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Sodium Ethyl Sulfate

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Wolfram - 19-11-2003 at 04:47

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..?

vulture - 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.

Theoretic - 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).

unionised - 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.

evil_lurker - 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?

Better Way?

bio2 - 19-11-2005 at 16:12

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!

doughnut - 19-11-2005 at 16:49

That, my son, is a gem.

evil_lurker - 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.

Antoncho - 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.

Aqueous Method

bio2 - 24-11-2005 at 20:20

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?

CherrieBaby - 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.

bio2 - 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.

cipi - 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]

Pyrosulfate

Metacelsus - 15-10-2014 at 08:53

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]

Metacelsus - 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]

Oscilllator - 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

Metacelsus - 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).

huegene - 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?

JJay - 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]

byko3y - 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]

Metacelsus - 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.

S.C. Wack - 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.

JJay - 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.

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.

JJay - 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.

S.C. Wack - 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]

JJay - 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.

JJay - 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.

S.C. Wack - 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.]

JJay - 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....

byko3y - 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).

JJay - 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.

[Edited on 4-1-2016 by JJay]

byko3y - 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.

JJay - 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.

JJay - 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]

JJay - 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.

S.C. Wack - 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.

JJay - 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]

[Edited on 17-3-2016 by JJay]

JJay - 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.

CaptainPike - 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!

UC235 - 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

JJay - 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....

JJay - 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]

JJay - 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....

JJay - 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]

WGTR - 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.

JJay - 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]

Metacelsus - 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.

WGTR - 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.

S.C. Wack - 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.

JJay - 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.

JJay - 10-4-2016 at 22:29

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

JJay - 15-4-2016 at 20:42

It's definitely sodium ethyl sulfate, and I think 90% is a good estimate of the purity. I did another run, this time neutralizing to pH 8 (instead of 9) and filtering the solution at 500 mL. I got slightly less product, and it appears that slight decomposition took place, so the product is light tan instead of nearly white. But after evaporation of the water, the product is crystalline (consistency of coarse sand). I don't know the purity, but I am guessing that it is higher than 95%.

S.C. Wack - 16-4-2016 at 10:43

What is this assessment based on? BTW barium chloride or recrystallization from alcohol would be a decent minimum effort...BTW2, old CRC Handbook says the sodium salt is a monohydrate, solubility 164 g/100 ml water at 17C. No K hydrate is mentiored and no solubility for the K salt is either, except it is soluble in water and alcohol.

JJay - 16-4-2016 at 11:34

I used an analysis that might be hard to reproduce, so I am looking into other methods of verification.... The melting point of the products have obtained appear to be less than 100C, and I am not exactly sure why. If the substance is heated on a water bath, it will melt and harden as the water is given off.

I am planning to do a recrystalization from methanol / ether.

[Edited on 16-4-2016 by JJay]

JJay - 17-4-2016 at 08:55

The crystaline substance was acidic, and the recrystallization was a disaster... I think it may have contained considerable sodium bisulfate....

It has appeared to me on multiple occasions that sodium ethyl sulfate decomposes in neutral pH solutions, but I have read a few places that it does not, and I can't figure out what could be acidifying it unless it is picking up carbon dioxide from the air.... Anyway, I started another run, and this time I'll aim for pH 9 and won't filter when the solution becomes saturated.

S.C. Wack - 17-4-2016 at 10:28

Desseigne and Giral seem to have no problem with distilling off excess ethanol and some added water from their neutral to slightly alkaline solution. Perhaps their simple procedure could be followed to dryness, if sodium sulfate all separates nicely first and why not throw in some vacuum too.

JJay - 17-4-2016 at 12:47

I wonder if there are impurities in my sulfuric acid that could be causing difficulties. It's one of the better-quality drain cleaners, but it's not distilled and titrates to 95%... I've also observed that with NaEtSO4 created with sodium bisulfate, though....

[Edited on 17-4-2016 by JJay]

JJay - 18-4-2016 at 01:43

Quote: Originally posted by S.C. Wack

Desseigne and Giral seem to have no problem with distilling off excess ethanol and some added water from their neutral to slightly alkaline solution. Perhaps their simple procedure could be followed to dryness, if sodium sulfate all separates nicely first and why not throw in some vacuum too.

