1. Chemical Bonding

2. Valence electrons & Lewis dot structures

3. How do valence electrons play a role in bonding and making compounds?
NaCl
MgO
MgBr2
AlP
BF3
Al2O3
CCl4

4. Three types of Bonds Ionic Covalent Metallic Metal + Non-metal
Two non-metals
Metallic
Metallic atoms in an element

5. Metallic Bonds Any metal in its’ elemental form
Nuclei held together by “sea” of
mobile electrons shared
by all atoms
Explains shiny, malleable
and ductile properties
of metals

6. Ionic Bonding Metal + Non-metal
Sometimes polyatomic ions replace either
Electrons transferred from
cation to anion
Electrostatic attraction
holds ions together or
(+) charges and (-) charges
attract

7. Drawing Lewis structures of Ionic Compounds
NaCl
MgO
MgBr2
AlP
Al2O3

8. Lewis Structures of Molecular Compounds – see packet
Count all of the valence electrons in the compound.
Draw a “shell” of the atoms.
Connect all atoms with single bonds.
Give every atom a full octet with lone pairs.
Check your structure shows the correct number of electrons counted in step one.

9. Double and Triple Bonds - packet
If the number of electrons shown on Lewis structure is too high:
Erase a lone pair from each atom
Add another bond between atoms
This decreases the electrons shown by 2e-
If this STILL is too many electrons, do it again.
Maximum of three bonds between any two atoms.

10. Incomplete Octets - packet
Beryllium
Has a full shell with 4e- (two bonds)
Ex: BeCl2
Boron, Aluminum
Has a full shell with 6e- (three bonds)
Ex: BF3 or AlCl3

11. Expanded Octets - packet
Any atom in the 3rd period or lower, when the CENTRAL atom in a Lewis structure can have 10 or 12 e- completing its octet.
If Lewis structure has too few electrons shown, add lone pairs to the central atom and expand its octet.
Examples: PF5, SF6, ClF3, XeF2

12. Lewis Structures of Polyatomic Ions
Count up the valence electrons from each atom and adjust it based on charge of ion.
Draw Lewis structure as normal.
Ex: NH4+, H3O+, ClF4+, NO3-, CO3-2 OR

13. Two Types of Covalent Bonds
Covalent Bond - each atom contributes 1e-
Coordinate Covalent Bond
- one atom contributes both e- in bond
- most often in polyatomic ions

14. Bond Polarity Describes how electrons are shared.
The sharing of e- can be equal or unequal.
Consider electronegativity differences.

15. Ionic Bonds - Review Electronegativity Difference is > 1.7
SHOULD NOT BE USED TO IDENTIFY IONIC MATERIAL - used to determine strength of bond & determine the degree of ionic character

16. Non-polar Covalent Bonds
Diatomic Molecules
Made from two identical nonmetals – equal attraction for e-
Therefore, equal sharing between atoms
0 difference in electronegativity
Ex: H2, N2, O2, F2, Cl2, Br2, I2
Can be single, double or triple bonds (varying number of electrons shared)

17. Polar Covalent Bonds Atoms in bond are two different nonmetals
Therefore, there is an unequal attraction for the e-
difference in electronegativity
The higher the difference, the more polar the bond, and this correlates to the strength of the bond
Electrons are attracted to one atom more than the other
E- spend more time around nucleus of atom of higher electronegativity
Can be single, double or triple bonds

18. Network Solids Special type of covalent bond
Limited to substances containing C & Si
C: Diamond, graphite, Buckyball
Si: Silicon carbide, Silicon dioxide (sand)

19. Energy is involved in Bonding!
Energy is RELEASED when bonds are formed.
Why? Nature proceeds to achieve lower energy states.
For ionic compounds, this is called LATTICE ENERGY.
High Lattice Energy = high attractions of ions

20. Bond Forming Energy Cont’d
For molecular compounds, this energy loss is represented as a drop in potential energy (PE)

21. Energy of breaking bonds
Energy is ABSORBED when bonds are broken.
Called Bond Energy
Equals the Lattice Energy for ionic compounds.
Bond strength correlates to bond energy.

22. Bond Type  Compound Properties
Metallic
Ionic
Covalent
Example
Potassium
Potassium Chloride (KCl)
Chlorine
Melting Pt.
63oC
770oC
- 101oC
Boiling Pt.
760oC
1500oC
- 34.6oC
Properties
soft, silvery solid
conductor as a solid
crystalline white solid
conductor when dissolved or molten
- greenish, yellow gas
- not a good conductor

23. Bonding Concept Map A “sea” of mobile electrons shared by all atoms
made of
Metallic
Bonds
hold
Positively charged
nuclei of metals
together