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Squall
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[*] posted on 22-8-2007 at 16:49
Barium Nitrate


I would like to know if it is possible to make barium nitrate from barium carbonate. I have heard of procedures involving ammonium nitrate but i have no details on how it is carried out.
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Xenoid
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[*] posted on 22-8-2007 at 17:17


First dissolve your barium carbonate in hydrochloric acid to give BaCl2 solution.
Barium nitrate is the least soluble of the common nitrates (9g/100g water) so mixing hot saturated BaCl2 solution with saturated ammonium, potassium, sodium or calcium nitrate solutions should allow you to crystallise out Barium Nitrate by double dissolution. Normally sodium nitrate is used.

If it's pottery grade barium carbonate, make sure you do the acid dissolution in a well ventilated area or preferably outside, as copious hydrogen sulphide is evolved from impurities.

Take note - soluble barium compounds are POISONOUS.

Regards,

Xenoid

[Edited on 22-8-2007 by Xenoid]
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[*] posted on 22-8-2007 at 17:35


Boiling strong solution of ammonium nitrate with the hydroxides or carbonates of the alkali or alkaline earth metals will give you the nitrates of those metals. Start with a slight excess of the NH4NO3, bring to a gentle boil - a simmer. Add water (deionized/distilled) as needed to keep the mix thin and fluid. After it stops giving off ammonia, add water to get things fully in solution, then add enough of the hydroxide/carbonate to be in slight excess. Continue to simmer for a few more minutes, filter while hot to remove the excess carbonate and trash. Exactly what should be done to get optimal crystallisation is going to depend on which nitrate you're dealing with.
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Fleaker
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[*] posted on 22-8-2007 at 23:05


@ Xenoid, would it not be easier to just dissolve BaCO3 in dilute HNO3 or am I missing something obvious?



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[*] posted on 22-8-2007 at 23:09


Nothing wrong there, its how I made my Sr(NO3)2, I imagine AN is more available than nitric acid however. (Ironically, not for me). Solubility may be an issue for concentrated solutions

There is also the dry method of heating a powdered mix of BaCO3 and AN, CO2 and NH3 and H2O is given off leaving the nitrate salt which is then recrystalized.

[Edited on 23-8-2007 by The_Davster]




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Xenoid
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[*] posted on 22-8-2007 at 23:36


Quote:
Originally posted by Fleaker
@ Xenoid, would it not be easier to just dissolve BaCO3 in dilute HNO3 or am I missing something obvious?


Yeah! Good one! Whoops!

I naturally assumed he would not have nitric acid (the holy grail), I don't either!
Whereas hydrochloric acid is fairly readily available in most hardware stores.

Also, once you have your Ba in solution as the chloride, it's in a somewhat more useful form than the carbonate. I made up a large batch of concentrated solution, a fairly smelly procedure with pottery barium carbonate.

Regards, Xenoid

[Edited on 22-8-2007 by Xenoid]
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Squall
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[*] posted on 23-8-2007 at 17:21


Thanks for the tips sounds like i will probably try the dissolving Carbonate in HCl first. And yes xenoid you assumed right nitric acid is a luxury that i have not come upon yet but would like to get it. Maybe I'll try to make it some day.
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[*] posted on 25-8-2007 at 14:19


after i dissolve the carbonate do i need to isolate BaCl2 from the solution before i add the nitrate solution. And can anyone refer me to where a can find solubility curves for barium compounds.
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[*] posted on 25-8-2007 at 15:56


BaCO3 + 2HCl --> BaCl2 + H2O + CO2

197.4g BaCO3 will react with 2 x 36.5g (73g)HCl to give 208.2g BaCl2

When the reaction is complete you will end up with a neutral solution of BaCl2 in water.

The amount of hydrochloric acid you use will depend on its strength, it will be written on the side of the container, it's usually about 300g/litre. I'll leave you to work out what volume of hydrochloric acid to use. You may need a bit more or less than theoretical, because of impurities. Make sure it has stopped "fizzing". You can of course use multiples of the above weights!

You can then use this, or perhaps "boil it down" a bit to concentrate it for the reaction with nitrate.

