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Author: Subject: Sodium Ethyl Sulfate
Wolfram
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[*] posted on 19-11-2003 at 04:47
Sodium Ethyl Sulfate


From http://www.rhodium.ws/chemistry/nitroalkane.html

Two moles of absolute ethanol (92 grams) is slowly dripped into a beaker containing one mole of 20% Oleum (H2SO4 containing 20% SO3), adjusting the rate so that the temperature is maintained at 45°C. When all the ethanol is added, the solution is neutralized with anhydrous sodium carbonate (Na2CO3), care being taken for the evolution of carbon dioxide. Yield 85% of theory.

Q1. Whould the yeld really drop to 0 if I tried with 96% H2SO4 insead. Making Oleum seems little dangerous.

Q2 Could 95% etOH be considered absolute ethanol or should it be dehydrated with CaCl maybee..?
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[*] posted on 19-11-2003 at 06:21


Add 96% H2SO4 to your ethanol, this binds the water but dilutes the H2SO4.

The next step is bubbling SO3 (by decomposition of CuSO4 or else) through the solution ethanol/H2SO4. This will concentrate your H2SO4, thereby consuming water AND making oleum if you add enough SO3.

Working with SO3 always involves certain hard to control hazards. I would avoid glassware for the SO3 production.




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thumbup.gif posted on 19-11-2003 at 07:33


"SO3 (by decomposition of CuSO4 or else)"
HUH? It takes a temperature over a thousand degrees to decompose CaSO4! There's a thread on oleum in this section that lists much easier preparation methods.
"I would avoid glassware for the SO3 production."
Why? I always considered glass the most chemically stable entity available to a home chemist. You mean that SO3 could displace the silicate ion andthe outer layer would turn into sulfates? But that would happen only with a vey thin outer layer.
I think the yield wouldn't drop to 0 if you used concentrated sulfuric acid (let it be CSA from now on) instead of oleum if you use it in excess. With a large excess the yield would still be high (maybe the reaction would be slower).
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[*] posted on 20-11-2003 at 13:58


Any water present will destroy the product. The equilibrium vastly favours acid and alcohol over ester and water.
You can't dry alcohol over CaCl2, it forms a complex (and dissolves).
Glass is innert but it's fragile.
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[*] posted on 19-11-2005 at 06:33


This is an old one to drag up...

But, what would happen if you used sodium pyrosulfate in the acid instead of oleum?

For that matter, why couldn't you use straight sodium pyrosulfate minus the acid?
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[*] posted on 19-11-2005 at 16:12
Better Way?


US3024263 Process for the Preparation of Anhydrous Ethyl Sulfuric Acid

2NaHSO4+EtOH=EtHSO4+Na2SO4+H20

Very simple and the hydrate of the sodium sulfate removes the water preventing hydrolysis.

Anyone ever try this procedure? I'm about to and was looking for input from those with experience.

Patent claims 96% EtOH, bisulfate monohydrate and concentrated or less sulfuric acid may be used!
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[*] posted on 19-11-2005 at 16:49


That, my son, is a gem.
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[*] posted on 19-11-2005 at 16:57


That ain't a gem, thats... well almost as good as a butt naked redhead spreadeagled on the bed waitin for ya.
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[*] posted on 22-11-2005 at 03:53


Usually HEtSO4 is made by mixing 95% H2SO4 with a 30% molar xcess of CuSO4-dehydrated EtOH and 3x xcess (counting on decahydrate, i.e. 0,3 moles per 1 mole of acid) finely ground Na2SO4.

Apply no cooling while adding, let chill at RT to ~40 C, than shake under a stream of cold water beelow 20 C. Let stand for several hrs with occasional shaking.

Titrate by dripping a weighed sample into dil. aq. NaOH until neutral pH (HEtSO4 takes 1 eq. NaOH, unreacted sulfuric - 2 eq's).

