Chapter 11 Intermolecular Forces

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Chapter 11 Intermolecular Forces
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11.1 Intermolecular Forces (IMF)

• IMF lt intramolecular forces (covalent, metallic,
  ionic bonds)
• IMF strength solids gt liquids gt gases
• Boiling points and melting points are good
  indicators of relative IMF strength.

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11.2 Types of IMF

• Electrostatic forces act over larger distances
  in accordance with Coulombs law
• Ion-ion forces strongest found in ionic
  crystals (i.e. lattice energy)

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1. Ion-dipole between an ion and a dipole (a
   neutral, polar molecule/has separated partial
   charges)

• Increase with increasing polarity of molecule and
  increasing ion charge.

Ex Compare IMF in Cl- (aq) and S2- (aq).
lt
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1. Dipole-dipole weakest electrostatic force exist
   between neutral polar molecules

• Increase with increasing polarity (dipole moment)
  of molecule

Ex What IMF exist in NaCl (aq)?
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• Hydrogen bonds (or H-bonds)
• H is unique among the elements because it has a
  single e- that is also a valence e-.
• When this e- is hogged by a highly EN atom (a
  very polar covalent bond), the H nucleus is
  partially exposed and becomes attracted to an
  e--rich atom nearby.

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• H-bonds form with H-XX', where X and X' have
  high EN and X' possesses a lone pair of e-
• X F, O, N (since most EN elements) on two
  molecules

F-H O-H N-H
F O N
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• There is no strict cutoff for the ability to
  form H-bonds (S forms a biologically important
  hydrogen bond in proteins).
• Hold DNA strands together in double-helix

Nucleotide pairs form H-bonds
DNA double helix
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• H-bonds explain why ice is less dense than water.

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Ex Boiling points of nonmetal hydrides

• Conclusions
• Polar molecules have higher BP than nonpolar
  molecules
• ? Polar molecules have stronger IMF
• BP increases with increasing MW
• ? Heavier molecules have stronger IMF

Boiling Points (ºC)

• NH3, H2O, and HF have unusually high BP.
• ? H-bonds are stronger than dipole-dipole IMF

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Inductive forces

• Arise from distortion of the e- cloud induced by
  the electrical field produced by another particle
  or molecule nearby.
• London dispersion between polar or nonpolar
  molecules or atoms
• Proposed by Fritz London in 1930
• Must exist because nonpolar molecules form liquids

Fritz London(1900-1954)
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• How they form
• Motion of e- creates an instantaneous dipole
  moment, making it temporarily polar.

• Instantaneous dipole moment induces a dipole in
  an adjacent atom
• Persist for about 10-14 or 10-15 second
• Ex two He atoms

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Geckos!

• Geckos feet make use of London dispersion forces
  to climb almost anything.
• A gecko can hang on a glass surface using only
  one toe.
• Researchers at Stanford University recently
  developed a gecko-like robot which uses synthetic
  setae to climb walls

http//en.wikipedia.org/wiki/Van_der_Waals27_forc
e
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• London dispersion forces increase with
• Increasing MW, of e-, and of atoms
  (increasing of e- orbitals to be distorted)
• Boiling points
• Effect of MW Effect of atoms
• pentane 36ºC Ne 246C
• hexane 69ºC CH4   162C
• heptane 98ºC
• ??? effect
• H2O 100C
• D2O 101.4C
• Longer shapes (more likely to interact with
  other molecules)
• C5H12 isomers 2,2-dimethylpropane 10C
• pentane 36C

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Summary of IMF
Van der Waals forces
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Ex Identify all IMF present in a pure sample of
each substance, then explain the boiling points.
BP(C) IMF Explanation
HF 20
HCl -85
HBr -67
HI -35
Lowest MW/weakest London, but most polar/strongest dipole-dipole and has H-bonds
Low MW/weak London, moderate polarity/dipole-dipole and no H-bonds
Medium MW/medium London, moderate polarity/dipole-dipole and no H-bonds
Highest MW/strongest London, but least polar bond/weakest dipole-dipole and no H-bonds
London, dipole-dipole, H-bonds
London, dipole-dipole
London, dipole-dipole
London, dipole-dipole
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11.3 Properties resulting from IMF

• Viscosity resistance of a liquid to flow
• Viscosity depends on
• -the attractive forces between molecules
• -the tendency of molecules to become entangled
• -the temperature

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11.3 Properties resulting from IMF

