Sciencemadness Discussion Board - Confusion on SO3 Synthesis

Confusion on SO3 Synthesis

SaccharinSlayer157 - 12-11-2022 at 22:29

Just a quick question about the importance of cations in certain reactions, hopefully I'm posting this in the right place...

So I recently attempted the SO3 outlined in this video. I cleaned my glassware really well to avoid the contamination he encountered, but I used sodium persulfate instead of the potassium salt because I found it for cheaper online. After running it, however, I found that I couldn't collect any product at all. Even after repeating it with different amounts of sulfuric acid and strong prolonged heating no SO3 ever came over for me to collect. In general I've always thought that alkali metal cations for salts can be pretty interchangeable for most reactions, but I'm wondering if this reaction might be an exception where sodium persulfate doesn't decompose nearly as readily as its potassium counterpart for some reason. I know I definitely formed some SO3 because I saw its characteristic fumes once the reaction flask was opened to the air, but it dissipated very quickly. I just want to know if this could be the reason for my failure so I need to obtain some of the potassium salt or if I'm overlooking something in my setup instead that I need to investigate further. I'm working with pretty small amounts and it's expensive and difficult to order any chemicals at the moment so I'd like to rule this out as a possibility before I devote any more time and money to it

metalresearcher - 13-11-2022 at 07:58

The largest online chemistry textbook tells that K2S2O8 decomposes below 100C while Na2S2O8 does not, it melts 'normally' at 180 C.

woelen - 13-11-2022 at 12:13

@metalresearcher: That does not need to make a big difference.
Persulfate decomposes in two stages. The first stage is loss of oxygen:
2 K2S2O8 ---heat---> 2 K2S2O7 + O2 (apparently 100 C is sufficient for this reaction to occur)

The same also happens for the sodium salt, but at a higher temperature.

The second stage is loss of SO3:
K2S2O7 ---->lots of heat---> K2SO4 + SO3

I do not know what heat is needed to get SO3, but it must be considerable. I can imagine that for Na2S2O7 the temperature, needed for making SO3 is higher than for K2S2O7, but I am quite sure that for K2S2O7 also quite a high temperature is needed. Do not expect this reaction to occur easily.

Admagistr - 13-11-2022 at 14:31

Quote: Originally posted by woelen

@metalresearcher: That does not need to make a big difference.
Persulfate decomposes in two stages. The first stage is loss of oxygen:
2 K2S2O8 ---heat---> 2 K2S2O7 + O2 (apparently 100 C is sufficient for this reaction to occur)

The same also happens for the sodium salt, but at a higher temperature.

The second stage is loss of SO3:
K2S2O7 ---->lots of heat---> K2SO4 + SO3

I do not know what heat is needed to get SO3, but it must be considerable. I can imagine that for Na2S2O7 the temperature, needed for making SO3 is higher than for K2S2O7, but I am quite sure that for K2S2O7 also quite a high temperature is needed. Do not expect this reaction to occur easily.

Yes Wilco,but SO3 at temperatures above 500C and at 1000C completely transform to SO2 and O2.I melted KAl(SO4)2 at 1000C for 6 hours and observed this.

SaccharinSlayer157 - 13-11-2022 at 21:29

I guess I'll just have to try again with the potassium salt and see whether that or my setup was the issue. Thanks for the insight, I'll post an update once I've run it again!

teodor - 14-11-2022 at 05:12

SaccharinSlayer157, trying to produce SO3 from a small amount of salt make sure everything is perfectly dry.

1g of Na2S2O7 gives a max of 0.3g of SO3 (and you never can get so a high yield) but 0.07 g of H2O is enough to spoil the result because it will bind with SO3 to form H2SO4.

Also, wet SO3 has a tendency to polymerization, so it could be hard to distill it in this form. If I would make such an experiment I would mix the salt with a tiny amount of P2O5 or red P to bind all free water (but I never did this experiment, so accept that recommendation with a grain of salt). Other aggressive drying agents should also work. I think they also will depress the reaction temperature.

[Edited on 14-11-2022 by teodor]

SaccharinSlayer157 - 2-12-2022 at 10:35

So I just finished my second much more careful run, cleaned and dried my glassware and the salt meticulously and even redistilled my sulfuric acid just in case. I got the mix molten and up to at least 500-600C for over an hour and still got nothing, barely even any oxygen evolution. It’s looking to me like the sodium salt is just too stubborn to decompose appreciably, so I’ve ordered some potassium persulfate and will try with that once it arrives.

clearly_not_atara - 4-12-2022 at 11:35

I've often wondered if there was a way to directly acidify pyrosulfate. Unfortunately, disulfuric acid is extremely strong (comparable to chlorosulfonic acid, pKa < -10) for the first deprotonation. I haven't found any reliable data for the second deprotonation. The only candidate I can think of is phosphotungstic acid H3PW12O40, but this would require a lot of tungsten, and solvent selection is non-obvious. HBr and HI are strong enough, but would be oxidized; HClO4 is strong enough, but could explode. HCl is significantly weaker than sulfuric acid.

