Sciencemadness Discussion Board

Removing trace impurities from H2O2

VeritasC&E - 1-9-2020 at 11:44


Hello,

I have some H2O2 at 9% concentration and I'd like to use it in small quantities for applications with puriss. reagents I would not like to contaminate.

How would you go about purifying it (in the sense of removing impurities other than water and H2O2 down to maybe <1ppm for any single impurity)?

I'm predominantly concerned about traces of metal ions and maybe stabilizer additives from the pre-dilution H2O2 solution it was probably made from.

I looked it up in a lab chemicals purification book but the process there was for concentrating H2O2, not riding it from impurities.

morganbw - 1-9-2020 at 11:53

I do not want to come off as being rude but sometimes if you need the good stuff you need to buy. Of course you may figure it out, I am just pointing out my thought process.

Metacelsus - 1-9-2020 at 12:18

What are the impurities? An ion-exchange resin might work.

VeritasC&E - 1-9-2020 at 12:20

Is there is no reasonable way to purify it and is it just better to produce it pure from scratch? Or do you mean in terms of time saving?

I like to purify stuff myself when possible. A point here is that the price per liter literally varies from ~$2 to ~$20 going between probably fairly clean H2O2 solution and puriss. H2O2. I can afford to do chemistry only at fairly priced consumables.

VeritasC&E - 1-9-2020 at 12:25

I'm not sure what they are. I know sometimes stabilizers are added to the more concentrated H2O2 it might have been made from, and other likely impurities would be ions from the water (say it's made from RO water for economical reasons), and maybe from the apparatus that was used to produce it (Fe, Ni, Cr for instance if stainless steel was used).

Do you know which stabilizers are commonly used in commercial H2O2 solutions?

DraconicAcid - 1-9-2020 at 12:27

I suspect hydrogen peroxide would be too prone to decomposition during purification to bother trying.

VeritasC&E - 1-9-2020 at 12:30

So we just have to rely on the pricing of big companies if we absolutely want it pure?


[Edited on 2-9-2020 by VeritasC&E]

unionised - 1-9-2020 at 23:55

I can't see any theoretical objections to sub distillation at home. as long as you have the kit.

VeritasC&E - 2-9-2020 at 01:42

Quote: Originally posted by unionised  
I can't see any theoretical objections to sub distillation at home. as long as you have the kit.


Should I vacuum distill the solution below boiling point with the goal in mind to get a much purer (albeit initially less concentrated) H2O2 solution in the receiver? And in a second step boil away enough H2O to reconcentrate the purified solution to the initial 9%?

Thank you for your contribution!

teodor - 2-9-2020 at 02:14

As far as I remember from this video - https://www.youtube.com/watch?v=nVAe__ToAOY
you can distil without vacuum H2O2 as 25-30% mixture with water. I didn't try it by myself and if I would do I would take a lot of precautions, there are some safety issues especially if you have little experience in this.

[Edited on 2-9-2020 by teodor]

unionised - 2-9-2020 at 02:15

I think you need to start by looking up "sub distillation" i.e. "sub boiling distillation".

zed - 2-9-2020 at 03:37

H2O2 is easy to concentrate. Problem is, it usually contains stabilizers that are easily concentrated too.

Easily concentrated, and hard to eliminate, when present in your final product. In years past, acetanilide was a common stabilizer. Might still be. I don't know. But, in some applications it presents a real problem.

The solution to this problem, is to start with H2O2 that doesn't contain stabilizing agents, that are gonna mess things up.

Oh yeah, and apparently H2O2 may be able to "etch" ordinary glasses. Thereby contaminating itself. Grrrr.

Gotta find a way to obtain "better" peroxide, if your work is critical.

An ancient paper on the subject, yet still so contemporary!
There may be numerous contaminants in commercial peroxide, and they may have to be contended with under some circumstances. Perhaps the author's difficulties, will provide illumination.
https://archive.org/details/interactionofpla00shaf/page/n1/m...



