Sciencemadness Discussion Board - oleum & SO3

oleum & SO3

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Organikum - 25-5-2003 at 14:35

Here we go
HARDCORE:
Oleum, smoking H2SO4, 50%SO3 in H2SO4, SO3

First I suggest everyone who wants to tinker with this may inform himself on the truely existing dangers of these compounds. They are not nasty - they are dangerous!

Madscientist passed H2SO4 over MgSO4 for oleum (did this really work?) he wrote and in more than one synth SO3 was asked for (acetic anhydride for example, nitric acid).
So here is a doable way to produce oleum and SO3:
- Ferrous sulfate septahydrate is made from iron (filings/steelwool) and H2SO4.
- Ferrous sulfate septahydrate dehydrates completely by heating >300°C.
- Anhydrous Ferrous sulfate decomposes to SO3 at 480°C and can be condensed to oleum (or the SO3 can be directly used of course).

If the oleum is what one is after it should be possible to decompose and condense in the same vessel without extern apparati necessary. The vessel has to be big and strong enough - thats understood.

madscientist - 25-5-2003 at 16:31

Organikum: I think my oleum attempt from long ago didn't actually work - I probably just thought it did (considering my relative ignorance back then, and the fact that I made way too many assumptions).

Organikum - 25-5-2003 at 18:26

xoo1246: Yes me too! The howto on the ferreous was only included for completeness and the sure to come question on it. Sooner or later.......

Madscientist: No problem for me. My complete errors and selfcaused mishappenings will be published soon. Five volumes. Open end. Working title: "Oh! Oh! Oh! - Oh shit!" Probably the "shit" has to be removed so.......
....oh, oh, oh.

I am a fish - 25-5-2003 at 23:31

Quote:

Originally posted by Organikum
- Ferrous sulfate septahydrate is made from iron (filings/steelwool) and H2SO4.
- Ferrous sulfate septahydrate dehydrates completely by heating >300°C.
- Anhydrous Ferrous sulfate decomposes to SO3 at 480°C and can be condensed to oleum (or the SO3 can be directly used of course).

Heating the ferrous sulphate must be done in an oxygen free atmosphere, otherwise it will oxidise into the ferric state.

Another good source of ferrous sulphate is 'Iron Mordant', whch is available from all good dyeing suppliers.

You are right, fish you are!

Organikum - 26-5-2003 at 00:07

The setup has to be closed against the outside what should be understood for the unhealthy properties of SO3 alone. For the ease of availability of the reactands I would say that it´s not necessary to purge with an inert gas, but it´s easier to sacrifice a part of the ferrous sulfate to the oxygen inside. There should be enough left to be converted to SO3 - the efficiency of these rather old and raw methods isn´t very good anyways I believe.

Exactly for such critics it got posted. Thanks fish you are!

[Edited on 26-5-2003 by Organikum]

Microtek - 26-5-2003 at 01:00

Industrially, the SO3 is absorbed in circulating concentrated H2SO4 which is then diluted with water. This is done to attenuate the violence of reacting SO3 directly with water.

vulture - 26-5-2003 at 04:53

Why not use CuSO4? That can't be oxidized any further.

higher temperatures?

Organikum - 26-5-2003 at 06:55

CuSO4 starts to decomposes at a temperatures of 560°C whereby FeSO4 is completely decomposed at 480°C. Thats how I understood it at least. In the original process costs will have been a factor also - I didn´t invent this, I just don´t know where this was taken from, if I knew it, I would have named it.
CuSO4 maybe worth a try, also I can´t see any advantages. as the FeSO4 scrubbing the oxygen seems to me favorable.

The humidity in the reactor should be driven out during heating the FeSO4 for becoming anhydrous what should be done in the same vessel anyways. I propose that vessel and FeSO4 will be in the same state of anhydrous condition so.

(Ah! The beauty of strong causal logic!)

btw. this is a concept for discussion, not a complete running process by now. So the formulation: "Why don´t you use..." is inadequate here.

a_bab - 26-5-2003 at 09:20

What about heating H2SO4 ? The white fumes are mainly because of SO3 IIRC.

rikkitikkitavi - 26-5-2003 at 09:53

no, a_bab, the white fumes of heated (boiling ) sulfuric acid is just sulfuric acid droplets. If H2SO4 is heated , it will split to SO3+ H2O at high temperatures and the SO3 will decompose into SO2 and O2 at these temperatures.

