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Author: Subject: Lab Prep of Diethyl Sulfate
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[*] posted on 10-4-2009 at 11:26


Just a quickie...I think the stuff that gets polished onto auto windshields (to make the water bead), contains diethylsulphate.
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[*] posted on 10-4-2009 at 15:54


Quote: Originally posted by garage chemist  
Fleaker, it's very interesting that you state that EtOH/Et2O + 20% oleum gives diethyl sulfate.
I once did an experiemnt where I reacted neat SO3 with an excess of absolute diethyl ether, and no diethyl sulfate was formed.
I slowly dripped Et2O onto solid SO3 with stirring and cooling.
After everything was homogenous, I poured all of it slowly into cold water, with lots of stirring. There only was a very small organic phase, and it floated to the top. It turned out to be nothing more than residual ether. No diethyl sulfate seemed to have ben formed.

Reading Ullmann on sulfonic acids I found that combining EtOH with SO3 by slowly introducing EtOH into oleum with cooling, ethionic acid (ethyl hydrogensulfate 2-sulfonic acid) is produced. This substance is miscible with water in all proportions, and was likely what I obtained in my Et2O + SO3 experiment.
The reaction proceeds by elimination of water from ethanol, followed by addition of SO3 to the ethylene giving carbyl sulfate (a cyclic sulfonate ester and anhydride) which hydrolyzes to ethionic acid.
http://v3.espacenet.com/publicationDetails/originalDocument?...
http://v3.espacenet.com/publicationDetails/originalDocument?...

How were you able to isolate diethyl sulfate from SO3 and Et2O or EtOH? Is it necessary to vacuum distill the crude mixture? Diluting the crude mixture with water did obviously not work for me.


garage chemist
I can assure you that it does work, having collected the product in quantity at the desired b.p. with reduced pressure distillation. Quenching the reaction mix gives an oily bottom layer that does not do much of anything in acidic cold water--it merely sits there. I tried by adding oleum to ether and by bubbling the ether through the oleum, and (oddly) there is less charring if you bubble the ether from another boiler into the oleum using N2 as the pusher gas. I think your mistake was cooling what should've been a very exothermic reaction. In any case, ether gave much less carbonization than the alcohol did when it came to obtaining product.

As I mentioned, any distillation, or any even moderate temperature increase will cause decomposition, particularly in the case of diethyl sulfate. In some of the poorer yielding runs done with alcohol and oleum, there was a carbonaceous ''cake'' that was stuck in the flask which was difficult to break up with a glass stirring rod. In my case, I isolated it every time by vacuum distillation except when I added ice cubes to a hot EtOH/SO3.H2SO4 mixture. The oleum was in great excess in that case. The methanol and ethanol used were all anhydrous, and the oleum was straight from the bottle.

If you don't believe my results, that is fine, I will not suggest you try to repeat them as it is an unnecessary danger. I didn't characterize my products save by boiling point, solubility in water/hydrolysis test with ammonia and then test with Ba(NO3)2, and only in 2 cases did I end up using product for alkylations. The one occasion that I did use product for alkylations failed miserably for some reason (perhaps too acidic). If you should choose to repeat an experiment, I highly suggest you repeat EtI from dry distillation of sodium sulfate, sodium ethyl sulfate and KI at ~20-50 torr. If ever I experiment with these agents again, that is the experiment I am most interested in doing.

Another thing to mention is that I never saw 2 layers appear ever during the course of a reaction--to get that would require the ice cube treatment, or it would require me distilling out the product.

Another note, when using the oleum, especially when using a 5 fold excess, I have noticed that SO3 can come over into the receiver and then DES or DMS will immediately char upon collecting on the other side. This happened in two cases--both times my distillate came over clear in the condenser but ended up brown when I went to work it up further.

I think making the hydrogen sulfate is preferable to making it all one pot, as the yields are generally BAD because at least half of the product gets turned into crap which is hell to remove from the flask. This happens if there is free sulfuric acid present but can mitigated if a hard enough vacuum is pulled.