[Edited on 2-9-2020 by zed]

VeritasC&E - 2-9-2020 at 04:14

Quote: Originally posted by unionised  
I think you need to start by looking up "sub distillation" i.e. "sub boiling distillation".



Isn't this just slow distillation under BP? Woudn't it be enhanced by vacuum? (which also also would allow a significantly reduced rate of decomposition)

Your suggestion is good but I just want to be sure I correctly understand the process to follow and how efficient that would be. How much H2O2 can I expect to carry away with the water vapour into the receiver flask (i.e. what concentration can I expect in my receiver starting with my 9% material)?

[Edited on 2-9-2020 by VeritasC&E]

VeritasC&E - 2-9-2020 at 05:29

In theory, separation by freezing could be a route (assuming H2O2 is a bad solvent)

In practice however the process looks it could be very tedious (very close MPs), and hard to perform without contaminating the solution that is supposedly being purified.

In theory:

Freeze a portion of the solution > Pass through pre-cleaned Frit > Dilute again with Distilled Water > Mix Well > Repeat

The main problems I could see:

1) How to minimize H2O2 loss in the ice?
2) Contamination during filtration
3) Energy & Time Cost





[Edited on 2-9-2020 by VeritasC&E]

unionised - 2-9-2020 at 08:14

Quote: Originally posted by VeritasC&E  
Quote: Originally posted by unionised  
I think you need to start by looking up "sub distillation" i.e. "sub boiling distillation".



Isn't this just slow distillation under BP? Woudn't it be enhanced by vacuum? (which also also would allow a significantly reduced rate of decomposition)

Your suggestion is good but I just want to be sure I correctly understand the process to follow and how efficient that would be. How much H2O2 can I expect to carry away with the water vapour into the receiver flask (i.e. what concentration can I expect in my receiver starting with my 9% material)?

[Edited on 2-9-2020 by VeritasC&E]


The important aspect is that you don't boil the material.
That way, you don't generate "spray" which carries over involatiles like stabilisers.
In principle, starting with 9% you end up with 9%.
But without any involatile materials.

VeritasC&E - 2-9-2020 at 11:31

Quote: Originally posted by unionised  
Quote: Originally posted by VeritasC&E  
Quote: Originally posted by unionised  
I think you need to start by looking up "sub distillation" i.e. "sub boiling distillation".



Isn't this just slow distillation under BP? Woudn't it be enhanced by vacuum? (which also also would allow a significantly reduced rate of decomposition)

Your suggestion is good but I just want to be sure I correctly understand the process to follow and how efficient that would be. How much H2O2 can I expect to carry away with the water vapour into the receiver flask (i.e. what concentration can I expect in my receiver starting with my 9% material)?

[Edited on 2-9-2020 by VeritasC&E]


The important aspect is that you don't boil the material.
That way, you don't generate "spray" which carries over involatiles like stabilisers.
In principle, starting with 9% you end up with 9%.
But without any involatile materials.


But isn't there a significant BP difference between H2O and H2O2? Do they form an azeotrope? In which case, would you know within which proportions and parameters?

[Edited on 2-9-2020 by VeritasC&E]

unionised - 2-9-2020 at 13:58

They might form an azeotrope, I think so, but I'd need to look it up.

However, an azeotrope is only relevant if you are boiling something and...

VeritasC&E - 2-9-2020 at 22:16

Quote: Originally posted by unionised  
They might form an azeotrope, I think so, but I'd need to look it up.

However, an azeotrope is only relevant if you are boiling something and...


Yes, we are not boiling.

What I mean is that unless they behave like such, I don't understand quite why I would get in the receiver flask the same concentration as in the original concentration when the BP of the two miscibles are so far appart. Per my limited understanding, I wouldn't expect that to happen. I'm not sure how to use their relative vapour pressures but I'd expect that they could be used to indicate what the concentration in the receiver might look like (I would expect it to be lower, or even somewhat higher, but not the same as in the initial solution).

