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Author: Subject: Synthesis of Sulfur Trioxide and Oleum: The persulfate method
garage chemist
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Synthesis of Sulfur Trioxide and Oleum: The persulfate method

This is a new synthetic pathway to sulfur trioxide, with completely OTC chemicals and without the need for fancy apparatus, like catalyst tubes and such.
The idea comes from the german wikipedia page about sulfur trioxide. The pictures and process description below are from me.

The only apparatus required is a ground- glass distillation setup, the same as you would use for the distillation of nitric acid.

Warning!
Sulfur trioxide is extremely corrosive. It instantly carbonizes any organic matter it touches, including skin. A drop placed on wood instantly makes a black spot, under strong fizzing and fuming. The same happens when a drop falls on the skin.

Sulfur trioxide also reacts with water with explosive violence.
When a drop of water falls into a flask containing SO3, the flask usually shatters because of the violent reaction and localized heating.

SO3 fumes in air very strongly.
The synthesis must be carried out outside or under a fume hood.

When sodium persulfate is heated, it loses oxygen to form sodium pyrosulfate.
2 Na2S2O8 ---> 2 Na2S2O7 + O2

The evolved oxygen contains a small amount of ozone, which is identifiable by its smell.

Sodium pyrosulfate further decomposes into sodium sulfate and sulfur trioxide, but the reaction only takes place when sulfuric acid is present.

Na2S2O7 ---(H2SO4)---> Na2SO4 + SO3

The sulfuric acid plays the role of a catalyst, it does not take part in the overall reaction, but without its presence, no SO3 is formed.

Experimental:

17,5g dry sodium persulfate were put into a 100ml round- bottom flask, and 2 ml concentrated sulfuric acid (of the highest obtainable strength, any water is very detrimental to the yield) were added.

The flask was attached to the distillation setup, and heated by means of a bunsen burner.
No water is run through the condenser, as the SO3 would crystallize in the condenser. If necessary, the receiver is immersed in ice water for cooling (not necessary with this small batch).
This was the setup:

On heating, the mixture starts to rapidly generate oxygen, with some ozone as byproduct.
When the oxygen evolution ceases, the mixture is heated stronger. It quickly melts into a clear liquid, and dense white fumes start to fill the setup.

Soon the first drops of liquid SO3 will fall into the receiver.
The thermometer registers a steam temperature between 45 and 50°C during the process. If it becomes significantly higher than that (e.g. over 60°C) heating should be reduced a bit, since some H2SO4 could start to come over.

Residues of grease in the apparature will be carbonized by the SO3, this is normal and nothing to worry about.
The SO3 will also usually be colored brown due to this.

The flask will reach a very high temperature during this step (around 300°C), this is necessary for the sodium pyrosulfate to decompose.

Care must be taken to avoid the crystallizing SO3 plugging the condenser outlet. This can be very dangerous. The solid SO3 can be molten by means of a hot air gun if it is found to crystallize at the condenser outlet.

When no more SO3 is coming over despite the reaction mixture still boiling, the reaction is over.

The SO3 will most likely have partially or completely solidified in the receiver.
Mine looked like this:

This shows the fuming of the SO3 (sorry, it is hard to capture it in a picture, and the smoke is constantly drawn away by the fume hood):

My yield was only 1,4g SO3 from the 17,5g persulfate, but this was mostly due to the very small batch size.
A yield of 100g SO3 from 200g persulfate is claimed from another person.
The ration of persulfate to H2SO4 can also be varied and its effects on yield studied.

The distillation flask can be emptied, recharged with fresh sodium persulfate and H2SO4 and distilled again to increase the amount of SO3 obtained.
The batches should not be larger than 40g persulfate at a time, because larger batches will be heated unevenly.

The combined portions of SO3 can be redistilled for better purity.
Pure SO3 boils at 44°C and solidifies at 16,8°C.

The still is left in the fume hood until no more fuming is observed (all the SO3 is turned into H2SO4 by atmospheric moisture), then it can be safely washed out with water.

SO3 will quickly polymerize upon storage, evidenced by transformation into a white amorphous mass (the polymerization starts within a few minutes to hours after preparation, depending on purity), but it depolymerizes again at ca. 66°C. Therefore, simple distillation gives normal SO3 again.

The SO3 can be dissolved in conc. H2SO4 to give oleum of any desired strength, but be warned, this process is exothermic, and with the very low boiling point of SO3 it is very likely to start boiling. The liquid SO3 should be slowly dripped into stirred, ice- cooled H2SO4 to avoid this.
Polymerized SO3 also dissolves in H2SO4, but slower and with less heat evolution.

