Problem 64 - General Chemistry

Question

In the following reactions, label the oxidizing agent and the reducing agent. a. \(\mathrm{Fe}_{2} \mathrm{O}_{3}(s)+3 \mathrm{CO}(g) \longrightarrow 2 \mathrm{Fe}(s)+3 \mathrm{CO}_{2}(g)\) b. \(\mathrm{PbS}(s)+4 \mathrm{H}_{2} \mathrm{O}_{2}(a q) \longrightarrow \mathrm{PbSO}_{4}(s)+4 \mathrm{H}_{2} \mathrm{O}(l)\)

Step-by-Step Solution

Answer

a. Oxidizing agent: \( \mathrm{Fe}_{2} \mathrm{O}_{3} \), Reducing agent: \( \mathrm{CO} \); b. Oxidizing agent: \( \mathrm{H}_{2} \mathrm{O}_{2} \), Reducing agent: \( \mathrm{PbS} \).

Step 1: Identify Changes in Oxidation States for (a)

In the reaction: \( \mathrm{Fe}_{2} \mathrm{O}_{3}(s)+3 \mathrm{CO}(g) \longrightarrow 2 \mathrm{Fe}(s)+3 \mathrm{CO}_{2}(g) \), identify the oxidation states. Iron (Fe) in \( \mathrm{Fe}_{2} \mathrm{O}_{3} \) is +3 and is reduced to 0 in \( \mathrm{Fe} \). Carbon (C) in CO is +2 and is oxidized to +4 in \( \mathrm{CO}_{2} \).

Step 2: Determine the Oxidizing and Reducing Agents for (a)

Since iron is reduced, \( \mathrm{Fe}_{2} \mathrm{O}_{3} \) is the oxidizing agent. Since carbon in CO is oxidized, \( \mathrm{CO} \) is the reducing agent.

Step 3: Identify Changes in Oxidation States for (b)

In the reaction: \( \mathrm{PbS}(s)+4 \mathrm{H}_{2} \mathrm{O}_{2}(a q) \longrightarrow \mathrm{PbSO}_{4}(s)+4 \mathrm{H}_{2} \mathrm{O}(l) \), the oxidation state of sulfur in \( \mathrm{PbS} \) goes from -2 to +6 in \( \mathrm{PbSO}_{4} \). Oxygen in \( \mathrm{H}_{2} \mathrm{O}_{2} \) goes from -1 to -2 in \( \mathrm{H}_{2} \mathrm{O} \).

Step 4: Determine the Oxidizing and Reducing Agents for (b)

Sulfur in \( \mathrm{PbS} \) is oxidized, so \( \mathrm{PbS} \) is the reducing agent. Oxygen in \( \mathrm{H}_{2} \mathrm{O}_{2} \) is reduced, making \( \mathrm{H}_{2} \mathrm{O}_{2} \) the oxidizing agent.

Key Concepts

Oxidizing Agent

An oxidizing agent is a substance that causes another substance to lose electrons during a chemical reaction. It does this by accepting electrons from the other substance. In the process, the oxidizing agent itself is reduced. In the first reaction provided:

• The oxidizing agent is Fe₂O₃, which means it gains electrons in the process, thus reducing the oxidation state of iron (Fe) from +3 to 0.

For example, in our step-by-step solution, oxygen is utilized to help reduce iron. It indicates that as the Fe in Fe₂O₃ undergoes reduction (gaining electrons), it effectively loses its positive oxidation state. Similarly, in the second reaction:

• The oxidizing agent is H₂O₂. Here, the oxygen in H₂O₂ gains electrons, which reduces its oxidation state from -1 to -2.

Thus, by accepting electrons, oxidizing agents are crucial players in oxidation-reduction reactions, powering the reduction of other elements.

Reducing Agent

A reducing agent is the opposite of an oxidizing agent. It donates electrons to another substance, causing the other substance to be reduced. As the reducing agent relinquishes electrons, it itself is oxidized, which elevates its oxidation state. In the first solution step:

• The reducing agent is CO. The carbon (C) in CO gives up electrons, changing its oxidation state from +2 to +4 when it forms CO₂.

This electron donation propels the process of oxidation for the carbon, highlighting CO's role as a reducing agent. In the second chemical reaction:

• PbS acts as the reducing agent. The sulfur (S) in PbS loses electrons, leading it to be oxidized from an oxidation state of -2 to +6.

This behavior exemplifies how reducing agents facilitate the reduction of other elements by offering their own electrons.

Oxidation States

Oxidation states, or oxidation numbers, are hypothetical charges that atoms would have if all bonds to atoms of different elements were fully ionic. They help in determining the electron transfer during a chemical reaction. When analyzing the first reaction:

• We observe the oxidation state change for Fe from +3 in Fe₂O₃ to 0 in pure Fe, indicating reduction.
• Similarly, for carbon in CO, its state moves from +2 to +4 in CO₂, illustrating oxidation.

By identifying these changes, we clarify which processes occurred. For the second reaction:

• The oxidation state of sulfur in PbS moves from -2 to +6, signifying its oxidation.
• Oxygen in H₂O₂, on the other hand, shifts from -1 to -2, pointing out its reduction.

Understanding oxidation states is essential, allowing chemists to track electron transfer and identify the roles of different agents in a reaction.