Thermodynamics

• enthalpy change of formation?

  the enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions
• first ionisation energy?

  the enthalpy change when 1 mole of 1+ ions is formed from 1 mole of gaseous atoms
• enthalpy of atomisation?


  the enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state
• bond dissociation enthalpy?


  enthalpy change when 1 mole of covalent bonds is broken into 2 gaseous atoms
• first electron affinity?

  the enthalpy change when each atom in one mole of gaseous atoms gain one electron to produce 1 mole of gaseous 1- ions
• lattice dissociation enthalpy?


  enthalpy change when one mole of an ionic crystal is broken into its constituent gaseous ions
• enthalpy of hydration?


  enthalpy change when 1 mole of gaseous ions become aqueous ions
• enthalpy of solution?


  enthalpy change when 1 mole of an ionic substance dissolves in a large enough amount of water to ensure that the dissolved ions are well separated and do not interact with one another
• order for Born - Haber cycle:

  • enthalpy of formation (downwards)
  • atomisation enthalpies x2
  • ionisation energy
  • electron affinity
  • lattice enthalpy of formation
• exothermic change:

  • enthalpy of formation
  • enthalpy of electron affinity
  • lattice formation enthalpy
• endothermic change:

  • enthalpy of atomisation
  • enthalpy of electron affinity
  • ionisation enthalpy
• endothermic change, energy increases
• exothermic change, energy decreases
• factors affecting strength of ionic bonding:

  • smaller ions
  • higher charged ions
• larger charge, greater attraction
• smaller size, stronger ionic bonding
• smaller difference between theoretical value and experimental value, closer to perfect ionic structure
• bigger difference between theoretical value and experimental value, more covalent character
• bigger lattice enthalpy, stronger ionic bonding
• more negative value for ΔHsol, more likely the solute is to dissolve
• if ΔHsol is large and positive, solute will not dissolve
• if ΔHsol is small and positive, solute may not dissolve, there is sufficient increase in entropy
• spontaneous reaction?


  occurs on its own record
• entropy?


  measure of how disorder in a system
• which state has the highest entropy?

  gas due to the random arrangement
• equation for entropy of system?

  entropy of products - entropy of reactants
• standard conditions?


  298K and 100kPa
• free energy change?


  measure used to predict whether a reaction is feasible
• feasible?


  once a reaction has started, it will carry onto completion, without any energy being supplied to it
• if ΔG is negative/equal to zero, reaction is forward and feasible
• if ΔG is positive then reaction isn't feasible
• ΔG = ΔH - TΔS
• y = mx + c

  • y is ΔG
  • m is -ΔS
  • c is ΔH
  • x is T
• exothermic reactions have negative ΔH value - heat energy is given out
• endothermic reactions have positive ΔH value - heat energy is absorbed
• second ionisation energy?


  enthalpy change when 1 mole of gaseous 2+ ions is formed from 1 mole of gaseous 1+ ions
• second electron affinity?


  the enthalpy change when each atom in one mole of gaseous 2- ions produce 1 mole of gaseous 1- ions
• lattice enthalpy?


  measure of ionic bond strength
• lattice enthalpy of formation?

  enthalpy change when 1 mole of a solid ionic compound is formed from its gaseous ions under standard conditions
• lattice enthalpy of dissociation?

  enthalpy change when 1 mole of a solid ionic compound is completely dissociated into it gaseous ions under standard conditions