3.3 Thermodynamic Relationships in Electrochemistry

Electrochemical cells are powerhouses of energy conversion. They transform chemical energy into electrical energy, or vice versa, through redox reactions. Understanding the thermodynamics behind these processes is crucial for optimizing battery performance and developing new energy technologies.

Gibbs energy, cell potential, and equilibrium constants are key concepts in electrochemical thermodynamics. These factors determine the spontaneity and efficiency of reactions in batteries and fuel cells. Temperature also plays a vital role, affecting cell potentials and reaction rates in ways that impact real-world applications.

Thermodynamic Relationships in Electrochemical Cells

Gibbs energy and cell potential

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• Gibbs free energy ($\Delta G$) measures the maximum amount of non-expansion work that can be extracted from a thermodynamically closed system (galvanic cell)
• $\Delta G$ is related to the cell potential ($E_{cell}$) and the number of electrons transferred ($n$) in an electrochemical reaction by the equation: $\Delta G = -nFE_{cell}$
  • $F$ is Faraday's constant, equal to 96,485 C/mol, representing the charge carried by one mole of electrons
• A negative $\Delta G$ indicates a spontaneous reaction (battery discharge), while a positive $\Delta G$ indicates a non-spontaneous reaction (electrolysis)
  • Example: In a lead-acid battery, the discharge reaction has a negative $\Delta G$, allowing the battery to provide electrical energy
• More negative cell potentials correspond to more spontaneous reactions, as they result in a more negative $\Delta G$
• Increasing the number of electrons transferred in a reaction makes $\Delta G$ more negative, favoring spontaneity
  • Example: The reduction of aluminum ions to aluminum metal ($Al^{3+} + 3e^- \rightarrow Al$) involves three electrons and is more spontaneous than the reduction of copper ions to copper metal ($Cu^{2+} + 2e^- \rightarrow Cu$)

Equilibrium constant and standard potential

• The equilibrium constant ($K$) is related to the standard cell potential ($E_{cell}^{\circ}$) by the Nernst equation: $E_{cell}^{\circ} = \frac{RT}{nF} \ln K$
  • $R$ is the universal gas constant (8.314 J/mol·K)
  • $T$ is the temperature in Kelvin
  • $E_{cell}^{\circ}$ is the cell potential under standard conditions (1 M concentrations, 1 atm pressure, 25℃)
• Rearranging the Nernst equation yields: $\ln K = \frac{nFE_{cell}^{\circ}}{RT}$, allowing calculation of $K$ from $E_{cell}^{\circ}$
• A larger equilibrium constant indicates a more favorable reaction and corresponds to a more positive standard cell potential
  • Example: The standard cell potential for the Daniell cell ($Zn|Zn^{2+}||Cu^{2+}|Cu$) is 1.10 V, corresponding to a large equilibrium constant and a favorable reaction

Temperature dependence of cell potentials

• The Gibbs-Helmholtz equation relates $\Delta G$ to the change in enthalpy ($\Delta H$) and temperature: $\Delta G = \Delta H - T\Delta S$, where $\Delta S$ is the change in entropy
• Substituting the relationship between $\Delta G$ and $E_{cell}$ yields: $-nFE_{cell} = \Delta H - T\Delta S$
• Rearranging the equation gives the temperature dependence of the cell potential: $E_{cell} = -\frac{\Delta H}{nF} + \frac{T\Delta S}{nF}$
• The temperature coefficient of the cell potential is given by: $\frac{dE_{cell}}{dT} = \frac{\Delta S}{nF}$, showing how $E_{cell}$ changes with temperature
  • Example: For the lead-acid battery, $\frac{dE_{cell}}{dT}$ is negative, meaning the cell potential decreases with increasing temperature

Entropy and enthalpy in electrochemistry

• The change in entropy ($\Delta S$) can be calculated from the temperature coefficient of the cell potential: $\Delta S = nF \frac{dE_{cell}}{dT}$
  • Example: If $\frac{dE_{cell}}{dT}$ is positive, $\Delta S$ is also positive, indicating an increase in disorder during the reaction
• The change in enthalpy ($\Delta H$) can be calculated using the Gibbs-Helmholtz equation and the relationship between $\Delta G$ and $E_{cell}$: $\Delta H = -nFE_{cell} + T\Delta S$
• Alternatively, $\Delta H$ can be calculated using the standard enthalpy of formation ($\Delta H_f^{\circ}$) of the products and reactants: $\Delta H = \sum \Delta H_f^{\circ}(products) - \sum \Delta H_f^{\circ}(reactants)$
  • Example: In a fuel cell, the enthalpy change can be calculated from the enthalpies of formation of the fuel (hydrogen) and the product (water)