Better than FeSO4 is Fe2(SO4)3 since +3 ions decomposes into SO3+SO2 at lower tempertures. Expect a large loss of SO3 due to eq. SO3 = SO2 + 1/2 SO2

/rickard

ziqquratu - 27-5-2003 at 16:16

I think that heating Sodim bisulfate may well work... I'm not exactly sure of the temperatures, but I do know they're accessable to a gas flame. If I remember, you dry the bisulfate in an oven, then heat it over a gas flame, and water gets produced. You keep the heat so that water is just getting produced, and you're getting solid sodium pyrosulfate, until water stops forming. Then crank up the heat to decompose the sodium pyrosulfate and get sodium sulfate and sulfur trioxide. then the best thing to do is dissolve the gas in cooled, concentrated H2SO4. It can work if you use a distillation setup - change flasks between the water production and the decomposition of the solid. Although I dont think you'd want to be too attached to the flask that you're heating - it may suffer some permanent damage if you're not careful. I dont actually have a reference for this, mind you - I just remember seeing it somewhere some time ago and thinking "hmm... must try this one day"... never got around to it... must remedy that. Oh, by the way, Sodium Bisulfate is commonly used for raising the pH of swimming pools... very cheap and easy to find

a_bab - 28-5-2003 at 08:59

Yes rikkitikkitavi, you are right. I was confused because of the lab method wich is using oleum to get SO3.

You are right aswell ziqquratu, a simple method could be heating NaHSO4 or KHSO4. The reaction will be in two stages, one at 200-300 degrees C, when the alkali pyrosulfate is formed, and then the other at 500 degrees C when SO3 is formed because of the thermal decomposition of the pyrosulfate.
The reactions are:

1. 2NaHSO4 --->Na2S2O7+H2O
2. Na2S2O7---->Na2SO4+SO3

[Edited on 28-5-2003 by a_bab]

vulture - 28-5-2003 at 10:46

NaHSO4 for raising the pH?

HSO4- has a pKa of 1,92 which means it will behave like an acid in solution. Especially in diluted solution.

Oops!!

ziqquratu - 28-5-2003 at 15:58

Sorry, I didnt see that one!! Sodium bisulfate makes pH lower!! I'd have simply edited my above post, but then anyone reading what vulture wrote would have no idea what was going on!! Thanks for pointing that one out

a short resumee

Organikum - 29-5-2003 at 00:09

- by principle the process will work and includes no dangers like sudden decomposition (explosion) which would render it useless for the amateur experimenter.
- precautions: drying the used apparatus which has to be a enclosed one.
Also the apparatus must be VERY acid resistant. Ceramics, porcellain, clay, not glass or metal (to stay on the safe side)
- Fumehood or outside with fan at least.
- Other starting compounds than FeSO4 seem to be possible.

Please correct me if I am wrong
thanks
ORG

Next step would be a layout for the apparatus - I thought on a modified KLIPP principle for keeping it small.

[Edited on 29-5-2003 by Organikum]

KABOOOM(pyrojustforfun) - 1-9-2003 at 19:55

destructive distillation of green vitriol(FeSO4.7H2O) is a very traditional way of making H2SO4 and was first done by hayyan (don't know its English spelling)
Organikum! are you sure the 480°C is the decomosition point for ferrous sulfate? it is the exact temp for ferric too.

is ferric hydrogen sulfate, anhydrous or hydrate? if it is anhydrous or can be dehydrated at a lower temp than its decompositin point, it's possible to make oleom without need of conc sulfuric acid! u simply add dilute H2SO4 to ferric sulfate/oxide/hydroxide.. recrystalize the salt and distill off.
2Fe(HSO4

3 => Fe2O3 + 3H2S2O7
it may lose H2O forming pyrosulfate first, (but my sence doesn't approve it.)
the pyrosulfate method is great for those who distill their own nitric acid. just dilute the wast acid => recrysalize the alkali metal hydrogen sulfate.

[Edited on 2-9-2003 by KABOOOM(pyrojustforfun)]

BromicAcid - 23-5-2004 at 14:07

We don't have a thread started for the reactions of SO3 but I figured this thread would do since I don't feel it necessary to start a new one.

"Sulfur trioxide reacts with KClO3 to form K2S3O10 and a red product assumed to be (ClO2)2S3O10 (mp 75C)" The red product has a violent decomposition at 100C. An anologous reaction occurs with KClO4.

On topic it appears that the decomposition temperature of ferric sulfate is significantly higher then 480C. Here are two passages from the Complete Tretise on Inorganic Chemistry that are relevent:

"The yield of fuming acid is poor because the temp. of dissociation of ferric sulfate is so high that a large proportion of of sulfur trioxide is decomposed at the same time."

And relating to that: "The dissociation pressure of the trioxide from ferric sulfate was found by G. Keppeler and J. d'Ans to be only 15mm. at 640C so that the temp of the retorts must be much higher then this."

