1. Liquids, Solids, and Intermolecular Forces
Chapter 11

2. Intermolecular Forces (IMF’s)
The state of a sample of matter (solid, liquid, gas) depends on the magnitude of intermolecular forces between the particles relative to the amount of thermal energy in the sample
Remember, thermal energy is kinetic (motion, vibrations, wiggling)

3. Review
Solids
Liquids
Gases

4. Liquids Properties of liquids
Liquids have high densities in comparison to gases
Liquids have indefinite shape and assume the shape of their container
Liquids have a definite volume and are not easily compressed

5. Solids Properties of solids
Solids have high densities in comparison to gases
Solids have a definite shape and do not assume the shape of their containers
Solids have a definite volume and are not easily compressed
Solids may be crystalline (ordered) or amorphous (disordered)

6. Properties of the Phases of Matter
A comparison of IMF’s
Properties of the Phases of Matter
Phase
Density
Shape
Volume
Strength of IMF’s
Gas
Low
Indefinite
Weak
Liquid
High
Definite
Moderate
Solid
STRONG

7. Q?? Why does pressure only affect the (l) and (g)??
Phase Changes
You can change the phase of a sample by changing the temp, pressure, or both
Add heat to a solid  liquid … add more heat (or increase the pressure)  gas
Take away heat from a gas (or decrease the pressure)  liquid … take away more  solid
Q?? Why does pressure only affect the (l) and (g)??

8. IMF’s
The strength of the intermolecular forces between molecules or atoms determine the phase (s, l, g)
Strong IMFs result in solids and liquids (high melting and boiling points)
Weak IMFs typically result in gases (low Tm & Tb)

9. IMFs vs Bonds
IMFs originate from the interactions between charges, partial charges, and temporary charges on molecules / atoms / ions
Similar concept as bonds but are MUCH WEAKER… they are not the bonds holding a compound / molecule together, they are the forces holding molecule A close to molecule B

10. Strengths Coulomb’s law: E = 1 * q1q2 4πɛo r
Bonds are strong because of large charged attractions at a close range
IMFs are much weaker due to weaker charges at longer distances

11. Water evaporates at 100⁰C
But, to break the O-H bond would take thousands of degrees

12. Different Types of IMFs
Dispersion Forces (London Force)
Dipole-Dipole Forces
Hydrogen Bonding
Ion-Dipole Forces
Three of these can potentially occur in all substances; H-Bonds can be found only in mixtures

13. #1: Dispersion Forces (London Forces)
Found in all molecules and atoms
Caused by the fluctuations in electrons around the atoms / within the molecules
Causes an instantaneous dipole (+ side / - side)
Think about your turn signal in your car…

14. Dispersion Forces
CAN INDUCE and stick to / attract to other neighboring molecules (IMF’s)

15. Magnitude of Dispersion Forces
The magnitude of the force depends on two things:
The size of the molecule (simply put: the Molar Mass) –the number of electrons
The shape / structure of the molecule

16. 1. Size
Typically, the larger the molecule, the more electrons and the more free “roaming” around they are able to do.
This opens the possibility of more electrons ‘clumping’ up and causing the +/-
As the electron cloud increases, the dispersion forces increase… causing a _________ B.P.

17. 2. Structure
The smaller the area of interaction, the lower the forces will be
The larger the area, the higher the forces

19. #2: Dipole-Dipole Forces
The dipole-dipole force exists in all molecules that are polar
Polar molecules have permanent dipoles that interact with perm dipoles of neighboring molecules
The (+) end of one molecule interacts with the (-) of a neighboring molecule

20. Polar Molecules Remember, all molecules have dispersion forces.
Polar molecules have Dipole-Dipole forces IN ADDITION to these dispersion forces
This raises their melting and boiling points
The strength (magnitude) of a dipole moment is measured in “Dipole moment” (D)
The greater the D, the larger the dipole-dipole forces

23. “LIKES DISSOLVE / MIX WITH LIKES”
Miscibility
The polarity of molecules also determines the miscibility of liquids
The ability to mix without separating into two phases
All polar liquids are miscible with other polar liquids
All nonpolar liquids are miscible with other nonpolar liquids
“LIKES DISSOLVE / MIX WITH LIKES”

