Want to dissolve it in water, filter out impurities, and then re-precipitate pure calcium carbonate.
Since calcium carbonate is not very soluble, I'd like to convert it into the BI carbonate (high solubility), filter, and then convert back to
carbonate and recover the precipitate.
Don't have a convenient source of CO2.
Is there a way to do this WITHOUT bubbling CO2 through a calcium carbonate / water slurry?
Any way to use something like SODIUM bicarbonate?
Any OTHER suggestions?
Thanks!White Yeti - 2-11-2011 at 12:46
If I were to obtain pure calcium carbonate, I would make it instead of purifying an impure sample. I made some relatively pure stuff not too long ago
by mixing some calcium chloride with some sodium bicarbonate.
I imagine if you mix pure calcium chloride with pure sodium bicarbonate, you'd get a very pure calcium carbonate sample.
Judging from your post, it seems like you're going for purity, not quantity. So here's my suggestion:
Take a sample of calcium chloride and dissolve as much as you can in boiling distilled water. Filter the solution and cool it down to as cold as you
can get it without freezing it. Collect the crystals and react them with an equally pure sample of sodium bicarbonate. It will bubble vigourously, so
be careful. The precipitate formed is calcium carbonate. Filter, wash with distilled water, and dry.Curious George - 2-11-2011 at 13:38
Thanks, Yeti, but the goal of this project is to recover pure calcium carbonate from an IMPURE sample.
Any suggestions for that?
Thanks!White Yeti - 2-11-2011 at 14:22
I suppose it also depends on what impurities are present. Some impurities can be dissolved alongside the calcium carbonate. You'll need to provide
more info on the composition of the sample in question.
What are you extracting the CaCO3 from? Eggshells? Chalk? If so, you could dissolve the CaCO3 in an acid (HCl, CH3COOH etc...) to get the
corresponding calcium salt, wait for fizzing to stop, filter, mix with a saturated NaHCO3 solution, wait for the fizzing to subside, filter, wash and
dry.
It would be incredibly time consuming [and wasteful], but you could give it a try
[Edited on 11-2-2011 by White Yeti]blogfast25 - 3-11-2011 at 02:52
WY: CaCO3 is notorious for co-precipitation problems, making it difficult to obtain pure (without some sodium contamination, for instance).
Precipitate from a fairly dilute solution for best purity results.
CG: as far as I know, dissolving CaCO3 in carbonated water is the only way to obtain Ca(HCO3)2.Think pure sparkling water, perhaps?ScienceSquirrel - 3-11-2011 at 03:35
A good route to very pure calcium carbonate would be to heat your calcium carbonate until it decomposes to calcium oxide, react with water to form a
solution of calcium hydroxide, filter and react with carbon dioxide to form the bicarbonate, filter again and then boil to precipitate the carbonate.
I suspect that this would be at least 99+% and probably a lot better!blogfast25 - 3-11-2011 at 07:10
That would get rid of sodium, I think. Takes a about 24 h in quite a hot furnace, though...White Yeti - 3-11-2011 at 11:43
WY: CaCO3 is notorious for co-precipitation problems, making it difficult to obtain pure (without some sodium contamination, for instance).
Precipitate from a fairly dilute solution for best purity results.
I said that already:
"I suppose it also depends on what impurities are present. Some impurities can be dissolved alongside the calcium carbonate. You'll need to provide
more info on the composition of the sample in question."WY
I know this is not the subject of the thread, but the OP should provide more info in order to get an informative answer. The least we need to know is:
-The source of the CaCO3
-The impurities in the sample
Suppose strontium and magnesium are present in the impure sample, there is no easy way that I can think of to remove these impurities due to the
similar behaviours of these three elements.Mr. Wizard - 3-11-2011 at 12:43
Calcium Carbonate is less soluble than Calcium Bicarbonate. There is a process that forms stalagmites and stalagmites in caves as the CO2 in the air
reacts with water and lime forming the more soluble bicarbonate, which is then carried with the water to the tip of a stalactite where is loses the
CO2 and precipitates the carbonate again.
Depending on the impurities, you might wash the powdered CaCO3 with pure water to dissolve any easily soluble matter. The remaining non soluble
material could then be placed in a cylinder of glass with a small amount of water and enough CO2 to form a solution. The action would be such that the
CO2 in the upper portion of a filter or porous cup dripped through the raw material, dissolving the formed bicarbonate, then dripping to the bottom
where gentle heat would drive the CO2 out of the bicarbonate and evaporate the water to an upper cooler area where it would fall again onto the raw
material.
