H2O2 By Electrolysis

What IS easier

You are going to purchase platinium coated titanium anode to make H2O2?? The 35% H2O2 is easy to get at "green" organic food suppliers. Also, a quick search on google will turn up the same type of suppliers that will sell almost any volume. It is one of the few chemicals that is actually being praised for its environmental freindliness as opposed to the panic that is risen about many other reagents.

Hydrogen peroxide from sulfuric acid (aq) electrolysis

The attached image is from some PDF file (don't ask for it as it has no other reference to this than what is shown). The acid is not consumed, and is in water. Anyone care to suggest how this can be turned into a practical procedure? I'd like to try to create a small amount of 90% for monopropellant experiments. Would I need split cell with a salt bridge; is temperature very important; is there a specific voltage range I must use, likewise for current density; what electrode materials are appropriate?

Edit: using graphite electrodes in a single pre-chilled cell with 10% H2SO4 and a supply of DC 12 V @ 5 A, very little bubbling at the cathode, and lots of bubbling at the anode; fast heating and production of nasty corrosive vapours (the vapours corroded the alligator clips holding the electrodes); contamination with grahite particles. So the simple guess doesn't work. The question is, how do I get it to work.

BTW, I diluted the acid with 35% H2O2 instead of water as I'm trying to get more than 35% concentration, so I might as well start from there.

[Edited on 3-3-2005 by Quince]

Thx for the reply.

I found many steps online but I search for something practical that someone else tried and worked with him.

I don't want to put a lot of effort then realize that nothing work

do you know any verified steps ?

Quote: Originally posted by Jstuyfzand

Very interesting, an awesome (Theoretical) yield of 400 grams of H2O2 for 1 kwh. I am just left wondering what would be suitable for the electrodes, it seems like almost anything decomposes the H2O2. Platinum is mentioned, but that is not very budget friendly.

budget?
did you notice the pressure needed ? it is 100 atmos this is the complex part.

Quote: Originally posted by ecos

Quote: Originally posted by Jstuyfzand   Very interesting, an awesome (Theoretical) yield of 400 grams of H2O2 for 1 kwh. I am just left wondering what would be suitable for the electrodes, it seems like almost anything decomposes the H2O2. Platinum is mentioned, but that is not very budget friendly. budget? did you notice the pressure needed ? it is 100 atmos this is the complex part.

Quote: Originally posted by XeonTheMGPony

100atm isn't that much, but 316L ss isn't too cheap.

Quote: Originally posted by ecos

Quote: Originally posted by XeonTheMGPony   100atm isn't that much, but 316L ss isn't too cheap. do you have an idea how to make a system that can stand this pressure?

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 714 times

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....?

I did a lot of reading and research.

my conclusion : H2O2 is very very very hard to be made at home from raw materials

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

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 :-).

Quote: Originally posted by Jstuyfzand

Quote: Originally posted by ecos   ...my conclusion : H2O2 is very very very hard to be made at home from raw materials Which is a shame, because the EU loves making 30% h2o2 regulated!

Quote: Originally posted by Sulaiman

but you can buy 12% and concentrate it if required.

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).

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 :-)

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 :-)

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]

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...

Distillation of K2S2O8 with 30% H2SO4 - Failure

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]

H2O2 by Ammonium Persulfate Hydrolysis & Vacuum Distillation

Hello Folks,

As promised, I finished electrolyzing the NH4HSO4 solution. And it is currently distilling.

Volumen: 500ml
Concentration: 400g/l
Current: 2A
Duration: ~30hrs
Anode: BDD (10cm^2)
Cathode: Titanium mesh wrapped in 2 turns Polyethylene cloth.
(Cathode was etched in 20% HCl for 10 hours before use)

As discussed previously, the Ammonium hydrogen Sulfate was made by boiling freshly distilled H2SO4 with DEF in the correct proportions.
After the electrolysis was finished, the solution was heated to ~50C and kept at that temperature for 6 hours to facilitate hydrolysis of (NH4)2S2O8 to H2O2. Interesting enough, the solution changed color from completely colorless to red/orange again - like in my first experiment. So I assume there is some Titanium contamination. likely from the holder of the anode which is made from Ti. At this stage I take this as something positive as I think it is the Ti-peroxo complex. So, in other words, there is H2O2 in solution :-).
During the 6 hours hydrolysis only very few tiny decomposition bubbles were observed.

