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DrMario
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In the morning I found the bottle to contain a clear solution.. and a ton of sh#t at the bottom. This is becoming a chore 
I'm not even sure how to practically crystalize the Fe(II) sulphate from this 1L of solution.
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woelen
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The reason why I say that it is hardly worth the effort to try to recover clean iron(II) sulfate from oxidized material is as follows:
- Iron(II) ion very easily is oxidized, especially in neutral to alkaline environments, when it is wet. Crystallizing from a solution cleanly only is
possible if air is excluded, otherwise it will become covered by a brown crust of basic salt, which contains a mix of iron(II) and iron(III).
- The material is dirt cheap. If you have to use other chemicals besides the ferrous sulfate (e.g. sulphuric acid in order to get a clear solution)
then the balance completely goes the wrong way.
With iron(II) chloride the situation is even worse. I once made that in solution and I had a nice blue/mint crystal. I took this out of the liquid,
and within seconds the crystal turned yellow/green, it was oxidized almost immediately.
Making Mohr's salt is fairly easy. Just add ammonium sulfate to a solution of ferrous sulfate. Add a few drops of H2SO4 as well. Allow any solid
material to settle, decant, and then allow to evaporate slowly. The acid and ammonium ions protect the iron(II) from aerial oxidation.
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DieForelle
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Interesting thread. I, too, found typical industrial ferrous sulfate to be surprisingly unmanageable when I ordered some to "fortify" my lawn without
adding any nitrogen, which is in any typical hardware store fertilizer and makes them grow faster instead of merely greener. I think I finally got a
pound or so dissolved in my 25 gallon sprayer by adding a bit of citric acid to help stabilize it, and filtering the larger precipitate that did form
with a couple layers of a tshirt. The water was still brownish-orange but didn't clog the Teejet nozzles. I had naively expected it to mostly
dissolve like any other salt.
(some of you might remember my story of first trying to buy this at a now defunct feed store in Newark Delaware - merely asking for "can I buy a bag
of iron sulfate" - and the hayseed counter sales lady asking "what kind o' drugs you makin'?" An example of people's bizarre paranoia about chemicals
and their uses. And no, I didn't look like a drug dealer! Maybe she was a big Breaking Bad fan. Anyhow, they didn't have it...problem was it more
of a horse feed store, apparently pure iron sulfate is more something used by cattle farmers. Not many of those in northern Delaware anymore.)
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Amos
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DrMario, the first time I had ever even dealt with ferrous sulfate(which was homemade, by the way) like I said, I managed to change the color back
after the inital oxidation, and I also kept a piece of steel wool in there 24/7, as well as a lot of sulfuric acid; kept this way it never went beyond
green in an entire week while I was inactive in the lab. Later, I removed the steel wool, added EVEN MORE acid, and was able to let it just
crystallize slowly in a container for over a month. The result was a crop of beautiful blue-green crystals underneath a golden-yellow solution, which
by then was likely highly concentrated sulfuric acid. I don't know why it was so easy for me; maybe homemade grades are somehow more stable?
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DrMario
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Quote: Originally posted by No Tears Only Dreams Now  | | DrMario, the first time I had ever even dealt with ferrous sulfate(which was homemade, by the way) like I said, I managed to change the color back
after the inital oxidation, and I also kept a piece of steel wool in there 24/7, as well as a lot of sulfuric acid; kept this way it never went beyond
green in an entire week while I was inactive in the lab. Later, I removed the steel wool, added EVEN MORE acid, and was able to let it just
crystallize slowly in a container for over a month. The result was a crop of beautiful blue-green crystals underneath a golden-yellow solution, which
by then was likely highly concentrated sulfuric acid. I don't know why it was so easy for me; maybe homemade grades are somehow more stable?
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Very interesting. Could you quantify, at least roughly, the amount of H2SO4 vs. original "ferrous" sulphate?
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Amos
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I think I initially prepared my ferrous sulfate using a 15-20% solution of sulfuric acid, and just periodically added more out of paranoia that I
would spoil it after that initial hiccup occurred. This video provides a few helpful tips that I made heavy use of during the process: https://www.youtube.com/watch?v=BhqcPqaL_KQ&index=35&...
