Gaseous State of Matter & Energetics – Key Vocabulary

Vocabulary flashcards covering major concepts from Units 4 (Gaseous State) and 5 (Energetics) of the A/L Chemistry resource book.

Ideal Gas

A hypothetical gas whose molecules occupy negligible volume, experience no intermolecular forces, and obey PV = nRT at all temperatures and pressures.

Ideal Gas Equation

The relationship PV = nRT linking the pressure, volume, amount (moles) and absolute temperature of an ideal gas.

Gas Constant (R)

The proportionality constant in the ideal-gas law; R = 8.314 J mol⁻¹ K⁻¹.

Boyle’s Law

For a fixed amount of gas at constant temperature, pressure is inversely proportional to volume (PV = constant).

Charles’s Law

For a fixed amount of gas at constant pressure, volume is directly proportional to absolute temperature (V ∝ T).

Avogadro’s Law

Equal volumes of all gases at the same temperature and pressure contain equal numbers of molecules; thus V ∝ n.

Molar Volume

Volume occupied by one mole of gas at a stated T and P; 22.414 dm³ mol⁻¹ at 0 °C and 1 atm.

Combined Gas Law

For a fixed amount of gas, P₁V₁/T₁ = P₂V₂/T₂ relates simultaneous changes in P, V and T.

Dalton’s Law of Partial Pressures

The total pressure of a non-reacting gas mixture equals the sum of the partial pressures of its components (PT = ΣPi).

Mole Fraction (χ)

The ratio of moles of a component to total moles in a mixture; Pi = χi PT.

Compressibility Factor (Z)

Dimensionless quantity Z = PV/RT for one mole; indicates the deviation of a real gas from ideal behaviour (Z = 1 for an ideal gas).

Real Gas

A gas whose molecules have finite volume and intermolecular forces, causing deviations from the ideal-gas law, especially at high P or low T.

van der Waals Equation

(P + a n²/V²)(V − nb) = nRT; an adjusted equation of state that corrects for molecular attractions (a) and finite size (b).

Critical Temperature (Tc)

The highest temperature at which a gas can be liquefied by pressure alone; above Tc no amount of pressure yields a liquid.

Root Mean Square Speed (crms)

The square root of the mean of the squared molecular speeds; crms = (3RT/M)½.

Maxwell–Boltzmann Distribution

The statistical distribution that describes the fraction of gas molecules possessing each speed at a given temperature.

Enthalpy (H)

A thermodynamic state function equal to the system’s internal energy plus PV; equals heat at constant pressure.

Endothermic Process

A change that absorbs heat from the surroundings; ΔH is positive.

Exothermic Process

A change that releases heat to the surroundings; ΔH is negative.

Standard Enthalpy Change (ΔH°)

Enthalpy change when reactants and products are in their standard states (1 bar, specified T, 1 mol dm⁻³ for solutions).

Specific Heat (c)

Heat required to raise the temperature of 1 g of substance by 1 °C (or 1 K).

Heat Capacity (C)

Heat required to raise the temperature of a given quantity of material by 1 °C; C = m c.

Hess’s Law

The total enthalpy change of a process is the same regardless of the pathway, because enthalpy is a state function.

Standard Enthalpy of Formation (ΔHf°)

Enthalpy change when 1 mol of a compound forms from its elements in their reference states under standard conditions.

Standard Enthalpy of Combustion (ΔHc°)

Enthalpy change when 1 mol of substance burns completely in oxygen under standard conditions.

Standard Enthalpy of Neutralization (ΔHneu°)

Enthalpy change when 1 mol of H⁺(aq) reacts with 1 mol of OH⁻(aq) to form H₂O(l) under standard conditions (≈ –57 kJ mol⁻¹).

Lattice Enthalpy (ΔHL°)

Enthalpy change when 1 mol of a solid ionic lattice is separated into its gaseous ions (or formed from them).

Born–Haber Cycle

A thermochemical cycle that uses Hess’s law to relate lattice enthalpy to atomization, ionization, electron gain and formation steps.

Entropy (S)

A measure of the disorder or randomness of a system; greater randomness corresponds to higher entropy (units J K⁻¹ mol⁻¹).

Standard Entropy Change (ΔS°)

Entropy change when reactants turn to products in their standard states; ΔS° = ΣS°(products) – ΣS°(reactants).

Gibbs Free Energy (G)

State function G = H – T S; at constant T and P, a negative ΔG indicates spontaneity, zero ΔG equilibrium, positive ΔG non-spontaneity.

Spontaneous Process

A change that proceeds without continuous external energy; characterised by ΔG < 0.

Non-spontaneous Process

A change that will not occur unless driven by external energy; ΔG > 0.

Reversible Process

An idealised change that can be reversed by an infinitesimal modification, leaving system and surroundings exactly as before.

Irreversible Process

A real process that cannot return both system and surroundings to their original states; involves finite gradients.

System (Thermodynamics)

The part of the universe chosen for study; separated from surroundings by a boundary.

Surroundings

Everything outside the defined system that can exchange matter or energy with it.

Open System

A system that exchanges both energy and matter with its surroundings.

Closed System

A system that exchanges energy but not matter with its surroundings.

Isolated System

A system that exchanges neither energy nor matter with its surroundings.

Intensive Property

A property independent of the amount of substance present (e.g., temperature, density).

Extensive Property

A property that depends on the amount of substance present (e.g., mass, enthalpy).

Standard State

Reference state of a pure substance at 1 bar (≈ 1 atm) and specified temperature; 1 mol dm⁻³ for solutions, ideal behaviour for gases.

Heat of Vaporization (ΔHvap°)

Enthalpy change when 1 mol of liquid converts to vapour at standard pressure.

Heat of Fusion (ΔHfus°)

Enthalpy change when 1 mol of solid melts to liquid at its melting point under standard pressure.

Heat of Sublimation (ΔHsub°)

Enthalpy change when 1 mol of solid converts directly to gas under standard conditions.

Hydration Enthalpy (ΔHhyd°)

Enthalpy change when 1 mol of gaseous ions becomes solvated by water to give 1 mol dm⁻³ aqueous ions under standard conditions.

Mean Square Speed (c²̅)

The average of the squares of molecular speeds in a gas; related to kinetic energy and pressure via PV = (1/3)Nm c²̅.