Diatomic molecules are molecules composed of only two atoms, of the same or different chemical elements. The prefix di- is of Greek origin, meaning "two". If a diatomic molecule consists of two atoms of the same element, such as hydrogen or oxygen it is said to be homonuclear. Otherwise, if a diatomic molecule consists of two different atoms, such as carbon monoxide or nitric oxide, the molecule is said to be heteronuclear; the only chemical elements that form stable homonuclear diatomic molecules at standard temperature and pressure are the gases hydrogen, oxygen and chlorine. The noble gases are gases at STP, but they are monatomic; the homonuclear diatomic gases and noble gases together are called "elemental gases" or "molecular gases", to distinguish them from other gases that are chemical compounds. At elevated temperatures, the halogens bromine and iodine form diatomic gases. All halogens have been observed as diatomic molecules, except for astatine, uncertain; the mnemonics BrINClHOF, pronounced "Brinklehof", HONClBrIF, pronounced "Honkelbrif", HOFBrINCl have been coined to aid recall of the list of diatomic elements.
Other elements form diatomic molecules when evaporated, but these diatomic species repolymerize when cooled. Heating elemental phosphorus gives diphosphorus, P2. Sulfur vapor is disulfur. Dilithium is known in the gas phase. Ditungsten and dimolybdenum form with sextuple bonds in the gas phase; the bond in a homonuclear diatomic molecule is non-polar. Dirubidium is diatomic. All other diatomic molecules are chemical compounds of two different elements. Many elements can combine to form heteronuclear diatomic molecules, depending on temperature and pressure; some examples include, gases carbon monoxide, nitric oxide, hydrogen chloride. Many 1:1 binary compounds are not considered diatomic because they are polymeric at room temperature, but they form diatomic molecules when evaporated, for example gaseous MgO, SiO, many others. Hundreds of diatomic molecules have been identified in the environment of the Earth, in the laboratory, in interstellar space. About 99% of the Earth's atmosphere is composed of two species of diatomic molecules: nitrogen and oxygen.
The natural abundance of hydrogen in the Earth's atmosphere is only of the order of parts per million, but H2 is the most abundant diatomic molecule in the universe. The interstellar medium is, dominated by hydrogen atoms. Diatomic elements played an important role in the elucidation of the concepts of element and molecule in the 19th century, because some of the most common elements, such as hydrogen and nitrogen, occur as diatomic molecules. John Dalton's original atomic hypothesis assumed that all elements were monatomic and that the atoms in compounds would have the simplest atomic ratios with respect to one another. For example, Dalton assumed water's formula to be HO, giving the atomic weight of oxygen as eight times that of hydrogen, instead of the modern value of about 16; as a consequence, confusion existed regarding atomic weights and molecular formulas for about half a century. As early as 1805, Gay-Lussac and von Humboldt showed that water is formed of two volumes of hydrogen and one volume of oxygen, by 1811 Amedeo Avogadro had arrived at the correct interpretation of water's composition, based on what is now called Avogadro's law and the assumption of diatomic elemental molecules.
However, these results were ignored until 1860 due to the belief that atoms of one element would have no chemical affinity toward atoms of the same element, partly due to apparent exceptions to Avogadro's law that were not explained until in terms of dissociating molecules. At the 1860 Karlsruhe Congress on atomic weights, Cannizzaro resurrected Avogadro's ideas and used them to produce a consistent table of atomic weights, which agree with modern values; these weights were an important prerequisite for the discovery of the periodic law by Dmitri Mendeleev and Lothar Meyer. Diatomic molecules are in their lowest or ground state, which conventionally is known as the X state; when a gas of diatomic molecules is bombarded by energetic electrons, some of the molecules may be excited to higher electronic states, as occurs, for example, in the natural aurora. Such excitation can occur when the gas absorbs light or other electromagnetic radiation; the excited states are unstable and relax back to the ground state.