The order of strength of extremely strong acids is solvent-dependent, so "aqueous" pKa values are not necessarily reliable; FSO3H is stronger than HClO4 in sulfuric acid but weaker in acetic acid.

I strongly suspect that KHS2O7 decomposes more readily than K2S2O7. So one protonation should be enough.

MaeBorowski - 24-12-2022 at 13:21

HCl is stronger than H2SO4 but here we are. Disulfuric acid isn't stable at elevated temperature and will decompose. There is very common method which describes the preparation of SO3 from Na2S2O7 and sulfuric acid but without much details

Bedlasky - 24-12-2022 at 22:50

Quote: Originally posted by clearly_not_atara

HClO4 is strong enough, but could explode.

HClO4 is weaker acid than H2S2O7. H₀(HClO₄) = -13, while H₀("H₂S₂O₇") = -14,44. I refer to "H2S2O7" as 50 mole % SO3 in H2SO4.

SaccharinSlayer157 - 25-1-2023 at 15:04

This is not going as planned...

There is no way this is how I'm going to get my SO3, I seriously hope someone can help by suggesting an easier or more reliable method.

Even with a short path setup and anhydrous EVERYTHING the potassium salt still takes ages to melt and decompose under very strong heating and my yield is reliably abysmal. So far all I've managed to do is convert less than 10ml of my initial sulfuric acid catalyst to what I assume to be a very dilute oleum over the course of several frustrating runs.

I've heard something about a bisulfate approach? Apparently that proceeds through the same pyrosulfate though so I'll probably encounter the same issues. Any other suggestions before I just cave and try to order P2O5 from firefox?

I ultimately need it to make oluem and then subsequently chlorofulfonic acid and I'm ready to move forward so any insight is greatly appreciated!

Deathunter88 - 26-1-2023 at 00:09

Quote: Originally posted by SaccharinSlayer157

This is not going as planned...

There is no way this is how I'm going to get my SO3, I seriously hope someone can help by suggesting an easier or more reliable method.

Even with a short path setup and anhydrous EVERYTHING the potassium salt still takes ages to melt and decompose under very strong heating and my yield is reliably abysmal. So far all I've managed to do is convert less than 10ml of my initial sulfuric acid catalyst to what I assume to be a very dilute oleum over the course of several frustrating runs.

I've heard something about a bisulfate approach? Apparently that proceeds through the same pyrosulfate though so I'll probably encounter the same issues. Any other suggestions before I just cave and try to order P2O5 from firefox?

I ultimately need it to make oluem and then subsequently chlorofulfonic acid and I'm ready to move forward so any insight is greatly appreciated!

I would just go with P2O5. The bad news is that firefox is basically dead. The good news is that P2O5 is now easily available on Ebay.

Mateo_swe - 26-1-2023 at 06:21

Have you read Len´s book "Small-scale Synthesis Of Laboratory Reagents"?
Chapter 20 - Sulfur Trioxide and Oleum

Good book with preparations and good explanation and pics.
Do a online search for it, its a good book.

Keras - 26-1-2023 at 07:41

Quote: Originally posted by Mateo_swe

Have you read Len´s book "Small-scale Synthesis Of Laboratory Reagents"?
Chapter 20 - Sulfur Trioxide and Oleum

Good book with preparations and good explanation and pics.
Do a online search for it, its a good book.

I’m planning to try the SO₃ synthesis mentioned in this book later this year. I have bought a suitable quartz round bottom flask, since standard borosilicate glass would not survive the 700 °C+ temperature needed to decompose sodium bisulphate. I just to need to buy one or two more blow torches to reach the proper temperature.

yobbo II - 26-1-2023 at 11:03

The following is from the Bell Jar, the first five years page 3 - 8:
A magazine dedicated to vacuum.

Low temperature chemical reactions:

The reaction
NaHCO3 = Na2CO3 + H2O + CO2
which normally takes place at 270° C will take place at
room temperature in a vacuum of about 1 mm Hg.

