Sciencemadness Discussion Board - the Permanganyl Ion MnO3+, salts and crystal systems

the Permanganyl Ion MnO3+, salts and crystal systems

quantumcorespacealchemyst - 14-1-2015 at 07:50

my prime concern is constructing novel/interesting Crystals out of/including this cation. i seek knowledge of how to work this ion into complex crystal systems such as those found in nature, the complex ones but in a pure form and lab controlled to make giant single crystals. ie: http://www.mindat.org/min-3182.html (H3O)3KCa(UO2)7(PO4)4O4 · 8H2O
i don't mean that MnO3+ is found readily in nature, on this planet, i mean that complex systems inherrant to nature may be possible with this cation if it were introduced into the formulation and so i seek how to engineer them.

please help me get information on salts/compounds that can be formed with the Permanganyl Ion MnO3+. preferably stable ones, and ones that can be isolated.

i am assuming, hopefully incorrectly that the preservation of the MnO3+ cation is dependant on a concentrated solution of sulphuric acid.

then it seems that an insoluble salt would have to be made to be extracted/filtered and even then, dissolving in water may degrade it back to permanganic acid or decompose to manganese oxide and oxygen or ozone

for reference on forum thread and backup data

"Mn2O7 can react further with sulfuric acid to give the remarkable cation MnO3+, which is isoelectronic with CrO3:[citation needed]

Mn2O7 + 2 H2SO4 → 2 [MnO3]+ + [HSO4]− + H2O

Mn2O7 decomposes near room temperature, explosively so at > 55 °C. The explosion can be initiated by striking the sample or by its exposure to oxidizable organic compounds."
http://en.wikipedia.org/wiki/Manganese_heptoxide

originally starting from KMnO4 there will be KHSO4 in solution
"Mn2O7 arises as a dark green oil by the addition of concentrated H2SO4 to KMnO4. The reaction initially produces permanganic acid, HMnO4 (structurally, HOMnO3), which is dehydrated by sulfuric acid to form its anhydride, Mn2O7.

2 KMnO4 +(cold) 2 H2SO4 → Mn2O7 + H2O + 2 KHSO4 "
and assume starting from permanganic acid, HMnO4, is ideal to leave out potassium ions for some reactions although i guess that isolating Mn2O7 can achieve this, but it will still probably have impurities, which cannot be washed out with water, as far as i know.

[Edited on 14-1-2015 by quantumcorespacealchemyst]

woelen - 15-1-2015 at 00:09

Interesting subject. I do not think that H2SO4 is sufficient to make permanganyl, but I can imagine that oleum can do this. I actually have some oleum (20% SO3 by weight, dissolved in H2SO4). I can take a little of this and add a tiny amount of finely powdered KmnO4 and see what happens. Maybe I get some interesting color change. I have no idea what color the MnO3(+) ion has (if it exists at all). I'll come back on this tomorrow or next weekend.

Isolating MnO3(+) in one of its compounds certainly is beyond my capabilities.

chornedsnorkack - 15-1-2015 at 00:37

Quote: Originally posted by woelen

Interesting subject. I do not think that H2SO4 is sufficient to make permanganyl, but I can imagine that oleum can do this. I actually have some oleum (20% SO3 by weight, dissolved in H2SO4). I can take a little of this and add a tiny amount of finely powdered KmnO4 and see what happens. Maybe I get some interesting color change. I have no idea what color the MnO3(+) ion has (if it exists at all). I'll come back on this tomorrow or next weekend.

Isolating MnO3(+) in one of its compounds certainly is beyond my capabilities.

Ask the question this way: what is the solubility of Mn(VII) at various concentrations of H2SO4/H2S2O7? What is the colour of the solution (Oleum is transparent and colourless)? And what is the condensed species in saturated solution?

At a certain concentration range, Mn2O7 precipitates as an immiscible liquid. But is there any concentration of oleum where the equilibrium solid would be a permanganyl salt?

turd - 15-1-2015 at 01:35

Quote: Originally posted by quantumcorespacealchemyst

i seek knowledge of how to work this ion into complex crystal systems

I don't think that word means what you think it means: https://en.wikipedia.org/wiki/Crystal_system.

Do you perhaps mean crystal structure (https://en.wikipedia.org/wiki/Crystal_structure)?

Even then your post doesn't make much sense. What does it mean to work an ion into a crystal structure? Do you want to crystallize this ion?

