Sciencemadness Discussion Board - oleum & SO3 - Powered by XMB 1.9.11 (Debug Mode)

oleum & SO3

Pages: 1 2

garage chemist - 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.

12AX7 - 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--.

Tim

Epiphany! Idea anyway...

ADP - 18-9-2005 at 16:21

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?

Taaie-Neuskoek - 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]

12AX7 - 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.

Tim

garage chemist - 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.

The_Davster - 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)

12AX7 - 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.

Tim

BromicAcid - 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.

12AX7 - 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.

Tim

Mr. Wizard - 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.

woelen - 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.

garage chemist - 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]

garage chemist - 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.

woelen - 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).

chromium - 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%

Eclectic - 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.)

Antwain - 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.

garage chemist - 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.

Engager - 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

Antwain - 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.

S.C. Wack - 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 2082 times

497 - 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.

12AX7 - 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).

Tim

497 - 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.

Lambda-Eyde - 20-11-2008 at 05:47

Hello !

I've been searching around this forum a little bit and read the bulk of this thread, but I still haven't found anyone mentioning synthesizing SO3 from potassium polysulfide and H2SO4 ? Just see this wonderful demonstration by Mabakken (I think he goes under the pseudonym "ScienceGeek" here at SM

) over at Youtube.

Here in Norway potassium polysulfide can be bought mixed with other sulfides of potassium as "Liver of sulfur" at the drugstore. Will these other compunds interfere with the reaction? What's the "worst case scenario"? Getting SO2 together with the SO3?

I'm surprised that I haven't seen this anywhere on the forum; am I missing something vital? Is potassium polysulfide virtually impossible to get ahold of in the US? Looking at the video, the reaction seems fairly straight-forward, provided you have adequate ventilation.

-Cheers

Picric-A - 20-11-2008 at 07:35

I think what you are refering to is where Mabakkens video of making SO3 using Potassium Persulphate, not polysulphide.
So that is probably the resason you havnt heard us talk about polysulphides. There are however many threads on using persulphates. UTSE!

hissingnoise - 20-11-2008 at 08:03

Welcome to SciMad, Lambda-Eyde. . .
It's a pity about the extensive carbonisation in SG's SO3; he must have used lubricant grease (or glycerin) on joints for that to happen.
Concentrated H2SO4 on joints might have been better.
I don't know if Glindemann rings would be stable at the temps used---if they are (breakdown~260*C) clean SO3 would come over.
IIRC, someone mentioned a putative electrolytic process for oleum preparation, but I can't seem now, to find it.

hissingnoise - 20-11-2008 at 09:08

Anyway, pyrosulphate NaS2O7 from heating the bi-sulphate is a less-costly route than the peroxydisulphate (full whack) one.

Lambda-Eyde - 20-11-2008 at 09:11

Oh dear, my bad. My mind must be playing tricks on me.

hissingnoise - 20-11-2008 at 09:43

It happens everyone. . .

SO3 using P2O5

Leander - 16-5-2009 at 10:33

I got hold of some P2O5, and I was kind of searching what to do with it. I allready used it to make my own rust remover. Works like a charm. :cool:

According to wikipedia:

Quote:

The desiccating power of P4O10 is strong enough to convert many mineral acids to their anhydrides. Examples: HNO3 is converted to N2O5; H2SO4 is converted to SO3; HClO4 is converted to Cl2O7.

I read the same in old chemistry books.

Is it really possible to convert H2SO4 to free oleum using P2O5? I'm really not gonna try this without more detailed information. I've seen actual flames when adding lumps of P2O5 to water.

Further, are there any lab methods known that confirm this and/or give an acceptable synthesis?

[Edited on 16-5-2009 by Leander]

BromicAcid - 16-5-2009 at 12:21

Conversion of sulfuric acid to sulfur trioxide by way of phosphorous pentoxide is indeed possible. Check out the thread on oleum production:

http://www.sciencemadness.org/talk/viewthread.php?tid=727

There are examples of using metaphosphoric acid to afford the conversion so phosphorous pentoxide should work like a charm, likely with a bit of heat to allow the separation of the volitile sulfur trioxide.

