States of Matter Liquids and Solids

1
States of Matter Liquids and Solids

• Chapter 11

2
Contents

• Comparison of gases, liquids, and solids
• Changes of state
• Phase transitions
• Phase diagrams
• Liquid state
• Properties of liquid surface tension and
  viscosity
• Intermolecular forces

• Solid State
• Classification of solids by type of attraction
• Crystalline solids crystal lattices and unit
  cell
• Structures of some crystalline solids
• Unit cell dimensions
• Determination of crystal structure by X-ray
  diffraction

3
Comparison of Gases, Liquids and Solids

• Gases are compressible fluids. Molecules or atoms
  move in random directions in almost empty space.
• Liquids are incompressible fluids. The molecules
  in liquid make random motions but they are
  tightly packed.
• Solids are also imcompressible. Particles in
  solid exist in close contact, vibrate about fixed
  points.

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Changes of State

• Change of a substance from one state to another
  is called a change of state or phase transition.
• Phase transitions
• Melting
• Freezing
• Vaporization
• Sublimation

5
Phase Transitions
6
Vapor Pressure

• The vapor pressure of a liquid is the partial
  pressure of the vapor over the liquid, measured
  at equilibrium.

7
Equilibrium Vapor Pressure

• Initial rate of vaporization is higher than rate
  of condensation.
• When the rates of vaporization and condensation
  have become equal, the vapor pressure remains
  unchanged liquid gets its equilibrium vapor
  pressure.

Initial
Equilibrium
8
Kinetic Energy of Molecules and Evaporation

• Fraction of molecules having KE greater than
  minimum can escape from liquid.
• Fraction of molecules having enough KE to escape
  increases with temperature.

9
Rates of Vaporization and Condensation with Time

• Rate of reaction is proportional to concentration
  of substance
• Rate of vaporization is constant. Rate of
  condensation is zero initially, but increases to
  rate of vaporization and equilibrium is reached.

10
Variation of Vapor Pressure with Temperature

• Vapor pressure of liquids increases
  rapidly(exponentially) with temperature.
• When the vapor pressure of liquid becomes equal
  to external pressure, liquid boils.

11
Boiling Point and Melting Point

• The temperature at which the vapor pressure of a
  liquid equals the pressure exerted on the liquid
  is called the boiling point.
• The temperature at which a pure liquid changes to
  a crystalline solid, or freezes, is called the
  freezing point. A dynamic equilibrium establishes
  during freezing
• Solid ? liquid

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Heat of Phase Transition

• Heat is added to heat ice at 20 0C at constant
  rate.
• Heat of fusion is the heat needed to melt a
  solid.
• H2O(s) ? H2O(l) ?Hfus6.01 kJ/mol
• Heat of vaporization is the heat needed to
  vaporize a liquid.
• H2O(l) ? H2O(g) ?Hvap 40.7 kJ/mol
• Specific heat of water (l) is 4.18 J/(g.K)

13
Heat Required to Freeze Water

• The heat of vaporization of ammonia is 23.4
  kJ/mol. How much heat is required to vaporize
  1.00 kg of ammonia? How many grams of water at
  00C could be frozen to ice at 00C by evaporation
  of this amount of ammonia? ?Hvap(water) 6.01
  kJ/mol
• Solution

14
Vapor Pressure-Temperature Relationship

• The vapor pressure of a liquid is a function of
  temperature.

15
Clausius-Clapeyron Equation

• log P ? (1/T) relation is linear. A line can be
  defined by two points

• By subtracting the two equations and rearranging
  yields the Clausius-Clapeyron equation.

16
Vapor Pressure of Water at a Given Temperature

• Carbon disulfide, CS2, has a normal boiling point
  of 460C and a heat of vaporization of 26.8
  kJ/mol. What is the vapor pressure of carbon
  disulfide at 350C?
• Solution

• P 529 torr

17
Phase Diagrams

• A phase diagram gives the stable form of a
  substance at a given temperature and pressure.
• Melting-Point Curve
• Solid ? liquid
• Boiling-point curve
• Liquid ? vapor
• Triple point
• Solid ? liquid ? vapor
• At and above the critical temperature and
  critical pressure only gas exist

18
Phase Diagram of CO2
19
Liquid State

• The intermolecular forces are much stronger in
  the liquids and solids than gases. Their size
  depends on
• Distance, electrical force falls of rapidly with
  distance
• Chemical composition which enables substances to
  be liquid or solid for a given temperature (room
  T)
• Intermolecular Forces
• Dipole-dipole forces
• London forces
• Hydrogen bonding
• Ion-induced dipole forces

20
Surface Tension

• Due to intermolecular forces, the molecules at
  the surface experience a net attraction toward
  the interior of the liquid.
• Surface tension is the energy required to
  increase the surface area of a liquid by a unit
  amount.

