In coordination chemistry, a ligand is an ion or molecule that binds to a central metal atom to form a coordination complex. The bonding with the metal involves formal donation of one or more of the ligand's electron pairs; the nature of metal–ligand bonding can range from covalent to ionic. Furthermore, the metal–ligand bond order can range from one to three. Ligands are viewed as Lewis bases, although rare cases are known to involve Lewis acidic "ligands". Metals and metalloids are bound to ligands in all circumstances, although gaseous "naked" metal ions can be generated in a high vacuum. Ligands in a complex dictate the reactivity of the central atom, including ligand substitution rates, the reactivity of the ligands themselves, redox. Ligand selection is a critical consideration in many practical areas, including bioinorganic and medicinal chemistry, homogeneous catalysis, environmental chemistry. Ligands are classified in many ways, including: charge, the identity of the coordinating atom, the number of electrons donated to the metal.
The size of a ligand is indicated by its cone angle. The composition of coordination complexes have been known since the early 1800s, such as Prussian blue and copper vitriol; the key breakthrough occurred when Alfred Werner reconciled isomers. He showed, among other things, that the formulas of many cobalt and chromium compounds can be understood if the metal has six ligands in an octahedral geometry; the first to use the term "ligand" were Alfred Stock and Carl Somiesky, in relation to silicon chemistry. The theory allows one to understand the difference between coordinated and ionic chloride in the cobalt ammine chlorides and to explain many of the inexplicable isomers, he resolved the first coordination complex called hexol into optical isomers, overthrowing the theory that chirality was associated with carbon compounds. In general, ligands are viewed as the metals as electron acceptors; this is because the ligand and central metal are bonded to one another, the ligand is providing both electrons to the bond instead of the metal and ligand each providing one electron.
Bonding is described using the formalisms of molecular orbital theory. The HOMO can be of ligands or metal character. Ligands and metal ions can be ordered in many ways. Metal ions preferentially bind certain ligands. In general,'hard' metal ions prefer weak field ligands, whereas'soft' metal ions prefer strong field ligands. According to the molecular orbital theory, the HOMO of the ligand should have an energy that overlaps with the LUMO of the metal preferential. Metal ions bound to strong-field ligands follow the Aufbau principle, whereas complexes bound to weak-field ligands follow Hund's rule. Binding of the metal with the ligands results in a set of molecular orbitals, where the metal can be identified with a new HOMO and LUMO and a certain ordering of the 5 d-orbitals. In an octahedral environment, the 5 otherwise degenerate d-orbitals split in sets of 2 and 3 orbitals. 3 orbitals of low energy: dxy and dyz 2 of high energy: dz2 and dx2−y2The energy difference between these 2 sets of d-orbitals is called the splitting parameter, Δo.
The magnitude of Δo is determined by the field-strength of the ligand: strong field ligands, by definition, increase Δo more than weak field ligands. Ligands can now be sorted according to the magnitude of Δo; this ordering of ligands is invariable for all metal ions and is called spectrochemical series. For complexes with a tetrahedral surrounding, the d-orbitals again split into two sets, but this time in reverse order. 2 orbitals of low energy: dz2 and dx2−y2 3 orbitals of high energy: dxy and dyzThe energy difference between these 2 sets of d-orbitals is now called Δt. The magnitude of Δt is smaller than for Δo, because in a tetrahedral complex only 4 ligands influence the d-orbitals, whereas in an octahedral complex the d-orbitals are influenced by 6 ligands; when the coordination number is neither octahedral nor tetrahedral, the splitting becomes correspondingly more complex. For the purposes of ranking ligands, the properties of the octahedral complexes and the resulting Δo has been of primary interest.