Over various short time scales after the excitation, transitions occur from higher to lower electronic states and to the ground state, in each transition results a photon is emitted. This emission is known as fluorescence. Successively higher electronic states are conventionally named A, B, C, etc.. The excitation energy must be greater than or equal to the energy of the electronic state in order for the excitation to occur. In quantum theory, an electronic state of a diatomic molecule is represented by 2 S + 1 Λ ( v
Hydrogen fluoride is a chemical compound with the chemical formula HF. This colorless gas or liquid is the principal industrial source of fluorine as an aqueous solution called hydrofluoric acid, it is an important feedstock in the preparation of many important compounds including pharmaceuticals and polymers. HF is used in the petrochemical industry as a component of superacids. Hydrogen fluoride boils near room temperature, much higher than other hydrogen halides. Hydrogen fluoride is a dangerous gas, forming corrosive and penetrating hydrofluoric acid upon contact with moisture; the gas can cause blindness by rapid destruction of the corneas. French chemist Edmond Frémy is credited with discovering anhydrous hydrogen fluoride while trying to isolate fluorine. Although Carl Wilhelm Scheele prepared hydrofluoric acid in large quantities in 1771, this acid was known in the glass industry before then. Although a diatomic molecule, HF forms strong intermolecular hydrogen bonds. Solid HF consists of zigzag chains of HF molecules.
The HF molecules, with a short H–F bond of 95 pm, are linked to neighboring molecules by intermolecular H–F distances of 155 pm. Liquid HF consists of chains of HF molecules, but the chains are shorter, consisting on average of only five or six molecules. Hydrogen fluoride does not boil until 20 °C in contrast to the heavier hydrogen halides which boil between −85 °C and −35 °C; this hydrogen bonding between HF molecules gives rise to high viscosity in the liquid phase and lower than expected pressure in the gas phase. Hydrogen fluoride is miscible with water, whereas the other hydrogen halides have large solubility gaps with water. Hydrogen fluoride and water form several compounds in the solid state, most notably a 1:1 compound that does not melt until −40 °C, 44 °C above the melting point of pure HF. Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in dilute aqueous solution; this is in part a result of the strength of the hydrogen–fluorine bond, but of other factors such as the tendency of HF, H2O, F− anions to form clusters.
At high concentrations, HF molecules undergo homoassociation to form polyatomic ions and protons, thus increasing the acidity. This leads to protonation of strong acids like hydrochloric, sulfuric, or nitric when using concentrated hydrofluoric acid solutions. Although hydrofluoric acid is regarded as a weak acid, it is corrosive attacking glass when hydrated; the acidity of hydrofluoric acid solutions varies with concentration owing to hydrogen-bond interactions of the fluoride ion. Dilute solutions are weakly acidic with an acid ionization constant Ka = 6.6×10−4, in contrast to corresponding solutions of the other hydrogen halides, which are strong acids. Concentrated solutions of hydrogen fluoride are much more acidic than implied by this value, as shown by measurements of the Hammett acidity function H0; the H0 for 100% HF is estimated to be between −10.2 and −11, comparable to the value −12 for sulfuric acid. In thermodynamic terms, HF solutions are non-ideal, with the activity of HF increasing much more than its concentration.
The weak acidity in dilute solution is sometimes attributed to the high H—F bond strength, which combines with the high dissolution enthalpy of HF to outweigh the more negative enthalpy of hydration of the fluoride ion. However, Paul Giguère and Sylvia Turrell have shown by infrared spectroscopy that the predominant solute species is the hydrogen-bonded ion pair, which suggests that the ionization can be described as a pair of successive equilibria: The first equilibrium lies well to the right and the second to the left, meaning that HF is extensively dissociated, but that the tight ion pairs reduce the thermodynamic activity coefficient of H3O+, so that the solution is less acidic. In concentrated solution, the additional HF causes the ion pair to dissociate with formation of the hydrogen-bonded hydrogen difluoride ion. + HF ⇌ H3O+ + HF−2The increase in free H3O+ due to this reaction accounts for the rapid increase in acidity, while fluoride ions are stabilized by strong hydrogen bonding to HF to form HF−2.
