A solvent is a substance that dissolves a solute, resulting in a solution. A solvent is a liquid but can be a solid, a gas, or a supercritical fluid; the quantity of solute that can dissolve in a specific volume of solvent varies with temperature. Common uses for organic solvents are in dry cleaning, as paint thinners, as nail polish removers and glue solvents, in spot removers, in detergents and in perfumes. Water is a solvent for the most common solvent used by living things. Solvents find various applications in chemical, pharmaceutical and gas industries, including in chemical syntheses and purification processes; when one substance is dissolved into another, a solution is formed. This is opposed to the situation. In a solution, all of the ingredients are uniformly distributed at a molecular level and no residue remains. A solvent-solute mixture consists of a single phase with all solute molecules occurring as solvates, as opposed to separate continuous phases as in suspensions and other types of non-solution mixtures.
The ability of one compound to be dissolved in another is known as solubility. In addition to mixing, the substances in a solution interact with each other at the molecular level; when something is dissolved, molecules of the solvent arrange around molecules of the solute. Heat transfer is involved and entropy is increased making the solution more thermodynamically stable than the solute and solvent separately; this arrangement is mediated by the respective chemical properties of the solvent and solute, such as hydrogen bonding, dipole moment and polarizability. Solvation does not cause a chemical chemical configuration changes in the solute. However, solvation resembles a coordination complex formation reaction with considerable energetics and is thus far from a neutral process. Solvents can be broadly classified into two categories: non-polar. A special case is mercury; the dielectric constant of the solvent provides a rough measure of a solvent's polarity. The strong polarity of water is indicated by its high dielectric constant of 88.
Solvents with a dielectric constant of less than 15 are considered to be nonpolar. The dielectric constant measures the solvent's tendency to cancel the field strength of the electric field of a charged particle immersed in it; this reduction is compared to the field strength of the charged particle in a vacuum. Heuristically, the dielectric constant of a solvent can be thought of as its ability to reduce the solute's effective internal charge; the dielectric constant of a solvent is an acceptable predictor of the solvent's ability to dissolve common ionic compounds, such as salts. Dielectric constants are not the only measure of polarity; because solvents are used by chemists to carry out chemical reactions or observe chemical and biological phenomena, more specific measures of polarity are required. Most of these measures are sensitive to chemical structure; the Grunwald–Winstein mY scale measures polarity in terms of solvent influence on buildup of positive charge of a solute during a chemical reaction.
Kosower's Z scale measures polarity in terms of the influence of the solvent on UV-absorption maxima of a salt pyridinium iodide or the pyridinium zwitterion. Donor number and donor acceptor scale measures polarity in terms of how a solvent interacts with specific substances, like a strong Lewis acid or a strong Lewis base; the Hildebrand parameter is the square root of cohesive energy density. It can not accommodate complex chemistry. Reichardt's dye, a solvatochromic dye that changes color in response to polarity, gives a scale of ET values. ET is the transition energy between the ground state and the lowest excited state in kcal/mol, identifies the dye. Another correlated scale can be defined with Nile red; the polarity, dipole moment and hydrogen bonding of a solvent determines what type of compounds it is able to dissolve and with what other solvents or liquid compounds it is miscible. Polar solvents dissolve polar compounds best and non-polar solvents dissolve non-polar compounds best: "like dissolves like".
Polar compounds like sugars or ionic compounds, like inorganic salts dissolve only in polar solvents like water, while non-polar compounds like oils or waxes dissolve only in non-polar organic solvents like hexane. Water and hexane are not miscible with each other and will separate into two layers after being shaken well. Polarity can be separated to different contributions. For example, the Kamlet-Taft parameters are dipolarity/polarizability, hydrogen-bonding acidity and hydrogen-bonding basicity; these can be calculated from the wavelength shifts of 3–6 different solvatochromic dyes in the solvent including Reichardt's dye and diethylnitroaniline. Another option, Hansen's parameters, separate the cohesive energy density into dispersion and hydrogen bonding contributions. Solvents with a dielectric constant (more relative
Uranium hexafluoride, colloquially known as "hex" in the nuclear industry, is a compound used in the process of enriching uranium, which produces fuel for nuclear reactors and nuclear weapons. Hex forms solid grey crystals at standard temperature and pressure, is toxic, reacts with water, is corrosive to most metals; the compound reacts mildly with aluminium, forming a thin surface layer of AlF3 that resists any further reaction from the compound. Milled uranium ore—U3O8 or "yellowcake"—is dissolved in nitric acid, yielding a solution of uranyl nitrate UO22. Pure uranyl nitrate is obtained by solvent extraction treated with ammonia to produce ammonium diuranate. Reduction with hydrogen gives UO2, converted with hydrofluoric acid to uranium tetrafluoride, UF4. Oxidation with fluorine yields UF6. During nuclear reprocessing, uranium is reacted with chlorine trifluoride to give UF6: U + 2 ClF3 → UF6 + Cl2 At atmospheric pressure, it sublimes at 56.5 °C. The solid state structure was determined by neutron diffraction at 77 K and 293 K.
