Water has an interesting combination of properties: it is a good sol-vent, it has a large thermal capacity, it is chemically stable (that is, it doesn't light up easily!), and it is in good supply. It also happens to freeze within the range of temperatures at which life is possible. If water did not have these properties, there would be no life-and no hockey. So far in our look at how ice is formed, we've used a nineteenth-century approach: heat is simply extracted until water freezes. But if we really want to find how to make perfect ice (or at least the best ice possible), it helps to take a look at the molecules that compose ice. Richard Feynman, one of the greatest and best-known physicists of the twentieth century, once made this important point: if all scientific knowledge were to be lost except for one principle, he'd like that principle to be the atomic hypothesis, which says that all matter is made of atoms. It's that important of a concept.
If there's one chemical formula most people remember from high school, it is H 2 0. A water molecule (see Fig I. 1) is made of two hydrogen atoms and one oxygen atom. Hydrogen is the most basic atom around, with one positive proton at its core and one negative electron circling on the outside. Oxygen is far more complex, with eight protons in the nucleus and eight electrons arranged in shells around it. The water molecule is roughly shaped like a boomerang: a central oxygen atom is flanked by smaller hydrogen atoms located on either side, forming a 105° angle. The diameter of the molecule is only about 0.5 nanometers, or less than a millionth of a millime-ter. The force holding the water molecule together is described as a covalent bond, in which the three atoms share electrons and be- come closely linked. This tendency to share electrons stems from the fact that atoms are "happiest" (that is, most stable) when they have a complete set of electrons in their outer shell. The reason behind this has to do with quantum mechanics-it can't be explained using classical physics.
There is no overall electric charge to a water molecule, as all atoms inside are neutral. But electrons, quantum-mechanically speaking, are not evenly distributed throughout the molecule, so some regions end up slightly positive and others slightly negative. In other words, the water molecule is polarized. This important aspect is responsible for some of water's unique properties. Polarization affects how neigh- boring molecules interact. As the positive part of one water molecule is attracted to the negative part of another, they stick together. In scientific jargon this is called hydrogen bonding. While they are fairly weaker than covalent bonds inside the molecule, hydrogen bonds are behind some of water's characteristic properties, such as its large heat capaCIty.
When two water molecules are attracted to each other, it takes some energy to keep them apart. One source that can do this is thermal energy. Above 100°C, there's enough thermal energy to break the hydrogen bonds completely, and each molecule goes its own way. Below 100°C, water molecules take up a more limited space, rubbing against neighbors while not being attached to any of them-that is, they are in a liquid state. Liquid water is therefore much denser than vapor, but it still retains fluid properties. As water is cooled below 15°C, the lower thermal energy allows molecules to stick even more closely together, promoting the formation of clusters and chains called polymers (see Fig. 1.2). At 10°C the long wiggly chains contain ten or so molecules. This clumping explains why the density starts to decrease below 4°C. As more molecules join these individual parties, space tends to form between clusters.
Polymers are an example of what physicists call short-range order-ing. Beyond a distance of a few molecular diameters, the positions of molecules are not related anymore, just like magnets cannot interact when moved apart. At O°C the vibrational energy in water is small enough to make long-range ordering possible, so that all molecules can act in concert and crystallize. But making the jump from a dis- ordered (liquid) to an ordered (solid) state comes at a price. For the molecules to arrange themselves into a perfect array, some molecules
must fit into new spaces and new bonds must be formed. The ad-ditional energy released from the formation of these bonds must be removed in order for ice to form, which explains the latent heat of fusion discussed earlier. Even when the temperature falls far below freezing, ice will not form as long as this extra heat remains. Indeed, liquid water can exist below O°C, though it has ro be fairly pure and conditions need ro be just right. This is an example of what is called a supercooled, metastable liquid. By the same token, water can exist above 100°C without necessarily turning into steam. When super-heated water does turn into steam, it does so with a bang, like in a kernel of corn when it explodes into popcorn. Metastable liquids
occur in nature as well. If tree sap could not exist in liquid form well below its official freezing point, no forests would exist in Canada and other places in the "frozen North."
When ice forms, the molecules arrange themselves a bit differently. One question physicists like to ask is, How many neighbors will a single water molecule have, or how "sociable" is it? In all, a water molecule can take no more than four companions: its oxygen atom can bond with two more hydrogen atoms (one on each side), while its own two hydrogen atoms can link with the oxygen of two more molecules. These five molecules form a tetrahedron, a shape popu-lar in elementary geometry textbooks. This configuration is the basicunit of ice's crystal structure (Fig. 1.3), which looks like a web of in- terlocking hexagons. Incidentally, this hexagonal structure is respon-
sible for the starlike shape of snowflakes, as the first few molecules to crystallize align themselves along a tetrahedron and the rest follow.
An ice crystal doesn't look the same in all directions, however.
Layers of molecules compose sheetlike structures. Molecules within a sheet are more tightly bound together than are those in two adjacent layers. It is therefore not surprising that the distance between layers is greater than that between hexagons: 0.734 nm and 0.452 nm, respectively. As we will see later, this sheetlike configuration is an important element in understanding the slipperiness of ice.
Hexagonal ice is not the only possible type of crystallization. More exotic forms, such as cubic ice, have been observed, although only under combinations of pressure and temperature that would not be present on the hockey rink.
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