Sciencemadness Discussion Board - Permanganates

Permanganates

Pages: 1 2 3 4 5

chief - 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]

Paddywhacker - 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.

12AX7 - 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.

Tim

Paddywhacker - 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.

Anders Hoveland - 23-6-2010 at 20:21

Mn2O7 + Cl2O7 --> 2MnO3(+) + 2ClO4(-)
The permanganyl ion is a powerful very reactive oxidizer.

DJF90 - 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.

Anders Hoveland - 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.

DerAlte - 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

Taoiseach - 3-8-2010 at 22:59

This is an awesome read. Thanks a bunch for your work DerAlte!

DerAlte - 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

Random - 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?

12AX7 - 13-9-2010 at 11:36

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

Tim

Random - 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?

Random - 25-12-2010 at 05:29

Actually, it was copper permanganate, I wrote the wrong compound name.

tnphysics - 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 CO2 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 O2 -- the O2 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/MnO2 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]

Mixell - 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?

The WiZard is In - 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.

plante1999 - 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!!

Random - 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.

Random - 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.

ScienceHideout - 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?

plante1999 - 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.

AJKOER - 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.

Chemistry Alchemist - 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

Chemistry Alchemist - 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?

blogfast25 - 27-8-2011 at 04:32

Quote: Originally posted by Chemistry Alchemist

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?

Again, this doesn't really belong in this thread.

To answer your question:

the reduction reaction 3 Mn2+(aq) + 2 Al(s) === > 3 Mn(s) + 2 Al3+(aq) can indeed proceed. Unfortunately manganese is a very reactive metal and any Mn formed would immediately be oxidised completely by water, first to Mn(OH)2, which would then in turn be oxidised again (by air oxygen) to MnO2. No manganese metal can be obtained this way. That's in contrast to for instance copper, which can be obtained that way but copper is far more 'noble' than manganese...

The correct equation for the reaction you were trying to write is in fact:

KMnO4 + 8 HCl === > KCl + MnCl2 + 5/2 Cl2 + 4 H2O

Potassium Manganate disporporting

Chemistry Alchemist - 22-9-2011 at 22:52

I done the casual reaction of KMnO4 + Break fluid, this Produces small amounts of Potassium Manganate when the potassium Permanganate decomposes due to the heat, i washed the remains in water to clean the Manganese oxides, and left nice deep green solution of manganate, now, i wanted to dry it to a powder so i poored it into a evaporating dish and left it in the sun, an hour later i came back and the solution had turned clear with a brown precipitate, what has happened here? is the brown precipitate dissolved manganese salt being oxidized in the air to form the Insoluble oxide and left a solution of potassium ____? please explain if possible

blogfast25 - 23-9-2011 at 04:35

Quote: Originally posted by Chemistry Alchemist

I done the casual reaction of KMnO4 + Break fluid, this Produces small amounts of Potassium Manganate when the potassium Permanganate decomposes due to the heat, i washed the remains in water to clean the Manganese oxides, and left nice deep green solution of manganate, now, i wanted to dry it to a powder so i poored it into a evaporating dish and left it in the sun, an hour later i came back and the solution had turned clear with a brown precipitate, what has happened here? is the brown precipitate dissolved manganese salt being oxidized in the air to form the Insoluble oxide and left a solution of potassium ____? please explain if possible

Manganates (Mn [+VI] or MnO42-

are inherently unstable, cannot be isolated and on attempting to isolate, disproportionate: Mn (VI) === > Mn (IV) + Mn (II), so brown MnO2 precipitates...

[Edited on 23-9-2011 by blogfast25]

Chemistry Alchemist - 23-9-2011 at 04:39

Thanks

what would be the potassium salt?

blogfast25 - 23-9-2011 at 05:23

K2MnO4. But as a solid this is almost impossible to synthesize. In solution it's emerald green (depending concentration, obviously...)

[Edited on 23-9-2011 by blogfast25]

Chemistry Alchemist - 23-9-2011 at 05:25

but isnt K2MnO4 a green solution?

rstar - 28-9-2011 at 04:31

Quote: Originally posted by blogfast25

disproportionate: Mn (VI) === > Mn (IV) + Mn (II), so brown MnO2 precipitates...

[Edited on 23-9-2011 by blogfast25]

They disproportionate like:
Mn(VI) = Mn(IV) + Mn(VII)

3MnO42- + 2H2O = MnO2 + 2MnO4- + 4OH-

blogfast25 - 28-9-2011 at 08:45

Quote: Originally posted by rstar

They disproportionate like:
Mn(VI) = Mn(IV) + Mn(VII)

3MnO42- + 2H2O = MnO2 + 2MnO4- + 4OH-

I only once made some K2MnO4 solution and didn't see it form any purple KMnO4.

rstar - 29-9-2011 at 02:26

Quote: Originally posted by blogfast25

I only once made some K2MnO4 solution and didn't see it form any purple KMnO4.