The arrangement of the d-orbitals on the central atom, has a strong effect on all the properties of the resulting complexes. E.g. the energy differences in the d-orbitals has a strong effect in the optical absorption spectra of metal complexes. It turns out that valence electrons occupying orbitals with significant 3 d-orbital character absorb in the 400–800 nm region of the spectrum; the absorption of light by these electrons can be correlated to the ground state of the metal complex, which reflects the bonding properties of the ligands. The relative change in energy of the d-orbitals as a function of the field-strength of the ligands is described in Tanabe–Sugano diagrams. In cases where the ligand has low energy LUMO, such orbitals participate in the bonding; the metal–ligand bond can be further stabilised by a formal donation of electron density back to the ligand in a process known as back-bonding. In this case a filled, c
Nickel chloride, is the chemical compound NiCl2. The anhydrous salt is yellow. Nickel chloride, in various forms, is the most important source of nickel for chemical synthesis; the nickel chlorides are deliquescent. Nickel salts have been shown to be carcinogenic to the lungs and nasal passages in cases of long-term inhalation exposure; the largest scale production of nickel chloride involves the extraction with hydrochloric acid of nickel matte and residues obtained from roasting refining nickel-containing ores. Nickel chloride is not prepared in the laboratory because it is inexpensive and has a long shelf-life. Heating the hexahydrate in the range 66-133.°C gives the yellowish dihydrate, NiCl2·2H2O. The hydrates convert to the anhydrous form upon heating in thionyl chloride or by heating under a stream of HCl gas. Heating the hydrates does not afford the anhydrous dichloride. NiCl 2 ⋅ 6 H 2 O + 6 SOCl 2 ⟶ NiCl 2 + 6 SO 2 + 12 HCl The dehydration is accompanied by a color change from green to yellow.
In case one needs a pure compound without presence of cobalt, nickel chloride can be obtained cautiously heating hexammine nickel chloride: Cl 2 hexammine nickel chloride → 175 − 200 ∘ C NiCl 2 + 6 NH 3 NiCl2 adopts the CdCl2 structure. In this motif, each Ni2+ center is coordinated to six Cl− centers, each chloride is bonded to three Ni centers. In NiCl2 the Ni-Cl bonds have "ionic character". Yellow NiBr2 and black NiI2 adopt similar structures, but with a different packing of the halides, adopting the CdI2 motif. In contrast, NiCl2·6H2O consists of separated trans- molecules linked more weakly to adjacent water molecules. Only four of the six water molecules in the formula is bound to the nickel, the remaining two are water of crystallization. Cobalt chloride hexahydrate has a similar structure; the hexahydrate occurs in nature as the rare mineral nickelbischofite. The dihydrate NiCl2·2H2O adopts a structure intermediate between the hexahydrate and the anhydrous forms, it consists of infinite chains of NiCl2.
The trans sites on the octahedral centers occupied by aquo ligands. A tetrahydrate NiCl2·4H2O is known. Nickel chloride solutions are acidic, with a pH of around 4 due to the hydrolysis of the Ni2+ ion. Most of the reactions ascribed to "nickel chloride" involve the hexahydrate, although specialized reactions require the anhydrous form. Reactions starting from NiCl2·6H2O can be used to form a variety of nickel coordination complexes because the H2O ligands are displaced by ammonia, thioethers and organophosphines. In some derivatives, the chloride remains within the coordination sphere, whereas chloride is displaced with basic ligands. Illustrative complexes include: Some nickel chloride complexes exist as an equilibrium mixture of two geometries. For example, NiCl22, containing four-coordinate Ni, exists in solution as a mixture of both the diamagnetic square planar and the paramagnetic tetrahedral isomers. Square planar complexes of nickel can form five-coordinate adducts. NiCl2 is the precursor to acetylacetonate complexes Ni22 and the benzene-soluble 3, a precursor to Ni2, an important reagent in organonickel chemistry.
In the presence of water scavengers, hydrated nickel chloride reacts with dimethoxyethane to form the molecular complex NiCl22. The dme ligands in this complex are labile. For example, this complex reacts with sodium cyclopentadienide to give the sandwich compound nickelocene. Hexammine nickel chloride complex is soluble when respective cobalt complex is not, which allows for easy separating of these close-related metals in laboratory conditions. NiCl2 and its hydrate are useful in organic synthesis; as a mild Lewis acid, e.g. for the regioselective isomerization of dienols:In combination with CrCl2 for the coupling of an aldehyde and a vinylic iodide to give allylic alcohols. For selective reductions in the presence of LiAlH4, e.g. for the conversion of alkenes to alkanes. As a precursor to nickel boride, prepared in situ from NiCl2 and NaBH4; this reagent behaves like Raney Nickel, comprising an efficient system for hydrogenation of unsaturated carbonyl compounds. As a precursor to finely divided Ni by reduction with Zn, for the reduction of aldehydes and nitro aromatic compounds.
