Sciencemadness Discussion Board
Not logged in [Login ]
Go To Bottom

Printable Version  
 Pages:  1  ..  4    6    8  ..  12
Author: Subject: H2SO4 by the Lead Chamber Process - success
12AX7
Post Harlot
*****




Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline

Mood: informative

[*] posted on 21-7-2008 at 12:31


Read the metaphosphoric acid thread. Glassy fused HPO3 has been used to prepare SO3.

Tim




Seven Transistor Labs LLC http://seventransistorlabs.com/
Electronic Design, from Concept to Layout.
Need engineering assistance? Drop me a message!
View user's profile Visit user's homepage View All Posts By User This user has MSN Messenger
497
International Hazard
*****




Posts: 778
Registered: 6-10-2007
Member Is Offline

Mood: HSbF6

[*] posted on 28-9-2008 at 13:10


Here's another patent on the decomposition of basic iron sulfate to SO3 and Fe2O3. It gives a little different reaction scheme than the other patent did. I tend to believe this patent as it is much newer.

4FeSO4 + H2SO4 + O2 --> Fe4O(SO4)5 + H2O

Then

Fe4O(SO4)5 --> 2Fe2O3 + 5SO3

The first step proceeds in a slurry of FeSO4 in semi-concentrated H2SO4 just below the boiling point of the mixture (which depends on water content of the H2SO4).
The second step occurs at 500-700*C, apparently without air.

So...

Mix 600g anhydrous FeSO4 (or equivalent hydrated) with 300g 33% H2SO4 battery acid (or equivalent concentrated) and heat with stirring to over 150*C in a stream of air, maybe with a heat gun. It should eventually harden into a solid cake of about 700g after all the water is driven off. Crush up the cake and load it into a makeshift retort and heat to ~600*C for a while. Yield: 300-350g SO3!

Sounds pretty good to me. Anything wrong with that process?
View user's profile View All Posts By User
DJF90
International Hazard
*****




Posts: 2266
Registered: 15-12-2007
Location: At the bench
Member Is Offline

Mood: No Mood

[*] posted on 28-9-2008 at 13:32


It looks too good to be true, but I guess the only way to find out is to try it :D
View user's profile View All Posts By User
Picric-A
International Hazard
*****




Posts: 796
Registered: 1-5-2008
Location: England
Member Is Offline

Mood: Fuming

[*] posted on 6-10-2008 at 10:25


Are you sure the ferrous sulphate can be oxidised that easily?
I would of thought it would take hours to oxidise it simply with air...
If not it is an extremly easy way to lots of SO3 :D

Since making that post i have found out that the oxidation of FeSO4 occurs rapidly at high Ph so Reacting FeSO4 with hot H2SO4 whilst bubbling air through it should oxidise it pretty quick!

[Edited on 6-10-2008 by Picric-A]
View user's profile View All Posts By User This user has MSN Messenger
not_important
International Hazard
*****




Posts: 3873
Registered: 21-7-2006
Member Is Offline

Mood: No Mood

[*] posted on 6-10-2008 at 17:23


High pH means neutral to alkaline, not acidic. Increased temperature will speed the reaction, though.
View user's profile View All Posts By User
Picric-A
International Hazard
*****




Posts: 796
Registered: 1-5-2008
Location: England
Member Is Offline

Mood: Fuming

[*] posted on 6-10-2008 at 23:09


Of course, sorry for that stupid mistake. Too tired last night:P
View user's profile View All Posts By User This user has MSN Messenger
Rosco Bodine
Banned





Posts: 6370
Registered: 29-9-2004
Member Is Offline

Mood: analytical

[*] posted on 27-10-2008 at 00:11


For the precursor desired for the pyrolysis ...

I am wondering if perhaps a synthesis of ferric sulfate
from ferrous sulfate via H2O2 and then partial hydrolysis,
then dehydration of precipitated copiapite may also work.

4 FeSO4 + 2 H2O2 + 2 H2SO4 ---> 2 Fe2(SO4)3 + 4 H2O

Fe2(SO4)3 + 2 H2O ----> 2 Fe(OH)SO4 + H2SO4

Since half the H2SO4 required for the first reaction is regenerated via hydrolysis, the algebraic sum would possibly adjust the initial reaction minimum H2SO4 stoichiometric requirement to 1 H2SO4

2 Fe(OH)SO4 + Fe2(SO4)3 + 17 H2O ----> Fe4(OH)2(SO4)5-17 H2O ( copiapite precipitate )

Fe4(OH)2(SO4)5 - 17 H2O -----> Fe4O(SO4)5 + 18 H2O
View user's profile View All Posts By User
Formatik
International Hazard
*****




