Module 3 - Lattice Enthalpy

Question 1

What is lattice enthalpy?

Answer 1

• Definition: the enthalpy change when one more of an ionic lattice is formed from its gaseous ions under standard conditions (298K, 100kPa)
• Equation Example:
  Na+ (g) + Cl- (g) → NaCl (s)
• Always exothermic (negative value) as energy is released when strong electrostatic forces form.

Question 2

Why is lattice enthalpy exothermic?

Answer 2

• Ionic compounds form due to strong electrostatic attractions between oppositely charged ions.
• Energy is released when these forces form a stable, ordered lattice.
• More exothermic values indicate stronger attractions and a more stable lattice.

Question 3

Factors affecting lattice enthalpy?

Answer 3

• Ionic Charge:
  → Higher charge = stronger attraction = more exothermic lattice enthalpy.
  → Example: MgO (Mg²⁺ & O²⁻) has a more exothermic lattice enthalpy than NaCl (Na⁺ & Cl⁻).
• Ionic Radius:
  → Smaller ions = stronger attraction = more exothermic lattice enthalpy.
  → Example: LiF (small ions) has a more exothermic lattice enthalpy than CsI (large ions).

Question 4

The Born-Haber cycle

Answer 4

• An enthalpy cycle used to calculate lattice enthalpy indirectly via Hess’s Law.
• Includes multiple enthalpy changes like formation, atomisation, ionisation energy, and electron affinity
• Allows experimental determination of lattice enthalpy.

Question 5

Hess’s law

Answer 5

• States that the total enthalpy change of a reaction is independent of the route taken.
• Helps apply enthalpy cycles like the Born-Haber cycle to determine unknown values.

Question 6

Key enthalpy changes in a Born-Haber cycle

Answer 6

• Enthalpy of Formation (ΔHf°):
  → The enthalpy change when one mole of an ionic compound forms from its elements in their standard states.
  → Example:
  Na(s) + 1/2Cl2(g) → NaCl(s)
• Enthalpy of Atomisation (ΔHat°):
  → The enthalpy change when one mole of gaseous atoms forms from the element in its standard state.
  → Example:
  1/2Cl2(g) → Cl(g)
• First Ionisation Energy (IE1):
  → The energy needed to remove one electron per atom from one mole of gaseous atoms to form gaseous +1 ions.
  → Always endothermic as energy is required to remove an electron.
  → Example:
  Na(g) → Na+(g) + e-
• First Electron Affinity (EA1):
  → The enthalpy change when one mole of gaseous atoms gains one electron per atom to form gaseous -1 ions.
  → Usually exothermic because the electron is attracted to the nucleus.
  → Example:
  Cl(g) + e- → Cl-(g)

Question 7

Second election affinity (EA2) and why it’s endothermic?

Answer 7

• The enthalpy change when one mole of gaseous -1 ions gains an additional electron to form gaseous -2 ions.
• Endothermic because energy is required to overcome the repulsion between the negative ion and the incoming electron.
• Example:
  O- (g) + e- → O2- (g)

Question 8

How Lattice Enthalpy Affects Ionic Compound Properties

Answer 8

• Melting & Boiling Points:
  → Higher lattice enthalpy = stronger ionic bonds = higher melting/boiling points.
  → Example: MgO has a much higher melting point than NaCl due to stronger attractions.
• Solubility in Water:
  → More exothermic lattice enthalpy means stronger bonds, making it harder to dissolve.
  → Solubility depends on a balance between lattice enthalpy and hydration enthalpy.