thermodynamics

• Enthalpy of atomisation

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

  • Na (s) → Na(g) [atH = +148 kJ mol-1]
  • ½ O2 (g) → O (g) [atH = +249 kJ mol-1]
• 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) [subH = +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) [dissH = +242 kJ mol-1]
  • CH4 (g) → CH3 (g) + H(g) [dissH = +435 kJ mol-1]
• For diatomic molecules the dissH of the molecule is the same as 2x atH 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 produces one mole of gaseous 2+ ions
• First electron affinity


  The enthalpy change that occurs when 1 mole of gaseous atoms gain 1 mole of electrons to form 1 mole of gaseous ions with a –1 charge
• The first electron affinity is exothermic for atoms that normally form negative ions
• Second electron affinity


  The enthalpy change when one mole of gaseous 1- ions gains one electron per ion to produce gaseous 2- ions
• The second electron affinity for oxygen is endothermic because it take energy to overcome the repulsive force between the negative ion and the electron
• 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 dissociation

  The standard enthalpy change when 1 mole of an ionic crystal lattice form is separated into its constituent ions in gaseous form
• Note the conflicting definitions and the sign that always accompanies the definitions
• Enthalpy of hydration


  The enthalpy change when one mole of gaseous ions become aqueous ions
• Enthalpy of hydration is always exothermic because bonds are made between the ions and the water molecules
• Enthalpy of solution


  The standard enthalpy change when one mole of an ionic solid dissolves in a large enough amount of water to ensure that the dissolved ions are well separated and do not interact with one another
• Enthalpy change of formation

  The energy transferred when 1 mole of the compound is formed from its elements under standard conditions (298 K and 100 kpa), all reactants and products being in their standard states
• The lattice enthalpy cannot be determined directly. We calculate it indirectly by making use of changes for which data are available and link them together in an enthalpy cycle the Born-Haber cycle
• Born-Haber cycle


  1. Mg (s) + Cl2(g) → MgCl2 (s)

  2. Na (s) + ½ Cl2(g) → NaCl (s)

  3. Ca (s) + ½ O2(g) → CaO (s)
• Factors affecting lattice enthalpies

  • The sizes of the ions
  • The charges on the ion
• The lattice enthalpies become less negative down any group
• Perfect ionic model


  Theoretical lattice enthalpies assumes a perfect ionic model where the ions are 100% ionic and spherical and the attractions are purely electrostatic
• When the negative ion becomes distorted and more covalent we say it becomes polarised. The metal cation is called polarising if it polarises the negative ion
• The Born Haber lattice enthalpy is the real experimental value
• When a compound shows covalent character, the theoretical and the born Haber lattice enthalpies differ. The more the covalent character the bigger the difference between the values
• Why does calcium chloride have the formula CaCl2 and not CaCl or CaCl3?
• Born-Haber cycle for CaCl, CaCl2, CaCl3

  Calculating the enthalpy of formation for each to determine the most stable form
• Free-energy change (G) and entropy change (S)

  Spontaneous processes can be endothermic due to an increase in entropy
• Entropy, S˚

  A description of the number of ways atoms can share quanta of energy. Higher entropy means more disordered systems
• Substances with higher entropy

  • Elements
  • Compounds
  • Simpler compounds
  • Complex compounds
  • Pure substances
  • Mixtures
  • Solids
  • Liquids
  • Gases
• ∆H

  Enthalpy change
• A reaction that is exothermic will result in products that are more thermodynamically stable than the reactants
• This is a driving force behind many reactions and causes them to be spontaneous (occur without any external influence)
• Some spontaneous reactions, however, are endothermic
• Entropy

  A description of the number of ways atoms can share quanta of energy