### Electrochemical Reactions and Gibbs Free Energy Calculation#### Given Reactions:1. \( 2 \text{Hg}^{2+} (\text{aq}) + 2 \text{e}^- \rightarrow \text{Hg}_2^{2+} (\text{aq}) \) - Standard electrode potential, \( E^\circ = 0.92 \, \text{V} \)2. \( \text{Hg}_2^{2+} (\text{aq}) + 2 \text{e}^- \rightarrow 2 \text{Hg} (\text{l}) \) - Standard electrode potential, \( E^\circ = 0.85 \, \text{V} \)### Task:**Calculate \(\Delta G^\circ\) for the following process at 25°C:**\[ \text{Hg}_2^{2+} (\text{aq}) \rightarrow \text{Hg}^{2+} (\text{aq}) + \text{Hg} (\text{l}) \]### Explanation:To calculate the Gibbs Free Energy (\(\Delta G^\circ\)), you will use the standard electrode potentials from the given half-reactions. Apply the formula:\[ \Delta G^\circ = -nFE^\circ_{\text{cell}} \]where:- \( n \) is the number of moles of electrons transferred.- \( F \) is the Faraday constant (approximately \( 96485 \, \text{C/mol} \)).- \( E^\circ_{\text{cell}} \) is the standard cell potential, which can be calculated by combining the given electrode potentials.**Note:** Enter your calculated \(\Delta G^\circ\) value in the box provided, in units of \(\text{kJ/mol}\). **Box: [ _____ ] kJ/mol**

Answered: Image /qna-images/answer/44cea6ca-1f86-42e5-88e6-ef80e64fbdf7.jpg

Answered: 2Hg²+ (aq) + 2e Hg₂²+ (aq) + 2e¯…

2Hg²+ (aq) + 2e Hg₂²+ (aq) + 2e¯ calculate AG Hg₂+ (aq) Mac Hg₂ 2+ (aq) 2Hg(1) = 0.92 V Eº = 0.85 V 0 for the following process at 25°C: Hg2+ (aq) + Hg(1) kJ/mol

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

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Question

### Electrochemical Reactions and Gibbs Free Energy Calculation

#### Given Reactions:

1. \( 2 \text{Hg}^{2+} (\text{aq}) + 2 \text{e}^- \rightarrow \text{Hg}_2^{2+} (\text{aq}) \)
- Standard electrode potential, \( E^\circ = 0.92 \, \text{V} \)

2. \( \text{Hg}_2^{2+} (\text{aq}) + 2 \text{e}^- \rightarrow 2 \text{Hg} (\text{l}) \)
- Standard electrode potential, \( E^\circ = 0.85 \, \text{V} \)

### Task:

**Calculate \(\Delta G^\circ\) for the following process at 25°C:**

\[ \text{Hg}_2^{2+} (\text{aq}) \rightarrow \text{Hg}^{2+} (\text{aq}) + \text{Hg} (\text{l}) \]

### Explanation:

To calculate the Gibbs Free Energy (\(\Delta G^\circ\)), you will use the standard electrode potentials from the given half-reactions. Apply the formula:

\[ \Delta G^\circ = -nFE^\circ_{\text{cell}} \]

where:
- \( n \) is the number of moles of electrons transferred.
- \( F \) is the Faraday constant (approximately \( 96485 \, \text{C/mol} \)).
- \( E^\circ_{\text{cell}} \) is the standard cell potential, which can be calculated by combining the given electrode potentials.

**Note:** Enter your calculated \(\Delta G^\circ\) value in the box provided, in units of \(\text{kJ/mol}\).

**Box: [ _____ ] kJ/mol**

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