Good luck, as few well defined Mn(VII) compounds have been isolated
c.f.: http://onlinelibrary.wiley.com/doi/10.1002/zaac.200600186/pd...

Have you even done a literature search?

[Edited on 15-1-2015 by turd]

woelen - 15-1-2015 at 03:37

Quote: Originally posted by chornedsnorkack

Mn2O7 is soluble in H2SO4, its solutions are dark green. I can imagine that with SO3, this reacts with formation of SO4(2-) and MnO3(+) ions and maybe these MnO3(+) ions in solution have a different color than Mn2O7 in solution.

Pure oleum is colorless and clear, like water. My 20% oleum indeed is clear like water. I also have a small amount of 65% oleum. That oleum has some impurities, it is pale brown.
In contact with air, oleum gives of intensely dense white fumes, which are amazingly corrosive. These fumes are SO3, which reacts with water vapor in air to form droplets of conc. H2SO4. It is not the nicest of compounds to play with

chornedsnorkack - 15-1-2015 at 05:16

Quote: Originally posted by woelen

Mn2O7 is soluble in H2SO4, its solutions are dark green.

They can obviously be dark green only if they are dark and concentrated. Dilute solutions of Mn2O7 in concentrated H2SO4 cannot possibly be dark green - they have to be something else, like light green.
On the other hand, dilute solutions of KMnO4 or HMnO4 in water or in dilute H2SO4 are not green - they are pink-violet.
Around which concentration of H2SO4 in water does the pink-violet colour of HMnO4 turn into green of MnO3+?

woelen - 15-1-2015 at 06:31

I say that the solutions are dark green, because they are so. Even at low concentrations, the solutions are very intensely colored. Intense colors also are observed with KMnO4, a very small amount already gives a strong color in water.

You need highly concentrated H2SO4 to observe the green color. Only little water is needed to lose the green color and get the usual purple color again. I never measured the precise amount, but I have the impression that if you add only 10% or so of the volume of water to the solution in H2SO4, then it becomes purple again.

Btw., I do not think that the green color is due to the presence of MnO3(+) ions. I think that the green color simply is the color of dissolved Mn2O7.

blargish - 15-1-2015 at 10:17

I have come across some sources that claim that when KMnO₄ is added to excess concentrated H₂SO₄, the resultant solution is a solution of permanganyl hydrogensulfate. They also state that the MnO₃⁺ ion is green in solution. Attached is an excerpt from Inorganic Chemistry by Holleman and Wiberg (2001)

I wonder if a similar effect could be achieved with perchloric acid to form a solution of permanganyl perchlorate. Now that seems like an exciting compound

Screen Shot 2015-01-15 at 1.10.07 PM.png - 116kB

woelen - 15-1-2015 at 11:22

I tried the experiment of adding KMnO4 to oleum (20% SO3 by weight). In this experiment I added a tiny amount of KMnO4 (just a few mg) to appr. 1 ml of oleum. This solution is dark green, but its color is more bright than the color of KMnO4 in normal conc. H2SO4 (the latter is more olive green).
The color of the solution is not stable. It slowly changes color from green to grey and after 10 minutes or so, the liquid is dark grey and also is somewhat turbid.
When a drop of this dark grey liquid is added to water, then there is an impressive hissing noise and the resulting liquid is purple/red like a solution of KMnO4, which has partially decomposed to hydrous MnO2. So, the SO3 does make a difference.

I indeed am inclined to think that green MnO3(+) is formed, but that this easily decomposes, one of the products being MnO2 or some sulfate of that in the extreme liquid. With water, the remaining MnO3(+) gives permanganate ion again and the reddish color is caused by very finely dispersed hydrous MnO2 which is brown/yellow at low concentrations.

Unfortunately I cannot try the experiment with HClO4. I only have 60% and 70% HClO4 and with these, KMnO4 just gives a purple solution.

[Edited on 15-1-15 by woelen]

deltaH - 15-1-2015 at 11:39

Quote: Originally posted by woelen

... products being MnO2 or some sulfate of that in the extreme liquid.
[Edited on 15-1-15 by woelen]

or perhaps the MnO3+ is stripped of another oxygen to ???MnO2[3+]??? I've never heard of such a species though, so forgive my imagination...

woelen - 15-1-2015 at 11:45

I do not think that another ion is formed. I saw that the liquid becomes turbid and i have the impression this turbidity was due to formation of very small bubbles of oxygen, which only very slowly move to the surface, due to their small size and due to the high viscosity of the oleum.

