water

hydrogen oxide (H₂O), the simplest chemical com-pound of hydrogen and oxygen (11.19 percent hydrogen and 88.81 percent oxygen by weight) that is stable under standard conditions. Molecular weight, 18.0160. Colorless, odorless, and tasteless liquid (deep water has a bluish color).

Water played a crucial role in the geological history of the earth, in the origin of life, and in the formation of the physical and chemical environment, the climate, and the weather of our planet. Living organisms could not exist without water. It is an essential component of almost all technological processes, both in agriculture and in industry.

Water in nature. Water is widespread in nature. The hydrosphere, which is the water envelope of the earth and which includes the oceans, seas, lakes, reservoirs, rivers, subterranean water, and moisture in the soil, contains approximately 1.4-1.5 billion cu km, of which approximately 90 million cu km is land water. Subterranean water accounts for 60 mil-lion cu km, glaciers for 29 million, lakes for 0.75 million, soil moisture for 75,000, and rivers for 1,200. Water exists in the atmosphere in the form of vapor, fog, clouds, drops of rain, and snow crystals, for a total of approximately 13,000-15,000 cu km. Glaciers permanently occupy approximately 10 percent of the land surface. In the northern and north-eastern USSR and in Alaska and northern Canada, there is always an underground layer of ice over an average area of about 16 million sq km (a total of about 0.5 million cu km). According to various estimates, the earth’s crust—the lithosphere—contains 1 to 1.3 billion cu km of water, which is close to the water content of the hydrosphere. In the earth’s crust a considerable amount of water is bound as a component of certain minerals and mineral rocks (gypsum, hydrated forms of silica, hydrosilicates, and so on). Large quantities of water (13-15 billion cu km) are concentrated in the deeper regions of the earth’s mantle. The water that was given off by the mantle during the process of the earth’s warming up at early stages of its formation was responsible, according to contemporary views, for the formation of the hydrosphere. The annual income of water from the mantle and magma beds is about 1 cu km. There are data to indicate that water is, at least in part, of“cosmic” origin: protons coming into the upper atmosphere from the sun and attracting electrons are transformed into hydrogen atoms which, by uniting with oxygen atoms, give H₂O. Water is a component of all living organisms, which together contain half as much water as all of the earth’s rivers. The amount of water in living organisms, excluding seeds and spores, varies between 60 and 99.7 percent by weight. In the words of the French biologist E. Du Bois-Reymond, a living organism is I’eau animee (“animated water”). All of the earth’s water is constantly intermingling and circulating in the atmosphere, the lithosphere, and the biosphere.

Under natural conditions, water always contains dissolved salts, gases, and organic substances. Their quantitative com-position varies with the source of the water and with environmental conditions. Water with a salt concentration of less than 1 g/kg is considered fresh; up to 25 g/kg is considered slightly salty; and above 25 g/kg is considered salt water.

The water with the lowest mineral content comes from atmospheric precipitation (on the average, about 10-20 mg/kg); the next lowest (50-1,000 mg/kg) is found in fresh-water lakes and rivers. The salt content of the ocean varies around 35 g/kg; seas have a lower mineral content (the Black Sea, 17-22 g/kg; the Baltic Sea, 8-16 g/kg; and the Caspian Sea, 11-13 g/kg). The mineral content of subterranean water near the surface under conditions of excess moisture may be as high as 1 g/kg; under arid conditions it reaches 100 g/kg; and in deep waters the mineralization varies within a broad range. The maximum concentration of salts is found in salt lakes (as high as 300 g/kg) and deep-lying subterranean water (about 600 g/kg).

Ions of HCO₃^-, Ca²⁺, and Mg²⁺ usually predominate in fresh water. As the total mineral content rises, the concentration of SO₄^2-, Cl^-, Na⁺, and K⁺ increases. In water with a high mineral content, Cl^- and Na⁺ ions predominate; less frequently Mg²⁺, and very rarely Ca²⁺. Other elements are present in very small quantities, but almost all natural elements of the periodic system are found in native water.

