What elements does calcium belong to? Calcium (chemical element). Being in nature and receiving

Calcium is located in the fourth major period, the second group, the main subgroup, the serial number of the element is 20. According to the periodic table of Mendeleev, the atomic weight of calcium is 40.08. The formula of the highest oxide is CaO. Calcium has a Latin name calcium, so the element's atom symbol is Ca.

Characteristics of calcium as a simple substance

Under normal conditions, calcium is a silvery-white metal. Having high chemical activity, the element is capable of forming many compounds of different classes. The element is valuable for technical and industrial chemical syntheses. The metal is widespread in the earth's crust: its share is about 1.5%. Calcium belongs to the group of alkaline earth metals: when dissolved in water, it produces alkalis, but in nature it occurs in the form of multiple minerals and. Sea water contains calcium in high concentrations (400 mg/l).

Pure sodium

The characteristics of calcium depend on the structure of its crystal lattice. This element has two types: cubic face-centric and volume-centric. The type of bond in the molecule is metallic.

Natural sources of calcium:

• apatites;
• alabaster;
• gypsum;
• calcite;
• fluorite;
• dolomite.

Physical properties of calcium and methods of obtaining the metal

Under normal conditions, calcium is in a solid state of aggregation. The metal melts at 842 °C. Calcium is a good electrical and thermal conductor. When heated, it first turns into a liquid and then into a vapor state and loses its metallic properties. The metal is very soft and can be cut with a knife. Boils at 1484 °C.

Under pressure, calcium loses its metallic properties and electrical conductivity. But then the metallic properties are restored and the properties of a superconductor appear, several times higher in their performance than the others.

For a long time it was not possible to obtain calcium without impurities: due to its high chemical activity, this element does not occur in nature in its pure form. The element was discovered at the beginning of the 19th century. Calcium as a metal was first synthesized by the British chemist Humphry Davy. The scientist discovered the peculiarities of the interaction of melts of solid minerals and salts with electric current. Nowadays, electrolysis of calcium salts (a mixture of calcium and potassium chlorides, a mixture of fluoride and calcium chloride) remains the most relevant method for producing metal. Calcium is also extracted from its oxide using aluminothermy, a common method in metallurgy.

Chemical properties of calcium

Calcium is an active metal that enters into many interactions. Under normal conditions, it easily reacts, forming the corresponding binary compounds: with oxygen, halogens. Click to learn more about calcium compounds. When heated, calcium reacts with nitrogen, hydrogen, carbon, silicon, boron, phosphorus, sulfur and other substances. In the open air, it instantly interacts with oxygen and carbon dioxide, and therefore becomes covered with a gray coating.

Reacts violently with acids and sometimes ignites. In salts, calcium exhibits interesting properties. For example, cave stalactites and stalagmites are calcium carbonate, gradually formed from water, carbon dioxide and bicarbonate as a result of processes within groundwater.

Due to its high activity in its normal state, calcium is stored in laboratories in dark, sealed glass containers under a layer of paraffin or kerosene. A qualitative reaction to calcium ion is the coloring of the flame in a rich brick-red color.

Calcium turns flames red

The metal in the composition of compounds can be identified by insoluble precipitates of some salts of the element (fluoride, carbonate, sulfate, silicate, phosphate, sulfite).

Reaction of water with calcium

Calcium is stored in jars under a layer of protective liquid. To conduct a demonstration of how the reaction of water and calcium occurs, you cannot simply take out the metal and cut off the desired piece from it. It is easier to use calcium metal in the laboratory in the form of shavings.

If there are no metal shavings and there are only large chunks of calcium in the jar, you will need pliers or a hammer. The finished piece of calcium of the required size is placed in a flask or glass of water. Calcium shavings are placed in a bowl in a gauze bag.

Calcium sinks to the bottom, and the release of hydrogen begins (first at the place where the fresh fracture of the metal is located). Gradually, gas is released from the surface of the calcium. The process resembles violent boiling, and at the same time a precipitate of calcium hydroxide (slaked lime) is formed.