It is a pity that I cannot do further research of this ion. It decomposes fairly quickly and the liquid in which it exists is so corrosive and heavily fuming that it is impossible for me to isolate anything from it.

(A nice and impressive demo is to drip oleum on paper tissue. The paper tissue chars at once and starts smoking while producing a loud hissing noise. Adding a drop results in complete charring in one second and the hissing and associated eating away of surrounding paper continues for a few more seconds.)

[Edited on 15-1-15 by woelen]

deltaH - 15-1-2015 at 11:53

Quote: Originally posted by woelen

I do not think that another ion is formed. I saw that the liquid becomes turbid and i have the impression this turbidity was due to formation of very small bubbles of oxygen, which only very slowly move to the surface, due to their small size and due to the high viscosity of the oleum.

[Edited on 15-1-15 by woelen]

Ah ok, now this is another matter

I was under the impression that it was a transparent grey and not turbid. I was actually also about to ask if you observed any bubbles, but you pre-empted my question!

Is it hypothetically possible that these bubbles contained ozone?

chornedsnorkack - 15-1-2015 at 13:20

Numbers for volatility I´ve encountered for oleum:
Boiling point:
Azeotropic concentration, 98,3 % acid - 330 degrees
100 % acid - 280 degrees
20 % SO3 - around 140 degrees
65 % SO3 - around 60 degrees
100 % SO3, liquid - 45 degrees, and also unstable to polymerization.

subsecret - 15-1-2015 at 15:07

Woelen, did you try adding the "fresh" permanganyl solution to water?

Molecular Manipulations - 15-1-2015 at 15:24

Quote: Originally posted by woelen

When a drop of this dark grey liquid is added to water, then there is an impressive hissing noise and the resulting liquid is purple/red like a solution of KMnO4, which has partially decomposed to hydrous MnO2. So, the SO3 does make a difference.

chornedsnorkack - 16-1-2015 at 00:17

How long does the olive-green solution of permanganate in 100 % H2SO4 last? Does it bubble oxygen to turn gray?

woelen - 16-1-2015 at 11:54

The olive-green solution in conc. H2SO4 (which is not 100% but somewhere between 96% and 98%, the rest being water) is stable. I kept it around for at least an hour and the color does not change.

chornedsnorkack - 16-1-2015 at 14:15

If the olive-green of concentrated H2SO4 permanganate solution turns into violet already by 10 % dilution with water, it is to be expected that 70 % perchloric acid also shows pink. 70 % perchloric acid is described as having Hammett H0 values around -7 - comparable to the acidity of 80 % H2SO4.
Found references at
https://books.google.ee/books?id=afkCAv3xwqQC&pg=PA43&am...
Is there any more reliable source than Google Books?
For H2SO4/H2O, the original source is quoted as Ryabova, 1966, table 2.1, page 37. Example values
80 % - -7,52
85 % - -8,29
90 % - -9,03
94 % - -9,59
96 % - -9,88
98 % - -10,27
99,0 % - -10,57
99,7 % - -11,01
99,9 % - -11,43
100,0% - - 11,94
For SO3/H2SO4, the original source is Vinnik, 1966, table 2.2, page 39. Again examples:
1,0 % - -12,43
2,0 % - -12,62
7,0 % - -13,04
10,0 % - -13,21
20,0 % - -13,62
36,0 % - -14,13
So... the H0 difference between 96 % H2SO4 and 20 % oleum is 3,74 units. Whereas the difference from 80 % to 96 % H2SO4 is 2,36 units.

[Edited on 16-1-2015 by chornedsnorkack]

quantumcorespacealchemyst - 17-1-2015 at 03:42

i had originally started by putting a small piece of terbium in water and hoping to get terbium hydroxide. warm water i read to speeds this up, i heated it and after a while realized the bubbles were most all nucleation. the outside tarnished and after days i decided to acidify it with 98% sulphuric acid to get the sulphate. it sat to dissolve and i got back to it a few days later. after leaving that for a week or so i decided to add some potassium permanganate (i forgot exact measurements) and some shavings of gold (fractions of a gram). on a hotplate after adding enough sulphuric acid to be somewhere from 2-4times the amount of liquid from before. when the permanganate was added, it made nice looking purple swirls (stirbar) which fizzed to clearness. it was heated moderatly and ceased when i felt it was far enough, having read that the heptoxide is form-able. there was no green precipitate at the time if i recall correctly but it did form on standing overnight at least, if not before that after standing.

