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blogfast25
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Anhydrous MnCl2 (Mn (II) chloride)?
Does anybody here have any experience with making anhydrous MnCl<sub>2</sub> (manganese (II) chloride) from any of the hydrated forms?
Can, for instance MnCl<sub>2</sub>.4 H<sub>2</sub>O be dried to anhydrous MnCl<sub>2</sub> without hydrolysis or
oxidation to +III or +IV? Or is dry chlorination of the metal or MnO the only real possibility?
[Edited on 16-7-2008 by blogfast25]
[Edited on 16-7-2008 by blogfast25]
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kilowatt
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Try heating the hydrated form strongly under circulating anhydrous HCl gas to prevent hydrolysis. You could set up a dehydrator flask with
concentrated H2SO4 and pump the wet HCl from the vessel with the salt back to that so it circulates around and the water ends up in the sulfuric acid.
This method can be used for preparing a number of anhydrous chlorides such as lithium, calcium, aluminum, iron, and more, so I imagine it would work
for manganese too.
I believe one of those little air pumps for aquariums may be suitable for HCl gas, but I have never tried it. They have no metal parts inside, just a
rubber diafragm (which may deteriorate after awhile) and plastic parts. Regardless, I would put it on the suction side so it only sees anhydrous HCl.
Having an actual pump like that would make the process much more convenient and efficient.
[Edited on 16-7-2008 by kilowatt]
The mind cannot decide the truth; it can only find the truth.
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blogfast25
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Looking through an old chembook, it suggests heating the double salt MnCl<sub>2</sub>. 2 NH<sub>4</sub>Cl. 2
H<sub>2</sub>O.
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ScienceSquirrel
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We used to do it the old fashioned way.
Reflux the hydrated salt in thionyl chloride!
Works every time....
[Edited on 17-7-2008 by ScienceSquirrel]
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not_important
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The hydrate chlorides dehydrate well. Slowly heat to 200 C in a slow stream of N2, or under vacuum.
Alternatively reflux in isopropanol + to;uene, xylene, or other azeotrope former, until no more water comes over; evaporate the alcohol to get alcohol
complexes related to the hydrates, heat to 120-150 C under reduced pressure.
Yet another way would be to add a few drops of hydrochloride acid and let sit sealed up for a day, then heat with stirring to 200 to 250 C, cool, mix
well with ~5% weight NH4Cl, heat this to 300 C under slightly reduced pressure for awhile, then slowly raise the temperature to 400 C. You can pump it
down cold, close off the exhaust tube, and then heat; use large diameter tubing and air condenser. Using larger amounts of NH4Cl chlorides that are
sensitive to hydrolysis, such as the hydrated lanthanide chlorides, can be dehydrated as well as by using SOCl2 but at much lower cost.
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blogfast25
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Yep, there are all sound ideas.
Does anybody know what exactly the role of ammonium chloride (salmiac) would be in this drying process? Does it have something to do with the fact
that it sublimes? Does it carry off the water as an azeotrope, somehow? I'm tempted to attempt the salmiac double salt route: it seems so simple -
adding the right amount of salmiac to an MnCl2 solution of known strength, slowly evaporating until the double salt is obtained and gently heating
further until the right weight loss is reached(?)
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not_important
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Don't even need the double chloride, MnCl2 is pretty stable against hydrolysis and not a lot of NHG4Cl is needed. If it were LaCl3 it would be a
different story. In any case simple mechanical mixing of the simple chlorides works.
The NH4Cl splits into NH3 and HCl, and thus functions similar to passing HCl gas over the drying salt. I've never tracked down anything very recent,
but the NH3 also may help kicking loose H2O from the hydrated salt. The method works with most reasonably non-volatile chlorides; and can be done
with NH4Br for bromides.
The process also can be done using oxides or carbonates and a goodly excess of NH4Cl, which can be handy. Take battery 'manganese oxide', wash it
well, heat it to redness in air, with stirring, for 10 or 15 minutes, cool, wash well again, and finally mix with a several-times excess of NH4Cl and
run the process.
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blogfast25
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I see, interesting...
I'll try with a mechanical mix of MnCl2 hydrate and ammonium chloride.
The purpose is to attempt reduction of anhydrous MnCl2 with Mg, as an alternative to the rather troubled manganese thermite reactions, which
invariably waste a lot of manganese through evaporation. An MnCl2/Mg reduction would run much cooler (perhaps too cool, going by my estimates...)
