Jmap science
Yet another ID of SimpleChemist-238
Posts: 40
Registered: 19-10-2013
Location: Florida
Member Is Offline
Mood: No Mood
|
|
Making barium nitrate from barium carbonate without nitric acid
I do not have any nitric acid but I do have potassium and strontium nitrates. How do I make barium nitrate from barium carbonate. I was wondering
if there was a reaction to make barium nitrate without nitric acid or nitrogen oxides.
I will use the barium nitrate to make nitric acid by addition of H2SO4. The flask would be chilled and filtered to make conc. nitric acid.
I have been working on a way to make conc. nitric acid from a conc. acid and potassium nitrate without distillation. I got the idea from
chemistry hideout. I want to improve on it.
|
|
|
TheChemiKid
Hazard to Others
Posts: 493
Registered: 5-8-2013
Location: ̿̿ ̿̿ ̿'̿'̵͇̿̿з=༼ ▀̿̿Ĺ̯̿̿▀̿ ̿ ༽
Member Is Offline
Mood: No Mood
|
|
This should be posted in beginnings, just search on google, and with a tiny bit of you can figure this out.
Add hydrochloric acid to make barium chloride, then add to a hot solution of sodium nitrate and let cool. This was stated on wikipedia.
This other method may not work, but you could try adding copper nitrate to the barium carbonate.
When the police come
\( * O * )/ ̿̿ ̿̿ ̿'̿'̵͇̿̿з=༼ ▀̿̿Ĺ̯̿̿▀̿ ̿ ༽
|
|
|
Brain&Force
Hazard to Lanthanides
Posts: 1302
Registered: 13-11-2013
Location: UW-Madison
Member Is Offline
Mood: Incommensurately modulated
|
|
Won't work. Barium carbonate is insoluble.
A similar synthesis and distillation has been done before with calcium nitrate and sulfuric acid. Don't waste your valuable barium.
At the end of the day, simulating atoms doesn't beat working with the real things...
|
|
|
bfesser
|
Thread Moved
22-1-2014 at 15:28 |
phlogiston
International Hazard
Posts: 1375
Registered: 26-4-2008
Location: Neon Thorium Erbium Lanthanum Neodymium Sulphur
Member Is Offline
Mood: pyrophoric
|
|
Dissolve in molten ammonium nitrate. Research well before attempting this, because melting ammonium nitrate can be dangerous.
-----
"If a rocket goes up, who cares where it comes down, that's not my concern said Wernher von Braun" - Tom Lehrer |
|
|
bismuthate
National Hazard
Posts: 803
Registered: 28-9-2013
Location: the island of stability
Member Is Offline
Mood: self reacting
|
|
How about NO2 and barium oxide from the decomposition of barium carbonate.
|
|
|
Jmap science
Yet another ID of SimpleChemist-238
Posts: 40
Registered: 19-10-2013
Location: Florida
Member Is Offline
Mood: No Mood
|
|
Would barium carbonate and potassium nitrate work?
|
|
|
TheChemiKid
Hazard to Others
Posts: 493
Registered: 5-8-2013
Location: ̿̿ ̿̿ ̿'̿'̵͇̿̿з=༼ ▀̿̿Ĺ̯̿̿▀̿ ̿ ༽
Member Is Offline
Mood: No Mood
|
|
No. Barium Carbonate is insoluble in water and basically anything else.
Go for the calcium nitrate instead of the barium nitrate, but if you really want to use the barium, then try either the NO2 or the
HCl/NaNO3 method.
When the police come
\( * O * )/ ̿̿ ̿̿ ̿'̿'̵͇̿̿з=༼ ▀̿̿Ĺ̯̿̿▀̿ ̿ ༽
|
|
|
HeYBrO
Hazard to Others
Posts: 289
Registered: 6-12-2013
Location: 'straya
Member Is Offline
Mood: 
|
|
if you have strontium and potassium nitrate, why don't you buy some? seems like the most logical solution, especially because barium is not the
friendliest ion to work with when soluble.
|
|
|
Jmap science
Yet another ID of SimpleChemist-238
Posts: 40
Registered: 19-10-2013
Location: Florida
Member Is Offline
Mood: No Mood
|
|
well I see your point
|
|
|
woelen
Super Administrator
Posts: 7976
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
If you have hydrochloric acid, then you could try making a solution of BaCl2 from BaCO3. Making very pure BaCl2 from the pottery grade BaCO3 is not
trivial, but if you intend to use your product for pyro-purposes, then you need to be less stringent on purity. You also do not need to isolate the
BaCl2, you can keep it in solution.
Here is my webpage on making BaCl2.2H2O:
http://woelen.homescience.net/science/chem/exps/BaCl2_2H2O/i...
To the solution of BaCl2 you add a solution of KNO3. Ba(NO3)2 is much less soluble than KNO3, KCl and BaCl2, so that should drop out. From a single
step you get mostly Ba(NO3)2 with a small amount of K-ions and Cl-ions in it, but that is not a problem for pyro-purposes. Otherwise, you need to
recrystallize your Ba(NO3)2 from quite a large amount of water.
This also works with NaNO3 instead of KNO3, but the resulting Ba(NO3)2 will contain some Na-ions in that case and if you want to use it for
pyro-purposes, that is not good at all.
