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?

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.

Quote: Originally posted by The Mad Plater

which manganese citrate? (II) or (III)? In my reference material, it's referred to simply as "manganese citrate"

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.

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.

@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.

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!

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?

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]