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to be the best [35]:

       Because of the Pauli Exclusion Principle, electrons with parallel (unpaired) spins tend to avoid each other, thus decreasing the electrostatic repulsion between them. This will be the situation when filling the first half of the shell. When electrons are forced to doubly occupy orbitals in the second half, their spins are constrained to be paired (antiparallel). Because they are no longer obliged to avoid each other, the [inter-electron] electrostatic repulsion increases.

      In Figure 2.5, the IE₁ are shown for the 2p and 3p block elements. Continuing the line of the p¹ to p³ configurations, a line parallel to the actual p⁴ to p⁶ values is obtained. The difference between the two represents the coulombic repulsion between pairs of electrons within the same orbital. For the 2p series, this amounts to about 430 kJ⋅mol⁻¹, while for the 3p series it is 250 kJ⋅mol⁻¹. Cann attributed the difference between the two series to the more diffuse 3p orbitals compared with the 2p orbitals. Thus, any paired 3p electrons are sharing a larger volume of space and therefore have less mutual repulsive forces.

      Figure 2.5 First ionization energy for the p-block elements of the 2nd and 3rd Periods (adapted from Ref. [35]).

To review, there is nothing exceptional about the “half-filled shell.” It is instead the interelectron repulsive forces between the electron pairs beyond the p³ configuration, which result in a lower ionization energy. To reinforce the point, as Rich and Suter added [36]:

       Likewise, when one compares the energy to remove an electron from the half-filled p subshell with that needed for a p² structure, nothing special is found.

      Rich and Suter then referred to the claimed stability of the “filled shell.” They continued [36]:

       Similarly, the large energy difference between electrons in 3s¹ and 2p⁶ configurations is readily explained by the difference in principal quantum number; this again indicates no more “extra” stability of a filled p shell than it does for a p⁵ or any other structure in which the electron being removed is at the lower principal number.

       3d-Series Metal Ionization Energies

The 1st and 2nd ionization 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,

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