Sciencemadness Discussion Board
Not logged in [Login - Register]
Go To Bottom

Printable Version  
 Pages:  1  2    4  ..  16
Author: Subject: Permanganates
Eclectic
International Hazard
*****




Posts: 899
Registered: 14-11-2004
Member Is Offline

Mood: Obsessive

[*] posted on 19-5-2007 at 10:21


Does fusion of MnO2 with NaNO3 work?

[Edited on 5-19-2007 by Eclectic]
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 19-5-2007 at 11:02
Permangantes


The last time I tried any fusion reactions was about 25 yrs ago, so no guarantees from the management.

It does, according to some scrappy notes I have from that era, but you have to add NaOH. So does KClO3. You get the green mangante in both cases. I suspect KClO4 would also but didn't have any available then. The trouble always is preserving the manganate. Permanganate and manganate decompose at between 200-250 C yet the fusion temps are all higher than this. I don't think you'd get anything without the alkali.

Another oddity is that the probable product of the nitrate oxidation is to nitrite, a reducing agent.
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 19-5-2007 at 11:52


Finally, I’ll throw in a method I tried earlier. It differs little from the above and still suffers from the same permanganate/chlorate problem. If anyone can find a solution to this problem, please post it! The method of using Ag is obviously impractical for nearly everyone.

Basically, you mix stoichimetric proportions DRY: 4 moles MnCO3, 2 moles Ca(ClO)2, 1 mole Na2CO3 and hope that

4MnCO3 + 3Ca(ClO)2 + Na2CO3 --> 2NaMnO4 + 3CaCl2 +Mn2O3 +5CO2

(Note: anyone can write an equation like this! The laws of physics determine whether it works. There is no doubt that all the Mn carbonate does not react and a black mess is also left, which may be Mn2O3 or MnO2. Also, the calcium is precipitated as carbonate, which makes it gray. K can replace Na, of course)

Add water and warm to 40-50C in a beaker on a water bath for a fair time, stirring frequently..

Chlorate is produced again, of course. It only avoids the messy step of actually making the NaOCl.

NOTE:- Further research shows that using Lithium instead of K or Na does allow a separation : the permanganate is very soluble at about 70g/100g aq but the chlorate is extremely soluble at about 400g/100g aq. Not too easy but practical, but most of us don’t have lithium (pyro enthusiasts might! I’ve got a few gms carbonate)

I do hope I did not disappoint too many. If there was an easy cheap method it would be used industrially, you bet. Electrolysis in KOH or K2CO3 on to a MnO2 anode should work but making a conductive one and holding it together is a challenge.

A tit bit for Mn enthusiasts. Have you ever seen a hypomanganate like K3MnO4? Azure blue in color. Very careful reduction of pure KMnO4 with sodium sulphite at around 0C in a dilute solution will take you through the steps MnO4- (Red/purple) --> MnO4-- (deep green) --> MnO4--- (sky blue) --> MnO2 (black ppt)
Regards,

DerAlte
View user's profile View All Posts By User
Eclectic
International Hazard
*****




Posts: 899
Registered: 14-11-2004
Member Is Offline

Mood: Obsessive

[*] posted on 19-5-2007 at 13:31


OK, so a nitrate/nitrite mixed melt with some -OH melting at about 125 C should work nicely, followed by electrolysis to bring up the oxidation state? Or disproportionation of the manganate?



[Edited on 5-19-2007 by Eclectic]
View user's profile View All Posts By User
The_Davster
A pnictogen
*******




Posts: 2853
Registered: 18-11-2003
Member Is Offline


[*] posted on 19-5-2007 at 14:10


Simple acidification of the manganate, even using CO2, will work causing the disporportionation.



View user's profile View All Posts By User
Eclectic
International Hazard
*****




Posts: 899
Registered: 14-11-2004
Member Is Offline

Mood: Obsessive

[*] posted on 19-5-2007 at 14:36


But I LIKE electricity! :D

I like the acidification with CO2 though. For making a strong NaMnO4 solution to prep a wide variety of permanganate salts, you could drive the excess Na out of solution as bicarbonate, and get rid of the very soluble nitrite at the same time. Excess NaNO3 should crystallize out by common ion effect, leaving fairly pure NaMnO4 solution.

