Chapter 11 Intermolecular Forces, Liquids, and Solids

1
Chapter 11Intermolecular Forces, Liquids, and
Solids
Chemistry, The Central Science, 10th
edition Theodore L. Brown H. Eugene LeMay, Jr.
and Bruce E. Bursten

• John D. Bookstaver
• St. Charles Community College
• St. Peters, MO
• ? 2006, Prentice Hall, Inc.

2
October 25 Chapter 11-Section 1 and 2

• HW p 476 1and 2
• 13 to 21 and 26 odd
• LAB REPORT DUE Tuesday Oct 30!!!!
• Ask any question by Monday!
• TEST ON CHAPTER 1011 Thursday November 1.
• PRINT SLIDES!

3
States of Matter

• The fundamental difference between states of
  matter is the distance between particles.

4
States of Matter

• Because in the solid and liquid states particles
  are closer together, we refer to them as
  condensed phases.

5
The States of Matter

• The state a substance is in at a particular
  temperature and pressure depends on two
  antagonistic entities
• The kinetic energy of the particles
• The strength of the attractions between the
  particles

6

• Converting a gas into a liquid or solid requires
  the molecules to get closer to each other
• cool or compress.
• Converting a solid into a liquid or gas requires
  the molecules to move further apart
• heat or reduce pressure.
• The forces holding solids and liquids together
  are called intermolecular forces.

7

• The covalent bond holding a molecule together is
  an intramolecular force.
• The attraction between molecules is an
  intermolecular force.
• Intermolecular forces are much weaker than
  intramolecular forces (e.g. 16 kJ/mol vs. 431
  kJ/mol for HCl).
• When a substance melts or boils the
  intermolecular forces are broken (not the
  covalent bonds).

8
Intermolecular Forces

• The attractions between molecules are not nearly
  as strong as the intramolecular attractions that
  hold compounds together.

9
Intermolecular Forces

• They are, however, strong enough to control
  physical properties such as boiling and melting
  points, vapor pressures, and viscosities.

10
Intermolecular Forces

• These intermolecular forces as a group are
  referred to as van der Waals forces.

11
Van der Waals Forces(Forces of attraction
between molecules)

• Dipole-dipole interactions
• Hydrogen bonding
• London dispersion forces

12
Ion-Dipole Interactions

• A fourth type of force, ion-dipole interactions
  are an important force in solutions of ions.
• The strength of these forces are what make it
  possible for ionic substances to dissolve in
  polar solvents.

13
Dipole-Dipole Interactions

• Molecules that have permanent dipoles are
  attracted to each other.
• The positive end of one is attracted to the
  negative end of the other and vice-versa.
• These forces are only important when the
  molecules are close to each other.

14
Dipole-Dipole Interactions

• The more polar the molecule, the higher is its
  boiling point.

15
How Do We Explain This?

• The nonpolar series (SnH4 to CH4) follow the
  expected trend.
• The polar series follows the trend from H2Te
  through H2S, but water is quite an anomaly.

16
Hydrogen Bonding

• The dipole-dipole interactions experienced when H
  is bonded to N, O, or F are unusually strong.
• We call these interactions hydrogen bonds, which
  are a special case of dipole-dipole attractions
• IT HAPPENS IN MOLECULES THAT CONTAIN A HYDROGEN
  ATOM DIRECTLY BONDED TO
• F, O or N

17
Why H bonds form?

• The Hydrogen atom consists of one proton and one
  electron. When its electron is attracted by the
  most electronegative elements of the table, what
  is left is an almost naked proton which in
  turns attracts the negative part of the molecule.
• The energies of H bonds range 4 to 25 kJ/mol.
  Regular covalent bond strength range between 150
  and 1100 kJ/mol.
• H bonds are of fundamental importance in
  biological molecules, they help to stabilize the
  structure of proteins, and are responsible for
  the way DNA carries genetic information

18

• Hydrogen bonds are also responsible for
• Ice floating, fish surviving winter and many
  broken pipes!
• Solids are usually more closely packed than
  liquids
• therefore, solids are more dense than liquids.
• Ice is ordered with an open structure to optimize
  H-bonding, therefore, ice is less dense than
  water.
• Ice density 0.917 g/ml
• Water density 1 g/ml
• Ice is an unusual solid because it has a larger
  volume in the solid state than in the liquid
  state!

