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Author: Subject: Sodium Ethyl Sulfate
JJay
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[*] posted on 2-1-2016 at 14:55


It looks as though the product recrystalized from ethanol will have at least one molecule of ethanol of crystalization... the commercial form is (according to Commercial Organic Analysis) the monohydrate. I wonder what form it is in if recrystalized from methanol.
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JJay
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[*] posted on 2-1-2016 at 15:17


Quote: Originally posted by S.C. Wack  
Myself with the excess ethanol there and all I'd dry by azeotrope, then toss dry zeolite in the distillate.


Sounds like a good idea.
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[*] posted on 2-1-2016 at 16:59


BTW Perrin/Armarego et al. 7th ed.'s full entry:

Sodium ethylsulfate [546-74-7] M 166.1. Recrystallise it three times from MeOH/Et2O and dry it in a vacuum. [Beilstein 1 H 326, 1 I 164, 1 III 1317, 1 IV 1325.]




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JJay
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[*] posted on 2-1-2016 at 17:45


That won't take long. I guess I'll pick up some ether on my way to the lab....
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[*] posted on 3-1-2016 at 01:34


JJay, alkyl sulfates are extensively researched compounds, because they are surfactants - unlike a regular carboxylic acid salts, alkyl sulfates of calcium and magnesium are water soluble, thus alkyl sulfates retain their surfactant activity in a hard water.
Pretty much the only way of purification (already mentioned) is to recrystallize the alkyl sulfate from alcohol-ether, although sodium sulfate admixture is not a problem because of low solublity in ethanol (0.4-0.5 g/100g), as well as sodium bisulfate is insoluble, provided the alcohol is anhydrous (dehydrated within reaction by Na2SO4). The reason to perform crystallization is to remove traces of inorganic sulfates and salts of other metals.
Quote: Originally posted by JJay  
It looks as though the product recrystalized from ethanol will have at least one molecule of ethanol of crystalization... the commercial form is (according to Commercial Organic Analysis) the monohydrate. I wonder what form it is in if recrystalized from methanol.

Alkyl sulfates are relatively stable at basic ph, they are even more stable in solid basic form.
Ethanol can be removed by vacuum or boiling (not higher than 100 °C).
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JJay
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[*] posted on 4-1-2016 at 04:16


I got delayed yesterday, and right now, I am distilling off the ethanol. I haven't seen any crystals form yet, so it looks like sodium ethyl sulfate may be highly soluble in ethanol. As the solution becomes more concentrated, it is turning pink due to a few drops of phenolphthalein indicator that I added when I was neutralizing the ethylsulfuric acid (the phenolphthalein probably wasn't necessary and hopefully won't be too hard to remove).

Sodium ethyl sulfate is a supposedly rather nonhazardous detergent-like product with a sweet taste, but I'm not going to verify the taste. There is a smell that is reminiscent of laundry detergent near the distillation apparatus.

Update: I've reduced the ethanol solution to around 170 mL, and it appears to be pretty concentrated and is likely anhydrous. I don't have time to finish the work up right now but hopefully will in the next couple of days.

20160104_040501.jpg - 794kB


[Edited on 4-1-2016 by JJay]
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[*] posted on 7-1-2016 at 23:52


I think you gonna need a vacuum if you want to obtain completely ethanol-free crystals, because ethyl sulfate does not seems to form crystals easily, while vapor pressure at low concentrations is proportional to concentration.
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JJay
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[*] posted on 8-1-2016 at 14:10


I have a high-vacuum pump and a water aspirator. I am pretty sure it is possible to drive off nearly all of the ethanol with boiling water temperatures, but recrystalizing from methanol/ether should do it. The question then is whether there is methanol in the crystal lattice.

I am currently working on constructing a new lab and will finish this experiment when it is functional.
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JJay
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[*] posted on 19-1-2016 at 15:43


I haven't finished building my lab yet, but I figured I would start the workup. I distilled off the remaining ethanol and discovered that the solution contained more water than I had expected. So I distilled off the water until crystals started to form, after which I transferred the solution to a dish and warmed it gently to drive off the remaining water. The sodium ethyl sulfate precipitated as flakes.

This is the crude, slightly wet product. I will have to finish building my lab before I do the final purification.

