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Electrochemistry: Introductory Concepts

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  • Learning Objectives

    Following this activity, students should be able to:

    • Recall basic concepts about redox chemistry
      • Identifying oxidation and reduction
      • Identify the oxidizing and reducing agent in a redox reaction
    • Interpret the relative magnitudes of standard reduction potentials
    • Identify the cathode and anode in an electrochemical cell
    • Calculate the standard potential difference of an electrochemical cell
    • Decipher line notation representing an electrochemical cell
    • Apply the Nernst equation to nonstandard conditions
    • Relating potentials across multiple reference electrodes

    Based on your recollection of general chemistry…

    Consider the reaction below and use it to complete the questions below.

    \[Pb^{2+}(aq) + Zn(s) \rightleftharpoons Zn^{2+}(aq) + Pb(s)\nonumber\]

    1. Define what is meant by oxidation and reduction. What is getting oxidized and reduced in the reaction (left to right) above?







    1. Electrochemical processes are often represented by half reactions. Write the two separate half reactions for the overall reaction above.


      Reduction half-reaction:


      Oxidation half-reaction:



    1. Define what is meant by oxidizing agent and reducing agent. What is acting as the oxidizing agent and what is meant by the reducing agent?


      Oxidizing agent


      Reducing agent


    1. In an electrochemical cell one electrode acts as the cathode and the other acts as the anode. How do you know what is what? Assuming the reaction above occurs as written from left to right, which electrode is the anode and which is the cathode?






    Standard Reduction Potentials

    The thermodynamic driving force for a particular electrochemical reaction is quantified by the Standard Reduction Potential (E°). Each redox pair (e.g., Ag and Ag+) has its own intrinsic tendency to donate or accept electrons, represented by E° (Table 1). Higher, more positive E° indicates reduction is more likely to occur and lower, more negative E° indicates oxidation is more thermodynamically favored.

    Table 1. Standard Reduction Potentials


    E° (V)

    \(Ag^+ + e^- \rightleftharpoons Ag(s)\)


    \(Cu^{2+} + 2e^- \rightleftharpoons Cu(s)\)


    \(2H^+ + 2e^- \rightleftharpoons H_2(g)\)


    \(Pb^{2+} + 2e^- \rightleftharpoons Pb(s)\)


    \(Cd^{2+} + 2e^- \rightleftharpoons Cd(s)\)


    \(Zn^{2+} + 2e^- \rightleftharpoons Zn(s)\)


    Standard reduction potentials are potentials for the specific reactions in which all species are in their standard state, which is indicated by the “°” symbol. This means 1 M concentration for aqueous solutions, 1 bar for gases and the pure substance for liquids or solids. Combining two half-reactions for an overall redox reaction allows one to calculate the overall standard potential of a cell (E°cell):

    \[E^\circ_{cell} = E^\circ_{red} - E^\circ_{ox} \hspace{30px} \textrm{or} \hspace{30px} E^\circ = E^\circ_+ - E^\circ_- \nonumber\]

    Where the reduction potential of the reduction half reaction (cathode) is E°red or E°+ and the reduction potential of the oxidation half reaction (anode) is E°ox or E°-. Use the values above to calculate the overall standard potentials of the reactions below. Use the sign of E° to predict if the reaction is spontaneous as written. (Recall Δ = -nFE°)



    E° (V)


    \(Pb^{2+}(aq) + Zn(s) \rightleftharpoons Zn^{2+}(aq) + Pb(s)\)



    \(Cu^{2+}(aq) + Zn(s) \rightleftharpoons Zn^{2+}(aq) + Cu(s)\)



    \(Pb^{2+}(aq) + 2Ag(s) \rightleftharpoons 2Ag^+(aq) + Pb(s)\)



    \(2H^+(aq) + Zn(s) \rightleftharpoons Zn^{2+}(aq) + H_2(g)\)



    Note that in the above calculations you do not multiply the standard reduction potentials by a factor even if a coefficient appears in the balanced redox reaction (e.g., 2 in the third reaction above). You also do not change the sign of the standard reduction potential when dealing the oxidation half reactions.  

