| Pages:
1
2
3
4
5
..
17 |
DerAlte
National Hazard
Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline
Mood: Disgusted
|
|
You need the KOH for sure. Conc. to induce the KMnO4 to precipitate off the anode (common ion effect) But just suspending the particles - I suppose
the OH- ions near the anode might just get them. As for graphite, does H2SO4 make a mess of them or only when it's really concentrated? PbO2 seems to
have a cult following, 12AX7. Unike most, I have made perchlorate sucessfuly with even graphite. You just have to tolerate the anode erosion and clean
up the mess. If it's inefficient, what does it matter? Carbon and electricity are cheap compared with the prices one sees chaged, if you can get it al
all.
Der Alte
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
I believe graphite forms an oxide (intercalated or something) when subject to anodic conditions. As I recall, it works with any sulfate, and probably
nitrate as well. Seems to me the only way you can possibly get perchlorate with graphite is by cheating the reaction with high voltage and current
density pushing past the erosion regardless.
Speaking of graphite oxide, a lot of times I've had the graphite sludge from my chlorate cell rise to the top due to adherent oxygen. Now, I would
ordinarily attribute this to hypochlorite decomposing slowly, but that only works when the smell of chlorine is strong. Sometimes it happens to
low-hypochlorite solutions. Graphite oxide as the erosion product, with a high oxidation potential (above chlorate, but below perchlorate,
persulfate, etc.), would seem to make sense, and if it's decomposing in suspension, that would explain the adherent oxygen bubbles.
Anyway, applying to this thread, you have to determine if the oxidation of whatever mechanism operates -- direct oxidation of manganite, production of
intermediate peroxide or superoxide, etc. -- if it's lower than graphite's erosion potential.
Tim
|
|
|
DerAlte
National Hazard
Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline
Mood: Disgusted
|
|
12AX7, both boiling nitric acid and I think I read somewhere, permanganates, will oxidize carbon to mellitic acid, which is benzenehexacarboxylic acid
- I'd draw it if I knew how to in this format. The anhydride of this is mellitic anhydride, C12O9, another weird 'oxide' of carbon.
Sure, you use a high voltage and high current density to make the perchlorate. The rods get eaten away like there's no tomorrow! To make the chlorate,
it's essential to make sure anode and cathode products mix - the action is chemical as well as electrolytic; for the perchlorate, this is not as
essential although hypochlorite, still produced, is capable of converting chlorate to perchlorate. I use different electrode configurations (Fe
cathode). Carbon is the sacrificial lamb. Current efficiency be damned! You are spending more effort to make your PbO2 anodes than it's worth (but I'd
love to try one - or Pt!)
With regard to permanganates, I'm working on it (still!)
Regards,
DerAlte
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
Mellitic acid is interesting stuff. I bet its first proton or two is strongly acidic (much as EDTA's first has a pKa in the unity range). And yes,
not to mention the anhydride being an aromatic oxide. Apparently it occurs naturally as the aluminum salt (mellite).
But I digress... maybe titanium doped with sun-dried shyte would work as an anode.  (Hey, don't poo-poo this idea -- excrement contains trace elements!)
Tim
|
|
|
ciscosdad
Hazard to Self
Posts: 76
Registered: 6-2-2007
Member Is Offline
Mood: Curious
|
|
Patent 3652417 (US) is interesting reading.
It refers to Alkaline Permanganate solutions used for descaling. It uses anionic surfactants to decrease the self destruction of the KMnO4 (as I
understand it). It may be applicable to synthesis if it shifts the equbilibrium appreciably.
I have found the boiling point of KOH 40% to be ~132 degC. Has anyone seen a table / graph for higher concentrations? I'm thinking of aerated
digestion of KOH and I expect the temp will need to be significantly over 132 DegC.
Perhaps a Eutectic mix of KOH/NaOH?
