Electrons are not found in neat circular orbits like planets. Instead, they occupy regions of space called orbitals, structured in distinct energy levels (shells) and sub-shells.
🔑 Key Principle
The electronic structure of an atom dictates its chemical properties. Electrons fill orbitals in order of increasing energy, with a maximum of two electrons per orbital, spinning in opposite directions.
Energy Levels and Sub-Shells
Electrons exist in principal energy levels designated by the principal quantum number, n (where n = 1, 2, 3, 4...). Each principal level contains sub-shells, which contain orbitals:
- An orbital is a region of space around the nucleus where there is a high probability (95%+) of finding an electron. Each orbital can hold a maximum of 2 electrons.
- Sub-shells are groups of orbitals of the same shape and energy. There are four types of sub-shells you need to know: s, p, d, and f.
| Sub-shell Type | Number of Orbitals | Max Number of Electrons | Orbital Shape |
|---|---|---|---|
| s | 1 | 2 | Spherical |
| p | 3 | 6 | Dumbbell |
| d | 5 | 10 | Complex cloverleaf |
| f | 7 | 14 | Highly complex |
This means the total capacity of each principal energy level increases as n increases:
- n = 1: s sub-shell (1 orbital) → holds 2 electrons
- n = 2: s, p sub-shells (1 + 3 = 4 orbitals) → holds 8 electrons
- n = 3: s, p, d sub-shells (1 + 3 + 5 = 9 orbitals) → holds 18 electrons
- n = 4: s, p, d, f sub-shells (1 + 3 + 5 + 7 = 16 orbitals) → holds 32 electrons
The Three Rules of Electron Filling
To determine the electron configuration of any atom, we apply three fundamental rules:
Electrons fill orbitals of the lowest energy level first before filling higher energy levels.
Electrons occupy orbitals of the same energy singly (with parallel spins) before pairing up. This minimises repulsion between negative electrons.
An orbital can hold a maximum of two electrons, and they must have opposite spins (represented by up and down arrows: ↑↓).
Writing Electron Configurations
To write configurations, we list the sub-shells in filling order with the number of electrons in each written as a superscript (e.g., 1s² 2s² 2p⁶).
The Filling Order: The 4s and 3d Overlap
Crucially, the 4s sub-shell has a lower energy than the 3d sub-shell. Therefore, according to the Aufbau Principle, the 4s sub-shell fills before the 3d sub-shell.
The overall filling order up to Kr is:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p
Sodium (Na, 11 electrons):
We fill the shells in order: 1s holds 2, 2s holds 2, 2p holds 6, leaving 1 electron for 3s.
\[ \text{Na: } 1\text{s}^2 2\text{s}^2 2\text{p}^6 3\text{s}^1 \]
Titanium (Ti, 22 electrons):
We fill: 1s² 2s² 2p⁶ 3s² 3p⁶. Next, the 4s orbital fills before 3d: 4s². This accounts for 20 electrons. The remaining 2 electrons go into 3d: 3d².
\[ \text{Ti: } 1\text{s}^2 2\text{s}^2 2\text{p}^6 3\text{s}^2 3\text{p}^6 3\text{d}^2 4\text{s}^2 \]
Note: It is standard practice at A-Level to write the 3d sub-shell before 4s in the final written configuration, grouping the shells with the same principal quantum numbers (n=3) together.
Shorthand Notation
To write shorthand configurations, we substitute the inner filled shells with the symbol of the corresponding **Noble Gas** (e.g., [Ne] or [Ar]).
- [Ne] = 1s² 2s² 2p⁶
- [Ar] = 1s² 2s² 2p⁶ 3s² 3p⁶
For example, Sodium's shorthand is: [Ne] 3s¹. Titanium's shorthand is: [Ar] 3d² 4s².
Exceptions to the Rule: Chromium and Copper
Chromium (Cr, Z=24) and Copper (Cu, Z=29) do not follow the expected Aufbau filling order. They are **stabilised** by having a half-filled or fully-filled d sub-shell:
Chromium (Cr)
Expected: [Ar] 3d⁴ 4s²
Actual: [Ar] 3d⁵ 4s¹
A half-filled d sub-shell (d⁵) and a half-filled s sub-shell (s¹) are particularly stable due to symmetrical charge distribution.
Copper (Cu)
Expected: [Ar] 3d⁹ 4s²
Actual: [Ar] 3d¹⁰ 4s¹
A fully-filled d sub-shell (d¹⁰) is highly stable, which outweighs the energetic cost of having only 1 electron in the 4s orbital.
AQA exam papers test Chromium and Copper configurations very frequently. Make sure you memorise these exceptions! If you write 3d⁴ 4s² or 3d⁹ 4s² in the exam, you will lose the mark.
Electron Configurations of Ions
When writing configurations for ions, we add or remove electrons from the neutral atom configuration:
- Anions (negative ions): Add electrons to the next available orbital according to standard filling rules.
- Cations (positive ions): Remove electrons from the orbital of the **highest principal energy level** first (highest n).
⚠️ CRITICAL RULE FOR TRANSITION METALS
Although the 4s orbital fills before the 3d orbital, electrons are lost from the 4s orbital before the 3d orbital when transition metal cations form. This is because once the 3d sub-shell starts filling, the 4s energy level shifts slightly above it, making 4s electrons the easiest to remove.
Iron atom (Fe):
\[ \text{Fe: } [\text{Ar}] 3\text{d}^6 4\text{s}^2 \]
Iron(II) ion (Fe²⁺):
We must remove two electrons. Following the rule, we remove the 4s electrons first.
\[ \text{Fe}^{2+}: [\text{Ar}] 3\text{d}^6 \]
Iron(III) ion (Fe³⁺):
We remove three electrons. We take the two 4s electrons, plus one from the 3d sub-shell.
\[ \text{Fe}^{3+}: [\text{Ar}] 3\text{d}^5 \]
Fe³⁺ has a very stable d⁵ (half-filled) sub-shell configuration, explaining why Fe²⁺ easily oxidises to Fe³⁺.
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