chemicalmixer
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Pearl Ash from KCl + Na2CO3 ?
From what I have gathered about the subject, it seems that there are several known viable routes for the home chemist to acheive the synthesis of
K2CO3, with some methods being more practical than others:
1) Partial thermal decomposition of potassium bicarbonate.
2) Thermal decomposition of cream of tarter/potassium bitartrate.
3) Partial carbonation of a KOH solution.
4) Boiling (NH4)2CO3 with KOH in solution.
5) Modified Solvay process, substituting trimethylamine for ammonia, and KCl for NaCl.
5) Modified Leblanc process, substituting K2SO4 for Na2SO4.
Actually, none of those methods seem particularly convenient for the average home experimenter, sans perhaps the first two (BTW, could someone please
describe how the thermal decomposition of potassium bitartrate proceeds?). Unfortunately, cream of tarter is expensive, and potassium bicarbonate or
hydroxide are not typically available at most consumer outlets. I realize that KHCO3 can be bought at most wine/brew supply stores, but then again
K2CO3 can also be ordered from specialty suppliers. I am more interested in producing K2CO3 from common grocery store items, purely for the academic
challenge. I thought that this might be a good beginner's experiment, since K2CO3 is a commonly used reagent.
I have an idea for a new, novel approach for synthesizing pearl ash from KCl (salt substitute) and Na2CO3 (washing soda), via double displacement
(2KCl + Na2CO3 --> K2CO3 + 2NaCl).
After a careful examination of the solubility properties of Na2CO3, KCl, NaCl, and K2CO3, from the available data at wikipedia's solubilty table (http://en.wikipedia.org/wiki/Solubility_table), I believe that I may have discovered a loophole, wherein the differing solubilities of Na2CO3,
KCl, NaCl, and K2CO3 in an H20 solution, and within a specific temperature range, might be exploited, in order to form NaCl and K2CO3 from salt
substitute and washing soda.
There appears to be, at between approximately 40 - 45 degrees celcius, a "sweet spot," at which NaCl becomes fairly less soluble in comparison to
Na2CO3, and also where KCl is more soluble than NaCl, meanwhile K2CO3 remains far more soluble, at most temperatures, than any of the other three
ionic compounds involved.
Might it be possible, with careful temperature control of a solution of Na2CO3 and KCl in water at about 43 degrees celsius, to slowly evaporate the
water from said saturated solution, and eventually crystallize out all or most of the NaCl from the solution, removing this salt by filtration while
the solution is still warm/hot, in order to leave behind K2CO3 in solution? Another idea I had was that, at the moment that a precipitate begins to
crystallize out of the solution, a few drops of just enough H20 could then be added to redissolve everything, and then a pinch of non-iodized kosher
salt could be added in order to seed and promote only the crystallization of NaCl out from the solution. Once the volume of the solution reaches the
theoretical one, filter out the crystallized salt.
Has this been tried before with any success?
[Edited on 20-3-2012 by chemicalmixer]
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weiming1998
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I don't know about the KCl+Na2CO3 process, but I have made K2CO3 using K2SO4(found in gardening stores as K fertilizer)+Na2CO3. It is more efficient
because Na2SO4 has a solubility of about 4g/100mls of water when refrigerated to about 4 degrees celsius, while K2CO3 has a solubility of about
100g/100ml at that temperature.
Your process will produce some K2CO3 with some careful temperature and water volume controls, that's for sure, but the solution that you extract will
have a higher percentage of contamination (mostly from unreacted Na2CO3+KCl/some NaCl dissolving in the solution instead of precipitating out.)
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chemicalmixer
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Interesting. I'd heard that using: K2SO4 + Na2CO3 --> K2CO3 + Na2SO4 works. I think the key here is that Na2SO4 becomes less soluble than K2SO4 at
around 0 degrees C, and so with careful cooling down to around that temperature, all or most of the Na2SO4 can be removed by filtration, leaving
behind a K2CO3 solution, which is very dilute.
