Name:
Lab Partner:
Date:
Molecular Shape
OBJECTIVES

Construct, draw and name molecules.
 Determine the polarity of a molecule.
 Apply VSEPR to the molecular shape of the molecules.
INTRODUCTION
Each atom in a molecule attempts to obtain a more stable electron configuration by sharing electrons with
neighboring atoms. The choice to share electrons or to donate or accept electrons depends on the electronegativity
differences of the bonding atoms. Differences that are less than 1.7 will result in a bond that is covalent while bonds
with differences greater than 1.7 are ionic in nature.
Atoms that share electrons, in most cases, will have a complete octet of electrons. The arrangement of these atoms is
based on the number of atoms in the molecule and the number of unpaired electrons that surrounds the central atom.
In order for the molecules to form, bonding orbitals form between the atoms. The central atom is usually the atom that
needs the most electrons to complete their octet.
For example the electron configuration of the carbon atom is [He] 2s2 2p2. In order for it to form a bond and make
CH4 it promotes one of the valence electrons in the s orbital to the vacant p orbital. When this occurs a bonding
orbital is formed. This bonding orbital is a hybrid of the s and the 3 p orbitals. We state the molecule has sp3
hybridization around the central atom.
Unpaired electrons around the central atom in the molecule repel the electrons that are being shared between the
molecule. This causes a distortion of the molecular shape and results in a polar molecule, where one region of the
molecule has a negative and/or positive charge. Polar molecules also result when there is an unequal pull on the
central atom by the terminal atoms. When the electronegativity differences range between .4 to 1.7 a bond is polar
covalent. The atom in the bond with the greater electronegativity will contain a partial negative region.
Electronegativity differences less than .4 are considered non-polar covalent.
MATERIALS
Molecular Model Kits
H – white
Cl – green
O – red
N – blue
lone electron pair – beige sphere
F – light green
C – black
P – purple
S – yellow
covalent bond – grey pear
Be – grey
B – beige “shiny”
Directions: Using the formulas construct, draw a 3D model, determine the name, draw Lewis dot structures, structural
formula, calculate electronegativity difference, identify direction of electron pull.
Formula and
shape
HCl
Drawing:
BeCl2
Drawing:
H2O
Drawing:
Name of the
molecule
Terminal and
Central Atoms’
Electronegativity
Difference
Direction of
Electron Pull
(Diagram)
Lone Electron Pairs
on the Central Atom
(yes/no)
Polar or Non-Polar
Formula and
shape
BH3
Drawing:
NH3
Drawing:
CH4
Drawing:
Name of the
molecule
Terminal and
Central Atoms’
Electronegativity
Difference
Direction of
Electron Pull
(Diagram)
Lone Electron Pairs
on the Central Atom
(yes/no)
Polar or Non-Polar
Formula and
shape
Name of the
molecule
Terminal and
Central Atoms’
Electronegativity
Difference
Direction of
Electron Pull
(Diagram)
Lone Electron Pairs
on the Central Atom
(yes/no)
Polar or Non-Polar
PCl5
Drawing:
SF6
Drawing:
ANALYSIS AND CONCLUSIONS
1. The bond angle for BH3 is 120°, CH4 is 109°, NH3 107° and H2O 105°. What influences the bond angles of
these molecules?
2. Why are their exceptions to the octet rule? Provide examples.
3. Why can PCl5 and SF6 exist but NCl5 and OF4 cannot?
4. Why can a non-polar molecule contain polar bonds between the central and terminal atoms? Provide
examples.
5. Which molecules will always be polar? Explain, provide examples and state the theory that supports your
answer.
6. What is the VSERP theory?