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Author: Subject: odd reaction of K azide with K permanganate
papaya
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[*] posted on 16-9-2013 at 14:14
odd reaction of K azide with K permanganate


Hello, while I was seeking for an easy quantitative determination method for alkali azides, I conducted an experiment which I didn't find described anywhere to see if permanganate is suitable for azide oxidation and thus also for titration. Described is a first test so sorry - it contains no numbers, was done just to look at it.

A test tube was half filled with water, to this I added a very small crop of KMnO4 - maybe 2-3 crystals, after everything is dissolved you have a pink to red not a strong solution . To this I added some crops of NaHSO4 to acidify solution (also small quantity) and after it dissolved I added a crop of my self-made potassium azide - dissolves immediately. First I noted nothing so I added one more crop, but it was not needed as I saw later. A reaction started lazily but accelerated and started to steadily bubble out a colorless odorless gas - must be nitrogen(a new wet method for pure N2 preparation?). In a minute solution discolored completely so I added more KMnO4 crystals - they also reacted and minutes later - again colorless solution (I took too much KN3). The reaction with permanganate was not really fast in order to use it for titrations, but given the small concentrations I find this normal. Also I added a very small amount of NaHSO4, so it was not really acidic and this can affect the speed(more test are needed to get optimal pH range).

Now the most fantastic part I want you to think about - the gas evolution didn't stop at the discoloration point, but continued steadily- I even waited for few minutes - still a flow of bubbles(though I'm not sure how long I had to wait, so I'm asking)! So may it be that some manganese compound (MnSO4,MnO2 ?) can catalytically decompose azide to nitrogen?? Searching google one finds many references on thermal decomposition catalysts (metal oxides) of alkali azides, but I found nothing about the same in aqua solutions.

I have 2 questions that I hope someone can answer - is the catalytic decomposition of azides in solution by manganese or other compounds known? If yes - please give references. And second - what easy titration method can you advise, does the one I tried to "develop" with permanganate feasible or even known?

Thanks.
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[*] posted on 16-9-2013 at 17:40


would expect the reaction to be
3 KN3 + KMnO4 + 2 H2O --> 4 KOH + MnO2

that would account for the discoloration. Did you see a blackish cloud momentarily form in the solution?
If not, one possibility could be that the MnO2 was forming some sort of soluble complex with azide, which might then be slowly oxidizing additional azide.

Formation of Mn(N3)3 is possible
just speculation

Once the permanganate is reduced downward (to +4 or possibly +3), any further reaction would be rate limited by the hydrolysis of the basic azide ion in water.

[Edited on 17-9-2013 by AndersHoveland]
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[*] posted on 16-9-2013 at 20:44


Quote: Originally posted by AndersHoveland  
would expect the reaction to be
3 KN3 + KMnO4 + 2 H2O --> 4 KOH + MnO2

[Edited on 17-9-2013 by AndersHoveland]


Where on earth did the nitrogen go?
I think you mean

6 KN3 + 2 KMnO4 + 4 H2O --> 8 KOH + 2 MnO2 + 9 N2

Problem is, no precipitate of MnO2 was observed, so manganese is likely soluble as a lesser oxidation state (+2 would account for the discoloration, as it's almost transparent, certainly so under these dilute conditions).


[Edited on 17-9-2013 by elementcollector1]




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[*] posted on 16-9-2013 at 21:26


Hello papaya,

apparently permanganate is even an old method for determining azide, permanganate will oxidize azide to N2 [1] . in the introduction of this article [1] (Microdetermination of chloride and azide by sequential titration, http://link.springer.com/article/10.1007/BF01216881) the author cited many references for different methods of titration. One of these references is the permanganate method, but unfortunately the article is published in 1948 and in an old journal (i can't find this journal online), but this is the reference for the titration of azide by permanganate

J.H. Van der Meulen, Rec. trav. chem. Pays-Bas 67, 600 (1948).

if someone (in the reference section) can help you to get this article.

conclusion: yes the method of titration of azide with permanganate is well known and old.

Dany.
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[*] posted on 16-9-2013 at 23:01


Dany, your answer is quite interesting, but it does not really answer the question of papaya. The strange thing is that the solution keeps bubbling, even when all permanganate is gone and the liquid is colorless already.

@papaya: Please try dissolving some NaHSO4 and add some of your KN3 to it and dissolve this. Allow to stand for a while. Does the liquid start bubbling? It should not.

If not, add a single drop of a solution of KMnO4. This drop will discolor quickly and bubbles will form. When the drop is gone, bubbling still continues with the same speed or is the speed of bubbling decreasing quickly after the color of the KMnO4 is gone?




