An epoxide is a cyclic ether with a three-atom ring. This ring approximates an equilateral triangle, which makes it strained, hence reactive, more so than other ethers, they are produced on a large scale for many applications. In general, low molecular weight epoxides are colourless and nonpolar, volatile. A compound containing the epoxide functional group can be called an epoxy, epoxide and ethoxyline. Simple epoxides are referred to as oxides. Thus, the epoxide of ethylene is ethylene oxide. Many compounds have trivial names, ethylene oxide is called "oxirane"; some names emphasize the presence of the epoxide functional group, as in the compound 1,2-epoxyheptane, which can be called 1,2-heptene oxide. A polymer formed from epoxide precursors is called an epoxy, but such materials do not contain epoxide groups; the dominant epoxides industrially are ethylene oxide and propylene oxide, which are produced on the scales of 15 and 3 million tonnes/year. The epoxidation of ethylene involves its reaction of oxygen according to the following stoichiometry: 7 H2C=CH2 + 6 O2 → 6 C2H4O + 2 CO2 + 2 H2OThe direct reaction of oxygen with alkenes is useful only for this epoxide.
Modified heterogeneous silver catalysts are employed. Other alkenes fail to react usefully propylene, though TS-1 supported Au catalysts can perform propylene epoxidation selectively. Aside from ethylene oxide, most epoxides are generated by treating alkenes with peroxide-containing reagents, which donate a single oxygen atom. Safety considerations weigh on these reactions because organic peroxides are prone to spontaneous decomposition or combustion. Metal complexes are useful catalysts for epoxidations involving hydrogen peroxide and alkyl hydroperoxides. Peroxycarboxylic acids, which are more electrophilic, convert alkenes to epoxides without the intervention of metal catalysts. In specialized applications, other peroxide-containing reagents are employed, such as dimethyldioxirane. Depending on the mechanism of the reaction and the geometry of the alkene starting material, cis and/or trans epoxide diastereomers may be formed. In addition, if there are other stereocenters present in the starting material, they can influence the stereochemistry of the epoxidation.
Metal-catalyzed epoxidations were first explored using tert-butyl hydroperoxide. Association of TBHP with the metal generates the active metal peroxy complex containing the MOOR group, which transfers an O center to the alkene. Organic peroxides are used for the production of propylene oxide from propylene. Catalysts are required as well. Both t-butyl hydroperoxide and ethylbenzene hydroperoxide can be used as oxygen sources. More for laboratory operations, the Prilezhaev reaction is employed; this approach involves the oxidation of the alkene with a peroxyacid such as m-CPBA. Illustrative is the epoxidation of styrene with perbenzoic acid to styrene oxide: The reaction proceeds via what is known as the "Butterfly Mechanism"; the peroxide is viewed as an electrophile, the alkene a nucleophile. The reaction is considered to be concerted; the butterfly mechanism allows ideal positioning of the O-O sigma star orbital for C-C Pi electrons to attack. Because two bonds are broken and formed to the epoxide oxygen, this is formally an example of a coarctate transition state.
Hydroperoxides are employed in catalytic enantioselective epoxidations, such as the Sharpless epoxidation and the Jacobsen epoxidation. Together with the Shi epoxidation, these reactions are useful for the enantioselective synthesis of chiral epoxides. Oxaziridine reagents may be used to generate epoxides from alkenes. Arene oxides are intermediates in the oxidation of arenes by cytochrome P450. For prochiral arenes, the epoxides are obtained in high enantioselectivity. Chiral epoxides can be derived enantioselectively from prochiral alkenes. Many metal complexes give active catalysts, but the most important involve titanium and molybdenum; the Sharpless epoxidation reaction is one of the premier enantioselective chemical reactions. It is used to prepare 2,3-epoxyalcohols from secondary allylic alcohols; this method involves dehydrohalogenation. It is a variant of the Williamson ether synthesis. In this case, an alkoxide ion intramolecularly displaces chloride; the precursor compounds are called halohydrins.
