Barium designation in the table. Barium

Barium is an alkaline earth metal that occupies position 56 in the periodic table of chemical elements. The name of the substance in translation 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 brittleness 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 a glass rod is lowered into barium and an open fire is brought up, then we will see a green flame. This method makes it possible to make a clear determination of even the minimum content of heavy metal impurities.

The crystal lattice of barium, which can be observed even outside the 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 an increase in temperature up to 365 0 С, and the other is able to withstand temperatures in the range of 375-710 0 С. The boiling point of barium is 1696 0 С.

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, which cannot be said about barium. An excellent medium for metal storage is paraffin oil, into which barium is directly immersed, or petroleum ether.

Barium reacts with oxygen, however, as a result of the reaction, its luster is lost, after which the metal first acquires a yellowish tint, then turns brown and eventually acquires a gray color. It is this appearance that is inherent in barium oxide. When the atmosphere is heated, barium becomes explosive.

The 56th element of the periodic system 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 gives an exceptionally pure metal, after which it becomes barium hydroxide. If metal salts are in contact with the aqueous medium, 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 conditions, the 56th element also reacts with ammonia, resulting in the formation of nitride. If the temperature is raised further, cyanide can be obtained.

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

Table 1. Properties of barium

Characteristic	Meaning
Atom properties
Name, symbol, number	Barium / Barium (Ba), 56
Atomic mass (molar mass)	137.327(7) a. e.m. (g/mol)
Electronic configuration	6s2
Atom 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 n.a.)	3.5 g/cm³
Melting temperature	1002 K
Boiling temperature	1910K
Oud. heat of fusion	7.66 kJ/mol
Oud. heat of evaporation	142.0 kJ/mol
Molar heat capacity	28.1 J/(K mol)
Molar volume	39.0 cm³/mol
The crystal lattice of a simple substance
Lattice structure	cubic body-centered
Lattice parameters	5.020Å
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 Gan. Then metal oxide was obtained. A few years later, Humphrey Davy succeeded in electrolyzing wet barium hydroxide with a mercury cathode to obtain an amalgam of metal, which he subjected to heating and evaporated mercury, thus obtaining metallic barium.

Obtaining metallic barium in modern laboratory conditions is carried out in several ways related to the atmosphere. The separation of barium is carried out in a vacuum due to the overly active reaction that is released when barium reacts with oxygen.

Barium oxide and chloride are obtained by metallothermic reduction under conditions of temperature increase up to 1200 0 С.

Also, a pure metal can be isolated from its hydride and nitride using thermal decomposition. In the same way, potassium is obtained. This process requires special capsules with complete sealing, 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.

Application of barium

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

1. In the manufacture of vacuum electronic devices. In this case, metallic barium, or its alloy with aluminum, is used as a getter. 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 capable of resisting corrosion. To do this, the metal, along with zirconium, is added to liquid metal coolants, which can significantly reduce the aggressive effect on pipelines. Such an application of barium has found a 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 atomic hydrogen energy. Barium chromate is actively used here during the production of hydrogen and oxygen using the thermochemical method.
7. in nuclear power. The metal oxide is an integral component of the process of making a certain type of glass, which is coated on uranium rods.
8. As a chemical current source. In this case, several compounds of barium can be used: fluoride, oxide and sulfate. The first compound is used in solid-state fluorine batteries as a component of a 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. Application of barium

Application area	Mode of application
Vacuum electronic devices	Barium metal, often in an alloy with aluminum, is used as a getter (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 an 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.
Ferroelectric 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).
Atomic hydrogen energy	Barium chromate is used in the production of hydrogen and oxygen by the 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%). In glassmaking for the nuclear industry, barium phosphate is also used.
Chemical current sources	Barium fluoride is used in solid-state fluorion batteries as a component of 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 expander in the production of lead-acid batteries .
Application in medicine	Barium sulfate, insoluble and non-toxic, is used as a radiopaque agent in medical examination of the gastrointestinal tract.

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 . Obtained by aluminothermy or decomposition of azide:
3BaO+2Al=Al 2 O 3 +3Ba
Ba(N 3) 2 \u003d 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. Tm.= 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, ignites when heated or when simply crushed. Vigorously interacts with halogens, when heated with hydrogen and sulfur.
Barium reacts vigorously with water and acids. Store, 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 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 to detect sulfate and carbonate ions, to purify vegetable and animal fats.
barium salts colorless crystals. substances. Soluble salts are highly toxic.
Chloride barium is obtained by the interaction of 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, a component of green pyrotechnic compositions. When heated, it decomposes to form barium oxide.
Sulfate barium is practically insoluble in water and in acids, therefore it is slightly toxic. used for bleaching paper, for fluoroscopy, barite concrete filler (protection from radioactive radiation).

