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). The simple substance is a soft, malleable alkaline earth metal of a silvery-white color. Has high chemical activity.
History of the discovery of bariumBarium was discovered as the oxide BaO in 1774 by Karl Scheele. In 1808, the English chemist Humphry Davy obtained barium amalgam by electrolysis of wet barium hydroxide with a mercury cathode; After the mercury evaporated when heated, it released 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 previously unknown “heavy earth,” which was later called barite (from the Greek βαρυς - heavy). And 34 years later, Humphry Davy, having subjected wet barite earth to electrolysis, obtained a new element from it - barium. It should be noted that in the same 1808, somewhat earlier than Davy, Jene Jacob Berzelius and his colleagues obtained amalgams of calcium, strontium and barium. This is how the element barium appeared.
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 natureThe earth's crust contains 0.05% barium. This is quite a lot - significantly more than, say, lead, tin, copper or mercury. It is not found in the earth in its pure form: barium is active, it belongs to the subgroup of alkaline earth metals and, naturally, is bound quite tightly 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. A small concentration of barium salts is contained in many mineral waters and sea water. The low content in this case is a plus, not a minus, because all barium salts, except sulfate, are poisonous.
Types of barium depositsBased on mineral associations, barite ores are divided into monomineral and complex. Complex complexes are divided into barite-sulfide (contain sulfides of lead, zinc, sometimes copper and iron pyrite, less often Sn, Ni, Au, Ag), barite-calcite (contain up to 75% calcite), iron-barite (contain magnetite, hematite, and in the upper zones goethite and hydrogoethite) and barite-fluorite (in addition to barite and fluorite, they usually contain quartz and calcite, and zinc, lead, copper and mercury sulfides are sometimes present in the form of small impurities).
From a practical point of view, hydrothermal vein monomineral, barite-sulfide and barite-fluorite deposits are of greatest interest. Some metasomatic strata 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.
Barium isotopesNatural 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 by the decay of uranium, thorium and plutonium.
Obtaining bariumThe metal can be obtained in different ways, in particular by electrolysis of a molten mixture of barium chloride and calcium chloride. It is possible to obtain barium by reducing it from its oxide using an aluminothermic method. To do this, witherite is fired with coal and barium oxide is obtained:
BaCO 3 + C → BaO + 2CO.
Then the mixture of BaO with aluminum powder is heated in vacuum to 1250°C. Reduced barium vapor condenses in the cold parts of the pipe in which the reaction takes place:
3BaO + 2Al → Al 2 O 3 + 3Ba.
It is interesting that the composition of ignition mixtures for aluminothermy often includes barium peroxide BaO 2.
It is difficult to obtain barium oxide by simply calcining witherite: 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 reduction with aluminum produce a soft (harder than lead, but softer than zinc) shiny white metal. It melts at 710°C, boils at 1638°C, and 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 known natural isotopes of barium. The most common of these is barium-138; its more than 70%.
Barium is very active. It self-ignites on impact and easily decomposes water to form soluble barium oxide hydrate:
Ba + 2H 2 O → Ba (OH) 2 + H 2.
An aqueous solution of barium oxide hydrate is called barite water. This “water” is used in analytical chemistry for the determination of CO 2 in gas mixtures. But this is already from the story about the use of barium compounds. Metallic barium finds almost no practical use. It is introduced in extremely small quantities 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 (gas absorber).
Metal barium is obtained from the oxide by reduction with aluminum in a vacuum at 1200-1250°C:
4BaO + 2Al = 3Ba + BaAl 2 O 4.
Barium is purified by vacuum distillation or zone smelting.
Preparation of barium titanium. It is relatively easy to obtain. Witherite BaCO 3 at 700...800°C reacts with titanium dioxide TiO 2, the result is exactly what is needed:
BaCO 3 + TiO 2 → BaTiO 3 + CO 2.
Basic prom. The method for obtaining barium metal 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 resulting barium vapors are condensed 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 the reduction of BaBeO 2 [synthesized by the fusion of Ba(OH) 2 and Be(OH) 2 ] at 1300°C with titanium, as well as the decomposition at 120°C of Ba(N 3) 2 formed during exchange p- tions of barium salts with NaN 3.
