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Cyclic Voltammetry: Principles, Setup, and Applications


Duck-shaped cyclic voltammogram
Duck-shaped cyclic voltammogram

Cyclic voltammetry is an electrochemical technique for measuring the current response of a redox active solution to a linearly cycled potential sweep between two or more set values. It is a useful method for quickly determining information about the thermodynamics of redox processes, the energy levels of the analyte and the kinetics of electronic-transfer reactions.

Like other voltammetric methods methods, cyclic voltammetry uses a three electrode system consisting of a working electrode, reference electrode, and counter electrode.

To perform cyclic voltammetry, the electrolyte solution is first added to an electrochemical cell along with a reference solution and the three electrodes. A potentiostat is then used to linearly sweep the potential between the working and reference electrodes until it reaches a preset limit, at which point it is swept back in the opposite direction.

This process is repeated multiple times during a scan and the changing current between the working and counter probes is measured by the device in real time. The result is a characteristic duck-shaped plot known as a cyclic voltammogram.

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Cyclic Voltammetry Theory


Cyclic voltammetry is a sophisticated potentiometric and voltammetric method. During a scan, the chemical either loses an electron (oxidation) or gains an electron (reduction) depending on the direction of the ramping potential.

The Potentiometry Principle

Potentiometry is a way of measuring the electrical potential of an electrochemical cell under static conditions (i.e. no current flow).

For a general reduction or oxidation (redox) reaction the standard potential is related to the concentration of the reactants (A) and products (B) at the electrode/solution interface according to the Nernst equation:

Nernst equation
Nernst equation

where E is the electrode potential, E0′ is the formal potential, R is the gas constant (8.3145 J·K-1·mol-1), T is temperature, n is the number of moles of electrons involved and F is the Faraday constant (96,485 C·mol-1).

The term [B]b/[A]a represents the ratio products to reactants, raised to their respective stoichiometric powers. This can be used in place of an activity term when the concentration is sufficiently low (< 0.1 mol·dm˗3).

Under standard conditions of temperature and pressure, the Nernst equation can be written as:

Nernst equation

An electrochemical reaction is reversible in nature when the kinetics of electron transfer are sufficiently fast such that the concentration of oxidised species and the concentration of reduced species is in equilibrium.

Introduction to Voltammetry

In the general sense, voltammetry is any technique where the current is measured while the potential between two electrodes is varied. Voltammetric methods include cyclic voltammetry, linear sweep voltammetry, and a number of variations such as staircase voltammetry, squarewave voltammetry and fast-scan cyclic voltammetry.

In voltammetry experiments, the current generated is the result of electron transfer between the redox species and the electrodes. This is carried through the solution by the diffusion and migration of ions.

Three electrode cell system
Three electrode cell as used in cyclic voltammetry

Although in principle voltammetry only requires two electrodes, in practice it is very difficult to maintain a constant potential while also passing current to counteract the redox events at the working electrode.

As a result, a three-electrode (working electrode, counter electrode, and reference electrode) cell is often used to separate the role of referencing the potential applied and balancing the current produced.

To measure and control the potential difference applied, as required for cyclic voltammetry, the potential of the working electrode is varied while the potential of reference electrode remains fixed by a electrochemical redox reaction with a well-defined value .

To keep the potential fixed, the reference electrode must contain constant concentrations of each component of the reaction, such as a silver wire and a saturated solution of silver ions.

Importantly, no current passes between the reference and working electrodes. The current observed at the working electrode is completely balanced by the current passing at the counter electrode, which has a much larger surface area.

The electron transfer between the redox species at the working electrode and counter electrode generates current that is carried through the solution by the diffusion and migration of ions. This forms a capacitive electrical double layer at the surface of the electrode called the diffuse double layer (DDL). The DDL is composed of ions and orientated electric dipoles that serve to counteract the charge on the electrode.

The measured current response is dependent on the concentration of the redox species (the analyte) at the working electrode surface, and is described by a combination of Faraday’s law and Fick’s first law of diffusion:

Fick’s first law of diffusion

where id is the diffusion-limited current, A is the electrode area, D0 is the diffusion coefficient of the analyte and (∂C0/∂x)0 is concentration gradient at the electrode surface.

