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Author: Subject: SO2 Generator
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[*] posted on 4-11-2007 at 06:24


I just tried it myself. Sulfur suspended in conc H2SO4, upon heating until the sulfuric acid nearly boiled there was a strong smell of SO2 at the opening of the test tube!
On the surface of the molten nearly black sulfur drop, one could see gas bubbles growing and rising to the surface of the H2SO4.
So the reaction DOES work, but needs powerful heat.




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[*] posted on 4-11-2007 at 06:25


Well i still ike steel shavings because A) They are cheap and often free. B) the ferric sulfate byproduct is usefull. This is appeals to me since I have 2 or 3 pounds of copper sulfate already.

@Electic-Having byproducts is green as long as you use them:D What is not green is dumping.:D




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[*] posted on 4-11-2007 at 06:42


A guy in a bar walks up to a woman and ask if she'll have sex with him for a million dollars. Sure, she says. Here's a $50 the guy says. "What kind of woman do you think I am?", asks the outraged woman. "We've already established that, now we're just haggling over the price."

The proper catalyst may reduce the thermal "price" of this reaction.
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[*] posted on 4-11-2007 at 15:40


I knew late Halfapint wasn't lying :D. Good stirring would surely be a great idea with this two-phase reaction. I bet that's why using pumice is suggested, for increasing the surface area between reactants. Maybe diatomaceous earth?

[Edited on 11/4/2007 by trilobite]
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[*] posted on 4-11-2007 at 19:51


*raises right hand* I swear that I've got my shit together.

*does blank test with same sulfuric acid, same (cleaned and dried) flask, same mantle, same voltage setting*
There is nothing wrong with my acid. It does not darken or release SO2 when heated to the point that it was both yesterday and today.

*washes and dries flask*
*takes first pictures with new digital camera*
*aggravated that the very bright balance display is overwhelmed by the flash at its lowest intensity*

Well, this is about what I did before. Probably a little more of each this time, and a more prolonged temperature ramp-up. Here we have around 1 g S + 5 ml H2SO4. Could not save smell to pdf, so illustrated as best I could with a potassium salt probably quite familiar on sight to woelen. The powdered sulfur is from a different source this time, but the result was the same.

One picture illustrates the black color of the mixture on heating to a medium temperature, hot enough to sublime the sulfur without any conversion to SO2. This did not occur yesterday because of a much faster temperature increase - fiddling with the camera really slowed me down today. Another shows reduction caused by SO2 at the higher temperature.

On cooling, the acid was light-but-not-faintly brown.

PS woelen - you mention heating till the acid fumed - in a container open to the air, this happens at a much lower temperature that you may think. Nowhere near the boiling point. I think it has something to do with the extreme affinity for water in the air, it reaches out to get it, as crazy and non-chemistspeak as that sounds, all of my observations on hot sulfuric acid show this to be true.

I've also noted the SO2 evolution and blackening on heating of drain cleaner H2SO4 (Rooto, etc.,) and thus my comments in the relevant thread quite some time ago, which were criticized without merit IMHO. I've never been sure what it is that isn't listed on the MSDS, maybe this is it. I assumed that it was something with carbon due to the blackening.

[Edited on 4-11-2007 by S.C. Wack]

Attachment: so2.pdf (411kB)
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[*] posted on 4-11-2007 at 20:39


Did you notice anything "violent and uncontrollable" about this as mentioned in Melsens' footnote? Maybe he was dropping H2SO4 into molten sulfur at over 300C. :o

What temp do you think this reaction is useful at, H2SO4 reflux temps, or maybe lower?
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[*] posted on 5-11-2007 at 06:02


Quote:
Originally posted by garage chemist
I just tried it myself. Sulfur suspended in conc H2SO4, upon heating until the sulfuric acid nearly boiled there was a strong smell of SO2 at the opening of the test tube!
On the surface of the molten nearly black sulfur drop, one could see gas bubbles growing and rising to the surface of the H2SO4.
So the reaction DOES work, but needs powerful heat.
Time to take out the test tube for me again :P. I hope to find some time for repeating the experiment this evening. Right now, I do think that I did not heat sufficiently strong, and that indeed the fuming I observed was at a lower temperature. For me, this issue must be settled, once and for all ;). I'll use a Duran/Jena-glass tube and heat until the acid starts boiling and see what happens. I'll also try adding a pinch of metal salts.



