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Eli25
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Separating MgCl2 from a solution of MgCl2 and CaCl2
I have a solution containing MgCl2 and CaCl2 with water. I need to separate MgCl2.How can I do it in the best and easiest way?
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annaandherdad
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Apart from fractional crystallization, you could add sodium sulfate to precipitate the calcium, since magnesium sulfate is much more soluble in water
than calcium sulfate. I'm not sure this will work, however, since the solubility of calcium sulfate is affected by the presence of the chlorides. In
any case, calcium sulfate is still slightly soluble in water, and even with the calcium gone, you'd end up with a mixture of sodium chloride and
magnesium sulfate.
Alternatively, you could precipitate both calcium and magnesium as carbonates, then dissolve the (mostly magnesium carbonate) in sulfuric acid to get
magnesium sulfate. This could be later converted to the chloride, by going through the carbonate again, if necessary. Both magnesium and calcium
carbonates are quite insoluble in water.
I've been working on separating sea water into its ingredients, just for fun. At a certain point you get a mixture of magnesium and calcium chlorides
(mostly). I'm using fractional crystallization. There is a very interesting thread here on SM on separating sodium and potassium chlorides, you
should check it out.
Any other SF Bay chemists?
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DistractionGrating
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Wouldn't another option be to add NaOH or KOH to precipitate Mg(OH)2, and then add HCl to the Mg(OH)2?
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gboneu
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That wouldn't work because both Mn(OH)2 and Ca(OH)2 are insoluble in water, and will both react with the acid to reform the chlorides.
The method i have in my book is the next.
If you have Calcium contamination un your Manganese chloride, precipitate all the manganese with Hydrogen sulfide (SH2) until no more comes out, then
filter off and wash with water to wash the soluble contaminants and acids. Then dissolve all the manganese sulfide in hydrochloric acid and heat up
the solution to get out all the SH2 remaining in solution.
Obviously a recristalization is recomended 
If you want to see if there are still some calcium ions you can test them with these reagents:
Get a small drop and touch a filter papaer, leave it to dry and test...
1)SODIUM RHODIZONATE (Gray-blue colour with Ca)
_Sodium Rhodizonate (1% Aqueous solution ) Spray 1
_ Ammonia (25% aqueous solution) Spray 2
2)ALIZARINE (Red-Violett colour with Ca)
_ Alizarine in Chloroform (2% solution) Spray 1
_2M sodium hydroxide solution. Spray 2_ Ammonia solution (25%) Spray 3
Note. If you dont have chloroform usea Alcoholic alizarine solution (saturated)
3) QUINALIZARIN (Violet colour with Ca)
_Quinalizarin solution (0.05%) in Ethanol (70% by volume) Spray 1
_Ammonia solution (25% aqueous solution) Spray 2
NOTE: If you have negative resoults try putting the papers in a ammonia rich atmosphere, the original analisis is using ammonia gas, bur the solution
has to work alsow, but it's always better to know 
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DistractionGrating
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Gboneu, I think you are confusing Manganese (Mn) with Magnesium (Mg).
You have a point, though, that Ca(OH)2 is not very soluble in water, but significantly more so than Mg(OH)2. I guess my idea's feasibility depends on
how much Ca is in the solution in question. I got my idea from the Hach Calcium Hardness EDTA Titration method, where KOH is used to precipitate the
Magnesium, leaving Calcium in solution.
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Chemosynthesis
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Coulometry.
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gboneu
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Owww.. (-.-)
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gboneu
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Ok, i found some another info in my book's, Magnesium carbonate is soluble in a solution of ammonium chloride and ammonia, where calcium strontium and
barium are not soluble (There carbonates) so you can extract the magnesium with this solution, Then maybe ammonium carbonate can be decomposed by
hydrochloric acid i think ... Sorry for the confusion, im was exited when i found that info xD
[Edited on 12-5-2015 by gboneu]
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blogfast25
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Quote: Originally posted by gboneu  | Ok, i found some another info in my book's, Magnesium carbonate is soluble in a solution of ammonium chloride and ammonia, where calcium strontium and
barium are not soluble (There carbonates) so you can extract the magnesium with this solution, Then maybe ammonium carbonate can be decomposed by
hydrochloric acid i think ... Sorry for the confusion, im was exited when i found that info xD
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I'm not convinced that this information is completely correct.
Magnesium hydroxide is indeed soluble in ammonium salt solutions due to:
Mg(OH)2 + 2 NH<sub>4</sub><sup>+</sup> === > Mg<sup>2+</sup> + 2 NH3 + 2 H2O
Magnesium carbonate, acc. Holleman's 'Inorganic Chemistry', is a basic carbonate: Mg(OH)2.3MgCO3.3H2O. Whether the ammonium salt trick also works on
that basic carbonate will depend on its solubility product K<sub>s</sub>.
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gboneu
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Here says that it work's with the carbonate.
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blogfast25
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Forgot to provide the link, maybe? I'm quite interested in seeing it. Ta.
