Thermodynamics and Enthalpy Changes Flashcards

A set of vocabulary flashcards covering key definitions in thermodynamics, including enthalpy changes, lattice energetics, entropy, and Gibbs free energy.

Enthalpy change

The heat change measured at constant pressure.

Standard conditions

Chosen conditions for measuring enthalpy changes, typically set at $100\,kPa$ and a stated temperature, usually $298\,K$.

Standard molar enthalpy of formation ($\Delta_f H^{\ominus}$)

The enthalpy change when one mole of a compound is formed from its constituent elements under standard conditions, with all reactants and products in their standard states.

Standard molar enthalpy change of combustion ($\Delta_c H^{\ominus}$)

The enthalpy change when one mole of substance is completely burnt in oxygen under standard conditions.

Standard enthalpy of atomisation ($\Delta_{at} H^{\ominus}$)

The enthalpy change which accompanies the formation of one mole of gaseous atoms from the element in its standard state under standard conditions.

First ionisation energy (first IE) ($\Delta_i H^{\ominus}$)

The standard enthalpy change when one mole of gaseous atoms is converted into a mole of gaseous ions each with a single positive charge (e.g., $Mg(g) \rightarrow Mg^+(g) + e^-$).

Second ionisation energy (second IE)

The energy change for the loss of a mole of electrons from a mole of singly positively charged gaseous ions (e.g., $Mg^+(g) \rightarrow Mg^{2+}(g) + e^-$).

First electron affinity ($\Delta_{ea} H^{\ominus}$)

The standard enthalpy change when a mole of gaseous atoms is converted to a mole of gaseous ions, each with a single negative charge.

Second electron affinity

The enthalpy change when a mole of electrons is added to a mole of gaseous ions each with a single negative charge to form ions each with two negative charges (e.g., $O^-(g) + e^- \rightarrow O^{2-}(g)$).

Lattice enthalpy of formation ($\Delta_L H^{\ominus}$)

The standard enthalpy change when one mole of solid ionic compound is formed from its gaseous ions.

Lattice enthalpy of dissociation

The standard enthalpy change when one mole of solid ionic compound dissociates into its gaseous ions; it has the same numerical value as lattice formation enthalpy but is always positive.

Enthalpy of hydration ($\Delta_{hyd} H^{\ominus}$)

The standard enthalpy change when water molecules surround one mole of gaseous ions.

Enthalpy of solution ($\Delta_{sol} H^{\ominus}$)

The standard enthalpy change when one mole of solute dissolves completely in sufficient solvent to form a solution in which the molecules or ions are far enough apart not to interact with each other.

Mean bond enthalpy

The enthalpy change when one mole of gaseous molecules each breaks a covalent bond to form two free radicals, averaged over a range of compounds.

Hess' law

States that the enthalpy change for a chemical reaction is always the same, whatever route is taken from reactants to products.

Born-Haber cycle

A thermochemical cycle that includes all the enthalpy changes involved in the formation of an ionic compound.

Polarisation

The distortion of the electron clouds of a large negative ion (anion) by a small, highly charged positive ion (cation), leading to a degree of covalent character in the bonding.

Feasible (or Spontaneous) reaction

A reaction that has a tendency to happen of its own accord at a particular temperature, determined by the Gibbs free energy change ($\Delta G < 0$).

Entropy ($S$)

A mathematical expression of the randomness or disorder of a system, measured in $J\,K^{-1}\,mol^{-1}$.

Gibbs free energy change ($\Delta G$)

A quantity that combines enthalpy and entropy factors to determine feasibility, calculated using the equation $\Delta G = \Delta H - T\Delta S$.

Thermodynamic stability

Determined by the Gibbs free energy change; a substance is thermodynamically unstable if its reaction with another substance (like oxygen) has a negative $\Delta G$.

Kinetic stability

Occurs when a reaction is thermodynamically feasible but has such a large activation energy barrier that the reaction rate is unmeasurable at room temperature.