Robert Robinson (chemist)
Sir Robert Robinson was a British organic chemist and Nobel laureate recognised in 1947 for his research on plant dyestuffs and alkaloids. In 1947, he received the Medal of Freedom with Silver Palm, he was born at Rufford House Farm, near Chesterfield, Derbyshire the son of James Bradbury Robinson, a maker of surgical dressings, his wife, Jane Davenport. Robinson went to school at the private Fulneck School, he studied Chemistry at the University of Manchester, graduating BSc in 1905. In 1907 he was awarded an 1851 Research Fellowship from the Royal Commission for the Exhibition of 1851 to continue his research at the University of Manchester, he was appointed as the first Professor of Pure and Applied Organic Chemistry in the School of Chemistry at the University of Sydney in 1912. He was at St Andrews University and was offered the Chair of Organic Chemistry at Manchester University. In 1928 he moved from there to be a professor at University College London where he stayed only two years, he was the Waynflete Professor of Chemistry at Oxford University from 1930 and a Fellow of Magdalen College, Oxford.
Robinson Close in the Science Area at Oxford is named after him, as is the Robert Robinson Laboratory at the University of Liverpool, the Sir Robert Robinson Laboratory of Organic Chemistry at the University of Manchester and the Robinson and Cornforth Laboratories at the University of Sydney. Robinson was a strong amateur chess player, he represented Oxford University in a friendly match with a team from Bletchley Park in December 1944. He was president of the British Chess Federation from 1950–53, with Raymond Edwards he co-authored the book The Art and Science of Chess, his synthesis of tropinone, a precursor of cocaine, in 1917 was not only a big step in alkaloid chemistry but showed that tandem reactions in a one-pot synthesis are capable of forming bicyclic molecules. He invented the symbol for benzene having a circle in the middle whilst working at St Andrews University in 1923, he is known for inventing the use of the curly arrow to represent electron movement, he is known for discovering the molecular structures of morphine and penicillin.
Robinson annulation has had application in the total synthesis of steroids. In 1957 Robinson founded the journal Tetrahedron with fifty other editors for Pergamon Press; the Structural Relationship of Natural Products He married twice. In 1912 he married Gertrude Maud Walsh. Following her death in 1954, in 1957 he married Mrs Stern Sylvia Hillstrom. Nobel Lecture Some Polycyclic Natural Products from Nobelprize.org website Biography Biography from Nobelprize.org website ABC Online Forum
Triflic acid known as trifluoromethanesulfonic acid, TFMS, TFSA, HOTf or TfOH, is a sulfonic acid with the chemical formula CF3SO3H. It is one of the strongest acids. Triflic acid is used in research as a catalyst for esterification, it is a hygroscopic, colorless viscous liquid and is soluble in polar solvents. Trifluoromethanesulfonic acid is produced industrially by electrochemical fluorination of methanesulfonic acid: CH3SO3H + 4 HF → CF3SO2F + H2O + 3 H2The resulting CF3SO2F is hydrolyzed, the resulting triflate salt is preprotonated. Alternatively, trifluoromethanesulfonic acid arises by oxidation of trifluoromethylsulfenyl chloride: CF3SCl + 2 Cl2 + 3 H2O → CF3SO3H + 5 HClTriflic acid is purified by distillation from triflic anhydride. Trifluoromethanesulfonic acid was first synthesized in 1954 by Robert Haszeldine and Kidd by the following reaction: In the laboratory, triflic acid is useful in protonations because the conjugate base of triflic acid is nonnucleophilic, it is used as an acidic titrant in nonaqueous acid-base titration because it behaves as a strong acid in many solvents where common mineral acids are only moderately strong.
