Question

Based on their positions in the periodic table, predict which atom of the following pairs will have the larger first ionization energy: (a) \(\mathrm{Cl}, \mathrm{Ar} ;\) (b) Be, Ca; (c) \(\mathrm{K}, \mathrm{Co}\); (d) S, Ge; (e) Sn, Te.

Short answer

Based on the positions in the periodic table, the elements with larger first ionization energy in each pair are: (a) Ar, (b) Be, (c) Co, (d) S, and (e) Te.

Step-by-step solution

Compare Ionization Energies for Cl and Ar

Looking at the periodic table, we can see that Cl and Ar are in the same period with positions: Cl (period 3, group 17) and Ar (period 3, group 18). Since they are in the same period, we just need to check their positions within that period. As the ionization energy increases from left to right across a period, Ar should have a larger first ionization energy than Cl.

Compare Ionization Energies for Be and Ca

Be and Ca are in the same group (group 2), but in different periods with positions: Be (period 2, group 2) and Ca (period 4, group 2). Since they are in the same group, we just need to check their positions within that group. As the ionization energy decreases from top to bottom within the same group, Be should have a larger first ionization energy than Ca.

Compare Ionization Energies for K and Co

K and Co are in the same period (period 4) but in different groups with positions: K (period 4, group 1) and Co (period 4, group 9). As they are in the same period, we only need to check their positions within that period. As ionization energy increases from left to right across a period, Co should have a larger first ionization energy than K.

Compare Ionization Energies for S and Ge

Looking at the periodic table, we find S and Ge in the same period (period 4) with positions: S (period 4, group 16) and Ge (period 4, group 14). Since they are in the same period, we only need to check their positions within that period. As ionization energy increases from left to right across a period, S should have a larger first ionization energy than Ge.

Compare Ionization Energies for Sn and Te

Sn and Te are in the same period (period 5) with positions: Sn (period 5, group 14) and Te (period 5, group 16). Since they are in the same period, we just need to check their positions within that period. As ionization energy increases from left to right across a period, Te should have a larger first ionization energy than Sn.

Summary of Results

(a) Ar has a larger first ionization energy than Cl. (b) Be has a larger first ionization energy than Ca. (c) Co has a larger first ionization energy than K. (d) S has a larger first ionization energy than Ge. (e) Te has a larger first ionization energy than Sn.

Key concepts

Periodic Table Trends

Understanding how elements behave is key to solving problems in chemistry, and the periodic table is a chemist’s roadmap. Ionization energy, which is the energy required to remove an electron from the atom in its gaseous state, exhibits notable trends in the periodic table.

Generally, ionization energy increases from left to right within a period. This happens because as we move across a period, the atomic number increases, adding more protons to the nucleus and therefore more positively charged nuclear pull. This increased pull makes it more difficult to remove an electron.

Conversely, ionization energy decreases from top to bottom within a group. As we move down a group, new electron shells are added, which means that the outer electrons are farther from the nucleus and experience a lesser nuclear attraction. Additionally, the effect of electron shielding, where inner electrons block some of the nuclear charge from reaching the outermost electrons, also contributes to this decrease.

The interplay of these trends helps predict which element in a pair might have a larger ionization energy based on its position in the periodic table.

Atomic Structure

The atomic structure of an element underlies its position in the periodic table and governs its properties, including ionization energy. Each atom consists of a dense nucleus made up of protons and neutrons, surrounded by a cloud of electrons in orbitals at various energy levels.

The arrangement of electrons in these orbitals is crucial as it impacts how tightly bound they are to the nucleus. Electrons in orbitals closer to the nucleus are more strongly attracted, making them harder to remove. As the number of protons increases, the effective nuclear charge felt by the valence electrons also increases, generally leading to higher ionization energies for elements with more protons, provided they are in the same period.

Furthermore, the distribution of electrons among the subshells within a period can lead to slight variations in ionization energies due to electron-electron repulsions in the same orbital, also known as Hund's Rule effects, which can impact the ease with which an electron can be removed.

Comparing Elements

When comparing elements to determine which has a higher ionization energy, the atomic structure and periodic table trends provide the necessary context. To improve students' abilities to predict ionization energies, consider the following focal points:

• Atomic number (amount of protons) – Higher atomic numbers within the same period usually correspond to higher ionization energies due to increased nuclear charge.
• Energy levels (electron shells) – Elements with more electron shells have their valence electrons further from the nucleus, which typically results in lower ionization energies.
• Position in the periodic table – Specifically, look at whether elements are in the same period or group to apply the correct trend.
• Electron configurations – These can cause anomalies in the general trend, such as the stability provided by full or half-full orbitals, which may result in unexpectedly higher ionization energies.

By focusing on these aspects and considering the underpinning principles of electron shielding and nuclear charge, one can make informed comparisons on the relative ionization energies of different elements.