    So, what about when we aren’t under standard conditions?

    The specificity of standard reduction potentials for standard conditions (i.e., only for solution concentrations of 1.0 M) may make them seem relatively useless. The Nernst equation allows us to use nonstandard potentials, which can be measured, for systems under nonstandard conditions.

    \[E_{cell} = E^\circ_{cell} - \dfrac{RT}{nF} \ln Q \nonumber\]

    Using a temperature of 25C, the constants R and F, and converting from a natural log to base 10 log scale, we can rewrite the Nernst equation as:

    \[E_{cell} = E^\circ_{cell} - \dfrac{0.05916V}{n} \log Q \nonumber\]

    Where n is the number of moles of electrons transferred in the overall redox reaction.

    1. What is Q? What form does it take?




    1. Write the expression of Q for the reaction 2Ag + + Cd ⇌ 2 Ag + Cd 2+




    1. Find cell for the reaction.



    1. How many moles of electrons are transferred?





    1. If the cell potential was measured to be 1.137 V and the silver ion concentration was 1.3 × 10-5 M, what is the concentration of Cd2+?



    A note about line notation

    Line notation is a convention used in electrochemistry to concisely represent electrochemical cells. Anytime there are two phases in contact with each other (e.g., between solid metal and a solution containing dissolved ions), a vertical line is used. The double line represents a salt bridge, which serves to maintain electroneutrality between the two sides of the electrochemical cell. For example, the reaction above (2Ag + + Cd ⇌ 2 Ag + Cd 2+) is represented in line notation (with nitrate as a spectator ion):

    \[Cd(s) \mid Cd(NO_3)_2(aq) \mid\mid AgNO_3(aq) \mid Ag(s)\nonumber\]

    Sometimes line notation can be quite detailed and contain information like concentration. For example:

    \[Cd(s) \mid Cd(NO_3)_2(1.2 \times 10^{-5}M) \mid\mid AgNO_3(2.7 \times 10^{-3}M) \mid Ag(s)\nonumber\]

    Can you calculate the potential of the cell represented by the line notation above?




    Reference Electrodes

    To make meaningful potentiometric measurements, we need to know exactly what’s going on with at least half the electrochemical cell. Analytical chemists do this by using reference electrodes that act as half of an electrochemical cell with a consistent potential. This allows the overall measured cell potential to be used as a way to measure an unknown concentration. Common reference electrodes are summarized in the table below.

    Reference electrode

    Line notation

    Half reaction

    E° (V)

    E (V)

    Standard Hydrogen (SHE)

    \(Pt \mid H_2(1\: bar) \mid H^+(1.0M)\) \(2H^+ + 2e^- \rightleftharpoons H_2(g)\)



    Silver-silver chloride (Ag/AgCl)

    \(Ag \mid AgCl \mid KCl(\textit{aq,sat'd})\) \(AgCl + e^- \rightleftharpoons Ag^+ + Cl^- \)



    Saturated Calomel (SCE)

    \(Pt \mid Hg \mid Hg_2Cl_2 \mid KCl (\textit{ag,sat'd})\) \(Hg_2Cl_2 + 2e^- \rightleftharpoons 2Hg + 2Cl^- \)



    1. In most cases the standard hydrogen electrode (SHE) is not physically used, but serves as a theoretical benchmark. All standard reduction potentials (E°) are given relative to the SHE. Why do you think SHE’s are not used?




    1. Note that in the line notation of both the SHE and SCE include platinum metal, but Pt does not participate in the electrode’s half reaction. Why do you think it is necessary to use a platinum wire as the electrode?




    1. Why is E° different than E for the Ag/AgCl electrode and SCE, but both are 0 V for the SHE?






    1. Because potential readings depend on the identity of the reference electrode, it is customary to define any reported potential as “vs.” a specified reference electrode. The potentials of the Ag/AgCl and SCE given are relative to a SHE. Use this information to convert potential readings between reference electrode scales.

    vs. SHE

    vs. Ag/AgCl

    vs. SCE

    0.200 V




    0.346 V




    -0.234 V


    -0.053 V


    -0.143 V








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