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elementcollector1
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So, on a scientific note, what happens if you add a chloride salt to max concentration (e.g. 37%) acid? Does it behave like even more concentrated
than 37%?
Also, does this process work for sulfates?
Elements Collected:52/87
Latest Acquired: Cl
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blogfast25
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Quote: Originally posted by elementcollector1  | So, on a scientific note, what happens if you add a chloride salt to max concentration (e.g. 37%) acid? Does it behave like even more concentrated
than 37%?
Also, does this process work for sulfates? |
As it happens, quite a few chlorides are insoluble in conc. HCl. AlCl3 (as a hydrate) precipitates on gassing a strong solution with HCl (to
saturation). Zirconyl chloride behaves also like that.
Also the formation of chloride complexes must be considered in some cases: copper, lead and silver form anions like CuCl4(2-) (tetrachlorocuprate
anions).
Re. sulphates, systems like an acid + salt of the same acid + water should really be considered three phase systems. The solubilities of the two
solutes can be surprising. BeSO4 is highy soluble in water (up to 44 % as BeSO4) at 20 C but it is almost insoluble in concentrated H2SO4.
[Edited on 22-2-2013 by blogfast25]
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DraconicAcid
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| Quote: Originally posted by blogfast25 | | Quote: Originally posted by elementcollector1 | So, on a scientific note, what happens if you add a chloride salt to max concentration (e.g. 37%) acid? Does it behave like even more concentrated
than 37%?
Also, does this process work for sulfates? |
As it happens, quite a few chlorides are insoluble in conc. HCl. AlCl3 (as a hydrate) precipitates on gassing a strong solution with HCl (to
saturation). Zirconyl chloride behaves also like that.
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So does sodium chloride. If you add a few drops of conc. HCl(aq) to a solution of saturated sodium chloride, you get salt precipitating out.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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AJKOER
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| Quote: Originally posted by elementcollector1 | So, on a scientific note, what happens if you add a chloride salt to max concentration (e.g. 37%) acid? Does it behave like even more concentrated
than 37%?
Also, does this process work for sulfates? |
Not sure to be honest, notice that a lot of this work is based on simulations and relatively recent. I would, however, not be surprised of an acid
could have an 'activity' level that is sustained in a reaction that would otherwise not be observed with an acid starting at its max concentration.
This, I would guess, could occur as the 'activity' level of H2O, normally around 1, could be at a much lower value (apparently somewhere between 0.3
and 0.4, for example, for very conc brine solutions which is far from unity).
-----------------------------------------------
To answer your second question, maybe, check out thiosulphate leaching. To quote " Ammoniacal thiosulphate solution allows the solubilization of gold
as stable anionic complex. Leaching of gold occurs at appreciable dissolution rates." See http://www.sciencedirect.com/science/article/pii/0304386X950... .
[Edited on 22-2-2013 by AJKOER]
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AJKOER
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For those interested in MgCl2 in cases where CaCl2 presents solubility issues, can prepare MgCl2 as follows:
MgSO4 + CaCl2 --> CaSO4 (s) + MgCl2
Don't have CaCl2, make Magnesium hypochlorite (from NaOCl + NaCl + MgSO4 actually forms the dibasic form, Mg(OCl)2.2Mg(OH)2 ) and then add H2O2, or
heat to disproportionate into MgCl2 and MgClO3.
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experimenter
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Why is the hydrochloric acid green ?
I have tried to make hydrochloric acid by distilling table salt and concentrated sulfuric acid (battery acid). I had few mililiters of deionised water
in the receiver to catch possible HCl fumes.
The experiment was a success, but the acid produced has a green color(?). Please, do you know if this is normal? I think it should be transparent.
Sometimes commercial hydrochloric acid is green due to iron impurities, but
in my case, distillation should have left back possible iron salts.
Another possibility is that the hydrochloric acid is contaminated with Cl2 gas. How can I find out if this is the case? How can I remove it? Thanks.
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Zyklon-A
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Perhaps the chloride ion was oxidized by an impurity in your battery acid (or salt).
If your sulfuric acid was pure, chlorine would not form.
I think one way to clear the color, is to put it in sunlight for a few hours.
AJKOER, I'm sure could explain how this works better than I could.
[Edited on 19-8-2014 by Zyklon-A]
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experimenter
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Thanks for your response. In an attempt to find the "culprit", I tried two experiments.
- I distilled commercial, low grade hydrochloric acid mixed with table salt. The green color appeared again in the distillate, especially during the
end of the distillation.
- Distilled the commercial, low grade hydrochloric acid by itself in order to check for impurities in it. No green color appeared in the distillate.
It left behind a very small amount of whitish salt only.
It seems that the impurity came from the salt. Some chemical that can volatilize with HCl?
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Texium
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| Quote: Originally posted by experimenter | Thanks for your response. In an attempt to find the "culprit", I tried two experiments.
- I distilled commercial, low grade hydrochloric acid mixed with table salt. The green color appeared again in the distillate, especially during the
end of the distillation.
- Distilled the commercial, low grade hydrochloric acid by itself in order to check for impurities in it. No green color appeared in the distillate.
It left behind a very small amount of whitish salt only.
It seems that the impurity came from the salt. Some chemical that can volatilize with HCl? |
If you think it
might be the salt, try to find some picking salt. It doesn't have additives like iodine compounds and free-flowing agents in it. Probably still
wouldn't be as good as real reagent grade NaCl, but it's a lot better than normal salt.
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