Cyclic Voltammetry: The Basics Explained
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.
- Introduction to potentiometry and voltammetry
- Cyclic voltammetry basics
- Setting up an electrochemical cell for cyclic voltammetry
- Cyclic voltammetry of ferrocene
Introduction to Potentiometry and Voltammetry
Cyclic voltammetry is a potentiometric method and one of a number of different types of voltammetry.
Potentiometry is a way of measuring the electrical potential of an electrochemical cell under static conditions (i.e. no current flow) and voltammetry is any technique where the current is measured while the potential between two electrodes is varied.
The potentiometry principle
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:
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 of the concentrations of products to reactants, raised to their respective stoichiometric powers, which 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:
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.
Voltammetry is a technique where the current is measured while the potential between two electrodes is varied. The current generated is a result of electron transfer between the redox species and the electrodes, which is carried through the solution by the migration of ions.
In practice, is very difficult to maintain a constant potential at the reference electrode while also passing current to counteract the redox events at the working electrode. As a result, a three-electrode cell is typically used to separate the role of referencing the potential applied and balancing the current produced.
To measure and control the potential difference applied, the potential of the working electrode is varied while the potential of reference electrode remains fixed by a well-defined value electrochemical redox reaction.
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.
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 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:
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 to provide enough ionic strength to the solution to obey the Nernst equation. The excess of electrolyte decreases the thickness of the diffuse double layer to ensure 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 Voltammetry Basics
Cyclic voltammetry is a sophisticated voltammetric method. A potentiostat is used to linearly sweep the potential between the working and reference electrodes until it reaches a preset limit. It is then swept back in the opposite direction, switching potentials. 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 resulting ‘duck-shaped’ plot is called a cyclic voltammogram, an example of is displayed in Figure 1.
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 DDL.
The current response decreases from linearity as the analyte is depleted and the DDL 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 and therefore does not satisfy the Nernst equation. 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.
The peak current, ip, of the reversible redox process is described by the Randles-Sevcik equation.
At 298 K, Randles-Sevcik equation is:
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).
How does cyclic voltammetry work?
When performing cyclic voltammetry, the applied potential causes the chemical being tested to undergo either oxidation or reduction depending on the direction of the ramping potential. Oxidation and reduction are electron transfer processes. When a chemical undergoes oxidation, it loses an electron and is said to be oxidised. Likewise, when a chemical undergoes reduction, it gains an electron and is said to be reduced.
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Setting up an Electrochemical Cell for Cyclic Voltammetry
Cyclic voltammetry uses an electrochemical cell consisting of five major components.
- The working electrode, where the compound of interest is reduced (Cn+ → C(n−1)+ ) or oxidised (Cn+ → C(n+1)+)
- The counter electrode, which completes the circuit with the potentiostat (see figure below)
- The reference electrode, used to measure the potential
- The studied solution containing the chemical to be studied
- 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.
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 set up of the electrochemical cell is relatively simple.
The working and counter electrodes sit in the electrochemical solution, and the reference electrode sits in the reference solution in a separate tube within the cell. The reference electrode tube should be approximately two thirds full - a syringe and needle can be used to add the solution.
Electrochemical glassware typically has holes for each electrode as well as gas intakes which 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 is kept out by flowing through a continuous stream of inert gas into the cell.
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A further requirement is to minimise any contamination with water. This can be done by heating the components in a glassware oven prior to use. Water can form reactive species when reduced or oxidised as well as reduce the potential range (see choice of solvent).
Typically there are three components in the electrochemical solution used in cyclic voltammetry.
- The compound of interest (10-3 – 10-5 M)
- An electrolyte (0.1 M)
- A solvent which dissolves both the compound of interest and the electrolyte
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 potential range desired.
Reference tables of the potential range of solvent and electrolyte pairs are available in various books on cyclic voltammetry . These ranges, however, are highly dependent on purity and dryness of both electrolyte and solvent; the highest purity solvent and electrolyte should be used for all measurements and all components should be oven dried before use.
Manual purification and drying can be done following standard procedures . Be aware when doing this that Grubbs purification apparatuses can add undesirable electroactive impurities .
A short table of potential ranges is listed below based on the values given by A.J. Bard and L.R. Faulkner . 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|
It is possible to study the electrical response if the compound of interest cannot be sufficiently dissolved in standard electrochemical solvents. To do this, coat the working electrode with the material by depositing it using a solvent (which does dissolve the compound).
Under these circumstances, the normal equations and mathematical proofs do not strictly apply because there is no free diffusion. However, it is still possible to get an approximation of the energy levels for insoluble materials (such as polymers) by using this technique and approximating the onset potential as the redox potential of that process.
Internal standards, usually ferrocene (see below), are often used to calculate the value of the oxidation and reduction potentials. They are compounds which oxidises or reduces in solution, ideally somewhat independently of the system (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.
Often these standards are studied immediately after the studied chemical using the same solutions, but recent reviews state that it is better to have the internal standard always present to prevent changes in position of voltammograms . This is particularly true for quasi reference electrodes where large shifts have been observed.
Cyclic voltammetry 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 .
Choice of reference electrode
Reference electrodes are designed so that an equilibrium is set up 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 set up 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 set up 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 the electrolyte, the solvent and estimating the volume of solution that you require for your experiment.
|Electrode||Standard reduction potential / eV|
|Normal Hydrogen Electrode||0.000 (by definition)|
|Standard Calomel Electrode||0.242|
|Ag / Ag+ 0.01 M (usually AgNO3) in CH3CN||Variable dependent on setup|
|Ag/AgCl, KCl(sat. in H2O) *||0.197|
* Note: the AgCl coats the silver electrode
The reference electrode is set up so that the reference solution is separated from the studied solution via a frit.
A frit is a porous glass membrane that allows liquid to flow through it at a slow rate. Frits should always be stored in liquid between uses to prevent degradation.
Never store a frit in air.
This arrangement allows an electrical connection which permits a measurement of voltage. The slow movement of liquid through the frit 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 favored 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.
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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. Further, because the point which is being referenced against can shift depending on what is in solution, it is vital 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 do reproduce the results of a standard reference electrode  and are much easier to set up, they are also much more susceptible to drift in its potential during recording of data .
A rather large standard deviation has also been reported, though this can be reduced by separating the electrode from the rest of the solution using a frit as is usually done for a non-quasi reference electrode (here the reference solution is the same as the studied solution)  .
Cyclic Voltammetry of Ferrocene
To aid in the explanation of what occurs during the measurement, we shall use the example of ferrocene (Fc).
First, a positively ramping potential (the forward sweep) is applied between the working and reference electrodes. As the potential increases, Fc 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 reacted material is called 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.
When the potential ramp switches direction, the process reverses (the reverse sweep). 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.
As with the forward sweep, a build up of material occurs near the electrode, eventually slowing down the diffusion of Fc+ and causing the negative current to decrease.
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