Sciencemadness Discussion Board - Iron(II) sulfate monohydrate from eBay - looks suspicious

Iron(II) sulfate monohydrate from eBay - looks suspicious

DrMario - 15-12-2014 at 14:00

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.

DrMario - 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]

dermolotov - 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...

Little_Ghost_again - 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

UnintentionalChaos - 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.

S.C. Wack - 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.

DrMario - 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.

DrMario - 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]

DrMario - 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?

j_sum1 - 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.

woelen - 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.

jamit - 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?

DrMario - 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.

DrMario - 16-12-2014 at 09:02

Quote: Originally posted by woelen

How about iron(II) chloride?

Amos - 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]

DrMario - 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.

DrMario - 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.

DrMario - 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.

DraconicAcid - 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]

DrMario - 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.

DrMario - 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]

DrMario - 16-12-2014 at 14:29

Let's try this image thing again:

The Volatile Chemist - 16-12-2014 at 14:43

Quote: Originally posted by woelen

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.

gdflp - 16-12-2014 at 14:54

This might be useful http://en.wikipedia.org/wiki/Ammonium_iron%28II%29_sulfate#P...

The Volatile Chemist - 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.

DrMario - 16-12-2014 at 23:27

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.

woelen - 16-12-2014 at 23:42

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.

DieForelle - 17-12-2014 at 00:42

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.)

Amos - 17-12-2014 at 06:31

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?

DrMario - 17-12-2014 at 07:38

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?

Very interesting. Could you quantify, at least roughly, the amount of H2SO4 vs. original "ferrous" sulphate?

Amos - 17-12-2014 at 07:53

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!

DrMario - 17-12-2014 at 08:46

Quote: Originally posted by No Tears Only Dreams Now

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&...

Okay, that's it! Ferrous sulphate (and perhaps ferrous anything) is too much trouble.

DrMario - 17-12-2014 at 10:56

At my wit's end. :mad:

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]

DraconicAcid - 17-12-2014 at 11:15

Quote: Originally posted by DrMario

At my wit's end. :mad:

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]

DrMario - 17-12-2014 at 11:25

Quote: Originally posted by DraconicAcid

Quote: Originally posted by DrMario

At my wit's end. :mad:

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.

DraconicAcid - 17-12-2014 at 11:32

Quote: Originally posted by DrMario

I thought a sufficiently acidic environment should prevent this from happening.

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.

Amos - 17-12-2014 at 12:33

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]

The Volatile Chemist - 17-12-2014 at 12:41

Quote: Originally posted by DrMario

At my wit's end. :mad:

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.

Texium - 17-12-2014 at 20:26

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.

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.

DrMario - 17-12-2014 at 22:49

Quote: Originally posted by DraconicAcid

Quote: Originally posted by DrMario

I thought a sufficiently acidic environment should prevent this from happening.

I guess it's the trappings of my Italian being my second mother tongue - it interferes with English sometimes.

DrMario - 17-12-2014 at 23:10

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?

DrMario - 13-1-2015 at 12:06

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.

The Volatile Chemist - 13-1-2015 at 12:20

They look really pretty, hopefully they stay similar to that.

Amos - 13-1-2015 at 13:06

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...

DrMario - 13-1-2015 at 15:31

Quote: Originally posted by The Volatile Chemist

They look really pretty, hopefully they stay similar to that.

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.

The Volatile Chemist - 14-1-2015 at 09:12

Quote: Originally posted by DrMario

Quote: Originally posted by The Volatile Chemist

They look really pretty, hopefully they stay similar to that.

The truth is, I prepared this batch three weeks ago. It was really very cold outside, so I used that instead of the fridge

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.

nimgoldman - 30-1-2023 at 07:22

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]

Texium - 30-1-2023 at 16:07

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.

chloric1 - 31-1-2023 at 06:49

Quote: Originally posted by woelen

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

nimgoldman - 31-1-2023 at 07:29

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.

chloric1 - 31-1-2023 at 09:12

It’s not critical so long as you have a slight excess of acid. You will be adding more iron into solution anyways.

As long as metallic iron is present ferric will be reduced to ferrous in acid solution

chloric1 - 31-1-2023 at 16:27

Quote: Originally posted by woelen