With a Ka = 5×1014, pKa −14.7±2.0, triflic acid qualifies as a superacid. It owes many of its useful properties to its great chemical stability. Both the acid and its conjugate base CF3SO−3, known as triflate, resist oxidation/reduction reactions, whereas many strong acids are oxidizing, e.g. perchloric or nitric acid. Further recommending its use, triflic acid does not sulfonate substrates, which can be a problem with sulfuric acid, fluorosulfuric acid, chlorosulfonic acid. Below is a prototypical sulfonation, which HOTf does not undergo: C6H6 + H2SO4 → C6H5 + H2OTriflic acid fumes in moist air and forms a stable solid monohydrate, CF3SO3H·H2O, melting point 34 °C; the triflate ligand is labile. Trifluoromethanesulfonic acid exothermically reacts with metal carbonates and oxides. Illustrative is the synthesis of Cu2. CuCO3 + 2 CF3SO3H → Cu2 + H2O + CO2Chloride ligands can be converted to the corresponding triflates: 3 CF3SO3H + Cl2 → 2 + 3 HClThis conversion is conducted in neat HOTf at 100 °C, followed by precipitation of the salt upon the addition of ether.
Triflic acid reacts with acyl halides to give mixed triflate anhydrides, which are strong acylating agents, e.g. in Friedel–Crafts reactions. CH3CCl + CF3SO3H → CH3COSO2CF3 + HCl CH3COSO2CF3 + C6H6 → CH3CC6H5 + CF3SO3HTriflic acid catalyzes the reaction of aromatic compounds with sulfonyl chlorides also through the intermediacy of a mixed anhydride of the sulfonic acid. Triflic acid promotes other Friedel–Crafts-like reactions including the cracking of alkanes and alkylation of alkenes, which are important to the petroleum industry; these triflic acid derivative catalysts are effective in isomerizing straight chain or branched hydrocarbons that can increase the octane rating of a particular petroleum-based fuel. Triflic acid reacts exothermically with alcohols to produce olefins. Dehydration gives the acid anhydride, trifluoromethanesulfonic anhydride, 2O. Triflic acid is one of the strongest acids. Contact with skin causes severe burns with delayed tissue destruction. On inhalation it causes fatal spasms and edema.
Addition of triflic acid to polar solvents can be dangerously exothermic
Aluminium chloride known as aluminium trichloride, is the main compound of aluminium and chlorine. It is white, but samples are contaminated with iron chloride, giving it a yellow color; the solid has boiling point. It is produced and consumed in the production of aluminium metal, but large amounts are used in other areas of the chemical industry; the compound is cited as a Lewis acid. It is an example of an inorganic compound that reversibly changes from a polymer to a monomer at mild temperature. AlCl3 adopts three different structures, depending on the state. Solid AlCl3 is a sheet-like layered cubic close packed layers. In this framework, the Al centres exhibit octahedral coordination geometry. In the melt, aluminium trichloride exists with tetracoordinate aluminium; this change in structure is related to the lower density of the liquid phase versus solid aluminium trichloride. Al2Cl6 dimers are found in the vapour phase. At higher temperatures, the Al2Cl6 dimers dissociate into trigonal planar AlCl3, structurally analogous to BF3.
The melt conducts electricity poorly, unlike more-ionic halides such as sodium chloride. The hexahydrate consists of octahedral 3+ centers and chloride counterions. Hydrogen bonds anions; the hydrated form of aluminium chloride has an octahedral molecular geometry, with the central aluminum ion surrounded by six water ligand molecules. This means that the hydrated form cannot act as a Lewis acid since it cannot accept electron pairs, thus this cannot be used as a catalyst in Friedel-Crafts alkylation of aromatic compounds. Anhydrous aluminium chloride is a most powerful Lewis acid, capable of forming Lewis acid-base adducts with weak Lewis bases such as benzophenone and mesitylene, it forms tetrachloroaluminate in the presence of chloride ions. Aluminium chloride reacts with calcium and magnesium hydrides in tetrahydrofuran forming tetrahydroaluminates. Aluminium chloride is hygroscopic, having a pronounced affinity for water, it fumes in moist air and hisses when mixed with liquid water as the Cl− ions are displaced with H2O molecules in the lattice to form the hexahydrate Cl3.