More searching.
|
|
|
ciscosdad
Hazard to Self
Posts: 76
Registered: 6-2-2007
Member Is Offline
Mood: Curious
|
|
More searching has turned up Patent No 7 056 424 (US)
This refers to the manufacture of MnO2 electrodes on a stainless steel substrate by initially applying a porous non conducting layer, which may be a
fabric, or some deposited porous ceramic or plastic.
What about a porous pot filled with Mercury! I love Mercury.
It might make the Electrochemical method a lot more appealing. The key to the method appears to be the electrodes.
I wish they would not use the convoluted legalise in these
things. Just getting my head around the chemistry is bad
enough
Patent No 3986 941 (US)
The process is for the electrolytic production of KMnO4 directly from a KOH and MnO2 slurry. (or NaMnO4)
Electrodes can be Stainless Steel (amongst other things)!
20%KOH (aq) at 80 - 90 Deg C. Much more friendly than the process I was thinking of.
Process time is less than a day assuming reasonable currents.
The patent even describes a run using a glass 1 litre beaker 
[Edited on 20-6-2007 by ciscosdad]
|
|
|
DerAlte
National Hazard
Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline
Mood: Disgusted
|
|
12AX7, your anode ideas, although a shade far fetched, may indeed have merit! Is not PBO2 a pleasing velvety brown color? But, I fear, sun-baked shyte
may have poor conduction characteristics. Liberal addition of titanium powder may help this. I am sure its oxygen overvoltage is suitably high!
Talking of sun-baked shyte reminds me of the layer of crud on the bottom of my graphite perchlorate cells – but there it is the pulverized remains
of many a carbon rod that gave it’s life for the cause…
Ciscosdad, nice thinking and research. Damn, before this thread dies of inanition, we’ll have a method, by hook or by crook. A clever and persistent
chemist can do anything… Sorry no data on KOH or NaOH/KOH mixtures
When I have time I’ll report on my latest effort and suggest a few ideas – at present I’m midway through testing…
Regards,
DerAlte
|
|
|
ciscosdad
Hazard to Self
Posts: 76
Registered: 6-2-2007
Member Is Offline
Mood: Curious
|
|
For those of you interested, the Patent Search site I'm using is:
http://patft.uspto.gov/netahtml/PTO/search-adv.htm
The patent Number I'm currently obsessing about is 3 986 941.
As I read it, the procedure can be:
In a (say) 2 litre steel container....
Make a slurry of MnO2 in 20% KOH solution so that there is 4.2 moles of KOH per Mole of MnO2. Add a catalytic quantity of KMnO4 to improve current
efficiency in the early stages (recommended but not essential as I understand it). Heat to 80 to 90 Deg C. Insert stainless steel elctrodes and pass
10A through the solution for ~18hrs while maintaining the temp.
Conversion is of the order of 98 - 99 %.
Little is said about subsequent extraction of the KMnO4 (the solution was apparently merely analysed to determine conversion ratios).
Perhaps it can be as simple as hot filtration followed by crystallization of the KMnO4. A quick bit of mental arithmetic implies a concentration of
~150g of KMnO4 in the approx 1Litre of solution.
To be determined:
1......Filtration and cooling will certainly recrystallize the KMnO4 , but is it soluble enough at 80 - 90 deg C to allow uncomplicated filtration of
the xs MnO2?
2......Does the solution need to be neutralized or acidified to work up?
3......Will the electrodes give problems with the MnO2 deposits mentioned elsewhere? Ref Patent No 7 056 424
4.......Current density required?
5.......Will agitaiton be necessary?
There is the possibility that the mother liquor (after KMnO4 recrystallization ) can be reused as is for the next cycle. Simply add approx 1 mole each
of MnO2 and KOH and repeat. It even has some KMnO4 left in it to help start the process.
If you guys have read the patent, this post is probably unnecessary, but I'm trying the get the small scale / improvised procedure straight in my own
mind. Have I missed anything?
