Sciencemadness Discussion Board - Anhydrous MnCl2 is easy to make?

Anhydrous MnCl2 is easy to make?

Altreon - 23-12-2025 at 03:15

On a whim, I decided to synthesize anhydrous MnCl₂ using 10g of pottery grade MnCO₃. After adding just enough HCl to clear up the solution, there were signs of obvious Fe³⁺ impurities, so I added 3g more of MnCO₃ to neutralize the soln. along with a few mL of 50% H₂O₂ (which somehow reacted very slowly) to ppt. out all the iron. The clear, filtered pink (with a very faint brown color) soln. was placed on a 195°C hotplate with magnetic stirring. I know of the many reports claiming MnCl₂ oxidizes very readily, and a stream of HCl or some ammonium double salt is required to dry it. For my random synthesis, I added 20mL of MeOH after the soln. reached 20mL, not expecting anything to change. Instead, the soln. immediately became cloudy and white when stirred, with a very faint bluish green color somewhat like window glass. I waited until the soln. went down to 15 mL, added 10mL of MeOH, and repeated this 2 more times. Every addition intensified the greenish blue color and reduced the cloudiness of the soln. I let it stand for a while until it reached 10mL, at which point it started forming a greyish white slush which slowly turned pink. All this time I've covered the beaker with a watchglass (which along with the beaker was coated in a white powder?) and tested the condensate with pH paper, showing it to be neutral for the whole process. After the solidifying slush, which i had to manually stir, became a very observably pink color, I unfortunately broke the beaker but was able to recover one of the chunks. I added that to another beaker and started heating at 230-260°C, causing it to quickly dry out as a purely light pink powder, showing no signs of oxidation.

I now ask, why does it seem that people go to great lengths to make setups for synthesizing MnCl₂ when it is so easy to make it from aqueous soln.?
The attached images are of the first MeOH addition, the last MeOH addition (ignore the tag on the beaker), the slush (first one is after heating removed most of the grey color + second is the pinkish state), and the dried MnCl₂.

[Edited on 23-12-2025 by Altreon]

[Edited on 23-12-2025 by Altreon]

Altreon - 23-12-2025 at 03:29

Just in case one is concerned about a basic MnCl₂ that somehow has the same color as MnCl₂, the residue was dissolved in water and the soln. remains perfectly clear.

bnull - 23-12-2025 at 07:14

Quote: Originally posted by Altreon

I now ask, why does it seem that people go to great lengths to make setups for synthesizing MnCl₂ when it is so easy to make it from aqueous soln.?

Are you sure you're not confusing it with anhydrous aluminum chloride? This one cannot be made by heating the hydrate. Manganese chloride dehydrates at 198 °C (Lange's Handbook 15th).

clearly_not_atara - 23-12-2025 at 09:16

I think it just comes down to temperature control. If you overheat it you may get some Mn3O4 contamination. If you are able to set your hotplate to the desired temperature it might be easier.

Altreon - 23-12-2025 at 20:05

Quote: Originally posted by bnull

Quote: Originally posted by Altreon

I now ask, why does it seem that people go to great lengths to make setups for synthesizing MnCl₂ when it is so easy to make it from aqueous soln.?

Are you sure you're not confusing it with anhydrous aluminum chloride? This one cannot be made by heating the hydrate. Manganese chloride dehydrates at 198 °C (Lange's Handbook 15th).

The scimad wiki page on MnCl₂ says this:

Quote:

To obtain the anhydrous form, the compound is heated in an oxygen-less or free atmosphere, like under inert gas or an organic solvent like methanol, or in a stream of hydrogen chloride.

Both that book and the wiki seem very misleading, not just because the organic solvent method is discussed with other, much less feasible methods, but especially because there are links on the wiki for threads on the NH₄Cl method and reports by users that you cannot just thermally dehydrate MnCl₂•4H₂O without any oxidation (directly contradicting that book), making this whole procedure sound a lot more complicated than it really is.

About the temperature control, I find it hard to believe that even an alcohol lamp with manual stirring would heat the slush beyond 260°C, which I held the temperature at for 1 hour without oxidation. The methanol inhibits thr formation of any hydrolysis products, which are the main oxidation catalyst (dry MnCl₂ has no stable oxidation products), so in theory this synthesis needs very little temperature control.

(I ask this because I am concerned as to why such a simple synthesis method for anhydrous MnCl₂ has been ignored by scimad forum posts. There is surely an interest in collecting anhydrous metal chlorides for the formation of complexes, exploring the metal's chemistry, etc., along with MnCl₂ being one of the more common ways to refine battery-grade MnO₂.)

[Edited on 24-12-2025 by Altreon]

bnull - 25-12-2025 at 14:54

Quote:

The book is not misleading, as it only gives the physical properties of manganese(ii) (manganous) chloride tetrahydrate, with no mention of oxidation or hydrolysis. The first file attached below is part of a table of physical constants of inorganic compounds (table 3.2) from Lange's Handbook of Chemistry, 15th edition. The two entries inside the rectangle are for manganous chloride, the first for the anhydrous salt and the second for the tetrahydrate. If you look at the values inside ovals in the same table, you will notice that there is a "d" next to each one of them, which denotes the temperature at which the substance decomposes. There is no "d" in the rectangle and no contradiction, as you can see.

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The methanol inhibits thr formation of any hydrolysis products, which are the main oxidation catalyst (dry MnCl2 has no stable oxidation products), so in theory this synthesis needs very little temperature control.