19
Hydrogen Bonding

• In water the H-O bond length is 1.0 Å.
• The OH hydrogen bond length is 1.8 Å.
• Ice has waters arranged in an open, regular
  hexagon.
• Each ? H points towards a lone pair on O.
• Each water forms H bonds with 4 other molecules

When water melts the structure collapses and the
molecules move closer, then the liquid is less
dense
20
Section 2 and 3

• London Dispersion Forces
• Comparing intermolecular forces
• Properties of liquids
• Viscosity
• Surface Tension
• HW p 476 3, 5
• 13 to 31 (odd)

21
London Dispersion Forces
Are due to momentary or instantaneous dipoles in
a molecule

• While the electrons in the 1s orbital of helium
  would repel each other (and, therefore, tend to
  stay far away from each other), it does happen
  that they occasionally wind up on the same side
  of the atom.

22
London Dispersion Forces

• At that instant, then, the helium atom is polar,
  with an excess of electrons on the left side and
  a shortage on the right side.

23
London Dispersion Forces

• Another helium nearby, then, would have a dipole
  induced in it, as the electrons on the left side
  of helium atom 2 repel the electrons in the cloud
  on helium atom 1.

24
London Dispersion Forces

• London dispersion forces, or dispersion forces,
  are attractions between an instantaneous dipole
  and an induced dipole.

25
London Dispersion Forces

• These forces are present in all molecules,
  whether they are polar or nonpolar, BUT THEY ARE
  THE ONLY ONES THAT EXIST IN NON POLAR
  MOLECULES!!!
• The tendency of an electron cloud to distort in
  this way is called polarizability.

26
Factors Affecting London Forces

• The polarizability of a molecule can be seen
  as a measure of the squishiness of its electron
  cloud. The greater the polarizability of the
  molecule, the more easily its electron cloud can
  be distorted to give a MOMENTARY dipole. LARGER
  MOLECULES TEND TO HAVE GREATER POLARIZABILITIES.
  Polarizability and dispersion forces increases
  with molecular mass

27
Factors Affecting London Forces

• 1.-The shape of the molecule affects the strength
  of dispersion forces long, skinny molecules
  (like n-pentane tend to have stronger dispersion
  forces than short, fat ones (like neopentane-
  dimethyl propane).This is due to the increased
  surface area in n-pentane.
• 2.The size of the molecule. Larger molecules have
  a greater number of electrons and those electrons
  are far away from the nucleus so they can be more
  easily polarized.

28
Summary of London Dispersion Forces

• Polarizability is the ease with which an electron
  cloud can be deformed.
• The larger the molecule (the greater the number
  of electrons) the more polarizable.
• London dispersion forces increase as molecular
  weight increases.
• London dispersion forces exist between all
  molecules.
• London dispersion forces depend on the shape of
  the molecule.
• Like all dipole-dipole forces are significant
  only if molecules are very close together.

29

• Predict the boiling point of Halogens and Noble
  Gases based on the molecular mass.

30
Factors Affecting London Forces

• The strength of dispersion forces tends to
  increase with increased molecular weight.
• Larger atoms have larger electron clouds, which
  are easier to polarize.

31
Which Have a Greater EffectDipole-Dipole
Interactions or Dispersion Forces?

• If two molecules are of comparable size and
  shape, dipole-dipole interactions will likely be
  the dominating force.
• If one molecule is much larger than another,
  dispersion forces will likely determine its
  physical properties (the gt molecular mass the
  stronger the attractions).

32
Summarizing Intermolecular Forces
33

• Examples determine the types of forces present
  in each
• H2O
• CCl4
• SO2
• LiF
• Ca(NO3)2 aqueous solution
• HF
• PCl3

34

• Examples determine the types of forces present
  in each
• H2O LDF, dipole-dipole, H-bonds
• CCl4 LDF
• SO2 LDF and dipole-dipole
• LiF ionic bonds
• Ca(NO3)2 aqueous solution ion-dipole forces
• HF LDF, dipole-dipole, H-bonds
• PCl3 LDF and dipole-dipole

35
Relative Strengths of Forces

1. Bonds ( ionic, covalent, metallic)
2. Ion-dipole forces
3. Hydrogen bonds
4. Dipole-dipole forces
5. London dispersion forces

