Figure 2.6 3rd ionization energy (IE3) for the 3d-block elements (adapted from Ref. [35]).
Group Trends in Ionization Energy
Proceeding down a group, the 1st ionization energy generally decreases. This is especially systematic for the noble gases.
Though the number of protons in the nucleus has increased, so has the number of shielding electron shells. In addition, as the sequential orbitals are filled, the electrons in the outermost shell occupy a larger volume of space and thus have lower interelectrons repulsion factors.
Successive Ionization Energies
There are also patterns in successive ionizations of an element [38]. One of the simplest examples is lithium:
Lithium has the electron configuration 1s22s1. Thus, the first electron to be removed is strongly shielded by the two 1s electrons. Then, to remove each of the 1s electrons requires very much greater energy. The lesser value for removing the second electron compared to the third can be accounted for by two factors: First, there are always electron–electron repulsions when two electrons occupy the same orbital; second, even within the same orbital, one electron does partially shield the other electron.
Electron Affinity
Much space is usually given to ionization energy and little to electron affinity (rarely, but more correctly, called electron attachment energy). Yet as mentioned earlier, atoms usually “want” to gain electrons and certainly not lose them! The following definition is parallel to that given for ionization energy.
The experimental 1st electron affinity is equal to the difference between the total electronic energy of the atom X and the total electronic energy of the ion X–, both in their ground states. That is, X(g) + e− → X−(g)
Sign Convention for Electron Affinity
For clarity, it is important to commence with a mention of the confusion over the sign convention for electron affinity. A proponent of the traditional sign convention (no longer in common use) was Wheeler, who contended that [39]:
With this convention, the electron affinity is positive for elements such as fluorine, for which energy is released when an electron is added to make an ion, while the widely quoted values for the alkaline earth metals and noble gases are negative.
This convention, however, is the opposite of that used for ionization energy. To remove the ambiguity, Brooks et al. proposed that the term “electron affinity” should be eliminated and, instead, the reverse process should be regarded as the 0th ionization energy [40]:
This format, which never gained wide acceptance, would correspond with the sign convention used here for electron affinity:
Period Patterns in Electron Affinity
If anything, the patterns for electron affinity are more interesting than those of ionization energy [41, 42]. The graph in Figure 2.7 shows the first electron attachment energies for the 1st, 2nd, and 3rd Periods.
As with ionization energy, there are the two factors involved: interelectron repulsion and exchange energy. There is still 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 [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 p1/2 and p3/2 while the d levels split into d3/2 and d5/2 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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