Sciencemadness Discussion Board - H2O2 decomp. catalysts

H2O2 decomp. catalysts

papaya - 30-6-2014 at 10:36

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?

Metacelsus - 30-6-2014 at 10:54

It might be the tetrachlorocuprate complex.

aga - 30-6-2014 at 12:05

Justa wild wild guess, but maybe there's a Reaction going on ...

papaya - 30-6-2014 at 12:18

a Reaction is going ??? Cannot be true...

papaya - 30-6-2014 at 13:00

Another try with CuSO4 + HCl instead of NaCl doesn't work!!! WTF?

woelen - 30-6-2014 at 14:05

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.

papaya - 30-6-2014 at 14:10

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]

papaya - 30-6-2014 at 14:30

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.

papaya - 30-6-2014 at 14:45

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 ?

Texium - 30-6-2014 at 15:06

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.

HgDinis25 - 30-6-2014 at 17:18

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.

Texium - 30-6-2014 at 18:45

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,

That would be reduction though

[Edited on 7-1-2014 by zts16]

Brain&Force - 30-6-2014 at 20:16

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.

HgDinis25 - 1-7-2014 at 03:41

Quote: Originally posted by zts16

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,

That would be reduction though

[Edited on 7-1-2014 by zts16]

Of course it is, too much into the night to think properly. My apologies.

papaya - 1-7-2014 at 03:49

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.

HgDinis25 - 1-7-2014 at 03:51

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.

papaya - 1-7-2014 at 03:55

Cu(II) ? You mean CuCl2? I don't have that, and CuSO4 doesn't react that much (not noticeable) as I Said..

HgDinis25 - 1-7-2014 at 04:01

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]

Texium - 1-7-2014 at 07:29

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.

woelen - 1-7-2014 at 10:24

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.

papaya - 1-7-2014 at 12:59

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]

woelen - 1-7-2014 at 13:08

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.

papaya - 1-7-2014 at 13:37

unfortunately don't have bromide. Nice page, but I thought Cu2+ is capable to oxidize bromide to bromine, not sure.

Texium - 1-7-2014 at 15:45

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.

HgDinis25 - 1-7-2014 at 15:57

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?

Brain&Force - 1-7-2014 at 17:01

Probably because the hydrogen peroxide is protonated. I guess. That seems probable, but I'm not sure. It also could be because the solution is resistant to oxidation.

vmelkon - 11-7-2014 at 17:13

Quote: Originally posted by Brain&Force

Probably because the hydrogen peroxide is protonated. I guess. That seems probable, but I'm not sure. It also could be because the solution is resistant to oxidation.

You mean when H2O2 gets a H+, it is more stable?

blogfast25 - 12-7-2014 at 04:52

I think H2O2 is much, much less prone to protonation than H2O. H2O2 is a very weak acid: pK = 11.75 (Wiki).

In H2O the oxygen atoms have 2 lone electron pairs available to share with a proton, in H2O2 only 1. [Edit: this is wrong and has been corrected in my post below this one]

For that reason, in watery solution any protons get lapped up by water, leaving little for H2O2, I think.

I could be wrong on that.

In alkaline conditions H2O2 would deprotonate slightly but going by the pKa not much either.

[Edited on 12-7-2014 by blogfast25]

[Edited on 12-7-2014 by blogfast25]

AJKOER - 12-7-2014 at 05:38

In my opinion, part of the reaction may be electrochemical based essential on a half reaction involving oxygen, of which, conc H2O2 is a good source. The NaCl serves both as a good electrolyte and as a supplier of chloride ions for ligand formation side reactions.

The reaction using dilute H2O2 and heat should also progress, but less vigorously.

To test my electrolyte hypothesis, replace NaCl with a small amount of ZnCl2 or NH4Cl being mindful of pH effects (that is, in comparison to NaCl and dilute HCl to form an equivalent pH), and see how the reaction proceeds.

[Edited on 12-7-2014 by AJKOER]

papaya - 12-7-2014 at 06:04

Quote: Originally posted by blogfast25

In H2O the oxygen atoms have 2 lone electron pairs available to share with a proton, in H2O2 only 1.
[Edited on 12-7-2014 by blogfast25]

How is this true?

blogfast25 - 12-7-2014 at 06:39

@papaya:

Look at the electronic structure of water: each hydrogen atom shares a pair of electrons with the oxygen atom in a so-called sigma molecular orbital (bonding orbital). But even after this bonding each oxygen atom still has two electron pairs left that are non-bonding (oxygen has 6 valence electrons).

When a water molecule protonates (to the hydronium ion - H3O+) one of these non-bonding pairs bonds to the extra proton.

