IB DP Chemistry - S2.1.3 Lattice structures and properties - Study Notes - New Syllabus - 2026, 2027 & 2028
IB DP Chemistry – S2.1.3 Lattice structures and properties – Study Notes – New Syllabus
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Structure 2.1.3 — Ionic Lattice Structures, Physical Properties, and Lattice Enthalpy
Structure 2.1.3 — Ionic Lattice Structures, Physical Properties, and Lattice Enthalpy
Ionic Compounds
Ionic compounds are formed from the electrostatic attraction between oppositely charged ions (cations and anions). These ions arrange themselves into a regular and repeating three-dimensional (3D) lattice structure.
Empirical Formula: The formula of an ionic compound represents the simplest ratio of ions that balances total positive and negative charges.
For example:
- Sodium chloride: \( \text{NaCl} \)
- Magnesium oxide: \( \text{MgO} \)
- Calcium chloride: \( \text{CaCl}_2 \)
Structure: 3D Giant Ionic Lattice
Each ion is surrounded by oppositely charged ions in a highly ordered array. The strong ionic bonds extend throughout the entire lattice, making the structure very stable.
Example: In \( \text{NaCl} \), each \( \text{Na}^+ \) is surrounded by 6 \( \text{Cl}^- \) ions and vice versa in a cubic arrangement.
Physical Properties of Ionic Compounds:
1. Volatility:
- Ionic compounds generally have low volatility (they do not evaporate easily).
- Strong ionic bonds require a large amount of energy to overcome.
2. Electrical Conductivity:
- In solid state: Do not conduct electricity (ions are fixed in position).
- Molten/aqueous state: Conduct electricity due to free-moving ions.
3. Solubility in Water:
- Many ionic compounds are soluble in water.
- Polar water molecules attract and stabilize ions, allowing dissolution.
- Solubility varies depending on lattice energy and hydration energy.
Lattice Enthalpy \( ( \Delta H_{\text{latt}} ) \):
Lattice enthalpy is a measure of the strength of the ionic bond in an ionic lattice. It is defined as:
The energy required to separate one mole of an ionic compound into its gaseous ions.
Factors Affecting Lattice Enthalpy:
- 1. Ionic charge: Greater charges → stronger electrostatic attraction → higher lattice enthalpy
- 2. Ionic radius: Smaller ions → ions are closer together → stronger attraction → higher lattice enthalpy
Examples:
| Compound | Ionic Charges | Lattice Enthalpy (kJ mol⁻¹) |
|---|---|---|
| \( \text{NaCl} \) | \( \text{Na}^+ \), \( \text{Cl}^- \) | -787 |
| \( \text{MgO} \) | \( \text{Mg}^{2+} \), \( \text{O}^{2-} \) | -3795 |
| \( \text{CaO} \) | \( \text{Ca}^{2+} \), \( \text{O}^{2-} \) | -3414 |
Conclusion: Higher lattice enthalpy means stronger ionic bonding and higher melting point. Compounds like \( \text{MgO} \), with high charges and small radii, have extremely strong ionic bonds.
Example
Which compound has lower volatility: sodium chloride \( (\text{NaCl}) \) or magnesium oxide \( (\text{MgO}) \)? Explain based on lattice enthalpy.
▶️Answer/Explanation
\( \text{MgO} \) has a much higher lattice enthalpy than \( \text{NaCl} \) because:
- \( \text{Mg}^{2+} \) and \( \text{O}^{2-} \) have higher charges than \( \text{Na}^+ \) and \( \text{Cl}^- \).
- \( \text{Mg}^{2+} \) and \( \text{O}^{2-} \) are smaller in size, so charge density is greater.
Thus, \( \text{MgO} \) has stronger ionic bonds → much lower volatility than \( \text{NaCl} \).
Example
Why is barium sulfate \( (\text{BaSO}_4) \) sparingly soluble in water despite being ionic?
▶️Answer/Explanation
The lattice enthalpy of \( \text{BaSO}_4 \) is very high because the sulfate ion is large and complex, leading to strong electrostatic interactions.
In addition, the hydration enthalpy (energy released when ions dissolve in water) is not enough to overcome the lattice enthalpy.
Result: \( \text{BaSO}_4 \) remains mostly undissolved → low solubility.
Example
Explain why \( \text{NaCl} \) conducts electricity when molten but not as a solid.
▶️Answer/Explanation
In solid \( \text{NaCl} \), the ions are locked in a rigid lattice and cannot move → no conductivity.
When molten, the lattice breaks down and ions become mobile → \( \text{Na}^+ \) and \( \text{Cl}^- \) can carry charge.
Thus, molten \( \text{NaCl} \) is a good conductor of electricity.
Example
Which compound has a higher lattice enthalpy: lithium fluoride \( (\text{LiF}) \) or cesium iodide \( (\text{CsI}) \)?
▶️Answer/Explanation
\( \text{LiF} \) has a much higher lattice enthalpy due to:
- \( \text{Li}^+ \) is much smaller than \( \text{Cs}^+ \).
- \( \text{F}^- \) is smaller than \( \text{I}^- \).
Smaller ions → stronger electrostatic attraction → higher lattice enthalpy.