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Taoiseach
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Damn I can't wait to know if this method really works 
Just heat the damn stuff somehow... put a few crystals in a testtube and heat over a candle. Or use a spoon over a lighter... be creative 
Btw. the crystals on your last pics look much more like HS!
[Edited on 11-2-2010 by Taoiseach]
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Myfanwy
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im just puryfiying it. i dissolved the 27g in 250ml boiling water,filtrated and cooled down.
no crystalls appeared at 0°C.
then i boiled 125ml away and cooled down. there was even no crystalls.
just after adding a cord the first crystalls settle on the bottom.
did N2H6SO4 sublimates in hot water? that would be a shame^^but would be very strange for a sulfate.
i think in 6 hours the whole glass will be full of crystalls.
then i heat the hydrazine sulfate in a test tube.
-
the HS dont really want to recrystallize. need help!
last time the whole liquid became a white mass, after the impulse when i opened the glass.
now very less crytsalls have formed.
the soltuion is ice cold.
[Edited on 11-2-2010 by Myfanwy]
you can see on the pics.
less than 1g crystallized.
HS dont like me and wants to stay in solution, or it sublimates in the boiling water.
i shouldnt had dissolve it in boiling water  
fail
[Edited on 11-2-2010 by Myfanwy]
 
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entropy51
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According to Merck Index hydrazine sulfate is soluble in 33 parts of water, freely soluble in hot water. So I think that no more than 3.8 gms would
stay in solution in the 125 mL of solution that you ended up with.
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rrkss
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Take a glass rod and scratch the bottom of your flask, maybe that will cause the crystals to crash out of solution usually works for me when I'm
having problems triggering crystalization otherwise reheat the flask gently in a hot water bath and allow it to gradually cool preferably in the hot
water bath without disturbance.
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Myfanwy
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good idea, but i got it now.
found that HS is nearly insoluble in Ethanol at 25°C, so i added about 300ml.
now i have it back!
If you having problems to crystallize the HS from the solution, add some ethanol and all will be fine =D
[Edited on 11-2-2010 by Myfanwy]
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rrkss
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Good idea using ethanol to crash it out of solution.
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Myfanwy
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i throw my HS away!
seems to be very unpure.
i heated this stuff, untill the glass from the test tube melts, but it doesnt melt..
BUT with sodium hydroxide a fuming liquid was prepared, that reacts violently with <25% HCl.
i have to try the whole synthesis again. i made something wrong.
can u guys try it , too?
i will make it again on tuesday.
im very dissapointed now.
sorry i failed it
but im sure hydrazine sulfate can prepared that way!
[Edited on 11-2-2010 by Myfanwy]
[Edited on 11-2-2010 by Myfanwy]
Attachment: hydrazine und HCl.wmv (335kB)
This file has been downloaded 1261 times
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woelen
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I have to agree with the remarks of len2. I have done quite a few experiments with TCCA and in all situations, where acid is added, the free cyanuric
acid is formed and this is only very sparingly soluble in cold water, but much better in warm water.
I'm very sure that indeed you made hydrazine or its sulfate, or even its isocyanurate, but I do not expect the product to be very pure. The reaction
you show in the small video also occurs with other molten/hot liquids containing lots of NaOH. Just mixing conc. NaOH-solution and conc. HCl will give
a very violent reaction as well (neutralisation of the acid and base is very exothermic).
Waht you could try is adding NaOH to the mix of urea and TCCA and trying to distill off dilute hydrazine hydrate. This distillation is somewhat risky
though, because you'll have a very viscous mass which may be hard to boil evenly. Hydrazine hydrate is fairly stable, hydrazine is not. Always assure
that water is present in sufficient quantities, otherwise you might have an explosion (anhydrous hydrazine is said to be very unstable, but I have no
personal experience with the anhydrous compound, I only have hands-on experience with the hydrate and salts of hydrazine).
So I think you can make hydrazine with TCCA and urea, but isolation of this or the sulfate of it is a hard job. The probem is that hydrazine sulfate
and cyanuric acid have similar solubility properties in water. Also on addition of a strong base they behave similarly (both dissolve in strong base).
