Sciencemadness Discussion Board - A attempt at Halfapint's dimethylsulfate synthesis..

A attempt at Halfapint's dimethylsulfate synthesis..

Klute - 2-12-2007 at 06:20

Discussion

Alkylating agents are key reagents in organic chemistry, reacting with O, S, and N to form corresponding alkylated compounds ( Ethers, Thiols, alkylamines, quaternary ammonium salts. Most of alkylating agents are quite noxious, and the safer ones are often less efficient. Methylating agents are often particulary usefull, being often easier to make and more reactive.
Although I already have some iodomethane at hand, the low vapor pressure makes this toxic compound a delicate one to handle.
Dimethylsulfate, although being at least as toxic as MeI if not more, has a much higher boiling point, and when sufficiently cooled, is more practical to manipulate. It still remains a dangerous reagent, and schlenk glassware, needles and septums are more than recommended when handling it. It is also a very efficient and reactive methylating agent, and is very often used in the litterature when a methylation is performed.
Conventional syntheses of this compound require oleum, pur SO3, SOCl2 or other noxious and difficult to make/obtain reagents. But, according to a russian article Foxy2 retreived some time ago at the Hive, DMS can be made by passing dry SO2 in MeOH containg anhydrous CuCl2. The way I see it the CuCl2 oxidizes the SO2 to SO2Cl2 in-situ, which immediatly reacts with the MeOH forming DMS:

SO2 + 2 CuCl2 ---> SO2Cl2 + 2CuCl
SO2Cl2 + 2 MeOH --> (MeO)2SO2 + 2 HCl

--> SO2 + 2CuCl2 + 2MeOH --> (MeO)SO2 + 2CuCl + 2HCl.

Apparezntly, when doing this in aqueous MeoH, mehtylsulfuric acid is produced.
Halfapint tried this procedure with apparently very conclusive results, although he only gave vague descrption. He claimed 2/3 of his MeOH turned into DMS, and then claimed 78% yields from the CuCl2, but as Antibody pointed out, the numbers don't work out, meaning he used a ridiculous amount of MeOH, which would never be enough to even cover the CuCl2. I decided I would explore this by myself.

Trimethylphosphate is also a efficient methylating agent, and apparently much safer to handle, but it syntheis requires either POCl3, which I would like to keep for other uses, or P2O5, which still requires pyrolysis to obtain ~50% yields. Compared to the materials needed for the DMS synth, I would like to keep the P2O5. Of course if the synthesis doesn't work as well as I would like it to, the TMP synthesis will performed.

Dupont published a bulletin on DMS handling and safety, I can only recommend anyone that plan on handling sucha reagent to read it thorouhgly: DMS Dupont bulletin

I uploaded the thread were Halfapint reported his results

Synthesis

Preparation of CuCl2 from copper metal

47g (0.74mol) of copper wire is denuded and cut into 1-5mm pieces with plyers.
250mL of 53% HNO3 are placed in a 500mL beaker with a stir bar, on a hotplate-stirrer outside, with a fan venting the top of the beaker. The copper is gradually added to the acid, bubbling starts immediatly, and the initially water-clear acid turns blue. Red fumes of NO2 start to appear very quickly, the beaker heats up to ~50°C. The copper is added in portions, to maintain a vigorous but controlled reaction. The addtion takes 30-40min, and calms down after 1 hours total reaction time, with a small amount of copper left. The beaker is then heated to 60°c for an hour.
The turquoise-blue solution is then left to cool down, and transfered to a 1000mL beaker, decanting the copper powder left. Roughly 400mL of 15% NaOH is slowly added, while vigorously stirring with a glass rod to break the light blue gelatinous Cu(OH)2 precipate. Once no more solid forms and pH is basic, the thick blue slurry is placed on the hotplate and heated to 70°c for 30 min. All of the light blue Cu(OH)2 is turned into the fine black CuO solid. The beaker is left to cool, the CuO settles down, and the clear supernatant is decanted. An equal volume of dH2O is added, the suspension stirred for 5 min, and left to settle, then the supernatant is again decanted. These operations are repeated another 3 times. The supernatant is still basic after the washes. The thick suspension is then gravity filtered (1H) to obtain a thick wet paste.

In a 1L beaker, 250mL of 30% HCl are placed with a stir bar, and then CuO paste gradually added with good stirring, using some dH2O to wash the sides now and then. It dissolved into a emeraude green solution. More HCl is added when the oxide seems to not dissolve anymore. After 1 Hours, a dark emeraude green solution was obtained.

This was filtered into a 1L RBF, and the HCl distilled using a NaOH wash bottle to neutralise the fumes. Once ~200mL were left, the hot solution was quickly transfered to a beaker, covered with cello to avoid the fumes, and left to cool. The clear green solid mass is then vacuum filtered, the green crystals being squeezed with cello on the buchner to get most of the solution out, and transfered to jointed erlenmeyer. The filtrate being transfered to a 250mL flask, and that's when I noticed that adding dH2O to a small amount of the green crystals turned the solution BLUE. I can only guess that the dimineralised water contains either sulfates or nitrates, something that turns the CuCl2 to another blue salt. So some CaCl2 was added to the filtrate to compensate for all the dH2O added during the workup. This meant that a certain amount of either CuSO4 or Cu(NO3)2 was present in the first crop of crystals. After a few minutes, the suspension was filtered into a 250mL RBF, and the remaining HCl distilled off. The dark green almost black thick residu was quickly transfered to a beaker where is immeditaly setted to a hard solid mass. An attempt at drying the first crop of crystals by applying vacuum and heating the erlenmeyer was made, but too much water was present, and refluxed on the sides, so the whole mass was placed in pyrex crystalizer and heated on the hot plate. After a few minutes, the green crystals started to take a bluish tint, and then turned to a dry light brown amorphous solid. The residu of the filtrate was broken down with a screw driver, the black solid tuning green after a few seconds of contact with the humid atmosphere. No blue color was noticed. The lumps were crushed with a pestle and mortar, and dried on a hotplate like the first crop. Afetr grinding the thoroughly dried solid, 100g of a very fine light brown powder was obtained, and immeditaly bottled.
A certain amount of another Cu2+ salt was present, but I had no means of knowing how much. I simply considered it to be 90% CuCl2. In any case, the other salts would make either H2SO4 or HNO3 during the reaction, hoping that wouldn't interfere too much.
Of course, this was alot of work to produce 100g of dry CuCl2. Next time, if ever, it will be bought.

