Sciencemadness Discussion Board

Transition metal permanganates/manganates.

Mixell - 8-5-2011 at 12:23

I got an interest in creating a few transition metal permanganates. Such as zinc/copper/indium/aluminium.
I thought about simple cation exchange, but as it seems, those metal permanganates are more soluble than potassium permanganate (which is the only source of prmanganate anion that I have).
Maybe there is a solvent that will allow a cation exchange due to the low solubility of transition metal permanganates in it or the high solubility of potassium permanganate (as I understand, most of the organic solvents can be ruled out, as the permanganate ion oxidizes them)?

Also, I know potassium hydroxide and elemental oxygen can convert MnO2 to potassium manganate, maybe the same thing can be done with aluminium hydroxide (according to Wikipedia, it is the only metal hydroxide that can survive high temperatures of the hydroxides that I have).

In addition to that, anyone knows the solubilities of zinc/copper/indium/aluminium mangantes in water at various temperatures (I know the manganate ion decomposes in aqueous media, but if the solution is quite alkaline this process is slow and sometimes even non-existent)?

Thank you in advance,
Michael.

Mixell - 8-5-2011 at 12:29

Just remembered that potassium perchlorate is very slightly soluble in water, does anybody know if any of the metal (the metals I have stated) perchlorate's are stable?

woelen - 8-5-2011 at 12:49

There are a few other permanganates that you can make from KMnO4 and other metals. If you have silver ntirate, then you can make AgMnO4 by dissolving AgNO3 and KMnO4 at somewhat elevated temperature and mixing these solutions. Crystals of AgMnO4 will separate.

Another permanganate you can make is CsMnO4 from CsCl and KMnO4. It simply separates as a fine crystal meal when you mix saturated solutions of CsCl and KMnO4.

A third permanganate you can make is Ba(MnO4)2. Simply mix solutions of BaCl2 and KMnO4 and a precipitate of Ba(MnO4)2 is formed. This permanganate is nearly insoluble in water.

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Most metal perchlorates are stable. Actually, perchlorate ion is a rather stable ion. In aqueous solution it is even more inert than sulfate ion. But if you have to start from KClO4 then you are not lucky at all. KClO4, CsClO4 and RbClO4 are the least soluble perchlorates, all others are very or even extremely soluble and with KClO4 as a starting point you won't be able to make any other metal perchlorate. You can distill HClO4 from a H2SO4/KClO4 mix under vacuum, but this is a very dangerous thing to do and only for the most experienced among us, who also have access to the required apparatus.

For an amateur, the only two good starting points for metal perchlorates are HClO4 (max. 70%) or to a lesser extent NH4ClO4. The latter can be put into reaction with metal carbonates or metal hydroxides for making the metal perchlorates, where ammonia is driven off. This only works for metals which do not readily form complex ions with ammonia. Unfortunately, HClO4 is hard to obtain for most of us. If you are in the EU, then obtaining NH4ClO4 should be no problem.

Mixell - 8-5-2011 at 13:08

I think I will just buy some perchloric acid, I intend KClO4 to form and drop out of the solution, leaving reasonably pure metal permanganate solution.
I don't got silver in any form, and caesium is very expensive...
So the perchlorate method is the optimal one.
Now all is left is to find a ride to the chemical store...

Edit- potassium metaperiodate is even less soluble in water, but the problem is that I will need to make the metal periodates in a very dilute solution, to avoid forming orthoperiodate's. And then I'll need to concentrate the solution to saturation or to isolate the solid product.

Woelen, do you know by any chance if transition metal metaperiodate's are stable on heating (~150C)?

[Edited on 8-5-2011 by Mixell]

blogfast25 - 9-5-2011 at 03:48

Note also that indium and aluminium are not transition metals.

Mixell - 9-5-2011 at 04:46

Yes, I know, they are poor metals, just instead of complicating the post, I put them in the same group.
Thats why I also refrained from using "transition-meta"l too much and using "metal" instead ;)

[Edited on 9-5-2011 by Mixell]

woelen - 9-5-2011 at 12:17

I strongly doubt that a route over periodic acid is a viable route. Periodates are not that stable and I expect many transition metal periodates to be so unstable that they will decompose well below 150 C.

The route over HClO4 looks better. An alternative route may be over NaClO4 or NH4ClO4, which give solutions of NaMnO4 or NH4MnO4 while KClO4 drops out of solution. The latter two might be more easily obtainable than HClO4. Keep in mind though that perchlorate and permanganate are isomorphic and easily form mix crystals (KMnO4 and KClO4 can form mix crystals in any ratio).

[Edited on 10-5-11 by woelen]

AndersHoveland - 11-5-2011 at 15:00

Quote: Originally posted by woelen  
But if you have to start from KClO4 then you are not lucky at all. KClO4, CsClO4 and RbClO4 are the least soluble perchlorates, all others are very or even extremely soluble and with KClO4 as a starting point you won't be able to make any other metal perchlorate. For an amateur, the only two good starting points for metal perchlorates are HClO4 (max. 70%) or to a lesser extent NH4ClO4. The latter can be put into reaction with metal carbonates or metal hydroxides for making the metal perchlorates, where ammonia is driven off. Unfortunately, HClO4 is hard to obtain for most of us. Obtaining NH4ClO4 should be no problem.


It is also easy to prepare your own NH4ClO4 if you have NaClO4. If you want to make the NaClO4 yourself, it does not even need to be separated out from the NaCl solution.

