FREE ACS General Chemistry Exam Study Guide 2026

This free ACS General Chemistry Exam study guide covers everything the ACS General Chemistry Exam tests — the standardized general chemistry final from the American Chemical Society’s Examinations Institute (ACS Exams), used as a course final at colleges nationwide.^[1][2]

The standard full-year and single-term forms are 70 multiple-choice questions in 110 minutes, organized around the ACS Anchoring Concepts Content Map — 10 big ideas spanning a full two-semester general chemistry sequence.^[4] It is norm-referenced: ACS reports your raw score and a national percentile rank, and your instructor decides how that maps to a grade — there is no single official pass mark.^[3]

It’s interactive, not a wall of text: every module has high-yield notes, worked examples, labeled diagrams, flashcards, and practice questions, so you learn by doing — not just reading. Drill gaps with our free ACS practice test and flashcards.

ACS Gen Chem Exam Snapshot

The ACS General Chemistry Exam is a secure, standardized, multiple-choice exam. Question counts and times depend on the form your instructor orders; the standard forms are below.^[1] ACS publishes no passing score — it provides national norms, and grading is up to the instructor.^[3]

That pace is about 94 seconds per question on a 70-item form — fast enough that you must recognize concepts quickly rather than derive everything from scratch. Because the exam spans two semesters, the winning strategy is broad, even review, not deep cramming of one unit.^[1]

Atoms & Periodicity

The Atoms concept is the foundation of everything else: the internal structure of atoms dictates their chemical and physical behavior. Master atomic structure, the quantum model, electron configurations, and periodic trends, and a large slice of the exam falls into place.^[4]

Atomic structure & isotopes

An atom is a tiny, dense nucleus of protons and neutrons surrounded by electrons. The number of protons (the atomic number) defines the element; isotopes are atoms of the same element with different numbers of neutrons, and therefore different mass numbers.^[6]Average atomic mass is the weighted average of an element’s isotopes by natural abundance.

Quantum numbers & electron configuration

Each electron is described by four quantum numbers: $n$ (principal — energy level), $\ell$ (azimuthal — orbital shape s/p/d/f), $m_\ell$ (magnetic — orientation), and $m_s$ (spin, $\pm\tfrac{1}{2}$). Electrons fill orbitals by the Aufbau principle (lowest energy first), Hund’s rule (singly fill degenerate orbitals first), and the Pauli exclusion principle (max 2 electrons per orbital).^[6]

Photons connect to atomic structure too: the energy of a photon is $E = h\nu = \frac{hc}{\lambda}$, and a hydrogen electron’s transitions follow $\Delta E = -2.18\times10^{-18}\,\text{J}\left(\frac{1}{n_f^2} - \frac{1}{n_i^2}\right)$. A larger drop in $n$ releases a higher-energy photon.^[9]

Periodic trends

Periodic trends are pure point-scoring once you internalize the direction. Across a period (left to right) atomic radius decreases while ionization energy and electronegativity increase, as the growing nuclear charge pulls electrons in. Down a group, radius increases and ionization energy and electronegativity decrease, as added shells move electrons farther out.^[6]

Bonding & Molecular Structure

Two anchoring concepts live here: Bonding (atoms interact via electrostatic forces to form bonds) and Structure & Function(a molecule’s geometry shapes its behavior). This is one of the most heavily tested clusters on the exam.^[4]

Ionic, covalent & polar bonds

An ionic bond transfers electrons from a metal to a nonmetal, leaving oppositely charged ions held by electrostatic attraction (NaCl). A covalent bondshares electrons between nonmetals. The electronegativity difference decides: a large difference (> ~1.7) is ionic, a moderate one (~0.4–1.7) is polar covalent, and a tiny one is nonpolar covalent.^[8]

Lewis structures & resonance

To draw a Lewis structure: total the valence electrons, place the least electronegative atom (never H) in the center, form single bonds, then add lone pairs to satisfy the octet (adding multiple bonds if an atom is short). Formal charge helps choose the best structure: $\text{FC} = (\text{valence } e^-) - (\text{nonbonding } e^-) - \tfrac{1}{2}(\text{bonding } e^-)$. Resonance structures differ only in where electrons are drawn — the nuclei stay put, and the true molecule is a hybrid of all contributors.^[8]

