Structure 1.3.5—Each orbital has a defined energy state for a given electron configuration and chemical environment, and can hold two electrons of opposite spin.

Structure 1.3.5—Each orbital has a defined energy state for a given electron configuration and chemical environment, and can hold two electrons of opposite spin.
Sublevels contain a fixed number of orbitals, regions of space where there is a high probability of
finding an electron.

What You’ll Learn:

• Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to deduce electron configurations for atoms and ions up to Z = 36
• Full electron configurations and condensed electron configurations using the noble gas core should be covered.
• Orbital diagrams, i.e. arrow-in-box diagrams, should be used to represent the filling and relative energy of orbitals.
• The electron configurations of Cr and Cu as exceptions should be covered.

Keywords

Each atomic orbital has a defined energy state for a given electron configuration and chemical environment. An orbital can hold a maximum of two electrons, which must have opposite spins. Sublevels are groups of orbitals within a given energy level, and they have a fixed number of orbitals where electrons have a high probability of being found.

To determine electron configurations for atoms and ions up to Z = 36 (krypton), we apply the Aufbau principle, Hund’s rule, and the Pauli exclusion principle:

1. Aufbau Principle: Electrons fill orbitals in order of increasing energy, starting from the lowest energy orbital.
2. Hund’s Rule: When filling degenerate orbitals (orbitals with the same energy) in a sublevel, electrons occupy each orbital singly with parallel spins before any orbital is doubly occupied.
3. Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers, meaning that each orbital can hold a maximum of two electrons with opposite spins.

To deduce electron configurations, follow the order of orbital filling: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

Full electron configurations list all the orbitals in the order they are filled, while condensed electron configurations use the noble gas core as a shorthand notation. For example, the electron configuration of chlorine (Z = 17) can be written as [Ne]3s²3p⁵, where [Ne] represents the electron configuration of neon (Z = 10).

Orbital diagrams, also known as arrow-in-box diagrams, visually represent the filling and relative energy of orbitals. In these diagrams, boxes represent orbitals, and arrows represent electrons. The direction of the arrow indicates the electron’s spin.

Chromium (Z = 24) and copper (Z = 29) are notable exceptions to the usual electron configurations due to their increased stability with half-filled or fully filled d subshells. Chromium’s electron configuration is [Ar]4s¹3d⁵, while copper’s electron configuration is [Ar]4s¹3d¹⁰, where [Ar] represents the electron configuration of argon (Z = 18).

Questions

1. What is the maximum number of electrons that can be present in a sublevel?
2. What is the Pauli exclusion principle and how does it apply to electron configurations?
3. Explain the difference between full and condensed electron configurations.
4. Why do chromium and copper have unusual electron configurations?
5. Describe the Aufbau principle and how it is used to determine electron configurations.
6. What is the difference between a sublevel and an orbital?
7. Explain the concept of degenerate orbitals and how they are filled.
8. How do you represent electron configurations using orbital diagrams?
9. Why is the electron configuration of chlorine written as [Ne]3s²3p⁵ instead of 1s²2s²2p⁶3s²3p⁵?
10. What is the significance of the noble gas core in condensed electron configurations?