What substances contain calcium? Calcium as a chemical element, its role

Calcium—element of the main subgroup of the second group, fourth period periodic table chemical elements of D.I. Mendeleev, with atomic number 20. Denoted by the symbol Ca (lat. Calcium). The simple substance calcium (CAS number: 7440-70-2) is a soft, chemically active alkaline earth metal, silver- white.

History and origin of the name

The name of the element comes from Lat. calx (in the genitive case calcis) - “lime”, “soft stone”. It was proposed by the English chemist Humphry Davy, who isolated calcium metal by the electrolytic method in 1808. Davy electrolyzed a mixture of wet slaked lime and mercuric oxide HgO on a platinum plate, which served as the anode. The cathode was a platinum wire immersed in liquid mercury. As a result of electrolysis, calcium amalgam was obtained. Having distilled mercury from it, Davy obtained a metal called calcium. Calcium compounds - limestone, marble, gypsum (as well as lime - a product of calcination of limestone) have been used in construction for several thousand years ago. Until the end of the 18th century, chemists considered lime to be a simple solid. In 1789, A. Lavoisier suggested that lime, magnesia, barite, alumina and silica are complex substances.

Being in nature

Due to its high chemical activity, calcium does not occur in free form in nature.

Calcium accounts for 3.38% of the mass of the earth's crust (5th most abundant after oxygen, silicon, aluminum and iron).

Isotopes

Calcium occurs in nature as a mixture of six isotopes: 40 Ca, 42 Ca, 43 Ca, 44 Ca, 46 Ca and 48 Ca, of which the most common is 40 Ca and accounts for 96.97%.

Of the six natural isotopes of calcium, five are stable. The sixth isotope 48 Ca, the heaviest of the six and very rare (its isotopic abundance is only 0.187%), was recently discovered to undergo double beta decay with a half-life of 5.3 x 10 19 years.

In rocks and minerals

Most of the calcium is contained in silicates and aluminosilicates of various rocks (granites, gneisses, etc.), especially in feldspar - Ca anorthite.

In the form of sedimentary rocks, calcium compounds are represented by chalk and limestones, consisting mainly of the mineral calcite (CaCO 3). The crystalline form of calcite - marble - is much less common in nature.

Calcium minerals such as calcite CaCO 3 , anhydrite CaSO 4 , alabaster CaSO 4 ·0.5H 2 O and gypsum CaSO 4 ·2H 2 O, fluorite CaF 2 , apatites Ca 5 (PO 4) 3 (F,Cl, OH), dolomite MgCO 3 ·CaCO 3 . The presence of calcium and magnesium salts in natural water determines its hardness.

Calcium, vigorously migrating in the earth's crust and accumulating in various geochemical systems, forms 385 minerals (the fourth largest number of minerals).

Migration in the earth's crust

In the natural migration of calcium, a significant role is played by “carbonate equilibrium”, associated with the reversible reaction of interaction of calcium carbonate with water and carbon dioxide with the formation of soluble bicarbonate:

CaCO 3 + H 2 O + CO 2 ↔ Ca (HCO 3) 2 ↔ Ca 2+ + 2HCO 3 -

(equilibrium shifts to the left or right depending on the concentration of carbon dioxide).

Biogenic migration plays a huge role.

In the biosphere

Calcium compounds are found in almost all animal and plant tissues (see also below). A significant amount of calcium is found in living organisms. Thus, hydroxyapatite Ca 5 (PO 4) 3 OH, or, in another entry, 3Ca 3 (PO 4) 2 ·Ca(OH) 2, is the basis of the bone tissue of vertebrates, including humans; The shells and shells of many invertebrates, eggshells, etc. are made of calcium carbonate CaCO 3. In living tissues of humans and animals there is 1.4-2% Ca (by mass fraction); in the body of a person weighing 70 kg, the calcium content is about 1.7 kg (mainly in the intercellular substance of bone tissue).

Receipt

Free metallic calcium is obtained by electrolysis of a melt consisting of CaCl 2 (75-80%) and KCl or CaCl 2 and CaF 2, as well as aluminothermic reduction of CaO at 1170-1200 °C:

4CaO + 2Al = CaAl 2 O 4 + 3Ca.

Properties

Physical properties

Calcium metal exists in two allotropic modifications. Up to 443 °C, α-Ca with a cubic face-centered lattice (parameter a = 0.558 nm) is stable; β-Ca with a cubic body-centered lattice of the α-Fe type (parameter a = 0.448 nm) is more stable. Standard enthalpy Δ H 0 transition α → β is 0.93 kJ/mol.

Chemical properties

In the series of standard potentials, calcium is located to the left of hydrogen. The standard electrode potential of the Ca 2+ /Ca 0 pair is −2.84 V, so that calcium actively reacts with water, but without ignition:

Ca + 2H 2 O = Ca(OH) 2 + H 2 + Q.

The presence of dissolved calcium bicarbonate in water largely determines the temporary hardness of water. It is called temporary because when water boils, bicarbonate decomposes and CaCO 3 precipitates. This phenomenon leads, for example, to the fact that scale forms in the kettle over time.

Application

Applications of calcium metal

Metallothermy

Pure metallic calcium is widely used in metallothermy for the production of rare metals.

Alloying of alloys

Pure calcium is used to alloy lead, which is used for the production of battery plates and maintenance-free starter lead-acid batteries with low self-discharge. Also, metallic calcium is used for the production of high-quality calcium babbits BKA.

Nuclear fusion

The isotope 48 Ca is the most effective and commonly used material for the production of superheavy elements and the discovery of new elements on the periodic table. For example, in the case of using 48 Ca ions to produce superheavy elements in accelerators, the nuclei of these elements are formed hundreds and thousands of times more efficiently than when using other “projectiles” (ions).) is used in the form and for the reduction of metals, as well as in the production of cyanamide calcium (by heating calcium carbide in nitrogen at 1200 °C, the reaction is exothermic, carried out in cyanamide furnaces).

Calcium, as well as its alloys with aluminum and magnesium, are used in backup thermal electric batteries as an anode (for example, calcium-chromate element). Calcium chromate is used in such batteries as a cathode. The peculiarity of such batteries is an extremely long shelf life (decades) in a suitable condition, the ability to operate in any conditions (space, high pressures), and a high specific energy in terms of weight and volume. Disadvantage: short lifespan. Such batteries are used where it is necessary to create colossal electrical power for a short period of time (ballistic missiles, some spacecraft etc.).