This is wrong. Methanol simply dissolves manganous chloride.

Quote:

I ask this because I am concerned as to why such a simple synthesis method for anhydrous MnCl2 has been ignored by scimad forum posts. There is surely an interest in collecting anhydrous metal chlorides for the formation of complexes, exploring the metal's chemistry, etc., along with MnCl2 being one of the more common ways to refine battery-grade MnO2.

No need to be concerned at all. (1) Collecting anhydrous metal chlorides for formation of complexes is a niche use. You're the first member (or person, for that matter) that I see actively pursuing it. (2) You don't need anhydrous manganous chloride to explore manganese chemistry. Any soluble manganese salt will do. Manganese sulfate is sold as fertilizer and can be used after purification. With relation to non-aqueous chemistry, there is always another way. (3) You only need anhydrous manganese salts if you really need them anhydrous, if the procedure requires it. Will the water of hydration make it harder to convert them back to dioxide? No, they're sooner or later going to be dissolved in water in the process. Is the hydration water going to mess up the production of permanganate? No.

***

Anyway. Mellor, volume XII, section 17 "Manganese Chloride" (p. 348-374, second attached document; you may want to look up the publications by the cited authors given in the references at the end of the section). The second paragraph of p. 351 says

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According to E. M. Farrer and S. U. Pickering, when the tetrahydrate is heated at 200° to 230°, it loses some chlorine, while retaining a mol. of water; and F. W. Krecke also found that the tetrahydrate loses hydrogen chloride, forming a basic chloride when heated. H. M. Dawson and P. Williams observed that the dihydrate loses its water of hydration at 198°, forming the anhydrous salt.

So far, no mention of oxidation on heating. Let's skip to the only paragraph in p. 356:

Quote:

M. Berthelot observed that when the chloride is heated in oxygen, the higher manganese oxides are formed. According to
J. Meyer and R. Nerlich, manganous ions are not oxidized by atm. oxygen; nor has the presence of sodium sulphite or arsenite any influence.

Interesting. The next lines say:

Quote:

Manganous salts in ammoniacal or alkaline soln. are oxidized to manganic hydroxide; thus, manganous chloride made alkaline with ammonia and mixed with ammonium chloride is oxidized by air to manganic hydroxide. In the presence of tartaric or oxalic acid, manganous salts are not precipitated by alkalies, and autoxidation takes place rapidly. In the presence of ammonia, the product is manganic hydroxide, but in the presence of sodium or potassium hydroxide, manganese dioxide is formed, probably because in presence of these stronger alkalies the complex compound is less stable, some manganous hydroxide being always present and oxidizing directly to the dioxide.

Manganous salts oxidize above neutral pH.

Quote:

J. Davy observed that when heated to redness in air [...]

So it is decomposed when heated to redness, which is about 400 °C above the dehydration temperature. Next, the first paragraph of p. 357.

Quote:

F. Kuhlmann found that when the chloride is heated in water vapour, hydrogen chloride and manganese tritatetroxide are produced. The solvent action of water has been previously described. The taste of the salt is described by J. L. Proust as being saline, but not unpleasant. L. Bruner found the hydrolysis of soln. of manganous chloride to be too small for measurement. According to H. W. Fischer, the salt is hydrolyzed when an aq. soln. is heated to 200° in a bomb, and a dark brown precipitate is formed.

Very interesting.

Let's consult another source. Chambers and Holliday, Modern Inorganic Chemistry (1975), chapter Transition Elements, section Manganese, subsection Oxidation state +2 (p. 389-390):

Quote:

This is the most common and stable state of manganese; the five d electrons half fill the five d-orbitals, and hence any transition of d electrons in a complex of manganese(II) must involve the pairing of electrons, a process which requires energy. Hence electron transitions between the split d-orbitals are weak for manganese(II), and the colour is correspondingly pale (usually pink). The stability of the d⁵ configuration with respect to either loss or gain of electrons also means that manganese(II) salts are not easily reduced or oxidised. Indeed, in oxidation state +2, manganese shows fewer 'transition-like' characteristics than any other transition metal ion; thus the aquo-ion [Mn(H₂O)₆]²⁺ is barely acidic, allowing formation of a 'normal' carbonate MnCO₃ which is insoluble in water and occurs naturally as "manganese spar'. The aquo-ion forms typical hydrated salts, for example MnSO₄.7H₂O, MnCl₂.xH₂O and double salts, for example (NH₄)₂Mn(SO₄)₂.6H₂O; dehydration of the simple hydrated salts, by heating, produces the anhydrous salt without decomposition.

Again, manganous chloride is stable on heating.

We know now that (a) manganous chloride loses chlorine in the form of hydrogen chloride when heated in water vapor and that (b) it is easily oxidized in the presence of alkalis. What if, instead of heating the chloride in water vapor, we put it in a flask with small diameter, hence a small surface exposed to the atmosphere, and heated it? If the surface area is small, less water vapor is able to leave the bulk of the salt than if a flask with larger diameter were used. Evaporation rate is smaller in a test tube than in a Petri dish of same volume. This is essentially observation (a). What happens if the manganous chloride you prepared is neutral or slightly alkaline and you heat it? Observation (b). A solution of manganous chloride must give a slightly acidic pH.

What can we suppose now? That heating manganous chloride tetrahydrate at about 200 °C produces the anhydrous salt without decomposition, as long as (1) the salt is spread on a thin layer with good ventilation so as to facilitate the expulsion of water vapor and (2) the salt had been made with a slight excess of acid. Hydrochloric acid is volatile, so using a slight excess of it when dissolving the carbonate is no problem at all.