And now I'm off to prepare a solution of ferrous sulfate by displacing copper with iron!
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DrMario
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Okay, that's it! Ferrous sulphate (and perhaps ferrous anything) is too much trouble.
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DrMario
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At my wit's end. 
I filtered the content of the bottle into a 2L beaker. I obtained a perfectly transparent green liquid, which I proceeded to heat to near boiling
point on a hotplate. After about an hour of heating... the liquid is opaque with beige crap deposited on the bottom!
This is bullsh#t.
[Edited on 18-12-2014 by DrMario]
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DraconicAcid
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| Quote: Originally posted by DrMario | At my wit's end.
I filtrated the content of the bottle into a 2L beaker. I obtained a perfectly transparent green liquid, which I proceeded to heat to near boiling
point on a hotplate. After about an hour of heating... the liquid is opaque with beige crap deposited on the bottom! |
That will continue to happen if it's exposed to air. 4 Fe2+ + O2 + 2 H2O --> 4 Fe(OH)2+.
[Edited on 17-12-2014 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 | At my wit's end.
I filtrated the content of the bottle into a 2L beaker. I obtained a perfectly transparent green liquid, which I proceeded to heat to near boiling
point on a hotplate. After about an hour of heating... the liquid is opaque with beige crap deposited on the bottom! |
That will continue to happen if it's exposed to air. 4 Fe2+ + O2 + 2 H2O --> 4 Fe(OH)2+.
[Edited on 17-12-2014 by DraconicAcid] |
I thought a sufficiently acidic environment should prevent this from happening.
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DraconicAcid
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It happens faster under basic conditions, but it will still happen in acidic ones. Keep the air out.
Oh, and you didn't filtrate the solution- you filtered it. The solution you got afterwards was the filtrate.
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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Amos
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This just doesn't make any sense! Do you recall when earlier I said I was going to make some ferrous sulfate solution by displacing the copper from
copper sulfate?
I put an excess of steel wool into some copper sulfate solution, added just a bit of HCl to get the reaction started, and was left with a very light
green solution. At this point the solution really shouldn't be more than a tiny bit below pH 7. So I boiled that solution down to about 1/3 of the
original volume, during which no color change occurred except a concentration of the green color. An hour in my freezer, and I got a nice little crop
of fluffy FeSO4 crystals; EVERYTHING was exposed to the air during this time; there wasn't even a lid on the container in the freezer. I don't
understand how one person can have such a negative experience with this compound when mine is perfectly content to remain unoxidized.
EDIT: I attached a picture; this stuff shows no signs at all of oxidation, IMO. I wish the crystals were a little more well-formed, though.
[Edited on 12-17-2014 by No Tears Only Dreams Now]
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The Volatile Chemist
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| Quote: Originally posted by DrMario | At my wit's end.
I filtrated the content of the bottle into a 2L beaker. I obtained a perfectly transparent green liquid, which I proceeded to heat to near boiling
point on a hotplate. After about an hour of heating... the liquid is opaque with beige crap deposited on the bottom! |
I know exactly how you feel. I did this too, but after noting the filtered solution was beige-ish, I went to take it off my hotplate, having decided
to give up, and dropped the 50mL beaker on the hotplate, shattering it and spilling a lot of almost boiling liquid on the ground and on the hotplate.
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Texium
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| Quote: Originally posted by No Tears Only Dreams Now | This just doesn't make any sense! Do you recall when earlier I said I was going to make some ferrous sulfate solution by displacing the copper from
copper sulfate?
I put an excess of steel wool into some copper sulfate solution, added just a bit of HCl to get the reaction started, and was left with a very light
green solution. At this point the solution really shouldn't be more than a tiny bit below pH 7. So I boiled that solution down to about 1/3 of the
original volume, during which no color change occurred except a concentration of the green color. An hour in my freezer, and I got a nice little crop
of fluffy FeSO4 crystals; EVERYTHING was exposed to the air during this time; there wasn't even a lid on the container in the freezer. I don't
understand how one person can have such a negative experience with this compound when mine is perfectly content to remain unoxidized.
