FREE AP Chemistry Study Guide 2026: All 9 Units, Reactions, Equilibrium & Acids

This free AP Chemistry study guide teaches to the College Board’s current AP Chemistry course — every one of the nine units, organized the way the exam is built.^[1] It covers the concepts and the math you need for both the multiple-choice and free-response sections, from moles and atomic structure all the way to free energy and electrochemistry.

The exam has two equally weighted halves — 60 multiple-choice questions and 7 free-response questions — so this guide teaches the science behind both, not just test tricks.^[2] It is interactive: every unit ends with a checkpoint quiz, key terms are hover-able, math and reactions are worked step by step, and concept questions let you learn by doing.

Read this guide unit by unit, test yourself at each checkpoint, then round out your free AP Chemistry prep with our practice questions and flashcards.

AP Chemistry is one of the 17 AP exams — explore our AP study guides to compare and prep across the whole family.

AP Chemistry Exam Snapshot

Section I and Section II are worth 50% each.^[2] Across the nine units, the College Board reports approximate weighting ranges, and they are not equal — two units carry far more of the multiple-choice section than the rest:

College Board reports each unit’s share of the multiple-choice section as a range, so the exact mix shifts a little each year.^[1] This guide teaches all nine units in order — but spend extra time on Units 3 and 8, which together are roughly a third of the multiple-choice questions.

1 · Atomic Structure & Properties

7–9% of the exam. The mole concept and stoichiometric foundations, the evidence for atomic structure (mass spectrometry and photoelectron spectroscopy), electron configurations, and the periodic trends that fall out of them.^[1]

Moles & the Mole Concept

The mole is the bridge between mass and number of particles. To find moles, divide mass by molar mass; to find particles, multiply moles by Avogadro’s number $6.022 \times 10^{23}$.

Electron Configuration & PES

Fill subshells in order of increasing energy: magnesium (12 electrons) is $\text{1s}^2\,\text{2s}^2\,\text{2p}^6\,\text{3s}^2$. A PES spectrum gives a peak for each occupied subshell — its position shows binding energy, its height the number of electrons. A filled s (2 electrons) and filled p (6) give a 1-to-3 height ratio.

Periodic Trends

Effective nuclear charge — the net pull on a valence electron after shielding — drives every trend. Across a period it rises, so atoms shrink and ionization energy and electronegativity increase.

2 · Compound Structure & Properties

7–9% of the exam. Bonding (ionic, covalent, metallic), Lewis structures, VSEPR geometry, hybridization, and how structure controls properties.^[1]

Types of Bonding

A metal + nonmetal usually forms an ionic bond (conducts when molten or dissolved); two nonmetals share electrons in a covalent bond. Higher bond order means a shorter, stronger bond: single < double < triple.

Lewis Structures & VSEPR

Draw the Lewis structure, count electron domains on the central atom, and apply VSEPR: domains repel and spread out, while lone pairs compress the angles. Watch for exceptions — boron trifluoride $\ce{BF3}$ has an incomplete octet, and $\ce{PCl5}$ has an expanded octet.

Hybridization & Bonds

A single bond is one sigma bond; a double bond is 1 sigma + 1 pi bond; a triple bond is 1 sigma + 2 pi. Hybridization follows the electron geometry: sp (linear), sp² (trigonal planar), sp³ (tetrahedral).

3 · Properties of Substances & Mixtures

18–22% of the exam — the single heaviest unit. Intermolecular forces, states of matter, solutions, and the physical properties that follow from them. Master this unit first.^[1]

Intermolecular Forces

All molecules have London dispersion forces; polar molecules add dipole-dipole; molecules with H bonded to F, O, or N add hydrogen bonding; an ion plus a polar molecule gives ion-dipole. Stronger forces mean higher boiling points and lower vapor pressure.

States of Matter & Solutions

Intermolecular forces explain why substances are solid, liquid, or gas at a given temperature and why some dissolve in others — “like dissolves like.” A liquid boils when its vapor pressure equals the external atmospheric pressure.

Physical Properties & Separation

Differences in intermolecular forces let us separate mixtures: distillation exploits boiling-point differences, chromatography exploits differing attractions to a stationary phase. A volatile liquid has weak forces, a high vapor pressure, and evaporates readily.