Merry Christmas.

Attachment: Lange's, manganese chlorides.pdf (30kB)
This file has been downloaded 74 times

Attachment: Mellor XII, Manganous chloride.pdf (1.6MB)
This file has been downloaded 103 times

Altreon - 25-12-2025 at 21:56

So it seems I'm receiving scientific enlightenment. I hope I will be able to internalize all the information I will be receiving today.
As a quick aside, I am unable to open either of your PDFs. This is a scimad-localized phenomenon. I admit I do not know anything about accessing textbooks online, and I find it impossible to check all the results from the Internet Archive copy of "Mellor."

Quote:

The book is not misleading, as it only gives the physical properties of manganese(ii) (manganous) chloride tetrahydrate, with no mention of oxidation or hydrolysis.

IMO, that is still misleading, at least in the context of this thread. Let's say, there is an amateur chemist, reading a thread on the dehydration of MnCl₂•4H₂O, and they find someone citing that it simply "decomposes" when heated. This hypothetical amateur chemist would assume that this means simple dehydration is possible, which has been shown by some other amateur chemists to be unfeasible.

Quote:

This is wrong. Methanol simply dissolves manganous chloride.

That may be the case, but this is how I think about it: MnCl₂ is in the highest possible oxidation state for a bulk manganese chloride. The only way for oxidation to happen is if the water of crystallization displaces HCl forming hydrolysis products involving oxygen, which can then be oxidized to Mn(OH)₃ and the other things we don't want. If the water of crystallization is being evaporated away with the MeOH, and the MeOH solvates the Mn²⁺ instead of H₂O, then no more oxidation can happen. That is unless Cl₂ can form, which I assume is an absurdity until the 600°C you cite.

Quote:

No need to be concerned at all. ...

Now this is where I have been enlightened. I didn't stop to consider that one of the few who care about this stuff, and that this self-centered thinking leads me nowhere and has pretty much no practical applications. Thank you for that. I do believe that this thread would at least be useful to the few people that do care about MnCl₂ for whatever other reason, like those linked on the wiki.

The rest of your post is a bunch of textbook sources. Just as in the first section you quoted, I ask, why, if the literature is so conclusive, does the scimad wiki and the scimad forums treat MnCl₂ on the same level as an actually hard-to-isolate salt like FeCl_2/3? Is it just a lack of interest? Is that enough for amateur chemists to fall under the same madness (that is, the existing MnCl₂ threads) without stopping to think about how they could make the procedure easier?

Another thing is that I personally trust first-hand accounts of the practicality and effectiveness of a procedure, more than information I have read through literature. It seems that some users in the wiki threads have tried direct thermal dehydration of MnCl₂•4H₂O, and it resulted in oxidation, while at least one seems to dehydrate it on a whim. What could be the cause of this? Are they all going mad over nothing?

To me it seems that simple dehydration of MnCl₂ is only possible in specialized laboratories. None of your sources seem to specify how they dehydrate MnCl₂, only that they can dehydrate it, and one of them groups MnCl₂ with a salt that definitely won't oxidize, MnSO₄. There's no guarantee that the sources aren't parroting each other. It's interesting that none of those sources mention MnCl₂•1H₂O and its hydrolysis, and that MnCl₂•4H₂O can melt which is implied to be a bad thing requiring temperature control to avoid in this source from 2009 making all of them seem at least somewhat unreliable.

Finally, your last suggestion would probably work, and I will not pretend to understand it completely. However, I doubt that amateur chemists who have tried dehydrating MnCl₂ don't already use an excess of HCl in preparing MnCl₂ from the hard-to-dissolve MnO₂ or MnCO₃ for the same volatility reasons. If it really was the slightly basic conditions causing oxidation of Mn²⁺, then my procedure shouldn't have worked either, because I refluxed the MnCl₂(impure) with MnCO₃ to entirely neutralize any Fe³⁺ complexes and thus indirectly removing all sources of acidity.

(belated) Merry christmas to you too.

[Edited on 26-12-2025 by Altreon]

Failure to purify MnCl2

Altreon - 25-12-2025 at 22:44

I repeated the procedure again, with twice the amount of MnCO₃. I added enough HCl to make the solution a somewhat clear brown color. I then added 5g MnCO₃, along with 50% H₂O₂. This is where the procedure deviates. No matter how long I refluxed, even after filtering and adding more MnCO₃ (totalling over 10g) or more H₂O₂, the solution was much more difficult to filter and the red(!) filtrate always had some of the brown MnCO₃. The last time I did this, the filtrate was very pure and easy to refilter. I tried refluxing with extra water, hoping to force hydrolysis, and with less water, hoping to "catch" more of the Fe³⁺, but I was eventually left with a solution that was very clearly not a pure pink (peach? it looks orange on camera). After boiling down the solution and refluxing with 2g MnCO₃ in a last ditch effort, the filtrate went back to being red with particles going through. I have no idea what I did wrong this time. The pH of the soln. stayed at 3-6 for a long time, meaning I probably didn't reflux the soln. enough, but the really red filtrate is neutral. The color was really unlike Mn²⁺ or Fe³⁺.

I added drops of conc. NH₃ at this point, creating a white ppt. that slowly turned brown, and turned very brown with H₂O₂. This stuff filtered into a dark brownish red soln., which is probably a colloidal Fe(OH)₃.

Attached images are of the cleanest the filtrate got, and the filtrate just before and after I added ammonia.

