5 Steps to a 5: AP Chemistry 2017 (2016)


Review the Knowledge You Need to Score High


Solids, Liquids, and Intermolecular Forces


Summary: In the chapter on Gases, we discussed the gaseous state. In this chapter, we will discuss the liquid and solid states and the forces that exist between the particles––the intermolecular forces. A substance’s state of matter depends on two factors: the average kinetic energy of the particles, and the intermolecular forces between the particles. The kinetic energy tends to move the particles away from each other. The temperature of the substance is a measure of the average kinetic energy of the molecules. As the temperature increases, the average kinetic energy increases and the particles tend to move farther apart. This is consistent with our experience of heating ice, for example, and watching it move from the solid state to the liquid state and finally to the gaseous state. For this to happen, the kinetic energy overcomes the forces between the particles, the intermolecular forces.

In the solid state, the kinetic energy of the particles cannot overcome the intermolecular forces; the particles are held close together by the intermolecular forces. As the temperature increases, the kinetic energy increases and begins to overcome the attractive intermolecular forces. The substance will eventually melt, going from the solid to the liquid state. As this melting takes place, the temperature remains constant even though energy is being added. The temperature at which the solid converts into the liquid state is called the melting point (m.p.) of the solid.

After all the solid has been converted into a liquid, the temperature again starts to rise as energy is added. The particles are still relatively close together, but possess enough kinetic energy to move with respect to each other. Finally, if enough energy is added, the particles start to break free of the intermolecular forces keeping them relatively close together and they escape the liquid as essentially independent gas particles. This process of going from the liquid state to the gaseous state is called boiling, and the temperature at which this occurs is called the boiling point (b.p.) of the liquid. Sometimes, however, a solid can go directly from the solid state to the gaseous state without ever having become a liquid. This process is called sublimation. Dry ice, solid carbon dioxide, readily sublimes.

These changes of state, called phase changes , are related to temperature, but sometimes pressure can influence the changes. We will see how these relationships can be diagrammed later in this chapter.

Keywords and Equations

No specific keywords or equations are listed on the AP exam for this topic.

Structures and Intermolecular Forces

Intermolecular forces are attractive or repulsive forces between molecules, caused by partial charges. The attractive forces are the ones that work to overcome the randomizing forces of kinetic energy. The structure and type of bonding of a particular substance have quite a bit to do with the type of interaction and the strength of that interaction. Before we start examining the different types of intermolecular forces, recall from the Bonding chapter that those molecules that have polar covalent bonding (unequal sharing of the bonding electron pair) may possess dipoles (having positive and negative ends due to charge separation within the molecule). Dipoles are often involved in intermolecular forces.

Ion–Dipole Intermolecular Forces

These forces are due to the attraction of an ion and one end of a polar molecule (dipole). This type of attraction is especially important in aqueous salt solutions, where the ion attracts water molecules and may form a hydrated ion, such as Al(H2 O)6 3+ . This is one of the strongest of the intermolecular forces.

It is also important to realize that this intermolecular force requires two different species––an ion and a polar molecule.

Dipole–Dipole Intermolecular Forces

These forces result from the attraction of the positive end of one dipole to the negative end of another dipole. For example, in gaseous hydrogen chloride, HCl(g), the hydrogen end has a partial positive charge and the chlorine end has a partial negative charge, due to chlorine’s higher electronegativity. Dipole–dipole attractions are especially important in polar liquids. They tend to be a rather strong force, although not as strong as ion–dipole attractions.

Hydrogen Bond Intermolecular Forces

Hydrogen bonding is a special type of dipole–dipole attraction in which a hydrogen atom is polar-covalently bonded to one of the following extremely electronegative elements: N, O, or F. These hydrogen bonds are extremely polar bonds by nature, so there is a great degree of charge separation within the molecule. Therefore, the attraction of the positively charged hydrogen of one molecule and the negatively charged N, O, or F of another molecule is extremely strong. These hydrogen bonds are, in general, stronger than the typical dipole–dipole interaction.

