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[*] posted on 26-10-2008 at 04:03
Concentrating H2O2


I'm looking for H2O2 in the region of 30% in the UK. I can buy otc at 6% and I've heard that it can be dessicated over H2SO4.

Firstly is anyone willing to share a commercial UK source of 20+% H2O2? I only need 100ml!

Secondly has anyone successfully concentrated H2O2 from 6% to 20+%? What method did you use and what is the test for concentration please?
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[*] posted on 26-10-2008 at 05:29


*edit* found it--> http://www.sciencemadness.org/talk/viewthread.php?tid=1325&a...

looks like i was wrong about the freezing..

check out the readily available chemicals thread just above this one, there are multiple sources in there. If for some reason they dont satisfy your needs try searching through hydroponic suppliers, H2O2 is sold as 'oxyplus' or something similar.

it is possible to purify your 6% solution, but is probably just excessive work. i remember something vague about freezing the peroxide out of low solutions, then freezing out the water or something... im not entirely sure.

dont use ether =P

there was a thread on this, or at least a subject covering this somewhere on these boards, have a wee search.

[Edited on 26-10-2008 by appetsbud]
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[*] posted on 26-10-2008 at 08:43


You can concentrate it with Na2SO4 also, but only up to 20% approximately. Take anhydrous Na2SO4, dissovle in your H2O2 (i dont know weather heat should release or not, but keep watch, cool if necessary). When you'll come to saturation, cool the solution in freezer and filter the crystals of Na2SO4*10H2O. Try to filter the salt as it forms, as many times as it is required, that should probably increase total yield of H2O2 (thats because if you filter only once at temperature minimum, you'll get very few free liquid(but a bit more concentrated), rest will remain on the crystals, that will fill all the solution)

[Edited on 26-10-2008 by Ebao-lu]

[Edited on 26-10-2008 by Ebao-lu]
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[*] posted on 26-10-2008 at 09:24


Does Na2SO4 form a peroxyhydrate like the carbonate (i.e. Electrasol product) does?

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[*] posted on 26-10-2008 at 11:14


I think, even if it forms some peroxohydrates, the content of H2O2 is low there. Otherwise, instead of H2O2 the water would result(sorry, ice), but that is not the case - the concentration of H2O2 in final solution is substantial. Anyway, thank you for this note. I did not analyse the crystals (nor H2O2 conc, maybe it is 15%, and maybe 30+%), and it is worth doing.
I remember only, that for most pyrotechnical purposes that H2O2 was OK (even i had a run-away with hexamethylenetetramine and citric acid)

[Edited on 26-10-2008 by Ebao-lu]
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[*] posted on 26-10-2008 at 12:33


Quote:
Originally posted by Ebao-lu
You can concentrate it with Na2SO4 also, but only up to 20% approximately. Take anhydrous Na2SO4, dissovle in your H2O2 (i dont know weather heat should release or not, but keep watch, cool if necessary). When you'll come to saturation, cool the solution in freezer and filter the crystals of Na2SO4*10H2O. Try to filter the salt as it forms, as many times as it is required, that should probably increase total yield of H2O2 (thats because if you filter only once at temperature minimum, you'll get very few free liquid(but a bit more concentrated), rest will remain on the crystals, that will fill all the solution)

[Edited on 26-10-2008 by Ebao-lu]

[Edited on 26-10-2008 by Ebao-lu]



http://webassign.net/graphics/solubi1.gif
As you can hopefully see by clicking on the link above the solubility of Sodium Sulfate increases as it gets colder, so what are you talking about precipitating it when it gets cold?

Pretty sure your method will not work, can you perhaps site your sources?

Anyways, from my Expirence with 3% H2O2, If you stick it in the freezer until untill it is as cold as the freezer allows it to get, and then drain of the remaining liquid, it is a lot more
concentrated - Enough to bleach your skin completely white. I don't know what percent is possible this method though, but I think it may be as high as 60%+ Of course, getting it to that concentration will happen the first time to freeze the water out of the mixture.
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[*] posted on 26-10-2008 at 13:00


Not to mention the fact that, unless it is analytic grade, sodium sulfate will contain traces of transition metals, which are going to decompose your peroxide.

(Technical sodium carbonate for example, was found to contain enough palladium traces to be able to catalyze suzuki couplings without any other palladium source.)

[Edited on 26-10-2008 by vulture]
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[*] posted on 26-10-2008 at 14:44


Interesting diagram! Perhaps, below zero this line goes down rapidly. I did not measure the temperature while the crystalls began to form. Maybe ice also precipitated, if it can form same type crystals(hardly). But i qotta plenty of solid at the end. I dont have any sources, i've done it myself
As for H2O freezing out, this method is good, but for high concentrations (>10% or such) requires low temperatures, not acheivble in usual freesers(-10C).
As fore traces of metals, small amounts of polyphosphorous acids as stabiliser should be used.
Actually, i used this method to produce H2O2 form Na2CO3*1.5H2O2 and 36%H2SO4 for car battaries. Both reagents are clean enough to make H2O2 and see not a single bubble of O2. After that the solution was immediately used, i didn't store it
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[*] posted on 26-10-2008 at 14:53


