Lanthanide

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The lanthanide (/ˈlænθənaɪd/) or lanthanoid (/ˈlænθənɔɪd/) series of chemical elements comprises the 15 metallic chemical elements with atomic numbers 57–71, from lanthanum through lutetium. These elements, along with the chemically similar elements scandium and yttrium, are often collectively known as the rare-earth elements.

The informal chemical symbol Ln is used in general discussions of lanthanide chemistry to refer to any lanthanide. All but one of the lanthanides are f-block elements, corresponding to the filling of the 4f electron shell. There is some dispute on whether lanthanum or lutetium is a d-block element, but lutetium is usually considered so by those who study the matter; it is included due to its chemical similarities with the other 14. All lanthanide elements form trivalent cations, Ln³⁺, whose chemistry is largely determined by the ionic radius, which decreases steadily from lanthanum to lutetium.

They are called lanthanides because the elements in the series are chemically similar to lanthanum. Since "lanthanide" means "like lanthanum", it has been argued that lanthanum cannot logically be a lanthanide, but IUPAC acknowledges its inclusion based on common usage.

In presentations of the periodic table, the f-block elements are customarily shown as two additional rows below the main body of the table, This convention is entirely a matter of aesthetics and formatting practicality; a rarely used wide-formatted periodic table inserts the 4f and 5f series in their proper places, as parts of the table's sixth and seventh rows (periods).

The 1985 International Union of Pure and Applied Chemistry "Red Book" (p. 45) recommends that "lanthanoid" is used rather than "lanthanide". The ending "-ide" normally indicates a negative ion. However, owing to wide current use, "lanthanide" is still allowed.

Lanthanides

Lan­thanum57La​138.91

Cerium58Ce​140.12

Praseo­dymium59Pr​140.91

Neo­dymium60Nd​144.24

Prome­thium61Pm​[145]

Sama­rium62Sm​150.36

Europ­ium63Eu​151.96

Gadolin­ium64Gd​157.25

Ter­bium65Tb​158.93

Dyspro­sium66Dy​162.50

Hol­mium67Ho​164.93

Erbium68Er​167.26

Thulium69Tm​168.93

Ytter­bium70Yb​173.05

Lute­tium71Lu​174.97

Primordial  From decay  Synthetic Border shows natural occurrence of the element

Etymology

Together with the two elements at the top of group 3, scandium and yttrium, the trivial name "rare earths" is sometimes used to describe all the lanthanides; a definition of rare earths including the group 3, lanthanide, and actinide elements is also occasionally seen, and rarely Sc + Y + lanthanides + thorium. The "earth" in the name "rare earths" arises from the minerals from which they were isolated, which were uncommon oxide-type minerals. However, these elements are neither rare in abundance nor "earths" (an obsolete term for water-insoluble strongly basic oxides of electropositive metals incapable of being smelted into metal using late 18th century technology). Group 2 is known as the alkaline earth elements for much the same reason.

The "rare" in the "rare earths" name has much more to do with the difficulty of separating out each of the individual lanthanide elements than scarcity of any of them. By way of the Greek "dysprositos" for "hard to get at," element 66, dysprosium was similarly named; lanthanum itself is named after a word for "hidden." The elements 57 (La) to 71 (Lu) are very similar chemically to one another and frequently occur together in nature, often anywhere from three to all 15 of the lanthanides (along with yttrium as a 16th) occur in minerals such as samarskite, monazite and many others which can also contain the other two group 3 elements as well as thorium and occasionally other actinides as well. A majority of the rare earths were discovered at the same mine in Ytterby, Sweden and four of them are named (yttrium, ytterbium, erbium, terbium) after the city and a fifth *(holmium) after Stockholm; scandium is named after Scandinavia, thulium after the old name Thule, and the immediately-following group 4 element (number 72) hafnium is named for the Latin name of the city of Copenhagen.

Samarskite (a mineral which is the source of the name of the element samarium) and other similar minerals in particular also have these elements in association with the nearby metals tantalum, niobium, hafnium, zirconium, vanadium, and titanium, from group 4 and group 5 often in similar oxidation states. Monazite is a phosphate of numerous group 3 + lanthanide + actinide metals and mined especially for the thorium content and specific rare earths especially lanthanum, yttrium and cerium. Cerium and lanthanum as well as other members of the rare earth series are often produced as a metal called mischmetal containing a variable mixture of these elements with cerium and lanthanum predominating; it has direct uses such as lighter flints and other spark sources which do not require extensive purification of one of these metals.

There are also rare earth-bearing minerals based on group 2 elements such as yttrocalcite, yttrocerite, yttrofluorite which vary in content of yttrium, cerium, and lanthanum in a particular as well as varying amounts of the others. Other lanthanide/rare earth minerals include bastnäsite, florencite, chernovite, perovskite, xenotime, cerite, gadolinite, lanthanite, fergusonite, polycrase, blomstrandine, håleniusite, miserite, loparite, lepersonnite, euxenite, all of which have a range of relative element concentration and may have the symbol of a predominating one such as monazite-ce; group 3 elements do not occur as native element minerals in the fashion of gold, silver, tantalum and many others on earth but may in lunar regolith. Very rare cerium, lanthanum, and presumably other lanthanide/group 3 halides, feldspars and garnets are also known to exist.

All of this is the result of the order in which the electron shells of these elements are filled—the outermost has the same configuration for all of them, and a deeper shell is progressively filled with electrons as the atomic number increases from 57 towards 71. For many years, mixtures of more than one rare earth were considered to be single elements, such as neodymium and praseodymium being thought to be the single element didymium and so on. Very small differences in solubility are used in solvent and ion-exchange purification methods for these elements which require a great deal of repeating to get a purified metal. The refined metals and their compounds have subtle and stark differences amongst themselves in electronic, electrical, optical, and magnetic properties which account for their many niche uses.

By way of examples of the term meaning the above considerations rather than their scarcity, cerium is the 26th most abundant element in the Earth's crust and more abundant than copper, neodymium is more abundant than gold; thulium (the second least common naturally occurring lanthanide) is more abundant than iodine, which is itself common enough for biology to have evolved critical usages thereof, and even the lone radioactive element in the series, promethium, is more common than the two rarest naturally occurring elements, francium and astatine, combined. Despite their abundance, even the technical term "lanthanides" could be interpreted to reflect a sense of elusiveness on the part of these elements, as it comes from the Greek λανθανειν (lanthanein), "to lie hidden". However, if not referring to their natural abundance, but rather to their property of "hiding" behind each other in minerals, this interpretation is in fact appropriate. The etymology of the term must be sought in the first discovery of lanthanum, at that time a so-called new rare-earth element "lying hidden" in a cerium mineral, and it is an irony that lanthanum was later identified as the first in an entire series of chemically similar elements and could give name to the whole series. The term "lanthanide" was introduced by Victor Goldschmidt in 1925.

Physical properties of the elements

Chemical element	La	Ce	Pr	Nd	Pm	Sm	Eu	Gd	Tb	Dy	Ho	Er	Tm	Yb	Lu
Atomic number	57	58	59	60	61	62	63	64	65	66	67	68	69	70	71
Image
Density (g/cm3)	6.162	6.770	6.77	7.01	7.26	7.52	5.244	7.90	8.23	8.540	8.79	9.066	9.32	6.90	9.841
Melting point (°C)	920	795	935	1024	1042	1072	826	1312	1356	1407	1461	1529	1545	824	1652
Boiling point (°C)	3464	3443	3520	3074	3000	1794	1529	3273	3230	2567	2720	2868	1950	1196	3402
Atomic electron configuration (gas phase)*	5d1	4f15d1	4f3	4f4	4f5	4f6	4f7	4f75d1	4f9	4f10	4f11	4f12	4f13	4f14	4f145d1
Metal lattice (RT)	dhcp	fcc	dhcp	dhcp	dhcp	**	bcc	hcp	hcp	hcp	hcp	hcp	hcp	fcc	hcp
Metallic radius (pm)	162	181.8	182.4	181.4	183.4	180.4	208.4	180.4	177.3	178.1	176.2	176.1	175.9	193.3	173.8
Resistivity at 25 °C (μΩ·cm)	57–80 20 °C	73	68	64		88	90	134	114	57	87	87	79	29	79
Magnetic susceptibility χmol /10−6(cm3·mol−1)	+95.9	+2500 (β)	+5530(α)	+5930 (α)		+1278(α)	+30900	+185000 (350 K)	+170000 (α)	+98000	+72900	+48000	+24700	+67 (β)	+183

* Between initial Xe and final 6s² electronic shells

** Sm has a close packed structure like most of the lanthanides but has an unusual 9 layer repeat

Gschneider and Daane (1988) attribute the trend in melting point which increases across the series, (lanthanum (920 °C) – lutetium (1622 °C)) to the extent of hybridization of the 6s, 5d, and 4f orbitals. The hybridization is believed to be at its greatest for cerium, which has the lowest melting point of all, 795 °C. The lanthanide metals are soft; their hardness increases across the series. Europium stands out, as it has the lowest density in the series at 5.24 g/cm³ and the largest metallic radius in the series at 208.4 pm. It can be compared to barium, which has a metallic radius of 222 pm. It is believed that the metal contains the larger Eu²⁺ ion and that there are only two electrons in the conduction band. Ytterbium also has a large metallic radius, and a similar explanation is suggested. The resistivities of the lanthanide metals are relatively high, ranging from 29 to 134 μΩ·cm. These values can be compared to a good conductor such as aluminium, which has a resistivity of 2.655 μΩ·cm. With the exceptions of La, Yb, and Lu (which have no unpaired f electrons), the lanthanides are strongly paramagnetic, and this is reflected in their magnetic susceptibilities. Gadolinium becomes ferromagnetic at below 16 °C (Curie point). The other heavier lanthanides – terbium, dysprosium, holmium, erbium, thulium, and ytterbium – become ferromagnetic at much lower temperatures.