Agricultural MnSO4 + NaHCO3 + agitation/air --> MnCO3
MnCO3+2NH4NO3 +heat --> Mn(NO3)2 + NH3 + NH5CO3 (volatile)

(If you try to crystalize out the Mn(NO3)2 from solution you get a double salt with some ammonium, not really a problem though if you are going to add NaOH and heat. The whole mess should dissolve in it's own water of crystallization below 100 C.)

[Edited on 5-19-2007 by Eclectic]

It looks like LiOH would work even better for making an extremely concentrated permanganate solution (70%+), and the carbonate has low solubility. :D

[Edited on 5-20-2007 by Eclectic]
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 20-5-2007 at 12:05


I Like electrolysis very much too, Eclectic! I guess everyone who has done the famous chlorate process also does..

Anodic oxidation of ferro manganese in KOH or K2CO3, gives manganate or permanganate (plus some MnO2, I believe). I tried it with an old piece of spiegeleisen (bright silvery stuff, Fe80/Mn20) a long time ago and that didn't work very well, if memory is OK.


Ferromanganese is an alloy 20Fe/80Mn which should be available on the web, as steelworks make it by the tons, but a Google serch revealed no sources. Mn metal is expensive but you can find it, but ferromanganese ought to be cheap. Anyone know a source by the kg instead of the ton?

I was hoping, by opening this thread, to see if some bright young spark had any radically new ideas on MnO4- production. This tired old brain is now out of ideas!

DerAlte
View user's profile View All Posts By User
Eclectic
International Hazard
*****




Posts: 899
Registered: 14-11-2004
Member Is Offline

Mood: Obsessive

[*] posted on 20-5-2007 at 12:52


If you are determined to play around with metallic manganese, you could probably electrowin your own from MnSO4 with an excess of MnCO3 on the bottom of the plating tank to neutralize acidity.

55lbs of agricultural MnSO4 can be had for $30-40, $2-3/lb in small quantities.

I wonder what you would get running electricity through a Mn(NO3)2 and NaNO3 solution?
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 5-6-2007 at 19:41


QUESTION

I am still messing around with KMnO4 synthesis. I noticed an old thread was resuscitated a day or so ago. It’s a hardy perennial!

I have a question that has been bugging me for a few days. Does anyone have any reputable NUMBERS for the solubility of NaMnO4 – repeat SODIUM permanganate? I’ve done a CRC and Googled it to hell and gone, and only land up with useless statements like “very soluble” or “extremely soluble” – hell, I know that! I have an old notebook from a past era that states 220g/100g aq at 25C. No reference – I’m bad at that. If anyone has any numbers, please post…

Desperate,

DerAlte
View user's profile View All Posts By User
not_important
International Hazard
*****




Posts: 3874
Registered: 21-7-2006
Member Is Offline


[*] posted on 6-6-2007 at 04:38


The notes I have are similar, only numbers are " greater than 300 g/100 cc" I suspect it is difficult to get a pure dry sample - "the sodium salt is so soluble and hygroscopic that its solubility is too difficult to be measured accurately". Note that it is commonly sold as a 40% by weight solution, which is only 400 g/l
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 6-6-2007 at 20:35


Thanks, not_important! It confirms what I have, near enough for government work.

The reason for asking was because I have been looking at various permanganates and chlorates, to see what ratio of solubilities exist to differentially separate the chlorate produced by the hypochlorite -> chlorate degradation ( disproportionation) inevitable in the process I have been suggesting.

I find the following ratios of solubility, permangante to chlorate - at approx RT -

K 9:10.1: Ba 65:34; Ca 338:209; Na ~260:96; Li 372:71; Ag 15:0.9 g/100g aq

Well, nobody is going to waste silver on a permanganate synthesis, nor are they likely to waste Rb/Cs, which also have rather insoluble permangantes, less than K.

Li looks a promising candidate, but that also sounds a bit too rare (not as rare as it used to be in my youth) so the next best bet is sodium.

