Question

Mercuric oxide dry-cell batteries are often used where a flat discharge voltage and long life are required, such as in watches and cameras. The two half-cell reactions that occur in the battery are

HgO(s) + H₂O(l) + 2 e-→ Hg(l) + 2 OH^-(aq)

Zn(s) + 2 OH^-(aq)→ ZnO(s) + H₂O(l) + 2 e-

a) Write the overall cell reaction. (b) The value of E°red for the cathode reaction is +0.098 V. The overall cell potential is +1.35 V. Assuming that both half-cells operate under standard conditions, what is the standard reduction potential for the anode reaction? (c) Why is the potential of the anode reaction different than would be expected if the reaction occurred in an acidic medium?

One Class Solution:-

a) Hg(s) +Zn (s) → Hg(l) + ZnO(s)

b) E⁰=E⁰(cathode)-E⁰(anode)=0.098-1.35= -1.25V

c) E_r red. is different from Zn ²⁺ (aq) + 2e-→ Zn(s), (-0.76 V) because in the battery the process happens in the presence of base and Zn²⁺ is stabilized as ZnO(s).

Stabilization of a reactant in a half-reaction decreases the driving force, so Er is more negative.