Cyclic voltammetry is a type of potentiodynamic electrochemical measurement. In a cyclic voltammetry experiment, the working electrode potential is ramped linearly versus time. Unlike in linear sweep voltammetry, after the set potential is reached in a CV experiment, the working electrode's potential is ramped in the opposite direction to return to the initial potential; these cycles of ramps in potential may be repeated as many times as needed. The current at the working electrode is plotted versus the applied voltage to give the cyclic voltammogram trace. Cyclic voltammetry is used to study the electrochemical properties of an analyte in solution or of a molecule, adsorbed onto the electrode. In cyclic voltammetry, the electrode potential ramps linearly versus time in cyclical phases; the rate of voltage change over time during each of these phases is known as the experiment's scan rate. The potential is measured between the working electrode and the reference electrode, while the current is measured between the working electrode and the counter electrode.
These data are plotted as current versus applied potential. In Figure 2, during the initial forward scan an reducing potential is applied. At some point after the reduction potential of the analyte is reached, the cathodic current will decrease as the concentration of reducible analyte is depleted. If the redox couple is reversible during the reverse scan the reduced analyte will start to be re-oxidized, giving rise to a current of reverse polarity to before; the more reversible the redox couple is, the more similar the oxidation peak will be in shape to the reduction peak. Hence, CV data can provide information about electrochemical reaction rates. For instance, if the electron transfer at the working electrode surface is fast and the current is limited by the diffusion of analyte species to the electrode surface the peak current will be proportional to the square root of the scan rate; this relationship is described by the Randles–Sevcik equation. In this situation, the CV experiment only samples a small portion of the solution, i.e. the diffusion layer at the electrode surface.
The utility of cyclic voltammetry is dependent on the analyte being studied. The analyte has to be redox active within the potential window to be scanned; the analyte displays a reversible CV wave, observed when all of the initial analyte can be recovered after a forward and reverse scan cycle. Although such reversible couples are simpler to analyze, they contain less information than more complex waveforms; the waveform of reversible couples is complex owing to the combined effects of polarization and diffusion. The difference between the two peak potentials, ΔEp, is of particular interest. ΔEp = Epa - Epc > 0This difference results from the effects of analyte diffusion rates. In the ideal case of a reversible 1e- couple, ΔEp is 57 mV and the full-width half-max of the forward scan peak is 59 mV. Typical values observed experimentally are greater approaching 70 or 80 mV; the waveform is affected by the rate of electron transfer discussed as the activation barrier for electron transfer. A theoretical description of polarization overpotential is in part described by the Butler–Volmer equation and Cottrell equation equations.
In an ideal system the relationships reduces to E p a − E p c = 56.5 mV n for an n electron process. Focusing on current, reversible couples are characterized by ipa/ipc = 1; when a reversible peak is observed, thermodynamic information in the form of a half cell potential E01/2 can be determined. When waves are semi-reversible, it may be possible to determine more specific information. Many redox processes observed by CV are non-reversible. In such cases the thermodynamic potential E01/2 is deduced by simulation; the irreversibility is indicated by ipa/ipc ≠ 1. Deviations from unity are attributable to a subsequent chemical reaction, triggered by the electron transfer; such EC processes can be complex, involving isomerization, association, etc. Adsorbed species give simple voltammetric responses: ideally, at slow scan rates, there is no peak separation, the peak width is 90mV for a one-electron redox couple, the peak current and peak area are proportional to scan rate; the effect of increasing the scan rate can be used to measure the rate of interfacial electron transfer and/or the rates of reactions that are coupled to electron transfer.