Other reactions which take place at much lower
temperatures in a vacuum are the following:
Cr2 O3 + 3C = 2Cr + 3CO
CaC2 + 2 NaCl = CaCl2 + 2Na (vapor) + 2C
2MgO + CaO + Si + 2Mg (vapor) + CaSiO3
FeO + C = Fe + CO
These reactions are of great commercial importance.
Any reaction evolving gas will proceed at lower
temperature in vacuum.

END OF QUOTE
___________________________________

Would this be applicable to making SO3?

You would need a water aspirator as you are not going to use your mechanical pump, though you could use a piston (the type used nowadays) pump from a fridge which I belived will pull a vacuum of about 1mm of Hg.

Yob

Keras - 27-1-2023 at 00:30

Quote: Originally posted by yobbo II

The following is from the Bell Jar, the first five years page 3 - 8:
A magazine dedicated to vacuum.

Any reaction evolving gas will proceed at lower
temperature in vacuum.

It's curious. The equation ∆G = ∆H - T∆S which relates the equilibrium constant to the temperature does not really explicitly depend on pressure. One could argue that removing the gaseous products help drive the reaction to completion, which is certainly true. Maybe that explains it. Maybe also S, and in a lesser extent H, depend on P. Maybe operating under reduced pressure increases ∆S, which automatically means that a lower T is needed to reach the same equilibrium.

As to the synthesis of SO₃, I’m not certain it would be helpful. You would end up with sulphur trioxide being sucked into your vacuum pump, which could lead to… much inconvenience.

[Edited on 27-1-2023 by Keras]

unionised - 27-1-2023 at 00:38

Quote: Originally posted by Keras

It's curious. The equation ∆G = ∆H - T∆S which relates the equilibrium constant to the temperature does not really explicitly depend on pressure.
[Edited on 27-1-2023 by Keras]

It implicitly depends on it. the entropy of a material- particularly a gas- is pressure dependent.

Keras - 27-1-2023 at 05:08

Quote: Originally posted by unionised

It implicitly depends on it. the entropy of a material- particularly a gas- is pressure dependent.

Does the entropy increases or decreases with pressure?

Junk_Enginerd - 5-2-2023 at 06:12

Quote: Originally posted by SaccharinSlayer157

This is not going as planned...

There is no way this is how I'm going to get my SO3, I seriously hope someone can help by suggesting an easier or more reliable method. !

Sure. I can vouch for sodium bisulfate. Easy to get, standard pool chemical. Then heat the crap out of it and SO3 comes out! By my estimate I heated it to 650°C, just beyond incandescence. It makes the glassware single use for sure, but that's what test tubes are for.

SaccharinSlayer157 - 10-2-2023 at 01:19

Hallelujah!

So I decided to bite the bullet and bought 500g of P2O5 off ebay.

Very much worth not having to deal with high temps and abysmal yields IMO (Plus I still have a bunch left over to play with later

). After reading a bit more and finding THIS thread with a helpful Engager post I found immediate success...

50ml 98% sulfuric acid with 100g phosphorous pentoxide at medium heat distilled directly into more H2SO4 yielded what I estimate to be ~50% free SO3 oleum in less than 45 min, no problem. Next stop, HSO3Cl!

Only "issue" was that if I didn't keep my condenser water warm enough (It's pretty cold in North Texas atm), the SO3 would crystallize in the receiving flask and plug it until I could warm the incoming distillate back up again. Got some lovely formations though I'm sure y'all will appreciate...

Also, a syrupy almost taffy like goop was left in the receiving flask. I assume this is mainly polyphosphoric acid, leftover sulfuric, and some other mixed phosphoric acids? I saved it and am keeping it dry for now. Anyone have any ideas on how this might be useful or is it really just waste?

Sulaiman - 10-2-2023 at 02:03

I've never seen SO₃ before, gas, liquid or solid.
Thanks for the photos.

The leftover in the pot would, I guess, be a very good general purpose dessicant.

BromicAcid - 10-2-2023 at 04:59

Boy, it can be nerve wracking when stuff starts freezing on the drip tip like that. I usually keep a heat gun handy for low melters but sometimes it's hard to melt out the tip with the void space. I mean, you obviously never want to get it blanked off and potentially build pressure in the setup. Nice crystals though.

Edit: I should mention that it's usually even more of a pain for me since the setup is usually under vacuum and thus the heat from the outside glass portion has an even more difficult (next to impossible time) getting that inner drip tip melted out.

[Edited on 2/10/2023 by BromicAcid]

Herr Haber - 10-2-2023 at 15:37

Quote: Originally posted by SaccharinSlayer157

Hallelujah!

Congratulations, beautiful pictures !!!