It has been shown that uranium hexafluoride is an oxidant and a Lewis acid, able to bind to fluoride. Polymeric uranium fluorides containing organic cations have been isolated and characterised by X-ray diffraction. UF6 is used in both of the main uranium enrichment methods — gaseous diffusion and the gas centrifuge method — because its triple point is at temperature 64.05 °C and only higher than normal atmospheric pressure. Fluorine has only a single occurring stable isotope, so isotopologues of UF6 differ in their molecular weight based on the uranium isotope present. All the other uranium fluorides are nonvolatile solids. Gaseous diffusion requires about 60 times as much energy as the gas centrifuge process: gaseous diffusion-produced nuclear fuel produces 25 times more energy than is used in the diffusion process, while centrifuge-produced fuel produces 1,500 times more energy than is used in the centrifuge process. In addition to its use in enrichment, uranium hexafluoride has been used in an advanced reprocessing method, developed in the Czech Republic.
In this process, used oxide nuclear fuel is treated with fluorine gas to form a mixture of fluorides. This mixture is distilled to separate the different classes of material. Uranium enrichment produces large quantities of depleted uranium hexafluoride, or DUF6, as a waste product; the long-term storage of DUF6 presents environmental and safety risks because of its chemical instability. When UF6 is exposed to moist air, it reacts with the water in the air to produce UO2F2 and HF both of which are corrosive and toxic. In 2005, 686,500 tonnes of DUF6 was housed in 57,122 storage cylinders located near Portsmouth, Ohio. Storage cylinders must be inspected for signs of corrosion and leaks; the estimated lifetime of the steel cylinders is measured in decades. There have been several accidents involving uranium hexafluoride in the US, including a cylinder-filling accident and material release at the Sequoyah Fuels Corporation in 1986; the U. S. government has been converting DUF6 to solid uranium oxides for disposal.
Such disposal of the entire DUF6 inventory could cost anywhere from $15 million to $450 million. Gmelins Handbuch der anorganischen Chemie, System Nr. 55, Teil A, p. 121–123. Gmelins Handbuch der anorganischen Chemie, System Nr. 55, Teil C 8, p. 71–163. R. DeWitt: Uranium hexafluoride: A survey of the physico-chemical properties, Technical report, GAT-280. Portsmouth, Ohio. August 1960. Ingmar Grenthe, Janusz Drożdżynński, Takeo Fujino, Edgar C. Buck, Thomas E. Albrecht-Schmitt, Stephen F. Wolf: Uranium, in: Lester R. Morss, Norman M. Edelstein, Jean Fuger: The Chemistry of the Actinide and Transactinide Elements, Dordrecht 2006. US-Patent 2535572: Preparation of UF6. December 1950. US-Patent 5723837: Uranium Hexafluoride Purification. March 1998. Simon Cotton: Uranium Hexafluoride. Uranium Hexafluoride – Physical and chemical properties of UF6, its use in uranium processing – Uranium Hexafluoride and Its Properties Import of Western depleted uranium hexafluoride to Russia Uranium Hexafluoride in www.webelements.com
Fluorine is a chemical element with symbol F and atomic number 9. It is the lightest halogen and exists as a toxic pale yellow diatomic gas at standard conditions; as the most electronegative element, it is reactive, as it reacts with all other elements, except for helium and neon. Among the elements, fluorine ranks 24th in universal 13th in terrestrial abundance. Fluorite, the primary mineral source of fluorine which gave the element its name, was first described in 1529. Proposed as an element in 1810, fluorine proved difficult and dangerous to separate from its compounds, several early experimenters died or sustained injuries from their attempts. Only in 1886 did French chemist Henri Moissan isolate elemental fluorine using low-temperature electrolysis, a process still employed for modern production. Industrial production of fluorine gas for uranium enrichment, its largest application, began during the Manhattan Project in World War II. Owing to the expense of refining pure fluorine, most commercial applications use fluorine compounds, with about half of mined fluorite used in steelmaking.
