STPM Chemistry 962 key terms

The weighted average mass of the atoms of an element on a scale where one atom of carbon-12 has a mass of exactly 12 units, taking natural isotopic abundance into account.

The sum of the relative atomic masses of all atoms shown in a formula unit of a molecular substance; for ionic compounds the equivalent sum is called the relative formula mass.

The amount of substance containing as many elementary entities as there are atoms in exactly 12 g of carbon-12, that is 6.02 x 10^23 particles.

The number of constituent particles, atoms, molecules or ions, present in one mole of a substance; its value is 6.02 x 10^23 per mole.

The mass of one mole of a substance in grams per mole, numerically equal to its relative molecular or formula mass.

The formula showing the simplest whole-number ratio of atoms of each element in a compound, without indicating how the atoms are arranged or how large the molecule is.

The formula showing the actual number of atoms of each element in one molecule of a compound; it is a whole-number multiple of the empirical formula.

The reactant that is completely consumed first in a chemical reaction, thereby determining the maximum amount of product that can be formed according to the stoichiometry of the equation.

The actual mass of product obtained expressed as a percentage of the theoretical maximum mass calculated from the balanced equation, reflecting losses and side reactions.

The maximum mass or amount of product predicted from a balanced equation when the limiting reactant is completely converted, assuming no losses and no competing reactions.

The minimum energy required to remove one mole of electrons from one mole of gaseous atoms or ions to form gaseous ions with a single positive charge, measured in kJ per mole.

The enthalpy change when one mole of electrons is added to one mole of gaseous atoms to form singly charged gaseous anions; a negative value indicates a favourable, exothermic process.

A measure of the tendency of an atom in a covalent bond to attract the shared pair of electrons towards itself, commonly expressed on the Pauling scale from about 0.7 to 4.0.

A measure of the size of an atom, usually taken as half the distance between the nuclei of two identical atoms joined by a single bond, or as the calculated radius of an isolated atom.

A regular variation in a property such as atomic radius, ionisation energy or electronegativity across a period or down a group, explained by changing nuclear charge, shielding and atomic size.

The rule that electrons occupy the lowest-energy available orbitals first when building up the electronic configuration of an atom, before filling higher-energy orbitals.

The rule stating that electrons fill orbitals of equal energy singly, with parallel spins, before any pairing occurs, so as to minimise electron-electron repulsion.

The principle that no two electrons in the same atom may have the same set of four quantum numbers, so an orbital can hold at most two electrons with opposite spins.

A bond formed when two atoms share one or more pairs of electrons, usually between non-metals, giving each atom a more stable electron arrangement.

The separation of electrical charge within a covalent bond arising from a difference in electronegativity between the bonded atoms, producing a dipole moment.

A strong intermolecular attraction between a hydrogen atom bonded to a highly electronegative atom such as N, O or F and a lone pair on another electronegative atom nearby.

Weak intermolecular attractions, including instantaneous dipole-induced dipole and permanent dipole-dipole interactions, that arise between molecules and are easily overcome by heating.

A model gas whose particles have negligible volume, exert no intermolecular forces and collide elastically, obeying the equation pV = nRT under all conditions.

The model explaining gas pressure and temperature in terms of the continuous random motion of particles and their collisions with the container walls, linking molecular speed to average kinetic energy.

The volume occupied by one mole of any gas, approximately 22.4 dm3 at standard temperature and pressure (273 K, 1 atm) or 24 dm3 at room temperature and pressure.

The enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions of 298 K and 1 atm pressure.

The principle that the total enthalpy change of a reaction is independent of the route taken, provided the initial and final states are the same, allowing indirect calculation of enthalpy changes.

The enthalpy change when one mole of a substance is completely burned in excess oxygen under standard conditions, usually an exothermic process with a negative enthalpy change.

The state of a reversible reaction in which the forward and reverse rates are equal, so the concentrations of reactants and products remain constant although both reactions continue.

The rule that when a system at equilibrium is disturbed by a change in concentration, pressure or temperature, the system shifts in the direction that opposes the change.

The ratio of the product of equilibrium concentrations of products to that of reactants, each raised to the power of its stoichiometric coefficient in the balanced equation, at a given temperature.

The negative logarithm to base ten of the hydrogen ion concentration in mol per dm3; a value below 7 indicates an acidic solution and above 7 an alkaline solution at 25 degrees Celsius.

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

The equilibrium constant for the partial ionisation of a weak acid in water, giving a measure of acid strength; a larger Ka means a stronger acid.

The equilibrium constant for the dissolution of a sparingly soluble ionic solid, equal to the product of the concentrations of its constituent ions raised to their stoichiometric powers in a saturated solution.

The reduction in the solubility of a sparingly soluble salt when a soluble compound containing one of its ions is added, shifting the dissolution equilibrium to the left.

The change in concentration of a reactant or product per unit time, commonly expressed in mol per dm3 per second and determined from the gradient of a concentration-time graph.

The minimum extra energy that colliding particles must possess for a reaction to occur, corresponding to the height of the energy barrier between reactants and products on an energy profile.

A substance that increases the rate of a reaction by providing an alternative pathway of lower activation energy, without being consumed or permanently changed in the overall reaction.

The proportionality constant k in a rate equation that relates reaction rate to reactant concentrations; its value depends on temperature and the presence of a catalyst.

A number assigned to an element in a compound representing the charge an atom would carry if electrons in its bonds were assigned to the more electronegative atom.

A reaction involving simultaneous oxidation, the loss of electrons, and reduction, the gain of electrons, which can be identified from changes in oxidation number.

The electromotive force of a cell in which the electrode under study is the cathode and a standard hydrogen electrode is the anode, measured under standard conditions.

The decomposition of a compound in molten or aqueous form by passing a direct electric current through it, with oxidation at the anode and reduction at the cathode.

An atom or group of atoms within a molecule that is responsible for its characteristic chemical reactions and physical properties, such as the hydroxyl or carboxyl group.

A family of organic compounds with the same functional group, similar chemical properties and a gradual change in physical properties, successive members differing by a CH2 unit.

The existence of compounds with the same molecular formula but different arrangements of atoms, giving different physical and chemical properties, including chain, position and functional group isomers.

The time taken for a component of a mixture to travel through a chromatographic column from injection to detection, used to identify substances by comparison with standards.

The ratio of the distance travelled by a spot of a substance to the distance travelled by the solvent front in paper or thin-layer chromatography; it is constant for a given substance under fixed conditions.

A separation technique in which components of a mixture distribute themselves between a stationary phase and a mobile phase, moving at different rates according to their relative affinities.

A volumetric procedure in which a solution of known concentration is added from a burette to a measured volume of analyte until the end point is signalled by an indicator or instrument.

The point in a titration at which the indicator changes colour or the instrument signals completion, ideally coinciding with the equivalence point where the reactants have reacted in exact stoichiometric proportions.

A solution of accurately known concentration, prepared by dissolving a precisely weighed mass of a pure primary standard in a volumetric flask and making up to a fixed volume.

Accuracy describes how close a measured value is to the true value, while precision describes how close repeated measurements are to one another, regardless of whether they are correct.

Unpredictable variation in measurements caused by uncontrolled factors such as reading fluctuations or environmental changes; it is reduced by taking the average of many repeated readings.

A consistent, one-directional deviation in measurements arising from faulty apparatus, poor technique or incorrect calibration, which shifts all readings by a similar amount and is not removed by averaging.

The digits in a measured value that carry meaning about its precision, including all certain digits plus one estimated final digit; the result of a calculation should not exceed the precision of the input data.