cause he prepared it by heating a mixture of sal ammoniac and lime.

It is found in sea water and in many mineral waters; it also occurs as tachydrite at Stassfurt; this is a compound with magnesium chloride, having the formula CaCl,, 2MgCl2. 12H2O.

Preparation.-Crude calcium chloride is a by-product in several chemical operations; for example, it is the residue in the preparation of ammonia from ammonium chloride; it is obtained in the preparation of potassium chlorate, in the ammonia-soda process, and in the preparation of carbon dioxide from limestone and hydrochloric acid. For many uses this crude product, which is in aqueous solution, is evaporated to dryness in an iron kettle, and then heated until it melts. The product is sent into commerce for dehydrating purposes as fused calcium chloride.

Purified calcium chloride is made by nearly saturating hydrochloric acid with marble, adding chlorine water to oxidize the iron and manganese compounds, and precipitating these by the addition of a slight excess of milk of lime (calcium hydrate). The clear, filtered solution, which is slightly alkaline, is carefully neutralized with hydrochloric acid. This solution is then evaporated to the crystallizing point or to dryness, according to the uses to which it is to be applied.

Properties. The crystallized salt occurs in large, hexagonal prisms, having the composition CaCl2.6H2O. On the application of heat it melts at 29°. These crystals rapidly deliquesce on exposure to the air, forming a thick, oily liquid, formerly known as oleum calcis.

When the crystals dissolve in water a considerable fall in temperature takes place. A temperature of -48° is obtained by mixing parts of the crystals with 1 part of snow. When heated to 200°, or when kept for some time over concentrated sulphuric acid, the crystals lose 4 molecules of water of crystallization, a white powder resulting, which possesses energetic dehydrating properties. Above 200° the remaining 2 molecules of water are driven off, and at 720° the salt becomes anhydrous.

The official salt is directed to be "calcium chloride, rendered anhydrous by fusion at the lowest possible temperature." described as occurring in "white, slightly translucent, hard fragments, odorless, having a sharp, saline taste, and very deliquescent."

The anhydrous salt is soluble in 1.5 parts of water, and in 8 parts of alcohol at 15°, in 1.5 parts of boiling alcohol, and very

freely soluble in boiling water; it is insoluble in ether. Since the aqueous solution is used for maintaining a temperature above that of boiling water, it is useful to know the boiling point of such a solution; that containing 50 parts of the anhydrous salt in 100 of water boils at 112°, that containing 200 per cent. of the salt boils at 158°, and that containing 325 per cent. boils at 180°.

The pure salt dissolves in water without residue, and has a neutral reaction. If, however, it be kept at or above the fusing point for some time, a slight decomposition takes place, so that it leaves a residue insoluble in water, and the solution has a faintly alkaline reaction.

Uses. Calcium chloride has some use in the laboratory as a reagent, but its chief value is as a desiccating agent, in the drying of gases, and in the concentration of liquids. Its solution is valuable, as above stated, for use in water-baths, where it is desired to maintain a constant temperature above that of boiling

water.

Calcium Bromide, CaBr. Calcii Bromidum, U. S. P.— Like the preceding salt, calcium bromide is found in certain mineral waters.

It is prepared by neutralizing hydrobromic acid with marble, adding bromine water to oxidize iron and manganese compounds, and precipitating these by the addition of a slight excess of milk of lime. The solution is filtered, and the filtrate carefully neutralized with hydrobromic acid. The resulting solution is then evaporated to dryness and carefully heated to 680°, whereby an anhydrous salt is obtained.

The Pharmacopoeia directs that the anhydrous salt be employed. In this state it occurs as white, granular, very deliquescent salt, without odor, and with a sharp, saline taste. At 15° it is soluble in 0.7 part of water, and in 1 part of alcohol; at the boiling point these liquids dissolve it freely. At 680° the saltmelts, above that temperature it is slowly decomposed with loss of bromine. The aqueous solution is neutral to litmus paper.

Uses. The principal use of this salt is in medicine. It is supposed to have some advantages over the other bromides.

Calcium Iodide, Cal, is prepared like the bromide, which it resembles in many particulars. Both salts are so extremely deliquescent that considerable care is necessary in order to keep them in the solid state.

Calcium Fluoride, CaF2, under the name of fluor-spar, occurs quite abundantly in nature. It is found especially in the limestone caves of the Castleton Valley in Derbyshire, England; considerable quantities are

also found in Saxony. It occurs in cubes and octohedra, and in some other forms belonging to the regular system.

It is also found in the ashes of some plants, in bones, in the enamel of the teeth, and in sea water and mineral springs.

In the pure condition calcium fluoride is colorless, but it is much oftener of a violet, blue, red, green, or brown color, owing to the presence of small quantities of impurities. At a red heat calcium fluoride fuses without decomposition; while hot it phosphoresces in the dark.