[Edited on 3-9-2020 by VeritasC&E]

unionised - 3-9-2020 at 00:00

Because, in principle, you evaporate all of the contents of the distillation flask and condense all of it in the receiver.

VeritasC&E - 3-9-2020 at 01:42

Quote: Originally posted by unionised  
Because, in principle, you evaporate all of the contents of the distillation flask and condense all of it in the receiver.


Is there no risk of overconcentrating the H2O2 in either the initial portion of the receiver or (more likely) in the very late portion of the source flask?

When only half of the solution has been carried over, do you know if the receiver will have an increased or rather a decreased H2O2 concentration?

I know 9% H2O2 is very safe, but I also know H2O2 vapours or more concentrated solutions can be dangerous, so it's important that I can visualize the dynamics of the process to know where H2O2 can get concentrated in case I need to do something to avoid it.

Thanks a lot for your contributions by the way!

[Edited on 3-9-2020 by VeritasC&E]

teodor - 3-9-2020 at 02:01

Could the method of purification by crystallisation with pure urea or sodium carbonate be as good as distillation?

Also, you will unable to remove some usual tap water organic or gases like NH3 with a distillation. By the way, the presence of organic matter and small quantities of resulting organic peroxides probably is something which should be considered when performing distillation, what do you think?



[Edited on 3-9-2020 by teodor]

[Edited on 3-9-2020 by teodor]

unionised - 3-9-2020 at 04:23

Quote: Originally posted by VeritasC&E  
Quote: Originally posted by unionised  
Because, in principle, you evaporate all of the contents of the distillation flask and condense all of it in the receiver.


Is there no risk of overconcentrating the H2O2 in either the initial portion of the receiver or (more likely) in the very late portion of the source flask?

When only half of the solution has been carried over, do you know if the receiver will have an increased or rather a decreased H2O2 concentration?

I know 9% H2O2 is very safe, but I also know H2O2 vapours or more concentrated solutions can be dangerous, so it's important that I can visualize the dynamics of the process to know where H2O2 can get concentrated in case I need to do something to avoid it.

Thanks a lot for your contributions by the way!

[Edited on 3-9-2020 by VeritasC&E]

When you are half way through, one container or the other will hold more than 9% H2O2
My guess is that water will evaporate preferentially at first leaving more concentrated H2O2 in the distillation flask and a more dilute solution in the receiver.

VeritasC&E - 3-9-2020 at 05:21

Quote: Originally posted by unionised  
Quote: Originally posted by VeritasC&E  
Quote: Originally posted by unionised  
Because, in principle, you evaporate all of the contents of the distillation flask and condense all of it in the receiver.


Is there no risk of overconcentrating the H2O2 in either the initial portion of the receiver or (more likely) in the very late portion of the source flask?

When only half of the solution has been carried over, do you know if the receiver will have an increased or rather a decreased H2O2 concentration?

I know 9% H2O2 is very safe, but I also know H2O2 vapours or more concentrated solutions can be dangerous, so it's important that I can visualize the dynamics of the process to know where H2O2 can get concentrated in case I need to do something to avoid it.

Thanks a lot for your contributions by the way!

[Edited on 3-9-2020 by VeritasC&E]

When you are half way through, one container or the other will hold more than 9% H2O2
My guess is that water will evaporate preferentially at first leaving more concentrated H2O2 in the distillation flask and a more dilute solution in the receiver.


Is there no danger towards the end of the process?

VeritasC&E - 3-9-2020 at 05:23

Quote: Originally posted by teodor  
Could the method of purification by crystallisation with pure urea or sodium carbonate be as good as distillation?

Also, you will unable to remove some usual tap water organic or gases like NH3 with a distillation. By the way, the presence of organic matter and small quantities of resulting organic peroxides probably is something which should be considered when performing distillation, what do you think?

