Author Topic: corrosion, dissolving metal, and the Al(Cu) couple  (Read 1905 times)

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ning

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corrosion, dissolving metal, and the Al(Cu) couple
« on: November 28, 2003, 10:54:00 PM »
This is ning's musing on the possibility of Al(Cu) couple replacing Al(Hg)....it is written in the style of an article. Comments from those in the know would be most appreciated.

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What do corrosion and dissolving metal reductions have in common?

How rust can help us find a substitute for the aluminum-amalgam


What is corrosion?

   Corrosion is rusting. But what is rusting? At its most basic, corrosion is the electrolytic transfer of electrons from a metal to another species capable of accepting them. A redox reaction, with the metal being oxidized, and something else being reduced. Usually in corrosion studies, dissolved oxygen is the species reduced, but also are H+ ions (acid corrosion), and water.

How does corrosion relate to dissolving metal reductions?

All dissolving metal reactions can be considered special cases of corrosion. In the Birch reduction, the metal is "pre-corroded"--that is, as the metal dissolves in ammonia, it loses its electrons, but the electrons are not consumed immediately, due to ammonia's special ability to stabilize free electrons in solution. Other dissolving metal reductions (Al-Hg, Clemmenson, etc.) proceed more normally, in that the metal dissolves only as its electrons are consumed in reduction. In an aluminum amalgam reduction, you could very well think of the ketone or imine as "rusting" the aluminum foil, and being reduced in the process.

So how does corrosion or a dissolving metal reaction work?

Corrosion always has two components: a dissolving process, where the dissolving metal ions leave their electrons to the bulk metal, and an electronation process, where the bulk metal gives electrons to a reducible species. With nothing else to electronate, water or other protic solvents can be reduced, causing the evolution of hydrogen. When sodium metal is dissolved in methanol, the methanol is reduced to OMe- and monatomic hydrogen. The hydrogen adsorbs to the metal surface until it finds a mate to join with and form an H2 molecule, then escapes as a gas. The OMe- ion is solvated by the polar methanol, and its charge is balanced by the Na+ ion that results from the metal dissolution. This is the usual way of producing alkoxide bases, in fact:

2 M- + 2 ROH   -->  2 M.H + 2 RO-  --> 2 M + 2 RO- + H2

If, however, there is a reducible species in protic solution, then there is no reason for the hydrogen produced to be wasted as a gas. The monatomic hydrogen adsorbed on the metal surface is a highly reactive free radical, and will readily attach to certain functional groups. The actual process for reduction of a double bond is something like this:

R=R + H.M.H  -->  HR-R.M.H   -->  HR-RH + M

   Metals have a rather low affinity for hydrogen, and would like very much to give them away to things like carbon and oxygen, which hold hydrogen much more tightly. So the end result is that the metal dissolves to a salt or oxide, while some other species is reduced, with the hydrogen being liberated from the solvent, usually water.
   When performed in aprotic solvent, there is no hydrogen source, so reduction consists of dimerization or condensation. The only species capable of receiving electrons are the species being reduced, so they become radicals and condense, much as the monatomic hydrogens do in the previous case. The Wurtz coupling of two alkyl halides is a good example of this type of reaction. The halide is stirred in inert solvent with sodium metal, and as the metal dissolves, it electronates the halides. The halogens are very electronegative and good leaving groups, so they grab the electrons and leave as ions. The remaining alkyl radicals then rapidly combine. The negative halogen ion's charge is balanced by the positive metal ion, leaving the solution neutral:

   2 M- + 2 R-Cl   -->   2 R. + 2 Cl-  -->  R-R + 2 Cl-

   When a dissolving metal reaction is performed, the pH is of some importance. Apparently, the aluminum amalgam reduction is done in alkaline conditions because in hydrogen generation is preferred in acid, no doubt due to the high concentration of H+ ions floating about, ready to receive electrons.


What is the difference between an amalgam, a couple, and just plain metal?

