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Author: Subject: Permanganates
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[*] posted on 17-5-2010 at 00:56


BaMnO4 can be done by heating Ba(NO3)2 with MnO2; this I did to 600-630 [Cels], still have it around ...
==> Then bubbling CO2 through the hot emulsion (in water) indeed gives permanganate ...
==> ... which has a slightly different color compared to KMnO4, more reddish ...

=================
Then I figured I could have much better yield by making NaMnO4 and ppt. Ba(MnO4)2 ... : NaMnO4 has it's fine high solubility in water ...
==> ... maybe I get onto it soon again ...

The ppt.-ing of Ba(MnO4)2 would serve as a cleaning-step as well ..., giving some 95 % purity in the first step ..., which would be useful ...; also less water would be involved via this route because of the high solubilities of both NaMnO4 and Ba(MnO4)2 ...

[Edited on 17-5-2010 by chief]
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[*] posted on 17-5-2010 at 15:26


The reaction should have produced water, causing the melt to bubble or foam vigorously.

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

Maybe the temperature of the melt was too low.
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[*] posted on 17-5-2010 at 17:49


I think the temperature was too low. Nitrate and chlorate tend to get cooking at higher temperatures. Because of its instability, chlorate probably has a steeper exponent vs. temp, being more prone to thermal runaway, but nitrate oxidations are often fairly well behaved.

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[*] posted on 17-5-2010 at 21:16


Yes. Next time I will mix and melt all of the hydroxides and nitrates, then add a little MnO2 and increase heat until something happens.

I'll get a stainless steel saucepan with a lid for that ... the cat wants its bowl back.
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[*] posted on 23-6-2010 at 20:21


Mn2O7 + Cl2O7 --> 2MnO3(+) + 2ClO4(-)
The permanganyl ion is a powerful very reactive oxidizer.
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[*] posted on 24-6-2010 at 06:54


Quote: Originally posted by Anders Hoveland  
Mn2O7 + Cl2O7 --> 2MnO3(+) + 2ClO4(-)
The permanganyl ion is a powerful very reactive oxidizer.


Funny that neither species has changed oxidation state then isn't it. Please stop posting your speculations, unless they've got hard evidence behind them (i.e. you've tried it or you have a paper reference for it - if so, attach the paper/pictures of your experiment). Otherwise they're not doing anyone any favors and you may even get some inexperienced person killed.
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[*] posted on 27-6-2010 at 13:18


"Mn2O7 can react further with sulfuric acid to give the remarkable cation MnO3+, which is isoelectronic with CrO3"
"Mn2O7 + 2 H2SO4 → 2 [MnO3]+[HSO4]− + H2O"
[wikipedia, Manganese heptoxide]

"Dichlorine hexoxide is a dark red fuming liquid at room temperature that crystallizes as an ionic compound, chloryl perchlorate, [ClO2]+[ClO4]−. Many other reactions involving Cl2O6 reflect its ionic structure, [ClO2]+[ClO4]−, including the following:
NO2F + Cl2O6 → NO2ClO4 + ClO2F "
[H. J. Emeleus, Alan George Sharpe (1963). Advances in Inorganic Chemistry and Radiochemistry. Academic Press. p. 65. ISBN 0120236052.]

Thus it should seem obvious that a perchlorate anion should be able to exist in solution with a permanganyl anion, at least in equilibrium. Adding one more thing, you might need just a tiny bit of HClO4 to catalyze the reaction between Mn2O7 and Cl2O7, since neither of these compounds ionize without an appropriate ionic solvent.
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[*] posted on 3-8-2010 at 20:52


Review of Manganates.
Many have tried to produce manganates without very satisfactory results. The difficulty of producing measurable amounts of manganate(vi) by the methods indicated by 99% of text books was the motivation which started this off. Too many text books regurgitate the same old facts (and myths). Some new facts (to me, at least) have emerged – or maybe some new myths! This is a review plus some extra stuff I have found out by experiment or literature search since my last posting:

SEE: http://www.sciencemadness.org/scipics/MnOXY.doc

For brevity I have not credited specific postings and/or references (in many cases all I have is brief notes in my notebook). If you have access, Kirk-Othmer, Mellor, Ullemann are all good sources of information.

It is long (28 pages), written in rambling Der Alte style, and contains theory, practice and experimental details, plus a large collection of useful data on Manganates from many sources. Mainly written for my own benefit as a compact collection of data on the subject, it may interest others too. It is too long to post - the format makes this too tedious.

Contents:
Introduction to the chemistry of Manganese, concentrating on its oxyanions.

