Calcium metal. Calcium in nature (3.4% in the Earth's crust)

History of calcium

Calcium was discovered in 1808 by Humphry Davy, who, by electrolysis of slaked lime and mercuric oxide, obtained calcium amalgam, as a result of the process of distilling mercury from which the metal remained, called calcium. In Latin lime sounds like calx, it was this name that was chosen by the English chemist for the discovered substance.

Calcium is an element of the main subgroup II of group IV of the periodic table of chemical elements D.I. Mendeleev, has an atomic number of 20 and an atomic mass of 40.08. The accepted designation is Ca (from the Latin - Calcium).

Physical and chemical properties

Calcium is a reactive soft alkali metal, silver- white. Due to interaction with oxygen and carbon dioxide the surface of the metal becomes dull, so calcium needs a special storage regime - a tightly closed container, in which the metal is poured with a layer of liquid paraffin or kerosene.

Calcium is the most well-known of the microelements necessary for humans; the daily requirement for it ranges from 700 to 1500 mg for a healthy adult, but it increases during pregnancy and lactation; this must be taken into account and calcium must be obtained in the form of preparations.

Being in nature

Calcium has very high chemical activity, therefore it is not found in nature in its free (pure) form. However, it is the fifth most common in the earth's crust; it is found in the form of compounds in sedimentary (limestone, chalk) and rocks (granite); feldspar anorite contains a lot of calcium.

It is quite widespread in living organisms; its presence has been found in plants, animals and humans, where it is present mainly in teeth and bone tissue.

Calcium absorption

An obstacle to the normal absorption of calcium from foods is the consumption of carbohydrates in the form of sweets and alkalis, which neutralize hydrochloric acid stomach, necessary for dissolving calcium. The process of calcium absorption is quite complex, so sometimes it is not enough to get it only from food, it is necessary additional reception microelement.

Interaction with others

To improve the absorption of calcium in the intestine, it is necessary, which tends to facilitate the process of calcium absorption. When taking calcium (in the form of supplements) while eating, absorption is blocked, but taking calcium supplements separately from food does not affect this process in any way.

Almost all of the body's calcium (1 to 1.5 kg) is found in bones and teeth. Calcium is involved in excitability processes nerve tissue, muscle contractility, blood clotting processes, is part of the nucleus and membranes of cells, cellular and tissue fluids, has anti-allergic and anti-inflammatory effects, prevents acidosis, activates a number of enzymes and hormones. Calcium is also involved in the regulation of cell membrane permeability and has the opposite effect.

Signs of calcium deficiency

Signs of calcium deficiency in the body are the following, at first glance, unrelated symptoms:

• nervousness, worsening mood;
• cardiopalmus;
• convulsions, numbness of extremities;
• slowing of growth and children;
• high blood pressure;
• splitting and brittleness of nails;
• joint pain, lowering the “pain threshold”;
• heavy menstruation.

Causes of calcium deficiency

Causes of calcium deficiency may include unbalanced diets(especially fasting), low calcium content in food, smoking and addiction to coffee and caffeine-containing drinks, dysbacteriosis, kidney disease, thyroid disease, pregnancy, lactation and menopause.

Excess calcium, which can occur with excessive consumption of dairy products or uncontrolled use of drugs, is characterized by extreme thirst, nausea, vomiting, loss of appetite, weakness and increased urination.

Uses of calcium in life

Calcium has found application in the metallothermic production of uranium, in the form of natural compounds it is used as a raw material for the production of gypsum and cement, as a means of disinfection (well-known bleach).

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% total number. Nuclear reactions Other isotopes of calcium were also obtained, for example radioactive 45 Ca.
To obtain a simple calcium substance, electrolysis of molten calcium salts or aluminothermy is used:
4CaO + 2Al = Ca(AlO 2) 2 + 3Ca

Physical properties:

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

Chemical properties:

Calcium is not as active as alkali metals, however, it must be stored under a layer 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 pure form white or colorless substance. 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. Slightly soluble in water, it can dissolve in excess sulfuric acid, forming hydrogen sulfate.