24. Immiscibility Opposites generally do not mix together
i.e. nonpolar + polar

25. Practice Which of the following molecules have dipole-dipole forces?
CO2 (b) CH2Cl2 (c) CH4
Draw the structure and determine if there are polar bonds present… try to see if the molecule is overall polar to have dipole-dipole forces

26. Practice Which of the following molecules have dipole-dipole forces?
C I4 (b) CH3Cl (c) HCl

27. #3: Hydrogen Bonding Polar molecules with hydrogen can bond to F,N,O
CAN make Hydrogen bonds:
HF NH3 H2O
H-bonds are STRONG for IMFs!!
LEADS TO HIGH MP’S AND BP’S!
Sort of the “super dipole-dipole” force

28. Why??
The large Δ in EN between H and N,O,F makes a very (+) end on the H and a very (-) end on the N, O, or F
Also, because these elements are still quite small, they are able to get pretty close to one another, increasing the strength of attraction (coulomb’s law)

30. What's going on here??

31. WARNING!!
H-BONDS ARE NOT BONDS!!! STILL ONLY ABOUT 2-5% THE STRENGTH OF A TYPICAL COVALENT BOND!!

32. Practice #1
One of the following is a liquid at room temperature. Which one?
Formaldehyde (CH2O)
Fluoromethane (CH3F)
Hydrogen peroxide (H2O2)
**TRY drawing them out!!

33. Explanation to Practice #1
All three have similar molar masses
Meaning similar dispersion forces
All three are also polar
BUT… the hydrogen peroxide has H-O bonds
So, there are hydrogen bonding in the H2O2
Note: the other two DO have H and N,O,F… but they need to be bonded directly to the H
Why??

34. Practice #2
What has the higher boiling point, HF or HCl??
Why??

35. THESE ARE THE STRONGEST INTERMOLECULAR FORCES
#4 Ion-Dipole Force
The ion-dipole force occurs when an ionic compound is mixed with a polar compound
Occurs when something ionic is put into water and dissociates into its ions
THESE ARE THE STRONGEST
INTERMOLECULAR FORCES

36. Ion-Dipole Force Think about NaCl in water
The Na+ and Cl- ions break apart and float around
The water SURROUNDS and attaches to the respective opposite charges

38. IMF’s in Action
Intermolecular forces is the reason and determination of phase (s, l, or g)
They also are the cause of other very important (for life) properties
Surface Tension
Viscosity
Capillary Action

39. Surface Tension
Why does a fly fishermen’s fly float on the surface of the water… the hook is metal and more dense
The Surface Tension of a liquid is the energy required to increase the surface are by unit amount
Water has a surface tension of 72.8 mJ/m2
Meaning it takes 72.8 mJ of energy to increase the surface area of water by one square meter

40. Why?? Liquids always like to minimize their surface area
This creates kind of a “skin” on the top of the liquid that resists penetration
For the fishermen’s fly to sink it would have to slightly increase the surface area of the water… which is resisted by the water

41. Surface Tension and IMFs
Surface Tension decreases with decreasing IMFs
You can float a paper clip on water but cannot float a paper clip on benzene (a nonpolar organic solvent) because there are no H-bonds AND the dispersion forces are not strong enough to combat the increase in surface area

42. Water Droplets
Surface Tension is also the reason why water “beads” up and forms almost perfect spheres of water – in space (with no gravity), large amounts of water will ball up perfectly.
Why? The sphere is the shape with the smallest ratio of surface are to volume… minimizing the number of molecules at the surface… minimizing the potential energy of the system

44. Viscosity Viscosity is the resistance for liquid to flow
Motor oil is more viscous than gasoline
Viscosity is greater in substances with stronger intermolecular forces because the molecules are strongly attracted to each other, preventing them from flowing past one another

45. Viscosity and Temperature
How does temperature affect viscosity?
Viscosity decreases with an increase in temperature
The added kinetic energy will temporary overcome the IMF’s and allow them to move more freely past one another.

