Chapter 1 Notes

Chapter 1 Notes - Chapter
1
 THE
BASICS


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Unformatted text preview: Chapter
1
 THE
BASICS
 BONDING
AND
MOLECULAR
STRUCTURE
 (1)  The
human
body
is
largely
composed
of
organic
compounds.
 Organic
Chemistry:
The
chemistry
of
the
 compounds
of
carbon
 (2)  Organic
chemistry
plays
a
central
role
in
medicine,
bioengineering
etc.
 Amoxicillin
 AZT
 Fluoxetine
 DDT
 Capsaicin
 Vitalism
 It
was
originally
thought
organic
compounds
could
be
made
only
by
living
things
by
 intervention
of
a
“vital
force”.
 Fredrich
Wöhler
disproved
vitalism
in
1828
by
making
the
organic
compound
urea
 from
the
inorganic
salt
ammonium
cyanate
by
evaporation:
 Inorganic
 Organic
 The
Structural
Theory
of
Organic
Chemistry
 Between
1858
and
1861,
Kekule,
Couper,
and
Butlerov
provided
the
foundations
 for
structural
theory.
 Central
Premises:
 (1)
Valency:
atoms
in
organic
compounds
form
a
 fixed
number
of
bonds.
 (2)
Carbon
can
form
one
or
more
bonds
to
other
carbons.
 (1)
Isomers
are
different
molecules
with
the
same
molecular
formula.
 (2)
Many
types
of
isomers
exist.
 




Example:
 Consider
two
compounds
with
molecular
formula
C2H6O.

These
compounds
cannot
be
 distinguished
based
on
molecular
formula.

However
they
have
different
structures.

The
 two
compounds
differ
in
the
connectivity
of
their
atoms.
 Isomers
 Constitutional
Isomers
 Constitutional
isomers
are
one
type
of
isomer.

They
are
different
compounds
that
 have
the
same
molecular
formula
but
different
connectivity
of
atoms.

They
often
 differ
in
physical
properties
(e.g.
boiling
point,
melting
point,
density)
and
chemical
 properties.
 In
Class
Problems:
 (1)
Draw
all
the
constitutional
isomers
of
C4H10O.
 Three
Dimensional
Shape
of
Molecules
 Virtually
all
molecules
possess
a
3‐dimensional
shape
which
is
often
not
accurately
 represented
by
drawings,
 It
was
proposed
in
1874
by
van’t
Hoff
and
le
Bel
that
the
four
bonds
around
carbon
 where
not
all
in
a
plane
but
rather
in
a
tetrahedral
arrangement
i.e.
the
four
C‐H
bonds
 point
towards
the
corners
of
a
regular
tetrahedron.
 Three
Dimensional
Shape
of
Molecules
 Consider
the
constitutional
isomers
ethyl
alcohol
and
dimethyl
ether:
 Ethyl
Alcohol
 Dimethyl
Ether
 Chemical
Bonds:
The
Octet
Rule
 Octet
Rule:
 (1)
Atoms
form
bonds
to
produce
the
electron
configuration
of
a
noble
gas
because
 the
electronic
configuration
of
noble
gases
is
particularly
stable.

 (2)
For
most
atoms
of
interest
this
means
achieving
a
valence
shell
configuration
of
8
 electrons
corresponding
to
that
of
the
nearest
noble
gas.

 (3)
Atoms
close
to
helium
achieve
a
valence
shell
configuration
of
2
electrons.
 (4)
Atoms
can
form
either
ionic
or
covalent
bonds
to
satisfy
the
octet
rule.
 Electronegativity
 (1)
Electronegativity
is
the
ability
of
an
atom
to
attract
electrons.
 (2)
Electronegativity
increases
from
left
to
right
and
from
bottom
to
top
in
the
 periodic
table
(noble
gases
excluded).
 (3)
Fluorine
is
the
most
electronegative
atom
and
can
stabilize
excess
electron
density
 the
best.
 Ionic
Bonds
 (1)  When
ionic
bonds
are
formed
atoms
gain
or
lose
electrons
to
achieve
the
electronic
 configuration
of
the
nearest
noble
gas;
there
is
a
transfer
ofelectrons
from
one
 atom
to
another.
 (2)  The
resulting
oppositely
charged
ions
attract
and
form
ionic
bonds.
 (3)  This
generally
happens
between
atoms
of
widely
different
electronegativities.
 Example:
Lithium
Fluoride
(a
salt)
 (1)
Lithium
loses
an
electron
(to
have
the
configuration
of
helium)
and
becomes
 positively
charged.

