Sciencemadness Discussion Board

Sodium methoxyde by electrolysis?

Tacho - 5-3-2004 at 11:13

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]

BromicAcid - 5-3-2004 at 11:43

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.

chemoleo - 5-3-2004 at 16:33

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 :)

Marvin - 5-3-2004 at 17:37

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.....

Tacho - 6-3-2004 at 03:21

Two flasks, salt bridge?

One flask ends up with methoxide solution, the other with garbage?

[Edited on 6-3-2004 by Tacho]

chemoleo - 6-3-2004 at 10:26

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 :P

HELP!

Tacho - 14-3-2004 at 12:57

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).

BromicAcid - 14-3-2004 at 18:35

Quote:

There is a stream of tiny bubbles coming from the catode. Nothing seems to happen in the anode (where methoxyde should form).


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.

Marvin - 14-3-2004 at 20:54

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.

Kolbe Synthesis

Turel - 14-3-2004 at 21:10

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.

Tacho - 15-3-2004 at 03:46

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)


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]

NaOOCMe/MeOH Solubility?

Turel - 15-3-2004 at 13:39

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?

Tacho - 16-3-2004 at 03:30

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.

Kolbe

Turel - 16-3-2004 at 07:41

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

Tacho - 16-3-2004 at 08:52

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.

Never worked with MeI

Turel - 17-3-2004 at 13:06

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?

BromicAcid - 17-3-2004 at 13:56

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.

It works! Very beautifull reaction.

Tacho - 22-3-2004 at 14:07

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.

Image017.jpg - 9kB

Electrolysis of Methanol Solutions

a123x - 17-3-2005 at 00:03

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?

garage chemist - 17-3-2005 at 09:43

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?

a123x - 17-3-2005 at 12:04

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'!:mad:

[Edited on 17-3-2005 by chemoleo]

PainKilla - 17-3-2005 at 12:22

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]

garage chemist - 18-3-2005 at 12:49

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...

S.C. Wack - 18-3-2005 at 13:44

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.

garage chemist - 19-3-2005 at 02:19

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]

Re: Sodium acetate in methanol

Dave Angel - 24-3-2005 at 16:33

Today I have been electrolysing a saturated solution of sodium acetate in methanol. I used a computer PSU's 5.5 V line which managed to get a current of between 300 and 400 mA flowing, the current increasing as the electrolysis proceeded. I didn't use a salt bridge for this - it was done in a single pyrex bowl with the cathode (copper) and the anode (graphite) as close together as possible. The 'apparatus' was lightly covered in clingfilm to limit evaporative losses of the alcohol although ingress of moisture wasn't prevented with this set up. The current was passed and plenty of bubbling was observed on the cathode, much less so at the anode as it's quite a hunk of graphite (a crucible) with large surface area.

After ca. 4 hours I came back to the set up and found that there was a good deal of white precipitate in the bowl, with no appreciable difference in level of methanol. This could of course be due to displacement of the remaining methanol by the precipitate as it forms... or something else. So, I added some straight methanol to the bowl and dissolved up most of the precipitate, allowing the solution to remain saturated. Checking the pH of the initial solution against this with universal indicator solution showed that the solution which had undergone electrolysis had a higher pH than it started off with.

Another 5 hours or so and yet more precipitate! The solution was filtered and pH of both filtrate and a saturated solution of the precipitate in methanol were checked, although the precipitate was still quite wet and the tests both gave a similar result to the previous one - both solutions had higher pHs than saturated sodium acetate solution, so some more accurate work is needed.

I hope that Tacho is right and that NaOMe is being produced! :)

What's next? I'm building a new set of electrodes which will be separated by 2-3 mm. The set up will likely a 2L 3-neck rb flask fitted with a condenser and drying tube, the electrodes and a stopper in the final neck. Stirring is probably a good idea to prevent build up of precipitate in any one place, especially around the electrodes and one could even heat the flask if desired.

With any luck, fractional crystallisation could be used to separate the product(s) from the acetate.

