The Mad Plater
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The Woes of Manganese Phosphating
Just yesterday I picked up my recent order of manganese carbonate and nickel nitrate, so I decided to give it a shot today, with a small scale test
(100mL).
That, instead of the planned electroless nickel tests - which would require a solid 8-10 hours to complete, and I just didn't have that much time to
spare today.
So, I went to the cupboard to grab the needed reagents and hardware - and I already knew there was trouble brewing, when I grabbed one of the brand
new 100mL beakers, straight out of the box, and found that it had a small crack originating from its rim.
BRAND. NEW. BEAKER. 
Sure, it was less than $2, but it still makes me angry.
A few minutes of torching later, the crack was esentially completely gone - but a small "scratch" remained on one side of the glass (just under the
curled rim), which I somehow missed while torching it, thus failing to close it up fully. Argh... gotta remelt that again later.
All right, time for the first attempt at mixing up the bath - doomed before even beginning, by a mathematical error, brought about by haste on my
part. (of course I didn't realize it until AFTER everything went wrong )
So, in go the reagents. Already, the wheels began to come off...
WTF? Why is the (brand new, freshly opened) nickel nitrate soaking wet? Well, that wouldn't be a problem in itself, but how am I supposed to "pour" a
tiny amount of that wet, aggregated lump in a controlled manner?
Also, why is this mixture so vividly green? The ready-to-work bath is supposed to end up a light pink color - can't see that possibly happening with
this...
Incidentally, yes, this is where I began to suspect that something might be amiss.
Adding the other reagents went on uneventfully - in went the phosphoric and citric acids, the iron powder, and the silt-like manganese carbonate - and
then it was time to heat up the mixture, to dissolve the iron and the manganese carbonate.
It was promising at first - there was some serious fizzing going on - but then it died down, despite bringing things to a roling boil.
Instead of the expected clear, light pink solution, it was a turbid green mess. That can't be right...
I stopped the heating, and after a while, the unreacted crap has settled down.
So, I was thus left with a clear green solution, with a bunch of sediment-shaped sediment at the bottom.
pH of ~6... what? It's supposed to be <3... Back to the drawing board then.
After rechecking my calculations against the (very sparse, and clear as mud )
source material, I found the error... it was supposed (???) to be ~1.6g/L of nickel nitrate, not 25g/L. No wonder that everything went wrong 
OK, time for take two.
This time, adding the manganese carbonate last, AFTER titrating for free acid. Also, using the correct amount of nickel nitrate.
The first stupid thing I've done this time, was adding the fine mesh iron powder to the neat 75% phosphoric acid on the bottom of the beaker.
I have no idea what the gaseous products of that reaction were, but they sure didn't smell healthy
All right then, time to dissolve the iron powder - all 0.25g of it. Several minutes of boiling later, and there's still a small amount (maybe 1-2%) of
unreacted iron left, as well a small quantity of some new, unidentified precipitate. (does life have to be so hard?)
Took a 10mL sample and titrated for free and total acid.
As expected, way too much free acid, so added the correct amount of manganese carbonate... well, it turns out that my previous initial estimate (for
that first test) was WAY off.
Several minutes of boiling and stirring later, and the (finely powdered!) manganese carbonate still hasn't fully dissolved. Really recalcitrant stuff,
that.
Also, bumping. Lots of it, because of the as-yet-unreacted powder collecting on the bottom of the beaker. The tiny beaker was doing quite a dance on
the improvised "hotplate" (a thick aluminium plate on top of a stove burner).
The color has changed, not to pink, but to a dirty yellow-ish with a hint of green. Still too much nickel nitrate?
Screw that, I don't have the time for this. In goes the test specimen (an old bearing roller).
OK, this looks promising - at least it started to blacken rather quickly.
Removed the specimen after half the "normal" processing time and miked it. Well, a partial success?
Good news: there is now a nice, very dark grey (near black) coating on the roller, as there should be.
Less good news: instead of increasing in size slightly (~20-30um), it decreased by 3um.
Almost certainly too much free acid - but despite that, THE DAMN CARBONATE JUST WON'T REACT!
Tried to continue the processing, just to see if the trend continues, but the bumping became even more extreme, ejecting much of the beaker's
remaining contents, and I decided to call it a day.