Other interesting information relating to things said upthread:

"P.G. Prelier proposed to heat the alkali or alkaline earth sulfates with sulfuric acid to form the hydrosulfate, and to distill the water from the hydrosulfate, and the fuming acid or sulfur trioxide from the pyrosulfate. R.W. Wallace proposed a modification of this process; adn W. Wolters said that the liberation of sulfur trioxide from the pyrosulfate occurs at a lower temp if some magnesium sulfate is present. This means that less sulfur trioxide is dissociated. W. Wolters also proposed to heat the pyrosulfate with anhydrous sulfuric acid so as to distill off the sulfur trioxide:
N2S2O7 + H2SO4 ---> 2NaHSO4 + SO3
A.M. Leon proposed to electrolyze the monohydrated sulfuric acid so that the water decomposes into hydrogen and oxygen, and the resulting sulfur trioxide dissolved in the electrolyte to form the fuming acid."

As for the last part though, I doubt graphite would survive and I would have to guess the electrodes would have to be made of something like Pt.

S.C. Wack - 23-5-2004 at 19:24

Wolters' patent with MgSO4: DE6091. With H2SO4: DE12295. Very uninformative. I think Wallace had a temp of 600C going.

GB237243 mentions a temp of 750-780 for the ferric sulfate deal.

GB 124988 is about sawdust and bisulfate.

In US342784 and 342785, clay and gypsum give cement and SO3. Like most thermal methods, expect at least half of the S as SO2. I think I saw this in Kirk-Othmer somewhere, but that doesn't mean it works! And US2528103 bluntly says so, but doesn't even get Cummings' name right. Anyways, it is supposed to be an improvement. I've seen both of Cummings' pats on espacenet before, but (no surprise) I was just there and they say that they don't exist. Of course uspto.gov says that they do.

Proteios - 23-5-2004 at 23:14

oks....my contribution.....

1) high temperature control without some reasonable kit is a bitch.
2) the temperatures you are talking are getting close to the softening point of pyrex (600-700 C).....

All my attempts to decomp. FeSO4 CuSO4 as a teeny failed. I only used glass.

Now with much more real experience:
I think your best bet is a metal kit. Air and oxidation of the FeII i think is a limited worry, given that at the temperature you are talking the SO3 = SO2 n O2 equilibrium will estabilsh fairly quickly. Plus you have to get the O2 into the matrix of the solid (rates)
How to best do this is really dependant upon all sort of unknown rates, but my guess is you will be best off with a straight tube, Copper i think is good and easily available. Use an air carrier, modest airflow (fish tank pump). This will get the SO3 away more quickly from the hot surface where decomposition/equilibration will occur. SO3 is a pretty good gas at even modest temperatures, consdense on cold glassware. How quickly SO3 eats the tube is also a worry, but also another good reason for using a carrier gas. It will dilute the SO3, it will cool the SO3 and it will move the SO3 away from the metal surface more quickly.
-there are a lot of guesses, assumptions and opinion here!!!!!!!

I also FAILED MANY times to make SO3 from SO2 and air, using a V 2O5 catalyst, with only the most primative of supports. Temperature control was a real bitch (propane gas torch), as was getting SO2 dry enough.
I only once succeded at making SO3, and that was at school (everything else was done at home), using an SO2 cylinder and a platinumised wool catalyst. Using relatively low heat (bearly off sooty flame on a bunsen) over a period of maybe 1 hr, I made maybe 2-5g of SO3 as a crytaline solid that readily melted with bodyheat.

In summary, probably a bent Cu tube, FeSO4 in bottom, air carrirer, roast to red heat with gas torch, have faith, that even though you cant see whats happening, it will work!
alternatively, maybe a pretty determined piece of kit for the SO2 oxidation!

Na2S2O8

chemoleo - 26-5-2004 at 19:10

A long time ago I generated SO3 from sodium persulphate, which I obtained from conrad.de, and which I coincidentially wanted to use for copper etching.
I just heated the persulphate in a testtube, and melted very quickly. Upon heating a bit further, VAST amounts of SO3 were produced, which quickly filled my room with a white fog. That lasted for quite a while, and surprisingly it wasn't particularly harmful. After a while, no more SO3 formed, and I was left with a clear molten salt... which is Na2S2O5? or possibly something different, with concommitant O2 evolution.

Realising then that this must be SO3, I fed this with some plastic tubing into H2O, but to my dismay it did not seemingly dissolve, and lowered the pH only a little. Nothing as impressive as bubbling Cl2 through NaOH.

Unfortunately, I never bubbled the SO3 through H2SO4 conc. - I wish I had done it, because it seems exceedingly easy to do, plus Na2S2O8 is serioulsy easy to get! Maybe someone else try it for me?

