Are flashcards effective for the ACS chemistry exam?

Yes. Flashcards use active recall — retrieving an answer from memory — which research shows is one of the most effective study methods, especially for the formula- and definition-heavy content tested on the ACS General Chemistry Exam, when used in short sessions spread over weeks.

What topics do the ACS General Chemistry flashcards cover?

All 10 ACS anchoring concepts: atomic structure and periodic trends, bonding and VSEPR geometry, intermolecular forces and gas laws, stoichiometry and redox, thermochemistry and Gibbs free energy, kinetics, equilibrium, acids/bases and buffers, solubility, and measurement and lab technique.

How should I study with these flashcards?

Because the ACS exam spans two semesters, review every concept rather than one unit. Mix the modes: flip to learn, type to test recall, match for speed, and quiz to self-check. Revisit the cards you miss, and pair the deck with our free practice test and study guide.

Are the ACS General Chemistry flashcards free?

Yes — 100% free, all four study modes, no paywall.

Electrons fill the lowest-energy orbitals first when building an atom's ground-state configuration — so 4s fills before 3d.

Pauli exclusion principle

No two electrons in an atom can have the same four quantum numbers; an orbital holds at most 2 electrons, with opposite spins.

Electrons singly occupy each orbital in a subshell, with parallel spins, before any orbital is doubly occupied.

The four quantum numbers

n (energy level), ℓ (orbital shape: s/p/d/f), mℓ (orientation), and mₛ (spin, ±½).

Atoms of the same element (same protons) with different numbers of neutrons, and so different mass numbers.

Atomic number vs. mass number

Atomic number = number of protons (defines the element); mass number = protons + neutrons.

Decreases across a period (left → right) and increases down a group.

Trend: ionization energy

Increases across a period and decreases down a group (it takes more energy to remove an electron from a smaller, tightly held atom).

Trend: electronegativity

Increases across a period and decreases down a group; fluorine is the most electronegative element.

E = hν = hc/λ — higher frequency (shorter wavelength) means higher energy.

Maximum electrons per subshell

s holds 2, p holds 6, d holds 10, f holds 14 (each orbital holds 2).

Electron configuration of a cation

Remove electrons from the highest n first — for transition metals the 4s electrons leave before the 3d (e.g., Fe²⁺ is [Ar]3d⁶).

Elements with a full valence shell, very stable and unreactive — all are monatomic gases at room temperature.

The abundance-weighted average mass of an element's naturally occurring isotopes.

Electrostatic attraction between oppositely charged ions formed when a metal transfers electrons to a nonmetal (e.g., NaCl).

A bond formed by two atoms sharing one or more pairs of electrons (typically between nonmetals).

Polar vs. nonpolar covalent bond

Unequal electron sharing (electronegativity difference ~0.4–1.7) is polar covalent; near-equal sharing is nonpolar.

Electron groups around a central atom arrange to minimize repulsion, which sets the molecule's geometry.

Bonded atoms + lone pairs on the central atom; it sets the electron-domain geometry and the hybridization.

Geometry: steric number 2 / 3 / 4

2 → linear (180°), 3 → trigonal planar (120°), 4 → tetrahedral (109.5°).

Hybridization by steric number

2 → sp, 3 → sp², 4 → sp³, 5 → sp³d, 6 → sp³d².

Oxygen has steric number 4 (2 bonds + 2 lone pairs); the lone pairs push the bonds to a bent ~104.5° shape.

Electron geometry vs. molecular geometry

Electron geometry counts all bonding pairs and lone pairs; molecular geometry describes only the arrangement of bonded atoms.

FC = (valence electrons) − (nonbonding electrons) − ½(bonding electrons); the best Lewis structure minimizes formal charges.

Structures that differ only in electron placement, not nuclear positions; the real molecule is a hybrid of all of them.

A sigma bond is head-on orbital overlap along the bond axis; a pi bond is side-to-side overlap above and below the axis.

