Liquids, solids, & intermolecular forces

1. Liquids, solids, & intermolecular forces Chapter 11

2. KMT meets liquids • Ideal gas is a gas even at absolute zero • Real gas condenses to liquid at low T/high P • Attractive forces exist between real gas molecules

3. Intermolecular attractions • Attractive forces exist between all atoms/molecules • Relative strength of attractions indicated by • Boiling point (higher b.p. = stronger attractions) • Vapor pressure (high v.p. = weaker attractions) • ∆Hvaporization (large ∆Hvap = stronger attractions)

4. Instantaneous or momentary dipoles • e– distribution is asymmetric –– just for a moment • Atom/molecule is polar –– just for a moment

5. Induced dipoles • Momentary dipole in one atom induces a dipole in a neighboring atom . . . which induces a dipole in another neighboring atom, and so on, causing a little ripple of dipoles

6. Dispersion force • Taken together, instantaneous & induced dipoles create an attractive force between molecules, called the dispersion force • Each dipole is tiny, but the constant ripple of countless dipoles throughout the substance makes this the primary attractive force between molecules • Even noble gas atoms show dispersion force between atoms

7. Polarizability • Magnitude of dispersion force depends on polarizability • Larger e– cloud = more polarizable • Dispersion force increases with increasing molar mass • Melting and boiling points of molecular substances generally increase as molar mass increases

8. Molar mass & boiling point • For compounds of similar structure, boiling point increases as molar mass increases

9. Polarizability • Polarizability is greater in elongated molecules than in compact ones of similar mass

10. Permanent dipoles • Polar molecules tend to arrange themselves +/– to maximize attractions • Extra ordering increases tendency to stick together in liquid state • Boiling point of a polar substance is higher than that of a nonpolar substance of similar mass.

11. Nonpolar/polar • Molecules have similar masses • Permanent dipoles increase b.p.

12. The van der Waals forces • Together, dispersion and pemanent dipole forces are known as the van der Waals forces • When comparing substances of comparable mass (±10%), the presence of a permanent dipole increases boiling point significantly • When comparing substances of different molar masses, the dispersion force (related to mass) is more important than the permanent dipole

13. Examples • Which would you expect to have the highest boiling point, and why: C3H8, CO2, CH3CN

14. Examples • Which would you expect to have the highest boiling point, and why: C3H8, CO2, CH3CN • masses similar (C3H8 = 44, CO2 = 44, CH3CN = 41) • CH3CN polar = highest bp • Actual values: C3H8 = 231K, CO2 = 195K, CH3CN =

15. Examples • Arrange these in order of increasing boiling point: Ne, He, Cl2, (CH3)2CO, O2, O3

16. Examples • Arrange these in order of increasing boiling point: Ne, He, Cl2, (CH3)2CO, O2, O3 • masses: Ne = 20, He = 4, Cl2 = 71, (CH3)2CO = 58, O2 = 32, O3 = 48 • Ordered by mass: He, Ne, O2, O3,(CH3)2CO, Cl2 • (CH3)2CO is polar & has large surface area = higher bp • Predict He, Ne, O2, O3,Cl2, (CH3)2CO • Actual values: He = 4K, Ne = 27K, O2 = 90K, O3 = 161K,Cl2 = 238K, (CH3)2CO = 329K

17. Then there’s hydrogen . . .

18. O–H bond is very polar, and atoms are very small • Dipoles are close together, so their attraction is very strong • H atom is covalently bonded to its own O and weakly bonded (dotted line) to the neighboring O • Weak bond to neighboring O is a hydrogen bond

19. Hydrogen bonding • Hydrogen bonding occurs only between molecules containing N–H, O–H, and F–H bonds • Hydrogen bonding is much stronger than ordinary dispersion/dipole → much higher boiling points than expected for their mass • Hydrogen bonds are not as strong as covalent bonds (15-40 kJ/mol, vs >150 kJ/mol)

20. Intermolecular forces

21. Substances that are not molecular • Ionic substances • Held together by lattice energy • Generally high mp & bp • Metallic substances • Metal cations in sea of electrons • Generally high mp & bp • Network covalent solids (e.g. diamond) • Melting = disrupt covalent bonds • VERY high mp & bp

22. Vaporization • At liquid surface, faster molecules have enough kinetic energy to escape (vaporize or evaporate) • As higher-energy molecules leave the liquid, average kinetic energy of the liquid decreases • Temperature of liquid decreases (evaporative cooling)

23. Vaporization • For liquid temperature to remain constant during evaporation, liquid must absorb energy from surroundings • Amount of energy liquid must absorb to keep temperature constant during evaporation = enthalpy (heat) of vaporization(∆Hvaporization) • Vaporization is endothermic, so ∆Hvap is positive

24. Example • How much energy is required to vaporize 2.35 g of diethyl ether, (C2H5)2O, at 298 K? ∆Hvap for diethyl ether at 298 K is 29.1 kJ/mol.

