AP CHEMISTRY
Atomic Theory & Electron Structure Test Date: 8/21/2026, Friday
Sequence Assessment Study Guide
Covers Tasks 1–24: mass & moles, atomic structure, isotopes, mass spectrometry, empirical/molecular formulas, hydrates, Coulomb’s law, electron configuration, periodic trends, photoelectron spectroscopy, and ions.
1. Significant Figures & Moles:
Key Ideas
● Significant figures = all digits known for certain, plus one estimated digit. ● Nonzero digits, zeros between nonzero digits, and zeros after a decimal + nonzero digit all count.
● Leading zeros (0.04) and trailing zeros without a decimal (4400) do NOT count. ● Exact (counted) quantities have infinite significant figures and are ignored when rounding.
● A mole is a counting unit for atoms/molecules, equal to Avogadro’s number. ● Read the question to decide the significant number of the answer
Key Equations / Values
Avogadro’s number = 6.02 × 10²³ particles
2. Atomic Structure & Models
Key Ideas
● Atoms contain protons (+1 charge, nucleus), neutrons (0 charge, nucleus), and electrons (≈0 mass, surrounding cloud with -1 charge).
● Protons define the element (atomic number); this never changes in chemical reactions. ● In a neutral atom, # protons = # electrons.
● Bohr model: electrons move in fixed circular orbits/energy levels.
● Schrödinger (quantum mechanical) model: electrons exist in probability clouds (orbitals), not fixed paths, considered more accurate, especially for complex atoms.
3. Isotopes & Relative Abundance
Key Ideas
● Isotopes = atoms of the same element (same protons) with different numbers of neutrons. ● Neutron number does not change an atom’s identity, only the proton number does.
● Relative abundance = the experimentally measured percentage of a given isotope in a natural sample.
4. Weighted Average Atomic Mass
Key Ideas
● The atomic mass on the periodic table is a weighted average of all naturally occurring isotopes.
● Multiply each isotope’s mass by its relative abundance (as a decimal), then sum. Key Equations / Values
Atomic mass = Σ (isotope mass × relative abundance)
5. Mass Spectrometry
Key Ideas
● A mass spectrometer ionizes atoms and measures how long they take to reach a detector. ● Lighter atoms travel faster; heavier atoms travel slower.
● The spectrum plots mass-to-charge ratio (m/z) vs. relative abundance, each isotope appears as its own peak.
● Unexpected peaks can reveal impurities in a sample.
6. Conservation of Mass & Law of Definite Proportions Key Ideas
● Law of conservation of mass: matter is neither created nor destroyed in a chemical reaction, atom counts stay the same, only their arrangement changes.
● Law of definite proportions: elements in a compound always combine in the same fixed ratio, regardless of sample size.
7. Empirical & Molecular Formulas
Key Ideas
● Empirical formula = simplest whole-number ratio of atoms in a compound. ● Molecular formula = actual number of atoms in a molecule; it’s a whole-number multiple of the empirical formula.
● Process: convert mass → moles (using molar mass) → find the mole ratio → reduce to simplest whole numbers.
● Use molar mass of the whole molecule to find how many empirical units fit into the molecular formula.
Key Equations / Values
mol = mass (g) ÷ molar mass (g/mol)
multiplier = molecular mass ÷ empirical formula mass
8. Hydrates
Key Ideas
● A hydrate is an ionic compound with water molecules trapped in its crystal structure. ● Heating drives off the water; the mass lost tells you how much water was present. ● Compare the mass of water lost to the molar mass of the anhydrous (water-free) compound to find moles of water per mole of compound.
Key Equations / Values
% mass lost = (mass of water lost ÷ total mass) × 100
% error = |experimental − expected| ÷ expected × 100
9. Coulomb’s Law & Electron Energy
Key Ideas
● Coulomb’s law: force between charged particles is proportional to the product of their charges and inversely proportional to the square of the distance between them. ● Electrons closer to the nucleus feel more attractive force and are more tightly held (lower/more negative energy).
● Electrons farther away feel less attractive force and are more easily removed (higher, less negative energy).
● Ball in a ditch analogy: removing an electron is like lifting a ball out of a ditch, electrons near the bottom (nucleus) need more energy to remove.
Key Equations / Values
Fc ∝ (q1 × q2) / r²
10. Electron Configuration
Key Ideas
● Aufbau principle: electrons fill the lowest-energy shells and subshells first. ● Subshell order (lowest to highest energy within a shell): s, p, d, f.
● Subshell capacities: s = 2, p = 6, d = 10, f = 14 electrons.
● d and f subshells are so high-energy they fill after the s subshell of the next primary shell (e.g., 4s fills before 3d).
11. Excitation & Emission
Key Ideas
● Ground state = an atom's lowest-energy, most stable electron arrangement. ● Excited state = electrons have absorbed energy and jumped to a higher shell before a lower shell is full.
● When an excited electron drops back down, it releases energy as light, this is emission (e.g., glowing neon signs).
● Electron energy is quantized: electrons can only occupy specific allowed energy levels, not values in between.
12. Periodic Trends
Key Ideas
● Effective nuclear charge = the net positive pull felt by an electron after accounting for shielding by inner electrons.
● Atomic radius: increases down a column (more shells); decreases across a row (higher nuclear charge pulls electrons in).
● Electronegativity & ionization energy: increase across a row, decrease down a column (opposite of atomic radius).
Key Equations / Values
Z − S (Z = protons, S = shielding/inner electrons)
13. Photoelectron Spectroscopy (PES)
Key Ideas
● PES measures the ionization (binding) energy required to eject each electron from an atom.
● The x-axis (binding energy) is logarithmic and decreases left to right; the y-axis shows the number of electrons at each energy.
● Peak height = number of electrons at that energy level; large gaps between peaks usually signal a change in primary shell.
● The lowest-energy peak(s) represent the valence electrons, easiest to remove.
14. Ions & Ionic Compounds
Key Ideas
● An ion forms when an atom gains or loses electrons; cations are positive (lost electrons), anions are negative (gained electrons).
● Atoms in the same group tend to form ions with the same charge (e.g., Group 1 = +1, Group 2 = +2, Group 17 = –1).
● Ionic compounds are always electrically neutral, total positive charge must balance total negative charge.
● Transition metals can form multiple charges; roman numerals in the compound name specify which one, e.g., mercury(II) sulfide.