• Surface tension energy required to increase the
  surface area of a liquid

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• 3. Cohesion attraction of molecules for other
  molecules of the same compound
• 4. Adhesion attraction of molecules for a
  surface

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• Meniscus curved upper surface of a liquid in a
  container a relative measure of adhesive and
  cohesive forces
• Ex

Hg
H2O
(cohesion rules)
(adhesion rules)
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Phase Changes

• Surface molecules are only attracted inwards
  towards the bulk molecules.
• Sublimation solid ? gas.
• Vaporization liquid ? gas.
• Melting or fusion solid ? liquid.
• Deposition gas ? solid.
• Condensation gas ? liquid.
• Freezing liquid ? solid.
• Energy Changes Accompanying Phase Changes
• Energy change of the system for the above
  processes are

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Intermolecular Forces Bulk and Surface
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Phase Changes

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

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Phase Changes

• Energy Changes Accompanying Phase Changes
• All phase changes are possible under the right
  conditions (e.g. water sublimes when snow
  disappears without forming puddles).
• The sequence
• heat solid ? melt ? heat liquid ? boil ? heat gas
• is endothermic.
• The sequence
• cool gas ? condense ? cool liquid ? freeze ? cool
  solid
• is exothermic.

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Phase Changes
Energy Changes Accompanying Phase Changes
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Phase Changes

• Heating Curves
• Plot of temperature change versus heat added is a
  heating curve.
• During a phase change, adding heat causes no
  temperature change.
• These points are used to calculate ?Hfus and
  ?Hvap.
• Supercooling When a liquid is cooled below its
  melting point and it still remains a liquid.
• Achieved by keeping the temperature low and
  increasing kinetic energy to break intermolecular
  forces.

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Phase Changes
Heating Curves
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Heating Curve Illustrated
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Phase Changes

• Critical Temperature and Pressure
• Gases liquefied by increasing pressure at some
  temperature.
• Critical temperature the minimum temperature for
  liquefaction of a gas using pressure.
• Critical pressure pressure required for
  liquefaction.

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Critical Temperature, Tc
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Transition to Supercritical CO2
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Supercritical CO2 Used to Decaffeinate Coffee
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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.

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Gas-Liquid Equilibration
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Vapor Pressure

• Explaining Vapor Pressure on the Molecular Level
• Dynamic Equilibrium the point when as many
  molecules escape the surface as strike the
  surface.
• Vapor pressure is the pressure exerted when the
  liquid and vapor are in dynamic equilibrium.

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Vapor Pressure

• Volatility, Vapor Pressure, and Temperature
• If equilibrium is never established then the
  liquid evaporates.
• Volatile substances evaporate rapidly.
• The higher the temperature, the higher the
  average kinetic energy, the faster the liquid
  evaporates.

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Liquid Evaporates when no Equilibrium is
Established
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Vapor Pressure
Volatility, Vapor Pressure, and Temperature
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Vapor Pressure

• 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).

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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.
• 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.

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Phase Diagrams

• Any temperature and pressure combination not on a
  curve represents a single phase.

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Phase Diagrams

• The Phase Diagrams of H2O and CO2
• Water
• The melting point curve slopes to the left
  because ice is less dense than water.
• Triple point occurs at 0.0098?C and 4.58 mmHg.
• Normal melting (freezing) point is 0?C.
• Normal boiling point is 100?C.
• Critical point is 374?C and 218 atm.
• Carbon Dioxide
• Triple point occurs at -56.4?C and 5.11 atm.
• Normal sublimation point is -78.5?C. (At 1 atm
  CO2 sublimes it does not melt.)
• Critical point occurs at 31.1?C and 73 atm.

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Phase Diagrams
The Phase Diagrams of H2O and CO2
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11.4 Phase Changes

• Processes
• Endothermic melting, vaporization, sublimation
• Exothermic condensation, freezing, deposition

I2 (s) and (g)
Microchip
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Water Enthalpy diagram or heating curve
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11.5 Vapor pressure
Pressure cooker 2 atm
Normal BP 1 atm
10,000 elev 0.7 atm
29,029 elev (Mt. Everest) 0.3 atm

• A liquid will boil when the vapor pressure equals
  the atmospheric pressure, at any T above the
  triple point.

