Sciencemadness Discussion Board - Permanganates

Permanganates

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ciscosdad - 8-10-2007 at 17:16

@Xenoid & DerAlte

Do you guys see any prospect of refinement of these methods? Do you think it feasible to do an air oxidation to Manganate followed by the electrolysis? At least there is some prospect of reusing/recharging the reaction mix after electrolysis if there is no Chloride or nitrite in there.
Of course accumulation of carbonate will probably limit the number of follow up steps.
The method may be marginally useful if we could make 30grams or so per cycle for half a dozen cycles.
The steps may be:
Air oxidation with KOH/MnO2 (maybe partial NaOH)
Electrolysis and separation by crystallization and filtration.
Evaporate with extra reagents
Air oxidize.
Etc

Your opinions?

Xenoid - 8-10-2007 at 17:39

Not sure about the air oxidation, I haven't tried it.

It seems to me that the electrolysis is something that can easily be done on an industrial scale, but is a lot of messing around for the home amateur. In the last electrolysis I carried out (above), I failed to take into account the extra frothing caused by the increased current. The lid I had made did not have an "o" ring so it did not seal perfectly and hot alkaline electrolyte ran down the outside of the cell and etched the polished surface of my hotplate.... :mad:

as well as putting brown stains all over it. Manganese compounds are such a pain, they seem to change oxidation state just by being looked at...

As I said before, the electrolysis is a lot of messing around for little return!

Possibly some refinements to the cell design could be made to make the operation more efficient but I don't know what!

Regards, Xenoid

ciscosdad - 8-10-2007 at 19:12

OK.

First refinement: The reactor is only 1/3 or 1/4 full. Possibly a steel pipe with a welded bottom. Use it as the anode? If so, it may well have to be Stainless.
We want 50 g or so of product per cycle, so 2 litres or so of solution. That makes a 5 - 6 litre pot. Say 150mm dia and 300 to 400 high. Requires a lid with vent pipe and provision for thermometer, stirrer and Cathode rod.

20A to 50A to produce 50g in a day or so. A higher current would be better so it could be done in 12 hrs or so and be supervised to avoid the thermal runaway Xenoid mentioned.
If the pot was well insulated, could the current maintain the required temp?
1 moles of MnO2 plus enough (~ 5 moles) KOH/NaOH mix to make ~20% alkali solution after oxidation.
Replace KOH and MnO2 to replace the amount extracted as KMnO4.
After a few cycles could some of the accumulating carbonates be precipitated with lime? Not all; we need to avoid calcium accumulation I would think.

Have I missed anything?

yangmeiqi0622 - 8-10-2007 at 23:14

Dear teacers:
I am so interested in manganate and have a lot of questions about it.
I had read the discussion about Mn(V), Mn(VI) and Mn(VII) above , but still not understand how to make Mn(V) stabilization in water liquor no matter Na3MnO4 or K3MnO4.
And KMnO4 decomposes into many different compounds after heating , such as K2MnO4,K3MnO4, MnO2 ……,but if I want to get K3MnO4 as main product, what's the proper temperature and heating time?
Thanks a lot

Xenoid - 9-10-2007 at 01:13

Quote:

Originally posted by yangmeiqi0622

.... but still not understand how to make Mn(V) stabilization in water liquor no matter Na3MnO4 or K3MnO4.

As far as I know Mn(V) (hypo-manganate) is not stable in solution, when I dissolve the solid blue Mn(V) from the fusion reaction it immediately hydrolises to produce a green Mn(VI) (manganate) solution and brown MnO2. I think the reaction is mentioned earlier in the thread.

2Na3MnO4 + 2H2O ---> Na2MnO4 + MnO2 + 4NaOH

Regards, Xenoid

DerAlte - 29-10-2007 at 11:10

MANGANATES - SYNOPSIS AND A FEW NEW RESULTS/COMMENTS: PART 1

Two different approaches have been considered in the above thread: direct chemical oxidation of Mn(iv) (usually as MnO2) in water solutions, and chemical/anodic oxidation by fusion/electrolysis.

The latter method has been proven practical by the splendid efforts of Xenoid, q.v.s., even if the yield appears low. I am sure it could be improved.

The usual methods of fusion to produce manganate by air or assisted oxidation have resulted in dismal failure in amateur hands. Very little Mn(vi) is ever produced, although this has been the industrial method stated in the literature of the last 100+ years.

Permanganate is never produced directly by fusion techniques, AFAIK.

This situation was resolved when Ciscosdad unearthed the Japanese patent, Attachment: US3986941A1.pdf (918.53 KiB) (courtesy not_important ). Therein methods using fusion and electrolysis were covered in good detail. This is an essential reference for those interested. Xenoid, acting upon this data, succeeded in obtaining KMnO4 crystals.

The key was to use a relatively low temperature fusion (250 - 350C) using alkali nitrate as oxidant and producing the hypomanaganate (Mn(v)O4)3- . I was very skeptical having with considerable difficulty produced this ion in solution (in about 5N NaOH), but it hydrolyses with water with great ease at a rather high pH, to the manganate (Mn(v)O4)2-. I have since tried this fusion, following the Japanese patent like Xenoid, and am totally convinced now!

See Xenoid’s posts above for his elegant methods.

Also, in the patent, it is possible to directly oxidize a MnO2 slurry electrolytically, in alkali hydroxide solution to KMnO4, under certain conditions. I had always assumed that this anodic oxidation could only be done with a massive MnO2 anode, and have, in the past, tried this crudely and gotten at least a coloration. It is surprising that it works with a suspension.

In view of this I wonder why fusion is even worthwhile. You may save a little electricity, but this is the least of problems. You use up oxidant unnecessarily and produce nitrite which has (assumedly) to he anodically oxidized to nitrate because otherwise KMnO4 will do it for you! In addition, since the hypomangante is hydrolyzed, (apparently, even in the very high pH hydroxide solution,

2MnO43- + 4H+ --> MnO2 + (MnO4)2- + 2H2O,

one half of the oxidation product is wasted in re-conversion to MnO2.

It complicates a process which otherwise would otherwise only contain KMnO4 and KOH (recyclable) at the end of electrolysis. In fact, KMnO4 ought to be precipitated because its solubility in KOH must be much lower than in water due to common ion effect. The only problem is the availability of cheap KOH. NaOH can be substituted and KCl or carbonate used to precipitate out the permanganate, NaMnO4 being extremely soluble.

Also noted in the Japanese patent was the use of a small amount of ‘catalytic’ KMnO4 to improve the conversion at start-up. Thus reminds one of the use of dichromate or other oxidant to help in the chlorate electrolytic cells. Has anyone any ideas as to why these ‘catalysts’ help? Or is it just folklore?

Next posting I will consider non-electrolytic methods for those who are allergic to electrolysis.

Regards, Der Alte

ciscosdad - 29-10-2007 at 14:42

Nice Summary DerAlte.
I am still of the opinion that the direct MnO2 to KMnO4 route has the most potential for amateur use. All we need to do is figure out what they are not telling us so the first embodiment process will work.
I am considering my own attempt in due course, and intend to try the upper level of the current density range and a relatively high starting KMnO4 content (if I can find any!). Xemoid seemed to be getting some premanganate but never enough to actually get any crystallization.

Xenoid - 29-10-2007 at 15:01

Quote:

Originally posted by ciscosdad
Nice Summary DerAlte.
I am still of the opinion that the direct MnO2 to KMnO4 route has the most potential for amateur use. All we need to do is figure out what they are not telling us so the first embodiment process will work.

Yes, I would agree. The direct route never seemed to "get of the ground" for me, but it would be a good route for amateurs. Plenty of experimentation still possible with this process, although I don't feel like it at the moment! I don't think I tried this method with NaOH because I was using 20% KOH drain cleaner solution at the time. It may have had some other chemicals in it which inhibited the reaction. Trying NaOH, and at higher concentrations may help. A nickel or monel metal anode might be worth trying also, I only tried SS.

Regards, Xenoid

[Edited on 30-10-2007 by Xenoid]

ciscosdad - 30-10-2007 at 13:54

@Xenoid
How likely is it that the drain cleaner you used had additives of some kind? I noticed that Hydroponics suppliers have 40% KOH as a pH increaser. I assume that is unlikely to have crap in it.

Xenoid - 30-10-2007 at 14:56

@ Ciscosdad
I have no idea what additives it may have had. It was TERGO brand, it may have been superceded because it is not listed on their web site, and the hardware store doesn't seem to stock it anymore. It "looked" clean and clear and the KOH content was listed as 192g/L which was slightly less than the 20% used in the first embodiment procedure. If you can get stronger KOH solution or better still, solid KOH it might be worth trying this procedure in a more alkaline environment.

Regards, Xenoid

DerAlte - 31-10-2007 at 08:53

Re: purity of alkali hydroxides.

Good points, Cisco & Xenoid. In a recent set of experiments I hope to report on, I used NaOH that was about 15 years old. Originally small half-spherical granules, it was now very moist and had encrustations. Without thinking straight (seem to do that too often) I weighed it as was and assumed it was 100%. On testing with titration, a found it was only about 70% NaOH. The rest was nearly all carbonate and I guess water.

It's not a strightforward titration against acid because of this. You need a high pH indicator because of the carbonate, to estimate the hydroxide. I didn't have one except for phenolphthalein. Bst way is to precipitate the carbonate with barium chloride and estimate the remining hydroxide.

In the US I think they use NaOH for drain stuff.

Regards,

Der Alte

Phosphor-ing - 31-10-2007 at 10:29

You can find this product in ACE hardware:

http://www.rootocorp.com/rooto/household_drain_opener.html

I beleive it is at least technical grade NaOH.

12AX7 - 31-10-2007 at 14:05

Hmm, I don't think we have that around here. Related products in that brand, but not that one.

DerAlte - 6-11-2007 at 10:58

MANGANATES – Part 2

Jottings from my notebook.

RAW MATERIALS (Emphasis on CRUD).

(1)Manganese sources:
MnCl2 (make or recover from Cl2 production); MnCO3; MnSO4 (fertilizer); MnO2, pyrolusite, (pottery / glass stores) or crud from spent Zn/C or (better) alkaline batteries - technical is around 77% typically. IN fact, just about anything containing manganese.

Convert salts to MnO2 using hypochlorite. I am a bit of a purist and don’t use the crud from batteries without purification, but if you do, at least wash well (boil) with water, or better, 5% acetic acid or dilute HCl (<.5%, cold only) to dissolve off metal oxides, zinc & ammonium salts, etc., or KOH. The resultant crud is by no means all MnO2 but contains various hydrated oxides such as MnO(OH) and possibly Mn2O3, and carbon, about 10% C by weight.

I used to think any MnO4- ions might attack this carbon, but as far as I can determine, in alkaline solution there is in fact no problem (Carbon is attacked in acid solution).

I have tried things like heating to bright red heat to oxidize carbon and convert all the Mn to Mn2O3 or Mn3O4 while expelling all ammonia compounds (Zn/C case). It’s messy and must be done outside. Not worth the effort. Floatation to remove carbon is only partly effective. Batery crud should be ground as fine as possible.

It is probably not necessary to go through a thorough purification via HCl leading to chlorine and MnCl2 as proposed earlier in this thread. Instead, use NaOCl which will oxidize the oxides (AFAIK) and hydroxide to MnO2. Boil with 10% NaOCl if you have it.

(2) NaOH or KOH. See prior posts. Not too difficult to obtain technical grade, rarely pure and if so relatively expensive.

(3) Na2CO3 – ‘washing soda’, easily obtained, essentially 98%+ pure, as deca (large ice-like crystals) or monohydrate (white powder). Easily dehydrated to anhydrous. Or heat NaHCO3 (99% pure as baking soda, NOT powder) to about 150-170C to expel water & CO2 to get it.
K2CO3 is used in pottery and glasswork & and also obtained from other sources. Purity unknown. Anhydrous (very deliquescent, though). (Not easy to make from the Na salt.)

(4) Hypochlorites.
a. Sodium salt in bleach (~5%, Clorox) or pool supply (10-15%).
b. KClO or any other hypochlorite must be made (You can get Ca(ClO)2 from pool suppliers and maybe LiClO in tub and spa sanitizer). Use Ca(ClO)2 as outlined earlier in this thread to make KClO.

An experiment made to determine the amount of NaCl in Clorox and pool hypochlorite showed clearly that roughly equal molecular amounts of NaCl are present in all Na products tested (5% and 10%). Clorox appeared to have somewhat over 5% NaClO when fresh (MSDS says ~6%), pool 10% stuff at least 10%. The calcium products seemed to have a bit less than quoted ‘available chlorine’ but age of sample might be the reason. Calcium products also are loaded with insoluble Ca(OH)2 and/or CaCO3. CaCl2 is also present. They don’t dissolve well but Ca(ClO)2 is reasonably soluble. Difficult to get concentrated solution due to insolubles that don’t easily settle and are difficult to filter.

Der Alte

DerAlte - 6-11-2007 at 11:13

MANGANATES 3 - A FEW NEW RESULTS/COMMENTS:
Jottings from my notebook.

(1) What is the composition of the brown substance precipitated from Mn++ salts by hypochlorite? Is it hydrated MnO2.xH2O?

Answer, yes, it’s hydrated. A weighed sample of 1g+- 0.005g (according to my weights. I have two sets which are mutually consistent) was heated carefully on a SS spoon to below red heat (est ~350C; decomp temp of MnO2 ~ 500C)) for about 35 mins. It changed color to a much deeper brown, almost black, the familiar color of MnO2, on cooling. Reweighing, the weight was 0.795 g. This is close to MnO2.H2O.

(2) Is a manganite A2MnO3, A=alkali metal, a stage in the oxidation of MnO2 to manganates in alkaline condition? Answer=no.

The following was tried.: A sample of 5g was weighed carefully (Hydrated MnO2). A solution (in excess of any expected reaction) of 3.4N NaOH was made up (meant to be 5N, but later tests showed the NaOH was only about 70%). This was heated to about 95-100C on a small hotplate for 4 hours. On weighing after washing, filtering and drying at 150C the weight was found to be less than the original (about 7% less, possibly less hydration). Any manganite should have increased it. The color was the same as the original hydrated MnO2.

