Barium reacts with water. Barium

Barium is an alkaline earth metal occupying position 56 in the periodic table of chemical elements. The name of the substance translated from ancient Greek means “heavy”.

Characteristics of barium

The metal has an atomic mass of 137 g/mmol and a density of about 3.7 g/cm 3 . It is very light and soft - its maximum hardness on the Mohs scale is 3 points. In the case of mercury impurities, the fragility of barium increases significantly.

The metal has a light silver-gray color. However, the metal is also famous for its green color, which is acquired as a result of a chemical reaction involving salts of the element (for example, barium sulfate). If we dip a glass rod into barium and bring it to an open flame, we will see a green flame. This method makes it possible to clearly determine even the minimum content of heavy metal impurities.

The crystal lattice of barium, which can be observed even outside laboratory conditions, has a cubic shape. It is worth noting that finding pure barium in nature is also appropriate. Today, there are two known modifications of the metal, one of which is resistant to temperature increases up to 365 0 C, and the other is able to withstand temperatures in the range of 375-710 0 C. The boiling point of barium is 1696 0 C.

Barium, along with other alkaline earth metals, exhibits chemical activity. It occupies a middle position in the group, leaving behind strontium and calcium, which can be stored in the open air, but this cannot be said about barium. An excellent medium for storing metal is paraffin oil, into which barium or petroleum ether is directly immersed.

Barium reacts with oxygen, however, as a result of the reaction, its shine is lost, after which the metal first acquires a yellowish tint, then becomes brown and eventually acquires a gray color. This is the appearance inherent in barium oxide. When the atmosphere heats up, barium becomes explosive.

The 56th element of the periodic table of Mendeleev also interacts with water, resulting in a reaction that is the opposite of the reaction with oxygen. In this case, the liquid is subject to decomposition. This reaction is carried out exclusively by pure metal, after which it becomes barium hydroxide. If metal salts come into contact with the aqueous environment, then we will not see any reaction, since nothing will happen. For example, its chloride is insoluble in water and an active reaction can only be observed when interacting with an acidic environment.

The metal easily reacts with hydrogen, but for this it is necessary to create certain conditions, namely, an increase in temperature. In this case, the output is barium hydride. Under conditions of increasing temperature, element 56 also reacts with ammonia, resulting in the formation of nitride. If the temperature is raised further, cyanide can be produced.

Barium solution has a characteristic blue color, which is obtained as a result of reaction with ammonia in a liquid aggregate state. If a platinum catalyst is added, barium amide is formed. However, the scope of application of this substance is far from wide - it is used exclusively as a reagent.

Table 1. Properties of barium

Characteristic	Meaning
Properties of the atom
Name, symbol, number	Barium / Barium (Ba), 56
Atomic mass (molar mass)	137.327(7) a. e.m. (g/mol)
Electronic configuration	6s2
Atomic radius	222 pm
Chemical properties
Covalent radius	198 pm
Ion radius	(+2e) 134 pm
Electronegativity	0.89 (Pauling scale)
Electrode potential	-2,906
Oxidation states	2
Ionization energy (first electron)	502.5 (5.21) kJ/mol (eV)
Thermodynamic properties of a simple substance
Density (at normal conditions)	3.5 g/cm³
Melting temperature	1 002 K
Boiling temperature	1 910 K
Ud. heat of fusion	7.66 kJ/mol
Ud. heat of vaporization	142.0 kJ/mol
Molar heat capacity	28.1 J/(K mol)
Molar volume	39.0 cm³/mol
Crystal lattice of a simple substance
Lattice structure	cubic body-centered
Lattice parameters	5.020 Å
Other characteristics
Thermal conductivity	(300 K) (18.4) W/(m K)
CAS number	7440-39-3

Obtaining barium

The metal was first obtained in the second half of the 18th century (in 1774) by chemists Karl Scheele and Johan Hahn. Then a metal oxide was obtained. A few years later, Humphry Davy succeeded in producing a metal amalgam by electrolysis of wet barium hydroxide with a mercury cathode, which he subjected to heat and evaporated the mercury, thus obtaining barium metal.

The production of barium metal in modern laboratory conditions is carried out in several ways associated with the atmosphere. Barium separation is carried out in a vacuum due to an overly active reaction that is released when barium reacts with oxygen.

Barium oxide and chloride are obtained by metallothermic reduction under conditions of increasing temperature to 1200 0 C.

Also, pure metal can be separated from its hydride and nitride using thermal decomposition. Potassium is obtained in the same way. To carry out this process, special capsules with complete sealing are required, as well as the presence of quartz or porcelain. It is also possible to obtain barium by electrolysis, by which the element can be isolated from molten barium chloride with a mercury cathode.

