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Author: Subject: oleum & SO3
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[*] posted on 18-9-2005 at 12:06


Hmm... the CaC2 idea is certainly worth a try.

But it might just react like this:

CaC2 + H2SO4 ---> C2H2 + CaSO4

Balanced equation, no water needed, and no SO3 produced.
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[*] posted on 18-9-2005 at 14:12


That makes more sense, since the remaining CaO is a strong alkali and would absorb any SO3. Thus, it's a simple displacement reaction between the very weak acid C2-- and the strong acid SO4--.

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[*] posted on 18-9-2005 at 16:21
Epiphany! Idea anyway...


So I was continuing my continual reading on the topic of chemisty today and I was thinking about borates and boric acid and decided to read about them. Well...
Quote:
from Wikipedia:

Boric acid is soluble in boiling water. When heated above 170°C it dehydrates, forming metaboric acid HBO2. Metaboric acid is a white, cubic crystalline solid and is only slightly soluble in water. It melts at about 236°C, and when heated above about 300°C further dehydrates, forming tetraboric acid or pyroboric acid, H4B4O7. Boric acid can refer to any of these compounds. Further heating leads to boron trioxide.

Anyway so how does this relate to the topic? Well this reminded me of metaphosphoric acid which also forms via condensation reactions. Perhaps pyroboric acid and tetraboric acid could be refluxed with H2SO4 at a temp ~ 330dC and the former would suck out the water converting back into H3BO3 while the H2O + SO3 equilibrium would be pushed toward excess SO3 which could be distilled out. Wikipedia says nothing of any dehydrating abilities but it's worth a try no?
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[*] posted on 24-9-2005 at 15:34


I've repeated Garage Chemist experiment, with succes!!

I heated 55ml 85% H3PO4 in a copper 'crucible', home made from a copper plate.
This was heated till no more bioling was observed, and the liquid went a bit blackish. Unfortunatly there was a small hole in the crucible, and I asked someone to repair it, but he did it with just normal solder, which happyly reacted with the (extremely) hot acid, and turned it greenish.
To check whether is was far enough I took a very small sample with a pasteur pipette, en let it cool down. It was solid as glass, so ok.

I poured the stuff after a bit of cooling into a dried 1L 3-neck RBF, and went for tea. After tea the stuf was rock-solid, and I added 30gr. H2SO4 (95-97%). A leibig condensor was attached, and a small erlenmeyer in an ica bath.
I heated with a propane torch, but that was not good enough, some SO3 was coming over, but not much. Later I used a 'campinggaz' heater, which worked a lot better. However, my lab-thermometer broke recently so I had to use another one which only goes to 50°C... not good, it needed to be replaced by a stopper rather soon. I stopped the destillation before all the SO3 came over, but I saw that some came over, which was the most important for me, I didn't want any water condensing, and withut proper temperature control (I now used my fingers to 'read' the temp.) I felt a bit uncomfortable.
My SO3 didn't become a liquid, but into a very fine cristaline form on the walls of the receiving flask, it looked very pretty, a bit like the sugarthings you can buy on a fair. (at least in NL, we call the "sugar spiders"...)
Upon opening thick wite clouds are emitted, really, really impressive. I didn't do any tests with it, nor did I try to determine the yield.

The only thing I did not like was that seemigly some of the stopps I used in the 3-neck had still some oil on them, so part of the SO3 is blackish, while some is snow-white. Some drops which fell out the condensor when I disassembled the destillation setup reacted vigoriousy with paper tissue, it eated the paper immidiatly away with a hissing sound, that was scary! I weared double latex gloves, which was reasonably resistant to SO3/oleum.

When I have a new thermometer I will distill the rest of the SO3 out of the H3PO4/H2SO4/HPO3/SO3/H2S2O7/H4PO7 mixture.

This is a very doable procedure imho, with no hard to obtain precursors!
I assume that after a complete destillation P3PO4 can be destilled off the H2SO4, and re-used to dehydrate to HPO3?

[Edited on 24-9-2005 by Taaie-Neuskoek]




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[*] posted on 24-9-2005 at 18:38


Quote:
Originally posted by Taaie-Neuskoek
I assume that after a complete destillation P3PO4 can be destilled off the H2SO4, and re-used to dehydrate to HPO3?


