Chemical and Physical Properties
Chemically, water is a compound of hydrogen and oxygen, having the formula H2O. It is chemically active, reacting with certain metals and metal oxides to form bases, and with certain oxides of nonmetals to form acids. It reacts with certain organic compounds to form a variety of products, e.g., alcohols from alkenes. Because water is a polar compound, it is a good solvent. Although completely pure water is a poor conductor of electricity, it is a much better conductor than most other pure liquids because of its self-ionization, i.e., the ability of two water molecules to react to form a hydroxide ion ion, atom or group of atoms having a net electric charge.
Positive and Negative Electric Charges
A neutral atom or group of atoms becomes an ion by gaining or losing one or more electrons or protons.
..... Click the link for more information. , OH−, and a hydronium ion, H3O+. Its polarity and ionization are both due to the high dielectric constant of water.
Water has interesting thermal properties. When heated from 0°C;, its melting point, to 4°C;, it contracts and becomes more dense; most other substances expand and become less dense when heated. Conversely, when water is cooled in this temperature range, it expands. It expands greatly as it freezes; as a consequence, ice is less dense than water and floats on it. Because of hydrogen bonding between water molecules, the latent heats latent heat, heat change associated with a change of state or phase (see states of matter). Latent heat, also called heat of transformation, is the heat given up or absorbed by a unit mass of a substance as it changes from a solid to a liquid, from a liquid to a gas,
..... Click the link for more information. of fusion and of evaporation and the heat capacity heat capacity or thermal capacity, ratio of the change in heat energy of a unit mass of a substance to the change in temperature of the substance; like its melting point or boiling point, the heat capacity is a characteristic of a substance.
..... Click the link for more information. of water are all unusually high. For these reasons, water serves both as a heat-transfer medium (e.g., ice for cooling and steam for heating) and as a temperature regulator (the water in lakes and oceans helps regulate the climate).
Structure of the Water Molecule
Many of the physical and chemical properties of water are due to its structure. The atoms in the water molecule are arranged with the two H-O bonds at an angle of about 105° rather than on directly opposite sides of the oxygen atom. The asymmetrical shape of the molecule arises from a tendency of the four electron pairs in the valence shell of oxygen to arrange themselves symmetrically at the vertices of a tetrahedron around the oxygen nucleus. The two pairs associated with covalent bonds (see chemical bond chemical bond, mechanism whereby atoms combine to form molecules. There is a chemical bond between two atoms or groups of atoms when the forces acting between them are strong enough to lead to the formation of an aggregate with sufficient stability to be regarded as
..... Click the link for more information. ) holding the hydrogen atoms are drawn together slightly, resulting in the angle of 105° between these bonds. This arrangement results in a polar molecule, since there is a net negative charge toward the oxygen end (the apex) of the V-shaped molecule and a net positive charge at the hydrogen end. The electric dipole gives rise to attractions between neighboring opposite ends of water molecules, with each oxygen being able to attract two nearby hydrogen atoms of two other water molecules. Such hydrogen bonding, as it is called, has also been observed in other hydrogen compounds. Although considerably weaker than the covalent bonds holding the water molecule together, hydrogen bonding is strong enough to keep water liquid at ordinary temperatures; its low molecular weight would normally tend to make it a gas at such temperatures.
Various other properties of water, such as its high specific heat, are due to these hydrogen bonds. As the temperature of water is lowered, clusters of molecules form through hydrogen bonding, with each molecule being linked to others by up to four hydrogen bonds, each oxygen atom tending to surround itself with four hydrogen atoms in a tetrahedral arrangement. Hexagonal rings of oxygen atoms are formed in this way, with alternate atoms in either a higher or lower plane than their neighbors to create a kinked three-dimensional structure.
Liquid Water
According to present theories, water in the liquid form contains three different molecule populations. At the highest temperatures single molecules are the rule, with little hydrogen bonding because of the high thermal energy of the molecules. In the middle range of temperatures there is more hydrogen bonding, and clusters of molecules are formed. At lower temperatures aggregates of clusters also form, these aggregates being the most common arrangement below about 15°C;. On the basis of these three population types and the transitions between them, many aspects of the anomalous behavior of water can be explained. For example, the tendency of water to freeze faster if it has been cooled rapidly from a relatively warm temperature than if it has been cooled at the same rate from a lower temperature is explained in terms of the greater number of irregularly shaped cluster aggregates in the cooler water that must find a suitable means of fitting together with a neighboring aggregate.
The discovery in the late 1960s of "superwater," or "polywater," helped to shed light on some aspects of the structure of water. This substance was thought by some to be a giant polymer of water molecules, 40 times denser and 15 times more viscous than ordinary water. Studies showed, however, that these new and unexplained properties were connected with the presence of contaminants in the water. Even so, the interaction of the water molecules with these other substances may be helpful in understanding the way in which water molecules interact with each other.