Hydrogen bonding explains why HF(aq) is a weak acid, while HCl(aq), HBr(aq), etc. are strong acids. The hydrogen bond between the hydrogen of one HF molecule and the fluorine of another “traps” the hydrogen, so it is much harder to break its bonds and free the hydrogen to be donated as an H+ . Hydrogen bonding also explains why water has such unusual properties—for example, its unusually high boiling point and the fact that its solid phase is less dense than its liquid phase. The hydrogen bonds tend to stabilize the water molecules and keep them from readily escaping into the gas phase. When water freezes, the hydrogen bonds are stabilized and lock the water molecules into a framework with a lot of open space. Therefore, ice floats in liquid water. Hydrogen bonding also holds the strands of DNA together.

Ion-Induced Dipole and Dipole-Induced Dipole Intermolecular Forces

These types of attraction occur when the charge on an ion or a dipole distorts the electron cloud of a nonpolar molecule and induces a temporary dipole in the nonpolar molecule. Like ion–dipole intermolecular forces, these also require two different species. They are fairly weak interactions.

London (Dispersion) Intermolecular Force

This intermolecular attraction occurs in all substances, but is significant only when the other types of intermolecular forces are absent. It arises from a momentary distortion of the electron cloud, with the creation of a very weak dipole. The weak dipole induces a dipole in another nonpolar molecule. This is an extremely weak interaction, but it is strong enough to allow us to liquefy nonpolar gases such as hydrogen, H2 , and nitrogen, N2 . If there were no intermolecular forces attracting these molecules, it would be impossible to liquefy them.

The Liquid State

At the microscopic level, liquid particles are in constant flux. They may exhibit short-range areas of order, but these do not last very long. Clumps of particles may form and then break apart. At the macroscopic level, a liquid has a specific volume but no fixed shape. Three other macroscopic properties deserve discussion: surface tension, viscosity, and capillary action. In the body of a liquid, the molecules are pulled in all different ways by the intermolecular forces between them. At the surface of the liquid, the molecules are only being pulled into the body of the liquid from the sides and below, not from above. The effect of this unequal attraction is that the liquid tries to minimize its surface area by forming a sphere. In a large pool of liquid, where this is not possible, the surface behaves as if it had a thin “skin” over it. It requires force to break the attractive forces at the surface. The amount of force required to break through this molecular layer at the surface is called the liquid’s surface tension . The greater the intermolecular forces, the greater the surface tension. Polar liquids, especially those that undergo hydrogen bonding, have a much higher surface tension than nonpolar liquids.

Viscosity , the resistance of liquids to flow, is affected by intermolecular forces, temperature, and molecular shape. Liquids with strong intermolecular forces tend to have a higher viscosity than those with weak intermolecular forces. Again, polar liquids tend to have a higher viscosity than nonpolar liquids. As the temperature increases, the kinetic energy of the particles becomes greater, overcoming the intermolecular attractive forces. This causes a lower viscosity. Finally, the longer and more complex the molecules, the more contact the particles will have as they slip by each other, increasing the viscosity.

Capillary action is the spontaneous rising of a liquid through a narrow tube, against the force of gravity. It is caused by competition between the intermolecular forces in the liquid and those attractive forces between the liquid and the tube wall. The stronger the attraction between the liquid and the tube, the higher the level will be. Liquids that have weak attractions to the walls, like mercury in a glass tube, have a low capillary action. Liquids like water in a glass tube have strong attractions to the walls and will have a high capillary action.

As we have noted before, water, because of its stronger intermolecular forces (hydrogen bonding) has some very unusual properties. It will dissolve a great number of substances, both ionic and polar covalent, because of its polarity and ability to form hydrogen bonds. It is sometimes called the “universal solvent.” It has a high heat capacity , the heat absorbed to cause the temperature to rise, and a high heat of vaporization , the heat needed to transform the liquid into a gas. Both of these thermal properties are due to the strong hydrogen bonding between the water molecules. Water has a high surface tension for the same reason. The fact that the solid form of water (ice) is less dense than liquid water is because water molecules in ice are held in a rigid, open, crystalline framework by the hydrogen bonds. As the ice starts melting, the crystal structure breaks and water molecules fill the holes in the structure, increasing the density. The density reaches a maximum at around 4°C; then the increasing kinetic energy of the particles causes the density to begin to decrease.

The Solid State

At the macroscopic level, a solid is defined as a substance that has both a definite volume and a definite shape. At the microscopic level, solids may be one of two types––amorphous or crystalline. Amorphous solids lack extensive ordering of the particles. There is a lack of regularity of the structure. There may be small regions of order separated by large areas of disordered particles. They resemble liquids more than solids in this characteristic. Amorphous solids have no distinct melting point. They simply get softer and softer as the temperature rises, leading to a decrease in viscosity. Glass, rubber, and charcoal are examples of amorphous solids.