Applications of barium

Considering all the properties that the 56th element of the periodic table has, barium is a fairly popular metal. So, it is used:

1. In the manufacture of vacuum electronic devices. In this case, barium metal, or its alloy with aluminum, is used as a gas absorber. And its oxide in the composition of a solid solution of oxides of other alkaline earth metals is used as an active layer of indirect channel cathodes.
2. As a material that can resist corrosion. For this purpose, metal along with zirconium is added to liquid metal coolants, which can significantly reduce the aggressive effect on pipelines. This use of barium found its place in the metallurgical industry.
3. Barium can act as a ferroelectric and piezoelectric. It is appropriate to use barium titanate, which acts as a dielectric during the manufacture of ceramic capacitors, as well as a material used in piezoelectric microphones and piezoceramic emitters.
4. In optical instruments. Barium fluoride is used, which has the form of single crystals.
5. As an integral element of pyrotechnics. Metal peroxide is used as an oxidizing agent. Barium nitrate and chlorate act as substances that give the flame a certain color (green).
6. In nuclear-hydrogen energy. Here, barium chromate is actively used in the production of hydrogen and oxygen using the thermochemical method.
7. In nuclear energy. The metal oxide is an integral component of the process of making the certain grade of glass that coats the uranium rods.
8. As a chemical current source. Several barium compounds can be used in this case: fluoride, oxide and sulfate. The first compound is used in solid-state fluorine batteries as a component of the fluoride electrolyte. The oxide has found its place in high-power copper oxide batteries as a component of the active mass. And the latter substance is used as an expander of the active mass of the negative electrode during the production of lead-acid batteries.
9. In medicine. Barium sulfate is an insoluble substance that is completely non-toxic. In this regard, it is used as a radiopaque material during studies of the gastrointestinal tract.

Table 2. Applications of barium

Application area	Mode of application
Vacuum electronic devices	Metal barium, often in an alloy with aluminum, is used as a gas absorber (getter) in high-vacuum electronic devices. Barium oxide, as part of a solid solution of oxides of other alkaline earth metals - calcium and strontium (CaO, SrO), is used as the active layer of indirectly heated cathodes.
Anti-corrosion material	Barium is added together with zirconium to liquid metal coolants (alloys of sodium, potassium, rubidium, lithium, cesium) to reduce the aggressiveness of the latter to pipelines and in metallurgy.
Ferro- and piezoelectric	Barium titanate is used as a dielectric in the manufacture of ceramic capacitors, and as a material for piezoelectric microphones and piezoceramic emitters.
Optics	Barium fluoride is used in the form of single crystals in optics (lenses, prisms).
Pyrotechnics	Barium peroxide is used for pyrotechnics and as an oxidizing agent. Barium nitrate and barium chlorate are used in pyrotechnics to color flames (green fire).
Nuclear-hydrogen energy	Barium chromate is used in the production of hydrogen and oxygen by thermochemical method (Oak Ridge cycle, USA).
High temperature superconductivity	Barium peroxide, together with oxides of copper and rare earth metals, is used to synthesize superconducting ceramics operating at liquid nitrogen temperatures and above.
Nuclear energy	Barium oxide is used to melt a special type of glass - used to coat uranium rods. One of the widespread types of such glasses has the following composition - (phosphorus oxide - 61%, BaO - 32%, aluminum oxide - 1.5%, sodium oxide - 5.5%). Barium phosphate is also used in glass melting for the nuclear industry.
Chemical current sources	Barium fluoride is used in solid-state fluorion batteries as a component of the fluoride electrolyte. Barium oxide is used in high-power copper oxide batteries as an active mass component (barium oxide-copper oxide). Barium sulfate is used as a negative electrode active mass extender in the production of lead-acid batteries .
Application in medicine	Barium sulfate, insoluble and non-toxic, is used as a radiocontrast agent in medical examinations of the gastrointestinal tract.

BARIUM, Ba (Latin Baryum, from the Greek barys - heavy * a. barium; n. Barium; f. barium; i. bario), - a chemical element of the main subgroup of group 11 of the Mendeleev periodic system of elements, atomic number 56, atomic mass 137.33. Natural barium consists of a mixture of seven stable isotopes; 138 Va (71.66%) predominates. Barium was discovered in 1774 by the Swedish chemist K. Scheele in the form of BaO. Metallic barium was first obtained by the English chemist H. Davy in 1808.

Obtaining barium

Barium metal is obtained by thermal reduction in a vacuum at 1100-1200°C of barium oxide powder. Barium is used in alloys - with lead (printing and antifriction alloys), aluminum and (gas absorbers in vacuum installations). Its artificial radioactive isotopes are widely used.