There's now some weird, dark gunge on the bottom of the beaker, and also the solution is an odd, dirty color - that's not supposed to happen, at least
not nearly this quickly
Lessons learned:
- use a MUCH bigger beaker next time, at least 250mL,
- CHECK AND RECHECK THE MATH!
- do not skip critical steps of the process?
- try making manganese citrate beforehand, and adding THAT to the bath, instead hoping for a reaction-in-place?
- try less nickel nitrate maybe?
- try adding the iron powder last?
- in any case, do not add the iron powder to neat acids!
Also, just for the lolz - earlier today, I was looking through the available glassware, and found... a 25mL, narrow neck erlenmeyer flask.
SO TINY
Laughed my ass off at that one. Seriously, WTF?
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The Mad Plater
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Manganese citrate
For today's experiment, I decided to investigate the production and solubility of manganese citrate, a rather obscure compound.
Can't buy that little bugger - even the big shots such as Aldrich don't sell it, for whatever reason.
Annoyingly, it also seems to be extremely poorly characterized in the literature - good luck finding any info on it, other than the absolute bare-bones essentials: formula, MW, appearance.
Certainly NOT any such "irrelevant" crap such as, say, solubility information. Nope. Nada.
I need to know the solubility, in order to successfully formulate my manganese phosphating bath concentrate.
OK, next problem - which manganese citrate? (II) or (III)?
In my reference material, it's referred to simply as "manganese citrate" throughout.
Since that text is rather dated, I presume that this is the old naming convention, and it means manganese(III) citrate then?
Well, then - I thought that I'm going to be clever, and use a quantitative approach, to verify if the expected reaction has occured, based on the
amount of mass lost to the gaseous reaction products.
(spoiler alert: it did not go as planned)
First attempt, on a test tube scale - I greatly underestimated the reaction rate at RT, based on the reaction rate of citric acid and
limescale as a reference point. (perhaps because limescale rather rarely appears as a finely divided powder?)
Stoich. mix of citric acid monohydrate + manganese carbonate, total of 1.5g.
Added all the water (~3ml) at once, and started heating right away. Result - within seconds, a column of foam extruding out of the test tube, carrying
much of the unreacted suspension with it. (even though the test tube only got
barely lukewarm...)
MAXIMUM FAILURE!
Second attempt - same as above, but no heating, and adding water dropwise. Uneventful, but painfully slow.
OK, reaction is over, test tube now full of highly porous mass of wet "cake".
Slight pink-ish color, but not much different than the starting carbonate, maybe a hint less brown.
Tried drying it over a flame (for weighing) - that went poorly. A few % of the product decomposed from the excessive heat.
In the end, I got ~1.45g - which seems to indicate a highly hydrated form of manganese citrate; anhydrous would have weighed ~1.1g.
Tried dissolving the product in DI water - well, it appeared roughly about as soluble as, say, sand or gravel. Which is to say, not at all.
Certainly less soluble than a few grams per 100mL, which is a total disaster.
Even worse, the fine dusty form of the product made a highly turbid, milky-looking suspension, which refused to settle in the slightest, even when
left undisturbed for over an hour.
But then, a wild idea appeared!
Well, I don't really care about the solubility of the product in pure water - what really matters is the solubility in highly acidic
solutions...
OK, added about 2-3mL of 75% orthophosphoric acid (which will be the primary component of the concentrate) - and in several seconds, the formerly
turbid suspension became completely clear. Success 
Unfortunately, in the end, it was only a qualitative, not quantitative analysis, so I haven't exactly narrowed down the solubility - but the results
were sufficiently promising, that I now feel comfortable attempting to concoct the bath concentrate on an 80mL scale; enough for 1L of the phosphating
bath. 
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aga
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Sounds interesting.
Can you please share a link to those references so we can all have a try ? Even some photos if it's a book or other printed material.
I got a lemon tree and some manganese metal ...
Also a few 100g of citric acid which might make it a bit quicker.
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The Mad Plater
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Quote: Originally posted by aga  | | Can you please share a link to those references so we can all have a try ? Even some photos if it's a book or other printed material.
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Sure, no problem.
There's just one tiny little catch - my main source material pertains to utilizing the abovementioned compound, not to producing it.