Crystalline solids display a very regular ordering of the particles in a three-dimensional structure called the crystal lattice . In this crystal lattice there are repeating units called unit cells . Figure 12.1 shows the relationship of the unit cells to the crystal lattice.

Figure 12.1   The crystal lattice for a simple cubic unit cell.

Several types of unit cells are found in solids. The cubic system is the most common type. Three types of unit cells are found in the cubic system:

  1. Thesimple cubic unit cell has particles located at the corners of a simple cube.
  2. Thebody-centered unit cell has particles located at the corners of the cube and in the center of the cube.
  3. Theface-centered unit cell has particles at the corners and one in the center of each face of the cube, but not in the center of the cube itself.

Figure 12.2 shows three types of cubic unit cells.

Figure 12.2   The three types of unit cell of the cubic lattice.

Five types of crystalline solid are known:

  1. Inatomic solids, individual atoms are held in place by London forces. The noble gases are the only atomic solids known to form.
  2. Inmolecular solids, lattices composed of molecules are held in place by London forces, dipole–dipole forces, and hydrogen bonding. Solid methane and water are examples of molecular solids.
  3. Inionic solids, lattices composed of ions are held together by the attraction of the opposite charges of the ions. These crystalline solids tend to be strong, with high melting points because of the strength of the intermolecular forces. NaCl and other salts are examples of ionic solids. Figure 12.3 shows the lattice structure of NaCl. Each sodium cation is surrounded by six chloride anions, and each chloride anion is surrounded by six sodium cations.
  4. Inmetallic solids , metal atoms occupying the crystal lattice are held together by metallic bonding. In metallic bonding , the electrons of the atoms are delocalized and are free to move throughout the entire solid. This explains electrical and thermal conductivity, as well as many other properties of metals.
  5. Incovalent network solids , covalent bonds join atoms together in the crystal lattice, which is quite large. Graphite, diamond, and silicon dioxide (SiO2 ) are examples of network solids. The crystal is one giant molecule.

Figure 12.3   Sodium chloride crystal lattice.

Phase Diagrams

The equilibrium that exists between a liquid and its vapor is just one of several that can exist between states of matter. A phase diagram is a graph representing the relationship of a substance’s states of matter to temperature and pressure. The diagram allows us to predict which state of matter a substance will assume at a certain combination of temperature and pressure. Figure 12.4 , on the next page, shows a general form of the phase diagram.

Note that the diagram has three general areas corresponding to the three states of matter––solid, liquid, and gas. The line from A to C represents the solid’s change in vapor pressure with changing temperature, for the sublimation equilibrium. The A-to-D line represents the variation in the melting point with varying pressure. The A-to-B line represents the variation of a liquid’s vapor pressure with varying pressure. The B point shown on this phase diagram is called the critical point of the substance, the point beyond which the gas and liquid phases are indistinguishable from each other. At or beyond this critical point, no matter how much pressure is applied, the gas cannot be condensed into a liquid. Point A is the substance’s triple point , the combination of temperature and pressure at which all three states of matter can exist together. The phase diagram for water is shown in Figure 12.5 .

Figure 12.4   A phase diagram.

For each of the phase transitions, there is an associated enthalpy change or heat of transition. For example, there are heats of vaporization, fusion, sublimation, and so on.

Figure 12.5   Phase diagram for H2 O.

Relationship of Intermolecular Forces to Phase Changes

The intermolecular forces can affect phase changes to a great degree. The stronger the intermolecular forces present in a liquid, the more kinetic energy must be added to convert it into a gas. Conversely, the stronger the intermolecular forces between the gas particles, the easier it will be to condense the gas into a liquid. In general, the weaker the intermolecular forces, the higher the vapor pressure. The same type of reasoning can be used about the other phase equilibria—in general, the stronger the intermolecular forces, the higher the heats of transition.

Example: Based on intermolecular forces, predict which will have the higher vapor pressure and higher boiling point, water or dimethyl ether, CH3 –O–CH3 .

Answer: Dimethyl ether will have the higher vapor pressure and the lower boiling point.

