Structure of ionic solids

Ionic solid

• A solid composed of positive and negative ions held together by strong electrostatic forces (ionic bonds) in a repeating 3D pattern called a crystal lattice

• Example: NaCl (table salt

Crystal lattice

• A highly ordered, repeating 3D arrangement of ions in an ionic solid

• Maximizes attractions and minimizes repulsions

• Gives ionic solids their regular geometric shapes

Lattice energy

• The energy released when gaseous ions form one mole of a solid ionic compound

• Higher lattice energy = stronger ionic bonds

• Increases with higher ion charges and smaller ion sizes

• Units: kJ/mol

Why ionic solids have high melting points

• Because it takes a lot of energy to overcome the strong electrostatic forces holding the ions in place in the lattice

• Example: NaCl melts at 801°C

Why ionic solids are brittle

• When a force is applied, layers of ions shift → like charges align → repulsion occurs → crystal shatters

• This is called cleavage

• Before: + – + – After force: + – + –

               – + – +    (shifted)    – + – +

               + – + –                      + – + –

  Now: + next to + → REPEL!

Electrical conductivity of ionic solids

• Solid state: ❌ No conductivity — ions are locked in place

• Molten (liquid): ✅ Conducts electricity — ions can move

• Dissolved in water (aqueous): ✅ Conducts electricity — ions are free to move

Coordination number

• The number of oppositely charged ions surrounding a given ion in a crystal lattice

• Depends on relative ion sizes

• In NaCl: each Na⁺ is surrounded by 6 Cl⁻ → coordination number = 6

Unit cell

• The smallest repeating unit in a crystal lattice that shows the full symmetry of the structure

• Stacking many unit cells forms the entire crystal

• For NaCl: face-centered cubic (FCC) structure

Ion size and packing

• Smaller ions pack more tightly → shorter distances between ions → stronger attraction → higher lattice energy

• Cations are smaller than their atoms (lost electrons)

• Anions are larger than their atoms (gained electrons)

NaCl structure

• Alternating Na⁺ and Cl⁻ ions in a 3D grid

• Each ion has a coordination number of 6

• Cubic unit cell

• Highly stable due to strong ionic bonding

CsCl structure

• Found in some ionic compounds where cation is large

• Coordination number = 8 (each Cs⁺ surrounded by 8 Cl⁻)

• Body-centered cubic unit cell

• Example: CsCl

Trend in lattice energy

• Lattice energy increases as:

• Ion charge increases (e.g., MgO > NaCl)

• Ion size decreases (e.g., LiF > KF)

• So Mg²⁺ and O²⁻ → very high lattice energy → very high mp

Born-haber cycle

• A thermochemical cycle used to calculate lattice energy indirectly using Hess’s Law

• Includes steps like atomization, ionization, electron affinity, etc

Why all ionic compounds don’t have the same structure

• Because the ratio of ion sizes determines how they pack

• If cation is small, fewer anions fit around it

• If cation is large, more anions can surround it

Stoichiometry in ionic crystals

• The formula of the compound (like NaCl or CaF₂) reflects the ratio of ions in the lattice

• In CaF₂: each Ca²⁺ is surrounded by 8 F⁻, but each F⁻ is shared among multiple Ca²⁺ → overall ratio = 1 Ca : 2 F

Interionic distance

• The distance between the nuclei of adjacent positive and negative ions

• Shorter distance → stronger attraction → higher lattice energy

Effect of charge on melting point

• Higher ion charges → stronger attractions → higher melting points

• NaCl (Na⁺Cl⁻): 801°C

• MgO (Mg²⁺O²⁻): 2852°C

Ionic radius trend

• Decreases across a period (left to right)

• Increases down a group

• Cations < neutral atom < anions

Packing efficiency

• How tightly ions are packed in the lattice

• Ionic solids are densely packed → efficient use of space

• But not 100% — there’s always some empty space