4 · Chemical Reactions

7–9% of the exam. Reaction types, net ionic equations, oxidation-reduction, and the stoichiometry of reactions and titrations.^[1]

Reaction Types & Net Ionic Equations

Recognize precipitation (a solid forms), acid-base (water forms), and redox reactions. To write a net ionic equation, split soluble strong electrolytes into ions and cancel the spectators. If every ion is a spectator, no net reaction occurred.

Oxidation–Reduction

Assign oxidation numbers to track electron flow. Oxidation is loss of electrons (number increases); reduction is gain (number decreases) — OIL RIG. In $\ce{SO4^2-}$, four oxygens at −2 and an overall −2 charge force sulfur to +6.

Stoichiometry & Titrations

The limiting reactantcaps the product. In a titration, moles of titrant equal volume × molarity, and at the equivalence point moles of acid and base match by the reaction’s mole ratio.

5 · Kinetics

7–9% of the exam. How fast reactions go: rates, rate laws and reaction order, mechanisms, activation energy, and catalysts.^[1]

Reaction Rates & Rate Laws

Average rate is the change in concentration over the time interval. The rate law $\text{rate} = k[A]^x[B]^y$ must be found by experiment. Use the method of initial rates: if doubling a reactant leaves the rate unchanged it is order 0, doubles it order 1, quadruples it order 2.

The graph that gives a straight line tells the order: concentration vs time → zero order, ln[A] vs time → first order, $1/[A]$ vs time → second order.

Mechanisms, Eₐ & Catalysts

A reaction proceeds over an activation-energy barrier; raising temperature gives more particles enough energy to react (the Arrhenius relationship). A catalyst offers a lower-barrier pathway without being consumed.

6 · Thermodynamics

7–9% of the exam. Energy and heat: systems and surroundings, heat capacity and calorimetry, enthalpy, and Hess’s law.^[1]

Heat, Energy & Calorimetry

The system is what you study; the surroundings are everything else. Heat is energy transferred because of a temperature difference; it flows from warmer to cooler. Heat absorbed by the system is positive q. Calorimetry uses $q = mc\Delta T$.

Enthalpy & Hess’s Law

Enthalpy change $\Delta H$ is negative for exothermic reactions (heat released) and positive for endothermic ones. By Hess’s law, the $\Delta H$ of an overall reaction is the sum of its steps, because enthalpy is a state function. An instant cold pack dissolving a salt is endothermic.

7 · Equilibrium

7–9% of the exam. Dynamic equilibrium, the equilibrium constant K and reaction quotient Q, Le Chatelier’s principle, and solubility (Ksp).^[1]

The Equilibrium Constant & Q

K is the ratio of products to reactants at equilibrium, each raised to its coefficient (pure solids and liquids excluded). Compare the reaction quotient Q to K: $Q < K$ shifts forward, $Q > K$ shifts in reverse, $Q = K$ is at equilibrium.

Le Chatelier & Solubility

Le Chatelier’s principle: a system shifts to partially offset a stress. Only temperature changes the value of K. For a sparingly soluble salt, a precipitate forms when the ion product exceeds the solubility product Ksp.

8 · Acids & Bases

11–15% of the exam — the second-heaviest unit. Bronsted-Lowry acids and bases, pH, strong vs weak acids, buffers, and titrations.^[1]

pH, Strong & Weak Acids

A Bronsted-Lowry acid donates a proton; a base accepts one. The two form a conjugate pair differing by one $\ce{H+}$. pH $= -\log[\ce{H+}]$: below 7 acidic, 7 neutral, above 7 basic, and pH + pOH = 14 at 25°C.

Buffers & Henderson-Hasselbalch

A buffer of a weak acid and its conjugate base resists pH change: $\text{pH} = \text{p}K_a + \log\dfrac{[\text{base}]}{[\text{acid}]}$. Equal amounts give pH = pKₐ; a 10-to-1 base-to-acid ratio gives pH one unit above pKₐ.

Acid–Base Titrations

Before the equivalence point of a weak acid + strong base, the solution is a buffer (pH changes slowly). The equivalence point sits above 7, so an indicator like phenolphthalein (changes near pH 8–10) is appropriate.

