Unit 3 · Properties of Substances and Mixtures

AP Chemistry Types of Solids

Ionic, molecular, covalent network, and metallic solids: what holds each together and why that sets its properties.

Every property question about a solid, whether it melts high or low, conducts, bends, or shatters, comes back to one thing: which particles make up the solid and what holds them together. This is topic 3.2 in the AP Chemistry course. Once you can classify a solid from its formula, the properties follow from the particle picture, and the exam asks you to explain them in exactly those terms.

The Four Types of Solids Compared

How to Identify the Type of Solid From a Formula

1. Only metal atoms? It is metallic: Cu, Fe, Na, brass.
2. A metal with nonmetals, or any polyatomic ion? It is ionic: NaCl, MgO, CaCO₃, and NH₄Cl, even though NH₄Cl has no metal.
3. C (diamond or graphite), Si, SiO₂, or SiC? These are the common covalent network solids in AP questions.
4. Any other compound of nonmetals only? It is molecular: H₂O, CO₂, I₂, sucrose (C₁₂H₂₂O₁₁).

The bonding behind steps 1 and 2 is covered in ionic vs covalent bonds, and the forces that hold molecular solids together in intermolecular forces.

Why Each Type of Solid Has Its Properties

• Ionic solids are brittle. A sharp blow shifts one layer of ions so that like charges line up next to each other, and the repulsion splits the crystal. Their high melting points come from strong Coulombic attraction, strongest for ions with large charges and small radii, the idea behind lattice energy.
• Molecular solids melt low. Melting only overcomes intermolecular forces; the covalent bonds inside each molecule stay intact. The stronger the intermolecular forces, the higher the melting point, so H₂O (hydrogen bonding) melts far above CH₄ (dispersion forces only).
• Covalent network solids melt very high. There are no separate molecules, so melting means breaking covalent bonds throughout the whole crystal. SiO₂ melts near 1,700 °C.
• Metals bend and conduct. The delocalized electrons move freely, carrying current and heat, and they keep holding the cations together when layers slide past one another, so metals are malleable instead of brittle. Alloys, which change these properties, are covered in the Unit 2 review.

Crystalline and Amorphous Solids

In a crystalline solid the particles repeat in a regular three-dimensional pattern, so it melts sharply at one temperature. An amorphous solid, such as glass, has no long-range order and softens over a range of temperatures. AP questions focus on the four types above; you are not expected to know specific crystal structures or unit cells.

Worked Example 1: Classifying Six Solids

Classify KBr, CO₂(s), SiC, Cu, NH₄Cl, and C₁₂H₂₂O₁₁ (sucrose).

• KBr: ionic. A metal and a nonmetal, made of K⁺ and Br⁻ ions.
• CO₂(s), dry ice: molecular. Separate CO₂ molecules held by dispersion forces, which is why it sublimes at −78 °C.
• SiC: covalent network. Every Si and C atom is covalently bonded into one continuous network.
• Cu: metallic. One metal element.
• NH₄Cl: ionic. Only nonmetals, but it contains the polyatomic ion NH₄⁺ paired with Cl⁻.
• Sucrose: molecular. Nonmetals only and no ions, so separate molecules held by hydrogen bonding and other intermolecular forces.

Worked Example 2: Ranking Melting Points of Four Solids

Rank SiO₂, NaCl, I₂, and CH₄ from highest to lowest melting point, and justify the order.

SiO₂ (about 1,700 °C) > NaCl (801 °C) > I₂ (114 °C) > CH₄ (−182 °C). SiO₂ is a covalent network, so melting breaks covalent bonds. NaCl is ionic, held by strong Coulombic attractions between ions. I₂ and CH₄ are both molecular and nonpolar, so only dispersion forces hold them; I₂ has far more electrons and a more polarizable electron cloud, so its dispersion forces are much stronger (see London dispersion forces).

Worked Example 3: Identifying a Solid From Its Properties

A shiny solid melts at 1,085 °C, conducts electricity as a solid, and can be drawn into wire. A second white solid melts at 801 °C, does not conduct as a solid, but conducts when dissolved in water. Identify the type of each.

The first conducts as a solid and is ductile, which only a metallic solid does; these are the properties of copper. The second conducts only once its particles are free to move in solution, the signature of an ionic solid; these are the properties of NaCl.

Common Types of Solids Mistakes

• Saying that melting ice or dry ice breaks covalent bonds. Only intermolecular forces are overcome; the molecules stay whole.
• Calling NH₄Cl or NH₄NO₃ molecular. Polyatomic ions make a solid ionic even without a metal.
• Saying ionic solids conduct. Only when molten or dissolved; the solid’s ions cannot move.
• Assuming every metal melts high. Metallic melting points vary widely: sodium melts at 98 °C and mercury is a liquid at room temperature.
• Explaining with “stronger bonds” without saying which kind. Name the interaction that has to be overcome: intermolecular forces, ionic bonds, covalent bonds, or metallic bonding.

The same particle reasoning explains vapor pressure and boiling points, the next topic in Unit 3. Particle-level explanations like these earn free-response points on every exam; to see how those points add up to a 1–5 score, use the AP Chem Score Calculator.

Related Resources

• AP Chem Score Calculator
• Unit 3 Review: Properties of Substances and Mixtures
• AP Chemistry Intermolecular Forces
• AP Chemistry Ionic vs Covalent Bonds
• AP Chemistry Lattice Energy
• AP Chemistry London Dispersion Forces
• AP Chemistry Real Gases
• AP Chemistry Full-Length Practice Test

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