Single, double, triple bonds

A single bond is 1 sigma; a double is 1 sigma + 1 pi; a triple is 1 sigma + 2 pi (shorter and stronger).

It is polar only if its individual bond dipoles do NOT cancel — CO₂ is nonpolar (symmetric), H₂O is polar (bent).

Lewis structure steps

Total valence electrons → least-electronegative atom (not H) in center → single bonds → distribute lone pairs to satisfy octets → add multiple bonds if needed.

Octet rule (and exceptions)

Atoms tend to gain a full 8-electron valence shell; exceptions include H (2), incomplete octets (B), and expanded octets (P, S).

Bond dipole direction

Points from the less electronegative atom toward the more electronegative atom.

Intermolecular forces (strongest to weakest)

Ion–dipole > hydrogen bonding > dipole–dipole > London dispersion.

A strong dipole–dipole attraction when H is bonded to N, O, or F; it gives water its high boiling point and surface tension.

London dispersion forces

Weak, temporary-dipole attractions present in ALL molecules; they grow with molar mass and polarizability.

Effect of stronger IMFs

Higher boiling point and melting point, higher viscosity and surface tension, and lower vapor pressure.

PV = nRT — relates pressure, volume, moles, and absolute temperature; R = 0.0821 L·atm/mol·K.

At constant temperature, pressure and volume are inversely related: P ∝ 1/V, so P₁V₁ = P₂V₂.

At constant pressure, volume is directly proportional to absolute temperature: V ∝ T.

P₁V₁/T₁ = P₂V₂/T₂ — temperature must be in kelvin.

Dalton's law of partial pressures

The total pressure of a gas mixture equals the sum of the partial pressures of its components.

Graham's law of effusion

Lighter gases effuse faster: r₁/r₂ = √(M₂/M₁).

Kinetic molecular theory

Gas particles are tiny, in constant random motion, with negligible volume and no IMFs; average kinetic energy ∝ absolute temperature.

A solution property that depends on the NUMBER of dissolved particles, not their identity.

Boiling-point elevation

A non-volatile solute raises the boiling point: ΔTb = i·Kb·m (m = molality, i = van't Hoff factor).

Freezing-point depression

A solute lowers the freezing point: ΔTf = i·Kf·m — why salt melts ice on roads.

Your FREE ACS General Chemistry Flashcards 2026 – 100+ Cards

Realistic, exam-style ACS General Chemistry flashcards across atoms, bonding, reactions, thermodynamics, kinetics, equilibrium, and acids/bases — flip, match, type, and quiz yourself.

How well do you know them?

