Sciencemadness Discussion Board - iron sulfate heptahydrate

iron sulfate heptahydrate

cthornbu - 29-9-2004 at 16:28

What is the thermal decomposition reaction for iron sulfate heptahydrate?

No need to start this thread two times!!

Mephisto - 29-9-2004 at 16:50

(1) 3 FeSO4 x aq  air-oxygen + heat  > FeO + Fe2(SO4

3

(2) Fe2(SO4

3   heat  > Fe2O3 + 3 SO3

According to "Chem. Unserer Zeit 16, 149–159 (1982)"

BromicAcid - 29-9-2004 at 18:03

You forgot the water in your first equation Mephisto, considering cthornbu mentioned that it was the heptahydrate twice. So, cthornbu, before the first step in Mephisto's reaction happens the water evaporates off leaving the ferrous sulfate unchanged, except its water of hydration.

However if you started from ferric sulfate start from the second equation, since you forgot to metion which oxidation state your iron was in, however it the iron (II) oxidation state is more common so that is what we're assuming you have.

JohnWW - 29-9-2004 at 19:40

I presume you mean ferrous sulfate, FeSO4.7H2O, called "copperas", forming green monoclinic crystals. in which 6 of the H2Os complex the Fe++, and one is bonded to the SO4--. It melts at 64ºC, and loses all its H2Os at 300ºC, presumably forming the anhydrous salt, although hydrolysed or partly hydrolysed species are a possibility. At much higher temperatures it decomposes to FeO (in the absence of air) and SO3.

John W.

Mephisto - 30-9-2004 at 01:00

Noticed this?

http://www.sciencemadness.org/talk/viewthread.php?tid=2617

If everybody posts only there for this topic, it would be easier to delete one of the double-threads.

Threads merged. Chemoleo

[Edited on 30-9-2004 by chemoleo]

Insoluble Ferric sulfate?

guy - 14-10-2004 at 20:56

I started with copper sulfate and used iron to displace it and it turned to ferrous sulfate. When i left it outside it turned red (ferric sulfate?) but it was insoluble. Also once I added H2O2 and it turned red but was soluble then it was later inslouble. What happened?

guy - 15-10-2004 at 14:17

Well, I looked at it for a while and I realized that the "ferric sulfate" must be copper powder that finally settled.

ordenblitz - 19-7-2005 at 18:42

I needed some Ferrous sulfate for another reaction so I went about making some. I first diluted some H2SO4 down to about a 40% solution, 200 ml in total, in a 1000 ml beaker. The actual concentration is not that important but should be above 15% and below about 60% for things to proceed smoothly. After it cooled to about room temperature I began stuffing in my iron source.
http://img149.imageshack.us/img149/1760/fewool4mz.jpg

I could have used iron turnings or some iron powder I have but I decided to go cheap, and it's perfectly acceptable to do it this way. The evolution of hydrogen began immediately and the temperature slowly started to rise. Every so often I would place the beaker in some cool water to mediate this. You need a big volume to work in as the foaming is energetic and even in this beaker rapid stirring was required to beat it back.
http://img149.imageshack.us/img149/9701/inh2so40sb.jpg

I continued this way until all steel wool on hand was consumed, about 14 pads. I actually had to stop anyway, as there was such an accumulation of a thick white precipitate in the now dark solution that I could hardly stir it any more. The FeSo4 was accumulating as ultra fine white crystals. The beaker was placed on the hot plate and heating begun. After reaching ~80º I began to add distilled water until I could no longer see any white crystals and the solution was perfectly black and not turbid. I brought the temperature back up to 70-80º and quickly poured into a vacuum filtration setup. I needed to add some more warm water to keep things flowing as the solution cooled somewhat and the crystals tried to clog the filter. The idea was to use as little water as possible to avoid much boiling to reduce for crystallization afterward. It was interesting to see the lovely blue solution coming from the bottom of the filter when such a black one went in the top. I don’t have a pic unfortunately, as I was a bit busy at the time. The beaker was placed in an insulated dewar to cool slowly thus encouraging large crystal growth with good purity.
http://img149.imageshack.us/img149/5783/crystal15lp.jpg
http://img149.imageshack.us/img149/1447/crystal25bp.jpg