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[*] posted on 5-11-2007 at 11:00


I did the test, and still, I think that the reaction hardly works. Yes, it does work, but it is horribly slow and extremely energy-intensive.

I took 2.0 ml of 96% acid, quality denoted as "rein", which means something like "pure". I think it is general decent lab grade, but not reagent grade. It is, however, perfectly colorless.

I took 100 mg of flowers of sulphur. This sulphur also is a lab-grade chemical of good purity (sulphur - LR grade, BDH).

In this setup there is a large excess of sulphuric acid.

Then I started heating. The sulphur melts, it first becomes yellow, then red, finally it becomes very dark red, almost black. The sulphur also collects into a single blob, while initially many smaller droplets were in the liquid. At a certain point, I had a still perfectly colorless liquid, with a single large black almost spherical blob in it. And then, suddenly, the blob started bumping, gas bubbles were produced at it. This reaction, however, is not particularly fast. I continued heating, until the sulphuric acid started boiling and INTENSE thick white smoke was produced. At this point, when the heat-source is taken away, you can see gas bubbles being formed at the almost black blob of sulphur. This situation was maintained for almost 10 minutes.

After these 10 minutes, the blob of sulphur hardly had become smaller. There was smell of SO2, but not very intense. Also after 10 minutes, the acid still was perfectly clear and colorless.

Conclusion: The reaction works, and may be interesting from a theoretical point of view, but for practical production of SO2 from S, it is totally useless. You need so much heat, and the conditions of the reaction are so harsh, that I absolutely think it is horrible.

===========================================================

I myself hardly could believe that there is a metal-catalyst for this reaction, but oke, I already had the tube with molten S and hot H2SO4 in front of me, so I added some V2O5 (a small pinch, the tip of a screw-driver, thinly covered) after this 10 minutes of heating. The V2O5 quickly dissolves, giving a red solution, but within 1 minute of little shaking, the color changed to bright blue. After yet another minute, the liquid become turbid and the color changed from blue to yellow. Finally, a bright ochre/yellow precipitate is formed, and the blob of sulphur still is almost black. When the V2O5 changes color to blue, then many small bubbles of SO2 are formed, but this is only for a short time. When it all is converted, the production of SO2 is as slow as before.

I also did a counter experiment, with only V2O5 added to H2SO4. This results in formation of a deep red solution, which even on boiling and intensely thick smoking/fuming remains red. No change to blue.

Conclusion: Vanadium does not catalyse the reaction. Vanadium in the +5 oxidation state is reduced by the S to vanadium in the +4 oxidation state (bright blue color is due to vanadyl, VO(2+)). The further step, the formation of the ochre/yellow precipitate, is not clear at all to me. Is this some sulfide of vanadium? I could not find info on such a compound. So, we have a new, interesting riddle, but absolutely no catalyst.
The sulphur is indeed the reductor, because without sulphur, vanadium is not reduced, the boiling hot solution remains deep red.

I did not bother trying other metals. I cannot see any catalytic mechanism.

If, however, some of you think it is appropriate to try another metal salt, then I could try (provided I have a salt of the metal), but I personally do not believe anything of this.




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[*] posted on 5-11-2007 at 11:14


As long as you have the setup, please try Fe, Cu, Hg, Br and NOx.
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[*] posted on 5-11-2007 at 13:31


Tried the Cu, Br and Nox, each with a fresh amount of H2SO4+S.

Adding a pinch of CuSO4.5H2O results in a nice pale blue solution, but besides that I can't see a significant difference.

Adding a pinch of KBr. This gives more interesting results. First, production of HBr (white fumes) and Br2. Lateron, there is a very foul smell, strongly sulphurous and really bad. This smell, however, disappears again, and furtheron, no interesting reaction seems to occur anymore. Small part of the bromine/bromide is reacting with the sulphur, and that's all. With the bromide, the liquid becomes yellow, and it remains yellow, even when the sulphuric acid starts boiling and intense white fumes are produced.

