Drawing Lewis Structures
and Predicting Molecular Geometry

Introduction

Lewis structures can be draw for all main group (representative) elements and compounds containing these elements connected through covalent bonds. A covalent bond is the sharing of electrons between two atoms. If the two bond atoms are identical or have similar electronegativities, the electrons are shared equally between the two atoms and the bond is considered to be non-polar. If the two atoms are differ significantly in their electronegativities, the electrons are shared unequally between the two atoms and the bond is considered to be polar. In a polar bond, the more electronegative element gains a partial negative charge, and the less electronegative element has a partial positive charge. Partial charges are typically shown using the small Greek letter "d".

A crude predictive tool that can be used for the lighter, main group elements is as follows. If atoms share an edge on the periodic table (for example, C & N), the difference in electronegativities is usually small enough that a bond between these elements is likely to be non-polar. If two atoms do not share an edge (for example, C & O), then any bond between these atoms is likely to be polar. Partial charges in this bond can be represented as d+C–Od-. When applying this approximation, hydrogen (H, electronegativity = 2.1) should be placed between B and C.

Pauling Electronegativities of Selected Elements

B
2.0		 C
2.5		 N
3.0		 O
3.5		 F
4.0
Al
1.5		 Si
1.8		 P
2.1		 S
2.5		 Cl
3.0

Definitions

Covalent Bond:
Sharing of two or more electrons between two atoms. If 2 e– are shared, it is a single bond. If 4 e– shared, it is a double bond, and if 6 e– are shared, it is a triple bond.
Charge Cloud:
Region of space occupied by two or more e– (Lone pairs or bonds). Note that single, double, and triple bonds all occupy only one region of space and therefore only count as one charge cloud.
Octet Rule:
In covalent compounds, all atoms (except H) prefer to have access to eight valence shell electrons. The common exceptions to this rule are:
B – typically only has 6 e–
P – can have 10 e– (in compounds such as PCl5)
S – can have up to 12 e– (in compounds such as SF6)
Electronegativity:
Measure of the ability of an atom to attract electrons. Greatest for F (4.0), lowest for Fr. Know the following trends: Going across row, E.N. increases. Going up column, E.N. increases. (Note that values for noble gases not defined, since these compounds don’t form covalent bonds.)

"Rules" for Drawing Lewis Structures

Using HCN as an example:

Atom # e- # Bonds desired
H 1 1
C 4 4
N 5 3
Grand Total 10 e-  
Framework C violates octet C still low Correct
H–C–N H – C – ··
N
··		 : H – C = ··
N
 		 : H – C º N :

VSEPR Model for Prediction of Molecular Geometry

Draw Lewis structure for molecule. Count the number of "charge clouds" immediately surrounding the atom of interest (central atom). Use following table to determine geometry. (Note that the "# Bound Atoms" + "# Lone Pairs" = "#Charge Clouds".

# Charge Clouds Angle # Bound Atoms # Lone Pairs Molecular Geometry Hybridization
2 180° 2 0 Linear sp
3 120° 3 0 Trigonal Planar sp2
2 1 Bent
4 110° 4 0 Tetrahedral sp3
3 1 Pyramidal
2 2 Bent