As distillation started, no decomposition bubbles were seen - only big boiling bubbles.

Pictures:

Boiling flask:


Temperature at beginning of distillation:


Pressure:


Electrode setup:


Cut PE cloth (believe was a bag a pillow came in):


I'll post about the results once distillation is complete.

Failure Again

Just tested the second ammonium persulfate solution. Please note that this is not the distilled H2O2 but the raw Persulfate solution. If there was any Persulfate ist was only present in trace amounts.

Here the picture of Copper immersed in the solution after some hours (no blue color visible):


I noticed during the electrolysis that a lot of gas was produced at the anode. I wonder, if there was still a lot of un-hydrolyzed Urea in the solution, which got oxidized to N2 and CO2 - that is at least my hypothesis.

Anyway, I switched back to NaHSO4 with the current setup. I had tried the electrolysis of NaHSO4 in a flowerpot cell in a prior run but had too much contamination. This was indicated by a lot of decomposition during the distillation (the red color in terracotta is probably from Iron).

Here a picture of the cell running again. It looks like there is a lot of gas coming from the anode, but most of the bubbles are stationary on the electrode and only few come off here and there. This can also be seen on the liquid surface, as there are very little bubbles there (@2A).


Next steps: I will test the cell liquor by the usual way once complete and then distill with/without prior hydrolysis and/or added Phosphate.

Here the results for the Persulfate content of the NaHSO4 electrolysis:


2.47g of Copper were dissolved. That means there were 10.1g Na2S2O8 in the 100ml sample -> 50.5g in the batch after ~24 hours of electrolysis.

Current efficiency: 23.6%

This is a little bit of a different angle, but I recently found some info on reduction of oxygen with electrochemical catalysis by phthalocyanine-transition metal complexes. As these usually use iron or cobalt (which are both easily obtained) and phthalocyanine (which is made from urea and phthalic anhydride), which make this a theoretically feasible approach for the amateur as well. The main difficulty to hydrogen peroxide is preventing the decomposition to water. Maybe an additive that would form an adduct, such as urea, could be the solution.

https://advanced.onlinelibrary.wiley.com/doi/full/10.1002/ad...
https://en.wikipedia.org/wiki/Oxygen_reduction_reaction
https://en.wikipedia.org/wiki/Phthalocyanine



P.S. The Wiley reference is behind a paywall, so I haven't read the full text.

[Edited on 9/5/2025 by chempyre235]

Here the result of the last distillation attempt - failure:

Before I started distillation, I heated the solution to ~60C (when small bubbles started forming) for a couple hours. Then I let the solution cool and started vacuum distillation.
At the end things became very foamy and some material spilled over into the collection flask. I still tested the distillate and observed some small bubbles, but I attributed them to the spilled-over Na2S2O8.

To exclude contamination as the culprit for poor Hydrogen Peroxide yield, I added commercial 3% H2O2 to the residue from the first distillation and distilled again. At least some H2O2 made it over. In other words, it looks like the hydrolysis step is the problem in this whole endeavor.

Boiling flask after first distillation:


H2O2 from distillation of purchased 3% solution & residue:


[Edited on 6-9-2025 by Mister Double U]

Experiment with Phosphate Addition - Small Success

Hello,

As mentioned, I also wanted to conduct a run and see which effect PO4[3-] has on the whole process.

NaHSO4 electrolysis was carried out as in the previous runs (500ml, 125g, 2A, wrapped cathode). After that 30ml of a solution containing ~8g Sodium Phosphate (mono & di) were added to the cell content. Unfortunately, the solution also contained some Na-EDTA (enema solution - Walmart).

Here the next steps:
1. A white precipitate formed, which I assume consisted mainly of the acid form of EDTA and was filtered off.
2. The volume of the solution was vacuum distilled down to ~250ml.
3. Vacuum was removed, and solution was heated to almost boiling and some foaming started - heat was then cut off and solution left to cool for ~2hrs.
4. 150ml of distilled water were added and vacuum distillation was repeated until nothing came over anymore.