Note: This does not mean that manganate cannot be produced in either much more highly concentrated NaOH solutions or by fusion.

MANGANATES – 4 - A FEW NEW RESULTS/COMMENTS
The manganate series.

Question: Under what conditions are MnO4-, MnO4-- and MnO4--- ions produced by chemical oxidation, specifically by hypochlorite?

Three solutions of NaOH in 10% NaOCl were made up. These were approx. 3.5M , 2M and 1M NaOH. On heating with a small quantity of MnO2(hydrated) in a test tube to just boiling, the following was observed:

3M solution – a darkish blue color, as sky near dusk. I do believe this is the hypomangante.

2M solution: dark green, manganate.

1M solution – permanganate coloration. Signs of mangante too, when a drop was placed on filter paper. Diluting 2:1 to ~0,5 N seemed to convert all to permanganate.

In addition, Na2CO3 with NaOCl was tried at an estimated pH of 11.5. Only permanganate coloration seen. The same happens with bicarbonate. It was concluded that the pH should be kept lower than about 13 for permanganate formation (1N NaOH is about pH 14, 0.5N about 13.7. 1N NaOH is 40g/L or ~4% w/w).
NaOCl 10% solution is 1.34 M in NaClO and will have a calculated pH of about 11. The MSDS says about 11.4. Since carbonic acid is a bit stronger than hypochlorous, adding carbonate to the solution will decrease the pH somewhat, (pKa HCO3- = 6.37, hypochlorous = 7.55). Carbonate thus acts as a buffer. These condition favor permanganate production rather than manganate.

The amount of free HClO in a 10% solution is low, (est ~ 10^-3) and it is usually assumed that unionized hypochlorite is responsible for bleaching and oxidation. In fact, diluting the solution increases the ‘hypochlorite’ odor. In the very basic conditions for manganate and hypomanganate, the presence of HClO is vanishingly small. So these reactions are probably with the ion ClO- rather than HClO.

If the reaction : 2MnO2 + 3KClO + K2CO3 --> 2KMnO4 + 3KCl + CO2

is what actually occurs, coupled with the disproportionation

3KClO -->2KCl + KClO3,

then from the Law of Mass Action, the equilibrium is shifted to the right of both reactions by increasing the concentration of KClO. However, in the permanganate case the factor is a power of 3/2, and for the chlorate case a factor a power of 3 of the concentration. Hence lower concentrations should favor permanganate. The reaction rates will be correspondingly slowed as well. Escape of CO2 as gas is also desirable, hence keeping the temperature high to expel it. Stirring will aid to keep the MnO2 circulating in suspension and contact with the reactants. Also, somewhat surprisingly, the ratio of products at equilibrium is in favor of KMnO4 production by a factor depending on the square root of the KCl concentration. So the presence of NaCl in the hypochlorite should not be deleterious, except that it has to be removed later.

Against efficient production of KMnO4 is the fact that the kinetics are against it. The chlorate production is a competing auto-oxidation. The Standard Electrode Potentials (SEPs) are close for both reactions but favor chlorate.

So, amidst all these conflicting suggestions the only way is to see what actually happens is to try it experimentally, using these ideas as a vague guide.

Der Alte

DerAlte - 6-11-2007 at 11:30

MANGANATES – 5 - EXPERIMENTAL RESULTS/COMMENTS
From the note book:

Several runs were conducted under slightly varying conditions. The average MnO2 weight used was about 10g.

Only one run was conducted with NaOH without any carbonate (except for that formed by age on the hydroxide). Normality was about 0.25 (1% solution), at temp ~85C . The reaction was so slow that impatience prevailed and the test was not completed.

The results with the five runs, using Na2CO3, conducted at from c. ~65 C to rapid heating at about the boiling point at somewhat over 100C, seemed to indicate that the main influence of temperature was to speed the reactions rather than markedly change the proportion of permanganate to chlorate produced.

Efforts to get the calcium salt by a similar process proved abortive (mere coloration).

One run was made with NaOCl with a lower proportion of NaCl present, made from the pool calcium hypochlorite. The others were done with 5% Clorox and 10% pool hypochlorite. Apart from easier removal of less NaCl, little difference was seen in the final results.

The Precipitated MnO2.H2O (so assumed) and the hypochlorite were in rough (~+-10%) stoichiometric ratio of 2:3. The Na2CO3 was dissolved in the hypochlorite solution , in about 50% excess, to maintain pH during the runs as carbonate and hypochlorite were depleted. Run lengths varied from 12 hours at C. 65C to 3 hrs at 95C.

The NaMnO4 produced was estimated from the amount of unreacted MNO2. This is tedious, like many of the processes required. It is not easy to measure the permanganate directly due to the presence of unreacted hypochorite so reduction titrations produce erroneous results. Provided the solution is kept sufficiently alkaline, all the Mn used should go to MnO4- ion, since this is the only stable form of Mn, apart from MnO2, in alkaline oxidizing conditions.

The results showed that the conversion based on available oxidant was between 19.5% and 25%. The rest of the hypochlorite assumed get converted by disproportionation to chloride and chlorate, which enables one to estimate the amount of chlorate produced.

Conversion to KMnO4 and extraction with acetone proved difficult. Keeping the temperature at 10C or lower and drying to remove water helps any degradation, but the solubility of KMnO4 is still poor, at around 10g/100cc. {Does NaMnO4 dissolve better, as it does in water?}. The mixture with around 3-4 times its weight of KClO3 was obtained via KCl by fractional crystallization and carefully drying in air and then a desiccator (silica gel).

Although the chlorate and permanganate are not isomorphous, microscopic examination showed the permanganate is included in crystals of chlorate and hence does not dissolve easily. Fine grinding helps. At best only 60-70% of the available KMnO4 is extracted , thus reducing a poor yield to a pathetic yield of around 14%.
I have one or two other ideas, but I need a rest from permanganates!

CONCLUSION:
No wet method tried has yet produced a useful yield. If tried, it should be a bit better using only potassium salts. Lithium might allow a separation of permanganate from chlorate in aqueous solution but is a bit more expensive unless recycled. The carbonate is poorly soluble, however.

The only stable manganate is the insoluble barium salt.

The electrolytic method using a slurry of MnO2 seems to offer the most promise (with catalytic amount of KMnO4 to start, as in the patent?). See Xenoid’s posts supra.

Regards,

Der Alte

Antwain - 6-11-2007 at 12:36

Quote:

Originally posted by DerAlte
MANGANATES 3 - A FEW NEW RESULTS/COMMENTS:
Jottings from my notebook.

Against efficient production of KMnO4 is the fact that the kinetics are against it. The chlorate production is a competing auto-oxidation. The Standard Electrode Potentials (SEPs) are close for both reactions but favor chlorate.

Isn't that thermodynamics? All the stuff above was kinetics.

You say temperature only effects the speeds of both reactions... What would happen, hypothetically, if you had a hot suspension of MnO2 and you ran hypochlorite into it with stirring? (slowly)

The concentration of ClO- would be low at any time, at least in comparison to the concentration of MnO2 (which your solution would be brimming with).

The only problem there (if it is a problem, I don't know as much about this as you) could be if permanganate can oxidise chloride under those pH conditions. Otherwise it should solve all problems?

DerAlte - 6-11-2007 at 22:13

@ Antwain

You said:

Quote:

Isn't that thermodynamics? All the stuff above was kinetics

Yes. Incorrent terminology. Kinetics relates to rate of change.

I am no expert! Just an amateur, an antique who has been interested in chemistry since age 10 or so, both on the theory side and the practical side. My interest in the transition metals is because many reactions are a challenge and the chemistry is complex.

I have always had a conceptual problem with the "concentration" of solids in a liquid phase. It seems to me that many texts gloss over this.
In equilibrium considerations the solid phase is assigned the arbitrary fixed concentration of unity, provided at least some ('enough') is present.

While this may be considered satisfactory for equilibrium cases, where the rate of forward and backward rates are equal, it does not make logical sense when considering kinetics. There the reaction rate is obviously dependent upon the surface area of the solid reactant.

Quote:

You say temperature only effects the speeds of both reactions... What would happen, hypothetically, if you had a hot suspension of MnO2 and you ran hypochlorite into it with stirring? (slowly)

The concentration of ClO- would be low at any time, at least in comparison to the concentration of MnO2 (which your solution would be brimming with)

Temperature affects reaction speed exponentially, concentration linearly. If my argument above is correct, reducing concentration should favor permangante. Increasing temperature should merely reduce reaction time, not the ratio of products. The net gain from reducing concentration of NaClO, ClO- or HClO, whatever, will result in a slower reaction rate, but only inversely linearly. Your idea might work to do this, hypothetically.

There is another factor I might as well mention now. The type of MnO2, using the suface area argument, affects the reaction rate. There sre several types of dioxide, of which the two chief are beta-MnO2 (pyroluste, AFAIK) abd gamma-MnO2, which is the grade used in batteries. The gamma version is more 'active ' chemically. See

http://www.uspatentserver.com/686/6863876.html.

I am now wondering how to make 'activated' dioxide chemically. Doing it electrolytically is not a wet method (cheating). I wonder which type the hydrated stuff, produced as above from hypochlorite and a Mn++ salt, would be called and how it rates against gamma for activity.

The gamma form has a larger effective surface area. The faster the permanganate reaction can be made to go versus the chlorate, the higher the yield.

Regards,

Der Alte.

MnO2 Types

ciscosdad - 6-11-2007 at 22:30

There is lots of info on MnO2 in Patents related to Batteries that use the stuff. If I remember correctly, at least one of the patents on the list I posted earlier in the thread was from that source. I was looking for hints on producing MnO2 from available Mn Salts.

DerAlte - 3-12-2007 at 21:48

I've been at it again. First a few thoughts...

MnO2 Types and reaction rates

@ciscosdad & Antwain et al:

The type of MnO2 used radically affects the reaction rate.

I have assumed above that the reaction can be expressed stoichiometrically as:

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

Or, more generally, for any manganate, n=1 to 3 (n=4?)

2MnO2(s) + (4-n)ClO- + 2nOH- --> 2MnO4(n-) + (4-n)Cl- + nH2O,

(in aqueous solution). Written this way, the X+ ions are mere spectator ions. The carbonate used in the KMnO4 production above provides OH- ions.

Notice that as n increases, i.e as [OH-] increases and [ClO-] decreases, it would be expected that the product would change from permanganate to manganate to hypomanganate, as I confirmed by experiment (see supra). In order to get manganate or hypomanganate, only NaOH will provide the required pH.

The reaction rate of production of MnO4- (when n=1) is equal to the depletion rate of MnO2 and the other reactants, which is proportional to a power of the LHS concentrations (activity, to be more exact). Note: this assumes that the above reaction is the rate determining reaction. As written this is a 7th order reaction, a bit hairy. It is certainly slow, but so are many ionic oxidations. We get

d[MnO2]/dt = -k* [MnO2(s)]^2 * [ClO-]^3 * [OH-]^2 mols/sec

where k is a reaction rate constant (not the equilibrium constant K).

This tells us that the rate is determined by a 3rd power of the concentration of ClO-, and as the square of the OH- concentration and the MnO2(s) concentration.

In equilibrium equations the concentration of solid phases is taken as unity, but in rate equations this is not admissible. Logic tells one that the effective concentration must be due to the exposed surface area of the solid. “Activation” must therefore increase this active area somehow, perhaps by etching the crystalline structure.

Consider the disproportionation next

3ClO- --> 2Cl- + ClO3-

It is a third order reaction so that

d[ClO-]/dt = -k1*[ClO-]^3 mols/sec, where k1 is the reaction rate constant, different, of course, from k above. However, it is still proportional to the cube of the ClO- concentration as for the permanganate reaction. So [ClO-] affects both equally. To keep the reaction rates up the higher the concentration the better. Both reactions deplete ClO- ion so the net effect is complicated.

All I can say is that chlorate appears to be produced at about 4-5 time the rate of permanganate in previous experiments.

Note that the permanganate reaction goes to completion, in theory, in excess ClO- because the carbon dioxide is eliminated from the system and the sodium ions the carbonate carries are used up in electrical neutralization of the permanganate ion. In contrast, the manganate and hypomanganates are true equilibrium reactions. The chlorate/hypochlorite should also reach an equilibrium since all components are in the liquid phase. Admittedly it must be far to the right.

What does an increase of temperature do to the reaction rates? It speeds both up, according to Arrhenius’s theory, at a rate proportional to exp(-E/RT). E is the (unknown) activation energy for the reactions. The best we can do is to assume that a common rate of increase is doubling per 10K but since we don’t know the rate we cannot make any firm conclusion. Raising the temp from 35C to 95C will roughly increase both rates by a factor of 2^6 =64. Because the permanganate reaction is dismally slow, it takes too long at too low a temperature.

In view of all this, then, the only approach to improving the yield of permanganate WRT chlorate seems to be to make the MnO2 as “active” as possible. This is achieved by activating it and using as much as possible to present a large surface, and also to keep it well stirred. The next two posts are concerned with a few recent experiments in the light of these thoughts.

Der Alte

DerAlte - 3-12-2007 at 21:54

Some results...

A test was undertaken to determine the reaction rate, crudely and qualitatively, by heating various samples of Mn compounds with the same strength carbonate and hypochlorite solution (in considerable excess). The rate was estimated by the relative coloration produced in a given time, one minute at 100C (actually, boiling) followed by cooling (when the rate drops radically). Only a small amount of the Mn compound was used in each case, with the about the same Mn content. The results were as follows:

(1) From Mn2O3 produced from Leclanche battery reclamation, fired at bright red heat to burn off carbon and ammonia compounds and treated with dilute acid to remove Zn compounds. Very slow, with only slight oxidation to permanganate.

(2) From MnCO3. Faster than (1) but still a very slow reaction. The carbonate is first oxidized to MnO2, with a black color.

(3) From newly precipitated Mn(OH)2. Black MnO2 is produced first. Then, slowly, a weak coloration of MnO4- ions was seen. The next slowest reaction overall.

(4) Using brown (hydrated) MnO2 precipitated previously and recovered from previous attempts. Faster than (3), as estimated by color produced.

