Sciencemadness Discussion Board

Storage of 35% H2O2

Mateo_swe - 22-7-2021 at 01:00

I recently bought 2L of 35% H2O2.
I currently have them in roomtemp and the plastic PE bottles are getting pressurized by the H2O2 so i have to release the pressure every day.
Do i need to store the bottles in a refridgerator for the H2O2 to stop develop pressure?
I know H2O2 can build up pressure when stored but releasing pressure every day seem way to often.
It say on the label the 35% H2O2 is stabilized.
Can it be that the bottles are not in complete darkness? and how much pressure should 35% H2O2 build up over time during storage?

I have had 35% H2O2 many years ago and that bottle didnt have this problem with built up pressure in the bottle, at least not in this amount that i have in these new bottles.
Any ideas?

teodor - 22-7-2021 at 02:04

Indeed, I have no problems with storing 25% H2O2 at room temperature in normal light conditions but my bottle with 10% H2O2 is building pressure every day after some day when I put some liquid back to the bottle from a beaker. I suspect that trace amount of transitional metals/oxides/salts is the main reason for this effect, not H2O2 concentration/light/temperature.

[Edited on 22-7-2021 by teodor]

[Edited on 22-7-2021 by teodor]

Bedlasky - 22-7-2021 at 02:12

Oxygen pressure is good thing, it prevents decomposition. I store H2O2 in brown glass bottle without issues.

Sulaiman - 22-7-2021 at 02:26

The vapour pressure of hydrogen peroxide is quite low,
so low that a sealed container does not need venting.
As above, I keep my 50% solution pressurised

Maybe something in the container is catalysing decomposition?
Try re-bottling in a clean container?

Mateo_swe - 22-7-2021 at 07:14

Hmm, the PE plastic bottles were brand new, unused and i washed them with water before i put the H2O2 in them.

So i can use a glass bottle to store the H2O2 and the built up pressure would not break the bottle or screw on cork?
The H2O2 was in different plastic bottles when purchased, but on arrival one bottle was looking like a ballon and the other had got a leak in the screw on cork from the pressure buildup.
Can a rough, bumpy shipping by truck make the H2O2 build up pressure like this?
I changed the bottles but the pressure buildup continues, i dont knoiw if i dare put the H2O2 in glass bottles.
It would be best to have it in glass reagent bottles but maybe they crack from the pressure.

karlos³ - 22-7-2021 at 09:18

Is the stuff stabilised?
I measured the density of some 30% H2O2 I bought like... uh... six or seven years ago or so, and according to the density it was(stabilised) still a good bit over 20%.
I think it keeps pretty well in that case.

Fyndium - 22-7-2021 at 10:45

Washed with.. tap or distilled/deionised water? It can make a huge difference. If I've read right, some metals like platinum can cause explosive decomposition, and lesser metals can likely do so at a more protracted rate. Yet still, storing it in a closed bottle, even with little decomposition would build up pressure until the bottle simply bursts. Since some sturdy bottles can withstand up to 10bar of pressure before going off, it can be quite the blast.

Now that I remember, my canister of H2O2 has got a pressure compensated cap, hence it doesn't seemingly build pressure. It was stored in cold over winter, and during hottest summer days the temp raised above 20C and some bubbles were lodged to the interiors of the can.

Panache - 22-7-2021 at 18:19

Storage for such items can be fraught. Repackaging can often worsen the problem. If you are considering repackaging glass cleaned with hot dilute nitric acid plus several hot rinses with validated pure water is your best bet. Afterward cap and store bottle in a mild steel tube. Open only with a socket wrench such that line of sight to closure is impaired. This would be the only safe method (yet still fraught) as you dont know what is causing your decomposition.

Mateo_swe - 22-7-2021 at 22:57

It says on the label:
Deka Fine Chemicals
Hydrogen Peroxide 35%, Pure, Stabilized.
It also says:
Store in room temperature 2-25°C, store in a dry, well ventilated environment, store with bottle tightly closed, store locked up.
Stable under recommended storage conditions.

