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Sodium Ethyl Sulfate

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JJay - 15-1-2017 at 15:55

You can't use the common ion effect to eliminate sulfates with ethylsulfate ions.

alking - 15-1-2017 at 17:37

Why not? Regardless, that's reported to work, I'm not just assuming that it does.

[Edited on 16-1-2017 by alking]

JJay - 15-1-2017 at 18:13

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]

alking - 15-1-2017 at 20:50

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.

JJay - 15-1-2017 at 22:19

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 683 times

[Edited on 16-1-2017 by JJay]

alking - 19-1-2017 at 14:09

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]

JJay - 19-1-2017 at 15:46

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 HEtSO4 does hydrolyze pretty quickly in water at room temperature. That seems unlikely.

alking - 19-1-2017 at 16:32

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.

alking - 20-1-2017 at 11:59

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]

JJay - 20-1-2017 at 14:14

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]

alking - 20-1-2017 at 20:18

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]

JJay - 20-1-2017 at 21:59

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 NaEtSO4 would produce diethyl sulfate, but the I think the presence of water would tend to favor the formation of hydrogen ethyl sulfate. NaEtSO4 is rarely prepared in anhydrous form.

[Edited on 21-1-2017 by JJay]

alking - 21-1-2017 at 16:20

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.

JJay - 21-1-2017 at 16:51

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]

JJay - 22-1-2017 at 13:59

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]

sulfuric acid is the king - 16-2-2017 at 08:08

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?

JJay - 16-2-2017 at 13:25

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.

sulfuric acid is the king - 16-2-2017 at 14:05

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/

JJay - 16-2-2017 at 15:20

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]

sulfuric acid is the king - 16-2-2017 at 15:32

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?

JJay - 16-2-2017 at 15:54

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.

sulfuric acid is the king - 16-2-2017 at 16:03

Where can i get solubility table for NaEtSO4?
Do you have more info now?

JJay - 16-2-2017 at 16:08

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]

sulfuric acid is the king - 18-2-2017 at 07:21

Quote: Originally posted by JJay

Oh and of course, here is a picture of the crude product.

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]

JJay - 18-2-2017 at 07:42

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.

sulfuric acid is the king - 18-2-2017 at 08:23

Hmmm so you have done it by acid alcohol methode...
Great.But can you tell me some details...Percentage by mass of the acid,was alcohol anhydrous?How long did reflux take?After reflux you neutralised acid with water?How much?Please tell me details...I am so obssesed with it i will not stop...It takes me days etc...I have improvised equipment but...

JJay - 18-2-2017 at 10:04

Reflux was 24 hours on a boilng water bath, 200 ml 95% sulfuric acid, 600 ml anhydrous ethanol.

The reflux is probably counterproductive; you might want to just heat it to 70 C for 10 minutes. I am still unsure as to the best neutralization method.

sulfuric acid is the king - 18-2-2017 at 13:16

Really?Nice then...I must try...
I added calcium carbonate but it was mess...
You have good expirience with buslulfate method?It should work fine...

JJay - 18-2-2017 at 16:19

Adding calcium carbonate is a mess... and you have to stir it to prevent wasting calcium carbonate. I have used the bisulfate method but never got nearly the yields claimed.

JJay - 18-2-2017 at 21:59

I'm halfway tempted to give this another shot knowing what I know now. I'd like to know the melting point of its anhydrous form (which I assume exists), and it would be good to know its solubility curve in water by temperature....

I imagine that good results detailing its solubility would be accepted for publication in a respectable journal.

sulfuric acid is the king - 19-2-2017 at 03:45

I found this "With excess sulfuric acid, the equilibrium can
be shifted to the right."
But how much excess?Next time i will try 2:1.
Yes i also need that physical properties,but it's very hard to find anything...

[Edited on 19-2-2017 by sulfuric acid is the king]

JJay - 19-2-2017 at 05:20

I've actually never tried using excess sulfuric acid, but I've seen that recommended in a couple of places.

sulfuric acid is the king - 21-2-2017 at 17:24

Ok,when you have calcium ethyl sulfate,how you add sodium carbonate,in solution or in powder form?
I have improvised equipment and everything,i do not have ph indicators etc,it's very dificult for me,but...
I plan to pour sodium sulfate already in solution to solution of calcium ethyl sulfate,and hope for calcium carbonate forming or "snowing" that can be only indicator for me,or?

