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cthornbu
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iron sulfate heptahydrate
What is the thermal decomposition reaction for iron sulfate heptahydrate?
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Mephisto
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No need to start this thread two times!!
(1) 3 FeSO<sub>4</sub> x aq <u> air-oxygen + heat </u><sub>></sub> FeO +
Fe<sub>2</sub>(SO<sub>4</sub> <sub>3</sub>
(2) Fe<sub>2</sub>(SO<sub>4</sub><sub>3</sub>
<u> heat </u><sub>></sub> Fe<sub>2</sub>O<sub>3</sub> + 3
SO<sub>3</sub>
According to "Chem. Unserer Zeit 16, 149–159 (1982)"
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BromicAcid
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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.
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JohnWW
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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.
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Mephisto
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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]
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guy
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Insoluble Ferric sulfate?
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?
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guy
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Well, I looked at it for a while and I realized that the "ferric sulfate" must be copper powder that finally settled.
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ordenblitz
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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]
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chloric1
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really bluish
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.
Fellow molecular manipulator
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ordenblitz
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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.
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12AX7
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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
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chloric1
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Bought copperas at the garden supply shop
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]
Fellow molecular manipulator
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12AX7
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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
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chloric1
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Yeh I know
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.
Fellow molecular manipulator
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MadHatter
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Purchased
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.
From opening of NCIS New Orleans - It goes a BOOM ! BOOM ! BOOM ! MUHAHAHAHAHAHAHA !
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chloric1
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yes indeed
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.
Fellow molecular manipulator
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chloric1
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Conclusive theories
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]
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jamit
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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.
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Magpie
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http://www.sciencemadness.org/talk/viewthread.php?tid=5529&a...
The single most important condition for a successful synthesis is good mixing - Nicodem
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jamit
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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?
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Magpie
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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?
The single most important condition for a successful synthesis is good mixing - Nicodem
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Texium
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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?
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Texium
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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.
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Texium
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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?
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Brain&Force
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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.
At the end of the day, simulating atoms doesn't beat working with the real things...
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