Usually gave me a quantative yield.
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[*] posted on 24-11-2005 at 20:20
Aqueous Method


Hey Antoncho perhaps you could elaborate on some ideas you presented long ago in another forum which I happened to save a copy of.

Did you ever develop this as it seems it would have definite advantages over the dry or wet method!

".....but i always wondered if one could improve the yield (maybee dramatically so) by taking advantage of the following gimmicks:

a) To your mixed NaNO2/NaEtSO4 add some EtOH and reflux for some time. Both reagents are slightly soluble in EtOH, so they will probably inter diffuse, thus providing a much larger rxn surface. Then EtOH is just distilled off.

b) Use microwave for (at least - even if the 'microwave effect' will bee of no use here) even and smooth heating. It's been SWIM's practice that someof his nitroalkane (MeNO2, in his case) decomposes/reacts from overheating somewhat on the falsk's walls........"

The dry method which was tried years ago only once by me didn't work out so well due to the scorching/heat transfer problem. Soon after I found a source of the ready made stuff. This was in the days before it was Listed.

Following is an excerpt from the French article translated by somebee which you may be familiar with.

PROCEDURE from the French

"...... One introduces into reaction flask:
137 g of water;
26.5 g (0.0625 X 3 moles) of technical 98% potassium carbonate.
One agitates to ensure dissolution, then adds: 320 g (4.5 moles) of technical 97% sodium nitrite;
6 ml of oleic alcohol, or cétyl-oleic alcohol (antifoaming agents).
Volume charged = 420 ml approximately.

One heats to 130° C, with the glycerin bath, while agitating.

Into the addition funnel, one places a quantity of sodium ethyl sulphate solution containing 444 g (3 moles);
Approximate volume: 750 ml.
Duration of addition 50 to 60 minutes.

The rate of addition is regulated so that the temperature of the mass in reaction remains within the limits indicated: 125 with 130° C. ................."

Any ideas or comments?? regarding this versus Antonchos ethanol solvent idea?
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[*] posted on 25-11-2005 at 08:13


If you are going to consider EtOH as the solvent you will need to reflux if you use microwaves because of the volatility of the solvent. You will need an oven modified for reflux.
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[*] posted on 25-11-2005 at 12:52


I think that the microwaving would really not promote even heating unless it was specially modified with a stirrer and true variable infinite type power control. The somewhat new inverter type micros could do this with a little tweaking for a temp sensor in a closed loop feedback circuit.

Simpler would be to immerse the flask all the way to the neck and lag the takeoff short adaptor. Maybe heat tape under it as well.

Problem is that the bath temp must be too high if only partially immersed and this causes the scorching. It also helps to distill
quicky to minimize any decomposition.
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[*] posted on 3-12-2011 at 10:49


Quote: Originally posted by Wolfram  


When all the ethanol is added, the solution is neutralized with anhydrous sodium carbonate (Na2CO3), care being taken for the evolution of carbon dioxide. Yield 85% of theory.



Once I have anhydrous ethyl sulfate I naturalize it with anhydrous sodium carbonate to get the sodium salt. No byproducts are formed besides carbon dioxide so what I can expect is a moist clumpy mass of sodium ethyl sulfate that will be contaminated with ethyl sulfate and sodium carbonate. I want to use a solvent to help the reaction proceed once the sodium salt begins to clump up. Ethanol will probably work. I can recrystallize from ethanol to purify.

Has anyone attempted this? What can I expect?



[Edited on 4-12-2011 by cipi]
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[*] posted on 15-10-2014 at 08:53
Pyrosulfate


Quote: Originally posted by evil_lurker  
This is an old one to drag up...

But, what would happen if you used sodium pyrosulfate in the acid instead of oleum?

For that matter, why couldn't you use straight sodium pyrosulfate minus the acid?


I am dehydrating some sodium bisulfate to sodium pyrosulfate right now to test this. Finally, this question will be answered!

