5 Steps to a 5: AP Chemistry 2020 - Moore J.T., Langley R.H. 2020

STEP 4 Review the Knowledge You Need to Score High
CHAPTER 12 Solids, Liquids, and Intermolecular Forces

IN THIS CHAPTER

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(H₂O)₆³⁺. 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. The simple cubic unit cell has particles located at the corners of a simple cube.

2. The body-centered unit cell has particles located at the corners of the cube and in the center of the cube.

3. The face-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. In atomic solids, individual atoms are held in place by London forces. The noble gases are the only atomic solids known to form.

2. In molecular 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. In ionic 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.

Figure 12.3 Sodium chloride crystal lattice.

4. In metallic 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. In covalent 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.

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.

Figure 12.4 A 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 12.5 Phase diagram for H₂O.

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.

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.

Experiments

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 18 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 25 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) Iron is composed of macromolecules held together by strong bonds.

(B) Iron is composed of atoms held together by delocalized electrons.

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

(D) Iron is 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) KNO₃ is composed of macromolecules held together by strong bonds.

(B) KNO₃ is composed of atoms held together by delocalized electrons.

(C) KNO₃ is composed of positive and negative ions held together by electrostatic attractions.

(D) KNO₃ is 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) Sand is composed of macromolecules held together by strong bonds.

(B) Sand is composed of atoms held together by delocalized electrons.

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

(D) Sand is 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) HCl(s) is composed of macromolecules held together by strong bonds.

(B) HCl(s) is composed of atoms held together by delocalized electrons.

(C) HCl(s) is composed of positive and negative ions held together by electrostatic attractions.

(D) HCl(s) is composed of molecules held together by intermolecular dipole—dipole interactions.

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

(A) Diamond is an ionic solid.

(B) Diamond is a metallic solid.

(C) Diamond is a molecular solid containing polar molecules.

(D) Diamond is a covalent network solid.

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

(A) Sulfur dioxide is an ionic solid.

(B) Sulfur dioxide is a metallic solid.

(C) Sulfur dioxide is a molecular solid containing polar molecules.

(D) Sulfur dioxide is 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). Which of the following is the best explanation for the fact that NH₃ does not follow the trend of the other hydrogen compounds?

(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 are present.

(B) Covalent bonds form.

(C) Hydrogen bonds form.

(D) Metallic bonds form.

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

(A) London dispersion forces are present.

(B) Covalent bonds are present.

(C) Hydrogen bonds are present.

(D) Metallic bonds are present.

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. The compounds 1-butanol, CH₃CH₂CH₂CH₂OH, and diethyl ether, CH₃CH₂OCH₂CH₃, are isomers (have the same chemical formula, C₄H₁₀O). However, the surface tension of 1-butanol is higher than that of diethyl ether. Which of the following best explains why the surface tension of 1-butanol is higher than that of diethyl ether?

(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. The compounds propanol, CH₃CH₂CH₂OH, ethylene glycol, HOCH₂CH₂OH, and butane, CH₃CH₂CH₂CH₃, are all similar in size. Which of the following most likely represents the relative solubilities of these compounds in water?

(A) CH₃CH₂CH₂OH < HOCH₂CH₂OH < CH₃CH₂CH₂CH₃

(B) CH₃CH₂CH₂CH₃ < HOCH₂CH₂OH < CH₃CH₂CH₂OH

(C) CH₃CH₂CH₂CH₃ < CH₃CH₂CH₂OH < HOCH₂CH₂OH

(D) CH₃CH₂CH₂OH < CH₃CH₂CH₂CH₃ < HOCH₂CH₂OH

13. Each of the following compounds has a similar structure, and the interionic distances are about the same size. Which compound is expected to have the highest lattice energy?

(A) LiCl

(B) MgO

(C) AlN

(D) NaF

14. The melting point of sodium fluoride, NaF, is 993°C, and the melting point of calcium fluoride, CaF₂, is 1,423°C. Which of the following best explains why the melting point of calcium fluoride is higher?

(A) Sodium is a more reactive element than calcium is.

(B) The sodium ion is smaller than the calcium ion.

(C) The sodium ion has a lower charge than the calcium ion.

(D) There are fewer fluoride ions in the formula of sodium fluoride.

15. Pure phosphoric acid, H₃PO₄, freezes at about 42°C. Which of the following best describes the interactions in solid phosphoric acid?

(A) The solid consists of a collection of molecules held together by London dispersion forces.

(B) The solid consists of a collection of atoms held together by delocalized electrons.

(C) The solid consists of positive and negative ions held together by ionic bonding.

