9: Voltaic Cells
- Page ID
- 516593
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1. Prerequisite Math & Theory
- Cell Potential: Ecell = Ecathode − Eanode. At standard states (1.0 M ion concentrations, 25°C): \[ E^\circ_\text{cell} = E^\circ_\text{cathode} - E^\circ_\text{anode} \]
- Spontaneous Direction: The metal species with the more positive (higher) standard reduction potential (E°) spontaneously undergoes reduction and serves as the cathode.
- Concentration Cells (Nernst Equation): When both half-cells utilize the same metal electrode and electrolyte species, E°cell = 0 V. The overall cell potential is driven solely by concentration differences: \[ E_\text{cell} = E^\circ_\text{cell} - \frac{0.0592\text{ V}}{n}\log Q \] where Q = [dilute] / [concentrated].
2. Required Technical Skills
- Connecting and zeroing Vernier Voltage Probes for low-voltage electrochemical measurements.
- Polishing metal electrodes (copper, zinc, and unknown strips) with steel wool to remove oxide layers prior to immersion.
- Performing volumetric serial dilutions to prepare a range of standard electrolyte concentrations.
3. Critical Safety
- Heavy metal waste (copper, zinc, and unknown metal salt solutions) must be collected in designated waste containers. Do not pour down the drain.
- To construct a semi-microscale Cu/Zn voltaic cell and measure its potential (Ecell) using a voltage probe.
- To measure the cell potential of two voltaic cells constructed with unknown metal electrodes and identify the unknown metals using standard reduction potential tables.
- To construct a series of copper concentration cells, generate a linear Nernst plot (Ecell vs. log Q), and experimentally determine the Nernst slope and electron transfer number (n).
INTRODUCTION
In electrochemistry, a voltaic (galvanic) cell is an electrochemical system in which a spontaneous oxidation-reduction reaction generates electrical energy. A voltaic cell consists of two separate half-cells: an oxidation half-cell occurring at the anode (negative terminal) and a reduction half-cell occurring at the cathode (positive terminal). Electrical neutrality between the two solutions is maintained by a salt bridge containing an inert electrolyte such as KNO3.
In Parts A and B of this experiment, you will construct semi-microscale voltaic cells in a 24-well plate using metal strips immersed in 0.10 M aqueous solutions of their respective metal nitrates. You will measure the cell potential of a standard Cu/Zn cell and two cells pairing copper with unknown metal electrodes (X and Y). By observing lead polarities and measured cell potentials, you will determine the reduction potentials of the unknown metals and identify them from an activity series.
In Part C, you will explore concentration cells and test the mathematical model of the Nernst equation. You will prepare a 10-fold serial dilution series of Cu2+ solutions (1.0 M, 0.10 M, 0.010 M, 0.0010 M, and 0.00010 M) and measure the potential of each dilute half-cell paired against a 1.0 M reference half-cell. By plotting measured cell potentials against log Q, you will construct a linear Nernst plot to determine the experimental slope and evaluate the electron transfer number (n).
Concentration cells demonstrate how concentration gradients drive spontaneous voltage across half-cells:
- Drive Mechanism: Thermodynamics favors equalizing concentrations. Electrons flow from the dilute half-cell (anode, where Cu(s) → Cu2+(aq) + 2 e−) to the concentrated reference half-cell (cathode, where Cu2+(aq) + 2 e− → Cu(s)).
- Linear Nernst Model: Because E°cell = 0 V, the theoretical potential simplifies to: \[ E_\text{cell} = -\frac{0.0592\text{ V}}{n}\log Q \] A plot of Ecell vs. log Q yields a straight line with a theoretical slope of \(-0.0296\text{ V/decade}\) at 25°C for a 2-electron transfer (n = 2).
Cell Potential:
\[ E_\text{cell} = E_\text{cathode} - E_\text{anode} \]
Standard Cell Potential:
\[ E^\circ_\text{cell} = E^\circ_\text{cathode} - E^\circ_\text{anode} \]
Nernst Equation at 25°C (298.15 K):
\[ E_\text{cell} = E^\circ_\text{cell} - \frac{0.0592\text{ V}}{n}\log Q \]
- 9.1: Voltaic Cells - Experiment
- This page details safety measures, necessary equipment, and chemicals for experiments with voltaic cells. It outlines procedures to determine the standard electrode potential of a Cu/Zn cell and compares two unknown cells to copper, utilizing a logic matrix to rank the unknown metals by reducing strength. Additionally, it provides steps for preparing and testing a copper concentration cell and highlights proper chemical disposal practices.
- 9.2: Voltaic Cells - Pre-lab
- This page teaches students how to calculate the theoretical standard cell potential for a galvanic cell with zinc and copper electrodes. It includes steps for writing a balanced equation, completing half-reaction tables with standard reduction potentials, and applying the Nernst equation with varying ion concentrations. A logic check helps assess the reduction potential of an unknown metal compared to copper, illustrating practical uses of standard reduction potentials.
- 9.3: Voltaic Cells - Data and Report
- This page details a laboratory exercise in electrochemistry, focusing on copper/zinc concentration cells and analyzing unknown metals. Participants measure cell potentials, identify unknowns using standard reduction potentials, and compare measured averages to theoretical values calculated via the Nernst equation. Post-lab questions prompt reflection on accuracy through percent error and percent difference calculations, reinforcing the importance of validating experimental models.