AP Chemistry Key Terms & Vocabulary

242 essential AP Chemistry terms, defined — aligned to the College Board CED.

A
Absolute (standard molar) entropy: The entropy of one mole of a substance in its standard state, always a positive value (unlike enthalpy of formation). Used to compute ΔS° for a reaction.
Acid-base indicator: A substance that changes color between its protonated and deprotonated forms as pH crosses its pKa. Pick an indicator whose pKa is near the equivalence-point pH for accurate endpoint detection.
Acid-base reaction: A reaction involving transfer of one or more protons (H+) between species.
Acid-base titration: Adding a measured volume of titrant to an analyte to reach a known stoichiometric point. Results are summarized in a titration curve of pH versus volume added.
Activation energy (Ea): The minimum energy a collision must have to reach the transition state and react. The energy difference from reactants to the transition state.
Amorphous solid: A solid whose particles lack long-range orderly arrangement (e.g., glass).
Analyte: The substance of unknown amount or concentration being measured in a titration.
Anode: The electrode where oxidation occurs in any electrochemical cell. Electrons leave the anode and travel through the external circuit.
Arrhenius equation: Relates the rate constant to temperature and activation energy, explaining why rate rises with temperature. (AP does not require Arrhenius calculations.)
Atomic mass unit (amu): A unit of mass equal to 1/12 the mass of a carbon-12 atom. The average mass of one particle in amu equals the molar mass in grams.
Atomic radius: A measure of atom size. It decreases across a period (rising Z_eff) and increases down a group (added shells).
Aufbau principle: Electrons fill the lowest-energy available orbitals first when building up the ground-state electron configuration of an atom.
Autoionization of water (Kw): Water self-ionizes into H₃O⁺ and OH⁻. Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C. Kw rises with temperature, so neutral pH shifts below 7 when water is heated.
Average atomic mass: The weighted average of an element's isotopic masses, each weighted by its natural relative abundance. This is the mass shown on the periodic table.
Avogadro's number: N_A = 6.022 x 10^23 mol^-1. The count of particles in one mole; it connects moles of a substance to the actual number of atoms, molecules, or formula units.
B
Balanced equation: An equation with equal numbers of each atom (and equal charge) on both sides, reflecting conservation of mass and charge.
Beer-Lambert law: A = epsilon*b*c. Absorbance is proportional to molar absorptivity, path length, and concentration. With path length fixed, absorbance tracks concentration directly.
Boiling point: The temperature at which vapor pressure equals external pressure. Rises with stronger intermolecular forces.
Bond angle: The angle between two bonds at a central atom. Roughly 180 deg (sp), 120 deg (sp2), and 109.5 deg (sp3), reduced by lone pairs.
Bond dipole: The separation of partial charge across a polar bond. Larger electronegativity differences in single bonds produce larger dipoles.
Bond energy: The energy required to break a bond and separate the atoms. Higher bond orders have larger bond energies.
Bond enthalpy: The average energy needed to break one mole of a particular bond. Breaking bonds absorbs energy; forming bonds releases it.
Bond length: The internuclear distance at minimum potential energy. Higher bond order and smaller atomic cores give shorter bonds.
Bond order: The number of bonding electron pairs between two atoms (1 = single, 2 = double, 3 = triple). Higher order means shorter, stronger bonds.
Bronsted-Lowry acid: A proton (H+) donor.
Bronsted-Lowry base: A proton (H+) acceptor.
Buffer capacity: How much acid or base a buffer can absorb before its pH changes substantially. Raising the concentration of both buffer components (keeping their ratio fixed) increases capacity without changing pH.
Buffer solution: A solution holding large, comparable amounts of a conjugate acid-base pair so it resists pH change. The acid neutralizes added base and the base neutralizes added acid.
C
Calorimeter: The insulated device used to measure heat exchange; in dissolution experiments a temperature rise signals an exothermic process, a drop an endothermic one.
Calorimetry: Measuring heat transfer by tracking temperature change in a known mass, typically using q = mcDeltaT to find heat absorbed or released.
Carboxylic acid: A common class of weak acid containing the –COOH group. It loses a proton to form a resonance-stabilized carboxylate ion, which is a weak base.
Catalyst: A substance that speeds a reaction by providing a lower-activation-energy path or more effective collisions, without being consumed overall.
Cathode: The electrode where reduction occurs in any electrochemical cell. Electrons arrive at the cathode from the external circuit.
