Liquids and solids

1
Chapter 10

• Liquids and solids

2
IMF
3
IntERmolecular forces

• Forces at work (not chemical bonds) between
  molecules in liquids, solids and real gases
• London dispersion forces (van der Waals)
• Dipole-dipole
• Hydrogen bonds

4
Those without dipoles.

• Most are gases at 25ºC.
• The only forces are London Dispersion Forces.
• These depend on size of atom.
• Large molecules (such as I2 ) can be solids even
  without dipoles.

5
Those with dipoles.

• Dipole-dipole forces are generally stronger than
  L.D.F.
• Hydrogen bonding is stronger than Dipole-dipole
  forces.
• No matter how strong the intermolecular force, it
  is always much, much weaker than the forces in
  bonds.
• Stronger forces lead to higher melting and
  freezing points.

6
London Dispersion ForcesExample
7
Dipole-Dipole ForcesExample
8
Hydrogen Bond Example
d
d-
d
9
Intermolecular forces

• What do they affect?
• During phase changes, IntERmolecular forces
  change
• Takes energy to overcome IMF forces
• - Thus affecting boiling points, melting points
• Gives liquids and solids important
  characteristics

10
100
Boiling Points
0ºC
-100
200
11
Liquids
12
Liquids

• Many of the properties due to internal attraction
  of atoms.
• Beading
• Surface tension
• Capillary action
• Stronger intermolecular forces cause each of
  these to increase.

13
Surface tension

• Molecules at the the top are only pulled inside.

• Molecules in the middle are attracted in all
  directions.

• Minimizes surface area.

14
Capillary Action

• Liquids spontaneously rise in a narrow tube.
• Inter molecular forces are cohesive, connecting
  like things.
• Adhesive forces connect to something else.
• Glass is polar.
• It attracts water molecules.

15
(No Transcript)
16
Beading

• If a polar substance is placed on a non-polar
  surface.
• There are cohesive,
• But no adhesive forces.
• And Visa Versa

17
Viscosity

• How much a liquid resists flowing.
• Large forces, more viscous.
• Large molecules can get tangled up.
• Cyclohexane has a lower viscosity than hexane.
• Because it is a circle- more compact.

18
Model of a liquid

• In motion but attracted to each other
• With regions arranged like solids but
• with higher disorder.
• with fewer holes than a gas.
• Highly dynamic, regions changing between types.

19
Solids
20
Solids

• Two major types.
• Crystalline- have a regular arrangement of
  components in their structure.
• Amorphous- those with much disorder in their
  structure.

21
Crystalline Solids

• Lattice- a three dimensional grid that describes
  the locations of the pieces in a crystalline
  solid.
• Unit Cell-The smallest repeating unit in of the
  lattice.
• 3 Types of crystalline solids
• Atomic
• Molecular
• Ionic

22
(No Transcript)
23
Properties of SolidsChart on pg 458

• Atomic
• Network solid hard, HMP, insulator
• - diamonds
• Metallic solid ranges, conductor
• - silver, iron, brass
• Group 8A solid lMP
• - argon
• Molecular
• Polar Nonpolar soft, LMP, insulator
• Ice, dry ice
• Ionic hard, HMP, insulator
• sodium chloride, calcium fluoride

24
Why is it an insulator?
E

• The space between orbitals make it impossible for
  electrons to move around

25
Ionic Solids

• The extremes in dipole dipole forces-atoms are
  actually held together by opposite charges.
• Huge melting and boiling points.
• Atoms are locked in lattice so hard and brittle.
• Every electron is accounted for so they are poor
  conductors-good insulators.

26
Carbon- A Special Atomic Solid

• Types of solid carbon.
• Amorphous- coal uninteresting.
• Diamond- hardest natural substance on earth,
  insulates both heat and electricity.
• Graphite- slippery, conducts electricity.

27
Diamond- each Carbon is sp3hybridized, connected
to four other carbons.

• Carbon atoms are locked into tetrahedral shape.
• Strong s bonds give the huge molecule its
  hardness.

28
Graphite is different.

• Each carbon is connected to three other
  carbons and sp2 hybridized.
• The molecule is flat with 120º angles in
  fused 6 member rings.
• The p bonds extend above and below the plane.

29
This p bond overlap forms a huge p bonding
network.

• Electrons are free to move through out these
  delocalized orbitals.
• The layers slide by each other.

30
Graphite layers of carbon
31
Types of Solids Examples

• Gold
• Atomic solid w/ metallic properties
• Carbon dioxide
• Molecular solid w/ nonpolar cov. properties
• Lithium fluoride
• Ionic solid
• Krypton-
• Atomic solid w/ nonpolar cov. Properties (LDF)

32
Structure Bonding in metals
33
Metallic bonding
Empty Molecular Orbitals
3p
Filled Molecular Orbitals
3s
2p
2s
1s
Magnesium Atoms
34
The 1s, 2s, and 2p electrons are close to
nucleus, so they are not able to move around.
Empty Molecular Orbitals
3p
Filled Molecular Orbitals
3s
2p
2s
1s
Magnesium Atoms
35
The 3s and 3p orbitals overlap and form molecular
orbitals.
Empty Molecular Orbitals
3p
Filled Molecular Orbitals
3s
2p
2s
1s
Magnesium Atoms
36
Electrons in these energy level can travel freely
throughout the crystal.
Empty Molecular Orbitals
3p
l l
Filled Molecular Orbitals
3s
2p
2s
1s
Magnesium Atoms
37
This makes metals conductors Malleable because
the bonds are flexible.
Empty Molecular Orbitals
3p
l l
Filled Molecular Orbitals
3s
2p
2s
1s
Magnesium Atoms
38
Phase Changes
39
Phases

• The phase of a substance is determined by three
  things.
• The temperature.
• The pressure.
• The strength of intermolecular forces.

