Diamond & Graphite
Section: 2. Atoms, Elements & Compounds | Syllabus: Cambridge AS Level Physics 9702
Giant Covalent Structures Giant Covalent Structure A continuous network of atoms bonded together by strong covalent bonds throughout the entire structure. Also called macromolecular or network structures.
Diamond and graphite are both allotropes of carbon - different structural forms of the same element. Despite being made of the same atoms, they have very different properties due to their different structures.
Allotropes Different structural forms of the same element in the same physical state. For carbon: diamond and graphite. Diamond Structure Atomic Arrangement Each carbon atom is bonded to 4 other carbon atoms All bonds are strong covalent bonds Atoms arranged in a tetrahedral (pyramid) shape Forms a giant 3D network extending throughout the entire crystal Very rigid structure with no free electrons Diamond Structure: Each C atom forms 4 single covalent bonds Tetrahedral arrangement All outer electrons used in bonding No free electrons available 3D giant structure continues throughout crystal Properties of Diamond 1.
Very High Melting Point (3550°C) Explanation Why: Strong covalent bonds hold every atom throughout the entire giant structure. To melt diamond, you must break many strong covalent bonds. Result: Enormous amount of energy needed, so diamond has an extremely high melting point.
2. Extremely Hard Explanation Why: Rigid 3D network of strong covalent bonds makes it very difficult to break or scratch the structure. Result: Diamond is the hardest naturally occurring substance. Used in cutting tools and drill bits.
3. Does NOT Conduct Electricity Explanation Why: All four outer electrons of each carbon atom are used in bonding. There are no free electrons available to move and carry electric charge. Result: Diamond is an electrical insulator.
4. Insoluble in Water Explanation Why: The covalent bonds throughout the structure are far too strong to be broken by water molecules. Result: Diamond does not dissolve in any solvent. Property Diamond Explanation Melting Point 3550°C (very high) Strong covalent bonds throughout giant structure Hardness Hardest natural substance Rigid 3D network of strong bonds Electrical Conductivity Does not conduct No free electrons (all used in bonding) Solubility Insoluble Bonds too strong to break Uses Cutting tools, drill bits, jewelry Extreme hardness and brilliance when cut Graphite Structure Atomic Arrangement Each carbon atom is bonded to 3 other carbon atoms Forms flat hexagonal layers (like honeycomb sheets) Strong covalent bonds within layers Weak intermolecular forces between layers One electron per carbon atom is delocalized (free to move) Layers can slide over each other easily Graphite Structure: Each C atom forms 3 covalent bonds Hexagonal layers (like chicken wire) Strong bonds WITHIN layers Weak forces BETWEEN layers One free electron per C atom (delocalized) Layers slide easily Properties of Graphite 1.
High Melting Point (3652°C) Explanation Why: Strong covalent bonds within the layers must be broken to melt graphite. Result: Very high melting point, though slightly higher than diamond. 2. Soft and Slippery Explanation Why: Weak intermolecular forces between layers allow layers to slide over each other easily.
Result: Graphite feels slippery and can be used as a lubricant. Also used in pencils - layers rub off onto paper. 3. DOES Conduct Electricity Explanation Why: Each carbon atom uses only 3 of its 4 outer electrons for bonding.
The fourth electron is delocalized (free to move throughout the structure). Result: Free electrons can carry electric charge, so graphite conducts electricity. This is unusual for a non-metal! 4. Low Density (compared to diamond) Explanation Why: Layers are relatively far apart with large spaces between them.
Result: Graphite is less dense than diamond. Property Graphite Explanation Melting Point 3652°C (very high) Strong covalent bonds within layers Hardness Soft, slippery Weak forces between layers allow sliding Electrical Conductivity DOES conduct Delocalized electrons free to move Density Lower than diamond Spaces between layers Uses Pencils, lubricants, electrodes Layers slide easily; conducts electricity Comparing Diamond and Graphite Feature Diamond Graphite Element Carbon (C) Carbon (C) Type Allotrope of carbon Allotrope of carbon Structure Type Giant covalent (3D network) Giant covalent (layered) Bonding per C atom 4 covalent bonds 3 covalent bonds Arrangement Tetrahedral 3D network Hexagonal flat layers Free electrons?
No (all 4 used in bonding) Yes (1 per atom delocalized) Melting Point 3550°C 3652°C Hardness Extremely hard Soft and slippery Electrical Conductivity Does NOT conduct DOES conduct Density Higher (3.5 g/cm³) Lower (2.3 g/cm³) Uses Cutting tools, drill bits, jewelry Pencils, lubricants, electrodes, batteries Why Different Properties?
Key Concept Diamond and graphite are both made of carbon atoms, but they have completely different properties because of their different structures and bonding arrangeme…
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