Electron orbitals and configuration

CIE A-Level Chemistry· Unit 2: Atomic structure· 20 min read

1. Quantum Structure: Shells, Subshells and Orbitals★★☆☆☆⏱ 4 min

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📘 Definition

Atomic Orbital

A 3-dimensional region of space around the nucleus where there is a 95% probability of finding an electron with a specific fixed energy.

Example:

A 1s orbital is a spherical region closest to the nucleus.

In the modern quantum model, electrons are arranged in principal shells (energy levels) labelled by the principal quantum number $n = 1, 2, 3...$ . Lower $n$ corresponds to lower energy and smaller distance from the nucleus. Each principal shell splits into subshells, which each contain a fixed number of orbitals.

Subshell letter	Number of orbitals	Maximum electrons	First appears at n=
s	1	2	1
p	3	6	2
d	5	10	3
f	7	14	4

📐 Worked Example

Calculate the maximum number of electrons that can occupy the $n = 3$ principal shell.

1
Identify all subshells present in $n = 3$ :
2
$3s, 3p, 3d$
3
Find the maximum electrons per subshell:
4
3s = 2, 3p = 6, 3d = 10
5
Sum the total number of electrons:
6
$2 + 6 + 10 = 18$

2. Rules for Filling Orbitals★★☆☆☆⏱ 5 min

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📘 Definition

Aufbau Principle

Electrons fill orbitals in order of increasing energy: lower energy orbitals are filled completely before higher energy orbitals are occupied.

Pauli Exclusion Principle: Each orbital holds a maximum of 2 electrons with opposite spins.
Hund's Rule: When filling degenerate (equal energy) orbitals (e.g. the three 2p orbitals), electrons fill each orbital singly first with parallel spins, before pairing up. This minimizes electron-electron repulsion.

For elements up to Z=36 (krypton), the order of filling orbitals by increasing energy is:

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p

📐 Worked Example

Draw the orbital filling diagram for neutral nitrogen (Z=7).

1
Total number of electrons = atomic number = 7.
2
Fill lowest energy orbitals first: 1s holds 2 paired electrons, 2s holds 2 paired electrons, leaving $7 - 4 = 3$ electrons for 2p.
3
Apply Hund's rule: 2p has 3 degenerate orbitals, so add one unpaired electron to each orbital with parallel spin.

Exam tip:

CIE examiners regularly test application of Hund's rule, so always check for single occupation before pairing electrons in orbitals.

3. Writing Electron Configurations★★★☆☆⏱ 5 min

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Electron configurations can be written in full, or abbreviated using noble gas core shorthand. For CIE exams, you are expected to write configurations for all atoms and ions up to Z=36.

📐 Worked Example

Write the full and abbreviated electron configuration for neutral iron (Z=26).

1
Total electrons = 26, fill in the energy order: $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{2} 3 d^{6}$
2
Rearrange to conventional order sorted by $n$ :
3
$1s^2 2s^2 2p^6 3s^2 3p^6 3d^6 4s^2$
4
For abbreviated form, replace the core electrons with the symbol of the previous noble gas (argon, Z=18):
5
$[Ar] 3d^6 4s^2$

✓ Quick check

Test your understanding of transition metal ion configurations:

What is the correct electron configuration of $F e^{2 +}$ ?
- [Ar] 3d^6
- [Ar] 3d^4 4s^2
- [Ar] 3d^6 4s^2
- [Ar] 3d^8
Reveal answer
[Ar] 3d^6 —
4s electrons are always lost before 3d electrons when transition metals form positive ions. Neutral Fe has 2 electrons in 4s, so both are lost to form Fe²⁺.

4. Exceptions to the Aufbau Principle★★★★☆⏱ 4 min

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For CIE 9701, you only need to remember two common exceptions to the standard filling order for first-row transition metals: chromium (Z=24) and copper (Z=29). These exceptions occur because a half-filled (5 electrons) or fully filled (10 electrons) 3d subshell has extra stability from symmetry and exchange energy.

📐 Worked Example

Write the correct electron configuration for copper (Z=29) and explain why it is an exception.

1
Standard filling would give the configuration $[A r] 3 d^{9} 4 s^{2}$ .
2
A fully filled 3d subshell (10 electrons) has extra stability compared to a $3 d^{9}$ arrangement.
3
One electron from 4s is promoted to 3d to achieve this stable configuration.
4
The correct configuration is:
5
$[Ar] 3d^{10} 4s^1$

Exam tip:

CIE almost always tests these two exceptions, so memorise: Cr = [Ar] 3d⁵ 4s¹, Cu = [Ar] 3d¹⁰ 4s¹.

5. Common Pitfalls

Wrong move:

Claims 4s is always written before 3d because it fills first

Why:

Confuses the order of filling with the conventional writing order sorted by principal quantum number n

Correct move:

4s fills before 3d, but 3d has lower n so it is written before 4s in the final configuration

Wrong move:

Writes chromium as [Ar] 3d⁴ 4s²

Why:

Forgets the extra stability of a half-filled 3d subshell, the key exception tested by CIE

Correct move:

The correct configuration for chromium is [Ar] 3d⁵ 4s¹

Wrong move:

Removes 3d electrons first when forming transition metal ions

Why:

Assumes 4s is always lower energy than 3d, which is only true before 3d starts filling

Correct move:

After 3d fills, 4s is higher energy, so 4s electrons are always lost first when ionising

Wrong move:

Defines an orbital as a fixed path electrons follow

Why:

Confuses the old Bohr model with the modern quantum model

Correct move:

An orbital is a region of space where there is ~95% probability of finding an electron of a given energy

Wrong move:

Pairs electrons in degenerate orbitals before all have one electron

Why:

Violates Hund's rule of maximum multiplicity

Correct move:

Fill all degenerate orbitals with one parallel-spin electron before pairing any electrons

6. Quick Reference Cheatsheet

Concept	Key CIE Exam Fact
Filling order	1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p
Writing order	Sort by n: 1s < 2s < 2p < 3s < 3p < 3d < 4s < 4p
Max electrons per subshell	s=2, p=6, d=10, f=14
Common exceptions	Cr: [Ar] 3d⁵ 4s¹; Cu: [Ar] 3d¹⁰ 4s¹
Ion electron loss	4s electrons lost before 3d electrons
Orbital definition	Region of 95% probability of finding an electron

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
MCQ on electron configuration of ions
2023 · 2
Write configurations for transition metals
2021 · 1
Orbital definition and shape question

What's Next

Electron configuration is the foundation of all modern chemistry, explaining periodic trends, bonding patterns, chemical reactivity and the unique properties of transition metals. This sub-topic is assessed in almost every CIE 9701 paper, appearing in both multiple choice and structured questions, so mastering the rules and exceptions is critical for achieving a high grade. Understanding how electron arrangement relates to ionisation energy will help you tackle periodic trend questions later in the course.