Applications of barium

Barium and its compounds are added to materials intended to protect against radioactive and x-ray radiation. Barium compounds are widely used: oxide, peroxide and hydroxide (to produce hydrogen peroxide), nitride (in pyrotechnics), sulfate (as a contrast agent in radiology, research), chromate and manganate (in the manufacture of paints), titanate (one of the most important ferroelectrics) , sulfide (in the leather industry), etc.

In 1808, Davy Humphrey obtained barium in the form of an amalgam by electrolysis of its compounds.

Receipt:

In nature, it forms the minerals barite BaSO 4 and witherite BaCO 3 . Prepared by aluminothermy or azide decomposition:
3BaO+2Al=Al 2 O 3 +3Ba
Ba(N 3) 2 =Ba+3N 2

Physical properties:

A silvery-white metal with a higher melting and boiling point and greater density than the alkali metals. Very soft. Melt = 727°C.

Chemical properties:

Barium is the strongest reducing agent. In air, it quickly becomes covered with a film of oxide, peroxide and barium nitride, and ignites when heated or simply crushed. Reacts vigorously with halogens and, when heated, with hydrogen and sulfur.
Barium reacts vigorously with water and acids. They are stored, like alkali metals, in kerosene.
In compounds it exhibits an oxidation state of +2.

The most important connections:

Barium oxide. A solid that reacts vigorously with water to form a hydroxide. Absorbs carbon dioxide, turning into carbonate. When heated to 500°C, it reacts with oxygen to form peroxide
Barium peroxide BaO 2, white substance, poorly soluble, oxidizing agent. Used in pyrotechnics, to produce hydrogen peroxide, bleach.
Barium hydroxide Ba(OH) 2, Ba(OH) 2 octahydrate *8H 2 O, colorless. crystal, alkali. Used for detection of sulfate and carbonate ions, for purification of vegetable and animal fats.
Barium salts colorless crystals substances. Soluble salts are highly poisonous.
Chloride barium is obtained by reacting barium sulfate with coal and calcium chloride at 800°C - 1100°C. Reagent for sulfate ion. used in the leather industry.
Nitrate barium, barium nitrate, green component of pyrotechnic compositions. When heated, it decomposes to form barium oxide.
Sulfate barium is practically insoluble in water and acids, therefore it is low-toxic. used for bleaching paper, for fluoroscopy, barite concrete filler (protection against radioactive radiation).

Application:

Barium metal is used as a component of a number of alloys and a deoxidizing agent in the production of copper and lead. Soluble barium salts are poisonous, MPC 0.5 mg/m 3 . See also:
S.I. Venetsky About rare and scattered. Stories about metals.

chemical element of the 2nd group of the periodic table, atomic number 56, relative atomic mass 137.33. Located in the sixth period between cesium and lanthanum. Natural barium consists of seven stable isotopes with mass numbers 130(0.101%), 132(0.097%), 134(2.42%), 135(6.59%), 136(7.81%), 137(11. 32%) and 138 (71.66%). Barium in most chemical compounds exhibits a maximum oxidation state of +2, but can also have a zero oxidation state. In nature, barium occurs only in the divalent state.History of discovery. In 1602, Casciarolo (Bolognese shoemaker and alchemist) picked up a stone in the surrounding mountains that was so heavy that Casciarolo suspected it was gold. Trying to isolate gold from a stone, the alchemist calcined it with coal. Although it was not possible to isolate gold, the experiment brought clearly encouraging results: the cooled calcination product glowed reddish in the dark. The news of such an unusual find created a real sensation in the alchemical community and an unusual mineral, which received a number of names - sun stone ( Lapis solaris ), Bolognese stone ( Lapis boloniensis ), Bologna phosphorus (Phosphorum Boloniensis) became a participant in various experiments. But time passed, and gold did not even think of standing out, so interest in the new mineral gradually disappeared, and for a long time it was considered a modified form of gypsum or lime. Only a century and a half later, in 1774, the famous Swedish chemists Karl Scheele and Johan Hahn carefully studied the “Bologna stone” and found that it contained some kind of “heavy earth”. Later, in 1779, Guiton de Morveau named this "land" Barot ( barote ) from the Greek word " barue » heavy, and later changed the name to barite ( baryte ). Under this name, barium earth appeared in chemistry textbooks of the late 18th and early 19th centuries. So, for example, in the textbook by A.L. Lavoisier (1789) barite is included in the list of salt-forming earthy simple bodies, and another name for barite is given - “heavy earth” ( terre pesante , lat. terra ponderosa). The still unknown metal contained in the mineral began to be called barium (Latin Barium ). In Russian literature of the 19th century. The names barite and barium were also used. The next known barium mineral was natural barium carbonate, discovered in 1782 by Withering and later named witherite in his honor. Barium metal was first prepared by Englishman Humphry Davy in 1808 by electrolysis of wet barium hydroxide with a mercury cathode and subsequent evaporation of mercury from barium amalgam. It should be noted that in the same 1808, somewhat earlier than Davy, barium amalgam was obtained by the Swedish chemist Jens Berzelius. Despite its name, barium turned out to be a relatively light metal with a density of 3.78 g/cm 3, so in 1816 the English chemist Clark proposed rejecting the name “barium” on the grounds that if barium earth (barium oxide) is indeed heavier than other earths (oxides), then the metal, on the contrary, is lighter than other metals. Clark wanted to name this element plutonium in honor of the ancient Roman god, ruler of the underground kingdom of Pluto, but this proposal did not meet with support from other scientists and the light metal continued to be called “heavy.”Barium in nature. The earth's crust contains 0.065% barium, it occurs in the form of sulfate, carbonate, silicates and aluminosilicates. The main minerals of barium are the already mentioned barite (barium sulfate), also called heavy or Persian spar, and witherite (barium carbonate). The world's mineral resources of barite were estimated in 1999 at 2 billion tons, a significant part of them concentrated in China (about 1 billion tons) and Kazakhstan (0.5 billion tons). There are large reserves of barite in the USA, India, Turkey, Morocco and Mexico. Russian barite resources are estimated at 10 million tons, its production is carried out at three main deposits located in Khakassia, Kemerovo and Chelyabinsk regions. The total annual production of barite in the world is about 7 million tons, Russia produces 5 thousand tons and imports 25 thousand tons of barite per year.Receipt. The main raw materials for the production of barium and its compounds are barite and, less commonly, witherite. By reducing these minerals with coal, coke or natural gas, barium sulfide and barium oxide are obtained, respectively:BaSO 4 + 4C = BaS + 4CO

