SCH4U
Bond Dipoles and Molecular Dipoles
Bond Dipoles
Recall that molecular compounds contain covalent bonds. Covalent bonds can be
classified as covalent or polar covalent based on the electronegativity (EN) of the
bonded atoms. Electronegativity is a relative measure of an atoms electron
attracting ability in a bond. By calculating the electronegativity difference (∆EN(,
bonds can be classified.
ΔEN = EN2 – EN1
where:
EN2 is the electronegativity of element with higher EN
EN1 is the electronegativity of element with lower EN
Although every bond is distinct, the following guidelines are used to classify
bonding:
COVALENT
0.0
e.g. N2
POLAR
COVALENT
0.5
IONIC
1.7
NH3
3.3
K3N
Recall that a polar covalent bond exists if the two bonding electrons are not equally
shared between the two atoms. A bond dipole exists if the two atoms have different
electronegativities (i.e. ∆EN > 0). Since the electrons are attracted towards one
atom, that atom has a partial negative charge (-) and the other atom has a partial
positive charge (+).
e.g. Propanol (isopropyl alcohol or rubbing alcohol) has the following structure:
H
H O H
H C C C H
H H H
Identify the types of bonds in each compound or group of compounds and rank them
according to increasing polarity. Indicate the bond dipole in each bond.
Electronegativity and Bond Dipole Questions:
1. a) Examine the periodic table and describe or sketch the general trend in
electronegativity for atoms.
b) What two aspects of atomic structure would explain this trend?
2. Draw the following bonds and label the electronegativity values of each atom.
Label the charge (+ or -) on each atom in the bond (if they exist) and classify the
bond as covalent, polar covalent or ionic.
a) H-Cl
b) C-H
c) N-O
d) I-Br
e) Mg-S
f) P-H
g) C-C
h) Al-H
3. Identify the types of bonds in each compound or group of compounds and rank
them according to increasing polarity. (State symbols indicate the state each
compound is in at room temperature.)
a) H2O (l),
H2 (g), CH4 (g), HF (g), NH3 (g) LiH (s), BeH2 (s)
b) PCl3 (l), LiI (s), I2 (s), RbF(s), AlCl3 (s)
c) CH3OH (l)
d) CHFCl2 (g)
Predicting Molecular Polarity
Bond polarity refers to the unequal distribution of charge in a bond. Similarly,
molecular polarity refers to the asymmetrical distribution of charge in a molecule.
Both bond polarity and molecular shape must be considered to determine if a
molecule is polar. Overall, molecular polarity can be determined with the vector
addition of all bond dipoles in a molecule.
1) Diatomic Molecules
With a diatomic molecule, there is only one bond so a bond dipole always results in
a molecular dipole.
e.g.
H – Cl
e.g. F
–F
2) Polyatomic Molecules
With larger molecules, there are two or more bonds to consider and thus one has to
consider the number and orientation of several bond dipoles. With these molecules,
a molecular dipole only exists if the bond dipoles do not cancel each other out.
e.g.
H 2S
e.g.
MgI2
SCH4U1
Molecular Shapes and Dipoles Problem Set
Part 1: VSEPR Shapes and Bond Polarity
1. i) Draw a Lewis structure for each compound.
ii) Draw and name the molecular shape of the following substances:
a) CH4
e) NF3
i) BrF5
m) BCl3
b) Cl2O
f) SiF4
j) LiH
n) MgI2
c) BeCl2
g) SeF6
k) H2S
o) NH4+
d) LiCl
h) PF5
l) PH3
p) H3O+
Part 2: Molecular Dipoles
1. For each of the following molecules:
i) Draw the Lewis structure and the molecular shape.
ii) Indicate the direction of any bond dipoles with vectors.
iii) Add the dipoles and determine if a molecular dipole exists.
a) LiF
b) CH4
c) HCl
d) Cl2
e) NH3
f) BCl3
g) BeH2
2. Which of the following molecules has a molecular dipole?
a) NaF (g)
b) MgO (g)
c) MgI2 (g)
h) BF2H
i) SFCl5
d) GaCl3 (g)