3.4 Electronic Structure of Atoms (Electron Configurations) | Chemistry: Atoms First 2e

Learning Objectives

Derive the predicted ground-state electron configurations of atoms
Identify and explain exceptions to predicted electron configurations for atoms and ions
Relate electron configurations to element classifications in the periodic table

Core Concepts & Principles

The specific arrangement of electrons in an atom's atomic orbitals dictates its chemical properties. Understanding how these orbitals fill requires looking at quantum numbers, subshell penetration, and electron-electron repulsions.

Orbital Energies and Penetration

As the principal quantum number ( $n$ ) increases, orbital size and energy increase. In multi-electron atoms, electron-electron repulsions cause subshells within the same shell to differ in energy in the order $s < p < d < f$ .

Shielding: Inner electrons repel outer electrons, shielding them from the full positive charge of the nucleus and making them higher in energy.
Penetration: $s$ -orbitals penetrate closer to the nucleus than $p$ , $d$ , or $f$ orbitals, making them more tightly bound and lower in energy. For small atoms, the $n$ trend dominates, but for larger atoms, overlaps occur (e.g., $4s$ fills before $3d$ ).

The Building-Up Rules

Aufbau Principle: Electrons are added sequentially to the lowest available energy subshells as atomic number increases.
Pauli Exclusion Principle: No two electrons in an atom can share the exact same set of four quantum numbers ( $n, l, m_l, m_s$ ). Consequently, any single orbital can hold a maximum of two electrons with opposite spins ( $m_s = +\frac{1}{2}$ and $-\frac{1}{2}$ ).
Hund’s Rule: When filling degenerate orbitals (orbitals of equal energy, like the three $p$ orbitals), electrons occupy them singly with parallel spins before pairing up.

Core Mental Model: Ground-State Filling

Electrons fill orbitals from lowest energy to highest energy ( $1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d \dots$ ).
Max capacity per subshell type: $s = 2$ , $p = 6$ , $d = 10$ , $f = 14$ .
Half-filled and completely filled subshells (e.g., $d^5$ , $d^{10}$ ) provide extra stability, causing occasional anomalies in transition metal configurations.

Valence vs. Core Electrons

Core Electrons: Inner-shell electrons corresponding to a noble gas configuration. They can be abbreviated using noble gas core notation (e.g., $[\text{Ne}]$ ).
Valence Electrons: Electrons in the outermost shell ( $highest\ n$ ). They govern chemical reactivity and bonding.

Electron Configurations of Ions

Cations (Positive Ions): Formed by removing electrons. For main group elements, electrons are removed from the highest energy subshell first. For transition metals, outermost $s$ -electrons are lost before $(n-1)d$ electrons.
Anions (Negative Ions): Formed by adding electrons following the standard Aufbau filling sequence.

Problem-Solving Routines & Methods

Deriving Electron Configurations and Ion States

1
Locate the element on the periodic table and identify its atomic number ( $Z$ ).
2
Follow the sequential subshell filling order ( $1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p dots$ ) or read the blocks directly from the periodic table.
3
Write out the full configuration with superscripts representing electron counts, or use noble gas core shorthand.
4
For ions, add or remove electrons from the appropriate outer subshells (remembering transition metals lose ns electrons first).

Pro-Tip: Watch out for transition metal anomalies like Cr (

[Ar]4s^13d^5

) and Cu (

[Ar]4s^13d^10

) which favor half-filled or full d-subshells.

Maximum Subshell Electron Capacity

2(2l + 1)

Calculates the maximum number of electrons an individual subshell can hold.

Variables & Constants

l

=angular momentum quantum number (

s=0, p=1, d=2, f=3

)

Practice & Concept Checks

Concept Check

What is the ground-state electron configuration of a neutral iron atom (

Fe

Z = 26

), and what is the configuration of the

Fe^{2+}

ion?

Concept Check

Why do chromium (

Cr

) and copper (

Cu

) violate the standard predicted Aufbau filling order?

Key Terms & Vocabulary

electron configurationQuantum Structure

The specific distribution of electrons among the atomic orbitals of an atom.

Example: 1s² 2s² 2p⁴ for oxygen

Aufbau principleQuantum Mechanics

A method of building ground-state electron configurations by sequentially adding electrons to the lowest available energy subshells.

Example: Filling 4s before 3d

Orbital diagramsQuantum Structure

Pictorial representations of electron configurations using boxes for orbitals and arrows for electron spins.

Example: Up and down arrows in a 1s box

Hund’s ruleQuantum Rules

States that the lowest-energy arrangement of electrons in degenerate orbitals maximizes the number of unpaired electrons with parallel spins.

Example: Placing single electrons in three separate 2p boxes before pairing

valence electronsAtomic Properties

Electrons occupying the outermost principal shell of an atom that determine chemical reactivity.

Example: The 3s¹ electron in sodium ([Ne]3s¹)

core electronsAtomic Properties

Inner-shell electrons that correspond to a noble gas configuration.

Example: The [Ne] core in sodium

valence shellAtomic Properties

The outermost electron shell of an atom containing the highest principal quantum number (n).

Example: The n = 3 shell in phosphorus

Main group elementsPeriodic Table

Elements in which the last added electron enters an s or p orbital in the outermost shell.

Example: Groups 1, 2, and 13–18

representative elementsPeriodic Table

Alternative terminology for main group elements, encompassing all s-block and p-block elements.

Example: Carbon, fluorine, potassium

Transition elements or transition metalsPeriodic Table

Metallic elements in which the last electron added enters a d orbital, or possessing partially filled d orbitals.

Example: Iron, copper, titanium

Inner transition elementsPeriodic Table

Metallic elements in which the last electron added occupies an f orbital, consisting of the lanthanides and actinides.

Example: Uranium, neodymium