Ref Item #4
Current density is quoted as 5 - 50 mA/cm2 (converted)
Current Concentration is 3 - 30 A/Litre
At 10A the area of the anode is: 250 - 2500 cm2.
I assume the iron rod (or whatever ) cathode area is not an issue, but the problem would be easily rectified by using another cylindrical electrode.
I'm hoping that the resistive heating of the cell will supply a significant amount of the heat required to maintain the required 90 Deg C. Insulation
and an idling hotplate under the reactor should do the job.
[Edited on 21-6-2007 by ciscosdad]
|
|
|
DerAlte
National Hazard
Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline
Mood: Disgusted
|
|
A few thoughts on the production of KMnO4
The paramount thing is to minimize the production of chlorate during the process. So I have been racking my brains as to how best do this. I was
looking at the possible kinetics of the process when it struck me that the equation I had given supra was in fact a combination of a fast reaction and
a slow one.
2MnCO3 + 5KClO + K2CO3 --> 2KMnO4 + 5KCl + 3CO2
On addition of the hypochlorite, the MnCO3 quickly turns (brownish) black due to (hydrated) MnO2 production:
MnCO3 + KClO -> MnO2 + KCl + CO2
This is followed by oxidation of Mn(IV) to Mn(VII) as permanganate, a very slow process as noted before. Hence the driving potential is quite small,
not the approx 0.34 volt redox potential from Mn++ to MnO4- but rather the 0.6 volt needed to convert MnO2 to Mn-. (The hypochlorite provides about
-0.89 volt to achieve this)
(Note: none of these values apply to the conc. solutions we are using. I am too lazy to apply the Nernst equation to get realistic values)
So there is actually no point in using the carbonate unless you have it. You can start with the chloride directly from the manganes purification (but
neatralize to acid to prevebt chlorine production first), or any Mn(II) you happen to have.
So, lets call it Mn++ + ClO- + 2e --> MnO2 + Cl- to keep it general.
The great thing about this is we can separate the MnO2 as a ppt. and use it for the following step. This has three advantages: (1) It avoids the
addition of unnecessary chloride ions because the dioxide is precipitated and can be separated and dried; (2) if you have MnO2 that is relatively pure
to start with, use that. (3) you can use cheap old bleach, NaOCl, for this step and not care about the extra load of NaCl it will introduce.
The MnO2 can be dried and weighed. It is likely hydrated to some extent as X(MnO2), Y(H2O). Dry at about 200C to minimize hydration, or even higher.
The (very!) slow reaction is then
2MnO2 + 3KClO + K2CO3 --> 2KMnO4 + 3KCl + CO2
Or, in general, 2MnO2(s) + 3ClO- + 2OH- --> 2MnO4- + 3Cl- + H2O
(We could have used hydroxide instead of carbonate, which in strong solution produces the OH- ions: carbonate is used to keep the pH so that
permanganate is produced instead of manganate – this is why, I guess, that CO2 is used to convert manganate to permanganate in the fusion process)
The burden of Cl- ions is thus reduced in the ratio 5:3. We haven’t done anything to avoid the disproportionation of the ClO- ions yet.
This side reaction is 3ClO- ClO3- + 2Cl-. (Alkaline solution understood throughout). The standard redox potential for this (also slow)
auto-oxidation is about 0.49 volt, which is actually less than the 0.6 volt needed for the Mn(IV) to MnO4- transition. So my original assumption that
the desired oxidation ought to beat the undesired were wrong.
My original assessment that about the same amount of chlorate was produced as permanganate bears this out. A word now about the kinetics. To produce
one ion of permanganate needs 3/2 ClO- ions; but for the chlorate, 3 ClO- ions. This should favor the permanganate in less concentrated solution, so
my assumption that the concentration should be high may also be at fault.