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[*] posted on 16-9-2013 at 23:13


Woelen,

my reply was for the last section of papaya question:

"does the one I tried to "develop" with permanganate feasible or even known?"

if somone can get the article that i mentionned earlier (J.H. Van der Meulen, Rec. trav. chem. Pays-Bas 67, 600 (1948). Maybe the author give explanations for the continuous bubbling of nitrogen.

Dany.

[Edited on 17-9-2013 by Dany]
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[*] posted on 17-9-2013 at 03:28


Quote: Originally posted by woelen  
The strange thing is that the solution keeps bubbling, even when all permanganate is gone and the liquid is colorless already.
The answer I think may be that the manganese is still in a higher oxidation state (+3 or +4) but complexed to azide so that is becomes soluble. This form of manganese is less reactive, so that it is not able to directly oxidize azide anions, but can oxidize hydrazoic acid. In aqueous solutions of sodium azide there is a small amount of HN3 in equilibrium. In fact, adding NaN3 to water gives off poisonous fumes. As the reaction makes itself for more basic, the equilibrium is shifted further to the left, so that HN3 is even less favorable. So it takes a longer time for all the azide to be oxidized. Just a theory.

Another possible theory may be that a small amount of hydrazine is created in the reaction, and the remaining hydrazine is gradually decomposed by the catalytic action of the manganese ions. Hydrazoic acid is certainly known to produce small amounts of ammonia when it decomposes.

[Edited on 17-9-2013 by AndersHoveland]
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[*] posted on 17-9-2013 at 04:07


@Dany: My excuse for my remark. I should have read your post more carefully.


@AndersHoveland: I have serious doubts about your explanations on manganese. A manganese(III) or manganese(IV) complex most likely would have a very strong color and they also would be very reactive. According to literature, manganese(III) and especially manganese(IV) are very unstable and very reactive, except in the case of the oxides. Coordination complexes of manganese(III) and manganese(IV) are very rare!




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[*] posted on 17-9-2013 at 04:18


Copper +1 is normally unstable in aqueous solution, but can be stabilized by chloride ligands.
Perhaps oxidizing power of Mn+3 is significantly reduced after complexing to azide ions.

Quote:
The existence of two different manganese (III)-cyanide complexes in cyanide solution is detected when manganese (II)-cyanide solutions are oxidized with hydrogen peroxide. A brown solution is obtained first which colour fades to give a nearly colourless solution. Manganese (III) is present both before and after this process.


http://www.tandfonline.com/doi/abs/10.1080/00032717808082224...

[Edited on 17-9-2013 by AndersHoveland]
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[*] posted on 17-9-2013 at 12:38


Hi all, first answering the questions:
@AndesHoveland
"Did you see a blackish cloud momentarily form in the solution? If not, one possibility could be that the MnO2 was forming some sort of soluble complex with azide, which might then be slowly oxidizing additional azide."

@elementcollector1
"Problem is, no precipitate of MnO2 was observed, so manganese is likely soluble as a lesser oxidation state (+2 would account for the discoloration, as it's almost transparent, certainly so under these dilute conditions)."

When KMnO4 first hit the acidified KN3 solution, for a very short time I can see that blackish cloud, which disappears in a second and never returns back - I also think this is MnO2 that forms at places where KMnO4 concentration is high.

@woelen
"Please try dissolving some NaHSO4 and add some of your KN3 to it and dissolve this. Allow to stand for a while. Does the liquid start bubbling? It should not."

No hydrazoic acid bubbles out, except the very first second when I throw KN3 crystals to NaHSO4 solution and everything stops after the dissolution is complete (very fast). I do anything in a relatively large volume of water (diluted solutions are formed) since this is also a safety concern.

@woelen
"If not, add a single drop of a solution of KMnO4. This drop will discolor quickly and bubbles will form. When the drop is gone, bubbling still continues with the same speed or is the speed of bubbling decreasing quickly after the color of the KMnO4 is gone? "

I won't say the reaction speed is much different after the discoloration, but it becomes more and more slow in time after that point .

@Dany - thanks for the reference, I definitely want to have that article(but how?)!

Now some more results from new experiments. I'll describe as short as I can:

To check the ability of MnO2 to react with KN3 I prepare some mgs of it by thermal decomposition of Mn(NO3)2 on the glass and collected that black powder. Two test tubes were filled with water, one was acidified with NaHSO4, to both was added nearly equal quantity of KN3 crystals. Then some of MnO2 was added to both tubes - first to neutral solution. To my surprize the neutral solution started to bubble at a VERY slow rate, which could be seen if you watch to the light source through the test tube, but definitely it was reacting. The acidified one reacted more readily but it's nothing compared to permanganate! Both tests were positive, but the neutral solution stopped to bubble completely at some point, while the acidic one still continued (at a slower rate). I want to point out that the color of the acidic solution especially was not really "white", that I could easily tell the color from the comparision with blank (water), but I cannot describe that very slight coloration - it's close to the color of very very diluted Cr(NO3)3, if you know what I mean. Also the "crude" MnO2 that initially was added became more "fine" powder, that will flow freely in the solution if you shake the test tube, the powder in neutral solution remained nearly the same. Then I decided to add some NaOH to neutral solution and see - nothing happened, no reaction. In all cases I never waited long enough for MnO2 to be digested completely (if it really is wasted and not just the case of catalysis).
Another set of tests with KMnO4: In alkali and acidic solutions.
In alkali solution no reaction was observed with azide and the color remained pink-red, I just want to add that in these conditions even NaNO2 didn't react! In acidic one, as I already wrote, the reaction is really fast, and IF I remember correctly - just after the pink color has gone I could see the same strange color I described above (though I'm not really sure on that). After a longer time I had completely clear "white" solution, which gave a few rare bubbles when shaken.

