Here are the results of my little investigation on H2O2's decomposition during boiling to concentrate it,
I did this because people including myself seemed not to agree on the importance of decomposition.
The investigation lacks precision in the seconth part but should anyways give a good hint.
I measured (100 ± 1) mL 3% pharmacy H2O2 and found its weight ; (100,99 ± 0,01)g. It gives a
density of 1,01 g/cm³ which is what you would expect so boiling 90% of it down should give roughly
30% H2O2. ( The 3% H2O2's density should be 1,0095 at 15°C, let's also not forget that there are a
few drops of acid stabilisants including phosphoric acid in it...).
I put (1000 ± 50) mL of it in an erlenmeyer and boiled it down in a distillation apparatus (I wanted to
analyze the distillate and keep the acid vapors of the stabilisants under control) until it reached 100
mL. I did observe small bubbles of decomposition from the bottom of the erlenmeyer while it was still
hot. I cooled it to 22°C measured it, 98,9 mL and measured it's mass; 109,91 g. It gives a density of
1,10 which is around 27,5%.
I sadly did not have the correct equipment to be very precise, but it still gives a good rough idea.
Considering 3-4 hours of boiling were necessary, I find this is an acceptable decomposition for the
time saved from bubbling air through it slowly. The results seemed better the first time (2 months
ago) I did it in a beaker without the distillation apparatus. I guess it is because the apparatus
increases a little the pressure and let some drops of water go back in the erlenmeyer so it is less efficient.
Anyways that was my little investigation, I would have liked to present more meaningful results but
it's still better than nothing. If someone has better equipment and want to investigate and share his
results, go for it! I know Zyklonb thought about investigating it, give us some news if you do.
[Edited on 8-2-2014 by alexleyenda]Zyklon-A - 8-2-2014 at 19:52
Yes I will be investigating this, I'm quite surprised that your results were positive twice in a row.
I haven't had much time to work in the lab all week. But hopefully I'll be able to find the time to do it tomorrow or the next day. I have a good
scale, but I broke my only graduated cylinder, however, I can easily decompose it and measure the oxygen gas generated. alexleyenda - 9-2-2014 at 07:40
Yes I will be investigating this, I'm quite surprised that your results were positive twice in a row.
I haven't had much time to work in the lab all week. But hopefully I'll be able to find the time to do it tomorrow or the next day. I have a good
scale, but I broke my only graduated cylinder, however, I can easily decompose it and measure the oxygen gas generated.
I wonder how you'll measure the gas released if you don't have a graduated cylinder :p a burette would also do but you probably don't have one either
as you would have used it instead of the cylinder to measure volumes.
[Edited on 9-2-2014 by alexleyenda]Zyklon-A - 9-2-2014 at 08:04
I have a test tube that I measured and marked with the graduated cylinder before it broke, so it will work, I will also buy a new graduated cylinder
soon...alexleyenda - 9-2-2014 at 11:20
I tried it, and I'm surprised and happy of the results!
I measured 0,22g of H2O2, put it in a test tube, put the test tube almost horizontally, put MnO2 in it, on top and made this setup:
I put the test tube straight so the MnO2 fell in the H2O2 and the reaction occured.
I lowered the cylinder in the water to equalize the pressure and noted that I recovered 22,5 mL of O2 in the graduated cylinder. After calculations
involving the perfect gas law (taking for granted the pressure would be 101,3 kPa) and some conversions, I get a concentration of roughly 29% rounded
up.
Once again many factors could have caused failures (the low precision of the scale vs the measures, the reaction is exothermic and can cause H2O vapor
to increase the pressure, the pressure is not equal everywhere...). However, according to the two results that are quite close considering my limited
precision, the concentration of the H2O2 is most probably from 27% to 29%.blogfast25 - 9-2-2014 at 13:33
Nice experiment.
And nice clamps: enjoy the new shininess because it won't last long, if you use them frequently! alexleyenda - 9-2-2014 at 14:48
I don't doubt it! The ring is already quite tarnished but the worst thing is certainly my "hotplate". I used it only one month and well... It looks
like it's not supposed to withstand conc. H2SO4, H2O2, KClO3, Iodine...
(Yes, I washed it)
It's hard to prevent overflow when the boiling is too intense like with H2SO4 and because the temperature control sucks. I use it with a circle I cut
in a frying pan to distribute the heat evenly. It cost me 30$ and it does the job of a cheap 90$ lab plate !
safely boiling down H2O2 to decompose it?
quantumcorespacealchemyst - 11-12-2014 at 02:29
how does one boil down a solution of H2O2 so that is all decomposes without a runaway decomposition? in the case of small amounts of H2SO4, [H2SeO4 in
this case specifically] the solution can hold a higher temperature when it boils down mostly, although the H2O2 i guess does that anyway (giving the
liquid the ability to reach a higher temperature, the more concentrated it becomes).
so if boiling down to 30%, how does one decompose it safely towards that point and if starting from that point without adding contaminants (by heating
method alone and/or other treatment)?
specific instance, reacting selenium in H2O2 to make selenium doixide which reacts to make selenic acid. starting with a 3-5%H2O2 or a 30-35%H2O2
concentration with dissolved HSeO4, how does one decompose the remainder so the acid can be obtained pure?
[Edited on 11-12-2014 by quantumcorespacealchemyst]CuReUS - 11-12-2014 at 05:26
