Sciencemadness Discussion Board - Verifying H2O2 Concentrations

Verifying H2O2 Concentrations

gravityzero - 6-1-2014 at 08:45

I've been messing around with Hydrogen Peroxide for the past week or so and testing methods for concentrating and verifying strengths of various concentrations.
I've been using 35% food grade for cleaning things around the house. I also keep 3% around the house for cuts and also use it sometimes during a cold.

I've seen methods on youtube, which use an inverted graduated cylinder filled in water to determine the amount of O2 gas collected from a reaction of H2O2 + MnO2.
From what I've been able to determine, the math works out to roughly every 3.33ml of O2 collected equates to 1% H2O2.

The first test I ran was with 3% H2O2 that has been sitting around the house for at least a few months, but really an undetermined amount of time.
This is store bought brand, within expiration and the label claims a stabilizer is used. I placed 1ml of the 3% in a test tube, followed by small amount of MnO2.
The stopper was quickly placed in the tube and the gas was allowed to bleed through a small tube into the cylinder for collection. The bubbles of O2 were not really large and they progressed into the collection tube at a very slow pace.
After letting this run for a couple hours, just over 9ml O2 was collected. I think this measurement was a success and matched with what could be expected for the materials used.

Feeling confident with the original test, I continued to move forward with testing of a 35% food grade H2O2. This peroxide was purchased around 3 months ago, but has been kept under vacuum and in a freezer.
Again, 1ml of H2O2 was placed in a test tube, followed by the addition of a small scoop of MnO2. What happened next reminds me of what I've seen in a demonstration of the "aladdin's lamp", but I could be wrong.
Immediately this reaction produced copious amounts of O2 gas. I felt a little concerned when securing the stopper, but I pressed on. The tube released a heavy amount of O2 gas into the collection cylinder for a few seconds, and then abruptly stopped.
I don't think any gas escaped the cylinder, but maybe some escaped before I could return the stopper. In all I collected around 75ml of O2 gas, which would equate to ~22% H2O2.
This could be the case, because I have no idea how long the peroxide might have stayed on the shelf at the store; it was purchased locally.

Here are some questions I have and I welcome any comments or suggestions.

One thing I would like to do is limit the speed of reaction when using higher concentrations.
I could use .5ml instead of 1ml, or I could use 1ml H2O2 and dilute it with several mls of distilled water. I don't think I got a good measurement from the second demonstration and I don't trust the determined percentage.

From what I understand, the MnO2 is used only as a catalyst. If that is the case, shouldn't I just add less MnO2 to reduce this vigorous reaction? Seems like the MnO2 could be reused as well, but I could be wrong.
If instead, the MnO2 is being converted during the reaction, it is possible that the low amount of gas collected from the concentrate was due to a sudden absence of MnO2. This would also account for the sudden stalling of O2 gas produced.
I was too shocked after the second reaction that I did not try adding more MnO2, but I probably should have tried.

I would like to try the concentrated reaction again, but do so in a more controlled experiment.
I wish the MnO2 were granular, but all I have access to is in finely powdered form.

Any comments are extremely appreciated.

[Edited on 6-1-2014 by gravityzero]

Zyklon-A - 6-1-2014 at 09:44

Yes the catalyst can be reused, but I don't think you need to, cause if you want the reaction to proceed slowly, just use a crumb of MnO2.
Or another way, if you have the glassware, would be to pour the H2O2 in a pressure equalized addition funnel over a erlenmeyer flask with a two holed stopper, allow the H2O2 to go through slowly then measure the O2 gas generated (over water is easiest), through the other hole.

[Edited on 6-1-2014 by Zyklonb]

gravityzero - 6-1-2014 at 10:05

Thanks for the help and suggestions. Any idea why the concentrated peroxide abruptly ceased to produce O2?

The weaker peroxide solution gradually decreased the amount of O2 produced over time, as would be expected.

Sedit - 6-1-2014 at 10:55

I thought H2O2 could be titrated using a form of iodimetry or however you spell that.

Zyklon-A - 6-1-2014 at 11:27

I think that the strong H2O2, just reacts quickly and then its over, I'm sure if it stops, the reaction is done.
Maybe because the 3% H2O2 does not get hot enough to react quickly, the 35% quickly heats up driving the reaction to so it is over quickly.

gravityzero - 6-1-2014 at 11:43

This crossed my mind as well. The reaction appears exothermic and the test tube was noticeably warm to the touch afterwards.

The speed of this reaction could be leading to bad results. I will probably try this again in the next few days, maybe even this afternoon. I will post my findings.

Once I derive a positive method for testing highly concentrated H2O2, I am going to try concentrating to 50% or more via rotovap.

I might give the titration SEDIT mentioned a try if my results from this procedure do not improve.