Theoretic - 27-5-2004 at 04:34

You probably got Na2SO4, Na2SO5 isn't stable at high temperatures.
I think you probably produced SO2 with a little bit of SO3 to create a fog and O2. The SO3 present dissolved and lowered the pH.
Anyway, that's my theory.

chemoleo - 27-5-2004 at 11:43

No way. Ever smelled SO2? You can detect it quite easily. And there was definitely no SO2 smell around (unless repeated ethanolic exposure screwed around my memories from more than a decade ago

).
Just try it out, this is an experiment that can be literally done in 10 minutes.

axehandle - 27-5-2004 at 11:53

My comments: When I tried superheating FeSO4 in an iron pipe I left it unattended, so it overheated and got yellow-hot. The entire flat was filled with a thin white fog, but it had almost no smell -- it was just irritating to the lungs (smelt _slightly_ like SO2, but only a little, almost undetectable). I think it was a mix of SO2 and SO3, mostly SO3.

I think chemoleo is right.

jimwig - 2-8-2004 at 14:48

it is my understanding that SO3 in air readily forms H2SO4 especially when inhaled and is quite irritating and painful to your lungs.

if you had any SO3 at all you should have been coughing and in pain.

NaHSO4 to Na2S2O7 by heat

axehandle - 2-8-2004 at 18:17

I'm right now keeping 1/2 kg of NaHSO4 in my electrical furnace at 400 degrees C in an attempt to convert it to Na2S2O7. I'll tell you how it goes...

On a more bizarre note, I'm going to take up pottery in about 10 minutes... going to make a couple of clay retorts. I'll post pictures, regardless of the result...

[Edited on 2004-8-3 by axehandle]

Clay

MadHatter - 2-8-2004 at 21:33

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

axehandle - 3-8-2004 at 19:11

Quote:

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

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

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

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

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

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

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

[Edited on 2004-8-4 by axehandle]

[Edited on 2004-8-5 by axehandle]

TheBear - 31-8-2004 at 06:43

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

garage chemist - 12-9-2004 at 04:32

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

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

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

The SO3 is produced by mixing H2SO4 and P2O5 and then distilling. The SO3 condenses without problems if you cool the receiving flask in ice water. That's the easiest way to produce SO3.

neutrino - 12-9-2004 at 17:10

Does anyone have any experience with the pyrosulfate + acid method?

axehandle - 12-9-2004 at 18:05

Quote:

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

Yes, I just forgot to write about it... I've got about 3kg worth of Na2S2O7 in my dessicator (box with CaCl2...). The comparison batch of about 500g which I kept in an open jar in the kitchen for several weeks is completely covered with water it has absorbed from the air. The lump in the dessicator is still completely dry.

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

I have yet to try a controlled thermal decomposition yielding SO2. I need some glassware which I can't afford to order right now to do it in a repeatable, conclusive manner.

[Edited on 2004-9-13 by axehandle]

ADP - 1-6-2005 at 10:30

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

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

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

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

neutrino - 1-6-2005 at 13:01

I wouldn't use a retort for this procedure for the simple fact that it is open to the air. While the SO3 would condense, a lot of it would still evaporate, forming a dangerous acid mist. Water vapor would also come into the receiver and form sulfuric acid.

chemoleo - 1-6-2005 at 13:30

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

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

ADP - 1-6-2005 at 15:08

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

However do you know of any in-town places that might supply it? I was planning on buying 500g online since I haven't seen it anywhere else.

ADP - 1-6-2005 at 15:16

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

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

Also perhaps I could avoid the hazards of the acid mist by placing the end of the retort directly in conc. sulfuric acid yielding oleum.

chloric1 - 2-6-2005 at 03:46

Yes this would be especially interesting since the MSDS shows the persulfate melts with decompostion at 180 degrees C.

THe lower temperature and the abundance of evolved oxygen should prevent sulfur dioxide from coming over. Sounds promising actually

[Edited on 6/2/2005 by chloric1]

ADP - 2-6-2005 at 09:58

Wow this is really getting me excited, One could easily use this method to recharge H2SO4 or convert to oleum. If pretty pure SO3 could be isolated in liquid, one could have a superior nitration but adding the SO3 to 70% HNO3 would easily make a 99%HNO3 and H2SO4 mixture. Can anyone say 95%+ yield...

ADP - 3-6-2005 at 20:41

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

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

I'll report back when I get the stuff in. I will order some next week.

BromicAcid - 3-6-2005 at 20:47

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

ADP - 4-6-2005 at 08:29

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

http://www.jtbaker.com/msds/englishhtml/p5983.htm

Cloner - 8-6-2005 at 09:53

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

I put a PH paper in the fumes and it was quite acidic when I stopped, but the product did not react with water violenly so I might just have had only a little bit of pyrosulphate.

chemoleo - 8-6-2005 at 10:27

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

S.C. Wack - 8-6-2005 at 11:16

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

It just so happens that a tumbler glass on a small cast-iron skillet on a gas stove provides a good temperature. Good temperature control is very important here, the range that this works within is small.

