FY13CE Chemistry key terms

The number of protons in the nucleus of an atom, which uniquely identifies the element. In a neutral atom it also equals the number of electrons.

The total count of protons plus neutrons in one atom's nucleus. It is a whole number written as a superscript beside the element symbol and differs between isotopes.

One of two or more atoms of the same element that have equal proton counts but different neutron counts, so they differ in mass while sharing nearly identical chemical behaviour.

The weighted average mass of an element's atoms compared with one-twelfth the mass of a carbon-12 atom, taking natural isotopic abundance into account. It is dimensionless and rarely a whole number.

The rule that electrons occupy the lowest-energy available orbitals first when an atom is built up, giving the familiar 1s 2s 2p 3s filling order and the ground-state configuration.

The arrangement of an atom's electrons among its shells, subshells and orbitals, usually written with superscripts such as 1s2 2s2 2p6. It governs chemical behaviour and position in the periodic table.

The statement that electrons fill degenerate orbitals singly, with parallel spins, before any pairing occurs. It explains the unpaired electrons in nitrogen and the paramagnetism of many transition-metal ions.

The electrons in an atom's outermost occupied shell that are chiefly responsible for bonding and chemical reactivity. Their count determines group behaviour and typical ion charge.

The minimum energy needed to remove one electron from a gaseous atom or ion. First ionisation energy generally rises across a period and falls down a group as atomic radius increases.

A measure of how strongly an atom in a bond attracts shared electrons. Values rise across a period and fall down a group; large differences between bonded atoms favour ionic character.

The electrostatic attraction between oppositely charged ions formed when one atom transfers electrons to another. Such bonds typically join a metal cation to a non-metal anion and give high melting points.

A bond formed when two atoms share one or more pairs of electrons. The shared pair is attracted to both nuclei, holding the atoms together with a definite bond length and energy.

A covalent bond in which the shared electrons are unevenly distributed because the two atoms differ in electronegativity, producing partial positive and partial negative charges at the bond ends.

A quantity describing the separation of charge in a polar molecule or bond, equal to charge times distance and measured in debye. It is a vector, so molecular geometry decides whether bond dipoles cancel.

A model predicting molecular shape by assuming electron pairs around a central atom repel and spread as far apart as possible. Lone pairs repel more strongly than bonding pairs, compressing bond angles.

The mixing of atomic orbitals on a central atom to give equivalent new orbitals that explain observed geometry, such as sp3 tetrahedral, sp2 trigonal planar and sp linear arrangements.

Attractions between separate molecules, including London dispersion, dipole-dipole and hydrogen bonding. They are much weaker than covalent bonds yet control boiling point, viscosity and solubility.

An unusually strong dipole attraction between a hydrogen atom bonded to N, O or F and a lone pair on a nearby electronegative atom. It explains water's high boiling point and ice's low density.

A hypothetical gas whose particles occupy negligible volume and exert no forces except elastic collisions. It obeys PV = nRT exactly; real gases approach this behaviour at low pressure and high temperature.

The amount of substance containing 6.022 x 10^23 specified entities, which is Avogadro's number. It links particle counts to measurable masses through molar mass.

The mass of one mole of a substance, expressed in grams per mole and numerically equal to its relative formula mass. It converts between mass and amount in stoichiometric calculations.

The reactant that is used up first in a reaction and therefore fixes the maximum amount of product formed. Other reactants are said to be in excess.

The actual mass of product obtained expressed as a percentage of the theoretical maximum calculated from stoichiometry. Losses arise from side reactions, incomplete reaction and handling.

The heat transferred at constant pressure during a reaction, given the symbol delta H. Negative values indicate exothermic processes; positive values indicate endothermic ones.

The principle that the total enthalpy change of a reaction is independent of the route taken, so indirect paths can be combined. It allows calculation of changes that cannot be measured directly.

The average energy required to break one mole of a particular bond in gaseous molecules. Summing broken minus formed bond enthalpies estimates the enthalpy change of a gas-phase reaction.

The experimental measurement of heat exchanged, usually by tracking a temperature change in a known mass of water or solution. The result is calculated with q = mc delta T.

The minimum extra energy reacting particles must possess for a collision to produce products. Raising temperature or adding a catalyst increases the fraction of collisions that succeed.

A substance that speeds a reaction by providing an alternative pathway of lower activation energy, without being consumed overall. It changes rate but not the position of equilibrium.

When a system at equilibrium is disturbed by a change in concentration, pressure or temperature, it shifts in the direction that opposes the change. It predicts the new equilibrium position.

The ratio of product concentrations to reactant concentrations, each raised to its stoichiometric power, at equilibrium. Its value depends on temperature and indicates how far a reaction proceeds.

A logarithmic scale of hydrogen ion concentration defined as pH = -log[H+]. Values below 7 are acidic, 7 is neutral at 25 degrees Celsius, and above 7 is alkaline.

An acid that dissociates almost completely in water, such as hydrochloric or nitric acid. Its solutions have low pH and conduct electricity well because of the abundant ions.

A solution that resists pH change when small amounts of acid or base are added, usually a weak acid with its conjugate base. It works by neutralising added acid or base.

The loss of electrons by a species, accompanied by an increase in oxidation number. In redox reactions oxidation is always coupled with reduction of another species.

The gain of electrons by a species, accompanied by a decrease in oxidation number. The oxidising agent is the species that is itself reduced.

A formal charge assigned to an atom in a compound by convention, assuming bonds are fully ionic. Tracking its changes identifies which species are oxidised and reduced.

The potential of a half-cell measured against the standard hydrogen electrode under standard conditions of one molar concentration and one atmosphere at 298 K. More positive values indicate greater tendency to be reduced.

An atom or group of atoms that gives an organic molecule its characteristic reactions, such as hydroxyl, carbonyl, carboxyl or amino. Members of a homologous series share one.

A family of organic compounds with the same functional group and general formula, where successive members differ by one CH2 unit. Members show gradually changing physical properties and similar chemistry.

One of two or more compounds with the same molecular formula but different arrangements of atoms. Structural isomers differ in connectivity; stereoisomers differ only in spatial arrangement.

A volumetric procedure in which a solution of known concentration is added from a burette until it exactly reacts with a measured volume of analyte. Indicator colour change marks the end point.

A solution whose concentration is accurately known, prepared by weighing a pure solid or diluting a stock solution. It is essential for reliable quantitative titrations.

A separation technique in which components distribute differently between a mobile phase and a stationary phase. Differences in travel rate allow identification and sometimes quantification of mixture components.

The energy needed to raise the temperature of one gram of a substance by one kelvin, measured in joules per gram per kelvin. Water's value of about 4.18 supports calorimetric calculations.