Update: I now have the pyrosulfate. I am going to powder it and add it to ethanol.

[Edited on 16-10-2014 by Cheddite Cheese]




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[*] posted on 17-10-2014 at 09:31


Update: At room temperature, there is no noticeable reaction, given by removing a portion of the mixture, filtering it, and evaporating.

I have started gently refluxing it. I have 330 g pyrosulfate (~1.5 mols) and am using 175 mL anhydrous ethanol (distilled over CaO).

The refluxing seems to be producing ether as a product, evidenced by smell. I'm running ice water through the condenser to try to condense everything.

Update: Refluxing has been going on for two hours now. I think the pyrosulfate has mostly reacted; I started with ~3 mm diameter chunks, which are now a fine powder (probably sodium sulfate).

Update: Filtering the solids out is taking a very long time.

[Edited on 17-10-2014 by Cheddite Cheese]




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[*] posted on 17-10-2014 at 17:58


I eagerly await the results of your experiment Cheddite! If your reaction does indeed works then it bodes well for using pyrosulfate for other reactions that normally require oleum - such as TNT :D
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[*] posted on 18-10-2014 at 10:39


So, the volume of filtrate has been much less than expected, and the solids are still mushy. Regular filtering is not doing the job. I'm going to try vacuum filtration (I don't really care if some solvent boils off).



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[*] posted on 18-12-2015 at 16:34


How did the vaccum filtration go?
By the way i was wondering, is sodium bisulfate soluble in anhydrous ethanol?




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[*] posted on 2-1-2016 at 00:45


U.S. patent 3024263 discusses an easy, convenient, and high-yielding method for making ethylsulfuric acid by refluxing sodium bisulfate monohydrate in 95% ethanol, exploiting the tendency of sodium sulfate to draw water out of the surrounding liquor to prevent hydrolysis of the product. I've read a few claims on this board and others that this procedure doesn't work, so I figured I'd check it out.

It is easy enough to run the reaction outlined in the patent with a large excess of ethanol in a single reaction vessel. I mixed 200 grams of tech grade 93% sodium bisulfate with 500 mL of anhydrous ethanol and brought it to reflux in a boiling water bath with mechanical stirring. There was no reaction visible as the mixture was brought to reflux in a boiling water bath, and then a pretty vigorous reaction began suddenly with a continuous stream of ethanol flowing from the condenser back into the reaction mixture. I removed the heat from the water bath about 10 minutes after the reaction started, and it continued for perhaps another 20 minutes and then dropped off quickly. The mixture took quite a while to cool, but when it had cooled, I was delighted to see that the sodium bisulfate prills had been replaced by an extremely fine precipitate.

I allowed the mixture to cool and then attempted to filter. The precipitate went right through the filter paper, so I allowed it to settle and separated it by decantation.

Internet reports on the solubilities of various salts of ethylsulfuric acid differ, but I reasoned that Commercial Organic Analysis by Alfred Henry Allen was a likely correct reference. So I measured out 100 grams of sodium carbonate monohydrate for neutralizing the acid, crossed my fingers, and added it to the filtrate. To my disappointment, nothing happened.

I was a bit dismayed, thinking that perhaps the patent's claims really were incorrect. But since sodium carbonate is highly insoluble in alcohol, I reasoned that perhaps catalytic amounts of water were necessary to drive the reaction and added about 10 mL. To my delight, the mixture began foaming and frothed for several minutes, and when the reaction ended, the pH measured 8.6. Slightly basic conditions are good for sodium ethyl sulfate; they prevent it from hydrolyzing in aqueous solutions.

I vacuum filtered the mixture with some difficulty, through celite, and now I have a slightly aqueous ethanol solution of sodium ethyl sulfate. I'm not really sure what the best way to work it up is, but I guess a good starting point is to remove the ethanol and recrystallize from water. Anyone have any better ideas?

