# Properties of Solids

> AP Chemistry · AP Chemistry 2024-2027
> Source: https://www.owlsprep.com/study/ap-chemistry-u3-properties-of-solids/

This module covers classification of the 4 core solid categories, links between bonding forces and physical properties, crystalline vs amorphous structures, and common AP exam property comparison questions.

**Prerequisites:** [Ionic, covalent, and metallic bonding fundamentals](https://www.owlsprep.com/study/ap-chemistry-u2-chemical-bonding-intro/); [Intermolecular force classification and strength trends](https://www.owlsprep.com/study/ap-chemistry-u3-intermolecular-forces/)

## Learning objectives

- Classify solids as ionic, molecular, covalent network, or metallic based on their composition and bonding
- Relate bulk physical properties including melting point, hardness, and conductivity to inter and intramolecular forces in solid lattices
- Predict and justify relative property rankings for unknown solid samples
- Distinguish between crystalline and amorphous solid structural variants

## Core Classification of Solid Types

All solid substances tested on the AP exam fall into four distinct categories, defined entirely by the dominant force holding their constituent particles together. The category of a solid can be predicted almost entirely from its chemical formula, no experimental data required.

**Crystalline vs Amorphous Solids** — Crystalline solids have a perfectly repeating long-range lattice, while amorphous solids have only short-range order with no consistent repeating structure.

| Solid Class | Constituent Particles | Dominant Binding Force |
| --- | --- | --- |
| Ionic | Cations + Anions | Electrostatic attraction |
| Molecular | Discrete neutral molecules | Intermolecular forces |
| Covalent Network | Neutral nonmetal atoms | Continuous covalent bonds |
| Metallic | Metal cations | Delocalized metallic bonding |

**Worked example:** Classify each of the following pure substances as one of the four solid types at standard conditions: KBr, C(diamond), H2O(s), Zn

1. 1. KBr is made of a metal cation (K+) and nonmetal anion (Br-), so it is an ionic solid
2. 2. Diamond is an extended lattice of covalently bonded carbon atoms, so it is a covalent network solid
3. 3. Ice is made of discrete H2O molecules held together by hydrogen bonds, so it is a molecular solid
4. 4. Zinc is a pure transition metal with delocalized valence electrons, so it is a metallic solid

**Check your understanding**

Test your classification skills

1. What type of solid is SiO2 (quartz)?

   - Ionic
   - Molecular
   - Covalent Network
   - Metallic

   *Why:* SiO2 forms a continuous extended lattice of Si-O covalent bonds, not discrete molecules.

> **Exam tip:** AP graders will not award points for classification alone: you must explicitly reference the bonding force to justify your answer.

## Bonding Strength and Melting Point Trends

The melting point of a solid is directly proportional to the strength of the force holding its constituent particles together. Breaking stronger forces requires more thermal energy, leading to a higher melting temperature.

**Worked example:** Rank the following solids from lowest to highest melting point: Sucrose (C12H22O11), MgO, Na, Diamond

1. 1. Sucrose is a molecular solid held together by weak intermolecular forces, so it has the lowest melting point ~186°C
2. 2. Sodium is a metallic solid with low charge density Na+ ions, so it melts at ~98°C? No, wait: sodium melts at 98°C which is lower than sucrose? Correct order: Na < Sucrose < MgO < Diamond
3. 3. MgO is an ionic solid with +2 and -2 charged ions, leading to very strong electrostatic attraction, melting point ~2800°C
4. 4. Diamond has continuous strong covalent bonds, so it has the highest melting point >3500°C

**Exam command terms**

AP FRQ command terms for this topic have strict grading expectations

- **Justify your prediction** — You must name the specific force that is stronger/weaker, not just state the trend *(Instead of 'MgO has higher melting point than NaCl', write 'MgO has higher melting point than NaCl because Mg2+ and O2- have higher charge density than Na+ and Cl-, leading to stronger electrostatic attraction in the ionic lattice')*

> **tip**
>
> Ionic solid melting points scale directly with the product of the ion charges, and inversely with the sum of ionic radii, following Coulomb's law.

## Electrical Conductivity of Solids

A solid will only conduct electricity if it contains free, mobile charge carriers that can move through the lattice when a voltage is applied. No mobile charge carriers = zero conductivity.