The problem with that process though, is that it would be very tedious to carry out, since you are dealing with very dilute solutions (solubility of
Na2SO4 is only 47.6 g/L @ 0 deg. C) in order to isolate just a little K2CO4. A large initial amount of water would be needed just to get the reactants
dissolved. I just don't see how this would be more effecient then what I had proposed.
By using a similar process, substituting KCl for K2SO4, and utilizing heat and evaporation to acheive a fractional crystallization of NaCl (instead of
freezing/cooling), better yields of K2CO3 may be had, owing to the fact that the concentration of salts in a water solution will generally be greater
at warm/hot temperatures. Also, the degree of difference between the solubilities of Na2CO3, K2CO3, NaCl, and KCl seem more favorable for separation,
in comparison to the first scheme which utilizes K2SO4.
[Edited on 20-3-2012 by chemicalmixer]
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bbartlog
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| Quote: | | at between approximately 40 - 45 degrees celcius, a "sweet spot," at which NaCl becomes fairly less soluble in comparison to Na2CO3, and also where
KCl is more soluble than NaCl |
What's the objection to higher temps? 'Sweet spot' would imply that higher temps become unfavorable for some reason but I don't see why.
Anyway you are wrong: KCl is more soluble than NaCl at this temperature in terms of *grams per liter*, but for your purposes you need to worry about
solubilities in terms of *moles per liter* and by this measure the solubility of KCl does not exceed that of NaCl until somewhere around 80C (just
eyeballing the chart, here: http://www.kentchemistry.com/links/Kinetics/SolubilityCurves...).
| Quote: | | Also, the degree of difference between the solubilities of Na2CO3, K2CO3, NaCl, and KCl seem more favorable for separation, in comparison to the first
scheme which utilizes K2SO4 |
Handwaving BS to justify your scheme. Solubility of K2SO4 at 0C is about 0.4 mole per liter; for Na2SO4 it is less than 0.3 mole per liter. KCl is
only marginally more soluble than NaCl even at 100C. The 4:3 ratio is actually not that great for fractional crystallization, but it's manageable,
whereas you have the very marginal 7:6 ratio even at 100C.
The less you bet, the more you lose when you win.
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weiming1998
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Quote: Originally posted by chemicalmixer  | Interesting. I'd heard that using: K2SO4 + Na2CO3 --> K2CO3 + Na2SO4 works. I think the key here is that Na2SO4 becomes less soluble then K2SO4 at
around 0 degrees C, and so with careful cooling down to around that temperature, all or most of the Na2SO4 can be removed by filtration, leaving
behind a K2CO3 solution, which is very dilute.
The problem with that process though, is that it would be very tedious to carry out, since you are dealing with very dilute solutions (solubility of
Na2SO4 is only 47.6 g/L @ 0 deg. C) in order to isolate just a little K2CO4. A large initial amount of water would be needed just to get the reactants
dissolved. I just don't see how this would be more effecient then what I had proposed.
By using a similar process, substituting KCl for K2SO4, and utilizing heat and evaporation to acheive a fractional crystallization of NaCl (instead of
freezing/cooling), better yields of K2CO3 may be had, owing to the fact that the concentration of salts in a water solution will generally be greater
at warm/hot temperatures. Also, the degree of difference between the solubilities of Na2CO3, K2CO3, NaCl, and KCl seem more favorable for separation,
in comparison to the first scheme which utilizes K2SO4.
*
I also wanted to add that, although it might be difficult to test for sodium impurities in the final K2CO3 product, maybe by first neutralizing the
K2CO3 solution with acetic acid, and then precipitating any residual Cl- ions that might be present, by adding lead acetate solution until no more
precipitate forms, filtering and weighing it, could give at best a rough determination of how effective the selective fractional crystallization is,
in getting mainly the NaCl out of the solution.
[Edited on 20-3-2012 by chemicalmixer] |
You do not need to actually dissolve all the K2SO4 and Na2CO3 in order to make K2CO3.(Think of the reaction between Ca(OH)2 and Na2CO3/K2CO3) The
relative insolublity of Na2SO4 at room temperature (less than 32oC because the solubility suddenly spikes up to 40g/100ml.) will drive this reaction
forward. Heating the water, then cooling could also be used to dissolve as much of the reactants as possible. As it cools, the sudden decrease in the
solubility of the Na2SO4 will cause it to crystallize out quickly.