Both the distillates from the first and second vacuum distillation were tested by the usual heat + Copper way and did contain small amounts of H2O2. Interestingly, judging from the bubbling they seemed to contain roughly the same amounts. One would think that the second (after hydrolysis) batch would contain more. I am guessing the solutions contained somewhere <0.5% H2O2. I did not conduct quantitative testing.

Lastly, the addition of PO4[3-] seemed to have helped. Almost no red/orange color was visible during any of the distillations - only a very faint hue during the hydrolysis step. Also, at the end of the distillations things did not get super foamy like in the previous run and the second one could be driven to the point where nothing came over anymore.

Hydrolysis (the cloudiness is actually very fine foam on top of liquid):


Peroxide decomposition on Copper:


At the moment I need to scratch my brain to see if there are any good ideas coming out of it regarding how to proceed. I am thinking about running another Potassium Persufate batch and then try to hydrolyze it in just the right amount of H2O so the KHSO4 can fully dissolve, but K2S2O8 can't. This way one would have a visual indicator of how the hydrolysis proceeds and see when it is complete.

Off track: I have also thought about doing the carbothermic reduction of BaSO4. Maybe the resulting BaO could then be oxidized to BaO2 in an oxidative melt with KClO3?

Anyway, best greetings!

Quote: Originally posted by Mister Double U

Hello chempyre235, I am not sure how easy it would be to make the Phthalocyanine. As you stated, you will need Phtalic Anhydride which doesn't seem to be too easy to make. Wikipedia says that it can be done with a Mercury catalyst in liquid media oxidizing Naphtalene. Alternatively, I read it can be oxidized in gas phase on a Vanadium catalyst. Then you still need to do the reaction & purification which yields the Phthalocyanine. This route does not seem easy to me - not saying it can't be done :-). Thank you for joining the discussion! It would be cool, if someone would start experimenting on the cathodic generation of H2O2 and then in the end we could see which process is better. I read somewhere that a carbon fiber cloth was used as cathode material in an NaOH electrolyte and then O2 was bubbled through it. As far as I remember, no additives were needed, but I might be wrong. Anyway, we'll get this done some time :-D

Quote: Originally posted by Mister Double U

chempyre235, You convinced me, it seems that Phthalic Anhydride can be made by the amateur. In general, I think these would be the required steps/questions to answer to get this to work: 1. Synthesis of Phthalic Anhydride 2. Purification of Phthalic Anhydride 3. Synthesis of Phthalocyanine-metal complex 4. Purification of Phthalocyanine-metal complex 5. Deployment of Phthalocyanine as an electrocatalyst (homogen or heterogene) 6. Choice of suitable electrolyte 7. Choice of electrode materials 8. Type of cell (divided/undivided) 9. Choice of process parameters (Current density, temperature, stirring, concentrations) 10. Dispersion method for Oxygen dissolution into the electrolyte 11. Recovery of product (extraction, distillation, crystallization) And there are probably more things to think about... Anyway, I do not want to rain on the parade, but I am trying to get the persulfate route to work. It would be nice, if someone could offer any help with that. At the moment, I am the only person in this thread running experiments. Comments of the type "you should do this <halfway thought through idea> not that" do not really help IMHO. On the other hand, if you think your proposed process has great potential, why don't you give it a try? I would like to see you succeed!

No, I get it. You've been doing excellent work and have made much headway on the persulfate process. The phthalocyanine process is a whole other animal. Great job sticking to this project! You have the persistence of @semiconductive, who had been working on nickel plating solutions.

Unfortunately, I don't have any kind of laboratory at the moment and so am bound to perform my chemical projects vicariously for the time being. I intend to make phthalocyanine in the future for its electrochromic properties, but it won't be for a while yet.

[Edited on 9/8/2025 by chempyre235]

https://onlinelibrary.wiley.com/doi/abs/10.1002/1521-3773(20020603)41:11%3C1922::AID-ANIE1922%3E3.0.CO;2-T
Thankfully, it does not appear to be a violent decomposition