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Author: Subject: oleum & SO3
MadHatter
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[*] posted on 2-8-2004 at 21:33
Clay


axehandle, IIRC, H2SO4 used to made by driving out the SO3 from
ferrous sulphate in ceramic retorts as you have described. This was
before modern catalytic methods but as I understand it, the retorts
were at red heat during the process.




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[*] posted on 3-8-2004 at 19:11


Quote:

...
before modern catalytic methods but as I understand it, the retorts
were at red heat during the process

Not red heat. The temperature is a measly 480 degrees for the decomposition of Fe2(SO4)3.

About my heating of NaHSO4: Something is very fishy here. Different sources give different temperatures for the formation of the pyrosulfate, so I kept the oven at slightly over 400 degrees C (fluctuating between 400 and 420 C) --- after approximately 4 hours of heating without anything interesting happening, the vat inside started billowing SO3 (or at least some annoying smoke that I'm at least sure wasn't SO2 (no smell)). My guess would be that what happened was that the temperature was a bit borderline w.r.t. SO3 formation and that once all formed H2O was driven out, the pyrosulfate started decomposing.

I'm right now keeping another batch at 300C to find out the right temperature for formation of the pyrosulfate without decomposition of same...

Still, is there anyone with hands on experience with heating NaHSO4 who'd like to post his observations?

Edit1: Have now had the NaHSO4 at 300C for about 8 hours. Water has evolved, but the mass hasn't solidified. I'm going to keep it at that temp. for 8 hours more and see what happens...

Edit2: Now it's been i the oven for what.. 30 hours or something. It's almost completely solidified. Probably 10 more hours should do it, although this isn't healthy for my electrical bill...


[Edited on 2004-8-4 by axehandle]

[Edited on 2004-8-5 by axehandle]




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[*] posted on 31-8-2004 at 06:43


Axehandle, what happened to this experiment? Did it work out?
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[*] posted on 12-9-2004 at 04:32


I'll soon get a lot of phosphorus pentaoxide and, as I know, it can dehydrate sulfuric acid to SO3.
I have made experiments with pure SO3 (I have purchased a little P2O5 before, but I didn't want to use it all to make SO3) and it is fun stuff. It fumes in air like you won't believe. A beaker with a little SO3 in it is at least as good as a burning KNO3- sugar mix.

A drop of liquid SO3 on wood makes a black charred spot instantly. It's extremely aggressive. I don't want to know what it does to skin.

I'll soon be able to make lots of it (100g P2O5 can make about 30- 35ml of liquid SO3), so do you know any useful application? I know that it can be used to make powerful nitration mixes (a mix of liquid SO3 and 100% HNO3 should make maximum quality NC!).

The SO3 is produced by mixing H2SO4 and P2O5 and then distilling. The SO3 condenses without problems if you cool the receiving flask in ice water. That's the easiest way to produce SO3.
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[*] posted on 12-9-2004 at 17:10


Does anyone have any experience with the pyrosulfate + acid method?
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[*] posted on 12-9-2004 at 18:05


Quote:

Axehandle, what happened to this experiment? Did it work out?

Yes, I just forgot to write about it... I've got about 3kg worth of Na<SUB>2</SUB>S<SUB>2</SUB>O<SUB>7</SUB> in my dessicator (box with CaCl<SUB>2</SUB>...). The comparison batch of about 500g which I kept in an open jar in the kitchen for several weeks is completely covered with water it has absorbed from the air. The lump in the dessicator is still completely dry.

It took 30 hours in a 300C kitchen oven to make the large batch.

I have yet to try a controlled thermal decomposition yielding SO<SUB>2</SUB>. I need some glassware which I can't afford to order right now to do it in a repeatable, conclusive manner.


[Edited on 2004-9-13 by axehandle]




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[*] posted on 1-6-2005 at 10:30


I'm interested in obtaining SO3 and after reading know 3 ways of doing so:

1) Chemoleo's method of thermal decomp of sodium persulfate.
2) P2O5 dehydration of sulfuric acid.
3) Ferric sulfate - Fe2(SO4)3 thermal decompostion

I most likely want to go with the thermal decomp of sodium persulfate because of that fact that it is cheap and the decompostion temp isn't as high as ferric sulfate's.

I have a distillation retort in which I plan on doing my procedure. The resulting SO3 gas will travel through the arm and should condense in a recieving apothecary jar in an ice bath. The liquid should then freeze or solidify into crystals if I am correct on my mp and bp info. Any thoughts? I believe that with my equipment (a distillation retort) that this procedure best fits what I have to work with.
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[*] posted on 1-6-2005 at 13:01


I wouldn't use a retort for this procedure for the simple fact that it is open to the air. While the SO<sub>3</sub> would condense, a lot of it would still evaporate, forming a dangerous acid mist. Water vapor would also come into the receiver and form sulfuric acid.
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[*] posted on 1-6-2005 at 13:30


I tried this once, condensing the SO3 from persulphate decomposition. Couldnt get it to work, it wouldnt visibly dissolve in H2O either. However, what should work is to dissolve the SO3 into conc H2SO4 directly. This is how it's done industrially too.
As to SO3 gas - at the time I reeked out my whole room and it didnt even make me cough (well a little). It was odd how harmless the smoke seemed.