Chemistry and compounds

Chemical element	La	Ce	Pr	Nd	Pm	Sm	Eu	Gd	Tb	Dy	Ho	Er	Tm	Yb	Lu
Atomic number	57	58	59	60	61	62	63	64	65	66	67	68	69	70	71
Ln3+ electron configuration*	4f0	4f1	4f2	4f3	4f4	4f5	4f6	4f7	4f8	4f9	4f10	4f11	4f12	4f13	4f14
Ln3+ radius (pm)	103	102	99	98.3	97	95.8	94.7	93.8	92.3	91.2	90.1	89	88	86.8	86.1
Ln4+ ion color in aqueous solution	—	Orange-yellow	Yellow	Blue-violet	—	—	—	—	Red-brown	Orange-yellow	—	—	—	—	—
Ln3+ ion color in aqueous solution	Colorless	Colorless	Green	Violet	Pink	Pale yellow	Colorless	Colorless	V. pale pink	Pale yellow	Yellow	Rose	Pale green	Colorless	Colorless
Ln2+ ion color in aqueous solution[]	—	—	—	—	—	Blood red	Colorless	—	—	—	—	—	Violet-red	Yellow-green	—

* Not including initial [Xe] core

The colors of lanthanide complexes originate almost entirely from charge transfer interactions between the metal and the ligand. f → f transitions are symmetry forbidden (or Laporte-forbidden), which is also true of transition metals. However, transition metals are able to use vibronic coupling to break this rule. The valence orbitals in lanthanides are almost entirely non-bonding and as such little effective vibronic coupling takes, hence the spectra from f → f transitions are much weaker and narrower than those from d → d transitions. In general this makes the colors of lanthanide complexes far fainter than those of transition metal complexes. f → f transitions are not possible for the f¹ and f¹³ configurations of Ce³⁺ and Yb³⁺ and thus these ions are colorless in aqueous solution.

Approximate colors of lanthanide ions in aqueous solution

Oxidation state

57

58

59

60

61

62

63

64

65

66

67

68

69

70

71

+2

Sm2+

Eu2+

Tm2+

Yb2+

+3

La3+

Ce3+

Pr3+

Nd3+

Pm3+

Sm3+

Eu3+

Gd3+

Tb3+

Dy3+

Ho3+

Er3+

Tm3+

Yb3+

Lu3+

+4

Ce4+

Pr4+

Nd4+

Tb4+

Dy4+

Effect of 4f orbitals

Going across the lanthanides in the periodic table, the 4f orbitals are usually being filled. The effect of the 4f orbitals on the chemistry of the lanthanides is profound and is the factor that distinguishes them from the transition metals. There are seven 4f orbitals, and there are two different ways in which they are depicted: as a "cubic set" or as a general set. The cubic set is f_z³, f_xz², f_yz², f_xyz, f_z(x²−y²), f_{x(x²−3y²)} and f_{y(3x²−y²)}. The 4f orbitals penetrate the [Xe] core and are isolated, and thus they do not participate in bonding. This explains why crystal field effects are small and why they do not form π bonds. As there are seven 4f orbitals, the number of unpaired electrons can be as high as 7, which gives rise to the large magnetic moments observed for lanthanide compounds.

Measuring the magnetic moment can be used to investigate the 4f electron configuration, and this is a useful tool in providing an insight into the chemical bonding. The lanthanide contraction, i.e. the reduction in size of the Ln³⁺ ion from La³⁺ (103 pm) to Lu³⁺ (86.1 pm), is often explained by the poor shielding of the 5s and 5p electrons by the 4f electrons.

Lanthanide oxides: clockwise from top center: praseodymium, cerium, lanthanum, neodymium, samarium and gadolinium.

The electronic structure of the lanthanide elements, with minor exceptions, is [Xe]6s²4fⁿ. The chemistry of the lanthanides is dominated by the +3 oxidation state, and in Ln^III compounds the 6s electrons and (usually) one 4f electron are lost and the ions have the configuration [Xe]4f^m. All the lanthanide elements exhibit the oxidation state +3. In addition, Ce³⁺ can lose its single f electron to form Ce⁴⁺ with the stable electronic configuration of xenon. Also, Eu³⁺ can gain an electron to form Eu²⁺ with the f⁷ configuration that has the extra stability of a half-filled shell. Other than Ce(IV) and Eu(II), none of the lanthanides are stable in oxidation states other than +3 in aqueous solution. Promethium is effectively a man-made element, as all its isotopes are radioactive with half-lives shorter than 20 years.

In terms of reduction potentials, the Ln^0/3+ couples are nearly the same for all lanthanides, ranging from −1.99 (for Eu) to −2.35 V (for Pr). Thus these metals are highly reducing, with reducing power similar to alkaline earth metals such as Mg (−2.36 V).

Lanthanide oxidation states

Ionization energies and reduction potentials of the elements

The ionization energies for the lanthanides can be compared with aluminium. In aluminium the sum of the first three ionization energies is 5139 kJ·mol⁻¹, whereas the lanthanides fall in the range 3455 – 4186 kJ·mol⁻¹. This correlates with the highly reactive nature of the lanthanides.

The sum of the first two ionization energies for europium, 1632 kJ·mol⁻¹ can be compared with that of barium 1468.1 kJ·mol⁻¹ and europium's third ionization energy is the highest of the lanthanides. The sum of the first two ionization energies for ytterbium are the second lowest in the series and its third ionization energy is the second highest. The high third ionization energy for Eu and Yb correlate with the half filling 4f⁷ and complete filling 4f¹⁴ of the 4f subshell, and the stability afforded by such configurations due to exchange energy. Europium and ytterbium form salt like compounds with Eu²⁺ and Yb²⁺, for example the salt like dihydrides. Both europium and ytterbium dissolve in liquid ammonia forming solutions of Ln²⁺(NH₃)_x again demonstrating their similarities to the alkaline earth metals.

The relative ease with which the 4th electron can be removed in cerium and (to a lesser extent praseodymium) indicates why Ce(IV) and Pr(IV) compounds can be formed, for example CeO₂ is formed rather than Ce₂O₃ when cerium reacts with oxygen.

Separation of lanthanides

The similarity in ionic radius between adjacent lanthanide elements makes it difficult to separate them from each other in naturally occurring ores and other mixtures. Historically, the very laborious processes of cascading and fractional crystallization were used. Because the lanthanide ions have slightly different radii, the lattice energy of their salts and hydration energies of the ions will be slightly different, leading to a small difference in solubility. Salts of the formula Ln(NO₃)₃·2NH₄NO₃·4H₂O can be used. Industrially, the elements are separated from each other by solvent extraction. Typically an aqueous solution of nitrates is extracted into kerosene containing tri-n-butylphosphate. The strength of the complexes formed increases as the ionic radius decreases, so solubility in the organic phase increases. Complete separation can be achieved continuously by use of countercurrent exchange methods. The elements can also be separated by ion-exchange chromatography, making use of the fact that the stability constant for formation of EDTA complexes increases for log K ≈ 15.5 for [La(EDTA)]⁻ to log K ≈ 19.8 for [Lu(EDTA)]⁻.

Coordination chemistry and catalysis

When in the form of coordination complexes, lanthanides exist overwhelmingly in their +3 oxidation state, although particularly stable 4f configurations can also give +4 (Ce, Tb) or +2 (Eu, Yb) ions. All of these forms are strongly electropositive and thus lanthanide ions are hard Lewis acids. The oxidation states are also very stable; with the exceptions of SmI₂ and cerium(IV) salts, lanthanides are not used for redox chemistry. 4f electrons have a high probability of being found close to the nucleus and are thus strongly affected as the nuclear charge increases across the series; this results in a corresponding decrease in ionic radii referred to as the lanthanide contraction.

The low probability of the 4f electrons existing at the outer region of the atom or ion permits little effective overlap between the orbitals of a lanthanide ion and any binding ligand. Thus lanthanide complexes typically have little or no covalent character and are not influenced by orbital geometries. The lack of orbital interaction also means that varying the metal typically has little effect on the complex (other than size), especially when compared to transition metals. Complexes are held together by weaker electrostatic forces which are omni-directional and thus the ligands alone dictate the symmetry and coordination of complexes. Steric factors therefore dominate, with coordinative saturation of the metal being balanced against inter-ligand repulsion. This results in a diverse range of coordination geometries, many of which are irregular, and also manifests itself in the highly fluxional nature of the complexes. As there is no energetic reason to be locked into a single geometry, rapid intramolecular and intermolecular ligand exchange will take place. This typically results in complexes that rapidly fluctuate between all possible configurations.

