Chapter 13: Intermolecular Forces, Liquids,
and Solids
The Kinetic-Molecular Theory of Matter
Recall the kinetic-molecular theory of gases (Chapter 12).
Kinetic-molecular theory applies to solids and liquids too:
1. Molecules are in constant motion in random directions.
2. Molecules move at different speeds, but the average speed
is proportional to temperature. All molecules have the
same average kinetic energy at a given temperature.
In gases, the distances between molecules are much larger than
the size of the molecules themselves. As such, most of the
motion of gas molecules is ______________________ (moving
from one location to another).
In solids, molecules are packed in a three-dimensional solid
lattice. They cannot easily change location, so their kinetic
energy is primarily due to ___________________ motion. This
motion is averaged about the molecule’s location in the lattice.
In liquids, molecules are much closer together than in gases, but
they are still able to change locations (unlike in solids). As
such, liquid molecules undergo significant amounts of
translational, rotational and vibrational motion.
Intermolecular Forces
The kinetic energy of molecules tends to keep them as separated
and disorganized as possible, so there must be some opposing
force(s) pulling them together – particularly in solids and
liquids. These attractions are intermolecular forces.
Intermolecular forces are similar to, but weaker than, those
involved in bonding. There are four kinds of intermolecular
forces, all of which are based on _________________________.
The strength of an intermolecular force is determined by the
strength and permanence of the dipoles involved (and by the
distance between the dipoles).
Type of Intermolecular Force
Approximate
Strength
(kJ/mol)
40-600
Ion-Dipole
+
-
+
5-25
Dipole-Dipole (includes hydrogen bonding)
-
+
-
+
2-10
Dipole-Induced Dipole
-
+
-
+
0.05-40
Induced Dipole-Induced Dipole
(aka London dispersion forces)
-
+
-
+
Hydrogen bonding is a special case of dipole-dipole attraction.
This is due to the very large dipole moment that occurs when a
hydrogen atom is bonded to a very small, very electronegative
atom (i.e. F, O or N). The large dipole moment results in a
particularly strong dipole-dipole attraction.
In hydrogen
bonding, both the atom bonded to H and the atom attracted
to H must be either F, O or N. They do not, however, have to
be the same element.
e.g.
Induced dipoles occur when two molecules approach closely
enough that the electron cloud of one is repelled by the electron
cloud of the other. The electrons move to the far end of the
molecule, setting up a temporary dipole.
These induced dipoles attract each other (and one induced
dipole can induce another to form).
Large atoms are most susceptible to forming induced dipoles.
(most polarizable) Why?
Intermolecular forces between neutral molecules are referred to
as a group as van der Waals forces.
Properties of Liquids
Liquids are much denser than gases. As in solids, there is
virtually no free volume between molecules in a liquid. As
such, they are virtually incompressible – a property used to
advantage in hydraulics.
The molecules in a liquid are in constant motion – translational,
vibrational and rotational. As in gases, the motion of any one
molecule is random, but the average speed is proportional to
temperature. The statistical picture for the energy distribution in
liquids is identical to that for ideal gases:
Enthalpy of Vaporization
When you leave wet dishes on a draining board overnight, they
will usually be dry the next morning. The water has evaporated
– even though your kitchen was 70-80˚C below the boiling point
of water! How is this possible?
At any given temperature, some of the liquid molecules will
have enough energy to escape the intermolecular forces holding
the molecules together.
Heating the liquid increases the rate of vaporization by
increasing the proportion of molecules having at least this much
energy. Since heating increases the rate of vaporization, this is
an _________________ process with a ___________ enthalpy.
This is consistent with the observation that rapid vaporization
cools any remaining liquid. The molecules with enough energy
to evaporate do so, reducing the average kinetic energy of the
remaining liquid molecules (thereby reducing the liquid’s
temperature).
During slow evaporation, heat is steadily
resupplied from the environment so we don’t notice this loss.
The enthalpy of vaporization is defined as the energy required
to vaporize a liquid.
Its opposite, the enthalpy of
condensation, is of equal magnitude but opposite sign.
e.g. H2O(l) → H2O(g)
∆H˚vap = 40.7 kJ/mol
H2O(g) → H2O(l)
∆H˚cond = –40.7 kJ/mol
Vapour Pressure
Because there will always be at least a few molecules in a liquid
with enough energy to evaporate, there will always be a layer of
gas molecules immediately above a liquid. The pressure exerted
by these gas molecules is referred to as vapour pressure.
When a liquid is placed in a sealed container, it will eventually
reach an equilibrium at which the rate of evaporation is the same
as the rate of condensation. In this situation, an equilibrium
vapour pressure (often shortened to “vapour pressure”) can be
measured. This is a physical property specific to a given liquid
at a given temperature, and it increases exponentially with
temperature.
When the vapour pressure
of a liquid is equal to the
atmospheric pressure, the
liquid begins to boil. The
pressure exerted by gas
molecules is equal to the
pressure of the atmosphere
pushing down on the
liquid’s surface. As such, bubbles of gas can form within the
liquid (not just at its surface), they rise, and we observe boiling.
