AP Chemistry Notes Chapter 10

Liquids, Solids and Intermolecular Forces

Any substance can exist as a solid, liquid or gas under the proper conditions.

The phase in which matter exists at standard conditions depends on intermolecular

forces of attraction or the attractive forces between individual particles of the substance.

Don’t be fooled into thinking that boiling points are “high” or freezing points are “low”
based on our familiarity with water, which is really one of a few substances that undergo
both those phase changes under “normal” conditions. Elemental nitrogen boils
vigorously at –196°C, and elemental iron freezes at nearly 1500°C.

It is assumed that you already know how solids, liquids and gases compare with regard to
definite shape and volume. Together, liquids and gases are known as fluids.

There are several important intermolecular forces. Some are stronger, others weaker, and
with a little knowledge and common sense, one can usually determine what forces are at
work between particles by making some simple observations about what phase a
substance is in at a particular set of conditions and what its freezing and boiling
temperatures are.

Categories of intermolecular forces:

1. Ion-ion attractions: Electrostatic attractions are very strong and virtually all ionic
compounds are solids at STP. Such compounds are collectively called salts.
2. Ion-dipole: Some molecules (of which water is the classic and most important
example) contain a permanent dipole, as discussed in Chapter 9. Such molecules
are strongly attracted to ionic charges, which is why water dissolves so many
ionic compounds well.
3. Dipole-dipole: For molecules with a permanent dipole, the center of positive
charge of one molecule is attracted to the center of negative charge of another.
Such attractions are not particularly strong unless the molecules are fairly close
together.
4. Hydrogen bonding: In molecules where H is bonded to O, N or F, there is a high
degree of polarization, resulting in particularly highly charged positive and
negative poles. In such molecules, the attraction between molecules is stronger
than simple dipole-dipole attractions. Hydrogen bonding explains many unusual
phenomena that will be discussed shortly.
5. London dispersion forces: Sometimes called van der Waal’s forces (although
technically, van der Waal’s forces also include dipole-dipole and hydrogen
bonding), these are the weakest of the intermolecular forces and exist between all
molecules, even non-polar ones. They are the result of induced dipolar charges.
Draw a diagram below to illustrate this.
As the molecular masses of the particles increase, London dispersion forces become
stronger. The more electrons there are, the greater the induced dipole. So we see that for
example, CH4 (molar mass 16g) is a gas at STP, but CCl4 (molar mass 152g) is a liquid.

The more attraction there is between particles, the higher the melting and boiling points
will be.

The different bonds range in strength as follows:

Ion-ion > ion-dipole > hydrogen bonding > dipole-dipole > London dispersion

Together, dipole-dipole, hydrogen bonding and London dispersion forces are known as
van der Waal’s forces.

Hydrogen Bonding

During hydrogen bonding, the more electronegative element (O, N, or F) pulls the shared
electrons strongly to itself, leaving a high partial positive charge on the hydrogen and a
high partial negative charge on the other atom.

Even in large molecules where there are –OH or –NH2 groups found, there will be
polarity, even if the main molecule itself is mainly non-polar.

Example: cyclohexane vs. glucose

This unusually high degree of dipole-dipole attraction explains some interesting things
about water in particular.

1. Capillary action:

2. Surface tension:

3. Unusually high boiling and freezing temperatures:

Because of its polarity and bond angle (104.5º), water also has the unusual characteristic
of expanding when it freezes, leaving the solid less dense than the liquid phase. This is
extremely important to life on earth for several reasons.

Other compounds that also demonstrate H-bonding between molecules dissolve VERY
well in water. Examples include NH3 and HF or HCl.

Vapor Pressure and Changes of State

Liquids can evaporate or change to gas at temperatures below the boiling point. The
greater the intermolecular forces, the lower the vapor pressure will be. Vapor pressure
(VP) also depends on temperature. As temperature increases, more molecules have the
kinetic energy to overcome the intermolecular forces and escape to the vapor phase. The
relationship between vapor pressure and temperature can be represented graphically and
mathematically by this simple formula:

Ln(Pvap) = -ΔHvap/R (1/T) + C where y = ln(Pvap), x = 1/T (in Kelvins), m = slope = -

ΔHvap/R. and b = y intercept = C. C is a constant particular to the liquid in question.

