Periodic Table, The: Past, Present, And Future

energies of the 3d-series metals, corresponding to the removal of each of the 3s² electrons, show a steady increase without any major deviations [37]. More interesting are the 3rd ionization energies, IE₃, of the 3d-series metals. With these subsequent ionization energies, it is the d electrons that are being “plucked off” one by one. As can be seen from Figure 2.6, it is an almost identically shaped plot to that for the 2p and 3p electrons in Figure 2.5, except that the greater coulombic repulsion between any electron pairs commences with the d⁶ configuration (instead of p⁴), as expected.

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