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Author: Subject: Iron(II) sulfate monohydrate from eBay - looks suspicious
DrMario
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[*] posted on 15-12-2014 at 14:00
Iron(II) sulfate monohydrate from eBay - looks suspicious


I got some iron(II) sulfate monohydrate from eBay (UK seller).
I added some water to a small sample of it, but instead of obtaining a green solution... I have a yellowish opaque liquid. Most of the powder has not dissolved.

Is this how iron sulfate monohydrate is supposed to behave? I would have thought that it would have become heptahydrate and then just dissolved forming a greenish solution.
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DrMario
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[*] posted on 15-12-2014 at 14:51


The more I think of it, the more I'm thinking I was sent this:

12 FeSO4 + 3 O2 → 4 Fe2(SO4)3 + 2 Fe2O3

That is, iron(II) sulphate that was exposed to air/oxygen a wee bit too much... (see http://en.wikipedia.org/wiki/Iron%28II%29_sulfate )

[Edited on 15-12-2014 by DrMario]
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[*] posted on 15-12-2014 at 15:23


Brown means Ferric Sulphate... Looks like some oxidation did occur.
I envy inorganic chemists in their dedication to their field and to the analytical process even before they carry on with their reactions.

Why not just make Iron Sulphate from Sulphuric Acid and Iron? Should be simple enough when you think about it...
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[*] posted on 15-12-2014 at 15:50


Quote: Originally posted by DrMario  
I got some iron(II) sulfate monohydrate from eBay (UK seller).
I added some water to a small sample of it, but instead of obtaining a green solution... I have a yellowish opaque liquid. Most of the powder has not dissolved.

Is this how iron sulfate monohydrate is supposed to behave? I would have thought that it would have become heptahydrate and then just dissolved forming a greenish solution.


Could you u2u me the seller? I have also ordered and there isnt that many sellers. I am hoping its not the same guy I have ordered from




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UnintentionalChaos
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[*] posted on 15-12-2014 at 15:56


Every batch of FeSO4*nH2O I have acquired (non-reagent grade) is usually grungy from air partial oxidation to basic ferric sulfate, which gives you the hydrated yellow glop when added to water. You can filter the solution which should give a pale green solution. A small amount of added sulfuric acid helps to stabilize the solution, and you can add some cleaned iron to reduce any traces of ferric contamination, which will appear brown. Once crystallized again and completely dry, the stuff is pretty resistant to oxidation.



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[*] posted on 15-12-2014 at 16:16


Quote: Originally posted by DrMario  
I have a yellowish opaque liquid.


The US OTC fertilizer product is a hazy weak yellow. Your product is probably fine.




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DrMario
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[*] posted on 15-12-2014 at 22:26


Quote: Originally posted by dermolotov  
Brown means Ferric Sulphate... Looks like some oxidation did occur.

That's exactly what I suspect (see my second post). In fact, there's even some reddish-brown precipitate!

BTW, I'm not an organic chemist. I'm actually not a chemist by training.
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DrMario
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[*] posted on 15-12-2014 at 22:31


Quote: Originally posted by UnintentionalChaos  
Every batch of FeSO4*nH2O I have acquired (non-reagent grade) is usually grungy from air partial oxidation to basic ferric sulfate, which gives you the hydrated yellow glop when added to water. You can filter the solution which should give a pale green solution. A small amount of added sulfuric acid helps to stabilize the solution, and you can add some cleaned iron to reduce any traces of ferric contamination, which will appear brown. Once crystallized again and completely dry, the stuff is pretty resistant to oxidation.


Very interesting, thank you!
I have no elemental iron, but I do have sulfuric acid.


Thanks again!

[Edited on 16-12-2014 by DrMario]
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[*] posted on 15-12-2014 at 22:32


Quote: Originally posted by S.C. Wack  
Quote: Originally posted by DrMario  
I have a yellowish opaque liquid.


The US OTC fertilizer product is a hazy weak yellow. Your product is probably fine.


Do you see some reddish-brown precipitate when you try to dissolve it in water?
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[*] posted on 15-12-2014 at 23:57


I am going to have to test mine. It was reagent grade once but could easily be a few decades old. I acquired it a couple of weeks ago. It doesn't have the fresh blue appearance that I expected of FeSO4. However it is in an amber jar so I just mught not be looking at it right. If it is partially oxidised, is there a simple restoration that can be done? It might be prudent to recrystallise.
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[*] posted on 16-12-2014 at 00:18


Recovering oxidized iron(II) sulfate is very difficult and is not worth the effort. If you want an iron(II) salt, then next time I would purchase Mohr's salt, Fe(NH4)2(SO4)2.12H2O. Mohr's salt is MUCH more stable in air than ferrous sulfate. It is not so easily oxidized.



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[*] posted on 16-12-2014 at 06:14


Quote: Originally posted by woelen  
Recovering oxidized iron(II) sulfate is very difficult and is not worth the effort. If you want an iron(II) salt, then next time I would purchase Mohr's salt, Fe(NH4)2(SO4)2.12H2O. Mohr's salt is MUCH more stable in air than ferrous sulfate. It is not so easily oxidized.