In addition, calcium compounds are included in medications for the prevention of osteoporosis, and in vitamin complexes for pregnant women and the elderly.-

Biological role of calcium

Calcium is a common macronutrient in the body of plants, animals and humans. In humans and other vertebrates, most of it is contained in the skeleton and teeth in the form of phosphates. The skeletons of most groups of invertebrates (sponges, coral polyps, shellfish, etc.). Calcium ions are involved in blood clotting processes, as well as in ensuring constant osmotic pressure of the blood. Calcium ions also serve as one of the universal second messengers and regulate a variety of intracellular processes - muscle contraction, exocytosis, including the secretion of hormones and neurotransmitters, etc. The calcium concentration in the cytoplasm of human cells is about 10−7 mol, in intercellular fluids about 10− 3 mol.

Calcium requirements depend on age. For adults, the required daily intake is from 800 to 1000 milligrams (mg), and for children from 600 to 900 mg, which is very important for children due to the intensive growth of the skeleton. Most of the calcium that enters the human body with food is found in dairy products; the remaining calcium comes from meat, fish, and some plant products (legumes are especially high in legumes). Absorption occurs in both the large and small intestines and is facilitated by an acidic environment, vitamin D and vitamin C, lactose, and unsaturated fatty acids. The role of magnesium in calcium metabolism is important; with its deficiency, calcium is “washed out” from the bones and deposited in the kidneys (kidney stones) and muscles.

Aspirin, oxalic acid, and estrogen derivatives interfere with the absorption of calcium. When combined with oxalic acid, calcium produces water-insoluble compounds that are components of kidney stones.

Calcium levels in the blood due to large quantity the processes associated with it are precisely regulated, and with proper nutrition, deficiency does not occur. Prolonged absence from the diet can cause cramps, joint pain, drowsiness, growth defects, and constipation. Deeper deficiency leads to constant muscle cramps and osteoporosis. Abuse of coffee and alcohol can cause calcium deficiency, since some of it is excreted in the urine.

Excessive doses of calcium and vitamin D can cause hypercalcemia, followed by intense calcification of bones and tissues (mainly affecting the urinary system). Long-term excess disrupts the functioning of muscle and nerve tissues, increases blood clotting and reduces the absorption of zinc by bone cells. The maximum daily safe dose for an adult is 1500 to 1800 milligrams.

Pregnant and breastfeeding women - from 1500 to 2000 mg.

Natural calcium compounds (chalk, marble, limestone, gypsum) and the products of their simplest processing (lime) have been known to people since ancient times. In 1808, the English chemist Humphry Davy electrolyzed wet slaked lime (calcium hydroxide) with a mercury cathode and obtained calcium amalgam (an alloy of calcium and mercury). From this alloy, having distilled off mercury, Davy obtained pure calcium.
He also proposed the name of a new chemical element, from the Latin "calx" denoting the name of limestone, chalk and other soft stones.

Finding in nature and obtaining:

Calcium is the fifth most abundant element in the earth's crust (more than 3%), forms many rocks, many of which are based on calcium carbonate. Some of these rocks are of organic origin (shell rock), showing the important role of calcium in living nature. Natural calcium is a mixture of 6 isotopes with mass numbers from 40 to 48, with 40 Ca accounting for 97% of the total. Nuclear reactions have also produced other isotopes of calcium, for example radioactive 45 Ca.
To receive simple substance calcium, electrolysis of molten salts or aluminothermy is used:
4CaO + 2Al = Ca(AlO 2) 2 + 3Ca

Physical properties:

A silver-gray metal with a cubic face-centered lattice, significantly harder than alkali metals. Melting point 842°C, boiling point 1484°C, density 1.55 g/cm3. At high pressures and temperatures of about 20 K it goes into the superconductor state.

Chemical properties:

Calcium is not as active as alkali metals, but it must be stored under a layer of mineral oil or in tightly sealed metal drums. Already at normal temperatures it reacts with oxygen and nitrogen in the air, as well as with water vapor. When heated, it burns in air with a red-orange flame, forming an oxide with an admixture of nitrides. Like magnesium, calcium continues to burn in an atmosphere of carbon dioxide. When heated, it reacts with other non-metals, forming compounds that are not always obvious in composition, for example:
Ca + 6B = CaB 6 or Ca + P => Ca 3 P 2 (also CaP or CaP 5)
In all its compounds, calcium has an oxidation state of +2.

The most important connections:

Calcium oxide CaO- ("quicklime") a white substance, an alkaline oxide, which reacts vigorously with water ("quenched") turning into a hydroxide. Obtained by thermal decomposition of calcium carbonate.

Calcium hydroxide Ca(OH) 2- ("slaked lime") white powder, slightly soluble in water (0.16g/100g), strong alkali. A solution (“lime water”) is used to detect carbon dioxide.

Calcium carbonate CaCO 3- the basis of most natural calcium minerals (chalk, marble, limestone, shell rock, calcite, Iceland spar). In its pure form, the substance is white or colorless. crystals. When heated (900-1000 C) decomposes, forming calcium oxide. Not p-rim, reacts with acids, is able to dissolve in water saturated with carbon dioxide, turning into bicarbonate: CaCO 3 + CO 2 + H 2 O = Ca(HCO 3) 2. The reverse process leads to the appearance of calcium carbonate deposits, in particular formations such as stalactites and stalagmites
It is also found in nature as part of dolomite CaCO 3 *MgCO 3

Calcium sulfate CaSO 4- a white substance, in nature CaSO 4 * 2H 2 O (“gypsum”, “selenite”). The latter, when carefully heated (180 C), turns into CaSO 4 *0.5H 2 O (“burnt gypsum”, “alabaster”) - a white powder, which, when mixed with water, again forms CaSO 4 *2H 2 O in the form of a solid, quite durable material. It is slightly soluble in water, but can dissolve in excess sulfuric acid, forming hydrogen sulfate.

Calcium phosphate Ca 3 (PO 4) 2- (“phosphorite”), insoluble, under the influence strong acids passes into more soluble calcium hydro- and dihydrogen phosphates. Feedstock for the production of phosphorus, phosphoric acid, phosphate fertilizers. Calcium phosphates are also included in apatites, natural compounds with the approximate formula Ca 5 3 Y, where Y = F, Cl, or OH, respectively, fluorine, chlorine, or hydroxyapatite. Along with phosphorite, apatites are part of the bone skeleton of many living organisms, incl. and man.