Which is why I had to do this the hard way.
Let's see...
R-TR-75-034, "Improved manganese phosphate coatings";
RE-TR-71-60, "A study of manganese phosphating reactions".
Both of the above are old, unclassified military publications. Google is your friend.
Then there is this:
Insight into the degradation of a manganese(III)-citrate complex in aqueous solutions (clicky)
That paper above is pretty much the most comprehensive source of information on these blasted compounds that I was able to locate.
Anything more than that, try google. Good luck with that, though - you'll need it.
Either I'm a complete and total moron, or this is a really obscure class of compounds.
Why would you want to make this, though? Interested in doing some manganese phosphating? (if so, search also for MIL-P-50002B, and MIL-HDBK-205A -
mandatory reading!)
AFAIK, there are no other significant uses for these compounds - industrial, as well as R&D - which is probably why there are no suppliers...
If you really want to try this, starting from raw manganese metal, I'd recommend first dissolving it in dilute acid (nitric should work), and adding
an (excess of?) alkali metal carbonate to create an insoluble precipitate of manganese carbonate, which is then amenable to subsequent workup.
(note: I haven't actually done any of the above listed prep work - I just cheated, and used pure, reagent grade manganese carbonate)
From there, just mix it with an aqueous solution of citric acid (1:1 molar ratio), and we're back in familiar territory.
After my previous "experiences" with this, I'd recommend adding the carbonate small amounts at a time, or - for a "set it and forget it" kind of setup
- use a dropping funnel for the citric acid solution.
Preferably, perform the reaction in a beaker - or even better, in an evaporating dish.
Using a flask of any kind is ill-advised.
In any case, use a generously oversized container.
No stirring is generally needed, at least until the gas production abates.
Note that the product will be very poorly soluble in water, unless the solution is first made acidic. The simplest way to achieve this, would be to
add an excess of citric acid.
Apparently, there are other synthetic routes possible, but if you want to do manganese phosphating, this is the way to go - manganese carbonate is a
required bath ingredient, no matter how you slice it. (also, manganese nitrate, in considerably smaller amounts)
In any case, I believe that I'm thoroughly done with this whole manganese citrate monkey business.
I already know what I need to create my manganese phosphating solutions, and the other physical/chemical properties are of no interest to me.
EDIT: Oh yeah, and once I'm done developing my recipe (a few weeks maybe?), I also intend to post an article detailing the whole of the process -
including photos of the setup, and the before-and-after of the results.
I suppose that should go into the "Prepublication" subforum?
[Edited on 27-11-2016 by The Mad Plater]
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wg48
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Mad Plater:
Which manganese citrate are you trying to synthesis?
I would have thought your phosphating bath would have Mn(II) in it not the III.
PS the paper on citrate sablising Mn(III) is interesting I will have to read it in detail when I have time, thanks.
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Texium
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Thread Moved
28-11-2016 at 08:23 |
The Mad Plater
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I'm just as confused here as you are, wg48. Maybe even more so, in fact.
In the old timey speak, wouldn't the Mn(II) citrate be called "manganous citrate" instead?
This is why I figured that I should probably be going after the Mn(III).
Or did I get that mixed up?
At the end of the day, I think that it might not matter very much in this application.
I suppose that the idea is really to get the right amount of citrate ions into the solution - and using Mn citrate is probably the most expedient way
of accomplishing that, without unduly upsetting the bath chemistry.
A similar thing happens in electroless nickel plating - citrate ions in the solution form complexes with the nickel, which shifts the reaction
equilibrium towards depositing more phosphorus and less nickel (very good in my case), and has some other beneficial effects for the resulting
coating.
Although in that particular case, sodium citrate is generally used, with the sodium becoming a spectator ion.
Back to the Mn(II) citrate - I tried preparing that first, by adding appropriately less than a 1:1 molar ratio of citric acid, but that was odd.
After the reaction had stopped (at close to neutral pH, no less), it appeared as though there was still unreacted Mn carbonate left.
I then added more citric acid, to reach the 1:1 molar ratio, and - sure enough - the reaction picked up again.
Now, I interpret that as the reaction conditions favoring Mn(III) citrate formation - since otherwise, there shouldn't have been much of a further
reaction?
But remember, I'm no chemist. That's barely an educated guess.