[Edited on 26-12-2025 by Altreon]

bnull - 26-12-2025 at 04:45

I think I see now what you mean by oxidation. Oxidation as in formation of an oxide, not oxidation as the state going from +2 to +3 or +4. Manganese still manganous but with an oxygen attached to it.

Quote:

As a quick aside, I am unable to open either of your PDFs. This is a scimad-localized phenomenon.

I uploaded them to a temporary repository (https://limewire.com/d/9kNr5#r6abBWhu9X). The link expires in a week. If you still can't open them, then there is something wrong with your browser or computer settings.

Quote:

IMO, that is still misleading, at least in the context of this thread.

If I told you that today is Friday, would that be misleading if I didn't tell you which day of the month is it? The table does the same. It tells that the salt dehydrates at 198 °C but does not give information as to how it should be heated, for how long, if in air or inert atmosphere, and what color it changes to. Why? Because it is a table of physical constants, for when you need to consult melting points, density, molar mass etc. of the substance you have at hand.

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Specialized stuff pops up here a few times a year and that's perfectly normal. It doesn't mean that no one else cares and you're losing your time, only that the number of amateurs interested in producing anhydrous metal chlorides is small; it is smaller than those looking for dichromate or organic dyes, anyway. My first post in the forum was a question about the existence and production of free allophanic acid (no and no, only salts and esters exist, and most of the information comes from a hard-to-find paper from the Soviet Union). If you can find an even rarer topic than allophanic acid, be my guest.

Quote:

The rest of your post is a bunch of textbook sources.

Mellor's Comprehensive Treatise isn't your run-of-the-mill textbook source. If you look closely, Mellor lists all the observations he could find, even when one observation contradicted another, and includes the original sources he used in the references. Comprehensive. The references for manganous chloride fill up 4 pages of small type. Can be a bother to find the publication you want but it is a fair trade-off.

Quote:

None of your sources seem to specify how they dehydrate MnCl2, only that they can dehydrate it

Which is why you need to check the references in Mellor. The details are there. Listed below are those I managed to find online.

K. M. Farrer and S. U. Pickering, Chem. News, 53. 279, 1886 (https://archive.org/details/s713id13691550/page/278/mode/2up)
F. W. Krecke, Journ. prakt Chem , (2), 5. 105, 1872 (https://advanced.onlinelibrary.wiley.com/share/2HH2UIZXMGRR2...)
H. M. Dawson and P. Williams, Journ. Phys. Chem., 4. 370, 1900 (https://babel.hathitrust.org/cgi/pt?id=uva.x002457627&se...)
J. Meyer and R. Nerlich. Zeit. anorg. Chem., 116. 117, 1921 (https://onlinelibrary.wiley.com/share/U72BHM8RX8IPIRFMBKPE?t...)
J. Davy, Phil. Trans., 102. 169, 1812 (https://royalsocietypublishing.org/rstl/article/doi/10.1098/...)

Altreon - 26-12-2025 at 06:05

I am now able to access the PDFs with the temporary link you have graciously provided me. The Mellor pdf confirms one of my suspicions, that I somehow magically managed to keep Mn³⁺ in soln., and that its solution in the presence of MnCl₂ is somehow extremely stable. It's also interesting that Mellor thinks that the monohydrate is a misprint and does not exist as a well defined hydrate, when the 2009 paper I linked describes it and its dehydrating capabilities.

I do not know what I said to imply I meant hydrolysis with the word "oxidation." I apologize for the confusion, but I am indeed referring to higher manganese oxyhydroxides forming (MnO isn't stable in air at >200°C anyways)

Quote:

If I told you that today is Friday ...

I see what you mean. There's no reason for me to continue debating on this point.

I do not know why you took the time to provide me with not only the sources, but direct links to the sources, but I gladly appreciate that. A troubling detail, for me, is that the 1st source describes the NH₄Cl method and describes the aerial oxidation of MnCl₂•1H₂O when heated, and its apparent robustness beyond 230°C. Both the 2nd and 4th sources are in German and I am currently in a hurry. The 3rd source appears irrelevant. The 5th source is absolutely ancient and they still refer to chlorine as oxymuriatic acid.

These sources seem to confirm the troubling procedures on the scimad wiki rather than your previous sources describing simple dehydration of MnCl₂, making anhydrous MnCl₂ again look as unreachable as FeCl₂. The narrative sources in Mellor seem to be at least somewhat contradictory to these sources, since Mellor seems to be under the simple-dehydration-is-possible camp.

When I successfully repeat my procedure, we will see if the monohydrate is robust enough to avoid displacement by methanol, and if I actually synthesized what I wanted.

Altreon - 26-12-2025 at 08:42

I have successfully synthesized 13.5g of what may or may not be anhydrous MnCl₂, accompanied with visually severe oxidation, despite my attempts. I will explain in detail tomorrow.

Altreon - 27-12-2025 at 06:17

After I repeated the procedure from my first post here, it seems that MnCl₂ shares a property with CoCl₂, CuCl₂, and possibly all transition metal chlorides where the first step in dehydration turns the substance rock-hard. For MnCl₂, it seems that resulted in very visible amounts of oxidation products forming at the bottom of the beaker. After breaking up the rock and spreading it over a wider beaker, I managed to stop all oxidation even when heating the pwdr. to 280°C, leaving me with an off-white slightly pinkish pwdr. If I spread the pwdr. out initially (which appears to be @bnull's suggestion, which I now understand) then maybe oxidation would have lessened.

