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Author: Subject: Sodium Ethyl Sulfate
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[*] posted on 11-8-2017 at 15:01


Thanks.
I think dilution is probably for less mess :D
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[*] posted on 16-8-2017 at 06:41


@JJay
Finally i have successfully synthisized Ca(EtSO4)2...
'I poured Na2CO3 solution into it and then CaCO3 formed...
I evaporated filtrate directly on hot plate and there was very little NaEtSO4 on the walls...
I am absolutrly unsatisfied with the yield,probaly the temp was to high it should be evaporated on water bath,but i think that would take very long time.
And adding Na2CO3 was unaccurate...
So what you do after making Ca(EtSO4)2?
How you know how much is enoug (Na2CO3),and then how you exactly evaporate it?
And yeah Na2CO3 is very bad and unsoluble,it's better K,but i don't have it,maybe i'll make it...
BTW this is my new account,couse i had major HDD failure,i lost some passwords...

EDIT(woelen): The original "sulfuric acid is the king" account is operational again.

[Edited on 21-8-17 by woelen]
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[*] posted on 16-8-2017 at 07:47


You can get a pretty good indicator of when the Ca(EtSO4)2 is neutralized when adding sodium carbonate solution does not produce a precipitate. Mix thoroughly and check the pH... I had the best luck with a pH of about 9... there was a lot of decomposition when evaporating the water with lower pH for some reason. I concentrated it on a hotplate, filtering any salts that precipitated out (calcium bicarbonate seems to be problematic here). Then I heated it on a hot water bath until the liquid hardened like rock candy when removed on a spatula and decanted onto a surface and allowed it to cool. It's somewhat easier to remove from aluminum foil than a glass plate.



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[*] posted on 4-9-2017 at 08:19


Few days ago i tried different ratios etc...
Today i tried to obtain Ca(EtSO4)2 powder but...It decomposed.
What have i done?
I poured EtHSO4 to water,neutralised with CaCO3,then filtered CaSO4 that was formed during neutralization.In the clear solution was some Ca(EtSO4)2...
Then i boiled it directly on hot plate.After half an hour or so i was left with white sludge...
Then i collected it,and tried to evaporate it again on hot plate,couse it was hard to dry on the sun or with hair drier...
There was some steam,and the smell was really pungent.
Maybe it decomposed to some sulfuric acid.I don't know what can it be...
But later when i added water,it was unsoluble.
What a disaster...
I did not know that even calcium salt of EtHSO4 is so sensitive..
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[*] posted on 4-9-2017 at 17:53


Sodium ethyl sulfate decomposes at around 130 C to produce ether. Some amount of ethyl alcohol is produced as well. I'm not exactly sure what the mechanism is. I suspect that heating it may produce some diethyl sulfate (nasty stuff) as well as probably ethene at higher temperatures, and there are numerous decomposition products that might be produced from the substances mentioned. Oh and carbon dioxide could displace ethylsulfuric acid. I did accidentally smell some sodium ethyl sulfate decomposing on a hotplate in an early experiment, and it was not something I wanted to smell again. It smelled like burning rubber.

I'd expect calcium ethyl sulfate to act similarly though it undoubtedly doesn't react quite the same.

The potassium salt is said to be easier to crystallize.

I read recently in an old edition of Systematic Organic Chemistry that it's possible to distill acetonitrile from a melt consisting of sodium ethyl sulfate and potassium cyanide. This was described as a "very general" reaction.




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[*] posted on 4-9-2017 at 22:10


Ca(EtSO4)2 >> Et2SO4 + CaSO4 is my guess. You didn't smell sulfur dioxide, and SO3 should not be stable, so the sulfate probably stuck to the ethyl groups.
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[*] posted on 4-9-2017 at 22:49


Diethyl sulfate smells good, but it is hazardous at levels barely above the odor threshold, so if you smell a pleasant peppermint-like odor coming from your reaction, you need better ventilation.