Starting with propylene chlorohydrin, most of the world's supply of propylene oxide arises via this route. An intramolecular epoxide formation reaction is one of the key steps in the Darzens reaction. In the Johnson–Corey–Chaykovsky reaction epoxides are generated from carbonyl groups and sulfonium ylides. In this reaction, a sulfonium is the leaving group instead of chloride. Electron-deficient olefins, such as enones and acryl derivatives can be epoxidized using nucleophilic oxygen compounds such as peroxides; the reaction is a two-step mechanism. First the oxygen performs a nucleophilic conjugate addition to give a stabilized carbanion; this carbanion attacks the same oxygen atom, displacing a leaving group from it, to close the epoxide ring. Epoxides are uncommon in nature, they arise via oxygenation of alkenes by the action of cytochrome P450. Ring-opening reactions dominate the reactivity of epoxides. Alcohols, amines and many other reagents add to epoxides; this reaction is the basi
The rate law or rate equation for a chemical reaction is an equation that links the reaction rate with the concentrations or pressures of the reactants and constant parameters. For many reactions the rate is given by a power law such as r = k x y where and express the concentration of the species A and B; the exponents x and y are the partial orders of reaction for A and B and the overall reaction order is the sum of the exponents. These are positive integers, but they may be zero, fractional, or negative; the constant k is the reaction rate constant or rate coefficient of the reaction. Its value may depend on conditions such as temperature, ionic strength, surface area of an adsorbent, or light irradiation. Elementary reactions and reaction steps have reaction orders equal to the stoichiometric coefficients for each reactant; the overall reaction order, i.e. the sum of stoichiometric coefficients of reactants, is always equal to the molecularity of the elementary reaction. However complex reactions may or may not have reaction orders equal to their stoichiometric coefficients.
This implies that the order and the rate equation of a given reaction cannot be reliably deduced from the stoichiometry and must be determined experimentally, since an unknown reaction mechanism could be either elementary or complex. When the experimental rate equation has been determined, it is of use for deduction of the reaction mechanism; the rate equation of a reaction with an assumed multi-step mechanism can be derived theoretically using quasi-steady state assumptions from the underlying elementary reactions, compared with the experimental rate equation as a test of the assumed mechanism. The equation may involve a fractional order, may depend on the concentration of an intermediate species. A reaction can have an undefined reaction order with respect to a reactant if the rate is not proportional to some power of the concentration of that reactant. Consider a typical chemical reaction in which two reactants A and B combine to form a product C: A + 2 B → 3 C; this can be written 0 = − A − 2 B + 3 C.
The prefactors -1, -2 and 3 are known as stoichiometric coefficients. One molecule of A combines with two of B to form 3 of C, so if we use the symbol for the number of moles of chemical X, − d d t = − 1 2 d d t = 1 3 d d t. If the reaction takes place in a closed system at constant temperature and volume, without a build-up of reaction intermediates, the reaction rate r is defined as r = 1 ν i d d t, where νi is the stoichiometric coefficient for chemical Xi; the reaction rate has some functional dependence on the concentrations of the reactants, r = f, this dependence is known as the rate equation or rate law. This law cannot be deduced from the chemical equation and must be determined by experiment. A common form for the rate equation is a power law: r = k x y … The constant k is called the rate constant; the exponents, which can be fractional, are called partial orders of reaction and their sum is the overall order of reaction. In a dilute solution, an elementary reaction is empirically found to obey the law of mass action.
This predicts that the rate depends only on the concentrations of the reactants, raised to the powers of their stoichiometric coefficients. The natural logarithm of the p
Thiosulfate is an oxyanion of sulfur. The prefix thio- indicates that the thiosulfate ion is a sulfate ion with one oxygen replaced by sulfur. Thiosulfate has a tetrahedral molecular shape with C3v symmetry. Thiosulfate occurs and is produced by certain biochemical processes, it dechlorinates water and is notable for its use to halt bleaching in the paper-making industry. Thiosulfate is useful in smelting silver ore, in producing leather goods, to set dyes in textiles. Sodium thiosulfate called hypo, was used in photography to fix black and white negatives and prints after the developing stage; some bacteria can metabolise thiosulfates. Thiosulfate is produced by the reaction of sulfite ion with elemental sulfur, by incomplete oxidation of sulfides, sodium thiosulfate can be formed by disproportionation of Sulfur dissolving in sodium hydroxide. Thiosulfates are stable only in neutral or alkaline solutions, but not in acidic solutions, due to disproportionation to sulfite and sulfur, the sulfite being dehydrated to sulfur dioxide: S2O2−3 + 2 H+ → SO2 + S + H2O This reaction may be used to generate an aqueous suspension of sulfur and demonstrate the Rayleigh scattering of light in physics.