36
Solution (a) Dipole-dipole attractions increase
in magnitude as the dipole moment of the molecule
increases. Thus, CH3CN molecules attract each
other by stronger dipole-dipole forces than CH3I
molecules do. (b) When molecules differ in their
molecular weights, the more massive molecule
generally has the stronger dispersion
attractions. In this case CH3I (142.0 amu) is
much more massive than CH3CN (41.0 amu), so the
dispersion forces will be stronger for CH3I. (c)
Because CH3CN has the higher boiling point, we
can conclude that more energy is required to
overcome attractive forces between CH3CN
molecules. Thus, the total intermolecular
attractions are stronger for CH3CN, suggesting
that dipole-dipole forces are decisive when
comparing these two substances. Nevertheless,
dispersion forces play an important role in
determining the properties of CH3I.
PRACTICE EXERCISE Of Br2, Ne, HCl, HBr, and N2,
which is likely to have (a) the largest
intermolecular dispersion forces, (b) the largest
dipole-dipole attractive forces?
Answers (a) Br2 (largest molecular weight), (b)
HCl (largest polarity)
37
(No Transcript)
38
PRACTICE EXERCISE In which of the following
substances is significant hydrogen bonding
possible methylene chloride (CH2Cl2) phosphine
(PH3) hydrogen peroxide (HOOH), or acetone
(CH3COCH3)?
Answer HOOH
39
Check The actual normal boiling points are H2
(20 K), Ne (27 K), CO (83 K), HF (293 K), and
BaCl2 (1813 K), in agreement with our predictions.
PRACTICE EXERCISE (a) Identify the intermolecular
forces present in the following substances, and
(b) select the substance with the highest boiling
point CH3CH3, CH3OH, and CH3CH2OH.
Answers (a) CH3CH3 has only dispersion forces,
whereas the other two substances have both
dispersion forces and hydrogen bonds (b) CH3CH2OH
40
Some Properties of Liquids

• Viscosity
• Viscosity is the resistance of a liquid to flow.
• A liquid flows by sliding molecules over each
  other.
• The stronger the intermolecular forces, the
  higher the viscosity. At higher temperatures
  viscosity decreases.

41
Viscosity is measured by

• Timing how long it takes for a liquid to flow
  through a thin tube under gravitational force.
• or
• By the rate at which steel spheres fall through
  the liquid

42
Surface Tension

• Surface tension results from the net inward
  force experienced by the molecules on the surface
  of a liquid.

43
Surface Tension

• Surface molecules are only attracted inwards
  towards the bulk molecules.

44

• Surface Tension
• Is due to an unbalance of intermolecular forces
  at the surface of the liquid. Molecules in the
  interior are attracted equally in all directions,
  but at the surface they experience an inward
  force.
• The inward pull reduces the surface area, making
  the molecules at the surface pack closely. This
  reduces the surface area. Spheres have the
  smaller surface per volume, then water drops
  asume spherical shape.
• .Surface tension is the amount of energy
  required to increase the surface area of a
  liquid.
• Water has high surface tension due to H bonding.

45

• Cohesive forces bind molecules to each other
  (intermolecular forces).
• Adhesive forces bind molecules to a surface.
  (water to glass or to paper)
• Meniscus is the shape of the liquid surface.
• If adhesive forces are greater than cohesive
  forces, the liquid surface is attracted to its
  container more than the bulk molecules.
  Therefore, the meniscus is U-shaped (e.g. water
  in glass).
• If cohesive forces are greater than adhesive
  forces, the meniscus is curved downwards (Hg)

46
(No Transcript)
47
Capillary Action

• Is the rise of a liquid up very narrow tubes.
  This phenomenon occurs because the adhesive
  forces between the liquid and the walls of the
  container (glass) tend to increase the surface
  area. The surface tension of the liquid pulls the
  molecules up in order to reduce the area and the
  liquid climbs until the adhesive and cohesive
  forces are balanced by the force of gravity.

48
Section 4 Phase changes Section 5 Vapor pressure

• HW P 479 Q 33-35-39-37
• Specific Heat of a Substance
• Energy Changes Accompanying Phase Changes
  /Heating and Cooling Curves
• Supercooled liquids

49
Phase Changes

• Examples Name each of the following phase
  changes
• solid ? gas
• liquid ? gas
• solid ? liquid
• gas ? solid
• gas ? liquid
• liquid ? solid

50
Phase Changes

• solid ? gas sublimation
• liquid ? gas vaporization (boiling)
• solid ? liquid melting (fusion)
• gas ? solid deposition
• gas ? liquid condensation
• liquid ? solid freezing (solidification)