Many of these features make lanthanide complexes effective catalysts. Hard Lewis acids are able to polarise bonds upon coordination and thus alter the electrophilicity of compounds, with a classic example being the Luche reduction. The large size of the ions coupled with their labile ionic bonding allows even bulky coordinating species to bind and dissociate rapidly, resulting in very high turnover rates; thus excellent yields can often be achieved with loadings of only a few mol%. The lack of orbital interactions combined with the lanthanide contraction means that the lanthanides change in size across the series but that their chemistry remains much the same. This allows for easy tuning of the steric environments and examples exist where this has been used to improve the catalytic activity of the complex and change the nuclearity of metal clusters.

Despite this, the use of lanthanide coordination complexes as homogeneous catalysts is largely restricted to the laboratory and there are currently few examples them being used on an industrial scale. Lanthanides exist in many forms other that coordination complexes and many of these are industrially useful. In particular lanthanide metal oxides are used as heterogeneous catalysts in various industrial processes.

Ln(III) compounds

The trivalent lanthanides mostly form ionic salts. The trivalent ions are hard acceptors and form more stable complexes with oxygen-donor ligands than with nitrogen-donor ligands. The larger ions are 9-coordinate in aqueous solution, [Ln(H₂O)₉]³⁺ but the smaller ions are 8-coordinate, [Ln(H₂O)₈]³⁺. There is some evidence that the later lanthanides have more water molecules in the second coordination sphere. Complexation with monodentate ligands is generally weak because it is difficult to displace water molecules from the first coordination sphere. Stronger complexes are formed with chelating ligands because of the chelate effect, such as the tetra-anion derived from 1,4,7,10-tetraazacyclododecane-1,4,7,10-tetraacetic acid (DOTA).

Samples of lanthanide nitrates in their hexahydrate form. From left to right: La, Ce, Pr, Nd, Sm, Eu, Gd, Tb, Dy, Ho, Er, Tm, Yb, Lu.

Ln(II) and Ln(IV) compounds

The most common divalent derivatives of the lanthanides are for Eu(II), which achieves a favorable f⁷ configuration. Divalent halide derivatives are known for all of the lanthanides. They are either conventional salts or are Ln(III) "electride"-like salts. The simple salts include YbI₂, EuI₂, and SmI₂. The electride-like salts, described as Ln³⁺, 2I⁻, e⁻, include LaI₂, CeI₂ and GdI₂. Many of the iodides form soluble complexes with ethers, e.g. TmI₂(dimethoxyethane)₃. Samarium(II) iodide is a useful reducing agent. Ln(II) complexes can be synthesized by transmetalation reactions. The normal range of oxidation states can be expanded via the use of sterically bulky cyclopentadienyl ligands, in this way many lanthanides can be isolated as Ln(II) compounds.

Ce(IV) in ceric ammonium nitrate is a useful oxidizing agent. The Ce(IV) is the exception owing to the tendency to form an unfilled f shell. Otherwise tetravalent lanthanides are rare. However, recently Tb(IV) and Pr(IV) complexes have been shown to exist.

Hydrides

Lanthanide metals react exothermically with hydrogen to form LnH₂, dihydrides. With the exception of Eu and Yb which resemble the Ba and Ca hydrides (non-conducting, transparent salt like compounds) they form black pyrophoric, conducting compounds where the metal sub-lattice is face centred cubic and the H atoms occupy tetrahedral sites. Further hydrogenation produces a trihydride which is non-stoichiometric, non-conducting, more salt like. The formation of trihydride is associated with and increase in 8–10% volume and this is linked to greater localization of charge on the hydrogen atoms which become more anionic (H⁻ hydride anion) in character.

Halides

The only tetrahalides known are the tetrafluorides of cerium, praseodymium, terbium, neodymium and dysprosium, the last two known only under matrix isolation conditions. All of the lanthanides form trihalides with fluorine, chlorine, bromine and iodine. They are all high melting and predominantly ionic in nature. The fluorides are only slightly soluble in water and are not sensitive to air, and this contrasts with the other halides which are air sensitive, readily soluble in water and react at high temperature to form oxohalides.

The trihalides were important as pure metal can be prepared from them. In the gas phase the trihalides are planar or approximately planar, the lighter lanthanides have a lower % of dimers, the heavier lanthanides a higher proportion. The dimers have a similar structure to Al₂Cl₆.

Some of the dihalides are conducting while the rest are insulators. The conducting forms can be considered as Ln^III electride compounds where the electron is delocalised into a conduction band, Ln³⁺ (X⁻)₂(e⁻). All of the diodides have relatively short metal-metal separations. The CuTi₂ structure of the lanthanum, cerium and praseodymium diodides along with HP-NdI₂ contain 4⁴ nets of metal and iodine atoms with short metal-metal bonds (393-386 La-Pr). these compounds should be considered to be two-dimensional metals (two-dimensional in the same way that graphite is). The salt-like dihalides include those of Eu, Dy, Tm, and Yb. The formation of a relatively stable +2 oxidation state for Eu and Yb is usually explained by the stability (exchange energy) of half filled (f⁷) and fully filled f¹⁴. GdI₂ possesses the layered MoS₂ structure, is ferromagnetic and exhibits colossal magnetoresistance

The sesquihalides Ln₂X₃ and the Ln₇I₁₂ compounds listed in the table contain metal clusters, discrete Ln₆I₁₂ clusters in Ln₇I₁₂ and condensed clusters forming chains in the sesquihalides. Scandium forms a similar cluster compound with chlorine, Sc₇Cl₁₂ Unlike many transition metal clusters these lanthanide clusters do not have strong metal-metal interactions and this is due to the low number of valence electrons involved, but instead are stabilised by the surrounding halogen atoms.

LaI is the only known monohalide. Prepared from the reaction of LaI₃ and La metal, it has a NiAs type structure and can be formulated La³⁺ (I⁻)(e⁻)₂.

Oxides and hydroxides

All of the lanthanides form sesquioxides, Ln₂O₃. The lighter/larger lanthanides adopt a hexagonal 7-coordinate structure while the heavier/smaller ones adopt a cubic 6-coordinate "C-M₂O₃" structure. All of the sesquioxides are basic, and absorb water and carbon dioxide from air to form carbonates, hydroxides and hydroxycarbonates. They dissolve in acids to form salts.

Cerium forms a stoichiometric dioxide, CeO₂, where cerium has an oxidation state of +4. CeO₂ is basic and dissolves with difficulty in acid to form Ce⁴⁺ solutions, from which Ce^IV salts can be isolated, for example the hydrated nitrate Ce(NO₃)₄.5H₂O. CeO₂ is used as an oxidation catalyst in catalytic converters. Praseodymium and terbium form non-stoichiometric oxides containing Ln^IV, although more extreme reaction conditions can produce stoichiometric (or near stoichiometric) PrO₂ and TbO₂.

Europium and ytterbium form salt-like monoxides, EuO and YbO, which have a rock salt structure. EuO is ferromagnetic at low temperatures, and is a semiconductor with possible applications in spintronics. A mixed Eu^II/Eu^III oxide Eu₃O₄ can be produced by reducing Eu₂O₃ in a stream of hydrogen. Neodymium and samarium also form monoxides, but these are shiny conducting solids, although the existence of samarium monoxide is considered dubious.

All of the lanthanides form hydroxides, Ln(OH)₃. With the exception of lutetium hydroxide, which has a cubic structure, they have the hexagonal UCl₃ structure. The hydroxides can be precipitated from solutions of Ln^III. They can also be formed by the reaction of the sesquioxide, Ln₂O₃, with water, but although this reaction is thermodynamically favorable it is kinetically slow for the heavier members of the series. Fajans' rules indicate that the smaller Ln³⁺ ions will be more polarizing and their salts correspondingly less ionic. The hydroxides of the heavier lanthanides become less basic, for example Yb(OH)₃ and Lu(OH)₃ are still basic hydroxides but will dissolve in hot concentrated NaOH.

Chalcogenides (S, Se, Te)

All of the lanthanides form Ln₂Q₃ (Q= S, Se, Te). The sesquisulfides can be produced by reaction of the elements or (with the exception of Eu₂S₃) sulfidizing the oxide (Ln₂O₃) with H₂S. The sesquisulfides, Ln₂S₃ generally lose sulfur when heated and can form a range of compositions between Ln₂S₃ and Ln₃S₄. The sesquisulfides are insulators but some of the Ln₃S₄ are metallic conductors (e.g. Ce₃S₄) formulated (Ln³⁺)₃ (S²⁻)₄ (e⁻), while others (e.g. Eu₃S₄ and Sm₃S₄) are semiconductors. Structurally the sesquisulfides adopt structures that vary according to the size of the Ln metal. The lighter and larger lanthanides favoring 7-coordinate metal atoms, the heaviest and smallest lanthanides (Yb and Lu) favoring 6 coordination and the rest structures with a mixture of 6 and 7 coordination.

