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Not all covalent substances are small molecules. Some carbon and silicon structures are giant — billions of atoms joined by covalent bonds into one enormous network — and these have utterly different properties from the gases and liquids of the last lesson: diamond is one of the hardest natural materials, while graphite (also pure carbon) is soft enough to write with. This lesson, part of Topic C2 of your OCR Gateway Combined Science course, covers giant covalent structures, metallic bonding, alloys and a brief look at polymers. Throughout, the skill being tested is the same one as in the bonding lessons: explaining properties from structure and bonding.
By the end of this lesson you should be able to describe the structures of diamond, graphite and graphene, explain their properties from structure, explain metallic bonding and the properties of metals and alloys, and describe polymers in outline.
This lesson develops AO1 (recalling the structures of diamond, graphite, graphene, metals and polymers) and AO3 (explaining the properties of each — hardness, conductivity, why alloys are harder than pure metals — by reasoning from structure and bonding).
A giant covalent structure (or macromolecule) is a huge network of atoms all joined by strong covalent bonds. Because there are so many strong covalent bonds to break, these substances have very high melting points — quite unlike simple molecular substances. The most important examples are forms of carbon.
In diamond, each carbon atom forms four covalent bonds to four other carbon atoms, building a rigid, three-dimensional lattice that extends throughout the crystal. This structure gives diamond its properties:
In graphite, each carbon atom forms only three covalent bonds, arranging the atoms into flat layers of hexagons. The layers are held to one another only by weak intermolecular forces, and the fourth outer electron of each carbon is delocalised (free to move). This very different structure gives very different properties:
Diamond and graphite are both pure carbon — they are allotropes of carbon (different structural forms of the same element). Their wildly different properties come entirely from how the atoms are arranged and bonded, which is a perfect illustration of structure → property.
Exam Tip: The key contrast: in diamond each carbon makes 4 bonds (rigid, hard, no free electrons → no conduction); in graphite each carbon makes 3 bonds in layers that slide (soft lubricant) with 1 delocalised electron per atom (conducts). A common misconception is that graphite does not conduct — it does, because of its delocalised electrons. Both are allotropes of carbon.
Graphene is a single layer of graphite — one atom thick. It is very strong, light and an excellent conductor (it has delocalised electrons like graphite), which makes it promising for electronics and as a reinforcement in composite materials. Its combination of great strength, low mass and electrical conductivity in a sheet one atom thick is what makes it so interesting to scientists.
Silicon dioxide (SiO2, found in sand and quartz) is another giant covalent structure, with each silicon bonded to oxygen in a hard, high-melting lattice, similar in form to diamond.
In a metal, the atoms are packed in a regular giant lattice, and the outer electrons become delocalised — they leave the individual atoms and are free to move throughout the whole structure. This leaves the atoms as positive metal ions sitting in a "sea" of delocalised electrons. The strong electrostatic attraction between the positive ions and the sea of electrons is the metallic bond.
This structure explains the properties of metals:
Exam Tip: Metals conduct because of delocalised electrons that are free to move — not because the ions move (a common misconception). State "the metallic bond = electrostatic attraction between positive metal ions and a sea of delocalised electrons" to score the bonding mark.
A pure metal is often too soft for many uses, because its layers of identical atoms can slide over each other easily. An alloy is a mixture of a metal with one or more other elements (often another metal). Alloys are harder than the pure metal because the different-sized atoms distort the regular layers, so the layers cannot slide over each other as easily. Examples include steel (iron with carbon) and brass (copper with zinc).
Exam Tip: An alloy is a mixture, not a compound — a frequent misconception. It is harder than the pure metal because the different-sized atoms disrupt the layers, stopping them sliding. This is a common 2-mark question.
A polymer is a very long-chain molecule made by joining together many small molecules, called monomers, in a repeating pattern. Poly(ethene), for example, is made from thousands of ethene molecules joined into one long chain. Polymers include many everyday plastics — poly(ethene) in plastic bags, poly(propene) in containers, PVC in pipes.
The atoms along the chain are joined by strong covalent bonds, but neighbouring chains are held to one another by intermolecular forces. Because polymer molecules are so large and long, these intermolecular forces add up to more than in a small molecule, so polymers are usually solid at room temperature and have melting points that are higher than small molecules but lower than giant ionic or giant covalent substances. Most polymers do not conduct electricity, which is why plastics are widely used as electrical insulators.
Exam Tip: A polymer is a long chain of many repeating units (monomers) joined by covalent bonds. Its intermolecular forces are stronger than a small molecule's because the chains are so long — so polymers are solids with higher melting points than small molecules, but still lower than giant structures.
Explain why graphite can be used as a lubricant but diamond cannot.
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