Ba acetate (OOСSN 3), - colorless. crystals; m.p. 490°C (with decomposition); dense 2.47 g/cm3; 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. Receive interaction. Ba(OH)2, BaCO3 or BaS with CH3CO2H. Used as a mordant when dyeing wool and calico.
Manganate(VI) BaMnO 4 - green crystals; does not decompose up to 1000°C. Obtained by calcination of a mixture of Ba(NO 3) 2 with MnO 2. A pigment (Cassel, or manganese green) commonly used for fresco painting.
Chromate(VI) BaСrO 4 - yellow crystals; m.p. 1380°C; - 1366.8 kJ/mol; sol. in non-org. k-tah, not sol. in water. Receive 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). Pyrconate BaZrO 3 - colorless. crystals; m.p. ~269°C; - 1762 kJ/mol; sol. in water and aqueous solutions of alkalis and NH 4 HCO 3, decomposes by strong inorg. to-tami. Receive interaction. ZrO 2 with BaO, Ba(OH) 2 or BaCO 3 when heated. Ba zirconate mixed with BaTiO 3 is a piezoelectric.
Bromide BaBr 2 - white crystals; m.p. 847°C; dense 4.79 g/cm3; -757 kJ/mol; well sol. in water, methanol, worse - in ethanol. Dihydrate crystallizes from aqueous solutions, turning into monohydrate at 75°C, into anhydrous salt - above 100°C. In aqueous solutions, interaction. with CO 2 and O 2 of air, forming BaCO 3 and Br 2. Get BaBr 2 interaction. aqueous solutions of Ba(OH) 2 or BaCO 3 with hydrobromic acid.
Iodide BaI 2 - colorless. crystals; m.p. 740°C (with decomposition); dense 5.15 g/cm3; . -607 kJ/mol; well sol. in water and ethanol. From hot water solutions, the dihydrate crystallizes (dehydrates at 150°C), below 30°C - the hexahydrate. Get BaI 2 interaction. aqueous solutions of Ba(OH) 2 or BaCO 3 with hydroiodic acid.
Physical properties of bariumBarium is a silvery-white malleable metal. If struck sharply, it breaks. There are two allotropic modifications of barium: α-Ba with a cubic body-centered lattice (parameter a = 0.501 nm) is stable up to 375 °C; β-Ba is stable above it.
Hardness on the mineralogical scale 1.25; Mohs scale 2.
Store barium metal in kerosene or under a layer of paraffin.
Chemical properties of bariumBarium is an alkaline earth metal. It oxidizes intensively in air, forming barium oxide BaO and barium nitride Ba 3 N 2 , and ignites with slight heating. Reacts vigorously with water, forming barium hydroxide Ba(OH) 2:
Ba + 2H 2 O = 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.
Nature Barium consists of seven stable isotopes since May. parts 130, 132, 134-137 and 138 (71.66%). The cross-section of thermal neutron capture 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).
Reacts easily with halogens to form halides.
When heated with hydrogen, it forms barium hydride BaH 2 , which in turn forms the Li complex with lithium hydride LiH.
Reacts when heated with ammonia:
6Ba + 2NH 3 = 3BaH 2 + Ba 3 N 2
When heated, barium nitride Ba 3 N 2 reacts with CO, 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 phosphide Ba 3 P 2 .
Barium reduces the oxides, halides and sulfides of many metals to the corresponding metal.
Applications of bariumAn alloy of barium with A1 (Alba alloy, 56% Ba) is the basis of getters (gas absorbers). To obtain the getter itself, barium is evaporated from the alloy by high-frequency heating in an evacuated flask of the device; as a result, the so-called barium is formed on the cold parts of the flask. barium mirror (or diffuse coating during evaporation in a nitrogen environment). The active part of the vast majority of thermionic cathodes is BaO. Barium is also used as a deoxidizing agent for Cu and Pb, and as an additive to antifriction agents. alloys, ferrous and non-ferrous metals, as well as alloys from which printing fonts are made to increase their hardness. Alloys of barium with Ni are used for the manufacture of spark plug electrodes in internal engines. combustion and in radio tubes. 140 Va (T 1/2 12.8 days) is an isotopic indicator used in the study of barium compounds.
Barium metal, often alloyed with aluminum, is used as a getter in high-vacuum electronic devices.
Anti-corrosion materialBarium 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 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%). Barium phosphate is also used in glass melting for the nuclear industry.
