Conductivity, Ionic Compounds, and Lattice Enthalpy

Requirement for Conductivity: Mobile charges are necessary for a material to conduct electricity.
Metals: Conductive due to metallic bonds and valence electrons that can move freely.
Insulators: Materials where electrons are not free to move, preventing electrical conductivity (e.g., air).
Voltage: Represents the 'push' on electrons. A high enough voltage can make even insulators conduct electricity.

Deionized Water: Does not conduct electricity well because it lacks mobile ions.
Tap Water: Conducts slightly due to the presence of ions like calcium, magnesium, iron, and chloride.
Salt Water: Conducts electricity very well due to a high concentration of mobile ions.

Solid Salt (e.g., NaCl): Does not conduct electricity because ions are locked in place within the crystal lattice and cannot move.
Salt Water (NaCl(aq)): Conducts electricity because ions are dissociated and mobile.

General Properties: Hard, brittle, high melting and boiling points, generally non-volatile (do not easily form a gas).
Conductivity: Do not conduct electricity when solid because ions are not free to move within the lattice structure.
Molten State: Conduct electricity when molten (liquid) because ionic bonds break, allowing ions to move freely.
Solubility: Can be insoluble in most nonpolar solvents like hexane.
Dissolution in Water: When an ionic solid dissolves in water, the ions are surrounded by water molecules (hydrated) and become mobile. This is represented as $Na^+(aq)$ , indicating the sodium ions are surrounded by water molecules.

Definition: The energy required to break the bonds in an ionic compound, separating it into gaseous ions at an infinite distance from each other.
Endothermic Process: Lattice enthalpy is always a positive value, indicating that energy is required to break the bonds.
Equation: For sodium chloride (NaCl):
$NaCl(s) \rightarrow Na^+(g) + Cl^-(g) \quad \Delta H = +790 \text{ kJ/mol}$
This equation shows one mole of solid NaCl turning into one mole of gaseous sodium ions and one mole of gaseous chloride ions.
Units: Measured in kilojoules per mole (kJ/mol).

Charge: Higher charges on ions result in stronger electrostatic attractions (e.g., $Mg^{2+}$ and $O^{2-}$ , which have a greater force of attraction than $Na^{+}$ and $Cl^{-}$ $).
Distance: Smaller ionic radii result in stronger attractions. The force of attraction is inversely proportional to the distance between the ions.
Electrostatic Attraction: The force of attraction between two charges is proportional to the product of the charges and inversely proportional to the square of the distance between them: F∝ q1q2/r2
Where:
- $F$ is the force of attraction.
- $q<em>1$ and $q</em>2$ are the magnitudes of the charges.
- $r$ is the distance between the charges.
Ionic Radii: Data booklets provide ionic radii. Smaller ions with higher charges result in greater lattice enthalpy.

Charge Comparison: Magnesium oxide (MgO) has a higher lattice enthalpy due to the +2 and -2 charges on the ions compared to +1 and -1 charges in NaCl.
Size Comparison: Within Group 1 metals, lattice enthalpy increases as the size of the ion decreases (e.g., LiF has a higher lattice enthalpy than CsI because Li+ and F- are smaller ions).