Where is calcium used? Calcium in nature (3.4% in the Earth's crust)

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 of various types, but calcium is characterized by a face-centered cubic lattice.

The sizes, shapes and relative positions of crystals in metals are emitted using 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.

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. Therefore, the chemical bond between the ions and the electron gas is stronger than that of 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 that 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. rt. Art.; distinguished by reaction:

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

Calcium vapor condenses on a cold surface.

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

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 large accumulations in nature in the form of carbonates (chalk, marble), sulfates (gypsum), and 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 place. 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. Hard water does not cook food well, 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 carbonic acid salts of 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 is eliminated either by adding chemicals 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 of 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 a pure metal, Ca is used as a reducing agent for U, Th, Cr, V, Zr, Cs, Rb and some rare earth metals and 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. Anti-fiction materials of the Pb - Na - Ca system, as well as Pb - Ca alloys used for the manufacture of electrical cable sheaths, have been widely used in technology. 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, and 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. 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 and oxalic acid in food, Ca absorption worsens. 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 bandages, and chalk (CaCO3), prescribed internally for increased acidity of gastric juice and for the preparation of tooth powder.

Home / Lectures 1st year / General and organic chemistry / Question 23. Calcium / 2. Physical and chemical properties

Physical properties. Calcium is a silver-white malleable metal that melts at a temperature of 850 degrees. C and boils at 1482 degrees. C. It is significantly harder than alkali metals.

Chemical properties. Calcium is an active metal. So, under normal conditions, it easily interacts with atmospheric oxygen and halogens:

2 Ca + O2 = 2 CaO (calcium oxide);

Ca + Br2 = CaBr2 (calcium bromide).

Calcium reacts with hydrogen, nitrogen, sulfur, phosphorus, carbon and other non-metals when heated:

Ca + H2 = CaH2 (calcium hydride);

3 Ca + N2 = Ca3N2 (calcium nitride);

Ca + S = CaS (calcium sulfide);

3 Ca + 2 P = Ca3P2 (calcium phosphide);

Ca + 2 C = CaC2 (calcium carbide).

Calcium reacts slowly with cold water, but very vigorously with hot water:

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

Calcium can remove oxygen or halogens from oxides and halides of less active metals, i.e. it has reducing properties:

5 Ca + Nb2O5 = CaO + 2 Nb;

• 1. Being in nature
• 3. Receipt
• 4. Application

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Calcium | directory Pesticides.ru

For many people, knowledge about calcium is limited only to the fact that this element is necessary for healthy bones and teeth. Where else it is contained, why it is needed and how necessary it is, not everyone has an idea. However, calcium is found in many familiar compounds, both natural and man-made. Chalk and lime, stalactites and stalagmites of caves, ancient fossils and cement, gypsum and alabaster, dairy products and anti-osteoporosis drugs - all this and much more is high in calcium.

This element was first obtained by G. Davy in 1808, and at first it was not particularly actively used. However, this metal is now the fifth most produced in the world, and the need for it is increasing year by year. The main area of ​​use of calcium is the production of building materials and mixtures. However, it is necessary to build not only houses, but also living cells. In the human body, calcium is part of the skeleton, makes muscle contractions possible, ensures blood clotting, regulates the activity of a number of digestive enzymes and performs other quite numerous functions. It is no less important for other living objects: animals, plants, fungi and even bacteria. At the same time, the need for calcium is quite high, which makes it possible to classify it as a macronutrient.

Calcium, Ca is a chemical element of the main subgroup of group II of the Mendeleev periodic system. Atomic number – 20. Atomic mass – 40.08.

Calcium is an alkaline earth metal. When free, malleable, fairly hard, white. By density it belongs to light metals.

• Density – 1.54 g/cm3,
• Melting point – +842 °C,
• Boiling point – +1495 °C.

Calcium has pronounced metallic properties. In all compounds the oxidation state is +2.

In air it becomes covered with a layer of oxide, and when heated it burns with a reddish, bright flame. It reacts slowly with cold water, but quickly displaces hydrogen from hot water and forms hydroxide. When interacting with hydrogen, it forms hydrides. At room temperature it reacts with nitrogen, forming nitrides. It also easily combines with halogens and sulfur, and reduces metal oxides when heated.

Calcium is one of the most abundant elements in nature. In the earth's crust its content is 3% of the mass. It occurs in the form of deposits of chalk, limestone, and marble (a natural type of calcium carbonate CaCO3). There are large quantities of deposits of gypsum (CaSO4 x 2h3O), phosphorite (Ca3(PO4)2 and various calcium-containing silicates.

Water

Hard water

Soft water

Soils

Loss of calcium from the soil occurs as a result of its leaching by precipitation. This process depends on the granulometric composition of the soil, the amount of precipitation, the type of plants, the forms and doses of lime and mineral fertilizers. Depending on these factors, calcium losses from the arable layer range from several tens to 200 – 400 kg/ha or more.

Calcium content in different types of soils

Podzolic soils contain 0.73% (of soil dry matter) calcium.

Gray forest – 0.90% calcium.

Chernozems – 1.44% calcium.

Serozems – 6.04% calcium.

In the plant, calcium is found in the form of phosphates, sulfates, carbonates, and in the form of salts of pectic and oxalic acids. Almost up to 65% of calcium in plants can be extracted with water. The rest is treated with weak acetic and hydrochloric acids. Most calcium is found in aging cells.

Functions of calcium

The effect of this element on plants is multifaceted and, as a rule, positive. Calcium:

• Strengthens metabolism;
• Plays an important role in the movement of carbohydrates;
• Affects the metamorphosis of nitrogenous substances;
• Accelerates the consumption of reserve proteins of seeds during germination;
• Plays a role in the process of photosynthesis;
• a strong antagonist of other cations, preventing their excess entry into plant tissues;
• Affects the physicochemical properties of protoplasm (viscosity, permeability, etc.), and therefore the normal course of biochemical processes in the plant;
• Calcium compounds with pectin substances glue the walls of individual cells together;
• Affects enzyme activity.

It should be noted that the influence of calcium compounds (lime) on enzyme activity is expressed not only in direct action, but also due to the improvement of the physico-chemical properties of the soil and its nutritional regime. In addition, liming of the soil significantly affects the processes of vitamin biosynthesis.

Lack (deficiency) of calcium in plants

Lack of calcium primarily affects the development of the root system. The formation of root hairs on the roots stops. The outer root cells are destroyed.

This symptom manifests itself both with a lack of calcium and with an imbalance in the nutrient solution, that is, the predominance of monovalent cations of sodium, potassium and hydrogen in it.

In addition, the presence of nitrate nitrogen in the soil solution increases the supply of calcium to plant tissues, and reduces the supply of ammonia.

Signs of calcium starvation are expected when the calcium content is less than 20% of the cation exchange capacity of the soil.

Symptoms Visually, calcium deficiency is determined by the following signs:

• The roots of plants have damaged tips with a brown color;
• The growing point becomes deformed and dies;
• Flowers, ovaries and buds fall off;
• The fruits are damaged by necrosis;
• The leaves are noted to be chlorotic;
• The apical bud dies and stem growth stops.

Cabbage, alfalfa, and clover are highly sensitive to the presence of calcium. It has been established that these same plants are also characterized by increased sensitivity to soil acidity.

