Question

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

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

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: \(\mathrm{H}_{2}\mathrm{O}_2 \rightarrow \mathrm{H}_{2}\mathrm{O}\)

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: \(2\mathrm{H}_{2}\mathrm{O}_2 + 4 e^{-} \rightarrow 2\mathrm{H}_{2}\mathrm{O} + \mathrm{O}_{2}\)

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 \(\mathrm{SO}_2\)##

Write half-reactions

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

Balance half-reactions

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

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).

Key concepts

Half-Reactions

In redox reactions, also known as oxidation-reduction reactions, we often break them down into two separate processes called half-reactions. This makes it easier to understand the transfer of electrons between substances. A half-reaction shows either the oxidation or reduction part of the reaction.
To start, identify which species is losing electrons (oxidation) and which is gaining electrons (reduction). By splitting the processes, each part focuses on either the electron donor or the electron acceptor.

• Oxidation Half-Reaction: This process involves the loss of electrons. The substance releasing electrons undergoes oxidation.
• Reduction Half-Reaction: This involves the gain of electrons. The substance accepting electrons is reduced.

For instance, in the oxidation of hydrazine (\( ext{N}_2 ext{H}_4 \) to \( ext{N}_2 \)), we write one half-reaction highlighting that transformation and another detailing the reduction of \( ext{H}_2 ext{O}_2 \) to \( ext{H}_2 ext{O} \).
Splitting these processes helps simplify the balancing of equations, allowing students to track electron changes more efficiently.

Oxidation-Reduction

Redox reactions are all about the movement of electrons from one reactant to another. The key terms here are oxidation and reduction. They always occur together because the electrons lost by one species must be gained by another.

• Oxidation: It's when a molecule, atom, or ion loses electrons. For example, in the hydrazine reaction, \( ext{N}_2 ext{H}_4 \) loses electrons to become \( ext{N}_2 \).
• Reduction: This is when a molecule gains electrons. In our reactions, \( ext{H}_2 ext{O}_2 \) is reduced to \( ext{H}_2 ext{O} \) by gaining electrons.

To determine which substance is oxidized or reduced, look at the changes in oxidation states of the atoms in the compounds involved. An increase in oxidation state indicates oxidation, while a decrease indicates reduction.
Recognizing these changes is crucial for writing half-reactions and ultimately balancing the equation. Understanding these concepts is essential for applying them in various chemical processes and reactions.

Balancing Chemical Equations

Balancing chemical equations ensures that the number of atoms and the charge are the same on both sides of the equation. For redox reactions, it's essential to balance not just the atoms, but also the charge, due to the transfer of electrons.
Here's a simple way to balance them:

• Write the unbalanced equation: Start with half-reactions that show either oxidation or reduction.
• Balance each half-reaction: Balance elements other than \( ext{H} \) and \( ext{O} \). Next, balance \( ext{O} \) using \( ext{H}_2 ext{O} \). Lastly, balance \( ext{H} \) using \( ext{H}^+ \) ions and add electrons (\( e^- \)) to balance the charge.
• Combine the half-reactions: After balancing individual half-reactions, sum them up ensuring that the number of electrons lost is equal to the number of electrons gained. These electrons will cancel out.

For example, when balancing the redox reaction between \( ext{SO}_2 \) and \( ext{H}_2 ext{O}_2 \), we make sure all sulfur and oxygen atoms are balanced along with the charges, resulting in a complete, balanced net ionic equation.
This process verifies the law of conservation of mass and charge, ensuring that everything is accounted for, which is fundamental in chemical reactions.