40
Vapor Pressure

• Vaporization - change from liquid to gas at
  boiling point.
• Evaporation - change from liquid to gas below
  boiling point
• Heat (or Enthalpy) of Vaporization (DHvap )- the
  energy required to vaporize 1 mol at 1 atm.

41

• Vaporization is an endothermic process - it
  requires heat.
• Energy is required to overcome intermolecular
  forces.
• Responsible for cool earth.
• Why we sweat. (Never let them see you.)

42
Condensation

• Change from gas to liquid.
• Achieves a dynamic equilibrium with vaporization
  in a closed system.
• What is a closed system?
• A closed system means matter cant go in or
  out.
• Put a cork in it.
• What the heck is a dynamic equilibrium?

43
Dynamic equilibrium

• When first sealed the molecules gradually escape
  the surface of the liquid.

44
Dynamic equilibrium

• When first sealed the molecules gradually escape
  the surface of the liquid.
• As the molecules build up above the liquid some
  condense back to a liquid.

45
Dynamic equilibrium

• When first sealed the molecules gradually escape
  the surface of the liquid.
• As the molecules build up above the liquid some
  condense back to a liquid.
• As time goes by the rate of vaporization remains
  constant but the rate of condensation
  increases because there are more molecules to
  condense.

46
Dynamic equilibrium

• When first sealed the molecules gradually escape
  the surface of the liquid
• As the molecules build up above the liquid some
  condense back to a liquid.
• As time goes by the rate of vaporization remains
  constant but the rate of condensation
  increases because there are more molecules to
  condense.
• Equilibrium is reached when

47
Dynamic equilibrium

• Rate of Vaporization Rate of Condensation
• Molecules are constantly changing phase Dynamic
• The total amount of liquid and vapor remains
  constant Equilibrium

48
Vapor pressure

• The pressure above the liquid at equilibrium.
• Liquids with high vapor pressures evaporate
  easily. They are called volatile.
• Decreases with increasing intermolecular forces.
• Bigger molecules (bigger LDF)
• More polar molecules (dipole-dipole)

49
Vapor pressure

• Increases with increasing temperature.
• Easily measured in a barometer.

50

• A barometer will hold a column of mercury 760 mm
  high at one atm

51

• A barometer will hold a column of mercury 760 mm
  high at one atm.
• If we inject a volatile liquid in the barometer
  it will rise to the top of the mercury.

52
Water

• A barometer will hold a column of mercury 760 mm
  high at one atm.
• If we inject a volatile liquid in the barometer
  it will rise to the top of the mercury.
• There it will vaporize and push the column of
  mercury down.

Patm 760 torr
Dish of Hg
53
Water Vapor

• The mercury is pushed down by the vapor pressure.
• Patm PHg Pvap
• Patm - PHg Pvap
• 760 - 736 24 torr

736 mm Hg
Dish of Hg
54
Temperature Effect
Energy needed to overcome intermolecular forces
T1
of molecules
Kinetic energy
55

• At higher temperature more molecules have enough
  energy - higher vapor pressure.

Energy needed to overcome intermolecular forces
Energy needed to overcome intermolecular forces
T1
T1
of molecules
T2
Kinetic energy
56
Mathematical relationship

• ln is the natural logarithm
• ln Log base e
• e Eulers number an irrational number like p
• DHvap is the heat of vaporization in J/mol

57
Mathematical relationship

• R 8.3145 J/K mol.
• Surprisingly this is the graph of a straight
  line. (actually the proof is in the book)

58
Changes of state

• The graph of temperature versus heat applied is
  called a heating curve.
• The temperature a solid turns to a liquid is the
  melting point.
• The energy required to accomplish this change is
  called the Heat (or Enthalpy) of Fusion DHfus

59
Heating Curve for Water
Steam
Water and Steam
Water
Water and Ice
Ice
60
Heating Curve for Water
Heat of Fusion
61
Melting Point

• Melting point is determined by the vapor pressure
  of the solid and the liquid.
• At the melting point the vapor pressure of the
  solid vapor pressure of the liquid

62
(No Transcript)
63

• If the vapor pressure of the solid is higher than
  that of the liquid the solid will release
  molecules to achieve equilibrium.

64

• While the molecules of condense to a liquid.

65

• This can only happen if the temperature is above
  the freezing point since solid is turning to
  liquid.

66

• If the vapor pressure of the liquid is higher
  than that of the solid, the liquid will release
  molecules to achieve equilibrium.

67

• While the molecules condense to a solid.

68

• The temperature must be above the freezing point
  since the liquid is turning to a solid.

69

• If the vapor pressure of the solid and liquid are
  equal, the solid and liquid are vaporizing and
  condensing at the same rate. The Melting point.

70
Boiling Point

• Reached when the vapor pressure equals the
  external pressure.
• Normal boiling point is the boiling point at 1
  atm pressure.
• Super heating - Heating above the boiling point.
• Supercooling - Cooling below the freezing point.

71
Phase Diagrams.

• A plot of temperature versus pressure for a
  closed system, with lines to indicate where there
  is a phase change.

72
D
D
Pressure
D
1 Atm
D
Temperature
73
Pressure
Temperature
74

• This is the phase diagram for water.
• The density of liquid water is higer than solid
  water.

Pressure
Temperature
75

• This is the phase diagram for CO2
• The solid is more dense than the liquid
• The solid sublimes at 1 atm.

Pressure
Liquid
Solid
1 Atm
Gas
Temperature