Electronic Configuration

Introduction

Electronic configuration describes how electrons are distributed among the orbitals of an atom. Three fundamental rules govern this distribution.

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The Three Rules

1. Aufbau Principle

“Building up” - Electrons fill orbitals in order of increasing energy.

$$1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p$$

The (n + l) Rule

Orbitals fill in order of increasing (n + l). If (n + l) is same, lower n fills first.

Orbital	n	l	n + l	Filling order
3d	3	2	5	Later
4s	4	0	4	Earlier
4p	4	1	5	After 3d
5s	5	0	5	After 4p

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2. Pauli’s Exclusion Principle

“No two electrons can have identical set of all four quantum numbers.”

Consequences:

• Maximum 2 electrons per orbital
• Must have opposite spins (↑↓)
• Maximum electrons in:
  • Subshell: 2(2l + 1)
  • Shell: 2n²

Subshell	Max electrons
s	2
p	6
d	10
f	14

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3. Hund’s Rule of Maximum Multiplicity

“Electrons occupy degenerate orbitals singly with parallel spins before pairing.”

For 3 electrons in p orbitals:

Correct: ↑ | ↑ | ↑ (parallel spins)

Wrong: ↑↓ | ↑ | (premature pairing)

Reason: Parallel spins in different orbitals minimize electron-electron repulsion.

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Writing Electronic Configurations

Standard Notation

Example: Oxygen (Z = 8)

$$1s^2 \, 2s^2 \, 2p^4$$

Noble Gas Core Notation

Example: Iron (Z = 26)

$$[\text{Ar}] \, 3d^6 \, 4s^2$$

where [Ar] = $1s^2 2s^2 2p^6 3s^2 3p^6$

Orbital Box Notation

For nitrogen (Z = 7):

1s	2s	2p
↑↓	↑↓	↑ | ↑ | ↑

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Electronic Configurations of Elements

First 30 Elements

Z	Element	Configuration
1	H	1s¹
2	He	1s²
3	Li	[He] 2s¹
4	Be	[He] 2s²
5	B	[He] 2s² 2p¹
6	C	[He] 2s² 2p²
7	N	[He] 2s² 2p³
8	O	[He] 2s² 2p⁴
9	F	[He] 2s² 2p⁵
10	Ne	[He] 2s² 2p⁶
11	Na	[Ne] 3s¹
18	Ar	[Ne] 3s² 3p⁶
19	K	[Ar] 4s¹
20	Ca	[Ar] 4s²
21	Sc	[Ar] 3d¹ 4s²
26	Fe	[Ar] 3d⁶ 4s²
29	Cu	[Ar] 3d¹⁰ 4s¹ (Exception!)
30	Zn	[Ar] 3d¹⁰ 4s²

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Exceptional Configurations

Why Exceptions Occur

Half-filled and fully-filled subshells have extra stability due to:

1. Exchange energy: More for parallel electrons
2. Symmetry: Even distribution of electrons

Common Exceptions

Interactive Demo: Visualize Electronic Configuration

See how electrons fill orbitals following the Aufbau principle, Pauli exclusion, and Hund’s rule. Explore exceptional configurations of Cr and Cu.

Other Exceptions (for reference)

Element	Z	Expected	Actual	Reason
Mo	42	[Kr] 5s² 4d⁴	[Kr] 5s¹ 4d⁵	Half-filled
Ag	47	[Kr] 5s² 4d⁹	[Kr] 5s¹ 4d¹⁰	Fully-filled
Au	79	[Xe] 6s² 4f¹⁴ 5d⁹	[Xe] 6s¹ 4f¹⁴ 5d¹⁰	Fully-filled

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Configurations of Ions

Cations (Positive ions)

Electrons are removed from outermost shell first, not the last-filled orbital.

Example: Fe²⁺ (Z = 26)

• Fe: [Ar] 3d⁶ 4s²
• Fe²⁺: [Ar] 3d⁶ (4s electrons removed first!)

Example: Cu⁺

• Cu: [Ar] 3d¹⁰ 4s¹
• Cu⁺: [Ar] 3d¹⁰

Anions (Negative ions)

Electrons are added to the lowest available orbital.

Example: O²⁻

• O: [He] 2s² 2p⁴
• O²⁻: [He] 2s² 2p⁶ = [Ne]

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Magnetic Properties

From Electronic Configuration

• Unpaired electrons → Paramagnetic (attracted by magnet)
• All paired → Diamagnetic (weakly repelled by magnet)

Number of unpaired electrons:

Ion	Configuration	Unpaired e⁻	Magnetic
Fe²⁺	[Ar] 3d⁶	4	Paramagnetic
Fe³⁺	[Ar] 3d⁵	5	Paramagnetic
Cu⁺	[Ar] 3d¹⁰	0	Diamagnetic
Zn²⁺	[Ar] 3d¹⁰	0	Diamagnetic

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Problem Solving

Example 1: Write Configuration

Problem: Write the electronic configuration of V (Z = 23) and V³⁺.

Solution:

• V: [Ar] 3d³ 4s²
• V³⁺: [Ar] 3d² (remove 2 from 4s, then 1 from 3d)

Example 2: Identify Element

Problem: An element has configuration [Kr] 4d¹⁰ 5s² 5p⁴. Identify it.

Solution:

• Kr has Z = 36
• Add: 10 + 2 + 4 = 16 more electrons
• Total Z = 36 + 16 = 52
• Element: Tellurium (Te)

Example 3: Compare Stability

Problem: Which is more stable: [Ar] 3d⁵ 4s¹ or [Ar] 3d⁴ 4s²?

Solution: [Ar] 3d⁵ 4s¹ is more stable because:

• d⁵ is half-filled (extra stability)
• This is why Cr has this configuration

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Quick Reference

Maximum Electrons

Subshell	Formula	Max
s	2(0)+1 = 1 orbital	2
p	2(1)+1 = 3 orbitals	6
d	2(2)+1 = 5 orbitals	10
f	2(3)+1 = 7 orbitals	14

Stable Configurations

• Fully filled: s², p⁶, d¹⁰, f¹⁴
• Half filled: p³, d⁵, f⁷
• Noble gas: Very stable

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Practice Problems

1. Write the electronic configuration of Mn²⁺ and calculate unpaired electrons.

2. Element X has configuration [Ar] 3d⁵. What is its atomic number and is it a ground state configuration?

3. Which has more unpaired electrons: Fe²⁺ or Fe³⁺?

4. Why is Cu⁺ colorless while Cu²⁺ is blue?

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What’s Next?

You’ve completed the Atomic Structure chapter! Key connections:

• Chemical Bonding - How atoms share/transfer electrons
• Periodic Properties - How configuration affects properties

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Related Topics

Within Atomic Structure

• Quantum Numbers — The foundation for understanding orbitals
• Orbital Shapes — Visualizing where electrons are

Chemistry Connections

• Chemical Bonding — How configurations determine bonding
• Periodic Properties — Configuration explains trends
• d-Block Elements — Transition metal configurations
• Coordination Compounds — d-orbital involvement

Physics Connections

• Magnetism — Unpaired electrons cause paramagnetism
• Bohr Model — Historical context for energy levels

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