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Author: Subject: Preparation of potassium hexafluoromanganate(IV) from household items
deltaH
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[*] posted on 19-11-2015 at 20:08


Quote: Originally posted by MolecularWorld  
@MrHomeScientist: I was well aware that the solution could be a fine suspension of manganese dioxide, which is why I tested it with additional peroxide, as was already noted in my original post and later posts. Additional peroxide did not decompose, so it's (probably) not manganese dioxide. I don't mind harsh, though I do mind harsh-and-useless, like blogfast's comment above. I very much appreciate Pok providing detailed explanations and references, in addition to voicing doubt. You and Pok are technically right that the yellow product 'could be anything' as the contaminants 'could be anything', though I've also explained above how unlikely it is that these products are contaminated with anything that could give a colored product. I've been researching it, and the only possibility I've found is if my drain-opener potassium hydroxide was contaminated with lead(II) hydroxide, it could form yellowish lead(II) oxide, but really, how likely is that?

@deltaH: I believe Mn2O3 would also decompose hydrogen peroxide. If so, per my above testing, it's not that.


Not necessarily, let's say that it's some kind of manganese oxide. The way it would decompose H2O2 is by acting as a heterogeneous catalyst. All reactions that happen on heterogeneous catalysts are surface reactions and the surfaces of these catalysts are subject to something called 'poisoning', usually by having some other species adsorb strongly/preferentially onto the surface and so make it inert towards the reagents it's meant to adsorb initially in a catalytic mechanism. It's a kind of passivation if you will.

Now these surfaces can be very sensitive to poisoning and you do have fluoride in there. I would imagine that the fluoride ions would adsorb onto the surfaces of manganese oxides very strongly.

So what you might have made could be a highly porous manganese oxide nanoparticles whose surfaces have been fluoridated and so passivated.

HF is also added at a couple percent strength to inhibit red fuming nitric acid (RFNA) in rockets so as to passivate the metal fuel tanks and prevent corrosion.





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[*] posted on 19-11-2015 at 20:43


For what it's worth, I attempted a very crude experiment which may or may not be relevant to the above theory. A small quantity of manganese(II) sulfate solution was added to dilute ammonia, resulting in a reddish-brown suspension. I'm not going to state what this suspension is, but I will think it loudly. The precipitate was allowed to settle, the supernatant decanted, and the solids divided among two beakers. To one was added 3% hydrogen peroxide. This resulted in much oxygen gas production, and the liquid turning black. To the other was added my dilute HF cleaning solution, which dissolved the brown particles to a faintly pink solution.. When hydrogen peroxide was added to this solution, no gas was evolved and the color did not change.

I'll be performing the reactions more closely relating to the original post over the next day or two.




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[*] posted on 19-11-2015 at 21:18


That's a nice experiment MW. It is interesting that even when dissolved (homogeneous mode), the fluoride binds the manganese ions so strongly that they lose their catalytic action.

It would also be interesting to see if you can passivate the black MnO2 with fluoride against H2O2 decomposition without dissolving it. Bear this in mind for future experiments if you can.

If, however, you can dissolve a fine suspension of MnO2 in excess fluoride cleaner, then I would say it is likely that you indeed have prepared a solution of MnF6(-2).

If that happens, you might want to try to saturate with potassium sulfate or chloride to see if you can precipitate something (exploiting the common ion effect to precipitate a dilute solution). Adding potassium hydroxide solution would not be convincing because one might argue that you simply precipitated an oxide again by raising the pH.

Anyway, lots of strategies and suggestions for investigation.




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[*] posted on 20-11-2015 at 00:01


A black suspension of freshly-precipitated manganese dioxide was prepared. Addition of hydrogen peroxide to this suspension produced much oxygen gas, with no color change. Addition of dilute hydrofluoric acid yielded no obvious reaction, even after sitting for several minutes. Addition of hydrogen peroxide to the manganese dioxide and hydrofluoric acid mixture produced no gas (!), but dissolved the manganese dioxide to a nearly colorless solution. Saturating this solution with potassium chloride gave no obvious reaction.

I must say, I was quite surprised that the peroxide produced no gas when added to the manganese dioxide suspension in dilute hydrofluoric acid, even though it agrees with the assertions I've been making in this thread. It seems fluorine chemistry is more interesting than even I expected.




