Definitions

At A Level these are the definitions you need to know, and will need to know and understand for university:

Enthalpy: Heat energy

ΔHθ Standard Enthalpy Change: The enthalpy change for a given reaction carried out under standard conditions or 298K and 1atm (100.1KPa). θ means the reaction is at standard pressure

ΔHθ Formation: The enthalpy change under constant pressure when one mole of a compound is formed from its elements with all reactants and products in their standard states

ΔHθ combustion: The enthalpy change under constant pressure when one mole of a compound is burned completely in excess oxygen with all reactants and products in their standard states

ΔHθ atomisation: The enthalpy change under constant pressure when one mole of gaseous atoms is formed from the elements in their standard states under standard conditions. This is positive: separating atoms takes in energy. Magnitude is related to the forces of attraction between atoms, so is affected by nuclear charge and atomic size

ΔHθ dissociation: The enthalpy change under constant pressure when the covalent bonds in a gaseous molecule are broken to form two free radicals

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

ΔHθ first ionisation: The enthalpy change under constant pressure to remove one mole of electrons from one mole of gaseous atoms to form one mole of unipositive gaseous ions

ΔHθ electron affinity: The enthalpy change under constant pressure when one mole of electrons is added to one mole of isolated gaseous ions to form one mole of ions with a single negative charge. The first electron affinity is exothermic because gaseous atoms have a strong affinity for electrons. The subsequent electron affinities are endothermic because energy is needed to overcome the repulsion between the electron and the negative ion

ΔHθ lattice enthalpy (formation): The enthalpy change under constant pressure when gaseous ions come together to form one mole or crystalline solid. Lattice formation is negative. The amount of energy depends on charges of ions, size of ions, type of lattice formed

ΔHθ lattice formation enthalpy (dissociation): The enthalpy change under constant pressure when one mole of crystalline solid is turned into its gaseous ions. Lattice dissociation is positive. The amount of energy depends on charges of ions, size of ions, type of lattice formed

ΔHθ solution: The enthalpy change at constant pressure when one mole of an ionic solid dissolves completely in water so the ions are far apart from each other and do not interact. The lattice is (theoretically) dissociated (endothermic), then gaseous ions formed are hydrated (exothermic). The more negative, the more soluble to salt

blank spaceΔHθ solution = ΣΔHθ dissociation + ΣΔHθ hydration

ΔHθ hydration: The enthalpy change at constant pressure when one mole of gaseous ions dissolves in water to form one mole of aqueous ions. Ions are attracted to water because it is a polar molecule, so form ion-dipole bonds, releasing energy. Amount of energy released depends on ionic radius and charge (charge/size ratio), and increases across the period for metallic ions and decreases down the group (for all ions)

 

Back to Contents: Chemistry: Thermodynamics

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