Chemistry 2e · Liquids and Solids
The Solid State of Matter
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In 30 seconds
A solid is the phase of matter with a definite shape and a definite volume. In a solid, particles — atoms, ions, or molecules — are held by attractive forces strong enough that they can only vibrate around fixed positions; they cannot slide past one another. That is why a solid keeps its shape and does not flow.
The key insight of this topic is that not all solids are alike, and the differences are predictable. The most important split is between crystalline solids (particles in a repeating three-dimensional pattern) and amorphous solids (particles packed irregularly). Among crystalline solids, chemists recognize four classes — ionic, molecular, covalent network, and metallic — and each class has a characteristic set of physical properties (melting point, hardness, conductivity) that follows directly from the type of bonding holding its particles together.
Why this matters
Every material you handle is a solid whose behavior is set by its internal structure. Metals are shaped into wires and sheets because of metallic bonding; diamond and quartz survive heat and scratching because they are covalent networks; table salt dissolves in water because it is an ionic lattice Regular repeating arrangement of particles in a crystal Full entry →. Pharmacists care whether an active ingredient is crystalline or amorphous, because the solid form controls how fast a tablet dissolves in the body. And in the lab, an unknown solid's melting point, hardness, and conductivity quickly identify what it is — and how it should be handled.
The college version
Core Concepts
Crystalline solids repeat a pattern; amorphous solids do not
A crystalline solid Solid whose particles form a repeating 3-D pattern Full entry → has long-range order: particles sit at the points of a repeating three-dimensional lattice. Table salt, diamond, quartz, and most metals are crystalline. An amorphous solid Solid with only short-range order, like frozen liquid Full entry → has only short-range order — particles packed irregularly, like a liquid frozen in place. Glass and many plastics are amorphous. This difference shows up in melting: a crystalline solid melts at one sharp temperature, while an amorphous solid softens gradually over a range because there is no single repeating structure to break apart.
Ionic solids: a lattice of alternating charges
Ionic solids (NaCl, MgO, CaF₂) are built from positive and negative ions arranged alternately so each ion is surrounded by opposite charges. The strong electrostatic attractions give these solids high melting points (NaCl melts at about 801 °C), hardness, and brittleness: when struck, like-charged layers are forced past each other, repel, and the crystal shatters. As solids they do not conduct electricity because the ions are locked in place — but when melted or dissolved in water, the freed ions carry charge and conduct well.
Molecular solids: molecules held by weak intermolecular forces
Molecular solids consist of discrete molecules held together only by intermolecular forces — dispersion forces, dipole–dipole interactions, or hydrogen bonds. Ice, table sugar, dry ice (solid CO₂), and naphthalene are examples. Because intermolecular forces are far weaker than ionic or covalent bonds, molecular solids are soft and melt at low temperatures (ice at 0 °C, dry ice at about −78 °C). They are electrical insulators.
Covalent network solids: one giant molecule
In a covalent network solid Atoms joined by covalent bonds into one giant network Full entry →, every atom is bonded to its neighbors by covalent bonds, forming a single continuous network. Diamond and quartz (SiO₂) are classic examples. Melting such a solid means breaking covalent bonds throughout the network, so these materials are extremely hard with very high melting points (diamond approaches 3550 °C) and are generally insulators. Graphite is the famous exception: its carbon atoms form flat sheets, so electrons move within each sheet — graphite conducts electricity and acts as a solid lubricant.
Metallic solids: cations in a sea of electrons
In a metallic solid Positive cores surrounded by a delocalized electron sea Full entry →, atoms release their valence electrons into a delocalized "electron sea" shared by the entire crystal, while the positive cores sit in a regular array. This model explains the defining metallic properties: electrical and thermal conductivity (mobile electrons), malleability and ductility (layers of cations slide past one another without breaking the electron sea), and luster. Metals vary widely in hardness — sodium is soft enough to cut with a knife, while tungsten is extremely hard.
Band theory: why some solids conduct
Band theory refines the electron-sea model by treating a crystal's overlapping valence orbitals as continuous bands of energy levels. In a conductor (a metal), a band is only partially filled, so electrons move freely. In a semiconductor (silicon, germanium), a small energy gap separates the filled band from the empty conduction band, so conductivity rises sharply when the solid is heated or doped. In an insulator (diamond), the gap is large and essentially no electrons reach the conduction band.
How It Works / Step-by-Step Process: Identifying an Unknown Solid
- Note the melting behavior: sharp and high, sharp and low, or gradual softening?
