Covalent structures Each carbon atom in diamond makes four covalent bonds with other carbon atoms in all directions in a giant covalent structure, so diamond is very hard. As all the bonds need to break for it to melt, this requires a great deal of energy. Diamond has a very high melting point. All the outer electrons of carbon are used so it does not conduct electricity. 4
Explain the difference in electrical properties between diamond and graphite.
5
Explain why each carbon atom in diamond can form four covalent bonds.
2.12 DID YOU KNOW? Diamond does, however, conduct thermal energy. It does this by transmitting the energy down its covalent bonds as it is in a rigid lattice.
Atomic structure and covalent bonding Carbon atoms have four electrons in the outer electron shell. In diamond every carbon atom is joined to four others in a three-dimensional tetrahedral lattice. The atoms join together by strong covalent bonds that involve electron sharing. This arrangement gives strength in all directions, and requires a large amount of energy to break the bonds, giving a very high melting point of 3350 °C at high pressure. No free electrons exist in the structure, so diamond does not conduct electricity. Diamonds are useful as cutting tools and are so rare and expensive that techniques were developed to create synthetic diamonds, known as industrial diamonds. These are not as optically perfect as natural cut diamonds but have the same useful properties and are used in many industries as diamond is the hardest substance.
HINTS & TIPS Carbon atoms in diamond make strong bonds in all directions which is why diamonds are hard.
Figure 2.43  The properties of industrial diamonds make them ideal as cutting tools if not as jewels. 6
Explain how the properties of diamond make it so useful as an abrasive for grinding other materials.
Google search: 'diamonds'
75054_P054_091.indd 79
79
5/24/16 8:52 PM