The fact that, even with the methanol, the pwdr. still oxidized, makes me willing to believe that the literature must be doing something quite unusual (i.e. the dry HCl and NH₄Cl methods rather than simple dehydration), and that it is probably a futile task to dehydrate MnCl₂•4H₂O without significant oxidation if even the MnCl₂•xMeOH•xH₂O adduct cannot stop the hydrolysis (which is equivalent to oxidation) completely. Another observation is that MnCl₂(aq.) added to an excess of MeOH causes a faint white gel to ppt. at r.t., indicating that MnCl₂ is so much more willing to coordinate with H₂O that it wouldn't even let itself get solvated by the MeOH, and that the MeOH additions probably don't completely dissociate the MnCl₂ aquo complex, leading to the unfortunate oxidation of my product.

Finally, my last failure was due to insufficient H₂O₂. It seems that trace Fe²⁺ is enough to slowly turn the solution acidic and very hard to filter (a pH paper displaying neutral pH after being dipped in the soln. became red after an hour).

I have analyzed the 2nd source you have graciously provided in further detail. It seems to suggest that the green complex that I've been associating with methanol all this time is actually a thermochromic aquo complex involved with the decomposition(!) of MnCl₂•4H₂O that apparently turns "rose-red" or "yellow" under some* heating conditions and can be formed by ethanolic solns(!). I have never observed any of these colors and crystals, so I cannot say if this article is accurate, but it also supports that MnCl₂ cannot just be heated to dehydration, further proving the side of fellow amateur chemists rather than Mellor and the rest of the old literature's descriptions. I am unable to translate the 4th source in full.

The attached images are of the oxidized MnCl₂ and the final product.

Altreon - 27-12-2025 at 06:20

(On the topic of obscure chemistry, I once wanted to synthesize sodium methylcarbamate from carbaryl, but there wasn't enough info online for me to proceed.)

bnull - 27-12-2025 at 16:28

Disclaimer: what follows is a conjecture based on my research of the literature related to the degradation of the hydrates of manganese(ii) chloride, henceforth called manganous chloride because this is how it is called in the older literature and it is such a nice name.

The production of anhydrous manganous chloride by heating the hydrated salt seems to be plagued by mixed results. Either it works smoothly and the anhydrous salt shows no signs of degradation or the product is contaminated by basic chloride and higher oxides of manganese. The quotations and sources given on previous posts present the same conflicting results, and members of the forum give anecdotal evidence for both outcomes, which makes me wonder: what are we missing?

I have read the posts from the manganous chloride thread and the sources given above. What I don't see is a precise description of the procedure adopted by the authors, apart from temperature and time. This lack of information is what makes the problem more mysterious than it deserves.

We know that (1) manganous chloride is prone to oxidation in basic media, (2) it reacts with water vapor, producing hydrogen chloride and a basic salt that may be subsequently oxidized, and (3) that it is supposed to become anhydrous at 198 °C. Assuming that the salt had been prepared by dissolving the carbonate in excess hydrochloric acid, (1) is out of the question, leaving us with observations (2) and (3).

Consider this. A portion of manganous chloride tetrahydrate is added to a clean beaker and submitted to heating at a suitable rate so that the final temperature is about 200 °C. As the salt loses water, it agreggates into a solid block. Water vapor that is liberated from the lower portions reacts with the salt. The salt further degrades as heating continues. This doesn't happen with other common hydrates used as desiccating agents. Magnesium sulfate, sodium sulfate, and copper sulfate dehydrate without tarnishing or degrading.

The solution is simple. Hydrated manganous chloride must be well ground and spread on a thin layer on the heating surface, and a means to remove water vapor from the proximity of the salt as soon as it is produced should be employed. This should reduce the time available for the reaction between water vapor and manganous chloride. The optimal thickness of the layer is probably related to the surface area and rate of air circulation.

Edit: I may be utterly wrong, of course.

[Edited on 28-12-2025 by bnull]

Altreon - 28-12-2025 at 02:23

Quote:

This doesn't happen with other common hydrates used as desiccating agents. Magnesium sulfate, sodium sulfate, and copper sulfate dehydrate without tarnishing or degrading.

I hypothesize that this problem occurs only with simple transition metal halides because of the volatility of HX at dehydration temperatures forcing an equilibrium with hydrolysis, coupled with irreversible oxidation (which does not occur with those metals).

Quote:

The solution is simple. Hydrated manganous chloride must be well ground and spread on a thin layer on the heating surface, and a means to remove water vapor from the proximity of the salt as soon as it is produced should be employed. This should reduce the time available for the reaction between water vapor and manganous chloride. The optimal thickness of the layer is probably related to the surface area and rate of air circulation.

Something to note is that MnCl₂•(2?<x≤4)H₂O is molten at its decomposition temperature until it resolidifies into MnCl₂•(0≤x≤2?), meaning one would have to first pour the molten tetrahydrate onto a large tray and grind the partially dehydrated material afterwards. Once the salt goes from a slurry to a dry rock, it can be broken down into a slightly clumpy pink powder which could probably be heated in a regular beaker without oxidation, as I have shown previously.

I realize this method sounds a lot like the oven-drying method for preparing anhydrous CaCl₂, which might not be coincidental. While I'm not particularly fond of the idea of drying manganese salts in an oven, someone who views this thread in the future could try to dehydrate pure MnCl₂•4H₂O in a large borosilicate dish and report their results.

