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ScienceSquirrel
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The way to make potassium permanganate on a small scale is by fusing manganese dioxide with potassium hydroxide and potassium nitrate.
Extract the cool solids with water to produce a solution of potassium manganate ( green)
Acidification converts it to potassium permanganate.
This easily makes potassium permanaganate on a gram scale and was a standard A level preparation.
I would be wary of trying to make HMnO4 as it readily dehydrates to Mn2O7.
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blogfast25
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In dilute solution permanganic acid (?), HMnO<sub>4</sub>, may be harmless but treating a pure permanganate with strong acid causes the
highly instable Mn<sub>2</sub>O<sub>7</sub> heptoxide to form. It's a recipe for explosions...
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chief
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| Quote: | Originally posted by ScienceSquirrel
The way to make potassium permanganate on a small scale is by fusing manganese dioxide with potassium hydroxide and potassium nitrate.
Extract the cool solids with water to produce a solution of potassium manganate ( green)
Acidification converts it to potassium permanganate.
This easily makes potassium permanaganate on a gram scale and was a standard A level preparation.
I would be wary of trying to make HMnO4 as it readily dehydrates to Mn2O7. |
Yes, only that I want to do it with the NaNO3 and NaOH, in order to have a NaMnO4-Solution, as concentrated, as I can make it !
(I wouldn't go via the H2SO4 (if doing the silver-route), because of the danger, but try to use Na2SO4.)
Can I use a mall amount of HNO3 for the acidification ? As I observed in the past, this readily turns green solutions into violet ones. But also
quantitatively ? Is the HNO3 only catalysator, and small amount sufficient ?
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blogfast25
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| Quote: |
Can I use a mall amount of HNO3 for the acidification ? As I observed in the past, this readily turns green solutions into violet ones. But also
quantitatively ? Is the HNO3 only catalysator, and small amount sufficient ? |
Well, I'm no expert on making permanganate. I've only obtained it inadvertently in alkaline oxidation of Mn(OH)<sub>2</sub> with
hypochlorite, as described on the previous page by DerAlte. Also by fusing MnO<sub>2</sub> with KOH I got small amounts of
K<sub>2</sub>MnO<sub>4</sub> but never managed to convert it to permanganate, it always just fell back to plain old
MnO<sub>2</sub>.
But acidification of manganate solutions should be fine: the danger of Mn<sub>2</sub>O<sub>7</sub> formation only occurs (I
believe) when treating pure (or maybe also very concentrated) permanganates with strong acids, for instance:
KMnO<sub>4</sub> + 1/2 H<sub>2</sub>SO<sub>4</sub> ---> 1/2 K<sub>2</sub>SO<sub>4</sub>
+ 1/2 Mn<sub>2</sub>O<sub>7</sub> + 1/2 H<sub>2</sub>O
If you're acidifying in normal concentrations, starting from a manganate (not permanganate) the risk of making the heptoxide inadvertently must be
very small...
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The_Davster
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The chance of making Mn2O7 starting from any aqueous solution is impractically small. The stuff is destroyed by any ammount of water, so all the
water solvent of your solution would have to be used up hydrating your sulfuric acid...not gonna happen! (unless you count taking your 100mL reaction
and dumping it in a couple litres of conc. H2SO4)
Also, KClO3 is a far superior oxidizer for fusing with KOH and MnO2, and it is completely impractical to not use an oxidizer at all. Atmospheric
oxidation of heated manganate is likely only feasible in industry with fancy heated grinders.
CO2 is the best acidification agent. No remainder ions to clean up and readily accessible from baking soda and vinegar.
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ScienceSquirrel
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I am not sure why potassium nitrate is preferred to potassium chlorate but you will need an oxidiser on a lab scale.
As I remember it the reagents melt to a black paste which solidifies as the reaction proceeds.
We used sulphuric acid to acidify the green manganate solution, you only need a small amount and it is fast and convenient.
Potassium permanganate is preferred over sodium permanganate as the potassium salt is a lot less soluble and readily crystallises.
It should also be noted that sodium permanganate cannot be prepared by this method.
http://en.wikipedia.org/wiki/Potassium_permanganate
http://en.wikipedia.org/wiki/Sodium_permanganate
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chief
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On the wikipedia-link they don't mention, how then the NaMnO4 is prepared. Since _this_ is, what I want, I'm gonna try some melt with Na2CO3, NaOH,
NaNO3, and the MnOx, and check for the green color.
