Quick questions on Ionic bonding and the properties of ionic compounds (QCE Chemistry Unit 1)

Metals lose electrons to form cations (positive ions); non-metals gain electrons to form anions (negative ions). Both species typically reach a noble-gas electron configuration (closed-shell stability).

In the solid state, ions are arranged in a regular repeating 3D pattern (a crystal lattice). The geometry depends on the relative sizes of the cation and anion (the radius ratio). Common QCE-relevant lattices:

Strong electrostatic attractions throughout the lattice; large lattice energy. Comparing within ionic compounds, higher charges and smaller ions give higher melting points. MgO (charges 2+/2-) melts above 2800 degrees C; NaCl (1+/1-) melts at 801 degrees C.

The lattice resists deformation, so the solid is hard. But under stress, one layer of ions can shift over another so that like charges align; the resulting repulsion shatters the crystal along that plane. This is why ionic crystals cleave cleanly under impact.

Water is a polar solvent that solvates ions effectively. Solubility depends on the balance between lattice energy (must be overcome to dissociate the solid) and hydration energy (released when ions are solvated). Compounds with comparable lattice and hydration energies tend to dissolve; compounds with very high lattice energy relative to hydration energy do not.

Predict and write the formula of the ionic compound formed between (a) aluminium and oxygen, (b) calcium and phosphate. [4 marks]

Lattice energies (in kJ mol$^{-1}$): $\text{NaCl}\,786$, $\text{MgO}\,3{,}795$, $\text{CaO}\,3{,}414$. Account for the differences. [4 marks]

A student tests solid and molten $\text{KBr}$ with a conductivity meter. (a) Describe the bonding. (b) Predict observations in both states.