CH3Notes - Chapter
Three
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Unformatted text preview: Chapter
Three
 Molecular
Shape
and
Structure
 CHEM
1211K
 Fall
2010
 1
 Objec5ves 
 •  Use
VSEPR
theory
to
understand
the
three‐dimensional
 shapes
of
molecules.
 –  Be
able
to
draw
Lewis
structures
and
idenGfy
electron
 arrangement
and
shape.
 –  Use
the
Lewis
structure
to
determine
whether
a
 molecule
is
polar.
 •  Use
Valence
Bond
theory
to
explain
how
covalent
 bonding
takes
place.
 –  Linear
combinaGon
of
atomic
orbitals
and
hybridizaGon
 •  Use
Molecular
Orbital
theory
to
explain
the
limitaGons
of
 VSEPR
and
valence
bond
theory.
 2
 Key
Ideas 
 •  Electronic
repulsions
between
electron
pairs
 determine
molecular
shapes.
 •  Chemical
bonds
can
be
discussed
in
terms
of
 two
quantum
mechanical
theories
that
 describe
the
distribuGon
of
electrons
in
 molecules.
 3
 Why
is
it
important? 
 •  The
shapes
of
molecules
determine
their
odors,
 tastes,
and
acGons
as
drugs.


 •  Molecular
shape
plays
an
essenGal
role
in
the
 reacGons
that
take
place
throughout
our
bodies
 and
are
necessary
for
life.
 •  It
also
affects
the
properGes
of
the
materials
 around
us,
including
their
physical
states
and
 solubiliGes.
 •  PercepGon,
thinking,
and
learning
depend
on
the
 shapes
of
molecules
and
how
they
change.
 4
 The
VSEPR
Model
 •  The
valence shell electron pair repulsion VSEPR)
model
uses
the
idea
of
electron
/
 electron
repulsion
to
explain
the
three
 dimensional
shape
of
molecules.
 –  Molecules
with
no
lone
pairs
on
the
central
atom
 –  Effects
of
lone
pairs
 –  Consequences
of
molecular
shape
 5
 The
Basic
VSEPR
Model
(3.1) 
 •  VSEPR
Rule
One:

Regions of high electron concentra8on repel one another. –  To minimize their repulsions, these regions move as far apart as possible while maintaining the same distance from the central atom. –  Regions
of
high
electron
concentraGon
=

 6
 The
Basic
VSEPR
Model
(3.1) 
 •  The
electron arrangement
of
the
molecule
is
 the
arrangement
of
the
“most
distant”
 locaGons
of
the
regions
of
electron
 concentraGon.
 –  Note
where
the
atoms
lie
to
idenGfy
the
shape
of
 the
molecule
and
give
it
the
corresponding
name.
 7
 The
Basic
VSEPR
Model
(3.1) 
 •  The
posiGons
that
two
to
seven
regions
of
 high
electron
concentraGon
take
around
a
 central
atom.
 Figure 3.2 8
 The
Basic
VSEPR
Model
(3.1) 
 9
 The
Basic
VSEPR
Model
(3.1) 
 •  VSEPR
Rule
Two:

There is no dis8nc8on between single and mul8ple bonds. –  A mul8ple bond is treated as a single region of high electron concentra8on.
 10
 The
Basic
VSEPR
Model
(3.1) 
 #
Regions
of
 high
electron
 concentra5on
 Name
 Bond
angles
 Example
 2
 Linear
 180o
 CO2
 3
 Trigonal
 planar
 120o
 BF3
 4
 Tetrahedral
 109.5o
 CH4
 5
 Trigonal
 bipyramidal
 90o,
120o,
 180o
 PCl5
 6
 Octahedral
 90o,

