Calculate the standard potential, E", for this reaction from its equilibrium constant at 298 K. X(s) + Y²+(aq) = X²+(aq)+Y(s) K = 1.67 × 10 %3D E = %3D
Calculate the standard potential, E", for this reaction from its equilibrium constant at 298 K. X(s) + Y²+(aq) = X²+(aq)+Y(s) K = 1.67 × 10 %3D E = %3D
Chemistry
10th Edition
ISBN:9781305957404
Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
Chapter1: Chemical Foundations
Section: Chapter Questions
Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...
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![**Calculate the Standard Potential**
Calculate the standard potential, \( E^* \), for this reaction from its equilibrium constant at 298 K.
\[ \text{X}(s) + \text{Y}^{2+}(aq) \rightleftharpoons \text{X}^{2+}(aq) + \text{Y}(s) \]
\[ K = 1.67 \times 10^5 \]
\[ E^* = \square \, \text{V} \]
**Instructions:**
To calculate the standard potential \( E^* \), use the Nernst equation and the relationship between the equilibrium constant \( K \) and the standard potential for the reaction at a given temperature.
\[ E^* = \frac{RT}{nF} \ln K \]
Where:
- \( R \) is the universal gas constant (\(8.314 \, \text{J/mol K}\))
- \( T \) is the temperature in Kelvin (\(298 \, \text{K}\))
- \( n \) is the number of moles of electrons transferred in the balanced equation
- \( F \) is the Faraday constant (\(96485 \, \text{C/mol}\))
- \( K \) is the equilibrium constant
**Graph/Diagram Explanation:**
There are no graphs or diagrams present. This is a text-based calculation problem.](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2Fd891009c-d394-441a-930b-8d5c1fe8da7b%2F144cf606-4852-4ae9-b7ad-9f6d0e0978d0%2Fcy0bjg.jpeg&w=3840&q=75)
Transcribed Image Text:**Calculate the Standard Potential**
Calculate the standard potential, \( E^* \), for this reaction from its equilibrium constant at 298 K.
\[ \text{X}(s) + \text{Y}^{2+}(aq) \rightleftharpoons \text{X}^{2+}(aq) + \text{Y}(s) \]
\[ K = 1.67 \times 10^5 \]
\[ E^* = \square \, \text{V} \]
**Instructions:**
To calculate the standard potential \( E^* \), use the Nernst equation and the relationship between the equilibrium constant \( K \) and the standard potential for the reaction at a given temperature.
\[ E^* = \frac{RT}{nF} \ln K \]
Where:
- \( R \) is the universal gas constant (\(8.314 \, \text{J/mol K}\))
- \( T \) is the temperature in Kelvin (\(298 \, \text{K}\))
- \( n \) is the number of moles of electrons transferred in the balanced equation
- \( F \) is the Faraday constant (\(96485 \, \text{C/mol}\))
- \( K \) is the equilibrium constant
**Graph/Diagram Explanation:**
There are no graphs or diagrams present. This is a text-based calculation problem.
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