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Lab 1: Cyclic Voltammetry

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  • Cyclic voltammetry (CV) is a technique used to study electrochemical reaction mechanisms that give rise to electroanalytical current signals. The method involves linearly varying an electrode potential between two limits at a specific rate while monitoring the current that develops in an electrochemical cell under conditions where voltage is in excess of that predicted by the Nernst equation.¹ Although CV is best at providing qualitative information about electrode reaction mechanisms, several quantitative properties of the charge transfer reaction can also be determined.


    Cyclic voltammetry involves applying a voltage to an electrode immersed in an electrolyte solution, and seeing how the system responds. In CV, a linear sweeping voltage is applied to an aqueous solution containing the compound of interest. The voltage is initially given by Equation \(\ref{1.2}\) (see below). After the voltage reaches a certain maximum value, the potential is reversed and the sign of vt reverses and Ei becomes the maximum voltage, \(E_\lambda \). The process can then be repeated in a periodic, or cyclic manner. The voltage after the potential sweep direction is switched is given by Equation \(\ref{1.3}\).

    As an important tool for studying mechanisms and rates of oxidation and reduction processes, CV provides the capability for generating a species during the forward scan and then probing its fate with the reverse scan or subsequent cycles, all within seconds. The unique aspect of cyclic voltammetry is the three electrodes used, which consist of a working electrode, a reference electrode, and a counter electrode. The working electrode can be seen as a medium whose reductive or oxidative power can be externally adjusted by the magnitude of the applied potential--as the potential is increased or decreased linearly versus time, it becomes a stronger oxidant or reductant, respectively. Therefore, the working electrode, which typically consists of a chemically inert conductive material such as platinum, acts as a donor or acceptor of electrons participating in the general electrode reaction,

    \[ O + ne^- \rightarrow R \]

    The reference electrode, typically AgCl or calomel, keeps the potential between itself and the working electrode constant. The potential is measured between the reference and working electrodes, and the current is measured between the working and counter electrodes. A counter electrode is employed to allow for accurate measurements to be made between the working and reference electrodes. The counter electrode's role is essentially to ensure that current does not run through the reference electrode, since such a flow would change the reference electrodes potential. A voltage sweep from Ei to Ef is produced using a signal generator, and the voltage is applied to the working electrode using a potentiostat. By sweeping the voltage slowly, information may be extracted from a graph of potential versus current going through the sample. Polarography utilizes this method of analysis where the limited current arising from a redox process in the solution during the sweep is used to quantitatively determine the concentration of species that are electrochemically active in solution.

    CV differs from polarography in two important ways. Firstly, the working electrode at which the reactions of interest occur has a constant area, not one which changes with time as in classical polarography. This electrode may be a solid such as graphite or platinum with a small surface area, or a stationary or hanging mercury drop. The latter type of electrode may have its surface renewed periodically. The second difference is that the potential of the working electrode is scanned rapidly over a wide potential range and then returned to its initial value using an applied potential signal which varies linearly with time between the initial value and the final value at the limit of the forward scan. Normally, this technique is applied so that currents due to reduction processes are observed during the forward scan, and those due to oxidation on the reverse scan. A typical potential against time profile applied to the working electrode and the resulting cyclic voltammogram are shown in Figure 1.1.

    Figure 1.1. (Left) Potential versus time program for cyclic voltammetry showing the forward and reversed linear potential ramp and (Right) Cyclic Voltammogram of measured current versus applied potential.

    Let us consider the response of the electron transfer reaction

    \[ A + ne^- \rightleftharpoons B^- \label{1.1} \]

    to the application of an electrode potential which is varying linearly with time. The electrode potential is given by the equation

    \[ E = E_i - vt \label{1.2} \]

    where Ei is the initial potential, v is the potential sweep rate (in volts s-1) and t is time after start of voltage sweep. At some time, λ, the direction of the potential sweep is switched, and the equation describing the electrode potential becomes

    \[ E = E_\lambda + v(t - \lambda) \label{1.3} \]

    where Eλ is the value of E at the switching point. Considering the initial sweep in the negative direction where reduction reactions are expected, it is clear that, if the sweep rate is sufficiently slow, the current against potential curve approaches that obtained by steady state measurements such as d.c. polarography. However, as v is increased, a peak develops on the i-E curve which becomes increasingly prominent (Figure 1.2). The peak is produced from the combined effects of high mass transfer rates in the non steady state followed by progressive depletion of the reactant in the diffusion layer.