Regards, Xenoid
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[*] posted on 25-8-2007 at 16:32


Thanks for the info i have worked out how much of everything i need i took a small scale approach and used about 5g of BaCO3 to start off. I then added hydrochloric acid which i calculated. Mine was 37% so i hope my math was right. I assumed that there was 37g in every 100ml not sure if thats right. Anyway i then made a solution of KNO3 also calculating the amount i would need. After mixing the two solutions and heating for a little bit i have a precipitant on the bottom of my beaker. I would like to know if there is a way to find out if it is Ba(NO3)2. And a note on the reaction mixing BaCO3 with HCL produces a nasty smell man it stunk up my garage.
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[*] posted on 25-8-2007 at 16:47


Quote:
Originally posted by Squall
And a note on the reaction mixing BaCO3 with HCL produces a nasty smell man it stunk up my garage.


Yeah! Most BaCO3 is made from BaSO4 (Barytes). BaS is an intermediate product and contaminates low quality BaCO3. H2S (hydrogen sulphide, aka. rotten egg gas) is produced when reacted with acids. H2S is very poisonous as are BaCl2 and BaNO3. Take care and wash your hands.
The precipitant may just be impurites. Cool the solution in a fridge at say 0 oC. You should get crystals forming.
You could mix it with charcoal and sulphur and see if you get a greenish flame, I assumed thats what you wanted it for in the first place.

Cheers, Xenoid
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[*] posted on 25-8-2007 at 17:12


That explains the smell. And yes i am washing my hands often, but thanks for looking out. The amount of precipitant is at least a couple of grams and isn't Ba(NO3)2 the least soluble in water. Anyway i test it out thanks for the help
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[*] posted on 25-8-2007 at 18:07


To test for barium, you could dissolve some of it in water and add SO4--, as Ba(SO4)2 is very insoluble it will ppt out. There are other insoluble sulfates (esp Sr) so this isn't conclusive, however.
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[*] posted on 26-8-2007 at 00:06


IIRC, adding a little H2O2 to the acid will stop quite a bit of the nasty H2S smell, it will cost you a little bit of product though.



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[*] posted on 25-11-2007 at 17:41


Well after my last failed attempt to make barium nitrate from carbonate. I have decided to try again. This time i was successful in producing a crystalline product. After waiting for the crystals to dry i decided to try them out. I simply mixed what i thought was Barium Nitrate with various fuels to see if the mixture would burn green. The first attempts were not successful the mixture failed to take light, but i could tell that an oxidizer was present. Thinking that my product was contaminated I proceeded with washing the crystals. After they had dried i tried the same thing. This time i mixed the red gum with the barium nitrate and i added a little KClO3 to the mix. It took light and had a faint greenish tint at first but as it continued to burn it became more yellow. This leads me to believe my mixture has a chloride contaminate in it, but with the yield i have its not worth trying to purify it.

Anyway while looking up info on barium nitrate i found a reference that said it could be prepared by dissolving barium carbonate in acetic acid then mixing with potassium or sodium nitrate and letting the solution evaporate and over time crystals of barium nitrate would form as the concentration became higher. I wonder if this would yield a cleaner product. If anyone has any insight on making Barium nitrate I will appreciate your help.
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[*] posted on 25-11-2007 at 18:01


If you are after the barium nitrate as a flame colourant, do not use sodium salts in the metathesis. Sodium emission lines are very strong - only a trace, and you will get a yellow flame. Stick with potassium or ammonium salts.



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[*] posted on 25-11-2007 at 18:31


I don't consider nitric acid "hard to get". If you can obtain conc sulfuric acid, KNO3 and all glass distillation setup then you have virtually unlimited nitric acid. Unless of course the small amount of HNO2 is bad for what you are doing. In fact, nitric acid was one of the first things I distilled.

At least it's easier than many other reagents.

Correct me if I'm wrong...

[Edited on 25-11-2007 by MagicJigPipe]




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[*] posted on 26-11-2007 at 19:02


I have always wanted to distill my sulfuric acid to get the dye out of it, but I currently don't have a condenser. Where do you recommend i buy one from. I have seen them sold on some websites going for 40 dollars a piece. IS that a good price for one or should i look for a better deal.
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[*] posted on 26-11-2007 at 21:16


$40 for what?

Depends on what kind you want and what you are distilling. For sulfuric acid I would use one with at least a 200mm jacket. A 200-300mm liebig (or a West if you can find one)(and diagonally to reduce the distance the vapors must travel to the condenser). I think that the fact that H2SO4 condenses at such a high temperature might compensate for the EXTREMELY high temps. It might be better to use a retort because you might get stuck joints at those temps with multiple pieces. Not to mention both joint grease and teflon decompose below h2so4's boiling point.

