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Cyclic Voltammetry Basic Principles, Theory and Setup

Duck-shaped cyclic voltammogram
Duck-shaped cyclic voltammogram.

Cyclic voltammetry is an electrochemical technique used to measure the current response of a redox active solution to a linearly cycled potential sweep. It is a useful method if you need to quickly find information about the thermodynamics of redox processes, the energy levels of the analyte and the kinetics of electronic-transfer reactions.

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

To perform cyclic voltammetry, you need to start by adding your electrolyte solution to an electrochemical cell, along with a reference solution and the three electrodes. After this, use a potentiostat to linearly sweep the potential between the working and reference electrodes. When the potentiostat reaches the pre-set limit, it will sweep back in the opposite direction.

The potentiostat will repeat this process multiple times during a scan. While doing so, it will record the changing current between the working and counter probes. The result is a characteristic duck-shaped plot known as a cyclic voltammogram.

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Basic Theory and Principles

Cyclic voltammetry is a sophisticated potentiometric and voltammetric method. During a scan, the chemical either loses an electron (oxidation) or gains an electron (reduction). This will depend 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 the products (B). This occurs at the electrode/solution interface and can be recognised according to the Nernst equation:

Nernst equation
Nernst equation

Here, 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 is the ratio of products to reactants, raised to their respective stoichiometric powers. You can use this 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

Introduction to Voltammetry

Voltammetry is any technique that involves measuring the current while varying the potential between two electrodes. Voltammetric methods include cyclic voltammetry and linear sweep voltammetry, as well as similar electrochemical techniques such as staircase voltammetry, squarewave voltammetry, and fast-scan cyclic voltammetry.

In voltammetry, the current is generated by electron transfer between the redox species and the two electrodes. The diffusion and migration of ions carries this current through the solution.

The Three Electrode System

In principle, cyclic voltammetry (and other types of voltammetry) only requires two electrodes. However, in practice it is difficult to keep a constant potential while measuring the resistance between the working electrode and the solution. This is made more difficult as you try to pass the necessary current while also passing current to counteract the redox events at the working electrode.

It is the three-electrode system that separates the role of referencing the potential applied from the role of balancing the current produced.

Three electrode cell system
Three electrode cell as used in cyclic voltammetry with an Ossila Potentiostat.

To measure and control the potential difference applied, the potentiostat varies the potential of the working electrode while the potential of reference electrode remains fixed by an 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.

It is important to note that minimal current passes between the reference and the 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 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 current response that you measure will be dependent on the concentration of the redox species (the analyte) at the working electrode surface. You can explain this by using 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/∂x0) 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.

Cyclic Voltammograms Explained

A cyclic voltammogram is the ‘duck-shaped’ plot generated by cyclic voltammetry.

Cyclic voltammogram for an electrochemically-reversible one-electron redox process
Example of a cyclic voltammogram for an electrochemically reversible one-electron redox process.

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

As the potential approaches several kT of the standard potential, the onset (Eonset) of oxidation is reached. Following this, the current exponentially increases (b) as the analyte begins its oxidation at the working electrode surface. For a reversible process, here the current rises initially as if there is no change in the concentration of oxidant. The current is dictated by the rate of diffusion of the oxidant to the electrode, as well as the proportion converted to the reduced form. This can be understood according to the Nernst equation. Gradually, as the scan continues, more oxidant is depleted. The concentration gradient adjusts to this. It is this change which causes a peak in the voltammogram. You can see how the decrease in current from depletion of the oxidant outweighs the increase from changing the proportion of oxidant oxidised at the electrode.

The current reaches peak maximum at point c (anodic peak current (ipa) for oxidation at the anodic peak potential (Epa). Here, 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. This occurs 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. This continues 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. The only difference is that it occurs 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. This is only 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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Cyclic Voltammetry of Ferrocene

Ferrocene (Fc) is a common internal standard for cyclic voltammetry. Its cyclic voltammogram can therefore be considered "typical". Like the general case which was described above, at the start of a cyclic voltammetry scan a positively ramping potential (the forward sweep) is applied between the working and reference electrodes. As the potential increases, ferrocene (Fc) physically close to the working electrode is oxidised (i.e. loses an electron). This converts it to Fc+, and the movement of the electrons creates a measurable 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.

Cyclic Voltammetry 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.
  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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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 you remove the oxygen, you can keep it out with a continuous stream of inert gas.

Risk of contamination

When you prepare an electrochemical cell, it is important to minimise the risk of any contamination with water. This is because water can form reactive species when reduced or oxidised. You can do this by heating the components in a glassware oven prior to using them. 

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

Your choice of which solvent and electrolyte to use should be dictated by the solubility of your 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, we recommend that you 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

You can use internal standards, usually ferrocene (see below), 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 you can use 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 [6]. This is particularly true for quasi reference electrodes where large shifts have been observed.

Why do we use supporting electrolyte for cyclic voltammetry?

Inert ions are added to the electrochemical solution in molar excess to the analyte in order to provide enough ionic strength to the solution. This is important for it obey the Nernst equation. The excess of electrolyte decreases the thickness of the diffuse double layer. This therefore means that the applied potential will decrease 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.

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. You can usually achieve inertness 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). You should 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). This is partly becuase 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 to ensure 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].

Applications of Cyclic Voltammetry

For applications, please see 'Cyclic Voltammetry Applications and Voltammograms'.

Similar Electrochemical Methods

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. Other types of voltammetry modify these methods, for example, with the use of a rotating electrode.

The types of voltammetry page gives more information on the advantages, disadvantages, and applications of each technique.

Cyclic voltammetry remains the most widely used voltammetric technique due to its speed, range of uses, and the ease with which the data can be analysed.


  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
  9. J. Heinze, Angew. Chemie Int. Ed. English, 1984, 23, 831–918
  10. G. A. Mabbott, J. Chem. Educ., 1983, 60, 697
  11. R. S. Nicholson, Anal. Chem., 1966, 38, 1406
  12. R. S. Nicholson and I. Shain, Anal. Chem., 1964, 36, 706–723
  13. R. S. Nicholson, Anal. Chem., 1965, 37, 1351–1355
  14. J. Heinze, B. A. Frontana-Uribe and S. Ludwigs, Chem. Rev., 2010, 110, 4724–4771

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