Electron Configuration
Shells, Subshells, Orbitals & the Aufbau Principle
Why Electron Configuration Matters
In Tier I we learned that electrons occupy shells around the nucleus. Now we go deeper. The precise arrangement of electrons — their electron configuration — is the single most important factor in determining how an element behaves chemically.
It explains why sodium reacts violently with water while neon is inert. It explains the entire structure of the periodic table, why elements in the same group have similar properties, and how and why chemical bonds form. Every trend you will study in Tier II flows from electron configuration.
Chemistry is governed by electrons — specifically by the electrons in the outermost shell. Electron configuration tells you exactly how many electrons an atom has and where they are. From that, you can predict almost everything about an element’s chemistry.
Shells & Subshells
Electrons do not orbit the nucleus freely — they occupy specific energy levels called shells (or principal energy levels), numbered 1, 2, 3, 4 outward from the nucleus. The higher the shell number, the greater the energy and the farther from the nucleus.
Each shell is divided into subshells, labelled s, p, d, and f. These differ in shape and energy. The number of subshells in a shell equals the shell number: Shell 1 has only an s subshell; Shell 2 has s and p; Shell 3 has s, p, and d; and so on.
| Shell (n) | Subshells present | Max electrons per subshell | Max electrons in shell |
|---|---|---|---|
| 1 (K) | 1s | s: 2 | 2 |
| 2 (L) | 2s, 2p | s: 2 · p: 6 | 8 |
| 3 (M) | 3s, 3p, 3d | s: 2 · p: 6 · d: 10 | 18 |
| 4 (N) | 4s, 4p, 4d, 4f | s: 2 · p: 6 · d: 10 · f: 14 | 32 |
s holds 2 electrons (1 orbital × 2) · p holds 6 (3 orbitals × 2) · d holds 10 (5 orbitals × 2) · f holds 14 (7 orbitals × 2). The pattern: 2, 6, 10, 14 — each increases by 4.
Orbitals
Each subshell is made up of orbitals — regions of space where there is a high probability of finding an electron. Every orbital can hold a maximum of two electrons, and those two electrons must have opposite spins (one spin-up ↑, one spin-down ↓).
Spherical shape. 1 orbital per subshell. Holds up to 2 electrons. Found in every shell (1s, 2s, 3s…). Lowest energy within each shell.
Dumbbell-shaped. 3 orbitals per subshell (px, py, pz), oriented along the x, y, and z axes. Holds up to 6 electrons. Present from Shell 2 onwards.
More complex shapes (cloverleaf). 5 orbitals per subshell. Holds up to 10 electrons. Present from Shell 3 onwards. Responsible for transition metal chemistry.
Most complex shapes. 7 orbitals per subshell. Holds up to 14 electrons. Present from Shell 4 onwards. Associated with the lanthanides and actinides.
Notice the 2p subshell: nitrogen has 3 electrons in 2p, and they each occupy a separate orbital with the same spin direction before any pairing begins. This is Hund’s Rule — covered in the next section.
The Three Rules
Three principles govern how electrons fill orbitals. Together they allow you to determine the electron configuration of any element.
Electrons fill orbitals starting from the lowest available energy level and working upward. The ground state of an atom has electrons in the lowest possible energy arrangement. You fill 1s before 2s before 2p before 3s, and so on — following the Aufbau (filling) order.
No two electrons in the same atom can have identical quantum numbers. In practice: each orbital holds at most two electrons, and they must have opposite spins (one ↑ and one ↓). You cannot have two spin-up electrons in the same orbital.
When filling orbitals of equal energy (within the same subshell), electrons occupy separate orbitals with the same spin direction before any pairing begins. Spreading out minimises electron-electron repulsion and gives the atom its lowest energy state.
The Aufbau Filling Order
The order in which subshells are filled does not follow the simple shell sequence — because subshell energies from different shells overlap. The 4s subshell, for example, fills before the 3d because it is slightly lower in energy. This is the single most important sequence to memorise in this lesson.
4s fills before 3d when building up an atom — but when electrons are removed (ionisation), 4s electrons are lost first. This is because once 3d is occupied, it drops below 4s in energy. So Fe is [Ar] 3d⁶ 4s² but Fe²⁺ is [Ar] 3d⁶ — the two 4s electrons are removed first, not 3d electrons.
Writing Electron Configurations
Electron configuration is written as a sequence of subshell labels, each with a superscript showing the number of electrons it contains. Read the Aufbau order, filling each subshell to capacity before moving to the next.
The format is: shell number + subshell letter + superscript (electron count). Example: 1s² means 2 electrons in the 1s subshell.
