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Chapter 10

Molecular Geometries and Bonding
Theories
– Lewis structures do not indicate the
molecular architecture – the shape of the
molecule.
– The shape and structure of a molecule
determines much of its physical and
chemical characteristics.
VSEPR Theory

Valence-shell
Electron Pair
Repulsion
– Electron pairs
(domains or
regions) repel
each other
completely.
– Balloon model.
Electron Regions

1.
2.
3.
The number of electron regions
around the central atom are counted
as:
Each single bond counts as a region.
Each lone pair counts as a region.
A multiple bond counts as a single
region.
Electron Regions

How many?
Electron Pair Geometry (EPG)



Can be from two to
six regions.
Thus, only five EPG’s
are possible.
Two regions produces
a linear EPG.
Electron Pair Geometry
Three regions
produces a
trigonal planar
geometry.
 Planar = 2D.
 Ex) BF3

Electron Pair Geometry
•Four regions becomes a three dimensional structure based
on the tetrahedron.
•Formally called tetrahedral with bond angles of 109.5o
Electron Pair Geometry
Tetrahedral is very common and
symmetrical.
 An example is CF4

Electron Pair Geometry

Five regions produces a trigonal
bipyramidal geometry with two sets of
bond angles.
Electron Pair Geometry

An example is PCl5
Electron Pair Geometry

Six regions produces an octahedral
geometry.
Electron Pair Geometry

An example is SF6
Molecular Geometry (MG)
This is based on the shape of the
electron pairs.
 When a molecule has no lone pairs, the
EPG = MG.
 If the molecules DOES have one or
more lone pairs, then the shape of the
atoms is determined based off of the
EPG.

Molecular Geometry
Molecular Geometry
Molecular Geometry
Examples

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Bent (120), SO2
Trigonal pyramidal, NH3
Bent (109.5), H2O
Seesaw, SF4
T-shaped, ClF3
Linear, I3Square pyramidal, BrF5
Square planar, XeF4
Sketching the Molecules


Simple = Ball and Stick figures
Representing the 3D shapes:
– Put as many of the molecules in the same plane
as possible including the central atom. Use
straight lines for bonds connected to atoms in
plane.
– For atoms in front of the plane, use a solid
wedge.
– For atoms behind the plane, use a hashed
wedge.
3D Representations
Lone Pairs
A non-bonding
pair will always
take up more
space.
 This compresses
the normal bond
angles.

Lone Pairs
Lone Pairs

This also explains
the MG’s for the
trigonal
bipyramidal
family.
Shapes of Larger Molecules

A molecule
like acetic
acid has
three central
atoms.
Shapes of Larger Molecules
Polarity
A molecule can
contain very polar
bonds, but can be
non-polar.
 An example is
CO2.

Polarity

On the other hand,
sometimes polar
bonds DO make a
molecule polar.
 An example of a
polar molecule is
H2O.
Polarity
Polarity
A molecule with a symmetrical
distribution of polar bonds will be nonpolar.
 A molecule with an un-symmetrical
distribution of polar bonds will be polar.

– presence of lone pairs
– different external atoms
Polarity
Polarity



Polar molecules are
attracted to other polar
molecules
Because water is a polar
molecule, other polar
molecules dissolve well in
water
– and ionic compounds as
well
Non-polar molecules do
NOT dissolve in water.
Valence Bond Theory


How can we explain the formation of the bonds in a
molecular compound?
A bond occurs when a valence orbital on one atom
overlaps with a valence orbital of another atom.
Valence Bond Theory

The H2 molecule – a closer look.
nuclear
repulsion
no interaction
minimum energy
Valence Bond Theory




Three (or more) atom molecules cannot be
explained by simple overlap of orbitals.
Fact: a bond generally forms between two half-filled
orbitals.
Fact: an s-type orbital is spherical, so it could form a
bond in any direction.
Fact: the three p-type orbitals are at 90 degree
angles to each other.
Valence Bond Theory
CH4 – has an EPG and
MG of tetrahedral with
bond angles of 109.5o.
 Valence diagram for C
and H before any bonding
is:

Valence Bond Theory
Solution: promote the paired electron
from the s orbital to the empty p orbital.
 Solution: mix the one s and three p
orbitals together to get a new set of four
orbitals all equal in energy. This is
called _____________________.

Valence Bond Theory

Each hybrid orbital has some s and some p
characteristics.
 Thus, they look different!
Types of Hybrids

Determined from the EPG.
EPG
Atomic
orbitals
Linear
s+p = sp
Trigonal
s+p+p =
planar
sp2
Tetrahedral s+p+p+p =
sp3
Hybrid
diagram
Examples
BeF2
BF3
CH4
Types of Hybrids
Atoms in the third period and beyond
have empty d orbitals that can
potentially be used for hybridization.
 PCl5 – requires five bonds, so need a
set of five orbitals.
 Once again, must first promote the s
electron to an empty d orbital.