This interaction between the acid and its own conjugate base is an example of homoassociation. At the limit of 100% liquid HF, there is self-ionization 3 HF ⇌ H2F+ + HF−2which forms an acidic solution; the acidity of anhydrous HF can be increased further by the addition of Lewis acids such as SbF5, which can reduce H0 to −21. Dry hydrogen fluoride dissolves low-valent metal fluorides, as well as several molecular fluorides. Many proteins and carbohydrates can be recovered from it. In contrast, most non-fluoride inorganic chemicals react with HF rather than dissolving. Hydrogen fluoride is produced by the action of sulfuric acid on pure grades of the mineral fluorite and as a side-product of the extraction of the fertilizer precursor phosphoric acid from various minerals. See hydrofluoric acid; the anhydrous compound hydrogen fluoride is more used than its aqueous solution, hydrofluoric acid. HF serves. A component of high-octane petrol called "alkylate" is generated in alkylation units that combine C3 and C4 olefins and iso-butane to generate petrol.
HF is a reactive solvent in the electrochemical fluorination of organic compounds. In this approach, HF is oxidized in the presence of a hydrocarbon and the fluorine replaces C–H bonds with C–F bonds. P
Diethyl ether, or ether, is an organic compound in the ether class with the formula 2O, sometimes abbreviated as Et2O. It is a colorless volatile flammable liquid, it is used as a solvent in laboratories and as a starting fluid for some engines. It was used as a general anesthetic, until non-flammable drugs were developed, such as halothane, it has been used as a recreational drug to cause intoxication. Most diethyl ether is produced as a byproduct of the vapor-phase hydration of ethylene to make ethanol; this process uses solid-supported phosphoric acid catalysts and can be adjusted to make more ether if the need arises. Vapor-phase dehydration of ethanol over some alumina catalysts can give diethyl ether yields of up to 95%. Diethyl ether can be prepared both in laboratories and on an industrial scale by the acid ether synthesis. Ethanol is mixed with a strong acid sulfuric acid, H2SO4; the acid dissociates in the aqueous environment producing hydronium ions, H3O+. A hydrogen ion protonates the electronegative oxygen atom of the ethanol, giving the ethanol molecule a positive charge: CH3CH2OH + H3O+ → CH3CH2OH2+ + H2OA nucleophilic oxygen atom of unprotonated ethanol displaces a water molecule from the protonated ethanol molecule, producing water, a hydrogen ion and diethyl ether.
CH3CH2OH2+ + CH3CH2OH → H2O + H+ + CH3CH2OCH2CH3This reaction must be carried out at temperatures lower than 150 °C in order to ensure that an elimination product is not a product of the reaction. At higher temperatures, ethanol will dehydrate to form ethylene; the reaction to make diethyl ether is reversible, so an equilibrium between reactants and products is achieved. Getting a good yield of ether requires that ether be distilled out of the reaction mixture before it reverts to ethanol, taking advantage of Le Chatelier's principle. Another reaction that can be used for the preparation of ethers is the Williamson ether synthesis, in which an alkoxide performs a nucleophilic substitution upon an alkyl halide, it is important as a solvent in the production of cellulose plastics such as cellulose acetate. Diethyl ether has a high cetane number of 85–96 and is used as a starting fluid, in combination with petroleum distillates for gasoline and Diesel engines because of its high volatility and low flash point.
Ether starting fluid is sold and used in countries with cold climates, as it can help with cold starting an engine at sub-zero temperatures. For the same reason it is used as a component of the fuel mixture for carbureted compression ignition model engines. In this way diethyl ether is similar to one of its precursors, ethanol. Diethyl ether is a common laboratory aprotic solvent, it has limited solubility in water and dissolves 1.5 g/100 g water at 25 °C. This, coupled with its high volatility, makes it ideal for use as the non-polar solvent in liquid-liquid extraction; when used with an aqueous solution, the diethyl ether layer is on top as it has a lower density than the water. It is a common solvent for the Grignard reaction in addition to other reactions involving organometallic reagents. Due to its application in the manufacturing of illicit substances, it is listed in the Table II precursor under the United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances as well as substances such as acetone and sulfuric acid.
