A Student Balances The Following Redox Reaction Using Half-reactions:${ Al + Mn^{2+} \longrightarrow Al^{3+} + Mn }$How Many Electrons Will Be Lost In Total?A. 1 B. 2 C. 5 D. 6

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Understanding Redox Reactions

Redox reactions, also known as oxidation-reduction reactions, involve the transfer of electrons between two species. These reactions are crucial in various chemical processes, including combustion, corrosion, and electrochemistry. In this article, we will focus on balancing redox reactions using half-reactions, with a specific example of the reaction between aluminum (Al) and manganese (Mn) ions.

The Importance of Balancing Redox Reactions

Balancing redox reactions is essential to ensure that the number of electrons lost by one species is equal to the number of electrons gained by another species. This balance is critical in predicting the outcome of a chemical reaction, including the formation of products and the release of energy. In the context of the given reaction, balancing the redox reaction will help us determine the total number of electrons lost.

Balancing the Redox Reaction

The given redox reaction is:

Al+Mn2+⟶Al3++Mn{ Al + Mn^{2+} \longrightarrow Al^{3+} + Mn }

To balance this reaction, we need to identify the oxidation and reduction half-reactions. The oxidation half-reaction involves the loss of electrons, while the reduction half-reaction involves the gain of electrons.

Oxidation Half-Reaction

In the oxidation half-reaction, aluminum (Al) loses three electrons to form aluminum ions (Al³⁺):

Al⟶Al3++3e−{ Al \longrightarrow Al^{3+} + 3e^- }

Reduction Half-Reaction

In the reduction half-reaction, manganese ions (Mn²⁺) gain two electrons to form manganese (Mn):

Mn2++2e−⟶Mn{ Mn^{2+} + 2e^- \longrightarrow Mn }

Balancing the Redox Reaction

To balance the redox reaction, we need to ensure that the number of electrons lost in the oxidation half-reaction is equal to the number of electrons gained in the reduction half-reaction. In this case, the oxidation half-reaction involves the loss of three electrons, while the reduction half-reaction involves the gain of two electrons. To balance the reaction, we need to multiply the reduction half-reaction by 1.5:

Mn2++3e−⟶Mn{ Mn^{2+} + 3e^- \longrightarrow Mn }

Now, the number of electrons lost in the oxidation half-reaction is equal to the number of electrons gained in the reduction half-reaction.

Determining the Total Number of Electrons Lost

To determine the total number of electrons lost, we need to add the number of electrons lost in the oxidation half-reaction to the number of electrons gained in the reduction half-reaction. In this case, the oxidation half-reaction involves the loss of three electrons, while the reduction half-reaction involves the gain of three electrons (1.5 x 2). Therefore, the total number of electrons lost is:

3e−+3e−=6e−{ 3e^- + 3e^- = 6e^- }

Conclusion

In conclusion, the total number of electrons lost in the given redox reaction is six. This is determined by balancing the redox reaction using half-reactions and ensuring that the number of electrons lost in the oxidation half-reaction is equal to the number of electrons gained in the reduction half-reaction.

Answer

The correct answer is D. 6.

Additional Tips and Resources

  • To balance redox reactions, it is essential to identify the oxidation and reduction half-reactions and ensure that the number of electrons lost in the oxidation half-reaction is equal to the number of electrons gained in the reduction half-reaction.
  • The half-reaction method is a powerful tool for balancing redox reactions, but it requires careful attention to detail and a thorough understanding of the underlying chemistry.
  • For more information on balancing redox reactions, consult a reliable chemistry textbook or online resource, such as the Khan Academy or Chemistry LibreTexts.

References

  • Petrucci, R. H., Harwood, W. S., & Herring, F. G. (2006). General chemistry: Principles and modern applications. Pearson Prentice Hall.
  • Atkins, P. W., & De Paula, J. (2010). Physical chemistry. Oxford University Press.
  • McMurry, J., & Fay, R. C. (2011). Chemistry. Pearson Prentice Hall.
    A Student's Guide to Balancing Redox Reactions: Q&A =====================================================

Understanding Redox Reactions

Redox reactions, also known as oxidation-reduction reactions, involve the transfer of electrons between two species. These reactions are crucial in various chemical processes, including combustion, corrosion, and electrochemistry. In this article, we will focus on balancing redox reactions using half-reactions, with a specific example of the reaction between aluminum (Al) and manganese (Mn) ions.

Q&A: Balancing Redox Reactions

Q: What is the purpose of balancing redox reactions?

A: The purpose of balancing redox reactions is to ensure that the number of electrons lost by one species is equal to the number of electrons gained by another species. This balance is critical in predicting the outcome of a chemical reaction, including the formation of products and the release of energy.

Q: How do I identify the oxidation and reduction half-reactions?

A: To identify the oxidation and reduction half-reactions, you need to determine which species is losing electrons (oxidation) and which species is gaining electrons (reduction). In the given reaction, aluminum (Al) loses three electrons to form aluminum ions (Al³⁺), while manganese ions (Mn²⁺) gain two electrons to form manganese (Mn).

Q: How do I balance the redox reaction using half-reactions?

A: To balance the redox reaction using half-reactions, you need to ensure that the number of electrons lost in the oxidation half-reaction is equal to the number of electrons gained in the reduction half-reaction. In this case, the oxidation half-reaction involves the loss of three electrons, while the reduction half-reaction involves the gain of two electrons. To balance the reaction, you need to multiply the reduction half-reaction by 1.5.

Q: What is the total number of electrons lost in the given redox reaction?

A: The total number of electrons lost in the given redox reaction is six. This is determined by balancing the redox reaction using half-reactions and ensuring that the number of electrons lost in the oxidation half-reaction is equal to the number of electrons gained in the reduction half-reaction.

Q: What are some common mistakes to avoid when balancing redox reactions?

A: Some common mistakes to avoid when balancing redox reactions include:

  • Failing to identify the oxidation and reduction half-reactions
  • Not ensuring that the number of electrons lost in the oxidation half-reaction is equal to the number of electrons gained in the reduction half-reaction
  • Not multiplying the reduction half-reaction by the correct factor to balance the reaction

Q: What are some additional tips for balancing redox reactions?

A: Some additional tips for balancing redox reactions include:

  • Using a systematic approach to identify the oxidation and reduction half-reactions
  • Ensuring that the number of electrons lost in the oxidation half-reaction is equal to the number of electrons gained in the reduction half-reaction
  • Verifying the balance of the reaction by checking the number of electrons lost and gained

Conclusion

In conclusion, balancing redox reactions using half-reactions is a powerful tool for predicting the outcome of chemical reactions. By identifying the oxidation and reduction half-reactions and ensuring that the number of electrons lost is equal to the number of electrons gained, you can determine the total number of electrons lost in a given redox reaction. Remember to avoid common mistakes and use a systematic approach to balance redox reactions.

Additional Resources

  • Khan Academy: Redox Reactions
  • Chemistry LibreTexts: Redox Reactions
  • Petrucci, R. H., Harwood, W. S., & Herring, F. G. (2006). General chemistry: Principles and modern applications. Pearson Prentice Hall.
  • Atkins, P. W., & De Paula, J. (2010). Physical chemistry. Oxford University Press.
  • McMurry, J., & Fay, R. C. (2011). Chemistry. Pearson Prentice Hall.