Question

(a) What is the trend in first ionization energies as one proceeds down the group 7 A 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 proceed down Group 7A elements, the first ionization energy decreases due to the increase in atomic radius. The larger atomic radius weakens the electrostatic force of attraction between the nucleus and the outer electron, making it easier to remove the electron and resulting in a lower ionization energy. (b) In the fourth period from $K$ to $Kr$, the first ionization energy generally increases as the atomic radii decrease. The increased number of protons in the nucleus lead to a stronger electrostatic force of attraction between the nucleus and the outer electrons. This causes an increase in ionization energy, as it becomes more difficult to remove an electron from the outermost shell.

Step-by-step solution

Define first ionization energy and atomic radius

First, let's understand the terms we are working with. The first ionization energy is the energy required to remove one electron from the outermost shell of an atom. Atomic radius refers to the size of an atom, measured as the distance between the nucleus and the outermost shell of electrons.

Analyze the trend in first ionization energies for Group 7A

As we proceed down the Group 7A elements, we notice that the first ionization energy decreases. This means that it becomes easier to remove an electron from the outermost shell of the atoms in this group as we move down the group.

Relate the trend in first ionization energies to atomic radii in Group 7A

The decrease in first ionization energy can be explained by the increase in atomic radius as we move down the group. As the atomic radius increases, the electrons in the outermost shell are farther from the nucleus, and the electrostatic force of attraction between the nucleus and the outer electron decreases. This weakens the hold of the nucleus on the outer electron, making it easier to remove and resulting in a lower ionization energy.

Analyze the trend in first ionization energies for the fourth period

When moving across the fourth period from K (potassium) to Kr (krypton), the first ionization energy generally increases. This means that it becomes more difficult to remove an electron from the outermost shell as we move across the period.

Relate the trend in first ionization energies to atomic radii in the fourth period

As we move across the fourth period, the number of protons in the nucleus increases, which increases the positive charge of the nucleus and the electrostatic force of attraction between the nucleus and the outer electrons. The atomic radii generally decrease across the period. The electrons are more tightly held by the nucleus due to the increased force of attraction. This causes an increase in ionization energy, as it becomes more difficult to remove an electron from the outermost shell.

Key concepts

Atomic Radius

The atomic radius is a key concept to understand when studying elements and their behaviors in the periodic table. The atomic radius is defined as the distance from the nucleus of an atom to the outermost shell of electrons. In simpler terms, you can think of it as the "size" of an atom.
A larger atomic radius means the outermost electrons are further away from the nucleus.
This distance can affect many properties of elements, including ionization energy, which is the energy required to remove an electron from an atom. Several factors influence the atomic radius, including:

• The number of electron shells: More shells mean a larger atomic radius.
• The effective nuclear charge: A stronger positive charge attracts electrons more tightly, potentially reducing the radius.

Understanding the atomic radius can help explain the trends in other elemental properties across periods and groups in the periodic table.

Group 7A Elements

Group 7A, also known as the halogens, consists of elements like fluorine, chlorine, bromine, iodine, and astatine. These elements are found in the second column from the right of the periodic table and are known for their high reactivity. As we look down the Group 7A elements, we observe several trends. One of these is the decrease in first ionization energy. First ionization energy decreases as you move down this group because the atomic radius increases. Here's why:

• As you go down the group, each element has an additional electron shell.
• These extra shells increase the distance between the nucleus and the outermost electrons.
• The further the electrons are from the nucleus, the weaker the attractive force, making it easier to remove an electron.

Thus, despite having more protons, the hold on the outermost electron is weaker, resulting in a lower ionization energy as we descend the group.

Periodic Trends

Periodic trends are patterns observed within the periodic table that help predict the behavior of elements. Two important trends involve ionization energy and atomic radius. When moving across a period, such as from potassium (K) to krypton (Kr), two key trends are observed:

• Increasing Ionization Energy: As the number of protons increases, the positive charge of the nucleus strengthens its pull on electrons, making it more difficult to remove an electron. Therefore, ionization energy generally increases across a period.
• Decreasing Atomic Radius: Despite having the same number of electron shells as other elements in the period, the increasing positive charge pulls the electron cloud closer, leading to a smaller atomic radius.

These trends are fundamental to understanding the reactivity and chemical properties of elements, providing a foundation for predicting how different elements might interact with one another.