I had a heck of a time getting this mass out of the beaker. I actually had to use a small soft face hammer and a screwdriver, ever so carefully so as not to crystallize my beaker. I decided not to boil down the remains for another crop as the yield was fantastic and I was interested in the best purity. I placed the broken mass in the filter funnel and gave them a quick wash with a dilute H2SO4 solution ~1% The result was placed in a perfectly dry dissector cabinet. There were two dishes full, one with the larger crystals and the other with broken up smaller ones. After one day of drying this is what the FeSO4o7H2O looked like.
http://img149.imageshack.us/img149/9664/1stday0ye.jpg

By the second day the water was being pulled off the crystals and the smaller ones (far dish) were starting to go pure white. I didn't know that they should lose their water that easily, but it was 0% humidity in the cabinet.
This is an interesting project with a pretty result.

[Edited on 20-7-2005 by ordenblitz]

2ndDay.JPG - 13kB

really bluish

chloric1 - 19-7-2005 at 19:02

Wow I have never seen ferrous salts that blue! When I made FeCl-4H2O pr the sulfate I always got more greenish color. the chloride was always pickle green in excess chloride but on washing with ice water a hint of aqua did show. Likewise ferrous sulfate monohydrate has been very pale green for me and more green for the heptahydrate with a bluish hint. I will have to play with ferrous salts soon. It is safer to heat the ferrous ammonium sulfate in the solution pass 60 degrees C, and the crystals are VERY striking I will show if I have time to make them.

ordenblitz - 19-7-2005 at 19:17

The concentrated solution was bluer than after the crystals had formed, whence it went a pale green. Out in the sun the large sparkling crystals were absolutely beautiful!
The first day of drying the crystals were greener then by the second and third day they went bluer. Then by the 6th day, the smaller ones lost their water and went snow white and crumbly. I had bottled the larger ones after the second.
I now have both flavors.

12AX7 - 19-7-2005 at 20:31

I've been doing something similar, but making zinc sulfate. My problem appears to be, after digesting the potmetal bar (an alloy of zinc with maybe 5-10% aluminum and 0.5-1% copper), it's totally black. A shame I don't have the filtering capability you do! The other problem is zinc sulfate is wildly variable in solubility, near boiling it's a saturated solution but slightly cooler and two inches of crystals drop out! I like your dewar idea... just need to stir it more next time, heheh

And by the way, any thoughts on seperating the Al(SO4)3?

Tim

Bought copperas at the garden supply shop

chloric1 - 17-8-2005 at 16:16

Well, I started messing with this a little and it is really different than the other garden chems. This stuff is actually pale green dry powder! Not brown, or yellow or black or whatever. So when dissolved in distilled water there is no true solution but instead an opaque muddy green suspension. This needs a little battery acid and some heat. I get a deep olive green liquid with black tinting. The dirty green liquid misleads, because if it surpasses saturation, beautiful blue-green ferrous sulfate heptahydrate will deposit. Alcohol can be added to facilitate crystallization. Acetone was found to work remarkably well. I even seen some heptahydrate converted to the white monohydrate when acetone was added at room temperature! More work will be done when time allows.

[Edited on 8/18/2005 by chloric1]

12AX7 - 17-8-2005 at 17:30

BTW if you want to keep down oxidation in the process, add some acid and iron... ferric sulfate is reduced to ferrous by the metal.

Tim

Yeh I know

chloric1 - 18-8-2005 at 16:53

I might do this as it may help saturate things for me. I just haven't yet and unreacted iron is just something more to filter out. Besides I am thinking of turning the mother liquor into ferric ammmonium alum with a littel H2O2. This would be after I collect some ferrous ammonium sulfate first.

Purchased

MadHatter - 19-8-2005 at 09:07

Chloric1, I also bought my ferrous sulphate from a garden supply shop. Very cheap
and very useful for neutralizing that pesky chlorate from a perchlorate electrolysis.

yes indeed

chloric1 - 19-8-2005 at 20:35

I might heat 1 gram of the unpurified salt to see how much sulfur trioxide comes off. If iI understand correctly half of the sulfur trioxide is reduced to the dioxide. Not exactly high yields but if you can buy several pounds for under 10 dollars what little SO3 can be collected pays in dividends! Some scientific suppliers sell 250 grams or so for $600! Must be a bitch to purify.