180o
 SF6
 7
 Pentagonal
 72o,
90o
180o
 bipyramidal
 IOF6‐
 11
 The
Basic
VSEPR
Model
(3.1) 
 •  What
if
there
is
more
than
one
central
atom?
 12
 The
Basic
VSEPR
Model
(3.1) 
 •  SecGon
summary:
 –  According to the VSEPR model, regions of high electron concentra8on take up posi8ons that maximize their separa8ons. –  Electron pairs in a mul8ple bond are treated as a single unit. –  The shape of the molecule is then iden8fied by the rela8ve loca8ons of its atoms. 13
 Molecules
with
Lone
Pairs
on
the
 Central
Atom(3.2) 
 •  VSEPR
Rule
Three:

All regions of high electron density, lone pairs AND bonds, are included in a descrip8on of the electronic arrangement, but only the posi8ons of atoms are considered when iden8fying the shape of molecules.
 14
 Molecules
with
Lone
Pairs
on
the
 Central
Atom(3.2) 
 •  VSEPR
Rule
Four:

The strengths of repulsions are in the order Lone pair – lone pair > lone pair – atom > atom – atom 15
 Molecules
with
Lone
Pairs
on
the
 Central
Atom(3.2) 
 •  Lone
pairs
push
the
atoms
bonded
to
the
 central
atom
closer
together
because
the
 electron
cloud
of
a
lone
pair
can
spread
 



over
a
larger
volume
than

 



a
bonding
pair.
 –  Example:

NH3
 16
 Molecules
with
Lone
Pairs
on
the
 Central
Atom(3.2) 
 Regions
of
 electron
 concentra5on
 Electronic
 arrangement
 Bonded
 pairs
 Lone
 pairs
 Shape
 Example
 2
 Linear
 2
 0
 Linear
 BeCl2
 3
 Trigonal
planar
 3
 0
 Triginal
planar
 BF3
 2
 1
 Bent
 NO2*
 4
 0
 Tetrahedral
 CCl4
 3
 1
 Trigonal
 pyramidal
 NH3
 2
 2
 Bent
 H2O
 4
 Tetrahedral
 17
 Molecules
with
Lone
Pairs
on
the
 Central
Atom(3.2) 
 Regions
of
 electron
 concentra5on
 Electronic
 arrangement
 Bonded
 pairs
 Lone
 pairs
 Shape
 Example
 5
 Trigonal
 bipyramidal
 5
 0
 Trigonal
 bipyramidal
 PCl5
 4
 1
 See‐saw
 IF4+
 3
 2
 T‐shaped
 ClF3
 2
 3
 Linear
 XeF2
 6
 0
 Octahedral
 SF6
 5
 1
 Square
 pyramidal
 BrF5
 4
 2
 Square
 planar
 XeF8

 14 6
 Octahedral
 Molecules
with
Lone
Pairs
on
the
 Central
Atom(3.2) 
 Figure 3.1 19
 Molecules
with
Lone
Pairs
on
the
 Central
Atom(3.2) 
 •  There
are
different
possible
locaGons
for
lone
 pairs
in
trigonal
bipyramidal
and
octahedral
 molecules.
 –  Axial lone pairs
lie
on
the
axis
on
the
molecule
 –  Equatorial lone pairs
lie
on
the
molecules’
 equator,
on
the
plane
perpendicular
to
the
 molecule’s
axis. Equitorial Axial 20
 Molecules
with
Lone
Pairs
on
the
 Central
Atom(3.2) 
 •  Trigonal
bipyramidal
 –  If
one
lone
pair
is
present,
should
it
be
axial
or
 equatorial?
Why?
 –  Two
lone
pairs?
Why?
 21
 Molecules
with
Lone
Pairs
on
the
 Central
Atom(3.2) 
 Figure 3.4 22
 Molecules
with
Lone
Pairs
on
the
 Central
Atom(3.2) 
 •  For
similar
reasons,
if
there
are
two
lone
pairs
 in
an
octahedral
molecule
they
must
be
on
 opposite
sides
of
the
central
atom.
 Figure 3.5 23
 Molecules
with
Lone
Pairs
on
the
 Central
Atom(3.2) 
 •  SecGon
summary:
 –  In a molecule that has lone pairs or a sinlge nonbonding electron on the central atom, the valence electrons contribute to the electron arrangement about the central atom but only bonded atoms are considered in the iden8fica8on of the shape. –  Lone pairs distort the shape of a molecule so as to reduce lone pair – bonding pair repulsion. 24
 Polar
Molecules
(3.3) 
 •  A
polar molecule
is
one
with
a
nonzero
dipole
 moment.
 •  Diatomic
molecules
are
polar
if
their
bonds
 are
polar.
 25
 Polar
Molecules
(3.3) 
 •  A
molecule
composed
of
ONLY
nonpolar
 bonds
is
a
nonpolar
molecule.
 •  Molecules
composed
of
polar
bonds
may
be
 polar
OR
nonpolar.
 26
 Polar
Molecules
(3.3) 
 •  If
dipoles
in
a
molecule
are
not
of
equal
 strength,
then
they
cannot
cancel
regardless
 of
shape.
 27
 Polar
Molecules
(3.3) 
 •  Example:

Is
CH2Cl2
polar
or
nonpolar?
 28
 Polar
Molecules
(3.3) 
 •  The
presence
of
lone
pairs
on
the
central
atom
 changes
the
shape
of
the
molecule
and
 prevents
dipoles
from
canceling.
 29
 Polar
Molecules
(3.3) 
 •  Nonpolar
molecules
are
symmetrical.
 –  There are identical atoms attached to the central atom. –  There are no lone pairs on the central atom. 
 •  Polar
molecules
are
asymmetrical.
 –  There are atoms of different elements attached to the central atom. –  Lone pairs are arranged asymmetrically around the central atom. 30
 Polar
Molecules
(3.3) 
 •  Example:

Predict
whether
each
of
the
 following
is
polar
or
nonpolar.
 –  BrF3
 –  SF4
 –  I3‐
 –  PCl4F
 31
 Polar
Molecules
(3.3) 
 •  SecGon
summary:
 –  A diatomic molecule is polar if its bond is polar. –  A polyatomic molecule is polar if it has polar bond arranged in space in such a way that the dipole moments associated with the bonds do not cancel. 32
 Valence‐Bond
Theory 
 •  Valence bond theory
describes
how
covalent
 bonding
takes
place.
 –  This
involves
the
overlap
of
atomic
orbitals
to
 form
mixed
or
hybridized
orbitals.
 33
 Sigma
and
Pi
Bonds
(3.4) 
 •  When
H2
forms,
the
single
 




1s
electron
of
each
H
atom
 




pair
and
their
atomic
 
orbitals
merge.
 •  The
resulGng
“sausage
–
 
shaped”
distribuGon
of
 
electrons
is
called
a
 
sigma bond
or
σ‐bond.
 Figure 3.8 34
 Sigma
and
Pi
Bonds
(3.4) 
 •  All
single
covalent
bonds
are
σ–bonds.
 –  The
greater
the
extent
of
orbital
overlap,
the
 stronger
the
bond.
 Figure 3.9 35
 Sigma
and
Pi
Bonds
(3.4) 
 •  Some
bonds
require
more
and
different
orbital
 overlap
than
occurs
in
σ–bonds.
 –  Consider
the
example
of
N2.
 Figure 3.11 Figure 3.10 36
 36
 Sigma
and
Pi
Bonds
(3.4) 
 •  The
π‐bond
has
a
single
nodal
plane
 containing
the
internuclear
axis.
 –  Though
there
is
electron
density
on
both
sides
of
 the
axis,
it is only one bond with the electron cloud in the form of two lobes.
 Part of Figure 3.11 37
 Sigma
and
Pi
Bonds
(3.4) 
 •  A
double
bond
is
composed
of
a
σ‐bond
PLUS
 a
π‐bond.
 •  A
triple
bond
is
composed
of
a
σ‐bond
PLUS
 TWO
π‐bonds.
 Figure 3.12 38
 Sigma
and
Pi
Bonds
(3.4) 
 •  SecGon
summary:
 –  In valence‐bond theory, we assume that bonds form when unpaired electrons in valence‐shell atomic orbitals pair. –  The atomic orbitals overlap end to end to from σ‐ bonds and side to die to form π‐bonds. 39
 Electron
Promo5on
and
the
 Hybridiza5on
of
Orbitals
(3.5) 
 •  If
we
try
to
apply
valence
bond
theory
to
CH4,
 we
run
into
trouble!
 •  The
electron
configuraGon
of
C
suggests
that
it
 should
only
be
able
to
form
two
bonds.