51

• Energy Changes Accompanying Phase Changes
• Indicate endo or exhotermic change
• Heat or Enthalpy ?H
• Sublimation ?Hsub
• Vaporization ?Hvap
• Melting or Fusion ?Hfus
• Deposition ?Hdep
• Condensation ?Hcon
• Freezing ?Hfre

52

• Energy Changes Accompanying Phase Changes
• Sublimation ?Hsub gt 0 (endothermic).
• Vaporization ?Hvap gt 0 (endothermic).
• Melting or Fusion ?Hfus gt 0 (endothermic).
• Deposition ?Hdep lt 0 (exothermic).
• Condensation ?Hcon lt 0 (exothermic).
• Freezing ?Hfre lt 0 (exothermic).

53
VAPORIZATION IS ENDOTHERMIC

• In hot climates drinking water is cooled by
  evaporating water from the surfaces of porous
  clay pots. As water evaporates it ABSORBS heat
  from the water inside the container which is
  maintained cool.
• Like cooling yourself off on a hot day by pouring
  water over your body. As water evaporates it
  absorbs heat

54
(No Transcript)
55
(No Transcript)
56
FREEZING IS EXOTHERMIC

• In freezing weather, citrus crops are sprayed
  with water to protect the fruit from frost
  damage. As the water freezes (around the
  fruit-outside the fruit!) it releases heat, which
  helps to prevent the fruit from freezing.

57
(No Transcript)
58
Phase Changes
59
Energy Changes Associated with Changes of State

• The heat added to the system at the melting and
  boiling points goes into pulling the molecules
  farther apart from each other.
• The temperature of the substance does not rise
  during the phase change.

60

• POTENTIAL ENERGY STORED ENERGY. The energy
  inside the substance.
• KINETIC ENERGY Associated with motion.
• Average KE TEMPERATURE

61
Endothermic Phase Changes

• If the substance is melting or boiling, heat is
  being absorbed, and is being used to change the
  state of matter.
• THE AVERAGE KINETIC ENERGY DOES NOT CHANGE!!!
  THE POTENTIAL ENERGY INCREASES.

62
Exothermic changes

• If the substance is undergoing condensation or
  freezing then heat energy is being released. The
  potential energy is decreasing and the
  TEMPERATURE REMAINS CONSTANT!!!

63
Heat or Enthalpy of Fusion ?Hfus

• Amount of heat needed to completely melt 1
  gram of substance at its melting point.
• For water the value is 334 J/g or 6.01 kJ/mol

64
Heat or Enthalpy of Vaporization ?Hvap

• Heat of vaporization amount of heat needed to
  completely convert 1 g of liquid to gas.
• For water the value is
• 2260 J/g or 40.7 kJ/mol

65
Enthalpies of Fusion vs Enthalpies of
Vaporization for several substances
66

• Generally heat of fusion (enthalpy of fusion) is
  less than heat of vaporization.
• it takes more energy to completely separate
  molecules (from liquid to gas), than partially
  separate them (from solid to liquid).

67
ProblemsHow to calculate heat

• Review of basic calorimetry problems
• Amount of Heat Q D H enthalphy
• C Specific Heat
• Q mass x D T x C
• This formula can be used when there is a change
  in T.
• During a phase change use the heat of fusion or
  the heat of vaporization

68

• How much energy is required to raise the
  temperature of 1800. g ice at 0C to 10C? DHfus
  6.01 kJ/mol, heat capacity of water is 75.2
  J/mol-K.

69
1800 g 100 moles of ice The enthalpy of fusion
(or heat of fusion) DH (6.01 kJ/mol)(100 mol)
601 kJ. To raise the water temperature 10C
requires q (75.2 J/mol-K)(100 mol)(10C )
75.2 kJ. Total energy 601 kJ 75 kJ 676 kJ
70
(No Transcript)
71
Problem 38 p 479

• Chlorofluorocarbons of CFCs were used as
  refrigerants.
• The heat of vaporization of CCl2F2 is 289 J/g.
  What mass of this substance must evaporate in
  order to freeze 100g of water initially at 18 0C?
  Heat of fusion for water is 334 J/g and its
  specific heat is 4.18 J/gC

72
Plan

• 1.Find the heat that water needs to release to
  cool down from 18 0C to 0 0C
• 2. Find the heat that the water needs to loose to
  freeze.
• 3. Add the results, that is the total amount of
  heat the water has to lose.
• 4. Use the heat of evaporation of CCl2F2 to find
  the mass needed to absorb the heat found in 3.

