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blogfast25
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Quote: Originally posted by Arthur Dent  |
[...]
Then the stuff got put away on the top shelf for a few months until now. So yesterday, I dropped the tea-colored Manganese chloride solution in a
flask and with a small funnel, poured the hydroxide dust in the flask. Instant blackness!
The solution right now is pitch-black! I know I have to "aerate" the solution a bit for the reaction to take place, but I don't have an air pump for
now. More to come during the weekend.
Robert
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Well, that IS decidely strange and you may be ‘onto something’.
When neutralising MnCl2 with NaOH, firstly Mn(OH)2.nH2O (off white) precipitates, which upon contact with air quickly goes brown by oxidation:
Mn(OH)2 === > MnO2 + 2 H+ + 2 e
½ [ O2 + 4 H+ +4 e === > 2 H2O ]
Mn(OH)2 + ½ O2 === > MnO2 + H2O
Given a little time and air and stirring, this reaction proceeds to completion.
What doesn’t happen is IMMEDIATE precipitation of something black: freshly precipitated MnO2 is always very dark brown, rather than black anyway and
made the way you did it starts off from off-white (Mn(OH)2). Even when adding a strong and alkaline oxidiser like hypochlorite (commercial bleach) you
get that typical colour transition, only much faster.
I’m also surprised you describe the remainder of the MnCl2 solution as ‘tea coloured’: it should really be pink or if weak, almost colourless.
One possibility is that the solution contains iron, in mixed state of oxidation: Fe2+ AND Fe3+. Co-precipitated that gives black magnetite, Fe3O4 (see
for instance ‘ferrofluids’)… The black would possibly mask anything else...
[Edited on 18-3-2011 by blogfast25]
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Arthur Dent
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Yeah I'm surprised also. Here is the list of the only reagents that were ever involved in this reaction:
1) Battery crud from Carbon Zinc batteries, thoroughly washed with distilled water.
2) Technical grade 32% Hydrochloric acid
3) Distilled water
4) Technical grade Sodium Hydroxide (small granules)
I can vouch for the relative purity of all the reagents. When I first prepared the manganese chloride some months ago, I expected to obtain a
tea-colored solution just like the 1st post of this thread, due to the Fe impurities, so to get rid of the iron, I precipitated some of the solution
into insoluble hydroxides with my NaOH, and I washed and dried the resulting precipitate. So far so good.
What I was expecting was that the dried Mn(OH)<sub>2</sub> would help precipitate the remaining Fe impurities and that with adequate
aeration of the solution, I would obtain a pale pink chloride solution with a black/brown Fe/Mn precipitate at the bottom.
Now it looks like there's an inordinate amount of MnO<sub>2</sub> in suspension in the solution, but in that case it would be opaque and
would have started settling, but it doesn't seem to settle at all. 2 days after the experiment, the solution is a very deep transparent brown, like
cola, or even darker.
Robert
[Edited on 18-3-2011 by Arthur Dent]
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blogfast25
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| Quote: Originally posted by Arthur Dent |
[…]
When I first prepared the manganese chloride some months ago, I expected to obtain a tea-colored solution just like the 1st post of this thread, due
to the Fe impurities, so to get rid of the iron, I precipitated some of the solution into insoluble hydroxides with my NaOH, and I washed and dried
the resulting precipitate. So far so good.
What I was expecting was that the dried Mn(OH)<sub>2</sub> would help precipitate the remaining Fe impurities and that with adequate
aeration of the solution, I would obtain a pale pink chloride solution with a black/brown Fe/Mn precipitate at the bottom.
Now it looks like there's an inordinate amount of MnO<sub>2</sub> in suspension in the solution, but in that case it would be opaque and
would have started settling, but it doesn't seem to settle at all. 2 days after the experiment, the solution is a very deep transparent brown, like
cola, or even darker.
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Not sure what you mean by “dried Mn(OH)<sub>2</sub>”: on drying (in air) that turns immediately to MnO<sub>2</sub>. Did
you mean you just allowed it to drip ‘dry’?
Have you tried filtering a small amount of the suspension? What does the filtrate look like?
I think peache’s method, assuming that’s what you’re trying to emulate to get rid of Fe, might be quite sensitive to experimental conditions…
Alternatively, precipitate everything with a small amount of hypochlorite, to a pH of 10 or so. This will ensure all Mn is as
MnO<sub>2</sub> and all Fe as Fe<sub>2</sub>O<sub>3</sub> hydrate. Then acidify with dilute
H<sub>2</sub>SO<sub>4</sub> (dilute HCl will also work, as will vinegar) to a pH of about 5. Allow to stand for a bit. The
iron re-enters solution, the Mn stays as MnO<sub>2</sub>. Filter and wash and redissolve MnO<sub>2</sub> in strong HCl.
Evaporate slowly to dryness.