Perhaps... I've tried very similar procedures before and found them disappointing, though. A large excess of ethanol could be used to chase the unwanted mineral salts out of solution, but there is still the problem of hydrolosis if the sodium ethyl sulfate solution is not sufficiently alkaline. Also, sources disagree on the temperature and duration of heating for forming ethylsulfuric acid, and I haven't found any sources that specify different melting points for the monohydrate and anhydrous sodium ethyl sulfate.

chemrox - 18-4-2016 at 13:59

Two comments: 1) it's a worthwhile venture as sodium ethyl sulfate goes for $134/25g. 2) this isn't "organic chemistry"

clearly_not_atara - 18-4-2016 at 14:55

I doubt that SES is actually worth $5/g. It's hardly useful for anything, other than selective ethylations of a few polyfunctional compounds. It *costs* $5/g, but what you're really paying for is the ability to order a chemical that isn't produced in large quantities, i.e. you're paying because it's special-made for you, the same way you would also pay (in terms of your time and energy) to make it yourself.

Similarly, McDonald's fries aren't worth what they charge for them -- you can buy frozen French fries a the store for way less. McDonald's fries cost what they do because of how and when they're given to you; you pay for the service of being handed a box of freshly-fried French fries, similar to paying for the service of making a small amount of a useless chemical just for you.

That's why, if you're a professional chemist, you're probably gonna just buy diethyl sulfate for $50/kg and partially hydrolyse it (I bet bicarbonate works). Much easier.

[Edited on 18-4-2016 by clearly_not_atara]

JJay - 18-4-2016 at 15:29

High-grade sodium ethyl sulfate actually is in demand and fetches a premium price since it is used for calibrating some equipment that monitors alcohol consumption... of course, the stuff that the home hobbyist produces in a garage isn't worth as much, but I'm pretty darn sure there is a market for tech grade NaEtSO4.

Also, not everyone wants to handle diethyl sulfate... and it's really not that much easier to make sodium ethyl sulfate from diethyl sulfate as it is from ethanol and sulfuric acid.

[Edited on 18-4-2016 by JJay]

JJay - 18-4-2016 at 15:33

Quote: Originally posted by chemrox

Two comments: 1) it's a worthwhile venture as sodium ethyl sulfate goes for $134/25g. 2) this isn't "organic chemistry"

This is a pretty long-running thread... it's a pretty simple organic compound, but I would have to say that Organic Chemistry is the best forum for it.

Oh and I just found the melting point of the monohydrate! It's 86 C, according to Commercial Organic Analysis.

[Edited on 18-4-2016 by JJay]

JJay - 23-4-2016 at 07:34

I just did another run that yielded 288 grams. pH 9 seems to be a safe level of alkalinity for NaEtSO4. If anyone wants some - it is a fairly crude tech grade but should work well in most synthetic reactions - feel free to message me. I can't say I'll let it go for a super cheap price, and I'm not interested in supplying chemicals to people who are making illegal drugs, but if you have a legitimate use for it, I'm sure I can quote a better price than you are likely to find elsewhere online.

The standardized solution markets for this stuff have completely unreal prices....

JJay - 29-12-2016 at 23:23

I revisited this a couple of months ago, attempting a larger batch, and discovered that upon repeated heating on a water bath with filtering, eventually, it will become much harder to concentrate the solution further, and evaporation at 100 C will not cause the precipitation of further salts. Upon cooling, a very slightly alkaline crystaline substance, melting at 86 C, will form on the sides and bottom of the flask and across the top of the solution. It's hard to know exactly when to stop removing the water... evaporating the water until the solution solidifies completely on cooling produces a similar product that visibly contains impurities in the form of a crust.

[Edited on 30-12-2016 by JJay]

sulfuric acid is the king - 11-1-2017 at 10:51

Quote: Originally posted by JJay

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

No,you are not!

