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DrMario
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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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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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dermolotov
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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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Little_Ghost_again
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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
Dont ask me, I only know enough to be dangerous
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UnintentionalChaos
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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.
Department of Redundancy Department - Now with paperwork!
'In organic synthesis, we call decomposition products "crap", however this is not a IUPAC approved nomenclature.' -Nicodem
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S.C. Wack
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The US OTC fertilizer product is a hazy weak yellow. Your product is probably fine.
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DrMario
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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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| 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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DrMario
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Do you see some reddish-brown precipitate when you try to dissolve it in water?
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j_sum1
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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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woelen
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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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jamit
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| 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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DrMario
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| 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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DrMario
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| 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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Amos
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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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DrMario
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| 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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DrMario
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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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DrMario
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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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DraconicAcid
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| 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]
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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DrMario
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| 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.
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Great, thank you.
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DrMario
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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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Let's try this image thing again:
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The Volatile Chemist
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| 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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gdflp
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This might be useful http://en.wikipedia.org/wiki/Ammonium_iron%28II%29_sulfate#P...
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The Volatile Chemist
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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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