Sciencemadness Discussion Board

Quick and easy source of H2O2?

Polesch - 22-4-2014 at 04:01

In many countries obtaining hydrogen peroxide has become increasingly difficult, and recent EU regulations have limited OTC peroxide solutions to a maximum of 12%. And for many people the pharmacy is the only OTC source, and will often be very expensive or/and only 3%.

There is a simple and easy solution to this, but as of yet it's only an idea. All my chemistry equipment have been set aside - so I'm hoping some of you active hobbyists could help with this.

A very common bleaching powder is sodium percarbonate (like Vanish OxiAction Crystal White Powder), which is an adduct of hydrogen peroxide and sodium carbonate. This chemical alone is incredibly useful as an oxidizer in a lot of chemistry, but it would be more practical if one could simply separate it to get pure H2O2 solutions.

The simplest way to do this - I think - is to add calcium chloride. Most of the carbonate will precipitate as CaCO3, leaving a solution of H2O2 and NaCl in water. Which again is more useful, for when the H2O2 will be used along with acids for example.

I guess you could further separate the two with vacuum distillation, but that's not necessary for my purposes.

Anyone have CaCl2, sodium percarbonate and a peroxidase/catalyst (and 5 minutes)?

See if this works? You should be able to end up with a ~12.5% H2O2 solution if the initial solution is saturated with sodium percarbonate (150 g/L).

macckone - 22-4-2014 at 05:28

Most of the percarbonate products around here will only give you
a 3% H2O2solution when dissolved.
They also aren't pure percarbonate, there is substantial carbonate
in most of them. Oxiclean brand for example is 30-45% percarbonate.
With the remainder being carbonate and some form of alcohol 1-3%.
http://www.ahprofessional.com/_downloads/msds/MSDS-1605-OxiC...

You would be better off using the freezing method
for concentrating the 3% off the shelf. Remember the percarbonate
is only soluble at 150g/L meaning you get a 15% solution of the
percarbonate, not peroxide.

One solution that may work is to attempt to vacuum dry distill
the percarbonate. The MSDS says it evaporate at 215F.

Another method, not necessarily safe would be attempting
to leach out the hydrogen peroxide with isopropyl alcohol.
Hydrogen peroxide is substantially soluble in various alcohols
while the carbonate is not. The 1-3% of other alcohol would
in theory carry over. But is heavy with high boiling point (c12-c15).
So after the alcohol extraction, you could vacuum distill
and get practically 100% hydrogen peroxide. But I would
do this outside and stand far far away. These things can
go badly fast. I also would not distill to dryness.

Bert - 22-4-2014 at 07:03
Quote:

and get practically 100% hydrogen peroxide. But I would do this outside and stand far far away. These things can go badly fast. I also would not distill to dryness.
Oh yes- Let's heat up some organic liquid fuel with peroxide in it... What could go wrong?

Perhaps this would be an alternative-


Carbamide peroxide?


[Edited on 22-4-2014 by Bert]

DoctorZET - 22-4-2014 at 08:32

I remember that if sodium oxide is heated up to 200-300*C in oxygen atmosphere, for 15-30 min ... will result sodium peroxide (95% +) with some impurities (sodium oxide, sodium superoxide...).

NOTE: =>Sodium peroxide has a very pale yellowish-grey color.

Next, add slowly the powder into distilled water (for 1 mol of sodium peroxide use almost 2 mol of water).
Then, distill the mixture and cool it.

NOTE:
=>a mixture about 50-50 of water and hidrogen peroxide is an eucthetic mixture with the boiling point at about 114*C (so, you don't need vaccum distilation)
=>Put the H2O2 solution in a dark glass bottle!

So, ~50% H2O2 might be produced very easy.

Zyklon-A - 22-4-2014 at 08:43
Quote:

Then, distill the mixture and cool it.
I don't think hydrogen peroxide can be distilled, it will decompose.
To purify it somewhat, perhaps add stoichiometric amounts of calcium iodate, which will precipitate calcium hydroxide and sodium iodate, both of which are almost insoluble.
It still wont be pure, but it will be more pure.