   All of them reduce by the same mechanism; however, coupling and amalgamation can significantly improve the efficiency of a dissolving metal reduction by separating the processes of electronation and dissolution into two different areas, and onto different metals. When two different metals are placed in intimate contact, a galvanic cell is produced. The difference in work function between the two materials makes one of them tend to absorb electrons from the other, and this metal becomes the cathode, where the electronation and reduction takes place. The other metal, usually the bulk, dissolves and produces the electrons. This galvanic cell has a certain potential, and if that potential is sufficiently more negative than that of the reducible species, electron transfer will occur.
Interestingly, even in homogeneous metal local stress and impurities cause certain parts of the metal to have a higher reduction potential than others, and those places are where the metal dissolves first. As those areas dissolve, they become pits where the local concentration of metal ions is high and access is difficult for the oxidizing species, causing further dissolution in a self-accelerating process. This is why a smooth piece of metal becomes pitted and scarred as it corrodes. A homogeneous piece of metal will rapidly develop more or less stable anodic and cathodic areas where metal dissolves and other species are reduced, respectively.
The advantages of a couple or amalgam over plain metal are several. Some metals have self-passivating oxides that quickly cover active areas and choke off dissolution. This is exactly why aluminum, a strong reducing agent used in pyrotechnics and flares, with a reduction potential of 0.83 volts above that required to break down water into hydrogen, doesn't burn on contact like sodium. Both aluminum oxide and aluminum hydroxide are not water soluble, and a protective layer rapidly forms preventing further oxidation. When aluminum is amalgamated, however, the oxide layer is broken and the bare aluminum surface effectively exposed. The mercury, being much more electronegative than the aluminum, generates an internal potential that sweeps up electrons as fast as the aluminum generates them, preventing negative charge from accumulating at the dissolving surface and slowing down the reaction. Also, the presence of a surface layer of electronegative metal like mercury or copper seems to help reduce the undesirable hydrogen generation side reaction. (Why?)
To state it simply: Coupling makes dissolving metal reductions work better by:
1.   Breaking the metal’s surface passivation
2.   Removing charge from the dissolution area, speeding up the reaction

So what would keep an aluminum-copper couple from working?

   Based on the information above, it would appear that the Al(Cu) couple is very feasible. Two objections that might arise questioning the possibility of an Al(Cu) couple are:
(1) "The difference in reduction potential between mercury and copper might lower the voltage of the Al(Cu) cell compared to the Al(Hg) cell, affecting its ability to reduce organic materials."
 This is a nonissue, as the voltage of the cell is not developed between the two metals in the couple, but rather between the metal composite and the reducible species in solution. In other words, the voltage of an aluminum-copper redox pair is not between Al and Cu, but rather between Al and Cu++ ion in solution. When all the Cu++ has been reduced to Cu and plated out onto the aluminum, the redox reaction is complete--the cell is discharged. There are no Cu++ or Hg++ ions in solution when a dissolving metal reduction is being performed with a couple, so this "cell voltage" in fact does not exist. There is no Al-Hg or Al-Cu redox pair present, only an Al-Organic redox pair.
(2) "Mercury is liquid, so while the aluminum is dissolving, the mercury can follow it, while copper is a solid and will flake off before the aluminum is done dissolving"
   This may or may not happen. It doesn't seem to occur in the zinc-copper couple, and zinc is also used as amalgam, so it seems a safe bet that it won't cause too much trouble with aluminum.
Ultimately, experiment is king. There are still many unknowns involved in guessing the efficiency of the Al(Cu) couple. There may be differences in the amount of hydrogen gas evolution permitted by Hg and Cu, which would affect efficiency, there may be different surface effects with hydrogen adsorption, the different work functions of the materials may affect the rate of dissolution in some significant way, as previously mentioned, there may be mechanical coupling difficulties. Some basic exploration is the only sure way of finding facts to fill in these unknowns.

[END]

ning

  • Guest
More insight
« Reply #1 on: November 30, 2003, 10:21:00 PM »
Not many comments, eh...

Here is an interesting bit of insight into the mechanism of amalgam/couple dissolving metal reduction.

The reaction at the electrode of a voltaic cell

(http://www.chemistryquestion.com/English/Questions/HighSchoolChemistry/3_VoltaCell.html)

seems the hydrogen overvoltage may be of some importance. Any bees out there with practical experience doing various types of amalgam and couple reductions who'd care to chip in some practical experience?

