UNIT: 9.11

  • Introduction

    • Focus on subunit 9.11: Electrolysis and Faraday’s Law.
    • Previous subunits covered galvanic/voltaic cells where redox reactions generate electrical power.
    • Today's topic: Electrolytic cells which operate the galvanic cell in reverse.

  • Definition of Electrolytic Cell

    • Functions by forcing the flow of electrons in the opposite direction compared to galvanic cells.
    • Requires an external power source to operate, as opposed to a galvanic cell where reactions occur spontaneously.

  • Why Use Electrolysis?

    • Enables processes like electroplating: Reverse a reaction where a metal electrode is losing mass and rather gain metal.
    • Example: In a galvanic cell, a copper electrode may lose mass, whereas in electroplating, we want to deposit copper on that electrode.

  • Setup of Galvanic vs. Electrolytic Cells

    • Galvanic Cell:

    • Components are in separate compartments.

    • Anode and cathode defined based on spontaneous electron flow facilitated by a salt bridge.

    • Example: Electrons flow from zinc to copper.

    • Electrolytic Cell:

    • Components are in the same compartment; hence, no salt bridge required.

    • Anode and cathode positions switch: Oxidation occurs at what is now defined as cathode.

    • Example: Reverse current flow, electrons now move from copper to zinc.

  • Key Differences

    • Galvanic Cell:

    • Favorable reaction, produces electrical energy, requires no external power.

    • Electrolytic Cell:

    • Unfavorable reaction (requires energy input), consumes electrical energy.

    • Always involves oxidation at the anode and reduction at the cathode.

  • Electrolysis Example:

    • Sample problem on electrolysis of water.
    • Part A: Balancing equation and determining half reactions for hydrogen and oxygen.
    • Part B: Calculating the standard cell potential which will yield a negative value. This is typical for electrolytic cells.

  • Understanding Cell Potential:

    • For an electrolytic reaction, the total recommended power supply needs to be greater than -1.53 volts (example)
    • Relation of cell potential to Gibbs Free Energy (ΔG): Negative cell potential yields positive ΔG, indicating non-favorable reaction.

  • Stoichiometry of Electrolysis

    • Example on determining the mass of chromium produced in electrolysis of chromium(II) nitrate solution.
    • Important concepts:
    • Units for current (Amperes, A) and charge (Coulombs, C).
    • Faraday’s constant relates moles of electrons to charge passed.

  • Calculation Steps

    • Convert amperes to coulombs over time.
    • Use dimensional analysis to convert charge to moles of electrons, then to moles of chromium.
    • Finally, convert moles of chromium to grams using the periodic table.
    • Final calculation yields a mass of 15 grams of chromium produced.

  • Shortcut Method for Calculations: ATM fee equation

    • Defines mass of deposited metal as related to current, time, molar mass, Faraday’s constant, and moles of electrons.
    • Useful for various forms of calculations where unknowns could vary (mass, current, time).

  • Multiple Choice Example:

    • Evaluating a simple Faraday's problem for the electroplating of nickel.
    • Demonstrated how to use dimensional analysis to derive the correct mass.

  • General Note About Faraday’s Law:

    • Total charge ($Q$) is the product of current ($I$) and time ($t$).
    • Further, 1 Faraday indicates moles of electrons transferred in electrolytic cell reactions.

  • Conclusion

    • Final notes on AP chemistry content wrapping up the curriculum.
    • Encouragement for upcoming examinations, wishing students success.