Open Question
What is the equilibrium constant for each of the reactions mentioned in Problem 66?
Ch.19 - Electrochemistry
Chapter 19, Problem 71
Calculate ΔG°_rxn and E°_cell for a redox reaction with n = 2 that has an equilibrium constant of K = 25 (at 25 °C).
Verified step by step guidance1
Identify the relationship between the equilibrium constant (K) and the standard Gibbs free energy change (ΔG°_rxn) using the equation: ΔG°_rxn = -RT ln(K), where R is the universal gas constant (8.314 J/mol·K) and T is the temperature in Kelvin (298 K for 25 °C).
Substitute the given values into the equation: ΔG°_rxn = -(8.314 J/mol·K)(298 K) ln(25).
Calculate the natural logarithm of the equilibrium constant, ln(25).
Use the calculated value of ln(25) to find ΔG°_rxn by completing the multiplication and division in the equation.
Relate ΔG°_rxn to the standard cell potential (E°_cell) using the equation: ΔG°_rxn = -nFE°_cell, where n is the number of moles of electrons transferred (n = 2) and F is Faraday's constant (96485 C/mol). Solve for E°_cell by rearranging the equation: E°_cell = -ΔG°_rxn / (nF).
Related Practice
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Calculate the equilibrium constant for the reaction between Ni2+(aq) and Cd(s) at 25 °C.
Textbook Question
Calculate the equilibrium constant for the reaction between Fe2+(aq) and Zn(s) (at 25 °C).
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Open Question
Calculate ΔG°rxn and E°cell for a redox reaction with n = 3 that has an equilibrium constant of K = 0.050 (at 25 °C).
Textbook Question
A voltaic cell employs the following redox reaction: Sn2+(aq) + Mn(s) → Sn(s) + Mn2+(aq) Calculate the cell potential at 25 °C under each set of conditions. c. [Sn2+] = 2.00 M; [Mn2+] = 0.0100 M
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Open Question
A voltaic cell employs the redox reaction: 2 Fe3+(aq) + 3 Mg(s) → 2 Fe(s) + 3 Mg2+(aq). Calculate the cell potential at 25 °C under each set of conditions. a. standard conditions. b. [Fe3+] = 1.0 × 10^-3 M; [Mg2+] = 2.50 M. c. [Fe3+] = 2.00 M; [Mg2+] = 1.5 × 10^-3 M.