Question

By using the data in Appendix E, determine whether each of the following substances is likely to serve as an exidant or a reductant: (a) \(\mathrm{Cl}_{2}\) (g), (b) \(\mathrm{MnO}_{4}^{-}\)(aq, acidic solution), (c) \(\mathrm{Ba}\) (s), (d) \(\mathrm{Zn}(\) s).

Short answer

Based on their standard electrode potentials, the given substances serve as follows: (a) \(\mathrm{Cl}_2\) (g) acts as an oxidant. (b) \(\mathrm{MnO}_4^{-}\)(aq, acidic solution) acts as an oxidant. (c) \(\mathrm{Ba}\) (s) acts as a reductant. (d) \(\mathrm{Zn}\) (s) acts as a reductant.

Step-by-step solution

Determine Standard Electrode Potentials

Using Appendix E, find the standard electrode potentials (E°) of each species in their respective reactions. For \(\mathrm{Cl}_{2}\) (g): The reaction for this species is, \(\mathrm{Cl}_{2}\) + 2 \(\mathrm{e}^{-}\) ⟶ 2 \(\mathrm{Cl}^{-}\) From Appendix E, the standard electrode potential for this reaction is E° = +1.36 V For \(\mathrm{MnO}_4^{-}\)(aq, acidic solution): In an acidic medium, the reaction for this species is, \(\mathrm{MnO}_4^{-}\) + 8 \(\mathrm{H}^{+}\) + 5 \(\mathrm{e}^{-}\) ⟶ \(\mathrm{Mn}^{2+}\) + 4 \(\mathrm{H}_2\mathrm{O}\) From Appendix E, the standard electrode potential for this reaction is E° = +1.51 V For \(\mathrm{Ba}\)(s): The reaction for this species is, \(\mathrm{Ba}\) ⟶ \(\mathrm{Ba}^{2+}\) + 2 \(\mathrm{e}^{-}\) From Appendix E, the standard electrode potential for this reaction is E° = -2.92 V For \(\mathrm{Zn}\) (s): The reaction for this species is, \(\mathrm{Zn}\) ⟶ \(\mathrm{Zn}^{2+}\) + 2 \(\mathrm{e}^{-}\) From Appendix E, the standard electrode potential for this reaction is E° = -0.76 V

Predict the Role of Each Substance as Oxidant or Reductant

Based on the standard electrode potentials (E°) obtained above, we can predict the nature of each substance. (a) For \(\mathrm{Cl}_2\) (g), the standard electrode potential is +1.36 V, which indicates a strong tendency to accept electrons, and thus, it acts as an oxidizing agent (or oxidant). (b) For \(\mathrm{MnO}_4^{-}\)(aq, acidic solution), the standard electrode potential is +1.51 V, which indicates a strong tendency to accept electrons. Hence, it serves as an oxidizing agent (or oxidant). (c) For \(\mathrm{Ba}\) (s), the standard electrode potential is -2.92 V, which indicates a strong tendency to lose electrons and be oxidized. So, \(\mathrm{Ba}\) (s) acts as a reducing agent (or reductant). (d) For \(\mathrm{Zn}\) (s), the standard electrode potential is -0.76 V, which indicates a tendency to lose electrons and be oxidized. Therefore, \(\mathrm{Zn}\) (s) serves as a reducing agent (or reductant). Final answers: (a) \(\mathrm{Cl}_2\) (g) - Oxidant (b) \(\mathrm{MnO}_4^{-}\)(aq, acidic solution) - Oxidant (c) \(\mathrm{Ba}\) (s) - Reductant (d) \(\mathrm{Zn}\) (s) - Reductant

Key concepts

Standard Electrode Potentials

Standard electrode potentials, represented by the symbol \( E^° \), are crucial in predicting the behavior of substances in redox reactions. For any half-reaction involving electron transfer, the standard electrode potential indicates how easily a substance gains or loses electrons under standard conditions. A more positive \( E^° \) implies a stronger tendency to gain electrons, acting as an oxidizing agent. Conversely, a more negative \( E^° \) suggests a stronger inclination to lose electrons, functioning as a reducing agent.

When looking at the standard electrode potentials found in Appendix E, we observed that chlorine gas \( \mathrm{Cl}_2 \) has an \( E^° \) of +1.36 V, indicating it is likely to act as an oxidizing agent. The permanganate ion \( \mathrm{MnO}_4^- \) in an acidic solution has an even higher \( E^° \) of +1.51 V, reinforcing its role as a powerful oxidizing agent.

On the other hand, barium metal \( \mathrm{Ba} \) and zinc \( \mathrm{Zn} \) exhibit negative standard electrode potentials, \( -2.92 \) V and \( -0.76 \) V respectively, making them good reducing agents as they have a tendency to donate electrons.

Oxidizing Agent

An oxidizing agent is a substance that gains electrons in a chemical reaction. This process is also called reduction. As the oxidizing agent accepts electrons, it facilitates the oxidation of another substance. Typically, oxidizing agents have positive standard electrode potentials, indicating their capability to accept electrons readily and drive oxidative reactions.

In our example, \( \mathrm{Cl}_2 \) and \( \mathrm{MnO}_4^- \) have positive \( E^° \) values (+1.36 V and +1.51 V respectively), marking them as strong oxidizing agents.

• \( \mathrm{Cl}_2 \): Known for its role in bleach and disinfection, where its oxidizing properties help in breaking down stains and killing bacteria.
• \( \mathrm{MnO}_4^- \): Used in various chemical industries for its power to oxidize different chemicals, hence, it's vital in cleaning and chemical synthesis.

Understanding the role of oxidizing agents can aid in grasping how substances undergo transformation and energy transfer in electrochemical processes.

Reducing Agent

A reducing agent is a substance that loses electrons during a chemical reaction, effectively undergoing oxidation to facilitate the reduction of another substance. Reducing agents typically have negative standard electrode potentials, making them prone to oxidizing by losing electrons.

In the context of our exercise, both \( \mathrm{Ba} \) and \( \mathrm{Zn} \) have negative standard electrode potentials (-2.92 V and -0.76 V respectively), which makes them strong reducing agents. They are eager to donate electrons to other substances, thus driving reduction processes.

• \( \mathrm{Ba} \): Commonly used in fireworks and in vacuum tubes where its reducing properties aid in the removal of traces of gases.
• \( \mathrm{Zn} \): Utilized in galvanization processes to protect iron or steel from rusting, its reducing nature allows it to oxidize preferentially, thus conserving the integrity of the protected metal.

Recognizing reducing agents is key in manipulating reactions in labs and in industrial applications, especially in the synthesis and stabilization of various compounds.

Electrochemistry

Electrochemistry is the branch of chemistry that deals with the relationship between electricity and chemical reactions. It's essential for understanding redox reactions, where electron transfer dictates the flow of electricity. In electrodes within a cell, redox reactions take place, resulting in the conversion of chemical energy into electrical energy and vice versa.

In electrochemical cells, oxidizing agents attract electrons and facilitate electricity flow, while reducing agents push electrons through the circuit. This interaction creates energy that can be harnessed for various applications, from powering batteries to running complex chemical syntheses.

• Batteries: Utilize redox reactions where oxidizing and reducing agents enable the conversion of stored chemical energy into electric energy to power devices.
• Electroplating: Uses electrochemical processes where a metal is deposited onto a surface, enhancing its characteristics or aesthetic appeal.

Understanding these concepts not only helps grasp theoretical aspects but also drives innovation in technology and industrial practices through advanced manipulation of electrochemical processes.