JTO Posted August 10, 2021 Share Posted August 10, 2021 Hi - this is easy to make but how are you supposed to separate the salts, since both are highly soluble in water (unlike when you make potassium chlorate via electrolysis)? Any ideas? I found this clip in Youtube but it makes zero sense to me because of the aforementioned reason. 1 Link to comment Share on other sites More sharing options...
Arthur Posted August 10, 2021 Share Posted August 10, 2021 I think a thorough search of relevant papers on the web would serve you well. Patents are usually available there as are proceedings of learned bodies and technical institutions.Web and literature search is the first place to look for almost everything THEN you have to work out how much you can use and what production method you want to use. Link to comment Share on other sites More sharing options...
AustralianPyromaniac Posted August 22, 2021 Share Posted August 22, 2021 I hadn't actually looked at research until just now, but I have looked at the solubility of the two salts and I just assumed it impossible (or at least impractical and inefficient) in the scale of a single run. Just run as long as possible, destroy side products, evaporate down the solution. The remaining chloride seems not to affect things to any great degree in my experience. A hydrometer can be used to find the exact concentration. In sodium chlorate manufacture (similar problem of high solubility) the liquor is recycled through multiple runs to produce a solution very high in chlorate and only moderately high in chloride. This causes the unfavorable precipitation of chlorate to be favored even with its higher solubility. A similar approach could be taken in the case of barium chlorate, but that video does not show that, so I would assume he is using the first method I mentioned. I think it is a reasonable question as the info isn't broadly available. A literature search with the university's database turned up nothing directly useful. Sodium chlorate is more broadly produced and probably better searching for that if you want more info JTO. The information should be generally applicable to both processes. 23F0F86F81DC1FE1567601E49193C1B7_extra_My_Chlorate.doc.pdf [see pg. 31-34] ie051018r.pdf How would you do it Arthur? AP. 1 Link to comment Share on other sites More sharing options...
sefrez Posted August 24, 2021 Share Posted August 24, 2021 As the concentration of barium chlorate increases, the solubility of barium chloride should go down (as is the case with a sodium chlorate cell.) I'm thinking it would be no problem, but a question I'd have is if somewhere through the process the mutual solubility would have that chloride solubility is low enough to cause anode corrosion problems. With a sodium chlorate cell (and especially potassium chlorate cell due to the low solubility of KClO3) your always able to have enough chloride in solution for things to work out. This makes the production process much simpler. I haven't done much research on barium compounds so I don't really know. Quiet "scared" to do barium chemistry. I'm bothered enough using it in compositions. Perhaps a magnesium sulfate scrubber (or two in series...) on the cell outlet would be a good idea. Link to comment Share on other sites More sharing options...
ZavodKirova Posted September 15, 2021 Share Posted September 15, 2021 It seems to me easier to get barium bromate if you are looking for practical applications for it. Barium chlorate is a rather dangerous oxidizer, more dangerous than potassium chlorate even. For this reason, I decided to do just that.Formula Best Green 80% Ba (BrO2)320 Red Gum+10% Al powder Link to comment Share on other sites More sharing options...
JTO Posted September 22, 2021 Author Share Posted September 22, 2021 As the concentration of barium chlorate increases, the solubility of barium chloride should go down (as is the case with a sodium chlorate cell.) I'm thinking it would be no problem, but a question I'd have is if somewhere through the process the mutual solubility would have that chloride solubility is low enough to cause anode corrosion problems. With a sodium chlorate cell (and especially potassium chlorate cell due to the low solubility of KClO3) your always able to have enough chloride in solution for things to work out. This makes the production process much simpler. I haven't done much research on barium compounds so I don't really know. Quiet "scared" to do barium chemistry. I'm bothered enough using it in compositions. Perhaps a magnesium sulfate scrubber (or two in series...) on the cell outlet would be a good idea.Anode corrosion is indeed a problem, especially since I use 10mm graphite (carbon) rods as anodes :-) They are very cheap and as they erode the carbon precipitate is really easy to remove. Just let the carbon particles precipitate fully, then move the clear electrolyte to another vessel with a syringe. If there are salts among the precipitate, bring the electrolyte to high temperature to dissolve the salts, and then let the carbon particles precipitate. If you want to salvage every last bit of the salts, add some water to the carbon slurry and then repeat the process. I use this this same carbon removal method with my potassium chlorate cell and it works great. I can let the barium chlorate cell run indefinitely, perhaps all or most of the chloride will have turned to chlorate once it has run long enough? Do you think temperature is an issue? I keep the barium chlorate cell in an unheated storage room because it produces quite a bit of chlorine gas which I do not want in our breathing air, and the temperature is getting lower by the day because it is autumn. At some point I plan to test the electrolyte by dipping a piece of paper in it, once it dries it should burn very aggressively if there is a lot of chlorate present. Link to comment Share on other sites More sharing options...
JTO Posted September 22, 2021 Author Share Posted September 22, 2021 "A hydrometer can be used to find the exact concentration."I don't have one, instead I have measured densities by drawing a known volume to a syringe and weighing it. To get better accuracy, I do it with the same syringe using water (1g/cc), and if there is an errors with the water, I use it as a multiplier when measuring the other density. I hadn't actually looked at research until just now, but I have looked at the solubility of the two salts and I just assumed it impossible (or at least impractical and inefficient) in the scale of a single run. Just run as long as possible, destroy side products, evaporate down the solution. The remaining chloride seems not to affect things to any great degree in my experience. A hydrometer can be used to find the exact concentration. In sodium chlorate manufacture (similar problem of high solubility) the liquor is recycled through multiple runs to produce a solution very high in chlorate and only moderately high in chloride. This causes the unfavorable precipitation of chlorate to be favored even with its higher solubility. A similar approach could be taken in the case of barium chlorate, but that video does not show that, so I would assume he is using the first method I mentioned. I think it is a reasonable question as the info isn't broadly available. A literature search with the university's database turned up nothing directly useful. Sodium chlorate is more broadly produced and probably better searching for that if you want more info JTO. The information should be generally applicable to both processes. 23F0F86F81DC1FE1567601E49193C1B7_extra_My_Chlorate.doc.pdf [see pg. 31-34] ie051018r.pdf How would you do it Arthur? AP. Link to comment Share on other sites More sharing options...
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