 (2)
Fluoride
gains
an
electron
(to
have
the
configuration
of
neon)
and
becomes
 negatively
charged.
 (3)
The
positively
charged
lithium
and
the
negatively
charged
fluoride
form
a
strong
 ionic
bond
(actually
in
a
crystalline
lattice).
 Sodium
Chloride‐
An
Ionic
Salt
 Covalent
Bonds
 (1)
Covalent
bonds
occur
between
atoms
of
similar
electronegativity
(close
to
each
 other
in
the
periodic
table).
 (2)
Atoms
achieve
octets
by
sharing
of
valence
electrons.
 (3)
Molecules
result
from
this
covalent
bonding.
 (4)
Valence
electrons
can
be
indicated
by
dots
(electron‐dot
formula
or
Lewis
structures)
 but
this
is
time‐consuming.
 (5)
The
usual
way
to
indicate
the
two
electrons
in
a
bond
is
to
use
a
line
(one
line
=
two
 electrons).
 Writing
Lewis
Structures
 (1)  Atoms
bond
by
using
their
valence
electrons.
 (2)
The
number
of
valence
electrons
is
equal
to
the
group
number
of
the
atom.
 
 (a)
Carbon
is
in
group
4A
and
has
4
valence
electrons.
 
 (b)
Hydrogen
is
in
group
1A
and
has
1
valence
electron.
 
 (c)
Oxygen
is
in
group
6A
and
has
6
valence
electrons.
 
 (d)
Nitrogen
is
in
group
5A
and
has
5
valence
electrons.
 (3)
To
construct
molecules
the
atoms
are
assembled
with
the
correct
number
of
valence
 electrons.
 (4)
If
the
molecule
is
an
ion,
electrons
are
added
or
subtracted
to
give
it
the
proper
 charge.
 (5)
The
structure
is
written
to
satisfy
the
octet
rule
for
each
atom
and
to
give
the
correct
 charge.
 (6)
If
necessary,
multiple
bonds
can
be
used
to
satisfy
the
octet
rule
for
each
atom.
 Lewis
Symbols
 (a)  The
chemical
symbol
represents
the
nucleus
and
the
core
electrons.
 (b)  The
dots
around
the
symbol
represent
the
valence
electrons.
 Some
examples:
 Lewis
Structures
for
Ionic
Compounds:
 Barium
Oxide
 Magnesium
Chloride
 Lewis
Structures
for
Some
Small
Molecules
 (1)
Hydrogen
Fluoride
 (2)
Water
 (3)
Ammonia
 (4)
Methane
 Lewis
Structures
with
Multiple
Covalent
Bonds
 (1)
Nitrogen
 (2)
Carbon
Dioxide
 Additional
Lewis
Structures
 (1)
Phosgene
 (2)
Hypochlorous












 Acid
 (3)
Chlorate
Ion
 (4)
Methanol
 In
Class
Problems:
 (1)  Draw
the
Lewis
structures
for
each
of
the
following
species.

The
atom
 connectivity
is
indicated.
 
 (a)
HN3


(hydrazoic
acid)

H

N

N

N
 
 (b)
HNO3

(nitric
acid)

H

O

N

O
 























































O
 
 
 
 (c)
CO3‐2
(carbonate
ion)









O


‐2
 















































O

C

O
 (d)
H2C2
(acetylene)


H

C

C

H
 Exceptions
to
the
Octet
Rule
 The
octet
rule
applies
only
to
atoms
in
the
second
row
of
the
periodic
table
(C,
O,
N,
F)
 which
are
limited
to
valence
electrons
in
the
2s
and
2p
orbitals
 In
second
row
elements
fewer
electrons
are
possible
 Example:
BF3
 In
higher
rows
other
orbitals
are
accessible
and
more
than
8
electrons
around
an
atom
 are
possible
 Example:
PCl5
and
SF6
 (1)
A
formal
charge
is
a
positive
or
negative
charge
on
an
individual
atom.
 (2)
The
sum
of
formal
charges
on
individual
atoms
is
the
total
charge
of
the
molecule
or
 ion.
 (3)
The
formal
charge
is
calculated
by
subtracting
the
assigned
electrons
on
the
atom
in
 the
molecule
from
the
electrons
in
the
neutral
atom.
 (4)
Electrons
in
bonds
are

evenly
split
between
the
two
atoms;
one
to
each
atom.
 (5)
Lone
pair
electrons
belong
to
the
atom
itself.
 Formal
charge
 Ammonium
ion
(NH4)+
 Nitrate
ion
(NO3)‐
 An
atom
will
always
have
the
same
formal
charge
depending
on
how
many
bonds
 and
lone
pairs
it
has
regardless
of
which
particular
molecule
it
is
in.

For
example
a
 singly
bonded
oxygen
with
3
lone
pairs
will
always
have
a
negative
charge
and
an
 oxygen
with
three
bonds
and
one
lone
pair
will
always
have
a
positive
charge.
 Knowing
these
forms
of
each
atom
is
invaluable
in
drawing
Lewis
structures
correctly
 and
rapidly
.
 In
Class
Problems:
 (1)

Draw
the
Lewis
structure
for
each
of
the
following
species
and
indicate
the
 formal
charge
associated
with
each
atom.