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Viscosity

• Viscosity is the resistance to flow that is
  exhibited by liquids and gases
• The viscosity of a liquid can be measured by the
  time a given quantity of liquid to flow through a
  capillary tube or, alternatively, by the time it
  takes a steel ball of given radius fall through
  column of liquid.

22
Dipole-Dipole Forces

• The dipole-dipole force is an attractive
  intermolecular force that results from the
  alignment of polar molecules such that positive
  end of one molecule is closer to the negative end
  of other molecule.

23
London (Dispersion) Forces

• Electrons are in constant motion. An
  instantaneous dipole on any atom sets
  polarization on neighbor atoms.
• London forces (dispersion forces) are the weak
  attractive forces between molecules resulting
  from the small, instantanous dipoles formed due
  to electron motion.
• London forces increase with molecular weight.

24
Geometry of Molecules and London Forces

• For molecules of about the same molecular weight,
  the more compact one is probably less
  polarizable, so the London forces are smaller.
• London forces increases in the following order
• pentene gt isopentane gt neopentane

25
Hydrogen Bonding

• Hydrogen bonding is a moderate attractive force
  that exists between a hydrogen atom covalently
  bonded to any one of the three very small and
  highly electronegative atoms, F, O and N.
  Hydrogen bond is about five times stronger than
  other dipole-dipole interactions. The strength of
  H-bond gives water some very unusual and special
  properties.

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Ion-Dipole Interaction

• Ion-dipole interaction is between an ion and a
  polar molecule. Since an ion can produce induced
  dipole on nonpolar molecule leading to an
  ion-induced dipole interaction. Since the charge
  on the ion does not flicker on and off like the
  instantaneous charges responsible for London
  forces, these attractions can be quite strong.

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Identifying Intermolecular Forces

• What kind of intermolecular forces are expected
  between for the following compounds?
• Propanol,CH3CH2CH2OH
• Carbon dioxide, CO2
• Sulfur dioxide, SO2

• Solution
• London force H-bonding
• London force
• Dipole-dipole and London force

28
Liquid Properties and Intermolecular Forces

• The surface tension and viscosity of liquids
  increase with the increase in intermolecular
  force. They both decrease as temperature
  increases.
• The vapor pressure of a liquid is low if the
  intermolecular force between its molecules is
  strong.
• The normal boiling temperature of a liquid is
  higher if the intermolecular force between its
  molecules is stronger.
• The heat of vaporization of a liquid is larger if
  the intermolecular force between its molecules is
  stronger.

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Comparison of Intermolecular Forces

• London(dispersion) forces are found in all
  substances. The strength of these forces increase
  with increased molecular weight and also depend
  on molecular shape.
• Dipole-dipole forces add to the effect of
  dispersion force
• H-bonding is the strongest type of intermolecular
  force
• None of these forces is as strong as ordinary
  ionic and covalent bonds
• Question Explain the trend in normal boiling
  temperatures.

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Properties of Some Liquids at 200C
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Solid State

• Solids can be classified in four different groups
  according to the forces holding them.
• Molecular solids molecules are held together by
  intermolecular forces. Ice, solid CO2.
• Metallic solids positive cores of atoms held
  together by sea of electrons. Iron, copper.
• Ionic solids cations and anions held together by
  electrostatic forces. NaCl, ZnS.
• Covalent network solids atoms held together by
  covalent bonds. Diamond

32
Crystalline Solids

• A crystalline solid has a well defined ordered
  structure in three dimensions.
• An amorphous solid has a disordered structure. A
  glass is an amorphous solid obtained by cooling
  rapidly that the units are frozen randomly.