Lime slaking

A piece of calcium floats up, caught up in hydrogen bubbles. After about 30 seconds, the calcium dissolves and the water turns cloudy white due to the formation of a hydroxide suspension. If the reaction is carried out not in a beaker, but in a test tube, you can observe the release of heat: the test tube quickly becomes hot. The reaction of calcium with water does not end with a spectacular explosion, but the interaction of the two substances proceeds vigorously and looks spectacular. The experience is safe.

If the bag with the remaining calcium is removed from the water and held in air, then after some time, as a result of the ongoing reaction, strong heating will occur and the remaining calcium in the gauze will boil. If part of the cloudy solution is filtered through a funnel into a glass, then when carbon monoxide CO₂ is passed through the solution, a precipitate will form. This does not require carbon dioxide - you can blow exhaled air into the solution through a glass tube.

Calcium compounds.

Sao– calcium oxide or quicklime, obtained by the decomposition of limestone: CaCO 3 = CaO + CO 2 is an oxide of an alkaline earth metal, so it actively interacts with water: CaO + H 2 O = Ca (OH) 2

Ca(OH) 2 – calcium hydroxide or slaked lime, therefore the reaction CaO + H 2 O = Ca(OH) 2 is called slaking of lime. If the solution is filtered, the result is lime water - this is an alkali solution, so it changes the color of phenolphthalein to crimson.

Slaked lime is widely used in construction. Its mixture with sand and water is a good binding material. Under the influence of carbon dioxide, the mixture hardens Ca(OH) 2 + CO 2 = CaCO3 + H 2 O.

At the same time, part of the sand and mixture turns into silicate Ca(OH) 2 + SiO 2 = CaSiO 3 + H 2 O.

The equations Ca (OH) 2 + CO 2 = CaCO 2 + H 2 O and CaCO 3 + H 2 O + CO 2 = Ca (HCO 3) 2 play a large role in nature and in shaping the appearance of our planet. Carbon dioxide in the form of a sculptor and architect creates underground palaces in the strata of carbonate rocks. It is capable of moving hundreds and thousands of tons of limestone underground. Through cracks in rocks, water containing carbon dioxide dissolved in it enters the limestone layer, forming cavities - caster caves. Calcium bicarbonate exists only in solution. Groundwater moves in the earth's crust, evaporating water under suitable conditions: Ca(HCO3) 2 = CaCO3 + H2O + CO 2 , This is how stalactites and stalagmites are formed, the formation scheme of which was proposed by the famous geochemist A.E. Fersman. There are a lot of castrum caves in Crimea. Science studies them speleology.

Calcium carbonate used in construction CaCO3- chalk, limestone, marble. You have all seen our railway station: it is decorated with white marble brought from abroad.

experience: blow through a tube into a solution of lime water, it becomes cloudy .

Ca(OH) 2 + CO 2 = CaCO 3 + N 2 ABOUT

Acetic acid is added to the formed precipitate, boiling is observed, because carbon dioxide is released.

CaCO 3 +2CH 3 COOH = Ca(CH 3 SOO) 2 +H 2 O + CO 2

THE TALE OF THE CARBONATE BROTHERS.

Three brothers live on earth
From the Carbonate family.
The older brother is a handsome MARBLE,
Glorious in the name of Karara,
An excellent architect. He
Built Rome and the Parthenon.
Everyone knows LIMESTONE,
That's why it's named like that.
Famous for his work
Building a house behind the house.
Both capable and able
Little soft brother MEL.
Look how he draws,
This CaCO 3!
Brothers love to frolic
Heat in a hot oven,
CaO and CO 2 are then formed.
This is carbon dioxide
Each of you is familiar with him,
We exhale it.
Well, this is SaO -
Hot-burnt quicklime.
Add water to it,
Mix thoroughly
So that there is no trouble,
We protect our hands
Well-kneaded LIME, but SLASHED!
Lime milk
The walls are whitewashed easily.
The bright house became cheerful,
Turning lime into chalk.
Hocus Pocus for the People:
You just have to blow through the water,
How easy it is
Turned into milk!
And now it's pretty clever
I get soda:
Milk plus vinegar. Ay!
Foam is pouring over the edge!
Everything is in worries, everything is in work
From dawn to dawn -
These brothers Carbonates,
These CaCO 3!

