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