Finally, given that chlorate is inevitable, is there a method to separate it? from the permanganate? We have to think of some solvent other than
water, and that means organic. Well, in another thread, mericad193724 in the thread "KMnO4 Synthesis" gave us the clue. Acetone dissolves KMnO4. Now
it seems the stuff he was using destroyed his product, but it should not have, as far as I know. Permanganate converts alcohols to Ketones, and IIRC
is then stable toward them. Acetone has a dielectric constant near 20, not too polar. But permanganate is a far stronger acid than chloric, although
both are strong. Question: does acetone dissolve KClO3 to any marked extent. As far as I can gather, it is almost insoluble, <0.1 %, versus KMnO4
at about 11%. Anyone got a comment on this – acetone should separate the two. I don’t have any acetone at present, but must try this on my impure
samples.
Sorry about the length of this but I just had too many thoughts and it helps to get it down.
Regards,
DerAlte
|
|
|
ciscosdad
Hazard to Self
Posts: 76
Registered: 6-2-2007
Member Is Offline
Mood: Curious
|
|
Acetone
The acetone looks good. What about a continuous extraction with a soxhlet apparatus? I assume the extraction would be done on the dry product.
Don't know about KClO3 solubility in the acetone, but I would expect it to be low.
Acetone has a low BP, so it should all happen at nice low temps.
I like the idea of prior generation of MnO2 from Mn++.
I personally would use MnSO4 for its easy availability.
|
|
|
not_important
International Hazard
Posts: 3873
Registered: 21-7-2006
Member Is Offline
Mood: No Mood
|
|
Simply exposing Mn(OH)2 to air will give MnO(OH), a mixed Mn(III) and Mn(IV) hydrated oxide, or even hydrated MnO2. So you could use the even cheaper
air to save some on the MnO2 prep.
You don't really want to dry the MnO2 that much, just vacuum filter suction should do well. Heating it reduces its reactivity, heating to 200 might
be enough to cause the lose of some oxygen.
Permanganate attacks acetone, although slowly. I suspect it may be through the enol form, a C=C double bond and a free OH group are prime targets for
permanganate. Washing the glassware with a strong acid, then rinsing well might reduce this effect, as it is base catalysed. Technical grade acetone
generally has reducible impurities in it as well. If you use it as a solvent, you don't want to have it sit around for too long to avoid too much
loss of the product.
See doi:10.1016/j.jpowsour.2005.03.178
and interestingly enough this
[Edited on 21-6-2007 by not_important]
Attachment: reaction_hazards.pdf (110kB)
This file has been downloaded 2788 times
|
|
|
DerAlte
National Hazard
Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline
Mood: Disgusted
|
|
Not_important, you always have something interesting (and important!) to say.
WRT to the keto-enol tautomerism of acetone, I had forgotten that. I am not much of an organic chemist. There’s no doubt KMnO4 eats up double C=C
bonds like Pacman, and also attacks -OH bands in alcohols. I am sure you’ve seen the impressive action against glycerol. Looking it up in my
organic text book, I find that KMnO4 can be used to create vicinal diols from alkenes, but only if you keep the temp low, below 25C. Above that you
get splitting into a couple of acids.
I had proposed to use the acetone on dried product and avoid water, which may catalyze the enol formation. I’ll give it a try. Keeping it cool may
help to slow any reactions.
The reason for drying the precipitated MnO2, xH2O was to be able to weigh the product. It’s easier, I decided, to then weigh the unreacted product
later, to determine the amount of permanganate produced, than to do a ferrous ammonium sulphate titration. (One has to assume all the reacted Mn goes
to permanganate, of course. It could go to Mn++ ions too, I suppose, but I don’t see how, in the presence of the hypochlorite)
As you can imagine, all this measurement is very tedious but necessary to get an idea of how worthwhile any method is. One then assumes that the
chlorate comes from the ureacted part of the hypochlorite, the rest to chloride.
AS to heating, CRC says MNO2 decomposes at 525C; permanganate at around 200C. So it should be pretty safe to heat the hydrated stuff to about 200C.
