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Author: Subject: SO3 from pyrosulfate and H2SO4
papaya
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[*] posted on 15-9-2015 at 11:36


Quote: Originally posted by ave369  
Today I tried this synthesis but failed it on the stage of making pyrosulfate. I chose a bad vessel for calcination and fouled the entire yield.

However, it seems that I made an interesting discovery. I'm still unsure if I understand correctly what happened, but I'll describe.

I had a small flask filled with a mixture of H2SO4 and HCl standing on a shelf in my lab for a long time. It was made by bubbling hydrogen chloride gas through 80% sulfuric acid. Today I tried to purify this acid by distilling HCl out of it.

I did it, and here's what I found: an addition of HCl completely suppresses the formation of the infamous white mist! Instead of it, azeotropic HCl vapors come out, which are easily absorbed by poking the retort's nose in water or even just condensed and dripped down, if the speed of distillation is slow enough. When all HCl was distilled away from the mixture, in the retort was 95% sulfuric acid! When I opened the retort's tubulus while it was still hot, it started fuming profusely and the retort filled itself with white mist, but while closed, no SOx at all emerged!
[Edited on 15-9-2015 by ave369]


Hmm, this is quite interesting, if your observations are right (getting higher concentration H2SO4 than initiall was there) then it seems HCL can pull water from conc. H2SO4 and put it into gas phase. Something similar happens in industry when benzene is used to get absolute alcohol (benzene forming azeotrope with water and leaving the dry alcohol). Now, what if you try using the same very dry benzene or even toluene to pull water from H2SO4 (add it to your 80% H2SO4 and distill together)? I doubt it though, but who knows?

[Edited on 15-9-2015 by papaya]
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ave369
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[*] posted on 15-9-2015 at 14:31


Obtaining toluene here in my town is a mess. A full day's work extracting it from an OTC solvent. Benzene is right out.

However, in my experiment, the concentration of sulfuric acid definitely increased. Maybe I was wrong about the number 95%, but not by a very wide margin: it's at least 94%. There can be no mistake. It has the density. And it does everything real concentrated H2SO4 does and 80% doesn't: it quickly chars wood, water dropped in it goes pop-sizzle. My normal boildowns never go that far: I usually stop at 80%, or at most 90% if I do the boiling down by remote hotplate control in an unmanned bath house, so the SOx vapors are less an issue. This boildown with HCl, this one I did at home, with absolutely none choking gas attacks.

[Edited on 15-9-2015 by ave369]




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[*] posted on 15-9-2015 at 15:01


Quote: Originally posted by papaya  

Now, what if you try using the same very dry benzene or even toluene to pull water from H2SO4 (add it to your 80% H2SO4 and distill together)? I doubt it though, but who knows?

That won't work. Under those conditions, both benzene and toluene will readily sulphonate and will simply contaminate, and add water to, the sulfuric acid.




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[*] posted on 15-9-2015 at 15:05


Quote: Originally posted by ave369  
Today I tried this synthesis but failed it on the stage of making pyrosulfate. I chose a bad vessel for calcination and fouled the entire yield.

However, it seems that I made an interesting discovery. I'm still unsure if I understand correctly what happened, but I'll describe.

I had a small flask filled with a mixture of H2SO4 and HCl standing on a shelf in my lab for a long time. It was made by bubbling hydrogen chloride gas through 80% sulfuric acid. Today I tried to purify this acid by distilling HCl out of it.

I did it, and here's what I found: an addition of HCl completely suppresses the formation of the infamous white mist! Instead of it, azeotropic HCl vapors come out, which are easily absorbed by poking the retort's nose in water or even just condensed and dripped down, if the speed of distillation is slow enough. When all HCl was distilled away from the mixture, in the retort was 95% sulfuric acid! When I opened the retort's tubulus while it was still hot, it started fuming profusely and the retort filled itself with white mist, but while closed, no SOx at all emerged!

I'm curious as to whether the concentrated acid still contains traces of chloride, do you have any silver nitrate to test it?




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papaya
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[*] posted on 15-9-2015 at 15:08


At lease benzene should not sulphonate? If I remember correctly an old procedure called to purify benzene from thiophene (very close physical properties to C6H6, cannot be separated easyly) is boiling technical grade benzene with conc. H2SO4, that destrois thiophene. I'm not sure on toluene, at least 80% sulfuric is not like 96% sulfuric, hardly it can do anything to toluene at that concentration, still needs to be tested.
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[*] posted on 15-9-2015 at 15:12


ave369 it seems that you have made a very useful discovery for those wishing to obtain 98% H2SO4 (bp = 338°C) starting with weaker acid.

