Water’s Properties

What exactly did we mean when we said that water is a common but unusual substance? Part of the answer was provided by the Greek philosopher Thales as long ago as the 6th century B.C., when he stated correctly that, of all substances, water was unusual in existing in all three states of matter—in the range of temperature and presure found on earth. For instance, the upper part of a lake may be solid ice, while beneath, it may exist as liquid; it also evaporates from the surface of the lake as gas, or invisible water vapour.

More than 2000 years later, when scientists began to investigate the physics and chemistry of matter, they examined water in relation to other compounds, and found that it was even more of a puzzle than Thales thought. But they could not account for water’s strange behaviour until more was known about the atom and the nature of the chemical bond.

Any atom consists of a relatively heavy, positively charged nucleus, surrounded by one or more light, negatively charged electrons. The negative and positive charges balance each other, so that the atom as a whole is electrically neutral. The electrons spin around the nucleus in orbits, and it is only the outer electrons orbit that are involved in chemical reactions; the nucleus plays no part at all.

The circles represent the outer orbit of six electrons. Hydrogen, which is the other element in water, has only one electron.

Atoms usually combine with other atoms to form molecules. The atoms may occur in pairs, as in oxygen and hydrogen may combine with the poisonous gas chlorine to form edible table salt .

Why is it that atoms combine to form molecules? The answer is that by doing so they become more stable—they form molecules that are less likely to turn into something else. Atoms are most stable when they contain eight electrons in their outermost orbit. Hydrogen is exceptional in that it needs only two electrons. Let us see how hydrogen attains stability. In the hydrogen molecule, each atom shares its single electron with the other, so that each is surrounded by two electrons.

In the case of the water molecule , two hydrogen atoms combine with one oxygen atom, each hydrogen atom sharing one electron with the oxygen.

The combination is stable because two electrons spin around the hydrogen and eight around the oxygen. This type of bond, which is called covalent, is very strong, and occurs in most organic, and many inorganic, compounds.

Another common type of link between atoms is the electrovalent bond, in which atoms lose or gain electrons, instead of sharing them as in the covalent bond. The chlorine atom, for example, has seven outer electrons, while the sodium atom has one outer electron, beneath which lies an orbit of eight electrons; both atoms, remember, are electrically neutral. But, when sodium and chlorine come together, the surplus sodium electron passes to the chlorine atom. The result is a molecule consisting of a positive sodium ion and a negative chlorine ion, which attract each other like the opposite poles of a magnet. The sodium chloride molecule is stable because both ions have eight electrons in their outer orbit. We shall see later how water dissolves electrovalent compounds by making the ions fall apart.

A solid has a definite volume, shape, and rigidity. A liquid has a definite volume, but shape is the same as its container. A gas always fills its container. It is not yet certain that water’s decrease in density, as temperature rises, is due to the increased separation of its molecules.

So much for the types of bond between atoms; now we must look at the forces that hold molecules together. Most molecules exert weak attractive forces on each other, called van der Wads’ forces. In solids, these hold the atoms together in a very compact pattern; the solid has a definite shape and volume, and also a certain rigidity. The molecules also vibrate slowly; this means that the solid contains heat energy , which can be measured as a definite temperature. If a solid is heated above a certain point, the molecules vibrate so fast that they become separated, overcoming to a certain extent the weak attractive forces between them. When this happens, the solid melts to form a liquid; this has definite volume, but almost no rigidity, the shape of the liquid assuming that of the containing vessel. Carrying the process further, if we heat the liquid sufficiently, the molecules vibrate so quickly that some of them escape from the surface, as a gas—a process called evaporation. Finally, the liquid molecules vibrate so fast that the van der Waals’ forces are no longer able to keep the molecules together; the molecules fly apart with great speed so that the whole liquid evaporates, or vaporizes, to form a gas. Gases have neither shape nor fixed volume, and always fill their containers.

The temperature at which a solid starts to turn into a liquid—the melting point—depends partly on how strong the forces of attraction. between the molecules are. This, in turn, depends on the molecular weight, and on how the atoms are arranged within the molecule. The same applies to the boiling point, which is the temperature at which the whole liquid turns into a gas. In general, we can say that the higher the molecular weight, the higher the melting and boiling points, and that different compounds with similar molecular structures behave in much the same way. The melting and boiling points of water should therefore be predictable by comparing the water molecule with molecules that appear to have the same basic structure, like H..S. But when we do this we find that water + + should theoretically freeze at about — 100°c and boil at — 80°c. Instead, water freezes at 0°c and boils at 100°c, as if it were a much heavier molecule, with an entirely different structure. There must obviously be something other than van der Waals’ forces holding water molecules together and delaying them from parting as the temperature rises.

