Periodic Table, The: Past, Present, And Future

by Geoffrey Rayner-canham;

  Tetrahedral Ions

  No, not those containing a tetrahedral bonding arrangement, but those species containing a tetrahedron of atoms. The yellow (not white [33]) allotrope of phosphorus, P4, provides the prototypical example. But just as cubane spawned pseudo-cubanes, so there are other molecules and ions adopting this very bond-strained tetrahedral shape. Some examples of these valence-isoelectronic ions are [Bi2Sn2]2−, [Sb2Pb2]2−, [Si4]4−, [Ge4]4−, [Sn4]4−, [Pb4]4−, and [Tl4]8−. Such species are examples of Zintl ions [34]. Zintl compounds are brittle, high-melting, intermetallic compounds, which contain polyatomic anions. First investigated in the 1930s, Zintl phases are formed by reacting a Group 1 or Group 2 metal with an element in any of Groups 13, 14, 15, or 16.

  Arsenic in Biological Systems

  One aspect often overlooked by inorganic chemists is that of substitution of one element for another in a biological organism or process. In this chapter, the substitution of one element by another in the same group will be the focus. For example, one bacterium can utilize arsenates instead of phosphates [35]. There has even been a computational study of whether such a substitution may be more widespread among bacteria [36].

  Group 16 (Chalcogens)

  With Group 16, there is the progression from nonmetals, oxygen and sulfur (not “sulphur” [37]); to metalloid, selenium; then to the two weak metals, tellurium and (radioactive) polonium. The allotropes of oxygen used to be dismissed as simply dioxygen and trioxygen (ozone). But no more. Tetraoxygen, O4, exists at very low pressures in the upper atmosphere [38]; while octaoxygen, O8, is formed as a dark red solid cubic structure under high pressure, low temperature [39].

  Sulfur on Io

  Sulfur has more allotropes than any other element [40] and even when it melts upon heating the chemistry continues to be complex [41]. Yet the most interesting sulfur chemistry is not on Earth, but on Jupiter’s moon, Io. Most people have seen NASA photos of the startingly multicolored moon, unique in the Solar System, which even possesses a sulfur lake [42]. One molecule identified in the atmosphere is disulfur oxide, S2O [43]. The Io surface colors — yellows, reds, blues, and greens — are almost certainly sulfur species. But what are they? Two possible contributors under the surface bombardment by particles from Jupiter’s radiation field are the cations: [S8]2+ (blue) and [S16]2+ (red) [44], but no one knows for sure at the date of writing.

  Selenium and Tellurium in Biological Systems

  Though sulfur-containing amino acids, such as cysteine, are well known, it is little appreciated that the next lower member of the Group, selenium, is also incorporated into biological systems. In fact, seleno-amino acids, specifically seleno-cysteine and seleno-methionine, have essential roles in living organisms, different than those of the sulfur analogues [45]. A phenomenon first observed in 1880, some plants even have a dependence on selenium-rich soils [46]. Descending Group 16 involves crossing from nonmetal to metalloid (see Chapter 5). Despite this change in element properties, there are indeed telluro-cysteine and telluro-methionine [47]. In fact, a certain fungus fed tellurium rather than sulfur quite happily synthesizes the tellurium analogues. Interestingly, a study has compared and contrasted the properties of seleno-cystine and telluro-cystine and their derivatives [48].

  Group 17 (Halogens)

  After the alkali metals, the halogens provide the second of the Groups that have a unitary classification, this time, of the nonmetals. These highly reactive nonmetals span the range from gases (pale yellow for fluorine, pale green for chlorine) to liquid (almost black, “oily” liquid with red-brown vapor for bromine) to solid (purple-black, metallic-looking solid that melts to a deep violet liquid then boils to a violet gas for iodine).

  Weakness of the Fluorine–Fluorine Bond

  There is one specific Group 17 feature that affects chemistry throughout the Periodic Table: the weakness of the F−F bond (see Table 7.3). The weakness of the F−F bond has to be contrasted with the strength of the (highly polar) bonds of fluorine to other elements. As an example, the F−C bond can be as strong as 544 kJ⋅mol−1. This bond energy difference provides a key factor in the ready formation of fluorocompounds, particularly those in high oxidation states.

  Table 7.3 Bond energies for the dihalogen molecules

  Dihalogen Molecule Bond Energy (kJ⋅mol−1)

  F−F 155

  Cl−Cl 240

  Br−Br 190

  I–I 149

  Interhalogen Compounds

  If there is one feature unique to the halogens, it must be the readiness to form interhalogen compounds (and inter-pseudo-halogen compounds, see Chapter 14). An interhalogen compound contains two different halogen atoms and no atoms of elements from any other group. The formulas are XYn, where n = 1, 3, 5, or 7, and X is the less electronegative of the two halogens. For n = 7, only iodine heptafluoride is known [49]. The halogen pairs also form polyatomic cations and anions that are favorite species for quizzes in general chemistry courses in which molecular geometries have to be deduced using Gillespie–Nyholm (VSEPR) Theory [50]. One fascinating aspect of the interhalogen compounds is the intermediate physical properties to those of the constituent halogens. As examples, chlorine monobromide is a yellow-brown gas at room temperature, while bromine trifluoride is a yellow-green liquid.