I have just performed an all sodium reaction of MnCO3. NaOCl, Na2CO3 as outlined earlier in this thread. So I now have a (well filtered) mixture of NaClO3, NaCl, NaMnO4 plus possibly some Na2CO3. I have given it a brisk boiling to convert all NaClO to NaClO3. At least there is no hypochlorite smell (but that is muted in alkaline solution - it needs actual HClO, as in acid solution, to really give that bleach smell).

I have evaporated the solution to precipitate that Na Cl, the major contaminant. Twice, right down to a fraction of the original volume. One more go has also precipitated considerable amounts of the Na ClO3 if the numbers above are correct. I should have the NaMnO4 as the major component of my solution, which is deep black and impentrable by a focussed flashlight. I shall add KCl tomorrow and - we'll see!

The results of an earlier all K run seemed to produce chlorate and permanganate almost equally. I have not estimate the amount of either except by viewing the crystals under a microscope - needles of the permanganate are interspersed with plates of chlorate.

Regards,

DerAlte
View user's profile View All Posts By User
JohnWW
International Hazard
*****




Posts: 2849
Registered: 27-7-2004
Location: New Zealand
Member Is Offline


[*] posted on 7-6-2007 at 01:38


As a cheaper substitute for permanganates, for industrial technical purposes such as water disinfection, how about making ferrates(VI) instead, in much the same ways, although preferably by electrolysis of a cold alkaline solution of Na or K ferrite(III)?

BTW An "holy grail" of iron chemistry would be to try to make, under more extreme such conditions, perferrates(VII) and FeO4. A theoretical study I have read somewhere suggests that they could exist. I wonder if anyone has succeeded.
View user's profile View All Posts By User
not_important
International Hazard
*****




Posts: 3874
Registered: 21-7-2006
Member Is Offline


[*] posted on 7-6-2007 at 05:07


Theory on FeO4(-) : Inorg. Chem. (1999) 38(22): 4942–4948
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 10-6-2007 at 23:16


PERMANGANATES

This will be my last contribution to this thread, unless there is something worth commenting on. The results of attempting via the all Na route still resulted in significant chlorate in the product on adding KCl and messing around with the very concentrated solutions was a pain.

I tried a one pot last effort, using MnCO3, calcium hypochlorite (HTH approximation as above) plus the calculated amount of sodium carbonate to precipitate the calcium AND satisfy the stoichiometric equation given earlier. This seemed as efficient as bothering to make sodium hypochlorite separately. The solid components I ground together in a mortar before adding water. Heated at about 55-60C for two hours followed by a 3/4 hour boiling to disproportionate remaining hypochlorite to chloride and chlorate. Filtered of the calcium carbonate and unreacted Mn products using glass filter. A titration of diluted solution (acidified by HCl) against ferrous sulphate gave the MnO4- in the product as 39% of the ideal yield.

Then reduced to half volume to precipitate NaCl, refiltered, and repeated to precipitate sodium chlorate. Very tedious! Adding a saturated solution of KCl actually gave a product that looked like KMnO4 but with very small elongated crystals when cooled to 0C for a while…

Haven’t yet got to drying but a small dried portion acted like KMnO4 when mixed with carbon and ignited. But then again, chlorate might do the same except this crackled in a way more reminiscent of permanganate. There’s still chlorate there, but how much I haven’t attempted to find out.

If 39% of the hypochlorite went to making MnO4- and the rest disproportionated, the ratio of MnO4- ion to ClO3- ion would be 3.9: 3.7 I calculate. That is, prior to attempts to separate. I think this effort to make permanganates has left me exhausted!

Regards,

DerAlte

[Edited on 12-6-2007 by DerAlte]
View user's profile View All Posts By User
JohnWW
International Hazard
*****




Posts: 2849
Registered: 27-7-2004
Location: New Zealand
Member Is Offline


[*] posted on 11-6-2007 at 14:16


With regard to my post above about ferrate(VI), FeO4(2-) and the possible Fe(VII) and Fe(VIII) oxidation states of Fe as FeO4(-) and FeO4, and possible uses as a cheaper substitute for permanganate, Solo has come though for me with the following post, with a PDF article, on a different thread in the References section at:

http://www.sciencemadness.org/talk/viewthread.php?tid=7204&p... :

posted on 12-6-07 at 08:19

Theoretical Studies on the Higher Oxidation States of Iron
M. Atanasov
Inorg. Chem. 38(22): 4942–4948 (1999)