Posts: 927
Registered: 25-3-2008
Member Is Offline

Mood: equilibrium

[*] posted on 28-10-2008 at 15:34
H2SO4 from ozone


I was reading the wiki entry on O3 and it mentioned the following reaction between elemental sulfur and ozone to form SA:

S + H2O + O3 = H2SO4

But no details. Does anyone know more about this like reaction time and conditions?
View user's profile View All Posts By User
12AX7
Post Harlot
*****




Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline

Mood: informative

[*] posted on 28-10-2008 at 20:36


Well, that's trivial, but ozone is rather harder to generate in quantity than SO2, or SO3 for that matter. And one could argue the SO3 is safer, pound for pound, than that much ozone.

Tim




Seven Transistor Labs LLC http://seventransistorlabs.com/
Electronic Design, from Concept to Layout.
Need engineering assistance? Drop me a message!
View user's profile Visit user's homepage View All Posts By User This user has MSN Messenger
Rosco Bodine
Banned





Posts: 6370
Registered: 29-9-2004
Member Is Offline

Mood: analytical

[*] posted on 28-10-2008 at 22:09


Using H2O2 instead of air oxidation for the process described
by 497 above ....

I'm still wondering if the following summary reaction derived from my reactions above wouldn't be the easiest way to form a precursor for pyrolysis to SO3.

4 FeSO4 + 2 H2O2 + H2SO4 + 15 H2O ---> Fe4(OH)2(SO4)5-17 H2O ( copiapite precipitate )

Fe4(OH)2(SO4)5 - 17 H2O -----> Fe4O(SO4)5 + 18 H2O

Fe4O(SO4)5 ----> 2 Fe2O3 + 5 SO3


[Edited on 29-10-2008 by Rosco Bodine]
View user's profile View All Posts By User
497
International Hazard
*****




Posts: 778
Registered: 6-10-2007
Member Is Offline

Mood: HSbF6

[*] posted on 29-10-2008 at 18:17


That would be nice if that worked. I don't see why it wouldn't. According wikipedia it can be prepared with nitric acid as the oxidizer. They call it ferric subsulfate or basic ferric sulfate. It is apperently used as some sort of medical treatment for certain skin problems.. I'll do some more looking around.
View user's profile View All Posts By User
Rosco Bodine
Banned





Posts: 6370
Registered: 29-9-2004
Member Is Offline

Mood: analytical

[*] posted on 29-10-2008 at 23:35


Yeah this sort of reminds me of the line of thinking which I had going when contemplating a pyrolysis precursor for
calcium cyanamide. If you can get right to the immediate precursor via some preliminary workup which eliminates
rotary kilns and other steps which are more convenient for
industry ....then you are a lot closer to a worthwhile lab scale method. The temperature and pH will probably affect the density of the precipitated copiapite ...and the only concern I have there is possible gelling .....but boiling and agitation would probably break that higher hydrate.
Sometimes those "superhydrates" are unstable transition
species which form nicely crystalline lower hydrates or
even anhydrous derivatives....and I agree it would be nice if this one behaves in that way. It would probably be the
easiest route to SO3 and oleum which has been proposed so far in any discussions here. So then SO3 would be
"pyrolitic distillate of copiapite anhydride" :D

[Edited on 30-10-2008 by Rosco Bodine]
View user's profile View All Posts By User
497
International Hazard
*****




Posts: 778
Registered: 6-10-2007
Member Is Offline

Mood: HSbF6

[*] posted on 29-10-2008 at 23:53


Well I found lots of stuff on using H2O2 to oxidize ferrous to ferric sulfate, but so far not much on making basic sulfate. Since O2 can do it under the right conditions, I see no reason H2O2 couldn't. It doesn't look like gels would be a problem, so far all everything I've found has said "crystalline" I haven't had a lot of time to work on this lately.. I'll try to post more soon.

US 2563623 is kind of interesting..
US 3529957 might be useful

Edit:
According to this, H2O2 will in fact work. They used it as a control for studying some bacterial oxidation crap, but it should still be useful data. Apparently the precipitate will have a greatly varying Fe:SO4 ratio depending on various conditions. At pH 2.5 (with 2%H2O2!) they got a ratio of 6:1 Fe:SO4, which is far from ideal.. But I think at a lower pH and more concentrated H2O2 (and maybe some other additives), the ratio will get much better. The decomposition should go similarly at different ratios right? It definately requires some testing though. If you wanted you might even be able to optimize the ratio to get the lowest decomposition point, I don't know.