Although I have not been as active on this forum as of late so there may be more relevant recent posts that someone might be able to point you toward.

garage chemist - 16-5-2009 at 16:07

I've made SO3 from H2SO4 and P2O5 several times when I didn't have my current equipment (tube furnace) available.
Mix 75ml conc. H2SO4 with 100g P2O5 in a round-bottom flask by shaking, let it stand tightly closed for 1-2 days and then distill the mixture over a free flame (strong heat is necessary).
SO3 collects as a distillate. Cool the receiver with ice, and don't use any water in the condenser, otherwise the danger of clogging
due to crystallizing SO3 exists (still, watch out for this!).
The procedure has to be done under a fume hood, like everything involving SO3.

This method has been developed by members of the german forum, I don't know where you could find a method in literature, but let me assure you that the above method does work.

Both oleum and solid SO3 can be directly converted to chlorosulfonic acid by gassing them with dry HCl. The reaction is very exothermic and rapidly melts the solid SO3.
Then the mixture is distilled to seperate ClSO3H from any H2SO4.

Leander - 17-5-2009 at 04:09

That sounds really great!

I didn't mentioned that my goal is to get oleum for the production of several nitroaromatics such as TNT. It wouldn't be strait forward at all to distill the product, and put myself in danger trying to obtain pure SO3, simply to add it to sulphuric acid again.

I haven't got any experience myself using H3PO4 in mixed acids. According to urbanski (Vol II, p341) nitration of cellulose by mixed acids containing H3PO4 gave better results then using H2SO4/HNO3, and a Nitrogen % of 14%.

H3PO4 doesn't hydrolyses the product as H2SO4 does, giving better yields and a more stable product. The only drawback is the fact that phosphoric acid can't absorb water as H2SO4 does. Therefore only mixtures with nearly zero H2O content or even free P2O5 are useful. As a result of that P2O5/HNO3 mixtures are not really interesting IMO. Also a lack of references using this method doesn't really encourages me.

Since aromatics are probably not bothered at all by mixed acids containing small amounts of phosphoric acid, wouldn't it be possible to use H2SO4/P2O5 strait away? The only danger I could think of is the formation of nitric anhydrides, when P2O5 is used in excess, but that's fixed by using proper molar ratio's. Further, the reaction rate might slow down a little.

Any idea's? :cool:

[Edited on 17-5-2009 by Leander]

DJF90 - 17-5-2009 at 05:33

I dont think phosphoric acid is strong enough to protonate the nitric and thus allow water to leave forming the nitronium ion (the electrophile in nitration reactions). H2SO4 has a pKa ~-2, whereas H3PO4 has a pKa ~ 3 (according to paulings rules).I think nitric acid should be somewhere between the two. Therefore I would expect P2O5 and HNO3 mixtures to be poor nitrating agents (except from the increased concentration of the nitric acid due to removal of the azeotropic (~32%) water by the P2O5).

garage chemist - 17-5-2009 at 17:22

No, I don't think this will work, since the P2O5/H2SO4 mix does not seem to contain any free SO3 at room temperature- it doesn't even fume in air! I think it contains mixed anhydrides of H2SO4 and H3PO4 instead.
The SO3 is only liberated upon strongly heating the mixture- despite the 45°C boiling point of monomeric SO3, the mixture must be heated far above 200°C to liberate most of its SO3. This strongly suggests that the SO3 is produced by some sort of pyrolysis process in the mixture instead of just being distilled out.

Also, the P2O5/H2SO4 mix is very viscous. You would have a hard time nitrating anything with it.
You should really try to make pure SO3 and add this to H2SO4.

UnintentionalChaos - 17-5-2009 at 17:58

Quote: Originally posted by DJF90

I dont think phosphoric acid is strong enough to protonate the nitric and thus allow water to leave forming the nitronium ion (the electrophile in nitration reactions). H2SO4 has a pKa ~-2, whereas H3PO4 has a pKa ~ 3 (according to paulings rules).I think nitric acid should be somewhere between the two. Therefore I would expect P2O5 and HNO3 mixtures to be poor nitrating agents (except from the increased concentration of the nitric acid due to removal of the azeotropic (~32%) water by the P2O5).

IIRC, you will generate some N2O5, which behaves as nitronium nitrate. The nitrating properties will differ from normal mixed acids, I imagine.