And if anybody here has the heat of formation of MnCl<sub>2</sub> (anh., at 298 K), I'd be grateful... NIST doesn't list it, no point
looking there.
[Edited on 17-7-2008 by blogfast25]
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not_important
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Electrolytic manganese metal from chloride electrolytes.
http://www.springerlink.com/content/r587344g25q6554x/
http://www.springerlink.com/content/x4x8n45210p72858/
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blogfast25
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Today I set out to make about 20 g of MnCl<sub>2</sub> hydrate and simultaneously (but separately) about 1 mol of
NH<sub>4</sub>Cl. The latter went smoothly.
But during the MnCl2 synth. something strange happened.
I used (for this 'pilot' occasion) pottery grade MnCO3, which is slightly contaminated with Fe (a pinch dissolved in HCl with KSCN added gives a
slight colouring of FeSCN<up>2+</sup> and possibly with Mn2O3 and/or
MnO2. It's beige in colour.
I dissolved the requisite amount of this pottery carbonate in the requisite amount of strong HCl (about 15 w%). Some acid insoluble residue was
filtered off and the solution remained very slightly turbid at pH ≈ 0.5.
The 200 or so ml was then simmered to about half the volume and the solution cleared up. I then added about 100 ml of 32 w% HCl with the purpose of
eliminating any residual Mn (IV).
And this is the strange thing: on heating, the solution, previously light orangey, turned a clear green... What could be the green colour?
I then simmered the solution until almost no water was left and finished off the drying process in an oven at about 150 C (302 F). The resulting
coarsely crystalline product is white to pink (the first time I've actually seen the famous 'manganese pink').
Later on today I'll test the finished product for solubility. But in theory I'm now ready to start dehydrating it by heating it with the salmiac. The
hydrate will be ground together with the salmiac to obtain an intimate mixture. But what temperature should I use? NH4Cl sublimes at 338 C, would
heating the mixture at about 275 C (527 F) be enough or would that be a slow boat to China?
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gsd
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I found following topics in "INORGANIC SYNTHESES" on this subject:
IS - 1, pp 28-33
11. ANHYDROUS RARE EARTH CHLORIDES
IS - 4, pp 104-111
36. ANHYDROUS METAL HALIDES
IS - 5, pp 153-156
43. ANHYDROUS METAL CHLORIDES
Links to all these issues of IS were posted previously in this forum by KMnO4
gsd
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DerAlte
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@blogfast25
One certainly cannot produce pure anhydrous MnCl2 by heating the hydrate with 4H20. At least I cannot! Even as crystals, it oxidizes and hyrolyses
unless kept acid. The method with NH4Cl sounds good and simple - preserving the acid environment during dehydration.
The green color is puzzling - AFAIK all Mn(II) salts are pink to reddish. Mn(III) salts are green but difficult to make and nearly impossible to
preserve.
The standard enthalpy of MnCl2 is -481 KJ/mol (CRC and Pauling agree). This is higher than MnO. The sulphide MnS has Ho = -214 Kj/mol. if you are
searching for a low value for your thermite. Al2S3 retails at -734 KJ/mol., a lot less than the oxide (-1676).
Regards
Der Alte
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Picric-A
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A while ago i made anydrous MnCl2 by slowly heating it in a U tube while passing dried hdyrogen chloride over it.
i passed the HCl vapaour that came out of the U tube into water to make more hydrochloric acid.
oh i made the HCl by heating sodium bisulphate with sodium chloride. H2SO4 and sodium chlride would also work, i just fell it is one hell of a waste
of conc H2SO4 lol 
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not_important
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| Quote: | Originally posted by blogfast25
...
Later on today I'll test the finished product for solubility. But in theory I'm now ready to start dehydrating it by heating it with the salmiac. The
hydrate will be ground together with the salmiac to obtain an intimate mixture. But what temperature should I use? NH4Cl sublimes at 338 C, would
heating the mixture at about 275 C (527 F) be enough or would that be a slow boat to China? |
Check , as already stated, Inorganic Synthesis #1, pp 28-33. This is the NH4Cl procedure, except they start with dry oxides. A temperature ramping
from 150 up to 250 C should remove the water, followed by heating to 340 C under reduced pressure to remove excess NH4Cl.