The HCl-water maximum boiling azeotrope boils at 108.6°C at a concentration of 20.2%. So if you add HCl to, say, 80% H2SO4, boiling should remove water in the distillate until the distillate temperature reaches ~108.6°C. Then it will constant temp/constant concentration (20.2%) boil until the HCl has all been evaporated.

If, when the concentration of H2SO4 reaches 98% and there is still HCl left, I imagine it would evaporate off very quickly.





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[*] posted on 15-9-2015 at 15:15


I believe that the standard procedure for removing thiophenes from benzene is by shaking with room temperature sulfuric acid, at least that's what Vogel calls for. When purifying toluene, both the toluene and sulfuric acid need to be chilled to prevent partial sulfonation of the toluene, I've done this before. And while the initial 80% sulfuric acid may not sulfonate the toluene or benzene, the more concentrated acid desired as the product will at boiling.



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[*] posted on 15-9-2015 at 15:25


I don't insist on anything, but still it's known that aromatic sulfonations are reversible, meaning that when benzene sulfonate is starting to form, there should be only a little water left, that's why oleum is preferred (higher yields), isn't it? Benzene at least can be re-used if this scheme works at all, hope organics people will correct me if wrong. And HCl find is a good find ave, congrats!

[Edited on 15-9-2015 by papaya]
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[*] posted on 15-9-2015 at 15:43


The issue I am seeing is that, even though the reaction is indeed reversible, I believe that it will inevitably contaminate the sulfuric acid with aromatic sulphonates. Here's an example : Benzene is added to 80% sulfuric acid and the mixture is refluxed with a Dean-Stark, or simply subjected to distillation. Initially, the less concentrated sulfuric acid may or may not be strong enough to sulphonate the benzene, so the benzene/water azeotrope will begin to be collected. As water is removed from the reaction mixture, the equilibrium will begin to shift towards the aromatic sulphonate side due to water being lost, though the loss of benzene mitigates this somewhat. As the aromatic sulphonate forms, the amount of volatile benzene drops and, after a certain point, disappears, however some sulphonate still remains. Seeing as this is dissolved in now nearly concentrated sulfuric acid, there is no way to shift the equilibrium back towards the benzene side to remove this, so I don't see any easy way to remove it. Do you agree, or am I missing something?



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[*] posted on 15-9-2015 at 22:58


Quote: Originally posted by gdflp  

I'm curious as to whether the concentrated acid still contains traces of chloride, do you have any silver nitrate to test it?


Yes, I do, and already did that during the distillation. The still bottoms do not contain trace chloride; I was distilling HCl out until the distilled liquid gave a very faint cloudiness with AgNO3 (and a profuse precipitation with Ba (OH)2, which meant it has much more H2SO4 than HCl).

Upd: repeated the experiment today. The effect was observed, but to a lesser degree (I concentrated sulfuric acid from 78 to 86%). Probably messed up the required proportion between H2SO4 and HCl: I've got that particular mixture by pure random chance. I'll conduct further experiments to come up with a definite methodics.

[Edited on 16-9-2015 by ave369]

[Edited on 16-9-2015 by ave369]




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[*] posted on 3-1-2016 at 15:27


Yes the novel HCl trick is a great find. I am so used to seeing the reverse. Adding concentrated sulfuric to concentrated HCl solution to get HCl gas. I was thrilled when I learned calcium chloride could release HCl gas from the acid.

What about adding dehydrated magnesium sulfate to 80% sulfuric acid and distilling? That would be simular to the magnesium nitrate concentration of nitric acid.




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[*] posted on 4-1-2016 at 05:09


chloric1, what HCl trick are you talkin bout? Dehydration of acidic chloride sources with either CaCl2 or H2SO4 or some other salt is a common way of HCl gas preparation.
Sulfuric acid can be dehydrated to oleum (SO3) with Na2SO4-MgSO4, but you need a high temperature (like 400 °C) and there's a problem with cleaning the device from moltent salt left after the reaction.
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[*] posted on 4-1-2016 at 16:10


Quote: Originally posted by byko3y  
chloric1, what HCl trick are you talkin bout? Dehydration of acidic chloride sources with either CaCl2 or H2SO4 or some other salt is a common way of HCl gas preparation.
Sulfuric acid can be dehydrated to oleum (SO3) with Na2SO4-MgSO4, but you need a high temperature (like 400 °C) and there's a problem with cleaning the device from moltent salt left after the reaction.