Polymorphism is common amongst the sesquisulfides. The colors of the sesquisulfides vary metal to metal and depend on the polymorphic form. The colors of the γ-sesquisulfides are La₂S₃, white/yellow; Ce₂S₃, dark red; Pr₂S₃, green; Nd₂S₃, light green; Gd₂S₃, sand; Tb₂S₃, light yellow and Dy₂S₃, orange. The shade of γ-Ce₂S₃ can be varied by doping with Na or Ca with hues ranging from dark red to yellow, and Ce₂S₃ based pigments are used commercially and are seen as low toxicity substitutes for cadmium based pigments.

All of the lanthanides form monochalcogenides, LnQ, (Q= S, Se, Te). The majority of the monochalcogenides are conducting, indicating a formulation Ln^IIIQ²⁻(e-) where the electron is in conduction bands. The exceptions are SmQ, EuQ and YbQ which are semiconductors or insulators but exhibit a pressure induced transition to a conducting state. Compounds LnQ₂ are known but these do not contain Ln^IV but are Ln^III compounds containing polychalcogenide anions.

Oxysulfides Ln₂O₂S are well known, they all have the same structure with 7-coordinate Ln atoms, and 3 sulfur and 4 oxygen atoms as near neighbours. Doping these with other lanthanide elements produces phosphors. As an example, gadolinium oxysulfide, Gd₂O₂S doped with Tb³⁺ produces visible photons when irradiated with high energy X-rays and is used as a scintillator in flat panel detectors. When mischmetal, an alloy of lanthanide metals, is added to molten steel to remove oxygen and sulfur, stable oxysulfides are produced that form an immiscible solid.

Pnictides (group 15)

All of the lanthanides form a mononitride, LnN, with the rock salt structure. The mononitrides have attracted interest because of their unusual physical properties. SmN and EuN are reported as being "half metals". NdN, GdN, TbN and DyN are ferromagnetic, SmN is antiferromagnetic. Applications in the field of spintronics are being investigated. CeN is unusual as it is a metallic conductor, contrasting with the other nitrides also with the other cerium pnictides. A simple description is Ce⁴⁺N³⁻ (e–) but the interatomic distances are a better match for the trivalent state rather than for the tetravalent state. A number of different explanations have been offered. The nitrides can be prepared by the reaction of lanthanum metals with nitrogen. Some nitride is produced along with the oxide, when lanthanum metals are ignited in air. Alternative methods of synthesis are a high temperature reaction of lanthanide metals with ammonia or the decomposition of lanthanide amides, Ln(NH₂)₃. Achieving pure stoichiometric compounds, and crystals with low defect density has proved difficult. The lanthanide nitrides are sensitive to air and hydrolyse producing ammonia.

The other pnictides phosphorus, arsenic, antimony and bismuth also react with the lanthanide metals to form monopnictides, LnQ. Additionally a range of other compounds can be produced with varying stoichiometries, such as LnP₂, LnP₅, LnP₇, Ln₃As, Ln₅As₃ and LnAs₂.

Carbides

Carbides of varying stoichiometries are known for the lanthanides. Non-stoichiometry is common. All of the lanthanides form LnC₂ and Ln₂C₃ which both contain C₂ units. The dicarbides with exception of EuC₂, are metallic conductors with the calcium carbide structure and can be formulated as Ln³⁺C₂²⁻(e–). The C-C bond length is longer than that in CaC₂, which contains the C₂²⁻ anion, indicating that the antibonding orbitals of the C₂²⁻ anion are involved in the conduction band. These dicarbides hydrolyse to form hydrogen and a mixture of hydrocarbons. EuC₂ and to a lesser extent YbC₂ hydrolyse differently producing a higher percentage of acetylene (ethyne). The sesquicarbides, Ln₂C₃ can be formulated as Ln₄(C₂)₃.

These compounds adopt the Pu₂C₃ structure which has been described as having C₂²⁻ anions in bisphenoid holes formed by eight near Ln neighbours. The lengthening of the C-C bond is less marked in the sesquicarbides than in the dicarbides, with the exception of Ce₂C₃. Other carbon rich stoichiometries are known for some lanthanides. Ln₃C₄ (Ho-Lu) containing C, C₂ and C₃ units; Ln₄C₇ (Ho-Lu) contain C atoms and C₃ units and Ln₄C₅ (Gd-Ho) containing C and C₂ units. Metal rich carbides contain interstitial C atoms and no C₂ or C₃ units. These are Ln₄C₃ (Tb and Lu); Ln₂C (Dy, Ho, Tm) and Ln₃C (Sm-Lu).

Borides

All of the lanthanides form a number of borides. The "higher" borides (LnB_x where x > 12) are insulators/semiconductors whereas the lower borides are typically conducting. The lower borides have stoichiometries of LnB₂, LnB₄, LnB₆ and LnB₁₂. Applications in the field of spintronics are being investigated. The range of borides formed by the lanthanides can be compared to those formed by the transition metals. The boron rich borides are typical of the lanthanides (and groups 1–3) whereas for the transition metals tend to form metal rich, "lower" borides. The lanthanide borides are typically grouped together with the group 3 metals with which they share many similarities of reactivity, stoichiometry and structure. Collectively these are then termed the rare earth borides.

Many methods of producing lanthanide borides have been used, amongst them are direct reaction of the elements; the reduction of Ln₂O₃ with boron; reduction of boron oxide, B₂O₃, and Ln₂O₃ together with carbon; reduction of metal oxide with boron carbide, B₄C. Producing high purity samples has proved to be difficult. Single crystals of the higher borides have been grown in a low melting metal (e.g. Sn, Cu, Al).

Diborides, LnB₂, have been reported for Sm, Gd, Tb, Dy, Ho, Er, Tm, Yb and Lu. All have the same, AlB₂, structure containing a graphitic layer of boron atoms. Low temperature ferromagnetic transitions for Tb, Dy, Ho and Er. TmB₂ is ferromagnetic at 7.2 K.

Tetraborides, LnB₄ have been reported for all of the lanthanides except EuB₄, all have the same UB₄ structure. The structure has a boron sub-lattice consists of chains of octahedral B₆ clusters linked by boron atoms. The unit cell decreases in size successively from LaB₄ to LuB₄. The tetraborides of the lighter lanthanides melt with decomposition to LnB₆. Attempts to make EuB₄ have failed. The LnB₄ are good conductors and typically antiferromagnetic.

Hexaborides, LnB₆ have been reported for all of the lanthanides. They all have the CaB₆ structure, containing B₆ clusters. They are non-stoichiometric due to cation defects. The hexaborides of the lighter lanthanides (La – Sm) melt without decomposition, EuB₆ decomposes to boron and metal and the heavier lanthanides decompose to LnB₄ with exception of YbB₆ which decomposes forming YbB₁₂. The stability has in part been correlated to differences in volatility between the lanthanide metals. In EuB₆ and YbB₆ the metals have an oxidation state of +2 whereas in the rest of the lanthanide hexaborides it is +3. This rationalises the differences in conductivity, the extra electrons in the Ln^III hexaborides entering conduction bands. EuB₆ is a semiconductor and the rest are good conductors. LaB₆ and CeB₆ are thermionic emitters, used, for example, in scanning electron microscopes.

Dodecaborides, LnB₁₂, are formed by the heavier smaller lanthanides, but not by the lighter larger metals, La – Eu. With the exception YbB₁₂ (where Yb takes an intermediate valence and is a Kondo insulator), the dodecaborides are all metallic compounds. They all have the UB₁₂ structure containing a 3 dimensional framework of cubooctahedral B₁₂ clusters.

The higher boride LnB₆₆ is known for all lanthanide metals. The composition is approximate as the compounds are non-stoichiometric. They all have similar complex structure with over 1600 atoms in the unit cell. The boron cubic sub lattice contains super icosahedra made up of a central B₁₂ icosahedra surrounded by 12 others, B₁₂(B₁₂)₁₂. Other complex higher borides LnB₅₀ (Tb, Dy, Ho Er Tm Lu) and LnB₂₅ are known (Gd, Tb, Dy, Ho, Er) and these contain boron icosahedra in the boron framework.

Organometallic compounds

Lanthanide-carbon σ bonds are well known; however as the 4f electrons have a low probability of existing at the outer region of the atom there is little effective orbital overlap, resulting in bonds with significant ionic character. As such organo-lanthanide compounds exhibit carbanion-like behavior, unlike the behavior in transition metal organometallic compounds. Because of their large size, lanthanides tend to form more stable organometallic derivatives with bulky ligands to give compounds such as Ln[CH(SiMe₃)₃].^[78] Analogues of uranocene are derived from dilithiocyclooctatetraene, Li₂C₈H₈. Organic lanthanide(II) compounds are also known, such as Cp*₂Eu.

Physical properties

Magnetic and spectroscopic

All the trivalent lanthanide ions, except lanthanum and lutetium, have unpaired f electrons. However, the magnetic moments deviate considerably from the spin-only values because of strong spin-orbit coupling. The maximum number of unpaired electrons is 7, in Gd³⁺, with a magnetic moment of 7.94 B.M., but the largest magnetic moments, at 10.4–10.7 B.M., are exhibited by Dy³⁺ and Ho³⁺. However, in Gd³⁺ all the electrons have parallel spin and this property is important for the use of gadolinium complexes as contrast reagent in MRI scans.

