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an effective nuclear charge on the periphery of each atom, which increases as the number of protons increases. In the 2nd Period, for example, the greatest EA₁ is that of fluorine. There are three exceptions to the negative EA₁: beryllium, nitrogen, and neon.

      Figure 2.7 Electron affinity (EA₁) hydrogen to calcium.

      •Beryllium has a positive EA₁ as an added electron would have to enter a 2p orbital where it would be shielded by the 2s² electrons. In fact, the electron repulsion must exceed the nuclear attraction:

      [He]2s² → [He]2s²2p¹

      •Nitrogen has a positive EA₁ as a result of the interelectronic repulsion being greater than the effective nuclear attraction:

      [He]2s²2p³ → [He]2s²2p⁴

      •Neon has a positive EA₁ as an added electron would have to enter a 3s orbital where it would be shielded from the nuclear attraction particularly by the 2s² and 2p⁶ electrons. In fact, the electron repulsion must exceed the nuclear attraction from the nucleus:

      [He]2s²2p⁶ → [He]2s²2p⁶3s¹

       Group Trends in Electron Affinities

      Down a group, as the atoms become larger and the nuclear attraction becomes less, so the electron affinities decrease. The trend is illustrated in Figure 2.8.

      The 2nd Period elements from boron to fluorine are clearly exceptions to the trends in their respective groups. Their electron attachment energies are significant deviations from the smooth progressions of the other members of their groups. That is, their electron attraction energy is significantly less than expected. For example, that of nitrogen is +7 kJ⋅mol⁻¹ while that for phosphorus is −72 kJ⋅mol⁻¹; similarly, that of oxygen is −141 kJ⋅mol⁻¹ while that for sulfur is −200 kJ⋅mol⁻¹. An accepted explanation is that the atoms are so small that the interelectron repulsion factor is exceptionally large and, as a result, the attraction for an additional electron is significantly reduced. The anomalous electron affinity of gold will be discussed later in the chapter.

      Figure 2.8 A plot of 1st electron affinities by period (adapted from Ref. [41]).

       Multiple Electron Affinities

      Just as there are multiple ionization energies, so there are the corresponding multiple electron affinities. However, whereas the atomic ionization energies are always positive, as discussed earlier, the 1st electron affinity is often negative. Nevertheless, the subsequent electron affinities are all positive as a result of the increasing electron–electron repulsions. This can be illustrated by the electron affinities of the nitrogen atom:

       Alkalide Ions

      As the formation of the Na⁻ ion is energetically favored, then compounds containing that ion should be feasible.

It was in 1974 that Dye et al. synthesized the first known compound containing the sodide ion [43]. The team realized 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 [ Скачать книгу