[Edited on 3-9-2020 by teodor]


Hello! Could you describe that method?

[Edited on 3-9-2020 by VeritasC&E]

Fyndium - 3-9-2020 at 06:27

I'm under impression H2O2 can be safely vacuum distilled to a high concentration.

VeritasC&E - 3-9-2020 at 08:42

Quote: Originally posted by Fyndium  
I'm under impression H2O2 can be safely vacuum distilled to a high concentration.


I think it can, but I think what is happening when you do that is that you distill away mainly H2O, not H2O2, which not only would concentrate the impurities in the H2O2 solution (the opposite of what I'd like to do), but also might be dangerous because if I understand correctly H2O2 of higher concentrations at some point can be a possible detonation hazard (possibly more so when you also concentrate impurities).

So vacuum distillation might or might not be a good solution for people who want a higher concentration (which is a different topic, though it is welcome to be debated here also since we're dealing with the same solution properties).
As per the removal of impurities at constant concentration or the as it could be called the "lab preparation of puriss. H2O2 solution from technnical H2O2 solution", which is more or less what I'd like to do, I still don't know for sure if vacuum distillation can help.

The suggested sub-distillation might be a solution but I'm still not sure I understand it correctly (it would have to mean that under boiling point the vapors of H2O2 and H2O behave somewhat differently than during a distillation above BP, but my chemical understanding is yet too limited to fully understand these behaviours).


[Edited on 3-9-2020 by VeritasC&E]

macckone - 3-9-2020 at 09:17

Fyndium,
if by safely, you mean with only moderate risk of explosion, then yes.

Hydrogen peroxide that needs to be ultra pure is prepared by vacuum distillation.
However depending on impurities, this can lead to explosive decomposition.

The industrial method would be using barium peroxide. Barium oxide is combined with hydrogen peroxide or oxygen.
Then barium peroxide is decomposed with sulfuric acid. The barium sulfate is insoluble. If you have pure barium oxide and pure sulfuric acid, you have pure hydrogen peroxide.

Freezing can reduce impurities.
Some impurties will precipitate at before the hydrogen peroxide portion freezes and others will freeze at the end.
So you need to do a freeze, filter, freeze, decant

VeritasC&E - 3-9-2020 at 09:59

Quote: Originally posted by macckone  
Fyndium,
if by safely, you mean with only moderate risk of explosion, then yes.

Hydrogen peroxide that needs to be ultra pure is prepared by vacuum distillation.
However depending on impurities, this can lead to explosive decomposition.

The industrial method would be using barium peroxide. Barium oxide is combined with hydrogen peroxide or oxygen.
Then barium peroxide is decomposed with sulfuric acid. The barium sulfate is insoluble. If you have pure barium oxide and pure sulfuric acid, you have pure hydrogen peroxide.

Freezing can reduce impurities.
Some impurties will precipitate at before the hydrogen peroxide portion freezes and others will freeze at the end.
So you need to do a freeze, filter, freeze, decant



Do you know if the H2O2 can act as a solvent?

And do you know if the sub-distillation described, carried under vacuum, would work to purify the solution? (i.e. if under the BP of water H2O2 vapor is slowly carried away with water vapors in significant relative proportions compared to relative amounts that are carried with the water vapors above the BP of water)

macckone - 3-9-2020 at 10:27

VeritasC&E,
You are correct, concentrated hydrogen peroxide can detonate especially with concentrated impurities.

Understanding the impurities is important:
http://www.h2o2.com/faqs/FaqDetail.aspx?fId=11#:~:text=Commo...)%20also%20are%20used.

It is interesting that barium is not a listed ion impurity for hydrogen peroxide in acs reagent chemicals.
Sulfate however is.
So the method of producing hydrogen peroxide from barium peroxide will not introduce excessive impurities.
Barium Peroxide solubility is half of the allowed solids for standard reagent but not ultra-trace.
Ultra trace can only be achieved through multiple distillation in teflon gear. Ion exchange won't even do it.
Carbon and Fluorine are allowed contaminants for ultra trace, metals, chloride and sulfate are not.

