QCE Chemistry
Electron Configuration up to Z=36: A Complete Guide for QCE Chemistry
Electron configuration is one of those topics where students feel confident right up until they hit the transition metals, and then everything unravels. The Cr and Cu exceptions, the counterintuitive order of 4s and 3d, and the rules for writing ion configurations are all favourite territory for QCAA assessors in the external assessment. Getting the foundations right from the start means you will not be caught out by any of them.
What the syllabus asks
For Unit 1 of QCE Chemistry, you need to be able to write full and condensed electron configurations for elements up to Z=36 (krypton). That means understanding the three rules that govern how electrons fill orbitals, knowing the correct filling order including the 4s before 3d overlap, recognising the special cases of chromium and copper, and applying all of this to ions as well as neutral atoms.
The idea, explained
Orbitals and sublevels
In the Schrödinger model of the atom, electrons do not orbit the nucleus in neat circles. Instead, they occupy orbitals, which are three-dimensional regions of space where there is a high probability of finding an electron. Orbitals are grouped into sublevels labelled s, p, d, and f.
- s sublevel: 1 orbital, holds up to 2 electrons
- p sublevel: 3 orbitals, holds up to 6 electrons
- d sublevel: 5 orbitals, holds up to 10 electrons
Every individual orbital can hold a maximum of 2 electrons, and those two electrons must have opposite spins.
The three rules
The Aufbau principle states that electrons fill sublevels from the lowest energy upward. For atoms up to Z=36, the correct filling order is:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p
Notice that 4s fills before 3d. This is because, in a neutral atom, the 4s sublevel sits at a slightly lower energy than 3d. Many students assume the number in front of the letter always determines the order, but that is not the case once sublevels from different principal energy levels start to overlap.
Hund's rule applies within a sublevel. When electrons are filling a p, d, or f sublevel, they occupy separate orbitals one at a time, with parallel spins, before any pairing occurs. This minimises electron-electron repulsion. Nitrogen (Z=7), for example, has the configuration 1s² 2s² 2p³, where each of the three 2p orbitals contains exactly one electron, not two electrons crammed into one orbital and the third orbital empty.
The Pauli exclusion principle states that no two electrons in the same atom can have identical quantum numbers. In practice, this means each orbital holds at most two electrons, and those two must have opposite spins.
Full and condensed configurations
A full configuration lists every sublevel from 1s outward. A condensed configuration replaces the inner-shell electrons with the symbol of the preceding noble gas in square brackets, making the notation much shorter.
Some examples to study:
- Na (Z=11): full 1s² 2s² 2p⁶ 3s¹, condensed [Ne] 3s¹
- Cl (Z=17): full 1s² 2s² 2p⁶ 3s² 3p⁵, condensed [Ne] 3s² 3p⁵
- Ca (Z=20): full 1s² 2s² 2p⁶ 3s² 3p⁶ 4s², condensed [Ar] 4s²
- Fe (Z=26): full 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s², condensed [Ar] 3d⁶ 4s²
- Br (Z=35): full 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁵, condensed [Ar] 3d¹⁰ 4s² 4p⁵
The Cr and Cu exceptions
Chromium (Z=24) and copper (Z=29) do not follow the straightforward Aufbau prediction:
- Cr expected: [Ar] 3d⁴ 4s², but actual: [Ar] 3d⁵ 4s¹
- Cu expected: [Ar] 3d⁹ 4s², but actual: [Ar] 3d¹⁰ 4s¹
In each case, one electron moves from 4s into 3d. A half-filled d-sublevel (3d⁵) and a fully-filled d-sublevel (3d¹⁰) are more stable than configurations that are one electron away from those arrangements. Manganese (Z=25), by contrast, is not an exception because its standard Aufbau configuration already gives a half-filled 3d sublevel: [Ar] 3d⁵ 4s².
Configurations of ions
For cations, remove electrons starting from the highest principal quantum number, not from the last sublevel filled. For transition metals, this means the 4s electrons leave before the 3d electrons.
- Fe²⁺: Fe is [Ar] 3d⁶ 4s², so removing two electrons gives [Ar] 3d⁶
- Cu²⁺: Cu is [Ar] 3d¹⁰ 4s¹, so removing two electrons gives [Ar] 3d⁹
For anions, add electrons to the next available sublevel. Cl⁻ gains one electron to give [Ne] 3s² 3p⁶, which is the same configuration as argon.
What the exam asks
QCAA assessors commonly ask students to write full or condensed configurations for a named element, to identify the number of unpaired electrons in a given configuration, or to write the configuration of an ion. You may also be asked to explain why Cr or Cu has an unexpected configuration. Practise writing configurations from memory for a range of elements across the periodic table, including several transition metals, so the filling order becomes automatic.
Common mistakes
- Filling 3d before 4s. Always fill 4s first in a neutral atom. The energy overlap means 4s is lower in energy than 3d for elements in this range.
- Removing 3d electrons first when forming transition metal cations. When writing ion configurations, remove electrons from the highest principal quantum number first, so 4s electrons go before 3d electrons.
- Ignoring Hund's rule and pairing electrons too early. Within any sublevel, spread electrons across all available orbitals before pairing. Pairing too soon gives the wrong number of unpaired electrons.
- Forgetting the Cr and Cu exceptions. These two elements appear regularly in external assessment questions. Memorise both actual configurations and be ready to explain the stability of half-filled and fully-filled d-sublevels.
- Writing the condensed configuration in the wrong order. The convention is to list the noble gas symbol first, then the remaining sublevels in order of principal quantum number (for example, [Ar] 3d⁶ 4s², not [Ar] 4s² 3d⁶). Check the order before finalising your answer.
If you want to work through practice questions and get immediate feedback on your configurations, Avocado's AI-native tutoring service has a dedicated lesson ready for you: Electron Configuration up to Z=36. Avocado's lessons are built specifically for Queensland students and aligned to the QCAA syllabus, so every question you practise is exactly the kind of question that matters for your external assessment.