The dissolved gases in native water include nitrogen, oxygen, carbon dioxide, the noble gases, and rarely hydrogen sulfide and hydrocarbons. The concentration of organic substances is small: on the average in rivers it is about 20 milli-grams per liter (mg/l); in subterranean water it is even less, and in the ocean it is about 4 mg/l. Water in marshes and petroleum deposits and water polluted by industrial and domestic sewage, which have a higher concentration of organic substances, are an exception. The qualitative composition of the organic substances is extremely varied and includes various products of the vital activity of the organisms inhabiting the water and the compounds formed upon the breakdown of their remains.

The salts in native water originated in substances that were formed during the chemical weathering of igneous rock (Ca²⁺, Mg²⁺, Na⁺, K⁺, and so on) and in substances discharged from the earth’s interior throughout its history (CO₂, SO₂, HCl, NH₃, and others). The composition of water de-pends on the diverse composition of these substances and conditions under which they reacted with water. The effects of living organisms are also of considerable significance on the composition of water.

Isotopic composition. Because of the existence of two stable isotopes of hydrogen (JH and 2H, usually designated H and D) and three of oxygen ¹⁶O, ¹⁷O, and ¹⁸O, nine isotopic forms of water are known. They are found in nature in the following average proportions (in molecular percent): H₂¹⁶O, 99.73; H₂¹⁷O, 0.04; H₂¹⁸O, 0.20; and HD¹⁶O, 0.03; as well as 10^-5 to 10^-15 percent (total) HD¹⁷O, HD¹⁸O, D₂¹⁶O, D₂¹⁷O, and D₂¹⁸O. Heavy water, D₂O, which contains deuterium, is of particular interest. In all of the earth’s water there is only 13-20 kg of“superheavy” water, containing a radioactive isotope of hydrogen—tritium (³H, or T).

Historical information. Because of its wide distribution and its role in human life, water has long been considered the source of life. The ancient philosophers’ notion that water was the origin of all things was reflected in Aristotle’s theory (fourth century B.C.) of the four elements (fire, air, earth, and water), according to which water was thought to be the carrier of cold and moisture. The concept of water as a single chemical element endured in science until as late as the end of the 18th century. In 1781-82 the English scientist H. Cavendish synthesized water for the first time by exploding a mixture of hydrogen and oxygen with an electric spark, and in 1783 the French scientist A. Lavpisier repeated these experiments and for the first time drew the correct conclusion that water is a compound of hydrogen and oxygen. In 1785, Lavoisier, together with the French scientist J. Meusnier, determined the quantitative composition of water. In 1800 the English scientists W. Nicholson and A. Carlisle separated water into its elements by means of an electrical current. Thus, the analysis and synthesis of water revealed its complex composition and permitted the determi-nation of its formula, H₂O. The study of the physical proper-ties of water had already begun before the determination of its composition in close conjunction with other scientific and technical problems. In 1612 the Italian scientist Galileo de-voted attention to the lower density of ice in comparison with liquid water as the reason for the buoyancy of ice. In 1665 the Dutch scientist C. Huygens proposed the adoption of the boiling and melting temperatures of water as the reference points for a thermometer scale. In 1772 the French physicist Deluc found that the maximum density of water occurs at 4° C; with the establishment of the metric system of weights and measures at the end of the 18th century, this observation was used in order to define the unit of mass and weight, the kilogram. In conjunction with the invention of the steam engine, the French scientists D. Arago and P. Dulong (1830) studied the pressure dependence of saturated water vapor on temperature. During the period from 1891 to 1897, D. I. Mendeleev derived the formulas for the dependence of water density on temperature. In 1910 the American scientist P. Bridgman and the German scientist G. Tammann discovered certain polymorphic modifications in ice at high pressures. In 1932 the American scientists E. Washburn and H. Urey discovered heavy water. The progress of physical methods of research made possible substantial advances in the study of the structure of water molecules and of the structure of ice crystals. In the last decade, scientists have de-voted particular attention to the structure of liquid water and of aqueous solutions.

Physical properties and structure. The most important physical constants for water are given in Table 1. (See the article WATER VAPOR regarding the pressure of saturated water vapor at various temperatures. See the article ICE regarding the polymorphic-modifications of water in the solid state.) The triple point of water, at which liquid water, ice, and water vapor are in equilibrium, occurs at a temperature of 0.01° C and a pressure of 6.03 x 10^-3 atmospheres.