Also on the same page:

Oxidation of formic acid

(http://www.chemistryquestion.com/English/Questions/HighSchoolChemistry/5_oxidation_formicacid.html)

those japanese high schools must be something scary....

Osmium

  • Guest
> If, however, there is a reducible species
« Reply #2 on: December 01, 2003, 01:34:00 AM »
> If, however, there is a reducible species in protic solution, then there is no
> reason for the hydrogen produced to be wasted as a gas. The monatomic hydrogen
> adsorbed on the metal surface is a highly reactive free radical, and will
> readily attach to certain functional groups.

The reason Al/Hg works is the overvoltage. H2 production on a Hg surface is a process that's not readily taking place unless you supply additional energy on top of the theoretically predicted. That's why this reaction rather reduces whatever is around than producing H2 gas.
Have a look at alkali electrolysis, same thing there. It only works with Hg, in this case the overvoltage is so big that Na will be reduced to the metal in an aqueous environment.


ning

  • Guest
Hmmm.
« Reply #3 on: December 01, 2003, 02:46:00 PM »
I thought that alkali electrolysis was done into mercury also because only the metal at the amalgam surface could react, and most of it went into the mercury so it couldn't react...

Osmium, do you or anybody else know about the relative reducing power of the zinc amalgam compared to the zinc-copper couple?

If I remember correctly, zinc's potential is below that of water decomposition, while aluminum's is above, so perhaps the overvoltage issue is of little concern for zinc, but critical for aluminum.

Hm.


Osmium

  • Guest
> I thought that alkali electrolysis was...
« Reply #5 on: December 02, 2003, 03:57:00 AM »
> I thought that alkali electrolysis was done into mercury also because only the
> metal at the amalgam surface could react, and most of it went into the mercury
> so it couldn't react...

No, the main reason Hg is used is that only with a Hg cathode Na metal can be produced. Most (if not all) other cathode materials would rather decompose the water into H2/O2 or H2/Cl2 in the case of NaCl solutions. H2 generation is energetically much favored over sodium reduction, the reason it isn't happening here is overvoltage. H2 does not like to be produced on a Hg surface. It requires more potential (voltage) to do so, that's why this phenomenon is called overvoltage.
Other reasons that Hg is such a suitable material include the 'solubility' of Na in Hg, and that the resulting alloy is still liquid and can be pumped as long as the Na content is low.

> If I remember correctly, zinc's potential is below that of water
> decomposition, while aluminum's is above,

No. In both cases the potential is below that of water. Both are soluble in dilute non-oxidising acids. The reason they do not spontaneously corrode is passivation, if it wasn't for the spontaneously forming oxide layer both metals would dissolve very quickly in water.

> so perhaps the overvoltage issue is of little concern for zinc, but critical
> for aluminum.

This doesn't make any sense. The Al or zinc do not care if they are alloyed with Hg, Cu, Ag or another noble metal, they have the desire to dissolve (a.k.a. negative potential). Both will readily supply their electrons to the more noble metal.


hypo

  • Guest
why?
« Reply #6 on: December 02, 2003, 05:44:00 AM »
stupid question: why is H2 generation hindered on Hg whereas
other reductions (eg imine->amine) are not?


Osmium

  • Guest
> why is H2 generation hindered on Hg ...
« Reply #7 on: December 02, 2003, 07:28:00 AM »
> why is H2 generation hindered on Hg

Dunno. Don't ask. It just is, period.

> whereas other reductions (eg imine->amine) are not?

This reaction works by electron transfer from the Hg surface to the molecule to be reduced. No H2 is generated in this case. The intermediate organic cation grabs a proton from the solvent (ROH or H2O) to finish the reduction.


Rhodium

  • Guest
overvoltage
« Reply #8 on: December 02, 2003, 08:10:00 AM »
The overvoltage thing occurs due to the hydrogen atoms formed on the metal surface being bound more strongly to Hg/Pb (both showing this phenomenon) than to other metals, so they prefer to stay there as opposed to reacting with protons in solution, evolving hydrogen gas. This could in turn be explained by HSAB theory, but that is too complex for me to describe here.

http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA3/MAIN/VOLTAGE/PAGE1.HTM


Organikum

  • Guest
there is this patent
« Reply #9 on: December 02, 2003, 08:37:00 AM »
posted somewhere here already, claiming that Al/Hg can be replaced by almost every Al-Metal couple whereby the metal has to have an lower overvoltage than Al has.
So Fe works fine is told or most Al-foil should do the job as it is alloyed with about 3% Fe mostly (Reynolds for example).