The temperature at which a liquid boils is referred to as its
boiling point; however, this property is pressure-dependent.
Reducing the pressure will decrease a liquid’s observed boiling
point while increasing pressure will increase the observed
boiling point. To allow us to compare boiling points of different
compounds, we therefore define the temperature at which the
vapour pressure of a liquid is equal to _______ as a liquid’s
normal boiling point.
Liquids with a high vapour pressure/low boiling point are
referred to as __________________.
e.g. Looking at the vapour pressure curves above,
(a) What is the approximate vapour pressure of ethanol at
room temperature (20˚C)?
(b) Are liquid ethanol and its vapour in equilibrium when the
temperature is 60˚C and the vapour pressure is 600 mmHg?
If not, does liquid evaporate to form more vapour, or does
vapour condense to form more liquid.
Plotting vapour pressure vs. temperature gives an exponential
curve (see previous page). Plotting the natural logarithm of the
vapour pressure versus the reciprocal of temperature gives a
linear graph. The equation of the line is the Clausius-Clapeyron
∆Hovap
equation:
ln P = + C
RT
This equation is more useful in its comparative form (i.e. the
equation for calculating the slope of the linear graph):
∆Hovap
P1
ln
= R
P2
1
1
T1 T2
Where P1 and P2 are vapour pressures (in the same units as each
other), ∆H˚vap is the enthalpy of vaporization (in J/mol), R is the
ideal gas constant (in J · mol-1 K-1) and T1 and T2 are
temperatures (in K).
e.g. If diethyl ether has a vapour pressure of 534 mmHg at
25.00˚C and a vapour pressure of 57.0 mmHg at -22.80˚C,
what is its enthalpy of vaporization?
Critical Temperature and Pressure
As the temperature of a liquid is increased, its vapour pressure
continues to increase, and it is necessary to have an increasingly
high atmospheric pressure in order to maintain the liquid state.
There comes a temperature at
which no amount of pressure is
able to compress gas to liquid; the
energy of the molecules is just too
high.
This is the critical
temperature. The point on a
phase diagram at which the
critical temperature and critical
pressure meet is referred to as the
critical point.
At temperatures higher than the critical temperature, high
pressures give dense gases called supercritical fluids. The
molecules in supercritical fluids are packed together almost as
tightly as in liquids but they have enough energy to overcome
intermolecular forces and move freely. As such, they can make
excellent solvents (especially supercritical CO2 – which has the
added advantage of being much more environmentally friendly
than most organic solvents).
When is it important to know a substance’s critical point?
e.g. Choosing a refrigerant. Refrigerators work by compressing
a gas to liquid. As the liquid is released into areas of lower
pressure, it evaporates, consuming energy and cooling the
fridge. If a potential refrigerant has a very low critical
temperature, this process will not work because it will be
impossible to condense the gas to a liquid. Freon-12, a
common refrigerant, has a critical temperature of ~112˚C.
Surface Tension
Molecules within a liquid are attracted to all of the molecules
around them via intermolecular forces. Molecules on the
surface of a liquid can only interact with the molecules
immediately below them and beside them. As such, the surface
molecules feel a net attraction toward the interior of the liquid
and act as a “skin” for the liquid. The energy required to break
through this surface “skin” is referred to as surface tension.
H
H
O
H
O
H
H
H
O
H
O
H
H
H
O
O
H
H
H
H
O
smooth surface;
lots of intermolecular attractions
H
H
O
H
O
H
H
H
O
O
H
H
O
H
H H
H
O
H
O
H
surface tension broken;
intermolecular attractions interfered with
Some insects take advantage of surface tension to literally walk
on water!
Surface tension is also the reason why liquids tend to bead –
spheres have the smallest surface area for a given volume.
Capillary Action
Capillary action is due to the same
intermolecular forces as surface tension;
however, the attractions are between
molecules of a polar liquid (often water)
and a polar solid (often glass). Glass
consists of polar Si-O bonds so, if a narrow
glass tube is placed in water, the water
molecules are attracted to the glass. Other
water molecules are attracted to the
initially “stuck” molecules, and are pulled
H
H
O
H
O
O
H O
O
H
O
O
Si
O
Si
H
O
O
H
O
Si
H
H
O O
O
H
H H
H
O
H
simplified structure for glass
(it's much more 3-dimensional!)
up by surface tension. This continues until the force of gravity
pulling water molecules down is equal to the force of surface
tension pulling water molecules up.
Capillary action is responsible for the meniscus typically
observed when using pipettes, burettes, etc..
Viscosity
Viscosity is defined as the resistance of liquids to flow. Liquids
with high viscosities tend to be thick and difficult to pour. This
is because the strong intermolecular forces in viscous liquids
cause the molecules to “stick together”. Viscous liquids tend to
fall into one of two categories (or both):
1. large molecules
2. polar molecules
Summary Exercises
1.
butane
(CH3CH2CH2CH3)
methanol
(CH3OH)
helium
(He)
(a) Rank the pure substances above in order of increasing
strength of intermolecular forces:
(b) Which of these pure substances would you expect to be
gases at 25˚C and 1 atm?