As far as I know, the permanganate-group forms under oxidizing and alcalic conditions, why should not then NaMnO4 form ? Maybe the temp has to be kept
below the decomposition of the NaNO3, so that no NaNO2 (said to be reducing ..., therefore bad for permanganate-group) can form ...
Besides I have some Ba(NO3)2, which decomposes not below 595 [Celsius]; maybe I'll have to use some eutectic of the both, or only Ba(NO3)2 at all ?
The Ba(NO3)2 can always be regenerated, using carbonate to ppt., then dissolving it in HNO3 again ....
[Edited on 15-7-2008 by chief]
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chloric1
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Ion exchange
IIRC there is a British patent for making calcium permanganate from potassium permanganate via ion exchange with zeolites. What it said was you set
up a column and run 30% calcium chloride solution through the zeolites to convert them to caclium zeolite. Afterwards you pass a warm solution of
potassium permanganate though said zeolite. Now they said you could condense and chill the effluent to precipitate unconverted potassium permanganate
as it is slightly soluble or rerun it through the zeolite a time or two more to get a more complete conversion. Ideally you could do this and add
sodium carbonate to precipitate the calcium but I am unsure about permanganate stability in alkaline solution. Maybe adding baking soda with vigorous
stirring then gradual heating so as not to raise the pH of the working solution too sharply.
Otherwise you could "charge' the zeolite with a few passes of concentrated sodium nitrate solution and then add your potassium permanganate to get a
more direct route. I just don't know if the zeolite will exchange sodium for potassium so readily. I still have yet to try this but it is intriging.
[Edited on 7/15/2008 by chloric1]
Fellow molecular manipulator
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ScienceSquirrel
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This is a slightly interesting link dealing with the production of sodium permanganate.
However after the initial stuff the trolls have got at it and it is garbage as far as I can see.
http://backyardchem.chemicalblogs.com/121_backyard_chemistry...
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chief
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Thanks for that link; that's exactly the easy way, that I like:
Heat the MnO2 + "Soda - Saltpeter", equal parts, to "dull red heat" for 16 - 48 hours.
Only: Since I have the MnOx from old Batteries (Mn3O4 ?)(of course washed within lots of water to remove the NH4Cl), maybe there is some residual
graphite within; may the graphite go off with the NaNO3 ? Would not be nice with a 1 kg-batch ..., also the explosion could be much stronger than
gunpowder, since the stuff would be pre-heated ?
Can I glow the graphite away, and at which temperature ? Some sources say at least, that the Mn3O4 may be oxidized to MnO2 by glowing on air.
Also MnO2 is known as an oxidizer, with metal-powders at least ...
And then: Does "Soda-Saltpeter" mean: "NaNO3", OR "Na2CO3 + NaNO3 (ratio?)" ?
[Edited on 15-7-2008 by chief]
[Edited on 15-7-2008 by chief]
[Edited on 15-7-2008 by chief]
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ScienceSquirrel
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Looks like you are in the running for a Darwin Award 
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chief
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No, no. Is it dangerous now ? I only mentioned the MnO2-as-oxidizer since I'm not quite sure, how the graphite would react in such a fine-powdery form
with an oxidizer ...
Besides anyhow the thing will run in my furnace, without me around ...
Why Darwin-Award ???
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chief
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Someone might also be careful about the Zn-Powder within alcaline cells ...
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ScienceSquirrel
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I think that soda - saltpetre is probably sodium nitrate.
Why are you messing around with manganese dioxide from batteries? Manganese dioxide is readily available from a variety of sources, as is sodium
nitrate.
I would buy the pure chemicals and try the reaction on a small scale to start with. More chance of the reaction going right and less chance of a nasty
accident.
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chief
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Yes, small scale, or in a distant, fireproof room ...
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not_important
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| Quote: | Originally posted by chief...
Can I glow the graphite away, and at which temperature ? Some sources say at least, that the Mn3O4 may be oxidized to MnO2 by glowing on air.
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Yes, you can remove the carbon through air roasting. However, some sources are wrong. The most stable oxide is the mixed one - Mn3O4 (MnO. Mn2O3)
which is formed by heating any of the other oxides in air above 950 C.
