barium chlorate
{{Use dmy dates|date=November 2024}}
{{chembox
| Watchedfields = changed
| verifiedrevid = 455115043
| ImageFile = Barium chlorate.svg
| ImageSize = 150px
| ImageFile1 = Bariumchloratepowder.jpg
| ImageAlt1 =
| ImageCaption1 =
| IUPACName = Barium dichlorate
| OtherNames = Chloric acid, barium salt
|Section1={{Chembox Identifiers
| CASNo_Ref = {{cascite|correct|??}}
| CASNo = 13477-00-4
| PubChem = 26059
| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| ChemSpiderID = 24273
| EC_number = 236-760-7
| UNNumber = 1445
| UNII = GRW9DUG818
| SMILES = [Ba+2].[O-]Cl(=O)=O.[O-]Cl(=O)=O
| InChI = 1/Ba.2ClHO3/c;2*2-1(3)4/h;2*(H,2,3,4)/q+2;;/p-2
| InChIKey = ISFLYIRWQDJPDR-NUQVWONBAT
| StdInChI_Ref = {{stdinchicite|correct|chemspider}}
| StdInChI = 1S/Ba.2ClHO3/c;2*2-1(3)4/h;2*(H,2,3,4)/q+2;;/p-2
| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
| StdInChIKey = ISFLYIRWQDJPDR-UHFFFAOYSA-L
| RTECS = FN9770000
}}
|Section2={{Chembox Properties
| Formula = Ba(ClO3)2
| MolarMass = 304.23 g/mol
| Appearance = white solid
| Density = 3.18 g/cm3, solid
| Solubility = 27.5 g/100 ml (20 °C)
| MeltingPtC = 413.9
| MeltingPt_notes = (decomposes)
| BoilingPt =
| MagSus = −87.5·10−6 cm3/mol
}}
|Section7={{Chembox Hazards
| Hazards_ref = {{Sigma-Aldrich|id=244554|name=Barium chlorate|accessdate=6 December 2024}}
| GHSPictograms = {{GHS03}}{{GHS07}}{{GHS09}}
| GHSSignalWord = Danger
| HPhrases = {{H-phrases|271|302|332|411}}
| PPhrases = {{P-phrases|210|220|221|261|264|270|271|273|280|283|301+312|304+312|304+340|306+360|312|330|370+378|371+380+375|391|501}}
| NFPA-H = 3
| NFPA-F = 0
| NFPA-R = 3
| NFPA-S = OX
| LD50 = 500.1 mg/kg
| LC50 = (4h) 1.5 mg/l - dust/mist
| PEL = 0.5 mg/m3 (Vacated)
| IDLH = 50 mg/m3
| ExternalSDS = [http://www.hummelcroton.com/msds/msdsp/baclo3_p.html Barium Chlorate MSDS]
}}
}}
Barium chlorate, Ba(ClO3)2, is the barium salt of chloric acid. It is a white crystalline solid, and like all soluble barium compounds, irritant and toxic. It is sometimes used in pyrotechnics to produce a green colour. It also finds use in the production of chloric acid.
Reactions
=Synthesis=
Barium chlorate can be produced through a double replacement reaction between solutions of barium chloride and sodium chlorate:
:{{chem2|BaCl2 + 2 NaClO3 -> Ba(ClO3)2 + 2 NaCl}}
After concentrating and cooling the resulting mixture, barium chlorate precipitates. This is perhaps the most common preparation, exploiting the lower solubility of barium chlorate compared to sodium chlorate.{{cn|date=September 2023}}
The above method does result in some sodium contamination, which is undesirable for pyrotechnic purposes, where the strong yellow colour of sodium can easily overpower the green of barium. Sodium-free barium chlorate can be produced directly through electrolysis:{{cite web| first = Tom | last= Perigrin | title = Barium Chlorate | publisher = GeoCities | url = http://www.geocities.com/CapeCanaveral/Campus/5361/chlorate/barium.html | accessdate = 2007-02-22 |archiveurl=https://web.archive.org/web/20071030002126/http://www.geocities.com/CapeCanaveral/Campus/5361/chlorate/barium.html|archivedate=2007-10-30}}{{Unreliable source?|date=November 2019}}
:{{chem2|BaCl2 + 6 H2O -> Ba(ClO3)2 + 6 H2}}
It can also be produced by the reaction of barium carbonate with boiling ammonium chlorate solution:{{cite book |last1=Brauer |first1=Georg |last2=Schmeisser |first2=M. |editor1-last=Riley |editor1-first=Reed F. |title=Handbook of Preparative Inorganic Chemistry |date=1963 |publisher=Academic Press |location=New York, London |isbn=9780121266011 |pages=314-315 |edition=2nd |url=https://archive.org/details/Handbook_of_Preparative_Inorganic_Chemistry_1_2_Brauer/page/n337/mode/2up |access-date=6 December 2024 |ref=Brauer |chapter=5. Chlorine, Bromine, Iodine}}
:{{chem2|2 NH4ClO3 + BaCO3 -> Ba(ClO3)2 + 2 NH3 + H2O + CO2}}
The reaction initially produces barium chlorate and ammonium carbonate; boiling the solution decomposes the ammonium carbonate and drives off the resulting ammonia and carbon dioxide, leaving only barium chlorate in solution.
=Decomposition=
When exposed to heat, barium chlorate alone will decompose to barium chloride and oxygen:
:{{chem2|Ba(ClO3)2 -> BaCl2 + 3 O2}}
=Chloric acid=
Barium chlorate is sometimes used to produce chloric acid.{{rp|312-313}}
Commercial uses
When barium chlorate is heated with a fuel, it burns to produce a vibrant green light, which is also a flame test for the presence of bariom ions. Because it is an oxidizer, a chlorine donor, and contains a metal ion, this compound produces a distinctive green colour.{{Citation needed|date=December 2024}} However, due to the instability of all chlorates to sulfur, acids, and ammonium ions, chlorates have been banned from use in class C fireworks in the United States. Therefore, more and more firework producers have begun to use more stable compounds such as barium nitrate and barium carbonate.{{Cite journal | author = Wilson, Elizabeth | title = What's That Stuff? Fireworks | journal = Chemical & Engineering News | volume = 79 | issue = 27 | page = 30 | date = July 2, 2001 | url = https://pubsapp.acs.org/cen/whatstuff/stuff/7927sci3.html }}
Environmental Hazard
Barium chlorate, like all oxidizing agents, is dangerous to human health and is also classed as toxic to the environment. It is very harmful to aquatic organisms if it is leached into bodies of water. Chemical spills of this compound, although not common, can pollute entire ecosystems and should be prevented.{{cite web | url = http://www.inchem.org/documents/icsc/icsc/eics0613.htm | title = Barium Chlorate | publisher = inchem.org }} It is necessary to dispose of this compound as hazardous waste. The Environmental Protection Agency (EPA) lists barium chlorate as hazardous.{{cite web | url =http://nj.gov/health/eoh/rtkweb/documents/fs/0183.pdf | title = Barium Chlorate | work = Hazardous Substance Fact Sheet | publisher = New Jersey Department of Health and Human Services }}