The product of the diffusion coefficient and concentration gradient can be thought of as the molar flux (mol·cm-2·s-1) of analyte to the electrode surface.

Inert ions are added to the electrochemical solution in molar excess to the analyte in order to provide enough ionic strength to the solution for it obey the Nernst equation. The excess of electrolyte decreases the thickness of the diffuse double layer so that the applied potential decreases to a negligible level within nanometers of the working electrode surface. The result is that the current response at the electrode surface is well defined.

Cyclic voltammograms

The ‘duck-shaped’ plot generated by cyclic voltammetry is called a cyclic voltammogram. An example is displayed in Figure 1.

Cyclic voltammogram for an electrochemically-reversible one-electron redox process
Figure 1. Cyclic voltammogram for an electrochemically-reversible one-electron redox process

In Figure 1, the scan starts at -0.4V and sweeps forward to more positive, oxidative potentials. Initially the potential is not sufficient to oxidise the analyte (Figure 1, a).

As the onset (Eonset) of oxidation is reached the current exponentially increases (b) as the analyte is being oxidised at the working electrode surface. Here the process is under electrochemical control with the current linearly increasing with increasing voltage with a constant concentration gradient of the analyte near the electrode surface within the diffuse double layer.

The current response decreases from linearity as the analyte is depleted and the diffuse double layer grows in size. The current reaches peak maximum at point c (anodic peak current (ipa) for oxidation at the anodic peak potential (Epa). The process is now under mixed control: more positive potentials cause an increase in current that is offset by a decreasing flux of analyte from further and further distance from the electrode surface.

From this point the current is limited by the mass transport of analyte from the bulk to the DDL interface, which is slow on the electrochemical timescale. This results in a decrease in current (d) as the potentials are scanned more positive until a steady-state is reached where further increases in potential no longer has an effect.

Scan reversal to negative potentials (reductive scan) continues to oxidise the analyte until the applied potential reaches the value where the oxidised analyte which has accumulated at the electrode surface can be re-reduced (e).

The process for reduction mirrors that for the oxidation, only with an opposite scan direction and a cathodic peak (ipc) at the cathodic peak potential (Epc) (f). The anodic and cathodic peak currents should be of equal magnitude but with opposite sign, provided that the process is reversible (and if the cathodic peak is measured relative to the base line after the anodic peak).

The Randles-Sevcik equation

The peak current, ip, of the reversible redox process is described by the Randles-Sevcik equation.[1]

At 298 K, the Randles-Sevcik equation is:

Randles-Sevcik equation

where n is the number of electrons, A the electrode area (cm2), C the concentration (mol·cm-3), D the diffusion coefficient (cm2·s-1), and v the potential scan rate (V·s-1).

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Experimental Setup


The experimental setup for cyclic voltammetry consists of an electrochemical cell containing five major components.

  1. The working electrode, where the compound of interest is reduced (Cn+ → C(n−1)+ ) or oxidised (Cn+ → C(n+1)+).
  2. The counter electrode, which completes the circuit with the potentiostat (see figure below).
  3. The reference electrode, used to measure the potential.
  4. The studied solution containing the chemical to be studied.
  5. The reference electrode solution (optional, see choice of reference electrode).

The potential of the studied solution is measured relative to the potential between the reference solution and reference electrode.

The Electrochemical Cell

An electrochemical cell is a device in which a chemical reaction generates an electrical response or, conversely, an electrical current is used to trigger a chemical reaction. The simplest possible electrochemical cell consists of two connected electrodes in an electrolyte solution. In cyclic voltammetry, three electrodes are used.

The physical setup of an electrochemical cell is relatively simple. The working and counter electrodes sit in an electrochemical solution, and the reference electrode sits in a separate tube within the cell containing the reference solution. The reference electrode tube should be approximately two thirds full - a syringe and needle can be used to add the solution.

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Degassing

In addition to having holes for each electrode, electrochemical glassware typically has gas intakes to allow for an inert gas (usually nitrogen or argon) to be bubbled through the solution to remove its oxygen. This process is known as degassing.