Conclusive theories

chloric1 - 24-8-2005 at 18:38

Ok here goes, with my copperas I tried to modify my purification to something closer to ordenblitz's method. I diluted battery acid to a strength range of 16 -19% then had the beaker heated to a temperature just so I barely handle the beaker with my bare hand (apporx 50 -55 degrees C). I was too lazy to find my thermometer so I used instinct. I then added as much of my copperas asI could. I used two common nails to maintain the ferrous state in the warm environment. After a couple hours of gettting my haircut and playing with my daughter I went to filter my solution, or more properly, muddy slop.
After everything began to cool, greenish blue iron sulfate heptahydrate began to deposit.

I included one photo of my salt. I believe some of the blue is created by artificial lighting hence my flash. Also, flourescent lighting seems to add to the bluer side of things. When I heated ferrous salt solutions in the absense of excess metallic iron, a blackish tint is assumed in solution that is not observed in solution with excess metal. I theorize that this could possibly be a tiny amount of the magnetic oxide, ferroferric oxide. This could be experimentaly determined with a neodymium magnet. But I have misplced my big super magnets and I want some cylinder or spherical magnets now anyways.

Enough about magnets, I just want to say I think the blue is from lighting effects, reaction conditions, and maybe even temperature. I alos believe the salt the has a bluer tint is more pure than the pickle green stuff you obtain from partially oxidized solution.

I changed the previous gigantic image of the crystals which blew the margins on this thread to a thumbnail linked to same immage-

[Edited on 5-8-2018 by Bert]

jamit - 4-2-2012 at 09:15

I bought some ferrous sulfate from eBay. Visually it looks pale green with some light brown tint. When I dissolve some in water you get a suspension of murky pale green, which over a period of hours turns increasingly pale yellowish-green. I added dil sulfuric acid to slow the oxidation and to convert any ferric back to ferrous. I then filtered but the solution was still yellowish green.

Here's my problem. I want to recrystallize ferrous sulfate but I think it's either too impure to recrystallize or it oxidizes to fast. Shouldn't the color of the solution be a clear pale green? Has anyone try purifying fertilizer ferrous sulfate? Any help or thought would be welcomed.

Magpie - 4-2-2012 at 09:21

http://www.sciencemadness.org/talk/viewthread.php?tid=5529&a...

jamit - 4-2-2012 at 15:39

Thanks magpie for the link... I already read that before I posted. My problem is more on purifying the ferrous sulfate I purchased from eBay. The color is pale green and seems like a higher purity then both garden supply brand and the one you buy from pottery store... At least based on visual appearance alone. The pottery and garden brand of ferrous sulfate has lots of brown particles mixed with pale green ferrous sulfate. The ones I purchased from eBay is basically pale green with very little, maybe a tint of brown particles. When I dissolved my ferrous sulfate I expected a light green solution but instead I get a greenish yellow color. I added a little dil sulfuric acid along with some steel wool to keep the oxidation to a minimum. I can't imagine that the solution I have now will crystalize into a green ferrous sulfate.

I've tried the sulfuric acid with steel wool before and yes it does works to make green ferrous sulfate, but my problem is why can't I just dissolve my ferrous sulfate purchased from eBay and recrystallize it to have green ferrous sulfate? I had the same problem with the garden brand and pottery, which is why I purchased the stuff on eBay. Any ideas what I'm doing wrong?

Magpie - 4-2-2012 at 17:30

Not really. If you have acidified the solution and added a bright nail I wouldn't know what else to do. Maybe someone else knows?

Texium - 12-9-2014 at 16:37

Sorry to pull up an old thread, but I was wondering, just how reactive is iron(II) sulfate? I accidentally made a nice big chunk of crystals (I was letting a partially oxidized solution fully oxidize so I could salvage some iron(III) sulfate, and to my surprise I found the crystals sitting in it after a few days), and I'd like to put them on display with my other compounds, but I wasn't sure how I should store them. From my experience working with solutions of it, it seems to oxidize very quickly, but what about for a dried chunk? I presume that it would take longer.

