Gases
Chapter 12
Gases are compressible
fluids.
This behaviour arises because
the gas molecules are in
constant random motion
through mostly empty space.
Intermolecular Forces: Liquids and Solids
Dr. Peter Warburton
[email protected]
http://www.chem.mun.ca/zcourses/1050.php
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Liquids
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Solids
Liquids are incompressible
fluids.
This behaviour arises because
the molecules of the liquid are in
constant random motion
without much empty space to
move around in.
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Solids are incompressible and
rigid.
This behaviour arises because
the molecules of the solid can
only vibrate because they
have almost no empty space
to move into.
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Ideal gas law
All molecules have intermolecular forces (IMFs)
We’ve treated gases in the past as
behaving ideally, so that
PV = nRT
However, this gas law ASSUMES that
molecules:
1) Have NO SIZE (no repulsive
intermolecular forces)
2) DO NOT have attractive
intermolecular forces
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Repulsive IMFs (real molecules have
size) will limit the compressibility of a
group of molecules.
SQUEEZE!
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van der Waals gas law
All molecules have intermolecular forces (IMFs)
One equation of state used to describe real gas
behaviour is the van der Waals (vdW) equation
Attractive IMFs cannot be overcome at
sufficiently low temperatures (molecules
have low average kinetic energies).
௔௧௧௥௔௖௧௜௩௘
ܽ݊ଶ
ܲ+ ଶ
ܸ
Lower T!
1)
2)
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௥௘௣௨௟௦௜௩௘
ܸ − ܾ݊ = ܴ݊ܶ
The nb term represents the
molecular size (repulsive forces)
The an2/V2 term represents the
attractive van der Waals forces
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Intermolecular forces
Intermolecular forces
The forces between molecules are
electrostatic. They depend on charges
(like charges repel, opposite charges
attract), and the distance between the
charges.
Larger charges (like those found on ions)
and smaller distances between
molecules tend to lead to stronger forces.
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At everyday conditions, the forces
between molecules tend to be
weakly attractive overall.
These intermolecular forces are
generally called van der Waals
forces. However, vdW forces can be
subdivided into two different groups.
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Dipole-dipole forces
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Dipole-dipole forces
We’ve already seen that some
molecules have a permanent
molecular dipole.
Since full charges are not involved in
molecular dipoles, these dipole-dipole
intermolecular interactions are
relatively weak as compared to ionic
bonds, where full charges are
involved.
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The larger the dipole moment (the
molecules are more polar), the
stronger the IMFs tend to be.
••
H - Cl :
••
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Dipole-dipole forces
London (dispersion) forces
Nonpolar molecules still interact with
each other despite their lack of
permanent dipoles.
Here we have three molecules
with similar molar masses. We
see the polar molecule has a
much higher boiling point due to
dipole-dipole IMFs
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London (dispersion) forces
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London (dispersion) forces
When a molecule with a instantaneous dipole
comes close to another molecule, the electrons
of the second molecule will try to move away
from the negative partial charge of the first
molecule, leading to the second molecule
having a temporary induced dipole.
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In a nonpolar molecule, “on average,”
the electrons do not prefer one part of the
molecule over the other. However, at
any given instant, they might not be
evenly distributed and so the molecule
ends up having a “instantaneous
dipole.”
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London forces tend to be very
weak because the partial charges
tend to be small and fleeting.
However, ALL chemical species
have London forces between
them.
Variations in the strength of London
forces depend on two factors.
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London (dispersion) forces
Polarizability
Polarizability – is the ability for
electrons to move freely within the
molecule. The more freely electrons
can move, the larger the induced
dipole can be.
Larger molecules and atoms are
more polarizable, and generally
have larger London forces.
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London (dispersion) forces
Weaker
London
forces
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Hydrogen bonding
Shape – The shape of a molecule plays a part in
determining how the electrons can move in a
molecule. More compact shapes are usually
more symmetrical and allow less contact
between molecules. They generally have smaller
induced dipoles with weaker London forces.
Stronger
London
forces
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A hydrogen bond
is an attractive
interaction between
a hydrogen atom
bonded to a very
electronegative
atom (usually O, N,
and F), and an
unshared electron
pair on another
electronegative
atom.
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Hydrogen bonding
Hydrogen bonding
Hydrogen bonds are really just a very
special case of dipole-dipole forces.
H-F, H-O, and H-N bonds are very polar
(larger partial charges)
Also, because the hydrogen is very small,
it is possible for another molecule to
approach it very closely (short distance).
