SectionC13_BB - Tagged

Page 1: Introduction to Cell Potentials

  • Title: Chemistry for Bioscientists I Section C13 Cell Potentials

  • Course: CHEM10021

  • University: The University of Manchester

  • Introduction to cell potentials in relation to bioscientists.

Page 2: Types of Electrode

  • Metal Electrode:

    • Composed of metal in contact with an ionic solution.

  • Redox Electrode:

    • Involves an inert metal in contact with a solution containing a species in two oxidation states.

  • Gas Electrode:

    • Consists of a gas in equilibrium with its ionic solution in the presence of an inert metal.

  • Insoluble Salt Electrode:

    • Involves a metal (M) coated with an insoluble salt (MX) immersed in a solution containing X- ions.

Page 3: Notation

  • Vertical Bar (|):

    • Denotes a barrier (e.g., Ag | AgCl).

  • Redox Electrode Notation:

    • Example: "M I Red, Ox" indicates a redox couple Red/Ox.

    • For instance, NAD+ redox couple as NAD+ + H+ + 2e- → NADH.

  • Double Vertical Bars (||):

    • Represent a connection, such as a salt bridge.

Page 4: Cell Potential

  • Electrolytic Cell Representation:

    • Example: Pt | NADH, NAD+, H+ || H2O2, H+, O | Pt.

  • Anode Reaction:

    • Left-hand half-cell: NAD+ + H+ + 2e- → NADH.

  • Cathode Reaction:

    • Right-hand half-cell: O2 + 2H+ + 2e- → 2H2O2.

  • Overall Cell Reaction:

    • NADH + O2 + H+ → H2O2 + NAD+.

Page 5: Understanding Cell Potential

  • Cell Potential (E):

    • Defined as the difference between the electrode potentials of the cathode and anode.

      • E = ERHS - ELHS.

  • Standard Cell Potential (E°):

    • E° = E° RHS - E° LHS.

  • Positive cell potential (E > 0) indicates a spontaneous cell reaction.

Page 6: Zero Current Potential

  • Defined as the maximum amount of non-expansion work (Wmax = ΔG).

  • Max work possible if the process is reversible.

  • Zero-current cell potential (cell emf) is the voltage measured between electrodes with no current flowing.

Page 7: Standard Hydrogen Electrode (SHE)

  • Standard Reference:

    • SHE is assigned the zero value for electrode potential.

  • Configuration: H2 (1 atm) & Pt | H+ (aq).

  • E° for SHE = 0V at all temperatures.

Page 8: Standard (Reduction) Potentials

  • Electrochemical Data:

    • Standard electrode potentials (E°) are tabulated for half-cell reactions.

  • E° indicates the zero-current cell potential relative to SHE.

Page 9: Standard Electrode Potentials (E°)

  • Strong Electron Acceptors:

    • F2 + 2e- → 2F- (E° = +2.87V).

    • Cl2 + 2e- → 2Cl- (E° = +1.36V).

    • O2 + 4H+ + 4e- → 2H2O (E° = +1.23V).

  • Strong Electron Donors:

    • Na+ + e- → Na (E° = -2.71V).

Page 10: Biological Standard Potentials (E°)

  • Half-Cell Reactions:

    • O2 + 4H+ + 4e- → 2H2O (E° = +0.82V).

    • UQ + 2H+ + 2e- → UQH2 (E° = +0.11V).

    • NAD+ + H+ + 2e- → NADH (E° = -0.32V).

Page 11: Cell Potentials and Thermodynamics

  • Relationship between electrochemistry and thermodynamics.

  • Gibbs Free Energy Calculation:

    • ΔG = Wmax = -nFE, where F = Faraday's constant (96485 C-mol⁻¹).

Page 12: Thermodynamic Relationships

  • Gibbs Free Energy:

    • G° = -RT lnK.

    • ΔG = -nFE°
      - lnK = (RT/E°).

Page 13: The Nernst Equation

  • Equation:

    • RT E = E° - (RT/nF) lnQ.

  • Applicable for half-cell reactions (reduction).

Page 14: Membrane Potential

  • Equation:

    • RT E = (Rln[A]out/[A]in)/F.

  • Membrane potential related to concentrations of species inside and outside the membrane.

Page 15: Summary and Conclusion

  • Title: Chemistry for Bioscientists I Section C13 Cell Potentials

  • Overall conclusion about the significance of cell potentials in biology.

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