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