BaSO 4 + 2CH 4 = BaS + 2C + 4H 2 O

BaCO 3 + C = BaO + 2CO

Barium metal is obtained by reducing it with aluminum oxide.

BaO + 2 Al = 3 Ba + Al 2 O 3

For the first time this process

cc carried out by the Russian physical chemist N.N. Beketov. This is how he described his experiments: “I took anhydrous barium oxide and, adding to it a certain amount of barium chloride, like flux, I put this mixture along with pieces of clay (aluminum) in a carbon crucible and heated it for several hours. After cooling the crucible, I found in it a metal alloy of a completely different type and physical properties than clay. This alloy has a coarse-crystalline structure, is very brittle, a fresh fracture has a faint yellowish sheen; analysis showed that at 100 hours it consists of 33.3 barium and 66.7 clay, or, otherwise, for one part of barium it contained two parts of clay...” Nowadays the reduction process with aluminum is carried out in a vacuum at temperatures from 1100 to 1250° C , while the resulting barium evaporates and condenses on the cooler parts of the reactor.

In addition, barium can be obtained by electrolysis of a molten mixture of barium and calcium chlorides.

Simple substance. Barium is a silvery-white malleable metal that shatters when struck sharply. Melting point 727° C, boiling point 1637° C, density 3.780 g/cm 3 . At normal pressure it exists in two allotropic modifications: up to 375° C stable a - Ba with a cubic body-centered lattice, stable above 375° C b-Ba . At elevated pressure, a hexagonal modification is formed. Metal barium has high chemical activity; it oxidizes intensively in air, forming a film containing BaO, BaO 2 and Ba 3 N 2, with slight heating or impact, it ignites.2Ba + O 2 = 2BaO; Ba + O 2 = BaO 2; 3Ba + N 2 = Ba 3 N 2,Therefore, barium is stored under a layer of kerosene or paraffin. Barium reacts vigorously with water and acid solutions, forming barium hydroxide or the corresponding salts:Ba + 2H 2 O = Ba(OH) 2 + H 2