Mn2O3 can be had by air ignition in the 500-900 C range, although conversion to that oxide can be very slow.
MnO2 starts to break down at temperatures below 300 C on up to over 500 C, depending on how it was originally made.
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Formatik
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| Quote: | Originally posted by blogfast25
In dilute solution permanganic acid (?), HMnO<sub>4</sub>, may be harmless but treating a pure permanganate with strong acid causes the
highly instable Mn<sub>2</sub>O<sub>7</sub> heptoxide to form. It's a recipe for explosions... |
It needs to be dehydrated also to form the compound. I've added some KMnO4 powder to reddish fuming HNO3 with a density of about 1.52, no Mn2O7
formed, and all of the acid fumes turned white. The same is in the well-known chlorine generator reaction of KMnO4 with conc. HCl, though the
permanganate can also oxidize the conc. acid to some ClO2, which is explosive.
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12AX7
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| Quote: | Originally posted by not_important
Mn2O3 can be had by air ignition in the 500-900 C range, although conversion to that oxide can be very slow.
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It works pretty well in the presence of charcoal. Getting Mn3O4 or MnO the same way takes considerably more heat, and gives off much less in return.
Tim
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chief
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I heatet to 530-555[Celsius] the Na2CO3-NaNO3-MnOx (latter fom batteries, washed and cooked 1/2 h in soda-sol.), equal parts of each component, 12
hours long. It was dry, fine milled, and mixed in the coffe-mill, before the heating.
After cooling down, with water: Not the slightest green, no color. Only the black MnOx now is brown, so maybe I at least oxidized the MnOx to MnO2 ...
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12AX7
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Did it fuse completely? Carbonates take a good bit more heat to work with. Try hydroxide.
Tim
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not_important
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Hmmm ... the nitrate would have melted at that temperature, but as it decomposes/gets used up the mixture will become solid. I think you need to heat
to 800-900 C at least, perhaps not mix all the nitrate into the original mixture but rather slowly add it to the fused mass with stirring.
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DerAlte
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Manganates Again
Since it seems that this old thread has been revived, this is a good time to put to rest a few myths I may have inadvertently engendered earlier.
Reference should be made to all the previous posts – it’s too difficult to sort out which are relevant.
Since last year I have repeated a few of the experiments. My conclusions are : (1) The production of permanganate by any wet method is far less then I
thought; (2) The fusion methods do work, though how well I have not determined. I have never had much success with manganate production before, like
mant others.
WRT (1), I have been consistently fooled by how intense the color of the permanganate (and manganate) ion is. An intensely colored solution can be
achieved by a few percent by weight. If both manganate and permanganate are present, an almost total absorption of most of the visible spectrum
occurs. The permanganate ion exhibits a broad absorption band around 520 nm, blocking out the entire green region and leaving a bit of the blue and
the yellow, but has less absorption in the red and blue violet – hence the well known magenta color. In fact, I can still see the color at 3ppm.
Manganate has a wide absorption band in the red region, centered on 620 nm, which is why it appears green. Add the two together and you are left with
attenuated passbands mainly in the blue/violet. Since my source of illumination has been a krypton filled filament lamp, very little light passes when
manganate and permanganate are both present ( this can be detected by filter paper chromatography).
In short: dense looking solutions of permanganates are not very concentrated.
Most of my efforts were concerned with pHs at a level where permanganate only can be produced (pH~ 11). In one effort that I thought particularly
promising, I used ~ 1N NaOH instead of Na2CO3 (pH ~ 14) to get a very dark mainly manganate solution – dark for the above reason.
I castigated the acetone extraction unfairly. I now think it works like a charm – there is just very little permanganate to be extracted, and it
does extract nearly all of it. But the acetone must be carefully dried (CaSO4) and so must the product from which the permanganate is extracted. Very
little of the permanganate gets reduced. Any MnO2 produced tends to color the remaining undissolved chlorate & chloride slightly brownish/red,
possibly due to methanol in the commercial acetone.
The titrations I performed earlier to determine the presence of oxidizer in solution were severely flawed. Hypochlorite, I have found, is not so easy
to destroy rapidly by heat as I supposed, and its effect cannot be distinguished from MnO4- using the Fe++ ion. I.e. remaining hypochlorite is still
present and masks the permanganate, giving an inflated number.