Degassing is important because molecular oxygen is electrochemically active, and if not removed will create unwanted redox processes. In addition, the products of this reaction (hydrogen peroxide) can also interact with the compound and further interfere with the results of the experiment.

Once the oxygen has been removed, it can be kept out with a continuous stream of inert gas.

Risk of contamination

When preparing an electrochemical cell, it is important to minimise the risk of any contamination with water, as water can form reactive species when reduced or oxidised. This can be done by heating the components in a glassware oven prior to use. 

Electrochemical solutions

The electrochemical solution used for cyclic voltammetry typically consists of three components.

  1. The compound of interest (10-3 – 10-5 M)
  2. An electrolyte (0.1 M)
  3. A solvent which dissolves both the compound of interest and the electrolyte

Solvent choice

The choice of solvent and electrolyte is dictated by the solubility of the studied chemical (so that it can be dissolved at the concentration needed) and the desired potential range.

Reference tables of the potential range of various solvent and electrolyte pairs are widely available [1] but these ranges are highly dependent on the purity and dryness of both the electrolyte and solvent. For the best results, choose a high purity solvent and electrolyte and oven dry all components before use.

Be aware when manually purifying and drying your components by standard procedures [2] that Grubbs purification apparatuses can add undesirable electroactive impurities [3].

Solvent reference table

A short table of potential ranges is listed below based on the values given by A.J. Bard and L.R. Faulkner [2]. Values are given relative to the Standard Calomel Electrode (SCE) (see choice of reference electrode).

Electrode Solvent Electrolyte Positive Range Relative to SCE / V Negative range Relative to SCE / V
Pt Water 1 M H2SO4 + 1.3 − 0.3
Pt Water pH 7 buffer + 1.0 − 0.7
Pt Water 1 M NaOH + 0.6 − 0.9
Hg Water 1 M H2SO4 + 0.3 − 1.1
Hg Water 1 M KCl + 0.0 − 1.9
Hg Water 1 M NaOH − 0.1 − 2.0
Hg Water 0.1 M Et4NOH − 0.1 − 2.4
C Water 1 M HClO4 + 1.5 − 0.2
C Water 0.1 M KCl + 1.0 − 1.3
Pt MeCN 0.1 M TBANF4 + 2.5 − 2.5
Pt DMF 0.1 M TBAP + 1.5 - 2.8
Pt Benzonitrile 0.1 M TBANF4 + 2.5 − 2.4
Pt THF 0.1 M TBAP + 1.4 − 3.1
Pt PC 0.1 M TEAP + 2.2 − 2.5
Pt CH2Cl2 0.1 M TBAP + 1.8 − 1.7
Pt SO2 0.1 M TBAP + 3.4 − 0.0
Pt NH3 0.1 M KI + 0.1 − 3.0

Insoluble chemicals

It is possible to study the electrical response of materials, like polymers, which cannot be sufficiently dissolved in standard electrochemical solvents. To do this, coat the working electrode with the material by depositing it with a solvent.

Due to the complex nature with which charges diffuse through the solid and the various distortions which occur within the deposited compound, the normal equations and mathematical proofs do not strictly apply under these circumstances. By approximating the onset potential as the redox potential of that process, however,  the technique still gives a good approximation of the energy levels for insoluble materials.

Internal standards

Internal standards, usually ferrocene (see below), are often used to calculate the value of the oxidation and reduction potentials. Internal standards are compounds which oxidise or reduce in solution, ideally somewhat independently of the system (although ferrocene does vary between solutions).

This oxidation or reduction provides a voltammogram which can be used to reference the position of the oxidation or reduction of the compound of interest.

It is common practice to study these standards immediately after the chemical of interest, using the same solutions. Recent reviews, however, suggest that it is better to always have the internal standard present in order to prevent changes in the position of the voltammograms [5]. This is particularly true for quasi reference electrodes where large shifts have been observed.

Three cell electrodes

Working and counter electrodes

The counter electrode and working electrode must be conductive so that charges can move to and from the solution, and they must not cause any chemical reaction in the solution. Inertness is usually achieved by making them out of unreactive material such as platinum.

