Understand Electron Configurations — Understand Structure...

SECTION 1

Historical Context & Motivation

For centuries, scientists asked a deceptively simple question: how are electrons arranged inside an atom? The answer turned out to be one of the most powerful ideas in all of chemistry. Understanding where electrons sit determines why sodium reacts violently with water, why neon refuses to react with anything, and why copper conducts electricity while sulfur does not. The story of electron configurations is really the story of scientists gradually decoding the internal architecture of matter itself.

1897

Discovery of the Electron

J.J. Thomson identifies the electron using cathode-ray tubes, proving that atoms contain negatively charged subatomic particles. His 'plum pudding' model embeds electrons in a sea of positive charge.

1913

Bohr's Quantized Orbits

Niels Bohr proposes that electrons orbit the nucleus only at fixed energy levels (n = 1, 2, 3…). This explains the line spectrum of hydrogen and introduces the concept of quantized energy.

1926

Schrödinger's Wave Equation

Erwin Schrödinger replaces fixed orbits with probability-based orbitals — 3-D regions of space where an electron is most likely to be found. The quantum mechanical model of the atom is born.

1925–1927

Pauli, Hund & Aufbau

Wolfgang Pauli establishes the exclusion principle, Friedrich Hund formulates rules for filling degenerate orbitals, and the Aufbau principle provides a step-by-step method for building up electron configurations.

1937–Present

Modern Periodic Table & Electron Configuration

The arrangement of the periodic table by atomic number is validated by electron configurations. Today, these configurations underpin our understanding of bonding, spectroscopy, and materials science.

The central question this lesson answers is: In what order do electrons fill the available orbitals of an atom, and how do we write that arrangement systematically? Mastering this will unlock your ability to predict reactivity, bonding behavior, and position on the periodic table for any element.

SECTION 2

Core Principles & Definitions

Before writing any electron configuration, you need to understand four foundational ideas. Each one is like a rule in a game — learn them, and the configurations fall into place naturally.

Energy Levels & Sublevels

Electrons exist in principal energy levels (n = 1, 2, 3…). Each level contains sublevels labeled s, p, d, f. The number of sublevels equals the principal quantum number n, so level 1 has only s, level 2 has s and p, level 3 has s, p, and d, and so on.

Aufbau Principle

The Aufbau principle (German for 'building up') states that electrons fill orbitals starting from the lowest available energy. This does not always follow the simple order 1→2→3 because some sublevels overlap in energy (e.g., 4s fills before 3d).

Pauli Exclusion Principle

The Pauli exclusion principle says no two electrons in an atom can have the same set of four quantum numbers. In practice, this means each orbital holds a maximum of two electrons with opposite spins (↑↓).

Hund's Rule

Hund's rule states that when filling orbitals of equal energy (degenerate orbitals), electrons occupy them singly with parallel spins before any orbital receives a second electron. Think of it as each seat on a bus getting one passenger before anyone doubles up.

Maximum Electrons per Sublevel

The s sublevel has 1 orbital → max 2 e⁻. The p sublevel has 3 orbitals → max 6 e⁻. The d sublevel has 5 orbitals → max 10 e⁻. The f sublevel has 7 orbitals → max 14 e⁻.

✦ KEY TAKEAWAY

Think of an atom like a multi-story hotel. The principal energy level is the floor number, and the sublevel (s, p, d, f) is the type of room on that floor. Guests (electrons) check into the cheapest available room first (Aufbau), each room holds at most two guests (Pauli), and guests prefer to spread out into separate rooms before sharing (Hund).

SECTION 3

Orbital Filling Order — Visual Explanation

The diagram below shows the Aufbau filling order using the diagonal rule. Each row represents a principal energy level, and the diagonal arrows trace the order in which sublevels are filled from lowest to highest energy. Follow the arrows starting from 1s, and you will generate the correct filling sequence for any element on the periodic table.

The diagonal arrows illustrate that sublevels from different principal energy levels can overlap in energy. Notice how 4s fills before 3d and 5s fills before 4d. The summary box on the right lists the complete filling sequence you should memorize.

A useful memory trick is the n + l rule: sublevels fill in order of increasing (n + l) value, where n is the principal quantum number and l is the sublevel number (s = 0, p = 1, d = 2, f = 3). When two sublevels have the same (n + l) value, the one with the smaller n fills first. For example, 3d has n + l = 3 + 2 = 5 and 4p has n + l = 4 + 1 = 5, but 3d fills first because its n is smaller. However, 4s (n + l = 4 + 0 = 4) fills before 3d (n + l = 5), which is why we see 4s appear before 3d in the sequence.

SECTION 4

Writing Electron Configurations

There are three main ways to represent the electron configuration of an atom. Each format has its own purpose, and the IB expects you to be comfortable with all of them.

