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) \(\mathrm{KI}+\mathrm{HCl}\)

Short answer

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

Step-by-step solution

Understanding the Reducer

Hydrogen peroxide (\( \mathrm{H}_2\mathrm{O}_2 \)) can act as both an oxidizing and reducing agent. However, in this question, we focus on its oxidizing behavior. It oxidizes substances that have a higher oxidation potential than it does.

Analyzing Options: Ozone (\(\mathrm{O}_3\))

Ozone (\(\mathrm{O}_3\)) is a strong oxidizing agent, even stronger than hydrogen peroxide. Thus, \(\mathrm{H}_2\mathrm{O}_2\) cannot oxidize \(\mathrm{O}_3\) because it's already at its highest oxidation state, +6.

Analyzing Options: Lead (II) Sulfide (\(\mathrm{PbS}\))

Lead (II) sulfide (\( \mathrm{PbS} \)) can be oxidized by hydrogen peroxide to form lead sulfate (\(\mathrm{PbSO}_4\)), as sulfur is oxidized from -2 in \(\mathrm{PbS}\) to +6 in \(\mathrm{PbSO}_4\).

Analyzing Options: Sodium Sulfite (\(\mathrm{Na}_2\mathrm{SO}_3\))

Sodium sulfite(\( \mathrm{Na}_2 \mathrm{SO}_3 \)) can be oxidized by \(\mathrm{H}_2 \mathrm{O}_2\) to sodium sulfate (\( \mathrm{Na}_2 \mathrm{SO}_4\)), as sulfur goes from a +4 to a +6 oxidation state.

Analyzing Options: Potassium Iodide with Hydrochloric Acid (\(\mathrm{KI} + \mathrm{HCl}\))

The iodide ion in \(\mathrm{KI}\) can be oxidized by \(\mathrm{H}_2\mathrm{O}_2\) in an acidic medium to iodine (\(\mathrm{I}_2\)), going from a -1 to 0 oxidation state.

Key concepts

Oxidizing Agents

In chemistry, oxidizing agents are substances that can accept electrons from other substances in a reaction. This means they help oxidize other compounds by taking away electrons. A classic example of an oxidizing agent is hydrogen peroxide (\( \mathrm{H}_2\mathrm{O}_2 \)). It can both gain electrons and cause other substances to release them.
Hydrogen peroxide acts as an oxidizing agent because it has a potential to gain electrons, going from its normal state to form water (\( \mathrm{H}_2\mathrm{O} \)) and releasing oxygen. Not all substances can be oxidized by the same agent—some substances, like ozone (\( \mathrm{O}_3 \)), are already at high oxidation states and thus resist further oxidation. In fact, ozone is a stronger oxidizing agent than hydrogen peroxide.
When examining the ability of hydrogen peroxide to act as an oxidizing agent in different reactions, we consider the oxidation potential—a measure of how easily a substance can acquire electrons. Substances with a lower oxidation maximum, like lead (II) sulfide (\( \mathrm{PbS} \)) or sodium sulfite (\( \mathrm{Na}_2\mathrm{SO}_3 \)), can be oxidized by hydrogen peroxide. By contrast, substances already stabilized in their highest oxidation state, such as ozone, remain unaffected.

Oxidation States

Oxidation state, sometimes known as oxidation number, is a concept used to describe the distribution of electrons among atoms in a chemical compound. It tells us how many electrons an atom gains or loses during a reaction, which helps in understanding redox reactions.
In the context of the question, oxidation states are important to determine what hydrogen peroxide can oxidize. For instance, sulfur in lead (II) sulfide (\( \mathrm{PbS} \)) has an oxidation state of -2, and it changes to +6 when it forms lead sulfate (\( \mathrm{PbSO}_4 \)). Similarly, sulfur in sodium sulfite (\( \mathrm{Na}_2\mathrm{SO}_3 \)) goes from +4 to +6 when oxidized to sodium sulfate (\( \mathrm{Na}_2\mathrm{SO}_4 \)).
By understanding oxidation states, you can also see why hydrogen peroxide cannot oxidize ozone (\( \mathrm{O}_3 \)). Ozone is already in a very high oxidation state, making further oxidation impossible without an exceedingly strong oxidizing agent. Thus, determining the oxidation state offers a clear picture of a chemical reaction's electron exchange process.

Redox Reactions

Redox reactions are processes in chemistry where oxidation and reduction occur simultaneously. In these reactions, one substance gains electrons (is reduced), while another loses electrons (is oxidized). The term 'redox' itself is a combination of "reduction" and "oxidation."
When hydrogen peroxide is involved in a redox reaction, it can act as either an oxidizing or reducing agent. In the context of the exercise, we focus on oxidation. For example, when hydrogen peroxide oxidizes sodium sulfite (\( \mathrm{Na}_2\mathrm{SO}_3 \)) to sodium sulfate (\( \mathrm{Na}_2\mathrm{SO}_4 \)), sulfur atoms lose electrons and their oxidation state increases from +4 to +6.
Redox reactions are essential as they depict not only the electron transfer but also the change in oxidation states of the substances involved. They occur everywhere, from biological systems within our bodies to industrial processes and environmental cycles. Understanding these reactions helps in decoding how compounds interact, transform, and recycle in nature and technology. Recognizing how substances like hydrogen peroxide are involved in redox reactions aids in comprehending their role in various significant chemical processes.