College-level general chemistry — atomic structure, bonding, reactions, kinetics, and equilibrium.
Study guide last reviewed September 25, 2026 · see what changed
Format reflects the current published structure — always confirm against AP Central before test day, since formats occasionally change.
| Section | Details | Time | Weight |
|---|---|---|---|
| Section I – Multiple Choice | 60 questions | 1 hr 30 min | 50% |
| Section II – Free Response | 7 questions (3 long, 4 short, calculator allowed) | 1 hr 45 min | 50% |
Straight answers about the AP Chemistry exam — format, credit, dates, and this site.
Section I – Multiple Choice: 60 questions (1 hr 30 min, 50% of the total score); Section II – Free Response: 7 questions (3 long, 4 short, calculator allowed) (1 hr 45 min, 50% of the total score). See the Exam Format tab above for the full breakdown, and always double-check the current year's format on AP Central before test day.
It's multiple-choice questions and written free-response or essay questions.
Yes — a graphing calculator is allowed (and expected) for this exam. Check AP Central for the current approved calculator list.
It depends entirely on the college — most award credit for a 3 or higher, but many selective schools require a 4 or a 5, and some don't grant credit at all for certain scores. Always check your target college's own AP credit policy page rather than assuming a 3 will count everywhere.
AP exams are administered during the College Board's annual testing window in the first two full weeks of May, with the exact date and time for each subject set by the College Board and announced the preceding fall. Check AP Central or your school's AP coordinator for this year's exact date.
Yes. The unit-by-unit study guide, worked examples, practice quiz, and full-length timed practice exam on this page are all free, with no login, signup, or paywall.
Unit-by-unit breakdown of what the course actually covers, in teaching order.
Electron configuration using the Aufbau principle, Pauli exclusion principle, and Hund's rule, and how orbital notation (s, p, d, f subshells) relates to the periodic table's structure. Periodic trends explained mechanistically: atomic radius decreases across a period (increasing nuclear charge) and increases down a group (added electron shells); ionization energy and electronegativity follow the inverse pattern, both tied to effective nuclear charge and electron shielding.
Ionic bonding (electron transfer, lattice energy) versus covalent bonding (electron sharing, bond polarity based on electronegativity difference). Lewis structures including resonance and formal charge, VSEPR theory predicting molecular geometry (linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral) from electron domain count, and orbital hybridization (sp, sp2, sp3) explaining how geometries arise from atomic orbital mixing.
The three types of intermolecular forces ranked by strength: hydrogen bonding (requires N-H, O-H, or F-H), dipole-dipole forces (between polar molecules), and London dispersion forces (present in all molecules, strength increasing with molecular size/surface area). Explains how these forces directly determine boiling/melting points, vapor pressure, viscosity, and solubility (like dissolves like), plus phase diagrams and the relationship between IMF strength and phase transitions.
Balancing chemical equations by conserving atoms, and stoichiometric calculations converting between grams, moles, and particles using molar mass and Avogadro's number. Reaction types (synthesis, decomposition, single replacement, double replacement, combustion) and solution chemistry: molarity calculations (M = mol/L), dilution (M1V1 = M2V2), and titration procedures for finding unknown concentrations using a known standard.
Rate laws (rate = k[A]^m[B]^n) determined experimentally from initial-rate data, and how reaction order affects the integrated rate law and half-life equations for zero-, first-, and second-order reactions. Collision theory explains why rate depends on temperature (more particles exceeding activation energy), concentration (more collision frequency), and catalysts (lowering activation energy via an alternate pathway) — reaction mechanisms and rate-determining steps tie these together.
Enthalpy (ΔH, heat absorbed/released at constant pressure) calculated via Hess's law or bond energies, entropy (ΔS, a measure of disorder) and the factors that increase it (more moles of gas, dissolving, temperature increase), and Gibbs free energy (ΔG = ΔH − TΔS) determining spontaneity — negative ΔG means spontaneous, and how temperature can flip a reaction's spontaneity depending on the signs of ΔH and ΔS.