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11.6 Phase diagrams CO2

• Lines 2 phases exist in equilibrium
• Triple point all 3 phases exist together in
  equilibrium (X on graph)
• Critical point, or critical temperature
  pressure highest T and P at which a liquid can
  exist (Z on graph)

Temp (ºC)

• For most substances, inc P will cause a gas to
  condense (or deposit), a liquid to freeze, and a
  solid to become more dense (to a limit.)

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Phase diagrams H2O

• For H2O, inc P will cause ice to melt.

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50

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11.7-8 Structures of solids

• Amorphous without orderly structure
• Ex rubber, glass
• Crystalline repeating structure have many
  different stacking patterns based on chemical
  formula, atomic or ionic sizes, and bonding

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Cubic Unit Cells in Crystalline Solids

• Primitive-cubic shared atoms are located only at
  each of the corners. 1 atom per unit cell.
• Body-centered cubic 1 atom in center and the
  corner atoms give a net of 2 atoms per unit cell.
• Face-centered cubic corner atoms plus half-atoms
  in each face give 4 atoms per unit cell.

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Common Lattice Structures
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Types of Crystalline Solids
Type Particles Forces Notable properties Examples
Atomic Atoms London dispersion Poor conductors Very low MP Ar (s),Kr (s)
A small (2 cm long) piece of rapidly melting
argon ice (the liquid is flowing off at the
bottom) which has been frozen by allowing a slow
stream of the gas to flow into a small graduated
cylinder which was immersed into a cup of liquid
nitrogen
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Molecular crystals Molecules (polar or non-polar) London dispersion, dipole-dipole, H-bonds Poor conductors Low to moderate MP SO2(s) CO2 (s), C12H22O11, H2O (s)
Sucrose (liq at room T)
Ice(liq at room T)
Carbon dioxide, dry ice(g at room T)
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Covalent (a.k.a. covalent network) Atoms bonded in a covalent network Covalent bonds Very hard Very high MP Generally insoluble Variable conductivity C (diamond graphite) SiO2 (quartz) Ge, Si, SiC, BN
Diamond
Graphite
SiO2
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Ionic Anions and cations Crystals shatter! Ion-ion (ionic bonding) High Lattice Energy Hard brittle High MP,BP Poor conductors Some solubility in H2O NaCl, Ca(NO3)2
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Metallic Metal cations in a diffuse, delocalized e- cloud Metallic bonds Usually face-centered or body centered Excellent conductors Malleable Ductile High but wide range of MP Cu, Al, Fe (hard) Alloys Pb, Au, Na (soft)
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Overall

• Physical properties depend on these forces. The
  stronger the forces between the particles,
• (a) the higher the melting point.
• (b) the higher the boiling point.
• (c) the lower the vapor pressure (partial
  pressure of vapor in equilibrium with liquid or
  solid in a closed container at a fixed
  temperature).
• (d) the higher the viscosity (resistance to
  flow).
• (e) the greater the surface tension (resistance
  to an increase in surface area).

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Practice

• Determine the type of solid and the forces
  holding the particles together
• SiO2 Covalent Network Covalent Bonds
• NaNO3 Ionic Electrostatic Att.
• C2H6 Molecular Dispersion
• CH3OH Molecular Dispersion, Dipole-Dipole,
  H-Bond
• C(diamond) Covalent Network Covalent Bonds
• Al Metallic Metallic
• Kr Atomic (Molecular) Dispersion
• H2O Molecular Dispersion, Dipole-Dipole, H-Bond

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Extra Material

• The following pages contain some additional
  material and review items

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Examples
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Ionic Solids

• stable, high melting points
• held together by strong electrostatic forces
  between oppositely charged ions
• larger ions are arranged in closest packing
  arrangement
• smaller ions fit in the holes created by the
  larger ions

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Cubic Unit Cells in Crystalline Solids

• Primitive-cubic shared atoms are located only at
  each of the corners. 1 atom per unit cell.
• Body-centered cubic 1 atom in center and the
  corner atoms give a net of 2 atoms per unit cell.
• Face-centered cubic corner atoms plus half-atoms
  in each face give 4 atoms per unit cell.

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Common Lattice Structures
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Calculations involving the Unit Cell

• The density of a metal can be calculated if we
  know the length of the side of a unit cell.
• The radius of an metal atom can be determined if
  the unit cell type and the density of the metal
  known
• Relationship between length of side and radius of
  atom
• Primitive 2r l FCC BCC
• E.g. Polonium crystallizes according to the
  primitive cubic structure. Determine its density
  if the atomic radius is 167 pm.
• E.g.2 Calculate the radius of potassium if its
  density is 0.8560 g/cm3 and it has a BCC crystal
  structure.