(5) Using MnCl2 directly. This was faster than (4), marginally, but still slow. (Brown MnO2.H2O is produced first). Since (2), (3) and (4) and (5) use hydrated MnO2 precipitates in effect, it was considered that these were not in activated form.

(6) Using ‘refined battery crud’ from alkaline cells. No attempt made to remove graphite, but it was treated with cold dilute HCl (c. 0.1 M) for a hour to get rid of ZnO and zinc compounds. This was very much faster than the previous results. Now it is known that cells use gamma MnO2, prepared electrolytically. How much of this remains in a spent cell I have no idea.

(7) Using previously precipitated and stored brown MnO2 as in (4), treated with dilute cold HCl as in (6). This was as fast as or faster than (5). The treated product was a different color, deeper brown, almost black, more like the battery product.

Conclusion: treatment with cold dilute HCl must produce an activated form. Whether this is gamma I do not know. Note that pyrolusite, generally the form available in pottery sources, is beta MnO2 and not very active.

Der Alte

DerAlte - 3-12-2007 at 21:59

Following up on the last posting, I used some ‘refined battery crud’ from alkaline cells (a weighed quantity) in a small batch, cooked it with the usual sodium hypochlorite (10%)/carbonate solution, and heated rapidly to near boiling for an hour, with occasional hand stirring.

The reaction took off quickly. Carbon dioxide bubbled off was and expelled, the suspension turning rapidly reddish. The graphite could be seen floating as a metallic sheen on the surface. In time it disappeared, being assumedly oxidized by either hypochlorite or produced permanganate. It does not matter which, since permanganate is reduced to MnO2 and merely recycled. Essentially, the oxidation uses hypochlorite. The final solution after settling the unused MnO2 was very dark, and a strong flashlight could only penetrated a layer about 1cm thick. Also, the time was about 1/3 of my other recent attempts.

The liquid was decanted and subjected to a series of evaporations and freezings to remove NaCl in large quantities, excess carbonate and some chlorate. After several cycles, ending in a freezing (so the solution was not saturated in NaCl, which had been partially separated as NaCl.2H2O) the liquid was allowed to settle for 24 hrs. and the somewhat viscous liquid poured off.

A drip of the liquid was placed on a microscope slide and left to evaporate at 40% RH, and observed at X50 with transmitted light. First to crystallize out were cubic crystals, followed by a stepped pyramidal form. These were assumed to be chloride and chlorate. The remaining purple liquid on the slide then took a long time to evaporate, but finally bunches of small brownish red needles like cactus spines deposited. It looked like KMnO4 crystals, but the NaMnO4 is hydrated (+3H2O, CRC) and might be expected to have a different shape. All these crystals occurred in the expected order of solubility and separately.

A hot, strong KCl solution was added to a small portion of the liquid and examined similarly under X50 magnification. The same cubic crystals appeared first, then elongated crystals with dark inclusions. I assume these were chlorate plus co-precipitated permanganate. The potassium chlorate being in excess seems to cover up and enclose the potassium permanganate, which may account for the poor extraction with acetone seen before.

I intended to dry off the solution, at moderate temperature, to see whether I could dissolve NaMnO4 with acetone in the same way as KMnO4. (A search gave no results of the solubility of the sodium salt). But disaster struck. I had left the solution on what I thought was moderate heat. I forgot about it and did something else. When I returned I found the heat was not as moderate as I thought and I had left it unattended too long. The entire bench was splattered with nasty red blobs and little remained in the Pyrex vessel being used. The mess was unbelievable!

So, does anyone have a figure for the solubility of sodium permanganate in acetone?

Der Alte

Antwain - 4-12-2007 at 16:43

Just a thought, but quickly precipitated solids usually have a much larger surface area and energy than 'dry' ones. How about preparing MnO2 in-situ by adding NaOCl to MnCl2 solution. If you have a manganese salt or MnO2 this can easily be prepared. By using the chloride you are not adding further unnecessary ions to the solution, just a slight excess of chloride. Well, actually a bit more than that because of the chloride from the reduced hypochlorite too.

I would try dissolving the MnO2 with HCl outside - unless you are partial to breathing chlorine

- and filtering then crystallising the MnCl2, then dissolving that in a decent amount of water and adding hypochlorite with good stirring.

Edit- actually, you may end up with less chloride, since you are not oxidising carbon.

[Edited on 5-12-2007 by Antwain]

cristiro - 31-3-2008 at 09:26

Bretfi method :cool:

Dear friends, I confirm that direct fusion of MnO2 and NaOH can lead to NaMnO4. I've done an empirical experiment which proved succesfully.
For this experiment was used:
- pottery grade MnO2,
- domestic use NaOH,
- tap water.
Description:
1. In a large tray was melted a teaspoon NaOH then powdered about the same quantity of MnO2 over melt then mixing for about 1/2 hour until a dark green liquid is formed. If cooled will crystalize like a glass film which breakes and care should be taken because small pieces when cooling can jump around; ( for this first step using a large tray is not a good idea). This product contains probably Na2MnO4, NaOH and MnO2; we call it Product A.
2. In a large tray, Product A is melted in a thin layer. MnO2 in excess is added until mixture becomes powdery and abrasive. Continue mixing for 1/2 hour so the mix to be well aerated while from time to time pulverize some water over powdery mix. Stop watering and keep heating for another 1/2 hour while mixing. Product B is formed, dark, which probably contains Na3MnO4, MnO2 etc.
3. Let the product B cool and then disolve it the cold water. Wait for an hour when the dark blue like plums is formed togeter with some mud residue. Decant the liquid and heat until a purple color is formed. Can be boiled untill concentrated. The purple color is also formed without further heating, after a while. What we obtained is NaMnO4 with very little impurities.
As stated before, this exp. was done empirically and a overwhelming positive result was not expected. Some of the parameters (time, quantity) might be slightly innacurate.

chief - 14-7-2008 at 07:20

above someone mentioned the following reactions:
$$$$$$$$$$$$$$$$$$$$$$$$$$$$$
KMnO4 + AgNO3 -> KNO3(aq) + AgMnO4(s) red precipitate : wiki claims 0.9g/100g water @ 20*C, but don't trust it

2AgMnO4 + BaCl2 -> Ba(MnO4)2(aq) + AgCl(s)

Ba(MnO4)2 + H2SO4 -> BaSO4(s) + 2HMnO4(aq)
$$$$$$$$$$$$$$$$$$$$$$$$$$$$$

I find this interesting: ppt-out the AgMnO4 from the watery solution of any melt or electrylytically processed melt; from there getting any desired HMnO4-salt ...

But: In a half- or not-at-all elecrolyzed melt from (K,Na)(NO3,CO3,OH) + MnOx there may be a lot of things going on .. : Is the AgMnO4 the only Salt to fall out with AgNO3 ?

That would seem to be the Kings way, most easy of all ...

ScienceSquirrel - 14-7-2008 at 07:36

The way to make potassium permanganate on a small scale is by fusing manganese dioxide with potassium hydroxide and potassium nitrate.
Extract the cool solids with water to produce a solution of potassium manganate ( green)

Acidification converts it to potassium permanganate.
This easily makes potassium permanaganate on a gram scale and was a standard A level preparation.
I would be wary of trying to make HMnO4 as it readily dehydrates to Mn2O7.

blogfast25 - 14-7-2008 at 07:59

In dilute solution permanganic acid (?), HMnO4, may be harmless but treating a pure permanganate with strong acid causes the highly instable Mn2O7 heptoxide to form. It's a recipe for explosions...

chief - 15-7-2008 at 00:55

Quote:

Originally posted by ScienceSquirrel
The way to make potassium permanganate on a small scale is by fusing manganese dioxide with potassium hydroxide and potassium nitrate.
Extract the cool solids with water to produce a solution of potassium manganate ( green)

Acidification converts it to potassium permanganate.
This easily makes potassium permanaganate on a gram scale and was a standard A level preparation.
I would be wary of trying to make HMnO4 as it readily dehydrates to Mn2O7.

Yes, only that I want to do it with the NaNO3 and NaOH, in order to have a NaMnO4-Solution, as concentrated, as I can make it !
(I wouldn't go via the H2SO4 (if doing the silver-route), because of the danger, but try to use Na2SO4.)
Can I use a mall amount of HNO3 for the acidification ? As I observed in the past, this readily turns green solutions into violet ones. But also quantitatively ? Is the HNO3 only catalysator, and small amount sufficient ?

blogfast25 - 15-7-2008 at 06:40

Quote:

Can I use a mall amount of HNO3 for the acidification ? As I observed in the past, this readily turns green solutions into violet ones. But also quantitatively ? Is the HNO3 only catalysator, and small amount sufficient ?

Well, I'm no expert on making permanganate. I've only obtained it inadvertently in alkaline oxidation of Mn(OH)2 with hypochlorite, as described on the previous page by DerAlte. Also by fusing MnO2 with KOH I got small amounts of K2MnO4 but never managed to convert it to permanganate, it always just fell back to plain old MnO2.

But acidification of manganate solutions should be fine: the danger of Mn2O7 formation only occurs (I believe) when treating pure (or maybe also very concentrated) permanganates with strong acids, for instance:

KMnO4 + 1/2 H2SO4 ---> 1/2 K2SO4 + 1/2 Mn2O7 + 1/2 H2O

If you're acidifying in normal concentrations, starting from a manganate (not permanganate) the risk of making the heptoxide inadvertently must be very small...

The_Davster - 15-7-2008 at 08:45

The chance of making Mn2O7 starting from any aqueous solution is impractically small. The stuff is destroyed by any ammount of water, so all the water solvent of your solution would have to be used up hydrating your sulfuric acid...not gonna happen! (unless you count taking your 100mL reaction and dumping it in a couple litres of conc. H2SO4)

Also, KClO3 is a far superior oxidizer for fusing with KOH and MnO2, and it is completely impractical to not use an oxidizer at all. Atmospheric oxidation of heated manganate is likely only feasible in industry with fancy heated grinders.

CO2 is the best acidification agent. No remainder ions to clean up and readily accessible from baking soda and vinegar.

ScienceSquirrel - 15-7-2008 at 09:52

I am not sure why potassium nitrate is preferred to potassium chlorate but you will need an oxidiser on a lab scale.
As I remember it the reagents melt to a black paste which solidifies as the reaction proceeds.
We used sulphuric acid to acidify the green manganate solution, you only need a small amount and it is fast and convenient.
Potassium permanganate is preferred over sodium permanganate as the potassium salt is a lot less soluble and readily crystallises.
It should also be noted that sodium permanganate cannot be prepared by this method.

http://en.wikipedia.org/wiki/Potassium_permanganate

http://en.wikipedia.org/wiki/Sodium_permanganate

chief - 15-7-2008 at 11:45

On the wikipedia-link they don't mention, how then the NaMnO4 is prepared. Since _this_ is, what I want, I'm gonna try some melt with Na2CO3, NaOH, NaNO3, and the MnOx, and check for the green color.
As far as I know, the permanganate-group forms under oxidizing and alcalic conditions, why should not then NaMnO4 form ? Maybe the temp has to be kept below the decomposition of the NaNO3, so that no NaNO2 (said to be reducing ..., therefore bad for permanganate-group) can form ...
Besides I have some Ba(NO3)2, which decomposes not below 595 [Celsius]; maybe I'll have to use some eutectic of the both, or only Ba(NO3)2 at all ? The Ba(NO3)2 can always be regenerated, using carbonate to ppt., then dissolving it in HNO3 again ....

[Edited on 15-7-2008 by chief]

Ion exchange

chloric1 - 15-7-2008 at 13:58

IIRC there is a British patent for making calcium permanganate from potassium permanganate via ion exchange with zeolites. What it said was you set up a column and run 30% calcium chloride solution through the zeolites to convert them to caclium zeolite. Afterwards you pass a warm solution of potassium permanganate though said zeolite. Now they said you could condense and chill the effluent to precipitate unconverted potassium permanganate as it is slightly soluble or rerun it through the zeolite a time or two more to get a more complete conversion. Ideally you could do this and add sodium carbonate to precipitate the calcium but I am unsure about permanganate stability in alkaline solution. Maybe adding baking soda with vigorous stirring then gradual heating so as not to raise the pH of the working solution too sharply.

Otherwise you could "charge' the zeolite with a few passes of concentrated sodium nitrate solution and then add your potassium permanganate to get a more direct route. I just don't know if the zeolite will exchange sodium for potassium so readily. I still have yet to try this but it is intriging.

[Edited on 7/15/2008 by chloric1]

ScienceSquirrel - 15-7-2008 at 14:39

This is a slightly interesting link dealing with the production of sodium permanganate.

However after the initial stuff the trolls have got at it and it is garbage as far as I can see.

http://backyardchem.chemicalblogs.com/121_backyard_chemistry...

chief - 15-7-2008 at 16:13

Thanks for that link; that's exactly the easy way, that I like:
Heat the MnO2 + "Soda - Saltpeter", equal parts, to "dull red heat" for 16 - 48 hours.

Only: Since I have the MnOx from old Batteries (Mn3O4 ?)(of course washed within lots of water to remove the NH4Cl), maybe there is some residual graphite within; may the graphite go off with the NaNO3 ? Would not be nice with a 1 kg-batch ..., also the explosion could be much stronger than gunpowder, since the stuff would be pre-heated ?

Can I glow the graphite away, and at which temperature ? Some sources say at least, that the Mn3O4 may be oxidized to MnO2 by glowing on air.

Also MnO2 is known as an oxidizer, with metal-powders at least ...

And then: Does "Soda-Saltpeter" mean: "NaNO3", OR "Na2CO3 + NaNO3 (ratio?)" ?

[Edited on 15-7-2008 by chief]

[Edited on 15-7-2008 by chief]

[Edited on 15-7-2008 by chief]

ScienceSquirrel - 15-7-2008 at 16:35

Looks like you are in the running for a Darwin Award

chief - 15-7-2008 at 23:00

No, no. Is it dangerous now ? I only mentioned the MnO2-as-oxidizer since I'm not quite sure, how the graphite would react in such a fine-powdery form with an oxidizer ...

Besides anyhow the thing will run in my furnace, without me around ...