This technique has been useful to study redox proteins, some of which adsorb on various electrode materials, but the theory for biological and non-biological redox molecules is the same. CV experiments are conducted on a solution in a cell fitted with electrodes; the solution consists of the solvent, in, dissolved electrolyte and the species to be studied. A standard CV experiment employs a cell fitted with three electrodes: reference electrode, working electrode, counter electrode; this combination is sometime
Mercury is a chemical element with symbol Hg and atomic number 80. It is known as quicksilver and was named hydrargyrum. A heavy, silvery d-block element, mercury is the only metallic element, liquid at standard conditions for temperature and pressure. Mercury occurs in deposits throughout the world as cinnabar; the red pigment vermilion is obtained by synthetic mercuric sulfide. Mercury is used in thermometers, manometers, sphygmomanometers, float valves, mercury switches, mercury relays, fluorescent lamps and other devices, though concerns about the element's toxicity have led to mercury thermometers and sphygmomanometers being phased out in clinical environments in favor of alternatives such as alcohol- or galinstan-filled glass thermometers and thermistor- or infrared-based electronic instruments. Mechanical pressure gauges and electronic strain gauge sensors have replaced mercury sphygmomanometers. Mercury remains in use in scientific research applications and in amalgam for dental restoration in some locales.
It is used in fluorescent lighting. Electricity passed through mercury vapor in a fluorescent lamp produces short-wave ultraviolet light, which causes the phosphor in the tube to fluoresce, making visible light. Mercury poisoning can result from exposure to water-soluble forms of mercury, by inhalation of mercury vapor, or by ingesting any form of mercury. Mercury is a silvery-white liquid metal. Compared to other metals, it is a fair conductor of electricity, it has a freezing point of −38.83 °C and a boiling point of 356.73 °C, both the lowest of any stable metal, although preliminary experiments on copernicium and flerovium have indicated that they have lower boiling points. Upon freezing, the volume of mercury decreases by 3.59% and its density changes from 13.69 g/cm3 when liquid to 14.184 g/cm3 when solid. The coefficient of volume expansion is 181.59 × 10−6 at 0 °C, 181.71 × 10−6 at 20 °C and 182.50 × 10−6 at 100 °C. Solid mercury can be cut with a knife. A complete explanation of mercury's extreme volatility delves deep into the realm of quantum physics, but it can be summarized as follows: mercury has a unique electron configuration where electrons fill up all the available 1s, 2s, 2p, 3s, 3p, 3d, 4s, 4p, 4d, 4f, 5s, 5p, 5d, 6s subshells.
Because this configuration resists removal of an electron, mercury behaves to noble gases, which form weak bonds and hence melt at low temperatures. The stability of the 6s shell is due to the presence of a filled 4f shell. An f shell poorly screens the nuclear charge that increases the attractive Coulomb interaction of the 6s shell and the nucleus; the absence of a filled inner f shell is the reason for the somewhat higher melting temperature of cadmium and zinc, although both these metals still melt and, in addition, have unusually low boiling points. Mercury does not react with most acids, such as dilute sulfuric acid, although oxidizing acids such as concentrated sulfuric acid and nitric acid or aqua regia dissolve it to give sulfate and chloride. Like silver, mercury reacts with atmospheric hydrogen sulfide. Mercury reacts with solid sulfur flakes. Mercury dissolves many metals such as silver to form amalgams. Iron is an exception, iron flasks have traditionally been used to trade mercury.
Several other first row transition metals with the exception of manganese and zinc are resistant in forming amalgams. Other elements that do not form amalgams with mercury include platinum. Sodium amalgam is a common reducing agent in organic synthesis, is used in high-pressure sodium lamps. Mercury combines with aluminium to form a mercury-aluminium amalgam when the two pure metals come into contact. Since the amalgam destroys the aluminium oxide layer which protects metallic aluminium from oxidizing in-depth small amounts of mercury can corrode aluminium. For this reason, mercury is not allowed aboard an aircraft under most circumstances because of the risk of it forming an amalgam with exposed aluminium parts in the aircraft. Mercury embrittlement is the most common type of liquid metal embrittlement. There are seven stable isotopes of mercury, with 202Hg being the most abundant; the longest-lived radioisotopes are 194Hg with a half-life of 444 years, 203Hg with a half-life of 46.612 days. Most of the remaining radioisotopes have half-lives.
