Question

Choose the element with the higher ionization energy from each pair. (a) \(\mathrm{Al}\) or In (b) \(\mathrm{Cl}\) or \(\mathrm{Sb}\) (c) \(\mathrm{K}\) or \(\mathrm{Ge}\) (d) \(\mathrm{S}\) or \(\mathrm{Se}\)

Short answer

The elements with the higher ionization energy in each pair are: (a) \text{Al}, (b) \text{Cl}, (c) \text{Ge}, (d) \text{S}.

Step-by-step solution

Understanding Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. The higher the ionization energy, the more difficult it is to remove an electron. Generally, ionization energy increases across a period from left to right and decreases down a group on the periodic table.

Comparing Elements in the Same Group

When comparing elements in the same group (column), the one with the higher period number (closer to the bottom of the periodic table) tends to have lower ionization energy, because its valence electrons are farther from the nucleus and more shielded by inner electrons.

Comparing Elements in the Same Period

When comparing elements in the same period (row), the one with the higher group number (further to the right on the periodic table) tends to have higher ionization energy, because its valence electrons are more tightly held by the higher nuclear charge.

Choice A: Aluminum (Al) vs Indium (In)

Aluminium (Al) and Indium (In) are both in group 13. Since Al is in period 3 and In is in period 5, Al has the higher ionization energy.

Choice B: Chlorine (Cl) vs Antimony (Sb)

Chlorine (Cl) is in period 3, group 17 and Antimony (Sb) is in period 5, group 15. Although Sb is further down, Cl is further to the right on the periodic table. Therefore, Cl has the higher ionization energy.

Choice C: Potassium (K) vs Germanium (Ge)

Potassium (K) is in period 4, group 1 while Germanium (Ge) is in period 4, group 14. Ge is to the right of K on the periodic table, therefore Ge has the higher ionization energy.

Choice D: Sulfur (S) vs Selenium (Se)

Sulfur (S) and Selenium (Se) are both in group 16. As S is in period 3 and Se is in period 4, S has the higher ionization energy.

Key concepts

The Periodic Table and Ionization Energy

The periodic table is not just a chart of all the chemical elements; it's a powerful tool for predicting various properties of elements, including ionization energy. Ionization energy refers to the energy needed to remove an electron from an atom in its gaseous state.

Understanding the layout of the periodic table helps us anticipate how easily an electron can be stripped from an atom. This is crucial because it affects the atom's reactivity and the formation of chemical bonds. Ionization energies tend to increase as you move from left to right across a period and usually decrease as you descend down a group. This trend occurs because of the effective nuclear charge and electron shielding, which we'll dissect further in the atomic structure.

For example, in the exercise provided, knowing that aluminum is to the left and higher in the periodic table compared to indium demonstrates why it has a higher ionization energy. The same concepts help us determine that among chlorine and antimony, chlorine has a higher ionization energy because it's located further to the right on the table.

Atomic Structure Influences on Ionization Energy

Delving into atomic structure sheds light on why ionization energy trends exist in the periodic table. An atom comprises a nucleus containing protons and neutrons, surrounded by electrons orbiting in 'shells.' The closer an electron is to the nucleus, the more strongly it is attracted to the positively charged protons.

Effective Nuclear Charge

The effective nuclear charge is the net positive charge experienced by valence electrons. As you move across a period, protons are added to the nucleus, ramping up the effective nuclear charge. This stronger attraction makes it harder to remove an electron, thereby increasing ionization energy.

Electron Shielding

Electron shielding occurs when inner electrons block the attraction between the nucleus and the outer electrons. Down a group in the periodic table, additional electron shells are added, increasing this shielding effect, and reducing the effective nuclear charge felt by the outer electrons, leading to lower ionization energy.

Thus, sulfur having fewer electron shells than selenium aligns with sulfur's higher ionization energy.

Electron Removal and Ionization Energy

Electron removal, or ionization, is an endothermic process. It demands energy to overcome the electrostatic forces that hold electrons in place. The concept of ionization energy becomes practical when elements undergo chemical reactions. Atoms with lower ionization energies tend to lose electrons more readily and form positive ions, like potassium in comparison to germanium in the given exercise.

When germanium's ionization energy is compared with potassium's, germanium's is higher because being in the same period, germanium has more protons creating a greater effective nuclear charge and, hence, more strongly attracts its electrons. This illustrates why elements with a higher atomic number in the same period tend to have higher ionization energies, influenced by both the atomic structure and the position on the periodic table.