Oxydative sulfuration of methanol with SO2/CuCl2

A 500mL 3-neck RBF equipped with: a gas admission tube shaped as a funnel at the end, a thermometer, a condenser with a gas outlet attached at the top, itself connected to a 5% NH4OH solution with a spatula tip of phenolphtaleine, and magnetic stirring, was charged with 175mL of dry MeOH, followed by the CuCl2 powder, added in portions with steady stirring. As soon as the brown powder was added, the MeOH turned to a dark green color. Alot of solid stayed in suspension. Considering the CuCl2 powder as being 90% pur, this represented 0.67mol of CuCl2.
A 250mL 3-neck RBF with magnetic stirring and a gas outlet was charged with 120mL 50% H2SO4, and a additon funnel was attached. The gas oulet was connected to a conc. H2SO4 wash bottle, itself connected to the admission tube of the reaction flask.
A solution of 39.2g of 97% sodium metabisulfite Na2S2O5 (0.2mol, giving 0.40mol SO2 gas, or 1.2 excess to the CuCl2) in 100mL dH2O was prepared in a erlenmeyer, and half of it charged into the addition funnel.
Once all the setup had been secured, checked and double-checked, the admission tube was lowered just above the stir bar, the reaction flask was immerged in a 40°c water bath, and the bisulfite solution slowly dripped in, at a rate that minimized the amount of bubbles at the exit wash bottle. The gassing took over 6 hours. It was very fast at first, 1/4 of the solution had been added afetr 1H, but then the quantity of SO2 per drop of bisulfite increased, surely because the H2SO4 solution was fully saturated. Very little SO2 escaped at the exit bottle. After 4h, the traces of CuCl2 powder in the admission tube from level variation clearly turned white. The whole reaction medium became darker, it seemed that most of solid dissolved at one point, but the very dark color of the medium made it hard to see. After 6h, the medium was a dark, viscous liquid. 15g of metabisulfite in 30mL dH2O was added at the end, and the dripping rate increased, to give 2 bubbles of SO2 per second in the H2SO4 wash bottle. Nearly all of it was taken up after a few minutes of pressure increase.
Once the addition was finished, the admission tube was redrawn above the surface of the medium, the exit wash bottle was emptied until the tube was just under the surface of the ammonium hydroxide solution to avoid suck back, and a stopper on the acid flask slightly opened to avoid over- or underpressure, and I went to sleep.
6h later, the suspension is now going to be carefully filtered, the excess MeOH removed, and -hopefully- the DMS fractionnated.

More news later.

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12AX7 - 2-12-2007 at 10:26

Alternate preparation for CuCl2:

a). Cu(I) oxide is obtained by anodic oxidation in a strong chloride solution, then oxidized to Cu(II). b). Alternately, a sulfate solution can be used, giving Cu(II) hydroxide directly.

a). 2Cu + 2OH- ---[CuCl4(3-)]--> Cu2O + H2O + 2e-
Cu2O + 4HCl + H2O2 --> 2CuCl2 + 3H2O
b). Cu + 2OH- ---[Cu(H2O)6(2+)]--> Cu(OH)2 + 2e-
Cu(OH)2 ---(Delta)--> CuO + H2O
CuO + 2HCl = H2O + CuCl2

Connect bulk copper metal (wire, plate, bar, etc.) to the anode of a ~2V power supply. Use any metal for the cathode; steel sheet is sufficient. Arrange the electrodes a few inches apart in a large beaker containing concentrated sodium chloride solution (>20% by weight). Place the cathode towards the surface. The anode can hang at any height. When current is applied, a yellow precipitate is produced, coloidial Cu2O, turning deeper orange as the process continues. Copper metal may deposit on the cathode as sponge, which should be cleared from the electrode. Thiourea might help prevent formation of a loose sponge, facilitating production.

When the desired amount of copper has been oxidized (which should be proportional to the charge applied), the Cu2O is washed (decanting may be easier than filtering as the particles are small) and dissolved in concentrated HCl, giving a brown solution. H2O2 (30%) is added until the color returns to a clear green color. The solution is evaporated, crystallized (as CuCl2.2H2O) and dehydrated (preferrably in a stream of hot, dry HCl), giving CuCl2(anh.).

Instead of H2O2, chlorine gas could be bubbled in, but hypochlorite nor any other chlorate cannot be used in solution, as that would leave ions behind.

Alternately, the Cu2O can be dried and roasted in air above 400°C, giving a dark oxide near CuO. This can be dissolved in acid directly with little need for an oxidizer.

Copper metal can also be dissolved directly in HCl, with the aid of an air bubbler. The process is slower than electrolysis or a more powerful oxidizer.

b). Copper is anodically oxidized in a solution of a sulfate. Sodium and ammonium sulfate are suitable. The cell should be operated at a high voltage, driving up the temperature (90-100C), so that the Cu(OH)2 produced is immediately dehydrated to black CuO. The CuO is filtered, washed and dissolved in HCl, giving a solution of CuCl2.

---

I do love copper, and I'm confused why anyone would bother wasting something as precious as raw fuming nitric acid on something so trivial. Sure, it needs to be oxidized, but not by much, and there are easier ways to go about it.

Tim

Klute - 2-12-2007 at 12:02

Thanks for the advice 12AX7, but I'll simply buy some next time. It's dirt cheap. I just didn't want to wait for it to arrive.
BTW, 53% isn't fuming nitric, it's totally clear, and costs 4E a liter at the local hardware store

Concerning the synthesis, I've ran into a problem.
An attempt at transfering the dark mixture with a canula to a 500mL still was fruitless as the canula got clogged up even though the solution had completly decanted. So, using as much precautions as possible (Clothing, googles, rubber gloves under neoprene gloves, gas mask for organic and inorganic vapors, strong draught, etc), the dark green/black mixture was decanted, the settled solids washed with 2x50mL MeOH, opening the flask as little as possible. Everything comming in contact with the fluid was immeditaly and thoroughly washed with 5% NH4OH which took a blue color from the Cu salts.
The MeOH was then distilled with a short vigreux, with a 10% NaOH wash bottle follwed by a 5% NH3 wash bottled connected to the vacuum inlet. A decent amount of bubbling accured until the vapors reached the condenser; then calmed down. At one point, when ~120mL of MeOh had been recovered, the thick syrupy fluid started boiling pretty vigorously, even when the mantle was removed and a cold water bath applied, alot of bubbling accured, and a white fog appeared in the NH3 wash bottle (HCl?). Once everyhting had cooled down, the receiver was detached, a brownish fog appearing in the flask. It was transfered to a brown bottle for neutralisation.
The wash bottles were deconnected, and vacuum applied, the MeOH distilled, and the residue started getting thicker and thicker, and then suddenly set into a thick greenish mass, with no more apparent liquid inside.

I'm thinking of adding some wtaer to try and dissolve the salts and see if any organic seperates. When water was added to the MeOH recovere dunder vacuum (~20mL), a very small amount of orgnaic liquid seemed to seperate, adding 5% NH4OH caused a white fog, and some heating up.

Halfapint simply mentions that a skin had appeared when most of the MeOH had been removed. I'm not sure what could have gone wrong.. Too much MeOH used?

I'm not sure I'm going to pursue in this direction, I'd rather try the TMP synth. I must admit I was scared shitless during the transfert and opening the distn setup, even with my protection. I took two showers, and changed gloves very regularily.

If anyone has suggestions, please speak up. I'm scared the CuCl will not dissolve in water, and I'm not sure if HCl would hydrolyse and formed DMS.

I guess making SO2Cl2 before hand an reacting that with MeOH might be more efficient, but that kind of hardcore chemistry isn't my thing. The activated carbon/Cl2/SO2 could be worth a try.

EDIT: Adding water caused some of the mass to dissolve in a greenish-blue solution. Very few oily droplets floated on the surface, probably some grease from the joints. No orgnaic layer athe bottom whatsoever. I neutralised everything with ammonia, chucked the blue washes in the dump jerrican, and all the setup in the base bath. I'm now scrubbing evrything out. This is something I'm not doing again. :mad:

I would like to know what failed though, and how come halfapint and antibody manadged to distill anything once the MeOH has been removed.
Next week I'm taking the P2O5 out.

