Question

The electron affinities, in \(\mathrm{kJ} / \mathrm{mol},\) for the group 1 \(\mathrm{B}\) and group 2 \(\mathrm{B}\) metals are as follows: (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 affinities of other groups as we proceed down the periodic table.]

Short answer

In short, the electron affinities of group 2B elements are greater than zero due to electron-electron repulsion in higher-energy shells resulting from full s orbitals. The electron affinities of group 1B elements become more negative as we move down the group because the increasing atomic radius and electron shielding make it easier to add an electron.

Step-by-step solution

(Step 1: Understanding Electron Affinity)

Electron affinity refers to the energy change that occurs when an electron is added to a neutral atom to form a negative ion. In other words, it's the energy released when an atom gains an electron. Elements with high electron affinity tend to accept electrons more readily, whereas elements with low electron affinity are less likely to accept electrons.

(Step 2: Identifying Group 1B and 2B Elements)

In the periodic table, group 1B elements include copper (Cu), silver (Ag), and gold (Au). Group 2B elements include zinc (Zn), cadmium (Cd), and mercury (Hg).

(Step 3: Explaining Electron Affinity in Group 2B Elements)

Group 2B elements have a greater electron affinity than zero because their outermost electron shell (s orbitals) is full, and adding an electron would require it to occupy a higher-energy shell. As a result, electron affinity is positive, which means energy is required to add an electron to these elements, making the process unfavorable. The electron-electron repulsion in the higher-energy shell due to the full s orbitals makes it challenging for the electron to be added.

(Step 4: Explaining Electron Affinity Trend in Group 1B Elements)

The electron affinities of Group 1B elements become more negative as we move down the group because the atomic radius increases down the group due to the increase in the number of electron shells. As a consequence, the valence electrons are further from the nucleus and are less strongly attracted to it, making it easier to add an electron. Additionally, electron shielding increases as we move down the group, which reduces the effective nuclear charge experienced by the incoming electron. This means that the incoming electron experiences less repulsion from the inner electrons, making it easier for it to be added to the atom. In conclusion, the electron affinity of group 2B elements is greater than zero because of the electron-electron repulsion experienced in the higher-energy shells. The electron affinities of group 1B elements become more negative as we move down the group due to the increase in atomic radius and electron shielding.

Key concepts

Understanding Group 1B Elements

Group 1B elements, which include copper (Cu), silver (Ag), and gold (Au), are fascinating components of the periodic table. These elements are known for their excellent electrical conductivity, malleability, and ductility.
They are essential in various industrial applications, such as electronics and jewelry making. The chemistry of Group 1B elements is influenced by the full d subshell, which contributes to their distinctive electron affinity trends as compared to other groups.

• Electron Affinity Trends: As we move down the group from copper to gold, the electron affinity becomes more negative. This arises because of the increase in atomic size leading to a bigger atomic radius. This increase permits an incoming electron to be added with greater ease due to a reduced effective nuclear charge after factoring in electron shielding.
• Noble Nature: These elements are often termed "coinage metals" due to their long-standing use in minting coins. Their resistance to corrosion and tarnishing is a noble trait, adding to their value and utility.
• Learn Through History: Understanding the practical applications of these elements throughout history paves the way for a deeper appreciation of their chemical properties.

Overview of Group 2B Elements

The group 2B elements—zinc (Zn), cadmium (Cd), and mercury (Hg)—are commonly referred to as the "zinc group" in the periodic table. These metals share similarities, such as having a full d^10 configuration. However, they stand out for their varied effects on electron affinity.

• Positive Electron Affinity: The electron affinities in this group are greater than zero, which means energy is required to add an electron. The full s orbitals mean any incoming electron must enter a higher energy p orbital, making the process energetically unfavorable due to electron-electron repulsion.
• Physical Properties: These metals are soft and have relatively low melting points, with mercury being uniquely liquid at room temperature.
• Applications and Uses: Zinc is vital in galvanization to protect iron from rusting. Cadmium finds uses in batteries, whereas mercury is used in certain types of electrical switches and lighting.

Understanding Periodic Table Trends

The periodic table is an organized map that reveals patterns and trends in element properties, including electron affinity. By studying these trends, we can predict how an element might behave in chemical reactions.

• General Trend for Electron Affinity: Electron affinity generally becomes more negative across a period (left to right) owing to increased nuclear charge, which strengthens the attraction for the added electrons.
• Down the Group Trends: As you move down a group, electron affinities tend to become less negative, owing to increased atomic radii and electron shielding effects. The impact of shielding is more pronounced in larger atoms, diminishing the effective nuclear charge felt by incoming electrons.
• Exceptions to the Rule: Electron affinity values are not universally predictable, as they depend on specific atomic configurations and the inherent stability of electron shells.