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[*] posted on 22-2-2009 at 03:14
Ir in perchlorate cell


I´ve now the chance to buy an Iridium foil so I´ve asked myself if it could be employed in a perchlorate cell, Ir should be a little bit more stable than platinum but many sources say that Pt should be stable enough which I found isn´t the case(and many others).

It is pure Ir, no alloy.

There are many Ir compounds, not unlike that it will also errode after a while as Pt does.

Has anybody experiences with an Ir anode?
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[*] posted on 22-2-2009 at 11:36


Hello Per,

AFAIK Pt is the most stable of all the Noble metals (the most Noble as it were). Perhaps I am wrong.
Platinum has been used to make Perchlorate in industry. ie. it is OK for the job. Excessive erosion probably comes from using it to reduce the Chlorate concentration to low levels (garage cells). If you use it to take Chloride solution all the way to Perchlorate which has low Chlorate concentration, that will cause even more Pt erosion. How much I don't know.
Perchlorate cells in industry are pH contolled (around neutral). I don't know why, but it may be to reduce anode erosion. pH controll in a Perchlorate cell will not increase current efficiency so (I guess) it may be done for erosion reasons.
Anodes do erode in industrial cells but is is considered economically OK at some grams per ton Perchlorate.
If you purchase one of those Ti anodes plated with a few micron of Pt and figure out how much Pt you have actually purchased you will be unpleasently surprised as to how little Pt you actually have.

What conditions are you using Pt (and the others) and getting erosion?

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[*] posted on 22-2-2009 at 11:57


All my attempts started with a almost concentrated NaCl solution and a platinum wire, 7-14V and temperatures of 40-65° in average.
The target was an easy, very long live and low cost cell, the efficiency wasn´t important, current isn´t so expensive but the cell shouldn´t be too sophisticated, so I don´t controlled the pH.
I always got perchlorates but the visible erosion of the wire increased with every run.

So whether the Ir is more unstable than the Pt the foil would be the wrong investment, thanks.
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[*] posted on 22-2-2009 at 12:10


Hello,

How much Chlorate, (or do you know) had you in the solution when the cell was considered 'finished'. How did you estimate/measure the amount of Chlorate contamination in end product?
What did you use to eliminate the last of the Chlorate (or did you eliminate it)?
How many KG's of Perchlorate (roughly) did the Pt wire make before it was eroded to nothing?


Also Ir is half the price of Pt. I think if it worked better than Pt it would be used all over the place.

Thanks,
Dann2

[Edited on 22-2-2009 by dann2]
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[*] posted on 22-2-2009 at 14:13


Actually iridium is IIRC "nobler" than platinum.

From every reference I have seen the durability of noble metal components in all anode schemes is multiplied several times by not using the noble metal itself as the anode, but in mixture with other materials or compounded
with other materials, or used in some bielectrode coating scheme. There is going to be erosion loss with any anode material and the economics have been studied carefully with regards to noble metal loss and found to be an operating expense wished to be reduced as much as possible, and there is also the current efficiency which
can be improved by composites which contain the noble metal as a component, beyond the efficiency which is realized using the noble metal alone. Therefore, if endurance and economics are any governing factors as they would be for any sizeable scale process, those factors are going to be unfavorable to the prospect of using anodes which are composed of the noble metal alone. Iridium has been used as a component in baked oxide titanium substrate anodes but I don't have the reference handy.
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[*] posted on 23-2-2009 at 08:20


Hello,

70/30 Pt/Ir alloy is the one to go for if you have access to such an alloy.
Pure Pt (easy to get) would be next best IMO.
How much pure Ir would erode I don't know.

There is some reading about the developement of the Pt (metal, metal alloy, and MMO) here:
http://www.geocities.com/CapeCanaveral/Campus/5361/chlorate/...

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[*] posted on 23-2-2009 at 08:57


Quote:
Originally posted by Per
All my attempts started with a almost concentrated NaCl solution and a platinum wire, 7-14V and temperatures of 40-65° in average.
I always got perchlorates but the visible erosion of the wire increased with every run.


If you start with a chlorate solution (bought or prepared using a graphite anode) erosion of platinum should be much reduced.
I'd hate to see a costly Pt. anode erode visibly. . .
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[*] posted on 1-3-2009 at 09:03


I undertook a few attempts in test tube and a few in a florence flask, with the test tubes I got just a few grams bit with the 500ml florence flask I harvested about 80g of the "product", I don´t know how much of if exactly was chlorate and perchlorate and how much chloride remained, I always added potassium carbonate to the finish solution to precipitate the perchloate, the hopefully perchlorate out.
I tested the product in a mix with charcoal and it always burned nicely;), so I assume that the most of it was precipitated perchlorate.
So as I sayd it was an very easy setup, the erosion of the Pt wire was in some experiments higher than in others, I tried also K2Cr2O7, adding HCl but erosion of the wire could never be prevent.

