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Author: Subject: Using [Cu(Cl)4]2- to activate Al for reductions
Klute
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[*] posted on 25-3-2007 at 13:56
Using [Cu(Cl)4]2- to activate Al for reductions


Thanks to Woelen's site where he talks about "activating" Al with tetrachlorcuprate complex, Bio had the idea of replacing Hg salts by this complex to acheive Al reductions.. It reduces H2O to H2 for sure, but the question was if it could reduce a compoud such as nitromethane for example. The Al could react only with the H2O, any remaining Cu2+ could complex the formed amine, alot of possible problems could happen. I did a few test-tube tests. here's my post from the Zonez.

Posted - 03/25/2007 : 2:34:42 PM
--------------------------------------------------------------------------------

I've FINALLY manadged to run a few tests... and the news are good...

About half a mL of a sat. fresh CuSO4 solution was introduced in a test tube, followed by enough sat. NaCl solution to cause a change in color, from nice deep CuSO4 blue to a light green/Feso4-style color. Then about 0.1 of shredded Al was added. After only a few seconds, light bubbling started, with the apparition of some black deposit. After perhaps a minute, the solution had turned transparent, the AL was bubbling (very fine bubbles, much less vigorous than an AL/Hg) and there was some voluminious black and dark brick-red deposits, by spots. They looked like mold. It wasn't a layer of the stuff, just some dark spots. The test tube had heated up to be warm on touch, about 30-35°C. Then about 0.5ml of pur Nitromethane was added, and the test tube shaked constantly. When shaking was stopped, the nitro seperated as a clear layer on top. The Al was already starting to decompose in grey sludge. The tube was shaked regulary for about 5-7 minutes. The nitro layer seemed to dimish, and the test tube seems to heat up more and the bubbling was more vigorous. This could just have been because the reaction with water was kicking in. No smell above the test tube, but as the CuSO4 was acidic (H2SO4 from preparetion), the possibly formed MeNH2 would be as a salt. It was left as is, shaken every minute or so, while another test tube was prepared.
The complex was prepared as in the first test, using roughtly the same amount. Al was introduced, and left to fizz until the solution had turned transparent. Again, the same black and dark red deposits. After 5 minutes with occasional shaking, the solution was decanted off, and 0.5mL of MeNO2 added. This covered the Al. To my releif, bubbling continued, slowly, gently. Some sludge started appearing, but didn't dissolve in the nitro, so the Al/black and red deposits/ sludge form a semi solid mass stuck to the bottom of the tube. Bubbles were evolving from every part of the Al. This tube was definitively hotter than the other one, but didn't bubble more violently.

A third test tube was prepared, using 2x the amount of Cu2+ and Cl- to prepare the complexe, but the same amount of Al. It was left as is, with occasional shaking.

Some 40% NaOh was added to the first test tube. The reaction had nearly died down altough there was alot of Al left, and a little nitro. There was alot of grey sludge. As soon as the base was added, the Al started bubbling and decomposing. No MeNH2 smell. After a few minutes, bubbling was violent, and the grey sludge turned dark green colour, you know that yucky sewer color. But, then, eureka, MeNH2 smell! It smelled just like a finish Al/Hg, although less stronger...

The second test tube was still slowly reacting, but bubbling a lot less. The nitro layer had dimished, and still no smell of MeNH2. Alot of grey sludge had evolved. Some more Cu2+ and Cl- solution was added, to see what the complexe would do on any formed amine (Seeing the smell in the first tube, i was sure it had worked here too, if not better.) The solution turned dark green/yellow, a bit like a old curry sauce or something, and the bubbling speeded up a little. After a minute, NaOH solution added. The Al started decomposing, no MeNH2 smell but another strange unplesant smell. Even after 5 mintues, there was still this unpleasant smell, and no amine.

In the third tube, there was much more deposits then the other, the solution had clared up.

Conclusions:

The Al is "activated" in such to be able to reduce compounds like MeNO2, although the reaction seems to be very slow. Apparently, the presence of water from the complex solution doesn't seem to prevent such reductions. But, for any quantitative transformations, more experiences need to be done.

The [Cu(Cl)4]2- is turned into some strange compound of the surface of the Al as dark-red spots. It doesn't look like metallic Cu IMHO. Not at all shiny, and dark bric-red color. Black spots are also present, although this *could* be a impurity in the Al that reacts. The AL used leaves nothing after reaction with 50% NaOH (no visible Fe, Cu, Sn, etc). The fact that the green color disappears after a few minutes and these deposits appears proves so. Also adding more complex forms more deposits. So the Cu cannot form complexes with formed or added amines for example. The fact that MeNH2 is smelled when no complex is added after is formation, but no smell is apparent when more complex is added suggests this. But, as the second reaction was run with much less water, it isn't possible to affirm the absence of MeNH2 smell was entirely du to the complex addition.