But one doesn't nkow how much water is left, unfortunately.
That doi abstract was interesting. Couldn’t download the pdf for some reason.
Regards,
DerAlte
|
|
|
DerAlte
National Hazard
Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline
Mood: Disgusted
|
|
Managed to download the pdf, Not_important. Germane to the problem. I assume it refers to KMnO4 and acetone in a highly acidic medium from the number
of H+ ions. What a hideous equation – too many prime factors in it!
Well I got my acetone, industrial solvent. I evaporated some on Al foil – very slight residue, probably organic (charred on rapid heat?) and a small
less volatile liquid residue, probably water.
I now have a small pile of glistening black acicular crystals. It worked.
I cooled the acetone to about -10C first and poured it over a gram or so of my previous dried solid product, a mixture of chlorate, chloride and
permanganate. This was in a small funnel with a glass plug. The permanganate dissolved readily without incident. The solution was just like a water
solution – very deep color. I re-poured a few times and finally gave it a small wash. Then evaporated at ambient (30C+ here) in a shallow dish in an
air stream from a fan – quite rapid. (Don’t mess with acetone indoors). The final result of this was a mixture of black goo in with some liquid
and solid and no smell of acetone or ketone. The liquid was water, I believe. It evaporated too, leaving little black needles with no sign of cubic or
plate type crystals typical of chloride or chlorate.
There were signs of MnO2 but not serious – slight brown stains. With permanganates this is par for the course. There may have been some reduction.
If I were going to do this in any quantity I would not leave the acetone solution hanging around, would keep everything cool and evaporate under
suction.
On the basis of this I claim to have made solid tolerably pure KMnO4 by a CRUD method – (Chemical Reagent from Utter Dross). I am not suggesting it
is very practical, however. Yields are low and acetone isn’t that cheap, unless you re-condense it.
Regards,
DerAlte.
|
|
|
ciscosdad
Hazard to Self
Posts: 76
Registered: 6-2-2007
Member Is Offline
Mood: Curious
|
|
Congratulations DerAlte!!!
Brilliant piece of work.
It seems a pity to evaporate the acetone off at the end and effectively lose it.
Do you think it will be feasible to vacuum distill to recover the acetone? What condition was the chlorate residue left in?
|
|
|
DerAlte
National Hazard
Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline
Mood: Disgusted
|
|
Ciscosdad, I don't see why one could not distill under reduced pressure. Acetone boils at about 59C, (IIRC) at atmospheric. Might want to use cooled
water in the condenser.
Flash point of acetone is -20C. almost as bad as ether at -40C ( - in the climate here ether has to be kept under refrigeration). I cooled the
solvent to minimize any reaction with it or impurities. Main impurity seems to be water, which could be removed by a drying agent such as anhydrous
CaSO4.
I find from a manual I have that KMnO4 is used as a test for organic impuities in acetone. It should not, apparently, reduce KMnO4, assumedly at room
temp.
The Chlorate (plus chloride) was heavily stained purple. of course. Some MnO2 obvious also in the insoluble remainder in the filter funnel.
I have a few more thoughts on this type of process but need to do a few more tests, if only to avoid making a fool of myself...
Regards,
DerAlte
|
|
|
Ballermatz
Harmless
Posts: 19
Registered: 17-7-2007
Member Is Offline
Mood: No Mood
|
|
Awesome work so far, "Alter" 
I found this VERY interesting article which provides detailed information on the production of various manganese components:
http://www.mrw.interscience.wiley.com/emrw/9780471238966/kir...
The industrial process is also outlined. One interesting fact is, that oxidation from Mn(+IV) to Mn(+VII) does not only go through Mn(+VI) but through
Mn(+V) as well. The conditions favouring formation of Mn(+V) do not favour the further oxidation to Mn(+VI) however! This is why the industrial
process is split into THREE steps; the first two are seperate roasting processes which bring the Mn from +IV to +V and then to +VI; the last one is
electrolytic oxidation to permanganate.
