Transition metal carbene complex
A transition metal carbene complex is an organometallic compound featuring a divalent organic ligand. The divalent organic ligand coordinated to the metal center is called a carbene. Carbene complexes for all transition metals have been reported. Many methods for synthesizing them and reactions utilizing them have been reported; the term carbene ligand is a formalism since many are not derived from carbenes and none exhibit the reactivity characteristic of carbenes. Described as M=CR2, they represent a class of organic ligands intermediate between alkyls and carbynes, they feature in some catalytic reactions alkene metathesis, are of value in the preparation of some fine chemicals. Metal carbene complexes are classified into two types; the Fischer carbenes named after Ernst Otto Fischer feature strong π-acceptors at the metal and being electrophilic at the carbene carbon atom. Schrock carbenes, named after Richard R. Schrock, are characterized by more nucleophilic carbene carbon centers, these species feature higher valent metals.
N-Heterocyclic carbenes were popularlized following Arduengo's isolation of a stable free carbene in 1991. Reflecting the growth of the area, carbene complexes are now known with a broad range of different reactivities and diverse substituents, it is not possible to classify a carbene complex with regards to its electrophilicity or nucleophilicity. Fischer carbenes are found with: low oxidation state metal center middle and late transition metals Fe, Mo, Cr π-acceptor metal ligands π-donor substituents on the carbene atom such as alkoxy and alkylated amino groups; the chemical bonding is based on σ-type electron donation of the filled lone pair orbital of the carbene atom to an empty metal d-orbital, π back bonding of a filled metal d-orbital to the empty p-orbital on carbon. An example is the complex 5Cr=CPh. Fischer carbenes can be likened to ketones, with the carbene carbon being electrophilic, much like the carbonyl carbon of a ketone. Like ketones, Fischer carbene species can undergo aldol-like reactions.
The hydrogen atoms attached to the carbon α to the carbene carbon are acidic, can be deprotonated by a base such as n-butyllithium, to give a nucleophile which can undergo further reaction. This carbene is the starting material for other reactions such as the Wulff-Dötz reaction. Schrock carbenes do not have π-accepting ligands; these complexes are nucleophilic at the carbene carbon atom. Schrock carbenes are found with: high oxidation state metal center early transition metals Ti, Ta π-donor ligands hydrogen and alkyl substituents on carbenoid carbon. Bonding in such complexes can be viewed as the coupling of a triplet state metal and triplet carbene; these bonds are polarized towards carbon and therefore the carbene atom is a nucleophile. An example of a Schrock carbene is the compound Ta3, with a tantalum center doubly bonded to a neopentylidene ligand as well as three neopentyl ligands. An example of interest in organic synthesis is Tebbe's reagent. N-Heterocyclic carbenes are common carbene ligands.
They are popular because they are more prepared than Schrock and Fischer carbenes. In fact many NHCs are isolated as the free ligand. Being stabilized by π-donating substituents, NHCs are powerful σ-donors but π-bonding with the metal is weak. For this reason, the bond between the carbon and the metal center is represented by a single dative bond, whereas Fischer and Schrock carbenes are depicted with double bonds to metal. Continuing with this analogy, NHCs are compared with trialkylphosphine ligands. Like phosphines, NHCs serve as spectator ligands that influence catalysis through a combination of electronic and steric effects, but they do not directly bind substrates. Carbenes without a metal ligand have been produced in the lab. Carbene radicals are long-lived reaction intermediates found with: low oxidation state metal center with singly occupied dz2 orbital middle and late transition metal, e.g. Co σ-donor and π-acceptor ligand π-acceptor substituents on the ligand such as carbonyl or sulfonyl groups.
The chemical bond present in carbene radicals is described as aspects of both Fischer and Schrock carbenes. The main applications of metal carbenes involves none of the above classes of compounds, but rather heterogeneous catalysts used for alkene metathesis in the Shell higher olefin process. A variety of related reactions are used to interconvert light alkenes, e.g. butenes and ethylene. Carbene-complexes are invoked as intermediates in the Fischer–Tropsch route to hydrocarbons. A variety of soluble carbene reagents the Grubbs' and molybdenum-imido catalysts have been applied to laboratory-scale synthesis of natural products and materials science. In the nucleophilic abstraction reaction, a methyl group can be abstracted from a Fischer carbene for further reaction; the characterization of 5W in the 1960s is cited as the starting point of the area, although carbenoid ligands had been implicated. Ernst Otto Fischer, for this and other achievements in organometalic chemistry, was awarded the 1973 Nobel Prize in Chemistry.