Application:

Barium metal is used as a component of a number of alloys, a deoxidizer 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. Metal stories.

BARIUM COMPOUNDS, in accordance with the position of barium in the alkaline earth subgroup of group II of the Mendeleev system, have a doubly charged ion Ba ∙∙ (except for barium peroxide BaO 2). Barium compounds are characterized by a high specific gravity, colorlessness if the anions are not colored, a green color of the flame, and a small amount of complex compounds. Technically, the most important are oxide and peroxide, insoluble salts: barium carbonate, sulphate and chromic acid, and soluble salts: barium nitrate, barium chloride, etc. Soluble salts of barium are poisonous. Quantitatively, barium is determined in the form of BaSO 4 , but in view of the extreme fineness of the precipitates obtained at low temperatures, it is necessary to conduct precipitation from a boiling solution slightly acidified with hydrochloric acid. If there is nitric acid in the solution, part of the precipitate goes into solution. In addition, the BaSO 4 precipitate can carry away part of the salts due to adsorption. To separate from strontium, barium is precipitated as BaSiF 6 . If the barium compounds are insoluble, then they are fused with potassium carbonate-sodium and, after washing the alloy with water, they are dissolved in acid. Barium compounds are most commonly found as the mineral barite; much less common is witherite - barium carbonate.

Barium oxide BaO- white solid, crystallizes in cubes, density 5.72-5.32, melting point 1580 °, forms a crystalline hydrate according to the formula:

BaO + 9H 2 O \u003d Ba (OH) 2 ∙ 8H 2 O.

Barium oxide is relatively well soluble: at 0 ° - 1.5 hours in 100 hours of water; at 10° - 2.2 hours, at 15° - 2.89 hours, at 20° - 3.48 hours, at 50° - 11.75 hours, at 80° - 90.77 hours. Oxide barium is obtained from barium nitrate by calcination; this results in a porous product suitable for the manufacture of peroxide from it. Heating is carried out in crucibles, in a muffle furnace, at first very carefully so that the crucibles do not burst. The release of nitrogen oxides begins after 4 hours, but for their final removal, the crucibles are ignited for several hours at white heat (nitrogen oxides by 30% can be used to obtain nitric acid). The product is very expensive, because expensive: starting material, crucibles that are only good for one time, fuel, etc. Extracting barium oxide from witherite (BaCO 3 \u003d BaO + CO 2) is much more difficult than burning lime, t to. the reverse addition of CO 2 occurs very easily; therefore, coal is mixed with witherite so that CO 2 passes into CO. If it is desirable to obtain a porous product, then it is necessary to strictly adhere to the firing temperature. To prevent sintering, barium nitrate, coal, tar or barium carbide are often added, i.e.

ВаСО 3 + Ba(NO 3) 2 + 2С = 2ВаО + 2NO 2 + 3СО

ЗВаСО 3 + ВаС 2 = 4ВаО + 5СО.

In addition, it is necessary to protect the product as much as possible from sintering with the walls of the crucible and from the influence of hot gases. Calcination in shaft furnaces gives a very pure product (95%) if the furnace is built of high quality material and heated with generator gas, which allows precise temperature control. In Italy, heating in electric furnaces is used, but, apparently, this produces "oxycarbide" and "barium", which, in addition to 80-85% barium oxide, contains 10-12% carbide and 3-5% barium cyanide.

Aqueous barium oxide, caustic barite Ba (OH) 2 , forms transparent monoclinic crystals

Va (OH) 2 ∙ 8H 2 0,

losing the last molecule of water only at dark red heat; with light red heat, BaO is obtained, and with incandescence in a stream of air, barium peroxide is obtained. A solution of caustic barium - a strong alkali - absorbs CO 2 from the air, forming insoluble CaCO 3. 100 g of the solution contains: at 0 ° - 1.48 g of BaO, at 10 ° - 2.17, at 15 ° - 2.89, at 20 ° - 3.36, at 50 ° - 10.5, at 80 ° - 4.76. Caustic barite is used to absorb CO 2, extract caustic alkalis from sulphate, extract sugar from molasses, etc. Caustic barite can be obtained by calcining witherite by passing water vapor, but it is easier to burn BaCO 3 and act on BaO with water; or a mixture of 60% BaO and 40% BaS, obtained by calcining BaSO 4 with coal, is dissolved in water, and Ba (OH) 2 is obtained not only from BaO, but also from a significant part of BaS due to hydrolysis:

2BaS + 2HOH = Ba(OH) 2 + Ba(SH) 2 .

The crystallized substance contains only 1% impurities. The old methods of adding iron or zinc oxides to BaS are no longer used. It is also proposed to obtain caustic barite by electrolysis of barium chloride or barium chlorate and perchlorate in the presence of a BaCO 3 precipitate, which is dissolved by the acid formed at the anode.