Ice
In ice, each molecule forms the maximum number of hydrogen bonds, resulting in crystals composed of open, hexagonal columns. Because these crystals have a number of open regions and pockets, normal ice is less dense than water. However, other forms of ice also exist at conditions of higher pressure, each of these different forms (designated ice II, ice III, etc.) having greater density and other distinct physical properties that differ from those of normal ice, or ice I. As many as eight different forms of ice have been distinguished in this manner. The higher pressures creating such forms cause rearrangements of the hexagonal columns in ice, although the basic kinked hexagonal ring is common to all forms.
When ice melts, it is thought that the fragments of these structures fill many of the gaps that existed in the crystal lattice, making water denser than ice. This tendency is the dominant one between 0°C; and 4°C;, at which temperature water reaches its maximum density. Above this temperature, expansion due to the increased thermal energy of the molecules is the dominant factor, with a consequent decrease in density.
Bibliography
See D. Eisenberg and W. Kauzmann, The Structure and Properties of Water (1969); A. K. Biswas, History of Hydrology (1970); C. Hunt and R. M. Garrels, Water: The Web of Life (1972); P. Ball, Life's Matrix: A Biography of Water (2000).
water
Inorganic compound composed of hydrogen and oxygen (H2O), existing in liquid, gas (steam, water vapour), and solid (ice) states. At room temperature, water is a colourless, odourless, tasteless liquid. One of the most abundant compounds, water covers about 75% of Earth's surface. Life depends on water for virtually every process, its ability to dissolve many other substances being perhaps its most essential quality. Life is believed to have originated in water (the world's oceans or smaller bodies), and living organisms use aqueous solutions (including blood and digestive juices) as mediums for carrying out biological processes. Because water molecules are asymmetric and therefore electric dipoles, hydrogen bonding between molecules in liquid water and in ice is important in holding them together. Many of water's complex and anomalous physical and chemical properties (high melting and boiling points, viscosity, surface tension, greater density in liquid than in solid form) arise from this extensive hydrogen bonding. Water undergoes dissociation to the ions H+ (or H3O+) and OH−, particularly in the presence of salts and other solutes; it may act as an acid or as a base. Water occurs bound (as water of hydration) in many salts and minerals. It has myriad industrial uses, including as a suspending agent (papermaking, coal slurrying), solvent, diluting agent, coolant, and source of hydrogen; it is used in filtration, washing, steam generation, hydration of lime and cement, textile processing, sulfur mining, hydrolysis, and hydraulics, as well as in beverages and foods. See also hard water; heavy water.
water
water [′wȯd·ər]
Water
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 SO42-, 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
| Table 1. Physical properties of water | |
|---|---|
| Note: 1 cal/(cm.sec-deg) = 418.68 W/(m.°K); 1 ohm-1 cm-1 = 100 Siemens/ min; 1 cal/(g-deg) = 4.186 kJ(kg.°K); 1 centipoise = 10-3 N-sec/m2; 1 dyne/cm = 10-3 N/m. | |
| Density (g/cm3) ice............... | 0.9168 (0°C) |
| liquid............... | 0.99987 (0°C) 1.0000 (3.98°C) 0.99823(20°C) |
| saturated vapor............... | 0.5977 kg/m3 (100°C) |
| Melting point............... | 0°C |
| Boiling point............... | 100°C |
| Critical temperature............... | 374.15°C |
| Critical pressure............... | 218.53kgf/cm2 |
| Critical density............... | 0.325 g/cm3 |
| Heat of fusion............... | 79.7 cal/g |
| Heat of vaporization............... | 539cal/g (100°C) |
| Specific heat conductivity [cal/(cm sec-deg)] ice................ | 5.6 X 10-3 (0°C) |
| liquid............... | 1.43 X10-3(0°C) 1.54X 10-3 (45°C) |
| saturated vapor............... | 5.51 X 10-5 (100°C) |
| Specific electrical conductivity (ohm-1cm-1) ice............... | 0.4 X 10-8(0°C) |
| liquid............... | 1.47 X 10-8 (0°C) 4.41 X10-3(18°C) 18.9 X 10-8 (50°C) |
| Specific heat [cal/(g-deg)] liquid............... | 1.00(15°C) |
| saturated vapor............... | 0.487 (100°C) |
| Dielectric constant ice............... | 74.6 (0°C) |
| liquid............... | 81.0(20°C) |
| saturated vapor............... | 1.007(145°C) |
| Viscosity (centipoises) liquid............... | 1.7921 (0°C) 0.284 (100°C) |
| Surface tension of liquid water at interface with air (dynes/cm)............... | 74.64 (0°C) 62.61 (80°C) |
| Index of refraction (D-line of sodium)............... | 1.33299 (20°C) |
| Speed of sound in water............... | 1,496m/sec(25°C) |
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.