Repetition: CaO– calcium oxide, quicklime;
Ca(OH) 2 – calcium hydroxide (slaked lime, lime water, milk of lime depending on the concentration of the solution).
The general thing is the same chemical formula Ca(OH) 2. Difference: lime water is a transparent saturated solution of Ca(OH) 2, and milk of lime is a white suspension of Ca(OH) 2 in water.
CaCl 2 - calcium chloride, calcium chloride;
CaCO 3 – calcium carbonate, chalk, shell marble, limestone.
L/R: collections. Next, we demonstrate a collection of minerals available in the school laboratory: limestone, chalk, marble, shell rock.
CaS0 4 ∙ 2H 2 0 - calcium sulfate crystal hydrate, gypsum;
CaCO 3 - calcite, calcium carbonate is part of many minerals that cover 30 million km 2 on earth.

The most important of these minerals is limestone. Shell rocks, limestones of organic origin. It is used in the production of cement, calcium carbide, soda, all types of lime, and in metallurgy. Limestone is the basis of the construction industry; many building materials are made from it.

Chalk It's not just tooth powder and school chalk. It is also a valuable additive in the production of paper (coated - top quality) and rubber; in construction and renovation of buildings - as whitewash.

Marble is a dense crystalline rock. There is a colored one - white, but most often various impurities color it in different colors. Pure white marble is rare and is mainly used by sculptors (statues by Michelangelo, Rodin. In construction, colored marble is used as a facing material (Moscow Metro) or even as the main building material of palaces (Taj Mahal).

In the world of interesting things “Taj Mahal MAUSOLEUM”

Shah Jahan of the Great Mughal dynasty kept almost all of Asia in fear and obedience. In 1629, Mumzat Mahal, Shah Jahan's beloved wife, died at the age of 39 during childbirth on a campaign (this was their 14th child, all of them boys). She was unusually beautiful, bright, smart, the emperor obeyed her in everything. Before her death, she asked her husband to build a tomb, take care of the children, and not marry. The saddened king sent his envoys to all the big cities, the capitals of neighboring states - to Bukhara, Samarkand, Baghdad, Damascus, to find and invite the best craftsmen - in memory of his wife, the king decided to erect the best building in the world. At the same time, messengers sent plans for all the best buildings in Asia and the best building materials to Agra (India). They even brought malachite from Russia and the Urals. The chief masons came from Delhi and Kandahar; architects - from Istanbul, Samarkand; decorators - from Bukhara; gardeners - from Bengal; the artists were from Damascus and Baghdad, and the well-known master Ustad-Isa was in charge.

Together, over 25 years, a chalk marble structure was built surrounded by green gardens, blue fountains and a red sandstone mosque. 20,000 slaves erected this miracle of 75 m (25-story building). Nearby I wanted to build a second mausoleum of black marble for myself, but I didn’t have time. He was overthrown from the throne by his own son (the 2nd, and he also killed all his brothers).

The ruler and master of Agra spent the last years of his life looking out of the narrow window of his prison. For 7 years my father admired his creation. When the father went blind, the son made him a system of mirrors so that the father could admire the mausoleum. He was buried in the Taj Mahal, next to his Mumtaz.

Those entering the mausoleum see cenotaphs - false tombs. The eternal resting places of the Great Khan and his wife are located downstairs in the basement. Everything there is encrusted with precious stones that glow as if they were alive, and the branches of fairy-tale trees, intertwined with flowers, adorn the walls of the tomb in intricate patterns. Crafted by the best carvers, turquoise-blue lapis lazuli, green-black jades and red amethysts celebrate the love of Shah Jahal and Mumzat Mahal.