Calcium fluoride CaF 2 - (natural:"fluorite", "fluorspar"), an insoluble substance of white color. Natural minerals have a variety of colors due to impurities. Glows in the dark when heated and under UV irradiation. It increases the fluidity (“fusibility”) of slags when producing metals, which explains its use as a flux.

Calcium chloride CaCl 2- colorless christ. It is well soluble in water. Forms crystalline hydrate CaCl 2 *6H 2 O. Anhydrous ("fused") calcium chloride is a good desiccant.

Calcium nitrate Ca(NO 3) 2- ("calcium nitrate") colorless. christ. It is well soluble in water. An integral part of pyrotechnic compositions that gives the flame a red-orange color.

Calcium carbide CaС 2- reacts with water, forming acetylene, for example: CaС 2 + H 2 O = С 2 H 2 + Ca(OH) 2

Application:

Metallic calcium is used as a strong reducing agent in the production of some difficult-to-reduce metals (“calciothermy”): chromium, rare earth elements, thorium, uranium, etc. In the metallurgy of copper, nickel, special steels and bronzes, calcium and its alloys are used to remove harmful impurities of sulfur, phosphorus, excess carbon.
Calcium is also used to bind small amounts of oxygen and nitrogen when obtaining high vacuum and purifying inert gases.
Neutron-excess 48 Ca ions are used for the synthesis of new chemical elements, for example element No. 114, . Another isotope of calcium, 45Ca, is used as a radioactive tracer in studies of the biological role of calcium and its migration in the environment.

The main area of application for numerous calcium compounds is the production of building materials (cement, building mixtures, plasterboard, etc.).

Calcium is one of the macroelements in living organisms, forming compounds necessary for the construction of both the internal skeleton of vertebrates and the external skeleton of many invertebrates, the shell of eggs. Calcium ions also participate in the regulation of intracellular processes and determine blood clotting. Lack of calcium in childhood leads to rickets, in old age - to osteoporosis. The source of calcium is dairy products, buckwheat, nuts, and its absorption is facilitated by vitamin D. If there is a lack of calcium, various drugs are used: calcex, calcium chloride solution, calcium gluconate, etc.
The mass fraction of calcium in the human body is 1.4-1.7%, the daily requirement is 1-1.3 g (depending on age). Excessive calcium intake can lead to hypercalcemia - deposition of its compounds in internal organs, and the formation of blood clots in blood vessels. Sources:
Calcium (element) // Wikipedia. URL: http://ru.wikipedia.org/wiki/Calcium (access date: 01/3/2014).
Popular library of chemical elements: Calcium. // URL: http://n-t.ru/ri/ps/pb020.htm (01/3/2014).

Ufa State Petroleum Technical University

Department of General and Analytical Chemistry

on the topic: “The element calcium. Properties, production, application"

Prepared by student of group BTS-11-01 Prokaev G.L.

Associate Professor Krasko S.A.

Introduction

History and origin of the name

Being in nature

Receipt

Physical properties

Chemical properties

Applications of calcium metal

Application of calcium compounds

Biological role

Conclusion

References

Introduction

Calcium is an element of the main subgroup of the second group, the fourth period of the periodic system of chemical elements of D.I. Mendeleev, with atomic number 20. It is designated by the symbol Ca (lat. Calcium). The simple substance calcium (CAS number: 7440-70-2) is a soft, reactive alkaline earth metal of a silvery-white color.

Calcium is called an alkaline earth metal and is classified as an S element. At the outer electronic level, calcium has two electrons, so it gives compounds: CaO, Ca(OH)2, CaCl2, CaSO4, CaCO3, etc. Calcium is a typical metal - it has a high affinity for oxygen, reduces almost all metals from their oxides, and forms a fairly strong base Ca(OH)2.

Despite the ubiquity of element No. 20, even chemists have not all seen elemental calcium. But this metal, both in appearance and in behavior, is not at all similar to alkali metals, contact with which is fraught with the danger of fires and burns. It can be safely stored in air; it does not ignite from water.

Elemental calcium is almost never used as a structural material. He's too active for that. Calcium easily reacts with oxygen, sulfur, and halogens. Even with nitrogen and hydrogen, under certain conditions, it reacts. The environment of carbon oxides, inert for most metals, is aggressive for calcium. It burns in an atmosphere of CO and CO2.

History and origin of the name

Calcium compounds - limestone, marble, gypsum (as well as lime - a product of calcination of limestone) have been used in construction for several thousand years ago. Until the end of the 18th century, chemists considered lime to be a simple solid. In 1789, A. Lavoisier suggested that lime, magnesia, barite, alumina and silica are complex substances.

Being in nature

Due to its high chemical activity, calcium does not occur in free form in nature.

Calcium accounts for 3.38% of the mass of the earth's crust (5th most abundant after oxygen, silicon, aluminum and iron).

Isotopes. Calcium occurs in nature as a mixture of six isotopes: 40Ca, 42Ca, 43Ca, 44Ca, 46Ca and 48Ca, among which the most common - 40Ca - accounts for 96.97%.

Of the six natural isotopes of calcium, five are stable. The sixth isotope, 48Ca, the heaviest of the six and very rare (its isotopic abundance is only 0.187%), was recently discovered to undergo double beta decay with a half-life of 5.3 ×1019 years.

In rocks and minerals. Most of the calcium is contained in silicates and aluminosilicates of various rocks (granites, gneisses, etc.), especially in feldspar - Ca anorthite.

In the form of sedimentary rocks, calcium compounds are represented by chalk and limestones, consisting mainly of the mineral calcite (CaCO3). The crystalline form of calcite - marble - is much less common in nature.

Calcium minerals such as calcite CaCO3, anhydrite CaSO4, alabaster CaSO4 0.5H2O and gypsum CaSO4 2H2O, fluorite CaF2, apatite Ca5(PO4)3(F,Cl,OH), dolomite MgCO3 CaCO3 are quite widespread. The presence of calcium and magnesium salts in natural water determines its hardness.

Calcium, vigorously migrating in the earth's crust and accumulating in various geochemical systems, forms 385 minerals (the fourth largest number of minerals).

Migration in the earth's crust. In the natural migration of calcium, a significant role is played by “carbonate equilibrium”, associated with the reversible reaction of interaction of calcium carbonate with water and carbon dioxide with the formation of soluble bicarbonate:

CaCO3 + H2O + CO2 ↔ Ca (HCO3)2 ↔ Ca2+ + 2HCO3ˉ

(equilibrium shifts to the left or right depending on the concentration of carbon dioxide).

