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Author: Subject: Permanganates
vmelkon
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[*] posted on 8-9-2012 at 12:20


I tried the NaClO and MnO2 and NaOH method but it doesn't look like anything has happened. What conditions does it need to work? Or is it just bullshit that appears on the Wikipedia page on NaMnO4.
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[*] posted on 8-9-2012 at 22:17


Have a look at http://www.sciencemadness.org/scipics/MnOXY.doc

The particular section you need is on pages 5-7 entitled WET METHODS. The yield is poor as explained there.

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[*] posted on 9-9-2012 at 17:25


Even if the yield is low, I would like to see the purple color of the damned thing.

I'm not exactly sure what concentrations you used. Did you use 15% NaClO? How much K2CO3 did you use in terms of grams? How many hours did you keep it at 60-70 °C? Did you do it in an open beaker?

Unfortunately, I don't have K2CO3. I'm thinking of using wood ash or Na2CO3.
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[*] posted on 9-9-2012 at 18:54


You can use 5% NaCl. Dissolve as much Na2CO3 in it as you can. For the Mn02 you could use the that from an unused alkaline cell if you don't have any pure enough. I think the .doc I wrote tells you how to make it (hydrated) reasonably pure. If not search the thread. But you will have to keep it at c. 60C for at least two hours and let it settle. Good luck!

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[*] posted on 26-4-2013 at 21:17


MnO2, Mn(VII) - and Mn(VI) - -
I was going to put this into the “Pretty Pictures” thread but decided it wasn’t pretty enough. Instead I’ll tag it on this link, because it illustrates many points WRT manganate production and the chemistry itself. I redid this experiment to show a 12-year old grand daughter.

Every compound used is OTC or made from trash.
The manganese sulphate comes from spent batteries, (process described elsewhere in these pages): also available as plant nutrient.
10-15% NaOCl from pool bleach.
Na2CO3 – use either washing soda powder (anhydrous or baking soda heated >120C, or monohydrate, do not use decahydrate; too much water).
NaOH – can sometimes be found OTC.



Procedure: (Only small amounts are needed). Test pH if you have a meter.

Make the following solutions:
(1) Saturated MnSO4 in H2O (slight pink to yellowish)
(2)Saturated Na2CO3 in 10% bleach. (pH ~ 12)
(3) saturated Na2CO3 in H2O (pH~11.8)
(4) NaOH (as much as possible) in 10% bleach; allow to cool to RT. ( about 10M; pH=14+; & make sure your meter can stand saturated NaOH without damage)
(5) Add about 100mg NaOH per 10ml of 10% bleach to get a solution of pH ~ 13.4
(6) Dilute NaOH in water about 1M. (4g/100ml.) (pH ~ 14)
(7)Saturate NaHCO3 in 10% bleach. (pH ~8.5 to 10+ - depends on bleach which contains hydroxide in small quantities)

Using small test tubes and a dropper,

(a)add a drop of MnSO4 to #3. Pinkish carbonate separates (not shown)
(b) add a drop of MnSO4 to #6. White hydroxide separates, quickly turning brown (not shown)
(c) add a drop of MnSO4 to #2. White carbonate formed rapidly turns black , as Mn02 is formed. Shake or stir and leave for ½ hour. You should get a dilute permangante solution as the MnO2 settles out. Picture first on right.
(d) add a drop of MnSO4 to #5 ; reaction is similar to (c) (2nd tube from right)
(e) add a drop of MnSO4 to #4 ; in this case managante (green) is formed (3rd tube from right)
(f) add a drop of MnSO4 to #7 ; in this case only manganese dioxide is formed ( the slight amount of red is due to a very small amount of NaOH added to the bleach to stabilize the NaOCl) (4th tube from right)
(g) perform reaction (d), Carefully pour the concetrated NaOH/bleach solution (4) down the side of the tube so it sinks to the bottom. Leave for about an hour. The middle section is actually blue due to an admix of manganate and permanganate but does not show up well in the picture (5th tube from right).

The dirty appearance of the tubes is due to MnO2, of course.

See http://www.sciencemadness.org/scipics/MnOXY.doc for an explanation of the chemistry involved.

Der Alte


[Edited on 27-4-2013 by DerAlte]
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[*] posted on 27-6-2013 at 04:32


Sticky...sure!