Many physical properties of water exhibit substantial irregularities. As is known, the properties of one type of com-pound with elements from the same group in Mendeleev’s periodic system vary in a regular fashion. In the row of hydrogen compounds with elements from Group VI (H₂Te, H₂Se, H₂S, and H₂O), the melting points and boiling points

become lower only for the first three; for water the melting point and boiling point are anomalously high. The density of water increases normally in the interval from 100° to 4° C, as with the vast majority of other liquids. However, after attaining a maximum value of 1.0000 g/cm³ at 3.98° C, the density decreases upon further cooling and upon freezing falls abruptly, whereas in most other substances crystallization is accompanied by an increase in density. Water is capable of considerable supercooling—that is, it can remain in the liquid state below the melting point (even at -30° C). The specific heat, heat of fusion, and heat of vaporization of water are abnormally high in comparison with other substances, and the specific heat is at a minimum at 40° C. The viscosity of water decreases rather than increases with an increase in pressure, as would be expected by analogy with other liquids. The compressibility of water is extremely small and decreases with an increase in temperature.

The anomalies in the physical properties of water are due to the structure of its molecules and to the peculiarities of the intermolecular interactions in liquid water and in ice. The three nuclei of a water molecule form an isosceles triangle, with the protons at the base and the oxygen at the top (Figure l,a). The distribution of electron density in the water molecule is such (Figure 1 ,b and c) that four charge poles are created: two positive, associated with the hydrogen atoms; and two negative, associated with the electron clouds of the unshared pairs of electrons on the oxygen atom. The four charge poles are located at the corners of a tetrahedron (Figure l,d). Because of this polarity, water has a large dipole moment (1.86 D), and the four charge poles permit each water molecule to form four hydrogen bonds with its neigh-boring (identical) molecules—for example, in ice crystals.

Figure 1. Structure of a water molecule: (a) geometry of the H₂O molecule (in the gaseous state), (b) electron orbits in the H₂O molecule, (c) electron configuration of the H₂O molecule (the unshared electron pairs are visible), (d) four charge poles situated at the corners of a tetrahedron in the H₂O molecule.

The crystal structure of ordinary ice is hexagonal (see Figure 2). It is“loose” and contains many“holes.” (If the water molecules were densely“packed” in ice crystals the density would be about 1.6 g/cm³.) In liquid water the bonds that are inherent in ice between H₂O molecules and their four neighbors (“short-range order”) are preserved to a substantial degree; however, the“looseness” of structure decreases upon melting, and molecules of“long-range order” fall into the“holes,” which leads to an increase in density. Upon further heating, the thermal motion of the molecules increases and the distance between them increases—that is, the water expands. This expansion is already predominant at 3.98° C, and thus the density of water decreases with an increase in temperature. Hydrogen bonds are approximately ten times stronger than the bonds caused by the intermolecular interactions characteristic of the majority of other liquids; thus, much more energy is required for the melting, evaporation, and heating of water than in the case of other liquids, which explains the abnormally high values for the heats of fusion and vaporization and for the specific heat of water. The hydrogen bonds break upon an increase in temperature, but a certain number of them are preserved even at 100° C. Water dissolved in organic solvents is composed of (H₂O)₂ aggregates, which form because of the hydrogen bonds.

Figure 2. Crystal structure of ice

Water as a solvent. Water is the universal solvent. Gases dissolve fairly readily in water if they are capable of entering into chemical interactions with it (ammonia, hydrogen sulfide, sulfur dioxide, and carbon dioxide). Other gases are not readily soluble in water. With a decrease in pressure and an increase in temperature, the solubility of gases in water decreases. At low temperatures and high pressures, many gases (argon, krypton, xenon, chlorine, hydrogen sulfide, hydrocarbons, and others) not only dissolve in water but also form crystal hydrates. In particular, propane at 10° C and 0.3 meganewtons per sq m (MN/m2), or 3 kilograms-force per sq cm (kgf/cm2), yields the crystal hydrate C3H8-17H2O. Such hydrates decompose with a decrease in pressure. The crystal hydrates formed at low temperatures from many gaseous substances contain water in the“holes” of their crystals (so-called clathrate compounds or inclusion complexes).

