Difficulty: Intermediate | Prerequisites: Atomic structure, electron configurations, valence electrons, octet rule, basic understanding of intermolecular forces.
This topic sits at the intersection of bonding theory and material properties. You are moving from individual molecules to extended networks of atoms, where bonding models (metallic, covalent network) explain bulk properties like hardness, conductivity, and melting point. You should already be comfortable with valence electrons, the octet rule, and the distinction between ionic, covalent, and metallic bonding. If those feel shaky, revisit your notes on electron configurations and periodic trends before diving in.
Copper, graphite, and diamond are all network solids, but their bonding models differ and that difference drives their physical properties. Metallic bonding (copper) gives conductivity and malleability; sp2 covalent sheets with delocalised pi electrons (graphite) give conductivity and softness; a full sp3 covalent network (diamond) gives extreme hardness and zero conductivity.
Network solid
A substance in which atoms are bonded together in a continuous, extended structure rather than existing as discrete molecules. Think of it as one giant molecule that keeps going in every direction.
Metallic bonding
Bonding model in which metal cations sit in a regular lattice and valence electrons are delocalised across the entire structure in molecular orbitals that span the system. In simple terms, the metal nuclei are fixed in place while the electrons flow freely like a sea around them.
Delocalised electrons
Electrons that are not confined to a single bond or atom but are shared across many atoms through extended molecular orbitals. These are the electrons responsible for electrical conductivity in metals and graphite.
sp2 hybridisation
A hybridisation state where one s orbital and two p orbitals mix to form three equivalent hybrid orbitals in a trigonal planar arrangement, leaving one unhybridised p orbital. This is the bonding setup in each carbon of graphite.
sp3 hybridisation
A hybridisation state where one s orbital and three p orbitals mix to form four equivalent hybrid orbitals in a tetrahedral arrangement. This is the bonding setup in each carbon of diamond.
London dispersion forces (LDFs)
Weak intermolecular forces arising from temporary dipoles in electron clouds. In graphite, LDFs are the only thing holding the 2D sheets together, which is why the sheets slide so easily.
Continuous band
When a very large number of molecular orbitals overlap in energy, they form a near-continuous range of energy levels. In metals, this band allows absorption and emission of photons across all visible wavelengths, which is why metals are shiny.
Structure: A lattice of regularly spaced nuclei and core electrons. Valence electrons occupy molecular orbitals that span the entire structure, making them delocalised.
Properties explained by bonding:
High melting point: strong metallic bonds throughout the lattice require significant energy to break.
Malleable and ductile: when the lattice is deformed, the metal cores move relative to each other, but the delocalised electrons redistribute to keep holding the nuclei together. The "electron glue" does not break when you bend the metal.
Shiny: the continuous band of molecular orbital energies allows copper to absorb and re-emit photons across a wide range of energies. Copper's particular band structure gives it a characteristic colour.
Conducts electricity: delocalised electrons are free to move through the structure under an applied voltage.
Structure: Two-dimensional sheets of sp2 hybridised carbon atoms. Each carbon forms three sigma bonds in a trigonal planar arrangement. The leftover (unhybridised) p orbital on each carbon overlaps with neighbours to form a band of delocalised pi molecular orbitals extending across the entire sheet.
Properties explained by bonding:
High melting point: covalent bonds within the sheets are strong.
Soft and slippery: sheets are held together only by London dispersion forces (LDFs), which are weak. The sheets slide over one another easily.
Shiny: the delocalised pi electrons form a continuous band, similar to metals, allowing photon absorption and emission.
Conducts electricity: the delocalised pi electrons are free to move within a sheet, carrying current.
Structure: A three-dimensional network of sp3 hybridised carbon atoms. Every carbon forms four sigma bonds in a tetrahedral arrangement. All bonds are localised covalent bonds with no delocalised electrons.
Properties explained by bonding:
High melting point: a dense 3D network of strong covalent bonds is very hard to disrupt.
Extremely hard: breaking diamond means breaking covalent bonds in every direction. There is no easy plane of cleavage (though it is brittle and can fracture along crystal planes).