[Edited on 5-18-09 by UnintentionalChaos]

DJF90 - 17-5-2009 at 19:10

Yes of course but only if there is excess P2O5. The [NO2][NO3] will of course be an effective nitrating agent, and if I'm doing the mechanism right in my head then HNO3 is the byproduct.

Oxidising sulphur dioxide

D S2 A - 7-1-2010 at 09:41

How can someone turn sulphur dioxide into trioxide without using high pressure vessels or huge deals of heat (>400 ºC)?

Best regards, Dúlio.

[Edited on 7-1-2010 by D S2 A]

Picric-A - 7-1-2010 at 11:05

Read the fucking thread!!!
Pass over V2O5 mixed with O2...

Ops!

D S2 A - 7-1-2010 at 11:11

Sorry. I forgot to mention that vanadium pentoxide is not available here...

Any suggestions?

Best regards, Dúlio.

Picric-A - 7-1-2010 at 11:13

Other catalyst, do some research!!!!

Hint - Fe2O3

D S2 A - 7-1-2010 at 11:15

Thank you, Picric-A . Do you have any other hint? Ferric oxide is not very efficient...

Best regards, Dúlio.

Picric-A - 7-1-2010 at 11:15

No, do some work yourself.

[Edited on 7-1-2010 by Picric-A]

D S2 A - 7-1-2010 at 11:19

I am doing that since a few hours ago. Thank you for the attention, Picric-A. In case of any good information: [ dulio@vista.aero ].

Best regards, Dúlio.

Picric-A - 7-1-2010 at 11:21

Thats obviously wrong as i did one google search now, 'oxidation of SO2' and it came up with a very good catalyst, oxidises efficiently, quickly with good yield.

D S2 A - 7-1-2010 at 11:26

Hum, I shall review it. I have notes here saying that Fe2O3 is not very good. Maybe it is a wrong notation. Thank you for advise, Picric-A.

Best regards, Dúlio.

Picric-A - 7-1-2010 at 11:28

Quote: Originally posted by D S2 A

Hum, I shall review it. I have notes here saying that Fe2O3 is not very good. Maybe it is a wrong notation. Thank you for advise, Picric-A.

Best regards, Dúlio.

Indeed Fe2O3 is a fairly bad catalyst for this but you said V2O5 is unavailable,
I said you were wrong becuase you said you have searched for an hour which is obviously wrong.

D S2 A - 7-1-2010 at 11:34

I am searching for methods to produce SO3 and produce oleum. I know the procedure with V2O5 but I want to avoid it. I shall not have unlimited resourses and equipment to produce SO3. I spent most of time on esterifications, not on oleum and related...

Best regards, Dúlio.

Picric-A - 7-1-2010 at 11:35

If you dont have the equiptment to make SO3/Oleum then you certainly dont have the equiptment to use/store it safely SO DONT!

D S2 A - 7-1-2010 at 11:40

I simply do not have equipment suitable to very high temperatures or pressurise the reaction vessel, Picric-A. The oleum shall be stored only for a day. The stumbling block for me is just related to elevated temperatures like 600 ºC.

Best regards, Dúlio.

Picric-A - 7-1-2010 at 11:42

600 can easily be reached by a bunsen burner... or any gas burner for that matter.

Learn to walk before you can run.

D S2 A - 7-1-2010 at 11:45

Even a candle goes above 1.000 ºC, Picric-A. The matter for me is how to build a reaction vessel able to resist such temperatures.

If it not bothers you, add me on MSN.

Best regards, Dúlio.

[Edited on 7-1-2010 by D S2 A]

Picric-A - 7-1-2010 at 11:49

Quote: Originally posted by D S2 A

Even a candle goes above 1.000 ºC, Picric-A.

Best regards, Dúlio.

[Edited on 7-1-2010 by D S2 A]

True.. but i would like to see you try heating half a gram of (say...) sand to 1000 degC with a candle...

Ps. I cannot go on MSN here (it is blocked by the school grr.) so if you want to contact me use U2U)

[Edited on 7-1-2010 by Picric-A]

D S2 A - 7-1-2010 at 11:51

It is just an example, Picric-A.

