I asked a question previously on here but still have some confusion. To briefly summarize, I made a voltaic cell for a school project with 0.1 M copper nitrate and varying concentrations of aluminum, zinc and nickel nitrates. My results were fairly consistent with the theoretical values from the Nernst equation for zinc, however, the aluminum and nickel cells produced substantially less voltage than their theoretical values. The answer I got from my previous question explained how the low voltage in the aluminum cell was a result of its oxide layer.

However, I still had some confusion over why there were differences between aluminum and nickel, and zinc. Firstly, it's my understanding that a metal's ability to form an oxide layer is related to its reactivity. This would explain aluminum's tendency to form such a strong oxide layer however, zinc has a higher reactivity than nickel so wouldn't it also form an oxide layer and its outputted voltage would be lowered similarly.

My next explanation for the differences between aluminum and nickel, and zinc was related to the ability of the oxides to either react or dissolve in the nitrate solutions, however, I found that nickel oxide is soluble in acids and aluminum oxide reacts with acids.

I was also wondering how I even got a voltage in the aluminum cell if the oxide layer does not allow any aluminum ions to be transferred.

I would appreciate some guidance from the chemistry wizards because at this point I am stumped.


2 Answers 2


There are several questions in your text.

First, you imagine that "a metal's ability to form an oxide layer is related to its reactivity". No. It is not. Aluminum is rather reactive with water. But sodium is more. And metallic sodium is not protected by any oxide layer. On the contrary. Sodium reacts dangerously with water. Aluminum is a strange metal, as it forms an oxide that has the same density as the metal. So this oxide can perfectly cover the surface of the metal.

Then you ask about the difference of reactivity between zinc and nickel. I cannot answer this question, as I don't know enough of the nickel chemistry.

Then you say that aluminum oxide reacts with acids. It is right, but this reaction is extremely slow. If some aluminum piece or foil is put into a solution of $\ce{HCl}$ made by mixing the concentrated solution ($36$%) by an equal amount of water, you will have to wait a couple of minutes before the oxide layer is dissolved according to : $$\ce{Al2O3 + 6 HCl -> 2 AlCl3 + 3 H2}$$ And this last reaction is so slow that the expected reaction producing hydrogen gas happens only two or three minutes later according to $$\ce{2 Al + 6 HCl -> AlCl3 + 3 H2}$$ In this case, the production of gas is violent and highly exothermic, but it only happens with delay.

At the end, you ask why "I even got a voltage in the aluminum cell if the oxide layer does not allow any aluminum ions to be transferred". This is a good question, for which I don't have a real answer. It may be due to some default of the alumina layer.

  • $\begingroup$ Thank you for your answer! In your previous answer on my last question you stated that "Aluminum ions cannot cross [the oxide] barrier" even after any attempts to remove it like the sanding process I used. The other response to this question by James Gaidis states that this layer was "fairly impermeable" but not perfectly. I was hoping for some clarification on the nature of aluminum oxide. Thanks! $\endgroup$
    – John
    Jan 20, 2021 at 17:15
  • $\begingroup$ On pristine aluminum, islands of oxide grow, connect, and build up by oxygen penetration thru the film already formed. pubs.acs.org/doi/10.1021/acsami.7b17224 Pitting is corrosion by destruction of the already-formed film and loss of aluminum. "By nature, aluminum is a reactive metal but it is also a passive metal." fractory.com/aluminium-corrosion So, the film is intact (until it isn't) over a wide range of conditions, but if it corrodes, either something gets in or aluminum gets out - not easy either way. $\endgroup$ Jan 21, 2021 at 14:56

Aluminum and nickel both develop passive oxide layers, different in style. The aluminum oxide layer is thick, quite adherent, and fairly impermeable (tho not perfectly) and continues to thicken in the atmosphere, but slower and slower. A thick panel of aluminum will become quite rough and non-shiny after exposure to air and water, especially after exposure to seawater (chloride!). The entire oxide layer does not have to dissolve completely before the corrosion potential of the aluminum exerts itself in your experiment. My own results with $HCl$ and aluminum foil show rapid reaction within seconds, not minutes. (In the nuclear power industry, large amounts of $Al(NO_3)_3$ are needed, but cannot be obtained simply by dissolving aluminum in $HNO_3$; the aluminum is dissolved first in $HCl$, then $HNO_3$ is added and the $HCl$ is boiled off. The pesky $Al_2O_3$ layer is oxidized by the $HNO_3$ and strengthened, whereas the $HCl$ bites right thru it.) An interesting experiment would be to see if your voltaic cells have a large current capacity, or whether the potential is developed thru a tiny imperfection in the oxide layer, and can provide only a tiny current, as if the aluminum electrode were very small.

Nickel develops a very thin adherent oxide layer, similar to the oxide layer on stainless steel, and is used in many corrosive environments for this reason. The cell potentials may well be affected by the anions involved, for nitrate could strengthen the oxide film, resisting the corrosion needed to provide electrons, while doing the same series with the chlorides, especially where voltages (at very low current) are measured could align more perfectly with your expectations. The tables of electric potentials are not developed from actual experiments, but from a totality of data on reaction heats and energies where reactions are assumed, not stifled by passivity.

Experiments with copper and zinc work nicely, perhaps because the layer of oxide on the zinc is even less perfect than on aluminum (but still fairly resistant to water corrosion in the mid-pH range). The chemistry in books seems so straightforward and almost perfectly logical (with a few explainable quirks, of course), but in practice, labwork is can be complicated and frustrating, and you can learn as much about chemistry by analyzing your "deficient" results as from getting a perfect match with theory.


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