EO

Experiment 8: Electrochemistry I - Galvanic and Electrolytic Cells

PURPOSE

To explore spontaneous and non-spontaneous electrochemical reactions through the construction and analysis of galvanic and electrolytic cells, and to experimentally determine Faraday’s constant using copper electrodes.

OBJECTIVES

Experimental Objectives

  • Construct electrochemical (galvanic) cells using various metal-metal ion half-cells and a salt bridge.

  • Measure cell potentials with a digital multimeter.

  • Create an electrolytic cell using copper electrodes in copper sulfate solution.

  • Drive a non-spontaneous reaction using an external power supply.

  • Measure mass changes in electrodes to determine Faraday’s constant.

  • Observe and analyze reactions between metal samples and metal ion solutions.

Learning Objectives

  • Understand the difference between spontaneous and non-spontaneous redox reactions.

  • Interpret electrochemical cell notation (line notation).

  • Apply standard reduction potentials to calculate cell voltages.

  • Correlate observed reactions with theoretical predictions using standard potentials.

  • Relate charge, current, time, and mass to experimentally determine Faraday’s constant.

UNDERSTANDING THE BACKGROUND

Galvanic Cells

  • Use spontaneous redox reactions to produce electricity.

  • Composed of two half-cells connected via a salt bridge and an external wire.

  • Anode (oxidation) is on the left; Cathode (reduction) is on the right.

  • Salt bridge maintains charge neutrality by allowing ion flow.

Line Notation Format

Example:
Fe(s) | Fe²⁺(aq) || Ag⁺(aq) | Ag(s)
Left = anode (oxidation), Right = cathode (reduction)

Electrolytic Cells

  • Use external voltage to drive non-spontaneous reactions.

  • Involve oxidation at the anode and reduction at the cathode.

  • The amount of substance deposited or dissolved is proportional to the current × time (Q = I × t).

  • Use gravimetric analysis to relate electron transfer to mass change and determine Faraday’s constant.

Key Equations

  • Cell Potential: E°cell = E°cathode − E°anode

  • Coulombs (Q): Q = current (A) × time (s)

  • Faraday’s Constant (F):
    =n⋅ΔmI⋅t⋅MCu​​

SUMMARY OF THE PROCEDURE

Part A: Electrochemical Reactions

  1. Place Cu, Pb, and Zn samples in separate wells.

  2. Add metal ion solutions to appropriate wells.

  3. Record visual changes over 10 minutes.

Part B: Galvanic Cells

  1. Clean and insert metal strips (Cu, Zn, Pb, Fe) in well plate with respective solutions.

  2. Prepare a salt bridge using KNO₃-soaked filter paper.

  3. Connect cells using alligator clips and measure voltages using a multimeter.

  4. Record cell voltages and determine which reactions are spontaneous.

Part C: Electrolytic Cell

  1. Clean, dry, and weigh two Cu electrodes.

  2. Submerge electrodes in CuSO₄ solution and connect to power supply.

  3. Run a constant current (0.5–0.7 A) for ~20 minutes.

  4. Reweigh electrodes to determine mass gain/loss.

  5. Repeat for three trials to determine Faraday’s constant.

ANALYZING THE DATA

Part A & B

  • Use redox tables to identify theoretical E° values for each pair.

  • Compare observed voltages to theoretical voltages; calculate % error.

  • Explain observed metal deposition/dissolution using cell potential data.

Part C

  • Calculate:

    • Charge transferred (Q = I × t)

    • Moles of Cu deposited: Δm / Molar Mass

    • Faraday’s constant from experimental data

  • Report average and standard deviation of Faraday’s constant from all trials.

  • Calculate % error from accepted value (F ≈ 96485 C/mol e⁻).

Discussion Prompts

  • Compare observed vs theoretical values.

  • Explain discrepancies (e.g., impurity, inaccurate salt bridge, voltage drift).

  • Discuss why Fe electrodes yield less accurate results.

  • Evaluate copper sulfate concentration change and mass changes at electrodes.