Mineral calcium poisoning results in interveinal chlorosis with whitish necrotic spots. They may be colored or have concentric rings filled with water. Some plants respond to excess calcium by growing leaf rosettes, dying shoots and dropping leaves. The symptoms are similar in appearance to iron and magnesium deficiency.

The source of calcium replenishment in the soil is lime fertilizers. They are divided into three groups:

• Hard calcareous rocks;
• Soft calcareous rocks;
• Industrial waste with high lime content.

Based on the content of CaO and MgO, hard calcareous rocks are divided into:

• limestones (55–56% CaO and up to 0.9% MgO);
• dolomitized limestones (42–55% CaO and up to 9% MgO);
• dolomites (32–30% CaO and 18–20% MgO).

Limestones

Dolomitized limestone

Marl

Chalk

Burnt lime

Slaked lime

Dolomite flour

Calcareous tuffs

Defecation dirt (defecation)

Shale ash cyclones

Dust from furnaces and cement factories

Metallurgical slags

Organic fertilizers. The Ca content in terms of CaCO3 is 0.32–0.40%.

Phosphorite flour. Calcium content – ​​22% CaCO3.

Lime fertilizers are used not only to provide soil and plants with calcium. The main purpose of their use is soil liming. This is a method of chemical reclamation. It is aimed at neutralizing excess soil acidity, improving its agrophysical, agrochemical and biological properties, supplying plants with magnesium and calcium, mobilizing and immobilizing macroelements and microelements, creating optimal water-physical, physical, air conditions for the life of cultivated plants.

Efficiency of soil liming

Simultaneously with satisfying the needs of plants for calcium as an element of mineral nutrition, liming leads to multiple positive changes in soils.

The effect of liming on the properties of some soils

Calcium promotes the coagulation of soil colloids and prevents their leaching. This leads to easier tillage and improved aeration.

As a result of liming:

• sandy humus soils increase their water absorption capacity;
• On heavy clay soils, soil aggregates and clumping are formed, which improve water permeability.

In particular, organic acids are neutralized and H-ions are displaced from the absorbing complex. This leads to the elimination of metabolic acidity and a decrease in hydrolytic acidity of the soil. At the same time, an improvement in the cationic composition of the soil absorption complex is observed, which occurs due to the replacement of hydrogen and aluminum ions with calcium and magnesium cations. This increases the degree of soil saturation with bases and increases the absorption capacity.

The effect of liming on the supply of nitrogen to plants

After liming, the positive agrochemical properties of the soil and its structure can be maintained for several years. This helps create favorable conditions for enhancing beneficial microbiological processes for the mobilization of nutrients. The activity of ammonifiers, nitrifiers, and nitrogen-fixing bacteria that live freely in the soil increases.

Liming helps to increase the proliferation of nodule bacteria and improve the supply of nitrogen to the host plant. It has been established that bacterial fertilizers lose their effectiveness on acidic soils.

The effect of liming on the supply of ash elements to plants

Liming helps supply the plant with ash elements, since it increases the activity of bacteria that decompose organic phosphorus compounds in the soil and promote the transition of iron and aluminum phosphates into calcium phosphate salts available to plants. Liming of acidic soils enhances microbiological and biochemical processes, which, in turn, increases the amount of nitrates, as well as digestible forms of phosphorus and potassium.

Effect of liming on the forms and availability of macroelements and microelements

Liming increases the amount of calcium, and when using dolomite flour - magnesium. At the same time, toxic forms of manganese and aluminum become insoluble and pass into the precipitated form. The availability of elements such as iron, copper, zinc, manganese is decreasing. Nitrogen, sulfur, potassium, calcium, magnesium, phosphorus and molybdenum become more available.

The influence of liming on the action of physiologically acidic fertilizers

Liming increases the effectiveness of physiologically acidic mineral fertilizers, especially ammonia and potash.

The positive effect of physiologically acidic fertilizers without the addition of lime fades, and over time can turn negative. So, in fertilized areas, yields are even less than in unfertilized areas. The combination of liming with the use of fertilizers increases their effectiveness by 25–50%.

When liming, enzymatic processes in the soil are activated, by which its fertility is indirectly judged.

Compiled by: Grigorovskaya P.I.

Page added: 05.12.13 00:40

Last update: 05/22/14 16:25

Literary sources:

Glinka N.L. General chemistry. Textbook for universities. Publisher: Leningrad: Chemistry, 1985, p. 731

Mineev V.G. Agrochemistry: Textbook. – 2nd edition, revised and expanded. – M.: Moscow State University Publishing House, KolosS Publishing House, 2004. – 720 p., l. ill.: ill. – (Classical university textbook).

Petrov B.A., Seliverstov N.F. Mineral nutrition of plants. A reference guide for students and gardeners. Ekaterinburg, 1998. 79 p.

Encyclopedia for children. Volume 17. Chemistry. / Head. ed. V.A. Volodin. – M.: Avanta +, 2000. – 640 p., ill.

Yagodin B.A., Zhukov Yu.P., Kobzarenko V.I. Agrochemistry / Edited by B.A. Yagodina. – M.: Kolos, 2002. – 584 pp.: ill (Textbooks and teaching aids for students of higher educational institutions).

Images (reworked):

20 Ca Calcium, licensed under CC BY

Calcium deficiency in wheat, by CIMMYT, licensed under CC BY-NC-SA

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Calcium and its role for humanity - Chemistry

Calcium and its role for humanity

Introduction

Being in nature

Receipt

Physical properties

Chemical properties

Application of calcium compounds

Biological role

Conclusion

Bibliography

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.

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 completely different from 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. The mechanical properties of elemental calcium do not make it a “black sheep” in the family of metals: calcium surpasses many of them in strength and hardness; it can be turned on a lathe, drawn into wire, forged, pressed.

And yet, 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

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 limestone firing) 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 - is 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 x 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 0.5h3O and gypsum CaSO4 2h3O, 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 the interaction of calcium carbonate with water and carbon dioxide with the formation of soluble bicarbonate:

CaCO3 + h3O + CO2 - Ca (HCO3)2 - Ca2+ + 2HCO3-

(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:

4CaO + 2Al = CaAl2O4 + 3Ca.

Physical properties

Calcium metal exists in two allotropic modifications. Up to 443 °C, ?-Ca with a cubic face-centered lattice (parameter a = 0.558 nm) is stable; higher stable is ?-Ca with a cubic body-centered lattice of the ?-Fe type (parameter a = 0.448 nm). Standard enthalpy?H0 transition? > ? is 0.93 kJ/mol.

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 actively reacts with water, but without ignition:

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

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

2Ca + 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,

3Ca + N2 = Ca3N2, Ca + 2C = CaC2,

3Ca + 2P = Ca3P2 (

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

2Ca + 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^,

Ca3N2 + 3H2O = 3Ca(OH)2 + 2Nh4^.

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 falls on 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 is heated by the sun's rays, a 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.

Applications of calcium metal

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.

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, Cah3 (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. Calcium compounds are widely used as an antihistamine.

Calcium chloride

Calcium gluconate

Calcium glycerophosphate

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

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.

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 (especially 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.

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.

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.

Products Calcium, mg/100 g

Sesame 783

Nettle 713

Forest mallow 505

Large plantain 412

Galinsoga 372

Sardines in oil 330

Ivy budra 289

Dog rose 257

Almond 252

Plantain lanceolist. 248

Hazelnut 226

Amaranth seed 214

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 nature in a free state. But calcium compounds - both natural and artificial - have acquired paramount importance.

