Tamil Nadu HSE First Year Chemistry key terms

The amount of substance containing exactly 6.022 × 10^23 elementary entities such as atoms, molecules or ions; it links mass in grams to particle count in calculations.

The mass of one mole of a substance in grams, numerically equal to its relative molecular or formula mass; used to convert between mass and moles.

The simplest whole-number ratio of atoms of each element present in a compound, obtained from percentage composition data by dividing moles by the smallest value.

The actual number of atoms of each element in one molecule of a compound; it is a whole-number multiple of the empirical formula, found using molar mass.

The reactant that is completely consumed first in a reaction and therefore fixes the maximum amount of product formed; other reactants are said to be in excess.

The quantitative relationship between amounts of reactants and products in a balanced chemical equation, based on mole ratios taken from the coefficients.

The digits in a measured value that carry meaningful information about its precision, including all certain digits plus one uncertain final digit; they govern rounding of results.

The rule that electrons occupy orbitals in order of increasing energy, so that lower-energy orbitals are filled before higher-energy ones in the ground state.

The rule that degenerate orbitals of a subshell are each singly occupied with parallel spins before any pairing occurs, minimising electron-electron repulsion.

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

A region of space around the nucleus where the probability of finding an electron is high; described by three quantum numbers and labelled s, p, d or f.

The set of numbers (principal, azimuthal, magnetic, spin) that specifies the energy, shape, orientation and spin state of an electron in an atom.

The statement that the position and momentum of an electron cannot both be measured exactly at the same time; the product of their uncertainties is at least h/4π.

The net positive charge experienced by an outer electron after accounting for shielding by inner electrons; it increases across a period and influences atomic radius and ionisation energy.

A measure of the size of an atom, usually taken as half the distance between neighbouring nuclei; it decreases across a period and increases down a group.

The energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state; it generally increases across a period and decreases down a group.

The enthalpy change when one electron is added to an isolated gaseous atom; a negative value means the process releases energy and the anion is stable.

The tendency of an atom in a molecule to attract the shared pair of electrons towards itself; it increases across a period and decreases down a group.

A covalent bond in which the shared electrons are unequally attracted by the two atoms because of an electronegativity difference, giving partial charges at the ends.

The mixing of atomic orbitals of similar energy on the same atom to give an equal number of equivalent hybrid orbitals with definite geometry, such as sp, sp2 and sp3.

The model that predicts molecular shape by arranging electron pairs around the central atom as far apart as possible, with lone pairs repelling more strongly than bond pairs.

The number of bonds between two atoms, equal to half the difference between bonding and antibonding electrons in molecular orbital theory; higher order means shorter, stronger bonds.

The product of the magnitude of a charge and the distance separating the charges in a polar molecule; it is a vector sum and indicates molecular polarity and symmetry.

A hypothetical gas whose molecules have negligible volume and no intermolecular forces, obeying PV = nRT exactly at all pressures and temperatures.

A modified gas equation, (P + an²/V²)(V − nb) = nRT, that corrects for intermolecular attraction (a) and finite molecular volume (b) in real gases.

The highest temperature above which a gas cannot be liquefied however great the pressure applied; above it the substance exists only as a supercritical fluid.

The pressure a single gas in a mixture would exert if it alone occupied the whole container at the same temperature; the total pressure is the sum of partial pressures.

A property whose value depends only on the current state of the system, not on the path taken to reach it; examples include enthalpy, entropy and internal energy.

The heat content of a system at constant pressure, defined as H = U + PV; its change equals the heat absorbed or released in a constant-pressure process.

The law that the enthalpy change of a reaction is the same whether it occurs in one step or several, because enthalpy is a state function.

A measure of the disorder or randomness of a system; it increases when a solid melts, a liquid vaporises or a gas expands, and drives the direction of spontaneous change together with enthalpy.

A thermodynamic quantity, G = H − TS, whose decrease at constant temperature and pressure predicts that a process is spontaneous; zero change means equilibrium.

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

The ratio of product concentrations to reactant concentrations, each raised to the power of its stoichiometric coefficient, at equilibrium; it is constant at a given temperature.

The same ratio as the equilibrium constant but calculated with the actual concentrations at any moment; comparing it with K tells which way the reaction must shift.

A solution that resists appreciable change in pH on adding a small amount of strong acid or base, usually a weak acid with its conjugate base or a weak base with its conjugate acid.

The suppression of ionisation of a weak electrolyte when a strong electrolyte providing a common ion is added, shifting the equilibrium towards the undissociated form.

The product of the molar concentrations of the ions of a sparingly soluble salt in a saturated solution, each raised to its stoichiometric power, at a given temperature.

A measure of the acidity of a solution defined as the negative logarithm to base ten of the hydrogen ion concentration; values below 7 are acidic at 25 °C.

Properties of dilute solutions that depend only on the number of solute particles, not their nature: relative lowering of vapour pressure, boiling point elevation, freezing point depression and osmotic pressure.

The number of moles of solute dissolved in one kilogram of solvent; unlike molarity it does not change with temperature because it uses mass.

The statement that the partial vapour pressure of each volatile component in a solution is its mole fraction multiplied by the vapour pressure of the pure component.

The sum of the powers of the concentration terms in the experimentally determined rate law; it may be zero, fractional or negative and is not the same as molecularity.

The proportionality constant k in a rate law, equal to the rate when all reactant concentrations are unity; its dependence on temperature follows the Arrhenius equation.

The minimum extra energy that colliding reactant molecules must possess for a reaction to occur; a larger value makes the reaction slower and more temperature-sensitive.

When a hydrogen halide adds to an unsymmetrical alkene, the hydrogen goes to the carbon carrying more hydrogens and the halogen to the more substituted carbon.

A reaction in which an electron-rich double or triple bond attacks an electrophile, forming a carbocation intermediate that is then captured by a nucleophile.

A positively charged carbon species with only three bonds and an empty p orbital; its stability rises from primary to secondary to tertiary because of hyperconjugation and inductive effects.

The special stability of a planar, cyclic, fully conjugated system with 4n + 2 π electrons; benzene with six π electrons is the standard example.

An atom or group of atoms that gives a family of organic compounds its characteristic chemical properties, such as the hydroxyl, carbonyl, carboxyl or amino group.

A family of organic compounds with the same functional group in which successive members differ by one CH2 unit and show a gradual change in physical properties.

The existence of two or more compounds with the same molecular formula but different arrangements of atoms, giving different structures or spatial configurations.

The permanent displacement of shared electrons along a sigma bond chain caused by an electronegative or electropositive atom, weakening with distance.

The systematic set of rules for naming organic compounds: select the longest chain, number to give the principal functional group the lowest locant, and list substituents alphabetically.