The anhydrous phase cannot be regained on heating as HCl is lost leaving aluminium hydroxide or alumina: Al6Cl3 → Al3 + 3 HCl + 3 H2OOn strong heating, aluminium oxide is formed from the aluminium hydroxide via: 2 Al3 → Al2O3 + 3 H2OAqueous solutions of AlCl3 are ionic and thus conduct electricity well. Such solutions are found to be indicative of partial hydrolysis of the Al3 + ion; the reactions can be described as: 3+ ⇌ 2+ + H+Aqueous solutions behave to other aluminium salts containing hydrated Al3+ ions, giving a gelatinous precipitate of aluminium hydroxide upon reaction with dilute sodium hydroxide: AlCl3 + 3 NaOH → Al3 + 3 NaCl Aluminium chloride is manufactured on a large scale by the exothermic reaction of aluminium metal with chlorine or hydrogen chloride at temperatures between 650 to 750 °C. 2 Al + 3 Cl2 → 2 AlCl3 2 Al + 6 HCl → 2 AlCl3 + 3 H2Aluminum chloride may be formed via a single displacement reaction between copper chloride and aluminum metal. 2 Al + 3 CuCl2 → 2 AlCl3 + 3 CuIn the US in 1993 21,000 tons were produced, not counting the amounts consumed in the production of aluminium.
Hydrated aluminium trichloride is prepared by dissolving aluminium oxides in hydrochloric acid. Metallic aluminum readily dissolves in hydrochloric acid ─ releasing hydrogen gas and generating considerable heat. Heating this solid does not produce anhydrous aluminium trichloride, the hexahydrate decomposes to aluminium hydroxide when heated: Al6Cl3 → Al3 + 3 HCl + 3 H2OAluminium forms a lower chloride, aluminium chloride, but this is unstable and only known in the vapour phase. AlCl3 is the most used Lewis acid and one of the most powerful, it finds application in the chemical industry as a catalyst for Friedel–Crafts reactions, both acylations and alkylations. Important products are ethylbenzene, it finds use in polymerization and isomerization reactions of hydrocarbons. The Friedel–Crafts reaction is the major use for aluminium chloride, for example in the preparation of anthraquinone from benzene and phosgene. In the general Friedel–Crafts reaction, an acyl chloride or alkyl halide reacts with an aromatic system as shown: The alkylation reaction is more used than the acylation reaction, although its practice is more technically demanding because the reaction is more sluggish.
For both reactions, the aluminium chloride, as well as other materials and the equipment, should be dry, although a trace of moisture is necessary for the reaction to proceed. A general problem with the Friedel–Crafts reaction is that the aluminium chloride catalyst sometimes is required in full stoichiometric quantities, because it complexes with the products; this complication sometimes generates a large amount of corrosive waste. For these and similar reasons, more recyclable or environmentally benign catalysts have been sought. Thus, the use of aluminium chloride in some applications is being displaced by zeolites. Aluminium chloride can be used to introduce aldehyde groups onto aromatic rings, for example via the Gattermann-Koch reaction which uses carbon monoxide, hydrogen chloride and a copper chloride co-catalyst. Aluminium chloride finds a wide variety of other applications in organic chemistry. For example, it can catalyse the "ene reaction", such as the addition of 3-but
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
Siegmund Gabriel was a German chemist. Gabriel went to school in Berlin. After studying a few semesters at the University of Berlin, Gabriel studied at the University of Heidelberg and received his Ph. D. for work with Robert Wilhelm Bunsen 1874. He was professor at the University of Berlin till 1921, he discovered the Gabriel Synthesis with his partner James Dornbush in 1887 and reported what became the Robinson–Gabriel synthesis in 1910. James Colman. "Siegmund Gabriel. A. Persönlicher Teil". Berichte der deutschen chemischen Gesellschaft. 59: A7–A26. Doi:10.1002/cber.19260590242
Phosphoryl chloride is a colourless liquid with the formula POCl3. It hydrolyses in moist air releasing phosphoric acid and fumes of hydrogen chloride, it is manufactured industrially on a large scale from phosphorus trichloride and oxygen or phosphorus pentoxide. It is used to make phosphate esters such as tricresyl phosphate. Like phosphate, phosphoryl chloride is tetrahedral in shape, it features three P−Cl bonds and one strong P=O double bond, with an estimated bond dissociation energy of 533.5 kJ/mol. On the basis of bond length and electronegativity, the Schomaker-Stevenson rule suggests that the double bond form is dominant, in contrast with the case of POF3; the P=O bond involves the donation of the lone pair electrons on oxygen p-orbitals to the antibonding combinations associated with phosphorus-chlorine bonds, thus constituting π bonding. With a freezing point of 1 °C and boiling point of 106 °C, the liquid range of POCl3 is rather similar to water. Like water, POCl3 autoionizes, owing to the reversible formation of POCl2+,Cl−.
