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Junk_Enginerd
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Can hydrogen peroxide be vacuum distilled straight from sodium percarbonate?
The way I've understood it, it's basically sodium carbonate hydrated by h2o2 instead of h2o. Could one simply distill it straight from the sodium
percarbonate? Either by dissolution in water followed by distilling, or maybe even dry, straight from the salt? Though I think any water in the system
would be desirable considering h2o2's temper...
Yes, I searched, but the threads are a mess and I saw no mention of this.
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Fery
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This is an interesting idea. But H2O has lower b.p. and will distill first with only very small partial fraction of H2O2.
How could be the content in the distillation flask stirred once it reaches too much solid state? Or how to stir the content if distilling directly the
commercial solid Na percarbonate without any additional water - maybe the commercial solid does not require stirring and if distilled slowly it could
work. Maybe the commercial solid melts at some low T in its onw water of crystallization and later during vacuum distillation it could solidify into 1
solid block if stirring not applied?
Also be very careful, once you overheat too much concentrated H2O2 it will decompose with evolution of significant heat which will autocatalytically
speed up the decomposition. Few years ago I used only 30% H2O2 to oxidize oleic acid with H2WO4 catalyst and it was hard to keep the T exactly at 100
C - once it reached about 103 C the temperature was quickly raising (autocatalytic decomposition) and I had to cool down the flask, once it fell below
100 C the T was falling down spontaneously due to insufficient reaction rate and required heating the flask to 100 C. I had to keep it at 100 C for
few hours, IIRC like half a day. It required a lot of manual interventions (almost permanent during whole time).
Do such experiments at small scale first, like 10 g of Na percarbonate and wear a face shield + protective clothes (anti shattered glass proof) + safe
glass on fumehood etc.
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Fery
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2nd idea - couldn't be the carbonate first decomposed with some weak acid (like citric, tartaric, acetic)? Some time ago I tried to obtain H2O2 from
Na percarbonate using NaHSO4 but the solubility of Na percarb. in H2O was quite low at low T (and it must be cooled to very low T during
neutralization) so I dissolved only about half of it in less water and the addition of NaHSO4 seemed to decompose most of H2O2 (very likely due to
local overheating) because I added commercial NaHSO4.xH2O in solid form not to introduce even more water into the reaction.
Unfortunately also these organic acid Na salts will very likely form hydrates (Na acetate forms trihydrate IIRC, I do not know what citrate either
tartarate). And citric acid could perhaps react with H2O2 as it has -OH group, no matter tertiary alcoholic group the H2O2 could maybe split the C-C
bonds in the acid.
Maybe you could have success with NaHSO4 when added not as solid but as conc. solution (but again very cold solution).
Just don't forget to cool the reaction in snow bath (I used snow which is much better that crushed ice but although almost all the H2O2 decomposed and
only tiny amount remained when added NaHSO4 in solid form - I tested it finally with tiny KMnO4 crystals which lost colors but after about 5th crystal
there was not decolorization anymore so really negligible concentration of resulting H2O2). Also my NaHSO4.xH2O was technical (swimming pool grade) so
maybe traces of Fe3+ catalyzed decomposition of H2O2 and maybe some chelaton could bind heavy metals and inhibit decomposition
(ethylenediaminetetraacetic acid or at least maybe gelatin like in hydrazine sythesis but gelatin could again react with H2O2...).
Anyway resulting mixture will be H2O2 with some fuel (Na salt of organic acid) so again very dangerous experiment and certainly strong vacuum during
distillation and heating to very low temperature... Perhaps better to experimentate with inorganic like NaHSO4 or better H2SO4. But H2SO4 dripping
would cause local overheating so again needs dilution (maybe H2SO4 concentration not more than 70% ???).
With H2SO4 + Na percarb. you maybe produce only Na persulfate which could be easily bought as it is used for etching PCBs in electronics.
What will happen when mixing Na percarb + Na persulf???
Just some untested ideas from my head... Be careful.
The goal is clear: to obtain H2O2 concentrated enough. As far the easiest way looks to be buying 10% H2O2 which is legal and vacuum concentrate it by
distilling out some of the water.
[Edited on 19-1-2024 by Fery]
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unionised
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I'm not really advocating this but...
If I got some "percarbonate" solution and neutralised it with HCl I'd have H2O2 in salty water.
Cool it and concentrate it and the salt would crystallise out.
I could distil that mixture to get H2O2
No acetate/ citrate whatever to act as a reducing agent.
And given that NaCl crystallises without water of hydration , I think it's unlikely to form a solvate with H2O2
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Junk_Enginerd
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Quote: Originally posted by Fery  | This is an interesting idea. But H2O has lower b.p. and will distill first with only very small partial fraction of H2O2.