Why Darwin-Award ???

chief - 16-7-2008 at 01:02

Someone might also be careful about the Zn-Powder within alcaline cells ...

ScienceSquirrel - 16-7-2008 at 02:41

I think that soda - saltpetre is probably sodium nitrate.

Why are you messing around with manganese dioxide from batteries? Manganese dioxide is readily available from a variety of sources, as is sodium nitrate.

I would buy the pure chemicals and try the reaction on a small scale to start with. More chance of the reaction going right and less chance of a nasty accident.

chief - 16-7-2008 at 04:04

Yes, small scale, or in a distant, fireproof room ...

not_important - 16-7-2008 at 04:39

Quote:

Originally posted by chief...
Can I glow the graphite away, and at which temperature ? Some sources say at least, that the Mn3O4 may be oxidized to MnO2 by glowing on air.
...

Yes, you can remove the carbon through air roasting. However, some sources are wrong. The most stable oxide is the mixed one - Mn3O4 (MnO. Mn2O3) which is formed by heating any of the other oxides in air above 950 C.

Mn2O3 can be had by air ignition in the 500-900 C range, although conversion to that oxide can be very slow.

MnO2 starts to break down at temperatures below 300 C on up to over 500 C, depending on how it was originally made.

Formatik - 16-7-2008 at 10:40

Quote:

Originally posted by blogfast25
In dilute solution permanganic acid (?), HMnO4, may be harmless but treating a pure permanganate with strong acid causes the highly instable Mn2O7 heptoxide to form. It's a recipe for explosions...

It needs to be dehydrated also to form the compound. I've added some KMnO4 powder to reddish fuming HNO3 with a density of about 1.52, no Mn2O7 formed, and all of the acid fumes turned white. The same is in the well-known chlorine generator reaction of KMnO4 with conc. HCl, though the permanganate can also oxidize the conc. acid to some ClO2, which is explosive.

12AX7 - 16-7-2008 at 11:59

Quote:

Originally posted by not_important

Mn2O3 can be had by air ignition in the 500-900 C range, although conversion to that oxide can be very slow.

It works pretty well in the presence of charcoal. Getting Mn3O4 or MnO the same way takes considerably more heat, and gives off much less in return.

Tim

chief - 17-7-2008 at 12:39

I heatet to 530-555[Celsius] the Na2CO3-NaNO3-MnOx (latter fom batteries, washed and cooked 1/2 h in soda-sol.), equal parts of each component, 12 hours long. It was dry, fine milled, and mixed in the coffe-mill, before the heating.
After cooling down, with water: Not the slightest green, no color. Only the black MnOx now is brown, so maybe I at least oxidized the MnOx to MnO2 ...

12AX7 - 17-7-2008 at 19:19

Did it fuse completely? Carbonates take a good bit more heat to work with. Try hydroxide.

Tim

not_important - 17-7-2008 at 19:47

Hmmm ... the nitrate would have melted at that temperature, but as it decomposes/gets used up the mixture will become solid. I think you need to heat to 800-900 C at least, perhaps not mix all the nitrate into the original mixture but rather slowly add it to the fused mass with stirring.

DerAlte - 17-7-2008 at 20:40

Manganates Again

Since it seems that this old thread has been revived, this is a good time to put to rest a few myths I may have inadvertently engendered earlier. Reference should be made to all the previous posts – it’s too difficult to sort out which are relevant.

Since last year I have repeated a few of the experiments. My conclusions are : (1) The production of permanganate by any wet method is far less then I thought; (2) The fusion methods do work, though how well I have not determined. I have never had much success with manganate production before, like mant others.

WRT (1), I have been consistently fooled by how intense the color of the permanganate (and manganate) ion is. An intensely colored solution can be achieved by a few percent by weight. If both manganate and permanganate are present, an almost total absorption of most of the visible spectrum occurs. The permanganate ion exhibits a broad absorption band around 520 nm, blocking out the entire green region and leaving a bit of the blue and the yellow, but has less absorption in the red and blue violet – hence the well known magenta color. In fact, I can still see the color at 3ppm.

Manganate has a wide absorption band in the red region, centered on 620 nm, which is why it appears green. Add the two together and you are left with attenuated passbands mainly in the blue/violet. Since my source of illumination has been a krypton filled filament lamp, very little light passes when manganate and permanganate are both present ( this can be detected by filter paper chromatography).

In short: dense looking solutions of permanganates are not very concentrated.

Most of my efforts were concerned with pHs at a level where permanganate only can be produced (pH~ 11). In one effort that I thought particularly promising, I used ~ 1N NaOH instead of Na2CO3 (pH ~ 14) to get a very dark mainly manganate solution – dark for the above reason.

I castigated the acetone extraction unfairly. I now think it works like a charm – there is just very little permanganate to be extracted, and it does extract nearly all of it. But the acetone must be carefully dried (CaSO4) and so must the product from which the permanganate is extracted. Very little of the permanganate gets reduced. Any MnO2 produced tends to color the remaining undissolved chlorate & chloride slightly brownish/red, possibly due to methanol in the commercial acetone.

The titrations I performed earlier to determine the presence of oxidizer in solution were severely flawed. Hypochlorite, I have found, is not so easy to destroy rapidly by heat as I supposed, and its effect cannot be distinguished from MnO4- using the Fe++ ion. I.e. remaining hypochlorite is still present and masks the permanganate, giving an inflated number.

Instead of using Fe++ ion, I repeated the measurement using glucose and ethyl alcohol as reducing agents in alkaline solutions in excess, and measured the MnO2 produced (carefully washed and dried) by weight. Reasonable consistency using these two methods showed that the manganese content implied that the oxidation efficiency to be no more than 2%. This also agreed roughly with the amount extracted by acetone in another run.

Somewhat disheartened by this, I did some detailed calculations on what might be theoretically expected. I included all the factors I could think of, such as concentration of ions, temperature. I spent some time seeking out accurate values for the Standard Electrode Potentials. For the reactions involved, these are none too certain but consistent to about a few %. The results of this suggested I was lucky to get ~2% product; the conversion to chlorate dominates.

The next post outlines some theoretical considerations.

Regards,

Der Alte

DerAlte - 17-7-2008 at 20:50

Attempt at Theoretical Reconciliation

*** Ignore the following if you are allergic to theory!

The half-reactions involved, ignoring spectator ions, are as follows (in alkaline standard conditions), written as reductions:

MnO4- + 2H2O(l) + 3e- --> MnO2(s) + 4OH-, Eo = 0.595 v ..(1)
MnO4-- + 2H20(l) + 2e- --> MnO2(s) + 4OH-, Eo= 0.60 v ..(2)
(Notice how close these SEPs are. But the one electron difference is crucial!)

ClO- + H2O(l) +2e- --> Cl- + 2OH- Eo=0.81 v ..(3)

The above are taken from CRC. We also need:

ClO3- + 2H2O(l) +4e- --> ClO- + 4OH- Eo=? (0.525 v) ..(4)

This value is not in CRC and has to be calculated from those that are. I used

ClO3- + 3H2O(l) +6e- --> Cl- + 6OH- Eo=0.62 v ..(5),

in conjunction with (3) above. The Eo for half reaction (4) is then obtained by subtracting the reaction (3) from (5) and taking account of the electron exchange ratios.
4Eo(4) = 6Eo(5) - 2Eo(3) so Eo(4) = 0.525 v.

The balanced ionic full reactions for the production of manganate and chlorate ion from hypochlorite can then be deduced:

3ClO- --> ClO3- + 2Cl- {from (3) & (4) above, eliminating e-}; dEo=0.265 v ..(6)

2MnO2 + 3ClO- + 2OH- --> 2 MnO4- + H2O + 3Cl- ; dEo= 0.215 v ..(7)

Plus the also ran case producing manganate
MnO2 + ClO- + 2OH- --> MnO4-- + H2O + Cl- ; dEo=0.21 v ..(8)

Or even the long odds for Mn(V) (hypo)manganate
2MnO2 + ClO- + 6OH- --> 2 MnO4--- + 3H2O + Cl-; dEo= -0.15 v. ..(9)

What does all this crapola mean? Basically, that the reactions tend to go to the right the higher dEo is. For the manganate case, dEo is negative, meaning a tendency to the left. The conversion to chlorate has the highest driving force.

The above values are at the standard conditions T=25C (298K), P=101.33kPa (1 std. atmosphere). I carried out the reactions at ~95C (368K), which modifies the values a bit. Also, the standard conditions are for solution concentrations as indicated by the equations, in mols/l.

For any other conditions (we can ignore pressure because the reaction are concerned with liquid/solid phases only) we have to use the Nernst equation

E = Eo - (RT/nF)lnQ = Eo – (0.0592T/nTs )log10(Q), in volts.

Where R= (molar) gas constant (J/deg/mol), T is temp (abs.), Ts is the Standard temp ( 298K for most SEP tables), F (faraday) = 96,485 coulomb/mol, n is the number of electrons exchanged in the reduction equation, Q = concentration ratio, products divided by reactants, as in the equilibrium equation. In fact, if K is the equilibrium constant, Q = K when E tends to zero – there is then no driving force either way. Hence:

If at equilibrium E = 0, Eo = (0.0592T/nTs )log10(K)

i.e. K = 10^(nEoTs/0.0592T) = 10^(nEo/0.0731) at 95C. Simple enough? It’s now in terms most amateur chemists can understand (Sorry, kewls.)

For (6) above, chlorate production @95C, n=4, Eo = 0.265 volts, K6 = 3.16E14
For (7), permanganate , n=3 per MnO4- ion, Eo=0.215 v., K7 = 6.66E8
For (8). Manganate, n=2 per MnO2-- ion, Eo=0.31, K8 = 5.57E5
For (9), Hypomanganate, n=1 per MnO4--- ion, Eo= -0.15, K9 = 7.87E-3.

What does all this mean? On the face of it, that chlorate is by far the most likely to be produced; then permanganate, then manganate, and hypomanganate has only a cat’s chance in hell of being produced.

But let’s look a bit deeper. Consider the chlorate case. The reaction assumed is
3ClO- --> ClO3- + 2Cl-

{First notice that this reaction does not depend apparently on alkalinity (i.e. [OH-]). This confused me at first, but actually the equation implicitly assumes the solution is fairly alkaline: in acid solution the reaction differs, because instead of ClO-, hypochlorite ion, the hypochlorite will be present as unionized HClO due to the weakness of the acid HClO}

We have
[ClO3-][Cl-]^2/[ClO-]^3 = K6 = 3.16E14
at equilibrium @95C, whenever that is attained. (It says nothing about the reaction rate, merely the final state. This might take forever, like for hydrogen and oxygen at room temperature. But we know it isn’t, by experiment!)

We have then, (forget the complication of actual activity, for the purists)
[ClO3-][Cl-]^2 = K6 [ClO-]^3 (@ equilibrium)

Typically, a 10% solution by wt. of NaClO has a concentration of 1.5M NaClO plus about the same of Cl-. Since it’s obvious this reaction goes to near completion, the final Cl- concentration will be ~ 2.5 M and that of NaClO3 about 0.5M. This gives the final concentration of ClO- at equilibrium, as about 1.6E-5.

Next look at the MnO4- reaction (7). Rewrite as:

MnO2 + (3/2) ClO- + OH- --> MnO4- + (1/2) H2O + (3/2) Cl-

The equilibrium equation is

[MnO4-] = K7 [MnO2][OH-][ClO-]^(3/2) / [H2O]^(1/2)[Cl-]^(3/2)
= K7 [OH-][ClO-]^(3/2) / [Cl-]^(3/2)

(Here the concentration of the dioxide and water are to be taken as 1, since the solid phase and liquid solvent are assumed always in excess.) If the pH is kept at ~11, then [OH-] will be about 1E-3. If [ClO-] is about 1.6E-5, determined by ClO3- production as above, then the concentration of MnO4- calculates as

[MnO4-] = 6.66E8 x 1E-3 x (1.6E-5 )^1.5 / 2.5^1.5 = 1.08E-2 mole/L = 1.28g/L. At pH ~11.5, this becomes 4.04g/l and at pH~12, 12.8g/L.

Pretty pathetic! But in line, roughly, with the experimental results.

The estimated production of MnO4- - can be calculated similarly as
[MnO4--] = K8 [OH-]^2 x [ClO-]/ [Cl-]  3.56E-6 mole/L @ pH~11. At pH~12, one gets 3.56E-4 mol/L and at pH ~13, 3.56E-2 mol/L

Thus in high pH conditions, manganate production becomes favorable. One might think that, per the above, permanganate production would be even more favorable; but, as pH rises, the following reaction takes place, in the presence of MnO2:

MnO2 + 4OH- + 2MnO4- --> 3MnO4-- + 2H2O

This can easily be shown experimentally by adding NaOH to a solution of permanganate in the presence of MnO2; the magenta solution turns green.

If the OH- concentration is very high (~ >5M or about 25% by weight) you can even go another step to Mn(V) (hypo)manganate:

MnO2 + 4OH- + MnO4-- (green) --> 2MnO4--- (light blue) + 2H2O

Correspondingly, small amounts of water readily hydrolyze hypomanganate to manganate and MnO2, and on further dilution permanganate and more MnO2 is produced.

Very similar reactions happen in the fusion case, where a liquid hydroxide and dissolved or fused oxidizer are present in very high concentration. A little on this in the next post.

{CAVEAT: the above is purely the work of Der Alte. As such, there are no guarantees from the management. The potential for BS exists, since Der Alte is very prone to errors in calculation, and has even been known to make logical errors due to senility. Criticisms and corrections form real physical chemists welcome}.

Regards,

Der Alte

DerAlte - 17-7-2008 at 21:14

Fusion

I have re-tried the age-old fusion reaction with an oxidizer (nitrate) and hydroxide. The proportions used were roughly according to

MnO2 + 2NaOH + NaNO3 --> Na2MnO4 + NaNO2 + H2O (gas)
(I.e. 2 mols NaOH to 1 mol NaNO3, (c 1:1 by weight) MnO2 in excess).

The usual recommended temperature is “dull red heat” which I take to be in the range 400 - 500C, which I believe is just visible in the dark. Time, 2-3 hrs.

