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Author: Subject: oleum & SO3
Taaie-Neuskoek
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[*] posted on 28-8-2005 at 14:36


Crap, good you pointed out the error... me in shame.
Furthermore the idea with the (discarted afterwards) metal pot was indeed the plan, heat NaHSO4 in glas till no more water comes off, then transfer to a can, and heat it up with a furnace to form SO3, and lead the latter through conc. H2SO4.

[Edited on 28-8-2005 by Taaie-Neuskoek]




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[*] posted on 28-8-2005 at 15:13


I think that you do not need to first heat the NaHSO4 in glass. This is not an oxidizer. I do not fully understand the remark about the persulfate. For this method of making SO3 no persulfate is needed, just heat NaHSO4 in a metal pot and heat until all water has gone. Next, heat much stronger in order to drive off SO3 and collect this in conc. H2SO4.

Keep your persulfate for other interesting experiments, it simply is a pity to use it for further attempts to make SO3. NaHSO4 is easy to get your hands on and it is much cheaper than persulfate.

[Edited on 28-8-2005 by woelen]




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[*] posted on 28-8-2005 at 15:34


Forget the persuphate thing, it was, let's say a massive typo... it's is getting too late for someone who is too tired. (lousy excuse...) (I've editted my post)
Anyway, would a simple furnace made of bricks, powered by burning charcoal and a fan feeding air from the downside would do for the design? On top of the thing a can with a pipe attached to it feeding to a water cooled glass cooler, and finally fed (bubbled) into conc. H2SO4. (A bit like the setup of Organikum in his benzene experiment, but then without the teflon tape on top of the can.)
When heating the NaHSO4 the first time, one can probably use the weight as an indication to see when the reaction is complete.
I think no extra air is needed to protect the tubing, it's a semi-consumable anyway, and air will also mean water in the climate I live in, and drying air is another extra pain.
Will copper piping work? How does hot SO3 behave in the presence of metals? Is it really strong oxidising everything into crap, or does it behave nicely...? Maybe there is just one way to find out...

[Edited on 28-8-2005 by Taaie-Neuskoek]




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[*] posted on 28-8-2005 at 19:11
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I don't think that either glass or metal could be recommended for this, certainly not glass that you don't want stressed. The pyrosulfate method came from a series of German patents, I've never heard of this method actually being used. Even though the iron sulfate method, allegedly used to 1900 to make Nordhausen oleum, requires higher temps and there are losses due to SO2 formation. espacenet doesn't have Wallace's DE2285 but does have the others. Of course they are not as informative as anyone would like. Just thought that I'd say that you might not want to use pyrosulfate alone. Apparently "M" was "R" back in the day.

DE3110 uses MgSO4 to lower the temp, by formation of a double salt:
Das Verfahren zur Darstellung des Schwefelsäureanhydrits beruht auf der Einwirkung des wasserfreien schwefelsauren Magnesiums (SO4Mg) auf wasserfreies saures schwefelsaures Natrium (SO4Na2 + SO3) bei einer Temperatur, welche noch erheblich unter der Dunkelrothglut liegt, bei welchem Vorgange die sogen. Doppelverbindung der beiden Metalle sich bildet und das Schwefelsäureanhydrit frei wird.

Statt der Natrium- kann die Kaliumverbindung und statt der Magnesiumverbindung auch die der übrigen sogen. Vitriole (SO4R + 7OH2) und des Calciums angewendet werden.

Als Vorfabrikation wird wasserfreies saures schwefelsaures Natrium, wie bekannt, durch Erhitzen von Glaubersalz mit Schwefelsäure und wasserfreies schwefelsaures Magnesium durch Erhitzen von Bittersalz dargestellt.

Die fernere Arbeit besteht in der Vereinigung der genannten wasserfreien Verbindungen unter einer Temperatur, bei welcher die Alkaliverbindung eben flüssig ist, und darauf folgender etwas stärkerer Erhitzung der Masse, wobei das Anhydrit frei wird. Die resultirende Verbindung des schwefelsauren Natriums mit schwefelsaurem Magnesium wird nach bekanntem Verfahren auf wässerigem Wege getrennt, um immer wieder zur Darstellung des Anhydrits zu dienen.

