Chapter 11: Intermolecular Forces, Liquids, and Solids
11.1
A Molecular Comparison of Liquids and Solids
In contrast to gases, liquids and solids are condensed phases, meaning that their particles are very
close together. Liquids and solids are not significantly compressible because the molecules have
little free space between them. The particles of a condensed phase stay together because the average
kinetic energy of the particles is smaller than the average energy of the attractions between them.
The strongest attractive forces occur in solids, where particles are essentially locked in place. The
attractive forces between liquid particles are strong enough to keep them together, but weak enough
(relative to the kinetic energies of the particles) to allow motion. This motion of liquid particles
relative to one another is what allows liquids to flow.
Table 11.1 highlights some of the features of the three different phases of matter.
11.2
Intermolecular Forces
Intermolecular forces are the forces that exist between molecules; they are the attractive forces that
hold together the particles of a liquid or a solid. They are typically much weaker than covalent or
ionic bonds; in other words, it takes much more energy to break the covalent bond between H and Cl
than it does to vaporize HCl.
The strength of the intermolecular forces in a compound are reflected in its melting point and boiling
point. The stronger the intermolecular forces, the more energy is required to overcome those forces
in order to effect a phase change.
Forces that hold together molecules and atoms (both electrically neutral species) are referred to
collectively as van der Waals forces. They include dipole-dipole attractive forces, hydrogen
bonding, and London dispersion forces. Other types of intermolecular forces between particles
include ion-ion attractive forces and ion-dipole attractive forces (important in solutions).
An ion-dipole force exists between an ion and the partial charge on the end of a polar molecule. A
polar molecule has an end with a partial negative charge and an end with a partial positive charge.
The positive ends will orient themselves toward a negative ion, and the negative ends will orient
themselves toward a positive ion. As with all electrostatic attraction, the magnitude of the attraction
increases with the size of the charge (both on the ion and on the dipole).
Figure 11. 3. Illustration of the preferred orientation of polar molecules
toward ions. The negative end of the polar molecule is oriented toward a
cation (a), the positive end toward an anion (b).
Dipole-dipole forces exist between neutral polar molecules. The positive end of one polar molecule
is attracted to the negative end of another. The greater the polarity of the molecules, the stronger the
attractions between them. For molecules of approximately the same size, boiling point increases with
increasing dipole moment.
Many nonpolar substances exist as condensed phases. The intermolecular forces holding nonpolar
atoms or molecules together are known as London dispersion forces. On average, a nonpolar
molecule has a symmetrical distribution of electrons, with no partial positive or negative charge. At
any given instant, though, the mobile electrons may be positioned in such a way as to create an
instantaneous dipole moment. We say the molecule has become polarized. At the instant this
happens, the instantaneous partial charge influences the motion of electrons in neighboring
molecules. The instantaneous partial positive charge attracts the electrons in a neighboring molecule
toward itself, and the instantaneous partial negative charge repels electrons in another neighboring
molecule. This influence causes induced dipoles in neighboring molecules. The instantaneous and
induced dipole moments cause attractions between nonpolar molecules.
Figure 11. 5. Two schematic representations of the instantaneous dipoles on
two adjacent helium atoms, showing the electrostatic attraction between them.
In general, the more easily a molecule becomes polarized, the stronger the London dispersion forces
generated between molecules. The ease with which a molecule becomes polarized is called the
polarizability. Polarizability increases with increasing size, owing to the fact that valence electrons
are more mobile when they are more distant and more effectively screened from the nucleus. A
comparison of the intermolecular forces between halogen molecules and between noble-gas atoms
illustrates this. Molecular shape also plays a role in determining polarizability and therefore the
strength of intermolecular forces.
It is important to note that London dispersion forces, or simply dispersion forces, exist between all
molecules regardless of polarity.
1. When the molecules have comparable molecular weights and shapes, dispersion forces are
approximately equal. In this case differences in the magnitudes of the attractive forces are
due to differences in the strengths of dipole-dipole attractions, with the most polar molecules
having the strongest attractions.
2. When molecules differ widely in their molecular weights, dispersion forces tend to be the
decisive ones. In this case differences in the magnitudes of the attractive forces can usually
be associated with differences in molecular weights, with the most massive molecule having
the strongest attraction.
Among the strongest of the van der Waals forces are hydrogen bonds. Hydrogen bonding occurs
when hydrogen is bound to a small, very electronegative atom. (The most significant hydrogen
bonding occurs in compounds with an N–H, O–H, or F–H bond.) The electronegative atom
effectively strips hydrogen of its only electron, leaving a nearly unshielded proton exposed. This
proton is powerfully attracted to the atoms on neighboring molecules that have unshared electron
pairs. Hydrogen bonds significantly impact the properties of some substances.
All other things being equal, we would expect a series of compounds to exhibit an increase in
strength of intermolecular attractions with increasing molecular mass because of dispersion forces.
