THE LANTHANIDES |
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JOOP A. PETERS, DELFT UNIVERSITY OF TECHNOLOGY, THE NETHERLANDS, AND DOUGLAS J. RABER, BOARD ON CHEMICAL SCIENCES & TECHNOLOGY, THE NATIONAL RESEARCH COUNCIL
Both of us are organic chemists by training, and until the end of the 1960s, we knew only vaguely where the lanthanides were located in the periodic table. For the most part, our teachers had treated this heretical family of elements only in the last hurried classes of the relevant courses. So our knowledge about these elements was limited to awareness of the very similar chemical properties of all these elements and the understanding that together with scandium and yttrium they are also called, mistakenly as it turns out, rare earths.
Things changed drastically after reading an intriguing paper by C. C. Hinckley [J. Am. Chem. Soc., 91, 51 (1969)], in which he reported on the use of paramagnetic Eu(III) complexes for the resolution of 1H NMR spectra. Soon after, we tried these complexes out on our samples and were very excited to see that our 60-MHz 1H nuclear magnetic resonance spectra changed from featureless heaps into well-resolved first-order spectra. But the use of lanthanide chelates as aids in NMR structure determination was not to be an area of major growth. The application of Eu(III) chelates for spectral resolution was soon superseded by the introduction of high magnetic fields and multidimensional techniques in NMR spectroscopy. Nevertheless, we remained fascinated by the structures, magnetic properties, and complexation reactions of the elements in this family.
For Mendeleyev, each discovery of a new rare-earth element meant a new puzzle, because each of them showed very similar chemical behavior that made it difficult to assign positions in his periodic table. This unique chemical similarity is due to the shielding of 4f valence electrons by the completely filled 5s2 and 5p6 orbitals. The beauty of this family of elements is that, although the members are very similar from a chemical point of view, each of them has its own very specific physical properties--including color, luminescent behavior, and nuclear magnetic properties. While exploring the possibilities of the latter for structural analysis, our paths crossed in the 1970s, and the lanthanides appeared to catalyze not only a fruitful professional collaboration but also a personal friendship.
Our mutual interests in the structures and magnetic properties of lanthanide complexes led to a long-distance collaboration that has now entered its fourth decade. Our intellectual curiosity has generated useful science, trans-Atlantic travel, and early forays into the use of e-mail to transfer data and manuscripts. Like the lanthanides used as phosphors in television screens, the experience has been enlightening.
At that time, studies of the coordination chemistry and the physical chemistry of the lanthanides were flourishing, but only a few, albeit important, practical applications existed. The major chemical one probably was the use of lanthanide-exchanged faujasite as a cracking catalyst used to refine crude oil into gasoline and other fuels. Such cracking catalysts are, from an economic point of view, probably the most successful catalysts ever developed. In many aspects of the automotive industry, lanthanides are becoming increasingly important. One can find them in catalytic converters and in alloys for very stable and powerful permanent magnets used in the antilock braking systems, air bags, and electric motors of the vehicles of the future. The importance of these applications is reflected in the distribution of rare earths by use: automotive catalytic converters, 15%; petroleum refining catalysts, 16%; and permanent magnets, 8%. At the disposal of the modern synthetic chemist is an array of lanthanide-based catalysts that exploits the high charge density and small differences in ionic radius among the lanthanides to achieve high reactivities and selectivities.
Most applications of lanthanides are based on their physical properties. First of all, there is the traditional use for staining glass and ceramics. Nowadays, one can find them everywhere. The TV screens and computer monitors that we use for our communication across the ocean have lanthanide phosphors. The glass fibers used for data transport contain lanthanides, and our offices and houses are illuminated with energy-saving tricolor lanthanide-based luminescent lamps. The new euro banknotes that were introduced in Europe in 2002 are protected against counterfeiting by Eu(III) compounds that luminesce in red, green, and blue upon excitation with UV light.
During our mutual visits in the early 1980s, we would daydream about potential applications of lanthanides in medicine. The first medical applications in this field became reality shortly after the development of magnetic resonance imaging and the introduction of this technique in medical diagnosis. MRI is an NMR technique that visualizes, with a very high resolution, the morphology of the body. The intensity of each voxel in a three-dimensional image reflects the intensity of the 1H NMR signal of the water in the corresponding part of body. The intensities of these signals and, consequently, the contrast of the images are dependent on magnetic relaxation of the nuclei.
