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A mad scientist is attempting to make copper metal via the reaction:   Cu2+ (aq) + Cr (s) → Cu (s) + Cr3+   a) Use the chart below to find the standard cell potential (E0cell) b) Based on the number of half reactions, what is the number of electrons (n) that would be exchanged in a complete reaction?   c) Calculate the cell potential for this reaction at 20.7 °C when [Cu2+] = 0.191 M and [Cr3+] = 0.061 M.

A mad scientist is attempting to make copper metal via the reaction:   Cu2+ (aq) + Cr (s) → Cu (s) + Cr3+   a) Use the chart below to find the standard cell potential (E0cell) b) Based on the number of half reactions, what is the number of electrons (n) that would be exchanged in a complete reaction?   c) Calculate the cell potential for this reaction at 20.7 °C when [Cu2+] = 0.191 M and [Cr3+] = 0.061 M.

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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
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A mad scientist is attempting to make copper metal via the reaction:

 

Cu2+ (aq) + Cr (s) → Cu (s) + Cr3+

 

a) Use the chart below to find the standard cell potential (E0cell)


b) Based on the number of half reactions, what is the number of electrons (n) that would be exchanged in a complete reaction?

 

c) Calculate the cell potential for this reaction at 20.7 °C when [Cu2+] = 0.191 M and [Cr3+] = 0.061 M.

### Electrode Potentials of Various Half-Reactions This table lists the standard electrode potentials (\(E^\circ\)) for various half-reactions involving copper (Cu) and chromium (Cr). | **Half-Reaction** | **\(E^\circ\) (V)** | |----------------------------------|-------------------| | \( \text{Cu}^+ + e^- \rightarrow \text{Cu (s)} \) | 0.52 | | \( \text{Cu}^{2+} + 2 e^- \rightarrow \text{Cu (s)} \) | 0.34 | | \( \text{Cu}^{2+} + e^- \rightarrow \text{Cu}^+ \) | 0.15 | | \( \text{Cr}^{3+} + e^- \rightarrow \text{Cr}^{2+} \) | -0.41 | | \( \text{Cr}^{3+} + 3 e^- \rightarrow \text{Cr (s)} \) | -0.74 | | \( \text{Cr}^{2+} + 2 e^- \rightarrow \text{Cr (s)} \) | -0.91 | ### Details: - **Half-Reaction:** The left side of the equation indicates the reactants, and the right side indicates the products. The electrons (\(e^-\)) involved are shown in the equation. - **\(E^\circ\) (V):** The standard electrode potential for each half-reaction, measured in volts (V). #### Explanation of Electrode Potentials: - A positive \(E^\circ\) value indicates a greater tendency for the species to gain electrons (reduce). - A negative \(E^\circ\) value indicates a lesser tendency to gain electrons (often reflecting a stronger tendency to lose electrons, i.e., to oxidize). These half-reactions and their corresponding \(E^\circ\) values are essential in understanding the electrochemical series and predicting the spontaneity of redox reactions.
Transcribed Image Text:### Electrode Potentials of Various Half-Reactions This table lists the standard electrode potentials (\(E^\circ\)) for various half-reactions involving copper (Cu) and chromium (Cr). | **Half-Reaction** | **\(E^\circ\) (V)** | |----------------------------------|-------------------| | \( \text{Cu}^+ + e^- \rightarrow \text{Cu (s)} \) | 0.52 | | \( \text{Cu}^{2+} + 2 e^- \rightarrow \text{Cu (s)} \) | 0.34 | | \( \text{Cu}^{2+} + e^- \rightarrow \text{Cu}^+ \) | 0.15 | | \( \text{Cr}^{3+} + e^- \rightarrow \text{Cr}^{2+} \) | -0.41 | | \( \text{Cr}^{3+} + 3 e^- \rightarrow \text{Cr (s)} \) | -0.74 | | \( \text{Cr}^{2+} + 2 e^- \rightarrow \text{Cr (s)} \) | -0.91 | ### Details: - **Half-Reaction:** The left side of the equation indicates the reactants, and the right side indicates the products. The electrons (\(e^-\)) involved are shown in the equation. - **\(E^\circ\) (V):** The standard electrode potential for each half-reaction, measured in volts (V). #### Explanation of Electrode Potentials: - A positive \(E^\circ\) value indicates a greater tendency for the species to gain electrons (reduce). - A negative \(E^\circ\) value indicates a lesser tendency to gain electrons (often reflecting a stronger tendency to lose electrons, i.e., to oxidize). These half-reactions and their corresponding \(E^\circ\) values are essential in understanding the electrochemical series and predicting the spontaneity of redox reactions.
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