Der Werth des ganzen Verfahrens für die Praxis liegt wesentlich darin, dafs das Freiwerden des Schwefelsäureanhydrits bei solcher verhältnifsmäfsig niedrigen Temperatur vor sich geht, dafs dazu Gefäfse und Apparate von allen in Betracht kommenden Materialien angewendet werden können, ohne dafs erhebliche Abnutzung stattfindet, und aufserdem in der Höhe der Ausbeute, welche bei guter Ausführung bis zur sogenannten theoretischen gesteigert werden kann...

It goes on to describe the illustration of the apparatus.

part 2, DE6091:
Die nach der Destillation von Natriumbisulfat mit Magnesiumsulfat verbleibende Verbindung der neutralen Salze wird durch Mühlen zerkleinert, dann das Pulver mit Schwefelsäure vermengt, durch Erhitzen das Wasser entfernt, und darauf durch stärkeres Erhitzen das Anhydrid abdestillirt. Es ist bei dieser Form des Verfahrens zweckmäfsig, nicht 1 Aequivalent Säure auf 1 Aequivalent Salze, sondern 1 Aequivalent Säure auf 2 Aequivalente Salze anzuwenden, da sonst die Hälfte der Säure als wasserhaltige fortgeht.

Nach dem beobachteten Verlauf der Reaction ist anzunehmen, dafs sich zu Anfang der Erhitzung eine Verbindung von 1 Molecul Nätriumsulfat mit 2 Moleculen Magnesiumsulfat bildet, Na2SO4 + 2 MgSO4, und dafs durch diese Verbindung erst bei stärkerem Erhitzen das noch vorhandene saure schwefelsaure Natrium zersetzt wird und unter Entweichen von Anhydrid das Doppelsalz Na2Mg(SO4)2 entsteht.

Eigenthümlich ist bei dieser Form der Anwendung der patentirten Hauptreaction die Benutzung des Doppelsalzes ohne vorherige Trennung desselben durch Krystallisation.

DE12295, the method that everyone already knows, by distillation from the pyrosulfate and H2SO4:
Wasserfreies saures schwefelsaures Alkali (R2S207) wird mit Schwefelsäurehydrat versetzt und hieraus, nach eintretender, theilweiser Umsetzung in saures schwefelsaures Alkali (RHSO4) und freies Anhydrit, letzteres abdestillirt.

R2S2O7 + H2SO4 = 2 RHSO4 + SO3.

Das zurückbleibende saure Salz wird durch Erhitzen wieder in wasserfreies saures Salz (Pyrosulfat) übergeführt und dient durch Wiederholung des Processes aufs neue zur Darstellung von Anhydrit.
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[*] posted on 29-8-2005 at 04:00


A-ha! MgSO4 does the trick!
They don't heat Na2S2O7 alone, for whatever reason, most likely too high temperatures needed.

Na2S2O7 can give off nearly quantitative amounts of SO3 but ONLY when either anhydrous MgSO4 or concentrated H2SO4 is added.

This is interesting now. It will be one of my next experiments.
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[*] posted on 29-8-2005 at 06:27


Thanks a lot, S.C. Wack, that sort of info is very, very usefull. As I still have the RBF with 100g persulphate/pyrosulphate in, I might test for persulphate activity, heat it more till no more activity if necessary, and to drive the water out, toss in 50mls of H2SO4, heat and distill of the SO3. Looks very nice, and the other method, with MgSO4 and Na2SO4 might be for some of us even better, one only need to have cheap precursors and sulphuric acid! Thanks again, mr S.C. Wack!!