If you got some sodium persulphate could you please test the following: heat a weighed amount (i.e. weigh the whole testube containing the Na2S2O8), then heat until SO3 formation and bubbling stops (it's a clear liquid). Then weigh again, and let us know. I'd like to know the mass loss that occurs, i.e. whether oxygen is evolved as well, or not.




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[*] posted on 1-6-2005 at 15:08


Hmm I wish I could do that as it seems a good experiment to test for that but I dont have any at the moment, I am planning on purchasing as soon as I can determine a good procedure for obtaining SO3 using that method.

However do you know of any in-town places that might supply it? I was planning on buying 500g online since I haven't seen it anywhere else.
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[*] posted on 1-6-2005 at 15:16


http://www.jtbaker.com/msds/englishhtml/s4730.htm

Looking at number 10 on there under dangerous decomposion info it says that oxides of sulfur and oxygen is released.

Also perhaps I could avoid the hazards of the acid mist by placing the end of the retort directly in conc. sulfuric acid yielding oleum.
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[*] posted on 2-6-2005 at 03:46


Yes this would be especially interesting since the MSDS shows the persulfate melts with decompostion at 180 degrees C.:D THe lower temperature and the abundance of evolved oxygen should prevent sulfur dioxide from coming over. Sounds promising actually

[Edited on 6/2/2005 by chloric1]




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[*] posted on 2-6-2005 at 09:58


Wow this is really getting me excited, One could easily use this method to recharge H2SO4 or convert to oleum. If pretty pure SO3 could be isolated in liquid, one could have a superior nitration but adding the SO3 to 70% HNO3 would easily make a 99%HNO3 and H2SO4 mixture. Can anyone say 95%+ yield... :P
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[*] posted on 3-6-2005 at 20:41


Alright as of now I cannot find a good source of sodium persulfate that I want to buy from. I am however going to try this production of SO3 using ferric sulfate Fe2(SO4)3 which decomposes at 480*C. I did a test tube test today and successfully melted KClO4 which has a 600*C+ melting point so I know that I should be able to get SO3 out of ferric sulfate.

One question, are there any special specifications to produce SO3 from ferric sulfate eg. oxygen free environment or will a test tube work? I plan on first doing a few test tube tests and then perhaps move it to my retort and attempt to bubble the SO3 through some conc. H2SO4, as well as isolate it in liquid form.

I'll report back when I get the stuff in. I will order some next week.
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[*] posted on 3-6-2005 at 20:47


Melting point of KClO4 according to my books is ~400C just be careful when decomposing your ferric sulfate not to get the temp too hot otherwise your SOx will dissociate and less SO3 will be produced, I can't see the need to facilitate an inert atmosphere as the gasses produced will drive out any atmosphere present and I can't see any constituents of the atmosphere interfering with this reaction.



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[*] posted on 4-6-2005 at 08:29


Hmm I guess I just got confused as the MSDS says that the boiling point occurs before the melting point at 400*C and 610*C. Oh well either way I will be careful. Thanks for the info.

http://www.jtbaker.com/msds/englishhtml/p5983.htm
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[*] posted on 8-6-2005 at 09:53


I recently heated sodium bisulphate to make pyrosulphate. But how can you tell where you are in the reaction?

I put a PH paper in the fumes and it was quite acidic when I stopped, but the product did not react with water violenly so I might just have had only a little bit of pyrosulphate.
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[*] posted on 8-6-2005 at 10:27


It's stopped when no more water comes off, as vapour, as it essentially is a dehydration reaction.



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[*] posted on 8-6-2005 at 11:16


The mp is a tip-off. If you are heating at the proper temperature, there will be a crust of pyrosulfate above the molten pyrosulfate, unless the layer is very thin. There should be no vapor of SO3, its too hot if there is. You may find it best to very slowly raise the heat until you see the SO3, then back it off a very little - this is the perfect temperature. It should solidify completely within a few seconds when poured out, due to the high mp. I recommend a cheap nonstick frying pan. The product falls out on cooling and is quite brittle. Bisulfate does not have these properties.

It just so happens that a tumbler glass on a small cast-iron skillet on a gas stove provides a good temperature. Good temperature control is very important here, the range that this works within is small.
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[*] posted on 9-6-2005 at 14:03


Quote:
Originally posted by chemoleo
I tried this once, condensing the SO3 from persulphate decomposition. Couldnt get it to work, it wouldnt visibly dissolve in H2O either. However, what should work is to dissolve the SO3 into conc H2SO4 directly. This is how it's done industrially too.
As to SO3 gas - at the time I reeked out my whole room and it didnt even make me cough (well a little). It was odd how harmless the smoke seemed.

If you got some sodium persulphate could you please test the following: heat a weighed amount (i.e. weigh the whole testube containing the Na2S2O8), then heat until SO3 formation and bubbling stops (it's a clear liquid). Then weigh again, and let us know. I'd like to know the mass loss that occurs, i.e. whether oxygen is evolved as well, or not.


