Electron Configuration and Atomic Structure Study Pack
Kibin's free study pack on Electron Configuration and Atomic Structure includes a 6-section study guide, 25 quiz questions, 30 flashcards, and 5 open-ended Explain review questions. Sign up free to track your progress toward mastery, plus upload your own notes and recordings to create personalized study packs organized by course.
Last updated May 27, 2026
Electron Configuration and Atomic Structure Study Guide
Break down electron configuration from the ground up — covering quantum numbers, the Pauli exclusion principle, Hund's rule, and the aufbau filling sequence. Learn to write nℓx notation, identify valence vs. core electrons, and understand anomalous cases like Cr and Cu.
Key Takeaways
- •Electrons occupy discrete energy levels called shells, which are subdivided into subshells (s, p, d, f) and further into individual orbitals, each defined by a unique set of four quantum numbers.
- •The Pauli exclusion principle states that no two electrons in an atom can share the same four quantum numbers, limiting each orbital to a maximum of two electrons with opposite spins.
- •The aufbau principle requires electrons to fill orbitals in order of increasing energy, following a specific diagonal filling sequence rather than simple shell-by-shell order.
- •Hund's rule specifies that when filling degenerate (equal-energy) orbitals within the same subshell, electrons occupy each orbital singly before any orbital receives a second electron, and all singly occupied orbitals have parallel spins.
- •Electron configuration notation uses the format nℓx — where n is the principal quantum number, ℓ is the subshell letter, and x is the electron count — to compactly describe how electrons are distributed across all occupied subshells.
- •Valence electrons, those in the outermost principal energy level, determine an element's chemical reactivity and bonding behavior, while core electrons remain shielded from full nuclear attraction.
- •Several elements, including chromium (Cr) and copper (Cu), display anomalous configurations because half-filled or fully filled d subshells confer extra stability, causing one electron to shift from the expected 4s² arrangement.
Quantum Numbers and the Address System for Electrons
Every electron in an atom is described by four quantum numbers that together specify its energy, the shape of its orbital, its spatial orientation, and its intrinsic spin — functioning like a unique four-part address that no two electrons can share.
Principal Quantum Number (n)
- •Indicates the main energy level, or shell, and takes positive integer values: 1, 2, 3, 4, and so on.
- •Larger values of n correspond to higher energy and greater average distance from the nucleus.
- •Each shell can hold a maximum of 2n² electrons total.
Angular Momentum Quantum Number (ℓ)
- •Defines the subshell and the shape of the orbital; its allowed values are 0 through (n − 1).
- •The value ℓ = 0 corresponds to an s subshell (spherical orbital), ℓ = 1 to p (dumbbell-shaped), ℓ = 2 to d, and ℓ = 3 to f.
Magnetic Quantum Number (mℓ)
- •Specifies the spatial orientation of an orbital within a subshell; its allowed integer values range from −ℓ to +ℓ.
- •An s subshell has only one orbital (mℓ = 0), a p subshell has three (mℓ = −1, 0, +1), a d subshell has five, and an f subshell has seven.
Spin Quantum Number (ms)
- •Describes the intrinsic angular momentum of an electron, restricted to only two values: +½ (spin-up) or −½ (spin-down).
- •Because ms has only two options, any single orbital can hold at most two electrons, and those two electrons must have opposite spins.
Orbital Energy Ordering and the Aufbau Filling Sequence
Orbitals do not increase in energy in a simple numerical order; the relative energies of subshells shift as electron–electron repulsion grows with atomic number, producing a diagonal filling pattern that must be memorized or reconstructed using the aufbau diagram.
General Energy Trend Across Subshells
- •Within the same principal shell, subshell energy increases in the order s < p < d < f.
- •Across different shells, subshell energies can cross: the 4s subshell is lower in energy than the 3d subshell in neutral atoms during filling, which is why 4s fills before 3d.
The Aufbau Principle
- •The aufbau principle (from the German word for 'building up') states that electrons enter the lowest available energy orbital before occupying higher-energy ones.
- •The diagonal filling order proceeds: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p, and so on.
- •A useful mnemonic is the n + ℓ rule: orbitals are filled in order of increasing (n + ℓ) value; when two orbitals have the same sum, the one with lower n fills first.
Energy Cross-Over and Electron Removal
- •Although 4s fills before 3d, once a transition metal is ionized, the 3d electrons are lower in energy than 4s, so 4s electrons are removed first during ionization — the filling order and the removal order are not identical.
Rules Governing Electron Placement: Pauli, Hund, and Orbital Diagrams
Three foundational rules work together to determine exactly how electrons are distributed among the available orbitals within a subshell, and orbital box diagrams make these rules visible by representing each orbital as a box and each electron as an arrow.
Pauli Exclusion Principle
- •No two electrons in the same atom can have an identical set of all four quantum numbers.
- •In practice, this means each orbital (defined by a unique n, ℓ, mℓ combination) holds at most two electrons, and those two must have opposite ms values.
Hund's Rule of Maximum Multiplicity
- •When electrons occupy a set of degenerate orbitals — orbitals of equal energy within the same subshell, such as the three 2p orbitals — each orbital receives one electron before any receives a second.
- •All singly occupied orbitals in a degenerate set carry parallel spins (all spin-up or all spin-down), because this arrangement minimizes electron–electron repulsion.
- •For example, carbon (Z = 6) places its two 2p electrons in separate 2p orbitals with parallel spins, not paired in the same orbital.
Orbital Box Diagrams
- •In an orbital box diagram, each box represents one orbital, and arrows (↑ for spin-up, ↓ for spin-down) represent individual electrons.
- •A filled orbital shows one ↑ and one ↓ arrow; a half-filled degenerate set shows all ↑ arrows in separate boxes before any ↓ arrows are added.
- •These diagrams make Hund's rule violations immediately visible and are used to predict magnetic properties: atoms with unpaired electrons are paramagnetic (attracted to magnetic fields), while those with all electrons paired are diamagnetic.
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Created by Kibin to help students review key concepts, prepare for exams, and study more effectively. This Study Pack was checked for accuracy and curriculum alignment using authoritative educational sources. See sources below.
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What is the maximum number of electrons that can occupy the third principal energy level (n = 3)?
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Quantum Numbers
Explain the four quantum numbers in your own words. What does each one describe about an electron, and how do they work together to give every electron a unique 'address' in an atom?
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