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Fantasma4500
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HClO4, less than 70% conc destillation
as i have just lately succesfully distilled what seems to be at this point +50mL methyl salicylate in full glass setup without any problems, running
it at low temperature and it all going smooth, i cannot help but consider destilling HClO4.. not anhydrous however, as this would be ridiculously
dangerous
im thinking about using KClO4 as the ClO4 source, with 138.5g for 1M of ClO4 it would require 2M of 'H' meaning 1M H2SO4, being 135.15mL 50% H2SO4 --
or 140 mL because sulfuric acid in this reaction is the expendable
what im really unsure about is .. if i just go straight 50% H2SO4 with KClO4 would i get 50% HClO4 or could i potentially pass the azeotrope?! i think
diluting it with 350 mL water giving me about 3M solution at least would be abit too aggressive dilution.. but if it goes wrong on this it does got
wrong
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DJF90
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Congratulations on the methyl salicylate distillation, very satisfying to obtain a pure oil with such a distinct odour. But steady on... theres a list
of things longer than my arm that I'd rather not distill, and perchloric acid is on there.
It forms a negative azeotrope with water, meaning that the anhydrous acid will come over first (so long as theres more perchloric present than the
azeotrope). You can avoid this by using excess water (such that the resulting solution before distillation is <70% HClO4. However, you'll obtain
the perchloric acid as the azeotropic strength regardless of how much water you add initially (after the fore-run of water).
As I mentioned, I'd avoid this unless absolutely necessary (and until you're more experienced...) You could get a pretty decent concentration by using
ion exchange on a sodium perchlorate solution.
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Fantasma4500
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hold on.. im getting some numbers against all odds....
i realized 100 mL H2SO4 going from 50% to 70% (HClO4 taking up the water) would give 17 mL water
so 1M HClO4 with 17 mL water would be 59M 'solution' meaning something bad happening
in order to get 9M solution i divided up the 9 with 59 and got 6.555, this multiplied with 17 mL gave 111,4mL water needed to be added
i didnt consider im using 140 mL H2SO4 here, meaning 17 mL multiplied with 1.4 giving 23,8mL water that can be dragged out from the H2SO4
so about 135 mL water added to 1M KClO4 and 1M 50% H2SO4 should give approx 9M solution of HClO4, which is just below 60% conc.
anyhow, just had to get that written down before i would manage to snap out of it.. you are saying that i will regardless of water i use, if i made 3%
HClO4 solution i would still distill out 70% HClO4..??
wouldnt this be slightly hygroscopic nearly before it would occur weaken itself, taking up moisture from sorroundings weakening its own concentration?
also, if it is that i can only happen to distill out azeotropic borderline HClO4, wouldnt accidents be prevented by distilling this into a flask
containing water already?
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Oscilllator
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An important point here is that although sulfuric acid is a strong acid, that is only the case for its first hydrogen. The second hydrogen behaves
like a weak acid, and so a much more likely equation for your reaction would be:
KClO4 + H2SO4 -> HClO4 + KHSO4
rather than:
2KClO4 + H2SO4 -> 2HClO4 + K2SO4
Therefore you should probably use a 1:1molar ratio of KClO4:H2SO4 rather than a 2:1 ratio.
Just as an example, this is why textbooks recommend using a 1:1Molar ratio of KNO3:H2SO4 when distilling nitric acid rather than a 2:1 ratio
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AvBaeyer
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Antiswat,
Do be VERY careful in trying to distill perchloric acid. Is this really necessary? Perhaps you could treat potassium perchlorate with sulfuric acid,
chill it really well and filter or decant from the poorly soluble potassium bisulfate. The supernatent should be relatively pure and perhaps could be
made quite concentrated without distillation. Is there a real need for pure, distilled perchloric acid?
Just my suggestion. I get really shaky when I hear about perchloric acid and distillation.