I am ultra inerested in this topic such a long time...
I tried to make NaEtSO4 many many times but i always failed at the major fact - it is soluble in ethanol/EtSO4,right?
I noticed when i neutralize acid with bicarbonate there was feezing,but i always thought that unsoluble part was NaEtSO4.
It was very suspicious that molar mass was very near,there was less product,and it reacted with weak acids like bicarbonate.
Now i know why...couse it was bicarbonate!!!

JJay - 11-1-2017 at 11:12

I'm not exactly sure how soluble it is offhand, but sodium ethyl sulfate is most definitely soluble in ethanol. I had little difficulty dissolving one part of sodium ethyl sulfate in ten parts methanol. If you're trying to prepare it from calcium ethyl sulfate, it's best to neutralize with saturated sodium carbonate solution and go slow so that you have less bicarbonate in solution at the end of the neutralization. Removing excess bicarbonate causes mechanical losses. Also, the carbon dioxide released by neutralizing with bicarbonate greatly increases the solubility of the calcium carbonate.

Although I'm not aware of it being mentioned in any ancient alchemical texts, it is a substance that was undoubtedly researched by alchemists... I think the alchemical term is something like "crystallized soda of wine oil."

Just don't try fusing with it with potash and phlogisticated nitre in a boiling bath of clarified spruce pitch and redistilling the heavy ethers... you might summon a demon or something.

[Edited on 11-1-2017 by JJay]

sulfuric acid is the king - 11-1-2017 at 13:41

I willl make it by neutralizing mix of ethanol/ethylsulfuric acid,then filter off unreacted bicarbonate,and evaporate liquid mix.
Will thermal decomposition occur?

JJay - 11-1-2017 at 13:45

I think so.

sulfuric acid is the king - 11-1-2017 at 16:11

Hmm...So i will need to use vacuum,or some recrystallization...
But i don't have any data on NaEtSO4,solubility table,some physical properties like boiling point etc... etc... It is very difficult to think what to do without some essential data.

JJay - 11-1-2017 at 16:43

The monohydrate melts at 86 C. I'm not sure about the anhydrous salt; while I have seen actual crystals of it, they melted in my filter funnel and then decomposed under heating in neutral conditions.

My observations are that heat and neutral conditions will destroy it in the presence of water, causing it to revert to sodium bisulfate and ethanol.

It's very highly soluble in water... I don't have the exact numbers in front of me, but it takes less than 100 mL of water to dissolve 100 grams of it. Crystals of the pure (or at least very nearly pure) substance will register very slight alkalinity when placed against damp pH paper. It's been stated that the pure monohydrate salt resembles cauliflower; that is a pretty accurate description.

I've attempted to produce it by similarly reacting ethylsulfuric acid with sodium carbonate but wasn't really sure how to recover all of the product or purify it. You could probably extract it with ethanol and then remove it under vacuum, but that's messier than it sounds without a rotovap, especially if you are trying to make more than a tiny quantity.

alking - 11-1-2017 at 18:01

Quote: Originally posted by JJay

The monohydrate melts at 86 C. I'm not sure about the anhydrous salt; while I have seen actual crystals of it, they melted in my filter funnel and then decomposed under heating in neutral conditions.

Did you put your filter into the freezer? Maybe reduce it further if you're not getting many crystals to begin with.

Quote:

I've attempted to produce it by similarly reacting ethylsulfuric acid with sodium carbonate but wasn't really sure how to recover all of the product or purify it. You could probably extract it with ethanol and then remove it under vacuum, but that's messier than it sounds without a rotovap, especially if you are trying to make more than a tiny quantity.

Wouldn't that basically be the same as the bisulfate route? You could probably put it in a dessicant chamber I would think, no need for a rotovape.

JJay - 12-1-2017 at 07:20

Quote: Originally posted by alking

Quote: Originally posted by JJay

The monohydrate melts at 86 C. I'm not sure about the anhydrous salt; while I have seen actual crystals of it, they melted in my filter funnel and then decomposed under heating in neutral conditions.

Did you put your filter into the freezer? Maybe reduce it further if you're not getting many crystals to begin with.

