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Author: Subject: Removal of surface oxide by copper(II) chloride
darkflame89
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shocked.gif posted on 5-9-2005 at 00:27
Removal of surface oxide by copper(II) chloride


Recently, while browsing through Woelen's website, I noticed a few curious reactions featured on his site regarding aluminium. We all know aluminium has a layer of surface oxide that is stable, non-porous and adheres to the metal. The aluminium oxide layer passivates it, making it unreactive to other chemical attacks. In fact, if you want to cause it to react with other substances say water, you need to remove the oxide layer by dipping it in say, sodium hydoxide soulution.

Yet, there seems to be another ion that can react and remove this oxide layer. That's the tetrachlorocopper(II) ion, [CuCl4]2- . For example, if you dip aluminium into HCl, the aluminium reacts sluggishly. But if you add a bit of CuSO4 to the solution, the reaction speeds up considerably, to the point of vigorous evolution of steam and much heat. The Cu2+ ions and the Cl- ions seem to work together to remove the oxide layer, then then Cu2+ ions catalyze the reaction between aluminium and H+ ions.

Now, I did a trial to verify this. On an aluminium pie pan, I added a drop copper(II) sulphate to the surface. Nothing happened. This is to be expected due to the oxide layer. On a separate surface, I added a pile of NaCl to the surface. Then, 1 drop of copper(II) sulphate solution was added on top of the NaCl pile. The CuSO4 seeped in, and the whole pile turned green due to the formation of the tetrachlorocopper(II) complex. After a while, the surface beneath the salt pile turned dark, and the mixture fizzed. Smoke(steam?) came off. I continuously added more drops of CuSO4 to replace the lost water and more fizzing, until abruptly, the entire piece of aluminium beneath the salt pile came off, eroded away. At the end of reaction, the products was washed off. Copper metal could be seen deposited. Some CuO appears to have formed as well( not too sure about this)

Well, of course, a few questions to be dealt with. The fizzing sound was caused by evolution of hydrogen. Since the oxide layer was removed, the aluminium beneath was rather reactive, and reacted with water present to form hydrogen gas. This process was probably aided by the aluminium and copper couple.

This example was backed by some on the Internet. Running a search via google, there appears to be some guy who found that this solution (CuSO4 + NaCl) was better way to etch aluminium than conventional means. Copper(II) chloride works just as well, its just that other copper salts were more easily obtainable and NaCl plain cheap. Another woman wrote to SAS too describing an experience similar to my adventure here. She said something about trying to grow a crystal garden on aluminium plate. To her surprise, when she tried copper(II) sulphate, she inadvertantly added sodium chloride and there was much fizzing. Her aluminium pan was similarly corroded.

It turns out there's a related thread here in Sciencemadness. Its in the Energetics thread, under exotic thermites. Polverone said something about igniting aluminium and copper(II) chloride. It seems like the mixture is easier to ignite.

Okay, to cut the long story short, does anyone know if the reaction mechanism on how copper(II) ions and chloride ions remove the aluminium oxide? Is it something to do with chloroaluminium complex? Incidentally, the above experiment that I tried, there was no aluminium hydroxide in the products.




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[*] posted on 5-9-2005 at 16:05


What's the pH of a CuSO4 solution? Maybe part of the solution lies there.
Also, I suppose oxy-chloro Al derivatives are possible, which then can react onwards... but that's sheer speculation.

I suppose you could try and investigate this, i.e. by using different transition metal salts, and different ions, and different ionisation states (i.e. CuI+)

[Edited on 6-9-2005 by chemoleo]




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[*] posted on 5-9-2005 at 22:46


I also tried this with cobalt chloride and aluminium. Cobalt also forms a tetrachloro ion at high chloride concentration (thjis is a deep blue ion), but this is not effective for oxidizing aluminium. Apparently the copper / chloride combination has something unique which makes the reaction with aluminium so vigorous.

A low pH favors the fast corrosion of aluminium, but that certainly is not necessary. A plain solution of table salt, mixed with CuSO4 gives a vigorous reaction with aluminium, especially if the concentration is somewhat higher. Of course, a solution of CuSO4 always is a little acidic, due to hydrolysis of the copper ions, but this effect only is very weak, especially when a lot of chloride is added (in that case almost all Cu(2+) ions are used in formation of CuCl4(2-) ions).