A solution of 4% holmium oxide in 10% perchloric acid, permanently fused into a quartz cuvette as a wavelength calibration standard

Crystal field splitting is rather small for the lanthanide ions and is less important than spin-orbit coupling in regard to energy levels. Transitions of electrons between f orbitals are forbidden by the Laporte rule. Furthermore, because of the "buried" nature of the f orbitals, coupling with molecular vibrations is weak. Consequently, the spectra of lanthanide ions are rather weak and the absorption bands are similarly narrow. Glass containing holmium oxide and holmium oxide solutions (usually in perchloric acid) have sharp optical absorption peaks in the spectral range 200–900 nm and can be used as a wavelength calibration standard for optical spectrophotometers, and are available commercially.

As f-f transitions are Laporte-forbidden, once an electron has been excited, decay to the ground state will be slow. This makes them suitable for use in lasers as it makes the population inversion easy to achieve. The Nd:YAG laser is one that is widely used. Europium-doped yttrium vanadate was the first red phosphor to enable the development of color television screens. Lanthanide ions have notable luminescent properties due to their unique 4f orbitals. Laporte forbidden f-f transitions can be activated by excitation of a bound "antenna" ligand. This leads to sharp emission bands throughout the visible, NIR, and IR and relatively long luminescence lifetimes.

Occurrence

The lanthanide contraction is responsible for the great geochemical divide that splits the lanthanides into light and heavy-lanthanide enriched minerals, the latter being almost inevitably associated with and dominated by yttrium. This divide is reflected in the first two "rare earths" that were discovered: yttria (1794) and ceria (1803). The geochemical divide has put more of the light lanthanides in the Earth's crust, but more of the heavy members in the Earth's mantle. The result is that although large rich ore-bodies are found that are enriched in the light lanthanides, correspondingly large ore-bodies for the heavy members are few. The principal ores are monazite and bastnäsite. Monazite sands usually contain all the lanthanide elements, but the heavier elements are lacking in bastnäsite. The lanthanides obey the Oddo-Harkins rule – odd-numbered elements are less abundant than their even-numbered neighbors.

Three of the lanthanide elements have radioactive isotopes with long half-lives (¹³⁸La, ¹⁴⁷Sm and ¹⁷⁶Lu) that can be used to date minerals and rocks from Earth, the Moon and meteorites.

Applications

Industrial

Lanthanide elements and their compounds have many uses but the quantities consumed are relatively small in comparison to other elements. About 15000 ton/year of the lanthanides are consumed as catalysts and in the production of glasses. This 15000 tons corresponds to about 85% of the lanthanide production. From the perspective of value, however, applications in phosphors and magnets are more important.

The devices lanthanide elements are used in include superconductors, samarium-cobalt and neodymium-iron-boron high-flux rare-earth magnets, magnesium alloys, electronic polishers, refining catalysts and hybrid car components (primarily batteries and magnets). Lanthanide ions are used as the active ions in luminescent materials used in optoelectronics applications, most notably the Nd:YAG laser. Erbium-doped fiber amplifiers are significant devices in optical-fiber communication systems. Phosphors with lanthanide dopants are also widely used in cathode ray tube technology such as television sets. The earliest color television CRTs had a poor-quality red; europium as a phosphor dopant made good red phosphors possible. Yttrium iron garnet (YIG) spheres can act as tunable microwave resonators.

Lanthanide oxides are mixed with tungsten to improve their high temperature properties for TIG welding, replacing thorium, which was mildly hazardous to work with. Many defense-related products also use lanthanide elements such as night vision goggles and rangefinders. The SPY-1 radar used in some Aegis equipped warships, and the hybrid propulsion system of Arleigh Burke-class destroyers all use rare earth magnets in critical capacities. The price for lanthanum oxide used in fluid catalytic cracking has risen from $5 per kilogram in early 2010 to $140 per kilogram in June 2011.

Most lanthanides are widely used in lasers, and as (co-)dopants in doped-fiber optical amplifiers; for example, in Er-doped fiber amplifiers, which are used as repeaters in the terrestrial and submarine fiber-optic transmission links that carry internet traffic. These elements deflect ultraviolet and infrared radiation and are commonly used in the production of sunglass lenses. Other applications are summarized in the following table:

Application	Percentage
Catalytic converters	45%
Petroleum refining catalysts	25%
Permanent magnets	12%
Glass polishing and ceramics	7%
Metallurgical	7%
Phosphors	3%
Other	1%

The complex Gd(DOTA) is used in magnetic resonance imaging.

Life science

Lanthanide complexes can be used for optical imaging. Applications are limited by the lability of the complexes.

Some applications depend on the unique luminescence properties of lanthanide chelates or cryptates). These are well-suited for this application due to their large Stokes shifts and extremely long emission lifetimes (from microseconds to milliseconds) compared to more traditional fluorophores (e.g., fluorescein, allophycocyanin, phycoerythrin, and rhodamine).

The biological fluids or serum commonly used in these research applications contain many compounds and proteins which are naturally fluorescent. Therefore, the use of conventional, steady-state fluorescence measurement presents serious limitations in assay sensitivity. Long-lived fluorophores, such as lanthanides, combined with time-resolved detection (a delay between excitation and emission detection) minimizes prompt fluorescence interference.

Time-resolved fluorometry (TRF) combined with fluorescence resonance energy transfer (FRET) offers a powerful tool for drug discovery researchers: Time-Resolved Fluorescence Resonance Energy Transfer or TR-FRET. TR-FRET combines the low background aspect of TRF with the homogeneous assay format of FRET. The resulting assay provides an increase in flexibility, reliability and sensitivity in addition to higher throughput and fewer false positive/false negative results.

This method involves two fluorophores: a donor and an acceptor. Excitation of the donor fluorophore (in this case, the lanthanide ion complex) by an energy source (e.g. flash lamp or laser) produces an energy transfer to the acceptor fluorophore if they are within a given proximity to each other (known as the Förster's radius). The acceptor fluorophore in turn emits light at its characteristic wavelength.

The two most commonly used lanthanides in life science assays are shown below along with their corresponding acceptor dye as well as their excitation and emission wavelengths and resultant Stokes shift (separation of excitation and emission wavelengths).

Donor	Excitation⇒Emission λ (nm)	Acceptor	Excitation⇒Emission λ (nm)	Stoke's Shift (nm)
Eu3+	340⇒615	Allophycocyanin	615⇒660	320
Tb3+	340⇒545	Phycoerythrin	545⇒575	235

Possible medical uses

Currently there is research showing that lanthanide elements can be used as anticancer agents. The main role of the lanthanides in these studies is to inhibit proliferation of the cancer cells. Specifically cerium and lanthanum have been studied for their role as anti-cancer agents.

One of the specific elements from the lanthanide group that has been tested and used is cerium (Ce). There have been studies that use a protein-cerium complex to observe the effect of cerium on the cancer cells. The hope was to inhibit cell proliferation and promote cytotoxicity. Transferrin receptors in cancer cells, such as those in breast cancer cells and epithelial cervical cells, promote the cell proliferation and malignancy of the cancer. Transferrin is a protein used to transport iron into the cells and is needed to aid the cancer cells in DNA replication. Transferrin acts as a growth factor for the cancerous cells and is dependent on iron. Cancer cells have much higher levels of transferrin receptors than normal cells and are very dependent on iron for their proliferation.

Cerium has shown results as an anti-cancer agent due to its similarities in structure and biochemistry to iron. Cerium may bind in the place of iron on to the transferrin and then be brought into the cancer cells by transferrin-receptor mediated endocytosis. The cerium binding to the transferrin in place of the iron inhibits the transferrin activity in the cell. This creates a toxic environment for the cancer cells and causes a decrease in cell growth. This is the proposed mechanism for cerium's effect on cancer cells, though the real mechanism may be more complex in how cerium inhibits cancer cell proliferation. Specifically in HeLa cancer cells studied in vitro, cell viability was decreased after 48 to 72 hours of cerium treatments. Cells treated with just cerium had decreases in cell viability, but cells treated with both cerium and transferrin had more significant inhibition for cellular activity.

Another specific element that has been tested and used as an anti-cancer agent is lanthanum, more specifically lanthanum chloride (LaCl₃). The lanthanum ion is used to affect the levels of let-7a and microRNAs miR-34a in a cell throughout the cell cycle. When the lanthanum ion was introduced to the cell in vivo or in vitro, it inhibited the rapid growth and induced apoptosis of the cancer cells (specifically cervical cancer cells). This effect was caused by the regulation of the let-7a and microRNAs by the lanthanum ions. The mechanism for this effect is still unclear but it is possible that the lanthanum is acting in a similar way as the cerium and binding to a ligand necessary for cancer cell proliferation.

Biological effects

Due to their sparse distribution in the earth's crust and low aqueous solubility, the lanthanides have a low availability in the biosphere, and for a long time were not known to naturally form part of any biological molecules. In 2007 a novel methanol dehydrogenase that strictly uses lanthanides as enzymatic cofactors was discovered in a bacterium from the phylum Verrucomicrobia, Methylacidiphilum fumariolicum. This bacterium was found to survive only if there are lanthanides present in the environment. Compared to most other nondietary elements, non-radioactive lanthanides are classified as having low toxicity.

