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Author: Subject: oleum & SO3
Organikum
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cool.gif posted on 25-5-2003 at 14:35
oleum & SO3


Here we go
HARDCORE:
Oleum, smoking H2SO4, 50%SO3 in H2SO4, SO3

First I suggest everyone who wants to tinker with this may inform himself on the truely existing dangers of these compounds. They are not nasty - they are dangerous!

Madscientist passed H2SO4 over MgSO4 for oleum (did this really work?) he wrote and in more than one synth SO3 was asked for (acetic anhydride for example, nitric acid).
So here is a doable way to produce oleum and SO3:
- Ferrous sulfate septahydrate is made from iron (filings/steelwool) and H2SO4.
- Ferrous sulfate septahydrate dehydrates completely by heating >300°C.
- Anhydrous Ferrous sulfate decomposes to SO3 at 480°C and can be condensed to oleum (or the SO3 can be directly used of course).

If the oleum is what one is after it should be possible to decompose and condense in the same vessel without extern apparati necessary. The vessel has to be big and strong enough - thats understood.

:D:D:D:D:D

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[*] posted on 25-5-2003 at 16:31


Organikum: I think my oleum attempt from long ago didn't actually work - I probably just thought it did (considering my relative ignorance back then, and the fact that I made way too many assumptions).



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[*] posted on 25-5-2003 at 18:26


xoo1246: Yes me too! The howto on the ferreous was only included for completeness and the sure to come question on it. Sooner or later.......

Madscientist: No problem for me. My complete errors and selfcaused mishappenings will be published soon. Five volumes. Open end. Working title: "Oh! Oh! Oh! - Oh shit!" Probably the "shit" has to be removed so.......
....oh, oh, oh. ;)




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[*] posted on 25-5-2003 at 23:31


Quote:
Originally posted by Organikum
- Ferrous sulfate septahydrate is made from iron (filings/steelwool) and H2SO4.
- Ferrous sulfate septahydrate dehydrates completely by heating >300°C.
- Anhydrous Ferrous sulfate decomposes to SO3 at 480°C and can be condensed to oleum (or the SO3 can be directly used of course).


Heating the ferrous sulphate must be done in an oxygen free atmosphere, otherwise it will oxidise into the ferric state.

Another good source of ferrous sulphate is 'Iron Mordant', whch is available from all good dyeing suppliers.




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Organikum
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[*] posted on 26-5-2003 at 00:07
You are right, fish you are!


The setup has to be closed against the outside what should be understood for the unhealthy properties of SO3 alone. For the ease of availability of the reactands I would say that it´s not necessary to purge with an inert gas, but it´s easier to sacrifice a part of the ferrous sulfate to the oxygen inside. There should be enough left to be converted to SO3 - the efficiency of these rather old and raw methods isn´t very good anyways I believe.

Exactly for such critics it got posted. Thanks fish you are! ;)

[Edited on 26-5-2003 by Organikum]




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[*] posted on 26-5-2003 at 01:00


Industrially, the SO3 is absorbed in circulating concentrated H2SO4 which is then diluted with water. This is done to attenuate the violence of reacting SO3 directly with water.
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[*] posted on 26-5-2003 at 04:53


Why not use CuSO4? That can't be oxidized any further.



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smile.gif posted on 26-5-2003 at 06:55
higher temperatures?


CuSO4 starts to decomposes at a temperatures of 560°C whereby FeSO4 is completely decomposed at 480°C. Thats how I understood it at least. In the original process costs will have been a factor also - I didn´t invent this, I just don´t know where this was taken from, if I knew it, I would have named it.
CuSO4 maybe worth a try, also I can´t see any advantages. as the FeSO4 scrubbing the oxygen seems to me favorable.

The humidity in the reactor should be driven out during heating the FeSO4 for becoming anhydrous what should be done in the same vessel anyways. I propose that vessel and FeSO4 will be in the same state of anhydrous condition so. ;)
(Ah! The beauty of strong causal logic!)

btw. this is a concept for discussion, not a complete running process by now. So the formulation: "Why don´t you use..." is inadequate here.