the flask had a rim of olive green moist precipitate, a banding of yellowy, slight olive plating/precipitate and olive with some brownish precipitate collected on the bottom with the stir bar. funny enough, the slight powder of gold seems to have not dissolved, while earlier it had seemed to, perhaps precipitated out.

around the time (at or after) starting the thread
i took pictures swirled it some and let it sit, and now most of the plating that was yellow seems to have fallen to the bottom and/or dissolved

the liquid was a light olive hue and seems the same now or somewhat darker but not by much.

while tentative to swirl the mix, i read that the explosive decomposition of manganese heptoxide is above 55degrees celcius and the flask has been at lower than standard room temperature and in an absence of oxidizable material (besides possibly gold which would seem to be slowly ionized, and moreover seems to be undissolved still)

photos
1,2 before swirling/letting sit overnight.
3,4 after settling overnight, swirling and settling for a few minutes.
5,6 after swirling and taking pictures of turbidity.
7 bottom view after settling a few minutes, gold colored powder/gold under olive/brown(slight). the bottom powder being denser, moves less with sway, and doesn't disperse much [maybe a salt rather than gold])

[Edited on 17-1-2015 by quantumcorespacealchemyst]

j_sum1 - 17-1-2015 at 04:16

I am reading but not understanding. What were you aiming to achieve by adding the gold? And how is this all related to the terbium?

quantumcorespacealchemyst - 17-1-2015 at 10:26

, yea, i had been trying to dissolve gold all sorts of unsuccessful ways. i was originally intending to make terbium hydroxide, and neutralize HAuI4 with it, and did not end up with any success making either. then i decided to turn the terbium to a sulphate, to see the cool florescence, i decided to see if a salt from permanganate or any more interesting species might result that would also be fluorescent. figuring the gold might be able (a long shot) to make a complex salt, double salt. i added the gold, figuring that with H2SO4 and KMnO4, it may dissolve.

i read that H2SO4 and KMnO4 dissolve gold together but not much more information besides that Mn2O7 can/does result. with a small amount, i figure it may not precipitate as it would attach to gold hopefully. i also read that gold sulphate does not form as a salt, but can exist ionically. i believe that was from looking into H2O2 and H2SO4. i actually didn't realize till much later that all my attempts with H2O2 were 3-5%, not 30% as i thought. around 2007 i had used 30% ecover H2O2, it turns out now that bleach substitute is watered down/regular 3-5% seventh generation 7% ecover, H2O2. good news for reformulation though

thank you all for the information so far. it turns out that MnO3Cl is not stable, and that MnO3F is not above 0°C. i don't know, please help me comprehend this, if fluorine is the most electronegative anion and that is what governs the most of the attraction/bonding in an ionic compound/salt such as MnO3+ F-. i also don't understand what, if possible with a ionic species, makes it possible to get past the valance electrons and to be covalent/stronger bonded. i mean i don't understand the orbitals and if/how they can be shuffled/edited/tweaked.

as far as the MnO3F, i read that it is stable under reflux (temperature not mentioned) in IF5

quoting Inorganic Reactions and Methods, The Formation of Bonds to Halogens, Part 2
(edited?) By J. J. Zuckerman
https://books.google.com/books?id=OcD4P1ZeXdAC&pg=PA274&...

"interestingly, MnO3F is stable under IF5 solution under reflux"

full related quote > "Permanganyl flouride, MnO3F, is unstable above 0°C but can be prepared^[63] by reacting KMnO4 with either HSO3F or HF;

KMnO4 + 2HF ---> MnO3F + KF + H2O (m)
KMnO4 + 2HSO3F ---> MnO3F + KSO3F + H2SO4 (n)

In addition to using HSO3F or HF, IF5 may be used^[64] ; interestingly, MnO3F is stable under IF5 solution under reflux. In all preparations the MnO3F needs to be further purified by vacuum distillation; if HF is present then a second treatment is needed using xs KMnO4 or KF, which reacts with the HF to form MnO3F or KHF2, respectively^[63]."