Quote:
Originally posted by Ebao-lu
Interesting diagram! Perhaps, below zero this line goes down rapidly. I did not measure the temperature while the crystals began to form. Maybe ice also precipitated, if it can form same type crystals(hardly). But i qot plenty of solid at the end. I don't have any sources, I've done it myself
As for H2O freezing out, this method is good, but for high concentrations (>10% or such) requires low temperatures, not achievable in usual freezers(-10C).
As fore traces of metals, small amounts of polyphosphorous acids as stabilizer should be used.
Actually, i used this method to produce H2O2 form Na2CO3*1.5H2O2 and 36%H2SO4 for car batteries. Both reagents are clean enough to make H2O2 and see not a single bubble of O2. After that the solution was immediately used, i didn't store it


It could also go up rapidly, pretty sure there are trends with salts that are exothermic and endothermic when they dissolve.

you know why you didn't see any O2 bubbles? because all of the CO2 bubbles got in the way! :D

How did you determine your H2O2 was more concentrated then before from your anhydrous Sodium Sulfate method?

I still say it doesn't work, also if it did, you'd still have a ton of Na2SO4 dissolved in the the peroxide solution.

EDIT: Also, Dry Ice and Acetone, or some other mixture that works well could help you get it cold enough to freeze out the rest of the water.


EDIT: Sorry about this, but I found conflicting Information on the solubility of Na2SO4. http://en.wikipedia.org/wiki/Image:Na2SO4_solubility.png

So, which one is right? I guess its solubility actually does have a dramatic decrease when it is at 0*C. Hmm, wonder which graph is wrong?

[Edited on 26-10-2008 by kclo4]

[Edited on 26-10-2008 by kclo4]
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[*] posted on 26-10-2008 at 20:38


Freezing low concentrated solutions doesn't result in either peroxide or water freezing out entierly, other already discussed concentration methods are preferable.

I've frozen 200 mL of 3% in an ordinary freezer and this froze in a few hours, left overnight it was then separated the liquid from the ice and the liquid density was 1.02g/cc which is 8%.

I've also done attempts with 35% H2O2 but this didn't freeze in the freezer after a few days because although below -20ºC and on the coldest setting, not enough to reach the freezing point of -33ºC, which would need a special refrigeration.

It's also been claimed that below 62%, the water will freeze first until it reaches 62%, the other claim is that the peroxide will supercool and that adding a crystal of frozen peroxide will cause it to freeze.
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[*] posted on 26-10-2008 at 21:44


I believe the Wikipedia graph is correct. Sodium sulfate decahydrate has relatively low solubility. Sodium sulfate (anhydrous) is not stable in cooler solution, hence the sudden reversal in the graph. Maybe the first plot is correct of the anhydrous salt, if the decahydrate weren't more stable; but because of this, it would have to be an extrapolation and may not be physically realizable.

I have a table of solubilities which agrees with the latter graph, so it's 2 to 1 now.

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not_important
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[*] posted on 27-10-2008 at 02:21


Wiki has it correct the sodium sulfate solubility curve is famous.

multiple salts solubility at Wiki
http://upload.wikimedia.org/wikipedia/commons/f/f9/Solubilit...

another example
http://www.vias.org/genchem/inorgcomp_sodiumsulfate.html

and one in context of saline lakes, see figure 23, mirabilite being natural decahydrate of sodium sulfate.
http://www.salinesystems.org/content/1/1/10
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[*] posted on 27-10-2008 at 03:53


Evaporate.
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kclo4
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[*] posted on 27-10-2008 at 16:05


Wow that is really helpful Zed. Want to elaborate on that a little more?
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[*] posted on 27-10-2008 at 23:48


Let's do a thought experiment.

H2O2 has a much higher BP than H20. !50C versus !00C.

It should be easy to concentrate H2O2 by evaporation.

A rotovap would be nice. That is, if you happen to have one.

A good vacuum pump would be nice.....You could boil off, aka vacuum distill, the excess water at room temperature.

Lowest tech method? Take 10 squeaky-clean dinner plates. Pour ~50 ml of 6% H202 onto each plate. Turn out the lights and go to sleep.

Depending upon ambient temperature, humidity, and air circulation.... Which will effect the rate of evaporation......You might have 100 ml of ~30% H2O2....When you wake up in the morning.

Or, maybe not. Sounds like an interesting home experiment, that you can perform for about a buck.
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[*] posted on 28-10-2008 at 00:03


Simple distillation or evaporation loses a lot of peroxide along with the water as peroxide concentrations increase. Don't have the charts handy, but I think that above 20% the H2O2 losses increase sharply.

