In simple terms
A friendly intro before the formal notes — no formulas yet.
The Atom's Electron Hotel
Electrons don't just fly around the nucleus randomly; they occupy specific 'addresses' defined by energy levels and orbital shapes. We can map out every electron's location using a few simple rules, which is key to understanding chemical behaviour.
Imagine an atom is a very exclusive apartment building for electrons. The floors are the principal energy shells (n=1, 2, 3...). On each floor, there are different types of apartments, which are the sub-shells (s, p, d). Each apartment has one or more rooms, which are the orbitals. By law, a room (orbital) can hold a maximum of two residents (electrons), and they must have opposite 'spins' to get along!
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Shells are numbered n=1, 2, 3... and contain sub-shells (s, p, d). Electrons fill these in order of increasing energy: 1s, then 2s, then 2p, and so on.
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An orbital is a region of space where an electron is most likely to be found. Each orbital can hold a maximum of two electrons, which must have opposite spins.
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The Aufbau principle states we build up the electron configuration by filling the lowest energy orbitals first. This gives the ground-state configuration for an atom or ion.
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The shape and orientation of orbitals dictate how atoms bond. An s orbital is spherical, while the three p orbitals are dumbbell-shaped along the x, y, and z axes.
Explore the concept
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2 more simulations for this topic — run them in the Simulations section below
Simulations
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2 simulations
- GeoGebraCore9701 1.3
Aufbau diagrams
Energy levels for every subshell: slide the atomic number and the electrons fill the orbitals one at a time.
Try this
- Tick Show Aufbau and Show Config.
- Slide the atomic number from 1 to 10 and watch 1s, 2s and then 2p fill.
- Stop at carbon, nitrogen and oxygen and look at how the 2p electrons are arranged.
- Go on to potassium and scandium: which fills first, 4s or 3d?
Look for Subshells fill in order of energy — 4s before 3d — and the configuration grows by one electron per element.
Adam Wolfe · GeoGebra · GeoGebra Terms of Service
- PhET9701 1.3
Models of the Hydrogen Atom
Shine light on a hydrogen atom and switch between the Bohr and Schrödinger models to compare orbits with orbital clouds.
Try this
- Choose the Schrödinger model and turn the light on.
- Watch the electron's state (n, l, m) change as it absorbs photons and compare an s cloud with a p cloud.
- Open the electron energy level diagram and see how the gaps change as n rises.
Look for An orbital is a region of probability, not an orbit: s clouds are spherical and p clouds are two lobes. The energy levels crowd together as n increases.
Simulation by PhET Interactive Simulations, University of Colorado Boulder · Licensed to MarkScheme (public licence CC BY-NC 4.0 since 2026-03-30)
Full topic notes
Formal explanation with the rigour you need for the exam.
Principal Quantum Shells and Sub-shells
Electrons in an atom are arranged in principal quantum shells, often just called energy levels or shells. These are numbered starting from the one closest to the nucleus. The higher the value of 'n', the further the shell is from the nucleus and the higher its energy. Each shell can hold a maximum of electrons.
Within each principal shell (from n=2 upwards), there are sub-shells, identified by the letters s, p, d, and f. These sub-shells have slightly different energies. For a given 'n', the energy of the sub-shells increases in the order s < p < d < f.
Shell 1 (n=1): Contains one sub-shell: 1s. Max electrons = 2.
Shell 2 (n=2): Contains two sub-shells: 2s, 2p. Max electrons = 8.
Shell 3 (n=3): Contains three sub-shells: 3s, 3p, 3d. Max electrons = 18.
Shell 4 (n=4): Contains four sub-shells: 4s, 4p, 4d, 4f. Max electrons = 32.
Atomic Orbitals: Shapes and Occupancy
An atomic orbital is a mathematical function that describes the wave-like behaviour of an electron. For our purposes, we can think of it as a three-dimensional region of space around the nucleus where there is a high probability (over 90%) of finding an electron. Each sub-shell is composed of one or more orbitals.
s sub-shell: Contains one spherical s orbital.
p sub-shell: Contains three dumbbell-shaped p orbitals, oriented along the x, y, and z axes (). These three orbitals are 'degenerate', meaning they have the same energy.
d sub-shell: Contains five d orbitals of more complex shapes. You are not required to know their shapes for AS Level.
Occupancy: Any single orbital can hold a maximum of two electrons. According to the Pauli Exclusion Principle, these two electrons must have opposite spins, often represented by arrows pointing up (↑) and down (↓).
Rules for Filling Orbitals
To determine the electronic configuration of an atom, we place its electrons into orbitals following three fundamental rules. The resulting arrangement is the 'ground state' configuration, which is the most stable, lowest-energy state for that atom.
Order of Filling: .
1. Aufbau Principle: Electrons fill orbitals in order of increasing energy. They occupy the lowest available energy level first.
2. Hund's Rule: For degenerate orbitals (like the three p-orbitals), electrons fill each orbital singly before any orbital gets a second electron.
3. Pauli Exclusion Principle: An orbital can hold a maximum of two electrons, and they must have opposite spins.
A very common mistake is to assume the third shell (3s, 3p, 3d) fills completely before the fourth shell (4s) begins. This is incorrect! The 4s sub-shell is at a lower energy level than the 3d sub-shell, so it fills first. Remember the sequence: ...3p, then 4s, then 3d.
Configurations of Ions and the Transition Metal Exceptions
Writing configurations for ions is straightforward. For anions (negative ions), you add electrons. For cations (positive ions), you remove electrons. A crucial rule for forming cations of d-block elements is that electrons are removed from the principal quantum shell with the highest 'n' value first. This means 4s electrons are removed before 3d electrons.
Worked examples
See the formulas applied — reveal one step at a time, like the exam.
Determine the full electronic configuration for a neutral silicon atom (Si, Z=14) and represent the valence electrons using the 'electrons in boxes' notation.
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Identify the number of electrons: A neutral silicon atom has a proton number (Z) of 14, so it has 14 electrons.
Deduce the electronic configuration of an iron(II) ion, Fe²⁺. The proton number of iron is 26.
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Write the configuration for the neutral Fe atom (26 electrons):
How it all connects
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Glossary
Key terms for this topic — skim now; the Check step will test them.
- Principal quantum shell
A main energy level in an atom, denoted by the integer 'n' (n=1, 2, 3...). Electrons in shells further from the nucleus have higher energy.
- Atomic orbital
A region of space around the nucleus that can hold up to two electrons with opposite spins. It represents a 3D volume where there is a high probability (typically >90%) of finding the electron.
Quick check
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Revision flashcards
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Key takeaways
Review these before you close the topic — retrieval beats re-reading.
Shell 1 (n=1): Contains one sub-shell: 1s. Max electrons = 2.
Shell 2 (n=2): Contains two sub-shells: 2s, 2p. Max electrons = 8.
Shell 3 (n=3): Contains three sub-shells: 3s, 3p, 3d. Max electrons = 18.
Shell 4 (n=4): Contains four sub-shells: 4s, 4p, 4d, 4f. Max electrons = 32.
Practice — then mark it
The whole point: a real Cambridge question, marked mark-by-mark.
At very high temperatures, phosphorus can form P2 molecules. P2 contains a triple bond, P≡P. Describe the formation of the P≡P bond in terms of orbital overlap.
Complete Table 1.1.
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