| Pages:
1
2 |
VeritasC&E
Hazard to Others
Posts: 176
Registered: 29-1-2018
Member Is Offline
Mood: No Mood
|
|
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
Dispenser of practical lab wisdom
Posts: 2168
Registered: 1-3-2013
Location: Over a mile high
Member Is Offline
Mood: Electrical
|
|
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
Hazard to Others
Posts: 176
Registered: 29-1-2018
Member Is Offline
Mood: No Mood
|
|
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
Dispenser of practical lab wisdom
Posts: 2168
Registered: 1-3-2013
Location: Over a mile high
Member Is Offline
Mood: Electrical
|
|
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
Hazard to Others
Posts: 176
Registered: 29-1-2018
Member Is Offline
Mood: No Mood
|
|
| 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
National Hazard
Posts: 924
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
| 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
Hazard to Others
Posts: 176
Registered: 29-1-2018
Member Is Offline
Mood: No Mood
|
|
| 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
International Hazard
Posts: 5128
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
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
Hazard to Others
Posts: 176
Registered: 29-1-2018
Member Is Offline
Mood: No Mood
|
|
| 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
National Hazard
Posts: 924
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
| 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 288 times
[Edited on 4-9-2020 by teodor]
[Edited on 4-9-2020 by teodor]
[Edited on 4-9-2020 by teodor]
|
|
|
VeritasC&E
Hazard to Others
Posts: 176
Registered: 29-1-2018
Member Is Offline
Mood: No Mood
|
|
| 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
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
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
International Hazard
Posts: 5128
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
| 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
International Hazard
Posts: 1596
Registered: 16-6-2014
Member Is Offline
|
|
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
Dispenser of practical lab wisdom
Posts: 2168
Registered: 1-3-2013
Location: Over a mile high
Member Is Offline
Mood: Electrical
|
|
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
International Hazard
Posts: 5128
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
| 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
Dispenser of practical lab wisdom
Posts: 2168
Registered: 1-3-2013
Location: Over a mile high
Member Is Offline
Mood: Electrical
|
|
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.
|
|
|
| Pages:
1
2 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
