Use tabulated electrode potentials to calculate 𝛥𝐺rxn° for each reaction at 25 °C. a. MnO2(s) + 4 H+(aq) + Cu(s)¡Mn2+(aq) + 2 H2O(l) + Cu2+(aq)

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Identify the half-reactions involved in the given redox reaction.

Write the reduction half-reaction and find its standard reduction potential (E°).

Write the oxidation half-reaction and find its standard oxidation potential (E°).

Calculate the standard cell potential (E°cell) by subtracting the oxidation potential from the reduction potential.

Use the formula 𝛥𝐺° = -nFE°cell to calculate the standard Gibbs free energy change, where n is the number of moles of electrons transferred and F is the Faraday constant.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Electrode Potentials

Electrode potentials, measured in volts, indicate the tendency of a chemical species to be reduced. Standard electrode potentials (E°) are measured under standard conditions (1 M concentration, 1 atm pressure, and 25 °C) and are used to predict the direction of redox reactions. A higher E° value suggests a greater likelihood of reduction, which is essential for calculating Gibbs free energy changes in electrochemical reactions.

Gibbs Free Energy (ΔG)

Gibbs free energy (ΔG) is a thermodynamic quantity that indicates the spontaneity of a reaction. It is calculated using the equation ΔG = ΔG° + RT ln(Q), where ΔG° is the standard Gibbs free energy change, R is the gas constant, T is the temperature in Kelvin, and Q is the reaction quotient. A negative ΔG indicates a spontaneous reaction, while a positive ΔG suggests non-spontaneity.

Nernst Equation

The Nernst equation relates the electrode potential of a half-cell to the concentrations of the reactants and products involved in the redox reaction. It is expressed as E = E° - (RT/nF) ln(Q), where E is the cell potential, E° is the standard potential, n is the number of moles of electrons transferred, and F is Faraday's constant. This equation is crucial for calculating the Gibbs free energy change in electrochemical reactions at non-standard conditions.

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