In Chapter 9, the several similarities of scandium to aluminum were discussed in the context of the (n) and (n + 10) relationship. The major use for scandium is in specialty alloys of aluminum, the Al3Sc micrograins imparting additional strength to the aluminum metal [4]. In fact, in some respects, scandium is closer in chemistry to aluminum than to any other element.
The 4f Elements
When the first rare earth elements were discovered, the question arose as to where they could be fitted in the eight-group Periodic Table. At the time, it was believed that the eight-group framework was the key “set in stone” feature of the Periodic Table: the elements themselves were the problem. It was not until 1882 that Bailey concluded that the known rare earth elements were not members of any of the eight Groups [5]. Then Bassett, in 1892, proposed that the lanthanoid elements (as we now call them) form their own series (and that the known actinoids formed a matching series). The key to the understanding of the lanthanoids came in 1921, when Bury postulated that they corresponded to the filling of the 4f orbitals.
Progressing to the present, what of their chemistry? Pimentel and Sprately summed up the opinion of most chemists to the chemistry of the 4f elements [6]:
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 lanthanides.
Yet the predominance of the +3 state is one of the very interesting things about the 4f elements. Nowhere else in the Periodic Table is it possible to study a sequential series of elements all in the same oxidation state. And there are many other interesting aspects to their chemistry as will be covered in the following.
Books specifically on the 4f elements (or the “rare earth” elements, which also encompass Group 3 elements — see Chapter 4) are very rare. One volume was part of a series on each segment of the Periodic Table [7]. An undergraduate textbook claiming to be on the d-block and f-block elements contained very little on the f-block, instead being almost entirely on the d-block [8]. Though the text by Cotton (mentioned earlier) [1] was comprehensive, it did not look beyond the confines of the f-block for relationships, nor did the similar book by Aspinall [9]. There has also been a compilation of studies of the coordination chemistry of the rare earth elements [10].
Yet in the research world, lanthanoids have been a burgeoning field. The metals and their alloys have become indispensable as magnetic materials [11]. In the 1990s, the primary interest in their compounds was as reagents in organic synthesis [12, 13]. Now the focus has become on the luminescent properties of the lanthanoid ions and their applications [14–16]. There is also an interest in a group of extremophile aerobic methanotrophs that require one of the early lanthanoids (lanthanum, cerium, praseodymium, or neodymium) for their metabolic pathways [17].
The Lanthanoids
To begin, as is recognized by the International Union of Pure and Applied Chemistry (IUPAC), the correct term is “lanthanoids” [18]. The ending “-ide” is accepted throughout chemistry nomenclature as referring to a negative ion, as in “oxide” or “sulfide.” The commonly accepted definition of a lanthanoid is therefore:
Figure 12.1 The lanthanoid elements as defined in this chapter.
A lanthanoid is any of the series of 15 consecutive chemical elements in the Periodic Table from lanthanum to lutetium (Figure 12.1).
Properties of the Elements
There is the inference that, for a particular property of the lanthanoid elements, a linear or smooth curve plot results. This is not necessarily true. Cater showed that for several parameters of high-temperature lanthanoid chemistry, the plots are much more uneven [19]. This discovery was revisited by Johnson, who came to the following conclusion [20]:
… the lanthanide elements behave similarly in reactions in which the 4f electrons are conserved, and very differently in reactions in which the number of 4f electrons changes.
Laing showed there were significant deviations from linearity for lanthanoid melting points. Though there is a general trend of increasing melting points of the lanthanoids with increasing atomic number, there are two exceptions to the rule: europium and ytterbium, which both have melting points well below that of the trend. He ascribed these anomalies to much weaker metallic bonds for these two elements [21].
In addition, Laing noted that the densities of the lanthanoid metals followed an even more linear relationship (see Figure 12.2), though again, with the exception of europium and ytterbium [21]. He accounted for these two deviations in terms of the electron-sea model of metallic bonding. For all other lanthanoids, the intermetallic forces involved 3+ ions and the intervening three “roaming” valence electrons. Laing argued that as europium and ytterbium favored the 2+ state (see in the following), then the intermetallic forces between the (theoretical) 2+ ions and two “roaming” electrons would be significantly less.