A large counter electrode surface area makes sure that the measured current corresponds to the current flow between the working and counter electrode [5].

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Choice of reference electrode

Reference electrodes are designed so that an equilibrium is setup with known potential between the metal wire and the surrounding solution. In cyclic voltammetry, all electrochemical processes occur relative to this potential.

The reference electrode is setup in the cell so that it is in a circuit with the reference electrode and working electrode in opposing directions. In one direction, the working electrode goes from the solid state into the solution and the reference electrode goes from the solution to the solid state.

The consequence of this (along with Kirchhoff's voltage law and zero solution resistivity) is that the measured potential is zero when the working electrode potential is equal to the reference electrode potential.

The most common reference electrodes are the standard calomel electrode, the normal hydrogen electrode, the silver/silver chloride (Ag/AgCl) electrode in saturated potassium chloride and the Ag/Ag+ (0.01M, usually AgNO3) electrode in acetonitrile. Their standard reduction potentials are listed below.

It should be noted that the Ag/Ag+ electrode is usually setup with the same electrolyte solution that is used in the studied solution. This is to minimise the junction potentials (the potential between the reference solution and the studied solution). Take this into consideration when choosing your electrolyte and your solvent as well as when estimating the volume of solution that you require for your experiment.

Electrode Standard reduction potential / eV
Normal Hydrogen Electrode 0.000 (by definition)[1]
Standard Calomel Electrode 0.242[1]
Ag / Ag+ 0.01 M (usually AgNO3) in CH3CN Variable dependent on setup[5]
Ag/AgCl, KCl(sat. in H2O) * 0.197[1]

* Note: the AgCl coats the silver electrode

The reference electrode is setup so that the reference solution is separated from the studied solution via a frit.

This arrangement allows an electrical connection which permits a measurement of voltage. The slow movement of liquid through the frit (a porous glass membrane that allows liquid to flow through it at a slow rate) reduces the mixing of the reference solution and the studied solution to a minimum.

Even with the use of a frit, however, some mixing to be expected. For this reason, an Ag/Ag+ electrode is sometimes favoured over the Ag/AgCl, KCl(sat. in H2O) as the Ag/AgCl, KCl(sat. in H2O) will slowly leak water over time and water impurities in the studied solution lead to the narrowing of the potential window.

In addition, the AgCl in solution may also be reactive to certain studied chemicals. A double frit can be employed to prevent this, with an interior reference solution and an exterior studied solution separated from the bulk studied solution. This prevents the studied solution near the working electrode from being contaminated with water.

Note: Frits should always be stored in liquid between uses to prevent degradation. Never store a frit in air.

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The quasi reference electrode

An alternative reference electrode in cyclic voltammetry experiments is the quasi reference electrode (also known as a pseudoreference electrode). This is a reference electrode (usually silver wire) which does not have a surrounding solution with ions to form the half equation.

Because the potential of this reference electrode is not defined by ions of known concentration, the use of an internal standard such as ferrocene is vital. In addition, because the point which is being referenced against can shift depending on the contents of the solution, it is important that the internal standard is present during the reduction / oxidation of the studied chemical.

There are some disadvantages to using a quasi reference electrode. While they can reproduce the results of a standard reference electrode and are much easier to setup, they are also much more susceptible to potential drift [6].

A large standard deviation has also been reported when using a quasi reference electrode. This can be reduced by separating the electrode from the rest of the solution using a frit (with the reference solution the same as the studied solution) [6] [7].

Cyclic Voltammetry Applications


Determining the reversibility of a reaction

Cyclic voltammetry can be used to determine the reversibility of a reaction. The reversibility of a reaction can be split into two measures: 

  • If an electron transfer can be reversed, and no side reactions occur (both reactions involving the species before or after electron transfer prevent chemical reversibility), it is said to be "chemically reversible".
  • If the rate of electron transfer is sufficient that equilibrium is maintained, then it is said to be "thermodynamically reversible"

If both thermodynamic and chemical reversibility are observed, the electron transfer is said to be reversible i.e. there is "practical reversibility".