Let us take another look at the structure of the water molecule. Although the oxygen and hydrogen atoms share their electrons, oxygen takes an unequal share; the electrons are closer to oxygen than to hydrogen. The oxygen end of the molecule is thus negative and the hydrogen end positive, so that molecules join up with the negative oxygen ends attracting the positive hydrogen ends. Water, in fact, exists as groups of linked molecules. These groups also form larger groups in a way that we do not yet understand.

These electrostatic forces linking water molecules are called hydrogen bonds. They are strong—the heat energy required to overcome the strength of the hydrogen bonds in one litre of water exceeds 300,000 calories. Because the molecules are held together strongly they have to absorb much heat before vibrating into the widely spaced arrangement found in the liquid or gas; the melting and boiling points are thus correspondingly high. However, recent research suggests that the behaviour of water molecules may be influenced by other forces as well, and there is some doubt as to the effect of temperature on the separation of water molecules. The exact values of water’s melting and boiling points gain significance only when we relate them to the temperatures found on earth. At sea-level, especially in the tropics, the air temperature is relatively not so very far below that of the boiling point of water. This means that considerable quantities of water evaporate from the sea, to be blown by winds across the land. High in the atmosphere, where the temperature is much lower than the freezing point of water, the water vapour turns to ice, and then falls as snow or hail, which later may turn to rain in the lower regions of the atmosphere.

The fact that hydrogen bonds are fairly strong also partly explains why water absorbs a great quantity of heat to show a small rise in temperature. It takes far more heat to raise the temperature of 10 lb. of water than it would to raise the temperature of 10 lb. of any other commonly occurring substance by the same amount. We say, therefore, that water has a high heat capacity, or specific heat. If the specific heat were not so high, the summer temperatures of lakes and oceans would be much higher than they are, and would make them unable to support aquatic life. The high specific heat also has a marked effect in regulating climate by preventing extremes of hot and cold. During the day, some of the heat from the sun is absorbed by water vapour near the ground, thus preventing the air from becoming too hot; during the cooler night, heat is given up by the vapour, which stops the air temperature from falling too much. Another effect of water’s high specific heat is that water cools and heats up about five times more slowly than land, so that coastal areas do not have such an extreme range of temperatures as inland areas in the same latitude. The difference between the specific heats of land and water also accounts for many daily and seasonal wind patterns. South-east Asia, for instance, is colder than the sea in winter, so that winds blow toward the sea. In summer, the land is relatively warmer, and the winds move onshore, bringing the long-awaited rain.

When ice melts or water boils, they absorb an appreciable amount of heat without showing any rise in temperature; this heat is used solely to change water from one state to another. The amount of heat absorbed on melting amounts to 80 calories per gram, and is called the latent heat of fusion. When water turns from liquid to gas, it absorbs 540 calories per gram, this being the latent heat of evaporation. Compared with other substances, water’s latent heats are extremely high, but this is what we should expect in view of what we know about hydrogen bonds. The change from solid to liquid involves a partial separation of the molecules against the attraction of the strong hydrogen bonds, and this requires heat energy. The change from liquid to gas, however, involves a complete breaking of the hydrogen bonds, and nearly seven times more heat is needed for this change of state. The latent heat absorbed while melting and boiling does not simply disappear; in the reverse process, when water vapour condenses into liquid, every gram gives out the same 540 calories, and when water freezes, every gram emits 80 calories.

As water is always changing from one state to another, heat is continuously being produced and absorbed at the earth’s surface. This has a great moderating effect on the earth’s climate. When the air temperature falls, water vapour condenses into liquid droplets, heat is given out, and the air temperature rises. When water evaporates on a hot day, it absorbs heat, and the surrounding air is cooled. Man makes use of water’s high latent heat of fusion, for instance, when he uses an ice-box cooler; as the ice melts it absorbs heat from both the air and the food inside the ice-box.

Almost all substances become heavier, or more dense, as they approach freezing point. The reason for this is that the molecules come closer together as the temperature falls. But water behaves differently: it is most dense at 4°c, instead of at its freezing point of 0°c. Below 4°c, the density of water decreases because the hydrogen bonds strongly influence the arrangement of the molecules so that they develop large-spaces between them. At 0°c, the molecules suddenly arrange themselves into a definite crystalline pattern, in which each oxygen atom is surrounded by four hydrogen atoms.