Every day tourists rush to Agra, wanting to see the true wonder of the world - the Taj Mahal mausoleum, as if floating above the ground.

CaCO 3 is a building material for the exoskeleton of mollusks, corals, shells, etc., and egg shells. (illustrations or Animals of the coral biocenosis” and display of a collection of sea corals, sponges, shell rock).

Calcium I Calcium (Ca)

chemical element of group II of the periodic system of chemical elements D.I. Mendeleev; belongs to alkaline earth metals and has high biological activity.

The atomic number of calcium is 20, the atomic mass is 40.08. Six stable isotopes of carbon with mass numbers 40, 42, 43, 44, 46, and 48 have been discovered in nature.

Calcium is chemically active, found in nature in the form of compounds - silicates (for example, asbestos), carbonates (limestone, marble, chalk, calcite, aragonite), sulfates (gypsum and anhydrite), phosphorite, dolomite, etc. It is the main structural element of bone tissue (see bone) , an important component of the blood coagulation system (blood clotting) , a necessary element of human food that maintains the homeostatic ratio of electrolytes in the internal environment of the body.

Among the most important functions in a living organism is its participation in the work of many enzyme systems (including those supporting muscles) in the transmission of nerve impulses, in the reaction of muscles to the nervous one and in changing the activity of hormones, which is realized with the participation of adenylate cyclase.

The human body contains 1-2 kg calcium (about 20 G by 1 kg body weight, in newborns about 9 g/kg). Of the total amount of calcium, 98-99% is found in bone and cartilage tissue in the form of carbonate, phosphate, compounds with chlorine, organic acids and other substances. The remaining amount is distributed in soft tissues (about 20 mg by 100 G tissue) and extracellular fluid. Blood plasma contains about 2.5 mmol/l calcium (9-11 mg/100 ml) in the form of two fractions: non-diffusing (complexes with proteins) and diffusing (ionized calcium and complexes with acids). Complexes with proteins are one of the forms of calcium storage. They account for 1/3 of the total amount of K. plasma. ionized K in blood is 1.33 mmol/l, complexes with phosphates, carbonates, citrates and anions of other organic acids - 0.3 mmol/l. There is an inverse relationship between ionized potassium and potassium phosphate in the blood plasma; however, with rickets, a decrease in the concentration of both ions is observed, and with hyperparathyroidism, an increase. In cells, the main part of phosphorus is associated with proteins and phospholipids of cell membranes and membranes of cell organelles. Regulation of transmembrane transfer of Ca 2+, in which specific Ca 2+-dependent is involved, is carried out by hormones of the thyroid gland (Thyroid gland) and parathyroid glands (Parathyroid glands) - parathyroid hormone and its antagonist calcitonin. The content of ionized K. in plasma is regulated by a complex mechanism, the components of which are (K. depot), liver (with bile), and calcitonin, as well as D (1,25-dioxy-cholecalciferol). increases the content of K. and reduces the content of K. phosphate in the blood, acting synergistically with vitamin D. It causes hypercalcemia by increasing the activity of osteoclasts and enhancing resorption, and increases the reabsorption of K. in the renal tubules. With hypocalcemia, parathyroid hormone increases significantly. , being an antagonist of parathyroid hormone, in case of hypercalcemia, it reduces the content of potassium in the blood and the number of osteoclasts, and increases the excretion of potassium phosphate by the kidneys. The pituitary gland also takes part in the regulation of calcium metabolism (see Pituitary hormones) , adrenal cortex (Adrenal glands) . Maintaining the homeostatic concentration of K. in the body is coordinated by the central nervous system. (mainly the hypothalamic-pituitary system (Hypothalamic-pituitary system)) and the autonomic nervous system.