199Hg and 201Hg are the most studied NMR-active nuclei, having spins of 1⁄2 and 3⁄2 respectively. Hg is the modern chemical symbol for mercury, it comes from hydrargyrum, a Latinized form of the Greek word ὑδράργυρος, a compound word meaning "water-silver" – since it is liquid like water and shiny like silver. The element was named after the Roman god Mercury, known for his mobility, it is associated with the planet Mercury. Mercury is the only metal for which the al
Mercury chloride is the chemical compound with the formula Hg2Cl2. Known as the mineral calomel or mercurous chloride, this dense white or yellowish-white, odorless solid is the principal example of a mercury compound, it is a component of reference electrodes in electrochemistry. The name calomel is thought to come from the Greek καλός beautiful, μέλας black; the black name is due to its characteristic disproportionation reaction with ammonia, which gives a “spectacular” black coloration due to the finely dispersed metallic mercury formed. It is referred to as the mineral horn quicksilver or horn mercury. Calomel was taken internally and used as a laxative, for example to treat George III in 1801, disinfectant, as well as in the treatment of syphilis, until the early 20th century; until recently, it was used as a horticultural fungicide, most notably as a root dip to help prevent the occurrence of clubroot amongst crops of the Brassicaceae family. Mercury became a popular remedy for a variety of physical and mental ailments during the age of "heroic medicine".
It was used by doctors in America throughout the 18th century, during the revolution, to make patients regurgitate and release their body from "impurities". Benjamin Rush was one particular well-known advocate of mercury in medicine and used calomel to treat sufferers of yellow fever during its outbreak in Philadelphia in 1793. Calomel was given to patients as a purgative or cathartic until they began to salivate and was administered to patients in such great quantities that their hair and teeth fell out. Shortly after yellow fever struck Philadelphia, the disease broke out in Jamaica. A war of words erupted in the press concerning the best treatment for yellow fever: bleeding. Anecdotal evidence indicates. Mormon prophet Joseph Smith's eldest brother Alvin Smith died in 1823 from mercury poisoning from calomel. Lewis and Clark brought along the wonder drug of the day, mercury chloride, as a pill, a tincture, an ointment. Modern researchers used that same mercury, found deep in latrine pits, to retrace the locations of their respective locations and campsites.
Mercury is unique among the group 12 metals for its ability to form the M–M bond so readily. Hg2Cl2 is a linear molecule; the mineral calomel crystallizes with space group I4/m 2/m 2/m. The unit cell of the crystal structure is shown below: The Hg–Hg bond length of 253 pm and the Hg–Cl bond length in the linear Hg2Cl2 unit is 243 pm; the overall coordination of each Hg atom is octahedral as, in addition to the two nearest neighbours, there are four other Cl atoms at 321 pm. Longer mercury polycations exist. Mercurous chloride forms by the reaction of elemental mercury and mercuric chloride: Hg + HgCl2 → Hg2Cl2It can be prepared via metathesis reaction involving aqueous mercury nitrate using various chloride sources including NaCl or HCl. 2HCl + Hg22 → Hg2Cl2 + 2HNO3Ammonia causes Hg2Cl2 to disproportionate: Hg2Cl2 + 2NH3 → Hg + HgCl + NH4Cl Mercurous chloride is employed extensively in electrochemistry, taking advantage of the ease of its oxidation and reduction reactions. The calomel electrode is a reference electrode in older publications.
