Английская Википедия:Barium chloride
Шаблон:Chembox Barium chloride is an inorganic compound with the formula Шаблон:Chem2. It is one of the most common water-soluble salts of barium. Like most other water-soluble barium salts, it is a white powder, highly toxic, and imparts a yellow-green coloration to a flame. It is also hygroscopic, converting to the dihydrate Шаблон:Chem2, which are colourless crystals with a bitter salty taste. It has limited use in the laboratory and industry.[1][2]
Preparation
On an industrial scale, barium chloride is prepared via a two step process from barite (barium sulfate).[3] The first step requires high temperatures.
The second step requires reaction between barium sulfide and hydrogen chloride:
or between barium sulfide and calcium chloride:
In place of HCl, chlorine can be used.[1] Barium chloride is extracted out from the mixture with water. From water solutions of barium chloride, its dihydrate (Шаблон:Chem2) can be crystallized as colorless crystals.[4]
Barium chloride can in principle be prepared by the reaction between barium hydroxide or barium carbonate with hydrogen chloride. These basic salts react with hydrochloric acid to give hydrated barium chloride.
Structure and properties
Шаблон:Chem2 crystallizes in two forms (polymorphs). At room temperature, the compound is stable in the orthorhombic cotunnite ([[Lead(II) chloride|Шаблон:Chem2]]) structure, whereas the cubic fluorite structure ([[calcium fluoride|Шаблон:Chem2]]) is stable between 925 and 963 °C.[5] Both polymorphs accommodate the preference of the large Шаблон:Chem2 ion for coordination numbers greater than six.[6] The coordination of Шаблон:Chem2 is 8 in the fluorite structure[7] and 9 in the cotunnite structure.[8] When cotunnite-structure Шаблон:Chem2 is subjected to pressures of 7–10 GPa, it transforms to a third structure, a monoclinic post-cotunnite phase. The coordination number of Шаблон:Chem2 increases from 9 to 10.[9]
In aqueous solution Шаблон:Chem2 behaves as a simple salt; in water it is a 1:2 electrolyteШаблон:Cln and the solution exhibits a neutral pH. Its solutions react with sulfate ion to produce a thick white precipitate of barium sulfate.
Oxalate effects a similar reaction:
When it is mixed with sodium hydroxide, it gives barium hydroxide, which is moderately soluble in water.
Шаблон:Chem2 is stable in the air at room temperature, but loses one water of crystallization above Шаблон:Cvt, becoming Шаблон:Chem2, and becomes anhydrous above Шаблон:Cvt.[4] Шаблон:Chem2 may be formed by shaking the dihydrate with methanol.[2]
Шаблон:Chem2 readily forms eutectics with alkali metal chlorides.[2]
Uses
Although inexpensive, barium chloride finds limited applications in the laboratory and industry.
Its main laboratory use is as a reagent for the gravimetric determination of sulfates. The sulfate compound being analyzed is dissolved in water and hydrochloric acid is added. When barium chloride solution is added, the sulfate present precipitates as barium sulfate, which is then filtered through ashless filter paper. The paper is burned off in a muffle furnace, the resulting barium sulfate is weighed, and the purity of the sulfate compound is thus calculated.
In industry, barium chloride is mainly used in the purification of brine solution in caustic chlorine plants and also in the manufacture of heat treatment salts, case hardening of steel.[1] It is also used to make red pigments such as Lithol red and Red Lake C. Its toxicity limits its applicability.Шаблон:Cn
Toxicity
Barium chloride, along with other water-soluble barium salts, is highly toxic.[10] It irritates eyes and skin, causing redness and pain. It damages kidneys. Fatal dose of barium chloride for a human has been reported to be about 0.8-0.9 g. Systemic effects of acute barium chloride toxicity include abdominal pain, diarrhea, nausea, vomiting, cardiac arrhythmia, muscular paralysis, and death. The Шаблон:Chem2 ions compete with the Шаблон:Chem2 ions, causing the muscle fibers to be electrically unexcitable, thus causing weakness and paralysis of the body.[2] Sodium sulfate and magnesium sulfate are potential antidotes because they form barium sulfate BaSO4, which is relatively non-toxic because of its insolubility in water.
Barium chloride is not classified as a human carcinogen.[2]
References
- ↑ 1,0 1,1 1,2 Шаблон:Cite book
- ↑ 2,0 2,1 2,2 2,3 2,4 Ошибка цитирования Неверный тег
<ref>; для сносокpubchemне указан текст - ↑ Шаблон:Greenwood&Earnshaw2nd
- ↑ 4,0 4,1 4,2 Ошибка цитирования Неверный тег
<ref>; для сносокsciencedirectне указан текст - ↑ Шаблон:Cite journal
- ↑ Wells, A. F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. Шаблон:ISBN.
- ↑ Шаблон:Cite journal
- ↑ Шаблон:Cite journal
- ↑ Шаблон:Cite journal
- ↑ The Merck Index, 7th edition, Merck & Co., Rahway, New Jersey, 1960.
External links
- International Chemical Safety Card 0614. (anhydrous)
- International Chemical Safety Card 0615. (dihydrate)
- Barium chloride's use in industry.
- ChemSub Online: Barium chloride.
Шаблон:Barium compounds Шаблон:Chlorides
- Страницы с ошибками в примечаниях
- Английская Википедия
- Chlorides
- Alkaline earth metal halides
- Barium compounds
- Inorganic compounds
- Pyrotechnic colorants
- Fluorite crystal structure
- Страницы, где используется шаблон "Навигационная таблица/Телепорт"
- Страницы с телепортом
- Википедия
- Статья из Википедии
- Статья из Английской Википедии