Buying chlorates is in the country I live a little bit difficult but this needs no more implementations, who cares if china is leading in sciences in a few decades?

Quote:
I'd hate to see a costly Pt. anode erode visibly

That´s why I stopped the program and looked for a better anode:)

[Bearbeitet am 1-3-2009 von Per]
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[*] posted on 1-3-2009 at 11:35


A polished Pt. surface will roughen and become matt as electrolysis proceeds but this isn't actually visible erosion.
Erosion does occur, but only AFAIK, to a very negligible extent.
A Pt. anode in a home set-up should last indefinitely.
You're unlikely to notice weight-loss at any time.
The problem with Ir. is its brittleness; it's used to harden platinum, but 70/30 alloys are fairly rare nowadays.
90/10 and 80/20 alloys are easier to get *and* are easier to manipulate.
IIRC, 70/30 was used at one time for syringe needles. . .
I'd continue with saturated NaCl solutions and use HCl to decompose residual chlorate in perchlorate solution.
The ammonium salt is more energetic than KClO4 but it is more soluble.

[Edited on 1-3-2009 by hissingnoise]
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[*] posted on 2-3-2009 at 11:35


Believe me, the Pt wire was really eroded, not just at the surface.

I never heared that HCl is able to decompose a chlorate solution but I know mixtures in which a HCl/chlorate solution is used as a very strong oxidiser.

NH4ClO4 is a good rocket propellant but I wouldn´t produce it with self made perchlorate solutions because of the hazards of possible NH4ClO3 formation, also the ammonium salt could not be precipitated out as easy as the potassium salt.
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[*] posted on 2-3-2009 at 12:54


Hello,

There is a good discription of how to make Perchlorate here
from GarageChemist making Perchlorate from Chlorate using Pt. It appears you don't get any corrosion by taking the Chlorate level low. If you attempt to go from Chloride to Perchlorate you will get erosion. Low levels of Chloride seem to be the Pt anode killer.

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[*] posted on 3-3-2009 at 08:20


Per, don't forget that contamination with even small amounts of chlorate renders the entire batch "chlorate" in a pyrotechnic sense. You may want to test the batch using any number of published methods to confirm if it is, in fact, contaminated.
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[*] posted on 3-3-2009 at 12:16


"I used OTC (well, not in germany, but in France at that time- nowadays very hard to find even there)"

Chlorates in germany can be ragarded as unavailable, so it is my intention to start with chloride which they can´t prohibit, come on, using chlorates would be really too easy.

"prevent reduction of chlorate at the cathode, which would produce chloride and cause anode erosion."

So I assume that even small amounts of chloride are poison for Pt anodes.

But I´ve my doubts that there are any anodes converting chloride into perchlorate without any erosion over a very long time.
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[*] posted on 3-3-2009 at 15:53


Hello Per,

I tried to look for pure Chlorate in France without success. The stuff in the garden centers was all (approx.) 60% Chlorate and expensive. Better to make your own and start from there.

The *"TOTALLY FORBIDDEN ZONE"* for most Anodes when going from Chloride to Perchlorate is [the region where Chloride concentration is getting low before any Perchlorate has started to form, to the region where most of the Chloride is gone and you have a 'Perchlorate cell']. Lead Dioxide seems to be the only candidate, but it does erode too. I think we need to pH controll the cells to stop the Lead Dioxide from being eroded too much.
Everyone has their own definition for what exactly is "eroded too much".
All Anodes erode.
It is a question of how many grams erosion per KG product. How much those few grams cost or contaminate (easy/hard to remove).
It is also a question of keeping conditions in the cell that are more favourable to low erosion.
An never ending subject.

If you are making Perchlorate with Pt you are better off making Chlorate, extracting solid Chlorate, then making Perchlorate. This is of course alot of work and you may prefer to go with the (somewhat high) Pt erosion and save labour.

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[*] posted on 3-3-2009 at 22:12


Quote:
Originally posted by dann2

I tried to look for pure Chlorate in France without success. The stuff in the garden centers was all (approx.) 60% Chlorate and expensive. Better to make your own and start from there.



You didn't look for the right stuff. The pure chlorate continues to be sold to this day, in green 5L canisters, as a 600g/L solution in water.
All the solid chlorate products contain 60% chlorate maximum (rest NaCl), it's required by law.
The solution is amazingly pure, it doesn't even get turbid when adding AgNO3.
I used it directly as feedstock for my perchlorate cell after adding 2g/L potassium dichromate. The Pt anode stays shiny indefinitely.

