Sciencemadness Discussion Board - SO2 Generator

SO2 Generator

Eclectic - 2-11-2007 at 06:55

Usually on this forum, we are obsessed with converting sulfur or sulfur dioxide to sulfuric acid and/or SO3. What if you want to go the other way? Anyone know of ways to use sulfuric acid as feedstock for a lab scale SO2 generator almost as convenient as an SO2 cylinder that can generate 1-40 lbs of SO2 in a reasonable amount of time without producing an excessive amount of heat or dilutant inert gasses?

Copper turnings in H2SO4 are supposed to work. Does anyone know if elemental sulfur reacts with hot H2SO4 the same way? Any other chemistry that might be useful? Catalysts?

(I've already thought about a sulfur/O2 burner. A 200 CF O2 welding tank could produce about 35 lbs of SO2 from burning sulfur, but getting rid of the heat could be a problem.)

[Edited on 11-2-2007 by Eclectic]

woelen - 2-11-2007 at 08:07

Don't you have access to sulfites or (meta)bisulfites. A lovely generator can be made by mixing H2SO4 with an equal volume of water, and then adding sulfite or (meta)bisulfite to this. Slight heating gives a smooth and not too fast generation of lots of SO2. You don't need strong heating, and that is a pleasant thing.

If you don't have sulfites, then indeed copper metal could be used, but the problem with this method is that a large excess of sulphuric acid is needed. Use of elemental sulphur buring in oxygen also can be used, but the problem with this is that it does not give pure SO2, but SO2 mixed with a lot of air or oxygen. A practical problem with that approach is to lead it through something. You don't have a pressurized source, because formation of SO2 from S and O2 does not lead to formation of more volume of gas from solid (each SO2 takes one O2). When copper/concentrated sulphuric acid, or sulfite/medium conc. sulphuric acid, is used, then you can have a closed flask, in which SO2 is formed.

Eclectic - 2-11-2007 at 08:48

I wanted to avoid the production of waste salts and excessive heat if possible. Does anyone know if hot H2SO4 reacts with elemental sulfur ?

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.

Maybe I should wait until winter and use it as a space heater.

[Edited on 11-2-2007 by Eclectic]

trilobite - 2-11-2007 at 09:14

Late Halfapint of the former Hive once mentioned heating sulfur with sulfuric acid as a SO2 generator that was well regulated. I suppose the reaction would be:

H2SO4 + S --> 2 SO2 + H2

I've never read about this way in any other source, but the person was most trustworthy in his descriptions. Of course hydrogen gas might pose a problem depending on application. One solution is to absorb the SO2 in aqueous alkali to make a sulfite and neutralize that later to release SO2, another is indeed liquification.

[Edited on 11/2/2007 by trilobite]

Eclectic - 2-11-2007 at 09:23

I was thinking in an excess of H2SO4, 2H2SO4 + S --> 2H2O + 3SO2 , but I don't know if it will actually work, or if there are nasty side reaction products to deal with. Maybe problems with thermal runaway?

trilobite - 2-11-2007 at 10:20

Hydrogen does sound a bit improbable for a reaction product, you may well be right. I don't think thermal runaway would be a problem since he emphasized that the rate of sulfur dioxide production was easily controlled by adjusting heating. To me this suggests an endothermic reaction with a large positive entropy term.

Xenoid - 2-11-2007 at 13:30

My abridged Mellors mentions heating any suitable reducing agent, copper turnings, sulphur or mercury, etc. with concentrated sulphuric acid. Copper is normally used;

"" The flask is about one third filled with copper turnings and sufficient concentrated sulphuric acid is added not quite to cover the copper. On heating, sulphur dioxide is evolved and may be passed through a wash bottle of concentrated sulphuric acid to dry it""

It also mentions SO2 can be formed by the action of "moderately concentrated"??? sulphuric acid on sodium sulphite or bisulphite.

Regards, Xenoid

chloric1 - 2-11-2007 at 14:54

steel shavings should work with hot H2SO4. If you use this or copper the salts can be harvested and purified for use or to sell on ebay.

Eclectic - 2-11-2007 at 14:58

Well yeah, but if sulfur works to cleanly reduce H2SO4, then you have NO byproducts to deal with! Green chemistry approach to generate acid rain produceing gas.

