Synthesis of Potassium Perchlorate

[sonlightcity]

Hello All!

 

I've been doing a lot of research recently on the synthesis of chlorates and perchlorates (specifically potassium). However, no where have I found a single comprehensive article on how the synthesis of potassium perchlorate is accomplished, and what the differences are between the production of the chlorate and the perchlorate. I know there is a lot of misinformation on the subject out there, and that many people make it more complicated than it needs to be. Therefore, I was hoping for some help on exactly how to produce potassium perchlorate and possible the why behind the different steps. I am looking for the EASIEST and QUICKEST way possible to produce it through electrolysis.

 

Here are a few of the questions on my mind.

 

Is the amount of amps you have running through your cell important the differences between producing chlorates and perchlorates?

 

How well will a titanium anode work for producing perchlorate?

 

I'm planning on using a 6v, 10amp battery for my cell. Will this work for perchlorate?

 

Also, can any answers please be explained in simple terms with correct grammer? I've had trouble reading posts on this topic in the past because they've either used terms that are hard to understand or because they've used incorrect grammer.

 

Thanks in advance for the help!

[John Cuthber]

"I've been doing a lot of research recently on the synthesis of chlorates and perchlorates (specifically potassium)."

Why?

"because they've used incorrect grammer."

LOL

Edited February 23, 2013 by John Cuthber

[sonlightcity]

"I've been doing a lot of research recently on the synthesis of chlorates and perchlorates (specifically potassium)."

Why?

"because they've used incorrect grammer."

LOL

Parially just out of curiosity, but mostly because I need potassium perchlorate to use in a propellant. It's more stable than potassium chlorate and I don't want it exploding on me. lol

[elementcollector1]

MMO, lead dioxide or platinum anodes will produce perchlorate. Dissolve chlorate at room temperature, and electrolyze: Perchlorate precipitates out due to lower solubility (at least in the case of potassium).

[ajkoer]

As an alternative to the above, consider the following preparation. Per Wikipedia (http://en.wikipedia.org/wiki/Chloric_acid ):

"It [Chloric acid] is stable in cold aqueous solution up to a concentration of approximately 30%, and solution of up to 40% can be prepared by careful evaporation under reduced pressure. Above these concentrations, and on warming, chloric acid solutions decompose to give a variety of products, for example:

8HClO₃ → 4HClO₄ + 2H₂O + 2Cl₂ + 3 O₂
3HClO₃ → HClO₄ + H₂O + 2 ClO₂

The decomposition is controlled by kinetic factors: indeed, chloric acid is never thermodynamically stable with respect to disproportionation."

So upon mild heating of solid NaClO3 and Oxalic acid dihydrate:

2 NaClO3 + H2C2O4.2H2O --> 2 HClO3 (g) + 2 H2O (g) + Na2C2O4 (s)

and condensing the vapors, one could theoretically have highly conc HClO3, and with careful evaporation under reduced pressure, the disproportionation to HClO4, with possibly around 33% to 50% of the Chloric acid converted into HClO4. However, the yield is most likely lower for several reasons including the need for an excess of NaClO3 as otherwise, Oxalic acid may reduce the HClO3 to explosive ClO2. I would still, however, consider this preparation as potentially dangerous, and exercise appropriate safety precautions, especially limiting the quantities involved. Also, be aware that although ClO2 has a more pleasing smell than Chlorine, it is many times more poisonous as measured by recommended exposure limits.

Edited April 30, 2013 by ajkoer

[John Cuthber]

As an alternative to the above, consider the following preparation. Per Wikipedia (http://en.wikipedia.org/wiki/Chloric_acid ):

 

"It [Chloric acid] is stable in cold aqueous solution up to a concentration of approximately 30%, and solution of up to 40% can be prepared by careful evaporation under reduced pressure. Above these concentrations, and on warming, chloric acid solutions decompose to give a variety of products, for example:

 

8HClO₃ → 4HClO₄ + 2H₂O + 2Cl₂ + 3 O₂

3HClO₃ → HClO₄ + H₂O + 2 ClO₂

 

The decomposition is controlled by kinetic factors: indeed, chloric acid is never thermodynamically stable with respect to disproportionation."

 

So upon mild heating of solid NaClO3 and Oxalic acid dihydrate:

 

2 NaClO3 + H2C2O4.2H2O --> 2 HClO3 (g) + 2 H2O (g) + Na2C2O4 (s)

 

and condensing the vapors, one could theoretically have highly conc HClO3, and with careful evaporation under reduced pressure, the disproportionation to HClO4, with possibly around 33% to 50% of HClO4 formed. However, the yield is most likely lower for several reasons including the need for an excess of NaClO3 as otherwise, Oxalic acid may reduce the HClO3 to explosive ClO2. I would still, however, consider this preparation as potentially dangerous, and exercise appropriate safety precautions especially limit the quantities involved.

Heating strong oxidants like sodium chlorate with strong reductants like oxalic acid is a recipe for disaster.

It's the sort of thing that proves the adage that a little learning is a dangerous thing.

[ajkoer]

John Cuthber:

 

Yes, I agree that it is a dangerous preparation, but having started a synthesis thread using Oxalic acid on various salts on Sciencemadness, the actual salt that did prove problematic was not the chlorate, but would you believe, the sulfate. Apparently, heating H2C2O4 and a sulfate gets to a point where the H2SO4 is so concentrated it literally shouts out of the test tube. Many chemists, including those with experience, are generally unaware of the nasty behavior of some common acids that are apparently very dangerous when highly concentrated (as they are not, for good cause, available).

[John Cuthber]

Once more, you seem to be talking hogwash.

Oxalic acid decomposes on heating to form a bunch of gases which may shoot stuff out of the test tube. (Especially since it's commonly encountered as the hydrate)

On the other hand, sulphuric acid can be boiled in a test tube (albeit carefully).

Oxalic and sulphuric acids can both be bought over the counter in the UK so describing them as "not available" is rather misleading.

 

However, perhaps the most serious error in your earlier post is that it simply won't work.

the reaction of oxalic acid and a chlorate doesn't give chloric acid.

It gives chlorine dioxide.

It's one of the standard methods for preparing that gas.

http://chemicals.etacude.com/c/chlorine_dioxide.php

 

 

Next time you are going to post something, perhaps you should consider not bothering unless you have learned a lot more chemistry.

 

Even if it had worked (and, of course it gives you a toxic explosive gas instead of working) it doesn't give the product that the OP was asking for so , to summarise.

Your suggestion was

Inappropriate since it didn't give the product asked for.

Inaccurate, since it didn't do what you said it would do and

Bloody dangerous as it gives rise to a rather treacherous product instead..

 

Next time you are going to post something, perhaps you should consider not bothering unless you have learned a lot more chemistry.