EDIT:

Garage chemist,

When you said you tried the distillation with Na2SO4, you reported the 15% yield based off of ethanol, correct? Your yield should be much higher if you based if off the hydrogen ethyl sulfate product. Obtaining that product is difficult as the workup from the bisulfate is lousy and lossy.


[Edited on 10-4-2009 by Fleaker]




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[*] posted on 10-4-2009 at 17:13


In what form was the sodium bisulfate that everyone used for their ethyl hydrogen sulfate preps? Was it really small pellets?

I have done the prep a few times, and got small yields of isolated sodium ethyl sulfate; I believe the reason why was, the sodium bisulfate should be in the form of a fine powder to ensure a higher surface area. Pulverizing your bisulfate may increase yields.




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[*] posted on 10-4-2009 at 19:15


I am not advocating making sodium ethyl sulfate. I am advocating the JACS article procedure which bypasses the isolation of ethyl hydrogen sulfate in any form.

S.C.Wack mentioned the US Navy patent which proceeds from ethyl hydrogen sulfate not the salt.

That patent cites an older patent that I franjly have not looked at.

AFAIK sodium ethyl sulfate is not involved.




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[*] posted on 12-3-2013 at 15:11
Diethyl Sulfate via Vacuum Distillation of Ethyl Sulfuric Acid


I know this is an old thread but I wanted to post my experience to save others, who might be considering attempting the JACS article method some time.

Diethyl sulfate, according to a number of patents and articles, can be made by distilling ethyl sulfuric acid under a vacuum. Ethyl sulfuric acid is added dropwise to hot, anhydrous sodium sulfate. The diethyl sulfate is distilled and the other product of the reaction, sulfuric acid, reacts with the sodium sulfate, forming sodium bisulfate. The 1924 JACS article (available in the first post on this thread) claims fair results by simply mixing ethyl alcohol with conc. sulfuric acid and dripping this mixture on hot, dry sodium sulfate while distilling under a vacuum (20-45 mmHg). My attempts to repeat this process using 95% Everclear and Rooto sulfuric acid were unsuccessful. The reason for this failure is clearly described in the two Navy patents (US 3,047,604 and 3,024,263). One of the patents states the presence of water causes "the decomposition of ethyl sulfuric acid as follows: C2H5HSO4 + H2O <--> C2H5OH + H2SO4." The other explains that "if water is present the ethyl sulfuric acid reacts with the water upon heating and ethyl alcohol distills over as the only recoverable product." This is exactly what I experienced.

The Navy patents recommend stirring but the JACS article doesn't mention it. Based on my experience, stirring is necessary, as attempts without stirring resulted in very low yields.


Experiemtal 1: 50 g of ethyl alcohol (95%) was mixed with 104.5 rooto sulfuric acid (96%). This mixture was stirred for 1 hour before being added dropwise to 90 g of anhydrous sodium sulfate. The apparatus was a 1000 mL flask, equipped with a two-neck adapter for vacuum distillation, with an addition funnel connected to a capillary tube (It was tried without a capillary tube also). The apparatus was evacuated using an Oakton Aspirator Pump with a valve connected to slow the vacuum. The valve was closed and the temp allowed to rise to 150C. Only alcohol was recovered. This method was repeated many times under many temperatures and pressures. It never produced diethyl sulfate, except when the alcohol/sulfuric acid mixture was dried with anhyd. sodium sulfate (about30-40g) using the navy patent method. The best result were obtained using nearly a full vacuum (15-20 mmHg) and a temp of 120-125C.

Experimental 2: 846 g of sodium bisulfate was refluxed in 550 mL of ethanol overnight. The mixture was allowed to cool to room temperature then further cooled in the fridge for two hours. It was filtered and the alcohol removed by distillation. The resulting anhydrous ethyl sulfuric acid (223 g) was carefully dripped onto 133 g of anhydrous sodium sulfate in a two-neck flask equipped for vacuum distillation and an addition funnel. The vacuum was an Oakton Aspirator Pump, 10-15 mmHg). The addition was started at 95C (upon reflection this was too low a temperature) and allowed to rise slowly to 126C. The yield was disappointing (7.75 mL/9 g).

