Water is almost unique among the more than 15 million known chemical substances in that its solid form is less dense than the liquid form. The plot at the right shows how the volume of water varies with the temperature.
The large increase (about 9%) at freezing shows why ice floats on water and why pipes burst when they freeze. The expansion between –4° and 0° is due to the formation of larger clusters. Above 4°, thermal expansion sets in as the thermal vibrations of the O—H bonds becomes more vigorous which pushes the molecules farther apart.
The other widely cited anomalous property of water is its high boiling point. As this graph shows, a molecule as light as H2O “should” boil at around –90°C. That is, it should exist in the world as a gas rather than as a liquid, if H-bonding were not present. Notice that H-bonding is also observed with fluorine and nitrogen.
Have you ever watched an insect walk across the surface of a pond? The water strider takes advantage of the fact that the water surface acts like an elastic film that resists deformation when a small weight is placed on it. (If you are careful, you can also “float” a small paper clip or steel staple on the surface of water in a cup.) This is all due to the surface tension of the water. A water molecule within the bulk of a liquid is attracted to neighboring molecules in all directions. But since these charges average out to “zero”, there is actually no net force on the molecule. For a molecule at the surface of a liquid, the situation is quite different. The surface molecules experience forces only sideways and downward, and this is what creates a stretched-membrane effect.
The distinction between molecules located at the surface and those deep inside a liquid is especially prominent in H2O because of water’s very strong hydrogen-bonding forces. The difference between the forces of a molecule at the surface and one in the bulk liquid gives rise to the liquid’s surface tension.
This drawing at the left highlights two H2O molecules, one at the surface, and the other in the bulk of the liquid. The surface molecule is attracted to its neighbors below and to either side, but there are no attractions above the surface. As a consequence, a molecule at the surface will tend to be drawn into the bulk of the liquid. But since there must always be some surface, the overall effect is to minimize the surface area of a liquid. The geometric shape that has the smallest ratio of surface area to volume is the sphere, so very small quantities of liquids tend to form spherical drops. As the drops get bigger, their weight deforms them into the typical tear shape.
Take a plastic mixing bowl from your kitchen, and splash some water around in it. You will probably observe that the water does not cover the inside surface uniformly, but remains dispersed into drops. The same effect is seen on a dirty windshield and turning on the wipers simply breaks hundreds of drops into thousands. By contrast, water poured over a clean glass surface will wet it, leaving a uniform film.
When a liquid is in contact with a solid surface, its behavior depends on the relative magnitudes of the surface tension forces and the attractive forces between the molecules of the liquid and of those comprising the surface the water is in contact with. If an H2O molecule is more strongly attracted to its own “kind”, then surface tension will dominate, increasing the curvature of the water drop. This is what happens at the interface between water and hydrophobic surfaces like plastic mixing bowls or windshields coated with oily residues. A clean glass surface, by contrast, has -OH groups sticking out of the surface that readily attach to water molecules through hydrogen bonding. This causes the water to spread out evenly over the surface, or to “wet” it. A liquid will wet a surface if the angle at which it makes contact with the surface is more than 90°. The value of this contact angle can be predicted by understanding the separate properties of the liquid and solid.
If we want water to wet a surface that is not ordinarily wettable, we add a detergent to the water to reduce its surface tension. A detergent is a special kind of molecule in which one end is attracted to H2O molecules but the other end is not; the latter ends stick out above the surface and repel each other, canceling out the surface tension forces due to the water molecules alone.
To a chemist, the term “pure” has meaning only in the context of a particular application or process. The distilled or de-ionized water we use in the laboratory contains dissolved atmospheric gases and occasionally some silica, but their small amounts and relative inertness make these impurities insignificant for most purposes. When water of the highest obtainable purity is required for certain types of exacting measurements, it is commonly filtered, de-ionized, and triple-vacuum distilled. But even this “chemically pure” water is a mixture of isotopic species: there are two stable isotopes of both hydrogen (H1 and H2, often denoted by D) and oxygen (O16 and O18) which give rise to combinations such as H2O18, HDO16, etc., all of which are readily identifiable in the infrared spectra of water vapor. (Interestingly, the ratio of O18/O16 in water varies enough from place to place that it is now possible to determine the source of a particular water sample with some precision.) And to top this off, the two hydrogen atoms in water contain protons whose magnetic moments can be parallel or antiparallel, giving rise to ortho- and para-water, respectively. The two forms are normally present in an o/p ratio of 3:1.
It has recently been found (Langmuir 2003, 19, 6851-6856) that freshly distilled water takes a surprisingly long time to equilibrate with the atmosphere, that it undergoes large fluctuations in pH and redox potential, and that these effects are greater when the water is exposed to a magnetic field. The reasons for this behavior are not clear, but one possibility is that dissolved O2 molecules, which are paramagnetic, might be involved.
Our ordinary drinking water, by contrast, is never chemically pure, especially if it has been in contact with sediments. Groundwater (from springs or wells) always contain ions of calcium and magnesium, and often iron and manganese as well; the positive charges of these ions are balanced by the negative ions carbonate/bicarbonate, and occasionally some chloride and sulfate. Groundwaters in some regions contain unacceptably high concentrations of naturally occurring toxic elements such as selenium and arsenic.
You might think that rain or snow would be exempt from contamination, but when water vapor condenses out of the atmosphere it always does so on a particle of dust which releases substances into the water, and even the purest air contains carbon dioxide which dissolves to form carbonic acid. Except in highly polluted atmospheres, the impurities picked up by snow and rain are too minute to be of concern.
Various governments have established upper limits on the amounts of contaminants allowable in drinking water; the best known of these are the U.S. EPA Drinking Water Standards.