Barium peroxide BaO 2 - white, mother-of-pearl intergrowths of the smallest crystals, very slightly soluble in water (only 0.168 hours in 100 hours of water). To obtain peroxide, barium oxide is heated in inclined tubes or in special muffles, which can be precisely kept at the desired temperature (500-600 °), and air purified from CO 2 and moisture is blown in. The purest peroxide is obtained in the form of square crystals of BaO 2 ∙ 8H 2 O, for which technical peroxide is first triturated with water, transferred to a solution by adding weak hydrochloric acid and precipitated with a solution of caustic barite, or simply add 10 times more 8% barite solution . The purest peroxide is a grayish-greenish sintered mass, insoluble in water, but interacting with carbonic anhydride. When heated, BaO 2 decomposes into BaO and oxygen. The elasticity of oxygen over BaO 2 at 555 ° - 25 mm, at 790 ° - 670 mm. Peroxide powder can ignite fibrous materials. On sale there are: the best grade - with 90% BaO 2 and the average - with 80-85%, with the main impurity being BaO. The content of BaO 2 is determined by titration with a 1/10 N-th KMnO 4 solution of BaO 2 in very weak cold hydrochloric acid (specific gravity 1.01-1.05), having previously precipitated barium ions with weak sulfuric acid. It is also possible to titrate the barium peroxide isolated from potassium iodide with sodium iodide sulphate. Barium peroxide is used to produce hydrogen peroxide (and at the same time get stronger whitewash "blancfix") and to prepare disinfectants.

Barium nitrite Ba (NO 2) 2 ∙ H 2 O - hexagonal colorless hexagonal prisms, melting point 220°. At 0 ° in 100 hours of water, 58 hours are dissolved, at 35 ° - 97 hours. It is obtained by adding a solution of sodium nitrite (360 hours of 96% NaNO 2 in 1000 hours of water) to a mixture of 360 hours of NaNO 2 and 610 hours BaCl 2 . At a high temperature, NaCl crystallizes, with further cooling - Ba (NO 2) 2.

Barium nitrate Ba (NO 3) 2 - colorless transparent octahedrons, melt at 375°; 100 hours of water are soluble at 10 ° - 7 hours, at 20 ° - 9.2 hours, at 100 ° - 32.2 hours. When heated, it first passes into barium nitrite, and then into barium oxide. It is used: 1) for the preparation of barium peroxide, 2) for green lights in fireworks, 3) for some explosives. It is produced: 1) by exchange decomposition when a theoretical amount of sodium nitrate is added to a hot solution of barium chloride (30 ° V) and subsequent recrystallization, 2) by the interaction of witherite or barium sulfide with nitric acid, 3) by heating calcium nitrate with technical barium carbonate.

Barium permanganate - manganese greens, Kassel greens, rosenstiel greens. BaMnO 4 - durable green paint suitable for fresco painting; obtained by calcining a mixture of compounds of barium (caustic barite, barium nitrate or barium peroxide) and manganese (dioxide or oxide).

Barium sulfide BaS - grayish porous mass, easily oxidized and attracts carbonic anhydride and water; decomposes with water. It is used for the manufacture of most barium compounds (lithopon, strong whitewash, etc.), for the extraction of sugar from molasses and the shearing of wool from skins (depilatorium). For mining, they use the calcination of a mixture of heavy spar with coal at 600-800 °:

BaSO 4 + 2C = 2CO 2 + BaS,

while at a higher temperature twice as much coal is wasted. The main condition is the close contact of coal and spar, which is achieved by grinding spar with 30-37% coal and water in rotating mills. The kilns are fired in rotary kilns, such as those used for cement or soda production, with a dusty chamber behind the short kilns to deposit smoke and soot. The resulting product contains 60-70% substances soluble in water, 20-25% - soluble in acids and 5% residue. The resulting product is thrown hot into water or into an aqueous solution of 1-2% NaOH (36 ° B), where half goes into aqueous oxide Ba (OH) 2, and the other into hydrosulphurous Ba (SH) 2. This solution is used directly for the preparation of barium compounds (lithopone, etc.) or for the extraction of sugar. When the residue reacts with hydrochloric acid, barium chloride is obtained. At old-type factories, calcination is carried out in fireclay retorts, evenly covered by flame. Well-dried slabs of coal and spar mixed with water are loaded into retorts. As soon as the flames of burning carbon monoxide disappear, the plates are removed so that they fall into hermetically sealed iron boxes.

Barium sulfate BaS 2 O 3 ∙ H 2 O It is formed from barium sulfide: 1) with free access of air and 2) with exchange decomposition with sodium sulphate. It is used to establish titers during iodometry.

