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Author: Subject: Sodium Ethyl Sulfate
sulfuric acid is the king
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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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[*] posted on 18-11-2017 at 04:55


I just did two runs to see how heat affects yields. During one run I boiled the solution down. During the other I evaporated it in a tray with a fan and a space heater.

Both had lousy yields, around 17%. Both had a lengthy plateau where the specific gravity of the solution remained at around 1.27-1.28 while lots of solids precipitated. I prefer the tray/fan/space heater route because it requires less supervision. It was also actually extremely easy except for the vacuum filtrations, but I had the assistance of a mechanical stirrer.

It seems plausible that the apparent decomposition and decreases in pH I had observed previously might be caused by impurities in drain cleaner sulfuric acid.

Heating the solution to 110 C doesn't seem to degrade the product much with the pH at around 9, but I think there is some kind of co-crystallization that occurs with a hydrate of sodium carbonate and sodium ethyl sulfate.

IMG_20171118_060047 - Copy.jpg - 408kB

As has been reported elsewhere, it resembles cauliflower.

[Edited on 18-11-2017 by JJay]

[Edited on 18-11-2017 by JJay]




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[*] posted on 18-11-2017 at 12:37


Great work JJay! This post was very validating for me, my result looked very similar on my one and only attempt so far. But due to an unfortunate rookie mistake involving an open stopcock on a sep funnel I lost nearly all of the resulting nitroethane I made with it and that was my only way of validating I had produced the proper compound. I will attempt again soon and let you all know if I find anything worthwhile during my efforts. I have a couple of ideas on improvements to its preparation that I want to try out.



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[*] posted on 22-11-2017 at 14:25


@JJay
Great.Thanks for the useful info.
@CaptainMolo
Which method did you use to obtain nitroethane?
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[*] posted on 22-11-2017 at 17:26


Ive got some video on my channel about sodium ethyl sulfate and nitroethane. Cheers nux.



My YouTube channel Nux's channel
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[*] posted on 26-11-2017 at 20:53


Hey it's Nux! I didn't know you were on here! It's Full Modern Alchemist from YouTube. Sorry I know that was off topic, but I had to say hi.

Also I've been thinking about this for a while and I wanted to run it by you guys to see what you thought.

So I was thinking I'd do the esterification by combining the Ethanol and Sulfuric Acid all at once, instead of dropping the acid into the Ethanol over time. Of course I'd be adding the Sulfuric Acid to the Ethanol and the Ethanol would be as anhydrous as possible and in excess to help avoid ether and diethyl sulfate formation.

Another modification is that I would have a bed of anhydrous Magnesium Sulfate on the bottom of the flask. I was thinking I could accomplish this by having the magnetic stirring offset to one side and slightly elevated, so that the stirbar is spinning just above this bed of sulfate salt. This would probably create a horizontal vortex just above the Magnesium Sulfate bed, which should be enough mixing to help keep homogeneity.

The thinking there is that in the 1800's journal article I read a while back, the esterification seems to proceed under it's own generated heat perfectly sufficiently, and is mostly complete in under 15 minutes. Another thing about fisher esterification, is that the equilibrium can be pushed forward with the removal of water, as per Le Chatelier's principle. The Magnesium Sulfate will accomplish this by complexing with the generated water molecules and sequestering them from the reaction mixture.

There's another strategy that involves a sohxlet extractor and a reflux column filled with molecular sieves for removing water from reactions like this, but the problem with that is the heat needed to reflux this mixture would probably drive more production of Diethyl Sulfate.

Anyway once this initial reaction is over, I would filter the Magnesium Sulfate out and then neutralize the filtrate with a chilled solution of Sodium Hydroxide, because I read that the rate of hydrolysis at lower temperatures is very slow, so the water in the solution shouldn't matter much. Once I had achieved a pH of around 8 or 9 I would then proceed to evaporate the excess water, and with any luck I would end up with relatively pure Sodium Ethyl Sulfate, avoiding the process of filtering out excess Carbonates that comes with neutralizing with those bases.

Anyway I am rambling because I've been out of the lab for several weeks and I'm slowly losing my mind with all the ideas I would like to try, so please excuse this wordy and possibly useless post. :D





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[*] posted on 6-4-2018 at 15:31
Any success


Any success with those optimizations captain, really interested in a synthesis of sodium ethyl sulfate but im no chemist so sifting through this page to come up with my own is doomed to failure. Any help would be appreciated, once ive got an idea of what to do ill give it a shot and report as detailed as possible here.

Also, first post here so sorry if attempting to resurrect this thread is considered inappropriate in any way.

Thanks!
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[*] posted on 7-4-2018 at 04:43


I am of the opinion that this thread is justified and appropriate; sodium ethyl sulfate has a number of niche uses. I've been wondering if it could be dry distilled with potassium iodide to make ethyl iodide.



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