Ionic compounds: Bonds, Structure & Properties

Electrostatic Magnitude: The strength of the ionic bond is determined by the magnitude of the charges on the ions; higher charges result in stronger attraction and more stable structures.

Energy Requirements: Because the electrostatic forces are strong and act in all directions throughout the giant lattice, a significant amount of thermal energy is required to disrupt the structure.

Coulombic Attraction: The force $F$ between two ions is proportional to the product of their charges $(q_1 q_2)$, meaning $Mg^{2+}$ and $O^{2-}$ have much stronger bonds than $Na^+$ and $Cl^-$.

Ions vs. Electrons: It is critical to distinguish that in ionic compounds, ions carry the charge during conduction, whereas in metals, delocalized electrons perform this role.

Bonding vs. Structure: While the 'bond' is the individual attraction between two ions, the 'structure' refers to the entire giant lattice containing trillions of ions.

The 'Energy' Keyword: When explaining melting points, always mention that 'large amounts of energy' are required to break the 'strong electrostatic forces'. Missing the energy component often loses marks.

State Matters: If a question asks why a substance doesn't conduct electricity, check if it is labeled as '(s)' for solid. If so, specify that 'ions are in fixed positions and cannot move'.

Charge Comparison: In comparisons (like $MgO$ vs $NaCl$), explicitly state that 'the charges on the ions are greater' (e.g., 2+ vs 1+) leading to 'stronger electrostatic attraction'.

The 'Moving Electrons' Error: A very common mistake is stating that ionic compounds conduct electricity because electrons move. Always specify that it is the ions that move to carry the charge.

Molecular Misconception: Students often refer to 'ionic molecules'. Ionic compounds do not form molecules; they form giant lattices. Avoid using the word 'molecule' when discussing $NaCl$ or $MgO$.

Melting vs. Dissolving: Ensure you understand that both melting (thermal energy) and dissolving (solvent interaction) achieve the same result: releasing ions from their fixed lattice positions.