Structure And Bonding Of Carbon
Overzicht
Sharpen a pencil and you are holding one of the softest solids in the laboratory. Look at a ring set with a diamond and you are looking at the hardest natural material on Earth. Yet nothing separates them at the level of the atom. Both are pure carbon: the same element, six protons in every nucleus, no impurity that matters. What differs is how the atoms are joined up, and that turns out to be everything.
This lesson takes that difference apart. You will count the bonds each carbon atom makes in a diamond and in a flake of pencil lead, hunt down the one spare electron that lets a pencil line carry a current, and work out why sheets held to each other by almost nothing are exactly what a lubricant needs. Then you will meet the carbon structures chemists have built since 1985: hollow cages of sixty atoms, sheets a single atom thick, and tubes stiff enough to be worth putting inside a tennis racket.
Doelstellingen
- The element carbon can form four covalent bonds.
- In diamond, each carbon atom forms four covalent bonds with other carbon atoms in a giant covalent structure, so diamond is very hard.
- In graphite, each carbon atom bonds to three others, forming layers. The layers are free to slide over each other because there are no covalent bonds between the layers and so graphite is soft and slippery. Students should be able to explain the properties of graphite in terms of weak forces between the layers.
- In graphite, one electron from each carbon atom is delocalised. These delocalised electrons allow graphite to conduct heat and electricity. Students should realise that graphite is similar to metals in that it has delocalised electrons.
- Carbon can also form fullerenes with different numbers of carbon atoms. Fullerenes can be used for drug delivery into the body, in lubricants, as catalysts, and in nanotubes for reinforcing materials, eg in tennis rackets. Students are only required to know that the structure of fullerenes is based on hexagonal rings of carbon atoms.
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- In graphite, how many other carbon atoms is each carbon atom covalently bonded to? A. One B. Two C. Three D. Four Answer: C
- Why is graphite soft and slippery? A. The covalent bonds within each layer are weak B. Only weak forces hold one layer to the next C. The carbon atoms are held together by ionic bonds D. Graphite contains delocalised electrons Answer: B
- Why does diamond not conduct electricity? A. Its atoms are held together by weak forces B. All four outer electrons of each carbon atom are used in covalent bonds C. It is transparent and colourless D. Its carbon atoms carry no charge Answer: B
- Which statement about fullerenes is correct? A. Their structure is based on hexagonal rings of carbon atoms B. They are giant covalent structures with no fixed formula C. Every fullerene contains exactly 12 carbon atoms D. Each carbon atom in a fullerene forms four covalent bonds Answer: A
- Graphite is similar to a metal because it: A. is shiny and can be bent into shape B. conducts electricity using delocalised electrons C. reacts with dilute acids to give hydrogen D. forms positive ions in solution Answer: B
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