[Edited on 22-4-2014 by Zyklonb]

DoctorZET - 22-4-2014 at 09:47

Arr ... at 120 *C, H2O2 has a half-life about 20-25 min ... and in a distillation tube, H2O2 and water exist in gas state for just a few seconds ... so it will not be a problem.

On the other hand, if the calcium iodate has some iodide impurities, the iodide ion will decompose very fast the H2O2 into H2O and O2:
H2O2 + I- --> H2O + OI-
H2O2 + OI- --> H2O + I- + O2
Moreover H2O2 should can oxidate IO3- to IO4- , but also one of this 2 ions can oxidate H2O2 to H2O and O2 ... and that's why I don't have the courage to dissolve Ca(IO3)2 into a H2O2 solution, and store it for a time longer than 4-5 minutes...:)

Zyklon-A - 22-4-2014 at 11:05

Quote:
Originally posted in Wikipedia
The boiling point of H2O2 has been extrapolated as being 150.2 °C, however, in practice hydrogen peroxide will undergo potentially explosive thermal decomposition if heated to this temperature. It may be safely distilled under reduced pressure via a variety of techniques.

Sure, it can be done, but good luck with that.
Iodate was just a suggestion, I wouldn't have tried that, there are probably other possibility's, but it's hard to find sodium salts that are insoluble were the corresponding calcium salt in more soluble. (Otherwise the reaction will keep sodium hydroxide in solution.) Potassium peroxide would be better, calcium perchlorate plus potassium hydroxide seems like a perfect match, as both products are nearly insoluble and potassium peroxide is easier to make. There are certainly much better ways to make hydrogen peroxide, and any potassium oxide is harder and more expensive than hydrogen peroxide.


[Edited on 22-4-2014 by Zyklonb]

DoctorZET - 22-4-2014 at 12:47
Quote:

"In aqueous solutions
In aqueous solutions hydrogen peroxide differs from the pure material due to the effects of hydrogen bonding between water and hydrogen peroxide molecules. Hydrogen peroxide and water form a eutectic mixture, exhibiting freezing-point depression; pure water has a melting point of 0°C and pure hydrogen peroxide of −0.43 °C, but a 50% (by volume) solution of the 2 freezes at -51°C. The boiling point of the same mixture is also depressed and occurs at 114°C, less than the average of the boiling points of pure water and hydrogen peroxide (125°C)."

But yes ... you're right, there are more ways to make purer and more concentrated H2O2 ...
For example:
1)Everybody can go to a pharmacy, buy some BaSO4 (used as X-ray contrast agent for an intestinal radiography).
2)But, instead of drinking an emulsion of BaSO4, you can heat it up with sodium metal, reduce it to BaSO3 (barium sulphite).
3)Then put it in water to remove the Na ions, and decant it.
4)After this, take the powder and mix it with some diluted HNO3, evaporate the water+SO2, and heat up the Ba(NO3)2 to decompose it into O2 , NO and BaO.
5)Heat the BaO in oxygen atmosphere up to 500*C to form BaO2 (the peroxide)
6)Dissolve it in water, then bubble CO2 in the solution (to form unsoluble BaCO3)
7)Repeat the steps (using the BaCO3) begining with the step "4)" down to here.

Good thing I have an ozone generator (cold plasma, 570W, 23-24kV).
If I let the output gas to bubble in a little beaker with water at 0*C for 4-5 hours, I can get up to 99% H2O2 . For me, this method is the best of all I could do ... :cool:

HgDinis25 - 22-4-2014 at 12:50

Read this thread:
http://www.sciencemadness.org/talk/viewthread.php?tid=17052

It discusses calcium peroxide formation (with a relatively easy process) and forming Hydrogen Peroxide from it.

Chemosynthesis - 22-4-2014 at 13:03

I'm with Bert.

