Chapter 10 Liquids and Solids

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Chapter 10Liquids and Solids
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Preview

• Intermolecular forces.
• Structure in liquid state.
• Structure in solid state.
• Types of crystalline solids. (Types of cells)
• Structure and bonding in metals.
• Molecular solid and ionic solid.
• Vapor pressure and changes of states.
• Phase diagram.

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The Three States of Matter
Introduction

• Gas Solid
• Density Low High
• Compressibility High Low
• Internal attraction Low High
• forces

Liquid lies somehow in between but is closer in
its structural properties to solid. H2O(s) ?
H2O(l) ?Hºfus 6.02 kJ/mol H2O(l) ?
H2O(g) ?Hºvap 40.7 kJ/mol
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The Three States of Matter
Introduction
The molecules are locked in place but can vibrate
about their positions.
The motions of the molecules increase, achieving
greater movement and disorder.
Molecules are far apart from each other and have
little interactions
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The Three States of Matter
Introduction
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Intermolecular Forces
Chapter 10 Section 1

• The force holding atoms together to form a
  molecule is intramolecular force.
• The force holding a collection of molecules
  together (aggregate) is intermolecular force.
• The changes in states are due the changes in
  intermolecular interaction.
• H2O(s) ? H2O(l) ?Hºfus 6.02
  kJ/mol
• H2O(l) ? H2O(g) ?Hºvap 40.7
  kJ/mol
• H2O(g) ? 2H(g) O(g) ?Hºdecomp. 934
  kJ/mol

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Intermolecular Forces
Chapter 10 Section 1

• Dipole-Dipole Forces
• It exists as a weak electrostatic attraction
  within polar molecules.
• In a condensed state, the molecules orient
  themselves to maximize electrostatic attractions
  and minimize the electrostatic repulsions.
• Dipole-dipole forces is 1 as strong as covalent
  and ionic bonds. They become weaker as the
  distance between the molecules gets larger.

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Intermolecular Forces
Chapter 10 Section 1

• Hydrogen Bonding
• It is associated with highly electronegative
  atoms, especially N, O, and F atoms.
• Two factors explain the strengths of hydrogen
  bonding
• The great polarity of the bond.
• The close approach of the dipole allowed by the
  smaller size of the atoms.

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Intermolecular Forces
Chapter 10 Section 1
The boiling points of the covalent hydrides of
the elements in groups 4A, 5A, 6A, and 7A
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Intermolecular Forces
Chapter 10 Section 1

• Many nitrogen-, oxygen-, and fluorine-containing
  molecules exhibit hydrogen bonding.

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Intermolecular Forces
Chapter 10 Section 1

• Ion-Dipole Interaction
• It is shown by attraction force between an ion
  and a polar molecule.

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Intermolecular Forces
Chapter 10 Section 1

• London Dispersion Forces
• It is shown by non-polar molecules and noble
  gases.
• We had assumed that the electron charge is
  uniformly distributed around the nucleus
  (electron cloud).
• However, nonsymmetrical electron distribution
  could develop on an atom. This would lead to
  temporary dipole within the atom, and a distorted
  electron would would result.

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Intermolecular Forces
Chapter 10 Section 1

• The instantaneous dipole in a molecule or atom
  can induce similar dipole in neighboring atoms.
  (Polarizability)
• This type of interaction is very weak and
  short-lived, yet it is significant especially for
  large atoms.

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Intermolecular Forces
Chapter 10 Section 1

• London Dispersion Forces in Nobel Gases
• Nobel gas can exist in a form of liquid or solid,
  but at very low temperatures.

• For the dispersion force to be strong, the
  motions of the atoms have to be slowed down.
• Dispersion force becomes stronger and more
  significant as the size of the atom increases
  (larger electron cloud). We say larger atoms
  exhibit higher polarizability.

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Intermolecular Forces
Chapter 10 Section 1

• London Dispersion Forces in Nonpolar Molecules

• Dispersion force becomes stronger and more
  significant as the molar mass of the molecule
  increases (larger number of electrons).

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Intermolecular Forces
Chapter 10 Section 1

• ? What type(s) of intermolecular forces exist
  between each of the following molecules?

HI
HBr is a polar molecule dipole-dipole forces.
There are also dispersion forces between HBr
molecules.
CF4
CH4 is nonpolar ONLY dispersion forces.
SO2
SO2 is a polar molecule dipole-dipole forces.
There are also dispersion forces between SO2
molecules.
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The Liquid State
Chapter 10 Section 2

• Liquid has
• Low compressibility.
• Lack of rigidity.
• Higher density as compared to gas.
• Tendency to bead as droplets (intermolecular
  forces).