Diamond, network crystal
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Crystal Lattices

• A crystal lattice is the geometric arrangement of
  lattice points in three dimensions.
• A unit cell of a crystal is the smallest boxlike
  unit from which the whole crystall is formed by
  stacking unit cells.

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Unit Cells of Seven Basic Crystal Systems
abc
ab
?
c
?
?
??90
b
a
Cubic
Tetragonal
Orthorhombic
Monoclinic
abc
ab
?120
????90
Hexagonal
Rhombohedral
Triclinic
35
Cubic Unit Cells

• Simple Cubic lattice points are at the corners.
• Body-Centered Cubic lattice points are at the
  center and corners.
• Face-Centered Cubic lattice points are at the
  centers of each face and corners

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Properties of Solids

• Melting Point Molecular solids have low melting
  points. Ionic solids and covalent network solids
  melt at very high temperature. Chemical bonds
  should be broken to melt. NaCl 8010C, MgO
  28000C. Metallic solids melts at a broad
  temperature range. Hg -390C, W 34100C.
• Hardness Molecular solids are soft, whereas
  covalent and ionic crystals are very hard but
  brittle. Metalic crstals are malleable.
• Electrical Conductivity Delocalized valence
  electrons of metals, but molecular, ionic and
  covalent crystals do not conduct electricity.

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Crystal Defects
Perfect crystal
Na
Cl-
Cl-
Na

• Real crystals may have defects or imperfections.
  Two kinds of defects may occur during
  crystallization
• Crystal planes may be misaligned,
• Sites in the crystal lattice may remain vacant.

Cl-
Cl-
Na
Na
Na
Cl-
Cl-
Na
Cl-
Cl-
Na
Na
Real crystal
Na
Cl-
Na
Cl-
Cl-
Na
Na
Cl-
Cl-
Na
Cl-
Cl-
Na
Na
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Structures of Some Crystalline Solids

• Molecular Solids Closest Packing (hexagonal and
  cubic close-pack structures)
• Metallic Solids body centered cubic structure,
• Ionic Solids body-centered cubic, face-centered
  cubic, simple cubic,
• Covalent Network Solids face-centered cubic

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Molecular Solids Closest Packing

• Maximum attraction is obtained when all atoms
  touch each other. Each atom has 6 closest
  neighbors.

• Hexagonal closed packed ABABA

• Cubic closed packed ABCABC, fcc.

• Coordination number 12
• 74 of space is occupied

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Metallic Solids

• Metals are mostly crytallize in hexagonal or
  cubic close-packed structures (copper and
  silver). Spheres occupy 74 of space.
  Coordination number is 12.
• Metals with body-centered cubic structure are
  also present (Cr, Fe). Spheres occupy 68 of
  space. Coordination number is 8.

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Ionic Solids

• Crystal is formed by cations and anions. For
  general formula MX there are three types of
  crystal structures
• Cesium chloride (CsCl) body-centered cubic
• Sodium chloride (NaCl) face-centered cubic
• Zinc blende (ZnS) polymorphic (two forms)
• Face-centered cubic structure, zinc blende or
  sphalerite
• Hexagonal crystal, wurtzite

42
Covalent Network Solids

• There are three allotropic forms of carbon
  diamond, graphite and C-60.
• Diamond sp3 hybrid orbitals form zinc blende
  structure
• Graphite sp2 hybrid orbitals form ? and
  unhybridized p orbital forms the ? bond.

43
Calculations with Unit Cells

• Potassium metal has a body-centered cubic
  structure. The density of metal is 0.86 g/cm3.
  Calculate the edge of unit cell.
• Solution

• Since there are 2 atoms in the unit cell mass
  of unit cell is 2x6.48x10-23 1.3x10-22 g.
• Volume of unit cell 1.3x10-22 g / 0.86 g.cm-3
  1.51x10-22cm3
• Edge of unit cell (1.51x10-22cm3)1/3 5.3x10-8
  cm.

44
X-Ray Diffraction
45
Types of Unit Cells
46
Unit Cells for Some Common Lattices
Body-centered cubic
Face-centered cubic
47
Crystal Structure of Diamond and Graphite
142 pm
335 pm
Zinc Blende
Planar ? and delocalized ?
48
Close Packed Atoms