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Balance the following redox reaction in acidic solution. 2+ CI (aq) + MnO4(aq) Cl₂(g) + Mn (aq) -

Balance the following redox reaction in acidic solution. 2+ CI (aq) + MnO4(aq) Cl₂(g) + Mn (aq) -

Chemistry
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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**Balancing a Redox Reaction in Acidic Solution** The task is to balance the following redox reaction in an acidic solution: \[ \text{Cl}^- (aq) + \text{MnO}_4^- (aq) \rightarrow \text{Cl}_2 (g) + \text{Mn}^{2+} (aq) \] **Current Equation:** \[ \text{Cl}^- (aq) + \text{MnO}_4^- (aq) \rightarrow \text{Cl}_2 (g) + \text{Mn}^{2+} (aq) \] **Explanation of Diagram:** The diagram contains a tool that appears to assist in balancing the equation, specifically for the electrons in the half-reactions. This tool includes symbols for square boxes and an electron symbol (\(e^-\)). It is likely used to track and balance electrons between the oxidation and reduction processes. - Arrows between boxes suggest indicating electron transfer. - It aids in ensuring that both mass and charge are balanced, a crucial step for redox reactions. To fully balance the redox equation, follow these general steps: 1. Separate the reactions into oxidation and reduction half-reactions. 2. Balance each half-reaction for atoms and charge, adding \(H_2O\), \(H^+\), and electrons (\(e^-\)) as needed. 3. Combine the half-reactions, ensuring electrons are balanced and simplify the equation. This approach ensures completeness and accuracy, following the redox balancing rules in acidic conditions.
Transcribed Image Text:**Balancing a Redox Reaction in Acidic Solution** The task is to balance the following redox reaction in an acidic solution: \[ \text{Cl}^- (aq) + \text{MnO}_4^- (aq) \rightarrow \text{Cl}_2 (g) + \text{Mn}^{2+} (aq) \] **Current Equation:** \[ \text{Cl}^- (aq) + \text{MnO}_4^- (aq) \rightarrow \text{Cl}_2 (g) + \text{Mn}^{2+} (aq) \] **Explanation of Diagram:** The diagram contains a tool that appears to assist in balancing the equation, specifically for the electrons in the half-reactions. This tool includes symbols for square boxes and an electron symbol (\(e^-\)). It is likely used to track and balance electrons between the oxidation and reduction processes. - Arrows between boxes suggest indicating electron transfer. - It aids in ensuring that both mass and charge are balanced, a crucial step for redox reactions. To fully balance the redox equation, follow these general steps: 1. Separate the reactions into oxidation and reduction half-reactions. 2. Balance each half-reaction for atoms and charge, adding \(H_2O\), \(H^+\), and electrons (\(e^-\)) as needed. 3. Combine the half-reactions, ensuring electrons are balanced and simplify the equation. This approach ensures completeness and accuracy, following the redox balancing rules in acidic conditions.
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