| IUPAC name
| Other names
Chloric acid, barium salt
|Molar mass||304.23 g/mol|
|Melting point||413.9 °C (777.0 °F; 687.0 K) (decomposes)|
| 27.5 g/100 ml (20 °C)|
37.9 g/ 100 ml (25 °C)
|Solubility||Slightly soluble in acetone, ethanol, methanol|
|Vapor pressure||~0 mmHg|
|Safety data sheet||Sigma-Aldrich (monohydrate)|
| Barium chloride|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Barium chlorate is is an extremely powerful oxidizer.
- Ba(ClO3)2 + H2SO4 → 2 HClO3 + BaSO4 (precipitates)
Since barium sulfate is very insoluble, this reaction is useful to obtain chloric acid of high purity: just filter the solution.
If the experiment is run dry (i.e. in a non-aqueous environment), the pure, very unstable chloric acid decomposes immediately to perchloric acid and chlorine dioxide. The latter will spontaneously ignite any combustible material (sugar, paper, dust). However, doing this reaction aqueously lets one synthesize stable solutions of chloric acid.
Barium chlorate always gives the green color to the flame if it is a component of a fuel-oxidizer mix. This is its main use in pyrotechnics.
Barium chlorate is a transparent to white salt that are poorly soluble in water and glycerol. It is even less soluble in cold water than potassium chlorate, which allows for easy conversion between two salts: in a cooled solution of BaCl2 and potassium chlorate, the double displacement reaction proceeds almost fully.
It can be found in some pyrotechnics.
The common precursor chem to barium chlorate is potassium chlorate, in form of the warm solution. You will also need ammonia and tartric acid. Bubble ammonia through a solution of tartaric acid until you get ammonium tartrate, or mix aqueous ammonia (fresh, so you know the concentration) with tartric acid. Add both solutions to barium carbonate and boil until ammonia and carbon dioxide stop emerging. You will notice that the insoluble, chalky barium carbonate dissolves. Ammonia will emerge and briefly liberate chloric acid, which will react with barium carbonate; the insoluble potassium bitartrate will precipitate.
- BaCO3 + 2 (NH4)2C6H2O6 + 2 KClO3 → Ba(ClO3)2 + H2O + CO2(gas) + 2 KHC6H2O6(prec.) + 4NH3(gas)
The same synthesis can be done with any other acid with a soluble ammonium salt and an insoluble potassium salt, such as hexanitritocobaltic. But tartric is just cheaper.
Do not use barium hydroxide, it will lead to an undesirable side reaction: formation of barium tartrate!
When mixed with combustible materials, even those normally slightly flammable (such as dust and lint), it will burn vigorously in combination and the fires are extremely hard to put out, as the chlorate provides the oxygen for the fire. Sulfur and red phosphorus, should be avoided in pyrotechnic compositions containing barium chlorate, as well as any acidic salts, as these mixtures are shock and friction sensitive and prone to spontaneous deflagration (in the safety head matches, such mixture is stabilized with glue).
Like all soluble barium compounds, this one is acutely toxic if ingested!
Barium chlorate should be stored in closed containers and away from any organic sources, as well as strong acidic vapors. Since it is not hygroscopic, it is not necessary to keep it air tight.
Barium chlorate can be neutralized in two steps. First, with a reducing agent, such as sodium metabisulfite, sodium bisulfite, sodium sulfite. Then, the toxic barium remains, and this is neutralized with sulfuric acid. Using sulfites allows combining both steps, since they oxidize to sulfates and neutralize barium in their oxidized form.