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.
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: Kw = [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.
REFERENCES
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Nekrasov, B. V. Osnovy obshchei khimii, vol. 1. Moscow, 1965.
Furon, R. Problemy vody na zemnom share. Moscow, 1966. (Translated from French.)
Krugovorot vody. Moscow, 1966.
Pounder, E. Fizika I’da. Moscow, 1967. (Translated from English.)
Vinogradov, A. P. Vvedenie v geokhimiiu okeana. Moscow, 1967.
Samoilov, O. la. Struktura vodnykh rastvorov elektrolitov i gidratatsiia ionov. Moscow, 1957.
Izotopnyi analiz vody, 2nd ed. Moscow, 1957.
Termodinamika i stroenie rastvorov. Moscow, 1959.
Kratkaia khimicheskaia entsiklopediia, vol. 1. Moscow, 1961. Pages 605-14.
Water makes possible turgor of tissues, the transport of nutrient substances and exchange products (blood, lymph, and plant sap), physical thermoregulation, and other processes of vital activity. Life probably began in a water environment. In the course of evolution, various water animals and water plants came out onto the land and adapted to a terrestrial form of life; nevertheless, water is still an essential component of the external environment for them as well.
Life without water is impossible. When there is a shortage of water, the vital activity of living organisms is disrupted. Only dormant forms of life-spores and seeds-withstand prolonged deprivation of water well. Plants droop and may die for lack of water; however, the sensitivity of different plants to water deprivation varies. Animals die quickly if deprived of water: a well-fed dog can survive without food for up to 100 days, but without water, it would survive less than 10 days. The water content of living organisms is high (see Table 2).
A living organism’s fluids-in intercellular spaces, lymph, blood, digestive juices, plant sap, and so on-contain free water. Water occurs in the bound state in the tissues of animals and plants - it does not flow out when an organ is cut. Water can cause the swelling of colloids and can bond with proteins and other organic compounds, as well as with ions that are components of cells and tissues (water of hydration). Water molecules that are inside the cell but are not components of the hydrated membrane of ions and molecules are immobile and are easier to draw into the general circulation of water in the organism than are hydrated water molecules.
| Table 2. Water content of various organisms and their organs and tissues | |
|---|---|
| Water content (percent) | |
| Land plants tip of the growing shoot............... | 91-93 |
| leaves............... | 75-86 |
| Seeds of cereals............... | 12-14 |
| Algae............... | 90-98 |
| Mosses and lichens............... | 5-7 |
| Jellyfish............... | 95-96 |
| Earthworms............... | 84 |
| Insects adults............... | 45-65 |
| larvae............... | 58-90 |
| Fish............... | 70 |
| Mammals (including man)............... | 63-68 |
| skeleton............... | 20-40 |
| muscles............... | 75 |
| liver............... | 75 |
| Human brain gray matter............... | 84 |
| white matter............... | 72 |
REFERENCES
Ziukov, A. M. Obmen vody v.organizme: Fiziologiia i patologiia. Kharkov [1929].Danilov, N. V. Fiziologicheskie osnovy pit’evogo rezhima. Moscow 1956.
Kravchinskii, B. D. Fiziologiia vodno-solevogo obmena zhidkostei tela. Leningrad, 1963.
Man’s physiological requirement for water, which is taken in by the organism by drinking and with food, is 3-6 liters per 24-hr period, depending on climatic conditions. A much greater quantity of water is necessary for sanitary and domestic necessities.
The removal of garbage and wastes by a sewer system is possible only when there is a sufficient level of water consumption served by centralized water-supply systems. The level of water consumption (in liters per capita per day) to a certain extent also characterizes the level of public-health measures in population centers (see Table 3).
