Electrochemistry Flashcards

Balancing Redox Reactions in Acidic Conditions

• Balancing the reaction: MnO4^- (aq) + Bi^{3+} (aq) \rightarrow Mn^{2+} (aq) + BiO3^- (aq)
• Identify oxidation states: Mn goes from +7 to +2 (reduction), Bi goes from +3 to +5 (oxidation).
• Balance in acidic conditions (using H_2O and H^+).
• Reduction half-reaction: 5e^- + 8H^+ + MnO4^- \rightarrow Mn^{2+} + 4H2O
• Oxidation half-reaction: 3H2O + Bi^{3+} \rightarrow BiO3^- + 6H^+ + 2e^-
• Multiply half-reactions to balance electrons:
  • Reduction: (5e^- + 8H^+ + MnO4^- \rightarrow Mn^{2+} + 4H2O) \times 2
  • Oxidation: (3H2O + Bi^{3+} \rightarrow BiO3^- + 6H^+ + 2e^-) \times 5
• Overall balanced reaction: 16H^+ + 2MnO4^- + 5Bi^{3+} \rightarrow 2Mn^{2+} + 5BiO3^- + 8H_2O

Balancing Redox Reactions in Basic Conditions

• Balancing the reaction: MnO4^- (aq) + C2O4^{2-} (aq) \rightarrow MnO2 (s) + CO_2 (g)
• Identify oxidation states: Mn goes from +7 to +4 (reduction), C goes from +3 to +4 (oxidation).
• Balance in basic conditions (using H_2O and OH^-.)
• Reduction half-reaction: 3e^- + 2H2O + MnO4^- \rightarrow MnO_2 + 4OH^-
• Oxidation half-reaction: C2O4^{2-} \rightarrow 2CO_2 + 2e^-
• Multiply half-reactions to balance electrons:
  • Reduction: (3e^- + 2H2O + MnO4^- \rightarrow MnO_2 + 4OH^-) \times 2
  • Oxidation: (C2O4^{2-} \rightarrow 2CO_2 + 2e^-) \times 3
• Overall balanced reaction: 2MnO4^- + 3C2O4^{2-} + 4H2O \rightarrow 2MnO2 + 6CO2 + 8OH^-

Electrochemical (Galvanic) Cells

• Galvanic cell: A device that uses the free energy of a reaction to produce an electromotive force (emf).
• Components:
  • Anode: Where oxidation occurs.
  • Cathode: Where reduction occurs.
  • Electrodes: Metal conductors in contact with the electrolyte solution.
  • Salt bridge: Allows ion transport to balance charge and maintain electroneutrality.

Salt Bridge

• Allows ions to transport from one side to another to balance charge.
• Driving force = emf

Potential

• Potential (V) = Potential Energy Difference / Charge
• Units: Joules per Coulomb (J/C)
• Current: Charge per second (C/s)
• 1 Coulomb (C) = 6.24 \times 10^{18} electrons

Electrochemical Cell Example

• Reaction: Zn(s) + CuSO4 (aq) \rightarrow ZnSO4 (aq) + Cu(s)
• Anode (oxidation): Zn(s) is oxidized to Zn^{2+} (aq).
• Cathode (reduction): Cu^{2+} (aq) is reduced to Cu(s).
• Cell Voltage: 1.10V (for 1M solutions at 25°C).
• Cell Notation: Zn(s)|ZnSO4(1M) || CuSO4(1M)|Cu(s)

Standard Hydrogen Electrode (SHE)

• Used as a reference electrode for measuring standard potentials.
• Platinum wire in HCl solution with H_2 gas.
• Standard conditions: [H^+] = 1M, Pressure of H_2 = 1 atm, Temperature = 25°C.
• E^° = 0V
• Reaction Example: Zn(s) + 2HCl(aq) \rightarrow H2(g) + ZnCl2 (aq)
• Cell Notation: Zn/ZnCl2 (1M) || H2 (g), 1 atm | HCl (1M) | Pt
• E_{cell} = 0.76V

Standard Reduction Potentials

• E{cell}^° = E{red}^° - E_{ox}^°
• Stronger oxidizing agents have more positive reduction potentials.
• Stronger reducing agents have more negative reduction potentials.
• Potentials are intensive properties: Do not multiply E^° if the reaction is multiplied.

Standard Reduction Potentials and Free Energy

• All potentials are measured against a Standard Hydrogen Electrode (S.H.E.).
• By convention, all half-reactions are reported as standard reduction potentials.
• Relationship between standard cell potential and Gibbs free energy:
  • \Delta G^° = -nFE^°_{cell}
  • n = number of moles of electrons transferred
  • F = Faraday's constant (96,485 C/mol)

Example Calculation

• Reaction: Cl2(g) + 2Br^-(aq) \rightarrow 2Cl^-(aq) + Br2(aq)
• E^°_{cell} = 1.36V - 1.09V = 0.27V
• \Delta G^° = -(2 \, mol \, e^-)(96,485 \, C/mol)(0.27 \, V) = -52 \times 10^3 \, J (Spontaneous)