Another idea that's probably less practical, but still interesting: an acid solution of FeSO4 is known to reduce N2O4 to NO and form Fe2(SO4)3. I wonder if the conditions were right it could be oxidized further to a basic sulfate (or subsulfate, oxysulfate, hydroxysulfate, or whatever else you want to call it)? It would be an interesting process because the NO could easily be recycled.

The other thing one could work on would be figuring out an easy effective high temperature air oxidation route that could be used instead. Still it's hard to see how that could be easier that mixing a couple solutions and having your precourser all ready to go..

[Edited on 30-10-2008 by 497]
View user's profile View All Posts By User
Rosco Bodine
Banned





Posts: 6370
Registered: 29-9-2004
Member Is Offline

Mood: analytical

[*] posted on 30-10-2008 at 01:35


The reactions I listed are valid and previously reported reactions, so I was already confident about the reactions before proposing them . The significant unknowns I have there are with regards to the rate of reaction, the most favorable pH , temperature and the physical form of the end product. It's like... this should work, but how well it will work I'm not sure, but it would be worth an experiment.
I think the concentration of the reactants would probably
be dilute and the solutions would be hot and perhaps even brought to boiling, basically an open simmering cauldron
reaction where the H2O2 simply accellerates what air would do given more time.

[Edited on 30-10-2008 by Rosco Bodine]
View user's profile View All Posts By User
497
International Hazard
*****




Posts: 778
Registered: 6-10-2007
Member Is Offline

Mood: HSbF6

[*] posted on 30-10-2008 at 01:44


According to the paper I linked the reaction time is apparently pretty fast.. But as far as optimum temp, pH, etc, that is anybodies guess. I seriously doubt there is much if any available information on these specific conditions. So experimentation is the name of the game as far as I can tell. And I certainly do think it is worth experimenting with, as you said before, this could be the easiest route to SO3 yet. And I really would like to have some SO3! :D
View user's profile View All Posts By User
Rosco Bodine
Banned





Posts: 6370
Registered: 29-9-2004
Member Is Offline

Mood: analytical

[*] posted on 30-10-2008 at 01:53
here's a couple of patents


US3078180 H2O2 oxidation of Ferrous Sulfate

Attachment: US3078180 H2O2 oxidation of Ferrous Sulfate.pdf (356kB)
This file has been downloaded 771 times

View user's profile View All Posts By User
Rosco Bodine
Banned





Posts: 6370
Registered: 29-9-2004
Member Is Offline

Mood: analytical

[*] posted on 30-10-2008 at 01:57
and the second patent


US3574599 Copiapite Basic Ferric Sulfate

A couple more of interest are

US2905533 Basic Ferric sulfate

US2413492 Purification Crystallization of Ferrous Sulfate

Attachment: US3574599 Copiapite Basic Ferric Sulfate.pdf (256kB)
This file has been downloaded 767 times

View user's profile View All Posts By User
Rosco Bodine
Banned





Posts: 6370
Registered: 29-9-2004
Member Is Offline

Mood: analytical

[*] posted on 30-10-2008 at 02:01


Looks like I should just post those other two patents also
so here they are

Attachment: US2905533 Basic Ferric sulfate.pdf (184kB)
This file has been downloaded 1205 times

View user's profile View All Posts By User
Rosco Bodine
Banned





Posts: 6370
Registered: 29-9-2004
Member Is Offline

Mood: analytical

[*] posted on 30-10-2008 at 02:02


and here's the last one

Attachment: US2413492 Purification Crystallization of Ferrous Sulfate.pdf (249kB)
This file has been downloaded 1858 times

View user's profile View All Posts By User
497
International Hazard
*****




Posts: 778
Registered: 6-10-2007
Member Is Offline

Mood: HSbF6

[*] posted on 30-10-2008 at 02:08


:P I saw all of those, except for the one about enameling, when I was searching earlier.. Some useful information but I didn't see anything too applicable or specific, at least in terms of synthesis of the basic sulfate via H2O2.

Edit;
Are there really that many people watching this thread? Each attachment you posted was downloaded about 10 times within a minute. Strange..

Edit2:
I just remembered, I neglected to post any info on an interesting patent I saw. It talked about using NaClO3 with iron sulfate to produce mixed oxysulfate salts. I didn't read the whole thing and I don't know how the decomposition properties are, but they looked interesting. Well, not too interesting to me since I have no easy source of chlorate, but maybe for someone else.

here it is.. ugh I'm accumulating such a huge mass of PDFs.. And they're all so disorganized, I feel sad just thinking about trying to organize them all.

