Thomas Winwood
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Regarding acquisition of pure sodium chloride from cooking salt...
I have a 1.5kg bag of cooking salt which I'd like to convert to sodium chloride.
"I hate to think what YOU put on your chips", I hear you say. My problem is the anti-caking agent in it - sodium hexacyanoferrate. How would
I go about seperating the useful chloride from the useless hexacyanoferrate?
My initial ideas consist of some solvent which will dissolve sodium chloride but not sodium hexacyanoferrate. Unfortunately I've not been able to
find any which fulfil these criteria...I might ask at school (they have all manner of compoundss, it's like the Holy Grail) but I'll ask
here first in the hope of getting a quicker answer.
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BromicAcid
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Dissolve as much sodium chloride as will dissolve in distilled water, bubble HCl gas through the solution, very pure NaCl will precipitate out, this
is the way it used to be done in lab experiments, it's a fun little experiment. There will be somewhat substantial looses though as you
won't be able to precipiate everything, plus it uses up HCl. Maybe just make a concentrated boiling solution of NaCl and cool to precipitate
NaCl.
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garage chemist
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The solubility of NaCl doesn't change with temperature. The diagram temperature vs. solubility is a flat horizontal line.
That's the reason why the HCl method was developed!
HCl can be made easily by heating NaCl (the "impure" stuff can be used) with NaHSO4 (can be got as "solid pH down" in pool shops).
I wouldn't bother to purify the salt, the sodium hexacyanoferrate is only a very small percentage and normal cooking salt can be used for all
syntheses.
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Thomas Winwood
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Thanks to both of you, especially for the instructions for producing HCl without electrolysis. That saves me the problem of reacting the hydrogen and
chlorine (although now I can just drive off the useless and difficult-to-store gases and be left with useful sodium hydroxide).
Once I get better equipment for storing gases (I might see if the school will give or sell some of their lab glass and/or reactants, and/or tell me
some good suppliers) I'll store the hydrogen and chlorine also.
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Thomas Winwood
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Now let's switch emphasis slightly...
I'm making a new post for a number of reasons:
1) It draws emphasis.
2) It wouldn't be right to make a whole new topic because the question is so related to the last one, but editing the above post would not draw
attention.
I read in my chemistry textbook that industrially an asbestos diaphragm is used to prevent the chlorine from rereacting with the NaOH. This was
confirmed by a Wikipedia hunt. Is this a problem on the small scale? What should I modify from the simple electrolytic cell with two copper electrodes
(although the book recommends one platinum and one steel, I haven't access to either material - what can I substitute?) to avoid the chlorine
rereacting with the NaOH?
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Marvin
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It seems like a great idea, to turn something so common into 2 things so useful, but its mostly a waste of time. If you need NaOH you can buy it
easily, if you want chlorine you can make it chemically.
For test tube amounts of chlorine no diaphram is needed if you have high current, say 12V at 10A. If you dont have a high current you arnt going to
make much anyway. Eats carbon rods like nothing on earth. If you have HCl, electrolysing this dilute for chlorine is easier.
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Thomas Winwood
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Fair enough. I'll find something more interesting to electrolyse.
I was intending on using copper wire for my electrodes over graphite for that reason exactly. (Copper is easier to get hold of too.)
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Marvin
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Copper will just go into solution, graphite is the next best thing to something like platinum. If you are trying to support carbon electrodes under
the surface of the liquid, you need to make sure the connection between carbon and copper wire is completely liquid tight.
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