[Edited on 2-12-2007 by Klute]

12AX7 - 2-12-2007 at 19:09

Oh, something else I was going to add: to the best of my knowledge, copper is blue in dilute solution. I don't know that sulfate is much of a ligand; chloride isn't either, but better nonetheless it seems. Any copper salt added to a strong chloride solution (HCl, NaCl, etc.) turns green, while any green copper solution (e.g., concentrated CuCl2) will turn blue on dilution. CuCl2.2H2O is also light blue when dry, a quite remarkable transformation from the deep green solution that produced it, and an even more dramatic change from CuCl2(anh.).

Nitric at the hardware store, BASTARD!

Tim

garage chemist - 2-12-2007 at 22:42

Very interesting work Klute.
One thing that I see that could have been the reason for the lack of product is the fact that DMS can't be distilled at atmospheric pressure, it decomposes. The boiling point of 188°C at atmospheric pressure is extrapolated, it cant be measured directly because decomposition is rapid.
The next time, use vacuum!

The separate phase that you said occured upon addition of water to the MeOH fraction might well be your product. Did you save it?

SO2Cl2 is not that bad, it hydrolyses very slowly, in sharp contrast to SOCl2. I'd consider its use as an alternative method.
However, the initial reaction of SO2Cl2 with alcohols (under formation of alkyl chlorosulfate) is VERY violent, and alkyl chlorosulfates are very powerful lachrymators.
Until now I have not been able to find out if the alkyl chlorosulfates are even able to further react with alcohols upon reflux. Ullmann gives an old method of dialkyl sulfate synthesis that consists in heating the mentioned alkyl chlorosulfates with dialkyl sulfites (from SOCl2 and the alcohol). He doesnt mention any reaction of alkyl chlorosulfates with alcohols- I take that as a bad sign, but experimentation concerning this is really needed.

not_important - 3-12-2007 at 01:24

Suter's Organic Chemistry of Sulfur reports on this reaction, but states that it gives "very small yields" It also states that alkyl chlorosulfate reaction with alcohols to give the alkyl chloride and the sulfonic acid.

Klute - 3-12-2007 at 12:45

Quote:

Originally posted by garage chemist
One thing that I see that could have been the reason for the lack of product is the fact that DMS can't be distilled at atmospheric pressure, it decomposes. The boiling point of 188°C at atmospheric pressure is extrapolated, it cant be measured directly because decomposition is rapid.
The next time, use vacuum!

Yes, vacuum was used after removing most of the MeOH, 20mmHg, and the residu turned to a solid in a matter of seconds..

In any case, this project will put aside for the moment... it's a pity it din't work out well, at least two people apparently reported sucess..

Garage chemsit, have you tried making SO2Cl2 from SO2 and Cl2? If the reaction is practical, i might give it a try..

I recall a patent or two on making dialkylsufates from SO2Cl2 and alcohols.. I'll dig the numbers out when I"m at home..

garage chemist - 4-12-2007 at 20:55

I read up on CuCl2, and found that simply heating the hydrate doesnt make the anhydrous salt due to hydrolysis- so you either had CuCl2 with residual water of hydration, or basic copper chloride. This could be a reason for the failure.
Dehydration of hydrated CuCl2 must be done in a stream of HCl gas. So if you can buy anhydrous CuCl2, you should definately do that.

Could it be possible to use a different Cu(II) salt instead of CuCl2, like anhydrous CuSO4, which is produced cleanly and without hydrolysis by heating the hydrate?

Also, SO3 (that the Cu(II) oxidation of SO2 is supposed to produce in situ) will only form methylsulfuric acid with methanol if no excess of SO3 is used. This decomposes to dimethyl sulfate and H2SO4 upon distillation in vacuum, but this certainly needs more time and heat than distillation of plain DMS.

I have never tried to make SO2Cl2 myself, but I plan on doing so at some time, and a user here (Natures Natrium) has done it and reported about it. He used the charcoal catalysed method (active charcoal packed in a liebig condenser with flowing water to remove the heat of reaction).

I have 100ml of SO2Cl2 here that I bought once, it was expensive though (more expensive than SOCl2, of which I have more than 250ml).
I also have 200ml DMS, but as you can see I am interested in its synthesis as well because I cant buy more of it.

It might be possible to force the reaction of methyl chlorosulfate with methanol by using a basic condensation reagent, like sodium methanolate. Natures Natrium tried to do this using lithium methanolate, but abandoned the reaction because the lachrymatory methyl chlorosulfate fumes forced him to stop (which shows that his setup was completely unsuitable for the preparation of something as dangerous as DMS- if he had used proper technique, he would never even have noticed that a lachrymator was being produced as the intermediate).

[Edited on 5-12-2007 by garage chemist]

Klute - 5-12-2007 at 08:53

Good remark on the CuCl2, now that you mention it, I've never read some reliable facts claiming that it can be deshydrated by simply heating. But, the light brown solid that slowly formed during heating returned to a light green very quickly if powdered and left a few minutes in humid atmosphere, and I don't recall and Cl2/other nasty fumes during the drying, I only weared the gas mask during the gassing (although there wasn't any SO2 smell, to wich I react very badly too, it makes me cough very badly so I didn't want to take chances) and the work up. I just considered it to be the anhydrous form. And there definatively was a white solid forming in the gas admission tube during gassing from the green deposits. But the whole reaction mixture didn't change aspect that much, just a thick dark suspension.
The CuCl2 that I can buy will surely be the hydrated form. ANd using CuSO4 would form H2SO4 instead of HCL, wich is normally driven out during the reaction. Swapping the two might make the recation slower. But I don't think having H2SO4 around during the distillation(/pyrolysis?) would damadge the DMS much, as it theorically can be made by distilling MeOH/H2SO4.

EDIT: Purification of Laboratory Chemical 3rd Ed., on CuCl2: "dehydrated by heating on a steam-bath under vacuum." The first portions of the brown solid I obtained was by using the vacuum. As there was alot of water genrated, I prefered conbtinuing under atmospheric pressure, and the product didn't change appearance that much..

US1641005 on diethylsulfate (and other dialkylsufates) from corresponding alcohol and sulfuryl chloride, generously reovered by PolytheneSam:

[Edited on 5-12-2007 by Klute]

Attachment: R2SO4_ROH_SO2Cl2_US1641005.pdf (147kB)
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garage chemist - 7-12-2007 at 06:28

Hm, the patent is interesting. They use two times the theoretical amount of sulfuryl chloride and separate the excess after the reaction by distillation, and the crude dialkyl sulfate is then washed with cold water.
I wonder how it comes that this doesnt give the alkyl chlorosulfate instead of the desired dialkyl sulfate?

The preparation of methyl chlorosulfate adds the alcohol to the SO2Cl2, btw.

Methyl chlorosulfate boils at 42°C (16mmHg) and with decomposition at 133°C (atmospheric pressure).
The smallest concentration that causes irritation is 2mg per cubic meter, and it becomes unbearable at 30-40mg per cubic meter, making it a very powerful irritant (as a comparison, chloroacetone becomes noticeable at 18 mg/m3 and unbearable at 100 mg/m3). It is also somewhat less toxic than dimethyl sulfate.

Sartoris Book "The war gases" gives a preparation of DMS from methanol and chlorosulfonic acid- ClSO3H and MeOH react cleanly to methylsulfuric acid and HCl, and the former gives DMS upon distillation in vacuum.
So if you can get or make chlorosulfonic acid (without oleum), you're set.

I am still doubtful whether the patent you posted actually works. It doesnt give a preparation example, which is a bad sign. We all know about bogus patents, like the hydrazine synthesis from nickel and urea, or the DPPP synthesis, and they all did not have any synthesis examples and were written in a similar style as your DMS patent.

