Endeavoring to identify the key features of the debate, Lavelle’s position largely related to electron configurations of the atoms. He argued that as atoms of neither lanthanum nor actinium possessed an electron in an f-orbital, instead possessing an electron in a d-orbital, then both elements needed to be considered as members of the d-block [6]:
However, placing lanthanum (La) and actinium (Ac) in the f-block is the only case where a pair of elements is placed in a group that results in their being part of a block with no outer electrons in common with that block.
This is shown in Table 4.2, together with the other relevant electron configurations.
The debate discussed in the previous section revolved largely around the relevance of placement of an element upon its ground-state electron configuration. For example, Laing noted that f-block member thorium has instead a ground-state configuration of [Rn]7s26d2 [12]. The unexpected electron occupancy of a 7p orbital for lawrencium (see Table 4.2) seemed to cause a particularly bitter exchange on the relevance of orbital occupancy for Group 3 membership. In conclusion, though certainly not in closure, to match the d-block electron configurations (Table 4.3), it would seem most logical to identify the Group 3 elements as Sc–Y–Lu–Lr.
Table 4.2 Electron configurations for the competing elements for membership of Group 3
Table 4.3 Comparative electron configurations for the Group 3 (inc. Lu and Lr), Group 4, and Group 5 elements
Commentary
In Chapter 5, the term “rare earth metals” will be introduced. By definition, this set of chemical elements includes scandium, yttrium, and all of the elements from lanthanum to lutetium inclusive. Thus, any decision of the membership of Group 3 does not affect the identity of the rare earth metals.
The membership of the f-group elements known as the lanthanoids will be affected, depending upon the definition of Group 3. There is no conflict if all 15 consecutive chemical elements are chosen, which have a common +3 ion charge and which are found in similar ores. Likewise, the actinoids will be a 15-member series. However, if the lower members of Group 3 are excluded from a “double membership” in the appropriate f-block series, then complexities will follow.
In this book, therefore, to avoid placing either of La–Ac or Lu–Lr in both Group 3 and in the f-series in an 18-column Periodic Table, the lower placements in Group 3 will be left empty. It is more important (in this Author’s view) to populate the f-series with 15 members and show the continuity of their properties.
References
1.L. S. Foster, “Why Not Modernize the Textbooks Also? I. The Periodic Table,” J. Chem. Educ. 16(9), 409–412 (1939).
2.W. B. Jensen, “The Position of Lanthanum [Actinium] and Lutetium [Lawrencium] in the Periodic Table,” J. Chem. Educ. 59(8), 634–636 (1982).
3.W. F. Luder, “The Atomic-Structure Chart of the Elements,” Can. Chem. Educ. 5(3), 13–16 (1970).
4.R. W. Clark and G. D. White, “The Flyleaf Periodic Table,” J. Chem. Educ. 85(4), 497 (2008).
5.W. B. Jensen, “The Periodic Table: Facts or Committees?” J. Chem. Educ. 85(11), 1491–1492 (2008).
6.L. Lavelle, “Lanthanum (La) and Actinium (Ac) Should Remain in the d-Block,” J. Chem. Educ. 85(11), 1482–1484 (2008).
7.L. Lavelle, “Response to ‘The Flyleaf Periodic Table’,” J. Chem. Educ. 85(11), 1491 (2008).
8.R. W. Clark, “Author of ‘The Flyleaf Periodic Table’ Responds,” J. Chem. Educ. 85(11), 1493 (2008).
9.W. B. Jensen, “Misapplying the Periodic Law,” J. Chem. Educ. 86(10), 1186 (2009).
10.L. Lavelle, “Response to Misapplying the Periodic Law,” J. Chem. Educ. 86(10), 1187 (2009).
11.P. C. L. Thorne and E. R. Roberts (translators), Fritz Ephraim’s Inorganic Chemistry (6th ed.), Oliver & Boyd, London (1954).
12.M. Laing, “More About the Periodic Table,” J. Chem. Educ. 86(10), 1189 (2009).
13.E. Scerri, “Which Elements Belong in Group 3?” J. Chem. Educ. 86(11), 1188 (2009).