K. plays an important role in the mechanism of muscle work (Muscular work) . It is a factor that allows muscle contraction: with an increase in the concentration of K ions in the myoplasm, K joins the regulatory protein, as a result of which it becomes able to interact with myosin; connecting, these two proteins form, and the muscle contracts. During the formation of actomyosin, ATP occurs, the chemical energy of which provides mechanical work and is partially dissipated as heat. The greatest skeletal contractility is observed at a calcium concentration of 10 -6 -10 -7 mole; when the concentration of K ions decreases (less than 10 -7 mole) muscle loses the ability to shorten and tense. K.'s effect on tissue is manifested in changes in their trophism, the intensity of redox processes, and in other reactions associated with the formation of energy. A change in the concentration of potassium in the fluid washing the nerve cell significantly affects its membranes for potassium ions and especially for sodium ions (see Biological membranes) , Moreover, a decrease in K level causes an increase in the permeability of the membrane for sodium ions and an increase in the excitability of the neuron. An increase in K concentration has a stabilizing effect on the nerve cell membrane. The role of K. in processes associated with the synthesis and release of mediators by nerve endings (Mediators) has been established. , providing synaptic transmission of nerve impulses.

The source of K. for the body is. An adult should receive 800-1100 per day from food mg calcium, children under 7 years old - about 1000 mg, 14-18 years old - 1400 mg, pregnant women - 1500 mg, nursing - 1800-2000 mg. Calcium contained in food products is represented mainly by phosphate, other compounds (carbonate, tartrate, K. oxalate and calcium-magnesium salt of phytic acid) - in much smaller quantities. The predominantly insoluble salts of potassium in the stomach are partially dissolved by gastric juice, then exposed to the action of bile acids, which convert it into an digestible form. K. occurs mainly in the proximal parts of the small intestine. an adult person absorbs less than half of the total amount of K supplied with food. K.'s absorption increases during growth during pregnancy and lactation. K.'s absorption is influenced by its relationship with fats, magnesium and phosphorus of food, vitamin D and other factors. With insufficient intake of fat, a deficiency of calcium salts of fatty acids, necessary for the formation of soluble complexes with bile acids, is created. Conversely, when eating excessively fatty foods, there are not enough bile acids to convert them into a soluble state, so a significant amount of unabsorbed calcium is excreted from the body. The optimal ratio of potassium and phosphorus in food ensures the mineralization of the bones of a growing organism. The regulator of this ratio is vitamin D, which explains the increased need for it in children.

The method of excretion of K. depends on the nature of the diet: if products with an acidic reaction (meat, bread, cereal dishes) predominate in the diet, the excretion of K. increases in the urine; products with an alkaline reaction (dairy products, fruits, vegetables) - in feces. Even a slight increase in its content in the blood leads to increased excretion of potassium in the urine.

Excess () K. or deficiency () of it in the body can be the cause or consequence of a number of pathological conditions. Thus, hypercalcemia occurs with excessive intake of calcium salts, increased absorption of calcium in the intestine, decreased excretion by the kidneys, increased consumption of vitamin D, and is manifested by growth retardation, anorexia, constipation, thirst, polyuria, muscle hypotonia, and hyperreflexia. With prolonged hypercalcemia, calcinosis develops , arterial, nephropathy. observed in a number of diseases accompanied by impaired mineral metabolism (see Rickets , Osteomalacia) , systemic bone sarcoidosis and multiple myeloma, Itsenko-Cushing's disease, acromegaly, hypothyroidism, malignant tumors, especially in the presence of bone metastases, hyperparathyroidism. Hypercalcemia is usually accompanied. Hypocalcemia, clinically manifested by tetany (Tetany) , may occur with hypoparathyroidism, idiopathic tetany (spasmophilia), diseases of the gastrointestinal tract, chronic renal failure, diabetes mellitus, Fanconi-Albertini syndrome, hypovitaminosis D. In case of K deficiency in the body, K drugs (calcium chloride, calcium gluconate, calcium lactate, calcium, calcium carbonate).