Conclusion: Mixing the sodium sulfate during the addition appears to be necessary. Though a small amount of diethyl sulfate was produced, the attention required and the equipment needed to continue these experiments led me to pursue other options. If I were to continue I would attempt stirring the sodium sulfate (and sodium bisulfate formed) during the addition. The claim made in the JACS article, that this process can be carried out on the lab scale by simply mixing alcohol and sulfuric acid and dripping it on hot sodium sulfate under a vacuum is not true. This was attempted as written and with many adjustments but, in every case, if the ethyl sulfuric acid was not completely anhydrous, no diethyl sulfate was produced.
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[*] posted on 21-3-2013 at 17:11


do you really need diethyl sulphate??? if you are using it as an ethylating agent, which I strongly suspect you are, why not go with ethyl chloride,bromide or iodide?? look up groves process...it doesnt matter what process you use, ultimately you're going to need to do a distillation to get your product, and ethanol,metal halide of your choice, and conc. hydrochloric acid are all you need. also, they are easier to manipulate,freindlier boiling points and nowhere near as fricken poisonous as diethlysulphate......you get that stuff on your bare skin and start panicking....have some conc. aq. ammonia on hand (no punn intended) in case you spill that stuff on your hands-it gets absorbed thru skin REAL quick...in fact it would make an excellent agent for taking someone out....seen films where some poor sucker goes to get in his car,touches the door handle, hey? whats that gooey paste on my fingers? oh, I dont feel too good, I better go to the.......thud....yes, its that bad. so , alkyl halide arent nearly as lethal, unless you drink it or breath it....
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[*] posted on 30-3-2013 at 21:01


Thanks for your concern. I am aware of the risks in dealing with this substance and am taking the necessary precautions, as I would recommend anyone do if they are interested in this process. There are many dangerous substances dealt with on this forum. It's always good to have warnings such as yours though, otherwise people may not realize how dangerous a process is. Thanks again.
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[*] posted on 1-4-2013 at 01:36


I'm not too suprised this doesn't work as advertised, but I'm amazed by what little product you did manage to obtain. My comment stems from personal experience with preparing diMethyl sulfate. I found that even following Painkilla's procedure on here my yield was never more than half of his. I have however managed to collect 20-25 mL over a couple small runs, which is waiting to be re-distilled under high vac. There were several sticking points for me; in the preparation of the methylsulfuric acid, the filtration of the sodium sulfate (Painkilla decanted...) was painfully slow. Like an hour for 100ml or so. During the distillation with further anhydrous sodium sulfate added, it was notable that around the temperature that the dimethyl sulfate distills, the reaction mixture was a homogenous amber-brown coloured solution. Another notable aspect, is the co-distillation of another liquid, immiscible with the heavier oily layer. This confuses me, as at 10 mbar (the pressure I was using), water's boiling point is below 0*C and even if some did condense, I'd expect evaporative cooling of the receiving vessel which was not observed. It is not dense enough to be sulfuric acid, though this has not been eliminated as a possibility. Another notable point is that methylsulfuric acid itself can be distilled if you bring the pressure down to 2 mmHg, so that might be a good place to start.

I intend on going back to this at some point, but other commitments in life prevent me doing so at the current time. Things I intend to investigate will be using excess methanol as a diluent in the preparation of the methylsulfuric acid (easier filtering of less viscous mixture), thoguh thorough removal of the excess alcohol is necessary otherwise dimethyl ether is produced in the reaction mixture at temperature, producing an increase in pressure from the desired 10-40mbar that we'd like.