[Edited on 13-5-2015 by blogfast25]
Turns out there are several Mg carbonates (Wiki), including a 'straight' MgCO3, with a K<sub>s</sub> of
3.5×10<sup>–8</sup>. That's probably soluble enough to soluble in ammonium salts...
[Edited on 13-5-2015 by blogfast25]
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gboneu
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The book where i got the info is the (I live in argentina it's in spanish)
https://books.google.com.ar/books?id=LoTwmQEACAAJ&dq=Pra...
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gboneu
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I search some info on using Chromatography but it's a little complicated to do, And it takes a lot of time, However if i have time i give it a try on
my lab to se what can i get.
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Eli25
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Hi guys. Thanks for your answers. Actually I want MgCl2 as the final product ...I just need to remove Ca from solution.....pleaseeeeee help me... I
need to keep my job by giving the best solution...
What if I add MgSo4 and remove CaSo4 as precipitate ? Can I have pure MgCl2 solution?
I'm working in a factory so the answer should be practical and economical....
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blogfast25
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| Quote: Originally posted by Eli25 | Hi guys. Thanks for your answers. Actually I want MgCl2 as the final product ...I just need to remove Ca from solution.....pleaseeeeee help me... I
need to keep my job by giving the best solution...
What if I add MgSo4 and remove CaSo4 as precipitate ? Can I have pure MgCl2 solution?
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Yes but CaSO4 is better described as 'poorly soluble' than 'insoluble'. Its solubility (acc. Wiki) is about 2 g/litre, so some residual CaSO4 will
stay in solution. It depends whether this relatively small contamination is acceptable to you or not.
You'd also need to know the concentration of CaCl2 quite well, as otherwise you may leave more Ca in solution or contaminate the MgCl2 with excess
MgSO4.
[Edited on 15-5-2015 by blogfast25]
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Chemosynthesis
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I am not sure how practical this is for you, but you may wish to consider trying a calcium specific ionophore such as ETH 129, ETH 5234, or ETH 1001.
These work in coulometric separations, and given the lipophilicity, could probably be utilized in a solvent extraction with something you have on
hand, though the latter may not be an option for you due to purity standards and the addition of an organic process stream.
Edit: and, probably sounding extremely callous here, but while I am sorry your job is at risk, that isn't really our problem or what the forum is for.
I wish you all the best, but if your job retention comes down to what strangers on an online website comprised of those who love amateur chemistry and
science suggest, you are in an extremely un-enviable position, which may not be due to your own fault, and I would highly recommend examining how to
improve your situation.
I do mean this as an attempt to be helpful and not just make light of your situation.
[Edited on 15-5-2015 by Chemosynthesis]
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blogfast25
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Ionophores? Blimey, occasionally one does learn something at SM!
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DistractionGrating
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Checking solubility tables, I see that Calcium molybdate is only sparingly soluble (0.004099 g / 100 mL), while Magnesium molybdate is much more
soluble (13.7 g / 100 mL). Now, I know nothing about working with molybdates, but would it be possible to add magnesium molybdate in order to
precipitate out calcium molybdate? If so, then of course, it would be helpful to know how much calcium is present in order to know how much magnesium
molybdate to add.
[Edited on 15-5-2015 by DistractionGrating]
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blogfast25
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| Quote: Originally posted by DistractionGrating | Checking solubility tables, I see that Calcium molybdate is only sparingly soluble (0.004099 g / 100 mL), while Magnesium molybdate is much more
soluble (13.7 g / 100 mL). Now, I know nothing about working with molybdates, but would it be possible to add magnesium molybdate in order to
precipitate out calcium molybdate? If so, then of course, it would be helpful to know how much calcium is present in order to know how much magnesium
molybdate to add.
[Edited on 15-5-2015 by DistractionGrating] |
The cost of the molybdate would not justify trying to 'save' that MgCl2. I'd bet good money on that being the case. 
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MrHomeScientist
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It always amuses me when people specify they want the "best and easiest" way to do something. Really? I never would have guessed.
Unfortunately I only know the worst and hardest ways. 
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DistractionGrating
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Fair enough, but does my idea have merit? Just trying to learn.
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blogfast25
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Assuming your solubility data are correct, then yes.
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Eli25
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Hiiiiii.... It was so nice seeing your comments...Actually I'm here to learn more and as we say in our country " every mind has a new thought", I'm
looking forward to gain more different ideas and suggestions....I'm not here to find the exact answer for those who said it's not our problem.....
Believe me your comments help me consider all different ways to solve my problem and I'm learning alot. I do appreciate your help.
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Eli25
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by the way where can I find the rightest values of solubility and Ksp?
I want to know if CaSo4 is more insoluble or CaCo3? to know weather adding MgCo3 works better or adding MgSo4 to my solution to gain purer MgCl2!!!!
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DistractionGrating
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In pure water, absent any other ions, CaSO4 is more soluble than CaCO3, but solubility is affected by the presence of other ions. For example, CaCO3
is more soluble in seawater than it is in freshwater. Then, there is the "common ion effect" that can make some substances much less soluble when a
common cation or anion is present. I'm still trying to figure this all out myself.
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