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ethan_c
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| Quote: | Originally posted by agent_entropy
Awesome, rogue chemist, I was wondering how critical the chromosulfuric acid cleaning was.
Did the procedure have any negative effect on the mortar and pestle? eg. stains, erosion, etc...? |
The reason they specifically recommend chromosulfuric (chromerge) cleaning is because that particular method is very good at dissolving and destroying
any traces of organic compounds. In a laboratory, there's a good chance that the mortar & pestle would be used for crushing solid organics as
well, or that one might decide to clean it with an alcohol or acetone or some other solvent, which of course would result in a rather rapid exothermic
reaction upon contact with the manganese heptoxide.
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The_Davster
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I figured that adding the KMnO4 slowly, as the procedure calls for, would slowly increase the Mn2O7 concentration, oxidizing any organic stuff in a
controlled fashion. Of course, this only works for the parts of the mortar and pestle in contact with the solution, but I figure the Mn2O7 fumes(as
seen in a pic I uploaded) would destroy anything organic above the the acids meniscus.
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agent_entropy
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So, hydrolyzing Mn2O7 in water would produce permanganic acid (right?). My searches for permanganic acid say that it only exists in dilute solution.
How dilute does it have to be?/How concentrated can it be? Any interesting uses for permanganic acid?
Also, as I understand it the plain manganate ion is MnO4 ^ -2, eg. permanganate but reduced by one electron. How come I never see any ions of the
type MnO3? (As I understand the nomenclature of polyatomic ions this should be the one called manganate, what's the deal with that?)
[Edited on 31-7-2006 by agent_entropy]
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Pyrovus
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Well, according to wikipedia, the H2SO4 solution of Mn2O7 gives the cation MnO3+.
I wonder if it would be possible to get this cation to participate in nucleophilic aromatic substitution reactions, in a manner analogous to nitration
(i.e. to compounds of the form R-MnO3 with R being an aromatic compound). Given how violently Mn2O7 likes to react with organic materials, it would
probably be difficult to avoid oxidising the starting material all the way to CO2 and H2O. However, if it is possible to do this, this could provide a
very interesting application of Mn2O7.
Never accept that which can be changed.
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The_Davster
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That would be kind of analogous to the reaction of iodine pentoxide and conc. H2SO4 to make (IO)2SO4, and pure iodosyl sulfate, without acid will do
SEAr on aryl comounds to produce the corresponding iodosoaryl compound. With acid, more complicated polyvalent iodine compounds are formed such as
diphenyl iodosobenzene +. An attempt at such a reaction on resorcinol lead to a spectacular runaway, big cloud of iodine vapour.
But back to the MnO3+, I imagine it could be used in nucleophilic aromatic substitution, but the direct metal-carbon may make things wonky somehow, I
don't know enough about that sort of chemistry. As the Mn2O7 is really MnO3+(manganosyl? what is the name for this ion?) permanganate technically,
perhaps like I2O5, it could hydrolyse in other acids making stuff like manganosyl sulfate(likely a better, less violent 'manganosylating' agent,
analogous to (IO)2SO4), manganosyl perchlorate, and all sorts of fun stuff like that.
OK, I made up too many words in that post.
EDIT: Next time I make Mn2O7 I am definatly going to attempt 'manganosyl perchlorate' by dissolving Mn2O7 in perchloric acid. At VERY low temp of
course
[Edited on 1-8-2006 by rogue chemist]
EDIT: Thanks Axt for the name of the MnO3+ ion.
I like the sound of permanganyl perchlorate.
[Edited on 1-8-2006 by rogue chemist]
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Axt
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CH3MnO3 rated a mention in "Overviews of Recent Research on Energetic Materials" by Sham et.al. but only a mention. Something about its
incompatibility with other energetic materials. Sorry for being vague but I dont have the book with me here. They refered to it as methylmangenese
trioxide.
Oh, MnO3 would be "permanganyl" rather then manganosyl. So I guess permanganylmethane would be another suitable name. though google knows nothing of
"permanganylbenzene" and the like.
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The_Davster
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After a bit of googling around(I'll do some scifinder searches tomorrow) it seems the permanganyl ion is a bit of a hot topic in reactive intermediate
studies. We have some of that research going on at my university, I may, time permitting, look a couple proffs up and have a 'hypothetical'
discussion about this sort of thing.
I did find some interesting information about permanganyl chloride on an incompatibility sheet. Apparently an explosion occured while attempting to
prepare it by adding conc. H2SO4 to a mix of KCl and KMnO4.
http://www.chemone.com/default/other/chemical%20incompatibil...
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JohnWW
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Permanganyl fluoride, MnO3F, has been made and is stable; like Mn2O7, but more volatile.
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franklyn
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Food for thought
| Quote: | Originally posted by kABOOM!
