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How to separate KCLO3 from KCLO4


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#1 Aspirin

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Posted 3 February 2005 - 01:57 PM

I Have NO idea

Any ideas.
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#2 Gilded

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Posted 3 February 2005 - 02:49 PM

You mean like you have a mix of both and you just want to get rid of the other?
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#3 Aspirin

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Posted 3 February 2005 - 03:44 PM

I have a mix of them and i want to separate them so i have only KCLO3 and KCLO4

Is this near to impossible
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#4 Gilded

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Posted 3 February 2005 - 06:23 PM

Hmmh. Perhaps a solubility separation method would be the best? I don't know how much their solubilities differ though, but just an idea. :P
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#5 Aspirin

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Posted 3 February 2005 - 06:33 PM

As far as i know

Potassium Perchlorate: 1.05g
Potassium Chlorate: 5g

This applies at 10 degrees celsius for 100 ml of water


But i don't know. Even at the lowest temp. some of both would dessolve so it would be rather impure.

What about if i eliminated 1 of the 2, how would that work
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#6 budullewraagh

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Posted 3 February 2005 - 09:22 PM

fractionally distill it. obtain crystallized product and remove
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#7 raivo

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Posted 4 February 2005 - 01:28 AM

It is possible to use that difference in solubilities to get almost pure product. You dissolve certain amount of mix in hot water. Then you cool that solution. If you choose quantities in certain well thought way then only one of salts will crystallise and another remains in solution. In practice some experimenting is needed because solubilities in such complex solutions are not exactly the same as in distilled water.
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#8 Hexaditon

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Posted 4 February 2005 - 01:44 AM

Do you mean fractionally crystallize it?
The chlorates and perchlorates will not distill at atmospheric pressure.... they are hardly in a liquid form before they decompose to liberate oxygen. I challenge you to produce a vacuum that can handle that sort of BP reduction. I will truly be impressed.

If you do mean fractionally crystallization it's already been the technique metioned ("solubility separation method" - Gilded).

Anyway off of that... I'll try to be helpful.
I have a chart for the solubility of KClO3 in water:
Temp oC : g/100mL h2o
0 : 3.3
10 : 5.1
20 : 7.3
30 : 10.2
40 : 13.9
60 : 24.0
80 : 37.7
100 : 57.7

I don't have a table for KClO4 but I have a couple values
@20oC it has a solubility of 1.8g/100mL h2o

Yeah.... I wouldn't say water is the best solvent for this application. Maybe you can look for charts for other solvents.

Allow me to heed warning on one thing though!
I do not know for what application you're trying to seperate the two nor how much impurities you have of each... (being more specific would allow us to be more helpful).

However there will STILL be TRACE amounts of KClO3 through solubility seperation (without more than huge losses in KClO4... if we can make it that far). So if you're worried about the KClO4 with a sulpher based composition I would probally say you'd be pretty safe from autoignition (use a tad bit of BORIC ACID [found at drug stores as a insect killer thing in powder] which acts as an inhibitor in the reaction that causes autoignition with KClO3 + S). If you're making something like Ammonium Perchlorate and worried about the impurities developing NH4ClO3.... Yeah I'd recommend scraping it for fireworks.

However, that's just me. All I'm saying is the impurities will still be there.
-Hexaditon
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#9 budullewraagh

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Posted 4 February 2005 - 01:53 AM

either way works, fractionally distilling or crystallizing. the former is of course more difficult to do than the latter, but more effective
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#10 Hexaditon

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Posted 4 February 2005 - 02:01 AM

Please explain to me as someone who doesn't understand - as a student: How would, physically... talk to me as if you're going to teach me something .... fractional distillation work any better in this case over fractional crystallization.

I will give you the benefit of the impossibility and even say that fractional distillation would be possible! (Which understanding anything about the process... the molecules would go through a chemical reaction under the theoretical thought conditions completly blowing your product; i.e. OXYGEN WILL BE LIBERATED .... no longer chlorate and perchlorate.....no longer a useful PHYSICAL seperation technique)

But I'm playing this "make believe" game for a second so you can explain to me how the former would be more effective.

-Hexaditon
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#11 budullewraagh

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Posted 4 February 2005 - 02:16 AM

think of le chatlier's principles. significant pressure will prevent the chlorate and perchlorate from decomposing.

now, think of this: we have our potassium chlorate and potassium perchlorate. if we were to fractionally crystallize it, there would be contaminants, ie there would be perchlorate in the chlorate and chlorate in the perchlorate. if we were to distill them instead, one would evaporate before the other, assuming we did not have a lump with a minimal surface area to volume ratio. then we would remove our product and heat to a greater temperature
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#12 Hexaditon

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Posted 4 February 2005 - 02:42 AM

Hmmm.... I must disagree with science to back me up.
First statement: You say that significant pressure will prevent the chlorate and perchlorate from decomposing. Yes at room temperature .... However these solids must be vaporized and in any pressure this will cause decomposition before it reached that state. Bringing up the pressure would bring up the boiling point (as well as melting point) and decomposition will still become the spontaneous reaction before it can ever vaporize. Make KClO4 vapor... and I think universities will be looking for you. It doesn't happen like that.