In the end, what really matters to me is the end result - the as-deposited manganese phosphate coating.
As I've mentioned before, I find the "behind the scenes" of the chemistry involved as uninteresting, and largely irrelevant.
All I want to do, is to mix up a bunch of reagents, and get a good result out of it.
Hopefully, my manganese nitrate should arrive this week, so that I can run what I hope to be the decisive medium-scale phosphating experiment
this weekend.
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wg48
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Mad Plater:
Yes that’s correct about old speak. However as you have apparently discovered as I have some papers call Mn(II) salts simply manganese with out
clearly saying that it is the Mn(II).
If the required oxidation state of the manganese is II for your mixture its probably very important that you use Mn(II) and not III. The Mn(III) may
react very differently in your phosphating solution than the II.
In general if you started with Mn(II) simply adding any amount of citric acid will not change it to Mn(III). Assuming you do not want one of the
acid salts of citric acid the correct molar ratio of Mn(II) to citrate is 3 to 2.
My understanding is that in electroless nickel the citrate has various functions principally to complex nickel to stabilize the bath, and to buffer
the ph of the bath. Its pH that mostly determines the phosphorous content of the plating. Having said that, all of the bath ingredients have some
effect on the nickel deposit.
PS I have not tried your electroless nickel recipe yet but it is on my to do list.
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mayko
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In my experience, dissolved citrates tend to form gels and goos, and solidify only grudgingly. If it's like copper citrate, I'd expect manganese
citrate to be only sparingly soluble in plain water, but possibly more soluble in the presence of other citrate ions. I might make a bit of this stuff
later to find out...
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WGTR
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I've been following this on a tiny smartphone screen, so I may have missed something. One thing I haven't noticed, though, is any characterization
you may have done on the deposits. Are you or someone else doing this? Otherwise, the whole question of plating baths and chemicals is just a random
shot in the dark. This is unfortunately from experience, as sometimes literature doesn't match up with results the way one would expect. Normally, I
use an SEM to determine rough elemental composition of samples, but there are other ways.
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The Mad Plater
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@wg48:
OK, now this is where I find my understanding to be seriously lacking.
Could you please briefly explain the practical significance of the oxidation state, as it pertains to this case?
Also, under what circumstances can it be altered?
As for my EN process (specifically, the replenishment chemistry) - it's still very much a WIP. It's on the backburner for a bit, though - I have a
more pressing need to get the affairs in order with the phosphating, first.
Until the first replenishment at least, that recipe should work OK on steel though.
That 5% NaCl immersion corrosion test gave very promising results - after a few hundred hours, other than a few spots where a folded-over burr
prevented plating coverage of the underlying substrate (yes, the sample had been of very poor prior quality, and hastily/poorly sandblasted... not the
bath's fault!), the nickel coating proved otherwise impervious to corrosion.
And even in the few spots where the sample did begin to corrode, the coating didn't become lifted/detached around the origin of corrosion.
Considering that this was the first "real" plated test sample, I'd call that a fairly decent result.
Also, eventually I want to reformulate the whole EN process to use maybe 3 or 4 premixed solutions combined in simple volumetric ratios, as opposed to
weighing the reagents each and every time.
Again, when I'm done with that, I hope that I can muster the time and effort to compile it into an article.
@mayko:
Yes, this largely agrees with both the experimental evidence, and the (very sparse) literature.
I recall seeing statements such as "soluble in water containing dissolved sodium citrate", as well as "sparingly soluble in water".
Sounds about right.
It sure appears to be decently soluble in a dilute phosphoric acid solution, in any case.
@WGTR:
This is a slippery subject.
Yes, I'm doing this on my own. With no access whatsoever to any "real" lab equipment (NMR, HPLC, mass spectrometry, etc.).
Additionally, as you can probably already tell, my lab technique is rather lacking, too.
This is just for my own use - I'd have to be totally nuts to try commercializing this in any way. (bringing wood to a forest, and all that)
Especially with things such as the Cr(VI)-containing final sealing step. Good luck with THAT - the ecoterrorists would skin you alive.
You also seem to greatly overestimate my (chemical) abilities - take a look the first part of this post.
At this point in time, I wouldn't be able to tell the difference between, say, Mn(II) and Mn(III), even if it bit me squarely in the ass.
