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calcium (kălˈsēəm) [Lat.,=lime], metallic chemical element; symbol Ca; at. no. 20; at. wt. 40.078; m.p. about 839℃; b.p. 1,484℃; sp. gr. 1.55 at 20℃; valence +2. Calcium is a malleable, ductile, silver-white, relatively soft metal with face-centered, cubic crystalline structure. Chemically it resembles strontium and barium; it is classed with them as an alkaline-earth metal in Group 2 of the periodic table. Calcium is chemically active; it tarnishes rapidly when exposed to air and burns with a bright yellow-red flame when heated, mainly forming the nitride. It reacts directly with water, forming the hydroxide. It combines with other elements, e.g., with oxygen, carbon, hydrogen, chlorine, fluorine, arsenic, phosphorus, and sulfur, forming many compounds.
Although lime (calcium oxide) has been known since ancient times, elemental calcium was first isolated by Sir Humphry Davy in 1808. Today, calcium metal is usually prepared by electrolysis of fused calcium chloride to which a little calcium fluoride has been added. It is used in alloys with other metals, such as aluminum, lead, or copper; in preparation of other metals, such as thorium and uranium, by reduction; and (like barium) in the manufacture of vacuum tubes to remove residual gases.
The metal is of little commercial importance compared to its compounds, which are widely and diversely used. The element is a constituent of lime (see calcium oxide), chloride of lime (bleaching powder), mortar, plaster, cement (see cement, concrete, whiting, putty, precipitated chalk, gypsum, and plaster of Paris. Tremolite, a form of asbestos, is a naturally occurring compound of calcium, magnesium, silicon, and oxygen. Calcium carbide reacts with water to form acetylene gas; it is also used to prepare calcium cyanamide, which is used as a fertilizer. The phosphate is a major constituent of bone ash. The arsenate and the cyanide are used as insecticides. Calcium bicarbonate causes temporary hardness in water; calcium sulfate causes permanent hardness. Generally, calcium compounds show an orange or yellow-red color when held in the Bunsen burner flame.
Although calcium is the fifth most abundant element in the earth's crust, of which it constitutes about 3.6%, it is not found uncombined. It is found widely distributed in its compounds, e.g., Iceland spar, marble, limestone, feldspar, apatite, calcite, dolomite, fluorite, garnet, and labradorite. It is a constituent of most plant and animal matter.
Calcium is essential to the formation and maintenance of strong bones and teeth; the recommended daily dietary allowance for all but young children ranges from 1,000 to 1,300 milligrams. In the human adult the bone calcium is chiefly in the form of the phosphate and carbonate salts. A sufficient store of vitamin D (see under vitamin) in the body is necessary for the proper utilization of calcium. Calcium also functions in the regulation of the heartbeat and in the conversion of prothrombin to thrombin, a necessary step in the clotting of blood.
(Ca), a chemical element in group IIA of Mendeleev’s periodic table. Atomic number, 20; atomic weight, 40.08; a silver-white, lightweight metal. The natural element consists of a mixture of six stable isotopes: 40Ca, 42Ca, 43Ca, 44Ca, 46Ca, and 48Ca, the most abundant of which is the first (96.97 percent).
Calcium compounds, such as limestone, marble, and gypsum (as well as lime, the product of limestone calcination), have been used in construction since antiquity. Chemists considered lime to be an element as recently as the end of the 18th century, but in 1789 A. Lavoisier suggested that lime, magnesia, baryta, alumina, and silica are compound materials. H. Davy electro-lyzed a mixture of moist slaked lime and mercuric oxide in 1808 using a mercury cathode to prepare a calcium amalgam from which he removed the mercury by distillation and obtained a metal that he named calcium (from the genitive calcis of the Latin calx, “lime”).