Full (Spdf) Notation

In full electron configuration notation, you list every sublevel in filling order with the number of electrons as a superscript. For example, sodium (Na, Z = 11) is written as 1s² 2s² 2p⁶ 3s¹. The superscript numbers must add up to the total number of electrons (2 + 2 + 6 + 1 = 11 for neutral Na).

Condensed (Noble-Gas Core) Notation

For larger atoms, writing every sublevel becomes tedious. Condensed notation replaces the inner electrons with the symbol of the preceding noble gas in square brackets. Sodium becomes [Ne] 3s¹ because neon (Ne) accounts for the first 10 electrons (1s² 2s² 2p⁶). This notation highlights the valence electrons — the outermost electrons that participate in chemical reactions.

Orbital Box (Arrow) Diagrams

An orbital box diagram shows each orbital as a box and each electron as an arrow (↑ or ↓). This representation makes Hund's rule and Pauli's exclusion principle visually obvious, since you can see electrons spreading out before pairing and each pair pointing in opposite directions.

📝 IB Tip

The IB syllabus expects you to write full and condensed electron configurations for elements up to Z = 36 (krypton). You should also be able to deduce the configuration of common ions (e.g., Na⁺, Cl⁻, Fe²⁺) by adding or removing electrons from the neutral atom's configuration.

SECTION 5

Electron Configuration & the Periodic Table

The periodic table is not just a list of elements — it is a map of electron configurations. Each block of the periodic table corresponds to the sublevel being filled. The table below summarizes this powerful relationship.

The four blocks of the periodic table and the sublevel each fills.

Block	Sublevel Being Filled	Groups (Columns)	Width (# Elements)	Example Elements
s-block	s	1–2	2	H, Li, Na, K, Ca
p-block	p	13–18	6	C, N, O, F, Ne
d-block	d	3–12	10	Sc, Ti, Fe, Cu, Zn
f-block	f	Lanthanides & Actinides	14	Ce, U

The s-block is 2 columns wide (max 2 electrons in an s sublevel), the p-block is 6 wide, the d-block is 10 wide, and the f-block is 14 wide. This is not a coincidence — it directly reflects electron capacity.

This connection is incredibly useful. If you know an element's position on the periodic table, you can immediately determine which sublevel is being filled and deduce its electron configuration without memorizing the entire diagonal chart. For example, chlorine (Cl) sits in Period 3, Group 17 — that is the 5th column of the p-block, so its outer electrons end in 3p⁵. Iron (Fe) sits in Period 4, the 6th column of the d-block, so its outer electrons involve 3d⁶.

SECTION 6

Worked Example — Iron (Fe, Z = 26)

Let's write the full electron configuration, the condensed notation, and the orbital diagram for a neutral iron atom. Iron is a transition metal in Period 4 of the d-block, which makes it a great example for practising the Aufbau filling order.

Electron Configuration of Iron (Fe)

Step 1 — Determine the Number of ElectronsIron has an atomic number Z = 26. Since the atom is neutral (no charge), it has 26 electrons to distribute among its orbitals.

26 electrons

Step 2 — Fill Sublevels Using the Aufbau OrderFollow the filling sequence: 1s² (2), 2s² (4), 2p⁶ (10), 3s² (12), 3p⁶ (18), 4s² (20), 3d⁶ (26). We reach 26 electrons by placing 6 electrons into the 3d sublevel after filling 4s.

1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶

Step 3 — Write the Condensed NotationThe noble gas before iron is argon (Ar, Z = 18). Argon accounts for 1s² 2s² 2p⁶ 3s² 3p⁶. Replace that portion with [Ar].

[Ar] 4s² 3d⁶

Step 4 — Apply Hund's Rule to the 3d SublevelThe 3d sublevel has 5 orbitals. With 6 electrons, by Hund's rule, fill each of the 5 orbitals with one electron (↑) first, then place the 6th electron as a second electron (paired ↑↓) in the first 3d orbital. The result is four unpaired electrons and one paired orbital, making iron paramagnetic.

4 unpaired electrons in the 3d sublevel

Step 5 — Verify the TotalAdd all superscripts: 2 + 2 + 6 + 2 + 6 + 2 + 6 = 26. ✓ This matches the atomic number, confirming the configuration is correct.

Fe: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ or [Ar] 4s² 3d⁶

⚠️ Watch Out — Ions of Transition Metals

When transition metals form cations, electrons are removed from the 4s sublevel first, not 3d. So Fe²⁺ is [Ar] 3d⁶ (not [Ar] 4s² 3d⁴). This is because once the atom has been formed, the 3d sublevel drops below 4s in energy, so 4s electrons are lost first.

SECTION 7

Exceptions & Important Patterns

While the Aufbau principle works beautifully for most elements, there are notable exceptions in the d-block. These arise because half-filled and fully filled d sublevels provide extra stability due to symmetrical electron distribution and favorable exchange energy.

The two most important electron configuration exceptions in the first row of the d-block.