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[*] posted on 26-5-2003 at 09:20


What about heating H2SO4 ? The white fumes are mainly because of SO3 IIRC.
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[*] posted on 26-5-2003 at 09:53


no, a_bab, the white fumes of heated (boiling ) sulfuric acid is just sulfuric acid droplets. If H2SO4 is heated , it will split to SO3+ H2O at high temperatures and the SO3 will decompose into SO2 and O2 at these temperatures.

Better than FeSO4 is Fe2(SO4)3 since +3 ions decomposes into SO3+SO2 at lower tempertures. Expect a large loss of SO3 due to eq. SO3 = SO2 + 1/2 SO2

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[*] posted on 27-5-2003 at 16:16


I think that heating Sodim bisulfate may well work... I'm not exactly sure of the temperatures, but I do know they're accessable to a gas flame. If I remember, you dry the bisulfate in an oven, then heat it over a gas flame, and water gets produced. You keep the heat so that water is just getting produced, and you're getting solid sodium pyrosulfate, until water stops forming. Then crank up the heat to decompose the sodium pyrosulfate and get sodium sulfate and sulfur trioxide. then the best thing to do is dissolve the gas in cooled, concentrated H2SO4. It can work if you use a distillation setup - change flasks between the water production and the decomposition of the solid. Although I dont think you'd want to be too attached to the flask that you're heating - it may suffer some permanent damage if you're not careful. I dont actually have a reference for this, mind you - I just remember seeing it somewhere some time ago and thinking "hmm... must try this one day"... never got around to it... must remedy that. Oh, by the way, Sodium Bisulfate is commonly used for raising the pH of swimming pools... very cheap and easy to find
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[*] posted on 28-5-2003 at 08:59


Yes rikkitikkitavi, you are right. I was confused because of the lab method wich is using oleum to get SO3.

You are right aswell ziqquratu, a simple method could be heating NaHSO4 or KHSO4. The reaction will be in two stages, one at 200-300 degrees C, when the alkali pyrosulfate is formed, and then the other at 500 degrees C when SO3 is formed because of the thermal decomposition of the pyrosulfate.
The reactions are:

1. 2NaHSO4 --->Na2S2O7+H2O
2. Na2S2O7---->Na2SO4+SO3

[Edited on 28-5-2003 by a_bab]
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[*] posted on 28-5-2003 at 10:46


NaHSO4 for raising the pH?

HSO4- has a pKa of 1,92 which means it will behave like an acid in solution. Especially in diluted solution.




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[*] posted on 28-5-2003 at 15:58
Oops!!


Sorry, I didnt see that one!! Sodium bisulfate makes pH lower!! I'd have simply edited my above post, but then anyone reading what vulture wrote would have no idea what was going on!! Thanks for pointing that one out
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[*] posted on 29-5-2003 at 00:09
a short resumee


- by principle the process will work and includes no dangers like sudden decomposition (explosion) which would render it useless for the amateur experimenter.
- precautions: drying the used apparatus which has to be a enclosed one.
Also the apparatus must be VERY acid resistant. Ceramics, porcellain, clay, not glass or metal (to stay on the safe side)
- Fumehood or outside with fan at least.
- Other starting compounds than FeSO4 seem to be possible.

Please correct me if I am wrong
thanks
ORG

Next step would be a layout for the apparatus - I thought on a modified KLIPP principle for keeping it small. :D

[Edited on 29-5-2003 by Organikum]




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[*] posted on 1-9-2003 at 19:55


destructive distillation of green vitriol(FeSO<sub>4</sub>.7H<sub>2</sub>O) is a very traditional way of making H<sub>2</sub>SO<sub>4</sub> and was first done by hayyan (don't know its English spelling)
Organikum! are you sure the 480°C is the decomosition point for ferrous sulfate? it is the exact temp for ferric too.