I could do some simple experiments with SO2Cl2 and MeOH this weekend that could shed some light onto how sulfuryl chloride actually interacts with alcohols.

One last idea, could the chlorosulfonic acid possibly be replaced by a combination of SO2Cl2 and H2SO4?

Klute - 7-12-2007 at 11:04

Chlorosulfonic acid is a very usefull reagent I unfortunaly cannot acces, though I'd would have alot to do with it...
Could one obtain it from methylchlorosulfonate? I guess a acidic hydrolysis wouldn't work as with traditional alkyl esters seeing the reactivity..

Concerning the patent, indeed I always take this kind of information with a rather large grain of salt... The litterature on this kind of reaction seems scarce too, those two points make it kind of dubious..
If you could try a few tests, that would really be appreciated. If these tests tend to suggests the recation works, I will surely give the SO2CL2 preperation a try. Otherwise, considering the negative aspects, I'll spend my time on something else.
Thanks again for your contribution.

garage chemist - 3-1-2008 at 01:38

Well, it turns out that DMS can be obtained simply by rapidly distilling 1 part (by weight) MeOH with 8-10 parts concentrated H2SO4 at atmospheric pressure.
I found this procedure here:
Organische Chemie by Kolbe, from 1854, page 249

Obviously, one would distill until only the excess of conc H2SO4 remains in the distilling flask.
The distillate separates into two phases, the upper one is discarded and the lower one (crude DMS) washed with water, dried with CaCl2 and redistilled over anhydrous BaO (can probably be dispensed with if the last distillation is done in vacuum).

Such preparations are the beauty of really old chemistry books.
You wont find such procedures anywhere in even the largest of todays chemistry lexica.

Marvin had found out about this process about more than three years ago:
https://sciencemadness.org/talk/viewthread.php?tid=1608&...
but it seemed to have gone unnoticed by most.

Apparently, DMS distills over largely unchanged if distilled rapidly at atmospheric pressure as in this preparation.
The preparation doesnt even mention vacuum at any stage, a sign that the decomposition might even be quite insignificant.

It also talks about reacting DMS with several salts, like NaCl to give methyl chloride, KF to give methyl fluoride, and mercury cyanide to give acetonitrile.

Analysis of the DMS product (from Ullmann):
The reaction with potassium iodide (one of the methyl groups reacts) takes place at room temperature in aqueous emulsion, the MeI is obtained in quantitative yield by distillation. Measuring the amount of MeI can be used to determine the purity of the DMS.

[Edited on 3-1-2008 by garage chemist]

Per - 3-1-2008 at 03:53

Very nice methode,
do you think there are an any chances that it works for ethanole, diethanole sulfate also?

To the CuCl2, I made it by immersing a copper rod into a solution of the right proportions of HCl and H2O2 and let standing this over night. I evaporated the water and then heated the salt to about 300°C and it gaves a dark brown powder which doesn´t seem to contain any water, also I think that the CuCl2 hydrate is blue.
Don´t think that heating under HCl is necessary.

Nicodem - 3-1-2008 at 04:28

You can not obtain CuCl2 by heating its hydrate in normal atmosphere. That has been said before. Try dissolving in water the powder you obtained and you will see what I mean. Also, CuCl2 is not "a dark brown powder".

I wanted to say, that if dimethyl sulfate is needed for methylations of phenols, then one could consider the use of KOSO2OMe which is easily prepared from OTC materials. I will dig out the reference for its use in phenol methylation if anyone is interested at all.

garage chemist - 3-1-2008 at 04:29

No. This does not work for diethyl sulfate.
Diethyl sulfate is best made by Mephistos method from H2SO4, Ethanol and sodium sulfate as seen in the link to the thread.
It can also be made from NaHSO4 and Ethanol as I have shown on VC (the synthesis which is no longer online).

woelen - 3-1-2008 at 23:35

I do not agree on the CuCl2 dehydration. I have tried it as well, and you can get anhydrous CuCl2 by careful and slow heating. If you take finely ground CuCl2.H2O, put that in a beaker and heat over a small flame of an alcohol burner, while constantly moving and swirling the beaker, then you get a nice brown powder, which dissolves in water without any turbidity. High concentration is green, lower concentration is light blue. A basic copper chloride definitely would not produce a clear solution.
But, it is important to do the heating very slowly and carefully. I also started from lab grade CuCl2, not from home-made material, it might be that this makes a difference as well. My CuCl2.2H2O has a beautiful blue/cyan color, not the grass-green stuff which I have seen many times in pictures from people who make it themselves.

Quote:

Originally posted by garage chemist
Ullmann gives an old method of dialkyl sulfate synthesis that consists in heating the mentioned alkyl chlorosulfates with dialkyl sulfites (from SOCl2 and the alcohol).

Do SOCl2 and alcohols really form alkyl sulfites?? I always was thinking that these produce the alkyl chloride, SO2 and HCl.

Per - 4-1-2008 at 07:41

I putted a small amount of my CuCl2 in 50ml water and it forms a clear, turquoise solution.

Quote:

not the grass-green stuff which I have seen many times in pictures from people who make it themselves.

Yes, the dark green stuff contains still much water, the turquoise green solid contains less but still crystal water and the brown powder should be free of water, my experience.

Quote:

It can also be made from NaHSO4 and Ethanol as I have shown on VC (the synthesis which is no longer online).

Great synthesis, if you have a vacuum pumpe.

, yes, I should really buy one.
But what´s the problem with the NaNSO4 synthesis, why isn´t it online anymore?

Klute - 4-1-2008 at 10:01

I'd like to confirm Woelen's observation, the light brown powder I obtained from slowly heating the hydrated CuCl2 on a hot plate on a pyrex dish dissolved quickly to a green/blue solution depending on the concentration. At first I thought the blue color was du to slight acid contamiantion in my cheap hardware distilled water, but it's obviously Cu2+ complexed by H2O ligands.
Also, when left in a humid atmospher (~70%), it reverts to a light green/blue powder over a few minutes, which than turns back to the same light brown powder upon heating. No specially noxious smells were noticed during deshydratation, not done in the hood.
So I'm pretty sure the CuCL2 was indeed that, though if I would ever try this again I would dry technical hydrated CuCl2.

garage chemist - 4-1-2008 at 10:19

Klute, if you find the time to try out the DMS synthesis from H2SO4 and MeOH, can you report on it?
It would be great to see this method confirmed. It would make DMS a simple OTC synthesis.

Or do you want to try the CuCl2 method again?

Klute - 6-1-2008 at 08:50

Are you talking of vacuum distilling a mixture of MeOH with a large excess of conc. H2SO4?
Unfortunaly, I've only heard about failure/terribles yields using this method, and was more interested in trying out what at first seemed to be a pretty easy/accesble synth for DMS. For my methylations needs, I have some MeI at hand, and I guess I will rather try out Halfapint's MeOH/P2O5 pyrolysis for Trimethylphosphate before the "classsical" DMS run. Also, I would prefer working with TMP for safety reasons, I'm still scared shitless when using MeI...
But it's not something I rull out, as conc. H2SO4 is much cheaper than phosphorus pentoxide. With 1mole of P2O5 (140g), Halfapint obtained roughly 45mL IIRC, by following US 2410118 (here)

Here is a link to his report (a the bottom of the page): Trimethyl Phosphate Alkylation of Phenolic Aldehydes

So if I can't obtain better yields than him, or don't consider TMP as efficient as either MeI or DMS (never used TMP for methylations before), I will surely give it a try, and of course report back. I would obviously be very interested if anyone else gave it a try, having a long to-do's list actually

Also, I've been very interested in your SO3 synthesis from peroxodisulfates, even if the yields is low, this could be a better way of producing small amounts of DMS through oleum, for the more delicate methylations. If I understood correctly, using 96% H2SO4 could actually give out oleum simply when the water present reacts with some SO3. Or would it be more adviseable to isolate the SO3, roughly weighing it, and adding it to a specific amout of conc. H2So4 to get a known-amount of oleum?

