1.5 Atomic Structure

Emission Spectra

A spectrograph plate and matching intensity trace for 9 elements. Click any line to get λ, ν and the photon energy in joules, worked step by step from c = λν and E = hν.

Line SpectraEmission vs AbsorptionE = hν9 Elements
Topic 1.5

Atomic Structure and Electron Configuration

Represent the ground-state electron configuration of an atom of an element or its ions using the Aufbau principle.

An atom is a small, dense, positively charged nucleus of protons and neutrons surrounded by negatively charged electrons. Everything about chemical behavior comes from the electrostatic attraction between those two, which the CED insists you describe with Coulomb's law: the force between charges is proportional to the product of the charges and inversely proportional to the square of the separation.

Electrons occupy shells (energy levels, labeled n) and subshells (s, p, d, f). Electrons in the outermost shell are valence electrons; the rest are core electrons. Three rules generate the ground-state configuration:

  • Aufbau principle — fill lowest-energy orbitals first: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s.
  • Pauli exclusion principle — at most 2 electrons per orbital, with opposite spins.
  • Hund's rule — within a set of equal-energy orbitals, place one electron in each before pairing any up.

Configurations can be written in full (1s²2s²2p⁶3s¹), in noble-gas shorthand ([Ne]3s¹), or as an orbital-box diagram with arrows.

For ions: add electrons to the next available orbital for anions; for cations, remove electrons from the shell with the highest n first — not simply the last orbital filled. Iron is [Ar]4s²3d⁶, so Fe²⁺ is [Ar]3d⁶ (the 4s electrons go first), and Fe³⁺ is [Ar]3d⁵.

Finally, EK 1.5.A.4 connects configuration to energy: how hard it is to pull an electron off depends on its distance from the nucleus and the effective nuclear charge it feels after shielding by inner electrons. That is the seed of both PES (1.6) and the periodic trends (1.7).

Key points

  • Core electrons shield; valence electrons react. Almost every explanation in Unit 1 is one of those two sentences.
  • When ionizing a transition metal, remove the highest-n electrons (4s) before the d electrons.
  • Coulomb’s law is the required justification: closer and/or higher charge means a stronger attraction and a higher ionization energy.

Equations

  • on the exam sheetPrinted on the equation sheet under Atomic Structure. Use it qualitatively — the AP Exam never asks for a numerical force.
    • the two interacting charges
    • separation between them

Common mistakes

  • Take 4s electrons off first when forming transition-metal cations. Fe²⁺ is [Ar]3d⁶, never [Ar]4s²3d⁴.
  • Hund’s rule means single occupancy first with parallel spins. Do not pair electrons in one p orbital while another sits empty.
  • Noble-gas shorthand uses the previous noble gas, never the one in the element’s own row.
  • “Ground state” matters. A configuration like 1s²2s¹2p¹ is a legitimate excited state, but it is not the ground state and will be marked wrong if the ground state was asked for.

Worked example

Write the ground-state electron configuration of (a) Br, (b) S²⁻, and (c) Fe³⁺.

(a) Br (Z = 35): [Ar] 4s² 3d¹⁰ 4p⁵ — 18 + 2 + 10 + 5 = 35 ✓

(b) S²⁻: S is [Ne] 3s² 3p⁴ (16 e⁻). Adding 2 electrons gives [Ne] 3s² 3p⁶, which is isoelectronic with argon.

(c) Fe³⁺: Fe is [Ar] 4s² 3d⁶. Remove the two 4s electrons first, then one 3d electron → [Ar] 3d⁵.

Full notes for topic 1.5 →