Figure 12.2 Densities of the elements from atomic number 56 to 72 (adapted from Ref. [20]).
However, though the electron-sea model can provide a very simplistic idea of metal behavior, it is incapable of being given any quantitative or even semiquantitative validity. Without going into the sophistication of band theory, the alternative is the soft-sphere model. Lang (not to be confused with Laing) has applied the soft-sphere model to the lanthanoid metals and showed that it is a good fit [22]. Nevertheless, it is more of a justification than an explanation, in that the calculation requires knowledge and use of the crystal packing factor. It does not explain why europium uniquely has the body-centered cubic packing rather than the more compact (and therefore denser) packing arrangements of the other lanthanoids.
Ion Charges of the Lanthanoids
The 3+ state predominates for all of the lanthanoids [23]. It is the 3+ oxidation state consistency that gives such a useful comparison across the series.
The Lanthanoid Contraction
Of importance is the 3+ ionic radii of the lanthanoids. As can be seen from Figure 12.3, the ionic radius decreases, almost linearly, from lanthanum to lutetium. This decrease is known as the lanthanoid contraction (or more commonly as the “lanthanide contraction”) [24]. The effect was first recognized and named by the Norwegian geochemist, Victor Goldschmidt. The contraction, defined in the following, is an important aspect of lanthanoid and, even postlanthanoid chemistry [25].
The lanthanoid contraction is the greater-than-expected decrease in ionic radii of the elements from lanthanum to lutetium, which results in smaller than otherwise expected ionic radii for the subsequent elements commencing with hafnium.
The lanthanoid contraction can be explained as follows [26]. It is the inner, filled, 5s25p6 electron “layer” that defines the ionic radius. The 4f electrons contribute little to the shielding. Thus, as the nuclear charge increases, there is a contraction of the 5s25p6 orbitals causing a radius reduction of the ions.
Figure 12.3 The lanthanoid contraction for the 3+ ions.
It is a common assumption that, having given it a special name, the lanthanoid contraction is greater than other contractions across periods. This is not true, as Lloyd has pointed out [27]. In fact, the decrease in radius from lanthanum(3+) to lutetium(3+) of 117 pm to 100 pm is less than that from calcium(2+) to zinc(2+) of 114 pm to 88 pm. From one ion to its neighbor, the average individual lanthanoid contraction is also less than that from scandium(3+) to gallium(3+) of 89 pm to 76 pm.
Post-Lanthanoid Effect of the Atom and Ion Contraction
As described in Chapter 8, a crucial consequence of the 14-element contraction is that the 6th Period transition metal series are of almost the same atom and ionic radius as their 5th Period analogue. As an example, the pair of zirconium and hafnium can be considered. Comparing these two elements, the number of protons (and number of electrons) has increased from 40 to 78. Yet the atomic radius decreases from 159 pm to 156 pm. Similarly, there is a small decrease, not an increase, in ionic radii from 86 pm (Zr4+) to 85 pm (Hf4+).
Had the lanthanoids not intervened, it i
s possible to make a rough estimate of the Hf4+ ion radius. This can be done by comparing the ionic radius of the 3+ ion preceding the lanthanoids, lanthanum (117 pm) with the element above it, yttrium (104 pm). Using approximately the same difference, without the lanthanoid contraction, Hf4+ ion would be expected to be about 10 pm larger than zirconium, instead of 1 pm smaller.
Other Lanthanoid Oxidation States
It is the variations in oxidation states from the +3 “norm,” which have provided the most interest [26]. When plotting out known oxidation states for elements, the question arises as to how extreme the conditions, or how unusual the ligands, that have been used in order to stabilize a specific oxidation state. Table 12.1 identifies those oxidation states that have a significant existence for simple compounds. In contradiction to this generality, praseodymium(V) is included (in parenthesis) as it is of intrinsic interest in the context of 4f0 configurations. It can be seen that the empty; half-filled; and filled f electron energy state plays a significant — but not exclusive role — in determining which other oxidation state(s) are feasible for a specific lanthanoid.
