1. Chapter 11 Intermolecular Forces, Liquids and Solids
CHEMISTRY The Central Science 9th Edition
Chapter 11 Intermolecular Forces, Liquids and Solids

2. 11.1: A Molecular Comparison of Liquids and Solids

3. Text, P. 409
The forces holding solids and liquids together are called intermolecular forces

4. 11.2: Intermolecular Forces
The covalent bond holding a molecule together is an intramolecular force
The attraction between molecules is an intermolecular force
Much weaker than intramolecular forces
Melting or boiling: the intermolecular forces are broken (not the covalent bonds)

5. The stronger the attractive forces, the higher the boiling point of the liquid and the melting point of a solid
Text, P. 409
(low boiling point)

6. Strongest of all intermolecular forces Found only in mixtures
Ion-Dipole Forces
Interaction between an ion and a dipole (a polar molecule such as water)
Strongest of all intermolecular forces
Found only in mixtures
Text, P. 410

7. Between neutral polar molecules
Dipole-Dipole Forces
Between neutral polar molecules
Oppositely charged ends of molecules attract
Weaker than ion-dipole forces
Dipole-dipole forces increase with increasing polarity
Strength of attractive forces is inversely related to molecular volume
Text, P. 410

8. London Dispersion Forces
Weakest of all intermolecular forces
Two adjacent neutral, nonpolar molecules
The nucleus of one attracts the electrons of the other
Electron clouds are distorted
Instantaneous dipole
Strength of forces is directly related to molecular weight
London dispersion forces exist between all molecules

9. London dispersion forces depend on the shape of the molecule
The greater the surface area available for contact, the greater the dispersion forces
Text, P. 412

10. Hydrogen Bonding
Special case of dipole-dipole forces
H-bonding requires H bonded to an electronegative element (most important for compounds of F, O, and N)

11. Hydrogen Bonding Text, P. 413
Boiling point increases with increasing molecular weight. The exception is water (H bonding)

12. Hydrogen Bonding
Text, P. 414

13. Text, P. 415
Solids are usually more closely packed than liquids (solids are more dense than liquids)
Ice is ordered with an open structure to optimize H-bonding (ice is less dense than water)

14. Text, P. 417

15. Sample Problems # 7, 9, 11, 13, 15, 17, 19

16. 11.3: Some Properties of Liquids
Viscosity
Viscosity is the resistance of a liquid to flow
Molecules slide over each other
The stronger the intermolecular forces, the higher the viscosity
Viscosity increases with an increase in molecular weight

17. Surface Tension
Surface molecules are only attracted inwards towards the bulk molecules
Molecules within the liquid are all equally attracted to each other

18. Surface tension is the amount of energy required to increase the surface area of a liquid
Cohesive forces bind molecules to each other (Hg)
Adhesive forces bind molecules to a surface (H2O)
If adhesive forces > cohesive forces, the meniscus is U-shaped (water in a glass)
If cohesive forces > adhesive forces, the meniscus is curved downwards (Hg in a barometer)

19. 11.4: Phase Changes Text, P. 420 (Exothermic) (Endothermic)

20. Generally heat of fusion (melting) is less than heat of vaporization (evaporation)
It takes more energy to completely separate molecules than to partially separate them
Text, P. 420

21. Plot of temperature change versus heat added is a heating curve
Heating Curves
Plot of temperature change versus heat added is a heating curve
During a phase change, adding heat causes no temperature change (equilibrium is established)
These points are used to calculate Hfus and Hvap
Remember: Q = m·Cp·ΔT

22. Added heat increases the temperature of a consistent state of matter
Text, P. 421
Added heat increases the temperature of a consistent state of matter
Energy used for changing molecular motion, no T change

23. Critical Temperature and Pressure
Gases are liquefied by increasing pressure at some temperature
Critical temperature: the maximum temperature for liquefaction of a gas using pressure
A high C.T. means strong intermolecular forces
Critical pressure: pressure required for liquefaction

24. Examples: # 31, 33, WDP # 48
Other WDP examples: # 44, 46, 50 and 51

25. 11.5: Vapor Pressure Explaining Vapor Pressure on the Molecular Level
Some of the molecules on the surface of a liquid have enough energy to escape to the gas phase
After some time the pressure of the gas will be constant at the vapor pressure (equilibrium is established)

26. Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface
Vapor pressure is the pressure exerted when the liquid and vapor are in dynamic equilibrium
Volatility, Vapor Pressure, and Temperature
If equilibrium is never established then the liquid evaporates
Volatile substances (high VP) evaporate rapidly
The higher the T, the higher the average KE, the faster the liquid evaporates (hot water evaporates faster than cold water)

