Electron Configurations and Orbital Diagrams in GOB Chemistry

Electron Orbital Diagrams

Introduction to Electron Orbital Diagrams

Electron orbital diagrams visually represent the arrangement of electrons within atomic orbitals. Understanding these diagrams is essential for predicting chemical behavior and bonding.

• Orbitals: Regions in an atom where electrons are likely to be found.

• Degenerate orbitals: Orbitals within the same subshell that have the same energy.

• Hund's Rule: Electrons fill degenerate orbitals singly first (with parallel spins) before pairing up.

Key Points

• Degenerate orbitals: Electrons in the same set of orbitals have same energy.

• Hund's Rule: Orbitals are first half-filled before they are totally filled.

Electron Orbital Diagram Table

Subshell	Number of Orbitals	Maximum Number of Electrons
s	1	2
p	3	6
d	5	10
f	7	14

Example

• Example: Properly fill in the orbitals of an atom that possesses 8 electrons within its set of orbitals.

• Application: Oxygen (Z = 8) electron configuration: $1s^2 2s^2 2p^4$

Ground State Electron Configurations

Introduction to Ground State Electron Configurations

Ground state electron configurations describe the arrangement of electrons in the lowest energy state of an atom, following specific principles.

• Auf Bau Principle: Electrons fill the lowest energy orbitals first before moving to higher energy orbitals.

• Electron configuration: Notation that shows the distribution of electrons among the orbitals.

Auf Bau Principle Diagram

• Electrons fill orbitals in the order: $1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d \rightarrow 4p$ ...

Periodic Table Blocks

• s-block: Groups 1A and 2A

• p-block: Groups 3A to 8A

• d-block: Transition metals

• f-block: Lanthanides and actinides

Example

• Example: Write the ground state electron configuration for Fluorine (Z = 9): $1s^2 2s^2 2p^5$

Unpaired vs. Paired Electrons

Introduction to Unpaired and Paired Electrons

Electrons in orbitals can be paired or unpaired, affecting the magnetic properties and reactivity of atoms.

• Unpaired Electron: An orbital contains one electron with its own spin.

• Paired Electron: An orbital contains two electrons with opposite spins.

Unpaired vs. Paired Electrons Table

Type	s orbital	p orbitals
Unpaired
Paired

Example

• Example: Determine the number of unpaired electrons in vanadium (Z = 23): $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^3$ (3 unpaired electrons in 3d).

Practice Problems and Applications

Practice: Hund's Rule Violations

• Identify electron configurations that violate Hund's Rule (e.g., pairing electrons in degenerate orbitals before each is singly occupied).

Practice: Auf Bau Principle Violations

• Identify electron configurations that violate the Auf Bau Principle (e.g., filling higher energy orbitals before lower ones).

Practice: Element Identification from Orbital Diagrams

• Given an orbital diagram, determine the element based on the number and arrangement of electrons.

Practice: Writing Electron Configurations

• Write the electron configuration and orbital diagram for elements such as Sulfur (Z = 16): $1s^2 2s^2 2p^6 3s^2 3p^4$

• Write the ground state electron configuration for Magnesium (Z = 12): $1s^2 2s^2 2p^6 3s^2$

Practice: Unpaired Electrons

• Determine which atoms have unpaired electrons and which have the most unpaired electrons.

Summary Table: Principles and Rules

Principle/Rule	Description	Example
Hund's Rule	Electrons occupy degenerate orbitals singly before pairing	$2p^3$: one electron in each p orbital
Auf Bau Principle	Electrons fill lowest energy orbitals first	$1s^2 2s^2 2p^6$
Pauli Exclusion Principle	No two electrons in an atom can have the same set of quantum numbers	Each orbital holds max 2 electrons with opposite spins

Additional info: The notes and questions are directly relevant to GOB Chemistry, focusing on electron configurations, orbital diagrams, and fundamental principles governing electron arrangement in atoms.