The 4f0 oxidation state is the expected state for lanthanum. For cerium, Ce(IV) is a common state, though highly oxidizing. In the next section, we will see that cerium(IV) has many similarities to members of Group 4. The oxidation of cerium(III) to cerium(IV) has relevance to geochemistry. In oxidizing waters, cerium is deposited as insoluble cerium(IV) oxide. This is one parameter by which the redox condition of ancient seas and oceans can be determined [29]. The cerium enrichment (as cerium(IV) oxide) compared with the other lanthanoids is known as the “positive” cerium anomaly.
Table 12.1 The 4f electron configurations corresponding to the common ion charges (adapted from Ref. [26])
Though an empty f shell would seem to be an obvious possibility for praseodymium, it seems that Pr(V) is, in fact, not a particularly favored oxidation state for the element. The only species obtained by the date of writing has been the ion [PrO2]+ under very low-temperature noble gas matrix isolation [28].
Several of the lanthanoids exhibit the +2 oxidation state [30], but it is only for europium and ytterbium that the +2 state is of major importance. In addition, terbium only “reluctantly” forms compounds in the +4 oxidation state [31], which is surprising considering it has a significantly lower 4th ionization energy [32].
As can be seen from Table 12.1, the progression: lanthanum(III); cerium(IV); and praseodymium(V) corresponds to the “empty-shell” 4f0 series, which would not be unexpected [33]. Similarly, ytterbium(II) and lutetium(III) correspond to the “full-shell” 4f14 configurations. The third isoelectronic set, europium(II); gadolinium(III); and terbium(IV) correspond to the “half-filled” 4f7 electron configuration. This so-called “stability of the half-filled shell” (see Chapter 2) is often discussed in the context of the main group elements [34] but it is also evident for the lanthanoids.
Restructuring the Lanthanoids
Though conventionally the lanthanoids are treated as a single continuous unit, attempts have been made to identify subcategories and possible rearrangements.
Figure 12.4 Cerium as a member of Group 4 in addition to the lanthanoids (modified from Ref. [35]).
Cerium as a Member of Group 4
Johansson et al. have singled out cerium on the basis of its +4 oxidation state to be better considered as a member of Group 4 [35]. This assignment is shown in Figure 12.4.
The Stacked Lanthanoid Arrangement
In geochemistry, the rare earth elements are classified as “light” or “heavy.” Of the lanthanoids, lanthanum to gadolinium are usually considered as “light” while terbium to lutetium are usually assigned as “heavy.” More of the “light” lanthanoids are in the Earth’s crust, while more of the “heavy” lanthanoids in the Earth’s mantle. This distinction comes about through the variation in ionic radii and hence the crystal structures in which the ion will fit. The trend of ionic radii results from the lanthanide contraction described earlier.
It was Ternström in 1976, who first proposed “stacking” the two half series on the basis of physical and chemical properties as: Ce–Gd and Tb–Lu [36]. Laing elaborated upon this concept, focusing upon chemical resemblances [21]. He noted (as stated earlier) that both europium and ytterbium form compounds in which they have a +2 oxidation state and therefore related to Group 2. Similarly, cerium commonly forms compounds in which it has the +4 oxidation state, and therefore should be associated with Group 4. Laing placed these three elements in their assigned groups, plus intervening lanthanum, gadolinium, and lutetium in Group 3. The other lanthanoids were then placed in order to complete each of the two subrows (Figure 12.5).
Though the early members fit, there is no evidence so far of any higher oxidation states for the later members of the lanthanoids. Laing subsequently changed his mind about the arrangement of the lanthanoids [37]. Instead, he devised a three-level sandwich that highlighted the +2 trio of [Ba–Eu–Yb], the +3 trio of [La–Gd–Lu], and the +4 trio of [Ce–Tb–Hf], with gadolinium being central, as shown in Figure 12.6.