VSEPR, geometry & hybridization

VSEPR theory predicts shape from the steric number (bonded atoms + lone pairs). Steric number sets the electron geometry; lone pairs then bend the molecular shape — so water (steric number 4, two lone pairs) is tetrahedral in electron geometry but bent (~104.5°) in shape, while ammonia is trigonal pyramidal.^[8]

Hybridization follows the steric number directly: 2 groups → $sp$, 3 → $sp^2$, 4 → $sp^3$, 5 → $sp^3d$, 6 → $sp^3d^2$. A molecule is polar when its bond dipoles do not cancel — symmetry matters, so $\ce{CO2}$ is nonpolar (linear, dipoles cancel) but $\ce{H2O}$ is polar (bent, they don’t).^[8]

States, Forces & Solutions

The Intermolecular Interactions anchoring concept explains the physical behavior of matter — why substances boil, dissolve, and change phase. This module also folds in gases and solutions, which lean on those same forces.^[4]

Intermolecular forces

Intermolecular forces are attractions betweenmolecules, far weaker than the bonds within them, but they set boiling point, melting point, viscosity, and solubility. Strongest to weakest: ion–dipole, hydrogen bonding(H bonded to N, O, or F), dipole–dipole, and London dispersion (present in everything).^[8]

Gases & the gas laws

The ideal gas law is $PV = nRT$ (use kelvin and $R = 0.0821\ \text{L·atm/mol·K}$). The simpler laws fall out of it: Boyle’s ($P \propto 1/V$), Charles’s ($V \propto T$), and the combined law $\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}$. Dalton’s law adds partial pressures, and Graham’s law of effusion gives $\frac{r_1}{r_2} = \sqrt{\frac{M_2}{M_1}}$.^[6]

Solutions & colligative properties

Colligative properties depend only on the number of dissolved particles, not their identity. A non-volatile solute lowers vapor pressure (Raoult’s law), raises the boiling point ($\Delta T_b = i K_b m$), and lowers the freezing point ($\Delta T_f = i K_f m$), where the van’t Hoff factor $i$ counts the particles each formula unit produces. Osmotic pressure is $\Pi = i M R T$.^[6]

Reactions, Stoichiometry & Thermochemistry

This module merges two anchoring concepts: Chemical Reactions (matter changes into new products) and Energy & Thermodynamics (energy is the currency of those changes). Stoichiometry and thermochemistry together are among the highest-yield, most calculation-heavy parts of the exam.^[4]

Stoichiometry & reaction types

Stoichiometry uses the mole ratios in a balanced equation. The workflow is always: balance → convert grams to moles (divide by molar mass) → apply the coefficient ratio → convert back. The limiting reactant runs out first and caps the theoretical yield; percent yield is $\frac{\text{actual}}{\text{theoretical}} \times 100\%$.^[6]

Thermochemistry & enthalpy

Enthalpy ($H$) is heat at constant pressure; $\Delta H < 0$ is exothermic (releases heat) and $\Delta H > 0$ is endothermic. Calorimetry uses $q = m c \Delta T$. Hess’s law lets you add the $\Delta H$ of steps (reversing a step flips its sign), and standard enthalpies of formation give $\Delta H^{\circ}_{rxn} = \sum \Delta H_f^{\circ}(\text{products}) - \sum \Delta H_f^{\circ}(\text{reactants})$.^[7]

Entropy & Gibbs free energy

Entropy ($S$) measures disorder; the second law says the entropy of the universe increases for any spontaneous process. The decision-maker is the Gibbs free energy: a process is spontaneous when $\Delta G = \Delta H - T\Delta S < 0$, at equilibrium when $\Delta G = 0$.^[7]

Redox & electrochemistry

In a redox reaction, one species is oxidized (loses electrons) and another is reduced (gains them) — OIL RIG. In a galvanic cell, oxidation happens at the anode and reduction at the cathode (AN OX, RED CAT), and the cell potential is $E^{\circ}_{cell} = E^{\circ}_{cathode} - E^{\circ}_{anode}$. Under non-standard conditions, the Nernst equation applies: $E = E^{\circ} - \frac{0.0592}{n}\log Q$ at 25 °C, and $\Delta G^{\circ} = -nFE^{\circ}_{cell}$.^[9]

