Problem 3 - Chemistry Matter and Change

Question

Identify the oxidizing agent and the reducing agent in the following equation. Explain your answer. $$ \mathrm{Fe}(\mathrm{s})+\mathrm{Ag}+(\mathrm{aq}) \rightarrow \mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{Ag}(\mathrm{s}) $$

Step-by-Step Solution

Answer

Fe is the reducing agent, and Ag⁺ is the oxidizing agent.

Step 1: Identify the Changes in Oxidation State

First, identify the elements that undergo a change in oxidation state. In the given reaction, Fe (iron) changes from e ightarrow ext{Fe}^{2+}e→Fe2+ and Ag (silver) goes from g^+ ightarrow ext{Ag} g+→Ag.

Step 2: Determine Oxidation and Reduction

Oxidation involves an increase in oxidation state, while reduction involves a decrease. For e ightarrow ext{Fe}^{2+}e→Fe2+, the oxidation state goes from 0 to +2, so Fe is oxidized. For g^+ ightarrow ext{Ag} g+→Ag, the oxidation state goes from +1 to 0, so Ag is reduced.

Step 3: Assign Oxidizing and Reducing Agents

The substance that is oxidized (Fe) acts as the reducing agent because it donates electrons. The substance that is reduced (Ag⁺) acts as the oxidizing agent because it accepts electrons.

Key Concepts

Oxidizing Agent

In the world of chemistry, an oxidizing agent is a substance that gains electrons in a reaction. This means it helps another substance lose electrons. When a substance loses electrons, we say it is being oxidized, and the one that gains those electrons is called the oxidizing agent.

In our reaction:

• The silver ion, \( ext{Ag}^+ \), accepts electrons, going from a +1 charge in its ionic state to a neutral metallic form, \( ext{Ag} \).
• By accepting these electrons, it causes the oxidation of iron.

The silver ion is, therefore, the oxidizing agent in the reaction because it gets reduced by gaining electrons.

Reducing Agent

The reducing agent is the opposite of the oxidizing agent. It donates electrons to another substance, causing it to be reduced. When the reducing agent gives away its electrons, it undergoes oxidation itself.

In our reaction:

• Iron, \( ext{Fe} \), starts with an oxidation state of 0 in its metallic form and loses electrons to become \( ext{Fe}^{2+} \) in aqueous form.
• This loss of electrons is what categorizes \( ext{Fe} \) as being oxidized.
• Since \( ext{Fe} \) donates electrons to \( ext{Ag}^+ \), it acts as the reducing agent.

Thus, the substance that is oxidized (iron in this case) and gives electrons away is termed as the reducing agent.

Oxidation State

The oxidation state is essentially a number that tells us how many electrons an atom is donating or accepting during a reaction. It is like an atom's "accounting system" for electrons.

In redox reactions, changes in oxidation states reveal which substances are oxidized and reduced:

• For iron, the oxidation state changes from 0 to +2, indicating that it loses two electrons.
• For silver, the oxidation state changes from +1 to 0, indicating it gains one electron.

Careful tracking of these changes in oxidation states helps identify the flow of electrons in the chemical reaction. Understanding this concept is crucial for mastering oxidation-reduction chemistry.

Redox Reaction

Redox is short for reduction-oxidation, a type of chemical reaction where the oxidation states of atoms are changed. In these reactions, one atom or molecule loses electrons while another gains them.

The essential parts of redox reactions include:

• Oxidation: Increase in oxidation state due to loss of electrons.
• Reduction: Decrease in oxidation state due to gain of electrons.

In the given reaction, iron undergoes oxidation by losing electrons, while silver undergoes reduction by gaining electrons.

Redox reactions are not only a fundamental concept in chemistry but are also vital in various industrial processes and biological systems, like respiration and photosynthesis.