Your solution might be quite basic, in which green K2MnO4 is stable, but as your solution will become acidic it will gradually change to purple KMnO4.

Try adding some acids

blogfast25 - 29-9-2011 at 05:02

Quote: Originally posted by rstar

Your solution might be quite basic, in which green K2MnO4 is stable, but as your solution will become acidic it will gradually change to purple KMnO4.

Try adding some acids

My solution was very alkaline: MnO2 + KOH + KClO3 fusion product, leached. But I haven't got any right now...

Waffles SS - 4-1-2012 at 00:53

This is possible to use air oxygen instead of potassium nitrate in first reaction?I want to use 50% potassium Hydroxide solution and MnO2 .

I decide to enter air oxygen in this boiling solution by air pump.Does it possible?

blogfast25 - 4-1-2012 at 09:27

Quote: Originally posted by Waffles SS

This is possible to use air oxygen instead of potassium nitrate in first reaction?I want to use 50% potassium Hydroxide solution and MnO2 .

I decide to enter air oxygen in this boiling solution by air pump.Does it possible?

I doubt that Mn (IV) can be oxidised to Mn (VI) in solution by air oxygen, In fact I'm pretty sure it isn't possible. It requires fusion at least, in which case vigourous aireation might do the trick.

Waffles SS - 4-1-2012 at 11:37

Thanks dear blogfast25,
What about pure oxygen or ozone(O3)?

blogfast25 - 5-1-2012 at 10:42

Ozone? Where are you gonna get that from?

Nah, good old fusing of MnO2/KOH with KNO3 or KClO3 does it nicely.

Waffles SS - 5-1-2012 at 12:56

Quote: Originally posted by blogfast25

Ozone? Where are you gonna get that from?

I have a ozone generator(its for swimming pool)

AndersHoveland - 5-1-2012 at 13:12

I have always found it interesting that ozone oxidizes chlorine, but not carbon tetrachloride.

DerAlte - 6-1-2012 at 08:12

@ Waffles SS

Quote:

Yes but not 50% solution. Commercial methods sparge air through MnO2/KOH mixtures but they use molten KOH at highish temperatures.

See the article I wrote and posted as http://www.sciencemadness.org/scipics/MnOXY.doc

Page 12ff.

Regards
Der Alte

Eddygp - 29-4-2012 at 13:03

Is it possible to synthesize in a relatively easy way calcium manganate? Or any other alkali earth, whose chemistry is probably the same...

blogfast25 - 30-4-2012 at 09:13

If manganate (MnO42-

- Mn [VI] - is what you mean then no. It's not very stable at all. Heating MnO2 with KOH and a strong oxidiser like chlorate or nitrate yields it but it decays quickly.

[Edited on 30-4-2012 by blogfast25]

Eddygp - 30-4-2012 at 12:15

Yes. Manganese, technetium and rhenium have always been quite interesting metals (at least to me). Oh, also chromium.

Random - 18-7-2012 at 11:07

Has anyone tried preparing calcium permanganate? Maybe it's easier to isolate, anyone knows its properties?

By the way, I have found that sodium permanganate is available as about 30% solution, so it should be definitelly possible to make it.

A very small amount is produced with the reaction of mno2 and bleach with sodium hydroxide. So I am thinking, if we would put MnO2 into hypochlorite cell and electrolyse it, maybe the continuous slow supply of Cl2 and hypochlorite in a basic solution would over time make a bigger amount of permanganate.

plante1999 - 18-7-2012 at 11:24

Quote: Originally posted by Random

Has anyone tried preparing calcium permanganate? Maybe it's easier to isolate, anyone knows its properties?

By the way, I have found that sodium permanganate is available as about 30% solution, so it should be definitelly possible to make it.

A very small amount is produced with the reaction of mno2 and bleach with sodium hydroxide. So I am thinking, if we would put MnO2 into hypochlorite cell and electrolyse it, maybe the continuous slow supply of Cl2 and hypochlorite in a basic solution would over time make a bigger amount of permanganate.

You need strong mixing of the MnO2 in the satured solution of sodium chloride, and it work, I already make some with the same process.

Random - 18-7-2012 at 11:28

Quote: Originally posted by plante1999

Quote: Originally posted by Random

You need strong mixing of the MnO2 in the satured solution of sodium chloride, and it work, I already make some with the same process.

Yeah, I thought stirring would be useful there. What concentration was it? Could it be useful solution for example for oxidation of alcohols?

plante1999 - 17-8-2012 at 10:57

I don't remember perfectly but yes it could be used to oxidize alcohols. Reaction with alcohol is fast and exothermic.