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I know exactly what you mean! One thing though is that it does seem like it oxidizes much more rapidly when
it is wet. From my experience with iron(II) sulfate, you need to dry it as quickly as possible once you get your crystals. After that, letting them
sit open to the air is no problem. I've had some open to the air for months with no signs of oxidation whatsoever. Solutions of it however will go
brown within hours. The silver lining though, is that if you let a solution evaporate, while the solution does get oxidized, large, well-formed
crystals of iron(II) sulfate will crystallize on the bottom of the container. Then you can remove the crystals, rinse them with water to clean them
(they clean up very nicely) and then once no more will crystallize, filter the solution, separating the iron(III) sulfate and oxide. The oxide can
then be dried, and the sulfate crystallized, and you have three decently pure compounds from one vessel. I think I will make a YouTube video of this
procedure at some point.
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DrMario
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| Quote: Originally posted by DraconicAcid |
It happens faster under basic conditions, but it will still happen in acidic ones. Keep the air out.
Oh, and you didn't filtrate the solution- you filtered it. The solution you got afterwards was the filtrate. |
I guess it's the trappings of my Italian being my second mother tongue - it interferes with English sometimes.
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DrMario
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I have one nagging question that has been bothering me since the first time I read the advice to add H2SO4: wouldn't an excess of sulphuric acid,
added to a solution that contains a mixture of Fe(II) and Fe(III) sulphate effectively "push" the ratio towards higher Fe(III)
sulphate content?
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DrMario
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And here it is, ferrous sulfate crystals:
It was easier than I thought... once I actually tried cooling the solution of iron(II) sulfate and H2SO4.
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The Volatile Chemist
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They look really pretty, hopefully they stay similar to that.
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Amos
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It's about time! I should be preparing quite a bit soon by the copper sulfate/steel wool method, if mine goes disastrously for once, I'll know which
thread to come to...
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DrMario
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The truth is, I prepared this batch three weeks ago. It was really very cold outside, so I used that instead of the fridge 
The crystals in the bottle look more-or-less the same, but I don't vouch that I managed to remove all the H2SO4 from them, so the acid might be
protecting the ferrous sulfate from oxidizing to ferric sulfate.
Maybe I'll take a photo of the bottle a couple of weeks from now and we can all compare the colour. A yellowing would indicate oxidation.
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The Volatile Chemist
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| Quote: Originally posted by DrMario |
The truth is, I prepared this batch three weeks ago. It was really very cold outside, so I used that instead of the fridge
The crystals in the bottle look more-or-less the same, but I don't vouch that I managed to remove all the H2SO4 from them, so the acid might be
protecting the ferrous sulfate from oxidizing to ferric sulfate.
Maybe I'll take a photo of the bottle a couple of weeks from now and we can all compare the colour. A yellowing would indicate oxidation.
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A bottle I ordered about 3 months ago has really oxidized, I'm kinds annoyed. I guess storing it in a little bit of residual
H2SO4 would help.
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nimgoldman
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I also tried recovering old iron(II) sulfate (technical-grade), which was considerably oxidized. Here are my two cents...
One issue is dissolving everything since using dilute H2SO4 won't dissolve all the iron(III) sulfate (it is more prone to hydrolysis than iron(II) so
requires more acid) and using too concentrated acid reduces solubility of iron(II) and may even oxidize it on heating:
2 FeSO4 + 2 H2SO4 --> Fe2(SO4)3 + SO2 + H2O
Concentrated acid also produces mixture of Fe(2+) and Fe(3+) with fresh iron so the comproportionation is possible only in dilute acid:
3 Fe + 4 H2SO4 --> Fe(2+) + 2 Fe(3+) + 4 SO4(2-) + 4 H2
To make things worse, iron forms various hydrated oxides some of which are resistant to mineral acids and some even colloidal. Here is an excerpt from
atomistry.com on Fe2O3.H2O (hydrated ferric oxide, a.k.a. Goethite):
| Quote: |
Boiling concentrated nitric acid has but little effect, and even concentrated hydrochloric acid only attacks it after prolonged digestion at the
boiling-point. After some hours of treatment with acetic acid at 100° C. a colourless colloidal solution is obtained, from which addition of a trace
of sulphuric acid effects the precipitation of the insoluble monohydrate. |
So the good strategy may be prolonged stirring and heating in dilute (e.g. 0.4M) sulfuric acid in a container protected from atmospheric oxygen. I
would add iron filings from time to time until there is virtually no iron(III) (tested by adding a drop of 10% potassium thiocyanate to diluted sample
- it not give strong red color and preferably should be colorless).