William T. G. Morton participated in a public demonstration of ether anesthesia on October 16, 1846 at the Ether Dome in Boston, Massachusetts. However, Crawford Williamson Long, is now known to have demonstrated its use as a general anesthetic in surgery to officials in Georgia, as early as March 30, 1842, Long publicly demonstrated ether's use as a surgical anesthetic on six occasions before the Boston demonstration. British doctors were aware of the anesthetic properties of ether as early as 1840 where it was prescribed in conjunction with opium. Diethyl ether supplanted the use of chloroform as a general anesthetic due to ether's more favorable therapeutic index, that is, a greater difference between an effective dose and a toxic dose. Diethyl ether increases tracheobronchial secretions. Diethyl ether could be mixed with other anesthetic agents such as chloroform to make C. E. mixture, or chloroform and alcohol to make A. C. E. Mixture. In the 21st century, ether is used; the use of flammable ether was displaced by nonflammable fluorinated hydrocarbon anesthetics.
Halothane was the first such anesthetic developed and other used inhaled anesthetics, such as isoflurane and sevoflurane, are halogenated ethers. Diethyl ether was found to have undesirable side effects, such as post-anesthetic nausea and vomiting. Modern anesthetic agents reduce these side effects. Prior to 2005 it was on the World Health Organization's List of Essential Medicines for use as an anesthetic. Ether was once used in pharmaceutical formulations. A mixture of alcohol and ether, one part of diethyl ether and three parts of ethanol, was known as "Spirit of ether", Hoffman's Anodyne or Hoffman's Drops. In the United States this concoction was removed from the Pharmacopeia at some point prior to June 1917, as a study published by William Procter, Jr. in the American Journal of Pharmacy as early as 1852 showed that there were differences in formulation to be found between commercial manufacturers, between international pharmacopoeia, from Hoffman's original recipe. The anesthetic and intoxicating effects of ether have made it a recreational drug.
Diethyl ether in anesthetic dosage is an inhalant which has a long history
CRC Handbook of Chemistry and Physics
The CRC Handbook of Chemistry and Physics is a comprehensive one-volume reference resource for science research in its 99th edition. It is sometimes nicknamed the "Rubber Bible" or the "Rubber Book", as CRC stood for "Chemical Rubber Company"; as late as the 1962–1963 edition the Handbook contained myriad information for every branch of science and engineering. Sections in that edition include: Mathematics and Physical Constants, Chemical Tables, Properties of Matter, Heat and Barometric Tables, Sound and Units, Miscellaneous. Earlier editions included sections such as "Antidotes of Poisons", "Rules for Naming Organic Compounds", "Surface Tension of Fused Salts", "Percent Composition of Anti-Freeze Solutions", "Spark-gap Voltages", "Greek Alphabet", "Musical Scales", "Pigments and Dyes", "Comparison of Tons and Pounds", "Twist Drill and Steel Wire Gauges" and "Properties of the Earth's Atmosphere at Elevations up to 160 Kilometers". Editions focus exclusively on chemistry and physics topics and eliminated much of the more "common" information.
22nd Edition – 44th Edition Section A: Mathematical Tables Section B: Properties and Physical Constants Section C: General Chemical Tables/Specific Gravity and Properties of Matter Section D: Heat and Hygrometry/Sound/Electricity and Magnetism/Light Section E: Quantities and Units/Miscellaneous Index 45th Edition – 70th Edition Section A: Mathematical Tables Section B: Elements and Inorganic Compounds Section C: Organic Compounds Section D: General Chemical Section E: General Physical Constants Section F: Miscellaneous Index 71st Edition – Current edition Section 1: Basic Constants and Conversion Factors Section 2: Symbols and Nomenclature Section 3: Physical Constants of Organic Compounds Section 4: Properties of the Elements and Inorganic Compounds Section 5: Thermochemistry and Kinetics Section 6: Fluid Properties Section 7: Biochemistry Section 8: Analytical Chemistry Section 9: Molecular Structure and Spectroscopy Section 10: Atomic and Optical Physics Section 11: Nuclear and Particle Physics Section 12: Properties of Solids Section 13: Polymer Properties Section 14: Geophysics and Acoustics Section 15: Practical Laboratory Data Section 16: Health and Safety Information Appendix A: Mathematical Tables Appendix B: CAS Registry Numbers and Molecular Formulas of Inorganic Substances Appendix B: Sources of Physical and Chemical Data IndexIn addition to an extensive line of engineering handbooks and references and textbooks across all scientific disciplines, CRC is today known as a leading publisher of books related to forensic sciences, forensic pathology and police sciences.