Best regards, Dúlio.

hissingnoise - 7-1-2010 at 13:01

This may be of interest.
http://www.google.com/patents?hl=en&lr=&vid=USPAT391...
In section 4 the patentee states that, by electrolysis, the concentration of dilute H2SO4 can be brought to 100% and above.
It is presumed that above 100% means oleum. . .
The downside is that platinum electrodes work best.

D S2 A - 8-1-2010 at 06:25

Unfortunately I do not have either equipment nor knowledge do deal with electrolysis. Any other idea?

Best regards, Dúlio.

Stoichiometric Magnesium Sulfate and Sodium Pyrosulfate

matt - 15-12-2011 at 05:23

Hi, this is my first post on this site. The process was carried out some 6 months ago, and I've since lost the photos and product.

1 mole (or proportion) of anhydrous Sodium Pyrosulfate (Na2S2O7) {prepared by thermal decomposition of Sodium Persulfate (Na2S2O8)} and 1 mole (or proportion) of anhydrous Magnesium Sulfate (MgSO4) {prepared by heating Epsom Salts (MgSO4.7H2O) at c. 250'C for several hours} were finely ground and mixed using mortar and pestle and the resulting white powder placed in a small retort. Heating over LPG burner flame resulted in a gradual rising, swelling and a change in colour to a metallic grey (similar to lead (a little whiter) or many amalgamated metals). Up to this point no fumes or condensation were observed. After effectively total conversion to this spongy-appearing, grey form, a clear liquid began to distill off and was collected in concentrated Sulfuric Acid (H2SO4). A weight gain corresponding to 0.8 mole (or proportion) of sulfur trioxide (SO3) was measured at the completion of the distillation run. When the receiver was removed from the end of the condenser, white fumes with an acrid smell developed from the clear liquid present at the end of both joints.

(The amounts involved in the run I describe correspond to an actual yield of 16g SO3 out of a theoretical 20g, I wont specify the weights of starting materials as I would be extrapolating, however I remember the yield values very clearly. The vessel was a 125ml retort [just like a retort is meant to look] with the condenser mounted in a 2 liter plastic bottle full of ice water [refreshed at about room temperature] and the burner was an LPG hotplate burner such as sold in camping stores for connection to a 9kg LPG or propane bottle.)

This is it for the moment, an attempt to convert to Thionyl Chloride failed and I lost what I had left. The system I was working with reaches a maximum internal temperature in the high 400'Cs as best i know and the late onset of distillation suggests to me the absence of any unrecognised water of crystallisation. I found this procedure to be accessible and productive to an extent beyond the others I came by in this forum. Owing to the fact that it remains merely an interpretation of the vague procedures outlined in the German patents already referenced in this thread, I felt that this was the right place to share this. I hope this is useful to somebody, even if I have misinterpreted the results.

This is only the beginning of a response to the stimulus I have gained from the pages of ScienceMadness and I will endeavour to share more of the developments which lead on specifically from ideas which arose out of this forum.

Thanks, Matt.

[Edited on 15-12-2011 by matt]

AndersHoveland - 19-3-2013 at 15:00

Sulphur trioxide reacts with nitrogen pentoxide in carbon-tetrachloride solution, with the formation of a crystalline precipitate melting at 124° to 128° C., which is probably (SO3)4.N2O5:

elementcollector1 - 19-3-2013 at 15:06

Quote: Originally posted by matt

1 mole (or proportion) of anhydrous Sodium Pyrosulfate (Na2S2O7) {prepared by thermal decomposition of Sodium Persulfate (Na2S2O8)} and 1 mole (or proportion) of anhydrous Magnesium Sulfate (MgSO4) {prepared by heating Epsom Salts (MgSO4.7H2O) at c. 250'C for several hours} were finely ground and mixed using mortar and pestle and the resulting white powder placed in a small retort. Heating over LPG burner flame resulted in a gradual rising, swelling and a change in colour to a metallic grey (similar to lead (a little whiter) or many amalgamated metals).[Edited on 15-12-2011 by matt]

A metal? There's only one metal in that mix as far as I can tell, and I don't think it'd form this easily.