Can I ask why its too difficult or not worth the effort to recover the oxidized Iron II sulfate. I recently purchased from ebay a bottle of iron II sulfate ACS, but discovered upon opening the bottle that it had oxidized to ferric sulfate... luckily the seller agreed to refund me... but i was allow to keep the ferric sulfate.

so is it not worth dissolving the ferric sulfate into water and then adding a few drops of sulfuric acid until the solution is green and then recrystallize it to get iron II sulfate?

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[*] posted on 16-12-2014 at 09:01


Quote: Originally posted by j_sum1  
I am going to have to test mine. It was reagent grade once but could easily be a few decades old. I acquired it a couple of weeks ago. It doesn't have the fresh blue appearance that I expected of FeSO4. However it is in an amber jar so I just mught not be looking at it right. If it is partially oxidised, is there a simple restoration that can be done? It might be prudent to recrystallise.


Please let me/us know the result of a simple dissolution test - i.e. try to dissolve a small sample of it in water and see what you get.
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[*] posted on 16-12-2014 at 09:02


Quote: Originally posted by woelen  
Recovering oxidized iron(II) sulfate is very difficult and is not worth the effort. If you want an iron(II) salt, then next time I would purchase Mohr's salt, Fe(NH4)2(SO4)2.12H2O. Mohr's salt is MUCH more stable in air than ferrous sulfate. It is not so easily oxidized.


How about iron(II) chloride?
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[*] posted on 16-12-2014 at 09:10


When I left my iron(II) sulfate solution out in the air overnight without a covering on it, I came back to find it a nasty brownish green color, but I was able to restore it just by gently heating it with some steel wool for a time.

[Edited on 12-16-2014 by No Tears Only Dreams Now]




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[*] posted on 16-12-2014 at 11:15


Quote: Originally posted by No Tears Only Dreams Now  
When I left my iron(II) sulfate solution out in the air overnight without a covering on it, I came back to find it a nasty brownish green color, but I was able to restore it just by gently heating it with some steel wool for a time.

[Edited on 12-16-2014 by No Tears Only Dreams Now]


Thanks a lot! I do have steel wool. I think I'll first add a bit of sulphuric acid and then later on the steel wool.

BTW, my solution is actually brown, at "higher concentrations" - though I find it hard to say "concentration" as the liquid is quite opaque and clearly not a real solution.
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[*] posted on 16-12-2014 at 13:13


Update: I dissolved some 30 g of mystery powder in hot water - obtained a really brown liquid with lots of precipitate - filtered out the precipitate and added 5 g of concentrated sulphuric acid to the resulting (still brown) liquid. After a couple of minutes... it became green-yellow! But not entirely translucent.

The solid on the filter is, I suspect, Fe2O3.

I suppose I should add some steel wool to the solution which is now light yellow green.
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[*] posted on 16-12-2014 at 13:18


One comment: even before adding the sulphuric acid, the solution was quite acidic already (indicator paper went red). I guess this is normal for a Fe(III) solution.
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[*] posted on 16-12-2014 at 13:25


Quote: Originally posted by DrMario  
One comment: even before adding the sulphuric acid, the solution was quite acidic already (indicator paper went red). I guess this is normal for a Fe(III) solution.


Yes- Fe(III) hydrolyzes significantly, forming Fe(OH)2+ type complexes (actually more complicated than that, but...) and hydronium ions.

ETA: My handy textbook (Kotz and Treichel) gives the Ka of Fe(III) as 6.3e-3, which is only slightly weaker than phosphoric acid, and significantly stronger than HF or acetic acid.

[Edited on 16-12-2014 by DraconicAcid]




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DrMario
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[*] posted on 16-12-2014 at 13:55


Quote: Originally posted by DraconicAcid  
Quote: Originally posted by DrMario  
One comment: even before adding the sulphuric acid, the solution was quite acidic already (indicator paper went red). I guess this is normal for a Fe(III) solution.


Yes- Fe(III) hydrolyzes significantly, forming Fe(OH)2+ type complexes (actually more complicated than that, but...) and hydronium ions.


Great, thank you.
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[*] posted on 16-12-2014 at 14:28


Here is the fruit of my labor, so far: about 100 mg of the mystery powder, dissolved in 1L water but UNfiltered this time. Then I added about 8 g of concentrated H2SO4. It became greenish and more translucent but still somewhat opaque. Finally, I added a bit of steel wool (2-3 g) into the bottle.

See picture for the result. I'll leave it like this overnight.
[img=http://s29.postimg.org/swedyrg1v/Fe_II_III.jpg]
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DrMario
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[*] posted on 16-12-2014 at 14:29


Let's try this image thing again:
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[*] posted on 16-12-2014 at 14:43


Quote: Originally posted by woelen  
Recovering oxidized iron(II) sulfate is very difficult and is not worth the effort. If you want an iron(II) salt, then next time I would purchase Mohr's salt, Fe(NH4)2(SO4)2.12H2O. Mohr's salt is MUCH more stable in air than ferrous sulfate. It is not so easily oxidized.

Could one add equimolar amounts of Ammonium sulfate to Iron(II) sulfate to make Mohr's salt? I have a very little amount of Mohr's salt, but about 30g of Iron(II) Sulfate.




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[*] posted on 16-12-2014 at 14:54


This might be useful http://en.wikipedia.org/wiki/Ammonium_iron%28II%29_sulfate#P...
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[*] posted on 16-12-2014 at 15:16


Indeed it is useful. I'll be trying that. I have a bit, but it's always 'better' when you make it.



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