Anyhow yes yes I assume we both have agreement on the fact that impurities would arise from fractional crystallization. However, you don't seem to realize .... such would also arise in fractional distillation. Notice the keyword >Fractional< one has a tendency to evaporate more so than the other... but both liquids in fractional distillation vaporize to some extent. This is why seperation of ethanol and methanol is so difficult. The boiling points are so similar you need a fractionating column to seperate them more effectively and even then there are great impurities.

It's the same problem going both ways.
I can only judge by all your posts... but I constantly find myself tutoring principles of distillation in one aspect or another. Should I stress reading the distillation chapter of Vogel's 3rd (pratical organic chemistry) again? I mean you don't have to read it... but I learned a great deal, many people have learned a great deal. You can too. It's a marvellous resource. That goes for anyone interested in understanding the principles behind distillation (among many many other things, that book is wonderful).

-Hexaditon
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#13 budullewraagh

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Posted 4 February 2005 - 03:00 AM

you do not understand le chatlier's principles then.
ask yourself: how many moles of gas does potassium chlorate produce on decomposition and how many moles of gas does potassium perchlorate produce on decomposition? how many moles of gas do both of these produce when evaporated?
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#14 Hexaditon

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Posted 4 February 2005 - 03:06 AM

Clearly the concept of physical seperation is surpassing you.
The whole point of seperation is we do not WANT any moles of gas produced upon any sort of decomposition. We do not want decomposition. Decomposition implies .... well... decomposition.

Follow the topic. We are trying to seperate KClO3 from KClO4. We are not trying to decompose or chemically alter any of the compounds... that would not be fractional distillation anyhow. Fractional distillation is a physical seperation technique. There is nothing chemical about it. Decomposition does not = Fractional distillation. How you drew such conclusions is beyond me... As if boiling an egg was a physical change.

There is no chemical equilibrium between a physical process and a chemical process. That is a duh. I'm sorry I mean... That should be intuitive. Le Chatelier's principle does not imply to this situation since this is NOT an chemical equilibrium issue.
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#15 Hexaditon

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Posted 4 February 2005 - 03:12 AM

In two words. Neither evaporate.
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#16 budullewraagh

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Posted 4 February 2005 - 03:14 AM

you misinterpret as you often do. you should learn to make a conscious effort to understand someone's ideas even if you do not agree with them.

decomposition of alkali chlorates and perchlorates yields more gas than vaporizing them. thus, they will not decompose and will rather just vaporize. according to le chatlier's principles, the reaction that produces the least moles of gas will be favored
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#17 Hexaditon

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Posted 4 February 2005 - 03:19 AM

Again.... there is no chemical equilibrium with a physical process and a chemical process. Vaporization and decomposition are not in equilibrium.

If you are referring to decomposition and staying in an instable state... I don't want to be rude and bluntly say wrong again.... but yeah. KClO3 -> KClO + O2 is NOT in equilibrium. It proceeds in only one direction. And again a principle based on equilibrium does not apply.

But if you have some quantitative groundbreaking information that states otherwise.

Produce it. I'm interested in seeing such numbers.
The logical is still lacking however.

-Hexaditon
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#18 budullewraagh

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Posted 4 February 2005 - 03:36 AM

synthesis and decomposition are in equilibrium. synthesis yields the product, which happens to be in gasous state. makes sense.
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#19 Hexaditon

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Posted 4 February 2005 - 03:55 AM

Does it?
Tell the forum, for aspirin's benefit at what conditions this would apply (temperature/pressure)
With much confidence you must know what you're talking about.
Inform us. Along with that could you tell us how that is still fractional distillation - a PHYSICAL seperation technique. And why you're at it.... why the equilibrium of the 'gasses' (which is never in equilibrium in the first place - it proceeds in one direction... synthesis does not happen.... it's not like boiling a potassium hypochlorite in water.... I'm not sure if that even happens NaOCl -> NaClO3 in boiling water.... perhaps its similar) be advantagous in seperating KClO3 and KClO4?

-Justin
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#20 budullewraagh

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Posted 4 February 2005 - 04:10 AM

wow, you really dont know kinetics.

this fractional distillation would require pressures beyond those potentially achievable with commonly found household items.
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