Quote: Originally posted by SaccharinSlayer157

Only "issue" was that if I didn't keep my condenser water warm enough (It's pretty cold in North Texas atm), the SO3 would crystallize in the receiving flask and plug it until I could warm the incoming distillate back up again. Got some lovely formations though I'm sure y'all will appreciate...

I'm sure you can find a heat GUN in Texas !
But these are so beautiful I'm happy you didnt have one at that moment

ErgoloidMesylate - 11-5-2023 at 15:18

Quote: Originally posted by SaccharinSlayer157

Wiki mentions using a tin salt at a lower temp.

Another two step method involving a salt pyrolysis starts with concentrated sulfuric acid and anhydrous tin tetrachloride:

Reaction between tin tetrachloride and sulfuric acid in a 1:2 molar mixture at near reflux (114°C):
SnCl4 + 2 H2SO4 → Sn(SO4)2 + 4 HCl
Pyrolysis of anhydrous tin(IV) sulfate at 150°C - 200°C:
Sn(SO4)2 → SnO2 + 2 SO3
The advantage of this method over the sodium bisulfate one is that it requires much lower temperatures and can be done using normal borosilicate laboratory glassware without the risk of shattering.

Raid - 25-5-2023 at 10:33

Is there a way of recovering the P2O5 after the oleum synthesis? Maybe like drying it at a very high temp or something?

Loptr - 25-5-2023 at 17:38

Quote: Originally posted by Raid

Is there a way of recovering the P2O5 after the oleum synthesis? Maybe like drying it at a very high temp or something?

It's not a reversible reaction as the oxygen is incorporated into the chemical structure.

Keras - 25-5-2023 at 21:56

Quote: Originally posted by Raid

Is there a way of recovering the P2O5 after the oleum synthesis? Maybe like drying it at a very high temp or something?

As far as I know, no. You can’t get back from phosphoric acid to phosphorus pentoxide. As you heat, the acid transforms into others forms (e.g. pyrophosphoric acid) and then polymerises. At a pinch, you maybe able to recover phosphorus, but that needs a special apparatus since phosphoric acids attack glass at high temperature (and metal also – so you'd need clay or something).

Tsjerk - 25-5-2023 at 22:06

In theory; yes. In practice; check out the lengthy elemental phosphorus thread...

You could separate the phosphate from the sulfate, reduce to phosphorus, oxidize to P2O5. But none of those steps is going to be easy.

Raid - 26-5-2023 at 05:56

hmm, I see.
Thanks for the info!

clearly_not_atara - 26-5-2023 at 14:00

I'm sure it isn't easy, but I've found a couple of reports that aluminum dihydrogen phosphate dehydrates to the metaphosphate Al(PO3)3 around 800-900 C and releases P2O5 around 1000-1200 C to leave AlPO4. So it seems like you theoretically might be able to avoid reducing phosphorus. As far as I can tell, this is not used in practice. The temperature control is probably tricky, but it's the only direct "synthesis" of P2O5 I've seen in the literature. Interestingly, if the partial pressure of oxygen is low, the gas released is actually P2O3 + O2.

The "azeotropic point" of metaphosphoric acid is at about 860 C and is slightly enriched in P2O5 compared to the nominal formula HPO3. But metaphosphoric acid is incredibly corrosive and at these temperatures would be even more so. So this is not useful.

BromicAcid - 26-5-2023 at 17:49

But the real question is, do you need to go back? There are multiple threads around here on using the meta phosphoric acids and pyro phosphoric acids as dehydrating agents:

One example
http://www.sciencemadness.org/talk/viewthread.php?tid=4409

It's even specifically called out that it can dehydrate sulfuric acid to sulfur trioxide. And would be recyclable (with some effort).

Loptr - 27-5-2023 at 09:08

Quote: Originally posted by BromicAcid

But the real question is, do you need to go back? There are multiple threads around here on using the meta phosphoric acids and pyro phosphoric acids as dehydrating agents:

One example
http://www.sciencemadness.org/talk/viewthread.php?tid=4409

It's even specifically called out that it can dehydrate sulfuric acid to sulfur trioxide. And would be recyclable (with some effort).

This was my thought exactly.

What about sodium pyrosulfate? Would that also be an option?

BromicAcid - 27-5-2023 at 13:23

Quote: Originally posted by Loptr

What about sodium pyrosulfate? Would that also be an option?

A cursory glance shows it could be. Take your sodium sulfate, treat with stoichiometric sulfuric acid to make the bisulfate then strip down, decompose to the pyrosulfate. Then workup to make your sulfur trioxide and take the residues and repeat. Or course that is a gross simplification but it's a reasonable start.