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Nicodem
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Quote: Originally posted by woelen  | | I have to agree with the remarks of len2. I have done quite a few experiments with TCCA and in all situations, where acid is added, the free cyanuric
acid is formed and this is only very sparingly soluble in cold water, but much better in warm water. |
I had many occasions to use TCCA at my job and I can assure you that it is quite stable in strong acids as long as there is nothing present that can
reduce it or otherwise react with it. This is one of its features along its increased electrophilicity that differs widely it from others typical
haloimides (for example, NCS is known to slowly decompose in strongly acidic media trough autooxidation). Like I said, I used TCCA/H2SO4 solutions
several times and as long as there is no organic substrate present it will not get reduced to cyanuric acid (particularly if there is no water present
to hydrolyse it to mixtures of chlorine oxides). I also used it for electrophilic chlorinations in dichloromethane with methanesulfonic or triflic
acids as catalyst and no unwanted decomposition happens (besides the intended reaction). Maybe you should tell which acid you used, in what solvent,
and what else was there present.
| Quote: | | I'm very sure that indeed you made hydrazine or its sulfate, or even its isocyanurate, but I do not expect the product to be very pure.
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I'm a surprised at all the optimistic replies even though no evidence is presented. My first thought would be that he isolated Na2SO4 decahydrate.
From the kewlish way Myfanwy describes his experiments it is impossible to reach any conclusion, except that his chances of dying from cancer some
time in future are now enhanced, but nothing really about the chemical aspects. Unless he buys a termometer for up to 300°C (for a dozen well spent
euros) and measures the melting point, or does any other qualitative test, there is not really much to say.
Judging from the smell? Take ammonia smell from urea hydrolysis and wishful thinking and there you have it, the perfect hydrazine smell! I'm not
claiming that you can not make hydrazine from urea and TCCA, actually I think it is possible, but taking a scientific approach is more likely to give
positive results. Myfanwy, it is nice that you are so motivated and eager to do experiments, but don't forget to keep it scientific so that others can
benefit from it as well.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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franklyn
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@ Myfanwy
Even if unsuccesful we celebrate your effort. Nothing ventured , nothing gained.
If you know exactly what the outcome of what you are doing is going to be , then
you are not doing science , which by definition is investigation of what is unknown.
Edison performed 17.000 experiments to develop his incendescent lightbulb, his
token remark of which was invention is 1 % inspiration and 99 % perspiration.
Adversity and setback is most of experience.
.
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woelen
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| Quote: | | I had many occasions to use TCCA at my job and I can assure you that it is quite stable in strong acids as long as there is nothing present that can
reduce it or otherwise react with it. This is one of its features along its increased electrophilicity that differs widely it from others typical
haloimides (for example, NCS is known to slowly decompose in strongly acidic media trough autooxidation). Like I said, I used TCCA/H2SO4 solutions
several times and as long as there is no organic substrate present it will not get reduced to cyanuric acid (particularly if there is no water present
to hydrolyse it to mixtures of chlorine oxides). I also used it for electrophilic chlorinations in dichloromethane with methanesulfonic or triflic
acids as catalyst and no unwanted decomposition happens (besides the intended reaction). Maybe you should tell which acid you used, in what solvent,
and what else was there present. |
Under the conditions you mention, indeed TCCA may be quite stable. I
probably was not very clear in my formulation. I wanted to say that in reactions in which TCCA is used as oxidizer in aqueous medium in the presence
of acid that cyanuric acid separates as an insoluble solid. I have not done experiments in non-aqueous solvents and then things might be completely
different. I have read about TCCA being dissolved in acetone, but to my opinion that kind of experiments is too dangerous.
| Quote: | I'm a surprised at all the optimistic replies even though no evidence is presented. My first thought would be that he isolated Na2SO4 decahydrate.
From the kewlish way Myfanwy describes his experiments it is impossible to reach any conclusion, except that his chances of dying from cancer some
time in future are now enhanced, but nothing really about the chemical aspects. Unless he buys a termometer for up to 300°C (for a dozen well spent
euros) and measures the melting point, or does any other qualitative test, there is not really much to say.