Per, it's nice that you found iridium to be an anode material that apparently can go directly from chloride to perchlorate- no other material apart from PbO2 seems to be able to do that. But I think you still see the problem with that: you can't produce a pure substance as product. You get an unknown mix of chlorate and perchlorate. And this is a problem because it's very difficult, if not impossible, to actually separate NaClO3 and NaClO4. Their solubilities are both so high that a separation cannot be achieved with reasonable yields.

The only option is to destroy chlorate chemically. Have you ever tried that once? It doesn't seem like it, or you wouldn't be saying things like
Quote:
Originally posted by PerNH4ClO4 is a good rocket propellant but I wouldn´t produce it with self made perchlorate solutions because of the hazards of possible NH4ClO3 formation, also the ammonium salt could not be precipitated out as easy as the potassium salt.

With reducers like acid + bisulfite, it's very easy to make completely chlorate-free perchlorate solutions. Using acid + bisulfite is essential even when using a pure NaClO3 solution to run the perchlorate cell with.
And precipitating NH4ClO4 from a NaClO4 solution works quite well, you just need to use a highly soluble ammonium salt like NH4NO3 as a concentrated solution, and cool to 0°C.
The residual perchlorate in the filtrate can be precipitated with KCl to give KClO4 byproduct.
Although it's possible to use NH4Cl as the ammonium salt, and repeatedly boil down the filtrate, filter the precipitating NaCl while very hot, and cool to crystallize more NH4ClO4, theoretically giving almost complete separation of NaCl and NH4ClO4, it is a lot of work.

Making pure chlorate with your iridium anode is also very difficult. How are you going to tell that there's no perchlorate in your product without performing an accurate analysis?

That's why you use SEPARATE CHLORATE AND PERCHLORATE CELLS.




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[*] posted on 4-3-2009 at 13:26


Hello Folks,

@GC. I will have to look a bit harder if I make it back again. I intend to, hopefully. Never though to look for the stuff in liquid form and was a bit short on time. Is it sold in garden stores? I can purchase it at the local hardware store in liquid form for 18 Euro for 500ml.:o I often felt like asking the shop attendent was it research grade! I have never purchase any of the stuff.

I think Per has not made any Chlorate or Perchlorate with the Ir. He was just asking would it work as he has an offer to purchase some of the stuff in foil form. It will erode quicker that Pt IMO but since it is half the price it may still be a winner. Foil is not a form I would like to purchase Ir (or Pt) for Chlorate or Perchlorate making.

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[*] posted on 7-3-2009 at 09:17


Destroying chlorates with reducers:
I would prefer the thermal decomposition of chlorate into perchlorate for this purpose, it might not be complete but sufficient for the most purposes.

So it seems that Ir could not be the better anode material, thx, I could save the money:)
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[*] posted on 7-3-2009 at 09:37


Chlorates cannot be thermally decomposed in solution!
Decomposition occurs in molten chlorate and KClO3 melts above 300*C.
It is really a disproportionation forming KClO4 and KCl.
Perchlorates have been prepared by this reaction. . .
It's actually a cheaper route for perchlorate, if you can get it to work effectively.
Chlorates by graphite anode, followed by thermal processing. . .
But even small amounts of dust falling into the melt can reduce chlorate to chloride.

[Edited on 7-3-2009 by hissingnoise]
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[*] posted on 12-3-2009 at 04:46


The other problem with thermal decomp is now you have a slab of perchlorate with significant KCl in it, requiring you to redissolve and filter/recrystallize the mass.

The thought of molten chlorate in a toaster oven or similar just sounds a bit dangerous to me, and you have to nail the temperature, or it doesn't work properly.
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[*] posted on 13-3-2009 at 08:46


The temperature must be controlled, otherwise the formed perchlorate would decompose further to pure chloride, a digital thermometer should be satisfactory for controlling the temp.
Dissolving the mass after that in hot water and filtering the chloride off shouldn´t be the problem.

Chlorates by graphite electrolysis, I have graphite anodes but unfortunately they erode so fast that I wasn´t able to prepare even a small amount of chlorate with them, may it´s the wrong graphite, it isn´t very substantial.
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[*] posted on 13-3-2009 at 13:37


It is possible to use acids to prepare perchlorate salts directly from chlorates, this will save time and effort of perchlorate electrolysis, but the yield suffers. Though this can also likely be used to recover more perchlorate from chlorate mixed with perchlorate, instead of destroying it.
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[*] posted on 15-3-2009 at 05:14


I know that from chromates and dichromates but I´ve never heard that chlorates can be converted with acids to perchlorates.
Have you references?
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[*] posted on 15-3-2009 at 12:28


The acid to use is aqueous HNO3 free of nitrogen oxides. I also doubt, but can't confirm this working with chloride-contaminated chlorate. Possibly not since HNO3 and chloride give NOCl, NO2. Don't use H2SO4 as mentioned below if you try this due to the high explosive hazard and lower yield, but I included it for informative and saftey purposes. Again, though yield will suffer (might be better to use for recovery) and you still need to separate KClO4 from KNO3 (not too hard, H2O sol.: 1.8 g/100g vs. 35.7g/100g at 20-25°) if starting from KClO3.

