AP Chemistry key terms

The electrostatic force between two point charges equals a constant times the product of the charges divided by the square of their separation; it explains why attractions strengthen as charges grow and distances shrink.

The net positive charge an outer electron experiences after inner electrons screen part of the nucleus's pull; it increases across a period and helps explain trends in atomic size and ionization energy.

The minimum energy needed to remove one electron from a gaseous atom or ion, always an endothermic quantity; successive removals cost more energy and reveal an element's electron shell structure.

A plot of ion abundance versus mass-to-charge ratio produced when a sample is ionized and separated; the peaks reveal isotopic masses and relative abundances, allowing average atomic mass to be calculated.

A technique that bombards atoms with X-rays and measures the kinetic energies of ejected electrons; the resulting binding-energy peaks show how many electrons occupy each subshell.

A measure of how strongly an atom in a bond attracts shared electrons; differences in this value between bonded atoms produce nonpolar, polar covalent, or ionic bonding and create partial charges.

A vector quantity equal to charge magnitude times separation distance that measures a molecule's overall polarity; it points from the positive to the negative end and depends on both bond polarities and molecular geometry.

A diagram showing valence electrons as dots and shared pairs as lines, drawn to satisfy the octet rule where possible and used to predict bonding, lone pairs, and formal charges in a species.

The model that predicts molecular shape by arranging electron domains around a central atom as far apart as possible; lone pairs occupy more space than bonds, compressing bond angles.

The mixing of an atom's valence orbitals into equivalent new orbitals that explain observed geometries; sp, sp2, and sp3 sets correspond to linear, trigonal planar, and tetrahedral electron-domain arrangements.

The description of a species as a weighted average of two or more valid Lewis structures that differ only in electron placement; the real structure is more stable than any single contributing structure.

A sigma bond forms from head-on orbital overlap along the internuclear axis and permits rotation; a pi bond forms from side-by-side overlap above and below that axis and blocks rotation, making multiple bonds rigid.

Relatively weak attractions between separate molecules or ions, including dispersion, dipole–dipole, and hydrogen bonding; they determine boiling points, vapor pressure, viscosity, and the physical states of substances.

An unusually strong dipole attraction between a hydrogen bonded to fluorine, oxygen, or nitrogen and a lone pair on a nearby electronegative atom; it explains water's high boiling point and ice's low density.

Temporary, instantaneous dipole attractions present in all substances that grow stronger with more electrons and greater polarizability; they dominate the intermolecular forces of nonpolar molecules and noble gases.

The pressure exerted by a gas in equilibrium with its own liquid or solid in a closed container; it rises with temperature and with weaker intermolecular forces, and it equals atmospheric pressure at the boiling point.

The absorbance of light by a solution is proportional to the concentration of the absorbing species and the path length, expressed as A = εbc; it underpins quantitative spectrophotometric analysis.

The reactant that is completely consumed first and therefore caps the amount of product formed; identifying it requires comparing the mole ratios actually available with those in the balanced equation.

A balanced equation that shows only the species that actually change during a reaction, omitting spectator ions; it is obtained by canceling ions that appear unchanged on both sides of the complete ionic equation.

A bookkeeping charge assigned to an atom in a compound by specific rules, such as giving shared electrons to the more electronegative atom; changes in it identify oxidation and reduction.

An equation relating reaction rate to reactant concentrations raised to experimentally determined powers; the exponents give the reaction order and are not generally taken from the balanced equation's coefficients.

The minimum kinetic energy colliding particles must possess for a reaction to occur; it sets the height of the energy barrier and appears in the exponential term of the Arrhenius equation.

A substance that increases reaction rate by providing an alternative pathway with lower activation energy while remaining chemically unchanged overall; it does not alter the equilibrium constant or the thermodynamics of the reaction.

The time required for half of a reactant's initial amount to be consumed; for a first-order reaction it is constant and equals 0.693 divided by the rate constant.

A proposed sequence of elementary steps by which an overall reaction occurs; the slowest step controls the rate, and the mechanism must sum to the balanced overall equation.

The heat exchanged at constant pressure when a reaction proceeds as written; it is negative for exothermic and positive for endothermic processes and can be summed over steps.

Because enthalpy is a state function, the enthalpy change of an overall reaction equals the sum of the enthalpy changes of any series of steps into which it is divided, including reversed or scaled steps.

The experimental measurement of heat exchanged using a device that isolates the system, usually calculated from mass, specific heat capacity, and temperature change via q = mcΔT.

The ratio of product concentrations to reactant concentrations, each raised to its stoichiometric coefficient, evaluated at equilibrium; its value at a given temperature indicates whether products or reactants are favored.

A term with the same form as the equilibrium constant but calculated from concentrations at any moment; comparing it with the equilibrium constant predicts the direction the reaction must shift.

If a system at equilibrium is disturbed by a change in concentration, pressure, or temperature, it shifts in the direction that partially offsets that disturbance and restores equilibrium.

The equilibrium constant for the dissolution of a sparingly soluble ionic solid, written as the product of the dissolved ion concentrations raised to their stoichiometric powers; it predicts precipitation when the ion product exceeds it.

A logarithmic measure of hydrogen ion activity in a solution, defined as the negative base-ten logarithm of the hydrogen ion concentration; values below 7 are acidic and above 7 basic at 25 °C.

A species that donates a proton to another species in a reaction; its conjugate base is what remains after donation, and water can act as either acid or base.

A solution containing a weak acid and its conjugate base, or a weak base and its conjugate acid, that resists pH change when small amounts of strong acid or base are added.

A thermodynamic measure of the dispersal of energy and matter among available microstates; it increases for the universe in any spontaneous process and generally grows with temperature, volume, and particle number.

A state function defined as G = H − TS; its change at constant temperature and pressure determines spontaneity, being negative for spontaneous, zero at equilibrium, and positive for nonspontaneous processes.

The voltage of an electrochemical cell when all dissolved species are at 1 M, gases at 1 bar, and the temperature is 298 K; a positive value indicates a spontaneous galvanic reaction.

The relationship that corrects a cell potential for nonstandard concentrations, E = E° − (RT/nF)lnQ; it shows how the voltage falls as products accumulate toward equilibrium.