The hydrides of group 4A, for example, have boiling points roughly proportional to their molecular
masses. The same general trend is observed for the hydrides of group 6A until we get to water
(H2O), where hydrogen bonding becomes significant. Hydrogen bonding is responsible for the
anomalously high boiling point of water.
Figure 11.7. Boiling points of the group 4A (bottom) and group 6A (top)
hydrides as a function of molecular weight.
Hydrogen bonding is also responsible for another remarkable property of water—its solid phase is
less dense than its liquid phase.
Figure 11. 10. (a) Hydrogen bonding between two water molecules. The distances
shown are those found in ice.
(b) The arrangement of H2O molecules in ice. Each hydrogen atom in one H2O
molecule is oriented toward a nonbonding pair of electrons on an adjacent H 2O
molecule. As a result, ice has an open, hexagonal arrangement of H 2O molecules.
11.3
Some Properties of Liquids
Two of the noteworthy properties of liquids that are governed by their intermolecular forces are
viscosity and surface tension.
Viscosity is a fluid's resistance to flow. The more strongly liquid molecules are attracted to one
another, the greater their resistance to flow and the greater their viscosity. In general, increasing
temperature decreases a fluid's viscosity. This is because at a higher temperature the molecules have
a higher average kinetic energy and are better able to move and overcome the attractive forces
between them.
Molecules of a liquid are attracted to each other. In the bulk of the liquid, there is no net "pull" on a
molecule by the surrounding molecules because the attractions in all directions cancel each other
out. At the surface of the liquid, however, the molecules experience a net pull downward, causing
the top layer of molecules to act like a skin, minimizing the surface area. Surface tension is what
makes water form droplets (for a given volume, a sphere has the minimum surface area) and enables
some insects to walk on water.
The intermolecular forces that bind similar molecules together are called cohesive forces.
Intermolecular forces that bind a substance to a surface are called adhesive forces. Water in a
graduated cylinder, for example, forms a meniscus, or curved surface. The water is drawn upward
along the inside of the vessel because the adhesive forces between the water and the glass are
stronger than the cohesive forces between water molecules. In mercury the cohesive forces between
the mercury atoms are stronger than the adhesive forces between mercury and glass. The result is the
inverted meniscus shown in Figure 11.16. Adhesive forces are responsible for liquids rising up very
narrow tubes (and water rising up the stems of plants) through capillary action.
11.4
Phase Changes
A solid melting to become a liquid and a liquid vaporizing to become a gas are examples of phase
changes. A phase change may occur between any two of the three states of matter. Every phase
change has an energy change associated with it. Melting (also referred to as fusion), vaporization,
and sublimation are endothermic processes. Freezing, condensation, and deposition (the opposite of
sublimation) are exothermic processes. The energy associated with melting a solid is called the heat
of fusion; the energy associated with vaporizing a liquid is called the heat of vaporization; and the
energy associated with the sublimation of a solid is called the heat of sublimation. For a given
substance, the heat of vaporization will always be greater than the heat of fusion. Imparting enough
energy to molecules to allow them to move past one another takes less heat than separating the
molecules completely.
We can follow the energy and temperature changes for a substance with a heating curve like the one
pictured in Figure 11.18.
Figure 11.18. Heating curve for the transformation of 1.00 mol of water from
–25° C to 125° at a constant pressure of 1 atm. Blue lines show the heating of
one phase from a lower temperature to a higher one. Red lines show the
conversion of one phase to another at constant temperature.
As a gas is heated, it becomes more difficult to liquefy by applying pressure. The highest
temperature at which a gas can be liquefied by application of pressure is the critical temperature.
The pressure required to liquefy the gas at the critical temperature is the critical pressure. Above
the critical temperature, a gas cannot be liquefied—no matter how great the pressure. A substance at
a temperature and pressure above its critical temperature and pressure is referred to as a supercritical
fluid. Supercritical fluids have properties intermediate between gases and liquids. This makes them
very useful as solvents for extraction of one material from another. Now widely used in the food
industry, supercritical fluids found their first large-scale application in the processing of
decaffeinated coffee.
11.5
Vapor Pressure
At a given temperature the molecules of a liquid have a certain tendency to escape from the liquid
surface into the vapor phase. In a closed, evacuated system the evaporation of liquid molecules will
result in a vapor pressure being established above the liquid. As evaporation occurs, so does the
reverse process, condensation. Both processes will continue to occur until a dynamic equilibrium is
established whereby both forward and reverse processes continue to occur at equal rates, and the
vapor pressure above the liquid remains constant. The vapor pressure over the liquid when this
equilibrium has been established is the equilibrium vapor pressure. A substance with relatively weak
intermolecular forces will have a greater tendency to escape into the vapor phase and will exhibit a
higher vapor pressure then a substance with stronger intermolecular forces. A substance with a high
vapor pressure is said to be volatile.
As temperature increases, the number of liquid molecules with enough energy to escape from the
liquid surface increases, resulting in a higher equilibrium vapor pressure.