The
atom
connectivity
is
indicated.
 (a)
NO3‐

(nitrate)


O

N

O
 
 






























O
 (b)
N3‐

(azide)


N

N

N
 Resonance
 Often
a
single
Lewis
structure
does
not
accurately
represent
the
true
structure
of
a
 molecule.

The
real
carbonate
ion
is
not
represented
by
any
of
the
structures
1,2
or
3.
 Experimentally
carbonate
is
known
to
have
all
bonds

equal
in
length
and
the
charge
is
 spread
equally
over
all
three
oxygens.

The
“real”
carbonate
ion
can
be
represented
by
a
 drawing
in
which
partial
double
bonds
to
the
oxygens
are
shown
and
a
partial
negative
 charge
exits
on
each
oxygen.

The
“real”
structure
is
a
resonance
hybrid
or
mixture
of
all
 three
Lewis
structures.

Double
headed
arrows
are
used
to
show
the
resonance
 contributors
to
the
experimental
structure.

NEVER
USE
EQUILIBRIUM
ARROWS!
 One
resonance
contributor
is
converted
to
another
by
the
use
of
curved
arrows
which
 show
the
movement
of
electrons.

The
use
of
these
arrows
serves
as
a
bookkeeping
device
 to
assure
all
structures
differ
only
in
position
of
electrons
 A
calculated
electrostatic
potential
map
of
carbonate
clearly
shows
the
electron
density
is
 spread
equally
among
the
three
oxygens.

Areas
which
are
red
are
more
negatively
 charged;
areas
of
blue
have
relatively
less
electron
density.
 Rules
for
Resonance
 (1)  Individual
resonance
structures
exist
only
on
paper.

The
real
molecule
is
a
hybrid
 (average)
of
all
contributing
forms.
 (2)  Resonance
forms
are
indicated
by
the
use
of
double‐headed
arrows.
 (3)  Only
electrons
are
allowed
to
move
between
resonance
structures.
 (4) The
position
of
nuclei
must
remain
the
same.
 (5) Usually
only
electrons
in
multiple
bonds
and
nonbonding
electrons
can
be
moved.
 Example:
3
is
not
a
resonance
form
because
an
atom
has
moved
 (6)
All
structures
must
be
proper
Lewis
structures.
 (7)
The
energy
of
the
actual
molecule
is
lower
than
the
energy
of
any
single
contributing
 form.

The
lowering
of
energy
is
called
resonance
stabilization.
 (8)
Equivalent
resonance
forms
make
equal
contributions
to
the
structure
of
the
real
 molecule.

In
general,
structures
with
equivalent
resonance
forms
tend
to
be
greatly
 stabilized.
 Example:
The
two
resonance
forms
of
benzene
contribute
equally
and
greatly
stabilize
it
 (9)
Unequal
resonance
structures
contribute
based
on
their
relative
stabilities.

 (10)
More
stable
resonance
forms
contribute
more
to
the
structure
of
the
real
molecule

 (1)
A
resonance
form
with
more
covalent
bonds
is
more
important
than
one
with
less.
 Example:

6
is
more
stable
and
more
important
because
it
has
more
total
covalent
 bonds
 Rules
to
Assign
Relative
Importance
of
 Resonance
Forms
 (2)
Resonance
forms
in
which
all
atoms
have
a
complete
valence
shell
of
electrons
are
 more
important.
 Example:

10
is
more
important
because
all
atoms
(except
hydrogen)
have
complete
 octets
 (3)
Resonance
forms
with
separation
of
charge
are
less
important.

Separation
of
 charge
cost
energy
and
results
in
a
less
stable
resonance
contributor.
 Example:

12
is
less
important
because
it
has
charge
separation
 (4)
Forms
with
negative
charge
on
highly
electronegative
atoms
are
more
important.

 Those
with
positive
charge
on
less
electronegative
atoms
are
also
more
important.
 Example:

Nitrate
Ion
 The
nitrate
ion
is
known
to
have
all
three
nitrogen‐oxygen
bond
lengths
the
same
and
 the
negative
charge
spread
over
all
three
atoms
equally.
 Resonance
theory
can
be
used
to
produce
three
equivalent
resonance
forms
.

Curved
 arrows

show
the
movement
of
electrons
between
forms.

When
these
forms
are
 hybridized
(averaged)
the
“true”
structure
of
the
nitrate
ion
is
obtained.
 In
Class
Problems:
 (1)  Draw
the
contributing
resonance
structures
for
each
of
the
following
species.
 (a)  Benzyl
cation
 (b)  Acetone
enolate


 Quantum
Mechanics
(Wave
Mechanics),
Atomic
 Orbitals,
and
Molecular
Orbitals
 Quantum
mechanics
(wave
mechanics)
provides
a
model
for
calculating
the
energies
 and
probability
locations
of
negatively
charged
electrons
in
the
field
of
positively
 charged
nuclei.