Arcuritech - 19-8-2012 at 19:57

I noticed that most (if not all) of this thread is about alkali metal permanganates. Does anyone have information on the permanganates of transition metals, poor metals, etc.? I imagine CuMnO4 would be an amazing oxidizer

vmelkon - 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.

DerAlte - 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.

Der Alte

vmelkon - 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.

DerAlte - 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!

Der Alte

DerAlte - 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]

Number 9 - 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.

learningChem - 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

blogfast25 - 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]

learningChem - 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...

blogfast25 - 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 (= G_postreaction – G_prereaction) 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):

MnO42- === > MnO4- + e- (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]

learningChem - 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)

blogfast25 - 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.

learningChem - 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...

blogfast25 - 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]

learningChem - 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...

blogfast25 - 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]

learningChem - 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

DerAlte - 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]

blogfast25 - 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]

DerAlte - 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 MnO4^n- state, you don't want to lose this moiety back to black crud!

DA

blogfast25 - 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.

DerAlte - 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

blogfast25 - 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]

DerAlte - 17-8-2013 at 11:52

Blogger:

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

Quote:

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

blogfast25 - 18-8-2013 at 09:12

Something I want to try shortly [big cough!] is to fuse 'MnO_x≈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'!

DerAlte - 18-8-2013 at 09:40

Quote: Originally posted by blogfast25

Something I want to try shortly [big cough!] is to fuse 'MnO_x≈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'!

Once a transition metal gets hold of you and tangles you in its swarm of outer electrons you are doomed!

Love to try that idea.

Der Alte

blogfast25 - 25-8-2013 at 10:38

I carried out two attempts to prepare potassium manganate (VI) with:

MnO2 (s) + 2 KNO3 (l) === > K2MnO4 (l) + 2 NO2 (g)

… and both failed. 5.0 g o MnO2 (pottery), 17.4 g of KNO3 (a 50 % excess) and 1 g of KOH (to ensure some alkalinity to protect any K2MnO4 formed) were mixed, ground and then fused together in a nickel crucible for 20 minutes on a medium-high Bunsen heat. In neither instances were significant amount of the emerald potassium manganate obtained, despite vigorous bubbling of the melt (what are these bubbles, O2?)

I then attempted:

MnO2 (s) + KOH (l) + KNO3 (l) === > K2MnO4 (l) + NO (g) + ½ H2O (g)

… in the same conditions and rather stupidly not realising at that point this equation isn’t balanced: ½ O is missing on the right hand side (*)!

Formulation: MnO2: 5.0 g; KOH: 4.1 g (a 20 % excess); KNO3: 9.5 g (a 50 % excess). This too melted easily and started bubbling right away, the latter which subsided completely after only about 5 minutes of heating. Heating was then stopped.

The crucible content was then leached with about 100 ml of approx. 1 M KOH solution. It was clear that much potassium manganate (VI) had formed with relatively little unreacted MnO2 in the leachate. The leachate was a comforting very deep green colour.

And then something unexpected happened: on hot filtering (normal filter paper) the manganate started to decompose to MnO2, ON THE FILTER! Finely formed MnO2 found its way through the filter, creating a mess in the filtrate. Even though I could still see green droplets passing into the filtrate, they too seemed to immediately revert to MnO2. But what is being oxidised here, since as no potassium permanganate is being formed? The filter itself?

A little bit of the unfiltered K2MnO4 bearing liquor was set aside. Simple attempts at acidifying this with dilute H2SO4 and acetic acid, to provoke the disproportionation to (IV) + (VII) were also unsuccessful.

(*) The correctly balanced equation is:

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

Or: MnO2 (s) + 1.333… KOH (l) + 0.666… KNO3 (l) === > K2MnO4 (l) + 0.666… NO (l) + 0.666… H2O (l), which isn't that far from the original unbalanced equation.

[Edited on 25-8-2013 by blogfast25]

Formatik - 25-8-2013 at 13:30

Quote: Originally posted by blogfast25

And the something unexpected happened: on hot filtering (normal filter paper) the manganate started to decompose to MnO2, ON THE FILTER! Finely formed MnO2 found its way through the filter, creating a mess in the filtrate. Even though I could still see green droplets passing into the filtrate, they too seemed to immediately revert to MnO2. But what is being oxidised here, since as no potassium permanganate is being formed? The filter itself?

