Question

Which of the following cannot be oxidized by \(\mathrm{H}_{2} \mathrm{O}_{2} ?\) (a) \(\mathrm{O}_{3}\) (b) \(\mathrm{PbS}\) (c) \(\mathrm{Na}_{2} \mathrm{SO}_{3}\) (d) \(\mathrm{Kl}+\mathrm{HCl}\)

Short answer

(a) \( \mathrm{O}_3 \) cannot be oxidized by \( \mathrm{H}_2\mathrm{O}_2 \).

Step-by-step solution

Analyze the Function of H2O2

Hydrogen peroxide, \( \mathrm{H}_2\mathrm{O}_2 \), acts as both an oxidizing and reducing agent, depending on the substances it interacts with. In this context, it is used as an oxidizing agent. We need to determine which compound cannot be oxidized by \( \mathrm{H}_2\mathrm{O}_2 \).

Evaluate the Oxidation Potential of Each Compound

To determine the answer, consider the oxidation states of the compounds: - \( \mathrm{O}_3 \) (ozone) is already in a high oxidation state (+2 for oxygen in \( \mathrm{O}_3 \)), making it less likely to be further oxidized.- \( \mathrm{PbS} \) can be oxidized to \( \mathrm{PbSO}_4 \).- \( \mathrm{Na}_2 \mathrm{SO}_3 \) can be oxidized to \( \mathrm{Na}_2 \mathrm{SO}_4 \).- \( \mathrm{Kl} + \mathrm{HCl} \) can react, and \( \mathrm{I}^- \) can be oxidized to \( \mathrm{I}_2 \).

Determine Which Cannot Be Oxidized

Compare the given options: - \( \mathrm{O}_3 \) is a strong oxidizer itself and unlikely to undergo oxidation by \( \mathrm{H}_2\mathrm{O}_2 \).- The other compounds, \( \mathrm{PbS}, \) \( \mathrm{Na}_2 \mathrm{SO}_3, \) and \( \mathrm{Kl} \), can each undergo oxidation because they are in forms that can be increased in oxidation state.

Key concepts

Oxidizing and Reducing Agents

An oxidizing agent is a substance that causes another substance to lose electrons, thus being reduced itself in the process. A reducing agent, on the other hand, causes another substance to gain electrons and is oxidized in the process. In the realm of redox chemistry, recognizing these agents is vital for predicting the outcomes of chemical reactions.

• Oxidizing agents accept electrons from other substances. Consequently, they become reduced. Typical oxidizing agents include oxygen, chlorine, and hydrogen peroxide.
• Reducing agents donate electrons to other substances and become oxidized as a result. Examples include metals like zinc or sodium.

Understanding the role of oxidizing and reducing agents helps in deciphering which substance will lose electrons and which will gain them during a chemical reaction. For instance, hydrogen peroxide (\( \mathrm{H}_2\mathrm{O}_2 \)) can act as both an oxidizing agent and a reducing agent, depending on its environment.

Oxidation States

Oxidation states, also known as oxidation numbers, are used to describe the degree of oxidation of an atom in a compound. These numbers vary based on the bonding between atoms and the type of elements involved. Identifying oxidation states is crucial in determining the direction of electron transfer in redox reactions.

• The oxidation state of a pure element is always zero. For instance, in molecules like \( \mathrm{O}_2 \) and \( \mathrm{N}_2 \), both elements have oxidation states of zero.
• In compound ions like \( \mathrm{H}_2\mathrm{O}_2 \), the sum of the oxidation states equals the charge of the molecule. Here, oxygen typically has an oxidation state of -2, but in \( \mathrm{H}_2\mathrm{O}_2 \), each oxygen has an oxidation state of -1.

By determining the oxidation states of the elements within a molecule, you can better understand their chemical behavior, including which species are likely to gain or lose electrons during reactions. Considerations of oxidation states are essential when evaluating compounds in reactions, such as knowing \( \mathrm{O}_3 \) is already highly oxidized and less likely to undergo further oxidation.

Hydrogen Peroxide Reactions

Hydrogen peroxide (\( \mathrm{H}_2\mathrm{O}_2 \)) is a versatile chemical compound frequently used in various applications, from disinfection to chemical synthesis. Its ability to act as both an oxidizing and a reducing agent makes it unique in redox chemistry.

• As an oxidizing agent, \( \mathrm{H}_2\mathrm{O}_2 \) can convert sulfides (\( \mathrm{PbS} \)) into sulfates (\( \mathrm{PbSO}_4 \)), and sulfites (\( \mathrm{Na}_2 \mathrm{SO}_3 \)) into sulfates (\( \mathrm{Na}_2 \mathrm{SO}_4 \)).
• It can also react with iodides, causing the conversion of \( \mathrm{I}^- \) to \( \mathrm{I}_2 \). This demonstrates its strong oxidizing power in such reactions.

However, not all substances react with \( \mathrm{H}_2\mathrm{O}_2 \). For instance, ozone (\( \mathrm{O}_3 \)) is already an oxidizing powerhouse, making it less susceptible to oxidation by hydrogen peroxide. Understanding these interaction paradigms showcases \( \mathrm{H}_2\mathrm{O}_2 \)'s dual-functionality, which is pivotal for many oxidation and reduction processes.