The Japanese recommendation (see thread above) for Hypomanganate is 0.5mol MnO2, 2.5 mol KOH, 0.5 mol KNO3, 2 hours fusion at c 300C. I used the sodium hydroxide and nitrate

That reaction is assumed to be:
2MnO2 + 6NaOH + NaNO3 --> 2Na3MnO4 + NaNO2 + 3H2O (g)
This is slightly a different molar ratio, 1:3:0.5 versus the recommended 1:5:1, but I used the recommended.

In all cases I used MnO2 that had been made from MnCl2, see earlier in the thread. This is likely hydrated as approx. MnO2.H2O

I did the fusions in identical small steel cans, well cleaned. The NaOH and nitrate fuse together at something like 250C - 270C I believe. I used no temperature measuring device, merely making sure that the interior of the container glowed just perceptibly in the dark in the manganate case, and the mix remained fairly fluid in the hypomanganate case. I wrapped the can in a sheath of glass cloth for the manganate case, and heated it with a small propane flame from a blow torch. I added MnO2 once the solids melte, then slowly raised the temperature to the dull red heat. The other case I heated on a camp stove, maintaining a sort of fluidity.

Both the NaOH and the nitrate that I used were damp and steam was emitted before the solids finally melted.

I stirred the mixes from time to time manually with a plated steel rod. Both turned pasty in time. I added the MnO2 in small batches. It fizzled on introduction to the fused liquids, especially in the manganate case. This may have been due to the fact that the MnO2 was probably hydrated and the reaction also produces steam.

On cooling, the resulting solid mass in the manganate case was a dirty brown with a hint of green. In the hypomanganate case, it was quite similar to that shown so well in Xenoid’s posts above, bluish with a hint of green and some blackish MnO2 specks.

I treated both messes with a small amount of water. Sure enough, both gave the deep green color of manganate.

Conclusion: With care, the fusion reactions work. I am now convinced. How well is TBD.

Regards,

Der Alte

[Edited on 17-7-2008 by DerAlte]

[Edited on 17-7-2008 by DerAlte]

chief - 21-7-2008 at 03:11

Success !
What I did:
==> I weighted in [5.72029375,16.9989416,7.9491552] gm of [Mn3O4,NaNO3,Na2CO3],
==> glowed it some minutes at 800 [Celsius],
==> with water (not directly (!): drops explode (!) by steam-generation, cool down first)
==> gives deep green solution

The solution, with a tiny amount of HNO3, gives violet solution, but a tiny bit more reddish than what I know from KMnO4. When boiling down, it at first became violet, but turned back to green.

The "Mn3O4" used was from batteries, washed, glowed and pre-oxidized at 500 [Celsius] with NaNO3 (whereafter it was more brown than black).

The equations I used were:
--------------------------------------
MnO2 + 2 O ==> MnO4
Mn2O3 + 5 O ==> 2 MnO4
Mn3O4 + 8 O ==> 3 MnO4
----------
NaNO3 ==> NaNO2 + O
---------
Mn3O4 + 8 NaNO3 ==> 3 MnO4 + 8 NaNO2
--------
So _at_least_ 8 parts of NaNO3 would have to be used, because the NaNO2 is known as "reducing", but thats an assumption. Question may be: What wants more oxygen (?): MnOx OR NaNO2, OR: How oxidizing is a melt of NaNO3 with x % NaNO2 in it ??
---------
Then the final assumption was:
Na2CO3 + MnO4 ==> Na2MnO4 + CO2 + X, no exact equation, and experiment rules.

What do you think about it ? Its very green, and now I try to concentrate it and find a way to either crystallize something (for microscope) or the measure somehow the quantity obtained ...

But I wanted to share now, so maybe someone wants to try himself ...

[Edited on 21-7-2008 by chief]

DerAlte - 21-7-2008 at 14:15

@chief

First, congratulations on having such a wonderfully sensitive and accurate balance, capable of weighing accurately to a few nanograms! It must be truly unique and worth a lot.

You have forgotten about the carbonate. It neither melts easily (858C) nor decomposes below that temp., and I have not managed to find a eutectic with NaNO3, which melts at 307C (CRC) and decomposes at around 500 – 600C (way below 800C). However, Na2CO3 probably dissolves in molten NaNO3 to some extent.

NaOH melts at 323C and a mixture with NaNO3 somewhat lower.

Now, the usual fusion of MnO2 with NaOH is assumed to be as follows:

2NaOH + MnO2 + ‘O’ --> Na2MnO4 + H2O. Since the ‘O’ is necessary, it has to be supplied by an oxidant. But notice that also the NaOH is acting as an oxidant (very unusual).

If the oxidant is NaNO3, (which is a very poor oxidant in alkaline solution (Eo ~0.01 V), but in the fused state the concentrations are much higher), we get

2NaOH + MnO2 + NaNO3 --> Na2MnO4 + H2O +NaNO2.

Other oxidants such as chlorate can be used; commercial processes allegedly use air.

With Na2CO3 instead of NaOH, one might get

Na2CO3 + MnO2 + NaNO3 --> Na2MnO4 + CO2 + NaNO2

There is a very large difference between Na2CO3 in solution – which is mildly alkaline – and dissolved in NaNO3. There are no OH- ions. The presence of OH- seems essential to avoiding decomposition of magnates and also nitrates. In general the hydroxide must always be in excess of stoichiometric.

Hence I am very surprised you managed to get oxidation at 800C. which is a temperature very noticeably cherry red in daylight, and also that you managed it with sodium carbonate. Further, you probably had only somewhat impure Mn2O3 and not dioxide.

However, you obviously did if you got green manganate color and changed it to red permanganate. It’s a sensitive test.

Making permanganate is surprisingly difficult for the amateur, it would seem. The best result, reported above by Xenoid, was around 17% IIRC, via the hypomanganate and electrolysis. I’d like to try the Japanese patent process of direct electrolytic oxidation of a slurry of MnO2, preferably in KOH rather than NaOH.

Regards,
Der Alte

chief - 22-7-2008 at 00:41

Direct electrolytic oxidation: There is, from literature, at least such a way to get Ba(MnO4)2:
By electrolytically dissolving a Mn-electrode in (hot ?) Ba(OH)2-solution.
That I have from a BASF-chemistry-book.
So a possibility would be:
==> Get Mn from MnOx by some reduction (hopefully not violent)
==> make the Ba(MnO4)2
==> react the Ba(MnO4)2 with the sulfate of the element X, to obtain BaSO4 (ppt.) and X(MnO4)y

But maybe the Mn-electrode can directly electrolytically be dissolved in NaOH-solution, and also yield the NaMnO4, without the way over Ba(MnO4)2 ?

By the way: The melt, in which I made the manganate, was initially thin flowing (at 500 [celsius]), and became stickkyer, probably as the Na2CO3 dissolved in the NaNO3. Also, at the 800 (+/- 10) [Celsius] it was foaming, but I dont know if from CO2 OR O2 -development.

[Edited on 22-7-2008 by chief]

[Edited on 22-7-2008 by chief]

chief - 22-7-2008 at 00:48

I just wonder, how to reduce _harmlessly_ the MnOx to Mn -powder.
From there it could be reacted with acid to Mn-Salt, and that could be electrolytically deposited to an electrode, which then could be dissolved in the X(OH)y-solution to get the X(MnO4)y

[Edited on 22-7-2008 by chief]

blogfast25 - 22-7-2008 at 05:04

There is no real "harmless" way of reducing manganese oxides: all have fairly high heats of formation and thus require fairly high heat to reduce them to the metal. Reductions with carbon, hydrogen, aluminium, lithium and I believe potassium have all been carried out in the past. Reduction of the dichloride (anh.) with Mg is also possible although thermocalcs show this wouldn't produce molten metal, rather powdered Mn in an MgCl2 slag pile.

I've made small quantities of manganese (lump, not powdered) metal by reducing various oxides (MnO, MnO2, Mn2O3) by means of Al powder but an easy process this ain't. If I get a bit better at it, I'll post a thread on it since as I'm working on it as I'm typing this.

Electrodepositing of the metal onto an electrode may be possible from aqueous solution of an Mn salt but I've no procedure for it.

Crude MnCl2 can be obtained by treating battery crud with strong HCl (caution: generates significant amounts of chlorine gas!) and separating the soluble MnCl2 from insolubles like graphite powder by filtration. DerAlte has a procedure to obtain quite pure MnCl2 by the same method but with some pre-purification of the battery crud. I believe it can be found higher up on this thread. It works very well, I've used it several times. Pure MnO2 can also be obtained from it that way, by re-oxidising the Mn2+ with strong bleach to the dioxide. Quite straightforward, OTC chemicals.

Your best bet for KMnO4 remains fusing relatively pure MnO2 with KOH in the presence of KClO3, leaching the fusion with cold water to extract K2MnO4 and carefully acidifying it.

High grade MnO2 is also available from several eBay sellers.

[Edited on 22-7-2008 by blogfast25]

DerAlte - 22-7-2008 at 19:42

As you’ve probably gathered, I am a manganese chemistry nut. (All transition metals, too). Manganese differs from most transition elements in producing no complexes (AFAIK), although it does form numerous organometallic compounds. It seems to an important trace element in the human diet, though what for I don’t know.

I agree with all the comments of Blogfast25. Years ago I did a whole series of thermite reactions when I was at college. (About 50 yrs ago). Not all were successful in producting useful product, but those for Fe, Mn, and Cr were (B, Si, and V were not, IIRC). Mn melts at a temperature less than Fe but Cr needs something like 1950C, IIRC.

If you try the manganese reaction, never use MnO2. The reaction is almost explosive. My first try was an excellent display, pyrotechnically, but not until I used Mn2O3 did I get a bit of Mn, an irregular mass about ½ inch diameter.. Mn3O4 might be better.

An electrolytic method is outlined in Brauer, q.v. Mn is pretty electropositive and consequently hard to produce electrolytically. I haven’t tried it but have it on my list of things I’d like to try (someday!).

Manganese is remarkable for its avidity for oxygen (or fluorine, like everything)..
Mn(II) salts easily hydrolyze in neutral or basic conditions, even in air, to mixed oxides/hydroxides; while all Mn(III) salts are powerful oxidizing agents and hard to produce. To keep any Mn(II) salts without degradation needs a little acid. AFAIK all are hygroscopic. The carbonate is insoluble and a good way to have Mn available, or the oxalate, Mn2O3 or MnO2..

MnO is difficult to prepare and oxidizes readily in air to Mn2O3. It is best made by heating the (insoluble) Mn(II) oxalate in a closed tube to exclude air:

MnC2O4 + heat (c. 350C) --> MnO + CO2 + CO (but let the gases escape).. I have made it this way and it is greenish.
Also you can make it by heating the carbonate, also excluding air:
MnCO3 + heat (c. 300C) --> MnO + CO2.
The MnO I produced this way seemed to be grey and possibly less pure than the oxalate method.

Another way quoted to produce the dioxide is to heat the nitrate. Heating the dioxide to > ~550C gradually produces the Mn(III) oxide: 4MnO2 --> 2 Mn2O3 + O2; but the (II,III) oxide is only produced at temperatures above about 1100C : 6Mn2O3 --> 4Mn3O4 + O2.

As for the other oxide, Mn2O7, don’t even think about it! A green liquid that explodes at about 95C.

Fairly pure MnO2 might be gotten from an unused alkaline battery, powered and well washed to remove KOH. But up to 50% by volume is still that nasty carbon, difficult to separate mechanically.

@Chief

How can you know the temperature is 800+-10C? Do you have a well controlled and metered furnace or a pyrometer?

I have searched for the decomposition temperature of Na2CO3 without success. I did find a footnote somewhere that suggested it began to decompose somewhere around 900C.

When it does, then the mode would be Na2CO3 --> Na2O + CO2.

Basic oxides often act like hydroxides – they are basic anhydrides; Mn2O3 if anything is slightly acidic.

Then we might have
2Na2O + Mn2O3 + ‘3O’ --> 2Na2MnO4 – but I doubt it at 800C.. NaNO3, Na2MnO4 and MnO2 all decompose at far lower temperatures.

Regards,

Der Alte

not_important - 22-7-2008 at 21:40

Quote:

Originally posted by DerAlte
...
Basic oxides often act like hydroxides – they are basic anhydrides; Mn2O3 if anything is slightly acidic.

Then we might have
2Na2O + Mn2O3 + ‘3O’ --> 2Na2MnO4 – but I doubt it at 800C.. NaNO3, Na2MnO4 and MnO2 all decompose at far lower temperatures.

Consider

2Na2CO3 + Mn2O3 + 3O => 2CO2 + 2Na2MnO4

Or

Na2CO3 + Mn2O3 + O => CO2 + Na2Mn2O5

Na2CO3 + Na2Mn2O5 + O2 => CO2 + 2Na2MnO4

Manganese forms complex cyanides similar to ferrocyanides, but not as stable. Complexes with thiocyanate can be formed as well.

anhydrous Mn(2+) halides form complexes with dry NH3, up to MnCl2.6NH3

Mn(3+) forms a number of complexes

oxalate 5 H2C2O4 + KMnO4 + K2CO3 => K3[Mn(C2O4)3] + 5H2O + 5CO2

or treat freshly ppt damp MnO2 with KHC2O4 at zero C, the adding alcohol to precipitate the complex as red-violet crystals of the trihydrate. Do these reactions in fairly dim red light.

Malonic acid forms similar complex, a bit more stable and of a greenish hue when solid.

blogfast25 - 23-7-2008 at 03:12

Quote:

Originally posted by DerAlte

I agree with all the comments of Blogfast25. Years ago I did a whole series of thermite reactions when I was at college. (About 50 yrs ago). Not all were successful in producting useful product, but those for Fe, Mn, and Cr were (B, Si, and V were not, IIRC). Mn melts at a temperature less than Fe but Cr needs something like 1950C, IIRC.

If you try the manganese reaction, never use MnO2. The reaction is almost explosive. My first try was an excellent display, pyrotechnically, but not until I used Mn2O3 did I get a bit of Mn, an irregular mass about ½ inch diameter.. Mn3O4 might be better.

