AP Chemistry – Comprehensive Vocabulary Review

A comprehensive set of vocabulary flashcards covering major concepts, terms, and principles from Units 1–8 of the AP Chemistry video lecture notes.

Atomic Number

The number of protons in the nucleus of an atom; identifies the element.

Mass Number

The total number of protons and neutrons in an atom’s nucleus.

Isotope

Atoms of the same element that contain different numbers of neutrons.

Mass Spectrometry

Experimental technique used to determine the masses and relative abundances of isotopes.

Mole

A quantity of substance containing Avogadro’s number (6.022 × 10²³) of representative particles.

Avogadro’s Number

6.022 × 10²³ particles per mole.

Molarity (M)

Moles of solute per liter of solution.

Percent Composition

The mass percent of each element in a compound.

Empirical Formula

The simplest whole-number ratio of atoms in a compound.

Molecular Formula

The actual number of atoms of each element in a molecule.

Coulomb’s Law

Describes the electrostatic force between charged particles; force increases with charge magnitude and decreases with distance.

Ionization Energy

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

Photoelectron Spectroscopy (PES)

Technique that measures the binding energies of electrons in atoms or molecules.

Subshell

Set of orbitals within an energy level that have the same angular momentum (s, p, d, f).

Orbital

Region of space where there is a high probability of finding an electron.

Aufbau Principle

Electrons occupy orbitals of lowest available energy first.

Pauli Exclusion Principle

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

Hund’s Rule

Electrons fill degenerate orbitals singly before pairing up, with parallel spins.

Valence Electrons

Electrons in the outermost shell; participate in bonding.

Cation

Positively charged ion formed by loss of electrons.

Anion

Negatively charged ion formed by gain of electrons.

Atomic Radius

Half the distance between nuclei of two identical atoms; decreases across a period, increases down a group.

Electronegativity

Ability of an atom in a covalent bond to attract shared electrons.

Electron Affinity

Energy change when a gaseous atom gains an electron.

Octet Rule

Atoms tend to gain, lose, or share electrons to attain eight valence electrons.

Bond Energy

Energy required to break one mole of a given bond in the gas phase.

Ionic Bond

Electrostatic attraction between cations (metals) and anions (non-metals).

Lattice Energy

Energy released when gaseous ions form an ionic solid; measure of ionic bond strength.

Metallic Bond

Attraction between metal cations and a “sea” of delocalized electrons.

Alloy

Mixture of metals; can be interstitial (different radii) or substitutional (similar radii).

Covalent Bond

Bond formed when two non-metals share electrons.

Sigma (σ) Bond

Single covalent bond formed by head-on orbital overlap.

Pi (π) Bond

Bond formed by sideways overlap of p orbitals; present in double and triple bonds.

Bond Order

Number of bonding electron pairs between two atoms (1 = single, 2 = double, 3 = triple).

Internuclear Distance

Separation between nuclei of bonded atoms at minimum potential energy (bond length).

Network Covalent Solid

Substance in which atoms are linked by an extensive lattice of covalent bonds (e.g., diamond, quartz).

Electrolyte

Substance that produces ions in solution and conducts electricity.

Lewis Structure

Diagram showing bonding and lone-pair electrons in a molecule or ion.

Resonance

Multiple valid Lewis structures differing only in electron placement.

Formal Charge

Hypothetical charge assigned to atoms in a Lewis structure; helps identify most stable resonance form.

Incomplete Octet

Stable molecule in which the central atom has fewer than eight electrons (e.g., B, Be compounds).

Expanded Octet

Molecule where the central atom has more than eight electrons (period 3 or higher elements).

VSEPR Theory

Predicts molecular geometry based on repulsion between electron groups around a central atom.

Linear Geometry

Two electron groups; bond angle 180°.

Trigonal Planar

Three electron groups; bond angle 120°.

Bent (Angular)

Three or four electron groups with lone pairs; bond angle <120° or <109.5°.

Tetrahedral

Four electron groups; bond angle 109.5°.

Trigonal Pyramidal

Four electron groups with one lone pair; bond angle ≈107°.

Trigonal Bipyramidal

Five electron groups; bond angles 90° & 120°.

