Question

(a) What is the trend in first ionization energies as one proceeds down the group 7A elements? Explain how this trend relates to the variation in atomic radii. (b) What is the trend in first ionization energies as one moves across the fourth period from \(\mathrm{K}\) to \(\mathrm{Kr}\) ? How does this trend compare with the trend in atomic radii?

Short answer

(a) As we move down the group 7A elements, the first ionization energy decreases due to the increase in atomic radii. As the atomic radius increases, the outermost electrons are further away from the nucleus and thus easier to remove. (b) Across the fourth period from K to Kr, the first ionization energy increases as the atomic radii generally decrease. A smaller atomic radius means that electrons are closer to the nucleus and more tightly held, making it harder to remove them.

Step-by-step solution

(a) Trend in first ionization energies for group 7A elements

As we move down the group 7A elements, the first ionization energy generally decreases. This means that it becomes easier to remove an electron from the elements as we go down the group.

Explanation of the trend in relation to atomic radii

This trend can be explained by the increase in atomic radii as we move down the group. As the atomic radius increases, the outermost electrons are further away from the nucleus, making it easier to remove them because they are less tightly held by the nuclear charge (things like electron shielding also contribute to this effect). Therefore, the increase in atomic radii leads to a decrease in first ionization energy down group 7A.

(b) Trend in first ionization energies across the fourth period from K to Kr

As we move across the fourth period from K to Kr, the first ionization energy generally increases. This means that it becomes more challenging to remove an electron from the elements as we go from left to right across the period.

Comparison of the trend in first ionization energies and atomic radii for K to Kr

As we move from K to Kr across the fourth period, the atomic radii generally decrease. A smaller atomic radius means that the electrons are closer to the nucleus and experience a stronger pull from the nuclear charge. As a result, these electrons are more tightly held, making it harder to remove them. This explains the increase in the first ionization energy across the fourth period, as it is harder to remove electrons from elements with smaller atomic radii.

Key concepts

Group 7A Elements

The elements that occupy Group 7A, also known as the halogens, follow a distinct trend in which their first ionization energy—the energy required to remove the outermost electron—decreases as you move down the group on the periodic table. For instance, removing an electron from fluorine (at the top) is generally harder than removing one from iodine (located further down).

This can be largely attributed to the increasing atomic radii. As the atoms get larger, the outermost electrons are further from the nucleus and are less tightly bound by the nuclear charge due to increased distance and the screening effect, which is the shielding of outer electrons from the nucleus by inner electrons. This makes it easier for an electron to escape, lowering the ionization energy. Understanding this concept is crucial for students in predicting reactivity and bonding behavior of the halogens.

Atomic Radii Variation

Atomic radii—measured as the distance from the nucleus to the boundary of the surrounding cloud of electrons—vary across the periodic table and are influenced by atomic number and electron configuration. In Group 7A, the atomic radii increase from top to bottom. This extends the distance between the outermost electrons and the nucleus, weakening their electrostatic interaction. As a result, these outer electrons are more easily removed, which reflects in lower ionization energies.

Contrastingly, across the fourth period, atomic radii decrease as we move from left to right. This happens because electrons are being added to the same energy level while the number of protons increases, creating a greater nuclear charge pulling the electron cloud in closer.

Fourth Period Elements

The fourth period of the periodic table features a trend in first ionization energy that increases from the element potassium (K) to krypton (Kr). A close examination reveals that despite electrons being added within the same principal quantum level, the effective nuclear charge experienced by the outer electrons increases. This is due to the overall increase in protons, which results in a stronger pull on the electron cloud. The increasing effective nuclear charge as one moves across the period explains the increase in ionization energy and the simultaneous decrease in atomic radius.

The implications of this are extensive, from determining element reactivity to predicting the types of chemical bonds an element is likely to form. In the context of learning about properties of elements, it's key to recognize how periodic trends like these are essential to understanding the chemical behavior of elements.

Electron Shielding Effect

The electron shielding effect is the phenomenon where inner shell electrons reduce the effective nuclear charge on the outer shell electrons. It significantly affects both atomic radius and ionization energy. As the number of electron shells increases, the effective nuclear charge felt by the outermost electrons diminishes due to the shielding provided by the inner layers of electrons.

Understanding the electron shielding effect is fundamental as it explains why atoms increase in size down a group and why the first ionization energy generally decreases. Despite an increase in the number of protons, the additional inner electrons shield the outer electrons from the full charge, making them easier to remove. This concept links the group and periodic trends in ionization energy and helps students to grasp the underlying principles governing element reactivity.