If white light is shone from below, blue light is seen from sideways and orange from above, due to the same mechanisms that color the sky at mid-day and dusk.. Thiosulfates react with halogens differently, which can be attributed to the decrease of oxidizing power down the halogen group: 2 S2O2−3 + I2 → S4O2−6 + 2 I− S2O2−3 + 4 Br2 + 5 H2O → 2 SO2−4 + 8 Br− + 10 H+ S2O2−3 + 4 Cl2 + 5 H2O → 2 SO2−4 + 8 Cl− + 10 H+ In acidic conditions, thiosulfate causes rapid corrosion of metals. Addition of molybdenum to stainless steel is needed to improve its resistance to pitting. In alkaline aqueous conditions and medium temperature, carbon steel and stainless steel are not attacked at high concentration of base, thiosulfate and in presence of fluoride ion; the natural occurrence of the thiosulfate group is restricted to a rare mineral sidpietersite, Pb4O22, as the presence of this anion in the mineral bazhenovite was disputed. Thiosulfate extensively forms complexes with transition metals hence a common use is dissolving silver halides in film photography developing.
Thiosulfate is used to extract or leach gold and silver from their ores as a less toxic alternative to cyanide. Thiosulfate is an acceptable common name; the external sulfur has an oxidation state of –2 while the central sulfur atom has an oxidation number of +6. The enzyme rhodanase catalyzes the detoxification of cyanide by thiosulfate: CN− + S2O2−3 → SCN− + SO2−3. Sodium thiosulfate has been considered as an empirical treatment for cyanide poisoning, along with hydroxocobalamin, it is most effective in a pre-hospital setting, since immediate administration by emergency personnel is necessary to reverse rapid intracellular hypoxia caused by the inhibition of cellular respiration, at complex IV. It activates TST in mitochondria. TST is associated with protection against obesity and type II diabetes. Tetrathionate Thiosulfuric acid Thiosulfate ion General Chemistry Online, Frostburg State University
Lewis acids and bases
A Lewis acid is a chemical species that contains an empty orbital, capable of accepting an electron pair from a Lewis base to form a Lewis adduct. A Lewis base is any species that has a filled orbital containing an electron pair, not involved in bonding but may form a dative bond with a Lewis acid to form a Lewis adduct. For example, NH3 is a Lewis base. Trimethylborane is a Lewis acid. In a Lewis adduct, the Lewis acid and base share an electron pair furnished by the Lewis base, forming a dative bond. In the context of a specific chemical reaction between NH3 and Me3B, the lone pair from NH3 will form a dative bond with the empty orbital of Me3B to form an adduct NH3•BMe3; the terminology refers to the contributions of Gilbert N. Lewis; the terms nucleophile and electrophile are more or less interchangeable with Lewis base and Lewis acid, respectively. However, these terms their abstract noun forms nucleophilicity and electrophilicity, emphasize the kinetic aspect of reactivity, while the Lewis basicity and Lewis acidity emphasize the thermodynamic aspect of Lewis adduct formation.
In many cases, the interaction between the Lewis base and Lewis acid in a complex is indicated by an arrow indicating the Lewis base donating electrons toward the Lewis acid using the notation of a dative bond—for example, Me3B←NH3. Some sources indicate the Lewis base with a pair of dots, which allows consistent representation of the transition from the base itself to the complex with the acid: Me3B +:NH3 → Me3B:NH3A center dot may be used to represent a Lewis adduct, such as Me3B•NH3. Another example is boron trifluoride diethyl etherate, BF3•Et2O. Although there have been attempts to use computational and experimental energetic criteria to distinguish dative bonding from non-dative covalent bonds, for the most part, the distinction makes note of the source of the electron pair, dative bonds, once formed, behave as other covalent bonds do, though they have considerable polar character. Moreover, in some cases, the use of the dative bond arrow is just a notational convenience for avoiding the drawing of formal charges.
