Chapter 19 Lattice Energies Flashcards
Define standard enthalpy change of lattice energy
The enthalpy change when 1 mole of an ionic lattice compound is formed from its gaseous ions, under standard conditions.
Standard enthalpy change of lattice energy exothermic or endothermic?
Exothermic, ionic bonds are formed, energy is released.
Define standard enthalpy change of atomisation
The enthalpy change when 1 mole of gaseous atoms is formed from its standard element, under standard conditions.
Standard enthalpy change of atomisation exothermic or endothermic?
Endothermic, bonds are broken, energy is absorbed.
Define first electron affinity
The enthalpy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous state -1 ions, under standard conditions.
First electron affinity exothermic or endothermic?
Exothermic, attraction is formed between added electron and nucleus.
Define second electron affinity
The enthalpy change when one mole of electron is added to one mole of gaseous state -1 ions to form one mole of gaseous state -2 ions, under standard conditions.
Second electron affinity exothermic or endothermic?
Endothermic, energy must be given to overcome electron repulsion.
Define standard enthalpy change of hydration
The enthalpy change when one mole of a specified gaseous ion dissolves in sufficient water to form a very dilute solution, under standard conditions.
Standard enthalpy change of hydration exothermic or endothermic?
Exothermic, ion-dipole bond is formed, energy is released.
Define standard enthalpy change of solution
The enthalpy change when 1 mole of an ionic solid dissolves in sufficient water to form a very dilute solution.
Standard enthalpy change of solution exothermic or endothermic?
Both.
Factors affecting ΔH⦵latt:
1) Ionic radius/Ion size…as ion size increases, ΔH⦵latt becomes less exothermic. Lower charge density so weaker ionic bonds.
2) Ionic charge…as ionic charge increases, ΔH⦵latt becomes more exothermic. Stronger ionic bonds formed.
Factors affecting first electron affinity:
1) Nuclear charge. Higher the nuclear charge, more exothermic the EA1.
2) Atomic radius. Higher the atomic radius, less exothermic the EA1.
3) Shielding. Higher the shielding, less exothermic the EA1.
Trend in first electron affinity down group 16 and 17 (and the exceptions)…
…first electron energy becomes less exothermic down the group. Oxygen and Fluorine are exceptions to this trend.
Why does oxygen and fluorine not fit the EA1 trend down the group?
They have a high electron density which causes electron repulsion and this will make the EA1 less exothermic as nuclear attraction is decreased.
Factors affecting ΔH⦵hyd:
1) Ionic charge. Higher the ionic charge, more exothermic the ΔH⦵hyd due to stronger ion-dipole bonds.
2) Ionic radius. Higher the ionic radius, less exothermic the ΔH⦵hyd due to lower charge density and weaker ion-dipole bonds.
Metal sulfate solubility down group 2
Decreases
Metal hydroxide solubility down group 2
Increases
An ionic compound is soluble if…
ΔH⦵hyd is higher than ΔH⦵latt. ΔH⦵hyd is sufficient to break ionic bonds.
An ionic compound is insoluble if…
ΔH⦵hyd is lower than ΔH⦵latt. ΔH⦵hyd is insufficient to break ionic bonds so energy must be given in, endothermic reaction, to dissolve the ionic compound.
Why do group 2 hydroxides get more soluble down the group?
ΔH⦵hyd and ΔH⦵latt both become less exothermic down the group but ΔH⦵latt becomes less exothermic more faster than ΔH⦵hyd, so at the bottom of the group, ΔH⦵hyd will be higher than ΔH⦵latt, making the salt soluble.
Why do group 2 sulfates get less soluble down the group?
ΔH⦵hyd and ΔH⦵latt both become less exothermic down the group but ΔH⦵hyd becomes less exothermic more faster than ΔH⦵latt, so at the bottom of the group, ΔH⦵hyd will be lesser than ΔH⦵latt, making the salt insoluble.
ΔH⦵sol = ?
ΔH⦵hyd - ΔH⦵latt.
(ΔH⦵sol + ΔH⦵latt = ΔH⦵hyd.)
If ΔH⦵sol is negative…
The ionic compound is more soluble.
If ΔH⦵sol is positive…
The ionic compound is less soluble.
How do group 2 carbonates and nitrates thermally decompose?
The metal cation polarises the carbonate or nitrate ion, weakening the C-O bond. If thermal energy is given in, the C-O bond breaks and a metal oxide and CO₂ is formed.
Thermal stability down group 2 for carbonates and nitrates?
Increases.
Why does thermal stability increase down the group?
The metal cation’s ionic radius increases down the group, which reduces the polarisation of the carbonate or nitrate ion. If the polarisation is less, the C-O bond is stronger so more heat energy is required to break the bond and thermally decompose the metal carbonate into metal oxide and CO₂.