Hmm, in a properly heat- and chemical-resistant flask, you could add some H3PO4, distill the water out of it, let it cool then add H2SO4, distill/sublimate SO3 off, distill remaining H2SO4, dehydrate H3PO4 and repeat. :) Or you could use a lot of HPO3 for a small amount of H2SO4, reducing the process by a step. :D

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[*] posted on 25-9-2005 at 03:10


Nice to hear of your success, Taaie- Neuskoek.

The SO3 crystallizing in the receiver is actually a good indicator of purity, as pure SO3 melts slightly above room temperature while high- percent oleum is liquid.

A nice thing to do is to put a few drops of liquid SO3 into an empty PE (not PET!) plastic bottle with its screwcap removed
and shoot very dense and impressive white
smoke rings by gently tapping the bottle.:D
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[*] posted on 31-1-2006 at 22:59


I came across an interesting reaction the other day
B2O3+K2SO4--> SO3+ K3BO3
Rxn occurs at red heat.

(It almost seems like on of those too good to be true reactions)




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[*] posted on 1-2-2006 at 07:16


Makes sense to me. Phosphate works, too, having diagonally similar behavior (glassy melts, low melting point (red heat), dehydrates to form molten anhydride, etc.). Silicate has too high a melting point, and I'm not sure if any other acidic neighbors (As2O5, SeO2, GeO2) have glassy, low-melting behavior. Alumina has an extreme melting point, and it's a sharp melting point, not glassy.

Iron sulfate still has a lower decomposition temperature though, AFAIK. Next to that, pyrophosphate is best.

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[*] posted on 1-2-2006 at 07:50


At red heat though a noticeable % of your sulfur trioxide will disproportionate though, will it not? I'm not at home to check the numbers right now but there is a temperature above which the losses get pretty tremendous.



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[*] posted on 1-2-2006 at 10:04


I recall someone making mention of that, even at pokey temperatures like used to decompose iron sulfate. For sure, burning out sulfate-rich ceramics (like something containing epsom salts) makes for a lot more stink than an in-your-face,nose-and-eyes burning sensation.

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[*] posted on 1-2-2006 at 11:11


Quote:
Originally posted by garage chemist
-snip-
The distillate, on pouring it into a beaker, fumed incredibly strong and emitted so much smoke that I had to turn my fume hood on maximum power. The exhaust pipe outside of my lab emitted a stream of white smoke which filled the garden.
As the liquid contacted some moisture in the beaker, a loud crackling noise was observed and the beaker erupted even more of the thick white smoke. You have to see it to believe how much a liquid can fume in air. It's a real spectacle.
heat. It can be reused indefinately for dehydrating H2SO4 to SO3.
-snip-

I've seen 10 gallons of Oleum dumped in a flat metal tray at a haz-mat training class. This was over 200 pounds of material. The display took place in an open desert area with a good breeze blowing away from the crowd. Within 15 minutes there was a plume of smoke stretching 5 miles away to the mountains. The smoke was so thick you couldn't see through it. When the trainees, in full protective gear, hit the pan with fire hoses as quickly as they could, there were tremendous thumping explosions, which were expected. This is the only way to stop the fuming, although it actually makes it worse for a short while. The smoke WAS truly a spectacle. I watched a sparrow fly through the plume and fall out of the air. This is truly nasty stuff. If you make this, expect a cloud.
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[*] posted on 15-3-2006 at 15:00


NaHSO4 + H2SO4 revisited

In response to the new thread on SO3 and oleum in the prepublication section, I did some tests with NaHSO4 and I did some research on this. I post the results here, in order to avoid cluttering of that thread, which is not really meant for discussion.

I found that NaHSO4 in fact is a hydrated salt, its precise formula is NaHSO4.H2O. When this is heated, then the molecule of water, which is in the crystal lattice, easily is lost.

I did a test in a test tube with appr. 500 mg of solid, kept the test tube almost horizontally, and then heating it with a propane torch. The NaHSO4.H2O easily melts and the liquid starts boiling vigorously. Water condenses in the colder parts of the test tube. This is the first stage of dehydration. What remains is molten NaHSO4, without the H2O.

