Most of you have probably carried out this reaction (and I do it whenever I have some leftover solution of Cu2+)
3Cu2+ + 2Al --> 3Cu + 2Al3+
but I wanted to ask a few questions about some things I noticed:
If there is only CuSO4 and Al, the reaction is very slow
However, if the solution is saturated with NaCl, the solution will turn green (Cu2+ goes to CuCl42-) and the
reaction is instantaneous, with generation of heat.
Does this happen because the newly formed Al3+ ion gets complexed by Cl- ions to give out AlCl4- instead
of reacting with water and forming the hydrated oxydes (which would, perhaps, slow down the reaction) or is there any other explanation? Or maybe the
presence of a strong electrolite like NaCl is enough to accelerate the electron exchange?
Thanks for your attention and let me know if I wrote or said anything wrong; I strive to improve my knowledge in chemistry as much as I can and I will
gladly welcome suggested literature articles.Brain&Force - 11-1-2014 at 11:51
I'm guessing this occurs because Cu2+ is a weak acid in itself, and a tiny amount of HCl is formed that destroys the oxide layer that
normally protects the metal. After it is destroyed the main reaction can proceed. AJKOER - 11-1-2014 at 16:08
In my opinion, the main reaction is electrochemical in nature. The aqueous NaCl serves as the electrolyte in one role, and in the creation of a
complex in another. With an excess of Aluminum, the reaction with CuSO4 provides Copper for the anode (actually a zone, and this electrode may not be
inert in this cell). I am not clear if any dissolved O2 plays a role, but if employing a boiled CuSO4 solution ( to remove gases) behaves differently,
it probably does. [EDIT] If adding a little H2O2 accelerates the reaction, another clue. Note, there is at least one study (see "Corrosion Rate
Monitoring of Cu Alloy Tubing in Flowing NaCl Solution.." at http://www.maneyonline.com/doi/abs/10.1179/00070598279827467... ) that looks at 'aerated NaCl solution' with Copper/Aluminum alloys.
To complicate things, there are also side reactions (like the formation of hydrogen). Chemical reactions, however, will slow down considerably as the
aqueous CuSO4 is diluted/cooled, but the electrochemical reaction will nevertheless proceed.
If this is a school project, check my references supplied in a prior thread on the so-called 'bleach battery' to get a taste on the possible
electrochemical reactions. The central piece, in my opinion, is to identified the REDOX reaction. And, to add complications, in a Galvanic cell there
can be more than one half reaction at each electrode. An example of half cell reactions:
2e- + Cu2+(aq) → Cu(s) (reduction)
Al(s) → Al3+(aq) + 3e- (oxidation)
Now, my comment is not totally different from Woelen's as he does states, to quote:
"Probably the aluminium/copper reaction is also fast, due to formation of a copper-aluminium electrolytic couple, which causes the aluminium to be
oxidized quickly, while hydrogen from the water is produced at the copper side of the connected metal-pair."
[Edited on 12-1-2014 by AJKOER]chornedsnorkack - 12-1-2014 at 03:35
What other ions catalyze or do not catalyze Al/Cu reaction?
Is the reaction with Cu(NO3)2 solution exactly as slow as with CuSO4 solution of equal Cu concentration?
If Cl- is saturated, the reaction gets violent. What could be the region of Cl- concentration where the reaction gets appreciably faster than in pure
CuSO4 solution, but not yet violent?
Does NaBr also catalyze Al dissolution in Cu salts?Metallus - 12-1-2014 at 06:38
What other ions catalyze or do not catalyze Al/Cu reaction?
Is the reaction with Cu(NO3)2 solution exactly as slow as with CuSO4 solution of equal Cu concentration?
If Cl- is saturated, the reaction gets violent. What could be the region of Cl- concentration where the reaction gets appreciably faster than in pure
CuSO4 solution, but not yet violent?
Does NaBr also catalyze Al dissolution in Cu salts?
I have both the sulphate and the nitrate and, without chloride ion, the reaction is always slow.
Anyways, I think that the first link gives out a reasonable explanation to the phenomenon. DraconicAcid - 12-1-2014 at 14:09
And yet, aluminum tends not to react with hydrochloric acid unless there's a catalytic amount of copper present....Metallus - 12-1-2014 at 23:49
And yet, aluminum tends not to react with hydrochloric acid unless there's a catalytic amount of copper present....
I did not know this. Do you know of any article where this was studied or observed? I'm interested a lot.
Also, is it only copper that acts as a catalyst or other metals (Such as cobalt salts) will do the same, as long as their reduction potential is
higher (I guess)?
ThanksMrHomeScientist - 13-1-2014 at 07:00
Yes post a reference to that if you have one handy, please. I've never had any trouble dissolving Al in hydrochloric acid without any copper around.