[Edited on 11-2-2007 by Eclectic]

garage chemist - 2-11-2007 at 23:36

Actually, water (not hydrogen!) is the byproduct of all methods utilizing reduction of H2SO4, and is troublesome because it dilutes the H2SO4 so that it may no longer reach reaction temperature. So if you make a SO2 generator with Cu or S and H2SO4, initially it will probably work well, but as soon as some SO2 has been generated, you will continue to boil the reaction mix and the SO2 production will get slower and slower despite the reagents being far from used up.

What you need is constant removal of water from the reaction mixture to keep the acid in the reaction mix concentrated. Running it in a distillation setup where the water condenses in the condenser and withdrawing the SO2 at the vacuum connection would be a good idea, for example.

If the H2SO4/S method actually works it would be a favorable method, simultaneous reduction of H2SO4 and oxidation of S to generate the SO2 and water as the only byproduct.
I could see some trouble though with evaporation of the sulfur, coating the inside of the distillation apparatus.
You better have some HNO3 at hand to distill in the apparatus (without water cooling, so that the hot HNO3 vapor reaches the whole inside surfaces) to oxidise the sulfur.
I used this cleaning method when my still was coated with sulfur on the inside from a synthesis of S2Cl2 and found it to be good and efficient. It emits vast amounts of red NOx vapors though.
Hot toluene can also work to dissolve the sulfur, and a wash with hot NaOH solution also dissolves sulfur (disproportionation to polysulfide and thiosulfate).

[Edited on 3-11-2007 by garage chemist]

Sauron - 3-11-2007 at 03:19

@Eclectic, see my PM. Hope this attachment helps.

[Edited on 3-11-2007 by Sauron]

[Edited on 3-11-2007 by Sauron]

Attachment: Pages from Mellor_ch57_1930.pdf (212kB)
This file has been downloaded 1427 times

Eclectic - 3-11-2007 at 05:34

Thanks for the info, guys. Has anyone actually tried this with H2SO4 and S? I haven't been able to find any references to this approach either. The practicality of the approach seems to be a matter of what temperature the reaction proceeds at a useful rate.
Below 100C, there would be water accumulation, but little sulfur fume coating the apparatus. At say 200C, I'd be reluctant to have a batch reaction where liquid reactants could convert to gas at an unknown rate.

I'd be much more comfortable injecting liquid S at a metered rate into the hot H2SO4 if that is the case, although at that temp water dilution would not be a problem

I guess it's time for a proof of concept small distillation setup scale trial.

2H2SO4 + S --> 2H2O + 3SO2 ???

[Edited on 11-3-2007 by Eclectic]

woelen - 3-11-2007 at 05:59

I tried the H2SO4 (96%) and S mix. It does not give SO2, not even when heated, such that the acid starts boiling (appr. 300 C). This simply does not work.

Eclectic - 3-11-2007 at 06:13

Thanks Woelen, that saved me quite a bit of time...

Maybe with a catalyst? A dash of CrO3 or V2O5? Hg?

[Edited on 11-3-2007 by Eclectic]

trilobite - 3-11-2007 at 08:29

Too bad it doesn't work, thanks for trying it out.

Eclectic - 3-11-2007 at 08:49

And yet, in the updated pdf provided by Sauron, there is a statement that it does, so maybe it WILL work with the appropriate catalyst? Test tube time...

woelen - 3-11-2007 at 14:07

I understood (from private communication on this subject) that sulphuric acid is not capable of producing SO2 from S, and this is perfectly confirmed by my experiments.
You need oleum for that, with free SO3, dissolved in H2SO4. With that, you can dissolve S (getting a blue solution), which on further heating decomposes, giving SO2. It is the SO3, which oxidizes the S, not the H2SO4. When all SO3 is used up, then the reaction stops.

Unfortunately I have no oleum (and if I had, I would not use it for making something simple as SO2). Using oleum for making SO2 is insane, oleum is very expensive and a very hard to obtain chemical (also difficult and dangerous to make yourself). If you can get oleum, then you certainly can get NaHSO3 or Na2SO3 and use that for making SO2.