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[*] posted on 5-9-2017 at 02:30


Ok,thanks guys...
I remember some very pleasent sweet smell but way before boiling clear solution,when i rised temperature of CaSO4,Ca(EtSO4)2 solution,before filtering...
After evaporating clear solution,and then heating white sludge (probably CaSO4 of decomposition and some left Ca(EtSO4)2,vapor in the air was very iritating and coused cough...
Yea it's probably bad stuff,but i will stop action for some time,and regenerate :)
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[*] posted on 18-9-2017 at 14:20


Quote: Originally posted by JJay  
It's somewhat easier to remove from aluminum foil than a glass plate.

So it's unreactive with Al?
This would be great.
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[*] posted on 18-9-2017 at 21:26


I'm not 100% sure that it can't react with aluminum, but it doesn't ordinarily... it very well may react with aluminum amalgam.



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[*] posted on 14-10-2017 at 07:36
Solubility questions.


I'm currently working out some optimizations to this synthesis, and I have some questions that I need help answering, as I can't find any good data.

Does anyone have access to solubility data for Sodium Ethyl Sulfate?

I found conflicting reports about it's solubility in alcohols and water.

Does anyone know how fast this compound hydrolyzes?




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[*] posted on 14-10-2017 at 10:59


It's extremely soluble in water... I'm not seeing the figures offhand, but I have seen a figure for its solubility in water published somewhere.... It's soluble in alcohols but not nearly as soluble as in water. I'd say roughly 1.5x weight equivalents of sodium ethyl sulfate (hydrated) will dissolve in 1 equivalent water but in alcohols it is more like 1 equivalent of sodium ethyl sulfate to 10 equivalents of methanol. It is not soluble in ether.

It hydrolyzes quite slowly in slightly alkaline conditions and somewhat more quickly but still pretty slowly in acidic conditions (it can take days to hydrolyze detectably in aqueous acid at room temperature but decomposes somewhat more quickly in acidic conditions with heating). A little bit of thermodynamic data has been published on ethyl sulfuric acid here: http://pubs.acs.org/doi/abs/10.1021/ja02248a014 but I don't think the effect of pH on rate constants, etc has been explored much. I've never seen any documentation stating that anyone else observed hydrolysis in alkaline conditions, though, so I can't rule out that the slight hydrolysis I've observed on a few occasions was caused by some impurity.

That's pretty much all the information I have on it, but a lot of early chemistry books discuss it (with varying degrees of accuracy), and it is likely one of the oldest synthetic substances known.




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[*] posted on 14-10-2017 at 16:13


Thanks JJay! I found the document you referenced on google books for those who can't access the full text via JACS.

That's very helpful, thank you!

If I find anything else of use, or I have a chance to get back in the lab soon, I'll report back here with my findings.

[EDIT]

The most interesting tidbit so far:

Quote:

This shows that there is no advantage in the artificial heating of the mixture of alcohol and sulfuric acid generally employed in the preparation of ethyl-sulfuric acid and its salts, the spontaneous heating to about 70° effecting a maximum production of ethyl sulfuric acid within 10 minutes.


[Edited on 15-10-2017 by CaptainMolo]




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[*] posted on 15-10-2017 at 03:13


My NaEtSO4 has little lower pH than Na2CO3 solution,so it goes from basic little bit to the neutral,by colour of improvised pH indicator...
Crystals are also little bit harder,and it burns paper if it stays on it long enough (few days or so)...
So can anybody confirm,can NaEtSO4 be little bit basic?
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[*] posted on 15-10-2017 at 04:14


I'm pretty sure that if completely pure, the pH of sodium ethyl sulfate solution is very close to neutral. But a little bit of excess sodium carbonate seems to stabilize it.

I hadn't realized that it burns paper... I'll have to check that out....

Edit: added word "solution"

[Edited on 15-10-2017 by JJay]




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[*] posted on 15-10-2017 at 05:53


I know what could it be...
Fact that the solution is basic but more to the neutral than Na2CO3 is because of excess Na2CO3.
Paper is burned probably due to wet NaEtSO4,over time maybe it decomposet to H2SO4 and Et...
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[*] posted on 15-10-2017 at 12:38


You might want to take a look at this discussion on sodium sulfate: https://chemistry.stackexchange.com/questions/57467/why-is-s...



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