1. © 2012 Pearson Education, Inc. Chapter 11 Liquids and Intermolecular Forces

2. Intermolecular Forces © 2012 Pearson Education, Inc. States of Matter The fundamental difference between states of matter is the distance between particles.

3. Intermolecular Forces © 2012 Pearson Education, Inc. A.Closer to the density of a gas B.Closer to the density of a solid C.Both are somewhat close to one another in density. D.They are significantly different in density.

4. Intermolecular Forces © 2012 Pearson Education, Inc. States of Matter Because in the solid and liquid states particles are closer together, we refer to them as condensed phases.

5. Intermolecular Forces © 2012 Pearson Education, Inc. The States of Matter The state a substance is in at a particular temperature and pressure depends on two antagonistic entities: –The kinetic energy of the particles. –The strength of the attractions between the particles.

6. Intermolecular Forces © 2012 Pearson Education, Inc. Intermolecular Forces The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together. These intermolecular attractions are, however, strong enough to control physical properties, such as boiling and melting points, vapor pressures, and viscosities. These intermolecular forces as a group are referred to as van der Waals forces.

7. Intermolecular Forces © 2012 Pearson Education, Inc. van der Waals Forces Dipole–dipole interactions Hydrogen bonding London dispersion forces

8. Intermolecular Forces © 2012 Pearson Education, Inc. London Dispersion Forces While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.

9. Intermolecular Forces © 2012 Pearson Education, Inc. London Dispersion Forces At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.

10. Intermolecular Forces © 2012 Pearson Education, Inc. London Dispersion Forces Another helium atom nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1. London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole. These forces are present in all molecules, whether they are polar or nonpolar. The tendency of an electron cloud to distort in this way is called polarizability.

11. Intermolecular Forces © 2012 Pearson Education, Inc. Factors Affecting London Forces The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane) tend to have stronger dispersion forces than short, fat ones (like neopentane). This is due to the increased surface area in n-pentane.

12. Intermolecular Forces © 2012 Pearson Education, Inc. Factors Affecting London Forces The strength of dispersion forces tends to increase with increased molecular weight. Larger atoms have larger electron clouds that are easier to polarize.

13. Intermolecular Forces © 2012 Pearson Education, Inc. A.CH 4 < CBr 4 < CCl 4 B.CCl 4 < CH 4 < CBr 4 C.CH 4 < CCl 4 < CBr 4 D.CBr 4 < CCl 4 < CH 4

14. Intermolecular Forces © 2012 Pearson Education, Inc. Dipole–Dipole Interactions Molecules that have permanent dipoles are attracted to each other. –The positive end of one is attracted to the negative end of the other, and vice versa. –These forces are only important when the molecules are close to each other.

15. Intermolecular Forces © 2012 Pearson Education, Inc. Dipole–Dipole Interactions The more polar the molecule, the higher its boiling point.

16. Intermolecular Forces © 2012 Pearson Education, Inc. Which Have a Greater Effect? Dipole–Dipole Interactions or Dispersion Forces If two molecules are of comparable size and shape, dipole–dipole interactions will likely be the dominating force. If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

17. Intermolecular Forces © 2012 Pearson Education, Inc. How Do We Explain This? The nonpolar series (SnH 4 to CH 4 ) follow the expected trend. The polar series follow the trend until you get to the smallest molecules in each group.

18. Intermolecular Forces © 2012 Pearson Education, Inc. A.SnH 4 is more polar than CH 4. B.SnH 4 is smaller in size than CH 4. C.SnH 4 has greater internal dispersion forces than in CH 4. D.SnH 4 is ionic in structure and CH 4 is molecular.

19. Intermolecular Forces © 2012 Pearson Education, Inc. Hydrogen Bonding The dipole–dipole interactions experienced when H is bonded to N, O, or F are unusually strong. We call these interactions hydrogen bonds.

20. Intermolecular Forces © 2012 Pearson Education, Inc. A.The non-hydrogen atom must have a nonbonding electron pair. B.The non-hydrogen atom must have low electronegativity. C.The non-hydrogen atom must have a large atomic size. D.The non-hydrogen atom must have a small electron affinity.

21. Intermolecular Forces © 2012 Pearson Education, Inc. Hydrogen Bonding Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

22. Intermolecular Forces © 2012 Pearson Education, Inc. Sample Exercise 11.1 Identifying Substances That Can Form Hydrogen Bonds In which of these substances is hydrogen bonding likely to play an important role in determining physical properties: methane (CH 4 ), hydrazine (H 2 NNH 2 ), methyl fluoride (CH 3 F), hydrogen sulfide (H 2 S)? Practice Exercise In which of these substances is significant hydrogen bonding possible: methylene chloride (CH 2 Cl 2 ), phosphine (PH 3 ), hydrogen peroxide (HOOH), acetone (CH 3 COCH 3 )?