Maybe we should get together: I'm a thermite nut (at least for the moment) and manganese is the one thermite that's hard to control in terms of actually obtaining the metal itself and not just pyrotechnics. I've produced w/o many problems Si, Fe, Cr, V, Ti, Co, Cu and often alloys of these as well, I would easily make Nb and Sc if I could get the oxides, but Mn remained the elusive one, at least up to recently.

I've been working on a workable solution for weeks using different oxides like MnO and Mn2O3 and combinations thereof, gradually making progress. Most of my Mn thermites now produce metal of good quality (visually speaking) but the yields remain frustratingly small (typically less than 30 % theor. yield). Yesterday I had a bit of a breakthrough with a 50 g reaction based on an equimolar blend of MnO and Mn2O3. Although yield was again poor (37 % theor.) there was at least one 7.3 g clean lump of metal, oblong, about 1 inch long. Today I'll be running another variant and shortly will open a thread on my experiments on making metallic manganese, using thermite, as well as ideas on the use of other reductants.

chief - 23-7-2008 at 03:15

"How can you know the temperature is 800+-10C? Do you have a well controlled and metered furnace or a pyrometer?"
-----
The furnace is conrolled via an old PC; PC switches the furnace on and off, depending on the temperature measured. Since the temperature oscillates a bit, from the switching, it is +/- 10 [Celsius](max), as I see. The temperature-probe was tested eg. at the melting points of Al, Ba(NO3)2 and others, and is quite exact: +/- 2-3 [Celsius].
The only inaccuracy comes from the temperature-distribution within the furnace: It's hotter on the ground, where the experimental substances rest (within stainless steel), than at the top, where the temperature-probe is.

I just can set the PC to hold any temperature, and it will do so, until the 9-V-Battery of the probe (RS232-connected) goes out of energy.
-------------------

I tried now to reduce the MnOx by glowing with charcoal (only raw pieces of C were used, so explosion was impossible; the stuff was mixed in an kitchen-engine to some sort of dough, and that was put wet into the furnace). That was glowed at 800 [Celsius] for a while, and then mixed into water (lot of charcoal was still present); HNO3 was added, but no really big H2-generation happened. So it didn't reduce the MnOx.
---------------------
Also I melted MnOx wit Na2CO3 at 880 [Celsius] for a while, looks unreacted, but I didnt test it yet.

[Edited on 23-7-2008 by chief]

blogfast25 - 23-7-2008 at 10:56

The thread I intended to start here on Mn2O3/Mn thermites has been spliced onto this one here.

DerAlte - 23-7-2008 at 17:42

Interesting, not_important. Perhaps I should have said that for all practical purposes (FAPP) manganese produces no stable complexes, unlike Co, Fe, Ni, Cu, etc., etc.. See Brauer for various cyanide and oxalate Mn complexes – nearly all unstable in either air or in water solution or at room temperature.

Now what the hell is Na2MnO5?
An ionic, assumedly… MnO4 can exist with Mn(V), (VI) or (VII) ie {Mn(+n)(4O(2-))} with charge n-8, n=5,6,7. Even K4MnO4 with Mn(IV) but I’ve yet to see a reference of an authoritative nature. Manganites, such as Na2MnO3, are said to exist, The Mn would there be in oxidation state (IV). Anyone got access to Mellor on Mn compounds?

The hypothetical MnO5 moiety , charge -2, requires n-10 = -2, i.e. n=8 – a very unusual valence state, although OsO4 is known. Are you pulling my leg, not_important?

Quote:

Mn(3+) forms a number of complexes
oxalate 5H2C2O4 + KMnO4 + K2CO3 => K3[Mn(C2O4)3] + 5H2O + 5CO2
or treat freshly ppt damp MnO2 with KHC2O4 at zero C, the adding alcohol to precipitate the complex as red-violet crystals of the trihydrate. Do these reactions in fairly dim red light.

That’s interesting. Reminiscent of the cobalti- complexes, Co(III) normally being unstable yet the complexes are very stable. In contrast cobalto- (Co(II)) complexes are not stable, but reducing agents, IIRC. Also, I believe I saw somewhere that Mn(III) salts can make fairly stable double salts of the alum type. Compare Fe+++ in ferric ammonium sulphate, a nice blue alum, quite stable.

Be that as it may, I cannot see myself trying to produce the above oxalate complex in freezing weather in a dim red light, as the climate I live in never gets to freezing (not recently) and I am a bit blind in dim red light!

@ chief-

You are lucky, sounds a nice set-up. I believe your 800C.

WRT production of manganese by reduction with charcoal, it can be done in a blast furnace – at c. 1500C in a CO atmosphere (standard steel making practice). But just heating C with MnOx will produce Mn3O4 from any dioxide, Mn2O3 mixture.

@blogfast25: It’s 50 yrs+ since I did a thermite so I’m relying on a fading memory plus a few marginal notes in an old chemistry text. However, I shall read the thread with interest….

A thought. Can you get or make the Mn3O4 oxide? It ought to be a bit slower than the rest, and has more Mn. Mn melts easily enough compared with Cr, V or Ti. It's a rather reactive metal, displacing H2 from hot water.

Edit - added: re Mn in the diet, see Wiki, subject superoxide dismutase if biochemically inclined...

Regards,
Der Alte

[Edited on 23-7-2008 by DerAlte]

blogfast25 - 24-7-2008 at 04:20

Quote:

@blogfast25: It’s 50 yrs+ since I did a thermite so I’m relying on a fading memory plus a few marginal notes in an old chemistry text. However, I shall read the thread with interest….

A thought. Can you get or make the Mn3O4 oxide? It ought to be a bit slower than the rest, and has more Mn. Mn melts easily enough compared with Cr, V or Ti. It's a rather reactive metal, displacing H2 from hot water.

Although I don't have the figures in front of me, off the top of my head I'd say that per mol of oxide Mn3O4 will generate possibly the most energy of all Mn thermites, because of its high oxygen content. But it does have the saving grace of producing the most reaction products of all (per mol, 3 mol Mn and 4/3 mol Al2O3) and that cools things down quite a bit, as the reaction products act as heat sponges.

Hausmannite is (or was?) used in industrial thermites in open magnesite lined crucibles. After initiating a charge, regularly aliquots of a mix of Mn3O4, Al and CaO are then added to keep the fire going and gradually fill the reactor with molten Mn (see Chemical Metallurgy: Principles and Practice, scroll down to page 391).

Couldn't Mn3O4 be prepared the same way as magnetite (Fe3O4), by co-precipitating equimolar amounts of Mn2+ and Mn3+? Or is the latter not stable enough in aqueous solution?

[Edited on 24-7-2008 by blogfast25]

not_important - 24-7-2008 at 07:10

Quote:

Originally posted by blogfast25
Couldn't Mn3O4 be prepared the same way as magnetite (Fe3O4), by co-precipitating equimolar amounts of Mn2+ and Mn3+? Or is the latter not stable enough in aqueous solution?

Mn(3+) is not stable in aqueous solution as simple compounds, only as complexes and even those tend to convert to Mn(2+) and Mn(4+) which precipitates out. Easier to do is mixing a solution of a Mn(2+) salt with one of KMnO4 in the proper ratio.

Mn3O4 is the most stable oxide of manganese when heated in air. Heat the metal, any of its oxides, hydroxides, carbonate, or a great many of the other salts in air above about 950 C and you eventually get Mn3O4. A standard prep was heating a higher oxide in air for 6 hours at 1000 C.

Quote:

Originally posted by DerAlte

Now what the hell is Na2MnO5?

You misread - Na2Mn2O5, or in old style mineralogy Na2O,2MnO2.

blogfast25 - 24-7-2008 at 09:20

Quote:

Easier to do is mixing a solution of a Mn(2+) salt with one of KMnO4 in the proper ratio.

Not_important, how would that work?

DerAlte - 24-7-2008 at 18:05

@not_important : duly noted, mea culpa.

@blogfast25
I am replying in this thread even if somewhat OT, or we’ll get cross-threaded and all screwed up, if we aren’t already!

Interesting that you are doing a whole set of thermites, just like I did long ago! I just remembered that in some of them (don’t know which) I added a layer of sand to the top for some reason; I ignited them with Mg ribbon and some mild Mg mixture in a small heap.

I must have used Mn2O3 because I can’t think of any reason for having Mn3O4; I must have made it by igniting MnO2, which was the likely chemical to have. As to the ‘thought’ of using Mn3O4, did a quick calculation of enthalpies of each reaction (using figures from CRC):

(1) MnO + 2/3 Al -- > Mn + 1/3 Al2O3 dH = -220.4 kJ/mol Mn produced
(2) MnO2 + 4/3 Al -- > Mn + 2/3 Al2O3 dH = -597.2 kJ/mol Mn
(3) ½ Mn2O3 + Al -- > Mn + ½ Al2O3 dH = -358.4 kJ/mol Mn
(4) 1/3 Mn3O4 + 8/9 Al -- > Mn + 4/9 Al2O3 dH = -282.2 kJ/mol Mn

So, if I haven’t made a mistake, MnO is the best bet, MnO2 the worst as too energetic, and we have found it to be highly pyrotechnic.

The figures are for Standard enthalpies, i.e. at 25C, solid state and 1 atm pressure. We need some extra heat to supply the latent heat of fusion of Mn, and any slag (Al2O3 melts at 2054C, I notice) plus the heat needed to bring reactants and products to the peak temperature, > melting point of Mn (around 1250C IIRC). Per unit Mn, MnO is also best in terms of Al used. You can make the MnO by igniting Mn carbonate in a closed crucible.

No idea how to make Mn3O4 other than roasting MnO2 or Mn2O3 to near white heat, as not_important says.

That process for making Fe3O4 sounds interesting. Does it merely produce a mixture or true magnetite? Could be tested magnetically.

Did not realize Mn had such a low BP.

Regards,
Der Alte

Formatik - 25-7-2008 at 05:09

I haven't seen it said here yet, but there is a nice simple procedure for KMnO4 from MnO2, KOH and KClO3 as the oxidant illustrated at versuchschemie:

http://www.versuchschemie.de/topic,10934,-Synthese+von+Kaliumpermanganat+aus+Braunstein.html

Some know already, but using MnO2 from batteries directly which also contain carbon should not be used because explosions have resulted by heating KClO3 with MnO2 containing carbon.

blogfast25 - 25-7-2008 at 05:47

@DerAlte

Yes, we're probably going to get shafted for being OT, so I'll keep it short.

Your reaction enthalpy values are correct. Interesting how you express them per mol of Mn produced, not per mol of oxide reduced.

MnO is by far the coolest running but it needs help to get to a molten slag/metal mix from one of the higher oxides. Yield is always quite low because the MP of alumina and BP of manganese being so close together.

I've made my MnO by decomposing MnCO3 at 400 - 500 C under a stream of CO2 (I had to build a CO2 generating "apparatus" for that - very a la Frankenstein movies - lol). The Mn2O3 is made be decomposing MnCO3 in the same conditions but in open air. Neither products are particularly pure but in thermite conditions they do exactly as "it says on the tin".

Provided I can get my hands on Mg powder, I'll now be gunning for the MnCl2 + Mg ---> Mn + MgCl2 reduction (ΔH = - 161 kJ/mol of MnCl2).

++++++++++++

The synthetic Fe3O4 is magnetite alright. It's the basis for homemade ferrofluids. Here's a bit of info and some links on ferrofluids. Mine worked well but not well enough to get the famous "hedgehog effect". The fluid was strongly magnetic though.

chief - 12-8-2008 at 08:57

I now continue to try a good NaMnO4-synthesis.

Today I dissolved MnOx in Sodium-MetaBisulfite-solution with HNO3 (quite _dilute_ solution of both ingredients), at the boiling point of H2O.

==> The MnOx from fresh precipitation dissolves quickly,
==> and also does the MnOx from batteries (glowed at 750 [Celsius] before), leaving only the black residual graphite-flakes.

It gives a discolorized solution, within which the Mn somehow is bound to the other elements.
And here the question is: How to make the best use of that way of dissolving the MnOx ?
Can somehow a Mn-salt be concentrated, for electrolytical Mn-plating (later to be electrolytically dissolved in a hydroxide-solution to get the XMnO4, as is said to funtoin at least for Ba(MnO4)2) ?

Or what could I do ?

The sodium-metabisulfite is sold and extensively used for making wine and food-conserves, since it gives SO2 on reaction with the acids in the food, thereby exterminating the bacteriae and conserving the food.

chief - 15-8-2008 at 02:29

I have to correct myself: The MnOx does not dissolve, but give a white (pale-yellow-brownish, but basically white) compound, that settles to the ground.
The rawer MnOx-parts stay undissolved (within timeframe of minutes).

The MnOx reracted with
==> Na2S2O5 (http://en.wikipedia.org/wiki/Sodium_metabisulfite)
==> HNO3

chief - 30-8-2008 at 00:34

The dissolution of MnOx sometimes occures, and at other times it doesn't:
KMnO4-crystals completely dissolve, nothing stays undissolved, and the severaln stages of reduction can be observed:
When shaking in a flask:
==> near the KMnO4 it's violet
==> 1/2 cm distance its brown and looks like MnO2
==> a little further and it looks whitish
==> then its dissolved.

So, obviously, also the freshly ppt. MnO2 (?) dissolves.

On the other hand I have tried several days old MnOx inthe automatic stirrer for hours, and it does not dissolve, at least not completely (but it looks like "not at all").

Anyhow, as I found suggested elsewhere, the Na2S2O5-discolorization may be used for titration of the MnO4, so conc.-determining is easily possible this way !!!

That it sometimes dissolves and sometimes not: Probably because of the modifications of the stuff, which have different chemical properties. This seems to be also an issue in the battery-making-busines.

chief - 30-8-2008 at 00:51

NOw, as second post, because it may be useful:
I found 2 patents of making the permanganate via oxidizing MnOx-solution with chlorine:

http://rapidshare.com/files/141248016/XMnO4-1.pdf.html
http://rapidshare.com/files/141248017/XMnO4-2.pdf.html

This way yields 50 % and is easy. I tried to adapt it and make the chlorine via electrolysis within the solution: The MnOx into water, added Na2CO3 and NaCl, waiting for the green or violet color. But until now the MnOx stayed unaffected. I let it run, maybe in the later perchlorate-stage it can react ...