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blogfast25
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@DerAlte:
I will soon enough find out whether I can reproduce the green colour or not. Explaining Mn (+III) in these conditions would be hard. Remember that the
solution was quite concentrated in Mn<sup>2+</sup> and with a large excess of Cl<sup>-</sup>. Are there no known complexes
like, say MnCl<sub>x</sub><sup>2-x</sup> ?
Thanks for the HoF on MnCl2 but Jeffrey from AmazingRust.com had already helped me out there. I had estimated it to be - 487 kJ/mol based on a
correlation between the HoF of some monoxides and the corresponding dichlorides, that turned out to be a good estimate.
As regards the MnS reduction with Al, that would be a really interesting proposition, particularly because the MP of Al2S3 is only about 1,100 C (well
below the BP of Mn), if it wasn't for the stinky nature of the process. I've used the oxidation of Al with S many times as a heat booster reaction for
the SiO2 reduction (with Al). Works very well... except you're inundated with H<sub>2</sub>S! The metal will stink to high
heaven unless you can re-melt it to get rid of inevitable slag inclusions. And MnS, AFAIK, can only be made using a soluble sulfide.
Since as you're the resident 'manganese nut', you wouldn't know of a pret-a-porter method for quantitatively determining Mn2+, preferably by
titrometry, would you?
@Picric-A:
For larger quantities, using the dry HCl method may not be enormously practical.
@not_important:
I don't have access to IS unless it's in the library here (doesn't appear to be the case).
I will use quite an excess of NH4Cl (about 1 mol of it per mol of MnCl2) and dry it at 275 C. As I have no vacuum either, it'll be a case of 300 - 400
C until constant weight.
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DerAlte
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@Blogfast25
WRT my statement that no Mn(II) compounds are green, I forgot that both oxide and sulphide can be! The hydrated ion (usually 4-6 H2O)) seems to be
from pink to red, but anhydrous salts are usually white, IIRC.
There are double salts like K2MnCl5, an Mn(III) compound (Brauer). About 25 yrs ago, during only of chemistry phases (they happened when my sons were
in high school) I made some Mn2(SO4)3 by heating MnO2 and conc, sulphuric acid. The liquid turns green but actually seperating the manganic sulphate I
did not manage. I’d like to try again but haven’t got any conc acid.
Now whether MnCl3 is formed during the oxidation of HCl is moot. If MnO2 is treated with cold conc. (30%) HCl ( at 0C) it dissolves without evolution
of chlorine to a very dark brown liquid, color of molasses, which some claim contains MnCl4. Heating a bit evolves Cl2, and at one stage it looks
greenish before all the color goes except maybe a slight pinkish. (Needs pretty pure MnO2 to be convincing)
Incidentally, I read (Mendeleev, Princ. Chem.?) recently that heating MnSO4 to 850C+ directly produces Mn304 plus, I assume, SO3.
| Quote: | | Since as you're the resident 'manganese nut', you wouldn't know of a pret-a-porter method for quantitatively determining Mn2+, preferably by
titrometry, would you? |
Nut yes, expert no! Titrometry is difficult to do accurately. I usually estimate as carbonate or dioxide, but titometrically the problem is that all
solutions of Mn++ have to be acidic (else they exhibit oxide/hydroxide formation) and the use of say, sodium carbonate has to neutralize the acid, or
by the use of hypochlorite for dioxide pptn., produces chlorine. You have to standardize the hypochlorite first as it’s notoriously unstable.
Further, the endpoint is difficult to discern due to slow pptn. and cloudy solution. So I usually use either of these in excess and weigh the
carbonate or dioxide produced, after careful drying, one thing I can do. (Also I broke my one burette recently!).
Carbonate seems best. The dioxide is always hydrated, and I assume MnO2.H20 but that’s pure assumption. I store Mn as carbonate, too. In A full
bottle it takes a long time to show signs of oxidation by turning brown, if well dried. MnO2 is also a good way.
Even so, both methods require the absence of metal ions other that alkali metals. Carbonates of most metals are also insoluble, as well as
oxides/hydroxides. But it’s close enough for government work, even if not pret-a-porter (nice phrase!).
Regards,
Der Alte
[Edited on 27-7-2008 by DerAlte]
Edited to add missing MnSO4 in above.
[Edited on 28-7-2008 by DerAlte]
[Edited on 28-7-2008 by DerAlte]
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blogfast25
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@DerAlte:
Not enormously keen on gravimetry on account of no decent furnace and scales limited to 0.01 g. I believe Mn<sup>2+</sup> may be titrated
with EDTA but I have no ready to use method.