With that being the case, I find "attempting to characterize the deposits" a rather daunting task.
Also, remember, I'm a mechanical engineer & machinist, not an R&D chemist. I do have many tools at hand, but they are all totally unsuitable
for this purpose.
Finally, I don't really care about the Mn citrate in its solid form - as long as it forms in the solution, and stays put, that's all that I need.
Clearly, there appears to be some "interesting" chemistry going on here - if only due to the scarcity of available literature - but I just don't find
it most of it relevant for my needs.
One more thing - at these "small" test scales, the reagents involved are substantially less costly than the amount of time it takes to perform the
tests.
Thus, I find it preferable to "just try it" in an applied form, rather than spend that time trying to discover some marginally useful information on a
compound I never intend on isolating.
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wg48
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Mad Plater:
Like a lot of science there are layers of explanation and rules that start relatively simply and approximately true or only true in a subset of all
possible conditions. A bit like spelling rules for English words. Explanations of how chemical elements combine are like that too.
Ultimately how elements combine in a compound is determined by the forces between the elements which depends on the particular state of those elements
and involves very complicated calculations of orbital theory and quantum mechanics. Oxidation states are an introductory theory.
For a starting point and probably most applicable to your scenario is this explanation http://www.chemguide.co.uk/inorganic/redox/oxidnstates.html.
If you are only interested in getting a recipe to work oxidation states are only important in order to get the correct state of a particular
ingredient for your recipe. For example Manganese (II) phosphate will have different characteristic compared to Manganese (III) ie different amounts
of manganese relative to phosphate, different solubilitys (the Mn(III) salt may be insoluble) and different chemical characteristics and stability.
Depending on the other ingredients in your solution Mn(III) may decompose to a Mn(II)salt and a precipitate of manganese dioxide.
[Edited on 30-11-2016 by wg48]
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mayko
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| Quote: Originally posted by mayko | | In my experience, dissolved citrates tend to form gels and goos, and solidify only grudgingly. If it's like copper citrate, I'd expect manganese
citrate to be only sparingly soluble in plain water, but possibly more soluble in the presence of other citrate ions. I might make a bit of this stuff
later to find out... |
So, I tried this the other day using trisodium citrate and manganese sulfate, adapting the RSC procedure discussed in this thread.
Trisodium citrate, dihydrate: C6H5Na3O7*2H2O, mw = 294.12 g/mol
Manganese sulfate, monohydrate: MnSO4*H2O, mw = 169.01 g/mol
2.95 g sodium citrate were massed (~0.01 mol); also, 1.98 g manganese sulfate (~0.012 mol). I forgot that the product is a 3:2 salt, so I flubbed the
stoichiometry; there is an excess of citrate rather than the excess of manganese I meant to have.
Possibly because of this extra citrate, no solid precipitated upon mixing the dissolved reactants, even after heating.
What I did find the next morning was that the clear, colorless solution had turned an orange-pink color. I tried heating it again to incite
precipitation... and the color disappeared! I set the beaker aside to cool and when I checked on it a day later, it had reverted to the orange-pink.
Here's a controlled experiment: I made a manganese sulfate solution of about the same concentration, and added it to two test tubes. It was clear and
colorless. I added some Mn-Cit solution to two tubes. One tube of each solution was brought to and held at a boil over the course of ~2 minutes.
Results:
L-R: Mn, no treatment; Mn, heat; Mn+Cit, no treatment; Mn+Cit, heat.
I'm going to keep adding MnSO4 to the remaining citrate solution to see if I can get the salt to drop out, but this unexpected thermochromism might be
more interesting than the solid salt!
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The Mad Plater
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I'm back with more findings.
The planned experiment for today was an investigation of electroless nickel replenishment chemistry.
However, other things came up, and there wasn't enough time left to do that.
So instead, I decided to try mixing up the phosphating bath concentrate, "test" recipe.
I still don't have the (required) manganese nitrate yet, but that's not a big problem - the nitrate salts are to be added last, in any case.
First, I assumed that I'd be dealing with Mn(II) citrate. As it turned out, that assumption seems to have been correct.