Ba + 2HCl = BaCl 2 + H 2

With halogens, barium forms halides, with hydrogen and nitrogen when heated, hydride and nitride, respectively.Ba + Cl 2 = BaCl 2; Ba + H 2 = BaH 2Barium metal dissolves in liquid ammonia to form a dark blue solution, from which ammonia can be isolated Ba(NH 3) 6 crystals with a golden luster, easily decomposing with the release of ammonia. In this compound, barium has zero oxidation state.Application in industry and science. The use of barium metal is very limited due to its high chemical reactivity; barium compounds are used much more widely. Barium alloy with aluminum Alba alloy containing 56% Ba the basis of getters (absorbers of residual gases in vacuum technology). To obtain the getter itself, barium is evaporated from the alloy by heating it in a evacuated flask of the device, as a result of which a “barium mirror” is formed on the cold parts of the flask. In small quantities, barium is used in metallurgy to purify molten copper and lead from impurities of sulfur, oxygen and nitrogen. Barium is added to printing and antifriction alloys; an alloy of barium and nickel is used to make parts for radio tubes and spark plug electrodes in carburetor engines. In addition, there are non-standard uses of barium. One of them is the creation of artificial comets: barium vapor released from a spacecraft is easily ionized by solar rays and turns into a bright plasma cloud. The first artificial comet was created in 1959 during the flight of the Soviet automatic interplanetary station Luna-1. In the early 1970s, German and American physicists, conducting research on the Earth's electromagnetic field, released 15 kilograms of tiny barium powder over Colombia. The resulting plasma cloud stretched along the magnetic field lines, making it possible to clarify their position. In 1979, jets of barium particles were used to study the aurora.Barium compounds. Divalent barium compounds are of greatest practical interest.

Barium oxide(

BaO ): intermediate product in the production of barium refractory (melting point about 2020° C ) white powder, reacts with water to form barium hydroxide, absorbs carbon dioxide from the air, turning into carbonate:BaO + H 2 O = Ba(OH) 2; BaO + CO 2 = BaCO 3Heated in air at a temperature of 500600° C , barium oxide reacts with oxygen to form peroxide, which upon further heating to 700° C goes back into the oxide, splitting off oxygen:2BaO + O 2 = 2BaO 2 ; 2BaO2 = 2BaO + O2This is how oxygen was obtained until the end of the 19th century, until a method for releasing oxygen by distilling liquid air was developed.

In the laboratory, barium oxide can be prepared by calcining barium nitrate:

2Ba(NO3)2 = 2BaO + 4NO2 + O2Now barium oxide is used as a water-removing agent, to produce barium peroxide and to make ceramic magnets from barium ferrate (for this, a mixture of barium and iron oxide powders is sintered under a press in a strong magnetic field), but the main use of barium oxide is the manufacture of thermionic cathodes. In 1903, the young German scientist Wehnelt tested the law of the emission of electrons by solids, discovered shortly before by the English physicist Richardson. The first of the experiments with platinum wire completely confirmed the law, but the control experiment failed: the flow of electrons sharply exceeded the expected one. Since the properties of the metal could not change, Wehnelt assumed that there was some kind of impurity on the surface of the platinum. After testing possible surface contaminants, he became convinced that the additional electrons were emitted by barium oxide, which was part of the lubricant of the vacuum pump used in the experiment. However, the scientific world did not immediately recognize this discovery, since its observation could not be reproduced. Only almost a quarter of a century later, the Englishman Kohler showed that in order to exhibit high thermionic emission, barium oxide must be heated at very low oxygen pressures. This phenomenon could only be explained in 1935. The German scientist Pohl suggested that electrons are emitted by a small impurity of barium in the oxide: at low pressures, part of the oxygen evaporates from the oxide, and the remaining barium is easily ionized to form free electrons, which leave the crystal when heated:2BaO = 2Ba + O 2 ; Ba = Ba 2+ + 2 e The correctness of this hypothesis was finally established in the late 1950s by Soviet chemists A. Bundel and P. Kovtun, who measured the concentration of barium impurity in the oxide and compared it with the flux of thermionic electron emission. Now barium oxide is the active part of most thermionic cathodes. For example, a beam of electrons that forms an image on a TV screen or computer monitor is emitted by barium oxide.

Barium hydroxide, octahydrate(

Ba(OH)2 8 H2O ). White powder, highly soluble in hot water (more than 50% at 80° C ), worse in cold (3.7% at 20° C ). Melting point of octahydrate 78° C , when heated to 130° C it becomes anhydrous Ba(OH ) 2 . Barium hydroxide is produced by dissolving the oxide in hot water or by heating barium sulfide in a stream of superheated steam. Barium hydroxide easily reacts with carbon dioxide, so its aqueous solution, called “barite water,” is used in analytical chemistry as a reagent for CO 2. In addition, “barite water” serves as a reagent for sulfate and carbonate ions. Barium hydroxide is used to remove sulfate ions from plant and animal oils and industrial solutions, to obtain rubidium and cesium hydroxides, as a component of lubricants.

Barium carbonate(

BaCO 3). In nature, the mineral is witherite. White powder, insoluble in water, soluble in strong acids (except sulfuric acid). When heated to 1000° C, it decomposes and releases CO 2: BaCO 3 = BaO + CO 2

Barium carbonate is added to glass to increase its refractive index and is added to enamels and glazes.