VeritasC&E - 3-9-2020 at 10:55

Quote: Originally posted by macckone  
VeritasC&E,
You are correct, concentrated hydrogen peroxide can detonate especially with concentrated impurities.

Understanding the impurities is important:
http://www.h2o2.com/faqs/FaqDetail.aspx?fId=11#:~:text=Commo...)%20also%20are%20used.

It is interesting that barium is not a listed ion impurity for hydrogen peroxide in acs reagent chemicals.
Sulfate however is.
So the method of producing hydrogen peroxide from barium peroxide will not introduce excessive impurities.
Barium Peroxide solubility is half of the allowed solids for standard reagent but not ultra-trace.
Ultra trace can only be achieved through multiple distillation in teflon gear. Ion exchange won't even do it.
Carbon and Fluorine are allowed contaminants for ultra trace, metals, chloride and sulfate are not.


Interesting information!

If one can chose impurities I'd prefer Sodium, Boron and Aluminum over Fluorine so glass over teflon for me (also my bank account can only support glassware distillation).

I'm still wondering what portions of gases H2O2 can constitute in sub-distillation of an H2O2 solution. Does sub-BP distillation really work, not for concentrating H2O2 but to separate it from non-volatile impurities?

teodor - 3-9-2020 at 11:19

Quote: Originally posted by VeritasC&E  
Quote: Originally posted by teodor  
Could the method of purification by crystallisation with pure urea or sodium carbonate be as good as distillation?

Also, you will unable to remove some usual tap water organic or gases like NH3 with a distillation. By the way, the presence of organic matter and small quantities of resulting organic peroxides probably is something which should be considered when performing distillation, what do you think?

[Edited on 3-9-2020 by teodor]


Hello! Could you describe that method?

[Edited on 3-9-2020 by VeritasC&E]


Hello. This is just invitation to discover something by yourself - I never tried that method, so I have no description. But, you can check https://en.wikipedia.org/wiki/Hydrogen_peroxide_-_urea and http://library.sciencemadness.org/library/books/chemical_rea... .

The main question is "what impurities are critical for your work and should not be present". Because you always have some impurities (dissolved CO2) and also trying to get of some of them could introduce some other.

So, method of distillation can get rid of metal ions but will well keep organic, NH3 etc. Some of them as a result (probably) of decomposition of stabiliser. I didn't study this topic and so I say only general things.

Also, any purification method should not be trusted but it is better to check whether the result has critical impurities or not (for some common H2O2 tests, see, for example, http://library.sciencemadness.org/library/books/chemical_rea... page 107).

I agree, that the method of distillation is the best to get rid of metals. But the danger you mean is already connected with impurities in initial mixture - pure H2O2 on pure glassware without scratches should not decompose, at least in concentration of 25-30%. That is what I think. So, if you can do your work with urea-H2O2 crystals why bother with distillation. If not - distil after initial purification. Hope people with experience can give better practical advice, but if not - just try carefully.

[Edited on 3-9-2020 by teodor]

[Edited on 3-9-2020 by teodor]

VeritasC&E - 3-9-2020 at 23:30

Quote: Originally posted by teodor  
Quote: Originally posted by VeritasC&E  
Quote: Originally posted by teodor  
Could the method of purification by crystallisation with pure urea or sodium carbonate be as good as distillation?

Also, you will unable to remove some usual tap water organic or gases like NH3 with a distillation. By the way, the presence of organic matter and small quantities of resulting organic peroxides probably is something which should be considered when performing distillation, what do you think?

[Edited on 3-9-2020 by teodor]


Hello! Could you describe that method?

[Edited on 3-9-2020 by VeritasC&E]


Hello. This is just invitation to discover something by yourself - I never tried that method, so I have no description. But, you can check https://en.wikipedia.org/wiki/Hydrogen_peroxide_-_urea and http://library.sciencemadness.org/library/books/chemical_rea... .

