For the cell shown, the measured cell potential, Ecell, is -0.3657 V at 25 °C. Pt(s) | H₂(g, 0.777 atm) | H*(aq, ? M) || Cd²+ (aq, 1.00 M) | Cd(s) The balanced reduction half-reactions for the cell, and their respective standard reduction potential values, Eº, are 2 H+ (aq) + 2 e- H₂(g) Eº = 0.00 V Cd²+ (aq) + 2 e Cd(s) E° -0.403 V Calculate the H+ concentration. 0.23 [H+] = Incorrect M
For the cell shown, the measured cell potential, Ecell, is -0.3657 V at 25 °C. Pt(s) | H₂(g, 0.777 atm) | H*(aq, ? M) || Cd²+ (aq, 1.00 M) | Cd(s) The balanced reduction half-reactions for the cell, and their respective standard reduction potential values, Eº, are 2 H+ (aq) + 2 e- H₂(g) Eº = 0.00 V Cd²+ (aq) + 2 e Cd(s) E° -0.403 V Calculate the H+ concentration. 0.23 [H+] = Incorrect M
Chemistry
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ISBN:9781305957404
Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
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Chapter1: Chemical Foundations
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![### Electrochemical Cell and Calculating Hydrogen Ion Concentration
For the cell shown, the measured cell potential, \(E_{\text{cell}}\), is \(-0.3657 \text{ V}\) at \(25 \degree \text{C}\).
**Cell Notation:**
\[ \text{Pt(s)} \mid \text{H}_2(\text{g}, 0.777 \, \text{atm}) \mid \text{H}^+(\text{aq}, \, ? \, \text{M}) \parallel \text{Cd}^{2+}(\text{aq}, 1.00 \, \text{M}) \mid \text{Cd(s)} \]
**Balanced Reduction Half-Reactions:**
\[ 2\text{H}^+ (\text{aq}) + 2\text{e}^- \rightarrow \text{H}_2 (\text{g}) \qquad E^\circ = 0.00 \, \text{V} \]
\[ \text{Cd}^{2+} (\text{aq}) + 2\text{e}^- \rightarrow \text{Cd(s)} \qquad E^\circ = -0.403 \, \text{V} \]
**Problem Statement:**
Calculate the \(\text{H}^+\) concentration.
**Initial Answer Provided:**
\[ [\text{H}^+] = 0.23 \, \text{M} \]
**Feedback:**
Incorrect
#### Explanation:
Given the cell potential and standard reduction potentials, you may use the Nernst equation to find the concentration of \(\text{H}^+\). The answer provided was found to be incorrect, indicating a need for re-evaluation of the calculation process. Further steps involve adjusting calculations for accuracy.](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F3f33cc2f-0221-41a8-a3a2-cc96bc7960a2%2F508a8522-075f-4258-a43c-b67e2fa6bb8e%2Ftoxzf48_processed.png&w=3840&q=75)
Transcribed Image Text:### Electrochemical Cell and Calculating Hydrogen Ion Concentration
For the cell shown, the measured cell potential, \(E_{\text{cell}}\), is \(-0.3657 \text{ V}\) at \(25 \degree \text{C}\).
**Cell Notation:**
\[ \text{Pt(s)} \mid \text{H}_2(\text{g}, 0.777 \, \text{atm}) \mid \text{H}^+(\text{aq}, \, ? \, \text{M}) \parallel \text{Cd}^{2+}(\text{aq}, 1.00 \, \text{M}) \mid \text{Cd(s)} \]
**Balanced Reduction Half-Reactions:**
\[ 2\text{H}^+ (\text{aq}) + 2\text{e}^- \rightarrow \text{H}_2 (\text{g}) \qquad E^\circ = 0.00 \, \text{V} \]
\[ \text{Cd}^{2+} (\text{aq}) + 2\text{e}^- \rightarrow \text{Cd(s)} \qquad E^\circ = -0.403 \, \text{V} \]
**Problem Statement:**
Calculate the \(\text{H}^+\) concentration.
**Initial Answer Provided:**
\[ [\text{H}^+] = 0.23 \, \text{M} \]
**Feedback:**
Incorrect
#### Explanation:
Given the cell potential and standard reduction potentials, you may use the Nernst equation to find the concentration of \(\text{H}^+\). The answer provided was found to be incorrect, indicating a need for re-evaluation of the calculation process. Further steps involve adjusting calculations for accuracy.
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