Sedit - 6-1-2014 at 11:59

Quote:

H2O2 oxidizes iodide to iodine in the presence of acid and molybdate catalyst.

The iodine formed is titrated with thiosulfate solution, incorporating a starch indicator.
H2O2 + 2 KI + H2SO4 ----> I2 + K2SO4 + 2 H2O
I2 + 2 Na2S2O3 ----> Na2S4O6 + 2 NaI

Reagents

Potassium iodide solution (1% w/v). Dissolve 1.0 grams KI into 100 mLs demineralized water. Store capped in cool place away from light. Yellow-orange tinted KI solution indicates some air oxidation to iodine, which can be removed by adding a 1-2 drops of dilute sodium thiosulfate solution.

Ammonium molybdate solution. Dissolve 9 grams ammonium molybdate in 10 mLs 6N NH4OH. Add 24 grams NH4NO3 and dilute to 100 mLs.
Sulfuric acid solution. Carefully add one part H2SO4-98% to four parts demineralized water.
Starch indicator.
Sodium thiosulfate solution (0.1N).

Weigh to the nearest 0.1 mg an amount of H2O2 equivalent to a titer of 30 mLs (0.06 g of H2O2) using a 5 mL beaker and medicine dropper. Transfer sample to Erlenmeyer flask.
Add to Erlenmeyer flask 50 mL of demineralized water, 10 mL of sulfuric acid solution, 10-15 mLs of potassium iodide solution, and two drops ammonium molybdate solution.
Titrate with 0.1 N sodium thiosulfate to faint yellow or straw color. Swirl or stir gently during titration to minimize iodine loss.

Add about 2 mL starch indicator, and continue titration until the blue color just disappears.
Repeat steps 2-4 on a blank sample of water (omitting the H2O2).
Calculation
Weight % H2O2 = (A - B) x (Normality of Na2S2O3) x 1.7
--------------------------------------...
Sample weight in grams

where: A = mLs Na2S2O3 for sample; B = mLs Na2S2O3 for blank

Sorry, can't help with part 2 of your question - statistics is not my strong point.

http://answers.yahoo.com/question/index?qid=20081211035810AA...

Pyro - 6-1-2014 at 12:04

Hmm, would this not work as H2SO4 will make I2 without the need for H2O2? HCl won't though.

DJF90 - 6-1-2014 at 12:35

The sulfuric acid is diluted, so competing oxidation is nil. Thats a nice titration procedure using ammonium molybdate as a catalyst. Cheers for that, Sedit. The titration procedure I was aware of uses potassium permanganate as titrant, but I don't like it as the Mn species that remains also catalyses the decomposition of peroxide.

gravityzero - 7-1-2014 at 10:36

Thanks for all the help and suggestions. At this point I am going to keep working with this cheap method of estimating concentration.

I may try another method too, but I always try to use what is first available. I have so many chemicals laying around as is and try to avoid total chaos.

I tried the concentrated solution again yesterday, but diluted it first and I used a very tiny amount of MnO2. I will have to meet somewhere in the middle because this time the reaction took all night.
This morning it appears that around 60ml was displaced, which equates to ~18%. It appears the reaction was still going, so I left it to run, but I doubt it does more than this.

I believe this reaction to be more correct than the first, but they will both be around the same. I'm going to try concentrating this solution to 35%, so I will probably rotovap off half of whatever is desired and that should get me in the ballpark.

I think this method is really fun and somewhat accurate for determining the concentration. As a final thought, I've decided that keeping the peroxide in a freezer is probably not a good idea. It appears that some escaped the tightly closed lid.
I am going to try storing only in a refrigerator, under vacuum, and see how this works out.

ScienceHideout - 7-1-2014 at 10:58

Assuming it is just water and peroxide, there is an easy trick I always use to measure exact concentration with little error

I do it all gravimetrically. For example, if I have some ammonium hydroxide and want to know the concentration, I use my 1 mL volumetric flask. First, I get it clean and dry and put it on my milligram balance. I tare the balance, take it out, and add exactly 1 mL of NH4OH. I then put it back on. If my balance reads .993, for example, it is .993 g/cm3. I then simply open up Ye Olde CRC handbook (you could probably look online, too) for a chart that relates density to concentration. There is a gazillion in there for every chemical you can think of!!!

If you are in lieu of a 1 mL volumetric flask and a milligram balance, you can scale it up sacrificing a bit of accuracy. For example, if you have a .01 g balance, use a 10 mL graduated cylinder--- 3 sig figs are what you are after! Hope this helps!

gravityzero - 7-1-2014 at 11:57

Thanks for that quick tip. I've used the scale to get around a few other things in similar fashion. I will say it totally escaped my mind.
Unfortunately I don't have that sensitive of a scale. I have an Ohaus capable of .01g, so I will try the 10ml option.

