mario840
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Dimethyl Sulfate, cheapest preparation, OTC !
Hi guys !
Chemistry is the science when always is a way to synthesis wanted chemical in cheaper way , but there is a point (line) when you can't cross them, so
i think for this chemical is this preparation, i looking in old patents and found magic "Patent 1,641,005 United States Patent Office".
Process decribed below is exceedingly simple and by far the cheapest method of making dimethyl sulfate and other alkyl sulfates.
When we add alcohol (methanol) to sulphuryl chloride then methyl hydrogen sulphate is formed instead dimethyl sulphate but when we ADD ONE
MOLE OF SULPHURYL CHLORIDE TO TWO MOLE OF METHANOL THE REACTION TAKES PLACE :
SO2Cl2 + 2 CH3OH --------> (CH3)2SO4 + 2HCl
A very good yield is obtainted of desired sulphate and recovery of sulphate and excess sulphuryl is really easy (just look at the patent).
What do you think about this ??
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bbartlog
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OK, but since when is SO2Cl2 OTC? Did I miss something there? 'OTC' is not the same as 'commercially available'.
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Jor
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Well, SO2Cl2 can quite easily be prepared by passing a mixture of SO2 and Cl2 over camphor (catalyst). I think garage chemist has done this.
Anyway, even if this would be a good method for dimethyl sulfate, you really want to use it? It is a very dangerous material, wich is more dangerous
than conc. HF in terms of skin contact AFAIK, and you can hardly smell it's vapours, wich easily cause pulmonary oedema. Added is it's very strong
carcinogenic property.
I would rather use something like MeBr (easy to make), or MeI (also quite easy, look in the republication subforum).
I known all MSDS's are exaggerated, but this is the most serious JTBaker MSDS I have seen around:
http://www.jtbaker.com/msds/englishhtml/d6960.htm
I'd rather pay a little extra for MeI than using something such as dimethyl sulfate (although MeI is still very dangerous and much more volatile).
[Edited on 20-12-2010 by Jor]
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mario840
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Sure SO2Cl2 like Jor said is simple to prepare , even buy, for example : thionyl chloride is the shit so for my opinion is really simple, yes this
chemical is very toxic but here is just pure science like in name of these forum, when you have experience with toxic chemical and voilate liquids and
proper equipment then this chemical won't hurt you , so everyone must to asnswer to themselfs: "can i handle this ? "
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bbartlog
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| Quote: | | SO2Cl2 can quite easily be prepared by passing a mixture of SO2 and Cl2 over camphor (catalyst) |
Fair enough! SO2, Cl2, camphor (in pure form) are *also* not really OTC in the usual sense of the word. I'm not disputing the possible value of the
proposed synthesis, nor saying it's infeasible. I'm saying that 'OTC' has a particular meaning (which I grant does vary from place to place): you can
buy the reagents needed at a bricks-and-mortar storefront, without needing special qualifications or credentials. It also doesn't mean that the
proposed reagents can be prepared via some other process from OTC inputs, else everything would be technically 'OTC' given enough expertise.
Certainly it would be nice if there were some shorthand to describe processes where the inputs were, while not strictly 'OTC', at least readily
available from internet suppliers and/or specialty shops. Welding gases, soapmaking supplies, and the various salts you can get from ceramic shops are
examples of things that are not normally considered OTC, but can still be sourced easily and cheaply enough without needing a corporate account with
Sigma-Aldrich.
Maybe that kind of stuff could be labeled 'OTI' (over-the-internet). Though I believe sulfuryl chloride still would not qualify.
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vulture
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I wouldn't call reacting SO2 and Cl2 easy. The preparation and handling of both is already quite hazardous per se, certainly if you're looking for a
decent yield.
One shouldn't accept or resort to the mutilation of science to appease the mentally impaired.
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madscientist
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I hope that by "proper equipment" you mean a glovebox. Alkylating agents like dimethyl sulfate and methyl iodide are nasty.
I weep at the sight of flaming acetic anhydride.