madscientist - 26-3-2013 at 16:24

Quote: Originally posted by AndersHoveland

Sulphur trioxide reacts with nitrogen pentoxide in carbon-tetrachloride solution, with the formation of a crystalline precipitate melting at 124° to 128° C., which is probably (SO3)4.N2O5:

The addition of sulfur trioxide to carbon tetrachloride is one of the early phosgene preparations.

matt - 8-4-2013 at 21:13

Yes, I recently repeated this procedure, on a larger scale (1 molar), this time only getting a 50% yield although on the upside my glassware didn't crack from the heat. The metallic color, I believe, is carbonisation of impurities in the pool chemical grade reagents I used, the melt returned to white on continued heating although I'm not sure if that contradicts my idea of carbonised impurities. IIRC, the article which listed results for various alkali and alkaline earth metal pyrosulfates nominated lithium as the most easily thermally decomposed. Large amounts of Lithium Sulfate might require some effort to arrange, but hopefully the Lithium itself could be recovered from post-reaction by-products by some sort of metathesis. From my understanding of pyrosulfates, an anhydrous lithium pyrosulfate would probably be more difficult to obtain than the oleum to be produced from the combination of Na2S2O7 and Li2SO4. This is the direction I hope to take it in next as neither 50% yields nor one-time sacrificial glassware seem efficient enough to me to consider oleum as a first choice where other reagents are suitable.

S.C. Wack - 9-4-2013 at 15:26

Quote: Originally posted by matt

IIRC, the article which listed results for various alkali and alkaline earth metal pyrosulfates nominated lithium as the most easily thermally decomposed.

Just last week I looked up the abstract for the lithium pyrosulfate -> SO3 temp. article by Spitsyn IIRC, CA 20435 (1960)...the long abstract gave the impression that Gmelin's may have taken the figure out of context, not surprising considering the quoted temp...there may be measurable losses....

halogen - 28-10-2013 at 11:56

"Bismuth sulfate (Bi2(S04)3) is found in the form of white hygroscopic
crystals, which readily decompose in water to give basic subsulfate and
dissolve without decomposition in aqueous sulfuric acid. Above 465°C, it is
converted to the oxide under the liberation of sulfur trioxide."

Organic Bismuth Chemistry Hitomi Suzuki

TYPO?

Agricola - 2-11-2013 at 11:02

For a preparation of SO3 from P4O10 and H2SO4, see the attached paper by Evans. Notice Evans calls SO3 and P4O10 "anhydrous sulfuric acid" and "anhydrous phosphoric acid". Today we would call them acid anhydrides not anhydrous acids.

Attachment: sulfur_trioxide_evans1848.pdf (358kB)
This file has been downloaded 1043 times

Magpie - 27-9-2017 at 09:38

Today I made 2-3ml of oleum using the procedure of CD-ROM-LAUFWERK as presented by garage chemist. This is an extremely facile method. Phosphoric acid is first dewatered to metaphosphoric acid (HPO3) under strong heat. This then dewaters the H2SO4 to SO3.

The hot phosphoric acid severely corrodes the 100 mL rbf but garage chemist says NaOH will remove this.

[Edited on 27-9-2017 by Magpie]

[Edited on 27-9-2017 by Magpie]

[Edited on 27-9-2017 by Magpie]

hissingnoise - 27-9-2017 at 10:08

Quote:

Phosphoric acid is first dewatered to metaphophoric acid (HPO3) under strong heat. This is then dewatered to SO3.

Didn't know (HPO₃) was such a strong desiccant, tbh, but NaOH will further corrode your already damaged RBF.

Melgar - 27-9-2017 at 11:07

Quote: Originally posted by hissingnoise

Didn't know (HPO₃) was such a strong desiccant, tbh, but NaOH will further corrode your already damaged RBF.

I've been using KOH a lot lately, since I've been able to get it for about the same price as NaOH. Although it's a stronger base, it has significantly less of a corrosive effect on glass compared to NaOH. Potassium salts are also much better than sodium salts for salting out an aqueous layer. There seems to be very little that potassium salts dissolve significantly in besides water.

clearly_not_atara - 27-9-2017 at 12:37

I don't think anything can actually reverse the damage metaphosphoric acid does to glass... you can remove the phosphate, at least, with alkali. In the future consider preparing your HPO3 in stainless steel.