Bibliography

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

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

3. Dotsenko VA. - Therapeutic and preventive nutrition. - Question. nutrition, 2001 - N1-p.21-25

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

www.e-ng.ru

World of Science

Calcium is a metal element of the main subgroup II of group 4 of the periodic table of chemical elements. It belongs to the alkaline earth metal family. The outer energy level of the calcium atom contains 2 paired s-electrons

Which he is able to energetically give away during chemical interactions. Thus, Calcium is a reducing agent and in its compounds has an oxidation state of +2. In nature, calcium is found only in the form of salts. The mass fraction of calcium in the earth's crust is 3.6%. The main natural calcium mineral is calcite CaCO3 and its varieties - limestone, chalk, marble. There are also living organisms (for example, corals), the backbone of which consists mainly of calcium carbonate. Also important calcium minerals are dolomite CaCO3 MgCO3, fluorite CaF2, gypsum CaSO4 2h3O, apatite, feldspar, etc. Calcium plays an important role in the life of living organisms. The mass fraction of calcium in the human body is 1.4-2%. It is part of teeth, bones, other tissues and organs, participates in the process of blood clotting, and stimulates cardiac activity. To provide the body with a sufficient amount of calcium, you should definitely consume milk and dairy products, green vegetables, and fish. The simple substance calcium is a typical silver-white metal. It is quite hard, plastic, has a density of 1.54 g/cm3 and a melting point of 842? C. Chemically, calcium is very active. Under normal conditions, it easily interacts with oxygen and moisture in the air, so it is stored in hermetically sealed containers. When heated in air, calcium ignites and forms an oxide: 2Ca + O2 = 2CaO. Calcium reacts with chlorine and bromine when heated, and with fluorine even in the cold. The products of these reactions are the corresponding halides, for example: Ca + Cl2 = CaCl2. When calcium is heated with sulfur, calcium sulfide is formed: Ca + S = CaS. Calcium can also react with other non-metals. Interaction with water leads to the formation of slightly soluble calcium hydroxide and the release of hydrogen gas :Ca + 2h3O = Ca (OH) 2 + h3. Calcium metal is widely used. It is used as a rosette in the production of steels and alloys, and as a reducing agent for the production of some refractory metals.

Calcium is obtained by electrolysis of molten calcium chloride. Thus, calcium was first obtained in 1808 by Humphry Davy.

worldofscience.ru

Calcium (Latin Calcium, symbolized Ca) is an element with atomic number 20 and atomic mass 40.078. It is an element of the main subgroup of the second group, the fourth period of the periodic table of chemical elements of Dmitry Ivanovich Mendeleev. Under normal conditions, the simple substance calcium is a light (1.54 g/cm3) malleable, soft, chemically active alkaline earth metal of silver-white color.

In nature, calcium is presented as a mixture of six isotopes: 40Ca (96.97%), 42Ca (0.64%), 43Ca (0.145%), 44Ca (2.06%), 46Ca (0.0033%) and 48Ca ( 0.185%). The main isotope of the twentieth element - the most common - is 40Ca, its isotopic abundance is about 97%. Of the six natural isotopes of calcium, five are stable; the sixth isotope 48Ca, the heaviest of the six and quite rare (its isotopic abundance is only 0.185%), was recently found to undergo double β-decay with a half-life of 5.3∙1019 years. Isotopes obtained artificially with mass numbers 39, 41, 45, 47 and 49 are radioactive. Most often they are used as an isotopic indicator in the study of mineral metabolism processes in a living organism. 45Ca, obtained by irradiating metallic calcium or its compounds with neutrons in a uranium reactor, plays an important role in the study of metabolic processes occurring in soils and in the study of the processes of calcium absorption by plants. Thanks to the same isotope, it was possible to detect sources of contamination of various types of steel and ultra-pure iron with calcium compounds during the smelting process.

Calcium compounds - marble, gypsum, limestone and lime (a product of limestone firing) have been known since ancient times and were widely used in construction and medicine. The ancient Egyptians used calcium compounds in the construction of their pyramids, and the inhabitants of the great Rome invented concrete - using a mixture of crushed stone, lime and sand. Until the very end of the 18th century, chemists were convinced that lime was a simple solid. It was only in 1789 that Lavoisier suggested that lime, alumina and some other compounds were complex substances. In 1808, calcium metal was obtained by G. Davy by electrolysis.

The use of calcium metal is associated with its high chemical activity. It is used for the recovery from compounds of certain metals, for example, thorium, uranium, chromium, zirconium, cesium, rubidium; for removing oxygen and sulfur from steel and some other alloys; for dehydration of organic liquids; for absorbing residual gases in vacuum devices. In addition, calcium metal serves as an alloying component in some alloys. Calcium compounds are used much more widely - they are used in construction, pyrotechnics, glass production, medicine and many other fields.

Calcium is one of the most important biogenic elements; it is necessary for most living organisms for the normal course of life processes. The adult body contains up to one and a half kilograms of calcium. It is present in all tissues and fluids of living organisms. The twentieth element is necessary for the formation of bone tissue, maintaining heart rate, blood clotting, maintaining normal permeability of outer cell membranes, and the formation of a number of enzymes. The list of functions that calcium performs in the bodies of plants and animals is very long. Suffice it to say that only rare organisms are able to develop in an environment devoid of calcium, and other organisms consist of 38% of this element (the human body contains only about 2% calcium).

Biological properties

Calcium is one of the biogenic elements; its compounds are found in almost all living organisms (few organisms are able to develop in an environment devoid of calcium), ensuring the normal course of life processes. The twentieth element is present in all tissues and liquids of animals and plants; most of it (in vertebrate organisms, including humans) is contained in the skeleton and teeth in the form of phosphates (for example, hydroxyapatite Ca5(PO4)3OH or 3Ca3(PO4)2Ca (OH)2). The use of the twentieth element as a building material for bones and teeth is due to the fact that calcium ions are not used in the cell. Calcium concentration is controlled by special hormones; their combined action preserves and maintains bone structure. The skeletons of most groups of invertebrates (mollusks, corals, sponges and others) are built from various forms of calcium carbonate CaCO3 (lime). Many invertebrates store calcium before molting to build a new skeleton or to ensure vital functions in unfavorable conditions. Animals receive calcium from food and water, and plants - from the soil and in relation to this element they are divided into calciphiles and calcephobes.

The ions of this important microelement are involved in blood clotting processes, as well as in ensuring constant osmotic pressure of the blood. In addition, calcium 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 (perhaps this circumstance is due to the fact that calcium replaces magnesium ions). Calcium 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, regulate exocytosis, including the secretion of hormones and neurotransmitters. Calcium affects the permeability of blood vessels - without this element, fats, lipids and cholesterol would settle on the walls of blood vessels. Calcium promotes the release of heavy metal salts and radionuclides from the body and performs antioxidant functions. Calcium affects the reproductive system, has an anti-stress effect and has an anti-allergic effect.

