quantumcorespacealchemyst
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theoretical equation, 8KMnO4 + 32H2SO4 + 16Ag + 64H2O
8KMnO4 + 32H2SO4 + 16Ag + 64H2O ---> 4K2SO4 + 8MnSO4 +16AgSO4 +4H2SO4 + 36H2O + 4O2
I am guessing that AgSO4 will be able to be separated first as a precipitate. K2SO4 seems to be the next least soluble, however, I don't know if it
precipitates well. I remember (possibly incorrectly) reading that it gets syrupy  .
Even though it seems there can be cancelling of water and sulfuric acid from both sides, I left it for reference if starting to rework/examine the
mechanism from the beginning.
I didn't get to do this reaction yet.
please let me know various techniques for separation/purification (for analysis and also for product purification).
[Edited on 9-2-2015 by quantumcorespacealchemyst]
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woelen
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This is a very theoretical equation. First, you could simplify it (at the left and the right you have water and sulphuric acid). Second, I do not
think that this reaction will occur. I expect you to get no AgSO4 at all and no oxygen. AgSO4 contains silver in oxidation state +2 and that is very
rare.
I would expect the reaction as follows (unbalanced, I leave that as exercise for you):
KMnO4 + H2SO4 + Ag --> Ag2SO4 + K2SO4 + MnSO4 + H2O
Use excess H2SO4 to assure that the liquid remains acidic and the reaction rate remains acceptably high.
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quantumcorespacealchemyst
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Thank you!
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Nicodem
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I seriously doubt you can get AgSO4 from aqueous chemistry. If it was so easy, this compound would have been prepared earlier and not only
5 years ago (DOI: 10.1002/anie.200906863).
Also, where did you get the information that AgSO4 is insoluble in water? I could not find any solubility data. It might just as well
decompose in the presence of water. Besides, if it would be stable and insoluble, the procedure for its preparation would probably avoid calling for
washing the crystals with anhydrous HF.
See also other articles about AgSO4: DOI: 10.1039/C2CE26282G, DOI:10.1016/j.apsusc.2012.12.097, DOI:10.1016/B978-0-08-097774-4.00222-9, and
DOI: 10.1002/anie.201000448.
You can easily create transient Ag(II) species in aqueous media by oxidation of Ag(I) salts with persulfates. These hydrated Ag(II) species are
unstable and utmost useful for SET oxidation catalysis, as in the Minisci reaction, for example, but I don't think you can prepare Ag(II) compounds in
such a manner (though this may depend on the ligands used). Perhaps it might be possible by an oxidation of Ag2SO4 in anhydrous
sulfuric acid or oleum with some persulfate (Ag(II) compounds are supposedly stable in concentrated sulfuric acid: DOI: 10.1039/C3CC43072C).
…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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woelen
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Silver(II) can indeed be prepared from persulfate and silver(I), but the compounds produced most likely are mixed silver(I)/silver(III) compounds. I
did experiments myself and made some Ag2O2, which is stable and can be kept around. I still have two grams or so of Ag2O2. This is a very energetic
compound (i.e. it makes a kind of flash powder with red P). It dissolves in moderately concentrated HNO3 giving intense brown solutions.
Making AgSO4 almost certainly is beyond what you can do as home chemist. Permanganate certainly is not a sufficiently strong oxidizer. Actually, it
goes the other way around. If you add some of the above-mentioned brown solution to a solution of MnSO4, then the liquid turns purple violet,
indicating formation of permanganate ion and the silver goes to oxidation state +1. If you add some of the Ag2O2 to concentrated HClO4, then the solid
dissolves while fizzling. It decomposes, giving oxygen, and the liquid becomes colorless, due to formation of silver(I). Apparently, nitric aacid can
form some complex which is stable, while perchloric acid can't.
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Sulaiman
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I've only been doing chemistry for less than a year so please forgive my ignorance,
could this reaction be happening first
6 KMnO4 + 9 H2SO4 → 6 MnSO4 + 3 K2SO4 + 9 H2O + 5 O3
and the ozone help change Ag(I) to Ag(II) ?