Pretty cool that metaphosphoric acid can make oleum, though. Chlorosulfonic acid and tosyl chloride become OTC, if only you can find a way to contain the rxn mixture. Come to think of it, it might be possible to go all the way to TsCl in situ by starting with a mixture of H2SO4 and TsOH, adding HPO3, and gassing with HCl. If produced successfully, tosyl chloride (s.g. 1.3) will float on top of the rxn mixture (s.g. ~1.8).

[Edited on 27-9-2017 by clearly_not_atara]

Melgar - 27-9-2017 at 17:21

Is tosyl chloride the same thing as para-toluenesulfonyl chloride? Because I'm pretty sure that's available on Amazon if anyone needs any.

amaming - 28-9-2017 at 15:53

the grit left on my flask after dehydrating the phosphoric acid stayed on despite me bombarding it with everything I could to try and clean it. If you want to do the meta phosphoric acid method to oleum, get yourself a 3" copper endcap or a flask you dont mind losing.

Magpie - 28-9-2017 at 18:24

I was able to remove the grit by refluxing with concentrated KOH solution for a few hours.

edit: No, this did not work. See pictures below:

[Edited on 29-9-2017 by Magpie]

Magpie - 28-9-2017 at 18:52

Today I made another batch of oleum, this time at 4X scale. I used a 600mL beaker for preparing the HPO3 and a 250mL flask for making the oleum. The yield was excellent at 29.6g vs an estimated 2-3g for the 1X scale.

I should caution that my ground glass joints leaked profusely. Therefore a good fume hood is mandatory.

2 pictures will be posted as soon as I figure out how with my new computer.

[Edited on 29-9-2017 by Magpie]

Magpie - 29-9-2017 at 17:07

The NaOH did Not remove the corrosion product from my flasks. See above picture.

softbeard - 29-9-2017 at 17:19

Quote: Originally posted by clearly_not_atara

I don't think anything can actually reverse the damage metaphosphoric acid does to glass... you can remove the phosphate, at least, with alkali. In the future consider preparing your HPO3 in stainless steel...

Yes, I agree. I ruined a large quartz test tube trying to boil phosphoric acid. Nothing can repair the damage.

Also, I tried to boil down concentrated phosphoric acid to HPO₃ in a stainless steel dish. The result was a bubbling green mass along with visible corrosion to the stainless dish.
You might be better off to try cast iron to boil & dehydrate your phosphoric acid acid accept some corrosion.

clearly_not_atara - 29-9-2017 at 17:39

My bad. Steel is attacked. I believe some varieties of stainless may be more chemically resistant than others but acid-resistant stainless may not be standard for consumer products.

I found a reference in which porcelain is used. Perhaps clay is the way to go?

http://anonym.to/http://www.science-chemistry.com/preparatio...

JJay - 30-9-2017 at 15:56

What is the appeal of using phosphoric acid to prepare oleum rather than simply cracking sodium bisulfate at 500 C or so?

Magpie - 3-10-2017 at 17:46

I think the main advantage is that you can use lower temperature. Garage chemist says heating to 680-880°C is needed for NaHSO4. I have done this using my tube furnace. With phosphoric acid/H2SO4 a bunsen burner can be used at something around the bp of H2SO4.

Incidentally my %yield at the 4X scale was approx 47%. I made another batch yesterday at scale 6X having a %yield of 50.7%.

Again, I will post some pictures if I can get this #$%& ing computer to assist me.

Attachment: 6X oleum1.odt (2MB)
This file has been downloaded 542 times

[Edited on 4-10-2017 by Magpie]

clearly_not_atara - 3-10-2017 at 18:08

Melgar: while it is certainly interesting that you can buy tosyl chloride on Amazon we have a relatively large number of questions from people who cannot simply buy things on Amazon. For them, metaphosphoric-acid -> oleum -> chlorosulfonate -> tosyl chloride is the most OTC preparation of tosyl chloride, I think.

Quote: Originally posted by JJay

What is the appeal of using phosphoric acid to prepare oleum rather than simply cracking sodium bisulfate at 500 C or so?

