Chemistry 11 Honours
Bonding Extension
Writing Lewis Structures
In Chem 11, you are told how the atoms are attached to each other. This is not the
case in AP Chemistry.
There are some general rules for skeletal structures:
1) Hydrogen atoms are terminal atoms
Example:
2) Central atoms are generally less electronegative than the terminal atoms
(hydrogen is an exception to this rule)
Example:
3) Acidic Hydrogen atoms are generally bonded to oxygen atoms
Example:
4) Molecules generally have compact, symmetrical structures
Example:
Allocating electrons:
1) Determine the total number of valence electrons
2) Write the skeletal structure, put two electrons in each bond
3) Give each terminal atoms an octet of electrons
4) Assign any remaining electrons to the central atom(s)
5) Form multiple bonds to give the central atom(s) an octet of electrons if
necessary
Examples:
a) Write a Lewis Structure for COCl 2
b) Write a Lewis Structure for ClO3

Resonance
When we draw the Lewis Structure for O3 we get the following structure:
But experimental data shows us that both bonds are identical. So instead we write
two Resonance Structures. The actual structure is a hybrid (a blend) of the two
structures.
Molecules That Do Not Follow The Octet Rule
1) Molecules with odd numbers of electrons
NO
NO2
ClO2
These types of molecules are generally not very stable.
2) Molecules with incomplete octets
Molecules where the central atom is Be, B or Al tend to have too few electrons to
give every atom a full valence shell.
BF3
3) Molecules with expanded valence shells
The central atoms of PCl5 and SF6 cannot obey the octet rule because it allows for
only four bonds between the central atom and the terminal atoms.
We use Expanded Valence Shells as shown below:
PCl5
SF6
Write a Lewis Structure for BrF5 :
Formal Charge
Number of Valence electrons in the uncombined atom
Minus
Number of lone pair electrons on the bound atom
Minus
Half the number of electron in bonds to the atom
Examples
(1)
(2)
Generally, the most plausible Lewis Structure is one with formal charges of zero on
all the atoms.
When non-zero charges are required, they should be as small as possible.
The total charges on an atom must be zero for a neutral molecule and equal the net
charge for an ion.
Molecular Geometry (Shape Of A Molecule)
Linear
Examples:
O2
CO2
Angular (V-Shaped)
Example:
H 2O
We can predict the shape of molecules using Valence Shell Electron Pair Repulsion
Theory (VSEPR Theory).
This theory is based on the idea that pairs of valence electrons in bonded atoms
repel each other.
Electron Groups – A collection of valence electrons, localized in a region around a
central atom that repels other groups of valence electrons.
An electron group can be:
-
a single unpaired electron
a lone pair of electrons
one bonding pair of electrons in a single covalent bond
two bonding pairs of electrons in a double covalent bond
three bonding pair of electrons in a triple covalent bond
Most commonly, there are two, three, four, five or six electron groups around a
central atom. This leads to different electron group geometries:
-
Two: Linear
Three: Trigonal Planar
Four: Tetrahedral
Five Trigonal Bipyramidal
Six: Octahedral
VSEPR Notation
Central Atom: A
Terminal Atoms attached to central atom: X
Lone Pairs of Electrons attached to central atom: E
-
So water is: AX 2 E 2 (A central atom with two terminal atoms and two lone
pairs around it)
Note:
Electron Group Geometry: Describes how the groups of valence electrons are
arranged around a central atom.
Molecular Geometry: Describes how the bonded atoms are arranged around a
central atom.
These two geometries are not necessarily the same.
Structures with No Lone Pairs
AX 2
Example: CO2
-
both the electron group geometry and the molecular geometry are linear
AX 3
Example: BF3 -
-
both the electron group geometry and the molecular geometry are trigonal
planar
AX 4
Example: CH 4
-
both the electron group geometry and molecular geometry are tetrahedral
AX 5
Example: PCl5
-
both the electron group geometry and molecular geometry are trigonal
bipyramidal
AX 6
Example: SF6
-
both the electron group geometry and molecular geometry are octahedral
Applying the VSEPR Method
1)
2)
3)
4)
Draw the Lewis Structure
Determine the Number of Electron Groups around the Central Atom
Identify the Electron-Group Geometry (Include Lone Pairs)
Identify the Molecular Geometry (Do not include Lone Pairs)
Structures with Lone Pair Electrons
AX 2 E
Example: SO2
-
electron group geometry is trigonal planar
molecular geometry is angular
AX 3 E
Example: NH 3
-
electron group geometry is tetrahedral
molecular geometry is trigonal pyramidal
AX 2 E 2
Example: H 2 O
-
electron group geometry is tetrahedral
molecular geometry is angular