Barium chloride

{{chembox

| Verifiedfields = changed

| Watchedfields = changed

| verifiedrevid = 476993736

| Name = Barium chloride

| ImageFile = Cotunnite structure.png

| ImageSize =

| ImageFile1 = Barium chloride.jpg

| ImageSize1 =

| OtherNames = {{ubl|Barium dichloride|Barium muriate|Muryate of Barytes{{Cite book | url=https://play.google.com/books/reader?printsec=frontcover&output=reader&id=nKQ-AAAAYAAJ&pg=GBS.PA64 |title = 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|publisher = R. Griffin & Company|year = 1834}}|Neutral barium chloride}}

|Section1={{Chembox Identifiers

| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}

| ChemSpiderID = 23540

| PubChem = 25204

| UNNumber = 1564

| UNII_Ref = {{fdacite|correct|FDA}}

| UNII = 0VK51DA1T2

| UNII2_Ref = {{fdacite|correct|FDA}}

| UNII2 = EL5GJ3U77E

| UNII2_Comment = (dihydrate)

| InChI = 1/Ba.2ClH/h;2*1H/q+2;;/p-2

| SMILES = [Ba+2].[Cl-].[Cl-]

| InChIKey = WDIHJSXYQDMJHN-NUQVWONBAL

| StdInChI_Ref = {{stdinchicite|correct|chemspider}}

| StdInChI = 1S/Ba.2ClH/h;2*1H/q+2;;/p-2

| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}

| StdInChIKey = WDIHJSXYQDMJHN-UHFFFAOYSA-L

| CASNo = 10361-37-2

| CASNo_Ref = {{cascite|correct|CAS}}

| CASNo2_Ref = {{cascite|correct|CAS}}

| CASNo2 = 10326-27-9

| CASNo2_Comment = (dihydrate)

| EINECS = 233-788-1

| RTECS = CQ8750000 (anhydrous)
CQ8751000 (dihydrate)

}}

|Section2={{Chembox Properties

| Formula = {{chem2|BaCl2}}

| MolarMass = 208.23 g/mol (anhydrous)
244.26 g/mol (dihydrate)

| Appearance = White powder, or colourless or white crystals (anhydrous)
Colourless rhomboidal crystals (dihydrate){{cite web | url=https://www.sciencedirect.com/topics/chemistry/barium-chloride | title=Barium Chloride - an overview | ScienceDirect Topics }}{{cite web | url=https://pubchem.ncbi.nlm.nih.gov/compound/Barium-chloride | title=Barium chloride }}

| Odor = Odourless

| Density = 3.856 g/cm³ (anhydrous)
3.0979 g/cm³ (dihydrate)

| Solubility = {{ubl|31.2 g/(100 mL) (0 °C)|35.8 g/(100 mL) (20 °C)|59.4 g/(100 mL) (100 °C)}}

| SolubleOther = Soluble in methanol, insoluble ethyl acetate, slightly soluble in hydrochloric acid and nitric acid, soluble in ethanol.Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990. The dihydrate of barium chloride is soluble in methanol, almost insoluble in ethanol, acetone and ethyl acetate.

| MeltingPtC = 962

| MeltingPt_notes = (960 °C, dihydrate)

| BoilingPtC = 1560

| BoilingPt_notes =

| MagSus = −72.6·10⁻⁶ cm³/mol

}}

|Section3={{Chembox Structure

| Coordination = {{ubl|Of the {{chem2|Ba(2+)}} cations:|8 (the fluorite polymorph)|9 (the cotunnite polymorph)|10 (the post-cotunnite polymorph at pressures of 7–10 GPa)}}

| CrystalStruct = PbCl₂-type orthorhombic (anhydrous)
monoclinic (dihydrate)

}}

|Section4={{Chembox Thermochemistry

| DeltaHf = −858.56 kJ/mol

| Entropy = 123.9 J/(mol·K)