So it should be stable and not build up pressure like this.
When i changed bottles i used brand new PE plastic bottles rinsed with tap water, there was no water left in bottles but might have been some very small amount.
Since the H2O2 was building up pressure even in the bottles it was shipped in, the problem seem to have been present from the start.
I have no steel cylinders to store glass reagent bottles in.
Can one buy these pressure compensated caps you have on your bottle Fyndium?

woelen - 22-7-2021 at 23:36

I am afraid that something is wrong with your H2O2 and that it has slight contamination with some catalyst. Even trace amounts, which have no effect at all in other reactions, can still lead to decomposition at a fairly high rate. In the past I had 30% H2O2 and this indeed did decompose, but very slowly. I had pressure buildup in weeks or months. After a few days of storage, I did not have any noticeable pressure builduip. Right now, I still have 10% H2O2 and this also is quite stable on storage. I have that bottle for more than two years now and it still is nearly as potent as when I purchased it.

For now, I certainly would NOT store the H2O2 in tightly sealed bottles. Keep the bottles loosely capped and assure that they remain upright in your storage cabinet. The oxygen outgassing is no problem for other nearby bottles and if you have the caps loosely sealed, you will have no noticeable vapors of water or H2O2.

I expect that your H2O2 will have lost its oxidizing properties within months if you have such severe pressure issues. I would convert the H2O2 to something more stable. E.g. with sodium borate or borax you can make sodium perborate, which is easily separated and has infinite shelf life. Or use it up in syntheses before it has decomposed.

unionised - 23-7-2021 at 02:32

Quote: Originally posted by Bedlasky  
Oxygen pressure is good thing, it prevents decomposition.

Bollocks.
At any pressure that you could plausibly contain, the effect will be tiny.


unionised - 23-7-2021 at 02:34

Quote: Originally posted by Mateo_swe  
H

So i can use a glass bottle to store the H2O2 and the built up pressure would not break the bottle or screw on cork?

It certainly will break the stopper or, the bottle if you allow it to build up..

S.C. Wack - 23-7-2021 at 16:21

3% peroxide may be found in dark plastic flip-top bottles. These are convenient containers for the same peroxide which has been concentrated to 35%; this concentrate is not to be expected to be perfectly stable nor is the cap expected to immediately release all pressure. As this can release some 120 volumes of oxygen, a concentrated peroxide which has developed much pressure which has been released many times over the years has actually lost little of its potency.

Sulaiman - 24-7-2021 at 06:22

I bought 1l of 50% hydrogen peroxide 15 months ago and transfered it to a 1l korken (Ikea) bottle.
IMG_20210724_225234.jpg - 9kB
It has been used a few times, and stored in a cupboard outdoors (25-30C night, 35-40C day).
So I just checked it's concentration by density: 47%

Unfortunately I did not test the concentration on arrival so no reference.

There may have been some significant decomposition, and the bottle may have vented (or leaked).
So
I would like to correct my post above to add:

your storage vesel should have some kind of pressure relief mechanism.
................................................................
P.S. When I was measuring the density I noticed a very small stream of tiny bubbles in the volumetric flask,
even though I had washed the flask with detergent, rinsed, dried, KOH, rinsed, HCl, rinsed, and finally a rinse with some of the H2O2.

Clearly H2O2 does not need much to decompose it.

[Edited on 25-7-2021 by Sulaiman]

Chem Science - 24-7-2021 at 10:10

I store my 70% peroxide between -10º and -15ºC, And at some time i decap the bottle to release the pressure, which is very little.
And the pressure is really low, i hear a slightly degassing noise only.
I agree with Sulaiman on the pressure relief mechanism, it's a very nice safeguard, especially in the case of there been other chemicals in the same space.

Mateo_swe - 28-7-2021 at 07:43

Quote: Originally posted by woelen  
I am afraid that something is wrong with your H2O2.
I expect that your H2O2 will have lost its oxidizing properties within months if you have such severe pressure issues.
I would convert the H2O2 to something more stable. E.g. with sodium borate or borax you can make sodium perborate, which is easily separated and has infinite shelf life. Or use it up in syntheses before it has decomposed.


Yes there is certainly something wrong with this H2O2, i still have to release the pressure in both my 2 bottles of 1L 35% H2O2 at least once every day to prevent the plastic bottles to break.
I expect the strength of the H2O2 to decline quite fast so i must do something.
I do have about a kilo of Borax so the best seem to convert it to some Sodium Perborate as this could be useful and stores well.
I will look up the procedure how its done.
Thanks to everyone for the advice and info on my problem.

