Chapter 11 Reading Guide Intermolecular Forces Dr. Baxley Tro 3 rd edition 1 Section 11.1: How do geckos stick to walls? What biological molecules have shapes and structures that depend on intermolecular forces? Section 11.2: How do properties of solids, liquids, and gases differ? Read over the differences between the molecular properties of solids, liquids, and gases. Which type of matter is usually the most dense? Is ice more dense than liquid water? Is this typical for most solids/liquids? In Table 11.1, how are the molecules spaced differently? After viewing Figure 11.2, why is it possible to own an air compressor but not a liquid or solid compressor? Answer Conceptual Connection 11.1. Why is the answer not (c)? Section 11.3: What holds molecules and atoms together in solids and liquids? Ignore the formula with q 1q 2/4πE 0r. Which force is usually stronger, an intermolecular force (IMF) or an intramolecular force (chemical bond)? At what temperatures do the intermolecular and intramolecular bonds of water break apart? Do dispersion forces occur in all atoms and molecules or in a small subset of substances? Why are dispersion forces called instantaneous or induced dipoles? Does this type of dipole last a long time? Why does neopentane have a lower boiling point than pentane, even though they have the same molar mass? In what type(s) of molecules do dipole-dipole forces exist? In what way are dipole-dipole forces different than London forces? Why does formaldehyde have a higher boiling point than ethane, even though they have a similar molar mass?
How do the molecules align in figure 11.7? 2 What property has a fairly linear relationship with boiling point for molecules with similar molar masses (Figure 11.8)? What does the term miscibility define? Answer the questions in Example 11.1 and For Practice 11.1. What is an important first step in answering a question of this type? What 3 elements must hydrogen be bonded to in order for a molecule to undergo hydrogen bonding? What 3 elements can a hydrogen attached to a N, O, or F atom form a hydrogen bond with? What is special about the 3 elements N, O, and F (think chapter 9)? What two physical properties are significantly impacted by the ability of a molecule to hydrogen bond with other molecules? Review Figures 11.10, 11,11, and 11.12 to determine how hydrogen bonding can be represented in different molecules. Why does ethanol have a much higher boiling point than dimethyl ether, even though they have the same molecular formulas? Answer questions in Example 11.2 and For Practice 11.2. Why are two of the molecules in the Example unable have hydrogen bonding between molecules? What is the strongest IMF discussed so far? What is an ion-dipole force, other than the last and not very interesting IMF? Read the summarizing bullets on page 495, and review Table 11.4. Answer Conceptual Connection 11.4. What concepts or skills are important in answering this kind of question? The following page (Chemistry and Medicine) is a great visualization and description of hydrogen bonding in DNA, something you looked at in the computer modeling lab. Do the base pairing IMFs in DNA fit with the existing definition of hydrogen bonds? Answer the Question in this feature box (about dispersion forces and DNA).
This is a nice video that summarizes the intermolecular forces, with examples: https://www.youtube.com/watch?v= E-tDn9qSg 3 Section 11.4: How do water skimmers float? What does the property of surface tension measure? Looking at Figure 11.18, note that the energy of attraction of surface molecules is spread over fewer particles (4) than the molecules in the middle of the solution (6). This means that each IMF ends up being stronger, creating what is called surface tension. Why does a paper clip float in water, but not in benzene? How does surface tension correlate with IMFs? What does the property of viscosity measure, and why do some substances have a higher viscosity than others? How does viscosity correlate with IMFs? What does the phenomenon of capillary action describe? Why does liquid water form a meniscus in a glass tube or cylinder, like a pipet? Section 11.5, What is vapor pressure? Examine Figure 11.23 after reading the first two paragraphs of this section. Why might some water molecules leave the bulk of the liquid while others remain behind? Why is the term vapor pressure used in this section? Figure 11.24 is important. The y-axis is a fraction from 0-1 (or percentage from 0-100) and the x axis is how much kinetic energy. The higher the temperature, the wider the distribution of energy (meaning that the upper limit is how further from zero on the x-axis). This means that the y-axis peak must be lower, because of the wide distribution. The dotted vertical line shows the minimum energy needed for a liquid molecule to move to the vapor phase. Is the percentage of low-temperature molecules that can move to the gas phase higher or lower than the percentage of high-temperature molecules?
How might IMFs impact the number of liquid particles with enough energy to move to the vapor phase? 4 Does vaporization of a liquid to a gas require the liquid to absorb or release energy? Is vaporization endothermic or exothermic? Why is the ΔH vap of water higher than the ΔH vap of acetone and diethyl ether, even though their molar masses are much higher than water? (see table 11.7) If the highest energy molecules move to the gas phase, what would happen to the average energy of the particles that are left behind? Solve Example 11.3 and For Practice 11.3. There is a video guide in the MC study area for chapter 11. Dynamic equilibrium is the first entry into this topic. Equilibrium helps describe why the liquid in an open jar will completely evaporate, but the liquid in a sealed jar will only partially evaporate. Figure 11.25 and 11.26 show that the rates of condensation and evaporation will eventually become the same in a sealed container. Why does the rate of condensation increase in these figures? Why does vapor pressure increase with an increase in temperature? It might be helpful to look at Figure 11.28 and back at 11.24. How is the normal boiling point defined? Why is the boiling point of water usually lower in Denver, CO than in SLO, CA? This is the best video I can find that describes the relationship between IMFs and boiling point. https://www.youtube.com/watch?v=08kggrqazxa IGNORE rest of this section on the Clausius-Clapeyron equation. Section 11.6: Can a solid turn directly into a gas? Define the terms sublimation, deposition, and fusion. Why does the temperature stay constant during melting as shown in Figure 11.34?
5 Why is the heat of fusion so much less than the heat of vaporization as shown in Figure 11.35? Section 11.7: How much heat is needed to melt ice and turn it to steam? Follow through the calculations shown in this section so that you can solve problem #83. Answer Conceptual Connection 11.6. Here is a video for explaining the heating curve: https://www.youtube.com/watch?v=6atx_k19css And another one for calculations: https://www.youtube.com/watch?v=ipcbwxgmnfm Section 11.8: What is a phase diagram? Phase diagrams show the temperature and pressure dependence of solid liquid gas transitions. Note the slope of the nearly vertical line between solid and liquid in water (Figure 11.38) and iodine or CO 2 in Figure 11.39. What is the triple point on the phase diagram? What is the critical point on the phase diagram? What changes can happen to cause a liquid to turn to a gas, or to a solid? The rest of this chapter is really cool and discusses amazing things like semiconductors and LEDs. Unit cells are interesting. We will not cover any of these topics in this class.