lattice enthalpy

• a lattice is a regular arrangement of particles in a solid, they can be ionic, covalent or metallic
• standard conditions 298K and 100kPa
• lattice enthalpy is the enthalpy change when one mole of a solid ionic compound is formed from its component gaseous ions under standard conditions
• standard enthalpy change of formation is the enthalpy change that occurs when one mole of compound is formed from its elements in their standard states under standard conditions
• the enthalpy change of atomisation energy is the enthalpy change that takes place when one mole of gaseous atoms is formed from the element in its standard state under standard conditions
• bond dissociation enthalpy is the energy needed to break one mole of a specific bond by homolytic fission, everything being in the gas state
• first ionisation energy is the energy required to remove one mole of the most loosely held electrons from one mole of gaseous atoms to produce one mole of gaseous ions each with a charge of 1+
• first electron affinity is the energy released when one mole of gaseous atoms each acquire an electron to form one mole of gaseous 1- ions
• the lattice formation enthalpies ae always negative (exothermic)
• enthalpy of atomisation is always endothermic (positive)
• first ionisation energy is always endothermic (positive)
• first electron affinity is always exothermic (negative)
• enthalpy change of solution is the enthalpy change when one mole of an ionic compound dissolves in water to give a solution of infinite dilution
• enthalpy change of hydration is the enthalpy change when one mole of gaseous ions dissolve in water to give an infinitely dilute solution of aqueous ions.
• hydration enthalpies are always negative (exothermic) as when the ions are surrounded by water molecules strong ion-dipole forces are formed
• ion-dipole attractions are electrostatic attractions between ions and polar water molecules
• lattice enthalpy depends on the charge on the ion as the greater the charge the greater the lattice enthalpy due to the stronger electrostatic bond, it also depends on the size of the ion as the bigger the ion the lower the lattice enthalpy as there is a larger charge distribution.
• lattice enthalpy cannot be measured directly it is found using a born-haber cycle
• the more exothermic the lattice enthalpy the higher the melting point of the ionic solid, so the stronger the ionic bonds in the compound
• diagram of born-haber cycle:

  A) enthalpy change of formation

  B) enthalpy change of atomisation

  C) first ionisation energy

  D) enthalpy change of atomisation

  E) electron affinity

  F) lattice enthalpy

  G) ionic compound

  
• in a born-haber cycle the arrow for first electron affinity points down as it is exothermic, if there is a second electron affinity the arrow points up as it is endothermic
• ionic compounds are insoluble in non-polar solvents like cyclohexane, but are soluble in polar solvents such as water
• for ionic compounds to dissolve energy is required to separate the ions and break up the lattice, this energy comes from solvation
• enthalpy change of solution depends on lattice enthalpy and hydration energy so if a substance dissolves it is exothermic as lattice enthalpy is smaller in magnitude than enthalpy of hydration
• steps of dissolving ionic compounds:

  1. breaking the ionic bonds which depends on the lattice enthalpy
  2. the formation of ion-dipole bonds which depends on hydration energy
• if the ion-dipole interactions are stronger or greater in magnitude than the ionic bonds that must be broken then the solid is likely to be soluble
• factors affecting hydration energy:

  • size of the ions- ion-dipole forces are strongest when the ions are small because the electric field intensity around smaller ions is greater
  • charge of the ions- the more highly charged an ion the more energy it releases
• entropy must be considered when trying to decide if a substance will dissolve and form a solution
• calculation for enthalpy change of solution:

  $\Delta H=$$-\Delta H\left(LE\right)+$$\Delta Hhyd\left(cations\right)+$$\Delta Hhyd\left(anions\right)$
• enthalpy of solution by itself is not a great indicator of solubility also need to take into account entropy and Gibbs free energy
• highly charged ions attract water molecules and order them, the entropy of the ions increases as they have broken free from the lattice but the entropy of water molecules massively decreases making dissolving unfavourable as the ions cannot offset the decrease in entropy of the solvent molecules
• if the effects of enthalpy and entropy are antagonistic temperature becomes important