1. Attractive Forces
The particles are attracted to each other by electrostatic forces.
The strength of the attractive forces depends on the kind(s) of particles.
The stronger the attractive forces between the particles, the more they resist moving.

2. Intermolecular Attractions
The strength of the attractions between the particles of a substance determines its state.
The stronger the attractive forces are, the higher will be the boiling point of the liquid and the melting point of the solid.

3. Why Are Molecules Attracted to Each Other?
Attractive forces between opposite charges.
+ ion to − ion
+ end of polar molecule to − end of polar molecule
H-bonding especially strong
Even nonpolar molecules will have temporary charges
Coulomb’s Law!
Larger charge = stronger attraction
Longer distance = weaker attraction
Intermolecular forces are weaker than bonding (intramolecular) forces

4. Kinds of Attractive Forces
dispersion forces.
dipole–dipole attractions.
hydrogen bonds.

5. Dispersion Forces
Changes in electron distribution in atoms and molecules result in a temporary dipole (partial charges).
The attractive forces caused by these temporary dipoles are called dispersion forces.
Aka London Forces
All molecules and atoms will have them.
As a temporary dipole is established in one molecule, it induces a dipole in all the surrounding molecules.

6. Dispersion Force

7. Size of the Induced Dipole
The magnitude of the induced dipole depends on several factors.
Polarizability of the electrons
Volume of the electron cloud
Larger molar mass = more electrons = larger electron cloud = increased polarizability = stronger attractions
Shape of the molecule
More surface-to-surface contact = larger induced dipole = stronger attraction

8. Effect of Molecular Size on Size of Dispersion Force

9. Effect of Molecular Shape on Size of Dispersion Force

10. Rank the following from least to greatest dispersion force attraction:
Bell Ringer 11/4
Rank the following from least to greatest dispersion force attraction:
Bromine
Chlorine
Iodine
Fluorine
Look at table 8.3 on page 368. How does the strength of the dispersion force affect the trends in melting and boiling point?
Larger molar mass = more electrons = larger electron cloud = increased polarizability = stronger attractions
Fluorine<Chlorine<Bromine<Iodine

11. Dipole–Dipole Attractions
Polar molecules have a permanent dipole.
Bond polarity and shape
Dipole moment
The always present induced dipole
The permanent dipole adds to the attractive forces between the molecules, raising the boiling and melting points relative to nonpolar molecules of similar size and shape.

13. Effect of Dipole–Dipole Attraction on Boiling and Melting Points

14. Attractive Forces and Solubility
Solubility depends, in part, on the attractive forces of the solute and solvent molecules.
Like dissolves like
Miscible liquids will always dissolve in each other
Polar substances dissolve in polar solvents.
Hydrophilic groups = OH, CHO, C═O, COOH, NH2, Cl
Nonpolar molecules dissolve in nonpolar solvents.
Hydrophobic groups = C—H, C—C

15. Immiscible Liquids
Pentane, C5H12 is a nonpolar molecule. Water is a polar molecule. The attractive forces between the water molecules is much stronger than their attractions for the pentane molecules. The result is that the liquids are immiscible.

16. Hydrogen Bonding
When a very electronegative atom is bonded to hydrogen, it strongly pulls the bonding electrons toward it.
O─H, N─H, or F─H
Because hydrogen has no other electrons, when its electron is pulled away, the nucleus becomes de-shielded, exposing the H proton.
The exposed proton acts as a very strong center of positive charge, attracting all the electron clouds from neighboring molecules.

17. Hydrogen Bonding

18. Stronger than dipole–dipole or dispersion forces
H–Bonds
Stronger than dipole–dipole or dispersion forces
have higher boiling points and melting points than similar substances that cannot.
not nearly as strong as chemical bonds.
2–5% the strength of covalent bonds

19. Ion–Dipole Attraction
In a mixture, ions from an ionic compound are attracted to the dipole of polar molecules.
The strength of the ion–dipole attraction is one of the main factors that determines the solubility of ionic compounds in water.

