I attempted to perform a double replacement between sodium carbonate and potassium sulfate.
Goal:
Fractional crystallization of potassium carbonate (purity not important)
Ingredients:
-Na2CO3 (washing soda) anhydrous white powder, kitchen quality. quite pure
-K2SO4 (sulfate of potash) anhydrous white prilled, fertilizer grade. trace impurities/organics.
Method:
I prepared saturated solutions of sodium carbonate and potassium sulfate, in separate beakers. Dissolving the carbonate was so exothermic no heating
was required and left a hot cloudy particulate that eventually faded but still noticeable (hopefully just microscopic bubbles)
Note that these salts were sourced from supermarkets, the sodium carbonate was anhydrous and quite pure (consulted manufacturers msds just to confirm
it contained no percarbonate aka napisan).
However the K2SO4 was fertilizer grade, and on inspection, contained traces of a black organic matter, smelled funky aswell, filtered easy though.
Manufacturer's MSDS didn't reveal anything out of the ordinary, (100% purity, appeared legitimate claim). Testing with CaCl2 and weighing CaSulfate
precipitate revealed 94-96% purity, okay then.
Both solutions were mixed in beaker @ 1:1 molar ratio, from the following assumption;
K2SO4 + Na2CO3 > K2CO3 + Na2SO4
Simple enough, first one to salt out loses. The steep solubility curve of sodium sulfate was promising. Just for reference, the solubilities out of a
chem book:
Na2SO4 = 7.9g/L @0degC
K2CO3 = 154g/L @0degC
Yet upon chilling... nothing precipitated
TLDR/conclusions: The whole exercise seemed so deceptively simple. I expected crystals of sodium sulfate to reveal themselves, after saturated
solutions of Na2Co3 and K2So4 were mixed and chilled, but no dice. I will try
again tomorrow just to be sure.
What's the catch? Could someguy kindly explain what happened here? I just need to know if this reaction is even possible before I reflect upon my
sanity/stupidity.
Thank you
Roxannebbartlog - 4-11-2013 at 08:45
Would be nice to have actual amounts...
Also, it is not the solubility of sodium sulfate versus potassium carbonate that is most important here, but the solubility of sodium sulfate versus
potassium sulfate. Sodium sulfate is slightly less soluble than potassium sulfate on a molar basis at 0C (I get 0.33M for sodium sulfate versus 0.45M
for potassium sulfate).
Regardless, it seems like it should work. In my experience with sodium sulfate, it is very prone to supersaturation; you may have to encourage the
formation of crystals in some way. Sometimes waiting a day or two also worked for me. bbartlog - 4-11-2013 at 10:12
On further research, there is another issue: potassium and sodium form a sulfate double salt,
K<sub>3</sub>Na(SO<sub>4</sub><sub>2</sub>.
That will probably crystallize out first. blogfast25 - 4-11-2013 at 13:57
@bbartlog:
You have a reference for this?bbartlog - 4-11-2013 at 19:18
p559 of Atherton Seidell's 1919 book 'Solubilities of Inorganic and Organic Compounds'.
Neither of these necessarily implies that precipitating Na<sub>2</sub>SO<sub>4</sub> alone is impossible - at any rate it
sounds like precipitating only the double salt requires fairly narrow conditions. But it's certainly a potential complication, and places further
limits on the degree to which sodium can be eliminated from the liquid by fractional crystallization alone.Anomalous - 5-11-2013 at 08:54
Several methods for the production of K2CO3 are discussed there. There are advocates for simply buying it, while others suggest using a Seltzer maker.
blogfast25 - 5-11-2013 at 13:38
Hmmm... Glaserite. Interesting.Roxanne - 6-11-2013 at 08:50
Thank you gentlemen for your replies
I think I shall pass on this tedious exercise and maybe attempt the other methods suggested. Chemistry is only a recent hobby i've picked up and as
such still an amateur.
<sub>2</sub>.
That will probably crystallize out first.