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Electrochemical Cells

Introduction

Consider the spontaneous reaction of Zn metal in a solution of Cu2+:

Zn(s) + Cu2+(aq) ---> Zn2+(aq) + Cu(s)     DeltaGo = -212.6 kJ/mol

If we just dump the reactants (Zn metal and Cu2+) together, they react to produce Zn2+, Cu metal, and heat. We can capture this energy by setting up an electrochemical cell to separate the two half reactions.

electrochemical cell
Zn(s) ---> Zn2+(aq) + 2e-   Cu2+(aq) + 2e- ---> Cu(s)

In this set-up Zn is being oxidized in one cell and Cu2+ is being reduced in the other cell. The Zn electrode appears corroded due to loss of Zn into solution, and the Cu electrode is being plated with Cu from solution. The salt bridge is a tube filled with a saturated KNO3 solution. It has frits on the ends that prevent mixing of the solutions but allowing ions to pass through. As the cell operates Zn2+ is produced in one cell and Cu2+ is removed from the other cell. Ions move through the salt bridge to keep the individual cells electrically neutral.

The driving force is the same for the case of just mixing the reactants. By separating the two half reactions, the electrons must travel through the wire and we can use the electrical energy. In the sketch above the electrical energy powers a light bulb.

When we first set up this electrochemical cell and complete the circuit, we could measure a voltage through the wire. As the reaction proceeds, the system approaches equilibrium and the voltage would eventually go to zero.

Batteries are electrochemical cells that convert chemical energy to electrical energy. The reactants in batteries are at non-equilibrium concentrations. As you use them, the reactants form products to approach equilibrium and the voltage drops until the battery is no longer usable.


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