AP Chemistry — Unit 1: Atomic Structure and Properties
Practice questions, answers, and key terms for Unit 1, aligned to the College Board CED.
Exam weighting: Unit 1 is 7-9% of the AP Chemistry exam.
What AP Chemistry Unit 1 covers
The College Board Course and Exam Description breaks Unit 1 into 8 topics:
- 1.1 Moles and Molar Mass
- 1.2 Mass Spectroscopy of Elements
- 1.3 Elemental Composition of Pure Substances
- 1.4 Composition of Mixtures
- 1.5 Atomic Structure and Electron Configuration
- 1.6 Photoelectron Spectroscopy
- 1.7 Periodic Trends
- 1.8 Valence Electrons and Ionic Compounds
AP Chemistry Unit 1 practice questions
What is Avogadro's number and what does it connect?
6.022 × 10²³ mol⁻¹ — it links moles to the actual count of particles. It bridges the particulate and macroscopic scales.
You weigh out 5.85 g of NaCl (molar mass 58.5 g/mol) for a titration. How many moles is that, and how did you get there?
0.100 mol — divide mass by molar mass: 5.85 g ÷ 58.5 g/mol. Units (g ÷ g/mol) leave mol, so divide, don't multiply.
Why can't chemists count atoms directly during lab work, and how does the mole solve this?
Atoms are far too small and numerous to count individually. The mole defines a fixed number of particles whose total mass equals the molar mass in grams, allowing counting by weighing.
What does a mass spectrum of a single element reveal?
Each isotope's mass number, plus its own relative natural abundance. You read masses off the x-axis and abundances off the y-axis.
How is average atomic mass calculated from a mass spectrum?
The average atomic mass = Σ (isotope mass × fractional abundance). It's a weighted average across all naturally occurring isotopes.
Chlorine's average atomic mass is 35.45 amu, yet no chlorine atom has that mass. Why?
35.45 amu is a weighted average of Cl-35 (~75% abundant) and Cl-37 (~25% abundant). Individual atoms are always whole-number mass isotopes; the average reflects the population.
State the Law of Definite Proportions.
A pure compound always has its elements in the same fixed mass ratio. That holds regardless of sample size or source.
What is an empirical formula, and how does it differ from a molecular formula?
The empirical formula shows the lowest whole-number ratio of atoms of each element in a compound. The molecular formula shows the actual number of atoms and may be a whole-number multiple of the empirical formula.
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Key terms in Unit 1
These 12 terms show up in the Unit 1 cards. Each one links to its definition in the AP Chemistry key-term reference.
Drill all of Unit 1 with spaced repetition
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