I was talking about ave369 adding HCl to 80% sulfuric acid and distilling off azeotropic hydrochloric acid to leave 93 -95% sulfuric acid behind. It is directly opposite of adding concentrated sulfuric acid to hydrochloric acid to get HCl gas. If Na2SO4 MgSO4 mix will liberate SO3 from sulfuric acid, it would surely concentrate 80% sulfuric acid to 93% or more.




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[*] posted on 5-1-2016 at 05:47


You don't need to add anything to make 93-95% sulfuric acid, simple distillation o dillute H2SO4 is enough.
No, sulfates can't simply dehydrate sulfuric acid. There's only few compounds that can do this, like polyphosphoric acid.
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[*] posted on 6-1-2016 at 15:04


It makes sense to me: the affinity of HCl for H2O is huge, comparable to H2SO4 itself.

Quote:
Upd: repeated the experiment today. The effect was observed, but to a lesser degree (I concentrated sulfuric acid from 78 to 86%). Probably messed up the required proportion between H2SO4 and HCl: I've got that particular mixture by pure random chance. I'll conduct further experiments to come up with a definite methodics.


I'm guessing it simply gets better the more HCl you use. It probably won't be possible to use too much HCl, as it has relatively little love for H2SO4; any "extra" will simply boil away (an unpleasant gas, but not as bad as H2SO4). The caveat is that you will, obviously, need to both generate dry HCl and trap the outgas.

The SO3 can be trapped as a complex; it forms complexes most famously with pyridine but also dioxane and DMF. Any of these is easier to handle than oleum...
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[*] posted on 7-1-2016 at 05:32


Are you kidding me? The affinity of H2SO4 for water is so huge it can dehydrate organic like sugar into carbon at room temperature.
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[*] posted on 7-1-2016 at 08:03


A common lab practice to generate HCl gas is to drip aqueous HCl onto concentrated sulfuric acid. The sulfuric acid strips the water from the HCl very effectively. Azeotropes exist because of affinity between two polar compounds. When another compound with any degree of affinity is introduced, the azeotrope shifts, and in some cases is destroyed.

Certainly the HCl you distill from the mixture cannot be its 20.2% azeotrope at atmospheric pressure, because if the affinity H2SO4 has for water. However, HCl is still acting as a volitile entrainer with some affinity for water. I suspect that for a given temperature, heating H2SO4/HCl/water mix is more effective at dehydrating the sulfuric acid than just heating H2SO4/water alone. However, the caveat is that it will reach an asymptotic equilibrium of concentration where the affinity for water is the same for both acids at that particular concentration, and no further concentration can be achieved. This will likely be substantially less than the 98% achievable by boiling.

Theoretically, one needs to add HCl of a concentration higher than the azeotrope that can be distilled from the desired concentration of sulfuric acid, to achieve that concentration. The theoretical maximum concentration of H2SO4 achievable with this method would be obtained by bubbling pure, gaseous HCl through aqueous sulfuric acid at some arbitrary temperature, preferentially less than that required to drive the water off using the old boiling method, but higher than the boiling point of the resulting HCl/water mixture. In this way I think it could be concentrated further than its 98% azeotrope, although it would take a long time unless the bath was simultaneously heated on the same terms as the boiling method.

In short, 31% hardware store HCl will only do so much. I fear that it may only serve to be a grossly inefficient way to achieve a concentration of H2SO4 much less than 98%, when boiling it could achieve an equal concentration at a reasonable temperature with far less apparatus. Concentrating beyond that would require the generation of HCl gas without conc. H2SO4 which involves more apparatus and reagents. At that rate, you may as well invest in a throwaway hotplate from WalMart and a $4 beaker, and just boil it down outside. It's free, and everything is totally reusable as long as you don't do anything stupid with the hot glass.

Besides, running all that HCl through the H2SO4 will concentrate metallic impurities that are common in hardware-store HCl and leave them in the sulfuric acid, while at the same time introducing traces of chloride which are very hard to remove.

[Edited on 7-1-2016 by Praxichys]




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