S2.1.3 Lattice strcutures and properties

Lattice Structure and Properties of Ionic Compounds

When you sprinkle salt on your food, have you ever wondered why it forms tiny crystals instead of just being a shapeless powder?

These crystals are not random—they reflect the underlying structure of ionic compounds.

Ionic Compounds and Their Lattice Structure

These ions are not scattered randomly; instead, they form a three-dimensional lattice structure.

What is a Lattice Structure?

The lattice structure gives ionic compounds their characteristic properties, such as high melting points and brittleness.

The formula of an ionic compound, such as NaCl, represents the simplest ratio of ions in the lattice (1:1 for $N{a}^{+}$ and $C{l}^{-}$).

Physical Properties of Ionic Compounds

Ionic compounds are known for their unique set of physical properties. Let’s break them down:

1.Volatility: Why Don’t Ionic Compounds Vaporize Easily?

Volatility refers to a substance’s ability to vaporize. Ionic compounds have low volatility because their lattice structure is held together by strong electrostatic forces.

To break these bonds and convert the solid into a gas, a significant amount of energy is required. This is why ionic compounds tend to have high melting and boiling points.

2.Electrical Conductivity: When Do Ionic Compounds Conduct Electricity?

To conduct electricity, a substance must have mobile charged particles. In ionic compounds:

• In the solid state: Ions are fixed in place within the lattice and cannot move, so the compound does not conduct electricity.
• In the molten (liquid) state or when dissolved in water (aqueous state): The lattice breaks down, freeing the ions to move and carry an electrical current.

3.Solubility: Why Are Ionic Compounds Soluble in Water?

Ionic compounds are generally soluble in polar solvents like water. This is because water molecules have partial positive and negative charges, which interact with the ions in the lattice:

• The partial positive charge on hydrogen atoms attracts anions (e.g., $C{l}^{-}$).
• The partial negative charge on oxygen atoms attracts cations (e.g., $N{a}^{+}$).

This interaction pulls the ions out of the lattice and into solution.

However, not all ionic compounds are soluble in water. If the forces holding the lattice together are stronger than the interactions with water molecules, the compound remains insoluble. For example, calcium carbonate ($CaC{O}_3$) is largely insoluble in water.

Lattice Enthalpy: A Measure of Bond Strength

The strength of the ionic bonds in a lattice can be quantified using lattice enthalpy. This is the energy required to separate one mole of an ionic solid into its gaseous ions under standard conditions.

Factors Affecting Lattice Enthalpy

1. Ion Radius:
   • Smaller ions are closer together in the lattice, leading to stronger electrostatic attractions and higher lattice enthalpy.
   • Example: NaF has a higher lattice enthalpy than KF because $N{a}^{+}$ is smaller than $K^{+}$.
2. Ion Charge:
   • Higher charges on the ions result in stronger attractions and higher lattice enthalpy.
   • Example: MgO has a much higher lattice enthalpy than NaCl because $M{g}^{2+}$ and $O^{2-}$ have greater charges than $N{a}^{+}$ and $C{l}^{-}$.

Summary of Key Properties

Reflection