Finally, adding NaNO3 gives brown NO2 above the liquid. Also, quite some gas is produced at the surface of the sulphur, the reaction definitely is faster than without NaNO3. But I have the strong impression that this really is due to oxidation of sulphur by nitrate/NOx and not by some catalytic reaction. With the NaNO3, the liquid remains colorless.

After these three experiments, I quit (it's 22.30 now, and I now want a nice red drink with some ethanol in it, which certainly works catalytically :D). I might try the iron tomorrow, if I find some time for it. The Hg I don't try, I don't want to poison myself with volatile mercury compounds or mercury metal vapor.

Altogether, things do not look promising. I think we simply should conclude that SO2 production from H2SO4 and S is not interesting, it is horribly slow and requires a LOT of heating.




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[*] posted on 5-11-2007 at 14:51


OK. Thanks for your efforts. I'm not able to run the tests myself right now or I would have. The NOx might be the thing that's different between the H2SO4 we have today and what was reported back in 1858. That and iron contamination.

[Edited on 11-5-2007 by Eclectic]
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[*] posted on 5-11-2007 at 18:06


The old literature as cited in Mellor reports that sulfur when dissolved in oleum of SO3 content 30% or higher, turns the solution blue. I have now identified the blue species as S2O3, formed by the reaction

S + SO3 -> S2O3

See Brauer p.380. Unstable crystalline solid blue-green. Soluble in oleum, color blue or brown depending on SO3 content. Can be stored for a few hours at below 15 C.

When heated decomposes into SO2, SO3 and S.

This is of course useless preperatively but does lay to rest one small mystery.

See Vogel and Partington, also J.Chem.Soc. 127 1514 (1925)

That is I.Vogel not A.Vogel.




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[*] posted on 5-11-2007 at 20:59


Quote:
Originally posted by woelen
...

Conclusion: Vanadium does not catalyse the reaction. Vanadium in the +5 oxidation state is reduced by the S to vanadium in the +4 oxidation state (bright blue color is due to vanadyl, VO(2+)). The further step, the formation of the ochre/yellow precipitate, is not clear at all to me. Is this some sulfide of vanadium? I could not find info on such a compound. So, we have a new, interesting riddle, but absolutely no catalyst.
The sulphur is indeed the reductor, because without sulphur, vanadium is not reduced, the boiling hot solution remains deep red.


I've read this,, thought it was one of the early Inorganic Synthesis but wasn't in the ones I have so can't give a reference.

The vanadium sulfides are dark, black powders or metallic black or dark grey crystals.

What I believe you made is anhydrous V(III) sulfate, which is described as a yellow micro-crystalline powder. The preparation I remember used H2SO4, S, and V2O5 heated together for some time. There was a suggested cooling process, which results in the remaining sulfur forming solid globules separated from the V2(SO4)3. The hydrated forms of the sulfate have the more common greenish colouration and are somewhat water soluble; they have some resemblance to the Cr(III) sulfate-hydrates.

------------------------------------------------

This has been a well done series of experiments, with people running checks by just heating one reagent, or less than the full set of reactants, to confirm that the combination was responsible for what was observed.
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[*] posted on 6-11-2007 at 00:25


Quote:
Originally posted by Sauron
The old literature as cited in Mellor reports that sulfur when dissolved in oleum of SO3 content 30% or higher, turns the solution blue. I have now identified the blue species as S2O3, formed by the reaction

S + SO3 -> S2O3

See Brauer p.380. Unstable crystalline solid blue-green. Soluble in oleum, color blue or brown depending on SO3 content. Can be stored for a few hours at below 15 C.

When heated decomposes into SO2, SO3 and S.

I also understood that a blue compound is formed in oleum, but this is not S2O3. Indeed, older literature describes it as S2O3, but in the very recent past, this has been shown to be false.