I'd say Peroxyacetone is far safer than this stuff. |
That's not very reassuring
Wiki states CCl4 solvates Mn2O7, so it's reasonable that CF4 will also.
I'm thinking a CF4 solution of Mn2O7 might be a good reagent to try with
Xenon Fluoride, to try to obtain higher oxides of Xenon. The two most notable
are XeF4 and XeF6 which yields XeO3 and HF in water. Mn2O7 has three double
bonded oxygens to each Manganese both sharing the seventh oxygen with a
single bond each. There are three Fluorides of Manganese, MnF2, MnF3, MnF4,
the question is what would be the reaction path and resulting compounding ?
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not_important
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Given that Freon-14 is a gas at STP, 'solution' might not be the proper term. I think it boils around -130 C, that cold usually doesn't lead to good
solutions.
Another thing is that highly fluoronated materials tend to for a 'third phase', not aqueous and not non-aqueous, the flourocarbons tend to stick to
their own company.
You're thinking of reacting a xenon fluoride with manganese heptoxide to get a higher manganese flouride and xenon trioxide?
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franklyn
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| Quote: | Originally posted by not_important
Given that Freon-14 is a gas at STP, 'solution' might not be the proper term. I think it boils around -130 C
You're thinking of reacting a xenon fluoride with manganese heptoxide |
I should have mentioned also that this would have to be done
at elevated pressure being that CF4 is a volatile refrigerant.
The objective would be to see if a super oxygenated Xenon
can be made. One or more of the Manganese Fluorides must
result regardless.
Nothing ventured nothing gained, this is experimental after all.
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woelen
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| Quote: | Originally posted by JohnWW
Permanganyl fluoride, MnO3F, has been made and is stable; like Mn2O7, but more volatile. |
Probably I have made this compound, but I have not isolated it.... too explosive and too toxic!
http://woelen.homescience.net/science/chem/exps/KMnO4+NaF+H2...
The experiment, however, is quite funny and impressive. Be careful though, if you want to repeat it. It is extremely corrosive, explosive and
extremely toxic, all at the same time!
EDIT(woelen): Changed link so that it works again.
[Edited on 8-3-13 by woelen]
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The_Davster
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I am not sure how you came to the conclusion that permananyl fluoride was present, it looks like frothy Mn2O7 too me. It definatly could be in there,
but I am not sure the visual evidence suggensts this.
I unfortunatly will not try this, HF is something I will not work with, I would sooner drip pure Mn2O7 into anhydrous perchloric acid.
I still do have designs to attempt a perchlorate.
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woelen
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The color of the compound is quite distinct from Mn2O7 in sulphuric acid. Mn2O7 in H2SO4 is dark olive green, this compound at a similar concentration
is much lighter and brighter green. I also read in literature that MnO3F is green, so that supports my hypothesis of this stuff being MnO3F, but as
you could have read on my site, I am not 100% sure about that.
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Axt
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| Quote: | Originally posted by Axt
CH3MnO3 rated a mention in "Overviews of Recent Research on Energetic Materials" by Sham et.al. but only a mention. Something about its
incompatibility with other energetic materials. Sorry for being vague but I dont have the book with me here. They refered to it as methylmangenese
trioxide. |
This is what was said:
<i>"<b>Methyl manganese trioxide</b> (CH<sub>3</sub>MnO<sub>3</sub> : This Useful oxidant and catalyst is highly exothermic whereby the fuel-laden methyl moiety is oxidized by the
unusually labile oxygen available at the high-valent metal centre. Unfortunately, this material is too reactive to be compatible with a variety of
sate-of-the-art formulations."</i>
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tom haggen
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I've always enjoyed this topic. However, the few times I tried experimenting with my own Mn2O7 it didn't really work out to well. I added very fine
potassium permaganate to sulfuric acid and I definitely saw the formation of the dark green Mn2O7, it was unmistakable. However, A large solid mass
was also formed which made it very hard to separate the Mn2O7. I assume that the solid mass was K2SO4 which is one of my most least favorite chemical
by products. Any way you guys are fucking nuts playing around with this stuff. Roguechemist, I am amazed that you still have 10 fingers because that
picture you posted almost made my computer screen explode it looked so volatile. Please excuse my shitty spelling I dont have microsoft word installed
at the moment so no spell checker.
[Edited on 16-9-2006 by tom haggen]
N/A
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ShadowWarrior4444
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Just to clarify, when KMnO4 is added to conc sulfuric acid, what is the resulting green solution? Is this a Mn2O7 disolved in the acid that slowly
seperates out, or is it a mix of permanganic acid with sulfuric?