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[*] posted on 20-11-2015 at 00:49


I'm thoroughly confused by these observations. Since the MnO2 only dissolved upon addition of H2O2 AND no gas was evolved, we must conclude that the manganese's oxidation state was raised and this formed a colourless ion? :o:o:o

If it was a fluoride, could it be MnF6(-) ??

Does anyone have literature on such an ion, is it colourless? Usually Mn(V) complexes are bright blue AFAIK.

If I were you, I'd really focus my efforts to try to crystallise the potassium salt of this mystery colourless ion. Sounds amazing.

[Edited on 20-11-2015 by deltaH]




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[*] posted on 20-11-2015 at 00:59


I couldn't find anything in the literature, though I didn't search very much.
I can tell you what didn't happen:

<strike>MnO2 + 6 HF + H2O2 >>> H2MnF6 + 4 H2O</strike>

...because, if such a simple reaction was possible, under such mild conditions, it would surely have been reported in the literature by now, right? Internet searches for "hexafluoromanganic acid" only turn up a few obscure patents, with little supporting documentation.

Edit: The above was written at 3:00am. I see the errors now.

[Edited on 21-11-2015 by MolecularWorld]




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[*] posted on 20-11-2015 at 01:10


Yes I think we can safely conclude from your observations that you are NOT dealing with MnF6(2-) ions. So far the only possibility I can think of is MnF6(-), which is rather crazy I must admit.

Your equation isn't possible because manganese isn't changing its oxidation state, yet the peroxide is getting reduced??? Not possible and it's not balanced.

You really should try to isolate a crystalline salt of whatever is there. It might be weakly coloured in concentrated form which might give us a clue as to what it is.

[Edited on 20-11-2015 by deltaH]




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[*] posted on 20-11-2015 at 01:19


Quote: Originally posted by deltaH  
]Your equation isn't possible because manganese isn't changing its oxidation state, yet the peroxide is getting reduced??? Not possible and it's not balanced.


It's not balanced? I agree the kinetics don't make sense.

Quote:
You really should try to isolate a crystalline salt of whatever is there. It might be weakly colored in concentrated form which might give us a clue as to what it is.


I added potassium chloride to the solution until some solids remained undissolved. I assumed this meant the solution was saturated with potassium chloride with no reaction, but I suppose it might have been a white precipitate...




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[*] posted on 20-11-2015 at 06:39


Quote: Originally posted by MolecularWorld  
I couldn't find anything in the literature, though I didn't search very much.
I can tell you what didn't happen:

MnO2 + 6 HF + H2O2 >>> H2MnF6 + 4 H2O

...because, if such a simple reaction was possible, under such mild conditions, it would surely have been reported in the literature by now, right? Internet searches for "hexafluoromanganic acid" only turn up a few obscure patents, with little supporting documentation.


What's the H2O2 for? MnO2 is already tetravalent. You are more likely to reduce the Mn4+ to Mn2+.

MnO<sub>2</sub> + 6HF &rarr; H<sub>2</sub>MnF<sub>6</sub> + 2H<sub>2</sub>O

Looks more likely
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[*] posted on 20-11-2015 at 07:05


Quote: Originally posted by deltaH  
Yes I think we can safely conclude from your observations that you are NOT dealing with MnF6(2-) ions. So far the only possibility I can think of is MnF6(-), which is rather crazy I must admit.

Your equation isn't possible because manganese isn't changing its oxidation state, yet the peroxide is getting reduced??? Not possible and it's not balanced.

You really should try to isolate a crystalline salt of whatever is there. It might be weakly coloured in concentrated form which might give us a clue as to what it is.
So crazy it might just be some "unreferenced speculation" that doesn't hold water. :P
Manganese(V) is the trickiest oxidation state to obtain out of all of them 2-7. I highly doubt that it can be achieved simply by adding a tiny bit of HF to the manganese dioxide/peroxide reaction. And, as you pointed out, Mn(V) solutions are known to be an intense blue color, which should be obvious if this was actually a Mn(V) compound by some miracle.




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[*] posted on 20-11-2015 at 08:10


To expand the thinking on some of the possible underlying chemistry, first note that the action of dilute HF on dilute H2O2 may form hypofluorous acid, HOF, in a manner noted by Watts as to the action of dilute H2O2 on dilute HCl forming HOCl. Per Wikipedia ( https://en.m.wikipedia.org/wiki/Hypofluorous_acid ) on pure
HOF says that it is a very unstable at RT and a yellow liquid above -170 C. Its salts, hypofluorites, exist and my speculation is that there could also be yellow (?). We may also consider, akin to basic chlorides and hypochlorites (for example, dibasic magnesium hypochlorite), such related fluorine salts.