Chemical bond: An attractive interaction between atoms strong enough to hold them together, distinct from weaker intermolecular forces. Types: ionic, covalent, metallic.
Chemical change: A transformation that produces new substances with different compositions. Evidence includes heat/light, gas formation, a precipitate, or a color change.
Chemical equation: A symbolic representation of a reaction showing reactants converting to products. It must be balanced so atoms and charge are conserved.
Chromatography: A separation technique exploiting differing strengths of intermolecular interactions between components, a mobile phase, and a stationary phase. Reveals relative polarities.
Collision model: The idea that reactions occur only when reactant particles collide with enough energy and correct orientation to break and form bonds.
Combustion: A redox subclass in which a species reacts with O2. Complete combustion of a hydrocarbon yields CO2 and H2O.
Common-ion effect: The solubility of a salt drops when the solution already contains one of its ions, because the extra ion shifts the dissolution equilibrium backward (Le Chatelier). Ksp itself does not change.
Complete ionic equation: A reaction written with all soluble strong electrolytes shown as their separated aqueous ions.
Concentration cell: A galvanic cell with identical electrodes but different ion concentrations in the two half-cells. Current flows in whichever direction moves the system toward equal concentrations.
Conjugate acid-base pair: Two species that differ by one proton (H⁺), such as CH₃COOH and CH₃COO⁻. The acid donates the proton; its conjugate base accepts it back.
Core electrons: Inner-shell electrons that are not in the outermost shell. They shield valence electrons from the full nuclear charge.
Coulomb's law: F is proportional to (q1*q2)/r^2. The force between charged particles grows with charge magnitude and shrinks with distance squared. It underlies ionization energy, atomic radius, and every periodic trend.
Coupled reaction: Pairing an unfavorable reaction with a favorable one that shares a common intermediate, so their ΔG° values sum to a negative total. This is how cells drive unfavorable steps using ATP.
Covalent bond: A bond formed by sharing valence electrons between atoms, typically between two nonmetals.
Covalent network solid: A solid where atoms are bonded in a continuous covalent network (e.g., diamond, SiO2). Very high melting points; rigid and hard.
Crystalline solid: A solid whose particles are arranged in a regular, repeating 3D lattice.
D
Dalton's law of partial pressures: The total pressure of a gas mixture equals the sum of the partial pressures of its components: P_total = P_A + P_B + P_C + ...
Dipole moment: A measure of net charge separation in a molecule. A molecule is polar overall only if its bond dipoles do not cancel by symmetry.
Dipole-dipole forces: Attractions between permanent dipoles of polar molecules. Add to dispersion forces, so polar molecules attract more strongly than nonpolar ones of similar size.
Dipole-induced dipole forces: Attractions where a polar molecule's dipole induces a temporary dipole in a nearby nonpolar molecule. Always attractive.
Distillation: A separation method based on differences in vapor pressure (and thus intermolecular forces); the more volatile component boils off first.
Dynamic equilibrium: A state where the forward and reverse reactions occur at equal rates, so no net change is observed even though both reactions keep running. Equilibrium means equal rates, not zero rates.
E
Effect of concentration change: Adding a species shifts equilibrium away from that species; removing one shifts toward it. Concentration stresses change Q, not K.
Effect of pressure/volume change: Decreasing volume (raising pressure) shifts equilibrium toward the side with fewer moles of gas; increasing volume shifts toward more moles. No shift if moles of gas are equal.
Effect of temperature change: Temperature is the only stress that changes the value of K. For an exothermic reaction, heating shifts toward reactants (K decreases); for endothermic, heating shifts toward products (K increases).
Effective nuclear charge: The net positive charge a valence electron actually feels after inner core electrons shield part of the nuclear charge. Higher Z_eff pulls electrons in tighter.
Electrolytic cell: An electrochemical cell that uses an external power source to drive a thermodynamically unfavored reaction (ΔG° > 0, E°cell < 0), as in electroplating.
Electromagnetic spectrum: The range of light by wavelength/frequency. Microwave drives molecular rotation, infrared drives vibration, and UV/visible drives electronic transitions.
Electron: A negatively charged subatomic particle occupying shells around the nucleus. Valence electrons drive bonding and reactivity.
Electron affinity: The energy change when a gaseous atom gains an electron. Generally more negative (more favorable) toward the upper right of the periodic table.
