Network covalent solids are materials in which atoms are joined by covalent bonds in a continuous lattice rather than as separate molecules. This bonding pattern makes many of them extremely hard, chemically stable, and difficult to melt. They are important in gemstones, cutting tools, electronics, glass, ceramics, and geological minerals.
Diamond, graphite, and silicon dioxide show how different network shapes can produce very different properties.
Understanding Chemistry: Network Covalent Solids
The arrangement of bonds matters as much as the type of atom. Carbon provides a clear example. In diamond, the bonds point in several directions, locking every atom into a rigid three-dimensional framework.
A force applied to the crystal must strain or break bonds throughout that framework. This is why diamond resists scratching.
It can still crack if hit in certain directions, because crystals have planes where splitting is easier. Hardness does not mean that a material cannot break.
Graphite is made only of carbon too, yet it behaves very differently. Within each sheet, carbon atoms are strongly joined. Between neighboring sheets, the attraction is much weaker.
The sheets can slide over one another when graphite is rubbed on paper or used as a dry lubricant. Pencil "lead" is actually a graphite and clay mixture, not the metal lead.
Small graphite layers detach and remain in the tiny spaces of the paper surface. This layered structure is a useful reminder that a substance cannot be predicted from its element name alone.
Bonding in graphite affects electricity as well. Each carbon atom has one electron that is not held in a single bond. These electrons can move through a sheet, so graphite conducts electricity along its layers.
Diamond has all of its outer electrons involved in strong bonds, so it usually does not conduct electricity. This difference is important in technology. Graphite is used for electrodes in batteries and electrolysis cells.
Diamond can be used where electrical insulation, heat transfer, or wear resistance is needed. The direction of a crystal can change its properties, especially for graphite.
Silicon dioxide forms the main structure of quartz and is a major part of sand. Its network contains silicon atoms linked through oxygen atoms. When molten silica cools slowly, atoms can settle into an ordered quartz crystal.
When it cools quickly, they may form glass instead. Glass has no repeating crystal pattern over long distances, but its atoms are still held by a connected covalent network. This explains why ordinary window glass is hard yet can shatter.
When studying these solids, connect structure to each property. Track how many bonds each atom makes, whether the bonds form sheets or a three-dimensional framework, and whether any electrons are free to move. These details explain melting behavior, hardness, electrical conduction, and the way a material breaks.
Key Facts
- Network covalent solids do not contain separate molecules, so there is no simple molecular formula for the whole solid.
- High melting points occur because many strong covalent bonds must be broken to change the solid into a liquid.
- Diamond: each carbon atom forms 4 covalent bonds in a tetrahedral 3D network.
- Graphite: each carbon atom forms 3 covalent bonds in flat hexagonal sheets, with weak forces between sheets.
- Silicon dioxide: each Si atom bonds to 4 O atoms and each O atom bonds to 2 Si atoms, giving an extended SiO2 network.
- Typical bond energy comparison: covalent network bonds are much stronger than intermolecular forces, so network solids usually melt at far higher temperatures than molecular solids.
Vocabulary
- Network covalent solid
- A solid made of atoms connected by covalent bonds in a continuous repeating structure.
- Covalent bond
- A chemical bond in which atoms share pairs of electrons.
- Lattice
- An ordered repeating arrangement of particles in a solid.
- Tetrahedral
- A shape in which one central atom is bonded toward the corners of a four-sided pyramid.
- Allotrope
- A different structural form of the same element, such as diamond and graphite for carbon.
Common Mistakes to Avoid
- Calling diamond a molecule is wrong because diamond is one continuous covalent network, not a collection of separate carbon molecules.
- Assuming all carbon solids have the same properties is wrong because diamond and graphite have different bonding arrangements and therefore different hardness and conductivity.
- Explaining graphite's slipperiness by weak covalent bonds within the sheets is wrong because the covalent bonds in each sheet are strong, while the forces between sheets are weak.
- Writing silicon dioxide as a small SiO2 molecule is misleading because SiO2 represents the atom ratio in an extended network, not an isolated molecule under normal solid conditions.
Practice Questions
- 1 Diamond has each carbon atom bonded to 4 other carbon atoms. How many covalent bonds are connected to 25 carbon atoms if each bond is shared by 2 atoms?
- 2 A sample of silicon dioxide contains 120 oxygen atoms. Using the formula ratio SiO2, how many silicon atoms are present?
- 3 Diamond and graphite are both made only of carbon. Explain why diamond is very hard while graphite is soft and slippery, using their bonding structures.