Hydrogen bonds are relatively strong
intermolecular forces
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Hydrogen bonding in water
Hydrogen bonding
This trend is generally true, except for NH3, H2O,
and HF because of the hydrogen bonds (stronger than
London forces) that can occur.
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With London forces, boiling points will increase with
molecular size (polarizability).
We EXPECT boiling points to follow the trend
CH4 < SiH4 < GeH4 < SnH4
and so on across the periodic table.
The hydrogen bonding in
ice creates an organized
crystal structure with
large gaps
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Liquid water has
enough energy to
overcome some of the
hydrogen bonds, and
the gaps “break down”
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Hydrogen bonding in water
Liquid water is most dense at
about 277 K, as the remaining
hydrogen bonds pack the
molecules into a smaller volume.
Hydrogen bond dimerization
Many molecules that
show hydrogen bonding
often form dimers, where
two of the molecules
hydrogen bond with each
other. This tends to
lower the enthalpy of
vaporization, since the
gas phase contains
many dimers where the
hydrogen bonds have
not been broken.
The net result is that
ice is less dense than
liquid water, and so it
floats (left beaker).
Most solids don’t have
the large gaps and
therefore the solid is
more dense than the
liquid. Parafin wax is
the example shown in
the right beaker.
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Hydrogen bonding and viscosity
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Inter vs intramolecular H bonding
Viscosity is a measure of a liquids resistance to
flow. More viscous liquids generally have
stronger IMFs, making it difficult for the
molecules to flow. Viscosity η is often measured
in units of centipoise, where 1 cP = 1 N s m-2
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We saw the structure
of ice, the formation of
dimers and the affect
of H bonding on
viscosity are the result
of intermolecular
hydrogen bonding
between different
molecules.
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Inter vs intramolecular H bonding
IMFs and bonding at a glance
However, it is
possible for one
part of a molecule
to form a hydrogen
bond with another
part of the SAME
molecule.
This is called
intramolecular
hydrogen
bonding.
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IMFs and bonding at a glance
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Problem
Which of the following substances would
you expect to have the highest boiling
point: C3H8, CO2 or CH3CN? Explain.
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Problem answer
Properties of liquids
CH3CN has dipole-dipole and London
IMFs, while the other two molecules are
non-polar and only show London IMFs, so
CH3CN likely has the highest boiling point
due to the stronger IMFs.
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Liquids are held together by the IMFs, so
the nature and strength of the IMFs affect
many properties of liquids, including
surface tension, viscosity,
enthalpy of vaporization, vapor
pressure, boiling points and
critical points
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Surface tension
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Surface tension
Surface tension arises from the differences in
IMFs the molecules experience at the surface
compared to the bulk liquid. It also tends to
decrease with increased T, since faster moving
molecules more easily overcome the IMFs.
Surface tension γ
(usually measured in J
m-2) is the work required
to increase the surface
area of a liquid.
Stronger IMFs would
require more work to
overcome them so we
can “spread out” the
surface.
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Manifestations of surface tension
Manifestations of surface tension
Surface tension is why things “get wet”. The
forces in a drop of liquid are called cohesive
forces while the forces at the surface are
adhesive forces. If adhesive forces are stronger,
the drop “spreads out” and wets a surface.
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Viscosity
Surface tension is also why we
see a meniscus when liquid is
in a narrow tube, and capillary
action where liquid is drawn up
a tube due to wetting
(adhesive forces). If adhesive
forces are stronger the
meniscus is concave due to
wetting. If cohesive forces are
stronger, the meniscus is
convex as the liquid tries to
form a “drop”.
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Enthalpy of vaporization
The enthalpy of vaporization is the amount
of heat required to transform one mole of a
liquid into a gas. In general, the stronger
the IMFs between molecules, the higher
the enthalpy of vaporization will be.
The opposite process to vaporization is
condensation, where the gas becomes a
liquid. Since enthalpy is a state function
∆Hvap = -∆Hcond
Viscosity is a measure of a
liquids resistance to flow.
More viscous liquids
generally have stronger
IMFs, making it difficult for
the molecules to flow. Since
faster moving molecules tend
to better overcome IMFs,
viscosity tends to decrease
with increasing temperature.
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Enthalpy of vaporization
Problem
How much heat is required to vaporize a
2.35 g sample of diethyl ether at 298K?
Dipole-dipole
Hydrogen bonding
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Problem answer
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Vapour pressure
Vapor pressure is the partial
pressure (the part of the total
pressure that comes from a given
substance) of the vapor (gas) above
the liquid (or solid) phase measured
at EQUILIBRIUM
at a GIVEN TEMPERATURE
0.923 kJ
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Equilibrium and vapour pressure
Equilibrium and vapour pressure
Vapor pressure should be measured
when the vapor pressure
has STOPPED CHANGING.
