O abc: Enthalpies of ionic lattices & solutions

Question 1

How is lattice enthalpy represented symbolically?

Answer 1

Δ_LEH

Question 2

Define lattice enthalpy.

Answer 2

The enthalpy change when one mole of ionic solid is formed by the coming together of gaseous anions and cations.

Usually 1 mole of each ion, except in cases like Mg(OH)₂

Question 3

Δ_LEH:

• Exothermic or endothermic?
• • or -?
• Which way does the arrow point?

Answer 3

• Exothermic (ionic bonds formed)
• Negative
• Gaseous ions to ionic solid (down)

The stronger the lattice, the more exothermic Δ_LEH is

Question 4

What factors increase lattice enthalpy?

Answer 4

Stronger ionic forces due to increased charge density:

• Increased charge
• Decreased ionic radii

Question 5

State and explain which of MgO and Na₂O has the more exothermic lattice enthalpy.

Answer 5

MgO

• Mg²⁺ has the same ionic radius as, but has a higher charge than, Na⁺
• So Mg²⁺ has a higher charge density
• So attracts O^2- more strongly
• So MgO has the more exothermic lattice enthalpy

Question 6

State the bonds broken and made when an ionic substance dissolves in water.

Answer 6

Broken: ionic + hydrogen bonds between water molecules

Made: ion-dipole

Question 7

Draw diagrams of the hydrated ions in NaCl_(aq)

Answer 7

Draw 4 water molecules around each ion

Question 8

What factor increases the size of a hydrated ion?

Answer 8

Increased charge density:

• Increased charge
• Decreased ionic radii

The higher an ion’s charge density, the more water molecules it attracts.

Question 9

How is the enthalpy change of hydration represented symbolically?

Answer 9

Δ_hydH

Δ_hydrationH

Question 10

Define the enthalpy change of hydration.

Answer 10

The enthalpy change when one mole of gaseous ions form a very dilute solution in water.

Values quoted depend on concentration of solution

Question 11

Δ_hydH:

• Exothermic or endothermic?
• • or -?

Answer 11

• Exothermic (gaseous ions form bonds with water)
• Negative

Question 12

What factors make the value of Δ_hydH more exothermic?

Answer 12

Higher charge density:

• Smaller ionic radii
• Greater charge

Question 13

What is the equivalent of Δ_hydH when considering solvents other than water?

Answer 13

Enthalpy of solvation: Δ_solvH

Question 14

How is enthalpy change of solution represented symbolically?

Answer 14

Δ_solutionH

Δ_solnH

Question 15

Define the enthalpy change of solution.

Answer 15

The enthalpy change when one mole of solute dissolves in water to form a very dilute solution.

Question 16

Write an expression which could be used to find the value of Δ_solnH experimentally.

Answer 16

Δ_solnH = Δ_hydH (cation) + Δ_hydH (anion) - Δ_LEH

Break ionic lattice into gaseous ions, then form solution

Question 17

Draw an enthalpy level diagram in which the value of Δ_solnH is negative.

Answer 17

Question 18

Draw an enthalpy level diagram in which the value of Δ_solnH is very positive.

Answer 18

Question 19

An endothermic value of Δ_solnH means that dissolving appears to be energetically unfavourable. If the value is small enough, the reaction can still happen. Why?

Answer 19

If the entropy of the system increases, the solute dissolves despite the change being endothermic.

Question 20

An ionic solid was placed in a non-polar solvent. Draw an enthalpy level diagram to represent this.

Answer 20

Question 21

Explain why the solubility of group 2 sulphates and carbonates decreases down the group.

Answer 21

• Solubility is determined by Δ_solnH = Δ_hydH - Δ_LEH
• Down group, ions remain same size but become larger, so charge density decreases
  
  • Δ_LEH + Δ_hydH both become less exothermic
• Δ_LEH decreases at lower rate because:
  
  • Cations become larger, approaching size of anions; lattice becomes more tightly packed
  • Electrostatic attraction increases
• Overall effect is that Δ_solnH becomes less exothermic, so solubility decreases

Question 22

Explain why the solubility of group 2 hydroxides increases down the group.

Answer 22

• Solubility is determined by Δ_solnH = Δ_hydH - Δ_LEH
• Down group, ions remain same size but become larger, so charge density decreases
  
  • So Δ_LEH + Δ_hydH both become less exothermic
• Δ_LEH decreases at higher rate because:
  
  • Cations become increasingly larger than OH^-; lattice becomes more loosely packed
  • Electrostatic attraction decreases
• Overall effect is that Δ_solnH becomes more exothermic, so solubility increases

Question 23

Suggest which of Na⁺ and Ca²⁺ has the most exothermic enthalpy of hydration.

Answer 23

• Ca²⁺ is larger than Na⁺, which would decrease Ca²⁺’s charge density, but it has double the charge, which would increase its charge density
• Charge has dominant effect since size doesn’t change by same degree
• So Ca²⁺ has higher charge density
• Attracts water more strongly / stronger ion-dipole interactions
• More exothermic enthalpy of hydration