thermodynamics definitions

• Enthalpy of formation ΔfH⦵

  • Enthalpy change when one mole of a substance is formed from its constituent elements with all substances in their standard states
  • Exothermic
  • eg. Na2O(s) = 2Na(s)+1/2O2(g) -> Na2O(s)
• Enthalpy of combustion ΔcH⦵

  • Enthalpy change when one mole of a substance undergoes complete combustion in oxygen with all substances in standard states
  • Exothermic
  • Eg. hydrogen = H2(g) + 1/2O2(g) -> H2O(l)
• Enthalpy of neutralisation ΔneutH⦵

  • Enthalpy change when one mole of water is formed in a reaction between an aid and alkali under standard conditions
  • Exothermic
  • eg. 1/2H2SO4(aq) + NaOH(aq) -> 1/2Na2SO4(aq) + H2O(l)
• Ionisation enthalpy ΔieH⦵

  • 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
  • second ionisation energy = enthalpy change when each ion in one mole of gaseous 1+ ions loses one electron to form one mole of gaseous 2+ ions
  • Endothermic
  • eg 1. Mg(g) -> Mg+(g) + e-
  • eg 2. Mg+(g) -> Mg2+(g) +e-
• Electron affinity ΔeaH⦵

  • 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)
  • 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)
  • eg 1. O(g) + e- -> O-
  • eg 2. O-(g) + e- -> O-2
• Enthalpy of atomisation ΔatH⦵

  • Enthalpy change when one mole of gaseous atoms is produced from an element in its standard state
  • Endothermic
  • 1/2I2(s) -> I(g)
• Hydration enthalpy ΔhydH⦵

  • enthalpy change when one mole of gaseous ions become hydrated
  • Exothermic
  • eg. Mg2+(g) + aq -> Mg2+(aq)
• Enthalpy of solution ΔsolH⦵

  • 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/Endo varies
  • eg. MgCl2(s) + aq -> Mg2+(aq) + 2cl-(aq)
• Bond dissociation enthalpy ΔdisH⦵

  • Enthalpy change when one mole of covalent bonds is broken in the gaseous state
  • Endothermic
  • eg. I2(g) -> 2I(g)
• Lattice enthalpy of formation ΔlefH⦵

  • enthalpy change when one mole of a solid ionic compound is formed into its constituent ions in the gas phase
  • exothermic
  • eg. Mg2+(g) + 2Cl-(g) -> MgCl2(s)
• lattice enthalpy of dissociation ΔledH⦵

  • Enthalpy change when one mole of a solid ionic compound is broken up into its constituent ions in the gas phase
  • endothermic
  • eg. MgCl2(s) -> Mg2+(g) + 2Cl-(g)
• Enthalpy of vaporisation ΔvapH⦵

  • Enthalpy change when one mole of a liquid is turned into a gas
  • Endothermic
  • eg. H2O(l) -> H2O(g)
• Enthalpy of fusion ΔfusH⦵

  • enthalpy change when one mole of a solid is turned into a liquid
  • endothermic
  • eg. Mg(s) -> Mg(l)