Read Sections 4.7, 4.11 and 4.12 before viewing the slide
show.
Unit 16
Shapes and Electrical Properties of Molecular
Compounds
•VSEPR Theory (4.11)
•Electronegativity (4.7)
•Identification of Compounds as Polar or Nonpolar
(4.12)
VSEPR Theory (4.11)
•The geometry of molecules is important in
determining some of their properties
•Electrons tend to repel each other because of their
negative charges
•In considering a Lewis structure for a molecule (or
ion) the geometry becomes apparent when one
considers that the electrons tend to get as far away
from each other as possible
•These observations are summarized in the Valence
Shell Electron Pair Repulsion theory (VSEPR).
Simply stated, electron pairs around a central atom
arrange themselves to minimize repulsion.
Molecular Shape from VSEPR (4.11)
Number of
Bonded
Atoms
Number of
Nonbonded
Pairs of
Electrons
Number of Sets of Electrons
(sum of bonded atoms and
nonbonded pairs)
Molecular Shape
Examples
Picture of Shape
2
3
0
0
2
3
Linear
Trigonal Planar
HCN, CO2
SO3, BF2Cl
O=C=O
2
1
3
Bent
SO2
4
0
4
Tetrahedral
CH4, CHCl3, CH2Cl2
3
1
4
Pyramidal
NH3
2
2
4
Bent
H2O
Application of the Previous Table (4.11)
•Application of the table on the previous slide involves drawing
the Lewis structure and determining the number of bonded
atoms and nonbonded pairs of electrons
•Notice the counting of atoms and nonbonded pairs only
involves those around the atom of interest, circled in the table
below.
Molecule
# of Bonded
Atoms
# of Nonbonded Molecular Shape
Pairs around
Atom of Interest
H▬C≡N|
2
0
Linear
..
H▬N ▬H
|
H
3
1
Pyramidal
More Examples (4.11)
Molecule
# of Bonded
Atoms
# of Nonbonded Molecular Shape
Pairs around
Atom of Interest
OC O
2
0
Linear
H  O H
2
2
Bent
4
0
Tetrahedral
: Br :
|
: Br  C  Br :
|
: Br :
Electronegativity (4.7)
•Recall that in ionic compounds electrons are essentially
transferred from a metal to a nonmetal
•Electrons are not necessarily shared equally between
two atoms in a covalent bond
•The determining factor as to which atom will attract
electrons more in a covalent bond is called the
electronegativity
•A scale has been set up for electronegativity with
fluorine having the highest electronegativity – the
strongest ability to draw electrons towards itself in a
chemical bond
•A section of the electronegativity table is on the next
slide
Electronegativity and the Periodic Table (4.7)
•The chart below gives the electronegativity values of some of the key
elements.
•A high value (F, for example) means a strong tendency to draw electrons
toward that element in a chemical bond
•Hydrogen is similar to carbon in terms of electronegativity
•The differences in electronegativity between metals and nonmetals are
typically large enough to result in a significant transfer of electrons to form
ionic compounds
1
2
13
14
15
16
17
18
•The nonmetal-nonmetal bonds H
Largest
are covalent – they involve
2.1
electron sharing, but not equally Li
Be
B
C
N
O
F
1.0
Na
0.9
K
0.8
1.5
Mg
1.2
Ca
1.0
2.0
Al
1.5
2.5
Si
1.8
3.0
P
2.1
As
2.0
3.5
S
2.5
Se
2.4
4.0
Cl
3.0
Br
2.8
I
2.5
Polar and Nonpolar Covalent Bonds (4.7)
•If the difference in electronegativity between the two atoms is less than 0.5, the
electrons are about equally shared and the bond is called a nonpolar covalent bond
•If the difference in electronegativity between the two atoms is between 0.5 and 2.0,
the electrons are drawn toward the more electronegative element and the bond is
called a polar covalent bond. Basically the negative and positive charges are
separated from each other.
•If the difference in electronegativity between the two atoms is greater than 2.0, the
bond is considered to be ionic because of a significant transfer of electrons.
Bond
Diff. in EN
Bond Type
N-O
3.5 - 3.0 = 0.5
polar
covalent
Al-Cl
3.0 – 1.5 = 1.5
polar
covalent
C-H
2.5 - 2.1 = 0.4
nonpolar
covalent
K-O
3.5 - 0.8 = 2.7
ionic
1
2
13
14
15
16
17
H
2.1
Li
1.0
Na
0.9
K
0.8
Be
1.5
Mg
1.2
Ca
1.0
B
2.0
Al
1.5
C
2.5
Si
1.8
N
3.0
P
2.1
As
2.0
O
3.5
S
2.5
Se
2.4
F
4.0
Cl
3.0
Br
2.8
I
2.5
18
Polar and Nonpolar Covalent Bonds (4.7)
•There are a couple of ways of representing in a polar covalent bond which
way the electrons are pulled
•Use the greek letter δ with a superscript to represent whether it is at the positive or
negative end of the bond
•Use an arrow such as
that points toward the more electronegative atom. The
length of the arrow may be used to indicate on a relative basis how big the difference in
electronegativity is. The arrow has a size and direction which together can be referred
to as the dipole moment.
Bond
Diff. in EN
N-O
3.5 – 3.0 = 0.5
Al-Cl
P-F
3.0 – 1.5 = 1.5
4.0 – 2.1 = 1.9
Greek Rep.

