Periodic Table, The: Past, Present, And Future

that, in the solid phase, there was little energy needed for the formation of the sodium cation–anion pair:

The key, then, was to find a way of keeping the two ions separated. To do this, Dye et al. caged the sodium ion in a bicyclic diaminoether, commonly known as 2,2,2-crypt. The synthesis was successful and gold-colored crystals of [Na(C₁₈H₃₆N₂O₆)]⁺⋅Na⁻ were produced. From the crystal structure, the radius of the sodide ion was calculated to be 217 pm, close to that of the iodide ion, and the sodide compound has a structure similar to that of the analogous iodide: [Na(C₁₈H₃₆N₂O₆)]⁺⋅I⁻. The preparation of anions of the other alkali metals followed [44]. Then in 1987, Concepcion and Dye synthesized a simpler compound of the sodide ion: [Li(diaminoethane)₂]⁺⋅Na⁻ [45].

Since then, simple stable compounds of both the sodide ion and the potasside ion have been synthesized [46]. Of note, the tradition of using the Latin-derived name for the anion was not followed as these anions should have been named “natride” and “kalide,” respectively. No explanation was stated, though perhaps it was to avoid confusion of “natride” with “nitride.”

A particularly intriguing compound is the so-called “inverse sodium hydride.” Sodium hydride itself, Na⁺H⁻, is a well-known reducing agent as a result of the “naked” hydride ion [47]. By “caging” the hydrogen ion, it has been possible to synthesize [H⁺]_cageNa⁻ [48].

The Auride Ion

Looking at the plot of electron affinities (Figure 2.8), gold stands out as an obvious candidate for anion formation.

In fact, the first evidence for the formation of an auride came in 1937 by the equimolar mixing of cesium and gold [49]. This transparent yellow compound was shown in 1959 not to be an alloy, but to be Cs⁺Au⁻, with a sodium chloride crystal structure. Since then, several other auride compounds have been synthesized [50], including tetramethylammonium auride, [N(CH₃)₄]⁺⋅Au⁻. The compound is isostructural to the corresponding bromide, which further illustrates the similarities between the auride and halide ions [51].

The Platinide Ion

At −205 kJ⋅mol⁻¹, EA₁ for platinum is close to that of gold. Thus, it should come as no surprise that there is an increasing chemistry of the platinide ion, Pt²⁻, including cesium platinide, Cs₂Pt [52].

Relativistic Effects on Atomic Properties

As an explanation for the significantly negative electron affinity, and other anomalous behavior, relativistic effects must be invoked [53]. These effects are rarely discussed in general chemistry [54], yet they are vital to the comprehension of many facets of atomic behavior [55]. Two of the contexts in which relativistic effects are discussed are the color of gold [56, 57] and the liquid phase of mercury at room temperature [58]. In this section, the focus will be on the relativistic explanation for the formation of auride and platinide ions and then in later chapters on some other relevant relativistic phenomena.

Though the electrons in all atoms experience some degree of relativistic effects, they only become important for the heavier elements. There are two significant factors that can be ascribed to relativistic effects [59] (Figure 2.9 shows both factors for the 5d, 6s, and 6p energy levels):

Figure 2.9 Nonrelativistic and relativistic energy levels for the 5d, 6s, and 6p orbitals (adapted from Ref. [59]).

•Changing in relative energy levels of atomic orbitals

s orbitals decrease substantially in energy and p orbitals decrease to a lesser extent when relativistic effects are taken into consideration. This results in increased shielding of the nucleus, causing d orbitals and f orbitals to increase in energy.

•Splitting of energy levels having l > 0 into two sublevels as a result of spin–orbit coupling

p levels split into p_1/2 and p_3/2 while the d levels split into d_3/2 and d_5/2 levels.

Platinum and Gold Electron Affinities

It is relativistic effects that can explain the high EA₁ for platinum and gold. The additional electron enters the 6s orbital:

Figure 2.10 Plot of ratio of relativistic to nonrelativistic atom radii for the 6s orbital (adapted from Ref. [60]).

As can be seen from Figure 2.10, the relativistic decrease in relative radius for an added 6s electron reaches a minimum at gold, with the value for platinum being not substantially different [60]. That is, there will be a greater effective nuclear charge on any additional 6s electron for platinum and gold than would be expected without taking relativistic effects into account.

Commentary

In this chapter, a mere selection of atomic periodic properties have been chosen for discussion. In this way, the Reader is not overwhelmed by endless tables and graphs of data. Those who wish to indulge should look elsewhere. This book is designed to make the many concepts of elemental relationships become alive and stimulating, not boring and soporific. The chapter has ended with an introduction to relativistic effects. This oft-overlooked aspect will not be simply a passing reference, but a topic that will be revisited in different contexts in later chapters.

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