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maxidastier
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Thank you!
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blogfast25
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No probs. Let us know how you get on. I'm really interested in making some dioxane but had never heard of the toluene sulphate method...
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maxidastier
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AFAIK you don't use the sulphate but the p-TsOH monohydrate.
I've one last question before I try it again: Since the p-TsOH is a catalyst, it doesn't get consumed right? So can I use the same amount of TsOH for
even more ethylene glycol if I continuously add fresh ethylene glycol?
By this way I also don't get that much waste at the end.
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blogfast25
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Quote: Originally posted by maxidastier  | Since the p-TsOH is a catalyst, it doesn't get consumed right? So can I use the same amount of TsOH for even more ethylene glycol if I continuously
add fresh ethylene glycol?
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Catalysts, chemically speaking, don't get consumed, that's correct. But catalyst recovery can be quite a kerfuffle, as I suspect will be the case
here. Just adding more and more glycol won't cut it, I'm fairly sure.
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maxidastier
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I didn't want to recover the catalyst, but if you say a continuous process isn't possible, ok then.
[Edited on 26-8-2011 by maxidastier]
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maxidastier
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Can HCl destroy the p-TsOH if it isn't removed immediately?
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blogfast25
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No.
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maxidastier
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This is now my second attempt to make p-TsOH from H2SO4 and Toluene.
I've with the help of two literature instructions. One is from Gattermann-Wieland and Experiments in Organic Chemistry by Fieser.
So, then: I've used 40ml 96% H2SO4 and 200ml Toluene and a Dean-Stark-Trap as Gatterman-Wieland suggests.
Everything went fine. I've got around 13ml of H2O after 4 hours, so should it be complete reaction, which means there still were aroung 115ml
unreacted Toluene left. I decided to collect this with the same apparatus, and collected aroung 100ml through the Dean-Stark-Trap.
Then I did something, which MIGHT be the reason for the failed attempt: I went upstairs to go to the toilet and drink something.
When I came back, the residue in the flask had become dark brown- I immediately removed the flask from the heating mantle.
I transferred the mixture in a beaker glass to cool down. I added around 1 ml of water to test, if there was some TsOH. Where the drops touched the
water, instantly became white and solid.
So I let it cool down and added 100ml conc. HCl. I heated until boiling and cool again to recrystallize and get mainly the para isomere.
But I didn't get white crystals. I only got a brown slurry, when I vacuum-filtrated it.
I thought okay, then another recrystallization, added another 50ml HCl, heated, cooled and still having brow-white slurry.
Is the dark brown colour after the reaction normal? Here, I can't see any brown: http://www.sciencemadness.org/talk/viewthread.php?tid=15522
Does that mean, the TsOH became completely desulfonated while I was upstairs and the heat was to high?
When I cleaned my flask and other stuff with water, there were white solid pieces swimming in it. This can't be TsOH, since it is soluble in water and
I couldn't dissolve by shaking or anything. What is it?
[Edited on 9-9-2011 by maxidastier]
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Nicodem
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| Quote: Originally posted by garage chemist | When no more water is collecting in the trap (all H2SO4 has reacted), one adds a small, carefully measured amount of water to the mix after cooling
which makes the p-TsOH crystallize from the excess toluene as the monohydrate. The byproduct o-TsOH does not form such a hydrate and stays in
solution.
The filtered crude p-TsOH*H2O is then purified by dissolving it in half its own weight of water and saturating the solution with HCl gas.
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For your needs you can skip the tedious recrystallization step. Impurities don't make much of a difference in your application.
Given you appear to be doing this out of practice/fun rather than the need for dioxane (or else you would just use the H2SO4 based procedures), you
might as well try using NaHSO4 as the catalyst (instead TsOH). It is available in any shop selling pool chemicals.
The idea of using a continuous process is a good one (reactive distillation of products as they form and feeding in fresh ethylene glycol), but I
think you should optimize the batch process first to get an understanding of the critical reaction parameters. Though with a continuous process you
would still end up doing a batch isolation process.
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maxidastier
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Thanks for your input!
But I still don't know what has happened to my TsOH...
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maxidastier
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It's really a mess. The problem is, after adding the water to make the p-TsOH crystallize as the monohydrate, it does not crystallize as long as the
solution is to hot, but if it's cool enough, after adding water, everything solidifies, the o-TsOH does not stay in solution. How can I get rid of it
then and moreover, how can I get rid of the toluene?
I didn't use HCl this time as you suggested. Now I have some clumpy, pink, wet crystals which dissolve in the air and there are little droplets where
the solid used to be.
What am I doing wrong?
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maxidastier
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Is there another way aside from potassium carbonate to create an organic layer containing mainly Dioxane and another layer containing the water, since
after the synthesis there would be a 1:1 mix of Dioxane/water?
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blogfast25
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I'm not sure why specifically it calls for (the more expensive) K2CO3, when I believe Na2CO3 should do just as well (my guess). But the principle of
'salting' cannot be avoided, AFAIK...
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maxidastier
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Yes, that is exactly my question: Could I use Na2CO3 instead?
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Nicodem
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| Quote: Originally posted by maxidastier | | It's really a mess. The problem is, after adding the water to make the p-TsOH crystallize as the monohydrate, it does not crystallize as long as the
solution is to hot, but if it's cool enough, after adding water, everything solidifies, the o-TsOH does not stay in solution. How can I get rid of it
then and moreover, how can I get rid of the toluene? |
What kind of analysis gave you the confirmation that ortho-toluenesulfonic acid also precipitated? Why don't you nevertheless follow up with the
isolation, do the vacuum filtration, wash the product, etc.? Why do you want to remove all the ortho isomer when it is obvious you need the product as
an acid catalyst?
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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blogfast25
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I see no good reason why not.
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Takron
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Could you use sodium bicarbonate instead of the carbonate?
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blogfast25
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No. It doesn't have any dehydrating power whatsoever.
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bbartlog
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And Na2CO3 is a pretty poor dehydrating agent compared to K2CO3.
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blogfast25
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I've never understood whether that is true or not: it seems counter-intuitive that a substance that froms no hydrates would be a better dessicant than
a substance that forms a relativly stable decahydrate.
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