CHEMISTRY FORM 2
- 1.1 Structure of the atom
- 1.2 Atomic Number and Mass Number
- 1.3 Isotopes
- 1.4 Energy levels and electron arrangement
- 1.5 Development of the Periodic Table
- 1.6 Relative Atomic Mass and Isotopes
- 1.7 Ion Formation
- 1.8 Chemical Formulae
- 1.9 Chemical Equations
- 2.1 Alkali metals (Group I elements)
- 2.2 Alkali Earth Metals (Group II elements)
- 2.3 Halogens (Group VII elements)
- 2.4 Noble gases (Group VIII elements)
- 2.5 Properties and Trends Across the Periodic Table
- 3.1 Bond
- 3.2 Ionic bond
- 3.3 Giant ionic structure
- 3.4 Covalent bond
- 3.5 Co-ordinate bond
- 3.6 Molecular structures
- 3.7 Giant covalent structures
- 3.8 Metallic Bond
- 3.9 Types of bond across a period
- 3.10 Oxides of elements in Period 3
- 3.11 Chlorides of Period 3 elements
- 4.1 What is a salt?
- 4.2 Types of salt
- 4.3 Solubility of salts in water
- 4.4 Methods of preparing salts
- 4.4.1 Reacting a Metal with an Acid
- 4.4.2 Reacting an Acid with a Base (Neutralization)
- 4.4.3 Reacting an Acid with a Carbonate (or hydrogencarbonate of metal)
- 4.4.4 Combining elements Directly (Direct Combination of elements)
- 4.4.5 Precipitation (Double decomposition)
- 4.5 Action of heat on salts
- 4.6 Uses of salts
- 5.1 Electrical conduction
- 5.2 Electrical conductivity of molten substances
- 5.3 Electrical conductivity of substances in aqueous state
- 5.4 Electrolysis
- 5.5 Applications of electrolysis
- 6.1 Allotropes of carbon
- 6.2 Chemical properties of carbon
- 6.3 Carbon (IV) oxide
- 6.4 Carbon (II) oxide (CO)
- 6.5 Large scale production of sodium carbonate and sodium hydrogencarbonate
- 6.6 Effect of carbon (II) oxide and carbon (IV) oxide on the environment
- 6.7 Carbon cycle
Chemical Bonding and Structure: Giant ionic structure
3.0 Chemical Bonding and Structure
3.3 Giant ionic structure
We have seen sodium chloride (NaCl) and copper (II) oxide as examples of ionic compounds.
Observe the demonstration involving strong heating of sodium chloride crystals, copper (II) oxide, and zinc oxide separately.
(courtesy Youtube-Strongly heating sodium chloride crystals by Joseph Rabari)
(courtesy Youtube-Strongly heating copper oxide by Joseph Rabari)
(courtesy Youtube-Strongly heating zinc oxide by Joseph Rabari)
- Is there any noticeable effect of heat on sodium chloride crystals, copper (II) oxide or zinc oxide?
- What can you say about the melting points of sodium chloride and copper (II) oxide?
- What can you say about the amount of heat energy required to melt sodium chloride or copper (II) oxide?
Answers to Questions 3.3(a)
Melting occurs when all the bonds are completely broken. It has been found that the amount of energy required to melt a sodium chloride crystal is over a billion trillion (1021) times larger than the energy required to break the bond between two ions (sodium ion and chloride ion).
From the information given, which of the following diagrams best represents the structure of sodium chloride crystal? Explain your choice.
Figure 3.3b:Test structures
Answers to Questions 3.3(b)
If sodium and chloride ions occurred in small separate units such as in (a), (b) and (d), the structure would have a low melting point because nothing ties such units together. They can be easily separated from one another. High melting point therefore indicates that trillions of the ions are bonded continuously (in a repeated pattern), one to the next. The resulting structure is therefore called a giant ionic structure (Figure 3.3(c)).
Figure 3.3(c): Giant ionic structure of sodium chloride
Note that, in this structure, each positive ion (Na+) is surrounded by negative ions (Cl-) and vice versa. Only then can the structure hold together. Why? This applies to all other ionic compounds. We have used balls (◯) to represent ions and sticks (__) to represent bonds. This is called a ball-and-stick model.
In the same manner as Figure 3.3(c), represent the giant ionic structures of copper (II) oxide (CuO) and lithium fluoride (LiF).
Answers to Questions 3.3(c)
It should be noted that in ionic bonding, the number of negative ions surrounding a positive ion and vice versa, does not depend on valency but size of ions. This is because an electric charge can attract any number of opposite charges around it.
Properties of ionic compounds
- They have giant ionic structures
- They have high melting and boiling points.
- They do not conduct electricity as solids, but conduct electricity in molten state or solution. This is the property used to identify ionic compounds.
In solids, the ions are locked together in ionic bonds so they are not free to move about. In molten state and solutions, the ions are free to move about and conduct electricity. This explains Property 3 above. Observe the demonstration with molten lead (II) bromide.
(courtesy Youtube-The electrolysis of lead bromide by David Read)
What range of temperatures is considered as low and high melting points?
Generally, temperatures a few hundred degrees above and below room temperature (25oC) are considered as low. High temperatures are around 1000 oC and higher.