View source for 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<ref>{{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}}</ref>|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)<br>CQ8751000 (dihydrate)
  }}
|Section2={{Chembox Properties
| Formula = {{chem2|BaCl2}}
| MolarMass = 208.23&nbsp;g/mol (anhydrous)<br>244.26&nbsp;g/mol (dihydrate)
| Appearance = White powder, or colourless or white crystals (anhydrous)<br>Colourless rhomboidal crystals (dihydrate)<ref name="sciencedirect">{{cite web | url=https://www.sciencedirect.com/topics/chemistry/barium-chloride | title=Barium Chloride - an overview &#124; ScienceDirect Topics }}</ref><ref name="pubchem">{{cite web | url=https://pubchem.ncbi.nlm.nih.gov/compound/Barium-chloride | title=Barium chloride }}</ref>
| Odor = Odourless
| Density = 3.856&nbsp;g/cm<sup>3</sup> (anhydrous)<br>3.0979&nbsp;g/cm<sup>3</sup> (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]], very slightly soluble in [[ethanol]].<ref>''Handbook of Chemistry and Physics'', 71st edition, CRC Press, Ann Arbor, Michigan, 1990.</ref><ref name="pubchem"></ref> The dihydrate of barium chloride is soluble in methanol, almost insoluble in ethanol, [[acetone]] and ethyl acetate.<ref name="pubchem"></ref>
| MeltingPtC = 962
| MeltingPt_notes = (960&nbsp;°C, dihydrate)
| BoilingPtC = 1560
| BoilingPt_notes =
| MagSus = −72.6·10<sup>−6</sup> cm<sup>3</sup>/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 = [[Lead(II) chloride|PbCl<sub>2</sub>]]-type [[orthorhombic]] (anhydrous)<br>[[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/m<sup>3</sup><ref name=PGCH>{{PGCH|0045}}</ref>
| REL = TWA 0.5 mg/m<sup>3</sup><ref name=PGCH/>
| IDLH = 50 mg/m<sup>3</sup><ref name=PGCH/>
| LD50 = 78 mg/kg (rat, oral)<br />50 mg/kg (guinea pig, oral)<ref name=IDLH>{{IDLH|7440393|Barium (soluble compounds, as Ba)}}</ref>
| LDLo = 112 mg/kg (as Ba) (rabbit, oral)<br />59 mg/kg (as Ba) (dog, oral)<br />46 mg/kg (as Ba) (mouse, oral)<ref name=IDLH/>
  }}
|Section8={{Chembox Related
| OtherAnions = {{ubl|[[Barium fluoride]]|[[Barium bromide]]|[[Barium iodide]]}}
| OtherCations = {{ubl|[[Beryllium chloride]]|[[Magnesium chloride]]|[[Calcium chloride]]|[[Strontium chloride]]|[[Radium chloride]]|[[Lead(II) chloride|Lead chloride]]}}
  }}
}}
'''Barium chloride''' is an [[inorganic compound]] with the [[chemical formula|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.<ref name="Ullman2005">{{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 }}</ref><ref name="pubchem"></ref>

==Preparation==
On an industrial scale, barium chloride is prepared via a two step process from [[barite]] ([[barium sulfate]]).<ref>{{Greenwood&Earnshaw2nd}}</ref> 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}}<ref name="sciencedirect"></ref>
In place of HCl, [[chlorine]] can be used.<ref name="Ullman2005"/> 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.<ref name="sciencedirect"></ref>

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 ([[Polymorphism (materials science)|polymorphs]]). At room temperature, the compound is stable in the [[Orthorhombic crystal system|orthorhombic]] [[cotunnite]] ([[Lead(II) chloride|{{chem2|PbCl2}}]]) structure, whereas the cubic [[fluorite structure]] ([[calcium fluoride|{{chem2|CaF2}}]]) is stable between 925 and 963 °C.<ref>{{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 }}</ref> Both [[Polymorphism (materials science)|polymorphs]] accommodate the preference of the large {{chem2|Ba(2+)}} ion for [[coordination number]]s greater than six.<ref>Wells, A. F. (1984) ''Structural Inorganic Chemistry'', Oxford: Clarendon Press. {{ISBN|0-19-855370-6}}.</ref> The coordination of {{chem2|Ba(2+)}} is 8 in the fluorite structure<ref>{{Cite journal | last1 = Haase | first1 = A. | last2 = Brauer | first2 = G.| doi = 10.1002/zaac.19784410120 | title = Hydratstufen und Kristallstrukturen von Bariumchlorid | journal = [[Zeitschrift für anorganische und allgemeine Chemie|Z. anorg. allg. Chem.]] | volume = 441 | pages = 181–195| year = 1978 }}</ref> and 9 in the cotunnite structure.<ref>{{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 = [[Journal of Physical Chemistry|J. Phys. Chem.]]| volume = 67 | issue = 10 | pages = 2132 | year = 1963 | doi = 10.1021/j100804a038 }}</ref> When cotunnite-structure {{chem2|BaCl2}} is subjected to pressures of 7–10&nbsp;GPa, it transforms to a third structure, a [[Monoclinic crystal system|monoclinic]] post-cotunnite phase. The coordination number of {{chem2|Ba(2+)}} increases from 9 to 10.<ref>{{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 BaCl<sub>2</sub>, BaBr<sub>2</sub> and BaI<sub>2</sub> under High Pressure | journal = [[Journal of Applied Crystallography|J. Appl. Crystallogr.]]| volume = 28 | issue = 4 | pages = 416 | year = 1995 | bibcode = 1995JApCr..28..416L }}</ref>

In aqueous solution {{chem2|BaCl2}} behaves as a simple [[Salt (chemistry)|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 + [[Sodium oxalate|Na2C2O4]] → 2 NaCl + [[Barium oxalate|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}}.<ref name="sciencedirect"></ref> {{chem2|BaCl2*H2O}} may be formed by shaking the dihydrate with [[methanol]].<ref name="pubchem"></ref>

{{chem2|BaCl2}} readily forms [[eutectics]] with [[alkali metal]] chlorides.<ref name="pubchem"></ref>

==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]].<ref name="Ullman2005" /> 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.<ref>''The Merck Index'', 7th edition, Merck & Co., Rahway, New Jersey, 1960.</ref> 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|[[Barium|Ba]](2+)}} ions compete with the {{chem2|[[potassium|K]]+}} ions, causing the muscle fibers to be electrically unexcitable, thus causing weakness and paralysis of the body.<ref name="pubchem"></ref> [[Sodium sulfate]] and [[magnesium sulfate]] are potential antidotes because they form barium sulfate BaSO<sub>4</sub>, which is relatively non-toxic because of its insolubility in water.

Barium chloride is not classified as a human carcinogen.<ref name="pubchem"></ref>

==References==
<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].

{{Barium compounds}}
{{Chlorides}}

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[[Category:Chlorides]]
[[Category:Alkaline earth metal halides]]
[[Category:Barium compounds]]
[[Category:Inorganic compounds]]
[[Category:Pyrotechnic colorants]]
[[Category:Fluorite crystal structure]]