Judging from the smell? Take ammonia smell from urea hydrolysis and wishful thinking and there you have it, the perfect hydrazine smell! I'm not
claiming that you can not make hydrazine from urea and TCCA, actually I think it is possible, but taking a scientific approach is more likely to give
positive results. Myfanwy, it is nice that you are so motivated and eager to do experiments, but don't forget to keep it scientific so that others can
benefit from it as well. |
Aren't you a little bit too harsh? I agree with you that things might look somewhat
unscientific, but didn't we all start this way? I actually think that these attempts are nice, even if no practical results are obtained. He shows
interest in chemistry, many other starters just go for the big BOOM and the F1r3 and Sm0ke. The reason why I think some hydrazine was made is because
Myfanwy reported the formation of a silver mirror from AgNO3 and his white solid.
The only real concern I have is that Myfanwy should indeed be more careful with being exposed to the chemicals. If he is working like this with all
his experiments and is exposed frequently, to all kinds of nasty fumes and gases, then he may experience bad health effects in the long run. So, my
advice to Myfanwy is to scale down experiments to test tube size first and when things look promising then scale up to larger batches. Besides that,
using smell for detecting stuff like hydrazine is not the wisest thing to do, it is better to find some chemical reactions which can detect its
presence or even can determine the amount in a quantitative way.
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len2
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| Quote: |
Also there is no need for acidic media
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Right, and where do you think the proton to protonate the chlorinated acid into its reduced form is going to come from, given urea is oxidized to
hydrazine?
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Nicodem
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I'm not sure I understand what you mean, especially the "protonate the chlorinated acid" part. Are you asking where the protons in cyanuric acid come
from? If so, I already mentioned it above, from the reaction with urea which gives monochlorourea, in analogy with the known reaction of TCCA (and
haloimides in general) with amides and sulfamides:
(CO-NCl)3 + 3 H2NCONH2 <==> (CO-NH)3 + 3 H2NCONHCl
The rearrangement however needs the presence of a base according to the classic (likely) mechanistic approach, which is furthermore supported by the
relative stability of monochlorourea in acidic media.
For example, if the rearrangement happens according to the "aza-Favorsky rearrangement", you can check the depicted mechanism at the Wikipedia entry (just change the methylene group neighbouring to the carbonyl for the -NH2 and the -CHCl- group for the -NHCl). The formed
N-aminocarbamate anion is unstable in neutral or acidic media and it decarboxylates to hydrazine.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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len2
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Right, and where do you think the proton to protonate the chlorinated acid into its reduced form is going to come from, given urea is oxidized
to hydrazine?
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Nicodem
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Len, there was no need to arrogantly copy-paste your question without answering my request for less vagueness. Particularly when I already answered it
in all details (bold text part included!).
Just what proton you can not account for? If you write down both reactions that I discussed above (and their sum) you can see that there are no
protons unaccounted for:
N-Chlorination step: (CO-NCl)<sub>3</sub> + 3 H<sub>2</sub>NCONH<sub>2</sub> <==>
(CO-NH)<sub>3</sub> + 3 H<sub>2</sub>NCONHCl (12 hydrogens on both sides)
Rearrangement step: H<sub>2</sub>NCONHCl + OH<sup>-</sup> => H<sub>2</sub>N-NH<sub>2</sub> +
Cl<sup>-</sup> + CO<sub>2</sub> (4 hydrogens on both sides)
-------------------------------------------------------------------------------------------------------------------------
Overall oxidation: (CO-NCl)<sub>3</sub> + 3 H<sub>2</sub>NCONH<sub>2</sub> + 3 OH<sup>-</sup> => 3
H<sub>2</sub>N-NH<sub>2</sub> + (CO-NH)<sub>3</sub> + 3 Cl<sup>-</sup> + 3 CO<sub>2</sub>
(15 hydrogens on both sides)
I counted and recounted and I still can not account for your attitude.