The main question is "what impurities are critical for your work and should not be present". Because you always have some impurities (dissolved CO2) and also trying to get of some of them could introduce some other.

So, method of distillation can get rid of metal ions but will well keep organic, NH3 etc. Some of them as a result (probably) of decomposition of stabiliser. I didn't study this topic and so I say only general things.

Also, any purification method should not be trusted but it is better to check whether the result has critical impurities or not (for some common H2O2 tests, see, for example, http://library.sciencemadness.org/library/books/chemical_rea... page 107).

I agree, that the method of distillation is the best to get rid of metals. But the danger you mean is already connected with impurities in initial mixture - pure H2O2 on pure glassware without scratches should not decompose, at least in concentration of 25-30%. That is what I think. So, if you can do your work with urea-H2O2 crystals why bother with distillation. If not - distil after initial purification. Hope people with experience can give better practical advice, but if not - just try carefully.

[Edited on 3-9-2020 by teodor]

[Edited on 3-9-2020 by teodor]



The urea-H2O2 possibility is a great new possibility, though it comes with a few limitations.

If somehow the binding could be but a transient step of H2O2 purification it would be absolutely awesome, but I'm not sure how that would be. My limited knowledge would want me to get rid of the urea by heating the crystals in distilled water but I guess the H2O2 would degrade long before the urea and its degradation products have degraded. My current chemical knowledge is too limited to see how this could be done. Do you consider this an "irreversible" binding or do have an idea how the urea and H2O2 could subsequently be cleanly separated as to recreate a pure H2O2/H2O solution?

A great purification process would be precipitation as H2O2-Urea followed by (if it works) sub-distillation, as these two are quite complementary (precipitation more likely will allow to separate H2O2 from any volatile impurities that sub-distillation wouldn't deal with efficiently).


_________

PS: Somehow the sciencemadness links aren't working for me, both links lead me to a page with the following content:

Not Found

The requested URL was not found on this server.

Apache/2.4.18 (Ubuntu) Server at library.sciencemadness.org Port 80

[Edited on 4-9-2020 by VeritasC&E]

unionised - 3-9-2020 at 23:59

I doubt that anyone has used the barium peroxide process for making H2O2 commercially for 50 years or more.
The solubility of BaSO4 is about 250 ppm. So, by the standards of high purity chemicals, that process doesn't produce a pure product.
All commercially produced H2O2 is distilled.

VeritasC&E - 4-9-2020 at 00:56

Quote: Originally posted by unionised  
I doubt that anyone has used the barium peroxide process for making H2O2 commercially for 50 years or more.
The solubility of BaSO4 is about 250 ppm. So, by the standards of high purity chemicals, that process doesn't produce a pure product.
All commercially produced H2O2 is distilled.


Hi Unionised!

Can you confirm that when distilling under BP the H2O2 concentration in the source flask won't increase to the point at which it will risk to explode my apparatus? And could you explain the science behind this? i.e. how it is, which I assume it would need to be, that under BP a higher amount of H2O2 vapors relative to H2O vapors carries over compared to above the BP of water?

teodor - 4-9-2020 at 03:17

Quote: Originally posted by VeritasC&E  

A great purification process would be precipitation as H2O2-Urea followed by (if it works) sub-distillation, as these two are quite complementary (precipitation more likely will allow to separate H2O2 from any volatile impurities that sub-distillation wouldn't deal with efficiently).
[Edited on 4-9-2020 by VeritasC&E]


I have no much experience with Urea chemistry, but I easily created a barium peroxide from H2O2 and BaOH, so I think you can try to do it also from H2O2-Urea (or, with some modifications, from percarbonate). Then, add H2SO4 to make 3% solution and distil. If I properly understand the video I posted here, there is 25% azeotrope and you don't need a vacuum. So, 3 steps. Or try to do it with H2O2->BaO2->H2O2 but I think it will be more contaminated because no crystallisation step, also fine powders like BaO2 always act as absorbents.