Thanks ScienceHideout

MrHomeScientist - 7-1-2014 at 12:24

I found the source document for the quote from Yahoo Answers Sedit posted above, from US Peroxide: http://www.h2o2.com/technical-library/analytical-methods/def...

Figured I'd post it as a more authorative source to cite in your work. Thanks for posting it originally, Sedit! I assume the ammonium molybdate is included to speed up the color changes in the iodometry?

This would make a good video for my channel, showcasing different methods of determining peroxide concentration. That question seems to come up quite often.

Edit: 500th post! At last I've become an International Hazard

[Edited on 1-7-2014 by MrHomeScientist]

gravityzero - 7-1-2014 at 12:27

Here is a link to a density conversion calculator specific for H2O2

http://h2o2.evonik.com/product/h2o2/en/about/calculations/pa...

ScienceHideout - 7-1-2014 at 12:45

Quote: Originally posted by gravityzero

Thanks for that quick tip. I've used the scale to get around a few other things in similar fashion. I will say it totally escaped my mind.
Unfortunately I don't have that sensitive of a scale. I have an Ohaus capable of .01g, so I will try the 10ml option.

Thanks ScienceHideout

Cool, let me know how that works out! Again, that is how I verify all my concentrations and it literally takes 2 minutes. Glad to help

blogfast25 - 7-1-2014 at 13:05

Quote: Originally posted by MrHomeScientist

I assume the ammonium molybdate is included to speed up the color changes in the iodometry?

It's the catalyst that speeds up the redox reaction, H2O2 + Iodide === > water + Iodine. That reaction has to be completed before the titration begins.

[Edited on 7-1-2014 by blogfast25]

MrHomeScientist - 7-1-2014 at 13:15

Got it, thanks blogfast. Does that mean you need to wait a few minutes before titrating, or does the catalyst make the conversion (nearly) instant? I've never done iodometry, myself.

MrHomeScientist - 31-1-2014 at 13:13

Another source describing iodometric titration of hydrogen peroxide, from Solvay Chemicals:

Attachment: H2O2 Conc via Iodometry.pdf (78kB)
This file has been downloaded 1127 times

This one is a little clearer to me than the previous procedure, plus it has background and safety information. It also answers my previous question about the catalyst (wait 5 minutes).

As long as the solutions are used soon after they are prepared, is the standardization (Part A) really necessary?

blogfast25 - 31-1-2014 at 14:15

Quote: Originally posted by MrHomeScientist

Another source describing iodometric titration of hydrogen peroxide, from Solvay Chemicals:

This one is a little clearer to me than the previous procedure, plus it has background and safety information. It also answers my previous question about the catalyst (wait 5 minutes).

As long as the solutions are used soon after they are prepared, is the standardization (Part A) really necessary?

Sodium thiosulphate isn't really a primary standard, so it can vary slightly in thiosulphate content from on source to another. Personally I think reagent grade sodium thiosulphate, dried for a few hours in a CaCl2 desiccator, can be used without standardisation. But for really high precision titrations, standardisation cannot be avoided.

Re. the five minutes, I doubt if it really takes that long but as a precaution, sure, have a short break.

The oxidation of iodide to iodine with H2O2 is the subject of a neat chemical clock experiment. To the mix is added some thiosulphate and some starch. The reaction between the thiosulphate and the iodine in much, much faster than the reaction between H2O2 and the iodide. When the thiosulphate is used up, the solution turns black in an instant.

It's really neat because it can be used also to demonstrate the effect of temperature and concentration (of reagents) or a catalyst on reaction rate. I did it with my daughter, which probably explains why she's studying astrophysics

[Edited on 31-1-2014 by blogfast25]

Lambda-Eyde - 31-1-2014 at 16:27

Hydrogen peroxide concentration is most conveniently determined by permanganate titration in my experience.

blogfast25 - 1-2-2014 at 06:50

Quote: Originally posted by Lambda-Eyde

Hydrogen peroxide concentration is most conveniently determined by permanganate titration in my experience.

Yes, that works well. And no indicator needed. But permanganate does require standardising (oxalic acid is one way)

Peroxide Concentration Experiments

MrHomeScientist - 28-2-2014 at 08:31

I tried two methods to determine the concentration of some of my 35% hydrogen peroxide - density, and measuring O₂ evolved from decomposition. The label on the bottle lists this as 35% concentration, but the MSDS says it ranges 35% - 40%. The bottle I used was only recently opened, and has been stored in a very cold fridge. I ended up with quite a difference between my two methods, which you'll see below.