[Edited on 29-8-2005 by Taaie-Neuskoek]




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[*] posted on 29-8-2005 at 06:41


Thats great news to hear. I as well will try with a little experimentation and come back with results. If one adds sulfuric acid with the pyrosulfate does it change into oleum when the SO3 is released? I know that oleum boils at a much lower temp than sulfuric acid so I'm guessing the SO3 would come right back out.
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[*] posted on 31-8-2005 at 05:10


50g NaHSO4 (0,4mol) ("pH minus" for swimming pool) were heated in a two-neck rbf until evolution of fumes ceased. At first, quite some gas was produced, most likely water vapor. Then it ceased and some fumes with resemblance to SO3 appeared. Heating was stopped at this point. The rbf was stoppered to exclude air moisture and left to cool down (it was very hot and it took long).
11ml H2SO4 (0,2mol) were added and the flask set up for distillation.
It was heated until the pyrosulfate melted. It mixed completely with the H2SO4 on swirling.
Then the heat was turned on high and after prolonged heating, the liquid started to boil. However, nothing was coming over, NOTHING, not even a drop, and no fumes crept through the condenser. The H2SO4 was just refluxing in the rbf because the high radiative heat losses didn't permit the distillation of the H2SO4 with my rather small burner.
The residue didn't fume in air.

This experiment was a complete failure.

I need to get some Bittersalz from the garden store and try the magnesium sulfate/sodium pyrosulfate process.
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[*] posted on 31-8-2005 at 08:22


What was the temperature of the bioling sulphuric acid? If the acid was just boiling, it should give off white fumes of small droplets of H2SO4. For how long did you let the boiling continue?



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[*] posted on 31-8-2005 at 09:05


Just to be sure, I suggest melting the bisulfate, weighing, and then weighing again when you think that you have pyrosulfate. It is easy to over/underheat the bisulfate, as I've said before.

I've read that quantitative conversion to pyrosulfate is not possible, by heating bisulfate.
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[*] posted on 1-9-2005 at 14:33


I did let the boiling continue for quite some time, about 10 minutes. The H2SO4/pyrosulfate mix was extremely hot, white fumes were at the top of the flask but they were just gaseous conc. H2SO4 and condensed on the walls, so I was basically refluxing it. The fumes didn't make it into the condenser, no matter how strong I heated.

After the mix had cooled down, I let it stand for some more time. After re- heating, some fumes got into the condenser and at quite a lower temperature, but the amount was too small for formation of visible drops of SO3 in the condenser.

I'm letting it stand a few days. Maybe the reaction needs time?

It would be very kind if someone who has got both MgSO4 and NaHSO4 would try out the other, more promising process and tell us about it!
I'd hate to buy a 5kg bag of MgSO4 (no smaller units available in the garden shop) and then find out it's useless!

[Edited on 1-9-2005 by garage chemist]
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[*] posted on 1-9-2005 at 16:29


Is there any more information available on that process? I can't find anything online or make sense of Babelfish's translation of that patent.

edit:

I just tried a test tube experiment with this. One gram each of NaHSO<sub>4</sub> and MgSO<sub>4</sub>.7H<sub>2</sub>O were heated in a test tube with a blow torch. At first, water vapor came off with a little SO<sub>2</sub> (probably from impurities in the bisulfate). This quickly changed to choking SOx vapors (mainly SO<sub>3</sub>, I would guess) and a white mist at a temperature well below red heat, probably several hundred C. Eventually, several drops of some clear liquid condensed on the cool part of the test tube (~room temp) and were collected in another test tube. I should note that drops of the liquid which condensed occasionally flowed back into the heated part of the tube, without any boiling. It would appear that SO<sub>3</sub> formed and reacted with the water that had already condensed in the cool part of the tube.

The collected liquid was heated on a water bath, but no bubbling was observed. A drop of water was then added, which immediately (and violently) boiled off. Another milliliter was added, forming a highly acidic solution. This solution was very concentrated, forming clear lines when poured into water.

[Edited on 2-9-2005 by neutrino]
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[*] posted on 4-9-2005 at 04:40


when i read the HPO3-thrade i had an idea:
what if HPO3 can dehydrate H2SO4?
i tested it in a duran test tube:
cooked down some H3PO4 until it was a half-crystaline very viscous mass, the glas startet to glow a bit dark red (but it wasnt moldable)
than let it cool down (the glas dont crack)
add some H2SO4 and re-heat carfuly, after some sec. white smoke startet to come out of the test tube
and to be sure it was SO3 i heat pure H2SO4 in a test tube until it boil but it taked some time until something comes out (coz the high bp. of 310°C, much condensation on the walls) and than the smoke wasnt that thick, so im sure it was SO3 :)
this is one of the perfekt methods to produce SO3 from H2SO4, u just cook the H3PO4 down, add H2SO4, collect the SO3 and so on :P
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[*] posted on 4-9-2005 at 10:07


Ladies and Gentlemen! An important breakthrough in the realm of amateur chemistry has been made right in the last few minutes!

