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
