|Lanthanides in the periodic table|
The lanthanide // or lanthanoid // series of chemical elements1 comprises the fifteen metallic chemical elements with atomic numbers 57 through 71, from lanthanum through lutetium.234 These fifteen lanthanide 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; lutetium, a d-block element, is also generally considered to be a lanthanide due to its chemical similarities with the other fourteen. All lanthanide elements form trivalent cations, Ln3+, whose chemistry is largely determined by the ionic radius, which decreases steadily from lanthanum to lutetium.
The lanthanide elements are the group of elements with atomic number increasing from 57 (lanthanum) to 71 (lutetium). They are termed lanthanide because the lighter elements in the series are chemically similar to lanthanum. Strictly speaking, both lanthanum and lutetium have been labeled as group 3 elements, because they both have a single valence electron in the d shell. However, both elements are often included in any general discussion of the chemistry of the lanthanide elements.
In presentations of the periodic table, the lanthanides and the actinides are customarily shown as two additional rows below the main body of the table,2 with placeholders or else a selected single element of each series (either lanthanum or lutetium, and either actinium or lawrencium, respectively) shown in a single cell of the main table, between barium and hafnium, and radium and rutherfordium, respectively. This convention is entirely a matter of aesthetics and formatting practicality; a rarely used wide-formatted periodic table inserts the lanthanide and actinide series in their proper places, as parts of the table's sixth and seventh rows (periods).
- 1 Etymology
- 2 Chemistry and compounds
- 3 Physical properties
- 4 Occurrence
- 5 Applications
- 6 Biological effects
- 7 See also
- 8 References
- 9 External links
Together with scandium and yttrium, the trivial name "rare earths" is sometimes used to describe all the lanthanides. This name arises from the minerals from which they were isolated, which were uncommon oxide-type minerals. However, the use of the name is deprecated by IUPAC, as the 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)citation needed. Cerium is the 26th most abundant element in the Earth's crust, neodymium is more abundant than gold and even thulium (the least common naturally occurring lanthanide) is more abundant than iodine,5 which is itself common enough for biology to have evolve critical usages for. 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 probably introduced by Victor Goldschmidt in 1925.6
|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*||5d1||4f15d1||4f3||4f4||4f5||4f6||4f7||4f75d1||4f9||4f10||4f11||4f12||4f13||4f14||4f145d1|
|Ln3+ electron configuration*7||4f08||4f1||4f2||4f3||4f4||4f5||4f6||4f7||4f8||4f9||4f10||4f11||4f12||4f13||
|Ln3+ radius (pm)9||103||102||99||98.3||97||95.8||94.7||93.8||92.3||91.2||90.1||89||88||86.8||86.1|
|Ln2+ ion colour in aqueous solution9||—||—||—||—||—||Blood red||Colourless||—||—||—||—||—||Violet-red||Yellow-green||—|
|Ln3+ ion colour in aqueous solution10||Colorless||Colorless||Green||Violet||Pink||Pale yellow||Colorless||Colorless||V. pale pink||Pale yellow||Yellow||Rose||Pale green||Colorless||Colorless|
|Ln4+ ion colour in aqueous solution9||—||Orange-yellow||Yellow||Blue-violet||—||—||—||—||Red-brown||Orange-yellow||—||—||—||—||—|
* Between initial [Xe] and final 6s2 electronic shells
The electronic structure of the lanthanide elements, with minor exceptions is [Xe]6s24fn. In their compounds, the 6s electrons and (usually) one 4f electron are lost and the ions have the configuration [Xe]4fm.13 The chemistry of the lanthanides differs from main group elements and transition metals because of the nature of the 4f orbitals. These orbitals are "buried" inside the atom and are shielded from the atom's environment by the 4d and 5p electrons. As a consequence of this, the chemistry of the elements is largely determined by their size, which decreases gradually from 102 pm (La3+) with increasing atomic number to 86 pm (Lu3+), the so-called lanthanide contraction. 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 which 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 Ln0/3+ 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).9
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)]–.914
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 and with the exception of SmI215 and cerium(IV) salts16 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,17 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, which typically results in complexes which will 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%.18 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 complex192021 and change the nuclearity of metal clusters.2223
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.24 It should be noted however, that 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.
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)93+ but the smaller ions are 8-coordinate, [Ln(H2O)83+. There is some evidence that the later lanthanides have more water molecules in the second coordination sphere.25 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).
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 Ln2+, 2I–, e–, include NdI2, DyI2 and TmI2. Many of the iodides form soluble complexes with ethers, e.g. TmI2(dimethoxyethane)3.26 Samarium(II) iodide is a useful reducing agent. Ln(II) complexes can be synthesized by transmetalation reactions.
Ce(IV) in ceric ammonium nitrate is a useful oxidizing agent. Otherwise tetravalent lanthanides are rare. The Ce(IV) is the exception owing to the tendency to form an unfilled f shell.
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 behaviour, unlike 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)3.27 Similarly complexes of cyclopentadienyl anion (Cp–), e.g. [Ln(C5H5)3, are far less common than the corresponding pentamethylcyclopentadienyl, e.g. [Ln(C5Me5)3Cl]. Analogues of uranocene are derived from dilithiocyclooctatetraene, Li2C8H8. Organic lanthanide(II) compounds are also known, such as Cp*2Eu.26
All the trivalent lanthanide ions, except 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.
Crystal field splitting is rather small for the lanthanide ions and is less important than spin-orbit coupling in regard to energy levels.28 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,29 and are available commercially.30
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.31 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.32
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.33
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.34
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).35 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 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.36 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 2011.37
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:5
|Petroleum refining catalysts||25|
|Glass polishing and ceramics||7|
As mentioned in the industrial applications section above, lanthanide metals are particularly useful in technologies that take advantage of their reactivity to specific wavelengths of light.38 Certain life science applications take advantage of the unique fluorescence properties of lanthanide ion complexes (Ln(III) 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)|
Due to their sparse distribution in the earth's crust and low aqueous solubility, the lanthanides have a low availability in the biosphere, and are not known to naturally form part of any biological molecules. Compared to most other nondietary elements, non-radioactive lanthanides are classified as having low toxicity.34
- The current IUPAC recommendation is that the name lanthanoid be used rather than lanthanide, as the suffix "-ide" is preferred for negative ions whereas the suffix "-oid" indicates similarity to one of the members of the containing family of elements. However, lanthanide is still favored in most (~90%) scientific articles and is currently adopted on Wikipedia. In the older literature, the name "lanthanon" was often used.
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- There exist other naturally occurred radioactive isotopes of lanthanides with long half-lives (144Nd, 150Nd, 148Sm, 151Eu, 152Gd) but they are not used as chronometers.
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- lanthanide Sparkle Model, used in the computational chemistry of lanthanide complexes
- USGS Rare Earths Statistics and Information
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