If it doesn't I also (maybe) will add NaNO3 to the solution, since I found patents of applying NO2 to MnOx-ore, to make it chemically dissolvable ... .
But care must be taken: Cl and N form a _very_ explosive and dangerous compound (NCl3 ?), that usually forms when electrolyzing NH4Cl !! Therefore the NH4Cl first must be washed out of any MnOx used from Batteries ! The NCl3 usually is yellowish-green (I never saw it personally) and settles on the ground.

Maybe this NCl3 might also form when electrolyzing NaCl and NaNO3 in the same bath ? I will think about that first, before adding the NaNO3.

DerAlte - 31-8-2008 at 22:45

@ chief

I have not been keeping up with the forum and have just seen your latest posts. I notice the patents are in German - not my best language. But they look interesting - I will read them and comment later.

Regards,

Der Alte

DerAlte - 5-9-2008 at 13:37

The two German patents above give methods for manufacturing NaMnO4 from MnO2 using NaOH and Cl2 gas. The most interesting is the first. This is a synopsis of the useful parts:

The introduction points out that permanganate can be made by fusing together MnO2, an alkali and an oxidant, the traditional method, and also by use of hypochlorite but the yield is poor (as I have found above).

It gives two examples: I paraphrase

EXAMPLE 1: 87 parts of hydrated manganese dioxide were suspended in 533 parts of 30% sodium hydroxide solution . The liquid was heated to 110-112 degrees and 106.5 parts chlorine gas passed into it over a 1 hour period. The yield of NaMnO4 was 50% of theoretical.

{In ½ hour the yield was 35% and in ¼ hour, 25%}

At room temperature under the same conditions the yield was very low. ; at 50 C the yield was 6.3%

For 10% NaOH, and temp 110C, yield = 10%, 35% NaOH 17% (?)

EXAMPLE 2: As in Ex 1, except 3 parts CuSO4.5H2O were added. This yielded 106.5 parts NaMnO4, 75% of theoretical. Using 4 parts ZnSO4.&H2O, same conditions, gave 92 parts, 65% theoretical.

………

This method sounds quite practical for the amateur prepared to produce and able to handle large quantities of Cl2 gas. No details were given of how easily the gas is absorbed or a suitable apparatus, except for mentioning that stirring was employed. In essence it appears to be equivalent to producing hypochlorite in situ. Whether this avoids the production of much chlorate is not said; but if the yields of permanganate are high enough, small amounts of chlorate will pose no problem when the salt is converted to the potassium salt. The catalytic action of the metal ion (Cu++ or Zn++) is especially interesting.

I am having a look at the probable reactions involved.

The chief problem would be the safe production of a steady stream of chlorine and making sure it is well absorbed.

Der Alte

Taoiseach - 16-9-2008 at 22:31

Nope it doesn't give the preparation from MnO2 but from the hydrate. Which is not the stuff you can buy in pottery stores.

The Cu/Zn salts seem to catalyze the reaction because they easily form manganites.

chief - 22-9-2008 at 06:03

2 nice patents: Using alkali-orthoplumbates, oxidizing with air:
http://rapidshare.de/files/40528249/MnO4viaPbO4-1.pdf.html
http://rapidshare.de/files/40528250/MnO4viaPbO4-2.pdf.html

The patents use the procedure to generate alternatingly oxygen or nitrogen from air:
==> oxidizing with air absorbs quantitatively the O, only N passes ("regeneration step")::
Na4PbO4 + MnO2 + O(air) ==> Na2MnO4 + Na2PbO3

==> blowing H2O-steam into it: gives only O
Na2MnO4 + Na2PbO3 + H2O ==> Na4PbO4 + MnO2+ H2O + O

But probably useful to adapt it to get easy oxidation via air, no extra oxidizers needed !

[Edited on 22-9-2008 by chief]

chief - 23-9-2008 at 03:38

Here now about the NaMnO4-making from
==> (NaNO3 OR NaOH) + Mn-oxides:
http://www.retrobibliothek.de/retrobib/seite.html?id=116000#...
(source unavailable at some times, but during european week-working-times it mostly functions)

There it's stated:
==> Heating to 400 [Celsius] the Mn-oxides with Chile-saltpeter or NaOH,in presence of air
==> "NaMnO4 is made same way like KMnO4" (!, as above someone said this couldn't be done)

chief - 10-10-2008 at 10:06

Success: Na2MnO4
==> from NaClO3 (not even clean,since difficult to crystallize without NaCl)
==> and NaOH
==> at 400 [Cels], 10-20 minutes

10 g NaOH + 10 g MnO2 (100% MnO2, pottery-grade) + 10 g of the not very clean NaClO3 gave a _very_ green solution, as green as known from dissolving KMnO4-thermal-decomp-educts; definately has some concentration !

What didn't work: KNO3 and NaNO3 just dont give the slightest green (KNO3 was tried without any hydroxide).
Also: NaClO3 + Soda (instead of NaOH) failed, at the 400 [Cels]; didn't try at higher temp.

So:One doesn't even have to crystallize the NaClO3 too extensively, just a raw crop from the electrolysis-cell does it !

[Edited on 10-10-2008 by chief]

Picric-A - 10-10-2008 at 10:43

How long was the mix heated and did you add the NaClO3 seperatly or just chuck all the reagents in and begin heating?
Nice work btw! Have you tried acidification yet?

chief - 10-10-2008 at 11:44

It was only in the furnace for 20 minutes, incl. heating-up-time;
the furnace was controlled to have 400 [Cels]; maybe the mixture needed up to 10 mins to reach this, and maybe part of it (bottom) got hotter by maybe 20 or 30 degrees, so local max-temp would have been 430 [Cels].

It was prepared in a old coffee-mill (10000 RPM) (which will never see any food again), everything was fine-powdered and well mixed; the MnO2 was from a pottery-supply, stated as "100 % MnO2"; maybe I'll try the battery-stuff too ... (thoroughly washed, filtrated, pre-glowed with some nitrate to remove the graphite first !)

[Edited on 10-10-2008 by chief]

Ah yes, the acidification: I preferred to ppt. out with Ba(NO3)2, to get the insoluble BaMnO4 ... to be continued ...

[Edited on 10-10-2008 by chief]

[Edited on 10-10-2008 by chief]

chief - 12-10-2008 at 07:12

... and the ppt, upon acidification, gives a violet-color; but it always is a slightly other violet (with Ba) than eg NaMnO4 OR KMnO4, somewhat reddish component when dilute.
Only: When acidification was with HNO3, at boiling-down the HNO3 gets too concentrated, and it all gets MnOx-ppt..
The other thing I tried was: NaMnO4 + MgSO4 (solution), as suggestedin a link from above (http://www.retrobibliothek.de/retrobib/seite.html?id=116000#...), gives also permanganate,but upon filtering the violet solution through the same filter as the green manganate-solution before, it turns into green again, perhaps some reaction with the MnO2 that was still in the filter (and that I left there to "better" the filtering-effect ...)

Also: Did the Na2MnO4 via chlorate, as above, several times now: It works. It doesn't harm, also, to set the furnace to 470 [Celsius]. Upon boiling down the green solution in the microwave the glass cracked quite cleanly around a ring where some already dry manganate was solidified on the wall (susceptibility of MnOx maybe high) .. maybe a way of cutting glass/bottles etc: Manganate around the bottle, somehow painted onto it ==> microwave for a few minutes ...

Now I think of how to mildly acidify it with CO2, don't want to abuse any acid for that purpose, since until now it's also an acid-free process to the manganate: Only MnO2 + NaOH + chlorate ...

Didn't find anything upon an easy-to-make CO2-generator at the forum here ...

[Edited on 12-10-2008 by chief]

[Edited on 12-10-2008 by chief]

Silverado7 - 27-10-2008 at 13:43

Quick question about potassium permanganate: If sulfur dioxide is dissolved in dilute potassium permanganate solution, the violet color disappears. Does this mean that the potassium permenganate present oxidized the sulfur dioxide into sulfur trioxide? Does the clear color mean sulfuric acid? Or did the sulfur trioxide combine with something somewhere. Like with the manganate, to permanganate. Not quite sure. HELP!?

Picric-A - 27-10-2008 at 13:45

It means the SO2 is reducing the Permanganate... you would definitly not get SO3 just left in soloution!!!
SO3 would form H2SO4 as soon as its made practically...

chief - 27-10-2008 at 15:12

Quote:

Originally posted by Picric-A
It means the SO2 is reducing the Permanganate... you would definitly not get SO3 just left in soloution!!!
SO3 would form H2SO4 as soon as its made practically...

But the Mn is still there! So either MnSO4 OR MnSO3 will be formed ? Maybe it's at least a way to oxidize SO2 to SO3 in Salt; later decompose the MnSO4 thermically, re-use the MnOx with another batch of cheap chlorate, and there is the H2SO4-engine !! Could this work that way ?

[Edited on 27-10-2008 by chief]

Picric-A - 27-10-2008 at 16:33

hmm, i said this in another post, i thought what could happen is the SO2 would be oxidised to SO3 but reduce the permanganate to manganate. the SO3 would form H2SO4 and this would oxidise the manganate back to permanaganate... where would it end? would the SO2 reduce the KMnO4 back to MnOx/MnSO4 and K2SO4?

chief - 28-10-2008 at 00:53

With Na2S2O5, a SO2-source (cheap available), the reaction goes on like this, as I stated earlier (previous page) and observed myself:

Quote:

Originally posted by chief
The dissolution of MnOx sometimes occures, and at other times it doesn't:
KMnO4-crystals completely dissolve, nothing stays undissolved, and the severaln stages of reduction can be observed:
When shaking in a flask:
==> near the KMnO4 it's violet
==> 1/2 cm distance its brown and looks like MnO2
==> a little further and it looks whitish
==> then its dissolved.

So the Mn is reduced from the +7 to the +2 stage, giving probably either the MnSO3 OR MnSO4. Which one ?

[Edited on 28-10-2008 by chief]

[Edited on 28-10-2008 by chief]

12AX7 - 28-10-2008 at 00:58

Quote:

Originally posted by chief
So the Mn is reduced from the +7 to the -2

Umm?

Quote:

stage, giving probably either the Mg

Umm?

Quote:

SO3 OR MgSO4. Which one ?

Manganese sulfate, if you didn't use too much excess sulfite. If you did, there may be some sulfite present. But these salts are ONLY formed if you evaporate the solution and crystallize it. Ions in solution are not compounds, they are ions in solution.

Tim

chief - 28-10-2008 at 01:18

@12AX7 MnSO3/4 (!), not MgSO3/4, of course ! Thanks for the correction, I edited it !

The reduction from +7 to +2:
KMnO4 : Mn ==> is +7 (K ==> +1 ; O4 ==> -8 ; ==> Mn ==> +7)
MnSO4: ah, yes : +2, not -2, I'm gonna edit it again.

Maybe I was with the Sulfur, which goes from neg. to pos: (H2S,S,SO2,SO3)

[Edited on 28-10-2008 by chief]

chief - 13-3-2009 at 10:38

Question: The electrolysis from manganate to permanganate:
==> Can it be done with PbO2-Anodes instead of Nickel ?
==> What about graphite ?

Besides: I reported above the manganate-formation at 800 [Cels] with Na2CO3 instead of NaOH ; I then found, that this is due to the decomposition of the Na2CO3 ... to NaO or someting . I think it's like that because the 780-800 [Cels], which are also the decomp-temp for the Na2CO3, were necessary for anything to get green: At lower temps no amount of time would do it (up to several hours; maybe a grinder/mixer can do after 48 hours ..., but I wouldn't bet on it) ; maybe just a mistake in the literature ...

I tried it with clean NaOH, and it works _well_ , _deep_ green solution ...

Formatik - 6-7-2009 at 12:08

Quote: Originally posted by Formatik

I've also tried it with the strongly hygroscopic conc. HClO4, where adding powdered KMnO4 didn't form Mn2O7 initially, but just a red mixture that discolors on contact with tissue. Though on standing longer the mixture does form small, shiny oily green droplets which have the reaction of Mn2O7.

Though as far as oxidation of HCl by KMnO4 goes, this is said by at least several references in Brethericks to have given a sharp explosion when conc. HCl was added to solid KMnO4, the explanation was that there is a remote possibility of the permanganate oxidizing the Cl2 to "chlorine oxide". I've added 0.4g conc. HCl to 1g KMnO4 dust in a clean, dry test tube, after some gas evolution a lit long match was inserted two times but nothing happened. Addition of liquid Cl2 to (moist or dry) KMnO4 would surely prove the hypothesis wether Cl2 is actually oxidized.

[Edited on 6-7-2009 by Formatik]

monoceros4 - 1-9-2009 at 07:39

I should have looked over the forums more thoroughly; I didn't know there was a dedicated thread on permanganates. I've wasted a lot of time recently on various attempts at manganates and permanganates. Complicating matters was that KMnO4 or K2MnO4 were out of consideration because I wanted the substance as an aid in detection of potassium as KClO4 with the microscope. I've been trying to keep notes of my progress using Google Docs; here's the document pertaining to all the manganate/permanganate attempts:

http://docs.google.com/Doc?docid=0AWsze3MPrfu6ZGh0eDUzZ3dfMj...

Quick summary: I succeeded in preparing some impure BaMnO4, but failed at converting it to Ba(MnO4)2 with CO2. Fusions of MnO2 with NaOH and added NaClO3, followed by extraction of the melt with dilute NaOH, gave Na2MnO4 solutions but I mostly succeeded only in destroying them when trying to convert them to some kind of soluble, crystallizable salt. The Na2MnO4 solution from the last attempt was divided, one half acidified to yield permanganate, and both solutions promptly frozen until I can figure out how to process them further.

I did succeed in preparing one well-crystallized salt, mixed crystals of Na2SO4.10H2O and Na2MnO4.10H2O in undetermined proportion, unstable in open air but stable in a humidor. It's enough for my purposes--all I need is a small, unmeasured amount of a manganate or permanganate for treating drops on a microscope slide--but not much use for large-scale preparation.

