Chemistry 20
Molecular Shape and Polarity (13)

The 3D shape of an atom influences its physical and chemical characteristics.
Review Information
1) Single chemical bond is the sharing of two electrons. (The overlapping of two half-filled orbitals)
2) Filled electron shells are not involved in bonding (unless one electron is transferred to an empty orbital by
hybridization – like in boron)
3) Orbitals will arrange themselves as far away from each other as possible, because negative charges repel
each other
a. This ideas is known as VSEPR – valence shell electron pair repulsion theory
Molecular Shape
Elements of the 2nd period
 All the groups/families have the same valence shell structure: they will bond (and be shaped) similarly
Group 1/Group 2/Group 17: Lithium/Beryllium/Chlorine
 These elements have one free electron
 If there are two atoms involved in the molecule, the shape will be LINEAR
o Ex:
HCl, LiF
 If there are more than two 2 atoms in the molecule, the shape is governed by the bonding orbitals of the
central atom.
o Ex:
H-Be-H
-Here the shape is linear, but determined by the Be atom
Group 13: Boron Family
 In this case the boron has 3 partially filled orbitals, therefore there are 3 places to bond
o Ex: BF3
o
o
Each F atom is equally distant (120º) from each other
This shape of molecules is called TRIGONAL PLANAR
Group 14: Carbon Family
 In this case the carbon has 4 partially filled orbitals, therefore there are 4 places to bond
o Ex: CH4
o
o
Each H atom is equally spaced at 109.5º from each other
This shape is TETRAHEDRAL
Group 15: Nitrogen Family
 In this case the nitrogen has 3 partially filled orbitals and one full orbital, therefore there are 3 places to
bond.
o Ex: NH3
o
o
The lone pair of electrons occupies slightly more space than the bonds do. This means the hydrogen
bonds are 107º away from each other.
This shape is TRIGONAL PYRAMIDAL
Group 16: Oxygen Family
 In this case the oxygen has 2 partially filled orbitals and 2 full orbitals, therefore there are 2 places to bond.
o Ex: H2O
o
o
The two pairs of electrons occupy more space than the bonds do. This means the angle between
hydrogen atoms is 104.5º.
This is a BENT molecule
Molecular Polarity
 The molecular polarity describes whether or not a molecule has dipoles.
 Molecules that have a dipole-dipole force are much more stable. (There is more energy needed to overcome
the intermolecular forces between the molecules.
Rules for Predicting Polarity
1) In a diatomic molecule, if the bond is nonpolar, the molecule is nonpolar. (Diatomic molecules have only one
bond)
2) Diatomic molecules that have a polar bond are polar molecules.
FOR MULTIATOMIC MOLECULES:
3) If all the bonds in the molecule are nonpolar, the molecule is nonpolar
4) If the bonds are polar, the molecule is polar or nonpolar depending upon the spatial arrangements of the
bonds.
a. If polar bonds are symmetrically arranged so that the polarity cancels, then the molecule is nonpolar
b. If polarity of bonds within a molecule add, the molecule will have a resultant polarity or dipole and
the molecule is polar

Ex: BeH2

NOTE: molecule that have different elements bonded to a central atom will have dipoles that may NOT
cancel
H-Be-H
(these bonds are polar, but they cancel each other)
Hydrogen Bonding
 Because hydrogen has only one electron, the positive proton will be attracted to the increased
electronegativity of another atom. This is called hydrogen bonding.
 This is a stronger force than a dipole-dipole bond.
Forces between Nonpolar molecules
 Nonpolar molecules have no attractive or repulsive force acting between them. BUT, when the molecules are
close together, a momentary dipole is created for a short time.
 This short-lived, weak attraction is known as a London Force.
 The weak forces of dipole/dipole and London forces are known as van der Waal forces.