21. Example Problems
Determine the kinds of intermolecular forces that are present:
Cl2
London dispersion CO
London dispersion, dipole-dipole
SO2
CH2Cl2 HF
London dispersion, hydrogen bonding
NaBr dissolved in water
London dispersion, ion-dipole

22. Surface Tension
The layer of molecules on the surface behave differently than the interior, because the cohesive forces on the surface molecules have a net pull into the liquid interior.

23. Factors Affecting Surface Tension
stronger the intermolecular forces=higher surface tension
Raising the temperature =lower surface tension.
increases KE of the molecules.
The increased molecular motion makes it easier to stretch the surface.

24. Factors Affecting Viscosity
Viscosity is the resistance of a liquid to flow.
stronger the intermolecular force=higher viscosity
spherical the molecular shape=lower viscosity
Molecules roll more easily.
Less surface-to-surface contact lowers attractions.
Higher temperature =lower viscosity

25. Phase Changes

26. Vaporization
If these high energy molecules are at the surface, they may have enough energy to overcome the attractive forces.
Therefore, the larger the surface area, the faster the rate of evaporation.
This will allow them to escape the liquid and become a vapor.

27. Distribution of Thermal Energy
Only a small fraction of the molecules in a liquid have enough energy to escape.
But, as the temperature increases, the fraction of the molecules with “escape energy” increases.
The higher the temperature, the faster the rate of evaporation.

28. Condensation
Some molecules of the vapor will lose energy through molecular collisions.
The result will be that some of the molecules will get captured back into the liquid when they collide with it.
Also some may stick and gather together to form droplets of liquid, particularly on surrounding surfaces.
We call this process condensation.

29. Evaporation versus Condensation
Vaporization and condensation are opposite processes.
In an open container, the vapor molecules generally spread out faster than they can condense.
The net result is that the rate of vaporization is greater than the rate of condensation, and there is a net loss of liquid.
However, in a closed container, the vapor is not allowed to spread out indefinitely.
The net result in a closed container is that at some time the rates of vaporization and condensation will be equal.

30. Bell Ringer 11/5
No Bell Ringer!

31. Effect of Intermolecular Attraction on Evaporation and Condensation
the weaker the attractive forces, the faster the rate of evaporation.
Liquids that evaporate easily are said to be volatile.
Liquids that do not evaporate easily are called nonvolatile.
For example, motor oil

32. Energetics of Vaporization
vaporization is an endothermic process.
Condensation is an exothermic process.
Vaporization requires input of energy to overcome the attractions between molecules.

33. Heat of Vaporization
The amount of heat energy required to vaporize one mole of the liquid is called the heat of vaporization, DHvap.
Sometimes called the enthalpy of vaporization
It is always endothermic; therefore, DHvap is +.
It is somewhat temperature dependent.
DHcondensation = −Dhvaporization
Units are kJ/mol

34. Dynamic Equilibrium rate of vaporization=rate of condensation
When two opposite processes reach the same rate so that there is no gain or loss of material, we call it a dynamic equilibrium.
This does not mean there are equal amounts of vapor and liquid; it means that they are changing by equal amounts.

35. Dynamic Equilibrium

36. as the temperature increases, the vapor pressure increases.
The pressure exerted by the vapor when it is in dynamic equilibrium with its liquid is called the vapor pressure.
Therefore, the weaker the attractive forces, the higher the vapor pressure.
The higher the vapor pressure, the more volatile the liquid.
as the temperature increases, the vapor pressure increases.

37. Vapor Pressure Curves

38. Dynamic Equilibrium
A system in dynamic equilibrium can respond to changes in the conditions.
When conditions change, the system shifts its position to relieve or reduce the effects of the change.

39. Changing the Container’s Volume Disturbs the Equilibrium
Initially, the rate of vaporization and condensation are equal and the system is in dynamic equilibrium.
When the volume is increased, the rate of vaporization becomes faster than the rate of condensation.
When the volume is decreased, the rate of vaporization becomes slower than the rate of condensation.

40. Boiling Point
temperature at which the vapor pressure equals external pressure is the boiling point.

41. Heating Curve of a Liquid
As you heat a liquid, its temperature increases linearly until it reaches the boiling point.
Once the temperature reaches the boiling point, all the added heat goes into boiling the liquid; the temperature stays constant.
Once all the liquid has been turned into gas, the temperature can again start to rise.

