1. Structures and Properties
Chemistry 11

2. Covalent Bonds and Molecules
When we consider covalent compounds, it is often important to consider the molecular geometry
This is vary with the number of lone and shared pairs as well as single, double and triple bonds
We will use Lewis Structures to model simple molecules

3. Pairs – Lone or Shared?
There are two types of electron pairs that we consider in molecules
Lone – these are pairs that belong entirely to one atom
Shared – these are pairs that are shared between two of the atoms in the molecule

4. Drawing Lewis Structures – Simple Molecules
Determine the total number of valance electrons in all atoms
Draw a skeleton structure where the atom with the largest number of unpaired electrons is the central atom
Place lone pairs around all of the atoms except the central atom to form octets (except hydrogen which has one pair)
If all valance electrons have not been used, add long pairs to central atom
If all valance electrons have been used but the central atom does not have an octet, move lone pairs to form double or triple bonds as required

5. Example - Formaldhyde O O H H C C H H CH2O Valance Electrons
H – 2x1
O – 6
Total - 12 O O H H C C H H

6. Co-ordinate Covalent Bonds
Normally in a covalent bond, each atom contributes one electron to the shared pair
However, it is possible that one atom contributes both electrons; this is known as a co-ordinate covalent bond
This behaviour explains how ammonium is formed

7. [ ] Example - Ammonium H + H N H H NH4+ Valance Electrons Total - 8
H – 4x1
+ – -1
Total - 8 H + [ ] H N H H

8. Resonance Structures
These occur when there is more than one possible Lewis Structure
These models will have the same chemical formula but different placing of bonding and lone pairs

9. Practice Problems
Page 189
Questions 1-5