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) $$ \mathrm{PbS}(s)+4 \mathrm{H}_{2} \mathrm{O}_{2}(a q) \longrightarrow \mathrm{PbSO}_{4}(s)+4 \mathrm{H}_{2} \mathrm{O}(l) $$

Short answer

(a) Oxidizing agent: Fe₂O₃, Reducing agent: CO; (b) Oxidizing agent: H₂O₂, Reducing agent: PbS.

Step-by-step solution

Identify Oxidation and Reduction

To find the oxidizing and reducing agents, we must first identify which species are oxidized and reduced. Oxidation is the loss of electrons, while reduction is the gain of electrons. In (a), Fe goes from an oxidation state of +3 in Fe₂O₃ to 0 in Fe, meaning iron is reduced. In (b), S in PbS goes from -2 to +6 in PbSO₄, indicating sulfur is oxidized.

Assign Oxidizing and Reducing Agents in (a)

In reaction (a), Fe₂O₃ is reduced to Fe, meaning Fe₂O₃ acts as the oxidizing agent. CO is oxidized to CO₂, indicating CO is the reducing agent.

Assign Oxidizing and Reducing Agents in (b)

In reaction (b), H₂O₂ is reduced to H₂O, so H₂O₂ is the oxidizing agent. PbS is oxidized to PbSO₄, thereby acting as the reducing agent.

Key concepts

Oxidizing Agent

An oxidizing agent is a substance that causes another substance to lose electrons in a chemical reaction. This means it itself gets reduced by gaining electrons. It’s like a chemist's version of a "good cop" who takes in the bad electrons! In reaction (a) of the exercise, we identified that iron in \( \text{Fe}_2\text{O}_3 \) was reduced. So, \( \text{Fe}_2\text{O}_3 \) acts as the oxidizing agent. In reaction (b), \( \text{H}_2\text{O}_2 \) was reduced to \( \text{H}_2\text{O} \), making \( \text{H}_2\text{O}_2 \) the oxidizing agent.
Some key characteristics of oxidizing agents include:

• They are often substances that contain oxygen or another electronegative element.
• They tend to have high oxidation states.
• They are essential in combustion, rusting, and many industrial processes.

When identifying oxidizing agents, look for substances that gain electrons during the reaction, as reflected in their decreased oxidation states.

Reducing Agent

The reducing agent is the substance that donates electrons to another substance, thus causing reduction by reducing the oxidation state of the other substance. It itself undergoes oxidation. You can think of it as a "generous friend" giving away its extra electrons!

In the reactions we're looking at, the reducing agents swap electrons with the oxidizing agents. For example, in reaction (a), carbon monoxide (\( \text{CO} \)) is oxidized to carbon dioxide (\( \text{CO}_2 \)). This means \( \text{CO} \) is the reducing agent because it provides electrons to \( \text{Fe}_2\text{O}_3 \).

In reaction (b), lead sulfide (\( \text{PbS} \)) is oxidized to lead sulfate (\( \text{PbSO}_4 \)), showing that \( \text{PbS} \) is the reducing agent.

Some important points about reducing agents include:

• They have high energy electrons available for donation.
• They often contain elements in low oxidation states.
• They play key roles in processes like metal extraction and biological metabolism.

Learning to spot reducing agents involves locating substances that lose electrons, signifying an increase in their oxidation states.

Oxidation States

Oxidation states, also known as oxidation numbers, are a clever way chemists track how many electrons are lost or gained by atoms in a reaction. They provide an imaginary charge based on a set of rules, not unlike a set scorecard.

Understanding oxidation states is crucial in correctly assigning oxidizing and reducing agents. Let's break down what these numbers mean:

• An increase in oxidation state means the species has lost electrons and is oxidized.
• A decrease in oxidation state means the species has gained electrons, undergoing reduction.

In step 1 of our solution, we observed:
1. In reaction (a), the oxidation state of iron changes from +3 (in \( \text{Fe}_2\text{O}_3 \)) to 0 (in free \( \text{Fe} \)), indicating reduction.
2. In reaction (b), sulfur's oxidation state changes from -2 (in \( \text{PbS} \)) to +6 (in \( \text{PbSO}_4 \)), showing oxidation.

Some rules to remember when determining oxidation states are:

• The oxidation state of a free element is always zero.
• For monoatomic ions, the oxidation state equals the ion's charge.
• Oxygen usually has an oxidation state of -2, while hydrogen is usually +1.

Keep practicing to become skilled at determining oxidation states, as they are a fundamental concept in understanding redox reactions.