One old preparation I read says distinctly not to use paper (Ausführliches Lehrbuch der pharmaceutischen Chemie, E.A. Schmidt, 848) to filter the potassium manganate solution, they used asbestos. Glass wool can be used instead. Pumice or fiberglass (which has the organic aspect destroyed) might also work for filtration.

blogfast25 - 25-8-2013 at 14:53

Thank you, Formatik, very useful. It makes sense, manganates being such fragile compounds. I'll also have a look at that used filter paper, for clues.

The German experimenter (Versuchschemie) only mentioned filtering, so I assumed he used paper. But one paper may not be equal to another...

[Edited on 25-8-2013 by blogfast25]

DerAlte - 25-8-2013 at 16:22

@Formatik & blogfast25 et al.:

Filter paper is a very definite NO for any manganate! I should have mentioned that in the MnOXY document. Cellolose is attacked even at high pH. I use filter paper as a cheap chromatograph to determine whether manganate is present during the electrolytic conversion to permanganate. Put a drop of the converting electrolyte on to filter paper and watch it spread out. Manganate and permangante separate, green outer, magenta inner. But in a few moments, only brown crud remains.

I use glass wool in the funnel. It traps MnO2 particles but may require a second pass, pouring the filtrate back over the funnel. The presence of MnO2 traps most of any remaining, but filtration may be slow.

You said: Formulation: MnO2: 5.0 g; KOH: 4.1 g (a 20 % excess); KNO3: 9.5 g (a 50 % excess). This too melted easily and started bubbling right away, the latter which subsided completely after only about 5 minutes of heating. Heating was then stopped.

5 mins. sounds an excessively short time to expect the reaction to complete, even with activated MnO2 instead of pottery grade. Was the temp. 'bright red heat' or merely enough to well melt the mix? Activated contains xH2O and can be expected to fizzle for a bit. Even if x is very small, steam should be given off as part of the MnO2/KOH reaction.

Old references talk about "hard" filter paper more resistant to oxidation, but I have never seen any.

Congratulations on trying out your idea, blogger. Not only do you blog fast, but you experiment fast!

Regards, DA

blogfast25 - 26-8-2013 at 05:08

DerAlte:

Good to see you still keep an eye on this thread, more about that later.

An additional observation: no NOx was ever seen, or smelled during fusion/reaction, so I need to assume the overall reaction was:

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

… as in fact described in ‘t’Jap patent’

. If so, I was probably not using adequate reagent quantities and will need to repeat at least once, with the right quantities.

For obvious reasons, I don’t know what the conversion was and there certainly was unreacted crud left but much K2MnO4 had formed too. Perhaps about 50 %? The difference with the ‘KNO3 only’ test was truly striking.

I accept that 5 minutes is a very short time but what’s a boy to do when after that time ALL steam evolution stops other than to assume it’s ‘game over’? Initially gas evolution was surprisingly vigorous and I had to stir to prevent over boiling. This wasn't water that was already there, no, this was being formed in situ, no question about it. Of course it's possible that more steam would have slowly and invisibly bled off over the course of a longer heat.

Can you confirm/infirm that I read correctly/incorrectly in your document that for manganate (VI) formation LOW temperature is to be preferred?

DA, I’ve finally managed till the wee small hours last night to trawl through this entire thread and your excellent summary of prior art. Boy, permanganates are still a lot harder to prepare than I thought (and I never had any illusions about it being easy to begin with). Hats off to you and Xenoid, especially for his work on electrolytic oxidation of MnO2.

It appears to me that in terms of Free Energy (G), the series Mn II/III/IV/V/VI/VII is a valley with the black crud at the lowest point and everything else wanting to roll back to that point, given half a chance!

I will continue some limited experimentation but need to prepare some ‘activated’ MnO2 first. I’ll use the MnSO4/NaClO method for that.

I’ll also want to make some Na manganate (V), like Xenoid did, just to be able to say: ‘I saw that!’

Thanks for the tip on filtration media. I should have known better. The German guy actually mentioned ‘harder’ filter paper. I’ll try and use (very clean!

) glass frit.

Oh, and one last thing. The little amount of green solution I managed to syphon off without ever seeing the filter paper, when gently acidified, once with HOAc, once with 1 M H2SO4, never yielded permanganate, not even a whiff, only brownish crud and some bubbles. Oxidation of water?

[Edited on 26-8-2013 by blogfast25]

DerAlte - 27-8-2013 at 20:32

@blogfast

Quote:

… as in fact described in ‘t’Jap patent’

I now believe most of what that says. I am always very sceptical about patents.
Exceptions are (1) yields – have you honestly got yields in any experimental process much greater than70-80%? (2) I have not yet successfully managed the electrolytic conversion of an alkaline suspension of ‘MnO2’ to MnO4- but do think it ought to be feasible, given a correct set of conditions and catalyst.

Quote:

Can you confirm/infirm that I read correctly/incorrectly in your document that for manganate (VI) formation LOW temperature is to be preferred?