Fluor-spar is extensively employed as a flux in many metallurgical operations, and the finer masses of crystals are made into vases and other ornaments.

CALCIUM AND OXYGEN.

Calcium Monoxide, CaO. Calx, U.S. P.-Lime is prepared by heating calcium carbonate. If a pure carbonate is used a correspondingly pure product results. In a small way this is accomplished by heating calc-spar or a pure marble in a crucible with a hole in the bottom in order that the furnace gases may pass through and carry off the carbon dioxide.

At the present time Such a one is shown

On a commercial scale lime-kilns are used. These were formerly constructed so as to be filled with limestone and fuel in alternate layers. Fire was then started at the bottom, and the temperature regulated by the air-supply below. Such a furnace, however, could only be emptied when cool. continuous furnaces are much more used. in Fig. 66. It is built on the slope of a hill, so as to bring the door at the top on a level with the ground to facilitate the introduction of the limestone. The heat is supplied by two fires, F, F, the finished product is removed at D, Fig. 67. Fig. 66 shows a section cut parallel with the side of the hill, while Fig. 67 shows a section cut at right angles to this, by which the draw-hole D is shown. These furnaces are used in the well-known lime districts of Chester Valley, near Philadelphia. At the ordinary pressure of the atmosphere, a temperature of 812° is required for the dissociation of limestone, but in practice a temperature of 925° is usually employed. If the limestone contain much magnesia it is liable to sinter and form a semi-fused mass, so that in burning this variety a lower heat must be maintained. Coal is the chief fuel used at the present time, although formerly much wood was employed.

Properties.-Pure lime occurs in white, amorphous masses. It often contains iron, and consequently is of a grayish or yellowish shade of color. When exposed to the air it gradually

FIG. 66.

absorbs moisture and carbon

dioxide, and crumbles to a

white powder. Lime is without odor, but has a sharp, caustic taste. It is soluble in about 750 parts of water at the ordinary temperatures, and in 1300 parts of boiling water; insoluble in alcohol. It is not affected by the highest degrees of heat, but the oxyhydrogen flame causes it to emit an intense white light.

When lime is treated with about one-third its weight of water, added drop by drop, it gradually becomes hot, swells to about twice its bulk, and then falls to a white powder.

The solubility and some of the other properties of lime are closely associated or iden FIG. 67.

Lime-kiln.

tical with those of calcium hydrate, since, in the presence of water, the oxide is first converted into hydrate.

Calcium Hydrate, Ca(OH)2, Calcium Hydroxide.-When lime is treated with one-third its weight of water, as stated above, the product is calcium hydrate. Considerable heat is developed in

this operation where the proportions are carefully adjusted a temperature of 150° is easily attained. The reaction involved in the formation of the hydrate is as follows:

CaO+ H2O = Ca(OH)2

The resulting product is a fine, white, impalpable powder. Its solution in water forms Liquor Calcis, U. S. P., or lime water. This solution has a saline and somewhat caustic taste and an alkaline reaction. On boiling it becomes turbid, owing to the lesser solubility of the hydrate in hot than in cold water. When the clear solution is exposed to the air, a pellicle forms on the surface, due to the formation of calcium carbonate, resulting from the absorption of carbon dioxide from the atmosphere. Calcium hydrate may be obtained by adding sodium or potassium hydrate to a concentrated solution of calcium chloride.

Crystals of calcium hydrate, in the form of tablets or small prisms, may be obtained by evaporating the clear solution of lime water in a vacuum over sulphuric acid.

Milk of lime is a mixture of calcium hydrate and water. It may be made of any desired strength by agitating the freshlyslaked lime with water.

Uses. Calcium oxide, and therefore also the hydrate, is largely used as a laboratory reagent. The oxide finds much use as a desiccating agent for gases. The hydrate is considerably employed in medicine. The hydrate is largely used in the manufacture of mortar.

Mortars and Cements.-When freshly slaked lime of the consistence of paste is mixed with an equal volume of water, and then 3 to 4 times as much sand is added as there was lime employed, the result is mortar.

The process of hardening which mortar undergoes is called "setting." This setting is sufficiently complete in a few days to give stability to a structure in which it is employed, but the hardening process continues for years. The peculiar action of mortar appears to be due to loss of water and subsequent absorption of carbon dioxide, rather than to the formation of calcium silicate. The value of a mortar depends to a great degree upon the thorough mixing of the ingredients in the process of manufacture. The nature of hydraulic cements will be explained under “Clays" in the section on alumina.

Calcium Dioxide, CaO2, is prepared by precipitating lime water with hydrogen dioxide. This precipitate has the composition CaO2.8H2O;