[Edited on 18-3-2011 by blogfast25]
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woelen
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I never obtained a totally black precipitate from manganous solutions, not even at high pH in the presence of strong oxidizers. I have observed the
brown color of hydrous MnO2 many times, but never pure black. There must be a strong contamination of the manganese in your case, but I am wondering
what it could be. Alkaline iron(II)/iron(III) mixes also are not totally black, they tend to have a dark grey/blue color and on longer standing they
turn brown.
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Arthur Dent
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| Quote: Originally posted by blogfast25 |
Not sure what you mean by “dried Mn(OH)<sub>2</sub>”: on drying (in air) that turns immediately to MnO<sub>2</sub>. Did
you mean you just allowed it to drip ‘dry’? |
Aw crud! There's my mistake. Yeah, I let it dry completely. Ugh.
My bad. So basically, I just chucked some MnO<sub>2</sub> and some iron oxide back into my solution. Back to square one. 
I think my best bet will be to start from scratch with new MnO<sub>2</sub> from a better source. I learn new thangs everyday! 
Robert
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blogfast25
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| Quote: Originally posted by Arthur Dent |
Aw crud! There's my mistake. Yeah, I let it dry completely. Ugh.
My bad. So basically, I just chucked some MnO<sub>2</sub> and some iron oxide back into my solution. Back to square one.
I think my best bet will be to start from scratch with new MnO<sub>2</sub> from a better source. I learn new thangs everyday!
Robert |
No, no, no, no! Recycle the stuff you’ve got! There’s nothing wrong with it. And because it hasn’t been calcined yet, it will re-dissolve easily
in HCl 32 %.
Alternatively try grinding it up and selectively leaching out the iron with a weak acid (or a weak solution of a strong acid) but the success of that
method will depend on the state of dehydration of the Fe2O3 and its age. Fairly dry Fe2O3 will only dissolve in fairly strong acid (but that also
dissolves MnO2!) It’s worth a shot: you can also try this selective leaching at BP to speed it up. Check the leachate for colour and with NaOH for
any Fe(OH)3 dropping out. Easy.
-----
Or how about another route? Dissolve your stuff in 50 % H2SO4, applying heat. Both oxides will go into solution, Fe as Fe2(SO4)3 and Mn as ruby red
Mn2(SO4)3 (yes, Mn [+III] sulphate, not [+II]!) Now add plenty of salt solution (NaCl). Mn3+ is as capable of oxidising Cl- as is Mn
[+IV]:
Mn3+ + e === > Mn2+
Cl- === > ½ Cl2 + e
So your Mn [+III] is reduced back to MnCl2 with liberation of Cl2.
Then neutralise the solution carefully stirring with NaOH to about pH ≈ 8 - 9. Then further neutralise with Na2CO3. The stable MnCO3 (off white)
precipitates. This DOES NOT OXIDISE in air (not even when wet) like Mn(OH)2 does, you can keep it for years. Filter and wash carefully. Dissolve in
HCl to convert back to MnCl2. Of course that doesn’t get rid of the Fe but FeCl3 is highly soluble in acetone. Wash your MnCl2 crystals with plenty
acetone to extract the soluble FeCl3…
[Edited on 19-3-2011 by blogfast25]
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Arthur Dent
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Since my solution of Manganese Chloride with Ferric impurities was in one of my nicest 500 ml boiling flasks, I decided to put it in a plain mason jar
'till I have time to experiment with it.
Since last week, the solution had turned back to its original dark tea-color, with a small black precipitate of MnO<sub>2</sub> at the
bottom of the flask. So to rinse and clean off the flask, I decided to drop in a good 100 ml of 10% hydrogen peroxide and chuck the whole thing in the
jar... Wow! Magic! 
After a lot of fizzing and evolution of oxygen, the solution became water clear and all traces of the precipitate dissolved! It looked like plain
water in the sun! That's pretty spiffy!
But very slowly, the liquid is getting a bit turbid and the Manganese Dioxide is slowly reforming a grayish suspension in the jar.
Now two intriguing questions...
1) If there are ferric impurities in my solution, the H<sub>2</sub>O<sub>2</sub>, shouldn't have affected the ferric chloride
at all and the solution should have remained somewhat yellow after the peroxide addition, shouldn't it ?
2) My beautiful boiling flask shows grayish stains inside... so I filled the flask with H<sub>2</sub>O<sub>2</sub> and thought
they would instantly dissapear, but they seem to be quite stubborn. I know that Manganese Dioxide is a bitch to clean off, so how should I deal with
these stains?
Robert
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Mixell
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You can try pouring some on the solution to another vessel (if you don't have any manganese), and reduce the iron and the manganese in the solution to
the metallic form using aluminium.