14.G. Leigh, “Periodic Tables and IUPAC,” Chem. Int. 31(1), 1–2 (2009).
15.E. Scerri, “Mendeleev’s Periodic Table Is Finally Completed and What to Do About Group 3?” Chem. Int. 34(4), 1–5 (2012).
16.W. B. Jensen, “The Positions of Lanthanum (Actinium) and Lutetium (Lawrencium) in the Periodic Table: An Update,” Found. Chem. 17, 23–31 (2015).
17.E. Scerri (task group chair), “Which Elements Belong in Group 3 of the Periodic Table?” Chem. Int. 38(2), 22–23 (2016).
18.E. Scerri and W. Parsons, “What Elements Belong in Group 3 of the Periodic Table?” In E. Scerri and G. Restrepo (eds.), Mendeleev to Oganesson: A Multidisciplinary Perspective on the Periodic Table, Oxford University Press, Oxford, 140–151 (2018).
Chapter 5
Categorizations of the Elements
As will become apparent throughout this book, chemists can be very casual about the use or misuse of chemical terms. This lack of clear definitions even applies to the categorization of the elements themselves. In this chapter, a range of types of categorizations will be introduced and their meanings clarified.
Is polonium a metalloid? What is a weak metal? Which are the noble metals? So many questions, and to many of them, no established definitive answers. In this chapter, definitive proposals will be given as to which elements belong to which categories.
Nonmetals, Metals, and “In-Betweens”
In beginning chemistry, elements are classified as “metals” or “nonmetals.” This is a considerable oversimplification, as will be shown in the following. Each element is an individual, but it is convenient for chemists to impose systems of categorization. However, all terminology needs clear definition. Following from the definition, there should be clear reasoning behind assignment of an element to a particular category.
Nonmetals
The nonmetals are a “motley crew.” Essentially, they are every element that does not fit into another category. They include all the elements that are in Group 18 and Group 17, plus the top two elements of Group 16 and Group 15, the top element of Group 14, and hydrogen as can be seen in Figure 5.1. The only contentious element placed here in this category is astatine. Often identified as a metalloid or even a metal, the predominance of evidence is that astatine is a nonmetal [1, 2].
Figure 5.1 Elements classified as nonmetals.
Metalloids
The realization that there was no specific line of demarcation between metals and nonmetals dates back to the late 19th century (even though many commercial Periodic Tables continue to show one) [3]. It was in the 1890s that Newth declared that there were elements with intermediate properties: the metalloids. That was the easy part. The hard part was deciding which elements should be classified as metalloids.
The name “metalloid” was traditionally used for these in-between elements, then the term “semimetal” became preferred. However, the term “semimetal” was subsequently appropriated to be defined in terms of semiconductor materials in general, not simply chemical elements. As a result, “metalloid” has regained its meaning as specifically pertaining to certain chemical elements to the left of the nonmetals in the Periodic Table.
Vernon compiled all the elements identified as metalloids in sources from 1947 until 2012 [4]. The most popular elements cited with their frequency in parentheses were the following nine elements: antimony (88%); arsenic (100%); astatine (40%); boron (86%); germanium (96%); polonium (49%); selenium (23%); silicon (95%); and tellurium (98%). Adapting previously suggested criteria, Vernon devised the following definition for a metalloid:
A metalloid is a chemical element that, in its standard state, has (a) the electronic band structure of a semiconductor or a semimetal, (b) an intermediate first ionization potential (say, 750–1,000 kJ/mol), and (c) an intermediate electronegativity (1.9–2.2, revised Pauling).
According to Vernon, there were six elements only which fitted his criteria for metalloid classification: antimony, arsenic, boron, germanium, silicon, an
d tellurium.
Hawkes plotted the electrical conductivity (as log10 in S⋅m−1) of the proposed metalloids together with a selection of metals and nonmetals along a scale (Figure 5.2). He observed that there were five elements whose electrical conductivity fitted into the gap between metals and nonmetals [5]. These elements were arsenic, boron, germanium, silicon, and tellurium.
Figure 5.2 A scale of electrical conductivity indicating the placement of some elements (adapted from Ref. [5]).
