Catalonia PAU Chemistry key terms

The arrangement of electrons among the orbitals of an atom, written following the Aufbau principle, Hund's rule and the Pauli exclusion principle; it explains chemical behaviour and position in the periodic table.

The rule that electrons occupy the lowest-energy orbitals available first, so 1s fills before 2s, and 2p before 3s, giving the ground-state configuration of an atom.

When degenerate orbitals are available, electrons occupy them singly with parallel spins before any pairing occurs, minimising electron-electron repulsion in the ground state.

The regular variation of atomic radius, ionisation energy, electron affinity and electronegativity across a period or down a group, caused by changing nuclear charge and shielding.

A measure of how strongly an atom in a bond attracts shared electrons; it increases across a period and decreases down a group and predicts bond polarity.

The minimum energy required to remove one mole of electrons from one mole of gaseous atoms or ions; successive values reveal shell structure and electron pairing.

An atom of the same element with the same number of protons but a different number of neutrons, so it has a different mass number but nearly identical chemical properties.

The weighted average mass of the naturally occurring isotopes of an element on the carbon-12 scale, calculated from isotopic abundances.

A diagram showing valence electrons as dots or lines to represent bonding pairs and lone pairs, used to predict molecular shape and polarity.

A model that predicts molecular geometry by arranging electron domains around a central atom as far apart as possible, distinguishing bonding pairs from lone pairs.

The mixing of atomic orbitals of similar energy to form equivalent hybrid orbitals (sp, sp2, sp3) that explain observed bond angles and molecular shapes.

The description of a molecule by two or more equivalent Lewis structures whose average represents the actual delocalised electron distribution, as in ozone or the carbonate ion.

A vector quantity measuring the separation of positive and negative charge in a polar bond or molecule; its sum determines whether a molecule is overall polar.

A strong intermolecular attraction between a hydrogen atom bonded to N, O or F and a lone pair on a nearby electronegative atom, raising boiling points and affecting solubility.

Attractive forces between molecules, including London dispersion, dipole-dipole and hydrogen bonding, which are much weaker than covalent bonds but govern physical properties.

A model gas whose particles have negligible volume and no intermolecular forces, obeying PV = nRT; real gases approach this behaviour at low pressure and high temperature.

The pressure a component gas would exert if it alone occupied the container volume; by Dalton's law the total pressure is the sum of partial pressures.

A physical property of a solution that depends on the number of dissolved particles, not their identity, such as boiling-point elevation, freezing-point depression and osmotic pressure.

The maximum amount of solute that dissolves in a given amount of solvent at a specified temperature, governed by like-dissolves-like interactions and by the solubility product for sparingly soluble salts.

When a system at equilibrium is disturbed by a change in concentration, pressure or temperature, it shifts in the direction that partially counteracts the change.

The ratio of products to reactants raised to their stoichiometric coefficients, using equilibrium concentrations; its value depends only on temperature for a given reaction.

The equilibrium constant expressed in partial pressures instead of concentrations, related to Kc by Kp = Kc(RT) raised to the change in moles of gas.

The equilibrium constant Ksp for the dissolution of a sparingly soluble ionic solid, equal to the product of the ion concentrations each raised to its stoichiometric coefficient.

The decrease in solubility of a sparingly soluble salt caused by adding a soluble salt that shares one of its ions, shifting the dissolution equilibrium to the left.

The negative base-10 logarithm of the hydrogen ion concentration in mol per litre; it measures acidity, with 7 neutral at 25 degrees Celsius.

The equilibrium constant Ka for the ionisation of a weak acid in water; a larger Ka (smaller pKa) indicates a stronger acid.

A solution of a weak acid and its conjugate base (or a weak base and its conjugate acid) that resists pH change on adding small amounts of acid or base.

The relation pH = pKa + log([conjugate base]/[weak acid]) used to calculate the pH of a buffer or the position of a titration midpoint.

A volumetric procedure in which a standard solution is added until the equivalence point is reached, allowing the concentration of an analyte to be determined from a pH or indicator curve.

The potential of a half-cell measured against the standard hydrogen electrode under standard conditions; combining two values gives the cell emf of a galvanic cell.

An electrochemical cell that converts chemical energy into electrical energy through a spontaneous redox reaction, with oxidation at the anode and reduction at the cathode.

An equation relating cell potential to standard potential and reaction quotient, E = Eo - (RT/nF) ln Q, showing how concentration changes affect voltage.

The use of an external electrical current to drive a non-spontaneous redox reaction in an electrolytic cell, producing chemical change at the electrodes.

An equation expressing reaction rate as a function of reactant concentrations, with exponents that define the reaction order and must be found experimentally.

The proportionality constant k in a rate law; it depends on temperature, as described by the Arrhenius equation, and on the presence of a catalyst.

The minimum energy that colliding particles must possess to react, corresponding to the energy barrier between reactants and the transition state.

A substance that increases reaction rate by providing an alternative pathway with lower activation energy, without being consumed in the overall reaction.

The heat exchanged at constant pressure, denoted H; the enthalpy change of a reaction equals the heat absorbed by the system under those conditions.

The enthalpy change of a reaction is independent of the route taken, so it can be calculated by adding the enthalpy changes of any sequence of steps.

The average energy required to break one mole of a particular bond in the gaseous state; it can be used to estimate reaction enthalpies.

The thermodynamic function G = H - TS; a negative change in Gibbs energy means a process is spontaneous at constant temperature and pressure.

A measure of the dispersal of energy or disorder of a system; it increases with temperature and with the number of accessible microstates, favouring spontaneity.

An atom or group of atoms that gives an organic compound its characteristic chemical properties, such as hydroxyl, carbonyl, carboxyl or amino groups.

The existence of compounds with the same molecular formula but different arrangements of atoms, giving structural isomers or stereoisomers with different properties.

The property of a molecule that is not superimposable on its mirror image, usually because a carbon atom is bonded to four different groups, giving enantiomers.

A reaction in which two molecules combine to form one product, typical of alkenes and alkynes where a pi bond opens to add atoms across the double or triple bond.

The joining of many small monomer molecules into a large chain, either by addition across double bonds or by condensation with loss of a small molecule such as water.