Periodic Table, The: Past, Present, And Future
Page 10
Main Group Elements
In the 18-Group or 32-Group Periodic Tables, the unfortunate main group elements are fissioned into two widely separated halves: Groups 1 and 2, and Groups 13 to 18. Yet there is no wide gap between beryllium and boron or between magnesium and aluminum. This is an artifact of our obsession with linearity following electron configurations rather than producing a design that has chemical pedagogical value.
Sanderson produced a design to solve this (and other) issues, which was updated by Jensen. Jensen commented [5]:
Despite its extraordinary advantages, Sanderson’s double-appendix table has seen virtually no use beyond his own writings. It is unclear whether this is due to resistance on the part of authors and publishers, who fear that any departure from the norm will diminish the sale of their textbooks, or to the fact that the use of the periodic table to correlate the facts of descriptive chemistry is so superficial in most textbooks that the very real limitations of the 18-column block table never become apparent.
A very specific advantage from this Author’s perspective is that zinc, cadmium, and mercury are unambiguously placed among the main group elements [5]. The Sanderson Periodic Table is shown in Figure 7.1.
Figure 7.1 Sanderson’s double-appendix Periodic Table design (from Ref. [5]).
Historical Background
It was the concept of repetitive properties that gave rise first to the proposal for triads of elements by Döbereiner, then the Law of Octaves by Newlands, followed by the periodic patterns of Mendeléev [6] and of Meyer [7]. As is apparent from his version of the Periodic Table of 1871 (Figure 7.2), Mendeléev was guided in part by the formulas of the oxides and the hydrides.
Figure 7.2 Mendeléev’s Periodic Table of 1871.
The Uniqueness Principle
As can be seen from Figure 7.2, and of relevance to this chapter, Mendeléev placed an underline beneath the 2nd Period. This underscore was to indicate a degree of difference between these elements and the ones beneath. It was apparent to him that the elements above the line did not wholly resemble the lower members of each Group.
This difference of the topmost member of each Group is known as the Uniqueness Principle. Rogers has identified three factors that cause significantly different behavior among most of the 2nd Period elements. Summarizing his statements, it can be stated that [8]:
•They have exceptionally small atomic radii
•They exhibit a maximum of four bonding directions
•The nonmetallic elements have an enhanced ability to form multiple (π) bonds
The theoretical underpinning of the Uniqueness Principle has been described by Kaupp [9]:
The similarity of the radial extent of 2s- and 2p-shells is a decisive factor that determines the special role of the 2p elements within the p-block . . . The overall small size of the 2p-elements can be appreciated from any tabulation of atomic, ionic, and covalent radii. . . . This of course does lead to overall coordination number preferences of the 2p-element, as is well known. The smaller radii and the high electronegativities of the 2p-elements are behind their strict obedience of the octet rule, in contrast to the apparent behavior of their heavier homologues . . .
The Uniqueness Principle can be illustrated by the formulas of the highest oxidation-state simple oxo-anions for the Group 14 and Group 15 elements. The top member has a preference for the delocalized trigonal-planar π-bonding system (Table 7.1), while the lower tetrahedrally coordinated tetra-oxo-anions are prone to polymerization.
Main Group Organometallic Compounds
Until the latter part of the 20th century, it was assumed that formation of organometallic compounds, particularly those with metal–carbon bonds, was an almost exclusive domain of the transition metals. This belief is no longer held. As Power has described [10]:
The new compounds that were synthesized highlighted the fundamental differences between their [the heavier main-group elements] electronic properties and those of the lighter elements to a degree which was not previously apparent. This has lead to new structural and bonding insights as well as a gradually increasing realization that the chemistry of the heavier main-group elements ever more resembles that of transition metal complexes than that of their lighter main-group congeners. The similarity is underlined by recent work, which has shown that many of the new compounds react with small molecules such as H2, NH3, C2H4 or CO under mild conditions and display potential for applications in catalysis.
