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nscheffield
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[*] posted on 18-4-2016 at 08:51
separation of carbonates


I bought powdered garden lime for my lab and was wondering if it is possible to purify it. It states it is about 55% calcium carbonate and 45% magnesium carbonate. my question is how do i separate the two carbonates?



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Marvin
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[*] posted on 18-4-2016 at 09:06


Start with your own ideas.
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nscheffield
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[*] posted on 18-4-2016 at 09:19


i considered converting the carbonates to bicarbonates, because one is more soluble in water than the other then, however it seems like a long process. thought about converting the carbonates to other salts but in the end both are very similar.



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[*] posted on 18-4-2016 at 09:36


Why do you want them separated?



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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[*] posted on 18-4-2016 at 09:41


Good Luck. Just consulted my Merck Index--both virtually insoluble in water, soluble in dilute acids.



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m1tanker78
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[*] posted on 18-4-2016 at 09:48


The sulfates have a wide gap in solubility.



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[*] posted on 18-4-2016 at 09:59


^^excellent point m1



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nscheffield
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[*] posted on 19-4-2016 at 06:35


If converted to sulfate how easily could they be converted to oxides or hydroxides? i like keeping my chemicals in the most easily convertable forms. (such as having lots sodium oxide/hydroxide and a variety of acids (H2SO4, HNO3, HCL) and just make what i need for a specific reaction and not store lots of "premade" compounds)




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[*] posted on 19-4-2016 at 07:03


I keep magnesium sulfate and calcium chloride on hand for use as drying agents as well as reagents, but I haven't bothered trying to separate garden lime and highly doubt it would be worth the effort.
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[*] posted on 19-4-2016 at 08:22


Looking at the composition you are probably dealing with Dolomite CaMg(CO3)2 a double carbonate that is much less reactive than CaCO3 mixed with MgCO3. A mechanical mixture would be easy to separate by leaching out the calcium carbonate with hydrochloric acid in which magnesium carbonate dissolved rather slowly. The double carbonate can't be separate so easily though it dissolves slowly in hydrochloric acid to give you a mixture of very soluble chlorides. It doesn't sound worth it to me.
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[*] posted on 19-4-2016 at 08:24


Wiki says the magnesium carbonate melts (trihydrate) at 165 C and decomposes at 350 C.

Calcium carbonate isn't going to melt much sooner than steel.

I'd just heat the mixture a lot and see what happens.

Looking at the densities, the Ca carbonate might float so you can skim it off, or sink so you can pour the liquid Mg carbonate off.

Might not be the best purification process ever conceived, but it's an idea.




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[*] posted on 19-4-2016 at 08:48


Quote: Originally posted by aga  
Wiki says the magnesium carbonate melts (trihydrate) at 165 C and decomposes at 350 C.

Calcium carbonate isn't going to melt much sooner than steel.

I'd just heat the mixture a lot and see what happens.

Looking at the densities, the Ca carbonate might float so you can skim it off, or sink so you can pour the liquid Mg carbonate off.

Might not be the best purification process ever conceived, but it's an idea.


That just might work.
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[*] posted on 19-4-2016 at 09:10


They are probably soluble in each other when you melt one. See IUPAC solubility data series and you will discover that almost all the substances are soluble in each other. Even elements, like Fe in Hg, or Fe in Na. So melting is bad idea for separation.

If you only want to separate Ca from Mg, and you don't care about carbonate ion, then the easiest way is to add hydrochloric acid in small portions (for example 5 to 10 times) to start converting one carbonate to chloride...and until you start converting another also to chloride. But test between each adding, and mix them well and give time to settle and decant. You can use flame test to know when another has started converting to chloride (Mg is white, Ca is red/orange). Or react with sulfuric acid and test for precipitate, if it exists, it's Ca, if not it's Mg. This works because one will always rather be some compound than the other.
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[*] posted on 19-4-2016 at 09:47


Interestingly, CaCO3 has been shown to crystallise in different ways under the influence of a magnetic field.

Perhaps melting then recrystallising in a magnetic field would help.

http://www.sciencedirect.com/science/article/pii/S0043135406...




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[*] posted on 19-4-2016 at 10:00


There is an industrial process for preparing magnesia from dolomite... I am not sure about the details, but they can't be hard to find.

I still think it would be tremendously less work to buy magnesium and calcium salts separately.
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[*] posted on 19-4-2016 at 10:05


Surely the Fun is in the Doing ?

Buying is much less fun.




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Chemist_Cup_Noodles
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[*] posted on 19-4-2016 at 10:21


Quote: Originally posted by aga  
Surely the Fun is in the Doing ?

Buying is much less fun.


Ah, precisely! Sometimes it gets hard to explain to friends how much more satisfying it is to make your own intermediaries. Syntheses can be so much more entertaining than just having whatever it is show up on your porch.




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[*] posted on 19-4-2016 at 11:16


Quote: Originally posted by Chemist_Cup_Noodles  
it gets hard to explain to friends how much more satisfying it is to make your own intermediaries

More seriously, how much very basic Knowledge has been lost due to people simply buying stuff ?

The current crop of youth (in my experience) seem very able to randomly press bits of hand-held glass, yet utterly lost if asked to do something as basic as starting a fire without the assistance of Google a lighter or a match.




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[*] posted on 19-4-2016 at 11:54


Dissolve it in vinegar to get the mixed acetates
Add dilute sulphuric acid to ppt the CaSO4.2H2O
Filter it off and boil it with dilute Na2CO3 to get CaCO3
Add Na2CO3 to the Mg sluphate/acetate mixture to get Mg CO3

The big problem is that CaSO4 doesn't ppt well- it's often sludgy and hard to filter.