[Edited on 28-12-2025 by Altreon]

bnull - 28-12-2025 at 10:55

Quote:

Something to note is that MnCl2•(2?<x≤4)H2O is molten at its decomposition temperature until it resolidifies into MnCl2•(0≤x≤2?), meaning one would have to first pour the molten tetrahydrate onto a large tray and grind the partially dehydrated material afterwards.

I forgot that. Good to see we got somewhere after all.

Do you have access to the supplementary information of the Polyachenok paper?

semiconductive - 29-12-2025 at 15:02

Quote:

That may be the case, but this is how I think about it: MnCl2 is in the highest possible oxidation state for a bulk manganese chloride. The only way for oxidation to happen is if the water of crystallization displaces HCl forming hydrolysis products involving oxygen, which can then be oxidized to Mn(OH)3 and the other things we don't want. If the water of crystallization is being evaporated away with the MeOH, and the MeOH solvates the Mn2+ instead of H2O, then no more oxidation can happen. That is unless Cl2 can form, which I assume is an absurdity until the 600°C you cite.

I'm trying to follow your reasoning and grasp what assumptions you are making:

Mn is one periodic table column less than Iron.
Therefore: I would expect Mn to be very similar to Fe.
When I look at possible Manganeese oxidation states, +2 to +7 are listed.
When I look at possible Iron oxidations states, I see +2 to +6 are listed.
They have only one oxidation state difference +6 vs. +7.

Mn is more willing to be oxidized than iron.

The most stable Mn oxidation state according to web pages I see, is +2.
(I'm not sure why this is.)

But:
When you talk about +2 being in the highest oxidation state for 'bulk' Manganese chloride; I don't understand why MnCl₃ can't form. FeCl₃ does form with Iron.

Since Manganese is more oxidize-able (easily able to lose electrons) than Iron, and +3 is a common state for Mn, why don't you expect a mixture of +2 and +3 Manganese oxidation states that depending on the relative stability of Manganous vs. Manganic chloride?

eg: Why ought we think that +2 is the 'highest' for 'bulk' Manganese Chloride?

The reason I ask is because you say "only" in the sentence that follows. I (usually) fail when I try to reason that way. Exceptions to rules seem to be as common as rules in chemistry ... ( Which is why I find it so frustrating. )

I know:

Mn, electron configuration: [Ar] 3d⁵ 4s²
Fe, electron configuration: [Ar] 3d⁶ 4s²

In either case, if the atom is ionized by at least +2, then both '4s' electrons go away, and we are left with either 5 or 6 'd' electrons. All 'd' electrons are roughly equal energy until we define a ligand (H₂O, CN, ?? R-OH ??) that it interacts with the 'd' electrons.

At that point, there are three possible choices for coordination complexes.

Planar, Octagonal, or Tetrahedral.

Why there aren't more choices, I am not certain. But, I think it has something to do with coordination complexes usually (or always) having 'sigma' bonds rather than distributed 'pi' bonds. ( I could be wrong, too. )

[Edited on 30-12-2025 by semiconductive]

DraconicAcid - 30-12-2025 at 10:51

Manganese(III) simply isn't stable without the right ligands. Cotton&Wilkinson say that MnF3 is highly reactive (easily hydrolyzed) and MnCl3 decomposes at -40 C (but can be kept in ethereal solution up to -10 C).

Complexes with oxalate, EDTA, acac, or sulphate are relatively stable.

semiconductive - 30-12-2025 at 13:13

Quote:

Cotton&Wilkinson say that MnF3 is highly reactive (easily hydrolyzed) and MnCl3 decomposes at -40 C (but can be kept in ethereal solution up to -10 C).

hmm...
Maybe I need to understand a vocabulary difference between stable and at-equilibrium, or Hydrolyze vs. Ionize.

As a rule, I know:
In water, a weak acid and a conjugate base have ionization constants whose product is the same as autoionization of water, Kw.

Ka = [H₃O⁺][X⁻] / [ HX ]
Kb = [ Hx ][ OH⁻ ]/ [ X⁻ ]

Ka·Kb = Kw

Kw changes with temperature ; but it has a definite equilbrium (non-zero) value at every temperature.

Is there a difference between having a very low equilibrium constant, and being 'unstable' ?

Based on general rules; I think that chloride attacking Mn is easier than Cl attacking Fe because periodic trends increase electron holding strength when moving along periodic columns to the right, and then decrease at the noble gas.

I expect -- Mn should be less electronegative than Fe, but Chlorine will have the same strength; hence I reason Cl ought to attack Mn faster/more strongly than Fe and the reaction to +3 itself ought to be possible/happen.

But, if MnCl₃, is 'unstable', does that mean that whenever chloride does oxidize Mn to +3 that MnCl₃ immediately reacts again to form MnCl₂[OH] or a different chemical that isn't reversible?

Or is Mn just a bizarre exception,

Does it take more energy to oxidize Mn to the +3 state using Chlorine even though it's periodic trend ought to make it weaker than Fe ?

Thanks.

DraconicAcid - 30-12-2025 at 13:25

Yes, if we look at trends, we would expect that MnCl3 would be perfectly stable, because we see with CrCl3, FeCl3, CoCl3 and NiCl3 vs CrCl2, FeCl2, CoCl2, and NiCl2, the stability of M(III) vs M(II) decreases going left to right. Manganese just doesn't care about that trend, and I don't have a good explanation for it.

MnF3 hydrolyzes readily, meaning it reacts with water to give MnF2(OH) and HF (and then presumably keeps going).