Eclectic - 3-11-2007 at 14:55

Woelen, did you do any tests with small amounts of Cr,Fe,V,Hg?

You of all people I would think would willing to actually try a quick experiment. (I don't have any test tubes, V2O5, or Hg with me right now).

Who told you H2SO4 is not capable of producing SO2 from S, and why would you believe hearsay? (I trust your experiment with pure acid and sulfur, I'm just questioning "someone informed me")

"Sulphur dioxide is also formed when sulphuric acid is heated with sulphur: 2H2SO4 + S=3SO2+2H2O, as shown by J. B. A. Dumas, and C. F. Anthon. H. L. F. Melsens
found it to be an advantage to mix the acid and sulphur with pumice-stone.
E. Hart recommended warming sulphur with fuming sulphuric acid, containing
30 per cent. SO3, when the blue soln. so formed is heated, sulphur dioxide mixed
with some trioxide is given off. The evolution of the dioxide ceases when all the
trioxide has been acted upon, and the sulphur melts. J. Knezaurek obtained
sulphur dioxide by heating sulphuric acid with charcoal: 2H2SO4+C=2H2O
+2SO2+CO2, for purposes where the admixed carbon monoxide and dioxide
will do no harm. W. L. Scott recommended using acid of sp. gr. 1-824 containing
74 per cent. SO3. If a more cone, acid is used, a portion is reduced to sulphur;
and if a more dil. acid is employed, some hydrogen sulphide is formed. The
washing liquid should be mixed with lead sulphate or coarsely powdered charcoal.
One of the commonest methods of preparing small quantities of sulphur dioxide
is to heat cone, sulphuric acid with copper: Cu+2H2SO4=SO2+2H2O+CuSO4
—vide 3. 21, 6. According to G. J. Warner, sulphuric acid is reduced to sulphur
dioxide at 160°, and S. Cooke showed that the reaction is accelerated by the
presence of platinum."

This is from Sauron's posted pdf above, page 2.

[Edited on 11-3-2007 by Eclectic]

woelen - 3-11-2007 at 15:58

What I said in my previous post is exactly the same as what you now tell me in this quote. SO2 is formed from oleum (fuming sulphuric acid, 30% SO3, according to your quote) and as soon as all SO3 is used up, no SO2 is formed anymore and the sulphur only melts.

So, again, pure H2SO4 is not capable of producing SO2 from sulphur, you need oleum (=fuming sulphuric acid with SO3 in it) for that. Your quote is not about H2SO4 alone, but about oleum.

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

Maybe the addition of metal salts changes things, but for one, I already know that it will not work:

CrO3 will not work. I did an experiment with this in sulphuric acid (CrO3 from K2Cr2O7), but when this is heated strongly, then it decomposes, giving oxygen and an insoluble chromium(III) compound (which is a riddle for me on its own): http://woelen.homescience.net/science/chem/exps/raw_material...

I can try with V2O5, but I also have doubts on that. I'll come back on that one after I did the experiment, but right now, I expect that either it does not react at all, or the sulphur reduces this to a vanadium(IV) species.

[EDIT: changed URL of link so that it works again]

[Edited on 9-5-12 by woelen]

Antwain - 3-11-2007 at 15:58

Quote:

E. Hart recommended warming sulphur with fuming sulphuric acid, containing
30 per cent. SO3, when the blue soln. so formed is heated, sulphur dioxide mixed

um, yeah. that would be oleum

Eclectic - 3-11-2007 at 16:12

"Sulphur dioxide is also formed when sulphuric acid is heated with sulphur: 2H2SO4 + S=3SO2+2H2O, as shown by J. B. A. Dumas, and C. F. Anthon. H. L. F. Melsens
found it to be an advantage to mix the acid and sulphur with pumice-stone."

It's a litany of different methods purported to produce SO2 from the 1850's. Just because oleum works, doesn't mean other method's DON'T. What's up with the latching on to one sentence and ignoring everything else?!?

