Electron Configuration and Atomic Structure Study Pack

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Last updated May 21, 2026

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Electron Configuration and Atomic Structure Study Guide

Unpack the rules governing how electrons fill atomic orbitals — including the Aufbau principle, Pauli exclusion principle, and Hund's rule — and see how these principles connect to periodic table structure. This pack covers s, p, d, and f subshells, notable exceptions like chromium and copper, and how core versus valence electrons influence reactivity and chemical bonding.

Key Takeaways

  • •Electrons occupy discrete energy levels called shells, which are subdivided into subshells (s, p, d, f) that differ in shape and energy, determining how electrons are distributed around a nucleus.
  • •The Aufbau principle states that electrons fill orbitals from lowest to highest energy, following a specific diagonal filling order that does not simply proceed shell by shell.
  • •The Pauli exclusion principle limits each orbital to a maximum of two electrons, which must have opposite spin quantum numbers (+½ and −½).
  • •Hund's rule requires that electrons occupy all orbitals within a subshell singly before any orbital receives a second electron, minimizing electron–electron repulsion.
  • •An atom's electron configuration directly determines its position in the periodic table: the period corresponds to the highest principal quantum number in use, and the block (s, p, d, or f) identifies the subshell being filled.
  • •Certain elements, notably chromium (Cr) and copper (Cu), deviate from predicted configurations because half-filled and fully filled d subshells confer extra stability.
  • •Core (inner) electrons shield valence electrons from the full nuclear charge, and the number of valence electrons governs an element's chemical reactivity and bonding behavior.

Quantum Numbers: The Address System for Electrons

Every electron in an atom is described by a unique set of four quantum numbers that specify its energy, the shape of the region it occupies, its spatial orientation, and its intrinsic spin.

Principal Quantum Number (n)

  • •Designates the main energy shell; n can be any positive integer (1, 2, 3, …).
  • •Higher n values correspond to greater average distance from the nucleus and higher potential energy.
  • •The maximum number of electrons a shell can hold is 2n², so shell 1 holds 2, shell 2 holds 8, shell 3 holds 18, and so on.

Angular Momentum Quantum Number (ℓ)

  • •Defines the subshell and the shape of the orbital; ℓ ranges from 0 to n − 1.
  • •ℓ = 0 corresponds to an s subshell (spherical), ℓ = 1 to p (dumbbell-shaped), ℓ = 2 to d (cloverleaf), and ℓ = 3 to f (complex multi-lobed shapes).

Magnetic Quantum Number (mₗ) and Spin Quantum Number (mₛ)

  • •mₗ specifies the orbital's orientation in space; it takes integer values from −ℓ to +ℓ, giving 1 orbital for s, 3 for p, 5 for d, and 7 for f subshells.
  • •mₛ describes the electron's intrinsic spin and can only be +½ (spin-up) or −½ (spin-down).
  • •No two electrons in the same atom can share all four identical quantum numbers — this is the Pauli exclusion principle and it limits each orbital to two electrons of opposite spin.

Rules Governing How Electrons Fill Orbitals

Three foundational principles work together to predict the ground-state electron configuration — the lowest-energy arrangement — of any atom.

Aufbau Principle and Diagonal Filling Order

  • •Electrons enter the lowest-energy orbital available before occupying higher-energy ones.
  • •Subshell energies do not increase in simple shell order; the 4s subshell, for example, lies lower in energy than 3d and fills first.
  • •The standard diagonal filling sequence is: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p.

Pauli Exclusion Principle

  • •Each orbital accommodates at most two electrons, and those two must have opposite spin quantum numbers.
  • •This constraint is why an s subshell holds 2 electrons, a p subshell holds 6, a d subshell holds 10, and an f subshell holds 14.

Hund's Rule

  • •When electrons are added to orbitals of equal energy (degenerate orbitals) within the same subshell, each orbital receives one electron before any orbital receives a second.
  • •All singly occupied orbitals within that subshell carry the same spin direction, which minimizes repulsion and produces a lower-energy state.
  • •For example, nitrogen (N, atomic number 7) has the configuration 1s² 2s² 2p³, with one electron in each of the three 2p orbitals rather than two electrons paired in one orbital.

About this Study Pack

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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