Distinguish between intermolecular and intramolecular attractions

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• Sublimation
• Deposition
• Condensation
• Evaporation
• Melting
• Freezing
• Freezing point
• Boiling point

Define

• Polar
• Nonpolar
• Dipole-dipole forces
• Ion-dipole forces
• Hydrogen bonding
• London dispersion forces

• Vapor pressure
• Viscosity
• Surface tension
• ?H of fusion
• ?H of vaporization
• ?H of sublimation

• Distinguish between intermolecular and
  intramolecular attractions
• Put a list of compounds in order of increasing
  melting point, boiling point, and vapor pressure
• Use and label the parts of a phase diagram
• Use the Clausius-Clapeyron equation to relate
  temperature to vapor pressure of a substance

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Solid, Liquid, or Gas

• What are three factors determine whether a
  substance is a solid, a liquid, or a gas
• The attractive intermolecular forces between
  particles that tend to draw the particles
  together.
• Temperature The kinetic energies of the
  particles (atoms, molecules, or ions) that make
  up a substance. Kinetic energy tends to keep the
  particles moving apart.
• Pressure pressure is increased or decreased as
  the volume of a closed container changes

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Types of Attractive Forces

• There are several types of attractive
  intermolecular forces
• Ionic
• Ion-dipole forces
• Dipole-dipole forces
• Hydrogen bonding
• Induced-dipole forces
• London dispersion forces

All of the intermolecular forces that hold a
liquid together are called cohesive forces.
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Ion-Dipole Forces

• An ion-dipole force is an attractive force that
  results from the electrostatic attraction between
  an ion and a neutral molecule that has a dipole.
• Most commonly found in solutions. Especially
  important for solutions of ionic compounds in
  polar liquids.
• Ion-dipole attractions become stronger as either
  the charge on the ion increases, or as the
  magnitude of the dipole of the polar molecule
  increases.

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Hydrogen Bonding
Bonding between hydrogen and more electronegative
neighboring atoms such as oxygen and nitrogen
Hydrogen bonding between ammonia and water
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Hydrogen Bonding in DNA
Thymine hydrogen bonds to Adenine
T
A
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Hydrogen Bonding in DNA
Cytosine hydrogen bonds to Guanine
C
G
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Dipole-Dipole Forces

• Dipole-dipole forces are attractive forces
  between the positive end of one polar molecule
  and the negative end of another polar molecule.
• They are much weaker than ionic or covalent bonds
  and have a significant effect only when the
  molecules involved are close together (touching
  or almost touching).

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Induced-Dipole Forces
Induced dipole forces result when an ion or a
dipole induces a dipole in an atom or a molecule
with no dipole. These are weak forces.
Ion-Induced Dipole Forces
An ion-induced dipole attraction is a weak
attraction that results when the approach of an
ion induces a dipole in an atom or in a nonpolar
molecule by disturbing the arrangement of
electrons in the nonpolar species.
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Dipole-Induced Dipole Forces
A dipole-induced dipole attraction is a weak
attraction that results when a polar molecule
induces a dipole in an atom or in a nonpolar
molecule by disturbing the arrangement of
electrons in the nonpolar species.
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London Dispersion Forces
The temporary separations of charge that lead to
the London force attractions are what attract one
nonpolar molecule to its neighbors.
London forces increase with the size of the
molecules.
Fritz London 1900-1954
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London Dispersion Forces
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London Forces in Hydrocarbons
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Boiling point as a measure of intermolecular
attractive forces
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Relative Magnitudes of Forces
The types of bonding forces vary in their
strength as measured by average bond energy.
Ionic bonds
Ion-dipole interactions
Strongest Weakest
Hydrogen bonding (12-16 kcal/mol )
Dipole-dipole interactions (2-0.5 kcal/mol)
Ion induced dipole interactions
Induced Dipole-dipole interactions
London forces (less than 1 kcal/mol)
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What Is a Liquid? No, really, what IS a
liquid??!!
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What Is a Liquid?

• A liquid is a state of matter in which a sample
  of matter
• is made up of very small particles (atoms,
  molecules, and/or ions).
• flows and can change its shape.
• is not easily compressible and maintains a
  relatively fixed volume.
• The particles that make up a liquid
• are close together with no regular arrangement,
• vibrate, move about, and slide past each other.