23. Intermolecular Forces © 2012 Pearson Education, Inc. A.Approximately 90° B.Approximately 109° C.Approximately 120° D.Approximately 180°

24. Intermolecular Forces © 2012 Pearson Education, Inc. Ion–Dipole Interactions Ion–dipole interactions (a fourth type of force) are important in solutions of ions. The strength of these forces is what makes it possible for ionic substances to dissolve in polar solvents.

25. Intermolecular Forces © 2012 Pearson Education, Inc. A.CH 3 OH in water, because CH 3 OH is a strong electrolyte and forms ions. B.Ca(NO 3 ) 2 in water, because Ca(NO 3 ) 2 is a strong electrolyte and forms ions. C.CH 3 OH in water, because CH 3 OH is a weak electrolyte and forms ions. D.Ca(NO 3 ) 2 in water, because Ca(NO 3 ) 2 is a weak electrolyte and forms ions.

26. Intermolecular Forces © 2012 Pearson Education, Inc. Summarizing Intermolecular Forces

27. Intermolecular Forces © 2012 Pearson Education, Inc. Sample Exercise 11.2 Predicting Types and Relative Strengths of Intermolecular Attractions List the substances BaCl 2, H 2, CO, HF, and Ne in order of increasing boiling point. Practice Exercise Identify the intermolecular attractions present in the following substances, and (b) select the substance with the highest boiling point: CH 3 CH 3, CH 3 OH, and CH 3 CH 2 OH.

28. Intermolecular Forces © 2012 Pearson Education, Inc. Intermolecular Forces Affect Many Physical Properties The strength of the attractions between particles can greatly affect the properties of a substance or solution.

29. Intermolecular Forces © 2012 Pearson Education, Inc. Viscosity Resistance of a liquid to flow is called viscosity. It is related to the ease with which molecules can move past each other. Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

30. Intermolecular Forces © 2012 Pearson Education, Inc. Surface Tension Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

31. Intermolecular Forces © 2012 Pearson Education, Inc. Shape of Water MeniscusShape of Hg Meniscus A. Yes, Inverted U Yes, downward U B.No changeYes, downward U C. No changeNo change D. Inverted UNo change

32. Intermolecular Forces © 2012 Pearson Education, Inc. A.Viscosity increases as intermolecular forces increase while surface tension decreases. Both viscosity and surface tension increase with increasing temperature. B.Viscosity decreases as intermolecular forces increase while surface tension increases. Both viscosity and surface tension increase with decreasing temperature C.Both viscosity and surface tension increase as intermolecular forces increase and temperature decreases. D.Both viscosity and surface tension decrease as intermolecular forces increase and temperature increases.

33. Intermolecular Forces © 2012 Pearson Education, Inc. Phase Changes

34. Intermolecular Forces © 2012 Pearson Education, Inc. Energy Changes Associated with Changes of State The heat of fusion is the energy required to change a solid at its melting point to a liquid. The heat of vaporization is defined as the energy required to change a liquid at its boiling point to a gas. The heat of sublimation is defined as the energy required to change a solid directly to a gas.

35. Intermolecular Forces © 2012 Pearson Education, Inc. A. No, because we are not dealing with state functions. B. No, because we need heat of melting. C. Yes, ΔH sub = ΔH fus + ΔH vap D. Yes, ΔH sub = ΔH fus – Δh vap

36. Intermolecular Forces © 2012 Pearson Education, Inc. A.Melting (or fusion) and endothermic B.Melting (or fusion) and exothermic C.Freezing and endothermic D.Freezing and exothermic

37. Intermolecular Forces © 2012 Pearson Education, Inc. Energy Changes Associated with Changes of State The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other. The temperature of the substance does not rise during a phase change.

38. Intermolecular Forces © 2012 Pearson Education, Inc. Sample Exercise 11.3 Calculating  H for Temperature and Phase Changes Calculate the enthalpy change upon converting 1.00 mol of ice at  25  C to steam at 125  C under a constant pressure of 1 atm. The specific heats of ice, liquid water, and steam are 2.03 J/g-K, 4.18 J/g-K, and 1.84 J/g-K, respectively. For H 2 O,  H fus = 6.01 kJ/mol and  H vap = 40.67 kJ/mol. Practice Exercise What is the enthalpy change during the process in which 100.0 g of water at 50.0  C is cooled to ice at  30.0  C? (Use the specific heats and enthalpies for phase changes given in Sample Exercise 11.3.)

39. Intermolecular Forces © 2012 Pearson Education, Inc. Vapor Pressure At any temperature some molecules in a liquid have enough energy to break free. As the temperature rises, the fraction of molecules that have enough energy to break free increases.