Calcium phosphate Ca 3 (PO 4) 2- (“phosphorite”), insoluble, under the influence of strong acids it turns 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 during the production of metals, which explains its use as a flux.

Calcium chloride CaCl 2- colorless christ. It is well soluble in water. Forms crystal 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. Component pyrotechnic compositions that give 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 ("calcethermy"): 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 calcium isotope, 45Ca, is used as a radioactive tracer in research biological role calcium and its migration in the environment.

The main area of ​​application of numerous calcium compounds is the production of building materials (cement, building mixtures, drywall, 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, and in the elderly - 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 - the deposition of its compounds in internal organs, 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).

Calcium

CALCIUM-I; m.[from lat. calx (calcis) - lime] Chemical element (Ca), a silver-white metal that is part of limestone, marble, etc.

◁ Calcium, oh, oh. K salts.

(lat. Calcium), a chemical element of group II of the periodic table, belongs to the alkaline earth metals. Name from lat. calx, genitive calcis - lime. Silver-white metal, density 1.54 g/cm 3, t pl 842ºC. At ordinary temperatures it is easily oxidized in air. In terms of prevalence in the earth's crust, it ranks 5th (minerals calcite, gypsum, fluorite, etc.). As an active reducing agent, it is used to obtain U, Th, V, Cr, Zn, Be and other metals from their compounds, to deoxidize steels, bronzes, etc. It is part of antifriction materials. Calcium compounds are used in construction (lime, cement), calcium preparations are used in medicine.