Table 7.1 Comparison of the formulas of the simple highest oxidation-state oxo-anions
The d-Block Contraction
Just as the Uniqueness Principle served to differentiate the chemistry of the 2nd Period elements from those in the subsequent periods, so the d-block contraction (sometimes called the scandide contraction) results in a greater similarity between some of the members of the 3rd Period and 4th Period [11].
The contraction is not actually a “contraction,” instead, the expected increase in effective ionic radii upon descending a p-block element is less when passing from the 3rd Period to the 4th Period than would be expected for a systematic trend. Table 7.2 shows that the increase in radius from Al3+ to Ga3+ is only 8.5 pm. For comparison, the Group 3 analogue, Sc3+, which like Al3+ has a noble gas ion configuration has a radius 21 pm greater. Similarly, the ionic radius of In3+ is 10 pm less than that of Y3+.
The accepted explanation for the d-block contraction is that the 3d electrons are poor shielders. Thus the increased effective nuclear charge results in an orbital contraction for gallium and its 3+ ion. The Uniqueness Principle is used to account for the exceptional differences between the 2nd Period and 3rd Period elements of the same Group. Relatedly, the d-block contraction is used to explain exceptional similarities between the 3rd Period and 4th Period elements of the same Group.
Table 7.2 A comparison of the ionic radii of the Group 3 and Group 13 ions
The 4th Period Anomaly
Following from the d-block contraction is the 4th Period Anomaly:
The 4th Period anomaly for the p-block elements is where the properties of the Group member of the 4th Period do not fit the trend for the other members of the Group.
Pyykkö [12] has referred to the phenomenon as secondary periodicity, a pattern first reported by Biron in 1915. According to Biron, descending a Group, many physical and chemical properties exhibit an alternating “sawtooth” pattern. Hildebrand rediscovered this phenomenon in the context of the Group 15 elements in 1941 [13]. He noted that the chemistry of nitrogen, arsenic, and bismuth was more focused toward the +3 oxidation state. By contrast, the chemistry of phosphorus and antimony revolved more around the +5 oxidation state.
Dasent wrote an article on compounds that a chemist would expect should exist, but which had not been synthesized at that time. He categorized the probable reasons for nonexistent compounds; category 3 being that of the 4th Period anomaly [14]:
. . . those whose instability is related to the reluctance of certain atoms of the first long period to assume their highest oxidation state.
It was Sanderson who provided a fuller account of the unique features that he found for 4th Period p-block elements, confirming the existence of this anomaly [15]. As exemplars of the 4th Period anomaly, he cited the difficulty in preparing AsCl5 (yet PCl5 and SbCl5 are well known) [16] and HBrO4 (yet HClO4 and H5IO6 are well known) [17].
In explanation, Pyykkö found that there were two different factors. The 4th Period anomaly resulted from the d-block contraction while the lesser 5th Period factor was a result of the combination of relativistic effects (see Chapter 2) and the lanthanoid contraction (see Chapter 12).
Group Trends
In textbook discussions of periodicity, trends in properties within each main group is the most common. An appropriate definition is:
Groups trends are the systematic variation of properties of elements and their compounds descending a specific group. Exceptions to such trends are usually indicative of a change in bonding type.
> The large majority of chemical elements are high-melting, unreactive metals. Were the Periodic Table completely filled with these elements, few young people would rush to become chemists! The fascination comes from the s-block and p-block elements where the curious student encounters exotica: highly reactive metals; a yellow solid; a green gas; and so on. This is the diverse world of the main group elements.
It is with the main groups that we have real groups — each with five or six elements for which patterns and trends can truly be traced. Yet each element is unique. Any discussion therefore needs a blend of pattern and individuality. Here such a blend will be attempted. To do so, the essential features of each element is provided. Why is this necessary? Many/most inorganic texts seem devoid of any sense that chemistry is anything other than theory and calculation. It is now 50 years since Davenport bemoaned the abandonment of the joy of inorganic chemistry [18]:
. . . the typical senior inorganic course leans heavily on theory, particularly bonding theory. Since so many of their teachers are children of the fabled Renaissance of Inorganic Chemistry (surely reports of its implied death were greatly exaggerated?) this is not surprising . . . — is it wise?