Biogenic migration. In the biosphere, calcium compounds are found in almost all animal and plant tissues (see also below). A significant amount of calcium is found in living organisms. Thus, hydroxyapatite Ca5(PO4)3OH, or, in another entry, 3Ca3(PO4)2·Ca(OH)2, is the basis of the bone tissue of vertebrates, including humans; The shells and shells of many invertebrates, eggshells, etc. are made of calcium carbonate CaCO3. In living tissues of humans and animals there is 1.4-2% Ca (by mass fraction); in a human body weighing 70 kg, the calcium content is about 1.7 kg (mainly in the intercellular substance of bone tissue).

Receipt

Free metallic calcium is obtained by electrolysis of a melt consisting of CaCl2 (75-80%) and KCl or from CaCl2 and CaF2, as well as aluminothermic reduction of CaO at 1170-1200 °C:

CaO + 2Al = CaAl2O4 + 3Ca.

A method has also been developed for producing calcium by thermal dissociation of calcium carbide CaC2

Physical properties

Calcium metal exists in two allotropic modifications. Stable up to 443°C α -Ca with cubic lattice, higher stability β-Ca with cubic body-centered lattice type α -Fe. Standard enthalpy ΔH0 transition α → β is 0.93 kJ/mol.

Calcium is a light metal (d = 1.55), silvery-white in color. It is harder and melts at a higher temperature (851 ° C) compared to sodium, which is located next to it in the periodic table. This is explained by the fact that there are two electrons per calcium ion in the metal. That's why chemical bond It has a stronger bond between ions and electron gas than sodium. At chemical reactions Calcium valence electrons are transferred to atoms of other elements. In this case, doubly charged ions are formed.

Chemical properties

Calcium is a typical alkaline earth metal. The chemical activity of calcium is high, but lower than that of all other alkaline earth metals. It easily reacts with oxygen, carbon dioxide and moisture in the air, which is why the surface of calcium metal is usually dull gray, so in the laboratory calcium is usually stored, like other alkaline earth metals, in a tightly closed jar under a layer of kerosene or liquid paraffin.

In the series of standard potentials, calcium is located to the left of hydrogen. The standard electrode potential of the Ca2+/Ca0 pair is −2.84 V, so that calcium reacts actively with water, but without ignition:

2H2O = Ca(OH)2 + H2 + Q.

Calcium reacts with active non-metals (oxygen, chlorine, bromine) under normal conditions:

Ca + O2 = 2CaO, Ca + Br2 = CaBr2.

When heated in air or oxygen, calcium ignites. Calcium reacts with less active non-metals (hydrogen, boron, carbon, silicon, nitrogen, phosphorus and others) when heated, for example:

Ca + H2 = CaH2, Ca + 6B = CaB6,

Ca + N2 = Ca3N2, Ca + 2C = CaC2,

Ca + 2P = Ca3P2 (calcium phosphide),

calcium phosphides of the compositions CaP and CaP5 are also known;

Ca + Si = Ca2Si (calcium silicide),

Calcium silicides of the compositions CaSi, Ca3Si4 and CaSi2 are also known.

The occurrence of the above reactions, as a rule, is accompanied by the release of a large amount of heat (that is, these reactions are exothermic). In all compounds with non-metals, the oxidation state of calcium is +2. Most of the calcium compounds with non-metals are easily decomposed by water, for example:

CaH2+ 2H2O = Ca(OH)2 + 2H2,N2 + 3H2O = 3Ca(OH)2 + 2NH3.

The Ca2+ ion is colorless. When soluble calcium salts are added to the flame, the flame turns brick-red.

Calcium salts such as CaCl2 chloride, CaBr2 bromide, CaI2 iodide and Ca(NO3)2 nitrate are highly soluble in water. Insoluble in water are fluoride CaF2, carbonate CaCO3, sulfate CaSO4, orthophosphate Ca3(PO4)2, oxalate CaC2O4 and some others.

It is important that, unlike calcium carbonate CaCO3, acidic calcium carbonate (bicarbonate) Ca(HCO3) 2 is soluble in water. In nature, this leads to the following processes. When cold rain or river water, saturated with carbon dioxide, penetrates underground and hits limestone, their dissolution is observed:

CaCO3 + CO2 + H2O = Ca(HCO3)2.

In the same places where water saturated with calcium bicarbonate comes to the surface of the earth and heats up sun rays, the reverse reaction occurs:

Ca(HCO3)2 = CaCO3 + CO2 + H2O.

This is how large masses of substances are transferred in nature. As a result, huge gaps can form underground, and beautiful stone “icicles” - stalactites and stalagmites - form in caves.

The presence of dissolved calcium bicarbonate in water largely determines the temporary hardness of water. It is called temporary because when water boils, bicarbonate decomposes and CaCO3 precipitates. This phenomenon leads, for example, to the fact that scale forms in the kettle over time.

calcium metal chemical physical

The main use of calcium metal is as a reducing agent in the production of metals, especially nickel, copper and stainless steel. Calcium and its hydride are also used to produce difficult-to-reduce metals such as chromium, thorium and uranium. Calcium-lead alloys are used in batteries and bearing alloys. Calcium granules are also used to remove traces of air from vacuum devices. Soluble calcium and magnesium salts cause overall water hardness. If they are present in water in small quantities, then the water is called soft. If the content of these salts is high, water is considered hard. Hardness is eliminated by boiling; to completely eliminate the water, it is sometimes distilled.

Metallothermy

Pure metallic calcium is widely used in metallothermy for the production of rare metals.

Alloying of alloys

Pure calcium is used to alloy lead used for the production of battery plates and maintenance-free starter lead-acid batteries with low self-discharge. Also, metallic calcium is used for the production of high-quality calcium babbits BKA.

Nuclear fusion

The 48Ca isotope is the most effective and commonly used material for the production of superheavy elements and the discovery of new elements on the periodic table. For example, in the case of using 48Ca ions to produce superheavy elements in accelerators, the nuclei of these elements are formed hundreds and thousands of times more efficiently than when using other “projectiles” (ions).

Application of calcium compounds

Calcium hydride. By heating calcium in a hydrogen atmosphere, CaH2 (calcium hydride) is obtained, which is used in metallurgy (metallothermy) and in the production of hydrogen in the field.

Optical and laser materials. Calcium fluoride (fluorite) is used in the form of single crystals in optics (astronomical objectives, lenses, prisms) and as a laser material. Calcium tungstate (scheelite) in the form of single crystals is used in laser technology and also as a scintillator.