Adding more complex *seems* to quicken the reaction, but causes more deposists that cover the Al surface. Apparently, seeing that the reactions continues when the complex solution is decanted when discolred, the "activation" is catalytic, and not stoechiometric. And the Al decomposes to Al(OH)3 as in an Al/Hg and doesn't passivate again.

So, I think that with more work and experiemnts, this complexe could be used to replace Hg salts in Al reductions. Although it isn't sure that it could reduces imines and such, and specific reaction conditions need to be found.
As for now I personaly don't have much avalible time, and other interests, I'll surely won't continue researching into this at the moment. I greatlt encourage other people interested to do so, it could be a great change, and avoid alot of people using toxic Hg salts.. And even if alot reagents and time are spent into this, that's were they are the best spent! Imagine if reductions could simply be done with AL, Cu and salt!

(Edited title)

[Edited on 26-3-2007 by chemoleo]
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[*] posted on 25-3-2007 at 16:38


A silly question perhaps -
but did you try MeNO2 on Al/NaOH on its own? How is the complex beneficial here?




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[*] posted on 25-3-2007 at 19:39


As a catalyst (like Hg, but by a clearly different mechanism).

Well okay, catalyst is a misnomer since the copper ions get reduced, but pffbt...guess that's why you are anal chemists :rolleyes: :P

You might try HCl + CuCl2.2H2O. Make it nice and concentrated to reduce the H2O side reaction. For that matter, does Cu dissolve in nitromethane (or others) any?

Tim

[Edited on 3-25-2007 by 12AX7]




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[*] posted on 25-3-2007 at 22:08


.....A silly question perhaps -
but did you try MeNO2 on Al/NaOH on its......

Methylamine produced from nitromethane
via Aluminum amalgam is well known.
Fe or Zn and HCl will do the same.

..... How is the complex beneficial here? ...

The complex reacts with surface Al oxide
thereby "activating" the aluminum

The black spots on the Al surface
are also seen with inferior Al found in
disposable baking pans. The spots spread
and reaction fizzles out.

As Al will react preferentially with water
instead alcohols try another test using methanol and little to no water. At least
saturated aqueous solutions mixed into
the MeOH and keep pH basic
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[*] posted on 26-3-2007 at 08:54


The black spots are definitively not the same as the Fe residu from impure Al, when reacted with Hg2+ or NaOH, this Al doesn't leave any residu... I really think it comes from the complex

The problem with the complexe is that it must be made in neutral/slightly acidic conditions, as adding base before reacting with the Al with precipitate Cu(OH)2... Of course, the base can be added once all the complex has reacted with the Al (incolor solution), but this might speed up the Al degradation. I guess only a very small amount of base should be present.

The Cu complex didn't seem to dissolve in MeNO2 when decanting the complex solution after reaction. MeOH should be beneficial, maybe it will react slower with the Al then water, and help achieve a homogeonus solution diluting the traces of water.

CuSO4.5H2O + NaCl is very easy to prepare, and if the CUSO4 is cleaned up well enough to remove most of the H2SO4 from it formation, less base will be required to atain neutral/basic pH. Adding HCl to CuCl2 and neutralising it would achieve the same result, with more water/ salts.

The most reasonnable path would be reacting the Al with the complex (I can't figure the reaction though, finding a minimal stoechiometric amount would be nice), decanting, possibly rinsing with a little MeOh, then adding the compoud to be reduced in MeOH... The reaction seems much slower than with Hg salts though, and I guess more exces would be necessary...

As I mentionned, I won't be spending the little available time I've got on this, I've got a few more projects I'm more interested in...
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[*] posted on 26-3-2007 at 13:26


.....The most reasonnable path would be reacting the Al with the complex (I can't figure the reaction though, finding a minimal stoechiometric amount would be nice), decanting, possibly rinsing with a little MeOh, then adding the compoud to be reduced in MeOH... The reaction seems much slower than with Hg salts though, and I guess more exces would be necessary.......

Yes, agreed entirely.