Would keeping them in a sealed container be enough to keep them looking the way that they're supposed to, or should I store them under mineral oil, or something else?

Texium - 14-9-2014 at 11:22

Alright, well, I put the crystal chunk in some mineral oil that I bubbled argon through since nobody replied. It's probably overkill, but I wanted to make sure that it stays nice-looking.
If this was a bad idea for any reason, please let me know.

Texium - 18-9-2014 at 19:32

Ok, I don't recommend storing this compound in mineral oil. As a test, I took some of the sample and put it in mineral oil, and I left some out dry in a plastic cup at around 25°C and about 35% humidity, and the sample exposed to the air is still the proper color while the sample in the oil has acquired a slight yellow tint to it over the past few days. I'm still confused about whether it oxidizes at all under ambient conditions when dry. Based on what I've observed, it doesn't, but it could need more time. Any advice?

Brain&Force - 18-9-2014 at 19:39

Yellow tint? How'd you source your mineral oil - could something be coordinating to it?

I've had no problems storing ferrous sulfate in a sealed vial. Purging the vial was found to be unnecessary.

Texium - 22-9-2014 at 15:47

Quote: Originally posted by Brain&Force

Yellow tint? How'd you source your mineral oil - could something be coordinating to it?

I've had no problems storing ferrous sulfate in a sealed vial. Purging the vial was found to be unnecessary.

I'm starting to realize that now, as the stuff that I had sitting out dry open to the air is completely clean. I'm making a fresh batch. My mineral oil came from Walgreen's.

Also, I noticed that there seem to be some side reactions involved with making it that don't seem to be pointed out usually. This time I tried using stoichiometric amounts of each reactant, and there was a lot of iron left over, as well as the presence of quite a bit of sulfur dioxide. I think that one or both of the following equations might have been occurring in addition to the expected one:

2Fe + 3H₂SO₄ --> Fe₂O₃ + 3SO₂ + 3H₂O

Fe + H₂SO₄ --> FeO + SO₂ + H₂O

Please correct me if I am blatantly mistaken.

blogfast25 - 23-9-2014 at 03:28

Quote: Originally posted by zts16

2Fe + 3H₂SO₄ --> Fe₂O₃ + 3SO₂ + 3H₂O

Fe + H₂SO₄ --> FeO + SO₂ + H₂O

Please correct me if I am blatantly mistaken.

Unfortunately you are.

Iron reacts with sulphuric acid (at least 'dilute') to form ferrous sulphate (FeSO4

and hydrogen.

Fe is incapable of reducing sulphate ions (oxidation number +6) to sulphur dioxide (oxidation number +4).

Once crystallised to the heptahydrate, the ferrous sulphate does tend to oxidise by air oxygen to ferric sulphate (Fe2(SO4

3

as well as ferric hydroxy sulphates and even ferric hydroxide, but those reactions are a rather slow process.

Amos - 23-9-2014 at 09:24

zts16, just make sure the environment the crystals are in is at a low pH, preferably maintained by sulfuric acid. I've had crystals left in sulfuric acid solution for over a month, exposed to the air, and not a hint of oxidation has been detected.

[Edited on 9-23-2014 by No Tears Only Dreams Now]

blogfast25 - 23-9-2014 at 12:21

Quote: Originally posted by No Tears Only Dreams Now

zts16, just make sure the environment the crystals are in is at a low pH, preferably maintained by sulfuric acid. I've had crystals left in sulfuric acid solution for over a month, exposed to the air, and not a hint of oxidation has been detected.

Give it time: eventually you'll see the first signs of oxidation. But it your conditions it is slow, for sure.

Texium - 23-9-2014 at 15:19

Quote: Originally posted by blogfast25

Quote: Originally posted by zts16

2Fe + 3H₂SO₄ --> Fe₂O₃ + 3SO₂ + 3H₂O

Fe + H₂SO₄ --> FeO + SO₂ + H₂O

Please correct me if I am blatantly mistaken.

Unfortunately you are.