27. Vapor pressure increases nonlinearly with increasing temperature
Text, P. 426

28. T is the Kelvin temperature R is the gas constant,
When Temperature changes from T1 to T2, Vapor Pressure changes from P1 to P2
These changes are related to ΔH by the equation,
Where
T is the Kelvin temperature
R is the gas constant,
ΔHvap is the molar heat of vaporization

29. This comes from the graph of P vs. inverse of T
Straight line
Negative slope
Equation:
C is a constant
Use the Clausius-Clapeyron Equation to
Predict the vapor pressure at a specified temperature
Determine the T at which a liquid has a specified VP
Calculate enthalpy of vaporization from measurements of VP’s at different temperatures

30. Vapor Pressure and Boiling Point
Liquids boil when the external pressure equals the vapor pressure
Normal BP: BP of a liquid at 1 atmosphere
Temperature of boiling point increases as pressure increases

31. The vapor pressure of water is 1
The vapor pressure of water is 1.0 atm at 373 K, and the enthalpy of vaporization is 40.7 kJ mol-1. Estimate the vapor pressure at temperature 363 and 383 K respectively.
Solution
where R (= J mol-1 K-1)
Using the Clausius-Clapeyron equation, we have:
P363 = 1.0 exp (- (40700/8.3145)(1/ /373)
= atm
P383 = 1.0 exp (- (40700/8.3145)(1/ /373)
= atm
Note that the increase in vapor pressure from 363 K to 373 K is atm, but the increase from 373 to 383 K is atm. The increase in vapor pressure is not a linear process.

32. Sample problems: # 45, WDP # 35
Other WDP examples: # 36 & 37

33. 11.6: Phase Diagrams
Phase diagram: plot of pressure vs. Temperature summarizing all equilibria between phases
Given a temperature and pressure, phase diagrams tell us which phase will exist

34. Vapor Pressure curve of the liquid (increase P, increase T)
Text, P. 428
Melting point curve: Increased P favors solid phase; Higher T needed to melt the solid at higher P
Beyond this point, liquid and gas phases are indistinguishable
Vapor Pressure curve of the liquid (increase P, increase T)
Stable at low T and high P
Stable at low P and high T
Triple Point: all 3 phases in equilibrium

35. The Phase Diagrams of H2O and CO2 Text, P. 429
Line slopes to the left: ice is less dense than water (why?) MP decreases with increased P
Text, P. 429
The Phase Diagrams
of H2O and CO2

36. Sample Problems: #49, 51

37. 11.7: Structures of Solids Unit Cells
Crystalline solid: well-ordered, definite arrangements of molecules, atoms or ions
The smallest repeating unit in a crystal is a unit cell
It has all the symmetry of the entire crystal
Three-dimensional stacking of unit cells is the crystal lattice
Close-packed structure

39. Unit Cells
Text, P. 431

40. Unit Cells
Text, P. 432
Primitive cubic: atoms at the corners of a simple cube
each atom shared by 8 unit cells

41. Unit Cells
Body-centered cubic (bcc): atoms at the corners of a cube plus one in the center of the body of the cube
corner atoms shared by 8 unit cells
center atom completely enclosed in 1 unit cell
Text, P. 432

42. Unit Cells
Face-centered cubic (fcc): atoms at the corners of a cube plus one atom in the center of each face of the cube
corner atoms shared by 8 unit cells
face atoms shared by 2 unit cells
Text, P. 432

43. Unit Cells 1 atom per cell 2 atoms per cell 4 atoms per cell
Text, P. 432

44. The Crystal Structure of Sodium Chloride
Text, P. 433
Two equivalent ways of defining unit cell:
Cl- (larger) ions at the corners of the cell, or
Na+ (smaller) ions at the corners of the cell

45. http://www. nytimes. com/2006/06/13/science/13find. html

46. Text, P. 435
11.8: Bonding in Solids

47. Covalent-Network Solids
Text, P. 437

48. Ionic Solids
The structure adopted depends on the charges and sizes of the ions
Text, P. 438

49. Various arrangements are possible
Metallic Solids
Various arrangements are possible
The bonding is too strong for London dispersion and there are not enough electrons for covalent bonds
The metal nuclei float in a sea of electrons
Metals conduct because the electrons are delocalized and are mobile
Close-packed structure
Text, P. 440

50. Amorphous solids (rubber, glass) have no orderly structure
IMFs vary in strength throughout the sample
No specific melting point
Sample Problems # 53, 69, 71, 73, 75

51. End of Chapter 11 Intermolecular Forces, Liquids and Solids