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Author: Subject: h202 concentration?
froot
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[*] posted on 20-1-2004 at 06:53


The Germans used hydrazine hydrate and hydrogen peroxide in manned rocket planes, this was their answer to cheaper aerial combat. They would shoot these poor buggers off in these 'planes' to go shoot down allied planes. The rocket engine only lasted a few minutes so he had to make the most of his trip and then try and safely glide the thing back to the field.
The project was cancelled when one crashed and the tank holding the hydrazine burst and half dissolved the poor pilot. Incredible!




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Saerynide
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[*] posted on 20-1-2004 at 10:28


Edit: Off topic, but....

Wouldnt he have died from the crash first? :o

[Edited on 20-1-2004 by Saerynide]
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[*] posted on 20-1-2004 at 15:05
Uses for h2o2 go WAY back.


I seem to remember that the Japanese used a couple mini-subs using h202 as fuel component in the subs that attacked pearl harbor and that the British had used a similarly designed sub (during WW1?) that used peroxide. I saw a layout about it in National Geographic.


Quote:
Originally posted by Saerynide
Edit: Off topic, but....
Wouldnt he have died from the crash first? :o
[Edited on 20-1-2004 by Saerynide]


I wonder how they explained that to the relatives?...(sending the body home in a jug)




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Saerynide
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[*] posted on 21-1-2004 at 00:20


Ewwwww.... =S
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wink.gif posted on 21-1-2004 at 07:26


Another way to concentrate H2O2 is to add an anhydrous salt that forms hydrates readily. For example, calcium chloride. 110 grams will absorb 108 grams of water (on formation of hexahydrate), and I bet it will settle out as (easily filtrable) crystals.
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[*] posted on 30-1-2004 at 01:13


what about freezing ?
it freezes at -35 celisus.
boiling h2o2 will cause degradation
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chloric1
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[*] posted on 30-1-2004 at 19:17


Break out the Dry ice and Acetone freezing mixture!And keep the peroxide out of the acetone!:D

[Edited on 1/31/2004 by chloric1]




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Hermes_Trismegistus
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[*] posted on 30-1-2004 at 19:29
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Quote:
Originally posted by simsalabin
boiling h2o2 will cause degradation


Boil at reduced pressure.




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[*] posted on 31-1-2004 at 02:35


Theoretic: kewls remember to use regular, and not glass filtering paper when filtering dehydrated solutions of 90% H<sub>2</sub>O<sub>2</sub>

But seriously - won't high conc. peroxide attack glass fibre even?




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[*] posted on 31-1-2004 at 14:09


Glass fibre is a no no as filter material. It's a tiny bit basic, so...

If you're going to distill H2O2, you're going to need to clean your glassware with concentrated HNO3. After that with H2SO5 and finally rinse it with diluted H2O2. Then add some phosphoric acid as stabilizer.




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jimwig
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[*] posted on 6-2-2004 at 13:30


Hey there ----mr a-bab, what about putting up that "90% H2O2" regime........
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[*] posted on 24-11-2004 at 21:33


The other day at the library I was looking though some book and actually ran across a reference that said solutions of hydrogen peroxide in ether are more stable then those in water... odd...

Anyway, I just found out some interesting information about the peroxide that I normally use. I buy it from a hair care place and it says on the bottle 40 volume Hydrogen Peroxide, I always assumed that either meant it was near 40% or perhaps it was done same as alcohol and 40 volume was like a proof and it was actually 20% so today I did some checking and found the actual concentration.

10 volume = 3% H2O2
20 volume = 6% H2O2

And etceteras, so, my H2O2 is only around 12%, very disappointing, I guess I might have to concentrate by freezing. Oh well, it's getting cold out. I really have no need for anything above 30% and solutions stronger then that kind of give me the creeps considering I've seen this 12% stuff I normally work with behave very strongly toward what I would consider mild conditions.

Aside from that, it was mentioned that one might use anhydrous salts to sap the water from a hydrogen peroxide solution, wouldn't many salts simply uptake the peroxide in place of water to form the hydrate, I know some salts do, but do they all?

Edit: One other thing, I've heard that barium peroxide can be precipiated from cold water, this seems to be supported by it forming a stable octohydrate and that hydrate having a temperature at which it looses its water of hydration, but does anyone have any experience with this?

[Edited on 11/25/2004 by BromicAcid]




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[*] posted on 25-11-2004 at 00:42


I went to the library last Sunday to look up the ozone via H2SO4 electrolysis refs. Not that I have Pt electrodes, just curious. The only peroxide sp. gr. tables that I had were in 10% increments, so I wrote down what was in one of the refs that I stumbled on while doing so. I think that this was at 16C: 20% = 1.0725, 25% = 1.0918, 30% = 1.1122, 35% = 1.1327, 40% = 1.1536, 45% = 1.1749.

I've made small amounts of the Ba and Na peroxide hydrates from the hydroxides and H2O2. Not sure how well these dry out, at least without P2O5 in the vacuum dessicator. But they do precipitate very nicely out of the H2O2 and are easily isolated quickly. I seem to remember reading somewhere that heating without vacuum gives loss of O as well as water, so I haven't tried that.

One of the books that I scanned and uploaded to axehandle, Oxidations in Organic Chemistry, mentions the preparation of anhydrous H2O2 in Et2O from the 30% peroxide if anyone ever needs such a thing.

