Bju90 Posted March 14 Posted March 14 500g Barium Carbonate (technical grade 99%+ purity advertised) 310ml Nitric Acid 68% (Technical Grade) diluted in 500ml water Gradually added the barium carbonate to the diluted nitric acid. Bubbling was noted. Continued stirring until bubbling subsided. Was left for 30mins+. I noted there was a small layer of clear ontop of a large volume of white suspended powder. Filtered off the clear layer, test strip showed its neutral ph. I figured the large volume, 500ml+, was unreacted Barium Carbonate. I added some more nitric acid and more bubbling was noted. I’ve stirred it and allowed time for the reaction to finish but the test strips are still indicating a very low ph. I’m confused, the white powder makes me think there is still unreacted barium carbonate but the high Ph suggest there is still nitric acid present? Also noted there was a slight smell of Sulfur indicating h2s. I’m in a well ventilated area using a respirator, although I’ll have to make sure it filters out h2s. Can trace amounts be harmful?
Bju90 Posted March 14 Author Posted March 14 Respirator is A1 + P2 so doesn’t look like it’s effective for filtering H2S so might have to rethink this. The smell was very weak and I had no ill reaction. Not sure if this is ok/expected to continue? The reacting has subsided. Extracted some of the white powder and added a drop of nitric acid - no reaction. Extracted some of the clear liquid and added a drop to bicarb soda - visible reaction. So looks like all of the barium carbonate has reacted… what is the white powder left? I’ve also used up all the barium carbonate I had, what to do now lol
mx5kevin Posted March 14 Posted March 14 For boiling water must add first in a beaker the nitric acid than the barium carbonate. Then gradually add the two components. And it should not be allowed to crystallize until the reaction is complete. At the end of the reaction, a pH test, followed by filtration and crystallization. If barium nitrate precipitates during the reaction, it is not good. But dissolving by boiling will fix it, what doesn't dissolve must be filtered. Since barium carbonate is poorly soluble, a small excess of it can be filtered out easily. During recrystallization, a large layer of water should remain on top. Heat the finished product thoroughly during drying. If you have used too much acid, you will need to heat it up into a dry powder and then recrystallize it in pure water without carbonate.
mx5kevin Posted March 14 Posted March 14 2 hours ago, Bju90 said: Respirator is A1 + P2 so doesn’t look like it’s effective for filtering H2S so might have to rethink this. The smell was very weak and I had no ill reaction. Not sure if this is ok/expected to continue? The reacting has subsided. Extracted some of the white powder and added a drop of nitric acid - no reaction. Extracted some of the clear liquid and added a drop to bicarb soda - visible reaction. So looks like all of the barium carbonate has reacted… what is the white powder left? I’ve also used up all the barium carbonate I had, what to do now lol I would dissolve this completely in hot water. If it has dissolved and you don't have any more barium carbonate, you can use ammonium bicarbonate. This theoretically won't form insoluble salts that cause pollution, it's used in baking, you can find it in baking powders (but this is not backing powder) line in the store. Then filter and crystallize the crystals in the solution. Filter and recrystallize again in pure water. Then, if you dry it, you get pure, and acid-free Ba(NO3)2.
Bju90 Posted March 14 Author Posted March 14 (edited) Hmm ok I think I made a mistake thinking the solubility of barium nitrate in water was much higher than it is, I think I understand whats happened now. Is your choice for Ammonium carbonate to produce ammonium nitrate because of its very high solubility in water so the barium will precipitate out first? I had considered along the same lines by using sodium carbonate. I need to order more chems anyway so it’s probably easier to store this and purchase some more Barium carbonate to finish and not have to worry about removing another nitrate salt Edited March 14 by Bju90
mx5kevin Posted March 14 Posted March 14 3 hours ago, Bju90 said: Hmm ok I think I made a mistake thinking the solubility of barium nitrate in water was much higher than it is, I think I understand whats happened now. Is your choice for Ammonium carbonate to produce ammonium nitrate because of its very high solubility in water so the barium will precipitate out first? I had considered along the same lines by using sodium carbonate. I need to order more chems anyway so it’s probably easier to store this and purchase some more Barium carbonate to finish and not have to worry about removing another nitrate salt Ammonium bicarbonate will not contaminating the product and can be easily removed (unless there was a very significant overdose of nitric acid). Sodium carbonate will contaminated the product, and it may not work in a colorful composition. Sodium contamination can cause a strong yellow color. I think what you made will be enough for a while. For boiling required a beaker. Or in worse case, wrap a glass jar in foil and heat it in a sand bath. Such an acidic solution cannot be placed in a metal container or heated. Where I live, you can buy NH4HCO3 any supermarket in a 15-gram bag. 1
Mumbles Posted March 14 Posted March 14 Your initial amounts didn't have enough nitric acid. 500g barium carbonate requires about 469g 68% nitric acid. You were about 30g short assuming your 68% concentration is correct. I'm assuming you used the density of 1.51 (pure nitric acid), instead of 1.41 (68% acid). The sulfur smell is small amounts of barium sulfide present being hydrolyzed. Any residual insoluble material after acidification is probably barium sulfate. Both are common trace impurities in barium carbonate. As for the pH, it doesn't take much excess acid to drop the pH fast. An extra 0.1mL will drop the pH to about 3 for instance.
Bju90 Posted March 15 Author Posted March 15 6 hours ago, Mumbles said: Your initial amounts didn't have enough nitric acid. 500g barium carbonate requires about 469g 68% nitric acid. You were about 30g short assuming your 68% concentration is correct. I'm assuming you used the density of 1.51 (pure nitric acid), instead of 1.41 (68% acid). The sulfur smell is small amounts of barium sulfide present being hydrolyzed. Any residual insoluble material after acidification is probably barium sulfate. Both are common trace impurities in barium carbonate. As for the pH, it doesn't take much excess acid to drop the pH fast. An extra 0.1mL will drop the pH to about 3 for instance. I’m going to be honest, I am slack and used chatgpt for the calculations. I understand chemistry at a first year undergrad level but my education is in Mech Engineering so I’m expanding my knowledge as I go. i can’t find the chat with the calc but I believe the idea was to use a slight excess of Barium Carbonate to ensure all of the nitric acid has reacts. Then filter out the excess Barium Carbonate since it’s insoluble.
Bju90 Posted March 15 Author Posted March 15 9 hours ago, mx5kevin said: Ammonium bicarbonate will not contaminating the product and can be easily removed (unless there was a very significant overdose of nitric acid). Sodium carbonate will contaminated the product, and it may not work in a colorful composition. Sodium contamination can cause a strong yellow color. I think what you made will be enough for a while. For boiling required a beaker. Or in worse case, wrap a glass jar in foil and heat it in a sand bath. Such an acidic solution cannot be placed in a metal container or heated. Where I live, you can buy NH4HCO3 any supermarket in a 15-gram bag. Ok that makes sense thanks for all the help! I think the only other thing that crossed my mind was the lower breakdown temp of Ammonium Nitrate when boiling.
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