Instead of using Fe++ ion, I repeated the measurement using glucose and ethyl alcohol as reducing agents in alkaline solutions in excess, and
measured the MnO2 produced (carefully washed and dried) by weight. Reasonable consistency using these two methods showed that the manganese
content implied that the oxidation efficiency to be no more than 2%. This also agreed roughly with the amount extracted by acetone in another run.
Somewhat disheartened by this, I did some detailed calculations on what might be theoretically expected. I included all the factors I could think of,
such as concentration of ions, temperature. I spent some time seeking out accurate values for the Standard Electrode Potentials. For the reactions
involved, these are none too certain but consistent to about a few %. The results of this suggested I was lucky to get ~2% product; the conversion to
chlorate dominates.
The next post outlines some theoretical considerations.
Regards,
Der Alte
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DerAlte
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Attempt at Theoretical Reconciliation
*** Ignore the following if you are allergic to theory!
The half-reactions involved, ignoring spectator ions, are as follows (in alkaline standard conditions), written as reductions:
MnO4- + 2H2O(l) + 3e- --> MnO2(s) + 4OH-, Eo = 0.595 v ..(1)
MnO4-- + 2H20(l) + 2e- --> MnO2(s) + 4OH-, Eo= 0.60 v ..(2)
(Notice how close these SEPs are. But the one electron difference is crucial!)
ClO- + H2O(l) +2e- --> Cl- + 2OH- Eo=0.81 v ..(3)
The above are taken from CRC. We also need:
ClO3- + 2H2O(l) +4e- --> ClO- + 4OH- Eo=? (0.525 v) ..(4)
This value is not in CRC and has to be calculated from those that are. I used
ClO3- + 3H2O(l) +6e- --> Cl- + 6OH- Eo=0.62 v ..(5),
in conjunction with (3) above. The Eo for half reaction (4) is then obtained by subtracting the reaction (3) from (5) and taking account of the
electron exchange ratios.
4Eo(4) = 6Eo(5) - 2Eo(3) so Eo(4) = 0.525 v.
The balanced ionic full reactions for the production of manganate and chlorate ion from hypochlorite can then be deduced:
3ClO- --> ClO3- + 2Cl- {from (3) & (4) above, eliminating e-}; dEo=0.265 v ..(6)
2MnO2 + 3ClO- + 2OH- --> 2 MnO4- + H2O + 3Cl- ; dEo= 0.215 v ..(7)
Plus the also ran case producing manganate
MnO2 + ClO- + 2OH- --> MnO4-- + H2O + Cl- ; dEo=0.21 v ..(8)
Or even the long odds for Mn(V) (hypo)manganate
2MnO2 + ClO- + 6OH- --> 2 MnO4--- + 3H2O + Cl-; dEo= -0.15 v. ..(9)
What does all this crapola mean? Basically, that the reactions tend to go to the right the higher dEo is. For the manganate case, dEo is negative,
meaning a tendency to the left. The conversion to chlorate has the highest driving force.
The above values are at the standard conditions T=25C (298K), P=101.33kPa (1 std. atmosphere). I carried out the reactions at ~95C (368K), which
modifies the values a bit. Also, the standard conditions are for solution concentrations as indicated by the equations, in mols/l.
For any other conditions (we can ignore pressure because the reaction are concerned with liquid/solid phases only) we have to use the Nernst equation
E = Eo - (RT/nF)lnQ = Eo – (0.0592T/nTs )log10(Q), in volts.
Where R= (molar) gas constant (J/deg/mol), T is temp (abs.), Ts is the Standard temp ( 298K for most SEP tables), F (faraday) = 96,485 coulomb/mol, n
is the number of electrons exchanged in the reduction equation, Q = concentration ratio, products divided by reactants, as in the equilibrium
equation. In fact, if K is the equilibrium constant, Q = K when E tends to zero – there is then no driving force either way. Hence:
If at equilibrium E = 0, Eo = (0.0592T/nTs )log10(K)
i.e. K = 10^(nEoTs/0.0592T) = 10^(nEo/0.0731) at 95C. Simple enough? It’s now in terms most amateur chemists can understand (Sorry, kewls.)