73
Supercooling

• When a liquid is cooled down so fast that the
  molecules do not have time to accommodate to
  their regular structure we get a supercooled
  liquid.
• It is basically a liquid below the FP of it. It
  is unstable and it freezes suddenly if a dust
  particle enters the liquid or by shaking or
  stirring it.

74
Vapor Pressure

• At any temperature, some molecules in a liquid
  have enough energy to escape.
• As the temperature rises, the fraction of
  molecules that have enough energy to escape
  increases.

75
Vapor Pressure

• The pressure exerted by the vapor of a liquid
  when vapor and liquid states are in dynamic
  equilibrium.
• Dynamic equilibrium 2 opposing processes
  occurring simultaneously with NO NET CHANGE
  OBSERVED!!!

76
Vapor Pressure

• As more molecules escape the liquid, the
  pressure they exert increases.

77
Vapor Pressure

• The liquid and vapor reach a state of dynamic
  equilibrium liquid molecules evaporate and
  vapor molecules condense at the same rate.

78
Vapor Pressure

• Explaining Vapor Pressure on the Molecular Level
• Some of the molecules on the surface of a liquid
  have enough energy to escape the attraction of
  the bulk liquid.
• These molecules move into the gas phase.
• As the number of molecules in the gas phase
  increases, some of the gas phase molecules strike
  the surface and return to the liquid.
• After some time the pressure of the gas will be
  constant at the vapor pressure.

79
Property Stronger forces mean
Viscosity
Surface tension
Melting point -freezing
Boiling point (condensation)
?Hfus
?Hvap
Vapor Pressure
80
Property Stronger forces mean
Viscosity higher
Surface tension higher
Melting point -freezing higher
Boiling point (condensation) higher
?Hfus higher
?Hvap higher
Vapor Pressure lower
81

• Vapor Pressure and Boiling Point
• Liquids boil when the external pressure equals
  the vapor pressure.
• Temperature of boiling point increases as
  pressure increases.
• Two ways to get a liquid to boil increase
  temperature or decrease pressure.
• Pressure cookers operate at high pressure. At
  high pressure the boiling point of water is
  higher than at 1 atm. Therefore, there is a
  higher temperature at which the food is cooked,
  reducing the cooking time required.
• Normal boiling point is the boiling point at 760
  mmHg
• (1 atm).

82
Boiling Point

• The temperature at which its vapor pressure
  equals the external pressure.
• The normal boiling point is the temperature at
  which its vapor pressure is 760 torr.

83
Volatile Liquids

• High vapor pressure Low attraction between
  molecules
• REMEMBER VAPOR PRESSURE DEPENDS ON TEMPERATURE!

84
Section 6,7 8

• Phase diagrams P480 Q 41 -42
• Critical Temperature and Pressure
• HW Q 41-42 51 to 55 odd only
• SOLIDS CLASSIFICATION AND BONDING

85

• Well begin with chapter 6 Atomic Structure
  next Monday. Print slides on Sunday.
• Structure of solids
• Bonding in solids
• HW 57, 71, 73, 75, 77

86
Phase Diagrams

• Phase diagram plot of pressure vs. Temperature
  summarizing all equilibria between phases.
• Given a temperature and pressure, phase diagrams
  tell us which phase will exist.
• Any temperature and pressure combination not on a
  curve represents a single phase.

87

• Features of a phase diagram
• Triple point temperature and pressure at which
  all three phases are in equilibrium.
• Vapor-pressure curve generally as pressure
  increases, temperature increases.
• Critical point critical temperature and pressure
  for the gas.
• Melting point curve as pressure increases, the
  solid phase is favored if the solid is more dense
  than the liquid.
• Normal melting point melting point at 1 atm.

88
Phase Diagrams
89
Liquid-vapor equilibrium

• The AB line is the liquid-vapor interface.
• It starts at the triple point (A), the point at
  which all three states are in equilibrium.

90

• Each point along this line is the boiling point
  of the substance at that pressure.

91
Critical Temperature

• The highest temperature at which a liquid can
  form/exist.
• Above the critical temperature no liquid can
  exist.
• The greater the intermolecular forces the greater
  the Tcrit

92
Critical Pressure

• The pressure required to liquify the gas at the
  critical temperature.
• When doing problems with Critical Temperatures
  and Pressure CHECK THE UNITS!!!!