Sulfur trioxide and Oleum are exceptionally versatile reagents in the laboratory. Some of the most important uses shall be mentioned here:

With methanol or dimethyl ether, the powerful methylating agent dimethyl sulfate is formed. It is isolated by fractional distillation in vacuum.
Diethyl sulfate is prepared analogously, by distilling SO3 into dry diethyl ether, distilling away the ether and distilling the diethyl sulfate over Na2SO4 in vacuum.

By distilling SO3 into SCl2, the important chlorinating agent thionyl chloride is produced. This can be used for the production of acetyl chloride from acetic acid, and subsequently acetic anhydride.

By leading a stream of dry HCl gas through oleum until gas uptake stops, chlorosulfonic acid can be isolated by distillation.
Chlorosulfonic acid is used for chlorosulfonations, by reaction with benzene, benzenesulfochloride (C6H5-SO2Cl) is obtained.
Chlorosulfonic acid can also be used for an alternative preparation of dimethyl sulfate, look into "The war gasses" by Mario Sartori for a detailed synthesis.

Oleum is also employed as a condensing agent in organic chemistry, for example for the condensation of chloral with chlorobenzene to form DDT, the known powerful insecticide.

I hope that this synthesis is of use for you.
I am further developing the process and start to make larger amounts of SO3, as soon as I get more persulfate.
My batch was so small since I had only a small amount of Na2S2O8 left. I'll make pics of a larger batch when I get more.