Barium fluoride is used in solid-state fluorine batteries as a component of the fluoride electrolyte.
Barium oxide is used in high-power 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 types of paper. In addition, barium sulfate is widely used in the production of white paint lithopone - a product of the reaction of solutions of barium sulfide with zinc sulfate:
BaS + ZnSO 4 → BaSO 4 + ZnS.
Both salts, which are white, precipitate, leaving pure water 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 has important uses. This is barium titanate BaTiO 3 - one of the most important ferroelectrics (ferroelectrics are polarized on their own, without the influence of an external field. They stand out among dielectrics in the same way as ferromagnetic materials among conductors. The ability for such polarization is retained only at a certain temperature. Polarized ferroelectrics differ higher dielectric constant), which are considered very valuable electrical materials.
In 1944, this class was replenished with barium titanate, the ferroelectric properties of which were discovered by the Soviet physicist B.M. Vulom. The peculiarity of barium titanate is that it retains ferroelectric properties over a very wide temperature range - from close to absolute zero to +125°C.
Barium has also found application in medicine. Its sulfate salt is used in the diagnosis of gastric diseases. BaSO 4 is mixed with water and given to the patient to swallow. Barium sulfate is opaque to X-rays, and therefore those parts of the digestive tract through which the “barium porridge” passes remain dark on the screen. This way the doctor gets an idea of the shape of the stomach and intestines and determines the place where an ulcer may occur.
The effect of barium on the human body
Routes of entry into the body.
The main route of entry of barium into the human body is food. Thus, some marine inhabitants are capable of accumulating 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 capable of accumulating barium from the soil 2-20 times. However, in areas where barium concentrations in water are high, drinking water may also contribute to total barium consumption. The intake of barium from the air is insignificant.
Health hazard.
Scientific epidemiological studies conducted under the auspices of WHO did not confirm the relationship between mortality from cardiovascular diseases and barium levels in drinking water. In short-term studies in volunteers, no harmful effects on the cardiovascular system were detected at barium concentrations up to 10 mg/l. True, in experiments on rats, when the latter consumed water even with a low barium content, an increase in systolic blood pressure was observed. This indicates a potential risk of increased blood pressure in people with long-term consumption of water containing barium (USEPA has such data).
USEPA data also suggests that even a single drink of water containing barium levels well above the maximum permissible levels can lead to muscle weakness and abdominal pain. It is necessary, however, to take into account that the standard for barium 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. Technologies for removing water: ion exchange, reverse osmosis, electrodialysis.
With the chemical formula BaSO 4. It is an odorless white powder, insoluble in water. Its whiteness and opacity, as well as its high density, determine its main areas of application.
History of the name
Barium belongs to the alkaline earth metals. The latter are so named because, according to D.I. Mendeleev, their compounds form an insoluble mass of earth, and the oxides “have an earthy appearance.” Barium is naturally found in the form of the mineral barite, which is barium sulfate with various impurities.
It was first discovered by Swedish chemists Scheele and Hahn in 1774 as part of the so-called heavy spar. This is where the name of the mineral came from (from the Greek “baris” - heavy), and then the metal itself, when in 1808 it was isolated in its pure form by Humphry Devi.
Physical properties
Since BaSO 4 is a salt of sulfuric acid, its physical properties are determined in part by the metal itself, which is soft, reactive and silvery-white. Natural barite is colorless (sometimes white) and transparent. Chemically pure BaSO 4 has a color from white to pale yellow, it is non-flammable, with a melting point of 1580°C.
What is the mass of barium sulfate? Its molar mass is 233.43 g/mol. It has an unusually high specific gravity - from 4.25 to 4.50 g/cm 3 . Given its insolubility in water, its high density makes it indispensable as a filler in aqueous drilling fluids.
Chemical properties
BaSO 4 is one of the most sparingly soluble compounds in water. It can be obtained from two highly soluble salts. Let's take an aqueous solution of sodium sulfate - Na 2 SO 4. Its molecule in water dissociates into three ions: two Na + and one SO 4 2-.
Na 2 SO 4 → 2Na + + SO 4 2-
Let us also take an aqueous solution of barium chloride - BaCl 2, the molecule of which dissociates into three ions: one Ba 2+ and two Cl -.