The
model
is

based
upon
the
wave
nature
of
electrons.
 A
hydrogen
atom
can
be
simply
described
as
an
electron
in
the
field
of
a
proton.

The
 application
of
the
principles
of
quantum
mechanics
to
this
system
provides
the
energies
 and
probability
locations
of
the
electron
in
the
ground
and
excited
states
of
the
 hydrogen
atom.

Each
state
is
mathematically
described
by
a
wave
function,
ψ.

The
 wave
functions
associated
with
each
of
these
states
provide
the
foundation
for
our
 pictorial
representation
of
the
various
atomic
orbitals‐
s,
p,
d,
and
f
orbitals.

The
square
 of
the
wave
function
(ψ2)is
related
to
the
probability
of
finding
an
electron
in
a
 particular
location
in
space.
 An
atomic
orbital
is
defined
as
a
region
of
space
where
there
is
a
90
–
95%
probability
of
 finding
an
electron.
 Atomic
orbitals
on
different
atoms
located
proximate
to
one
another
interact
(overlap)
 to
form
molecular
orbitals.

Molecular
orbitals
provide
the
foundation
for
our
 understanding
of
molecular
structure.
 The
“Shapes”
of
Some
Atomic
Orbitals
 (1)
The
1s
and
2s
atomic
orbitals
have
spherical
symmetry
centered
around
the
 nucleus.

The
2s
orbital
is
higher
in
energy
as
compared
to
the
1s
orbital
and
contains
a
 nodal
surface
(Ψ
=
0).

Each
s
orbital
can
contain
a
maximum
of
2
electrons.
 (2)
The
2p
atomic
orbital
has
a
nodal
plane
of
symmetry‐
the
node
passing
through
the
 nucleus.

There
are
three
2p
orbitals
which
are
perpendicular
(orthoganol)
to
each
 other‐

2px,
2py,
and
2pz.

The
three
2p
orbitals
are
degenerate
(equal
in
energy).

Each
 p
orbital
can
contain
a
maximum
of
2
electrons.
 (3)
The
2p
orbitals
are
higher
in
energy
than
the
1s
or
2s
atomic
orbitals.
 (4)
Generalization:
The
greater
the
number
of
nodes
the
higher
the
energy.
 Atoms
can
be
assigned
electronic
configuration
using
the
following
rules:
 (1)
Aufbau
Principle:
The
lowest
energy
orbitals
are
filled
first.
 (2)
Pauli
Exclusion
Principle:
A
maximum
of
two
spin
paired
electrons
may
be
placed
in
 each
orbital.
 (3)
Hund’s
Rule:
One
electron
is
added
to
each
degenerate
orbital
before
a
second
 electron
is
added.
 Electronic
Configurations
of
Some
Second
Row
Elements
 Molecular
Orbitals
(MOs)
 A

simple
model
of
bonding
is
illustrated
by
forming
molecular
H2
from
two
H
atoms.

 The
following
diagram
illustrates
the
relationship
between
the
two
H
atoms
as
a
 function
of
distance:
 Region
I:

The
total
energy
of
two
isolated
hydrogen
atoms.

The
hydrogen
atoms
do
 not

“feel”
each
other ’s
presence.
 Region
II:
As
the
two
hydrogen
atoms
approach
one
another
their
1s
orbitals
begin
to
 overlap
initiating
the
formation
of
a
bond;
the
energy
of
the
system
is
lowered.
 Region
III:
at
0.74
Å

the
attractive
overlap
between
the
two
1s
orbitals
is
balanced
by
 the
repulsive
nuclear
interaction;
this
minimum
represents
the
bond
length
of
H2.
 Region
IV:
energy
of
system
rises
as
the
repulsion
of
the
two
nuclei
predominates.
 (1)  The
overlap
of
1s
orbitals
on
each
of
the
hydrogen
atoms
results
in
two
molecular
 orbitals‐
one
of
lower
energy
(the
bonding
molecular
orbital‐
addition
of
the
wave
 functions)
and
one
of
higher
energy
(the
anti‐bonding
molecular
orbital‐
 subtraction
of
the
wave
functions).

Generalization:

The
combination
of
n
atomic
 orbitals
results
in
the
formation
of
n
molecular
orbitals.
 (2)  The
molecular
orbitals
(MO’s)
that
form
encompass
both
nuclei.