Oleum, even if you can buy it, is horribly expensive from any of the known chemical suppliers (over 100$per Liter). Therefore, a good synthesis for SO3 is necessary. [Edited on 14-3-2006 by garage chemist] Edit by Dav: Topic name extended to aid differentiation between this method of synthesis, and Fleaker and NERV's catalytic V2O5 one. [Edited on 3-8-2007 by The_Davster] Magpie lab constructor Posts: 5939 Registered: 1-11-2003 Location: USA Member Is Offline Mood: Chemistry: the subtle science. That's very nice work gc. Thanks for putting it into publication. I hope to make use of it someday as I've always wanted some thionyl chloride. Aside: I enjoyed the blowup of your pictures. Especially the incongruity of the bottles labeled in German along side the ad for Chilly's bar & grill. The single most important condition for a successful synthesis is good mixing - Nicodem woelen Super Administrator Posts: 7764 Registered: 20-8-2005 Location: Netherlands Member Is Offline Mood: interested GC, this is a very nice result. I have a few ideas, which might be helpful for increasing the yield or using cheaper chems. Could adding a small amount of H4P2O7 be beneficial for the yield? H4P2O7 can be prepared easily by heating H3PO4 (85%) until no water is coming off anymore and the liquid becomes somewhat syruppy. I can imagine that adding 1 ml of this, 2 ml of 96% H2SO4 and 20 to 25 grams of Na2S2O7 can give a somewhat higher yield. Another thing. Why don't you use the much cheaper and easier to obtain NaHSO4 (pH-minus for swimming pools, over here I pay just EUR 20 or so per 2.5 kilos). This then first needs to be preheated, until it has lost all its water. The then dry Na2S2O7 is mixed with the H2SO4 (and maybe H4P2O7) and then you perform the procedure as described above. These are just my two ideas. Unfortunately I have no opportunity to do this at my home (no suitable room and outside to many little children's hands around ), but I'm very interested in your results. The art of wondering makes life worth living... Want to wonder? Look at https://woelen.homescience.net garage chemist chemical wizard Posts: 1803 Registered: 16-8-2004 Location: Germany Member Is Offline Mood: No Mood I do not want to add phosphoric acid, as it will attack the glass. The 100ml rbf you see in the first picture has already been badly attacked by an experiment with HPO3 + H2SO4 SO3 synthesis. For the NaHSO4: I wish that it would work! But unfortunately it doesn't. I have 1,5kg of NaHSO4, it would be perfect if I could make a lot of SO3 with it. The problem is to decmpose it just to the pyrosulfate and not further. It seems like just H2SO4 evaporates when NaHSO4 is heated. I already tried that once. Maybe I just added too much H2SO4, so that too much water was carried into the reaction... I'll probably have to try it again as my supply of sodium persulfate is running very low. My nearest electronics shop only sells in 100g portions, and the electronic supplier who sells 2kg bags of it has a rather high price. @ woelen: if you have a fume hood you are perfectly able to do this experiment. It's quite simple really. And the fuming can also be kept low if a good still like mine is used. [Edited on 15-3-2006 by garage chemist] [Edited on 15-3-2006 by garage chemist] woelen Super Administrator Posts: 7764 Registered: 20-8-2005 Location: Netherlands Member Is Offline Mood: interested I understand your problem with the phosphoric acid. This possible optimization can be ruled out then.  Quote: Maybe I just added too much H2SO4, so that too much water was carried into the reaction... What I meant is first heating the NaHSO4 (without any H2SO4), such that it looses water and Na2S2O7 remains behind. Next, to this Na2S2O7 you add a small amount of H2SO4 and then you perform your steps. So, there is a two-step process, first dehydrating the NaHSO4 and next making SO3. What I understand from your reply is that you mix NaHSO4 and H2SO4 and then start heating that mix. I'll try tonight what happens if I heat NaHSO4. I'm curious whether this only looses water or also looses H2SO4. The art of wondering makes life worth living... Want to wonder? Look at https://woelen.homescience.net garage chemist chemical wizard Posts: 1803 Registered: 16-8-2004 Location: Germany Member Is Offline Mood: No Mood I meant that when NaHSO4 is heated alone, it will mainly lose H2SO4 instead of H2O, I see that my reply was formulated ambiguously. It seems that the decomposition of NaHSO4 into pyrosulfate and decomposition of the pyrosulfate cannot be carried out as separate steps. When pyrosulfate is formed, it seems like it rapidly decomposes at the same time, making only H2SO4. I got a lot of white fumes when heating NaHSO4. But I'll try it again. I think that with correct temperature and heating time, it will be possible to get it to work. Polverone Now celebrating 18 years of madness Posts: 3164 Registered: 19-5-2002 Location: The Sunny Pacific Northwest Member Is Offline Mood: Waiting for spring Nice work, garage chemist. You can see an initial version of the PDF here: http://www.sciencemadness.org/member_publications/SO3_and_ol... If you are going to post an updated writeup soon, after getting some more persulfate to experiment with, I will hold off on linking to the file on the publications index page. If it will be a while, I will just add the link now. PGP Key and corresponding e-mail address garage chemist chemical wizard Posts: 1803 Registered: 16-8-2004 Location: Germany Member Is Offline Mood: No Mood Thanks for putting the article into PDF format, it looks really nice. I don't think that I will post a second synthesis with larger amounts very soon, as I first have to obtain more persulfate and this will take a while. You can add the link now. Polverone Now celebrating 18 years of madness Posts: 3164 Registered: 19-5-2002 Location: The Sunny Pacific Northwest Member Is Offline Mood: Waiting for spring It has been added. Thank you for the writeup. I had feared that "prepublication" might never be used again! PGP Key and corresponding e-mail address garage chemist chemical wizard Posts: 1803 Registered: 16-8-2004 Location: Germany Member Is Offline Mood: No Mood I already got the persulfate. The electronics supplier delivered much faster than usual. Now I have 2kg Na2S2O8, good for theoretically over 600g SO3. I think it is worth investigating the use of only a very small amount of H2SO4 as catalyst, since H2SO4 always contains about 4% of water. It's Oleum time! garage chemist chemical wizard Posts: 1803 Registered: 16-8-2004 Location: Germany Member Is Offline Mood: No Mood Now I have the time to research SO3 production. Yesterday, two experiments. First one: usage of H2SO4 as catalyst. 23,8g sodium persulfate (0,1 mol) were added to a 100ml round- bottom flask. 1ml conc. H2SO4 was added. Heated, a lot of oxygen was evolved, and a bit of smoke. It melted into a liquid, and a very small amount of SO3 distilled over, just a few drops. I heated for 10 minutes on strongest flame, nothing more. It was left to cool down a bit and 2ml H2So4 added, then reheated and a bit more SO3 distilled over. Added another 2ml of H2So4 and heated for 20 minutes, the flask nearly glowed red. If it hadn't been made of Duran, it would have melted. Total yield: 1,7g SO3 (theoretically, 8g should form). PATHETIC!!! Next experiment: using MgSO4 as catalyst. MgSO4 hydrate (Epsom salts) were heated in a quartz dish for 10 minutes on hottest flame, until everything turned into a white powder (the last mol of crystal water of MgSO4 does not begin to split off until the temperature reaches 200°C, so intense heating is necessary). 3g anhydrous MgSO4 were added to 23,8g sodium persufate and heated in the distilation setup. Again lots of oxygen, but this time a bit of the MgSO4 got blown out of the flask. This was due to too rapid heating. The mix melted, and SO3 slowly distilled off. It took maybe 30 minutes of highest heat until SO3 evolution stopped. Yield this time: 2,8g of again theoretically 8g. A bit better, but still bad. Conclusion: The MgSO4 method is definately the way to go, the yields with H2SO4 are too low to be of any use. The next experiment will use 6g MgSO4 and again 23,8g sodium persulfate. If this gives a higher yield, another experiment with even more MgSO4 will be made. When the optimal amount of MgSO4 has been found, it will be tested if additional drying of the sodium persulfate in an oven before synthesis does help (specifications for technical grade sodium persulfate specify up to 0,5% water content). EDIT: if you choose to try this synthesis yourself, you have to get some advice first. For example, the condenser must not be cooled by water, otherwise the SO3 will solidify in there. Only the receiver must be cooled. The SO3 has a very dangerous tendency to solidify in the outlet of the condenser, which can lead to pressure buildup. It has to be melted periodically with a hot air gun (without removing the receiver!). The apparatus must be absolutely dry. If possible, dry it in an oven at over 100°C. The vacuum connection of the receiver should be connected to a drying tube if possible (I didn't use a drying tube, it still worked, but some moisture can make it into your SO3 and decrease its percentage). SO3 reacts explosively with water. A piece of solid SO3, thrown on water, hisses loudly and swims around like a piece of sodium, with lots of fuming. Liquid SO3 makes a sharp crack each time a drop hits water. Water added to SO3 produces powerful explosions and scatters the contents of the flask. [Edited on 9-4-2006 by garage chemist] www.versuchschemie.de Das aktivste deutsche Chemieforum! freachem Harmless Posts: 14 Registered: 5-12-2006 Member Is Offline Mood: fluid Hi Could Fe(SO3)3 be used for SO3 production 12AX7 Post Harlot Posts: 4803 Registered: 8-3-2005 Location: oscillating Member Is Offline Mood: informative No, that doesn't exist. By a long shot. You need Fe2(SO4)3, and is covered in other threads. Tim Seven Transistor Labs LLC http://seventransistorlabs.com/ Electronic Design, from Concept to Layout. Need engineering assistance? Drop me a message! ciscosdad Hazard to Self Posts: 76 Registered: 6-2-2007 Member Is Offline Mood: Curious SO3 Lovely Work GC. Has anyone pursued the route via Sodium Bisulfate? The figures I have seen are that it dehydrates at ~60C over a period of 4 hours or so to give the anhydrous product. Maybe the secret to getting to the pyrosulfate is this initial gentle dehydration, then the much more vigorous heating to get the Na2S2O7 after all the water of crystallization has been driven off. Along with the use of Magnesium Sulfate, this will give quite easily accessible SO3. Sauron International Hazard Posts: 5351 Registered: 22-12-2006 Location: Barad-Dur, Mordor Member Is Offline Mood: metastable @G C Nice thread! SO3 is always of interest. The procedure you have described looks like a simple and low cost way to generate a small amount of SO3 (c. 5 g from 40 g sodium persulfate). One use you have not suggested would be to add this SO3 to ordinary conc H2SO4 to prepare 100% H2SO4 - not oleum but water free. Such dry sulfuric acid is normally prepared by adding the calculated amount of oleum of a known % SO3 to ordinary acid to cancel the water content, Since oleum is expensive this is somewhat onerous. Your method is inexpensive as sodium persulfate is$23 a Kg, (Acros).

