Of course! Let's analyze the given redox reaction step by step to identify the reducing agent and the oxidizing agent.
The reaction provided is:
[tex]\[ \text{Mn} (s) + \text{Fe}^{2+}(aq) \longrightarrow \text{Mn}^{2+}(aq) + \text{Fe} (s) \][/tex]
### Step 1: Determine the oxidation states of each element before and after the reaction.
- Manganese (Mn):
- Before the reaction: Mn is in its elemental form ([tex]\(\text{Mn} (s)\)[/tex]), so its oxidation state is 0.
- After the reaction: Mn is now in the form of [tex]\(\text{Mn}^{2+}\)[/tex], so its oxidation state is +2.
- Iron (Fe):
- Before the reaction: Fe is in the form of [tex]\(\text{Fe}^{2+}\)[/tex], so its oxidation state is +2.
- After the reaction: Fe is in its elemental form ([tex]\(\text{Fe} (s)\)[/tex]), so its oxidation state is 0.
### Step 2: Identify which element is oxidized and which is reduced.
- Oxidation: This process involves the loss of electrons.
- Mn goes from an oxidation state of 0 to +2, meaning Mn loses 2 electrons. Therefore, Mn is oxidized.
- Reduction: This process involves the gain of electrons.
- Fe goes from an oxidation state of +2 to 0, meaning Fe gains 2 electrons. Therefore, Fe is reduced.
### Step 3: Determine the reducing agent and the oxidizing agent.
- Reducing Agent: The substance that is oxidized (loses electrons) and, in doing so, reduces another substance.
- In this reaction, Mn (s) loses electrons, so [tex]\(\text{Mn} (s)\)[/tex] is the reducing agent.
- Oxidizing Agent: The substance that is reduced (gains electrons) and, in doing so, oxidizes another substance.
- In this reaction, [tex]\(\text{Fe}^{2+}(aq)\)[/tex] gains electrons, so [tex]\(\text{Fe}^{2+}(aq)\)[/tex] is the oxidizing agent.
### Conclusion
Based on the analysis:
- The reducing agent is [tex]\( \text{Mn} (s) \)[/tex].
- The oxidizing agent is [tex]\( \text{Fe}^{2+}(aq) \)[/tex].
Thus, we have systematically identified the reducing agent and the oxidizing agent for the given redox reaction.