[Edited on 17-1-2015 by quantumcorespacealchemyst]

[Edited on 17-1-2015 by quantumcorespacealchemyst]

quantumcorespacealchemyst - 23-1-2015 at 00:51

i just realized that the equation from wikipedia seems to be incorrect

http://en.wikipedia.org/wiki/Manganese_heptoxide
"Mn2O7 can react further with sulfuric acid to give the remarkable cation MnO3+, which is isoelectronic with CrO3:[citation needed]

Mn2O7 + 2 H2SO4 → 2 [MnO3]+ [HSO4]− + H2O "

Mn2O7 + 2H2SO4 ---> 2[MnO3]+ + 2[SO4]-2 + H2O doesn't charge balance
Mn2O7 + H2SO4 ---> 2[MnO3]+ + [SO4]-2 + H2O seems to work

since MnO3 does that bridge mechanism, with the O3Mn-O-MnO3
and sulphuric acid can form to oleum> pyrosulphuric/disulphuric acid
with HO3S-O-SO3H,
do you think, assuming water is removable somehow from formation of permangany, that they can trade endings forming two O3Mn-O-SO3H ?

reconfiguring with pyrosulphuric acid in the formula,
Mn2O7 + H2S2O7 ---> 2H2SMnO7 ?
O3Mn-O-MnO3 + HO3S-O-SO3H ---> HO3S-O-MnO3 + O3Mn-O-SO3H

chornedsnorkack - 23-1-2015 at 02:09

Quote: Originally posted by quantumcorespacealchemyst

i just realized that the equation from wikipedia seems to be incorrect

http://en.wikipedia.org/wiki/Manganese_heptoxide
"Mn2O7 can react further with sulfuric acid to give the remarkable cation MnO3+, which is isoelectronic with CrO3:[citation needed]

Mn2O7 + 2 H2SO4 → 2 [MnO3]+ [HSO4]− + H2O "

How is this incorrect?

Quote: Originally posted by quantumcorespacealchemyst

Mn2O7 + 2H2SO4 ---> 2[MnO3]+ + 2[SO4]-2 + H2O doesn't charge balance
Mn2O7 + H2SO4 ---> 2[MnO3]+ + [SO4]-2 + H2O seems to work

It balances. Does not mean it works. HSO4- is much weaker acid than H2SO4.

Quote: Originally posted by quantumcorespacealchemyst

since MnO3 does that bridge mechanism, with the O3Mn-O-MnO3
and sulphuric acid can form to oleum> pyrosulphuric/disulphuric acid
with HO3S-O-SO3H,
do you think, assuming water is removable somehow from formation of permangany, that they can trade endings forming two O3Mn-O-SO3H ?

Yes. MnO3+HSO4-. The question is then only, how far is permanganyl hydrogen sulphate an ion pair, and how far it is covalently bonded.

Are the olive green solutions of permanganate in sulphuric acid solutions of MnO3+ cation, or solutions of molecular and undissociated Mn2O7? When that olive green solution is electrolyzed, does Mn travel to anode (as MnO4-), or cathode (as MnO3+), or does it remain stationary as undissociated Mn2O7 while H2SO4 alone is electrolyzed (as HSO4- on anode and H3SO4+ on cathode)?

Quote: Originally posted by quantumcorespacealchemyst

reconfiguring with pyrosulphuric acid in the formula,
Mn2O7 + H2S2O7 ---> 2H2SMnO7 ?

That one does not balance hydrogens. Permanganyl hydrogen sulphate is HMnSO7, not H2MnSO7.

Quote: Originally posted by quantumcorespacealchemyst

O3Mn-O-MnO3 + HO3S-O-SO3H ---> HO3S-O-MnO3 + O3Mn-O-SO3H

quantumcorespacealchemyst - 25-1-2015 at 11:26

No, the first one, its 2H2SO4 but there is only 1HSO4- and H2O. leaving
4 hydrogens on the left, 3 on the right

that equation is from Wikipedia and needs correction. I don't know the correct mechanism.

You are right about permanganyl hydrogen sulphate being HSMnO7 H2SMnO7

And thanks for the tip on analysis by electrolysis. Do you know if these ions are stable to current? I am guessing platinum or iridium electrodes are ideal.

Furthermore, I forgot before that the gold was minimal dust left over from experiments, I dumped it in after adding the KMnO4 to the sulphuric acid, terbium sulphate, as it was left over and to light to weigh on the milligram balance.

I recovered it with some powder and realized again how little there was.
In most, the step was probably a 50:50 ratio of 98% H2SO4 to water with a little, maybe .6g terbium added, later 1/2 to 2/3 H2SO4 by volume added and then about .5-1g KMnO4.

The step I did yesterday was take a mystery liquid, please take this seriously, and add it with precaution, outside with a snow bath. i will describe the liquid below. It turned color to partly darkish amber and overnight turned darker amber. all precipitate dissolved save for the gold dust (I assume most to all gold recovery) and some grains of something I don't know.