The calcium content in the body of an adult (weighing 70 kg) is 1.7 kg (mainly in the intercellular substance of bone tissue). The need for this element depends on age: for adults the required daily intake is from 800 to 1,000 milligrams, for children from 600 to 900 milligrams. For children, it is especially important to consume the required dose for intensive bone growth and development. The main source of calcium in the body is milk and dairy products; the rest of calcium comes from meat, fish, and some plant products (especially legumes). Absorption of calcium cations occurs in the large and small intestines; absorption is facilitated by an acidic environment, vitamins C and D, lactose (lactic acid), and unsaturated fatty acids. In turn, aspirin, oxalic acid, and estrogen derivatives significantly reduce the digestibility of the twentieth element. Thus, when combined with oxalic acid, calcium produces water-insoluble compounds that are components of kidney stones. The role of magnesium in calcium metabolism is great - with its deficiency, calcium is “washed out” from the bones and deposited in the kidneys (kidney stones) and muscles. In general, the body has a complex system for storing and releasing the twentieth element; for this reason, the calcium content in the blood is precisely regulated, and with proper nutrition, deficiency or excess does not occur. A long-term calcium diet can cause cramps, joint pain, constipation, fatigue, drowsiness, and growth retardation. A prolonged lack of calcium in the diet leads to the development of osteoporosis. Nicotine, caffeine and alcohol are some of the causes of calcium deficiency in the body, as they contribute to its intensive excretion in the urine. However, an excess of the twentieth element (or vitamin D) leads to negative consequences - hypercalcemia develops, the consequence of which is intense calcification of bones and tissues (mainly affecting the urinary system). A long-term calcium surplus disrupts the functioning of muscle and nerve tissues, increases blood clotting and reduces the absorption of zinc by bone cells. Osteoarthritis, cataracts, and blood pressure problems may occur. From the above we can conclude that the cells of plant and animal organisms need strictly defined ratios of calcium ions.

In pharmacology and medicine, calcium compounds are used for the manufacture of vitamins, tablets, pills, injections, antibiotics, as well as for the manufacture of ampoules and medical utensils.

It turns out that a fairly common cause of male infertility is a lack of calcium in the body! The fact is that the head of the sperm has an arrow-shaped formation, which consists entirely of calcium; with a sufficient amount of this element, the sperm is able to overcome the membrane and fertilize the egg; if there is insufficient amount, infertility occurs.

American scientists have found that a lack of calcium ions in the blood leads to weakened memory and decreased intelligence. For example, from the well-known US magazine Science News, it became known about experiments that confirmed that cats develop a conditioned reflex only if their brain cells contain more calcium than blood.

The compound calcium cyanamide, highly valued in agriculture, is used not only as a nitrogen fertilizer and a source of urea - a valuable fertilizer and raw material for the production of synthetic resins, but also as a substance with which it was possible to mechanize the harvesting of cotton fields. The fact is that after treatment with this compound, the cotton plant instantly sheds its leaves, which allows people to leave the cotton picking to machines.

When talking about foods rich in calcium, dairy products are always mentioned, but milk itself contains from 120 mg (cow) to 170 mg (sheep) calcium per 100 g; cottage cheese is even poorer - only 80 mg per 100 grams. Of the dairy products, only cheese contains from 730 mg (Gouda) to 970 mg (Emmenthal) of calcium per 100 g of product. However, the record holder for the content of the twentieth element is poppy - 100 grams of poppy seeds contain almost 1,500 mg of calcium!

Calcium chloride CaCl2, used, for example, in refrigeration units, is a waste product of many chemical technological processes, in particular large-scale soda production. However, despite the widespread use of calcium chloride in various fields, its consumption is significantly lower than its production. For this reason, for example, near soda factories, entire lakes of calcium chloride brine are formed. Such storage ponds are not uncommon.

In order to understand how much calcium compounds are consumed, it is worth giving just a couple of examples. In steel production, lime is used to remove phosphorus, silicon, manganese and sulfur; in the oxygen-converter process, 75 kilograms of lime are consumed per ton of steel! Another example comes from a completely different area - the food industry. In sugar production, crude sugar syrup is reacted with lime to precipitate calcium sucrose. So, cane sugar usually requires about 3-5 kg ​​of lime per ton, and beet sugar - a hundred times more, that is, about half a ton of lime per ton of sugar!

“Hardness” of water is a number of properties that calcium and magnesium salts dissolved in it give water. Stiffness is divided into temporary and permanent. Temporary or carbonate hardness is caused by the presence of soluble hydrocarbonates Ca(HCO3)2 and Mg(HCO3)2 in water. It is very easy to get rid of carbonate hardness - when water is boiled, bicarbonates turn into water-insoluble calcium and magnesium carbonates, precipitating. Permanent hardness is created by sulfates and chlorides of the same metals, but getting rid of it is much more difficult. Hard water is dangerous not so much because it prevents the formation of soap suds and therefore washes clothes worse; what is much worse is that it forms a layer of scale in steam boilers and boiler systems, thereby reducing their efficiency and leading to emergency situations. What’s interesting is that they knew how to determine the hardness of water back in Ancient Rome. Red wine was used as a reagent - its coloring substances form a precipitate with calcium and magnesium ions.

The process of preparing calcium for storage is very interesting. Calcium metal is stored for a long time in the form of pieces weighing from 0.5 to 60 kg. These “ingots” are packed in paper bags, then placed in galvanized iron containers with soldered and painted seams. Tightly closed containers are placed in wooden boxes. Pieces weighing less than half a kilogram cannot be stored for a long time - when oxidized, they quickly turn into oxide, hydroxide and calcium carbonate.

Story

Calcium metal was obtained relatively recently - in 1808, but humanity has been familiar with compounds of this metal for a very long time. Since ancient times, people have used limestone, chalk, marble, alabaster, gypsum and other calcium-containing compounds in construction and medicine. Limestone CaCO3 was most likely the first building material used by humans. It was used in the construction of the Egyptian pyramids and the Great Wall of China. Many temples and churches in Rus', as well as most of the buildings of ancient Moscow, were built using limestone - a white stone. Even in ancient times, a person, by burning limestone, received quicklime (CaO), as evidenced by the works of Pliny the Elder (1st century AD) and Dioscorides, a doctor in the Roman army, to whom he introduced calcium oxide in his essay “On Medicines.” the name “quicklime”, which has survived to this day. And all this despite the fact that pure calcium oxide was first described by the German chemist I. Then only in 1746, and in 1755, the chemist J. Black, studying the firing process, revealed that the loss of limestone mass during firing occurs due to the release of carbon dioxide gas:

CaCO3 ↔ CO2 + CaO

The Egyptian mortars that were used in the Giza pyramids were based on partially dehydrated gypsum CaSO4 2H2O or, in other words, alabaster 2CaSO4∙H2O. It is also the basis of all the plaster in the tomb of Tutankhamun. The Egyptians used burnt gypsum (alabaster) as a binder in the construction of irrigation structures. By burning natural gypsum at high temperatures, Egyptian builders achieved its partial dehydration, and not only water, but also sulfuric anhydride was split off from the molecule. Subsequently, when diluted with water, a very strong mass was obtained that was not afraid of water and temperature fluctuations.

The Romans can rightfully be called the inventors of concrete, because in their buildings they used one of the varieties of this building material - a mixture of crushed stone, sand and lime. There is a description by Pliny the Elder of the construction of cisterns from such concrete: “To build cisterns, take five parts of pure gravel sand, two parts of the best slaked lime and fragments of silex (hard lava) weighing no more than a pound each, after mixing, compact the bottom and side surfaces with the blows of an iron rammer " In Italy's humid climate, concrete was the most resilient material.