In general, the donor–acceptor bond is viewed as somewhere along a continuum between idealized covalent bonding and ionic bonding. Classically, the term "Lewis acid" is restricted to trigonal planar species with an empty p orbital, such as BR3 where R can be an organic substituent or a halide. For the purposes of discussion complex compounds such as Et3Al2Cl3 and AlCl3 are treated as trigonal planar Lewis acids. Metal ions such as Na+, Mg2+, Ce3+, which are invariably complexed with additional ligands, are sources of coordinatively unsaturated derivatives that form Lewis adducts upon reaction with a Lewis base. Other reactions might be referred to as "acid-catalyzed" reactions; some compounds, such as H2O, are both Lewis acids and Lewis bases, because they can either accept a pair of electrons or donate a pair of electrons, depending upon the reaction. Lewis acids are diverse. Simplest are those, but more common are those. Examples of Lewis acids based on the general definition of electron pair acceptor include: the proton and acidic compounds onium ions, such as NH4+ and H3O+ high oxidation state transition metal cations, e.g. Fe3+.
Again, the description of a Lewis acid is used loosely. For example, in solution, bare protons do not exist; some of the most studied examples of such Lewis acids are the boron trihalides and organoboranes, but other compounds exhibit this behavior: BF3 + F− → BF4−In this adduct, all four fluoride centres are equivalent. BF3 + OMe2 → BF3OMe2Both BF4− and BF3OMe2 are Lewis base adducts of boron trifluoride. In many cases, the adducts violate the octet rule, such as the triiodide anion: I2 + I− → I3−The variability of the colors of iodine solutions reflects the variable abilities of the solvent to form adducts with the Lewis acid I2. In some cases, the Lewis acid is capable of binding two Lewis base, a famous example being the formation of hexafluorosilicate: SiF4 + 2 F− → SiF62− Most compounds considered to be Lewis acids require an activation step prior to formation of the adduct with the Lewis base. Well known cases are the aluminium trihalides, which are viewed as Lewis acids. Aluminium trihalides, unlike the boron trihalides, do not exist in the form AlX3, but as aggregates and polymers that must be degraded by the Lewis base.
A simpler case is the formation of adducts of borane. Monomeric BH3 does not exist appreciably, so the adducts of borane are generated by degradation of diborane: B2H6 + 2 H− → 2 BH4−In this case, an intermediate B2H7− can be isolated. Many metal complexes serve as Lewis acids, but only after dissociating a more weakly bound Lewis base water. 2+ + 6 NH3 → 2+ + 6 H2O The proton is one of the strongest but is one of the most complicated Lewis acids. It is convention to ignore the fact that a proton is solvated (bound to solvent
A chemical reaction is a process that leads to the chemical transformation of one set of chemical substances to another. Classically, chemical reactions encompass changes that only involve the positions of electrons in the forming and breaking of chemical bonds between atoms, with no change to the nuclei, can be described by a chemical equation. Nuclear chemistry is a sub-discipline of chemistry that involves the chemical reactions of unstable and radioactive elements where both electronic and nuclear changes can occur; the substance involved in a chemical reaction are called reactants or reagents. Chemical reactions are characterized by a chemical change, they yield one or more products, which have properties different from the reactants. Reactions consist of a sequence of individual sub-steps, the so-called elementary reactions, the information on the precise course of action is part of the reaction mechanism. Chemical reactions are described with chemical equations, which symbolically present the starting materials, end products, sometimes intermediate products and reaction conditions.