The intermolecular attractions between the
molecules in a liquid cause the surface area of
the droplet to be minimized (spherical shape).
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The Liquid State
Chapter 10 Section 2

• Surface Tension
• It is the resistance of a liquid to an increase
  in its surface area.
• Liquid with larger intermolecular forces (such
  as those with polar molecules) will have higher
  surface tension.

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The Liquid State
Chapter 10 Section 2

• Capillary Action
• It is the rising of liquid a narrow tube.
• Cohesive forces (Intermolecular).
• Adhesive forces (between the liquid molecules and
  the container).
• In a glass tube (container made of polar
  substances)
• Water forms a concave meniscus.
• cohesive lt adhesive
• Mercury forms a convex meniscus.
• cohesive gt adhesive

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The Liquid State
Chapter 10 Section 2

• Viscosity
• It is the measure of liquids resistance to
  flow.
• Liquids with larger intermolecular interaction
  tend to be more viscous.
• Liquids with more molecular complexity are highly
  viscous.

Gasoline (CH3 (CH2)8 CH3) vs.
Grease (CH3 (CH2)25 CH3)
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Introduction to Solid-State Structure
Chapter 10 Section 3

• Categories of Solid
• Crystalline Solid its components are highly
  arranged.
• Amorphous Solid its components have considerable
  disorder.

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Introduction to Solid-State Structure
Chapter 10 Section 3

• Crystalline Solids
• Lattice is 3-D system of points (in defined
  positions) that make up a substance.
• Unit cell is the smallest repeating unit in a
  lattice.

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Introduction to Solid-State Structure
Chapter 10 Section 3

• Crystalline Solids
• Lattice is 3-D system of points (in defined
  positions) that make up a substance.
• Unit cell is the smallest repeating unit in a
  lattice.

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Introduction to Solid-State Structure
Chapter 10 Section 3
Unit cell
Space-filling unit cell
No. of atoms per unit cell
Lattice
(8 1/8 1)
1 atom / unit cell
(8 1/8 1 2)
2 atoms / unit cell
(8 1/8 6 1/2 4)
4 atoms / unit cell
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X-Ray Diffraction
Chapter 10 Section 3

• X-ray diffraction instrument can be used to
  analyze the structure of various chemical solid
  materials.

Diffraction pattern
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X-Ray Diffraction
Chapter 10 Section 3

• Diffraction is the scattering of light beams from
  a regular array of points in which the spacings
  (d) between the components are comparable with
  the wavelength of the light (?). The angle of
  incident / reflection is given by (?).

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X-Ray Diffraction
Chapter 10 Section 3

• X-rays scattered from two Different atoms may
  (a) reinforce (constructive interference), or (b)
  cancel (destructive interference) one another.

• The difference in distances traveled is an
  integral number (n) of wavelength (?).
• The diffraction pattern can be used to determine
  the interatomic distances.

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X-Ray Diffraction
Chapter 10 Section 3

• xy yz n? (In-phase light)
• xy yz 2d sin? (From trigonometry)
• 2d sin?
  n? (Bragg Equation)

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X-Ray Diffraction
Chapter 10 Section 3

• Question 43 on page 476.
• A topaz crystal has an interplanar spacing of
  1.36 Å. Calculate the wavelength of the X ray
  that should be used if the ? is 15.0º. Assume n
  1.
• 2d sin? n?

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Classification of Solids
Chapter 10 Section 3
Atomic Solid
Ionic Solid
Molecular Solid
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Classification of Solids
Chapter 10 Section 3
Examples
Copper
Graphite, diamond, fullerene
Argon, xenon
Ice, sugar, oil
Salt
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Structure and Bonding in Metals
Chapter 10 Section 4

• Characteristics of metals
• High thermal and electrical conductivity.
• Malleability.
• Ductility.
• Metals have special type of nondirectional
  covalent bonding.

• The closest packing model describes the metal
  atoms as layers of uniform, hard spheres. These
  layers dont lie directly on top pf each other.

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Structure and Bonding in Metals
Chapter 10 Section 4

• Two main types of packing
• aba closest packing the spheres in every other
  layer occupy the same vertical positions.
• abc closest packing the spheres in every fourth
  layer occupy the same positions.
• These arrangements have the most efficient use of
  the available space.

