The Ullmann reaction or Ullmann coupling is a coupling reaction between aryl halides and copper. The reaction is named after Fritz Ullmann. Ullmann-type reaction is the copper-promoted conversion of aryl halides to aryl ethers, aryl thioethers, aryl amines. Ullmann-type reactions are comparable to Buchwald-Hartwig reactions; the mechanism of the Ullmann reaction is extensively studied. Complications arise because the reactions are heterogeneous those starting with metallic copper. Radical intermediates are not observed by Electron spin resonance. Oxidative addition/reductive elimination sequence seem but in one-electron steps in some cases. Copper has been observed but is invoked in this area of catalysis. Initial organocopper intermediates are speculated to be the copper species ArCu and CuX or the copper species ArCuX. In the case of Ullmann-type reactions, the reaction involves copper copper amides; the copper reagent can be generated in situ from the aryl copper metal. Copper sources are effective.
A number of innovations have been developed with regards to copper reagents. These copper compounds subsequently react with the aryl halide in a net metathesis reaction: Ar-X + ROCu → Ar-OR + CuX Ar-X + RSCu → Ar-SR + CuX Ar-X + 2 RNHCu → Ar-NHR + CuXIn the case of C-N coupling, kinetic studies implicate oxidative addition reaction followed by reductive elimination. A typical example of classic Ullmann biaryl coupling is the conversion of ortho-chloronitrobenzene into 2,2'-dinitrobiphenyl with a copper - bronze alloy; the traditional version of the Ullmann reaction requires harsh reaction conditions, the reaction has a reputation for erratic yields. Because of these problems many improvements and alternative procedures have been introduced; the classical Ullmann reaction is limited to electron deficient aryl halides and requires harsh reaction conditions. Modern variants of the Ullman reaction employing palladium and nickel have widened the substrate scope of the reaction and rendered reaction conditions more mild.
Yields are still moderate, however. In organic synthesis this reaction is replaced by palladium coupling reactions such as the Heck reaction, the Hiyama coupling, the Sonogashira coupling. Biphenylenes had been obtained before with reasonable yields using 2, 2 diiodobiphenyl or 2, 2 diiodobiphenylonium ion as starting material. Closure of 5-membered ring are more favorable more facile, but larger rings have been made using this approach. Ullmann synthesis of biaryl compounds can be used to generate chiral products from chiral reactants. Nelson and collaborators worked on the synthesis of asymmetric biaryl compounds and obtained the thermodynamically controlled product; the diastereometric ratio of the products is enhanced with bulkier R groups in the auxiliary oxazoline group. Unsymmetrical Ullmann reactions are pursued but have been achieved when one of the two coupling components is in excess. Copper-thiophene-2-carboxylate, a copper reagent used in the Ullmann reaction Wurtz-Fittig reaction, a similar reaction useful for alkylbenzenes synthesis
The density, or more the volumetric mass density, of a substance is its mass per unit volume. The symbol most used for density is ρ, although the Latin letter D can be used. Mathematically, density is defined as mass divided by volume: ρ = m V where ρ is the density, m is the mass, V is the volume. In some cases, density is loosely defined as its weight per unit volume, although this is scientifically inaccurate – this quantity is more called specific weight. For a pure substance the density has the same numerical value as its mass concentration. Different materials have different densities, density may be relevant to buoyancy and packaging. Osmium and iridium are the densest known elements at standard conditions for temperature and pressure but certain chemical compounds may be denser. To simplify comparisons of density across different systems of units, it is sometimes replaced by the dimensionless quantity "relative density" or "specific gravity", i.e. the ratio of the density of the material to that of a standard material water.