I've done this, 1.975g of ammoniumpersulphate was added to a reaction tube, and heated till the bubbling began, a white, non-irritating but inducing a headache smoke was released, the tube was further heated till the whole tube was liquid.
Re-weiged, 1.823 grams was left, meaning that 0.152g had dissapeared.
molar weight of (NH4)2S2O8 is 228.19g/mol.
2(NH4)2S2O8 --> 2(NH4)2SO4 + 2SO3 + O2.
Start was 8.66mM, and assuming the weight loss is (2SO3:O2) it means there is 2.10mM SO3 escaped from the solution, which is a bit dissapointing IMHO, but if it works to make oleum it is worth a shot.
I tried to dissolve the SO3 in conc H2SO4, but that didn't work, the gas evolution was to vigorious, and the SO3 came out on all small holes.




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[*] posted on 17-8-2005 at 12:49


Here is a very nice link. I still plan on using sodium bisulfate to attempt to produce SO3 via pyrosulfate decomposition.

http://en.wikipedia.org/wiki/Sulfur_trioxide
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[*] posted on 17-8-2005 at 13:40
PCB etching agent, as SO3 source


Quote:
Originally posted by Chemoleo on 31-5-2005 at 09:43 PM:
Thread: Ammonium Persulphate production:
https://sciencemadness.org/talk/viewthread.php?tid=3930


....However, heat releases from sodium persulphate SO3 (try it - melt it, and get copious amounts of white fumes)

Na2S2O8 --> Na2SO5 (which decomposes further) + SO3
PCB etching agent, as SO3 source:

Persulphates (Potassium, Sodium and Ammonium salts), are also sold in electronic stores as a PCB etching agent. It works cleaner, and has finer and neater etching properties than Iron(III)chloride, however, it's more expensive though.

Please also read Chemoleo's post on this thread: Oleum & SO3:
Chemoleo, posted on 27-5-2004 at 04:10 AM.

[Edited on 18-8-2005 by Lambda]
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[*] posted on 17-8-2005 at 14:12


Interesting, Taiie.
I must have overlooked your post.
Could you try the same thing with sodium persulphate, I have a feeling that the decomposition of ammonium persulphate might be less straight forward. Some funny odd products could form, reducing the yield during decomposition.
Also your scale must be quite accurate if you are dealing with such small amounts...but then you have a labscale me thinks.

Also, the decomposition of the persulphate does not necessarily stop at the sulphate, it may stop ealrier at the S2O5(2-), meaning no O2 is released. That would bring up your yield to 3.8 mmoles, so ~ 50%. Which brings to an interesting stoichiometry of 1:1 product vs unreacted reactant, which could indeed mean some stable other things are formed.

Anyway, pls try it with the Na2S2O8.




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[*] posted on 17-8-2005 at 14:22


Does anyone know the decompostion temp of sodium bisulfate? I know it's melting point is 58dC. Also after the decomposition to pyrosulfate, what temp is required to make the pyrosulfate decompose?
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[*] posted on 17-8-2005 at 16:18


Quote:
Originally posted by chemoleo
Interesting, Taiie.
I must have overlooked your post.
Could you try the same thing with sodium persulphate, I have a feeling that the decomposition of ammonium persulphate might be less straight forward. Some funny odd products could form, reducing the yield during decomposition.
Also your scale must be quite accurate if you are dealing with such small amounts...but then you have a labscale me thinks.

Also, the decomposition of the persulphate does not necessarily stop at the sulphate, it may stop ealrier at the S2O5(2-), meaning no O2 is released. That would bring up your yield to 3.8 mmoles, so ~ 50%. Which brings to an interesting stoichiometry of 1:1 product vs unreacted reactant, which could indeed mean some stable other things are formed.

Anyway, pls try it with the Na2S2O8.


If I get my Na2S2O8, I will.
Sodium persulphate is hardly used in etching (sp??) here (NL) because of all the eco-tax on it... I have been able to get my hands on a few 100g via a friend, but it is certainly not OTC anymore, everybody uses Fe2Cl3 here.
The scale could handle up to 0.1mg, so was pretty precise, unfortunatly it isn't mine, I did the experiment at my work. (My scale can handle 1mg up to 40 grams...)
About the end products, a liquid was left, which cristalised after some time.

I have done some more efforts to make SO3, one by letting react SO2 and NO2, SO2 was generated by reacting copper and H2SO4, and NO2 by adding starch to 65%HNO3.
These 2 were created in a rbf and a erlenmeyer, both connected to each other and to a vigreux column, and a cooler, where the tip of the cooler was placed in a beaker, placed in an ice bath.
This quick-and-dirty setup yielded nothing but a bit of smoking red stuff, and finally in an NO2 intoxication for me... not really pleasant.
(I did go to a docter, but she couldn't hear anything unusual inside, but my lungs felt like they were hurt the rest of the week...)
I still wonder how it happened, I created a very good draft with a couple of fans, and smelled hardly any NO2...
I will try to repeat this experiment (with full-face mask) in the future.




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