PS:

Quote:

Do SOCl2 and alcohols really form alkyl sulfites?? I always was thinking that these produce the alkyl chloride, SO2 and HCl.

I remember a procedure in Vogels for this. I was quite surprised myself. But I don't think I would use SOCl2 even if I had some on making DMS

.

[Edited on 6-1-2008 by Klute]

garage chemist - 6-1-2008 at 14:16

The H2SO4/MeOH mix is not distilled in vacuum, but at atmospheric pressure.
This could be the key in reaching the required reaction temperature.
As I have written before in this thread, there is evidence that the decomposition of DMS might be insignificant.
The old chemistry book does not mention any vacuum in the synthesis and even purification of DMS, and gives the boiling point as 188°C, clearly indicating that no vacuum is being used.

[Edited on 6-1-2008 by garage chemist]

franklyn - 6-1-2008 at 23:06

I have no personal basis for contributing to this thread, just this item I noted
related in Fownes Manual of Chemistry Theoretical & Practical
http://books.google.com/books/pdf/Fownes__Manual_of_Chemistr...

which I summarized with comments here , it's from 1878 so you know it's good

--
From page 560 _
Methylsulfuric acid is prepared by 1 part ( weight ) Methyl alcohol slowly mixed
with 2 parts of concentrated Sulfuric acid, and the whole is heated to boiling,
( slowly indeed , one assumes it's the acid that's added to the alcohol ) and
left to cool, after which it is diluted with water, and neutralized with Barium
carbonate. The solution is filtered from the insoluble sulfate, and evaporated,
first in a water bath ( warmed ) and afterwards in a vacuum to the proper
degree of concentration. ( pray tell what ) The salt crystallizes in beatiful,
square, colorless tables, ( flakes ) containing (CH3)2Ba(SO4)2.2H2O, which
effloresce ( crystallize ) in dry air, and are very soluble in water.
By exactly precipitating the metal sulfate from this substance with dilute sulfuric
acid, and leaving the filtered liquid to evaporate in air, Methylsulfuric acid may be
procured in the form of a sour, ( DO NOT TASTE ) syrupy liquid, or in minute
acicular ( needle shaped ) crystals, very soluble in water and alcohol. It is very
instable, being easily decomposed by heat.
Potassium Methysulfate, CH3KSO4 crystallizes in small, nacreous ( iridescent )
deliquescent rhombic tables. The lead salt is very soluble.

DiMethylsulfate (CH3)2SO4 is prepared by distilling 1 part of Methyl alcohol with
8 or 10 parts of strong Sulfuric acid: ( i.e. able to burn holes in shoe leather

)
the distillation may be carried nearly to dryness. The oleaginous ( oily ) liquid
found in the receiver is agitated with water, and purified by rectification ( distilled )
from powdered anhydrous baryta ( BaO ). The product is a colorless, oily liquid,
of alliaceous odor ( garlic pew ), having a density of 1.324, and boiling at 188 ºC
( 370 ºF ). It has neutral pH, and is insoluble in water, but decomposed by it,
slowly in the cold, rapidly and with violence at boiling, into Methylsulfuric acid
and Methyl alcohol. Anhydrous lime and baryta ( CaO , BaO ) have no action on
it, their hydrates however, and those of Potassium and Sodium, decompose it
instantly, with production of the Methylsulfate salt and Methyl alcohol. When
this Methylsulfate salt is heated with common salt, it yields Sodium sulfate and
Methyl chloride. With Mercuric cyanide, or Potassium cyanide, it forms the
Sulfate salt and Methyl cyanide. With dry formate, it yields Sodium sulfate and
Methyl formate.

Methyl Sulfite (CH3)2SO3 formed by the action of Sulfur dichloride S2Cl2, on
Methyl alcohol, as a fragrant liquid ( that can't be good for you ) having a
density of 1.045 and boiling at 121 ºC ( 250 ºF )

Methyl Sulfonic acid etc.

_____________________

Ethylsulfuric acid is prepared as is Methylsulfuric acid above , it can instead be -

From page 573 _
- " neutralized with chalk, whereby much calcium sulfate is produced. The mass
is placed upon a cloth filter, drained , and pressed; and the clear solution is
evaporated to a small bulk by the heat of a water bath, filtered from a little
sulfate, and left to crystalize.
The product is Calcium Ethylsulfate in beautiful colorless, transparent crystals,
containing Ca(C2H5)2(SO4)2.2H2O. These dissolve in an equal weight of water,
and effloresce in a dry -?- bulk.
It is annoying that Google has gone to the trouble of copying books as this
and it is done so sloppily that much of it is out of focus and illegible.
Given the nature of these reagents it's possible it was deliberate though it
doesn't really matter as it's available in many libraries search worldcat
http://www.worldcatlibraries.org
Barium Ethylsulfate, equally soluble and still more beautiful may be produced by
substituting Barium Carbonate for the chalk. From this salt the acid may be
procured by exactly precipitating the metal sulfate with dilute sulfuric acid, and
evaporating the filtered solution in a vacuum at the temperature of air. It forms
a sour ( REALLY DO NOT TASTE ) syrupy liquid, in which sulfuric acid cannot be
recognized by the ordinary reagents, and is very easily decomposed by boiling.
The Lead salt resembles the Barium compound. The Potassium salt C2H5KSO4
easily made by decomposing Calcium Ethylsulfate with Potassium Carbonate,
is anhydrous, deliquescent(?) in the air, very soluble, and crystallizes well.

Neutral Ethyl sulfate ( diethylsulfate )
is formed by passing the vapor of sulfuric oxide ( sulfer trioxide ?) into perfectly
anhydrous ( diethyl ) ether. A syrupy liquid is produced, which, when shaken
with 4 volumes of water and 1 volume of ether, separates into two layers, the
lower containing ethylsulfate acid and various other compounds, while the upper
layer of an ethereal solution of neutral ethyl sulfate. At a gentle heat the ether
is volatilized, and the ethyl sulfate remains as a odorless liquid. It cannot be
distilled without decomposition
( I cringe at the comment that it is " odorless " and wonder what became of
the author.)

Ethyl sulfites

Ethylsulfinic acid etc.

.

Klute - 7-1-2008 at 09:15

I'm surprised this atmospheric "pyrolysis" hasn't be mentionned before IIRC. I don't remember anything on the subject at the Hive or elsewhere.
The flask might get damadged after a few runs though, if they advise to distill to near dryness, your distilling excess sulfuric acid!
I'd be pleased if this worked.

not_important - 7-1-2008 at 21:55

Regarding Fownes Manual of Chemistry Theoretical & Practical in particular, and and problems with Google books, sometimes there are solutions. In this case http://www.archive.org/details/manualofelementa00fownrich has a scanned version where on page 527 it end with "effloresce in a dry atmosphere."