Practical reversibility does not mean that no side reactions occur, or that the system cannot be perturbed from equilibrium. It only requires equilibrium, and a lack of side reactions on the time scale of the experiment. For example, if the oxidised compound degrades slowly, and the scan rate (the rate of potential change) of the cyclic voltammogram is sufficient that this happens in minimal compound, it remains considered as reversible. In addition, if thermodynamic equilibrium is not maintained at very fast scan rates, but is maintained at the scan rate used, it is still considered "reversible".

To determine if a reaction is reversible for a given scan rate, two parameters are needed:

  1. The difference in the potential difference between the anodic peak current and the cathodic peak current
  2. The height of the anodic and cathodic peaks relative to the baseline.2 

A variety of methods are used for determining the ratio of the cathodic and anodic peaks. Note: the baseline after the electron transfer is used for the reverse peak.3 

For an oxidation, the height of the peak cathodic current (ipc) can be hard to determine, and likewise for the peak anodic current for a reduction. To determine the ratio of the anodic and cathodic peaks from a single experiment, without fitting the curve, the Nicholson parameter is often used.3,4 

In this equation the isp is the absolute current at the switching potential, ipc is the absolute current at the cathodic peak potential, and ipa is the absolute anodic peak potential; the (…)0 notation says that the currents are taken relative to the i=0 line. The Nicholson parameter only gives accurate results when the switching potential is 60/n mV past the peak anodic potential for an oxidation, and -60/n mV for a reduction.3,4

When the currents discussed are measured, they must be done such that the charging currents are neglected. In cyclic voltammetry, approximately constant charging currents are observed for each linear sweep. The direction of the charging currents reverse as the direction of sweep changes. It is this effect which leads to hysteresis in the voltammogram when no reduction or oxidation occurs (i.e. in the absence of redox active species).

To first approximation, these currents can be subtracted by running a voltammogram without redox active species present, and subtracting the results from the actual voltammogram. Alternatively, they may be subtracted by looking at the steady current when no redox reactions occur, and subtracting this for forward and reverse runs. While subtraction does improve the results, it is not perfect and does produce errors for faster scan rates, specifically when the uncompensated resistance is considered. An example where problems occur is at fast scan rates accessible at ultramicroelectrodes.1 

In addition, large uncompensated resistance resulting from the movement of charges in solution may make the system look quasireversible or irreversible when it is not, and this applies generally to voltammograms. The size of the potential expended on uncompensated resistance increases with the scan rate. More accurate results may be obtained from fitting the system with more complex equations,8,9 or using specially designed instruments.1

For a reversible reaction, ipa/ipc=1, and Epa-Epc=0.52/n V for an oxidation, and -0.52/n V for a reduction where n is the number of electrons in the oxidisation or reduction.2 

A quasireversible system is where the scan-rate and the electron transfer rate compete, such that thermodynamic equilibrium is never reached. A voltammogram which still has both anodic and cathodic peaks, but does not have Epa-Epc=0.52/n is considered quasi-reversible. 

If the reverse peak is missing, an electron transfer is classified as irreversible. Note if ipa/ipc≠1 the reaction is not chemically reversible. This suggests that another reaction is taking place in addition to the oxidation and reduction of the compound (i.e. O+n e-⇌R). When additional chemical reactions occur, there is said to be coupled reactions.

Determining the formal reduction potential of a species

The formal potential is a measure of the potential of a cell where both oxidised and reduced species are present at equal concentration. The formal potential is defined via potentiometry which involves measuring the potential difference in a galvanic cell. 

In a galvanic cell, there is a solution of one redox pair with an electrode on the left, and another solution of another redox pair on the right. The solutions are connected such that ions can move between the solutions in order to maintain a charge balance. A potential difference can be measured between these two solutions and is given a positive sign if reduction occurs spontaneously at the right electrode, and oxidation at the left electrode, when the electrodes are allowed to discharge. If the oxidation were to occur at the right electrode, and reduction at the left electrode, then a negative sign would be given. 

In a standard three electrode setup, the reference electrode is considered the "left electrode", and the potential of other species is measured relative to it. The measured potential difference is often referred to as the electromotive force or EMF, but this is no longer a recommended term.1 This potential difference is related to the Gibbs free energy, which corresponds to the overall reaction when there is oxidation at the left electrode and reduction at the right electrode.

