P- B-H Cycles =)

Question 1

Enthalpy of hydration of Mg2+ ions: −1920
Enthalpy of hydration of Na+ ions: −406
/
Explain why there is a difference between the hydration enthalpies of the magnesium and sodium ions. (2)

Answer 1

• Mg2+ - smaller + ↑ charged
• attracts H2O ↑ strongly

Question 2

Enthalpy of hydration-
Mg2+ (g): -1920
Ca2+ (g): -1650
Suggest why the value for Ca2+ is less exothermic than that of Mg2+. (2)

Answer 2

• Ca2+ - larger ion
• weaker attraction to (Oδ- in) water

Question 3

Define the term electron affinity for chlorine. (2)

Answer 3

• enthalpy change for formation of 1 mole of Cl- ions from chlorine atoms
• atoms & ions in gaseous states

Question 4

Suggest why aluminium hydroxide is insoluble in water. (1)

Answer 4

• high lattice enthalpy
• // strong covalent bonds

Question 5

Define bond dissociation enthalpy as applied to chlorine. (2)

Answer 5

• enthalpy change to break the bond in 1 mol of Cl2 molecules
• to form 2 mols of gaseous chlorine atoms

Question 6

A calculation of the enthalpy of lattice formation of silver iodide based on a perfect ionic model gives a smaller numerical value than the value calculated (-869kJmol-1).
Explain this difference. (2)

Answer 6

• AgI contains covalent character
• forces holding the lattice tgt are stronger

Question 7

Give an equation that represents the process when the standard enthalpy of atomisation of iodine is measured. (1)

Answer 7

1/2 I2 (s) –> I (g)

1 mol of gaseous atoms formed!

Question 8

The enthalpy of lattice formation for caesium iodide obtained by experiment is -585kJmol-1.
The value obtained by calculation using the perfect ionic model is –582 kJmol–1.
Deduce what these values indicate about the bonding in caesium iodide. (1)

Answer 8

almost perfectly ionic

Question 9

Units for molar enthalpy of vaporisation (∆Hvap)

Answer 9

kJmol-1