The problem was that the ether was evaporating much faster than the methanol in my filter funnel. Cooling the funnel is an interesting idea, but I think if I'd used a more concentrated solution containing less ether, I would have had an easier time.

Quote:

Wouldn't that basically be the same as the bisulfate route? You could probably put it in a dessicant chamber I would think, no need for a rotovape.

In my experience, the bisulfate route works but doesn't work nearly as well as advertised.

[Edited on 12-1-2017 by JJay]

alking - 12-1-2017 at 11:12

What kind of yields did you get with the bisulfate route? Or do you not know since the product is impure? Sorry if you mentioned it already, I don't want to read the whole thread.

JJay - 12-1-2017 at 11:34

I'm not really sure... I haven't tried it in quite some time, but the yield was nowhere close to the quantitative yield claimed in the patent... maybe 20%?

alking - 12-1-2017 at 13:58

Oh, that is pretty bad. Did you try the sulfuric acid route?

JJay - 12-1-2017 at 14:24

Yeah....

sulfuric acid is the king - 13-1-2017 at 11:17

Quote: Originally posted by JJay

I'm not really sure... I haven't tried it in quite some time, but the yield was nowhere close to the quantitative yield claimed in the patent... maybe 20%?

Hmm,so your picture of crude product is not by bisulfate route?
Very interesting,wich method do you use then?
About recrystalizing...If you just use vacuum to evaporate ethanol,how can it be recrystalization?Or?I am confused little bit,sorry about that

Right now i have no time and equipment,i must plan and improvise,then i will share my results with you,no problem.

[Edited on 13-1-2017 by sulfuric acid is the king]

alking - 13-1-2017 at 13:35

If your ethanol is anhydrous then you can concentrate and freeze it to recrystallize. I don't remember what the bisulfate route is, but I think there is both water in it (in the alcohol I mean) and that it generates water. If so you will have to dry it first. Assuming it works as advertised I would use a drying agent to dry the alcohol, filter off any solids which should remove any sodium sulfates, concentrate until it starts to look cloudy, then cool. That should give you fairly pure crystals.

[Edited on 13-1-2017 by alking]

sulfuric acid is the king - 13-1-2017 at 16:15

Bisulfate route is from the patent.
The legend says even hydrated bisulfate takes water from the alcohol,and makes sulfate decahydrate...

JJay - 13-1-2017 at 19:20

It didn't seem to work as well as reported for me, although it did produce some product. At least one person reported fantastic success with it earlier in this thread. Maybe someone can get it to work?

[Edited on 14-1-2017 by JJay]

sulfuric acid is the king - 14-1-2017 at 05:16

But which method do you use then?

[Edited on 14-1-2017 by sulfuric acid is the king]

alking - 14-1-2017 at 09:28

Do you have sulfuric acid of bisulfate? They both seem to work well enough and the cost is about the same, just do w/e is easiest for you.

sulfuric acid is the king - 14-1-2017 at 15:24

I have both.
Btw i wonder which method JJay use,'couse of materials...

JJay - 14-1-2017 at 15:29

I've tried both. Obviously, sulfuric acid is the king

sulfuric acid is the king - 14-1-2017 at 16:47

Aha,thanks.
But do i need oleum or just concetrated acid,what about ethanol,must be anhydrous or can i with just 95%?
And where can i find that procedure little bit detailed?

JJay - 14-1-2017 at 17:35

I've never tried it with oleum, but that's probably the best way if you can get it and have the means to handle it.

Antocho said he got a quantitative yield of ethylsulfuric acid using this method: http://www.sciencemadness.org/talk/viewthread.php?tid=1113&a...

It's been suggested that sodium ethyl sulfate doesn't hydrolyze under heat in neutral conditions. That doesn't seem to be the case....

The method I've had most luck with is an adaptation of Cohen's synthesis of potassium ethyl sulfate using sodium carbonate instead of potassium carbonate: http://www.prepchem.com/synthesis-of-potassium-ethyl-sulfate...

Use a drying tube and add the sulfuric acid very slowly with a dropping funnel and stirring in a ventilated area. Side products of making ethylsulfuric acid can include diethyl ether and diethyl sulfate if you add the sulfuric acid too fast, so using a fume hood is not a bad idea. Should you mess up, don't spend too much time enjoying the pleasant peppermint smell; that's a potent carcinogen.