(Edit: Nevermind, it is apparently permanganyl sulfate, though the uses of this compound are not entirely clear yet. Perhaps a few more seconds of
research will do it, heh.) It does raise the question of perhaps forming a permanganyl nitrate or persulfate, although I am aware that adding H2O2
will result in decomp to Ozone, so forming Permanganyl Monopersulfate would likely not work using that method. It seems that permanganyl fluoride is
stable, so I have my doubts about any other anions being able to stabilize it. It may be interesting to attempt a reaction between it and a
fluoropolymer.
On a related note, I’ve been testing the oxidizing power of this in solution on Sulfur. (As inspired by woelen's thread on higher oxidation states
of S, Te, and Se.)
The procedure was as such: KMnO7 added to H2SO4 in a small borosilicate culture tube. Elemental sulfur was then immediately added after. Results so
far: No explosive reaction, the granules of sulfur bubbled slightly which ceased quickly (chalked up to impurities,) then the green around each sulfur
granule gradually changed to purple, and the sulfur granules themselves were observed to turn white. Only mild heating of the tube was noted. Nothing
nearly as violent as I would expect from the Mn2O7 oil.
Something else I just remembered from an earlier experiment:
Adding 30% H2O2 to Mn2O7 produced a thick brown smoke that was quite obviously more dense than air. I seem to recall Mn vapor appearing a wee bit
purple, and this looked distinctly like NO2, I didn't detect any scent, but then I didn’t any long whiffs either. Is it possible that rapid
formation of hot O3 reacted with the N2 in the atmosphere? (This would certainly be an incredibly circuitous method for obtaining HNO3.)
P.S. I am not insane, I'm just very good friends with strong oxidizers. *smirk*
[Edited on 5-1-2008 by ShadowWarrior4444]
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ScienceGeek
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Sounds very interesting what you've done here!
I don't know if Ozone is capable of oxidising di- Nitrogen molecules, but it is able to oxidise Nitrogen Oxides.
You can test this, though! (This is one way I can think of, there probably is a better one):
Try to isolate the gas, and freeze it. If it goes colorless, it probably is Nitrogen Dioxide! Don't know if this test can act ambiguously...
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-jeffB
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| Quote: | Originally posted by ShadowWarrior4444
Adding 30% H2O2 to Mn2O7 produced a thick brown smoke that was quite obviously more dense than air. |
I don't believe I tried that, but when I put Mn2O7 on a paper towel, it eventually "poofed" into a mixture of purple smoke-rings and wisps of
tissue-like brown MnO2 (I assume). The latter floated for quite a long time before settling out all over everything. The purple vapor was incredibly
acrid, smelling like a cross between chlorine and hydrogen chloride, but with a strange fruity note. I figured it wasn't a good idea to take too many
sniffs.
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The_Davster
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| Quote: | Originally posted by ShadowWarrior4444
Just to clarify, when KMnO4 is added to conc sulfuric acid, what is the resulting green solution? Is this a Mn2O7 disolved in the acid that slowly
seperates out, or is it a mix of permanganic acid with sulfuric?
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The green colour is the result of the permanganyl cation. However if you observed the formation of shiny green drops on the surface of the liquid,
this is the pure liquid Mn2O7, and not just a solution. There is another thread on Mn2O7 in which I detail the preparation of this pure.
| Quote: | Originally posted by ShadowWarrior4444
Adding 30% H2O2 to Mn2O7 produced a thick brown smoke that was quite obviously more dense than air. I seem to recall Mn vapor appearing a wee bit
purple, and this looked distinctly like NO2, I didn't detect any scent, but then I didn’t any long whiffs either. Is it possible that rapid
formation of hot O3 reacted with the N2 in the atmosphere? (This would certainly be an incredibly circuitous method for obtaining HNO3.)
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Catch it in a large overturned jar, it is just a very light fluffy MnO2 that can stay suspended in the air for a while. If you add a drop of either
Mn2O7 solution or the pure Mn2O7 to a piece of paper, when it catches fire this same MnO2 fume is given off, and it is a lot less dispersed and is
easier to identify.
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ShadowWarrior4444
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| Quote: | | Quote: | Originally posted by ShadowWarrior4444
Adding 30% H2O2 to Mn2O7 produced a thick brown smoke that was quite obviously more dense than air. I seem to recall Mn vapor appearing a wee bit
purple, and this looked distinctly like NO2, I didn't detect any scent, but then I didn’t any long whiffs either. Is it possible that rapid
formation of hot O3 reacted with the N2 in the atmosphere? (This would certainly be an incredibly circuitous method for obtaining HNO3.)