CALCIUM (lat. Calcium), Ca (read “calcium”), chemical element with atomic number 20, located in the fourth period in group IIA of Mendeleev’s periodic system of elements; atomic mass 40.08. Belongs to the alkaline earth elements (cm. ALKALINE EARTH METALS).
Natural calcium consists of a mixture of nuclides (cm. NUCLIDE) with mass numbers of 40 (in a mixture by mass of 96.94%), 44 (2.09%), 42 (0.667%), 48 (0.187%), 43 (0.135%) and 46 (0.003%). Outer electron layer 4 configuration s 2 . In almost all compounds the oxidation state of calcium is +2 (valence II).
The radius of the neutral calcium atom is 0.1974 nm, the radius of the Ca 2+ ion is from 0.114 nm (for coordination number 6) to 0.148 nm (for coordination number 12). The energies of sequential ionization of a neutral calcium atom are, respectively, 6.133, 11.872, 50.91, 67.27 and 84.5 eV. According to the Pauling scale, the electronegativity of calcium is about 1.0. In its free form, calcium is a silvery-white metal.
History of discovery
Calcium compounds are found everywhere in nature, so humanity has been familiar with them since ancient times. Lime has long been used in construction (cm. LIME)(quick and quenched) which for a long time believed simple substance, "earth". However, in 1808 the English scientist G. Davy (cm. DAVY Humphrey) managed to obtain a new metal from lime. To do this, Davy subjected to electrolysis a mixture of slightly moistened slaked lime with mercury oxide and isolated a new metal from the amalgam formed on the mercury cathode, which he called calcium (from the Latin calx, genus calcis - lime). In Russia for some time this metal was called “liming”.
Being in nature
Calcium is one of the most common elements on Earth. It accounts for 3.38% of the mass of the earth's crust (5th most abundant after oxygen, silicon, aluminum and iron). Due to its high chemical activity, calcium does not occur in free form in nature. Most calcium is found in silicates (cm. SILICATES) and aluminosilicates (cm. ALUMINUM SILICATES) various rocks (granites (cm. GRANITE), gneisses (cm. GNEISS) and so on.). In the form of sedimentary rocks, calcium compounds are represented by chalk and limestones, consisting mainly of the mineral calcite (cm. CALCITE)(CaCO 3). The crystalline form of calcite - marble - is much less common in nature.
Calcium minerals such as limestone are quite common (cm. LIMESTONE) CaCO3, anhydrite (cm. ANHYDRITE) CaSO 4 and gypsum (cm. GYPSUM) CaSO 4 2H 2 O, fluorite (cm. FLUORITE) CaF 2, apatites (cm. APATITE) Ca 5 (PO 4) 3 (F,Cl,OH), dolomite (cm. DOLOMITE) MgCO 3 ·CaCO 3 . The presence of calcium and magnesium salts in natural water its rigidity is determined (cm. HARDNESS OF WATER). 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 from calcium carbonate CaCO 3.
Receipt
Metallic calcium is obtained by electrolysis of a melt consisting of CaCl 2 (75-80%) and KCl or from CaCl 2 and CaF 2, as well as aluminothermic reduction of CaO at 1170-1200 °C:
4CaO + 2Al = CaAl 2 O 4 + 3Ca.
Physical and chemical properties
Calcium metal exists in two allotropic modifications (see Allotropy (cm. ALLOTROPY)). Up to 443 °C, a-Ca with a cubic face-centered lattice (parameter a = 0.558 nm) is stable; b-Ca with a cubic body-centered lattice of the a-Fe type (parameter a = 0.448 nm) is more stable. Melting point of calcium is 839 °C, boiling point is 1484 °C, density is 1.55 g/cm3.
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.
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:
Ca + 2H 2 O = Ca(OH) 2 + H 2.
Calcium reacts with active non-metals (oxygen, chlorine, bromine) under normal conditions:
2Ca + O 2 = 2CaO; Ca + Br 2 = CaBr 2.
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 + H 2 = CaH 2 (calcium hydride),
Ca + 6B = CaB 6 (calcium boride),
3Ca + N 2 = Ca 3 N 2 (calcium nitride)
Ca + 2C = CaC 2 (calcium carbide)
3Ca + 2P = Ca 3 P 2 (calcium phosphide), calcium phosphides of the compositions CaP and CaP 5 are also known;
2Ca + Si = Ca 2 Si (calcium silicide); calcium silicides of the compositions CaSi, Ca 3 Si 4 and CaSi 2 are also known.
The occurrence of the above reactions is usually accompanied by the release of large quantity heat (i.e. 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:
CaH 2 + 2H 2 O = Ca(OH) 2 + 2H 2,
Ca 3 N 2 + 3H 2 O = 3Ca(OH) 2 + 2NH 3.
Calcium oxide is typically basic. In the laboratory and technology it is obtained by thermal decomposition of carbonates:
CaCO 3 = CaO + CO 2.
Technical calcium oxide CaO is called quicklime.
It reacts with water to form Ca(OH) 2 and release a large amount of heat:
CaO + H 2 O = Ca(OH) 2.
Ca(OH)2 obtained in this way is usually called slaked lime or milk of lime (cm. LIME MILK) due to the fact that the solubility of calcium hydroxide in water is low (0.02 mol/l at 20°C), and when it is added to water, a white suspension is formed.
When interacting with acidic oxides, CaO forms salts, for example:
CaO + CO 2 = CaCO 3; CaO + SO 3 = CaSO 4.
The Ca 2+ ion is colorless. When calcium salts are added to the flame, the flame turns brick-red.
Calcium salts such as CaCl 2 chloride, CaBr 2 bromide, CaI 2 iodide and Ca(NO 3) 2 nitrate are highly soluble in water. Insoluble in water are fluoride CaF 2, carbonate CaCO 3, sulfate CaSO 4, average orthophosphate Ca 3 (PO 4) 2, oxalate CaC 2 O 4 and some others.
It is important that, unlike the average calcium carbonate CaCO 3, acidic calcium carbonate (bicarbonate) Ca(HCO 3) 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 falls on limestone, their dissolution is observed:
CaCO 3 + CO 2 + H 2 O = Ca(HCO 3) 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(HCO 3) 2 = CaCO 3 + CO 2 + H 2 O.
This is how large masses of substances are transferred in nature. As a result, huge holes can form underground (see Karst (cm. KARST (natural phenomenon))), and beautiful stone “icicles” - stalactites - form in the caves (cm. STALACTITES (mineral formations)) and stalagmites (cm. STALAGMITES).
The presence of dissolved calcium bicarbonate in water largely determines the temporary hardness of water. (cm. 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 of calcium and its compounds
Calcium metal is used for metallothermic production of uranium (cm. URANIUM (chemical element)), thorium (cm. THORIUM), titanium (cm. TITANIUM (chemical element)), zirconium (cm. ZIRCONIUM), cesium (cm. CESIUM) and rubidium (cm. RUBIDIUM).
Natural calcium compounds are widely used in production binding materials(cement (cm. CEMENT), gypsum (cm. GYPSUM), lime, etc.). The binding effect of slaked lime is based on the fact that over time, calcium hydroxide reacts with carbon dioxide in the air. As a result of the ongoing reaction, needle-shaped crystals of calcite CaCO3 are formed, which grow into nearby stones, bricks, and other building materials and, as it were, weld them into a single whole. Crystalline calcium carbonate - marble - is an excellent finishing material. Chalk is used for whitewashing. Large quantities limestone is consumed in the production of cast iron, as it makes it possible to convert refractory impurities of iron ore (for example, quartz SiO 2) into relatively low-melting slag.
As disinfectant Bleach is very effective (cm. BLEACHING POWDER)- “bleach” Ca(OCl)Cl - mixed chloride and calcium hypochloride (cm. CALCIUM HYPOCHLORITE) with high oxidizing ability.
Calcium sulfate is also widely used, existing both in the form of an anhydrous compound and in the form of crystalline hydrates - the so-called “semi-aqueous” sulfate - alabaster (cm. ALEVIZ FRYAZIN (Milanese)) CaSO 4 ·0.5H 2 O and dihydrate sulfate - gypsum CaSO 4 ·2H 2 O. Gypsum is widely used in construction, in sculpture, for making stucco and various art products. Plaster is also used in medicine to fix bones during fractures.
Calcium chloride CaCl 2 is used along with table salt to combat icing road surfaces. Calcium fluoride CaF 2 is an excellent optical material.
Calcium in the body
Calcium is a biogenic element (cm. BIOGENIC ELEMENTS), constantly present in the tissues of plants and animals. An important component of the mineral metabolism of animals and humans and the mineral nutrition of plants, calcium performs various functions in the body. Composed of apatite (cm. APATITE), as well as sulfate and carbonate, calcium forms the mineral component of bone tissue. The human body weighing 70 kg contains about 1 kg of calcium. Calcium participates in the functioning of ion channels (cm. ION CHANNELS) transporting substances through biological membranes in the transmission of nerve impulses (cm. NERVOUS IMPULSE), in blood clotting processes (cm. BLOOD CLOTTING) and fertilization. Calciferols regulate calcium metabolism in the body (cm. CALCIFEROLS)(vitamin D). A lack or excess of calcium leads to various diseases- rickets (cm. RICKETS), calcinosis (cm. CALCINOSIS) etc. Therefore, human food should be required quantities contain calcium compounds (800-1500 mg of calcium per day). Calcium content is high in dairy products (such as cottage cheese, cheese, milk), some vegetables and other foods. Calcium preparations are widely used in medicine.