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The aba Closest Packing
Chapter 10 Section 4

• The resulting structure from the aba packing is
  the hexagonal closest packed (hcp) structure (Mg,
  Zn, Ca)
• The resulting unit cell is hexagonal prism cell.

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The abc Closest Packing
Chapter 10 Section 4

• The resulting structure from the abc packing is
  the cubic closest packed (ccp) structure (Al, Fe,
  Cu, Co, Ni, Ca).
• The resulting unit cell is faced-centered cubic
  cell.

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Closest Packing Model
Chapter 10 Section 4

• In both types of packing (hcp and ccp), each
  sphere has 12 equivalent neighboring spheres
• 6 in the same layers,
• 3 in the layer above, and
• 3 in the layer below.
• For the body-centered cubic (bcc) unit cell, each
  sphere is surrounded by 8 neighboring spheres
  (alkali metals).

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Volume of Unit Cells
Chapter 10 Section 3
Unit cell
Space-filling unit cell
No. of atoms per unit cell
Lattice
(8 1/8 1)
1 atom / unit cell
(8 1/8 1 2)
2 atoms / unit cell
(8 1/8 6 1/2 4)
4 atoms / unit cell
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Volume of Unit Cells
Chapter 10 Section 4

• Simple Cubic Cell
• d 2r
• Volume d3 8r3

d
r
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Volume of Unit Cells
Chapter 10 Section 4

• Body-Centered Cubic Cell
• c 2d 4r
• b2 d2 d2
• c2 d2 b2
• c2 d2 2d2 3d2
• 3d2 16r2
• d 4r/v3
• Volume 64r3/(3)3/2

c
b
d
r
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Volume of Unit Cells
Chapter 10 Section 4

• Face-Centered Cubic Cell
• d2 d2 b2
• b 4r
• b2 16r2
• 2d2 16r2
• d2 8r2
• d rv8
• Volume d3 83/2r3

b
d
r
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Bonding in Metals
Chapter 10 Section 4

• Bonds in metals are strong and nondirectional.
• The Electron Sea Model It is difficult to
  separate metal atoms, but it is relatively easy
  to move them.

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Bonding in Metals (MO Model)
Chapter 10 Section 4

• MO model can be used to explain the bonding in
  metals.
• Li2 molecule

Levels become continuum
Band Model
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Bonding in Metals (MO Model)
Chapter 10 Section 4

• Mg metal crystal.
• ? The electrons in the 1s, 2s, and 2p orbitals
  are localized.
• ? The 3s and 3p orbitals can overlap to form
  MOs. Electrons in valence MOs can travel through
  crystal (Conduction band).

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Alloys
Chapter 10 Section 4

• An alloy is a substance that contains a mixture
  of elements and has metallic properties.
• There are two main types of alloys
• Substitutional alloys some of the host metal
  atoms are replaced by other metal atoms.
• Brass 67 Cu and 33 Zn.
• Interstitial alloys some holes in the closest
  packed metal are occupied by small atoms.
• Steel C atoms in the holes between Fe atoms.

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Alloys
Chapter 10 Section 4

• Steel is one of the most common alloys
• Mild steels less than 0.2 C (chains, cable).
• Medium Steels 0.2-0.6 C (structural steel
  beams).
• High-Carbon Steels 0.6-1.5 C (cutlery,
  springs).
• Alloy Steels mixed interstitial and
  substitutional alloys.

• Stainless steel is an alloy steel with minimum
  10.5 Cr.

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Network Atomic Solids
Chapter 10 Section 5

• Network atomic solids have strong directional
  bonds. Their conduction to electricity is not as
  good as metals.
• Carbon network solids (diamond and graphite).
• Silicon network solids (rock, sands, soils,
  quartz, and glass)

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Carbon Network Solids
Chapter 10 Section 5

• Tetrahedral arrangement (sp3).
• Extremely hard, poor electrical and thermal
  conductor, and colorless.

• Layers of carbon atoms in form of rings (sp2).
• slippery, conduct heat and electricity, and
  black.

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Carbon Network Solids
Chapter 10 Section 5

• The large gap between the filled orbitals and the
  empty ones accounts for the insulation property
  for diamond.

Diamond
Metal
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Carbon Network Solids
Chapter 10 Section 5

• In graphite, the unhybridized 2p orbitals on the
  C atoms result in p bonds in which the electrons
  are delocalized. This explains why graphite is an
  electrical conductor.
• The layer structure explains the slipperiness of
  graphite.