Rare-earth magnet

From Wikipedia, the free encyclopedia

https://en.wikipedia.org/wiki/Rare-earth_magnet 

 

Ferrofluid on glass, with a rare-earth magnet underneath

Rare-earth magnets are strong permanent magnets made from alloys of rare-earth elements. Developed in the 1970s and 1980s, rare-earth magnets are the strongest type of permanent magnets made, producing significantly stronger magnetic fields than other types such as ferrite or alnico magnets. The magnetic field typically produced by rare-earth magnets can exceed 1.4 teslas, whereas ferrite or ceramic magnets typically exhibit fields of 0.5 to 1 tesla.

There are two types: neodymium magnets and samarium–cobalt magnets. Rare-earth magnets are extremely brittle and also vulnerable to corrosion, so they are usually plated or coated to protect them from breaking, chipping, or crumbling into powder.

The development of rare-earth magnets began around 1966, when K. J. Strnat and G. Hoffer of the US Air Force Materials Laboratory discovered that an alloy of yttrium and cobalt, YCo₅, had by far the largest magnetic anisotropy constant of any material then known.

The term "rare earth" can be misleading, as some of these metals can be as abundant in the Earth's crust as tin or lead, but rare earth ores do not exist in seams (like coal or copper), so in any given cubic kilometre of crust they are "rare". The major source is currently China. Some countries classify rare earth metals as strategically important, and recent Chinese export restrictions on these materials have led some to initiate research programs to develop strong magnets that do not require rare earth metals.

Neodymium magnets (small cylinders) lifting steel balls. As shown here, rare-earth magnets can easily lift thousands of times their own weight.

Explanation of strength

The rare-earth (lanthanide) elements are metals that are ferromagnetic, meaning that like iron they can be magnetized to become permanent magnets, but their Curie temperatures (the temperature above which their ferromagnetism disappears) are below room temperature, so in pure form their magnetism only appears at low temperatures. However, they form compounds with the transition metals such as iron, nickel, and cobalt, and some of these compounds have Curie temperatures well above room temperature. Rare-earth magnets are made from these compounds.

The greater strength of rare-earth magnets is mostly due to two factors:

• First, their crystalline structures have very high magnetic anisotropy. This means that a crystal of the material preferentially magnetizes along a specific crystal axis but is very difficult to magnetize in other directions. Like other magnets, rare-earth magnets are composed of microcrystalline grains, which are aligned in a powerful magnetic field during manufacture, so their magnetic axes all point in the same direction. The resistance of the crystal lattice to turning its direction of magnetization gives these compounds a very high magnetic coercivity (resistance to being demagnetized), so that the strong demagnetizing field within the finished magnet does not reduce the material's magnetization.

• Second, atoms of rare-earth elements can have high magnetic moments. Their orbital electron structures contain many unpaired electrons; in other elements, almost all of the electrons exist in pairs with opposite spins, so their magnetic fields cancel out, but in rare-earths there is much less magnetic cancellation. This is a consequence of incomplete filling of the f-shell, which can contain up to 7 unpaired electrons. In a magnet it is the unpaired electrons, aligned so they spin in the same direction, which generate the magnetic field. This gives the materials high remanence (saturation magnetization J_s ). The maximal energy density B·H_max is proportional to J_s², so these materials have the potential for storing large amounts of magnetic energy. The magnetic energy product B·H_max of neodymium magnets is about 18 times greater than "ordinary" magnets by volume. This allows rare-earth magnets to be smaller than other magnets with the same field strength.

Magnetic properties

Some important properties used to compare permanent magnets are: remanence (B_r), which measures the strength of the magnetic field; coercivity (H_ci), the material's resistance to becoming demagnetized; energy product (B·H_max), the density of magnetic energy; and Curie temperature (T_C), the temperature at which the material loses its magnetism. Rare-earth magnets have higher remanence, much higher coercivity and energy product, but (for neodymium) lower Curie temperature than other types. The table below compares the magnetic performance of the two types of rare-earth magnets, neodymium (Nd₂Fe₁₄B) and samarium-cobalt (SmCo₅), with other types of permanent magnets.

Magnet	preparation	Br (T)	Hci (kA/m)	B·Hmax (kJ/m3)	TC (°C)
Nd2Fe14B	sintered	1.0–1.4	750–2000	200–440	310–400
Nd2Fe14B	bonded	0.6–0.7	600–1200	60–100	310–400
SmCo5	sintered	0.8–1.1	600–2000	120–200	720
Sm(Co,Fe,Cu,Zr)7	sintered	0.9–1.15	450–1300	150–240	800
Alnico	sintered	0.6–1.4	275	10–88	700–860
Sr-ferrite	sintered	0.2–0.4	100–300	10–40	450
Iron (Fe) bar magnet	annealed	?	800	?	770

Types

Samarium-cobalt

Samarium–cobalt magnets (chemical formula: SmCo₅), the first family of rare-earth magnets invented, are less used than neodymium magnets because of their higher cost and lower magnetic field strength. However, samarium–cobalt has a higher Curie temperature, creating a niche for these magnets in applications where high field strength is needed at high operating temperatures. They are highly resistant to oxidation, but sintered samarium–cobalt magnets are brittle and prone to chipping and cracking and may fracture when subjected to thermal shock.

Neodymium

Neodymium magnet with nickel plating mostly removed

Neodymium magnets, invented in the 1980s, are the strongest and most affordable type of rare-earth magnet. They are made of an alloy of neodymium, iron, and boron (Nd₂Fe₁₄B), sometimes abbreviated as NIB. Neodymium magnets are used in numerous applications requiring strong, compact permanent magnets, such as electric motors for cordless tools, hard disk drives, magnetic holddowns, and jewelry clasps. They have the highest magnetic field strength and have a higher coercivity (which makes them magnetically stable), but they have a lower Curie temperature and are more vulnerable to oxidation than samarium–cobalt magnets.

Corrosion can cause unprotected magnets to spall off a surface layer or to crumble into a powder. Use of protective surface treatments such as gold, nickel, zinc, and tin plating and epoxy-resin coating can provide corrosion protection; the majority of neodymium magnets use nickel plating to provide a robust protection.

Originally, the high cost of these magnets limited their use to applications requiring compactness together with high field strength. Both the raw materials and the patent licenses were expensive. However, since the 1990s, NIB magnets have become steadily less expensive, and their lower cost has inspired new uses such as magnetic construction toys.

Hazards

The greater force exerted by rare-earth magnets creates hazards that are not seen with other types of magnet. Magnets larger than a few centimeters are strong enough to cause injuries to body parts pinched between two magnets or a magnet and a metal surface, even causing broken bones. Magnets allowed to get too near each other can strike each other with enough force to chip and shatter the brittle material, and the flying chips can cause injuries. Starting in 2005, powerful magnets breaking off toys or from magnetic construction sets started causing injuries and deaths. Young children who have swallowed several magnets have had a fold of the digestive tract pinched between the magnets, causing injury and in one case intestinal perforations, sepsis, and death.

A voluntary standard for toys, permanently fusing strong magnets to prevent swallowing, and capping unconnected magnet strength, was adopted in 2007. In 2009, a sudden growth in sales of magnetic desk toys for adults caused a surge in injuries, with emergency room visits estimated at 3,617 in 2012. In response, the U.S. Consumer Product Safety Commission passed a rule in 2012 restricting rare-earth magnet size in consumer products, but it was vacated by a US federal court decision in November 2016, in a case brought by the one remaining manufacturer. After the rule was nullified, the number of ingestion incidents in the country rose sharply, and is estimated to exceed 1,500 in 2019.

Applications

Since their prices became competitive in the 1990s, neodymium magnets have been replacing alnico and ferrite magnets in the many applications in modern technology requiring powerful magnets. Their greater strength allows smaller and lighter magnets to be used for a given application.

Common applications

Neodymium magnet balls

Common applications of rare-earth magnets include:

• computer hard disk drives
• wind turbine generators
• speakers / headphones
• bicycle dynamos
• MRI scanners
• fishing reel brakes
• permanent magnet motors in cordless tools
• high-performance AC servo motors
• traction motors and integrated starter-generators in hybrid and electric vehicles
• mechanically powered flashlights, employing rare earth magnets for generating electricity in a shaking motion or rotating (hand-crank-powered) motion
• industrial uses such as maintaining product purity, equipment protection, and quality control
• capture of fine metallic particles in lubricating oils (crankcases of internal combustion engines, also gearboxes and differentials), so as to keep said particles out of circulation, thereby rendering them unable to cause abrasive wear of moving machine parts

Other applications

Other applications of rare-earth magnets include:

• Linear motors (used in maglev trains, etc.)
• Stop motion animation: as tie-downs when the use of traditional screw and nut tie-downs is impractical.
• Diamagnetic levitation experimentation, the study of magnetic field dynamics and superconductor levitation.
• Electrodynamic bearings
• Launched roller coaster technology found on roller coaster and other thrill rides.
• LED Throwies, small LEDs attached to a button cell battery and a small rare earth magnet, used as a form of non-destructive graffiti and temporary public art.
• Neodymium magnet toys
• Electric guitar pickups
• Miniature figures, for which rare-earth magnets have gained popularity in the miniatures gaming community for their small size and relative strength assisting in basing and swapping weapons between models.

