In the zinc-copper cell, Zn(s) | Zn^+2(1M) || Cu^+2(1M) | Cu(s), which electrode is negative?
Cu^+2
Cu(s)
Zn(s)***
Zn^+2
Zn⁰(s) + Cu⁺²(aq) => Zn⁺²(aq) + Cu⁰(aq)
Oxidation => Zn(s) => Zn⁺²(aq) + 2e¯ (The loss of 2e¯ means Zn⁺²(aq) has moved into solution leaving 2e¯ behind in the Zn(s) electrode which are responsible for the anodic negative charge on the anode.
Reduction => Cu⁺²(aq) + 2e¯ => Cu⁰(s) (The gain of 2e¯ by the Cu⁺²(aq) in solution leaves the copper electrode deficient in electrons and is responsible for the cathodic positive charge on the cathode.
In the zinc-copper cell, the electrode that is negative is the zinc electrode (Zn(s)). To determine this, it is important to understand the concept of electrode potentials. The electrode potential is a measure of the tendency of an electrode to undergo an oxidation or reduction reaction.
In this cell, the zinc electrode (Zn(s)) is connected to the negative terminal of the cell and is referred to as the anode. It is where the oxidation reaction takes place:
Zn(s) → Zn^+2(aq) + 2e^-
On the other hand, the copper electrode (Cu(s)) is connected to the positive terminal of the cell and is referred to as the cathode. It is where the reduction reaction takes place:
Cu^+2(aq) + 2e^- → Cu(s)
Since oxidation occurs at the anode and reduction occurs at the cathode, and the overall cell reaction is the oxidation and reduction reactions occurring simultaneously, the electrode where oxidation occurs (the anode) is considered to be negative. Therefore, in the zinc-copper cell, the zinc electrode (Zn(s)) is negative.