Barium sulfate BaSO 4 , heavy spar (“strong”, “mineral”, “new”, etc. whitewash), pure white, earthy, very heavy powder, practically insoluble in water and acids (solubility: at 18 ° in 1 liter of water - 2 .3 mg). Natural grind directly. The best colorless varieties are called "flower" spar; ultramarine is added to yellowish and pinkish. Sometimes the heavy spar is ground and heated with hydrochloric acid to remove the iron; or spar is fused with Na 2 SO 4 and separated from the alloy by the action of water. Artificially it is obtained: 1) as waste in the preparation of hydrogen peroxide; 2) from barium chloride by interaction: a) with sulfuric acid, which gives a rapidly precipitated precipitate, b) with sodium sulphide Na 2 SO 4 or with magnesium sulphide salt MgSO 4, which gives a slowly falling and high-covering powder; during production, it is important to clean the sulfuric acid clean; 3) from witherite; if it is very pure, it can be ground directly by the action of H 2 SO 4 , but with the addition of 2% HCl; if witherite contains impurities, it is first dissolved in hydrochloric acid and then precipitated. Barium sulphate is used by Ch. arr. for coloring wallpaper colored paper, cardboard and especially for photographic papers, for light oil paints and varnish paints from coal, in the manufacture of artificial ivory and rubber, for mixing with food introduced into the stomach during radiography.

Barium carbonate BaCO 3 - mineral witherite (rhombic crystals) or artificially obtained in the form of the smallest sediment (specific gravity 4.3); more difficult to dissociate upon calcination than CaCO 3 ; at 1100° CO 2 pressure is only 20 mm. It is used to extract other barium compounds, in the manufacture of bricks and terracotta, porcelain, artificial marble and barite crystal. It is artificially prepared: 1) from a crude solution of barium sulfide by injecting carbonic anhydride; 2) heating barium sulfate with potash at 5 atm pressure; 3) upon decomposition of barium saccharate with carbonic anhydride.

Barium acetate Ba (C 2 H 3 O 2) 2 ∙ H 2 O - easily soluble crystals used in dyeing; are produced by the interaction of sodium sulphide or carbonate with acetic acid.

Barium fluoride BaF 2 - white powder, slightly soluble in water, melts at 1280°, obtained by dissolving barium carbonate or caustic in HF or boiling cryolite with aqueous barium oxide.

Barium chloride l 2 ∙ 2Н 2O- colorless flat rhombic plates (specific gravity 3.05), stable in air, sour in taste, poisonous; when heated, it is relatively easy to lose the first particle of water and much more difficult to lose the second; anhydrous BaCl 2 right. system melts at 962°. 100 hours of solution contains anhydrous salt:

ВаСl 2 is used for the manufacture of "durable" white and for the conversion of vitriol contained in ceramic products into insoluble BaSO 4; it is extracted from barite by calcining it with coal and calcium chloride in soda furnaces at 900-1000 ° in a reducing flame, and a 70% solution of calcium chloride can also be used, but solid calcium chloride is better:

BaSO 4 + 4C \u003d BaS + 4CO;

BaS + SaSl 2 \u003d YOUl 2 + CaS.

Properly produced, an almost black porous product with 50-56% BaCl 2 is obtained. After systematic leaching, the salt is crystallized (previously, a jet of carbonic anhydride is passed through) until hydrogen sulfide is completely removed and evaporated in vessels varnished inside. Crystals are separated by centrifugation. If anhydrous BaCl 2 is needed, then the salt is heated in vessels with stirrers to obtain very small crystals, which are then calcined, and 95% BaCl 2 is obtained. It is possible to obtain BaCl 2 by adding BaS powder to hydrochloric acid in closed vessels, from where it is necessary to remove the released hydrogen sulfide into the factory pipe or burn it to SO 2 using the latter for sulfuric acid. Of course, it is much more advantageous to act with hydrochloric acid on BaCO 3 .

Barium chlorate Ba(C lO 3) 2 ∙ H 2O- monoclinic prisms, highly soluble in cold and even better in hot water. Easily explodes when heated and on impact if mixed with a combustible substance. It is used in pyrotechnics for green flames. It is produced by electrolysis at 75° of a saturated solution of BaCl 2 , with a platinum anode and a graphite cathode.

Barium is an element of the main subgroup of the second group, the sixth period of the periodic system of chemical elements of D. I. Mendeleev, with atomic number 56. It is designated by the symbol Ba (lat. barium). A simple substance is a soft, ductile silver-white alkaline earth metal. Possesses high chemical activity.

History of the discovery of barium

Barium was discovered in the form of oxide BaO in 1774 by Karl Scheele. In 1808, the English chemist Humphrey Davy produced a barium amalgam by electrolysis of wet barium hydroxide with a mercury cathode; after evaporating the mercury on heating, he isolated barium metal.