However I was wrong above: the oxygen atoms in H2O2 also have two non-bonding electron pairs. My bad.

blogfast25 - 12-7-2014 at 06:56

Quote: Originally posted by AJKOER

In my opinion, part of the reaction may be electrochemical based essential on a half reaction involving oxygen, of which, conc H2O2 is a good source. The NaCl serves both as a good electrolyte and as a supplier of chloride ions for ligand formation side reactions.

I think we can all more or less agree that the tetrachloro cuprate ion is involved here and it would be useful to replace the sodium with ammonium or potassium.

That NaCl is a good electrolyte is irrelevant here: the moment you've got a copper salt in there the solution becomes an electrolyte. But if some electrolytic effect is taking place it is not two half reactions taking place on separate electrodes, requiring an electrolyte to close the circuit.

The effect of pH is truly striking and hard to explain.

[Edited on 12-7-2014 by blogfast25]

papaya - 12-7-2014 at 07:34

The effect of pH may be is due to

H2O2 <=> H+ + HOO-

equilibrium - it's known that even pure peroxide solutions are more stable in acidic solutions than in basic, I also already provided 2 links to relevant papers in this thread, from which it seems there's possibility for many mechanisms working in parallel (including free radical ones).

blogfast25 - 12-7-2014 at 11:02

Quote: Originally posted by papaya

The effect of pH may be is due to

H2O2 <=> H+ + HOO-

equilibrium - it's known that even pure peroxide solutions are more stable in acidic solutions than in basic, I also already provided 2 links to relevant papers in this thread, from which it seems there's possibility for many mechanisms working in parallel (including free radical ones).

It's possible but the Ka of that equilibrium is very, very small: Ka = 10<sup>-11.75</sup> (compare to acetic acid e.g.: pKa = 4.76). Even at pH, say 6, the vast majority of the H2O2 will be as H2O2, not as HO2(-). You can calculate the % of H2O2 present as HO2(-) as a function of pH very easily: it's very small peanuts.

The first deprotonation of H2O2 is actually weaker than the first deprotonation of boric acid!

[Edited on 12-7-2014 by blogfast25]

blogfast25 - 13-7-2014 at 05:22

While Papaya’s find is interesting, it was always unlikely that this would have gone unnoticed by other chemists so far. In fact a superficial google provides plenty references:

http://www.chemicalforums.com/index.php?topic=7170.0

Video showing the decomposition of H2O2 by CuCl2 solution (no NaCl or other chloride):

https://www.youtube.com/watch?v=-Tm2x8iZQCE

An article that presents a partial explanation for the phenomenon and provides activation energies:

http://link.springer.com/article/10.1007%2Fs007060170004#pag...

Atomistry.com article that in fact states that NaCl on its own has a stabilising effect on H2O2, as does very dilute H2SO4 (same source):

http://oxygen.atomistry.com/catalytic_decomposition_hydrogen...

So, not ‘new’ but possibly not yet fully explained either. And no references to pH, other than the H2SO4 claim by atomistry.com...

--------

On a side note, here an interesting link on the oxidation of Rochelle's salt to formate and CO2, using hydrogen peroxide as oxidiser and CoCl2 solution as catalyst:

http://www.nuffieldfoundation.org/practical-chemistry/involv...

It appears to be rather colourful too.

[Edited on 13-7-2014 by blogfast25]

AJKOER - 13-7-2014 at 10:26

Quote: Originally posted by blogfast25

…

--------

On a side note, here an interesting link on the oxidation of Rochelle's salt to formate and CO2, using hydrogen peroxide as oxidiser and CoCl2 solution as catalyst:

http://www.nuffieldfoundation.org/practical-chemistry/involv...

It appears to be rather colourful too.

[Edited on 13-7-2014 by blogfast25]

Actually, in my opinion, it is not so much CoCl2 that is responsible for the heightened power of H2O2 in this reaction, but I suspect, it is the formation of the more reactive singlet oxygen. The latter is formed by the action of H2O2 on HOCl or a hypochlorite in a basic ionic medium. For example, passing Cl2 gas into a cold (4 C) conc solution of NaOH in 30% H2O2.

The hydrolysis of CoCl2 supplies HCl that is turned into Cl2 and water (and HOCl) by the H2O2.

The same mechanism may also occur with this thread's topic of interest. There is a possible manner to confirm the presence of Singlet oxygen, run the reaction in the dark and with small additions of H2O2, is there luminescence?

[Edited on 13-7-2014 by AJKOER]

blogfast25 - 13-7-2014 at 10:32

Quote: Originally posted by AJKOER

Actually, in my opinion, it is not so much CoCl2 that is responsible for the heightened power of H2O2 in this reaction, but I suspect, it is the formation of the more reactive singlet oxygen.

Opinions we're not short of. Evidence however is fairly scarce (and all that really matters in science).