is ferric hydrogen sulfate, anhydrous or hydrate? if it is anhydrous or can be dehydrated at a lower temp than its decompositin point, it's possible to make oleom without need of conc sulfuric acid! u simply add dilute H<sub>2</sub>SO<sub>4</sub> to ferric sulfate/oxide/hydroxide.. recrystalize the salt and distill off.
2Fe(HSO<sub>4</sub>;)<sub>3</sub> => Fe<sub>2</sub>O<sub>3</sub> + 3H<sub>2</sub>S<sub>2</sub>O<sub>7</sub>
it may lose H<sub>2</sub>O forming pyrosulfate first, (but my sence doesn't approve it.)
the pyrosulfate method is great for those who distill their own nitric acid. just dilute the wast acid => recrysalize the alkali metal hydrogen sulfate.

[Edited on 2-9-2003 by KABOOOM(pyrojustforfun)]
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[*] posted on 23-5-2004 at 14:07


We don't have a thread started for the reactions of SO3 but I figured this thread would do since I don't feel it necessary to start a new one.

"Sulfur trioxide reacts with KClO3 to form K2S3O10 and a red product assumed to be (ClO2)2S3O10 (mp 75C)" The red product has a violent decomposition at 100C. An anologous reaction occurs with KClO4.

On topic it appears that the decomposition temperature of ferric sulfate is significantly higher then 480C. Here are two passages from the Complete Tretise on Inorganic Chemistry that are relevent:

"The yield of fuming acid is poor because the temp. of dissociation of ferric sulfate is so high that a large proportion of of sulfur trioxide is decomposed at the same time."

And relating to that: "The dissociation pressure of the trioxide from ferric sulfate was found by G. Keppeler and J. d'Ans to be only 15mm. at 640C so that the temp of the retorts must be much higher then this."

Other interesting information relating to things said upthread:

"P.G. Prelier proposed to heat the alkali or alkaline earth sulfates with sulfuric acid to form the hydrosulfate, and to distill the water from the hydrosulfate, and the fuming acid or sulfur trioxide from the pyrosulfate. R.W. Wallace proposed a modification of this process; adn W. Wolters said that the liberation of sulfur trioxide from the pyrosulfate occurs at a lower temp if some magnesium sulfate is present. This means that less sulfur trioxide is dissociated. W. Wolters also proposed to heat the pyrosulfate with anhydrous sulfuric acid so as to distill off the sulfur trioxide:
N2S2O7 + H2SO4 ---> 2NaHSO4 + SO3
A.M. Leon proposed to electrolyze the monohydrated sulfuric acid so that the water decomposes into hydrogen and oxygen, and the resulting sulfur trioxide dissolved in the electrolyte to form the fuming acid."

As for the last part though, I doubt graphite would survive and I would have to guess the electrodes would have to be made of something like Pt.




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[*] posted on 23-5-2004 at 19:24


Wolters' patent with MgSO4: DE6091. With H2SO4: DE12295. Very uninformative. I think Wallace had a temp of 600C going.

GB237243 mentions a temp of 750-780 for the ferric sulfate deal.

GB 124988 is about sawdust and bisulfate.

In US342784 and 342785, clay and gypsum give cement and SO3. Like most thermal methods, expect at least half of the S as SO2. I think I saw this in Kirk-Othmer somewhere, but that doesn't mean it works! And US2528103 bluntly says so, but doesn't even get Cummings' name right. Anyways, it is supposed to be an improvement. I've seen both of Cummings' pats on espacenet before, but (no surprise) I was just there and they say that they don't exist. Of course uspto.gov says that they do.
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[*] posted on 23-5-2004 at 23:14


oks....my contribution.....

1) high temperature control without some reasonable kit is a bitch.
2) the temperatures you are talking are getting close to the softening point of pyrex (600-700 C).....

All my attempts to decomp. FeSO4 CuSO4 as a teeny failed. I only used glass.

