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ecos
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I have attached a file that talks about the steps of manufacturing hydrogen peroxide under pressure of oxygen gas. it reference to a patent with
number 766,091 that has the detailed steps of the electrolysis process.
Unfortunately, I couldn't find this old patent to see what is exactly written inside 
Attachment: US1108752.pdf (198kB)
This file has been downloaded 710 times
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XeonTheMGPony
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got a general lay out figured just need to run material compatibilities.
general process will be O2 bottle, and H2 bottle, feed into two separate resavor chambers filled with pure water for the gas absorption phase then
these two fluids under pressure into the catylest chamber, the product will then be metered out via needle valve and then pressure regulator.
need to find more into on the catylist tube and how the diffusion works.
probably be a month or so as got lots to do here, but so far it seems simple enough using sch 80 ss 316l nipples and swagelok fittings (http://swagelok.com/en/product)
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ecos
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Did any try to get h2o2 from sodium peroxide?
It seems easy. Just oxidation of sodium and then mix water
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Melgar
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Quote: Originally posted by Jstuyfzand  | I do wonder, what does the pressure do in this process?
As far as I know, Which is not alot though, pressure speeds up reactions between gasses.
Maybe the pressure is not necessary, it might improve the efficiency.
That would be a big requirement to perform this process gone, it seems (Kind of) straight forward.
"Making H2O2 by electrolyzing Sulphuric acid and Bisulfates"
Anyone....? |
Increasing pressure on a gas favors the formation of larger molecules with more bonds. Diamonds, for example, only form in nature when the pressure
on graphite is so great that it can be relieved somewhat by rearranging its bonds into a diamond crystal structure that takes up less space. The
Haber process also takes advantage of this principle to generate ammonia from nitrogen and hydrogen, turning four molecules (1xN2, 3xH2) into two
molecules (2xNH3).
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Jstuyfzand
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interesting, thank you
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ecos
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I did a lot of reading and research.
my conclusion : H2O2 is very very very hard to be made at home from raw materials
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Jstuyfzand
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| Quote: Originally posted by ecos | I did a lot of reading and research.
my conclusion : H2O2 is very very very hard to be made at home from raw materials
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Which is a shame, because the EU loves making 30% h2o2 regulated!
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ecos
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yes, I agree.
the industrial method uses anthraquinone process. is it possible to buy anthraquinone (such as 2-ethylanthraquinone or the 2-amyl derivative) ?
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Mister Double U
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Experiment: Hydrogen Peroxide by Sulfuric Acid Electrolysis
Breathing some fresh air into this threat.
I prepared a ~50% H2SO4 solution and started electrolyzing.
Current: 3 [A]
Cathode: Titanium mesh
Anode: Boron Doped Diamond
Current Density on Anode: ~250 [mA/cm^2]
Here some pictures:
720g ice cubes & 750g H2SO4 (Drain Cleaner):
The setup:
Electrolysis after 2 hours:
Electrolysis after 12 hours:
I was not expecting to see any colors in this experiment. Because of this I assumed that the H2SO4 was contaminated, and I distilled the whole batch.
I also reduced the amount of water in the acid a little, so it should now have ~55%.
Here the residue in the distillation flask:
There is some green color visible in the residue and my guess is that this is Iron(ii) Sulfate. I heard somewhere that sulfuric acid plants use iron
piping, as high concentration acid does not readily attack iron metal. So, this would make sense.
I then restarted electrolysis. To my great astonishment, the same orange color appeared again. After a day, it is as red as the 12hr picture above.
I did find an article about some unspecified Titanium-peroxo complex, which is orange or even red in sulfuric acid:
"pH Effect on the Optical Properties of Peroxo-Titanium Complex"
"Interesting optical properties of peroxo-titanium hydrogen peroxide complex against pH changes were studied for the first time using an analytical
UV-VIS spectrophotometer. A freshly prepared peroxo titanium complex with pH value of 2.21 exhibited an orange color. This color changed to light
orange, cloudy yellow, and then translucent pale-yellow as its pH value increased to 3.9, 5.8 and 6.7 or above. UV-VIS spectra show that the fresh
complex had an absorption band rising at wavelength around 400 nm to higher energy, but had no maximum peak. A new absorption peak appears at 245nm
for the cloudy yellow samples. This is similar to that of colloidal TiO2, which suggests that TiO2 particles might form upon increasing pH value to or
above 5.80. The formation of TiO2 particles was accelerated in the pH value range from 6 to 9, but not in the acidic environment. In H2SO4 acidified
environment, the color turned red-orange instead, and a strong absorption at 397nm was observed only at pH =0.99. Based on the experimental
observations, a model for the color-forming species and possible applications by using the techniques generated from the present study were proposed."