Woelen, I'm fairly sure that acidic CrO3 is going to oxidize sulfur, probably all the way to SO3. The question is whether or not the resulting Cr2(SO4)3 will in turn be oxidized by hot H2SO4 to a higher oxidation state, producing SO2 and becoming available to attack more elemental sulfur, hence acting as a catalyst. In your experiments with Cr(III) in hot H2SO4, were you in a position to smell if any SO2 was evolved? You didn't indicate that there was bubbling, so probably not. But maybe other transition elements would catalyze the reaction?

If Dumas and Co. got sulfur to react with H2SO4, and you didn't, it's likely that their H2SO4 had a contaminant that was acting as a catalyst, yes? Dumas was a contemporary of Faraday in the 1800's., and I doubt they were using a reagent grade acid.

A trace of HNO3, Br, or Cl might facilitate the reaction. Eye of newt, hair of dog...

[Edited on 11-3-2007 by Eclectic]

S.C. Wack - 3-11-2007 at 17:48

No. You're not going to like what I'm going to say, BTW.

My first thought when I saw this thread was: "What is so hard about mixing S and H2SO4 in a test tube and finding out for yourself?"

Then I thought about all the old literature available to anyone with the slightest inclination to read it, and I thought some other things that are probably best left unsaid.

I also thought that it would make sense that the recommended methods for making it are the recommended methods for good reason.

Now it's just not Mellor that gives that quote, it's all over.

I have no idea why it does not work for woelen. Like I said this is easy to test and it sure is. Quicker than writing this post. Over here with colorless 96% technical acid and sulfur powder, it works fine. The acid soon turns blackish the evolution of gas, easily noted by bubbles at the exit tube just under the surface of water in a test tube. The was no boiling or anything in the flask (it was more convenient for me to use a jointed flask with gas tube and a mantle than dig though my stuff for stoppers and tubing) No analysis of the gas was made, but it is plentiful enough and it contains an unpleasant amount of SO2.

I don't see why Melsens would talk of 12 L cast iron flask scale if it didn't work.

Eclectic - 3-11-2007 at 18:07

Am I bovvered? Do I look bovvered? (Obscure reference to the Catherine Tate Show)

No test tubes handy today, nor my 50 lb bag of sulfur prills.

I'm all for experimentation, and figured I'd ask and try and stimulate some. Cast Iron Flask may point to iron as a catalyst?

woelen - 4-11-2007 at 04:09

Quote:

I have no idea why it does not work for woelen. Like I said this is easy to test and it sure is. Quicker than writing this post. Over here with colorless 96% technical acid and sulfur powder, it works fine. The acid soon turns blackish the evolution of gas, easily noted by bubbles at the exit tube just under the surface of water in a test tube. The was no boiling or anything in the flask (it was more convenient for me to use a jointed flask with gas tube and a mantle than dig though my stuff for stoppers and tubing) No analysis of the gas was made, but it is plentiful enough and it contains an unpleasant amount of SO2.

What flask are you using? It really does not work for me. I only see the sulphur melting and nothing else. At a certain point the liquid becomes so hot, that the acid starts emitting white fumes, at that point I stopped. The acid remains colorless, the sulphur forms droplets in the acid.

First, the droplets are viscous and yellow:

At much higher temperature, the droplets have a deep red color:

I also cannot understand why the acid would become black-ish, as S.C. Wack describes. My acid does not change color, it may become slightly darker, but that's all. Even after 10 minutes of heating bloody hot, things are not changing.

[EDIT: Changed URL's of picture so that they work again]

[Edited on 9-5-12 by woelen]

not_important - 4-11-2007 at 05:36

I suspect that the reaction of S and H2SO4 may be catalysed by transition metals or NOx, both of which would not be unexpected in 19th century H2SO4. A quick test of this would be to add the smallest crystal of KNO3 to the non-reacting S+H2SO4, or a trace of Fe, V, or Mn. I'd try it myself, but it's late enough that the neighbors might complain about infernal visitations.

garage chemist - 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.

chloric1 - 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

What is not green is dumping.

Eclectic - 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.

trilobite - 4-11-2007 at 15:40

I knew late Halfapint wasn't lying

. 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]

S.C. Wack - 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)
This file has been downloaded 734 times

Eclectic - 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.

What temp do you think this reaction is useful at, H2SO4 reflux temps, or maybe lower?

woelen - 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

. 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.

woelen - 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.