Occurrence in nature. Calcium is the fifth most abundant element in the earth’s crust (after oxygen, silicon, aluminum, and iron): 2.96 percent by weight. It migrates vigorously and accumulates in various geochemical systems, forming 385 minerals (the fourth largest number of minerals formed by an element). The earth’s mantle contains little calcium and the earth’s core probably contains even less (iron meteorites contain 0.02 percent). Calcium predominates in the lower part of the earth’s crust, accumulating in basic rocks. A large amount of calcium is found in the feldspar anorthite, Ca[Al2Si2O8]; basic rocks contain 6.72 percent and acidic rocks, such as granites, contain 1.58 percent. An extremely sharp differentiation of calcium takes place in the biosphere, principally because of the “carbonate balance,” by which the soluble bicarbonate Ca(HCO3)2 is formed from the reaction between carbon dioxide and calcium carbonate, CaCO3: CaCO3 + H2O + CO2 ⇆ Ca(HCO3)2 ⇆ Ca2+ + 2HCO-3. This reaction is reversible and the basis for the redistribution of the element. Calcium dissolves in water with a high CO2 content, but at a low CO2 content the mineral calcite (CaCO3) precipitates to form large deposits of limestone, chalk, and marble.
Biogenic migration also plays a very large role in the history of the element. Calcium is chief among the metallic elements in living matter. Organisms are known that contain more than 10 percent calcium (more than carbon) and build their skeleton from calcium compounds (mainly CaCO3); these include calcif-erous algae, many mollusks, echinoderms, corals, and rhizopods. Deposits of the skeletons of marine animals and plants lead to the accumulation of colossal masses of algal, coral, and other limestones, which, buried deep within the earth and mineralized, are transformed into various types of marble.
Large territories with damp climates (wooded zones, tundra) are characterized by a lack of calcium, which is readily leached out of the soils. This phenomenon is a cause of low soil fertility, low productivity and poor growth in domestic animals, and, frequently, of diseases of the skeleton. Liming of soils and supplemental feeding of domestic animals and birds are therefore of great importance. Conversely, CaCO3 is poorly soluble in dry climates, and steppe and desert regions are therefore rich in calcium. Gypsum (CaSO4⋅2H2O) is frequently enriched in solonchak and salt lakes.
Rivers carry large quantities of calcium to the oceans; however, the calcium does not remain in the ocean water (average content, 0.04 percent) but concentrates in the skeletons of organisms and settles to the bottom after their death (mostly in the form of CaCO3). Limestone sediments are widely distributed over the bottom of all oceans at depths to 4, 000 m (CaCO3 dissolves at greater depths and organisms at those depths frequently suffer a calcium deficit).
Underground waters play an important role in the migration of calcium. In places they vigorously leach out CaCO3 in masses of limestone, which results in the formation of karst and of caverns, stalactites, and stalagmites. Not only calcite but also calcium phosphates (for example, the Karatau phosphorite deposits in Kazakhstan) and dolomite (CaCO3 ⋅MgCO3) were widely deposited in the oceans of past geological epochs. Gypsum was deposited upon the evaporation of lagoons.
In the course of geologic history the biogenic formation of carbonate increased and the chemical precipitation of calcite decreased. The Precambrian oceans (more than 600 million years ago) contained no animals with limestone skeletons; these animals first became widespread in the Cambrian (corals, sponges). An explanation is found in the high CO2 content of the Precambrian atmosphere.
Physical and chemical properties. The crystal lattice of the a-form of calcium (stable at ordinary temperatures) is face-centered cubic; a = 5.56 Å. Its atomic radius is 1.97 Å; the ionic radius of Ca2+, 1.04 Å. Its density is 1.54 g per cm3 (20°C). The hexagonal β-form is stable above 464°C; melting point, 851°C; boiling point, 1482°C; temperature coefficient of linear expansion, 22 × 10-6 (0°-300°C); thermal conductivity at 20°C, 125.6 watts per m°K, or 0.3 calories per cm.sec«°C); specific heat capacity (0°-100°C), 623.9 joules per kg.°K), or 0.149 cal per g-°C); specific electrical resistivity at 20°C, 4.6 × 10 -8 ohm-m, or 4.6 × 10”6 ohm • cm; temperature coefficient of electrical resistivity, 4.57 × 10-3 (20°C); modulus of elasticity, 26 giga-newtons per m2 (2, 600 kilograms-force per mm2); tensile strength, 60 meganewtons per m2 (6 kgf/mm2); elastic limit, 4 MN/m2 (0.4 kgf/mm2); yield point, 38 MN/m2 (3.8 kgf/ mm2); relative elongation, 50 percent; Brinell hardness, 200-300 MN/m2 (20-30 kgf/mm2). Calcium of sufficiently high purity is ductile. It may be pressed, rolled, and cut.