{Aside: I am not sure your use of the word 'infirm' is the reverse of confirm! Infirm describes my current condition, unfortunately}.

Yes. Higher temperature favors hypomanganate with KOH. From the original sourece: (but this is with direct oxidation using air):

The ﬁrst step in roasting processes is the formation of K3MnO4 from MnO2 ore. This is promoted by high temperature and high KOH and low H2O concentration. The second step oxidizes Mn(V) to Mn(VI). A lower temperature and control of moisture in the air is used.

Note also, from the same source:

Aqueous potassium permanganate solutions are not perfectly thermodynamically stable at 25C, because MnO2, not MnO4-, is the thermodynamically stable form of manganese in water. Thus permanganate tends to oxidize water with the evolution of oxygen and the deposition of manganese dioxide, which acts to further catalyze the reaction.

Which is why black crud is often the result…

Quote:

It appears to me that in terms of Free Energy (G), the series Mn II/III/IV/V/VI/VII is a valley with the black crud at the lowest point and everything else wanting to roll back to that point, given half a chance!

Absolutely. Any experiment to produce Mn X seems to start with xMnO2.yMn2O3.zH2O and end up with the same, with only x,y,z changed!

One way to convert manganate to permanganate without disproportionation and the black plague is to use hypochlorite;
2MnO4-- + ClO- + 2H+  Cl- + H2O + 2MnO4-
Compare anodic oxidation:
MnO4--  MnO4- + e-

I’ve got a lot more but enough for now,
Regards,
Der Alte

blogfast25 - 28-8-2013 at 07:47

In the recesses of my chemical cavern I found some 'MnO2' I bought years ago, as a 'high grade' (definitely not 'pottery'). It's very brown (and very fine), rather than black and I believe it's a synthetic grade (it dissolves completely in HCl and appears also Fe free). So next week end I'll try some longer, relatively 'cool' fusions with KOH + KNO3 quantities in accordance with some experimenters here, using that material. Target: K2MnO4. Any specific recommendations welcome.

Thanks DA.

[Edited on 28-8-2013 by blogfast25]

DerAlte - 28-8-2013 at 13:33

The brown stuff is what one usually gets when one prepares "MnO2" chemically. It contains xH2O, x order <2. Expect more fizzle! Also, try heating a small quantity carefully to reduce water content. If it turns black, or darker, then x is fairly large.

Electrolytic MnO2 is quoted as 98% MnO2. I have never tried making it, so I don't know what color it is.

Der Alte

blogfast25 - 28-8-2013 at 16:55

I'll have look at heating a bit of it tomorrow. A bit of water doesn't do this thing any harm though, does it?

Any recommended formulation for best chances with K2MnO4?

Ta.

[Edited on 29-8-2013 by blogfast25]

DerAlte - 28-8-2013 at 20:00

Blogfast wrote:

Quote:

A bit of water doesn't do this thing any harm though, does it?

A source I have says

Quote:

liquid-phase oxidation: KOH, (O2)air, H2O , MnO2 ore -- > K2MnO4; electrolysis -- > KMnO4;

Of course this is with air oxidation but many industrial processes use very concentrated KOH solution :

Quote:

The USSR process (118–120) is discontinuous, uses turbine-agitated, low pressure reactors having a volume of 4 m3 each, and processes 2000–2500 L/ batch. Preconcentrated molten potassium hydroxide (70–80%) is added to the reactor with a quantity of 78–80% MnO2ore (<0.1mm particle size) resulting in a 1:5 molar ratio of MnO2: KOH. Air, or O2, is introduced below the liquid level by a sparging device at such a rate that a positive pressure of 186 – 216 kPa (1.9 – 2.2 atm) is maintained. The temperature is kept at 250 – 320C for the duration of the reaction, which requires approximately 4–6 h for completion. The reaction mixture, which reportedly remains ﬂuid during the entire time, is then emptied through a siphon.

From which I would guess that the presence of water is not deleterious – it is produced in the reaction anyrate. I have used highly concentrated NaOH and added nitrate and MnO2xH2O and heated, driving off water and the reaction seems to go as if you merely melted the hydroxide.

Der Alte

The Volatile Chemist - 23-3-2014 at 16:48

As for buying permanganate in the Eastern US: http://www.lowes.com/pd_112505-677-PF65N_0__?Ntt=potassium&a... Lowe's sells cheap 5lbs containers of it...

Brain&Force - 24-3-2014 at 09:15

I wonder where you would find it on the west coast - it's not available where I live. I have to resort to aquarium suppliers, which supply it only as a solution.

vmelkon - 1-4-2014 at 13:55

Quote: Originally posted by The Volatile Chemist

As for buying permanganate in the Eastern US: http://www.lowes.com/pd_112505-677-PF65N_0__?Ntt=potassium&a... Lowe's sells cheap 5lbs containers of it...