Take the manganese/iron powder and pour that into the first vessel and heat maybe to speed up the reaction (can be done with the aluminium too). And
stop when all of the iron in the first vessel will be reduced and you should have a pure manganese chloride solution (presumingly you had enough
iron/manganese powder).
I haven't done that myself, but by comparing the reduction potentials, it should work.
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blogfast25
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Robert:
1) Yes. Fe3+ cannot affect H2O2 (neither oxidise nor reduce it). But don’t stare yourself blind to the colour: in dilute solutions the colour is
quite deceptive. Just the water from your H2O2 solution could have diluted it enough to affect colour.
2) MnO2 is a bitch to clean off but strong HCl will get rid of it, each time and every time, fresh, wet, old, dry, whatever.
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blogfast25
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Mixell:
Even if that does work, you're then contaminating your solution with aluminium and will have to separate these two (Mn and Al).
[Edited on 27-3-2011 by blogfast25]
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Mixell
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Well, I'll suggest using aluminium sheets (to easily separate from the power, this one worked fine for me).
Or to wash with concentrated alkali, either way, you will have an aluminium free powder.
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blogfast25
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No, no, no. For Fe3+ to plate out, the Al has to go into solution:
Al(s) + Fe3+(aq) === > Al3+(aq) + Fe(s), a redox reaction: the Al supplies the electron needed to reduce the Fe3+, the Al therefore inevitably
oxidises to aqueous Al3+. That's how it is...
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Mixell
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Well, thats what I meant...
You separate a part of the solution, insert aluminium- manganese and iron come out.
Insert the manganese/iron powder to the other part of the solution (not the aluminium one) - manganese goes in, iron comes out. And you receive a pure
solution of manganese chloride.
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Arthur Dent
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@ blogfast25: You're right about the HCl to clean my flask, I filled it with 32% HCl and will leave it in for a couple of days, because I really want
it ultra clean since i'll eventually use this flask in my home brewing and strong spirits distillation apparatus.
Just to make sure, i'll probably do a conc. sulphuric acid bath after that to be on the safe side. And much of my "distillation" glassware will also
be treated in such manner. I do have some brand new glassware that has never been exposed to anything more than distilled water so that will also be
used.
As for the Manganese solution... i'll put it on a shelf till I find some time to tinker with it.
Thanks guys!
Robert
--- Art is making something out of nothing and selling it. - Frank Zappa ---
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blogfast25
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Robert: good!
Mixell: Al also plates out Mn2+:
Mn2+(aq) + 2e === > Mn(s) … -1.185 V (red)
Al(0) === > Al3+(aq) ... +1.662 V (ox)
Cell pot. = Ered + Eox = - 1.185 + 1.662 > 0, ergo following Nernst: ΔG < 0, reaction proceeds.
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Mixell
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Again, that is what I mean...
In the second vessel AL plates out Fe and Mn, then the Fe,Mn mix is added to the first vessel, where the Mn plates out Fe.
I hope thats clear enough.
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blogfast25
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I see. I wouldn't put my noney on that working though...
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Mixell
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Well, I could try it, if I'll find the time and nerve to handle the awful smell of chlorine, I see no reason for why it would not work, but who
knows...
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elementcollector1
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Wait, slow down. Are you saying aluminum metal literally plates out manganese metal and iron metal? Or a compound of each? If it does the metal, can I
melt it to a good-sized lump without oxygen?
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blogfast25
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| Quote: Originally posted by elementcollector1 | | Wait, slow down. Are you saying aluminum metal literally plates out manganese metal and iron metal? Or a compound of each? If it does the metal, can I
melt it to a good-sized lump without oxygen? |
The problem is that manganese is actually quite reactive and freshly prepared metal reacts even with cold water (slowly). Actual plating out of Mn
from aqueous media is therefore impractical.
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elementcollector1
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Aw, darn. Is there another way to get manganese metal (short of thermite)?
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blogfast25
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There's always the method used by E. Glatzel:
http://www.sciencemadness.org/talk/viewthread.php?tid=10249&...
But I think you won't fancy that much either...
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elementcollector1
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You're right, that sounds pretty beyond my reach. So, when the manganese reacts with cold water (assume ice water), what are the products?
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blogfast25
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First it would react to Mn(OH)2, which would then quickly air-oxidise to MnO2.
Like with many other elements, there is no quick'n easy way to home make manganese. But there's always the shops...
If you do want to have a shot at it, a fairly large, slow burning (use coarse materials) MnO2 thermite with lots of CaF2 (to slow 'burn' rate even
more) is probably your best bet. But even that make take many attempts to get it really right.
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elementcollector1
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Durn. My aluminum be powdered too fine for a 'coarse' reaction.
I did make beautiful bubblegum-colored MnCl2 chunks just now, will aluminum reduce those?
For that manganese dioxide thingy, would that be a good way to make a MMO electrode?
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