Energetics + Thermodynamics

• Energy diagrams

  A) exothermic reaction

  B) endothermic reaction

  C) activation energy

  D) energy released

  E) activation energy

  F) energy absorbed

  G) reactants

  H) products

  I) reactants

  J) products

  
• thermochemistry = study of heat changes during chemical reactions

  enthalpy change (ΔH) = heat change at constant pressure (kJ/mol)

  mean bond enthalpies = energy required to break 1 mole of a covalent bond in the gas state averaged over a variety of different molecules -> reactants - products

  exothermic = heat energy released into surroundings causing temperature to increase (make bonds)

  • products have less energy than reactants -> given out

  endothermic = heat energy is taken in from surroundings causing temperature to decrease (break bonds)

  • products have more energy than reactants -> taken in
• standard conditions = done under specific conditions


  temperature = 298K (25°C)

  pressure = 100kPa

  concentration = 1 mol/dm3
• standard definitions

  standard enthalpy of formation = Δ͐fH = the enthalpy change when 1 mole of a substance is formed from it's elements with all reactants & products in their standard states under standard conditions

  standard enthalpy of combustion = Δ͐cH = enthalpy change when 1 mole of a substance is burned completely in excess oxygen with all reactants and products in standard states under standard conditions

  standard enthalpy of neutralisation = the enthalpy change when 1 mole of water is formed from the reaction between an acid and an alkali/base under standard conditions
• temperature = average kinetic energy of all particles in a system

  • particles move faster -> average KE increases -> temp increases
  • temp independent of the no. of particles present
  • measured with a thermometer

  heat = measure of total energy of all particles present in a given amount of substance

  • is dependent on the amount of substance present
  • heat moves from high temp to low temp
• enthalpy formula - no instrument to measure heat


  q = mcΔT

  q = heat energy change (J) m = mass (kg) ΔT = change in temp (°C/K)

  ΔH = -q/n

  q = heat energy change (J) n = no. of moles ΔH = enthalpy change per mole

  assumptions made in calculations:

  • no loss of heat energy to the surroundings
  • calculated energy will be lower than actual value
  • no change in mass
  • some mass will evaporate
  • container has negligible heat capacity
  • assume complete combustion
• calorimetry = technique used to find the heat energy change that occurs during a reaction

  1. transfer 25cm3 of 2 mol/dm3 HCl to measuring cylinder
  2. transfer 25cm3 of 2 mol/dm3 NaOH to a clean, dry polystyrene cup & place inside a beaker
  3. stir the NaOH with thermometer & record temp to 1dp. Start stopwatch
  4. every minute for 3 minutes stir the solution, measure the temp & record.
  5. at 4th minute, add the 25cm3 of HCl from plastic cup
  6. stir mixture but DO NOT RECORD THE TEMP
  7. continue to stir the mixture & measure the temp at the 5th minute & then every minute for a further 8 minutes.
• Hess' Law

  'overall energy change for a reaction is independent of the route taken'

  arrows up -> standard enthalpy of formation

  arrows down -> standard enthaply of combustion
• first ionisation energy, Δ(IE1)H = enthalpy required to remove an electron from 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions (endothermic)

  F(g) -> F+(g) + e-

  second ionisation energy Δ(IE2)H = enthalpy required to remove an electron from 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions

  F+(g) -> F2+(g) + e-
• first electron affinity, Δ(EA1)H = enthalpy when 1 mole of gaseous atoms gain an electron to form 1 mole of gaseous 1- ions (exothermic)

  second electron affinity, Δ(EA2)H = enthalpy change when 1 mole of gaseous 1- ions gain an electron to form 1 mole of gaseous 2- ions

  • every successive affinity is endothermic except 1st affinity - repulsion between -ve charges
• enthalpy of atomisation, Δ(at)H = enthalpy change when 1 mole of gaseous atoms are formed from an element in its standard state

  Ni(s) -> Ni(g)

  1/2F2(s) -> F(g)

  1/4P4(s) -> P(g)

  bond disassociation enthalpy, Δ(BDE)H = enthalpy change when 1 mole of a covalent bond is broken in the gas state

  F2(g) -> 2F(g)

  O2(g) -> 2O(g)
• lattice enthalpy, Δ(LE)H = enthalpy change when 1 mole of an ionic compound is formed from its ions in the gas state (exothermic)

  Na+(s) + Cl-(g) -> NaCl(s)

  lattice dissociation enthalpy, Δ(LDE)H = enthalpy change when 1 mole of an ionic compound is broken down into its gaseous ions (endothermic)

  NaCl(s) -> Na+(s) + Cl-(g)
• The perfect ionic model = theoretically predicting the lattice enthalpy of an ionic compound


  assumes: bonding is 100% ionic, ions are perfectly spherical point charges

  polarisation/distortion of electron density -> electrons almost being shared -> has a degree of covalency

  if actual & theoretical value for LE are similar -> assume bonding almost perfectly ionic

  large difference in values -> degree of covalency

  smaller ions have higher electron density -> higher LE as electrostatic attraction are stronger between +ve & -ve charges
• Born-Haber cycles

  down arrow = exothermic

  up arrow = endothermic

  if you go in same direction as arrow -> add otherwise subtract
• entropy, S = a measure of disorder within a system or substance



  • randomness of components in a reaction can determine if it's feasible or spontaneous
  • factors that affect entropy & in priority:
  • state (solid < liquid < gas)
  • no. of moles of substance
  • size of molecule

  A) S(products)

  B) S(reactants)

  
• enthalpy of hydration, Δ(hyd)H = enthalpy change when 1 mole of gaseous ions are dissolved in water (exothermic)

  partial charges in H2O interact with ions

  enthalpy of solution, Δ(sol)H = enthalpy change when 1 mole of a solid ionic compound is dissolved is water

  can be exothermic or endothermic depending on strength of attraction
• Gibbs free energy = determines if a reaction is feasible using enthalpy & entropy values


  • for a reaction to be feasible -> ΔG has to be negative
  • some reactions only occur at 1 temperature
  • to find temperature at which a reaction is feasible at, let ΔG = 0

  ΔG = ΔH - TΔS

  A) always -ve

  B) -ve when TΔS>ΔH -> feasible at high temps

  C) -ve when TΔS<ΔH -> feasible at low temps

  D) always +ve

  E) feasible at all temps

  F) feasible at high temps

  G) feasible at low temps

  H) not feasible at any temp

  
• ΔG graphically

  ΔG = -ΔS(T) + ΔH -> y = mx + c

  A) never feasible

  B) always feasible

  C) feasible at low temps

  D) feasible at high temps

  
• feasible v spontaneous

  spontaneous = has sufficiently low activation energy for the reaction to take immediately

  feasible = ΔG <= 0

  ΔG also used for change of state 'reactions'