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Catch it in a large overturned jar, it is just a very light fluffy MnO2 that can stay suspended in the air for a while. If you add a drop of either
Mn2O7 solution or the pure Mn2O7 to a piece of paper, when it catches fire this same MnO2 fume is given off, and it is a lot less dispersed and is
easier to identify. |
I have exposed a piece of paper to the Mn2O7 oil and observed the vapor--this vapor was decidedly different in color, relative density toward air, and
scent. The vapor product of Mn2O7 and H2O2 left a solid black material behind (likely Mn, or MnO2), giving off thick brown smoke. I should note that
this was pure Mn2O7 and pure 30% H2O2. *would certainly have recognized the products of Mn2O7/organic combustion.*
This was also formed not as an instantaneous plume, but as gradually rising brown smoke within the culture tube.
| Quote: | Originally posted by ScienceGeek
Sounds very interesting what you've done here!
I don't know if Ozone is capable of oxidising di- Nitrogen molecules, but it is able to oxidise Nitrogen Oxides.
You can test this, though! (This is one way I can think of, there probably is a better one):
Try to isolate the gas, and freeze it. If it goes colorless, it probably is Nitrogen Dioxide! Don't know if this test can act ambiguously...
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I am aware of industrial processes that utilize ozone generated via electrical arc to oxidize atmospheric N2 which is then collected as nitric
acid--this was one of the first industrial processes for producing HNO3. It can even be accomplished quite nicely on a lab scale using a neon sign
transformer, a few borosilicate tubes, and a wee bit of knowledge about setting up corona discharges.
As for testing the substance via freezing, I'll attempt to do that by sealing in the culture tube with PTFE. Though, any procedure for an NO2 test
that doesn’t involve glass, a freezer, and waiting, would be appreciated. *can tend to be overly paranoid though*
Now that I think about it, this may indeed give an ambiguous result--if it is a metal vapor it will likely condense yielding a colorless tube. Also, a
test responsive to any oxidizer may give a false positive, as well.
[Edited on 5-2-2008 by ShadowWarrior4444]
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Trotsky
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I didn't feel like creating a new thread for this, but since I had a clean batch of concentrated sulfuric acid, I thought I'd make some MnO7 and video
tape it. Well, that was done and I ran out, so I made some more and wasn't video taping it.
I should have been!
I took a chunk of wax a bit smaller than an ES Tylenol, so it probably weighted 800mg or so. I did not bother to weigh it- I just wanted to see if it
would start it on fire too.
I dropped it on my MnO7 mixture and waited. Nothing happened. So- with gloves and face protection, I grabbed a channel lock and with a foot or so of
thick steel wire, I pushed that wax a little bit.
The resulting boom was enough to crack the window I was near and generated a bit of ringing in one ear (almost normal now).
I've done the toluene and xylene thing with it- this FAR surpassed that.
I'll probably attempt to duplicate this under controlled settings and far more carefully, but wow, this was something special.
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AndersHoveland
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One idea would be to place a tube of Mn2O7 at the center of a liquid fuel. The detonation of the Mn2O7 would suddenly cause mixture of the resulting
oxygen by-product with the fuel. My guess is that there would be a slight delay before the oxygen-fuel mixture spontaneously ignited. The initial
force of the explosion would blow out any flame, quickly thereafter the remnants of the Mn2O7 would spontaneously ignite. Placing the whole thing a
thick walled container to help contain the initial small detonation and fuel-oxygen mixture would likely also help.
A similar concept has been demonstrated with pyrophoric rocket fuels. A combination a pyrophoric fuel and oxidizer are placed in numerous tubes. A
small explosive at the center suddenly mixes them, and the rocket fuel combustion becomes part of the explosion.
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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DraconicAcid
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Urg. I remember making Mn2O7 once in solution. I mixed the KMnO4 with conc. H2SO4, and saw the lovely green solution. Now, I could have sworn that
Cotton and Wilkinson had described the oxide as being soluble in carbon tetrachloride or chloroform, so I figured that dichloromethane (the most
common solvent in that lab) would work fine. Big mistake- huge cloud of brown MnO2-coloured smoke billowed out of the mixture, and a huge mess was
created. I later checked the book, and found that it was "carbon tetrachloride or chlorofluorocarbons". Not chloroform. And definitely not
dichloromethane.
I repeated the reaction once to show off to a date, and was rewarded with the sight of the dichloromethane actually catching fire, and a bigger mess
than I had remembered.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Trotsky
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"A similar concept has been demonstrated with pyrophoric rocket fuels. A combination a pyrophoric fuel and oxidizer are placed in numerous tubes. A
small explosive at the center suddenly mixes them, and the rocket fuel combustion becomes part of the explosion. "
That's a cool idea, actually. A 6x6x8 inch block of wax, 9 holes drilled into it (6.5 inches deep, so not all the way through). 9 6" test tubes
filled with Mn2O7, stuck into the wax slightly, with the open end covering the open hole. Keep it wax-side up, until ready to set off.
Flips over, the Mn2O7 will quickly (if it's not too cold out, it REALLY thickens up and slows it's reactivity if it's under -5C) fill the wax holes.
If the quantities are right, it should produce one hell of a detonation. Have your camera ready, I want to see it on youtube
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