Thermodynamics

• enthalpy of formation - enthalpy change when one mole of a substance is formed from its constituent elements with all substances in their standard states under standard conditions, exothermic reaction
• enthalpy of combustion:

  • enthalpy change when one mole of a substance undergoes complete combustion in oxygen with all substances in standard states under standard conditions
  • exothermic
• enthalpy of neutralisation:

  • enthalpy change when one mole of water is formed in a reaction between an acid and an alkali under standard conditions
  • exothermic
• 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
• second ionisation energy:

  • enthalpy change when each atom in one mole of gaseous 1+ atoms loses one electron to form one mole of gaseous 2+ ions
  • endothermic
• 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 for many non-metals
• 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
• enthalpy of atomisation:

  • enthalpy change when one mole of gaseous atoms is formed from an element in its standard state
  • endothermic
• hydration enthalpy:

  • enthalpy change when one mole of gaseous ions become hydrated (dissolved in water)
  • exothermic
• 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
  • can be exothermic or endothermic
• bond dissociation enthalpy:

  • enthalpy change when one mole of covalent bonds is broken in the gaseous state
  • endothermic
• lattice enthalpy of formation:

  • enthalpy change when one mole of a solid ionic compound is formed from its constituent ions in the gas phase
  • exothermic
• 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
• enthalpy of vaporisation:

  • enthalpy change when one mole of a liquid is turned into a gas
  • endothermic
• enthalpy of fusion:

  • enthalpy change when one mole of a solid is turned into a liquid
  • endothermic
• theoretical lattice enthalpies can be calculated from data assuming there is a perfectly ionic model
• perfect ionic model:

  • ions are perfectly spherical
  • charge is evenly distributed in the sphere
• the actual experimental value for lattice enthalpy is different to the theoretical value:

  • experimental compound is not perfectly ionic and has some covalent character
  • the positive ion distorts the charge distribution in the negative ion
  • the positive ion polarises the negative ion
  • more polarisation = more covalent character
• lattice enthalpy value tells us how much a substance is purely ionic
• the bigger the difference between the experimental and theoretical lattice enthalpies, the more polarisation between the ions which results in greater covalent character for the compound
• • entropy is the measure of disorder in a system
  • disorder is the number of ways that energy is shared out between particles
• more disorder = higher level of entropy
• solids have the lowest level of disorder/entropy as the particles are fixed into rows
• increasing disorder/entropy:

  1. solids
  2. liquids
  3. gases
• • the number of particles affects the entropy change
  • if a reaction is in the same state but more moles are produced, the entropy increases
• entropy is measured in J mol-1 K-1
• 2nd law of thermodynamics: over time entropy will naturally increase
• 3rd law of thermodynamics: the entropy of a substance is zero at absolute zero and increases with temperature
• calculating entropy change:

  • ΔS = [sum products] - [sum reactants]
• entropy increases when:

  • moles of gas increases
  • temperature increases
  • ionic solids dissolve in water
• • in reactions where entropy increases, ΔS is positive
  • in reactions where entropy decreases, ΔS is negative
• if there is an equal number of gas moles on both sides of the equation, ΔS is close to zero
• enthalpy of solution calc:

  • gas ions to ionic solid is lattice enthalpy of formation
  • gas ions to dissolved ions is enthalpy of hydration
  • ionic solid to dissolved ions is enthalpy of solution
  • enthalpy of solution = hydration enthalpies - lattice enthalpy of formation
• • entropy change = products - reactants
  • units = J K-1 mol-1
• Gibbs free energy tells us if a reaction is feasible or not
• a reaction is feasible if Gibbs energy is negative or zero
• even if a reaction is calculated to be feasible, a reaction may not be observed:

  • activation energy too high
  • rate of reaction too slow
• gibbs calc:

  A) enthalpy change in kJ mol-1

  B) temperature in K

  C) entropy change in J K-1 mol-1

  
• entropy graph:

  • entropy increases with temperature
  • there are big changes in entropy in state changes - melting/boiling
  • the entropy change from liquid to gas is greater than for solid to liquid due to the larger amount of disorder in gases compared to liquids and solids
• entropy graph:

  A) boiling

  B) melting

  C) boiling

  D) melting

  E) vibrate more

  F) disorder increases