¨~~
Do such experiments at small scale first, like 10 g of Na percarbonate and wear a face shield + protective clothes (anti shattered glass proof) + safe
glass on fumehood etc. |
Right, so it sort of becomes a dry distillation anyway, because you'll have to mostly remove the water before you get even get any peroxide out...
Hmm. Maybe if you add water and also a drying agent? Either something like calcium chloride, or maybe some zeolite/molecular sieve could be helpful?
Not sure if something like calcium chloride would have a greater affinity for H2O vs H2O2.
Yes, I'm aware that heating it past ~70°C or so is inadvisable, I reckon vacuum is necessary for that reason.
Oh absolutely. I'm not planning to try it at all unless there's a route that at least seems probable to work anyway.
| Quote: Originally posted by Fery | 2nd idea - couldn't be the carbonate first decomposed with some weak acid
The goal is clear: to obtain H2O2 concentrated enough. As far the easiest way looks to be buying 10% H2O2 which is legal and vacuum concentrate it by
distilling out some of the water.
[Edited on 19-1-2024 by Fery] |
From what I understand, adding any acid to sodium percarbonate generally results in instant decomposition of both carbonate and peroxide, with O2
being generated rather than H2O2. I reckon it works with sulfuric acid, perhaps because it is also a powerful dehydrating agent? I'd like a route that
doesn't involve sulfuric acid though, since that's quite a premium chemical for me. I've got 1 L conc. sulfuric acid, and that's all I think I'll ever
have since I bought it just before it was thoroughly banned in EU.
Getting a hold of even 3% is a little difficult. Maybe some specialty webshop has 10%... It's not easy to buy in Sweden at least. It's also often
contaminated with detergents and whatnot, not to mention expensive.
In the end I'm doing it mostly as a challenge, I don't have any specific need for H2O2 at the moment.
| Quote: Originally posted by unionised | I'm not really advocating this but...
If I got some "percarbonate" solution and neutralised it with HCl I'd have H2O2 in salty water.
Cool it and concentrate it and the salt would crystallise out.
I could distil that mixture to get H2O2
No acetate/ citrate whatever to act as a reducing agent.
And given that NaCl crystallises without water of hydration , I think it's unlikely to form a solvate with H2O2
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My understanding is that adding acid to sodium percarbonate decomposes both the carbonate and the peroxide, generating CO2, O2, and H2O. I could be
wrong though, I'm not sure of it.
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Sir_Gawain
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Also, highly concentrated hydrogen peroxide oxidizes HCL to chlorine.
“Alchemy is trying to turn things yellow; chemistry is trying to avoid things turning yellow.” -Tom deP.
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bnull
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If it's just for the sake of doing it, you could try to leach the peroxide from the percarbonate with a solvent in which sodium carbonate is
insoluble. Ethanol, for example. The ordinary 96% ethanol is good enough, as the water seems to help the extraction.
Quod scripsi, scripsi.
B. N. Ull
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unionised
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Interestingly, if you neutralise the HCl with Na2CO3 it's no longer there.
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Sir_Gawain
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Hydrogen peroxide is miscible with diethyl ether while water is not. You'll be able to obtain very concentrated hydrogen peroxide, at only the cost of
both arms and left eye.
[Edited on 1-19-2024 by Sir_Gawain]
“Alchemy is trying to turn things yellow; chemistry is trying to avoid things turning yellow.” -Tom deP.
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unionised
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| Quote: Originally posted by bnull | | If it's just for the sake of doing it, you could try to leach the peroxide from the percarbonate with a solvent in which sodium carbonate is
insoluble. Ethanol, for example. The ordinary 96% ethanol is good enough, as the water seems to help the extraction. |
What do you propose to do with a solution of H2O2 in ethanol?
https://www.sciencedirect.com/science/article/abs/pii/S03043...
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bnull
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Fly some rockets to London, perhaps?
Now, seriously. Just ignore my suggestion. No one likes homemade T-Stoff. First, you can't concentrate the solution to 30% or more, or it blows up.
Second, you can't vacuum-distill it, or else the concentration rises and then it blows up (I've checked my Bretherick's: he's right). If you use a
different solvent, you'll get either a peroxide (THF, ether, acetone) or phosgene (chloroform) or flames (about everything else).
Extraction of H2O2 from the percarbonate, whatever the means, is not worth the risk and effort. You can still use the
percarbonate as an in situ source of peroxide. At least it is stable.
Quod scripsi, scripsi.
B. N. Ull
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clearly_not_atara
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You could probably get some decent results with a mixture of water and urea (20-30% urea) at < 0 C. The urea will complex H2O2 pretty well and the
solubility of sodium carbonate should be pretty low in this mixture.
[Edited on 04-20-1969 by clearly_not_atara]
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Tsjerk
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Or use sulfuric acid and don't worry about what happens first. Or does sulfate also react in an irreversible manner?