A cationic species instead is formed, something like S8(2+) or S4(2+), with a very peculiar structure. I did experiments with this kind of cations, which I could prepare for Te and Se, but not for S. For Te, a beautiful red/pink ion is formed, and for Se, an olive green ion is formed. Te already forms such an ion in only warm H2SO4. Se requires the acid to be much hotter, and S requires the presence of SO3.

For more information on this, read "Chemistry of the Elements" from Earnshaw and Greenwood.




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[*] posted on 6-11-2007 at 01:31


Brauer is not 19th century, and neither is the J.Chem.Soc. reference I cited (which is in Brauer.)

S2O3 is formed from neat SO3 and S according to the eqn I gave.

As S2O3 is blue-green and decomposes on heating to form SO2 which will evolve, SO3 which will likely return to soln and S which will redisolve, and the medium is the H2SO4-SO3 system, I would not lightly dismiss the proposition that S2O3 is responsible for the blue color.

[Edited on 6-11-2007 by Sauron]

Attachment: ct9252701514.pdf (786kB)
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[*] posted on 6-11-2007 at 03:18


Interesting, literature does not agree on this. According to the book "Chemistry of the Elements", S2O3 does not exist at all. It is mentioned in the book, but only as a false entity, which was believed to exist. Who is right? I don't have the resources to check (actually, I cannot even make that blue stuff). I think this is quite common in chemistry, even simple mixes and reagents can lead to strange compounds, whose structure certainly is not yet established. This also makes the chemistry of simple inorganic compounds so interesting for me.



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[*] posted on 6-11-2007 at 05:51


Well, I have 65% oleum, and sulfur, so, as soon as my fume hood is installed, I will try the procedure described in details inthe above attached J.Chem.Soc. paper and in Brauer and even though the stuff is unstable it will be stable enough to photograph and weigh. It was characterized, and its molecular weight determined with care. Not some fleeting ionic species but a solid substantial compound.

The color in solution in the 19thcentury lit. varied from blue to brown depending on the strength of the oleum and according to later authors those variations were probably due to the great variables in purity of the reagents. The article I attached above describes the pains that were taken to purify the sulfur and the SO3, the apparatus used, the reactions of the product of the interaction of those two and the removal of excess SO3 by decantation and reduced pressure. All methodologisies meticulously recorded.




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[*] posted on 6-11-2007 at 14:11


Sauron, I tried making the blue compound with H2SO4, mixed with P4O10 and then adding S. No positive result. Maybe I used insufficient P4O10, but I think that the SO3 simply is not concentrated enough in such a situation.
I REALLY would like to have oleum, that is one of the few chemicals, which I feel I am badly missing, but till now, I have no affordable source for this chemical. It is sooooo expensive...

----------------------------------------------------------------------------

Finally, I did the Fe experiment with H2SO4 and S. I took some FeCl3.6H2O and added this to a large excess amount of 96% H2SO4. This results in formation of HCl (as expected) and on heating, much more HCl is formed. The liquid becomes off-white and turbid. Anhydrous Fe2(SO4)3 is formed. This is almost white. To this mix, I added some sulphur. The sulphur simply melts, and when it has become really hot, the familiar almost black drop of sulphur is formed, with slow formation of SO2. The iron does not add any catalytic effect.

Then I added some NaNO3 to the hot mix of Fe2(SO4)3, H2SO4 and S. Immediately brown NO2 is formed, and also there is more formation of gas at the surface of the blob of sulphur. This, however, only lasts for a short time. I think the nitrate quickly is used up in a redox reaction between the nitrate and the sulphur.

Altogether, I am now really convinced that there is no nice metallic catalyst, nor halogen, nor NoX as catalyst, which makes production of SO2 from S and H2SO4 feasible/economical. This is something, which I already was thinking in advance, but it is good to let experimental outcome decide. What I did learn is that metal salts can form really funny anhydrous sulfates, which are remarkably insoluble. The ferric sulfate (plus acid) also is insoluble in water at first glance. Only after many hours of standing, it dissolved, giving a near colorless solution.




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[*] posted on 6-11-2007 at 16:10


I have no confidence in the ability of P2O5 to wrest water away from H2SO4 and liberate SO3. At best, you might end up with a mix of anhydrous H2SO4, H3PO4 and some excess P2O5.