This reagent promotes homo-coupling reactions, 2RX → R-R where R = aryl, vinyl. As a catalyst for making dialkyl arylphosphonates from phosphites and aryl iodide, ArI:ArI + P3 → ArP2 + EtINiCl2-dme is used due to its increased solubility in comparis
Vaska's complex is the trivial name for the chemical compound trans-carbonylchlorobisiridium, which has the formula IrCl2. This square planar diamagnetic organometallic complex consists of a central iridium atom bound to two mutually trans triphenylphosphine ligands, carbon monoxide, a chloride ion; the complex was first reported by J. W. DiLuzio and Lauri Vaska in 1961. Vaska's complex can undergo oxidative addition and is notable for its ability to bind to O2 reversibly, it is a bright yellow crystalline solid. The synthesis involves heating any iridium chloride salt with triphenylphosphine and a carbon monoxide source; the most popular method uses dimethylformamide as a solvent, sometimes aniline is added to accelerate the reaction. Another popular solvent is 2-methoxyethanol; the reaction is conducted under nitrogen. In the synthesis, triphenylphosphine serves as both a ligand and a reductant, the carbonyl ligand is derived by decomposition of dimethylformamide via a deinsertion of an intermediate Ir-CH species.
The following is a possible balanced equation for this complicated reaction. IrCl33 + 3 P3 + HCON2 + C6H5NH2 → IrCl2 + Cl + OP3 + Cl + 2 H2OTypical sources of iridium used in this preparation are IrCl3·xH2O and H2IrCl6. Studies on Vaska's complex helped provide the conceptual framework for homogeneous catalysis. Vaska's complex, with 16 valence electrons, is considered "coordinatively unsaturated" and can thus bind to one two-electron or two one-electron ligands to become electronically saturated with 18 valence electrons; the addition of two one-electron ligands is called oxidative addition. Upon oxidative addition, the oxidation state of the iridium increases from Ir to Ir; the four-coordinated square planar arrangement in the starting complex converts to an octahedral, six-coordinate product. Vaska's complex undergoes oxidative addition with conventional oxidants such as halogens, strong acids such as HCl, other molecules known to react as electrophiles, such as iodomethane. Vaska's complex binds O2 reversibly: IrCl2 + O2 ⇌ IrCl2O2The dioxygen ligand is bonded to Ir by both oxygen atoms, so-called side-on bonding.
In myoglobin and hemoglobin, by contrast, O2 binds end-on, attaching to the metal via only one of the two oxygen atoms. The resulting dioxygen adduct reverts to the parent complex upon heating or purging the solution with an inert gas, signaled by a colour change from orange back to yellow. Infrared spectroscopy can be used to analyse the products of oxidative addition to Vaska's complex because the reactions induce characteristic shifts of the stretching frequency of the coordinated carbon monoxide; these shifts are dependent on the amount of π-back bonding allowed by the newly associated ligands. The CO stretching frequencies for Vaska's complex and oxidatively added ligands have been documented in the literature. Vaska's complex: 1967 cm−1 Vaska's complex + O2: 2015 cm−1 Vaska's complex + MeI: 2047 cm−1 Vaska's complex + I2: 2067 cm−1Oxidative addition to give Ir products reduces the π-bonding from Ir to C, which causes the increase in the frequency of the carbonyl stretching band; the stretching frequency change depends upon the ligands that have been added, but the frequency is always greater than 2000 cm−1 for an Ir complex
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
Hapticity is the coordination of a ligand to a metal center via an uninterrupted and contiguous series of atoms. The hapticity of a ligand is described with the Greek letter η. For example, η2 describes a ligand. In general the η-notation only applies. In addition, if the ligand coordinates through multiple atoms that are not contiguous this is considered denticity, the κ-notation is used once again; when naming complexes care should be taken not to confuse η with μ, which relates to bridging ligands. The need for additional nomenclature for organometallic compounds became apparent in the mid-1950s when Dunitz and Rich described the structure of the "sandwich complex" ferrocene by X-ray crystallography where an iron atom is "sandwiched" between two parallel cyclopentadienyl rings. Cotton proposed the term hapticity derived from the adjectival prefix hapto placed before the name of the olefin, where the Greek letter η is used to denote the number of contiguous atoms of a ligand that bind to a metal center.