For (6) above, chlorate production @95C, n=4, Eo = 0.265 volts, K6 = 3.16E14
For (7), permanganate , n=3 per MnO4- ion, Eo=0.215 v., K7 = 6.66E8
For (8). Manganate, n=2 per MnO2-- ion, Eo=0.31, K8 = 5.57E5
For (9), Hypomanganate, n=1 per MnO4--- ion, Eo= -0.15, K9 = 7.87E-3.
What does all this mean? On the face of it, that chlorate is by far the most likely to be produced; then permanganate, then manganate, and
hypomanganate has only a cat’s chance in hell of being produced.
But let’s look a bit deeper. Consider the chlorate case. The reaction assumed is
3ClO- --> ClO3- + 2Cl-
{First notice that this reaction does not depend apparently on alkalinity (i.e. [OH-]). This confused me at first, but actually the equation
implicitly assumes the solution is fairly alkaline: in acid solution the reaction differs, because instead of ClO-, hypochlorite ion, the
hypochlorite will be present as unionized HClO due to the weakness of the acid HClO}
We have
[ClO3-][Cl-]^2/[ClO-]^3 = K6 = 3.16E14
at equilibrium @95C, whenever that is attained. (It says nothing about the reaction rate, merely the final state. This might take forever,
like for hydrogen and oxygen at room temperature. But we know it isn’t, by experiment!)
We have then, (forget the complication of actual activity, for the purists)
[ClO3-][Cl-]^2 = K6 [ClO-]^3 (@ equilibrium)
Typically, a 10% solution by wt. of NaClO has a concentration of 1.5M NaClO plus about the same of Cl-. Since it’s obvious this reaction goes to
near completion, the final Cl- concentration will be ~ 2.5 M and that of NaClO3 about 0.5M. This gives the final concentration of ClO- at equilibrium,
as about 1.6E-5.
Next look at the MnO4- reaction (7). Rewrite as:
MnO2 + (3/2) ClO- + OH- --> MnO4- + (1/2) H2O + (3/2) Cl-
The equilibrium equation is
[MnO4-] = K7 [MnO2][OH-][ClO-]^(3/2) / [H2O]^(1/2)[Cl-]^(3/2)
= K7 [OH-][ClO-]^(3/2) / [Cl-]^(3/2)
(Here the concentration of the dioxide and water are to be taken as 1, since the solid phase and liquid solvent are assumed always in excess.) If the
pH is kept at ~11, then [OH-] will be about 1E-3. If [ClO-] is about 1.6E-5, determined by ClO3- production as above, then the concentration of MnO4-
calculates as
[MnO4-] = 6.66E8 x 1E-3 x (1.6E-5 )^1.5 / 2.5^1.5 = 1.08E-2 mole/L = 1.28g/L. At pH ~11.5, this becomes 4.04g/l and at pH~12, 12.8g/L.
Pretty pathetic! But in line, roughly, with the experimental results.
The estimated production of MnO4- - can be calculated similarly as
[MnO4--] = K8 [OH-]^2 x [ClO-]/ [Cl-] 3.56E-6 mole/L @ pH~11. At pH~12, one gets 3.56E-4 mol/L and at pH ~13, 3.56E-2 mol/L
Thus in high pH conditions, manganate production becomes favorable. One might think that, per the above, permanganate production would be even more
favorable; but, as pH rises, the following reaction takes place, in the presence of MnO2:
MnO2 + 4OH- + 2MnO4- --> 3MnO4-- + 2H2O
This can easily be shown experimentally by adding NaOH to a solution of permanganate in the presence of MnO2; the magenta solution turns green.
If the OH- concentration is very high (~ >5M or about 25% by weight) you can even go another step to Mn(V) (hypo)manganate:
MnO2 + 4OH- + MnO4-- (green) --> 2MnO4--- (light blue) + 2H2O
Correspondingly, small amounts of water readily hydrolyze hypomanganate to manganate and MnO2, and on further dilution permanganate and more MnO2 is
produced.
Very similar reactions happen in the fusion case, where a liquid hydroxide and dissolved or fused oxidizer are present in very high concentration. A
little on this in the next post.
{CAVEAT: the above is purely the work of Der Alte. As such, there are no guarantees from the management. The potential for BS exists, since Der Alte
is very prone to errors in calculation, and has even been known to make logical errors due to senility. Criticisms and corrections form real physical
chemists welcome}.