And this is just a guess but I've heard that even carbon can oxidise to CO2 in boiling conc. sulfuric acid so your dyes (and possibly any other organic material) might decompose to many different byproducts.

P.S. I get a lot of stuff on ebay and I haven't experienced any of the negative things that people speak about here in reference to ebay. I suppose you need to know what to look out for when buying used.

[Edited on 26-11-2007 by MagicJigPipe]

[Edited on 26-11-2007 by MagicJigPipe]




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[*] posted on 27-11-2007 at 01:12
Solubility of Barium Salts


Squall, this is the data from CRC, 62nd edition, 1981 - 1982. This is grams per 100 ml at
the listed temperature in C:


BaCl2 ______ 37.5000 gr. @ 26 ______ 59.0000 gr. @ 100
BaCO3 ______ 0.0022 gr. @ 18 _______ 0.0065 gr. @ 100
Ba(NO3)2 ____ 8.7000 gr. @ _0 ______ 34.2000 gr. @ 100


Hope this helps you out.




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[*] posted on 27-11-2007 at 01:44


Squall,
As Twospoons said, sodium is the problem with getting a yellow flame. Using barium acetate and potassium nitrate might work, the much weaker purple of potassium having less effect on the barium green.

However, there may be another way. Sodium acetate is soluble in alcohol to the extent of a few grams per 100 cc, sodium nitrate is very slightly soluble in alcohol, the barium compounds are almost completely insoluble. If you used the barium acetate and sodium nitrate method, then stirred the barium nitrate crystals with alcohol several times, much of the sodium would be washed away. Washing the original set of crystals, then dissolving them in DW and recrystallising (meaning there is some liquid left) and washing that batch with alcohol should clean it up pretty well.

Treating barium carbonate with nitric does solve most of the contamination problem. If you can't get nitric acid and can't distill your own, then isothermal distillation using a nitrate mixed with sulfuric or phosphoric acid, even pretty impure, over onto barium carbonate will give you a reasonably pure barium nitrate; this will take awhile though, because isothermal distillation isn't fast.
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[*] posted on 27-11-2007 at 19:10


I appreciate all of your advice, but I really need to do some reading on distilling before i attempt the procedure.
But anyway here are the specs of a condenser that United Nuclear sells:

Highly efficient, and perfect for distilling water, making or concentrating acids, extracting oils or perfumes from
plants, etc. Sealed inner tube, large capacity water jacket provides efficient cooling.

Far more resistant to breakage than fragile coil-type condensers.

Top quality heat resistant Borosilicate (Pyrex type) glass construction.

Large end fits a standard #4 rubber stopper, and the upper & lower cooling water inlets fit standard 1/4" ID
plastic or latex tubing.

Overall length is 18" (460mm), cooling section is 12" (300mm), and the jacket diameter is
1-1/8" (30mm).

Sells for well over $100 in most laboratory supply catalogs.

This is what their ad says about the condenser. It seems like a good buy but, then again i have never owned a condenser before so i am just guessing here.
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[*] posted on 28-11-2007 at 02:05
Available Ingredients


Sqall, what is available to you ? I asume you have no choice but to work wth what you
have available OTC.




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[*] posted on 28-11-2007 at 03:19


Squall,
sounds a lot like the one here http://www.sciplus.com/category.cfm/subsection/4/category/42...

Distilling nitric acid takes a bit of care, rubber stoppers can make it tricky.

I use something similar to UI-3830 here http://www.uicoglass.com/distillingah.htm which minimizes the number of joints to deal with.

Isothermal distillation is slow, but gets you away from the need to heat stuff, and to worry about what hot nitric acid is going to do to your apparatus.
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[*] posted on 5-12-2014 at 19:28


I just brought some Barium carbonate from a ceramic and pottery supplier and want to turn it into Barium nitrate.
Instead of a nitrate salt i wish to use 62% Nitric acid because I have a large amount stored.
Is this concentration too strong and needs to be diluted or is it fine to use?
And just to double check, can I dissolve the Barium carbonate in the Nitric acid until neutral to litmus and then boil down with low heat outside until crystals form and then cool the solution to obtain full yield.
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