Types of Hybrids
EPG
Atomic
orbitals
Hybrid
diagram
Examples
Trigonal
bipyramidal
s+p+p+p+d
= sp3d
PCl5
Octahedral
s+p+p+p+d+d =
sp3d2
SF6
Molecules with Lone Pairs

Ex) NH3

Ex) H2O

Ex) BrF3
Multiple Bonds
Two types of bonds are possible.
 1. Sigma (s) bonds have a cylindrical
shape of electron density along the
central axis between the two nuclei.

s bond
Multiple Bonds

2. Pi (p) bonds have
an electron density
above and below the
central axis.
– Are formed by the
overlap of two parallel
half-filled p-type
orbitals.
Multiple Bonds
The majority of bonds are sigma bonds.
 When a double bond is present, the first
bond is a sigma and the second is a pi.

Pi bonds
Multiple Bonds

For any pi bonds, you MUST use an unhybridized half-full p-type orbital.

Ex) C2H4

Ex) CO2
Multiple Bonds
Pi Bond Significance
Sigma bonds have free-rotation about
the central axis.
 Ex) C2H4Cl2
 Pi bonds have NO free-rotation due to
the fact that they must overlap above
and below the central axis.
 Ex) C2H2Cl2

Pi Bond Significance
Isomers
When two compounds share the exact
same formula but are different either
structurally or spatially, then they are
said to be isomers.
 Structural isomers

– C5H12
– C2H6O
Isomers
Geometric isomers are different
spatially.
 This can occur for our carbon-carbon
double bond.
X
Y
Y
Y
C=C
C=C
Y
X
X
X

Trans
Cis
Isomers
The last molecule in your packet has
three possible structures. One is
structural and two are geometric
isomers.
 One other geometry can have cis/trans
isomerism – is it tetrahedral or square
planar?
 Ex) CH2Cl2 or Pt(NH3)2Cl2?

Limitations of V.B. Theory
Valence Bond Theory does not
adequately explain molecules with
resonance structures nor some other
observed properties.
 Ex) O2 or molecular oxygen is
paramagnetic (unpaired electrons).

Lewis structure for O2
Molecular Orbital Theory
A more sophisticated and complex
model of bonding.
 Atomic orbitals from each atom
contribute to new MO’s.
 Like atomic orbitals, each MO can hold
up to two electrons.
 A MO, though, is spread out over the
entire molecule.

MO Theory
For each atomic orbital contributed we
get one MO.
 Half of the MO’s become bonding and
the other half become anti-bonding.

– Waveforms add either constructively or
destructively like light!

For the n=1 period, each atom
contributes a 1s atomic orbital.
MO Theory
MO Theory
MO Theory
The Bond Order in MO theory is found
by: BO = ½ (Bonding e- - Anti-bonding e-).
 Any bond order = 0 implies that the
molecule is not possible.
 Odd number of electrons will produce
half-integer BO’s.

MO Theory
Period 2 elements have both the 2s and
2p atomic orbitals to contribute towards
MO’s.
 Thus, two atoms from period 2 will have
how many atomic orbitals total?
 How many MO bonding orbitals will be
produced? Anti-bonding?

MO Theory
The two 2s orbitals overlap just like the
two 1s orbitals did in period 1.
 This produces the s2s and s2s* MO’s.
 The six 2p orbitals overlap differently.

– Two will overlap end-on-end and produce a
s2p type MO.
– Four will overlap sideways and produce
two p MO’s that are equal in energy.
MO Theory

The energy
level
diagram
produced
for all of
these new
MO’s is:
MO Theory

Diagram assumes that no 2s-2p orbitals
interactions occur.
 For B2 , C2 , and N2 the interactions cause the
s2p and p2p order to trade places on the
diagram.
 Since these are filled for O2 , F2 , and Ne2 the
diagram can be written with those two always
reversed to simplify.
MO Theory
MO Theory


It is possible for some molecules and ions to be
paramagnetic – one or more unpaired electrons.
Most will be diamagnetic – all paired electrons.
n=2 Diatomic Molecules
n=2 Diatomic Molecules
Molecular Oxygen



According to MO theory, it has a BO = 2 and it is
paramagnetic!
As liquid O2 is poured between the poles of a
magnetic, it will have a strong attraction.
Clip.
Heteronuclear Diatomics

For period 2, we can mix and match other
elements and apply the same MO diagram.
 Simply add up the total valence electrons that
each contribute and place in the diagram.
 For ions, a positive charge means we would
decrease by one electron and a negative
charge means we would increase by one
electron.
Heteronuclear Diatomics



More electronegative
element has lower
energy orbitals.
Can produce bond
orders with ½ values.
B.O. = ___________
Polyatomic Molecules



When many atoms are combined together, the
atomic orbitals of all the atoms are combined to make
a set of molecular orbitals, which are delocalized over
the entire molecule
Gives results that better match real molecule
properties than either Lewis or valence bond theories
This is why resonance structures cannot be
explained by valence bond theory.
Ozone, O3

The structure of O3 includes two
resonances.
M.O. showing delocalized pi bonds