I am very pleased to tell you that the highly useful chemical OLEUM in every desired SO3 concentration can now be mass-produced in every garage/basement laboratory. The apparatus is a simple distillation setup with ground- glass joints.
Credits go to CD-ROM-LAUFWERK for having made the discovery that Metaphosphoric Acid is able to dehydrate conc. Sulfuric Acid to Sulphur Trioxide in high yields.


My last experiment was the following:

13,5ml (0,2mol) 85% Phosphoric Acid were heated in a 100ml beaker until boiling. When the boiling subsided, the heat was turned on maximum until the bottom part of the beaker was glowing faintly red. This was continued for 5 minutes.
The resulting liquid (HPO3)n was allowed to cool covered for a short time, but only to the point where it was still pourable (use gloves, it's really hot!).
It was poured into a dried 100ml round- bottom flask and quickly stoppered.

The beaker was attacked somewhat by the hot (HPO3)n and had a frosted appearance after the (HPO3)n residues had been washed out. However, the frosted appearance could be removed by boiling some NaOH solution in this beaker.
An iron crucible is useless, it gets dissolved rapidly.


To the (HPO3)n in the flask was added 4ml of conc. H2SO4 (a bit less than 0,1mol).
Then the flask was fitted with a dried NS 14,5/23 distillation bridge WITHOUT running cooling water through the condenser (otherwise the SO3 will solidify in there and clog it). The condenser had a length of 160mm. The receiver was immersed in ice water in order for the SO3 to condense there and not in the condenser.
The mixture was heated with a bunsen burner.
After some heating and swirling, the (HPO3)n mixed completely with the H2SO4.
Then the heating was put on maximum.
The liquid very soon started boiling and a colorless liquid began distilling at about 40- 60°C steam temperature. A lot of heat is needed in order to effect the complete reaction between (HPO3)n and H2SO4. About 1,5-2ml collected in the receiver.

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.

This liquid is definately oleum of a very high concentration, if not pure sulphur trioxide.
Redistilling the liquid will yield pure SO3 in liquid form.

The residue from the distillation can again be turned into (HPO3)n by heating to red heat. It can be reused indefinately for dehydrating H2SO4 to SO3.

EDIT: the text in bold.

[Edited on 4-9-2005 by garage chemist]
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[*] posted on 4-9-2005 at 10:46


Excellent work Garage Chemist! This is a very usefull procedure!
As an iron crucible is required, would this one also be needed for the last step, or would it be severely attacted by the hot H2SO4 and SO3?

How damaged was the RBF after the destillation of the SO3?




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[*] posted on 4-9-2005 at 11:00


I used a rbf which was already damaged, as I knew that it would be attacked somewhat. But the damage is limited because you don't have to heat to red heat like in the production of the (HPO3)n.
Also, the (HPO3)n has already pulled the water from the H2SO4 and turned into the much less aggressive H3PO4 when it is at the temperature at which all of the SO3 gets expelled.

You can therefore use a glass apparatus for the generation of the SO3.

If you heat too high, the H3PO4 will eliminate water vapor which will drip into the liquid SO3 in the receiver and cause it to EXPLODE (yes explode, dripping water into SO3 has the same effect as throwing a chunk of cesium into water).

[Edited on 4-9-2005 by garage chemist]
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[*] posted on 4-9-2005 at 13:46


Did I say "iron resists"? That's nonsense as I found out today.

Iron is resistant to cold 85% phosphoric acid, but it rapidly reacts once hot.

The most resistant vessel for production of (HPO3)n so far is a quartz crucible, the attack was only slight.

Anyways, let's continue the talk about preparation of (HPO3)n in the thread with this name.

This thread shall be reserved for the reaction of H2SO4 with (HPO3)n and especially for reports about successful preparation of oleum/SO3.

