energy

• Enthalpy of atomisation


  The enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state
• Enthalpy of atomisation


  • Na (s) → Na(g) [atH = +148 kJ mol-1]
  • ½ O2 (g) → O (g) [atH = +249 kJ mol-1]
• Enthalpy of sublimation


  The enthalpy change for a solid metal turning to gaseous atoms, numerically the same as the enthalpy of atomisation
• Enthalpy of sublimation


  • Na (s) → Na(g) [subH = +148 kJ mol-1]
• First ionisation enthalpy


  The enthalpy change required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous ions with a +1 charge
• Second ionisation enthalpy


  The enthalpy change to remove 1 mole of electrons from one mole of gaseous 1+ ions to produces one mole of gaseous 2+ ions
• First electron affinity


  The enthalpy change that occurs when 1 mole of gaseous atoms gain 1 mole of electrons to form 1 mole of gaseous ions with a –1 charge
• First electron affinity


  • O (g) + e- → O- (g) [eaH] = -141.1 kJ mol-1
• Second electron affinity


  The enthalpy change when one mole of gaseous 1- ions gains one electron per ion to produce gaseous 2- ions
• Second electron affinity


  • O – (g) + e- → O2- (g) [eaH = +798 kJ mol-1]
• Lattice Enthalpy (LEH)


  The standard enthalpy change when 1 mole of an ionic crystal lattice is formed from its constituent ions in gaseous form
• Lattice Enthalpy


  • Na+ (g) + Cl- (g) → NaCl (s) [LEH = -787 kJ mol-1]
• Enthalpy of Hydration (hydH)

  The enthalpy change when one mole of gaseous ions become aqueous ions
• Enthalpy of Hydration


  • X+ (g) + aq → X+ (aq) [For Li+ hydH = -519 kJ mol-1]
  • X- (g) + aq → X- (aq) [For F- hydH = -506 kJ mol-1]
• Enthalpy of solution (solH)

  The standard enthalpy change when one mole of an ionic solid dissolves in a large enough amount of water
• Enthalpy of solution


  • NaCl (s) + aq → Na+ (aq) + Cl-(aq) [solH]
• Enthalpy of formation (fH)

  The standard enthalpy change when 1 mole of a compound is formed from its elements under standard conditions
• Enthalpy of formation


  • Na (s) + ½Cl2 (g) → NaCl (s) [fH = - 411.2 kJ mol-1]
• Lattice Enthalpy


  • Can be used as a measure of ionic bond strength
  • Cannot be determined directly, calculated indirectly using Born-Haber cycle
• Born-Haber cycle


  Applies Hess's law to calculate lattice enthalpy indirectly using other enthalpy changes
• Factors affecting lattice enthalpy

  • Sizes of the ions (larger ions = less negative lattice enthalpy)
  • Charges on the ions (higher charges = more negative lattice enthalpy)
• Trends in lattice enthalpies
• Entropy (S)


  A description of the number of ways atoms can share quanta of energy. Higher disorder = higher entropy
• Substances with higher entropy
• Change in entropy (ΔS)


  • Positive if there is an increase in disorder (e.g. solid/liquid to gas, increase in number of molecules)
  • Negative if there is a decrease in disorder (e.g. gas to solid/liquid, decrease in number of molecules)
• Calculating ΔS


  ΔS = Σ S°products - Σ S°reactants
• Gibbs free energy change (ΔG)

  Combines the effects of enthalpy and entropy into one value
• Elements in their standard states do not have zero entropy. Only perfect crystals at absolute zero (T = 0 K) will have zero entropy
• Standard entropy change (ΔS°)


  Σ S°products - Σ S°reactants
• Unit of entropy


  J K-1 mol-1
• At 0K substances have zero entropy. There is no disorder as particles are stationary
• Activation Energy (EA)


  The energy barrier that must be overcome for a reaction to occur
• ΔS° = 558.1 J K-1 mol-1 = 558 J K-1 mol-1 (3 S.F.) for the reaction: 2Fe2O3 (s) + 3C (s) 4Fe (s) + 3CO2 (g)
• Gibbs Free Energy Change (ΔG)


  A term that combines the effect of enthalpy and entropy into one number
• Relationship between ΔG, ΔH and ΔS

  ΔG = ΔH - TΔS
• For any spontaneous change, ΔG will be negative
• A reaction that has increasing entropy (+ve ΔS) and is exothermic (-ve ΔH) will make ΔG be negative and will always be feasible
• For the reaction Al2O3(s) + 3C(s) → 2Al(s) + 3CO(g): ΔH° = +1336 kJ mol-1, ΔS° = +581 J K-1 mol-1, ΔG = +1163 kJ mol-1
• The reaction N2(g) + O2(g) → 2NO(g) will be feasible at temperatures above 7200K
• The temperature for methane to melt is 91K