The
electrons
are
 not
restricted
to
the
vicinity
of
one
nucleus
or
another.
 (3)  Each
of
the
resulting
MO’s
has
a
maximum
of
2
spin‐paired
electrons.
 (4)  The
mathematic
operation
by
which
wave
functions
are
added
or
subtracted
is
 called
the
linear
combination
of
atomic
orbitals
(LCAO).
 The
energy
of
electrons
in
the
bonding
molecular
orbital
is
substantially
less
than
the
 energy
of
electrons
in
the
individual

(isolated)
H
atoms.
 The
energy
of
electrons
in
the
antibonding
molecular
orbital
is
substantially
more
 than
the
energy
of
the
electrons
in
the
individual
(isolated)
H
atoms.
 The
two
electrons
derived
from
the
hydrogen
atoms
occupy
the
bonding
molecular
 orbital.

The
anti‐bonding
molecular
orbital
is
empty
(virtual
molecular
orbital).
 The
Structure
of
Methane
and
Ethane:
 sp3
Hybridization
 The
structure
of
methane
with
its
tetrahedral
geometry
can
be
accomodated
by
the
 combination
(hybridization)
of
the
2s
and
the
three
2p
atomic
orbitals
to
form
four
sp3
 hybrized
atomic
orbitals.

The
four
sp3
orbitals
can
then
be
combined
with
the
1s
 orbitals
of
four
hydrogens
to
give
the
molecular
orbitals
of
methane.
 An
sp3
orbital
looks
like
a
p
orbital
with
one
lobe
greatly
extended.

This
hybrid
orbital
 can
be
looked
upon
as
a
“polarized”
p
atomic
orbital.


 The
extended
sp3
lobe
can
then
effectively
overlap
with
the
hydrogen
1s
to
form
a
strong
 bond.
 These
type
of
bond
have
cylindrical
symmetry
when
viewed
along
the
bond
axis;
these
 type
of
bond
are
referred
to
as
sigma
(σ)
bonds.

 A
variety
of
representations
of
methane
show
its
tetrahedral
nature
and
electron
 distribution:
 a.  calculated
electron
density
surface
 b.  ball‐and‐stick
model
c.
a
typical
 c.  3‐dimensional
drawing
 Ethane
(C2H6)
 The
carbon‐carbon
bond
is
made
from
overlap
of
two
sp3
orbitals
to
form
a
carbon‐ carbon
σ
bond.

Six
hydrogen
1s
atomic
orbitals
combine
(overlap)
with
the
remaining
 sp3
hybridized
orbitals.

The
molecule
is
approximately
tetrahedral
around
each
of
the
 carbon
atoms.
 The
representations
of
ethane
show
the
tetrahedral
arrangement
around
each
of
the
 carbon
atoms.
 a.  calculated
electron
density
surface
 b.  ball‐and‐stick
model
 c.  typical
3‐dimensional
drawing
 Generally
there
is
relatively
“free”
rotation
about
the
carbon‐carbon
σ
bonds.

Very
little
 energy
(13‐26
kcal/mol)
is
required
to
rotate
around
the
carbon‐carbon
bond
of
ethane.
 The
Structure
of
Ethene
(Ethylene):
 sp2
Hybridization
 Ethene
(C2H2)
contains
a
carbon‐carbon
double
bond
and
is
in
the
class
of
organic
 compounds
called
alkenes.

Another
example
of
the
alkenes
is
propene
(propylene).
 Overall
Structure
of
Ethene:
 The
geometry
around
each
carbon
is
called
trigonal
planar.

All
atoms
directly
 connected
to
each
carbon
are
in
a
plane.

The
bonds
point
towards
the
corners
of
a
 regular
triangle.

The
bond
angles
are
approximately
120o.
 There
are
three
σ
bonds
around
each
carbon
of
ethene
and
these
are
formed
by
using
 sp2
hybridized
orbitals.

The
three
sp2
hybridized
orbitals
come
from
mixing
one
2s
and
 two
2p
atomic
orbitals.

One
p
orbital
is
left
untouched
(unhybridized).
 The
sp2
orbitals
are
arranged
in
a
trigonal
planar
arrangement.

The
p
orbital
is
 perpendicular
(orthoganol)
to
the
plane.
 Overlap
of
sp2
orbitals
in
ethylene
results
in
formation
of
a
σ
framework.

One
sp2
 orbital
on
each
carbon
overlaps
to
form
a
carbon‐carbon
σ
bond;
the
remaining
sp2
 orbitals
form
bonds
to
hydrogen.

The
leftover
p
orbitals
on
each
carbon
overlap
to
 form
a
bonding
π
bond
(lower
energy)
and
an
anti‐bonding
π*
bond
(higher
energy)
 between
the
two
carbons.

These
types
of
π
interactions
result
from
parallel
overlap
of
 p
orbitals
above
and
below
the
plane
of
the
σ
bond.

A
π
bond
is
characterized
by
nodal
 symmetry;
it
has
a
nodal
plane
passing
through
the
two
bonded
nuclei
and

between
 the
two
lobes
of
the
π
molecular
orbital.
 The
bonding
π
orbital
results
from
overlap
of
p
orbital
lobes
of
the
same
sign.