But thats from memory, a search might help.

moo

  • Guest
Generalization?
« Reply #10 on: December 02, 2003, 11:34:00 AM »
posted somewhere here already, claiming that Al/Hg can be replaced by almost every Al-Metal couple whereby the metal has to have an lower overvoltage than Al has.

I'd be very interested to see exactly in which reductions can this replacement be made succesfully. In electrochemistry, especially before proper controlled potential electrolysis, it was commonplace to select the electrode material so that the hydrogen overvoltage was not much higher than the reduction potential of the compound to be reduced, so that there was a narrow margin in potential between the desired reduction and hydrogen evolution. This resulted in selectivity: when the cell voltage was kept in such a limits that hydrogen evolution was negligible, the experimenters could be sure that there was no side-reactions due to higher potential reductions occurring. And even if something changed so that the potential at the working electrode got higher, hydrogen evolution would eat more current than the side-reaction.

My point being that you must have the metal selected according to your substrate and can not expect to get similar results with different metals. Dissolving metal reductions have parallel aspects to electrochemical reductions.

Then, on the other hand some metals have catalytic properties. Nickel, for example, being a hydrogenation catalyst could be expected to behave a more or less different to some other metals when used in metal couples.

By the way, some time ago I read about the experiments on nitrostyrene reduction with Al/Hg, using which Sunlight came to a 35% yield of 2C-H. Compare his method (

Post 108626 (missing)

(dormouse: "2C-H with Al/Hg , 35 % yields  -sunlight", Novel Discourse)
) with the method the german scientists used (

Post 459028 (missing)

(Rhodium: "Nitrostyrenes to phenethylamines using Al/Hg", Novel Discourse)
) and note that Sunlight used a temperature of 60-80 degrees due to problems with solubility, whereas the original procedure used a temperature of 20 degrees and a different substrate, with a better solubility maybe. No wonder the yield was unexpectedly low, eh? The idea of adding THF cosolvent to the mixture is, in my opinion, a good idea.


Osmium

  • Guest
> My point being that you must have the...
« Reply #11 on: December 03, 2003, 05:11:00 AM »
> My point being that you must have the metal selected according to your substrate
> and can not expect to get similar results with different metals. Dissolving
> metal reductions have parallel aspects to electrochemical reductions.

Correct. This is like an electrolysis taking place and the Hg serves as the cathode, while the choice for the dissolving metal determines the potential difference (a.k.a. voltage). If you take a metal which doesn't have enough negative potential then the electricity produced might not be enough to overcome the overvoltage and make the desired reaction happen.


ning

  • Guest
So I guess the question is...
« Reply #12 on: December 05, 2003, 04:48:00 PM »
So I guess the question is, will a couple of Al-Cu blow off most of its "power" as hydrogen, or will it reduce imines?

I don't know what potential is required to reduce an imine, but if zinc can do it, aluminum certainly should. The only question is: is the potential of aluminum TOO high, thus blowing its load as hydrogen instead of inseminating our ketones with N-methylated goodness?

Does anyone have a ketone to spare?


lugh

  • Guest
Reference Articles: Overvoltage + Al couples
« Reply #13 on: December 05, 2003, 05:10:00 PM »
While in a science library recently, SWIL did some research on the original topic of this thread, the proposed aluminum-copper couple. Sorry to say, there's nothing in the chemical literature, which probably means it won't work  ;D  Aluminun-tin and Aluminum-nickel are mentioned  ;)  Here's the articles from Paquette's Encyclopedia of Organic Reagents on Aluminum and Aluminum Amalgum, by Emmanuil I Troyansky, of the Institute of Organic Chemistry, Russian Acadamy of Sciences, Moscow, Russia:



As far as the concept of hydrogen overvoltage, in general the lower the melting point of the metal, the greater the overvoltage, thus mercury has the highest; this article from Thorpe's goes into far more detail:



8)