Next, I heated the test tube as a whole, driving of all H2O. I still had the liquid in the test tube. This liquid is molten NaHSO4. Now, after much more heating, I again obtain a liquid in the relatively cooler parts of the test tube (but these still are quite hot). A drop of this liquid quickly chars paper tissue (it becomes black almost at once) and this liquid can withstand a lot of heat, before it evaporates. It almost certainly is concentrated H2SO4, with possibly tiny amounts of water in it. The liquid does not fume in contact with air. Inside the test tube there was some white mist.

I continued heating for 10 minutes in the flame of the torch. This gives me a little bit more of the colorless and very corrosive liquid, but the NaHSO4 (or whatever remains) does not solidify. At this point the test tube was very hot and I stopped heating, being afraid that the test tube with the ultrahot and corrosive liquid salt cracks or melts.

I let the test tube cool down. The molten salt solidifies to a white crust. The drops of liquid become oily, but they do not solidify. Next, I dripped a few drops of 96% H2SO4 in the white solid and then started heating again. All solid quickly melts and mixes with the H2SO4. There is some boiling (water from the H2SO4) and then a mobile colorless liquid is obtained. This liquid was heated much stronger for 10 minutes or so, but this only gives oily drops inside the relatively cool parts of the test tube.

From all this, I conclude that even dehydrating NaHSO4, getting Na2S2O7 is very hard and requires really high temperatures. Also, NaHSO4 probably decomposes giving H2SO4 and Na2SO4 instead of Na2S2O7 and H2O (as Garage Chemist mentioned in the thread in the prepublication section). I'm afraid that most people over here, who did tests with NaHSO4.H2O regarded the first boiling as the dehydration of the salt giving Na2S2O7, but this only is the boiling away of water of crystallization.

Here is two MSDSes, which tell that sodium bisulfate is NaHSO4.H2O instead of NaHSO4.
http://www.jtbaker.com/msds/englishhtml/s3050.htm
http://www.physchem.ox.ac.uk/MSDS/SO/sodium_hydrogen_sulfate...

Anhydrous NaHSO4 apparently is a much more difficult to obtain chemical and it is really expensive. What many sellers call "anhydrous" in reality is NaHSO4.H2O (same CAS number), it only is "anhydrous" in the sense that it is a dry free flowing powder.




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[*] posted on 15-3-2006 at 15:29


I did similar experiments today, I heated 90g of sodium bisulfate in a 100ml rbf with a 100W heating mantle.
It boils gently for about an hour, during which the temperature rises.
At the end, all steam evolution stops, and the liquid is extremely hot. I added a drop of H2SO4 to it, and the drop boiled violently!
When the temperature was somewhat lower, I added 7ml of H2SO4 to it, and heated it again.
Eventually it started to boil, but no SO3 was generated. Absolutely none!
Adding an extra 3ml of H2SO4 didn't change anything.

It seems like the preparation of sodium pyrosulfate is difficult, requiring either the use of sodium persulfate at modest temperatures (somewhat expensive) or sodium bisulfate and extremely high temperatures (difficult to reach the temperatures, also dangerous).

EDIT: Ulman has this to say about sodium bisulfate:

2. Sodium Hydrogensulfate

Sodium hydrogensulfate occurs as the monohydrate, NaHSO4 · H2O, in the system sodium sulfate – sulfuric acid – water, or exists as the solid phase NaHSO4 at a sulfuric acid concentration of 62 %. The monohydrate is converted to the anhydrous salt at 58.45 ± 0.05 °C.
The thermal decomposition represented by the equation
2 NaHSO4 <----> Na2S2O7 + H2O
takes place near the melting point, which can be determined only approximately (ca. 183 °C) with a water vapor pressure of 2500 Pa (25 mbar). Conversion to sodium disulfate is complete after heating for ca. 4 h at 240 – 260 °C. Sodium disulfate decomposes above 400 °C to form sodium sulfate with liberation of sulfur trioxide.

[Edited on 15-3-2006 by garage chemist]
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[*] posted on 16-3-2006 at 15:11


Hmm, what about decomposition of the NaHSO4 in Vacuum?
This will drastically lower the partial pressure of the water vapor over the melt and therefore shift the equilibrium to the right side.
I should try that out.
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[*] posted on 25-6-2006 at 05:00


I now have some P2O5, and I did the experiment of adding this to conc. H2SO4 and heating this. The P2O5 dissolves quickly. After that, I continued heating. I obtained a colorless and quite mobile liquid, but no fumes at all. I used appr. 0.5 gram P2O5 and 1 ml of H2SO4.