2. The graph below shows vapour pressure curves for carbon
disulfide (CS2) and nitromethane (CH3NO2).
(a) What type of intermolecular forces exist in the liquid phase
of each compound?
(b) What are the vapour pressures of CS2 and CH3NO2 at 40˚C?
(c) What are the normal boiling points of CS2 and CH3NO2?
(d) At what temperature does CS2 have a vapour pressure of 600
mmHg?
(e) At what temperature does CH3NO2 have a vapour pressure
of 60 mmHg?
Properties of Solids
As in liquids, the molecules in solids are packed tightly together.
Unlike in liquids, the molecules in solids do not experience
translational motion – only vibrational. Solid molecules are held
tightly in place by strong intermolecular forces.
There are two main categories of solids:
• Crystalline solids, in which the molecules are ordered, and
• Amorphous solids, in which the molecules are disordered.
Crystalline Solids
Crystalline solids are made up of repeating units built upon each
other like bricks in a wall. An individual “brick” is called the
_____________.
Mathematicians have determined that there are seven possible
shapes for a unit cell (aka seven crystal systems). These include
cubic unit cells and hexagonal unit cells:
Crystalline solids can be analyzed by x-ray crystallography, in
which an x-ray is passed through a crystal. The crystal acts as a
diffraction grating (the x-rays can pass through gaps in the
crystal structure but not through the atoms themselves), and
analysis of the resulting diffraction pattern allows a chemist to
determine the structure of the crystal.
There are four different kinds of crystalline solids, each
distinguished by the type of chemical bonding involved:
• Ionic solids: positive and negative ions held together by
electrostatic attractions
• Metallic solids: metal atoms held together by electrostatic
attractions between nuclei and a ‘sea’ of electrons
• Molecular solids: molecules held together by van der Waals
attractions
• Network solids: covalently bonded networks of atoms
For a more detailed look at each type of solid, see (and learn!)
Table 13.6 in Kotz.
Amorphous Solids
The molecules in amorphous solids have no long-range order
and can be thought of as frozen liquids. The most widely
known amorphous solid is glass.
Interestingly enough,
crystallization results in the failure of glass.
Melting Points and Enthalpies of Fusion
Just like liquids, the physical properties of solids can be
explained using intermolecular forces. Converting a solid to a
liquid requires input of enough energy to overcome the
intermolecular forces holding each atom/molecule in the solid in
place. This process can be called either melting or fusion.
Thus, melting points (or fusion points) are higher for solids
with stronger intermolecular forces.
In general, ionic solids and transition metals have the highest
melting points; polar molecular solids have intermediate melting
points; and nonpolar molecular solids have low melting points.
The enthalpy of fusion is defined as the energy required to
liquefy a solid. Its opposite, the enthalpy of crystallization, is
of equal magnitude but opposite sign.
e.g. H2O(s) → H2O(l)
∆H˚fus = 6.02 kJ/mol
H2O(l) → H2O(s)
∆H˚cryst = –6.02 kJ/mol
Solids with high enthalpies of fusion have high melting points
(and vice versa).
Sublimation
Some solids convert directly from the solid to the gas phase
without passing through the liquid stage. This is called
sublimation. The most familiar examples are probably carbon
dioxide (“dry ice”) and iodine.
The enthalpy of sublimation is defined as the energy required
to sublime a solid.
Phase Diagrams
The equilibrium vapour pressure curves we looked at earlier this
chapter can be extended to describe the complete range of states
of matter. This is a phase diagram, and each state of matter is
referred to as a ________________.
The phase diagram to the right is that
for carbon dioxide (CO2). Note that
the slope of the solid-liquid boundary
line is positive, indicating that the
liquid is less dense than the solid.
Also, note the point where the three
phases (gas, liquid and solid) meet.
This is the _________________.
Below the triple point, it is not possible for a substance to exist
in the liquid phase. For CO2, this means that liquid CO2 cannot
exist at pressures below 5 atm or temperatures below -57˚C.
According to its phase diagram, the critical temperature for CO2
is _________ and the critical pressure is _________.
The phase diagram to the
right is that for water. Note
the unusual negative slope
of the solid-liquid boundary
line, indicating that the solid
is less dense than the liquid.
This unusual property is due
to the strongly hydrogen
bonded structure of the solid
phase of water.
This relatively low density of ice (compared to liquid water) is
what allows fish to survive in partially frozen lakes and rivers
over the winter. If ice was more dense than water, lakes would
freeze from the bottom up, and there would be no liquid water
left for the fish. Because ice is less dense than water, lakes
essentially freeze from the top down, so there is often a layer of
liquid water left at the bottom.
Important Concepts from Chapter 13
• kinetic-molecular theory of matter
• intermolecular forces (four categories, hydrogen bonding)
• properties of liquids
o vapour pressure
o boiling vs. evaporation
o enthalpy of vaporization (Clausius-Clapeyron equation)
o surface tension
o capillary action
o viscosity
• properties of solids
o five different types of solids (see Table 13.6 in Kotz)
o enthalpy of fusion
o enthalpy of sublimation
• phase diagrams
o predicting changes of state under different conditions
o critical point (and supercritical fluids)
o triple point