BaCl 2 → Ba 2+ + 2Cl -
Mix an aqueous solution of sulfate and a mixture containing chloride. Barium sulfate is formed as a result of the combination into one molecule of two ions with the same charge and opposite sign.
Ba 2+ + SO 4 2- → BaSO 4
Below you can see the complete equation for this reaction (called molecular).
Na 2 SO 4 + BaCl 2 → 2NaCl + BaSO 4
As a result, an insoluble precipitate of barium sulfate is formed.
Commercial barite
In practice, the starting material for obtaining commercial barium sulfate, intended for use in drilling fluids when drilling oil and gas wells, is, as a rule, the mineral barite.
The term "primary" barite refers to commercial products, which include raw material (obtained from mines and quarries), as well as products of simple beneficiation by methods such as washing, precipitation, separation in heavy media, and flotation. Most raw barite requires refinement to a minimum purity and density. The mineral used as a filler is crushed and sifted to a uniform size so that at least 97% of its particles are up to 75 microns in size, and no more than 30% are less than 6 microns. Primary barite must also be dense enough to have a specific gravity of 4.2 g/cm 3 or higher, but soft enough not to damage the bearings.
Obtaining a chemically pure product
Mineral barite is often contaminated with various impurities, mainly iron oxides, which color it in different colors. It is processed carbothermically (heating with coke). The result is barium sulfide.
BaSO 4 + 4 C → BaS + 4 CO
The latter, unlike sulfate, is soluble in water and easily reacts with oxygen, halogens and acids.
BaS + H 2 SO 4 → BaSO 4 + H 2 S
To obtain a highly pure output product, sulfuric acid is used. The barium sulfate produced by this process is often called blancfix, which is French for “white fixed.” It is often found in consumer products such as paints.
In laboratory conditions, barium sulfate is formed by combining barium ions and sulfate ions in solution (see above). Since sulfate is the least toxic barium salt due to its insolubility, waste containing other barium salts is sometimes treated with sodium sulfate to bind all the barium, which is quite toxic.
From sulfate to hydroxide and back
Historically, barite was used to produce barium hydroxide Ba(OH) 2, necessary in sugar refining. This is generally a very interesting compound that is widely used in industry. It is highly soluble in water, forming a solution known as barite water. It is convenient to use for binding sulfate ions in various compositions by forming insoluble BaSO 4 .
We saw above that when heated in the presence of coke, it is easy to obtain water-soluble barium sulfide - BaS - from sulfate. The latter, when interacting with hot water, forms hydroxide.
BaS + 2H 2 O → Ba(OH) 2 + H 2 S
Barium hydroxide and sodium sulfate, taken in solutions, when mixed, will give an insoluble precipitate of barium sulfate and sodium hydroxide.
Ba(OH) 2 + Na 2 SO 4 = BaSO 4 + 2NaOH
It turns out that natural barium sulfate (barite) is industrially first converted into barium hydroxide, and then serves to produce the same sulfate when purifying various salt systems from sulfate ions. The reaction will proceed in exactly the same way when purifying a solution of copper sulfate from SO 4 2 ions. If you make a mixture of barium hydroxide + copper sulfate, the result will be copper hydroxide and insoluble barium sulfate.
CuSO 4 + Ba(OH) 2 → Cu(OH) 2 + BaSO 4 ↓
Even in a reaction with sulfuric acid itself, its sulfate ions will be completely bound with barium.
Use in drilling fluids
About 80% of the world's production of barium sulfate, purified and crushed barite, is consumed as a component of drilling fluids in the creation of oil and gas wells. Adding it increases the density of the fluid pumped into the well in order to better resist high reservoir pressure and prevent breakthroughs.
When a well is drilled, the bit passes through various formations, each of which has its own characteristics. The greater the depth, the greater the percentage of barite that should be present in the solution structure. An additional advantage is that barium sulfate is a non-magnetic substance, so it does not interfere with various downhole measurements using electronic devices.
Paint and paper industry
Most synthetic BaSO 4 is used as a component of the white pigment for paints. Thus, blancfix mixed with titanium dioxide (TiO 2) is sold as a white oil paint used in painting.
The combination of BaSO 4 and ZnS (zinc sulfide) produces an inorganic pigment called lithopone. It is used as a coating for certain types of photographic paper.
More recently, barium sulfate has been used to brighten paper intended for inkjet printers.
