A concentration cell is an electrochemical cell with identical electrodes in electrolyte solutions of different concentrations; since both half-cells are the same, E° = 0, and the cell potential comes entirely from the concentration difference as the system drives toward equilibrium (Q → K).
A concentration cell is a galvanic cell where both electrodes are made of the same material (say, two silver electrodes) sitting in solutions of the same electrolyte at different concentrations (like 0.10 M and 1.0 M AgNO₃). Because the two half-reactions are identical, the standard cell potential E° is exactly zero. So where does the voltage come from? Entirely from the concentration difference.
The cell runs spontaneously in whatever direction evens out the concentrations. In the dilute compartment, oxidation occurs (the metal electrode dissolves, adding ions and raising the low concentration). In the concentrated compartment, reduction occurs (ions plate onto the electrode, lowering the high concentration). Electrons flow from the dilute side (anode) to the concentrated side (cathode). This is the purest possible demonstration of the CED's big idea in 9.10.A.1: cell potential is a driving force toward equilibrium, and the farther the system is from equilibrium, the bigger the voltage. Once the concentrations equalize, Q = K = 1, E = 0, and the cell is dead.
Concentration cells live in Unit 9 (Thermodynamics and Electrochemistry) under Topic 9.10, supporting learning objective 9.10.A. That LO asks you to explain how deviations from standard conditions change cell potential, and a concentration cell is the ultimate stress test of that idea. There's no E° to lean on, so any voltage you measure is pure proof that nonstandard concentrations (Q ≠ 1) create a driving force. The CED is also explicit that you should reason with Q and equilibrium-distance arguments, not Le Châtelier's principle, because an operating cell is not at equilibrium (9.10.A.2). If you can correctly predict which electrode is the anode in a concentration cell and explain why, you've genuinely understood what cell potential means.
Keep studying AP® Chemistry Unit 9
Electrochemical Cell (Unit 9)
A concentration cell is just a galvanic cell stripped down to one variable. Same electrodes, same electrolyte, so the only thing generating voltage is the concentration gap. It's the limiting case that shows voltage doesn't require different metals, just a system that isn't at equilibrium.
Reaction Quotient Q and Equilibrium (Unit 7)
Everything about a concentration cell runs on Q. For identical half-cells, Q is the ratio of dilute to concentrated ion concentration, and the cell pushes Q toward K = 1. The farther Q is from 1, the larger the voltage, which is Unit 7's equilibrium logic wearing an electrochemistry costume.
Salt Bridge (Unit 9)
A concentration cell still needs a salt bridge or semipermeable membrane to maintain charge neutrality. As the dilute side gains positive ions from oxidation and the concentrated side loses them to reduction, ions migrate through the bridge to keep both compartments neutral so electrons keep flowing.
Gibbs Free Energy and Spontaneity (Unit 9)
Since ΔG = −nFE, a positive cell potential means the mixing process is thermodynamically spontaneous. A concentration cell converts the free energy released by evening out concentrations directly into electrical work, tying the electrochemistry of 9.10 back to the thermodynamics earlier in Unit 9.
Concentration cells show up in multiple-choice questions that hand you two identical electrodes (often Ag in AgNO₃ or Cu in CuSO₄) at different concentrations and ask you to identify the cell type, name the anode and cathode, or predict the direction of spontaneous electron flow. The classic trap is direction. Remember that electrons flow FROM the dilute compartment (where oxidation raises the low concentration) TO the concentrated one. Particulate-representation questions are common too, asking you to match a drawing of electron flow to the correct concentrations. On free-response questions, the move is justification. Be ready to explain that E° = 0 because the half-cells are identical, that the voltage exists because Q ≠ 1, and that the cell potential shrinks to zero as the concentrations approach equality. Use Q-based reasoning, not Le Châtelier, since the CED explicitly says equilibrium arguments don't apply to a cell that isn't at equilibrium.
A standard galvanic cell uses two different half-reactions (like Zn and Cu), so its voltage comes from a difference in reduction potentials and E° ≠ 0. A concentration cell uses the same half-reaction on both sides, so E° = 0 and the voltage comes only from the concentration difference. If a question gives you identical electrodes and identical electrolytes at different molarities, you're looking at a concentration cell, and the standard reduction potential table won't help you.
A concentration cell has identical electrodes and electrolytes, so its standard cell potential E° is exactly zero.
The voltage in a concentration cell comes entirely from the concentration difference, because the system is not at equilibrium (Q ≠ 1).
Oxidation happens in the dilute compartment and reduction happens in the concentrated compartment, so electrons flow from the dilute side to the concentrated side.
The cell runs in the direction that equalizes the two concentrations, and the voltage drops to zero once Q reaches K.
Explain concentration cells using Q and distance-from-equilibrium reasoning, not Le Châtelier's principle, because an operating cell is not at equilibrium.
The bigger the gap between the two concentrations, the farther Q is from 1 and the larger the cell potential.
It's an electrochemical cell where both electrodes are the same material in solutions of the same electrolyte at different concentrations, like two silver electrodes in 0.10 M and 1.0 M AgNO₃. The cell generates voltage purely from the concentration difference as it drives toward equilibrium.
Because both half-cells use the identical half-reaction, their standard reduction potentials cancel exactly. Any measured voltage comes from nonstandard conditions, meaning Q ≠ 1, not from a difference in electrode materials.
The dilute side. Oxidation there dissolves the electrode and raises the low ion concentration, while reduction on the concentrated side plates out ions and lowers the high concentration. Electrons flow from the dilute electrode to the concentrated one.
No, and the CED says so directly (9.10.A.2). An operating cell is not at equilibrium, so Le Châtelier doesn't apply. Instead, argue that the cell potential is a driving force toward equilibrium and that the farther Q is from K, the larger the voltage.
A regular galvanic cell gets its voltage from two different half-reactions with different reduction potentials, so E° ≠ 0. A concentration cell uses the same half-reaction on both sides (E° = 0) and gets all its voltage from the concentration gap between compartments.
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