Question

Listed below are two atomic properties of the element germanium. Refer only to the periodic table on the inside front cover and indicate probable values for each of the following elements, expressed as greater than, about equal to, or less than the value for Ge. $$\begin{array}{lcc} \hline \text { Element } & \text { Atomic Radius } & \begin{array}{c} \text { First lonization } \\ \text { Energy } \end{array} \\ \hline \mathrm{Ge} & 122 \mathrm{pm} & 762 \mathrm{kJ} / \mathrm{mol} \\ \mathrm{Al} & ? & ? \\ \mathrm{In} & ? & ? \\ \mathrm{Se} & ? & ? \\ \hline \end{array}$$

Short answer

Al has smaller atomic radius and lower first ionization energy than Ge. In has larger atomic radius and lower first ionization energy than Ge. Se has smaller atomic radius and higher first ionization energy than Ge.

Step-by-step solution

- Comparing with Aluminum (Al)

Aluminum (Al) is located above and to the left of Germanium (Ge) in the periodic table. Thus, relative to Ge, Al has a smaller atomic radius and a lower first ionization energy.

- Comparing with Indium (In)

Indium (In) is located below and to the left of Germanium (Ge) in the periodic table. Therefore, relative to Ge, In has a larger atomic radius and a lower first ionization energy.

- Comparing with Selenium (Se)

Selenium (Se) is located above and to the right of Germanium (Ge) in the periodic table. Thus, compared to Ge, Se has a smaller atomic radius and a higher first ionization energy.

Key concepts

Atomic Radius

Understanding the concept of atomic radius is crucial when analyzing trends in the periodic table. The atomic radius is the average distance from the nucleus of an atom to its outermost electron cloud. As we move down a group in the periodic table, the atomic radius increases.
This occurs because entering a new period means adding an extra electron shell, making the atoms larger.​

• For example, Indium (In), which is below Germanium (Ge) in the periodic table, is expected to have a larger atomic radius than Ge.
• Aluminum (Al), located above Ge, should have a smaller atomic radius than Ge.

When moving across a period from left to right, the atomic radius generally decreases.
This is due to the increase in the positive charge of the nucleus, pulling electrons closer and reducing the atomic size.
Thus, Selenium (Se), which is to the right of Ge, will likely have a smaller atomic radius compared to Ge.

Ionization Energy

Ionization energy is the energy required to remove an electron from a neutral atom. It indicates how strongly an atom holds onto its electrons.
Generally, ionization energy increases across a period and decreases down a group.
As we move from left to right across a period, the nucleus has a stronger attraction for electrons due to an increase in the number of protons, leading to higher ionization energies.

• For instance, Selenium (Se), being to the right of Germanium (Ge), will possess a higher ionization energy than Ge.

In contrast, moving down a group, the outer electrons are further from the nucleus due to additional electron shells, decreasing the attractive force and thus decreasing the ionization energy.

• Hence, Indium (In) below Ge, will have a lower ionization energy than Ge.
• Similarly, Aluminum (Al), located above Ge, should also have a lower ionization energy than Ge.

Group and Period Trends

The periodic table is structured in such a way that certain trends in atomic properties can be observed both vertically in groups and horizontally in periods. Understanding these general trends can simplify predictions regarding atomic size, electron affinity, and metallic character.
Moving down a group, atoms generally increase in size due to additional electron shells. This increase in size means that electrons are farther from the nucleus, affecting both atomic radius and ionization energy.
In any given group:

• Atomic radius increases.
• Ionization energy generally decreases.

In a period, these trends change due to the increasing number of protons in the nucleus. Each successive element in a period has a smaller atomic radius owing to stronger nuclear attraction for electrons, while the ionization energy tends to rise as we move to the right.
These trends are critical for understanding chemical reactivity and bonding, allowing predictions of atomic interactions based on position alone in the periodic table.