Note, if attempting this reaction with concentrated reagents, like in the case of concentrated HCl and H2O2 (see, for example, discussion at https://sites.google.com/site/unusualchemistry/halogen-hydro... ), and assuming parallel products, then in place of chlorine in water, one would expect fluorine water (F2 and H2O, where the intermediate HOF quickly decomposes, releasing O2 and HF fumes). However, there is some discussion, for H2O2 that HF actually acts as a stabilizer (see, for example, discussion by KMnO4 in this old SM thread with a cited source at:
http://www.sciencemadness.org/talk/viewthread.php?tid=14490 and also https://books.google.com/books?id=rpfsCAAAQBAJ&pg=PA156&... ).

Now, if no HOF is created, then a modified Fenton reaction between H2O2 and whatever Manganese ion is first created (in place of the usual Fe(ll) ) could take place with favorable pH. However with HOF, then a related Fenton-type reaction may still proceed, where any created HOF supplants HOCl, for example, as seen in biological systems. Source, see "Fenton chemistry in biology and medicine*" by Josef Prousek available at https://www.google.com/url?sa=t&source=web&rct=j&...;)698ab943000000.pdf&ved=0CB0QFjAAahUKEwjCjvjHqrDHAhVKz4AKHZrKDuc&usg=AFQjCNERAlvJfhkQp1Z7LrFk1zdFyLaPMQ&sig2=QUzbx2NQWIZHre_kitaEbQ . To quote reaction 15 on page 2330:

"For Fe(II) and Cu(I), this situation can be generally depicted as follows [20,39],

Fe2+/Cu+ + HOX → Fe3+/Cu2+ + HO• + X- (15)

where X = Cl, ONO, and SCN. "

In the current context, without or with any HOF, the transition metals Fe and Cu are replaced by Mn and HOX is either HOOH or possibly HOF. Note, this reaction shows how the valence state of the transition metal can increase.

[Edit] I should also note on the possible issue of impurities, I recall reading there may be some Fe in the MnO2.

Also, in the presence of any formed hydroxide radical and the flouride ion, to quote a source:

"Well known examples of this are the reactions of .OH with halide and pseudo-halide ions (X−) where the first product is HOX.− although the measured product is .X2−.

.OH + X− → HOX.−

HOX.− → X. + OH- "

Now, in the case of the flourine radical, F., one reaction I would expect is the formation of F2, but not a gaseous release in dilute conditions, with its subsequent reaction with water:

F. + F.→ F2

F2 + H2O = HF + HOF

And then, various other reactions including the possible Fenton reaction noted above, reforming .OH, or a reaction with H2O per a source above:

HOF + H2O → HF + H2O2

where the hydrogen peroxide could feed a Fenton reaction, or its self-decomposition:

2 HOF → 2 HF + O2

[Edited on 20-11-2015 by AJKOER]

[Edited on 21-11-2015 by AJKOER]
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[*] posted on 20-11-2015 at 09:33


Quote: Originally posted by Boffis  
What's the H2O2 for? MnO2 is already tetravalent. You are more likely to reduce the Mn4+ to Mn2+.

MnO<sub>2</sub> + 6HF &rarr; H<sub>2</sub>MnF<sub>6</sub> + 2H<sub>2</sub>O

Looks more likely

I added H2O2 because the MnO2 didn't dissolve in HF alone. When H2O2 was added to the suspension of manganese dioxide in ~1% HF, it dissolved without producing gas. For comparison, Mn2O3 dissolved in HF without H2O2, also without producing gas. (All this is in my posts above.)

I haven't found any references to the reaction of manganese dioxide, hydrofluoric acid, and hydrogen peroxide, but I did find one obscure reference to the existence of the acid I mentioned above. Note that in my tests, MnO2 didn't dissolve in dilute HF alone.

Chemical News and Journal of Physical Science, Volume 4, 1869, William Crookes.png - 84kB

These additional experiments are causing my HF supply to run low, so I'll be postponing study of this line of experiments until after I do the 'control' test, and attempt to precipitate the originally desired complex, as stated on the first page.