Electron configuration: The distribution of an atom's electrons among shells and subshells (e.g., 1s^2 2s^2 2p^6). It explains an element's chemistry and its place on the periodic table.
Electronegativity: An atom's tendency to attract shared bonding electrons. Increases left to right across a period and decreases down a group; it predicts bond polarity.
Elementary reaction: A single-step reaction occurring in one collision. Its rate law can be written directly from the stoichiometry of the colliding particles.
Empirical formula: The chemical formula giving the lowest whole-number ratio of atoms of each element in a compound (e.g., CH2O for glucose).
Endothermic process: A process that absorbs energy from the surroundings; the system's energy increases and surroundings cool. Has a positive enthalpy change.
Endpoint: The point in a titration where the indicator changes color, signaling the equivalence point has (ideally) been reached. A well-chosen indicator makes endpoint and equivalence point nearly coincide.
Energy diagram: A diagram representing the endothermic or exothermic nature of a process by showing the relative energies of initial and final states.
Enthalpy (H): The heat content of a system at constant pressure. Most AP reactions occur at constant pressure, where the enthalpy change equals the heat of reaction.
Enthalpy from bond energies: DeltaH = (bonds broken) - (bonds formed). If more energy is released forming bonds than required to break them, the reaction is exothermic.
Enthalpy of formation method: DeltaH(rxn) = sum DeltaHf(products) - sum DeltaHf(reactants), using tabulated formation values to find a reaction's enthalpy.
Enthalpy of fusion: The energy absorbed per mole to melt a solid; the same magnitude is released when the liquid freezes.
Enthalpy of reaction (DeltaH): The heat released (negative) or absorbed (positive) by a reaction at constant pressure, scalable by the moles of reacting substance.
Enthalpy of vaporization: The energy absorbed per mole to boil a liquid; condensation releases the same magnitude (opposite sign).
Enthalpy- vs entropy-driven: When ΔH° < 0 and ΔS° > 0 a process is favored at all temperatures; when both are positive it needs high T (entropy-driven); when both are negative it needs low T (enthalpy-driven).
Entropy (S): A measure of how dispersed matter and energy are in a system. Entropy rises when particles spread out (solid→liquid→gas), gas volume increases, or more moles of gas form.
Enzyme: A biological catalyst that binds reactants to orient them favorably or lower activation energy, often forming a catalyst-bound intermediate.
Equilibrium constant (K): The value of the reaction quotient at equilibrium. For aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d / [A]^a[B]^b. K depends only on temperature.
Equivalence point: The point in a titration where moles of titrant equal moles of analyte. Strong-strong titrations are neutral here; weak acid titrations are basic and weak base titrations are acidic because the conjugate is present.
Excess reactant: The reactant left over after the limiting reactant is used up.
Exothermic process: A process that releases energy to the surroundings; the system's energy decreases and surroundings warm. Has a negative enthalpy change.
F
Faraday's constant (F): The charge of one mole of electrons, 96,485 C/mol. It converts between moles of electrons and electric charge in electrochemical calculations.
Faraday's laws of electrolysis: Relate charge passed to the amount of substance deposited or consumed at an electrode, via the reaction stoichiometry and number of electrons transferred. Charge q = current × time (I = q/t).
First law of thermodynamics: Energy is conserved: it is neither created nor destroyed in chemical or physical processes, only transferred as heat or work.
First-order reaction: Rate proportional to [A]; a plot of ln[A] vs. time is linear with slope -k. Form: ln[A]t - ln[A]0 = -kt. Radioactive decay is first order.
Formal charge: A bookkeeping charge assigned to an atom in a Lewis structure assuming equal electron sharing. The structure minimizing formal charges is usually the best model.
Free energy and equilibrium (ΔG° and K): ΔG° = −RT ln K, equivalently K = e^(−ΔG°/RT). Negative ΔG° gives K > 1 (products favored); positive ΔG° gives K < 1 (reactants favored).
Free energy of dissolution (ΔG°dissolution): The Gibbs free energy change for a salt dissolving. A negative value means dissolution is thermodynamically favored; it links solubility to enthalpy and entropy of the process.
G
Galvanic (voltaic) cell: An electrochemical cell that runs a thermodynamically favored redox reaction to produce a positive voltage and deliver electrical energy. ΔG° < 0 and E°cell > 0.
