Section 6.5 – Molecular Geometry
 The properties of
molecules depend
on the bonding
and the molecular
geometry, the 3dimensional
arrangement of the
atoms in space.
Molecular Polarity
 This is the uneven
distribution of
molecular charge,
and it is
determined by the
polarity of each
bond, along with
the geometry of the
molecule.
Two Theories – based on evidence
 VSEPR Theory:
Accounts for
molecular bond
angles.
 Hybridization:
Describes the orbitals
that contain the
valence electrons of a
molecule’s atoms.
VSEPR Theory
Valence
Shell
Electron
Pair
Repulsion
Repulsion between the
valence-shell electrons
surrounding an atom
causes these sets to be
oriented as far apart as
possible.
Activity
CH4
NH3
H2O
• Draw the Lewis structures for all 3 molecules
• Sets of electrons that remain together in
bonds or in lone pairs are referred to as
electron domains. Electron domains prefer to
be as far apart as possible from each other
within a molecule.
• How many electron domains are located
around the central atom of each molecule?
Activity
CH4
NH3
H2O
• Collect supplies – 1 set of gumdrops and
toothpicks for each pair of students
• Build a 3-D model of each molecule – SHOW
ME YOUR MODEL!!!!
• Remember – electron domains want to be as
far apart as possible.
• How do lone pairs affect the shape of each
molecule?
Diatomic Molecules
 Diatomic molecules are
composed of two atoms, so
the geometry is always linear
 the molecular polarity is
determined by the
electronegativity differences
between the atoms.
H2: non-polar HCl: polar
Shorthand for Describing
 For molecules containing more
than two atoms, we can use the
following symbols with
subscripts:
A – the central atom
B – number of bonds
on the central atom
E – number of lone
pairs on central atom
(for atoms that have double or
triple bonds, it is treated as a
single B for geometry)
The Basis For VSEPR Theory
 That one must consider the
locations of all electron
pairs of the valence
electrons in the molecule.
 Polyatomic ions are treated
the same way.
The Basis For VSEPR Theory
 The following examples do
not have lone pairs that
influence the geometry of
the molecule.
Linear – AB2
 Central atom with two
single bonds, no lone
pairs. Because the
valence electron pairs
in the bonds repel each
other, the bonds are as
far apart as possible
(180°).
 Ex: BeH2
Trigonal Planar – AB3
The 3 A-B bonds
stay furthest apart
by pointing to the
corners of an
equilateral triangle,
giving 120° angles
between the bonds.
Ex.: BH3
Tetrahedral – AB4
Octet rule is followed
here. The distance
between the A-B
bonds is maximized if
each bond points to
the corners of a
tetrahedron, giving
bond angles of 109.5°
between the bonds.
Ex.: CH4
Trigonal-bipyramidal – AB5
 120° angles
Ex.: PCl5
between
bonds within
the trigonal
plane, 90°
bond angles
between the
axial bond
and those in
the plane.
Octahedral – AB6
6 bonds to
the central
atom, all
equidistant
from each
other. 90°
bond
angles.
Ex.: SF6
Lone Pairs Do Occupy Space and Influence
Geometry
 But our description of the
molecular geometry refers
to the positions of the atoms
only.
 A summary of the shapes of
various molecules is in
Table 6-5, p. 186.
 Different sizes of B groups
may distort some bond
angles that are given in the
table.
VSEPR and Unshared Electron Pairs
 One must always write
out the Lewis structure
for a molecule to decide
on the proper geometry,
the chemical formula of
something does not tell
you about lone pairs
around the central atom.
Bent – AB2E
 2 bonds to central atom
with one lone pair. The
lone pair bends what
one would expect to be
linear. The lone pair
takes up more space
than a bond and shoves
the bonded atoms
closer together than the
120° for trigonal planar.
Bent – AB2E2
The addition of a
second lone pair
forces the
bonding atoms
even closer
together than
what one expects
from tetrahedral.
Trigonal Pyramidal – AB3E
 With 3 bonds one
expects trigonal
planar with 120°
between the bonds,
but the lone pair
bends the plane
away from the pair,
forcing the atoms
closer together.
Molecular Polarity
 Reflecting on the
geometries, we can
now see why lone
pairs on the central
atom make a molecule polar – it changes the
geometry of the molecule and creates an
uneven “tug of war” across the molecule.
Intermolecular Forces
 These are the forces of
attraction that occur
between molecules. They
vary in strength but are
generally weaker than
regular bonding (ionic,
covalent, or metallic).
Melting and Boiling Points
 Usually are a good measure of the force of attraction. The
higher the boiling point, the stronger the forces between
particles.
Dipole Force
The strongest intermolecular
forces exist between polar
molecules. Each one acts
as a dipole, created by
equal but opposite charges
that are separated by a
short distance.
The direction of the arrow
is pointed to the negative pole,
the crossed tail indicates the + side
Dipole-dipole forces
 These are the forces of
attraction between polar
molecules.
 Example:
bp for F2 is -188°C
bp for HF is 20°C
bp for HCl is -85°C
Which is the stronger
dipole-dipole force?
For Molecules With More Than 2 Atoms
 The molecular polarity
depends on both the
polarity and the
orientation of each bond.
Induced Dipoles
 The electrons of a nonpolar
molecule can be temporarily
attracted by a polar molecule.
This is weaker than a regular
dipole-dipole force.
This can be very important in the
solubility of gases in water.
Hydrogen Bonding
 A very special type of
dipole-dipole force in
which a hydrogen atom
is bonded to a highly
electronegative atom is
attracted to an unshared
pair of electrons of an
electronegative atom of
another molecule.
Usually Represented By
Dotted Lines
Explains High bp’s of Some Compounds
 Gives the H atom a large
positive charge, and it’s
small size allows it to
come very close to the
unshared pair of electrons
on an adjacent molecule.
Extremely Important in Biochemistry
Stereoisomers
 Isomers – same chemical formula, different structures
(isopropyl alcohol vs 2-propanol)
 Stereoisomers – structures are mirror images.
WTF?????
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London Dispersion Forces
 This is a weak attractive
force resulting from the
imbalance of electrons and
the creation of an
instantaneous dipole.
Important for noble gases
and nonpolar molecules.
Increased Force
 With the increased number
of electrons in the
interacting atoms or
molecules, thus with
increasing atomic or molar
mass.
Assignment – Due Wed. EOP
Section 6.5 Worksheet
Mixed Review Worksheet
Molecular Geometry
Worksheet
6.5 Textbook Problems