Electron configuration rules

IB Chemistry SL· Structure 2.2: Electron configuration· 40 min read

1. Pauli Exclusion Principle★★☆☆☆⏱ 10 min

📘 Definition

Pauli Exclusion Principle

No two electrons in the same atom can have an identical set of four quantum numbers. This limits every atomic orbital to a maximum of two electrons, which must have opposite spin.

Example:

A single 1s orbital can hold 0, 1, or 2 electrons, never 3.

Spin is an intrinsic quantum property of electrons, categorized as up ( $+ \frac{1}{2}$ ) or down ( $- \frac{1}{2}$ ). Opposite spins allow two electrons to occupy the same orbital without violating the principle.

📐 Worked Example

Which rule is violated by the electron configuration $1 s^{3} 2 s^{2} 2 p^{6}$ ?

1
1. Recall the maximum number of electrons per orbital from the Pauli principle:
2
$1 \text{ orbital} = 2 \text{ electrons maximum}$
3
1. The 1s subshell contains only one orbital, so it can hold at most 2 electrons. The given configuration has 3 electrons in 1s.
4
1. This violates the Pauli exclusion principle, as three electrons would require two to share the same set of quantum numbers.

Exam tip:

When asked to identify rule violations on Paper 1, always count electrons per orbital first.

2. Aufbau Principle★★☆☆☆⏱ 15 min

📘 Definition

Aufbau Principle

Electrons fill atomic orbitals starting from the lowest available energy level, moving to higher energy levels only after lower levels are full.

Example:

4s fills before 3d for period 4 elements, because 4s has lower energy than 3d when empty.

The energy order of orbitals can be found using the $n + l$ rule: orbitals with lower $(n + l)$ values fill first. If two orbitals have the same $(n + l)$ value, the orbital with lower $n$ fills first.

Filling order for the first 36 elements: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p
When transition metals form ions, 4s electrons are lost before 3d electrons

📐 Worked Example

Use the $n + l$ rule to order 3s, 3p, 3d, 4s by filling energy

1
1. Calculate $n + l$ for each orbital (l=0 for s, 1 for p, 2 for d):
2
- 3s: $n = 3, l = 0$ → $n + l = 3$
3
- 3p: $n = 3, l = 1$ → $n + l = 4$
4
- 4s: $n = 4, l = 0$ → $n + l = 4$
5
- 3d: $n = 3, l = 2$ → $n + l = 5$
6
1. Order by increasing $n + l$ , then lower $n$ for equal $n + l$ :
7
$3s < 3p < 4s < 3d$

3. Hund's Rule of Maximum Multiplicity★★★☆☆⏱ 15 min

📘 Definition

Hund's Rule

For degenerate (equal energy) orbitals, electrons fill each orbital singly with parallel spin before any orbital receives a second paired electron.

Example:

Three electrons in the 2p subshell occupy all three p orbitals singly, rather than pairing in one.

Single occupation with parallel spin minimizes electron-electron repulsion, resulting in a lower energy, more stable configuration. This rule explains why many atoms have unpaired electrons, leading to paramagnetic behavior.

📐 Worked Example

Draw a valid orbital box diagram for nitrogen (Z=7) per Hund's rule

1
1. Fill lower energy orbitals first: 1s holds 2 electrons, 2s holds 2 electrons, leaving 3 electrons for the three degenerate 2p orbitals.
2
1. Add one electron to each 2p orbital, all with the same (parallel) spin.
3
We do not pair any electrons in the 2p orbitals, as this would increase repulsion and violate Hund's rule.

✓ Quick check

Test your understanding of the three rules:

Which rule is violated if two electrons are paired in a 2p orbital before all three 2p orbitals have one electron?
- Pauli Exclusion Principle
- Aufbau Principle
- Hund's Rule
- No rule is violated
Reveal answer
Hund's Rule —
Correct! Hund's rule requires single occupation of all degenerate orbitals before pairing.

4. Common Pitfalls

Wrong move:

Writing 3d before 4s in the configuration of neutral potassium

Why:

Orders orbitals by n value instead of following the Aufbau filling order

Correct move:

Write configurations in filling order: [Ar] 4s¹ for potassium, not [Ar] 3d¹

Wrong move:

Placing 3 electrons in a single p orbital

Why:

Forgets the Pauli exclusion principle limits each orbital to 2 electrons

Correct move:

Distribute electrons across all available degenerate p orbitals, max 2 per orbital

Wrong move:

Removing 3d electrons before 4s when writing transition metal ion configurations

Why:

Confuses filling order with ionization order; 4s fills first but loses electrons first

Correct move:

Fe²+ has configuration [Ar] 3d⁶, not [Ar] 4s² 3d⁴

Wrong move:

Writing copper's configuration as [Ar] 4s² 3d⁹

Why:

Ignores the extra stability of full d-subshells that causes an exception to Aufbau

Correct move:

The correct stable configuration is [Ar] 4s¹ 3d¹⁰

5. Quick Reference Cheatsheet

Rule	Core Statement
Pauli Exclusion	Max 2 electrons per orbital, opposite spin
Aufbau Principle	Fill lowest energy orbitals first
Hund's Rule	Single parallel electrons in degenerate orbitals first
Filling order (Z < 36)	1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p
Common exceptions	Cr = [Ar]4s¹3d⁵, Cu = [Ar]4s¹3d¹⁰
Ionization rule	Lose 4s electrons before 3d electrons

6. Frequently Asked

Why are chromium and copper exceptions to the Aufbau principle?

Half-filled ( $d^{5}$ ) and fully filled ( $d^{10}$ ) d-subshells have extra quantum mechanical stability, so one electron is promoted from 4s to 3d to achieve this lower energy configuration.

When this came up on past exams

AI-estimated based on syllabus patterns — cross-check with official past papers for accuracy. Use only as revision-focus signals.

2022 · 1
Identify violation of Hund's rule
2023 · 2
State rules for electron configuration

Going deeper

syllabusOfficial IB Chemistry SL SyllabusStructure 2 reference

What's Next

Now that you have mastered the core rules for electron configuration, you are ready to apply these rules to write full, condensed, and orbital box configurations for any element up to Z=36, which is the required scope for IB Chemistry SL. Electron configuration is the foundation for all subsequent topics in IB Chemistry, including periodic trends, bonding, and chemical reactivity, as the arrangement of valence electrons directly determines an element's chemical behavior. Common exceptions like chromium and copper are regularly tested on both Paper 1 and Paper 2, so ensure you memorize these key cases.