I had originally attempted dissolving silver in the liquid prior to addition of the mystery liquid. the silver was caked woth some grains which seemed to leave intact after decanting to remove the silver.

the mystery liquid was originally around 50ml of 3-5% H2O2. into it was bubbled the HCl gas produced by reacting 1.68g 1,4-dibromobenzene with >/~1ml CH3COCl, (1.318g AlCl3, 3.5hrs).
a bit of acetyl chloride efficient condensing and the volatile oil produced (assuming 1,4-dibromoacetophenone).

nonetheless, the actual liquid seems to have been HCl. the H2O2 being exhausted to send Cl2 to NaOH solution dissolving in remaining water and dehydration over time leaving concentrated solution which decomposed sodium hydrogen carbonate. solution was partly cloudy maybe due to acetyl chloride and or 1,4 dibromoacetophenone contamination.

it is all dissolved though, all the green KMnO7 is dissolved it seems

[Edited on 25-1-2015 by quantumcorespacealchemyst]

quantumcorespacealchemyst - 25-1-2015 at 23:33

this stuff is really dark now.

[Edited on 26-1-2015 by quantumcorespacealchemyst]

chornedsnorkack - 26-1-2015 at 09:38

Quote: Originally posted by quantumcorespacealchemyst

No, the first one, its 2H2SO4 but there is only 1HSO4- and H2O. leaving
4 hydrogens on the left, 3 on the right

that equation is from Wikipedia and needs correction. I don't know the correct mechanism.

In this case, you would also have 2 sulphurs on the left, 1 on the right.
The equation would balance if number 2 applied to entire MnO3+HSO4-, not just MnO3+. Could the equation be presented in a way that is ambiguous or incorrect?

quantumcorespacealchemyst - 28-1-2015 at 16:13

i guess it may be Mn2O7 + H2SO4 ---> H+, MnO4-, MnO3+, HSO4-

also, while i am unsure now, it seems the mixture lightened a tiny amount. i am unsure what it is. as far as i know its is from an interaction of
Potassium permanganate KMnO4,
Terbium sulfate Tb2(SO4)3,
sulphuric acid H2SO4,
and Hydrochloric acid HCl.

there was a green precipitate (with some brownish (MnO2?) it seemed) that was insoluble most of the time until addition of HCl. i read HCl acts on MnO2 to Make the chloride (III) but it evolves Cl2, which wasn't observed. also though, the contact of MnCl3 with Cl- makes different negative ions of Manganese chlorides, although this would need the Cl2 made in the first place. the flask was in snow and it is possible, i think that the cl2 was dissolved as it was made back to HCl and HOCl, itself possibly decayed to HCl, so this seems possible. also this put the green stuff, the Mn2O7, in solution.

also the amount of HCl added did not seem sufficient to act on the Mn2O7/HMnO4.

[Edited on 29-1-2015 by quantumcorespacealchemyst]

mnso4 industrial route

learner1112 - 8-2-2015 at 10:33

mso4 industrial route

i am trying to figure out an economical industrial route of processing mnso4 from natural ore. essentially mno2. the first process is to burn sulphur and inject the so2 formed into the crushed ore and water solution. the slurry thus obtained is kept at ph4-4.5 by adding lime. after the reaction the slurry is filtered and the liquid which is pink in colour is obtained. now the problem is how to seperate the mnso4 present in the solution. can anyone guide me, as when i try to heat the solution its ph goes down and it becomes acidic with evolution of so2 gas and probably the mnso4 is converted baqck to mno2

blogfast25 - 8-2-2015 at 19:11

It would help if you quit with the 'txtinglish' and started using proper chemical symbols.

But you seem to have it all wrapped up: SO2 indeed reduces a slurry of MnO2 to MnSO4. Filter off insoluble/residue and you have a solution of crude MnSO4. This solution will only be faintly pink if it is very concentrated.

Your pH goes down mainly because you're removing excess SO2. It will not revert back to MnO2 upon concentrating.

Evaporate to dryness to obtain MnSO4.H2O, which is faintly pink. But your product thus obtained is likely to be contaminated with iron though and thus darker in colour.

Keeping the pH at 4/4.5 during the SO2 addition step could help keeping most of the iron as insoluble Fe(OH)3 and reduce the level of iron contamination in the final product.