It's my understanding that if you combine metaphosphoric acid and concentrated sulfuric acid you can obtain a solution containing dissolved SO3 (as polysulfates) without ever handling SO3 gas. This is useful because of the considerable risks associated with accidental release of the latter. It may be possible to use oleum containing phosphoric acid in some applications of oleum, such as the decomposition of citric acid to acetonedicarboxylic acid, or similar.

I'm less clear on how much such a solution has to be heated in order to produce a stream of SO3 gas. However, I suspect it is much lower than the ~400 C decomposition temperature of Na2S2O7. Since SO3 is already hard to handle it is nice not to have to handle it at 400 C! Producing SO3 (g) at lower temperatures makes it much easier to do everything that isn't making oleum, such as the formation of Py*SO3 and other similar complexes, since these may be unstable/char at very high temperatures.

Melgar - 4-10-2017 at 09:20

Quote: Originally posted by clearly_not_atara

Melgar: while it is certainly interesting that you can buy tosyl chloride on Amazon we have a relatively large number of questions from people who cannot simply buy things on Amazon. For them, metaphosphoric-acid -> oleum -> chlorosulfonate -> tosyl chloride is the most OTC preparation of tosyl chloride, I think.

Probably. I just thought it would be helpful to point that out in case anyone needed it for anything. I've known it was available for a while, but never could come up with a reason to buy it. All the reactions that I'm aware of it being used in, also require a bunch of other stuff that's inaccessible to me.

Magpie - 4-10-2017 at 15:10

I made oleum at 10X scale today, 60g as predicted. I have about 125g of oleum now at 50% SO3. Thie should be enough to make a nice batch of thionyl chloride.

Anyone else having trouble loading pictures with Windows 10?

[Edited on 5-10-2017 by Magpie]

Attachment: 6X oleum1.odt (2MB)
This file has been downloaded 523 times

NEMO-Chemistry - 8-10-2017 at 06:27

Quote: Originally posted by Magpie

I made oleum at 10X scale today, 60g as predicted. I have about 125g of oleum now at 50% SO3. Thie should be enough to make a nice batch of thionyl chloride.

Anyone else having trouble loading pictures with Windows 10?

[Edited on 5-10-2017 by Magpie]

I upgraded to win 10, i have no end of problems including pic problems. I have no idea of the solution, in the end i simply downloaded VM ware and installed windows 8.1 on a virtual machine. I find I use that more than win 10 now.

Win 10 is said to be much better than win 8.1, but personally i just cant get on with win 10.

This is the one i use, its free. kind of defeats the point in some ways, but might help you out

https://www.virtualbox.org/

[Edited on 8-10-2017 by NEMO-Chemistry]

SWIM - 8-10-2017 at 11:02

Quote: Originally posted by Magpie

About your procedure: You mentioned some serious joint leakage.
Were the joints assembled dry, or were they wetted with concentrated H2SO4?

Magpie - 8-10-2017 at 13:46

Quote: Originally posted by NEMO-Chemistry

Quote: Originally posted by Magpie

I am having the same problem with my IOS i-pad so conclude it is related to the formum software limit of 8MGB for picture uploads. Funny I never had this problem before. I will try again after I reduce the picture bytes.

Hi swim,

I use con H2SO4 for joint sealant for Oleum. Evenually it still leaks after it starts boiling. I will try Krytox next.

[Edited on 8-10-2017 by Magpie]

SWIM - 8-10-2017 at 14:25

@Magpie

Damn.

Well, I need to rig up some kind of fume hood anyway. (Yes, I've been looking at the appropriate threads for ideas and advice)

Thanks for mentioning that leakage problem.

I'm sure it has saved some of us an unpleasant surprise.
Certainly has in my case.

Magpie - 8-10-2017 at 16:04

picture test

This works! Used Paint to reduce size of a picture that wouldn't post before resizing.

[Edited on 9-10-2017 by Magpie]

Magpie - 8-10-2017 at 16:21

Quote: Originally posted by Magpie

Here's the picture I wanted to show showing a large amount of H2SO4 mist (smoke) being formed due to uneven cooling. I would heat with a heat gun to clear the condenser of frozen SO3 then when not enough SO3 was condensing I would turn on the ice water. When I held the condenser water at 11°C this smoke was nearly eliminated.

