CHEM 1001: Chemical Dynamics Study Notes

Chapter 18: Electrochemistry

  • Introduction to Electrochemistry

    • Definition: The study of chemical processes that cause electrons to move, forming a current.
    • Relevance to batteries and energy production.

  • Batteries and Their Dynamics

    • Quote by Demetri Martin: "Batteries are the most dramatic object. Other things stop working or they break, but batteries… they die."
    • Quote by Don DeLillo: "The smoke alarm went off in the hallway upstairs, either to let us know the battery had just died or because the house was on fire."

  • Voltaic (Galvanic) Cells

    • Purpose: Generating electricity from spontaneous chemical reactions.
    • Key Components:
      • Zinc (Zn) metal dipped in a solution of Copper(II) ions ( ext{Cu}^{2+}).
      • Redox reactions: Involve the transfer of electrons, where oxidation and reduction occur simultaneously.
    • Electrical current is the flow of electric charge, resulting from movement of electrons or ions.
    • Reaction Example:
      • Oxidation: Zn(s) → ext{Zn}^{2+}(aq) + 2e−
      • Reduction: ext{Cu}^{2+}(aq) + 2e− → Cu(s)
      • Overall reaction:
        ext{Zn(s)} + ext{Cu}^{2+}(aq)
        ightarrow ext{Zn}^{2+}(aq) + ext{Cu(s)}

  • Electrochemical Cell Overview

    • Cell Parts:
      • Anode (Oxidation occurs):
        • Half-reaction: ext{Zn}(s) → ext{Zn}^{2+} + 2 e−
      • Cathode (Reduction occurs):
        • Half-reaction: ext{Cu}^{2+} + 2 e− → ext{Cu}(s)
    • Salt Bridge:
      • Core function: Allows ionic movement to maintain charge neutrality. Contains KNO₃(aq).
    • Diagram Structure:
      • Cell Diagram Format:
        • Oxidation on left, reduction on right:
        • extZn(s)∣extZn2+(aq)∣∣extCu2+(aq)∣extCu(s)ext{Zn(s)} | ext{Zn}^{2+}(aq) || ext{Cu}^{2+}(aq) | ext{Cu(s)}
    • Mnemonic: An Ox Red Cat - Oxidation occurs at Anode, Reduction at Cathode.

  • Key Differences: Voltaic vs. Electrolytic Cells

    • Voltaic Cells:
      • Driven by spontaneous reactions, generating current without external force;
      • Ecell > 0.
    • Electrolytic Cells:
      • Require external energy source to drive nonspontaneous reactions.
      • Ecell < 0.
    • Connection to spontaneity analyzed via Ecell calculation.

  • Calculating Ecell for Redox Reactions

    • Steps:
      1. Assign oxidation states (O.S.) of involved species.
      2. Write a balanced equation.
      3. Use standard reduction potentials from Table 18.1 to calculate E° cell.
    • Oxidation States (O.S.) Definition: A fictitious charge of an atom in a compound that represents its electron distribution.
    • Rules for O.S. Calculation:
      1. Charges must sum to total of the species.
      2. Group 1 metals: O.S. = +1, Group 2 metals: O.S. = +2.
      3. H: O.S. is usually +1.
      4. Common O.S. for elements (F, Cl, O, N, etc.) provided.

  • Balancing Redox Reactions

    • Process involves:
      1. Assigning O.S. for all atoms.
      2. Splitting into half-reactions.
      3. Balancing atoms in sequential order:
        • Other atoms > O > H > Charge (using e−).
    • Example Balancing Steps:
      • Initial Reaction: ext{UO}_2^{+} + ext{Cr}_2O_7^{2-}
        ightarrow ext{UO}_2^{2+} + ext{Cr}^{3+}
      • Balancing demonstration provided in sequential steps linked to oxidation and reduction.

  • Reduction Potentials

    • Definition: High E° means the species has a strong tendency to gain electrons; low E° indicates a strong tendency to lose electrons.
    • Table of standard electrode potentials provided for common half-reactions, indicating their strength as oxidizing or reducing agents.
    • Example:
      • ext{F}_2 + 2e^{-}
        ightarrow 2 ext{F}^{-}, ext{E°} = 2.87 V signals a strong oxidizer.

  • Standard Hydrogen Electrode (SHE)

    • Reference electrode defined as: 2 ext{H}^{+} + 2e^{-}
      ightarrow ext{H}_2(g)
    • Assigned E° = 0 V, allowing relative comparison of all other electrodes.
    • In cell diagrams: extPt(s)∣extH+(aq)∣extH2(g)∣∣(otherelectrode)ext{Pt(s)} | ext{H}^+(aq) | ext{H}_2(g) || (other electrode)

  • Calculating E°cell

    • In a spontaneous voltaic cell: E°cell = E°(cathode) - E°(anode) > 0
    • Spontaneous reactions involve ingestion of electrons at cathode and produce electrons at anode.
    • Numeric examples provided, reinforcing calculations related to spontaneous and non-spontaneous reactions.

  • Equilibrium and Thermodynamics

    • Relation between E°cell, ΔG° and equilibrium constant (K):
      • extΔG°=−nFE°cellext{ΔG°} = -nFE°_{cell}
      • E°cell=RTnFextlnKE°_{cell} = \frac{RT}{nF} ext{ln} K at standard conditions
    • Nernst Equation: Ecell=E°cell−RTnFextlnQE_{cell} = E°_{cell} - \frac{RT}{nF} ext{ln} Q
      • Demonstrates effects of concentration and temperature on cell potential and spontaneity.

  • Applying the Nernst Equation

    • Concentration cells: Demonstration given how to determine cell potentials when concentrations differ across half-cells.
    • Example provided illustrating practical calculation of concentration and its impact on an electrochemical cell's performance.

  • Tutorial Practices

    • Series of tutorial problems that engage students in calculating ΔG, K, E°cell under various chemical reactions and conditions featuring both spontaneous and non-spontaneous conditions.
    • Applications of Nernst equations to solve for unknown concentrations in unknown solutions.

  • Concluding Topics in Electrochemistry

    • Relation of electrochemistry to real-world applications: Energy storage, battery technology, and the role in environmental interfaces.
    • Importance of understanding these principles for advancements in renewable energy.