Method 1: Density
I have access to a nice electronic balance that is accurate down to 4 decimal places, as well as class A 10 mL volumetric flasks. I used both to determine the density of my peroxide. Very simply, I filled the flasks to the line and weighed them on the scale. This had to be done pretty quickly because the cold peroxide caused condensation on the flask, which lead to an ever-increasing weight reading. I took a sample of the peroxide and left it out on the bench to warm up, and at several intervals took out 10 mL to make a density measurement (also measuring liquid temperature each time). I then used the online calculator mentioned earlier in the thread to get the H₂O₂ concentration.

I could not get a good HTML table in this post to save my life, so the results are tabulated in the attached spreadsheet at the bottom. For this method, Sample 1 is somewhat off I think because it is outside the temperature range specified on the calculator, so I omitted it when averaging.

Average Concentration: 34.28%

Method 2: Gas Evolution
I set up the following apparatus to capture the gas evolved when decomposing my peroxide: a stoppered 250 mL filter flask with a tube leading from the sidearm into an inverted 100 mL graduated cylinder in a water bath. I wanted to use smaller glassware, but this was smallest flask available at the time.

I started by taking 0.6mL of peroxide (previously calculated to yield ~80mL of gas) and added this to 10 mL of de-ionized water (DI) in the flask. This was to slow down fizzing so I could get the stopper on without losing too much gas. I added a small spatula full of solid KI and quickly and firmly stoppered the flask. Gas generation began immediatly and the solution turned yellow from iodine. It took a few minutes to overcome the water pressure and make its way through the tube, but once it started bubbling I got one big bubble about every 15 seconds. In the picture, you can see a bubble rising in the cylinder. I let this run for 2 hours to ensure gas generation was complete, and repeated one more time with exactly the same setup.

I then calculated the concentration:

*Gas Law (P & R values from STP, V & T measured):
n = PV/RT = (101.3 kPa)(0.088 L) / (8.314 J/K*mol)(294.82 K) = 0.0036 mol O₂

*Stoichiometry, from 2H₂O₂ --> 2H₂O + O₂:
0.0036 mol O₂ * (2 mol H₂O₂ / 1 mol O₂) = 0.0073 mol H₂O₂
0.0073 * (1 / 0.6 mL) * (1000 mL / 1 L) = 12.12 M

*M to % (Equation 10.9 from this site - I used the density calculator for 35% peroxide at measured room temp. of 71 F):
C_% = C_Mm / 10d = (12.12 M)(34 g/mol) / 10(1.1301 g/mL) = 36.46 %

The second run produced exactly the same volume. Thus,

Average Concentration: 36.46%

Attachment: Peroxide Testing Data.xlsx (10kB)
This file has been downloaded 489 times

Conclusion
So, you can see there is about a 2% difference in my results for each method. Possible sources of error:
- Method 1: Condensation on the cold volumetric flasks lead to higher weight. Attempted to fix by weighing quickly.
- Method 1: Leaving the peroxide on the bench to warm up exposed it to light, possibly leading to some decomposition. It was only exposed for ~20 minutes total.
- Method 2: Some gas may have escaped while stoppering, but this should be very minimal.
- Method 2: Large diameter tube meant only one large bubble at a time could escape into the cylinder. After the experiment, there was still some pressure pushing the water down in the tube, but not enough to make it into the cylinder. This means volume measurements were fairly coarse, and that extra bit in the tube may need to be accounted for as well.
- Method 2: Unsure what density to use in last equation. Perhaps I should allow a sample to warm to room temp then measure it directly.
- Method 2: Barometric pressure as reported by local weather was somewhat less than STP (stormy day), so this may influence a bit.
- Method 2: Other math error?

Any thoughts, particularly a review of the math in Method 2, are welcomed.

blogfast25 - 28-2-2014 at 09:29

I think 2 % difference is actually rather good.

For the density measurement I think it would better to do it at 20 C. The VFs are calibrated to that temperature.

Now try the third method: iodometry!

copperastic - 17-3-2014 at 16:17

MrHomescientist what is the ideal gas law? I dont really get it (Ive looked it up).
Thanks

HgDinis25 - 17-3-2014 at 16:38

Quote: Originally posted by copperastic

MrHomescientist what is the ideal gas law? I dont really get it (Ive looked it up).
Thanks

Here you can find some information about the Ideal Gas Law:
http://en.wikipedia.org/wiki/Ideal_gas_law

It's based on the following equation: PV =nRT
P - Pressure | V - Volume | n - number of moles | R - Constant | T - Temperature

It allows you to calculate, for instance, the pressure a determined gas will occupy, or how much moles of gas you have, provided all the other variables.
It's a derivation from Avogadro's findings wich lead to the famous:
V = n . Vm | Vm = 22.4, at STP

Here you have an online calculator wich will help you understand the Ideal Gas Law better:
http://www.chemicool.com/idealgas.html