I derive from this, that the Titanium cathode is slowly dissolved and forms this beautiful color with H2O2 present in the solution. I will let the
cell run another day and then attempt to vacuum distill the content - let's hope something is there :-).
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j_sum1
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One thing that may be contributing to the colour...
Ti(III) makes an amber complex with hydrohen peroxide.
Thanks for reviving the thread with your experiment. I am going to have to give it a more thorough read.
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Sulaiman
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but you can buy 12% and concentrate it if required.
CAUTION : Hobby Chemist, not Professional or even Amateur
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Precipitates
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Or even lower concentrations of hydrogen peroxide (e.g., 3-9%) where 12% is a regulated explosives precursor (i.e., the UK).
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Mister Double U
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Update On: Hydrogen Peroxide by Sulfuric Acid Electrolysis - Failure
Today I distilled the first half of cell liquor. I was not able to detect any H2O2 in the distillate :-(.
Here some pics:
Cell liquor after ~3 days of electrolysis - very dark red color:
Etched cathode - lower part which was in the liquid lost thickness:
Max vacuum the pump could reach:
Distillation setup:
Liquid in flask before distillation:
Liquid in flask after distillation:
Before I started the distillation I tested if both the distillation flask and also the receiving flask could hold the vacuum without imploding. This
was done by wrapping them in a thick towel and then placing in a heavy-duty plastic crate.
After the first half of the distillation the liquid in the distillation flask went to almost colorless. During the second half, the temperature was
increased, and the liquid became greenish. Does anyone know what that green color could be?
Anyway, I still got the second half of the batch. I want to wrap the distillation flask in aluminum foil, so less hydrogen peroxide condenses on the
way up. The setup is a basically a very simple rectification column. Also, I want to improve the cooling on the receiving side, so I can distill
quicker (first half took almost 2 hours).
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Mister Double U
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Failure again.
Just distilled the second half with the improvements mentioned before. I also added 100ml of distilled water to the acid in hope I could distill
something over at a lower temperature / concentration. But this did not work. I tested the obtained solution by heating with an oxidized copper wire.
No fizzling was observed. I repeated that with a bought 3% H2O2 and it fizzled vigorously. So, no H2O2 at all.
That leaves us to think about future improvements:
1. Get a greater vacuum. My pump creates a vacuum of 55 cm-Hg which should be 180 mbar. Commercially this distillation is done by lower pressure.
This website (https://chemistrypage.in/hydrogen-peroxide-properties/) states that it is distilled at 26 mm (I assume Hg as it did not state) this would
correspond to ~31 mbar. The website also states that water jet pumps can get down to these kinds of absolute pressures.
2. Get rid of the Titanium contamination by using a different cathode material.
3. Use a diaphragm to reduce reduction at the cathode.
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Mister Double U
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Sodium Bisulfate Electrolysis for H2O2 - Minimal Success
Hello Folks,
Did another run of electrolysis with Sodium Bisulfate in a clay pot.
The anode compartment contained 400 ml and initially 100g of NaHSO4. Then more Bisulfate was added to a total of 200g over the course of a 2-day run.
Current was 2 A and anode was BDD again.
To test how much Sodium Persulfate was created, 50 ml of the anode liquor were taken (spilled the first time) and 50 ml of H2O were added. Into this
21 g of copper wire were dropped and left for a couple hours. After that the solution was heated to drive the reaction to completion. The copper
weighed 18.6 g afterwards - which means 2.4 g were dissolved (0.038 mol). This means that 9 g of NaS2O8 were in the sample. In other words, in the
whole 400 ml anode liquor were 8*9 = 72g. So current efficiency was not very good, but I must mention that I did this in my garage and the cell was at
~40 C.
After this initial test, the remaining 300 ml of anode liquor were vacuum distilled, but this time heated with a water bath. The distillate weighed
118 g. 38 g of this was taken and heated on a hot plate with some copper in it to watch for O2 bubbles. There were some, but a lot less than a 3%
commercial solution. 100g of 30% HCl were added to the remaining 80 g of the distillate, and this was also left with copper wire in it for a couple of
hours. 0.35 g of copper was dissolved. That means there was 0.18 g H2O2 in that sample (considering there was some headspace with air in the beaker
maybe even less).
Anyway, here the pics:
Foaming distillation flask:
Receiving flask with some decomposition bubbles (hard to see):
50 ml of anode liquor + 50 ml of water with copper after some hours:
Distillate heated with copper - some bubbles visible:
Conclusion: Generation of Persulfate somewhat successful for the high temperature, but only traces of H2O2 survived distillation. Again, the pressure
might need to be lowered further to get real amounts of H2O2.