AvB
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TheChemiKid
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If you want to go with that route, you should try with barium or calcium perchlorate, barium is ideal, but calcium is cheaper. Use a similar
procedure, just modify it slightly.
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woelen
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Don't you have anything else than potassium perchlorate? The trouble with this salt is that it is only very sparingly soluble, much less so than any
other common potassium salt. Combining it with dilute H2SO4 (even 50% can be considered dilute in this context) gives no HClO4. If you want to make
HClO4 from KClO4 then your only alternative is through (nearly) anhydrous HClO4 which can be prepared from conc. H2SO4 and KClO4.
Distilling HClO4 is very dangerous. I do not need to elaborate on anhydrous HClO4. Distilling the azeotrope, however, also is not easy. The azeotrope
has quite a high boiling point (appr. 200 C) and at that temperature there may also be some decomposition in possibly dangerous by-products (Cl2, O2
and oxides of chlorine).
If you have NaClO4, then you can make decent HClO4 by mixing a hot saturated solution of NaClO4 with excess conc. HCl. NaCl will precipitate, what
remains is a solution of HClO4, HCl and remains of NaCl in water. Boiling this down to 150 C or so drives off virtually all HCl and a lot of water. On
cooling down some more sodium ions may be precipitated as NaClO4. In this way you can get quite pure better than 50% HClO4 with only a little amount
of sodium ions in it in the form of dissolved NaClO4.
I also tried making HClO4 from NH4ClO4 and HNO3, but that experiment was a total failure. NaClO4 plus HCl works quite well, but NH4ClO4 plus HNO3 does
not give any interesting results. I merely boiled off water and azeotropic HNO3 and was left with an impure very concentrated solution of mainly
NH4ClO4 with just a small amount of acid in it. On cooling down this stuff nearly solidified, a wet slurry was formed.
Unfortunately, I consider KClO4 the endpoint on the route of usefulness for the perchlorate ion. The only interesting use of KClO4 is as a solid
oxidizer in pyro-applications. All my soluble waste perchlorate ends up as KClO4. When I have done experiments with perchlorates then I keep that
waste in a special perchlorate solution waste bottle. Once I have a few 100's of ml of perchlorate solution waste, I add a solution of KCl, such that
KClO4 precipitates. In this way I recover the perchlorate and it finds its final destination in KClO4 which I use for nice demo experiments for
visitors and kids.
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Fantasma4500
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i see that its very popular with suggesting calcium and barium perchlorates --- but these are quite hard to get hold of, and it may aswell be somewhat
hard to get to make..
but if its very well diluted down it shouldnt be a problem.. also the solubility of KClO4 is ofcourse higher, as you rise the temperature.. but it
sounds very unrealistic to me that distilling 3% HClO4 would give me a small bit of azeotropic HClO4, but if this does not happen then it would be
very possible to distill weak concentration HClO4..
also, again.. i think asking whether it HAS to be done is murdering the spirit of home chemistry, how many times have people watched a movie and
definitely agreed with themselves that it was an action that HAD to be done?!
as for KClO4's possibilities of reactions.. Ca(OH)2 is barely soluble, but yet it can be boiled with K2CO3 to yield KOH and CaCO3, this relies on
boiling it so that it can react at decent speeds, and if you think about H2SO4 reacting with KClO4, then 0.7g / 100 mL (approx) should be perfectly
possible, especially with heating it
but again.. if it goes to near / past azeotropic concentrations then it would by any means be dangerous, in which HClO4 is not if its below, which
makes it possible to destill without a ultra high end lab
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Metacelsus
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I've done the distillation (under vacuum) of perchloric acid using sodium perchlorate and sulfuric acid, and an excess of water. The acid produced was
then further distilled to the azeotropic concentration, again under vacuum. Both distillations proceeded smoothly.
In case you're wondering, I didn't use the acid myself; I sent it to plante1999 in exchange for other chemicals.