Does not conduct electricity: every electron is locked in a localised bond between two atoms. No free electrons, no current.
Translucent: without a continuous band of energy levels, diamond does not absorb visible light the way metals do. Light passes through.
Students often think graphite is soft because its bonds are weak. They are not. The covalent bonds within each sheet are very strong. Graphite is soft because the sheets are held together by weak LDFs and can slide past each other.
Students sometimes assume that all network solids are hard. Graphite is a network solid and it is soft, precisely because of its 2D sheet structure.
A common error is thinking diamond does not conduct electricity because it has weak bonds. The bonds in diamond are strong. Diamond is an insulator because its electrons are all localised in covalent bonds, leaving none free to carry a current.
Students may confuse "delocalised electrons" with "free electrons." Delocalised electrons are still bound to the system, they just occupy molecular orbitals that span many atoms rather than sitting between two.
Expect questions asking you to explain a physical property (conductivity, hardness, lustre) in terms of the bonding model. The answer always links back to whether electrons are delocalised or localised, and whether the structure is 2D or 3D.
Comparing copper, graphite, and diamond is a classic exam setup. Be ready to fill in a properties table and justify each entry from bonding.
Hybridisation is testable: know that graphite is sp2 (trigonal planar, leftover p orbital for pi bonding) and diamond is sp3 (tetrahedral, all bonds localised).
"Why is graphite a conductor but diamond is not?" appears frequently. The answer: graphite has delocalised pi electrons from unhybridised p orbitals; diamond does not.
True or false: Diamond conducts electricity because it has strong covalent bonds. (False: strong bonds do not equal conductivity. Electrons must be delocalised.)
Fill in the blank: Graphite is soft because its 2D sheets are held together by ______ , which are weak. (London dispersion forces / LDFs)
True or false: Copper is malleable because its delocalised electrons can redistribute when the lattice is deformed. (True)
Fill in the blank: Carbon atoms in graphite are ______ hybridised, while carbon atoms in diamond are ______ hybridised. (sp2; sp3)
True or false: Both graphite and copper conduct electricity, but for different structural reasons. (True: copper via metallic bonding and a sea of delocalised electrons; graphite via delocalised pi electrons in its sheets.)
Q: Explain why copper is shiny in terms of its electronic structure.
A: The molecular orbitals in copper overlap to form a continuous band of energy levels. This allows copper to absorb photons across a wide range of energies and re-emit them, producing a shiny appearance.
Q: Graphite conducts electricity but diamond does not. Both are made entirely of carbon. Explain.
A: In graphite, each carbon is sp2 hybridised, leaving one unhybridised p orbital per atom. These p orbitals overlap to form delocalised pi molecular orbitals across each sheet, allowing electrons to move and carry current. In diamond, each carbon is sp3 hybridised with all four valence electrons locked in localised sigma bonds. No electrons are free to move, so diamond is an insulator.
Q: Why is graphite soft and slippery while diamond is one of the hardest known materials?
A: Graphite consists of 2D sheets held together by weak London dispersion forces. The sheets slide over each other easily. Diamond is a 3D network of covalent bonds in every direction, with no weak plane to slip along.
Q: A student claims that metals are malleable because their bonds are weak. Correct this statement.
A: Metallic bonds are not weak. Metals are malleable because when the lattice is deformed, the delocalised electrons redistribute around the shifted nuclei, maintaining the bonding. The electron sea acts as a flexible glue that holds the structure together even as it changes shape.
This material connects directly to molecular orbital theory and band theory, which you will encounter in more depth later. Understanding why some materials conduct and others do not is the foundation for semiconductor physics and materials science. The hybridisation concepts here (sp2 vs sp3) reappear throughout organic chemistry when you discuss molecular geometry and reactivity.
Network solid, metallic bonding, covalent network solid, delocalised electrons, electron sea model, sp2 hybridisation, sp3 hybridisation, London dispersion forces, LDFs, van der Waals forces, band theory, continuous band, copper properties, graphite properties, diamond properties, electrical conductivity in solids, malleability, allotropes of carbon, CEM 141, general chemistry bonding