3.1.8 Thermodynamics

• Bond dissociation enthalpy △H diss

  The enthalpy change when one mole of a covalent bond in a gaseous diatomic molecule is broken to form two mole of gaseous atoms

  e.g. Cl2(g) -> 2Cl(g)
• Enthalpy of atomisation of an element AH°at

  This is the enthalpy change when one mole of gaseous atoms are formed from an element in its standard state.

  e.g. Na(s) ->Na(g)

  Note that for elements that exist as diatomic gaseous molecules

  AH*at = ½ × AH* diss

  .e.g. ½ Cl2(g) -> CI(g)
• Ionisation Enthalpy △H*i

  The enthalpy changes when one mole of electrons is removed from a gaseous particle to form one mole of gaseous positive ions

  The first ionisation enthalpy applies to the removal of one mole of electrons from a mole of gaseous atoms

  e.g. Mg(g)->Mg+(g)+e-

  The second ionisation enthalpy applies to the removal of one mole of electrons from a mole of singly charged gaseous positive ions (+1) to form one mole of doubly charged gaseous positive ion (2+)

  e.g. Mg+(g)->Mg2+(g)+e-
• Electron affinity △H*ea

  This is the enthalpy change when one mole of electrons is added to one mole of gaseous atoms or negative ions.

  e.g. O(g)+ e- -> O-(g)

  The electron affinity is exothermic for most atoms because of the overall attraction between the positive nucleus and the added electron.

  Second electron affinities are endothermic as energy must be supplied to overcome the repulsion between the negative ion and extra electron

  e.g. O-(g)+ e- -> O2-(g)
• Lattic Enthalpy △H*L

  The enthalpy of lattice dissociation is the enthalpy change when one mole of a solid ionic compound is completely dissociated into its gaseous ions under standard conditions. This is always endothermic because energy must be supplied to overcome the ionic bonding

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

  The enthalpy of lattice formation is the enthalpy change when one mole of a solid ionic compound is formed from its gaseous ions under standard conditions. This exothermic s energy is released when the ionic bonds form.

  Cl-(g) + Na+(g)-> NaCl(s)
• Enthalpy of Sumblimation △H*sub

  This is the enthalpy change when one mole substance is converted from a solid to gaseous atoms

  e.g. Ca(s)-> Ca(g)

  The enthalpy of sublimation for a solid metallic element is the same as the enthalpy of atomisation
• Enthalpy of Vaporisation △H*vap

  The enthalpy changes when one mole substance is converted from a liquid to a gas.

  e.g. Br2(l)->Br2(g)

  Enthalpies of vaporisation are always endothermic because energy must be supplied to overcome the forces or bonds that hold particles together in a liquid.
• Enthalpy of Solution △H*sol

  This is the enthalpy change when one mol of a solid ionic compound is dissolved in excess water to form a dilute solution.

  e.g. NaCl(s)->Na+(aq) + Cl-(aq)
• Enthalpy of Hydration △H*hyd

  This is the enthalpy change when one mol of gaseous ions is dissolved in excess water to form a dilute solution

  e.g. Cl-(g)->Cl-(aq)

  Hydration enthalpies are always very exothermic as a result of the strong attraction between the ion and the polar water molecules.
• ΔH solution = ΔH lattice diss + ΔH hydration
• Gibbs free-energy change

  For a reaction to be feasible ΔG must be zero or negative

  ΔS is the entropy change unit ( J/mol/K ) calculated : ΔS = Products - Reactants. ΔS is the gradient

  ΔG is Gibbs free energy change unit ( kj/mol ) calculated: Δ G = Δ H - T Δ S

  ΔH is either the enthalpy of formation or combustion unit: kJ/mol and is also the + c in y = mx + c
• If ΔS is negative and ΔH is negative the reaction will be feasible if Δ H > T Δ S

  If ΔS is positive and ΔH is positive the reaction will be feasible if Δ H < T Δ S

  If ΔS is positive and ΔH is negative the reaction will always be feasible

  If ΔS is negative and ΔH is positive the reaction will never be feasible