Relaxation can be enhanced by paramagnetic compounds, and the lanthanide ion Gd(III) with its seven unpaired electrons is the paramagnetic champion of the periodic table. This ion is ideal for improving the contrast in MRI scans. Gd(III) chelates such as Gd(DTPA) and Gd(DOTA) have been developed [Prog. Nucl. Magn. Reson. Spectrosc., 28, 283 (1996)] that have a very low toxicity, even at the relatively high doses in which they are applied. These contrast agents are as safe as an aspirin, and they have contributed to the success of MRI in clinical diagnostics. Nowadays, about 30% of MRI scans are performed after administration of a Gd(III)-based contrast agent.
Currently available contrast agents distribute rather unselectively over the extracellular space. Consequently, much of the research in this field is aimed at the development of more specific contrast agents by linking Gd(III) chelates to moieties that can bring about an accumulation of the contrast agents in the region of interest. The increased knowledge of receptors allows the design of contrast agents for molecular imaging--contrast agents that respond to, for example, functional groups that are overexpressed due to disease. Moreover, smart contrast agents are being developed that can report parameters such as concentration of enzymes, pH, pO2, and temperature.
The radioisotopes of the lanthanides show a variety of radiation characteristics that are suitable for applications ranging from diagnostics with positron-emitting tomography to radiotherapy. In the latter case, targeting is essential to limit damage to healthy tissue. Excellent results have been obtained with DOTA-type chelates conjugated to cyclic tumor-targeting peptides. The radiodiagnostic techniques are in a way complementary to the MRI contrast techniques. The sensitivity of the radiodiagnostic reagents is much higher than that of the MRI contrast agents (nM and mM concentrations are required, respectively), but on the other hand, the resolution of radiodiagnostic techniques is much lower.
The luminescent properties of the lanthanides also have been utilized in medical diagnosis. A variety of luminescent bioassays and sensors have been developed that take advantage of the unique luminescent properties of these elements, such as a relatively long-lived emission.
The applications of lanthanides are numerous, but we have presented only a small and quite personal selection. We expect that many new and exciting applications will emerge in the future. There is, however, no need to worry that these elements will be depleted in the near future. The first black stone containing lanthanides was found by the Swedish army lieutenant Arrhenius in 1787 near Ytterby, a small village not far from Stockholm. The ruins of the resulting mine were designated by ASM International as a historical landmark in 1989. The lanthanides are sometimes called the rare earths, a name that dates to the 18th century and misleadingly suggests that their abundance is low. At present, the world reserve of rare earths is estimated to be about 150 million metric tons, in the form of rare-earth oxides, while the annual consumption is only about 120,000 tons.
The lanthanides have played an important role in our lives. They initiated our collaboration and friendship, and they helped us in our careers, during which the lanthanides have gained significantly in importance and will undoubtedly continue to do so. Our conclusion is a simple one: Teachers and textbooks can no longer treat the lanthanides as an insignificant footnote to the periodic table.
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RAINBOW Electron micrographs of lanthanum chloride crystals.
PHOTO COURTESY OF ROB BERDAN
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Joop A. Peters is universitair hoofddocent at Delft University of Technology, the Netherlands, and Douglas J. Raber is senior scholar with the Board on Chemical Sciences & Technology at the National Research Council in Washington, D.C. Peters and Raber have maintained a long-term research collaboration ever since they met in 1979.
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LANTHANIDES AT A GLANCE
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Name: From the Greek lanthaneis, to lie hidden. Also called rare earths.They are lanthanum (La), cerium (Ce), praseodymium (Pr), neodymium (Nd), promethium (Pm), samarium (Sm), europium (Eu), gadolinium(Gd), terbium (Tb), dysprosium (Dy), holmium (Ho), erbium (Er), thulium (Tm), ytterbium (Yb), and lutetium (Lu).
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Occurrence: Despite their name, the lanthanides are not particularly rare. La, Ce, and Nd are more common than lead.
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Behavior: The lanthanides are all very reactive and electropositive. The chemistry is dominated by the +3 oxidation state. Despite the high charge, the large size of the Ln(III) ions results in low charge densities. Lanthanide compounds are predominantly ionic in character.
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Appearance: Silvery white, but they tarnish in air.
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Uses: Lanthanide silicides improve strength of low alloy steels; other lanthanides are used for lenses and magnets.