| Quote: Originally posted by woelen | | Under the conditions you mention, indeed TCCA may be quite stable. I probably was not very clear in my formulation. I wanted to say that in reactions
in which TCCA is used as oxidizer in aqueous medium in the presence of acid that cyanuric acid separates as an insoluble solid. I have not done
experiments in non-aqueous solvents and then things might be completely different. I have read about TCCA being dissolved in acetone, but to my
opinion that kind of experiments is too dangerous. |
But cyanuric acid also separates in organic solvents as well, but only after enough TCCA gets consumed in the reaction, so it is not like it is much
different in this respect. This is actually a nice indication of the reaction progress, but only in solvents where TCCA dissolves completely. In
acetonitrile, alcohols and acetone it is well soluble, but in CH2Cl2 it has lower solubility. The only difference is that in water cyanuric acid
precipitates in larger crystals as dihydrate, while in organic solvents it drops out as fine white powder of anhydrous form. Only in reactions carried
in H2SO4 the cyanuric acid side product does not precipitate until you quench the reaction by pouring it over ice - cyanuric acid is soluble in
concentrated H2SO4! And of course in aqueous basic media where it is also soluble.
TCCA in acetone is no problem as long as the media is neutral (no acid or base must be present). You can make an experiment and put a little solution
of TCCA in acetone in some dark place at room temperature and check how much time it takes for a precipitate to form. Obviously when using acetone as
solvent one expects the reaction of TCCA to be much faster with the substrate in comparison to the solvent.
| Quote: | | Aren't you a little bit too harsh? I agree with you that things might look somewhat unscientific, but didn't we all start this way? I actually think
that these attempts are nice, even if no practical results are obtained. He shows interest in chemistry, many other starters just go for the big BOOM
and the F1r3 and Sm0ke. The reason why I think some hydrazine was made is because Myfanwy reported the formation of a silver mirror from AgNO3 and his
white solid. |
Yes, maybe just a little too harsh, but I was taking into account also his other posts in other threads. Kewlish or not, I don't want him to die
poisoned by HCN, H2S, hydrazine or whatever just because he is unable to mature faster. OK, maybe for a teenager it is not easy to get motivated for
science in the absence of such stuff as cyanides and other supermegacool dangerous stuff. I can understand that - I used to be young as well. But
still he should realize that this is an amateur science forum where not all are pleased to read about such adventures and that some members
expect a certain level is maintained. It is all about equilibrium - a little push here, a little push there, and he will get into the desired
direction and maybe become a successful researcher some day in the future (if he survives long enough).
I also suspect there might have been some hydrazine formed there, but obviously it must have been very little or else he would have had less troubles
crystallizing it out before all the Na2SO4.10H2O starts to (this salt is not particularly soluble in cold water). But since he does not report
scientifically we don't even know if he added enough H2SO4 to protonate the eventual hydrazine. He used excess urea and a huge excess of NaOH. I think
it would make more sense starting with stoichiometric ratios to avoid several practical problems and also because there is no need for using any
excesses (I also do not see much point in adding gelatin in the first trials). He also was adding NaOH into the putative monochlorourea solution
instead of the opposite (hydrazine can get oxidized by N-chloroamides, especially in basic media!).
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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len2
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| Quote: |
N-Chlorination step: (CO-NCl)3 + 3 H2NCONH2 <==> (CO-NH)3 + 3 H2NCONHCl (12 hydrogens on both sides)
Rearrangement step: H2NCONHCl + OH- => H2N-NH2 + Cl- + CO2 (4 hydrogens on both sides)
-------------------------------------------------------------------------------------------------------------------------
Overall oxidation: (CO-NCl)3 + 3 H2NCONH2 + 3 OH- => 3 H2N-NH2 + (CO-NH)3 + 3 Cl- + 3 CO2 (15 hydrogens on both sides)
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So now you have protonation of your chloramine to give you the acid in basic aqueous solution.
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Nicodem
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Len, again you do your best to be incomprehensible. There are no chloramines involved. Do you perhaps mean TCCA or N-chlorourea? None of these needs
to get protonated. And which acid do you mean? Cyanuric acid is nearly insoluble in water so it does not go the next steps. So you probably mean CO2?