CORDIC PDF copy of the 8th edition, published in 1920 Handbook of Chemistry and Physics online Tables Relocated or Removed from CRC Handbook of Chemistry and Physics, 71st through 87th Editions
Argonium, an ion combining a proton and an argon atom, can be made in an electric discharge, was the first noble gas molecular ion to be found in interstellar space. Argonium is isoelectronic with hydrogen chloride, its dipole moment is 2.18 D for the ground state. The binding energy is 369 kJ mol−1; this is smaller than that of H+3 and many other protonated species, but more than that of H+2. Lifetimes of different vibrational states vary with isotope and become shorter for the more rapid high-energy vibrations. For ArH+ v=1 2.28, v=2 1.2, v=3 0.85, v=4 0.64, v=5 0.46 ms For ArD+ 9.09, 4.71, 3.27, 2.55, 2.11 msThe force constant in the bond is calculated at 3.88 mdyne/Å2. ArH+ + H2 → Ar + H+3 ArH+ + C → Ar + CH+ ArH+ + N → Ar + NH+ ArH+ + O → Ar + OH+ ArH+ + CO → Ar + COH+But the reverse reaction happens: Ar + H+2 → ArH+ + H. Ar + H+3 → *ArH+ + H2Ar+ + H2 has a cross section of 10−18 m2 for low energy, it drops off for energies over 100 eV Ar + H+2 has a cross sectional area of 6×10−19 m2 for low energy H+2, but when the energy exceeds 10 eV yield reduces, more Ar+ and H2 is produced instead.
Ar + H+3 has a maximum yield of ArH+ for energies between 0.75 and 1 eV with a cross section of 5×10−20 m2. 0.6 eV is needed to make the reaction proceed forward. Over 4eV more Ar+ and H starts to appear. Argonium is produced from Ar+ ions produced by cosmic rays and X-rays from neutral argon. Ar+ + H2 → *ArH+ + H 1.49 eVWhen ArH+ encounters an electron, dissociative recombination can occur, but it is slow for lower energy electrons, allowing ArH+ to survive for a much longer time than many other similar protonated cations. ArH+ + e− → Ar + HBecause ionisation potential of argon atoms is lower than that of the hydrogen molecule, the argon ion reacts with molecular hydrogen, but for helium and neon ions, they will strip an electron from a hydrogen molecule. Ar+ + H2 → ArH+ + H Ne+ + H2 → Ne + H+ + H He+ + H2 → He + H+ + H Artificial ArH+ made from earthly argon contains the isotope 40Ar rather than the cosmically abundant 36Ar. Artificially it is made by an electric discharge through an argon-hydrogen mixture.
Brault and Davis were the first to detect the molecule using infrared spectroscopy to observe vibration-rotation bands. The UV spectrum has two absorption points resulting in the ion breaking up; the 11.2 eV conversion to the B1Π state has a low dipole and so does not absorb much. A 15.8 eV to a repulsive A1Σ+ state is at a shorter wavelength than the Lymann limit, so there are few photons around to do this in space. ArH+ occurs in interstellar diffuse atomic hydrogen gas. For argonium to form, the fraction of molecular hydrogen H2 must be in the range 0.0001 to 0.001. Different molecular ions form in correlation with different concentrations of H2. Argonium is detected by its absorption lines at 617.525 GHz, 1234.602 GHz. These lines are due to the isotopolog 36Ar1H+ undergoing rotational transitions; the lines have been detected in the direction of the galactic centre SgrB2 and SgrB2, G34.26+0.15, W31C, W49, W51e, however where absorption lines are observed, argonium is not to be in the microwave source, but instead in the gas in front of it.
