Sciencemadness Discussion Board

ferrous sulfate

Magpie - 4-5-2007 at 11:59

I am getting materials together to make a few grams of quinoline via the Skraup synthesis. My procedure calls for the use of "ferrous sulfate" to moderate a normally violent reaction.

What I have available is pottery grade ferrous sulfate heptahydrate. I've looked at Vogel, OrgSyn, and Cummings to see if any call out a specific hydrate or lack thereof. None do. Each just says "ferrous sulfate."

In looking at Fisher and VWR catalogs it seems that the heptahydrate is all that is commonly available. So, it seems that the common nomenclature used in the procedures assumes the usual hydration without actually specifying what it is. It seems like I have run onto this before.

I feel I need to understand this before proceeding for a couple of reasons:

1) the water of hydration for FeSO4*7H2O is about 45% of its weight.

2) the procedure specifies the use of dry glycerine indicating that excess water added will reduce yield.

So, my question to the forum is: In these quinoline procedures do you think that "ferrous sulfate" actually means ferrous sulfate in its commonly available form, ie, FeSO4*7H2O?

woelen - 4-5-2007 at 12:29

Yes, you may assume that ferrous sulfate is FeSO4.7H2O. The anhydrous sulfate is a white solid, which is very unreactive (it does not dissolve in water). Anhydrous ferrous sulfate is hard to make, it cannot be made from hydrated ferrous sulfate, due to an internal redox reaction (iron goes to +3, sulfate forms SO2).

If you use pottery grade ferrous sulfate, then first try to clean it a little bit. Usually pottery grade ferrous sulfate contains brown clumps, which are basic ferric sulfate. These do not dissolve and should be removed.

Magpie - 4-5-2007 at 12:33

Thank you Woelen. I shall proceed to the lab.

12AX7 - 4-5-2007 at 13:12

Sure of that? I've left out crystals that partially dehydrated without turning to brown mush.

Tim

obsessed_chemist - 4-5-2007 at 13:38

A little off-topic, but...

I've contemplated a procedure for purifying garden-quality copperas (ferrous sulfate, heptahydrate). AFAIK, this source tends to be acidic from excess sulfuric acid used in it's synthesis (unsure if pottery-grade copperas is similar in this regard?).

Firstly, ferrous hydroxide, an insoluble solid, can be prepared from lye and copperas:

FeSO4 + 2OH[−] → Fe(OH)2 + SO4[2−]

After washing the product several times with distilled water, it could then be added to a saturated solution of copperas, refluxed gently, filtered through clean kieselgur, boiled, and finally recrystalized by slowly cooling the mother liquid.

[Edited on 5/4/2007 by obsessed_chemist]

UnintentionalChaos - 4-5-2007 at 16:17

Glycerin isn't that hard to find fairly dry. If you buy "swan" brand glycerin, I'd consider using it straight up. It's meant as a skin care product, so I do not doubt the purity, and it is really dry when fresh. I realized this the other day when I tried placing a drop of it on KMnO4 which to my great confusion caused absolutely no reaction, even when left for several minutes. It eventually occured to me that some water must be necessary to initiate it and with a drop of water, the pile started fuming and then burst into flames as expected.

not_important - 4-5-2007 at 17:20

Quote:
Originally posted by obsessed_chemist
A little off-topic, but...

I've contemplated a procedure for purifying garden-quality copperas (ferrous sulfate, heptahydrate). AFAIK, this source tends to be acidic from excess sulfuric acid used in it's synthesis ...

[Edited on 5/4/2007 by obsessed_chemist]


Adding ferrous hydroxide will produce basic sulfates, and moving the pH higher may decrease the stability of the Fe(II) around oxygen. Just recrystallise the FeSO4 without adding hydroxide.

UnintentionalChaos - 4-5-2007 at 18:32

Also, your synthesis uses H2SO4 in it, does it not? So why bother removing leftover acid?

[Edited on 5-4-07 by UnintentionalChaos]

12AX7 - 4-5-2007 at 18:51

Recrystallization with a hot filtering step is all that is required. Wash and dry the crystals to remove more impurities (including free acid).

Tim

obsessed_chemist - 6-5-2007 at 18:17

^ - I was unaware that 'basic' iron sulfates existed. As I said, just an idea I had swimming around for a while; thanks for debunking the myth guys. I stand corrected.

woelen - 7-5-2007 at 11:23

Quote:
Originally posted by 12AX7
Sure of that? I've left out crystals that partially dehydrated without turning to brown mush.

Tim

There may be some efflorescence, but making real anhydrous FeSO4 is a completely different thing. It is my experience that making anhydrous salts from hydrated ones is very difficult, except for a few alkali metal salts. For almost all other metal salts, there is (at least partial) hydrolysis, which results in loss of acid and formation of basic salts. For some salts with strongly oxidizing anions, or strongly reducing cations, things even become worse, due to decomposition and/or internal redox reactions. Ferrous sulfate is an example of the latter.

ordenblitz - 7-5-2007 at 19:52

I made mine by accident.
I placed my freshly crystallized and lovely blue FeSO4*7H2O in an argon flushed glove box to dry. By the second day it had turned completely white and fallen to dust losing all of its water in the perfectly dry atmosphere.

UnintentionalChaos - 7-5-2007 at 20:22

IIRC that is most likely the monohydrate, not the anhydrous form.

EDIT: found where I read that http://www.sas.org/tcs/weeklyIssues_2005/2005-07-15/feature2...

It says you may have had a combination of tetrahydrate and monhydrate. You could conduct the mistake again as an experiment and weigh the sample before and after to get a definitive answer.

[Edited on 5-7-07 by UnintentionalChaos]

cnidocyte - 13-1-2011 at 16:45

I made some ferrous sulfate a while ago by adding conc. H2SO4 to steel wool, at first I thought it didn't work and I left it overnight to react fully but when I was cleaning up the next day I noticed a pile of green crystals at the bottom of the solution. It was definitely ferrous sulfate but for some reason it reeked. It had a bit of an H2S smell to it but it was something else, something I'd never smelled before. Wasn't a pleasant smell and even after cleaning with water the flask still reeked of the stuff. Is that what ferrous sulfate smells like or was it impurities I was smelling?

Also I noticed this stuff barely dissolved at all in cold water but dissolved rapidly in hot water and made a brownish yellow solution.

Magpie - 13-1-2011 at 17:03

I don't know what the smell is - is it a metallic smell?

This thread may help:

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

Picric-A - 14-1-2011 at 05:05

The dissoloution of the ferrous sulphate probably reduced the H2SO4 to H2S and other noxious sulphur compounds which is probably reliable for the smell...
I always find pure FeSO4 has a metallic smell, cant quite describe it but if you smell it you will know what i mean...