Sciencemadness Discussion Board - Anhydrous MnCl2 (Mn (II) chloride)?

Anhydrous MnCl2 (Mn (II) chloride)?

blogfast25 - 16-7-2008 at 09:50

Does anybody here have any experience with making anhydrous MnCl2 (manganese (II) chloride) from any of the hydrated forms? Can, for instance MnCl2.4 H2O be dried to anhydrous MnCl2 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]

kilowatt - 16-7-2008 at 10:00

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]

blogfast25 - 16-7-2008 at 12:17

Looking through an old chembook, it suggests heating the double salt MnCl2. 2 NH4Cl. 2 H2O.

ScienceSquirrel - 16-7-2008 at 16:04

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]

not_important - 16-7-2008 at 18:54

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.

blogfast25 - 17-7-2008 at 06:22

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(?)

not_important - 17-7-2008 at 07:54

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.

blogfast25 - 17-7-2008 at 08:30

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 MnCl2 (anh., at 298 K), I'd be grateful... NIST doesn't list it, no point looking there.

[Edited on 17-7-2008 by blogfast25]

not_important - 17-7-2008 at 18:16

Electrolytic manganese metal from chloride electrolytes.

http://www.springerlink.com/content/r587344g25q6554x/

http://www.springerlink.com/content/x4x8n45210p72858/

blogfast25 - 26-7-2008 at 07:57

Today I set out to make about 20 g of MnCl2 hydrate and simultaneously (but separately) about 1 mol of NH4Cl. 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+

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?

gsd - 26-7-2008 at 09:04

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

DerAlte - 26-7-2008 at 09:36

@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

Picric-A - 26-7-2008 at 14:52

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

not_important - 26-7-2008 at 22:52

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.

blogfast25 - 27-7-2008 at 05:03

@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 Mn2+ and with a large excess of Cl-. Are there no known complexes like, say MnClx2-x ?

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 H2S! 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.

DerAlte - 27-7-2008 at 14:42

@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]

blogfast25 - 28-7-2008 at 07:31

@DerAlte:

Not enormously keen on gravimetry on account of no decent furnace and scales limited to 0.01 g. I believe Mn2+ 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 CO2. 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]

not_important - 28-7-2008 at 07:37

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)

blogfast25 - 28-7-2008 at 08:24

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... :cool:

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 MnAc2 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]

blogfast25 - 30-7-2008 at 07:56

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 Fe3+ 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?

DerAlte - 30-7-2008 at 22:48

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

blogfast25 - 31-7-2008 at 10:12

Yes, CuCl2 gave me the idea that there might be some MnClx2-x 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)...

woelen - 10-8-2008 at 14:00

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]

chloric1 - 10-8-2008 at 17:57

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?

blogfast25 - 11-8-2008 at 08:39

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 MnCO3 (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 Mn2O3 (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 MnCl2 in sufficient quantities and reliably, the first test reaction MnCl2 + Mg --> Mn + MgCl2 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.

woelen - 11-8-2008 at 09:21

Blogfast25, even MINUTE quantities of manganese in higher oxidation state give a deep green/brown color. The almost black liquid in the test tube on my webpage was made from 2 ml of 30% HCl and just 3 little specks of solid KMnO4, eack speck having a size of just 2 mm diameter, or even less. If your MnCO3 only contains a few thenths of percent of manganese in higher than +2 oxidation state, then you'll get a dark green/brown liquid when you dissolve this. On heating this color will fade though. The presence of iron will make the liquid appear yellowish, due to formation of the FeCl4(-) complex, which has an intense yellow color.

blogfast25 - 11-8-2008 at 11:16

Woelen:

I accept what you say but "2 ml of 30% HCl and just 3 little specks of solid KMnO4, eack speck having a size of just 2 mm diameter" is not that dilute a solution. If Mn(III) is responsible at very low concentrations, wouldn't that in itself point to some strongly coloured complex of Mn3+? Surely the Mn3+ ion, even with water ligands cannot be that strongly coloured?

[Edited on 11-8-2008 by blogfast25]

woelen - 11-8-2008 at 11:30

I do not say that no Mn(3+) complex is formed in the experiments (both yours and mine). I am saying that no such Mn(2+) complexes are formed.

A nice test would be to add a tiny pinch of Na2SO3 to your green solutions. I expect them to either turn pale pink, or yellowish (the latter if iron is present in your solutions).