Rare-earth-free permanent magnets

The United States Department of Energy has identified a need to find substitutes for rare-earth metals in permanent-magnet technology and has begun funding such research. The Advanced Research Projects Agency-Energy (ARPA-E) has sponsored a Rare Earth Alternatives in Critical Technologies (REACT) program, to develop alternative materials. In 2011, ARPA-E awarded 31.6 million dollars to fund Rare-Earth Substitute projects.

Recycling efforts

The European Union's ETN-Demeter project (European Training Network for the Design and Recycling of Rare-Earth Permanent Magnet Motors and Generators in Hybrid and Full Electric Vehicles) is examining sustainable design of electric motors used in vehicles. They are, for example, designing electric motors in which the magnets can be easily removed for recycling the rare earth metals.

The European Union's European Research Council also awarded to Principal Investigator, Prof. Thomas Zemb, and co-Principal Investigator, Dr. Jean-Christophe P. Gabriel, an Advanced Research Grant for the project "Rare Earth Element reCYCling with Low harmful Emissions : REE-CYCLE", which aimed at finding new processes for the recycling of rare earth.

 

Atomoxetine

From Wikipedia, the free encyclopedia

https://en.wikipedia.org/wiki/Atomoxetine 

 

Atomoxetine

Clinical data
Trade names	Strattera, others
Other names	(R)-N-Methyl-3-phenyl-3-(o-tolyloxy)propan-1-amine
AHFS/Drugs.com	Monograph
MedlinePlus	a603013
License data	US DailyMed: Atomoxetine
Pregnancy category	AU: B3
Routes of administration	By mouth
ATC code	N06BA09 (WHO)
Legal status
Legal status	AU: S4 (Prescription only) CA: ℞-only UK: POM (Prescription only) US: ℞-only
Pharmacokinetic data
Bioavailability	63 to 94%
Protein binding	98%
Metabolism	Liver, via CYP2D6
Elimination half-life	4.5-19 hours
Excretion	Kidney (80%) and faecal (17%)
Identifiers

IUPAC name
CAS Number	83015-26-3  82248-59-7 (HCl)
PubChem CID	54841
IUPHAR/BPS	7118
DrugBank	DB00289
ChemSpider	49516
UNII	ASW034S0B8
KEGG	D07473
ChEBI	CHEBI:127342
ChEMBL	ChEMBL641
CompTox Dashboard (EPA)	DTXSID9044297
ECHA InfoCard	100.120.306
Chemical and physical data
Formula	C17H21NO
Molar mass	255.361 g·mol−1
3D model (JSmol)	Interactive image

Atomoxetine, sold under the brand name Strattera, among others, is a medication used to treat attention deficit hyperactivity disorder (ADHD). It may be used alone or along with psychostimulants. Use of atomoxetine is only recommended for those who are at least six years old. It is taken by mouth. It was approved for medical use in the United States in 2002. In 2018, it was the 162nd most commonly prescribed medication in the United States, with more than 3 million prescriptions.

Common side effects of atomoxetine include abdominal pain, loss of appetite, nausea, feeling tired, and dizziness. Serious side effects may include angioedema, liver problems, stroke, psychosis, heart problems, suicide, and aggression. There is a lack of data regarding its safety during pregnancy; as of 2019, its safety during pregnancy and for use during breastfeeding is not certain.

Atomoxetine is a norepinephrine reuptake inhibitor and is believed to work by increasing norepinephrine and dopamine levels in the brain.

Medical uses

Attention deficit hyperactivity disorder

Atomoxetine is approved for use in children, adolescents, and adults. However, its efficacy has not been studied in children under six years old. Its primary advantage over the standard stimulant treatments for ADHD is that it has little known abuse potential. While it has been shown to significantly reduce inattentive and hyperactive symptoms, the responses were lower than the response to stimulants. Additionally, 40% of participants who were treated with atomoxetine experienced residual ADHD symptoms.

While its efficacy may be less than that of stimulant medications, there is some evidence that it may be used in combination with stimulants. Doctors may prescribe non-stimulants including atomoxetine when a person has bothersome side effects from stimulants; when a stimulant was not effective; in combination with a stimulant to increase effectiveness; or when there is concern about the abuse potential of psychostimulants in a patient with a history of drug use disorder.

Unlike α₂ adrenoceptor agonists such as guanfacine and clonidine, atomoxetine's use can be abruptly stopped without significant discontinuation effects being seen.

The initial therapeutic effects of atomoxetine usually take 2–4 weeks to become apparent. A further 2–4 weeks may be required for the full therapeutic effects to be seen. The maximum recommended total daily dose in children and adolescents over 70 kg and adults is 100 mg.

Other

Atomoxetine may be used in those with ADHD and bipolar disorder although such use has not been well studied. Some benefit has also been seen in people with ADHD and autism.

Contraindications

Contraindications include:

• Hypersensitivity to atomoxetine or any of the inactive ingredients in the product
• Symptomatic cardiovascular disease including:

-moderate to severe hypertension

-atrial fibrillation

-atrial flutter

-ventricular tachycardia

-ventricular fibrillation

-ventricular flutter

-advanced arteriosclerosis

• Severe cardiovascular disorders
• Pheochromocytoma
• Concomitant treatment with monoamine oxidase inhibitors
• Narrow angle glaucoma
• Poor metabolizers (due to the metabolism of atomoxetine by CYP2D6)

Adverse effects

Common side effects include abdominal pain, loss of appetite, nausea, feeling tired, and dizziness. Serious side effects may include angioedema, liver problems, stroke, psychosis, heart problems, suicide, and aggression. A 2020 meta-analysis found that atomoxetine was associated with anorexia, weight loss, and hypertension, rating it as a "potentially least preferred agent based on safety" for treating ADHD. As of 2019, safety in pregnancy and breastfeeding is not clear; a 2018 review stated that, "[b]ecause of lack of data, the treating physician should consider stopping atomoxetine treatment in women with ADHD during pregnancy."

The U.S. Food and Drug Administration (FDA) has issued a black box warning for suicidal behavior/ideation. Similar warnings have been issued in Australia. Unlike stimulant medications, atomoxetine does not have abuse liability or the potential to cause withdrawal effects on abrupt discontinuation.

Incidence of adverse effects:

Very common (>10% incidence) adverse effects include:

• Nausea (26%)
• Xerostomia (Dry mouth) (20%)
• Appetite loss (16%)
• Insomnia (15%)
• Fatigue (10%)
• Headache
• Cough
• Vomiting (in children and adolescents)

Common (1–10% incidence) adverse effects include:

• Constipation (8%)
• Dizziness (8%)
• Erectile dysfunction (8%)
• Somnolence (sleepiness) (8%)
• Abdominal pain (7%)
• Urinary hesitation (6%)
• Tachycardia (high heart rate) (5–10%)
• Hypertension (high blood pressure) (5–10%)
• Irritability (5%)
• Abnormal dreams (4%)
• Dyspepsia (4%)
• Ejaculation disorder (4%)
• Hyperhidrosis (abnormally increased sweating) (4%)
• Vomiting (4%)
• Hot flashes (3%)
• Paraesthesia (sensation of tingling, tickling, etc.) (3%)
• Menstrual disorder (3%)
• Weight loss (2%)
• Depression
• Sinus headache
• Dermatitis
• Mood swings

Uncommon (0.1–1% incidence) adverse effects include:

• Suicide-related events
• Hostility
• Emotional lability
• Aggression
• Psychosis
• Syncope (fainting)
• Tremor
• Migraine
• Hypoaesthesia
• Seizure
• Palpitations
• Sinus tachycardia
• QT interval prolongation
• Increased blood bilirubin
• Allergic reactions

Rare (0.01–0.1% incidence) adverse effects including:

• Raynaud's phenomenon
• Abnormal/increased liver function tests
• Jaundice
• Hepatitis
• Liver injury
• Acute liver failure
• Urinary retention
• Priapism
• Male genital pain

Overdose

Atomoxetine is relatively non-toxic in overdose. Single-drug overdoses involving over 1500 mg of atomoxetine have not resulted in death. The most common symptoms of overdose include:

• Gastrointestinal symptoms
• Somnolence
• Dizziness
• Tremor
• Abnormal behaviour
• Hyperactivity
• Agitation
• Dry mouth
• Tachycardia
• Hypertension
• Mydriasis

Less common symptoms:

• Seizures
• QTc interval prolongation

The recommended treatment for atomoxetine overdose includes use of activated charcoal to prevent further absorption of the drug.

Interactions

Atomoxetine is a substrate for CYP2D6. Concurrent treatment with a CYP2D6 inhibitor such as bupropion, fluoxetine, or paroxetine has been shown to increase plasma atomoxetine by 100% or more, as well as increase N-desmethylatomoxetine levels and decrease plasma 4-hydroxyatomoxetine levels by a similar degree.

Atomoxetine has been found to directly inhibit hERG potassium currents with an IC₅₀ of 6.3 μM, which has the potential to cause arrhythmia. QT prolongation has been reported with atomoxetine at therapeutic doses and in overdose; it is suggested that atomoxetine not be used with other medications that may prolong the QT interval, concomitantly with CYP2D6 inhibitors, and caution to be used in poor metabolizers.