A test with MnSO4 : To NaHSO4+KN3 water solution few crops of MnSO4 were added - a very short "rush" of bubbles and then nothing. This I can only explain by additional "acid" coming from some hydrolysis, but no further reaction is really observed. Also the color stayed clear "white".

To put short - reaction proceeds well in acidic solutions, both with MnO2 and KMnO4 (MUCH MORE FASTER), some transitional color is observed (well in case of MnO2 I didn't wait long to see if this color passes out, in the case of KMnO4 it MIGHT be just the color of extremely dilute permanganate after discoloration point.. or NOT, these subjective things!), no reaction or coloration is observed with MnSO4 in acidic and neutral solutions (in basic one it gave Mn(OH)2 with no gas evolution).

[EDIT] I just want to ask - how this reaction could have been used for titration, while the reaction still continues at a noticeable rate after discoloration ?




[Edited on 17-9-2013 by papaya]
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[*] posted on 17-9-2013 at 12:52


That paper is available online if you have subscription to the journal:

table of contents of the right volume:
http://onlinelibrary.wiley.com/doi/10.1002/recl.v67:8/issuet...

And the PDF:
http://onlinelibrary.wiley.com/doi/10.1002/recl.19480670802/...




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[*] posted on 17-9-2013 at 13:16


Can I suggest making a solution of manganese(III) cyanide, and then try reacting this with a solution of NaOH and NaN3 ? If what I suspect is true, then there should be slow gradual bubbling, similar to what you observed. My guess is, the higher the pH, the slower the reaction rate.

Woelen, manganese(III) cyanide has no color, apparently. Cyanide and azide are both considered pseudohalogens, their chemistry as ligands is comparable.

[Edited on 17-9-2013 by AndersHoveland]
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[*] posted on 17-9-2013 at 13:31


I don't possess of any cyanide, so impossible.
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[*] posted on 17-9-2013 at 22:19


here's the article, but unfortunately the text is in German.

Dany.








[Edited on 18-9-2013 by Dany]

Attachment: J.H. Van der Meulen, Rec. trav. chem. Pays-Bas 67, 600 (1948).pdf (127kB)
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[*] posted on 18-9-2013 at 07:35


Thanks Dany, I tried to understand what it's written there with an online translator, but not really understood the most important part - they prepare solution of 1M H2SO4 1M H3PO4 and 1M MnSO4, then add permanganate? What for? Shouldn't it yield MnO2 precipitate? Also in the example following that passage he speaks something probably about titration of remaining permanganate with KJ, then the latter with thiosulfate..
I'm lost - I was looking for something easy but this seems to be more complicated and even not reliable(after so many steps precision is lost, not to say the article begins with "it's known that permanganate-azide is not always stoichometric" or like that. :P)
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[*] posted on 18-9-2013 at 09:32


A summary (I am Dutch, but I fully understand German):

Oxidation of azides and HN3 with permanganate in dilute H2SO4 does not result in a reaction with a precise stoichiometry, there are large fluctuations. This is because part of the azide is converted to N2, but also different nitrogen-oxygen species are formed in side-reactions.
Manganese in oxidation state above +2 and below +7 oxidizes azide or HN3 quantitatively to N2. So, one has to bring the manganese in the permanganate to a lower oxidation state, but this has to be done in such a way, that the total oxidation amount (normality) of the solution remains the same. E.g. one can bring down the oxidation state to +4, but then some other manganese is brought up from +2 to +4, such that the total equivalent oxidizer remains the same. With KMnO4 and Mn(2+) only, however, non-satisfactory results are obtained. The hydrous MnO2 formed from these does not react smoothly. A good result was obtained with appr. 1 mole of H2SO4, appr. 1 mole of H3PO4 and appr. 1 mole of MnSO4, all dissolved in the same liter of water (not three separate solutions!). When permanganate is added to this liquid, this produces a clear dark brown solution, which quickly and quantitatively oxidizes azide and HN3.