Lucky you.
I guess 5 lbs is something like 2.5 kg.
I had to buy mine from England and it was 70$ for 800 g, shipping included to Canada.

//EDIT: woops, the actual amount is 46.35 $USD
26.50$ (800 g of KMnO4) + 19.84$ shipping.

[Edited on 2-4-2014 by vmelkon]

plante1999 - 1-4-2014 at 14:06

You gotta love shipping to Canada.

DoctorZET - 16-4-2014 at 11:47

Bubbling with CO2 a boiling solution of potassium manganate, should work :

3 K2MnO4 (aq) + 2 CO2 --(boiling > 100*C)--> 2 KMnO4 (aq) + 2 K2CO3 (aq) + MnO2(s)

The reason is that H2CO3 (carbonic acid) is a stronger acid than [H2MnO4] (the theoretic formula of manganic acid) is .
We know this because, K2MnO4 has stronger basic properties than K2CO3...

So, I think that the reaction mechanism is:

2K⁺ + MnO4^2- + [H2CO3](aq) --(activation energy > 100*C)--> 2K⁺ + CO3^2- + [H2MnO4](aq)

...("H2CO3" and "H2MnO4" are placed in brackets, because this are unstable compounds and exist only in aquous solutions with H3O⁺ anions and the respective cations associated)

[H2MnO4] + 2 H2O <--> 2H3O⁺ + MnO4^2- <--> 3 H2O + [MnO3]

"MnO3" is a very unstable oxide of manganese .

[MnO3] <--> MnO2 (solid) + [O]^2-

The atomic oxygen [O] makes a nucleophilic attack on the an other manganate ion , a double oxigen-manganese bond is formed, but, a single O-Mn bond is breaked ...

[O]^2- + MnO4^2- + 2H3O⁺ --> (MnO)O4^2- + 2H3O⁺ --> MnO4^- + 3H2O

As soon the permanganate ion is formed, a potassium ion get closer ( there was 2 K⁺ before the oxidation of the manganate ion to the permanganate ion):

K⁺ + MnO4^- = KMnO4 (aq)

I'm now thinking if I also can use calcium manganate instead of the potassium one .

Random - 8-4-2015 at 15:09

How would you form calcium manganate?

Kagutsuchi - 22-6-2015 at 10:06

Does anyone know anything about aluminum permanganate? I'd try to prepare some but I'm unsure if it will worth it or it will be completely useless or even decompose instantly.
I'd try the
2KMnO4+dilute H2SO4----->2HMnO4+K2So4
3HMnO4+Al----->Al(MnO4)3+1.5H2 method.

P.S.:I'd be happy to hear if you tried it in pyrotechnics.

[Edited on 22-6-2015 by Kagutsuchi]

softbeard - 25-11-2015 at 09:40

Quote: Originally posted by Kagutsuchi

2KMnO4+dilute H2SO4----->2HMnO4+K2So4
3HMnO4+Al----->Al(MnO4)3+1.5H2 method.

Hey Kagutsuchi, I don't know the specifics of a theoretical 'aluminum permanaganate' but I'm sure the 2nd equation you've written is nonsense. There is no way the MnO₄^- anion is going to survive with a reductant like Al metal around. Much less yielding H₂ gas in the process.
You'll get instant reduction of the MnO₄^- to Mn₃O₄ or Mn⁺⁺ by Al metal, depending on mainly the pH.
Furthermore, I really doubt Al(MnO₄)₃ exists as a compound. I would venture that even if it does exist, it would be impossible to prepare in an aqueous medium.

MolecularWorld - 25-11-2015 at 10:33

I can only see it in snippets, but the Kirk-Othmer Encyclopedia of Chemical Technology seems to suggest aluminum permanganate exists, can be crystallized, is unstable above 80*C, and can be formed from the reaction of cold solutions of potassium permanganate and aluminum sulfate: aluminum permanganate stays in solution, potassium aluminum sulfate precipitates (I couldn't actually see the equation).

gatosgr - 9-3-2016 at 02:41

Have you found an easy way to turn decomposed KMnO4 from exposure to air back to KMnO4?

[Edited on 9-3-2016 by gatosgr]

clearly_not_atara - 18-3-2016 at 15:47

Quote:

Have you found an easy way to turn decomposed KMnO4 from exposure to air back to KMnO4?

From the first page:

Quote:

The industrial method of heating MnO2 with KOH and using air as an oxidizer is doomed to failure.