The rest of the fluorite is converted into corrosive hydrogen fluoride en route to various organic fluorides, or into cryolite, which plays a key role in aluminium refining. Molecules containing a Carbon–fluorine bond have high chemical and thermal stability. Pharmaceuticals such as atorvastatin and fluoxetine contain C-F bonds, the fluoride ion inhibits dental cavities, so finds use in toothpaste and water fluoridation. Global fluorochemical sales amount to more than US$15 billion a year. Fluorocarbon gases are greenhouse gases with global-warming potentials 100 to 20,000 times that of carbon dioxide. Organofluorine compounds persist in the environment due to the strength of the carbon–fluorine bond. Fluorine has no known metabolic role in mammals. Fluorine atoms have nine electrons, one fewer than neon, electron configuration 1s22s22p5: two electrons in a filled inner shell and seven in an outer shell requiring one more to be filled; the outer electrons are ineffective at nuclear shielding, experience a high effective nuclear charge of 9 − 2 = 7.
Fluorine's first ionization energy is third-highest among all elements, behind helium and neon, which complicates the removal of electrons from neutral fluorine atoms. It has a high electron affinity, second only to chlorine, tends to capture an electron to become isoelectronic with the noble gas neon. Fluorine atoms have a small covalent radius of around 60 picometers, similar to those of its period neighbors oxygen and neon; the bond energy of difluorine is much lower than that of either Cl2 or Br2 and similar to the cleaved peroxide bond. Conversely, bonds to other atoms are strong because of fluorine's high electronegativity. Unreactive substances like powdered steel, glass fragments, asbestos fibers react with cold fluorine gas. Reactions of elemental fluorine with metals require varying conditions. Alkali metals cause; some solid nonmetals react vigorously in liquid air temperature fluorine. Hydrogen sulfide and sulfur dioxide combine with fluorine, the latter sometimes explosively. Hydrogen, like some of the alkali metals, reacts explosively with fluorine.
Carbon, as lamp black, reacts at room temperature to yield fluoromethane. Graphite combines with fluorine above 400 °C to produce non-stoichiometric carbon monofluoride. Carbon dioxide and carbon monoxide react at or just above room temperature, whereas paraffins and other organic chemicals generate strong reactions: fully substituted haloalkanes such as carbon tetrachloride incombustible, may explode. Although nitrogen trifluoride is stable, nitrogen requires an electric discharge at elevated temperatures for reaction with fluorine to occur, due to the strong triple bond in elemental nitrogen. Oxygen does not combine with fluorine under ambient conditions, but can be made to react using electric discharge at low temperatures and pressures. Heavier halogens react with fluorine as does the noble gas radon. At room temperature, fluorine is a gas of diatomic molecules, pale yellow, it has a characteristic halogen-like biting odor detectable at 20 ppb. Fluorine condenses into a bright yellow liquid at −188 °C, a transition temperature similar to those of oxygen and nitrogen.
Fluorine has two solid forms, α- and β-fluorine. The latter crystallizes at −220 °C and is transparent and sof
In organic chemistry, an alkane, or paraffin, is an acyclic saturated hydrocarbon. In other words, an alkane consists of hydrogen and carbon atoms arranged in a tree structure in which all the carbon–carbon bonds are single. Alkanes have the general chemical formula CnH2n+2; the alkanes range in complexity from the simplest case of methane, where n = 1, to arbitrarily large and complex molecules, like pentacontane or 6-ethyl-2-methyl-5- octane, an isomer of tetradecane. IUPAC defines alkanes as "acyclic branched or unbranched hydrocarbons having the general formula CnH2n+2, therefore consisting of hydrogen atoms and saturated carbon atoms". However, some sources use the term to denote any saturated hydrocarbon, including those that are either monocyclic or polycyclic, despite their having a different general formula. In an alkane, each carbon atom is sp3-hybridized with 4 sigma bonds, each hydrogen atom is joined to one of the carbon atoms; the longest series of linked carbon atoms in a molecule is known as its carbon skeleton or carbon backbone.