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chemicalmixer
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| Quote: | | Quote: | | Quote: Originally posted by bbartlog |
What's the objection to higher temps? 'Sweet spot' would imply that higher temps become unfavorable for some reason but I don't see why.
|
At about 40 deg. C, the difference in solubility between NaCl and Na2CO3 is the greatest (according to wiki data), and at temperatures higher than
this, NaCl continues to increase slightly in solubility, meanwhile as Na2CO3 starts becoming increasingly less soluble.
I didn't realize that the solubilty in moles is what's important.
[Edited on 20-3-2012 by chemicalmixer] |
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Hexavalent
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Cream of tartar is potassium bitatrate, not KHCO3.
I can buy 2kg of potassium carbonate from the local lawn and garden centre for £3, labeled as carbonate of potash.
"Success is going from failure to failure without loss of enthusiasm." Winston Churchill
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chemicalmixer
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| Quote: Originally posted by Hexavalent | Cream of tartar is potassium bitatrate, not KHCO3.
I can buy 2kg of potassium carbonate from the local lawn and garden centre for £3, labeled as carbonate of potash. |
I edited my originial post to clarify what I meant. I realize that potassium bitartrate and potassium bicarbonate are two entirely different
compounds; what I intended to say is that, AFAIK, either compound may be subjected to heat in order to obtain K2CO3. I know that I read somewhere
(probably this forum) that potassium bitartrate can be used to conveniently prepare K2CO3 by way of thermal decomposition, however I am
unsure of the reaction scheme involved, or if this in fact would work.
Also, I have never seen "carbonate of potash" sold in garden centers here in the US, however apparently the bicarbonate is sold (I think at about 85%
purity, along with inert ingredients) for use in gardens as some kind of mild fungicide or something (Bonide Remedy).
[Edited on 20-3-2012 by chemicalmixer]
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bbartlog
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'NaCl doesn't become much more soluble, meanwhile Na2CO3 starts becoming less soluble.'
Yes, but it doesn't matter, because Na2CO3 is still vastly more soluble than NaCl at 100C.
Anyway, a more fundamental problem with your scheme is that there is no common ion effect to reduce the amount of NaCl present in the solution below
the level of normal saturation. If you start with stoichiometric Na2CO3 and KCl and reduce the volume of the solution at 100C, any liquid that remains
will end up saturated with NaCl. Now, at some point it will also be saturated with K2CO3, so I guess once you reach 100ml of water at 100C with 156g
of K2CO3 and 40g of NaCl, you've done as well as you can and can get 80% K2CO3 by weight.
The less you bet, the more you lose when you win.
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chemicalmixer
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Revised K2CO3 method:
I have been pondering a new way to make K2CO3 or even KHCO3 from common grocery/department store items for a while. I came across something pretty
interesting:
2KCl + 3MgCO3 + 9H2O + CO2 --> 2[MgKH(CO3)2 * 4H2O] + MgCl2
Apparently, when a KCl solution with MgCO3 suspended in it is pressurized with CO2 gas to a couple of atmospheres and agitated, a crystalline
power-like double salt is formed which can be filtered from the remaining MgCl2 solution, and rinsed with cold water to remove traces of MgCl2.
Finally, the double salt is decomposed by boiling, where it is converted to insoluble MgCO3, CO2 gas, and K2CO3 solution, which can be filtered, and
some of the by-product reused.
More info can be found at this link:
http://books.google.com/books?id=FQ4KAAAAIAAJ&pg=PA383&a...
Or, by doing a google search; there's a good deal of info about it.