This reaction instead almost always produces potassium manganate, K2MnO4. Unfortunately Mn(VI) is more stable than Mn(VII), despite being far less interesting. Stupid quantum mechanics demons. On the other hand:

https://en.wikipedia.org/wiki/Potassium_ferrate

Quote:

Edmond Frémy (1814 – 1894) later discovered that fusion of potassium hydroxide and iron(III) oxide in air produced a compound that was soluble in water.

In the case of iron, Fe(VI) is more stable than Fe(V) which is more stable than Fe(IV), although none of these is particularly stable. As a result only potassium ferrate is produced when iron oxide is heated with KOH. But luckily, Fe(VI) is a stronger oxidizing agent than Mn(VII), which means that this should happen:

2FeO4(2-) + 6MnO4(2-) + 5K+ + 5 H2O >> Fe2O3 + 6MnO4- + 5 KOH + 5 OH-

Excess ferrate can be removed, presumably, by simply allowing it to decompose.

urenthesage - 15-5-2016 at 11:06

Quote: Originally posted by MadHatter

Given the recent CPSC(assholes) victory against FireFox, I'm sure that any convenient
method will be appreciated. KMnO4 sales by pyro suppliers are now limited to 1 LB a
year. There's still some OTC sources for now but we don't know how long that'll last.

So dont go to a pyro supplier. I got mine at a water depot, 5 lbs for about 45$. Farmers use it to take iron out of well water for cattle. They didnt even give me a second glance when I bought it.

KMnO4 from manganese metal electrodes: A paper and my results.

Romain - 1-9-2016 at 07:50

While looking for a method to prepare permanganates, I came across a paper from 1921 titled "The Electrolytic Production of Sodium and Potassium Permanganates from Ferromanganese", by Wilson et al.. I could get the full-text access with my uni library account and found very interesting information.

I'll just give the main points of the paper and technical details that are relevant for those who want to try it. And then I'll outline my two runs with a similar setup and the results I got.

The paper

On the electrolytic cell the authors say:
"The electrolyses were carried
out in a cell which consisted of a cylindrical glass jar (12.5
cm. in diameter by 15 cm. high), containing within it a
porous porcelain cup (5 em. in diameter by 12.5 em. high)
which served as a diaphragm. In this cell were placed the
two electrodes, the cathode (of 16-gage sheet iron, 3 em. by
15 em.) being within the cup, and the ferromanganese anode
standing in the jar outside, and about 3 em. away from
the cup. The freshly cast anodes had the dimensions 2.5
em. square by 15 em. in height. Electrical contact with the
anode was made by clamping a strip of brass, carrying a
binding post, against a freshly ground surface near the top."

On the anolyte and catholyte they used:
"It was found that the carbonate electrolyte (as compared
with the hydroxide) gives the purest product-uncon-
taminated with manganate-at the best efficiencies and for
the least expenditure of power. Hence, this electrolyte is
recommended for technical operation."

"The anolyte was 12 per cent sodium carbonate solution
(previously found to give about the optimum results), and
the catholyte was 8 per cent sodium hydroxide solution."

On anode composition:
"An experiment with spiegel iron (about 40
per cent Mn) showed practically zero yield of per-
manganate."

They state the cathode was simply iron and the current density was around 5 to 10 A/sq. dm.

In total they did 180 runs with varying conditons so this paper is just pure gold. If some of you want more of it I can paste the whole of it here (7 pages) or upload it somewhere (where?)

My two test-runs

My setup consists of a 600 ml beaker with a terracotta flower pot (acts as a porous membrane) in it. The hole of the flower pot was plugged with rubber stopper so that the flower pot can hold the anolyte without leaks
The beaker contains the catholyte. The levels of the electrolytes were adjusted to be at the same height to prevent the electrolytes from mixing.

The anode was a pile of 99.9% manganese metal flakes (ordered on OnyxMet.com. Btw they sell Mn flakes for 16 bucks/kg and ferromanganese 86% for 8 bucks/kg) that I put in the terracotta pot. Electrical contact was made with a long manganese flake touching the pile of Mn and connected to a power supply (+). I didn't use ferromanganese because I didn't have it, but it appears to work fine with pure Mn.

The cathode was a stainless steel plate.