Calcium carbide. Calcium carbide CaC2 is widely used for the production of acetylene and for the reduction of metals, as well as in the production of calcium cyanamide (by heating calcium carbide in nitrogen at 1200 °C, the reaction is exothermic, carried out in cyanamide furnaces).

Chemical current sources. Calcium, as well as its alloys with aluminum and magnesium, are used in backup thermal electric batteries as an anode (for example, calcium-chromate element). Calcium chromate is used in such batteries as a cathode. The peculiarity of such batteries is an extremely long shelf life (decades) in a suitable condition, the ability to operate in any conditions (space, high pressures), high specific energy in terms of weight and volume. Disadvantage: short lifespan. Such batteries are used where it is necessary to create colossal electrical power for a short period of time (ballistic missiles, some spacecraft, etc.).

Fireproof materials. Calcium oxide, both in free form and as part of ceramic mixtures, is used in the production of refractory materials.

Medicines. In medicine, Ca drugs eliminate disorders associated with a lack of Ca ions in the body (tetany, spasmophilia, rickets). Ca preparations reduce hypersensitivity to allergens and are used to treat allergic diseases (serum sickness, sleepy fever, etc.). Ca preparations reduce increased vascular permeability and have an anti-inflammatory effect. They are used for hemorrhagic vasculitis, radiation sickness, inflammatory processes (pneumonia, pleurisy, etc.) and some skin diseases. Prescribed as a hemostatic agent, to improve the activity of the heart muscle and enhance the effect of digitalis preparations, as an antidote for poisoning with magnesium salts. Together with other drugs, Ca preparations are used to stimulate labor. Ca chloride is administered orally and intravenously.

Ca preparations also include gypsum (CaSO4), used in surgery for plaster casts, and chalk (CaCO3), prescribed internally for increased acidity of gastric juice and for preparing tooth powder.

Biological role

Calcium is a common macronutrient in the body of plants, animals and humans. In humans and other vertebrates, most of it is contained in the skeleton and teeth in the form of phosphates. The skeletons of most groups of invertebrates (sponges, coral polyps, mollusks, etc.) consist of various forms of calcium carbonate (lime). Calcium ions are involved in blood clotting processes, as well as in ensuring constant osmotic pressure of the blood. Calcium ions also serve as one of the universal second messengers and regulate a variety of intracellular processes - muscle contraction, exocytosis, including the secretion of hormones and neurotransmitters, etc. The calcium concentration in the cytoplasm of human cells is about 10−7 mol, in intercellular fluids about 10− 3 mol.

Most of the calcium that enters the human body with food is found in dairy products; the remaining calcium comes from meat, fish, and some plant products (legumes are especially high in legumes). Absorption occurs in both the large and small intestines and is facilitated by an acidic environment, vitamin D and vitamin C, lactose, and unsaturated fatty acids. The role of magnesium in calcium metabolism is important; with its deficiency, calcium is “washed out” from the bones and deposited in the kidneys (kidney stones) and muscles.

Due to the large number of processes associated with it, the calcium content in the blood is precisely regulated, and with proper nutrition, a deficiency does not occur. Prolonged absence from the diet can cause cramps, joint pain, drowsiness, growth defects, and constipation. Deeper deficiency leads to constant muscle cramps and osteoporosis. Abuse of coffee and alcohol can cause calcium deficiency, since some of it is excreted in the urine.

Products Calcium, mg/100 g

Sesame 783

Nettle 713

Large plantain 412

Sardines in oil 330

Ivy budra 289

Dog rose 257

Almond 252

Plantain lanceolist. 248

Hazelnut 226

Watercress 214

Soybeans dry 201

Children under 3 years old - 600 mg.

Children from 4 to 10 years old - 800 mg.

Children from 10 to 13 years old - 1000 mg.

Adolescents from 13 to 16 years old - 1200 mg.

Youth 16 and older - 1000 mg.

Adults from 25 to 50 years old - from 800 to 1200 mg.

Pregnant and breastfeeding women - from 1500 to 2000 mg.

Conclusion

Calcium is one of the most abundant elements on Earth. There is a lot of it in nature: mountain ranges and clay rocks are formed from calcium salts, it is found in sea and river water, and is part of plant and animal organisms.

Calcium constantly surrounds city dwellers: almost all main building materials - concrete, glass, brick, cement, lime - contain this element in significant quantities.

Naturally, having such chemical properties, calcium cannot exist in a free state in nature. But calcium compounds - both natural and artificial - have acquired paramount importance.

References

1.Editorial Board: Knunyants I. L. (chief editor) Chemical encyclopedia: in 5 volumes - Moscow: Soviet Encyclopedia, 1990. - T. 2. - P. 293. - 671 pp.

2.Doronin. N.A. Calcium, Goskhimizdat, 1962. 191 pp. with illustrations.

.Dotsenko V.A. - Therapeutic and preventive nutrition. - Question. nutrition, 2001 - N1-p.21-25

4.Bilezikian J. P. Calcium and bone metabolism // In: K. L. Becker, ed.

5.M.H. Karapetyants, S.I. Drakin - General and inorganic chemistry, 2000. 592 pp. with illustrations.

Calcium is very common in nature in the form of various compounds. In the earth's crust, it ranks fifth, accounting for 3.25%, and is most often found in the form of limestone CaCO3, dolomite CaCO3*MgCO3, gypsum CaSO4*2H2O, phosphorite Ca3(PO4)2 and fluorspar CaF2, not counting a significant proportion of calcium in composition of silicate rocks. Seawater contains an average of 0.04% (wt) calcium

Physical and chemical properties calcium

Calcium is in the subgroup of alkaline earth metals of group II of the periodic table of elements; serial number 20, atomic weight 40.08, valency 2, atomic volume 25.9. Calcium isotopes: 40 (97%), 42 (0.64%), 43 (0.15%), 44 (2.06%), 46 (0,003%), 48 (0.185%). Electronic structure of the calcium atom: 1s2, 2s2p6, 3s2p6, 4s2. The atomic radius is 1.97 A, the ion radius is 1.06 A. Up to 300°, calcium crystals have the shape of a cube with centered faces and a side size of 5.53 A, above 450° they have a hexagonal shape. The specific gravity of calcium is 1.542, melting point 851°, boiling point 1487°, heat of fusion 2.23 kcal/mol, heat of vaporization 36.58 kcal/mol. Atomic heat capacity of solid calcium Cр = 5.24 + 3.50*10В-3 T for 298-673° K and Cp = 6.29+1.40*10В-3T for 673-1124° K; for liquid calcium Cp = 7.63. The entropy of solid calcium is 9.95 ± 1, gaseous at 25° 37.00 ± 0.01.
The vapor elasticity of solid calcium was studied by Yu.A. Priselkov and A.N. Nesmeyanov, P. Douglas and D. Tomlin. The values of the saturated vapor pressure of calcium are given in table. 1.