[Edited on 9-2-2015 by blogfast25]

MnO2 + H2SO4 ---> MnSO4 + H2O + O2

quantumcorespacealchemyst - 9-2-2015 at 02:07

I read that MnO2 + H2SO4 ---> MnSO4 + H2O + 1/2O2

Is this correct? If so, why not pulverize the ore, and react the MnO2 from Mn+4 to Mn+2 with sulfuric acid turning O-2 to 1/2O2, filter and evaporate?

MnO2 + H2SO4 ---> MnSO4 + H2O + O2

quantumcorespacealchemyst - 9-2-2015 at 02:38

If there is SO2 it may be S+4 turning to S+6, giving 2e- to Mn+4 (MnO2), making Mn+2 (MnO) with the O-2 attaching to (now) SO2+2--->SO3. Ideally, SO3 + H2O---> H2SO4, which reacts with MnO forming MnSO4 and H2O.

[Edited on 9-2-2015 by quantumcorespacealchemyst]

With the lime, I am unsure what it is exactly, if it is CaO, Ca(OH)2, carbonates or with different beginning parts, Al, Mg, or other stuff.

If they switch anion/cations and are all soluble, I don't know which ones are present in salt crystals and or if they double salt. I recently asked a teacher I had learned from in school a similar question about the nature of precipitating soluble ion solutions and was told that the "activity coefficient" of the ions are needed and some complex mathematics, the "specific ion theory" being what is needed, it seems (http://en.wikipedia.org/wiki/Specific_ion_interaction_theory).

in this case if using Calcium salts, it seems that the low solubility of Calcium hydroxide and especially Calcium carbonate can be used to separate concentrated solutions of MnSO4. I don't know how pure it can be made by that technique.

I hope some people can clarify, verify all this. I haven't done this yet. it seems interesting.

I thought the lime was unnecessary, and then I read what you wrote about SO2 being driven out of solution on heating and now I am puzzled too. Does this happen when you don't add lime? I don't know how soluble SO2 is in water, and if the SO2 is from boiling it out of solution, then I don't know why the pH lowers. It seems odd if it is reversal of Mn+2 as if the 2e- (Mn+2 -> Mn+4 +2e-) comes from the Mn+2, where does it go? if the H2SO4 somehow turns to H2SO3 (S+6 -> S+4), turning to SO2 and H2O with heating, the pH lowering seems odd. I don't know, I hope someone here does.

[Edited on 9-2-2015 by quantumcorespacealchemyst]

[Edited on 9-2-2015 by quantumcorespacealchemyst]

blogfast25 - 9-2-2015 at 10:17

quantumcorespacealchemyst:

It's hard to respond to your garbled junk because although it contains truth it also contains stuff you clearly haven't digested very well, a constant in your posts. I suggest very strongly to start your career as a chemical experimenter by cutting your teeth on simple and satisfying experiments, rather than constantly over-reaching and ending up very confused.

I believe the lime (which would be Ca(OH)2 in this case) is there mainly to control pH. Mn bearing ores often contain iron which you want to keep out of your end-product.

Fe(OH)3 is very, very insoluble, requiring quite low pH to enter solution. Keeping pH at 4 to 4.5 during leaching with the SO2 would keep the iron insoluble, or at least that's the idea.

In contrast, using actual H2SO4 would definitely dissolve all or part of the iron.

[Edited on 9-2-2015 by blogfast25]

learner1112 - 9-2-2015 at 12:35

exactly the lime is used to controll the ph. and so2 is used instead of h2so4 because that way it is cheaper.

quantumcorespacealchemyst - 10-2-2015 at 08:49

why does it give off SO2 on heating with pH lowering?

blogfast25 - 10-2-2015 at 09:25

During the reactive leaching, SO2 is absorbed by the lime in the slurry to form calcium sulphite (CaSO3

and further down the road calcium bisulphite (Ca(HSO3

2

. This combination forms the (primitive) buffer system (pH controller). Once pH is below 5 (or so) the reduction of the MnO2 starts and MnSO4 starts forming and solubilising.

When the reactive leaching step has dissolved all (or most anyway) of the manganese the slurry will still reek of excess SO2. After filtering, this excess sulphite will partly oxidise (with air) to sulphate, part of it will simply escape. Nothing mysterious here...

Come to think of it, it's probably recommended to blow air through the slurry prior to filtering, to flush out the excess SO2 and/or oxidise it to sulphate.

[Edited on 10-2-2015 by blogfast25]

Surprise~

quantumcorespacealchemyst - 10-2-2015 at 09:52

I checked it yesterday, it turns out 14 days exactly since the red/brown picture.