[Edited on 2-1-2016 by JJay]
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[*] posted on 2-1-2016 at 02:18


JJay, filtering the acidic methylation agent through paper - nice move. If you just could perform quantitative analysis of the product before and after filtering, then you should have know that working ph for paper is 3-10. The book can be downloaded via Commercial Organic Analysis.
Because titration via barium salt is unreliable and benzidine is carcinogenic, the most convenient way to measure ethylsulfuric content is to hydrolyze the crystalline ethylsulfate with concentrated HCl.
UPD: can't find the link to google books, adding a backup link here

[Edited on 3-1-2016 by byko3y]
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[*] posted on 2-1-2016 at 07:52


Yes, when I was trying this procedure, I too learned the effects on paper. Instead of filtration, settling and decantation worked reasonably well.

A final note on my attempt with the pyrosulfate, which I forgot to update: The solids that I obtained had large amounts of unreacted sulfate/bisulfate (evidenced by precipitation with calcium chloride). It appeared that there was less than 40% sodium ethyl sulfate, but I was unable to get it dry enough to get an accurate mass.




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[*] posted on 2-1-2016 at 12:27


Quote: Originally posted by JJay  
remove the ethanol and recrystallize from water.


Or maybe you've got it backwards and you actually don't want water anywhere near your product unless you want it in solution, IDK really. It's been a while and nothing was weighed and analyzed, but IIRC one plan came together which included a vacuum desiccator and OTC CaCl2 and methanol, giving a product that seemed decent enough and remained a powder in storage.




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[*] posted on 2-1-2016 at 13:32


Quote: Originally posted by S.C. Wack  
Quote: Originally posted by JJay  
remove the ethanol and recrystallize from water.


Or maybe you've got it backwards and you actually don't want water anywhere near your product unless you want it in solution, IDK really. It's been a while and nothing was weighed and analyzed, but IIRC one plan came together which included a vacuum desiccator and OTC CaCl2 and methanol, giving a product that seemed decent enough and remained a powder in storage.


Who knows really... all I know for sure is that there is a tremendous amount of contradictory information on the Internet on this one particular product. I think drying the ethanol solution with sodium sulfate first is a good idea. After that, removing the ethanol should produce anhydrous sodium ethyl sulfate... depending on how soluble it is in ethanol, there may have been some quantity in the second filter cake, but I think it was mostly sodium bicarbonate.

I'll definitely attempt the workup later... while many hobbyists have tried to make sodium ethyl sulfate, most have been unsuccessful, and it looks like it is rarely isolated in pure form by amateurs.
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[*] posted on 2-1-2016 at 13:56


Quote: Originally posted by byko3y  
JJay, filtering the acidic methylation agent through paper - nice move. If you just could perform quantitative analysis of the product before and after filtering, then you should have know that working ph for paper is 3-10. The book can be downloaded via Commercial Organic Analysis.
Because titration via barium salt is unreliable and benzidine is carcinogenic, the most convenient way to measure ethylsulfuric content is to hydrolyze the crystalline ethylsulfate with concentrated HCl.


There was very little methylation agent involved... while there is nearly always some quantity of methanol in ethanol, in this case, the quantity was very small indeed.
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[*] posted on 2-1-2016 at 14:20


Myself with the excess ethanol there and all I'd dry by azeotrope, then toss dry zeolite in the distillate.

Quote: Originally posted by JJay  
while many hobbyists have tried to make sodium ethyl sulfate, most have been unsuccessful, and it looks like it is rarely isolated in pure form by amateurs.


A lot of what gets posted is garbled bumblings, and most success by most members is never mentioned AFAIK. In this particular case the salt is often being made in solution for nitroethane, and there's not much incentive to mention making it. I have a good feeling about the product recrystallized from methanol (or denatured alcohol maybe? Something dry.) and evaporated, and don't remember bringing it up before.

[Edited on 2-1-2016 by S.C. Wack]




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