I conducted a few experiments on the drying side.
#1:
6.4 g of a mixture of the MnCl2. x H2O (I write 'x' because I have reason to believe the product is tetrahydrate, mixed with dihydrate and possibly
some anhydrous too. I didn't recrystallise) and NH4Cl in a ratio of roughly 1 mol MnCl2 to 1 mol NH4Cl, was heated in an open steel crucible on a
medium propane gas flame. Temperature unknown but almost certainly above 338 C because fumes of NH4Cl started to evolve right away. NH4Cl elution
stopped after about 20 mins and after 30 mins the weight loss was 3.1 g. This could possibly tie in with loss of both NH4Cl and H2O. The product was
slightly discoloured (greyish black) near the bottom.
I continued heating, checking weight loss every 15 mins. The product continued to lose weight and the discolouration became increasingly strong. After
60 mins the product had lost about 3.5 g of weight and was a mixture of off-white blotches and grey/black matter. Presumably oxidation to (I presume)
Mn2O3 was taking place.
#2:
6.3 g of the same mixture was heated in a propane fired oven at about 290 C. The product continued to lose weight up to about 90 mins (1.7 g) but no
discolouration took place and I saw no fumes of NH4Cl evade either. The weight loss doesn't even account for the amount of NH4Cl and it's safe to
assume most of the weight loss is water.
#3:
The remaining 4.6 g of product from #3 was heated in the same conditions as #1 but this time under inert atmosphere, i.e. a stream of homemade,
dry CO<sub>2</sub>. Total weight loss from #2 and #3 was 2.7 g in 30 mins (of #3). As in #1, the product started fuming off NH4Cl
right away and the fuming stopped after about 20 mins.
#4:
6.5 g of the MnCl2 hydrate/NH4Cl mixture was heated in the same conditions as #3 (under inert atmosphere). Fuming started right away and stopped after
about 20 mins. After 30 mins the weight loss is about 3.4 g. The product was slightly discoloured near the bottom but overall is off-white.
Conclusion: for my purposes method #2+3 is probably the best. I assume the product from #2 is dry MnCl2 + dry NH4Cl. Batches could be
prepared in advance and stored in a desiccator. When needed, the required amount could then be fumed off under CO2.
The final product does not appear to be deliquescent although the sample from #3 had picked up a little weight over 24 h (covered loosely with kitchen
foil). It would best stored in a dessicator or prepared immediately prior to use.
I'll have to get one of those ammonium/ammonia aquarium test kits to check for residual NH4Cl...
[Edited on 28-7-2008 by blogfast25]
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not_important
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You need to run this in glass or ceramic, the NH4Cl will corrode ferrous metals when hot (in effect you're boiling HCl in them)
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blogfast25
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Yes, the steel is already corroding. Problem is that I have very little Pyrex glassware. But a small (200 ml) conical flask will do. I could even
recover most of the ammonium chloride with a cold trap or two... 
And the end product (from #2+3) tests surprisingly positive for Fe, much more so than the MnCl2 hydrate. So it must have picked up some from the
crucible, at least that's the most likely source of contamination...
**********
I repeated the procedure for converting the MnCO3 to hydrated MnCl2 yesterday and the green colour appeared again after adding HCl to the concentrated
raw MnCl2 solution. Reducing the liquor further and the solution becomes greener and greener, ending up a kind of emerald green. Yet when completely
evaporated, the product is nicely pink, with some whiter areas too.
I know for sure the MnCO3 contains significant quantities of Fe, maybe that's the cause as FeCl2 is definitely green (but the Fe in the carbonate is
most likely Fe (+III), can Mn (+II) reduce that to Fe (+II)? I didn't think so...). I'll now eliminate the Fe by converting the MnCO3 to acetate,
using vinegar. At 0.77 M [HAc] this will not dissolve any Fe oxides and careful filtration should yield an MnAc<sub>2</sub> solution
that's free of Fe. Then convert back to carbonate and then chloride.
[Edited on 28-7-2008 by blogfast25]
[Edited on 29-7-2008 by blogfast25]
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blogfast25
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Well, well. The green colour doesn't appear to be caused by the presence of iron. I separated the Fe into one fraction, obtaining another that must be
free of Fe (or at least almost). Both were then treated the same to obtain the MnCl2 (hydrate). In both cases the green colour appeared in the very
concentrated MnCl2 solution. The MnCl2 fraction containing the Fe tests very mildly positive for Fe<sup>3+</sup> with H2O2 and KSCN, the
Fe-free MnCl2 doesn't give any reaction at all.