I make all of my peroxide now, in not small amounts, from drugstore 3%. I hoard it when it is on sale, 3 pints for $1. I used to just let it evaporate because with the stabilizers and all, if you use big clean (buffed to a polish) glass mixing bowls then dust is not a problem. Even 35% can have a bunch of dust on it and not bubble at all. Yields were high. Metal is another story. But then I found that heating below 70C will concentrate to 45% with little loss of O2.

I've tried the freezing yet I use heating. But that's me.

Don't know if this has been mentioned before, but the volume thing is based on how much O can be released from 1 ml of the peroxide; e.g. 1 ml of 3% H2O2 can release 10 ml of O.

Off topic, but the only other thing that happened at the library Sunday, other than the peroxide refs (and being required to present photo ID in addition to my library card!) was accidentally stumbling onto an interesting Na production lecture demonstration from JCE. Looks like Cyrus was onto something when he mentioned using light bulbs in the unconventional Na thread, sort of, but he got no love. I uploaded it to axehandle as na_from_nano3_and _a_light_bulb.pdf. Someone might want to see exactly how much Na a light bulb could produce if left on long enough. I paid too much for mine to not use it, so it's up to someone else.

[Edited on 25-11-2004 by S.C. Wack]
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[*] posted on 25-11-2004 at 21:34


Ive shaken H2O2 with ether, I think you would have a hard time hurting yourself with the mix, unless there is crystals floarting around in it. Though apart from burning myself with the extract I havnt pursued it further. Oh, I was using 50% H2O2 in ether as well, so more "dangerous".

Another thought, Is pure silica gel inert towards H2O2? Another attempt I made to concentrate the 50% was by pouring it over silica gel containing blue/pink indicator The prills exploded as it catalysed the breakdown of the H2O2.

By soaking the prills in H2SO4, to try and kill the indicator, washing them with water and drying for 1.5hr at 150° there was no more blue colour. These prills only slowly attacked the H2O2. So im thinking the indicator is the problem, but it could also be explained by residual H2SO4 stabalising the mix (which it DOES do with the gel).

So someone may want to try adding H2O2 to pure silica gel (pull it out of a shoebox or something), sodium silicate works to concentrate H2O2, but its soluble. Imagine pouring H2O2 into the top of a tube filled with silica gel and concentrated H2O2 exiting out the bottom, this seems to good to work :(
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[*] posted on 25-11-2004 at 22:32


If you need 100% h2o2 I would recommend making the 90% pure and adding ether potassium superoxide or sodium peroxide to it they would chemical react with the remaining water to make h2o2 the I would let it evaporate in a still however ko2 and na2o2 are expensive and in order to make yourself you have to take pure na or k and burn it with pure oxygen
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[*] posted on 26-11-2004 at 00:09


The indicating silica gel uses cobalt chloride to do its indication, and of course cobalt compounds effectively catalyze H2O2 decomposition. Adding alkali metal peroxides or superoxides to H2O2 to bring it to 100% sounds like a recipe for disaster. H2O2 decomposes more rapidly in alkaline conditions, and of course it wouldn't be pure even if it didn't explode because it would have sodium or potassium compounds left in it.



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[*] posted on 26-11-2004 at 06:17


i heard before something about bubbling NO2 into conc H2O2 to form HNO3...can that happen? or do you need a platinum catalyst?



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[*] posted on 26-11-2004 at 12:24


NO2 dissolves in water giving a mixture of nitric (III) and nitric (V) acids, i.e. HNO2 and HNO3 respectively. H2O2 oxidizes the HNO2 to HNO3 - no need of any catalyst IIRC. As a side note, but wouldn't that be a 'waste' of H2O2. HNO2 is oxidised to HNO3 by bubbling air through the solution.



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[*] posted on 26-11-2004 at 14:08


i see. bubbling air?



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[*] posted on 28-11-2004 at 15:22


Dry air, specifically.
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[*] posted on 28-11-2004 at 15:28


eh...so sorta using pressure or something?



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[*] posted on 28-11-2004 at 19:26


It's just the oxygen in the air oxidizing the nitrous acid thus formed and any nitrogen monoxide to allow more of the nitrogen dioxide entering the solution to be absorbed as nitric acid.



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[*] posted on 29-11-2004 at 11:55


You'd want to bubble air from a common air pump (fish aquarium pump would work fine) through a drying tube containing calcium sulfate or a similar desiccant, into your NO2 contaminated acid. Make sure that the tubing that you're using is acid resistant as well (PTFE).
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[*] posted on 1-12-2004 at 18:27


Stirring a solution of H2O2 during freezing seems to be a good thing in that it will allow for better heat conduction in the solution and result in smaller pieces of ice which could be filtered from the solution and reduce the risk of a solid sheet forming across the top and breaking a beaker.

However by that same process the freezing point will be depressed somewhat due to the motion and if the stirring is too high could an even higher percentage of the peroxide end up in the ice? How about adding a piece of ice once it gets sufficiently low in temperature to function as a seed crystal?

I haven't found answers to these little questions in my little excursion into peroxide concentration and I was hoping someone could help.

Found another hydrogen peroxide FAQ




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[*] posted on 1-12-2004 at 21:54


A rotovap does a fine job of concentrating 30% H2O2 to better than 70%. There is a much bigger diference in the boiling points than there is in the freezing points.
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