1) The first run was done with 1M KOH in both cathode and anode compartment. A voltage of about 3V with a current of 100 to 200 mA was applied to the cell for a few hours. The catholyte stayed crystal clear with bubbles of gas being evolved at the cathode (oxygen I guess). The anolyte turned a purple color at first but within a few minutes turned instead a deep green/brownish color which indicates that potassium manganate was produced. I didn't process this batch and insted changed the composition of the anolyte:

2) The second run was done in the same conditions except:
- The anolyte used was about 10-15% Na2CO3. Catholyte was still KOH soln. because I didn't want to throw it away and thought it should have no influence on the run.
- The current supplied was about 500 mA and the voltage 5V, because I wanted to produce NaMnO4 in high enough conc. to be able to precipitate KMnO4 upon adding KCl and I don't have time to wait a week for the run to "complete". The temperature of the anolyte was about 25°C for the whole run (room temp 16°C). Total run time was about 50 hours.
The anolyte turned very dark purple and a fine mist of permanganate is produced so I used plastic wrap to contain it.

Results
Concentration: I titrated a sample of the anolyte with Mohr's salt (a stable source of Fe2+ ions) solution by dripping the anolyte into acidified mohr's salt solution until the purple colors stays. The anolyte was found to be about 0.4M in conentration with an error of say 5 or 10 % because I don't have very precise titration instruments (such as a burette).

Extraction of the permanganate
I filtered the anolyte on a GLASS filter funnel (no paper) and added conc. KCl solution to the filtrate. the undissolved solids appear to be MnO2 but I'll have to check that. The filtrate is currently evaporating at about 80°C on a hotplate with a fan blowing over it. Dark needle-like crystals appear to be forming upon cooling. I'll report back when I have separated the crystals.

I plan to do a more carefully controlled test on a 1 liter anolyte scale but I cannot go ahead right now because I don't have enough manganese left (the pieces are too small to work with) so I'll have to order some more Mn. I might try to cast a manganese metal plate.

First picture: The cell. Second picture: Mn flakes I used piled up in the flower pot.

Update: The crystals that formed from the evaporated anolyte are highly impure as evidenced by the presence of white crystals and brown crud among the needles of KMnO4. The crystals were dissolved in a small amount of water and the solution was heated up to boiling. It was filtered while hot on a preheated filter funnel (to remove MnO2, again). The filtrate yielded 0.37 g of small (and pure) needles of KMnO4 upon cooling.

(More KMnO4 could be extracted from the anolyte, but this run was just to prove that the method works and I didn't try to be particularly efficient in the extraction.)

The product was tested by adding a few drops of conc. H2SO4 to a crystal to form Mn2O7 and the reaction with ethanol was as expected: it ignited!

[Edited on 1-9-2016 by Romain]

[Edited on 1-9-2016 by Romain]

[Edited on 1-9-2016 by Romain]

NEMO-Chemistry - 6-12-2017 at 19:02

I havnt read all of this post, only the first 6 pages. I did however find a webpage with an experiment to make it.
The method mentioned is as follows.

"Making Potassium Permanganate The synthesis of this compound is comprised of a few steps that demonstrate "redox" or reduction-oxidation reactions. You need 7 grams of potassium nitrate, 1 gram of manganese dioxide, 2 grams of potassium hydroxide and a few milliliters of sodium bicarbonate. Protective eye wear, a small glass vial, a 50 mililiter beaker, a small hammer, mortar and pestle and a ventilation hood are recommended. Start the experiment outdoors or under a ventilated fume hood. Mix 7 grams of potassium nitrate and 1 gram of manganese dioxide in the vial. Using a torch, heat the vial gradually until the two chemicals melt together. Keep heat on the molten mixture for several minutes. Add 2 grams of potassium hydroxide to the mixture and promptly re-heat the vial until a green boiling substance appears. Continue boiling mixture for 5 to 7 minutes. Take the torch off the boil and let the vial cool down. After the mixture is a green solid, use the hammer to smash the substance into smaller pieces. Use the mortar and pestle to grind the pieces to a powder. Pour the powder into the beaker and dissolve in 50 mililiters of distilled water. After the solution turns green, pour off the mixture that has risen to to top. Add the sodium bicarbonate in small increments while stirring steadily until the solution takes on a purple color. Adding too much sodium bicarbonate will result in a light pink color that signifies a destruction of permanganate."

I noticed some of this mentioned in the first few pages, so its likely the website is wrong. The site is here
https://sciencing.com/potassium-permanganate-experiments-123...

I am intrigued enough however to have a go at this, so i have ordered the reagents needed. Even if its a bust i think its worth trying the method as stated just to see what happens. The description of the colour makes me think the yield is appalling.

I know its an old thread, but will post the results in next couple of weeks.
Which reminds me, i got an early Christmas present, a tablet! not super duper by a long way, but it take pics easy enough for me to use in the lab, i also downloaded a chemistry lab book app

, it works by speech and typing as well as being able to add pics.

Should be a handy little device for the lab, i have downloaded loads of free chem apps like periodic table etc. Will be nice to have a small device with all the info i need on it.

[Edited on 7-12-2017 by NEMO-Chemistry]

Pages: 1 2 3 4 5