encyclopedic Dictionary. 2009 .

Calcium is a chemical element of group II with atomic number 20. periodic table, denoted by the symbol Ca (Latin Calcium). Calcium is a soft alkaline earth metal with a silvery-gray color.

Element 20 of the periodic table The name of the element comes from lat. calx (in genitive case calcis) - “lime”, “soft stone”. It was proposed by the English chemist Humphry Davy, who isolated calcium metal in 1808.
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.
Calcium is one of the most common elements on Earth. Calcium compounds are found in almost all animal and plant tissues. It accounts for 3.38% of the mass of the earth's crust (5th place in abundance after oxygen, silicon, aluminum and iron).

Finding calcium 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). The content of the element in sea water is 400 mg/l.

Isotopes

Calcium occurs in nature as a mixture of six isotopes: 40Ca, 42Ca, 43Ca, 44Ca, 46Ca and 48Ca, of which the most common, 40Ca, accounts for 96.97%. Calcium nuclei contain magic number protons: Z = 20. Isotopes
40
20
Ca20 and
48
20
Ca28 are two of the five nuclei that exist in nature with twice the magic number.
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%), undergoes double beta decay with a half-life of 1.6 1017 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).

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 found in the skeleton and teeth. Calcium is found in bones in the form of hydroxyapatite. From various forms Calcium carbonate (lime) constitutes the “skeletons” of most groups of invertebrates (sponges, coral polyps, mollusks, etc.). Calcium ions are involved in blood clotting processes, and also serve as one of the universal second messengers inside cells and regulate a variety of intracellular processes - muscle contraction, exocytosis, including the secretion of hormones and neurotransmitters. The calcium concentration in the cytoplasm of human cells is about 10−4 mmol/l, in intercellular fluids it is about 2.5 mmol/l.

Calcium requirements depend on age. For adults aged 19-50 years and children aged 4-8 years inclusive, the daily requirement (RDA) is 1000 mg (contained in approximately 790 ml of milk with 1% fat content), and for children aged 9 to 18 years inclusive - 1300 mg per day (contained in approximately 1030 ml of milk with a fat content of 1%). During adolescence, consuming enough calcium is very important due to the rapid growth of the skeleton. However, according to research in the United States, only 11% of girls and 31% of boys aged 12-19 years achieve their needs. In a balanced diet, most of the calcium (about 80%) enters the child’s body with dairy products. The remaining calcium comes from grains (including whole grain bread and buckwheat), legumes, oranges, greens, and nuts. In "dairy" products based on milk fat ( butter, cream, sour cream, cream-based ice cream) contain virtually no calcium. The more in dairy product milk fat, the less calcium it contains. Calcium absorption in the intestine occurs in two ways: transcellular (transcellular) and intercellular (paracellular). The first mechanism is mediated by the action active form vitamin D (calcitriol) and its intestinal receptors. It plays a big role in low to moderate calcium intake. With a higher calcium content in the diet, intercellular absorption begins to play a major role, which is associated with a large gradient of calcium concentration. Due to the transcellular mechanism, calcium is absorbed to a greater extent in the duodenum (due to the highest concentration of calcitriol receptors there). Due to intercellular passive transfer, calcium absorption is most active in all three sections of the small intestine. Paracellular absorption of calcium is promoted by lactose (milk sugar).

Calcium absorption is inhibited by some animal fats (including cow's milk fat and beef fat, but not lard) and palm oil. The palmitic and stearic fatty acids contained in such fats are split off during digestion in the intestines and, in their free form, firmly bind calcium, forming calcium palmitate and calcium stearate (insoluble soaps). In the form of this soap, both calcium and fat are lost in the stool. This mechanism is responsible for decreased calcium absorption, decreased bone mineralization, and decreased indirect measures of bone strength in infants using palm oil (palm olein) based infant formulas. In such children, the formation of calcium soaps in the intestines is associated with hardening of stools, a decrease in stool frequency, as well as more frequent regurgitation and colic.

Blood calcium concentration due to its importance for large number vital important processes precisely adjustable and proper nutrition and adequate consumption of low-fat dairy products and vitamin D deficiency does not occur. Long-term deficiency of calcium and/or vitamin D in the diet increases the risk of osteoporosis and causes rickets in infancy.

Excessive doses of calcium and vitamin D can cause hypercalcemia. The maximum safe dose for adults aged 19 to 50 years inclusive is 2500 mg per day (about 340 g of Edam cheese).

Calcium compounds- limestone, marble, gypsum (as well as lime - a product 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, subjecting a mixture of wet slaked lime and mercuric oxide to electrolysis with a mercury cathode, prepared calcium amalgam, and by distilling mercury from it, he obtained a metal called “calcium” (from the Latin. Calx, genus. case calcis - lime).

Placing electrons in orbitals.

+20Sa… |3s 3p 3d | 4s

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.

Crystal lattices of metals can be various types, however, calcium is characterized by a face-centered cubic lattice.