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vulture
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I usually handle MeI in an efficient fumehood, wearing two pairs of nitrile gloves with a NaHCO3 solution standby. I've never handled dimethyl
sulfate, but from what I've read it's quite more hazardous.
One shouldn't accept or resort to the mutilation of science to appease the mentally impaired.
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ScienceSquirrel
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Dimethyl sulphate is OK in a good hood with a high chimney at a research lab and you want good gloves as well.
I did a lot of alkylations with it at one time.
Making and handling it in a home lab in a residential area.
I will give it a miss.
Home science is OK within reason, but dimethyl sulphate, phosgene, etc are very hard to handle.
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Sandmeyer
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I also put my vote on "Should be avoided at home"... But I don't agree with madscientist about the glove-box. I mean it is not Ebola... If I was to
choose between DMS and MeI on lab scale I'd go for MeI any day...
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Sydenhams chorea
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Ive handled DMS at home and although its dangers are not to be underestimated, it can be done using adequate protection.
Not that it doesn't make me nervous, and I'd rather not attempt distilling it either.
To the original poster: this thread is a good read:
http://www.sciencemadness.org/talk/viewthread.php?tid=9570
| Quote: |
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. |
Il n'y a point de sots si incommodes que ceux qui ont de l'esprit.
François de La Rochefoucauld.
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Natures Natrium
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Here is some info. Nothing conclusive, as I decided on careful consideration that I did not have the proper facilities to handle the DMS. It serves,
I think, as more of an example of what does not work.
http://www.sciencemadness.org/talk/viewthread.php?tid=4139#p...
The idea behind the synthesis is solid, IMO, but a safe and working procedure still needs to be devised.
EDIT: Also, just wanted to mention that having synthesized both S2Cl2 and SO2Cl2 in the home lab, I would definitely prefer the later over the
former.
S2Cl2 required passing chlorine gas through molten sulfur, and to this day the thermometer adapter is stuck to the distillation adapter, and no amount
of ultrasonic cleaning, heating of joints, or chemical cleansers could get them to seperate.
SO2Cl2, on the other hand, was a mild, room temperature reaction. I simply combined the gases in as close to 1:1 molar as possible, and passed them
through a condenser filled with activated carbon. Details here: Hmm, I actually cant find the thread detailing it. Odd.
Ah, I apparently only detailed it in a U2U. If there is some interest I can post it.
[Edited on 21-12-2010 by Natures Natrium]
\"The man who does not read good books has no advantage over the man who cannot read them.\" - Mark Twain (1835-1910)
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sonogashira
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"If there is some interest I can post it."
Yes, please do...
Thanks!
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Natures Natrium
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"First, the reagents/reactants/catalysts. As you know, my homemade "lecture" bottle of sulfur dioxide was my source of same, and proved to be quite
convenient. My source of dry chlorine was obtained by simply dropping Muriatic acid on pool bleach, and running the gas through a large test tube
filled with calcium chloride pellets. The catalytic carbon used was the same I mentioned in the SCl2->SOCl2 thread. That is to say, it is ultra-dry
aquarium activated carbon, boiled with calcium chloride solution for an hour, decanted, dried, and heated under vacuum for several hours. I can
honestly say that I think all this treatment was a grand waste of time, and that trying to get this stuff anhydrous afterward is nearly impossible.
However, the next thing I did is what I believe actually made the stuff work well. I simply poured about 10-20mL of my SCl2 directly onto the carbon
in its glass container (in a good draft of air, lots of SO2 and HCl given off). This thoroughly eliminated any trace of water in the carbon. This was
used as is for the reaction.
Second, apparatus. I used rubber and copper tubing to funnel the gases. In particular, I used a bronze 3-way nipple to mix the gases, just prior to
entering the condenser (reaction tube). The three way junction was inserted into a rubber stopper which was plugged into the head of the condenser.
