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General Chemistry II · Intermolecular Forces Liquids Solids

Types of Solids

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On this page 8 sections
  1. In 30 seconds
  2. Why this matters
  3. The college version
  4. Eli explains
  5. Worked example
  6. Key takeaway
  7. Study tools
  8. Sources & references

In 30 seconds

Solids fall into two broad classes — crystalline (atoms/ions/molecules in a regular repeating lattice) and amorphous (disordered, no long-range order). Crystalline solids are further classified by the force holding them together: molecular (IMFs), ionic (electrostatic ion–ion), metallic (delocalized electron sea), and covalent-network (a continuous covalent bond network). The bonding type directly sets the solid's melting point, hardness, brittleness, and conductivity.

Why this matters

  • Materials selection: whether a part must conduct, withstand heat, or be lightweight determines which solid class is used (copper wires vs. ceramic insulators vs. diamond tool bits).
  • Electronics: silicon (covalent-network) is the backbone of semiconductors.
  • Drugs: most pharmaceuticals are molecular solids; their crystal packing affects solubility and bioavailability.
  • Geology: mineral hardness (Mohs scale) reflects the bonding type of the crystal.

The college version

Core Concept

Solids fall into two broad classes — crystalline (atoms/ions/molecules in a regular repeating lattice) and amorphous (disordered, no long-range order). Crystalline solids are further classified by the force holding them together: molecular (IMFs), ionic (electrostatic ion–ion), metallic (delocalized electron sea), and covalent-network (a continuous covalent bond network). The bonding type directly sets the solid's melting point, hardness, brittleness, and conductivity.

Key Ideas

Molecular solids

  • Held by IMFs (dispersion, dipole–dipole, H-bonds). Examples: dry ice (CO₂), ice (H₂O), sugar, iodine.
  • Soft, low melting points, generally nonconducting.

Ionic solids

  • Lattice of cations and anions held by Coulombic attraction. Examples: NaCl, CaF₂, MgO.
  • Hard, brittle, high melting points; conduct only when molten or dissolved (mobile ions).

Metallic solids

  • Metal cations immersed in a "sea" of delocalized valence electrons. Examples: Cu, Fe, Al.
  • Malleable/ductile, lustrous, high thermal and electrical conductivity, variable melting points.

Covalent-network solids

  • Atoms linked by covalent bonds in a continuous 3-D (or 2-D) network. Examples: diamond, graphite, quartz (SiO₂), silicon.
  • Very hard, very high melting points, poor conductors (except graphite, which conducts along its sheets).

Amorphous vs. crystalline

  • Crystalline: long-range periodic order; sharp melting point.
  • Amorphous: disordered (glass, rubber, many plastics); soften over a temperature range.

Equations and Variables

  • Coulombic lattice energy (ionic): E ∝ q₁q₂/r, where q₁, q₂ = ion charges, r = interionic distance. Larger charges and smaller ions → stronger lattice → higher melting point (e.g., MgO melts far higher than NaCl).
  • Metallic bonding: no single equation; conductivity arises from delocalized electrons that move freely under an applied field.
  • Conductivity types: ionic conduction requires mobile ions (molten/aqueous); metallic conduction requires mobile electrons (present at all temperatures in the solid).

How It Works

  1. Identify the bonding that holds the solid together — IMFs (molecular), ion–ion (ionic), electron sea (metallic), or covalent network (network).
  2. Translate bonding to properties. Stronger, more directional bonds → harder, higher-melting solids.
  3. Explain brittleness of ionic solids. Sliding a layer past its neighbor puts like charges next to like charges, which repel and shatter the crystal.
  4. Explain malleability of metals. The electron sea is not directional, so layers of cations can slide past each other without breaking bonds.
  5. Explain conductivity. Electrons flow freely in metals; in ionic solids, current needs ions that can move, which only happens when melted or dissolved.

Worked Example

Classify each solid and predict its properties: dry ice (CO₂), table salt (NaCl), copper wire (Cu), and diamond (C).

  1. CO₂ — molecular. Held by weak dispersion forces → sublimes at −78 °C (very low "melting"), soft, nonconducting.
  2. NaCl — ionic. Strong ion–ion lattice → high melting point (801 °C), hard but brittle, conducts when molten or in solution.
  3. Cu — metallic. Electron sea → malleable/ductile, excellent conductor, melting point 1085 °C.
  4. Diamond — covalent-network. Every carbon bonded to four others → extreme hardness and a very high melting point (>3500 °C), but an electrical insulator (all electrons locked in bonds).