The term is employed to refer to ligands containing extended π-systems or where agostic bonding is not obvious from the formula. Ferrocene - bisiron Uranocene - bisuranium W32 - the first compound to be synthesized with a dihydrogen ligand. IrCl2 - the dioxygen derivative which forms reversibly upon oxygenation of Vaska's complex; the η-notation is encountered in many co-ordination compounds: Side-on bonding of molecules containing σ-bonds like H2: W32 Side-on bonded ligands containing multiple bonded atoms, e.g. ethylene in Zeise's salt or with fullerene, bonded through donation of the π-bonding electrons: K. H2O Related complexes containing bridging π-ligands: Co26 and 2 Dioxygen in bis,Note that with some bridging ligands, an alternative bridging mode is observed, e.g. κ1,κ1, like in 3VV3 contains a bridging dinitrogen molecule, where the molecule is end-on coordinated to the two metal centers. The bonding of π-bonded species can be extended over several atoms, e.g. in allyl, butadiene ligands, but in cyclopentadienyl or benzene rings can share their electrons.
Apparent violations of the 18-electron rule sometimes are explicable in compounds with unusual hapticities: The 18-VE complex Fe2 contains one η5 bonded cyclopentadienyl, one η1 bonded cyclopentadienyl. Reduction of the 18-VE compound 2+, results in another 18VE compound:. Examples of polyhapto coordinated heterocyclic and inorganic rings: Cr3 contains the sulfur heterocycle thiophene and Cr3 contains a coordinated inorganic ring; the hapticity of a ligand can change in the course of a reaction. E.g. in a redox reaction: Here one of the η6-benzene rings changes to a η4-benzene. Hapticity can change during a substitution reaction: Here the η5-cyclopentadienyl changes to an η3-cyclopentadienyl, giving room on the metal for an extra 2-electron donating ligand'L'. Removal of one molecule of CO and again donation of two more electrons by the cyclopentadienyl ligand restores the η5-cyclopentadienyl; the so-called indenyl effect describes changes in hapticity in a substitution reaction. Hapticity must be distinguished from denticity.
Polydentate ligands coordinate via multiple coordination sites within the ligand. In this case the coordinating atoms are identified using the κ-notation, as for example seen in coordination of 1,2-bisethane, to NiCl2 as dichloronickel. If the coordinating atoms are contiguous, the η-notation is used, as e.g. in titanocene dichloride: dichlorobistitanium. Molecules with polyhapto ligands are fluxional known as stereochemically non-rigid. Two classes of fluxionality are prevalent for organometallic complexes of polyhapto ligands: Case 1, typically: when the hapticity value is less than the number of sp2 carbon atoms. In such situations, the metal will migrate from carbon to carbon, maintaining the same net hapticity; the η1-C5H5 ligand in Fe2 rearranges in solution such that Fe binds alternatingly to each carbon atom in the η1-C5H5 ligand. This reaction is degenerate and, in the jargon of organic chemistry, it is an example of a sigmatropic rearrangement. A related example is Bisiron. Case 2, typically: complexes containing cyclic polyhapto ligands with maximized hapticity.
Such ligands tend to rotate. A famous example is ferrocene, Fe2, wherein the Cp rings rotate with a low energy barrier about the principal axis of the molecule that "skewers" each ring; this "ring whizzing" explains, inter alia, why only one isomer can be isolated for Fe2. In this case, the rotamers are not degenerate, but the rotational barriers have low energies of activation
A cyclopentadienyl complex is a metal complex with one or more cyclopentadienyl groups. Cyclopentadienyl ligands invariable bind to metals as pentahapto bonding mode; the metal–cyclopentadienyl interaction is drawn as a single line from the metal center to the center of the Cp ring. Biscyclopentadienyl complexes are called metallocenes. A famous example of this type of complex is ferrocene, which has many analogues for other metals, such as chromocene and nickelocene; when the Cp rings are mutually parallel the compound is known as a sandwich complex. This area of organometallic chemistry was first developed in the 1950s. Bent metallocenes are represented by compounds of the type; some are catalysts for ethylene polymerization. Metallocenes are thermally stable, find use as catalysts in various types of reactions. Mixed-ligand Cp complexes containing Cp ligand and one or more other ligands, they are more numerous. One studied example is the Fp dimer. Monometallic compounds featuring only one Cp ring are known as half sandwich compounds or as piano stool compounds, one example being cyclopentadienylmanganese tricarbonyl.
