Module 5

Lesson 1.3  Oxidizing and Reducing Agents



Key Concepts


To introduce the concept of oxidizing agents and reducing agents, return to the example of the copper coil in the silver nitrate solution. The balanced redox reaction is shown below:


\( \mathrm { Cu(s) + 2 AgNO_3(aq) \rightarrow Cu(NO_3)_2(aq) + 2 Ag(s) } \)


You will recall from earlier in this lesson that the copper atom loses electrons in this reaction.


What is causing the copper atom to undergo oxidation? The answer is that the silver ions are causing the oxidation. Therefore, the silver ions are classified as the oxidizing agent (OA). Strong oxidizing agents have a strong affinity for electrons and promote the oxidation (loss of electrons) in another substance during a reaction.


Oxidizing Agent (OA):     a substance that causes oxidation in another substance while itself undergoes reduction by gaining electrons

Recall that silver ions gain electrons in this reaction.


What is causing the silver ion to undergo reduction? The answer is that the solid copper atoms are causing the reduction. Therefore, solid copper is classified as the reducing agent (RA). Strong reducing agents have a weak affinity for electrons and tend to cause reduction in another substance during a reaction. In other words, reducing agents tend to lose electrons to other substances.


Reducing Agent(RA):     a substance that causes reduction in another substance while itself undergoing oxidation by losing electrons


 
ProcessHalf-reaction
Reduction\( \mathrm { Ag^+(aq) + 1e^- \rightarrow Ag(s) } \)
Oxidation\( \mathrm { Cu(s) \rightarrow Cu^{2+}(aq) + 2 e^- } \)           
Reducing Agent\( \mathrm { Cu(s) } \)
Oxidizing Agent\( \mathrm { Ag^+(aq) } \)


Fig. 1




Learning Tip

Generally, metallic ions are oxidizing agents and solid metals are reducing agents. However, some metallic ions can act either as oxidizing agents or as reducing agents. The "middle" charge of multivalent ions can either gain or lose electrons as represented by the following half-reactions for chromium(II) ions.

\( \mathrm { Cr^{2+}(aq) + 2e^- \rightarrow Cr(s) } \)
\( \mathrm { Cr^{2+}(aq) } \)    shows a gain of electrons/OA 
\( \mathrm { Cr^{2+}(aq) \rightarrow Cr^{3+}(aq) + 1e^- } \)
\( \mathrm { Cr^{2+}(aq) } \)    shows a loss of electrons/RA

Probably, you have noticed that the oxidizing agents and reducing agents are usually different entities. However, this is not always the case. Some chemical species can act as both oxidizing agent and reducing agent at the same time. These species appear on both sides of the redox table. For example, consider the Cu1+ ion. Copper(I) ions can either lose or gain electrons and, therefore, can act as either reducing agents or oxidizing agents.



Read "Oxidizing and Reducing Agents" on pages 568 and 569 of the textbook.

Check Your Understanding


For each of the following net reactions, write the reduction and oxidation half-reactions and identify the reducing agent and the oxidizing agent. Click on the link below to check your work.

  1. \( \mathrm { Cr(s) + Sn^{2+}(aq) \rightarrow Cr^{2+}(aq) + Sn(s) } \)

  2. \( \mathrm { Fe(s) + 2 H^+(aq) \rightarrow Fe^{2+}(aq) + H_{2}(g) } \)

  3. \( \mathrm { Sn^{4+}(aq) + 2 Cr^{2+}(aq) \rightarrow 2 Cr^{3+}(aq) + Sn^{2+}(aq) } \) 

  1. Process Half-reaction
    Reduction \( \mathrm { Sn^{2+}(aq) + 2e^- \rightarrow Sn(s) } \)
    Oxidation \( \mathrm { Cr(s) \rightarrow Cr^{2+}(aq) + 2e^- } \)
    Reducing Agent \( \mathrm { Cr(s) } \)
    Oxidizing Agent \( \mathrm { Sn^{2+}(aq) } \)


  2. Process Half-reaction
    Reduction \( \mathrm { 2 H^+(aq) + 2e^- \rightarrow H_2(g) } \)
    Oxidation \( \mathrm { Fe(s) \rightarrow Fe^{2+}(aq) + 2e^- } \)
    Reducing Agent \( \mathrm { Fe(s) } \)
    Oxidizing Agent \( \mathrm { H^+(aq } \)


  3. Process Half-reaction
    Reduction \( \mathrm { Sn^{4+}(aq) + 2e^- \rightarrow Sn^{2+}(aq) } \)
    Oxidation \( \mathrm { Cr^{2+}(aq) \rightarrow Cr^{3+}(aq) +1e^- } \)
    Reducing Agent \( \mathrm { Cr^{2+}(aq) } \)
    Oxidizing Agent \( \mathrm { Sn^{4+}(aq) } \)