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Silicon Network Solids (Silica)
Chapter 10 Section 5

• They are the main constituents of earth crust.
• They involve chains with Si-O bonds.
• Silica (SiO2) is the most common type of silicon
  network solids.
• Does the structure of SiO2 look like that of CO2?

Silica network (quartz)
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Silicon Network Solids (Quartz)
Chapter 10 Section 5

• In silica, SiO2, which is the basic component of
  quartz, the inability of the Si atoms to form
  efficient p bonds with oxygens leads to
  tetrahedral network structure, SiO4, instead of
  the discrete linear structure (like CO2).
• Still the empirical formula of quartz is SiO2.

Silica network (quartz)
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Silicon Network Solids (Silicates)
Chapter 10 Section 5

• The silicates are found mostly in rocks, clays,
  and soils.
• ? O/Si ratio gt 21.
• ? Contains Si/O anions.
• ? Cations are needed to balance the excess
  negative charge. This means that silicate is a
  salt.

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Silicon Network Solids (Glass)
Chapter 10 Section 5

• Glass is formed by heating the silica above its
  melting point (1600ºC) and then cooled rapidly.
  This process causes an amorphous solid called
  glass to form.
• Glass contains a great deal of disorder.
• Glass can be mixed with other substances to form
  special classes of glass, such as Pyrex, optical
  objects, and windows.

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Silicon Network Solids (Glass)
Chapter 10 Section 5
55
Silicon Network Solids (Ceramics)
Chapter 10 Section 5

• Ceramics are made from clays (silicates).
• hardened by firing at high temperatures.
• strong and brittle.
• extremely thermally resistant.
• chemically inert.

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Semiconductors
Chapter 10 Section 5

• Silicon networks is tetrahedral (like diamond),
  but the energy gap between filled and empty MOs
  is relatively smaller. Few electrons can cross
  it.
• (Semiconductors).
• Conduction property increases with the increase
  in temperature.
• n-type and p-type semiconductors can be produced.

57
Molecular Solids
Chapter 10 Section 6
58
Molecular Solids
Chapter 10 Section 6

• In molecular solids, each molecule occupies a
  lattice position. Molecular solid is
  charecterized by strong covalent intramolecular
  bonding and weak intermolecular forces.
• Ice is made of H2O molecules held together by
  means of hydrogen bonding.

59
Molecular Solids
Chapter 10 Section 6

• In molecular solids, each molecule occupies a
  lattice position. Molecular solid is
  charecterized by strong covalent intramolecular
  bonding and weak intermolecular forces.
• Ice is made of H2O molecules held together by
  means of hydrogen bonding.

60
Molecular Solids
Chapter 10 Section 6

• In molecular solids, each molecule occupies a
  lattice position. Molecular solid is
  charecterized by strong covalent intramolecular
  bonding and weak intermolecular forces.
• Ice is made of H2O molecules held together by
  means of hydrogen bonding.
• Dry ice (solid CO2) is another example of
  molecular solid.

61
Molecular Solids
Chapter 10 Section 6

• Yellow sulfur crystals (with S8 molecules) and
  white phosphorous (with P4 molecules).
• They dont have dipole moments. Thus, the
  intermolecular forces are dispersion forces.
• At 25ºC
• CO2-based crystal is gas.
• P4, S8, and Cl2-based crystals are solid.

62
Molecular Solids
Chapter 10 Section 6
63
Ionic Solids
Chapter 10 Section 7

• Ionic molecules
• very stable.
• high-melting substances.
• electrostatic attraction.
• when melted, they conduct electricity.
• The closest packed model can be used to extend
  the structure of ionic molecules to ionic solids.

64
Ionic Solids
Chapter 10 Section 7

• The larger ions are packed with the hcp or ccp
  arrangements. The resulting holes are filled by
  the oppositely charged ions.
• Types of holes in ionic solids
• Trigonal holes.
• Tetrahedral holes.
• Octahedral holes.
• The hole size order
• trigonal lt tetrahed. lt octahed.

65
Ionic Solids
Chapter 10 Section 7

• Trigonal holes are so small that they can not be
  occupied to form binary ionic compounds.
• Tetrahedral or octahedral arrangements are
  determined by the relative sizes of the cations
  and anions.
• Tetrahedral holes in the ccp packed structure.
• Example ZnS.
• Ionic radius Zn2 70 pm , S2 180 pm.

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Ionic Solids
Chapter 10 Section 7

• For the fcc unit cell in the ccp packed
  structure, number of tetrahedral holes is as
  twice as the number of packed anions (S2).
• Half of these holes in ZnS is occupied by Zn2
  cations to achieve neutrality.