ACS General Chemistry Exam Flashcards

Aufbau principle: Electrons fill the lowest-energy orbitals first when building an atom's ground-state configuration — so 4s fills before 3d.
Pauli exclusion principle: No two electrons in an atom can have the same four quantum numbers; an orbital holds at most 2 electrons, with opposite spins.
Hund's rule: Electrons singly occupy each orbital in a subshell, with parallel spins, before any orbital is doubly occupied.
The four quantum numbers: n (energy level), ℓ (orbital shape: s/p/d/f), mℓ (orientation), and mₛ (spin, ±½).
Isotopes: Atoms of the same element (same protons) with different numbers of neutrons, and so different mass numbers.
Atomic number vs. mass number: Atomic number = number of protons (defines the element); mass number = protons + neutrons.
Trend: atomic radius: Decreases across a period (left → right) and increases down a group.
Trend: ionization energy: Increases across a period and decreases down a group (it takes more energy to remove an electron from a smaller, tightly held atom).
Trend: electronegativity: Increases across a period and decreases down a group; fluorine is the most electronegative element.
Energy of a photon: E = hν = hc/λ — higher frequency (shorter wavelength) means higher energy.
Maximum electrons per subshell: s holds 2, p holds 6, d holds 10, f holds 14 (each orbital holds 2).
Electron configuration of a cation: Remove electrons from the highest n first — for transition metals the 4s electrons leave before the 3d (e.g., Fe²⁺ is [Ar]3d⁶).
Effective nuclear charge: The net positive charge an electron feels after shielding by inner electrons; it rises across a period, pulling electrons in.
Noble gas (Group 18): Elements with a full valence shell, very stable and unreactive — all are monatomic gases at room temperature.
Average atomic mass: The abundance-weighted average mass of an element's naturally occurring isotopes.
Ionic bond: Electrostatic attraction between oppositely charged ions formed when a metal transfers electrons to a nonmetal (e.g., NaCl).
Covalent bond: A bond formed by two atoms sharing one or more pairs of electrons (typically between nonmetals).
Polar vs. nonpolar covalent bond: Unequal electron sharing (electronegativity difference ~0.4–1.7) is polar covalent; near-equal sharing is nonpolar.
VSEPR theory: Electron groups around a central atom arrange to minimize repulsion, which sets the molecule's geometry.
Steric number: Bonded atoms + lone pairs on the central atom; it sets the electron-domain geometry and the hybridization.
Geometry: steric number 2 / 3 / 4: 2 → linear (180°), 3 → trigonal planar (120°), 4 → tetrahedral (109.5°).
Hybridization by steric number: 2 → sp, 3 → sp², 4 → sp³, 5 → sp³d, 6 → sp³d².
Why is water bent?: Oxygen has steric number 4 (2 bonds + 2 lone pairs); the lone pairs push the bonds to a bent ~104.5° shape.
Electron geometry vs. molecular geometry: Electron geometry counts all bonding pairs and lone pairs; molecular geometry describes only the arrangement of bonded atoms.
Formal charge: FC = (valence electrons) − (nonbonding electrons) − ½(bonding electrons); the best Lewis structure minimizes formal charges.
Resonance structures: Structures that differ only in electron placement, not nuclear positions; the real molecule is a hybrid of all of them.
Sigma vs. pi bond: A sigma bond is head-on orbital overlap along the bond axis; a pi bond is side-to-side overlap above and below the axis.
Single, double, triple bonds: A single bond is 1 sigma; a double is 1 sigma + 1 pi; a triple is 1 sigma + 2 pi (shorter and stronger).
Is a molecule polar?: It is polar only if its individual bond dipoles do NOT cancel — CO₂ is nonpolar (symmetric), H₂O is polar (bent).
Lewis structure steps: Total valence electrons → least-electronegative atom (not H) in center → single bonds → distribute lone pairs to satisfy octets → add multiple bonds if needed.
Octet rule (and exceptions): Atoms tend to gain a full 8-electron valence shell; exceptions include H (2), incomplete octets (B), and expanded octets (P, S).
Bond dipole direction: Points from the less electronegative atom toward the more electronegative atom.
Intermolecular forces (strongest to weakest): Ion–dipole > hydrogen bonding > dipole–dipole > London dispersion.
Hydrogen bonding: A strong dipole–dipole attraction when H is bonded to N, O, or F; it gives water its high boiling point and surface tension.
London dispersion forces: Weak, temporary-dipole attractions present in ALL molecules; they grow with molar mass and polarizability.
Effect of stronger IMFs: Higher boiling point and melting point, higher viscosity and surface tension, and lower vapor pressure.
Ideal gas law: PV = nRT — relates pressure, volume, moles, and absolute temperature; R = 0.0821 L·atm/mol·K.