 Instead,
we
know
that
it
almost
always
forms
 four
bonds!
How?

Why?
 40
 Electron
Promo5on
and
the Hybridiza5on
of
Orbitals
(3.5) 
 •  For
CH4,
we
are
combining
one
s
and
three
p
 orbitals.


 –  The
four
orbitals
interfere
with
one
another
and
 produce
new
paoerns
where
they
intersect,
like
 waves
in
water.
 –  The
new
paoerns
represent
the
hybrid orbitals.
 41
 Electron
Promo5on
and
the
 Hybridiza5on
of
Orbitals
(3.5) 
 •  The
four
hybrid
orbitals
are
called
sp3 hybrids
 because
they
are
formed
from
one
s‐orbital
 and
three
p‐orbitals.
 2s and 2p orbitals Set of 4 sp3 hybrid orbitals Hybridization 42
 42 Electron
Promo5on
and
the
 Hybridiza5on
of
Orbitals
(3.5) 
 Figure from Whitten, et al. 43
 Electron
Promo5on
and
the
 Hybridiza5on
of
Orbitals
(3.5) 
 •  To
form
four
bonds,
we
want
C
to
have
four
 unpaired
electrons.


 –  We
can
achieve
this
by
promo8ng
one
of
the
 bonding
electrons
(a
2s)
into
an
empty
2p
orbital.
 44
 Electron
Promo5on
and
the
 Hybridiza5on
of
Orbitals
(3.5) 
 •  SecGon
summary:
 –  Hybrid orbitals are constructed on an atom to reproduce the electron arrangement characteris8cs of the experimentally determined shape of a molecule. –  The promo8on of electrons will occur if, overall, it leads to a lowering of energy by permiRng the forma8on of more bonds. 45
 Other
Common
Types
of
 Hybridiza5on
(3.6) 
 •  We
can
use
different
hybridizaGon
schemes
to
 describe
other
arrangements
of
electron
pairs.
 1
s
orbital
+
1
p
orbital
=
2
sp
hybridized
orbitals 
 1
s
orbital
+
2
p
orbitals
=
3
sp2
hybridized
orbitals 
 1
s
orbital
+
3
p
orbitals
=
4
sp3
hybridized
orbitals 
 1
s
orbital
+
3
p
orbitals
+
1
d
orbital
=
 
 5
sp3d
hybridized
orbitals 
 1
s
orbital
+
3
p
orbitals
+
2
d
orbitals
=
 
 6
sp3d2
orbitals 
 46
 Other
Common
Types
of
 Hybridiza5on
(3.6) 
 Regions of Electronic Hybridization High Electron Arrangement Density 2 Linear sp Trigonal 3 sp2 planar 4 Tetrahedral sp3 Trigonal 5 sp3d bipyramidal 6 Octahedral sp3d2 47
 Other
Common
Types
of
 Hybridiza5on
(3.6) 
 Figure 3.16 Figure 3.17 Figure 3.18 48
 Other
Common
Types
of
 Hybridiza5on
(3.6) 
 •  Example:

What
is
the
hybridizaGon
of
the
 central
atom
in
each
molecule?
 –  ICl3
 –  AsI5
 –  PH3
 –  C2H4

 49
 Other
Common
Types
of
 Hybridiza5on
(3.6) 
 •  SecGon
summary:


 –  A hybridiza8on scheme is adopted to match the electron arrangement of the molecule. –  Valence‐shell expansion requires the use of d‐ orbitals. 50
 Characteris5cs
of
Mul5ple
 
 Bonds
(3.7) 
 •  Double
bonds
are
rarely
found
between
atoms
 of
elements
in
Period
3
and
later,
but
they
 form
readily
between
Period
2
elements
C,
N,
 and
O.
 51
 Characteris5cs
of
Mul5ple
 
 Bonds
(3.7) 
 •  The
presence
of
C
to
C
double
bonds
strongly
 influences
the
shape
of
a
molecule.
 Figure 3.19 52
 Characteris5cs
of
Mul5ple
 
 Bonds
(3.7) 
 53
 Characteris5cs
of
Mul5ple
 
 Bonds
(3.7) 
 •  Remember
from
chapter
2
that
a
double
bond
 between
to
atoms
is
stronger
than
a
single
 bond,
but
not
twice
as
strong.


 –  Based on what you now know about σ‐ and π‐ bonds, can you explain this? 54
 Characteris5cs
of
Mul5ple
 
 Bonds
(3.7) 
 •  SecGon
summary:
 –  Mul8ple bonds are formed when an atom forms a σ‐bondby using an sp or sp2 hybrid orbital and one or more π‐bonds by using unhybridized p‐ orbitals. –  The side‐by‐side overlap that forms a π‐bond makes a molecule resistant to twis8ng, results in bonds weaker than σ‐bonds, and prevents atoms with large radii from forming mul8ple bonds. 55
 Molecular
Orbital
Theory 
 •  Molecular orbital theory
has
proved
to
be
the
 most
successful
theory
of
the
chemical
bond.
 –  Electrons
occupy
molecular orbitals
that
spread
 throughout
the
enGre
molecule.
 56
 Molecular
Orbitals
(3.9) 
 •  The
addiGon
of
wavefuncGons
is
called
 “forming
a
linear
combinaGon”
and
the
 molecular
orbital
described
is
called
a
linear combina8on of atomic orbitals (LCAO).
 Figure 3.25 57
 Molecular
Orbitals
(3.9) 
 •  Because
the
electron
now
occupies
a
greater
 volume
(in
the
molecular
orbital)
than
it
did
 when
it
was
confined
to
an
atomic
orbital,
it also has a lower kine8c energy.
 –  A
combinaGon
of
atomic
orbitals
that
results
in
an
 overall
lowering
of
energy
is
called
a
bonding orbital.
 58
 Molecular
Orbitals
(3.9) 
 •  An
important
part
of
MO
theory
is
that:
 N atomic orbitals combine to form N molecular orbitals 59
 Molecular
Orbitals
(3.9) 
 •  Since
two
atomic
orbitals
were
used
to
form
 H2
(one
from
each
H),
we
must
form
two
 LCAO‐MO.
 –  In
the
second
MO,
the
two
atomic
orbitals
 interfere
destruc1vely.
 Figure 3.26 60
 Molecular
Orbitals
(3.9) 
 •  If
an
electron
occupies
a
molecular
orbital
 orbital
resulGng
from
destrucGve
interference,
 then
it
is
largely
excluded
from
the
 internuclear
region
and
has a higher energy than when it occupies one of the atomic orbitals alone.
 61
 Molecular
Orbitals
(3.9) 
 Figure 3.27 62
 Molecular
Orbitals
(3.9) 
 •  SecGon
summary:
 –  Molecular orbitals are built from linear combina8ons of atomic orbitals. –  When atomic orbitals interfere construc8vely, they give rise to bonding orbitals. –  When they interfere destruc8vely, they give rise to an8bonding orbitals. –  N atomic orbitals combine to give N molecular orbitals. 63
 Electron
Configura5ons
of
 Diatomic
Molecules
(3.10) 
 •  Once
molecular
orbitals
have
been
constructed,
valence
 electrons
can
be
placed
in
them
using
the
same
building
 up
principle
that
is
used
for
atomic
orbitals.
 1. Electrons
are
accommodated
in
the
lowest‐energy
molecular
 orbitals,
then
in
orbitals
of
increasingly
higher
energy.
 2. According
to
the
Pauli
exclusion
principle,
each
molecular
 orbital
can
accommodate
up
to
two
electrons.