Taaie-Neuskoek - 9-6-2005 at 14:03

Quote:

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

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

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

ADP - 17-8-2005 at 12:49

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

http://en.wikipedia.org/wiki/Sulfur_trioxide

PCB etching agent, as SO3 source

Lambda - 17-8-2005 at 13:40

Quote:

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

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

Na2S2O8 --> Na2SO5 (which decomposes further) + SO3

PCB etching agent, as SO3 source:

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

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

[Edited on 18-8-2005 by Lambda]

chemoleo - 17-8-2005 at 14:12

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

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

Anyway, pls try it with the Na2S2O8.

ADP - 17-8-2005 at 14:22

Does anyone know the decompostion temp of sodium bisulfate? I know it's melting point is 58dC. Also after the decomposition to pyrosulfate, what temp is required to make the pyrosulfate decompose?

Taaie-Neuskoek - 17-8-2005 at 16:18

Quote:

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

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

Anyway, pls try it with the Na2S2O8.

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

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

chemoleo - 17-8-2005 at 16:25

You can still get Na2S2O8 from www.conrad.de. I actually thought that's better for the environment than FeCl3.

As to the latter part of your post - you realise you are essentially doing what is done in the
lead chamber process?
Batchwise synth might work better, but then I suppose you won't get SO3 from it to make oleum. Hmm.

Lambda - 17-8-2005 at 16:46

Quote:

Originally posted by Taaie-Neuskoek
If I get my Na2S2O8, I will.
Sodium persulphate is hardly used in etching (sp??) here (NL) because of all the eco-tax on it. .... I have been able to get my hands on a few 100g via a friend, but it is certainly not OTC anymore, everybody uses Fe2Cl3 here....

Taaie-Neuskoek, howmany tons of this etchant would you like me to arrange for you ?

It's desirable properties, are very much appreciated by many, and for this reason, it is still OTC all the way. The high price, being the only drawback, compared to the cheap Iron(III)chloride.

[Edited on 18-8-2005 by Lambda]

Taaie-Neuskoek - 17-8-2005 at 17:18

Quote:

As to the latter part of your post - you realise you are essentially doing what is done in the lead chamber process

Yes, actually, the whole idea of making SO3 that way was extracted from that thread.
I am not really interested in making sulfuric acid as I have more than enough of the stuff, but SO3 or at least oleum was the goal. The problem with this method is that your product will be contaminated with NO2, which has to be distilled of or something.
A batchwise setup will indeed probably be the best, but in a large container the minimal of liquid will be hard to get out I fear.
I do get you point, the gasses need time to react, and they do that well when they spend a little time with each other...
BTW, can SO3 maybe be made by adding acetic anhydride to H2SO4? SO3 to glacial acetic acid yields acetic anhydride IIRC.

ADP - 20-8-2005 at 15:56

I have read that sodium persulfate decomposes in alcohol. What does it decompose into?

Nicodem - 21-8-2005 at 00:22

To acetic acid and sodium hydrogen sulphate (NaHSO4).

ADP - 23-8-2005 at 17:05

Interesting, thank you for the information, I would assume ammonium persulfate follows the same reaction giving NH4HSO4.

ADP - 23-8-2005 at 19:10

Sorry for the double post but I was curious about a method for making SO3 involving NaHSO4. If heated to anhydrous and partially decomposed by heat. It produces H2O and Sodium pyrosulfate (Not sure of the nomenclature). The pyrosulfate, if heated then produces SO3 and NaSO4.

Does anyone have any information on this method other than the sulfur trioxide article on wikipedia? That is where I got my information BTW. Sodium bisulfate is much easier to come by that persulfate. Also I am curious of decomposition temperatures.

BromicAcid - 23-8-2005 at 19:13

That method is covered somewhat extensively earlier in this thread, axehandle even made an attempt at it with some interesting results and there is some other relevent information here, please read the thread.

Taaie-Neuskoek - 24-8-2005 at 15:10

I honorated Chemoleo's request, and repeated the experimetn with Na2S2O8.

11,662g was heated till it was mostly molten, not everything reacted in my opinion, but the erlenmeyer was very, very hot, and the reaction didn't continue very much more... It was very different from the reaction with the ammoniumsalf, that one was almost self-sustaining.
Anyways, re-weighing yielded 10,795g, meaning that around 0,867 (+/- 0,005g) has reacted to SO3, which is a rather poor yield, as the reaction mechanism Na2S2O8 --> Na2SO5 + SO3 would yield in a 100% conversion something like 3,76g.
It is too late now (1:00AM) to do the experiment again, I'll try to reproduce this, or to get better yield.
Keep in mind though that this probably ís reproducable within a few days, but that sodiumpersulphate is unstable, and will decompose if no stabiliser is added.
Ammoniumpersulphate is stable.