As noted in the above posts, High Test Peroxide (HTP) is high concentration hydrogen peroxide, and it's an energetic propellant and explosive. Distillation, even under vacuum without the application of heat, is not advisable. Organic peroxides have extremely low bond enthalpies, and are not stable. You are best suited to freezing out your water to get 30-35% peroxide, in my opinion.

macckone - 22-4-2014 at 15:51

Quote: Originally posted by Bert  



Quote:

and get practically 100% hydrogen peroxide. But I would do this outside and stand far far away. These things can go badly fast. I also would not distill to dryness.
Oh yes- Let's heat up some organic liquid fuel with peroxide in it... What could go wrong?

[Edited on 22-4-2014 by Bert]


I believe your signature states very clearly what could go wrong. BOOM.
I recommended freezing 3% as the best way to do it.

macckone - 22-4-2014 at 16:03

Quote: Originally posted by Chemosynthesis  
I'm with Bert.

As noted in the above posts, High Test Peroxide (HTP) is high concentration hydrogen peroxide, and it's an energetic propellant and explosive. Distillation, even under vacuum without the application of heat, is not advisable. Organic peroxides have extremely low bond enthalpies, and are not stable. You are best suited to freezing out your water to get 30-35% peroxide, in my opinion.


I am with you on the freezing.
That was my actual recommendation.

I specifically said distillation is not safe.
Distillation has to be done under vacuum in small quantities and
no one allowed near the apparatus. To get higher percentage
it is actually distilled in commercial application.

http://www.dtic.mil/dtic/tr/fulltext/u2/819081.pdf
http://www.diyspaceexploration.com/preparing-manufacture-hyd...

But because it can be done from a half a mile away in a plant
doesn't mean you want to do it in your back yard.

Zyklon-A - 22-4-2014 at 16:24

I'm currently boiling some hydrogen peroxide to test the decomposition, I've heard many people here say that it will decompose far to much, but one other person said it worked with ~90% yields. We shall see.

macckone - 22-4-2014 at 17:18

Quote: Originally posted by Zyklonb  
I'm currently boiling some hydrogen peroxide to test the decomposition, I've heard many people here say that it will decompose far to much, but one other person said it worked with ~90% yields. We shall see.


I hope you are doing that under vacuum and outside with plenty of room for boom.
If you do achieve concentrated peroxide it can decompose in an uncontrolled manner.

Zyklon-A - 22-4-2014 at 20:14

No, I only started with 3% and decreased its volume by about 90%, which can give a maximum of 30% hydrogen peroxide. It's done now, but I have not checked its concentration, maybe tomorrow...

macckone - 22-4-2014 at 20:57

Did you do vacuum at all?
Glad to know it didn't decompose in an uncontrolled manner.

Zyklon-A - 23-4-2014 at 13:43

No vacuum, I just watched it closely.

DoctorZET - 23-4-2014 at 14:05

I think that any mixture of H2O2 and water (up to 50% H2O2) can be distilled without problems, because it forms an azeotrope with the boiling point at about 114*C (not 150, as the pure H2O2).
This should be right, because, remember : both hydrogen peroxide and water (hidrogen oxide) have a lot of hydrogen bonds, so H2O molecules will be very "sticky" for the H2O2 ones.

macckone - 23-4-2014 at 16:02

Quote: Originally posted by DoctorZET  
I think that any mixture of H2O2 and water (up to 50% H2O2) can be distilled without problems, because it forms an azeotrope with the boiling point at about 114*C (not 150, as the pure H2O2).
This should be right, because, remember : both hydrogen peroxide and water (hidrogen oxide) have a lot of hydrogen bonds, so H2O molecules will be very "sticky" for the H2O2 ones.

Do you have a reference with percentage of the azetrope?
I don't think it forms a true constant boiling azetrope.

blogfast25 - 24-4-2014 at 12:31

For what it's worth, the method to freeze out some water from H2O2 solutions:

http://www.sciencemadness.org/talk/viewthread.php?tid=29523#...

vmelkon - 24-4-2014 at 13:07

Quote: Originally posted by Zyklonb  
No, I only started with 3% and decreased its volume by about 90%, which can give a maximum of 30% hydrogen peroxide. It's done now, but I have not checked its concentration, maybe tomorrow...


