Topic B5 - Energetics II
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Lattice Enthalpy (def)
Formation of 1 mole of an ionic compound from its gaseous ions under standard conditions
Atomisation Enthalpy (def)
Formation of 1 mole of gaseous atom from elements in its standard states under standard conditions
1st Electron Affinity (def)
One electron is added to a gaseous atom to form 1 mole of gaseous 1- ion
2nd Electron Affinity (def)
One electron is added to a gaseous 1- ion to form 1 mole of gaseous 2- ion
Enthalpy of Hydration (def)
Enthalpy change when one mole of gaseous ions interact with/ disolve in water molecules to form hydrated aqueous ions
Enthalpy of solution (def)
enthalpy change when one mole of a solute dissolves in a solvent
Is lattice enthalpy exothermic or endothermic?
exothermic
Is atomisation enthalpy exothermic or endothermic?
endothermic
Is 1st electron affinity endothermic or exothermic?
exothermic
Is 2nd electron affinity endothermic or exothermic?
endothermic
Is enthalpy of hydration exothermic or endothermic?
exothermic
Why can MgCl3 not form from the elements magnesium and chlorine?
• third ionisation energy of magnesium is very positive/ endothermic • not compensated for by more negative electron affinities/ lattice enthalpy
What does the experimental lattice enthalpy and the theoretical lattice enthalpy being similar tell you?
there is very little covalency in the compound (i.e. positive ion is not very polarising and negative ion is not very polarisable)
Why is it difficult to predict whether the enthalpy change of solution becomes more exothermic or less exothermic down the group from MgF2 to MgI2?
• halide ion radii gets larger down the group • hence LE is less exothermic down group (as halide ions has less attraction to Mg2+) • hydration enthalpy is less exothermic down group (halide ion has less of an attraction to polar water molecules) • difficult to predict whether LE or hydration enthalpy has a bigger effect