This proved to be a mild error, as it seemed to me that the SO2 and Cl2 were capable of combining on the surface of the rubber, which the sulfuryl
chloride then reacted with producing a white mist that inhibited further catalysis. I eliminated this problem by inserting a glass tube into the
rubber stopper until it seated against the end of the nipple of the 3-way bronze junction. I think that as an improvement it would be best to
thoroughly mix the gasses before they enter the catalysis tube, however something other than rubber hose would probably have to be used. Anyways, the
catalysis tube was in fact my 300mm lieberg condenser, and I used teflon tape across the glass joint at the end to hold the carbon in. This too was a
mistake that had to be rectified, for two reasons. One, The sulfuryl chloride was capable of leeching through the tape, thereby escaping in small
quantities. Secondly, since the activated carbon was outside the cooling zone of the condenser, it easily became too hot and hindered the yield
severely. Both problems were eliminated by using a glass rod to hold the activated carbon back up in the cooling zone of the condenser. Glass wool
would work much better, I believe. The end of the condenser was attached to a vacuum adapter, and to this was attached a 500mL RBF in ice water. A
hose was attached from the vacuum adapter that led to first an empty "suckback" bottle, and from there to a tube leading to the bottom of a 500mL
flask filled with saturated NaHCO3. I suppose that NaOH would be a better wash, but hydroxides are becoming scarce around these parts.
So, the chlorine was started until the receiving flask has hints of yellow-green, and then the sulfur dioxide was started at such a rate as to match
the chlorine by volume. It is very important to have cold water running through the condenser, and that all the activated carbon is inside the cooling
zone. Otherwise, the initial good yields will quickly drop as the carbon appears to get hot enough to reverse the catalysis. (And the carbon does get
quite hot without cooling. As an aside, its weird to have a vapor phase catalysis that needs efficient cooling instead of efficient heating. ^_^)
I wish I could comment on the camphor route, but I have never seen pure chemical camphor for sale (although I am sure I could find or extract it if I
was so inclined). I suspect that the camphor method has probably a better yield from starting reagents, as both gases get quickly absorbed in the
SO2Cl2 formed. I should note that I attempted to form extra SO2Cl2 by adding some of my activated carbon to the receiving flask, adding a sitrbar, and
bubbling in first chlorine, then SO2. I think I had limited success, but ultimately the reaction was very slow.
I eliminated the excess chlorine from the solution (there was quite a bit) by storing it over a small pool of mercury. After a month or so in the
freezer the mercury had turned to a light grey powder (HgCl2, another nice product), and the SO2Cl2 had become water clear. I distilled the SO2Cl2
before use in a chlorination process, and found that even on gentle distillation the stuff disassociates in the vapor phase to a small degree,
resulting in the distillate being a light yellow-green color."
\"The man who does not read good books has no advantage over the man who cannot read them.\" - Mark Twain (1835-1910)
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aliced25
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According to recent research, methyl hydrogen sulfate is formed within an hour when sulfuric acid and methanol are mixed, pretty much quantitively at
that. So with methyl hydrogen sulfate has anyone actually tried Sartori's much talked about thermal decomposition of at 130-140C to dimethyl sulfate?
It would be of significant interest, methyl sulfuric acid would be the product of acid hydrolysis of alkali (or ammonia) methyl sulfate salts made by
reacting methanol with alkali bisulfate (or sulfamic acid) and would put a useful (albeit quite nasty - but aren't they all) methylating agent within
reach of all but the most dimwitted.
Attachment: Guzowski.etal.Understanding.and.Control.of.Dimethyl.Sulfate.in.a.Manufacturing.Process.Kinetic.Modeling.of.a.Fischer.Est (1.5MB)
This file has been downloaded 848 times
Attachment: Foster.Sulfonation.and.Sulfation.Processes.pdf (315kB)
This file has been downloaded 16906 times
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DJF90
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I've tried the thermal decomposition of methyl hydrogen sulfate, as per sartori. Painkilla (a member here) originally did the work and I tried to
replicate it, with limited success. I got about half the yield he did, and I amounted approx 20 mL over three runs, which given the hassle of working
with such a nasty substance (decontamination of apparatus, phase split of the biphasic distillate etc.) was disappointing. I tried in the region of
about 10-40mbar (different pressure each time) with approximately the same result. I've seen that OPRD paper before, too. I've seen data to suggest
that methylsulfuric acid can be distilled unchanged at less than 2 mbar; I've not tried this yet though. I suspect the presence of impurities in the
in situ prepared methylsulfuric acid (i.e. excess methanol, water or sulfuric acid) are partly to blame for the problems. Excess methanol
gives dimethyl ether under the thermal decomp. conditions, which in turn raises the system pressure and messes up the reaction (I've had this happen
before actually). The presence of water should not be too problematic, as it is generated in the reaction itself. But removal of water can only
benefit the decomposition, as it limits the amount of methanol that can form at equilibrium (c.f. the above mentioned problem).