Boyle's law: At constant temperature, pressure and volume are inversely related: P ∝ 1/V, so P₁V₁ = P₂V₂.
Charles's law: At constant pressure, volume is directly proportional to absolute temperature: V ∝ T.
Combined gas law: P₁V₁/T₁ = P₂V₂/T₂ — temperature must be in kelvin.
Dalton's law of partial pressures: The total pressure of a gas mixture equals the sum of the partial pressures of its components.
Graham's law of effusion: Lighter gases effuse faster: r₁/r₂ = √(M₂/M₁).
Kinetic molecular theory: Gas particles are tiny, in constant random motion, with negligible volume and no IMFs; average kinetic energy ∝ absolute temperature.
Colligative property: A solution property that depends on the NUMBER of dissolved particles, not their identity.
Boiling-point elevation: A non-volatile solute raises the boiling point: ΔTb = i·Kb·m (m = molality, i = van't Hoff factor).
Freezing-point depression: A solute lowers the freezing point: ΔTf = i·Kf·m — why salt melts ice on roads.
Raoult's law: The vapor pressure of a solvent over an ideal solution is lowered: P = X(solvent)·P° (X = mole fraction).
Van't Hoff factor (i): The number of particles a solute produces in solution: NaCl gives i ≈ 2, glucose i = 1.
Like dissolves like: Polar/ionic solutes dissolve in polar solvents; nonpolar solutes dissolve in nonpolar solvents.
Critical point: The temperature and pressure above which a distinct liquid phase can no longer exist (a supercritical fluid).
Stoichiometry: Using the mole ratios in a balanced equation to relate amounts of reactants and products.
The mole: The amount of substance containing Avogadro's number (6.022 × 10²³) of particles; it bridges mass and number of particles.
Moles from mass: n = mass ÷ molar mass.
Limiting reactant: The reactant consumed first; it caps the amount of product (theoretical yield).
Percent yield: (actual yield ÷ theoretical yield) × 100%.
Molarity: Concentration = moles of solute ÷ liters of solution (mol/L, M).
Dilution equation: M₁V₁ = M₂V₂ — moles of solute are conserved when you add solvent.
Oxidation vs. reduction (OIL RIG): Oxidation Is Loss of electrons; Reduction Is Gain of electrons. They always occur together.
Oxidizing vs. reducing agent: The oxidizing agent is reduced (gains electrons); the reducing agent is oxidized (loses electrons).
Combustion reaction: A hydrocarbon + O₂ → CO₂ + H₂O, releasing heat (exothermic).
Precipitation reaction: Two soluble salts react to form an insoluble solid (precipitate) — predicted with solubility rules.
Acid–base (neutralization) reaction: An acid + a base → a salt + water (H⁺ + OH⁻ → H₂O).
Avogadro's number: 6.022 × 10²³ — the number of particles in one mole.
Enthalpy (ΔH): Heat exchanged at constant pressure; ΔH < 0 is exothermic (releases heat), ΔH > 0 is endothermic (absorbs heat).
Calorimetry equation: q = mcΔT — heat = mass × specific heat × temperature change.
Hess's law: The overall ΔH of a reaction equals the sum of the ΔH of its steps, because enthalpy is a state function.
ΔH from heats of formation: ΔH°rxn = ΣΔHf°(products) − ΣΔHf°(reactants).
Standard enthalpy of formation of an element: Zero — for an element in its standard state (e.g., O₂(g), N₂(g)).
Entropy (ΔS): A measure of disorder / dispersal of energy and matter; gases have higher entropy than liquids or solids.
Second law of thermodynamics: The total entropy of the universe (system + surroundings) increases for any spontaneous process.
Gibbs free energy: ΔG = ΔH − TΔS; a process is spontaneous when ΔG < 0, at equilibrium when ΔG = 0.
When is a reaction spontaneous at all temperatures?: When ΔH < 0 (exothermic) and ΔS > 0 (more disorder) — both favor a negative ΔG.
When is a reaction never spontaneous?: When ΔH > 0 (endothermic) and ΔS < 0 (less disorder) — ΔG is positive at every temperature.
Sign of ΔS for fewer moles of gas: Negative — converting more moles of gas to fewer (e.g., N₂ + 3H₂ → 2NH₃) decreases disorder.
First law of thermodynamics: Energy is conserved: ΔU = q + w (internal energy change = heat added + work done on the system).
ΔG° and the equilibrium constant: ΔG° = −RT ln K — a negative ΔG° means K > 1 (products favored).
Rate law: rate = k[A]ᵐ[B]ⁿ; the orders m and n are found experimentally, NOT from the balanced coefficients.
Reaction order: The exponent on a reactant in the rate law; the overall order is the sum of the exponents.
Activation energy (Eₐ): The minimum collision energy needed for a reaction to occur — the barrier on the energy profile.
How does a catalyst work?: It provides an alternate pathway with lower Eₐ, speeding both directions; it is not consumed and does not change ΔH or K.
First-order half-life: t½ = 0.693/k — constant, independent of the starting concentration.
Arrhenius equation: k = A·exp(−Eₐ/RT); raising temperature or lowering the activation energy Eₐ increases the rate constant k.