[Edited on 19-1-2016 by JJay]

20160119_142336.jpg - 823kB
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[*] posted on 16-3-2016 at 13:41


Today I heated about 100 grams of crude sodium ethyl sulfate in 450 mL of methanol with stirring. This formed a cloudy suspension... I suspect the contaminants are sodium bicarbonate and sodium sulfate, but I don't really know how much contamination there is, and I'm not really sure how much methanol to use.

This suspension is extremely hard to filter... the suspended particles form a layer on the filter that is almost impermeable, and the filtrate quickly gels in the flask under vacuum.

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[*] posted on 16-3-2016 at 14:54


Sounds like you need more solvent. Hot solvent.

BTW my earlier comments were unrelated to where sodium bisulfate is the source of sulfuric acid; just the textbook route, where step 2 is removal of sulfate with CaCO3. It's a tedious step if you're making a lot, but when done carefully, everything else is simple.




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JJay
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[*] posted on 16-3-2016 at 15:15


I ended up adding about 150 mL more solvent. The impurities are still very hard to filter. I ended up filtering through paper, changing the paper several times and then switching to a glass frit. Some spillage and loss occurred during the many filtrations. The filtrate is still a bit cloudy but is much better than it was.

Update: I added a little diatomaceous earth (actually bed bug killer; I probably shouldn't call it celite) to the fritted filter and swished it around with some methanol, and now the filtrate is coming through it clear, with a slight yellow tinge.

[Edited on 17-3-2016 by JJay]

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[Edited on 17-3-2016 by JJay]
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[*] posted on 16-3-2016 at 19:28


Recrystalization didn't go so great... I've seen crystals several times, but filtering them out of solution is not easy. They tend to disintegrate in the filter funnel. I think it's from the ether evaporating and the remaining methanol melting the crystals.

I'm think I'm going to have to use a more proper ice bath... perhaps ice / calcium chloride hexahydrate... and probably use a higher concentration of ether. Gravity filtration would probably also help keep the ether from evaporating, and I think it would work fine... some of the crystals were more than a quarter inch across.

So it looks like I'll need to get some more ether... and perhaps make some more sodium ethyl sulfate.
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[*] posted on 19-3-2016 at 09:30


Quote: Originally posted by evil_lurker  
That ain't a gem, thats... well almost as good as a butt naked redhead spreadeagled on the bed waitin for ya.


Yeah, um… I've definitely got some chemistry going on now, thank you, evil for that nice visual wake-up call!
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[*] posted on 19-3-2016 at 11:04


I feel like there is an absurd amount of work going on here without having ever looked at prior work on the stuff.

http://www.sciencemadness.org/talk/viewthread.php?tid=15837
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[*] posted on 19-3-2016 at 14:12


I read somewhere that it is a good substance for practicing recrystalizations. Of course, that probably wouldn't be necessary for making nitroethane....
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JJay
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[*] posted on 20-3-2016 at 22:33


...I just realized that Aldrich sells sodium ethyl sulfate for $50/gram, and I just burned through 100 grams of it. Haha!

[Edited on 21-3-2016 by JJay]
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[*] posted on 9-4-2016 at 14:04


Quote: Originally posted by S.C. Wack  
Sounds like you need more solvent. Hot solvent.

BTW my earlier comments were unrelated to where sodium bisulfate is the source of sulfuric acid; just the textbook route, where step 2 is removal of sulfate with CaCO3. It's a tedious step if you're making a lot, but when done carefully, everything else is simple.


I tried a variation on Cohen's textbook procedure for potassium ethyl sulfate (in Practical Organic Chemistry) with 200 mL of 95% sulfuric acid and 600 mL anhydrous ethanol. It was a lot of work, but right now I am looking at about 500 mL of crude, saturated NaEtSO4 solution. It really wasn't that hard, and that's actually consistent with his yields....
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[*] posted on 10-4-2016 at 11:10


It looks like I am the only one interested in this topic....