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[*] posted on 21-4-2016 at 05:28


Quote: Originally posted by aga  

More seriously, how much very basic Knowledge has been lost due to people simply buying stuff ?


That's honestly where I've gotten 60% of my current knowledge of chemistry, as I've only taken high school honors chem and AP Chemistry. I guess I'm one of the "current crop of youth", but at least I know a couple ways to start a fire. (you rub sticks just like in the movies, right? /s)

But @unionised, do you think plain vinegar might be a little too dilute? At 5% it would take a lot of vinegar if he's got a likely several pound bag of garden lime to work with. I'm not saying he needs glacial but it would help speed things up. It is true that calcium sulfate can be annoyingly hard to filter off, but for the most part you can try to decant the majority your solution off of it and through a filter, and then pipette out some more solution. Then you can just put the sludge alone in a filter and wash it a good bit.




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[*] posted on 21-4-2016 at 05:35


Gravity filtering CaSO4 is not easy, but you can usually decant, vacuum filter, or wrap it in something like an old t-shirt and wring it out.
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JJay
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[*] posted on 21-4-2016 at 05:41


Quote: Originally posted by unionised  
Dissolve it in vinegar to get the mixed acetates
Add dilute sulphuric acid to ppt the CaSO4.2H2O
Filter it off and boil it with dilute Na2CO3 to get CaCO3
Add Na2CO3 to the Mg sluphate/acetate mixture to get Mg CO3

The big problem is that CaSO4 doesn't ppt well- it's often sludgy and hard to filter.



Are you sure you can boil CaSO4 with dilute Na2CO3 to get CaCO3? I'm just not seeing why that would work.

Edit: Scratch that - I do see how that could work, but I'm not sure it's practical. You'd have to separate the CaCO3 from unreacted CaSO4 (perhaps by forming the acetate or chloride) and then re-precipitate it by forming the carbonate if you want a pure product.


[Edited on 21-4-2016 by JJay]
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Chemist_Cup_Noodles
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[*] posted on 21-4-2016 at 15:01


Quote: Originally posted by JJay  
Gravity filtering CaSO4 is not easy, but you can usually decant, vacuum filter, or wrap it in something like an old t-shirt and wring it out.


Hmm, are you sure that the weave in the t-shirt wouldn't be too large? I think you might end up squeezing out some CaSO4. Then again, I have never tried it (and maybe you have), so maybe it could work.

I've been doing some fierce googling though, and I soon noticed that calcium sulfate is somewhat soluble in glycerol. But alas, after I checked some experimental data it is about as soluble in low glycerin concentration solutions as it is in just water. However, what I haven't seen mentioned is the somewhat slight, but still notable acidic nature of CaSO4. So with enough water and patience, in a 100oC solution with the slightly basic Na2CO3 you might be able to get a good reaction. I'm not sure if your water would still boil at 100oC depending on at what elevation you live and how much of the salts actually dissolve. If you keep an eye on the pH you could probably tell when all of your CaSO4 is reacted back to CaCO3. Idk, do some math with the pKa's of each substance involved if you have the time and energy (I don't rn). But seeing how much more soluble Na2SO4 is in water, and hot water at that, than CaCO3, the two should separate rather well and render you a fairly pure product. CaCO3 is one of those quirky compounds that has a downward sloping solubility curve.




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JJay
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[*] posted on 21-4-2016 at 19:17


Quote: Originally posted by Chemist_Cup_Noodles  
Quote: Originally posted by JJay  
Gravity filtering CaSO4 is not easy, but you can usually decant, vacuum filter, or wrap it in something like an old t-shirt and wring it out.


Hmm, are you sure that the weave in the t-shirt wouldn't be too large? I think you might end up squeezing out some CaSO4. Then again, I have never tried it (and maybe you have), so maybe it could work.

I've been doing some fierce googling though, and I soon noticed that calcium sulfate is somewhat soluble in glycerol. But alas, after I checked some experimental data it is about as soluble in low glycerin concentration solutions as it is in just water. However, what I haven't seen mentioned is the somewhat slight, but still notable acidic nature of CaSO4. So with enough water and patience, in a 100oC solution with the slightly basic Na2CO3 you might be able to get a good reaction. I'm not sure if your water would still boil at 100oC depending on at what elevation you live and how much of the salts actually dissolve. If you keep an eye on the pH you could probably tell when all of your CaSO4 is reacted back to CaCO3. Idk, do some math with the pKa's of each substance involved if you have the time and energy (I don't rn). But seeing how much more soluble Na2SO4 is in water, and hot water at that, than CaCO3, the two should separate rather well and render you a fairly pure product. CaCO3 is one of those quirky compounds that has a downward sloping solubility curve.


I've wrung out CaSO4 filtered through an old T-shift before. Initially, some particles of CaSO4 might make it through the T-shirt, but a caked layer will build up on the T-shirt and eventually the liquid flowing through will be clear. You might have to filter it twice through the filter cake, or you could filter it through something finer once most of the CaSO4 has been removed.

The problem you are going to run into is CaCO3 caking around the CaSO4 and preventing it from reacting... it will be virtually impossible to completely react the CaSO4... or so I believe. So you'll probably need to separate some low solubility calcium salts.



[Edited on 22-4-2016 by JJay]
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[*] posted on 23-4-2016 at 01:42


Conversion of carbonates to hydrogen carbonates has the advantage of being more easily reversible than conversion to other salts.
I see that CaCO3 solubility in 0,35 mbar CO2 solution is 0,47 mM, with pH of 8,27.
In 1000 mbar CO2 solution, at saturation 6,6 mM Ca is dissolved, with pH of 5,96.
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