Thus when you simply fuse MnO2 with KOH (oxidizing using air or KNO3 or whatever else), yields can never exceed 60% of theoretical. Only if you fuse
twice, first favouring formation of Mn(V) and then formation of Mn(VI) you will get better yields.
It is also interesting that decomposition of KMnO4 is highly dependant on pH. Thus if you solve the fused mass in boiling water and the mass is too
alkaline, decomposition will be greatly accelerated.
Direct oxidation to permanganate by fusing with KNO3 (similar to direct oxidation from Cr2O3 to dichromate) is impossible because KMnO4 decomposes at
240°C; much lower than the melting point of KNO3. NH4NO3 would be perfect - except it will most probably explode due to formation of highly instable
ammonium permanganate 
One encouraging fact is that electrolytic oxidation of manganate to permanganate does NOT require any special gadgets like nickel anodes or
diaphragms. Even industrial electrolysis uses plain steel electrodes and most processes dont need a diaphragm.
|
|
|
DerAlte
National Hazard
Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline
Mood: Disgusted
|
|
@ Ballermatz: Thank you for the compliment. I have been very interested in Mn compounds for years. The range of colors, oxidations states, oxyanions
and general chemistry of the transition elements has fascinated me since youth.
Welcome to the forum (though it’s not my place as a novice member to say!). I found your post very interesting.
I have never got the fusion with KOH to work other than to give a faint green manganate color. Even with the aid of oxidants. (Look for other
permanganate threads on this site with the search engine for several efforts in this direction).
Yet older authorities claim that any Mn compound, fused with KOH and an oxidant like nitrate on a Pt foil will produce a very sensitive test for Mn
due to permanganate formation! Must be the magic of Pt! Matches its price! No industrial process I have ever read about ever manages Permangante in
one step, they always get manganate which is electrolytically oxidized to permanganate.
I long ago decided they must use a very special process. KOH melts at 406C; KNO3 @ 337C; KClO3 @ 368C with decomp.; and KClO4 at 525C. All these are
way above the decomp temp of KMnO4, (240C, you quote) or manganate (190C). As for NH4NO3, no go, I’m afraid. It explodes at 210C. Worse, NH4MnO4
decomposes at 70C! (Figures from CRC).
AS far as the oxidation states go, in alkaline solution the direct step Mn(IV) to Mn(VII) has standard reduction potential of 0.60Volt versus about
0.62v for Mn(IV) to Mn(VI). However, the intermediate hypomanganate Mn(V) state MnO4--- has a potential of 0.96 volts to MnO2. I.e. permanganate is
slightly easier to produce than manganate and much more so than hypomanganate. The ion MnO--- can be made by carefully reducing KMnO4 at ~ ph 14 with
a sulphite at about 0C via the manganate. It is a light blue color and quite unstable. I have done this – the sequence purple Mn(VII) to green
Mn(VI) to blue Mn(V) to black (MnO2) is quite remarkable.
I am still tooling around with these processes but am temporarily sidetracked. Hope to report some more later.
Regards.
DerAlte.
|
|
|
Ballermatz
Harmless
Posts: 19
Registered: 17-7-2007
Member Is Offline
Mood: No Mood
|
|
Hi "Alter"
Permanganates all have relatively low decomposition temperatures; KMnO4 seems to be the most thermally stable on my list; much more stable than say
barium or copper permanganate which decompose below 100°C. Decomposition of the formed manganate or permanganate seems to be the most likely reason
for the reported failures to produce them by fusing MnO2 with hydroxides. However, a liquid phase oxidation process IS used industrially, and involves
fusing MnO2 ore with a large excess of KOH (1:5 molar ratio) for 4-5h. However, the temperature is said to be 250°C, yet the mixture is said to be
liquid at all times (WITHOUT water). So they must use some kind of trick to lower the melting point of KOH.