Application in the chemical industry and non-ferrous metallurgy
In the production of polypropylene and polystyrene, BaSO 4 is used as a filler in a proportion of up to 70%. It has the effect of increasing the resistance of plastics to acids and alkalis and also imparting opacity to them.
It is also used to produce other barium compounds, particularly barium carbonate, which is used to make LED glass for television and computer screens (historically in cathode ray tubes).
Molds used in metal casting are often coated with barium sulfate to prevent adhesion to the molten metal. This is what is done in the manufacture of anode copper plates. They are cast into copper molds coated with a layer of barium sulfate. Once the liquid copper has solidified into a finished anode plate, it can be easily removed from the mold.
Pyrotechnic devices
Since barium compounds emit green light when burned, salts of this substance are often used in pyrotechnic formulas. Although nitrate and chlorate are more common than sulfate, the latter is widely used as a component of pyrotechnic strobes.
X-ray contrast agent
Barium sulfate is a radiopaque contrast agent used to diagnose certain medical problems. Since such substances are opaque to x-rays (they block them as a result of their high density), the areas of the body in which they are localized appear as white areas on x-ray film. This creates the necessary distinction between one (diagnosed) organ and other (surrounding) tissues. The contrast will help the doctor see any special conditions that may exist in that organ or body part.
Barium sulfate is taken by mouth or rectally with an enema. In the first case, it makes the esophagus, stomach or small intestine opaque to X-rays. This way they can be photographed. If the substance is administered through an enema, the colon or intestines can be seen and recorded with x-rays.
The dose of barium sulfate will be different for different patients, depending on the type of test. The drug is available in the form of a special medical barium suspension or in tablets. Different tests that require contrast and X-ray equipment require different amounts of suspension (in some cases the drug may be taken in tablet form). Contrast material should only be used under the direct supervision of a physician.
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 Ba ∙∙ ion (except for barium peroxide BaO 2). Barium compounds are characterized by a high specific gravity, colorlessness if the anions are not colored, a green flame color and a small number of complex compounds. Technically, the most important are oxide and peroxide, insoluble salts: barium carbonate, sulfate and chromate, and soluble salts: barium nitrate, barium chloride, etc. Soluble barium salts are poisonous. Barium is determined quantitatively in the form of BaSO 4 , but due to the extreme fineness of the precipitation obtained at low temperatures, it is necessary to precipitate 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 may entrain some of the salts due to adsorption. To separate it from strontium, barium is precipitated in the form of BaSiF 6 . If barium compounds are insoluble, then they are fused with potassium-sodium carbonate and, after washing the alloy with water, dissolved in acid. Barium compounds are most often 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 = Ba(OH)2 ∙ 8H 2 O.
Barium oxide is relatively highly soluble: at 0° - 1.5 parts in 100 parts 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 produces a porous product suitable for making peroxide from it. Heating is carried out in crucibles in a muffle furnace, very carefully at first so that the crucibles do not burst. The release of nitrogen oxides begins after 4 hours, but to completely remove them, the crucibles are heated for several hours at white heat (30% of nitrogen oxides can be used to produce nitric acid). The product is very expensive, because expensive: starting material, crucibles, which are only suitable for one time, fuel, etc. Extraction of barium oxide (BaCO 3 = BaO + CO 2) from witherite is much more difficult than burning lime, t because the reverse addition of CO 2 occurs very easily; Therefore, coal is mixed with witherite so that CO 2 turns into CO. If it is desired to obtain a porous product, then the firing temperature must be strictly adhered to. To prevent sintering, barium nitrate, coal, tar or barium carbide are often added, i.e.
BaCO 3 + Ba(NO 3) 2 + 2C = 2BaO+ 2NO 2 + 3CO
ZBaCO 3 + BaC 2 = 4BaO + 5CO.
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 produces a very pure product (95%) if the furnace is built from high quality material and heated by generator gas, which allows precise temperature control. In Italy, heating is used in electric furnaces, but apparently this produces “oxycarbide” and “barium”, which, in addition to 80-85% barium oxide, contains 10-12% carbide and 3-5% barium cyanide.
