Thermodynamics

• Hess’s law definition

  Enthalpy change is independent of the route taken
• Standard enthalpy of formation (ΔH°f)


  Enthalpy change when one mol of a compound is formed from its constituent elements in there standard states under standard conditions
• Standard Enthalpy of atomisation (ΔH°at)


  One mole of GASEOUS atoms are formed from consituent elements in standard states under standard conditions
• Is atomisation endo (takes up E) or exo ( releases E)

  Endothermic as bonds are broken therefore energy is needed
• Ionisation energy (ΔH°IE)


  One mole of e- removed from one mole of gaseous atoms

  Endo bc e- is attracted to +ve nucleus
• Electron affinity (ΔH°EA)

  One mole of gaseous atoms gains one mole of electrons
• Why is the first EA exo and the second is endo ?

  • Eg, Cl(g) +e- —> Cl- is exo as e- is attracted to the positive Cl (1st)
  • However O-(g) +e- —> O^2- (g) is endo as they are both -ve and so they repel
• Lattice formation Enthalpy (ΔH°LF)

  • Enthalpy change when one mole of a solid ioniccompound is formed from its free gaseous ions under standard conditions
  • exothermic as opposite charged ions attract
• Lattice dissociation Enthalpy (ΔH°LD)


  Enthalpy change when one mol of solid ionic compounds are broken down into free gaseous ions under standard conditions

  Opposite of ΔH°LF therefore endo
• Enthalpy of hydration (ΔH°hyd)


  Enthalpy change when one mole of gaseous ions —> aqueous ions .

  Exo as H20 is polar and has both positive and negative dipoles to cause attraction to either ion

  
• Chemical reactions occur when particles of substances collide
• For a reaction to occur successfully, collisions must have energy greater than or equal to the activation energy and the particle orientation must be correct
• Reaction conditions can be altered to provide particles with more energy, increasing the likelihood of a collision occurring with sufficient energy to react and therefore increasing the rate of reaction
• Not all molecules in a substance have the same amount of energy; their energies are distributed in a pattern called the Maxwell-Boltzmann distribution
• Changing reaction conditions alters the shape of the Maxwell-Boltzmann distribution curve, affecting the number of particles with energy greater than the activation energy
• When a substance is heated, thermal energy is converted to kinetic energy, causing molecules to move faster and further, leading to more frequent collisions with greater energy and an increased rate of reaction
• Increasing the reaction temperature results in a shift of the Maxwell-Boltzmann distribution to the right, with a greater proportion of molecules having energy greater than or equal to the activation energy
• Increasing the concentration of a sample or pressure leads to molecules being packed closer together, increasing the likelihood of collisions with energy greater than the activation energy and thus increasing the rate of reaction
• A catalyst increases the rate of reaction by providing an alternative reaction path with a lower activation energy, without being used up in the reaction
• The Maxwell-Boltzmann distribution curve remains unchanged in shape with a catalyst, but the position of the activation energy is shifted to the left, allowing a greater proportion of molecules to have sufficient energy to react
• Enthalpy of atomisation

  The enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state
• Enthalpy of sublimation

  The enthalpy change for a solid metal turning to gaseous atoms, numerically the same as the enthalpy of atomisation
• Enthalpy of sublimation


  • Na (s) -> Na(g) [ΔsubH = +148 kJ mol-1]
• Bond dissociation enthalpy (bond energy)


  The standard molar enthalpy change when one mole of a covalent bond is broken into two gaseous atoms (or free radicals)
• Bond dissociation enthalpy

  • Cl2 (g) -> 2Cl (g) [ΔdissH = +242 kJ mol-1]
  • CH4 (g) -> CH3 (g) + H(g) [ΔdissH = +435 kJ mol-1]
• For diatomic molecules

  The ΔdissH of the molecule is the same as 2x ΔatH of the element
• First ionisation enthalpy

  The enthalpy change required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous ions with a +1 charge
• Second ionisation enthalpy

  The enthalpy change to remove 1 mole of electrons from one mole of gaseous 1+ ions to produce one mole of gaseous 2+ ions
• First electron affinity

  The enthalpy change when 1 mole of gaseous atoms gain 1 mole of electrons to form 1 mole of gaseous ions with a -1 charge
• First electron affinity

  • O (g) + e- -> O- (g) [Δea 1H = -141.1 kJ mol-1]
• Second electron affinity

  The enthalpy change when one mole of gaseous 1- ions gains one electron per ion to produce gaseous 2- ions
• Second electron affinity

  • O- (g) + e- -> O2- (g) [Δea 2H = +798 kJ mol-1]
• Enthalpy of lattice formation

  The standard enthalpy change when 1 mole of an ionic crystal lattice is formed from its constituent ions in gaseous form
• Enthalpy of lattice formation

  • Na+(g) + Cl- (g) -> NaCl (s) [ΔLattH = -787 kJ mol-1]
• Enthalpy of lattice dissociation

  The standard enthalpy change when 1 mole of an ionic crystal lattice form is separated into its constituent ions in gaseous form
• Enthalpy of lattice dissociation

  • NaCl (s) -> Na+(g) + Cl- (g) [Δ LattH = +787 kJ mol-1]
• Enthalpy of hydration

  The enthalpy change when one mole of gaseous ions become aqueous ions
• Enthalpy of hydration

  • X+ (g) + aq -> X+ (aq) [For Li+ ΔhydH = -519 kJ mol-1]
  • X- (g) + aq -> X- (aq) [For F- ΔhydH = -506 kJ mol-1]
• Enthalpy of solution

  The standard enthalpy change when one mole of an ionic solid dissolves in a large enough
• This process always gives out energy (exothermic, -ve) because bonds are made between the ions and the water molecules