POCl3 reacts with water to give hydrogen chloride and phosphoric acid: O=PCl3 + 3 H2O → O=P3 + 3 HClIntermediates in the conversion have been isolated, including pyrophosphoryl chloride, P2O3Cl4. Upon treatment with excess alcohols and phenols, POCl3 gives phosphate esters: O=PCl3 + 3 ROH → O=P3 + 3 HClSuch reactions are performed in the presence of an HCl acceptor such as pyridine or an amine. POCl3 can act as a Lewis base, forming adducts with a variety of Lewis acids such as titanium tetrachloride: Cl3PO + TiCl4 → Cl3POTiCl4The aluminium chloride adduct is quite stable, so POCl3 can be used to remove AlCl3 from reaction mixtures, for example at the end of a Friedel-Crafts reaction. POCl3 reacts with hydrogen bromide in the presence of Lewis-acidic catalysts to produce POBr3. Phosphoryl chloride can be prepared by many methods. Phosphoryl chloride was first reported in 1847 by the French chemist Adolphe Wurtz by reacting phosphorus pentachloride with water; the commercial method involves oxidation of phosphorus trichloride with oxygen: 2 PCl3 + O2 → 2 POCl3A related reaction include the oxidation of phosphorus trichloride with potassium chlorate: 3 PCl3 + KClO3 → 3 POCl3 + KCl The reaction of phosphorus pentachloride with phosphorus pentoxide.
6 PCl5 + P4O10 → 10 POCl3The reaction can be simplified by chlorinating a mixture of PCl3 and P4O10, generating the PCl5 in situ. The reaction of phosphorus pentachloride with boric acid or oxalic acid: 3 PCl5 + 2 B3 → 3 POCl3 + B2O3 + 6 HCl PCl5 + 2 → POCl3 + CO + CO2 + 2 HCl Reduction of tricalcium phosphate with carbon in the presence of chlorine gas: Ca32 + 6 C + 6 Cl2 → 3 CaCl2 + 6 CO + 2 POCl3The reaction of phosphorus pentoxide with sodium chloride is reported: 2 P2O5 + 3 NaCl → 3 NaPO3 + POCl3. In one commercial application, phosphoryl chloride is used in the manufacture of phosphate esters. Triarylphosphates such as triphenyl phosphate and tricresyl phosphate are used as flame retardants and plasticisers for PVC. Trialkylphosphates such as tributyl phosphate are used as liquid–liquid extraction solvents in nuclear reprocessing and elsewhere. In the semiconductor industry, POCl3 is used as a safe liquid phosphorus source in diffusion processes; the phosphorus acts. In the laboratory, POCl3 is a reagent in dehydrations.
One example involves conversion of primary amides to nitriles: RCNH2 + POCl3 → RCN + "PO2Cl" + 2 HClIn a related reaction, certain aryl-substituted amides can be cyclised using the Bischler-Napieralski reaction. Such reactions are believed to proceed via an imidoyl chloride. In certain cases, the imidoyl chloride is the final product. For example and pyrimidones can be converted to chloro- derivatives such as 2-chloropyridines and 2-chloropyrimidines, which are intermediates in the pharmaceutical industry. In the Vilsmeier-Haack reaction, POCl3 reacts with amides to produce a "Vilsmeier reagent", a chloro-iminium salt, which subsequently reacts with electron-rich aromatic compounds to produce aromatic aldehydes upon aqueous work-up
Phosphorus pentoxide is a chemical compound with molecular formula P4O10. This white crystalline solid is the anhydride of phosphoric acid, it is dehydrating agent. Phosphorus polymorphs; the most familiar one, a metastable form, shown in the figure, comprises molecules of P4O10. Weak van der Waals forces hold these molecules together in a hexagonal lattice; the structure of the P4O10 cage is reminiscent of adamantane with Td symmetry point group. It is related to the corresponding anhydride of phosphorous acid, P4O6; the latter lacks terminal oxo groups. Its density is 2.30 g/cm3. It boils at 423 °C under atmospheric pressure; this form can be made by condensing the vapor of phosphorus pentoxide the result is an hygroscopic solid. The other polymorphs are polymeric, but in each case the phosphorus atoms are bound by a tetrahedron of oxygen atoms, one of which forms a terminal P=O bond involving the donation of the terminal oxygen p-orbital electrons to the antibonding phosphorus-oxygen single bonds.
