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Tacho
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Sodium methoxyde by electrolysis?
If I do electrolisys fo a NaOH solution in methanol, the Na that go to the anode will react with methanol to make sodium methoxyde?
EditbyC: title
[Edited on 1-6-2005 by chemoleo]
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BromicAcid
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The Na+ will go to the cathode and turn to sodium metal wich will instantly react to form sodium methoxide. However on the other hand the OH-
that's going to the anode will undoubtably form a water molecule through some process I would guess and destroy the sodium methoxide. For some
time I was facinated by a similar reaction. LiCl in methanol, the Cl2 gas produced would react with methanol and the lithium would make the alkoxide
and you have fertile grounds that may lead to a wilamson ether synthesis which would generate a molecule of LiCl as a byproduct which would recycle
back into the reaction. Of course you would end up with quite the odd mix at the end if you kept adding LiCl as it depleated.
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chemoleo
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Just a thought - my initial guess was, just like yours Bromic, that the product H2O would backreact with the alcoholate. However, don't you think
it's possible that nascent H2O would be electrolysed itself, into H2 and O2?
In that case you would remove the unwanted reaction product right away... hence leaving you with sodium methoxide after all.
Just a thought 
Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
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Marvin
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OH- will be oxidised, not reduced, so it doesnt automatically form water. What might well scupper the reaction though, is this could oxidise ethanol
to aldehyde/acetic acid. With a decent membrane this might be useful route to ethoxide, but the resistance of the cell is probably quite high.
Overall, I think finding magnesium metal, and converting the sodium hydroxide to alkoxide that way would be far easier. Now if only there was an easy
way to regenerate the magnesium metal.....
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Tacho
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Two flasks, salt bridge?
One flask ends up with methoxide solution, the other with garbage?
[Edited on 6-3-2004 by Tacho]
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chemoleo
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So are you going to try it? Just obtained a Pt coated Ti electrode... might be useful ey?
In fact, another one I wanted to try for some time is trichloroacetic acid, to get hexachloroethane (Kobe reaction). Dont ask me what I will do with
it though 
Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
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Tacho
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HELP!
Anybody knows a test to confirm that I have sodium methoxyde in methanol solution?
I have been doing the electrolysis for the last 5 hours in two flasks connected with a glass salt bridge. The liquid in the anode flask has some
petroleum ether on top to prevent atmosferic water contamination.
BTW, NaOH dissolves much better in methanol than ethanol. The resistance is high though: 24V renders only 17mA. There is a stream of tiny bubbles
coming from the catode. Nothing seems to happen in the anode (where methoxyde should form). The anode is copper and the catode is copper and carbon
(don't ask).
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BromicAcid
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| Quote: |
There is a stream of tiny bubbles coming from the catode. Nothing seems to happen in the anode (where methoxyde should form).
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Like I said before the sodium methoxide should be forming at the cathode. You know the saying, REDCAT, reduction occurs at the cathode? Oxidation
involves loss of electrons, reduction involves gain. The reaction:
Na+ +e- ---> Na(s)
Involves sodium gaining an electron, therefore it is reduction, therefore it is occuring at the cathode. Those bubbles you see, they're hydrogen
hopefully from the sodium being produced there reacting with the methanol and producing the desired sodium methoxide. So don't throw away the
'junk' that's forming at the chathode.
By the way, I don't know of any quantitative tests to determine if what you're producing is the methoxide, sorry.
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Marvin
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Part of the problem is that sodium methoxide will be in equilibrium anyway, so really I think we need a test war water.
If you made a solution of magnesium ethoxide (and filtered), you could add this until you nolonger got a ppt. Compair with the same solution
unelectrolysed (same taken from same batch), and the difference would be due to loss of water by electrolysis. Hopefully.
I'm still worried about the aldehydes/acids that form.
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Turel
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Kolbe Synthesis
Organic acids will also be electrolyzed. The anions will get oxidized to a carbonyl radical which will immediately decompose to CO2 and the R radical.
When this is exploited to join two R functions, it is called the Kolbe Synthesis or Kolbe Reaction.
Any oxidized methanol to formic acid would be further oxidied to CO2 and H2O or be electrolyzed to CO2 and H2.
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Tacho
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| Quote: | Originally posted by BromicAcid
Like I said before the sodium methoxide should be forming at the cathode. You know the saying, REDCAT, reduction occurs at the cathode? Oxidation
involves loss of electrons, reduction involves gain. The reaction:
Na+ +e- ---> Na(s)
(snip)
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Ouch!