So, here's the setup (with the nitrates omitted for the time being):
- 100mL beaker, in a 60C water bath,
- 30mL of DI water (finished concentrate volume = 80mL),
- 7.74g citric acid monohydrate,
- 6.35g MnCO3,
- 25.6g 75% orthophosphoric acid,
- and finally 9.33g more MnCO3, for a total of 15.86g.
First step: dissolve the citric acid in water. Easy enough.
Second step: add the carbonate. Slowly.
Adding more carbonate caused vigorous gassing, right until the very end.
This would suggest that Mn(II) citrate is being formed, as opposed to the Mn(III) salt, since otherwise the reaction would've stopped before all the
carbonate had been added.
At this point, the solution contained a large amount of insoluble precipitate, of a very light pink color - considerably lighter than the
MnCO3.
Third step: add the phosphoric acid.
This caused some gassing, due to a small amount of unreacted carbonate remaining from the previous reaction.
All the insoluble precipitate disappeared without a trace, having apparently become soluble again.
The solution turned a pinkish orange color. Or maybe it was orangish pink instead?
Fourth step: add the remaining MnCO3.
At first, everything was going well - the MnCO3 dissolving with vigorous gassing.
However, after more than about half of it had been added, the appearance of increasing amounts of an insoluble precipitate was noted.
I added water to a total of 80mL - a bit too much, but it'll evaporate anyway - and it seems to have improved the situation somewhat, but not by much.
There's still a bit of gassing going on, but it's really slow, even after cranking the temperature up to 75-80C.
Argh... at this point, it's getting very late, and I need to wrap it up for today.
Later on, I'll recheck the math, and try to salvage this mess if at all possible.
First thing to try - adding even more water, maybe it's a solubility issue?
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The Mad Plater
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Well, I'll be...
Well, folks, I finally managed to make the manganese(II?) citrate, and the synthetic route involved is bracingly direct.
Just allow a stoich mixture of citric acid and hydrogen peroxide to react with pure manganese dioxide, and behold!
As of right now, I did that at a very small scale (0.5g), because I ran out of H2O2.
Of course I also managed to spill some of the reaction mixture before all of the MnO2 reacted, which messed up the stoichiometry.
So then I added some more acid-peroxide mix to compensate for the spillage, erring towards the side of excess acid.
I'm at a loss here, since I expected the MnO2 to be effectively useless for the purpose of making Mn citrate.
The synthetic route itself is also making me very confused - the manganese in MnO2 is already pretty much oxidized to hell and gone, and
attempting to oxidize it any further wouldn't be of any help.
I just completely fail to understand how a strong oxidizer can act as a reducing agent.
BTW, just for the record - my previous attempt involved MnCO3 + citric acid + NaOH, and I did manage to (briefly) make some Mn citrate that
time as well, but I totally messed up when trying to precipitate it from the mixture.
My understanding is that the whitish powder which resulted was actually Mn(OH)2, instead of the citrate salt.
In fact, looking back at it now, I'm pretty sure that this method could never be made to work, at least with the reagents and equipment feasibly
available.
That's a big advantage of this new method - it doesn't introduce large amounts of impurities into the reaction mixture, in the form of spectator ions
(Na citrate), which enormously simplifies the workup: Just Remove Water.
Or don't even bother with that, since for my needs, a dilute aqueous solution is perfectly fine.
EDIT: Hmm, this is odd.
I tried evaporating the remaining water, which was rather uneventful - a pure white powder resulted, with no obvious signs of decomposition -
but then, weirdness happened
Upon attempting to redissolve the white powder, it was found to be most recalcitrant, refusing to dissolve completely even in an amount of water
>10x greater than what I started out with.
Boiling didn't help, either.
Well, I guess I'll just leave it alone for a day or two, and see if anything else happens.
This definitely warrants further investigation. Next time, I'll put some of the Mn citrate solution in an improvised desiccator, and see what happens
then - will the cold-crystallized powder have the same properties?
Also, perhaps more directly, I'll try adding the Mn citrate solution straight to the phosphating bath concentrate, and survey the resulting damage
(?).
EDIT2: A few hours later, it appears that some more of the white powder has indeed redissolved at RT.
It's as if the dissolution is just extremely slow, for whatever reason.
Why did it not dissolve quickly in boiling water though?
[Edited on 13-12-2016 by The Mad Plater]
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