It is unfortunate that your method does not scale up (as far as you have stated.)

You are quite right to advise baking out the glassware. Personally I would suggest positive pressure dry N2 atmosphere and a bubbler (maybe Hg) on the other side of a cold trap, so product is solated from atmospheric moisture.
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While experimenting with tellurium, I came across a nice observation. When Te is added to concentrated H2SO4, then first a red/purple solution is obtained, but on continued heating, just below the boiling point of H2SO4 copious amounts of white crystals are produced. According to literature, these crystals are 2TeO2.SO3, which decompose above 500 C, giving solid TeO2 and vapor of SO3.

TeO2 in turn can be dissolved in moderately concentrated hydrochloric acid or sulphuric acid and is easily reduced to elemental Te again, which separates as a black coarse precipitate.

I have done the experiments, except the real making of SO3 (I do not have suitable equipment at the time to handle and isolate such a dangerous compound). I dissolved Te-metal in concentrated H2SO4, heated till all was converted to a beautiful crystalline white solid. I added this white solid to water (together with acid, sticking to it) and obtained a colorless solution, containing tellurium(IV) and recovered most of the Te as a black precipitate.

Would this be a suitable method of making SO3, assuming that the Te can be recycled? I myself have no real equipment to fully investigate this, due to (1) the extremely corrosive properties of SO3 and (2) the risk of being exposed to volatile Te-compounds, making me smell like hell for a long time. Someone with a true fume hood and suitable micro distillation apparatus, however, could give it a try on gram scale. I just post the idea, which came up in my mind.