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clearly_not_atara
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^ you could also use NaH2PO4 so that the sodium carbonate becomes bicarbonate instead of making lots of bubbles
[Edited on 04-20-1969 by clearly_not_atara]
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unionised
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| Quote: Originally posted by Tsjerk |
Or use sulfuric acid and don't worry about what happens first. Or does sulfate also react in an irreversible manner? |
Sodium sulphate forms a set of hydrates.
I suspect it would also form solvates with H2O2.
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Junk_Enginerd
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Oh I will not be mixing it with anything that can be regarded as fuel lol. At least not in the process of synthesising it. Maybe
afterward...
| Quote: Originally posted by Tsjerk |
Or use sulfuric acid and don't worry about what happens first. Or does sulfate also react in an irreversible manner? |
Sulfuric acid works, that's a standard procedure detailed in the smwiki. I'm just looking for an alternative since sulfuric acid is quite precious to
me, and I don't want to spend it unless there's no alternative. I have a bit that I bought before it became super regulated in the EU...
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bnull
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| Quote: Originally posted by Junk_Enginerd | | Oh I will not be mixing it with anything that can be regarded as fuel lol. At least not in the process of synthesising it. Maybe
afterward... |
| Quote: Originally posted by Junk_Enginerd | | I'm just looking for an alternative since sulfuric acid is quite precious to me, and I don't want to spend it unless there's no alternative. I have a
bit that I bought before it became super regulated in the EU... |
What about phosphoric acid? It's relatively weak, nonvolatile, noncombustible. They even use it as stabilizer for the
H2O2. I suppose it is still sold in Sweden without much hassle. I'm not that familiarized with the EU regulations.
By the way, the urea-peroxide is a high explosive. It's rather insensitive, or so it seems.
Quod scripsi, scripsi.
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Tsjerk
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At least it does in the presence of NaCl, the 4Na2SO4.2H2O2.NaCl also seems to be quite stable, up to 180 degrees. Couldn't find anything about a
phosphate adduct.
[Edited on 22-1-2024 by Tsjerk]
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clearly_not_atara
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It's not really a concern whether complexes form. The existence of the complex is not the major barrier to utilization of carbonate peroxide. The
problem is that alkaline solutions catalyze the decomposition of hydrogen peroxide. So it is not an issue if sulfate or whatever forms a complex. It
may even be more convenient that way.
But the vigorous bubbling caused by the acidification of carbonate may also promote the decomposition process. This is why I suggested NaH2PO4.
Urea peroxide is shock-insensitive and stable below 60 C. It cannot be distilled at atmospheric pressure but it is considered safe enough in solution
to be used in eardrops, even though the ear is extremely sensitive to changes in pressure. At any rate it is probably more useful than percarbonate.
[Edited on 04-20-1969 by clearly_not_atara]
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teodor
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I believe the answer to the original question is dependent on one very important number: the vapour pressure of H2O2 of the percarbonate crystalls.
But I think it is not "hydradet" by H2O2, the crystall doesn't contain H2O2 blocks, the structure formula is not Na2CO3 * nH2O2, that is only
equivalent formula, in reality it has CO4 groups.
[Edited on 23-1-2024 by teodor]
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bnull
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| Quote: Originally posted by teodor | I believe the answer to the original question is dependent on one very important number: the vapour pressure of H2O2 of the percarbonate crystalls.
But I think it is not "hydradet" by H2O2, the crystall doesn't contain H2O2 blocks, the structure formula is not Na2CO3 * nH2O2, that is only
equivalent formula, in reality it has CO4 groups.
[Edited on 23-1-2024 by teodor] |
According to the ILO, the vapor pressure is negligible and the formula is 2Na2CO3*3H2O2.
One idea (that will not blow you up into tiny bits): "neutralize" the percarbonate with NaH2PO4 as clearly_not_atara suggested,
then cool it in the fridge. You can remove most of the of the salts (sodium triphosphate solubility goes from about 30 g/L at 20°C to 5 g/L at 0°C).
There's an interesting old thread here in the forum where someone was trying to concentrate 3% peroxide by fractional feezing. There is also a paper about phase equilibrium of hydrogen peroxide-water solutions, with a plot of freezing point as function of concentration.
[Edited on 23-1-2024 by bnull]
Note: phosphates are very alkaline and they form complexes with peroxide (patent CA2216061C). Maybe NaHSO4 is a better choice, with just a little excess to lower the pH.
[Edited on 23-1-2024 by bnull]
Quod scripsi, scripsi.
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Tsjerk
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| Quote: Originally posted by bnull |
You can remove most of the of the salts (sodium triphosphate solubility goes from about 30 g/L at 20°C to 5 g/L at 0°C).
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Disodium hydrogen phosphate, as the goal is to keep the CO2 as the bicarbonate.
