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.