    Figure 1.2. Effect of potential sweep rate on the PE curves in a linear potential sweep experiment adjacent to the electrode. It should be noted that since E is a linear function of t, that the potential axis is also a time axis.

    In order to relate the observed current to the reactant concentration cA, one must know how cA varies with distance from the electrode, x, and with time, t. This variation is described by Fick's Law for mass transfer by diffusion:

    \[ \dfrac{\partial c_A}{\partial t} = D_A \dfrac{\partial^2c_A}{\partial x^2} \label{1.4} \]

    where DA is the diffusion coefficient of reactant A (units: cm2/s). Solution of this second order partial differential equation requires specification of boundary and initial conditions, and is described in textbooks on electroanalytical chemistry (1,2). If the electron transfer reaction is sufficiently fast in reaction (1) to maintain a Nernstian equilibrium at the electrode surface (i.e., the reversible case), then the peak current, ip (in amperes), for a negative sweep is given by

    \[ i_p = (2.69 \times 10^5) \; n^{3/2} \; SD_A^{1/2} \; v^{1/2} \; C_A, \label{1.5} \]

    where n is equal to the number of electrons gained in the reduction, S is the surface area of the working electrode in cm2, \(D_A\) is the diffusion coefficient, v is the sweep rate, and \(C_A\) is the molar concentration (mole/cm3) of A in the bulk solution.

    The peak potential is independent of sweep rate and is related to the polarographic half wave potential by

    \[ E_p = E_{1/2} - \dfrac{0.0285}{n} \label{1.6} \]

    Furthermore, the shape of the peak is defined by the potential difference between the peak and the position when the current is one half of that at the peak:

    \[ E_p = E_{p/2} - \dfrac{0.0565}{n} \label{1.7} \]

    Thus, using the measured values of ip, Ep, and Ep/2, one can determine n and DA for a given electrode area and sweep rate. If the electrode reaction is slow so that the surface concentrations of \(A\) and \(B\0 are no longer related by the Nernst equation at a given sweep rate, then the peak characteristics change such that

    \[ E_p - E_{p/2} > \dfrac{0.0565}{n} \label{1.8} \]

    and the peak potential now depends on sweep rate. Using the appropriate boundary conditions to describe the rate of reaction (Equation \(\ref{1}\)) in the forward direction, (Equation \(\ref{1.4}\)) may be solved to obtain expressions for \(i_p\) and \(E_p\) (Equation \ref{1.2}) which are now much more complicated.

    After the direction of the potential sweep is reversed, a second current peak is observed corresponding to oxidation of the product B. When reaction (Equation \(\ref{1.1}\)) is reversible, implying that \(B\) is stable, the height of this peak is equal to that observed on reduction, but with the current flowing in the opposite direction. The method of estimating peak currents is illustrated in Figure 1.3, where \(E_{pc}\) is the peak potential in the cathodic sweep.

    Figure 1.3: Cyclic voltammogram

    The cyclic voltammogram in the above figure for the process A ± ne ↔ B assumes that only A is initially present in the solution. \(\Delta{E}_p\) is then defined as

    \[ \Delta{E}_p = |E_{pc} - E_{pa}|. \label{1.9}\]

    When the electrode process is reversible, ΔEp = 0.059/n and it is independent of sweep speed. As v is increased to the stage at which Nernstian equilibrium for reaction (1) cannot be maintained, ΔEp increases with increasing sweep rate, and the shape and position of the peaks depend on both v and the kinetic parameters of the electrode reaction.

    Reaction Mechanisms

    One of the major uses of cyclic voltammetry is in the rapid qualitative elucidation of electrode reaction mechanisms. Organic molecules often undergo a rapid chemical reaction with the solvent or some other constituent of the solution after the electron transfer process. The product of this reaction is usually reduced or oxidized at a much different potential. The resulting reaction scheme, referred to as the ECE mechanism, can be written

    \[\begin{align} A + ne^- &\rightarrow B^{n-} \label{1.10} \\[5pt] B + Z &\rightarrow C \label{1.11} \\[5pt] C + me^- &\rightarrow D^{m-} \label{1.12} \end{align}\]

    where \(Z\) is the solvent or some other species. The first and third reactions are labeled \(E\) since they involve the electrode, and the second step (or any other chemical step) is labeled \(C\). Hence, the above three-step mechanism is referred to as the ECE reaction mechanism. It is possible to garner information about the (non-electrode-dependent) rate constant for step 2 (Equation \(\ref{1.12}\)) via cyclic voltammetry.