I dumped the liquid in a bucket of water (after letting it cool down somewhat). This reaction is quite spectacular, on contact with the water it made a fairly loud crackling noise.

What temperature do I need to get the SO3 out of the colorless liquid? The liquid I had was already quite warm (I estimate it at 200 C or so).




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[*] posted on 25-6-2006 at 06:05


Maybe this is of some help:

Boiling points of oleum:
247C - 5%
200C - 13.5%
150C - 24%
100C - 40%
75C - 55%
50C - 89%
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[*] posted on 25-6-2006 at 06:45


"What temperature do I need to get the SO3 out of the colorless liquid?"

I've done this about 30 years ago to make 50g or so of SO3. Really, really hot, near the boiling point of H2SO4.
Too hot to use a mercury lab thermometer. Use all pyrex apparatus, and cool the condenser with a stream of air, not water. I think I used a hot air gun for heating, 700C or so. (Cracking the glass could be really disasterous.)
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[*] posted on 11-9-2007 at 05:50


Today I successfully made some SO3. I dehydrated ~100mL of H3PO4, not at red heat but pretty bloody hot, for ~1/2 hour. I haven't washed out my beaker yet (when I do I will be trying to save the phosphoric acid in it, I tested today that it doesn't matter if you wash it in with water, just takes longer to dehydrate) but I believe it was attacked savagely. I am also not sure of whether it has enough integrity left for another try, I will find that out too. I added 20mL of sulfuric acid (a lower ratio than garage_chemist, but that was the ratio I calculated???) I obtained perhaps 10g or perhaps more of very nice SO3. Then I got greedy and added another 20mL sulfuric acid through the top and perhaps 1/4 mL ran down the condenser into my product but thats ok because I will redistill after another batch or 2 anyway. Not much more SO3 came across,however the temperature rose on the vapor thermometer while I wasn't paying attention... A drop of water condensed on the thermometer and fell back in resulting in a fairly decent 'something' - half way between bumping and an explosion, which almost certainly dumped some liquid into my distillate (again, redistill, so I don't care + its still a solid, albeit a 'wet' solid)

@ garage_chemist - If you can remember back to your experiment posted on 5/9/05 (pg4)- what kind of glassware did you use, mine was pretty well attacked, i'm sure. Also when you said steam was coming off, do you remember if it was just steam, mine looked like smoke for ages, and produced a vapor resembling nitric acid fumes (in low concentration) only slightly more choking somehow... it was still producing this when I decided to turn off the heat so I didn't have my beaker give out over an LPG cylinder. Finally, do you have any idea what volume contraction took place (may have been difficult to observe with only 14mL). Mine was down to 70-85mL when I used it... cant be more accurate because the liquid level was obscured by dense white in the beaker.

Btw, my phosphoric acid was lab quality, but I inherited it over a decade ago and the bottle it was in isn't at all etched so I don't think it was cation contamination, although it looked a hell of a lot like sodium phosphate(s) I have made before. Actually one other thing, was your (HPO4)n still colourless? while most of the white stayed in my beaker the acid acquired a granular white colour.

If you cant remember any of this, doesn't matter, it was 2 years ago after all.
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[*] posted on 11-9-2007 at 08:14


I used a quartz dish for concentrating the H3PO4 which was attacked, but not too severely.
Yes, the "HPO3" solidifies when cooling and becomes somewhat crystalline.
I also observed some "smoke" during the concentration, as you described.
And the mixture (or was it the HPO3 itself?) became white as well at some point, probably due to silicon compounds from the attacked quartz becoming SiO2 again.




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[*] posted on 11-10-2007 at 17:22


Quote:
Originally posted by woelen
I now have some P2O5, and I did the experiment of adding this to conc. H2SO4 and heating this. The P2O5 dissolves quickly. After that, I continued heating. I obtained a colorless and quite mobile liquid, but no fumes at all. I used appr. 0.5 gram P2O5 and 1 ml of H2SO4.

I dumped the liquid in a bucket of water (after letting it cool down somewhat). This reaction is quite spectacular, on contact with the water it made a fairly loud crackling noise.

What temperature do I need to get the SO3 out of the colorless liquid? The liquid I had was already quite warm (I estimate it at 200 C or so).


