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Author: Subject: Conversion of Manganese Sulphate to Acetate
CharlesWood
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[*] posted on 18-7-2022 at 23:48
Conversion of Manganese Sulphate to Acetate


The basic reaction (from the Crystal People) is

2CH3COOH + MnSO4 = Mn(CH3COO)2↓ + H2SO4

They seem to not bother with the H2SO4 and crystallize directly using low temperature. I assume leaving reasonably strong sulphuric Acid behind?

Do I get any advantage by adding CaCO3 to the mix to neutralise the H2SO4?
Will the potential presence of Ca(CH3COO)2 cause problems in getting pure Mn(CH3COO)2 by fractional crystalization?
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mysteriusbhoice
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[*] posted on 19-7-2022 at 02:26


Dude just react manganese sulfate with sodium carbonate or bicarbonate to get manganese carbonate. Then filter off the manganese carbonate then react with acetic acid. The yield is higher using this.
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[*] posted on 19-7-2022 at 05:47


Maybe even direct reaction with sodium or calcium acetate would work although these might leave contaminated manganese acetate.



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Texium
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[*] posted on 19-7-2022 at 07:01


Quote: Originally posted by mysteriusbhoice  
Dude just react manganese sulfate with sodium carbonate or bicarbonate to get manganese carbonate. Then filter off the manganese carbonate then react with acetic acid. The yield is higher using this.
Agreed. And you won’t need to use an excess of concentrated acetic acid. Going through the carbonate is far more efficient and practical. You can also make extra carbonate to save and use to make other manganese salts with ease.



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Fulmen
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[*] posted on 19-7-2022 at 13:28


On the other hand such precipitated compounds can be hard to work with in larger amounts. Which in turn makes it hard to wash out the sodium. So the direct route sounds very interesting if you want large amounts of pure product.



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SplendidAcylation
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[*] posted on 19-7-2022 at 14:07


I feel like a bit of an idiot here, but can someone explain how this reaction is even possible?

The equilibrium should lie so far to the left as to make this unfeasible, or am I missing something obvious? :o
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[*] posted on 19-7-2022 at 18:01


I tried to make mangense(II) acetate once via MnSO4 and Na2CO3, but it turned brown during evaporation of water.
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[*] posted on 19-7-2022 at 19:14


Quote: Originally posted by SplendidAcylation  
I feel like a bit of an idiot here, but can someone explain how this reaction is even possible?

The equilibrium should lie so far to the left as to make this unfeasible, or am I missing something obvious? :o

That was my thought too. But I guess, if manganese (II) acetate is extremely insoluble (I have not looked it up), then it would precipitate and be removed from the equilibrium. But is it really insoluble?? I am going to be a bit skeptical here.
Quote: Originally posted by Bedlasky  
I tried to make mangense(II) acetate once via MnSO4 and Na2CO3, but it turned brown during evaporation of water.

Desiccation via vacuum is probably a better option. Mn(II) does tend to oxidise if oxygen is present.
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SplendidAcylation
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[*] posted on 20-7-2022 at 02:05


Quote: Originally posted by j_sum1  
Quote: Originally posted by SplendidAcylation  
I feel like a bit of an idiot here, but can someone explain how this reaction is even possible?

The equilibrium should lie so far to the left as to make this unfeasible, or am I missing something obvious? :o

That was my thought too. But I guess, if manganese (II) acetate is extremely insoluble (I have not looked it up), then it would precipitate and be removed from the equilibrium. But is it really insoluble?? I am going to be a bit skeptical here.


Manganese acetate actually seems to be more soluble than the sulfate! :P
Although the data are not exactly at the same temperature:
Manganese acetate: 700g/L at 20c (tetrahydrate)
Manganese sulfate: 700g/L at 70c

So it seems like manganese acetate is a good bit more soluble...

Anyway, even if it were insoluble, would that actually push the equilibrium forward?
If you think about, for instance, magnesium carbonate, it is an insoluble salt of a weak acid, yet its insolubility doesn't prevent it from reacting with sulfuric acid, the opposite also applies, bubbling CO2 into a solution of magnesium sulfate isn't going to produce a precipitate of magnesium carbonate...

In this case, the volatile nature of CO2 seems to indicate that possibly the escape of CO2 is what pushes the neutralization reaction forward, but you could come up with another example where the acid isn't so volatile, such as, for instance:
3MgSO4 (aq) + 2H3PO4 (aq) ----> 3H2SO4 (aq) + Mg3(PO4)2 (s)

I'm pretty sure this wouldn't take place either, because the acid-base equilibrium tends to predominate, and sulfuric acid is a much stronger acid than phosphoric acid.

If this reaction did take place, it'd be a nice way of making sulfuric acid!

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Fulmen
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[*] posted on 20-7-2022 at 02:13


According to this the solubility of the sulfate is a bit weird: https://en.wikipedia.org/wiki/Solubility_table#M

It doesn't list the acetate however, so no real conclusion.




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[*] posted on 21-7-2022 at 23:16


Mn(II) acetate is well soluble.
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[*] posted on 22-7-2022 at 00:04


It might be possible in glacial acetic acid, if the Mn salt isn't soluble in that .
But I would still ppt the hydroxide or carbonate and then dissolve that in acetic acid.
If you really need to remove Na afterwards, you can recrystalise the stuff.
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[*] posted on 22-7-2022 at 10:41


The source specifies 70% or glacial acid, I have a feeling that just might work.

And I have to stress that large precipitation reactions suck the big felota. Any more than a few grams at the time and you're looking at a bucket of unworkable slugdge. You really need some solubility in order to grow dense particles that can be filtered.




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