Explanation: Water is a polar substance with strong intermolecular hydrogen bonds. Dimethyl ether is a polar material with weaker intermolecular forces (dipole–dipole). It will take much more energy to vaporize water; thus, water has a lower vapor pressure and higher boiling point.


The concept of intermolecular forces is important in the separation of the components of a mixture.

Common Mistakes to Avoid

  1. Don’t confuse the various types of intermolecular forces.
  2. The melting point and the freezing point are identical.
  3. Hydrogen bonding can occur only when a hydrogen atom is directly bonded to an N, O, or F atom.
  4. When moving from point to point in a phase diagram, pay attention to which phase transitions the substance exhibits.
  5. In looking at crystal lattice diagrams, be sure to count all the particles, in all three dimensions, that surround another particle.

 Review Questions

Use these questions to review the content of this chapter and practice for the AP Chemistry exam. Below are 14 multiple-choice questions similar to what you will encounter in Section I of the AP Chemistry exam. To make these questions an even more authentic practice for the actual exam, time yourself following the instructions provided.

Multiple-Choice Questions

Answer the following questions in 20 minutes. You may not use a calculator. You may use the periodic table and the equation sheet at the back of this book.

1 . Which of the following best describes Fe(s)?

(A) composed of macromolecules held together by strong bonds

(B) composed of atoms held together by delocalized electrons

(C) composed of positive and negative ions held together by electrostatic attractions

(D) composed of molecules held together by intermolecular dipole–dipole interactions

2 . The best description of the interactions in KNO3 (s) is which of the following?

(A) composed of macromolecules held together by strong bonds

(B) composed of atoms held together by delocalized electrons

(C) composed of positive and negative ions held together by electrostatic attractions

(D) composed of molecules held together by intermolecular dipole–dipole interactions

3 . Sand is primarily SiO2 (s). Which of the following best describes the interactions inside a grain of sand?

(A) composed of macromolecules held together by strong bonds

(B) composed of atoms held together by delocalized electrons

(C) composed of positive and negative ions held together by electrostatic attractions

(D) composed of molecules held together by intermolecular dipole–dipole interactions

4 . At sufficiently low temperatures, it is possible to form HCl(s). What best describes the interactions in this solid?

(A) composed of macromolecules held together by strong bonds

(B) composed of atoms held together by delocalized electrons

(C) composed of positive and negative ions held together by electrostatic attractions

(D) composed of molecules held together by intermolecular dipole–dipole interactions

5 . Which of the following best describes diamond, C(s)?

(A) an ionic solid

(B) a metallic solid

(C) a molecular solid containing polar molecules

(D) a covalent network solid

6 . What type of solid is solid sulfur dioxide, SO2 (s)?

(A) an ionic solid

(B) a metallic solid

(C) a molecular solid containing polar molecules

(D) a covalent network solid

7 . The approximate boiling points for hydrogen compounds of some elements in the nitrogen family are: (SbH3 – 15°C), (AsH3 – 62°C), (PH3 – 87°C), and (NH3 – 33°C). The best explanation for the fact that NH3 does not follow the trend of the other hydrogen compounds is:

(A) NH3 is the only one to exhibit hydrogen bonding.

(B) NH3 is the only one that is water-soluble.

(C) NH3 is the only one that is nearly ideal in the gas phase.

(D) NH3 is the only one that is a base.

8 . Why is it possible to solidify argon at a sufficiently low temperature?

(A) London dispersion forces

(B) covalent bonding

(C) hydrogen bonding

(D) metallic bonding

9 . Which of the following best describes why diamond is so hard?

(A) London dispersion forces

(B) covalent bonding

(C) hydrogen bonding

(D) metallic bonding

10 . A sample of a pure liquid is placed in an open container and heated to the boiling point. Which of the following may increase the boiling point of the liquid?

(A) The moles of liquid are increased.

(B) The size of the container is increased.

(C) A vacuum is created over the liquid.

(D) The container is sealed.

11 . Which of the following best explains why 1-butanol, CH3 CH2 CH2 CH2 OH, has a higher surface tension than its isomer, diethyl ether, CH3 CH2 OCH2 CH3 ?

(A) the higher density of 1-butanol

(B) the lower specific heat of 1-butanol

(C) the lack of hydrogen bonding in 1-butanol

(D) the presence of hydrogen bonding in 1-butanol

12 . Pick the answer that most likely represents the substances’ relative solubilities in water.