The
 antibonding
π*
orbital
results
from
overlap
of
p
orbital
lobes
of
opposite
sign.

The
 antibonding
orbital
has
one
node
connecting
the
two
nuclei
and
another
node
between
 the
two
carbons.

The
bonding
π
orbital
is
lower
in
energy
than
the
antibonding
orbital.

 In
the
ground
state
two
spin
paired
electrons
are
in
the
bonding
orbital.

The
antibonding

 π*orbital
can
become
occupied
if
an
electron
is
promoted
(excited)
from
a
lower
level
 (e.g.
by
absorption
of
light).

 The
σ
molecular
orbital
is
lower
in
energy
than
the
π
molecular
orbital.

The
ground
 state
electronic
configuration
of
ethene
is
shown
as
follows:
 Restricted
Rotation
and
the
Double
Bond
 There
is
a
large
energy
barrier
to
rotation
(approximately
264
kJ/mol)
around
the
 double
bond.

This
corresponds
to
the
strength
of
a
π
bond.

In
comparison,
the
 rotational
barrier
of
a
carbon‐carbon
single
bond
is
13‐26
kJ/mol.

This
rotational
 barrier
results
because
the
p
orbitals
must
be
parallel
to
one
another
for
maximum
 overlap.


Rotation
of
one
of
the
p
orbitals
90o
breaks
the
π
bond.
 Cis‐trans
isomers
 Cis‐trans
isomers
are
the
result
of
restricted
rotation
about
double
bonds.

These
 isomers
have
the
same
atom
connectivity
and
differ
only
in
the
arrangement
of
atoms
 in
space.

This
places
them
in
the
broader
class
of
stereoisomers.

The
molecules
below
 do
not
superpose
on
each
other.

One
molecule
is
designated
cis
(substitutent
groups
 on
same
side
of
the
double
bond)
and
the
other
is
trans
(substituent
groups
on
 opposite
sides
of
the
double
bond).
 Cis‐trans
isomerism
is
not
possible
if
one
carbon
of
the
double
bond
has
two
identical
 groups
attached
to
it.
 The
Structure
of
Ethyne
(Acetylene):
 sp
Hybridization
 Ethyne
(acetylene)
is
a
member
of
a
group
of
compounds
called
alkynes;
alkynes
have
 carbon‐carbon
triple
bonds.

Propyne
is
another
typical
alkyne
 The
arrangement
of
atoms
around
each
carbon
is
linear
with
bond
angles
180o.
 The
carbons
of
ethyne
are
sp
hybridized.

One
2s
and
one
2p
atomic
orbitals
are
 mixed
(hybridized)
to
form
two
sp
orbitals.

Two
p
orbitals
are
left
untouched
 (unhybridized).
 The
two
sp
orbitals
on
a
particular
carbon
atom
are
oriented
180o
relative
to
each
 other
around
the
carbon
nucleus.

The
two
p
orbitals
are
perpendicular
to
the
axis
 that
passes
through
the
center
of
the
sp
orbitals.
 In
ethyne
the
sp
orbitals
on
the
two
carbons
overlap
to
form
a
σ
bond.

The
remaining
 sp
orbitals
overlap
with
two
hydrogen
1s
orbitals.

The

p
orbitals
on
each
carbon
 overlap
to
form
two
π
bonds.

The
triple
bond
consists
of
one
σ
and
two
π
bonds.
 Bond
Lengths
of
Ethyne,
Ethene
and
Ethane
 The
carbon‐carbon
bond
length
is
shorter
and
stronger
as
more
bonds
hold
the
carbons
 together.

The
carbon‐hydrogen
bond
lengths
also
get
shorter
the
greater
the
s
character
 of
the
orbital
forming
the
bond.

(The
electrons
in
2s
orbitals
are
held
more
tightly
to
the
 nucleus
than
electrons
in
2p
orbitals.)

A
hybridized
orbital
with
more
percent
s
character
 is
held
more
closely
to
the
nucleus
than
an
orbital
with
less
s
character
 The
sp
orbitals
associated
with
the
carbon
atoms
of
ethyne
have
50%
s
character;
the
sp2
 hybridized
orbitals
of
ethene
have
33%
s
character;
the
sp3
orbitals
of
ethane
have
only
 25%
s
character.

 In
Class
Problems:
 Norethindrone
is
a
steroidal
oral
contraceptive.

The
structure
is
drawn
below.
 (1)  Indicate
the
hybridization
at
each
of
the
atoms
designated
by
the
arrows.
 (2)  Indicate
the
geometry
at
the
circled
atoms.