[Edited on 9-2-2015 by Sulaiman]
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quantumcorespacealchemyst
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thank you, this information is great. i goofed here referencing Ag2O solubility as if it was AgO. I overlooked Ag making +1 salts and wrote this whole
thing thinking that the oxide was a +2 (I even overlooked that the silver oxide page on wikipedia said it was +1/Ag2O) . now i am interested in the Ag(II) ion.
Sulaiman, from what i gathered from wikipedia, the first reaction of Potassium permanganate with Sulfuric acid produces Manganese heptoxide, Mn2O7,
which in concentrated Sulfuric acid, Manganese heptoxide, turns to the Permanganyl ion MnO3+. The equation they had looked incorrect, as it was
unbalanced in charge and/or atoms. using MnO4- and Hydronium seems to balance it. i am unsure of it and worked out an equation based on it (leaving
the waters in the beginning, which can be cancelled)
while the reaction of Mn2O7 to get permanganyl (MnO3+) seemsto occur. it seems to need very concentrated sulfuric acid, oleum, to occur. [woelen and
others on here https://www.sciencemadness.org/whisper/viewthread.php?tid=61173#pid388933]i have ~98%H2SO4, so i have to see. nonetheless, it seems
that permangate (+7) likes to turn +4 and +2, needing 2 permanganates to do so. 2Mn+7 +8e- ---> Mn+4 + Mn+2. i may be wrong about this. it seems
correct from what i have gleaned.
heres the corrected, hopefully, equation
2KMnO4 + 2H2SO4 + 2H2O + 8Ag ---> K2SO4 + MnO + MnSO4 + 4Ag2O + 1/2O2 +4H2O
using the mechanism (if correct)
2KMnO4 + H2SO4 + H2O ---> 2K+ + Mn2O7 + HSO4- + O-2 + H3O+
and
Mn2O7 + H2SO4 + H2O ---> MnO4- + MnO3+ + HSO4- + H3O+
then using this bit, i found was crucial
MnO2 + H2SO4 ---> MnSO4 + H2O + 1/2O2
[Mn+4 + 2O-2 + SO4-2 +2H3O+ --->Mn+4 + SO4-2 + 3H2O + 1/2O2]
(H2SO4 + 2H2O <---> 2H3O+ + SO4-2)
i figured here that Mn+2 is more basic than Ag+ and grabs the SO4-2 ion after the 2K+ grabs an SO4-2, being more basic than the Mn+2. the Ag+ i
figured is more acidic and grabs the O-2, which is more basic than the SO4-2. the last Mn+2 grabs an O-2.
i don't know if this logic is correct as the Mn+2 ion has a charge of 2, and would possibly be more acidic, looking at it the other way, with K+ being
next least and Ag+ the least acidic. SO4-2 being more basic than O-2 in this logic, being SO4-2 will displace an Oxide. in this case possibly, 2MnSO4
+ 4Ag2O + K2O + 4H2O + 1/2O2 which turns to 2MnSO4 + 4Ag2O + 2KOH + 2H2O + 1/2O2. i don't know though as i think the potassium ion making a strong
base and the sulfate ion making a strong acid neutralize better than the manganese salt.
nonetheless, there seems to be a manganese sulfate mol at least in there and this makes me wonder about complex salts. i read MnSO4 makes a metal aquo
complex, [Mn(H2O)6]2+, i guess with SO2-2 as the anion. i don't know if Silver does this and if potassium doesn't being an alkali metal.
[Edited on 10-2-2015 by quantumcorespacealchemyst]
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Molecular Manipulations
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Quote: Originally posted by Sulaiman  | I've only been doing chemistry for less than a year so please forgive my ignorance,
could this reaction be happening first
6 KMnO4 + 9 H2SO4 → 6 MnSO4 + 3 K2SO4 + 9 H2O + 5 O3
and the ozone help change Ag(I) to Ag(II) ?
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Could it happen? Yeah, it's possible, every reaction is an equilibrium reaction.
That reaction won't produce much ozone though, yields are very bad.
Is that the likely mechanism?
Do you have a reference for the oxidation of silver (I) to silver (II) via ozone?
Think about the kinetics of the reaction you propose (regarding silver (I) oxidation).
An ozone molecule decomposes into oxygen and a free radical (endothermic), this radical must then collide with two Ag+ ions at
the exact same time with the proper orientation without oxidizing MnSO4 first. Sounds unlikely to me.