Gibbs free energy (ΔG°): The thermodynamic quantity that decides favorability: ΔG° = ΔH° − TΔS°. A negative ΔG° means the process is thermodynamically favored.
H
Half-equivalence point: The point halfway to equivalence in a weak-acid titration, where [HA] = [A⁻]. Here pH = pKa, making it a direct way to read off pKa from a curve.
Half-life: The time for a reactant's concentration to drop by half. For a first-order reaction it is constant: t(1/2) = 0.693/k.
Half-reaction: One half of a redox process showing either the oxidation or the reduction separately; balanced half-reactions combine into the full redox equation.
Heat (q): Energy transferred between bodies due to a temperature difference, occurring as particle collisions move energy from warmer to cooler matter.
Heat transfer equation: q = mcDeltaT, relating heat to mass, specific heat capacity, and temperature change. The core relationship for calorimetry calculations.
Henderson-Hasselbalch equation: pH = pKa + log([A⁻]/[HA]). It gives buffer pH from the pKa and the ratio of conjugate base to acid. When the ratio is 1, pH = pKa.
Hess's law: The enthalpy change of an overall reaction equals the sum of the enthalpy changes of its steps, because energy is conserved regardless of pathway.
Hess's law rules: Reversing a reaction flips the sign of DeltaH; multiplying a reaction by a factor multiplies DeltaH by that factor; adding reactions adds their DeltaH values.
Hybridization: Mixing of valence orbitals to describe electron arrangement around a central atom: sp (180 deg), sp2 (120 deg), sp3 (109.5 deg).
Hydrogen bonding: A strong intermolecular attraction when H bonded to N, O, or F is drawn to a lone pair on N, O, or F in another molecule. Explains water's unusually high boiling point.
Hydronium ion (H₃O⁺): The aqueous form of the proton donated by an acid. Its concentration sets the pH; 'hydrogen ion' (H⁺) is used interchangeably.
I
ICE table: A bookkeeping tool listing Initial, Change, and Equilibrium amounts for each species. It organizes the algebra needed to solve for equilibrium concentrations or K.
Ideal gas law: PV = nRT, relating pressure, volume, moles, and Kelvin temperature of a gas. R = 0.08206 L*atm/(mol*K).
Initial rates method: Comparing reaction rates at the start of trials with different starting concentrations to determine the order with respect to each reactant.
Integrated rate law: An equation relating concentration to time, used to find which linear plot (of [A], ln[A], or 1/[A]) matches the data and thus the reaction order.
Intermolecular forces: Attractions between separate molecules (not within them). They govern boiling point, melting point, and vapor pressure, and are weaker than chemical bonds.
Interstitial alloy: An alloy where small atoms (e.g., carbon in steel) fill gaps between larger metal atoms, making the lattice more rigid and less malleable.
Ion-dipole forces: Attractions between an ion and a polar molecule, such as Na+ surrounded by water. Typically stronger than dipole-dipole forces; key to dissolving salts.
Ionic bond: The electrostatic attraction between oppositely charged ions, typically formed between a metal and a nonmetal via electron transfer.
Ionic solid: A crystalline lattice of cations and anions arranged to maximize attraction and minimize repulsion. Result: high melting points, brittle, conducts only when molten or dissolved.
Ionization energy: The energy required to remove an electron from a gaseous atom or ion. It rises across a period and falls down a group, tracking effective nuclear charge and distance.
Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons, giving them different masses.
K
Ka (acid ionization constant): The equilibrium constant for a weak acid ionizing: Ka = [H₃O⁺][A⁻]/[HA]. A larger Ka means a stronger weak acid that ionizes more.
Ka × Kb relationship: For any conjugate acid-base pair, Ka × Kb = Kw, and equivalently pKa + pKb = 14 at 25°C. A stronger acid always has a weaker conjugate base.
Kb (base ionization constant): The equilibrium constant for a weak base reacting with water: Kb = [OH⁻][HB⁺]/[B]. A larger Kb means a stronger weak base.
Kc: Equilibrium constant written in terms of molar concentrations. Used for reactions in solution or when concentrations are given.
Kinetic control: When a thermodynamically favored reaction does not proceed at a measurable rate, usually because of high activation energy. A slow reaction is not the same as one at equilibrium.
Kinetic energy: Energy of motion. For a particle, KE = (1/2)mv^2; average kinetic energy of gas particles is proportional to Kelvin temperature.