You'll want a tent over those plates, the dust falling into them will not be good for the H2O2.
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[*] posted on 28-10-2008 at 00:14


Indeed. It isn't perfect. But, I just "Googled up" a nice link. As luck would have it, this very topic was discussed on Science Madness, about a year ago....By folks that have more practical experience than I do.

https://sciencemadness.org/talk/viewthread.php?tid=1325&...
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[*] posted on 28-10-2008 at 00:16


You can concentrate 6% H2O2 by vacuum distilling away H2O until you get the volume down to something less than the theoretical volume for the concentration you will be seeking. I'm not sure what the percentage loss will be
due to carryover and decomposition, a drop of phosphoric acid added to the dilute solution at the start will help
with stability during heating and boiling it down. I have done it before and you probably easily get a 40+% concentration this way..stronger if you keep going IIRC. Density measurement will tell you the concentration by referencing a table.

[Edited on 28-10-2008 by Rosco Bodine]
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[*] posted on 29-12-2008 at 02:16


Quote:
Originally posted by Ebao-lu
Interesting diagram! Perhaps, below zero this line goes down rapidly. I did not measure the temperature while the crystalls began to form. Maybe ice also precipitated, if it can form same type crystals(hardly). But i qotta plenty of solid at the end. I dont have any sources, i've done it myself
As for H2O freezing out, this method is good, but for high concentrations (>10% or such) requires low temperatures, not acheivble in usual freesers(-10C).
As fore traces of metals, small amounts of polyphosphorous acids as stabiliser should be used.
Actually, i used this method to produce H2O2 form Na2CO3*1.5H2O2 and 36%H2SO4 for car battaries. Both reagents are clean enough to make H2O2 and see not a single bubble of O2. After that the solution was immediately used, i didn't store it


Mr Ebao-Lu, you need to learn to make the difference between CO2 and O2 bubbles :P.

[Edited on 29-12-2008 by Quantum_Dom]
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[*] posted on 29-12-2008 at 09:26


you too

If you realy want to count bubbles, find another hobby. I mean, in general the process of decomposion was not significant because if there were any H2O2 decomposion catalysts, my mixture would had made a runaway, with heating, and no traces of H2O2 would be there after the gas has evaluated. But my solution was feasible for HMTD. Or you think that the process of CO2 evaluation is a catalyst for H2O2 decomposion, because there are bubbles?
funny logic..


[Edited on 29-12-2008 by Ebao-lu]
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[*] posted on 30-12-2008 at 12:38


As far as I know H2O2 can be concentrated by vacuum distillation with the aid of a good column. Of course this is dangerous, especially with larger amounts and/or higher concentrations.
If ether is too dangerous to extract H2O2 from its dilute solutions, why not try to use DCM instead? DCM is able to extract nearly pure HNO3 from dilute HNO3 solutions. It doesn´t form peroxides, has a low boiling point and so on.
I think one should give it a trial.....
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[*] posted on 30-12-2008 at 17:38


Quote:
Originally posted by grind
As far as I know H2O2 can be concentrated by vacuum distillation with the aid of a good column. Of course this is dangerous, especially with larger amounts and/or higher concentrations.


Boiling and evaporating dilute H2O2 itself at standard atm pressure will concentrate it. Boiling 200mL of the stabilized 3% in a beaker to a volume of about 20mL on a hot plate has gotten me a result of d=1.08g/cc, 22.5%. Enough to turn the skin white. A more careful heating should have gotten a higher concentration, but also takes longer. In Ber. 27 [1894] 3307/12, Wolffenstein got varying concentrations even up to 49.5% from heating 132.2g 3% on a water-bath at 68 mm Hg, or even evaporating 3% in a beaker 64.7% (56% yield). That is also a reason for the fire hazard of allowing dilute H2O2 to evaporate.

Quote:
If ether is too dangerous to extract H2O2 from its dilute solutions, why not try to use DCM instead? DCM is able to extract nearly pure HNO3 from dilute HNO3 solutions. It doesn´t form peroxides, has a low boiling point and so on.
I think one should give it a trial.....


You will likely need massive amounts of DCM also. DCM/H2O2 mixture is also detonable, but probably less dangerous than that ether mixture.
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[*] posted on 30-12-2008 at 18:29


Quote:
Originally posted by grind
As far as I know H2O2 can be concentrated by vacuum distillation with the aid of a good column. ...


You don't distill the H2O2 itself, just remove water. A column helps reduce H2O2 losses, unless your dilute peroxide is very cheap it's worth using. It doesn't need to be that good of a column, at 40 mmHg (a decent aspirator pump) water boils at about 34 C and H2O2 at 77 C and a 40+ degree difference isn't too hard to separate. Concentration to the range of 30 to 50 percent H2O2 is not difficult and is a decent compremise between easy, safety, and bulk/concentration of the product.
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[*] posted on 31-12-2008 at 03:24


Quote:
Originally posted by not_important
You don't distill the H2O2 itself, just remove water.

That was what I meant. H2O2 remains in the residue.
By the way, there is a book "Applications of hydrogen peroxide and derivatives" (from where I got my informations). There you can find detailed descriptions about manufacture, purification, properties and applications of H2O2.
Here is the link:

http://ifile.it/fc8lt5o

Another idea for concentration of dilute solutions: azeotropic water removal...
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[*] posted on 7-1-2009 at 14:31


You can also buy large or small quantities of 50% H2O2 at any hydroponics store for pretty cheap..
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