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[*] posted on 5-9-2005 at 23:37


Okay, didn't think about the pH part, will try it today. About trying it with other transition metals, you mean to see whether other metal+chloride combos work, or whether the copper+ chloride combo work for other metal oxides?

I tried the experiment again with iron, just iron, no rust whatsoever. Same thing copper sulphate and NaCl, left the mixture out for 10 mins, but nothing happened.




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[*] posted on 6-9-2005 at 00:23


There are three parameters, which can be varied:

1) The metal to be corroded.
2) The counter ion for the chloride.
3) The counter halide for the copper.

What I have done is (2). I changed the copper chloride to cobalt chloride but I still used aluminium metal. That had no effect.

I also did (3). Instead of a copper / chloride combination, I used a concentrated copper / bromide combination. This also works and makes the aluminium corrode much faster. I'll try this evening with copper (II) and fluoride. However, I do not expect this to work, but I'll let you know.

I did not do (1). As I understand from your post, you did (1) by replacing aluminium with iron metal.




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[*] posted on 6-9-2005 at 08:31


I don't think 1 can be varied much because we're dealing with a specific situation where the metal is covered with an only somewhat inert metal oxide layer. Magnesium wouldn't substitute well because it and its oxide are more reactive. The refractory metals (titanium, zirconium, et al) have stronger, less reactive oxides. Iron has very little surface oxide and it's more reactive in low pH, I would say.

Would nickel ions work? I have an impure solution of NiCl2 that turns blue (like CuCl2, which is a likely impurity..) on dilution.

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[*] posted on 7-9-2005 at 10:34


I did the test with sodium fluoride as well and that indeed did not result in fast corrosion of aluminium.

So, summarizing:

a) Cu(2+) and F(-) do not corrode Al-metal at high speed.
b) Cu(2+) and Cl(-) are most effective.
c) Cu(2+) and Br(-) also corrode Al-metal, but not as fast as the combination Cu(2+)/Cl(-).
d) Combination Cu(2+)/I(-) is not possible. The Cu(2+) oxidizes the iodide, forming solid CuI and I2.

I also tried combinations of Cu(2+) and the pseudohalogens CN(-) and also SCN(-).
The copper / cyanide combination gives a pale blue/green precipitate, probably a mix of copper (I) cyanide and copper (II) hydroxide.
Thiocuanate gives a mix of white copper (I) thiocyanate and some black stuff (I think that is copper (II) thiocyanate).




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[*] posted on 7-9-2005 at 19:10


This get more and more interesting. If there were a chloroaluminium complex, then there would be no need for copper(II) ions to be present. Perhaps there's something to do with the covalent properties between aluminium and chlorine, bromine and iodine? Fluorine does not count here as it is strongly ionic.

Perhaps, a mixture of iodine and iodide ions plus CuSO4 should do the trick? I tried such an experiment with iodine and aluminium. The iodine was dissolved in alcohol as in tincture. The CuSO4, iodine, iodide and aluminium took a night to react fully. When the reaction was over, some excess aluminium was covered with copper and a few bubbles of hydrogen evolved from the aluminum. The reaction was repeated, but without adding CuSO4. The reaction took a week(!) to complete.

Next step, to verify the chloroaluminium complex, we should try shaking up aluminium hydroxide with an extremely concentrated solution of NaCl. Not HCl, because of the acid-base reaction. Example, NaOH can be added to aluminium sulphate and the gelatinous ppt appears. Can this ppt be filtered off?I'm not sure about this one. If yes, the filter and test the residue. If no, then add HCl to the test tube until the pH of the solution is neutral. Then, lots of NaCl can be added and see whether this ppt dissolves.

The experiment can be repeated by using Al2O3 instead. Same thing, added conc. NaCl solution to the Al2O3 and see if it dissolved.




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[*] posted on 8-9-2005 at 08:19


I don't think it does. I precipitated Al(OH)3 from NaOH and HCl solutions, using the opposite to neutralize it.

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[*] posted on 14-9-2005 at 22:20


Hard anodised aluminium seems to be unaffected.
I had some used pcb etchant (ammonium persulfate + copper sulfate) that I'd finished with, so I added some table salt and threw in some aluminium scrap I had laying about. The unanodised aluminium reacted fast, getting covered in spongey copper and bubbling furiously (what fun!). The hard-anodised piece was unaffected, except on the cut edges.




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