9 · Applications of Thermodynamics

7–9% of the exam. Entropy, Gibbs free energy and thermodynamic favorability, and electrochemistry.^[1]

Entropy & Gibbs Free Energy

Entropy measures the dispersal of matter and energy; making more moles of gas raises it. A reaction is favorable when $\Delta G = \Delta H - T\Delta S$ is negative. When $\Delta H$ and $\Delta S$ oppose, temperature decides; the crossover is $T = \Delta H / \Delta S$.

Electrochemistry

A galvanic cell turns a spontaneous redox reaction into electricity: oxidation at the anode, reduction at the cathode, electrons flowing anode → cathode. The standard cell potential $E^\circ = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$; a positive value means spontaneous, and $\Delta G^\circ = -nFE^\circ$.

How to Use This Study Guide

A study guide is a map, not the whole territory — use it alongside official College Board practice and lots of problem solving. AP Chemistry rewards fluency, so spaced, mixed practice beats one long cram. Lead with Units 3 and 8 (intermolecular forces and acid–base), which together are about a third of the multiple-choice section, then shore up the rest.

AP Chemistry Glossary

Quick definitions for the terms you’ll see most across the nine AP Chemistry units:

Activation energy

The minimum energy colliding particles need to react — the height of the energy barrier between reactants and the transition state. A catalyst lowers it.

Avogadro's number

$6.022 \times 10^{23}$, the number of particles in one mole. Multiply moles by this number to get atoms, molecules, or ions.

Bronsted-Lowry acid

A proton ($\ce{H+}$) donor. When it donates a proton it becomes its conjugate base.

Bronsted-Lowry base

A proton ($\ce{H+}$) acceptor. When it accepts a proton it becomes its conjugate acid.

Buffer

A solution of a weak acid and its conjugate base that resists pH change. Its pH is given by Henderson-Hasselbalch: $\text{pH} = \text{p}K_a + \log\frac{[\text{base}]}{[\text{acid}]}$.

Catalyst

A substance that speeds a reaction by providing a lower-activation-energy pathway. It is not consumed and does not change $\Delta H$ or the equilibrium position.

Conjugate acid-base pair

Two species that differ by exactly one proton, such as $\ce{NH4+}$ and $\ce{NH3}$. A strong acid has a weak conjugate base.

Covalent bond

A bond formed by sharing electron pairs between two nonmetal atoms. Single, double, and triple bonds share one, two, or three pairs respectively.

Effective nuclear charge

The net positive pull a valence electron feels after core electrons shield it from the full nuclear charge. It drives the periodic trends in radius and ionization energy.

Electron configuration

The arrangement of an atom's electrons among its subshells, written in order of increasing energy, for example magnesium is $\text{1s}^2\,\text{2s}^2\,\text{2p}^6\,\text{3s}^2$.

Electronegativity

A measure of how strongly an atom attracts the shared electrons in a bond. It increases up and to the right of the periodic table; fluorine is the most electronegative element.

Enthalpy

The heat content of a system at constant pressure. $\Delta H$ is negative for exothermic reactions (heat released) and positive for endothermic ones.

Entropy

A measure of the dispersal of matter and energy, or the number of accessible microstates. Forming more moles of gas increases entropy (positive $\Delta S$).

Equilibrium constant (K)

The ratio of product to reactant concentrations at equilibrium, each raised to its coefficient. A large K favors products; pure solids and liquids are excluded.

Equivalence point

The point in a titration where moles of added titrant exactly neutralize the analyte. Its pH depends on the salt formed (above 7 for a weak acid + strong base).

Galvanic cell

An electrochemical cell that uses a spontaneous redox reaction to produce electricity. Oxidation occurs at the anode, reduction at the cathode.

Gibbs free energy

$\Delta G = \Delta H - T\Delta S$. A negative $\Delta G$ means the process is thermodynamically favorable.

Hess's law

The enthalpy change of an overall reaction equals the sum of the enthalpy changes of the steps, because enthalpy is a state function.

Hybridization

The mixing of atomic orbitals to form equivalent hybrid orbitals: sp (linear), sp² (trigonal planar), and sp³ (tetrahedral) match the electron geometry.

Hydrogen bonding

A strong intermolecular attraction when hydrogen is bonded to fluorine, oxygen, or nitrogen. It explains water's unusually high boiling point.