Nah, it will be around 17% conc.
I had 473 mL and distilled until I was left with 50 mL.
Density was between 1.07 and 1.08 g/mL.
I did it twice and had the same result.

blargish - 26-4-2014 at 10:02

Quote: Originally posted by HgDinis25  
Read this thread:
http://www.sciencemadness.org/talk/viewthread.php?tid=17052

It discusses calcium peroxide formation (with a relatively easy process) and forming Hydrogen Peroxide from it.


I did a bunch of experiments on this conversion, starting by synthesizing CaO2 (calcium peroxide) from dilute H2O2, CaCl2·2H2O, and NH3; and reconverting it back to H2O2 via acidification.

Forming the CaO2 was kind of easy as it used cheap and easy to find reagents. Getting a pure sample was more difficult. My most pure sample, about 97% and essentially anhydrous, (based on experimentation with the decomposition of the peroxide) only had like a 30% yield based on the starting CaCl2. I was getting yields of up to 90% of about 74% purity with another method of synthesis. I did not examine many different methods of synthesis and I'm sure this can be improved.

Once you have this CaO2, in theory, you could concentrate this right up to whatever concentration desired by adding the desired stoichiometric amount of acid to the CaO2 and precipitating the calcium ion with sulfuric acid. I did tests with nitric acid and through titration determined that only 12 percent of H2O2 was lost during the reconversion.

The real problem with this is that you're left with a soluble nitrate impurity at the end (anion depends on acid used) which kinda sucks, but there may be a way of removing it that I do not know of.

quantumcorespacealchemyst - 21-11-2014 at 01:14

if i recall correctly there is a member here who has a hydrogen peroxide generator (read in a post somewhere). if we can get him to help, that will be great.

not so quick and easy, and a thread to start if this is not welcome here, is

http://en.wikipedia.org/wiki/Hydrogen_peroxide

Manufacture

Previously, hydrogen peroxide has been prepared industrially by hydrolysis of the ammonium peroxydisulfate, which was itself obtained via the electrolysis of a solution of ammonium bisulfate (NH
4HSO
4) in sulfuric acid.

(NH4)2S2O8 + 2 H2O → H2O2 + 2 (NH4)HSO4

Today, hydrogen peroxide is manufactured almost exclusively by the anthraquinone process, which was formalized in 1936 and patented in 1939. It begins with the reduction of an anthraquinone (such as 2-ethylanthraquinone or the 2-amyl derivative) to the corresponding anthrahydroquinone, typically via hydrogenation on a palladium catalyst; the anthrahydroquinone then undergoes to autoxidation to regenerate the starting anthraquinone, with hydrogen peroxide being produced as a by-product. Most commercial processes achieve oxidation by bubbling compressed air through a solution of the derivatized anthracene, whereby the oxygen present in the air reacts with the labile hydrogen atoms (of the hydroxy group), giving hydrogen peroxide and regenerating the anthraquinone. Hydrogen peroxide is then extracted and the anthraquinone derivative is reduced back to the dihydroxy (anthracene) compound using hydrogen gas in the presence of a metal catalyst. The cycle then repeats itself.[13][14]



The simplified overall equation for the process is deceptively simple:[13]

H2 + O2 → H2O2

The economics of the process depend heavily on effective recycling of the quinone (which is expensive) and extraction solvents, and of the hydrogenation catalyst.