A practical note that doesn't come across well in the notes (I need to improve my note-taking): I cracked a lot of cheap porcelain crucibles doing this. Thermal stress was no doubt sometimes to blame: my sources of heat are poor, comprising only an alcohol lamp (wretched), a couple of different propane torches (one the standard pencil torch, another with a fairly broad flame), and an electric range. But even using the electric range and gradual heating to the maximum value, the porcelain cracked, sometimes extensively. I know that alkaline melts attack porcelain but I'd underestimated how much it would weaken it.

I'll have to read up on this thread. I probably could have avoided some rookie mistakes had I known it existed.

JohnWW - 1-9-2009 at 07:54

Quote: Originally posted by monoceros4

(cut)
A practical note that doesn't come across well in the notes (I need to improve my note-taking): I cracked a lot of cheap porcelain crucibles doing this. Thermal stress was no doubt sometimes to blame: my sources of heat are poor, comprising only an alcohol lamp (wretched), a couple of different propane torches (one the standard pencil torch, another with a fairly broad flame), and an electric range. But even using the electric range and gradual heating to the maximum value, the porcelain cracked, sometimes extensively. I know that alkaline melts attack porcelain but I'd underestimated how much it would weaken it.(cut)

More resistant materials than porcelain (or borosilicate glass) should be used as crucibles for reactions involving fused or hot concentrated alkalis. The reaction mixture would become contaminated with silica and alumina. The first step upwards would be fused zirconia, or failing that, graphite or nickel; or better still, platinum if it could be afforded.

monoceros4 - 1-9-2009 at 08:11

Quote: Originally posted by JohnWW

I suppose I could try to find a used platinum crucible but then I'd have to go without paying rent for at least a couple months. Vitreous carbon is somewhat cheaper. I have a small nickel crucible but was hoping to manipulate larger amounts of material. Would magnesia hold up? It's reasonably priced.

I didn't worry about silica or alumina contamination, partly because it's highly unlikely that any significant Si or Al would find its way into a manganate or permanganate crystallization, partly because nobody expects permanganates to be all that pure anyway.

If I mess with this again I hope to avoid methods involving fusion. Even when it works out it's so damn tedious.

[Edited on 1-9-2009 by monoceros4]

monoceros4 - 2-9-2009 at 22:17

I have made some progress in an unexpected direction and prepared BaMnO4 from MnO2 without needing to melt or electrolyze anything.

User Paddywhacker in another thread (that I shouldn't have created) suggested to me that a good procedure for making barium manganate, BaMnO4, would be useful because of the compound's utility as a mild oxidant. I was interested because BaMnO4 is a route to barium permanganate, thence to other permanganates. I tried a fusion method the first time and got poor results.

I think I've done better this time and with a very simple method requiring no special equipment. Finding references to "hydrothermal" methods of preparing potassium manganate, i.e. operating in aqueous solution at high temperature and pressure, I reasoned that BaMnO4 ought to be even easier to prepare that way because the insolubility of the compound would help the reaction along. To test this I heated 2.0 g of MnO2, 9.0 g of impure, homecooked Ba(OH)2.8H2O, and 3.0 g of KNO3 (I wasn't sure of the necessity of this--maybe air is enough?) moistened with a little water in a loosely closed container in an ordinary pressure cooker and lots of water at 15 psi overpressure. The result was a greyish-green, somewhat heterogeneous mass that, after exhaustive washing with cold water, drying, and grinding, yielded 6.1 g of olive-green powder. The weight is not far off theoretical but the substance is definitely not pure. While I saw only a few black particles, indicating that most of the MnO2 had reacted, there was more whitish matter that was probably BaCO3 from the impure Ba(OH)2. (This was later confirmed: a pinch of the solid fizzed briefly in weak acid.)

I'm in the process of trying to assay the product. My attempt to determine "active oxygen" by boiling a small sample with strong HCl and collecting the Cl2 in dilute NaOH failed; my hastily improvised apparatus was too leaky and a fair amount of chlorine escaped, judging from the smell; also the NaOH I used was not strong enough to deal with all the HCl vapor also coming off. The HCl solution was free from sediment and bright yellow, indicating Fe contamination no doubt from the crude MnO2. Now the plan is to determine Ba, after removal of Mn as MnO2 and Fe as Fe(OH)3, by precipitation of barium chromate and iodometric estimation of Cr(VI) in the BaCrO4. I haven't worked out what method might be best for determining Mn. I'm leaning towards removing the Fe by shaking with BaCO3, removing the Ba by precipitation of BaSO4, then determining Mn as MnNH4PO4.H2O. I wouldn't consider a gravimetric method but I did once succeed in determining Mg by precipitating MgNH4PO4.6H2O--my one stab at a gravimetric procedure--so I might be able to carry it off. If I had any sodium bismuthate I'd consider oxidation to permanganate; I have ammonium persulfate but Vogel states that it gives poor results in quantitative work.

One last note of possible interest: in trying to precipitate MnO2 from a sample, I decided to use NaOCl, which was a mistake. As noted above, NaOCl is capable of at least partial oxidation of Mn to permanganate, and that is what happened. Even prolonged boiling did not destroy all of it probably because there was always plenty of excess NaOCl around. Adding NH4Cl destroyed the excess hypochlorite and decolored the solution.

In any case, I'm fairly optimistic of the results--at least in getting as far as barium manganate. I see no reason why this preparation could not be scaled up considerably, either, which is definitely not true of fusion methods without going to a lot of trouble. The next trick will be to attempt preparing barium permanganate by gassing a suspension of BaMnO4 with carbon dioxide.

It might also be possible to dispense with the barium hydroxide and use a mixture of barium carbonate and sodium hydroxide. That would save me the effort of preparing Ba(OH)2.8H2O first; it's not terribly difficult but it is a tedious process.

not_important - 3-9-2009 at 03:03

By the atached diagram, it doesn't look as if air would do the job under aqueous conditions.

For alkaline fusions, iron or nickel is acceptable. There will be some contamination, especially if chlorates or perchlorates are added as oxidiser - halogens are bad news even under alkaline conditions.

Formatik - 3-9-2009 at 07:23

Muthmann (Ber. 26 (1893) 1016) has a procedure for Ba(MnO4)2 starting from Ba(OH)2, KMnO4 and Ba(NO3)2. The product so made is said to be free of potassium. The same procedure appears in Vanino's book, Handbuch der präparativen Chemie.

monoceros4 - 3-9-2009 at 07:39

Quote: Originally posted by not_important

By the atached diagram, it doesn't look as if air would do the job under aqueous conditions.

Aqueous standard conditions, to be sure. But the hydrothermal reactions are under conditions decidedly nonstandard. Have a look at this abstract:

Quote:

Chemische Technik (Leipzig, Germany) 17(8):493-494 (1965)
The effect of reaction parameters on the formation of potassium manganate(VI) from manganese dioxide and potassium hydroxide
Teske, Klaus; Lehmann, Hans Albert
In the system MnO2-KOH-O2, K3MnO4 alone is only formed at a K/Mn ratio of >3 and temps, of >500°. At lower temps. and K/Mn ratios, some K2MnO4 always formed. K3MnO4 was oxidized at 100-300° for 4 hrs. in O and a PH2O of 23 mm. (PH2O = pressure of H2O vapor). From ∼ zero at 100° the conversion to K2MnO4 was practically complete at 150-200°, and almost completely reversed at 300°. In air, the range of max. conversion was narrowed (175-200°), but increased (175-300°) by raising the PH2O to 107 mm. With increasing PH2O at 250° for 4 hrs., the conversion was almost complete at PH2O of 82 mm., and raised only slightly at PH2O of 150 mm. Addn. of an equimolar amt. of KOH lowered the initiation of the thermal decompn. of K2MnO4 from 500 to 200°, and caused almost complete decompn. at 250°. The addn. of more KOH caused only a slight further lowering of the decompn. temp. In the presence of moist O and KOH, K2MnO4 decompd. at 200-75°. The results may be partially explained by the equil.: 2K3MnO4 + H2O + 1/2O2 2K2MnO4 + 2KOH.

That will give you an idea of the conditions needed to effect oxidation to manganate. The phase diagram above is based off standard electrode potentials: 25 deg. C, 1 atm, 1M solutions. None of these is the case with a reaction at high temperature, increased pressure, and highly concentrated solutions used. Nor is the paper above the only reference to air-oxidation of MnO2 to manganate at high temperatures. A Japanese patent (JP 55085425) describes production of K2MnO4 at 275 deg. C and 3 atm with air the only oxidant.

Bear in mind also that preparing BaMnO4 instead of K2MnO4 means that, even if the equilibrium concentration of manganate is very small, the formation of barium manganate need not be hindered because the substance is highly insoluble.

Taoiseach - 5-9-2009 at 02:43

Very interesting stuff monoceros!

Quote:

there was more whitish matter that was probably BaCO3 from the impure Ba(OH)2. (This was later confirmed: a pinch of the solid fizzed briefly in weak acid.)

You had MnO2, KNO3 and Ba(OH)2 in a tightly closed vessel. I wonder where the CO2 would come from to form the carbonate?

That fizzling must be due to BaCO3 that was already present as impurity in your Ba(OH)2. Also keep in mind that KNO3 should be reduced to KNO2 and nitrite is decomposed by HCl. A solid containing a small amount of NO2- gives quite some fizzling with HCl but no visible NOx whereas a commercial sample of pure nitrite produces copious red fumes. I noticed this during my nitrate->nitrite reduction experiments.

If you find preparation of pure Ba(OH)2 is too hasslesome you could try MnO2+Ba(NO3)2+NaOH.

Anyways keep up the good work, you might be close to finding a working permanganate synthesis.

-----------------------------------------------------------------------------

Here's my short summary - the gist of this lengthy thread:

1. Direct oxidation with hypochlorite: It works but yield is low and permanganate has to be extracted with dry acetone (plenty of chloride & chlorate has to be removed)

2. Direct electrolytic oxidation of pyrolusite/psilomelane (pottery grade MnO2): Doesn't seem to work at all

3. Electrolytic oxidation of manganate made by fusion of MnO2 with KNO3: Works but rather low yield

4. Fusion of MnO2 with KClO3 and subsequent disproportionation of manganate: So far this is the only proven procedure which gives good yield at acceptable effort & expense: http://www.versuchschemie.de/htopic,10934,kaliumpermanganat....

monoceros4 - 5-9-2009 at 05:37

Quote: Originally posted by chief

(thoroughly washed, filtrated, pre-glowed with some nitrate to remove the graphite first !)

That last item seems to me an unnecessary precaution. Why waste the KNO3? The graphite won't interfere anything and it will be left behind anyway when the melt is extracted.

Quote: Originally posted by chief

Ah yes, the acidification: I preferred to ppt. out with Ba(NO3)2, to get the insoluble BaMnO4

For me that's going in the wrong direction. BaMnO4 is not the ultimate goal, KMnO4 is, and so a method for making BaMnO4 from KMnO4 is not useful to me.

DerAlte - 7-9-2009 at 09:15

monoceros4 - nice to see someone else trying to get permanganates. I found it especially interesting that you need NaMnO4 and not the usual lab KMnO4. Barium manganate, being almost insoluble, is probably the only stable manganate - oxygen from the air seems to convert all others to permanagante and MnO2 except in strong alkaline solution.

You are certainly innovative! I like your test for potassium and I don't know why more amateurs do not use the microscope for a quick analysis of reaction products. Often the crystal forms can tell you what you have and very roughly, how much. I use this all the time - it only needs a drip and a bit of patience, plus very rudimentary knowlege of crystallography.

The solubility of NaMnO4 makes it very difficult to crystallize. I believe Na2MnO4 is near impossible. Your idea of a mixed sulphate/manganate is neat. Never realized the salts were isomorphous. This mirrors the formation of an alum from Mn(III) sulphate which can be isolated, although Mn(III) salts are very unstable.

Quote:

That last item seems to me an unnecessary precaution. Why waste the KNO3? The graphite won't interfere anything and it will be left behind anyway when the melt is extracted.

I think chief is right. Try taking some KMnO4 and mixing it with graphite. Apply a match or flame. Even in solution graphite is slowly oxidized by permanganate if you boil it.

Regards, Der Alte

Formatik - 3-2-2010 at 12:18

Quote: Originally posted by Formatik

... Though as far as oxidation of HCl by KMnO4 goes, this is said by at least several references in Brethericks to have given a sharp explosion when conc. HCl was added to solid KMnO4, the explanation was that there is a remote possibility of the permanganate oxidizing the Cl2 to "chlorine oxide". ...

Finally understand now. For the preparation of Cl2 using KMnO4 and HCl, the KMnO4 should have so much water over it, just so that it's covered. Because action of hydrochloric acid onto the dry crystals easily forms small amounts of gaseous MnO3Cl, which can't be entierly removed by washing with water. In light, the permanganyl chloride also decomposes into Cl2O and MnO2. This all being found out by Kümmel, Wobig (Z. Elektrochem. 15 [1909] 252).

Paddywhacker - 16-5-2010 at 21:43

An equimolar amount of NaOH and KOH forms an eutectic that melts well below 200 degrees. So I tried that with MnO2, and tried to oxidize it to manganate with equimolar KNO3 and NaNO3 ... equimolar to keep the Na/K ratio constant and the melt molten.

The MnO2 reacted only slowly, and the nitrate seemed to have no effect. The melt turned green on the surface from reacting with oxygen, but the bulk stayed black/grey.

The photo is after adding the nitrate. You can see the dark bulk of the melt with blue manganate on the surface.
The melt bubbled slowly like thick boiling mud. Most of the MnO2 did not go into solution after 40 minutes of heating.

My reason for doing this was to make the manganate, then stir in zinc oxide to form a zincate/manganate melt. Then to cool, dissolve up in water, filter, and precipitate zinc manganate by lowering the pH with CO2. ZnMnO4 is a mild oxidizer like BaMnO4 with a very easy workup ... you just filter it off.

Edit --- another attempt to upload the photo.
The quantities were NaOH/KOH/MnO2/NaNO3/KNO3 44/61/47.2/23/27.5

[Edited on 17-5-2010 by Paddywhacker]

Sdc10045.jpg - 119kB

[Edited on 17-5-2010 by Paddywhacker]

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