Abstract
Density functional theory (DFT) and multiconfiguration self-consistent field (MCSCF) calculations on the oxo FeO42- (FeVI) and the hypothetical oxo FeO4- (FeVII), and FeO4 (FeVIII) and peroxo FeO2(O-O)z [z ) -2 (FeIV),z ) -1 (FeV), z ) 0 (FeVI)], Fe(O-O)2z [z ) -2 (FeII), z ) -1 (FeIII), z ) 0 (FeIV)], and FeO(O-O)2z [z ) -2 (FeIV), z ) -1 (FeV), z ) 0 (FeVI)] clusters are presented and discussed. The results show the potential of stabilizing FeVII and FeVIII in tetrahedral oxo coordination. On the basis of absolute electronegativities calculated using DFT, it is predicted that FeO4 will be rather oxidizing, even stronger than Cl2 and O2. On the basis of a comparison between total bonding energies of M1M2FeVIO4 (M1, M2 ) Li, K), MFeVIIO4 (M ) Li, K), and FeVIO4 clusters, possible synthetic routes for electrochemical preparation of FeO4- and FeO4 species are discussed.

Attachment: Theoretical Studies on the Higher Oxidation States of Iron .pdf (141.03 KiB)
http://www.sciencemadness.org/talk/viewthread.php?action=att...
This file has been downloaded 1 times

[Edited on 12-6-07 by JohnWW]
View user's profile View All Posts By User
12AX7
Post Harlot
*****




Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline

Mood: informative

[*] posted on 11-6-2007 at 17:53


I would think the Cs or Ba salt would be most stable; maybe cold anodic oxidation of a suspension of BaFeO4?

Hmm, how does the solubility of Cs2FeO4 compare to the K salt...

Tim




Seven Transistor Labs LLC http://seventransistorlabs.com/
Electronic Design, from Concept to Layout.
Need engineering assistance? Drop me a message!
View user's profile Visit user's homepage View All Posts By User This user has MSN Messenger
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 12-6-2007 at 19:05


Above, On 10-6-07 I wrote:

"A titration of diluted solution (acidified by HCl) against ferrous sulphate gave the MnO4- in the product as 39% of the ideal yield."

Looking over my calculations, I find I omitted the 7H2O of the ferrous sulphate.

The corrected figure is 21% yield based on Mn, pretty poor. 39% didn't sound too bad! This accounts for the difficulty of separation.

DerAlte
View user's profile View All Posts By User
ciscosdad
Hazard to Self
**




Posts: 76
Registered: 6-2-2007
Member Is Offline

Mood: Curious

[*] posted on 12-6-2007 at 20:59


Fascinating Thread DerAlte.

What about some other means of reducing the xs ClO that will not produce ClO3?
IIRC Urea will reduce it forming N2 and CO2.
If the Urea does not interfere with the MnO4 that may work. It also has the advantage of not introducing any extra ion species to the final mix (with care).
Can anyone else think of some selective reducing agent should the Urea be unworkable?

On further reading about Permanganates, it seems very unlikely that any reducing agent will be unreactive with it.
Sigh


[Edited on 13-6-2007 by ciscosdad]

[Edited on 13-6-2007 by ciscosdad]
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 13-6-2007 at 18:42


Ciscosdad, thanks for the compliment. But I suspect that KMnO4 would oxidize urea all too easily.

You raise an interesting point. In acidic solution both permanganate ion and hypochlorite are much stronger oxidants than in alkaline solution, but still almost equal in redox potential. In alkaline solution, hypochlorite is marginally the better oxidant – hence the difficulty in carrying out the suggested synthesis rapidly and efficiently. Hypochlorite would just as easily oxidize itself to chlorate as do what we want.

Regards,

DerAlte
View user's profile View All Posts By User
ciscosdad
Hazard to Self
**




Posts: 76
Registered: 6-2-2007
Member Is Offline

Mood: Curious

[*] posted on 14-6-2007 at 15:40


Interesting Patent.

http://www.patentopedia.us/process_manufacturing_chlorine_di...

