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.
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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
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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):
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Glossary
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Quick check
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Revision flashcards
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What is a 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.
Key takeaways
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Shell 1 (n=1): Contains one sub-shell: 1s. Max electrons = 2.
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Shell 2 (n=2): Contains two sub-shells: 2s, 2p. Max electrons = 8.
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Shell 3 (n=3): Contains three sub-shells: 3s, 3p, 3d. Max electrons = 18.
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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.
9701/22 · Q3(f)(i)
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.
9701/22 · Q1(b)(i)
Complete Table 1.1.
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