[Edited on 23-1-14 by woelen]
|
|
|
Jmap science
Yet another ID of SimpleChemist-238
Posts: 40
Registered: 19-10-2013
Location: Florida
Member Is Offline
Mood: No Mood
|
|
Thanks
|
|
|
markx
National Hazard
Posts: 645
Registered: 7-8-2003
Location: Northern kingdom
Member Is Offline
Mood: Very Jolly
|
|
Well, the method that I used, was to add a slight stoichiometric excess of BaCO3 to ammonium nitrate solution and then boil, boil, boil until no more
ammonia fumes evolved. Ammonium carbonate will decompose in boiling water and will therefore be expelled from the solution. What you are left with is
Ba(NO3)2 solution. The boiling will take a LONG time though....as I remember I kept the kettle simmering for 2-3 days. But in the end you are left
with a nice crop of crystals. Filter, purify and dry them as needed. Beware though, that water soluble barium salts are very poisonous. Take the
neccesary precautions when handling them.
[Edited on 23-1-2014 by markx]
Exact science is a figment of imagination.......
|
|
|
TheChemiKid
Hazard to Others
Posts: 493
Registered: 5-8-2013
Location: ̿̿ ̿̿ ̿'̿'̵͇̿̿з=༼ ▀̿̿Ĺ̯̿̿▀̿ ̿ ༽
Member Is Offline
Mood: No Mood
|
|
Why did you take away the picture?
When the police come
\( * O * )/ ̿̿ ̿̿ ̿'̿'̵͇̿̿з=༼ ▀̿̿Ĺ̯̿̿▀̿ ̿ ༽
|
|
|
Brain&Force
Hazard to Lanthanides
Posts: 1302
Registered: 13-11-2013
Location: UW-Madison
Member Is Offline
Mood: Incommensurately modulated
|
|
Woelen's procedure sounds good to me, but if you take the KCl contaminated crystals and dissolve them in sulfuric acid, wouldn't you be making aqua
regia with the nitrate? Would distillation cause the nitrosyl chloride to escape and decompose?
At the end of the day, simulating atoms doesn't beat working with the real things...
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Brain&Force:
Good thinking. Always expect interactions!
However, I suspect the only way to form some NOCl here would be by the action of conc H2SO4 on the dry chloride and nitrate salts.
|
|
|
Zyklon-A
International Hazard
Posts: 1547
Registered: 26-11-2013
Member Is Offline
Mood: Fluorine radical
|
|
This is rather off topic, but I want to make nitric acid from Barium nitrate, (Jmap science gave me the idea). I have 98% sulfuric acid, and pyro
grade Ba(NO3), I mixed them in stoichiometric proportions, and waited.
The problem was that even though Ba(NO3) is insoluble, there was just way to much solid in proportion to the liquid. It just looked like
there wasn't any liquid at all. The HNO3 was fuming out from the heat generated, what should I do. I'm not asking to be spoon fed here,
tried a few different ideas, and have worked on this since the day this topic was started, but nothing has worked.
|
|
|
Metacelsus
International Hazard
Posts: 2531
Registered: 26-12-2012
Location: Boston, MA
Member Is Offline
Mood: Double, double, toil and trouble
|
|
Dissolve the barium nitrate in water first. Or, you could distill the nitric acid (however, you might as well do this with any other nitrate.)
|
|
|
Zyklon-A
International Hazard
Posts: 1547
Registered: 26-11-2013
Member Is Offline
Mood: Fluorine radical
|
|
The whole reason that I want to use barium is because I don't have a distillation setup. I know I'll have to get one at some point, but I can't afford
to drop the required $100- $300, at least not now. Nitric acid >95% is not required, but I'd like it to be as con. as possible.
|
|
|
servo
Harmless
Posts: 16
Registered: 15-12-2018
Member Is Offline
|
|
Quote: Originally posted by woelen  |
To the solution of BaCl2 you add a solution of KNO3. Ba(NO3)2 is much less soluble than KNO3, KCl and BaCl2, so that should drop out. From a single
step you get mostly Ba(NO3)2 with a small amount of K-ions and Cl-ions in it, but that is not a problem for pyro-purposes. Otherwise, you need to
recrystallize your Ba(NO3)2 from quite a large amount of water.
[Edited on 23-1-14 by woelen] |
Can you use Ammonium Nitrate instead?
what I did is added an excess of ammonium nitrate solution (hot) to a hot saturated solution of Barium Chloride
what i noticed is an immediate crystalline sort of precipitation and I kept boiling the solution for about 2 hours during which more and more ppt
were formed finally I took it off the heat let it cool down filter the crystals
Now my question is are the crystals really Barium Nitrate? or its Barium Chloride that did not react with ammonium nitrate , I took excess of amonium
nitrate so that if they both react All of Barium chloride reacts
I know topic is old but chemistry never gets old
[Edited on 27-11-2021 by servo]
[Edited on 27-11-2021 by servo]
|
|
|
Fulmen
International Hazard
Posts: 1693
Registered: 24-9-2005
Member Is Offline
Mood: Bored
|
|
@Zyklon: Just wash it out with water and recover the dilute acid. I think it's the best you can hope for, tried the same with calcium nitrate and it
was literally a smoking mess. The stiff paste was unworkable and too corrosive to handle.
We're not banging rocks together here. We know how to put a man back together.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
