Structure and bonding – Part 2
1. EXPLAIN THE DIFFERENCE BETWEEN SIMPLE AND GIANT MOLECULAR STRUCTURES.
Simple molecular crystals have strong covalent bonds between atoms but have weak forces of attraction between molecules, resulting in low melting and boiling temperatures. Examples include carbon dioxide, iodine and methane.
Giant structures of atoms and molecules, such as silicon dioxide (silica), diamond and graphite, have very high melting and boiling points due to strong covalent bonds throughout their three-dimensional network.
2. WHAT ARE ALLOTROPES? GIVE EXAMPLES.
Allotropes are different forms of the same element existing in the same physical state. This causes them to have different physical properties but the same chemical properties. For example, the element carbon can exist in different forms, including graphite, diamond, fullerene and charcoal. Other elements, such as phosphorus, sulphur, oxygen and silicon, also exist as allotrophes.
3. HOW DO THE STRUCTURES OF SODIUM CHLORIDE, GRAPHITE AND DIAMOND INFLUENCE THEIR PROPERTIES?
Diamond and graphite are giant molecular or macromolecular crystals. Diamond consists of carbon atoms tetrahedrally arranged and bonded by strong covalent bonds. Four carbon atoms are joined in a tetrahedral arrangement which is repeated throughout to give a three-dimensional structure.
Diamond is the hardest natural substance and has high melting (about 3500°C) and boiling points due to the strong covalent bonds throughout its 3D network. These bonds require a lot of energy to break. It is also a non-conductor of electricity since it has no free or mobile electrons to carry the electric charge.
Graphite consists of carbon atoms arranged in hexagonal rings and in layers. Each carbon atom is bonded to three other atoms arranged hexagonally in layers. These layers are held together by weak bonds which enable them to slide over each other. There are strong covalent bonds, however, between the carbon atoms in each layer. Since the carbon atom is bonded to only three others, it means that each carbon atom has a fourth electron not involved in bonding; that is, a free mobile electron. This causes graphite to conduct heat and electricity. Graphite is used as a lubricant since it is soft. It also has high melting and boiling points due to the strong covalent bonds throughout its giant molecular structure. Graphite and diamond are composed of carbon atoms but their structures are different, hence these solids are allotropes. They show the same chemical properties since they have the same element carbon, but the difference in their structure causes them to have different chemical properties.
Sodium chloride is an ionic solid in a crystal lattice with a giant ionic structure. Each sodium ion is surrounded by six chloride ions and each chloride ion is surrounded by six sodium ions. Strong forces of attraction exist between the negatively and positively charged ions, resulting in high melting and boiling points. These require large amounts of energy to break. The regular repeating lattice structure causes sodium chloride to be brittle as the regular pattern is disrupted. Sodium chloride is readily soluble in water (polar solvent) and conducts electricity when molten or in solution as the ions become mobile. Sodium chloride does not conduct in the solid state.
4. HOW DO THE PROPERTIES OF IONIC AND MOLECULAR SOLIDS DIFFER?
Molecular solids refer to covalent compounds which form solids, for example, iodine and dry ice. The properties of these solids differ from ionic solids.
5. WHY DO IONIC COMPOUNDS NOT CONDUCT ELECTRICITY IN THE SOLID STATE?
Ionic compounds are composed of ions (anions and cations). These are held together by strong electrostatic forces of attraction. In the solid state, these ions are not free to move (not mobile), hence they cannot carry an electric charge. When these ionic compounds are molten or dissolved in water, however, the ions are now mobile and can, therefore, conduct an electric charge.