Question

Arrange each set of elements in order of increasing ionization cnergy. a. \(\mathrm{F}, \mathrm{Cl}, \mathrm{Br}\) b. \(\mathrm{Na}, \mathrm{Cl}, \mathrm{Al}\) c. \(\mathrm{Na}, \mathrm{K}, \mathrm{Cs}\) d. As, \(\mathrm{Ca}, \mathrm{Br}\)

Short answer

a. \[ Br < Cl < F \] b. \[ Na < Al < Cl \] c. \[ Cs < K < Na \] d. \[ Ca < As < Br \]

Step-by-step solution

Understanding Ionization Energy

Ionization energy is the energy required to remove an electron from an atom in the gaseous state. Elements with higher ionization energy hold onto their electrons more tightly.

Trend in the Periodic Table

Ionization energy generally increases across a period (left to right) and decreases down a group (top to bottom).

Arrange Set (a) \( F, Cl, Br \)

Fluorine (F), Chlorine (Cl), and Bromine (Br) are in Group 17. Ionization energy decreases down the group. Increasing order: \[ Br < Cl < F \]

Arrange Set (b) \( Na, Cl, Al \)

Sodium (Na), Chlorine (Cl), and Aluminum (Al) span across Period 3. Ionization energy increases as you move from left (Na) to right (Cl). Increasing order: \[ Na < Al < Cl \]

Arrange Set (c) \( Na, K, Cs \)

Sodium (Na), Potassium (K), and Cesium (Cs) are in Group 1. Ionization energy decreases down the group. Increasing order: \[ Cs < K < Na \]

Arrange Set (d) As, Ca, Br

Arsenic (As), Calcium (Ca), and Bromine (Br) are different in terms of position on Periodic Table. Noting the trends, we arrange them: Increasing order: \[ Ca < As < Br \]

Key concepts

Periodic Table Trends

Periodic table trends are key patterns that help us understand various elements' properties. One important trend is ionization energy, which refers to the energy required to remove an electron from an atom. Across a period from left to right, ionization energy generally increases. This is because atoms have more protons, increasing the positive charge and pulling electrons closer. This means more energy is needed to remove an electron. Conversely, as you move down a group (top to bottom), ionization energy decreases. This happens because additional electron shells are added, increasing the distance between the nucleus and the outermost electrons. As a result, these electrons are held less tightly, making them easier to remove.

Element Groups

Element groups refer to columns in the periodic table that share similar properties. For example, Group 1 elements are called alkali metals, such as Sodium (Na), Potassium (K), and Cesium (Cs). They have a single electron in their outermost shell, making them highly reactive. The ionization energy for alkali metals decreases as you go down the group because the outer electron is farther from the nucleus, making it easier to remove.
Group 17 elements, known as halogens, include Fluorine (F), Chlorine (Cl), and Bromine (Br). They have seven electrons in their outer shell and are very electronegative. Their ionization energy decreases down the group for similar reasons as the alkali metals. Knowing these group characteristics can help predict the behaviors and reactivities of elements in the same group.

Atomic Structure

Atomic structure is fundamental to understanding ionization energy. Each atom consists of a nucleus containing protons and neutrons, surrounded by electrons in different shells or energy levels. The number of protons (atomic number) directly influences the ionization energy. More protons mean a stronger positive charge, pulling electrons closer. Thus, removing an electron requires more energy. The electron configuration also plays a role. Electrons in inner shells shield the outer electrons from the nucleus's full attraction, impacting ionization energy.
The concept of electron shielding explains why ionization energy decreases going down a group. As more electron shells are added, the outermost electrons are less tightly bound to the nucleus, making them easier to remove. Understanding these basic principles of atomic structure helps explain why different elements have varying ionization energies.