67
Ionic Solids
Chapter 10 Section 7

• For the fcc unit cell in the ccp packed
  structure, number of octahedral holes is the same
  as the number of packed ions.
• Example NaCl
• The Na ions fill all the four octahedral holes.

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Types and Properties of Solids
Chapter 10 Section 7
69
Vapor Pressures
Chapter 10 Section 8

• Vaporization (Evaporation)
• Change of state from liquid to gas.
• Molecules try to escape the liquids and form a
  gas.
• Energy is absorbed by the liquid to overcome the
  intermolecular forces to form a gas.
• It is an endothermic process.
• Heat (or Enthalpy) of Vaporization ?Hvap It is
  the energy required to vaporize 1 mole of a
  liquid at a pressure of 1 atm.
• ?Hvap for H2O _at_ 100 ºC is 40.7 kJ/mol

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Vapor Pressures
Chapter 10 Section 8

• Liquid in a close container
• Evaporation of the molecules at a constant rate
  takes place and the amount of liquid decreases.
• The condensation rate (returning to the liquid
  state) of the molecules begins to increase.

• When the evaporation and condensation processes
  exactly balance each other, the system is then at
  equilibrium. The liquid stops decreasing and the
  system is said to be highly dynamic.

71
Vapor Pressures
Chapter 10 Section 8

• Liquid in a close container
• Evaporation of the molecules at a constant rate
  takes place and the amount of liquid decreases.
• The condensation rate (returning to the liquid
  state) of the molecules begins to increase.

• When the evaporation and condensation processes
  exactly balance each other, the system is then at
  equilibrium. The liquid stops decreasing and the
  system is said to be highly dynamic.
• The vapor pressure at equilibrium is known as the
  vapor pressure.

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Vapor Pressures
Chapter 10 Section 8

• Measuring the vapor pressure of a liquid using a
  simple barometer.

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Vapor Pressures
Chapter 10 Section 8

• Patmosphere Pvapor PHg Column
• Pvapor Patmosphere PHg Column

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Vapor Pressures
Chapter 10 Section 8
Weaker intermolecular forces
More volatile
Higher vapor pressure
The hydrogen bonding here overcomes the effect of
molar masses
75
Effect of Temperature on Vapor Pressures
Chapter 10 Section 8

• Vapor pressure increases significantly with
  temperature.

• T1 gt T2
• A specific amount of KE is needed to overcome
  the intermolecular forces in the liquid phase.
• At the higher temperature, T2, the fraction of
  molecules having the minimum KE to convert into
  gas increases.

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Effect of Temperature on Vapor Pressures
Chapter 10 Section 8
77
Effect of Temperature on Vapor Pressures
Chapter 10 Section 8
78
Effect of Temperature on Vapor Pressures
Chapter 10 Section 8
y mx b
y
x
m
b
79
Effect of Temperature on Vapor Pressures
Chapter 10 Section 8

• From the adjacent plot, predict whether water or
  diethyl ether has larger enthalpy of vaporization
  (?Hvap).
• mH2O gt m(C2H5)2O
• ?HvapH2O gt ?Hvap(C2H5)2O

80
Effect of Temperature on Vapor Pressures
Chapter 10 Section 8

• Clausius-Clapeyron equation can be obtained from
  the previous relationship
• Example 10.6
• The vapor pressure and heat of vaporization of
  water at 25ºC is 23.8 torr and 43.9 kJ/mol,
  respectively. Calculate the vapor pressure of
  water at 50ºC.

81
Vapor Pressure of a Solid
Chapter 10 Section 8

• Sublimation a process in which the substance
  goes directly from the solid to the gaseous
  states.

Dry Ice (CO2)
Iodine (I2)
Moth Balls
82
Changes of State
Chapter 10 Section 8

• The changes in the state of a matter can be
  represented by a heating curve, which is a plot
  of temperature vs. time for a process when an
  energy is added in a constant rate to a specific
  amount of a substance.

83
Changes of State
Chapter 10 Section 8
Boiling point, heat of vaporization (?Hvap)
H2O gas with slope different than that for water
The plateau here is 7 time longer than that at
the melting point
The molecule is completely changed to vapor
Melting point, heat of fusion (?Hfus)
Water molecules gain more energy, and gradually
the hydrogen bonding becomes weaker and weaker
Ice molecules vibrate
The molecule is completely changed to liquid
84
Changes of State
Chapter 10 Section 8
85
Definition of Melting Point
Chapter 10 Section 8

• Below 0ºC, liquid water is supercooled.
• Above 0ºC, the vapor pressure curve of ice is
  extrapolated.
• The melting point is when a substance has
  identical vapor pressures at the solid and liquid
  phases.
• Ice is more temperature-dependent.

