Barium chloride

Last updated
Barium chloride
Cotunnite structure.png
Barium chloride.jpg
Names
Other names
  • Barium dichloride
  • Barium muriate
  • Muryate of Barytes [1]
  • Neutral barium chloride
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.030.704 OOjs UI icon edit-ltr-progressive.svg
EC Number
  • 233-788-1
PubChem CID
RTECS number
  • CQ8750000 (anhydrous)
    CQ8751000 (dihydrate)
UNII
UN number 1564
  • InChI=1S/Ba.2ClH/h;2*1H/q+2;;/p-2 Yes check.svgY
    Key: WDIHJSXYQDMJHN-UHFFFAOYSA-L Yes check.svgY
  • InChI=1/Ba.2ClH/h;2*1H/q+2;;/p-2
    Key: WDIHJSXYQDMJHN-NUQVWONBAL
  • [Ba+2].[Cl-].[Cl-]
Properties
BaCl2
Molar mass 208.23 g/mol (anhydrous)
244.26 g/mol (dihydrate)
AppearanceWhite powder, or colourless or white crystals (anhydrous)
Colourless rhomboidal crystals (dihydrate) [2] [3]
Odor Odourless
Density 3.856 g/cm3 (anhydrous)
3.0979 g/cm3 (dihydrate)
Melting point 962 °C (1,764 °F; 1,235 K) (960 °C, dihydrate)
Boiling point 1,560 °C (2,840 °F; 1,830 K)
  • 31.2 g/(100 mL) (0 °C)
  • 35.8 g/(100 mL) (20 °C)
  • 59.4 g/(100 mL) (100 °C)
Solubility Soluble in methanol, insoluble ethyl acetate, slightly soluble in hydrochloric acid and nitric acid, very slightly soluble in ethanol. [4] [3] The dihydrate of barium chloride is soluble in methanol, almost insoluble in ethanol, acetone and ethyl acetate. [3]
−72.6·10−6 cm3/mol
Structure
PbCl2-type orthorhombic (anhydrous)
monoclinic (dihydrate)
  • Of the Ba2+ cations:
  • 8 (the fluorite polymorph)
  • 9 (the cotunnite polymorph)
  • 10 (the post-cotunnite polymorph at pressures of 7–10 GPa)
Thermochemistry
Std molar
entropy (S⦵298)
123.9 J/(mol·K)
−858.56 kJ/mol
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
Highly toxic, corrosive
GHS labelling:
GHS-pictogram-skull.svg
Danger
H301, H302, H332
P261, P264, P270, P271, P301+P310, P304+P312, P304+P340, P312, P321, P330, P405, P501
NFPA 704 (fire diamond)
NFPA 704.svgHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
3
0
0
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
78 mg/kg (rat, oral)
50 mg/kg (guinea pig, oral) [5]
112 mg/kg (as Ba) (rabbit, oral)
59 mg/kg (as Ba) (dog, oral)
46 mg/kg (as Ba) (mouse, oral) [5]
NIOSH (US health exposure limits):
PEL (Permissible)
TWA 0.5 mg/m3 [6]
REL (Recommended)
TWA 0.5 mg/m3 [6]
IDLH (Immediate danger)
50 mg/m3 [6]
Safety data sheet (SDS) NIH BaCl
Related compounds
Other anions
Other cations
Supplementary data page
Barium chloride (data page)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
X mark.svgN  verify  (what is  Yes check.svgYX mark.svgN ?)

Barium chloride is an inorganic compound with the formula Ba Cl 2. 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 BaCl2·2H2O, which are colourless crystals with a bitter salty taste. It has limited use in the laboratory and industry. [7] [3]

Contents

Preparation

On an industrial scale, barium chloride is prepared via a two step process from barite (barium sulfate). [8] The first step requires high temperatures.

BaSO4 + 4 C → BaS + 4 CO

The second step requires reaction between barium sulfide and hydrogen chloride:

BaS + 2 HCl → BaCl2 + H2S

or between barium sulfide and calcium chloride:

BaS + CaCl2 → CaS + BaCl2 [2]

In place of HCl, chlorine can be used. [7] Barium chloride is extracted out from the mixture with water. From water solutions of barium chloride, its dihydrate (BaCl2·2H2O) can be crystallized as colorless crystals. [2]

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.