The configuration of the outer electron shell of the calcium atom is 4s2, so that calcium is divalent in its compounds. Calcium is very active chemically. At ordinary temperatures it reacts very readily with atmospheric oxygen and moisture; it is therefore stored in hermetically sealed vessels or under mineral oil. When heated in air or in oxygen it ignites, resulting in the formation of the basic oxide CaO. The calcium peroxides CaO2 and CaO4 also occur. Calcium reacts with cold water rapidly at first, but the reaction is then retarded because of the formation of a film of Ca(OH)2. Calcium reacts vigorously with hot water and with acids, giving off H2 (this does not occur with concentrated HNO3). Calcium reacts with flourine at lower temperatures and with chlorine and bromine above 400°C, yielding CaF2, CaCl2, and CaBr2, respectively. In the molten state these halides react with calcium to form the so-called subcompounds CaF and CaCl, in which the calcium is formally univalent. Heating calcium with sulfur yields calcium sulfide (CaS), which adds additional sulfur to give polysulfides, such as CaS2 and CaS4. Calcium reacts with dry hydrogen at 300°-400°C to give the hydride CaH2, an ionic compound in which hydrogen is the anion. Calcium and nitrogen react at 500°C to give the nitride Ca3N2. The reaction of calcium with ammonia at low temperatures gives the complex ammoniate Ca[NH3]6. Heating calcium with silicon, phosphorus, or graphite in the absence of air gives calcium silicides (Ca2Si, CaSi, and CaSi2), calcium phosphide (Ca3P2), and calcium carbide (CaC2), respectively. Calcium forms intermetallic compounds with aluminum, silver, gold, copper, lithium, magnesium, lead, and tin.
Production and uses. Calcium is produced industrially by two methods: (1) by heating a briquetted mixture of CaO and powdered aluminum at 1200°C in a vacuum of 0.01-0.02 mm Hg; the calcium vapors, evolved according to the equation 6CaO + 2A1 = 3CaO ⋅ A12O3 + 3Ca, condense on a cold surface; (2) by electrolyzing a melt of CaCl2 and KCl, using a liquid copper-calcium cathode, to give a copper-calcium alloy (65 percent calcium) from which the calcium is distilled at 950°-1000°C in a vacuum of 0.1-0.001 mm Hg.
The pure metal calcium is used as an agent for reducing uranium, thorium, chromium, vanadium, zirconium, cesium, rubidium, and some rare-earths from their compounds. Metallic calcium is also used for deoxidizing steels, bronzes, and other alloys, removing sulfur from petroleum products, dehydrating organic liquids, removing nitrogenous impurities from argon, and as a gas absorber in electrovacuum instruments. Antifrictional materials based on a lead-sodium-calcium system and lead-calcium alloys for the fabrication of electric-cable sheaths are used widely in industry. A calcium-silicon-calcium alloy (silicocalcium) is used as a deoxidizing agent and degasifier in the production of high-quality steels. The uses of calcium compounds are described in separate articles.
A. IA. FISHER and I. A. PEREL’MAN
Calcium in the organism. Calcium is one of the biogenic elements, that is, elements that are necessary for the normal progress of the life processes. It is present in all tissues and fluids in animals and plants. Only rarely are organisms capable of developing in an environment lacking in calcium. In some organisms the calcium content is as high as 38 percent; in man, the content is 1.4-2 percent. The cells of plants and animals require strictly fixed proportions of Ca2+, Na+, and K+ ions in the extracellular fluids. Plants obtain calcium from the soil. Plants are divided into the calcephiles and calcephobes, depending on their behavior toward the element. Animals obtain calcium from food and water.