[Edited on 22-2-2017 by sulfuric acid is the king]

JJay - 21-2-2017 at 20:21

That might work well, actually, but I've never tried it.

I have used sodium carbonate and sodium bicarbonate solution, saturated, sometimes adding powder first, and finishing off with saturated sodium carbonate to a pH of at least 9. The pH seems to decrease slightly when solutions of sodium ethyl sulfate are concentrated over heat unless they are already acidic... in which case the pH decreases rapidly. Exactly why that happens, I'm not sure, but it seems to be some sort of hydrolysis.

Concentrating the solution under vacuum might keep that from happening. The crystals of impurities obtained while concentrating a solution of sodium ethyl sulfate, for me, have had different properties... some have been deliquescent, some have had basically neutral pH, some have been sludges that are hard to filter, some flakes, some crusty substances that creep up the side of the flask, etc. In general, if it remains a solid above 86 C, it's an impurity and should be removed.

As the solution gets concentrated, any excess calcium salts will slowly precipitate out. Calcium carbonate forms a complex with carbon dioxide that is actually pretty soluble in water, and there can be a surprisingly large amount of this, especially if you used sodium bicarbonate for the neutralization. When the solution starts to get very concentrated, unless you somehow magically have absolutely 0 unwanted sodium salts in solution, the excess sodium salts will precipitate out. When this started happening, you'll want to remove the water slowly--I've usually used a boiling water bath for this but low heat on a hotplate works ok too--and remove the salts periodically while checking the specific gravity. It will climb very slowly at times with the elimination of of a lot of salt material, and at other times it will climb very quickly with the elimination of little salt material. Once you have a saturated solution, the specific gravity will be around 1.33-1.34, and removing water at 100 C won't change the volume of the solution much or cause precipitation of salts. In fact, it will look like nothing is happening unless you allow the solution to cool, in which case a crystalline mass of sodium ethyl sulfate will form around the outer parts of the beaker, encasing the solution in a manner reminiscent of partially frozen ice cubes.

sulfuric acid is the king - 22-2-2017 at 04:47

I use marble as a calcium carbonate source,i hope that's not a problem,i tested it disolving in hydrochloric acid it fully disolved,so it's chloride now.But i don't know if using crushed marble is good,when i added it to solution of sulfuric acid etc. it formed clumps...Maybe couse i was to fast adding it.

JJay - 22-2-2017 at 04:54

That probably won't work well... the calcium carbonate needs to be very finely powdered and added slowly with strong stirring so that it doesn't get coated with nonreactive calcium sulfate. You can use finely powdered chalk if it is actually calcium carbonate. It is also easy to prepare it from calcium chloride and sodium carbonate.

Edit: Oh, crushed marble. That might work, actually. It needs to be very finely powdered.

[Edited on 22-2-2017 by JJay]

sulfuric acid is the king - 22-2-2017 at 14:58

@JJay
This is from the first try that i mentioned before,solution with misterios ultra soluble salt is evaporated by half and now at 10C crystals!
What do you think?

JJay - 22-2-2017 at 15:46

I would check the specific gravity of the solution to be sure, but it looks like that might be sodium ethyl sulfate.

sulfuric acid is the king - 22-2-2017 at 15:46

Nah...It's to easy... :/
I just touched crystals to get them out,poured remained liquid first,and they are liquidy...Hah...Probably that's just some sulphate monohydrate or something...

sulfuric acid is the king - 22-2-2017 at 16:00

What i know so far is that it has no pungent smell like sulphate when i made it by h2so4/nacl methode,substance is nearly without smell,and i think smell is more deep,basic.

sulfuric acid is the king - 22-2-2017 at 16:02

[Edited on 23-2-2017 by sulfuric acid is the king]

JJay - 22-2-2017 at 16:05

That very well could be.

