Electrochemical Cells and Potentials Notes

Electrochemistry Overview

• Electrochemistry: Study of chemical processes that cause electrons to move.
• Electrochemical cells: Devices that convert chemical energy into electrical energy.

Key Concepts

• Movement of Electrons: Essential for various chemical reactions and energy conversion technologies.
• Galvanic Cells: Type of electrochemical cell that generates electrical energy from spontaneous chemical reactions.

Standard Cell Potentials

• Standard Cell Potential, E°: Voltage for oxidation/reduction reactions under standard conditions.

• Example Reaction:
  Cl2(g) + 2Br^- → 2Cl^- + Br2(l)
  with
  E° = 0.30V

• Half-Cell Reactions:

  • Reduction example:
    Cl_2(g) + 2e^- → 2Cl^- ext{; } E° = 1.36V
  • Oxidation example:
    2Br^- → Br_2(l) + 2e^- ext{; } E° = -1.06V

• Total Cell Potential Calculation:

  • Sum of two half-cell potentials:
    E° = E°{reduction} + E°{oxidation}
  • Example:
    E° = 1.36 + (-1.06) = 0.30V

Conditions for Standard Potentials

• Voltages are measured under standard conditions (1 M concentration, 1 atm pressure) at 298 K.
• Standard Hydrogen Electrode (SHE) is the reference point (0 V).

Reaction Spontaneity

• Positive E° indicates a spontaneous reaction, while negative E° indicates non-spontaneity.

Voltaic Cell Structure

• Components: Anode, cathode, salt bridge, and external circuit.
• Example Cell Reaction:
  Zn(s) + Cu^{2+} → Zn^{2+} + Cu(s)
  where
  E° = 0.76 + 0.34 = 1.10V
• Salt Bridge: Maintains electrical neutrality and allows ion movement.

Key Considerations in Experiments

• Multiple factors affect cell voltage readings:
  • Non-standard states of reagents
  • Temperature deviation from 298 K
  • Presence of contaminants
  • Internal resistances
  • Current load during measurement

Safety Precautions

• Always wear eye protection and gloves when handling solutions and electrodes.

Measurement Procedure

1. Prepare the apparatus using clean equipment and agar gel as the salt bridge.
2. Fill half-cells with appropriate metal ion solutions.
3. Connect half-cells to a voltmeter to measure voltage.
4. Record measured voltages, identify anode/cathode, and write down half-reactions.

Cell Reaction Examples

• Copper/Lead (Cu/Pb):
  • Measure and record values.
• Silver/Lead (Ag/Pb):
  • Determine voltage and reactions.
• Zinc/Lead (Zn/Pb):
  • Conduct voltage measurement.

Developing a Reduction Potential Table

• Set standard for one reaction (e.g., Pb^{2+} + 2e^- → Pb at 0.000V).
• Calculate other half-reaction potentials based on measured voltages.

Adjustment of Measured Values

• Compare adjusted half-cell potentials against accepted standard values from reference tables.
• Calculate differences to analyze results.

Waste Disposal

• Dispose of ionic solutions and agar appropriately. Clean and return all equipment after use.