40. Intermolecular Forces © 2012 Pearson Education, Inc. A.Increases B.Decreases

41. Intermolecular Forces © 2012 Pearson Education, Inc. Vapor Pressure As more molecules escape the liquid, the pressure they exert increases. The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.

42. Intermolecular Forces © 2012 Pearson Education, Inc. A.CBr 4 because dispersion forces between its molecules are greater than in CCl 4. B.CBr 4 because polar forces between its molecules are smaller than in CCl 4. C.CCl 4 because polar forces between its molecules are greater than in CBr 4. D.CCl 4 because dispersion forces between its molecules are smaller than in CBr 4.

43. Intermolecular Forces © 2012 Pearson Education, Inc. Vapor Pressure The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure. The normal boiling point is the temperature at which its vapor pressure is 760 torr.

44. Intermolecular Forces © 2012 Pearson Education, Inc. A.260 torr B.460 torr C.660 torr D.760 torr

45. Intermolecular Forces © 2012 Pearson Education, Inc. Vapor Pressure The natural log of the vapor pressure of a liquid is inversely proportional to its temperature. This relationship is quantified in the Clausius–Clapeyron equation: ln P = −  H vap /RT + C, where C is a constant.

46. Intermolecular Forces © 2012 Pearson Education, Inc. Sample Exercise 11.4 Relating Boiling Point to Vapor Pressure Use Figure 11.25 to estimate the boiling point of diethyl ether under an external pressure of 0.80 atm. Practice Exercise At what external pressure will ethanol have a boiling point of 60  C?

47. Intermolecular Forces © 2012 Pearson Education, Inc. Phase Diagrams Phase diagrams display the state of a substance at various pressures and temperatures, and the places where equilibria exist between phases. The liquid–vapor interface starts at the triple point (T), at which all three states are in equilibrium, and ends at the critical point (C), above which the liquid and vapor are indistinguishable from each other. Each point along this line is the boiling point of the substance at that pressure. The interface between liquid and solid marks the melting point of a substance at each pressure. Below the triple point the substance cannot exist in the liquid state. Along the solid–gas line those two phases are in equilibrium; the sublimation point at each pressure is along this line.

48. Intermolecular Forces © 2012 Pearson Education, Inc. A.Freezing B.Melting C.Vaporization D. Condensation

49. Intermolecular Forces © 2012 Pearson Education, Inc. Phase Diagram of Water Note the high critical temperature and critical pressure. –These are due to the strong van der Waals forces between water molecules. The slope of the solid– liquid line is negative. –This means that as the pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid.

50. Intermolecular Forces © 2012 Pearson Education, Inc. Phase Diagram of Carbon Dioxide Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO 2 sublimes at normal pressures.

51. Intermolecular Forces © 2012 Pearson Education, Inc. Sample Exercise 11.5 Interpreting a Phase Diagram Use the phase diagram for methane, CH 4, shown in Figure 11.30 to answer the following questions. (a) What are the approximate temperature and pressure of the critical point? (b) What are the approximate temperature and pressure of the triple point? (c) Is methane a solid, liquid, or gas at 1 atm and 0  C? (d) If solid methane at 1 atm is heated while the pressure is held constant, will it melt or sublime? (e) If methane at 1 atm and 0  C is compressed until a phase change occurs, in which state is the methane when the compression is complete? Practice Exercise Use the phase diagram of methane to answer the following questions. (a) What is the normal boiling point of methane? (b) Over what pressure range does solid methane sublime? (c) Liquid methane does not exist above what temperature?

52. Intermolecular Forces © 2012 Pearson Education, Inc. Liquid Crystals Some substances do not go directly from the solid state to the liquid state. In this intermediate state, liquid crystals have some traits of solids and some of liquids.

53. Intermolecular Forces © 2012 Pearson Education, Inc. Liquid Crystals Unlike liquids, molecules in liquid crystals have some degree of order. In nematic liquid crystals, molecules are only ordered in one dimension, along the long axis. In smectic liquid crystals, molecules are ordered in two dimensions, along the long axis and in layers.

54. Intermolecular Forces © 2012 Pearson Education, Inc. Liquid Crystals In cholesteryl liquid crystals, nematic-like crystals are layered at angles to each other.

55. Intermolecular Forces © 2012 Pearson Education, Inc. Sample Exercise 11.6 Properties of Liquid Crystals Which of these substances is most likely to exhibit liquid crystalline behavior? Practice Exercise Suggest a reason why decane does not exhibit liquid crystalline behavior. CH 3 CH 2 CH 2 CH 2 CH 2 CH 2 CH 2 CH 2 CH 2 CH 3

56. Intermolecular Forces © 2012 Pearson Education, Inc. Liquid Crystals These crystals can exhibit color changes with changes in temperature.