DAT General Chemistry: Types of Crystalline Solids

DAT High-Yield Concept

Four Main Solid Types: Crystalline solids are categorized by their structural units and attractive forces into ionic, molecular, covalent-network, and metallic solids. (Long-range geometric order contrasts with amorphous solids, which are covered in depth in their own lesson).
Electrical Conductivity Rule: Metallic solids conduct in solid and liquid states; ionic solids conduct only when molten or dissolved in aqueous solution; molecular solids do not conduct in any phase.
Relative Melting Points: Molecular solids have the lowest melting points due to weak intermolecular forces. Covalent-network solids (e.g., diamond $\ce{C}$ , quartz $\ce{SiO2}$ ) exhibit exceptionally high melting points due to continuous covalent bonding throughout the lattice.
Mechanical Properties: Ionic solids are hard and brittle; metallic solids are malleable and ductile; covalent-network solids are extremely hard; molecular solids are soft.

1. Ionic Solids

Ionic solids consist of positively charged cations and negatively charged anions arranged in a rigid, three-dimensional repeating lattice. The structural integrity is maintained by strong coulombic (electrostatic) attractions described by Coulomb's law, where lattice energy increases with higher ion charges and smaller ionic radii.

Because electrostatic forces act in all directions, ionic solids possess high melting and boiling points (e.g., $\ce{NaCl}$ melts at $\pu{801 degC}$ ). Mechanical stress causes planes of ions to shift relative to one another; when like-charged ions align, intense electrostatic repulsion forces the crystal to cleave, making ionic solids hard but brittle.

In the solid state, ions are fixed in lattice positions, so ionic solids are electrical insulators. However, when melted or dissolved in water, the lattice breaks down, releasing mobile charge carriers (ions) that conduct electricity.

The crystal structure of sodium chloride, NaCl, a typical ionic compound. The __ purple spheres represent sodium cations, Na + , and the __ green spheres represent chloride anions, Cl − . (Source: Benjah-bmm27 ( talk · contribs ) · Public domain · Wikimedia Commons)

2. Molecular Solids

Molecular solids are composed of discrete neutral molecules (or monoatomic noble gas atoms) held together by non-covalent intermolecular forces (IMFs): London dispersion forces, dipole-dipole interactions, or hydrogen bonds. Crucially, phase changes involve overcoming these relatively weak IMFs rather than breaking covalent intramolecular bonds.

Because IMFs are substantially weaker than ionic or covalent bonds, molecular solids generally feature low-to-moderate melting points (typically below $\pu{300 degC}$ ), low mechanical hardness, and high volatility. Nonpolar molecular solids such as dry ice ( $\ce{CO2}$ ), iodine ( $\ce{I2}$ ), phosphorus ( $\ce{P4}$ ), and sulfur ( $\ce{S8}$ ) rely purely on dispersion forces. Polar molecular solids like ice ( $\ce{H2O}$ ) and sucrose ( $\ce{C12H22O11}$ ) feature dipole-dipole or hydrogen bonding, producing slightly higher melting points.

Because valence electrons are localized strictly within individual covalent bonds and no free ions exist, molecular solids are non-conductors of electricity in all physical states.

3. Covalent-Network Solids

Covalent-network solids consist of atoms held together by a continuous network of directional covalent bonds extending throughout the entire crystal sample. The entire solid functions essentially as one giant macro-molecule.

Common examples tested on the DAT include diamond ( $\ce{C}$ ), quartz ( $\ce{SiO2}$ ), silicon carbide ( $\ce{SiC}$ ), and boron nitride ( $\ce{BN}$ ). In diamond, each carbon atom is $sp^3$ -hybridized in a tetrahedral lattice with bond angles of $109.5^\circ$ . Overcoming these continuous covalent bonds requires immense thermal energy, yielding extraordinarily high melting points (often $> \pu{2000 degC}$ ) and supreme hardness.

Most covalent-network solids lack mobile electrons and are electrical insulators. A key exception is graphite (an allotrope of carbon), where $sp^2$ -hybridized carbon sheets feature delocalized $\pi$ -electrons that allow electrical conduction along the two-dimensional planes.

4. Metallic Solids

Metallic solids consist of metal cations suspended within a delocalized "sea of mobile valence electrons." The non-directional electrostatic attraction between the metal cations and the surrounding electron cloud constitutes the metallic bond.

Because valence electrons are free to drift throughout the entire structure, metallic solids exhibit high electrical and thermal conductivity in both solid and liquid states. The non-directional nature of metallic bonding allows atomic planes to slide past one another without catastrophic repulsive interactions, making metals malleable (capable of being hammered into sheets) and ductile (capable of being drawn into wires).

Melting points of metallic solids vary widely depending on atomic size and the number of valence electrons contributed to the sea: mercury ( $\ce{Hg}$ ) is a liquid at room temperature (melts at $\pu{-39 degC}$ ), whereas tungsten ( $\ce{W}$ ) melts at $\pu{3422 degC}$ .

Diagnostic Test: Electrical Conductivity & Solid Type

Conducts electricity as a solid: Metallic solid (or graphite).
Does NOT conduct as a solid, but DOES conduct when liquid/molten or dissolved: Ionic solid.
Does NOT conduct in any phase: Molecular solid (if low melting point) or Covalent-Network solid (if extremely high melting point).