Over the past 50 years, it has been superseded by the silver/silver chloride electrode. Although the mercury electrodes have been abandoned due to the dangerous nature of mercury, many chemists believe they are still more accurate and are not dangerous as long as they are handled properly; the differences in experimental potentials vary little from literature values. Other electrodes can vary by 70 to 100 millivolts. Mercurous chloride decomposes into mercury elemental mercury upon exposure to UV light. Hg2Cl2 → HgCl2 + HgThe formation of Hg can be used to calculate the number of photons in the light beam, by the technique of actinometry. By utilizing a light reaction in the presence of mercury chloride and ammonium oxalate, mercury chloride, ammonium chloride and carbon dioxide are produced. 2HgCl2 + 2C2O4 + Light → Hg2Cl2 + 2 + 2CO2This particular reaction was discovered by J. M. Eder in 1880 and reinvestigated by W. E. Rosevaere in 1929. Mercury bromide, Hg2Br2, is a light yellow, whereas Hg2I2, is greenish in colour.
Both are poorly soluble. Mercury fluoride is unstable in the absence of a strong acid. Mercurous chloride is toxic, although due to its low solubility in water it is less dangerous than its mercuric chloride counterpart, it was used in medicine as a diuretic and purgative in the United States from the late 1700s through the 1860s. Calomel was a common ingredient in teething powders in Britain up until 1954, causing widespread mercury poisoning in the form of pink disease, which at the time had a mortality rate of 1 in 10; these medicinal uses were discontinued when the compound's toxicity was discovered. It has found uses in cosmetics as soaps and skin lightening creams, but these preparations are now illegal to manufacture or import in many countries including the US, Canada and the European Union. A study of workers involved in the production of these preparations showed that the sodium salt of 2,3-dimercapto-1-propanesulfonic acid was effective in lowering the body burden of mercury and in decreasing the urinary mercury concentration to normal levels.
International Chemical Safety Card 0984 National Pollutant Inventory - Mercury and compounds Fact Sheet NIOSH Pocket Guide to Chemical Hazards
Electrochemistry is the branch of physical chemistry that studies the relationship between electricity, as a measurable and quantitative phenomenon, identifiable chemical change, with either electricity considered an outcome of a particular chemical change or vice versa. These reactions involve electric charges moving between an electrolyte, thus electrochemistry deals with the interaction between electrical energy and chemical change. When a chemical reaction is caused by an externally supplied current, as in electrolysis, or if an electric current is produced by a spontaneous chemical reaction as in a battery, it is called an electrochemical reaction. Chemical reactions where electrons are transferred directly between molecules and/or atoms are called oxidation-reduction or reactions. In general, electrochemistry describes the overall reactions when individual redox reactions are separate but connected by an external electric circuit and an intervening electrolyte. Understanding of electrical matters began in the sixteenth century.
During this century, the English scientist William Gilbert spent 17 years experimenting with magnetism and, to a lesser extent, electricity. For his work on magnets, Gilbert became known as the "Father of Magnetism." He discovered various methods for strengthening magnets. In 1663, the German physicist Otto von Guericke created the first electric generator, which produced static electricity by applying friction in the machine; the generator was made of a large sulfur ball cast inside a glass globe, mounted on a shaft. The ball was rotated by means of a crank and an electric spark was produced when a pad was rubbed against the ball as it rotated; the globe could be used as source for experiments with electricity. By the mid—18th century the French chemist Charles François de Cisternay du Fay had discovered two types of static electricity, that like charges repel each other whilst unlike charges attract. Du Fay announced that electricity consisted of two fluids: positive, electricity; this was the two-fluid theory of electricity, to be opposed by Benjamin Franklin's one-fluid theory in the century.
In 1785, Charles-Augustin de Coulomb developed the law of electrostatic attraction as an outgrowth of his attempt to investigate the law of electrical repulsions as stated by Joseph Priestley in England. In the late 18th century the Italian physician and anatomist Luigi Galvani marked the birth of electrochemistry by establishing a bridge between chemical reactions and electricity on his essay "De Viribus Electricitatis in Motu Musculari Commentarius" in 1791 where he proposed a "nerveo-electrical substance" on biological life forms. In his essay Galvani concluded that animal tissue contained a here-to-fore neglected innate, vital force, which he termed "animal electricity," which activated nerves and muscles spanned by metal probes, he believed that this new force was a form of electricity in addition to the "natural" form produced by lightning or by the electric eel and torpedo ray as well as the "artificial" form produced by friction. Galvani's scientific colleagues accepted his views, but Alessandro Volta rejected the idea of an "animal electric fluid," replying that the frog's legs responded to differences in metal temper and bulk.