Dimensions, shape and mutual arrangement crystals in metals are emitted by metallographic methods. The most complete assessment of the structure of the metal in this regard is provided by microscopic analysis of its thin section. A sample is cut out of the metal being tested and its surface is ground, polished and etched with a special solution (etchant). As a result of etching, the structure of the sample is highlighted, which is examined or photographed using a metallographic microscope.

Kaltsy - light metal(d = 1.55), silvery-white. It is harder and melts at a higher temperature high 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. During chemical reactions, calcium valence electrons are transferred to atoms of other elements. In this case, doubly charged ions are formed.

Calcium has great chemical activity towards metals, especially oxygen. In air, it oxidizes more slowly than alkali metals, since the oxide film on it is less permeable to oxygen. When heated, calcium burns, releasing enormous amounts of heat:

Calcium reacts with water, displacing hydrogen from it and forming a base:

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

Due to its high chemical reactivity to oxygen, calcium finds some use in obtaining rare metals from their oxides. Metal oxides are heated together with calcium shavings; The reactions result in calcium oxide and metal. The use of calcium and some of its alloys for the so-called deoxidation of metals is based on this same property. Calcium is added to the molten metal and it removes traces of dissolved oxygen; the resulting calcium oxide floats to the surface of the metal. Calcium is included in some alloys.

Calcium is obtained by electrolysis of molten calcium chloride or by the aluminothermic method. Calcium oxide, or slaked lime, is a white powder, it melts at 2570 °C. It is obtained by calcining limestone:

CaCO3 = CaO + CO2^

Calcium oxide is a basic oxide, so it reacts with acids and acid anhydrides. With water it gives the base - calcium hydroxide:

CaO + H2O = Ca(OH)2

The addition of water to calcium oxide, called slaking of lime, occurs with the release of a large amount of heat. Some of the water turns into steam. Calcium hydroxide, or slaked lime, is a white substance, slightly soluble in water. An aqueous solution of calcium hydroxide is called lime water. This solution has fairly strong alkaline properties, since calcium hydroxide dissociates well:

Ca(OH)2 = Ca + 2OH

Compared to hydrates of alkali metal oxides, calcium hydroxide is a weaker base. This is explained by the fact that the calcium ion is doubly charged and attracts hydroxyl groups more strongly.

Slaked lime and its solution, called lime water, react with acids and acid anhydrides, including carbon dioxide. Lime water is used in laboratories for the discovery of carbon dioxide, since the resulting insoluble calcium carbonate causes cloudiness in the water:

Ca + 2OH + CO2 = CaCO3v + H2O

However, if carbon dioxide is passed in for a long time, the solution becomes clear again. This is explained by the fact that calcium carbonate is converted into a soluble salt - calcium bicarbonate:

CaCO3 + CO2 + H2O = Ca(HCO3)2

In industry, calcium is obtained in two ways:

By heating the briquetted mixture of CaO and Al powder at 1200 °C in a vacuum of 0.01 - 0.02 mm. Hg Art.; distinguished by reaction:

6CaO + 2Al = 3CaO Al2O3 + 3Ca

Calcium vapor condenses on a cold surface.

By electrolysis of a melt of CaCl2 and KCl with a liquid copper-calcium cathode, a Cu - Ca alloy (65% Ca) is prepared, from which calcium is distilled off at a temperature of 950 - 1000 ° C in a vacuum of 0.1 - 0.001 mmHg.

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

Calcium is one of the most common elements in nature. The earth's crust contains approximately 3% (wt.). Calcium salts form in nature large clusters in the form of carbonates (chalk, marble), sulfates (gypsum), phosphates (phosphorites). Under the influence of water and carbon dioxide, carbonates go into solution in the form of bicarbonates and are transported by groundwater and river water over long distances. When calcium salts are washed away, caves can form. Due to the evaporation of water or an increase in temperature, calcium carbonate deposits can form in a new location. For example, stalactites and stalagmites form in caves.