ADP - 25-8-2005 at 17:49

I got aquired some sodium bisulfate today and did a test tube experiment to see how a torch effected it. I first melted it evenly until all was a liquid and heated gradually more until it was boiling rapidly. I heated even more and then white smoke started puffing out. I waved my hand over it to move it my way to smell and it was a choking smell that made me cough like NO2 without the alkaline smell. It must have been SO3. My idea next is to dry the bisulfate to pyrosulfate and dry the pyrosulfate.

When I aquire some pyrosulfate I'll add about 5 grams to a test tube and a small amount of H2SO4 to cover it. Perhaps this would act as a good intermediant for the SO3.

Also I'd like to add in the beginning I tested about 5 grams of NaHSO4

SO3

gilbert pinkston - 25-8-2005 at 17:57

years ago i was under the mistaken idea oleum was needed to add the last NO2 to toluene so i tryed to make it in many ways one way was the iron sulphate way....after being sure i got the water of crystalization out i strongly heated it with white sand and lead the vapors into the strongest (cold) HNO3 (1.52) the SO3 was caught and SO2 was converted to SO3 with the release of nitrogen oxides after a while i noticed white solids in the acid some was collected and put in water where it fizzed and spun around and dissolved solid sulphur trioxide?

neutrino - 25-8-2005 at 18:12

How cold was the acid? SO3 melts at 16.8*C.

gilbert pinkston - 25-8-2005 at 18:16

dont really remember it was cold outside.......late winter early spring i think

Taaie-Neuskoek - 26-8-2005 at 12:24

Ok, I tried again with Na2S2O8, but then with daylight.
The result is more or less the same, starting was 8,725g, after it all went liquid 8,048g was left.
I also tried (NH4)2S2O8 again, starting with 10,740g, after heating 9,902g was left.
These results are almost the same as I got before, so at least it is reproducable.

I also added about 100g of (Na)persulphate to a 1L RBF, this flask was subsequently heated, and the liberated SO3 was bubbled through 30ml of conc H2SO4 in a 100ml graduated cylinder. The damp wasn't really absorbed or something, nor did the H2SO4 heat up. I could see some mist coming out of the cylinder, so had not really the strong impression something was absorbed.
After 10-15 min of bubbling I stopped, and took the H2SO4 and poured it into an erlenmeyer, but it wasn't smoking at all...
The H2SO4 was of lab-quality, no drain opener or something.
Should the contact suface be a lot more larger, should the H2SO4 be heated, or cooled, or whatever went wrong?? There wás a lot of SO3 present there...

[Edited on 26-8-2005 by Taaie-Neuskoek]

woelen - 26-8-2005 at 14:38

How do you know there is a lot of SO3? You wrote there is some mist. From 100 grams of Na2S2O8 I would expect a lot of thick fume, not some mist.

You heated the Na2S2O8 and more than 90% of the initial weight remains. Does the remaining compound still contain Na2S2O8? You can check out by adding a spatula full of this to a precipitate of Ni(OH)2. If the precipitate turns black at once, then still there is a lot of S2O8(2-) in the remaining stuff. If the Ni(OH)2 remains light green, then probably the Na2S2O8 just decomposes, giving oxygen and Na2S2O7.

Quote:

Ok, I tried again with Na2S2O8, but then with daylight.
The result is more or less the same, starting was 8,725g, after it all went liquid 8,048g was left.

Based on this observation, I only can conclude that you have the following reaction almost quantitatively:

2Na2S2O8 ---> 2Na2S2O7 + O2.

Perform the computation yourself and you'll see that your results are very close to the theoretical loss of mass. You loose just a few percent more and only that very small amount probably is released as SO3, the remaining part I expect to be O2.

That can explain, why in your last batch of 100 grams the bubbles are not absorbed by the H2SO4 and why the H2SO4 is not fuming after this treatment. Most likely you were just bubbling oxygen through H2SO4

Quote:

11,662g was heated till it was mostly molten, not everything reacted in my opinion, but the erlenmeyer was very, very hot, and the reaction didn't continue very much more... It was very different from the reaction with the ammoniumsalf, that one was almost self-sustaining.
Anyways, re-weighing yielded 10,795g, meaning that around 0,867 (+/- 0,005g) has reacted to SO3, which is a rather poor yield, as the reaction mechanism Na2S2O8 --> Na2SO5 + SO3 would yield in a 100% conversion something like 3,76g.

In this one you probably made 0.784 grams of oxygen and at most 80 mg or so of SO3. Probably you made even less SO3, because you may have 'lost' some material in your reweighing action.

[Edited on 26-8-2005 by woelen]

CD-ROM-LAUFWERK - 26-8-2005 at 14:53

if u only melt the S2O8 salts, there comes at first a bit white smoke, but that ISNT everything!
after all stuff is molten u must heat it strongly to gain the SO3, and than much is coming out
pleasy try again Taaie-Neuskoek and now whit full heat power untill no more white fumes coming out
i dont see a real advantage of using (NH4)2S2O8 instead Na2S2O8...
the decompositon temp. is everytime high ...like pyrosulfate...
the way i use to make SO3 is just the dehydration of H2SO4 whit P2O5... its the fastest and easyest way because u can use a normal destillation-apperatur and dont need a super-heat-resistant apperatur!