[Edited on 10-2-2015 by Molecular Manipulations]
-The manipulator
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blogfast25
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QCSA:
Again, so few kernels of truth mixed into word salad of undigested half truths, irrelevancies and plain nonsense and hard to disentangle.
This experimenter really needs to review his learning strategy and start working on simpler 'beginner' problems and preparations.
Re. Ag(II) compounds, if reactions can be found where the second ionisation energy can be offset for instance by extra lattice energy then the overall
ΔG of reaction can be made negative.
Example: Cu(III) compounds;
3 KCl(s) + CuCl(s) + 3 F2(g) ===> K3CuF6(s) + 2 Cl2(g)
K3CuF6 oxidises water to O2 and gets reduced to Cu(II).
Ag(II) compounds are likely to suffer the same fate, making aqueous preparation impossible.
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blogfast25
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And this means what exactly? That every pretty and balanced reaction equation someone can come up with represents a possible reaction? Possible in
what way and why?
Why for instance write the oxygen as 5 O<sub>3</sub> and not 7.5 O<sub>2</sub>? Ozone is an endothermic substance: Standard
Enthalpy of Formation = + 142.7 kJ/mol. In and of itself that would mean (any entropic effects neglected) the formation of O2 would be favoured over
the formation of ozone.
Trust me, no ozone can be prepared that way.
Incidentally, reacting conc. H2SO4 with KMnO4 causes the potentially explosive Mn2O7 to be formed. With dilute H2SO4 you end up with
a simple acidic solution of KMnO4 in dilute H2SO4.
[Edited on 10-2-2015 by blogfast25]
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Sulaiman
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I'm not qualified to debate this
but when I did the 'storm in a bottle' experiment I used 96% H2SO4
https://www.youtube.com/watch?v=Rm42H6fRmfQ
and with potassium permanganate something was produced
that violently reacted with the alcohol above the H2SO4.
judging from the number of flashes I guess the yield was quite low.
[Edited on 10-2-2015 by Sulaiman]
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Molecular Manipulations
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| Quote: Originally posted by blogfast25 |
And this means what exactly? That every pretty and balanced reaction equation someone can come up with represents a possible reaction?
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No, that's not my point, perhaps I should have said: Every reaction that happens is an equilibrium
reaction.
| Quote: Originally posted by blogfast25 |
Trust me, no ozone can be prepared that way.
Incidentally, reacting conc. H2SO4 with KMnO4 causes the potentially explosive Mn2O7 to be formed. |
Yes we all know that manganese heptoxide is formed upon addition of conc. sulfuric acid. But a little known fact is that it may decompose into ozone
and oxygen above 0°C. Vulture posted that 11 years ago: | Quote: Originally posted by vulture | | You'd need to store it below -10C. Otherwise it will slowly decompose yielding ozone and oxygen. This can be troublesome in closed vessels as high
ozone concentrations are prone to detonation. |
It's right there on Wikipedia's KMnO4 page.
Obviously formation of oxygen is thermodynamically favored, but as I'm sure you're well aware, thermodynamics and equilibrium tied together. Gibbs
Free Energy can be used to predict the Keq of a reaction.
And if your "the formation of O2 would be favoured over the formation of ozone." argument was valid, then ozone couldn't be prepared using any method,
because the formation of O2 would be favoured over the formation of ozone. But of course this isn't the case. The main reason oxygen is produced
nearly pure in most reactions has to due with kinetics mostly, which happen because of thermodynamics, thus oxygen may be prepared containing ozone at
equilibrium, ozone in quantities much less than or much greater at than equilibrium, it depends on the reaction's kinetics.
Now when ozone is made, the energy must come from somewhere, but the total reaction is exothermic, so perhaps it comes from that.
[Edited on 10-2-2015 by Molecular Manipulations]
-The manipulator
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blogfast25
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| Quote: Originally posted by Molecular Manipulations |
It's right there on Wikipedia's KMnO4 page.
Obviously formation of oxygen is thermodynamically favored, but as I'm sure you're well aware, thermodynamics and equilibrium tied together. Gibbs
Free Energy can be used to predict the Keq of a reaction. |
Yes, it is. Live and learn.