Y'all have a great start into your week :-)
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Belowzero
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I don't have much to contribute at this time but I am very interested in what you are doing.
Keep it up and thanks for sharing this with us! Keep 'm coming.
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Mister Double U
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Thank you, Belowzero!
I think with my current equipment I am almost at an end point here.
One more thing I want to try is to vacuum distill some 3% commercial solution to see if that also decomposes to almost nothing - just to confirm that
it is a pressure problem.
Best greetings!
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Mister Double U
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Persulfuric Acid
Hello,
I thought about the problem of generating H2O2 by electrolysis of Sulfates, and I think it is fair to say that there are 3 major aspects to it:
1. Electrolytic generation of Persulfate.
2. Hydrolysis of Persulfate to Hydrogen Peroxide.
3. Isolation/Distillation to get H2O2 in solution without anything else.
Since my last post, I have done some small experiments regarding the Persulfate generation.
I wanted to see, how effective it would be to make Peroxodisulfuric Acid. I have read somewhere, that this would better hydrolyze to H2O2 than the
peroxo salts. All experiments were conducted at ambient temperatures ~25C.
Here some photos:
Setup - 2l beaker with the anode compartment (400ml) being a flowerpot to prevent cathodic reduction (anode is again BDD).
The concentration of the acid was ~60%; current was 2A at ~5.5V.
Cathode was a copper spiral plated with nickel:
Anode liquor after completed run (~2days):
Interesting is the yellow color of the anolyte. I initially though it might be iron, but the color faded over the course of a day. In other words, I
do not know what it is.
Result:
50ml of the anolyte were diluted with another 50ml of H2O and some Copper was left in the solution until no more reaction took place. The weight loss
of the copper was 0.3g. That means that there were 0.3[g]/63.5[g/mol]*194[g/mol] = 0.91g H2S2O8 in the sample (7.3g in the whole annolyte. This result
seemed not very encouraging.
I did remember though that a couple years back, when I tried to make Sodium Persulfate, that I had better results with lower concentrations of the
salt vs adding more NaHSO4 over the course of the electrolysis. So, after this, I repeated the experiment with ~30% H2SO4. Run time was ~1day.
Results 2:
100ml of the anolyte were taken and a piece of Copper introduced. Weight loss of Cu was 1.44g -> There were 17.6g H2S2O8 in the whole anolyte.
So, in one day the amount produced was more than double compared to the 2-day-run with the 60% acid. Astonishing is that in literature (e.g. The
manufacture of Chemicals by Electrolysis) it is mentioned that concentration needs to be rather higher than lower.
Anyway, I want to do another run with 20% acid and see how that works out.
Best Greetings :-)
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MrDoctor
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Mister Double U, if the vacuum is insufficient, have you considered a low-pressure steam distillation? perhaps the peroxide could be entrained in
something to help it distill over? Also, doesnt this reaction require a divided cell? if so, in leu of a membrane, according to mysteriusbhoice,
encasing an electrode in a "solid electrolyte" can produce results comparable to using a divided cell provided leakage of the oxidation/reduction
products doesnt poison the reaction entirely, it becomes vastly more efficient. ive been meaning to test out my new desktop kiln and try sintering
glass frit loaded with a secondary solid electrolyte like calcium sulfate, onto carbon rods to test it out, as i think neru demonstrated this working
to make HCl by electrolysis on an occasion.
Back to the issue of distillation though, if you cant try anything else at the moment, i have a suggestion. very slowly add an azeotropic entrainer of
water, like toluene, benzene, hexane etc, to your distillate such that it would under vacuum be superheated, it may need to be pre-heated or even
pulled in as a gas also pre-boiling under that vacuum, and you would likely want it to boil as violently as possible that you can manage to still
condense again. the reasoning here is that while peroxide doesnt form an azeotrope, it does co-distill with water. and if you produce a superheated
steam, that vapor will have a sort of relative humidity for the peroxide, below 100%. like how theres a balance of how much water dissolves into air
and increasing temperature raises that value, i would think that while azeotropically boiling peroxide solution normally leaves the peroxide behind,
if the steam produced is of a higher temperature than that azeotropic mixture boils at, at that pressure, it should also then pull out peroxide too if
hopefully there is some convenient overlap.
at 180mbar water should boil at 56C or so, though the presence of sulfuric acid could be complicating things there.
Lastly, if you have any other pumps, as long as the volumetric flow rate is higher than your primary pump, or you include a check valve venting to the
air so no pressure builds up between the two, you can connect them in series to drop pressure further. Some pumps more than others work on a P-delta
basis.
The main pump will output very little air once at vacuum pressures, at which point a second pump helps.