The best way to make KMnO4 is by oxidation of Mn(ll), e.a. manganese sulphate with potassium peroxodisulphate. NaMnO4 is not stable as a solid but a solution is a good alternative. Most oxidizers tend to stop when stable MnO2 is formed. Heating KOH with MnO2 works but isn't easy since temperatures far above 600 degree Celcius (temp. of a butane/propane burner) are needed to afford high yields.
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[*] posted on 6-8-2013 at 21:21


Can the amount of KNO3 needed be calculated using this?

2 MnO2 + 4 KOH + 2 KNO3 --> 2 K2MnO4 + 2 KNO2
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[*] posted on 7-8-2013 at 12:19


Quote: Originally posted by learningChem  
Can the amount of KNO3 needed be calculated using this?

2 MnO2 + 4 KOH + 2 KNO3 --> 2 K2MnO4 + 2 KNO2


The equation isn't balanced: 14 O on the left, only 12 O on the right. Also: 4 H on the left, none on the right.

To figure it out, determine which oxidation takes place and balance it, then determine which reduction takes place and balance it. Then balance the two against each other.

But a balanced reaction equation still doesn't mean things actually happen that way...

[Edited on 7-8-2013 by blogfast25]




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[*] posted on 8-8-2013 at 13:26


Quote:
The equation isn't balanced: 14 O on the left, only 12 O on the right. Also: 4 H on the left, none on the right.


My bad, should have been :

2 MnO2 + 4 KOH + 2 KNO3 --> 2 K2MnO4 + 2 KNO2 + 2 H2O

Still, that doesn't affect the KNO3 part does it?

Right - I don't know how KNO3 works as an oxidizer - that's what I'm asking...
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[*] posted on 9-8-2013 at 08:27


@learningChem:


Firstly, the reduction of nitrate is likely to go to NO [+II oxidation state], and not to nitrite [+III] like you implied, acc.:

3 MnO2 + 4 KOH + 2 KNO3 === > 3 K2MnO4 + 2 NO(*) + 2 H2O

This saves a bit on nitrate, with respect to your equation.

Chemical reactions proceed (in the right conditions) if the Gibbs Free Energy change ΔG (= Gpostreaction – Gprereaction) of the system is negative, or preferably: STRONGLY negative.

The oxidation of MnO2 with air oxygen:

MnO2 + 2 KOH +1/2 O2 === > K2MnO4 + 2 H2O

… is likely to have a ΔG that is much less negative than those oxidations assisted by oxidisers like nitrates or chlorates, which explains why the air oxidation isn’t really practical. The values of ΔG for many reactions (including these discussed) can be calculated more or less easily, depending on availability of G values for reagents and reaction products.

To further oxidise manganate to permanganate only electrolytic oxidation is practical if you actually want to obtain, solid, pure KMnO4 (although hard to do at the home level, apparently):

MnO<sub>4</sub><sup>2-</sup> === > MnO<sub>4</sub><sup>-</sup> + e<sup>-</sup> (oxidation)

(*) NO will immediately air oxidise to the brown fumes of nasty, chlorine like NO2, so careful with that! NO +1/2 O2 === > NO2



[Edited on 9-8-2013 by blogfast25]




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[*] posted on 9-8-2013 at 18:21


Thanks Blogfast!

For what it's worth, I did a couple of small tests melting KOH, adding the nitrate to the molten KOH, and then adding MnO2. I got 'some' (a bit) of manganate, but didn't see any NO2 (I think)
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[*] posted on 10-8-2013 at 03:43


@ LC:

To make this work you probably need to go 'all in': mix required amounts of MnO2, KOH and KNO3, grind to high degree of homogeneity and heat on full.




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[*] posted on 10-8-2013 at 12:59


I thought about doing that but wasn't too keen on grinding KOH since it gets wet fast. Maybe it can be ground if I put it inside a plastic bag?

What I did try was the versuchschemie procedure. Dissolve KOH and chlorate (or nitrate) in water, add MnO2, mix and dry.

Also, regarding the versuchschemie procedure : google seems to give a usable translation, and the guy says he got 45g permanganate from 40g MnO2, which is pretty reasonable as yields go?

My attemps so far only produced a tiny amount of permanganate needles though...
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[*] posted on 11-8-2013 at 04:52


KOH isn't near as hygroscopic as you seem to think. I've ground down mixtures with KOH many times, in fact I'm working on come mixtures of Cr2O3/KOH/KNO3 right now. Small amounts of water in the mix may even be beneficial to the process anyway.

The 'mix-as-solution' then heat to dry method does give an amazing level of intimacy of mixing but it's more hassle too. I think thorough mechanical mixing/grinding must work as well.