Not a method to be attempted in the backyard I guess
:(

I'm still looking for details of the the method that Condy used in the last part of the 1800's. The encyclopedia references I've seen say that the method was simple enough that he had problems with patent infringers. There's got to be some clues there.

I wonder if a pot of 40%KOH/MnO2 kept at ~200 DegC and with air bubbled through for a few days would be possible?
If the pot was a tallish tube, well insulated and heated by resistance wire, the heating costs should not be excessive.
The fumes/mist from the hot KOH would be a problem, but not insurmountable. Evaporation could be checked on by inspection and (careful ) topping up

Just mental doodling. I need to think more on this.
I also need to check if urea solution decolourises KMnO4.
View user's profile View All Posts By User
12AX7
Post Harlot
*****




Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline

Mood: informative

[*] posted on 14-6-2007 at 16:14


As I recall, I once tried to oxidize urea with permanganate, getting nowhere, having let it sit for a week or two at room temperature. I don't remember if I did anything to the pH, but it shouldn't need any if it's as good an oxidizer as it claims to be!

Tim




Seven Transistor Labs LLC http://seventransistorlabs.com/
Electronic Design, from Concept to Layout.
Need engineering assistance? Drop me a message!
View user's profile Visit user's homepage View All Posts By User This user has MSN Messenger
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 14-6-2007 at 23:44


ciscosdad , I believe Condy used the time honored fusion with KOH. and oxidation by air. Apparently even the alchemists knew of the process.

Regards,

DerAlte
View user's profile View All Posts By User
ciscosdad
Hazard to Self
**




Posts: 76
Registered: 6-2-2007
Member Is Offline

Mood: Curious

[*] posted on 17-6-2007 at 15:27
Condys process


DerAlte,
You are almost certainly right about the basic method, but as they say, the devil is in the details. There should be a patent listed where he specifies exactly what he does to get from Pyrolusite to the solid crystals. Buggered if I can find it though.

One further thought re Urea: It will be of llimited use even if Permanganate does not oxidise it as there will likely be a significant amount of ClO3 formed during the initial synthesis (ClO + heat). So we're still stuck with the main impurity.

I still like the idea of the steel reactor with liquid KOH and MnO2 +O2. Next is a boiling point/concentration graph for KOH to see what temps will work.
More doodling!
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 17-6-2007 at 20:27


ciscosdad, you have it. We do not want to reduce the hypochlorite but the chlorate. And chlorate is a weaker oxidant than hypochlorite.

I have been doing some thinking and have come up with a method to somewhat ameliorate the situation. I will hold my water on that one until I have completed my latest experiment.

If any are curious as to how I manage to do these somewhat lengthy tria runs, the answer is I am retired and have been for years (and years and years!) I happen to be alone at present, which makes it even easier... and chemistry is still fun, even at my age.

Notice that (provided you start with potassium hydroxide and DON"T use chlorate as an oxidant) the traditional fusion does not give any products that interfere with the production of the manganate. Not so with the wet methods here. Chlorate production is inevitable and must be minimized. Other oxidants such as persulphates might work but it seems pointless to use a more expensive one to make a permanganate.

The other method that I would like to try is electrolysis in a divided cell using a MnO2 anode. I have tried this in a hit and miss fashion and you do get the purple coloration. The problem is making a suitably conductive anode. Any ideas, anyone? The batteries using MNO2 use graphite. Electrolysis is a favorite of mine because almost anyone with a bit of knowledge can do it, and it allows us to achieve 'reactions' that otherwise require high energy reactants. Just what the amateur requires!

Regards,

DerAlte
View user's profile View All Posts By User
12AX7
Post Harlot
*****




Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline

Mood: informative

[*] posted on 17-6-2007 at 21:10


How about some anode in KOH with a suspension of MnO2?

Would graphite break down (I know it makes a mess in sulfuric acid), needing say, PbO2? (Hmm, but that would want to try to make K2PbO3, at least without voltage...)

Tim




Seven Transistor Labs LLC http://seventransistorlabs.com/
Electronic Design, from Concept to Layout.
Need engineering assistance? Drop me a message!
View user's profile Visit user's homepage View All Posts By User This user has MSN Messenger
 Pages:  1  2    4  ..  16

  Go To Top