The use of other dessicants - CuSO4 would be nice as a self indicating one. The use of mechanical stirring is also to be considered, as the magnetic follower may be reducing the supposedly granular sodium sulfate hydrate to a fine powder, complicating the filtration. Finally, distillation of the methylsulfuric acid (high vacuum, something like 0.1mbar) seems like a potentially good way to go. It would provide a solid starting point for the preparation, and shouldn't be too difficult.

Edit: I also posted here: http://www.sciencemadness.org/talk/viewthread.php?tid=9570&p...


[Edited on 1-4-2013 by DJF90]
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[*] posted on 2-4-2013 at 11:05


I look forward to hearing any results. Proper equipment is a must if one even considers this reaction. I wouldn't waste my time unless you have a ground-glass setup, mechanical stirring, strong vacuum and experience. It was foolish for me to even try but I can often say that ;)
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[*] posted on 3-4-2013 at 10:27


Its a good job I won't be wasting my time then...
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[*] posted on 9-12-2013 at 16:35


I know that this thread is about making ethyl sulfate. I became intrigued by this compound when I noted that my 1960 lab manual (Brewster et al) called for ethyl sulfate for making phenetole via a Williamson ether synthesis. There is no warning about the hazards of using ethyl sulfate, but then this manual was published in 1960. It seems that since that time there has been a huge amount of information accumulated on the toxicity and carcinogeneity of many commonly used reagents, eg, benzene, CCl4, etc. About the only toxicity that Brewster warned about was that of cyanides. I don't think the word carcinogen was used in the whole manual.

In reviewing the procedure for phenetole I thought it might be fun to do as it is so simple and produces a nice smelling ether. But I have no ethyl sulfate and after reading this thread really don't want to make it. This brought me to the question asked twice before in this thread but not answered:

Quote: Originally posted by Jor  

Alkyl halides are less nasty. Are there many reactions where alkyl halides don;t work, and dialkyl sulfates do?


Quote: Originally posted by psychronizer  
do you really need diethyl sulphate??? if you are using it as an ethylating agent, which I strongly suspect you are, why not go with ethyl chloride,bromide or iodide??


I can see why Brewster et al would specify ethyl sulfate for 4 reasons:

(1) cost: $21.50/500g (Sigma) vs $36.40/500g (Sigma) for ethyl bromide. EtI is $150/500g (Sigma)

(2) pedagogic reasons as they already had experiments using a bromide and an iodide in the manual.

(3) they weren't aware of (Et)2SO4 toxicity and its easy transmission by skin contact.

(4) they just copied the procedure from some other trusted source.

Boiling points are (Et)2SO4: 209°C (d); EtI: 72°C; EtBr: 38°C. Perhaps in some cases this is important. For making phenetole I don't see where it is.

So, does anyone have any answers for why, on a lab scale, we wouldn't just substitute an ethyl halide for ethyl sulfate - especially in a home setting where we can't easily obtain ethyl sulfate by just getting it from the stockroom?

[Edited on 10-12-2013 by Magpie]




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[*] posted on 10-12-2013 at 00:33


The use of (m)ethyl sulfate is largely driven by cost, especially on scale. Sometimes the halides aren't strong enough electrophiles to undergo a reasonable rate of reaction. Sometimes the sulfates don't cut it either, so the "big guns" have to come out (triflates or oxonium salts) - MeOSO3F was the strongest (aka "magic methyl") but it is now a COSHH banned substance like 2-naphthylamine and benzidine. Another consideration is the "hardness" of the electrophile, especially when you have competing nucleophilic sites (e.g. alkylation of enolates: C- vs O- alkylation).

So the upshot I guess is that you can get away with using an alkyl halide. Ethyl bromide is a little low boiling but it should react relatively fast and can be pepped up with the addition of something like 10mol% NaI. The alkylations are commonly run in acetone with K2CO3, generally at reflux.