0.5g P2O5 to 1ml of H2SO4? You got to be joking ;) To get sulfur trioxide you have to add 1 mole of P2O5 per mole of H2SO4 this is about 2.66g of P2O5 for 1 ml of 100% H2SO4. I used 150g P2O5 + 75 ml H2SO4 and distilled off SO3, and got 72% yield counting on H2SO4. Mixture was heated with standart lab spirit-lamp. Quantities of reagents must be taken from following reaction equation:

P2O5 + H2SO4 => 2HPO3 + SO3

I added less then calculated P2O5 mass (200g) so yield is reduced respectively. Dont forget to count loss of P2O5 from reaction with water if your H2SO4 is not 100%. To count amount of P2O5 lost due to the reaction with water you must use hydration scheme of P2O5. Below 20C P2O5 reacts with 1 molecule of water to form (HPO3)x - polymeric crystaline mass, between 20 and 100C P2O5 reacts with 2 molecules of water to form H2P2O7, and on boiling around 100C it finaly forms H3PO4 reacting with 3 molecules of water.

By the way N2O5 also forms only if you get 1 mole P2O5 for 2 moles of 100% HNO3 (P2O5 + 2HNO3 => N2O5 + 2HPO3).So your problem is that you have taken too small amount of P2O5 - that's why you got nothing.

Also i must give you a tip: If your apparatus is sealed from air moisture, fumes of SO3 are colorless and you will not see them, but if you will use good condenser liquid SO3 will be seen as clear transparent liquid, dropping to reciever flask. SO3 is somethat hard to condense in right way, if you run water through condenser with temp around 20C to make sure SO3 is still liquid and maximum cooling is acchived, some SO3 escapes uncondensed. If you use cold water in condenser your SO3 will freeze somethere in it and block the tubbing, so good ballance is needed.

Here is photo of my frozen SO3 drops, with needle like crystall layer at their surface:

[Edited on 12-10-2007 by Engager]

SO3.jpg - 81kB




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[*] posted on 11-10-2007 at 20:06


@ Engager, Nope.... I made SO3 using dehydrated H3PO4...... and this was not P4O10, so clearly under the right conditions the metaphosphoric acid will dehydrate H2SO4.

SO3 is a pain in the bum is what it is. Freezing in the condenser at the same time as fuming out the distant vent. Next time I attempt to make it (and it will probably HAVE to be SO2 + O2, since the damage to the beaker from heating H3PO4 was too extreme and the acid can not be recycled, its too full of crud) it is going to be condensed by going through glass tubing melted to a Pasteur pipette so that it is blown hard out the small hole at the (in)side of a flask sitting in dry ice/acetone. I will just waste excess oxygen so that no SO2 comes across. But that is a plan for the distant future.
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[*] posted on 30-3-2008 at 23:55


If anyone wonders what Gmelin's Handbuch has to say about thermal decomposition of sulfates, wonder no more. Not that it will be of any importance. It may be more misleading than anything else - I have a feeling that very little SO3 is given off by Li pyrosulfate at said temperature.

Attachment: gmelin_thermal_sulfate_dec.pdf (116kB)
This file has been downloaded 1264 times

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[*] posted on 7-9-2008 at 19:32


I have never seen this mentioned before so I thought I'd share it and see what you thought.

In cold anhydrous Et2O (maybe other solvents?)

Na2S2O3 + 2HCl > H2S2O3*2Et2O + 2NaCl
H2S2O3 > H2S + SO3

How feasible is this? It should result in very pure SO3 if the ether can be separated.
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[*] posted on 8-9-2008 at 05:43


What about H2S + SO3 <--> SO2 + H2O + S ? For that matter, the driving force is solid sulfur, besides that SO3 is an oxidizer.

Is thiosulfuric acid even stable under any condition? Seems to me it would love to polymerize as-is, forming sulfur and sulfurous acid (which is in equilibrium with SO2, thus getting at the same thing as above, just in solution).

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[*] posted on 8-9-2008 at 20:00


Quote:

What about H2S + SO3 <--> SO2 + H2O + S ?


Yeah you're probably right about that.

Thiosulfuric acid is stable up to about 0*C where it supposedly decomps to H2S and SO3. I'm not sure why it wouldn't go to SO2 + H2O + S straight away, you would think that would be the outcome.
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