Question

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

Short answer

In the given equation, the oxidizing agent is $A{g}^{+}$ and the reducing agent is Fe in its neutral state. This is because Silver (Ag) gains electrons and goes from its 1+ state ($A{g}^{+}$) to its neutral (0) state, while Iron (Fe) loses electrons and goes from its neutral (0) state to its 2+ state ($F{e}^{2+}$).

Step-by-step solution

Determine the oxidation states of the elements

We start by determining the oxidation states of Iron (Fe), Silver (Ag), and their respective ions in the equation. Before the reaction, Fe is in its neutral (0) state, and Ag is in its 1+ state ($A{g}^{+}$). After the reaction, Fe becomes its 2+ state ($F{e}^{2+}$), and Ag returns to its neutral (0) state.

Identify the element that gains electrons

In the equation, Silver (Ag) goes from its 1+ state ($A{g}^{+}$) to its neutral (0) state. Since the loss of positive charge indicates the gain of electrons, Silver has gained electrons.

Identify the element that loses electrons

Iron (Fe) goes from its neutral (0) state to its 2+ state ($F{e}^{2+}$). Since the gain of positive charge indicates the loss of electrons, Iron has lost electrons.

Identify the oxidizing agent and the reducing agent

Based on the information from steps 2 and 3: - Silver (Ag) gains electrons and therefore is reduced in the reaction. Thus, $A{g}^{+}$ is the oxidizing agent. - Iron (Fe) loses electrons and therefore is oxidized in the reaction. Thus, Fe (in its neutral state) is the reducing agent. In conclusion, the oxidizing agent in the given equation is $A{g}^{+}$, and the reducing agent is Fe (in its neutral state).

Key concepts

Oxidizing Agent

In a chemical reaction, an oxidizing agent plays a crucial role by gaining electrons from another substance. This process, known as reduction, involves the oxidizing agent itself undergoing a decrease in oxidation state.

An oxidizing agent essentially "accepts" electrons during the reaction. For instance, in our equation, silver ions ($A{g}^{+}$) act as the oxidizing agent. They accept electrons from iron in its metallic form. This transforms the silver ions from $A{g}^{+}$ to neutral silver ($Ag$), indicating that the silver ions have been reduced.

Remember:

• Oxidizing agents cause oxidation by being reduced themselves.
• They gain electrons and decrease in oxidation state.

Reducing Agent

A reducing agent is a substance that donates electrons to another substance in a chemical reaction. This action is responsible for reducing the oxidation state of the receiving molecule while the reducing agent itself becomes oxidized.

In our discussed equation, iron in its neutral state acts as the reducing agent. It loses electrons, embodying the typical behavior of a reducing agent, which results in its oxidation. By giving up electrons, iron transforms into $F{e}^{2+}$, experiencing an increase in oxidation state.

Here's what to remember about reducing agents:

• They cause reduction by being oxidized themselves.
• They donate electrons and increase in their oxidation state.

Oxidation State

The oxidation state, or oxidation number, indicates the degree of oxidation of an atom in a molecule. It represents the hypothetical charge an atom would possess if the compound was composed of ions.

In our equation, the oxidation states are vital for identifying the transfer of electrons:

• Iron ($Fe$) starts in a neutral state with an oxidation number of 0, then moves to $F{e}^{2+}$, indicating it has lost two electrons.
• Silver ($A{g}^{+}$) begins with a +1 oxidation state and changes to 0, showing it has gained an electron.

Paying attention to changes in oxidation states helps decode the electron transfer that occurs during the reaction.

Electron Transfer

Electron transfer is crucial in oxidation-reduction reactions. It is the movement of electrons from one reactant to another, which leads to changes in their oxidation states.

In any redox reaction, such as our example, one atom's electrons are transferred to another atom, either partially or entirely. For example:

• Iron loses two electrons, becoming oxidized to $F{e}^{2+}$.
• Silver gains electrons, being reduced to $Ag$.

Monitoring electron flow allows us to identify oxidizing and reducing agents, thus completing the understanding of the redox process.