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23.1: Direct Redox Reactions

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  • Gold and silver are widely used metals for making jewelry. One of the reasons these metals are employed for this purpose is that they are very unreactive. They do not react in contact with most other metals, so they are more likely to stay intact under challenging conditions. Who wants their favorite piece of jewelry to fall apart on them?

    Direct Redox Reactions

    When a strip of zinc metal is placed into a blue solution of copper (II) sulfate (figure below), a reaction immediately begins as the zinc strip begins to darken. If left in the solution for a longer period of time, the zinc will gradually decay due to oxidation to zinc ions. At the same time, the copper (II) ions from the solution are reduced to copper metal (see second figure below), which causes the blue copper (II) sulfate solution to become colorless.

    Figure 23.1.1: Copper sulfate solution.
    Figure 23.1.2: Reaction of zinc metal in copper sulfate solution.

    The process that occurs in this redox reaction is shown below as two separate half-reactions, which can then be combined into the full redox reaction.

    \[\begin{array}{ll} \text{Oxidation:} & \ce{Zn} \left( s \right) \rightarrow \ce{Zn^{2+}} \left( aq \right) + 2 \ce{e^-} \\ \text{Reduction:} & \ce{Cu^{2+}} \left( aq \right) + 2 \ce{e^-} \rightarrow \ce{Cu} \left( s \right) \\ \hline \text{Full Reaction:} & \ce{Zn} \left( s \right) + \ce{Cu^{2+}} \left( aq \right) \rightarrow \ce{Zn^{2+}} \left( aq \right) + \ce{Cu} \left( s \right) \end{array}\]

    Why does this reaction occur spontaneously? The activity series is a listing of elements in descending order of reactivity. An element that is higher in the activity series is capable of displacing an element that is lower on the series in a single-replacement reaction. This series also lists elements in order of ease of oxidation. The elements at the top are the easiest to oxidize, while those at the bottom are the most difficult to oxidize. The table below shows the activity series together with each element's oxidation half-reaction.

    Table 23.1.1: Activity Series of Metals
    Element Oxidation Half-Reaction
    Most active or most easily oxidized Lithium \(\ce{Li} \left( s \right) \rightarrow \ce{Li^+} \left( aq \right) + \ce{e^-}\)
    Potassium \(\ce{K} \left( s \right) \rightarrow \ce{K^+} \left( aq \right) + \ce{e^-}\)
    Barium \(\ce{Ba} \left( s \right) \rightarrow \ce{Ba^{2+}} \left( aq \right) + 2 \ce{e^-}\)
    Calcium \(\ce{Ca} \left( s \right) \rightarrow \ce{Ca^{2+}} \left( aq \right) + 2 \ce{e^-}\)
    Sodium \(\ce{Na} \left( s \right) \rightarrow \ce{Na^+} \left( aq \right) + \ce{e^-}\)
    Magnesium \(\ce{Mg} \left( s \right) \rightarrow \ce{Mg^{2+}} \left( aq \right) + 2 \ce{e^-}\)
    Aluminum \(\ce{Al} \left( s \right) \rightarrow \ce{Al^{3+}} \left( aq \right) + 3 \ce{e^-}\)
    Zinc \(\ce{Zn} \left( s \right) \rightarrow \ce{Zn^{2+}} \left( aq \right) + 2 \ce{e^-}\)
    Iron \(\ce{Fe} \left( s \right) \rightarrow \ce{Fe^{2+}} \left( aq \right) + 2 \ce{e^-}\)
    Nickel \(\ce{Ni} \left( s \right) \rightarrow \ce{Ni^{2+}} \left( aq \right) + 2 \ce{e^-}\)
    Tin \(\ce{Sn} \left( s \right) \rightarrow \ce{Sn^{2+}} \left( aq \right) + 2 \ce{e^-}\)
    Lead \(\ce{Pb} \left( s \right) \rightarrow \ce{Pb^{2+}} \left( aq \right) + 2 \ce{e^-}\)
    Hydrogen \(\ce{H_2} \left( g \right) \rightarrow 2 \ce{H^+} \left( aq \right) + 2 \ce{e^-}\)
    Copper \(\ce{Cu} \left( s \right) \rightarrow \ce{Cu^{2+}} \left( aq \right) + 2 \ce{e^-}\)
    Mercury \(\ce{Hg} \left( l \right) \rightarrow \ce{Hg^{2+}} \left( aq \right) + 2 \ce{e^-}\)
    Silver \(\ce{Ag} \left( s \right) \rightarrow \ce{Ag^+} \left( aq \right) + \ce{e^-}\)
    Platinum \(\ce{Pt} \left( s \right) \rightarrow \ce{Pt^{2+}} \left( aq \right) + 2 \ce{e^-}\)
    Least active or most difficult to oxidize Gold \(\ce{Au} \left( s \right) \rightarrow \ce{Au^{3+}} \left( aq \right) + 3 \ce{e^-}\)

    Notice that zinc is listed above copper on the activity series. which means that zinc is more easily oxidized than copper. That is why copper (II) ions can act as an oxidizing agent when put into contact with zinc metal. Ions of any metal that is below zinc, such as lead or silver, would oxidize the zinc in a similar reaction. These types of reactions are called direct redox reactions because the electrons flow directly from the atoms of one metal to the cations of the other metal. However, no reaction will occur if a strip of copper metal is placed into a solution of zinc ions, because the zinc ions are not able to oxidize the copper. In other words, such a reaction is nonspontaneous.


    • The activity series of metal reactivities is given.
    • Parameters for spontaneous reactions between metals are described.


    • CK-12 Foundation by Sharon Bewick, Richard Parsons, Therese Forsythe, Shonna Robinson, and Jean Dupon.