Hydrous barium oxide, caustic barite Ba(OH) 2 , forms transparent monoclinic crystals
Ba(OH) 2 ∙ 8H 2 0,
losing the last molecule of water only at dark red heat; When heated with light red heat, BaO is obtained, and when heated 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 solution contains: at 0° - 1.48 g 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 sulfuric acids, extract sugar from molasses, etc. Caustic barite can be obtained by calcining witherite by passing water steam, 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 perchlorate and barium chloride in the presence of a BaCO 3 precipitate, which is dissolved by the acid formed at the anode.
Barium peroxide BaO 2 - white, pearl-like aggregates of tiny crystals, very slightly soluble in water (only 0.168 parts in 100 parts of water). To obtain peroxide, barium oxide is heated in inclined pipes or in special muffles, which can be precisely kept at the desired temperature (500-600°), and air purified from CO 2 and moisture is pumped in. The purest peroxide is obtained in the form of square crystals of BaO 2 ∙ 8H 2 O, for which they first grind technical peroxide with water, transfer it into solution by adding weak hydrochloric acid and precipitate it with a solution of caustic barite or simply add 10 times more quantity of 8% barite solution . The purest peroxide is a grayish-greenish sintered mass, insoluble in water, but reacts with carbonic anhydride. When heated, BaO 2 decomposes into BaO and oxygen. The elasticity of oxygen over BaO 2 at 555° is 25 mm, at 790° - 670 mm. Peroxide powder may cause fibrous materials to ignite. On sale there are: the best variety - with 90% BaO 2 and the average - with 80-85%, with the main impurity being BaO. The BaO 2 content is determined by titrating a 1/10 N 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. You can also titrate barium peroxide isolated from potassium iodide with sodium iodide. Barium peroxide is used to produce hydrogen peroxide (and at the same time, stronger white "Blanfix" is obtained) and for the preparation of disinfectants.
Barium nitrate Ba(NO 2) 2 ∙ H 2 O - hexagonal colorless hexagonal prisms, melting point 220°. At 0°, 58 parts of water dissolve in 100 parts, at 35° - 97 hours. It is obtained by adding a solution of sodium nitrate (360 parts of 96% NaNO 2 in 1000 parts of water) into a mixture of 360 parts of NaNO 2 and 610 hours BaCl 2 . At high temperatures, NaCl crystallizes, with further cooling - Ba(NO 2) 2.
Barium nitrate Ba(NO 3) 2 - colorless transparent octahedra, melt at 375°; in 100 parts of water, it is soluble at 10° - 7 hours, at 20° - 9.2 hours, at 100° - 32.2 hours. When heated, it turns first into barium nitrate, and then into barium oxide. Used: 1) for the preparation of barium peroxide, 2) for green lights in fireworks, 3) for some explosives. It is obtained: 1) by exchange decomposition when adding a theoretical amount of sodium nitrate to a hot solution of barium chloride (30° Ве) and subsequent recrystallization, 2) by the interaction of witherite or barium sulfide with nitric acid, 3) by heating calcium nitrate with commercial barium carbonate.
Barium permanganate - manganese greens, Kassel greens, rosenstiel greens. BaMnO 4 - durable green paint, suitable for fresco painting; is obtained by calcining a mixture of barium compounds (caustic barite, barium nitrate or barium peroxide) and manganese (dioxide or oxide).
Barium sulfide BaS - a grayish porous mass that easily oxidizes and attracts carbonic anhydride and water; decomposes with water. It is used for the manufacture of most barium compounds (lithopone, strong white, etc.), for isolating sugar from molasses and removing wool from skins (depilatory). For mining, they use calcination of a mixture of heavy spar with coal at 600-800°:
BaSO4 + 2C = 2CO2+BaS,
whereas 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. Firing is carried out in rotary kilns, similar to those used for cement or soda production, and a dust chamber must be placed behind the short kilns to allow smoke and soot to settle. The resulting product contains 60-70% substances soluble in water, 20-25% soluble in acids and 5% residue. The resulting hot product is thrown into water or into an aqueous solution of 1-2% NaOH (36° Ве), where half turns into aqueous Ba(OH)2 oxide, 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. In old-type factories, calcination is carried out in fireclay retorts, uniformly engulfed in flame. Well-dried slabs of coal and spar mixed with water are loaded into the retorts. As soon as the flames of burning carbon monoxide disappear, the slabs are removed so that they fall into hermetically sealed iron boxes.