I want to know if you are able to reproduce my success!
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[*] posted on 4-9-2005 at 14:55


Excellent stuff, although it sounds fairly dangerous and all. But that's nothing to stop anyone :D Still I'd have preferred some dry method, whereby SO3 would be fed into H2SO4.

Anyway.

I was wondering, does this reaction work with H4P2O7 as well, or only the HPO3 form? Because then you could simply dehydrate H3PO4 to H4P2O7, add the correct (stoichiometric) amount of H2SO4 accounting for ~4% water, heat up to ~200 deg C, collect the distillate (H2S2O7, oleum), then heat to 250 deg C in the very same container, collect the water (while H4P2O7 reforms), add some more H2SO4 and repeat.

So the question is, at what temp does the oleum come over?

Also, if H4P2O7 is workable, then the mixing between the two acids should be better as both are liquids.




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[*] posted on 4-9-2005 at 15:20


I wasn't able to tell when the H3PO4 was H4P2O7, the evolution of water vapor just gradually slackened and then ceased entirely.

However, in the Organikum, a method for polyphosphoric acid is given: heat only to 150°C, but in a vacuum and for 6 hours.
This also doesn't attack the glass.
However, this polyphosphoric acid is crystalline.

My (HPO3)n was liquid (but extremely viscous).

[Edited on 4-9-2005 by garage chemist]
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[*] posted on 5-9-2005 at 12:24


Now I tried it again with a less agressively dehydrated version of phosphoric acid.
It was heated to 350°C until no more bubbles formed. This did NOT attack the quartz crucible at all.
The liquid was also much less viscous and could be easily poured. At room temp, it was thick, but less viscous than e.g. honey, it was still pourable.
(In contrast to this, my HPO3 from the previous batch was not pourable at room temperature)

It worked, but yield was terrible (2,5g SO3 instead of the expected 18g). I used 0,5mol H3PO4 and 0,2mol H2SO4. I also had to heat very high, until the H2SO4 was refluxing, and then the 2,5g SO3 very slowly crept out.

So one needs to use real metaphosphoric acid, pyrophosphoric and polyphosphoric acid don't work.
You have to heat the HPO3 to red heat for some time in order for acceptable SO3 yields to be possible.

[Edited on 5-9-2005 by garage chemist]
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[*] posted on 5-9-2005 at 16:37


Thats a good info, GC, thanks for the experiment. I guess one has to put a lot of effort into the HPO3 before it can be used for SO3 production. Thats great that this method can be effectivily used too. Do you think larger production of SO3 is possible using this method? I mean like 20 - 30 mL and more?
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[*] posted on 6-9-2005 at 02:45


Large batches of SO3 are definately possible, however you will need a large quartz crucible (200ml) and large amounts of phosphoric acid.

Another approach would be to use diammonium phosphate (available as fertilizer, I haven't seen it in garden stores but you could order it online, it surely won't raise any eyebrows because of its inert nature). On heating, this also eventually gives metaphosphoric acid, but ammonia is given off instead of water.
The ammonia smell could be a good indicator of when the reaction has finished, though!
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[*] posted on 6-9-2005 at 11:13


what about heating in a vakuum up to 350°C?
would this form HPO3?
i dont think that glas is atacked at this temperaturs, always when i destilled SO3 out of a P2O5-H2SO4 mixture the temperature is >310°C (heat untill the H2SO4 reflux) and the glas is not atacked, and that after a few times :P
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[*] posted on 13-9-2005 at 18:18


I'm sure there are many dehydrating agents that can be used for the procedure. The only issue with this method for most people is the availablity of most dehydrating agents. P2O5 and most other phosphorus containing dehydrating agents like PCl3 are generally very hard to get if i'm not mistaken. Are there any other dehydrating agents in addition to HPO3 that are fairly easy to come by that could used for the job?
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[*] posted on 18-9-2005 at 10:59


Sorry for the double post but I was doing some thinking. What happens when you react calcium carbide with sulfuric acid? Calcium sulfate, carbon, and hydrogen gas? If not, I was thinking one could reflux H2SO4 with Ca2C and the H2O - SO3 equalibrium would be pulled toward the SO3 as the H2O reacted with the Ca2C to form acetylene.
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