Chemistry_Keegan
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Potassium Iodide From Iodized Table Salt
So I would really like some potassium iodide, and couldn't help but notice that it was one of the ingredients in the table salt that we buy (Windsor
Iodized Table Salt). I am aware that you can buy tablets and such of potassium iodide, and it would almost certainly be much cheaper to do so, but I
would still rather get it from items that I already have, just for the fun of it. As for the ingredients in the table salt, they are as follows:
Sodium Chloride
Calcium Silicate
Sucrose
Potassium Iodide
So, does anyone know how I could go about separating the potassium iodide from this mixture?
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Vargouille
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It would be an exercise in futility. For every one gram of iodized salt, there is 77 mcg of iodide (0.000077 g). To get even 30 g of potassium iodide, you would have to get 100% recovery from over 400 26-oz containers of iodized
salt, at a cost of over $3000. For comparison, you can buy about 30 g of pure potassium iodide for $7.00.
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Chemistry_Keegan
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Oh, okay, scrap that idea. But let's say we had a mixture of just sodium chloride and potassium iodide (in much larger quantities) how would the
potassium iodide be extracted, just out of curiosity?
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Mixell
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There is barely any iodine in that salt.
Quote from wikipedia: Edible salt can be iodised by spraying it with a potassium iodate solution. 60 ml of potassium iodate, costing about US$1.15 (in
2006), is required to iodise a ton of salt.
Of-course it speak about iodate, but you can clearly see the scale.
You can try dissolving in it water, and adding copper ions, you'll get an insoluble and oxygen sensitive copper (I) iodide and elemental iodine.
You can also try oxidizing the iodide to elemental iodine and leeching it with a relatively non-polar and water immiscible solvent (like chloroform or
toluene).
http://en.wikipedia.org/wiki/Solvent
http://en.wikipedia.org/wiki/List_of_water-miscible_solvents
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Vargouille
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Potassium iodide is much more soluble in acetone than sodium chloride is (1.31 g/100 mL for KI; 0.000033 g/ 100 mL for NaCl).
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blogfast25
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A far more interesting idea, chemically speaking, is to extract iodide from kelp. Gather a rather large amount of kelp, dry it, burn it and leach the
ashes with hot water.
Oxidising the iodide to iodine then allows to separate it from other halides.
It takes a lot of kelp for a bit of iodide/iodine but at least you get an interesting lesson in history into the bargain!
[Edited on 21-2-2013 by blogfast25]
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Bezaleel
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Peach once did this, with a photograph novel-type post, somewhere on this forum. Great to read/watch through, I remember.
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Mixell
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I'd imagine you wouldn't be able to leach potassium iodide from sodium chloride/potassium iodide mixture. I think that acetone-soluble sodium iodide
will be leached and potassium chloride will remain.
But seriously, if you just need the salt, you can buy it easily, it's a pretty common substance.
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Vargouille
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If it's purely a mixture of NaCl and KI powders, then that won't happen, because no ion exchange would happen. If, however, it's a residue from a
solution of NaCl and KI, then that would occur because it's actually a mixture of NaCl, NaI, KCl, and KI in varying proportions. All depends on how
they're mixed.
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Henry
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and it would almost certainly be much cheaper to do so
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Mixell
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Well, potassium iodide is a bit soluble in acetone. And I'd imagine sodium chloride is soluble somewhere in the milligram range too (http://www.sciencemadness.org/talk/files.php?pid=198445&... although it doesn't mention 100% acetone).
Also, having potassium iodide in the solution might promote the corrosion of the NaCl crystal lattice and ion exchange at the surface of it.
That might take a while, but if you leave it for days/weeks (even without stirring, although stirring should really fasten the process) it should
work.
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Vargouille
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I mentioned the solubility of NaCl in acetone (based on the value on the Wiki) previously: 0.000033 g/ 100 mL. More in the microgram scale, methinks.
The change from potassium iodide to sodium iodide could happen, but it would be astoundingly slow due to the extremely low solubility of NaCl.
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