What it covers: The structure of atoms — moles, mass spectrometry, electron configuration, photoelectron spectroscopy, and periodic trends.
Exam weight: About 7–9% of the AP Chemistry exam.
The big question: How does the arrangement of electrons in an atom explain the periodic trends and the chemical behavior of elements?
Big Ideas covered: Structure & Properties (SAP), Scale, Proportion & Quantity (SPQ), Chemical Effects (CE).
Key topics at a glance
Moles & Molar Mass
The mole (6.022 × 10²³ particles) bridges atomic-scale counting and lab-scale mass. Molar mass converts grams ↔ moles.
Mass Spectrometry
A mass spectrum shows isotopes and their relative abundances. Average atomic mass = weighted average of isotope masses.
Composition
Percent composition and empirical formulas come from mass data. Mixtures are analyzed using chromatography and spectroscopy.
Electron Configuration
Electrons fill orbitals following the Pauli exclusion principle (max 2 per orbital, opposite spins) and Hund's rule (fill singly first).
Photoelectron Spectroscopy
PES peaks show binding energy (x-axis, decreasing left to right) and relative electron count (peak height) — direct evidence for electron configuration.
Periodic Trends
Atomic radius ↓ across a period, ↑ down a group. Ionization energy and electronegativity do the opposite — driven by effective nuclear charge.
Effective Nuclear Charge
Zeff = actual nuclear charge − shielding. Increases across a period as protons are added without new shielding electrons.
Ionic Compounds
Atoms transfer valence electrons to reach stable configurations, forming oppositely charged ions held by Coulomb's law electrostatic attraction.
The key terms you must know
Mole / Avogadro's number — 6.022 × 10²³ particles; the bridge between atoms and grams.
Isotope — atoms of the same element with different neutron counts and masses.
Average atomic mass — the weighted average of isotope masses, weighted by abundance.
Electron configuration — the arrangement of electrons among orbitals (e.g., 1s² 2s² 2p⁶).
Pauli exclusion principle / Hund's rule — the rules that govern how orbitals fill.
Photoelectron spectroscopy (PES) — peak position = binding energy; peak height = electron count.
Effective nuclear charge (Zeff) — net positive pull felt by valence electrons; drives every periodic trend.
Atomic radius / ionization energy / electronegativity — the three core periodic trends, all explained by Zeff and shielding.
Valence electrons — outermost electrons; determine bonding behavior and ion charge.
Coulomb's law — force ∝ (charge₁ × charge₂) / distance² — the physics underneath ionic bonding and periodic trends.
Key themes to remember
Everything in this unit eventually reduces to Coulomb's law. Periodic trends, ionic bonding, and even orbital energies all come down to charge and distance.
The mole is a counting unit, not a substance. It works the same way for atoms, molecules, or ions — always 6.022 × 10²³ particles.
PES is evidence, not just a diagram. Treat every PES question as "what experimental data supports electron configuration?"
Trends move in opposite directions down a group vs. across a period — know which direction goes with which property.
Valence electrons are the story. Almost every later unit (bonding, IMFs, acids/bases) depends on understanding valence electron behavior first introduced here.
Common exam traps
Don't confuse atomic mass with mass number. Mass number is a whole number (protons + neutrons) for one isotope; atomic mass on the periodic table is a weighted average across all isotopes.
PES binding energy increases right to left on the graph, not left to right. Always check the axis direction before reading a PES spectrum.
Ionization energy and atomic radius are inversely related, but they're not measuring the same thing. Don't substitute one explanation for the other on a free-response question — name the correct property.
Shielding only comes from inner-shell electrons. Electrons in the same subshell shield each other very little — that's why Zeff still increases across a period.
"More electrons" doesn't always mean "bigger atom." Down a group, more electrons does mean a bigger atom (new shell). Across a period, more electrons in the same shell makes the atom smaller (more nuclear pull, same shielding).