Thermodynamics and Enthalpy Changes Flashcards

A set of vocabulary flashcards based on lecture notes covering thermodynamics, specific enthalpy definitions, Born-Haber cycles, entropy, and Gibbs free energy.

Enthalpy change

The heat change measured at constant pressure.

Standard conditions

Chosen conditions for measuring enthalpy changes, specified as $100\,kPa$ 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, 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.

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.

Second ionisation energy (second IE)

The enthalpy change refers to the loss of a mole of electrons from a mole of singly positively charged ions, represented as $M^+(g) \rightarrow M^{2+}(g) + e^-$.

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.

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; this process results in energy being given out and is always negative.

Lattice enthalpy of dissociation

The standard enthalpy change when one mole of solid ionic compound dissociated into its gaseous ions; this value 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).

Feasible or Spontaneous reaction

Terms used to describe reactions which could take place of their own accord, governed by enthalpy change and entropy factors.

Entropy ($S$)

The randomness or disorder of a system, expressed mathematically. Entropy usually increases with temperature and as matter changes from solid to liquid to gas.

Gibbs free energy change ($\Delta G$)

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

Thermodynamically unstable

A substance or reaction where the Gibbs free energy change ($\Delta G$) is negative, indicating the reaction is feasible.

Kinetically stable

A state where a reaction is thermodynamically feasible but occurs so slowly in practice (due to a large activation energy barrier) that no change is measured.