Chlorate is a better oxidiser (in these conditions) than nitrates but old reports confirm KNO3 can oxidise Pyrolusite (MnO2) to manganate (VI), so it's a matter of 'getting it right'.

Which versuchschemie post are you referring too? Work out that reported yield as a percentage Actual Yield, it's much more meaningful as a number (hint: it's about 62 %, which is neither great nor too bad).

And how are you carrying out the final oxidation step, VI to VII?



[Edited on 11-8-2013 by blogfast25]




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[*] posted on 11-8-2013 at 13:06


This is the procedure I was referring to :

http://www.versuchschemie.de/topic,10934,87e5b2c81f17192950a...

Yes, the yield is ~60% with respect to MnO2.

Quote:
And how are you carrying out the final oxidation step, VI to VII?


I'm bubbling CO2 into the manganate solution. I think my biggest failure is in the KOH/MnO2 fusion step though. I'm going to re-read DerAlte's .doc summary now...

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[*] posted on 11-8-2013 at 13:35


How does he separate the KMnO4 from the MnO2 (formed in the acidification step)?

Using CO2 as an acid here is quite clever: with stronger acids the manganate tends to disintegrate back to 100 % MnO2.

The yield here is much lower than 60 %, because 3 moles of MnO2 only give 2 moles of KMnO4. Granted, you get some of your MnO2 back but it should be calculated on a 3 MnO2/2 KMnO4 basis.

Beware that KMnO4 containing residual MnO2 is often considered worthless. Apparently solutions of it degrade too quickly to be of much use. Or so I've read... ;)

[Edited on 12-8-2013 by blogfast25]




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[*] posted on 11-8-2013 at 18:04


Quote:
How does he separate the KMnO4 from the MnO2 (formed in the acidification step)?


Filter the MnO2, concentrate the solution, and let the permanganate crystalize.

Quote:
Using CO2 as an acid here is quite clever: with stronger acids the manganite tends to disintegrate back to 100 % MnO2.


I see. Another advantage may be the fact that you get K2CO3, which is pretty soluble in water? K2CO3 112g/100ml 20C - KMnO4 6.3g/100ml

Picture : I pooled the solutions from 3 or 4 different runs (using nitrate and chlorate) and evaporatedd most of the water. You can see permanganate needles...plus a lot of other crap =P


per1.jpg - 206kB
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[*] posted on 11-8-2013 at 21:51


@LC & Blogger:

Nice to see some action in this old thread!

One point I would make is that using CO2 is no different than any other acid - you still lose in the disproportionation

3MnO4 -- + 4H+ < --> 2MnO4 - + MnO2 + H2O

Anodic oxidation or Cl2 oxidation avoid this.

WRT the MnOXY.doc I wrote way back, I'd make the following correction. Use of PbO2 plus nitric acid as oxidizers of MnO2 or of Mn ++ , if it works at all, is about as useless as using air. I retried it and maybe got a faint coloration, that's all. Yet all the literature seems to say it works. If you follow references back, eg from Wiki, you'll just get other references, going back to the dawn of chemistry!

Of course, I do not claim it cannot work - if conditions are very precise or extreme, such as during an exact phase of the moon.

As to whether "... that KMnO4 containing residual MnO2 is often considered worthless. Apparently solutions of it degrade too quickly to be of much use..", I vouchsafe no opinion. However, the following experiment is instructive.

Convert a manganate solution to a permanganate by adding enough CO2 or H+. Carefully titrate this against hydroxide to increase the pH and the colors are reversed as per above eqn. However if you take permanganate, add some MnO2 and try to reverse it at increased pH, you will not get manganate, although you have the same components available. I think the reason for this has to do with the surface state (activity) of the MnO2 in the two cases.

@LC Looks like you are almost there. Try recrystallization.

Regards, DA

[Edited on 13-8-2013 by DerAlte]
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[*] posted on 12-8-2013 at 08:16


@DerAlte:

What I seem to have understood, perhaps erroneously, is that using CO2 as an 'acid' prevents local areas of very low pH to form (which could arise using stronger acids), which would cause more MnO2 to drop out than otherwise expected.

Re. the stability of KMnO4 solutions, it's a bit of a bullshit subject. One peer reviewed paper I found showed practically no loss of titer of 0.1 N standardised solutions over a 2 year period, provided the starting point KMnO4 is MnO2 free. Make of it what you will, I guess...