Group 1 (Alkali Metals)
Group 1, solely consisting of metals, is one of the few groups to actually show systematic changes descending the series. In this case, for example, chemical reactivity and density increases down the Group.
Sodium, the Second Atypical Alkali Metal
The “abnormality” of lithium is commonly discussed, yet the difference of sodium, also, from the heavier alkali metals is often overlooked [19]. All of potassium, rubidium, and cesium form dioxides(−1), that is, MO2. Instead, sodium forms a dioxide(−2), Na2O2. As another example of the difference, potassium, rubidium, and cesium form triiodide(−en1) compounds, MI3, whereas lithium and sodium do not.
Zmaczynski [19] has pointed out that sodium compounds with di- and trinegative anions tend to be highly hydrated, such as Na2SO4⋅10H2O, Na2CO3⋅10H2O, and Na2HPO4⋅12H2O. By contrast, the potassium (and rubidium and cesium) compounds are all anhydrous: K2SO4, K2CO3, and K2HPO4.
So significant are the differences of lithium and sodium from the heavier alkali metals, that Smith, in his classic 1917 text, Introduction to Inorganic Chemistry, discussed lithium and sodium separately from potassium, rubidium, and cesium [20]. A century later, in 2018, a review article by Restrepo of phenomenological studies included two relevant fragmented Periodic Tables. These Tables, one by Restrepo et al. and the other by Leal et al., show lithium and sodium as a separate unit in chemical behavior from the lower three heavier alkali metals [21].
Naked Radii and Hydrated Radii
As chemists, the term “ionic radius” is very clearly defined. Data tables list the values. In the world of biochemistry, the value is larger and fluid. The hydrated ionic radius of an ion is significantly larger than that of the “naked” ion. And it is the inverse order for the alkali metal ion. This results in a free hydrated sodium ion radius of 276 pm compared with 116 pm for the naked ion, while the hydrated radius for potassium is 232 pm compared with 152 for the naked ion. The reason for this can be explained in terms of charge densities. The charge density of the sodium ion is about twice that of the potassium ion. That is, the sodium ion will attract more polar water molecules to it in hydration shells than will potassium. Even though both ions can shed some of the water molecules to pass through passages, in general, the potassium ion will actually pass through many cell wall channels more readily than the sodium ion [22].
Group 2 (Alkaline Earth Metals)
This Group is the first one encountered in which there is only a smooth transition of properties if the first member of the Group is ignored. Thus, from magnesium to barium, chemical reactivity and density increase. Beryllium has a higher density than magnesium, and it exhibits weak metal behavior such as forming beryllates in very basic conditions.
Dolomite: The Mystery Mineral
Containing both calcium and magnesium in precisely equimolar proportions, dolomite has the formula: CaMg(CO3)2. Massive sedimentary deposits occur on Earth, including those in the Dolomite Alps in Italy. Yet, until recently, when chemists tried to synthesize the compound in the laboratory, all they obtained was a mixture of crystals of magnesium carbonate and calcium carbonate. Many hypotheses — some quite bizarre — were proposed to explain how it must have formed. Only in 2013 was this mineral laboratory synthesized by a reasonable pathway [23]. The stability of this mineral can be accounted for by the slightly different sizes of cation sites, for the related mineral ankerite has a composition: Ca(Fe(II),Mg,Mn(II))(CO3)2 where the other ions will only substitute for the magnesium ion, not the calcium ion.
Biological Roles of Strontium and Barium
It is rarely realized that Group 2 provides the greatest number of elements with biological roles: magnesium, calcium, strontium, and barium. The roles of magnesium and calcium are well-documented, thus, the focus here will be on strontium and barium. Some algae selectively concentrate these ions to form crystals of barium sulfate and strontium sulfate [24]. However, what is of crucial importance is the incorporation of strontium ion into human bone, hydroxoapatite, Ca5(PO4)3(OH). Bone formation favors incorporation of strontium over calcium by a very large factor. Presumably the larger strontium ion (132 pm) fits “more snugly” in the crystal lattice than the smaller calcium ion (114 pm). Perhaps if the concentration of strontium had been much higher in the geological past, vertebrates might have normally had strontium-containing bones. In the second half of the 20th century, incorporation of radioactive strontium-90 from weapons tests was feared. Now, the addition of (natural) strontium ion to diet and the incorporation into bone is being proposed as a means of combating osteoporosis [25].