How embarassing... Sheer stupidity. Thanks Bromic.
Then, my methoxide, if any, was not protected from air humidity and is probably ruined in the... cathode.
Marvin, I don´t have magnesium, so I can´t test your Idea.
Turel, if I undertand right, you are saying that I don´t have to worry about formation of aldehydes/acids?
It occurred to me that sodium methoxide could be extracted from methanol/NaOH with some non-polar solvent, I´ll research that, but would apreciate
any comments.
[Edited on 15-3-2004 by Tacho]
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Turel
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NaOOCMe/MeOH Solubility?
If sodium acetate is appreciably soluble in ethanol, you might have better results with your salt bridge. I checked my CRC and google, and I cannot
seem to find data on NaOOCMe solubility in alcohols. I don't have my chem dictionary on hand, so I honestly do not know if it is.
Charge the catholyte chamber with pure alcohol, and the anolyte chamber with NaOOCMe in alcohol solution, seperated by a salt bridge. Apply voltage.
The sodium ions should be able to pass through your salt bridge, to the cathode where they are reduced to elemental sodium in the catholyte chamber.
The nascent sodium reacts with the alcohol and produces sodium alkoxide and hydrogen gas. The alkoxide ion is prevented from traversing to the other
half cell and getting oxidized by the salt bridge.
In the anolyte chamber, the acetate ion undergoes Kolbe-like degradation, becoming oxidized and decomposing to CO2 and a methyl radical. Methyl
radicals join to form ethane, which evaporates away.
What do you think?
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Tacho
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I put some crystals in methanol and they seemed very happy to dissolve completely and quickly.
I like the idea very much because sodium acetate is probably neutral or acidic, and sodium methoxide should be basic. That will give me an easy
indication of what is going on in the flasks.
I don't know about pure methanol, I think it won't be conductive enough. I'll try.
Incidentally, I bought this sodium acetate to make some pure ethane by aqueous electrolysis. BTW, electrolysis in water gives ethane and CO2.
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Turel
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Kolbe
Just as it should. The acetate ion is oxidized, giving a radical on the oxygen that was once negative. The molecule simply collapses, with the oxygen
radical extending an electron to the positive polar center of C in the carboxyl, and that carbon losing it's bond with the R entity, in this
case, a methyl group. So the decomposition products are CO2 and CH3 radical, which combine to form ethane.
Sodium acetate is very slightly basic in water solution.
If you are worried about the conductivity of pure methanol, poison the charge with electrolyte. A small amount of sodium acetate, or sodium hydroxide
preferably, will give a significant increase in conductivity, with a low to zero decrease in product quality. You don't need much electrolyte,
because as the reaction progresses, sodium ions in the catholyte chamber will accumulate, continuously lowering the resistance of the catholyte half
cell.
-T
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Tacho
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| Quote: | Originally posted by Turel
Just as it should. The acetate ion is oxidized, giving a radical on the oxygen that was once negative. The molecule simply collapses, with the oxygen
radical extending an electron to the positive polar center of C in the carboxyl, and that carbon losing it's bond with the R entity, in this
case, a methyl group. So the decomposition products are CO2 and CH3 radical, which combine to form ethane.
(snip)
-T |
Wow!... If you say so.
Turel, I know this is off topic but, since you seem to be rather involved with chemistry, you wouldn't, by any chance, have anecdotal information
on the true dangers of methyl iodide, would you? Not MSDS data sheet stuff, but real world evaluation.
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Turel
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Never worked with MeI
I have never personally worked with methyl iodide, so I could not provide a personal assay of it's behaviors. It is a strong methylating agent,
and is known to be carcinogenic. It should have a lower vapor pressure than methyl chloride, and I would thus assume it would be less volatile.
For what it is worth, I would take precautionary measures to not get it on my skin, and to not breathe it or get it in my eyes. Basically, avoid
exposure as well as you can. I don't think MeI poisoning is an immediate issue, but rather a problem with cumulative exposures.
The reason I have more information on cyanides and NOx and N2H4 is because I work with them. But I have never had to work with MeI, and I try to avoid
pure methylating agents in the lab. I don't know how much longer I can avoid their use because I have a few experiments that require their use,
as I cannot get the methylated reagents directly. So I suppose we will see?
Anyone else care to comment on MeI?