Figure 5.3 The commonly accepted metalloids.
It is this Author’s preference to adopt the Hawkes list of five elements as metalloids (Figure 5.3). Antimony, which made Vernon’s list but not Hawkes’s list, would seem better accommodated in the category in the next subsection (in the following).
Though the categorization earlier is the one that will be adopted here, there is one other candidate for consideration as a metalloid: radon. It has been argued by Stein that radon behaves as an ionic cation in aqueous solution and therefore should be treated as metalloid [6]. An interesting proposal, but one that has not garnered any significant support.
Chemically Weak Metals
To reiterate the point, there are no rigid boundaries in properties across the Periodic Table. Just as we have “invented” an additional category of metalloids for those elements that have properties between metal and nonmetals, so there is now a need for a category between metalloids and “true” metals. A “true” metal has essentially cationic behavior, these metals can also be found as parts of polyatomic anions. A definition is:
There is a subgroup of the metals, the chemically weak metals (or amphoteric metals [7]) those closest to the metalloid borderline, that exhibit some chemical behavior more typical of the metalloids, particularly formation of anionic species in basic solution.
The nine elements in this category are aluminum, antimony, beryllium, bismuth, gallium, lead, polonium, tin, and zinc. As an example of anionic species, the pH dependency of zinc ion species is shown in Table 5.1 and compared with that of a “true” metal, magnesium.
Just as zinc ion in very basic solution forms soluble zincates, the other chemically weak metals similarly form aluminates, beryllates, gallates, stannates, plumbates, antimonates, bismuthates, and polonates. The weak metals are shown in Figure 5.4.
The term “chemically weak metals” defines this cluster of elements according to chemical criteria, differentiating them from “normal” metals. To confuse matters, the terms post-transition metal and poor metal are sometimes used in the literature. However, these categories refer to Periodic Table locations. That is, post-transition metal refers to all the metals of Groups 12 to 16, while poor metals are specifically the metals of the p-block elements (Groups 13 to 16). Also, Habashi devised a category of less typical metals that partially overlap with the category here of chemically weak metals [8].
Table 5.1 Variation of species with pH for aluminum and zinc ions
Figure 5.4 The chemically “weak” metals.
Metals
Though in the teaching of chemistry the focus is usually upon the p-block elements, in fact, about 80% of the naturally occurring elements are metals [9]. All the s-block, d-block, f-block, and the lower left part of the p-block are metals.
The old saying: “I know one when I see one” is often used as a criterion for a metal. As gold prospectors have found to their disappointment “It ain’t necessarily so.” Among several compounds that have a metallic luster, yellow metallic-looking iron(II) disulfide, FeS2, mineral name, pyrite, well deserves its appellation of “fool’s gold.” Chemists sometimes refer to a metal by a rather tautological argument as an element containing a metallic bond: that is, bonding throughout the crystal structure involving delocalized electrons [10, 11].
Sometimes metals are defined by a combination of properties, including ductility. Ductility is a measure of a material’s ability to undergo significant plastic deformation before rupture. Its opposite, brittleness, is defined as a material that breaks without significant plastic deformation. It is certainly true that some metals, such as gold and lead, are highly ductile, but then other metals, such as beryllium, manganese, uranium, and chromium are very brittle. Similarly, the usually associated term of malleability is true for some elements classed as metals, but not for others.
High three-dimensional electrical conductivity (thus excluding carbon as graphite) is possibly the best superficial indicator of a metallic element [12]. Hawkes has pointed out that under extremes of pressure, the atoms of most elements can be forced into close enough proximity to result in delocalized metallic bonding [13]. Thus in any definition of metals in terms of electrical conductivity, it is important to add “under ambient conditions” or “at SATP.” From the best electrical conductor (silver) to the worst (plutonium and manganese) among metals, one is looking at a factor of 102 in conductivity difference. Nevertheless, even the worst conducting metals exceed the electrical conductivity of nonmetals and metalloids by a factor of 105.
