1911 Encyclopædia Britannica/Water
WATER. Strictly speaking, water is the oxide of hydrogen which is usually stated to have the formula H2O (see below), but in popular use the term is applied to a great variety of different substances, all of which agree, however, in being the water of the chemist modified differently in the several varieties by the nature or proportion of impurities. In all ordinary waters, such as are used for primary purposes, the impurities amount to very little by weight—as a rule to less than 110th of 1%.
Of all natural stores of water the ocean is by far the most abundant, and from it all other water may be said to be derived. From the surface of the ocean a continuous stream of vapour is rising up into the atmosphere to be recondensed in colder regions and precipitated as rain, snow or sleet, &c. Some 811ths of these precipitates of course return directly to the ocean; the rest, falling on land, collects into pools, lakes, rivers, &c., or else penetrates into the earth, perhaps to reappear as springs or wells. As all the saline components of the ocean are nonvolatile, rain water, in its natural state, can be contaminated only with the ordinary atmospheric gases—oxygen, nitrogen and carbon dioxide. Rain water also contains perceptible traces of ammonia, combined as a rule, at least partly, with the nitric acid, which is produced wherever an electric discharge pervades the atmosphere.
Lake waters, as a class, are relatively pure, especially if the mountain slopes over which the rain collects into a lake are relatively free of soluble components. For example, the water of Loch Katrine (Scotland) is almost chemically pure, apart from small, but perceptible, traces of richly carboniferous matter taken up from the peat of the surrounding hills, and which impart to it a faint brownish hue, while really pure water is blue when viewed through a considerable thickness.
River water varies very much in composition even in the same bed, as a river in the course of its journey towards the ocean passes from one kind of earth to others; while, compared with spring waters, relatively poor in dissolved salts, rivers are liable to be contaminated with more or less of suspended matter.
Spring waters, having been filtered through more or less considerable strata of earth, are, as a class, clear of suspended, but rich in dissolved, mineral and organic matter, and may also contain gases in solution. Of ordinarily occurring minerals only a few are perceptibly soluble in water, and of these calcium carbonate and sulphate and common salt are most widely diffused. Common salt, however, in its natural occurrence, is very much localized; and so it comes that spring and well waters are contaminated chiefly with calcium carbonate and sulphate. Of these two salts, however, the former is held in solution only by the carbonic acid of the water, as calcium bicarbonate. But a carbonate-of-lime water, if exposed to the atmosphere, even at ordinary temperatures, loses its carbonic acid, and the calcium carbonate is precipitated. The stalactites (q.v.) which adorn the roofs and sides of certain caverns are produced in this manner. Many waters are valuable medicinal agents owing to their contained gases and salts (see Mineral Waters).
In addition to its natural components, water is liable to be contaminated through accidental influxes of foreign matter. Thus, for instance, all the Scottish Highland lochs are brown through the presence in them of dissolved peaty matter. Rivers flowing through, or wells sunk in, populous districts may be contaminated with excrementitious matter, discharges from industrial establishments, &c. The presence of especially nitrogenous organic matter is a serious source of danger, inasmuch as such matter forms the natural food or soil for the development of micro-organisms, including those kinds of bacteria which are now supposed to propagate infectious diseases. Happily nature has provided a remedy. The nitrogenous organic matter dissolved in (say) a river speedily suffers disintegration by the action of certain kinds of bacteria, with formation of ammonia and other (harmless) products; and the ammonia, again, is no sooner formed than, by the conjoint action of other bacteria and atmospheric oxygen, it passes first into (salts of) nitrous and then nitric acid A water which contains combined nitrogen in the form of nitrates only is, as a rule, safe organically; if nitrites are present it becomes liable to suspicion; the presence of ammonia is a worse symptom, and if actual nitrogenous organic matter is found in more than microscopic traces the water is possibly (not necessarily) a dangerous water to drink.
All waters, unless very impure, become safe by boiling, which process kills any bacteria or germs that may be present.
Of the ordinary saline components of waters, soluble magnesium and calcium salts are the only ones which are objectionable sanitarily if present in relatively large proportion. Calcium carbonate is harmless; but, on the other hand, the notion that the presence of this component adds to the value of a water as a drinking water is a mistake. The farinaceous part of food alone is sufficient to supply all the lime the body needs, besides, it is questionable whether lime introduced in any other form than that of phosphate is available for the formation of, for instance, bone tissue.
The fitness of a water for washing is determined by its degree of softness. A water which contains lime or magnesia salts decomposes soap with formation of insoluble lime or magnesia salts of the fatty acids of the soap used. So much of the soap is simply wasted; only the surplus can effect any detergent action. Several methods for determining the hardness of a water have been devised. The most exact method is to determine the lime and magnesia gravimetrically or by alkalimetry; or by Clark’s soap test, but this process frequently gives inaccurate results. In this method, which, however, is largely used, a measured volume of the water is placed in a stoppered bottle, and a standard solution of soap is then dropped in from a graduated vessel, until the mixture, by addition of the last drop of soap, has acquired the property of throwing up a peculiar kind of creamy froth when violently shaken, which shows that all the soap-destroying components have been precipitated. The volume of soap required measures the hardness of the water. The soap-solution is referred to a standard by means of a water of a known degree of hardness prepared from a known weight of carbonate of lime by converting it into neutral chloride of calcium, dissolving this in water and diluting to a certain volume. The hardness is variously expressed. On Clark’s scale it is the grains of calcium carbonate per gallon of 70,000 grains; in Germany the parts of lime per 100,000 of water, and in France the parts of calcium carbonate per 100,000. On the English scale, a water of 15° and over is hard, between 5° and 15° moderately hard, and of less than 5° soft.