The Mn(3+) ion does not have such a strong color. I made a manganese(III) compound in another experiment in a solution, which does not contain chloride, but sulfate only at very low pH. This ion is reddish/brown. Here follows a picture of a solution of Mn(3+) in the presence of lots of sulfate, but no chloride:

EDIT by woelen: Changed link, so that it works again.

[Edited on 12-6-12 by woelen]

chloric1 - 12-8-2008 at 02:19

Interesting

Well, I would say that many transition metals form intensely colored chloride complexes and manganese is no exception. I guess this is the reason why transition metal nitrates and sulfates are usualy preferred over chlorides for precipitating hydroxides, carbonates, oxalates etc. But some metals like chromium form complexes even with sulfate where only a portion of the sulfate precipitates with addition of barium chloride

.

[Edited on 8/12/2008 by chloric1]

woelen - 12-8-2008 at 03:43

Indeed, most transition metals form complexes with chloride, but at lower oxidation states, the formation of complexes is less pronounced. Manganese in oxidation state +2 is not at all strong in formation of complexes, but manganese in oxidation state +3 is much more so.

Chromium is a completely different story. The metal in oxidation state +3 has the special property that it not only forms complexes, but these complexes also are very stable, and once they are formed, the ligands are not easily replaced. For instance, I have chromium(III)sulfate and also potassium chromium(III) sulfate (chrome alum). Both are dark purple solids and solutions are purple with a blue/grey hue. When barium salt is added to such a solution, all sulfate is precipitated. However, when such a solution is boiled for a while, then the solution turns green. The sulfate ions then replace some of the water ligands and in that case, barium salts only precipitate part of the sulfate. It takes weeks before such a green solution has reverted to the purple/blue solution after cooling down. A similar thing happens when dichromate ion is reduced by sulphur dioxide in acidic medium. The sulphur dioxide is transformed into sulfate, and this sulfate immediately is coordinated to the chromium(III), formed from the dichromate. Hence, the solution becomes green. When dichromate is reduced by e.g. nitrous acid in excess acidic medium, then a purple/blue solution is obtained.

chloric1 - 12-8-2008 at 15:39

I think only chromic nitrate and perchlorate are the only common non-complexing salts. Woelin, am I correct that coordinated transistion metal compounds are covalent by nature? Or is the ionic/covalent character on a sliding scale based on valence and ligand properties?

DerAlte - 12-8-2008 at 21:43

@Woelen

How can your be sure that what you have in your picture is Mn+++? The Mn+++ ion is a powerful oxidizer, about the same as MnO4- in acid solution, pH~0 (SEP~1.5 for both)... It is unstable at higher pH, IIRC.

Have you tried dissolving MnO2 in conc. HCl at 0C as I mentioned above? It dissolves easily to produce a very deep brownish red solution which is said to be MnCl3 (or possibly MnCl4 even); raising the temp causes chlorine to be emitted and there is a transient stage at which a green coloration can be seen. I have read somewhere that the ‘MnCl4’ can be extracted with ether to give a green solution. (You can’t keep ether here; it virtually boils unless refrigerated.)

Mn+++ is far less stable than Fe+++. Like Fe(III) it is said to produce a (red) alum with ammonium, also unstable, more so the violet Fe(III) version.

As I mentioned above, MnO2 heated with conc. (98%) H2SO4 produces a dark green solution (plus O2). This is said to be Mn2(SO4)3. So what color is the Mn(III) ion?
Regards,

Der Alte

woelen - 12-8-2008 at 22:42

According to literature, Mn(3+) can be made by careful reduction of permanganate with malonic acid. It also can be made by careful oxidation of manganese(II) ions by any oxidizer and in the absence of chloride ions, and in the presence of malonic acid. I'm not sure why the malonic acid is needed, but apparently it only reduces to the +3 state quickly and the final step is going much more slowly.

I have also done oscillating reactions with malonic acid and manganese(II)/(III) systems and then the solution oscillated between colorless (+2 oxidation state) and the color shown above (+3 oxidation state). All these things make me quite sure that this is the color of manganese(III) in aqueous solution, without the presence of chloride.

When some of the above reddish solution is added to concentrated hydrochloric acid, then a green/brown solution is obtained, as shown in my webpage from a few posts before.

If you have malonic acid yourself, please try what I have described above. It is very interesting...