All 5 carbon atoms of a Cp ligand are bound to the metal in the vast majority of M–Cp complexes. This bonding mode is called η5-coordination; the M–Cp bonding arises from overlap of the five π molecular orbitals of the Cp ligand with the s, p, d orbitals on the metal. This π bonding is significant, hence these complexes are referred to as π-complexes. All of the transition metals, that is, group 4 to 10 metals, employ this coordination mode. In rare cases, Cp binds to metals via only one carbon center; these types of interactions are described as σ-complexes because they only have a σ bond between the metal and the cyclopentadienyl group. Typical examples of this type of complex are group 14 metal complexes such as CpSiMe3. An example of both is.. It is probable. Still rarer, the Cp unit can bond to the metal via a three-carbons. In these η3-Cp complexes, the bonding resembles that in allyl ligands; such complexes, sometimes called "slipped Cp complexes", are invoked as intermediates in ring slipping reactions.
The compounds are prepared by salt metathesis reactions of alkali-metal cyclopentadienyl compounds with transition metal chlorides. Sodium cyclopentadienide and lithium cyclopentadienide are used. Trimethylsilylcyclopentadiene cyclopentadienylthallium are alternative sources. For the preparation of some robust complexes, e.g. nickelocene, cyclopentadiene is employed in the presence of a conventional base such as KOH. When only a single Cp ligand is installed, the other ligands carbonyl, halogen and hydride. Most Cp complexes are prepared by substitution of preformed Cp complexes by replacement of halide, CO, other simple ligands. Variations of Cyclopentadienyl complexes A pair of cyclopentadienyl ligands can be covalently linked giving rise to so-call ansa metallocenes; the angle between the two Cp rings is fixed. Rotation of the rings about the metal-centroid axis is stopped as well. A related class of derivatives give rise to the constrained geometry complexes. In these cases, a Cp ligand; such complexes have been commercialized for the production of polypropylene.
Pentamethylcyclopentadiene gives rise to pentamethylcyclopentadienyl complexes. These ligands are more basic, which results in distinctive properties. Constrained geometry complexes are related to ansa-metallocenes except that one ligand is not Cp-related. Cp metal complexes are used as stoichiometric reagents in chemical research. Ferrocenium reagents are oxidants. Cobaltocene is a soluble reductant. Derivatives of Cp2TiCl2 and Cp2ZrCl2 are the basis of some reagents in organic synthesis. Upon treatment with aluminoxane, these dihalides give catalysts for olefin polymerization; such species are called Kaminsky-type catalysts. Yamamoto, A.. Organotransition Metal Chemistry: Fundamental Concepts and Applications. New York, NY: Wiley-Interscience. P. 105. Shriver, D.. Inorganic Chemistry. New York, NY: W. H. Freeman. King, R. B.. B.. "Organometallic chemistry of the transition metals XXI. Some π-pentamethylcyclopentadienyl derivatives of various transition metals". J. Organomet. Chem. 8: 287–297. Doi:10.1016/S0022-328X91042-8
Samarium is a chemical element with symbol Sm and atomic number 62. It is a moderately hard silvery metal that oxidizes in air. Being a typical member of the lanthanide series, samarium assumes the oxidation state +3. Compounds of samarium are known, most notably the monoxide SmO, monochalcogenides SmS, SmSe and SmTe, as well as samarium iodide; the last compound is a common reducing agent in chemical synthesis. Samarium has no significant biological role but is only toxic. Samarium was discovered in 1879 by the French chemist Paul-Émile Lecoq de Boisbaudran and named after the mineral samarskite from which it was isolated; the mineral itself was earlier named after a Russian mine official, Colonel Vassili Samarsky-Bykhovets, who thereby became the first person to have a chemical element named after him, albeit indirectly. Although classified as a rare-earth element, samarium is the 40th most abundant element in the Earth's crust and is more common than metals such as tin. Samarium occurs with concentration up to 2.8% in several minerals including cerite, samarskite and bastnäsite, the last two being the most common commercial sources of the element.