Thus a relative density less than one means. The density of a material varies with pressure; this variation is small for solids and liquids but much greater for gases. Increasing the pressure on an object decreases the volume of the object and thus increases its density. Increasing the temperature of a substance decreases its density by increasing its volume. In most materials, heating the bottom of a fluid results in convection of the heat from the bottom to the top, due to the decrease in the density of the heated fluid; this causes it to rise relative to more dense unheated material. The reciprocal of the density of a substance is called its specific volume, a term sometimes used in thermodynamics. Density is an intensive property in that increasing the amount of a substance does not increase its density. In a well-known but apocryphal tale, Archimedes was given the task of determining whether King Hiero's goldsmith was embezzling gold during the manufacture of a golden wreath dedicated to the gods and replacing it with another, cheaper alloy.
Archimedes knew that the irregularly shaped wreath could be crushed into a cube whose volume could be calculated and compared with the mass. Baffled, Archimedes is said to have taken an immersion bath and observed from the rise of the water upon entering that he could calculate the volume of the gold wreath through the displacement of the water. Upon this discovery, he leapt from his bath and ran naked through the streets shouting, "Eureka! Eureka!". As a result, the term "eureka" entered common parlance and is used today to indicate a moment of enlightenment; the story first appeared in written form in Vitruvius' books of architecture, two centuries after it took place. Some scholars have doubted the accuracy of this tale, saying among other things that the method would have required precise measurements that would have been difficult to make at the time. From the equation for density, mass density has units of mass divided by volume; as there are many units of mass and volume covering many different magnitudes there are a large number of units for mass density in use.
The SI unit of kilogram per cubic metre and the cgs unit of gram per cubic centimetre are the most used units for density. One g/cm3 is equal to one thousand kg/m3. One cubic centimetre is equal to one millilitre. In industry, other larger or smaller units of mass and or volume are more practical and US customary units may be used. See below for a list of some of the most common units of density. A number of techniques as well as standards exist for the measurement of density of materials; such techniques include the use of a hydrometer, Hydrostatic balance, immersed body method, air comparison pycnometer, oscillating densitometer, as well as pour and tap. However, each individual method or technique measures different types of density, therefore it is necessary to have an understanding of the type of density being measured as well as the type of material in question; the density at all points of a homogeneous object equals its total mass divided by its total volume. The mass is measured with a scale or balance.
To determine the density of a liquid or a gas, a hydrometer, a dasymeter or a Coriolis flow meter may be used, respectively. Hydrostatic weighing uses the displacement of water due to a submerged object to determine the density of the object. If the body is not homogeneous its density varies between different regions of the object. In that case the density around any given location is determined by calculating the density of a small volume around that location. In the limit of an infinitesimal volume the density of an inhomogeneous object at a point becomes: ρ = d m / d V, where d V is an elementary volume at position r; the mass of the body t
Ethers are a class of organic compounds that contain an ether group—an oxygen atom connected to two alkyl or aryl groups. They have the general formula R -- O -- R ′, where R ′ represent the alkyl or aryl groups. Ethers can again be classified into two varieties: if the alkyl groups are the same on both sides of the oxygen atom it is a simple or symmetrical ether, whereas if they are different, the ethers are called mixed or unsymmetrical ethers. A typical example of the first group is the solvent and anesthetic diethyl ether referred to as "ether". Ethers are common in organic chemistry and more prevalent in biochemistry, as they are common linkages in carbohydrates and lignin. Ethers feature C–O–C linkage defined by a bond angle of about 110° and C–O distances of about 140 pm; the barrier to rotation about the C–O bonds is low. The bonding of oxygen in ethers and water is similar. In the language of valence bond theory, the hybridization at oxygen is sp3. Oxygen is more electronegative than carbon, thus the hydrogens alpha to ethers are more acidic than in simple hydrocarbons.