In 1774, the Swedish chemist Carl Wilhelm Scheele and his friend Johan Gottlieb Hahn investigated one of the heaviest minerals, heavy spar BaSO 4 . They managed to isolate the previously unknown "heavy earth", which was later called barite (from the Greek βαρυς - heavy). And after 34 years, Humphry Davy, having subjected wet barite earth to electrolysis, obtained from it a new element - barium. It should be noted that in the same 1808, a little earlier than Davy, Jene Jacob Berzelius and his co-workers obtained amalgams of calcium, strontium and barium. This is how the element barium was born.

Ancient alchemists calcined BaSO 4 with wood or charcoal and obtained phosphorescent "Bolognese gems". But chemically, these gems are not BaO, but barium sulfide BaS.

origin of name

It got its name from the Greek barys - "heavy", since its oxide (BaO) was characterized as having an unusually high density for such substances.

Finding barium in nature

The earth's crust contains 0.05% barium. This is quite a lot - much more than, say, lead, tin, copper or mercury. In its pure form, it does not exist in the earth: barium is active, it is included in the subgroup of alkaline earth metals and, naturally, it is quite firmly bound in minerals.

The main minerals of barium are the already mentioned heavy spar BaSO 4 (more often called barite) and witherite BaCO3, named after the Englishman William Withering (1741 ... 1799), who discovered this mineral in 1782. Barium salts are found in a small concentration in many mineral waters and sea water. The low content in this case is a plus, not a minus, because all barium salts, except for sulfate, are poisonous.

Types of barium deposits

By mineral associations, barite ores are divided into monomineral and complex. Complex complexes are subdivided into barite-sulfide (contain lead, zinc, sometimes copper and iron pyrite sulfides, less often Sn, Ni, Au, Ag), barite-calcite (contain up to 75% calcite), iron-barite (contain magnetite, hematite, and goethite and hydrogoethite in the upper zones) and barite-fluorite (except for barite and fluorite, they usually contain quartz and calcite, and zinc, lead, copper, and mercury sulfides are sometimes present as small impurities).

From a practical point of view, hydrothermal vein monomineral, barite-sulfide and barite-fluorite deposits are of the greatest interest. Some metasomatic sheet deposits and eluvial placers are also of industrial importance. Sedimentary deposits, which are typical chemical sediments of water basins, are rare and do not play a significant role.

As a rule, barite ores contain other useful components (fluorite, galena, sphalerite, copper, gold in industrial concentrations), so they are used in combination.

Isotopes of barium

Natural barium consists of a mixture of seven stable isotopes: 130 Ba, 132 Ba, 134 Ba, 135 Ba, 136 Ba, 137 Ba, 138 Ba. The latter is the most common (71.66%). Radioactive isotopes of barium are also known, the most important of which is 140 Ba. It is formed during the decay of uranium, thorium and plutonium.

Obtaining barium

The metal can be obtained in various ways, in particular, by electrolysis of a molten mixture of barium chloride and calcium chloride. It is possible to obtain barium by restoring it from the oxide by the aluminothermic method. To do this, witherite is fired with coal and barium oxide is obtained:

BaCO 3 + C → BaO + 2CO.

Then a mixture of BaO with aluminum powder is heated in vacuum to 1250°C. Vapors of reduced barium condense in the cold parts of the tube in which the reaction takes place:

3BaO + 2Al → Al 2 O 3 + 3Ba.

It is interesting that barium peroxide BaO 2 is often included in the composition of ignition mixtures for aluminothermy.

Obtaining barium oxide by simple calcination of witherite is difficult: witherite decomposes only at temperatures above 1800°C. It is easier to obtain BaO by calcining barium nitrate Ba (NO 3) 2:

2Ba (NO 3) 2 → 2BaO + 4NO 2 + O 2.

Both electrolysis and aluminum reduction produce a soft (harder than lead, but softer than zinc) shiny white metal. It melts at 710°C, boils at 1638°C, its density is 3.76 g/cm 3 . All this fully corresponds to the position of barium in the subgroup of alkaline earth metals.

There are seven natural isotopes of barium. The most common of these is barium-138; it is more than 70%.

Barium is highly active. It self-ignites on impact, easily decomposes water, forming a soluble barium oxide hydrate:

Ba + 2H 2 O → Ba (OH) 2 + H 2.

An aqueous solution of barium hydroxide is called barite water. This "water" is used in analytical chemistry to determine CO 2 in gas mixtures. But this is already from the story about the use of barium compounds. Metallic barium finds almost no practical application. In extremely small quantities, it is introduced into bearing and printing alloys. An alloy of barium and nickel is used in radio tubes, pure barium is used only in vacuum technology as a getter (getter).