Then theres the problem of how to perform the reaction. All of the attempts I made were following Painkilla's procedure. I think this is not the
optimal way to do the thermal decomp. Considering dimethyl sulfate is unstable at the reaction temperature (it distills as it forms). Under reduced
pressure at about 40 mbar decomposition was characterised by a pressure rise, likely from formation of volatile and non-condensible dimethyl ether. 10
mbar seems to give a reasonable result. I wasn't willing to go much lower as the methylsulfuric acid may have began to pass over unchanged. Based on
my personal experience, I'd like to try again with several changes. Firstly, the use of anhydrous copper sulfate to dry the methanol-sulfuric acid
mixture prior to distillation would be advantageous as it is self indicating when the water content falls below a certain level (no longer turns
blue). Other drying agents might also be employed; I've previously tried sodium sulfate as per Painkilla, and also magnesium sulfate. Both were slow
to filter through a dried glass frit. Use of mechanical stirring (as opposed to magnetic stirring) might have prevented the drying agent being ground
to such a fine powder, so I'd like to try that at some point also. Then theres the purification of the intermediate methyl sulfuric acid by high
vacuum distillation. 2 mbar is the literature data, but I know I can hit 0.3 mbar on my schlenk line (so sayeth McLeod) so thats what I'll use (my
pump is rated to 1 micron but thats an unrealistic expectation on a system). This should rid the methylsulfic acid of any excess methanol (it'll be
kept under high vac for a bit before heating is commenced).
Once the methylsulfuric acid is in hand, I'd like to try the thermal decomp several different ways. Firstly I'll do it as Painkilla (and myself) have
done previously. Then I'll try my preferred method, which is akin to the preparation of diethyl sulfate in JACS 1924 p999-1001; dropping the
pre-formed methylsulfuric acid onto anhydrous sodium sulfate heated to the appropriate temperature. I like this method because the bulk reactant and
dimethyl sulfate product are only in contact with the high temperatures for a short period of time; its a flash pyrolysis.
Whilst on the topic, I won't bore everyone with the very real dangers associated with the target material. What I will say is that a solution should
be on hand to neutralise any spills whenever this material is being encountered, preparation or use. Conc. ammonia should not be used due to the
violent reaction that ensues. 5% ammonia has been found adequate for decontaminating glassware. 2M KOH is also adovated by some parties. The solution
I'd prefer to use is a 5:1:1 mixture of methanol, water and conc. ammonia. The methanol ensures miscibility with the alkylating agent, enabling
efficient neutralisation with the contained ammonia.
A note pertaining to cleanup: When following Painkilla method (and probably any other method based on this approach), the stillpot remains are stongly
acidic, likely a mixture of sodium bisulfate and conc. sulfuric acid (consider the balanced equation). This is an absolute bitch to safely neutralise
and decontaminate. The best approach I've found so far is to pour the mixture into a neutralising solution (I've been using 5% ammonia with added NaOH
to maintain alkalinity) and then drown the mostly empty flask in the same solution. This is every exothermic, so use a large volume of neutralising
solution.
Breifly talking of other methods... I'd like to hear of any first hand (or otherwise) accounts of using the SO2-CuCl2-MeOH system. I can see how and
why it would work, but I've seen nothing conclusive online. Someone mentioned it here in passing, mentioning someone at the hive had used it with
success. I'd like further confirmation that this actually works before trying it myself.