Unless I'm wrong on this it's either another contaminant causing this or a concentrated solution of MnCl2 in strong HCl is indeed green in colour... A
case for Inspector Woelen?
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DerAlte
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The green color is a bit OT but I found this lurking in my files:
 .
Doesn't sound like a contaminant. Pyrolusite (pottery grade MnO2) does contain Fe, I've found.
It says that halogens do form complexes with Mn(ii) but feebly. The pink color is really a H20 complexed ion. Elsewhere I have seen that CN forms the
ion complex
Mn(CN)6- - - - .
(Compare CuCl2; this is bright green in conc. HCl. bluish other wise.)
Hell, you might even be able to extract it with EDTA, if the source is correct. {IIRC from some lecture notes from Laval U. Que., Canada}
Regards,
Der Alte
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blogfast25
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Yes, CuCl2 gave me the idea that there might be some MnCl<sub>x</sub><sup>2-x</sup> complexes. In all likelihood I've
inadvertently created conditions in which these can exist. Obviously on further evaporation of the HCl, they break down and MnCl2 crystallises out.
I'll see if temperature has an effect on the colour next time. Other than that, I won't pursue this any further...
My Mg powder is on its way so I'll be testing the reduction reaction 'shortly'...
Pottery grade MnO2 I bought a while back contained an estimated 20 w% of Fe (expressed as Fe2O3)...
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woelen
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I have done quite some testing about this MnCl4(2-) complex or other similar complexes, but I am inclined to say that these do not exist, at least not
under the conditions which are representative for the experimental conditions, described in this thread.
I indeed noticed the jaune-vert color as shown in DerAlte's post, but I am quite sure that this is due to oxidation of a tiny fraction of the
manganese(II) to some higher oxidation state. Even a tiny amount of oxidation results in a green/yellow/brown color.
The following webpage contains the result of my experiments and based on those I come to my conclusion.
http://woelen.homescience.net/science/chem/exps/manganese/in...
So, I think that the green/yellow color, which is ascribed to MnCl4(2-) ions in reality is just due to oxidation of a tiny fraction of manganese(II)
in the concentrated acid.
EDIT by woelen: Changed link, so that it works again.
[Edited on 12-6-12 by woelen]
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chloric1
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Blogfast- You know your "green manganese" reminds me of an experience I had 4 years ago. I want chlorine gas AND divalent manganese. I started with
potassium permanganate and added 32% HCl. A vigorous reaction of coarse ensued and I got copious chlorine, probably along with vaporized HCl mixed
in. I let the acid solution set open air for a couple days in midwinter(-5 Celsius) and checked it to be an intensely green solution. I proceeded to
heat this and it was reduced to the pale pink manganese(II). If I remember corectly, I posted it in this forum. I am unsure if I had a trivalent
manganese but its vulnerability to mild heating seems to be indicative. Maybe trivalent manganese is stablized somewhat by excess chloride ion. Is
there a simple oxidometric test for manganic(III) ion?
Fellow molecular manipulator
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blogfast25
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In my conditions I thought that trivalent Mn as the source of the green colour can be more or less safely rejected because of the starting material. I
use pottery grade MnCO<sub>3</sub> (for now) as the source of MnCl2 hydrate. The pottery grade isn't 100 % soluble in 32 w% HCl and a
brown-red residue always remains. I know this to contain Fe (as Fe2O3) but it might also contain Mn<sub>2</sub>O<sub>3</sub>
(which is also brownish-red). It's possible that during (RT) treatment of the carbonate with HCl some of the suspected Mn (III) oxide does dissolve
but the quantities could, IMHO, never justify that strength of the green colour. But Woelen's experiments appear to show otherwise, so now I'm not so
sure... It's a little "unresolved" I feel...
Your conditions are very different from mine.
Despite continuing minor difficulties in producing anhydrous MnCl<sub>2</sub> in sufficient quantities and reliably, the first test
reaction MnCl<sub>2</sub> + Mg --> Mn + MgCl<sub>2</sub> was quite successful. The first results were reported in the Exotic thermites & analogs section, here (post of 10/08/08 - 7.30). Further developments will be posted there.
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