Carbyne Carbene radical
Three-center four-electron bond
The 3-center 4-electron bond is a model used to explain bonding in certain hypervalent molecules such as tetratomic and hexatomic interhalogen compounds, sulfur tetrafluoride, the xenon fluorides, the bifluoride ion. It is known as the Pimentel–Rundle three-center model after the work published by George C. Pimentel in 1951, which built on concepts developed earlier by Robert E. Rundle for electron-deficient bonding. An extended version of this model is used to describe the whole class of hypervalent molecules such as phosphorus pentafluoride and sulfur hexafluoride as well as multi-center π-bonding such as ozone and sulfur trioxide. While the term "hypervalent" was not introduced in the chemical literature until 1969, Irving Langmuir and G. N. Lewis debated the nature of bonding in hypervalent molecules as early as 1921. While Lewis supported the viewpoint of expanded octet, invoking s-p-d hybridized orbitals and maintaining 2c–2e bonds between neighboring atoms, Langmuir instead opted for maintaining the octet rule, invoking an ionic basis for bonding in hypervalent compounds.
In a 1951 seminal paper, Pimentel rationalized the bonding in hypervalent trihalide ions via a molecular orbital description, building on the concept of the "half-bond" introduced by Rundle in 1947. In this model, two of the four electrons occupy an all in-phase bonding MO, while the other two occupy a non-bonding MO, leading to an overall bond order of 0.5 between adjacent atoms. More recent theoretical studies on hypervalent molecules support the Langmuir view, confirming that the octet rule serves as a good first approximation to describing bonding in the s- and p-block elements. Triiodide Xenon difluoride Krypton difluoride Radon difluoride Argon fluorohydride Bifluoride Hydrogen Bonding Carboxylates Amides Ozone Azide Allyl anion The σ molecular orbitals of triiodide can be constructed by considering the in-phase and out-of-phase combinations of the central atom's p orbital with the p orbitals of the peripheral atoms; this exercise generates the diagram at right. Three molecular orbitals result from the combination of the three relevant atomic orbitals, with the four electrons occupying the two MOs lowest in energy – a bonding MO delocalized across all three centers, a non-bonding MO localized on the peripheral centers.
Using this model, one sidesteps the need to invoke hypervalent bonding considerations at the central atom, since the bonding orbital consists of two 2-center-1-electron bonds, the other two electrons occupy the non-bonding orbital. In the natural bond orbital viewpoint of 3c–4e bonding, the triiodide anion is constructed from the combination of the diiodine σ molecular orbitals and an iodide lone pair; the I− lone pair acts as a 2-electron donor, while the I2 σ* antibonding orbital acts as a 2-electron acceptor. Combining the donor and acceptor in in-phase and out-of-phase combinations results in the diagram depicted at right. Combining the donor lone pair with the acceptor σ* antibonding orbital results in an overall lowering in energy of the highest-occupied orbital. While the diagram depicted in Figure 2 shows the right-hand atom as the donor, an equivalent diagram can be constructed using the left-hand atom as the donor; this bonding scheme is succinctly summarized by the following two resonance structures: I—I···I− ↔ I−···I—I, which when averaged reproduces the I—I bond order of 0.5 obtained both from natural bond orbital analysis and from molecular orbital theory.