As for preparation of highest grade H2O2 ("spectroscopic purity") Brauer recommends persulfate-steam method but I have no idea weather it is doable at all, just can attach here the German article he mentioned (I don't know German, may be people who knows can say whether is it something useful. The Brauer says the same method as for D2O2 in the article could be applied for making H2O2 of highest purity).



Attachment: 10.1002@cber.19390720925.pdf (700kB)
This file has been downloaded 36 times

[Edited on 4-9-2020 by teodor]

[Edited on 4-9-2020 by teodor]

[Edited on 4-9-2020 by teodor]

VeritasC&E - 4-9-2020 at 03:35

Quote: Originally posted by teodor  
Quote: Originally posted by VeritasC&E  

A great purification process would be precipitation as H2O2-Urea followed by (if it works) sub-distillation, as these two are quite complementary (precipitation more likely will allow to separate H2O2 from any volatile impurities that sub-distillation wouldn't deal with efficiently).
[Edited on 4-9-2020 by VeritasC&E]


I have no much experience with Urea chemistry, but I easily created a barium peroxide from H2O2 and BaOH, so I think you can try to do it also from H2O2-Urea. Then, add H2SO4 to make 3% solution and distil. If I properly understand the video I posted here, there is 25% azeotrope and you don't need a vacuum. So, 3 steps. Or try to do it with H2O2->BaO2->H2O2 but I think it will be more contaminated because no crystallisation step, also fine powders like BaO2 always act as absorbents.

As for preparation of highest grade H2O2 ("spectroscopic purity") Brauer recommends persulfate-steam method but I have no idea weather it is doable at all, just can attach here the German article he mentioned (I don't know German, may be people who knows can say whether is it something useful. The Brauer says the same method as for D2O2 in the article could be applied for making H2O2 of highest purity).


[Edited on 4-9-2020 by teodor]

[Edited on 4-9-2020 by teodor]


I'll look into it and try methods out when I have time (which may be in a while). I'll report back here if I successfully apply one if these processes to start with H2O2 with impurities and end up with H2O2 with extremely low contamination. Please share any ideas, data or recommendations on the subject in the mean time.

AJKOER - 5-9-2020 at 14:29

A suggestion that may serve your purposes.

First, create zinc peroxide (see https://en.wikipedia.org/wiki/Zinc_peroxide), whose prep per Wikipedia has been described as:

"Zinc hydroxide is reacted with a mixture of hydrochloric acid and hydrogen peroxide and precipitated with sodium hydroxide also containing hydrogen peroxide to ensure a higher yield of zinc peroxide."

Although, here I would employ pure NH3 (aq) in place of NaOH, and thoroughly rinse the product with recently boiled distilled water.

Next, try treating the damp peroxide salt with CO2 gas and warming, followed by capturing and condensing the vapors.

This is based on the following source: 'Biofunctionalized zinc peroxide (ZnO2) nanoparticles as active oxygen sources and antibacterial agents" at https://pubs.rsc.org/en/content/articlelanding/2017/ra/c7ra0... ), to quote:

"Zinc peroxide dissociates in aqueous acidic media into zinc ions (Zn2+) and hydrogen peroxide (H2O2) while the hydrogen peroxide can be immediately converted into water and oxygen in presence of metal salts or metal oxide surfaces which are provided by the nanoparticles"

Note: Here I am not recommending a nanoparticle preparation path, also try avoiding strong light or lab rich UV illumination, and do not expect a long shelf life for the created H2O2.

Here is a link to a prior SM thread dedicated to Zinc peroxide at https://www.sciencemadness.org/whisper/viewthread.php?tid=14... .