 (3)
What
is
the
empirical
formula?
 (4)
What
is
the
molecular
formula?
 (5)
What
is
the
molecular
weight?
 Summary
of
Concepts
from
Quantum
Mechanics
 Atomic
Orbital(AO):
region
in
space
around
a
nucleus
where
there
is
a
high
probability
 of
finding
an
electron
 Molecular
Orbital
(MO):
results
from
overlap
of
atomic
orbitals
 Bonding
Orbitals:
overlap
of
AOs
of
the
same
sign
 Antibonding
Orbitals:

overlap
of
AOs
of
opposite
sign
 When
two
atomic
orbitals
overlap
to
form
two
molecular
orbitals
the
energy
of
the
 bonding
combination
is
less
than
the
energies
of
the
individual
atomic
orbitals
and
the
 anti‐bonding
combination
is
greater
than
the
energies
of
the
individual
atomic
orbitals.

 The
combination
of
n
atomic
orbitals
results
in
the
formation
of
n
molecular
orbitals.
 Hybridized
atomic
orbitals
are
created
by
mixing
unhybridized
atomic
orbital
wave
 functions.

For
instance:
 (1)  Four
sp3
orbitals
are
obtained
from
mixing
one
s
and
three
p
orbitals
.

The
geometry
 of
the
four
hybrid
orbitals
is
tetrahedral.

sp3
hybrids
are
used
to
describe
the
 geometry
of
the
methane
molecule.

 (2)  Three
sp2
orbitals
are
obtained
from
mixing
one
s
and
two
p
orbitals.

The
geometry
 of
the
three
orbitals
is
trigonal
planar.

The
remaining
unhybridized

p
orbital
can
be
 used
to
make
a
π
bond.

sp2
hybrids
are
used
to
describe
the
geometry
and
carbon‐ carbon
bond
rotation
restriction
of
ethene.
 (3)  Two
sp
orbitals
are
obtained
from
mixing
one
s
and
one
p
orbital.


The
geometry
of









 the
two
orbitals
is
linear.

The
two
remaining
unhybridized
p
orbitals
can
be
used
to
 make
two
π
bonds.

sp
hybrids
are
used
to
describe
the
geometry
of
ethyne.
 Sigma
(σ)
bonds
have
cylindrical
symmetry
when
viewed
along
the
bond
axis.
 Pi
(π)
bonds
have
nodal
symmetry
and
result
from
sideways
overlap
of
two
p
orbitals
.

 Molecular
Geometry:
The
Valence
Shell
Electron
 Pair
Repulsion
(VSEPR)
Model
 This
is
a
simple
theory
to
predict
the
geometry
of
molecules.

All
sets
of
valence
 electrons
are
considered
including:
 (1)  Bonding
pairs
involved
in
single
or
multiple
bonds.
 (2)  Non‐bonding
pairs
which
are
unshared
 Electron
pairs
repel
each
other
and
tend
to
be
as
far
apart
as
possible
from
each
 other.
 Non‐bonding
electron
pairs
tend
to
repel
other
electrons
more
than
bonding
pairs.
 The
geometry
of
the
molecule
is
determined
by
the
number
of
sets
of
electrons
by
 using
simple
geometrical
principles.
 Methane
 The
valence
shell
of
methane
contains
four
pairs
or
sets
of
electrons.

To
be
as
far
 apart
from
each
other
as
possible
they
adopt
a
tetrahedral
arrangement
(bond
 angle
109.5o).

The
molecule
methane
is
therefore
tetrahedral.
 Ammonia
 When
the
bonding
and
nonbonding
electrons
are
considered
there
are
4
pairs
or
sets
 of
electrons.

In
order
for
the
electron
pairs
to
be
as
far
away
from
each
other
as
 possible,
the
electron
pairs
should
be
at
the
corners
of
a
regular
tetrahedron.


 However,
because
the
non‐bonding
electron
pair
is
“more
repulsive”
than
the
 bonding
electron
pairs
the
actual
location
of
the
electron
pairs
is
described
by
a
 “distorted”
tetrahedral
arrangement.

Taking
into
account
only
the
positions
of
the
 nuclei,
the
molecule
is
described
as
trigonal
pyramidal.

The
actual
bond
angles
are
 about
107o
and
not
109.5o.


 Water
 There
are
four
pairs
or
sets
of
electrons
including
2
bonding
pairs
and
2
non‐bonding
 pairs.

Again,
the
position
of
the
electron
pairs
would
be
at
the
corners
of
a
regular
 tetrahedon.

However,
because
the
non‐bonding
electron
pairs
are
more
repulsive
 than
the
bonding
electron
pairs
the
actual
location
of
the
electron
pairs
is
described
 by
a
“distorted”
tetrahedral
arrangement.

Taking
into
account
only
the
positions
of
 the
nuclei,
the
molecule
is
described
as
angular.

The
bond
angle
is
about
105o.

 Boron
Trifluoride
 Three
sets
of
bonding
electrons
are
farthest
apart
in
a
trigonal
planar
arrangement
 (bond
angle
120o).