Barium sulfate(

BaSO 4). In nature barite (heavy or Persian spar) the main mineral of barium white powder (melting point about 1680° C ), practically insoluble in water (2.2 mg/l at 18° C ), dissolves slowly in concentrated sulfuric acid.

The production of paints has long been associated with barium sulfate. True, at first its use was of a criminal nature: crushed barite was mixed with lead white, which significantly reduced the cost of the final product and, at the same time, deteriorated the quality of the paint. However, such modified whites were sold at the same price as regular whites, generating significant profits for dye plant owners. Back in 1859, the Department of Manufactures and Domestic Trade received information about the fraudulent machinations of Yaroslavl factory owners who added heavy spar to lead white, which “deceives consumers about the true quality of the product, and a request was also received to prohibit the said manufacturers from using spar in the production of lead white.” " But these complaints came to nothing. Suffice it to say that in 1882 a spar plant was founded in Yaroslavl, which in 1885 produced 50 thousand pounds of crushed heavy spar. In the early 1890s, D.I. Mendeleev wrote: “...Barite is mixed into the mixture of white at many factories, since white brought from abroad contains this mixture to reduce the price.”

Barium sulfate is part of lithopone, a non-toxic white paint with high hiding power, widely in demand on the market. To make lithopone, aqueous solutions of barium sulfide and zinc sulfate are mixed, during which an exchange reaction occurs and a mixture of fine-crystalline barium sulfate and zinc sulfide lithopone precipitates, and pure water remains in the solution.

BaS + ZnSO 4 = BaSO 4 Ї + ZnS Ї

In the production of expensive grades of paper, barium sulfate plays the role of a filler and weighting agent, making the paper whiter and denser; it is also used as a filler for rubber and ceramics.

More than 95% of the barite mined in the world is used to prepare working solutions for drilling deep wells.

Barium sulfate strongly absorbs x-rays and gamma rays. This property is widely used in medicine for diagnosing gastrointestinal diseases. To do this, the patient is given a suspension of barium sulfate in water or a mixture of it with semolina “barium porridge” to swallow and then x-rayed. Those parts of the digestive tract through which the “barium porridge” passes appear as dark spots in the picture. This way the doctor can get an idea of ​​the shape of the stomach and intestines and determine the location of the disease. Barium sulfate is also used to make barite concrete, used in the construction of nuclear power plants and nuclear plants to protect against penetrating radiation.

Barium sulfide(

BaS ). Intermediate product in the production of barium and its compounds. The commercial product is a gray friable powder, poorly soluble in water. Barium sulfide is used to produce lithopone, in the leather industry to remove hair from hides, and to produce pure hydrogen sulfide. BaS a component of many phosphors substances that glow after absorbing light energy. This is what Casciarolo obtained by calcining barite with coal. By itself, barium sulfide does not glow: it requires the addition of activating substances - salts of bismuth, lead and other metals.

Barium titanate(

BaTiO 3). One of the most industrially important compounds of barium white refractory (melting point 1616° C ) a crystalline substance insoluble in water. Barium titanate is obtained by fusing titanium dioxide with barium carbonate at a temperature of about 1300° C: BaCO 3 + TiO 2 = BaTiO 3 + CO 2

Barium titanate one of the best ferroelectrics ( cm. Also FERROELECTRICS), very valuable electrical materials. In 1944, Soviet physicist B.M. Vul discovered extraordinary ferroelectric abilities (very high dielectric constant) of barium titanate, which retained them in a wide temperature range - from almost absolute zero to +125°

C . This circumstance, as well as the great mechanical strength and moisture resistance of barium titanate, have contributed to its becoming one of the most important ferroelectrics, used, for example, in the manufacture of electrical capacitors. Barium titanate, like all ferroelectrics, also has piezoelectric properties: it changes its electrical characteristics under pressure. When exposed to an alternating electric field, oscillations occur in its crystals, and therefore they are used in piezoelements, radio circuits and automatic systems. Barium titanate was used in attempts to detect gravitational waves.Other barium compounds. Nitrate and chlorate (Ba(ClO 3) 2) barium an integral part of fireworks, the addition of these compounds gives the flame a bright green color. Barium peroxide is a component of ignition mixtures for aluminothermy. Tetracyanoplatinate( II) barium (Ba[Pt(CN ) 4 ]) glows under the influence of x-rays and gamma rays. In 1895, German physicist Wilhelm Roentgen, observing the glow of this substance, he suggested the existence of a new radiation, later called X-ray. Now tetracyanoplatinate ( II ) barium covers the luminous screens of devices. Barium thiosulfate ( BaS2O 3) gives the colorless varnish a pearly tint, and by mixing it with glue, you can achieve a complete imitation of mother-of-pearl.Toxicology of barium compounds. All soluble barium salts are poisonous. Barium sulfate used in fluoroscopy is practically non-toxic. The lethal dose of barium chloride is 0.80.9 g, barium carbonate is 24 g. When poisonous barium compounds are ingested, a burning sensation in the mouth, pain in the stomach, salivation, nausea, vomiting, dizziness, muscle weakness, shortness of breath occur. , slow heart rate and drop in blood pressure. The main method of treating barium poisoning is gastric lavage and the use of laxatives.

