Suppose a procedure calls upon the use of 30% (aq) H2O2 and all one has access to is a dilute 6% solution.
Is it possible to distill the water off until the desired concentration is reached?
If neccessary, either vacuum distillation will be attempted, or use of a rotary evaporator, i.e. the excess water will be removed in vacuo.
I see no reason why this wont work, other than it says H2O2 should be stored in a refrigerator as it gradually decomposes over time when stored at
room temperature. Hence, this implies that its decomposition is catalyzed by heat but that is not saying that it cant be heated for the brief interval
required to increase the concetration of peroxide appreciably.
30/6 = 5, so whatever amount the procedure requires one is simply going to use 5 times that amount and distill off the water until the desired
concentration is attained.
Before people say, UTFSE, ive already tried that but could not see an answer.
I think the London bombers who failed to set-off their ruck-sack bombs got some H2O2 from the barbers which they boiled on the stove until they got
the required concentration (no vacuum). However, they failed to get the mixture concentrated enough for use in explosives.
Obviously, the 30% here is for organic chemistry and not energetic materials. However, i'd still be interested to know if there are any potential
risks with something like this? Clearly I wont be adding acetone to the flask (why would I intentionally do something like that?)
Okay, thanks for reading this.
Duke - 10-3-2008 at 08:21
I'm pretty sure a mild heating of the solution will bring it close to 30%. I think the steam turns white when you reach that point or close to it.ScienceGeek - 10-3-2008 at 08:40
Why would the steam turn white? I couldn't find any information on whether Hydrogen Peroxide forms an azeotrope with water (which it probably does),
but there's no reason the steam would turn white...?MagicJigPipe - 10-3-2008 at 09:58
I don't think it forms an azeotrope. Also, I think that higher concentrations of H2O2 are more stabile and less prone to decomposition. If you are
in the US, ~25-30% H2O2 can be purchased as "Klean-Strip Wood Bleach" (I even see it being sold on Ebay in gallon size jugs of both solutions). It
usually comes in a box with 2 237ml bottles (1 is 45-50% nearly saturated NaOH solution).
H2O2 forms a 63% azeotrope with water and it can be concentrated to 63% by freezing the water out. I'm very skeptical.
[Edited on 10-3-2008 by MagicJigPipe]
ScienceGeek - 10-3-2008 at 10:48
Quote:
Originally posted by MagicJigPipe
Also, I think that higher concentrations of H2O2 are more stabile and less prone to decomposition.
Could you please elaborate?
Methods for Concentration of Hydrogen Peroxide To Obtain It in Anhydrous Form - Russian Journal of Applied Chemistry
msp2 - 10-3-2008 at 13:04
Abstract: Methods for concentrating hydrogen peroxide and obtaining this compound in anhydrous form, based on removal or binding of water from the
surface of solution in an open vessel or in a closed volume and also on binding of water directly in a H2O2 solution, are described.
I think someone has mentioned this technique before on this forum IIRC.
At least this answers Drunkguy's question. MagicJigPipe - 10-3-2008 at 18:03
Yes, this is one of the most imformative articles I have ever obtained from this board.
Also, I can't remember where I heard it but I'm pretty sure that H2O2 is subject to faster decomposition at lower concentrations. However, according
to the article I just read, the amount of catalytic material it takes to decompose H2O2 decreases with higher concentrations. It's a catch-22. I
think as long as you keep the higher concentrations in a closed, clean container then they should keep longer then severly diluted solutions.crazyboy - 10-3-2008 at 18:36
Yes this is totally possible I have done it. I take a bottle of 3% H2O2 from the pharmacy (about 930ml) then I boil it down in an old pot at 85
Celsius until it is reduced to about 93ml although hydrogen peroxide may form an azetrope well above 30% concentration it must be evaporated or
concentrated by another method to get it much past 30% because it will begin to decompose shortly thereafter. I believe I had a chart with temp
corresponding to concentration in regards to hydrogen peroxide.
I can vouch for this method but not the exact concentration however I did a HMTD synth and it performed perfectly maybe someone can figure it out by
some density/molar weight theorems.leu - 10-3-2008 at 18:47
That very informative article has been posted already:
Posts: 1874
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posted on 7-11-2007 at 08:47 PM
Methods for Concentration of Hydrogen Peroxide To Obtain It in Anhydrous Form
K. V. Titova, V. P. Nikol’skaya, V. V. Buyanov, and I. P. Suprun
Russian Journal of Applied Chemistry, Vol. 75, No. 12, pp. 1903-1906., 2002
Abstract-
Methods for concentrating hydrogen peroxide and obtaining this compound in anhydrous form, based on removal or binding of water from the surface of
solution in an open vessel or in a closed volume and also on binding of water directly in a H2O2 solution, are described.
Attachment: Methods for Concentration of Hydrogen Peroxide To Obtain It in Anhydrous Form .pdf (38.21 KiB)
This file has been downloaded 36 times
Despite the claim in the first post in this thread; if one had been truly diligent about using the search engine, this post would have been found S.C. Wack - 11-3-2008 at 03:45
If only there was a thread with a title like H2O2 concentration, which this may be merged into at some point. With the unstoppable flow of (mostly
lame) questions (and no shortage of a certain 3 or 4 people who answer them all anyways), you really think that this never came up before?
What is the nomograph for?
MagicJigPipe is in error in thinking that the less concentrated peroxide is more prone to decompose. The Russians were a long ways away from being
first to quantitatively study various methods of concentration. They dismiss previous authors as "These studies are of only historical interest", yet
do not show them to be in error. But the article shows that my comments here and elsewhere on the subject are non-bullshit so oh well.
APC may not be the best place for finding science facts.MagicJigPipe - 11-3-2008 at 22:07
Okay, I want to read up on the phenomena of H2O2 becoming less stable (to non-catalyzed decomposition) at higher concentrations. Could you direct me
to the reference where you acquired that information because I can't find anything (aside from Wikipedia) that discusses it.
Apparently it would have to decompose slower than 1% per year (the fact that a percentage is used seems to support what you say) to be more stabile
than commercial, concentrated solutions.
The only thing I have been able to find is a line in the Wikipedia "Hydrogen peroxide" page with no apparent citation, reading:
" Hydrogen peroxide available at drug stores is three percent solution. In such small concentrations, it is less stable, and decomposes faster. It is
usually stabilized with acetanilide, a substance that has toxic side effects in significant amounts."
Of course, I am skeptical of anything on Wikipedia.S.C. Wack - 12-3-2008 at 02:46
My reference is only my own observations. For sure I was never as careful as I could have been, nor am I talking distillation. Just concentration by
low heat and 3% USP peroxide. I admit that this is rather flimsy scientific evidence to say that you are in error and I apologise for wording this as
though it was a fact, or indeed for contradicting you when I have only vague half-forgotten, half-ass data from some hack, especially when all
calculations were based on density only.
Two quotes from Kirk-Othmer: "The stability of pure hydrogen peroxide solutions increases with increasing concentration", but then in the next
paragraph - "The decomposition is zero-order with respect to hydrogen peroxide concentration". Both of these contradict what I said, but - how is
stability affected by concentration, but the rate of auto-decomposition not?
