The electron affinities, in \(\mathrm{kJ} / \mathrm{mol}\), for the group \(1 \mathrm{~B}\) and group \(2 \mathrm{~B}\) metals are $$ \begin{array}{|c|c|} \hline \mathrm{Cu} & \mathrm{Zn} \\ -119 & >0 \\\ \hline \mathrm{Ag} & \mathrm{Cd} \\ -126 & >0 \\ \hline \mathrm{Au} & \mathrm{Hg} \\ -223 & >0 \\ \hline \end{array} $$ (a) Why are the electron affinities of the group \(2 \mathrm{~B}\) elements greater than zero? (b) Why do the electron affinities of the group \(1 \mathrm{~B}\) elements become more negative as we move down the group? [Hint: Examine the trends in the electron affinity of other groups as we proceed down the periodic table. \(]\)

Electron affinity is defined as the energy change that occurs when an electron is added to a neutral atom to form a negatively charged ion. It is an exothermic process, which means energy is released when an electron is added to an atom. However, in some cases, energy must be added to force an electron to join an atom, and that makes the electron affinity of those elements positive.

Group 2B elements have a completely filled outer s subshell and a partially filled d subshell. When adding an electron to a group 2B element, the additional electron has to enter the same d subshell, as per Hund's rule. This increases electron-electron repulsion because the electron-electron distance within the d subshell is relatively short, making it harder for the added electron to stay within. This repulsion causes the energy change to be positive when an electron is added to a Group 2B element. Therefore, the electron affinities of Group 2B elements are greater than zero because adding an electron increases electron-electron repulsion within the d subshell.

Similar to ionization energies, electron affinities also exhibit periodic trends in the periodic table. However, the trends for electron affinity do not have a continuous progression down a group. It is generally observed that electron affinity becomes more negative as we move across a period from left to right and tends to be less negative as we move down a group. This is due to the increased nuclear charge and smaller atomic size across a period and the increased atomic size and electron shielding down a group.

Group 1B elements have a completely filled s subshell and a partially filled d subshell with a single unpaired electron. When an electron is added to a group 1B element, it occupies the same d subshell, but it occupies a different d orbital, resulting in a relatively weaker electron-electron repulsion compared to Group 2B elements. While moving down the periodic table, the trend of less negative electron affinity reverses within some groups, including group 1B. In this specific case, the increase in atomic size and electron shielding dominates over the positive features leading to more negative electron affinities down the group. This corresponds to the stronger attraction between the nucleus and the added electron and hence, less energy is required to attach the electron to the atom. To sum up, the electron affinities of Group 1B elements become more negative as we move down the group because the increase in atomic size and electron shielding prevails over the electron-electron repulsion, resulting in a stronger attraction between the nucleus and the added electron.