AP Chemistry Unit 1 Cheat Sheet: Atomic Structure & Properties
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AP Chemistry Unit 1: Atomic Structure & Properties
TLDR
Understand atoms from the inside out — protons, neutrons, and electrons drive everything in chemistry.
Learn how electron configurations and periodic trends explain reactivity.
Master the formulas: E = hν = hc/λ, c = λν, and average atomic mass.
Know how to read and interpret PES (Photoelectron Spectroscopy) graphs.
✅ Download the full AP Chemistry Unit 1 Cheat Sheet (PDF) above
Why This Unit Matters
Atomic structure is the blueprint for all chemical behavior. From ionization energy to bonding strength, every reaction on the AP Chemistry exam traces back to how electrons are arranged around the nucleus.
A strong grasp of this unit helps you reason through trends instead of memorizing them. When you know why fluorine attracts electrons more than sodium, everything else falls into place.
1. The Atom — Small but Powerful
Atoms consist of a dense nucleus (protons and neutrons) surrounded by fast-moving electrons.
Protons (p⁺): define the element (atomic number).
Neutrons (n⁰): stabilize the nucleus; vary between isotopes.
Electrons (e⁻): control bonding and reactivity.
Average Atomic Mass:
multiply each isotope’s mass by its fractional abundance and sum them.
Example:
(0.75 × 35) + (0.25 × 37) = 35.5 amu for chlorine.
Exam Tip: expect calculation-based isotope questions early in multiple choice.
2. Light, Energy & the Quantum Model
Light acts as both a wave and a particle.
E = hν = hc/λ connects energy, frequency, and wavelength.
High frequency = short wavelength = high energy.
Photoelectric effect: electrons ejected only if photon energy ≥ threshold.
Photoelectron Spectroscopy (PES):
Each peak = electrons in a subshell (height = # of electrons, position = energy).
Sudden gaps show shell transitions.
Tutor Tip: PES graphs are like “electron fingerprints” — interpret patterns, not just numbers.
3. Electron Configuration
Electrons occupy orbitals by increasing energy:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p...
Rules to remember:
Aufbau: fill lowest energy first.
Pauli: two electrons per orbital, opposite spins.
Hund: fill singly before pairing.
Example:
Oxygen → 1s² 2s² 2p⁴
Shortcut: use noble gas core — Cl → [Ne]3s²3p⁵
4. Periodic Trends
Atomic structure explains the periodic table’s patterns.
Trend | Across Period | Down Group |
|---|---|---|
Atomic radius | ↓ | ↑ |
Ionization energy | ↑ | ↓ |
Electronegativity | ↑ | ↓ |
Electron affinity | more negative | less negative |
Mnemonic: “Rabbits Down, Energy Across” → Radius ↓, IE ↑, EN ↑
Concept Link: Ionization energy increases as electrons are removed; a large jump signals entering a new shell.
5. Quantum & Atomic Models
Bohr Model: electrons orbit nucleus in fixed energy levels (works for H only).
Quantum Mechanical Model: electrons exist in orbitals with probabilities, not fixed paths.
Heisenberg Uncertainty: can’t know both position and momentum.
Schrödinger Equation: predicts electron distributions.
Mnemonic: “Bohr orbits, Schrödinger clouds.”
Common Pitfalls
Mixing up mass number (whole number) vs atomic mass (decimal, average).
Forgetting successive ionization energy increases.
Misinterpreting PES — peak height = # of electrons, not energy.
Writing wrong configuration for ions (always remove from outer shell first).
Thinking energy and wavelength increase together (they’re inversely related).
Tutor Tip: AP Chemistry isn’t about memorizing equations — it’s about connecting them.
If you understand that smaller atomic radius → stronger nuclear attraction → higher ionization energy → higher electronegativity, you can reason through any question on trends.
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