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THORIUM
SETH H. GRAE, THORIUM POWER
Thorium was discovered by the Swedish chemist Jöns J. Berzelius in 1828. He named it after Thor, the Norse god of thunder and war, because of its power. Thorium is the most efficient nuclear reactor fuel, so there have been many attempts to produce reactor fuel from thorium, including current unrelated efforts by Thorium Power and by institutes in India.
Thorium oxide has the highest melting point of all oxides, which can provide a safety benefit as a nuclear reactor fuel. Thorium is estimated to be three times as abundant as uranium, holding more untapped energy than all oil, coal, natural gas, and uranium combined, with vast reserves in the U.S., Australia, India, Canada, and many other countries. Thorium is also responsible for most of the internal heating of Earth.
Thorium is stockpiled around the world. When rare-earth elements are separated from monazite sands for use as a fuel-cracking catalyst in the petroleum industry, a by-product is thorium. Although hundreds of tons of thorium are stockpiled from this process, only small amounts are used in commercial processes and products. In the U.S., a commercial use of thorium was small quantities in the wicks of camping lanterns. I have been told that Coleman stopped this use when it found out that people were testing Geiger counters on the wicks. In addition, small quantities of various forms of thorium have been produced for nonenergy uses such as ceramics, magnesium-thorium alloys, and welding electrodes.
I first heard of thorium in 1992 from Alvin Radkowsky, who designed the thorium fuel that Thorium Power is now developing. Alvin was the original chief scientist for the U.S. Naval Reactors Program for Adm. Hyman G. Rickover, and he also headed the design team for the first nuclear power plant on land, at Shippingport, Pa. That reactor used thorium in its first core and was the first reactor under former president Dwight D. Eisenhower's Atoms for Peace program, which began 50 years ago.
When Radkowsky's former professor, Edward Teller, asked him in 1986 if he could come up with a fuel design to address proliferation concerns, Radkowsky turned back to thorium as the answer. Radkowsky passed away last year, but his work led to the fuel design now being developed and tested in Russia to dispose of weapons-grade plutonium and to stop the production of new weapons-suitable plutonium. Radkowsky spent most of his career in the U.S. Navy program during the Cold War and was amused by how times change; his technology is being tested in Russia to dispose of plutonium from warheads.
The project has been headed in Russia by academician Nikolai N. Ponomarev-Stepnoi, vice president of the Kurchatov Institute, where the first Soviet atomic bomb was developed. The work in Russia has determined that the thorium fuel design will dispose of more plutonium per year in a reactor than any other method, and current estimates are that the costs will prove to be significantly less than any other method. In addition, fuels that burn plutonium without thorium actually produce large quantities of new weapons-suitable plutonium in the spent fuel; this is not the case with thorium fuel. The fuel also results in dramatically less toxic spent fuel, with much less spent fuel for the same amount of electricity generated.
Reps. Jim Gibbons (R-Nev.) and Curt Weldon (R-Pa.) have taken up the cause of thorium fuel for plutonium disposition in Russia. They have traveled to Moscow to speak with scientists and engineers and witnessed the thorium fuel ampules in a test reactor in Moscow. These congressmen are leading the effort to secure government funding for the project.
Promising new thorium fuel technologies eventually may lead governments and the nuclear power industry to uncover the untapped energy potential of thorium for peaceful energy uses.
Seth H. Grae is president of Thorium Power, a Washington, D.C.-based company developing nuclear fuel designs to stop the production of weapons-grade plutonium and eliminate plutonium stockpiles. He holds a B.A. from Brandeis University, a J.D. from American University, and an L.L.M. and M.B.A. from Georgetown University.
THORIUM AT A GLANCE
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Name: Named for Thor, the Norse god of thunder.
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Atomic mass: 232.04.
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History: Discovered in 1828 by Swedish chemist Jöns Jakob Berzelius.
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Occurrence: Occurs naturally as a weakly radioactive isotope. Much of Earth's internal heat has been attributed to thorium and uranium decay. Thorium is primarily obtained from the minerals thorite and thorianite.
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Appearance: Silvery white metal.
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Behavior: Radioactive with a high melting point. The powdered metal is a fire hazard. The pure metal resists corrosion but tarnishes in air when contaminated with the oxide.
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Uses: Used in fabricating portable gas lamps, filament wire, breeder reactor fuel, gaslight mantles, and crucibles. Thorium oxide is used as a catalyst in the production of sulf uric acid and the conversion of ammonia to nitric acid.
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