I thought about saving sodium sulfate when it is generated as a side product but so far have always decided it isn't worth it.

sulfuric acid is the king - 10-8-2017 at 05:39

@JJay
Finally i've got some glass..

Tell me,after you refluxed acid and alcohol,and made EtHSO4,you poured it to cold water,and then neutralised with CaCO3,or was it directly after reflux?
Some literature says that first you need to pour it into cold water,then neutralize,but that's nonsense to me,why would you then need high percentage of alcohol and acid...

JJay - 10-8-2017 at 06:27

I did some more research on it, and it looks like the highest yields are possible by just heating equimolar amounts of ethanol and sulfuric acid to 70 C for 10 minutes without any refluxing. The workup I had most success with was pouring the ethylsulfuric acid into water and neutralizing with calcium carbonate then neutralizing with sodium carbonate, boiling off most of the water, filtering, and concentrating on a water bath until it freezes when it cools. I tried neutralizing it directly but then it is a pain to deal with all of the sodium sulfate in its various hydration states, and the process is not incredibly high yielding unless you use oleum. Someone suggested neutralizing directly and extracting with boiling ethanol. I haven't tried that and am not sure if ethanol will dissolve hydrated sodium sulfate or not, and most procedures that call for sodium ethyl sulfate expect it to be in monohydrate form. I'm sure direct neutralization can be gotten to work, though, if you think the calcium carbonate / sodium carbonate route is too much work.

sulfuric acid is the king - 10-8-2017 at 07:34

Thanks.
Can you explain me why is high yield of both (near anhydrous) needed during the reflux when you later pour it directly to water?

JJay - 10-8-2017 at 07:45

The higher the yield, the less impurities you have to eliminate. The starting materials are cheap.

sulfuric acid is the king - 10-8-2017 at 09:19

I know that but i am confused about water.Why ethanol and sulfuric acid need to be near anhydrous during formation of EtHSO4,but then they can be simply poured into water?
Why is water undesirable at first step if you put it later,that confuses me.

JJay - 10-8-2017 at 09:32

There is an equilibrium between ethanol + sulfuric acid and ethyl sulfuric acid + water. The reaction that forms ethyl sulfuric acid is reversible, but hydrolysis is slow. Yet a high concentration of water greatly reduces the rate of formation of ethyl sulfuric acid. A lot of esterification reactions have similar kinetics (see this one, for example: https://www.youtube.com/watch?v=Ah5ds_3s5BI).

sulfuric acid is the king - 10-8-2017 at 10:44

So once formed EtSO4 can be poured into water without any danger of hydrolysis?
So it's only dangerous during the process of formation?
I think that second step (pouring it into the water) ruins the yield.

JJay - 10-8-2017 at 10:55

I've read that it takes days for ethylsulfuric acid to hydrolyze in water at room temperature. I'm not really sure what the purpose of pouring it in water is, though.

This paper discusses the effect of water in some detail: https://books.google.com/books?id=zplJAQAAMAAJ&pg=PA456&...

sulfuric acid is the king - 11-8-2017 at 15:01

Thanks.
I think dilution is probably for less mess

sulfuricacidistheking 2 - 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]

JJay - 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.

sulfuric acid is the king - 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..

JJay - 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.

clearly_not_atara - 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.

JJay - 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.

sulfuric acid is the king - 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

sulfuric acid is the king - 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.

JJay - 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.

Solubility questions.

CaptainMolo - 14-10-2017 at 07:36

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?

JJay - 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.

CaptainMolo - 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]

sulfuric acid is the king - 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?

JJay - 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]

sulfuric acid is the king - 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...

JJay - 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...

JJay - 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.

As has been reported elsewhere, it resembles cauliflower.

[Edited on 18-11-2017 by JJay]

[Edited on 18-11-2017 by JJay]

CaptainMolo - 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.

sulfuric acid is the king - 22-11-2017 at 14:25

@JJay
Great.Thanks for the useful info.
@CaptainMolo
Which method did you use to obtain nitroethane?

nux vomica - 22-11-2017 at 17:26

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

CaptainMolo - 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.

Any success

CalAzeotrope - 6-4-2018 at 15:31

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!

JJay - 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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