@chloric: I indeed am inclined to say that coordinate bonds have a covalent nature. The ligand shares a free electron pair with the metal core. But there also is kind of sliding scale. Some coordinate bonds are weak (labile), the electron pair is only slightly shared with the metal core, while others are really strong and then one cannot distinguish anymore between a normal bond and a coordinate bond.

A nice example is the ammonium ion. This is formed by means of a coordinate bond between the free electron pair of NH3 and the H(+) ion. So, it is formed thus:

H3N: + H(+) --> [H3N:H](+)

Once the bond is made, one cannot at all distinguish between the newly added H-atom and the other three, all bonds are equally strong and one cannot tell anymore which H-atom is from the H(+). If lateron a strong base is added, and NH3 is split off again, then there only is a 25% chance that the original H-atom from the H(+) ion is not in that NH3-molecule.

blogfast25 - 13-8-2008 at 08:37

I see the 'manganese wars' are still raging - :cool:

A little sick of working with various (and variable) sources of Mn (pottery MnCO3, battery MnO2) I'm now purchasing 250 g of MnSO4. H2O (reagent grade, only about 5 bucks) as a relatively pure source of Mn, for conversion to anhydrous MnCl2 (via carbonate). This should be Fe-free and I'll test the green solution obtained the way as Woelen suggests (although with metabisulphite).

And tomorrow I'll test the following dehydration procedure: dissolve (more or less) equimolar amounts of MnCl2 and NH4Cl, then crystallise and grind up. Then fume off the NH4Cl directly at about 400 C in a stream of dry CO2.

The various batches of anhydrous MnCl2 seem to differ a little in deliquescence, some remain quite dry even in open air, others seem to 'go liquid' quite quickly. I suspect the level of Fe contamination may be causing this: I imagine anhydrous FeCl3 to be highly deliquescent.

chloric1 - 13-8-2008 at 13:58

That should work nicely blogsfast. Hopefully, at least some of the carbon dioxide will combine with the ammonia to leave the HCl free. If not the CO2 will at least keep the pH lower so the manganese chloride won't hydrolsize.

blogfast25 - 14-8-2008 at 07:45

Thanks, Chloric, fingers crossed (FYI I ran some dichromate boosted thermites - see the relevant thread).

While making the first 90 g batch of this MnCl2 hydrate/NH4Cl co-crystallised mixture, I ran a couple of tests that seem to confirm Woelen's ideas.

I dissolved two 'pinches' of pottery MnCO3 in separate test tubes in excess HCl and boiled both up. Before boiling both are a reddish/brown but on heating they turn a kind of kaki green, which on cooling disappears to a light orangy-ish. On heating, the green reappears. Adding some metabisulphite (Na2S2O5

to one and the colour disappears altogether. Adding NaOH to the (hot) second one and whitish-pink Mn(OH)2 precipitated.

The fact that the colour disappears on adding a reduction agent does indeed strongly suggest the colour is due to a labile higher oxidation state than +II. But which one?

DerAlte - 14-8-2008 at 20:21

@blogfast25 & Chloric: Nice to see, in the thermite thread, some interesting work being done way above the ke3wl level! Makes me want to try some myself!

And, WRT this ion color thing we seem to have stumbled on, I found the following, which may help explain diverse things we & Woelen have discused:

Quote:

• Mn3+ causes red and green colors in octahedral sites. Muscovite mica from Brazil containing is red as is Mn3+ in beryl from Utah, synthetic orthopyroxene, and piemontite from Whitewater, California. Andalusite containing Mn3+ is green. In the amphibole, tremolite, from New York, it produces a violet color.
• Mn2+ usually results in a pink color in octahedral sites. Rhodonite from Minas Gerais, Brazil, is a pyroxenoid containing Mn2+and has the typical pink color of Mn2+ minerals. Rhodocrosite from Colorado has a high concentration of Mn2+ and a bright red color. At lower concentrations, Mn2+ causes pale pink color. When the Mn2+ is in a tetrahedral site, then yellow-green color results such as is the case with willemite.
• Fe2+ in forsterite from San Carlos, Arizona, and in phosphophyllite from Bolivia is the ion responsible for the green color. In some minerals with high concentrations of Fe2+, such as fayalite or orthopyroxene, the color is brown.
• Fe2+ in the square planar site of gillespite or eudialyte produces a rasberry red color.
• Fe2+ in the eight-coordinated site of pyrope garnet from Tanzania produces the near-red color.
• Fe3+ in octahedral sites causes only pale color when the Fe3+ ions are isolated from each other by intervening silicate ions, etc. Pale purple color is found in phosphates such as strengite and sulfates such as coquimbite. Yellow-green can be found in ferric silicates such as andradite garnet from Italy.
• Fe3+ is in the tetrahedral site of plagioclase feldspar from Lake County, Oregon, produces a pale yellow color. In an unusual variety of diopside containing Fe3+ in a tetrahedral site, it produces bright orange color in thin section.
• Co2+ in synthetic olivine and cobaltian calcite from the Kakanda Mine, Zaire, causes a typical reddish color. In tetrahedral sites, Co2+ causes blue color such is found in some spinels.
• Ni2+ in synthetic olivine has the green color typical of Ni2+ in an octahedral site. If all the nickel is forced in to the larger M2 site by appropriate chemical substitution (in this case in a LiScSiO4 olivine), the color is yellow, typical of Ni2+ in large, distorted sites.