Probably a more useful version of this equation involves 2 pressures and 2 temperatures.
It allows a scientist to calculate the new vapor pressure at a new temperature, or the
temperature needed to achieve a certain vapor pressure (such as calculating a boiling
point at a different pressure or elevation.) The equation looks like this:

Ln PvapT1 = ΔHvap 1 - 1
PvapT2 T2 T1

Phase Changes

You should know the names for all possible phase changes between solid, liquid and gas
phases.

Heat of Fusion: ΔHfus is the heat required to melt one mole of a solid. Melting is always
endothermic, leaving ΔHfus positive. Conversely, freezing is exothermic and ΔHfus is
negative.

Heat of Vaporization: ΔHvap is the heat required to vaporize one mole of liquid to gas.
Vaporization is endothermic, and condensing is exothermic. This is one of the reasons
that steam burns can be so severe.

Specific Heat Capacity: c is the heat required to change the temperature of 1 gram of a
substance by 1ºC. It is different for different substances, and is even different for
different phases of the same substance. For example, c for water is 1.0 cal/gºC but for
steam and ice, c is only about 0.5cal/gºC.

To calculate heat energy change within a given phase for a substance ΔH = mcΔT.

A Heating Curve (or Cooling Curve) has this basic profile. Fill in the missing
information.
Triple Point Diagrams tell us a lot about a substance’s phase change conditions.

Triple point diagrams look like this. Fill in the missing information.

Two important features of the diagram are the triple point and the critical point.

Critical Temperature is the highest temperature at which a substance can exist as a

liquid. Beyond that temperature, the gas can no longer be compressed to liquid phase
regardless of pressure.

Critical Pressure is the pressure required to liquefy the gas at the critical temperature.

Compare triple point diagrams for water and carbon dioxide.

Vapor Pressure: All substances possess a vapor pressure, that is, pressure caused by
evaporating (or subliming) molecules. Vapor pressure depends on intermolecular
attractions (viscosity) and temperature. Non-polar liquids like ether or gasoline have
higher vapor pressures than polar liquids like water. The vapor pressure of solids is
lower still.

Tungsten metal is said to have a vapor pressure of

Vapor pressure is directly proportional to temperature.

• When a liquid is heated until its vapor pressure equals the pressure of the
atmosphere above the liquid, the boiling point is reached.

The molecules have so much kinetic energy and are moving so fast that they cannot
remain in the liquid phase, even if they are not at the liquid surface. The result is the
formation of vapor bubbles in the liquid that rise to the surface and burst.

• A liquid’s normal boiling point is the boiling temperature at 1 atmosphere of

pressure.
• Change the pressure, change the boiling point.

Solids

When molecules slow down enough so that the intermolecular forces of attraction are
strong enough to keep the molecules from moving past each other, the substance freezes
to solid form.

Non-polar substances (like the components of air) have very low freezing (and boiling)
points.

Polar and ionic substances have higher freezing/melting points.

If allowed to cool slowly enough, most substances will freeze into some sort of crystals.
Crystals are regular geometric arrangements of particles that have flat cleavable surfaces
and straight edges. There are many different crystal shapes.

If substances are cooled too rapidly for crystals to form, the material becomes an
“amorphous” solid. Amorphous literally means “without form,” but here is simply means
non-crystalline.

Glass is an example of an amorphous solid.

Some solids are classified as metallic or covalent network solids. Here, there are no
simple intermolecular forces holding particles together.

1. Metallic bonding: This type of bonding occurs in metals and metal alloys
(solutions). Think of this as a group of positive nuclei all sharing a collective
group of electrons. Scientists call this the “sea of electrons” or “electron gas”
bonding model. Since electrons do not adhere to a particular nucleus, but are free
to move around from one nucleus to another, most metals have the properties of
electrical conductivity and malleability. Metallic bonding is typically quite
strong, making metals hard solids at room temperature.

A second theory of metallic bonding is the MO (molecular orbital) theory,

which suggests that the s or s and p orbitals of neighboring metallic atoms form
overlapping molecular orbitals, allowing the collective sharing of outer electrons
by groups of metallic atoms. It may be thought of as being similar to the sharing
of delocalized electrons in molecules through the formation of resonance orbitals.

2. Covalent network solids: Within molecules there are covalent bonds holding the
atoms of the molecule together. Covalent bonds are extremely strong; much
stronger than intermolecular forces. In covalent network solids, atoms are bonded
covalently in a 3-dimensional array, making for a very strong crystalline solid.
The hardest substances known to man are covalent network solids. They include
diamond, silica (silicon dioxide or quartz) and silicon carbide. A diamond may be
thought of as one huge molecule of carbon atoms.