I recently tried the alkylation of vanillin with isopropyl bromide (MeCN with K2CO3 and NaI). TLC indicated the formation of a second species but it was very slow even at reflux for 10 hrs. I've seen some people mention that vanillin is tricky to alkylate and that potassiumw carbonate is not a strong enough base, but I used a pKa predictor and its sufficiently strong enough (about 4 units difference IIRC). Maybe I'll give it another go sometime with KOH.
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[*] posted on 21-9-2017 at 17:05


Diethyly sulfate was prepared following the 1924 JACs procedure by Lynn and Shoemaker. My yield was poor at 4.8% based on ethanol. This was 8.5g. The JACs expexted yield was 32.6g. I attribute this discrepancy to poor temperature control and feed rate control on my part. Surely this can be improved.

Shown below is a picture of the oil layer below ethanol in the distillatiion receiving flask.

Edit: %yield would instead be 9.6% as 2 moles of ethanol are required for one mole Et2SO4.

Thanks JJ. This is really an easy prep. I can't see what all the fuss was for upthread.

IMG_2110.JPG - 115kB

[Edited on 22-9-2017 by Magpie]




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[*] posted on 21-9-2017 at 17:48


Very cool. I don't suggest smelling it, but it has a clean, pleasant odor reminiscent of peppermint but sweeter and less pungent.



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[*] posted on 11-10-2017 at 11:21


My second organic synthesis ever was diethylsulfate.

I found a recipe in a 1940s edition of the Encyclopedia Americana which just called for sulfuric acid and ethanol to be distilled under vacuum and tried it out as I had that stuff left over from my first experiment making ether.

It went as described in the text, and I got a two phase product in the receiver that looked a good deal like the result pictured above. The lower phase was slightly viscous. I do not think either phase was, or contained, significant ether as the distillation was under vacuum and the condensers were cooled with tap water.

I just made it to make it and had no use for it at that time so I eventually threw it out. It was not characterized properly by using in a reaction.

I would like to find that reference and give it another shot and properly characterize it and calculate the yield, but given the information on this thread about the hazards of this substance I am reluctant to do so until I can work up at least some sort of half-assed fume hood.

Sure I could put a 2-valve pressure equalizing dropping funnel between the receiving flask and the rest of the rig, which would let me separate the product without opening the apparatus.

And I suppose I could neutralize the residue in the still by cooling it off, cautiously adding ammonia solution to the pot via another pressure equalizing funnel and then heating it to flush the condenser and the funnel I use a receiver with ammonia gas to neutralize residues in there. (the funnel would have both valves closed, isolating the product)

But there'd still be that inevitable moment where I'd have to take that flask off and cap it.

And accidents do happen. Joint leakage, implosion, just dropping the damn stuff.

I have that encyclopedia in storage, and will post the proportions here when I can dig it out.
I think it just called for ethanol, but it might have specified anhydrous. (I had both and it was a very long time ago.)

The yield was certainly well below 50% but I think it was well over 10%. That's the best I can do in describing it without resorting to wishful thinking.





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[*] posted on 12-11-2017 at 18:47


Today I attempted to make the full scale prep of diethyl sulfate using the 1924 procedure of Lynn and Shoemaker. However, a series of equipment problems caused me to abort.

My first mistake was to use a heating mantle instead of the specified oil bath to heat the pot, a 1 liter rbf. My mantle was a fake, made from a flattened 500 ml mantle. As the starting material is dry Na2SO4 heat transfer was terrible.

Second mistake was to use a thermometer that I couldn’t read through the neck wall of the rbf (my TC was broken). So it seemed like I could never get the temperature up to the required 155 - 165 deg C. So I insulated the flask heavily and turned the heat way up high.


Eventually the thermometer bulb broke. It was over its maximum temp so burst due to over pressure in the bulb.

I had put so much heat to the flask the bottom flattened due to melting.

After getting proper equipment I will try again.

[Edited on 13-11-2017 by Magpie]

[Edited on 13-11-2017 by Magpie]

[Edited on 13-11-2017 by Magpie]




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