 
NO

 
Al  Cl

 
P F
Arrow Rep.
NO
Al  Cl
PF
Polar and Nonpolar Molecules (4.12)
•A molecule itself may be polar or nonpolar depending upon the geometric arrangement of its
bonds.
•A polar molecule has separate centers of positive and negative charges, similar to north and south poles of a magnet.
•A nonpolar molecule has the centers of positive and negative charges at the same place, similar to a magnet in which
the poles have collapsed on top of each other.
•The polarity of the bonds and the geometry of the molecule must both be considered to
determine whether or not a molecule is polar or nonpolar.
CH4
NH3
•The red represents areas of positive charge, the blue the negative charge
•In CH4, the red positive charge is evenly distributed throughout the molecule indicating a
nonpolar molecule.
•In NH3, the negative and positive charges are separated from each other indicating a polar
molecule.
Methods of Telling Polarity (4.12)
•A quick way to determine something about polarity of a molecule is to:
1. Draw the Lewis structure
2. If the central atom has any nonbonded pairs around it, it will be a polar molecule.
3. If the central atom has only atoms around it and they are all the same, it is a
nonpolar molecule.
4. If the central atom has only atoms around it and any of them are different, it is a
polar molecule.
• On the previous slide, CH4 has four atoms all the same around it – it is
nonpolar.
• NH3 has a nonbonded pair of electrons and will be polar
• These guidelines are general. Sometimes electronegativity differences
will be small enough that a molecule, even with different atoms around
the central atom, will be virtually nonpolar. An example of this is
compounds containing only carbon and hydrogen – these are
considered nonpolar regardless of geometry.
• There are also molecules with 5, 6, or 7 atoms and/or nonbonded
electrons bonded to the central atom. Those don’t strictly follow the
rules above. Those we won’t worry about in this course.
A Simulation
•If you would like to get some practice on molecule polarity using the PHET
simulations, go to:
•http://phet.colorado.edu
•Click on the chemistry icons in the upper left.
•Choose the Molecule Polarity simulation
•You can Save It and then run it or Run it Now
•A reasonable way of using this simulation is to focus simply on changing
the electronegativity of the atoms and watching what happens to the arrow.
The arrow represents what we call the dipole moment and points toward
the negative charge in the molecule.
•Take a look at the 2-atom, 3-atom, and molecule tabs to get a feel for how
the polarity idea works.
Importance of Polarity (4.12)
•Properties of compounds depend to a large extent on the
polarity or nonpolarity of the compound. Properties affected
include:
•boiling point
•melting point
•vapor pressure
•solubility
•Some of these properties and their relationship to polarity will
be explored later.