Well, the reason for that second reaction being written as it is, was in that I was not sure if the decarboxylation was a discrete reaction step or
not, so I wrote the reaction as a neutral media outcome (which was just fine for I was answering your question about protons). Now I know better that
the N-aminocarbamate salt formed during the rearrangement does not decarboxylate under the basic reaction conditions, but only after adding H2SO4 (or
whatever acid of choice). So I was wrong and the second reaction should actually be two discrete reactions:
Rearrangement step: H<sub>2</sub>NCONHCl + 2 OH<sup>-</sup> => H<sub>2</sub>N-NHCOO<sup>-</sup> +
Cl<sup>-</sup> + H<sub>2</sub>O
Decarboxylation step: H<sub>2</sub>N-NHCOO<sup>-</sup> + H<sup>+</sup> <=>
H<sub>2</sub>N-NH<sub>2</sub> + CO<sub>2</sub>
I did a literature search on the topic of the rearrangement of N-haloureas into hydrazines and to my surprise realized that what appeared as a rather
obscure reaction is not that obscure after all. In one instance it was even given a name, the Schestakow rearrangement (or Shestakov; after the author
of the seminal papers). The reaction was found to proceed in three steps, the same ones I mentioned above, that is
N-halogenation/rearrangement/decarboxylation (see J. Am. Chem. Soc., 76, 2572–2574). The N-halogenation reagents most commonly
used are NaClO and NaBrO, but other halogenating reagents can be used as well. The rearrangement of N-chlorourea only occurred when heating under
highly basic conditions (so I was wrong in saying that the reaction might work without using excess NaOH and alkali carbonates are not necessarily
basic enough). Heat is not always used though. The decarboxylation of the so formed N-aminocarbamate salt requires neutralization, so no hydrazine
forms until the final addition of the acid. N,N-Disubstituted ureas also undergo the reaction, so the mechanism appears to be that of a Hofmann
rearrangement and not Favorsky (though this does not prove or disprove that both mechanisms are operating at the same time where possible, which is
however limited only to N-monosubstituted ureas). I could find no examples of rearrangement of N,N’-disubstituted ureas which cannot
undergo a Hofmann rearrangement, but could potentially undergo a Favorsky rearrangement. So either nobody bothered trying the reaction on
N,N’-disubstituted ureas or the rearrangement cannot proceed through a Favorsky-like mechanism. This raises an interesting question and
maybe when I’ll have some free lab time I’ll check the reaction on N,N’-dicyclohexylurea (I have some leftovers from couplings with
DCC). In fact, a lot of work on using this reaction on substituted ureas was done due to the potential industrial use in the synthesis of carbidopa,
so chances are that N,N’-dicyclohexylureas were simply neglected and thus never tried.
I currently do not have the time to acquire and read all the papers (some I do not even have access to), but for those interested, here are the
references (I screened Beilstein, CA and Spresi).
Seminal papers and patent by Schestakow (the papers are in Russian):
DE164755
Zhurnal Russkago Fiziko-Khimicheskago Obshchestva, 35, 1903; 858.
Zhurnal Russkago Fiziko-Khimicheskago Obshchestva; 37, 1905, 5.
Other references for plain urea to hydrazine:
DE729105
J. Am. Chem. Soc., 76, 2572–2574.
Chemistry and Industry (London, United Kingdom),1954, 1452.
Journal of the South African Chemical Institute, 9 1956, 37.
Substituted ureas to substituted hydrazines:
Heterocycles, 12, 1979, 1571-1574.
Tetrahedron, 43, 1987, 891-894.
Indian J. Chem., A19, 1980, 825-828.
Chem. and Pharm. Bull., 31, 1983, 423-428.
Org. Prep. and Proced. Int., 17, 1985, 1-9.
Acta Chem. Scand., B29, 1975, 93-98.
Org. Synth., (1987) 173-182.
J. Chem. Soc. Perkin Trans. 1, 1990, 1319-1329.
Journal fuer Praktische Chemie (Leipzig), 316, 1974, 347-348.
Liebigs Ann. Chem., 1990, 949-952.
Journal of Organic Chemistry, 36, 1971, 1949-1951.