Collision theory: A collision leads to reaction only if the molecules have enough energy (≥ Eₐ) AND the correct orientation.
Rate-determining step: The slowest elementary step in a mechanism; it controls the overall reaction rate.
Effect of temperature on rate: Higher temperature increases the fraction of molecules with energy ≥ Eₐ, so the rate rises sharply.
Effect of concentration / surface area: Higher reactant concentration or surface area increases collision frequency, raising the rate.
Integrated first-order rate law: ln[A] = ln[A]₀ − kt — a plot of ln[A] vs. time is linear for a first-order reaction.
Chemical equilibrium: The state where forward and reverse reaction rates are equal, so concentrations stay constant (dynamic, not static).
Equilibrium constant (K): Products over reactants, each raised to its coefficient: K > 1 favors products, K < 1 favors reactants.
Reaction quotient (Q): The same ratio as K but at any moment: Q < K shifts forward, Q > K shifts reverse, Q = K is equilibrium.
Le Chatelier's principle: A disturbed equilibrium shifts to partly counteract the change and restore balance.
Effect of pressure on a gas equilibrium: Increasing pressure (decreasing volume) shifts toward the side with fewer moles of gas.
Effect of temperature on K: Treat heat as a reactant/product: raising T raises K for endothermic reactions and lowers K for exothermic ones.
Does a catalyst change K?: No — a catalyst speeds the approach to equilibrium but does not change the equilibrium position or K.
pH: pH = −log[H⁺]; below 7 acidic, 7 neutral, above 7 basic. At 25 °C, pH + pOH = 14.
Strong vs. weak acid: A strong acid ionizes completely (HCl); a weak acid only partially ionizes, described by its Ka.
Kw and its value: The ion-product of water, Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C; Ka·Kb = Kw for a conjugate pair.
Buffer: A solution of a weak acid + its conjugate base that resists pH change when small amounts of acid/base are added.
Henderson–Hasselbalch equation: pH = pKa + log([A⁻]/[HA]); a buffer works best when pH ≈ pKa.
Equivalence point of a titration: Where moles of added titrant exactly neutralize the analyte; an indicator marks it by changing color.
Solubility product (Ksp): The equilibrium constant for dissolving a slightly soluble salt; a smaller Ksp means a less soluble compound.
Common-ion effect: Adding an ion already present in the equilibrium shifts it toward the solid, decreasing solubility.
Conjugate acid–base pair: Two species differing by one H⁺ (e.g., NH₄⁺/NH₃); the stronger the acid, the weaker its conjugate base.
Galvanic (voltaic) cell: An electrochemical cell that uses a spontaneous redox reaction to generate electricity (ΔG < 0, E°cell > 0).
Anode vs. cathode (AN OX, RED CAT): Oxidation occurs at the ANode; reduction occurs at the CAThode.
Standard cell potential: E°cell = E°cathode − E°anode; a positive E°cell means a spontaneous reaction.
Nernst equation: E = E° − (0.0592/n)·log Q at 25 °C — gives the cell potential under non-standard conditions.
ΔG° and cell potential: ΔG° = −nFE°cell (F = 96,485 C/mol); a positive E° gives a negative ΔG° (spontaneous).
Standard hydrogen electrode (SHE): The reference half-cell, assigned E° = 0 V; all other standard reduction potentials are measured against it.
Significant figures: × and ÷: The answer keeps the fewest significant figures of any factor used.
Significant figures: + and −: The answer keeps the fewest decimal places of any value used.
Accuracy vs. precision: Accuracy is closeness to the true value; precision is reproducibility — they are independent.
Systematic vs. random error: A systematic error (uncalibrated balance) biases every reading the same way; random error scatters readings unpredictably.
Titration: A technique to find an unknown concentration by adding a measured titrant until the equivalence point (indicator color change).
Distillation: Separates liquids by boiling point — the lower-boiling component vaporizes and is collected first.
Gravimetric analysis: Quantifies an analyte by the mass of a precipitate that is filtered, dried, and weighed.
Three levels of representation (Visualization): Particulate (atoms/molecules), macroscopic (what you observe), and symbolic (formulas, equations, graphs).
Blank titration: A titration run with all reagents except the analyte, used to correct for impurities in the reagents.
Why calibrate a pH meter with two buffers?: To check linearity — calibrating at two known pH values verifies the meter's response across a range.
Reflux: Heating with a vertical condenser so vapor condenses and returns, allowing a long reaction without losing volatile reactants.
Exact vs. measured numbers: Exact (counted or defined) numbers have unlimited significant figures; measured numbers carry uncertainty in the last digit.