Seesaw

Trigonal bipyramidal with one lone pair.

T-Shaped

Trigonal bipyramidal with two lone pairs.

Octahedral

Six electron groups; bond angles 90°.

Square Pyramidal

Octahedral with one lone pair.

Square Planar

Octahedral with two lone pairs; bond angles 90°.

Polarity

Unequal sharing of electrons in a bond or molecule producing partial charges.

Dipole Moment

Measure of bond or molecular polarity; product of charge and distance.

Dipole-Dipole Forces

Attractions between polar molecules’ partial charges.

Hydrogen Bond

Strong dipole interaction where H is bonded to N, O, or F and attracted to a lone pair on another molecule.

London Dispersion Forces

Weak attractions from temporary dipoles; present in all molecules.

Polarizability

Ease with which electron cloud can be distorted, enhancing dispersion forces.

Melting Point

Temperature at which a solid becomes a liquid.

Boiling Point

Temperature at which vapor pressure equals external pressure.

Vapor Pressure

Pressure exerted by a vapor in equilibrium with its liquid; decreases with stronger IMF.

Chromatography

Separation technique using differing affinities for stationary and mobile phases.

Distillation

Separation of liquid mixtures based on different boiling points.

Kinetic Molecular Theory

Model describing gases as particles in constant, random motion with negligible volume and elastic collisions.

Ideal Gas Law

PV = nRT; relates pressure, volume, temperature, and moles of an ideal gas.

Maxwell-Boltzmann Distribution

Graph showing range of molecular speeds in a gas sample.

Effusion

Process by which gas molecules escape through a tiny opening.

Dalton’s Law of Partial Pressures

Total pressure equals sum of individual gas partial pressures in a mixture.

Electromagnetic Spectrum

Range of all types of EM radiation, characterized by wavelength and frequency.

Planck’s Constant (h)

6.626 × 10⁻³⁴ J·s; relates energy and frequency of photons.

Beer-Lambert Law

A = εbc; absorbance proportional to molar absorptivity, path length, and concentration.

Synthesis Reaction

Two or more reactants combine to form one product.

Decomposition Reaction

Single compound breaks into two or more products.

Acid-Base (Neutralization) Reaction

Acid reacts with base to produce salt and water.

Redox Reaction

Reaction involving transfer of electrons; changes oxidation states.

Combustion Reaction

Hydrocarbon reacts with O₂ to produce CO₂ and H₂O, releasing heat.

Precipitation Reaction

Mixing of solutions forms an insoluble solid (precipitate).

Net Ionic Equation

Equation showing only species that actually change during reaction.

Spectator Ion

Ion that remains unchanged during a reaction; omitted from net ionic equation.

Solubility Rule (Nitrates)

All nitrate (NO₃⁻), alkali metal, and ammonium salts are soluble in water.

Limiting Reactant

Reactant that is completely consumed first, limiting product formation.

Percent Yield

Actual yield divided by theoretical yield, multiplied by 100%.

Combustion Analysis

Technique determining empirical formula by measuring CO₂ and H₂O produced on burning.

Gravimetric Analysis

Quantitative determination via mass of a precipitate.

Oxidation State

Apparent charge of an atom in a compound, based on electron bookkeeping rules.

Rate Law

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

Rate Constant (k)

Proportionality constant in a rate law; specific to reaction and temperature.

Reaction Order

Exponent of concentration term in rate law indicating dependence on that reactant.

Integrated Rate Law

Mathematical expression linking concentration and time for a given reaction order.

Half-Life (t½)

Time required for half of a reactant to be consumed.

Collision Theory

Reactions occur when molecules collide with sufficient energy and proper orientation.

Activation Energy (Ea)

Minimum energy required to initiate a reaction.

Reaction Mechanism

Sequence of elementary steps that describes the pathway from reactants to products.

Elementary Step

Single molecular event in a mechanism with its own rate law.

Intermediate

Species produced in one elementary step and consumed in another.

Rate-Determining Step

Slowest elementary step controlling overall reaction rate.

Catalyst

Substance that increases reaction rate by providing an alternative pathway with lower Ea, and is regenerated.

First Law of Thermodynamics

Energy is conserved; ΔE = q + w.