Chemical reactions happen at a characteristic reaction rate at a given temperature and chemical concentration. Reaction rates increase with increasing temperature because there is more thermal energy available to reach the activation energy necessary for breaking bonds between atoms. Reactions may proceed in the forward or reverse direction until they go to completion or reach equilibrium. Reactions that proceed in the forward direction to approach equilibrium are described as spontaneous, requiring no input of free energy to go forward. Non-spontaneous reactions require input of free energy to go forward. Different chemical reactions are used in combinations during chemical synthesis in order to obtain a desired product. In biochemistry, a consecutive series of chemical reactions form metabolic pathways; these reactions are catalyzed by protein enzymes. Enzymes increase the rates of biochemical reactions, so that metabolic syntheses and decompositions impossible under ordinary conditions can occur at the temperatures and concentrations present within a cell.
The general concept of a chemical reaction has been extended to reactions between entities smaller than atoms, including nuclear reactions, radioactive decays, reactions between elementary particles, as described by quantum field theory. Chemical reactions such as combustion in fire and the reduction of ores to metals were known since antiquity. Initial theories of transformation of materials were developed by Greek philosophers, such as the Four-Element Theory of Empedocles stating that any substance is composed of the four basic elements – fire, water and earth. In the Middle Ages, chemical transformations were studied by Alchemists, they attempted, in particular, to convert lead into gold, for which purpose they used reactions of lead and lead-copper alloys with sulfur. The production of chemical substances that do not occur in nature has long been tried, such as the synthesis of sulfuric and nitric acids attributed to the controversial alchemist Jābir ibn Hayyān; the process involved heating of sulfate and nitrate minerals such as copper sulfate and saltpeter.
In the 17th century, Johann Rudolph Glauber produced hydrochloric acid and sodium sulfate by reacting sulfuric acid and sodium chloride. With the development of the lead chamber process in 1746 and the Leblanc process, allowing large-scale production of sulfuric acid and sodium carbonate chemical reactions became implemented into the industry. Further optimization of sulfuric acid technology resulted in the contact process in the 1880s, the Haber process was developed in 1909–1910 for ammonia synthesis. From the 16th century, researchers including Jan Baptist van Helmont, Robert Boyle, Isaac Newton tried to establish theories of the experimentally observed chemical transformations; the phlogiston theory was proposed in 1667 by Johann Joachim Becher. It postulated the existence of a fire-like element called "phlogiston", contained within combustible bodies and released during combustion; this proved to be false in 1785 by Antoine Lavoisier who found the correct explanation of the combustion as reaction with oxygen from the air.
Joseph Louis Gay-Lussac recognized in 1808 that gases always react in a certain relationship with each other. Based on this idea and the atomic theory of John Dalton, Joseph Proust had developed the law of definite proportions, which resulted in the concepts of stoichiometry and chemical equations. Regarding the organic chemistry, it was long believed that compounds obtained from living organisms were too complex to be obtained synthetically. According to the concept of vitalism, organic matter was endowed with a "vital force" and distinguished from inorganic materials; this separation was ended however by the synthesis of urea from inorganic precursors by Friedrich Wöhler in 1828. Other chemists who brought major contributions to organic chemistry include Alexander William Williamson with his synthesis of ethers and Christopher Kelk Ingold, among many discoveries, established the mechanisms of substitution reactions. Chemical equations are used to graphically illustrate chemical reactions, they consist of chemical or structural formulas of the reactants on the left and those of the products on the right.
They are separated by an arrow which indicates the type of the reaction.
Hydrogen is a chemical element with symbol H and atomic number 1. With a standard atomic weight of 1.008, hydrogen is the lightest element in the periodic table. Hydrogen is the most abundant chemical substance in the Universe, constituting 75% of all baryonic mass. Non-remnant stars are composed of hydrogen in the plasma state; the most common isotope of hydrogen, termed protium, has no neutrons. The universal emergence of atomic hydrogen first occurred during the recombination epoch. At standard temperature and pressure, hydrogen is a colorless, tasteless, non-toxic, nonmetallic combustible diatomic gas with the molecular formula H2. Since hydrogen forms covalent compounds with most nonmetallic elements, most of the hydrogen on Earth exists in molecular forms such as water or organic compounds. Hydrogen plays a important role in acid–base reactions because most acid-base reactions involve the exchange of protons between soluble molecules. In ionic compounds, hydrogen can take the form of a negative charge when it is known as a hydride, or as a positively charged species denoted by the symbol H+.
