The first ionization energy of an atom is a measure of the energy change associated with removing an electron from the atom to form a cation. For example, the first ionization energy of Cl(g), 1251 kJ/mol, is the energy change associated with the process

The positive ionization energy means that energy must be put into the atom to remove the electron.

Most atoms can also gain electrons to form anions. The energy change that occurs when an electron is added to a gaseous atom is called the electron affinity because it measures the attraction, or affinity, of the atom for the added electron. For most atoms, energy is released when an electron is added. For example, the addition of an electron to a chlorine atom is accompanied by an energy change of −349 kJ/mol, the negative sign indicating that energy is released during the process. We therefore say that the electron affinity of Cl is −349 kJ/mol.*

It is important to understand the difference between ionization energy and electron affinity: Ionization energy measures the ease with which an atom loses an electron, whereas electron affinity measures the ease with which an atom gains an electron.

The greater the attraction between an atom and an added electron, the more negative the atom's electron affinity. For some elements, such as the noble gases, the electron affinity has a positive value, meaning that the anion is higher in energy than are the separated atom and electron:

The fact that the electron affinity is positive means that an electron will not attach itself to an Ar atom; the Arion is unstable and does not form.

FIGURE 7.11 shows the electron affinities for the s- and p-block elements of the first five periods. Notice that the trends are not as evident as they are for ionization energy. The halogens, which are one electron shy of a filled p subshell, have the most-negative electron affinities. By gaining an electron, a halogen atom forms a stable anion that has a noble-gas configuration (Equation 7.5). The addition of an electron to a noble gas, however, requires that the electron reside in a higher-energy subshell that is empty in the atom (Equation 7.6). Because occupying a higher-energy subshell is energetically unfavorable, the electron affinity is highly positive. The electron affinities of Be and Mg are positive for the same reason; the added electron would reside in a previously empty p subshell that is higher in energy.

The electron affinities of the group 5A elements are also interesting. Because these elements have half-filled p subshells, the added electron must be put in an orbital that is already occupied, resulting in larger electron-electron repulsions. Consequently, these elements have electron affinities that are either positive (N) or less negative than the electron affinities of their neighbors to the left (P, As, Sb). Recall that in Section 7.4 we saw a discontinuity in the trends for first ionization energy for the same reason.

Electron affinities do not change greatly as we move down a group (Figure 7.11). For F, for instance, the added electron goes into a 2p orbital, for Cl a 3p orbital, for Br a 4p orbital, and so forth. As we proceed from F to I, therefore, the average distance between the added electron and the nucleus steadily increases, causing the electron-nucleus attraction to decrease. However, the orbital that holds the outermost electron is increasingly spread out, so that as we proceed from F to I, the electron-electron repulsions are also reduced. As a result, the reduction in the electron-nucleus attraction is counterbalanced by the reduction in electron-electron repulsions.


Which of the groups shown here has the most negative electron affinities? Why does this make sense?

FIGURE 7.11 Electron affinity in kJ/mol for selected s- and p-block elements.


What is the relationship between the value for the first ionization energy of a Cl(g) ion and the electron affinity of Cl(g)?