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Author: Subject: SO3 from Iron Sulphate
Eclectic
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[*] posted on 12-9-2007 at 13:32


If you order ferric sulfate in a 50lb bag, expect to get the NONAHYDRATE, not the monohydrate.

Dehydration in an oven should not be too difficult, unless the material dissolves in it's own water of crystallization.

Edit: So far so good at 300F in a glass tray. My supplier was totally unreceptive to the idea that he should make a price adjustment or at least edit his product description. :mad:

[Edited on 9-12-2007 by Eclectic]
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Sauron
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[*] posted on 12-9-2007 at 16:04


Yeah the stuff is very hygroscopic

I reckon a big sack once opened will need to be repacked immediately into airtight bottles.

I do have a drying oven by Memmert, good sized one, thermostated and exhausted.
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Eclectic
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[*] posted on 12-9-2007 at 17:58


It doesn't seem to be very deliquescent.

5 gallon plastic buckets with Gammaseal lids are convenient for storing this sort of thing.

http://freckleface.com/shopsite_sc/store/html/gammaseals.htm...

Edit: No problem with melting or fuming while drying in an oven at up to 450F. Probably no predrying is needed. I bet this stuff would make a great alcohol dehydration catalyst. ;)

[Edited on 9-12-2007 by Eclectic]

OK, yes, it is deliquescent. Not as bad as calcium chloride, but a few granules left on the counter in an air conditioned environment will start to get damp after 4-5 hours.

[Edited on 9-13-2007 by Eclectic]
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ciscosdad
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[*] posted on 8-10-2007 at 18:37
Ferrous vs Ferric


Has anyone wondered why the traditional method of vitriol production used the ferrous salt?
There seems to be absolutely no reference to the ferric at all in the preps.
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[*] posted on 8-10-2007 at 19:36


Quote:
Originally posted by ciscosdad
Has anyone wondered why the traditional method of vitriol production used the ferrous salt?
There seems to be absolutely no reference to the ferric at all in the preps.


Ferrous sulphate occurs naturally as copperas or melanterite. It was also made readily from iron pyrites (fool's gold, FeS2) a very common mineral, by the action of air and moisture, the resulting solution being treated with scrap iron.

Ferric sulphate is normally made from ferrous sulphate. It could be considered as an "intermediate compound" in the production of sulphuric acid from ferrous sulphate.

Regards, Xenoid
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[*] posted on 8-10-2007 at 19:39


The ferrous salt gives SO3 in H2SO4 (oleum) and some losses to SO2.

The ferric salt gives neat SO3.




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[*] posted on 8-10-2007 at 19:50


Thanks guys.
I understand that ferric is better as there is no loss of SO3 oxidising the ferrous to ferric during roasting. It just seems strange that the old chemists didn't pick up on the refinement given the obvous improvement in yield. The copperas oxidises easily enough with air. Perhaps that was one of the trade secrets that never got published.

@Sauron. Sorry to jiggle your elbow here, but we are waiting with bated breath for your account of Oleum production using this method.:P
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[*] posted on 8-10-2007 at 19:58


Everything awaits installation of fume hood, and that awaits sufficient funds to start. Fume hood and scrubber will cost me $7000 (it's a big hood!) and this has been a crappy year for $$.



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ciscosdad
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[*] posted on 18-10-2007 at 15:37
Ferrous vs Ferric


Ref p244 - 5 of the Knapp "Chemical Technology" vol 1
Right at the end of the Oleum article.
(posted by S.C.Wack earlier in this thread).
It implies that they lightly air roasted the ferrous sulfate to form (basic) ferric sulfate. They were certainly aware of the advantages.
Now I know. Those old guys were not so silly eh?
I particularly appreciate the older generation of books that focus on the "how to" and the "what happens". The modern obsession with the "why is it so" is depriving us of most of the fun in Chemistry.

BTW. Thanks a million for the book S.C.Wack. I am reading it cover to cover. If you are aware of anything similar, please post that too.
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[*] posted on 19-10-2007 at 08:45


I think I might give this a shot on a large scale. Anyone think a stainless stock pot and some stainless tubing would be up to the task? (I'm envisioning building a fire around it, then having the SO3 lead into sulfuric). I think the heating vessel is the biggest problem with the whole method.

[Edited on 19-10-2007 by Fleaker]




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[*] posted on 19-10-2007 at 13:17


An old steel CO2 tank might be better.
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[*] posted on 21-10-2007 at 15:04


Charcoal and a draft may be a better heat source than wood.
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[*] posted on 4-2-2008 at 13:31


I made some anhydrous ferric sulfate today, and want to warn people that the reaction of Fe2O3 and H2SO4 can be violent, something I never saw mentioned anywhere.

Three different sources say that ferric sulfate can be made by strong heating of ferric oxide with conc. sulfuric acid, and this sounds like a better method of production than aqueous methods since only little water (that arising from the reaction) has to be driven off afterwards.
Well, the reaction even drives out the water by itself, in a spectacular manner!

Here's what I did:
I mixed 30g Fe2O3 (commercial red pigment) and 60g conc. H2SO4 (roughly stochiometric amounts) together in a beaker, forming a paste without reacting.
I then proceeded to heat this with a flame while stirring. As it became hot, at one point the entire mass suddenly started to bubble uncontrollably and then immediately erupted into a fountain of steam!
The solid/liquid contents of the beaker stayed in there, however. The mix was very hot and of a thick pasty consitence after that, and I continued to heat it while stirring until a dry brownish powder resulted, giving off some more steam in the process.
Its weight corresponded very well to the expected weight of neutral anhydrous ferric sulfate from this reaction, and I put it into a jar for storage.

The violent reaction makes this method of preparation unsuitable for direct upscaling. One would have to heat the Fe2O3/H2SO4 slurry in portions of no more than 100g each.

The final dehydration could then be done in an oven.

For making up to 100g ferric sulfate at a time, the method I described is very fast and convenient.

How the product performs in the SO3 synthesis will be the next subject of investigation.

[Edited on 4-2-2008 by garage chemist]




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[*] posted on 4-2-2008 at 14:23


Well, if I go that route, I will do the 100 g Ferric sulfate in a 1000ml beaker. I wonder at what temperature your reaction "took off"? I wonder if a hot spot from your flame far exceeded the temperature of startup. When you said pigment grade iron oxide I imagined distinctive red oxide as apposed to the more orangy more hydrated variety. Orangy is a word because I just now invented it.:D:D



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[*] posted on 4-2-2008 at 14:42


There are many different types of iron oxide reds available, some of them not even pure Fe2O3 (I hope mine is pure!).
I have two different ones, one (source somewhat dubious and long gone) is more brightly red, the other (from a ceramics supplier) is more brownish. I used the bright red one in the synth because I had less than 100g left of it and wanted to use it up before opening the bag of the brownish one.

I originally wanted to heat the reaction mix as uniformly as possible and therefore stirred very well. So the whole reaction mix was brought close to the takeoff temperature instead of only one hotspot, which may have contributed to the remarkably violent reaction.




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[*] posted on 4-2-2008 at 15:56


OK, well I have some bright red iron oxide and I will try this in 1 gram Fe2O3 vs. drops of H2SO4 scale in a test tube to just see for myself how violent it really is. Hopefully I can find a suitable thermometer to measure the reaction temperature.



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