Kinetic molecular theory: A model linking gas macroscopic behavior to particles in constant random motion with negligible volume and forces, and average kinetic energy set by temperature.
Kp: Equilibrium constant written in terms of partial pressures of gases instead of concentrations. Used for gas-phase reactions.
Ksp (solubility-product constant): The equilibrium constant for a salt dissolving, written as the product of dissolved ion concentrations raised to their coefficients. A smaller Ksp means a less soluble salt.
L
Law of definite proportions: Any pure sample of a given compound always contains its constituent elements in the same fixed mass ratio.
Law of mass action: The rule that the equilibrium expression is products over reactants, each raised to its stoichiometric coefficient. It defines how Kc and Kp are constructed from a balanced equation.
Le Chatelier's principle: If a system at equilibrium is stressed (concentration, pressure/volume, or temperature change), it shifts in the direction that partially counteracts the stress. The core tool for predicting qualitative shifts.
Lewis diagram: A representation showing valence electrons as bonding pairs and lone pairs around atoms, built by a set fixed rules to satisfy octets where possible.
Limiting reactant: The reactant that is fully consumed first and therefore caps how much product can form. Compare reactants mole-to-mole, never mass-to-mass.
London dispersion forces: Attractions from temporary, fluctuating dipoles present in all molecules. Strengthen with more electrons, larger electron clouds, and greater contact area.
Lone pair: A valence electron pair not shared in a bond. Lone pairs occupy space and compress bond angles in VSEPR predictions.
M
Magnitude of K: How large K is signals where equilibrium lies. K >> 1 means products dominate (reaction nearly complete); K << 1 means reactants dominate (reaction barely proceeds).
Manipulating K (adding reactions): When reactions are added, the overall K is the product of the individual K values. This lets you build K for a multistep process.
Manipulating K (reversing a reaction): When a reaction is reversed, K becomes 1/K (its reciprocal). Direction of the equation flips the constant.
Manipulating K (scaling coefficients): When all coefficients are multiplied by a factor n, K is raised to the power n. Doubling an equation squares its K.
Mass spectrometry: A technique that separates ions by mass-to-charge ratio. For a single element, the spectrum reveals which isotopes are present and each isotope's relative abundance.
Maxwell-Boltzmann distribution: A curve describing the spread of particle energies in a sample. Raising temperature shifts more particles above the activation energy, speeding the reaction.
Metallic bond: Bonding in which metal atoms release valence electrons into a shared, delocalized 'sea of electrons' surrounding positive metal cores.
Metallic solid: A lattice of positive metal cores in a sea of delocalized electrons. Explains conductivity, malleability, and ductility of metals.
Mixture: Matter containing two or more types of particles whose relative proportions can vary; can be homogeneous or heterogeneous.
Molar absorptivity: The constant epsilon in the Beer-Lambert law describing how strongly a species absorbs light at a specific wavelength.
Molar enthalpy of phase transition: The energy per mole to drive a phase change at constant temperature. The temperature of a pure substance stays constant during the change.
Molar heat capacity: The heat needed to raise the temperature of one mole of a substance by 1 degree C (or 1 K).
Molar mass: The mass of one mole of a substance in g/mol. Numerically equal to the average particle mass in amu, so it converts between grams and moles via n = m/M.
Molar solubility: The number of moles of a salt that dissolve per liter to form a saturated solution. It is calculated from Ksp using an ICE setup and the salt's dissolution stoichiometry.
Molarity: Concentration in moles of solute per liter of solution: M = n_solute / L_solution. The most common lab concentration unit.
Mole: The SI unit for amount of substance. One mole contains exactly Avogadro's number of particles, bridging the particle scale and the lab scale.
Mole fraction: The ratio of moles of one component to total moles in a mixture (X_A = moles A / total moles). Used to find partial pressures.
Mole ratio: The proportion between two substances given by their coefficients in a balanced equation, used to convert moles of one species to moles of another.
Molecular equation: A reaction written with all reactants and products as complete, neutral formula units, ignoring dissociation into ions.
Molecular geometry: The 3D arrangement of atoms in a molecule (linear, trigonal planar, tetrahedral, bent, trigonal pyramidal, etc.), set by bonding and lone pairs around the central atom.
Molecular solid: A solid of discrete covalent molecules held by weak intermolecular forces (e.g., ice). Low melting points; does not conduct electricity.