Figure 11. 21. Distribution of kinetic energies of surface molecules of a
hypothetical liquid at two temperatures. Only the fastest molecules have
sufficient kinetic energy to escape the liquid and enter the vapor, as shown by
the shaded areas. The higher the temperature, the larger the fraction of
molecules with enough energy to escape.
One measure of a compound's intermolecular forces is its boiling point. A liquid boils when its
vapor pressure equals the external pressure acting on the liquid's surface. The boiling point of a
liquid at 1 atm pressure is called its normal boiling point. This explains why cooking can take
longer at high elevations; the atmospheric pressure is lower at higher altitudes, so water boils at a
lower temperature. When the atmospheric pressure is less than 760 torr, water boils at less than
100°C.
Figure 11. 22. Vapor pressure of four common liquids, shown as a
function of temperature.
The temperature at which the vapor pressure is 760 torr is the normal
boiling point of each liquid.
11.6
Phase Diagrams
A phase diagram is a way to give a graphic summary of the conditions under which equilibria exist
between the solid, liquid, and gaseous states of matter for a substance. It allows us to predict the
phase of a substance that is stable at any given temperature and pressure.
The phase diagram for water is actually somewhat unusual in that the phase boundary line between
solid and liquid slopes backward, indicating that the melting point of water decreases with
increasing pressure. Carbon dioxide's phase diagram is much more typical. Increasing pressure
generally increases melting point.
Figure 11. 25. Phase diagram of (a) H2O and (b) CO2. The axes are
not drawn to scale in either case.
11.7
Structures of Solids
Solids may be amorphous or crystalline. Amorphous solids, such as glass and plastic, do not have
any long-range order and are somewhat heterogeneous in terms of their intermolecular forces. By
contrast, crystalline solids are highly ordered. This section will focus on crystalline solids.
A crystalline solid exists as a highly ordered network of a repeating unit called a unit cell. The unit
cell is the smallest portion of the solid that adequately conveys its structure and symmetry. A crystal
lattice is a three-dimensional array of points called lattice points, each of which is identical. We can
picture a crystalline solid as a collection of unit cells, each positioned at a lattice point.
Figure 11. 30.
The simplest unit cells are the cubic unit cells, in which the angles are all 90° and the lengths of all
sides are equal.
Sodium chloride crystallizes with a face-centered unit cell. Look at the salt, and locate a unit cell.
Notice that you can locate the unit cell either by positions of sodium ions (blue spheres) or by the
positions of the chloride ions (green spheres). It is important to remember when identifying the unit
cell of a crystalline solid that all lattice points must be identical. That is, we don't identify sodium
chloride as face-centered cubic, because there are sodium ions at the corners and chloride ions at the
faces. We identify it by the locations of either sodium ions or chloride ions, not both.
Figure 11. 34. Portion of the crystal lattice of NaCl, illustrating
two ways of defining its unit cell.
When an atom or an ion resides at a lattice point (with the exception of the center position), it does
not reside entirely in a single unit cell. An atom that sits at the corner of a unit cell also sits at the
corner of seven other unit cells. Only one eighth of the atom resides in any one particular unit cell.
11.8
Bonding in Solids
There are several varieties of crystalline solids. These can be classified according to the types of
forces between particles in solids.
Molecular solids are compounds that consist of molecules or atoms held together by intermolecular
forces such as dipole-dipole interactions and dispersion forces. They tend to be fairly soft and are
generally poor thermal and electrical conductors. Examples of molecular solids include ice, Dry
IceTM (solid carbon dioxide) and sugar.
Covalent-network solids are, in effect, giant molecules held together by covalent bonds. Carbon
exists as two common allotropes that are covalent-network solids: diamond and graphite. (It also
exists in the form of fullerenes, a class of carbon allotropes that includes so-called buckyballs.) This
type of solid tends to be very hard and possesses high melting points. They are also poor thermal and
electrical conductors.
Ionic solids consist of ions held together by electrostatic attractions. They are hard and brittle, have
high melting points, and are poor thermal and electrical conductors. Examples of ionic solids are
salts, including sodium chloride. An ionic solid, as you saw in the preceding section, exists as a
highly ordered lattice arrangement of symmetrical unit cells. Some of the common unit cells are
shown in Figure 11.42.
Figure 11. 42. (a) CsCl; (b) ZnS (zinc
blende); (c) CaF2 (fluorite).
Metallic solids consist of metal atoms held together by metallic bonding, which consists of a "sea" of
delocalized valence electrons. Metallic solids range in hardness from soft to very hard. They are
excellent thermal and electrical conductors, owing to their highly mobile valence electrons. They are
also malleable and ductile.
Figure 11. 45. A cross section of a metal. Each sphere represents the nucleus
and inner-core electrons of a metal atom. The surrounding colored "fog"
represents the mobile sea of electrons that binds the atoms together.