Figure 12.5 The two-row lanthanoids according to Laing, with the surrounding elements shaded (modified from Ref. [21]).
Figure 12.6 Laing’s gadolinium-centered lanthanoid series (from Ref. [37]).
Dendrogram Restructuring
Horovitz and Sârbu used a cluster analysis to develop a similar double-row set [38]. The dendrogram that relied largely upon a variety of numerical values for each atom/ ion is shown here (Figure 12.7).
From the structure of the dendrogram, Horovitz and Sârbut devised a segment of the Periodic Table. This arrangement (Figure 12.8) largely matches Laing’s two-row lanthanoid pattern shown in Figure 12.5 (though Laing contested the differences [39]), except that the Eu–Yb pair are shown at the right-hand, not left-hand, end. Noticeably, the two-row sequence is fragmented.
Figure 12.7 A dendrogram for the lanthanoids (adapted from Ref. [38]).
Figure 12.8 Lanthanoids arranged according to the cluster analysis of Horovitz and Sârbut (adapted from Ref. [38]).
The clustering of samarium with europium and ytterbium is particularly significant. Together with europium and ytterbium, samarium is the only other lanthanoid to form stable 2+ compounds. In fact, samarium(II) iodide (Kagan’s reagent) is an important mild reducing agent in organic chemistry [40].
The gap between promethium and samarium according to the cluster analysis is interesting in that it seems to be on the borderline of some ion-packing arrangements. For the lanthanoid(III) fluorides, from lanthanum to promethium, the crystal structure is based upon the nine-coordinate lanthanum(III) fluoride tricapped trigonal prism arrangement. Then from samarium to lutetium, the crystal structure is, by contrast, based upon the eight-coordinate yttrium(III) fluoride bicapped trigonal prism. Similarly, for the lanthanoid(III) iodides, from lanthanum to promethium, the crystal structure is based upon the eight-coordinate plutonium(III) bromide bicapped trigonal prism. Then from samarium to lutetium, the crystal structure is based upon the six-coordinate iron(III) chloride octahedral structure.
External Lanthanoid Relationships
Surprisingly, all the lanthanoid(III) (and yttrium(III)) ions have been shown to be essential cofactors in certain enzymes in methylotrophic bacteria. They bind 108 times more strongly to the ligand sites than calcium, which is the usual metal ion in such enzymes [41]. Nevertheless, it is with the two “unusual” oxidation states of lanthanoids for which there are interesting similarities with other elements.
Similarities of Europium(II) with Strontium (and Calcium)
The europium(II) ion behaves very similarly to an alkaline earth ion; for example, its carbonate, sulfate, and chromate are insoluble, as are those of the heavier alkaline earth metals. The ionic radius of europium(II) is actually very similar to that of strontium, and, as might be expected, several europium(II) and strontium compounds are isostructural.
In the study of lanthanoid-containing minerals, the pro
portion of europium can be significantly different from that of the other lanthanoids [42]. In anaerobic geothermal conditions, europium(III) can be reduced to europium(II). Then, though significantly larger than the calcium ion, europium(II) can replace calcium in minerals. This ion exchange process is known as the europium anomaly and it is said to be “positive” if europium is enriched with respect to the other lanthanoids and “negative” if it is depleted.
Similarities of Cerium(IV) with Zirconium(IV) and Hafnium(IV)
Whereas europium has a lower than normal oxidation state, cerium has a higher than normal oxidation state of +4. Cerium(IV) behaves like zirconium(IV) and hafnium(IV) of Group 4. For example, all three of these ions form insoluble fluorides and phosphates. The similarity can be seen from a comparison of acid–base behavior under strongly oxidizing conditions (Table 12.2).
The ready oxidation of cerium(III) to cerium(IV) has significant geochemical implications. For example, the similar ion sizes of cerium(IV) and zirconium(IV) resulted in incorporation of Ce4+ ions into zircons, key minerals in the context of the earliest Earth’s rocks [43].
Periodic Table, The: Past, Present, And Future Page 18