(D) The solid consists of molecules held together by intermolecular dipole—dipole interactions

16. Which of the following liquids will probably have the highest viscosity at a given temperature?

(A) CH₃CH₂CH₂OH

(B) PCl₃

(C) CCl₄

(D) CH₃CH₂CH₂CH₃

17. Most molecular species vaporize as simple molecules. However, acetic acid, CH₃COOH, is an exception. In the gaseous state, acetic acid molecules exist as pairs known as dimers. What is the most probable cause of this behavior?

(A) London dispersion forces

(B) covalent bonding

(C) hydrogen bonding

(D) metallic bonding

18. The above diagram represents the heating curve for a pure crystalline substance. Both solid and liquid are present in which region of the heating curve?

(A) between points A and B

(B) between points B and C

(C) between points C and D

(D) between points D and E

Answers and Explanations

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 answer 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 exists as molecules that 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 many strong covalent bonds between the carbon atoms.

10. D—The size of the container and the number of moles are 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 —OH groups can form hydrogen bonds. The hydrogen bonding in 1-butanol is stronger than the dipole—dipole attractions in diethyl ether.

12. C—Butane is nonpolar, whereas the other two compounds can form hydrogen bonds through the —OH groups. The more —OH groups, the more hydrogen bonding, and the more soluble in water (where hydrogen bonding also occurs).

13. C—If the structures are similar and the interionic distances are about the same, then the only factor remaining that will influence the lattice energy is the magnitude of the charges. The ionic charges are Li⁺, Cl⁻, Mg²⁺, O²⁻, Al³⁺, N³⁻, Na⁺, and F⁻. The largest magnitude charges are in AlN, which will lead to the greatest attraction between the ions, as indicated by the highest lattice energy.

14. C—For ionic compounds, the key factors are the sizes and the charges on the ions. Smaller ions lead to higher melting points. Higher charges lead to higher melting points. Answer C deals with size; however, if this were the key factor, the melting point of sodium fluoride would be higher, which is not the case. This leads to C being the correct answer. The higher charge on calcium (+2 versus +1) leads to a higher lattice energy, which results in a higher melting point.

15. D—This answer describes a solid consisting of discrete polar molecules. Answer A refers to a nonpolar molecule. Answer B refers to a metallic solid. Answer C might apply; however, even though phosphoric acid is an acid that forms ions in aqueous solution, there is no water present for ionization.

16. A—Compound A can form hydrogen bonds. Compound B has dipole—dipole interactions. Compounds C and D are both nonpolar; therefore, their strongest intermolecular force is London dispersion force. For relatively small molecules such as these, the stronger the intermolecular force, the higher the viscosity. Hydrogen bonding in the strongest of the three intermolecular forces present in these molecules.

17. C—The carbonyl, C=O, and —OH groups can participate in hydrogen bonds, as illustrated by the dotted lines in the following diagram:

The hydrogen bonds hold the two molecules together as a dimer.

18. B—The solid begins to melt at B and finishes melting at C. Once the solid begins to melt, both solid and liquid are present.

Free-Response Question

You have 15 minutes to answer the following question. You may use a calculator and the tables in the back of the book.

Question

Write a brief explanation concerning each of the following observations.

(a) An aqueous solution in a glass buret has a meniscus that is concave, whereas in a plastic buret, the meniscus is convex.

(b) The surface tension of liquid CBr₄ is greater than the surface tension of liquid CCl₄.

(c) At the same temperature, liquid HF has a higher viscosity than liquid HCl.

(d) When heated, glass does not melt; it becomes soft.

Answer and Explanation

(a) The concave surface results from capillary action, and the convex surface results from surface tension. Surface tension is the result of strong intermolecular forces between the solution and the glass; these forces are greater than those within the solution causing a convex surface. The convex surface results from a stronger interaction within the solution than between the solution and the plastic.

Give yourself 1 point for this answer.

(b) The greater surface tension results from stronger intermolecular forces in CBr₄. Both CCl₄ and CBr₄ are nonpolar molecules; therefore, the intermolecular forces are London dispersion forces. For simple molecules, such as these, the one with more electrons, CBr₄, will have the stronger London dispersion forces and hence the greater surface tension.

Give yourself 1 point for this answer.

(c) Viscosity, the resistance to flow, is related to the strength of attraction between the molecules. Both HF and HCl are polar molecules with dipole—dipole forces; however, HF exhibits an especially strong form of dipole—dipole force known as hydrogen bonding. Since hydrogen bonding is stronger than normal dipole—dipole forces, HF has a higher viscosity.

Give yourself 1 point for this answer.

(d) Glass is a liquid with very high viscosity (amorphous solid). Heating any liquid gives the molecules more kinetic energy, which lowers the viscosity. In the case of glass, the decrease in viscosity begins with the material becoming softer and softer until it flows freely.

Give yourself 1 point for this answer.

Total your points. There are 4 points possible.

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.