Sulaiman
International Hazard
Posts: 3561
Registered: 8-2-2015
Location: 3rd rock from the sun
Member Is Online
|
|
Barium chloride or nitrate more usefull ?
To test for sulfates I can buy barium chloride or barium nitrate,
Which has the more general use as an indictator and as a general reagent ?
|
|
|
Detonationology
Hazard to Others
Posts: 362
Registered: 5-5-2015
Location: Deep South
Member Is Offline
Mood: Electrophillic
|
|
Barium nitrate can be used to make some really beautiful green compositions in pyrotechnics.
“There are no differences but differences of degree between different degrees of difference and no difference.” ― William James
|
|
|
j_sum1
Administrator
Posts: 6230
Registered: 4-10-2014
Location: Unmoved
Member Is Offline
Mood: Organised
|
|
Why not get barium hydroxide and make the chloride or nitrate as needed?
Calcium chloride is an acceptable substitute for testing sulfates.
[Edited on 9-11-2015 by j_sum1]
|
|
|
Upsilon
Hazard to Others
Posts: 392
Registered: 6-10-2013
Member Is Offline
Mood: No Mood
|
|
Barium carbonate is cheaper, just as an FYI.
|
|
|
woelen
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
The commercially available barium carbonate (from potteries) is very impure. It contains a lot of BaSO4 and BaS as well. Making pure and clear
solutions from this is hard. I once used this to make clean and pure BaCl2.2H2O, but this was not really easy. It was much more involved than simply
dissolving some of the material in dilute HCl.
I would go for the Ba(NO3)2. It is sufficiently soluble for aqueous chemistry tests and experiments, and it also is quite interesting for colorful
pyrotechnics experiments.
[Edited on 9-11-15 by woelen]
|
|
|
Upsilon
Hazard to Others
Posts: 392
Registered: 6-10-2013
Member Is Offline
Mood: No Mood
|
|
What makes it hard to deal with those impurities? In theory, on addition of HCl the BaSO4 would not be attacked and the BaS would form H2S and BaCl2.
Then you would just have to filter off the BaSO4.
|
|
|
woelen
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Theory and practice differ a lot in many cases. The sulfide is partially oxidized by oxygen from the air, leading to formation of colloidal sulphur,
which remains suspended in the liquid and which is VERY hard to filter. Evaporating the liquid to dryness hence yields a very impure product, which is
yellow and has a dirty smell of sulphurous compounds.
http://woelen.homescience.net/science/chem/exps/BaCl2_2H2O/i...
The procedure is not really difficult, but it also is not as easy as dissolving some BaCO3 in dilute HCl. It is a long and tedious process.
|
|
|
MolecularWorld
Hazard to Others
Posts: 110
Registered: 30-10-2015
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by Upsilon  | | What makes it hard to deal with those impurities? In theory, on addition of HCl the BaSO4 would not be attacked and the BaS would form H2S and BaCl2.
Then you would just have to filter off the BaSO4. |
Suppose you wanted to react 100g of barium carbonate with an 8% barium sulfide impurity (not atypical for pottery grade), you would generate 1.6g, or
1.05 liters, of hydrogen sulfide. That is enough to stink up about two million cubic meters of air (@0.47ppb), and enough to render 105 cubic meters of air (a 7x5x3 meter room) unsafe (@10ppm).
Edit: Slight correction to calculations (probably still wrong). The point is, even a little hydrogen sulfide can be a big problem if you're not
expecting it or don't know how to deal with it.
[Edited on 9-11-2015 by MolecularWorld]
|
|
|
careysub
International Hazard
Posts: 1339
Registered: 4-8-2014
Location: Coastal Sage Scrub Biome
Member Is Offline
Mood: Lowest quantum state
|
|
| Quote: Originally posted by MolecularWorld | | Quote: Originally posted by Upsilon | | What makes it hard to deal with those impurities? In theory, on addition of HCl the BaSO4 would not be attacked and the BaS would form H2S and BaCl2.
Then you would just have to filter off the BaSO4. |
Suppose you wanted to react 100g of barium carbonate with an 8% barium sulfide impurity (not atypical for pottery grade), you would generate 1.6g, or
1.05 liters, of hydrogen sulfide. That is enough to stink up about two million cubic meters of air (@0.47ppb), and enough to render 105 cubic meters of air (a 7x5x3 meter room) unsafe (@10ppm).