Always check both Google and the Internet Archives, either or both may have a less than perfect scan of a book, but often using both version you can patch together a decent copy of the book.

franklyn - 8-1-2008 at 16:00

Quote:

Originally posted by not_important
Always check both Google and the Internet Archives,

@ not_important,
that's an outstanding idea, I had not imagined to search archive.net
It also gives one a choice of format in which to view. As a text file
it leaves no doubt as to content, and less than 3MB compared to the
clunky 106 MB pdf. Curiously viewing this text as html in the browser
disallows highlighting to copy, one must first save that viewed file as
a text file, open that and then copy from that.
The edition of Fownes provided ( 1869 ) pre-dates the 1978 edition
I reference , so it is not word for word correspondent.

I have just discovered that Google maintains 4 separately prepared editions
of this book I have reviewed them and some leave very much to be desired.

http://books.google.com/books?id=kPc4AAAAMAAJ&printsec=f...
Download link _
http://books.google.com/books/pdf/Fownes__Manual_of_Chemistr...
64.6 MB - RECOMMENDED - OCR Text very good

http://books.google.com/books?id=ceMHAAAAIAAJ&printsec=f...
http://books.google.com/books/pdf/Fownes__Manual_of_Chemistr... XmLWI
60.8 MB - DO NOT BOTHER WITH THIS - page photographs very badly done

http://books.google.com/books?id=7awEAAAAYAAJ&printsec=f...
http://books.google.com/books/pdf/Fownes__Manual_of_Chemistr...
67.7 MB - DO NOT BOTHER WITH THIS - A mixed format of OCR with some bad photographs

This is the 1885 edition ( also embodies 'Watts' Physical and Inorganic Chemistry )
http://books.google.com/books?id=Xfg4AAAAMAAJ&printsec=f...
Download link _
http://books.google.com/books/pdf/Fownes_Manual_of_Chemistry...
45.4 MB - NOTE - Methyl on page 588 is essentially the same, Ethyl on page 603 is heavily abridged
. . . . . . . . that's not to say it is not a worthwhile reference.

For convenience one can do the following :
open notepad and copy and paste the following at the very top _

[InternetShortcut]
url=http://books.google.com/books/pdf/Fownes_Manual_of_Chemistry__Theoretical_.pdf?id=Xfg4AAAAMAAJ&output=pdf&sig=QvqKzrx9g7EJVR5FT8cakMzdkEA

name this file and save it with the extension .url
you have now made a shortcut to quickly download the 1885 pdf anytime,
that you can save instead of the whole reference.
- The url after ? above MUST continue as an unbroken line in notepad or this
will not work ! -

Note that whatever you enter after url= ? can be any internet address.

.

not_important - 12-1-2008 at 08:36

Soda siphons are filled to 100 to 130 psi, but are not meant to be heated. You'd do better making a metal shell to take the pressure, and placing a glass container inside it to hold the reaction mix.

However note that this reaction can be done in a flow reactor using K2CO3 coated with PEG as the catalyst and reaction phase, and the reactants as gas. The process can be run at atmospheric pressure. I suspect a mixed phase batch reaction, with DMC in the gas phase and the other reactant as liquid, could also be done.

garage chemist - 12-1-2008 at 10:57

It would be nice if this thread can be reserved for DMS synthesis discussion.
Maybe you could open a "Green methylating agents" thread to discuss things like dimethyl carbonate, methyl tosylate, methyl mesylate etc...
Just a suggestion.

Some old bones

franklyn - 21-1-2008 at 11:12

http://www.sciencemadness.org/talk/viewthread.php?tid=1113&a...

Attachment: pat3024263.pdf (256kB)
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Fleaker - 21-1-2008 at 11:48

I did a quick and dirty attempt in the lab the other day using 20% oleum and methanol. I also tried the bisulfate+MeOH to get the sodium methyl salt.

I just used a small amount, excess oleum, and about 2 mL of MeOH. I did it in a long Pyrex test tube immersed in cool water. Upon adding the first half of the MeOH, there was a strong exotherm when the tube was put against my gloved hand. Measuring the temperature with a probe said it was 64*C, and this was after I had removed it from the cool water! This is highly exothermic and produces a lot of vapor, most likely H2SO4 droplets. I let the rest of the reaction mix cool down to 18C and added more MeOH and it went up to 59*C, but there was much less vapor produced. I added the alcohol using a pasteur pipet.

There were two layers formed, the bottom sulfuric acid layer, and the oily top layer which did not wet glass. I have not had a chance to do any formal qualitative testing on it, but I do not think it is methanol...
My reasoning is that I never saw any methanol condensing on the walls of the reaction tube. I used a tall test tube because I was expecting to see MeOH condense on the walls but there was none, so I can only assume that it reacted to something much less volatile that MeOH. You would think that a MeOH mixture would have some vapor pressure at 64*C, but there were no droplets on the walls of the test tube. I will probably run a larger batch and see what I can get by rotovaping.

With the bisulfate method, I did not notice much indication of a reaction but rather saw methanol condensing. I heated it to about 50*C as the reaction did not seem terribly exothermic by itself (very little measured temp difference). I do not think it works as well as the oleum method, where there was definitely something happening.

Klute - 21-1-2008 at 13:46

Thanks for sharing!!

It definatively looks like the oleum produced some DMS, so using Garagechemist's peroxodisulfate mehtod of making some oleum and reacting that with MeOH could be the easiest method for producing small amounts of DMS. Plus oleum can be usefull for tons of other things.
I guess adding the oleum to excess MeOH would give less of an exothermic reaction, although adding oleum might be a bit less practical than the other way around.

My To-do's list has just got bigger

garage chemist - 21-1-2008 at 17:05

You should try pouring the oleum/MeOH reaction mix into ice water and see if any oil settles to the bottom. Any oil would likely be crude DMS.

Again, I would be most interested in hearing about experiments with the H2SO4 method.
One could easily do that on a small scale: put 10g MeOH in a flask, carefully add 50ml conc H2SO4 with swirling and rapidly distill the mix over a strong flame.
The distillate should then consist of two phases.

Fleaker - 21-1-2008 at 20:37

I think I already have DMS available, so I'm mainly just doing this to do it. I have no use for the methylating agent. I used oleum instead of trying with sulfuric because I have plenty of oleum, and I was fairly certain I'd see something. I'll see what refluxing with conc. sulfuric does.

The ice water idea is a good one. I will try that.

As far as distilling the mix, I'm really prone to thinking that there will be no methanol present. It appeared to be a fairly complete reaction, but it's all speculation...

Klute - 22-1-2008 at 12:36

I'd equally be very interested to see how the atm distn goes!

At the risk of sounding a little selfish, I'm on dozens of other projects at the moment, chemical or not, and as I'm not in desperate need for methylating agents, this reaction isn't a priority. Even if it can't be too long at the scale you suggest GC, doing long days of work, having a minimum social life (my girlfriend isn't a chemistry fan, she's even jealous sometimes

) leaves very little time in the lab.

Again, thanks for sharing Fleaker, especially if you're not in the need of DMS. Considering the price of oleurm, i can only say this a very generous experiment

Thumbs up!

[Edited on 22-1-2008 by Klute]

Nicodem - 22-1-2008 at 13:07

Quote:

Originally posted by garage chemist
You should try pouring the oleum/MeOH reaction mix into ice water and see if any oil settles to the bottom. Any oil would likely be crude DMS.