Galvani refuted this by obtaining muscular action with two pieces of the same material. In 1800, William Nicholson and Johann Wilhelm Ritter succeeded in decomposing water into hydrogen and oxygen by electrolysis. Soon thereafter Ritter discovered the process of electroplating, he observed that the amount of metal deposited and the amount of oxygen produced during an electrolytic process depended on the distance between the electrodes. By 1801, Ritter observed thermoelectric currents and anticipated the discovery of thermoelectricity by Thomas Johann Seebeck. By the 1810s, William Hyde Wollaston made improvements to the galvanic cell. Sir Humphry Davy's work with electrolysis led to the conclusion that the production of electricity in simple electrolytic cells resulted from chemical action and that chemical combination occurred between substances of opposite charge; this work led directly to the isolation of sodium and potassium from their compounds and of the alkaline earth metals from theirs in 1808.
Hans Christian Ørsted's discovery of the magnetic effect of electric currents in 1820 was recognized as an epoch-making advance, although he left further work on electromagnetism to others. André-Marie Ampère repeated Ørsted's experiment, formulated them mathematically. In 1821, Estonian-German physicist Thomas Johann Seebeck demonstrated the electrical potential in the juncture points of two dissimilar metals when there is a heat difference between the joints. In 1827, the German scientist Georg Ohm expressed his law in this famous book "Die galvanische Kette, mathematisch bearbeitet" in which he gave his complete theory of electricity. In 1832, Michael Faraday's experiments led him to state his two laws of electrochemistry. In 1836, John Daniell invented a primary cell which solved the problem of polarization by eliminating hydrogen gas generation at the positive electrode. Results revealed that alloying the amalgamated zinc with mercury would produce a higher voltage. William Grove produced the first fuel cell in 1839.
In 1846, Wilhelm Weber developed the electrodynamometer. In 1868, Georges Leclanché patented a new cell which became the forerunner to the world's first used battery, the zinc carbon cell. Svante Arrhenius published
A pH meter is a scientific instrument that measures the hydrogen-ion activity in water-based solutions, indicating its acidity or alkalinity expressed as pH. The pH meter measures the difference in electrical potential between a pH electrode and a reference electrode, so the pH meter is sometimes referred to as a "potentiometric pH meter"; the difference in electrical potential relates to the pH of the solution. The pH meter is used in many applications ranging from laboratory experimentation to quality control; the rate and outcome of chemical reactions taking place in water depends on the acidity of the water, it is therefore useful to know the acidity of the water measured by means of a pH meter. Knowledge of pH is critical in many situations, including chemical laboratory analyses. PH meters are used for soil measurements in agriculture, water quality for municipal water supplies, swimming pools, environmental remediation. Advances in the instrumentation and in detection have expanded the number of applications in which pH measurements can be conducted.
The devices have been miniaturized. In addition to measuring the pH of liquids, specially designed electrodes are available to measure the pH of semi-solid substances, such as foods; these have tips suitable for piercing semi-solids, have electrode materials compatible with ingredients in food, are resistant to clogging. Potentiometric pH meters measure the voltage between two electrodes and display the result converted into the corresponding pH value, they comprise a simple electronic amplifier and a pair of electrodes, or alternatively a combination electrode, some form of display calibrated in pH units. It has a glass electrode and a reference electrode, or a combination electrode; the electrodes, or probes, are inserted into the solution to be tested. The design of the electrodes is the key part: These are rod-like structures made of glass, with a bulb containing the sensor at the bottom; the glass electrode for measuring the pH has a glass bulb designed to be selective to hydrogen-ion concentration.