I had to pay an arm and a leg for this 2 liters (4 Kg) 65% oleum from Merck, local agent raped me badly. About $1500, so $750 a liter or $375 a Kg. Ouch. Hence my interest in making oleum or SO3 from iron sulfates.




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[*] posted on 6-11-2007 at 16:15


P2O5 dehydration of H2SO4 works (I know from personal experience), but only near the boiling point. Maybe there is appreciable dissociation into H2O and SO3 at those temps?

[Edited on 11-6-2007 by Eclectic]
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[*] posted on 6-11-2007 at 18:19


Even if so, useless in contest of a compound that falls apart at anything over 15 C.

Anyway, are you telling us that conc H2SO4 with sufficient P2O5 at the 290 C bp of H2SO4, will allow SO3 to distill out, say into a receiver of conc H2SO4 w/cooling, to produce oleum?

That would be nice if it were true. There'd be no need for pyrolysis of iron sulfates, etc. By employing an excess of P2O5 to theoretically remove all the water in conc H2SO4 including 1 mol H2O from every mol H2SO4 would require about 20+ mols P2O5 per L conc H2SO4 and produce about 19 mols phosphoric acid. Heat this to the bp of H2SO4 and you have a mix of pyrophosphoric acid and excess P2O5 - and if you are right, 18 mols of SO3. 1440 g SO3? Like I said, this would be nice, but...

At a slightly higher temperature (300 C) pyrophosphoric acid changes to metaphosphoric acid.

Problems: you can't do this in a glass vessel. The phosphoric acid when hot will eat the glass.

We are talking about 1840 g H2SO4 + >2840 g P2O5 giving up 1440 g SO3 and leaving behind >3240 g of pyro- or metaphosphoric acid and a little P2O5 excess that ensured complete removal of water.

So, what is your experience? Did you actually get any free SO3 to come out of this



[Edited on 7-11-2007 by Sauron]




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[*] posted on 6-11-2007 at 18:32


Quote:
Originally posted by Eclectic
I wanted to avoid the production of waste salts and excessive heat if possible.


Thing is, with HCl you could buy at the hardware store and the sodium metabisulphite you could buy at a winemaker's supply store, your waste salt would be the kid you could sprinkle on you supper afterwards. Don't know about the heat involved, though (I live in an apartment, so doing experiments with stuff that smells bad is a major no-no). These experiments with oleum, steel and copper, chromates, P2O5, etc. are interesting in their own right, but if you just want to make SO2, metabisulphite and HCl sounds like the way to go. I suppose one advantage to using oleum is that by the time you got any SO2, you wouldn't be worrying about nasty smells anyway.

Quote:

A pure O2/sulfur burner is probably the way to go if not, but the cooling would be a bitch in order to be able to liquify the resulting gas.[Edited on 11-2-2007 by Eclectic]

As a kid I tried burning sulphur outside at around maybe -20 C (real temperature, not wind chill) in the hope of liquifying the product, but I didn't have the proper setup and I soon got cold and gave up.

[Edited on 6/11/2007 by nitroglycol]




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[*] posted on 6-11-2007 at 20:06


@Sauron, Yes as I recall I got about 50 ml SO3 from 98% H2SO4 and an old bottle of P2O5 that had partially hydrated and was otherwize useless. Air cooled short path distillation setup heated with a heat gun really hot. I was worried about the glass cracking.

I think it was a 250ml flask, but it was a long time ago. Excess H2SO4, and no attempt to get high yield.

[Edited on 11-6-2007 by Eclectic]
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[*] posted on 6-11-2007 at 20:29


Very excellent!

I'll give this a try on the basis of half the P2O5 that would take all the water out. Hopefully the remaining sulfuric acid will keep the phosphoric from eating the glass.




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[*] posted on 12-4-2012 at 14:54



What impurities would Camden tablets (Sodium Metabisulphite for wine making) contain.
Would it effect SO2 making using the tablets + Sulphuric acid
Could other acids be used instead of Sulphuric?
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