PS: the way whit P2O5 is my invention, not garage chemist's, i just posted it in a forum

PPS: the bit white stuff when u melt the persulfate is coming on Na2S2O8 and on
(NH4)2S2O8

woelen - 26-8-2005 at 14:58

From your reply I understand that the SO3 actually is not from the initial reaction with decomposition to S2O7(2-), but at a much higher temperature, where S2O7(2-) decomposes to SO4(2-) and SO3?

Then, what is the advantage of using Na2S2O8 over using NaHSO4? NaHSO4 first decomposes giving water and then at a much higher temp. you can get SO3?

Taaie-Neuskoek - 26-8-2005 at 15:05

Quote:

You wrote there is some mist. From 100 grams of Na2S2O8 I would expect a lot of thick fume, not some mist.

Woelen, thanks for your reaction, however, the SO3(?) was released over 10-15 minutes, and there was a constant reaction going on in the flask, clearly emitting cloudy stuff. It doesn't very fast, as it does with ammoniumpersulphate, so there was nothing like a huge cloud, as it was released over time.

What would be the cloudy stuff then? IIRC in an MSDS was written that during decomposition of xpersulphate irritating/corroding vapours were emitted, and no SO2 smell is noticable, as chemoleo also wrote in his post on this topic. However, when I did the first experiment, it was indoors, and the vapour didn't dissapear, but floated in the room around, untill it was so dilute that it couldn't be seen anymore.
I'll try the decompostion again in a testtube and a glowing fling or any O2 is formed when I have a bit more time. I don't have any Li(OH)2, only the sulphate... I'll try to find other ways of detecting persulphate, and also measure the density of the sulphuric acid solution, as far as I can do that in a good way.

EDIT: to CD-ROM LAUFWERK: I don't have any P2O5, nor do I have a cheap source for it.
I'll try to heat it more strongly, fortunatly I didn't chuck away the clump of salf which is still in the RBF...
However, I assume that I am getting already close to the softening point of pyrex at those temperatures... what makes things a lot more complicated, as I am carefull with my glassware, my budget is not endless, if I may use an understatement.
Can I get to those temps with a normal propane (campinggaz) torch?

Quote:

i dont see a real advantage of using (NH4)2S2O8 instead Na2S2O8...

There is, the the sodiumsalt is unstable and will deteriorate after some time, the ammoniumsalt however, is stable.

[Edited on 26-8-2005 by Taaie-Neuskoek]

woelen - 26-8-2005 at 15:33

If you have NiSO4.xH2O, then dissolve some of this in water, add a solution of NaOH and you get a green suspension. Add some persulfate to this and it will turn black at once. This is a VERY sensitive reaction for persulfate.

Try to heat some Na2S2O8 in a test tube and test for O2 as you already suggested yourself. Keep on heating for 15 minutes, as you did in the other experiments.

After the heating dissolve the molten and solidified residue in a small amount of water and add part of it to a suspension of nickel hydroxide. Also keep some and check for acidity. If it is very acid, then you also have an indication for Na2S2O7.

The small amount of white fume you get most likely indeed is some SO3, but only a small amount. It is remarkable how a small amount of just a few tens of mg can give a lot of fume with many chemicals, so I'm not surprised if you see quite some fumes, even if you just get a small amount of SO3.

I'm looking forward at your results tomorrow. Now the light is dimmed at Woelen's place

, it is past 1:30 now....

chemoleo - 26-8-2005 at 21:08

I should say, when I did those experiments a long time ago, I melted the Na2S2O8 until it melted, directly under a bunsen flame, in a test tube, for an extended period of time. The room was filled with white fog, which I found non-irritant.
Maybe it's a temp issue indeed. Damn i wish I could try it myself.

ADP - 27-8-2005 at 10:39

Well if it is a temperature thing, I'm curious at what temperature Na2S2O7 will decompose. I've made some today via bisulfate and have a distillation retort but if the temperature required is too much then I must find some alternate route.