Patents: US5300688, ES2026759, HU9902484, PL162497, PL163125
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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len1
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Now I can agree with some of what you say. The reaction is indeed an analogue of Hoffman rearrangement (I have never seen it referred to by any other
name). This typically occurs with hypochlorites although some chloramines work as well (reagents of type (RR')-NCl are called chloramines of which
chlorimides to which TCCA belongs are a subgroup). (TCCA is not a good chloramine for this reaction for reasons below).
The protonation remark is not see if you can balance hydrogens, but a thermodynamic one. What drives the reaction of TCCA with water (or with amines
in non aqueous media) is the formation of cyanuric acid, which can not happen in aqueous basic conditions. you rightly step away from that now, by
dividing the reaction into two parts, the first part being in non-basic conditions. However this is all quite unnecessary, bacause acidic
conditions, which is what I wrote about originally, speed up the reaction substantially - and in these the Hoffman rearrangemnt of urea doesnt work.
In basic media in which the Hoffman rearrangement of urea is normally carried out cyanuric acid and TCCA decomposes violently. Just add TCCA to NaOH
solution and wait. This accounts for much of what this chap observed.
It is not true to say that hydrazine does not form until the neutralization stage - I can vouch for the fact that hydrazine is present as soon as the
hypochlorite has reacted with the urea. This maybe is due to other reaction pathways.
For all these reasons direct reaction with TCCA is not a good way to hydrazine.
[Edited on 17-2-2010 by len1]
[Edited on 17-2-2010 by len1]
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woelen
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I purchased some hyhdrazine hydrate, it is from Acros Organics. The bottle is labeled "64% hydrazine" and this closely corresponds to N2H4.H2O. The
label also says that the bottle must be stored in a refrigerator between 2 C and 8 C. The last thing makes me a little bit worried. Is this really
necessary?
I have read many MSDS's and they all say that hydrazine hydrate is stable (as opposed to anhydrous hydrazine, which indeed is mentioned as unstable).
Some MSDS's say that it must be stored in a tightly closed container at ambient temperature. What should I believe? Why would Acros suggest storage at
low temperature? I don't like to have chemicals around which are unstable at room temperature.
[Edited on 18-2-10 by woelen]
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len1
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At work we have a whole pile of chemicals that the bottle says need to be stored in that magical 2C- 8C: chlorotrimethylsilane, iodomethane, methyl
borate, butyllythium, to name a few.
They are all air and moisture sensitive compounds, sealed therefore air tight. For a few of these, CH3ClSi for instance I think the temperature
indication is just to keep the pressure down inside the bottle. They will still be stable at 15-25C.
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turd
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I have a plastic bottle of ~20% N2H4 stored at room temperature. It produces a few bubbles and slight pressure if left standing for a year or so. It
says "repackage after 5 years" on the label. Subjectively I'd say it's about as stable as my 35% H2O2 solution, which also produces a few bubbles in a
similar time frame.
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12AX7
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And mix them and you'll get "a few bubbles" in a few microseconds
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woelen
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@12AX7: Actually, I tried that
I mixed 1 ml of 10% H2O2 with 0.5 ml of 64% N2H4 (added drop by drop). Nothing happens when this is done. The liquids mix and that's all.
Next, I added one granule of cobaltous sulfate (appr. 1 cubic mm, one crystal). At once a cloud of steam is produced and no liquid is left in the test
tube.
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watson.fawkes
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| Quote: Originally posted by woelen | | Next, I added one granule of cobaltous sulfate (appr. 1 cubic mm, one crystal). At once a cloud of steam is produced and no liquid is left in the test
tube. |
Why did you choose to use cobalt? Cobalt oxide is used for catalytic oxidation of ammonia to NO2 in
nitric acid plants; it's the only catalyst other than platinum that's been viable in industry. So it doesn't particularly surprise me, but it does
intrigue me, since this is not a gas-phase interaction.
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woelen
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The reason I used cobalt is that this is a good catalyst at high pH for redoxreactions in which H2O2 is involved. In such environments the cobalt is
oxidized to hydrous Co2O3 and this in turns catalyzes the reaction. I have noticed this behavior in many redox reactions with H2O2 and since then it
is my favorite catalyst when I want to try things like this.
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