The Gibbs free energy can be calculated using the following formula:

The "standard potential" is the potential where all species are present with unit activity. In order to make a set of comparable standard potentials, the standard potential where a set standard is oxidised was created. The "standard reduction potential" is the "standard potential" where the reference electrode is the normal hydrogen electrode.

It is possible to calculate the standard potential of a cell (Ecell0), from the standard reduction potential of the redox couples on the left (Eleft0) and the right (Eright0). The standard potential, i·e. the potential where all compounds are present under standard conditions, is given by the following:

The EMF of a cell with non-unit activity (E) may be related to the standard potential (E0) using the following formula, where ax is the activity of compound x, n the number of electrons involved in the oxidation/reduction, and T is the temperature:

O+n e-⇌R

Because the standard potential requires knowledge of the activity constant of the oxidised and reduced compound, it is quite complex to calculate. As such the formal potential is frequently used instead. The formal potential (E0') is given by the following formula, where symbols are defined as previously defined, [X] is the concentration of compound X, and X is its activity constant.


To approximate the formal potential with cyclic voltammetry, the polarographic half wave potential E1/2 is often calculated. For a reversible (or quasireversible) reaction the polarographic half wave potential E1/2 can be approximated as the average of the peak cathodic and anodic currents.3,5 The halfwave potential can be calculated more accurately by polarographic type methods; however, in cyclic voltammetry, for a reversible reaction, an accuracy of the order of mV is observed.3,5

The halfwave potential (E1/2) is related to the formal potential (E0'), which in turn is related to the standard potential (E0). The relationship is shown below, where am is the activity of species m, m the concentration, m the activity coefficient, and Dm the diffusion constant.



In practice, the difference between E1/2 and E0' is minimal.3 As such, it is quite common to approximate E1/2=E0'.

Measuring electron transfer kinetics

The kinetics of the electron transfer are often characterised in terms of the standard rate constant k0. k0 is the rate constant of both oxidation and reduction when the electrode is at the formal potential i.e. when E=E0'.

The standard rate constant can be calculated using the "method of Nicholson",1 if the diffusion constant of the species in solution is known. In his 1965 paper,6 Nicholson calculated the separation of peak cathodic and anodic current for a reduction for various values of ψ:

Here, Do is the diffusion constant of the oxidised species, DR of the reduced species, v the scan rate, n the number of electrons involved in the oxidation/reduction, α is the transfer coefficient (see Butler-Volmer equation), T the temperature, F the faradays constant, and R the gas constant.

From the graph plotted in the paper, it is possible to look up a value of for a known separation of anodic peaks. The now known value of can be used to calculate a value of k0 if all other constants are known. In addition, the diffusion constants of oxidation and reduction are usually found to be similar. It is therefore claimed6 to be possible to approximate:

It should be noted these results are only reliable6 for 0.3<α<0.7, but this is commonly true. For a reduction, as opposed to an oxidation, the same process applies, but would require different formulations of the equations.

A problem with this method is that the uncompensated resistance may dominate the cyclic voltammogram.1,6 Uncompensated resistance is the resistance to charge flowing outside the electrochemical process at the electrode. The uncompensated resistance is particularly an issue when reversibility is approached, and as such resistance in solution should be minimised.

Determining the energy levels of semiconducting polymers

A commonly used method for determining the energy levels of an organic semi-conducting polymers is to deposit the polymer on the working electrode and measure the electrical response. Depositing a solid on an electrode forms what is referred to as a modified electrode.1 The onset potential is usually calculated, and this value assigned as the energy of the HOMO and LUMOs. 

The evidence for this method comes from a paper7 to test the ability of Valence Effective Hamiltonian (VEH) to calculate redox potentials. In the paper, the onset potential and gas phase ionisations energies were observed to correlate with VEH calculated HOMO and LUMO. However, the absolute values show absolute error of the order of 0.1 V, in opposite directions for the two experimental techniques. Based on this evidence, the onset potentials are often used to calculate the HOMO and LUMO energies more generally. Here HOMO is being used to mean the highest energy electron in the solid state, and LUMO the lowest energy excited electron (different from a theoretical description of HOMO and LUMO).