Mateo_swe - 28-7-2021 at 08:49

I found the SM thread about preparation of sodium perborate and also the Prepchem preparation of sodium perborate.
The prepchem procedure uses 3% H2O2 and apparently it seems to weak for good results, at least it seem so reading the SM thread.
My H2O2 is probably still much better than 3%, i havent measured its density yet.
Do you think this Prepchem procedure will work if i re-calculate the amount of H2O2 needed depending on my current strength of the H2O2?

The Prepchem Sodium Perborate preparation:

Na2B4O7 + 2NaOH + 4H2O2 -> 4NaBO3 + 5H2O

24g of borax and 5g of sodium hydroxide are dissolved in 150 ml of warm water. The solution is cooled to room temperature and add 283 ml of 3% hydrogen peroxide are added slowly. The solution is cooled additionally with stirring by immersing the beaker in ice water and dropping 20 grams of ice into the solution. After a few minutes, fine crystals of sodium perborate begin to separate. The solution is stirred frequently for the next 20 minutes, then the crystals are collected by filtration, washed with two successive portions of 25 ml each of ethyl alcohol and then two successive portions of 25 ml each of ethyl ether. Finally, sodium perborate is dried and preserved in a stoppered bottle. Sodium perborate obtained by described method contains four molecules of water (sodium perborate tetrahydrate), which at 70° C becomes hygroscopic sodium perborate monohydrate.

Some things i wonder about the above prepchem method, can IPA wash be used instead of ethyl alcohol wash?
IPA is anhydrous, or very close to anhydrous and i have a lot of it compared to ethyl alcohol that i only have in anhydrous form and its expensive.
Also is the second ethyl ether wash needed?
I do have diethyl ether, but i rather skip the ether wash if its not needed.

Prepchem Sodium Perborate preparation:
https://www.prepchem.com/synthesis-of-sodium-perborate/

SM Sodium Perborate thread:
http://www.sciencemadness.org/talk/viewthread.php?tid=71240#...

woelen - 29-7-2021 at 00:48

I think that with much higher concentration than 3% your preparation works better. With 3% you get dilute solutions and then it is more difficult to get crystals of sodium perborate.

If you have K2Cr2O7 or K2CrO4 (or CrO3), then you also could try making the interesting compound K3CrO8. I would not make more than 20 grams or so of that, but it is a really interesting prep. This stuff also stores well if you rinse it well after preparation: https://woelen.homescience.net/science/chem/exps/K3CrO8_synt...
With K3CrO8 you can do many interesting experiments, it has energetic properties.

Another interesting peroxo complex is the following: https://woelen.homescience.net/science/chem/exps/Cr_ammine_p...
This one should only be made in small quantities and always keep in mind that this material may ignite unexpectedly (e.g. because of static discharge). But as long as your H2O2 is strong, it is a really interesting prep.

I also made calcium peroxide from concentrated H2O2 in the past (from Ca(OH)2 and conc. H2O2). Yield was low though and I found it not really satisfying. The compound is impure and only somewhat energetic.

If you have vanadium pentoxide or some vanadate, you can make an energetic and interesting peroxovanadate, which also kan be kept around indefinitely: https://woelen.homescience.net/science/chem/exps/KVO4/index....
I still have the KVO4 from that experiment (it's many years old now) and it still is as good as when I made it.

----------------------------------------------------

Rinsing with IPA works well. It works as well as rinsing with ethanol in my experience. A final rinse with ether is very pleasant, because of the quick drying. You could also use acetone for a final rinse, but keep in mind that acetone is more reactive and certain chemicals may react with that. But if IPA or ethanol does not react, then I also expect acetone to be non-reactive, at least with dry crystalline inorganic compounds.


[Edited on 29-7-21 by woelen]

Bedlasky - 29-7-2021 at 01:08

Woelen: Calcium peroxide is the worst metal peroxide for synthesis. It is unstable and goes through filter during filtration. SrO2 and BaO2 are stable and more easy to prepare. I made SrO2 myself and analyse the product. It had very good purity.

woelen - 29-7-2021 at 05:38

I tried making CaO2, simply because I had Ca(OH)2 and CaO, while I did not have the corresponding Sr- and Ba-compounds. Sometimes you just have to use what you have, even if not optimal. But I agree with you, making CaO2 is not a very satisfying thing. For that reason I also did not make a webpage about that experiment.