It turns out that humanity has long been aware of calcium compounds, which they widely consumed. However, until the end of the 18th century, chemists considered lime to be a simple solid; only on the threshold of the new century did the study of the nature of lime and other calcium compounds begin. So Stahl suggested that lime was a complex body consisting of earthy and watery principles, and Black established the difference between caustic lime and carbonic lime, which contained “fixed air.” Antoine Laurent Lavoisier classified calcareous earth (CaO) as an element, that is, as a simple substance, although in 1789 he suggested that lime, magnesia, barite, alumina and silica are complex substances, but it will be possible to prove this only by decomposing the “stubborn earth” (calcium oxide). And the first person to succeed was Humphry Davy. After the successful decomposition of potassium and sodium oxides by electrolysis, the chemist decided to obtain alkaline earth metals in the same way. However, the first attempts were unsuccessful - the Englishman tried to decompose lime by electrolysis in air and under a layer of oil, then calcined the lime with metallic potassium in a tube and carried out many other experiments, but to no avail. Finally, in a device with a mercury cathode, he obtained an amalgam by electrolysis of lime, and from it metallic calcium. Quite soon, this method of obtaining metal was improved by I. Berzelius and M. Pontin.

The new element received its name from the Latin word “calx” (in the genitive case calcis) - lime, soft stone. Calx was the name given to chalk, limestone, generally pebble stone, but most often lime-based mortar. This concept was also used by ancient authors (Vitruvius, Pliny the Elder, Dioscorides), describing the burning of limestone, slaking lime and preparing mortars. Later, in the circle of alchemists, “calx” denoted the product of firing in general - in particular metals. For example, metal oxides were called metallic limes, and the firing process itself was called calcination. In ancient Russian prescription literature the word kal (dirt, clay) is found, so in the collection of the Trinity-Sergius Lavra (XV century) it is said: “find feces, from it they create the gold of the crucible.” It was only later that the word feces, which is undoubtedly related to the word "calx", became synonymous with the word dung. In Russian literature of the early 19th century, calcium was sometimes called the base of calcareous earth, liming (Shcheglov, 1830), calcification (Iovsky), calcium, calcium (Hess).

Being in nature

Calcium is one of the most common elements on our planet - the fifth in quantitative content in nature (of non-metals, only oxygen is more common - 49.5% and silicon - 25.3%) and third among metals (only aluminum is more common - 7.5% and iron - 5.08%). Clarke (the average content in the earth's crust) of calcium, according to various estimates, ranges from 2.96% by mass to 3.38%, we can definitely say that this figure is about 3%. The outer shell of the calcium atom has two valence electrons, the connection of which with the nucleus is rather weak. For this reason, calcium is highly chemically reactive and does not occur in free form in nature. However, it actively migrates and accumulates in various geochemical systems, forming approximately 400 minerals: silicates, aluminosilicates, carbonates, phosphates, sulfates, borosilicates, molybdates, chlorides and others, ranking fourth in this indicator. When basaltic magmas melt, calcium accumulates in the melt and is included in the composition of the main rock-forming minerals, during the fractionation of which its content decreases during the differentiation of magma from basic to acidic rocks. For the most part, calcium lies in the lower part of the earth's crust, accumulating in basic rocks (6.72%); there is little calcium in the earth's mantle (0.7%) and, probably, even less in the earth's core (in iron meteorites similar to the core, the twentieth element is only 0.02%).

True, the clarke of calcium in stony meteorites is 1.4% (rare calcium sulfide is found), in medium-sized rocks it is 4.65%, and acidic rocks contain 1.58% calcium by weight. The main part of calcium is contained in silicates and aluminosilicates of various rocks (granites, gneisses, etc.), especially in feldspar - anorthite Ca, as well as diopside CaMg, wollastonite Ca3. In the form of sedimentary rocks, calcium compounds are represented by chalk and limestones, consisting mainly of the mineral calcite (CaCO3).

Calcium carbonate CaCO3 is one of the most abundant compounds on Earth - calcium carbonate minerals cover approximately 40 million square kilometers of the earth's surface. In many parts of the Earth's surface there are significant sedimentary deposits of calcium carbonate, which were formed from the remains of ancient marine organisms - chalk, marble, limestone, shell rocks - all this is CaCO3 with minor impurities, and calcite is pure CaCO3. The most important of these minerals is limestone, or rather limestones - because each deposit differs in density, composition and amount of impurities. For example, shell rock is limestone of organic origin, and calcium carbonate, which has fewer impurities, forms transparent crystals of limestone or Iceland spar. Chalk is another common type of calcium carbonate, but marble, a crystalline form of calcite, is much less common in nature. It is generally accepted that marble was formed from limestone in ancient geological eras. As the earth's crust moved, individual deposits of limestone became buried under layers of other rocks. Under the influence of high pressure and temperature, the process of recrystallization occurred, and the limestone turned into a denser crystalline rock - marble. Bizarre stalactites and stalagmites are the mineral aragonite, which is another type of calcium carbonate. Orthorhombic aragonite is formed in warm seas - huge layers of calcium carbonate in the form of aragonite are formed in the Bahamas, the Florida Keys and the Red Sea basin. Also quite widespread are calcium minerals such as fluorite CaF2, dolomite MgCO3 CaCO3, anhydrite CaSO4, phosphorite Ca5(PO4)3(OH,CO3) (with various impurities) and apatites Ca5(PO4)3(F,Cl,OH) - forms of calcium phosphate, alabaster CaSO4 0.5H2O and gypsum CaSO4 2H2O (forms of calcium sulfate) and others. Calcium-containing minerals contain isomorphically replacing impurity elements (for example, sodium, strontium, rare earth, radioactive and other elements).

A large amount of the twentieth element is found in natural waters due to the existence of a global “carbonate equilibrium” between poorly soluble CaCO3, highly soluble Ca(HCO3)2 and CO2 found in water and air:

CaCO3 + H2O + CO2 = Ca(HCO3)2 = Ca2+ + 2HCO3-

This reaction is reversible and is the basis for the redistribution of the twentieth element - with a high carbon dioxide content in waters, calcium is in solution, and with a low CO2 content, the mineral calcite CaCO3 precipitates, forming thick deposits of limestone, chalk, and marble.

A considerable amount of calcium is part of living organisms, for example, hydroxyapatite Ca5(PO4)3OH, or, in another entry, 3Ca3(PO4)2 Ca(OH)2 - the basis of the bone tissue of vertebrates, including humans. Calcium carbonate CaCO3 is the main component of the shells and shells of many invertebrates, eggshells, corals and even pearls.

Application

Calcium metal is used quite rarely. Basically, this metal (as well as its hydride) is used in the metallothermic production of difficult-to-reduce metals - uranium, titanium, thorium, zirconium, cesium, rubidium and a number of rare earth metals from their compounds (oxides or halides). Calcium is used as a reducing agent in the production of nickel, copper and stainless steel. The twentieth element is also used for deoxidation of steels, bronzes and other alloys, for removing sulfur from petroleum products, for dehydrating organic solvents, for purifying argon from nitrogen impurities and as a gas absorber in electric vacuum devices. Calcium metal is used in the production of antifriction alloys of the Pb-Na-Ca system (used in bearings), as well as a Pb-Ca alloy used for the manufacture of electrical cable sheaths. Silicocalcium alloy (Ca-Si-Ca) is used as a deoxidizing agent and degassing agent in the production of quality steels. Calcium is used both as an alloying element for aluminum alloys and as a modifying additive for magnesium alloys. For example, the introduction of calcium increases the strength of aluminum bearings. Pure calcium is also 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. With the help of calcium, the carbon content in cast iron is regulated and bismuth is removed from lead, and the steel is purified from oxygen, sulfur and phosphorus. Calcium, as well as its alloys with aluminum and magnesium, are used in thermal electric backup batteries as an anode (for example, calcium chromate element).