- Test electrical conductivity in the solid state.
- Test whether the solid dissolves in water and whether that solution conducts.
- Note mechanical behavior: does it shatter (brittle), bend (malleable), or scratch easily?
- Match the evidence pattern to one of the four classes and confirm against known melting points.
Common Confusions
| Do not confuse | With | Difference |
|---|---|---|
| "Solid" and "crystalline" | Amorphous solids | All crystals are solids, but glass and many plastics are amorphous |
| "Hard" and "high melting point" | Soft metals | Sodium is soft yet melts near 98 °C; hardness and melting point differ |
| Ionic solids "not conducting" | Being electrical insulators | Ionic solids conduct when molten or dissolved — ions are just immobile in the solid |
| Graphite (a network solid that conducts) | All network solids being insulators | Graphite's layered structure lets electrons move within sheets; diamond does not conduct |
Eli explains
The same idea, in plain words
Explain it like I’m 10
A solid is like a school gym where everyone stays in their assigned spot and only wiggles in place. Some groups stand in a neat grid — a crystal; others are jumbled randomly — amorphous. Strong glue, like diamond, makes the group hard to break apart; weak glue, like ice, melts easily.
Worked example
Example 1: Classifying a solid from its properties
A white solid melts sharply at 801 °C, does not conduct as a solid, conducts when melted and when dissolved in water, and shatters when struck. What class of solid is it?
Reasoning. A sharp melting point rules out an amorphous solid, and lack of solid-state conductivity rules out a metal. A melting point near 800 °C rules out a molecular solid, while a covalent network solid would not conduct when melted. Conductivity appears exactly when the solid melts or dissolves — the signature of ions becoming mobile. The evidence matches an ionic solid; NaCl itself melts at 801 °C.
Example 2: Predicting relative melting points
Rank ice, NaCl, and diamond from lowest to highest expected melting point, and explain the trend.
Reasoning. Ice is molecular: melting breaks only hydrogen bonds between water molecules, so it melts at 0 °C. NaCl is ionic: melting breaks electrostatic attractions, so it melts near 801 °C. Diamond is a covalent network: melting would break covalent bonds throughout the crystal, so it melts near 3550 °C. The order is ice < NaCl < diamond, tracking the strength of the forces overcome — intermolecular forces, then ionic bonds, then a full network of covalent bonds.
Key takeaways
- A solid has definite shape and volume; its particles vibrate about fixed positions.
- Crystalline solids show long-range order and a sharp melting point; amorphous solids soften gradually.
- Four classes of crystalline solids: ionic, molecular, covalent network, metallic — properties follow from the bonding.
- Ionic solids: high melting point, hard, brittle; nonconducting as solids but conducting when molten or dissolved.
- Molecular solids: low melting points, soft, insulating; held together by intermolecular forces, not chemical bonds.
- Covalent network solids: very hard, very high melting points (diamond, quartz); graphite is a conducting exception.
- Metallic solids: conductive, malleable, ductile — the electron-sea model explains all three.
- Band theory: conductor (no gap), semiconductor (small gap), insulator (large gap).
Check yourself
5 review questions from the chapter. Try each one, then open the answer.
Name the four classes of crystalline solids with one example each.
Show answer
Ionic (NaCl), molecular (ice), covalent network (diamond), metallic (copper).
Why does glass soften gradually while table salt melts at one sharp temperature?
Show answer
Salt is crystalline — one repeating lattice breaks at a single temperature. Glass is amorphous — its varied local environments fail gradually.
Why do ionic solids conduct electricity when melted but not when solid?
Show answer
In the solid, ions are locked in the lattice; when melted, ions migrate freely and carry charge.
What structural feature makes metals malleable and ductile?
Show answer
The electron sea holds the structure together while layers of positive cores slide past one another.
Why does diamond have a far higher melting point than ice?
Show answer
Diamond is one giant covalent network — melting means breaking covalent bonds throughout the crystal. Ice is held together only by hydrogen bonds between separate molecules.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- crystalline solid
- Solid whose particles form a repeating 3-D pattern
- amorphous solid
- Solid with only short-range order, like frozen liquid
- lattice
- Regular repeating arrangement of particles in a crystal
- ionic solid
- Alternating cations and anions held by electrostatic forces
- molecular solid
- Discrete molecules held together by intermolecular forces
- covalent network solid
- Atoms joined by covalent bonds into one giant network
- metallic solid
- Positive cores surrounded by a delocalized electron sea
- band gap
- Energy gap between filled and empty electron bands
Sources & references
This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.
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