Sulaiman - 29-7-2021 at 05:54

You could pipe the oxygen generated by your bottle of decomposing peroxide to a gas collector and see if it really is decomposing quickly - or maybe less than you fear.
eg a typical 'gas collection by water displacement' setup.
https://www.google.com/search?q=gas+collection+by+water+disp...
You could then estimate how urgent the problem is.

PS The cost and effort of converting to perborate seems high to me,
'oxy' laundry powder or similar is less expensive (I bought mine as sodium percarbonate)
and I think almost as useful?

[Edited on 29-7-2021 by Sulaiman]

teodor - 29-7-2021 at 07:26

I suspect that gas collection by water setup will not work well for low gas flow. O2 will go easily through water driven by the difference of partial pressure in the H2O2 bottle and in the air.

Once I used decomposing H2O2 to prepare BaO2 hydrate, the procedure is not complex except may be drying.

Bedlasky - 29-7-2021 at 11:12

Woelen: I once did small project to make several metal hydroxides - I made MgO2, CaO2, SrO2 and ZnO2. I mixed soluble metal salt solution, hydrogen peroxide and NaOH to get peroxide. You don't need to isolate metal hydroxide, you can make it in situ. And CaO2 synthesis was the worst from all. Even MgO2 synthesis was easier than this (despite the fact that MgO2 is less stable than CaO2). CaO2 and MgO2 were heavily contaminated with hydroxides. Maybe I revesit this project some time (but just Sr, Ba and Zn peroxides) and try some energetic mixtures with these peroxides. I am also interested in CuO2, but this experiment is better to do in winter, because of instability of CuO2.

[Edited on 29-7-2021 by Bedlasky]

teodor - 29-7-2021 at 11:45

Also there is hydrogen peroxide - urea complex which is probably is the accessible way to protect H2O2 from decomposition and use it as H2O2 later.

karlos³ - 29-7-2021 at 12:09

Quote: Originally posted by teodor  
Also there is hydrogen peroxide - urea complex which is probably is the accessible way to protect H2O2 from decomposition and use it as H2O2 later.

I have way too much urea and I worry about decomposition of my H2O2.

What would you suggest teo?
Do you have some references that seemed to you like they are simple and harmless to do, and how to revert to the free peroxide without distillation, etc?

teodor - 29-7-2021 at 13:27

Well, Karlos, my experience with this complex was only its usage. For long time it was the only one form of concentrated H2O2 available for my home experiments. I was able to buy it quite cheap in any household shop (back in SU). I remember it had a quite good shelf life as for 30% H2O2 and I was able to handle it with my bare hands.
I doubt there is any practical method to remove urea except converting e.g. to BaO2 or some other hydrated metal peroxide. I understand that you don't want urea always when you want H2O2 in your reaction mixture but it is the price of simplicity and harmless life, I believe you understand me here.

Of course there are some threads here about it from men who had to say more, for example this one: http://www.sciencemadness.org/talk/viewthread.php?tid=23477

And if you like to ask me personally: "urea is dissolved in 30% hydrogen peroxide (molar ratio 2:3) at temperatures below 60 °C. upon cooling this solution, hydrogen peroxide - urea precipitates in the form of small platelets" (https://doi.org/10.1021/ja01851a007). Don't credit or trust me, it's from wikipedia.

I am also was always interested in the method to remove urea from H2O2 water solution, by the way, it sounds like a good challenge.
As well as purification of H2O2 by recrystallisation of urea complex.

[Edited on 29-7-2021 by teodor]

Mateo_swe - 30-7-2021 at 08:31

Sodium Perborate could be useful ta have in the future and i dont think its part of any "oxy-action" cleaning compound.
Those usually contain Sodium Percarbonate, at least the ones i have looked at.
So i will try make some Sodium Perborate from the H2O2 before it become useless.

Nice webpage with chemical and physics experiments woelen.
I might try a few of those experiments.