More recent theoretical investigations suggest the existence of a novel type of donor-acceptor interaction that may dominate in triatomic species with so-called "inverted electronegativity". Molecules of theoretical curiosity such as neon difluoride and berylium dilithide represent examples of inverted electronegativity; as a result of unusual bonding situation, the donor lone pair ends up with significant electron density on the central atom, while the acceptor is the "out-of-phase" combination of the p orbitals on the peripheral atoms. This bonding scheme is depicted in Figure 3 for the theoretical noble gas dihalide NeF2; the valence bond description and accompanying resonance structures A—B···C− ↔ A−···B—C suggest that molecules exhibiting 3c–4e bonding can serve as models for studying the transition states of bimolecular nucleophilic substitution reactions. Hypervalent molecule
A hypervalent molecule is a molecule that contains one or more main group elements bearing more than eight electrons in their valence shells. Phosphorus pentachloride, sulfur hexafluoride, chlorine trifluoride, the chlorite ion, the triiodide ion are examples of hypervalent molecules. Hypervalent molecules were first formally defined by Jeremy I. Musher in 1969 as molecules having central atoms of group 15–18 in any valence other than the lowest. Several specific classes of hypervalent molecules exist: Hypervalent iodine compounds are useful reagents in organic chemistry Tetra-, penta- and hexavalent phosphorus and sulfur compounds Noble gas compounds Halogen polyfluorides N-X-L nomenclature, introduced collaboratively by the research groups of Martin and Kochi in 1980, is used to classify hypervalent compounds of main group elements, where: N represents the number of valence electrons X is the chemical symbol of the central atom L the number of ligands to the central atomExamples of N-X-L nomenclature include: XeF2, 10-Xe-2 PCl5, 10-P-5 SF6, 12-S-6 IF7, 14-I-7 The debate over the nature and classification of hypervalent molecules goes back to Gilbert N. Lewis and Irving Langmuir and the debate over the nature of the chemical bond in the 1920s.
Lewis maintained the importance of the two-center two-electron bond in describing hypervalence, thus using expanded octets to account for such molecules. Using the language of orbital hybridization, the bonds of molecules like PF5 and SF6 were said to be constructed from sp3dn orbitals on the central atom. Langmuir, on the other hand, upheld the dominance of the octet rule and preferred the use of ionic bonds to account for hypervalence without violating the rule. In the late 1920s and 1930s, Sugden argued for the existence of a two-center one-electron bond and thus rationalized bonding in hypervalent molecules without the need for expanded octets or ionic bond character. In the 1940s and 1950s, Rundle and Pimentel popularized the idea of the three-center four-electron bond, the same concept which Sugden attempted to advance decades earlier; the attempt to prepare hypervalent organic molecules began with Hermann Staudinger and Georg Wittig in the first half of the twentieth century, who sought to challenge the extant valence theory and prepare nitrogen and phosphorus-centered hypervalent molecules.
The theoretical basis for hypervalency was not delineated until J. I. Musher's work in 1969. In 1990, Magnusson published a seminal work definitively excluding the significance of d-orbital hybridization in the bonding of hypervalent compounds of second-row elements; this had long been a point of contention and confusion in describing these molecules using molecular orbital theory. Part of the confusion here originates from the fact that one must include d-functions in the basis sets used to describe these compounds, the contribution of the d-function to the molecular wavefunction is large; these facts were interpreted to mean that d-orbitals must be involved in bonding. However, Magnusson concludes in his work that d-orbital involvement is not implicated in hypervalency. A 2013 study showed that although the Pimentel ionic model best accounts for the bonding of hypervalent species, the energetic contribution of an expanded octet structure is not null. In this modern valence bond theory study of the bonding of xenon difluoride, it was found that ionic structures account for about 81% of the overall wavefunction, of which 70% arises from ionic structures employing only the p orbital on xenon while 11% arises from ionic structures employing an s d z 2 hybrid on xenon.
The contribution of a formally hypervalent structure employing an orbital of sp3d hydridization on xenon accounts for 11% of the wavefunction, with a diradical contribution making up the remaining 8%. The 11% sp3d contribution results in a net stabilization of the molecule by 7.2 kcal mol-1, a minor but significant fraction of the total energy of the total bond energy. Other studies have found minor but non-negligible energetic contributions from expanded octet structures in SF6 and XeF6. Despite the lack of chemical realism, the IUPAC recommends the drawing of expanded octet structures for functional groups like sulfones and phosphoranes, in order to avoid the drawing of a large number of formal charges or partial single bonds. Both the term and concept of hypervalency still fall under criticism. In 1984, in response to this general controversy, Paul von Ragué Schleyer proposed the replacement of'hypervalency' with use of the term hypercoordination because this term does not imply any mode of chemical bonding and the question could thus be avoided altogether.
