high pressures, and beginning in 1912, an American physicist, Percy W. Bridgman, carried the matter much further.

In this way, it was found that there were many forms of ice.

In any solid there is an orderly arrangement of molecules and there is always the possibility of a variety of different arrangements under different conditions. Some arrangements are more compact than others and these would be favored by high pressures and low temperatures.

Thus, under ordinary temperatures and pressures, ordinary ice (which we can call Ice I) is the only variety that can exist. As the pressure is increased, however, two other forms are found, Ice II at temperatures below ‑35o C., and Ice III at temperatures between ‑35 o C. and ‑20 o C.

If the pressure is raised still further, Ice V is formed. (There is no Ice IV; it was reported but proved to be a case of mistaken observation and was dropped; but not before Ice V had been reported.)

If the pressure is raised still further, Ice VI and Ice VII are formed. Whereas all other forms of ice exist only at 0o C. and below, Ice VI and Ice VII can exist at temperatures above 0 o C., though only at enormous pressures.

In fact, at a pressure of 20,000 kilograms per square centimeter (one and a half million times the pressure of the atmosphere), Ice VII will exist at temperatures above 100 o C., the boiling point of liquid water under ordinary conditions.

All these high‑pressure forms of ice are denser than liquid water, as you would expect, for the holes have been squeezed out of them. Indeed, of all known forms of ice, only Ice i, the ordinary variety, is less dense than liquid water.

It would follow that if any of the forms of ice other than Ice i could form in the oceans, they would sink to the bottom and gradually accumulate.

In one of his excellent novels, Kurt Vonnegut hypothesized a mythical "Ice IX," which could exist at the ocean bottoms and which would form spontaneously if only some small quantity existed as a "seed." The hero had such a small piece and, of course, it got into the ocean to bring about the final catastrophe.

Is there really a chance of that? No. Any form of ice but Ice i can only exist at enormous pressures. Even the least high‑pressure ices (Ice II and Ice III) can exist only at pressures more than two thousand times that of the atmosphere. If such pressures could be attained at the bottom of the ocean (they can't), a further requirement would be that the temperature be well below ‑20 o C. (It isn't.)

You can see, further, that no form of ice other than Ice I could exist in someone's pocket. If any other ice were formed and the high pressure required to form it were removed, the ice would instantly expand to Ice i with explosive violence.

That still leaves one thing to discuss. Though solid forms of a substance can (and often do) exist in a variety of crystalline forms, liquid and gaseous forms do not. In liquids and gases there is not, generally speaking, any orderly array of molecules, and one does not find varieties of disorder.

But in 1965, a Soviet scientist, B. V. Deryagin, studied liquid water in very thin capillary tubes and found some of it to possess most unusual properties. For one thing, its density was 1.4 times that to be expected of ordinary water. Its boiling point was extraordinarily high and it could be heated up to 500 o C. before ceasing to be liquid. It could be cooled down to ‑40 o C. before turning into a glassy solid.

The report was largely disbelieved in the West, where there is almost automatically skepticism toward any unusual finding that emerges outside the charmed circle of nations prominent in nineteenth‑century science.

However, when Americans repeated Deryagin's work, they found, much to their own surprise, that they got the same results and could even see droplets of the anomalous form of liquid water‑droplets so small they could be made out only under a microscope.

What was behind this?

Water molecules, while slipping and sliding around each other, do tend to take up the hydrogen bond orientation, as in ice. This happens over very small volumes and for very brief periods, but it is enough to make liquid water behave as though it consisted of submicroscopic particles of ice that form and unform with superspeed.

The "ice" never forms over a volume large enough and for a time long enough to make the holes significant and cause water to be as un‑dense as ice, but it does keep the water molecules far enough apart to allow hydrogen bonds to form and unform. Liquid water is therefore less dense than it might be.

Suppose, though, that pressure is placed on water in such a way that molecules are forced closer together while in the hydrogen bond orientation. With neighboring molecules unusually close, the hydrogen bond would be much stronger than ordinary and would, indeed, approach an ordinary chemical bond in strength. Molecule after molecule would fall into place and, thanks to the unusually strong hydrogen bond attractions. they would make up a kind of giant molecule built up out of the small water-molecule units.

When small units build up a giant molecule in this fashion, the small units are said to "polymerize" and the giant molecule is a "polymer." The new form of water was therefore spoken of as "polymerized water" or, for short, "polywater."

In polywater, the molecules are in orderly array, something as in ice, but in much more compact fashion, and certainly without the holes. Not only does this compact array of water molecules produce a substance considerably more dense than ice, but considerably more dense than ordinary liquid water as well.

What's more, because the molecules are held more tightly together, it takes a much higher temperature than 100 o C. to tear them apart and make polywater boil. It also takes a much lower temperature than 0o C. to force the molecules apart into the less compact array of ordinary ice. Other unusual properties of polywater are also easily explained on the basis of the compact array of molecules.

Apparently, polywater does not form under ordinary increases in pressure, but does form in the constricted volume of tiny capillary tubes. Biologists at once began to wonder whether within the constricted volume of tissue cells, polywater also formed; and whether some of the properties of life could not be most easily explained in terms of polywater.

I wish I could end the matter here, with this glamorous discovery and the still more glamorous speculation, but I can't. The trouble is that many chemists remain skeptical of the whole business.

It is possible, after all, that investigators have been mis
led by the chance of solution of the glass from , the tubes in
which the polywater was being studied. If it were not pure
liquid water they were studying but tiny volumes of glass
solution, all bets were off.

Indeed, one chemist recently prepared a solution of silicic acid (something which could form when water is in contact with glass) and reported it to possess the very properties of polywater.

So it may be that polywater is a false alarm, after all.