MnCl3 decomposes; 2 MnCl3(s) == 2 MnCl2(s) + Cl2(g) has a high equilibrium constant above -40 C. It doesn't need water to do so.

If you can find a copy of Cotton & Wilkinson's Advanced Inorganic Chemistry, I highly recommend it.

bnull - 30-12-2025 at 14:16

It's got to do with the electronic configuration of the ions. The electronic configuration of a neutral manganese atom is [Ar] 3d5 4s2; seven electrons in the valence subshell (or orbital; I forgot how they are called today). Manganese(ii) is [Ar] 3d5. Hence, the 3d subshell is half-filled and stable and the subshell 4s is empty, making Mn²⁺ very stable. From Modern Inorganic Chemistry again (https://archive.org/details/Modern_Inorganic_Chemistry/page/...):

Quote:

(t)his is the most common and stable state of manganese; the five d electrons half fill the five d-orbitals, and hence any transition of d electrons in a complex of manganese(II) must involve the pairing of electrons, a process which requires energy. Hence electron transitions between the split d-orbitals are weak for manganese(II), and the colour is correspondingly pale (usually pink). The stability of the d5 configuration with respect to either loss or gain of electrons also means that manganese(II) salts are not easily reduced or oxidised.

So far, so good.

Manganese(iii) is [Ar] 3d4. The 4s subshell is empty as before, but the 3d subshell has now four electrons instead of five and is not as stable as 3d5. It needs to gain one electron to become manganese(ii). Apparently, the ion that most easily donates one electron to Mn³⁺ is another Mn³⁺, so one ion becomes Mn²⁺ and the other becomes Mn⁴⁺, which ends up as the dioxide. Everything is stable once more.

semiconductive - 30-12-2025 at 18:22

That's a helpful clue!

Let's see: I need to consider 5 electrons, 5 d orbitals.
I'll ignore Square Planar geometry because that requires 8 electrons 5 d orbitals.

That leaves Octahedral and Tetrahedral shapes.
In either case, there are only two energy levels with five orbitals.

Octahedral:
_ _ x²+y², z²
_ _ _ xy, xz, yz

Tetrahedral;
_ _ _ ???? (I don't remember names.)
_ _

Octahedral is more likely (stable) since there are more low energy levels present.

Without explanation:
Kotz and Purcell, show "typical" Mn²⁺ and Fe³⁺ examples as high spin octahedral, which means one electron fills every orbital before pairing begins.

In such a situation, electrons will not be able to transition normally because they would have to 'spin flip' in order to jump from either high or low energy band to the other:

AKA: There are two abnormal energy barriers involved, 1 -- spin flipping, 2 -- pairing.
This makes both the color-less-ness (pale-ness) and the chemical stability of Mn reasonable to understand.

Minor detail --- doesn't filling in the orbitals in high-spin vs. low-spin order (pairing) actually depend on the ligand present ?

With Iron, I've tried Me-O-H, H-C-N, Et-O-H, and (of course) H₂O is always causing me trouble.

Ferrocyanide is pale yellow, Ferricyanide is dark reddish (I haven't made it myself, yet), but these are both diamagnetic, therefore low-spin. Which means they are A-typical ?

[ Fe(CN)₆ ]⁴⁻ 4K⁺

Side Note: I have a Jaz™ Spectrometer, although I need some help calibrating it since the Ocean Insights software won't run on my Raspberry PI. eg: I would happily make some color measurements with a bit of help, and I don't mind paying someone to help me solve the math in a way that benefits the whole scientific community by making calculations numerically accurate (Hint: They aren't right now!!!):

https://physicsdiscussionforum.org/integration-of-planck-s-b... )

But: In this thread we have H₂O and Me-O-H competing.
I see in the first posts that pictures show mostly pinkish and clear colors.

Therefore:
The most helpful clue I missed is the near color-less ness of the solution.
I'm not sure how I would have judged "pink" as the same as colorless, but there was also some iron contamination which initially caused color to appear;

Altreon chose to get rid of Fe using oxygen, which is more than I would have known to do as an amateur.

Good job!

DraconicAcid - 30-12-2025 at 19:34

Quote: Originally posted by semiconductive

Minor detail --- doesn't filling in the orbitals in high-spin vs. low-spin order (pairing) actually depend on the ligand present ?

With Iron, I've tried Me-O-H, H-C-N, Et-O-H, and (of course) H₂O is always causing me trouble.

Ferrocyanide is pale yellow, Ferricyanide is dark reddish (I haven't made it myself, yet), but these are both diamagnetic, therefore low-spin. Which means they are A-typical ?

Yes- cyanide is very high on the spectrochemical series, meaning it causes a large difference in energy between the stabilized and destabilized d orbitals.
https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystal_Field_Theory/Crystal_Field _Theory

Most Fe(II) complexes will be high spin and paramagnetic; [Fe(bipy)3](2+) and [Fe(CN)6](4-) are low-spin and diamagnetic. For similar reasons, most NiX4(2-) complexes are tetrahedral; the tetracyano complex is square planar.

High-spin Mn(II) complexes are pale in colour because their d-d transitions are technically symmetry-forbidden, and only happen when the complex is distorted, IIRC.

If you want to calibrate a spectrometer, you'll probably want to make and do measurements on some lanthanide complexes. Transitions between d orbitals vary greatly depending on the metal-ligand distances (which are constantly changing due to vibrations), so regular transition metals give broad peaks in the UV-vis spectrum. Splitting between f orbitals doesn't change that much, which means that the peaks are much sharper (and that different complexes of the same metal won't have dramatically different colours).

semiconductive - 30-12-2025 at 22:38

I'm open to all kinds of calibration techniques; Though I don't want to hijack this thread;
Do lanthanide standards of color come in pre-sealed optical cells that won't spoil over time?