They are far less acidic than hydrogens alpha to carbonyl groups, however. Depending on the groups at R and R′, ethers are classified into two types:Simple ethers or symmetrical ethers. Mixed ethers or asymmetrical ethers. In the IUPAC nomenclature system, ethers are named using the general formula "alkoxyalkane", for example CH3–CH2–O–CH3 is methoxyethane. If the ether is part of a more-complex molecule, it is described as an alkoxy substituent, so –OCH3 would be considered a "methoxy-" group; the simpler alkyl radical is written in front, so CH3–O–CH2CH3 would be given as methoxyethane. IUPAC rules are not followed for simple ethers; the trivial names for simple ethers are a composite of the two substituents followed by "ether". For example, ethyl methyl ether, diphenylether; as for other organic compounds common ethers acquired names before rules for nomenclature were formalized. Diethyl ether is called "ether", but was once called sweet oil of vitriol. Methyl phenyl ether is anisole, because it was found in aniseed.
The aromatic ethers include furans. Acetals are another class of ethers with characteristic properties. Polyethers are compounds with more than one ether group; the crown ethers are examples of small polyethers. Some toxins produced by dinoflagellates such as brevetoxin and ciguatoxin are large and are known as cyclic or ladder polyethers. Polyether refers to polymers which contain the ether functional group in their main chain; the term glycol is reserved for low to medium range molar mass polymer when the nature of the end-group, a hydroxyl group, still matters. The term "oxide" or other terms are used for high molar mass polymer when end-groups no longer affect polymer properties; the phenyl ether polymers are a class of aromatic polyethers containing aromatic cycles in their main chain: Polyphenyl ether and Poly. Many classes of compounds with C–O–C linkages are not considered ethers: Esters, carboxylic acid anhydrides. Ether molecules cannot form hydrogen bonds with each other, resulting in low boiling points compared to those of the analogous alcohols.
The difference in the boiling points of the ethers and their isomeric alcohols becomes lower as the carbon chains become longer, as the van der Waals interactions of the extended carbon chain dominates over the presence of hydrogen bonding. Ethers are polar; the C–O–C bond angle in the functional group is about 110°, the C–O dipoles do not cancel out. Ethers are more polar than alkenes but not as polar as alcohols, esters, or amides of comparable structure; the presence of two lone pairs of electrons on the oxygen atoms makes hydrogen bonding with water molecules possible. Cyclic ethers such as tetrahydrofuran and 1,4-dioxane are miscible in water because of the more exposed oxygen atom for hydrogen bonding as compared to linear aliphatic ethers. Other properties are: The lower ethers are volatile and flammable. Lower ethers act as anaesthetics. Ethers are good organic solvents. Simple ethers are tasteless. Ethers are quite stable chemical compounds which do not react with bases, active metals, dilute acids, oxidising agents, reducing agents.
They are of low chemical reactivity, but they are more reactive than alkanes. Epoxides and acetals are unrepresentative classes of ethers and are discussed in separate articles. Important reactions are listed below. Although ethers resist hydrolysis, their polar bonds are cloven by mineral acids such as hydrobromic acid and hydroiodic acid. Hydrogen chloride cleaves ethers only slowly. Methyl ethers afford methyl halides: ROCH3 + HBr → CH3Br + ROHThese reactions proceed via onium intermediates, i.e. +Br−. Some ethers undergo rapid cleavage with boron tribromide to give the alkyl bromide. Depending on the substituents, some ethers can be cloven with a variety of reagents, e.g. strong base. When stored in the presence of air or oxygen, ethers tend to form explosive peroxides, such as diethyl ether peroxide; the reaction is accelerated by light, metal catalysts, aldehydes. In addition to avoiding storage conditions to form peroxides, it is recommended, when an ether is used as a solvent, not to distill it to dryness, as any peroxides that may have formed, being less volatil
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
Geranium is a genus of 422 species of flowering annual and perennial plants that are known as the cranesbills. They are found throughout the temperate regions of the world and the mountains of the tropics, but in the eastern part of the Mediterranean region; the long, palmately cleft leaves are broadly circular in form. The flowers have five petals and are coloured white, purple or blue with distinctive veining. Geraniums will grow in any soil as long. Propagation is by seed, or by division in autumn or spring. Geraniums are eaten by the larvae of some Lepidoptera species including brown-tail, ghost moth, mouse moth. At least several species of Geranium are gynodioecious; the species Geranium viscosissimum is considered to be protocarnivorous. The genus name is derived from the Greek γέρανος or γερανός ‘crane’; the English name ‘cranesbill’ derives from the appearance of the fruit capsule of some of the species. Species in the genus Geranium have a distinctive mechanism for seed dispersal; this consists of a beak-like column.
