Review for Chapter 11: Intermolecular Forces and Liquids and Solids
1. Substances exist in one of three states: gas, liquid, or solid.
2. A gas has a low density and is very compressible. The molecules are free to move about and assume the volume and
shape of the container.
3. A liquid has high density and is only slightly compressible. The molecules are able to move and slide past one
another. A liquid has a definite volume but assumes the shape of the container.
4. A solid has a high density and is nearly incompressible because the molecules are already packed closely together. The
molecules are fixed in position but can vibrate about those positions. A solid has a definite volume and a definite shape.
5. Molecules are held together by two types of attractive forces: intramolecular forces (forces within one molecule –
chemical bonds) and intermolecular forces (forces between molecules). Intramolecular forces are generally much stronger
than intermolecular forces.
6. There are several different types of intermolecular forces as described below…
• ion-dipole forces: an ion and a polar molecule are attracted to each other
• dipole-dipole forces: polar molecules are attracted to each other
• dipole-induced dipole forces: a polar molecule induces a dipole in another molecule leading to attractive forces
• dispersion forces: temporary dipoles are induced that result in attractive forces between atoms or molecules
7. Van der Waals forces include dipole-dipole, dipole-induced dipole, and dispersion forces (but not ion-dipole forces).
8. Dispersion forces are sometimes called London forces and tend to increase with increasing molecular mass.
9. Hydrogen bonds are a special type of dipole-dipole interaction between a hydrogen atom (-H) in a polar bond (such as
N-H, O-H, or F-H) and an electronegative O, N, or F atom. Hydrogen bonding can be quite strong and raises the boiling
point of compounds. Without hydrogen bonding, water would be a gas at room temperature.
10. Liquids assume a geometry that minimizes the total surface area. Surface tension is the amount of energy required to
increase the surface of a liquid by a unit area. Strong intermolecular forces lead to high surface tensions.
11. Cohesion is the intermolecular attraction between like molecules.
Adhesion is the intermolecular attraction between unlike molecules.
For water, the adhesion of water molecules to glass molecules is greater than the cohesion of water molecules to each other
and so water will be pulled up by capillary action in a glass tube. In mercury, the cohesion between mercury molecules is
greater than the adhesion between the mercury and glass molecules so mercury will show at depression beneath a capillary
tube dipped in mercury.
12. Viscosity is a measure of a fluid’s resistance to flow. Intermolecular interactions such as hydrogen bonding or molecular
entanglements increase the resistance to flow and thus increase the viscosity.
13. Water is an unusual substance. Its solid phase (ice) is less dense than its liquid phase, which is the opposite of most
substances. This occurs because the highly ordered 3-D structure of ice keeps the H2O molecules from getting too close to
one another.
14. Solids can be divided into two categories: crystalline and amorphous.
Crystalline solids have rigid and long-range order.
Amorphous solids do not have a well-defined molecular order (Glass is a good example).
15. A unit cell is the basic structural unit of a crystalline solid, which is repeated to form a three-dimensional structure. The
lattice points of the unit cell may be occupied by atoms, ions, several atoms, or entire molecules.
16. There are seven types of unit cells (Figure 11.15): simple cubic, tetragonal, orthorhombic, rhombohedral, monoclinic,
triclinic, and hexagonal.
17. For cubic unit cells, there are three types of structures:
• simple cubic (contains a total of 1 complete sphere)
• body-centered cubic (contains a total of 2 complete spheres)
• face-centered cubic (contains a total of 4 complete spheres).
The relationship between the unit cell edge length and the radius of atoms for each structure is summarized in
Figure 11.22 and can be calculated if the density of the substance and crystal structure are known.
18. X-ray diffraction can provide information about the structure of solids. The Bragg equation, 2d sin  = n  , relates the
angle, , at which X-rays are diffracted to the distance, d, between planes in the crystal and to a multiple, n, of the X-ray
wavelength, .
19. There are four main types of crystals: ionic, covalent, molecular, and metallic.
20. Ionic crystals are composed of charged ions and are held together by electrostatic attraction (ionic bonding). These
crystals are generally hard, brittle, have a high melting point, and are poor conductors of heat and electricity.
Examples: NaCl, LiF, MgO, CaCO3.
21. Covalent crystals are composed of atoms held together by covalent bonds in an extensive three-dimensional network.
These crystals are generally hard, have a high melting point, and are usually poor conductors of heat and electricity.
Examples: C (diamond), C (graphite), SiO2 (quartz).
22. Molecular crystals are composed of molecules arranged at the lattice points of a crystal structure and held together by
dispersion forces, dipole-dipole forces, and/or hydrogen bonds. These crystals are generally soft, have low melting points,
and are poor conductors of heat and electricity. Examples: Ar, CO 2, I2, H2O, C12H22O11
23. Metallic crystals are composed of metal atoms arranged into a crystalline lattice structure held together by delocalized
“seas” of bonding electrons. These crystals range from soft to hard, have low to high melting points, and are good
conductors of heat and electricity. Examples: Na, Mg, Fe, Cu
24. Phase changes are transformations from one phase (gas, liquid, solid) to another, occurring when energy is added or
removed. These include evaporation/vaporization (liquid  gas), condensation (gas  liquid), melting (solid  liquid),
freezing (liquid  solid), sublimation (solid  gas), and deposition (gas  solid).
25. The Clausius-Clapeyron equation relates the vapor pressure P of a liquid to the absolute temperature T
ln P = - (∆Hvap) + C
RT
where ln P is the natural logarithm of the vapor pressure, ∆Hvap is the molar heat of vaporization, R is the gas constant
(R = 8.314 J/(K-mol)), and C is a constant.
26. The molar heat of sublimation is the sum of the molar heats of fusion and vaporization at a given temperature:
∆Hsub = ∆Hfus + ∆Hvap
27. A phase diagram shows the relationships among the phases of a single substance, where each region represents a pure
phase and the boundary lines between the regions show the temperatures and pressures at which the two phases are in
equilibrium.
28. At the triple point, where the three boundary lines meet on a phase diagram, all three phases can coexist and are in
equilibrium. The triple point for water is 0.006 atm and 0.01°C.
29. The critical temperature, Tc, is the temperature above which a substance exists as a gas and cannot be made to liquefy
no matter how much pressure is applied. The critical temperature of water is 374.4°C.