Your hypochlorite oxidation method is clever in this regard, because it does not need such high temperatures. However it suffers from unwanted
auto-oxidation of hypochlorite to chlorate. I remember reading somewhere that NaOCl can be composed into NaCl and O using catalysts like cobalt oxide,
but I dont know to which extend it avoids the chlorate formation.
"I long ago decided they must use a very special process. KOH melts at 406C; KNO3 @ 337C; KClO3 @ 368C with decomp.; and KClO4 at 525C. All these are
way above the decomp temp of KMnO4, (240C, you quote) or manganate (190C). As for NH4NO3, no go, I’m afraid. It explodes at 210C. Worse, NH4MnO4
decomposes at 70C! (Figures from CRC). "
Do you have decomp. temperatures for the hypomanganate at hand? I would guess they`re much higher because the industrial process uses 400°C during
the first roasting (conversion to hypomanganate, Mn(+V)) and 200°C in the second, which forms the manganate. The first roasting kiln is fed with a
slurry of KOH and MnO2 ore, but from the description it becomes obvious that the water is evaporated during the process. It seems like its only used
to produce an intimate mixture and be able to spray the slurry into hot air. During the second roasting, however, water is continously sprayed onto
the mass to keep it wet at all times.
This whole process sounds feasible for the amateur chemist as well. The first roasting process sounds a bit difficult to copy because it involves
spraying a highly corrosive slurry into a rotating kiln. However I imagine that spreading out a intimate KOH/MnO2 mixture on a baking sheet so that a
very thing layer is formed, then spraying water onto it and treating it with a heatgun might do the trick. It will be necessary to heat the baking
sheat by an additional heat source because it has a large surface that will cool down rapidly. The hot air (>300°C) from the heat gun will provide
the necessary O2 without cooling the mixture. Heat until it is dry, maybe mix yet again, wet with water and repeat the step. Important question is:
How can the formation of K3MnO4 be observed? If figure it would be blue in aqueous solution but how does it look when dry?
In a second roasting step, one would carefully reduce the heat to 190°C and spray water onto it to keep it wet during the process. Dont forget that
the whole process will be a matter of 4h or more! The hypomanganate formation is a matter of minutes but turning it into manganate requires more time.
Final oxidation to permanganate could still be accomplished using chlorine from TCCS or calcium hypochlorite.
|
|
|
Ballermatz
Harmless
Posts: 19
Registered: 17-7-2007
Member Is Offline
Mood: No Mood
|
|
I just found that an eutecticum of 90%mol NaOH and Na2CO3 has a melting point of 286°C so there is hope
Best regards
Der Ballermatz
|
|
|
DerAlte
National Hazard
Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline
Mood: Disgusted
|
|
Most fascinating! I believe the hypomangante Mn(V)O4--- forms blue crystals (according to Bauer). I've never made them. The solurion in 30% NaOH at
around 0c is light blue and on keeping rapidly decomposes, especially on heating. Must try to repeat this.
In this context many of the older references talk about manganites of the general structure xK2O.yMnO2 which are generally insoluble and black in
color, hence difficult to distinguish from the dioxide.
There appear to be several forms of crystalalline MnO2, possibly with different reactivity. I do know that technical grade, probably pyrolusite with
75-90% MnO2, is less reacive in the processes described than freshly made MnO2 as above. The color of MnO2 is generally blackish brown as
precipitated, probably due to hydration. The precipitated stuff is also very finely powdered giving a large surface area.
It is obvious that the commerial processes use considerable art and experience and chemistry alone will not lead one to them!
I have some further comments but will reserve those for later, afyer I have done one further experiment to test the idea.
Regards,
DerAlte
|
|
|
chief
National Hazard
Posts: 630
Registered: 19-7-2007
Member Is Offline
Mood: No Mood
|
|
Hello. I'm new here; why not make the KMnO4 from Ba(MnO4)2 by mixing solutions of the latter and of K2SO4: BaSO4 would precipitate, leaving KMnO4 in
the solution.