    The standard potential for reaction \(\ref{1.12}\) is generally different from that for reaction \ref{1.10}. A typical current-potential curve for such a system is shown in Figure 1.4. The current on the reverse sweep will depend on the sweep rate and the rate constant for reaction \ref{1.11}, which is assumed to take place under pseudo first order conditions (\(c_Z \gg c_B\)). For very fast sweep rates, very little \(B\) will react to form \(C\), and the voltammogram will have the same appearance as the reversible case, with reduction and oxidation peaks at I and II, respectively. As \(v\) is decreased, peak II diminishes more rapidly and peak I less rapidly than the usual v1/2 dependence would predict because the chemical step removing species B becomes important and peak I has a contribution from reaction \(\ref{1.12}\). In addition, a peak develops at III due to oxidation of D.


    Figure 1.4. Cyclic voltammogram at an intermediate sweep rate for a system with an ECE mechanism.

    It should also be noted that the current-potential curves on the second and successive sweeps are not the same as that observed on the first. At very slow sweep rates, peak II disappears completely and peak I then corresponds to the process

    \[ A + (m + n)e^- \rightarrow D \label{1.13} \]

    The variation of peak I with sweep rate is shown in Figure 1.5 for the case that \(n = m = 2\).


    Figure 1.5. Variation of measured current for peak I from Figure 1.4 vs. potential scan rate, \(v\).

    The rate constant k for the chemical reaction \(\ref{1.13}\)) can be obtained from analysis of data obtained in the transition region of Figure 1.5 or from the ratio of peaks I and II at intermediate sweep rates.

    Cyclic voltammetry can be applied to the analysis of many other reaction mechanisms including those with dimerization of the product of electron transfer, with preceding chemical steps, catalytic processes, etc.

    The Apparatus

    The instrument used for this experiment is the BAS Epsilon potentiostat. It is controlled from a computer running Windows. A number of different electrochemical techniques are available in the Epsilon software, including cyclic voltammetry (current vs. potential for a linear potential sweep), chronoamperometry, time base, bulk electrolysis (current vs. time at a constant potential), and chronpotentiometry (potential vs. time at a fixed current).


    Figure 1.5.5: UCD cyclic voltammetry Instrument (including potentiostat)

    To begin an experiment, make sure that the Epsilon unit is turned on, and doubleclick the Epsilon icon on the desktop. Select New from the File menu or click the New icon. This will generate a menu that lists the available techniques. (This list can also be generated by selecting Select NEW Experiment from the Experiment menu or by using the F2 key.)


    Figure 1.6. Computer Program

    Highlight Cyclic Voltammetry (CV), and click Select to confirm the selection. An experiment window containing an empty axis set is displayed (Figure 1.7), and the appropriate parameters are set in the various dialog boxes.


    Figure 1.7: Computer Program

    The potential limits and the scan rate for CV are set using the Change Parameters dialog box (Figure 1.8) in either the Experiment menu or the pop-up menu (the pop-up menu is accessed with the right mouse button).


    Figure 1.8. Change Parameters dialog box for cyclic voltammetry.