Edit: Slight correction to calculations (probably still wrong). The point is, even a little hydrogen sulfide can be a big problem if your not
expecting it or don't know how to deal with it.
[Edited on 9-11-2015 by MolecularWorld] |
<p>On the strength of this argument, I can say that the barium carbonate currently being sold by Seattle Pottery is pretty pure.</p>
<p>Treating 1000 grams of it with pool grade HCl (patio) produced a distinct hydrogen sulfide smell but hardly anything I would describe as a
"stench" and not noticeable beyond several feet away. I would not want to do that indoors, the smell would linger, but I was expecting (from reports)
a far greater H2S release.
|
|
|
careysub
International Hazard
Posts: 1339
Registered: 4-8-2014
Location: Coastal Sage Scrub Biome
Member Is Offline
Mood: Lowest quantum state
|
|
Indeed, it is much, much cheaper. The cheapest barium hydroxide I can find is more than 100 times more expensive that the cheapest barium carbonate.
|
|
|
Sulaiman
International Hazard
Posts: 3561
Registered: 8-2-2015
Location: 3rd rock from the sun
Member Is Online
|
|
I am as far as possible trying to not generate toxic waste
so I like j_sum1's suggestion of calcium chloride substitution (thanks)
but
I just ordered 100g of BaCl2.2H2O claimed >99.8% purity
as solubility of earth alkaline sulphates decreases with increasing atomic number,
so more sensitive than Calcium salts,
and barium chloride is relatively cheap.
I almost bought barium and strontium nitrates as they would be nice for pyrotechnics, maybe another time.
[Edited on 10-11-2015 by Sulaiman]
|
|
|
gdflp
Super Moderator
Posts: 1320
Registered: 14-2-2014
Location: NY, USA
Member Is Offline
Mood: Staring at code
|
|
Be aware that calcium salts are not as sensitive as barium salts. ~2g CaSO<sub>4</sub> will dissolve in one liter of water as compared
with ~2mg for BaSO<sub>4</sub>. Strontium salts may be a good compromise, as they are non toxic and the solubility of
SrSO<sub>4</sub> is ~100mg per liter.
|
|
|
j_sum1
Administrator
Posts: 6230
Registered: 4-10-2014
Location: Unmoved
Member Is Offline
Mood: Organised
|
|
| Quote: Originally posted by gdflp | | Be aware that calcium salts are not as sensitive as barium salts. ~2g CaSO<sub>4</sub> will dissolve in one liter of water as compared
with ~2mg for BaSO<sub>4</sub>. Strontium salts may be a good compromise, as they are non toxic and the solubility of
SrSO<sub>4</sub> is ~100mg per liter. |
Agreed. It really depends on how sensitive a test you need for your application.
If you test with CaCl2 first, you might have a need for only a very small amount of Ba(NO3)2. That answers some of your toxicity concerns as well as
some of your acquisition concerns. Synthesising barium nitrate from pottery grade carbonate might be a PITA but if all you need is a few grams, it
should be reasonably do-able.
|
|
|
gdflp
Super Moderator
Posts: 1320
Registered: 14-2-2014
Location: NY, USA
Member Is Offline
Mood: Staring at code
|
|
Synthesizing barium nitrate from pottery grade barium carbonate is actually much easier than synthesizing barium chloride because nitric acid is an
oxidizing acid. The nitric acid oxidizes barium sulfide either to barium sulfate or sulfur dioxide which removes the issue of colloidal sulfur
contamination. I have found that it is necessary to boil the solution for a few minutes with a decent excess of nitric acid to fully oxidize the
sulfur in some instances.
|
|
|
careysub
International Hazard
Posts: 1339
Registered: 4-8-2014
Location: Coastal Sage Scrub Biome
Member Is Offline
Mood: Lowest quantum state
|
|
Hot toluene dissolves sulfur, according to this:
http://media.rsc.org/Classic%20Chem%20Demos/CCD-73.pdf
it is about 10 g per 100 mL at 50 C.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