The main sources of barium entering the human body are food (especially seafood) and drinking water. According to the recommendation of the World Health Organization, the barium content in drinking water should not exceed 0.7 mg/l; in Russia there are much more stringent standards of 0.1 mg/l.

Yuri Krutyakov

LITERATURE Figurovsky N.A. The history of the discovery of elements and the origin of their names. M., Nauka, 1970
Venetsky S.I. About rare and scattered. Tales of Metals. M.,neMetallurgy, 1980
Popular library of chemical elements. Under. ed.neI.V.Petryanova-Sokolova M., Science, 1983
Information and analytical review of the state and prospects of the global and domestic markets of non-ferrous, rare and precious metals. Issue 18. Barite. M., 2002

Group IIA contains only metals – Be (beryllium), Mg (magnesium), Ca (calcium), Sr (strontium), Ba (barium) and Ra (radium). The chemical properties of the first representative of this group, beryllium, differ most strongly from the chemical properties of the other elements of this group. Its chemical properties are in many ways even more similar to aluminum than to other Group IIA metals (so-called “diagonal similarity”). Magnesium, in its chemical properties, also differs markedly from Ca, Sr, Ba and Ra, but still has much more similar chemical properties with them than with beryllium. Due to the significant similarity in the chemical properties of calcium, strontium, barium and radium, they are combined into one family called alkaline earth metals.

All elements of group IIA belong to s-elements, i.e. contain all their valence electrons on s-sublevel Thus, the electronic configuration of the outer electronic layer of all chemical elements of this group has the form ns 2 , Where n– number of the period in which the element is located.

Due to the peculiarities of the electronic structure of group IIA metals, these elements, in addition to zero, can have only one single oxidation state equal to +2. Simple substances formed by elements of group IIA, when participating in any chemical reactions, are only capable of oxidation, i.e. donate electrons:

Me 0 – 2e — → Me +2

Calcium, strontium, barium and radium have extremely high chemical reactivity. The simple substances formed by them are very strong reducing agents. Magnesium is also a strong reducing agent. The reduction activity of metals obeys the general laws of the periodic law of D.I. Mendeleev and increases down the subgroup.

Interaction with simple substances

with oxygen

Without heating, beryllium and magnesium do not react with either atmospheric oxygen or pure oxygen due to the fact that they are covered with thin protective films consisting of BeO and MgO oxides, respectively. Their storage does not require any special methods of protection from air and moisture, unlike alkaline earth metals, which are stored under a layer of liquid inert to them, most often kerosene.

Be, Mg, Ca, Sr, when burned in oxygen, form oxides of the composition MeO, and Ba - a mixture of barium oxide (BaO) and barium peroxide (BaO 2):

2Mg + O2 = 2MgO

2Ca + O2 = 2CaO

2Ba + O 2 = 2BaO

Ba + O 2 = BaO 2

It should be noted that when alkaline earth metals and magnesium burn in air, a side reaction of these metals with air nitrogen also occurs, as a result of which, in addition to compounds of metals with oxygen, nitrides with the general formula Me 3 N 2 are also formed.

with halogens

Beryllium reacts with halogens only at high temperatures, and the rest of the Group IIA metals - already at room temperature:

Mg + I 2 = MgI 2 – Magnesium iodide

Ca + Br 2 = CaBr 2 – calcium bromide

Ba + Cl 2 = BaCl 2 – barium chloride

with non-metals of groups IV–VI

All metals of group IIA react when heated with all nonmetals of groups IV–VI, but depending on the position of the metal in the group, as well as the activity of the nonmetals, varying degrees of heating are required. Since beryllium is the most chemically inert among all group IIA metals, when carrying out its reactions with non-metals, significant use is required. O higher temperature.