An example of the use of this reagent:
Take appr. 5 ml of the above mentioned liquid.
Add 10 ml of appr. 5 N (which is 2.5 M) H2SO4
Add 25.00 ml of a 0.100 N solution of KMnO4 (which is 0.0200 M) ----> Be very precise with this! Stir or shake to get the liquids well mixed.
Add 20 ml of an appr. 0.1 N (which is 0.1 M) NaN3 solution ---> Concentration and volume is not needed precisely, but weighed amount of solid must be recorded as precise as possible! So for testing your NaN3 you should take appr. 0.002 mole of NaN3 and dissolve this in appr. 20 ml of water. Measure the weight as accurately as possible.
When the production of gas has stopped after adding the solution of NaN3, then allow to stand for another 5 minutes (stir occasionally)
The solution has excess KMnO4 (appr. 5 ml of an excess if your NaN3 is 100% pure)
Add 5 ml of a 0.5 N solution of KI (which is 0.5 M)
Titrate with thiosulfate and starch until the iodine just disappears.

From this you can compute exactly how much NaN3 you had and this gives a good estimate of the purity of your product.


-----------------------------------------------------------------

As you see, there is the additional step of iodometric determination of the excess amount of KMnO4 and this most likely is done because the brown solution of Mn(x+) with x between 3 and 6 is not as good as self-indicating as permanganate itself.

The use of H3PO4 is indeed very important. It forms a very special complex with manganese(III), which in solid form has a beautiful purple color and in solution is brown. It is called manganese violet.

I now also can understand the continued bubbling of your solution after all KMnO4 has disappeared. If oxo-species of nitrogen are formed as byproducts and you still have azide in solution, then these oxospecies can show secondary slower reactions with remaining azide.



[Edited on 18-9-13 by woelen]




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[*] posted on 18-9-2013 at 12:05


I did the experiment with permanganate and azide as well. My azide is of reagent quality, but the behavior is exactly as what papaye describes.

I took 3 ml of 2 M H2SO4 and added a pinch of NaN3
The NaN3 dissolves, without producing any bubbles.
I added a solution of KMnO4 and immediately swirled to mix the solutions

When this is done, the color of the KMnO4 does not disappear at once, the color slowly fades, and there is weak bubbling. This process takes ten seconds or so, and then suddenly, the bubbling becomes much more vigorous. It is as if after an initial soft reaction a booster is switched on and the reaction proceeds at 10 times speed, but the remarkable thing is that the discoloring of the permanganate does not speed up.
After a few tens of seconds, the color of the KMnO4 is gone and the liquid is clear and colorless. The bubbling goes on for quite some time after this. When the liquid is shaken, the bubbling is somewhat more vigorous.




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[*] posted on 18-9-2013 at 12:43


Thanks woelen for the translation and also for replicating the experiment - now things became more clear. If I understood correctly in the presence of phosphoric acid the reaction mixture remains homogeneous for the Mn+3( why not +4 as in MnO2?) state (by complexing?) and thus everything reacts smoothly down to the +2 state. You mentioned "oxo-species of nitrogen as byproducts", I'm just curious what that could be?
Also one question which I want to be answered precisely - if we take into account the more complicated pathways of oxidation by Mn+7 (compared to Mn+3/+4), but assume we let the reaction long enough that ALL the intermediate products are oxidized to N2 - isn't at the end a stoichiometric amount of KMnO4 used? Because if YES - I have an idea how to overcome the KMnO4 "fast decoloration" problem - by addition of Fe(III) sulphate to KMnO4 solution - it forms dark red complex with azide so we still can monitor azide concentration by "painting" it red(I think the color of permanganate + color of that complex will still appear red)!
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[*] posted on 19-9-2013 at 13:28


Something I overlooked in Bagal (transl.)- "Permanganate cannot be used, because there are three different reaction pathways depending on permanganate/hydrazoic acid ratio"

2KMnO4 + 10HN3 + 3H2SO4 -> K2SO4 + 2MnSO4 + 8H2O +15N2

4KMnO4 + 8HN3 + 2H2SO4 -> 2K2MnO4 + 2MnSO4 + 6H2O +12N2 + O2

2KMnO4 + 6HN3 + 3H2SO4 -> 2K2SO4 + 2MnSO4 + 6H2O +9N2 + O2

Note that some oxygen can be evolved in the total reaction, so this method is useless even for gasometric determination!
However it is also mentioned, that the oxidation with nitrous acid is quantitative and a method described - "to the solution of NaN3 an excess of NaNO2 is added, the mixture is transferred into the excess amount of sulfuric acid, then the excess of the acid is determined by a standard acid-base titration with 0.1N NaOH (in the presence of phenolphtalein), with 1% accuracy."

Seems there's no possible one-step way to go, but at least the nitrite route or some modification of it (like determination of the excess nitrite by permanganate if I have no pH indicators?) are most promising.
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