42. Heating Curves
While heating a substance in one phase, the temperature changes. In order to calculate the amount of heat needed to change the temperature substance, use:
q = mass × Cs × DT
While in a phase change use heats of vaporization or fusion to calculate the heat needed for the phase change.

43. Energetics of phase changes
Suppose that 1.15 grams of rubbing alcohol (C2H8O) evaporates from a 65.0g aluminum block. If the Al block is initially at 25°C, what is the final temperature of the block after the vaporization of the alcohol? Assume that the only heat required for the vaporization comes from the Al block and that the alcohol vaporizes at 25°C.

44. Clausius–Clapeyron Equation
Relationship between vapor pressure and temperature.
A graph of ln(Pvap) versus 1/T is a straight line.
The slope of the line × J/mol ∙ K = Dhvap.
In J/mol

45. Clausius–Clapeyron Equation: Two-Point Form
Just two measurements of vapor pressure and temperature.
It can also be used to predict the vapor pressure if you know the heat of vaporization and the normal boiling point.
Remember, the vapor pressure at the normal boiling point is 760 torr.

46. Example: Clausius-Clapeyron
Benzene has a heat of vaporization of kJ/mol and a normal boiling point of 80.1°C. At what temperature does benzene boil when the external pressure is 445 torr?

47. Sublimation and Deposition
Molecules in the solid have thermal energy that allows them to vibrate.
Surface molecules with sufficient energy may break free from the surface and become a gas; this process is called sublimation.
The capturing of vapor molecules into a solid is called deposition.
The solid and vapor phases exist in dynamic equilibrium in a closed container at temperatures below the melting point.
Therefore, molecular solids have a vapor pressure.
sublimation
solid gas
deposition

48. Sublimation

49. Melting = Fusion
As a solid is heated, its temperature rises and the molecules vibrate more vigorously.
Once the temperature reaches the melting point, the molecules have sufficient energy to overcome some of the attractions that hold them in position and the solid melts (or fuses).
The opposite of melting is freezing.

50. Heating Curve of a Solid
As you heat a solid, its temperature increases linearly until it reaches the melting point.
q = mass × Cs × DT
Once the temperature reaches the melting point, all the added heat goes into melting the solid.
The temperature stays constant.
Once all the solid has been turned into liquid, the temperature can again start to rise.
Ice/water will always have a temperature of 0 °C at 1 atm.

51. Energetics of Melting
When the high energy molecules are lost from the solid, it lowers the average kinetic energy.
If energy is not drawn back into the solid its temperature will decrease; therefore, melting is an endothermic process, and freezing is an exothermic process.
Melting requires input of energy to overcome the attractions among molecules.

52. Heat of Fusion
The amount of heat energy required to melt one mole of the solid is called the heat of fusion, DHfus.
Sometimes called the enthalpy of fusion
It is always endothermic; therefore, DHfus is +.
It is somewhat temperature dependent.
DHcrystallization = −DHfusion
Generally much less than Dhvap
DHsublimation = DHfusion + DHvaporization

53. Heats of Fusion and Vaporization

54. Heating Curve of Water

55. Segment 1
Heating 1.00 mole of ice at −25.0 °C up to the melting point, 0.0 °C
q = mass × Cs × DT
Mass of 1.00 mole of ice = 18.0 g
Cs = 2.09 J/mol ∙ °C

56. Segment 2 Melting 1.00 mole of ice at the melting point, 0.0 °C
q = n ∙ DHfus
n = 1.00 mole of ice
DHfus = 6.02 kJ/mol

57. Segment 3
Heating 1.00 mole of water at 0.0 °C up to the boiling point, °C
q = mass × Cs × DT
Mass of 1.00 mole of water = 18.0 g
Cs = 2.09 J/mol ∙ °C

58. Segment 4 Boiling 1.00 mole of water at the boiling point, 100.0 °C
q = n ∙ DHvap
n = 1.00 mole of ice
DHfus = 40.7 kJ/mol

59. Segment 5 Heating 1.00 mole of steam at 100.0 °C up to 125.0 °C
q = mass × Cs × DT
Mass of 1.00 mole of water = 18.0 g
Cs = 2.01 J/mol ∙ °C