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Figure 11.31

• Length of sides a, b, and c as well as angles a,
  b, g vary to give most of the unit cells. Return
  to unit cells

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Unit Cells in Crystalline Solids

• Metal crystals made up of atoms in regular arrays
  the smallest of repeating array of atoms is
  called the unit cell.
• There are 14 different unit cells that are
  observed which vary in terms of the angles
  between atoms some are 90, but others are not.
  Go to Figure 11.31

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Packing of Spheres and the Structures of Metals

• Arrays of atoms act as if they are spheres. Two
  or more layers produce 3-D structure.
• Angles between groups of atoms can be 90 or can
  be in a more compact arrangement such as the
  hexagonal closest pack (see below) where the
  spheres form hexagons.
• Two cubic arrays one directly on top of the other
  produces simple cubic (primitive) structure.
• Each atom has 6 nearest neighbors (coordination
  number of 6) nearest neighbor is where an atom
  touches another atom.
• 54 of the space in a cube is used.
• Offset layers produces a-b-a-b arrangement since
  it takes two layers to define arrangement of
  atoms.
• BCC structure an example.
• Coordination is 8.

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Packing of Spheres and the Structures of Metals

• FCC structure has a-b-c-a-b-c stacking. It takes
  three layers to establish the repeating pattern
  and has 4 atoms per unit cell and the
  coordination number is 12.

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Metallic Crystals

• can be viewed as metals atoms (spheres) packed
  together in the closest arrangement possible
• closest packing- when each sphere has 12
  neighbors
• 6 on the same plane
• 3 in the plane above
• 3 in the plane below

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Bonding of Metals

• the highest energy level for most metal atoms
  does not contain many electrons
• these vacant overlapping orbitals allow outer
  electrons to roam freely around the entire metal

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Bonding of Metals

• these roaming electrons
• form a sea of electrons
• around the metal atoms
• malleability and ductility
• bonding is the same in every direction
• one layer of atoms can slide past another without
  friction
• conductivity
• from the freedom of electrons to move around the
  atoms

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Metal Alloys

• substance that is a mixture of elements and has
  metallic properties
• substitutional alloy
• host metal atoms are replaced by other metal
  atoms
• happens when they have similar sizes
• interstitial alloy
• metal atoms occupy spaces created between host
  metal atoms
• happens when metal atoms have large difference in
  size

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Examples

• Brass
• substitutional
• 1/3 of Cu atoms replaced by Zn
• Steel
• interstitial
• Fe with C atoms in between
• makes harder and less malleable

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Chapter 11 Overview

• Changes of State
• Phase transitions
• Phase Diagrams
• Liquid State
• Properties of Liquids Surface tension and
  viscosity
• Intermolecular forces explaining liquid
  properties
• Solid State
• Classification of Solids by Type of Attraction
  between Units
• Crystalline solids crystal lattices and unit
  cells
• Structures of some crystalline solids
• Calculations Involving Unit-Cell Dimensions
• Determining the Crystal Structure by X-ray
  Diffraction

Exam on Friday We will begin Chp 14 Thursday
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Comparison of Gases, Liquids and Solids

• Gases are compressible fluids. Their molecules
  are widely separated.
• Liquids are relatively incompressible fluids.
  Their molecules are more tightly packed.
• Solids are nearly incompressible and rigid. Their
  molecules or ions are in close contact and do not
  move.

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Phase Transitions

• Melting change of a solid to a liquid.
• Freezing change a liquid to a solid.
• Vaporization change of a solid or liquid to a
  gas. Change of solid to vapor often called
  sublimation.
• Condensation change of a gas to a liquid or
  solid. Change of a gas to a solid often called
  deposition.

H2O(s) ? H2O(l) H2O(l) ? H2O(s) H2O(l) ? H2O(g)
or H2O(s) ? H2O(g) H2O(g) ? H2O(l) or H2O(g) ?
H2O(s)
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Vapor Pressure

• In a sealed container, some of a liquid
  evaporates to establish a pressure in the vapor
  phase.
• Vapor pressure partial pressure of the vapor
  over the liquid measured at equilibrium and at
  some temperature.
• Dynamic equilibrium

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Temperature Dependence of Vapor Pressures

• The vapor pressure above the liquid varies
  exponentially with changes in the temperature.
• The Clausius-Clapeyron equation shows how the
  vapor pressure and temperature are related. It
  can be written as

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Clausius Clapeyron Equation

• A straight line plot results when ln P vs. 1/T is
  plotted and has a slope of ?Hvap/R.
• Clausius Clapeyron equation is true for any two
  pairs of points.
• Write the equation for each and combine to get

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Using the Clausius Clapeyron Equation

• Boiling point the temperature at which the vapor
  pressure of a liquid is equal to the pressure of
  the external atmosphere.
• Normal boiling point the temperature at which the
  vapor pressure of a liquid is equal to
  atmospheric pressure (1 atm).

