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Author: Subject: barium hydroxide [Ba(OH)2]
Magpie
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[*] posted on 11-3-2006 at 19:07
barium hydroxide [Ba(OH)2]


S.C. Wack has discussed preparation of barium compounds under the ferrate thread. But I think that barium hydroxide is important enough to merit its own thread.

I needed to prepare a few grams of barium hydroxide. The starting compound I have is BaCO3, readily availabe from pottery suppliers.

I tried 2 methods: (1) ignition w/carbon at 1100C to form BaO, and (2) formation of Ba(NO3)2 w/HNO3 then ignition at 500C to BaO. Addition of water to BaO gives Ba(OH)2 or a hydrate thereof. Here is a summary of my results:

1. Ignition of BaCO3 to BaO with finely ground carbon seemed to be successful. Heated to 1025C (highest my muffle furnace will do) for 1 hour. Weight loss was 24.4% vs theoretical 26.8%.
2. Conversion of BaO to Ba(OH)2 not successful due to considerable absorbtion of CO2 from the atmosphere.
3. Conversion of BaCO3 to Ba(NO3)2 using HNO3 was successful using the theoretical amount of HNO3. Yield was 99.6%.
4. Ignition of Ba(NO3)2 to BaO was not successful. The crucible contents frothed over the edge. A glassine material formed in the bottom of the crucibles. Since my furnace temperature indicator doesn't seem to be accurate in the 500C area I heated the furnace until it just stared to show redness (meter indication: 700C).

So, in playing around with the compounds I found that Ba(OH)2 agressively goes after the CO2 in the air. Everything I tested fritzed CO2 when treated w/10%HCl.

My next attempt will be to convert the BaCO3 to BaCl2 and then convert this to Ba(OH)2 using NaOH per the method in Henderson & Fernelius (our library). This includes some steps to prevent absorption of atmospheric CO2.

The attached photo shows the BaO formed after igniting the BaCO3 at 1025C for an hour.

[Edited on 12-3-2006 by Magpie]

[Edited on 12-3-2006 by Magpie]

[Edited on 12-3-2006 by Magpie]

[Edited on 30-1-2007 by chemoleo]

calcined.jpg - 37kB




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12AX7
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[*] posted on 11-3-2006 at 19:17


It's the strongest alkaline earth, no? Beware of glass and porcelain, it'll probably tend to dissolve things, aqueous or glassy state. (Barium is a strong flux, though not much used for toxicity reasons.)

Tim




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Magpie
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[*] posted on 11-3-2006 at 19:25


Yes, it chewed into my porcelain crucibles fairly well. I had a feeling it was fluxing with the wall but thanks for the confirmation. You can see this in the photo.

The Ba(NO3)2 glassine product won't come out of the crucibles. One became brittle and broke when I was cleaning it - this is getting expensive! :mad:

This will be a big advantage of the next method - no ignition. ;)




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[*] posted on 11-3-2006 at 20:29


You can easily precipitate the hydrated form of barium peroxide from a solution of barium chloride with a concentrated H<sub>2</sub>O<sub>2</sub> solution. The peroxide could then be heated to decompose the peroxide to the oxide and subsequently hydrated. Although I think the peroxide is still somewhat resistant to thermal decomposition with respect to many other peroxides. There is a preparation around here for barium peroxide and I have additional properties if you decide to act on this possible method of preparation.



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[*] posted on 12-3-2006 at 05:14


Quote:

You can easily precipitate the hydrated form of barium peroxide from a solution of barium chloride with a concentrated H2O2 solution.

yes, but only in a alkaline solution.
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[*] posted on 12-3-2006 at 10:44


Yes, indeed. Hence my reference to finding the procedure that states such a thing. I intended for the statement to be made as an overview, that's all. Additionally it should be noted that trying to pyrolyze barium peroxide might turn into quite the mess, wouldn't it be quite corrosive at high temperatures.



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[*] posted on 12-3-2006 at 15:15


In reference to the glassine product left in the crucibles after igniting the Ba(NO3)2: It completely softened and turned white after I left it covered in water overnight. So I believe that the glassine material was BaO. What I found today was likely Ba(OH)2*8H2O.



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[*] posted on 13-3-2006 at 08:00


BaO is converted to a less reactive form at high temperatures, due to the particles fusing together. This BaO reacts only slowly with water.

A solution of Ba(OH)2 becomes turbid in air in a matter of minutes, due to CO2 being absorbed. You should filter it to get out the barium carbonate.
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[*] posted on 13-3-2006 at 10:46


Quote:
A solution of Ba(OH)2 becomes turbid in air in a matter of minutes, due to CO2 being absorbed. You should filter it to get out the barium carbonate.


Tell me about it. :P When I filter my Ba(OH)2 solutions they develop a carbonate layer on top in a matter of *seconds*. Thankfully I'm reacting them with stronger acids than H2CO3, so it shouldn' t matter a lot...
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[*] posted on 17-3-2006 at 19:15


Making Ba(OH)2*8H20 from BaCl2 according to the method in Henderson & Fernelius worked well and I ended up with some nice looking, but wet, crystals. Not able to leave well enough alone I tried to dry these by wrapping them in a paper towel for 24 hours (per another book in our library, Walton IIRC). This didn't work out too well as about half of the crystals effloresced. They also fizzed with 10%HCl whereas the non-effloresced crystals did not. (I thought about separating the crystals by hand then quickly decided that was a bad idea.)