If
two
 electrons
are
present
in
one
orbital,
they
must
be
paired.
 3. If
more
than
one
molecular
orbital
of
the
same
energy
Is
 available,
the
electrons
enter
them
singly
and
adopt
parallel
 spins
(Hund’s
rule)
 64
 Electron
Configura5ons
of
 Diatomic
Molecules
(3.10) 
 Figure 3.28 65
 Electron
Configura5ons
of
 Diatomic
Molecules
(3.10) 
 Figure 3.31 Typical MO diagram for the homonuclear 66
 diatomic molecules Li2 through N2 Electron
Configura5ons
of
 Diatomic
Molecules
(3.10) 
 •  We
can
write
an
electron
configuraGon
for
the
 molecular
orbitals
just
as
we
did
for
the
 atomic.
 
N2:


 67
 Electron
Configura5ons
of
 Diatomic
Molecules
(3.10) 
 •  Example:
Deduce
the
ground‐state
electron
 configuraGon
of
the
O2
molecule.
 68
 Electron
Configura5ons
of
 Diatomic
Molecules
(3.10) 
 •  The
molecular
orbital
configuraGon
can
be
used
 to
calculate
the
bond order,
the
net
number
of
 bonds
 Bond
order
(b)
=
½
(number
of
e‐
in
bonding
 orbitals
–
number
of
e‐
in
anG‐bonding
orbitals) 
 b
=
½
(Ne
–
Ne*) 
 69
 Electron
Configura5ons
of
 Diatomic
Molecules
(3.10) 
 •  Example:

Calculate
the
bond
order
for
O2.
 70
 Electron
Configura5ons
of
 Diatomic
Molecules
(3.10) 
 •  SecGon
summary:
 –  The ground‐state electron configura8ons of diatomic molecules are deduced by forming molecular orbitals from all the valence‐shell atomic orbitals of the two atoms and adding the valence electrons to the molecular orbitals in order of increasing energy, in accord with the building‐up principle. 71
 Bonding
in
Heteronuclear
 Diatomic
Molecules
(3.11) 
 •  
Because
heteronuclear
bonds
are
polar,
the
 electrons
are
shared
unequally
by
the
two
 atoms.
 –  We
must
add
weighted
coefficients
to
the
 wavefuncGon
equaGon.
 72
 Bonding
in
Heteronuclear
 Diatomic
Molecules
(3.11) 
 •  The
relaGve
values
of
cA2
and
cB2
determine
 the
type
of
bond:
 –  Nonpolar
covalent
bond
 –  Polar
covalent
bond
 –  Ionic
bond
 73
 Bonding
in
Heteronuclear
 Diatomic
Molecules
(3.11) 
 Figure 3.33 74
 Bonding
in
Heteronuclear
 Diatomic
Molecules
(3.11) 
 •  SecGon
summary:
 –  Bonding in heteronuclear diatomic molecules involves an unequal sharing of the bonding electrons. –  The more electronega8ve element contributes more strongly to the bonding orbitals. –  The less electronega8ve element contributes more strongly to the an8bonding orbitals. 75
 Orbitals
in
Polyatomic
 
 Molecules
(3.12) 
 Figure 3.40 76
 ...
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