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[*] posted on 20-11-2015 at 11:56


I attempted deltaH's control experiment. The procedure in the original post was performed, with half the reagents, side-by-side with a similar reaction, with ~1.5% sulfuric acid substituting for the hydrofluoric acid. Both reactions appeared similar: no obvious reaction until the peroxide was added, then much gas, followed by lightening of the color of the solution. I noted two differences: the sulfuric acid mix gave a small amount of dark precipitate, which quickly redissolved to give a nearly colorless solution, while the hydrofluoric acid mix gave no precipitate and left a noticeably yellow-orange solution. A picture of the result is below, a video clip of the reactions is available, but filming had to be aborted due to the overproduction of fluoride foam.

control.jpg - 16kB

I also prepared a new batch of !hexafluoromanganate, using the same procedure and half the quantities of reactants specified in my original post. I made one modification to the procedure: all of the reactants were chilled to near 0*C before mixing, to slow any hydrolysis that might take place. The peroxide was also added slowly, in increments and with stirring, to ensure no excess was used.

<iframe sandbox width="280" height="160" src="//www.youtube.com/embed/FjGKL_UTqA0?rel=0" frameborder="0" allowfullscreen></iframe>

This resulted in a yellow solution identical to that produced in my first post. This was quickly added to 500ml acetone, which resulted in a voluminous, light-pink-orange precipitate and a small amount of gas evolution. The pink may not be the true color of the precipitate: there were a few droplets of permanganate on the top of the beaker walls, which mixed in as it was poured, coloring the otherwise yellow stuff pink.

!hexafluoromanganate.jpg - 18kB

There may be a risk of forming acetone peroxide with this procedure. Unfortunately, I couldn't think of any other way to extract the complex before it could hydrolyze. Both the fluoride solution and the acetone were ice-cold. A small quantity of the pink stuff was heated in a flame and did not explode.

The supernatant was decanted and the ~100ml of thin pink slurry placed into two improvised desiccators with calcium chloride, one of which is being kept near 0*C while drying.

[Edited on 20-11-2015 by MolecularWorld]




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[*] posted on 20-11-2015 at 15:49


Quote: Originally posted by AJKOER  
HOF...flourine radical...speculation

Speculation is a very friendly word. This is totally senseless. Fluorine radicals? This would mean that this simple reaction could produce elemental fluorine. HOF is extremely unstable and could never be formed by mixing these chemicals together.

@MolecularWorld: as your MnO2 didn't dissolve in (extremely diluted) HF, you can say that the reference does NOT describe what you have done. So there is no reason to quote this reference. You should find something else that deals with your reaction and isn't 150 years old. And in your video you didn't "isolate" a compound. Isolation means that you make the pure compound and not a slurry of it.



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[*] posted on 20-11-2015 at 16:01


Quote: Originally posted by Pok  
...make the pure compound and not a slurry of it.


I'm working on it. I don't have a centrifuge or vacuum-filtration apparatus, so all I can do is let it settle, and try to remove the liquid by evaporation.

I may also attempt another batch, careful to ensure there is no excess permanganate (tricky, as I also don't want an excess of peroxide) to get a better idea of the color of the precipitate.

Any comments on why the version with sulfuric acid instead of hydrofluoric gave a much lighter colored (almost colorless) product?




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[*] posted on 20-11-2015 at 20:48


Quote: Originally posted by Pok  
Quote: Originally posted by AJKOER  
HOF...flourine radical...speculation

Speculation is a very friendly word. This is totally senseless. Fluorine radicals? This would mean that this simple reaction could produce elemental fluorine. HOF is extremely unstable and could never be formed by mixing these chemicals together.
.....


OK, anytime you are combining three reagents, I would suggest one considers all reaction pairs to explain products.

Also, more from Wikipedia on HOF, to quote:

"Hypofluorous acid in acetonitrile (generated in situ by passing gaseous fluorine through "wet" acetonitrile) serves as a highly electrophilic oxygen-transfer agent.[2][3] Treating phenanthroline with this reagent yielded the previously elusive 1,10-phenanthroline dioxide,[4] more than 50 years after the first unsuccessful attempt.[5]"

Yes, so the presence of HOF is speculative as it is certainly unstable in H2O, reacting to form HF and H2O2. But under dilute conditions, perhaps the equilibrium shift more back to HOF and H2O, as has been reported for the action of dilute HCl on H2O2 forming hypochlorous acid and water, or not as fluorine is a little more unique.