The hydrogen cation is written as though composed of a bare proton, but in reality, hydrogen cations in ionic compounds are always more complex. As the only neutral atom for which the Schrödinger equation can be solved analytically, study of the energetics and bonding of the hydrogen atom has played a key role in the development of quantum mechanics. Hydrogen gas was first artificially produced in the early 16th century by the reaction of acids on metals. In 1766–81, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance, that it produces water when burned, the property for which it was named: in Greek, hydrogen means "water-former". Industrial production is from steam reforming natural gas, less from more energy-intensive methods such as the electrolysis of water. Most hydrogen is used near the site of its production, the two largest uses being fossil fuel processing and ammonia production for the fertilizer market. Hydrogen is a concern in metallurgy as it can embrittle many metals, complicating the design of pipelines and storage tanks.
Hydrogen gas is flammable and will burn in air at a wide range of concentrations between 4% and 75% by volume. The enthalpy of combustion is −286 kJ/mol: 2 H2 + O2 → 2 H2O + 572 kJ Hydrogen gas forms explosive mixtures with air in concentrations from 4–74% and with chlorine at 5–95%; the explosive reactions may be triggered by heat, or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C. Pure hydrogen-oxygen flames emit ultraviolet light and with high oxygen mix are nearly invisible to the naked eye, as illustrated by the faint plume of the Space Shuttle Main Engine, compared to the visible plume of a Space Shuttle Solid Rocket Booster, which uses an ammonium perchlorate composite; the detection of a burning hydrogen leak may require a flame detector. Hydrogen flames in other conditions are blue; the destruction of the Hindenburg airship was a notorious example of hydrogen combustion and the cause is still debated. The visible orange flames in that incident were the result of a rich mixture of hydrogen to oxygen combined with carbon compounds from the airship skin.
H2 reacts with every oxidizing element. Hydrogen can react spontaneously and violently at room temperature with chlorine and fluorine to form the corresponding hydrogen halides, hydrogen chloride and hydrogen fluoride, which are potentially dangerous acids; the ground state energy level of the electron in a hydrogen atom is −13.6 eV, equivalent to an ultraviolet photon of 91 nm wavelength. The energy levels of hydrogen can be calculated accurately using the Bohr model of the atom, which conceptualizes the electron as "orbiting" the proton in analogy to the Earth's orbit of the Sun. However, the atomic electron and proton are held together by electromagnetic force, while planets and celestial objects are held by gravity; because of the discretization of angular momentum postulated in early quantum mechanics by Bohr, the electron in the Bohr model can only occupy certain allowed distances from the proton, therefore only certain allowed energies. A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the Schrödinger equation, Dirac equation or the Feynman path integral formulation to calculate the probability density of the electron around the proton.
The most complicated treatments allow for the small effects of special relativity and vacuum polarization. In the quantum mechanical treatment, the electron in a ground state hydrogen atom has no angular momentum at all—illustrating how the "planetary orbit" differs from electron motion. There exist two different spin isomers of hydrogen diatomic molecules that differ by the relative spin of their nuclei. In the orthohydrogen form, the spins of the two protons are parallel and form a triplet state with a molecular spin quantum number of 1. At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form known as the "normal form"; the equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but because the ortho form is an excited state and has a higher energy
Malachite green is an organic compound, used as a dyestuff and controversially as an antimicrobial in aquaculture. Malachite green is traditionally used as a dye for materials such as silk and paper. Despite its name the dye is not prepared from the mineral malachite, the name just comes from the similarity of color. Malachite green is classified in the dyestuff industry as a triarylmethane dye and using in pigment industry. Formally, malachite green refers to the chloride salt Cl, although the term malachite green is used loosely and just refers to the colored cation; the oxalate salt is marketed. The anions have no effect on the color; the intense green color of the cation results from a strong absorption band at 621 nm. Malachite green is prepared by the condensation of benzaldehyde and dimethylaniline to give leuco malachite green: C6H5CHO + 2 C6H5N2 → C6H5CH2 + H2OSecond, this colorless leuco compound, a relative of triphenylmethane, is oxidized to the cation, MG: C6H5CH2 + HCl + 1⁄2 O2 → Cl + H2OA typical oxidizing agent is manganese dioxide.
