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papaya
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H2O2 decomp. catalysts
Minutes ago I evidenced a spectacular reaction when adding little amounts of CuSO4 and NaCl to 30% hydrogen peroxide - it almost ejected from test
tube, which became HOT! Further tests showed that CuSO4 or NaCl when used alone don't catalyze decomposition of peroxide to any perceptible level,
however when mixed in tiny amounts the volcano ejects. I knew that Cu[(NH3)4](OH)2 is a potent catalyst for this particular reaction, but copper +
chloride? What causes copper to change it's behaviour in the presence of chloride? What other less known catalysts do you know?
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Metacelsus
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It might be the tetrachlorocuprate complex.
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aga
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Justa wild wild guess, but maybe there's a Reaction going on ...
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papaya
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a Reaction is going ??? Cannot be true... 
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papaya
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Another try with CuSO4 + HCl instead of NaCl doesn't work!!! WTF?
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woelen
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This is an interesting find. I'll try it tomorrow (now it is very late over here already). What if you take NaCl + CuSO4 + a very small amount of HCl?
Another thing may be that there was some impurity in your first experiment (maybe in the salt you use). Try to repeat it.
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papaya
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I also think there may be impurity in salts, BUT both salts if used ALONE with H2O2 will not give any decomposition as I said. don't know..
[Edited on 30-6-2014 by papaya]
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papaya
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Woelen, conc. HCl (1ml) + CuSO4 (maybe 50mg) + NaCl (about 50mg) + H2O2(2ml) didn't work! First dissolved CuSO4 into HCL, then added H2O2 - no
reaction. Then NaCl was added - a little decomposition occurred that nearly(some slow bubbling still goes) stopped after all NaCl dissolved. This is
NOTHING compared to NaCl+CuSO4 only.
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papaya
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SHIT, YOU'LL NOT BELIEVE!
New, more clever experiment: about 50-100mgs of CuSO4 and NaCl dissolved in 0.5ml of water. After dissolution 2mls of 30% H2O2 added at once. Rapid
decomposition sets up, the speed is accelerating with the temperature increase (self-heating). I waited for the rapid boiling to set up, then added
about 1ml of conc. HCL to the mixture - reaction ceased at once!
where is my nobel award ? 
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Texium
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I just tried this, and I didn't get a reaction. It was probably because my peroxide was only 3%. I don't have any concentrated stuff right now, I used
up all that I had on other stuff. I found it quite interesting though how when neutralized with Na bicarb it instantly formed a dark brown solution. I
wasn't expecting that at all.
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HgDinis25
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It seems to be a pH dependent decomposition. As soon as you acidified it (using HCl), the reaction stopped.
Or perhaps your H2O2 is oxidising the Copper(II) Chloride (formed by the ions interchange of CuSO4 and NaCl) to Copper(I) Chloride, wich is insoluble
in water, thus decomposing the H2O2. Copper(I) Chloride is insoluble in water but very soluble in HCl. When you added the HCl it might have stopped
any Copper(I) ions to form (no precipitation), thus decreasing decomposition rate. You could try the reaction out using only Copper(II) Chloride.
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Texium
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Quote: Originally posted by HgDinis25  |
Or perhaps your H2O2 is oxidising the Copper(II) Chloride (formed by the ions interchange of CuSO4 and NaCl) to Copper(I) Chloride,
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That would be reduction though
[Edited on 7-1-2014 by zts16]
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Brain&Force
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Whenever I place a copper sulfate crystal in hydrogen peroxide solution (3%) it always decomposes slowly. I thought all transition metals (save for
group 12) decompose hydrogen peroxide.
At the end of the day, simulating atoms doesn't beat working with the real things...
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HgDinis25
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Of course it is, too much into the night to think properly. My apologies.
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papaya
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Few articles on topic:
http://link.springer.com/article/10.1007%2Fs007060170004
https://www.sciencedirect.com/science/article/pii/S016201340...
Don't see why pH influences so much.
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HgDinis25
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Have you tried using Copper(II) alone, instead of the mixture of NaCl and CuSO4? If it gives the same reaction, your riddle is solved.
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papaya
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Cu(II) ? You mean CuCl2? I don't have that, and CuSO4 doesn't react that much (not noticeable) as I Said..
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HgDinis25
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| Quote: Originally posted by papaya | | Cu(II) ? You mean CuCl2? I don't have that, and CuSO4 doesn't react that much (not noticeable) as I Said.. |
I mean Cu(II) Chloride. CuSO4 and NaCl are probably donating the halide ions and the Cu(II) ions required for the rapid decomposition of H2O2.
Actualy, your first link describes the decomposition using Cu(II) and Halide ions. So, you should try the Cu(II) Chloride alone, and see what you get.
Edit: You may want to make the Cu(II) Chloride by reaction of Calcium Chloride and Copper SUlfate. It will make insoluble Calcium Sufate, leaving you
with CuCl2 in solution. But you'll have to use vacuum filtration or you won't be able to filter the Calcium Sulfate.
[Edited on 1-7-2014 by HgDinis25]
[Edited on 1-7-2014 by HgDinis25]
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Texium
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Yeah, I've tried making Cu(II) chloride that way, and it requires quite a bit of evaporation and recrystallization to get it decently pure, partly
because calcium sulfate is still slightly soluble.
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woelen
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I did an experiment with very pure CuCl2.2H2O (analytical reagent grade with foreign metal impurities and foreign anion impurities in the order of
magnitude of 0.01% or less) from Merck.
I prepared a solution of this CuCl2.2H2O in distilled water. I took appr. 200 mg and dissolved this in 2 ml of water. To this I added 1 ml of 12%
H2O2. This starts fizzling immediately, fairly vigorously. I added a single drop of 35% HCl (also reagent grade) and the fizzling immediately stops.
So, indeed, the reaction is pH-dependent. At low pH, the decomposition of H2O2 ceases.
We here have a nice other example of what a combination of chloride and copper(II) can do. Another interesting thing is that copper(II) alone or
chloride alone does not react with aluminium, but if they are combined (e.g. from CuCl2, or a combination of NaCl and CuSO4), then a violent reaction
occurs with metallic Al.
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papaya
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It is interesting to me here to see how a ligand can change the behavior of the cation - what if instead of chlorine we take something else (besides
ammonia which is well known). Any interesting suggestions? I may try urotropine, but it's still a nitrogen base.
[Edited on 1-7-2014 by papaya]
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woelen
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Try with bromide. Do you have NaBr or KBr? If so, then dissolve some of this and add a little CuSO4. This is a nice experiment in its own, it gives
you deep purple complexes of copper(II) and bromide ion:
http://woelen.homescience.net/science/chem/exps/copper_halog...
An interesting experiment is adding H2O2 to these copper complexes.
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papaya
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unfortunately don't have bromide. Nice page, but I thought Cu2+ is capable to oxidize bromide to bromine, not sure.
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Texium
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| Quote: Originally posted by papaya | | unfortunately don't have bromide. Nice page, but I thought Cu2+ is capable to oxidize bromide to bromine, not sure. |
Um, nope. It makes an interesting complex though.
And you can find sodium bromide at the pool supply store if there's one near you. They sell it as hot tub disinfectant, I bought a 4lb container.
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HgDinis25
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I wonder why such decomposition is pH dependent...
It ocurred to me that the cause for reaction stopping could be the increase of Cl anions concentration and that the pH didn't have anything to do with
it. Perhaps using a different type of H3O+ source?
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