A process to produce hydrogen peroxide directly from the elements has been of interest for many years. Direct synthesis is difficult to achieve as, in terms of thermodynamics, the reaction of hydrogen with oxygen favours production of water. Systems for direct synthesis have been developed; most of which are based around finely dispersed metal catalysts.[15][16] However none of these have yet reached a point where they can be used for industrial-scale synthesis.

perhaps we can manage the former and the latter. the latter cites
http://onlinelibrary.wiley.com/doi/10.1002/9783527627547.ch8...
and
http://www.sciencemag.org/content/323/5917/1037

the latter mentions an Au-Pd alloy

greenlight - 21-11-2014 at 02:14

If you only require a small amount of high concentration Hydrogen peroxide, the fibreglass repair kits sold at hardware stores contain about 50 ml of 50% H202 used as the "hardener" after the hole is patched.
I suppose you could use this if your absolutely stuck and need H202, but it sucks having to buy a whole heap of other things you can't use like fibreglass resin as it come in a pack.
I don't think these would be banned in your country.

NO AZEOTROPE WITH WATER

quantumcorespacealchemyst - 21-11-2014 at 03:10

according to Craig W. Jones in the book "Applications of Hydrogen Peroxide and Derivatives"
1999
http://books.google.com/books?id=QLePlhTsE0oC&pg=PA14&am...

"hydrogen peroxide and water do not form azeotropic mixtures and can be completely separated by distillation. most workers, however, obtain 100% w/w hydrogen peroxide by fractional crystallization of highly concentrated solutions"

page 14

also according to chapter 1, introduction to the production and properties of hydrogen peroxide

http://www.google.com/url?sa=t&rct=j&q=&esrc=s&a...

"Higher
strengths can be achieved as hydrogen peroxide does not form an azeotrope
with water, but a number of technical safety requirements must be observed.
"

page 9

attached is the pdf. for backup

also according to "Process for the production of non-aqueous hydrogen peroxide solutions and their use
US 4686010 A"

http://www.google.com/patents/US4686010

"The production of solutions of hydrogen peroxide in phenol or its derivatives, e.g. hydrocarbyl substituted phenols, halo substituted phenols or phenol ethers, is carried out in a single step. Practically no loss of hydrogen peroxide occurs since a total distillation of hydrogen peroxide together with phenol or phenol derivative is avoided. Simultaneously the solutions obtained are practically free from water. The mixture of phenol or phenol derivative and aqueous hydrogen peroxide is treated with a material that boils below the boiling point of hydrogen peroxide, phenol or phenol derivative or forms an azeotrope with water that boils below the boiling point of hydrogen peroxide, phenol or phenol derivative and the water removed as an azeotrope. The solution of hydrogen peroxide in phenol or phenol derivative which remains behind is suitable for carrying out oxidation reactions and above all, also for hydroxylation reactions."

DISTILLATION

according to http://en.wikipedia.org/wiki/Hydrogen_peroxide

"The boiling point of H2O2 has been extrapolated as being 150.2 °C, approximately 50 degrees higher than water; however, in practice hydrogen peroxide will undergo potentially explosive thermal decomposition if heated to this temperature. It may be safely distilled under reduced pressure via a variety of techniques."

specifically, "Riley, edited by Georg Brauer ; translated by Scripta Technica, Inc. Translation editor Reed F. (1963). Handbook of preparative inorganic chemistry. Volume 1 (2nd ed. ed.). New York, N.Y.: Academic Press. p. 140. ISBN 978-0121266011."

also, just to add something from the 1st book, as it isn't downloadable (or i haven't figured how), "hydrogen peroxide and it's highly concentrated aqeous (>65% m/m) are soluble in a range of organic solvents, such as carboxylic esters"

[Edited on 21-11-2014 by quantumcorespacealchemyst]

Attachment: perxoide-property.pdf (2MB)
This file has been downloaded 1677 times

[Edited on 21-11-2014 by quantumcorespacealchemyst]

[Edited on 21-11-2014 by quantumcorespacealchemyst]

[Edited on 21-11-2014 by quantumcorespacealchemyst]

greenlight - 21-11-2014 at 04:01

There is also this thread from 2006 which discusses distilling Hydrogen peroxide to 90%+ from dilute solution under reduced pressure and has some links to boiling points and pressures.

http://www.sciencemadness.org/talk/viewthread.php?tid=7030