Emission lines are found in the Crab Nebula. In the Crab Nebula ArH+ occurs in several spots revealed by emission lines; the strongest place is in the Southern Filament. This is the place with the strongest concentration of Ar+ and Ar2+ ions; the column density of ArH + in the Crab Nebula is between 1013 atoms per square centimeter. Possible the energy required to excite the ions so that can emit comes from collisions with electrons or hydrogen molecules. Towards the Milky Way centre the column density of ArH+ is around 2×1013 cm−2. Two isotopologs of argonium 36ArH+ and 38ArH+ are known to be in a distant unnamed galaxy with z=0.88582, on the line of sight to the blazar PKS 1830−211. Electron neutralization and destruction of argonium outcompletes the formation rate in space if the H2 concentration is below 1 in 10−4. Using the McMath solar Fourier transform spectrometer at Kitt Peak National Observatory, James W. Brault and Sumner P. Davis observed ArH+ vibration-rotation infrared lines for the first time.
J. W. C. Johns observed the infrared spectrum. Argon facilitates the reaction of tritium with double bonds in fatty acids by forming an ArT+ intermediate; when gold is sputtered with an argon-hydrogen plasma, the actual displacement of gold is done by ArH+. Argonium is the name of a major tectonic scarp on the edge of the Rheasilvia impact basin in the Quadrangle Av-12 on 4 Vesta. Argonium is an erroneous name for Argon
Nitrogen is a chemical element with symbol N and atomic number 7. It was first discovered and isolated by Scottish physician Daniel Rutherford in 1772. Although Carl Wilhelm Scheele and Henry Cavendish had independently done so at about the same time, Rutherford is accorded the credit because his work was published first; the name nitrogène was suggested by French chemist Jean-Antoine-Claude Chaptal in 1790, when it was found that nitrogen was present in nitric acid and nitrates. Antoine Lavoisier suggested instead the name azote, from the Greek ἀζωτικός "no life", as it is an asphyxiant gas. Nitrogen is the lightest member of group 15 of the periodic table called the pnictogens; the name comes from the Greek πνίγειν "to choke", directly referencing nitrogen's asphyxiating properties. It is a common element in the universe, estimated at about seventh in total abundance in the Milky Way and the Solar System. At standard temperature and pressure, two atoms of the element bind to form dinitrogen, a colourless and odorless diatomic gas with the formula N2.
Dinitrogen forms about 78 % of Earth's atmosphere. Nitrogen occurs in all organisms in amino acids, in the nucleic acids and in the energy transfer molecule adenosine triphosphate; the human body contains about 3% nitrogen by mass, the fourth most abundant element in the body after oxygen and hydrogen. The nitrogen cycle describes movement of the element from the air, into the biosphere and organic compounds back into the atmosphere. Many industrially important compounds, such as ammonia, nitric acid, organic nitrates, cyanides, contain nitrogen; the strong triple bond in elemental nitrogen, the second strongest bond in any diatomic molecule after carbon monoxide, dominates nitrogen chemistry. This causes difficulty for both organisms and industry in converting N2 into useful compounds, but at the same time means that burning, exploding, or decomposing nitrogen compounds to form nitrogen gas releases large amounts of useful energy. Synthetically produced ammonia and nitrates are key industrial fertilisers, fertiliser nitrates are key pollutants in the eutrophication of water systems.
Apart from its use in fertilisers and energy-stores, nitrogen is a constituent of organic compounds as diverse as Kevlar used in high-strength fabric and cyanoacrylate used in superglue. Nitrogen is a constituent including antibiotics. Many drugs are mimics or prodrugs of natural nitrogen-containing signal molecules: for example, the organic nitrates nitroglycerin and nitroprusside control blood pressure by metabolizing into nitric oxide. Many notable nitrogen-containing drugs, such as the natural caffeine and morphine or the synthetic amphetamines, act on receptors of animal neurotransmitters. Nitrogen compounds have a long history, ammonium chloride having been known to Herodotus, they were well known by the Middle Ages. Alchemists knew nitric acid as aqua fortis, as well as other nitrogen compounds such as ammonium salts and nitrate salts; the mixture of nitric and hydrochloric acids was known as aqua regia, celebrated for its ability to dissolve gold, the king of metals. The discovery of nitrogen is attributed to the Scottish physician Daniel Rutherford in 1772, who called it noxious air.