Soluble calcium and magnesium salts cause overall water hardness. If they are present in water in small quantities, then the water is called soft. With a high content of these salts (100 - 200 mg of calcium salts in 1 liter in terms of ions), the water is considered hard. In such water, soap does not foam well, since calcium and magnesium salts form insoluble compounds with it. Doesn't boil well in hard water food products, and when boiled it forms scale on the walls of steam boilers. Scale conducts heat poorly, causes increased fuel consumption and accelerates wear of the boiler walls. Scale formation is a complex process. When heated, acidic salts carbonic acid calcium and magnesium decompose and turn into insoluble carbonates:

Ca + 2HCO3 = H2O + CO2 + CaCO3v

The solubility of calcium sulfate CaSO4 also decreases when heated, so it is part of the scale.

Hardness caused by the presence of calcium and magnesium bicarbonates in water is called carbonate or temporary hardness, since it is eliminated by boiling. In addition to carbonate hardness, there is also non-carbonate hardness, which depends on the content of calcium and magnesium sulfates and chlorides in the water. These salts are not removed by boiling, and therefore non-carbonate hardness is also called permanent hardness. Carbonate and non-carbonate hardness add up to total hardness.

To completely eliminate hardness, water is sometimes distilled. To eliminate carbonate hardness, water is boiled. General hardness can be eliminated or by adding chemical substances, or using so-called cation exchangers. When using the chemical method, soluble calcium and magnesium salts are converted into insoluble carbonates, for example, milk of lime and soda are added:

Ca + 2HCO3 + Ca + 2OH = 2H2O + 2CaCO3v

Ca + SO4 + 2Na + CO3 = 2Na + SO4 + CaCO3v

Removing hardness using cation exchange resins is a more advanced process. Cation exchangers are complex substances (natural compounds of silicon and aluminum, high-molecular organic compounds), the composition of which can be expressed by the formula Na2R, where R is a complex acid residue. When filtering water through a layer of cation exchange resin, Na ions (cations) are exchanged for Ca and Mg ions:

Ca + Na2R = 2Na + CaR

Consequently, Ca ions pass from the solution into the cation exchanger, and Na ions pass from the cation exchanger into the solution. To restore the used cation exchanger, it is washed with a solution table salt. In this case, the reverse process occurs: Ca ions in the cation exchanger are replaced by Na ions:

2Na + 2Cl + CaR = Na2R + Ca + 2Cl

The regenerated cation exchanger can be used again for water purification.

In the form of pure metal, Ca is used as a reducing agent for U, Th, Cr, V, Zr, Cs, Rb and some rare earth metals and their connections. 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. Antifiction materials of the Pb - Na - Ca system, as well as Pb - Ca alloys used for the manufacture of shells, have been widely used in technology electrical cables. The alloy Ca - Si - Ca (silicocalcium) is used as a deoxidizer and degasser in the production of high-quality steels.

Calcium is one of the biogenic elements necessary for the normal functioning 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, Na and K ions in extracellular environments. Plants obtain Ca from the soil. Based on their relationship to Ca, plants are divided into calcephiles 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 ions transmit excitation to the muscle fiber, causing it to contract, increase the strength 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 bound to 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 CaCO3 (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 Ca content in the blood of humans and higher animals is regulated by hormones of the parathyroid and thyroid glands. Vitamin D plays a key role in these processes. Ca absorption occurs in the anterior section of the small intestine. The absorption of Ca deteriorates with a decrease in acidity in the intestine and depends on the ratio of Ca, phosphorus and fat in food. Optimal ratios Ca/P in cow's milk about 1.3 (in potatoes 0.15, in beans 0.13, in meat 0.016). If there is an excess of P and oxalic acid Ca absorption deteriorates. Bile acids accelerate its absorption. The optimal Ca/fat ratio in human food is 0.04 - 0.08 g. Ca per 1 g. fat Ca excretion occurs mainly through the intestines. Mammals lose a lot of Ca in milk during lactation. With disturbances in phosphorus-calcium metabolism, rickets develops in young animals and children, and changes in the composition and structure of the skeleton (osteomalacia) develop in adult animals.

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. Ossocalcinol (15% sterile suspension of specially prepared bone powder in peach oil) has been proposed for tissue therapy.

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