[Edited on 29-7-07 by woelen]

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Fleaker
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Woelen, this may be unfounded speculation, but I would worry about volatilization of TeO2 at those red hot temperatures. I am interested in giving it a shot, but it's no less thermally intensive than the method that I will very soon show you all: the long awaited Sulfur trioxide via vanadium pentoxide catalyst.

I'm planning on starting a new thread here in prepublication on this method as this current thread deals with the persulfate route.

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I tried the persulfate route (NaHSO4 ---> H2O----->SO3 + NaSO4) in a tube furnace using all glass at 350C to drive off water than 550-600C to obtain SO3 and Iron oxide catalyst under positive N2 pressure at first and then O2 gas at 500C. I definitely got some SO3. Maybe 2-4g. Im thinking that most of got converted to H2O4 as someone already mentioned. I got the gamma form so I was happy with that. Has anyone seen a blue material contaiminate their product? It has low bp since under vacuum, was stripped fairly quickly.

There is a patent out there with a Ytterbium/La/FeO catalyst (easy to make) that catalyzes SO2 to SO3 in the presence of air fully to SO3. Havent tried it but the materials for the catalyst are cheap.
alphacheese
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 Quote: Originally posted by ciscosdad Has anyone pursued the route via Sodium Bisulfate? The figures I have seen are that it dehydrates at ~60C over a period of 4 hours or so to give the anhydrous product. Maybe the secret to getting to the pyrosulfate is this initial gentle dehydration, then the much more vigorous heating to get the Na2S2O7 after all the water of crystallization has been driven off.

According to US patent 6767528 sodium bisulfate decomposes to produce sodium pyrosulfate and water at about 240° to 250° C. Sodium pyrosulfate then decomposes to give sodium sulfate and sulfur trioxide at close to 460° C. It doesn’t discuss the need for the H2SO4 catalyst because the SO3 is mixed with the H2O from the decomposing bisulfate in this patented process.
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After thinking about the procedure for a while I have produced a couple of suggestions:
1) Heat the persulphate on its own to evolve oxygen and remove any water content. This should produce the pyrosulphate:
2Na2S2O8 => 2Na2S2O7 + O2
2) Use some of the SO3 that you have already produced to make 100% H2SO4 as Sauron has suggested. Then use this 100% H2SO4 as the catalyst for this reaction. This should eliminate most of the water that would have otherwise been present.
3) Possibly use a mixture of anhydrous MgSO4 and 100% H2SO4 as the catalyst. I dont know if it will make any difference but surely its worth a try.

I would also like to congratulate you on writing up this experiment, I too think that SO3 production for the home chemist is important as it is a very useful chemical. I myself have some sodium persulphate that I used to etch a PCB for my physics coursework project, and would like to use it to produce some SO3 but unfortunately I have yet to get any labware, and doubt if I will have the money to do so anytime soon
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Perhaps the low yield of the H2SO4 catalyzed reaction is because alot of SO3 is consumed forming oleum? How hard is it to drive SO3 out of oleum?

One shouldn't accept or resort to the mutilation of science to appease the mentally impaired.
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No so difficult at all vulture; heat will drive it out, and so will vacuum.

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garage chemist
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In the meantime I have succeeded in making SO3 from NaHSO4.
It worked extremely well, I got 23,8g SO3 from 100g pool pH-minus.
It's simply a matter of enough heat.
From 680-880°C, plain sodium pyrosulfate (from NaHSO4 at 480°C) gives off all its SO3.
And the best part: not a bit of it is decomposed to SO2. Because decomposition of SO3 requires a catalyst, like for example iron compounds, which are lacking here.

I just have to write a documentation.
Oleum has now become a simple OTC preparation.

[Edited on 22-3-2008 by garage chemist]

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DJF90
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Thats amazing news! Any idea when the write-up will be complete? I can't wait to read this
garage chemist
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I already did the writeup in the german forum, I just didn't get around to doing it here.
I'll do it this weekend, promised.

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 Sciencemadness Discussion Board » Special topics » Prepublication » Synthesis of Sulfur Trioxide and Oleum: The persulfate method Select A Forum Fundamentals   » Chemistry in General   » Organic Chemistry   » Reagents and Apparatus Acquisition   » Beginnings   » Responsible Practices   » Miscellaneous   » The Wiki Special topics   » Technochemistry   » Energetic Materials   » Biochemistry   » Radiochemistry   » Computational Models and Techniques   » Prepublication Non-chemistry   » Forum Matters   » Legal and Societal Issues