chornedsnorkack
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Permanganic acid strength, stability, solubility
Permanganic acid is well known to be weaker than concentrated sulphuric acid. And while it has limited solubility in concentrated sulphuric acid, it
is not miscible. Contact of solid permanganates with concentrated sulphuric acid is liable to precipitate dimanganese heptoxide.
What is the concentration necessary to cause thus precipitation? Would moderately concentrated sulphuric acid, like 80 %, also form Mn2O7, or would
HMnO4 stay in solution till excess KMnO4 remains solid as KMnO4?
Also, is concentrated HNO3 strong enough acid to separate Mn2O7?
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blogfast25
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Precipitate? Mn2O7 is an oily liquid. And it's very dangerous. I'd stay well clear if I were you.
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chornedsnorkack
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It is not miscible with liquid sulphuric acid (and also happens to be denser). Is there a different term for separation of an immiscible liquid from a
solution?
Yes. In which case, it also is important to stay on the desired side of the solubility limits, and know what they are.
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blogfast25
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KMnO4 is used in acid conditions all the time, w/o problems. I doubt very much that anything less than 80 % H2SO4 will liberate the heptoxide. Could
be wrong on that, of course. Textbooks, as far as I can tell, always mention conc. H2SO4 in this context. Too much water and the Mn will stay as
dissociated 'HMnO4'.
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Dan Vizine
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You won't get it with 80% H2SO4.
It is kind of nice, in small amounts, to get carbon out of glass frits.
"All Your Children Are Poor Unfortunate Victims of Lies You Believe, a Plague Upon Your Ignorance that Keeps the Youth from the Truth They
Deserve"...F. Zappa
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blogfast25
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Interestingly, I've just come across a Brauer entry for Mn(III) sulphate and it's pretty bonkers. Page 1467 of Vol. 2 (Ed. 1963) is basically a
controlled demolition job on Mn2O7. The heptoxide is prepared by dropping powdered KMnO4 into conc. H2SO4 and then applying controlled heating (danger
of explosion is elaborated on briefly) to decompose the heptoxide to Mn2O3, which then crystallises as the Mn(III) sulphate! (green crystals,
apparently)
Explosion proof fume hood strongly recommended! 
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chornedsnorkack
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Does Mn(SO4)2 stay safely in solution at 80 % H2SO4?
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blogfast25
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Did you mean Mn2(SO4)3 (Mn(III) sulphate)? If so, according to the method presented in Brauer, it crystallises out on cooling.
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chornedsnorkack
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Quote: Originally posted by blogfast25  |
Did you mean Mn2(SO4)3 (Mn(III) sulphate)? If so, according to the method presented in Brauer, it crystallises out on cooling. |
No, I did mean Mn(SO4)2.
When permanganate is reduced, under acid conditions H2MnO4 and H3MnO4 are unstable to dismutation. MnO2 is not. So, when small amounts of carbon are
oxidized by an excess of HMnO4 or MnO3HSO4, does removing C precipitate merely produce MnO2 precipitate instead, or is Mn(SO4)2 fully soluble in
sulphuric acid under these conditions?
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blogfast25
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I've no knowledge of Mn(IV) salts, only MnO2. The latter does dissolve slowly in conc. H2SO4 and it yields Mn(III) sulphate, not Mn(IV) sulphate.
There's a thread on it somewhere, by me and woelen.
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chornedsnorkack
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| Quote: Originally posted by blogfast25 |
Did you mean Mn2(SO4)3 (Mn(III) sulphate)? If so, according to the method presented in Brauer, it crystallises out on cooling. |
Then does a cold 80 % sulphuric acid solution of permanganate leave Mn2(SO4)3 precipitate behind wherever it is reduced?
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Arcuritech
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The only way I know that produces clean (metal cation free) HMnO4(aq) is to hydrolyze pure Mn2O7 in room temperature
water.
edit: HMnO3(aq) corrected to HMnO4(aq)
[Edited on 2014-12-4 by Arcuritech]
"If we knew what we were doing, it wouldn't be called research." -Albert Einstein
"There are few things -- whether in the outward world, or, to a certain depth, in the invisible sphere of thought -- few things hidden from the man
who devotes himself earnestly and unreservedly to the solution of a mystery." -Nathaniel Hawthorne ("Roger Chillingworth")
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WGTR
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I use Mn2O7 as prepared below to oxidize graphite:
To study the effect of oxidation level we used KMnO4:acid (H2SO4:H3PO4) at the weight ratios of 1:80, 1:40, and 1:20 (hereon denoted as KS-80, KS-40,
and KS-20, respectively, where the weight ratio of H2SO4:H3PO4 was fixed at 9:1) and the weight ratio of graphite:sulfuric acid was maintained at
1:100 for all the samples. Oxidation time was fixed at 6 hours, 1 day, 2 days, and 3 days, respectively, to evaluate the oxidation level effects on
GO. The samples were gently sonicated for 5 minutes prior to FESEM viewing (Huang et al. 3445).
I keep the mixture close to room temperature by stirring and careful addition of the reagents, to keep away from that magic 55C temperature where
Mn2O7 supposedly detonates. I've never verified what it does above that temperature, nor do I intend to in the future. I
haven't had any problems personally. The general procedure above is the most commonly used synthetic method for partial graphite oxidation. It is a
modified Hummers Method.
Huang, NM et al. “Simple Room-Temperature Preparation of High-Yield Large-Area Graphene Oxide.” International Journal of Nanomedicine 6 (2011):
3443–3448. PMC. Web. 4 Dec. 2014.
Attachment: Simple room-temperature preparation of high-yield large-area graphene oxide.pdf (1.5MB)
This file has been downloaded 474 times
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woelen
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As far as I know, HMnO4 does not exist in the pure free form. If you were to get pure HMnO4, then it would fall apart to water and Mn2O7 and H3O(+)
and MnO4(-) ions and there will be some equilibrium. Such a mix would hardly contain any HMnO4. Most of it will be Mn2O7 and H3O(+) with MnO4(-). If
more water were added, you would get more HMnO4, which falls apart to H3O(+) and MnO4(-) with water.
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chornedsnorkack
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| Quote: Originally posted by woelen | | As far as I know, HMnO4 does not exist in the pure free form. If you were to get pure HMnO4, then it would fall apart to water and Mn2O7 and H3O(+)
and MnO4(-) ions and there will be some equilibrium. |
Outside water stability field?
I have read references where pure solid HMnO4 WAS produced. (And so was some crystal hydrate, I think HMnO4.2H2O).
So... what do you normally get when dilute permanganic acid aqueous solutions are frozen? Ice I is well known for being poor solid solvent. And what
happens when permanganic acid solutions are vacuum evaporated, below freezing point?
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woelen
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Maybe at some low temperature it is possible to obtain HMnO4, but in practice, at room temperature, you either get Mn2O7, or H3O(+) and MnO4(-).
I myself have done quite a few experiments with Mn2O7 and acidified permanganate at all kinds of concentration of H2SO4 and I never obtained HMnO4.
From literature (diverse textbooks) I also concluded that HMnO4 is not stable (either forming Mn2O7, split into ions, or loss of oxygen, forming MnO2,
water and oxygen).
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