I am analyzing a reaction between water sulfate copper and aluminum. A chunk of $\ce{Al}$, e.g., aluminum foil. I realize that a layer of $\ce{Al2O3}$ forms on the surface of the $\ce{Al}$, passivizing it, and I realize that the $\ce{NaCl}$ (or rather just the $\ce{Cl^-}$) cleans the aluminum oxide off the surface of the chunk of aluminum, but I don't understand why. I'd appreciate greatly if somebody could offer me an explanation.

Furthermore, (I didn't exactly include this in the topic of my question,) why is it that the reaction looks like this:

$$\ce{6H2O + 3CuSO4 + 5Al \rightarrow 3H2 + Al2(SO4)3 + 3Al(OH)2 + 3Cu}$$

(or something along the lines of that) rather than:

$$\ce{3CuSO4 + 2Al -> Al2(SO4)3 + 3Cu}$$ ?

Something along these lines has to be the case, because the reaction yields gas..

  • 2
    $\begingroup$ You should split your first equation into two, one where Al reduces water, one where Al reduces Cu. Both reactions happen more or less independently, there is no absolute stochiometric relation between them. $\endgroup$
    – Karl
    Nov 15, 2016 at 21:44
  • $\begingroup$ Would that imply that clean Al would react with H2O to produce H2 gas equally as vigorously as when it's reacting with CuSO4 + H2O solution and forming H2 gas, if at the same temperature? Also, thanks for the suggested edit. $\endgroup$ Nov 15, 2016 at 21:49
  • $\begingroup$ No, that's what I meant by "more or less". I would not remove the first equation, just add a third that describes the evolution of H2 from water + Al. There is surely some interesting interaction between the two parts. Btw. Al-sulfate is soluble in water. $\endgroup$
    – Karl
    Nov 15, 2016 at 22:00

2 Answers 2


Aluminum resists corrosion in neutral or slightly off-neutral water because of the very insoluble Al2O3 film on the metal. If you break this film, it will corrode the bare metal and reform.

But if you scratch the Al and attach a cathode (a less active metal), you have a galvanic cell, and H2 can be evolved from this cathode as Al dissolves. The simplest example of this is when you put Al in H2O in contact with a drop of mercury. Scratch the Al thru the Hg; the Hg amalgamates with the Al (attaching itself as cathode), allowing H2 to be evolved as Al dissolves elsewhere.

The experiment with CuSO4 is similar: scratch the Al; some Cu will be deposited and act as a cathode and the rest of the Al will eventually dissolve in the H2O. If you are too lazy to scratch the Al (I'm being facetious!), you can add a little Cl- ion to the H2O, which will begin corrosion of the Al, deposition of Cu, galvanic cell production, dissolution of all the Al.

There is an electrochemical example posed as a test question: would it be better to make a ship out of aluminum with copper rivets or out of copper with aluminum rivets? An aluminum boat? Nah! It would corrode in seawater. Copper would survive longer. But NO! The copper boat would have a huge cathode and tiny aluminum anodes (the rivets), which would corrode rapidly and the copper plates would fall apart. On the other hand, an aluminum boat would not last forever in seawater, but the tiny cathodes would limit corrosion current, and the large aluminum anodes would have corrosion distributed over the whole ship, so it would survive longer than the other way. Of course, any intelligent person would use copper rivets on copper plates, and aluminum rivets on aluminum, but it's just to make a point.


Chlorides are known to be an aggressive ion for metal corrosion, this is thought to be because of the ability of chlorides to destabilise the passivation film on metals which protect the metals from corrosion. What is known as pitting corrosion can occur, where a breakdown of the passivation layer at specific points of the metal surface leads to corrosion there by other oxidising agents in the aqueous medium, leading to the formation of corrosion pits. Exposed metal surfaces due to the breakdown of the passivation layer releases metal ions, which complex with the chloride anions; the metal-chloride complex reacts with water to form hydrochloric acid and a metal-hydroxide complex, reducing the pH of the medium close to the metal surface, further accelerating corrosion. (A reference can be found here: http://sassda.co.za/stainless- steel-and-corrosion/).

Aluminium is an "amphoteric" metal which reacts with alkalis; water first oxidises Al to Al (III), producing hydrogen gas (probably the effervescence that you observe) due to the reduction of water. Al (III) ions then complex with hydroxide ions to form a soluble complex in solution. In fact, when an alloy of Cu and Al is placed in NaOH, the Al is seen to dissolve away, leaving the copper behind in a nanoporous structure, a material that has potential for catalytic applications.

As for your question as to whether Al can react as vigorously with water as copper (II) sulfate solution, I think yes. But if your Al isn't as "clean" as you think and still has a layer of passivation on the surface, sulfate ions are known to also destabilise passivation layers of metals so that might enhance the corrosion (in fact some sources claim sulfate to be a better corrosive agent than chloride).

Sorry that I don't have any references, most of this is based on literature that I have read quite a while ago. Also, I am just a beginner learner in chemistry and corrosion science, so please help to improve my answer and correct any mistakes!


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