Question

Which would you expect to be the stronger reducing agent, \(\mathrm{Cr}^{2+}\) or \(\mathrm{Fe}^{2+}\) ?

Short answer

Comparing the standard reduction potentials for their respective oxidation half-reactions, we find that \(\mathrm{Fe}^{2+}\) has a more negative reduction potential (-0.77 V) than \(\mathrm{Cr}^{2+}\) (+0.41 V). This indicates a higher tendency to lose electrons and undergo oxidation. Therefore, \(\mathrm{Fe}^{2+}\) is the stronger reducing agent.

Step-by-step solution

Understand the concept of reduction potentials

A reduction potential is a measure of the tendency of a chemical species to be reduced or to accept electrons. The more positive the reduction potential, the more likely a chemical species is to accept electrons and undergo reduction. On the other hand, a more negative reduction potential indicates a greater tendency for a chemical species to lose electrons and undergo oxidation. In the context of this problem, a stronger reducing agent is the one with a higher tendency to lose electrons and undergo oxidation, meaning it will have a more negative reduction potential.

Find the standard reduction potentials of \(\mathrm{Cr}^{2+}\) and \(\mathrm{Fe}^{2+}\)

To compare the reducing strength of the two ions, we need to look at their standard reduction potentials. These can be found in a standard reduction potentials table or the Nernst equation may be used. From the standard reduction potential table, we find the following half-reactions and their respective reduction potentials: 1. \(\mathrm{Cr}^{3+} + e^- → \mathrm{Cr}^{2+}\), E° = -0.41 V 2. \(\mathrm{Fe}^{3+} + e^- → \mathrm{Fe}^{2+}\), E° = +0.77 V However, in order to compare their reducing abilities, we need to reverse the half-reactions to make them oxidation reactions (since reducing agents undergo oxidation): 1. \(\mathrm{Cr}^{2+} → \mathrm{Cr}^{3+} + e^-\), E° = +0.41 V 2. \(\mathrm{Fe}^{2+} → \mathrm{Fe}^{3+} + e^-\), E° = -0.77 V

Compare the standard reduction potentials

Now that we have the standard reduction potentials for both ions as their corresponding oxidation half-reactions, we can compare their potentials: 1. \(\mathrm{Cr}^{2+} → \mathrm{Cr}^{3+} + e^-\), E° = +0.41 V 2. \(\mathrm{Fe}^{2+} → \mathrm{Fe}^{3+} + e^-\), E° = -0.77 V We can see that \(\mathrm{Fe}^{2+}\) has the more negative reduction potential (-0.77 V) when compared to \(\mathrm{Cr}^{2+}\) (+0.41 V).

Conclusion

Since \(\mathrm{Fe}^{2+}\) has a more negative reduction potential than \(\mathrm{Cr}^{2+}\), it has a higher tendency to lose electrons and undergo oxidation. Therefore, we can conclude that \(\mathrm{Fe}^{2+}\) is the stronger reducing agent.

Key concepts

Reduction Potentials

Reduction potentials are vital in understanding how chemical species behave in redox reactions, where there is a transfer of electrons between species. The reduction potential indicates how likely a chemical species is to gain electrons.
The standard reduction potential is represented by E°, quantified in volts, and is measured under standard conditions.
A more positive value signifies a stronger tendency of a chemical species to gain electrons (thus being reduced), while a more negative value suggests a greater inclination to lose them (thus being oxidized).

• A species that is easily reduced (highly positive E°) will typically not be a strong reducing agent, but rather a strong oxidizing agent.
• Conversely, a species with a highly negative E° is more likely to be a strong reducing agent.

In summary, knowing a chemical species' reduction potential helps predict its behavior in electron transfer processes.

Oxidation Reactions

Oxidation reactions are processes where a chemical species loses electrons, frequently paired with reduction reactions where another species gains these electrons. This combined process is known as a redox reaction.
A reducing agent is a substance that donates electrons to another species by undergoing oxidation itself.
When looking at oxidation reactions, the focus is on the electron donation from the reducing agent.

• For example, in the oxidation of \({\mathrm{Cr}^{2+}}\), the species loses electrons to form \({\mathrm{Cr}^{3+}}\).
• In this transformation, the electron loss signifies that \({\mathrm{Cr}^{2+}}\) acts as a reducing agent.

Recognizing the oxidation process is essential in identifying the nature and strength of reducing agents in chemical reactions.

Electrochemical Series

The electrochemical series is a chart of chemical species arranged by their standard reduction potentials. It serves as a reference for predicting how different substances will behave in redox reactions.
This series is organized from the most positive reduction potentials at the top to the most negative at the bottom.

• Species located higher up, with high positive reduction potentials, typically act as strong oxidizing agents.
• Conversely, those lower down with negative potentials are often potent reducing agents.

Utilizing the electrochemical series can simplify the prediction of which chemical species will act as reducing or oxidizing agents in a given reaction.

Electron Transfer

Electron transfer is a key concept in redox reactions, representing the movement of electrons from one chemical species to another. This process is fundamental to understanding how reducing and oxidizing agents function.
During a redox reaction:

• The reducing agent loses electrons, undergoing oxidation.
• The oxidizing agent gains electrons, undergoing reduction.
• This electron exchange influences the energy of the reaction.

Understanding electron transfer is crucial for grasping the principles behind how chemical reactions proceed and how energy is distributed in these processes.

Chemical Species

Chemical species refer to atoms, molecules, ions, or different substances participating in chemical reactions. Identifying and understanding the role of chemical species in a reaction is essential for predicting reaction outcomes.
In redox reactions:

• The reducing agent is the species that donates electrons.
• The oxidizing agent is the species that accepts electrons.

Each chemical species has unique properties, like reduction potentials, that determine how they interact in reactions. These properties help chemists predict which species will act as a reducing or oxidizing agent.