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Author: Subject: Separation of NaCl and KCl
kyro8008
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[*] posted on 24-8-2005 at 11:24
Separation of NaCl and KCl


Hi, I was wondering if anyone could help. I was shopping today and found some "Lo-Salt" and it said low Sodium content, so I looked to see what it contained out of interest.

It contains 66% Potassium Chloride and 33% Sodium Chloride. I was wondering if there was an easy way to separate these two compounds as the KCl would be useful to me, and as it was so cheap would be a good source of it.

I cant think of anyway chemically to separate them, but alternatively some sort of way to separate by their weight difference might work? Not sure.

Any help would be appreciated, thanks!
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chemoleo
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[*] posted on 24-8-2005 at 11:50


Let me guess, is it for precipitation with NaClO3?

Anyways, I think there was a thread on this already... couldn't find it though. Maybe try google. But essentially, you can do careful crystallisation... and that's about the only practical one. YOu could use the fact that one salt is better soluble in i.e. Ethanol than the other. It sure is bothersome.

[Edited on 24-8-2005 by chemoleo]




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chromium
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[*] posted on 24-8-2005 at 11:53


Solubility of NaCl is almost independent of temperature. KCl on contrary is much more soluble in hot than in cold water. You have to saturate very hot water with mix of NaCl and KCl, then let it cool. Precipitate is almost pure KCl.
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kyro8008
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[*] posted on 26-8-2005 at 13:00


Thanks for the quick responce, I did as you said and it seemed to go well. I did a flame test and unfortunetly did'nt see much lilac, but ive always found K flame tests a pain to see anyway.
Apart from molten electrolysis is there another way to find if it is potassium?

Thankyou.
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12AX7
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[*] posted on 26-8-2005 at 13:35


Purify it again if you want, sodium shows up at very low precentages.

There would be Na present for two reasons: 1. evaporation during crystallization, causing NaCl to fall out, 2. excess NaCl after draining the KCl crystals.

BTW, does anyone know how the equilibrium works on a situation like this? For instance, boil water with an excess NaCl + KCl mixture. KCl is more soluble, so it dissolves in preference to NaCl, but does it "salt out" and prevent much (if any) NaCl from dissolving? Or does the NaCl cause it to salt out? Does it dissolve (and salt out likewise, if any) depending on the ratio of surface area?

Tim




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chromium
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[*] posted on 26-8-2005 at 14:09


Flame test is so sensitive to sodium that even minutest amount of it will turn flame yellow no matter what other substances there are. So its no help here. It will probably work if you recrystallise your KCl some more times.

There are ways to make sure your salt is KCl and not NaCl - such as exploring crystals or measuring specific gravity - but why to test? Just use it!
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chromium
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[*] posted on 26-8-2005 at 15:41


to 12AX7: NaCl and KCl both dissolve in remarkable quantities. If temperature is 100C and we take 100g of water then liquid phase over mix of both salts will contain 27.5g of NaCl and 35.3g KCl after equilibrum is reached. If temperature is 10C those numbers are 30.6g and 12.9g respectively. So obviously NaCl will not precipitate if hot solution cools down.

I have lot of tables that describe equilibrum in various complex solutions. As much as i know those are just measurement data. To be honest i do not know can what theory lies behind this.
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[*] posted on 27-8-2005 at 00:03


Just like to ask, regarding the conversion of NaClO3 to KClO3, whether i can use KMnO4 to convert instead of the KCl, as KCl is such a pain to seperate from NaCl....I have pure KMnO4 so I would easy to just add KMnO4 instead of going through all the steps of separating KCl.



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[*] posted on 27-8-2005 at 00:24


Just use the salt mix as is. No need to purify as the NaCl just stays in solution and KClO3 will precipitate. Recrystallize the KClO3 few times if you want it free of sodium salts.
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[*] posted on 27-8-2005 at 16:30


A couple of other sources for KCl that might be easier to purify:

From a garden supply/ agricultural store -- they may have it as "Muriate of Potash". I have usually seen it with the granules a rusty brown color due to the the impurity [or anti-caking agent] which appears to be an iron oxide or hydroxide, and which is insoluable so easier to separate. It is also slightly magnetic. Price usually around $1.20 (US) a pound.

If you live in an area that gets snow in the winter look for KCl as an ice melter -- considered more "environmently friendly" around lawns and gardens than NaCl. I have found this to be a better source -- purer, and cheaper, about $0.40 to $0.45 a pound. Probably can use "as is" for KClO3 manufacture.




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ChemicalBlackArts
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[*] posted on 27-8-2005 at 18:56


If you can procure any sodium cobaltinitrite (Na3Co(NO2)6), that can be used to seperate any potassium ions from solution, as it will precipitate the wholly insolube complex K3Co(NO2)6. The compound is awfully useful in qualitative analysis.



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[*] posted on 27-8-2005 at 19:55


Fascinating, but not helpful in our case. Purified cobalt compounds are outside the ameteur chemist's price range, I'm afraid.
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[*] posted on 27-8-2005 at 21:50


Ok ok ok, perchloric acid. :P

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[*] posted on 2-9-2005 at 04:48


Just thought I would throw this one in:
KCl is radioactive
NaCl is not
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[*] posted on 2-9-2005 at 06:11


KCl radioactive? You must be kidding, we must all be irradiated to burnt chops by now by all the Lo-sodium salt.:o



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[*] posted on 2-9-2005 at 06:32


No, he's not kidding. KCl is radioactive.

Not radioactive enough to worry, though. You get more rads from other sources around you.
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[*] posted on 2-9-2005 at 06:43


Yeah but if I sprinkle it on my food and eat it, the potassium will come into VERY close contact with my body. I suppose much will enter the blood stream and exit through the kidneys.

Stick with high sodium salt.

I've tried putting solid KCl next to a Giger counter I got 40 to 50 counts per minute. The background count was about 8 counts per minute.
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neutrino
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[*] posted on 2-9-2005 at 07:05


That's not really that much of a difference, IMO.

The radioactivity is from the .012% of naturally-occuring K-40.
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[*] posted on 2-9-2005 at 17:57


Quote:
Originally posted by neutrino
Fascinating, but not helpful in our case. Purified cobalt compounds are outside the ameteur chemist's price range, I'm afraid.


Actually, in a rather recently produced chemistry set that I acquired, it contains 10g of a 25% CoCl2 mix with NaCl. I had 2 bottles of this so the last few days I have worked it up to a pure cobalt(II) nitrate solution which is currently evaporating to get the crystals.

EDIT: Checked brauer, the compound mentioned which would ppt potassium is a rather simple synth from cobalt nitrate. When I get the cobalt nitrate crystals I may give a synth of that a try.:)

[Edited on 3-9-2005 by rogue chemist]




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[*] posted on 3-9-2005 at 04:17


"Yeah but if I sprinkle it on my food and eat it, the potassium will come into VERY close contact with my body. I suppose much will enter the blood stream and exit through the kidneys.

Stick with high sodium salt. "
Be sure to avoid food that contains carbon too.:)

On a practical note, while buying cobalt to prepare pure potassium compounds would work, it would be easier to just buy KCl.
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