Barium metal is obtained from oxide by aluminum reduction in vacuum at 1200-1250°C:

4BaO + 2Al \u003d 3Ba + BaAl 2 O 4.

Barium is purified by vacuum distillation or zone melting.

Preparation of barium titanium. Getting it is relatively easy. Witherite BaCO 3 at 700 ... 800 ° C reacts with titanium dioxide TYu 2, it turns out just what you need:

BaCO 3 + TiO 2 → BaTiO 3 + CO 2.

Main prom. a method for obtaining metallic barium from BaO is its reduction with A1 powder: 4BaO + 2A1 -> 3Ba + BaO * A1 2 O 3. The process is carried out in a reactor at 1100-1200°C in an Ar atmosphere or in a vacuum (the latter method is preferable). The molar ratio of BaO:A1 is (1.5-2):1. The reactor is placed in a furnace so that the temperature of its "cold part" (the formed barium vapors condense in it) is about 520 ° C. By distillation in vacuum, barium is purified to an impurity content of less than 10 ~ 4% by weight, and when using zone melting - up to 10 ~ 6%.

Small amounts of barium are also obtained by reduction of BaBeO 2 [synthesized by fusion of Ba (OH) 2 and Be (OH) 2] at 1300 ° C with titanium, as well as by decomposition at 120 ° C Ba (N 3) 2, formed during exchange p- cations of barium salts with NaN 3 .

Acetate Ba (OOCHN 3), - colorless. crystals; m.p. 490°С (decomp.); dense 2.47 g/cm 3 ; sol. in water (58.8 g per 100 g at 0°C). Below 25 ° C, trihydrate crystallizes from aqueous solutions, at 25-41 ° C - monohydrate, above 41 ° C - anhydrous salt. Get interaction. Ba (OH) 2, VaCO 3 or BaS with CH 3 CO 2 H. Used as a mordant when dyeing wool and chintz.

Manganate(VI) BaMnO 4 - green crystals; does not decompose up to 1000°C. Obtained by calcining a mixture of Ba(NO 3) 2 with MnO 2 . A pigment (kassel or manganese green) commonly used for fresco painting.

Chromate (VI) ВаСrO 4 - yellow crystals; m.p. 1380°C; - 1366.8 kJ/mol; sol. in inorg. to-max, not sol. in water. Get interaction. aqueous solutions of Ba (OH) 2 or BaS with alkali metal chromates (VI). Pigment (barite yellow) for ceramics. MPC 0.01 mg / m 3 (in terms of Cr0 3). Pirconate ВаZrО 3 - colorless. crystals; m.p. ~269°С; - 1762 kJ/mol; sol. in water and aqueous solutions of alkalis and NH 4 HCO 3, decomposed by strong inorg. to-tami. Get interaction. ZrO 2 with BaO, Ba(OH) 2 or BaCO 3 when heated. Ba zirconate mixed with ВаТiO 3 -piezoelectric.

Bromide BaBr 2 - white crystals; m.p. 847°C; dense 4.79 g/cm 3 ; -757 kJ/mol; well sol. in water, methanol, worse - in ethanol. From aqueous solutions, the dihydrate crystallizes, turning into a monohydrate at 75 ° C, into an anhydrous salt - above 100 ° C. In aqueous solutions, the interaction. with CO 2 and O 2 of air, forming VaCO 3 and Br 2. Get BaBr 2 interaction. aqueous p-ditch Ba (OH) 2 or VaCO 3 with hydrobromic acid.

Iodide BaI 2 - colorless. crystals; m.p. 740°С (decomp.); dense 5.15 g/cm 3 ; . -607 kJ/mol; well sol. in water and ethanol. From hot water solutions, the dihydrate crystallizes (dehydrated at 150 ° C), below 30 ° C - hexahydrate. Get VaI 2 interaction. water p-ditch Ba (OH) 2 or VaCO 3 with hydroiodic acid.

Physical properties of barium

Barium is a silvery-white malleable metal. It breaks on a sharp blow. There are two allotropic modifications of barium: α-Ba with a cubic body-centered lattice is stable up to 375 °C (parameter a = 0.501 nm), β-Ba is stable above.

Hardness on a mineralogical scale 1.25; on the Mohs scale 2.

Barium metal is stored in kerosene or under a layer of paraffin.

Chemical properties of barium

Barium is an alkaline earth metal. It oxidizes intensively in air, forming barium oxide BaO and barium nitride Ba 3 N 2, and ignites when heated slightly. Vigorously reacts with water, forming barium hydroxide Ba (OH) 2:

Ba + 2H 2 O \u003d Ba (OH) 2 + H 2

Actively interacts with dilute acids. Many barium salts are insoluble or slightly soluble in water: barium sulfate BaSO 4, barium sulfite BaSO 3, barium carbonate BaCO 3, barium phosphate Ba 3 (PO 4) 2. Barium sulfide BaS, unlike calcium sulfide CaS, is highly soluble in water.