[Edited on 30-10-2008 by 497]

Attachment: PRODUCTION_OF_IRON_OXIDE_AND_IRON_FREE_O-1.pdf (177kB)
This file has been downloaded 885 times

View user's profile View All Posts By User
Rosco Bodine
Banned





Posts: 6370
Registered: 29-9-2004
Member Is Offline

Mood: analytical

[*] posted on 30-10-2008 at 02:10


You kind of have to pick the pieces parts from each one
and then interpolate :P

Yeah I saw all that interest and figured I better give it up :D

It looks to me like there would need to be some additional
H2SO4 in that sodium chlorate oxidation.....and it also
seems that reaction is adaptable to the use of ordinary bleach instead of sodium chlorate. Of course you end up
with salt solution as a byproduct.

[Edited on 30-10-2008 by Rosco Bodine]
View user's profile View All Posts By User
497
International Hazard
*****




Posts: 778
Registered: 6-10-2007
Member Is Offline

Mood: HSbF6

[*] posted on 30-10-2008 at 02:33


And here's another patent. I'm not sure how useful it is, but it does talk a little about alternative oxidizers.. Maybe ammonium persulfate would be useful? Or various peroxyhydrates?

Anyway, it was nice corresponding semi-instantly with you Rosco, but alas it is 2:30 AM and I am very tired. Talk to you later.

Yes bleach might be a good idea, have to try that. As long as you could do it without getting too much sodium caught along with the Fe. I considered that a while ago but dismissed it for some reason, can't remember why... :o

[Edited on 30-10-2008 by 497]

[Edited on 30-10-2008 by 497]

Attachment: Process_of_preparing_a_preferred_ferric_.pdf (228kB)
This file has been downloaded 1903 times

View user's profile View All Posts By User
497
International Hazard
*****




Posts: 778
Registered: 6-10-2007
Member Is Offline

Mood: HSbF6

[*] posted on 30-10-2008 at 13:46


Would this be the reaction?

4FeSO4 + 2NaOCl + 3H2SO4 = Fe4(OH)2(SO4)5 + 2NaHSO4 + 2HCl

I wonder if something would have to be added to neutralize the HCl?

Edit:
I just found some patents on ferrate synthesis.. interesting stuff. FeO4-- is supposed to be a stronger oxidizer than MnO4--. It is said to be produced by oxidizing Fe3+ salts with concentrated hypochlorite. It releases O2 in acid solutions, maybe you could use it to synth basic sulfate? It looks useful for many other things too. Maybe via this:

3FeSO4 + Na2FeO4 + 3H2SO4 --> Fe4(OH)2(SO4)5 + Na2SO4 + 2H2O

Then you wouldn't have to deal with chlorine and other crap in there. It should be a strong enough oxidizer at least..

[Edited on 30-10-2008 by 497]
View user's profile View All Posts By User
Rosco Bodine
Banned





Posts: 6370
Registered: 29-9-2004
Member Is Offline

Mood: analytical

[*] posted on 30-10-2008 at 19:03


I'll have to check further but I would expect that the
basic ferric sulfate is going to require moderately basic to
neutral or only very slightly acidic conditions to form ....
maybe pH 10 to pH 6.5 for example.

Your first equation proposed above would be too acidic.

You are going to have intermediate hydrates and hydrolysis
reactions to consider ....so it is going to be an algebraic
stoichiometry. I'll have to work it out later for the bleach.

There was a method I think in one of the lead oxide related
threads where bleach was used to precipitate lead oxide from lead salt solutions and the method for iron may be similar.

[Edited on 30-10-2008 by Rosco Bodine]
View user's profile View All Posts By User
497
International Hazard
*****




Posts: 778
Registered: 6-10-2007
Member Is Offline

Mood: HSbF6

[*] posted on 30-10-2008 at 20:35


Quote:

would expect that the basic ferric sulfate is going to require moderately basic to neutral or only very slightly acidic conditions to form .... maybe pH 10 to pH 6.5 for example.


I don't know about that, according to that paper I linked to, they got a precipitate that was like 99:1 Fe:SO4 when precipitated at pH 9, a 20:1 ratio at 6, and a 6:1 ratio at 2.5 (with H2O2).So you'd have to have it below pH 1 probably, to get any decent ratio? Unless there's some other factors that could be changed? I wonder how different oxidizers effect the precipitate differently?

Here's the important part of the paper I was talking about..

[Edited on 30-10-2008 by 497]

Attachment: precipitated iron.doc (94kB)
This file has been downloaded 819 times

View user's profile View All Posts By User
 Pages:  1  ..  4    6    8  ..  12

  Go To Top