The number of carbon atoms may be considered as the size of the alkane. One group of the higher alkanes are waxes, solids at standard ambient temperature and pressure, for which the number of carbon atoms in the carbon backbone is greater than about 17. With their repeated –CH2 units, the alkanes constitute a homologous series of organic compounds in which the members differ in molecular mass by multiples of 14.03 u. Alkanes are not reactive and have little biological activity, they can be viewed as molecular trees upon which can be hung the more active/reactive functional groups of biological molecules. The alkanes have two main commercial sources: natural gas. An alkyl group abbreviated with the symbol R, is a functional group that, like an alkane, consists of single-bonded carbon and hydrogen atoms connected acyclically—for example, a methyl or ethyl group. Saturated hydrocarbons are hydrocarbons having only single covalent bonds between their carbons, they can be: linear wherein the carbon atoms are joined in a snake-like structure branched wherein the carbon backbone splits off in one or more directions cyclic wherein the carbon backbone is linked so as to form a loop.
According to the definition by IUPAC, the former two are alkanes, whereas the third group is called cycloalkanes. Saturated hydrocarbons can combine any of the linear and branching structures. Alkanes are the acyclic ones, corresponding to k = 0. Alkanes with more than three carbon atoms can be arranged in various different ways, forming structural isomers; the simplest isomer of an alkane is the one in which the carbon atoms are arranged in a single chain with no branches. This isomer is sometimes called the n-isomer; however the chain of carbon atoms may be branched at one or more points. The number of possible isomers increases with the number of carbon atoms. For example, for acyclic alkanes: C1: methane only C2: ethane only C3: propane only C4: 2 isomers: n-butane and isobutane C5: 3 isomers: pentane and neopentane C6: 5 isomers: hexane, 2-methylpentane, 3-methylpentane, 2,2-dimethylbutane, 2,3-dimethylbutane C12: 355 isomers C32: 27,711,253,769 isomers C60: 22,158,734,535,770,411,074,184 isomers, many of which are not stable.
Branched alkanes can be chiral. For example, 3-methylhexane and its higher homologues are chiral due to their stereogenic center at carbon atom number 3. In addition to the alkane isomers, the chain of carbon atoms may form one or more loops; such compounds are called cycloalkanes. Stereoisomers and cyclic compounds are excluded; the IUPAC nomenclature for alkanes is based on identifying hydrocarbon chains. Unbranched, saturated hydrocarbon chains are named systematically with a Greek numerical prefix denoting the number of carbons and the suffix "-ane". In 1866, August Wilhelm von Hofmann suggested systematizing nomenclature by using the whole sequence of vowels a, e, i, o and u to create suffixes -ane, -ene, -ine, -one, -une, for the hydrocarbons CnH2n+2, CnH2n, CnH2n−2, CnH2n−4, CnH2n−6. Now, the first three name hydrocarbons with single and triple bonds, it is impossible to find compounds with more than one IUPAC name. This is because shorter chains attached to longer chains are prefixes and the convention includes brackets.
Numbers in the name, referring to which carbon a group is attached to, should be as low as possible so that 1- is implied and omitted from names of organic compounds with only one side-group. Symmetric compounds will have two ways of arriving at the same name. Straight-chain alkanes are sometimes indicated by the prefix "n -". Although this is not necessary, the usage is still common in cases where there is an important difference in properties between the straight-chain and branched-chain isomers, e.g. n-hexane or 2- or 3-met
A carboxylic acid is an organic compound that contains a carboxyl group. The general formula of a carboxylic acid is R–COOH, with R referring to the rest of the molecule. Carboxylic acids occur widely. Important examples include acetic acid. Deprotonation of a carboxyl group gives a carboxylate anion. Important carboxylate salts are soaps. Carboxylic acids are identified by their trivial names, they have the suffix -ic acid. IUPAC-recommended names exist. For example, butyric acid is butanoic acid by IUPAC guidelines. For nomenclature of complex molecules containing a carboxylic acid, the carboxyl can be considered position one of the parent chain if there are other substituents, for example, 3-chloropropanoic acid. Alternately, it can be named as a "carboxy" or "carboxylic acid" substituent on another parent structure, for example, 2-carboxyfuran; the carboxylate anion of a carboxylic acid is named with the suffix -ate, in keeping with the general pattern of -ic acid and -ate for a conjugate acid and its conjugate base, respectively.