(A) CH3 CH2 CH2 CH3 < CH3 CH2 CH2 OH < HOCH2 CH2 OH

(B) CH3 CH2 CH2 OH < CH3 CH2 CH2 CH3 < HOCH2 CH2 OH

(C) CH3 CH2 CH2 CH3 < HOCH2 CH2 OH < CH3 CH2 CH2 OH

(D) HOCH2 CH2 OH < CH3 CH2 CH2 OH < CH3 CH2 CH2 CH3

13 . What is the energy change that accompanies the conversion of molecules in the gas phase to a liquid?

(A) heat of condensation

(B) heat of deposition

(C) heat of sublimation

(D) heat of fusion

14 . Which of the following explains why the melting point of sodium chloride (NaCl 801°C) is lower than the melting point of calcium fluoride (CaF2 1423°C)?

(A) Sodium is more reactive than calcium is.

(B) The chloride ion is smaller than the fluoride ion.

(C) The charge on a sodium ion is less than the charge on a calcium ion.

(D) The ratio of anions to cations is lower in sodium chloride.

Answers and Explanations for the Multiple-Choice Questions

1 . B —This answer describes a metallic solid.

2 . C —This answer describes an ionic solid.

3 . A —This answer describes a covalent network solid.

4 . D —This answers describes a solid consisting of discrete polar molecules. Even though HCl(aq) is a strong acid with ions in solution, there is no water here to lead to ionization.

5 . D —Each of the carbon atoms is covalently bonded to four other carbon atoms.

6 . C —Sulfur dioxide molecules are polar.

7 . A —Hydrogen bonding occurs when hydrogen is directly bonded to F, O, and in this case N.

8 . A —Argon is a noble gas; none of the other bonding choices is an option.

9 . B —Diamond is a covalent network solid with a large number of strong covalent bonds between the carbon atoms.

10 . D —The size of the container or the number of moles is irrelevant. Sealing the container will cause an increase in pressure that will increase the boiling point. A decrease in pressure will lower the boiling point.

11 . D —The compound with the higher surface tension is the one with the stronger intermolecular force. The hydrogen bonding in 1-butanol is stronger than the dipole-dipole attractions in diethyl ether.

12 . A —The sequence for these similar molecules is nonpolar, then one hydrogen bond, then two hydrogen bonds.

13 . A —This change is condensation, so the energy is the heat of condensation.

14 . C —The only applicable factor listed is the charge difference. The chloride ion is larger than the fluoride ion. The ion ratio is not important nor is the reactivity of the elements.

 Rapid Review

  • The state of matter in which a substance exists depends on the competition between the kinetic energy of the particles (proportional to temperature) and the strength of the intermolecular forces between the particles.
  • The melting point is the temperature at which a substance goes from the solid to the liquid state and is the same as the freezing point.
  • The boiling point is the temperature at which a substance goes from the liquid to the gaseous state. This takes place within the body of the liquid, unlike evaporation, which takes place only at the surface of the liquid.
  • Sublimation is the conversion of a solid to a gas without ever having become a liquid. Deposition is the reverse process.
  • Phase changes are changes of state.
  • Intermolecular forces are the attractive or repulsive forces between atoms, molecules, or ions due to full or partial charges. Be careful not to confuse intermolecular forces with intramolecular forces, the forces within the molecule.
  • Ion–dipole intermolecular forces occur between ions and polar molecules.
  • Dipole–dipole intermolecular forces occur between polar molecules.
  • Hydrogen bonds are intermolecular forces between dipoles in which there is a hydrogen atom attached to an N, O, or F atom.
  • Ion-induced dipole intermolecular forces occur between an ion and a nonpolar molecule.
  • London (dispersion) forces are intermolecular forces between nonpolar molecules.
  • Liquids possess surface tension (liquids behaving as if they had a thin “skin” on their surface, due to unequal attraction of molecules at the surface of the liquid), viscosity (resistance to flow), and capillary action (flow up a small tube).
  • Amorphous solids have very little structure in the solid state.
  • Crystalline solids have a great deal of structure in the solid state.
  • The crystal lattice of a crystalline solid is the regular ordering of the unit cells.
  • Know the five types of crystalline solid: atomic, molecular, ionic, metallic, and network.
  • Phase changes can be related to the strength of intermolecular forces.