Consider these elements: P, Ca, Si, S, Ga. a. Write the electron configuration for each element. b. Arrange the elements in order of decreasing atomic radius. c. Arrange the elements in order of increasing ionization energy. d. Use the electron configurations in part a to explain the differences between your answers to parts b and c.

Verified step by step guidance

1. Step 1: Write the electron configuration for each element. Use the periodic table to determine the number of electrons for each element and fill the orbitals in the order of increasing energy (using the Aufbau principle).
2. Step 2: Arrange the elements in order of decreasing atomic radius. Recall that atomic radius generally decreases across a period from left to right and increases down a group in the periodic table.
3. Step 3: Arrange the elements in order of increasing ionization energy. Remember that ionization energy generally increases across a period from left to right and decreases down a group.
4. Step 4: Compare the electron configurations from Step 1 to explain the differences in atomic radius. Consider the number of electron shells and the effective nuclear charge experienced by the outermost electrons.
5. Step 5: Use the electron configurations to explain the differences in ionization energy. Consider the stability of the electron configuration and the energy required to remove an electron from the outermost shell.

Key Concepts

Electron Configuration

Electron configuration describes the distribution of electrons in an atom's orbitals. It is represented using a notation that indicates the energy levels and sublevels occupied by electrons, following the Aufbau principle, Pauli exclusion principle, and Hund's rule. Understanding electron configurations is essential for predicting an element's chemical behavior and properties, including its atomic radius and ionization energy.

Atomic Radius

Atomic radius is a measure of the size of an atom, typically defined as the distance from the nucleus to the outermost electron shell. Atomic radius generally increases down a group in the periodic table due to the addition of electron shells, while it decreases across a period due to increased nuclear charge, which pulls electrons closer to the nucleus. This concept is crucial for comparing the sizes of different elements.

Ionization Energy

Ionization energy is the energy required to remove an electron from an atom in its gaseous state. It generally increases across a period due to higher effective nuclear charge and decreases down a group as the outer electrons are further from the nucleus and experience less attraction. Understanding ionization energy is important for explaining trends in reactivity and the stability of electron configurations.