The fruit capsule consists of five cells, each containing one seed, joined to a column produced from the centre of the old flower. The common name ‘cranesbill’ comes from the shape of the unsprung column, which in some species is long and looks like the bill of a crane. However, many species in this genus do not have a long beak-like column. Confusingly, "geranium" is the common name of members of the genus Pelargonium, which are in the Geraniaceae family and are grown as horticultural bedding plants. Linnaeus included all the species in one genus, but they were separated into two genera by Charles L’Héritier in 1789. Other former members of the genus are now classified in Erodium, including the plants known as filarees in North America; the term "hardy geranium" is applied to horticultural Geraniums to distinguish them from the Pelargoniums, which are not winter-hardy in temperate horticulture. However, not all Geranium species are winter-hardy; the shape of the flowers offers one way of distinguishing between the two genera Geranium and Pelargonium.
Geranium flowers have five similar petals, are thus radially symmetrical, whereas Pelargonium flowers have two upper petals which are different from the three lower petals, so the flowers have a single plane of symmetry. A number of geranium species are cultivated for pharmaceutical products; some of the more grown species include: All the above species are perennials and winter-hardy plants, grown for their attractive flowers and foliage. They most have a mounding habit, with palmately lobed foliage; some species have spreading rhizomes. They are grown in part shade to full sun, in well-draining but moisture retentive soils, rich in humus. Other perennial species grown for their flowers and foliage include: G. argenteum, G. eriostemon, G. farreri, G. nodosum, G. procurrens, G. pylzowianum, G. renardii, G. traversii, G. tuberosum, G. versicolor, G. wallichianum and G. wlassovianum. Some of these are not winter-hardy in cold areas and are grown in specialized gardens like rock gardens. Geranium'Johnson's Blue' is a hybrid between G. himalayense, with G. pratense.
The following hybrid cultivars have gained the Royal Horticultural Society's Award of Garden Merit:- List of cranesbill species Pelargonium graveolens, from which Geranium essential oil is distilled ITIS list of Geranium species Geranium Taxonomic Information System Preparing Geraniums for Winter
In chemistry, bases are substances that, in aqueous solution, release hydroxide ions, are slippery to the touch, can taste bitter if an alkali, change the color of indicators, react with acids to form salts, promote certain chemical reactions, accept protons from any proton donor or contain or displaceable OH− ions. Examples of bases are the hydroxides of the alkaline earth metals; these particular substances produce hydroxide ions in aqueous solutions, are thus classified as Arrhenius bases. For a substance to be classified as an Arrhenius base, it must produce hydroxide ions in an aqueous solution. Arrhenius believed; this makes the Arrhenius model limited, as it cannot explain the basic properties of aqueous solutions of ammonia or its organic derivatives. There are bases that do not contain a hydroxide ion but react with water, resulting in an increase in the concentration of the hydroxide ion. An example of this is the reaction between water to produce ammonium and hydroxide. In this reaction ammonia is the base.
Ammonia and other bases similar to it have the ability to form a bond with a proton due to the unshared pair of electrons that they possess. In the more general Brønsted–Lowry acid–base theory, a base is a substance that can accept hydrogen cations —otherwise known as protons. In the Lewis model, a base is an electron pair donor. In water, by altering the autoionization equilibrium, bases yield solutions in which the hydrogen ion activity is lower than it is in pure water, i.e. the water has a pH higher than 7.0 at standard conditions. A soluble base is called an alkali if it releases OH − ions quantitatively. However, it is important to realize. Metal oxides and alkoxides are basic, conjugate bases of weak acids are weak bases. Bases can be thought of as the chemical opposite of acids. However, some strong acids are able to act as bases. Bases and acids are seen as opposites because the effect of an acid is to increase the hydronium concentration in water, whereas bases reduce this concentration.