Molecular structure and acid strength: Acid strength can be predicted from structure: features that stabilize the conjugate base (electronegativity, inductive effects, resonance) make the acid stronger. More resonance/oxygen atoms on an oxyacid increase strength.
Molecularity: The number of particles colliding in an elementary step (unimolecular, bimolecular, etc.). Three-particle collisions are rare.
N
Nernst equation: E = E° − (RT/nF) ln Q, describing how cell potential changes under nonstandard conditions. Qualitatively: Q < 1 raises E above E°, Q > 1 lowers it; E reaches zero at equilibrium (Q = K).
Net ionic equation: A reaction showing only the species that actually change, with spectator ions removed. It captures the real chemistry of the process.
Neutral solution: A solution where [H₃O⁺] = [OH⁻]. At 25°C that is pH 7, but at other temperatures neutral pH differs because Kw changes.
Neutralization reaction: When an acid and base react. Strong acid + strong base gives H⁺ + OH⁻ → H₂O and a neutral solution; the pH afterward depends on which reagent (if any) is in excess.
Neutron: A neutral subatomic particle in the nucleus. Changing neutron count changes the isotope, not the element.
Nonpolar covalent bond: A covalent bond between atoms of similar electronegativity, so bonding electrons are shared roughly equally (e.g., C-H bonds are effectively nonpolar).
O
Octet rule: The tendency of main-group atoms to acquire eight valence electrons (a full s and p shell) through bonding. A guideline with notable exceptions.
Overall order: The sum of the exponents on all reactant concentrations in the rate law.
Oxidation: Loss of electrons, marked by an increase in oxidation number.
Oxidation number: An assigned charge-like value for an atom in a compound, used to track electron transfer and identify what is oxidized or reduced.
Oxidation-reduction reaction: A reaction (redox) involving transfer of one or more electrons, shown by changes in oxidation numbers. Electrons flow from the species oxidized to the species reduced.
P
Partial pressure: The pressure one gas in a mixture would exert alone. P_A = P_total * X_A, where X_A is the mole fraction of A.
Particulate representation of equilibrium: A drawing showing relative numbers of reactant and product particles. At equilibrium those numbers stay constant over time and reflect the magnitude of K.
Percent composition: The percent by mass that each element contributes to a compound. Used to derive empirical formulas from elemental analysis data.
Percent ionization: The fraction of a weak acid or base that ionizes, expressed as a percent. It is computed from Ka (or Kb) and the initial concentration; dilution increases percent ionization.
Periodicity: The recurring pattern of element properties as a function of atomic number, explained by repeating ground-state electron configurations.
pH: A measure of acidity equal to −log[H₃O⁺]. Each unit drop in pH means ten times more hydronium; lower pH is more acidic.
pH and solubility: A salt is more soluble at lower pH when one of its ions is a weak base (like F⁻, OH⁻, or a carbonate). Adding H⁺ removes that ion, pulling more solid into solution; salts with neutral anions like Cl⁻ are unaffected.
Photoelectron spectroscopy (PES): A technique that measures the energy needed to eject electrons from each subshell. Peak position maps to subshell binding energy; peak height is proportional to the number of electrons in that subshell.
Physical change: A change in a substance's properties or phase with no change in its chemical composition. Melting ice or dissolving salt are physical changes.
Pi bond: A bond from side-by-side orbital overlap, found in double and triple bonds. Weaker than a sigma bond and prevents rotation, enabling geometric isomers.
pKa: The negative log of Ka (pKa = −log Ka). A smaller pKa means a stronger acid. At the half-equivalence point of a titration, pH = pKa.
pKb: The negative log of Kb (pKb = −log Kb). A smaller pKb means a stronger base.
Planck's equation: E = h*nu, where h = 6.626 x 10^-34 J*s. A photon's energy is proportional to its frequency; absorbing it raises a species' energy by exactly that amount.
pOH: A measure of basicity equal to −log[OH⁻]. At 25°C, pH + pOH = 14.
Polar covalent bond: A covalent bond between atoms of unequal electronegativity; the more electronegative atom gains a partial negative charge, creating a bond dipole.
Polarizability: How easily a molecule's electron cloud distorts to form a temporary dipole. Increases with electron count and cloud size, strengthening dispersion forces.