With the ΔG for:
3 O<sub>2</sub> < === > 2 O<sub>3</sub>
... being positive that clearly shows (Nernst) that that equilibrium lies firmly to the left (K = a<sub>O3</sub><sup>2</sup> /
a<sub>O2</sub><sup>3</sup> << 1). This explains why mainly Mn2O7 is the reaction product, with only small amounts of
ozone.
| Quote: Originally posted by Molecular Manipulations | | And if your "the formation of O2 would be favoured over the formation of ozone." argument was valid, then ozone couldn't be prepared using any method,
because the formation of O2 would be favoured over the formation of ozone. |
Here you are confused: the equilibrium constant is small, so small concentrations of O<sub>3</sub>, prepared in whatever way, don't
violate that.
| Quote: Originally posted by Molecular Manipulations | | But of course this isn't the case. The main reason oxygen is produced nearly pure in most reactions has to due with kinetics mostly, which happen
because of thermodynamics, thus oxygen may be prepared containing ozone at equilibrium, ozone in quantities much less than or much more than
equilibrium, it depends on the reaction. |
Careful with mixing up thermodynamics and kinetics. Thermodynamics makes NO pronouncements about kinetics whatsoever.
Interesting nonetheless. Thanks.
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Sulaiman
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I just tried potassium permanganate in 96% sulfuric acid
and held moist filter paper above the mixture ....
no ignition and no ozone smell
(I didn't try to get much of a smell due to the H2SO4)
I did see a distinct green colouration of the fluid.
My shed/lab and contents are at c5C at the moment so I may try with warmer acid
BUT
I'm now concerned about this Mn2O7
1) how dangerous is it?
2) how can I proceed at minimal risk from Mn2O7?
The thought of an explosion in 96% H2SO4 isn't encouraging.
(I did a quick forum search and found nothing useful)
any advice?
[Edited on 10-2-2015 by Sulaiman]
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Molecular Manipulations
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| Quote: |
Thermodynamics makes NO pronouncements about kinetics whatsoever.
|
Yes. Thermodynamics predicts which direction a reaction will go, kinetics attempts to explain how it gets there, what steps are involved.
What I'm saying is, the ozone doesn't come from oxygen after reaching equilibrium, rather it's produced as a byproduct along with oxygen, not the
result of an equilibrium. How exactly this happens I've no idea. Just like electrolysis of sulfuric acid with Pt anodes, ozone is produced with
oxygen, but in concentrations much, much higher than the normal equilibrium would bring it. The energy in this case comes from the electricity, and
may be catalyzed by the Pt.
Sulaiman
1) Manganese heptoxide is quite nasty. I'd avoid it altogether, but I know that's not the answer you want.
2) Use very small amounts is the best advice, dilute and dispose of as soon as you're done.
I wouldn't expect ozone to catch moist filter paper on fire at all, so that's no surprise to me.
-The manipulator
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blogfast25
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Yes, be careful. AVOID any organics because Mn heptoxide is one of the most powerful oxidisers. A grain of organic dust could set it off.
http://en.wikipedia.org/wiki/Manganese_heptoxide
[Edited on 10-2-2015 by blogfast25]
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Sulaiman
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o.k. thanks.
I only used c5ml acid and a small spatula of permanganate.
but the thought of even 5ml (and glass shards) exploding at me did make me put a shield between it and me.
(too many almost nasty accidents doing experiments as a youth)
I have already diluted and disposed of the mixture.
(I added this part after the post below, not realising it had been posted)
I just read the wikipedia article;
"Mn2O7 decomposes near room temperature, explosively so at > 55 °C."
so
I don't think I'll bother with warm acid !
In fact I will avoid H2SO4 + KMnO4 for as long as my memory works.
Thanks.
[Edited on 10-2-2015 by Sulaiman]
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blogfast25
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| Quote: Originally posted by Sulaiman | o.k. thanks.
I only used c5ml acid and a small spatula of permanganate.
but the thought of even 5ml (and glass shards) exploding at me did make me put a shield between it and me.
(too many almost nasty accidents doing experiments as a youth)
I have already diluted and disposed of the mixture. |
Try 1.0 ml acid and a small pinch of KMnO4. Please? Scale up slightly later if you want.
