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JJay
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You can't use the common ion effect to eliminate sulfates with ethylsulfate ions.
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Why not? Regardless, that's reported to work, I'm not just assuming that it does.
[Edited on 16-1-2017 by alking]
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JJay
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Where is that reported? I think you're confused.
You can't even use the common ion effect to crash out bisulfate with sulfate. They are different ions.
[Edited on 16-1-2017 by JJay]
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You could be right. I can't remember where I read it now, it was probably a thread here on it, someone who did it claimed this iirc.
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JJay
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I just ran across this paper on the kinetics of the reaction between ethyl alcohol and sulfuric acid. It contradicts some other research I have seen,
but this paper was published by ACS. While the researchers didn't look at the effect of water on the reaction, they did look at temperature, and
apparently, heating for an extended period of time actually hurts the yield.... Also, they used equimolar quantities to achieve a 60% yield, which is
actually good compared to what has often been reported for this reaction.
They stated that the reaction is complete in 10 minutes at 70 C with 95% sulfuric and 99.9% grain alcohol, and their suggestion is to add sulfuric
acid to ethanol at a sufficient rate to attain and then maintain that temperature.
Attachment: EthylsulfuricReaction.pdf (1.6MB)
This file has been downloaded 574 times
[Edited on 16-1-2017 by JJay]
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Interesting. I also notice that they see no appreciable hydrolysis when diluting the HEtSO4, even when leaving it for 24+ hours.
edit: It should also be noted, to clarify on what you just said JJay, that they did not heat the solution to 70C for 10 minutes, they simply put it in
a bath that is at 70C for 10 minutes, so for some of that time the solution is warming up. The quantities tested were rather small so that length of
time may be negligible, but with such a small time frame of 10 minutes it may not.
[Edited on 19-1-2017 by alking]
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JJay
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Hmm... looking at their experimental results, it looks like no hydrolysis takes place at 60 C, and the reaction is complete in around 20 minutes....
Others have stated that HEtSO<sub>4</sub> does hydrolyze pretty quickly in water at room temperature. That seems unlikely.
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Yeah, I wouldn't think it would based on this, their methods seem pretty thorough. They did say somewhere near the end of the paper that it *does*
hydrolyze when water is added which seems to contradict what they said earlier. It may have involved heating though, I can't remember. However even
then they said the hydrolysis was slow and negligible, so unless you're adding water and then going to abandon the project for days it shouldn't
matter.
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So I've tried this twice now, two different ways. Unfortunately I did not write down my yields, though I did figure them out at some point, but I
think the first time it was around 60-70%, and the way I did it last night, which was a HUGE pain in the ass, I will never do that way again, yielded
only 30-40% assuming my solution is relatively pure (I evapped it down and left it as a saturated solution with a small bit of precipitated salts,
density is ~1.29g/ml, or what should be about 49% NaHEtOH by weight).
Both methods were with H2SO4, 98%, Anhydrous or near anhydrous EtOH, and Na2SO4. The first time, iirc, I did not heat it and let it sit overnight
stirring in the Na2SO4, however I forgot to account for the h2o created in the process so I could actually see getting much higher yields if I had.
Yesterday I tried heating it to 40-50C for 30 minutes, then I added Na2SO4 and let it sit for an hour, which likely was not enough time to do too
much. The first time the Na2SO4 was also ground into a fine powder.
The main difference was the workup however. The first time decanted the solution, neutralized it with Na2CO3, and evaporated it to dryness. I think
once I filtered off precipitate to aid in the drying process as it began to bump too much. The final drying was done in a vacuum as I'm not sure if
NaHEtSO4 can be dried completely with heat or not, I didn't want to risk it. From there I dissolved it in MeOH, filtered the undissolved carbonates
and sulfates, warmed it up again as the NaHEtSO4 would begin to crystallize as I sucked it through the filter, and then allowed it to cool to w/e my
freezer allows. This yielded very pure, nearly translucent, flaky and uniform crystals. It crystallized very very easily. You definitely do not need
ether for this, and any alcohol should work just as well. I think 100g/100ml is a good rule of thumb, it's *very* soluble in alcohols while hot, and
quite insoluble when cold.
Last night instead of doing what I know works I had some calcium carbonate lying around so I figured I'd take others recommendation and try that. What
a waste of time, a huge mess, and for... what? I don't know. Why would anyone possibly do this method? I spent about 2 hours tediously adding CaCO3 to
a big sludgy mess. I had upwards of a liter of precipitated calcium salts to filter off and wash, and then about a gallon of water to boil down as
opposed to ~3-500ml. Jeez what a huge pain in the ass. Fuck, and such a huge mess too, so many different dishes and whatnot used. Sure it works, but I
would equate this method to relieving my bladder by cutting it open with a box cutter instead of simply urinating. Not to mention you're still going
to have some carbonate and sulfate contaminates, albeit few. By the time I started boiling that gallon of liquid down I could have have a ~200ml
solution of MeOH in the fridge already forming crystals and i still have a mess to clean up as I didn't expect that to take all night as it did.
edit: To clarify the difference in yield I'm sure is mostly due to the method of production, not the extraction part, although I would be surprised if
I did not lose some in all that calcium snot. The reason I had so much water was not only because it would turn to a thick soup and become hard to
stir/neutralize, but because I also filtered twice to make it more manage and because the calcium salts are so voluminous it required quite a bit more
water to then wash it. I'm sure I did lose some in that.
