Sciencemadness Discussion Board

Peroxide titration.

Aurum - 31-7-2013 at 19:00

I'm trying to ascertain the strength of some out of date peroxide solutions (3% - 12%) and am atemting the following.

5H<sub>2</sub>O<sub>2</sub> + 2KMnO<sub>4</sub> + 6NaHSO<sub>4</sub> => 2MnSO<sub>4</sub> + K<sub>2</sub>SO<sub>4</sub> + 5O<sub>2</sub> + 8H<sub>2</sub>O + 3Na<sub>2</sub>SO<sub>4</sub>

The method I'm using is to add a teaspoon of sodium bisulphate to a glass of water and then half a teaspoon of the peroxide. I then want to add a half teaspoon of a solution of permangante that will consume 1% of the peroxide. So number of measures of permanganate = percentage H2O2. I've calculated that a 1% H2O2 solution is 0.293 molar (10g/34) and I think I need a solution of KMnO4 at 0.117 molar to titrate with. I think I need 18g per liter to get this solution.

Can anyone confirm the equation and the calculations?

Thanks in advance,


[Edited on 1-8-2013 by Aurum]

My calculations,

5H<sub>2</sub>O<sub>2</sub> + 2KMnO<sub>4</sub> + 3H<sub>2</sub>SO<sub>4</sub> => 2MnSO<sub>4</sub> + K<sub>2</sub>SO<sub>4</sub> + 5O<sub>2</sub> + 8H<sub>2</sub>O

Molar weights,

H<sub>2</sub>O<sub>2</sub> 34
KMnO<sub>4</sub> 158
H<sub>2</sub>SO<sub>4</sub> 120

Grams H<sub>2</sub>O<sub>2</sub> in 1% solution = 10g
Molarity of 1% H<sub>2</sub>O<sub>2</sub> = 10/34 = 0.294

Molarity required of KMnO<sub>4</sub> = 0.294/5*2 = 0.118
Weight KMnO<sub>4</sub> = 158 * 0.118 = 18.6 grams.

All above in 1L solutions.



[Edited on 1-8-2013 by Aurum]

<!-- bfesser_edit_tag -->[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: merged sequential posts]

[Edited on 2.8.13 by bfesser]

Oopsy_daisy - 31-7-2013 at 23:57

IMO the easiest way to do this is to use battery crud (MnO2) from a zinc carbon battery then set up a pneumatic trough apparatus in which a known volume of hydrogen peroxide has been placed, the decomposition catalyst is then added and the flask closed. The gas reciever is a graduated cylinder. With the ideal gas law you can then calculate the amount of moles of gas in the cylinder and then you can calculate the concentration of your original sample of H2O2.

2 H2O2-> 2 H2O+ O2

amazingchemistry - 1-8-2013 at 18:10

That would of course depend on how closely O2 follows the ideal gas law under the temperature and pressure of interest. In my opinion the easiest and most accurate way is through titrimetry. You can do what you're doing (reducing with KMnO4) as in here: You need to be aware that this particular reaction is susceptible to interferentce by other reducing agents present. Alternatively you could use iodometry, as in here: but you would need sodium thiosulfate and potassium or sodium iodide instead of permanganate. This reaction is slower (unless you have access to molybdate catalysts) but can be used for smaller concentrations of peroxide.

Aurum - 1-8-2013 at 21:07

Thanks for the suggestions but I'm trying to make it as simple as possible - it's for a hair salon!!!

So, what I want is a simple test, I.E.
1, get a glass of warm water.
2, add 1 teaspoon of PH down (NaHSO4).
3, add 1 teaspoon of peroxide.
4, add teaspoons of purple liquid until it goes purple.

Number of spoonfuls added = strength of peroxide in percentage.

BTW, I tried it today and all the bottles appear to be 1% stronger than stated. Maybe that is the way they are manufactured or I've got something wrong in my calculations.


Artemus Gordon - 2-8-2013 at 19:54

I'm a total noob, but the only thing I can find wrong is that under molar weights, you wrote H2SO4 when you meant NaHSO4.

Since your results are close to what you expected, I think it's not unreasonable to assume that peroxides intended for salon use might be intentionally made a bit strong to allow for warehouse time. What I would do is go buy a small bottle of 3% H2O2 at the busiest drugstore I can find (to try to get the freshest stuff). If it comes out to anything like 3-5% you're probably golden.

amazingchemistry - 5-8-2013 at 21:18

Are your results repeatable? Did you repeat them? A random error of 1% is expected. Specially when we don't have an accurate measure of the weights and volumes involved. I don't think you should draw too many conclusions yet.

Aurum - 6-8-2013 at 04:42

I tested a few bottles today and all the unopened bottles appear to be at full strength. A couple of bottles that only had about 20% left were at about half strength and were thrown away. So, overall I think it's fairly accurate.

BTW, it was brand new bottles that are about 1% over the stated concentration. Note, 1% over on a 3% bottle is really 33% error!! However, I'm now sure this is intentional.

Thanks everyone.


mayko - 6-8-2013 at 06:01

Densitometry is an alternative that I've found to work well for this, being slightly less messy in my opinion than gas generation and measurement.