[Edited on 10-2-2015 by quantumcorespacealchemyst]

blogfast25 - 10-2-2015 at 11:03

It would help of course if we knew what the picture was and what exactly you did to get that solution.

quantumcorespacealchemyst - 10-2-2015 at 11:56

1st page

H2SO4, Terbium, KMnO4,
then HCl(impurity?it was cloudy)

the pictures of the flask show the green/yellow insolubles.
HCl put it all in solution. didn't see bubbles.

darkyellow/lightbrown color, darkened alot overnight, and after.
picture of that on bottom of 1st page.
14 days, checked it and took a picture^^^ very light/clear

the HCl was made by gassing dilute H2O2 with HCl from a reaction. it sat a long time and possibly absorbed acetyl chloride and possibly some 1,4-dibromoacetophenone.
it was about a quarter it's original volume (the H2O2) when removed from the reaction vessel about a month later. it was cloudy colorless and decomposed baking soda.

[Edited on 10-2-2015 by quantumcorespacealchemyst]

chornedsnorkack - 22-2-2015 at 05:38

Quote: Originally posted by quantumcorespacealchemyst

I read that MnO2 + H2SO4 ---> MnSO4 + H2O + 1/2O2

Is this correct? If so, why not pulverize the ore, and react the MnO2 from Mn+4 to Mn+2 with sulfuric acid turning O-2 to 1/2O2, filter and evaporate?

1) Oxidizing molecular oxygen is a difficult process, unless the oxidant is extremely strong, like ferrates. MnO2 reaction with H2SO4 would be sluggish
2) SO2 is a cheap substance. Why spend money to oxidize it to pure H2SO4 and then discard the oxygen again, if you can just use SO2 directly to reduce MnO2 much faster?

quantumcorespacealchemyst - 22-2-2015 at 14:00

yea

blogfast25 - 22-2-2015 at 14:19

Quote: Originally posted by chornedsnorkack

1) Oxidizing molecular oxygen is a difficult process, unless the oxidant is extremely strong, like ferrates. MnO2 reaction with H2SO4 would be sluggish

Oxidising molecular oxygen (O2

cannot be oxidised. The only real oxidation states for oxygen are -2 and -1. There are NO positive oxidation states for oxygen.

Going from molecular oxygen (oxidation state 0) to -1 or -2 is a reduction, not an oxidation. Throwing ferrate at it doesn't change that.

[Edited on 22-2-2015 by blogfast25]

j_sum1 - 22-2-2015 at 15:12

Quote: Originally posted by blogfast25

Oxidising molecular oxygen (O2

cannot be oxidised. The only real oxidation states for oxygen are -2 and -1. There are NO positive oxidation states for oxygen.

FOOF???

blogfast25 - 22-2-2015 at 16:06

Quote: Originally posted by j_sum1

FOOF???

C'mon j_sum1. You're right of course but PLEASE look at context. We can also 'oxidise' O to +3 by stripping it of 3 electrons in a machine but on meeting any other atom or ion that state would immediately be reduced. And that's not what he was talking about.

Even in O2F2, which is highly unstable, the bonds are more covalent and clearly unstable.

I'm guessing he ['chorned'] simply got a bit confused between oxidation and reduction.

[Edited on 23-2-2015 by blogfast25]

chornedsnorkack - 22-2-2015 at 21:51

Quote: Originally posted by blogfast25

I'm guessing he ['chorned'] simply got a bit confused between oxidation and reduction.

[Edited on 23-2-2015 by blogfast25]

No - forgot to spell out clearly oxidation TO molecular oxygen.

j_sum1 - 22-2-2015 at 22:17

Quote: Originally posted by blogfast25

Quote: Originally posted by j_sum1

FOOF???

I find the mere concept of FOOF hilarious -- it is just so extreme. And that article never fails to make me laugh. Satan's kimchi indeed.
In any case your assessment of the situation stands. You specified "real" oxidation states of Oxygen. Playing with FOOF is so far out in imagination land that I think it is unlikely to be something real chemists would ever really encounter.

quantumcorespacealchemyst - 15-10-2015 at 21:23

it tuns out, the H2SO4 i was using was 2N, nowhere near the 98% o 95% i thought i had (wrote one o those somewhere). that explains my constant puzzlement about it's peculia properties and failed reactions.

the reaction here evaporated on standing over time leaving a dark green oil, presumably KMn2O7 which i brought outside to dunk in a coldwater container and then added a little KOH.