Polyprotic acid: An acid with more than one ionizable proton (like H₂SO₄ or H₃PO₄). Its titration curve shows one equivalence-point step per acidic proton, each with its own pKa.
Potential energy curve: A graph of potential energy vs. internuclear distance. Its minimum gives the equilibrium bond length, and its depth gives the bond energy.
Pre-equilibrium approximation: A method used when the first step is fast and reversible (not rate limiting) to derive the rate law in terms of observable reactants.
Precipitate: An insoluble solid that forms and separates out when two solutions are mixed.
Precipitation reaction: A reaction in which mixed aqueous ions form an insoluble or sparingly soluble solid (a precipitate). All Na+, K+, NH4+, and NO3- salts stay soluble.
Proton: A positively charged subatomic particle in the nucleus. The number of protons (atomic number) defines the element.
Pure solids and liquids (in K): Solids and pure liquids are left out of equilibrium expressions because their effective concentrations stay constant regardless of amount. Only gases and aqueous species appear in K or Q.
Pure substance: Matter composed of only one type of atom, molecule, or formula unit, with a fixed composition throughout.
Q
Q vs K comparison: When Q < K the reaction shifts forward (makes products); when Q > K it shifts reverse (makes reactants); when Q = K the system is at equilibrium. This predicts reaction direction.
R
Rate constant (k): The proportionality constant in the rate law. It is temperature dependent and its units depend on the overall reaction order.
Rate law: An equation expressing rate as proportional to reactant concentrations raised to powers: rate = k[A]^m[B]^n. Determined experimentally, not from coefficients.
Rate-determining step: The slowest elementary step in a mechanism; it sets the overall rate law from its molecularity.
Rate-limiting step: Another name for the slowest step that controls the overall reaction rate.
Reaction coordinate: The axis tracking the progress of a reaction from reactants through the transition state to products on an energy profile.
Reaction energy profile: A diagram plotting potential energy along the reaction coordinate, showing activation energy and the overall energy change.
Reaction intermediate: A species produced in one elementary step and consumed in a later one, present only while the reaction proceeds. Detecting it supports a proposed mechanism.
Reaction mechanism: The series of elementary steps that add up to the overall balanced reaction. May include reactants, intermediates, products, and catalysts.
Reaction order: The power to which a reactant's concentration is raised in the rate law. The overall order is the sum of those powers.
Reaction quotient (Q): The ratio of product to reactant concentrations (Qc) or partial pressures (Qp) at any moment, using the same expression form as K. Comparing Q to K tells you which way a reaction will shift.
Reaction rate: How fast reactants convert to products per unit time, tracked as the change in a species' concentration over time. Rates of each species are linked by stoichiometry.
Real gas deviation: Real gases depart from PV = nRT because particles have actual volume (matters at high pressure) and attract one another (matters near condensation).
Reduction: Gain of electrons, marked by a decrease in oxidation number.
Resonance: When two or more equivalent Lewis structures can be drawn, the true structure is a hybrid (average) of them. Required to model molecules like ozone or nitrate accurately.
Reversible reaction: A reaction that can proceed in both directions, written with a double arrow (⇌). Both forward and reverse processes occur simultaneously and can reach equilibrium.
S
Salt bridge: A connection that lets ions flow between half-cells to keep each side electrically neutral. Without it, charge builds up and current stops.
Second-order reaction: Rate proportional to [A]^2; a plot of 1/[A] vs. time is linear with slope k. Form: 1/[A]t - 1/[A]0 = kt.
Shielding: The reduction in attraction a valence electron feels from the nucleus because core electrons repel it and block part of the nuclear charge.
Sigma bond: A bond from head-on orbital overlap along the internuclear axis. Stronger than a pi bond and allows free rotation.
Solubility: The extent to which a solute dissolves in a solvent. 'Like dissolves like': substances with similar intermolecular forces tend to be miscible.
Solubility equilibrium: The dynamic balance between a dissolving solid and its dissolved ions in a saturated solution. Treated like any reversible reaction.
Solute: The substance dissolved in a solution, usually present in the smaller amount.
Solution: A homogeneous mixture whose macroscopic properties are uniform throughout. Can be solid, liquid, or gas.
Solvent: The substance that does the dissolving, usually present in the larger amount.
Specific heat capacity (c): The heat needed to raise 1 gram of a substance by 1 degree C (or 1 K). Higher c means a smaller temperature change for the same heat.