In terms of thermal conductivity, calcium approaches sodium and potassium, at temperatures of 20-100° the coefficient of linear expansion is 25 * 10v-6, at 20° the electrical resistivity is 3.43 μ ohm/cm3, from 0 to 100° the temperature coefficient of electrical resistance is 0.0036. Electrochemical equivalent 0.74745 g/a*h. Calcium tensile strength 4.4 kg/mm2, Brinell hardness 13, elongation 53%, relative contraction 62%.
Calcium has a silvery-white color and shines when broken. In air, the metal is covered with a thin bluish-gray film of nitride, oxide and partially calcium peroxide. Calcium is flexible and malleable; it can be processed on a lathe, drilled, cut, sawed, pressed, drawn, etc. The purer the metal, the greater its ductility.
In the voltage series, calcium is among the most electronegative metals, which explains its high chemical activity. At room temperature, calcium does not react with dry air, at 300° and above it intensively oxidizes, and with strong heating it burns with a bright orange-reddish flame. In humid air, calcium gradually oxidizes, turning into hydroxide; It reacts relatively slowly with cold water, but vigorously displaces hydrogen from hot water, forming hydroxide.
Nitrogen reacts with calcium noticeably at a temperature of 300° and very intensively at 900° to form the nitride Ca3N2. With hydrogen at a temperature of 400°, calcium forms the hydride CaH2. Calcium does not bind to dry halogens, with the exception of fluorine, at room temperature; intensive formation of halides occurs at 400° and above.
Strong sulfuric (65-60° Be) and nitric acids have a weak effect on pure calcium. Of aqueous solutions of mineral acids, hydrochloric acid is very strong, nitric acid is strong, and sulfuric acid is weak. In concentrated NaOH solutions and soda solutions, calcium is almost not destroyed.

Application

Calcium is increasingly used in various industries. Recently he acquired great value as a reducing agent in the production of a number of metals. Pure uranium metal is obtained by reducing uranium fluoride with calcium metal. Calcium or its hydrides can be used to reduce titanium oxides, as well as oxides of zirconium, thorium, tantalum, niobium and other rare metals. Calcium is a good deoxidizer and degasser in the production of copper, nickel, chromium-nickel alloys, special steels, nickel and tin bronzes; it removes sulfur, phosphorus, and carbon from metals and alloys.
Calcium forms refractory compounds with bismuth, so it is used to purify lead from bismuth.
Calcium is added to various light alloys. It helps improve the ingot surface, fine grain size and reduce oxidation. Bearing alloys containing calcium are widely used. Lead alloys (0.04% Ca) can be used to make cable sheaths.
Calcium is used for the dehydration of alcohols and solvents for the desulfurization of petroleum products. Alloys of calcium with zinc or with zinc and magnesium (70% Ca) are used to produce high-quality porous concrete. Calcium is part of antifriction alloys (lead-calcium babbit).
Due to the ability to bind oxygen and nitrogen, calcium or calcium alloys with sodium and other metals are used for the purification of noble gases and as a getter in vacuum radio equipment. Calcium is also used to produce hydride, which is a source of hydrogen in the field. With carbon, calcium forms calcium carbide CaC2, which is used in large quantities to produce acetylene C2H2.

History of development

Dewi first obtained calcium in the form of an amalgam in 1808, using the electrolysis of wet lime with a mercury cathode. In 1852, Bunsen obtained an amalgam with a high calcium content by electrolysis of a hydrochloric acid solution of calcium chloride. In 1855, Bunsen and Matthiessen obtained pure calcium by electrolysis of CaCl2 and Moissan by electrolysis of CaF2. In 1893, Borchers significantly improved the electrolysis of calcium chloride by using cathode cooling; Arndt in 1902 obtained by electrolysis a metal containing 91.3% Ca. Ruff and Plata used a mixture of CaCl2 and CaF2 to reduce the electrolysis temperature; Borchers and Stockham obtained a sponge at a temperature below the melting point of calcium.
The problem of electrolytic production of calcium was solved by Rathenau and Suter, proposing the method of electrolysis with a touch cathode, which soon became industrial. There have been many proposals and attempts to produce calcium alloys by electrolysis, especially on a liquid cathode. According to F.O. Banzel, calcium alloys can be obtained by electrolysis of CaF2 with the addition of salts or fluoroxides of other metals. Poulene and Melan prepared a Ca-Al alloy on a liquid aluminum cathode; Kügelgen and Seward obtained a Ca-Zn alloy on a zinc cathode. The production of Ca-Zn alloys was studied in 1913 by W. Moldenhauer and J. Andersen, and they also prepared Pb-Ca alloys on a lead cathode. Koba, Simkins and Gire used a 2000 A lead cathode cell and produced an alloy with 2% Ca at a current efficiency of 20%. I. Tselikov and V. Wasinger added NaCl to the electrolyte to obtain an alloy with sodium; R.R. Syromyatnikov mixed the alloy and achieved 40-68% current output. Calcium alloys with lead, zinc and copper are produced by electrolysis on an industrial scale
The thermal method of producing calcium has attracted considerable interest. Aluminothermic reduction of oxides was discovered in 1865 by H.H. Beketov. In 1877, Malet discovered the interaction of a mixture of calcium, barium and strontium oxides with aluminum when heated. Winkler tried to reduce the same oxides with magnesium; Biltz and Wagner, reducing calcium oxide with aluminum in a vacuum, obtained a low yield of metal. Gunz in 1929 achieved better results. A.I. Voinitsky in 1938 reduced calcium oxide in the laboratory with aluminum and silicon alloys. The method was patented in 1938. At the end of the Second World War, the thermal method received industrial application.
In 1859, Caron proposed a method for producing sodium alloys with alkaline earth metals by the action of metallic sodium on their chlorides. Using this method, calcium (and barin) is obtained in an alloy with lead. Before the Second World War, the industrial production of calcium by electrolysis was carried out in Germany and Fraction. In Bieterfeld (Germany), in the period from 1934 to 1939, 5-10 tons of calcium were produced annually. The US need for calcium was covered by imports, which amounted to 10-25 g per year in the period 1920-1940. Since 1940, when imports from France ceased, the United States began to produce calcium itself in significant quantities by electrolysis; at the end of the war they began to obtain calcium using the vacuum-thermal method; according to S. Loomis, its output reached 4.5 tons per day. According to Minerale Yarbook, Dominium Magnesium in Canada produced calcium per year:

Information on the scale of calcium release over recent years none.

