Sciencemadness Discussion Board - KMnO4 + K2S2O5 = ???

KMnO4 + K2S2O5 = ???

rstar - 3-11-2011 at 04:02

Hi Geeks,
I just wanna know, what is produced in the following reaction:
KMnO4 + K2S2O5 = ??

I think maybe mixture of MnSO4 and K2SO4 will be formed. If so, how could I separate MnSO4 from K2SO4 ??

blogfast25 - 3-11-2011 at 05:35

KMnO4 is a powerful oxidiser and thiosulphate is easily oxidised. Half reactions in acid conditions are likely to be:

KMnO4 + 8 H+ + 5 e === > Mn2+ + K+ + 4 H2O

K2S2O5 + 3 H2O === > K2SO4 + SO4(2-) + 6 H+ + 4 e

[Edited on 3-11-2011 by blogfast25]

ScienceSquirrel - 3-11-2011 at 05:39

It is not thiosulphate but disulphite, your chemistry is correct though

blogfast25 - 3-11-2011 at 07:04

Quote: Originally posted by ScienceSquirrel

It is not thiosulphate but disulphite, your chemistry is correct though

Ooops. Off to the bathroom to get the egg off my face

rstar - 3-11-2011 at 08:33

So, the reaction is likely to be:
KMnO4 + K2S2O5 + H+ = K2SO4 + MnSO4 + H2O eh,... right ???

I think its not... right.
However, is there any method to separate K2SO4 from MnSO4 ??

Endimion17 - 3-11-2011 at 08:36

I think the metabisulphite (

) will reduce the permanganate down to the manganese(II). That's a common laboratory cleanup method.

[Edited on 3-11-2011 by Endimion17]

ScienceSquirrel - 3-11-2011 at 08:40

It will be reduced to manganese II as you say.
He should then precipitate the manganese as say the carbonate and then redissolve it in sulphuric acid to make the sulphate.

blogfast25 - 3-11-2011 at 09:24

Quote: Originally posted by ScienceSquirrel

It will be reduced to manganese II as you say.
He should then precipitate the manganese as say the carbonate and then redissolve it in sulphuric acid to make the sulphate.

An expensive way to lay waste to two interesting chemicals, to obtain one that's basically dirt cheap!

Save that permanganate for the finer things in life, is what I say...

Quote: Originally posted by rstar

So, the reaction is likely to be:
KMnO4 + K2S2O5 + H+ = K2SO4 + MnSO4 + H2O eh,... right ???

I think its not... right.

Multiply the top one by 4, the bottom one by 5. Get rid of excess H2O and H+ on either side. Enjoy!

[Edited on 3-11-2011 by blogfast25]

ScienceSquirrel - 3-11-2011 at 09:29

A good way to make manganese sulphate is manganese dioxide, oxalic acid and sulphuric acid.
The above method is just a pointless and mucky waste.

[Edited on 3-11-2011 by ScienceSquirrel]

blogfast25 - 3-11-2011 at 10:45

Squirrel:

Any ideas for the titrometric determination of Mn2+? Chelometry perhaps?

AndersHoveland - 3-11-2011 at 15:01

This is basically what will happen (unbalanced equation):
(2)KMnO4 + K2S2O5 --> (2)K2SO4 + MnO2

If excess K2S2O5 is used, the MnO2 is likely to react further and you will probably get manganese sulfite.

[Edited on 3-11-2011 by AndersHoveland]

blogfast25 - 4-11-2011 at 06:31

The only time I've seen permanganate being reduced to MnO2 in acid conditions was when oxidising alcohol (to acetic acid). That was in the presence of an excess alcohol.

In the standardisation titration of KMnO4 with oxalic acid I was warned to titrate quite slowly to avoid MnO2 being formed. It didn't form.

For the most part, the reduction of KMnO4 in acid conditions goes down all the way to Mn (II).

rstar - 4-11-2011 at 07:49

Quote: Originally posted by AndersHoveland

If excess K2S2O5 is used, the MnO2 is likely to react further and you will probably get manganese sulfite.

Yes, I had K2S2O5 in excess, but is the sulfite of manganese soluble in water ?? Never heard of such compound.

woelen - 4-11-2011 at 08:00

Indeed, in acidic solution, permanganate is reduced to manganese(II). This reaction, however does not give pure sulfate, part of the sulfite is oxidized to dithionate ion. So you get a mix of sulfate and dithionate. The dithionate ion is nearly as stable as sulfate, even in acidic solution.