However, compounds of the twentieth element are used much more widely. And first of all we are talking about natural calcium compounds. One of the most common calcium compounds on Earth is CaCO3 carbonate. Pure calcium carbonate is the mineral calcite, and limestone, chalk, marble, and shell rock are CaCO3 with minor impurities. Mixed calcium and magnesium carbonate is called dolomite. Limestone and dolomite are used mainly as building materials, road surfaces, or soil deacidifiers. Calcium carbonate CaCO3 is necessary for the production of calcium oxide (quicklime) CaO and calcium hydroxide (slaked lime) Ca(OH)2. In turn, CaO and Ca(OH)2 are the main substances in many areas of the chemical, metallurgical and mechanical engineering industries - calcium oxide, both in free form and as part of ceramic mixtures, is used in the production of refractory materials; Colossal volumes of calcium hydroxide are needed by the pulp and paper industry. In addition, Ca(OH)2 is used in the production of bleach (a good bleaching and disinfectant), Berthollet salt, soda, and some pesticides to control plant pests. A huge amount of lime is consumed in the production of steel - to remove sulfur, phosphorus, silicon and manganese. Another role of lime in metallurgy is the production of magnesium. Lime is also used as a lubricant in drawing steel wire and neutralizing waste pickling fluids containing sulfuric acid. In addition, lime is the most common chemical reagent in the treatment of drinking and industrial water (together with alum or iron salts, it coagulates suspensions and removes sediment, and also softens water by removing temporary - bicarbonate - hardness). In everyday life and medicine, precipitated calcium carbonate is used as an acid neutralizer, a mild abrasive in toothpastes, a source of additional calcium in diets, a component of chewing gum, and a filler in cosmetics. CaCO3 is also used as a filler in rubbers, latexes, paints and enamels, as well as in plastics (about 10% by weight) to improve their heat resistance, stiffness, hardness and workability.

Calcium fluoride CaF2 is of particular importance, because in the form of a mineral (fluorite) it is the only industrially important source of fluorine! Calcium fluoride (fluorite) is used in the form of single crystals in optics (astronomical objectives, lenses, prisms) and as a laser material. The fact is that glasses only made of calcium fluoride are permeable to the entire spectrum region. Calcium tungstate (scheelite) in the form of single crystals is used in laser technology and also as a scintillator. No less important is calcium chloride CaCl2 - a component of brines for refrigeration units and for filling tires of tractors and other vehicles. With the help of calcium chloride, roads and sidewalks are cleared of snow and ice; this compound is used to protect coal and ore from freezing during transportation and storage; wood is impregnated with its solution to make it fire-resistant. CaCl2 is used in concrete mixtures to accelerate the onset of setting and increase the initial and final strength of concrete.

Artificially produced calcium carbide CaC2 (by calcination of calcium oxide with coke in electric furnaces) is used to produce acetylene and to reduce metals, as well as to produce calcium cyanamide, which, in turn, releases ammonia under the action of water vapor. In addition, calcium cyanamide is used to produce urea - a valuable fertilizer and raw material for the production of synthetic resins. 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 (more than a cubic meter of hydrogen can be obtained from 1 kilogram of calcium hydride), which is used to fill balloons, for example. In laboratory practice, calcium hydride is used as an energetic reducing agent. The insecticide calcium arsenate, which is obtained by neutralizing arsenic acid with lime, is widely used to combat cotton weevil, codling moth, tobacco worm, and Colorado potato beetle. Important fungicides are lime sulfate sprays and Bordeaux mixtures, which are made from copper sulfate and calcium hydroxide.

Production

The first person to obtain calcium metal was the English chemist Humphry Davy. In 1808, he electrolyzed a mixture of wet slaked lime Ca(OH)2 with mercury oxide HgO on a platinum plate that served as an anode (a platinum wire immersed in mercury acted as a cathode), as a result of which Davy obtained calcium amalgam by removing mercury from it , the chemist obtained a new metal, which he called calcium.

In modern industry, free metallic calcium is obtained by electrolysis of a melt of calcium chloride CaCl2, the share of which is 75-85%, and potassium chloride KCl (it is possible to use a mixture of CaCl2 and CaF2) or by aluminothermic reduction of calcium oxide CaO at a temperature of 1,170-1,200 °C. The pure anhydrous calcium chloride required for electrolysis is obtained by chlorinating calcium oxide when heated in the presence of coal or by dehydrating CaCl2∙6H2O obtained by the action of hydrochloric acid on limestone. The electrolytic process takes place in an electrolysis bath, into which dry calcium chloride salt, free of impurities, and potassium chloride, necessary to lower the melting point of the mixture, are placed. Graphite blocks are placed above the bath - the anode, a cast iron or steel bath filled with a copper-calcium alloy, acts as a cathode. During the electrolysis process, calcium passes into the copper-calcium alloy, significantly enriching it; part of the enriched alloy is constantly removed; instead, an alloy depleted in calcium (30-35% Ca) is added, at the same time chlorine forms a chlorine-air mixture (anode gases), which subsequently goes to the chlorination of lime milk. The enriched copper-calcium alloy can be used directly as an alloy or sent for purification (distillation), where metallic calcium of nuclear purity is obtained from it by distillation in vacuum (at a temperature of 1,000-1,080 ° C and a residual pressure of 13-20 kPa). To obtain high-purity calcium, it is distilled twice. The electrolysis process is carried out at a temperature of 680-720 °C. The fact is that this is the most optimal temperature for the electrolytic process - at a lower temperature, the calcium-enriched alloy floats to the surface of the electrolyte, and at a higher temperature, calcium dissolves in the electrolyte with the formation of CaCl. During electrolysis with liquid cathodes from alloys of calcium and lead or calcium and zinc, alloys of calcium with lead (for bearings) and with zinc (for producing foam concrete - when the alloy reacts with moisture, hydrogen is released and a porous structure is created) are directly obtained. Sometimes the process is carried out with a cooled iron cathode, which only comes into contact with the surface of the molten electrolyte. As calcium is released, the cathode is gradually raised and a rod (50-60 cm) of calcium is pulled out of the melt, protected from atmospheric oxygen by a layer of solidified electrolyte. The “touch method” produces calcium heavily contaminated with calcium chloride, iron, aluminum, and sodium; purification is carried out by melting in an argon atmosphere.

