Thermodynamics

• Standard enthalpy of atomisation - Enthalpy change when one mole of gaseous atoms is formed from an element in its standard states
• Mean bond enthalpy - Enthalpy change when one mole of gaseous molecules each break a covalent bond to form two free radicals (averaged over a range of compounds)
• Mean bond enthalpy = 2 x standard enthalpy of atomisation
• Standard enthalpy of formation - enthalpy change when one mole of compounds is formed from its elements under standard conditions
• Standard enthalpy of combustion - Enthalpy change when one mole of a compound is completely burned in oxygen under standard conditions
• First ionisation enthalpy - Enthalpy change when one mole of electrons is removed from one mole of gaseous atoms to give one mole of gaseous ions
• Second ionisation enthalpy - Enthalpy change when one mole of electrons is removed from one mole of gaseous 1+ ions to give one mole of gaseous 2+ ions
• First electron affinity - Enthalpy change when one mole of gaseous atoms converted into a mole of gaseous ions with a single negative charge (under standard conditions)
• Second electron affinity - Enthalpy change when one mole of electrons is added to a mole of gaseous ions each with a single negative charge to form one mole of ions each with two negative charge
• Lattice formation enthalpy - Enthalpy change when one mole of solid ionic compound is formed from its gaseous ions
• Lattice dissociation enthalpy - Enthalpy change when one mole of solid ionic compound dissociates into its gaseous ions
• Enthalpy of hydration - Enthalpy change when one mole of gaseous ions is converted into one mole of aqueous ions
• Enthalpy of solution - Enthalpy change when one mole of solute dissolves in enough solvent to form a solution where ions are so far apart they do no interact with each other
• UP arrows on born-haber cycle is positive
• DOWN arrows on born-haber cycle is negative
• LHS = RHS on a born-haber cycle
• Lattice formation enthalpy will either be a really large positive or negative number
• Second electron affinity is always positive because the second electron is more easily removed from the second shell than the first due to repulsion of the electron pairs
• Factors that determine how exothermic a lattice is:

  • Charge on the ions
  • Size of the ions (ionic radius)
• The bigger the charge and the smaller the ion the greater the charge density
• Charge on the ions:

  • Greater the charge an ion has the greater its attraction to an oppositely charged ion
• Size of the ions:

  • The smaller the ion the greater the ionic attraction so smaller radius
• A positive ion which is small and highly charged is very polarising
• A negative ion which is large and highly charged is very polarisable
• Highly charged = Highly polarising
• Theoretical Lattice Enthalpies:

  Model name - Perfect ionic model

  Ion type - Ions are point charges

  Bonding nature - purely ionic

  SO NO COVALENT CHARACTER CONSIDERED
• Experimental bond enthalpies:

  Model name: Born-Haber

  Ion type: Ions are polarisable

  Bonding nature: Covalent character

  COVALENT CHARACTER CONSIDERED
• If no covalent character present:

  Theoretical Lattice Formation Enthalpy = Experimental Lattice formation enthalpy
• If covalent character is present:

  Theoretical lattice formation enthalpy is less exothermic than the lattice formation calculated in the experimental model
• Ionic + covalent character = stronger bonding (more exothermic or endothermic lattice enthalpy)
• Covalent character:

  • If a positive ion is strongly polarising
  • The electron cloud in the negative ion becomes distorted
  • Some of the electron density is shared
  • Covalent character is present
• Entropy - The measure of disorder of the system
• Increase in entropy:

  solid (least) -> liquid -> gas (most)
• Entropy change = All products - All reactants
• Feasible - If the reaction can happen or not
• Positive gradient = -ΔS
• Negative gradient = ΔS
• ΔG = ΔH - TΔS

  ΔG - kJ mol-1

  ΔH - kJ mol-1

  T - Kelvin

  ΔS - kJ K-1 mol-1
• Why enthalpy of hydration becomes less exothermic down group 1:

  • Li+ has a smaller ionic radius than K+
  • Electrostatic attraction between Li+ and lone pair on oxygen increases
• What factor changed:

  A) Increased temperature