The yield may have been 40-50%, I doubt it was over 50 though. I haven't actually done the math to determine the percentage yield, but I used 5M of
H2SO4 and 6M of EtOH, I yielded ~575ml with a density of 1.25-1.29g (say 1.27), so you can do the math from there. According to antocio in another
post a 52% solution has a density of 1.33g/ml which I believe is fully saturated and thus hard to obtain w/o precipitation or careful measurements.
edit2: Actually I just did the math and I have a yield of roughly 48%, 2.4M from 5M H2SO4. Its likely lower than that as the solution should have
slight contaminants, so maybe 45-47%.
[Edited on 20-1-2017 by alking]
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JJay
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You can obtain a fully saturated solution on a boiling water bath by filtering the precipitated salts periodically. It's hard to avoid heavy
mechanical losses in doing this. It is much harder to reduce the volume of the fully saturated solution on a water bath, but as the water is driven
off, when removing the solution from the bath, you can see the sodium ethyl sulfate crystallize when the solution cools.
Do you know if the form that is recrystallized from MeOH is anhydrous or the monohydrate?
Edit: I should also add that a lot of the salts you are filtering off are probably sodium sulfate from decomposition. I don't think a lot of diethyl
sulfate is created, but I don't really suggest doing this in your kitchen.
[Edited on 21-1-2017 by JJay]
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Are you referring to the CaCO3 method or if you just neutralize with NaCO3 from the start? I thought I said earlier that you should be able to do
that, but you said that it doesn't work? What you're saying though is what i experienced the first time when I just neutralized and dried it. I
shouldn't have experienced any significant losses because I dissolved everything in MeOH afterward, not just my final product but anything that was
filtered in between too.
As to if it's a hydrate or not I'm not sure, how could I find that out? I can test the mp if I had a reference, but didn't you say earlier that you
could only find mention of one and it didn't say which it was?
edit: Also, I thought to make diethy sulfate you had to dehydrate the EtSO4 before boiling it? If it's boiling off in a water bath it should be fine.
I don't know if it's safe to boil it all the way down, but that's why I stopped early and finished in a vacuum. Wouldn't the sodium salt end up with
another breakdown product before 2EtSO4 formed anyway though or do you think you could make it directly from it?
[Edited on 21-1-2017 by alking]
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JJay
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The monohydrate has a melting point of 86 C, according to Commercial Organic Analysis. I have verified that by experimentation, or at least I am
almost completely positive that the material I tested was sodium ethyl sulfate monohydrate. It melted at 86-90 C.
I don't think I ever said that either of those methods doesn't work. My concern with simply neutralizing with carbonate or bicarbonate, drying, and
then extracting with an alcohol is that the usual form of sodium ethyl sulfate is the monohydrate, and sodium sulfate present is not likely to be
anhydrous and will try to bind up to 10 moles of water per mole. The water undoubtedly changes the solubility of sodium sulfate in the alcohol. More
to the point: it almost undoubtedly changes the shape of the solubility curve of sodium sulfate in alcohol. This problem is by no means eliminated by
neutralizing with chalk before forming the sodium ethyl sulfate, though, due to (I believe) decomposition, which is likely temperature dependent.
If one recrystallizes repeatedly from anhydrous ethanol, supposedly, and I'd have to do some digging to find the reference, the sodium ethyl sulfate
contains an ethanol of crystallization. I have not been able to determine whether this occurs with methanol, but it seems likely. I haven't seen any
data on the melting point of the ethanol adduct or the anhydrous form. The recommendation for recrystallizing from ether and methanol comes from none
other than Armarego and Chai's Purification of Laboratory Chemicals, who said to perform this recrystallization three times then dry under vacuum. I
believe the purpose of the multiple recrystallizations is to eliminate water and mineral salts and that the purpose of the ether is to ensure that the
crystals form quickly to keep unwanted hydroxyls out of the crystal lattice to the extent possible.
It seems plausible to me that dry distillation of anhydrous NaEtSO<sub>4</sub> would produce diethyl sulfate, but the I think the presence
of water would tend to favor the formation of hydrogen ethyl sulfate. NaEtSO<sub>4</sub> is rarely prepared in anhydrous form.
