Electron Configuration- Rules, Exceptions, and Examples

What Is Electron Configuration?

Electron configuration is the arrangement of electrons around an atom's nucleus. It tells you exactly where each electron sits in an atom, organized by energy levels and sublevels.

You need this to predict chemical behavior, understand bonding, and figure out why elements react the way they do. Skip it and you're basically guessing.

The Three Rules You Must Know

These aren't suggestions. Every electron configuration follows these rules. Memorize them.

1. Aufbau Principle

Electrons fill the lowest energy orbitals first. That's it. You work your way up from 1s to higher levels.

The order follows this sequence:

• 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p

Notice 4s fills before 3d. Don't ask why — just remember it. The math behind orbital energies is more complicated than this simplified rule suggests, but it works for most purposes.

2. Hund's Rule

When filling degenerate orbitals (like the three p orbitals), electrons go in singly first. They spread out before pairing up.

Example: Nitrogen (1s² 2s² 2p³) has one electron in each 2p orbital, not two electrons in one orbital and none in the others. Maximize parallel spins. That's what Hund's rule demands.

3. Pauli Exclusion Principle

Each orbital holds maximum two electrons. And those two electrons must have opposite spins. You can't squeeze a third electron into the same orbital — physics won't allow it.

This is why electrons are described with four quantum numbers, and why no two electrons in an atom can have identical sets.

Reading the Notation

The notation looks like this: 1s² 2s² 2p⁶

Here's how to decode it:

• The number before the letter = the principal energy level (n)
• The letter = the sublevel (s, p, d, f)
• The superscript = how many electrons are in that sublevel

So 2p⁶ means six electrons in the p sublevel of the second energy level. Simple.

Orbital Capacity Cheat Sheet

Every s holds 2, every p holds 6, every d holds 10, every f holds 14. That's the maximum — electrons fill these spots according to the rules above.

How to Write Electron Configurations: Getting Started

Let's do this step by step with Phosphorus (atomic number 15).

Step 1: Write out the filling sequence starting from 1s.

Step 2: Add electrons equal to the atomic number. Phosphorus has 15 electrons.

Step 3: Count electrons as you go.

• 1s² — 2 electrons
• 2s² — 4 electrons total
• 2p⁶ — 10 electrons total
• 3s² — 12 electrons total
• 3p³ — 15 electrons total. Stop here.

Phosphorus configuration: 1s² 2s² 2p⁶ 3s² 3p³

That's the long way. You'll also see shorthand notation using noble gas cores. For phosphorus: [Ne] 3s² 3p³. [Ne] represents 1s² 2s² 2p⁶ — the configuration of neon. Saves space.

Exceptions to the Rules

Here's where it gets messy. The simplified Aufbau principle fails for certain elements. Chemistry has exceptions.

Chromium (Cr) and Copper (Cu)

These are the most famous problem children.

Chromium should be [Ar] 4s² 3d⁴. But it's actually [Ar] 4s¹ 3d⁵.

Copper should be [Ar] 4s² 3d⁹. But it's actually [Ar] 4s¹ 3d¹⁰.

Why? Half-filled and fully-filled d sublevels have extra stability. One electron moves from 4s to 3d to achieve this. It's not a mistake — it's lower energy in practice.

Other elements show similar behavior. Molybdenum, silver, and gold follow comparable patterns. Don't panic when your calculated configuration doesn't match reality — check for d-block anomalies first.

The d and f Block quirks

Elements in the d-block (transition metals) don't always follow the simple filling order perfectly. Electron repulsion, shielding effects, and relativistic influences in heavier elements create deviations.

For f-block elements (lanthanides and actinides), the situation gets worse. 4f and 5f electrons fill in unpredictable ways. Most general chemistry courses just ask you to memorize these configurations or use a reference table.

Common Mistakes to Avoid

• Forgetting 4s fills before 3d. This trips up almost everyone at first.
• Pairing electrons too early. Hund's rule exists for a reason — follow it.
• Ignoring exceptions. Chromium and copper aren't typos. They're real configurations.
• Misreading superscripts. 2p⁶ is not the same as 2p2 with a superscript 6 floating around. Keep your notation clean.

Why This Matters

Electron configuration determines:

• Chemical reactivity and bonding behavior
• Oxidation states an element can achieve
• Whether an element is diamagnetic or paramagnetic
• Where elements sit on the periodic table

Get the configuration wrong and you can't explain why sodium loses one electron but carbon doesn't. You can't predict ion formation or molecular geometry. It's foundational stuff — learn it properly.

Quick Reference: Some Common Configurations

Notice gold's configuration above. Another exception — you'd expect [Xe] 6s² 4f¹⁴ 5d⁹. But gold is [Xe] 6s¹ 4f¹⁴ 5d¹⁰. Relativistic effects in heavy atoms shift orbital energies in ways the simple rules can't predict.

The Bottom Line

Electron configuration follows specific rules, but those rules have documented exceptions. Learn the principles first. Then learn where they break down. That's how chemistry actually works — models with known limitations.

Practice with neutral atoms until you're fast. Then tackle ions. Then memorize the d and f block anomalies. That's the progression. No shortcuts.