[Edited on 12-8-2013 by blogfast25]




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[*] posted on 12-8-2013 at 16:59


I suspect one would have to be fairly careful to maintain 0.1N KMnO4 over 2 yrs. Permanganate has a nasty habit of dropping MnO2 for no apparent reason, especially at lower pH where it's oxidizing power rapidly increases. See the Pourbaix diagrams.

CO2 in solution is sufficiently acidic for manganate to not exist, and even bicarbonates will cause the change. The critical pH is about 13. The point I was making was that once you have done the difficult step of converting MnO2 to the MnO4n- state, you don't want to lose this moiety back to black crud!

DA
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[*] posted on 13-8-2013 at 05:25


If the critical pH is 13, then it would indeed be very easy to overshoot and end up with more black crud than you bargained for. That makes gassing with CO2 a good idea, if disproportionation is your thing ;) [not yours personally of course].

I'll see if I can find that paper on the stability of standardised KMnO4 solutions.




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[*] posted on 16-8-2013 at 08:57


Quote: Originally posted by DerAlte  

WRT the MnOXY.doc I wrote way back, I'd make the following correction. Use of PbO2 plus nitric acid as oxidizers of MnO2 or of Mn ++ , if it works at all, is about as useless as using air. I retried it and maybe got a faint coloration, that's all. Yet all the literature seems to say it works. If you follow references back, eg from Wiki, you'll just get other references, going back to the dawn of chemistry!

Of course, I do not claim it cannot work - if conditions are very precise or extreme, such as during an exact phase of the moon.

As to whether "... that KMnO4 containing residual MnO2 is often considered worthless. Apparently solutions of it degrade too quickly to be of much use..", I vouchsafe no opinion.


@Blogger

Alzheimer's strikes again!

The reaction of MnO4- with Mn++ ions produces MnO2 by one of those nasty disproportionations that bite you in the butt with transitional metals, a well known (?) but easily forgotten fact.

So, if MnO4- ions are produced with Mn++ in excess, the black crud results. This explains my recent failure and my (long ago) prior conviction that the production of MnO4- does in fact result.

I would be very interested if someone would verify this reasoning by attempting the oxidation of 'pure' MnO2 by this method using PbO2 and nitric acid. I'd try it but I have no HNO3 and don't fancy making any at present due to ill health.

Regards,

DA
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[*] posted on 17-8-2013 at 09:42


DerAlte:

I think we might be getting our wires crossed here. I was specifically referring to the German experimenter's (slightly erroneous equation) deliberate disproportionation:


3 K2MnO4 + 2 CO2 + H2O -------> 2 KMnO4 + MnO2 x H2O + 2 K2CO3

... which is indeed a dispropotionation reaction of the type 3 [VI] === > 2 [VII] + 1 [IV]

and how local areas of low pH (when using strong acid, instead of CO2) might cause more of the permanganate to be destroyed than with CO2. I could be entirely wrong on the latter and the by-product K2CO3 may be the real reason for the choice of acid.

[Edited on 17-8-2013 by blogfast25]




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[*] posted on 17-8-2013 at 11:52


Blogger:

My fault! Up above I added - as an aside -

Quote:
WRT the MnOXY.doc I wrote way back, I'd make the following correction. Use of PbO2 plus nitric acid as oxidizers of MnO2 or of Mn ++ , if it works at all, is about as useless as using air. I retried it and maybe got a faint coloration, that's all. Yet all the literature seems to say it works. If you follow references back, eg from Wiki, you'll just get other references, going back to the dawn of chemistry!

Of course, I do not claim it cannot work - if conditions are very precise or extreme, such as during an exact phase of the moon.

and later
Quote:
Alzheimer's strikes again!

The reaction of MnO4- with Mn++ ions produces ... etc

_ a bit astray of the topic in question.

I have no quibble with what you say. I interjected a confusing ad lib and and that's why wires got crossed.

Regards, Der Alte
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[*] posted on 18-8-2013 at 09:12


Something I want to try shortly [big cough!] is to fuse 'MnOx≈2' with an excess of KNO3 and just a bit of KOH (to ensure alkalinity):

MnO2 + 2 KNO3 ===> K2MnO4 + 2 NO2

I have reason to believe that with a 50 %w stoichiometric excess (of KNO3) the conversion [IV] to [VI] may be near 100 %.

Then maybe leach with 1 M KOH and acidify to [VII] + [IV]? Maybe with acetic acid ≈ 2 M, slowly and with vigorous stirring?

Sooner or later anyone always catches the 'permanganate virus'! :D




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