Ba(OH)2 + 2 HCl → BaCl2 + 2 H2O
BaCO3 + 2 HCl → BaCl2 + H2O + CO2

Structure and properties

BaCl2 crystallizes in two forms (polymorphs). At room temperature, the compound is stable in the orthorhombic cotunnite (PbCl2) structure, whereas the cubic fluorite structure (CaF2) is stable between 925 and 963 °C. [9] Both polymorphs accommodate the preference of the large Ba2+ ion for coordination numbers greater than six. [10] The coordination of Ba2+ is 8 in the fluorite structure [11] and 9 in the cotunnite structure. [12] When cotunnite-structure BaCl2 is subjected to pressures of 7–10 GPa, it transforms to a third structure, a monoclinic post-cotunnite phase. The coordination number of Ba2+ increases from 9 to 10. [13]

In aqueous solution BaCl2 behaves as a simple salt; in water it is a 1:2 electrolyte[ clarification needed ] and the solution exhibits a neutral pH. Its solutions react with sulfate ion to produce a thick white solid precipitate of barium sulfate.

BaCl2 + Na2SO4 → 2 NaCl + BaSO4

This precipitation reaction is used in chlor-alkali plants to control the sulfate concentration in the feed brine for electrolysis.

Oxalate effects a similar reaction:

BaCl2 + Na2C2O4 → 2 NaCl + BaC2O4

When it is mixed with sodium hydroxide, it gives barium hydroxide, which is moderately soluble in water.

BaCl2 + 2 NaOH → 2 NaCl + Ba(OH)2

BaCl2·2H2O is stable in the air at room temperature, but loses one water of crystallization above 55 °C (131 °F), becoming BaCl2·H2O, and becomes anhydrous above 121 °C (250 °F). [2] BaCl2·H2O may be formed by shaking the dihydrate with methanol. [3]

BaCl2 readily forms eutectics with alkali metal chlorides. [3]

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. [7] It is also used to make red pigments such as Lithol red and Red Lake C. Its toxicity limits its applicability.[ citation needed ]

Toxicity

Barium chloride, along with other water-soluble barium salts, is highly toxic. [14] 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 Ba 2+ ions compete with the K + ions, causing the muscle fibers to be electrically unexcitable, thus causing weakness and paralysis of the body. [3] 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. [3]

Related Research Articles

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<span class="mw-page-title-main">Barium carbonate</span> Chemical compound

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<span class="mw-page-title-main">Lead(II) chloride</span> Chemical compound

Lead(II) chloride (PbCl2) is an inorganic compound which is a white solid under ambient conditions. It is poorly soluble in water. Lead(II) chloride is one of the most important lead-based reagents. It also occurs naturally in the form of the mineral cotunnite.

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<span class="mw-page-title-main">Manganese(II) chloride</span> Chemical compound

Manganese(II) chloride is the dichloride salt of manganese, MnCl2. This inorganic chemical exists in the anhydrous form, as well as the dihydrate (MnCl2·2H2O) and tetrahydrate (MnCl2·4H2O), with the tetrahydrate being the most common form. Like many Mn(II) species, these salts are pink, with the paleness of the color being characteristic of transition metal complexes with high spin d5 configurations.

<span class="mw-page-title-main">Strontium chloride</span> Chemical compound

Strontium chloride (SrCl2) is a salt of strontium and chloride. It is a 'typical' salt, forming neutral aqueous solutions. As with all compounds of strontium, this salt emits a bright red colour in flame, and is commonly used in fireworks to that effect. Its properties are intermediate between those for barium chloride, which is more toxic, and calcium chloride.

<span class="mw-page-title-main">Copper(II) chloride</span> Chemical compound

Copper(II) chloride, also known as cupric chloride, is an inorganic compound with the chemical formula CuCl2. The monoclinic yellowish-brown anhydrous form slowly absorbs moisture to form the orthorhombic blue-green dihydrate CuCl2·2H2O, with two water molecules of hydration. It is industrially produced for use as a co-catalyst in the Wacker process.

