Chap11-Part2

Chap11-Part2 - Combined
Gas
Law:
...

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Unformatted text preview: Combined
Gas
Law:
 ‐ It's
the
Gas
law
that
combines
all
the
other
gas
laws:
Boyles
law
 
Charles
Law
and
Gay‐Lussac's
law
 PV=
nRT
 PV/T
=
nR
 PV/T
=
k
(for
a
fixed
amount
of
gas)
 
 
 
 
 
 
P1V1/T1
=
P2V2/T2
 Review
Session
Problem
 Solid
potassium
chlorate
(KClO3)
decomposes
to
produce
solid
 potassium
chloride
(KCl)
and
oxygen
gas
(O2)
according
to
the
balanced
 chemical
equaPon:
2KClO3
(s)

2KCl
(s)
+
3O2
(g)
 What
volume
of
gas,
measure
at
23.4oC
and
75.4
kPa,
will
be
produced
 when
32.5
g
of
potassium
chlorate
is
decomposed?
 Given,
 P
=
75.4kPa


(75.4kPa)
(1
atm)/
(101.325kPa)
=
0.745
atm
 T
=
23.4oC
=
273.1+23.4
=
296.5
K
 32.5g
of
KClO3
decomposed

?
g
of
oxygen
produced
 2
moles
of
KClO3
produce
3
moles
of
Oxygen
 2(122.55g)
of
KClO3
produce
3(32g)
of
oxygen
 32.5g
of
KClO3
produce
(3X32)
(32.5)/
(2X
122.55)
=
12.73g
of
Oxygen
 n
=
12.73/32
=
0.398
mol
 V=
(0.398)
(0.08206)
(296.5)
/
0.745
=
13
L
 1)  Ion‐dipole
 2)  dipole
–
dipole
 3)  H‐bonding
 4)  London
dispersion
interacPon
 The
strengths
of
intermolecular
interacPons
are
almost
always
less
than
the

 strengths
of
covalent
bonds.
For
example,
less
than
16
kJ/mol
is
needed
to

 overcome
the
intermolecular
acracPons
between
HCl
molecules,
but
431
kJ/mol
is

 needed
to
break
the
covalent
bond.

When
a
substance
undergoes

 a
change
from
solid
to
liquid
to
gas,
the
covalent
bonds
usually
remain
intact.


 Boiling
and
melPng
points
are
largely
determined
by
the
strength
of

 intermolecular
interacPons.

Above,
the
covalent
bond
is
represented
as
a
 solid
line
whereas
the
intermolecular
interacPon
is
represented
as
a
dashed
line.
 Three
types
of
non‐covalent
intermolecular
interacPons
are
known
to
exist

 between
neutral
molecules:
 1)
dipole‐dipole
interacPons
 2)
London
dispersion
interacPons
 3)
hydrogen‐bonding
 The
three
types
of
intermolecular
acracPons
are
called
as
van
der
 Waals
interacPons.
 There
is
a
fourth
known
interacPon
important
in
soluPons,
called
the
 4)
ion‐dipole
interacPon
 Other
non‐covalent
interacPons
have
also
been
idenPfied
and
characterized,
 ofen
in
the
context
of
biochemistry,
but
these
interacPons
will
not
be

 discussed
here.
 1)
Dipole‐Dipole
Interac>ons
 Even
though
a
molecule
may
not
have
a
net
charge,
the
molecule

 may
have
a
“permanent”
dipole
(one
end
posiPve
and
the
other

 negaPve).


In
this
case,
there
exists
the
possibility
of
molecular

 acracPon
if
the
two
molecules
arrange
themselves
as
follows:
 +
 ‐
 ‐
 +
 Molecules
do
not
need
to
be
charged
to
acract
each
other.

They
 only
have
to
be
“polarized”.
 1)
Dipole‐Dipole
Interac>ons
 There
are
both
acracPve
and
 repulsive
interacPons.
 Two
molecules
acracPng
each
 other
spend
more
Pme
together
 than
two
molecules
repelling
 each
other.
The
overall
effect

 is
a
net
acracPon.


 For
molecules
of
approximately
 equal
mass
and
size,
the
strengths
 of
intermolecular
interacPons
 increase
with
increasing
polarity.
 Even
though
the
molecular
weight
and
size
are
about
equal,
 the
boiling
point
increases
with
increasingdipole
moment.
 This
result
implies
that
the
intermolecular
interacPons
are
 stronger
for
molecules
with
higher
dipole
moments.
 
The
more
polar
the
molecule,
the
higher
is
its
boiling
point.
 2)
London
Dispersion
Interac>ons
 Even
though
some
molecules
are
nonpolar
on
the
average,
they
can
sPll
 have
temporary
asymmetries
in
their
electron
distribuPons
and
therefore
 “instantaneous”
dipole
moments.