Now with much more real experience:
I think your best bet is a metal kit. Air and oxidation of the FeII i think is a limited worry, given that at the temperature you are talking the SO3 = SO2 n O2 equilibrium will estabilsh fairly quickly. Plus you have to get the O2 into the matrix of the solid (rates)
How to best do this is really dependant upon all sort of unknown rates, but my guess is you will be best off with a straight tube, Copper i think is good and easily available. Use an air carrier, modest airflow (fish tank pump). This will get the SO3 away more quickly from the hot surface where decomposition/equilibration will occur. SO3 is a pretty good gas at even modest temperatures, consdense on cold glassware. How quickly SO3 eats the tube is also a worry, but also another good reason for using a carrier gas. It will dilute the SO3, it will cool the SO3 and it will move the SO3 away from the metal surface more quickly.
-there are a lot of guesses, assumptions and opinion here!!!!!!!


I also FAILED MANY times to make SO3 from SO2 and air, using a V 2O5 catalyst, with only the most primative of supports. Temperature control was a real bitch (propane gas torch), as was getting SO2 dry enough.
I only once succeded at making SO3, and that was at school (everything else was done at home), using an SO2 cylinder and a platinumised wool catalyst. Using relatively low heat (bearly off sooty flame on a bunsen) over a period of maybe 1 hr, I made maybe 2-5g of SO3 as a crytaline solid that readily melted with bodyheat.


In summary, probably a bent Cu tube, FeSO4 in bottom, air carrirer, roast to red heat with gas torch, have faith, that even though you cant see whats happening, it will work!
alternatively, maybe a pretty determined piece of kit for the SO2 oxidation!
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[*] posted on 26-5-2004 at 19:10
Na2S2O8


A long time ago I generated SO3 from sodium persulphate, which I obtained from conrad.de, and which I coincidentially wanted to use for copper etching.
I just heated the persulphate in a testtube, and melted very quickly. Upon heating a bit further, VAST amounts of SO3 were produced, which quickly filled my room with a white fog. That lasted for quite a while, and surprisingly it wasn't particularly harmful. After a while, no more SO3 formed, and I was left with a clear molten salt... which is Na2S2O5? or possibly something different, with concommitant O2 evolution.

Realising then that this must be SO3, I fed this with some plastic tubing into H2O, but to my dismay it did not seemingly dissolve, and lowered the pH only a little. Nothing as impressive as bubbling Cl2 through NaOH.

Unfortunately, I never bubbled the SO3 through H2SO4 conc. - I wish I had done it, because it seems exceedingly easy to do, plus Na2S2O8 is serioulsy easy to get! Maybe someone else try it for me? :)




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smile.gif posted on 27-5-2004 at 04:34


You probably got Na2SO4, Na2SO5 isn't stable at high temperatures.
I think you probably produced SO2 with a little bit of SO3 to create a fog and O2. The SO3 present dissolved and lowered the pH.
Anyway, that's my theory.:)
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[*] posted on 27-5-2004 at 11:43


No way. Ever smelled SO2? You can detect it quite easily. And there was definitely no SO2 smell around (unless repeated ethanolic exposure screwed around my memories from more than a decade ago :D).
Just try it out, this is an experiment that can be literally done in 10 minutes.




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[*] posted on 27-5-2004 at 11:53


My comments: When I tried superheating FeSO4 in an iron pipe I left it unattended, so it overheated and got yellow-hot. The entire flat was filled with a thin white fog, but it had almost no smell -- it was just irritating to the lungs (smelt _slightly_ like SO2, but only a little, almost undetectable). I think it was a mix of SO2 and SO3, mostly SO3.

I think chemoleo is right.




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[*] posted on 2-8-2004 at 14:48


it is my understanding that SO3 in air readily forms H2SO4 especially when inhaled and is quite irritating and painful to your lungs.

if you had any SO3 at all you should have been coughing and in pain.
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[*] posted on 2-8-2004 at 18:17
NaHSO4 to Na2S2O7 by heat


I'm right now keeping 1/2 kg of NaHSO4 in my electrical furnace at 400 degrees C in an attempt to convert it to Na2S2O7. I'll tell you how it goes...

On a more bizarre note, I'm going to take up pottery in about 10 minutes... going to make a couple of clay retorts. I'll post pictures, regardless of the result... :)


[Edited on 2004-8-3 by axehandle]




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