Why the wet method I originally proposed has only a cat’s chance in hell of working (theoretical recantation).

How to make respectable reagents and MnO2 from battery crud.

Other non-fusion ways of directly producing permanganates.

Fusion methods – the key part of the review. – lots of good data.

The Japanese patent (US #3986941) – a critical examination

The results of a recent series of fusions that I have performed.

Industrial Practice: – if you look at nothing else, look at this. It’s fascinating and well worth it, and shows why amateurs have such problems.
Finally, some experiments on electrolysis and some caveats.

If you want to merely investigate the manganates in solution, by all means use Na salts but don’t expect to extract and purify the product. It takes the patience of Job. If you want to keep a manganate, produce the Barium salt from the Potassium or sodium manganate..

Overall conclusion: To produce measurable quantities of permanganate with acceptable yield, only one method is really worthwhile. Use the liquid melt process with potassium salts only. Forget all sodium based processes; forget all wet processes.

Regards, Der Alte
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[*] posted on 3-8-2010 at 22:59


This is an awesome read. Thanks a bunch for your work DerAlte!
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[*] posted on 13-8-2010 at 17:38


Thanks, Taoiseach , for your kind words.

In keeping with my philosophy of recycling crud and using OTC products, the whole process of making potassium permanganate can be done using the following starting ingredients:

a piece of lead pipe, a SS or Ni bowl or cup, some SS and mild steel strips, terracotta pots, carbon rods/slabs (ex-Zn/C cells, or welding rods), KCl (‘salt substitute’), sodium bisulphate (‘pH Down’, for pools), Clorox bleach (optional), some used alkaline cells, and a battery charger, PSU or current regulated power supply.

The only thing that may be difficult to obtain may be potassium nitrate which you may find as fertilizer – (most other nitrates can be converted to KNO3). Preparing nitrates is difficult otherwise. As a substitute, potassium chlorate, easily made from KCl, is suggested. Ingenuity and a lot of time and patience helps, too.

Note: Using the italicized methods is cheating, under the CRUD rules!

(1) 50% sulphuric acid from NaHSO4 by electrolytic methods; or use what you have, or get battery acid.

(2) H2SO4 to produce MnSO4 from battery crud;

(3) KOH by electrolysis of KCl in a divided cell, at the same time making KClO at a carbon anode; keep temp low, to avoid chlorate.

{For chlorate, use an undivided cell, allow temp. rise and circulation, following the recommendations of hundreds of posts here}.

Boil down KOH produced in the cathode cell by heating in oven to > ~ 160C – this should solidify on cooling to 2KOH.H2O eutectic, about 88% KOH. {The water may actually help the conversion to manganate}. Or buy KOH.

(4) Use the KClO you made (or Clorox) to make MnO2. Dry well at 100C.

(5) Produce K2MNO4 (and/or K3MnO4) by the liquid fusion method outlined. Use KNO3:MnO2:KOH in molar ratio 1:1:5. It may be possible to reduce the nitrate to 0.5 mol. and still get acceptable results. (I have not tried chlorate). Heat at 300 - 400C, or keep the melt liquid for at least 2-3 hours.

(6) Dissolve in minimum H2O and electrolyze at 2-5N KOH, leaving excess MnO2 in the solution, carefully following guidelines in the review posted previously.

(7) Or if you have an OTC source, just buy the permanganate! – added for the faint of heart, or those who just need it for organic oxidations.

- and so, anyone for ferrates, chromates, vanadates?

Regards, Der Alte
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[*] posted on 13-9-2010 at 10:49


I read in Patnaik's P. Handbook of inorganic chemicals that reaction between Mn(OH)2 and KOH yields potassium manganate that can be turned into potassium permaganate. Is this true?
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[*] posted on 13-9-2010 at 11:36


With enough oxygen, or with an oxidizer, yes. This is demonstrated somewhere earlier in the thread.

Tim




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[*] posted on 25-11-2010 at 05:29


I saw that copper nitrate exists in encylopedia of organic reagents. Maybe we could make easier permanganates form copper?
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[*] posted on 25-12-2010 at 05:29


Actually, it was copper permanganate, I wrote the wrong compound name.
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[*] posted on 29-12-2010 at 20:37


Quote: Originally posted by 12AX7  
I believe graphite forms an oxide (intercalated or something) when subject to anodic conditions. As I recall, it works with any sulfate, and probably nitrate as well. Seems to me the only way you can possibly get perchlorate with graphite is by cheating the reaction with high voltage and current density pushing past the erosion regardless.