The concept itself has been criticized by Ronald Gillespie who, based on an analysis of electron localization functions, wrote in 2002 that "as there is no fundamental difference between the bonds in hypervalent and non-hypervalent molecules there is no reason to continue to use the term hypervalent."Fo
An oxide is a chemical compound that contains at least one oxygen atom and one other element in its chemical formula. "Oxide" itself is the dianion of an O2 -- atom. Metal oxides thus contain an anion of oxygen in the oxidation state of −2. Most of the Earth's crust consists of solid oxides, the result of elements being oxidized by the oxygen in air or in water. Hydrocarbon combustion affords the two principal carbon oxides: carbon monoxide and carbon dioxide. Materials considered pure elements develop an oxide coating. For example, aluminium foil develops a thin skin of Al2O3 that protects the foil from further corrosion. Individual elements can form multiple oxides, each containing different amounts of the element and oxygen. In some cases these are distinguished by specifying the number of atoms as in carbon monoxide and carbon dioxide, in other cases by specifying the element's oxidation number, as in iron oxide and iron oxide. Certain elements can form many different oxides, such as those of nitrogen.
Due to its electronegativity, oxygen forms stable chemical bonds with all elements to give the corresponding oxides. Noble metals are prized because they resist direct chemical combination with oxygen, substances like gold oxide must be generated by indirect routes. Two independent pathways for corrosion of elements are oxidation by oxygen; the combination of water and oxygen is more corrosive. All elements burn in an atmosphere of oxygen or an oxygen-rich environment. In the presence of water and oxygen, some elements— sodium—react to give the hydroxides. In part, for this reason and alkaline earth metals are not found in nature in their metallic, i.e. native, form. Cesium is so reactive with oxygen that it is used as a getter in vacuum tubes, solutions of potassium and sodium, so-called NaK are used to deoxygenate and dehydrate some organic solvents; the surface of most metals consists of hydroxides in the presence of air. A well-known example is aluminium foil, coated with a thin film of aluminium oxide that passivates the metal, slowing further corrosion.
The aluminum oxide layer can be built to greater thickness by the process of electrolytic anodizing. Though solid magnesium and aluminum react with oxygen at STP—they, like most metals, burn in air, generating high temperatures. Finely grained powders of most metals can be dangerously explosive in air, they are used in solid-fuel rockets. In dry oxygen, iron forms iron oxide, but the formation of the hydrated ferric oxides, Fe2O3−x2x, that comprise rust requires oxygen and water. Free oxygen production by photosynthetic bacteria some 3.5 billion years ago precipitated iron out of solution in the oceans as Fe2O3 in the economically important iron ore hematite. Oxides have a range of different structures, from individual molecules to polymeric and crystalline structures. At standard conditions, oxides may range from solids to gases. Oxides of most metals adopt polymeric structures; the oxide links three metal atoms or six metal atoms. Because the M-O bonds are strong and these compounds are crosslinked polymers, the solids tend to be insoluble in solvents, though they are attacked by acids and bases.
The formulas are deceptively simple. Many are nonstoichiometric compounds; some important gaseous oxides. Examples of molecular oxides are carbon monoxide. All simple oxides of nitrogen are molecular, e.g. NO, N2O, NO2 and N2O4. Phosphorus pentoxide is a more complex molecular oxide with a deceptive name, the real formula being P4O10; some polymeric oxides depolymerize when heated to give molecules, examples being selenium dioxide and sulfur trioxide. Tetroxides are rare; the more common examples: ruthenium tetroxide, osmium tetroxide, xenon tetroxide. Many oxyanions are known, such as polyoxometalates. Oxycations are rarer, some examples being nitrosonium and uranyl. Of course many compounds are known with other groups. In organic chemistry, these include many related carbonyl compounds. For the transition metals, many oxo complexes are known as well as oxyhalides. Conversion of a metal oxide to the metal is called reduction; the reduction can be induced by many reagents. Many metal oxides convert to metals by heating.