ScienceSquirrel - 4-11-2011 at 08:53

Quote: Originally posted by blogfast25

Squirrel:

Any ideas for the titrometric determination of Mn2+? Chelometry perhaps?

The standard way is EDTA with Eriochrome Black T as indicator, a typical procedure is here;

http://fa.kfda.go.kr/standard/egongjeon_standard_view.jsp?Se...

blogfast25 - 4-11-2011 at 11:21

Thanks Squirrel. I've got Na2EDTA, not EDTA (and Erio black T), I assume that should work too. I was thinking more of a back titration with ZnSO4:

(known excess) Na2EDTA + Mn2+ === > Na2EDTAMn2+

Titrate excess Na2EDTA with ZnSO4 standard.

That avoids having to keep the solution warm: it could be warmed for some time to allow the complex to form, then cooled and titrated. What do you think?

blogfast25 - 4-11-2011 at 13:29

Talking about manganese, I wanted to test:

Quote: Originally posted by ScienceSquirrel

A good way to make manganese sulphate is manganese dioxide, oxalic acid and sulphuric acid.

(Or MnO2 + C2O4H2 + H2SO4 === > MnSO4 + 2 CO2 + 2 H2O)

... with some battery crud. As luck would have it I had just came into two defunct ‘Heavy Duty’ 6 V, non-alkaline batteries. These contain 4 large 1.5 cells and I got exactly 100 g of battery crud out of one of them (that was at least 95 % of it).

According to a report I found on recycled battery crud, the analysed batch contained 27 % Mn, 9 % Zn and of course loads of graphite, some iron and other crap.

The oxalic/sulphuric route would present a nice alternative to using HCl or even the SO2/sulphite method of Nurdrage.

To break up the material and extract water-solubles (zinc, ammonium) I added (w/o thinking too much) 200 ml of 1 M H2SO4 and immediately got a bit of fizz, not much. I then simmered the slurry until the lumps had broken up sufficiently for about 15 min then filtered on coffee filter (filtered well) and washed with another 200 ml of hot water. The first filtrate was lightly yellow coloured and clear.

Now I know that manganese oxides do react with sulphuric acid. Freshly precipitated MnO2 dissolves quickly in strong H2SO4 (with O2 evolution), Mn2O3 dissolves in conc. H2SO4 to Mn2(SO4)3. But I didn’t expect battery crud manganese oxide(s) to react much with 1 M H2SO4. And yet it did... Perhaps the fact that the main oxides (mainly Mn2O3 and MnO(OH) in a spent cell, according to SM’s Der Alte) are relatively ‘young’ makes them more reactive?

I first tested the first filtrate for Fe3+ with KSCN and it tested very positive. Then I added 1 M NaOH to a second, identical sample of filtrate and a thick, off-white precipitate dropped out. Adding thin bleach turned it brown/black and it dissolved readily in conc. H2SO4: this was Mn(OH)2.nH2O, lots of it.

The filtrate and wash water were then combined quantitatively and dilute Na2CO3 was added, resulting in much, much precipitation of (iron contaminated), beige coloured MnCO3. Surprisingly little CO2 evolved during this step, clearly much of the acid had reacted away. This was filtered off, washed with copious amounts of water and dried. It still not bone dry yet but based on the last reading there was about 12 g MnCO3 there: that’s really quite a lot.

The remaining battery sludge also tested still strongly positive for MnOx with strong HCl, evolving quite a bit of chlorine.

Der Alte recommends washing the crud with 0.1 (or less) M HCl but my own past experience shows that already at this low concentration some Mn oxide goes into solution and a modicum of chlorine is released.

Tomorrow I will try to simmer the crud with 1 M acetic acid, prior to adding the oxalic acid/sulphuric acid mix.

blogfast25 - 5-11-2011 at 13:17

So another 100 g of the same battery crud was now simmered for about ½ hour in 200 ml of 1 M acetic acid, then filtered and washed with 3 x 100 ml of hot water.

The first 75 ml of filtrate were tested for Fe3+ with KSCN and tested totally negative. It also tested positive for Mn2+ but much less so than the previous run with 1 M H2SO4. Due to the presence of Mn2+ I couldn’t test for Zn2+

To the filter cake, transferred to a Pyrex beaker, was added 100 ml of water, 50 g of conc. H2SO4 and 63 g of oxalic acid dihydrate dissolved in 100 ml of hot water. The oxalic acid solution was added slowly over a period of 10 - 15 min because strong effervescence threatened a boil-over, strong agitation prevented the foam from coming over the edge of the beaker.