Though he did not recognise it as an different chemical substance, he distinguished it from Joseph Black's "fixed air", or carbon dioxide. The fact that there was a component of air that does not support combustion was clear to Rutherford, although he was not aware that it was an element. Nitrogen was studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, Joseph Priestley, who referred to it as burnt air or phlogisticated air. Nitrogen gas was inert enough that Antoine Lavoisier referred to it as "mephitic air" or azote, from the Greek word άζωτικός, "no life". In an atmosphere of pure nitrogen, animals died and flames were extinguished. Though Lavoisier's name was not accepted in English, since it was pointed out that all gases are mephitic, it is used in many languages and still remains in English in the common names of many nitrogen compounds, such as hydrazine and compounds of the azide ion, it led to the name "pnictogens" for the group headed by nitrogen, from the Greek πνίγειν "to choke".
The English word nitrogen entered the language from the French nitrogène, coined in 1790 by French chemist Jean-Antoine Chaptal, from the French nitre and the French suffix -gène, "producing", from the Greek -γενής. Chaptal's meaning was that nitrogen is the essential part of nitric acid, which in turn was produced from nitre. In earlier times, niter had been confused with Egyptian "natron" – called νίτρον in Greek – which, despite the name, contained no nitrate; the earliest military and agricultural applications of nitrogen compounds used saltpeter, most notably in gunpowder, as fertiliser. In 1910, Lord Rayleigh discovered that an electrical discharge in nitrogen gas produced "active nitrogen", a monatomic allotrope of nitrogen; the "whirling cloud of brilliant yellow light
Methylidyne called carbyne, is an organic compound with the chemical formula CH•. Methylidyne is the simplest carbyne, it is a reactive gas, destroyed in ordinary conditions but is abundant in the interstellar medium. In October 2016, astronomers reported that the basic chemical ingredients of life – the carbon-hydrogen molecule, the carbon-hydrogen positive ion, the carbon ion – are the result of ultraviolet light from stars, rather than in other ways, such as the result of turbulent events related to supernovae and young stars, as thought earlier; these results have given new light to the formation of organic compounds in the early development of life on earth. The trivial name carbyne is the preferred IUPAC name; the systematic names methylidyne, hydridocarbon, valid IUPAC names, are constructed according to the substitutive and additive nomenclatures, respectively. Methylidyne is viewed as methane with three hydrogen atoms removed. By default, this name pays no regard to the radicality of methylidyne.
When the radicality is considered, the radical states with one unpaired electron are named methylylidene, whereas the radical excited states with three unpaired electrons are named methanetriyl. As an odd-electron species, CH is a radical; the ground state is a doublet. The first two excited states are a doublet; the quartet lies at 71 kJ above the ground state. Reactions of the doublet radical with non-radical species involves insertion or addition, whereas reactions of the quartet radical involves only abstraction. • + H2O → • + H2 or • 3• + H2O → + • Methylidyne-like species are implied intermediates in the Fischer–Tropsch process, the hydrogenation of CO into hydrocarbons. Methylidyne entities are assumed to bond to the catalyst's surface. A hypothetical sequence is: MnCO + 1/2 H2 → MnCOH MnCOH + H2 → MnCH + H2OA molecular example of an MnCH is HCCo39. MnCH + 1/2 H2 → MnCH2The methylene ligand is poised couple to CO or to another methylene, thereby growing the C–C chain; the methylylidyne group can exhibit both Lewis basic character.
Such behavior is only of theoretical interest. Methylidine can be prepared from bromoform. Methylene group Methylene bridge