Metals are "won" from their oxides by chemical reduction, i.e. by the addition of a chemical reagent. A common and cheap reducing agent is carbon in the form of coke; the most prominent example is that of iron ore smelting. Many reactions are involved, but the simplified equation is shown as: 2 Fe2O3 + 3 C → 4 Fe + 3 CO2Metal oxides can be reduced by organic compounds; this redox process is the basis for many important transformations in chemistry, such as the detoxification of drugs by the P450 enzymes and the production of ethylene oxide, converted to antifreeze. In such systems, the metal center transfers an oxide ligand to the organic compound followed by regeneration of the metal oxide by oxygen in the air. Metals that are lower in the reactivity series can be reduced by heating alone. For example, silver oxide decomposes at 200 °C: 2 Ag2O → 4 Ag + O2 Metals that are more reactive displace the oxide of the metals that are less reactive. For example, zinc is more reactive than copper, so it displaces copper oxide to form zinc oxide: Zn + CuO → ZnO + Cu Apart from metals, hydrogen can displace metal oxides to form hydrogen oxide
In chemistry, a nitride is a compound of nitrogen where nitrogen has a formal oxidation state of −3. Nitrides are a large class of compounds with a wide range of applications; the nitride ion, N3−, is never encountered in protic solution because it is so basic that it would be protonated immediately. Its ionic radius is estimated to be 140 pm. Like carbides, nitrides are refractory materials owing to their high lattice energy which reflects the strong attraction of "N3−" for the metal cation. Thus, titanium nitride and silicon nitride are used as cutting hard coatings. Hexagonal boron nitride, which adopts a layered structure, is a useful high-temperature lubricant akin to molybdenum disulfide. Nitride compounds have large band gaps, thus nitrides are insulators or wide bandgap semiconductors; the wide band gap material gallium nitride is prized for emitting blue light in LEDs. Like some oxides, nitrides can absorb hydrogen and have been discussed in the context of hydrogen storage, e.g. lithium nitride.
Classification of such a varied group of compounds is somewhat arbitrary. Compounds where nitrogen is not assigned −3 oxidation state are not included, such as nitrogen trichloride where the oxidation state is +3. Only one alkali metal nitride is stable, the purple-reddish lithium nitride, which forms when lithium burns in an atmosphere of N2. Sodium nitride remains a laboratory curiosity; the nitrides of the alkaline earth metals have the formula. Examples include Be3N2, Mg3N2, Ca3N2, Sr3N2; the nitrides of electropositive metals hydrolyze upon contact with water, including the moisture in the air: Mg3N2 + 6 H2O → 3 Mg2 + 2 NH3 Boron nitride exists as several forms. Nitrides of silicon and phosphorus are known, but only the former is commercially important; the nitrides of aluminium and indium adopt diamond-like wurtzite structure in which each atom occupies tetrahedral sites. For example, in aluminium nitride, each aluminium atom has four neighboring nitrogen atoms at the corners of a tetrahedron and each nitrogen atom has four neighboring aluminium atoms at the corners of a tetrahedron.
This structure is like hexagonal diamond. Thallium nitride, Tl3N is known, but thallium nitride, TlN, is not. For the group 3 metals, ScN and YN are both known. Group 4, 5, 6 transition metals, the titanium and chromium groups all form nitrides, they are chemically stable. Representative is titanium nitride. Sometimes these materials are called "interstitial nitrides." Nitrides of the Group 7 and 8 transition metals decompose readily. For example, iron nitride, Fe2N decomposes at 200 °C. Platinum nitride and osmium nitride may contain N2 units, as such should not be called nitrides. Nitrides of heavier members from group 11 and 12 are less stable than copper nitride, Cu3N and Zn3N2: dry silver nitride is a contact explosive which may detonate from the slightest touch a falling water droplet. Many metals form molecular nitrido complexes; the main group elements form some molecular nitrides. Cyanogen and tetrasulfur tetranitride are rare examples of a molecular binary nitrides, they dissolve in nonpolar solvents.