According to Mitscherlich (Pogg. Ann. 25 [1832] 298) who added finely powdered KClO3 into a dish containing conc. H2SO4, portion after portion, using 1 pt. H2SO4 and 1 pt. KClO3, this is warmed a bit. And ClO2, KHSO4 and KClO4 form, the latter two can easily be separated by crystallization. The Pogg. Ann. ref is here. If 50 mL conc. H2SO4 are added slowly up to 2 to 5g KClO3, and if heating here is avoided KClO4 forms without explosion at a yield of 11% (J. Am. Soc. 44[1922]143). Also according to Sérullas (Ann. Chim. Phys. [2] 45 [1830] 272,273) by preparing with conc. H2SO4, usually explosions occur.

Heating KClO3 or NaClO3 with dilute HNO3 on the sand bath forms KClO4 or NaClO4, according to:

8 KClO3 + 6 HNO3 = 2 KClO4 + 6 KNO3 + 3 Cl2 + 3 H2O + 13/2 O2.

As opposed to decomp. with H2SO4, this reaction with aq. HNO3 occurs smoothly since Cl and O come off ununited and not as ClO2; thus no explosion (Penny, Lieb. Ann. 37 [1841] 204; J.pr. Ch. 23 [1841] 296). By multiple evaporation of KClO3 with the regular conc. HNO3 or anhydrous HNO3, KClO4 in a yield of 30% is obtained, but almost none if fuming HNO3 is used, due to the reducing nitrogen oxides (J. Am. Soc. 44[1922]143).

Boiling KClO3 with 85% pure H3PO4 until the yellow color has disappeared gets a yield of 15% KClO4, chromium trioxide boiling with KClO3: yield 11% KClO4. No KClO4 from treating 30% HClO3 with KClO3 (J.Am.Soc. 44[1922]143).

And if you're wondering about HCl, far as I know, HCl will not work. The chlorate might just oxidize it. E.g. addition of HCl onto aq. alkali chlorate is known to yield the following by formation of chloric acid: 2 HClO3 + 2 HCl = 2 ClO2 + Cl2 + 2 H2O (Luther, MacDougall, Z. phys. Ch. 62[1908] 199). Haven't seen any data indicating perchlorate forms.

The J. pr. Ch. ref is here. For those who can't read German below is a translation of the procedure. I didn't modernize the translation much so note the funky equation, "chloric acid" (ClO2), and use of the word atoms as opposed to moles:

With the intention to investigate the action of nitric acid upon potassium chlorate, a certain mass of the salt was mixed in a retort with a measured quantity of acid and then the mixture was heated in a sand-bath. As soon as it became warm, chlorine and oxygen were formed but not as a compound, the potassium chlorate slowly disappeared. The solution was evaporated to dryness, and then the remaining salt was found to be a mixture of potassium perchlorate and nitrate, in a ratio of 3 equivalents of the latter to one of the first. The author expresses the resulting reaction as follows: 4 KClO6 + 3 NO5 = KClO8 and 3 KNO6 + Cl3 + O13.

The action of nitric acid onto potassium chlorate differs than that of sulfuric acid on the same salt. Through nitric acid the salt is decomposed smoothly as chlorine and oxygen form unbound, whereas through sulfuric acid, the gases develop in the condition of a compound, in that they form the dangerous exploding compound, chloric acid. So nitric acid should be used, since with this the operation can be done without a violent detonation, which occurs easily with sulfuric acid.

The action of nitric acid onto sodium chlorate is the same as by potassium chlorate. The liberated chlorine and the oxygen are in the form of a mixture, and every 4 atoms of the salt will get 3 atoms of sodium nitrate and 1 atom of sodium perchlorate. Sodium perchlorate is very soluble and crystallizes in rhomboides, etc. The paper also deals with reaction of iodates and bromates with HNO3.

[Edited on 15-3-2009 by Formatik]
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[*] posted on 21-3-2009 at 07:52


Interesting, not for practical purposes but the mechanism. It seems that the formed ClO2 oxidises the chlorate to form perchlorate, and so breaks down into Cl2 and oxygen.

As strong nitric acid can´t be used because of the reducing effect of the nitrogen oxides, an addition of a small amount of chlorates to this acid may be a good method for preparing nitrogen oxide free nitric acid:)




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[*] posted on 8-4-2009 at 20:15


There's more info on that with detailed results here, and also some other interesting information about potassium perchlorate preparations in general.
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