**Worked example:** Explain why solid NaCl does not conduct electricity, but molten NaCl at 900°C conducts very well

1. 1. In solid NaCl, Na+ and Cl- ions are locked in a rigid repeating lattice with no freedom of movement
2. 2. No mobile charge carriers exist in the solid state, so conductivity is zero
3. 3. When melted, the ionic lattice breaks apart, and individual Na+ and Cl- ions are free to move through the liquid phase
4. 4. These free moving ions act as charge carriers, allowing the molten sample to conduct electricity efficiently

**Check your understanding**

1. Which of the following solids will conduct electricity at standard conditions?

   - Solid KCl
   - Solid Copper
   - Solid Sucrose
   - Solid Diamond

   *Why:* Metallic copper has delocalized free electrons that act as mobile charge carriers in the solid state.

## Alloys and Special Solid Structures

Alloys are homogeneous mixtures of two or more elements, where at least one is a metal. They are classified as substitutional (atoms of the second metal replace host metal atoms in the lattice) or interstitial (small nonmetal atoms fit in gaps between host metal atoms).

**Worked example:** Classify brass (mixture of Cu and Zn, similar atomic radii) and steel (mixture of Fe and small C atoms) as substitutional or interstitial alloys

1. 1. Zinc and copper have nearly identical atomic radii, so Zn atoms replace Cu atoms directly in the copper lattice: brass is a substitutional alloy
2. 2. Carbon atoms are far smaller than iron atoms, so C atoms fit in the gaps between Fe atoms in the iron lattice: steel is an interstitial alloy

## Common pitfalls

- **Wrong:** Claiming molecular solids have low melting points because weak covalent bonds are broken during melting
  - Why it fails: Melting molecular solids only breaks weak intermolecular forces between discrete molecules, not the strong intramolecular covalent bonds inside individual molecules
  - Correct: Explicitly reference weak intermolecular forces (LDF, dipole-dipole, H-bonding) as the cause of low melting points for molecular solids
- **Wrong:** Stating ionic solids conduct electricity in their solid crystalline form
  - Why it fails: Ions are locked in a rigid lattice with no freedom of movement, so no mobile charge carriers exist
  - Correct: Note ionic solids only conduct when molten or fully dissolved in a polar solvent, when ions are free to move independently
- **Wrong:** Confusing covalent network solids with large molecular solids like C60 fullerene
  - Why it fails: Fullerene is made of discrete C60 molecules held together by weak LDFs, not a continuous extended covalent lattice
  - Correct: Check if the structure is a single unbroken lattice or made of separate discrete molecules to distinguish the two classes
- **Wrong:** Ignoring graphite as a special covalent network solid variant
  - Why it fails: Graphite has a 2D layered structure with weak interlayer LDFs, so it cleaves easily even though its melting point is extremely high
  - Correct: Explicitly note the layered structure of graphite when discussing its unusual softness and conductivity
- **Wrong:** Ranking metallic melting points only by group number
  - Why it fails: Metallic bonding strength depends on the charge density of the metal cation and number of delocalized electrons, not just group placement
  - Correct: Reference Coulomb's law for metallic bonding to justify relative melting point trends for different metals

## Cheatsheet

| Solid Type | Dominant Force | Typical Melting Point | Solid State Conductivity | Common Example |
| --- | --- | --- | --- | --- |
| Ionic | Electrostatic ion attraction | High (> 500°C) | Zero | NaCl, MgO |
| Molecular | Intermolecular forces | Low (< 300°C) | Zero | Ice, Sucrose, Dry Ice |
| Covalent Network | Continuous covalent bonds | Extremely High (> 2000°C) | Zero (except doped Si) | Diamond, Quartz, Graphite |
| Metallic | Delocalized metallic bonding | Medium to High | Excellent | Copper, Iron, Sodium |

## What's next

Mastering properties of solids is a critical foundational skill for Unit 3, as you will extend these bonding-property relationships to liquids, phase changes, and solutions in upcoming modules. This classification framework is tested heavily in long-form AP FRQ questions that require particulate-level reasoning to compare two substances, a skill that makes up roughly 40% of your total exam score. You will also apply these rules to predict solubility trends and colligative property behavior later in the course.

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