For example, the conjugate base of acetic acid is acetate. Carboxylic acids are polar; because they are both hydrogen-bond acceptors and hydrogen-bond donors, they participate in hydrogen bonding. Together the hydroxyl and carbonyl group forms the functional group carboxyl. Carboxylic acids exist as dimers in nonpolar media due to their tendency to "self-associate". Smaller carboxylic acids are soluble in water, whereas higher carboxylic acids have limited solubility due to the increasing hydrophobic nature of the alkyl chain; these longer chain acids tend to be rather soluble in less-polar solvents such as ethers and alcohols. Hydrophobic carboxylic acids react aqueous sodium hydroxide to give water soluble sodium salts. For example, enathic acid has a small solubility in water, but its sodium salt is soluble in water: Carboxylic acids tend to have higher boiling points than water, not only because of their increased surface area, but because of their tendency to form stabilised dimers through hydrogen bonds.
For boiling to occur, either the dimer bonds must be broken or the entire dimer arrangement must be vaporised, both of which increase the enthalpy of vaporization requirements significantly. Carboxylic acids are Brønsted -- Lowry acids, they are the most common type of organic acid. Carboxylic acids are weak acids, meaning that they only dissociate into H3O+ cations and RCOO− anions in neutral aqueous solution. For example, at room temperature, in a 1-molar solution of acetic acid, only 0.4% of the acid are dissociated. Electron-withdrawing substituents, such as -CF3 group, give stronger acids. Electron-donating substituents give weaker acids Deprotonation of carboxylic acids gives carboxylate anions; each of the carbon–oxygen bonds in the carboxylate anion has a partial double-bond character. The carbonyl carbon's partial positive charge is weakened by the -1/2 negative charges on the 2 oxygen atoms. Carboxylic acids have strong sour odors. Esters of carboxylic acids tend to have pleasant odors, many are used in perfume.
Carboxylic acids are identified as such by infrared spectroscopy. They exhibit a sharp band associated with vibration of the C–O vibration bond between 1680 and 1725 cm−1. A characteristic νO–H band appears as a broad peak in the 2500 to 3000 cm−1 region. By 1H NMR spectrometry, the hydroxyl hydrogen appears in the 10–13 ppm region, although it is either broadened or not observed owing to exchange with traces of water. Many carboxylic acids are produced industrially on a large scale, they are pervasive in nature. Esters of fatty acids are the main components of lipids and polyamides of aminocarboxylic acids are the main components of proteins. Carboxylic acids are used in the production of polymers, pharmaceuticals and food additives. Industrially important carboxylic acids include acetic acid and methacrylic acids, adipic acid, citric acid, ethylenediaminetetraacetic acid, fatty acids, maleic acid, propionic acid, terephthalic acid. In general, industrial routes to carboxylic acids differ from those used on smaller scale because they require specialized equipment.
Carbonylation of alcohols as illustrated by the Cativa process for production of acetic acid. Formic acid is prepared by a different carbonylation pathway starting from methanol. Oxidation of aldehydes with air using cobalt and manganese catalysts; the required aldehydes are obtained from alkenes by hydroformylation. Oxidation of hydrocarbons using air. For simple alkanes, this method is inexpensive but not selective enough to be useful. Allylic and benzylic compounds undergo more selective oxidations. Alkyl groups on a benzene ring are oxidized to the carboxylic acid, regardless of its chain length. Benzoic acid from toluene, terephthalic acid from para-xylene, phthalic acid from ortho-xylene are illustrative large-scale conversions. Acrylic acid is generated from propene. Base-cata
In organic chemistry, an alkene is an unsaturated hydrocarbon that contains at least one carbon–carbon double bond. The words alkene and olefin are used interchangeably. Acyclic alkenes, with only one double bond and no other functional groups, known as mono-enes, form a homologous series of hydrocarbons with the general formula CnH2n. Alkenes have two hydrogen atoms fewer than the corresponding alkane; the simplest alkene, with the International Union of Pure and Applied Chemistry name ethene, is the organic compound produced on the largest scale industrially. Aromatic compounds are drawn as cyclic alkenes, but their structure and properties are different and they are not considered to be alkenes. Like a single covalent bond, double bonds can be described in terms of overlapping atomic orbitals, except that, unlike a single bond, a carbon–carbon double bond consists of one sigma bond and one pi bond; this double bond is stronger than a single covalent bond and shorter, with an average bond length of 1.33 ångströms.