Then the solution is filtered, concentrated and cooled to give reasonably pure iron(II) sulfate crystals.
The residues on filter should be undissolved iron(III) compounds. I've read that some iron oxides are resistant even to conc. HCl and people used
5-10% acetic, citric or oxalic acids to dissolve all the oxides. Still, this may keep iron(III) sulfate undissolved.
Another strategy might be to oxidize and dissolve all the residue with conc. sulfuric and oxidizer. This way we can prepare fresh iron(III) hydroxide
that can be separated, dissolved, reduced...
An iteresting approach would be bubbling SO2 gas through the iron(II,III) mixture - it may even help dissolving the oxides and it may be faster than
using elemental iron.
As woelen noted, it's a lot of work. I have about a kilo of the old iron(II) sulfate as well as Mohr's salt that I use preferably as it is much more
stable. I don't know what to do with the former, maybe I will make iron(III) sulfate from it instead (useful oxidizing agent) or ferric ammonium
sulfate (FAS) which is iron(III) counterpart of the Mohr's salt.
[Edited on 30-1-2023 by nimgoldman]
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Texium
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| Quote: Originally posted by Texium | | From my experience with iron(II) sulfate, you need to dry it as quickly as possible once you get your crystals. After that, letting them sit open to
the air is no problem. I've had some open to the air for months with no signs of oxidation whatsoever. Solutions of it however will go brown within
hours. The silver lining though, is that if you let a solution evaporate, while the solution does get oxidized, large, well-formed crystals of
iron(II) sulfate will crystallize on the bottom of the container. Then you can remove the crystals, rinse them with water to clean them (they clean up
very nicely) and then once no more will crystallize, filter the solution, separating the iron(III) sulfate and oxide. |
About 8 years later (!) I have some fresh insights on this topic. My iron(II) sulfate I talked about here did eventually oxidize. It
started from the top of the vial where it was most exposed to fresh air and humidity, and worked its way through until everything was brown. This took
several years though, and it was stored in a vial with a crappy cork stopper. I think ambient humidity plays a big role in how long it takes for
oxidation to take place. Commercial samples from reputable suppliers that I observed in Texas (where it is quite humid) seemed to oxidize very
quickly. Meanwhile, bottles of similar age that I've seen here in dry Utah have looked very nice.
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chloric1
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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. |
Sorry but no. It’s so easy. I have purchased “copperas” from a fertilizer co-op store, took it home, and made up 20% sulfuric acid solution
added the brownish yellow iron salt and boiler with regular nails for about an hour or two and let it cool and got beautiful blue green crystals
Fellow molecular manipulator
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nimgoldman
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| Quote: Originally posted by chloric1 | | I have purchased “copperas” from a fertilizer co-op store, took it home, and made up 20% sulfuric acid solution added the brownish yellow iron
salt and boiler with regular nails for about an hour or two and let it cool and got beautiful blue green crystals |
How much acid solution did you use per gram of crude iron salt?
I managed to reduce all the iron(III) in solution with iron filings (to the point the thiocyanate test went colorless), but there was still a very
fine white-ish precipitate that wouldn't dissolve - after a day or two it settled and became yellowish-brownish. I plan to just filter this out but I
am still interested whether it is some kind of insoluble iron compound similar to calcined chromium(III) oxide - something completely insoluble under
normal conditions. But maybe it's just some dirt.
I would be worried that 20% H2SO4 will promote the decomposition of FeSO4 into Fe2(SO4)3 and SO2, but given that SO2 is also reducing agent the
reaction can go in reverse.
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