Two
such
atoms
can
then
experience

 acracPve
interacPons
as
for
molecules
with
permanent
dipoles.
 
The
temporary
dipole
on
one
atom
can
induce
a
temporary
dipole
on

 the
adjacent
atom,
causing
the
forces
of
acracPon.
This
interacPon
called
 as
‘London
dispersion
force’
 2)
London
Dispersion
Interac>ons
 4 He
 
 2
 
While
the
electrons
in
the
1s
orbital
of
helium
would
repel
 each
other
(and,
therefore,
tend
to
stay
far
away
from
 each
other),
it
does
happen
that
they
occasionally
wind
up
 on
the
same
side
of
the
atom.
 2)
London
Dispersion
Interac>ons
 
At
that
instant,
then,
the
helium
atom
is
polar,
with
an
 excess
of
electrons
on
the
lef
side
and
a
shortage
on
the
 right
side.
 2)
London
Dispersion
Interac>ons
 
Another
helium
nearby,
then,
would
have
a
dipole
 induced
in
it,
as
the
electrons
on
the
lef
side
of
helium
 atom
2
repel
the
electrons
in
the
cloud
on
helium
atom
1.
 2)
London
Dispersion
Interac>ons
 
London
dispersion
forces,
or
dispersion
forces,
are
 acracPons
between
an
instantaneous
dipole
and
an
 induced
dipole.
 2)
London
Dispersion
Interac>ons
 •  These
forces
are
present
in
all
molecules,
whether
they
 are
polar
or
nonpolar.
 •  The
tendency
of
an
electron
cloud
to
distort
in
this
way
is
 called
polarizability.
 2)
London
Dispersion
Interac>ons
 Dispersion
forces
depend
on
the
molecular
polarizability,
which
describes
 how
easily
the
electron
cloud
can
be
rearranged.


 More
polarizable
molecules
have
stronger
dispersion
forces.


 Polarizabilites
tend
to
increase
with
molecular
size
because
there
are
more

 electrons
and
they
are
farther
from
the
nuclei.
 2)
London
Dispersion
Interac>ons
 Molecular
shapes
also
influence
the
strengths
of
dispersion
forces.
 C5H12
 CH3 CH3 H3C CH3 H3C CH3 N‐pentane
has
stronger
dispersion
forces
and
a
higher
boiling
point
 because
its
more
linear
shape
allows
for
greater
intermolecular
interacPons.
 Which
Have
a
Greater
Effect?
 Dipole‐Dipole
InteracPons
or
Dispersion
Forces
 •  All
molecules
experience
dispersion
forces,
whether
they
are
 polar
or
not.

For
polar
molecules,
the
strength
of
the
 dispersion
interacPons
is
usually
much
greater
than
the
 strength
of
the
dipole‐dipole
interacPons.
 •  If
two
molecules
are
of
comparable
size
and
shape,
dipole‐ dipole
interacPons
will
likely
the
dominaPng
force.
 •  If
one
molecule
is
much
larger
than
another,
dispersion
forces
 will
likely
determine
its
physical
properPes.
 Effect
of
London
Dispersion
Forces
&
Polarizability
 on
Physical
Proper>es
 Fritz
London
(1900‐1954)
 ‐ TheorePcal
Physicist
 ‐ 
London
Dispersion
forces
 ‐ 
ElectromagnePc
properPes
of
Superconductors
 ‐ His
studies
supported
Superfluidity
concept


 (images
from
Google)
 3)
Hydrogen
Bonding
 Boiling
points
generally
increase

 with
molecular
size
because
of

 the
higher
London
dispersion

 forces.
 Water
deviates
considerably
from

 this
trend.
 The
deviaPon
is
explained
by
the

 high
propensity
of
water
to

 hydrogen
bond
with
itself.
 Hydrogen
bonds
are
stronger
 than
dipole‐dipole
or
dispersion
 forces
but
weaker
than
covalent
 bonds.