Water is a weak electrolyte, dissociating according to the equation H₂O H⁺ + OH^-, in which the ion production serves as a quantitative indicator of the electrolytic dissociation: K_w = [H⁺] [OH^-], where [H^-] and [OH^-] are the concentrations of the respective ions in gram ion per liter; Kw is 10^-14 at 22° C and 72 x l0^-14 at 100° C, which corresponds to an increase in dissociation with an increase in temperature.

Since it is an electrolyte, water dissolves many acids, bases, and mineral salts. Such solutions conduct electric cur-rent because of the dissociation of dissolved substances with the formation of hydrated ions (hydration). Many substances enter into an exchange reaction with water when they are dissolved in it; this is called hydrolysis. Those organic substances that contain polar groups (—OH, —NH₂, —COOH, and others) and whose molecular weight is not too high dis-solve in water. Water itself is readily soluble (or mixes well in all proportions) only in a limited number of organic solvents. However, water is almost always present in organic substances as an insignificant admixture and is capable of radically altering the physical constants of the substances.

The water in any natural reservoir contains various substances, principally salts, in solution. Because of water’s great solvent power, it is extremely difficult to obtain it in the pure state. The electrical conductivity of water usually serves as a measure of its purity. Distilled water obtained from ordinary water—and even distillate that has been dis-tilled a second time—has an electrical conductivity 100 times greater than absolutely pure water. The purest water is produced by synthesis in special apparatus, using carefully purified oxygen and hydrogen.

In recent years much information has been collected about the substantial change in the properties of industrial and dis-tilled water that occurs when it is passed through magnetic fields of optimal (very low) strength at a specific velocity. These changes are temporary and disappear gradually and spontaneously after 10-25 hours. It has been noted that, after such“magnetic treatment,” absorption and the processes of crystallization of substances dissolved in the water are speeded up and that the moistening capacity of the water and other properties also change. Although the theoretical explanations of these phenomena are still lacking, the principles have already been widely applied to prevent the formation of scum in steam boilers and to improve the processes of flotation, elimination of suspended matter from water, and so on.

Formation and dissociation. The formation of water during the interaction between hydrogen and oxygen is accompanied by the release of 286 kilojoules per mole (kJ/mole), or 58.3 kilocalories per mole (kcal/mole), of heat at 25° C (for liquid water). The reaction 2H₂ O₂ = 2H₂O proceeds very slowly below a temperature of 300° C; above 550° C it is explosive. The presence of a catalyst (for example, platinum) permits the reaction to take place at ordinary temperatures. Both the slow combustion of hydrogen in oxygen and their explosive reaction are chain reactions, which take place with the participation of free radicals.

Chemical properties. Under ordinary conditions water is a relatively stable compound. The breakdown of H₂O molecules (thermal dissociation) becomes noticeable only above 1500° C. The breakdown of water takes place both under the action of ultraviolet and radioactive radiation (photodissociation and radiolysis respectively). In the latter case, hydrogen peroxide and many free radicals are also formed in addition to H₂ and O₂. A characteristic chemical property of water is its ability to enter into addition reactions and the hydrolytic dissociation of the reacting substances. Reducers act on water mainly at high temperatures. Only the most active of these, such as the alkaline and alkaline-earth metals, react with water even at room temperature with the release of hydrogen and the formation of hydroxides: 2Na 2H₂O = 2NaOH H₂; Ca 2H₂O = Ca(OH)₂ H₂. Magnesium and zinc react with boiling water; aluminum reacts with water when the oxide film has been removed from its surface. Less active metals react with water at red heat: 3Fe 4H₂O = Fe₃O₄ 4H₂. The slow reaction of many metals and their alloys with water takes place at ordinary temperatures. By using water containing the oxygen isotope ¹⁸O it has been demonstrated that in the corrosion of iron in a moist atmosphere the“rust” receives oxygen specifically from water and not from the air. The noble metals (gold, silver, platinum, palladium, ruthenium, and rhodium), as well as mercury, do not react with water.

