Cell Potentials and Galvanic Cells

  • Reduction Potentials & Spontaneity

    • Focus on electrochemistry, specifically reduction potentials.
    • Review of galvanic cells: look for false statements in context.

  • Galvanic Cells

    • Define galvanic (voltaic) cells: systems converting chemical energy from spontaneous redox reactions into electrical energy.
    • Oxidation-Reduction (redox) reactions involve species losing (oxidation) and gaining (reduction) electrons.

  • Half-Reactions

    • Split overall reactions into half-reactions: oxidation and reduction components.
    • Example:
      • Fe²⁺ + 2e⁻ → Fe (reduction).
      • Mg → Mg²⁺ + 2e⁻ (oxidation).
    • Number of electrons transferred in reactions: 2 (one for each element involved).
    • Identify oxidation (Mg) and reduction (Fe).

  • Electrode Types

    • Anode: where oxidation occurs (negative electrode).
    • Cathode: where reduction occurs (positive electrode).
    • In this example:
    • Anode is magnesium strip.
    • Cathode is iron strip.

  • Electron Flow

    • Electrons flow from anode (Mg) to cathode (Fe).
    • Mass changes:
    • Mg electrode: decreases in size (oxidation).
    • Fe electrode: increases in size (reduction).

  • Electrode & Cell Potentials

    • Definition of electrode potentials and cell potentials.
    • Standard cell potential, E°_cell, relates to spontaneity of redox reactions.
    • E°cell > 0: spontaneous; E°cell < 0: non-spontaneous.
    • E°_cell is the voltage generated by the galvanic cell.

  • Positive Cell Potential

    • For a galvanic cell to function spontaneously, the cell potential must be positive.
    • Relate to Gibbs Free Energy, ΔG:
    • ΔG < 0 for spontaneity; E°_cell > 0 (opposite signs).

  • Cell Construction

    • Conventional setup: anode on left, cathode on right.
    • Connections made via wire/voltmeter to measure E°_cell.

  • Standard Hydrogen Electrode

    • Defined as 0 volts; all other potentials compared against this.

  • Half Cell Potentials

    • E°_(half cell) measured under standard conditions (1 M ion concentration, 298 K).
    • Values extracted from tables refer to reductions, where electrons are reactants.
    • Example: Zn²⁺ + 2e⁻ → Zn has a potential of -0.76 V.

  • Identifying Oxidizing and Reducing Agents

    • Stronger oxidizing agent: higher (more positive) E°_(half cell).
    • Stronger reducing agent: lower (more negative) E°_(half cell).

  • Combining Half-Reactions

    • Identify which reaction is the oxidizing agent (more positive E°) and which is the reducing agent (more negative E°).
    • Flip the oxidation reaction (keeping E° the same) to write as an oxidation process.
    • Balance electron transfer between half-reactions.

  • Example of Combination

    • Half reactions:
    • E.g., Ag extsuperscript{+} + e extsuperscript{-} → Ag, E° = 0.80 V (cathode).
    • Zn → Zn²⁺ + 2e extsuperscript{-}, E° = -0.76 V (anode).
    • Multiply to balance electrons and add reactions.

  • Key Takeaway

    • Understanding and applying these concepts are vital for predicting reactions and spontaneity in electrochemical systems.
    • Further practice with combining half-reactions and understanding E° values will be crucial for the upcoming material on free energy change.