These minerals are found in China, the United States, India, Sri Lanka and Australia. The major commercial application of samarium is in samarium–cobalt magnets, which have permanent magnetization second only to neodymium magnets; the radioactive isotope samarium-153 is the active component of the drug samarium lexidronam, which kills cancer cells in the treatment of lung cancer, prostate cancer, breast cancer and osteosarcoma. Another isotope, samarium-149, is a strong neutron absorber and is therefore added to the control rods of nuclear reactors, it is formed as a decay product during the reactor operation and is one of the important factors considered in the reactor design and operation. Other applications of samarium include catalysis of chemical reactions, radioactive dating and X-ray lasers. Samarium is a rare earth metal having a density similar to those of zinc. With the boiling point of 1794 °C, samarium is the third most volatile lanthanide after ytterbium and europium. At ambient conditions, samarium assumes a rhombohedral structure.
Upon heating to 731 °C, its crystal symmetry changes into hexagonally close-packed, however the transition temperature depends on the metal purity. Further heating to 922 °C transforms the metal into a body-centered cubic phase. Heating to 300 °C combined with compression to 40 kbar results in a double-hexagonally close-packed structure. Applying higher pressure of the order of hundreds or thousands of kilobars induces a series of phase transformations, in particular with a tetragonal phase appearing at about 900 kbar. In one study, the dhcp phase could be produced without compression, using a nonequilibrium annealing regime with a rapid temperature change between about 400 and 700 °C, confirming the transient character of this samarium phase. Thin films of samarium obtained by vapor deposition may contain the hcp or dhcp phases at ambient conditions. Samarium are paramagnetic at room temperature, their corresponding effective magnetic moments, below 2µB, are the 3rd lowest among the lanthanides after lanthanum and lutetium.
The metal transforms to an antiferromagnetic state upon cooling to 14.8 K. Individual samarium atoms can be isolated by encapsulating them into fullerene molecules, they can be doped between the C60 molecules in the fullerene solid, rendering it superconductive at temperatures below 8 K. Samarium doping of iron-based superconductors – the most recent class of high-temperature superconductors – allows enhancing their transition temperature to 56 K, the highest value achieved so far in this series. Freshly prepared samarium has a silvery luster. In air, it oxidizes at room temperature and spontaneously ignites at 150 °C; when stored under mineral oil, samarium oxidizes and develops a grayish-yellow powder of the oxide-hydroxide mixture at the surface. The metallic appearance of a sample can be preserved by sealing it under an inert gas such as argon. Samarium is quite electropositive and reacts with cold water and quite with hot water to form samarium hydroxide: 2 Sm + 6 H2O → 2 Sm3 + 3 H2 Samarium dissolves in dilute sulfuric acid to form solutions containing the yellow to pale green Sm ions, which exist as 3+ complexes: 2 Sm + 3 H2SO4 → 2 Sm3+ + 3 SO2−4 + 3 H2 Samarium is one of the few lanthanides that exhibit the oxidation state +2.
The Sm2+ ions are blood-red in aqueous solution. The most stable oxide of samarium is the sesquioxide Sm2O3; as many other samarium compounds, it exists in several crystalline phases. The trigonal form is obtained by slow cooling from the melt; the melting point of Sm2O3 is rather high and therefore melting is achieved not by direct heating, but with induction heating, through a radio-frequency coil. The Sm2O3 crystals of monoclinic symmetry can be grown by the flame fusion method from the Sm2O3 powder, that yields cylindrical boules up to several centimeters long and about one centimeter in diameter; the boules are transparent when are orange otherwise. Heating the metastable trigonal Sm2O3 to 1900 °C converts it to the more stable monoclinic phase. Cubic Sm2O3 has a