The macromolecular form can be made by heating the compound in a sealed tube for several hours, maintaining the melt at a high temperature before cooling the melt to the solid. The metastable orthorhombic, "O"-form, adopts a layered structure consisting of interconnected P6O6 rings, not unlike the structure adopted by certain polysilicates; the stable form is a higher density phase orthorhombic, the so-called O' form. It consists of a 3-dimensional framework, density 3.5 g/cm3. The remaining polymorph is a glass or amorphous form. P4O10 is prepared by burning tetraphosphorus with sufficient supply of oxygen: P4 + 5 O2 → P4O10For most of the 20th century, phosphorus pentoxide was used to provide a supply of concentrated pure phosphoric acid. In the thermal process, the phosphorus pentoxide obtained by burning white phosphorus was dissolved in dilute phosphoric acid to produce concentrated acid. Improvements in filter technology is leading to the "wet phosphoric acid process" taking over from the thermal process, obviating the need to produce white phosphorus as a starting material.
The dehydration of phosphoric acid to give phosphorus pentoxide is not possible as on heating metaphosphoric acid will boil without losing all its water. Phosphorus pentoxide is a potent dehydrating agent as indicated by the exothermic nature of its hydrolysis: P4O10 + 6 H2O → 4 H3PO4 However, its utility for drying is limited somewhat by its tendency to form a protective viscous coating that inhibits further dehydration by unspent material. A granular form of P4O10 is used in desiccators. Consistent with its strong desiccating power, P4O10 is used in organic synthesis for dehydration; the most important application is for the conversion of primary amides into nitriles: P4O10 + RCNH2 → P4O92 + RCNThe indicated coproduct P4O92 is an idealized formula for undefined products resulting from the hydration of P4O10. Alternatively, when combined with a carboxylic acid, the result is the corresponding anhydride: P4O10 + RCO2H → P4O92 + 2OThe "Onodera reagent", a solution of P4O10 in DMSO, is employed for the oxidation of alcohols.
This reaction is reminiscent of the Swern oxidation. The desiccating power of P4O10 is strong enough to convert many mineral acids to their anhydrides. Examples: HNO3 is converted to N2O5. Between the commercially important P4O6 and P4O10, phosphorus oxides are known with intermediate structures. Phosphorus pentoxide. Just like sulfur trioxide, it reacts vigorously with water and water-containing substances like wood or cotton, liberates much heat and may cause fire due to the exothermic nature of such reactions, it is corrosive to metal and is irritating – it may cause severe burns to the eye, mucous membrane, respiratory tract at concentrations as low as 1 mg/m3. In Anthony Burgess' The Wanting Seed, phosphorus pentoxide is a prized compound. In Detective Comics #825, Batman notices that phosphorus pentoxide was at the scene of a fire, indicating that the villain Dr. Phosphorus was involved. In Aldous Huxley's Point Counter Point, to his assistant Illidge, Lord Edward bemoans societal loss of phosphorous pentoxide.
In Aldous Huxley's Brave New World, Henry Foster tells Lenina about the recovery of phosphorus pentoxide. In The Tunnel, a victim was consumed by a fire started with phosphorus pentoxide. Eaton's reagent OSHA Spec sheet Definition Website of the Technische Universität Darmstadt and the CEEP about Phosphorus Recovery