Taaie-Neuskoek - 28-8-2005 at 05:20

Ok, tried again, 1,300g Na2S2O8 was heated for a long time, and a stead flow of mist came out of the small erlenmeyer. This was contuniued, untill the steam almost dissapeared, and the salt had gone from molten till solid state again.
This took at least 10min, maybe more, (didn't clock) but it took a looong tine before everything was converted.
Reweighted there was 1,007g left, which means 0,293g has dissapeared, or a 56% conversion from NaS2O8 via Na2S2O7 to Na2SO4.
After cooling down a little water was added, which was still acidic. Woelen's test with Ni(OH)2 showed clearly that no perchlorate was present. (pos. control showed a very clear colourchange.)
This good and bad news, it does probably work, but it involves very high temperatures, my propane torch could handle a gram, maybe 10gram, but no more. Also using glass in not very nice, as the temps involved may damage your glass very well. It is also probably just a waste of persulphate to use it for this purpose, as at those temperature FeSO4 is much, much cheaper... pity, it looked so nice and easy.

chemoleo - 28-8-2005 at 05:56

Well it is nice and easy. From what I know the vitriol (FeSO4) decomposition isn't exactly easy to do, it needs very high temps, contains water (thus u cant get oleum from this, which is the idea o this thread - water-free SO3!) etc.
Anyway I am glad you were able to verify this.

What I wonder - do you get Na2SO4 or Na2SO5? Both exist, and for the latter I don't know the decomposition temperature. This would change your yields of course. But then, is the Na2SO5 active with the Ni test? It likely is. So yes, we have evolution of O2 as well.
How come however you get a 56% yield and yet a negative test? Conversion to pyrosulfate, which doesn't break down any further?

[Edited on 28-8-2005 by chemoleo]

ADP - 28-8-2005 at 06:13

I didn't think that pyrosulfate decompostition temp was that high. If it is higher than around 550dc then one might as well use ferric sulfate (Fe2(SO4)3). Upon heating to 480dC it decomposes into Iron (III) Oxide and SO3.

FeSO4 is too high as well is it not?

woelen - 28-8-2005 at 07:18

Quote:

What I wonder - do you get Na2SO4 or Na2SO5? Both exist, and for the latter I don't know the decomposition temperature. This would change your yields of course. But then, is the Na2SO5 active with the Ni test? It likely is. So yes, we have evolution of O2 as well.

Forget about Na2SO5. This stuff is extremely unstable and you'll never get that. In fact, persulfate itself already is quite unstable, especially the sodium salt (look at JT Baker's MSDS).

So, based on Taaie-Neuskoek's observations, the only conclusion, which can be drawn here is that Na2S2O8 first decomposes to Na2S2O7 and that this happens even at moderate temperature. The decomposition of Na2S2O7 to Na2SO4 and SO3 requires a much higher temperature, such that it can only be done with difficulty and quite some risk in glass apparatus.

Quote:

How come however you get a 56% yield and yet a negative test? Conversion to pyrosulfate, which doesn't break down any further?

This of course is quite well possible, although I think Taaie-Neuskoek made some computational error on the yield.

All Na2S2O8 was decomposed to Na2S2O7 (otherwise he would have a positive test for persulfate) and from the latter, part is decomposed to Na2SO4 and SO3. Taaie-Neuskoek started with 1.300 grams. From this he had 87 mg of oxygen gas. He was left with 1.213 grams of Na2S2O7. From this, theoretically he could have 437 mg of SO3, instead, he had 1.007 grams of solid remaining, so he had 206 mg of SO3, which is a yield of 47% from the theoretical maximum.

Altogether, I think that Na2S2O8 is not the best starting point for SO3. It indeed is fairly expensive and not easy to find for everybody. I myself have a nice, but not very cheap, source for this stuff, so I followed all this with great interest, but now I think it is time to go for another synthetical means for SO3. If you want to make SO3 from a cheap source, then NaHSO4 probably is better. This also decomposes to Na2S2O7, but this stuff can be obtained at low price (over here it is available for appr. EUR 20 per 3 kilos as pH-minus for swimming pools). A good procedure may be to heat a lot of NaHSO4, first letting the water boil away and then heat it at much higher temperature in a metal pot. Probably the metal pot will be spoiled after this excercise.

Taaie-Neuskoek - 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]

woelen - 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]

Taaie-Neuskoek - 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]

Wolters

S.C. Wack - 28-8-2005 at 19:11

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.

garage chemist - 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.

Taaie-Neuskoek - 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]

ADP - 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.

garage chemist - 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.

Taaie-Neuskoek - 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?

S.C. Wack - 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.

garage chemist - 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]

neutrino - 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 NaHSO4 and MgSO4.7H2O were heated in a test tube with a blow torch. At first, water vapor came off with a little SO2 (probably from impurities in the bisulfate). This quickly changed to choking SOx vapors (mainly SO3, 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 SO3 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]

CD-ROM-LAUFWERK - 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

garage chemist - 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]

Taaie-Neuskoek - 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?

garage chemist - 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]

garage chemist - 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!

chemoleo - 4-9-2005 at 14:55

Excellent stuff, although it sounds fairly dangerous and all. But that's nothing to stop anyone

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.

garage chemist - 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]

garage chemist - 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]

ADP - 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?

garage chemist - 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!

CD-ROM-LAUFWERK - 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

ADP - 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?

ADP - 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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