Each carbon of the double bond uses its three sp2 hybrid orbitals to form sigma bonds to three atoms. The unhybridized 2p atomic orbitals, which lie perpendicular to the plane created by the axes of the three sp² hybrid orbitals, combine to form the pi bond; this bond lies outside the main C–C axis, with half of the bond on one side of the molecule and half on the other. With a strength of 65 kcal/mol, the pi bond is weaker than the sigma bond. Rotation about the carbon–carbon double bond is restricted because it incurs an energetic cost to break the alignment of the p orbitals on the two carbon atoms; as a consequence, substituted alkenes may exist as one of called cis or trans isomers. More complex alkenes may be named with the E–Z notation for molecules with three or four different substituents. For example, of the isomers of butene, the two methyl groups of -but-2-ene appear on the same side of the double bond, in -but-2-ene the methyl groups appear on opposite sides; these two isomers of butene are different in their chemical and physical properties.
Twisting to a 90° dihedral angle between two of the groups on the carbons requires less energy than the strength of a pi bond, the bond still holds. The carbons of the double bond become pyramidal, which allows preserving some p orbital alignment—and hence pi bonding; the other two attached. This contradicts a common textbook assertion that the two carbons retain their planar nature when twisting, in which case the p orbitals would rotate enough away from each other to be unable to sustain a pi bond. In a 90°-twisted alkene, the p orbitals are only misaligned by 42° and the strain energy is only around 40 kcal/mol. In contrast, a broken pi bond has an energetic cost of around 65 kcal/mol; some pyramidal alkenes are stable. For example, trans-cyclooctene is a stable strained alkene and the orbital misalignment is only 19°, despite having a significant dihedral angle of 137° and a degree of pyramidalization of 18°. Trans-cycloheptene is stable at low temperatures; as predicted by the VSEPR model of electron pair repulsion, the molecular geometry of alkenes includes bond angles about each carbon in a double bond of about 120°.
The angle may vary because of steric strain introduced by nonbonded interactions between functional groups attached to the carbons of the double bond. For example, the C–C–C bond angle in propylene is 123.9°. For bridged alkenes, Bredt's rule states that a double bond cannot occur at the bridgehead of a bridged ring system unless the rings are large enough. Following Fawcett and defining S as the total number of non-bridgehead atoms in the rings, bicyclic systems require S ≥ 7 for stability and tricyclic systems require S ≥ 11. Many of the physical properties of alkenes and alkanes are similar: they are colourless and combustable; the physical state depends on molecular mass: like the corresponding saturated hydrocarbons, the simplest alkenes, ethene and butene are gases at room temperature. Linear alkenes of five to sixteen carbons are liquids, higher alkenes are waxy solids; the melting point of the solids increases with increase in molecular mass. Alkenes have stronger smells than the corresponding alkane.
Ethylene is described to have a "sweet" odor, for example. The binding of cupric ion to the olefin in the mammalian olfactory receptor MOR244-3 is implicated in the smell of alkenes. Strained alkenes, in particular, like norbornene and trans-cyclooctene are known to have strong, unpleasant odors, a fact consistent with the stronger π complexes they form with metal ions including copper. Alkenes are stable compounds, but are more reactive than alkanes, either because of the reactivity of the carbon–carbon pi-bond or the presence of allylic CH centers. Most reactions of alkenes involve additions to this pi bond. Alkenes serve as a feedstock for the petrochemical industry because they can participate in a wide variety of reactions, prominently polymerization and alkylation. Alkenes react in many addition reactions. Most of these addition reactions follow the mechanism of electrophilic addition. Examples are hydrohalogenation, halohydrin formation, hydroboration, dichlorocarbene addition, Simmons–Smith reaction, catalytic hydrogenation, epox
In chemistry and manufacturing, electrolysis is a technique that uses a direct electric current to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially important as a stage in the separation of elements from occurring sources such as ores using an electrolytic cell; the voltage, needed for electrolysis to occur is called the decomposition potential. The word "electrolysis" was introduced by Michael Faraday in the 19th century, on the suggestion of the Rev. William Whewell, using the Greek words ἤλεκτρον "amber", which since the 17th century was associated with electric phenomena, λύσις meaning "dissolution". Electrolysis, as a tool to study chemical reactions and obtain pure elements, precedes the coinage of the term and formal description by Faraday. 1785 – Martinus van Marum's electrostatic generator was used to reduce tin and antimony from their salts using electrolysis. 1800 – William Nicholson and Anthony Carlisle, decomposed water into hydrogen and oxygen.