17.12.2019

The Far Cry series continues to delight its players with stability. After so much time, it becomes clear what you need to do in this game. Hunting, survival, capture...

16.12.2019

When creating a residential design, special attention should be given to the interior of the living room - it will become the center of your “universe”....

15.12.2019

It is impossible to imagine building a house without the use of scaffolding. Such structures are also used in other areas of economic activity. WITH...

14.12.2019

Welding appeared as a method of permanently joining metal products a little more than a century ago. At the same time, it is impossible to overestimate its importance at the moment. IN...

14.12.2019

Optimizing the surrounding space is extremely important for both small and large warehouses. This greatly simplifies the work and provides...

13.12.2019

Metal tiles are metal roofing materials. The surface of the sheets is coated with polymer materials and zinc. Natural tiles are imitated by the material...

13.12.2019

Testing equipment has been widely used in various fields. Its quality must be impeccable. To achieve this goal, the devices are equipped...

Calcium(calcium), ca, chemical element Group II of the Mendeleev periodic system, atomic number 20, atomic mass 40.08; silver-white light metal. The natural element is a mixture of six stable isotopes: 40 ca, 42 ca, 43 ca, 44 ca, 46 ca and 48 ca, of which 40 ca is the most common (96.97%).

Ca compounds - limestone, marble, gypsum (as well as lime - a product of calcination of limestone) were already used in construction in ancient times. Until the end of the 18th century. chemists considered lime to be a simple solid. In 1789 A. Lavoisier suggested that lime, magnesia, barite, alumina and silica are complex substances. In 1808 Davy By subjecting a mixture of wet slaked lime with mercury oxide to electrolysis with a mercury cathode, he prepared an amalgam ca, and by distilling mercury from it, he obtained a metal called “calcium” (from the Latin calx, genitive calcis - lime).

Distribution in nature. In terms of prevalence in the earth's crust, ca ranks 5th (after O, si, al and fe); content 2.96% by weight. It migrates vigorously and accumulates in various geochemical systems, forming 385 minerals (4th place in the number of minerals). There is little ca in the Earth's mantle and probably even less in the earth's core (in iron meteorites 0.02%). ca predominates in the lower part of the earth's crust, accumulating in the main rocks; most of the ca is contained in the feldspar anorthite ca; the content in basic rocks is 6.72%, in acidic rocks (granites, etc.) 1.58%. In the biosphere, an exceptionally sharp differentiation of ca occurs, associated mainly with “carbonate equilibrium”: when carbon dioxide interacts with carbonate caco 3, soluble bicarbonate Ca (HCO 3) 2 is formed:

CaCO 3 + h 2 o + co 2<=>Ca (HCO 3) 2<=>ca 2+ + 2hco 3 -.

This reaction is reversible and is the basis for the redistribution of ca. With a high co2 content in waters, ca is in solution, and with a low co2 content, the mineral calcite CaCO3 precipitates, forming thick deposits of limestone, chalk, and marble.

Biogenic migration also plays a huge role in the history of ca. In living matter of the elements - metals - ca is the main one. Organisms are known that contain more than 10% ca (more carbon), building their skeleton from ca compounds, mainly from CaCO 3 (calcareous algae, many mollusks, echinoderms, corals, rhizomes, etc.). The burial of the skeletons of marine animals and plants is associated with the accumulation of colossal masses of algae, coral and other limestones, which, plunging into the depths of the earth and mineralizing, turn into various types of marble.

Vast areas with a humid climate (forest zones, tundra) are characterized by a deficiency of ca - here it is easily leached from the soil. This is associated with low soil fertility, low productivity of domestic animals, their small size, and often skeletal diseases. Therefore, liming of soils, feeding of domestic animals and birds, etc. are of great importance. On the contrary, in dry climates CaCO 3 is difficult to dissolve, so the landscapes of steppes and deserts are rich in ca. In salt marshes and salt lakes it often accumulates gypsum caso 4 · 2h 2 o.

Rivers bring a lot of ca to the ocean, but it does not linger in ocean water (average content 0.04%), but is concentrated in the skeletons of organisms and, after their death, is deposited to the bottom mainly in the form of CaCO 3. Calcareous silts are widespread at the bottom of all oceans at depths of no more than 4000 m(At great depths, CaCO 3 dissolves; organisms there often experience a deficiency of Ca).

Groundwater plays an important role in ca migration. In limestone massifs, in places they vigorously leach CaCO 3, which is associated with the development karst, formation of caves, stalactites and stalagmites. In addition to calcite, in the seas of past geological eras there was widespread deposition of phosphates ca (for example, the Karatau phosphorite deposits in Kazakhstan), dolomite CaCO 3 · mgco 3, and in lagoons during evaporation - gypsum.

Over the course of geological history, biogenic carbonate formation increased and chemical precipitation of calcite decreased. In the Precambrian seas (over 600 million years ago) there were no animals with calcareous skeletons; they became widespread since the Cambrian (corals, sponges, etc.). This is associated with the high content of co 2 in the Precambrian atmosphere.

Physical and chemical properties. Crystal lattice of a-shape ca (stable at ordinary temperature), face-centered cubic A= 5.56 å. Atomic radius 1.97 å, ionic radius ca 2+, 1.04 å. Density 1.54 g/cm 3(20 °C). Above 464 °C, the hexagonal b-form is stable. t pl 851°c, t kip 1482 ° c; temperature coefficient of linear expansion 22? 10 -6 (0-300 ° c); thermal conductivity at 20 °C 125.6 W/(m? K) or 0.3 cal/(cm? sec° C); specific heat capacity (0-100 °C) 623.9 j/(kg? TO) or 0.149 cal/(G? °c); electrical resistivity at 20°c 4.6? 10 -8 ohm? m or 4.6? 10 -6 ohm? cm; temperature coefficient of electrical resistance 4.57? 10 -3 (20 °c). Modulus of elasticity 26 Gn/m 2 (2600 kgf/mm 2); tensile strength 60 Mn/m 2 (6 kgf/mm 2); elastic limit 4 Mn/m 2 (0,4 kgf/mm 2), yield strength 38 Mn/m 2 (3,8 kgf/mm 2); relative elongation 50%; Brinell hardness 200-300 Mn/m 2 (20-30 kgf/mm 2). Coatings of sufficiently high purity are plastic, easily pressed, rolled, and amenable to cutting.