Group 13 (Triels)
Is Group 13 really a group? At the top is boron whose chemistry is dominated by unique cluster species. Then comes aluminum, which would be happier in Group 3 (see Chapter 9). Next is gallium with its near room temperature melting point. And at the bottom, there is thallium that likes to masquerade as a Group 11 element or as a lower Group 1 element (see Chapter 10).
Boron Is Not Boring, It’s Unique
As soon as a chemist sees an icosohedron, boron comes immediately to mind. This beautiful and symmetrical molecule is the centerpiece of what makes this element unique. Of course, now a plethora of open- and closed cluster molecules and ions are known. Initially, these other species were believed to be simply fragments of an icosahedron. It was in 1971 that Wade showed that this was not true: instead, they were arranged into three families (closo-, nido-, and arachno-). Subsequently refined by Mingos, the criteria for these skeleta are now known as the Wade–Mingos rules [26]. The rules provide a straightforward and elegant rationalization of the shapes of “electron-deficient” cluster compounds in terms of the number of skeletal electron pairs (SEPs) these molecules.
Group 14 (Tetrels)
Just as Groups 1, 17, and 18 are regarded as epitomizing the smooth change in properties descending the respective group, Groups 14 to 16 represent the “discontinuity” groups. These are the groups that span the range of element behavior from nonmetal, through metalloid, to metal. In these cases, though there are sometimes similarities in chemical formula of compounds, there is little that can be chosen to select for group trend. In fact, group individuality is more interesting.
Graphite: The Forgotten Allotrope
Though the diamond and fullerene allotropes of carbon have taken the limelight in recent years, here the focus will be on oft-forgotten graphite. Graphite, with its layer structure of conjugated aromatically bonded atoms, has the ability to trap other atoms and molecules between the carbon sheets [27]. These are known as intercalation compounds:
Graphite intercalation compounds (GICs) are complex materials having a formula CXm where the ion Xn+ or Xn− is inserted between the oppositely charged carbon layers. Typically, m is much less than 1.
GICs are of interes
t in providing the electrode framework in battery systems. One specific example, of the many known species, is that between graphite and potassium. Molten potassium is absorbed into the black graphite layers to give a bronze-colored ionic solid with limiting composition of [K]+[C8]− [28].
Cubane: A Whole New Field of Inorganic Chemistry
Over the history of organic chemistry, cubane, C8H8, was considered simply a hypothetical molecule. With 90° bond angles, no one thought it could actually be synthesized, that is, until it was in 1964 [29]. Not only was it synthesizable but, when produced, it was a stable molecule. Why mention this in a book that is essentially inorganic chemistry? The pseudo-cubane structure is one that permeates cluster inorganic chemistry, and by its name, recognizes the simplest structure from which they are all derived. For example, there are the thallium–oxygen pseudo-cubanes, such as Tl4(OCH3)4 [30]. Silicon forms pseudo-cubanes, Si8(SitBuMe2)8. Phosphorus forms pseudo-cubanes where it alternates with boron, or aluminum, or nitrogen, or carbon, such as P4(CtBu)4. But of all the pseudo-cubanes, one must take top billing: that of the iron–sulfur pseudo-cubanes that are such crucial redox systems in so many biochemical processes (Figure 7.3) [31].
Figure 7.3 The common iron–sulfur pseudo-cubane core of many biological redox systems.
Group 15 (Pnictogens)
As for Group 14, the elements of Group 15 span a wide range of behaviors. And there are always surprises awaiting discovery. As an example, nitrogen is cited as having a single allotrope, N2. However, we see things from the perspective of our own SATP world. Under the conditions of the very low pressure at the edge of the Earth’s atmosphere, tetranitrogen, N4, is to be found [32].