Effervescence gradually slowed down and after full addition of the oxalic acid the slurry was simmered for a bit more, then hot filtered. Here’s what the filtrate looked like:

The filtrate tested strongly positive for Fe3+. The filter cake was washed with 2 x 100 ml of hot water. The washed filter cake was treated with hot 36 % HCl and no effervescence was observed or Cl2 smelled. But when emptying that beaker I di get a distinct whiff of Cl2. It would appear that the filter cake is largely free of Mn oxides.

I wonder about the zinc: if it was still in the cake after the acetic acid wash, would it combine with oxalate to form relatively insoluble zinc oxalate?

The filtrate is now being de-ironised using the method demonstrated by Nurdrage, here.

blogfast25 - 6-11-2011 at 07:55

Here’s the solution after treatment with Nurdrage’s method, contrast with the photo above:

It’s pinkish and clear. Its pH is 6.7 and it reacts not one iota with KSCN.

I believe Nurdrage’s explanation of what happens is somewhat needlessly complicated, at least if we assume that the iron present is as Fe [III]. In essence a good dollop of fresh and carefully washed Fe(OH)3.nH2O is added to the Fe3+/Mn2+ bearing solution, enough so that any remaining acidity in the solution is reacted away with the hydroxide and making sure some undissolved Fe(OH)3.nH2O remains.

At that point, and slightly simply put, the two main equilibria that rule are:

Fe(OH)3.nH2O(s) < === > Fe3+(aq) + 3 OH-(aq) … Ks = [Fe3+] x [OH-]3

And: 2 H2O(l) < === > H3O+(aq) + OH-(aq) … Kw = [H3O+] x [OH-]

The high pH and the extremely low solubility product of ferric hydroxide of 2.8 E-39, as opposed to the lesser insolubility of Mn(OH)2 of (Wiki) 0.0003221 g/100 ml, causes ferric oxide (but perhaps also basic ferric sulphates) to drop out, while the manganese remains a spectator ion. Filtering then yields an Fe3+ -free MnSO4 solution.

I’ve set aside the hydrated ferric oxide filter cake for future use. All in all it’s a neat separation method.

The final volume of filtrate obtained was about 300 ml which was gently boiled down without bumping, until almost no liquid remained and crystals of (presumably) MnSO4.4H2O had formed. These are now drying on a hot plate. Weighing tomorrow…

[Edited on 6-11-2011 by blogfast25]

rstar - 6-11-2011 at 08:33

Quote: Originally posted by ScienceSquirrel

A good way to make manganese sulphate is manganese dioxide, oxalic acid and sulphuric acid

I currently don't have Oxalic acid but have the other two. Is their any other reducer which can be used instead of Oxalic acid ??
And also does the H2SO4 really needs to be Concentrated ??

blogfast25 - 6-11-2011 at 13:13

Quote: Originally posted by rstar

I currently don't have Oxalic acid but have the other two. Is their any other reducer which can be used instead of Oxalic acid ??
And also does the H2SO4 really needs to be Concentrated ??

Sodium sulphite and SO2 work, metabisulphite will do it too. Also HCl (but that gives loads of Cl2 during the reduction!)

No, the sulphuric acid doesn't need to be conc. but for diluted H2SO4 you'll need to adjust strength.

Here’s the product, 34 g of pink MnSO4.4H2O:

But that’s only the equivalent of 8.4 g Mn or 13.3 g MnO2, so actual percentage yield from 100 g of crud was quite low.

rstar - 6-11-2011 at 20:27

Quote: Originally posted by blogfast25

Sodium sulphite and SO2 work, metabisulphite will do it too.

I have K2S2O5, it will too work.

So what will the equation be ?? Is it:
2MnO2 + K2S2O5 + H2SO4 = K2SO4 + 2MnSO4 + H2O

If that's so, then I will need to separate the K2SO4 off,
but......... HOW ??

AndersHoveland - 7-11-2011 at 02:47

K2SO4 is less soluble. Fractional crystallization.

solubility of potassium sulfate at 0degC is 8 g/100mL
solubility of manganese sulfate at 0degC is 52.9 g/100mL

[Edited on 7-11-2011 by AndersHoveland]

blogfast25 - 7-11-2011 at 06:56

Easier to precipitate the Mn as MnCO3, a stable, complete precursor to most Mn [+II] salts. On heating in air it gives MnO2.