 •  The
nonpolar
series
(SnH4
to
CH4)
follow
the
expected
trend.
 •  The
polar
series
follows
the
trend
from
H2Te
through
H2S,
but
water
is
quite
an
 anomaly.
 3)
Hydrogen
Bonding
 Hydrogen
bonding
is
a
special
type
of
dipole‐dipole

 interacPon.
 These
bonds
occur
between
the
hydrogen
atom
in
 a

 polar
bond
(H‐F,
H‐O,
H‐N,
etc.)
and
an
unshared

 electron
pair
on
a
nearby
electronegaPve
atom

 (F,
O,
N,
etc.).
 The
posiPve
side
of
the
bond
dipole
consists
of
the

 nearly
bare
proton
of
the
hydrogen
nucleus.
The

 negaPve
side
of
the
bond
dipole
consists
of
the

 electronegaPve
atom
in
the
other
molecule.
 Because
hydrogen
is
small,
it
can
come
close
to
its

 molecular
partner
and
increase
the
interacPon
 strength.
 Water
&
Ammonia:

Solid
lines
show
 covalent
bonds
and
dashed
lines
 show
hydrogen
bonds.
 3)
Hydrogen
Bonding
 Unlike
for
most
substances,
solid
water
is
less
dense

 than
liquid
water.

This
phenomenon
can
be
explained

 by
hydrogen
bonding.
 In
ice,
the
molecules
maximize
their
hydrogen
bonding,
 leading
to
an
open,
ordered
arrangement.
In
water,
 the
arrangement
collapses,
although
hydrogen
bonding
 is
sPll
very
prominent.

 Effect
of
Hydrogen
Bonding
on
Physical
Proper>es
 Importance
of
H‐Bonding:
 ‐ Stabilizing
the
structure
of
Proteins
 ‐ Holds
the
double
stranded
DNA

 ‐ Carrying
GenePc
InformaPon
with
DNA
 ‐ Drug
delivery
systems
 Books
on
Hydrogen
Bonding
 4)
Ion‐Dipole
Forces
 These
forces
exist
between
a
charged
ion
and
the
parPal
charge
on
the
end
 of
a
polar
molecule.

They
are
parPcularly
important
for
ions
dissolved
in
 polar
liquids;
e.g.,
salt
water.
 4)
Ion‐Dipole
Forces
 Iden>fying
Intermolecular
Forces
 London
dispersion
forces
are
present
in
all
cases.

The
strengths
of
the
 other
interacPons
generally
increase
from
lef
to
right.
 What
kind
of
intermolecular
acracPve
force
is
illustrated
in
the
following
examples?
 What
kind
of
intermolecular
acracPve
force
is
illustrated
in
the
following
examples?
 hydrogen
bond
 What
kind
of
intermolecular
acracPve
force
is
illustrated
in
the
following
examples?
 hydrogen
bond
 London
dispersion
 What
kind
of
intermolecular
acracPve
force
is
illustrated
in
the
following
examples?
 hydrogen
bond
 London
dispersion
 ion‐dipole
 What
kind
of
intermolecular
acracPve
force
is
illustrated
in
the
following
examples?
 hydrogen
bond
 London
dispersion
 ion‐dipole
 dipole‐dipole
 Which
member
of
the
following
pairs
has
stronger
dispersion
forces?
 (a)
CH3CH2SH
or
CH3CH2CH2SH
 (b)
CH3CH2CH2Cl
or
(CH3)2CHCl?
 Which
member
of
the
following
pairs
has
stronger
dispersion
forces?
 (a)
CH3CH2SH
or
CH3CH2CH2SH
 (b)
CH3CH2CH2Cl
or
(CH3)2CHCl?
 (larger)
 Which
member
of
the
following
pairs
has
stronger
dispersion
forces?
 (a)
CH3CH2SH
or
CH3CH2CH2SH
 (larger)
 (b)
CH3CH2CH2Cl
or
(CH3)2CHCl?
 (less
spherical)
 substance 
 
 























structure 
 
 
boiling
point
 ethylene
glycol 
 
 















HOCH2CH2OH
 






198oC
 ethylene
glycol
dimethyl
ether







CH3OCH2CH2OCH3

 
83oC
 ethyl
alcohol
 
 ethyl
methyl
ether 























CH3CH2OH 
 
 















CH3CH2OCH3 
 








78oC
 








11oC
 Explain
why
replacing
the
hydrogen
atoms
on
the
oxygen
atoms
with
CH3
 results
in
a
lower
boiling
point.
 substance 
 
 























structure 
 
 
boiling
point
 ethylene
glycol 
 
 















HOCH2CH2OH
 






198oC
 ethylene
glycol
dimethyl
ether







CH3OCH2CH2OCH3

 
83oC
 ethyl
alcohol
 
 ethyl
methyl
ether 























CH3CH2OH 
 
 















CH3CH2OCH3 
 








78oC
 








11oC
 Explain
why
replacing
the
hydrogen
atoms
on
the
oxygen
atoms
with
CH3
 results
in
a
lower
boiling
point.
 The
effect
cannot
be
one
of
molecular
size
leading
to
higher
dispersion
forces,
 which
would
have
the
opposite
effect.


 The
effect
is
probably
due
to
decreased
hydrogen
bonding.
 ...
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