chloric1
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HCl and BaCl2 solubility
OK processing pottery grade barium carbonate by first converting to the chloride. The first batch had a stubborn yellow color that seems to like to
stick to deposited crystals of BaCl2.2H2O. I added boilling water to the crystals only get an opaque taupe colored liquid. Out of frustration, I
figured that the impurities where HCl soluble and BaCl2 was less so. What surprised me was the what happened as I added the concetrated acid. It
clouded up imediately as if I had precipitated AgCL. I only added 300ml of conc HCl to saturated BaCl2 solution. I am happy it is alot whiter.   NOw I only
need to wait for larger crystals for easy separation.
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12AX7
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Yep that happens, I noticed it when I was making some BaCl2 solution. Add HCl, precipitation. Good old common ion effect, no?
Mine is pure white, well it would be except for a stain of yellow (iron) that I accidentially got in there, oh well. Did you notice sulfides in your
BaCO3? Stinky...
Tim
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chloric1
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Sulfide for sure
OH GOD The sulfides are bad enough that I have to take special precuations for
Hydrogen Sulfide toxicity. I also believe this is common ion effect but somehow this seems a little more pronounced. What really is astounding is
the particle size seems relatively large. What I mean as the precipitate settles within a couple hours and moving or upsetting the container does not
move the sediment much except for the loose crystals. This is opposed to barium carbonate precipitate which moves with fluidic properties even after
4 or 5 hours of boiling and digestion. I believe HCl also affects Ammonium, sodium , and potassium chlorides the same way. I would gather that
strontium would behave simularly as well. BUt I know Magnesium will not precipitate and calcium chloride I have not tried. I think calcium remains
dissolved as well. I need to buy 5 pounds of strontium carbonate to confirm or reasses my theories.
Besides I am going to need sizable amounts of strontium chloride of respectable purity that I can make strontium chromate with.
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woelen
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I made nicely pure BaCl2 as follows:
Dissolve potteries BaCO3 in dilute HCl (10% by weight). Do this outside, there is quite some H2S in the gas mix produced. Use only a slight excess
amount of HCl. The liquid remains cloudy. Heating a little bit helps speed up the dissolving of the BaCO3, especially when almost all acid is used,
the dissolving is slow.
Add a small amount of H2O2. This makes the liquid more cloudy. Remains of H2S are oxidized to S and/or SO4(2-). Both precipitate out, as elemental S
and as BaSO4. In this way, you very effectively remove any sulphur contamination. Any excess of H2O2 is no problem, it will be removed/destroyed in a
further stage.
Leave liquid alone for a day. The liquid slowly becomes clear and white stuff settles at the bottom. I did not get crystals of BaCl2.2H2O at the
bottom, but if that happens, then add a little more water until these are dissolved and let liquid settle again for a day or so.
The clear liquid can be decanted and concentrated by heating. This takes quite some time, but on a hot-plate it can be done quite well. At a certain
point, on cooling down, crystals of BaCl2 are formed. Any H2O2 is decomposed/evaporated by the heat. It may be possible to evaporate to dryness, but
I'm not sure, whether BaCl2.2H2O is stable on heating or looses HCl and becomes a basic compound. I don't think so, but I did not take the risk of
spoiling my sample.
I have a nice white sample of BaCl2.2H2O. It dissolves in distilled water without cloudyness.
Precipitating the BaCl2 with concentrated HCl is a trick, which works, but don't you get a lot of coprecipitated HCl? I can imagine that the product
still is quite acidic, even after drying and that a recrystallization from distilled water is necessary to get rid of trapped HCl.
[Edited on 22-12-05 by woelen]
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chloric1
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woelen you are right about recrystallizing to remove HCl. Since the initial BaCO3 source is so impure I am concerned about sodium as well. Some of
this surely would have coprecipitated with the initial BaCl2 was precipitated. So when I get a chance I will use boiling water and allow to cool.
MOst if not all sodium will stay in solution as long as HCl is not added.
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12AX7
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I've melted BaCl2, I didn't notice any HCl fumes so it can't be too much of a problem. Brauer says all the alkaline earths can be prepared in the
same manner, which doesn't involve dehydrating the hydrate salt.
BTW how would you seperate Ba, Sr and Ca? They all precipitate SO4(2-), so that's no good except to remove slightly soluble CaSO4 with boiling water.
How sensitive is SrCl2 to common ion effect?
Tim (haven't prepared SrCl2 yet, will sooner or later)
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chloric1
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Mysterious Canary yellow color
Woelin,
Tried a little H2O2 to the more contaminated and the brown and gray stuff was oxidized to a pretty canary yellow material. Expecting bright yellow
crystals, I got a clear yellow filtrate and white BaCl2 instead. Even more intresting is that the yellow is somewhat soluble in Denatured alcohol and
HCL. I may try more drastic measures to figure out what the yellow stuff is. It looks very much like chromate but will not react with BaCl2. It
might be FeCl3 so I will heat it with very concentrated NaOH. If it is not FeCl2 then I will burn it and note any color changes or carbonizing etc.
If it is organic, I will mix solution with activated charcoal and see if it goes clear.
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woelen
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That sounds interesting. The BaCO3 I used in my experiments resulted in formation of a very pale yellow precipitate. This can easily be explained as a
mix of BaSO4 and S, formed by oxidation of the sulfide.
What color has your crude BaCO3? Mine is white, not brown or grey. Did you use excess H2O2. If not, then you may have formed some polysulfide. A
polysulfide solution is deep yellow. However, it is unstable in acidic solutions, giving sulphur and H2S.
If the BaCO3 you have is brown, then an iron contamination may be a possible source of the yellow color, but I doubt that BaCO3 contains iron
impurities, because that would result in strong coloration in ceramics and that is precisely the thing where BaCO3 is used. Any contamination,
resulting in strong colors (e.g. iron, manganese, chromium, nickel, cobalt) are very unlikely to my opinion, or the BaCO3 would be quite crappy for
its intended purpose.
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chloric1
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Well yes you have a point about the heavy metals. I have concluded it is something organic. It is stable towards very acidic conditions and can only
be removed from the barium chloride by repeated recrystallizations and washings. This is where I am at now. My barium chloride was in three distinct
batches. Each with increasing levels of contamination. The first batch was dried on the steam bath and was offwhite. A solution was made in boiling
distilled water and about 10% of the material precipitated. I could have redissolved this and processed it for a few more days but I want to finish
so I may attend to other projects so I threw this batch out. Insidentally this batch came from processing the "tails" so nothing was really expected
anyways. I pulled the first portion of the second batch off the steam bath and it was more crystaline and actually sparkles and shimmers in white
light. A lot better My third batch is still largely in solution and the
crystals which are present are almost transparent! The solution is water
clear. The third batch will be the best. The second batch may be converted into barium nitrate or barium hydroxide.
Oh Woelen, my carbonate is alomst as white as flour but in my halogen light there is a distinctive bluish-grey tint. And i am a good boy as I am
washing all my glassware in magnesium sulfate solution. I have all my washing in a 1500 ml beaker. When the BaSO4 matures, I wish to collect and
save it for later use and discard the magnesium salts.
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