46. Capillary Action
Capillary Action – the ability of a liquid to flow against gravity up a narrow tube
Caused by two forces:
Adhesion –
Cohesion –

47. Capillary Action
Capillary Action – the ability of a liquid to flow against gravity up a narrow tube
Caused by two forces:
Adhesion – the attraction between liquid molecules and the sides of the glass, root, etc. (to other things)
Cohesion – the attraction between molecules within a liquid (to itself)

48. QUESTION:
What is a meniscus?? Why
does it happen??

49. Properties Affected by IMFs
Phase Changes…
Vaporization – the process by which thermal energy can overcome IMFs and produce a phase change from liquid to gas
Removing energy will decrease their average kinetic energy and bring them back down as a liquid (condensation)

50. Open System
Both vaporization and condensation occur at EQUAL rates (as some go into the gaseous phase, others fall back in the liquid phase)
Average kinetic energy (some are lower than the average, others are higher)
In an open beaker / container
The evaporation occurs more… the gases escape into the atm and do not come back

51. Closed System
In the same scenario, the beaker or bottle is not capped / sealed
The two processes still occur at the same rate
Reach a happy “dynamic equilibrium”

52. Other Liquids What happens if you increase the temperature?
What happens to the vaporization process / temperature if the liquid is not water… but something with much weaker IMFs?

53. Volatility If a liquid evaporates easily it is volatile
Those that do not evaporate easily are called nonvolatile
Example: Acetone (nail polish remover) is more volatile than water.
Motor oil is essentially nonvolatile at room temp
Does not evaporate much or at all at room temp.

54. **Summary**
The rate of vaporization increases with increasing temperature
The rate of vaporization increases with increasing surface area!!
The rate of vaporization increases with decreasing strength of IMFs

55. Endo- or Exo- ?? Vaporization vs. condensation
Q: endo- or exo- ?
If no additional energy is put into a system, what happens to the E as the liquid is let to evaporate?

56. Heat of Vaporization
The amount of heat required to vaporize one mole of a liquid to gas is called the heat of vaporization (ΔHvap)
The ΔHvap is always positive because the process is always endothermic
It is temperature / pressure dependent

57. Heat of Vaporization The ΔHvap for water at 100 ⁰ C is +40.7 kJ/mol
When a liquid condensates, that same amount of energy is released
The ΔH for water condensing at 100⁰C is: ΔH = - ΔHvap = kJ/mol
(releases 40.7 kJ per mole of H2O)
 (-) means leaves the system

58. Practice #1
Calculate the mass of water (in grams) that can be vaporized at its boiling point with 155 kJ of heat.
68.6 grams H2O

59. Practice #2
(pg 480)
Calculate the amount of heat (in kJ) required to vaporize 2.58 kg of water at its boiling point.

60. More Practice (good TQ)
Suppose that 0.48 g of water at 25⁰C condenses on the surface of a 55-g block of aluminum that is initially at 25⁰C. If the heat released during condensation goes only toward heating the metal, what is the final temperature (in ⁰C) of the metal block?
The specific heat capacity of aluminum is J/g ⁰C)

61. Vapor Pressure – Back to a SEALED System!!
Once the liquid/gas is allowed to reach this dynamic equilibrium:
The pressure of the gas with its liquid is called vapor pressure
The vapor pressure depends on the IMFs!!
The weaker the IMFs, the more volatile the liquid, the more molecules in the gaseous state, the higher the vapor pressure!!

62. Volume and Vapor Pressure
What happens to the vapor pressure in these scenarios?

63. Systems WANT Equilibrium!!
When a system in dynamic equilibrium is disturbed, the system will respond so to minimize the disturbance and reestablish its state of equilibrium

64. Practice
What happens to the vapor pressure of a substance when its surface area is increased at constant temperature?
The vapor pressure increases
The vapor pressure remains the same
The vapor pressure decrease
B!! The VP is independent of surface area… with the increase in surface area DOES cause an increase in vaporization… but also an increase in the condensation!!

65. Vapor Pressure and Temp
When the temperature of a substance is increased, the vapor pressure rises
The thermal energy enables more molecules to break free and enter the gaseous state
A small change in T will cause a big increase… but this curve relationship does level out and flatten
What do we call this??
The boiling point!!!