Kinetics, Equilibrium, Acids & Bases

Two more anchoring concepts anchor this module: Kinetics (how fast a reaction goes) and Equilibrium(where a reversible reaction settles). Acid–base chemistry and solubility are equilibrium problems, so they live here too.^[4]

Reaction rates & kinetics

A rate law, $\text{rate} = k[A]^m[B]^n$, gives how rate depends on concentration; the orders $m$ and $n$ come from experiment, not the coefficients. A first-order reaction has a constant half-life $t_{1/2} = \frac{0.693}{k}$, and the Arrhenius equation $k = A e^{-E_a / RT}$ links the rate constant to temperature and the activation energy.^[8]

A catalyst lowers $E_a$ via an alternate pathway, speeding both directions without being consumed — and crucially, it does not change $\Delta H$ or the position of equilibrium. The slowest elementary step (the rate-determining step) controls the overall rate.^[8]

Chemical equilibrium

At equilibrium, forward and reverse rates are equal and the equilibrium constant $K$ is fixed at a given temperature: $K > 1$ favors products, $K < 1$ favors reactants. Compare the reaction quotient $Q$ to $K$ to predict the direction of shift, and use Le Chatelier’s principle for disturbances. Raising temperature shifts an endothermic equilibrium toward products (raising $K$).^[6]

Acids, bases & buffers

pH is $-\log[\ce{H+}]$, and at 25 °C $\text{pH} + \text{pOH} = 14$ with $K_w = 1.0\times10^{-14}$. Strong acids and bases ionize completely; weak ones only partially, described by $K_a$ or $K_b$ (with $K_a K_b = K_w$). A bufferresists pH change; its pH is the Henderson–Hasselbalch equation $\text{pH} = \text{p}K_a + \log\frac{[\ce{A-}]}{[\ce{HA}]}$.^[6]

Solubility & Ksp

For a slightly soluble salt $\ce{A_xB_y}$, the solubility product is $K_{sp} = [\ce{A}]^x[\ce{B}]^y$. A smaller $K_{sp}$ means a less soluble compound, and the common-ion effect (adding an ion already in the equilibrium) pushes it toward the solid, lowering solubility.^[6]

Measurement, Data & Lab Skills

The last two anchoring concepts — Experiments, Measurement & Data and Visualization— make sure you can handle units, precision, lab technique, and the link between what you observe (macroscopic) and what’s happening to particles (particulate).^[4]

Units, sig figs & uncertainty

Significant figures communicate precision. When multiplying or dividing, keep the fewest significant figures of any factor; when adding or subtracting, keep the fewest decimal places. Distinguish accuracy (closeness to the true value) from precision (reproducibility), and recognize that a systematic error (like an uncalibrated balance) biases every reading the same way, while a random error scatters them.^[6]

Lab techniques & error

Know the standard general-chemistry techniques: titration (find concentration via an indicator color change at the equivalence point), distillation (separate by boiling point), gravimetric analysis (mass of a precipitate), and calorimetry (measure heat). On the Visualization concept, be ready to move between the three levels of representation — particulate (atoms and molecules), macroscopic (what you see), and symbolic (formulas, equations, graphs).^[4]

How the ACS Exam Is Scored

The ACS General Chemistry Exam is norm-referenced, not pass/fail. ACS computes your raw score (number correct) and converts it to a national percentile rank using a large, multi-year sample of students who took the same form. Instructors also receive the mean, standard deviation, and reliability statistics for their class.^[3]

Crucially, ACS does not set a passing score. Your instructor decides how the percentile maps to a course grade — some use the national median as a benchmark, others curve to their own class. Any “you need X% to pass” figure you see online is not an ACS standard, so confirm your course’s grading policy with your instructor.^[2][3]

How to Use This Study Guide

The ACS exam spans two semesters, so reward yourself with a broad, structured review rather than a last-night cram. Pair every reading session with active practice — that is where the score gains live:

ACS Gen Chem Glossary

Quick definitions for the terms you’ll meet most across the ACS General Chemistry Exam’s anchoring concepts:

ACS General Chemistry Exam

A standardized, secure general chemistry final exam from the American Chemical Society's Examinations Institute (ACS Exams). Standard full-year and single-term forms have 70 multiple-choice questions in 110 minutes; results are reported against national percentile norms rather than a fixed passing score.

activation energy

The minimum energy a collision must have for a reaction to occur (Eₐ); catalysts speed reactions by lowering it.

anchoring concept

One of the 10 big ideas in the ACS Anchoring Concepts Content Map (ACCM) that organizes the general chemistry curriculum: Atoms, Bonding, Structure & Function, Intermolecular Interactions, Chemical Reactions, Energy & Thermodynamics, Kinetics, Equilibrium, Experiments/Measurement/Data, and Visualization.

Aufbau principle

The rule that electrons fill the lowest-energy orbitals first when building up an atom's ground-state electron configuration (so 4s fills before 3d).

buffer

A solution of a weak acid and its conjugate base (or vice versa) that resists pH change; its pH follows the Henderson–Hasselbalch equation.

catalyst

A substance that speeds up a reaction by providing a lower-activation-energy pathway without being consumed; it does not change ΔH or the position of equilibrium.

colligative property

A solution property that depends on the number of dissolved particles, not their identity — boiling-point elevation, freezing-point depression, vapor-pressure lowering, and osmotic pressure.

covalent bond

A bond formed when two atoms share one or more pairs of electrons; equal sharing is nonpolar, unequal sharing is polar covalent.

electronegativity

A measure of how strongly an atom attracts the shared electrons in a bond. It increases across a period and decreases down a group; fluorine is the most electronegative element.

enthalpy

The heat content of a system at constant pressure (H); a reaction's enthalpy change ΔH is negative for exothermic and positive for endothermic reactions.

entropy

A measure of the disorder or dispersal of energy and matter in a system (S); the second law of thermodynamics states the entropy of the universe increases for any spontaneous process.

equilibrium constant

Keq, the ratio of product to reactant concentrations (each raised to its coefficient) at equilibrium; K > 1 favors products, K < 1 favors reactants.

galvanic cell

An electrochemical cell that uses a spontaneous redox reaction to produce electricity; oxidation occurs at the anode and reduction at the cathode.

Gibbs free energy

The thermodynamic quantity ΔG = ΔH − TΔS that predicts spontaneity: a process is spontaneous when ΔG is negative, at equilibrium when ΔG is zero.

Hess's law

The principle that the overall enthalpy change of a reaction is the sum of the enthalpy changes of its steps, because enthalpy is a state function.

Hund's rule

The rule that electrons singly occupy each orbital in a subshell before any orbital is doubly occupied, keeping unpaired spins parallel.

hybridization

The mixing of atomic orbitals to form equivalent hybrid orbitals (sp, sp², sp³, sp³d, sp³d²) that match a molecule's geometry.

hydrogen bond

An especially strong dipole–dipole attraction between a hydrogen atom bonded to N, O, or F and a lone pair on another such atom. It gives water its high boiling point.

ideal gas law

PV = nRT — relates pressure, volume, moles, and absolute temperature of a gas through the gas constant R (0.0821 L·atm/mol·K).

intermolecular force

An attractive force between molecules (London dispersion, dipole–dipole, hydrogen bonding, ion–dipole), much weaker than the bonds within a molecule but controlling phase, boiling point, and solubility.

ionic bond

A bond formed by the electrostatic attraction between oppositely charged ions, created when electrons transfer from a metal to a nonmetal (e.g., NaCl).

isotope

Atoms of the same element (same number of protons) that have different numbers of neutrons, and therefore different mass numbers.