Barium sulphate BaS 2 O 3 ∙ H 2 O is formed from barium sulfide: 1) with free access of air and 2) during exchange decomposition with sodium sulfate. Used to establish titers during iodometry.
Barium sulfate BaSO 4 , heavy spar (“strong”, “mineral”, “new”, etc. white), pure white, earthy, very heavy powder, practically insoluble in water and acids (solubility: at 18° in 1 liter of water - 2 .3 mg). Natural directly grind. The best colorless varieties are called "flower" spar; ultramarine is added to yellowish and pinkish ones. Sometimes heavy spar is ground and heated with hydrochloric acid to remove 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 during the preparation of hydrogen peroxide; 2) from barium chloride by interaction: a) with sulfuric acid, which gives a rapidly precipitating precipitate, b) with sodium sulfur Na 2 SO 4 or with magnesium sulfur salt MgSO 4, which gives a slowly falling powder with high covering power; During production, it is important to wash the sulfuric acid clean; 3) from witherite; if it is very pure, it can be crushed 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 sulfate is used in Ch. arr. for coloring colored wallpaper paper, cardboard and especially for photographic papers, for light oil paints and coal varnish paints, 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 minute sediment (specific gravity 4.3); dissociates more difficult upon calcination than CaCO 3 ; at 1100° the CO 2 pressure is only 20 mm. It is used for the extraction of other barium compounds, in the manufacture of bricks and terracotta, porcelain, artificial marble and barite crystal. It is prepared artificially: 1) from a crude solution of barium sulfide by introducing carbonic anhydride; 2) heating barium sulfate with potash at 5 atm pressure; 3) during the decomposition of barium sucrose by carbonic anhydride.
Barium acetate Ba (C 2 H 3 O 2) 2 ∙ H 2 O - easily soluble crystals used in dyeing; are obtained by reacting sodium sulphide or carbon dioxide with acetic acid.
Barium fluoride BaF 2 - white powder, slightly soluble in water, melts at 1280°, obtained by dissolving carbon dioxide or caustic barium in HF or boiling cryolite with aqueous barium oxide.
Barium chloride VaS l 2 ∙ 2Н 2O- colorless flat rhombic plates (specific gravity 3.05), stable in air, taste sour, 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. systems melts at 962°. 100 parts of solution contains anhydrous salt:
BaCl 2 is used for the manufacture of “strong” white and for converting 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 be used, but solid calcium chloride is better:
BaSO 4 + 4C = BaS + 4CO;
BaS + CaSl 2 = VaSl 2 + CaS.
When properly produced, an almost black porous product with 50-56% BaCl 2 is obtained. After systematic leaching, the salt is crystallized (a stream of carbonic anhydride is first passed through) until hydrogen sulfide is completely removed and evaporated in vessels varnished inside. The 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. BaCl 2 can be obtained by adding BaS powder to hydrochloric acid located in closed vessels, from where it is necessary to remove the released hydrogen sulfide into a factory pipe or burn it to SO 2 using the latter for sulfuric acid. Of course, it is much more profitable to act with hydrochloric acid on BaCO 3.
Barium perchlorate Ba(C lO 3) 2 ∙ N 2O- monoclinic prisms, highly soluble in cold and even better in hot water. Explodes easily when heated and on impact if mixed with a flammable substance. Used in pyrotechnics for green flames. It is produced by electrolysis at 75° of a saturated BaCl 2 solution, with a platinum anode and a graphite cathode.
BARIUM (Latin Barium), Ba, a chemical element of group II of the short form (group 2 of the long form) of the periodic system; refers to alkaline earth metals; atomic number 56, atomic mass 137.327. There are 7 stable nuclides in nature, among which 138 Ba predominates (71.7%); about 30 nuclides are obtained artificially.
Historical reference. Barium in the form of oxide was discovered in 1774 by K. Scheele, who discovered a previously unknown “earth”, later called “heavy earth” - barite (from the Greek βαρ?ς - heavy). In 1808, G. Davy obtained barium metal in the form of an amalgam by electrolysis of molten salts.
Prevalence in nature. The barium content in the earth's crust is 5·10 -2% by mass. Due to high chemical activity, it is not found in free form. Main minerals: barite BaSO 4 and witherite BaSO 3. World production of BaSO 4 is about 6 million tons/year.