The net effect would be a collection of carbonate at the bottom and non reacting material left behind. The amount of CO2 would have to be enough to
form the bicarbonate, but not so much as to preclude driving it out. Maybe with the right design it could be similar to a solar powered still.
A Sohxlet extractor comes to mind.Curious George - 3-11-2011 at 13:17
Mr Wizard, I think you are on to something!
Other than using gaseous CO2, is there any other way to get this effect? Any other way to generate bicarbonate ions?Mr. Wizard - 3-11-2011 at 18:26
Other than using gaseous CO2, is there any other way to get this effect? Any other way to generate bicarbonate ions?
You could use carbonated water as the source for the CO2, Schweppervescence ;-) I would guess the cheapest bottles of soda water would do the job,
preferably the ones made with pure water and CO2. Drying a drop or two on a glass plate will tell you if it has any soluble solids.
An educated guess would be that bicarbonate is always part of the equilibrium with water, CO2 and carbonates. Curious George - 4-11-2011 at 05:36
I'll give it a try -- THANKS!LHcheM - 8-11-2011 at 02:21
well, I think you need to do that in three steps...
Suppose the most troublesome impurities (Mg, Sr, Ba) are present:
1) Dissolve in HCl / HNO3, filter, and add excess hydroxide, Calcium hydroxide and Magnesium hydroxide precipitate out. Filter to get the residue.
2) Without drying the residue, dissolve it in HCl / HNO3 again. Add sulphite ions, Magnesium sulphite remains in solution, Calcium sulphite can be
filtered out.
3) Add the calcium sulphite into a solution of sodium bicarbonate, calcium carbonate, calcium bicarbonate, sulphur dioxide and sodium salts will form.
Evaporate the solution to a certain dryness, filter off the calcium carbonate (this is to increase your yield, as Calcium has high polarizing power,
its bicarbonate is not stable and decomposes to calcium carbonate as the solution gets concentrated)
Tedious, but see if it would work Acetic Acid - 8-11-2011 at 06:57
I think I've got it...
CaCl2 + 2NaHCO3 --> CaCO3 + CO2 + H2O + 2NaCl
CaCl2 is Damp Rid moisture absorber, or drive way de-icer perhaps. As you said these would be somewhat impure so you would want to recrystallize the
CaCl2.
NaHCO3 is baking soda. It's generally pretty pure so you shouldn't need to recrystallize.
Mix pure CaCl2 with stoichiometric amount of baking soda. Water may be necessary to start the reaction. Mixture will fizz due to production of gaseous
CO2.
CO2 will be lost as a gas. Water and NaCl will exist as a solution. The extremely insoluble CaCO3 will precipitate out. Filter the mixture to retrieve
the pure CaCO3.
CaCl2 is Damp Rid moisture absorber, or drive way de-icer perhaps. As you said these would be somewhat impure so you would want to recrystallize the
CaCl2.
NaHCO3 is baking soda. It's generally pretty pure so you shouldn't need to recrystallize.
Mix pure CaCl2 with stoichiometric amount of baking soda. Water may be necessary to start the reaction. Mixture will fizz due to production of gaseous
CO2.
CO2 will be lost as a gas. Water and NaCl will exist as a solution. The extremely insoluble CaCO3 will precipitate out. Filter the mixture to retrieve
the pure CaCO3.
Yay for figuring things out on my own! :->
If you read the entire thread, you wouldn't have to repeat something someone already said.Acetic Acid - 8-11-2011 at 14:02
Why is everyone on the forums like this? Shooting down the information I can offer and making me feel bad for posting it. Much appreciated. Panache - 8-11-2011 at 16:52
Acetic don't misinterpret wy, he's in no way trying to shoot you down. It's great that you figured It out however when you posted your reply it had
already been mentioned as a possible solution earlier in the thread. It's important you read the thread before replying so that you don't repeat what
has already been suggested.
Cool. Don't feel bad.Acetic Acid - 9-11-2011 at 06:01
Understood, thank youWhite Yeti - 9-11-2011 at 12:15
I'm sorry if it sounded like I was trying to shoot you down, that was not my intention.