Mnemonic: Remember "I-M-C-M" for the 4 Crystalline Classes:
• Ionic = Ions + Coulombic forces (Conducts when molten/aq)
• Molecular = Molecules + IMFs (Low melting point, Insulator)
• Covalent-Network = Continuous covalent bonds (Extremely hard & High melting point)
• Metallic = Mobile electron sea (Conducts as solid, Malleable)

Summary Comparison Table

Solid Type	Structural Units	Primary Forces	Properties	Electrical Conductivity	Examples
Ionic	Cations and anions	Ionic / Electrostatic attraction	Hard, brittle, high melting points	Solid: No Molten/Aq: Yes	$\ce{NaCl}$ , $\ce{CaF2}$ , $\ce{MgO}$ , $\ce{KBr}$
Molecular	Molecules or atoms	IMFs (Dispersion, Dipole-Dipole, H-bonds)	Soft, low melting points ( $< \pu{300 degC}$ )	Solid: No Liquid: No	$\ce{H2O}$ (ice), $\ce{CO2}$ , $\ce{I2}$ , $\ce{C12H22O11}$
Covalent-Network	Atoms	Continuous Covalent Bonds	Very hard, extremely high melting points	Solid: No (except graphite) Liquid: No	Diamond ( $\ce{C}$ ), Quartz ( $\ce{SiO2}$ ), $\ce{SiC}$ , $\ce{BN}$
Metallic	Metal cations in electron sea	Metallic bonding	Malleable, ductile, variable melting points	Solid: Yes Liquid: Yes	$\ce{Cu}$ , $\ce{Fe}$ , $\ce{Na}$ , $\ce{Au}$ , $\ce{W}$

Worked Examples & Exam Strategy

Worked Example 1: Classification & Conductivity Prediction

Problem: Classify the solid form of each species: $\ce{SiO2}$ , $\ce{I2}$ , $\ce{Cu}$ , and $\ce{KBr}$ . State whether solid $\ce{KBr}$ conducts electricity.

Solution & Reasoning:

$\ce{SiO2}$ (quartz): Consists of a continuous 3D arrangement of covalently bonded silicon and oxygen atoms $\rightarrow$ Covalent-network solid.
$\ce{I2}$ : Consists of nonpolar diatomic molecules held in a lattice by London dispersion forces $\rightarrow$ Molecular solid.
$\ce{Cu}$ : Consists of copper cations surrounded by delocalized valence electrons $\rightarrow$ Metallic solid.
$\ce{KBr}$ : Consists of $\ce{K^+}$ cations and $\ce{Br^-}$ anions held by electrostatic attraction $\rightarrow$ Ionic solid. Solid $\ce{KBr}$ does not conduct electricity because its ions are locked into fixed lattice positions.

Common Exam Trap: Confusing covalent-network solids like quartz ( $\ce{SiO2}$ ) with molecular solids like dry ice ( $\ce{CO2}$ ). Remember that $\ce{SiO2}$ forms an extended 3D covalent lattice, whereas $\ce{CO2}$ consists of individual molecules bound only by weak dispersion forces.

Worked Example 2: Relative Melting Point Ranking

Problem: Rank the following solids in order of increasing melting point: $\ce{CH4}$ , $\ce{H2O}$ , $\ce{NaCl}$ , $\ce{C_{(diamond)}}$ .

Solution & Reasoning:

$\ce{CH4}$ : Nonpolar molecular solid bound only by weak dispersion forces (lowest melting point, sublimates at $\pu{-182 degC}$ ).
$\ce{H2O}$ : Polar molecular solid held by hydrogen bonds (melts at $\pu{0 degC}$ ).
$\ce{NaCl}$ : Ionic solid held by strong electrostatic forces (melts at $\pu{801 degC}$ ).
$\ce{C_{(diamond)}}$ : Covalent-network solid bound by continuous covalent bonds throughout the crystal (highest melting point, $> \pu{3500 degC}$ ).

Ranking: $\ce{CH4}$ < $\ce{H2O}$ < $\ce{NaCl}$ < $\ce{C_{(diamond)}}$ .

Common Exam Trap: Assuming ionic compounds always have the highest melting points. Covalent-network solids (e.g., diamond, $\ce{SiO2}$ ) generally melt at significantly higher temperatures than ionic solids because melting requires breaking thousands of strong covalent bonds across the entire lattice.

Critical Thinking Questions

Diamond and graphite are both allotropes of carbon and both are covalent-network solids. Why does graphite conduct electricity while diamond acts as an insulator?
Why does magnesium oxide ( $\ce{MgO}$ , melting point $\pu{2852 degC}$ ) have a much higher melting point than sodium chloride ( $\ce{NaCl}$ , melting point $\pu{801 degC}$ )?
Explain mechanically why striking a piece of sodium chloride with a hammer shatters it, whereas striking a piece of copper wire deforms it without fracturing.

References

Sources returned by Google Search grounding during lesson verification.

12.7: Types of Crystalline Solids- Molecular, Ionic, and Atomic - Chemistry LibreTexts — chem.libretexts.org
10.5 The Solid State of Matter - OpenStax Chemistry 2e — openstax.org

Types of Crystalline Solids

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