This bottle contains both liquid bromine Br2(l), the darker phase at the bottom of the bottle and gaseous bromine Br2(g), the lighter phase above the liquid. The circles show microscopic views of both liquid bromine and gaseous bromine.
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More Properties of a Liquid

• Surface Tension The resistance to an increase
  in its surface area (polar molecules, liquid
  metals).

• Capillary Action Spontaneous rising of a liquid
  in a narrow tube.

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Even More Properties of a Liquid

• Viscosity Resistance to flow

• High viscosity is an
• indication of strong
• intermolecular forces

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Evaporation
Evaporation is the change of a liquid to a gas.

Microscopic view of a liquid. Microscopic view after evaporation.

• When a liquid is heated sufficiently or when the
  pressure on the liquid is decreased sufficiently,
  the forces of attraction between molecules do not
  prevent them from moving apart, and the liquid
  evaporates to a gas.
• Example The sweat on the outside of a cold glass
  evaporates when the glass warms.
• Example Gaseous carbon dioxide is produced when
  the valve on a CO2 fire extinguisher is opened
  and the pressure is reduced.

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Condensation
Condensation is the change from a vapor to a
condensed state (solid or liquid).

• When a gas is cooled sufficiently or, in many
  cases, when the pressure on the gas is increased
  sufficiently, the forces of attraction between
  molecules prevent them from moving apart, and the
  gas condenses to either a liquid or a solid.
• Example Water vapor condenses and forms liquid
  water (sweat) on the outside of a cold glass or
  can.
• Example Liquid carbon dioxide forms at the high
  pressure inside a CO2 fire extinguisher.

Microscopic view of a gas. Microscopic view after condensation.
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Vapor Pressure
The vapor pressure of a liquid is the equilibrium
pressure of a vapor above its liquid (or
solid) The vapor pressure of a liquid varies with
its temperature, as the following graph shows for
water. The line on the graph shows the boiling
temperature for water.
As the temperature of a liquid or solid increases
its vapor pressure also increases. Conversely,
vapor pressure decreases as the temperature
decreases.
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Vapor Pressure Revealed

• When a solid or a liquid evaporates to a gas in a
  closed container, the molecules cannot escape.
• Some of the gas molecules will eventually strike
  the condensed phase and condense back into it.
• When the rate of condensation of the gas becomes
  equal to the rate of evaporation of the liquid or
  solid, the amount of gas, liquid and/or solid no
  longer changes.
• The gas in the container is in equilibrium with
  the liquid or solid.                         

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Factors That Affect Vapor Pressure

• Types of Molecules the types of molecules that
  make up a solid or liquid determine its vapor
  pressure. If the intermolecular forces between
  molecules are

• relatively strong, the vapor pressure will be
  relatively low.
• relatively weak, the vapor pressure will be
  relatively high.

substance vapor pressure at 25oC
diethyl ether 0.7 atm
bromine 0.3 atm
ethyl alcohol 0.08 atm
water 0.03 atm

• Surface Area the surface area of the solid or
  liquid in contact with the gas has no effect on
  the vapor pressure.

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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?
334 K
33.0 kJ/mol
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Solids
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Types of Solids

• Amorphous solids considerable disorder in their
  structures (glass).

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Types of Solids

• Crystalline Solids highly regular arrangement
  of their components

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Metal Alloys

• Substitutional Alloy some metal atoms replaced
  by others of similar size.
• brass Cu/Zn

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Metal Alloys(continued)

• Interstitial Alloy Interstices (holes) in
  closest packed metal structure are occupied by
  small atoms.
• steel iron carbon

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Network Atomic Solids
Some covalently bonded substances DO NOT form
discrete molecules.
Graphite, a network of covalently bonded carbon
atoms
Diamond, a network of covalently bonded carbon
atoms
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Molecular Solids
Strong covalent forces within molecules
Weak covalent forces between molecules
Sulfur, S8
Phosphorus, P4
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Phase Transitions
H2O(s) ? H2O(l) H2O(l) ? H2O(s) H2O(l) ? H2O(g)
H2O(g) ? H2O(l) H2O(s) ? H2O(g) H2O(g) ? H2O(s)

• Melting change of a solid to a liquid.
• Freezing change a liquid to a solid.
• Vaporization change of a liquid to a gas.
• Condensation change of a gas to a liquid.
• Sublimation Change of solid to gas
• Deposition Change of a gas to a solid.

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Water phase changes
Temperature remains constant during a phase
change.
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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 6.02 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.
130 kJ
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Phase Diagrams

• 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.

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Phase changes by Name
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Water
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Carbon dioxide
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Carbon
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Sulfur