86
Vapor Pressures of Ice and Water
Chapter 10 Section 8

• An apparatus that allows solid and liquid water
  to interact only through the vapor state

87
Vapor Pressures of Ice and Water
Chapter 10 Section 8

• Case I
• With T when the vapor pressure of the solid is
  greater than that of the liquid.
• As solid releases more vapor the liquid absorbs
  more vapor for both states to achieve equilibrium.

The liquid begins to increase and solid begins
to decrease. The net effect is conversion of all
solid ice into liquid. Eventually, liquid water
will exist in equilibrium with water vapor. The
T here must be above the melting point.
88
Vapor Pressures of Ice and Water
Chapter 10 Section 8

• Case II
• With T when the vapor pressure of the solid is
  less than that of the liquid.
• This case is the opposite to Case I.

The supercooled liquid begins to decrease and
solid begins to increase. The net effect is
conversion of all liquid water into solid
ice. Eventually, solid ice will exist in
equilibrium with water vapor. The T here must be
below the melting point.
89
Vapor Pressures of Ice and Water
Chapter 10 Section 8

• Case III
• With T when the vapor pressure of the solid and
  liquid are the same.
• Here, both states have an identical vapor
  pressure and can coexist (both exist).

This temperature represents the freezing point..
Both states are in equilibrium with each other.
90
Melting Point and Boiling Point
Chapter 10 Section 8

• Melting Point is the temperature at which the
  solid and liquid states have the same vapor
  pressure under condition where the total pressure
  is 1 atm.

• Normal Boling Point is the temperature at which
  the vapor pressure of the liquid is exactly 1
  atm.
• In this close system, no bubbles can form within
  the liquid as long as the vapor pressure of water
  is less than 1 atm.

91
Supercooling of Water
Chapter 10 Section 8

• It is when the water can be cooled below 0ºC at 1
  atm without being solidified.
• This can happen when the water molecules did not
  achieve the degree of organization necessary to
  form ice at 0ºC.
• As soon as that order of organization occurs,
  the liquid crystallizes and releases energy.

92
Phase Diagram
Chapter 10 Section 9

• Represents phases for a closed system as a
  function of temperature and pressure.
• P 1.00 atm
• Melting point (Tm).
• Normal boiling point (Tb).

The phase diagram for water
93
Phase Diagram
Chapter 10 Section 9

• Ps 2 torr 0.0025 atm
• Sublimation process will take place at
  10 ºC. It occurs when the ice vapor pressure is
  equal to the external pressure (2 torr in this
  case).

Ps
The phase diagram for water
94
Phase Diagram
Chapter 10 Section 9

• P3 4.58 torr 0.0060 atm
• Triple point is the temperature and pressure at
  which solid, liquid and gas coexist.
• For H2O, this point is at 4.58 torr and 0.01ºC.
• The reason is because the three states have
  identical vapor pressures that are equal to the
  external pressure.

Ps
The phase diagram for water
95
Phase Diagram
Chapter 10 Section 9

• Pc 218 atm
• Critical point is the point at which the
  transition from liquid to vapor involves fluid
  intermediate region
• Critical temperature is the temperature above
  which the vapor cant be liquefied, no matter
  what pressure is applied.
• Critical pressure is the pressure required to
  liquefy the vapor at the critical temperature.

Ps
The phase diagram for water
96
Phase Diagrams of H2O and CO2
Chapter 10 Section 9
The phase diagram for water
The phase diagram for CO2
97
Phase Diagrams of H2O and CO2
Chapter 10 Section 9
? Fire extinguishers CO2 is compresses as liquid
inside the cylinder at high P. Once the CO2 is
released through the hose it is directly
converted into vapor. The liquid/vapor transition
is highly endothermic reaction, so cooling will
result.
The phase diagram for CO2
98
Phase Diagram of H2O
Chapter 10 Section 9

• Unlike most of the substances, when solid ice is
  compressed, it is converted into water. This is
  because the density of ice is less than water at
  the melting point.

99
Phase Diagram of H2O
Chapter 10 Section 9

• Application of that special physical properties
  of water
• Ice skating (more pressure from the skaters
  blade causes the ice to convert into water).
• Pipe freezing causes them to crack.
• Ice formed on lakes floats.