Iron reacts with sulphuric acid (at least 'dilute') to form ferrous sulphate (FeSO4

3

as well as ferric hydroxy sulphates and even ferric hydroxide, but those reactions are a rather slow process.

So what then is causing the smell of sulfur dioxide? It's quite strong: far more pungent than the faint fumes normally present over concentrated sulfuric acid. I kept the reaction quite cool to avoid boiling away the acid. The acid that I have is lab grade, so there shouldn't be anything else in there other than the acid, the iron, and whatever else was in the nail which shouldn't be very much.
I'm not saying this in an argumentative way, but rather because it still puzzles me.

Brain&Force - 23-9-2014 at 15:41

There are often sulfur impurities in iron, and adding acid in an oxidizing environment will release not sulfur dioxide, but hydrogen sulfide.

Texium - 23-9-2014 at 18:03

I've smelled both of those gases before, and this is definitely mostly sulfur dioxide. There might be some hydrogen sulfide, but only a trace amount.

Amos - 23-9-2014 at 19:47

If you're trying to compare sulfuric acid and sulfur dioxide smells, I'm not sure you should do so. They both smell very different, and I believe that it is sulfur TRIoxide coming off of the acid, but you can correct me if I'm wrong. Iron sulfide impurities are fairly easily oxidized to ferrous sulfate at higher temperatures.

blogfast25 - 24-9-2014 at 03:20

Quote: Originally posted by zts16

[ It's quite strong: far more pungent than the faint fumes normally present over concentrated sulfuric acid. I kept the reaction quite cool to avoid boiling away the acid. The acid that I have is lab grade, so there shouldn't be anything else in there other than the acid, the iron, and whatever else was in the nail which shouldn't be very much.
I'm not saying this in an argumentative way, but rather because it still puzzles me.

How strong is your concentrated sulphuric acid? 95? 98? Higher?

And why use conc. H2SO4 to dissolve iron, when dilute works perfectly fine?

Dissolving most metals in acid does produce strange, sulphurous smells. A nail is of course not pure iron either...

Texium - 24-9-2014 at 09:11

I didn't use concentrated sulfuric acid. I have 93%, and I diluted a stoichiometric amount of it. I'm not sure exactly what concentration I diluted it to.
After it didn't fully react, I doubled the amount of acid and the reaction proceeded, but the nail was still not fully dissolved.
I dissolved another nail in HCl, and it reacted completely and there was only a faint odor of hydrogen sulfide.

blogfast25 - 24-9-2014 at 10:46

Quote: Originally posted by zts16

I didn't use concentrated sulfuric acid. I have 93%, and I diluted a stoichiometric amount of it. I'm not sure exactly what concentration I diluted it to.
After it didn't fully react, I doubled the amount of acid and the reaction proceeded, but the nail was still not fully dissolved.
I dissolved another nail in HCl, and it reacted completely and there was only a faint odor of hydrogen sulfide.

As I wrote above, with diluted sulphuric acid and iron you can't get SO2: there's simply nothing there to reduce the sulphate ions to sulphurous oxide.

Little_Ghost_again - 24-9-2014 at 11:48

Fabulous crystals! Another experiment on my to do in the future list! Interesting reading as well. I have great respect for the shear amount of knowledge floating around on here.

Texium - 24-9-2014 at 14:24

Quote: Originally posted by blogfast25

Quote: Originally posted by zts16

As I wrote above, with diluted sulphuric acid and iron you can't get SO2: there's simply nothing there to reduce the sulphate ions to sulphurous oxide.

Very well then. I will just assume that the nail is too impure for me to determine what is actually happening in this reaction and leave it at that.

aga - 24-9-2014 at 14:27

Nope.

Get some totally Pure Iron and do comparative tests.

Scientifically.

Then post the results so we can All learn something new.

[Edited on 24-9-2014 by aga]

Texium - 24-9-2014 at 15:33

Quote: Originally posted by aga

Nope.

Get some totally Pure Iron and do comparative tests.

Scientifically.

Then post the results so we can All learn something new.

I plan on doing that to make sure that my sulfuric acid is good. It should be, but you never really know.
I don't currently have any pure iron unfortunately (hence my use of nails in these reactions)