Electron Configuration: Rules, Notation, and Exceptions

What Is Electron Configuration?

Electron configuration describes how electrons are arranged in an atom's orbitals. It determines an element's chemical properties, reactivity, and position in the periodic table.

Every element has a unique electron configuration, written using a notation that specifies the shell, subshell, and number of electrons in each subshell.

Learning Goals: By the end of this guide, you should be able to:

1. Write electron configurations for any element (up to Period 4).
2. Apply the three key rules (Aufbau, Hund, Pauli).
3. Use the periodic table as a filling order guide.
4. Handle transition metal exceptions (Cr, Cu).
5. Write configurations for ions.

The Three Rules

1. Aufbau Principle (Building Up)

Electrons fill the lowest energy orbitals first, before moving to higher energy ones.

The filling order is:

Memory trick: Follow the diagonal arrow pattern, or simply read across the periodic table from left to right, top to bottom.

2. Pauli Exclusion Principle

Each orbital can hold a maximum of 2 electrons, and they must have opposite spins (one spin-up ↑, one spin-down ↓).

3. Hund's Rule

When filling orbitals of the same energy (degenerate orbitals, like the three $2p$ orbitals), electrons occupy them singly first, all with the same spin, before pairing up.

Why? Electrons repel each other. Spreading them across orbitals minimises repulsion.

Orbital Capacity

Writing Electron Configurations

Notation

The configuration is written as a series of subshells with superscripted electron counts:

Example: Iron (Fe, Z = 26)

Shorthand (Noble Gas Core)

Replace the completed inner shells with the noble gas symbol in brackets:

Where $[Ar] = 1s^2\ 2s^2\ 2p^6\ 3s^2\ 3p^6$.

Using the Periodic Table as a Guide

The periodic table is organised by electron configuration:

Reading across Period 4: $K$ fills $4s^1$, $Ca$ fills $4s^2$, then $Sc$ through $Zn$ fill $3d^1$ to $3d^{10}$, then $Ga$ through $Kr$ fill $4p^1$ to $4p^6$.

Transition Metal Exceptions

Two notable exceptions exist due to the extra stability of half-filled and fully-filled $d$ subshells:

One electron "moves" from $4s$ to $3d$ because the symmetrical electron distribution in half-filled and fully-filled $d$ subshells provides extra stability through exchange energy.

Electron Configurations of Ions

Forming Positive Ions (Cations)

Remove electrons from the highest principal quantum number first:

• $Fe: [Ar]\ 4s^2\ 3d^6$
• $Fe^{2+}: [Ar]\ 3d^6$ (remove 2 electrons from $4s$ first, NOT $3d$)
• $Fe^{3+}: [Ar]\ 3d^5$ (remove 2 from $4s$, then 1 from $3d$)

Critical: For ions, $4s$ electrons are removed before $3d$ electrons, even though $4s$ fills first. This is because in a charged ion, $3d$ is lower in energy than $4s$.

Forming Negative Ions (Anions)

Add electrons to the next available subshell:

• $O: 1s^2\ 2s^2\ 2p^4$
• $O^{2-}: 1s^2\ 2s^2\ 2p^6$ (electrons added to complete the $2p$ subshell)

Worked Examples

Example 1: Write the Configuration of Sulfur (S, Z = 16)

Follow the filling order:

Check: $2 + 2 + 6 + 2 + 4 = 16$ ✓

Shorthand: $[Ne]\ 3s^2\ 3p^4$

Example 2: Configuration of $Cu^{2+}$

Start with Cu: $[Ar]\ 4s^1\ 3d^{10}$ (exception)

For $Cu^{2+}$: Remove 2 electrons. First from $4s$ (1 electron), then from $3d$ (1 electron):

Example 3: Is $Fe^{2+}$ or $Fe^{3+}$ More Stable?

$Fe^{2+}: [Ar]\ 3d^6$ — no special stability.

$Fe^{3+}: [Ar]\ 3d^5$ — half-filled $d$ subshell, which is extra stable.

This explains why $Fe^{3+}$ compounds are often more stable and why $Fe^{2+}$ is easily oxidised to $Fe^{3+}$.

Common Mistakes

1. Removing $3d$ electrons first when forming ions — For transition metal ions, always remove $4s$ electrons first, then $3d$. The filling order and ionisation order are different.

2. Forgetting the Cr and Cu exceptions — These are the two most commonly tested exceptions. Memorise: Cr is $4s^1\ 3d^5$ (not $4s^2\ 3d^4$), Cu is $4s^1\ 3d^{10}$ (not $4s^2\ 3d^9$).

3. Violating Hund's rule — In the $2p$ subshell with 3 electrons, write $2p^1_x\ 2p^1_y\ 2p^1_z$ (one in each), not $2p^2_x\ 2p^1_y$ (pairing before spreading).

4. Confusing shells with subshells — Shell 3 contains $3s$, $3p$, and $3d$ subshells. They have different energies despite being in the same shell.

5. Not checking the total electron count — Always add up all superscript numbers. The total must equal the atomic number (for neutral atoms) or atomic number minus charge (for ions).

Exam Tips (A-Level / AP / IB)

• The periodic table is the filling order diagram. Read across each period: s-block (1–2 electrons), d-block (1–10 electrons), p-block (1–6 electrons).
• For "explain the exception" questions, mention exchange energy and the stability of half-filled or fully-filled subshells.
• When writing ion configurations, explicitly state: "Remove electrons from $4s$ before $3d$."
• Know that electron configuration explains periodic trends: ionisation energy, atomic radius, and electronegativity all correlate with electron arrangement.

Frequently Asked Questions

Why does 4s fill before 3d?

In neutral atoms, the $4s$ orbital is slightly lower in energy than $3d$ due to better penetration of the $4s$ orbital toward the nucleus. However, once the atom is ionised, $3d$ becomes lower in energy, which is why $4s$ electrons are removed first.

What are degenerate orbitals?

Degenerate orbitals have the same energy level. The three $2p$ orbitals ($2p_x$, $2p_y$, $2p_z$) are degenerate. Hund's rule applies to degenerate orbitals: fill them singly before pairing.

How do you write the electron configuration of a transition metal ion?

Write the configuration of the neutral atom first, then remove electrons from the $4s$ subshell before removing from $3d$. For example, $Fe^{3+}$: start with $Fe = [Ar]\ 4s^2\ 3d^6$, remove 2 from $4s$, then 1 from $3d$ → $[Ar]\ 3d^5$.

Why are half-filled subshells extra stable?

Half-filled subshells ($d^5$, $p^3$) have maximum exchange energy — a quantum mechanical stabilisation that occurs when electrons with the same spin occupy different orbitals. The symmetrical distribution also minimises electron-electron repulsion.

Related Topics

• Atomic Models — How our understanding of electron arrangement evolved from Bohr to quantum mechanics.
• Periodic Trends — How electron configuration drives patterns in ionisation energy, electronegativity, and atomic radius.
• Orbital Hybridisation — How filled orbitals combine to explain molecular geometry.