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The "role" of chemical equilibrium in electrochemical cells.

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I do not understand why a reaction (oxidation and reduction) can continue producing electricity if equilibrium (at a given temperature) has already been reached. For example, the conversion of Fe2+(aq) ion to Fe3+(aq) ion is necessary for the anode half reaction to continue producing electrons while the continued conversion of CO3+ to CO2+ is necessary to accept the electrons. All species involved in the reaction are AQUEOUS, so why does the build-up of product NOT cause the reaction to stop?

My chemistry book lists several half reactions that can be used to generate electricity. I understand that under aqueous conditions, the product or reactant may be solid/gas etc and therefore may not figure into the equilibrium constant. However, several of the half reactions listed contain ONLY aqueous species.

I will attempt to describe a scenario to which my question applies: Anode reaction: Fe2+(aq) ===== Fe3+(aq) + e-. Cathode reaction: CO3+(aq) + e- ===== CO2+(aq). ALL SPECIES ARE AQUEUOS!!!

Before the wires are connected, both reactions should have already reached equilibrium, even if a salt bridge is in place, right? Why then, do both half reactions continue producing products? Does the equilibrium constant suddenly change when electricity is being produced?

Also, what happens chemically in the above scenario when the cells finally "die" (stop producing an electron flow)?

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Dear student,
You are right that reactions are in equilibrium. But if you look at the products in Anode half reaction there ...

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