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. Nonpolar vs. polar N2 NO O2
µ = 0 (nonpolar)
mass 28 g/mol
bp 77 K
µ = D (polar)
mass 30 g/mol
bp 121 K
mass 32 g/mol
bp 90 K
Based on mass alone, expect bp of NO to be between bp of N2 and bp of O2
Permanent dipole of NO makes it stickier so bp is higher than expected

13. 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

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

15. 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 =

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

17. 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

18. Then there’s hydrogen . . .

19. 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

20. 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)

21. Intermolecular forces
Dispersion force (momentary/induced dipoles)
Present in all atoms & molecules
bp ↑ as molar mass increases
Permanent dipole force
Present in polar molecules
bp ↑ compared to nonpolar molecule of similar mass
Hydrogen bonding
Present in molecules that contain N-H, O-H, or F-H bonds
bp ↑↑ than predicted from mass alone

22. Intermolecular forces

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

24. 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)

25. 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

26. 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.

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

28. 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)

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

30. 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!

31. 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

32. 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

33. 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

34. The critical point

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

36. Clausius-Clapeyron equation
P (vapor pressure) can be in any unit
R must be 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)

37. 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.

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

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

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

41. Heats of melting & freezing
Melting is endothermic
Energy to melt one mole of solid = heat of fusion (∆Hfusion)
∆Hfusion << ∆Hvaporization

42. Heating curve
Add energy
Temperature

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

45. 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

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

47. A phase diagram for iodine

48. A phase diagram for carbon dioxide

49. A phase diagram for water

50. Types of solids

51. 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

52. 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

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

54. Atomic substances
Atoms in network covalent solid held together by covalent bonds
Examples: C (subl 3652 °C), SiC (subl 2700 °C)