It should be noted that the reaction of metals with carbon can form carbides of different natures. There are carbides that belong to methanides and are conventionally considered derivatives of methane, in which all hydrogen atoms are replaced by metal. They, like methane, contain carbon in the -4 oxidation state, and when they are hydrolyzed or interact with non-oxidizing acids, one of the products is methane. There is also another type of carbides - acetylenides, which contain the C 2 2- ion, which is actually a fragment of the acetylene molecule. Carbides such as acetylenides, upon hydrolysis or interaction with non-oxidizing acids, form acetylene as one of the reaction products. The type of carbide - methanide or acetylenide - obtained when a particular metal reacts with carbon depends on the size of the metal cation. Metal ions with a small radius usually form metanides, and larger ions form acetylenides. In the case of metals of the second group, methanide is obtained by the interaction of beryllium with carbon:

The remaining metals of group II A form acetylenides with carbon:

With silicon, group IIA metals form silicides - compounds of the type Me 2 Si, with nitrogen - nitrides (Me 3 N 2), with phosphorus - phosphides (Me 3 P 2):

with hydrogen

All alkaline earth metals react with hydrogen when heated. In order for magnesium to react with hydrogen, heating alone, as in the case of alkaline earth metals, is not enough; in addition to high temperature, increased hydrogen pressure is also required. Beryllium does not react with hydrogen under any conditions.

Interaction with complex substances

with water

All alkaline earth metals react actively with water to form alkalis (soluble metal hydroxides) and hydrogen. Magnesium reacts with water only when boiled due to the fact that when heated, the protective oxide film MgO dissolves in water. In the case of beryllium, the protective oxide film is very resistant: water does not react with it either when boiling or even at red-hot temperatures:

with non-oxidizing acids

All metals of the main subgroup of group II react with non-oxidizing acids, since they are in the activity series to the left of hydrogen. In this case, a salt of the corresponding acid and hydrogen are formed. Examples of reactions:

Be + H 2 SO 4 (diluted) = BeSO 4 + H 2

Mg + 2HBr = MgBr 2 + H 2

Ca + 2CH 3 COOH = (CH 3 COO) 2 Ca + H 2

with oxidizing acids

− diluted nitric acid

All metals of group IIA react with dilute nitric acid. In this case, the reduction products, instead of hydrogen (as in the case of non-oxidizing acids), are nitrogen oxides, mainly nitrogen oxide (I) (N 2 O), and in the case of highly dilute nitric acid, ammonium nitrate (NH 4 NO 3):

4Ca + 10HNO3 ( razb .) = 4Ca(NO 3) 2 + N 2 O + 5H 2 O

4Mg + 10HNO3 (very blurry)= 4Mg(NO 3) 2 + NH 4 NO 3 + 3H 2 O

− concentrated nitric acid

Concentrated nitric acid at ordinary (or low) temperature passivates beryllium, i.e. does not react with it. When boiling, the reaction is possible and proceeds predominantly in accordance with the equation:

Magnesium and alkaline earth metals react with concentrated nitric acid to form a wide range of different nitrogen reduction products.

− concentrated sulfuric acid

Beryllium is passivated with concentrated sulfuric acid, i.e. does not react with it under normal conditions, but the reaction occurs at boiling and leads to the formation of beryllium sulfate, sulfur dioxide and water:

Be + 2H 2 SO 4 → BeSO 4 + SO 2 + 2H 2 O

Barium is also passivated by concentrated sulfuric acid due to the formation of insoluble barium sulfate, but reacts with it when heated; barium sulfate dissolves when heated in concentrated sulfuric acid due to its conversion to barium hydrogen sulfate.

The remaining metals of main group IIA react with concentrated sulfuric acid under any conditions, including in the cold. Reduction of sulfur can occur to SO 2, H 2 S and S depending on the activity of the metal, reaction temperature and acid concentration:

Mg + H2SO4 ( conc. .) = MgSO 4 + SO 2 + H 2 O

3Mg + 4H 2 SO 4 ( conc. .) = 3MgSO 4 + S↓ + 4H 2 O

4Ca + 5H 2 SO 4 ( conc. .) = 4CaSO 4 +H 2 S + 4H 2 O

with alkalis

Magnesium and alkaline earth metals do not interact with alkalis, and beryllium easily reacts both with alkali solutions and with anhydrous alkalis during fusion. Moreover, when a reaction is carried out in an aqueous solution, water also participates in the reaction, and the products are tetrahydroxoberyllates of alkali or alkaline earth metals and hydrogen gas:

Be + 2KOH + 2H 2 O = H 2 + K 2 - potassium tetrahydroxoberyllate

When carrying out a reaction with a solid alkali during fusion, beryllates of alkali or alkaline earth metals and hydrogen are formed

Be + 2KOH = H 2 + K 2 BeO 2 - potassium beryllate

with oxides

Alkaline earth metals, as well as magnesium, can reduce less active metals and some nonmetals from their oxides when heated, for example:

The method of reducing metals from their oxides with magnesium is called magnesium.