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aliced25
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Is it possible to isolate the Sodium Methyl Sulfate (per http://www.sciencemadness.org/talk/viewthread.php?tid=15837), purify the piss out of it, then acidify it with strong H2SO4 (carefully using the
theoretical quantity of acid, or thereabouts), thereby generating Sodium Sulfate and almost anhydrous Methyl Hydrogen Sulfate on Sodium Sulfate, with
some H2SO4 in the pot? That might remove some of the problem with getting the original water out, getting the pure precursor into the pot (without
high-vacuum purification) and having a water absorbent in there (Na2SO4) to keep the reaction going as water is produced?
| Quote: | | I suspect the presence of impurities in the in situ prepared methylsulfuric acid (i.e. excess methanol, water or sulfuric acid) are partly to blame
for the problems. Excess methanol gives dimethyl ether under the thermal decomp. conditions, which in turn raises the system pressure and messes up
the reaction (I've had this happen before actually). The presence of water should not be too problematic, as it is generated in the reaction itself.
But removal of water can only benefit the decomposition, as it limits the amount of methanol that can form at equilibrium (c.f. the above mentioned
problem). |
Without additional methanol or water, the reaction should be pushed to the right once it is started by removal of the product, shouldn't it? Neither
hydrolysis or methanolysis (or the production of the dimethyl ether) should be as problematic if they aren't being formed? Just thinking, not taking
the piss. Extra methanol isn't going to help, this is a basic decomposition 2MeHSO4 ==> (Me)2SO4 + H2SO4, as H2SO4 is already in the system with
Na2SO4 (from the decomposition of the sodium methyl sulfate, then it can't hurt (and according to the paper it shouldn't).
[Edited on 3-6-2013 by aliced25]
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IPN
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Just wanted to second this:
| Quote: |
The solution I'd prefer to use is a 5:1:1 mixture of methanol, water and conc. ammonia.
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I keep the same stuff in a washbottle in the fumehood as I do a lot of work with DMS and it works well for quenching any spills, used glassware before
washing, etc.
Also don't fear the reagents you use, just show respect.  You create more hazards
by being overly nervous.
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KrysHalide
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I've tried the CuCl/MeOH/SO2Cl2 in the past, with no success whatsoever even after a few attemps.. Do not put alot of expectations on it.
Never regret thy fall,
O Icarus of the fearless flight
For the greatest tragedy of them all
Is never to feel the burning light.
Oscar Wilde
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SM2
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Prey tell, why is Di-Ethyl Sulphate on order of more than almost 2 magnitude LESS toxic than DMS. Apparently Di Ethyl sulphate is relatively benign.
"Old men who speak of victory
shed light upon their stolen life
they - drive by night- and act as if they're
moved by unheard music." B. Currie
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Nicodem
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Perhaps because it is a couple of orders less electrophilic?
BTW, it is called "diethyl sulfate", not "Di-Ethyl Sulphate". Chemical names are not capitalized except for grammatical reasons (like when starting a
sentence). You also don't use the "-" thing after "di". The USA/UK orthographic difference in the f/ph usage is a national choice, but most scientist
nowadays use the "sulfate"-like writing (it helps to know this when doing keyword searches).
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SM2
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porpoer ununciation
Nicodem,
Is there an authoritative guide which firmly establishes the most common name of the reagent. Like a Merck Index. Not interested in IUPAC, just most
.something so the TV news dummies don't say "pseudo effehDREEN, or drine". Is there such an english guide to your knowledge. I just hate it when
peptide bioligists say "ppipper-ah-deen" when it's really "pie-pear-ah DEAN". Is there not an authority?