I got the impression that dimethyl sulfate forms only during the distillation while at room temperature one would mostly have a mixture of methyl sulfuric acid and H2SO4? Kind of like the thermolysis of sodium methylsulfate which gives dimethyl sulfate and Na2SO4. If dimethyl sulfate is indeed the main ester in the mix already at room temperature, then it could potentially be extracted with CH2Cl2 with which H2SO4 is immiscible. Washing the combined extracts once with water should remove most of the methyl sulfuric acid and methanol traces, leaving only dimethyl sulfate.

Fleaker - 22-1-2008 at 14:42

Excellent idea as usual Nicodem. A rotary evaporator would make short work of the DCM solvent.

How hot must it be for the methyl sulfuric acid to convert completely to the ester?

PainKilla - 22-1-2008 at 21:33

I've had mediocre results at synthesizing DMS. There has been extensive discussion of this topic at other forums and so far what I've tried has only lead me to one conclusion: vacuum distillation is the only really viable method.

- I've tried distilling from various methyl sulfuric acid salts, which according to an old Chem. Ber. reference (IIRC) are supposed to liberate DMS upon warming, and especially so under vacuum. Nothing happened.

- I've tried distilling straight methanol/sulfuric under a few different conditions, using different proportions, and drying agents. This also has led to mixed success (ranging from 10-20% or so, from what I remember)... under vacuum at least.

The method I have found works the best is attached.

A couple of notes:

I am sure that the methyl sulfuric acid formed has traces of sulfuric acid in it... I can't decant well enough to eliminate all traces of sodium sulfate, so as a result water is inherently present during distillation. Additionally, I am not sure how well the patent actually does work to begin with... but it does seems to work better than straight distillation, and given the cost of the materials, the relatively low yield of obtained dimethyl sulfate is decent in that regard.

The dimethyl sulfate obtained works well for further reactions, though care should be taken to fully neutralize it as some decomposition was noted after a few months of storage at room temperature in a sealed vial (stored over anhydrous magnesium sulfate). I don't know what happened here... in any case, just a forewarning if you plan on storing this for an extended period of time.

Attachment: Dimethyl Sulfate Synthesis.rtf (4kB)
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Klute - 23-1-2008 at 22:10

Thanks for that very well written document Painkilla! I take it the yield claimed is what you usualy obtained, and I think it is very reasonable, owing, as you say, to the cheapness of the reactants and the usefullness of the product.

Concerning the distn, a few questions: did you use a short coluum? How marked is the transition from DMS and excess H2SO4? Was there any damadge to the still flask?

PainKilla - 24-1-2008 at 06:07

I used regular 19/22 Pyrex glassware with a 200mm condensor; only one was needed. As far as the rest of the setup, I typically tried to shorten the distance that the vapor must travel, as using larger flasks seemed to give a lesser yield. I don't know if this is a fault of my vacuum or a tendency of the reaction.

I did do this specific reaction a few times (overall I've tried making DMS using various methods probably 20-25 times), as dimethyl sulfate was at one point a reagent that was rather difficult for me to buy. The reaction is rather quick - it's possible to do within two hours, plus another hour for setup/cleanup. The glassware used still survives to this day, no problems of any sort. In fact, the only time something did break, was indirectly caused by the reaction: I stepped on a condensor. :*(

The transition is rather strange. When you initally mix the Na2SO4 and CH3HSO4, the Na2SO4 forms a hard solid in the bottom of the flask, which is why it's important to added Na2SO4 to an already stirred mixture of CH3HSO4. After this, heat was applied (via an oil bath) and nothing seems to happen until the temperature is about 70-80*C. At this point you can see an oily substance condensing, and it seems to co-condense with methanol or water - makes no sense to me though, since they should both be quite gaseous at those pressures.

The distillation continues and the temperature raised until no more oily DMS condenses. A vapor begins to come off the reaction flask, I suppose it is sulfuric acid vapors. In any case, it does not affect the distillation in any way. The Na2SO4 then seems to dissolve in the mixture - also an interesting curiosity.

The liquid in the reciever is most interesting: it's oily in the condensor but upon looking at it during the first few minutes it's in the reciever, it looks like regular water. Eventually, DMS oils out of it. The most interesting thing is, this liquid, does not get neutralized by water, and is soluble in it. So it's not any sort of acid... and probably not methanol. But how can it be water?

In any case, the reaction is indeed cheap and great for making small, usable amounts of DMS.

Methyl iodide is easier to make and yields higher, but since you almost always need to use an excess of CH3I, the increased yield and volatility really isn't worth it.

I should note I did this distillation outside with gloves and the like. (Although the reaction can be left unattended.)

sonogashira - 25-1-2008 at 08:48

PainKilla, did you not purify the methyl sulfuric acid other than by decanting it from the sodium sulfate? Could it be vacuum distilled perhaps, or is it not necessary to purify it for use in the next step?

Quote:

Originally posted by Nicodem
I wanted to say, that if dimethyl sulfate is needed for methylations of phenols, then one could consider the use of KOSO2OMe which is easily prepared from OTC materials. I will dig out the reference for its use in phenol methylation if anyone is interested at all.

Any chance of getting this reference Nicodem? I would be most interested

PainKilla - 25-1-2008 at 09:48

There are actually quite a few ways to make methyl sulfuric acid and/or it's salts, however, I only briefly explored the other routes, mainly it's salts, simply because they involve much more work so as a result I took a liking to this simple and quick method that seems to largely work as advertised.

The procedure I posted isn't meant to be high-yielding, it's rather just a simple and efficient way to get DMS. I've never needed large amounts of it so there was never a reason to optimize the synthesis to yield especially well given the ease and price of the reagents.

I did try filtering the methyl sulfuric acid, it however tore my cellulose filter apart during filtration, so something more resistant would have to be used.

If you distill methyl sulfuric you are just going to get dimethyl sulfate (albeit in lower yields).

There may be other ways to dry it, but this is the most commonly accessible. I would recommend a patent and literature search... There is a huge amount of patent "literature" on the topic (though how much of that actually works in practice is unknown to me - it seems very few actually work), and older books on chemistry almost always have some "unique" method of synthesized DMS.

Edit: I should not however, discourage anyone from retrying the experiments with methyl sulfuric acid salts, I suspect as there was some water (and thus sulfuric acid) present in my CH3HSO4, the reaction would not have gone quite as advertised... that is:

2CH3HSO4 + Ba(OH)2 ---> (CH3)2Ba(SO4)2 + 2H2O... but:
H2SO4 + Ba(OH)2 --> BaSO4 + 2H2O

The second reaction would likely occur faster, and the water formed would serve to hydrolyze more acid... so I think you really need to have very, very, dry methyl sulfuric for this to work to any appreciable degree.

The sodium salt however was made via different method:

CH3HSO4 + NaCl --> CH3NaSO4 + HCl

However this salt did not work as a DMS source, as advertised, under vacuum.

[Edited on 25-1-2008 by PainKilla]

Klute - 25-1-2008 at 10:12

So you distiled the MeHSO4 directly from the Na2SO4? That's surely were your water came from,a s it woill loose the water of hydration upon heating. But I recall this method worked anyway, so maybe the Na2SO4 hold juste enough to the water for it not to hydrolyse the methyl sulfuric acid or the DMS, especiially at those temperatures. Have you tried deacnting the MeHSO4 from the sulfate before distilling? This was brought up a few times too.

It's a pleasure to read such detailled descriptions. Makes it seem so easy

PainKilla - 25-1-2008 at 10:25

No, first you make the methyl sulfuric acid (at least, in a much more appreciable concentration), then decant from the (hydrated) sodium sulfate as per the patent, then add Na2SO4 (just because it seems to help) and distill.