67. Boiling Point
The boiling point is the temperature at which its vapor pressure equals the external pressure.
This is the temperature where molecules, EVEN THOSE IN THE MIDDLE, have the energy to break free (overcome the IMFs between them and their neighbors) ….
Not just those on the surface bouncing and springing out

68. The Normal Boiling Point
The normal boiling point is the temperature that the vapor pressure equals 1 atm (standard)
This changes!! How?
Increasing or decreasing the pressure (deviating from the P = 1 atm) will change this boiling point (no longer “normal” boiling point)

69. Supercritical Points
Sealed system… but NOW increasing the Temperature
The critical temperature (Tc) and critical pressure (Pc) are the conditions needed to cause this change – SUPERCRITICAL FLUID

70. Supercritical Fluids
Supercritical Fluids have properties of both liquids and gases.
Kind of an in between phase
They are particularly good solvents… dissolving a large range of molecules.
Supercritical CO2 dissolves caffeine, making it easy to be removed for de-caf

71. Phase Changes Involving Solids
Think about a block of ice
The thermal energy is significantly less than that of liquids or gases
However, at the surface, molecules may still have enough energy to break free from their neighboring IMFs holding them together and float off into the gaseous phase
What is this called?

72. Sublimation
Sublimation - The molecules on the surface of a solid that have enough energy to go directly to the gaseous phase
Some of the water vapor molecules may also collide with the surface of the ice and be ‘pulled’ in / captured by the IMFs of the solids… this process is known as deposition

73. Open System
Sublimation usually occurs at a faster rate than deposition because the gaseous particles float away and never come back
Think about ice cubes left in the freezer for too long… they shrink… why?
Notice the gradual growth of ice crystals on the walls around the freezer??

74. Interesting
Freezer burn is when foods are left in the freezer too long and dry out
The frozen water molecules are left to sublime and never return

75. Heating Curves Thinking back to our block of ice…
If we heat the cube up (from, say -10⁰C) the molecules move faster and faster and the temperature will increase
However, you soon realize the temperature stops increase even though you are adding in more heat energy. What is happening?

76. Heating Curve for Water, H2O

78. Terms for Heating Curves
Fusion – the phase change between solid and liquid (aka: melting and freezing)
Vaporization – (we already know) is the change between liquid and gas (aka: evaporation and condensation)

80. Heat of Fusion (ΔHfus)
The fastest way to cool a drink is by adding ice cubes to it
They melt and pull absorb energy (endothermic)
H2O(s)  H2O(l) ΔHfus = 6.02 kJ/mol
Freezing is the opposite (exothermic) – releases energy to cool and freeze itself
H2O(l)  H2O(s) ΔH = -ΔHfus = kJ/mol

81. Compare ΔHfus and ΔHvap
Typically, the ΔHfus is much less than the ΔHvap
This is because the solid and liquid phases are much closer together and require less energy to exhibit this phase change.
Gases are much more energized and will usually require a much higher ΔH (in this case, ΔHvap)

82. Mathematical Relevance

84. Five Parts
When there is NOT a phase change taking place, you can calculate the heat (q) using
q = mCΔT
Like Parts 1, 3, and 5

85. q = nΔHfus Like Part (2) q = nΔHvap Like Part (4)
Five Parts
When there IS A CHANGE taking place, you need to use the ΔH fus/vap
Remembering, these values are PER MOLE
q = nΔHfus Like Part (2)
q = nΔHvap Like Part (4)

86. Phase Diagrams
Phase Diagrams are maps of the phases of a substance as a function of pressure (y-axis) and temperature (on the x-axis)

87. Important Parts Regions – solid, liquid, or gas
Lines – each line represents a T and P that the two phases are at dynamic equilibrium. The two phases are equally stable and coexist.
The Triple Point – a unique set of conditions where all three phases are actually equally stable and in equilibrium
The Critical Point – the set of P and T above which a supercritical fluid exists

89. Phase Diagram for CO2

90. Phase Diagram for Iodine

91. WATER IS SPECIAL!!