▼ Part A For the reaction 2Co³+ (aq) + 2Cl(aq)→2Co²+ (aq) + Cl₂ (g). E = 0.483 V what is the cell potential at 25 °C if the concentrations are [Co³+] = 0.425 M. [Co²+] = 0.384 M, and [Cl] = 0.337 M, and the pressure of Cl₂ is PC₁₂ = 8.30 atm ? Express your answer with the appropriate units. ► View Available Hint(s) E= HA Value Units S ?
▼ Part A For the reaction 2Co³+ (aq) + 2Cl(aq)→2Co²+ (aq) + Cl₂ (g). E = 0.483 V what is the cell potential at 25 °C if the concentrations are [Co³+] = 0.425 M. [Co²+] = 0.384 M, and [Cl] = 0.337 M, and the pressure of Cl₂ is PC₁₂ = 8.30 atm ? Express your answer with the appropriate units. ► View Available Hint(s) E= HA Value Units S ?
Chemistry
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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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![**Understanding the Nernst Equation in Electrochemistry**
The Nernst equation is crucial in electrochemistry for calculating cell potentials under non-standard conditions. The equation is given by:
\[ E = E^o - \frac{2.303 \, RT}{nF} \log_{10} Q \]
where:
- \( E \) is the potential in volts.
- \( E^o \) is the standard potential in volts.
- \( R \) is the gas constant.
- \( T \) is the temperature in kelvins.
- \( n \) is the number of moles of electrons transferred.
- \( F \) is the Faraday constant.
- \( Q \) is the reaction quotient.
For a common reference temperature of 25°C or 298 K, the equation can be simplified to:
\[ E = E^o - \left( \frac{0.0592}{n} \right) \log Q \]
*The reaction quotient \( Q \) is generally expressed as:*
\[ Q = \frac{[\text{products}]^z}{[\text{reactants}]^v} \]
This table assumes standard electrode potentials are measured with solutions at 1.00 M concentration and gases at 1.00 atm. The Nernst equation enables the calculation of \( E \) under varying concentration and pressure conditions.
**Example Problem**
*Consider the reaction:*
\[ 2\text{Co}^{3+} (\text{aq}) + 2\text{Cl}^{-} (\text{aq}) \rightarrow 2\text{Co}^{2+} (\text{aq}) + \text{Cl}_2(\text{g}) \]
*With \( E^o = 0.483 \, \text{V} \), find the cell potential at 25°C, given:*
- \([\text{Co}^{3+}] = 0.425 \, M\)
- \([\text{Co}^{2+}] = 0.384 \, M\)
- \([\text{Cl}^{-}] = 0.337 \, M\)
- \( P_{\text{Cl}_2} = 8.30 \, \text{atm} \)
*Submit your calculated cell potential with the appropriate units.*](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2Fb0be78df-234d-4653-8235-8227b9098990%2Ff8ba32c9-8861-4a9e-bebf-be7e05a26c0c%2Fprx53z_processed.jpeg&w=3840&q=75)
Transcribed Image Text:**Understanding the Nernst Equation in Electrochemistry**
The Nernst equation is crucial in electrochemistry for calculating cell potentials under non-standard conditions. The equation is given by:
\[ E = E^o - \frac{2.303 \, RT}{nF} \log_{10} Q \]
where:
- \( E \) is the potential in volts.
- \( E^o \) is the standard potential in volts.
- \( R \) is the gas constant.
- \( T \) is the temperature in kelvins.
- \( n \) is the number of moles of electrons transferred.
- \( F \) is the Faraday constant.
- \( Q \) is the reaction quotient.
For a common reference temperature of 25°C or 298 K, the equation can be simplified to:
\[ E = E^o - \left( \frac{0.0592}{n} \right) \log Q \]
*The reaction quotient \( Q \) is generally expressed as:*
\[ Q = \frac{[\text{products}]^z}{[\text{reactants}]^v} \]
This table assumes standard electrode potentials are measured with solutions at 1.00 M concentration and gases at 1.00 atm. The Nernst equation enables the calculation of \( E \) under varying concentration and pressure conditions.
**Example Problem**
*Consider the reaction:*
\[ 2\text{Co}^{3+} (\text{aq}) + 2\text{Cl}^{-} (\text{aq}) \rightarrow 2\text{Co}^{2+} (\text{aq}) + \text{Cl}_2(\text{g}) \]
*With \( E^o = 0.483 \, \text{V} \), find the cell potential at 25°C, given:*
- \([\text{Co}^{3+}] = 0.425 \, M\)
- \([\text{Co}^{2+}] = 0.384 \, M\)
- \([\text{Cl}^{-}] = 0.337 \, M\)
- \( P_{\text{Cl}_2} = 8.30 \, \text{atm} \)
*Submit your calculated cell potential with the appropriate units.*
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