What makes you think singlet oxygen plays a part here?

AJKOER - 13-7-2014 at 10:43

Blogfast:

Here is a reference on singlet oxygen formation http://www.researchgate.net/publication/15390248_Singlet_mol... [Edit] and a quote from the abstract citing a pH link to the formation of singlet oxygen in the presence of chloride ions:

"ABSTRACT A study of the pH profile of the decomposition of aqueous hypochlorite has revealed the evolution (onset at pH 8) of single (1 delta g) molecular oxygen (singlet spin state dioxygen) detected spectroscopically (1268 nm), prior to the appearance of chlorine (onset at pH 5.5). The possible mechanism of the singlet state dioxygen evolution is presented, and the origin of its chloride ion dependence is discussed, especially in reference to chloride ion dependence of singlet molecular oxygen evolution in biological systems. Recent epidemiological analyses of the correlation of human cancer with chlorinated water supplies focus attention on the singlet oxygen mechanisms of DNA lesion formation."

Here is a text quote relating to preparation:

"The generation of singlet molecular oxygen by the chemiluminescent reaction of hydrogen peroxide with aqueous hypochlorite is well known and is now a standard method (5-9)"

The full article may be available at https://www.google.com/url?sa=t&source=web&rct=j&...

Interestingly, the cobalt chloride catalyzed reaction proceeds slowly with just H2O2, but with CoCl2, somehow the oxidation (via oxygen) proceeds more rapidly.

[Edited on 14-7-2014 by AJKOER]

blogfast25 - 13-7-2014 at 13:08

Interesting as that may be, it says sweet FA about any role CoCl2 could play in there. Or CuCl4(2-) for that matter.

I have no idea how that Rochelle's salt oxidation with H2O2 is catalysed by CoCl2 but I could propose [hypothesise] that CoCl2 forms some activated complex with the tartrate anion. That would be as good or bad a guess as invoking singlet oxygen.

We really don't know (yet).

AJKOER - 13-7-2014 at 15:25

Actually, there is some depth on the employment of singlet oxygen in organic synthesis. See, for example, "Singlet Oxygen in Organic Synthesis", (McKerrall, 2011), link: https://www.google.com/url?sa=t&source=web&rct=j&...

Also, "Singlet Oxygen as a Reagent in Organic Synthesis - Handbook of Synthetic Photochemistry", link: https://www.google.com/url?sa=t&source=web&rct=j&...

And, "SINGLET OXYGEN: A REAGENT IN ORGANIC - IUPAC", link: https://www.google.com/url?sa=t&source=web&rct=j&...

But a search on CoCl2 is far less expansive, see, for example, "Handbook of Reagents for Organic Synthesis, Reagents for ...,
http://books.google.com/books?id=3jLhNq1B2xsC&pg=PA671&a...
Philip L. Fuchs - 2011 - ‎Science
Table 1 Cobalt-catalyzed trimethylsilylmethylmagnesium-promoted styrylation of alkyl halides cat [CoCl2(dpph)] + Alkyl–X Ar Me3SiCH2MgCl ArAlkyl ether, ..."

which is not convincing evidence, but possibly more supportive for a singlet oxygen argument.

A well discussed and important reaction involving singlet oxygen in inorganic chemistry occurs in Atmospheric Chemistry. To quote, for example,

"Nitrous oxide diffuses from the earth to the troposphere, where it is gradually introduced into the stratosphere by turbulent diffusion. There it is exposed to ultraviolet radiation that breaks, it into its components, atmospheric nitrogen (N2) and singlet oxygen (o). These single oxygen atoms react further with nitrous oxide to produce significant amounts of nitric oxide (NO). During the chain of chemical alterations that follows, nitric oxide steals an atom from the ozone, combining with it to become nitrogen dioxide (N02) and free oxygen (o2). "

Source: Proc.Nat.Acad.Sci.USA Vol.69,No.9,pp.2369-2372,September1972 , title "Newly Recognized Vital Nitrogen Cycle", by Harold Johnston, link: http://www.ncbi.nlm.nih.gov/pmc/articles/PMC426942/

So little known singlet oxygen, with a short (under 75 minutes) but apparently active life span, is contributing to the depletion of the protective ozone layer that keeping us all alive.

[Edited on 14-7-2014 by AJKOER]

blogfast25 - 14-7-2014 at 05:10

The Nuffield link on tartrate oxidation with H2O2 and CoCl2 provides a tentative explanation for that reaction:

"Cobalt(ll) ions are pink. The hydrogen peroxide initially oxidises the cobalt(II), Co2+, to cobalt(lll), Co3+, which is green. The cobalt(III) bonds to the tartrate ion, allowing the oxidation to take place. The CO3+ is then reduced back to CO2+ and the pink colour returns."