  Group 18 (Aerogens)

  The aerogens/noble gases (Group 18) used to be hailed as exemplars of periodicity with the systematic trends in melting and boiling points. But no more. There seems to be nothing systematic — no “patterns and trends” — about their chemical properties.

  Some Xenon Compounds

  It is still a common belief that aerogen chemistry is limited to bonds with fluorine or oxygen. Here the focus will be on compounds with other elements, to make the point of the now known diversity of aerogen — particularly xenon — chemistry. Xenon is still by far the most chemistry-rich member of the Group [51]. To stabilize bonds between xenon and less electronegative elements, electron-withdrawing groups on the bonded species are required. These species are most usually fluorine substituted [52]. The prototypical example is the pentafluorophenylxenon(II), [(C6F5) Xe]+, ion [53]. This cation is prepared similarly to that of [(C6F5)I]2+. The electron-withdrawing power of the pentafluorophenyl group is so strong that even chloride ion can be induced to complete the two coordination to form C6F5XeCl [54]. However, the most interesting ion including xenon has to be its coordination as a ligand to gold, [AuXe4]2+ [55] (previously mentioned in Chapter 5).

  A Selection of Compounds of Other Aerogens

  There seem to be some similarities with krypton in that krypton forms an analogous species to xenon, that of C6F5KrCl [56]. However, at low temperatures, argon seems to form some unique compounds, such as the now well-established HArF [57]. For highly radioactive radon, there is currently only RnF2 and RnO3, which illustrates the common feature that oxides can often be in a higher oxidation state than fluorine [58]. No “real” compounds of helium and neon have been synthesized to the date of writing.

  Periodic Trends

  The systematic progressions of formulas of hydrides and oxides across each Period was one of the crucial factors in Mendeléev’s development of the Periodic Table (see Figure 7.4).

  Such a progression is still a fundamental basis of why the Periodic Table is still important today. The definition of trends across a period is best stated as:

  Periodic properties are those systematic patterns observed across a Period. Such patterns are commonly trends in chemical formulas of the compounds formed by the elements.

  Figure 7.4 The top part of a Periodic Table published by Mendeléev in 1871 (note that superscripts, not subscripts, were used to identify atom ratios).

  Bonding Trends in Main Group Highest Oxidation-State Oxides

  It seems therefore appropriate to conclude this chapter with a more detailed examination of periodicity in the formulas of the main group oxides. Table 7.4, as did Mendeléev’s table, only displays the highest oxidation-state oxides. The formulas of
the highest (oxidation state) oxides correlate with the group number of the nonoxygen element; that is, +1 (Group 1), +2 (Group 2), +3 (Group 13), +4 (Group 14), +5 (Group 15), +6 (Group 16), and +7 (Group 17). The one exception is oxygen difluoride, the only oxide in which the other element has a higher electronegativity than oxygen.

  Though there is a smooth progression in formulas (except for the halides), this hides sudden breaks in phase and behavior at room temperature. These changes can be related to the progression of bonding types from ionic to covalent, with the network covalent region marking the borderline between the two bonding categories [59].

  The location of the network covalent species shifts diagonally from the 2nd to the 3rd Period. The common explanation is that the location is reflective of the electronegativity, which itself crosses the Periodic Table on a diagonal. The corresponding network covalent molecule in the 4th Period is beneath that in the 3rd, perhaps a reflection of the similarity resulting from the d-block contraction.

  As is not uncommon in inorganic chemistry, things do not always fit neat patterns. Specifically, at room temperature, “N2O5” has an ionic structure: [NO2]+[NO3]−. The two lower members of Group 16 do not form simple XO3 molecules, instead, “sulfur trioxide” is a trimer, S3O9, containing alternating sulfur and oxygen atoms to form a six-membered ring while “selenium trioxide” is an analogous tetramer, Se4O12. As another “breakdown” of pure periodicity, though chlorine forms a heptaoxide, bromine only forms a pentaoxide, a fact which might be ascribed to the 4th Period anomaly.

  Table 7.4 Bonding categories for the 2nd, 3rd, and 4th Period highest oxidation-state oxides

  Commentary

  Periodic patterns and trends are the fundamental basis of the Periodic Table. And it is not all about Groups and Periods, as Rogers has pointed out [60]. However, in this Author’s view, sometimes periodicity is raised to almost mythical status in which patterns and trends are “cherry-picked” to illustrate near-perfect sequences. As shown in this chapter, there are many species that “stubbornly” refuse to fit how they “should.” As chemists, we should not be afraid to teach the limitations of periodicity and sometimes revel in the uniqueness of each element and its compounds.

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