To find us again, just search “Career Employer ACS General Chemistry Exam”

By Career EmployerUpdated Sep 12, 2026

100+ free ACS General Chemistry Exam flashcards — 4 ways to study

ACS General Chemistry Exam Flashcard of the DayNew daily

The classic card. Do you know this one?

Click Study Flashcards above to open the flashcard hub — hundreds of ACS General Chemistry cards you can flip, match, type, or quiz yourself on. Every card is drawn from the exam’s 10 anchoring concepts, so you study exactly what the ACS General Chemistry Exam tests.^[1] Pair them with our free practice test and study guide.

ACS Flashcard Study Modes

Flip mode lets you read a front like Hess’s law and check the definition at your own pace. Type mode reverses that: you read the definition and type the term, which is where a front such as Gibbs free energy stops being vaguely familiar. Match is a timed term-to-definition game for speed, and Quiz builds multiple choice questions from the same 121 cards.

Free ACS General Chemistry flashcards from Career Employer — active recall for the ACS exam

Why Flashcards Work for the ACS Exam

Energy & Thermodynamics is the largest section at 19 cards, and it mixes thermochemistry with electrochemistry: Hess’s law, Enthalpy (ΔH) and Entropy (ΔS) sit next to the Calorimetry equation, the Nernst equation, and cards linking ΔG° and cell potential to a Galvanic (voltaic) cell. Forces, Gases & Solutions follows with 18 cards on gas relationships and intermolecular behavior, including Boyle’s law, the Ideal gas law, Raoult’s law, and solution and phase ideas like Hydrogen bonding, Critical point, and Like dissolves like.

Bonding & Structure carries 17 cards on how atoms connect and what shape results. You drill Ionic bond and Covalent bond definitions, then geometry and electron bookkeeping through VSEPR theory, Steric number, Formal charge, Resonance structures, and the card that asks Why is water bent? Equilibrium & Acids/Bases adds 16 cards covering Chemical equilibrium, Reaction quotient (Q), and Solubility product (Ksp) alongside the acid-base side: pH, Buffer, Kw and its value, and the Common-ion effect.

Atoms & Periodicity holds 15 cards on electron arrangement and trends, from Aufbau principle and Hund’s rule to Effective nuclear charge and Trend: atomic radius, plus quantitative fronts like Energy of a photon and Average atomic mass. Reactions & Stoichiometry contributes 13 cards on the calculation backbone of the exam, including The mole, Avogadro’s number, Limiting reactant, Percent yield, Molarity, and the Dilution equation.