"Old men who speak of victory
shed light upon their stolen life
they - drive by night- and act as if they're
moved by unheard music." B. Currie
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aliced25
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Found an interesting article where it was describing the alkylating efficacy of methyl sulfate half-esters (http://www.ncbi.nlm.nih.gov/pmc/articles/PMC1765482/) which suggests the half-ester 'might' be useful for methylation of amines, if nothing else
(although I'd be interested to hear from Nicodem on this):
| Quote: | | When using dimethylamine as a common acceptor, methyl sulfate is quite reactive compared with other methyl donors, surpassing the trimethylsulfonium
ion by a factor of >100 |
Now that might just be for the amine group, but IIRC Nicodem used the trimethylsulfonium to alkylate a phenolic OH, so what would be the prospects of
this working for that purpose? I do recall an old article (Lewis, et al, 1929 'Methylation of Phenol by Dimethyl Sulfate' Ind. Eng. Chem., 22(1) -
attached) where they used various conditions, and with 1 mol of DMS and 2 of phenol, they got a 90% yield of anisole, which means the methyl sulfuric
acid must be capable of methylating the phenolic OH.
Seems strange to me that their conditions required minimal water and this article is talking about running the reaction in water. Any advice? The
half-ester would be considerably nicer to work with at home (much less mutagenic for a start and a one-step production procedure), although the longer
reaction time would put most off. Anyone got any suggestions? Say, for the sake of argument, for Rasberry Ketone
[Edited on 15-6-2013 by aliced25]
Attachment: Lewis.etal.Methylation.of.Phenol.by.DimethylSulfate.pdf (417kB)
This file has been downloaded 880 times
From a Knight of the Realm: "Animated movies are not just for kids, they're also for adults who do a lot of drugs." Sir Paul McCartney
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sonogashira
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Unless one specifically needs dimethyl sulphate, and instead requires merely a means of methylation, I would suggest looking into the method of
preparation of methyl bromide described here (http://www.orgsyn.org/orgsyn/pdfs/CV2P0279.pdf), in the notes section. I have successfully used this method, using 10% by weight of the amounts
stipulated, to methylate a tertiary amine in acetone to give a quantitative yield of the ammonium bromide, upon thoroughly gassing until the presence
of a white precipitate, then leaving the flask, stoppered, to age overnight, where I found a flask full of crystals the following day. I dried the gas
over sodium hydroxide, due to the unavailability of potassium hydroxide, and it seems to have worked admirably. Being a gas, one may lead the excess
safely outside with a length of pipe. The generation is easily controlled by the addition, or removal of a water bath, and the addition, or not, of
further sulphuric acid/methanol mixture to the sodium bromide. I did this outside, in my shed, on ca. 10 gram scale. A rubber stopper that I was using
has been blackened, but aside from that the experiment was uneventful. As regards disposal: I allowed the mixture to cool, added water, leaving it to
dissolve the remaining solid (which increases in volume during the experiment, especially-so as the flask is heated to dryness) added a dash of sodium
hydroxide to neutrality, and threw it on the dirt. It then began to rain, so I worried no more. I used a 500ml Buchner flask for the generation, 50ml
H2SO4 etc. (as per instructions). I used a huge excess above the demand of my requirements, but the ingredients are cheap, and I wanted to make sure
of success. If conducted again, I would use a series of Dreschel bottles to methylate different things, making full use of the gas, for which I have
more uses than the air. I highly recommend doing this in a shed, rather than a house, lest anything unexpected happens - but it is a very easily
controlled reaction, can be conducted cleanly, safely, and leads to very satisfactory, if not excellent results, and for practical pursuits, surely
favourable to dimethyl sulphate for use, or for preparation. Given the boiling point of 3 C, it may be worthwhile to condense some, and store it in a
propane cylinder, or similar.
[Edited on 21-6-2013 by sonogashira]
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Nicodem
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Quote: Originally posted by aliced25  | | Now that might just be for the amine group, but IIRC Nicodem used the trimethylsulfonium to alkylate a phenolic OH, so what would be the prospects of
this working for that purpose? |
The use of methyl sulfates for the methylation of phenols has already been discussed at this forum and some references were given. There are
discussions also about the use of ethyl sulfates - one recent thread even discusses the members experience in the N-ethylation of nitrite
with an ethyl sulfate. The forum abounds in information about alkylations with alkyl sulfates.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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