Element	Expected Configuration	Actual Configuration	Reason
Chromium (Cr, Z = 24)	[Ar] 4s² 3d⁴	[Ar] 4s¹ 3d⁵	A half-filled 3d sublevel (3d⁵) is more stable; one 4s electron moves to 3d.
Copper (Cu, Z = 29)	[Ar] 4s² 3d⁹	[Ar] 4s¹ 3d¹⁰	A fully filled 3d sublevel (3d¹⁰) is more stable; one 4s electron moves to 3d.

✦ KEY TAKEAWAY

Think of a half-filled or fully filled d sublevel like a perfectly balanced seesaw — it's in its lowest-energy, most stable state. Nature will sacrifice one 4s electron to achieve this symmetry, much like you might rearrange a deck of cards so all suits are evenly distributed. The IB expects you to know the chromium and copper exceptions specifically.

Beyond these exceptions, several general patterns emerge from electron configurations. Elements in the same group share the same valence electron configuration (just with different principal quantum numbers), which is why they exhibit similar chemical properties. For instance, all Group 1 metals end in ns¹, and all Group 17 halogens end in ns²np⁵. Recognizing these patterns means you rarely need to build configurations from scratch during an exam.

SECTION 8

Electron Configurations of Ions & Connection to Higher-Level Concepts

Neutral atoms often gain or lose electrons to form ions. Writing electron configurations for ions is straightforward once you understand two rules: for anions (negative ions), add the appropriate number of electrons to the neutral configuration; for cations (positive ions), remove electrons — but remember that for transition metals, you remove from the highest principal energy level (typically 4s) before the 3d sublevel.

How electron configurations connect to broader IB Chemistry topics.

Concept	Structure 1.3 (This Lesson)	Where It Leads
Valence electrons	Identified from the outermost energy level in the electron configuration.	Predict bonding behavior, Lewis structures (Structure 2.1), and VSEPR shapes.
Ionization energy trends	Linked to how tightly valence electrons are held (electron shielding, sublevel energy).	Explains periodic trends in Structure 3.1 and the successive ionization energy graph.
Transition metal ions	Involves removing 4s electrons before 3d, creating variable oxidation states.	Underlies color, catalysis, and complex ion formation in Reactivity 3 and HL topics.
Magnetic properties	Unpaired electrons in the orbital diagram determine paramagnetism vs. diamagnetism.	Connected to advanced materials science and coordination chemistry (HL).

At Higher Level, you will encounter the concept of orbital hybridization (sp, sp², sp³), where atomic orbitals merge to form new hybrid orbitals that better explain molecular geometry. You will also explore crystal field theory, which explains why transition metal complexes display vivid colors based on d-orbital splitting. Both of these advanced ideas are built directly on the electron configuration foundations you are learning now.

SECTION 9

Practice Problems

PROBLEM 1 — CONCEPTUAL

Explain why the 4s sublevel fills before the 3d sublevel, even though 4 is a higher principal quantum number than 3. Reference the Aufbau principle in your answer.

PROBLEM 2 — BASIC CALCULATION

Write the full electron configuration and the condensed (noble-gas core) notation for a neutral phosphorus atom (P, Z = 15).

PROBLEM 3 — INTERMEDIATE

Write the electron configuration for the Fe²⁺ ion. Explain why electrons are removed from 4s before 3d when forming the cation.

PROBLEM 4 — APPLIED

An element has the condensed electron configuration [Ar] 4s² 3d¹⁰ 4p⁴. (a) Identify the element. (b) State which block of the periodic table it belongs to. (c) Predict how many electrons it would need to gain to achieve a noble-gas configuration, and write the configuration of the resulting ion.

PROBLEM 5 — CRITICAL THINKING

Copper's actual electron configuration is [Ar] 4s¹ 3d¹⁰ rather than the expected [Ar] 4s² 3d⁹. (a) Explain the energetic reason for this anomaly. (b) Predict the electron configuration of Cu⁺ and Cu²⁺. (c) Based on their configurations, predict which ion is more stable and justify your reasoning in terms of d-sublevel filling.

SUMMARY

Lesson Summary

Electron configurations describe how electrons are distributed among an atom's energy levels and sublevels. The Aufbau principle dictates that electrons fill from lowest to highest energy, the Pauli exclusion principle limits each orbital to two electrons with opposite spins, and Hund's rule requires electrons to spread out across degenerate orbitals before pairing. The filling order is 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p and so on, with sublevels filling in order of increasing (n + l) value. Configurations can be written in full notation or condensed (noble-gas core) notation.

The periodic table's s-, p-, d-, and f-blocks directly correspond to the sublevel being filled, making the table itself a roadmap for writing configurations. Key exceptions exist for chromium ([Ar] 4s¹ 3d⁵) and copper ([Ar] 4s¹ 3d¹⁰) due to the extra stability of half-filled and fully filled d sublevels. When forming cations of transition metals, electrons are removed from 4s before 3d. Mastering these principles enables you to predict an element's reactivity, bonding patterns, and position on the periodic table.