1. Chemical Bonding Why & How Atoms Combine
Ionic (Electrovalent) Bonds Ionic Compounds
Covalent Bonds
Simple Covalent Structures
Giant Covalent Structures
Metallic Bonds Metals
How Structures Determine Physical Properties of Substances

2. Electron Configurations of the Group 0 Elements (Rare/Noble Gases)
WHY DO ATOMS COMBINE?
Chemical bonds form when atoms react to achieve a stable electron configuration
Electron Configurations of the Group 0 Elements (Rare/Noble Gases)
Helium Neon Argon Krypton 2
2.8
2.8.8
These atoms are stable & generally do not combine with other atoms  elements exist as monoatomic gases.
This is because the atoms have a stable electron configuration with
happy atoms
a) a full outermost electron shell OR
b) eight electrons in the outermost electron shell
All other atoms need to achieve stability by attaining a stable electron configuration in the outermost shell.
This is done by gaining, losing or sharing electrons with other atoms.

3. Loss of Electrons
Let us consider the sodium atom.
Na 2.8.1
How do you think the Na atom achieves stability?
For sodium to have a stable octet of electrons in its outermost shell it can do one of 2 things: pick up 7 new electrons or give one up. Which is easier? 
When the sodium atom loses an electron, it becomes positively charged  Na
(an atom is electrically neutral and the number of protons in an element never changes, so after losing an electron sodium will have one more positively charged proton than it does negatively charged electrons and becomes a sodium ion ). 

4. Gain of Electrons
Now let us consider the chlorine atom.
Cl 2.8.7
How do you think the Cl atom achieves stability?
For chlorine to have a stable octet of electrons in its outermost shell it can do one of 2 things: lose 7 electrons (which is a very difficult thing to do) or pick up an electron. Which is easier?
When a chlorine atom gains an electron, it becomes negatively charged  Cl–
(an atom is electrically neutral and the number of protons in an element never changes, so after gaining an electron chlorine will have one more negatively charged electron than it does positively charged protons and becomes a chloride ion). 

5. Formation of Ionic Bonds
Thus chlorine and sodium are a perfect match for each other.  One needs an electron and the other wants to lose an electron. So when a chlorine atom and a sodium atom come together there is a transfer of an electron from the sodium to the chlorine atom. The sodium atom loses an electron to form a positive ion (cation) while the chlorine atom gains the electron to form a negative ion (anion).
The opposite charges on the Na+ and Cl– ions will cause a strong electrostatic attraction between the ions and an ionic bond is formed.  This is why the formula of sodium chloride is NaCl.

6. Na atom 2.8.1
Cl atom 2.8.7
Loses an electron
Gains an electron from Na
Na+ ion 2.8
Cl– ion 2.8.8 + –
electrostatic
attraction
NaCl
Na+
Cl–

7. So what determines the type of atoms which react to form ionic bonds and therefore ionic compounds?
Elements in Group I, II and III of the Periodic Table (Metals) tend to lose electrons to form positive ions (cations)
Elements in Group IV, V, VI and VII of the Periodic Table (Non-metals) tend to gain electrons to form negative ions (anions)
Thus, ionic bonds are formed between metals and non-metals.
Examples of ionic compounds: MgO, CaCl2, KOH

8. Crystal Lattice Structure of NaCl
A crystal lattice of NaCl consists of a large number of sodium and chloride ions held together by strong electrostatic attraction.
Each sodium ion is surrounded by 6 chloride ions and each chloride ion is surrounded by 6 sodium ions. Thus we say that the coordination number for this lattice is 6.

9. Sharing of Electrons
The individual atoms of chlorine have only seven outermost electrons, but want eight. By contributing one electron to each other, the two atoms achieve a stable electron configuration. A chlorine molecule is formed.

10. Cl atom 2.8.7 Cl atom 2.8.7 sharing of electrons or Clx o xx oo
Cl – Cl (Cl2 molecule)

11. So what determines the type of atoms which react to form covalent bonds?
Covalent bonding occurs between non-metallic elements which have 4, 5, 6, or 7 electrons in the outermost shell. When two non-metals come together, both want to gain electrons so they contribute an equal number of electrons to each other. One covalent bond is made up of two electrons.
Now complete Exercise 2 – Covalent Bonding in simple molecules.
Double covalent bonds are found in carbon dioxide, oxygen and ethene molecules.
Triple covalent bonds are found in nitrogen and ethyne molecules.