The
three
fluorides
lie
at
the
corners
of
an
equilateral
triangle
 Beryllium
Hydride
 Two
sets
of
bonding
electrons
are
farthest
apart
in
a
linear
arrangement
(bond
angles
 180o).
 Carbon
Dioxide
 There
are
only
two
sets
of
electrons
around
the
central
carbon
and
so
the
molecule
is
 linear
(bond
angle
180o).

Electrons
in
multiple
bonds
are
considered
as
one
set
of
 electrons.
 A
summary
of
the
results
also
includes
the
geometry
of
other
simple
molecules
 In
Class
Problems:
 Using
VSEPR
predict
the
geometry
of
each
of
the
following:
 (1)  BH4‐
 (2)  BH3
 (3)  H2S
 Representations
of
Structural
Formulas
 Dot
formulas
are
more
cumbersome
to
draw
than
dash
formulas

and
condensed
 formulas.

Lone‐pair
electrons
are
often
(but
not
always)
drawn
in,

especially
 when
they
are
crucial
to
the
chemistry
being
discussed.
 Dash
formulas
 Each
dash
represents
a
pair
of
electrons.

This
type
of
representation
is
meant
to
 emphasize
connectivity
and
does
not
represent
the
3‐dimensional
nature
of
the
molecule.


 For
instance,
the
dash
formulas
of
n‐propyl
alcohol
appear
to
have
90o
angles
for
carbons
 which
actually
have
tetrahedral
bond
angles
(109.5o).

There
is
relatively
free
rotation
 around
single
bonds
so
the
dash
structures
below
are
all
equivalent.
 Constitutional
isomers:

Constitutional
isomers
have
the
same
molecular
formula
but
 different
connectivity.

n‐Propyl
alcohol
(above)
is
a
constitutional
isomer
of
isopropyl
 alcohol
(below).
 Condensed
Structural
Formulas
 In
these
representations,
some
or
all
of
the
dash
lines
are
omitted.

In
partially
 condensed
structures
all
hydrogens
attached
to
an
atom
are
simply
written
after
it
 but
some
or
all
of
the
other
bonds
are
explicitly
shown.

In
fully
condensed
structure
 all
bonds
are
omitted
and
atoms
attached
to
carbon
are
written
immediately
after
it.

 For
emphasis,

branching
groups
are
often
written
using
vertical
lines
to
connect
them
 to
the
main
chain.
 Bond‐Line
Formulas
 A
further
simplification
of
drawing
organic
molecules
is
to
completely
omit
all
carbons
 and
hydrogens
and
only
show
heteroatoms
(e.g.
O,
Cl,
N)
explicitly.

Each
intersection
or
 end
of
line
in
a
zig‐zag
represents
a
carbon
with
the
appropriate
amount
of
hydrogens.


 Heteroatoms
with
attached
hydrogens
must
be
drawn
in
explicitly.
 Cyclic
compounds
are
condensed
using
a
drawing
of
the
corresponding
polygon.
 Multiple
bonds
are
indicated
by
using
the
appropriate
number
of
lines
connecting
the
 atoms.
 In
Class
Problems:
 (1)
Represent
the
following
condensed
structural
formulas
as
a
bond‐line
formula:
 
 CH2=C(CH3)CH(OH)CH3
 (2)
Represent
the
following
dash
formula
as
a
condensed
structural
formula:
 Three‐Dimensional
Formulas
 Since
virtually
all
organic
molecules
have
a
3‐dimensional
shape
it
is
often
important
to
 be
able
to
convey
this.

The
conventions
for
this
are:
 (1)
Bonds
that
lie
in
the
plane
of
the
paper
are
indicated
by
a
simple
line.
 (2)
Bonds
that
come
forward
out
of
the
plane
of
the
paper
are
indicated
by
a
solid
 wedge.
 (3)
Bonds
that
go
back
out
of
the
plane
of
the
paper
are
indicated
by
a
dashed
wedge
 (4)
Generally
to
represent
a
tetrahedral
arrangement
of
atom
(a)
two
of
the
bonds
are
 drawn
in
the
plane
of
the
paper
about
109o
apart
and
(b)
the
other
two
bonds
are
 drawn
in
the
opposite
direction
to
the
in‐
plane
bonds

but
right
next
to
each
other.
 (5)
Trigonal
planar
arrangements
of
atoms
can
be
drawn
in
3‐dimensions
in
the
plane
 of
the
paper.

Bond
angles
should
be
approximately
120o.

These
can
also
be
drawn
 side‐on
with
the
central
bond
in
the
plane
of
the
paper,
one
bond
forward
and
one
 bond
back
shown
as
follows:
 (6)
Linear
arrangements
of
atoms
are
always
best
drawn
in
the
plane
of
the
paper.
 ...
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This note was uploaded on 06/19/2009 for the course CHEM 2311 taught by Professor Tyson during the Fall '07 term at Georgia Tech.

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