Entropy and Enthalpy

• Enthalpy change of lattice formation, Δlatt H θ
• Enthalpy change of lattice breaking, Δlatt H θ
• Enthalpy change of formation, Δf H θ
• Enthalpy change of hydration, Δhyd H θ
• Enthalpy change of solution, Δsol H θ
• Enthalpy change of atomisation, Δat H θ
• Bond dissociation enthalpy, Δbond H
• Enthalpy change of solution depends on two factors: enthalpy change of lattice breaking, which is endothermic, and enthalpy change of hydration, which is exothermic.
• If ΔG < 0, a reaction will occur spontaneously.
• The elements are turned into gaseous atoms, this is the enthalpy of atomisation, which is endothermic.
• If the enthalpy change of solution, Δ sol H θ , is negative, an ionic compound will be soluble, with a more negative enthalpy change of solution leading to a more soluble compound.
• Chemical reactions tend to occur when the enthalpy change is negative, i.e., they are exothermic, but some endothermic reactions also occur.
• The gaseous atoms are turned to cations (IE) or anions (EA)

  • ionisation energy - endothermic
  • electron affinity - exothermic.
• Entropy is represented by the symbol S and has the unit J K –1 mol –1 , it is the disorder in a system, or a measure of the freedom possessed by particles within a system.
• The temperature when a reaction becomes feasible is when the value of ΔG changes from positive to negative, at this point ΔG = 0.
• The gaseous ions are combined to form a solid, lattice enthalpy, which is exothermic.
• The formation of ionic compounds from their elements depends on the enthalpy changes involved in a number of steps, which are combined in an energy cycle known as the Born-Haber cycle.
• Particles that can move freely in any direction have a much higher entropy than particles that are constrained, S(gas) > S(liquid) > S(solid).
• If the enthalpy change of hydration is greater than the enthalpy change of lattice breaking, the compound will dissolve.
• The balance between enthalpy, entropy and temperature is the Gibbs free energy change, ΔG.