Question

What element behaves as the oxidizing agent in the following equation and what element behaves as the reducing agent? $$ \mathrm{Sn}^{2+}+2 \mathrm{Ag} \longrightarrow \mathrm{Sn}+2 \mathrm{Ag}^{+} $$

Short answer

The oxidizing agent is \( \mathrm{Sn}^{2+} \), and the reducing agent is \( \mathrm{Ag} \).

Step-by-step solution

Identify Oxidation States (Reactants)

First, determine the oxidation states of each element in the reactants. For \( \mathrm{Sn}^{2+} \), the oxidation state is +2. For \( \mathrm{Ag} \), the elemental form, the oxidation state is 0.

Identify Oxidation States (Products)

Determine the oxidation states of each element in the products. \( \mathrm{Sn} \) in elemental form has an oxidation state of 0. For \( \mathrm{Ag}^{+} \), the oxidation state is +1.

Determine Changes in Oxidation State

Examine the changes in oxidation state from reactants to products. \( \mathrm{Sn}^{2+} \) changes from +2 to 0 (reduction), while \( \mathrm{Ag} \) changes from 0 to +1 (oxidation).

Identify the Reducing and Oxidizing Agents

The element that gets oxidized serves as the reducing agent, while the element that gets reduced serves as the oxidizing agent. Since \( \mathrm{Ag} \) is oxidized, it is the reducing agent, and \( \mathrm{Sn}^{2+} \) is reduced, so it is the oxidizing agent.

Key concepts

Oxidizing Agent

An oxidizing agent is a substance that brings about oxidation in another substance by accepting electrons from it. In other words, it is reduced during the reaction. The oxidizing agent plays a crucial role by

• Gaining electrons
• Decreasing its own oxidation state

In our problem, the oxidizing agent is identified by looking at the element that undergoes a reduction in its oxidation state. For instance, in the equation given, the oxidation state of \( \mathrm{Sn}^{2+} \) changes from +2 to 0, indicating that it gains electrons and is reduced in the process. Therefore, \( \mathrm{Sn}^{2+} \) acts as the oxidizing agent because it facilitates the oxidation of another substance.

By accepting electrons, the oxidizing agent ensures that the overall reaction proceeds through the transfer of electrons. This electron acceptance is vital because it drives forward many chemical processes.

Reducing Agent

The reducing agent is the substance that donates electrons to another element or compound, causing the latter to be reduced. In this donation process, the reducing agent itself becomes oxidized. The reducing agent is characterized by:

• Losing electrons
• Increasing its oxidation state

In our specific reaction, \( \mathrm{Ag} \) is initially in its elemental form with an oxidation state of 0. It gives away electrons, increasing its oxidation state to +1, leading to its identification as the reducing agent.

A handy way to remember this is by noting that the reducing agent "reduces" another substance by giving up electrons. This electron loss is critical for the oxidation-reduction process to occur, balancing the transfer and acceptance of electrons in the reaction.

Oxidation States

Oxidation states, or oxidation numbers, provide a way to keep track of electrons in a chemical reaction. They help us understand which elements gain or lose electrons. A few guidelines to remember when determining oxidation states include:

• The oxidation state of an atom in its elemental form is 0.
• The oxidation state of a monatomic ion is equal to its charge.
• In compounds, specific rules are followed for common elements.

In the given reaction:

• The oxidation state of \( \mathrm{Sn}^{2+} \) is +2, while \( \mathrm{Ag} \) is 0.
• In the products, \( \mathrm{Sn} \), in its elemental form, has an oxidation state of 0 and \( \mathrm{Ag}^{+} \) is +1.

By comparing these states before and after the reaction, we can identify which components are oxidized and reduced. Understanding oxidation states is essential in solving oxidation-reduction problems as it indicates electron movement in reactions.

Electron Transfer

Electron transfer is the fundamental process that occurs in oxidation-reduction (redox) reactions. It involves the movement of electrons from one species to another. This transfer is crucial to determining the role of substances in a redox reaction, as either oxidizing or reducing agents. In our example,

• \( \mathrm{Sn}^{2+} \) gains electrons, causing a reduction.
• \( \mathrm{Ag} \) loses electrons, resulting in oxidation.

This exchange allows the chemical reaction to proceed, with electrons moving from the reducing agent to the oxidizing agent. Understanding electron transfer helps in figuring out how redox reactions are balanced and enable energy changes in chemical systems.