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Ch.20 - Electrochemistry
Brown - Chemistry: The Central Science 15th Edition
Brown15th EditionChemistry: The Central ScienceISBN: 9780137542970Not the one you use?Change textbook
All textbooksBrown 15th EditionCh.20 - ElectrochemistryProblem 59
Chapter 20, Problem 59

Using the standard reduction potentials listed in Appendix E, calculate the equilibrium constant for each of the following reactions at 298 K:
(a) Fe(s) + Ni2+(aq) → Fe2+(aq) + Ni(s)
(b) Co(s) + 2 H+(aq) → Co2+(aq) + H2(g)
(c) 10 Br-(aq) + 2 MnO4-(aq) + 16 H+(aq) → 2 Mn2+(aq) + 8 H2O(l) + 5 Br2(l)

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1
Identify the half-reactions involved in the redox process. For the given reaction, the half-reactions are: \( \text{Br}^- \rightarrow \text{Br}_2 \) and \( \text{MnO}_4^- \rightarrow \text{Mn}^{2+} \).
Look up the standard reduction potentials (E°) for each half-reaction from Appendix E. The standard reduction potential for \( \text{Br}_2 + 2e^- \rightarrow 2\text{Br}^- \) and \( \text{MnO}_4^- + 8H^+ + 5e^- \rightarrow \text{Mn}^{2+} + 4H_2O \) are needed.
Calculate the standard cell potential (E°_cell) using the formula: \( E°_\text{cell} = E°_\text{cathode} - E°_\text{anode} \). Identify which half-reaction is the reduction (cathode) and which is the oxidation (anode).
Use the Nernst equation to relate the standard cell potential to the equilibrium constant (K): \( E°_\text{cell} = \frac{RT}{nF} \ln K \), where \( R \) is the gas constant, \( T \) is the temperature in Kelvin, \( n \) is the number of moles of electrons transferred, and \( F \) is Faraday's constant.
Rearrange the Nernst equation to solve for the equilibrium constant \( K \): \( \ln K = \frac{nFE°_\text{cell}}{RT} \). Substitute the known values to find \( K \).

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Key Concepts

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

Standard Reduction Potentials

Standard reduction potentials are measured voltages that indicate the tendency of a chemical species to gain electrons and be reduced. Each half-reaction has a specific potential, and these values are typically listed in tables. The more positive the potential, the greater the species' ability to be reduced. These values are essential for calculating the overall cell potential and determining the direction of electron flow in redox reactions.
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Nernst Equation

The Nernst equation relates the cell potential to the concentrations of the reactants and products at non-standard conditions. It is expressed as E = E° - (RT/nF)ln(Q), where E° is the standard cell potential, R is the gas constant, T is the temperature in Kelvin, n is the number of moles of electrons transferred, F is Faraday's constant, and Q is the reaction quotient. This equation is crucial for calculating the equilibrium constant from the standard potentials.
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Equilibrium Constant (K)

The equilibrium constant (K) quantifies the ratio of the concentrations of products to reactants at equilibrium for a given reaction at a specific temperature. It is derived from the standard Gibbs free energy change and can be calculated using the relationship K = e^(nFE°/RT). Understanding K is vital for predicting the extent of a reaction and how changes in conditions affect the position of equilibrium.
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Related Practice
Textbook Question

A cell has a standard cell potential of +0.177 V at 298 K. What is the value of the equilibrium constant for the reaction

(a) if n = 1?

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Textbook Question

At 298 K a cell reaction has a standard cell potential of +0.17 V. The equilibrium constant for the reaction is 5.5 × 105. What is the value of n for the reaction?

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Textbook Question

From each of the following pairs of substances, use data in Appendix E to choose the one that is the stronger reducing agent: (d) BrO3-1aq2 or IO3-1aq2

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