The configuration of the outer electron shell of the atom is ca 4s 2, according to which ca in compounds is 2-valent. Chemically ca is very active. At normal temperatures, ca easily interacts with oxygen and moisture in the air, so it is stored in hermetically sealed containers or under mineral oil. When heated in air or oxygen, it ignites, giving the basic oxide cao. The peroxides ca - cao 2 and CaO 4 are also known. Ca reacts quickly with cold water at first, then the reaction slows down due to the formation of a film of ca (oh) 2. ca reacts vigorously with hot water and acids, releasing h 2 (except for concentrated hno 3). It reacts with fluorine in the cold, and with chlorine and bromine - above 400 ° C, giving caf 2, cacl 2 and cabr 2, respectively. These halides in the molten state form so-called subcompounds with ca - caf, caci, in which ca is formally monovalent. When heated, ca with sulfur turns out calcium sulfide cas, the latter adds sulfur, forming polysulfides (cas 2, cas 4, etc.). Interacting with dry hydrogen at 300-400 °C, ca forms the hydride cah 2 - an ionic compound in which hydrogen is an anion. At 500 °C ca and nitrogen give nitride ca 3 n 2; the interaction of ca with ammonia in the cold leads to complex ammonia ca 6. When heated without air access with graphite, silicon or phosphorus, ca gives respectively calcium carbide cac 2, silicides casi 2 and phosphide ca 3 p 2. ca forms intermetallic compounds with al, ag, au, cu, li, mg, pb, sn, etc.

Receipt and application. In industry, ca is obtained in two ways: 1) by heating a briquetted mixture of cao and al powder at 1200 ° C in a vacuum of 0.01-0.02 mmHg st.; released by the reaction: 6cao +2al = 3 CaO? l 2 o 3 + 3Ca vapors condense on a cold surface; 2) by electrolysis of the cacl 2 and kcl melt with a liquid copper-calcium cathode, an alloy cu - ca (65% ca) is prepared, from which ca is distilled off at a temperature of 950-1000 ° C in a vacuum of 0.1-0.001 mmHg st.

In the form of a pure metal, ca is used as a reducing agent for u, th, cr, v, zr, cs, rb and some rare earth metals from their compounds. It is also used for deoxidation of steels, bronzes and other alloys, for removing sulfur from petroleum products, for dehydrating organic liquids, for purifying argon from nitrogen impurities and as a gas absorber in electric vacuum devices. Widely used in technology antifriction materials pb-na-ca systems, as well as pb-ca alloys used for the manufacture of electrical cable sheaths. The ca-si-ca alloy (silicocalcium) is used as a deoxidizing agent and degassing agent in the production of high-quality steels. For information on the use of K compounds, see the relevant articles.

A. Ya. Fischer, A. I. Perelman.

Calcium in the body . ca - one of nutrients necessary for the normal course of life processes. It is present in all tissues and fluids of animals and plants. Only rare organisms can develop in an environment devoid of ca; in some organisms the ca content reaches 38%; in humans - 1.4-2%. Cells of plant and animal organisms require strictly defined ratios of ca 2+, na + and K + ions in extracellular environments. Plants obtain ca from the soil. According to their relation to ca, plants are divided into calciphiles And calcephobes. Animals obtain ca from food and water. ca is necessary for the formation of a number of cellular structures, maintaining normal permeability of outer cell membranes, for fertilization of eggs of fish and other animals, and activation of a number of enzymes. Ca 2+ ions transmit excitation to the muscle fiber, causing it to contract, increase the force of heart contractions, increase the phagocytic function of leukocytes, activate the system of protective blood proteins, and participate in its coagulation. In cells, almost all ca is found in the form of compounds with proteins, nucleic acids, phospholipids, in complexes with inorganic phosphates and organic acids. In the blood plasma of humans and higher animals, only 20-40% of ca can be associated with proteins. In animals with a skeleton, up to 97-99% of all ca is used as a building material: in invertebrates mainly in the form of caco 3 (mollusk shells, corals), in vertebrates - in the form of phosphates. Many invertebrates store ca before molting to build a new skeleton or to ensure vital functions in unfavorable conditions.

The content of ca in the blood of humans and higher animals is regulated by hormones of the parathyroid and thyroid glands. Vitamin D plays a critical role in these processes. Absorption of ca occurs in the anterior part of the small intestine. The absorption of ca deteriorates with a decrease in acidity in the intestines and depends on the ratio of ca, P and fat in food. The optimal ca/p ratio in cow's milk is about 1.3 (in potatoes 0.15, in beans 0.13, in meat 0.016). With an excess of P or oxalic acid in food, the absorption of ca worsens, Bile acids accelerate its absorption. Optimal Ca/fat ratios in human food are 0.04-0.08 G ca by 1 G fat Ca excretion occurs mainly through the intestines. Mammals during the period lactation lose a lot of ca with milk. With disturbances of phosphorus-calcium metabolism in young animals and children, rickets, in adult animals - changes in the composition and structure of the skeleton ( osteomalacia).

I. A. Skulsky.

In medicine, the use of ca drugs eliminates disorders associated with a lack of ca 2+ ions in the body (tetany, spasmophilia, rickets). CA preparations reduce hypersensitivity to allergens and are used to treat allergic diseases (serum sickness, urticaria, angioedema, hay fever, etc.). CA preparations reduce increased vascular permeability and have an anti-inflammatory effect. They are used for hemorrhagic vasculitis, radiation sickness, inflammatory and exudative processes (pneumonia, pleurisy, endometritis, etc.) and some skin diseases. Prescribed as hemostatic agents, to improve the activity of the heart muscle and enhance the effect of digitalis preparations; as weak diuretics and as antidotes for poisoning with magnesium salts. Together with other drugs, ca drugs are used to stimulate labor. Calcium chloride is administered orally and intravenously. Ossocalcinol (15% sterile suspension of specially prepared bone powder in peach oil) has been proposed for tissue therapy. Ca preparations also include gypsum (caso 4), used in surgery for plaster bandages, and chalk (CaCO 3), prescribed internally for increased acidity of gastric juice and for the preparation of tooth powder.

Lit.: Brief chemical encyclopedia, vol. 2, M., 1963, p. 370-75; Rodyakin V.V., Calcium, its compounds and alloys, M., 1967; Kaplansky S. Ya., Mineral exchange, M. - L., 1938; Vishnyakov S.I., Metabolism of macroelements in farm animals, M., 1967.