A reaction between an acid and a base is called neutralization. In a neutralization reaction, an aqueous solution of a base reacts with an aqueous solution of an acid to produce a solution of water and salt in which the salt separates into its component ions. If the aqueous solution is saturated with a given salt solute, any additional such salt precipitates out of the solution; the notion of a base as a concept in chemistry was first introduced by the French chemist Guillaume François Rouelle in 1754. He noted that acids, which at that time were volatile liquids, turned into solid salts only when combined with specific substances. Rouelle considered that such a substance serves as a "base" for the salt, giving the salt a "concrete or solid form". General properties of bases include: Concentrated or strong bases are caustic on organic matter and react violently with acidic substances. Aqueous solutions or molten bases dissociate in ions and conduct electricity. Reactions with indicators: bases turn red litmus paper blue, phenolphthalein pink, keep bromothymol blue in its natural colour of blue, turn methyl orange yellow.
The pH of a basic solution at standard conditions is greater than seven. Bases are bitter in taste; the following reaction represents the general reaction between a base and water to produce a conjugate acid and a conjugate base: B + H2O ⇌ BH+ + OH−The equilibrium constant, Kb, for this reaction can be found using the following general equation: Kb = /In this equation, both the base and the strong base compete with one another for the proton. As a result, bases that react with water have small equilibrium constant values; the base is weaker. Bases react with acids to neutralize each other at a fast rate both in alcohol; when dissolved in water, the strong base sodium hydroxide ionizes into hydroxide and sodium ions: NaOH → Na+ + OH−and in water the acid hydrogen chloride forms hydronium and chloride ions: HCl + H2O → H3O+ + Cl−When the two solutions are mixed, the H3O+ and OH− ions combine to form water molecules: H3O+ + OH− → 2 H2OIf equal quantities of NaOH and HCl are dissolved, the base and the acid neutralize leaving only NaCl table salt, in solution.
Weak bases, such as baking soda or egg white, should be used to neutralize any acid spills. Neutralizing acid spills with strong bases, such as sodium hydroxide or potassium hydroxide, can cause a violent exothermic reaction, the base itself can cause just as much damage as the original acid spill. Bases are compounds that can neutralize an amount of acids. Both sodium carbonate and ammonia are bases, although neither of these substances contains OH− groups. Both compounds accept H+ when dissolved in protic solvents such as water: Na2CO3 + H2O → 2 Na+ + HCO3− + OH− NH3 + H2O → NH4+ + OH−From this, a pH, or acidity, can be calculated for aqueous solutions of bases. Bases directly act as electron-pair donors themselves: CO32− + H+ → HCO3− NH3 + H+ → NH4+A base is defined as a molecule that has the ability to accept an electron pair bond by entering another atom's valence shell through its possession of one electron pair. There are a limited number of elements that have atoms with the ability to provide a molecule with basic properties
The Jmol applet, among other abilities, offers an alternative to the Chime plug-in, no longer under active development. While Jmol has many features that Chime lacks, it does not claim to reproduce all Chime functions, most notably, the Sculpt mode. Chime requires plug-in installation and Internet Explorer 6.0 or Firefox 2.0 on Microsoft Windows, or Netscape Communicator 4.8 on Mac OS 9. Jmol operates on a wide variety of platforms. For example, Jmol is functional in Mozilla Firefox, Internet Explorer, Google Chrome, Safari. Chemistry Development Kit Comparison of software for molecular mechanics modeling Jmol extension for MediaWiki List of molecular graphics systems Molecular graphics Molecule editor Proteopedia PyMOL SAMSON Official website Wiki with listings of websites and moodles Willighagen, Egon. "Fast and Scriptable Molecular Graphics in Web Browsers without Java3D". Doi:10.1038/npre.2007.50.1