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bnull
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| Quote: Originally posted by Tsjerk | | Quote: Originally posted by bnull |
You can remove most of the of the salts (sodium triphosphate solubility goes from about 30 g/L at 20°C to 5 g/L at 0°C).
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Disodium hydrogen phosphate, as the goal is to keep the CO2 as the bicarbonate. |
Oops. Let me rephrase it:
You can remove most of the of the salts (disodium hydrogen phosphate solubility goes from about 7 g/L at 20°C to 1 g/L at 0°C).
Nevermind. It is less soluble anyway. Thanks for pointing the mistake.
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teodor
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But making a solution will spoil the whole process which will now be the same as concentrating H2O2 solution by distillation. If percarbonate contains
traces of transitional metals (and it likely contains Fe) the distillation is hard to do without decomposition because it requires not only solution
but also glass be completely free from transitional metal ions.
So, the interesting possibility mentioned in the original question, to get H2O2 vapours without mixing it in solution with impurities is a less
explored way and I think this possibility deserves some attention.
And again, I don't thing that n * Na2CO3 * m H2O2 is a structural formula, it is equivalent formula, it looks like all those O atoms are connected to
C atoms according to the crystall structure.
[Edited on 23-1-2024 by teodor]
Look here: https://en.wikipedia.org/wiki/Sodium_percarbonate#/media/Fil...
All O atoms are bound to C. It doesn't contain any H2O2, it just produce H2O2 "in-situ" when it is dissolving in water.
[Edited on 23-1-2024 by teodor]
But, here H2O2 is captured as a molecule:
https://en.wikipedia.org/wiki/Hydrogen_peroxide_-_urea
According to this article in the range 70-90C it is a liquid, solution of urea in pure H2O2 at atm pressure. So, just apply some vacuum ...
[Edited on 23-1-2024 by teodor]
Also, trying di-ethyl ether as a solvent for H2O2 is not a good idea - the oxidation product, if formed, is very explosive.
[Edited on 23-1-2024 by teodor]
[Edited on 23-1-2024 by teodor]
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bnull
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| Quote: Originally posted by teodor | And again, I don't thing that n * Na2CO3 * m H2O2 is a structural formula, it is equivalent formula, it looks like all those O atoms are connected to
C atoms according to the crystall structure.
[Edited on 23-1-2024 by teodor]
Look here: https://en.wikipedia.org/wiki/Sodium_percarbonate#/media/Fil...
All O atoms are bound to C. It doesn't contain any H2O2, it just produce H2O2 "in-situ" when it is dissolving in water.
[Edited on 23-1-2024 by teodor] |
"The structure of PBS [sodium perborate] corresponds to a six membered heterocyclic dianion. PCS [sodium percarbonate], on the other hand, is
not a true percarbonate, but is a perhydrate (Na2CO3*1.5H2O), as shown in Figure 2.9. [...] Its rather
confusing name has arisen because of historical uncertainties over its structure and PCS should be considered as a solid form of hydrogen peroxide
like UHP [urea-hydrogen peroxide]." Craig W. Jones. Applications of Hydrogen Peroxide and Derivatives, pp. 41-42
"Apart from sodium and carbonate ions, the structure determination has shown the presence of hydrogen peroxide molecules which are
hydrogen-bonded to the carbonate ions, and therefore the compound is not a true 'percarbonate.' " Corrondo et al. X-Ray crystal structure of the industrial bleaching agent ‘sodium percarbonate’(sodium carbonate–hydrogen peroxide (2/3))
As for the Wikipedia image, it's only a (wrong) model for the crystalline structure at 100 K. The right one is described in Acta Cryst. (2003). B59, 596-605.
Diethyl ether was suggested by Sir_Gawain to illustrate the dangers of using organic solvents to leach the peroxide. I still don't know why I
suggested ethanol.
From the ILO-WHO Internation Chemical Safety Cards for sodium percarbonate: "If the temperature exceeds 50°C a self-accelerating
decomposition can occur, releasing heat, oxygen and steam."
Urea-hydrogen peroxide decomposes at the melting point. I don't know the rate of decomposition (there's not much research about the topic) but it may
account for the 15°C range of the melting point. And from the same Wikipedia article, "(b)ecause of the tendency for thermal decomposition, which
accelerates at temperatures above 82 °C, it should not be heated above 60 °C, particularly in pure form."
Solution seems to be the logical path, given that PCS has insignificant vapor pressure, decomposes on heating, and sulfuric acid is almost a luxury
item in Sweden. He may dissolve the PCS in the least amount of water, add the phosphate (which will also precipitate any iron impurities), decant the
liquid, and vacuum-distill it. If he is really worried about transition metals, he can add a pinch of sodium silicate in the water before making the
solution. Of course, there's always dust in the air to screw it all.
[Edited on 23-1-2024 by bnull]
Quod scripsi, scripsi.
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