Speaking of graphite oxide, a lot of times I've had the graphite sludge from my chlorate cell rise to the top due to adherent oxygen. Now, I would ordinarily attribute this to hypochlorite decomposing slowly, but that only works when the smell of chlorine is strong. Sometimes it happens to low-hypochlorite solutions. Graphite oxide as the erosion product, with a high oxidation potential (above chlorate, but below perchlorate, persulfate, etc.), would seem to make sense, and if it's decomposing in suspension, that would explain the adherent oxygen bubbles.

Anyway, applying to this thread, you have to determine if the oxidation of whatever mechanism operates -- direct oxidation of manganite, production of intermediate peroxide or superoxide, etc. -- if it's lower than graphite's erosion potential.

Tim


As the formation of CO<sub>2</sub> thermodynamically should occur at a much lower potential, I am virtually certain that the ability of graphite to survive at all at such potentials is purely kinetic.

Graphite oxide should not evolve O<sub>2</sub> -- the O<sub>2</sub> is from the electrolysis.

As for the fusion methods, I believe that hypomanganate will form if enough base is present. For manganate, I believe you would want a MOH/MnO<sub>2</sub> ratio of 2:1 molar. Any more base will cause the formation of hypomanganate. These are all based on the thermal stability of anhydrous hypomanganates and the relevant reduction potentials.

BTW, the 2-stage fusion probably is needed only when no oxidizer (other than air) is employed. The first fusion probably uses

4MnO2 + O2 + 12KOH -> 4K3MnO4 + 6H2O.

This explains the need for high temperatures, as it is probably largely driven by loss of water. The second fusion probably uses either

4K3MnO4 + O2 + 2H2O -> 4K2MnO4 + 4KOH or

MnO2 + 2K3MnO4 +O2 -> 3K2MnO4.

Probably the first of the two, as it is mentioned somewhere that the mixture must be kept wet and no mention is made of unreacted MnO2 being left after the first fusion.

I would use a divided cell for manganate -> permanganate.

[Edited on 30-12-2010 by tnphysics]

This is my first post BTW. I have not actually done any of these things (and am not a chemist), so no guarantees.

[Edited on 30-12-2010 by tnphysics]




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[*] posted on 13-1-2011 at 14:11


Just heated NaOH and MnO2 in a 3 to 1 mole ratio, a dark green solid was formed.
I assume the reaction was as follow:
12NaOH + 4MnO2 + O2 --> 4Na3MnO4 + 6H2O ?

Also, when I dissolved the dark green solid in water, the solution turned green, after dissolving more of the substance it turned black (or very dark green/blue) and small brown flakes appeared at the bottom of it (MnO2 may be?).
Can someone correct me if I'm wrong or explain what exactly happened when I dissolved it in water?
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[*] posted on 14-1-2011 at 12:33


Quote: Originally posted by Mixell  
Just heated NaOH and MnO2 in a 3 to 1 mole ratio, a dark green solid was formed.
I assume the reaction was as follow:
12NaOH + 4MnO2 + O2 --> 4Na3MnO4 + 6H2O ?

Also, when I dissolved the dark green solid in water, the solution turned green, after dissolving more of the substance it turned black (or very dark green/blue) and small brown flakes appeared at the bottom of it (MnO2 may be?).

Can someone correct me if I'm wrong or explain what exactly happened when I dissolved it in water?


http://tinyurl.com/4k546cd

I would go with unreacted MnO2.
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[*] posted on 14-2-2011 at 16:02


tutorial:make Potassium permanganate


there are 2 months I did this:

receppe for making K permanganate

i take a new duracel AA alcaline battery and i open it i take 1/5 of the impure manganese dioxide and i filter all the rest of manganese dioxide , i keep the KOH for later use i take 1/3 of the solution and i ad very tiny amont of bleech to make sodium hydroxide and potassium hypochloride, than i ad this solution to the 1/5 manganese dioxide and 2/3 KOH solution . 2 week latter i ave got a verry depp pink solution with somme insolube matter , so i filter it and then i lets it a -5to-20 (shed). 1mont in alf later , now i ave a very little amont of wather and a small amont of black violet crystal withch a letts dry and i tested with glyserine, and ive got fire!!

hop this guide will help someone,the yield are praticaly insinifien 'arround 0.7gram but for someone canot get essayli k permangane this is a good option.


thanks!!
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[*] posted on 16-2-2011 at 05:58


If you got potassium permanganate, you actually had a nice yield as it crystalized I think. The solution had to be quite concentrated then. Maybe this would work with sodium hydroxide too, just sodium permanganate would stay in the solution. It's worth to try, it looks like the solution has to be left in air few weeks for reaction to proceed. Maybe the resulting sodium permanganate solution could still be used for some oxidations.
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[*] posted on 17-2-2011 at 14:36