}}

|Section7={{Chembox Hazards

|ExternalSDS = [https://pubchem.ncbi.nlm.nih.gov/compound/Barium-chloride NIH BaCl]

|MainHazards = Highly toxic, corrosive

|NFPA-H = 3

|NFPA-F = 0

|NFPA-R = 0

| GHSPictograms = {{GHS06}}

| GHSSignalWord = Danger

| HPhrases = {{H-phrases|301|302|332}}

| PPhrases = {{P-phrases|261|264|270|271|301+310|304+312|304+340|312|321|330|405|501}}

| FlashPt = Non-flammable

| PEL = TWA 0.5 mg/m³{{PGCH|0045}}

| REL = TWA 0.5 mg/m³

| IDLH = 50 mg/m³

| LD50 = 78 mg/kg (rat, oral)
50 mg/kg (guinea pig, oral){{IDLH|7440393|Barium (soluble compounds, as Ba)}}

| LDLo = 112 mg/kg (as Ba) (rabbit, oral)
59 mg/kg (as Ba) (dog, oral)
46 mg/kg (as Ba) (mouse, oral)

}}

|Section8={{Chembox Related

| OtherAnions = {{ubl|Barium fluoride|Barium bromide|Barium iodide}}

| OtherCations = {{ubl|Beryllium chloride|Magnesium chloride|Calcium chloride|Strontium chloride|Radium chloride|Lead chloride}}

}}

Barium chloride is an inorganic compound with the formula {{chem2|BaCl2|auto=1}}. 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 {{chem2|BaCl2*2H2O}}, which are colourless crystals with a bitter salty taste. It has limited use in the laboratory and industry.{{cite book |author=Kresse, Robert |author2=Baudis, Ulrich |author3=Jäger, Paul |author4=Riechers, H. Hermann |author5=Wagner, Heinz |author6=Winkler, Jocher |author7=Wolf, Hans Uwe |chapter=Barium and Barium Compounds |editor=Ullman, Franz |title=Ullmann's Encyclopedia of Industrial Chemistry |date=2007 |publisher=Wiley-VCH |doi=10.1002/14356007.a03_325.pub2|isbn=978-3527306732 }}

Preparation

On an industrial scale, barium chloride is prepared via a two step process from barite (barium sulfate).{{Greenwood&Earnshaw2nd}} The first step requires high temperatures.

:{{chem2|BaSO4 + 4 C → BaS + 4 CO}}

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

:{{chem2|BaS + 2 HCl → BaCl2 + H2S}}

or between barium sulfide and calcium chloride:

:{{chem2|BaS + CaCl2 → CaS + BaCl2}}

In place of HCl, chlorine can be used. Barium chloride is extracted out from the mixture with water. From water solutions of barium chloride, its dihydrate ({{chem2|BaCl2*2H2O}}) can be crystallized as colorless crystals.

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.