Source:- http://minerals.caltech.edu/COLOR_Causes/Metal_Ion/

I am surprised that Mn(III) occurs in minerals, except that it is very stable with oxygen. Not being a mineralogist(? sp.) , I have no idea of the quoted mineral compositions. I added the stuff for other transition metal for general interest. There's a lot more on that site for those interested.

Regards

Der Alte

blogfast25 - 15-8-2008 at 03:39

Quote:

Originally posted by DerAlte
@blogfast25 & Chloric: Nice to see, in the thermite thread, some interesting work being done way above the ke3wl level! Makes me want to try some myself!

Regards

Der Alte

What's the "ke3wl level"??

ScienceSquirrel - 15-8-2008 at 03:43

Quote:

Originally posted by blogfast25

Quote:

What's the "ke3wl level"??

Blowing things up and making a general nuisance of yourself

not_important - 15-8-2008 at 07:59

DerAlte beat me to the punch. Mn(II) in a tetrahedral environment is both a different colour, yellow-green, and much more intense, by several orders of magnitude, than octahedral Mn(II). However all the references I've got state that the tetrahedral complexes are not stable in donar solvents (H2O, ROH, RCO2H).

Mn(III) is reasonable stable in non-aqueous environments and where it is not in solution; Mn(OH)2 is oxidised by air to MnO(OH) and other Mn(III) hydrated oxides, the white Mn(II) turning to brown. Alkaline conditions favour this, acid inhibits it.

Colour changes on heating need not indicate changes in oxidation state, many complexes change colour when heated and revert when cooled.

Given that so many posters have said that their pottery grade MnO2 contains iron, I think that the possibility that some of the colour effects may be coming from Fe contamination, especially given the weakness of Mn(II) absorption bands.

blogfast25 - 15-8-2008 at 09:41

Quote:

Originally posted by not_important

Colour changes on heating need not indicate changes in oxidation state, many complexes change colour when heated and revert when cooled.

Granted, but how to explain the colour change on addition of a reducing agent (here basically SO2

woelen - 15-8-2008 at 13:05

I have the opinion that the only explanation is that Mn(III) or higher oxidation state is reduced to Mn(II). This also perfectly matches the remarks made above, that the green Mn(II) complexes are not stable in water.

blogfast25 - 16-8-2008 at 05:18

@ Woelen:

Hmmm... doesn't appear that clear cut to me...

DerAlte - 16-8-2008 at 20:31

@blogfast25

Re: What's the "ke3wl level"??

I think ScienceSquirrel described it well. It was meant to be a compliment to both you and Chloric for conducting scientific investigation into high energy reactions.

Incidentally, since I am at least 2 generations removed from them, I misspelled "Kewl", not being au fait withe their prole talk. The correct kewlese is, I believe, k3wl. I make no apologies to them, of course.

WRT Woelen's last comment, that Laval(?) slide I gave earlier suggests that any of the complexes of Mn(II) have a low K (equilibrium const.) of formation but can be 'extracted' by EDTA. The 'jaune-vert' (yellow-green) MnCl4 - - one might be stable in excess HCl.

But I think we are chasing a red herring here. Mn can produce about every color under the sun, and many of the transition metals do likewise.

Regards,

Der Alte

blogfast25 - 17-8-2008 at 04:35

Thanks. Der Alte...