When the system is balanced only mild heating is required with the bunsen burner. No insulation of the pot or still head is needed. Under this condition there is virtually no SO3 leakage at the joints. The smoke seen above is due to irregular heating/cooling in the condenser and is spewing out the vacuum adapter tublature.

[Edited on 9-10-2017 by Magpie]

[Edited on 9-10-2017 by Magpie]

Magpie - 10-10-2017 at 19:46

new confirm test

[Edited on 11-10-2017 by Magpie]

Magpie - 19-11-2017 at 15:03

Quote: Originally posted by Magpie

I made oleum at 10X scale today, 60g as predicted. I have about 125g of oleum now at 50% SO3.

I am in the final reagent preparation stage for making thionyl chloride and chlorosulfonic acid. As I never had an assay of the oleum I made I was skeptical that it was actually 50%. So I set out to measure its boiling point. This I determined to be 63°C which confirms a concentration of about 70%. This was very pleasing but I was skeptical so determined its mp using a salt brine coolant. This result was also very pleasing at -2°C, which confirms about 70%. I was also considering doing a titration but this is not necessary now.

aga - 19-11-2017 at 15:21

Serious stuff Magpie.

Excellent work, yet again.

From a 1673 alchemy parchment, the words/pictures deciphered as: "With the vitriolic dragon, take care to not touch it, lest it bite"

Keras - 29-9-2024 at 23:25

Making oleum by pyrolysis of sodium bisulphate.
The drops of liquid SO₃ fall directly into concentrated sulphuric acid in the RBF.

Attachment: Oleum.mp4 (5.9MB)
This file has been downloaded 63 times

Random - 4-10-2024 at 04:21

Look

Screenshot_2024-10-04_14-20-54.png - 305kB

Screenshot_2024-10-04_14-19-58.png - 295kB

Alkoholvergiftung - 4-10-2024 at 09:03

Keras. Ive read in an old journal that an reaktion of Sodiumbisulfate and Magnesiumsulfat yields SO3 it lower temperatures.
Na2SO4 + 2MgSO4 =MgNa2(SO4)2 + SO3

Keras - 5-10-2024 at 06:27

Quote: Originally posted by Random

Look

Yeah, so you'd end up with a partially or completely eaten RBF. Given the price and availability of quartz RBF, I will avoid that :p
But it can be attempted in a metal can, I suppose (I have powdered glass I could use for that). No temperature is given, though. [EDIT: Correction: 1400 °C, that way too high for my blow torches anyway. The pyrolysis of sodium bisulphate is complete at 850 °C. This is caused apparently by Na₂SO₄ being a weak base.]

[Edited on 5-10-2024 by Keras]

Keras - 5-10-2024 at 06:32

Quote: Originally posted by Alkoholvergiftung

Keras. Ive read in an old journal that an reaktion of Sodiumbisulfate and Magnesiumsulfat yields SO3 it lower temperatures.
Na2SO4 + 2MgSO4 =MgNa2(SO4)2 + SO3

Interesting, but I can buy kilotonnes of sodium bisulphate as pH- pool powder, whereas both sodium and magnesium sulphate are much more difficult to find in bulk and probably more expensive. Also, an indication of the temperature would be handy.

Jedenfalls, wenn Sie diesen Artikel wiederfinden können, wurde ich natürlich interessiert. :p

Alkoholvergiftung - 5-10-2024 at 08:08

https://dingler.bbaw.de/articles/mi233mi05_15.html#:~:text=J...

Dont know the patent he writes of,maybe there is an temperature for this reaction.
If you want lower temperatures make it like the old vitriol burners in the middleage. Iron III sulfate. 480C.

[Edited on 5-10-2024 by Alkoholvergiftung]

Alkoholvergiftung - 5-10-2024 at 23:42

Oh and other methode would be mixing 1 part Fe2O3 powder with 2 parts magnesiumsulfat powder and heat it to 600 C. The mix will Fe2O3 should work with Sodiumsulfate too. The weas product is Sodiumferrit if you drop it in water you get concentrated NaOH and Fe2O3 again.

Alkoholvergiftung - 6-10-2024 at 10:54

Keras.
https://dingler.bbaw.de/articles/mi255mi02_8.html#:~:text=Er...
That one maybe interesting for you.

Pages: 1 2