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BromicAcid
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I used methyl iodide a coupld times in my orgo class. It came in a tiny reagent bottle maybe 35 ml and had a sterotypical skull and crossbones on it.
We just wore the regular gloves and goggles and took it out with the mechanical pippets under the fume hood like we did everything else. It was a
liquid and I'm assuming if I would have misshandled it I could have been in trouble. But it didn't fume in the air, it didn't behave
unnaturally, just handle with care.
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Tacho
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It works! Very beautifull reaction.
A solution of sodium acetate with some phenolphtalein in methanol, when electrolysed, produces a beautifull red color that slowly covers the copper
electrode and tints the electrolyte.
I assume this is due to methoxyde formation.
I attach a picture of the setup.
Notice that the picture is just to show the setup I used. The reaction in the picture is my failed electromethylation, hence the brown color in the
cathode.
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a123x
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Electrolysis of Methanol Solutions
Today I was trying to make magnesium methoxide by refluxing a mixture of Mg powder and methanol. Things were proceding excessively slow and after a
few hours very little of the MG seemed to have dissolved. So I was thinking about other possible methods for producing methoxide ions that
wouldn't require elemental sodium. One though I had was running current through a solution of sodium hydroxide in methanol. However, I then
rapidly concluded that this might produce methoxide ions but that they likely would then be oxidized at the positive electrode of the cell. Then it
started me thinking about what methoxide ions would be oxidized to. I'm thinking that it'd end up forming formaldehyde but I really have no
idea as I don't really know all that much about electrolysis driven reactions. Any thoughts as to what the results of electrolyzing a NaOH in
methanol solution would be?
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garage chemist
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Magnesium only reacts with methanol if the water content of the methanol is below 1%. Try drying it with anhydrous K2CO3 or better CaO or even better
CaC2.
What do you need a methoxide ion solution for? Why does everyone want methoxide?
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a123x
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I think the methanol should be rather dry. I got it in the form of gas line antifreeze which I would think would be anhydrous since its purpose is to
absorb water in a car's fuel system.
Anyway, I want methoxide for trying to convert dichlorobenzene to dimethoxybenzene or at least 1 methoxy-4chlorobenzene.
Edit by Chemoleo: SEARCH next time before you post new threads, a123x. Why do other people have to do it
for you just because you are too lazy? It's not the first time either!
Neither are the mods your personal 'thread merging servants'!
[Edited on 17-3-2005 by chemoleo]
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PainKilla
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I was just reading in my organic chem book that Cl substituion is near impossible without a nitro or other group there. I think it would be best to
nitrate the bezene, (hoping it doesnt nitrate in more than one spot) and from there attemp the methoxidation. once again, hopefully the nitrate
won't be affected.
I do know for sure though that in the p-chlronitrobenzene the Cl can be substituted for a methoxy! I am not sure if itll work with two Cl's and
the nitro next to them (i remember it not working well) However, you should then have the chlromethoxybenzene. And of course then simply make the
aldehyde  ... (which is pretty easy from there)
[Edited on 17-3-2005 by PainKilla]
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garage chemist
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Today I started a little experiment to see how the reaction between Methanol and Magnesium goes.
I put some coarse Mg powder (0,5mm grain diameter) into a pear shaped 25ml flask and added 5ml of reagent grade Methanol. I heated it and could not
see any hydrogen evolution. To remove the water that was apparently preventing the reaction, I added two pinhead sized amounts of sodium, which
rapidly dissolved. The Methoxide would then react with residual water. I again heated the mixture and after a few minutes, I could see a very slow
stream of hydrogen bubbles coming from the magnesium. It stopped almost completely when the methanol cooled down again.
So we see that even totally anhydrous methanol reacts only slowly with magnesium and the mixture has to be refluxed for several hours to produce
usable amounts of methoxide. But it works!
Now, can someone tell me what reactions are possible with Methoxide? Replacement of halogen with a methoxy group can't be the only application...
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S.C. Wack
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A Grignard initiator will work here as well. A little iodine will immediately move the reaction forward at STP. The alkoxides are useful in many
reactions where a strong base/nucleophile is a good thing.
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garage chemist
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Today I looked after the reaction and there was a white precipitate. Apparently magnesium methoxide isn't too soluble in methanol.
But I used only a very small amount of methanol, with more methanol the methoxide will surely stay in solution.
[Edited on 19-3-2005 by garage chemist]
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