The lanthanides or rare earths occupy a peculiar place in the periodic table. Being close to the bottom they seem to sit in its outer reaches, a region seldom reached by school chemistry courses, and indeed often even at undergraduate level. And their peculiarly similar chemical properties – at least in terms of classical chemistry – mean that there was a time when they were looked on with contempt, the great spectroscopist George Pimentel once writing in a textbook that “Lanthanum has only one important oxidation state in aqueous solution, the +3 state. With few exceptions, this tells the whole boring story about the other 14 Lanthanides” (1). Yet at the same time the lanthanides are technologically crucial, each element occupying a specific niche that no other can fill, and their importance has grown over the past five decades. To understand this perplexing situation – similarity of chemistry in spite of extraordinary individuality – we need to understand the lanthanide elements and what makes them tick.
The story of the rare earths goes back about 250 years to the chance discovery of a particularly dense rock by a young amateur geologist, Carl Axel Arrhenius, in a mine near the village of Ytterby (pronounced Ütterbü) on the island of Vaxholm, not far from the Swedish capital Stockholm. Puzzled to have found something that he was unable to identify, he sent a sample to the leading chemical analyst of the day, Johannes Gadolin, who worked at the University of Åbo (now Turku) in what is today Finland. After numerous experiments, Gadolin declared that the mineral must contain a hitherto unknown element, which he named Yttria in honour of the place of its discovery.
The initial report spawned a flurry of work on the dense minerals to be found in the east of Sweden. Over the course of almost 150 years, these minerals were found to contain no fewer than 15 new elements, many of which would be given names associated with the original discovery (scandium, ytterbium, yttrium, erbium, terbium, thulium, holmium, gadolinium). The difficulty lay in their separation, as the elements proved maddeningly similar in their properties. For example, all of them would form precipitates with the same reagents – fluoride or oxalic acid; all would dissolve in mineral acids such as nitric. The differences between them were so small that a complete separation might take as many 40,000 crystallization steps, and their pale colours added to the confusion. Yet even while chemists expressed general frustration at this unruly gaggle of elements, spectroscopy showed that each had a unique signature, absorbing light of very narrow ranges of wavelengths.
In the lanthanides, the normally accessible outer electrons that give most elements their distinctive chemistry (in this case the 4f electrons) are, for quantum mechanical reasons, buried deep inside the core of the atom. The result is that the radius of the atom contracts as you go from the light, early elements (La, Ce, Pr….) across to the heavier, later lanthanides (…Tm, Yb, Lu). It is these size differences that enable separation, although it remains exceptionally difficult. Furthermore, the chemical similarity of the rare earths means that they are found together in mineral deposits, which are typically enriched in either the LRE or the HRE fractions. The rare earths are by no means rare in the earth crust: their abundance significantly exceeds that of the noble metals such as gold and platinum, and even elements we take for granted such as iodine. The problem is not their abundance – it is their chemical similarity. The Manhattan project provided two solutions to the problem of separation. The first was the deployment at the Ames Lab in Iowa of ion exchange resins, porous beads of a polymer studded with negatively charges. On passing a rare earth solution onto a column of the resin, the ions would travel down the column at rates that depended on their size. This method was very quick, but being a batch process, it did not scale up well. A competing project at the Oak Ridge National Laboratory in Tennessee developed solvent extraction. By adding a binding agent, such as tributylphosphate (TBT), to a mixture of rare earths dissolved in nitric acid, the ions became slightly soluble in kerosene, to an extent that depended on the size of the ion. Thus, by flowing the two solutions past each other, it was possible to separate the ions from each other in large quantities. Today, solvent extraction is the commercial method for separating lanthanides. It is a technically demanding and often messy process that involves the use of large volumes of nitric acid and proprietary binding agents. Environmental and Health and Safety regulation has made separation expensive in the West; this has contributed to the technology moving to China, where costs are lower.
If the rare earths are chemically similar, their electronic structure is another matter. It is the 4f electrons that give the ions their remarkable optical and magnetic properties. The mutual interaction of these electrons (there will be between 1 and 13) give the energy levels of the lanthanides a ferocious complexity that tests modern spectroscopic and computational methods to their limit. But this very complexity and distinctiveness provides unparalleled opportunities for applications. The reactivity of the metal cerium toward oxygen has seen it used for well over a century in the alloy ferrocerium that is used in cigarette lighters. The oxide produced – ceria, CeO2 – has a hardness that makes it ideal for the polishing of glass, jewellery and lenses; it is known simply as “rouge”. Yet this oxide is also the ceramic that helps clean up the exhaust from our vehicles in catalytic converters. The magnetism of the neodymium ion helps to pin the electrons of the iron in the alloy neodymium-iron-boron, giving permanent magnets with an energy product 1000 times greater than in conventional ferrites. Rare earth magnets are to found in electric motors, regenerative brakes, and wind turbine generators. In biomedicine, the fact that the magnetism of the gadolinium ion affects the hydrogen nuclei of water has led to its use in contrast reagents for MRI imaging.
A specific optical transition of the erbium ion produces light that sits at the precise wavelength – 1.3 microns – where transmission of light through optical fibres is at a maximum.
Thus erbium is the element that drives much of the optical communications around the globe. The energy levels of Nd, Er, and other rare earths allow the construction of lasers by embedding the ions in a matrix of transparent aluminium/yttrium oxide. If you have a personalized iPad, the letters were probably engraved with one of these. And embedding rare earth ions into other ceramic matrices yields the phosphors for our television screens, our low energy fluorescent and LED lights.
These examples only scratch the surface of the myriad applications to which the rare earths are put. Today we can truly see the truth in the prophetic words of the nineteenth century chemist and spectroscopist Sir William Crookes who became obsessed with these elements: “These elements perplex us in our researches, baffle us in our speculations, and haunt us in our very dreams. They stretch like an unknown sea before us – mocking, mystifying, and murmuring strange revelations and possibilities” (2).
References 1. George C. Pimentel and Richard D. Spratley, Understanding Chemistry, Holden-Day, San Francisco, 1971, pp. 862-863. 2. Sir William Crookes, Address to the British Association, 1887.
ANDREA SELLA Department of Chemistry, University College London, 20 Gordon Street, London WC1H 0AJ E-mail: firstname.lastname@example.org
This article is based on a presentation by Andrea Sella at the ECG 2013 Distinguished Guest Lecture and Symposium held in the Science Room at Burlington House on Wednesday, March 20th 2013. The periodicity and some properties of the lanthanoids are shown on p 20.
Royal Society of Chemistry Environmental Chemistry Group