Ionic solids are held together by strong electrostatic attractions that strongly influence phase-change behavior and mechanical and electrical properties. Understanding these properties requires linking macroscopic observations to what ions can and cannot do inside a crystal lattice.

What makes an ionic solid “ionic”

Ionic solids form when cations and anions arrange into an extended, repeating structure that maximizes attractions and minimizes repulsions.

Two complementary models of the NaCl ionic lattice are shown: an expanded model that makes the alternating ion pattern and coordination easier to see, and a space-filling model that better represents how ions pack in a real crystal. The color-coding distinguishes cations from anions, reinforcing that the “particles” in an ionic solid are ions in a continuous 3D array rather than discrete molecules. Source

Because there are no discrete molecules to “separate,” many properties depend on disrupting the entire lattice rather than just weakening a few intermolecular attractions.

Strength of ion–ion attraction (what controls it)

The strength of attraction between ions increases with:

• Greater ionic charge magnitude (e.g., $2^+$ with $2^-$ stronger than $1^+$ with $1^-$)

• Smaller ionic size (shorter ion–ion distance)

A larger-magnitude lattice energy generally corresponds to higher melting/boiling points and lower volatility.

This proportionality explains why ionic solids built from small, highly charged ions tend to be especially “high-melting.”

Melting point, boiling point, and vapor pressure

Why melting and boiling points are high

To melt an ionic solid, enough energy must be supplied to partially overcome the ion–ion attractions that lock ions into fixed lattice positions. To boil it, those attractions must be overcome to a much greater extent, separating ions into the gas phase (often requiring extremely high temperatures and sometimes accompanied by decomposition).

Key macroscopic results:

• High melting points: significant energy needed before ions can move past one another

• High boiling points (when boiling is feasible): separating ions completely is energetically costly

• Trends track attraction strength: higher charges and smaller ions typically increase these temperatures

Why vapor pressure is low

Vapor pressure reflects how easily particles escape into the gas phase. In ionic solids, ions are strongly stabilized in the lattice, so very few particles have enough energy to escape under ordinary conditions.

• Stronger ion–ion attractions ⟶ lower vapor pressure

• Low vapor pressure aligns with the idea that ionic solids are generally nonvolatile

Brittleness (why ionic crystals shatter)

Ionic solids are typically brittle rather than malleable. Under stress (e.g., a hammer blow), lattice layers may be forced to shift slightly. If that shift brings like-charged ions closer together:

A schematic of the NaCl lattice highlights that small displacements can convert locally favorable cation–anion attractions into unfavorable like-charge neighbors. The diagram makes the brittleness mechanism concrete: once like charges align, electrostatic repulsion rises sharply and the crystal fractures rather than plastically deforming. Source

• cation–cation and/or anion–anion repulsions increase sharply

• the local repulsion can exceed the lattice’s ability to rearrange

• the crystal cleaves along planes and fractures

This brittleness is a direct consequence of an ordered lattice of alternating charges: small displacements can create strongly unfavorable like-charge contacts.

Electrical conductivity (when ions are mobile)

Ionic substances conduct electricity only when they contain mobile charged particles.

Solid ionic compounds: do not conduct

In the solid state:

• ions are fixed in lattice positions (they vibrate but do not translate)

• charges cannot flow through the crystal

• result: poor electrical conductivity

Molten or dissolved ionic compounds: conduct

When ionic solids are molten or dissolved in water, ions become free to move:

The NaCl lattice is depicted as a regular 3D arrangement of alternating ions, with an expanded view that makes the geometry and nearest-neighbor relationships more visible. This supports the conductivity idea by emphasizing that in the solid, ions occupy fixed lattice positions; electrical conduction requires disrupting that rigid arrangement so ions can migrate. Source

• molten ionic liquid: ions move within the liquid, carrying charge

• aqueous solution: ions are separated and stabilized by water, allowing migration toward electrodes

• result: good electrical conductivity (an electrolyte)

Conductivity depends qualitatively on how many ions are present and how easily they move, but the core requirement is the same: ionic mobility.