Natural barium has seven stable isotopes since May. ch. 130, 132, 134-137 and 138 (71.66%). The cross section of capture of thermal neutrons is 1.17-10 28 m 2 . External configuration electron shell 6s 2 ; oxidation state + 2, rarely + 1; ionization energy Ba° -> Ba + -> Ba 2+ resp. 5.21140 and 10.0040 eV; Pauling electronegativity 0.9; atomic radius 0.221 nm, ionic radius Ba 2+ 0.149 nm (coordination number 6).

Easily reacts with halogens to form halides.

When heated with hydrogen, it forms barium hydride BaH 2 , which, in turn, with lithium hydride LiH gives the Li complex.

Reacts on heating with ammonia:

6Ba + 2NH 3 = 3BaH 2 + Ba 3 N 2

Barium nitride Ba 3 N 2 reacts with CO when heated, forming cyanide:

Ba 3 N 2 + 2CO = Ba(CN) 2 + 2BaO

With liquid ammonia, it gives a dark blue solution, from which ammonia can be isolated, which has a golden sheen and easily decomposes with the elimination of NH 3. In the presence of a platinum catalyst, ammonia decomposes to form barium amide:

Ba (NH 2) 2 + 4NH 3 + H 2

Barium carbide BaC 2 can be obtained by heating BaO with coal in an arc furnace.

With phosphorus it forms the phosphide Ba 3 P 2 .

Barium reduces the oxides, halides and sulfides of many metals to the corresponding metal.

Application of barium

An alloy of barium with A1 (alba alloy, 56% Ba) is the basis of getters (getters). To obtain the getter itself, barium is evaporated from the alloy by high-frequency heating in an evacuated flask of the device; barium mirror (or diffuse coating during evaporation in a nitrogen atmosphere). The active part of the overwhelming majority of thermionic cathodes is BaO. Barium is also used as a Cu and Pb deoxidizer, as an additive to antifrictions. alloys, ferrous and non-ferrous metals, as well as alloys, from which typographic fonts are made to increase their hardness. Barium alloys with Ni are used for the manufacture of electrodes for glow plugs in internal engines. combustion and in radio tubes. 140 Va (T 1/2 12.8 days) is an isotope indicator used in the study of barium compounds.

Barium metal, often in an alloy with aluminum, is used as a getter in high vacuum electronic devices.

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.

Barium fluoride is used in the form of single crystals in optics (lenses, prisms).

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).

Barium chromate is used in the production of hydrogen and oxygen by the thermochemical method (Oak Ridge cycle, USA).

Barium oxide, together with oxides of copper and rare earth metals, is used to synthesize superconducting ceramics operating at liquid nitrogen temperatures and above.

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%). In glassmaking for the nuclear industry, barium phosphate is also used.

Barium fluoride is used in solid state fluorine batteries as a component of the fluoride electrolyte.

Barium oxide is used in powerful copper oxide batteries as a component of the active mass (barium oxide-copper oxide).

Barium sulfate is used as a negative electrode active mass expander in the production of lead-acid batteries.

Barium carbonate BaCO 3 is added to the glass mass to increase the refractive index of the glass. Barium sulfate is used in the paper industry as a filler; the quality of paper is largely determined by its weight, barite BaSO 4 makes the paper heavier. This salt is necessarily included in all expensive grades of paper. In addition, barium sulfate is widely used in the production of white lithopone paint, a product of the reaction of solutions of barium sulfide with zinc sulfate:

BaS + ZnSO 4 → BaSO 4 + ZnS.

Both salts, having a white color, precipitate, pure water remains in the solution.

When drilling deep oil and gas wells, a suspension of barium sulfate in water is used as a drilling fluid.

Another barium salt finds important uses. This is barium titanate BaTiO 3 - one of the most important ferroelectrics (ferroelectrics are polarized on their own, without exposure to an external field. Among dielectrics, they stand out in the same way as ferromagnetic materials among conductors. The ability for such polarization is maintained only at a certain temperature. Polarized ferroelectrics differ higher dielectric constant), which are considered very valuable electrical materials.

In 1944, this class was supplemented by barium titanate, the ferroelectric properties of which were discovered by the Soviet physicist B.M. Vulom. A feature of barium titanate is that it retains ferroelectric properties in a very wide temperature range - from close to absolute zero to +125°C.