Atomic oxygen converts water into hydrogen peroxide: H₂O + O = H₂O₂. Fluorine also splits water at ordinary temperatures: F₂ + H₂O = 2HF O. Simultaneously, H₂O₂, ozone, fluorine oxide (F₂O), and molecular oxygen (O₂) are formed. At room temperature, chlorine and water give hydrochloric and hypochlorous acids: Cl₂ + H₂O = HCl+ HClO. Bromine and iodine react with water in a similar fashion under these conditions. At high temperatures (100° C for chlorine, 550° C for bromine) the reaction proceeds with the liberation of oxygen: 2Cl₂ 2H₂O = 4HC1 O₂. Phosphorus reduces water and forms metaphosphoric acid (only in the presence of a catalyst under pressure and at high temperature): 2P 6H₂O = 2HPO₃ 5H₂. Water does not react with nitrogen and hydrogen, but with carbon at high temperatures it gives water vapor: C + H₂O = CO + H₂. The reaction can be used both in the industrial production of hydrogen and in the conversion of methane: CH₄ H₂O = CO + 3H₂ (1200°-1400° C). Water reacts with many basic and acidic oxides to form the corresponding bases and acids. The addition of water to unsaturated hydrocarbons provides the basis for the industrial method of producing alcohols, aldehydes, and ketones. Water participates in many chemical processes as a catalyst. For example, the reaction of alkaline metals or hydrogen with halogens, and many oxidation reactions, will not proceed without the presence of small quantities of water.

Water that is chemically bonded with a substance of which it is a part and that is undetectable in the form of H₂O is called water of constitution; the H₂O molecules form only at the moment of decomposition of the substance—for example, as a result of strong heating: Ca(OH)₂ = CaO + H₂O. Water that is part of a number of crystalline substances—for example, aluminum alums, K₂SO₄ • Al₂(SO₄)₃ • 24H₂O—and is detectable in these crystals by means of X-ray crystallography is called water of crystallization or crystal hydrate water. Water that is absorbed by solid substances that have high porosity and large surface areas (for example, activated carbon) is called water of adsorption. Free water that occu-pies small channels (for example, in the soil) is called hygroscopic (capillary) water. Structurally free water, which fills in the holes in certain structures, such as in minerals is also distinguished. It is possible to detect water qualitatively in the form of a condensate formed upon heating the specimen under examination; by heating while weighing the specimen, quantitative results are obtained (thermogravimetric analysis). Water in organic solvents can be detected by dyeing with colorless copper sulfate, which, when added to water, forms the blue crystal hydrate CuSO₄-5H₂O. It is often possible to separate and analyze water quantitatively by azeotropic distillation of the water with benzene, toluene, or other liquid in the form of an azeotropic mixture; after the separation of the mixture upon cooling, the volume of sepa-rated water is measured.

Use in industry. It is impossible to think of any other substance that has as varied and wide use as water. It is the chemical reagent that participates in the production of oxygen, hydrogen, alkalies, nitric acid, alcohols, aldehydes, hydrated lime, and many other very important products. It is a necessary component in the setting and hardening of binding materials such as cement, gypsum, and lime. Water is used in many industrial processes as a technological component: in cooking, solution, dilution, lixiviation, and crystallization. In engineering, water serves as an energy carrier and heat carrier (steam heating and water cooling) and as the working medium in steam engines, and it is used in the transmission of pressure (in particular, in hydraulic transmissions and in presses, as well as in petroleum extraction) or power (hydraulic machine drives). Sprayed through a nozzle under great pressure, water washes away soil or rock.

The demands made of water in industry are extremely varied. Water of special purity is needed for the development of the newer branches of industry (the production of semiconductors and phosphors, atomic engineering, and so on). Therefore, special attention is currently being devoted to problems of water treatment and purification. According to some estimates, the total volume of material (ore, coal, oil, minerals, and so on) processed yearly in the entire world is about 4 billion cu m (4 cu km); for the same period the consumption of fresh water—that is, water from water-supply sources—in the USSR alone was 37 billion cu m in 1965. The rapid increase in the use of water poses a new and important problem for mankind—the struggle against the depletion and pollution of the water resources of the planet.