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Oscilllator
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Just an idea: I have some Magnsesium perchlorate and I am planning to prepare perchloric acid from it using oxalic acid. The insoluble calcium oxalate
will precipitate out, leaving behind perchloric acid. My only concern is that oxalic acid is a reducing agent, but if I keep the perchloric acid below
70% concentration the oxidising power of the perchloric acid should be minimal.
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woelen
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I do not think that oxidation of the oxalic acid will be a problem, but the weakness of oxalic acid will be. If some HClO4 has formed and some
magnesium oxalate precipitated, then you get an equilibrium in which oxalic acid exists mainly as free acid:
H(+) + HC2O4(-) <---> H2C2O4 will be mainly to the right
So, you only get a tiny amount of magnesium oxalate and H(+), and a lot of Mg(2+), ClO4(-) and free H2C2O4.
If you have Mg(ClO4)2, then you could try precipitating MgCO3 with Na2CO3 and boil down the resulting solution of NaClO4 and add conc. HCl to that
solution and then boil down to drive off water and HCl. It will be hard though to get a solution of NaClO4 free of Mg(2+). I would add a little excess
Na2CO3 to be sure that all Mg(2+) is precipitated and only Na(+) remains. Some excess CO3(2-) is no real problem, it will be destroyed by the HCl and
the associated Na(+) ions will precipitate.
In the above preparation it is important to make your solution of NaClO4 as concentrated as possible, so if there is quite some water, boil it down.
And it is important to use HCl as concentrated as possible, otherwise no NaCl precipitates. Use excess HCl (e.g. two times stoichiometric amount).
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Oscilllator
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Hmm I understand the problem here, but I do think I will at least try it on a small scale first to confirm it doesn't work, since my way (if it works)
will be much easier. Even then I might just try distillation with sulfuric acid, since Cheddite Cheese has described success.
One other reason I want to isolate relatively pure HClO4 is because I want to make some NH4ClO4, for some small homemade rockets 
[Edited on 21-8-2014 by Oscilllator]
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Fantasma4500
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if NH4ClO4 is the goal, and you have NaClO4, then simply go with NaClO4 + NH4Cl
just boil it together and scoop any crystals out appearing when the solution is boiling hot, when it cools down relatively pure NH4ClO4 will ppt out
forming nice rectangular crystals, or well depending on how quickly it will cool down and probably some other things i dont know the name of
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hissingnoise
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| Quote: | | I do not think that oxidation of the oxalic acid will be a problem, but the weakness of oxalic acid will be. |
IDK, I'd be very wary of organics in HClO4 . . . ?
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woelen
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Organics and HClO4 in dilute form are not a problem. As I have stated often before, dilute HClO4 is not oxidizing at all, even less so than H2SO4.
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hissingnoise
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Yes woelen, but "dilute" can mean anywhere between anhydrous and non-reacting?
And the azeotrope, I'd imagine, is very oxidising, despite its ~30% water content . . .
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woelen
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With dilute in this context, I mean what can be achieved by mixing oxalic acid with KClO4 or NaClO4. But even the azeotrope is remarkable
non-oxidizing. I once heated 70% HClO4 with KI and Na2SO3 to around 100 C. With KI, no iodine was formed, the liquid only became very pale yellow, but
that may also be due to aerial oxidation at the temperature of 100 C or so. With Na2SO3 I only obtained a lot of SO2, which was driven off at the high
temperature. No oxidation of sulfite.
If even fairly strong reductors like I(-) or SO2 are not oxidized by hot 70% HClO4, then I call it non-oxidizing.
Only when 70% HClO4 is heated to its boiling point at 200 C there is the tendency of oxidizing, but even at that temperature the reaction still is
only of moderate speed. Things become different with burning paper, on which 70% HClO4 is dripped. Where the drops are absorbed by the paper, it burns
more fiercely. But then we are talking of temperatures of 500 C or more.
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unionised
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I'd worry about the solution a bit.
But I would worry a whole lot more about the ppt of Mg Oxalate , contaminated with perchlorate.
