Cyclic Voltammetry Basics, Setup, and Applications
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.
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 pre-set 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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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) 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:
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 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:
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 similar electrochemical techniques such as staircase voltammetry, squarewave voltammetry and fast-scan cyclic voltammetry.
When performing voltammetry, a current is generated as 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.
The Three Electrode System
Although in principle cyclic voltammetry (and other types of 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 system is often used to separate the role of referencing the potential applied and balance 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 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:
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.
The ‘duck-shaped’ plot generated by cyclic voltammetry is called a cyclic voltammogram.
In the example cyclic voltammogram shown, 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 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.
At 298 K, the 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).
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Cyclic Voltammetry of Ferrocene
Ferrocene (Fc) is commonly used as an internal standard for cyclic voltammetry, and its cyclic voltammogram can therefore be considered "typical". As described above for the general case, 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.
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.
The experimental setup for cyclic voltammetry consists of an electrochemical cell containing 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.
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 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.
The electrochemical solution used for cyclic voltammetry typically consists of three components.
- 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 desired potential range.
Reference tables of the potential range of various solvent and electrolyte pairs are widely available  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  that Grubbs purification apparatuses can add undesirable electroactive impurities .
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 . 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 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, 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 . 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 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.
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 .
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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)|
|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 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 .
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)  .
Cyclic Voltammetry Applications
Cyclic voltammetry is a versatile electrochemical method and the most commonly used type of voltammetry. It can be used to probe qualitative and quantitative information about the electrochemical properties of chemicals.
Common applications of cyclic voltammetry include:
- Determining the reversibility of an electrochemical reaction
- Determine the formal reduction potential of a species
- Measure electron transfer kinetics
- Determine the energy levels of semiconducting polymers
- Characterising a coupled reaction
Note: the equations in this section assume linear diffusion. Most electrodes can be approximated as having linear diffusion on short time scales. The larger the electrode, and the shorter the time observed, the more the linear diffusion approximation applies.
Determining the reversibility of a reaction
What determines the reversibility of a reaction?
The reversibility of a reaction can be split into two metrics, chemical reversibility and thermodynamic reversibility.
- If an electron transfer can be reversed without any side reactions (either before or after) then 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 necessarily mean that no side reactions ever 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, a reaction would still be considered reversible if the oxidised compound slowly degrades over time, providing that a sufficient scan rate (rate of potential change) is used. Similarly, even if thermodynamic equilibrium is not maintained at very fast scan rates, the reaction is still considered reversible if it is maintained during the cyclic voltammetry experiment.
How to determine the reversibility of a reaction
Cyclic voltammetry can be used to measure two parameters, which can together can be used to determine if a reaction is reversibly for a given scan rate. These are:
- The difference in the potential difference between the anodic peak current and the cathodic peak current
- The height of the anodic and cathodic peaks relative to the baseline  (note: for the reverse peak, the baseline after the electron transfer is used )
The Nicholson parameter
For an oxidation, the height of the peak cathodic current (ipc) can be hard to determine, and likewise, the peak anodic current can be difficult to determine for a reduction. As a result, a variety of methods are used for determining the ratio of the cathodic and anodic peaks. To do so from a single experiment without fitting the curve, the Nicholson parameter is often used.  
In this equation, 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 simply indicates 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.  
When measuring the currents required to calculate the Nicholson parameter, the charging currents need to be neglected. In cyclic voltammetry, these are approximately constant for each linear sweep, but their direction changes depending on the direction of the sweep. It is this effect which leads to hysteresis in the cyclic voltammogram when no reduction or oxidation occurs (i.e. in the absence of redox active species).
To first approximation, the charging currents can be neglected simply by running a voltammogram without redox active species present and subtracting the results from the actual voltammogram. Alternatively, they can be subtracted by looking at the steady current when no redox reactions occur, and subtracting this for forward and reverse runs.
Though they do improve the quality of the results obtained, neither of these approaches are perfect. Specifically, both methods are prone to errors at faster scan rates, such as those which are accessible at ultramicroelectrodes . At these rates, the size of the potential expended on uncompensated resistance becomes more significant.
Uncompensated resistance is the resistance to charge flowing outside the electrochemical process at the electrode. Large uncompensated resistance resulting from the movement of charges in solution can make the system look quasireversible or irreversible, even when it is not.
More accurate results can be obtained from fitting the system with more complex equations   or by using specially designed instruments. 
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. 
A quasireversible system is where the scan-rate and the electron transfer rate compete, such that thermodynamic equilibrium is never reached. A cyclic 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 that if ipa/ipc≠1, then 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.
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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, but can be approximated with cyclic voltammetry.
How to determine the formal potential with potentiometry
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. 
The Gibbs free energy
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 and standard reduction potential
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) can 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:
The formal potential
Because the standard potential requires knowledge of the activity constant of the oxidised and reduced compound, it is quite complicated to calculate. As such, the formal potential (E0') is frequently used instead. This 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.
How to approximate the formal potential with cyclic voltammetry
To approximate the formal potential with cyclic voltammetry, the polarographic half wave potential E1/2 is often calculated. For a reversible (or quasireversible) reaction, this can be approximated as the average of the peak cathodic and anodic currents.  
The halfwave potential can be calculated more accurately by polarographic type methods, but in cyclic voltammetry, an accuracy of the order of mV is observed for a reversible reaction.  
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] is the concentration, γm is the activity coefficient, and Dm is the diffusion constant.
In practice, the difference between E1/2 and E0' is minimal.  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'.
How to calculate the standard rate constant with the method of Nicolson
The standard rate constant can be calculated using the "method of Nicholson", if the diffusion constant of the species in solution is known. In his 1965 paper, 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 is the scan rate, n is the number of electrons involved in the oxidation/reduction, α is the transfer coefficient (see Butler-Volmer equation), T is the temperature, F is the faradays constant, and R is the gas constant.
It is possible to look up a value of ψ for a known separation of anodic peaks, and if all other constants are known, this value can then be used to calculate a value of k0. In addition, the diffusion constants of oxidation and reduction are usually found to be similar, so it is sometimes considered valid  to approximate:
It should be noted these results are only reliable  for 0.3<α<0.7, but this is commonly true.
For a reduction, as opposed to an oxidation, the same basic process applies.
A problem with this method is that the uncompensated resistance (resistance to charge flowing outside the electrochemical process at the electrode) can dominate the cyclic voltammogram.   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 organic semi-conducting polymers with cyclic voltammetry 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. The onset potential is usually then calculated, and this value is assigned as the energy of the HOMO and LUMOs.
The evidence for this method comes from a paper  which tested 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,  discourages the use of this method, claiming that onset potential "are not standard potentials, and do not posses thermodynamic significance".
Assignment and characterisation of coupled reaction
Cyclic voltammetry can also be used to assign coupled reactions to the electrochemically formed species. By examining the shape and plotting the dependence on peak position, and ipa/ipc on scan rate, the processes which occur can be characterised.  
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.
- Sweep type methods
- Polarography like methods
- Conventional polarography
- Normal pulse voltammetry
- Reverse pulse voltammetry
- Differential pulse voltammetry
- Squarewave pulse voltammetry
- Other types of voltammetry
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.
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