Hydrolysis of MG gives an alcohol: Cl + H2O → C6H5C2 + HClThis alcohol is important because it, not MG, traverses cell membranes. Once inside the cell, it is metabolized into LMG. Only the cation MG is colored, whereas the leuco and alcohol derivatives are not; this difference arises because only the cationic form has extended pi-delocalization, which allows the molecule to absorb visible light. The leuco form of malachite green was first prepared by Hermann Fischer in 1877 by condensing benzaldehyde and dimethylaniline in the molecular ratio 1:2 in the presence of sulfuric acid. Malachite green is traditionally used as a dye. Kilotonnes of MG and related triarylmethane dyes are produced annually for this purpose. MG is active against the oomycete Saprolegnia, which infects fish eggs in commercial aquaculture, MG has been used to treat Saprolegnia and is used as an antibacterial, it is a popular treatment against Ichthyophthirius multifiliis in freshwater aquaria. The principal metabolite, LMG, is found in fish treated with malachite green, this finding is the basis of controversy and government regulation.
See Antimicrobials in aquaculture. MG has been used to catch thieves and pilferers; the bait money, is sprinkled with the anhydrous powder. Anyone handling the contaminated money will find that on upon washing the hands, a green stain on the skin that lasts for several days will result. Numerous niche applications exploit the intense color of MG, it is used as a biological stain for microscopic analysis of cell tissue samples. In the Gimenez staining method, basic fuchsin stains bacteria red or magenta, malachite green is used as a blue-green counterstain. Malachite green is used in endospore staining, since it can directly stain endospores within bacterial cells. Malachite green can be used as a saturable absorber in dye lasers, or as a pH indicator between pH 0.2–1.8. However, this use is rare. Leuco-malachite green is used as a detection method for latent blood in forensic science. Hemoglobin catalyzes the reaction between LMG and hydrogen peroxide, converting the colorless LMG into malachite green.
Therefore, the appearance of a green color indicates the presence of blood. In 1992, Canadian authorities determined that eating fish contaminated with malachite green posed a significant health risk. Malachite green was classified a Class II Health Hazard. Due to its low manufacturing cost, malachite green is still used in certain countries with less restrictive laws for non aquaculture purposes. In 2005, analysts in Hong Kong found traces of malachite green in eels and fish imported from China and Taiwan. In 2006, the United States Food and Drug Administration detected malachite green in seafood imported from China, among others, where the substance is banned for use in aquaculture. In June 2007, the FDA blocked the importation of several varieties of seafood due to continued malachite green contamination; the substance has been banned in the United States since 1983 in food-related applications. It is banned in the UK, also. Animals metabolize malachite green to its leuco form. Being lipophillic, the metabolite is retained in catfish muscle longer than is the parent molecule.
The LD50 is 80 mg/kg. Rats fed malachite green experience "a dose-related increase in liver DNA adducts" along with lung adenomas. Leucomalachite green causes an "increase in the number and severity of changes"; as leucomalachite green is the primary metabolite of malachite green and is retained in fish muscle much longer, most human dietary intake of malachite green from eating fish would be in the leuco form. During the experiment, rats were fed up to 543 ppm of leucomalachite green, an extreme amount compared to the average 5 ppb discovered in fish. After a period of two years, an increase in lung adenomas in male rats was discovered but no incidences of liver tumors. Therefore, it could be concluded that malachite green caused carcinogenic symptoms, but a direct link between malachite green and liver tumor was not established. Cho, Bongsup P.. "Synthesis and Characterization of N-Demethylated Metabolites of Malachite Green and Leucomalachite Green". Chem. Res. Toxicol. 16: 285–294. Doi:10.1021/tx0256679.
PMID 12641428. Plakas, S. M.. "Uptake, tissue distribution, metabolism of malachite green in the channel catfish". C