The Henderson & Fernelius procedure for final drying of Ba(OH)2*8H20 to Ba(OH)2 was to place it in a dessicator with - you guessed it - effloresced Ba(OH)2! :D

Overall, I'd have to say that working with barium salts kicked my ass. My BaCO3 and Ba(NO3)2 ignitions weren't too practical. And when I did get Ba(OH)2 I had to use air purification and sealing to prevent absorbtion of CO2. Although I now think that just working quickly would have worked just as well. But there is more to it than that. The Ba(OH)2 adheres something fierce to all glassware. Then when it absorbs the CO2 you have white stubborn stains. This was removeable with H2SO4 cleaning solution, however.

[Edited on 18-3-2006 by Magpie]




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[*] posted on 20-12-2015 at 14:19
Huge swing in BaOH2 in water


I am pretty sure I mentioned this elsewhere on this forum, but BaOH2 is like 100 times more soluble in boiling water than cold water. Boil water to dissolve barium hydroxide, filter out carbonate, then cool to crystallize octahydrate. Crystals remind me of marble.



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[*] posted on 21-12-2015 at 00:51


I have had trouble with making barium hydroxide by decomposing carbonate/nitrate. The best method I found was double decomposition with carbonate free alkali. This gives a hydroxide precipitate which will still absorb CO2 pretty rapidly. Dissolve the precipitate in hot water, filter and cool. Most of the hydroxide will precipitate out.



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[*] posted on 31-8-2018 at 23:41


I made some barium hydroxide from barium carbonate using an idea posted by clearly_not_atara (who probably did not intend using impure pottery grade starting materials). He suggested making barium hydroxide by dissolving barium carbonate in dilute acetic acid and then adding sodium hydroxide.

I simply extracted an excess of barium carbonate with 1L of 5% household vinegar in a well-ventilated area, filtered, and precipitated barium hydroxide with sodium hydroxide. Then I heated the mixture to boiling, causing the barium hydroxide to dissolve and did a hot filtration through a coffee filter on top of a cotton ball, placing the flask in a boiling water bath and guiding the steam around the funnel with aluminum foil. At times I used a heat gun to make the mixture in the funnel boil.

Once the filtration was complete, I chilled the filtrate in the refrigerator and then vacuum filtered. The mixture wasn't easy to filter, and to make sure things stayed cold, I dropped a couple of ice cubes in the filter funnel. I also covered it with plastic wrap to keep out carbon dioxide to a large extent. I washed the filter cake a few times with ice cold water.

Here is a picture of the crude product. It is contaminated with a little bit of an extremely fine, brown, silty material that passes through the filter paper (perhaps sulfur?).

baryta.jpg - 84kB

The yield I'm sure is not very good, and it will be even worse once I recrystallize, but I'm looking forward to having some barium hydroxide on hand for tests and experiments.





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[*] posted on 1-9-2018 at 21:39


Quote: Originally posted by Magpie  
Yes, it chewed into my porcelain crucibles fairly well. I had a feeling it was fluxing with the wall but thanks for the confirmation. You can see this in the photo.

The Ba(NO3)2 glassine product won't come out of the crucibles. One became brittle and broke when I was cleaning it - this is getting expensive! :mad:

This will be a big advantage of the next method - no ignition. ;)


It might be better to consider ceramic crucibles as single use, disposable. They are cheap to buy in bulk and not worth the effort to clean. Platinum ones on the other hand...
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[*] posted on 1-9-2018 at 22:03


After researching the matter, I think the brown, silty material is most likely iron pyrite. Who knows for sure... I'm thinking about trying the synthesis again on larger scale with hydrochloric acid in a 5-gallon bucket outside.



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[*] posted on 2-9-2018 at 07:46


I'd be surprised if the pottery grade material doesn't contain at least some Strontium which has very similar chemistry. It also precipitates its hydroxide as the octahyrate from water, but it has lower solubility than barium. So, any impurities are likely concentrated through this method. Strontium nitrate on the other hand, is way more soluble than Barium Nitrate. Perhaps the nitrate can be recrystallized as a precursor to the hydroxide.
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[*] posted on 2-9-2018 at 12:34


Wouldn't a hot filtration remove most of the strontium hydroxide?



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[*] posted on 4-9-2018 at 07:55


Ignition of Barium salts (carbonate or nitrate) requires very high temperatures and is not at all easy. I have found that the best method is to add carbonate free caustic alkali to a solution of a soluble barium salt and filter the precipitate off. It can be washed or recrystallised to purify it. Carbonate free alkali can be prepared by dissolving sodium or preferably potassium hydroxide in alcohol and allowing to settle. Any chloride or carbonate will not dissolve. This can then be added to boiled out water to give carbonate free alkali.



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[*] posted on 4-9-2018 at 08:10


It is possible to differentiate qualitatively between barium hydroxide and strontium hydroxide with flame testing. There's also a test based on rhodizonic acid, which can be made by nitric acid oxidation of inositol.

(I should probably mention that in the procedure above I boiled down the solution/mixture a couple of times.)



[Edited on 4-9-2018 by JJay]




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[*] posted on 13-9-2018 at 07:19


Quote: Originally posted by 12AX7  
It's the strongest alkaline earth, no? Beware of glass and porcelain, it'll probably tend to dissolve things, aqueous or glassy state. (Barium is a strong flux, though not much used for toxicity reasons.)

Tim


Just recently I blurred my plastic juice jug all over just because of handling it into an acidic barium salt batch!
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[*] posted on 18-9-2018 at 07:14


Hi Magpie,

Are you in the US?

I have maybe 75 or 100 g of BaO. You can have whatever you need for the cost of postage.





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[*] posted on 20-9-2018 at 05:15


Off topic:
Was it barium chromate that had that disgusting yellow/green coloration reminishing of snot?
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[*] posted on 29-9-2018 at 06:57


No thats iron phosphate lol
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[*] posted on 29-9-2018 at 08:50


BTW it is said the nitrate requires 900C for some hours, then crushing and reheating.



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