Spectator ion: An ion present in solution that appears unchanged on both sides of a reaction and is cancelled out in the net ionic equation.
Spectrophotometer: An instrument that measures light absorbance by a solution, usually set to the wavelength of maximum absorbance for greatest sensitivity.
Standard cell potential (E°cell): The voltage of a cell under standard conditions, found from the standard reduction potentials of the two half-reactions. A positive E°cell marks a thermodynamically favored reaction.
Standard enthalpy of formation (DeltaHf): The enthalpy change to form one mole of a compound from its elements in their standard states.
Standard entropy change (ΔS°): The entropy change for a reaction, calculated as ΣS°(products) − ΣS°(reactants) from absolute molar entropies. A positive value means the products are more disordered.
Standard Gibbs free energy of formation (ΔG°f): The free energy change to form one mole of a compound from its elements in standard states. ΔG°reaction = ΣΔG°f(products) − ΣΔG°f(reactants).
Standard reduction potential: The tendency of a half-reaction to gain electrons, measured in volts relative to the standard hydrogen electrode. Used to rank oxidizing/reducing strength and compute E°cell.
Stoichiometry: The use of balanced-equation coefficients as mole ratios to relate amounts of reactants and products. The basis for predicting product yield.
Strong acid: An acid that ionizes completely in water (HCl, HBr, HI, HClO₄, H₂SO₄, HNO₃), so [H₃O⁺] equals the initial acid concentration and pH is found directly.
Strong base: A base that fully dissociates in water (Group I and II hydroxides). [OH⁻] equals the initial concentration for Group I, or double it for Group II.
Substitutional alloy: An alloy where atoms of comparable size replace host atoms in the lattice (e.g., zinc substituting for copper in brass).
Surface catalysis: Catalysis where a reactant binds to (or bonds with) a solid surface, creating new bound intermediates and elementary steps.
T
Thermal equilibrium: The state reached when bodies in contact share the same average kinetic energy, and therefore the same temperature, so net heat transfer stops.
Thermodynamically favored: A process with ΔG° < 0, meaning products are favored at equilibrium (K > 1). Preferred over the word 'spontaneous,' which wrongly suggests 'fast' or 'sudden.'
Thermodynamics: The study of energy and its transfer as heat and work, used to predict the direction and energetics of physical and chemical changes.
Titrant: The solution of known concentration added during a titration; it reacts specifically and quantitatively with the analyte.
Titration: A technique that adds a titrant of known concentration to an analyte until they react completely, used to find the analyte's amount or concentration.
Titration curve: A plot of pH against volume of titrant. Its shape reveals equivalence points, buffer regions, and (for weak acids) the pKa at the half-equivalence point.
Transition state: The high-energy arrangement at the peak of the energy profile, between reactants and products, where bonds are breaking and forming.
V
Valence electrons: The outermost-shell electrons. They determine an element's bonding behavior, typical ionic charge, and reactivity.
Vapor pressure: The pressure of vapor in equilibrium with its liquid. Lower when intermolecular forces are strong, since fewer molecules can escape to the gas phase.
VSEPR theory: Valence Shell Electron Pair Repulsion: electron pairs around a central atom arrange to minimize repulsion, which predicts molecular geometry and bond angles.
W
Wavelength-frequency relationship: c = lambda * nu. Wavelength and frequency are inversely related, with c = 3.00 x 10^8 m/s the speed of light.
Weak acid: An acid that only partially ionizes, leaving most molecules intact. Its [H₃O⁺] is far below the initial concentration, so pH must be found from Ka and an equilibrium calculation.
Weak base: A base that only partially ionizes in water (like ammonia), producing far less OH⁻ than its initial concentration. pH is found from Kb.
Z
Zeroth-order reaction: Rate independent of concentration; a plot of [A] vs. time is linear with slope -k. Form: [A]t - [A]0 = -kt.
Δ
ΔG° = ΔH° − TΔS°: The master equation tying favorability to enthalpy, entropy, and temperature. When ΔH° and ΔS° have opposing signs, temperature decides whether ΔG° is negative.
ΔG° = −nFE°: The equation linking free energy to cell potential, where n is moles of electrons and F is Faraday's constant. A positive E° gives a negative ΔG° (favored cell).

Where these terms show up

Every AP Chemistry term below belongs to one of the units on the exam. Each unit page has sample questions that put the vocabulary to work:

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