I used anhydrous ethanol (it makes a huge difference) at 75% of the volume suggested by Cohen. Also, I heated the flask in a boiling water bath for five hours and evaporated off excess water on a boiling water bath, decanting from the precipitated salts as the water is removed. The solution *must* be alkaline, at least pH 9, when heating is started, or the entire product will be destroyed. I suspect that heating causes some decomposition even under alkaline conditions but not everyone thinks so, and I have not rigorously verified this. You can get a reasonably pure product if you filter all of the salts from a saturated solution of the sodium ethyl sulfate before removing the rest of the water. Crystals taken early on in evaporating the saturated solution will be translucent or perhaps even clear when wet if you use extremely pure reagents.

Oh and sodium ethyl sulfate smells like laundry detergent.

[Edited on 15-1-2017 by JJay]

alking - 14-1-2017 at 19:49

Quote: Originally posted by JJay

It's been suggested that sodium ethyl sulfate doesn't hydrolyze under heat in neutral conditions. That doesn't seem to be the case....

What makes you say this?

JJay - 14-1-2017 at 20:04

Read the thread. I've ended up with acidic product after starting with slightly neutral solution and heating.

sulfuric acid is the king - 15-1-2017 at 09:00

@JJay
Thanks so much.
Is there any method to know when is sulfuric acid totally used?
And what about some analytical method to confirm that i actually obtained right stuff (NaEtSO4)?

[Edited on 15-1-2017 by sulfuric acid is the king]

alking - 15-1-2017 at 09:48

Quote: Originally posted by JJay

Read the thread. I've ended up with acidic product after starting with slightly neutral solution and heating.

Yeah, but you said there was H2O left after you evaporated off your ethanol so that was probably some sodium bisulfate wouldn't you think? I wouldn't conclude that it necessarily hydrolyzed from that alone.

JJay - 15-1-2017 at 09:52

I don't think the sulfuric acid is ever totally used unless the water is removed somehow. I saw a paper on the reaction kinetics a while back which suggested that it's hard to get much above 40% in equillibrium with ethanol and sulfuric acid, although someone did reportedly manage to get 70% starting with completely anhydrous reagents. I don't remember the name of the paper offhand.... Water is a major yield killer as it slows the reaction considerably and pushes the equillibrium in the wrong direction.

You can use the melting point and check other properties like its solubilities. I think you can also check its purity with a GCMS, and it will give off ether if refluxed in sulfuric acid.

[Edited on 15-1-2017 by JJay]

alking - 15-1-2017 at 09:53

Quote: Originally posted by sulfuric acid is the king

@JJay
Thanks so much.
Is there any method to know when is sulfuric acid totally used?
And what about some analytical method to confirm that i actually obtained right stuff (NaEtSO4)?

[Edited on 15-1-2017 by sulfuric acid is the king]

What do you mean, totally used as in you get a quantitative yield on the H2SO4? The oleum route should do it. Supposedly wet EtOH and H2SO4 can be used if you dry the solution with enough Na2SO4 to hold all of the water. For that matter the bisulfate route is supposed to do the same if you use an excess of bisulfate, but JJay reports mixed results on that, I'm not sure the reason.

Analytical wise you could test the melting point (86C according to JJay), the PH (again, slightly alkaline according to JJay), it smells like laundry detergent as he said, and visually if you were to recrystalize it you should see that the crystals are uniform. Likewise if your product fully dissolves in anhydrous alcohol, and does not leave a cloudy suspension, then you can assume that it is the correct product as no other products in either of these reactions has appreciable solubility in alcohols. This should be a straightforward synthesis, just give it a shot.

[Edited on 15-1-2017 by alking]

alking - 15-1-2017 at 09:58

Quote: Originally posted by JJay

I don't think the sulfuric acid is ever totally used unless the water is removed somehow. I saw a paper on the reaction kinetics a while back which suggested that it's hard to get much above 40% in equillibrium with ethanol and sulfuric acid, although someone did reportedly manage to get 70% starting with completely anhydrous reagents. I don't remember the name of the paper offhand.... Water is a major yield killer as it slows the reaction considerably and pushes the equillibrium in the wrong direction.

