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.