The Giant Ionic Lattice
Ionic compounds do not exist as discrete molecules. Instead, they arrange themselves into a giant three-dimensional lattice structure: a regular repeating array of alternating cations and anions extending indefinitely in all directions. The chemical formula (e.g. NaCl) represents only the simplest whole-number ratio of ions. It is an empirical formula, not a molecular formula.
Never Say "Molecule" for Ionic Compounds
You must completely eradicate the term "molecule" when discussing ionic compounds. A molecule implies a discrete, finitely bounded, covalently bonded unit. Stating "water dissolves molecules of NaCl" demonstrates a fundamental misunderstanding. Always use "formula units" or "ions in a lattice".
Physical Properties. All Traceable to the Lattice
| Property | Observation | Explanation (Mark-Scheme Language) |
|---|---|---|
| Melting / Boiling Points | Very high | Large amounts of thermal energy needed to overcome the strong electrostatic attractions between oppositely charged ions throughout the giant lattice. |
| Electrical Conductivity (solid) | Does NOT conduct | Ions are locked in fixed positions within the lattice. They cannot migrate to carry a current. |
| Electrical Conductivity (molten / aqueous) | DOES conduct | Ions are free to move and migrate toward electrodes, carrying charge through the liquid. |
| Solubility in Water | Generally soluble | Polar water molecules (ion-dipole interactions) surround and stabilise the free ions, overcoming the lattice forces. The ion is said to be hydrated. |
| Brittleness | Hard but brittle | A mechanical force displaces a layer of ions, causing like charges to align. The resulting repulsion shatters the crystal along the plane. |
Conductivity Terminology
When explaining why molten or aqueous ionic compounds conduct, writing "electrons are free to move" will result in zero marks. The current in an ionic liquid is carried by the physical migration of massive charged ions, not subatomic electrons.
The ONLY acceptable mark-scheme phrase is: "the ions are free to move".
Why Ionic Compounds Are Brittle (Cleavage Under Stress)
Melting Point Explanation
When explaining high melting points, you must explicitly state: "energy is needed to overcome the strong electrostatic attractions between oppositely charged ions in the giant lattice." Do NOT write "intermolecular forces". Ionic compounds have no molecules and therefore zero intermolecular forces.
Solubility in Non-Polar Solvents
Ionic compounds are generally insoluble in non-polar solvents (e.g. hexane, cyclohexane). Non-polar molecules lack the partial charges needed to form ion-dipole interactions, so they cannot overcome the strong lattice forces holding the ions together.
Lattice Enthalpy
Lattice Enthalpy (ΔHlat)
Lattice enthalpy is the enthalpy change when one mole of an ionic compound is separated into its constituent gaseous ions under standard conditions. It is always endothermic (positive \(\Delta H\)) because energy must be input to overcome the electrostatic attractions.
\(\text{NaCl(s)} \rightarrow \text{Na}^+(g) + \text{Cl}^-(g) \quad \Delta H_{lat} > 0\)
Lattice enthalpy is a measure of the strength of the ionic bond in a compound. A higher lattice enthalpy means stronger electrostatic attractions between the ions, which in turn means a higher melting point and greater stability.
Coulomb's Law
The electrostatic force F between two ions is governed by Coulomb's Law. While you do not need to calculate F numerically, you must understand qualitatively how the two key factors affect lattice enthalpy:
Coulomb's Law
\[ F \propto \frac{q^+ \times q^-}{r^2} \]
↑ Ionic Charge (q)
Higher charge = stronger attraction = greater lattice enthalpy.
MgO (2+ / 2−) has a much higher lattice enthalpy than NaCl (1+ / 1−).
↓ Ionic Radius (r)
Smaller ions = closer together = greater lattice enthalpy.
NaF has a higher lattice enthalpy than NaI because F⁻ is much smaller than I⁻.
Charge Density
Charge density combines both factors into one concept: it is the ratio of ionic charge to ionic volume. Ions with high charge density (small radius, high charge) create the strongest electrostatic fields and thus the highest lattice enthalpies.
Why MgO Melting Point > NaCl
To explain why MgO has a much higher melting point than NaCl, structure your answer using both charge and radius:
- 1. \(\text{Mg}^{2+}\) has a higher charge than \(\text{Na}^+\) → stronger electrostatic attraction.
- 2. \(\text{Mg}^{2+}\) has a smaller ionic radius than \(\text{Na}^+\) → ions packed closer together.
- 3. \(\text{O}^{2-}\) has a higher charge than \(\text{Cl}^-\) → stronger electrostatic attraction.
- Conclusion: MgO has a vastly greater lattice enthalpy → substantially higher melting point.
Ionic vs Covalent Character (Polarisation)
Bonding exists on a continuum between purely ionic and purely covalent. Covalent character arises when a small, highly charged cation polarises (distorts) the electron cloud of a large anion, pulling electron density toward itself. This creates a degree of electron sharing within the ionic lattice.
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