600 mL anhydrous denatured ethanol was placed in a 1L flask, and it was fitted with a 2-neck Claisen adapter. 200 mL of 95% sulfuric acid was placed in an addition funnel and it was fitted to one neck. A reflux condenser fitted with a drying tube was placed in the other neck. The sulfuric acid was allowed to drip into the ethanol slowly over 12 hours, and then the flask was placed in a water bath, which was slowly heated to boiling over 1h and maintained at boiling temperature for 2.5h. Very little reflux took place. The solution turned yellow, then orange, then dark red, then almost black (it is thought likely that this occurred due to decomposition of the denaturant). The contents of the flask were cooled slightly and poured into 3.5L water in a polypropylene bucket. Freshly precipitated calcium carbonate was added in small portions until no further reaction was evident, and then the mixture was allowed to settle and decanted from the precipitate. 1L water was mixed with the precipitate, the mixture was allowed to settle, and the liquid was decanted and the supernaturants were combined. Most of the red-orange impurity stuck to the sides of the bucket or to the precipitate, leaving only a yellow color. Sodium carbonate was added in small portions until the solution measured pH 9, after which the mixture was filtered and the filtrate concentrated to 1L in a PTFE-coated pan on a hot plate, then decanted into a 1L beaker. The solution was further concentrated on a water bath until liquid withdrawn on a small spatula immediately hardened at room temperature. Yield: 275 grams. Purity is not known; assuming 90% purity, this represents a yield of about 40% on sulfuric acid.


[Edited on 11-4-2016 by JJay]
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[*] posted on 10-4-2016 at 18:40


I'm personally interested in what you're doing, although I haven't tried making sodium ethyl sulfate yet. It's a useful intermediate to diethyl sulfate, which I may be making (if I feel that I can do this safely), on my quest for a total synthesis of an ionic liquid.



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[*] posted on 10-4-2016 at 19:45


Sodium ethyl sulfate is also useful for making diethyl sulfide and as a mild alkylating reagent. And some people have used it for making nitroethane, which is hard to buy and might have some legitimate uses.

Sodium ethyl sulfate is also expensive... the cheapest I have found is $4/gram for 95% purity... analytical grades cost in excess of $50/gram.

[Edited on 11-4-2016 by JJay]
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[*] posted on 10-4-2016 at 19:47


I once considered making diethyl sulfate in much the same way, and decided against it. For me, it's too much of a risk. It's safe if everything goes right, but very bad if an accident happens.



As below, so above.
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[*] posted on 10-4-2016 at 20:29


Yeah, I've read through all of the posts here on DES, as well as from several other sources. I'd be using it as an alkylating agent to make something like 1,3-diethyl imidazolium hydrogen sulfate. The fact that it's such a strong alkylating agent means that it's good at doing other things too, like growing an extra arm out the back of your head. Naturally, if I worked with it at all, I'd be working in a fume hood.

If sodium ethyl sulfate is a strong enough alkylating agent as is, that would certainly be a plus, as I'd be hydrolyzing the ethyl sulfate anion to hydrogen sulfate after akylation anyway. I haven't seen this approach used in the literature so far, though. I guess it wouldn't hurt to try it and see (famous last words). Methyl hydrogen sulfate is probably a much stronger alkylating agent than the corresponding ethyl one, and might be a better one to start with.

There are some posts where I've seen people heating sodium ethyl sulfate strongly enough to make DES, apparently without realizing what they had just made. I couldn't help but cringe a bit. In any case, I'm hoping to craft a one-pot reaction, that doesn't involve moving carcinogenic materials from one piece of glassware to another.





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[*] posted on 10-4-2016 at 21:24


Denatured alcohol is not always recommendable; at hardware stores here this is over 50% methanol.

The method I used was similar to if not from Mann and Saunders (great, underappreciated book), where 12 g. CaCO3 is added over 20 minutes. This is what I mean by tedious, and it is necessary. They recommend a fine sieve; a flour sifter is handy.

Cohen adds "chalk ground into a thin paste with water." and it is unclear if this is better.




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[*] posted on 10-4-2016 at 21:37


The denatured alcohol that I used contained no methanol, but I'd probably suggest using non-denatured if you have any available. For the neutralization, I used a freshly precipitated paste made by mixing sodium bicarbonate and calcium chloride solutions in a 20L bucket. The supernaturant was decanted from the precipitate, which was washed 3x with an abundance of water and allowed to settle to jelly-like suspension. I have read that when care is not taken in neutralizing with calcium carbonate, up to a 10x excess may be required, but the freshly precipitated chalk reacted pretty close to quantitatively when added in small portions, stirring with a PTFE-coated rod.
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