As far as minimal amount goes I would try to dissolve
the copper sulfate as much as possible in the hot
MeOH then add water by drops. Immersing the Al
into a say 2-5% solution or maybe stronger will
provide more than a catalytic amount so that
should be enough Cu complex. Also let it react a few minutes
befor decanting.

When using HgCl2with thicker(say0.30mm) than foil aluminum the reaction sometimes won't go unless the temperature is elevated. Try keeping the "activating" solution covering the Al at 40-50 deg. before decanting or I suppose you could also suck it out. If the reaction gets sluggish raise to reflux temp (65deg
for MeOH). Don't let the activated Al contact the air
for more than a few second.

Adding the MeNO2 in portions will help, then the formed
amine will keep the pH up. In the beginning only a very small amount of base is needed to make the solution basic.
Certainly a lot less than would react with only say a couple
percent of the Al.





..
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[*] posted on 1-4-2007 at 02:14


Quote:
Originally posted by Klute
So, I think that with more work and experiemnts, this complexe could be used to replace Hg salts in Al reductions.[Edited on 26-3-2007 by chemoleo]


i hope you are aware that the tetrachlorocuprate(II) complex does not seem to work in a catalytic manner - see this experiment.
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[*] posted on 1-4-2007 at 02:55


I disagree. Although some copper metal is precipitated, the solution remains green, it does not become colorless: "When the reaction comes to an end, then the dark green liquid can be decanted from the precipitate at the bottom ...". Although, it isn't clear if the aluminum or copper is the limiting reactant in this reaction. Certainly, a disproportionate amount of hydrogen is produced compared to the amount of copper precipitated.

It may be that every amount of copper has some action towards oxidizing the aluminum before being precipitated from the solution as copper metal. That would be a cross between catalyst and reactant. But how?

It could be that the reaction mechanism is such that tetrachlorocopper ions truely catalyse the reaction, and the catalyst is then "poisoned" by reduction to Cu(0), either by plating onto the aluminum or reduction by transient hydrogen (H or H2). It could also be that the Cu(0) plating is the true catalyst, and the aluminum gets eaten away from under it; but the reaction begins almost immediately, rather than not at all or delay due to the Al2O3 layer, so this seems unlikely. This would be easy to test by rubbing aluminum with a copper wire under salt solution.

Certainly, it isn't AlCl4 or other chloro-species, as the strong chloride solution doesn't do anything.

Tim




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[*] posted on 1-4-2007 at 14:17


In my experiments, the solution discolored itseld pretty quickly (1-3min), and the Al kept on bubbling. But the reaction ws very slow, that's why i was surprised by the term "vigorous reaction"... Possible explanation is that I used less CuSO4, as I used freshly made solution, from recrystalization of home made CuSO4.5H2O, so i didn't now the exact amount of copper sulfate in there.. it was a near saturated solution, though.
But i don't understand why the fact that in Woelen's experience the solution stays green implies that it isn't a catalytic reaction...

My guess is that just a certain amount of complex is needed to destroy the Al2O3 layer, and that the rest could partially, or totally, be reduced by the Al. As in woelen experience, the solution stays green when the reaction is complete, I would think that the Al reacts only partially with the complexe, and of course more so with the water. Without any maths or whatever, I find it strange that the 2 balls of Al didn't reduce all the Cu2+ from the little CuSO4 introduced.. Then again, i just eyeball it, not calculating the weights and stoechiometrics, etc
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[*] posted on 1-4-2007 at 16:02


.....It could also be that the Cu(0) plating is the true catalyst, and the aluminum gets eaten away from under.......

If the stirring was very vigorous the "spots" should be
sloughed off leaving the activated Al to react .


.....My guess is that just a certain amount of complex is needed to destroy the Al2O3 layer, and that the rest could partially, or totally, be reduced by the Al..........

Yes, and if the deposits are removed by stirring then all
thats required is enough complex as is needed to react
with most of the surface Al oxide.

I have seen some published procedures that use as
much as 17moles to1 (Al to substrate) and two to three
times excess is normally enough in reducttive aminations.
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[*] posted on 2-4-2007 at 02:27


I did all the experiments without any organic compound in the liquid. It might be that the presence of this organic compound has an adverse effect on the speed of the reaction.

I like the idea of using this reaction in the reduction of organics, but of course, it might be that this cannot simply be applied as such and a little more research is needed. I can though, that without the organic, the reaction is quite vigorous, especially if it is somewhat scaled up (self heating accelerates the reaction in that case).