12. H atom can lose, gain or share single electron
Now consider the hydrogen atom. There are three ways this atom can achieve stability. • + H2 + • + – + • +
H atom can lose, gain or share single electron
H+ ion
H– ion
Exists in metal hydrides e.g. CaH2
Exists in acids e.g. HCl

13. Both atoms need to gain electrons – equal sharing of electrons
Non-metal & Non-metal
Metal & Non-metal
Ionic or Covalent?
One atom needs to gain electrons & one atom needs to lose electrons – transfer of electrons
Both atoms need to gain electrons – equal sharing of electrons
How many electrons?
How many atoms of each element required so that both achieve stability?

14. Physical Properties of Ionic & Simple Covalent Compounds
Ionic Compounds
Simple Covalent Compounds
Usually liquids or gases at room temperature
Usually solids at room temperature
High melting & boiling points due to strong electrostatic attraction between ions
Low melting & boiling points due to weak intermolecular forces (Van Der Waal’s forces) between molecules
Ionic compounds with very high melting points are refractory materials e.g. MgO is used to line the insides of industrial furnaces
Note: strong covalent bonds exist between atoms of a molecule
Conduct electricity in molten (liquid) or solution (aqueous) state as ions are mobile but not in solid state where ions are in fixed positions
Do not conduct electricity due to absence of mobile charged particles
Generally soluble in water but not in organic solvents
Generally soluble in organic solvents but insoluble in water

15. Simple covalent compounds (many molecules)
We have learnt about covalent compounds with simple molecular structures.
weak Van Der Waal’s forces between molecules o C
strong covalent bonds within the molecule
Now let us look at macromolecular structures in which atoms are joined together with strong covalent bonds to form one large molecule.

16. Giant Molecular Structures (Giant Covalent Structures)
Tetrahedral Structures
1. Silicon dioxide (silica) Each silicon atom is joined to 4 oxygen atoms in a tetrahedral arrangement by strong covalent bonds.
silicon dioxide
2. Silicon Each silicon atom is joined to 4 other silicon atoms in a tetrahedral arrangement by strong covalent bonds.
Learn how to draw this.
diamond/silicon
3. Diamond Each carbon atom is joined to 4 other carbon atoms in a tetrahedral arrangement by strong covalent bonds.

17. Giant Molecular Structures (Giant Covalent Structures)
Allotropes of Carbon
Different forms of an element are known as allotropes.
Diamond, fullerene (buckyballs) and graphite are allotropes of carbon.
Graphite consists of carbon atoms arranged in hexagonal layers. Each carbon atom is bonded to three other carbon atoms. Strong covalent bonds hold carbon atoms together within layers while Weak Van der Waal’s forces exist between layers.

18. Giant Molecular Structures (Giant Covalent Structures)
Physical properties of diamond
Physical properties of graphite
Soft and slippery material due to weak Van Der Waal’s forces between hexagonal layers – layers can slide over each other easily  soft and slippery
Used as a lubricant and as pencil lead
Hard, rigid structure due to the strong covalent bond network between all carbon atoms.
Used in cutting tools and drill bits
High melting & boiling points due to the strong covalent bond network between atoms
Low melting & boiling points due to weak Van Der Waal’s forces between hexagonal layers
Cannot not conduct electricity as all four outermost electrons of each atom are utilised in covalent bonding. These electrons are not free to move.
Can conduct electricity parallel to layers as electrons found between layers are free to move (delocalised) due to weak Van der Waals forces between layers.
Used as electrodes in batteries.
Insoluble in all solvents
Silicon and silicon dioxide have similar properties as diamond due to their similar structures.

19. Metals Structure/Bonding Physical Properties + e
Metals consist of atoms closely packed together in orderly layers.
Metals are soft, malleable (can be flattened) & ductile (can be pulled into wires) as the orderly layers of atoms can slide over each other easily.
force
Metals are good conductors of electricity in both solid and molten state due to the mobility of the electrons.
Metals consist of positive ions in a sea of electrons. + e
Metallic bonding due to strong forces of attraction between positive ions and free electrons.
Metals have relatively high melting and boiling points.

20. Alloys
An alloy is a mixture formed when a small amount of another substance (metal or non-metal) is added to a metal.
pure metal
alloy
Why the need for alloys?
Alloys are stronger than the parent metal. The size of the atoms of the substance added is different from the size of the atoms of the parent metal. These new atoms disrupt the orderly layers of the parent metal atoms making it harder for the atoms to slide over each other.