A modern electrochemical review,8 however, discourages the use of this method, claiming,8 that onset potential "are not standard potentials, and do not posses thermodynamic significance". 

Assignment and characterisation of coupled reaction

Cyclic voltammetry may also be used to assign coupled reactions to the electrochemically formed species. By examining the shape, and by plotting the dependence on peak position, and ipaipc on scan rate, the processes which occurs may be characterised.2,3

Cyclic Voltammetry of a Polymer


Cyclic Voltammetry of Ferrocene

Ferrocene (Fc) is commonly used as an internal standard for cyclic voltammetry.

First, a positively ramping potential (the forward sweep) is applied between the working and reference electrodes. As the potential increases, Fc physically close to the working electrode is oxidised (i.e., loses an electron), converting it to Fc+. The movement of the electrons creates an electrical current.

As un-reacted Fc diffuses to the electrode and continues the oxidation process, the electrical current is increased and there is a build up of Fc+ at the electrode. This build up of Fc+ and depletion of Fc is called the the diffusion layer, and effects the rate at which un-reacted material can reach to the electrode.

Once the diffusion layers reaches a certain size, the diffusion of Fc to the electrode slows down, resulting in a decrease in the oxidation rate and thus a decrease in electrical current.

Cyclic voltammogram of Ferrocene
Cyclic voltammogram of ferrocene in the Ossila Potentiostat PC software

When the potential ramp switches direction, the process reverses and the reverse sweep begins. Fc+ close to the working electrode reduces (i.e., gains an electron), converting it back to Fc. The electrical current flows in the opposite direction, creating a negative current. The Fc+ diffuses to the electrode, reducing to Fc and resulting in a increase in the negative current.

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Other Voltammetric Techniques


Broadly speaking, voltammetric techniques can be categorised as being either sweep type or polarography-like. The former refers to methods like cyclic voltammetry where the solution is not stirred after each set potential, and the latter refers to techniques where it is. Additional methods modify other techniques, e.g. with the use of a rotating electrode.

Although cyclic voltammetry remains the most widely used voltammetric technique (primarily due to its speed, range of uses, ease with which the data can be analysed), each technique has different advantages, disadvantages, and applications.

  

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References


  1. Sevćik, A. Collection of Czechoslovak Chemical Communications 1958, 13, 349
  2. A. L. Bard and L. Faulkner Electrochemical methods: Fundamentals and Applications, 2nd ed. John Wiley & Sons 2001
  3. W. L. G. Armarego and C. L. L. Chai Purification of Laboratory Chemicals, 7th ed. Butterworth-Heinemann 2012
  4. L. J, L. B, and P. G Advanced Practical Organic Chemistry, 3rd edition. Manipal: Routledge 2013
  5. J. L. Brédas, R. Silbey, D. S. Boudreaux, and R. R. Chance Chain-Length Dependence of Electronic and Electrochemical Properties of Conjugated Systems: Polyacetylene, Polyphenylene, Polythiophene, and Polypyrrole J. Am. Chem. Soc., vol. 105, no. 22, pp. 6555–6559, 1983
  6. N. Elgrishi, K. J. Rountree, B. D. McCarthy, E. S. Rountree, T. T. Eisenhart, and J. L. Dempsey A Practical Beginner’s Guide to Cyclic Voltammetry J. Chem. Educ., vol. 95, no. 2, pp. 197–206, 2018
  7. G. A. Snook, A. S. Best, A. G. Pandolfo, and A. F. Hollenkamp Evaluation of a Ag/Ag+ reference electrode for use in room temperature ionic liquids Electrochem. commun., vol. 8, no. 9, pp. 1405–1411, 2006
  8. V. M. Hultgren, A. W. A. Mariotti, A. M. Bond, and A. G. Wedd Reference potential calibration and voltammetry at macrodisk electrodes of metallocene derivatives in the ionic liquid [bmim][PF6] Anal. Chem., vol. 74, no. 13, pp. 3151–3156, 2002

Contributing Authors


  • Harry Robson
  • Max Reinhardt
  • Chris Bracher
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