Note: Lasers and crystals for Doppler/Raman scattering are very sensitive to temperature changes in the detector ( on top of being expensive ). Without the commerical software working perfectly, I'm not going to be able to use it. FTIR is kind of the same thing. It's not really complicated, electronically, but it's expensive with easily destroyed parts that aren't very practical for an amateur/hobbiest.

I figure I best start by being able to characterize a plain (low cost) optical spectrometer for temperature changes in it's sensor and in the liquid being analyzed.

If I had an accurate way to integrate the Planck distribution (quickly) then I could do simple experiments like placing ligands in an oscillating magnetic field in the low kilohertz through megahertz frequencies; Magnetic fields can induce resonant stresses on paramagnetic ligand bonds and I would only need to fight Planck distribution noise / thermal effects to extract useful data.

Being able to compute the Planck distribution accurately (and fast) is what I need to get beyond my present roadblock. But -- commercial tools like Mathematica, vs. others, give me contradictory answers for numeric intergrals.

But, I'm not a mathematician, and got frustrated when my attempt at algebraically approximate the Planck integral partially failed a few years ago. I don't know exactly what I did wrong; and NASA's software to do the same job is only 1% accurate, (which is garbage.) I'm used to 0.01% accuracy in electronics and don't think 1% accuracy is enough to get much useful chemical data with.

MnCl2

Altreon - 8-1-2026 at 19:10

Well, I didn't expect this thread to get as much attention as it did. Not that anyone is interested in the actual MnCl₂, rather it seems to have spiralled into some theoretical discussion that I am too ignorant to participate in.

To bnull: I can't seem to find any links to the supplementary paper, so I can't provide it. I'd like to try out this synthesis of MnCl₂ myself, but I am currently trying to obtain a 10cm borosilicate dish for this process. I will be preparing ≈30g of anhydrous MnCl₂ with excess HCl.

I've recently been working on anhydrous NiCl₂ as well, and it seems that both compounds I've prepared produce a cloudy suspension in MeOH despite them dissolving completely in water. NiCl₂ only produces a faint green color in the organic solvents. Does anyone have a pure sample of either compound who can observe their solubility in organics? AFAIK they are similarly covalent to CoCl₂ and CuCl₂, which both dissolve easily and form strongly colored complexes with alcohols, so NiCl₂ and MnCl₂ should behave similarly.

[Edited on 9-1-2026 by Altreon]

bnull - 8-1-2026 at 19:57

Quote:

To bnull: I can't seem to find any links to the supplementary paper, so I can't provide it. I'd like to try out this synthesis of MnCl₂ myself, but I am currently trying to obtain a 10cm borosilicate dish for this process. I will be preparing ≈30g of anhydrous MnCl₂ with excess HCl.

I'm not surprised. Odds are that it wouldn't have enough details on the process anyway. Thanks for trying.

[Edited on 9-1-2026 by bnull]

Fery - 8-1-2026 at 21:26

Quote:

Manganese(II) chloride is commercially available as MnCl2 . 4 H2O containing 36.4% water. Under an argon atmosphere most of the water is lost at 350 C, but absolutely anhydrous material is only obtained after prolonged heating at 400 C.
Our method is as follows: Finely ground MnCl2 . 4 H2O (50 g) was heated in vacuo at 200 C for 6 h, the cake finely ground, and then further heated in
vacuo at 220 C for 6 h.

https://sci-hub.st/10.1016/S0020-1693(00)90464-X

Altreon - 9-1-2026 at 00:02

Quote:

Manganese(II) chloride is commercially available as MnCl2 . 4 H2O containing 36.4% water. Under an argon atmosphere most of the water is lost at 350 C, but absolutely anhydrous material is only obtained after prolonged heating at 400 C.
Our method is as follows: Finely ground MnCl2 . 4 H2O (50 g) was heated in vacuo at 200 C for 6 h, the cake finely ground, and then further heated in
vacuo at 220 C for 6 h.

This source seems quite dubious. Not only does it not mention hydrolysis in any way, despite the absurd temperatures, it seemingly implies the existence of a hemihydrate, which has not been supported by any of the sources quoted in this thread. Now, I must admit that I cannot confirm or deny a hemihydrate existing, as I cannot accurately measure yield because of the MnCO₃ used for purification. It may be the case that my product is a hemihydrate, but I highly doubt that these guys would happen to be the only ones who would mention a dehydration issue this severe.

bnull - 10-1-2026 at 05:56

Have you checked parts I through VI of the same series?

Altreon - 12-1-2026 at 16:35

I have, and they seem to use harsher dehydration conditions as the articles progress (from 110-130°C 8h in III, to 100-200°C 6h in V, and finally to 200-220°C 12h in VII, all under vacuum)
I'm not sure why they seem to be the only ones having such a hard time dehydrating MnCl₂, even compared to other sources already mentioned here which just use dry HCl at regular atmospheric pressure in some way.

VII mentions

Quote:

Painstaking investigations of a large number of dehydration-without-decomposition methods, which will be the subject of another publication, ...

but I cannot seem to find that article.

[Edited on 13-1-2026 by Altreon]

bnull - 13-1-2026 at 03:26

The series goes up to part 14 at least, https://doi.org/10.1039/DT9850000135.