Another reason for stipulating ambient conditions is because the stable allotrope of tin below 13°C, gray α-tin, is nonelectrically conducting. On the other hand, under readily obtainable pressures, iodine becomes electrically conducting. A more specific physical criterion for a metal is the temperature dependence of the electrical conductivity. The conductivity of metals decreases with increasing temperature, whereas that of nonmetals increases.
Supermetals
Metals are commonly accepted as being hard (except mercury), dense, high-melting point, and generally unreactive. At the far left of the conventional Periodic Table, there are metals that are soft, low-melting point, low-density, and highly chemically reactive: the alkali metals. As the alkali metals are so chemically reactive, they deserve their own subcategory: the supermetals. If the emphasis is on the high chemical reactivity alone, should the category be broadened to include the three low-density, water-reactive alkaline earth metals? These positive attributes are contradicted by their high-melting points and hardness. The category of supermetals is therefore clearly delineated as being just the alkali metals.
Main Group Appellations
Numerous categorizations have been applied to the chemical elements. The names for Group 1 (alkali metals); Group 2 (alkaline earth metals); Group 17 (halogens); and Group 18 (noble gases) have been long accepted. It is curious that for Groups 1 and 2, the term “metals” is appended — using the metal/nonmetal categorization. At the other end, Group 18 is defined as “gases” using the solid/liquid/gas categorization.
For Group 18, at the time of the discovery of the noble gases (when they were more commonly called the “inert gases”), many chemists considered that the so-called “octet rule” precluded compound formation [14]. Now, at least 500 noble gas compounds have been characterized [15]. Adding to the irrelevancy of the “noble gas” name, the latest member, oganesson (element 118), is predicted to be a solid or a liquid at room temperature with a boiling point of between 50°C and 110°C [16]. However, unless very long-lived isotopes are synthesized, it is unlikely that the value can be experimentally confirmed. With ever more noble gas compounds being synthesized, including the exotic Na2He — actually (Na+)2He(e−)2 [17] — the term “noble” seems ever more inappropriate. As a result, the term aerogen is starting to be used. This term first appeared in print in a paper by Noyes, in which he gave credit for the name to Hembold of the University of Oregon [18]. It is now quite widely used in the contemporary literature (see, e.g., Ref. [19]).
Two more group names have become widely adopted. These are pnictogens for Group 15 and chalcogens for Group 16. The first proposed name for Group 16 was amphigens — based on this group’s ability to form both acidic and basic compounds — by Berzelius in the early 1800s [20]. The term “chalcogen” was invented by a member of Blitz’s research group at the University of Hannover, Germany, about 1932 [21]. Though as a group name, it should apply to all members of Group 16, chalcogen seems to have become commonly used for
all Group 16 except for oxygen [22]. Similarly, geochemists refer to chalcophile (“chalcogen-loving”) elements. These are the metals and heavier nonmetals that have a low affinity for oxygen and instead react with sulfur, the “heavier” chalcogen, to form insoluble sulfides [23].
The term pnictogen is more recent. The name was devised by van Arkel in the 1950s [24]. International Union of Pure and Applied Chemistry (IUPAC) originally rejected pnictogens as the name for Group 15. In 1970, it was pronounced by IUPAC that if group names were needed, they should be triels for Group 13, tetrels for Group 14, and pentels for Group 15 [25]. Fernelius et al. explained [26]:
There is strong, though not unanimous, sentiment within the IUPAC Commission on the Nomenclature of Inorganic Chemistry to adopt some systematic method for arriving at group names. The following has been suggested.
Family
B, Al, Ga, In, Tl — triels
C, Si, Ge, Sn, Pb — tetrels
N, P, As, Sb, Bi — pentels
Pentels never seemed to be adopted; instead pnictogens (mentioned earlier) became the commonly used term for Group 15 elements. Nevertheless, triels and tetrels have gained in popularity in recent years to describe the Group 13 and Group 14 elements, respectively [27], even in Chemistry International, IUPAC’s own magazine [28].
d-Block Metal Appellations
There have been numerous names associated with parts of the d-block elements. The issue of which are encompassed by the term “transition metals” (formerly the “transitional metals”) will be discussed in Chapter 8. Here are some selected examples that have a specific purpose.
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