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Author: Subject: Permanganates
ciscosdad
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[*] posted on 8-10-2007 at 17:16
Refinement


@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?
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Xenoid
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[*] posted on 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...:o

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
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[*] posted on 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?
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[*] posted on 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
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[*] posted on 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
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[*] posted on 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
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[*] posted on 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.
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[*] posted on 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]
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[*] posted on 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.
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[*] posted on 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
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[*] posted on 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
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[*] posted on 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.
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[*] posted on 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.



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[*] posted on 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.

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[*] posted on 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.

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[*] posted on 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
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[*] posted on 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?
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[*] posted on 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.
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[*] posted on 6-11-2007 at 22:30
MnO2 Types


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.
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[*] posted on 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.

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[*] posted on 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.

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[*] posted on 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?

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[*] posted on 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]
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[*] posted on 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.




Things that we dont understand completly and things that completly dont understand.
I did a reaction to find an answer and I ended up with so many questions.
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[*] posted on 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 ...
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