Properties. The configuration of the outer electron shell of the barium atom is 6s 2; in compounds exhibits an oxidation state of +2, rarely +1; Pauling electronegativity 0.89; atomic radius 217.3 nm, radius of the Ba 2+ ion 149 nm (coordination number 6). The ionization energy of Ba 0 → Ba + → Ba 2+ is 502.8 and 965.1 kJ/mol. The standard electrode potential of the Ba 2+ /Ba pair in an aqueous solution is -2.906 V.
Barium is a silvery-white malleable metal; t pl 729 °C, t KIK 1637 °C. At normal pressure, the barium crystal lattice is body-centered cubic; at 19 °C and 5530 MPa, a hexagonal modification is formed. At 293 K, the density of barium is 3594 kg/m 3, thermal conductivity is 18.4 W/(m·K), electrical resistance is 5·10 -7 Ohm·m. Barium is paramagnetic; specific magnetic susceptibility 1.9·10 -9 m 3 /kg.
Barium metal quickly oxidizes in air; it is stored in kerosene or under a layer of paraffin. Barium reacts at ordinary temperatures with oxygen, forming barium oxide BaO, and with halogens, forming halides. By calcining BaO in a stream of oxygen or air at 500 °C, BaO 2 peroxide is obtained (decomposes to BaO at 800 °C). Reactions with nitrogen and hydrogen require heating; the reaction products are Ba 3 N 2 nitride and BaH 2 hydride. Barium reacts with water vapor even in the cold; It dissolves vigorously in water, giving Ba(OH)2 hydroxide, which has the properties of alkalis. Barium forms salts with dilute acids. Of the most widely used barium salts, those that are soluble in water are: BaCl 2 chloride and other halides, Ba(NO3)2 nitrate, Ba(ClO3)2 chlorate, Ba(OOCH3)2 acetate, BaS sulfide; poorly soluble - BaS0 4 sulfate, BaCO 3 carbonate, BaCrO 4 chromate. Barium reduces the oxides, halides and sulfides of many metals to the corresponding metal. Barium forms alloys with most metals; sometimes alloys contain intermetallic compounds. Thus, in the Ba - Al system BaAl, BaAl 2, BaAl 4 were found.
Soluble barium salts are toxic; BaSO 4 is practically non-toxic.
Receipt. The main raw material for barium production is barite concentrate (80-95%) BaSO 4, which is reduced with coal, coke or natural flammable gas; the resulting barium sulfide is processed into other salts of this element. BaO is obtained by calcination of barium compounds. Technically pure barium metal (96-98% by weight) is obtained by thermal reduction of BaO oxide with Al powder. By distillation in vacuum, barium is purified to an impurity content of less than 10-4%, and by zone melting - to 10-6%. Another way to obtain barium from BaO is electrolysis of the oxide melt. Small amounts of barium are obtained by reducing beryllate BaBeO 2 at 1300 °C with titanium.
Application. Barium is used as a deoxidizer of copper and lead, as an additive to antifriction alloys, ferrous and non-ferrous metals, as well as to alloys used for the manufacture of printing fonts in order to increase their hardness. Alloys of barium and nickel are used to make electrodes for spark plugs in internal combustion engines and radio tubes. An alloy of barium with aluminum - Alba, containing 56% Ba, is the basis of getters. Metal barium is a material for anodes in chemical current sources. The active part of most thermionic cathodes is barium oxide. Barium peroxide is used as an oxidizing agent, bleaching agent, and in pyrotechnics; previously it was used to regenerate oxygen from CO 2 . Barium hexaferrite BaFe 12 O 19 is a promising material for use in information storage devices; BaFe 12 O 19 is used for the manufacture of permanent magnets. BaSO 4 is introduced into drilling fluids during oil and gas production. Barium titanate BaTiO 3 is one of the most important ferroelectrics. Nuclide 140 Va (β-emitter, T 1/2 12.8 days) is an isotopic tracer used to study barium compounds. Since barium compounds absorb X-ray and γ-radiation well, they are included in the protective materials of X-ray installations and nuclear reactors. BaSO 4 is used as a contrast agent for X-ray studies of the gastrointestinal tract.
Lit. : Akhmetov T. G. Chemistry and technology of barium compounds. M., 1974; Tretyakov Yu.D. and others. Inorganic chemistry. M., 2001.
D. D. Zaitsev, Yu. D. Tretyakov.
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.