Le Chatelier's principle

When a system at equilibrium is disturbed, it shifts to partly counteract the change and restore equilibrium.

limiting reactant

The reactant that is completely consumed first and therefore caps the amount of product that can form (the theoretical yield).

mole

The SI unit for amount of substance; one mole contains Avogadro's number (6.022 × 10²³) of particles. It is the bridge between the mass of a sample and the number of atoms or molecules it contains.

pH

A logarithmic measure of acidity, pH = −log[H⁺]. Below 7 is acidic, 7 is neutral, above 7 is basic; at 25 °C, pH + pOH = 14.

quantum number

One of four values (n, ℓ, mℓ, mₛ) that together specify an electron's energy, orbital shape, orientation, and spin. No two electrons in an atom share all four (the Pauli exclusion principle).

rate law

An expression, rate = k[A]ᵐ[B]ⁿ, giving how a reaction's rate depends on reactant concentrations; the exponents (orders) are determined experimentally.

significant figures

The meaningful digits in a measurement that convey its precision; calculations carry the appropriate number of significant figures to avoid overstating accuracy.

solubility product

Ksp, the equilibrium constant for the dissolution of a slightly soluble ionic solid; a smaller Ksp means a less soluble compound.

steric number

The number of electron groups around a central atom — bonded atoms plus lone pairs. It sets the electron-domain geometry in VSEPR.

stoichiometry

Using the mole ratios in a balanced equation to relate the amounts of reactants and products — the quantitative heart of the Chemical Reactions anchoring concept.

VSEPR theory

Valence Shell Electron Pair Repulsion theory, which predicts molecular geometry by assuming electron groups around a central atom arrange to minimize repulsion.

Free ACS Gen Chem Study Materials & Resources

Everything you need to prepare for the ACS General Chemistry Exam is free here — no paywall, no sign-up. This guide is the foundation; pair it with the rest of our free ACS study materials:

• ACS General Chemistry Practice Test — realistic, concept-weighted questions with explanations.
• ACS General Chemistry Flashcards — active-recall decks for the high-yield facts and formulas.

ACS Gen Chem Study Guide FAQ

References

1. 1.ACS Examinations Institute. “ACS Exams — General Chemistry Exams (Assessment Materials).” uwm.edu/acs-exams. ↑
2. 2.ACS Examinations Institute. “About ACS Exams.” uwm.edu/acs-exams. ↑
3. 3.ACS Examinations Institute. “National Norms and Score Reporting.” uwm.edu/acs-exams. ↑
4. 4.Holme, T.; Murphy, K.. “The ACS Exams Institute Undergraduate Chemistry Anchoring Concepts Content Map I: General Chemistry.” J. Chem. Educ. 2012, 89 (6), 721–723. ↑
5. 5.ACS Examinations Institute. “General Chemistry Official Study Guide.” uwm.edu/acs-exams. ↑
6. 6.National Institute of Standards and Technology (NIST). “Periodic Table, Atomic Structure, and Periodic Trends.” nist.gov. ↑
7. 7.National Institute of Standards and Technology (NIST). “NIST Chemistry WebBook (Thermochemistry and Reference Data).” nist.gov. ↑
8. 8.International Union of Pure and Applied Chemistry (IUPAC). “Compendium of Chemical Terminology (Gold Book).” iupac.org. ↑
9. 9.National Institute of Standards and Technology (NIST). “CODATA Recommended Values of the Fundamental Physical Constants.” physics.nist.gov. ↑

Sources for the concept answers

Every answer in the ACS Gen Chem concept questions above is drawn from an official or authoritative primary source:

1. National Institute of Standards and Technology (NIST). “The Mole, Avogadro's Number, and Stoichiometry.” nist.gov, accessed 20 June 2026.
2. International Union of Pure and Applied Chemistry (IUPAC). “Catalyst and Activation Energy (Gold Book).” iupac.org, accessed 20 June 2026.
3. International Union of Pure and Applied Chemistry (IUPAC). “Rate Law, Order of Reaction, and the Arrhenius Equation (Gold Book).” iupac.org, accessed 20 June 2026.
4. National Institute of Standards and Technology (NIST). “Atomic Structure and Quantum Numbers.” nist.gov, accessed 20 June 2026.
5. International Union of Pure and Applied Chemistry (IUPAC). “Chemical Bonding: Ionic, Covalent, and Polar Bonds (Gold Book).” iupac.org, accessed 20 June 2026.
6. International Union of Pure and Applied Chemistry (IUPAC). “Molecular Geometry and VSEPR (Gold Book).” iupac.org, accessed 20 June 2026.
7. International Union of Pure and Applied Chemistry (IUPAC). “Intermolecular Forces and Hydrogen Bonding (Gold Book).” iupac.org, accessed 20 June 2026.

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