And I'd worry even more about distilling perchloric acid, it's possible to produce anhydrous perchloric acid -and you don't want to do that.
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woelen
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With oxalic acid and magnesium perchlorate no combinations can be achieved which would make me worry. The dangers of perchlorates really are
exaggerated. Only covalent perchlorates are dangerous, but in no way will you get any covalent perchlorate from oxalic acid and magnesium perchlorate.
I myself made quite a few organic perchlorate salts and all of them are amazingly stable, more so than corresponding nitrates. Only when they are
heated in a flame, they ignite.
I mistreated perchloric acid quite a lot, with reductors added to it (e.g. KI, Zn, Cr, Na2SO3, organic amines) and none of these resulted in reduction
of the perchlorate ion. Really, perchlorate ion is remarkable inert, even in 70% HClO4, as long as the temperature does not become too high. I would
not heat a solid organic perchlorate complex to 100 C in larger quantities than a few 100's of mg, but in aqueous solution I do not worry to boil the
solution. In this way I prepare the solid perchlorates, by boiling solutions of them, so that water is driven off and then letting them cool down so
that crystals of the solid appear.
Mixing 70% HClO4 with pure amines also is not something I do (the simple acid/base reaction also is quite exothermic), but mixing 50% HClO4 with 50%
solutions of the amine in water, carefully dripping one liquid in the other, while stirring, is not a problem for me.
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chornedsnorkack
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Quote: Originally posted by woelen  |
I mistreated perchloric acid quite a lot, with reductors added to it (e.g. KI, Zn, Cr, Na2SO3, organic amines) and none of these resulted in reduction
of the perchlorate ion. |
That´s interesting. So you have never found anything that does reduce cold dilute perchloric acid or neutral perchlorate solutions?
Some substances that have been alleged to reduce perchlorate are Ti(III), V(II), V(III), Cr(II) and Mo(III). What did Cr give for you? Could you
separate, say, Cr(ClO4)2?
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unionised
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Really?
http://en.wikipedia.org/wiki/PEPCON_disaster
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woelen
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I have done the experiment of adding Cr-powder to 50% HClO4. The Cr dissolves with production of a lot of hydrogen gas. I do not remember the color of
the resulting soluition anymore. I doubt whether it was green or dark violet/grey. I think it was dark violet/grey, but I also heated in order to
speed up the reaction and when heating, the liquid turned green, but I am not 100% sure of that anymore. If you wish, I can try it again.
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chornedsnorkack
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Look at obvious similarities and contrasts between nitric and perchloric acids.
Both form a high boiling azeotrope at about 70 % (nitric acid at 68 %, perchloric at 72 %).
Both are significantly unstable against oxidizing oxygen.
But the contrast: nitric acid is kinetically easy to reduce. In dilute solutions metals reduce it to NO, not H2. Nitric acid readily decays to NO2 and
oxygen on standing. Yet this decay does not release much energy in absence of an external reductant, and nitric acid alone is not liable to thermal
runaway explosions.
Whereas perchloric acid is kinetically sluggish oxidant - its dilute solutions are reduced to H2 by metals, unlike nitric acid. Yet the decay of
perchloric acid releases so much energy that concentrated perchloric acid can and will explode by itself.
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Fantasma4500
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but ----------- why would HClO4 at 30% distill off perfectly azeotropic HClO4 from start and then thereafter distill the somehow remaining water over?
and this concentration would it be capable of somehow by mistake going past the azeotrope using low concentration of H2SO4?????
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AvBaeyer
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There is a great little book which I found in my personal library on perchloric acid and its salts which would be a worthwhile read for those so
interested:
Alfred A. Schilt, "Perchloric Acid and Perchlorates," published by G.F. Smith Chemical Company (1979).
Much of the discussion in this thread is dealt with in that book.
As I write this there is one copy available on the ABE books site. It may also be available in various libraries.
AvB
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