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Author: Subject: oleum & SO3
chemoleo
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[*] posted on 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.




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[*] posted on 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 ?:P

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]
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[*] posted on 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.




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[*] posted on 20-8-2005 at 15:56


I have read that sodium persulfate decomposes in alcohol. What does it decompose into?
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[*] posted on 21-8-2005 at 00:22


To acetic acid and sodium hydrogen sulphate (NaHSO4).



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[*] posted on 23-8-2005 at 17:05


Interesting, thank you for the information, I would assume ammonium persulfate follows the same reaction giving NH4HSO4.
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[*] posted on 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.
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[*] posted on 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.



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[*] posted on 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.




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[*] posted on 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
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[*] posted on 25-8-2005 at 17:57
SO3


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?
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[*] posted on 25-8-2005 at 18:12


How cold was the acid? SO<sub>3</sub> melts at 16.8*C.
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[*] posted on 25-8-2005 at 18:16


dont really remember it was cold outside.......late winter early spring i think
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[*] posted on 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]




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[*] posted on 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]




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[*] posted on 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
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[*] posted on 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?




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[*] posted on 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]




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[*] posted on 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....




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[*] posted on 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.




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[*] posted on 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.
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[*] posted on 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.




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[*] posted on 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]




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[*] posted on 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?
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[*] posted on 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.




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