Determination of K. content in blood serum, urine and feces serves as an auxiliary diagnostic test for some diseases. Direct and indirect methods are used to study biological fluids. Indirect methods are based on preliminary precipitation of K. with ammonium oxalate, chloranilate or picrolenate and subsequent gravimetric, titrimetric or colorimetric determination. Direct methods include complexometric titration in the presence of ethylenediaminetetraacetate or ethylene glycoltetraacetate and metal indicators, for example murexide (Greenblatt-Hartman method), fluorexone, acid chromium dark blue, calcium, etc., colorimetric methods using alizarin, methylthymol blue, o-cresolphthalein complexone, glyokeal -bis-2-hydroxyanyl; fluorimetric methods; flame photometry method; atomic absorption spectrometry (the most accurate and sensitive method, allowing to determine up to 0.0001% calcium); method using ion-selective electrodes (allows you to determine the activity of calcium ions). The content of ionized calcium in blood serum can be determined using the data) of the concentration of total calcium and total protein using the empirical formula: percentage of calcium bound to protein = 8() + 2() + 3 G/100 ml.

Bibliography: Kostyuk P.G. Calcium and Cellular, M., 1986, bibliogr.; Laboratory methods of research in the clinic, ed. V.V. Menshikova, s. 59, 265, M., 1987; Regulation of calcium ions, ed. M.D. Kursky et al., Kyiv, 1977; Romanenko V.D. calcium metabolism, Kyiv, 1975, bibliogr.

II Calcium (Ca)

chemical element of group II of the periodic table D.I. Mendeleev; atomic number 20, atomic mass 40.08; has high biological activity; is an important component of the blood coagulation system; part of bone tissue; Various calcium compounds are used as medicines.

1. Small medical encyclopedia. - M.: Medical encyclopedia. 1991-96 2. First aid. - M.: Great Russian Encyclopedia. 1994 3. Encyclopedic Dictionary of Medical Terms. - M.: Soviet Encyclopedia. - 1982-1984.

Synonyms:

See what "Calcium" is in other dictionaries:

- (Ca) yellow shiny and viscous metal. Specific gravity 1.6. Dictionary of foreign words included in the Russian language. Pavlenkov F., 1907. CALCIUM (new Latin calcium, from Latin calx lime). Silver colored metal. Dictionary of foreign words,... ... Dictionary of foreign words of the Russian language

CALCIUM- CALCIUM, Calcium, chemical. element, symbol Ca, shiny, silvery-white crystalline metal. fracture, belonging to the group of alkaline earth metals. Ud. weight 1.53; at. V. 40.07; melting point 808°. Sa is one of the very... Great Medical Encyclopedia

- (Calcium), Ca, chemical element of group II of the periodic system, atomic number 20, atomic mass 40.08; refers to alkaline earth metals; melting point 842shC. Contained in the bone tissue of vertebrates, mollusk shells, and eggshells. Calcium... ... Modern encyclopedia

The metal is silvery-white, viscous, malleable, and quickly oxidizes in air. Melting rate pa 800-810°. Found in nature in the form of various salts that form deposits of chalk, limestone, marble, phosphorites, apatites, gypsum, etc. dor... ... Technical railway dictionary

- (Latin Calcium) Ca, a chemical element of group II of the periodic table, atomic number 20, atomic mass 40.078, belongs to the alkaline earth metals. Name from Latin calx, genitive calcis lime. Silvery white metal,... ... Big Encyclopedic Dictionary

- (symbol Ca), a widespread silvery-white metal from the ALKALINE EARTH group, first isolated in 1808. Found in many rocks and minerals, especially limestone and gypsum, as well as bones. In the body it promotes... Scientific and technical encyclopedic dictionary

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).

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 location. For example, stalactites and stalagmites form in caves.

Soluble calcium and magnesium salts cause overall water hardness. If they are present in water in small quantities, then the water is called soft. With a high content of these salts (100 - 200 mg of calcium salts in 1 liter in terms of ions), the water is considered hard. In such water, soap does not foam well, since calcium and magnesium salts form insoluble compounds with it. 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.