RogueRose - 10-1-2018 at 11:59
I ended up getting a 5 new kg tub of bisulfate for about $6 (not a bad price!!) but I didn't see that it was only 93.6% pure. After checking the MSDS
it says that the other percentage is Na2CO3. This stuff is for pool ph control.
I'm not sure if it is worth trying to purify the bisulfate or if it is fine as is (maybe most applications it is fine as-is, such as decomp for SO2/3
production).
I thought it might be nice to have a few hundred grams of the purified product on hand so I thought maybe I'd give it a go.
it seems that NaHSO4 is soluble by about .25g/ml @ 25 deg C and 1g/ml @ 100C and fairly linear. That is a pretty interesting solubility and
temperature plot IMO.
Na2CO3 is less soluble as temp increases and is about .43g/ml @ 100C.
I was figuring that boiling DH2O with some biusulfate, filter quickly while hot and allow to cool slowly (maybe use a thermos??). Is there a time
frame I should shoot for to allow the crystals to be "selective" in their choice between the sulfate or carbonate? Are we looking at 2 hours, 4, 8,
12, 18, 24, hours, etc. I was planning on about 2 - 4L of liquid so about 200 - 400g of Bisulfate.
I can also do some cooling under vacuum if that could help the process. I'm not sure I've seen how crystal formation is effected under vacuum.
Thanks in advance!
happyfooddance - 10-1-2018 at 12:15
Sodium bisulfate is an acidic salt, it will readily react when dissolved with sodium carbonate, producing sodium sulfate and carbon dioxide gas. You
should be able to separate sodium sulfate from bisulfate by fractional crystallization, though I have not done it with these two chemicals.
Additionally, you could add an amount of sulfuric acid equal to twice the molar amount of carbonate. Two moles of sulfuric react with one mole of
sodium carbonate to give two moles of sodium bisulfate.