Soon i was planning on testing how low i could go by putting my portable recirculated 12v bucket aspirator (-500mmHg) in series with my fridge
compressor (-720mmHg), when i have some free time to burn again, plus i need to assemble my new digital pressure gauge as its possible my mechanical
ones are no good at extreme vacuum
[Edited on 18-8-2025 by MrDoctor]
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Mister Double U
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MrDoctor,
I was thinking about bleeding some air into the system to have the air carry H2O/H2O2 over. I am just not sure, if I would be able to condense
anything.
Anyway, as this setup would make things much more complicated, I decided to buy a cheap (60 USD on Ebay) rotary vane vacuum pump - should arrive some
time this week.
Not sure how long it will last, but I got a small smelting oven from Vevor and it is of a not too bad medium quality.
Concerning a divided cell, that is what the flowerpot was for. Anode compartment in the pot, cathode compartment outside.
Regarding the idea about a 'solid electrolyte' around the electrode:
This idea will work if a suitable material is found. For example, I am using CaCl2 as an additive in a Chlorate cell to form a Ca(OH)2 diaphragm on
the cathode during the run. In the sulfuric acid case this diaphragm would most likely have to be deposited in a separate run beforehand. Like coating
the cathode in Ba(OH)2 and then converting to the sulfate as it has very low solubility. Not saying it can't be done, but there is some work to it.
Flowerpots on the other hand work great in acidic environments and are cheap. Another idea would be to wrap some PE cloth like Tyvec around the
cathode - I have worked with that before as well (not sure though how it would hold up in peroxo-acids).
Lastly, I do appreciate your ideas, and some of them seem to require a lot of different equipment/technologies. It seems that you have a lot of
equipment at hand. Have you tried running some experiments? E.g. buying some 3% H2O2 and distilling/steam distilling it with your setup? The results
would ceartainly drive this threat forward :-).
Best Greetings!
[Edited on 18-8-2025 by Mister Double U]
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bariumbromate
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yes the vevor pumps get a good vacuum but take awhile to get there (2-5 min)
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Mister Double U
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Thanks for the info, bariumbromate!
I am looking forward to trying it out then :-)
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Mister Double U
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Potassium Peroxodisulfate
Hello,
Last week I decided to prepare some K2S2O8. I am not sure if it will be a good precursor to H2O2, but I wanted to actually see a product to keep my
motivation going.
The electrolysis was carried out in an undivided cell without any additives.
Starting point was a solution prepared from 140g KHSO4 and 350g H2O (28.5%).
I = 2A
T = 30 - 40C (depending on daytime)
t = 24h (duration of electrolysis
Here some photos.
Setup:
After completed run:
Vacuum filtration:
Final product:
After the run was completed, the product was vacuum filtered and washed twice with ice cold water. It was then spread on a piece of paper and left to
dry in ambient air for 3 days. The yield was 81g of a very fine crystalline product, which was almost free flowing and had no odor.
The current efficiency for this run was 33.5%.
To test the purity of the product, 10g were dissolved in 100ml of water. The solution needed to be heated to ~50C for all K2S2O8 to dissolve. Then
dissolution of a piece of Copper was measured as in the previous experiments.
2.4g were dissolved. For that amount of Cu to dissolve, it takes ~10g of Persulfate. In other words, the product was almost pure.
Next will be mixing the remaining product with H2SO4 and vacuum distillaton...
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Mister Double U
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Distillation of K2S2O8 with 30% H2SO4 - Failure
Attempted the distillation of ~70g Potassium Persulfate with ~200ml 30% H2SO4 with my new vacuum pump.
Unfortunately, a lot of decomposition was visible in the boiling flask and just water made it over to the receiver.
Max vacuum the pump could reach was 27 inHg. Which converts to ~100mbar absolute pressure.
The distillation temperature was 65C.
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Mister Double U
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Distillation of a purchased 3% Hydrogen Peroxide Solution
After the previous failures, I wanted to see, if H2O2 can be distilled at 100mbar without decomposition or if the pressure is just not low enough.
Boiling started around 35C. From there the temperature slowly rose to 45C and remained there for the rest of the distillation (+/- 5C, my thermometer
is not very accurate).
It did not look like any decomposition was happening at all. I only saw big 'boiling' bubbles and no small foaming 'decomposition' bubbles.
After the first half distilled over, I heated the liquid with some copper wire.
Gas evolution was visible:
When the rest of the liquid was distilled over, the same test was performed as previously.
Vigorous gas evolution was visible:
Second fraction kept bubbling for several minutes after being taking of the hot plate:
Conclusion: Hydrogen Peroxide can be distilled at this pressure without decomposition. Since Persulfates were clearly created in previous experiments,
the problem lies withing the hydrolysis / distillation (contamination) step.
[Edited on 23-8-2025 by Mister Double U]
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