Another method for producing calcium - metallothermic - was theoretically justified back in 1865 by the famous Russian chemist N. N. Beketov. The aluminothermic method is based on the reaction:

6CaO + 2Al → 3CaO Al2O3 + 3Ca

Briquettes are pressed from a mixture of calcium oxide and powdered aluminum, they are placed in a chromium-nickel steel retort and the resulting calcium is distilled off at 1,170-1,200 °C and a residual pressure of 0.7-2.6 Pa. Calcium is obtained in the form of steam, which is then condensed on a cold surface. The aluminothermic method for producing calcium is used in China, France and a number of other countries. The USA was the first to use the metallothermic method of producing calcium on an industrial scale during the Second World War. In the same way, calcium can be obtained by reducing CaO with ferrosilicon or silicoaluminium. Calcium is produced in the form of ingots or sheets with a purity of 98-99%.

Pros and cons exist in both methods. The electrolytic method is multi-operational, energy-intensive (40-50 kWh of energy is consumed per 1 kg of calcium), and is also not environmentally friendly, requiring a large amount of reagents and materials. However, the calcium yield with this method is 70-80%, while with the aluminothermic method the yield is only 50-60%. In addition, with the metallothermic method of obtaining calcium, the disadvantage is that it is necessary to carry out repeated distillation, and the advantage is low energy consumption and the absence of gas and liquid harmful emissions.

Not long ago, a new method for producing calcium metal was developed - it is based on the thermal dissociation of calcium carbide: carbide heated in a vacuum to 1,750 °C decomposes to form calcium vapor and solid graphite.

Until the middle of the 20th century, calcium metal was produced in very small quantities, as it found almost no application. For example, in the United States of America during the Second World War, no more than 25 tons of calcium were consumed, and in Germany only 5-10 tons. Only in the second half of the 20th century, when it became clear that calcium is an active reducing agent for many rare and refractory metals, a rapid increase in consumption (about 100 tons per year) and, as a consequence, production of this metal began. With the development of the nuclear industry, where calcium is used as a component of the metallothermic reduction of uranium from uranium tetrafluoride (except in the United States, where magnesium is used instead of calcium), the demand (about 2,000 tons per year) for element number twenty, as well as its production, has increased manifold. At the moment, China, Russia, Canada and France can be considered the main producers of calcium metal. From these countries, calcium is sent to the USA, Mexico, Australia, Switzerland, Japan, Germany, and the UK. Prices for calcium metal rose steadily until China began producing the metal in such quantities that there was a surplus of the twentieth element on the world market, causing the price to plummet.

Physical properties

What is calcium metal? What properties does this element, obtained in 1808 by the English chemist Humphry Davy, have, a metal whose mass in the body of an adult can be up to 2 kilograms?

The simple substance calcium is a silvery-white light metal. The density of calcium is only 1.54 g/cm3 (at a temperature of 20 °C), which is significantly less than the density of iron (7.87 g/cm3), lead (11.34 g/cm3), gold (19.3 g/cm3) or platinum (21.5 g/cm3). Calcium is even lighter than such “weightless” metals as aluminum (2.70 g/cm3) or magnesium (1.74 g/cm3). Few metals can “boast” a density lower than that of the twentieth element - sodium (0.97 g/cm3), potassium (0.86 g/cm3), lithium (0.53 g/cm3). The density of calcium is very similar to rubidium (1.53 g/cm3). The melting point of calcium is 851 °C, the boiling point is 1,480 °C. Other alkaline earth metals have similar melting points (albeit slightly lower) and boiling points - strontium (770 °C and 1,380 °C) and barium (710 °C and 1,640 °C).

Metallic calcium exists in two allotropic modifications: at normal temperatures up to 443 ° C, α-calcium is stable with a cubic face-centered lattice like copper, with parameters: a = 0.558 nm, z = 4, space group Fm3m, atomic radius 1.97 A, ionic Ca2+ radius 1.04 A; in the temperature range 443-842 °C, β-calcium with a body-centered cubic lattice of the α-iron type is stable, with parameters a = 0.448 nm, z = 2, space group Im3m. The standard enthalpy of transition from the α-modification to the β-modification is 0.93 kJ/mol. The temperature coefficient of linear expansion for calcium in the temperature range 0-300 °C is 22 10-6. The thermal conductivity of the twentieth element at 20 °C is 125.6 W/(m K) or 0.3 cal/(cm sec °C). The specific heat capacity of calcium in the range from 0 to 100 ° C is 623.9 J/(kg K) or 0.149 cal/(g °C). The electrical resistivity of calcium at a temperature of 20° C is 4.6 10-8 ohm m or 4.6 10-6 ohm cm; temperature coefficient of electrical resistance of element number twenty is 4.57 10-3 (at 20 °C). Calcium elastic modulus 26 H/m2 or 2600 kgf/mm2; tensile strength 60 MN/m2 (6 kgf/mm2); the elastic limit for calcium is 4 MN/m2 or 0.4 kgf/mm2, the yield strength is 38 MN/m2 (3.8 kgf/mm2); relative elongation of the twentieth element 50%; Calcium hardness according to Brinell is 200-300 MN/m2 or 20-30 kgf/mm2. With a gradual increase in pressure, calcium begins to exhibit the properties of a semiconductor, but does not become one in the full sense of the word (at the same time, it is no longer a metal). With a further increase in pressure, calcium returns to the metallic state and begins to exhibit superconducting properties (the temperature of superconductivity is six times higher than that of mercury, and far exceeds all other elements in conductivity). The unique behavior of calcium is similar in many ways to strontium (that is, the parallels in the periodic table remain).

The mechanical properties of elemental calcium do not differ from the properties of other members of the family of metals, which are excellent structural materials: high-purity calcium metal is ductile, easily pressed and rolled, drawn into wire, forged and amenable to cutting - it can be turned on a lathe. However, despite all these excellent qualities of a construction material, calcium is not one - the reason for this is its high chemical activity. True, we should not forget that calcium is an irreplaceable structural material of bone tissue, and its minerals have been a building material for many millennia.

Chemical properties

The configuration of the outer electron shell of the calcium atom is 4s2, which determines the valency 2 of the twentieth element in compounds. Two electrons of the outer layer are relatively easily split off from the atoms, which turn into positive doubly charged ions. For this reason, in terms of chemical activity, calcium is only slightly inferior to alkali metals (potassium, sodium, lithium). Like the latter, calcium, even at ordinary room temperature, easily interacts with oxygen, carbon dioxide and moist air, becoming covered with a dull gray film of a mixture of CaO oxide and Ca(OH)2 hydroxide. Therefore, calcium is stored in a hermetically sealed container under a layer of mineral oil, liquid paraffin or kerosene. When heated in oxygen and air, calcium ignites, burning with a bright red flame, forming the basic oxide CaO, which is a white, highly fire-resistant substance with a melting point of approximately 2,600 °C. Calcium oxide is also known in engineering as quicklime or burnt lime. Calcium peroxides - CaO2 and CaO4 - were also obtained. Calcium reacts with water to release hydrogen (in a series of standard potentials, calcium is located to the left of hydrogen and is capable of displacing it from water) and the formation of calcium hydroxide Ca(OH)2, and in cold water the reaction rate gradually decreases (due to the formation of a poorly soluble layer on the metal surface calcium hydroxide):

Ca + 2H2O → Ca(OH)2 + H2 + Q

Calcium reacts more energetically with hot water, rapidly displacing hydrogen and forming Ca(OH)2. Calcium hydroxide Ca(OH)2 is a strong base, slightly soluble in water. A saturated solution of calcium hydroxide is called lime water and is alkaline. In air, limewater quickly becomes cloudy due to the absorption of carbon dioxide and the formation of insoluble calcium carbonate. Despite such violent processes occurring during the interaction of the twentieth element with water, yet, unlike alkali metals, the reaction between calcium and water proceeds less energetically - without explosions or fires. In general, the chemical activity of calcium is lower than that of other alkaline earth metals.