[Edited on 21-1-2017 by JJay]
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Oh, that makes sense. Wouldn't any sodium sulfate contamination be pretty minor though? Even if the hydrate increases its solubility there still
shouldn't be enough present to matter for most applications I would think.
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JJay
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I don't know... you may be right, but I have several books and articles stating that best practice is to neutralize with calcium or barium carbonate
before producing an alkali ethyl sulfate. These books may all be wrong.
There are a simple few tests that could be used to test for impurities. One would be to dissolve the ethyl sulfate in some in water and then add a few
drops of barium chloride solution and see if any precipitate is formed (indicating sulfate, carbonate, or perhaps hydroxide impurities). You can add a
few drops of sulfuric acid and look for a precipitate to determine if there is alkali earth metal contamination.
I believe that the common ion effect *can* be used to precipitate sodium sulfate from solution with sodium ethyl sulfate since both share sodium ions
and sodium ethyl sulfate is more soluble.
I've been meaning to do some rigorous testing of different methods of preparing sodium ethyl sulfate, but right now I'm quite limited in the sorts of
materials I can work with, and I am reluctant to handle sulfuric acid or any barium compounds.
When I have time and a good location for doing experiments, I hope to revisit this.
[Edited on 22-1-2017 by JJay]
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JJay
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Oh and just FYI, according to my notes, the last time I made it, the yield was a dismal 18%, but the final specific gravity measurement of the
solution was 1.34. Boiling it down until it hardened like rock candy when cooled resulted in an ivory substance with a wavy crystal pattern. It's very
hard to scrape off of the glass if you don't get to it while it is still a bit soft. It's been said elsewhere that sodium ethyl sulfate crystals
resemble cauliflower; I really can't think of a more apt description. The molten sodium ethyl sulfate was water-clear except for a yellowish tint,
thought to be a dye introduced with the sulfuric acid.
[Edited on 23-1-2017 by JJay]
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sulfuric acid is the king
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I am back 
I had no time for foruming...
Ok,today i made EtSO4 by acid/ethanol mehod,i neutralized it with sodium carbonate,feezzing etc started...It settled at the bottom...
Then i added distilled watter and mixed,mixed,mixed...
After that i put it in the freezer on <10C.And now i have unexpected - clear solution?!
What is going on?!I expected some crystals or whatever?
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JJay
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I'm not sure... do you know how much water was in your starting reactants? Temperature at neutralization could also matter.
I've never seen a clear writeup in a published document stating how to produce a pure product using that method despite seeing several procedures
stating how to produce ethylsulfates through using calcium or barium. While I don't know what the reason is, I'm sure there's a reason.
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sulfuric acid is the king
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Near anhydrous,about 15-20C before adding distilled water which was at room temperature.
I will save that solution for future tests...
Interesting thing: "Solubility in Water almost transparency"
http://www.tcichemicals.com/eshop/de/de/commodity/E0277/
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JJay
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Baffling... I've tried that method a couple of times under similar conditions, and I did get crystals, but I was never sure if they were sodium
sulfate or sodium ethyl sulfate. The amount of water added could matter too, of course....
[Edited on 16-2-2017 by JJay]
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sulfuric acid is the king
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You got crystals,but i got nothing,and now temperature of the solution is little above zero,and nothing.I added water half a volume of first
solution.Even in more saturated,much more it should settle at the bottom,and temperature is very low,so something ultra soluble is in there.
You are not sure what you got in simillar method?Why you just don't melt and see melting point?
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JJay
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I didn't know the melting points for all of the sodium sulfate hydrates or for sodium ethyl sulfate at the time; this was well over a year ago. Come
to think of it, I didn't use a 50/50 mix of ethanol... it was more like 20/80, and I neutralized with saturated sodium carbonate. I don't remember all
of the details, though, and this was before I was keeping a detailed lab notebook.
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sulfuric acid is the king
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Where can i get solubility table for NaEtSO4?
Do you have more info now?
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JJay
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I don't know where you could find a solubility table for it; I had a hard time even finding the melting point with a clear statement of whether the
NaEtSO4 was a hydrate or not.
I do know this: it is extremely soluble in water.
[Edited on 17-2-2017 by JJay]
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sulfuric acid is the king
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From which method was this?
Method with alcohol sulphuric doesn't work for me...
I tryed neutralising it with calcium carbonate,and then filtrate again with sodium carbonate to form unsoluble calcium carbonate,nothing happened...
Fail after fail...Sad 
[Edited on 18-2-2017 by sulfuric acid is the king]
[Edited on 18-2-2017 by sulfuric acid is the king]
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JJay
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Calcium carbonate neutralization then precipitation of the calcium salts with sodium carbonate. The product in the picture was actually not as pure as
I had thought; it contained some mineral salts that caused it to look creamy when melted. I had a lot of failures before I managed to obtain much
product that seemed to have the expected physical properties. It's not that hard but it takes a long time and is not trivial.
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