Now it doesn't let me edit, but I will try in the future mixing Sodium Hydroxide, bleach (strong, made by electrolysis of NaCl) and MnO2. After that I will leave it in the air, stir the mixture every day to dissolve some more oxygen and add 1mL of bleach sometime. Then I will try to take a photo of the solution almost every day to see how the reaction is proceeding. Maybe nobody tried this on few months length, but I believe bleach and dissolved oxygen will oxidize MnO2 a little bit each day, till I will have concentrated solution. If a weak solution of permanganate is possible in a month, then in a few months will be stronger. If this experiment will proceed succesfully, that means we can use this on a big scale, add like a few litres of bleach, 0.5kg NaOH and 0.5kg MnO2, leave it and add some fresh bleach every week. Maybe this should be done whole year to get a nice solution, but if we work on a larger scale, we will have much more permanganate than we can use.
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[*] posted on 14-3-2011 at 17:24


I have KMnO4 from 2 different lab suppliers. When I reduce one of them, it turns from purple to red to orange to yellow. The other takes the usual blue green yellow. Any idea why?



hey, if you are reading this, I can't U2U, but you are always welcome to send me an email!


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[*] posted on 22-3-2011 at 15:45


Quote: Originally posted by ScienceHideout  
I have KMnO4 from 2 different lab suppliers. When I reduce one of them, it turns from purple to red to orange to yellow. The other takes the usual blue green yellow. Any idea why?


vanadium impurity i think.




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[*] posted on 9-6-2011 at 07:49


Two facts that may add value to this thread.

First, please note the extract below relating to the increase in reaction rates in a particular oxidization using dichlorine monoxide over Cl2 and HOCl:

"Dichlorine monoxide exists in very low equilibrium concentrations in dilute HOCl solutions, nevertheless, it is a kinetically significant reactant. For example, although tetracyanonickelate(II) can be oxidized by chlorine in aqueous solution to trans-NiIII(CN)4(H2O)2-, the second-order rate constant at 25C for oxidation with Cl2O is 40 times greater than for Cl2 and 2.6X 10^7 greater than for HOCl (32)."

LINK to source:
http://www.scribd.com/doc/30121142/Dichlorine-Monoxide-Hypoc...

The second fact alludes to a previously cited patent involving a chlorination approach. This implies potential feasibility by substituting for Chlorine an oxide of Chlorine, although I have not read the patent or have the pertinent chemical equation. However, lets consider replacing KClO with Cl2O. Per the current reaction:

2MnO2 + 3XClO + X2CO3 --> 2XMnO4 + 3XCl + CO2 (g), (X = alkali metal)

and replacing XClO with Cl2O (or think of it as ClClO):

2MnO2 + 3 Cl2O + X2CO3 --> 2XMnO4 + 3Cl2 (g) + CO2 (g), (X = alkali metal)

with no salt separation issues (except perhaps for excess X2CO3, so choose carbonate salt accordingly) and no chlorate concerns (depending on Cl2O generation, see below).

SIMPLE LOW COST PREPARATION OF Cl20 (Caution, mildly explosive in 25% concentration range with sensitivity to light, heating, shock and organic compounds):

1. Prepare dilute HClO by combining Bleach (NaClO, for example) with a weak acid. These could include acetic, ascorbic, boric or even carbonic acid, but reaction time for the latter is much slower. Using Acetic (Ac) acid:

HAc + NaClO --> NaAc + HClO

2. Distill solution half way as most of the HClO vapor and the gaseous Cl2O, is brought over first.

3. Reheat HClO from Step 2 and pass vapors over a drying agent (best is Ca(NO3)2 but dehydrated NaCl may also work). Dichlorine Mono-oxide can also be dissolved in CCl4. Use the Cl2O vapors directly as storing the gas is not recommended for safety reasons.
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[*] posted on 2-8-2011 at 07:12


i tried to disproportionate Potassium permanganate with hxdrogen peroxide, but i got really poor results, with about 0.5 or a gram - 1 gram, i had to use 100mls of hydrogen peroxide and it still didnt fully react, the solution was left purple and i just waisted my H2O2
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[*] posted on 25-8-2011 at 19:50


i want manganese metal for some time now, is it possible to add HCl to the KMnO4 = KCl + MnCl2 + Cl2 and then with the KCl and MnCl2 solution, add aluminum to it to precipitate the manganese? is this a good method to use?



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