Ammonium Perchlorate
To a boiling solution that contains as much sodium perchlorate and ammonium chloride as will dissolve, allow this solution to cool to room temperature, then cool to around 15degC. Crystals of Ammonium Perchlorate will start precipitating out of solution. NH4ClO4 is only about 10% as soluble as NaClO4.

Ammonium Perchlorate may also be prepared by adding sodium perchlorate to ammonium hydroxide and then bubbling carbon dioxide into the mixture. Sodium bicarbonate will precipitate out, since its solubility is reduced in alkaline solution. The solution will then contain dissolved ammonium perchlorate. This is basically a modification of the Solvay process, but this method is not the most ideal, since the solubility of NH4ClO4 in the solution will not be substantially greater than that of the bicarbonate.

In the manufacture of ammonium perchlorate, sodium perchlorate is used as the starting material, with the ammonium ion being contributed by such starting materials as ammonium chloride, or ammonium sulfate. Reaction of sodium perchlorate and ammonium chloride:

NaCIO4 + NH4Cl ---> NH4CIO4 + NaCl
(dissolved) (dissolved) (solid precipitate) (dissolved)

separated by fractional crystallization. The solubility of sodium chloride varies only slightly with temperature, whereas that of ammonium perchlorate is temperature-dependent. Thus, a saturated solution can be cooled with very little sodium chloride precipitating as contaminant.
This reaction takes advantage of the fact that ammonium perchlorate has an unusually low solubility. Almost all other ammonium salts are more soluble than their sodium counterparts.

One other note worth mentioning, sodium ferrate can be used to oxidize chlorate (or a limited quantity of chloride) to perchlorate. Note that the usual method of preparing ferrate from hypochlorite is problematic, as the sodium ferrate must then be separated from the chloride ions, and the sodium salt is harder to isolate from this solution than the potassium salt.

Preparation of Sodium Ferrate
Twenty-four parts of Fe2O3 and 40 parts of Na2O by weight were well mixed in a porcelain mortar in the absence of water or carbon dioxide. The mixture was transferred to a refractory vessel and placed in a furnace at 150°C through which dry oxygen flowed. In the next 30 minutes the temperature was increased to 450°C and yielded sodium ferrate(IV) after one hour.
(4)Na2O + Fe2O3 + (1/2) O2 —> (2)Na4FeO4
The product contained only around fifty percent sodium ferrate(lV). Sodium ferrate(lV) disproportionates in water or alkaline solutions according to the reaction:
(3)Na4FeO4 +(8)H2O —> Na2FeO4 + (2)Fe(OH)3 + (10)NaOH


[Edited on 11-5-2011 by AndersHoveland]

copper permangante

bismuthate - 28-9-2013 at 08:40

i want to make some copper permangante. could i use a CuSO4+KMn04 method?

eidolonicaurum - 4-1-2014 at 02:47

Check the solubilities of copper sulphate, ptassium permanganate, potassium sulphate, and copper permanganate. If one chemical has a lowest solubility, the equilibrium will be more on that side (poor English i know). If it will not work in water, try other solvents.

An even easier method?

Brain&Force - 8-1-2014 at 17:37

This picture claims to show permanganic acid in solution. As far as I know the acid is unstable and reduces to lower manganese compounds, but I have no clue beyond that. Can someone confirm this? And if so, how could you make pure permanganic acid?
If this acid is stable, it may be possible to synthesize most metal permanganates by neutralizing the metal oxide. On the other hand it is likely unstable at significant concentrations.

woelen - 9-1-2014 at 00:14

Permanganic acid is quite stable in dilute solution, but you cannot obtain it in the pure state.

Making permanganates from other metals than potassium can also be done through Mn2O7, but this route is dangerous (risk of explosion). You could try it at a small scale.

Mn2O7 can be made by reacting concentrated H2SO4 with KMnO4. A dark oil separates. This oil can be added to water. That gives you a solution of HMnO4. This solution then can be reacted with a metal carbonate or metal oxide.

Please note: Mn2O7 is very unstable and very easily explodes, sometimes even without apparent reason!! Use face and body protection and thick gloves! Do not make more than a few tenths of milliliters of Mn2O7 in one batch! Work behind a shield of plexi glass.

chornedsnorkack - 9-1-2014 at 01:43

Quote: Originally posted by woelen  
Permanganic acid is quite stable in dilute solution, but you cannot obtain it in the pure state.

Making permanganates from other metals than potassium can also be done through Mn2O7, but this route is dangerous (risk of explosion). You could try it at a small scale.

Mn2O7 can be made by reacting concentrated H2SO4 with KMnO4. A dark oil separates. This oil can be added to water. That gives you a solution of HMnO4.


Will Mn2O7 and saturated HMnO4 solution in water separate as immiscible liquids when a small amount of water is added to excess of Mn2O7? And what is the concentration of HMnO4 saturated with respect to Mn2O7?

woelen - 9-1-2014 at 02:25

I do not know. I never made so much Mn2O7 that this was in excess relative to some added water. I know from literature that HMnO4 is not stable when the concentration goes above a certain value. It decomposes to MnO2, O2 and H2O. So, I expect that addition of a small amount of H2O to Mn2O7 will lead to complete decomposition of Mn2O7. Addition of a lot of water to Mn2O7 leads to formation of a dilute solution of HMnO4. The dilute solution is stable, because HMnO4 is a strong acid and splits in H(+) and MnO4(-) ions, which are stable in solution. At very high concentration there is molecular covalently bonded HMnO4 and this decomposes.