    1. Potential values are entered in mV, and the Scan Rate in mV/s.
    2. If the Apply Open Circuit Potential for Initial E box is checked, then the open circuit potential will automatically be measured and used as the Initial Potential.
    3. When the experiment is started, the cell is held at the Initial Potential for the number of seconds defined by the Quiet Time.
    4. There are two gain stages for the current-to-voltage converter. The default values of these stages that are used for a given current Full Scale value are determined by the software. However, they can be adjusted manually using the Filter / F.S. dialog box. This dialog box is also used to change the analog Noise Filter Value settings from the default values set by the software. The Full Scale value should always be at the 10 mA/V setting at the start of the experiment, and then adjusted to a more convenient range depending on the maximum current observed in the experiment.
    5. The default condition of the cell is that the cell is On (i.e., the electronics are connected to the electrodes) during the experiment, and is Off between experiments. THIS OPTION SHOULD NOT BE CHANGED SINCE CONNECTING OR DISCONNECTING THE ELECTRODES WHEN THE CELL IS ON CAN RESULT IN DAMAGE TO THE POTENTIOSTAT, THE CELL, AND/OR THE USER!
    6. A series of identical experiments on the same cell can be programmed using the MR (Multi-Run) option.
    7. Clicking the IR-COMP button activates the iR compensation option (compensates for the drop in voltage due to the resistance of the solution).
    8. Clicking Exit will exit the dialog box without saving any changes made to the parameter values. Any changes can be saved by clicking Apply before exiting.
    9. Range of allowed parameter values:
    • Potential = -3275 - +3275 mV
    • Scan Rate = 1 - 10,000 mV/s (also see below)
    • Quiet Time = 0 - 100 s
    • The # of Segments is limited by the total number of data points that can be stored (32,000) (note that in this initial version, the potential resolution of the current measurement is fixed at 1 mV).
    1. Once the parameters have been set, the experiment can be started by clicking Run (either in this dialog box, in the Experiment menu, in the pop-up menu, on the Tool Bar, or using the F5 key).

    The Experiment

    I) Reduction of the Ferricyanide Anion

    The properties of the reaction

    \[ Fe(CN)_6^{-3} + e^{-} \rightarrow Fe(CN)_6^{-4} \label{1.14}\]

    are investigated in this part of the experiment. The solution used consists of 4 x 10-3 M K3Fe(CN)6 with 1M KNO3 as background electrolyte. The working electrode is a Pt disc and the reference electrode is the Ag/AgCl system. The Pt electrode should be cleaned by gently rubbing the electrode surface on fine damp paper cloth. Usually the electrode has to be cleaned when the cyclic voltammogram is no longer symmetrical and peaks are getting flat. Sometimes you have to repeat the cleaning procedure. The same method of cleaning applies to the glassy carbon electrode used in part II. The cleaning operation must be carried out under the supervision of the teaching assistant.

    The next step is to determine the current which flows in the presence of the ferricyanide reactant. The cell is assembled and filled with the solution of 4 mM Fe(CN)6 and 1M KNO3 so that the ends of the electrodes are immersed. The solution in the cell is now deoxygenated by purging with argon gas for 10 minutes. While deoxygenation is taking place, the scan parameters can be set. The working (Pt) electrode should be disconnected during this procedure. Use the following initial settings:

    Initial Potential +700 mV
    Switching Potential -200 mV
    Final Potential +700 mV
    # of Segments 2
    Scan Rate 250 mV/s
    Quiet Time 10 seconds
    Full Scale 10 mA/V

    When deoxygenation is complete, connect the electrodes:

    • Reference electrode (SCE) to the WHITE wire;
    • Working electrode (Pt) to the BLACK wire;
    • Counter electrode (platinum wire) to the RED wire.

    To begin the experiment, click Run (either in this dialog box, in the Experiment menu, in the pop-up menu, on the Tool Bar, or using the F5 key).

    Repeat the experiment at scan rates of 160, 100, 50, and 20 mV/s using the same solution. Between each scan, initial conditions at the electrode surface are restored by stirring the solution for 10 seconds. Be sure no gas bubbles remain on the electrode surface. Wait about 1 minute after stirring, and then record the current-potential curve at a new scan rate using the same procedure described above. Note: You do not need to recalibrate the X Y recorder. The same sensitivity is used to record all curves at the given concentration on one piece of graph paper.

    Dilute the working solution containing 4 mM ferricyanide with the 1 M KNO3 solution to obtain a 2mM ferricyanide solution. Fill the cell with this solution, degas it and record the current-potential curves at the above scan rates. Compare the characteristics of the current-voltage curves for this solution with that for the more concentrated solution.

    Record the characteristics of the current-voltage curve by measuring the values of ipc, ipa, Epc, Epa, and Epc/2. Tabulate the measured data and plot ipc and ipa against v1/2. Determine the diffusion coefficient of the ferricyanide ion in the given solutions (eq. 5). To do this you need to measure the surface area of the working electrode. Calculate E1/2 according to eq. (6). Test the reversibility of the reaction using the values of Epc, Epc/2, Epa and Epa/2 (eq. 7 and 8).

