Concentration Cells

What if you built a cell where both electrodes were made of the same metal, dipped into solutions of the same ion at different concentrations? E° would be zero - the same half-reaction on both sides cancels perfectly. Yet the cell still produces a voltage. How?

The Concept

A concentration cell uses identical electrodes and identical ions but at different concentrations in the two half-cells. The driving force comes entirely from the concentration difference, not from a difference in reduction potentials.

The system is trying to equalize concentrations - just like how gases diffuse from high to low concentration. Nature drives toward equilibrium, and the electron flow is the mechanism.

How It Works

Consider a copper concentration cell:

• Anode (dilute side): Cu(s) | Cu2+(0.010 M)
• Cathode (concentrated side): Cu2+(1.0 M) | Cu(s)

At the anode (dilute side): Copper metal oxidizes to add Cu2+ ions to the dilute solution, increasing its concentration.

Cu(s) -> Cu2+(aq) + 2e-

At the cathode (concentrated side): Cu2+ ions are reduced to solid copper, decreasing the concentration.

Cu2+(aq) + 2e- -> Cu(s)

Net effect: The dilute solution gets more concentrated, and the concentrated solution gets more dilute. The system drives toward equal concentrations on both sides.

Calculating the Voltage

Since E° = 0 (identical half-reactions), the Nernst equation gives:

E = 0 - (0.0592/n)log(Q)

For the copper concentration cell with [Cu2+]dilute = 0.010 M and [Cu2+]concentrated = 1.0 M:

Q = [Cu2+]anode / [Cu2+]cathode = 0.010 / 1.0 = 0.010

E = -($\frac{0.0592}{2}$)log(0.010) = -(0.0296)(-2) = +0.0592 V

The voltage is small but positive, confirming the reaction is spontaneous. As the concentrations equalize, Q approaches 1, log(Q) approaches 0, and E approaches 0.

Key Features of Concentration Cells