Build electron configurations from orbitals
Use shell, subshell and orbital models to write ground-state configurations and box notation for atoms and ions up to krypton.
Section 1 · Atomic structure
AS Level · Cambridge International Chemistry 9701, 2025–2027 syllabus. Electronic configurations are limited to ground-state atoms and ions of elements H to Kr. Written practice is self-checked; it does not replace cumulative exam practice or human subject-teacher review.
Before you begin
- Use atomic symbols and the Periodic Table, and distinguish an atom from an ion.
Work on paper. Try each question before revealing help, and explain why your method works.
Quick readiness check
How many electrons can one orbital hold?
What you will learn
- Relate shells, subshells and orbitals to electron occupancy
- Write full and shorthand ground-state configurations for atoms and ions
- Use orbital box notation and recognise s and p shapes
- Identify free radicals
Shells contain subshells, which contain orbitals
The principal quantum number n identifies a shell. A shell can contain subshells labelled s, p and d at the levels considered here. An orbital is a region associated with an electron's probability distribution; each orbital holds up to two electrons with opposite spins.
The s subshell has one orbital and holds two electrons; p has three orbitals and holds six; d has five orbitals and holds ten. The s orbital is spherical. Each p orbital has two lobes oriented along one of three perpendicular axes. These are schematic shapes, not paths followed by electrons.
Fill orbitals in energy order, then count
For the assessed ground-state configurations, the relevant increasing subshell order is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p. Fill the lower-energy subshells first and use box notation to show at most two opposite-spin electrons per orbital. In equal-energy p or d orbitals, place one electron in each before pairing; inter-electron repulsion also helps explain configurations such as chromium and copper.
Add superscripts to check that the total equals the electron number. For an ion, first adjust the total using the charge. When a transition-metal atom forms a positive ion, remove 4s electrons before 3d electrons. A free radical is a species with one or more unpaired electrons. Configuration assessment here is limited to ground-state atoms or ions of H to Kr.
Pause and explain
How many orbitals are in a p subshell?
Worked example
Write the full and shorthand ground-state electron configurations of Fe²⁺ (Z = 26).
Show the worked solution
- Neutral Fe has 26 electrons; Fe²⁺ has lost two, so it has 24.
- Fill to [Ar] 3d⁶ 4s⁰ after removing the two 4s electrons from neutral Fe.
- Expand [Ar] to write the full configuration and check the total is 24.
Answer Full: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶. Shorthand: [Ar] 3d⁶.
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Test this topic →From guided practice to a new situation
Write your answer and reasoning first. Use a hint only if you are stuck. The model solution is for self-checking; your written work is not automatically marked.
Write the shorthand ground-state configuration of S²⁻ (Z = 16) and state the number of unpaired electrons.
Give me a hint
Adjust the electron number for the 2− charge, then fill the subshells.
Compare my reasoning
- S²⁻ has 16 + 2 = 18 electrons.
- 18 electrons give the same arrangement as argon.
- Check whether any occupied subshell is partially filled.
1s² 2s² 2p⁶ 3s² 3p⁶, or [Ar]. All electrons are paired, so there are no unpaired electrons.
Look for these in your work
- Electron total adjusted for charge
- Configuration totals 18
- Pairing conclusion stated
Write the full ground-state configuration of Cr (Z = 24) and identify the syllabus-relevant feature of its configuration.
Give me a hint
Fill the lower subshells, then account for the observed 3d and 4s arrangement.
Compare my reasoning
- The argon core accounts for 18 electrons.
- The remaining six electrons are 3d⁵4s¹ in ground-state chromium.
- Check the total: 18 + 5 + 1 = 24.
1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹. Chromium is an example where configuration is explained using subshell energy and inter-electron repulsion rather than a simplistic fixed filling rule.
Look for these in your work
- Ground-state configuration correct
- Electron total checked
- Configuration limitation explained
Draw box notation for the three 2p orbitals of N (Z = 7), then state whether an isolated N atom is a free radical under the syllabus definition.
Give me a hint
Use Hund's rule across the three equal-energy p orbitals before pairing.
Compare my reasoning
- Nitrogen is 1s²2s²2p³.
- Put one electron in each 2p orbital.
- A free radical has one or more unpaired electrons.
2p: [↑] [↑] [↑]. Nitrogen has three unpaired electrons, so it is a free radical by the syllabus definition.
Look for these in your work
- Three orbitals shown
- One electron placed in each
- Unpaired-electron definition applied
Common mistakes
- Putting more than two electrons in an orbital
- Removing 3d electrons before 4s when forming a positive transition-metal ion
- Treating an orbital as a fixed circular path
- Calling a species a radical when every electron is paired
What can you explain now?
How many electrons can one orbital hold?
Compare with the explanation
Two with opposite spins
The maximum occupancy of an orbital is two electrons with opposite spins.
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Make it stick
How do you move from electron number to a ground-state configuration, and what do boxes reveal that superscripts can hide?
If you needed the solution, close it and solve the problem again from a blank page. A correct answer is stronger when you can explain the reason for each step.