Both undergo polymerization. S4N4 is unstable with respect to the elements, but less so that the isostructural Se4N4. Heating S4N4 gives a polymer, a variety of molecular sulfur nitride anions and cations are known. Related to but distinct from nitride is pernitride, N2−2
In chemistry, the term transition metal has three possible meanings: The IUPAC definition defines a transition metal as "an element whose atom has a filled d sub-shell, or which can give rise to cations with an incomplete d sub-shell". Many scientists describe a "transition metal" as any element in the d-block of the periodic table, which includes groups 3 to 12 on the periodic table. In actual practice, the f-block lanthanide and actinide series are considered transition metals and are called "inner transition metals". Cotton and Wilkinson expand the brief IUPAC definition by specifying; as well as the elements of groups 4 to 11, they add scandium and yttrium in group 3, which have a filled d subshell in the metallic state. Lanthanum and actinium in group 3 are, classified as lanthanides and actinides respectively. English chemist Charles Bury first used the word transition in this context in 1921, when he referred to a transition series of elements during the change of an inner layer of electrons from a stable group of 8 to one of 18, or from 18 to 32.
These elements are now known as the d-block. In the d-block the atoms of the elements have between 10 d electrons; the elements of groups 4–11 are recognized as transition metals, justified by their typical chemistry, i.e. a large range of complex ions in various oxidation states, colored complexes, catalytic properties either as the element or as ions. Sc and Y in group 3 are generally recognized as transition metals. However, the elements La–Lu and Ac–Lr and group 12 attract different definitions from different authors. Many chemistry textbooks and printed periodic tables classify La and Ac as group 3 elements and transition metals, since their atomic ground-state configurations are s2d1 like Sc and Y; the elements Ce–Lu are considered as the "lanthanide" series and Th–Lr as the "actinide" series.</ref> The two series together are classified as f-block elements, or as "inner transition elements". Some inorganic chemistry textbooks include Ac with the actinides; this classification is based on similarities in chemical behaviour and defines 15 elements in each of the two series though they correspond to the filling of an f subshell, which can only contain 14 electrons.
A third classification defines the f-block elements as La–Yb and Ac–No, while placing Lu and Lr in group 3. This is based on the Aufbau principle for filling electron subshells, in which 4f is filled before 5d, so that the f subshell is full at Yb, while Lu has an s2f14d1 configuration; however La and Ac are exceptions to the Aufbau principle with electron configuration s2d1, so it is not clear from atomic electron configurations whether La or Lu should be considered as transition metals. Zinc and mercury are excluded from the transition metals, as they have the electronic configuration d10s2, with no incomplete d shell. In the oxidation state +2 the ions have the electronic configuration …d10. However, these elements can exist in other oxidation states, including the +1 oxidation state, as in the diatomic ion Hg2+2; the group 12 elements Zn, Cd and Hg may therefore, under certain criteria, be classed as post-transition metals in this case. However, it is convenient to include these elements in a discussion of the transition elements.
For example, when discussing the crystal field stabilization energy of first-row transition elements, it is convenient to include the elements calcium and zinc, as both Ca2+ and Zn2+ have a value of zero, against which the value for other transition metal ions may be compared. Another example occurs in the Irving–Williams series of stability constants of complexes; the recent synthesis of mercury fluoride has been taken by some to reinforce the view that the group 12 elements should be considered transition metals, but some authors still consider this compound to be exceptional. Although meitnerium and roentgenium are within the d-block and are expected to behave as transition metals analogous to their lighter congeners iridium and gold, this has not yet been experimentally confirmed. Early transition metals are on the left side of the periodic table from group 3 to group 7. Late transition metals are on the right side of the d-block, from group 8 to 11; the general electronic configuration of the d-block elements is d1–10n s0–2.
The period 6 and 7 transition metals add f0–14 electrons, which are omitted from the tables below. The Madelung rule predicts that the typical electronic structure of transition metal atoms can be written as ns2dm where the inner d orbital is predicted to be filled after the valence-shell's s orbital is filled; this rule is however only approximate – it only holds for some of the transition elements, only in their neutral ground state. The d-sub-shell is denoted as d - sub-shell; the number of s electrons in the outermost s sub-shell is one or two except palladium, with no electron in that s-sub shell in its ground state. The s-sub-shell in the valence shell is represented as e.g. 4s. In the periodic table, the transition metals are present in eight groups, with some authors including some elements in groups 3 o
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