Regards,
Der Alte
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DerAlte
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Fusion
I have re-tried the age-old fusion reaction with an oxidizer (nitrate) and hydroxide. The proportions used were roughly according to
MnO2 + 2NaOH + NaNO3 --> Na2MnO4 + NaNO2 + H2O (gas)
(I.e. 2 mols NaOH to 1 mol NaNO3, (c 1:1 by weight) MnO2 in excess).
The usual recommended temperature is “dull red heat” which I take to be in the range 400 - 500C, which I believe is just visible in the dark.
Time, 2-3 hrs.
The Japanese recommendation (see thread above) for Hypomanganate is 0.5mol MnO2, 2.5 mol KOH, 0.5 mol KNO3, 2 hours fusion at c 300C. I used the
sodium hydroxide and nitrate
That reaction is assumed to be:
2MnO2 + 6NaOH + NaNO3 --> 2Na3MnO4 + NaNO2 + 3H2O (g)
This is slightly a different molar ratio, 1:3:0.5 versus the recommended 1:5:1, but I used the recommended.
In all cases I used MnO2 that had been made from MnCl2, see earlier in the thread. This is likely hydrated as approx. MnO2.H2O
I did the fusions in identical small steel cans, well cleaned. The NaOH and nitrate fuse together at something like 250C - 270C I believe. I used no
temperature measuring device, merely making sure that the interior of the container glowed just perceptibly in the dark in the manganate case, and the
mix remained fairly fluid in the hypomanganate case. I wrapped the can in a sheath of glass cloth for the manganate case, and heated it with a small
propane flame from a blow torch. I added MnO2 once the solids melte, then slowly raised the temperature to the dull red heat. The other case I heated
on a camp stove, maintaining a sort of fluidity.
Both the NaOH and the nitrate that I used were damp and steam was emitted before the solids finally melted.
I stirred the mixes from time to time manually with a plated steel rod. Both turned pasty in time. I added the MnO2 in small batches. It fizzled on
introduction to the fused liquids, especially in the manganate case. This may have been due to the fact that the MnO2 was probably hydrated and the
reaction also produces steam.
On cooling, the resulting solid mass in the manganate case was a dirty brown with a hint of green. In the hypomanganate case, it was quite similar to
that shown so well in Xenoid’s posts above, bluish with a hint of green and some blackish MnO2 specks.
I treated both messes with a small amount of water. Sure enough, both gave the deep green color of manganate.
Conclusion: With care, the fusion reactions work. I am now convinced. How well is TBD.
Regards,
Der Alte
[Edited on 17-7-2008 by DerAlte]
[Edited on 17-7-2008 by DerAlte]
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chief
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Posts: 630
Registered: 19-7-2007
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Success !
What I did:
==> I weighted in [5.72029375,16.9989416,7.9491552] gm of [Mn3O4,NaNO3,Na2CO3],
==> glowed it some minutes at 800 [Celsius],
==> with water (not directly (!): drops explode (!) by steam-generation, cool down first)
==> gives deep green solution
The solution, with a tiny amount of HNO3, gives violet solution, but a tiny bit more reddish than what I know from KMnO4. When boiling down, it at
first became violet, but turned back to green.
The "Mn3O4" used was from batteries, washed, glowed and pre-oxidized at 500 [Celsius] with NaNO3 (whereafter it was more brown than black).
The equations I used were:
--------------------------------------
MnO2 + 2 O ==> MnO4
Mn2O3 + 5 O ==> 2 MnO4
Mn3O4 + 8 O ==> 3 MnO4
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NaNO3 ==> NaNO2 + O
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Mn3O4 + 8 NaNO3 ==> 3 MnO4 + 8 NaNO2
--------
So _at_least_ 8 parts of NaNO3 would have to be used, because the NaNO2 is known as "reducing", but thats an assumption. Question may be: What wants
more oxygen (?): MnOx OR NaNO2, OR: How oxidizing is a melt of NaNO3 with x % NaNO2 in it ??
---------
Then the final assumption was:
Na2CO3 + MnO4 ==> Na2MnO4 + CO2 + X, no exact equation, and experiment rules.
What do you think about it ? Its very green, and now I try to concentrate it and find a way to either crystallize something (for microscope) or the
measure somehow the quantity obtained ...
But I wanted to share now, so maybe someone wants to try himself ...
[Edited on 21-7-2008 by chief]
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