    The next step is to determine the current which flows in the absence of the ferricyanide reactant. The cell is assembled and filled with 1M KN03. Degas the solution and record the current-voltage curve at a 10 mV/s scan rate using higher sensitivity. Does any current flow through the cell? Explain.

    The effect of the supporting electrolyte on the shape of the current-voltage curve is examined by studying a 4mM ferricyanide solution in 1M Na2S04. Prepare the solution for the experiment as described above but only use scan rates of 250, 100, and 50 mV/s. Compare the current-voltage curves. Discuss the differences. (Hint: note the difference in the charge on the ions of the supporting electrolyte. Ref. 2, Chapter 1).

    When you are finished, rinse all three electrodes with deionized water and immerse them in the cell filled with deionized water.

    II) The Oxidation of Acetaminophen

    Acetaminophen (AAPH), the active ingredient in Tylenol, undergoes a complex oxidation process in an aqueous solution. This process is pH-dependent. When AAPH is oxidized at a carbon electrode, N-acetyl-p-quinoneimine (NAPQ) is formed:


    The product can be protonated, and then undergo the elimination of acetamide to form benzoquinone (BQ):


    The rate of the first reaction is obviously pH-dependent. Finally, benzoquinone may be reduced to hydroquinone:


    The extent to which NAPQ is converted to BQ depends on the rate of reaction #16. This system will be investigated as a function of pH to obtain evidence for the above reaction mechanism and to determine the best conditions for analyzing for AAPH.

    The working electrode in this experiment is glassy carbon. It must be prepared by mechanical polishing with your laboratory instructor's help. The reference electrode is the Ag/AgCl system. Make up three solutions of AAPH by pipetting 5 mL of the stock solution stored in the refrigerator into each of three 100 mL volumetric flasks. These solutions are made up to the mark with solutions of different pH, either buffer at pH 2.2, buffer at pH 6, or 1.8 M sulfuric acid. The final concentration of AAPH should be 3.5 mM. Fill the cell with the solution of highest pH and begin the deoxygenation procedure with the electrodes disconnected. Set the initial potential at 0.0 V. The positive limit should be set at +1.0 V (switching potential), the negative limit at -0.2 V (final potential), and the direction switch to POS. Once deoxygenation is complete, record current-voltage curves at scan rates of 40, 100, and 250 mV/s. Repeat this experiment for the two remaining AAPH solutions.

    Identify the peaks on each curve by noting how the peak current varies with v1/2. Which buffer system is appropriate for the quantitative determination of AAPH and why?

    Accurately weigh a Tylenol tablet, then crush it into a fine powder using a mortar and pestle, then drop that into a 250 mL volumetric flask, dissolve it in a few mL of the buffer chosen from above (if pH 2.2 was chosen as the most appropriate for determining AAPH, then that buffer solution would be used), and then make the solution up to the mark with the same buffer solution that was chosen in the previous step as the most appropriate for quantitative determination of AAPH. Pipette a 5 mL aliquot of this solution into a 50 mL volumetric flask, and fill this up to the mark with the chosen buffer. Deoxygenate the solution and record the current-voltage curve for this system at one of the three scan rates used above. Compare the peak height for the AAPH oxidation step with that found for the standard solution at the same pH and scan rate.

    In order to accurately determine the AAPH concentration in Tylenol, current-voltage curves for two more standard solutions which cover the concentration range of the unknown must be recorded. Decide what the appropriate concentration range is, prepare the necessary solutions, and obtain the corresponding curves. Measure the peak heights and plot peak height against concentration for the scan rate used (calibration curve). Report the percent AAPH in the original Tylenol tablet.


    1. Heineman, W. R.; Kissinger, P. T. Large-Amplitude Controled-Potential Techniques. In Laboratory Techniques in Electroanalytical Chemistry; Kissinger, P. T., Heineman, W. R., Eds.; Monographs in Electroanalytical Chemistry and Electrochemistry; Marcel Dekker: New York, NY, 1984; pp 51-127.
    2. Bard, A. J.; Faulkner, L. R. Electrochemical Methods; Wiley: New York, NY, 1980; Chapter 6.
    3. Compton, R. G.; Banks, C. E. Understanding Voltammetry; World Scientific: Hackensack, NJ, 2007.