60. Water – An Extraordinary Substance
Water is a liquid at room temperature.
Most molecular substances with similar molar masses are gases at room temperature.
For example, NH3, CH4
This is due to H-bonding between molecules.
Water is an excellent solvent, dissolving many ionic and polar molecular substances.
It has a large dipole moment.
Even many small nonpolar molecules have some solubility in water.
For example, O2, CO2
Water has a very high specific heat for a molecular substance.
Moderating effect on coastal climates
Water expands when it freezes at a pressure of 1 atm.
About 9%
Making ice less dense than liquid water

61. Crystal Lattice
When allowed to cool slowly, the particles in a liquid will arrange themselves to give the maximum attractive forces.
Therefore, minimize the energy.
The result will generally be a crystalline solid.
The arrangement of the particles in a crystalline solid is called the crystal lattice.

62. Classifying Crystalline Solids
Molecular solids are solids whose composite particles are molecules.
Ionic solids are solids whose composite particles are ions.
Atomic solids are solids whose composite particles are atoms.
Nonbonding atomic solids are held together by dispersion forces.
Metallic atomic solids are held together by metallic bonds.
Network covalent atomic solids are held together by covalent bonds.

63. Types of Crystalline Solids
Stop here

64. The lattice sites are occupied by molecules.
Molecular Solids
The lattice sites are occupied by molecules.
CO2, H2O, C12H22O11
The molecules are held together by intermolecular attractive forces.
Dispersion forces, dipole–dipole attractions, and H-bonds
Because the attractive forces are weak, they tend to have low melting points.
Generally < 300 °C

65. Ionic Solids Lattice sites are occupied by ions.
They are held together by attractions between oppositely charged ions.
Nondirectional
Therefore, every cation attracts all anions around it, and vice versa.

66. Nonbonding Atomic Solids
Noble gases in solid form
Solid held together by weak dispersion forces
Very low melting
Tend to arrange atoms in closest-packed structure
Either hexagonal cP or cubic cP
Maximizes attractive forces and minimizes energy

67. Metallic Atomic Solids
Solid held together by metallic bonds
Strength varies with sizes and charges of cations
Coulombic attractions
Melting point varies
Mostly closest-packed arrangements of the lattice points
Cations

68. Closest-Packed Crystal Structures in Metals

69. Metallic Bonding
Metal atoms release their valence electrons.
Metal cation “islands” fixed in a “sea” of mobile electrons

70. Network Covalent Solids
Atoms attach to their nearest neighbors by covalent bonds.
Because of the directionality of the covalent bonds, these do not tend to form closest–packed arrangements in the crystal.
Because of the strength of the covalent bonds, these have very high melting points.
Generally > 1000 °C
Dimensionality of the network affects other physical properties.

71. The Diamond Structure: A Three-Dimensional Network
The carbon atoms in a diamond each have four covalent bonds to surrounding atoms.
sp3
Tetrahedral geometry
This effectively makes each crystal one giant molecule held together by covalent bonds.
You can follow a path of covalent bonds from any atom to every other atom.

72. Properties of Diamond Very high melting point, ~3800 °C Very rigid
Need to overcome some covalent bonds
Very rigid
Due to the directionality of the covalent bonds
Very hard
Due to the strong covalent bonds holding the atoms in position
Used as abrasives
Electrical insulator
Thermal conductor
Best known
Chemically very nonreactive

73. The Graphite Structure: A Two-Dimensional Network
In graphite, the carbon atoms in a sheet are covalently bonded together.
Forming six-member flat rings fused together
Similar to benzene
Bond length = 142 pm
sp2
Each C has three sigma and one pi bond.
Trigonal-planar geometry
Each sheet a giant molecule
The sheets are then stacked and held together by dispersion forces.
Sheets are 341 pm apart.

74. Properties of Graphite
Hexagonal crystals
High melting point, ~3800 °C
Need to overcome some covalent bonding
Slippery feel
Because there are only dispersion forces holding the sheets together, they can slide past each other.
Glide planes
Lubricants
Electrical conductor
Parallel to sheets
Thermal insulator
Chemically very nonreactive