On immersion in the solution to be tested, hydrogen ions in the test solution exchange for other positively charged ions on the glass bulb, creating an electrochemical potential across the bulb. The electronic amplifier detects the difference in electrical potential between the two electrodes generated in the measurement and converts the potential difference to pH units; the magnitude of the electrochemical potential across the glass bulb is linearly related to the pH according to the Nernst equation. The reference electrode is insensitive to the pH of the solution, being composed of a metallic conductor, which connects to the display; this conductor is immersed in an electrolyte solution potassium chloride, which comes into contact with the test solution through a porous ceramic membrane. The display consists of a voltmeter, which displays voltage in units of pH. On immersion of the glass electrode and the reference electrode in the test solution, an electrical circuit is completed, in which there is a potential difference created and detected by the voltmeter.
The circuit can be thought of as going from the conductive element of the reference electrode to the surrounding potassium-chloride solution, through the ceramic membrane to the test solution, the hydrogen-ion-selective glass of the glass electrode, to the solution inside the glass electrode, to the silver of the glass electrode, the voltmeter of the display device. The voltage varies from test solution to test solution depending on the potential difference created by the difference in hydrogen-ion concentrations on each side of the glass membrane between the test solution and the solution inside the glass electrode. All other potential differences in the circuit do not vary with pH and are corrected for by means of the calibration. For simplicity, many pH meters use a combination probe, constructed with the glass electrode and the reference electrode contained within a single probe. A detailed description of combination electrodes is given in the article on glass electrodes; the pH meter is calibrated with solutions of known pH before each use, to ensure accuracy of measurement.
To measure the pH of a solution, the electrodes are used as probes, which are dipped into the test solutions and held there sufficiently long for the hydrogen ions in the test solution to equilibrate with the ions on the surface of the bulb on the glass electrode. This equilibration provides a stable pH measurement. Details of the fabrication and resulting microstructure of the glass membrane of the pH electrode are maintained as trade secrets by the manufacturers. However, certain aspects of design are published. Glass is a solid electrolyte; the pH-sensitive glass membrane is spherical to simplify manufacture of a uniform membrane. These membranes are up to 0.4 millimeters in thickness, thicker than original designs, so as to render the probes durable. The glass has silicate chemical functionality on its surface, which provides binding sites for alkali-metal ions and hydrogen ions from the solutions; this provides an ion-exchange capacity in the range of 10−6 to 10−8 mol/cm2. Selectivity for hydrogen ions arises from a balance of ionic charge, volume requirements versus other ions, the coordination number of other ions.
Electrode manufacturers have developed compositions that suitably balance these factors, most notably lithium glass. The silver chloride electrode is most used as a reference electrode in pH meters, although so
A salt bridge, in electrochemistry, is a laboratory device used to connect the oxidation and reduction half-cells of a galvanic cell, a type of electrochemical cell. It maintains electrical neutrality within the internal circuit, preventing the cell from running its reaction to equilibrium. If no salt bridge were present, the solution in one half cell would accumulate negative charge and the solution in the other half cell would accumulate positive charge as the reaction proceeded preventing further reaction, hence production of electricity. Salt bridges come in two types: glass tube and filter paper. One type of salt bridge consists of a U-shaped glass tube filled with a inert electrolyte; the electrolyte is so chosen that it does not react with any of the chemicals used in the cell the anion and cation have similar conductivity, hence similar migratory speed. The electrolyte is gelified with agar-agar to help prevent the intermixing of fluids which might otherwise occur; the conductivity of a glass tube bridge depends on the concentration of the electrolyte solution.