1808 – Potassium, barium and magnesium were discovered by Sir Humphry Davy using electrolysis. 1821 – Lithium was discovered by the English chemist William Thomas Brande, who obtained it by electrolysis of lithium oxide. 1833 – Michael Faraday develops his two laws of electrolysis, provides a mathematical explanation of his laws. 1875 – Paul Émile Lecoq de Boisbaudran discovered gallium using electrolysis. 1886 – Fluorine was discovered by Henri Moissan using electrolysis. 1886 – Hall–Héroult process developed for making aluminium 1890 – Castner–Kellner process developed for making sodium hydroxide Electrolysis is the passing of a direct electric current through an ionic substance, either molten or dissolved in a suitable solvent, producing chemical reactions at the electrodes and a decomposition of the materials. The main components required to achieve electrolysis are: An electrolyte: a substance an ion-conducting polymer that contains free ions, which carry electric current in the electrolyte.
If the ions are not mobile, as in most solid salts electrolysis cannot occur. A direct current electrical supply: provides the energy necessary to create or discharge the ions in the electrolyte. Electric current is carried by electrons in the external circuit. Two electrodes: electrical conductors that provide the physical interface between the electrolyte and the electrical circuit that provides the energy. Electrodes of metal and semiconductor material are used. Choice of suitable electrode depends on chemical reactivity between the electrode and electrolyte and manufacturing cost; the key process of electrolysis is the interchange of atoms and ions by the removal or addition of electrons from the external circuit. The desired products of electrolysis are in a different physical state from the electrolyte and can be removed by some physical processes. For example, in the electrolysis of brine to produce hydrogen and chlorine, the products are gaseous; these are collected. 2 NaCl + 2 H2O → 2 NaOH + H2 + Cl2A liquid containing electrolyte is produced by: Solvation or reaction of an ionic compound with a solvent to produce mobile ions An ionic compound is melted by heatingAn electrical potential is applied across a pair of electrodes immersed in the electrolyte.
Each electrode attracts ions. Positively charged ions move towards the electron-providing cathode. Negatively charged ions move towards the electron-extracting anode. In this process electrons are either released. Neutral atoms gain or lose electrons and become charged ions that pass into the electrolyte; the formation of uncharged atoms from ions is called discharging. When an ion gains or loses enough electrons to become uncharged atoms, the newly formed atoms separate from the electrolyte. Positive metal ions like Cu2+deposit onto the cathode in a layer; the terms for this are electroplating and electrorefining. When an ion gains or loses electrons without becoming neutral, its electronic charge is altered in the process. In chemistry, the loss of electrons is called oxidation. Oxidation of ions or neutral molecules occurs at the anode. For example, it is possible to oxidize ferrous ions to ferric ions at the anode: Fe2+ → Fe3+ + e−Reduction of ions or neutral molecules occurs at the cathode, it is possible to reduce ferricyanide ions to ferrocyanide ions at the cathode: Fe3-6 + e− → Fe4-6Neutral molecules can react at either of the electrodes.
For example: p-Benzoquinone can be reduced to hydroquinone at the cathode: + 2 e− + 2 H+ → In the last example, H+ ions take part in the reaction, are provided by an acid in the solution, or by the solvent itself. Electrolysis reactions involving H+ ions are common in acidic solutions. In aqueous alkaline solutions, reactions involving OH− are common. Sometimes the solvents themselves are reduced at the electrodes, it is possible to have electrolysis involving gases. Such as when using a Gas diffusion electrode; the amount of electrical energy that must be added equals the change in Gibbs free energy of the reaction plus the losses in the system. The losses can be arbitrarily close to zero, so the maximum thermodynamic efficiency equals the enthalpy change divided by the free energy change of the reaction. In most cases, the electric input is larger than the enthalpy change of the reaction, so some energy is released in the form of heat. In some cases, for instance, in