The potential advantage to be hoped for would be the higher temperature-stabilities of the Ba-compounds, although yet unknown to me about the
Ba(MnO4), but Ba(NO3)2 is stable up to ~ 595 [Celsius]. Thereby it might be possible to easily create the Ba(MnO4)2 and after the reaction with the
K2SO4 reuse the BaSO4 by cooking it with concentrated Na2CO3-solution (which works!, although not quantitatively) and going from the carbonate again.
Only thing not solved for me so far: Ba(NO3)2 also does not melt below the 595 [Celsius];
so how then make the Ba(MnO4) from MnO2 (from old cells, it's the by easiest source !!!) and Ba(NO3)2 or a combination with other nitrates ??
Another route would be to anodically oxidize Mn in a solution of Ba-hydroxide (which gives at least the manganate)
You see: the way would be to somehow create Ba(MnO4)2, and from there on it should be possible to create a lot of manganates by just mixing the
solutions of Ba(MnO4)2 with the sulfates of the elements -- and Ba(MnO4) should be easier to create because of the higher temperature-stability of the
compounds.
Any ideas ??
|
|
|
chief
National Hazard
Posts: 630
Registered: 19-7-2007
Member Is Offline
Mood: No Mood
|
|
Another Idea from me about the Ba(No3)2-melt: should it not decompose immediately above the 595 [Celsius] (it somehow seemingly "boils" [releasing
fumes] at 607-615 [Celsius], but I didn't measure too exactly) one might trade of the decomposition-time of it versus the creation/re-decomposition
time of the thereby created Ba(MnO4)2 but temperature-control.
How is it: do nitrates decompose at a specified temperature or is it a temperature-range with temperatue-dependent decomposition-rate ? If it were the
latter, a half-life-period exist (adjustable by the temperature). And it might be the same for the permanganate. so if these facts were known: 1-time
purchase of a Ba-salt (carbonate or whatever), and just by putting K2SO4 and nitric acid into it obtaining the KMnO4 !!
|
|
|
Ballermatz
Harmless
Posts: 19
Registered: 17-7-2007
Member Is Offline
Mood: No Mood
|
|
"The potential advantage to be hoped for would be the higher temperature-stabilities of the Ba-compounds, although yet unknown to me about the
Ba(MnO4), but Ba(NO3)2 is stable up to ~ 595 [Celsius]"
What we need is actually the opposite - a nitrate that decomposes below 200°C so that the formed manganate will not be decomposed. Ammonium nitrate
would do the trick but explosive compounds will be formed.
It also doesnt work because batrium permanganate has an even lower dec. temp. (100°C) than KMnO4. Ba(No3)2 decomposes >550°C - much higher than
KNO3.
What we really need is a mixture that brings the melting point of KOH or NOH below 200°C.
|
|
|
chief
National Hazard
Posts: 630
Registered: 19-7-2007
Member Is Offline
Mood: No Mood
|
|
Is it really the case that Ba(MnO4)2 decomposes at such a low temperature? Is it not the general trend for Ba-compounds to be quite stable (see the
oxide)?
Besides: Manganese nitrate is said to decompose at above 140 [Celsius] ...
I could have bet the Ba(MnO4) is at least stable up to 500 [Celsius]
|
|
|
DerAlte
National Hazard
Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline
Mood: Disgusted
|
|
I have Ba(Mno4)2 decomposing @ 200C (CRC)
@chief : The double decomp of potassium sulphate with barium permanganate would work like a charm -IF you have barium permanganate. Since the thrust
of this thread is to find an alternative to the fusion process by a wet process at <100C. how do you propose to make Ba(Mno4)2? See my posting for
6-6-2007 above in this thread.
More later!
Regards,
DerAlte
|
|
|
| Pages:
1
2
3
4
5
..
17 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