I also should note that initial tests using a copper anode produced less than satisfactory results. Whilst there is a build up of blackness on the cathode over time which does not affect operation, a copper anode quickly becomes covered in a green layer which prevents any current flowing. Graphite seems to be doing a good job for now.

a123x - 31-5-2005 at 21:42

Sodium amalgam to make sodium methoxide. I recently read that sodium amalgam can be produced by electrolyzing a solution of sodium chloride, sodium hydroxide, and I would assume sodium acetate with a mercury cathode. Rather than release hydrogen, it forms the amalgam. I figure that this might allow for rather convenient production of sodium alkoxides by simply electrolyzing a solution of sodium acetate in an anhydrous alcohol using a mercury cathode. This will form the amalgam which, when the current is stopped, will react with the alcohol to form the alkoxide. Thus, with a small amount of mercury, one could make very large amounts of sodium methoxide since the mercury is recycleable in the process. Too bad mercury is toxic(although I imagine it shouldn't contaminate the methoxide much). I know molten gallium and liquid gallium/indium/tin alloys can dissolve aluminum to make a substance similar to aluminum amalgam and should be able to dissolve sodium but the thing I wonder is whether gallium can be used as a cathode to form such an alloy starting with a solution of sodium salt. I realize that gallium is fairly pricy but only a small amount is needed and I tend to prefer pricy to toxic, even if the toxicity is generally exagerated.

BromicAcid - 1-6-2005 at 07:20

A decent idea that has been discussed before, the major hurdle being that the solid amalgam formed only contains a few precent of sodium and therefore this process would have to be repeated many times before any major results were obtained, not to mention drying the amalgam before reacting with methanol otherwise adhering water would decompose the methoxide and additionally it takes some time for these alloys to react entirely with methanol.

a123x - 11-6-2005 at 22:06

I was actually thinking along the lines of making the sodium amalgam in a methanol solution and then just stopping the current to allow it to react.

Also, one of the main uses for methoxides is in nucleophillic substitutions with halogen groups on aromatic rings. So, I had a thought. Rather than make a duel chamber electrolytic cell with a salt bridge to use in making methoxide directly by an electric current(as opposed to the indirect method of making the amalgam and reacting with methanol), why not have a single chamber which results in the methoxide being reacted before it can reach the anode and be oxidized. For example, have a cell containing sodium acetate in methanol and also bromovanillin/bromoanisaldehyde/some other bromoaromatic dissolved in the methanol. Include a bit of copper powder as the catalyst. It'll need to be heated to fairly high temperatures as well and have a reflux condensor. The methoxide is formed at the cathode and reacts with the bromoaromatic rather than being oxidized at the anode. I assume that some will reach the anode and some will react with the bromoaromatic to form NaBr(which I'm guessing is insoluble in methanol, if it is soluble than there are issues with Br2 forming). If NaBr is soluble then a duel chamber cell might be used with one containing a bromoaromatic and the other containing the initial aromatic which will react with bromine formed to give the bromoaromatic. I figure if built properly it would ultimately give a low maintanence cell which would need little more than a fresh supply of methanol, aromatic, and electricity to produce methoxylated aromatics.

Now, I'm sure there's something wrong with this idea so feel free to let me know what it is.

jarynth - 4-9-2008 at 16:48

After reading the thread about Na production by Castner cell, I succeeded in electrolyzing NaOH with a Ni welding electrode as the anode. The main obstacle that remains is the removal of liquid Na above the cathode during the process. Instead, the Na/NaOH solution was simply allowed to solidify. The presence of sodium metal was confirmed by the observation of small yellow burning and fuming beads on contact with water.

Now, the analogous reaction with anhydrous MeOH (or EtOH, etc) would yield NaOEt (which dissolves) contaminated with NaOH (less soluble in the anhydous alcohol).

This procedure involves one simple electrolysis step, anhydrous alcohol and a means to separate NaOH and the alkoxide from the alcohol, thus should prove no harder than the alternatives presented above. The question remains, how to separate the sodium compounds and store the purified alkoxide.

@a123x: what happens at the cathode is always nice and smooth, but to skip an accurate analysis of the anodic processes would be unforgivable. With halogenated benzenes you'd get biphenyls and whatnot at the cathode: polymeric mess. Plus, methoxylations are more successfully carried out in weakly ionizing solvents, so you'd get only a tiny current flowing. Most importantly, the NaOAc is a weak base and will probably cleave the halogen before you even reach reflux (but we'll need an expert opinion on this). The number of electrolysis products increases exponentially with the number of species in your flask. The reaction won't be as clean as you purport.

[Edited on 4-9-2008 by jarynth]

[Edited on 4-9-2008 by jarynth]