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Chapter 22: Problem 83

Hydrogen peroxide is capable of oxidizing (a) hydrazine to \(\mathrm{N}_{2}\) and \(\mathrm{H}_{2} \mathrm{O},(\mathbf{b}) \mathrm{SO}_{2}\) to \(\mathrm{SO}_{4}^{2-},(\mathbf{c}) \mathrm{NO}_{2}^{-}\) to \(\mathrm{NO}_{3}^{-},(\mathbf{d}) \mathrm{H}_{2} \mathrm{~S}(g)\) to \(\mathrm{S}(s),(\mathbf{e}) \mathrm{Fe}^{2+}\) to \(\mathrm{Fe}^{3+}\). Write a balanced net ionic equation for each of these redox reactions.

Short Answer

Expert verified
The short answer: 1. Reaction with Hydrazine: \[\mathrm{N}_{2}\mathrm{H}_{4} + 2\mathrm{H}_{2}\mathrm{O}_2 \rightarrow \mathrm{N}_2 + 4\mathrm{H}^{+} + 2\mathrm{H}_{2}\mathrm{O} + \mathrm{O}_{2}\] 2. Reaction with \(\mathrm{SO}_2\): \[\mathrm{SO}_2 + 2\mathrm{H_2O} + 2\mathrm{H}_{2}\mathrm{O}_{2} \rightarrow \mathrm{SO_{4}^{2-}} + 4\mathrm{H}^{+} + 2\mathrm{H}_{2}\mathrm{O} + \mathrm{O}_{2}\] 3-5. Reactions with \(\mathrm{NO}_{2}^-\), \(\mathrm{H}_{2} \mathrm{~S}(g)\), and \(\mathrm{Fe}^{2+}\) can be calculated similarly following the same steps.

Step by step solution

01

Write half-reactions

First, identify the oxidation state changes of the given elements in hydrazine to form a balanced half-reaction for oxidation and reduction processes. Hydrazine (\(\mathrm{N}_{2}\mathrm{H}_{4}\)) oxidizes to nitrogen gas (\(\mathrm{N}_{2}\)) and water (\(\mathrm{H}_{2}\mathrm{O}\)): Oxidation: \(\mathrm{N}_{2}\mathrm{H}_{4} \rightarrow \mathrm{N}_2\) Reduction: \(\mathrm{H}_{2}\mathrm{O}_2 \rightarrow \mathrm{H}_{2}\mathrm{O}\)
02

Balance half-reactions

Balance the half-reactions by adjusting the coefficients and adding electrons (\(e^-\)) as needed: Oxidation: \(\mathrm{N}_{2}\mathrm{H}_{4} \rightarrow \mathrm{N}_2 + 4\mathrm{H}^{+} + 4 e^{-}\) Reduction: \(2\mathrm{H}_{2}\mathrm{O}_2 + 4 e^{-} \rightarrow 2\mathrm{H}_{2}\mathrm{O} + \mathrm{O}_{2}\)
03

Combine balanced half-reactions

Combine the balanced half-reactions to obtain the balanced net ionic equation: \(\mathrm{N}_{2}\mathrm{H}_{4} + 2\mathrm{H}_{2}\mathrm{O}_2 \rightarrow \mathrm{N}_2 + 4\mathrm{H}^{+} + 2\mathrm{H}_{2}\mathrm{O} + \mathrm{O}_{2}\) ##Reaction with \(\mathrm{SO}_2\)##
04

Write half-reactions

Oxidation: \(\mathrm{SO}_2 \rightarrow \mathrm{SO_{4}^{2-}}\) Reduction: \(\mathrm{H}_{2}\mathrm{O}_{2} \rightarrow \mathrm{H}_{2}\mathrm{O}\)
05

Balance half-reactions

Oxidation: \(\mathrm{SO}_2 + 2\mathrm{H_2O} \rightarrow \mathrm{SO_{4}^{2-}} + 4\mathrm{H}^{+} + 2e^{-}\) Reduction: \(2\mathrm{H}_{2}\mathrm{O}_{2} + 2e^{-} \rightarrow 2\mathrm{H}_{2}\mathrm{O} + \mathrm{O}_{2}\)
06

Combine balanced half-reactions

\(\mathrm{SO}_2 + 2\mathrm{H_2O} + 2\mathrm{H}_{2}\mathrm{O}_{2} \rightarrow \mathrm{SO_{4}^{2-}} + 4\mathrm{H}^{+} + 2\mathrm{H}_{2}\mathrm{O} + \mathrm{O}_{2}\) To be continued for the other 3 reactions (c, d, e).

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