AP Chemistry Study Guide – Core Vocabulary

A comprehensive set of vocabulary flashcards covering fundamental terms and concepts from the AP Chemistry study guide, suitable for exam review.

Mole

The SI unit for amount of substance, equal to 6.022 × 10²³ particles.

Avogadro’s Number

6.022 × 10²³, the number of entities in one mole of substance.

Molar Mass

Mass of one mole of a substance, expressed in g mol⁻¹.

Mass Spectrometry

Technique that separates ions by mass-to-charge ratio to determine isotopic abundance.

Pure Substance

Matter with uniform composition that cannot be separated by physical means; includes elements and compounds.

Element

Pure substance consisting of only one kind of atom.

Compound

Pure substance composed of two or more different atoms chemically bonded.

Law of Definite Proportions

A compound always contains the same elements in the same ratio by mass.

Mixture

Physical combination of substances with variable composition.

Homogeneous Mixture

Mixture with uniform composition throughout (solution).

Heterogeneous Mixture

Mixture with non-uniform composition and distinct phases.

Proton

Positively charged subatomic particle found in the nucleus.

Neutron

Uncharged subatomic particle in the nucleus with mass ~1 amu.

Electron

Negatively charged subatomic particle occupying orbitals around the nucleus.

Atomic Number (Z)

Number of protons in an atom’s nucleus; defines the element.

Mass Number (A)

Total number of protons plus neutrons in the nucleus.

Isotope

Atoms of the same element with different numbers of neutrons.

Atomic Mass

Weighted average mass of all naturally occurring isotopes of an element.

Ion

Atom or group of atoms with a net electric charge.

Cation

Positively charged ion formed by loss of electrons.

Anion

Negatively charged ion formed by gain of electrons.

Electron Shell (n)

Principal energy level where electrons reside around a nucleus.

Subshell (l)

Division of a shell (s, p, d, f) with specific shape and energy.

Orbital

Region in a subshell with a maximum of two electrons having opposite spins.

Aufbau Principle

Electrons occupy lowest-energy orbitals available first.

Hund’s Rule

Electrons fill degenerate orbitals singly with parallel spins before pairing.

Pauli Exclusion Principle

No two electrons in an atom can have the same set of four quantum numbers.

Paramagnetic

Substance with unpaired electrons attracted to a magnetic field.

Diamagnetic

Substance with all electrons paired; weakly repelled by a magnetic field.

Ionization Energy

Energy required to remove an electron from a gaseous atom or ion.

Coulomb’s Law

Electrostatic force is proportional to product of charges divided by distance squared.

Photoelectron Spectroscopy (PES)

Technique that measures binding energies of electrons to deduce electron configuration.

Periodic Trend

Predictable variation of properties (radius, IE, EN) across periods and groups.

Alkali Metals

Group 1 elements; soft, highly reactive, form +1 ions.

Alkaline Earth Metals

Group 2 elements; reactive, form +2 ions, shiny metals.

Transition Metals

d-block elements capable of multiple oxidation states and colored compounds.

Chalcogens

Group 16 elements, typically form –2 ions (e.g., O, S, Se).

Halogens

Group 17 elements; highly reactive nonmetals forming –1 ions or diatomic molecules.

Noble Gases

Group 18 elements; inert gases with full valence shells.

Ionic Bond

Electrostatic attraction between oppositely charged ions formed by electron transfer.

Covalent Bond

Chemical bond formed by sharing electron pairs between atoms.

Metallic Bond

Attraction between metal cations and a sea of delocalized electrons.

Electronegativity

Atom’s ability to attract shared electrons in a bond.

Dipole

Separation of charge in a molecule producing partial positive and negative ends.

London Dispersion Forces

Weak, temporary attractions from instantaneous dipoles in all molecules.

Hydrogen Bond

Strong dipole-dipole attraction between H bonded to N, O, or F and another electronegative atom.

Ideal Gas Law

Relationship PV = nRT describing an ideal gas.