Calcium actively combines with halogens, forming compounds of the CaX2 type - it reacts with fluorine in the cold, and with chlorine and bromine at temperatures above 400 ° C, giving CaF2, CaCl2 and CaBr2, respectively. These halides in the molten state form with calcium monohalides of the CaX type - CaF, CaCl, in which calcium is formally monovalent. These compounds are stable only above the melting temperatures of dihalides (they disproportionate upon cooling to form Ca and CaX2). In addition, calcium actively interacts, especially when heated, with various non-metals: with sulfur, when heated, calcium sulfide CaS is obtained, the latter adds sulfur, forming polysulfides (CaS2, CaS4 and others); interacting with dry hydrogen at a temperature of 300-400 °C, calcium forms the hydride CaH2 - an ionic compound in which hydrogen is an anion. Calcium hydride CaH2 is a white salt-like substance that reacts violently with water to release hydrogen:

CaH2 + 2H2O → Ca(OH)2 + 2H2

When heated (about 500° C) in a nitrogen atmosphere, calcium ignites and forms nitride Ca3N2, known in two crystalline forms - high-temperature α and low-temperature β. Nitride Ca3N4 was also obtained by heating calcium amide Ca(NH2)2 in vacuum. When heated without air access with graphite (carbon), silicon or phosphorus, calcium gives, respectively, calcium carbide CaC2, silicides Ca2Si, Ca3Si4, CaSi, CaSi2 and phosphides Ca3P2, CaP and CaP3. Most of the calcium compounds with non-metals are easily decomposed by water:

CaH2 + 2H2O → Ca(OH)2 + 2H2

Ca3N2 + 6H2O → 3Ca(OH)2 + 2NH3

With boron, calcium forms calcium boride CaB6, with chalcogens - chalcogenides CaS, CaSe, CaTe. Polychalcogenides CaS4, CaS5, Ca2Te3 are also known. Calcium forms intermetallic compounds with various metals - aluminum, gold, silver, copper, lead and others. Being an energetic reducing agent, calcium displaces almost all metals from their oxides, sulfides and halides when heated. Calcium dissolves well in liquid ammonia NH3 to form a blue solution, upon evaporation of which ammonia [Ca(NH3)6] is released - a golden-colored solid compound with metallic conductivity. Calcium salts are usually obtained by the interaction of acid oxides with calcium oxide, the action of acids on Ca(OH)2 or CaCO3, and exchange reactions in aqueous solutions of electrolytes. Many calcium salts are highly soluble in water (CaCl2 chloride, CaBr2 bromide, CaI2 iodide and Ca(NO3)2 nitrate), they almost always form crystalline hydrates. Insoluble in water are fluoride CaF2, carbonate CaCO3, sulfate CaSO4, orthophosphate Ca3(PO4)2, oxalate CaC2O4 and some others.

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 with a silvery-white color. Due to interaction with oxygen and carbon dioxide, the surface of the metal becomes dull, so calcium requires a special storage regime - a tightly closed container, in which the metal is filled 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 food is the consumption of carbohydrates in the form of sweets and alkalis, which neutralize the hydrochloric acid of the stomach, which is necessary to dissolve calcium. The process of calcium absorption is quite complex, so sometimes it is not enough to get it only from food; additional intake of the microelement is necessary.

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 the processes of excitability of nervous 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, and 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).

Although calcium is very widespread on the globe, it does not occur in a free state in nature.

Before we learn how you can get pure calcium, let's get to know the natural calcium compounds.

Calcium is a metal. In the periodic table of Mendeleev, calcium (Calcium), Ca has atomic number 20 andlocated in group II. This is a chemically active element; it easily interacts with oxygen. It has a silvery-white color.

Natural calcium compounds

We find calcium compounds almost everywhere.

Calcium carbonate, or calcium carbonate – it is the most common calcium compound. Its chemical formula is CaCO 3. Marble, chalk, limestone, shell rock - all these substances contain calcium carbonate with a small amount of impurities. There are no impurities at all in calcite, the formula of which is also CaCO 3.

Calcium sulfate also called calcium sulfate. The chemical formula of calcium sulfate is CaSO 4. The mineral gypsum known to us is the crystalline hydrate CaSO 4 2H 2 O.

Calcium phosphate, orthophosphoric acid calcium salt. This is the material from which the bones of humans and animals are built. This mineral is called tricalcium phosphate Ca 3 (PO 4) 2.

Calcium chlorideCaCl 2, or calcium chloride, occurs in nature in the form of crystalline hydrate CaCl 2 6H 2 O. When heated, this compound loses water molecules.

Calcium fluoride CaF 2, or calcium fluoride, can be found naturally in the mineral fluorite. And pure crystalline calcium difluoride is called fluorspar.

But natural calcium compounds do not always have the properties that people need. Therefore, man has learned to artificially transform such compounds into other substances. Some of these artificial compounds are even more familiar to us than natural ones. An example is slaked Ca(OH) 2 and quicklime CaO, which have been used by humans for a very long time. Many building materials such as cement, calcium carbide, bleach also contain artificial calcium compounds.

What is electrolysis

Probably almost all of us have heard about the phenomenon called electrolysis. We will try to give the simplest description of this process.

If you pass an electric current through aqueous solutions of salts, new chemical substances are formed as a result of chemical transformations. The processes that occur in a solution when an electric current is passed through it are called electrolysis. All these processes are studied by a science called electrochemistry. Of course, the electrolysis process can only take place in a medium that conducts current. Aqueous solutions of acids, bases and salts are such a medium. They are called electrolytes.

The electrodes are immersed in the electrolyte. The negatively charged electrode is called the cathode. A positively charged electrode is called an anode. When an electric current passes through an electrolyte, electrolysis occurs. As a result of electrolysis, the constituents of dissolved substances are deposited on the electrodes. At the cathode they are positively charged, at the anode they are negative. But secondary reactions can occur on the electrodes themselves, resulting in the formation of a secondary substance.

We see that with the help of electrolysis, chemical products are formed without the use of chemical reagents.

How do you get calcium?

In industry, calcium can be obtained by electrolysis of molten calcium chloride CaCl 2.

CaCl 2 = Ca + Cl 2

In this process, a bath made of graphite serves as the anode. The bath is placed in an electric oven. An iron rod that moves across the width of the bath, and also has the ability to rise and fall, is the cathode. The electrolyte is molten calcium chloride, which is poured into the bath. The cathode is lowered into the electrolyte. This is how the electrolysis process begins. Molten calcium is formed under the cathode. When the cathode rises, calcium solidifies where it touches the cathode. So gradually, in the process of raising the cathode, calcium builds up in the form of a rod. Then the calcium rod is knocked away from the cathode.

Pure calcium was first obtained by electrolysis in 1808.

Calcium is also obtained from oxides using aluminothermic reduction .

4CaO + 2Al -> CaAl 2 O 4 + Ca

In this case, calcium is obtained in the form of steam. This steam then condenses.

Calcium has high chemical activity. That is why it is widely used in industry for the recovery of refractory metals from oxides, as well as in the production of steel and cast iron.