Measurement & Lab rounds out the technique side with 12 cards on Titration, Distillation, Gravimetric analysis, and the error vocabulary tested by Accuracy vs. precision and Systematic vs. random error. Kinetics closes the deck with 11 cards on Rate law, Reaction order, Activation energy (Eₐ), the Arrhenius equation, Rate-determining step, and the card that asks How does a catalyst work?

General chemistry is definition- and formula-dense — the periodic trends, the gas laws, the thermodynamics equations, the equilibrium rules — so spaced flashcards are the most efficient way to keep it all fresh. Used alongside our practice test and study guide, they turn review time into measurable progress.

ACS General Chemistry Flashcards by Concept

The cards are organized by the ACS anchoring concepts. Because the exam spans two full semesters, drill every group — equilibrium, thermodynamics, and kinetics are where most students lose the most points:^[2]

ACS General Chemistry flashcards by concept
Concept group	What the cards cover
Atoms & Periodicity	Quantum numbers · electron configuration · Aufbau/Hund/Pauli · isotopes · periodic trends
Bonding & Structure	Ionic/covalent/polar bonds · Lewis structures · resonance · VSEPR · hybridization · polarity
Forces, Gases & Solutions	Intermolecular forces · gas laws (PV = nRT) · colligative properties · solubility
Reactions & Stoichiometry	The mole · balancing · limiting reactant · percent yield · redox (OIL RIG) · molarity
Thermo & Electrochemistry	Enthalpy · Hess's law · entropy · ΔG = ΔH − TΔS · galvanic cells · Nernst equation
Kinetics & Equilibrium	Rate laws · activation energy · catalysts · Keq · Le Chatelier · pH · buffers · Ksp
Measurement & Lab	Significant figures · accuracy vs. precision · titration · distillation · error types

How to Get the Most Out of These Flashcards

Start with Energy & Thermodynamics. At 19 cards it is the biggest block, and its vocabulary feeds electrochemistry and equilibrium, so locking in Enthalpy (ΔH) and Entropy (ΔS) early pays off twice.
Type-drill the equation cards. Fronts like the Nernst equation and Gibbs free energy reward exact recall, and typing the term from its definition exposes the ones you only recognize by sight.
Use Match for the short trend and law cards. Pairs such as Boyle’s law, Charles’s law, and Trend: atomic radius sort quickly under time pressure and build the recall speed a timed exam needs.
Move to the practice test once Quiz is steady. When multiple choice on Bonding & Structure and Kinetics stops surprising you, test full-length and send the misses back to the study guide.
Keep a repeatable cadence. Two domains per sitting works well for 121 cards: flip a fresh domain, then Type-drill the previous one so review and new material always overlap.

ACS General Chemistry Flashcards FAQ

Hundreds of free ACS General Chemistry flashcards, organized across the exam's 10 anchoring concepts — atoms and periodicity, bonding and structure, intermolecular forces, reactions and stoichiometry, thermodynamics, kinetics, equilibrium, acids/bases, and measurement. They're free with no account required.

References

1.ACS Examinations Institute. “ACS Exams — General Chemistry Exams.” uwm.edu/acs-exams. ↑
2.Holme, T.; Murphy, K.. “Undergraduate Chemistry Anchoring Concepts Content Map I: General Chemistry.” J. Chem. Educ. 2012, 89 (6), 721–723. ↑
3.National Institute of Standards and Technology (NIST). “Periodic Table, Constants, and Chemistry WebBook.” nist.gov. ↑

Career Employer

Career Employer is the ultimate resource to help you get started working the job of your dreams. We cover topics from general career information, career searching, exam preparation with free study materials, career interviewing, and becoming successful in your career of choice.

All Posts

Career Employer’s Editorial Process

Here at Career Employer, we focus a lot on providing factually accurate information that is always up to date. We strive to provide correct information using strict editorial processes, article editing, and fact-checking for all of the information found on our website. We only utilize trustworthy and relevant resources. To find out more, make sure to read our full editorial process page here.