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quantumcorespacealchemyst
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I was mistaken about the Ag(ii) and reformulated above the possible reaction. Any thoughts?
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blogfast25
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Where did you reformulate? Quote, please...
[Edited on 15-2-2015 by blogfast25]
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quantumcorespacealchemyst
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heres the corrected, hopefully, equation
2KMnO4 + 2H2SO4 + 2H2O + 8Ag ---> K2SO4 + MnO + MnSO4 + 4Ag2O + 1/2O2 +4H2O
using the mechanism (if correct)
2KMnO4 + H2SO4 + H2O ---> 2K+ + Mn2O7 + HSO4- + O-2 + H3O+
and
Mn2O7 + H2SO4 + H2O ---> MnO4- + MnO3+ + HSO4- + H3O+
then using this bit, i found was crucial
MnO2 + H2SO4 ---> MnSO4 + H2O + 1/2O2
[Mn+4 + 2O-2 + SO4-2 +2H3O+ --->Mn+4 + SO4-2 + 3H2O + 1/2O2]
(H2SO4 + 2H2O <---> 2H3O+ + SO4-2)
i figured here that Mn+2 is more basic than Ag+ and grabs the SO4-2 ion after the 2K+ grabs an SO4-2, being more basic than the Mn+2. the Ag+ i
figured is more acidic and grabs the O-2, which is more basic than the SO4-2. the last Mn+2 grabs an O-2.
i don't know if this logic is correct as the Mn+2 ion has a charge of 2, and would possibly be more acidic, looking at it the other way, with K+
being next least and Ag+ the least acidic. SO4-2 being more basic than O-2 in this logic, being SO4-2 will displace an Oxide. in this case possibly,
2MnSO4 + 4Ag2O + K2O + 4H2O + 1/2O2 which turns to 2MnSO4 + 4Ag2O + 2KOH + 2H2O + 1/2O2. i don't know though as i think the potassium ion making a
strong base and the sulfate ion making a strong acid neutralize better than the manganese salt.
nonetheless, there seems to be a manganese sulfate mol at least in there and this makes me wonder about complex salts. i read MnSO4 makes a metal
aquo complex, [Mn(H2O)6]2+, i guess with SO2-2 as the anion. i don't know if Silver does this and if potassium doesn't being an alkali metal.
[Edited on 17-2-2015 by quantumcorespacealchemyst]
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DraconicAcid
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If anything, the most likely reaction is:
MnO4- + 8 H+ + 5 Ag -> Mn2+ + 5 Ag+ + 4 H2O.
| Quote: | i figured here that Mn+2 is more basic than Ag+ and grabs the SO4-2 ion after the 2K+ grabs an SO4-2, being more basic than the Mn+2. the Ag+ i
figured is more acidic and grabs the O-2, which is more basic than the SO4-2. the last Mn+2 grabs an O-2.
i don't know if this logic is correct as the Mn+2 ion has a charge of 2, and would possibly be more acidic, looking at it the other way, with K+
being next least and Ag+ the least acidic. SO4-2 being more basic than O-2 in this logic, being SO4-2 will displace an Oxide. |
This is gibberish. The potassium ions are not going to "grab" anything- they are just going to hang out in aqueous solution. If you are going to see
any oxides precipitate, it's simply due to the insolubility of transition metal oxides, not some fanciful acidity or basicity of the ions.
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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quantumcorespacealchemyst
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so then the relative solubilities of different salts do not affect crystallization? if the formation of an insoluble precipitates, then with less
water, would this be the same(K2SO4), or does the possibility not occur because other ions keep it from forming (doesn't happen with oxides though,
which just form).
I found out from a chemist recently that Activity Coefficients govern the salt formation of soluble ions. The method to figure this out seems to be
Specific Ion Interaction Theory.
the thing is that is for soluble salts, with a salt that has a solubilty difference as much as K2SO4 to others, one might suppose it can be isolated
by reducing the water (evaporation and cooling enough to break supersaturation). i say that because it seems that way, but the reality is what i am
inquiring about.
[Edited on 22-2-2015 by quantumcorespacealchemyst]
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