For heavier elements, chemists use a shorthand called the noble gas configuration. Replace the core electrons (everything up to the previous noble gas) with the noble gas symbol in brackets, then write only the remaining outer electrons.
Valence Electrons
Valence electrons are the electrons in the outermost (highest-numbered) shell of an atom. They are the electrons that participate in chemical bonding — forming ionic bonds by transfer, or covalent bonds by sharing. The core electrons (inner shells) are shielded and generally do not participate in reactions.
For main-group elements (Groups 1–2 and 13–18), the number of valence electrons equals the group number:
| Group | Example | Configuration (outer) | Valence e⁻ | Typical ion |
|---|---|---|---|---|
| 1 | Na | 3s¹ | 1 | Na⁺ (loses 1) |
| 2 | Mg | 3s² | 2 | Mg²⁺ (loses 2) |
| 13 | Al | 3s² 3p¹ | 3 | Al³⁺ (loses 3) |
| 14 | Si | 3s² 3p² | 4 | Shares (4 bonds) |
| 15 | P | 3s² 3p³ | 5 | P³⁻ (gains 3) |
| 16 | S | 3s² 3p⁴ | 6 | S²⁻ (gains 2) |
| 17 | Cl | 3s² 3p⁵ | 7 | Cl⁻ (gains 1) |
| 18 | Ar | 3s² 3p⁶ | 8 | No ion (stable) |
Notable Exceptions
Two important exceptions to the standard Aufbau filling order arise among the transition metals. They occur because a completely full or exactly half-full d subshell confers extra stability — enough to “steal” an electron from the adjacent s subshell.
For Tier II, focus on learning the standard Aufbau order and being able to write configurations for elements Z=1 through Z=36 (up to krypton). Know the chromium and copper exceptions by name. The deeper quantum mechanical reasons for these anomalies are explored in Tier IV.
Configuration Builder
Select an element to see its full electron configuration, noble gas shorthand, and valence electron count.
Worked Examples
Method: Follow the Aufbau order, filling each subshell to capacity before moving to the next. Count electrons against Z.
1s² (2) → 2s² (4) → 2p⁶ (10) → 3s² (12) → 3p³ (15)
1s² 2s² 2p⁶ 3s² 3p³
Valence electrons: 5 (in 3s² 3p³) · Group 15
1s² 2s² 2p⁶ 3s² 3p⁶ (18, = Ar) → 4s² (20)
[Ar] 4s²
Valence electrons: 2 (in 4s²) · Group 2
[Ar] (18) → 4s² (20) → 3d⁶ (26)
[Ar] 3d⁶ 4s²
Valence electrons: 2 (in 4s²) — transition metals counted by 4s for main bonding
Method: Start with the neutral atom’s configuration. Remove electrons for cations (from outermost shell first); add electrons for anions.
1s² 2s² 2p⁶ — identical to neon’s configuration.
1s² 2s² 2p⁶ 3s² 3p⁶ — identical to argon’s configuration.
[Ar] 3d⁶
[Ar] 3d⁵ — a half-filled d subshell, which is notably stable.
Practice Questions
Q1. What is the maximum number of electrons that can occupy the 3d subshell?
Q2. Which set of subshells is filled in the correct Aufbau order?
Q3. Hund’s Rule states that electrons in degenerate (equal-energy) orbitals will:
Q4. An element has the configuration [Ne] 3s² 3p³. How many valence electrons does it have, and which element is it?
Q5. When Fe (Z=26) forms an Fe²⁺ ion, which electrons are removed first?
Key Takeaways
- Electron configuration describes the precise arrangement of electrons in an atom’s shells, subshells, and orbitals — and governs all chemical behaviour.
- Shells (n=1,2,3…) contain subshells (s, p, d, f). Subshell capacities: s=2, p=6, d=10, f=14 electrons.
- Each orbital holds at most 2 electrons with opposite spins (Pauli Exclusion Principle).
- Aufbau Principle: fill from lowest energy upward. Critical crossover: 4s fills before 3d.
- Hund’s Rule: within a subshell, electrons occupy separate orbitals (same spin) before pairing.
- Noble gas shorthand replaces the inner-electron core with [noble gas symbol] for efficiency.
- Valence electrons (outermost shell) determine reactivity and bonding. Group number = valence electrons for main-group elements.
- Exceptions: Cr ([Ar] 3d⁵ 4s¹) and Cu ([Ar] 3d¹⁰ 4s¹) gain stability from half-filled and fully-filled d subshells.
- For ion configurations: remove electrons from outermost shell first (4s before 3d for transition metals).