E.g. Determine normal boiling point of chloroform
if its heat of vaporization is 31.4 kJ/mol and it
has a vapor pressure of 190.0 mmHg at
25.0C. E.g.2. The normal boiling point of
benzene is 80.1C at 26.1C it has a vapor
pressure of 100.0 mmHg. What is the heat of
vaporization?
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Energy of Heat and Phase Change

• Heat of vaporization heat needed for the
  vaporization of a liquid.
• H2O(l) ?H2O(g) DH 40.7 kJ
• Heat of fusion heat needed for the melting of a
  solid.
• H2O(s) ?H2O(l) DH 6.01 kJ
• Temperature does not change during the change
  from one phase to another.

E.g. Start with a solution consisting of 50.0 g
of H2O(s) and 50.0 g of H2O(l) at 0C. Determine
the heat required to heat this mixture to 100.0C
and evaporate half of the water.
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Phase Diagrams

• Graph of pressure-temperature relationship
  describes when 1,2,3 or more phases are present
  and/or in equilibrium with each other.
• Lines indicate equilibrium state two phases.
• Triple point- Temp. and press. where all three
  phases co-exist in equilibrium.
• Critical temp.- Temp. where substance must always
  be gas, no matter what pressure.

• Critical pressure- vapor pressure at critical
  temp.
• Critical point- point where system is at its
  critical pressure and temp.

85
Properties of Liquids

• Surface tension the energy required to increase
  the surface area of a liquid by a unit amount.
• Viscosity a measure of a liquids resistance to
  flow.
• Surface tension The net pull toward the interior
  of the liquid makes the surface tend to as small
  a surface area as possible and a substance does
  not penetrate it easily.
• Viscosity Related to mobility of a molecule
  (proportional to the size and types of
  interactions in the liquid).

• Viscosity decreases as the temperature increases
  since increased temperatures tend to cause
  increased mobility of the molecule.

86
Intermolecular Forces

• Intermolecular forces attractions and repulsions
  between molecules that hold them together.
• Intermolecular forces (van der Waals forces) hold
  molecules together in liquid and solid phases.
• Ion-dipole force interaction between an ion and
  partial charges in a polar molecule.
• Dipole-dipole force attractive force between
  polar molecules with positive end of one molecule
  is aligned with negative side of other.
• London dispersion Forces interactions between
  instantaneously formed electric dipoles on
  neighboring polar or nonpolar molecules.
• Polarizability ease with which electron cloud of
  some substance can be distorted by presence of
  some electric field (such as another dipolar
  substance). Related to size of atom or molecule.
  Small atoms and molecules less easily polarized.

87
Boiling Points vs. Molecular Weight

• Hydrogen bonds the interaction between hydrogen
  bound to an electronegative element (N, O, or F)
  and an electron pair from another electronegative
  element. Hydrogen bonding is the dominate force
  holding the two DNA molecules together to form
  the double helix configuration of DNA.

88
Comparisonof Energies for Intermolecular Forces
Interaction Forces Approximate Energy
Intermolecular
London 1 10 kJ
Dipole-dipole 3 4 kJ
Ion-dipole 5 50 kJ
Hydrogen bonding 10 40 kJ
Chemical bonding
Ionic 100 1000 kJ
Covalent 100 1000 kJ
89
Structure of Solids

• Types of solids
• Crystalline a well defined arrangement of
  atoms this arrangement is often seen on a
  macroscopic level.
• Ionic solids ionic bonds hold the solids in a
  regular three dimensional arrangement.
• Molecular solid solids like ice that are held
  together by intermolecular forces.
• Covalent network a solid consists of atoms held
  together in large networks or chains by covalent
  networks.
• Metallic similar to covalent network except
  with metals. Provides high conductivity.
• Amorphous atoms are randomly arranged. No
  order exists in the solid. Example glass