Calcium is essential for the formation of a number of cell structures, for the maintenance of the normal permeability of the outer cell membranes, for the fertilization of the egg cells of fish and other animals, and for the activation of a number of enzymes. Ca2+ ions transmit stimuli to muscle fiber and bring about its contraction, they increase the strength of cardiac contractions, they increase the phagocytic function of leukocytes, they activate the immunoprotein system in blood, and they participate in blood coagulation. Almost all of the calcium in cells is present in the form of compounds with proteins, nucleic acids, phospholipids, and complexes with inorganic phosphates and organic acids. In the blood plasma of man and the higher animals as little as 20-40 percent of the calcium may be combined with proteins. In animals with skeletons as much as 97-99 percent of the calcium is used as structural material: in invertebrates, mainly in the form of CaCO3 (the shells of mollusks and corals), and in vertebrates, mainly as phosphates. Many invertebrate animals store calcium for the construction of a new skeleton prior to shedding their shells or for providing for the vital functions in unfavorable conditions.
The calcium content of human blood and the blood of other higher animals is regulated by hormones of the parathyroid and thyroid glands. Vitamin D plays an important role in these processes. Calcium absorption takes place in the anterior section of the small intestine. Calcium assimilation decreases when the acidity of the intestinal tract is lowered and is a function of the calcium-phosphorus-fat ratio in the food. The optimum ratios of calcium to phosphorus are about 1.3 in cow’s milk, 0.15 in potatoes, 0.13 in beans, and 0.016 in meat. In cases of excess phosphorus or oxalic acid in the food, calcium absorption decreases. Absorption is accelerated by the bile acids. For man, the optimum proportion of calcium to fat in food is 0.04-0.08 g Ca to 1 g fat.
The elimination of calcium takes place mainly through the intestinal tract. Mammals lose large quantities of the elementwith their milk during lactation. Upsets in phosphorus-calciummetabolism in young animals and children lead to rickets, andin adults, to changes in the composition and structure of theskeleton (osteomalacia).
I. A. SKUL’SKII
In medicine. The use of calcium preparations eliminates disturbances related to a deficiency of Ca2+ ions in the body (for example, in tetanus, spasmophilia, and rickets). Calcium preparations reduce increased sensitivity to allergens and are used in the treatment of allergic conditions (serum allergy, nettle rash, angioneurotic edema, hay fever). They lower increased permeability of the blood vessels and have an anti-inflammatory effect. They are used for hemorrhaging vasculitis, radiation sickness, inflammatory and exudative processes (pneumonia, pleuritis, endometritis), and certain skin conditions. They are prescribed to control bleeding, improve the activity of heart muscle, and reinforce the action of digitalis preparations. They are used as weak diuretics and as antitoxins in cases of magnesium-salt poisoning. Jointly with other drugs, they are used to stimulate labor. Calcium chloride is prescribed orally and intravenously. Ossocal-cinol (a 15-percent sterile suspension of a specially prepared bone meal in persic oil) is suggested for tissue therapy.
Calcium preparations also include gypsum (CaSO4), which is used in surgery for casts, and chalk (CaCO3), which is prescribed for internal use in cases of elevated gastric acidity and for the preparation of tooth powder.
REFERENCESKratkaia khimicheskaia entsiklopediia, vol. 2. Moscow, 1963. Pages 370-75.
Rodiakin, V. V. Kal’tsii, ego soedineniia i splavy. Moscow, 1967.
Kaplanskii, S. Ia. Mineral’nyi obmen. Moscow-Leningrad, 1938.
Vishniakov, S. I. Obmen makroelementov u sel’skokhoziaistvennykh zhi-votnykh. Moscow, 1967.