<span class="mw-page-title-main">Chromium(III) chloride</span> Chemical compound

Chromium(III) chloride (also called chromic chloride) is an inorganic chemical compound with the chemical formula CrCl3. It forms several hydrates with the formula CrCl3·nH2O, among which are hydrates where n can be 5 (chromium(III) chloride pentahydrate CrCl3·5H2O) or 6 (chromium(III) chloride hexahydrate CrCl3·6H2O). The anhydrous compound with the formula CrCl3 are violet crystals, while the most common form of the chromium(III) chloride are the dark green crystals of hexahydrate, CrCl3·6H2O. Chromium chlorides find use as catalysts and as precursors to dyes for wool.

<span class="mw-page-title-main">Nickel(II) chloride</span> Chemical compound

Nickel(II) chloride (or just nickel chloride) is the chemical compound NiCl2. The anhydrous salt is yellow, but the more familiar hydrate NiCl2·6H2O is green. Nickel(II) chloride, in various forms, is the most important source of nickel for chemical synthesis. The nickel chlorides are deliquescent, absorbing moisture from the air to form a solution. Nickel salts have been shown to be carcinogenic to the lungs and nasal passages in cases of long-term inhalation exposure.

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<span class="mw-page-title-main">Barium chlorate</span> Chemical compound

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<span class="mw-page-title-main">Barium bromide</span> Chemical compound

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Radium compounds are compounds containing the element radium (Ra). Due to radium's radioactivity, not many compounds have been well characterized. Solid radium compounds are white as radium ions provide no specific coloring, but they gradually turn yellow and then dark over time due to self-radiolysis from radium's alpha decay. Insoluble radium compounds coprecipitate with all barium, most strontium, and most lead compounds.

References

  1. ↑ Chemical Recreations: A Series of Amusing and Instructive Experiments, which May be Performed with Ease, Safety, Success, and Economy ; to which is Added, the Romance of Chemistry : An Inquiry into the Fallacies of the Prevailing Theory of Chemistry : With a New Theory and a New Nomenclature. R. Griffin & Company. 1834.
  2. 1 2 3 4 "Barium Chloride - an overview | ScienceDirect Topics".
  3. 1 2 3 4 5 6 7 8 "Barium chloride".
  4. ↑ Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
  5. 1 2 "Barium (soluble compounds, as Ba)". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  6. 1 2 3 NIOSH Pocket Guide to Chemical Hazards. "#0045". National Institute for Occupational Safety and Health (NIOSH).
  7. 1 2 3 Kresse, Robert; Baudis, Ulrich; Jäger, Paul; Riechers, H. Hermann; Wagner, Heinz; Winkler, Jocher; Wolf, Hans Uwe (2007). "Barium and Barium Compounds". In Ullman, Franz (ed.). Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH. doi:10.1002/14356007.a03_325.pub2. ISBN   978-3527306732.
  8. ↑ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN   978-0-08-037941-8.
  9. ↑ Edgar, A.; Zimmermann, J.; von Seggern, H.; Varoy, C. R. (2010-04-15). "Lanthanum-stabilized europium-doped cubic barium chloride: An efficient x-ray phosphor". Journal of Applied Physics. 107 (8). AIP Publishing: 083516–083516–7. Bibcode:2010JAP...107h3516E. doi:10.1063/1.3369162. ISSN   0021-8979.
  10. ↑ Wells, A. F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN   0-19-855370-6.
  11. ↑ Haase, A.; Brauer, G. (1978). "Hydratstufen und Kristallstrukturen von Bariumchlorid". Z. anorg. allg. Chem. 441: 181–195. doi:10.1002/zaac.19784410120.
  12. ↑ Brackett, E. B.; Brackett, T. E.; Sass, R. L. (1963). "The Crystal Structures of Barium Chloride, Barium Bromide, and Barium Iodide". J. Phys. Chem. 67 (10): 2132. doi:10.1021/j100804a038.
  13. ↑ Léger, J. M.; Haines, J.; Atouf, A. (1995). "The Post-Cotunnite Phase in BaCl2, BaBr2 and BaI2 under High Pressure". J. Appl. Crystallogr. 28 (4): 416. Bibcode:1995JApCr..28..416L. doi:10.1107/S0021889895001580.
  14. ↑ The Merck Index, 7th edition, Merck & Co., Rahway, New Jersey, 1960.
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