Summary A level chemistry revision notes (physical)

, SECTION 2 – Definitions
Enthalpy change Definition Exo/endothermic Example

Enthalpy of formation Enthalpy change when one mole of a substance is formed from its Exothermic (-ve) for e.g. Na2O(s) 2 Na(s) + ½ O2(g) ® Na2O(s)
(DfHº): constituent elements with all substances in their standard states most substances

Enthalpy of Enthalpy change when one mole of a substance undergoes complete Exothermic (-ve) e.g. hydrogen H2(g) + ½ O2(g) ® H2O(l)
combustion (DcHº) combustion in oxygen with all substances in standard states

Enthalpy of Enthalpy change when 1 mole of water is formed in a reaction between an Exothermic (-ve) e.g. H2SO4 + NaOH ½H2SO4(aq) + NaOH(aq) ® ½Na2SO4(aq) + H2O(l)
neutralisation (DneutHº) acid and alkali under standard conditions

Ionisation enthalpy First ionisation energy = enthalpy change when each atom in one mole of Endothermic (+ve) e.g. magnesium Mg(g) ® Mg+(g) + e-
(DieHº) gaseous atoms loses one electron to form one mole of gaseous 1+ ions.

Second ionisation energy = enthalpy change when each ion in one mole of Endothermic (+ve) e.g. magnesium Mg+(g) ® Mg2+(g) + e-
gaseous 1+ ions loses one electron to form one mole of gaseous 2+ ions.

Electron affinity First electron affinity = enthalpy change when each atom in one mole of Exothermic (-ve) for e.g. oxygen O(g) + e- ® O–(g)
(DeaHº) gaseous atoms gains one electron to form one mole of gaseous 1– ions. many non-metals

Second electron affinity = enthalpy change when each ion in one mole of Endothermic (+ve) e.g. oxygen O–(g) + e- ® O2–(g)
gaseous 1– ions gains one electron to form one mole of gaseous 2– ions. as adding –ve
electron to –ve ion

Enthalpy of Enthalpy change when one mole of gaseous atoms is produced from an Endothermic (+ve) e.g. iodine ½ I2(s) ® I(g)
atomisation (DatHº) element in its standard state.

Hydration enthalpy Enthalpy change when one mole of gaseous ions become hydrated Exothermic (-ve) e.g. magnesium ions Mg2+(g) + aq ® Mg2+(aq)
(DhydHº) (dissolved in water).

Enthalpy of solution Enthalpy change when one mole of an ionic solid dissolves in an amount Varies e.g. magnesium chloride MgCl2(s) + aq ® Mg2+(aq) + 2 Cl–(aq)
(DsolHº) of water large enough so that the dissolved ions are well separated and do
not interact with each other.

Bond dissociation Enthalpy change when one mole of covalent bonds is broken in the Endothermic (+ve) e.g. I-I bond I2(g) ® 2 I(g)
enthalpy (DdisHº) gaseous state.

Lattice enthalpy of Enthalpy change when one mole of a solid ionic compound is formed from Exothermic (-ve) e.g. magnesium chloride Mg2+(g) + 2 Cl–(g) ® MgCl2(s)
formation (DLEFHº) into its constituent ions in the gas phase

Lattice enthalpy of Enthalpy change when one mole of a solid ionic compound is broken up Endothermic (+ve) e.g. magnesium chloride MgCl2(s) ® Mg2+(g) + 2 Cl–(g)
dissociation (DLEDHº) into its constituent ions in the gas phase

Enthalpy of Enthalpy change when one mole of a liquid is turned into a gas Endothermic (+ve) e.g. H2O(l) H2O(l) ® H2O(g)
vaporisation (DvapHº):

Enthalpy of fusion Enthalpy change when one mole of a solid is turned into a liquid Endothermic (+ve) e.g. Mg(s) Mg(s) ® Mg(l)
(DfusHº):



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