That part of the hardness of a water which is actually owing to carbonate of lime (or magnesia) can easily be removed in two ways, (1) By boiling, the free carbonic acid goes off with the steam, and the carbonate of lime, being bereft of its solvent, comes down as a precipitate which can be removed by filtration, or by allowing it to settle, and decanting off the clear supernatant liquor. (2) A method of Clark’s is to mix the water with just enough of milk of lime to convert the free carbonic acid into carbonate. Both this and the original carbonate of lime are precipitated, and can be removed as in the first case.
From any uncontaminated natural water pure water is easily prepared. The dissolved salts are removed by distillation; if care be taken that the steam to be condensed is dry, and if its condensation be effected within a tube made of a suitable metal (platinum or silver are best, but copper or block tin work well enough for ordinary purposes), the distillate can contain no impurities except atmospheric gases, which latter, if necessary, must be removed by boiling the distilled water in a narrow-necked flask until it begins to “bump,” and then allowing it to cool in the absence of air. This latter operation ought, strictly speaking, to be performed in a silver or platinum flask, as glass is appreciably attacked by hot water. For most purposes distilled water, taken as it comes from the condenser, is sufficiently pure. The preparation of absolutely pure water is a matter of great difficulty. Stas, in his stoichiometric researches, mixed water with potassium manganate, and distilled after twenty-four hours; the product being re distilled and condensed in a platinum tube just before it was required.
Pure water, being so easily procured in any quantity, is used largely as a standard of reference in metrology and in the quantitative definition of physical properties. Thus a “gallon” is defined as the volume at 62° F. of a quantity of water whose uncorrected mass, as determined by weighing in air of 30-in. pressure and 62° F of temperature, is equal to 10 ℔ avoirdupois. The kilogramme in like manner is defined as the mass of 1 cubic decimetre of water, measured at the temperature corresponding to its maximum density (4° C.). The two fixed points of the thermometer correspond—the lower (0° C., or 32° F.) to the temperature at which ice melts, the upper (100° C., or 212° F.) to that at which the maximum tension of steam, as it rises from boiling water, is equal to 760 mm. or 30-in. mercury pressure. 30 in. being a little more than 760 mm., 212° F. is, strictly speaking, a higher temperature than 100° C., but the difference is very trifling. Specific heats arc customarily measured by that of water, which is taken as =1. All other specific heats of liquids or solids (with one exception, formed by a certain strength of aqueous methyl alcohol) are less than 1. The temperate character of insular climates is greatly owing to this property of water Another physio graphically important peculiarity of water is that it expands on freezing (into ice), while most other liquids do the reverse. 11 volumes of ice fuse into only 10 volumes of water at 0° C.; and the ice-water produced, when brought up gradually to higher and higher temperatures, again exhibits the very exceptional property that it contracts between 0° and 4° C. (by about 110000 of its volume) and then expands again by more and more per degree of increase of temperature, so that the volume at 100° C. is 1·043 times that at 4° C.
In former times water was viewed as an “element,” and the notion remained in force after this term (about the time of Boyle) had assumed its present meaning, although cases of decomposition of water were familiar to chemists. In Boyle’s time it was already well known that iron, tin and zinc dissolve in aqueous hydrochloric or sulphuric acid with evolution of a stinking inflammable gas. Even Boyle, however, took this gas to be ordinary air contaminated with inflammable stinking oils. This view was held by all chemists until Cavendish, before 1784, showed that the gas referred to. if properly purified, is free of smell and constant in Its properties, which are widely different from those of air—the most important point of difference being that the gas when kindled in air burns with evolution of much heat and formation of water. Cavendish, however, did not satisfy himself with merely proving this fact qualitatively; he determined the quantitative relations, and found that it takes very nearly 1000 volumes of air to burn 423 volumes of “hydrogen” gas; but 1000 volumes of air, again, according to Cavendish, contain 210 volumes of oxygen; hence, very nearly, 2 volumes of hydrogen take up 1 volume of oxygen to become water. This important discovery was only confirmed by the subsequent experiments of Humboldt and Gay-Lussac, which were no more competent than Cavendish’s to prove that the surplus of 3 units (423 volumes instead of 420) of hydrogen was an observational error. More recent work, e.g. of Morley, Leduc and Scott, has shown that the ratio is not exactly 2·1. The gravimetric composition was determined by Berzelius and Dulong, and later by Dumas by passing pure hydrogen over red-hot copper oxide. It has also been determined by several other variations and methods (see Hydrogen).
The molecular weight of liquid water has attracted much attention, for it was perceived long ago that its high boiling point, refractive index and other properties were not consistent with the simple formula H2O. Cryoscopic measurements led to the probable formula (H2O)2, whilst the surface tension leads to (H2O)4. The question has been considered by H. E. Armstrong, who suggests that the simple molecule, H2O, which he calls hydrone, condenses in liquid water to form cyclic or chained compounds, containing tetravalent oxygen, resembling in structure the polymethylenes or paraffins.