Interestingly, the quote from Wikipedia suggests to me, that perhaps the HOF created in situ (possibly concentrated?) is not so unstable so as not to be able to effect any reaction. [Edit] Or, is it the highly concentrated and unstable H2O2 it forms with water?

Also, if you are forming hydroxyl radicals as transition metals tend do in Fenton-type reactions, in the presence of a halogen ion (including the flouride ion), you can get halogen radicals, which are generally more detectable upon further reacting with another halogen ion. For example:

.Cl + Cl- → .Cl2-

I did edit my post to make it clear that under these dilute conditions I am not expecting gaseous F2 or HF.

[Edited on 21-11-2015 by AJKOER]
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[*] posted on 20-11-2015 at 21:00


I'd venture that the version with sulfuric was reduced to MnSO4 (colourless).

You might have recovered crude MnF2 (see wiki) from the fluoride media. Permanganate is not particularly stable under acidic conditions, so I doubt that's causing your light pink colour. MnF2 is described as being a light pink (looks almost white on photos).

I think what happened in both beakers was roughly similar, except that the counter ion varied, i.e. manganese dioxide oxidised hydrogen peroxide to oxygen, itself being reduced to a manganese(II) salt.




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[*] posted on 20-11-2015 at 21:27


I agree the sulfuric acid mix probably resulted in a solution of manganese(II) sulfate. But, I'd thought manganese(II) sulfate and manganese(II) fluoride were effectively the same color, very pale pink, the color being due to the manganese(II), not the anion. The concentration of manganese is also nearly identical in the two beakers in the above photo, so manganese(II) fluoride would have to be a much stronger and more yellow color than manganese(II) sulfate for that to explain the difference in coloration.

I just looked at my precipitate again. It appears to be turning a darker pink. It's still wet with acetone/water mix, so this suggests to me that is may have actually been (partly or mostly) potassium hexafluoromanganate(IV), but is slowly hydrolyzing to MnF3, which is dark purple-pink. I'm currently trying to arrange for some sort of vacuum filtration, so that I can attempt my acetone precipitation procedure again (ensuring neither permanganate or peroxide are in excess), and then pump the precipitate dry, to see if it's yellow and stays yellow (the complex is supposed to be quite stable when dry).




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[*] posted on 21-11-2015 at 00:59


Quote:
... (partly or mostly) potassium hexafluoromanganate(IV), but is slowly hydrolyzing to MnF3, which is dark purple-pink.


What you describe is not hydrolysis, but a reduction, manganese (IV) being reduced to manganese (III). What was the reducing agent?

Also, your crude product might be a mix of many things, you would need to recrystallise it first before doing anything else.




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[*] posted on 21-11-2015 at 06:14


That was based on the reference cited by Pok here. The pink is darker than any picture of manganese(II) fluoride I can find, and while I agree Mn(III) and Mn(IV) are unlikely to be formed, and permanganate wouldn't last long in the presence of acid and acetone, these are the only explanations I can think of. That or color-changing mystery contaminants. Of course, it's not pure MnF3: Wikipedia states that it too is prone to "hydrolysis", to MnF2, and there would also be some colorless potassium salts.

If my product contains an unstable complex prone to hydrolysis/reduction/decomposition, recrystallization could be problematic. I'll retain the pink stuff for future study, but I'm now focusing my efforts on preparing a new batch of precipitate, and drying it as quickly as possible.

[Edited on 21-11-2015 by MolecularWorld]




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[*] posted on 21-11-2015 at 10:24


Quote: Originally posted by MolecularWorld  
Rather, my baseless assertions have been countered by... baseless assertions. My hypothesis is that the main reason concentrated hydrofluoric acid is used in the literature preparations is to precipitate the complex before it can hydrolyze, but that small amounts of unstable potassium hexafluoromanganate(IV) could form in dilute solution/suspension.


You really haven't got the foggiest idea with regards to what is 'evidence' and 'burden of proof', do you? Nor do you seem to care: too busy putting out triumphant and misleading vids on the UToobs, the preferred medium of the brain dead.

You are beneath contempt, Sir.

Make a note to self: 'must not give up the day job yet'.