At concentrations below saturation, an increase in concentration increases conductivity. Beyond-saturation electrolyte content and narrow tube diameter may both lower conductivity; the other type of salt bridge consists of a filter paper soaked with a inert electrolyte potassium chloride or sodium chloride because they are chemically inert. No gelification agent is required. Conductivity of this kind of salt bridge depends on a number of factors: the concentration of the electrolyte solution, the texture of the filter paper and the absorbing ability of the filter paper. Smoother texture and higher absorbency equates to higher conductivity. A porous disk or other porous barriers between the two half-cells may be used instead of a salt bridge. Liquid junction potential Ion transport number
Silver chloride electrode
A silver chloride electrode is a type of reference electrode used in electrochemical measurements. For environmental reasons it has replaced the saturated calomel electrode. For example, it is the internal reference electrode in pH meters and it is used as reference in reduction potential measurements; as an example of the latter, the silver chloride electrode is the most used reference electrode for testing cathodic protection corrosion control systems in sea water environments. The electrode functions as a redox electrode and the equilibrium is between the silver metal and its salt—silver chloride; the corresponding half-reactions can be presented as follows: Ag + + e − ↽ − − ⇀ Ag AgCl + e − ↽ − − ⇀ Ag + Cl − or can be written together: AgCl + Ag + e − ↽ − − ⇀ Ag + e − + Cl − + Ag + which can be simplified: AgCl ↽ − − ⇀ Ag + + Cl − This reaction is characterized by fast electrode kinetics, meaning that a sufficiently high current can be passed through the electrode with the 100% efficiency of the redox reaction.
The reaction has been proven to obey these equations in solutions of pH values between 0 and 13.5. The Nernst equation below shows the dependence of the potential of the silver-silver chloride electrode on the activity or effective concentration of chloride-ions: E = E 0 − R T F ln a Cl − The standard electrode potential E0 against standard hydrogen electrode is 0.230 V ± 10 mV. The potential is however sensitive to traces of bromide ions which make it more negative. Commercial reference electrodes consist of a plastic tube electrode body; the electrode is a silver wire, coated with a thin layer of silver chloride, either physically by dipping the wire in molten silver chloride, chemically by electroplating the wire in concentrated hydrochloric acid or electrochemically by oxidising the silver in a chloride solution. A porous plug on one end allows contact between the field environment with the silver chloride electrolyte. An insulated lead wire connects the silver rod with measuring instruments.
A voltmeter negative lead is connected to the test wire. The electrode body contains potassium chloride to stabilize the silver chloride concentration; when working in seawater, this body can be removed and the chloride concentration is fixed by the stable salinity of the water. The potential of a silver:silver chloride reference electrode with respect to standard hydrogen electrode depends on the electrolyte composition. Notes to the Table: The table data source is, except where a separate reference is given. Elj is the potential of the liquid junction between the given electrolyte and the electrolyte with the activity of chloride of 1 mol/kg; the electrode has many features making it suitable for use in the field: Simple construction Inexpensive to manufacture Stable potential Non-toxic componentsThey are manufactured with saturated potassium chloride electrolyte, but can be used with lower concentrations such as 1 mol/kg potassium chloride. As noted above, changing the electrolyte concentration changes the electrode potential.
Silver chloride is soluble in strong potassium chloride solutions, so it is sometimes recommended the potassium chloride be saturated with silver chloride to avoid stripping the silver chloride off the silver wire. Silver chloride electrodes are used by many applications of biological electrode systems such as biomonitoring sensors as part of electrocardiography and electroencephalography, in transcutaneous electrical nerve stimulation to deliver current; the electrodes were fabricated from solid materials such as silver, brass coated with silver and nickel. In today's applications, most biomonitoring electrodes are silver/silver chloride sensors which are fabricated by coating a thin layer of silver on plastic substrates and the outer layer of silver is converted to silver chloride; the principle of silver/silver chloride sensors operation is the conversion of ion current at the surface of human tissues to electron current to be delivered through the lead wire to the instrument to read. An im