I was thinking that this would be a far superior method for making K2CO3 from commonly available items when compared to the alternatives (Leblanc,
double displacement, etc.). The one think that would make this process simple to perform would be a Seltzer Maker! Either the
old-fashioned glass bottle type that uses the small CO2 cartridges, or the newer fancier type that uses a large disposable CO2 canister (which are
expensive, but the device can be easily modified with certain fittings to accept the cheaper, refillable paintball gun CO2 tanks). Also, not only
could K2CO3 be easily synthesized, but the seltzer maker could be used to perform a Solvay or even Merseburg Process [CaSO4 + (NH4)2CO3 --> CaCO3 +
(NH4)2SO4] with janitor's ammonia.
They only challenges/complications that I can think of are:
1) There would need to be a means of installing a magnetic stirrer below the seltzer bottle "reaction" vessel, and:
2) There would need to be a means of weighing the mass of the seltzer bottle as the weight of the reactants/products changes.
[Edited on 23-2-2013 by chemicalmixer]
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blogfast25
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Sorry to rain on the parade but this is likely to become on of these projects that doesn't lead to anything because the thread starter will eventually
just buy the desired product, like everyone else. K2CO3 is fairly OTC. Garden centers will stock it. eBay's pavements are probably lined with potash
sellers.
And exploiting 'loopholes' and 'sweetspots' in solubility systems of generally well soluble compounds is a recipe for heavily contaminated products
that require endless recrystallising...
Note also that the title is a little misleading: 'pearl ash' specifically refers to K2CO3 obtained from firing cream of tartar. In the heady days of
the 'soda wars' and the Leblanc process, pearl ash was the creme-de-la-creme in terms of K2CO3 quality. Linguistic erosion now makes it an old term
for potassium carbonate.
[Edited on 23-2-2013 by blogfast25]
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chemicalmixer
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^^ I understand what you're implying, but I think that it would be an interesting experiment. I'm more into inorganic chemistry at the moment, and I
would like to learn how to synthesize an array of ordinary, yet less available inorganic compounds, starting from more commonly available ones. I
mean, an ordinary seltzer maker has a practical use if you want to make soda also, but wouldn't it be amusing to use it to make K2CO3 from KCl salt
substitute, and magnesium carbonate from baking soda and epsom salt from the grocery store? Also, with plaster of paris and janitor's ammonia, you
could use the soda maker bottle to make ammonium sulfate via the Merseburg Process. I'm aware that most of these things can be found as fertilizers,
etc., but it's usually only in 2 to 5 lbs amounts or more, unless one resorts to ebay or internet/mail-order science reagent chemicals.
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chemicalmixer
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| Quote: | | [quote=blogfast25]Note also that the title is a little misleading: 'pearl ash' specifically refers to K2CO3 obtained from firing cream of tartar. In
the heady days of the 'soda wars' and the Leblanc process, pearl ash was the creme-de-la-creme in terms of K2CO3 quality. Linguistic erosion now makes
it an old term for potassium carbonate. |
Oops, I wasn't aware of the history of the term "Pearl Ash." To be honest, where I live in the US, it is kind of difficult to find manageable amounts
of either K2CO3, or KHCO3. I realize that wine-making/beer-brewing supply stores - including internet ones - will usually carry KHCO3 as a pH adjunct,
and that it could be heated like baking soda is heated in order to afford the carbonate.
Also, I think there is a product called Bonide "Remedy" that is an impure (>85%) K2CO3 which is used as some sort of fungicide. I haven't seen
that about at any stores though, plus it is impure.
KCl can be bought easily in small amounts (80ish grams) as a salt substitute, and epsom salts and baking soda, found at the dollar store, can be used
to make the MgCO3. The seltzer making systems are probably available now at Wal*Mart (I haven't checked yet), because I have been seeing them
advertised in TV commercials in the US. i'm just wondering if they could be modified somehow to install a computer-fan DIY magnetic stirrer somewhere
near the base of where the charging bottle sits.
If you look at the design of this particular soda-maker, you can see that there exists a gap in between the bottom of the charging bottle, and the
base of the appliance:
That space seems plenty big enough to accommodate the placement of a cheap, DIY magnetic stirrer to agitate whatever mixture is in the bottle.
[Edited on 23-2-2013 by chemicalmixer]
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