Scientifically justified hygienic norms for the maximum permissible content of chemical substances in water are of great importance in averting the danger of direct or indirect harmful effects of water on the health and sanitary conditions under which the population lives. These norms are the basis of state standards for the quality of drinking water—COST (All-Union State Standard) 2874—and are mandatory in the planning and use of mains carrying industrial and drinking (municipal) water. In the interests of public health, the quality standards for drinking water were revised in the 1960’s in all socialist countries, in the USA, and in France. International standards for drinking water were promulgated by the World Health Organization in 1963; in the USSR the drafting of a plan for new quality standards for drinking water was completed in 1968.
| Table 3. Norms for household and drinking-water consumption | |
|---|---|
| Level of provision of amenities in residential construction | Per capita water consumption (liters per day; dally average for the year) |
| Buildings with water supplied from hydrants (without sewage systems)............... | 30-50 |
| Buildings with interior water pipes and sewage systems (without baths)............... | 125-150 |
| Buildings with water pipes, sewage systems, baths, and water heaters operating on solid fuel............... | 150-180 |
| Buildings with water pipes, sewage systems, and centralized hot-water supply systems............... | 275-400 |
The indexes of water supply from an epidemiological point of view are (1) the total quantity of bacteria, which are grown in a nutrient medium (agar) at a temperature of 37° C (not more than 100 per milliliter), and (2) the number of intestinal bacilli, which are grown in a dense nutrient medium, concentrated on membrane filters (not more than 3 per liter). When liquid mediums are used the number of titers of intestinal bacilli must be no less than 300. According to the 1968 GOST plan, gram-negative nonsporiferous bacilli (facultative anaerobes), which are capable of fermenting glucose to form acid and gas at a temperature of 35°-37° C in 24 hours, are included among bacteria of the group of intestinal bacilli.
The natural composition of water has long attracted attention as a possible cause of widespread noninfectious diseases. The content of chlorides, sulfates, and products of the breakdown of organic substances (ammonia, nitrites, and nitrates) was regarded only as an indirect indicator of water pollution by household sewage that was dangerous to public health. Regions with a shortage or excess of one or another trace element in their water have been found thanks to the application of new research methods. Distinct changes have been observed in the flora and fauna of these regions. As a result of insufficient or excess intake of trace elements into the organism obtained from water and from food, characteristic diseases have been observed among the population. The development of endemic fluorosis is caused by an insufficient level of fluorine in drinking water; a direct relationship has been discovered between the concentration of fluorine in water and the frequency and severity of tooth decay. The fluorine in drinking water also has an effect on phosphorus-calcium exchange and on the calcification of bones. A small range of concentrations, from the toxic to the physiologically beneficial, is characteristic of fluorine in drinking water. In this regard, it has been established that the fluorine level in drinking water should not exceed 0.7-1.0 mg/l (up to 1.2 in water fluoridation), depending on climatic conditions. The content of nitrates in water was long regarded as an indirect indicator of the domestic pollution of water. However, the presence of high concentrations of nitrates was discovered in natural subterranean water and even in artesian water-bearing horizons (the Moldavian SSR, the Tatar ASSR, and the Vladivostok region). The use in babies’ formula of water containing high concentrations of nitrates causes methemoglobinemia of varying seriousness. Methemoglobinemia caused by nitrates in water is also found in older children and thus can take on the proportions of an endemic disease (see Table 4).
| Table 4. Indexes of the harmfulness off chemical substances (natural and added in the treatment process) in drinking water | |
|---|---|
| Maximum concentration in water (mg/l) | |
| Lead............... | 0.1 |
| Arsenic............... | 0.05 |
| Fluorine............... | 0.7-1.5 |
| Beryllium............... | 0.0002 |
| Molybdenum............... | 0.5 |
| Nitrates (by nitrogen content)............... | 10.0 |
| Polyacrylamide............... | 2.0 |
| Strontium............... | 2.0 |
The chemical substances found in water also include substances that, in small concentrations, change the organoleptic properties of water (odor, taste, transparency, and so on). In native waters, chemical substances (common mineral salts, iron, manganese, copper, zinc, and so on), residual quantities of compounds used as reagents in the treatment of water, and industrial contamination of reservoirs are most frequently responsible for changes in the organoleptic properties of water.
Indicators that ensure favorable organoleptic properties in water are listed in Table 5.
| Table 5. Indexes off beneficial organoleptic properties off water at given content off natural substances or substances added in the purification process | |
|---|---|
| Maximum content In water (mg/l) | |
| Turbidity (on standard scale)............... | 1.5 |
| Iron............... | 0.3 |
| Manganese............... | 0.5 |
| Copper............... | 1.0 |
| Zinc............... | 5.0 |
| Chlorides............... | 350 |
| Sulfates............... | 500 |
| Dry residue............... | 1,000 |
| Tripolyphosphate............... | 5.0 |
| Hexametaphosphate............... | 5.0 |
REFERENCES
Rukovodstvo po kommunal’noi gigiene, vol. 2. Moscow, 1962.Vernadskii, V. I. Biogeokhimicheskie ocherki: 1922-1932 gg. Moscow-Leningrad, 1940.
Mezhdunarodnye standarty pit’evoi vody, 2nd ed. Moscow, 1964. (Translated.)
S. N. CHERKINSKII [5-50O-1]
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