I agree with 12AX7 that this most likely is a mixed catalytic/normal chemical reaction. If it were completely catalytic, then no copper (II) ions would be used at all, if it were complete standard, then for each 3 mols of copper(II) 2 mols of Al are used. The reality is somewhere in between (the production of hydrogen shows that). When I did the experiment, I was quite surprised about that reaction. Such common compounds, and such unexpected results.




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[*] posted on 3-4-2007 at 00:38


Well, actually if a compound gets consumed during the course of a reaction, then it can not be called a catalyst by any definition. Since the tetrachloroaluminate anion gets reduced to elementary copper it is not a catalyst. It fits the definition of "activator" or something similar. Like I2 in the activation of Mg turnings for a Grignard preparation, or like HgCl2 in the activation of aluminium by amalgamation…



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[*] posted on 5-4-2007 at 10:51


It is commonly known that mixture CuSO4+NaCl dissolves Cu, giving
soluble (in reaction conditions) [CuCl3]- and similar complexes.
After adding water Cu2Cl2 is precipitated.
Propably Cu(II) must be first reduced to Cu(I), to gain "clearier" machanism of reaction. I think that MeNO2 reducion process is a side reaction in this conditions. And as was mentioned - Cu do not
amalgamate Al.
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[*] posted on 5-4-2007 at 15:21


.......I think that MeNO2 reducion process is a side reaction in this conditions........

Yet the Aluminum continued evolving Hydrogen after the
copper solution was removed and the nitromethane added.
This points to the Al being responsible for the reduction.

Maybe a side reaction occurs simultaneously with the
reduction which occurs on the aluminum surface.
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[*] posted on 5-4-2007 at 19:28


Look at this excerpt from an old SM thread on the reduction of imines.

Certainly adds creedence that the CuO is the real "activator"
as suggested by 12AX7.

..............................................................................................
...............About catalytic hydrogenations, there is a patent, which has dissapeared from the US patent office website which details a catalyst which could potentially bypass the need for fancy and expensive catalysts made with rare metals like platinum, palladium, and rhodium.

It is made by reacting copper sulfate pentahydrate with barium or calcium hydroxide to get a........... catalyst consisting of cupric oxide........... and calcium or barium sulfate.

The patent is US No. 2,828,343, but like I said, I couldn't find it on the website. I would like to know what kinds of pressures this catalyst requires andhow much of it is needed per mole of imine reduced......................

BTW

I was able to find the patent on espacenet for but for some
odd reason I am unable to copy the link.





[Edited on by bio2]
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[*] posted on 5-4-2007 at 19:47


Quote:
Originally posted by bio2...
The patent is US No. 2,828,343, but like I said, I couldn't find it on the website. I would like to know what kinds of pressures this catalyst requires andhow much of it is needed per mole of imine reduced......................


Like this?

Attachment: US2828343A1.pdf (353kB)
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[*] posted on 7-4-2007 at 10:39


Quote:
Originally posted by bio2
.......I think that MeNO2 reducion process is a side reaction in this conditions........

Yet the Aluminum continued evolving Hydrogen after the
copper solution was removed and the nitromethane added.
This points to the Al being responsible for the reduction.

Maybe a side reaction occurs simultaneously with the
reduction which occurs on the aluminum surface.


Before we go too far with speculations two facts should be noted:
- Aluminium reduces water to H2 as well as it reduces many other compounds, but (un)fortunately the oxide layer prevents water and other substances to reach the its surface. Hence the apparent inertness of aluminium toward water and other oxidants (like the organic substrates to be reduced).
- Aluminium does dissolve in water as well if the oxide coating is etched. This can happen by various mechanisms, either by H3O+ or OH- ions dissolving the hydrated oxide layer (like dissolving Al in HCl or NaOH), by amalgamation and probably other mechanisms as well.

I would say that with the tetrachlorocuprate the creation of a short circuited galvanic cell from the Cu plating is why we get the aluminium to dissolve. The chloride ligand is probably needed only to facilitate the etching of the oxide layer so that the Cu2+ gets reduced and form tiny Cu "electrodes" (aluminium also forms a complex with the chloride just like Cu2+). On this Cu "electrodes" while in contact with Al the redox potential is that of Al (being essentially a galvanic cell) thus reductions can occur on its surface. But it is not to generalize that every reduction that works with aluminium amalgam will also work with cupper etched aluminium – the overvoltage of the system will be different due to different surfaces on which the reduction occurs. Just my hypothesis, while it is on others to do the experimental. :P
A simple test for part of the hypothesis would be to try this with another ligand for aluminium cations instead of chloride.




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