It is no secret that the need to switch from petroleum-based, CO2-producing energy sources to alternatives, in particular those based on electrochemical principles, is ever more pressing. For that reason alone, this topic is an important goal of general undergraduate instruction in chemistry. One of the most important experimental methods to investigate redoxsystems for using as rechargeables is cyclic voltammetry (CV).

The experimental CV results are sometimes difficult to interpret, given the plethora of possible processes (Table 1), and also the need to consider the diffusion of the reacting species to and from the electrode surface on which the reaction takes place.

Nevertheless, there are systems that lend themselves as subjects for an introduction to the use of CV in understanding electrochemical processes.

Therefore, in this paper we present and interpret three relative simple but instructive systems to understand the principles of CV, a method in electrochemistry for determining electrode reaction mechanisms, standard electron transfer rate constants, and also diffusion coefficients. We propose and describe procedures that show the behavior of reversible, quasi-reversible and irreversible systems (the first two are necessary condition for rechargeables), restricting the discussion to only the first case in Table 1, that of simple electron transfer. On base of the theory of Matsuda and Ayabe [1] and Nicholson and Shain [2] the experimenter can calculate the electron transfer rate constant, the decisive factor describing the electrode process, from the CV in a very simple way.

PowerPoint Slide

Larger image(png format)

View current table in a new window

View next table

Many texts on CV are available [see, e.g., 3-10]. Here we will only present a basic outline that is necessary to understand the presented experiments.

A CV is obtained by measuring the current between the working and the counter electrode as a function of the potential (normalized to the potential of the reference electrode). To do this, the experimenter uses a three-electron setup and varies the potential of an electrode (the “working” electrode), which is immersed in an unstirred solution, and measures the resulting current.

A triangular potential sweeps the potential of the working electrode between the starting potential to the switching potential and back again. The scan rate v (in mV/s) is an important parameter, as will be shown below.

The current flows in or out of the working electrode to or from a counter electrode. The potential of the working electrode is controlled versus a reference electrode, e.g., a saturated calomel or a silver/silver chloride electrode. The reference electrode passes no current.

All these requirements can be fulfilled by a potentiostat.

The rate of an electron transferred between the electrode and the solution depends on the potential.

Here are the main points to consider when starting out with CV:

•  According to the IUPAC recommendation anodic peaks point upward, cathodic downward.

•  The standard redox potential E0 can be calculated

where EP are the peak potentials.

•  If the experimenter wants to determine the standard potential E0 the scan rate should be slowed down to minimize ΔEp. However, there is a lower limit to the scan rate: This is set by the ability to maintain convection-free conditions. This means that scan rates lower than 1 mV/s are often not useful. (Calculation of E0 requires the potential of the reference electrode to compare E0 with tabulated values which are relative to the normal hydrogen electrode NHE).

•  If the current peaks appear to be sliding apart as a function of scan rate, the process is quasi-reversible (or, in an extreme case, irreversible).

•  If an organic, high resistance electrolyte is used, the applied potential is not identical with the theoretical potential between the working and reference electrodes. The resistance of the solution causes a so-called IR-drop. This means, that the potential of the working electrode drifts away with I∙R and is unreliable. For this reason a highly conducting supporting electrolyte, which is not electroactive in the potential range being studied, is added to the solution. In the experiments described here this is not necessary because the electrolyte itself is highly conductive.

Two qualities, parameterized by the standard electron rate constant k0 and the mass transport to or from the working electrode mtransport, determine the CV. For a derivation of these parameters we refer readers to, e.g., Ref [10]. We limit ourselves here to a qualitative description that will allow an initial interpretation of the results.

The electron transfer rate constant for a reduction and oxidation process is a function of the applied potential and can be described as

and

where k0 is the standard electron-transfer rate constant in cm/s at the standard potential E0, a is the so-called transfer coefficient (a measure of the symmetry of the activation energy barrier for the oxidation and reduction processes), n is the number of the transferred electrons, F is the Faraday constant, E the applied potential, R the ideal gas constant, and T the absolute temperature.

The exponential dependence of k on the potential E results in a steep rise in the current. This leads to a depletion of the concentration of the corresponding species at the electrode. Now, diffusion is the only process in an unstirred solution by which the reactant can move to the electrode surface. (Migration is mainly controlled by the electrolyte and not by the electroactive substances). As diffusion is slow (diffusion coefficients are in the range of 10-5 cm2/s) the current does not increase exponentially with E as (2) and (3) indicate but decreases after the species reacted on the electrode surface and so a depletion layer results.

The mass transport is given by eq. (5):

In each case the slowest process determines the electrochemical behavior:

•  If k0 >> mtrans then the electrode process is reversible and diffusion controlled.

•  In the intermediate, the so-called quasi-reversible case, diffusion and electron transfer are in the same order (k0 ≈ mtrans).

•  The process will be rate-determined if the mass transport is faster than the electron transfer (k0 0.67·v1/2 and with v = 0.06 V/s k0 > 0.16 cm/s.

The data in Figure 2 are plotted according to the Randles-Sevcik (eq. 8). The graph yields the linear relationship indicative of Ip ˜ v1/2.

View next figure

The Randles-Sevcik plot verifies the linearity between the peak current and the square root of the scan rate indicating reversible electron transfer (the analogous behavior of the cathodic peak is not shown here).

Chemicals and procedure

4 mg potassium hexacyanoferrate (III) (SigmaAldrich, 31254) in 10 mL 0.1 M aqueous Na2SO4 (corresponds to 1/1000 mol potassium hexacyanoferrate (III)).

Five different scan rates (2, 5, 10, 20, 50 mV/s), scan range: 0.5 V à -0.3 V à 0.5 V

Hazard: No hazardous chemicals.

The currents peaks for the oxidation and reduction slide apart with increasing scan rate from 75 mV at 2 mV/s to 160 mV at 50 mV/s.

Using the theory of Nicholson and Shain we find that the values of Λ change from 1.51 to 0.185 with increasing scan rate.

In Table 4 we estimate k0 using the Nicholson and Shain method. The average value is 2.2∙10-3 cm/s (for D ≈ 10-5 cm2/s).

PowerPoint Slide

Larger image(png format)

View current table in a new window

View previous table

View next figure

Chemicals and procedure

0.32 mg N,N,N’,N’-tetramethyl-p-phenylenediamine (TMPD, Wurster’s Blue, SigmaAldrich, T 7394) in 10 mL 2M H2SO4 (corresponds to 1/500 mol TMPD). Scan rate 20 – 50 mV/s.

Hazard: No hazardous chemicals.

Electron transfer in TMPD can be quasi-reversible or irreversible depending on the electrode used (Figure 4). Here we measured a peak current drift for Pt of 110 mV to 160 mV (scan rate 20 - 50 mV/s) yielding k0 ≈ 10-3cm/s, whereas for graphite it is between 250 mV and 350 mV (scan rate 20 - 50 mV/s) yielding ko