Other notable drug interactions include:

• Antihypertensive agents, due to atomoxetine acting as an indirect sympathomimetic
• Indirect-acting sympathomimetics, such as pseudoephedrine, norepinephrine reuptake inhibitors, or MAOIs
• Direct-acting sympathomimetics, such as phenylephrine or other α₁ adrenoceptor agonists, including pressors such as dobutamine or isoprenaline and β₂ adrenoceptor agonists
• Highly plasma protein-bound drugs: atomoxetine has the potential to displace these drugs from plasma proteins which may potentiate their adverse or toxic effects. In vitro, atomoxetine does not affect the plasma protein binding of aspirin, desipramine, diazepam, paroxetine, phenytoin, or warfarin

Pharmacology

Pharmacodynamics

Atomoxetine (and metabolites)

Site	ATX	4-OH-ATX	N-DM-ATX
SERT	77	43	ND
NET	5	3	92
DAT	1,451	ND	ND
5-HT1A	>1,000	ND	ND
5-HT1B	>1,000	ND	ND
5-HT1D	>1,000	ND	ND
5-HT2	2,000	1,000	1,700
5-HT6	>1,000	ND	ND
5-HT7	>1,000	ND	ND
α1	11,400	20,000	19,600
α2A	29,800	>30,000	>10,000
β1	18,000	56,100	32,100
M1	>100,000	>100,000	>100,000
M2	>100,000	>100,000	>100,000
D1	>10,000	>10,000	>10,000
D2	>10,000	>10,000	>10,000
H1	12,100	>100,000	>100,000
MOR	ND	422	ND
DOR	ND	300	ND
KOR	ND	95	ND
σ1	>1,000	ND	ND
GABAA	200	>30,000	>10,000
NMDA	3,470a	ND	ND
Kir3.1/3.2	10,900b	ND	ND
Kir3.2	12,400b	ND	ND
Kir3.1/3.4	6,500b	ND	ND
hERG	6,300	20,000	5,710
Values are Ki (nM). The smaller the value, the more strongly the drug binds to the site. All values are for human receptors unless otherwise specified. arat cortex. bXenopus oocytes. Additional sources:

Atomoxetine inhibits the presynaptic norepinephrine transporter (NET), preventing the reuptake of norepinephrine throughout the brain along with inhibiting the reuptake of dopamine in specific brain regions such as the prefrontal cortex, where dopamine transporter (DAT) expression is minimal. In rats, atomoxetine increased prefrontal cortex catecholamine concentrations without altering dopamine levels in the striatum or nucleus accumbens; in contrast, methylphenidate, a dopamine reuptake inhibitor, was found to increase prefrontal, striatal, and accumbal dopamine levels to the same degree. In mice, atomoxetine was also found to increase prefrontal catecholamine levels without affecting striatal or accumbal levels.

Atomoxetine's status as a serotonin transporter (SERT) inhibitor at clinical doses in humans is uncertain. A PET imaging study on rhesus monkeys found that atomoxetine occupied >90% and >85% of neural NET and SERT, respectively. However, both mouse and rat microdialysis studies have failed to find an increase in extracellular serotonin in the prefrontal cortex following acute or chronic atomoxetine treatment. Supporting atomoxetine's selectivity, a human study found no effects on platelet serotonin uptake (a marker of SERT inhibition) and inhibition of the pressor effects of tyramine (a marker of NET inhibition).

Atomoxetine has been found to act as an NMDA receptor antagonist in rat cortical neurons at therapeutic concentrations. It causes a use-dependent open-channel block and its binding site overlaps with the Mg²⁺ binding site. Atomoxetine's ability to increase prefrontal cortex firing rate in anesthetized rats could not be blocked by D₁ or α₂-adrenergic receptor antagonists, but could be potentiated by NMDA or an α₁-adrenergic receptor antagonist, suggesting a glutaminergic mechanism. In Sprague Dawley rats, atomoxetine reduces NR2B protein content without altering transcript levels. Aberrant glutamate and NMDA receptor function have been implicated in the etiology of ADHD.

Atomoxetine also reversibly inhibits GIRK currents in Xenopus oocytes in a concentration-dependent, voltage-independent, and time-independent manner. K_ir3.1/3.2 ion channels are opened downstream of M₂, α₂, D₂, and A₁ stimulation, as well as other G_i-coupled receptors. Therapeutic concentrations of atomoxetine are within range of interacting with GIRKs, especially in CYP2D6 poor metabolizers. It is not known whether this contributes to the therapeutic effects of atomoxetine in ADHD.

4-Hydroxyatomoxetine, the major active metabolite of atomoxetine in CYP2D6 extensive metabolizers, has been found to have sub-micromolar affinity for opioid receptors, acting as an antagonist at μ-opioid receptors and a partial agonist at κ-opioid receptors. It is not known whether this action at the kappa-opioid receptor leads to CNS-related adverse effects.

Pharmacokinetics

Orally administered atomoxetine is rapidly and completely absorbed. First-pass metabolism by the liver is dependent on CYP2D6 activity, resulting in an absolute bioavailability of 63% for extensive metabolizers and 94% for poor metabolizers. Maximum plasma concentration is reached in 1–2 hours. If taken with food, the maximum plasma concentration decreases by 10-40% and delays the t_max by 1 hour. Drugs affecting gastric pH have no effect on oral bioavailability.

Atomoxetine has a volume of distribution of 0.85 L/kg, with limited partitioning into red blood cells. It is highly bound to plasma proteins (98.7%), mainly albumin, along with α₁-acid glycoprotein (77%) and IgG (15%). Its metabolite N-desmethylatomoxetine is 99.1% bound to plasma proteins, while 4-hydroxyatomoxetine is only 66.6% bound.

The half-life of atomoxetine varies widely between individuals, with an average range of 4.5 to 19 hours. As atomoxetine is metabolized by CYP2D6, exposure may be increased 10-fold in CYP2D6 poor metabolizers.

Atomoxetine, N-desmethylatomoxetine, and 4-hydroxyatomoxetine produce minimal to no inhibition of CYP1A2 and CYP2C9, but inhibit CYP2D6 in human liver microsomes at concentrations between 3.6-17 μmol/L. Plasma concentrations of 4-hydroxyatomoxetine and N-desmethylatomoxetine at steady state are 1.0% and 5% that of atomoxetine in CYP2D6 extensive metabolizers, and are 5% and 45% that of atomoxetine in CYP2D6 poor metabolizers.

Atomoxetine is excreted unchanged in urine at <3% in both extensive and poor CYP2D6 metabolizers, with >96% and 80% of a total dose being excreted in urine, respectively. The fractions excreted in urine as 4-hydroxyatomoxetine and its glucuronide account for 86% of a given dose in extensive metabolizers, but only 40% in poor metabolizers. CYP2D6 poor metabolizers excrete greater amounts of minor metabolites, namely N-desmethylatomoxetine and 2-hydroxymethylatomoxetine and their conjugates.

Major metabolites of atomoxetine in humans.

Pharmacogenomics

Chinese adults homozygous for the hypoactive CYP2D6*10 allele have been found to exhibit two-fold higher AUCs and 1.5-fold higher maximum plasma concentrations compared to extensive metabolizers.

Japanese men homozygous for CYP2D6*10 have similarly been found to experience two-fold higher AUCs compared to extensive metabolizers.

Chemistry

Atomoxetine, or (−)-methyl[(3R)-3-(2-methylphenoxy)-3-phenylpropylamine, is a white, granular powder that is highly soluble in water.

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  Strattera 60-mg capsule back

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  Strattera 60-mg capsule front with Lilly logo

Synthesis

Original synthesis of atomoxetine, as patented by Eli Lilly and Company

Detection in biological fluids

Atomoxetine may be quantitated in plasma, serum or whole blood in order to distinguish extensive versus poor metabolizers in those receiving the drug therapeutically, to confirm the diagnosis in potential poisoning victims or to assist in the forensic investigation in a case of fatal overdosage.

History

Atomoxetine is manufactured, marketed, and sold in the United States as the hydrochloride salt (atomoxetine HCl) under the brand name Strattera by Eli Lilly and Company, the original patent-filing company and current U.S. patent owner. Atomoxetine was initially intended to be developed as an antidepressant, but it was found to be insufficiently efficacious for treating depression. It was, however, found to be effective for ADHD and was approved by the FDA in 2002, for the treatment of ADHD. Its patent expired in May 2017. On 12 August 2010, Lilly lost a lawsuit that challenged its patent on Strattera, increasing the likelihood of an earlier entry of a generic into the US market. On 1 September 2010, Sun Pharmaceuticals announced it would begin manufacturing a generic in the United States. In a 29 July 2011 conference call, however, Sun Pharmaceutical's Chairman stated "Lilly won that litigation on appeal so I think [generic Strattera]'s deferred."

In 2017 the FDA approved the generic production of atomoxetine by four pharmaceutical companies.

Society and culture

Brand names

In India, atomoxetine is sold under brand names including Axetra, Axepta, Attera, Tomoxetin, and Attentin. In Australia, Portugal, Italy and Romania, atomoxetine is sold under the brand name Strattera. In Iran, atomoxetine is sold under brand names including Stramox. In 2017, a generic version was approved in the United States.

Research

There has been some suggestion that atomoxetine might be a helpful adjunct in people with major depression, particularly in cases with concomitant ADHD.