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DerAlte
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@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
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chief
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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]
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chief
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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]
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blogfast25
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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
MgCl<sub>2</sub> 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 MnCl<sub>2</sub> can be obtained by treating battery crud with strong HCl (caution: generates significant amounts
of chlorine gas!) and separating the soluble MnCl<sub>2</sub> from insolubles like graphite powder by filtration. DerAlte has a procedure
to obtain quite pure MnCl<sub>2</sub> 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 MnO<sub>2</sub> can also be obtained from it that way, by
re-oxidising the Mn<sup>2+</sup> with strong bleach to the dioxide. Quite straightforward, OTC chemicals.
Your best bet for KMnO<sub>4</sub> remains fusing relatively pure MnO<sub>2</sub> with KOH in the presence of
KClO<sub>3</sub>, leaching the fusion with cold water to extract K<sub>2</sub>MnO<sub>4</sub> and carefully
acidifying it.
High grade MnO<sub>2</sub> is also available from several eBay sellers.
[Edited on 22-7-2008 by blogfast25]
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DerAlte
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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
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not_important
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| 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.
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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.
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blogfast25
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| 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.
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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. 
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chief
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"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]
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blogfast25
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The thread I intended to start here on Mn2O3/Mn thermites has been spliced onto this one here.
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DerAlte
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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]
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blogfast25
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| 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.
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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 Mn<sup>2+</sup> and
Mn<sup>3+</sup>? Or is the latter not stable enough in aqueous solution?
[Edited on 24-7-2008 by blogfast25]
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not_important
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| Quote: | Originally posted by blogfast25
Couldn't Mn3O4 be prepared the same way as magnetite (Fe3O4), by co-precipitating equimolar amounts of Mn<sup>2+</sup> and
Mn<sup>3+</sup>? Or is the latter not stable enough in aqueous solution?
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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.
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blogfast25
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| Quote: | Easier to do is mixing a solution of a Mn(2+) salt with one of KMnO4 in the proper ratio.
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Not_important, how would that work?
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DerAlte
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@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
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Formatik
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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.
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blogfast25
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@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.
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chief
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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.
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chief
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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
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chief
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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.
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chief
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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.
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DerAlte
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@ 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
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DerAlte
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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
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Taoiseach
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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.
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chief
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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]
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chief
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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)
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