You can use the melting point and check other properties like its solubilities. I think you can also check its purity with a GCMS, and it will give off ether if refluxed in sulfuric acid.

[Edited on 15-1-2017 by JJay]

I read a paper on the H2SO4, EtOH, and Na2SO4 route that claimed 'quantitative yields.' The oleum paper claimed 86% yields. Antiocio(sp?) also claimed 'near quantitative' yields with the Na2SO4 route. I would go with that. The way you do it according to the paper is to drip in the H2SO4 slowly so that it doesn't get too hot to minimize side products. After it's returned to room temperature you grind your Na2SO4 into a fine a powder as possible to maximize surface area and add it slowly into the solution while it is stirring to evenly disperse the salts. Then you leave it overnight, ~12 hours, or w/e. The Na2SO4 should grab onto all of the water during that time and as it does the reaction will go to completion.

JJay - 15-1-2017 at 10:04

Quote: Originally posted by alking

Quote: Originally posted by JJay

You can get quantitative yields with oleum. The route with Na2SO4 looks nice in theory, and I certainly intend no disrespect towards Antoncho, but call me skeptical. If you'd like to try it and report back, by all means do so.

Also, I have no idea what paper you're talking about.

[Edited on 15-1-2017 by JJay]

alking - 15-1-2017 at 10:54

Here you go. https://www.google.com/patents/US3047604

I might give it a shot if I can find some free time just for curiosities sake, it looks easy enough.

JJay - 15-1-2017 at 13:10

Why would sodium sulfate be more hydrophilic than ethanol, sulfuric acid or even ethyl sulfate? Also, there's the problem of surface area; the water molecules can really only attach to the sodium sulfate on the outside of the sodium sulfate particles, requiring a large excess of desiccant, right?

I'm not saying that I'm 100% sure it can't work, but I was disappointed by the results of testing another of Mr. Leatherman's patents that made similar claims, so I'm not convinced that it actually does work.

alking - 15-1-2017 at 15:17

Quote:

[quote=472024&tid=1113&author=JJay]Why would sodium sulfate be more hydrophilic than ethanol, sulfuric acid or even ethyl sulfate? Also, there's the problem of surface area; the water molecules can really only attach to the sodium sulfate on the outside of the sodium sulfate particles, requiring a large excess of desiccant, right?

I don't know, that's just what the patent claims. Even if EtOH or H2SO4 were more hydrophobic if Na2SO4 is more hydrophobic than HEtSO4 as Na2SO4 grabs some of the water the equilibrium would constantly shift and any free EtOH and H2SO4 would form, further shifting that equilibrium so that Na2SO4 could grab more water and so on until it's shifted all the way to the right. There's likely traces of EtOH and H2SO4 remaining at least, but theoretically it makes sense.

[Edited on 15-1-2017 by alking]

sulfuric acid is the king - 15-1-2017 at 15:20

Simply.When i said totally used i thought reacted to form ethylsulfuric acid,diethyl sulfate,diethyl ether...
Why i ask that?To know what to expect,if i neutralize mix i will get more NaHSO4 contamination etc...

alking - 15-1-2017 at 15:28

Oh, and I forgot to address the later question. That's why you grind it into a powder and add it as a fine stream, because of the surface area issue. They address that in the paper.

alking - 15-1-2017 at 15:32

Quote: Originally posted by sulfuric acid is the king

Simply.When i said totally used i thought reacted to form ethylsulfuric acid,diethyl sulfate,diethyl ether...
Why i ask that?To know what to expect,if i neutralize mix i will get more NaHSO4 contamination etc...

You'll have bisulfate contaminant regardless, but as I said earlier it's easy to filter off. Someone mentioned as well that once your H2O/EtOH is saturated with NaHEtSO4 that all the sulfates crash out (besides the product of course). Filter at that point and you're good. From there dry it and optionally recrystallize it.

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