Energetics- enthalpy change definitions

Enthalpy of formation

• Enthalpy change when one mole of a substance is formed from its constituent elements with all substances in their standard states

• exothermic (-ve) for most substances

• E.g 2Na(s) + ½ O2(g) —> Na2O(s)

Enthalpy of combustion

• Enthalpy change when one mole of a substance undergoes complete combustion in oxygen with all substances in standard states

• Exothermic (-ve)

• E.g H2(g) + ½ O2 (g) —> H2O (l)

Enthalpy of neutralisation

• Enthalpy change when one mole of WATER is formed in a reaction between an acid and alkali under standard conditions

• Exothermic (-ve)

• e.g ½ H2SO4(aq) + NaOH (aq) —> ½ Na2SO4(aq) +H2O(l)

Ionisation Enthalpy (two definitions- 1st and 2nd ionisation energies)

FIRST ionisation energy= Enthalpy change when each atom in one mole of gaseous atoms loses one electron to form one mole of gaseous 1+ ions.

• endothermic (+ve)

• E.g Mg(g) —> Mg+(g) + e-

SECOND ionisation energy= Enthalpy change when each ion in one molecule of gaseous 1+ ions loses one electron to form one mole of gaseous 2+ ions

• endothermic (+ve)

• E.g Mg+(g) —> Mg2+(g) +e-

Electron affinity (2 definitions- first and second)

FIRST electron affinity= Enthalpy change when each atom in one mole of gaseous atoms gains one electron to form one mole of gaseous 1- ions.

• exothermic (-ve) for many non-metals

• E.g O(g) + e- —> O- (g)

SECOND electron affinity= Enthalpy change when each ion in one mole of gaseous 1- ions gains one electron to form one mole of gaseous 2- ions

• endothermic (+ve) as adding -ve electron to -ve ion

• E.g O-(g) + e- —> O2- (g)

Enthalpy of atomisation

Enthalpy change when one mole of gaseous atoms is produced from an element in its standard state

• endothermic (+ve)

• E.g iodine : ½ I2(s) —> I(g)

Hydration Enthalpy

Enthalpy change when one mole of gaseous ions become hydrated (dissolved in water)

• exothermic (-ve)

• E.g Mg2+ (g) + aq —> Mg2+ (aq)

Enthalpy of solution

Enthalpy change when one mole of an ionic solid dissolves in an amount of water large enough so that the dissolved ions are well separated and do not interact with each other

• exo/endothermic varies

• E.g MgCl2(s) + aq —> Mg2+ (aq) + 2Cl- (aq)

Bond dissociation enthalpy

Enthalpy change when one mole of covalent bonds is broken in the gaseous state

• endothermic (+ve)

• E.g I-I bond (iodine): I2(g) —> 2 I(g)

Lattice enthalpy of formation

Enthalpy change when one mole of a solid ionic compound is formed into its constituent ions in the gas phase

• exothermic (-ve)

• E.g Mg2+(g) + 2Cl-(g) —> MgCl2(s)

Lattice enthalpy of dissociation

Enthalpy change when one mole of a solid ionic compound is broken up into its constituent ions in the gas phase

• endothermic (+ve)

• E.g MgCl2(s) —> Mg2+(g) + 2Cl-(g)

Enthalpy of vaporisation

Enthalpy change when one mole of a liquid is turned into a gas

• endothermic (+ve)

• E.g H2O(l) —> H2O(g)

Enthalpy of fusion

Enthalpy change when one mole of a solid is turned into a liquid

• endothermic (+ve)

• E.g Mg(s) —> Mg(l)