How it works

  1. Identify the bonding that holds the solid together — IMFs (molecular), ion–ion (ionic), electron sea (metallic), or covalent network (network).
  2. Translate bonding to properties. Stronger, more directional bonds → harder, higher-melting solids.
  3. Explain brittleness of ionic solids. Sliding a layer past its neighbor puts like charges next to like charges, which repel and shatter the crystal.
  4. Explain malleability of metals. The electron sea is not directional, so layers of cations can slide past each other without breaking bonds.
  5. Explain conductivity. Electrons flow freely in metals; in ionic solids, current needs ions that can move, which only happens when melted or dissolved.

Common confusions

  • "All solids have high melting points." — Wrong. Molecular solids (dry ice, ice) melt or sublime at low temperatures because they're held only by IMFs.
  • "Ionic solids conduct electricity as solids." — Wrong. Their ions are locked in place; conduction requires melting or dissolving to free the ions.
  • "Metals are strong because they have strong covalent bonds." — Wrong. Metals are held by delocalized electrons (metallic bonding), which is why they are malleable rather than brittle.
  • "Glass is a crystalline solid." — Wrong. Glass is amorphous; it lacks long-range order and softens over a temperature range.
  • "Diamond and graphite are the same because both are pure carbon." — Wrong. Their bonding networks differ (3-D vs. sheet-like), giving opposite hardness and conductivity.

Quick review

  • Molecular (IMFs) → soft, low mp. Ionic (ion–ion) → hard, brittle, high mp.
  • Metallic (electron sea) → malleable, conductive. Covalent-network → very hard, very high mp.
  • Ionic solids conduct when molten/aqueous; metals conduct as solids.
  • Lattice energy ↑ with charge, ↓ with size.
  • Amorphous = no long-range order.
Eli, the EliExplains learning guide

Eli explains

The same idea, in plain words

Explain it like I’m 10

Imagine four ways to build a sandcastle wall. One uses stickers (molecular solids — weak and easy to knock over). One uses a neat alternating pattern of positive and negative magnets (ionic — strong, but shove a layer sideways and matching magnets suddenly face each other and repel, cracking the wall). One pours everything into a pool of glue (metallic — you can bend and shape it and it still holds, and the "glue" carries electricity). The last welds every grain to its neighbors with steel bars (covalent-network — nearly indestructible). (The analogy's limit: the "glue" in a metal is actually a shared sea of electrons, and the "magnets" are electrostatic charges, but the four different ways of holding things together map one-to-one onto the four solid types.)

Worked example

Worked Example

Classify each solid and predict its properties: dry ice (CO₂), table salt (NaCl), copper wire (Cu), and diamond (C).

  1. CO₂ — molecular. Held by weak dispersion forces → sublimes at −78 °C (very low "melting"), soft, nonconducting.
  2. NaCl — ionic. Strong ion–ion lattice → high melting point (801 °C), hard but brittle, conducts when molten or in solution.
  3. Cu — metallic. Electron sea → malleable/ductile, excellent conductor, melting point 1085 °C.
  4. Diamond — covalent-network. Every carbon bonded to four others → extreme hardness and a very high melting point (>3500 °C), but an electrical insulator (all electrons locked in bonds).

Key takeaways

  • ### High-Yield Facts
  • Four crystalline types: molecular (IMFs), ionic (ion–ion), metallic (electron sea), covalent-network (continuous bonds).
  • Ionic solids are hard but brittle (like-charge repulsion on slip).
  • Metals are malleable and ductile (nondirectional electron sea).
  • Ionic solids conduct only when molten or dissolved; metals conduct as solids.
  • Covalent-network solids (diamond, quartz) are extremely hard with very high melting points.
  • Lattice energy ↑ with charge, ↓ with ion size (MgO > NaCl).
  • Amorphous solids (glass) have no long-range order and soften over a range.

Keep learning

Ready to build on this? Continue to the next lesson.

Practice General Chemistry II

This lesson has no separate scored set. Practice draws from the subject’s question bank.

Study tools & related lessonsYou’ll learn to · Related

You’ll learn to

  • Distinguish crystalline from amorphous solids.
  • Classify solids as molecular, ionic, metallic, or covalent-network based on their bonding.
  • Predict typical properties (melting point, hardness, conductivity) from the solid type.
  • Explain why ionic solids are brittle and metallic solids are malleable and conductive.

Sources & references

  1. OpenStax, *Chemistry 2e*, "10.5 The Solid State of Matter." https://openstax.org/books/chemistry-2e/pages/10-5-the-solid-state-of-matter
  2. OpenStax, *Chemistry 2e*, "10.6 Lattice Structures in Crystalline Solids." https://openstax.org/books/chemistry-2e/pages/10-6-lattice-structures-in-crystalline-solids
  3. OpenStax, *Chemistry 2e* (book home). https://openstax.org/details/books/chemistry-2e
  4. PubChem, "Water" (compound 962). https://pubchem.ncbi.nlm.nih.gov/compound/962
  5. American Chemical Society. https://www.acs.org/

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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