Problem 16 - Introductory Chemistry Atoms First

Question

For each redox reaction, indicate which substance is the oxidizing agent and which is the reducing agent. \(4 \mathrm{Fe}+3 \mathrm{O}_{2} \rightarrow 2 \mathrm{Fe}_{2} \mathrm{O}_{3}\)

Step-by-Step Solution

Answer

In the given redox reaction \(4 \mathrm{Fe} + 3 \mathrm{O}_2 \rightarrow 2 \mathrm{Fe}_2 \mathrm{O}_3\), the oxidizing agent is O2, and the reducing agent is Fe.

Step 1: Write down the given reaction

The given redox reaction is: \(4 \mathrm{Fe} + 3 \mathrm{O}_2 \rightarrow 2 \mathrm{Fe}_2 \mathrm{O}_3\)

Step 2: Determine the oxidation states of the elements before and after the reaction

Before the reaction, the oxidation state of each element is: - Fe: 0 (as it is in its pure elemental form) - O: 0 (as it is in its pure elemental form) After the reaction, the oxidation state of each element in the Fe2O3 compound is: - Fe: +3 (because it is bound to 3 oxygen atoms, each with an oxidation state of -2) - O: -2 (typical oxidation state of oxygen in compounds)

Step 3: Identify the element being oxidized and the element being reduced

Comparing the oxidation states before and after the reaction, we can see that: - Fe has increased its oxidation state from 0 to +3, which means it has lost electrons and is being oxidized. - O has decreased its oxidation state from 0 to -2, which means it has gained electrons and is being reduced.

Step 4: Identify the oxidizing agent and the reducing agent

The oxidizing agent is the substance that causes the oxidation of another substance (i.e., it gains electrons from the other substance). In this case, it is the O2 that causes the oxidation of Fe, so O2 is the oxidizing agent. The reducing agent is the substance that causes the reduction of another substance (i.e., it loses electrons to the other substance). In this case, it is the Fe that causes the reduction of O2, so Fe is the reducing agent.

Step 5: Conclusion

In the given redox reaction, the oxidizing agent is O2, and the reducing agent is Fe.

Key Concepts

Oxidizing Agent

In a redox reaction, the oxidizing agent is the substance that accepts electrons. By gaining electrons, it causes another substance to lose electrons. This means it's facilitating oxidation in another element or compound. It plays a crucial role in driving the redox reaction process forward.
For an easy way to remember, think of the oxidizing agent as the electron tip jar; it collects electrons from others.

• In the example of the redox reaction with iron (Fe) and oxygen (\(\mathrm{O}_2\)):
  \(\mathrm{O}_2\) gains electrons from \(\mathrm{Fe}\), increasing its oxidation state from 0 to -2.
• Thus, \(\mathrm{O}_2\) acts as the oxidizing agent since it facilitates the oxidation of \(\mathrm{Fe}\).

Remember, an oxidizing agent gets reduced itself as it helps in the oxidation of another substance.

Reducing Agent

The reducing agent in a redox reaction is the molecule, ion, or atom that donates electrons to another species. This means it encourages another substance to gain electrons, which is also known as reduction. As it gives away electrons, it "reduces" the other substance but is oxidized itself in the process.
Just imagine the reducing agent as the donor at a blood drive, but here it donates electrons instead.

• Taking the redox reaction between iron (\(\mathrm{Fe}\)) and oxygen (\(\mathrm{O}_2\)), iron loses electrons as its oxidation state increases from 0 to +3.
• This makes \(\mathrm{Fe}\) the reducing agent because it provides electrons to oxygen, enabling its reduction.

Reduce means to donate here, making the reducing agent essential to such chemical transformations.

Oxidation States

Understanding oxidation states is central to grasping redox reactions as they help track the transfer of electrons. An oxidation state (or oxidation number) identifies the degree of oxidation of an atom in a substance. It reflects how many electrons an atom has gained, lost, or appears to use in bonding.
Visualize oxidation states like scoring points in a game. Positive points (oxidation numbers) indicate electrons lost, while negative ones suggest electrons gained.

• In a given reaction: before, both \(\mathrm{Fe}\) and \(\mathrm{O}_2\) elements have an oxidation state of 0 due to their pure elemental form.
• When they react to form \(\mathrm{Fe}_2\mathrm{O}_3\), \(\mathrm{Fe}\)'s state changes to +3 and \(\mathrm{O}\)'s to -2.

This shows that iron is oxidized (loses electrons) and oxygen is reduced (gains electrons), emphasizing the significance of oxidation states in detecting which component undergoes oxidation or reduction in redox reactions.