:{{chem2|Ba(OH)2 + 2 HCl → BaCl2 + 2 H2O}}

:{{chem2|BaCO3 + 2 HCl → BaCl2 + H2O + CO2}}

Structure and properties

{{chem2|BaCl2}} crystallizes in two forms (polymorphs). At room temperature, the compound is stable in the orthorhombic cotunnite (Lead(II) chloride) structure, whereas the cubic fluorite structure (calcium fluoride) is stable between 925 and 963 °C.{{cite journal | last1=Edgar | first1=A. | last2=Zimmermann | first2=J. | last3=von Seggern | first3=H. | last4=Varoy | first4=C. R. | title=Lanthanum-stabilized europium-doped cubic barium chloride: An efficient x-ray phosphor | journal=Journal of Applied Physics | publisher=AIP Publishing | volume=107 | issue=8 | date=2010-04-15 | pages=083516–083516–7 | issn=0021-8979 | doi=10.1063/1.3369162 | bibcode=2010JAP...107h3516E }} Both polymorphs accommodate the preference of the large {{chem2|Ba(2+)}} ion for coordination numbers greater than six.Wells, A. F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. {{ISBN|0-19-855370-6}}. The coordination of {{chem2|Ba(2+)}} is 8 in the fluorite structure{{Cite journal | last1 = Haase | first1 = A. | last2 = Brauer | first2 = G.| doi = 10.1002/zaac.19784410120 | title = Hydratstufen und Kristallstrukturen von Bariumchlorid | journal = Z. anorg. allg. Chem. | volume = 441 | pages = 181–195| year = 1978 }} and 9 in the cotunnite structure.{{Cite journal | last1 = Brackett | first1 = E. B. | title = The Crystal Structures of Barium Chloride, Barium Bromide, and Barium Iodide | last2 = Brackett | first2 = T. E. | last3 = Sass | first3 = R. L. | journal = J. Phys. Chem.| volume = 67 | issue = 10 | pages = 2132 | year = 1963 | doi = 10.1021/j100804a038 }} When cotunnite-structure {{chem2|BaCl2}} is subjected to pressures of 7–10 GPa, it transforms to a third structure, a monoclinic post-cotunnite phase. The coordination number of {{chem2|Ba(2+)}} increases from 9 to 10.{{Cite journal | last1 = Léger | first1 = J. M. | last2 = Haines | first2 = J. | last3 = Atouf | first3 = A. | doi = 10.1107/S0021889895001580 | title = The Post-Cotunnite Phase in BaCl₂, BaBr₂ and BaI₂ under High Pressure | journal = J. Appl. Crystallogr.| volume = 28 | issue = 4 | pages = 416 | year = 1995 | bibcode = 1995JApCr..28..416L }}

In aqueous solution {{chem2|BaCl2}} behaves as a simple salt; in water it is a 1:2 electrolyte{{cln|What on this Earth is "1:2 electrolyte"??? Not all readers of this paragraph are experts in electrochemistry, so please, clear this jargon!|date=February 2023}} and the solution exhibits a neutral pH. Its solutions react with sulfate ion to produce a thick white solid precipitate of barium sulfate.

:{{chem2|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:

:{{chem2|BaCl2 + Na2C2O4 → 2 NaCl + BaC2O4}}

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

:{{chem2|BaCl2 + 2 NaOH → 2 NaCl + Ba(OH)2}}

{{chem2|BaCl2*2H2O}} is stable in the air at room temperature, but loses one water of crystallization above {{cvt|55|C|F}}, becoming {{chem2|BaCl2*H2O}}, and becomes anhydrous above {{cvt|121|C|F}}. {{chem2|BaCl2*H2O}} may be formed by shaking the dihydrate with methanol.

{{chem2|BaCl2}} readily forms eutectics with alkali metal chlorides.

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. It is also used to make red pigments such as Lithol red and Red Lake C. Its toxicity limits its applicability.{{cn|date=September 2023}}

Toxicity

Barium chloride, along with other water-soluble barium salts, is highly toxic.The Merck Index, 7th edition, Merck & Co., Rahway, New Jersey, 1960. 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 {{chem2|Ba(2+)}} ions compete with the {{chem2|K+}} ions, causing the muscle fibers to be electrically unexcitable, thus causing weakness and paralysis of the body. 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.

References

External links

[http://www.inchem.org/documents/icsc/icsc/eics0614.htm International Chemical Safety Card 0614]. (anhydrous)
[http://www.inchem.org/documents/icsc/icsc/eics0615.htm International Chemical Safety Card 0615]. (dihydrate)
[https://web.archive.org/web/20060521123455/http://www.solvayvishnubarium.com/products/bariumchloride/0%2C%2C4872-2-0%2C00.htm Barium chloride's use in industry].
[http://chemsub.online.fr/name/barium_chloride.html ChemSub Online: Barium chloride].

Category:Chlorides

Category:Alkaline earth metal halides

Category:Barium compounds

Category:Inorganic compounds

Category:Pyrotechnic colorants

Category:Fluorite crystal structure