Boyle’s Law

At constant T and n, P is inversely proportional to V (P₁V₁ = P₂V₂).

Charles’s Law

At constant P and n, V is proportional to T (V₁/T₁ = V₂/T₂).

Gay-Lussac’s Law

At constant V and n, P is proportional to T (P₁/T₁ = P₂/T₂).

Avogadro’s Law

At constant P and T, V is proportional to moles (V₁/n₁ = V₂/n₂).

Dalton’s Law of Partial Pressures

Total pressure equals sum of partial pressures of component gases.

Kinetic Molecular Theory (KMT)

Model describing gas particles in constant random motion with energy proportional to temperature.

Molarity (M)

Moles of solute per liter of solution.

Molality (m)

Moles of solute per kilogram of solvent.

Chromatography

Separation technique using differential affinity between mobile and stationary phases.

Solubility

Maximum amount of solute that dissolves in a given solvent at specific conditions.

Henry’s Law

Solubility of a gas in a liquid is proportional to its partial pressure (p = kH C).

Beer-Lambert Law

Absorbance A = ε b c; relates light absorption to concentration and path length.

Endothermic Process

Reaction or change that absorbs heat from surroundings (ΔH > 0).

Exothermic Process

Reaction or change that releases heat to surroundings (ΔH < 0).

Calorimetry

Measurement of heat changes using insulated devices like coffee-cup or bomb calorimeters.

Hess’s Law

Total enthalpy change equals sum of enthalpies of individual steps in a reaction pathway.

Equilibrium Constant (K)

Ratio of product to reactant concentrations at equilibrium, each raised to stoichiometric coefficients.

Le Châtelier’s Principle

A system at equilibrium shifts to counteract applied stress (changes in T, P, or concentration).

Solubility Product Constant (Ksp)

Equilibrium constant for the dissolution of a sparingly soluble ionic compound.

Buffer

Solution that resists pH change, containing a weak acid/base and its conjugate.

Henderson–Hasselbalch Equation

pH = pKa + log([A⁻]/[HA]); estimates pH of a buffer.

Gibbs Free Energy (ΔG)

Energy available to do work; ΔG < 0 indicates a spontaneous process.

Entropy (S)

Measure of molecular disorder or energy dispersal in a system.

Galvanic (Voltaic) Cell

Electrochemical cell that converts spontaneous redox energy into electrical work.

Electrolytic Cell

Cell that uses electrical energy to drive a non-spontaneous redox reaction.

Standard Reduction Potential (E°)

Voltage associated with a reduction half-reaction under standard conditions relative to SHE.

Arrhenius Acid

Substance that produces H⁺ in aqueous solution.

Arrhenius Base

Substance that produces OH⁻ in aqueous solution.

Brønsted-Lowry Acid

Proton (H⁺) donor in a chemical reaction.

Brønsted-Lowry Base

Proton (H⁺) acceptor in a chemical reaction.

pKa

Negative log of an acid’s Ka; lower pKa means stronger acid.

Activation Energy (Ea)

Minimum energy required for reactants to form products.

Catalyst

Substance that lowers activation energy and increases reaction rate without being consumed.

Rate Law

Equation expressing reaction rate as function of reactant concentrations and rate constant.

Reaction Order

Exponent of a reactant concentration in the rate law; indicates effect on rate.

Half-Life (t½)

Time required for half the quantity of a reactant or radioactive isotope to disappear.

Rate-Determining Step

Slowest elementary step in a reaction mechanism controlling overall rate.

Buffer Capacity

Amount of acid or base a buffer can neutralize before significant pH change.

Nernst Equation

E = E° – (RT/nF) ln Q; relates cell potential to reaction quotient under non-standard conditions.

Salt Bridge

Ionic pathway in electrochemical cells that maintains electrical neutrality by allowing ion flow.

Lattice Energy

Energy released when one mole of an ionic solid forms from gaseous ions; measure of bond strength.