Question

The following table gives the electron affinities, in \(\mathrm{kJ} / \mathrm{mol}\), for the group \(1 \mathrm{~B}\) and group \(2 \mathrm{~B}\) metals: (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.] \begin{tabular}{|c|l|} \hline \(\mathrm{Cu}\) & \multirow{2}{*} { \(\mathrm{Zn}\) \(-119\)} & \(>0\) \\ \hline \(\mathrm{Ag}\) & \(\mathrm{Cd}\) \\ \(-126\) & \(>0\) \\ \hline \(\mathrm{Au}\) & \(\mathrm{Hg}\) \\ \(-223\) & \(>0\) \\ \hline \end{tabular}

Short answer

(a) The electron affinities of group \(2 \mathrm{~B}\) elements are greater than zero because they have a stable electron configuration with completely filled s and d orbitals, and adding an electron requires an input of energy to overcome the energy barrier and stability of the filled orbitals. (b) The electron affinities of group \(1 \mathrm{~B}\) elements become more negative as we move down the group due to an increase in atomic size and a decrease in the effective nuclear charge experienced by the added electron, which leads to a greater release of energy when forming a negative ion.

Step-by-step solution

Part (a) Explanation: Positive Electron Affinities of Group 2B Elements

Electron affinity measures the energy released when an atom gains an electron to form a negative ion. For group \(2 \mathrm{~B}\) elements, the electron affinity values are greater than zero indicating that energy is absorbed in the process of adding an electron. The following reasons account for this observation: 1. Group \(2 \mathrm{~B}\) elements have completely filled s and d orbitals in their outermost energy levels. Adding an electron would require filling the next energy level, which is at a significantly higher energy. 2. Due to the completely filled outer energy levels, these elements have a relatively stable electron configuration as the s and d orbitals create spherical, symmetrical electron clouds, resulting in minimal repulsion between electrons. To conclude, the electron affinities of group \(2 \mathrm{~B}\) are positive because they have a stable electron configuration and adding an electron requires an input of energy to overcome the energy barrier and stability of the filled orbitals.

Part (b) Explanation: Increasingly Negative Electron Affinities of Group 1B Elements

As we move down the periodic table, the electron affinities of group \(1 \mathrm{~B}\) elements become more negative. This trend can be explained by examining the atomic size and effective nuclear charge experienced by the added electron: 1. As we progress down the group, the atomic size increases due to the increase in the number of electron shells. This increased distance between the nucleus and the valence electrons weakens the electrostatic attraction between the added electron and the nucleus. 2. Additionally, the effective nuclear charge experienced by the added electron decreases as we move down the group due to increased electron shielding provided by the inner-shell electrons. 3. As a result, the energy released on adding an electron becomes more negative for group \(1 \mathrm{~B}\) elements as we move down the periodic table. The weaker electrostatic attraction and reduced effective nuclear charge of the added electron lead to a greater release of energy when forming a negative ion. In summary, the increasingly negative electron affinities in group \(1 \mathrm{~B}\) elements are due to the increase in atomic size and decrease in effective nuclear charge experienced by the added electron as we move down the group.

Key concepts

Electron Affinity Trends

Electron affinity refers to the amount of energy released when an electron is added to a neutral atom in the gaseous state to form a negative ion. This property follows distinct trends on the periodic table. Generally, electron affinities become more negative as you move from left to right across a period, which is due to the increasing nuclear charge with relatively stable shielding effect, resulting in a stronger attraction for the additional electron.

However, there are exceptions to this trend where the electron affinity is less negative or even positive. For instance, group 2B elements have electron affinities greater than zero because adding electrons to these atoms requires energy. Conversely, the values become progressively more negative as you move down a group, particularly within group 1B elements. This gradual change is largely due to variations in atomic size and effective nuclear charge, which will be further explained in the following sections.

Group 1B Elements

Group 1B elements, consisting of copper (Cu), silver (Ag), and gold (Au), exhibit a pattern where their electron affinities become more negative with increasing atomic number. This is related to their unique electron configurations, where the elements have a single s electron outside a completed d shell. As you move down the group, the outermost electrons are farther from the nucleus, and this increased distance means a reduction in energy is required to add another electron, leading to more negative electron affinity values.

For example, gold has a more negative electron affinity than copper due to its larger atomic radius and reduced electrostatic attraction to the nucleus for the incoming electron.

Group 2B Elements

Group 2B elements include zinc (Zn), cadmium (Cd), and mercury (Hg). Unlike group 1B, these elements have electron affinities that are positive. This is because they possess a fully filled d subshell, making them more stable and less inclined to accept additional electrons. Adding an electron to these elements means placing it into a higher energy level, which isn't energetically favorable without an input of energy.

The positive electron affinity values reflect the energy needed to surpass this stability barrier. Thus, for group 2B elements, the process of gaining an electron is endothermic, absorbing rather than releasing energy.

Periodic Table Groups

The periodic table is structured in rows (periods) and columns (groups) with elements placed according to their atomic number. Groups are particularly significant when discussing elements' chemical properties because elements within the same group generally share similar valence electron configurations, lending to common chemical behaviors.

Group 1B and group 2B are part of the transition metals, often characterized by their ability to form positive ions with incomplete d orbitals. Understanding the properties of each group facilitates the comprehension of trends such as electron affinity, which can be explained by the elements' position on the periodic table.

Atomic Size

Atomic size, or atomic radius, is a concept that describes the size of an atom, typically measured from the nucleus to the boundary of the surrounding electron cloud. Across a period, atomic size decreases due to the increase in nuclear charge, which pulls the electron cloud closer to the nucleus.

Down a group, atomic size increases as each successive element has an additional electron shell, making the atom larger. This increase in atomic size contributes to the trend of electron affinities becoming more negative as you move down group 1B, because the further the valence electrons are from the nucleus, the less energy is required to attach an additional electron.

Effective Nuclear Charge

Effective nuclear charge (Zeff) is the net positive charge experienced by an electron in an atom. Despite the increase in the number of protons in the nucleus as we move across a period, the effective nuclear charge felt by outer electrons doesn't increase proportionally, due to shielding by inner electrons.

Down a group, inner electron shells provide added shielding, reducing the Zeff experienced by valence electrons, which in turn leads to a weaker attraction for additional electrons and contributes to the trend of more negative electron affinities in group 1B elements lower down the group. An understanding of Zeff helps explain variations in properties such as electron affinity, atomic size, and ionization energy.