This equilibrium means that the rate of
molecules leaving the liquid (or solid)
phase is BALANCED EXACTLY by the
rate of the molecules in the vapour
joining the liquid (or solid) phase.
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Temperature and vapor pressure
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Temperature and vapor pressure
The vapor pressure depends on how
easily molecules can overcome
attractive intermolecular forces that
keep it in the liquid (or solid) phase.
Higher temperatures mean each
molecule, on average, has more
kinetic energy that COULD ALLOW
the molecule to “escape” from the
attractive forces.
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Temperature and vapor pressure
Temperature and vapor pressure
In this set of vapor
pressure curves,
the MOST
VOLATILE
(weakest IMFs)
liquid is on the left,
while the LEAST
VOLATILE
(strongest IMFs) is
on the right.
Since all groups of molecules AT THE
SAME TEMPERATURE have the
SAME AVERAGE KINETIC ENERGY,
then molecules that have WEAKER
intermolecular forces are MORE
VOLATILE (have GREATER vapour
pressures at the GIVEN temperature)
than NONVOLATILE molecules with
STRONGER intermolecular forces.
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Problem
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Problem answer
Equilibrium is established between liquid
hexane and its vapor at 298 K. A sample
of the vapor is found to have a density of
0.701g L-1. If C6H14 has a molar mass of
87.1766 g mol-1 then what is the vapor
pressure of hexane in Torr at this
temperature.
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151 Torr
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Clausius-Clapeyron equation
Temperature and vapor pressure
By looking at the
vapor pressure
curves, it seems
obvious there is
some sort of
exponential
relationship
between vapor
pressure and
temperature!
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The exponential relationship between
vapor pressure and temperature can also
be explored in a logarithmic (inverse of
exponential) manner. The relationship is
−∆‫ܪ‬௩௔௣ 1
ܲଶ
1
݈݊
=
−
ܶଶ ܶଵ
ܲଵ
ܴ
This is the Clausius-Clapeyron equation!
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Clausius-Clapeyron equation
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Problem
When we plot the
natural logarithm of
vapor pressure as a
function of inverse
temperature, we get a
straight line!
Comparison of any two
points on this straight
−∆‫ܪ‬௩௔௣ 1
ܲଶ
1
݈݊
=
−
line lead to the
ܴ
ܲଵ
ܶଶ ܶଵ
Clausius-Clapeyron
equation.
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A handbook lists the vapor pressure of
methanol as 100 mmHg at 21.2 ºC and we
saw on slide 41 its enthalpy of vaporization
is 38.0 kJ mol-1. What’s the vapor
pressure at 25.0 ºC.
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Problem answer
Boiling point
The boiling point of a
liquid is the temperature
where the vapor
pressure IS THE SAME
AS the external
pressure.
Therefore, if the external
pressure changes, the
boiling point
temperature ALSO
changes.
121 mmHg
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Normal boiling point
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Critical point
The normal boiling
point of a liquid is the
temperature where
the vapor pressure
IS THE SAME AS an
external pressure of
exactly
1 ATMOSPHERE.
The density of a liquid tends to DECREASE
with temperature, as the molecules move
farther apart on average (same mass in a
larger volume) because the IMFs are less
successful at holding them together.
However, since vapor pressure INCREASES
with temperature, the density of the vapor in a
closed container will increase with more
molecules in the vapor.
1 atm = 760 mmHg
= 101.325 kPa
= 1.01325 bar
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Critical point
Critical point
If we are at the critical point temperature, we can
turn a gas into a liquid by increasing the external
pressure to the critical pressure.
If we are below the critical temperature, then
condensation occurs at some pressure below
the critical pressure.
If we are above the critical temperature, we
cannot condense the gas to a liquid with ANY
amount of pressure. The molecules move too
fast for the IMFs “to take hold”.
At some set of conditions, the density of the
liquid and the density of the vapor become
the same and the surface tension becomes
zero. Something “very strange” happens at
this critical point: we can’t tell the liquid and
vapor apart from each other any more!
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Critical point
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Problem
Arrange the following in the expected order
of increasing boiling point: Ne, He, Cl2,
(CH3)2CO, O2, O3.
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Problem answer
He < Ne < O2 < O3 < Cl2 < (CH3)2CO
O3 comes before Cl2 even though its
slightly polar and Cl2 is nonpolar because
Cl2 is much heavier (more polarizable) and
so it has London forces that become
stronger than the slight dipole-dipole
forces in ozone.
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