Intermolecular forces

Attractions between molecules (London dispersion, dipole-dipole, hydrogen bonding, ion-dipole) that determine boiling point, vapor pressure, and solubility.

Ionic bond

A bond formed when a metal transfers electrons to a nonmetal, producing oppositely charged ions held together by electrostatic attraction.

Ionization energy

The energy required to remove an electron from a gaseous atom. It increases across a period and decreases down a group as atoms get larger.

Le Chatelier's principle

A system at equilibrium shifts to partially offset an applied stress (added reactant, pressure, or temperature change). Only temperature changes the value of K.

Limiting reactant

The reactant that runs out first and so caps how much product can form. Compare available moles to the mole ratio to identify it.

London dispersion forces

Weak attractions from temporary, instantaneous dipoles caused by shifting electron clouds. Present in all molecules and stronger with more electrons.

Molar mass

The mass of one mole of a substance in grams per mole, equal numerically to the sum of the atomic masses in its formula. Mass divided by molar mass gives moles.

Mole

The SI unit for amount of substance. One mole contains Avogadro's number of particles, $6.022 \times 10^{23}$. Moles connect mass to the number of particles.

Net ionic equation

An equation showing only the ions and substances that actually react, after soluble strong electrolytes are split into ions and spectator ions are canceled.

Oxidation

The loss of electrons, which raises an atom's oxidation number. Remember OIL RIG: Oxidation Is Loss.

pH

A measure of acidity: $\text{pH} = -\log[\ce{H+}]$. Below 7 is acidic, 7 is neutral, above 7 is basic; at 25°C pH + pOH = 14.

Photoelectron spectroscopy (PES)

A technique that measures the energy needed to remove electrons from an atom. Each peak is a subshell; peak height is proportional to the number of electrons in that subshell.

Pi bond

A bond from side-by-side overlap of p orbitals above and below the bond axis. A double bond has one pi bond; a triple bond has two.

Rate constant

The proportionality factor k in a rate law. It depends on temperature and activation energy but not on concentration.

Rate law

An equation relating reaction rate to reactant concentrations, $\text{rate} = k[A]^x[B]^y$. The orders x and y come from experiment, not the coefficients.

Reaction quotient (Q)

The same ratio as K but at any moment. If Q < K the reaction shifts forward; if Q > K it shifts in reverse; if Q = K it is at equilibrium.

Reduction

The gain of electrons, which lowers an atom's oxidation number. Remember OIL RIG: Reduction Is Gain.

Sigma bond

A bond from head-on (end-to-end) orbital overlap along the bond axis. Every single bond is one sigma bond.

Solubility product (Ksp)

The equilibrium constant for a sparingly soluble salt dissolving into its ions. A precipitate forms when the ion product Q exceeds Ksp.

Specific heat capacity

The energy needed to raise one gram of a substance by one degree Celsius, used in $q = mc\Delta T$. Water's is unusually high.

Standard cell potential

$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$. A positive value means a spontaneous (galvanic) cell.

Vapor pressure

The pressure of vapor in equilibrium with its liquid. Stronger intermolecular forces lower vapor pressure; a liquid boils when its vapor pressure equals the external pressure.

VSEPR theory

Valence Shell Electron Pair Repulsion: electron domains around a central atom spread out to minimize repulsion, which predicts molecular geometry and bond angles.

Free AP Chemistry Study Materials & Resources

Everything you need to prepare for AP Chemistry is free here — no paywall, no sign-up. This guide is the foundation; pair it with the rest of our free AP Chemistry study materials for active recall, timed practice, and last-minute review:

• AP Chemistry Practice Test — exam-style questions across all nine units, with explanations.
• AP Chemistry Flashcards — active-recall decks for the laws, formulas, reactions, and definitions you have to know cold.

AP Chemistry Study Guide FAQ

References

1. 1.College Board. “AP Chemistry Course and Exam Description.” College Board. ↑
2. 2.College Board. “AP Chemistry Exam — AP Students.” College Board. ↑
3. 3.College Board. “AP Chemistry Course Overview.” College Board. ↑
4. 4.College Board. “AP Chemistry Course at a Glance.” College Board. ↑

Sources for the concept answers

Every answer in the AP Chemistry concept questions above is drawn from an official primary source:

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