93
Supercritical Fluid Extraction

• Substances at temperature and pressures higher
  than the critical T and P( P several hundred atm)
  are supercritical fluids. At these conditions the
  they behave like special gases because their
  densities are similar to the liquids.
  Supercritical fluids can be good solvents and
  lowering the P or increasing T changes the
  solubility and that can be used to separate
  mixtures.
• Supercritical CO2 is used to decaffeinate coffee.

94
Critical Temperature and Pressure
95
End of the liquid line Tc and Pc

• It ends at the critical point (B) above this
  critical temperature and critical pressure the
  liquid and vapor are indistinguishable from each
  other.

96
Liquid-Solid equilibrium

• The AD line is the interface between liquid and
  solid.
• The melting point at each pressure can be found
  along this line.

97
Solid-Gas equilibrium

• Below A the substance cannot exist in the liquid
  state.
• Along the AC line the solid and gas phases are in
  equilibrium the sublimation point at each
  pressure is along this line.

98
Phase Diagram of Water

• Note the high critical temperature and critical
  pressure
• These are due to the strong van der Waals forces
  between water molecules.

99

• Remember Le Chateliers principle, increasing the
  pressure shifts the equilibrium to the side that
  occupies less volume. Since ice occupies more
  volume than liquid when we increase the pressure
  ice melts! (the equilibrium shifts to the water).
  Most substances are the opposite because the
  solid phase is more dense that the liquid phase.

100
Phase Diagram of Water

• The slope of the solidliquid line is negative,
  tilted to the left.
• This means that as the pressure is increased at a
  temperature just below the melting point, water
  goes from a solid to a liquid.

101

• For most substances the solid form is denser than
  the liquid form and an increase in P favors the
  solid, so to change it to liquid more T is
  needed. This results in an slope towards the
  right (positive slope).

102

• The slope to the left is abnormal and is due to
  the fact that water is less dense in the solid
  state (why???), so increasing the pressure favors
  the liquid state.

103
Phase Diagram of Carbon Dioxide

• The low critical temperature and critical
  pressure for CO2 make supercritical CO2 a good
  solvent for extracting nonpolar substances (such
  as caffeine).

104
Phase Diagram of Carbon Dioxide

• Carbon dioxide cannot exist in the liquid state
  at pressures below 5.11 atm CO2 sublimes at
  normal pressures.

105

• Water
• Why does the melting point curve slope to the
  left?
• What are the temperature and pressure at the
  triple point?
• What are the normal freezing and boiling points?
• What are the critical temperature and pressure?
• What change occurs at 50?C as the pressure is
  decreased from 1.0 atm to 0.0010 atm?

106

• Water
• Why does the melting point curve slope to the
  left?
• ice is less dense than water
• What are the temperature and pressure at the
  triple point?
• 0.0098?C and 4.58 mmHg
• What are the normal freezing and boiling points?
• Freezing 0 ?C and Boiling 100 ?C
• What are the critical temperature and pressure?
• 374?C and 218 atm
• What change occurs at 50?C as the pressure is
  decreased from 1.0 atm to 0.0010 atm?
• vaporization

107

• Carbon Dioxide
• At what temperature and pressure does the triple
  point occur?
• What is the normal sublimation point?
• What is the critical point?
• What change occurs at 30. atm as you move from
  -60C to 0C?

108

• Carbon Dioxide
• At what temperature and pressure does the triple
  point occur?
• -56.4?C and 5.11 atm
• What is the normal sublimation point?
• -78.5?C
• What is the critical point?
• 31.1?C and 73 atm
• What change occurs at 30. atm as you move from
  -60C to 0C?
• melting

109
Structures of Solids

• Unit Cells
• Crystalline solid well-ordered, definite
  arrangements of molecules, atoms or ions.
• Crystals have an ordered, repeated structure.
• The smallest repeating unit in a crystal is a
  unit cell.
• Unit cell is the smallest unit with all the
  symmetry of the entire crystal.
• Three-dimensional stacking of unit cells is the
  crystal lattice.

110

• Unit Cells

111

• Unit Cells
• Three common types of unit cell.
• Primitive cubic, atoms at the corners of a simple
  cube,
• each atom shared by 8 unit cells
• Body-centered cubic (bcc), atoms at the corners
  of a cube plus one in the center of the body of
  the cube,
• corner atoms shared by 8 unit cells, center atom
  completely enclosed in one unit cell
• Face-centered cubic (fcc), atoms at the corners
  of a cube plus one atom in the center of each
  face of the cube,
• corner atoms shared by 8 unit cells, face atoms
  shared by 2 unit cells.