The equilibrium constant expressions Kc and Kp, and how reaction quotient Q compared to K predicts which direction a reaction will shift. Le Chatelier's principle applied to concentration changes, pressure/volume changes (only affects gas-phase equilibria with unequal moles), and temperature changes (treating heat as a reactant or product depending on exothermic/endothermic). ICE tables as the standard method for calculating equilibrium concentrations.
Brønsted-Lowry acid-base theory (proton donors/acceptors) and conjugate acid-base pairs, pH/pOH calculations, and Ka/Kb values determining relative acid/base strength (with pKa as a quick-reference scale). Buffer solutions (a weak acid/conjugate base pair) resist pH change, calculated via the Henderson-Hasselbalch equation, and titration curves show distinct features for strong-strong, weak-strong, and polyprotic acid-base titrations including the equivalence point and buffer region.
Electrochemistry: galvanic (voltaic) cells generating current from spontaneous redox reactions, with standard reduction potentials (E°) used to calculate cell potential and predict spontaneity, versus electrolytic cells using external current to drive non-spontaneous reactions (electroplating, electrolysis). The relationship between ΔG°, E°, and the equilibrium constant K ties this unit back to thermodynamics and equilibrium.
Concrete, fully worked examples — this is what separates a page reviewers see as genuinely useful from one that just lists topics.
Electron configuration: write the configuration for Fe (Z=26), and identify how many unpaired electrons it has.
Following the Aufbau order: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶. Applying Hund's rule to the 3d⁶ subshell (5 orbitals, 6 electrons) gives 4 orbitals singly-occupied and one doubly-occupied, for 4 unpaired electrons — this is exactly why iron is strongly magnetic.
VSEPR: predict the molecular geometry of SF₄ (sulfur has one lone pair).
Sulfur has 5 electron domains (4 bonding + 1 lone pair), giving a trigonal bipyramidal electron geometry, but the lone pair occupies an equatorial position (to minimize repulsion), leaving a molecular shape called seesaw. Electron-domain geometry and molecular geometry differ whenever lone pairs are present — a distinction MC questions test directly.
Stoichiometry & titration: 25.0 mL of 0.100 M NaOH neutralizes 20.0 mL of HCl. Find the HCl concentration.
Moles NaOH = 0.0250 L × 0.100 M = 0.00250 mol. Since NaOH + HCl react 1:1, moles HCl = 0.00250 mol. Concentration = 0.00250 mol / 0.0200 L = 0.125 M.
Rate law from data: doubling [A] quadruples the rate while [B] is held constant; doubling [B] has no effect on rate. Find the rate law.
Since rate ∝ [A]² (doubling A ×4's the rate means order 2) and rate is independent of [B] (order 0): rate = k[A]². Overall reaction order = 2. Determining order from experimental data, rather than from the balanced equation's coefficients, is the entire point of this unit.
ICE table: for N₂O₄ ⇌ 2NO₂, Kc = 4.6×10⁻³ at a given temperature, starting with 0.050 M N₂O₄ and no NO₂.
Setting up ICE with x as the amount reacted: Kc = (2x)²/(0.050−x) ≈ 4x²/0.050 (approximating x small) = 4.6×10⁻³ → x² ≈ 5.75×10⁻⁵ → x ≈ 7.6×10⁻³ M. Always check afterward that x is small enough (<5% of initial) to justify the approximation — otherwise the full quadratic is required.
Buffer pH: find the pH of a buffer with 0.30 M acetic acid (Ka = 1.8×10⁻⁵) and 0.20 M sodium acetate.
Using Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]) = −log(1.8×10⁻⁵) + log(0.20/0.30) = 4.74 + (−0.176) = 4.56. The buffer's pH sits close to but not exactly at the pKa, shifted by the ratio of the two components — not by their absolute concentrations.
The errors that cost the most points across the whole exam, not just one unit.
Unofficial practice questions written for this study guide — not official AP content and not from College Board. Use them to self-check, then use the real past free-response questions below for authentic exam practice.
Unofficial practice, timed to feel like the real thing. Section I pulls from this page's full multiple-choice bank; Section II gives you free-response prompts to write out and self-score against a rubric. Bank size is still growing toward full exam length — timing is paced per-question, not pulled from this subject's official duration.