25. Liquid-vapor equilibrium When rate of vaporization = rate of condensation in a closed sysem, system has reached equilibrium

26. Vapor Pressure • Pressure exerted by vapor in dynamic equilibrium w its liquid = vapor pressure of that liquid • Vapor pressure depends only on type of liquid & temperature • As long as both phases are present, amount of liquid in container does not affect vapor pressure • Liquids with high vapor pressure at room temperature are volatile (evaporate easily)

27. Vapor pressure curves Vapor pressure always increases as temperature increases

28. Vapor pressure and boiling • In open container, evaporation occurs only at surface • As temperature increases, evaporation increases • At some point, evaporation begins to occur throughout the liquid instead of just at the surface: boiling!

29. Vapor pressure & boiling • Vapor bubbles form throughout liquid • Bubbles rise to surface, burst, release vapor • All energy is used to convert liquid to vapor, so temperature remains constant while liquid boils

30. Boiling point • Boiling begins when the liquid’s vapor pressure matches the external pressure of the atmosphere • The temperature at which this occurs is the boiling point • When the external atmospheric pressure = 1 atm, the boiling point is called the normal boiling point

31. The critical point • Liquid heated in a rigid sealed container does not boil • Vapor pressure and vapor density increase • Liquid density decreases • Vapor & liquid densities become equal & meniscus disappears • This point is called the critical point

32. The critical point

33. Vapor pressure and temperature • Clausius-Clapeyron equation shows relationship between vapor pressure and temperature

34. Clausius-Clapeyron equation • P (vapor pressure) can be in any unit • R must be 8.3145 J/mol K • ∆Hvaporization is usually given in kJ/mol but must be converted to J/mol to agree with R • T is in Kelvins (duh)

35. Example • The vapor pressure of methanol is 100 mm Hg at 21.2 °C. What is its vapor pressure at 25.0 °C? ∆Hvap for methanol is 38.0 kJ/mol.

36. Example • The normal boiling point of isooctane is 99.2 °C and its ∆Hvap is 35.76 kJ/mol. What is the vapor pressure of isooctane at 25.0 °C?

37. Clausius-Clapeyron equation • Plot of ln P vs 1/T gives straight line w slope –∆Hvap/R

38. Changes of state • Liquid ↔ gas • Vaporization/boiling and condensation • Solid ↔ liquid • Melting (fusion) and freezing • Solid ↔ gas • Sublimation and deposition

39. Temperature Add energy Heating curve

40. Temperature Add energy Heating curve (g) boiling boiling point condensing (l) melting melting/freezing point freezing (s)

42. Part of a cooling curve for water • The dotted line shows supercooling • The water remains liquid below 0 °C • At the bottom of the dotted line, crystallization begins • Crystallization releases energy; temperature returns to freezing temperature • Temperature remains constant until freezing is completed

43. Phase diagram • A graphical representation of the conditions of temperature & pressure under which various phases of a substance exist

44. A phase diagram for iodine

45. A phase diagram for carbon dioxide

46. A phase diagram for water

47. Types of solids

48. Molecular substances • Molecular solids held together by • Dispersion • Dipole • Hydrogen bonding • Relatively low mp & bp • For molecules of similar structure, boiling point increases as molar mass increases

49. Ionic substances • Ions held together by lattice forces • Coulomb’s law: • Attraction of oppositely charged ions increases with increased charge and/or decreased ion size • Which has a higher mp, NaF or MgO? • NaF mp 993 °C, MgO mp 2852 °C • NaCl or KI? • NaCl mp 801 °C, KI mp 681 °C

50. Atomic substances • Noble gas atoms held together only by dispersion forces • Metals atoms held together by metal cations in sea of electrons