[Edited on 5-9-2020 by AJKOER]

unionised - 6-9-2020 at 01:45

Quote: Originally posted by VeritasC&E  
Quote: Originally posted by unionised  
I doubt that anyone has used the barium peroxide process for making H2O2 commercially for 50 years or more.
The solubility of BaSO4 is about 250 ppm. So, by the standards of high purity chemicals, that process doesn't produce a pure product.
All commercially produced H2O2 is distilled.


Hi Unionised!

Can you confirm that when distilling under BP the H2O2 concentration in the source flask won't increase to the point at which it will risk to explode my apparatus? And could you explain the science behind this? i.e. how it is, which I assume it would need to be, that under BP a higher amount of H2O2 vapors relative to H2O vapors carries over compared to above the BP of water?

No, I can't confirm that, which is why I said this earlier

Quote: Originally posted by unionised  

When you are half way through, one container or the other will hold more than 9% H2O2
My guess is that water will evaporate preferentially at first leaving more concentrated H2O2 in the distillation flask and a more dilute solution in the receiver.

RogueRose - 6-9-2020 at 14:08

Quote: Originally posted by unionised  
I think you need to start by looking up "sub distillation" i.e. "sub boiling distillation".


Thanks for mentioning this procedure. I tried looking for some good pictures of setups to see how they really work, but there are hardly any high resolution pics and the ones that are available you can't see much. It looks fairly simple - radiant heat above the liquid and then a water chilled condensing tube next to it (I can't tell if it is shielded from heat though...) and angled down so the droplets that form run down into a collection tube. I'm not sure if there is anything that I'm missing in the process but does that about sum it up?

macckone - 8-9-2020 at 07:30

Unionized,
The anthrohydroquinone process is not 'home lab friendly'.
Not that it can't be done.
Heat barium oxide in a stream of air and add sulfuric acid is definitely easier.

ACS grade allowable residue after evaporation is 0.002%.
Barium sulfate solubility is 0.0002448% aka 2.5 ppm not 250ppm.
Specifically at 20C it is 0.0002448 g/100 mL.


unionised - 8-9-2020 at 08:39

Quote: Originally posted by macckone  
Unionized,
The anthrohydroquinone process is not 'home lab friendly'.
Not that it can't be done.
Heat barium oxide in a stream of air and add sulfuric acid is definitely easier.

ACS grade allowable residue after evaporation is 0.002%.
Barium sulfate solubility is 0.0002448% aka 2.5 ppm not 250ppm.
Specifically at 20C it is 0.0002448 g/100 mL.


"The anthrohydroquinone process is not 'home lab friendly'."
Had anyone suggested that it was?

Good catch on the PPM.
But 2.5ppm is still not in the realms of high purity.

Heating BaO in air is easy
Buying commercial H2O2 and cleaning it up is
(1) easier
(2) What the OP actually asked about

I'm not sure if I'd rather try sub distilling H2O2 or H2SO4; neither is exactly friendly.
But, if you want a high purity product you would need to do one or the other (or some other process for cleaning it up).

macckone - 8-9-2020 at 17:35

The requirement for ACS grade H2O2 is 20ppm.
The ultra trace is of course lower but that is only going to be necessary if you are doing trace metal analysis.

ACS grade specs:
Assay 29.0 - 32.0%
Color(APHA) 10
Residue after evap: 0.0002% (20ppm)
Titratable Acid: 0.0006 meq/g
Chloride: 2ppm
Nitrate: 2ppm
Phosphate: 2ppm
Sulfate: 5ppm
Ammonium: 5ppm
Heavy metal (as Pb): 1ppm
Iron: 0.5ppm

Ultratrace type
Assay 25-35%
Chloride 3ppm
Nitrate 2ppm
Phosphate 2ppm
Sulfate 5ppm

Most cations are 1ppb
Boron, Iron, Silicon, and Sodium are 5ppb

You cannot distill in glass and reach the correct trace levels.
You have to use Teflon.
There are specific ion exchange resins that are compatible with high test H2O2 that are used to remove cations.