Water thus expands on freezing, and it does so with great force, as some of us know to our cost when we find steel pipes burst by frost. At temperatures far below 0°c, water may freeze in our body’s cells and tear them apart. But against these disadvantages we must weigh the advantages. Freezing water splits rock into fragments that fall on to the soil, contributing to its mineral content. In winter, farmers welcome the frost and ice, which break up large clods of earth into smaller soil crumbs. The fact that water expands on freezing also means that ice is lighter than liquid water, and therefore floats. Imagine what would happen if this were not so. In winter, the Great Lakes of America and Canada, and the upper reaches of the Mississippi and Hudson, for example, would freeze; ice would sink to the bottom and would not be melted by the sun next summer. Each year, more ice would sink, until finally the amount of liquid water would be appreciably reduced. These waterways would no longer be navigable, and the amount of water available for domestic, agricultural, and industrial use, and for the disposal of sewage, would no longer be sufficient. There would also be no fishing.

Our next property of water is its high power of cohesion—the power of sticking to itself, also largely due to the strong attractive forces of the hydrogen bonds. It means, for example, that a long column of water does not easily break; a column of fairly pure water needs a force of roughly 2000 pounds per square inch to rupture it. This is very important, because all plants contain columns of rising water in which are dissolved foods for growth. If these columns broke, all plant life would cease, and all animal and human life with it.

Because water has a high cohesive power, it also has the highest surface tension of any common liquid. Surface tension gives water a ‘skin,’ as we may have noticed when watching insects walk across the surface of a pond. The reason is that molecules below the surface exert forces of attraction on each other in all directions, and these cancel out. At the surface, however, there is no upward pull to balance the downward pull of the molecules beneath, so the surface is pulled downward.

Let us consider the effect surface tension has on a drop of water. Strong equal forces pull inward all over its surface, and so the drop tends to become a sphere, this being the shape with the smallest surface area. The force of surface tension is so strong that a raindrop resists spreading out at high velocities; one can therefore think of raindrops as being fairly hard. Because of this, rain has a marked effect on many bare soils that lack a shock-absorbing mat of vegetation. A heavy downpour gouges holes in these soils and also closes the minute openings of the soil channels. The drops then lie on the surface without being absorbed. In time, the holes may merge to form gulleys along which flows water, at the same time carrying -away valuable topsoil.

A further special property of water is that it adheres to, or wets, many substances. At the side of a glass vessel, for example, water climbs up a short distance because the force of attraction between the water and glass is stronger than the cohesive force below the water surface. In a tube with a very narrow bore, this adhesion makes water rise a considerable distance—a property called capillarity. This is very important in nature, for water is thus able to travel through the network of narrow channels in the soil for a short way, to feed plant roots.

One reason why water adheres is that it forms hydrogen bonds with many substances. Glass, for example, consists largely of silica , whose oxygen atoms join up with the hydrogen ends of the water molecules. In the same way, water strongly wets those soils that have a large component of clay, because this also contains oxygen atoms combined with silicon.

Finally, we examine water as a remarkable solvent, for it dissolves more substances than any other liquid. To explain what happens when it dissolves electrovalent compounds , we can return to the molecule of sodium chloride. In the solid state, salt consists of oppositely charged sodium and chlorine ions held together by electrostatic forces, with small spaces between them. Since water can penetrate between these spaces at the surface of the salt, the water molecules arrange themselves so that their negative ends point toward the positive sodium and their positive ends toward the negative chlorine. This neutralizes the electrostatic forces between the salt ions to a strength 80 times less than when the salt was solid. There is then little force left to keep the ions together, and they fall apart; in other words, the salt dissolves. In the same way, water dissolves over 40 different salts in the sea, and a great many in soil; as a cheap industrial solvent, it is unequalled.

It is due to the fact that water has these special properties that the human body is composed of over 60 per cent water. In another post we shall examine in more detail the varied roles that water plays in the body. The same principles apply also to plants, on which we all rely directly or indirectly for food.

We are now nearer to understanding why water is basic to life. It is a part of our environment as important as the sun itself and the air we breathe. There is no substitute; no other substance has such unusual properties that make it float when frozen, break up rock and soil, moderate our climate—all properties that make the earth habitable. From the time life began, plants and animals evolved in the presence of water; they are now entirely dependent on it, and man is equally vulnerable.