Eclectic - 5-11-2007 at 11:14

As long as you have the setup, please try Fe, Cu, Hg, Br and NOx.

woelen - 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

). 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.

Eclectic - 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]

Sauron - 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.

not_important - 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.

woelen - 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.

Sauron - 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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woelen - 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.

Sauron - 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.

woelen - 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...

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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.

Sauron - 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.

Eclectic - 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]

Sauron - 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]

nitroglycol - 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]

Eclectic - 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]

Sauron - 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.

dann2 - 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?
Dann2

weiming1998 - 12-4-2012 at 17:31

Quote: Originally posted by dann2

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?
Dann2

I don't know about impurities, but they wouldn't really matter that much. Most impurities probably won't affect the SO2 made, but that depends on what acids you use. If there is chlorides in your winemaking tablets, for example, and you use sulfuric acid, the resulting evolution of HCl gas will ruin your SO2.

Sulfuric acid isn't actually a good candidate acid for generating SO2, as other soluble gases generated might contaminate the SO2, and also the fact that SO2 dissolves in sulfuric acid more than water. So your best acid choice would be HCl solution, as that prevents the formation of hydrogen halides.

But almost all acids work (well, apart from a few exceptions that are extremely weak acids/oxidizing acids like HNO3, which will oxidize the SO2 produced to sulfuric acid, and HClO/H2CO3, which almost certainly doesn't work) by this equation: Na2S2O5(s/aq)+2H+(aq?)==>2SO2(g)+H2O+2Na+(aq). The SO2 exits the solution as a gas, making this reaction favourable, so even weak acids (like acetic acid) works. But not a dilute solution though (if you used vinegar, the yield will not be great compared to if you used 25% acetic acid.)

dann2 - 13-4-2012 at 03:46

Thanks for reply.
I might get around to trying it some time as I need to clear away all Chlorate from Perchlorate that I have using SO2.

woelen - 8-5-2012 at 23:49

After reading over this whole thread I now think that the formation of SO2 from H2SO4 and S maybe due to organic impurities in the S. This also could account for the black coloration of the acid, reported by some people. I did not know of that kind of impurities back in 2007, but recently I have seen some posts about sulphur from gardening suppliers containing organic impurities, leading to charring when the sulphur is heated. I also read that some sources of sulphur may contain small amounts of CaCO3, which also may cause bubbling (CO2).

I still am convinced that very pure sulphur and pure sulphuric acid only react with difficulty making SO2.

elementcollector1 - 30-1-2013 at 10:12

I saw you mentioned on the first page that your CrO3 decomposed to oxygen and a green Cr(III) compound. That was chromium (III) oxide. How did you make your CrO3 from K2Cr2O7 and H2SO4, anyway?

Also, for production of SO2 from sulfuric acid, what about electrolysis? This could even work for sulfates.

woelen - 31-1-2013 at 00:57

I did not isolate CrO3. This compound simply is formed when chromates or dichromates are mixed with conc. H2SO4.

E.g. K2Cr2O7 + 2H2SO4 --> 2KHSO4 + H2O + 2CrO3

This reaction also works with very concentrated solutions. E.g. if you prepare a saturated solution of Na2Cr2O7 and pour in conc. H2SO4, then a dark red slurry of CrO3 is formed. This material, however, is very hard to isolate.

For my experiment there was no need to isolate the CrO3, it simply was in solution and reacted with the sulphur.

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Electrolysis of solutions of sulfates or aqueous sulphuric acid does not lead to formation of SO2. At the anode you get oxidation of the anode material, or formation of oxygen if the anode is inert and at the cathode you get hydrogen gas.

elementcollector1 - 31-1-2013 at 15:43

What makes it hard to isolate? Does filtration or decanting the excess liquid not work for whatever reason?

UnintentionalChaos - 31-1-2013 at 16:07

Quote: Originally posted by elementcollector1
What makes it hard to isolate? Does filtration or decanting the excess liquid not work for whatever reason?

Hygroscopic, corrosive, highly water soluble, sets a number of organic solvents on fire on contact. You need to remove the fairly concentrated sulfuric acid from the product somehow.