[Edited on 21-11-2015 by blogfast25]




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[*] posted on 21-11-2015 at 10:33


Quote: Originally posted by AJKOER  
Its salts, hypofluorites, exist and my speculation is that there could also be yellow (?).
[Edited on 21-11-2015 by AJKOER]


Do hypofluorite salts really exist? All my google-fu on hypofluorites yields some organic, definitely non-salty compounds. And one source claims that HOF does not exhibit any acidity whatsoever, and its "acid" name is a misnomer.

Who is right? Summoning Woelen here, our resident fluorine expert.

(mumbles to herself: Crazy, crazy people. Experimenting with H *cking F like it's nothing...)

[Edited on 21-11-2015 by ave369]




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[*] posted on 21-11-2015 at 10:37


Another borderline-trolling comment from blogfast. Past the one sentence in my first post, I never claimed the videos proved anything. Rather, like the images and text descriptions, they simply serve to provide information for consideration. And since nobody's attempting to replicate my procedure, I thought being able to actually see the reaction take place would be helpful.

As for proof, I'm working on it. There's only so much I can do in a limited amount of time with an improvised setup.

I was wrong to state the products of the reaction without further testing, and I've agreed that my products may not be what I've originally stated. Further berating me over that is pure trolling, and truly beneath contempt.

Edit: blogfast: I'm also confused as to why you quoted me stating my hypothesis, then went on to complain about 'evidence' and 'proof'. By definition, a hypothesis is "made on the basis of limited evidence". Yes, I should have titled this thread "Attempted preparation of potassium hexafluoromanganate(IV) from household items", and shouldn't have declared the identity of my product based mostly on color: get over it.

[Edited on 21-11-2015 by MolecularWorld]




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[*] posted on 21-11-2015 at 14:46


I also am not convinced at all by the results, posted in this thread, but it is interesting to see that no strictly anhydrous conditions are needed.

Tomorrow I'll try with 48% HF, 50% H2O2 and KMnO4. Right now it is near midnight, time to go asleep, but this is a nice Sunday-afternoon project, especially with the bad weather we will have tomorrow ;)

And yes, I will be careful, VERY careful! I'll work in PP test tubes and use micro quantities.

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

I also am not impressed by the results with the use of acetone. You really think that a beast like MnF6(2-) can coexist with acetone? Acetone is too strongly reducing. I think you just made hydrous MnO2 or maybe hydrous Mn2O3.

I bet that if you replace your HF with acetic acid, that your results will be very similar. HF is a weak acid and hence the formation of colorless (or very pale pink) Mn(2+) is going very slowly or does not complete at all. Reduction of permanganate results in formation of hydrous MnO2 or Mn2O3 (the latter is yellow/brown when very finely suspended and not too concentrated) when there is insufficient free acid. Only at low pH (in the presence of excess strong acid) you get the nearly colorless Mn(2+) ion. So, repeat your experiment with dilute acetic acid and then you'll get quite similar results as with dilute HF.

I even do not expect too much of my experiment I will try tomorrow, but it is interesting enough to try it.


[Edited on 21-11-15 by woelen]




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[*] posted on 21-11-2015 at 15:17


Quote: Originally posted by woelen  
Tomorrow I'll try with 48% HF, 50% H2O2 and KMnO4.
Thanks. I'll be very interested to see your results.
Quote:
I also am not impressed by the results with the use of acetone. You really think that a beast like MnF6(2-) can coexist with acetone? Acetone is too strongly reducing. I think you just made hydrous MnO2 or maybe hydrous Mn2O3.

Acetone is used in the published procedures to wash the product.
My acetone precipitate does not decompose peroxide.
As noted above, I do think my product was reduced (to Mn(III)) by prolonged contact with the acetone/water mix.
Quote:
I bet that if you replace your HF with acetic acid, that your results will be very similar. HF is a weak acid and hence the formation of colorless (or very pale pink) Mn(2+) is going very slowly or does not complete at all. Reduction of permanganate results in formation of hydrous MnO2 or Mn2O3 (the latter is yellow/brown when very finely suspended and not too concentrated) when there is insufficient free acid. Only at low pH (in the presence of excess strong acid) you get the nearly colorless Mn(2+) ion. So, repeat your experiment with dilute acetic acid and then you'll get quite similar results as with dilute HF.
At the suggestion of deltaH, I did a 'control' experiment with sulfuric acid, and produced a differently colored product.
I agree a test with a different weak acid would be a better control, and I will try that.

[Edited on 21-11-2015 by MolecularWorld]




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