Granted, my HCl is hardware store muriatic acid so it's possible there is some contamination, but I've never heard of this before. Iron is the usual
contaminant you hear about.DraconicAcid - 13-1-2014 at 08:26
I don't have a reference, but I have noticed that when I show the reaction of aluminum foil with hydrochloric acid, it just sits there making a fool
of me in front of my class. But if I add a drop of copper(II) sulphate sol'n, the reaction goes nicely. Zinc won't react without some copper
present, which is why I thought to try it.
The same aluminum foil will also not react with 0.1 mol/L silver nitrate solution until a bit of copper is added.Metallus - 13-1-2014 at 09:43
Ok, I just wanted to confirm your observations and so I made 3 solutions of 10 ml diluite HCl (5-10%):
I then added a small strip of aluminium foil (0,5cm x 1,5cm) to each solution and started measuring the time each solution took to totally dissolve
the aluminium:
1 solution: 17:07 min
2 solution: 10:02 min
3 solution: 12.37 min
The acid alone takes its time before totally dissolving the Al, but the presence of copper did speed up the reaction by a lot. I wanted to test if the
presence of another metal influenced the reaction and, since I had some leftover cobalt (II) chloride, I decided to test it out: it turns out that
cobalt acts as a catalyst too, even if it is less "active" than copper.
Since the HCl used is the one bought at the store, the presence of Fe is obvious but I used the same acid for every solution, so the operating
conditions were always the same.
When I have time I'll try out if chromium or manganese act as catalysts as well.
Thanks for sharing your results. This is interesting for me DraconicAcid - 13-1-2014 at 10:31
Perhaps the iron also acts as a catalyst, because my lab-grade hydrochloric acid did not show any bubbling at all with just aluminum foil.Metallus - 13-1-2014 at 10:47
Perhaps the iron also acts as a catalyst, because my lab-grade hydrochloric acid did not show any bubbling at all with just aluminum foil.
What concentration of the acid did you use?
I remember we did this experiment during the inorganic chemistry course in lab and we used a 2N solution of HCl to dissolve a strip of aluminium foil,
and the reaction started after 2 mins, if I recall correctly.
The acid I'm using is most likely even less than 5% (It's more contamined water rather than diluite HCl) and that's why the times of reaction are so
long.
30-37% HCl destroys the aluminium on the spotDraconicAcid - 13-1-2014 at 10:57
I've used 6 M and 1 M for the demo. It might have started after a couple of minutes, but I didn't want to pause the lecture for five minutes while
waiting to see if something was going to happen.woelen - 14-1-2014 at 06:14
Hydrochloric acid does react with aluminium, also without copper (or any other metal) added. The reaction, however, needs time to get started. Even if
you add Al-foil to 30% HCl it takes some time before the reaction becomes vigorous, but once it is going, it can be very violent.
When a little copper is added, then the reaction betweel Al-metal and HCl starts at once and also immediately becomes very violent. Copper certainly
has a strong influence. It is the combination of copper ion and chloride ion which does the trick. The same is true for copper ion and bromide ion.blogfast25 - 14-1-2014 at 13:05
When a little copper is added, then the reaction betweel Al-metal and HCl starts at once and also immediately becomes very violent. Copper certainly
has a strong influence. It is the combination of copper ion and chloride ion which does the trick. The same is true for copper ion and bromide ion.
So to be clear: is it copper metal or Cu(II) that has the significant effect?Metallus - 14-1-2014 at 13:09
Cu(II) that gets complesserd as CuCl42-
[Edited on 14-1-2014 by Metallus]AJKOER - 15-1-2014 at 11:36
Per this source, "Copper deposition during the corrosion of aluminum alloy 2024 in sodium chloride solutions", Journal of Materials Science, July 2002
(link: http://www.deepdyve.com/lp/springer-journals/copper-depositi... ) first page only, it appears to me that the process is basically a galvanic cell
(as I noted previously) and, yes, Fe has a role (particles containing Iron are said to act as cathodic sites fostering the reduction of the cupric ion
dissolved in the NaCl solution).
A quote:
"The plating or cementation of copper from solution as an electrochemical displacement reaction appears to be a major contributor to the pitting
corrosion of 2024 aluminum alloy."
"These observations confirm a mechanism contributing to pitting corrosion in copper-rich aluminum alloys involving the plating or cementation of Cu2+
from solution as an electrochemical displacement reaction, resulting in nucleation and growth of Cu clusters on the aluminum alloy surface, and
causing additional aluminum dissolution and pitting around the Cu deposits."
The acid alone takes its time before totally dissolving the Al, but the presence of copper did speed up the reaction by a lot. I wanted to test if the
presence of another metal influenced the reaction and, since I had some leftover cobalt (II) chloride, I decided to test it out: it turns out that
cobalt acts as a catalyst too, even if it is less "active" than copper.
I liked your experiment but unless you carefully controlled the amounts of Cu and Co you can't really conclude that.
