chemical energetics

• Standard conditions: temperature of 298 K, pressure of 100 kPa, solution concentrations of 1 mol dm^-3. All products and reactants are in their standard states
• Lattice formation enthalpy: enthalpy change when one mole of a substance is formed from its gaseous ions
• Electron Affinity:

  • First electron affinity: energy released when one mole of gaseous atoms each gains an electron to form one mole of 1- ions
  • Second electron affinity: energy required to add one electron to each ion in one mole of gaseous 1- ions to form one mole of gaseous 2- ions
• Born-Haber Cycles:

  • Used to calculate lattice enthalpy
  • Enthalpy changes involved: lattice enthalpy of formation or dissociation, enthalpy change of atomisation, enthalpy change of formation, first ionisation energy, first electron affinity
• Entropy Change, ΔS:

  • Entropy is a measure of the degree of disorder in a system
  • Entropy increases from solid to liquid to gas, and aqueous substances have higher entropy than solids
  • Entropy change of a reaction can be calculated using the equation: ΔS = ΣSϴ(products) - ΣSϴ(reactants)
• Gibbs Free Energy Change, ΔG:

  • Calculated using the equation: ΔG = ΔH - TΔS
  • A reaction is spontaneous when ΔG is less than or equal to 0
  • To find the minimum temperature for spontaneity: ΔH - TΔS < 0
• When ΔH is negative and ΔS is positive, the reaction is spontaneous at all temperatures
• Lattice dissociation enthalpy: enthalpy change when one mole of an ionic compound is broken down to form its gaseous ions
• Factors affecting lattice enthalpy: ionic charge and ionic radius
• Standard enthalpy change of atomisation (ΔHӨat): enthalpy change when one mole of gaseous atoms are formed from the element in its standard state
• Standard enthalpy change of hydration (ΔHӨhyd): enthalpy change when one mole of gaseous ions dissolves in water
• Standard enthalpy change of solution (ΔHӨsol): enthalpy change when one mole of an ionic solid dissolves in water
• When ΔH is positive and ΔS is negative, the reaction is never spontaneous
• When ΔH and ΔS are positive, the reaction is spontaneous at high temperatures
• The entropy of a system increases as it becomes more disordered or random.
• When ΔH and ΔS are both negative, the reaction is spontaneous at low temperatures
• A reaction with a large negative value of ΔG will be highly exothermic and have a very small equilibrium constant
• A reaction with a large positive value of ΔG will be endothermic and have a very large equilibrium constant
• Increasing temperature causes increased molecular motion and disorder, leading to higher entropy values.
• Increasing temperature favors reactions that increase disorder (increase entropy) and decrease favorable interactions between molecules (decrease enthalpy)
• At standard conditions, the free energy change (ΔG°) can be calculated using the equation: ΔG° = ΔH° - TΔS°
• Reactions involving condensation (loss of water) decrease entropy due to loss of freedom of movement.