Electron Configurations and Periodic Trends: Study Notes for GOB Chemistry

Electron Configurations

Introduction to Electron Configurations

Electron configurations describe the arrangement of electrons in the orbitals of an atom. The order in which sublevels are filled is determined by their increasing energy, which can be visualized using the periodic table. Understanding electron configurations is essential for predicting chemical properties and reactivity.

• Electron configuration is the notation that shows the distribution of electrons among the atomic orbitals.

• Electrons fill orbitals in order of increasing energy, following the Aufbau principle.

• The periodic table is divided into blocks (s, p, d, f) that correspond to the type of sublevel being filled.

Example: The electron configuration for carbon is $1s^2 2s^2 2p^2$.

Blocks of the Periodic Table

The periodic table is organized into blocks based on the type of atomic orbital that is being filled with electrons.

• s block: Groups 1A (1) and 2A (2); final electrons enter s sublevel.

• p block: Groups 3A (13) to 8A (18); final electrons enter p sublevel.

• d block: Transition elements; final electrons enter d sublevel.

• f block: Inner transition elements; final electrons enter f sublevel.

Orbital Diagrams

Introduction to Orbital Diagrams

Orbital diagrams use boxes to represent atomic orbitals and arrows to represent electrons. They provide a visual way to show how electrons are distributed among orbitals and their spins.

• Each box represents an orbital; each arrow represents an electron.

• Electrons in the same orbital have opposite spins (Pauli exclusion principle).

• Orbitals within the same sublevel are filled singly before pairing (Hund's rule).

Example: The orbital diagram for carbon (6 electrons):

• 1s: two paired electrons

• 2s: two paired electrons

• 2p: two unpaired electrons (each in a separate box)

Step-by-Step Guide to Drawing Orbital Diagrams

1. Draw boxes to represent the occupied orbitals.

2. Place a pair of electrons with opposite spins in each filled orbital.

3. Place remaining electrons in the last occupied sublevel in separate orbitals, with parallel spins.

Example: For nitrogen (atomic number 7):

• 1s: two paired electrons

• 2s: two paired electrons

• 2p: three unpaired electrons (each in a separate box, all with the same spin)

Electron Configurations for Periods 1 and 2

Period 1: Hydrogen and Helium

Element	Atomic Number	Orbital Diagram	Electron Configuration
H	1	1s1	$1s^1$
He	2	1s2	$1s^2$

Period 2: Lithium to Neon

Element	Atomic Number	Orbital Diagram	Electron Configuration	Abbreviated Electron Configuration
Li	3	1s22s1	$1s^2 2s^1$	[He]$2s^1$
Be	4	1s22s2	$1s^2 2s^2$	[He]$2s^2$
B	5	1s22s22p1	$1s^2 2s^2 2p^1$	[He]$2s^2 2p^1$
C	6	1s22s22p2	$1s^2 2s^2 2p^2$	[He]$2s^2 2p^2$
N	7	1s22s22p3	$1s^2 2s^2 2p^3$	[He]$2s^2 2p^3$
O	8	1s22s22p4	$1s^2 2s^2 2p^4$	[He]$2s^2 2p^4$
F	9	1s22s22p5	$1s^2 2s^2 2p^5$	[He]$2s^2 2p^5$
Ne	10	1s22s22p6	$1s^2 2s^2 2p^6$	[He]$2s^2 2p^6$

Key Terms and Principles

• Aufbau Principle: Electrons occupy the lowest energy orbitals first.

• Pauli Exclusion Principle: No two electrons in the same atom can have the same set of four quantum numbers; each orbital holds a maximum of two electrons with opposite spins.

• Hund's Rule: Electrons fill degenerate orbitals singly before pairing up.

Summary Table: Electron Configuration Notation

Notation	Description
$1s^2$	1 = principal energy level; s = type of orbital; 2 = number of electrons in that orbital
[He]	Abbreviated notation using the noble gas core

Additional Info

• Valence electrons are the electrons in the outermost energy level and are important for chemical bonding and reactivity.

• Electron configurations can be used to explain periodic trends such as atomic size, ionization energy, and metallic character (covered in later sections).