Barium has also been used in medicine. Its sulfate salt is used in the diagnosis of gastric diseases. BaSO 4 is mixed with water and allowed to be swallowed by the patient. Barium sulfate is opaque to x-rays, and therefore those parts of the digestive tract, through which the "barium porridge" goes, remain dark on the screen. So the doctor gets an idea about the shape of the stomach and intestines, determines the place where an ulcer can occur.

The effect of barium on the human body

Routes of entry into the body.
The main way barium enters the human body is through food. Thus, some marine inhabitants are able to accumulate barium from the surrounding water, and in concentrations 7-100 (and for some marine plants up to 1000) times higher than its content in sea water. Some plants (soybeans and tomatoes, for example) are also able to accumulate barium from the soil by 2-20 times. However, in areas where the concentration of barium in the water is high, drinking water can also contribute to the total barium intake. The intake of barium from the air is negligible.

Health hazard.
In the course of scientific epidemiological studies conducted under the auspices of WHO, data on the relationship between mortality from cardiovascular diseases and the content of barium in drinking water have not been confirmed. In short-term studies in volunteers, there was no adverse effect on the cardiovascular system at barium concentrations up to 10 mg/l. True, in experiments on rats, when the latter consumed water even with a low content of barium, an increase in systolic blood pressure was observed. This indicates the potential danger of an increase in blood pressure in humans with prolonged use of water containing barium (USEPA has such data).
The USEPA data also show that even a single drink of water with a barium content far in excess of the maximum allowable levels can lead to muscle weakness and abdominal pain. However, it is necessary to take into account that the barium standard established by the USEPA quality standard (2.0 mg/l) significantly exceeds the value recommended by WHO (0.7 mg/l). Russian sanitary standards set an even more stringent MPC value for barium in water - 0.1 mg/l. Water removal technologies: ion exchange, reverse osmosis, electrodialysis.

Barium- an element of the main subgroup of the second group, the sixth period of the periodic system of chemical elements of D. I. Mendeleev, with atomic number 56. It is designated by the symbol Ba (lat. Barium). A simple substance is a soft, ductile silver-white alkaline earth metal. Possesses high chemical activity. History of the discovery of barium

1 element of the periodic table Barium was discovered in the form of oxide BaO in 1774 by Karl Scheele. In 1808, the English chemist Humphrey Davy produced a barium amalgam by electrolysis of wet barium hydroxide with a mercury cathode; after evaporating the mercury on heating, he isolated barium metal.
In 1774, the Swedish chemist Carl Wilhelm Scheele and his friend Johan Gottlieb Hahn investigated one of the heaviest minerals, heavy spar BaSO4. They managed to isolate the previously unknown "heavy earth", which was later called barite (from the Greek βαρυς - heavy). And after 34 years, Humphry Davy, having subjected wet barite earth to electrolysis, obtained from it a new element - barium. It should be noted that in the same 1808, a little earlier than Davy, Jene Jacob Berzelius and his co-workers obtained amalgams of calcium, strontium and barium. This is how the element barium was born.

Ancient alchemists calcined BaSO4 with wood or charcoal and obtained phosphorescent "Bolognese gems". But chemically, these gems are not BaO, but barium sulfide BaS.
It got its name from the Greek barys - "heavy", since its oxide (BaO) was characterized as having an unusually high density for such substances.
The earth's crust contains 0.05% barium. This is quite a lot - much more than, say, lead, tin, copper or mercury. In its pure form, it does not exist in the earth: barium is active, it is included in the subgroup of alkaline earth metals and, naturally, it is quite firmly bound in minerals.
The main minerals of barium are the already mentioned heavy spar BaSO4 (more often called barite) and witherite BaCO3, named after the Englishman William Withering (1741 ... 1799), who discovered this mineral in 1782. In a small concentration of barium salts, many mineral waters and sea water. The low content in this case is a plus, not a minus, because all barium salts, except for sulfate, are poisonous.

56 ← Barium→ Lantan
Atom properties
Name, symbol, number	Barium / Barium (Ba), 56
Atomic mass (molar mass)	137.327(7)(g/mol)
Electronic configuration
Atom radius
Chemical properties
covalent radius
Ion radius
Electronegativity	0.89 (Pauling scale)
Electrode potential
Oxidation states
Ionization energy (first electron)	502.5 (5.21) kJ/mol (eV)
Thermodynamic properties of a simple substance
Density (at n.a.)
Melting temperature
Boiling temperature
Oud. heat of fusion	7.66 kJ/mol
Oud. heat of evaporation	142.0 kJ/mol
Molar heat capacity	28.1 J/(K mol)
Molar volume	39.0 cm³/mol
The crystal lattice of a simple substance
Lattice structure	cubic body-centered
Lattice parameters
Other characteristics
Thermal conductivity	(300 K) (18.4) W/(m K)