Voltammetry

Voltammetry is a category of electroanalytical methods used in analytical chemistry and various industrial processes. In voltammetry, information about an analyte is obtained by measuring the current as the potential is varied. The analytical data for a voltammetric experiment comes in the form of a voltammogram, which plots the current produced by the analyte versus the potential of the working electrode.

Voltammetry is a category of electroanalytical methods used in analytical chemistry and various industrial processes. In voltammetry, information about an analyte is obtained by measuring the current as the potential is varied. The analytical data for a voltammetric experiment comes in the form of a voltammogram, which plots the current produced by the analyte versus the potential of the working electrode. These curves I = f(E) are called voltammograms. The potential is varied arbitrarily, either step by step or continuously, and the resulting current value is measured as the dependent variable. The opposite, i.e., amperometry, is also possible but not common. The shape of the curves depends on the speed of potential variation, (nature of driving force) and whether the solution is stirred or quiescent (mass transfer). Most experiments control the potential (volts) of an electrode in contact with the analyte while measuring the resulting current (amperes). To conduct such an experiment, at least two electrodes are required. The working electrode, which makes contact with the analyte, must apply the desired potential in a controlled way and facilitate the transfer of charge to and from the analyte. A second electrode acts as the other half of the cell. This second electrode must have a known potential to gauge the potential of the working electrode from; furthermore it must balance the charge added or removed by the working electrode. While this is a viable setup, it has a number of shortcomings. Most significantly, it is extremely difficult for an electrode to maintain a constant potential while passing current to counter redox events at the working electrode. To solve this problem, the roles of supplying electrons and providing a reference potential are divided between two separate electrodes. The reference electrode is a half cell with a known reduction potential. Its only role is to act as reference for measuring and controlling the working electrode's potential and it does not pass any current. The auxiliary electrode passes the current required to balance the observed current at the working electrode. To achieve this current, the auxiliary will often swing to extreme potentials at the edges of the solvent window, where it oxidizes or reduces the solvent or supporting electrolyte. These electrodes, the working, reference, and auxiliary make up the modern three-electrode system. There are many systems which have more electrodes, but their design principles are similar to the three-electrode system. For example, the rotating ring-disk electrode has two distinct and separate working electrodes, a disk, and a ring, which can be used to scan or hold potentials independently of each other. Both of these electrodes are balanced by a single reference and auxiliary combination for an overall four-electrode design. More complicated experiments may add working electrodes, reference, or auxiliary electrodes as required. In practice it can be important to have a working electrode with known dimensions and surface characteristics. As a result, it is common to clean and polish working electrodes regularly. The auxiliary electrode can be almost anything as long as it doesn't react with the bulk of the analyte solution and conducts well. A common voltammetry method, polarography, uses mercury as a working electrode e.g. DME and HMDE, and as an auxiliary electrode. The reference is the most complex of the three electrodes; there are a variety of standards used. For non-aqueous work, IUPAC recommends the use of the ferrocene/ferrocenium couple as an internal standard. In most voltammetry experiments, a bulk electrolyte (also known as a supporting electrolyte) is used to minimize solution resistance. It is possible to run an experiment without a bulk electrolyte, but the added resistance greatly reduces the accuracy of the results. With room temperature ionic liquids, the solvent can act as the electrolyte. The supporting electrolyte also minimises the effect of migration-controlled and ensures that the reaction is diffusion-controlled. E= E^0 - \frac{R T}{z F} \ln Q Where: E : Reduction potential E^0 : Standard potential R : Universal gas constant T : Temperature in kelvin z : Ion charge (moles of electrons) F : Faraday constant Q : Reaction quotient This equation describes how the changes in applied potential will alter the concentration ratio. However, the Nernst equation is limited, as it is modeled without a time component and voltammetric experiments vary applied potential as a function of time. Other mathematical models, primarily the Butler-Volmer equation, the Tafel equation, and Fick's law address the time dependence. The Butler–Volmer equation relates concentration, potential, and current as a function of time. It describes the non-linear relationship between the electrode and electrolyte voltage difference and the electrical current. It helps make predictions about how the forward and backward redox reactions affect potential and influence the reactivity of the cell. This function includes a rate constant which accounts for the kinetics of the reaction. A compact version of the Butler-Volmer equation is as follows: {\displaystyle j=j_{0}\cdot \left\{\exp \left[{\frac {\alpha _{\rm {a}}zF\eta }{RT}}\right]-\exp \left[-{\frac {\alpha _{\rm {c}}zF\eta }{RT}}\right]\right\}} Where: j : electrode current density, A/m 2 (defined as j = I/S) j_0 : exchange current density, A/m 2 E : electrode potential, V {\displaystyle E_{\rm {eq}}} : equilibrium potential, V T : absolute temperature, K z : number of electrons involved in the electrode reaction F : Faraday constant R : universal gas constant \alpha_c : so-called cathodic charge transfer coefficient, dimensionless \alpha_a : so-called anodic charge transfer coefficient, dimensionless \eta : activation overpotential (defined as {\displaystyle \eta =E-E_{\rm {eq}}} ). At high overpotentials, the Butler–Volmer equation simplifies to the Tafel equation. The Tafel equation relates the electrochemical currents to the overpotential exponentially, and is used to calculate the reaction rate. The overpotential is calculated at each electrode separately, and related to the voltammogram data to determine reaction rates. The Tafel equation for a single electrode is: {\displaystyle \eta =\pm A\cdot \log _{10}\left({\frac {i}{i_{0}}}\right)} Where: the plus sign under the exponent refers to an anodic reaction, and a minus sign to a cathodic reaction \eta : overpotential, V A: "Tafel slope", V i : current density, A/m 2 i_0 : "exchange current density", A/m 2 . As the redox species are oxidized and reduced at the electrodes, material accumulates at the electrode/electrolyte interface. Material accumulation creates a concentration gradient between the interface and the bulk solution. Fick's laws of diffusion is used to relate the diffusion of oxidized and reduced species to the faradaic current used to describe redox processes. Fick's law is most commonly written in terms of moles, and is as follows: {\displaystyle J=-D{\frac {d\varphi }{dx}}} Where: J: diffusion flux (in amount of substance per unit area per unit time) D: diffusion coefficient or diffusivity (in area per unit time) φ: concentration (in amount of substance per unit volume) x: position (in length) ===Voltammetric sensors=== A number of voltammetric systems are produced commercially for the determination of species that are of interest in industry and research. These devices are sometimes called electrodes but are actually complete voltammetric cells, which are better referred to as sensors. These sensors can be employed for the analysis of organic and inorganic analytes in various matrices. ===The oxygen electrode=== The determination of dissolved oxygen in a variety of aqueous environments, such as sea water, blood, sewage, effluents from chemical plants, and soils is of tremendous importance to industry, biomedical and environmental research, and clinical medicine. One of the most common and convenient methods for making such measurements is with the Clark oxygen sensor, which was patented by L.C. Clark, Jr. in 1956. ==See also== Current–voltage characteristic Neopolarogram ==References== ==External links== http://new.ametek.com/content-manager/files/PAR/App%20Note%20E-4%20-%20Electrochemical%20Analysis%20Techniques1.pdf [https://web.archive.org/web/20061221122151/http://www.prenhall.com/settle/chapters/ch37.pdf]

Text derived from https://en.wikipedia.org/w/index.php?oldid=1353400870 under CC BY-SA 4.0; modifications have been made.