112
(No Transcript)
113
Solids

• We can think of solids as falling into two
  groups
• Crystallineparticles are in highly ordered
  arrangement.
• Specific melting points

114
Solids

• Amorphousno particular order in the arrangement
  of particles.
• Melt at a range of temperatures not at specific
  temperature
• Example Glass

115
Crystalline Solids

• Because of the order in a crystal, we can focus
  on the repeating pattern of arrangement called
  the unit cell.

116
The Crystal Structure of Sodium Chloride
117

• The unit cell is the smallest repeating unit
  that has all of the symmetry characteristic of
  the way atoms/ions or molecules are arranged in
  the crystal.
• It reflects the Stoichiometry of the solid.

118
Bonding in Solids

• There are four types of solid
• Molecular (formed from molecules) - usually soft
  with low melting points and poor conductivity.
• Covalent network (formed from atoms) - very hard
  with very high melting points and poor
  conductivity.
• Ionic (formed from ions) - hard, brittle, high
  melting points and poor conductivity.
• Metallic (formed from metal atoms) - soft or
  hard, high melting points, good conductivity,
  malleable and ductile.

119
(No Transcript)
120

• Molecular Solids
• Intermolecular forces dipole-dipole, London
  dispersion and H-bonds.
• Weak intermolecular forces give rise to low
  melting points.
• Room temperature gases and liquids usually form
  molecular solids at low temperature.
• Efficient packing of molecules is important
  (since they are not regular spheres).

121
Covalent-Network Solids

• Forces covalent bonds.
• Atoms held together in large networks.
• Examples diamond, graphite, quartz (SiO2),
  silicon carbide (SiC), and boron nitride (BN).
• They tend to be hard and have high melting
  points.

122
Covalent-Network andMolecular Solids

• Diamonds are an example of a covalent-network
  solid in which atoms are covalently bonded to
  each other.

123
Diamond

• each C atom has a coordination number of 4 each
  C atom is tetrahedral there is a
  three-dimensional array of atoms.
• Diamond is hard, and has a high melting point
  (3550 ?C).
• contain orbitals or bands of delocalized
  electrons that belong not to single atoms but to
  each crystal as a whole

124
Covalent-Network andMolecular Solids

• Graphite is an example of a molecular solid in
  which atoms are held together with van der Waals
  forces.
• They tend to be softer and have lower melting
  points.

125

• In graphite
• each C atom is arranged in a planar hexagonal
  ring
• layers of interconnected rings are placed on top
  of each other
• the distance between C atoms is close to benzene
  (1.42 Å vs. 1.395 Å in benzene)
• the distance between layers is large (3.41 Å)
• electrons move in delocalized orbitals (good
  conductor).

126
Silicon dioxide

• Has a high melting point - around 1700C. Very
  strong silicon-oxygen covalent bonds have to be
  broken throughout the structure before melting
  occurs.
• It is hard. This is due to the need to break the
  very strong covalent bonds.
• Doesn't conduct electricity. There aren't any
  delocalized electrons. All the electrons are held
  tightly between the atoms, and aren't free to
  move.
• It is insoluble in water and organic solvents.
  There are no possible attractions which could
  occur between solvent molecules and the silicon
  or oxygen atoms which could overcome the covalent
  bonds in the giant structure.
• Giant covalent structures are arranged in a
  continuous lattice. This structure is very strong
  because of the strong forces between the
  molecules.

127

• Ionic Solids
• Ions (spherical) held together by electrostatic
  forces of attraction.
• There are some simple classifications for ionic
  lattice types.

128

• NaCl Structure
• Each ion has a coordination number of 6.
• Face-centered cubic lattice.
• Cation to anion ratio is 11.
• Examples LiF, KCl, AgCl and CaO.
• CsCl Structure
• Cs has a coordination number of 8.
• Different from the NaCl structure (Cs is larger
  than Na).
• Cation to anion ratio is 11.

129
Metallic Solids

• Metals are not covalently bonded, but the
  attractions between atoms are too strong to be
  van der Waals forces.
• In metals, valence electrons are delocalized
  throughout the solid.

130
Metals

• Closely packed lattice with delocalize electrons
  throughout
• The metal nuclei float in a sea of electrons.
• Metals conduct because the electrons are
  delocalized and are mobile

131
(No Transcript)