[Twigg]

I've noticed that the oxidation of Cu metal either as a reaction intermediate or as a reactant does not produce consistent compounds. I rely on qualitative observations to distinguish the products, but the difference between CuO and Cu2O is quite obvious--colored black and red respectively. Nevertheless, I suspect that in some cases I think the product is a mixture of both, as it turns up brown. Here are some cases I've personally observed:

Aluminum wire in dilute copper sulfate solution produces red Cu2O in clumps on the wire itself
Zinc wire or powder in ~0.8M copper sulfate or copper chloride solution produces black CuO in clumps on the wire or mixed with the powder
A steel nail in ~1M copper sulfate produces both a red-brown powder (some mixture of a-Fe2O3, y-Fe2O3, Cu2O, and/or CuO) on the nail itself; the nail had a reddish almost copper colored coating.

And I've read that the direct oxidation of copper metal produces a mixture of both oxides.

What is the cause for the different products?

[AWK]

All metals in your post form metallic copper (not oxides) with copper sulfate.

[Twigg]

In aqueous solution the copper is instantaneously oxidized, so you'll only yield an oxide and never even see metallic copper form. Theoretically you're right though.

[ptryon]

A relatively simple explanation is that Copper has two common oxidation states: +2 and +1: Copper(II) ions will form CuO where as Copper(I) ions will form Cu₂O. However, this does not explain why in some systems copper(I)oxide is formed and in other systems copper(II)oxide is formed...

In the solutions that you mention the copper ions will be in the +2 oxidation state. The main reaction at the metal surface is the reduction of copper; Each copper ion gains 2 electrons from the metal to become a copper atom. For most pre-16 chemistry courses this is generally enough detail. However, to explain the preference for different oxides of copper we need to add a level of detail:

Oxidation occurs rapidly as the newly formed copper atoms immediately react with dissolved oxygen molecules:

4Cu + O₂  2Cu₂O

And further oxidation may occurs in some cases...

2Cu₂O + O₂  4CuO

Why does the first step of oxidation only occur in the presence of some metals? I don't know- but I will take a guess. Remember that at the same time the other metal (zinc, iron or aluminium in your examples) is being oxidized. Perhaps the more reactive metals like aluminium (which are higher up the electrochemical series) compete with the oxidation reaction for copper. This would match your observation that copper is only partially oxidized to copper(i) ions in the presence of Aluminium. I suspect that the actual explanation might be much more complex.

[Twigg]

Ptryon, your explanation clears things up a lot! Thanks. It occurred to me though that it might just've been that aluminum formed a good passivation layer and zinc didn't. I'll have to try it with stainless steel or another material with a quick passivation layer.

[ptryon]

No problem, good luck with your investigations