I have tried doing this with MgSO4 and the results were a bit worse.

[Edited on 25-1-2008 by PainKilla]

Klute - 25-1-2008 at 11:32

Sheesh, I've just realized you explained it explicitly in the document... Read it quickly last night, half didn't stay in

S.C. Wack - 26-1-2008 at 12:35

Old news- the formation of DMS from H2SO4 and methanol was confirmed a long time ago in Compt. rend., and the same dudes also wrote of an interesting lithium methyl sulfate pyrolysis. Both articles are on Gallica. Haven't tried the pyrolysis method yet even though I've got nothing better to do with my LiCO3.

But I have tried the straight up 8-10 parts thing a few times (24/40, 19/22). I'd recommend a VERY short path distillation apparatus and the flask top and head wrapped up good in something, because although it CAN be distilled at 760 mm., the problem is pot temperature (if like me you want to get your product anytime soon). DMS high pot temp NO like. My yields sucked regardless.

Go vacuum.

Klute - 27-1-2008 at 05:44

Could you give the ref for these articles, if you've got them at hand obviously? I can translate them, they could be helpfull.

Nicodem - 27-1-2008 at 08:11

I think S. C. Wack is talking about the attached paper.

Quote:

Originally posted by sonogashira
Any chance of getting this reference Nicodem? I would be most interested

Well, one reference describing the methylation of plain phenol with MeOSO3H is the example 2 of patent US2490842. Also note the other examples where there is also described a methylation method for phenols that uses 1/2 equivalents of dimethyl sulfate (both Me groups get transferred). The reaction temperature used is considerably higher than those in classical methylations with dimethyl sulfate (that is: unlike those described in Vogel's or Organikum which work already at room temperature, but transfer only one Me group). This method using MeOSO3H or 1/2 eq. (MeO)2SO2 is thus not suitable for those phenols which are unstable in basic media at 110-120°C.
This patented method was checked by Antoncho on hydroquinone and posted as post No 256342 on the late Hive forum. About 56% yield of p-dimethoxybenzene was obtained, which is not bad considering the substrate and conditions used.
There is also at least one (but maybe more) reference(s) from literature using MeOSO3Na or MeOSO3K, but I can't find them at the moment (I will post them as soon as I find them).

[Edited on 27/1/2008 by Nicodem]

Attachment: metal methylsulfates.pdf (153kB)
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S.C. Wack - 27-1-2008 at 08:37

I uploaded the salt one before in another post here long ago IIRC. It's barium and strontium as well. The one using acid mentions getting a 44.4% yield from 3 moles H2SO4 to 1 of methanol, which I was unable to obtain. I have a retort - which would have been the better apparatus, but I don't trust the glassblower that much. Both are from 1919, in different volumes.

http://gallica.bnf.fr/ark:/12148/bpt6k3121p
page 1204.
http://gallica.bnf.fr/ark:/12148/bpt6k31220
page 655.

Fleaker - 11-3-2008 at 04:41

I tried this reaction again last night on a larger scale using 50 mL of MeOH and excess oleum.
The reaction is very exothermic, even with dripping in using an addition funnel. It is sputters and sprays when the funnel is open full tilt, I believe this is because the MeOH penetrates the interface between the formed DMS and the heavier oleum.

Upon distillation with aspirator vacuum, the bulk of the product came over from 115-124*C, then a second portion at 180-188, the temperature just shot up afterwards and I stopped the distillation. I wonder if I was distilling SO3 into my product!? In the reaction vessel, there was a good bit of carbonization by the hot sulfuric acid.

About 75 mL of product collected in the receiving flask, but it was a bilayer: a clear, slight yellow bottom layer with a top layer floating above. Both phases were quite mobile. I think there is a significant portion hydrogen methyl sulfate. I will try and take a GC/MS tomorrow evening.

I'm going to redistill the DMS over anhy. Na2SO4 with a little bit of oleum left.

I plan on trying this for diethyl sulfate by using chlorosulfonic acid sometime soon.

sonogashira - 3-4-2008 at 11:23

Quote:

Nicodem
This patented method was checked by Antoncho on hydroquinone and posted as post No 256342 on the late Hive forum. About 56% yield of p-dimethoxybenzene was obtained, which is not bad considering the substrate and conditions used.

I'm sorry to always be asking for assistance but i really have no idea where where i could find this "reference." I would be most interested if someone could be of help

Klute - 4-4-2008 at 09:47

There you go:

Attachment: 000256342.html (55kB)
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DJF90 - 19-10-2012 at 13:58

Sorry to dig up an old post but I figured I'd share some recent (partial) success and also ask for advice.

So I've been trying to make dimethyl sulfate. I figured I'd start with Painkilla's excellent text document. I tried the reaction on half the scale. The sodium sulfate was well dried and finely powdered, and added to the mixture of methanol and sulfuric acid. The appearance of the mixture changed gradually over time, becoming more opaque (white?) and thick, almost with the appearance of how I'd imagine a fine chalk suspension to look. Leaving this for a couple of hours (2 hours), it was allowed to settle for a little time (I had little patience though it helped very little). A glass sintered funnel (por 3) and buchner flask were dried with a heat gun under evacuation, left to cool, and the mixture filtered (under full vacuum, about 20mbar). This took OVER AN HOUR due to the fine nature of the sodium sulfate.

100g of the methylsulfuric acid* were transferred to a dried 100ml RBF and the appropriate amount (22.5g according to Painkilla's document) of finely powdered anhydrous sodium sulfate was added with vigorous magnetic stirring. The thick suspension was placed under high vacuum (0.5mbar) for a few minutes, and then a bleed was introduced via another tap on the schlenk line, bringing the system pressure to 10mbar. Heating was commenced, and strong bumping began at a bath temperature of 95*C. At a bath temperature of 120*C, distillate began to pass over, and the evolution of distillate increased at a bath temperature of 145*C. Heating was maintained and increased as the amount of distillate dwindled. A final temperature of 200*C (vapour temperature of 120*C). A biphasic mixture was obtained, which was poured into a measuring cylinder, the upper layer removed, a volume of water added, agitated, and removed, leaving 11ml of slightly milky dimethyl sulfate (a 25% yield) which was dried over anhydrous magnesium sulfate.

A subsequent run was made using 20g anhydrous magnesium sulfate to dry the mixture of methanol and sulfuric acid, and no additive during the distillation. 83g of methylsulfuric acid was obtained on filtration (less than before) and on distillation at a pressure of ~25mbar gave a biphasic mixture yielding 7ml dimethyl sulfate after washing. The bath temperatures more closely match what Painkilla describe though, and the distillation in practice was "easier" (for want of a better description). In total I now have 18ml of crude dimethyl sulfate.

However, I wish to still accumulate and find a better method of preparing it. The distillations take an hour or two each, but the clean up after takes longer, carefully dismantling the appratus and placing each piece in a base bath and leaving it there for a few hours (overnight) to hydrolyse any dimethyl sulfate. I've thought about possible problems and my main concern is with the preliminary drying step. When adding the magnesium sulfate (a much more rapid drying agent), it initially clumped together, but it appears to have been ground to a fine powder by the mag stirrer bar. I think if I were to try this again, the mag stirring would be cut off before addition of solid.

I'm also interested in trying the CuCl2/SO2 method, as I feel this would be fruitful, especially if it can be easily adapted to produce diethyl sulfate. I feel the troubles lie in the workup of the reaction mixture, so any advice here is welcomed.