Oxidation and reduction reactions have been know for millennia but were not understood until the 17th century. The terms come from metallurgy. Most metals do not naturally exist in their metallic forms (except gold and silver), but were extracted from rocks and minerals. As such the ores were “reduced” to a small amount of metal from a large amount of ore. It was noted that the metals would react with oxygen, form a new substance, and hence were oxidized.
We now understand that redox (oxidation-reduction) reactions involve the transfer of electrons. Consider, for instance, the single displacement reaction between Copper ions, Cu2+(aq), and Zinc metal, Zn(s). The “(aq)” designation on Cu stands for “aqueous” and means that the ion is dissolved in water. The “(s)” designation on Zn means that the Zinc metal is a solid. One could add Zinc metal to an aqueous solution of Cu(NO3)2 and these react spontaneously according to the chemical reaction represented by the molecular and net-ionic equations below:
Cu(NO3)2 (aq) + Zn(s) Cu(s) + Zn(NO3)2 (aq)(molecular equation)
Cu2+(aq) + Zn(s) Cu(s) + Zn2+(aq)(net-ionic equation)
Electrons were exchanged in this reaction, making it a redox reaction. To make the electron exchange more apparent, we can break this reaction into “half reactions”:
Zn(s) Zn2+(aq) + 2e- (Zinc metal gave up electrons)
Cu2+(aq) + 2e- Cu(s) (Copper ion gained electrons)
Substances that gain electrons are said to be reduced and substances that give up electrons are oxidized. Therefore, in the above reaction, Zn is oxidized and Cu is reduced. Another way of looking at the above reaction is to consider what the Cu2+ ion is doing to the Zn. Cu2+ is causing the Zn to be oxidized, so Cu is acting as an oxidizing agent. Conversely, Zn is causing Cu2+ to be reduced, so Zn is a reducing agent. Reactions such as that between Zn and Cu only go in one direction. In other words, we will not see the reverse reaction occur:
Zn2+(aq) + Cu(s) Zn(s) + Cu2+(aq) [non-spontaneous or DOES NOT OCCUR]
We can summarize these by saying that Zn metal is able to reduce Cu2+, but Cu metal is not able to reduce Zn2+. In other words, Zn is a stronger reducing agent than Cu.
Part I: The stockroom of the virtual lab contains nitrate solutions of Cu2+, Mg2+, Zn2+, and Pb2+ ions, and the corresponding metal solids – Cu, Mg, Zn, and Pb. Your main task is to order Cu, Mg, Zn and Pb from strongest to weakest reducing agent. You can do this by combining any pair consisting of a solid and a solution and determining whether a reaction occurs. For example, add the solid in 0.1 g increments to 10 mL aliquot of your solution. Observe what visual changes occur in the reaction beaker and check its contents in the information chart at the left of the workbench before and after mixing reactants. If nothing happens, add up to two more 0.1 g increments of your solid. If significant changes occur, then the added metal can reduce the ion. Run enough reactions to determine the proper order of reducing power.
What do you expect to be true of the strongest reducing agent? What do you expect to be true of the weakest reducing agent?
Part II: The virtual lab stockroom also contains a solution of Ag+ ion and Ag metal.
Write a balanced chemical reaction for Cu metal reducing Ag2+ ion. (Hints: Don’t forget to balance charges in your reaction. How many electrons does Ag give up as it goes from Ag to Ag+? How many electrons does Cu gain as it goes from Cu2+ to Cu? How is this difference in number of electrons reflected in your balanced chemical reaction?)
Perform experiments to determine how strong Ag is as a reducing agent. Tabulate your results as in Part I. Where does Ag lie relative to Cu, Mg, Zn and Pb?