Liquids and Solids

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CHAPTER 9

• Liquids and Solids

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Description of Liquids Solids

• Solids liquids are condensed states
• atoms, ions, molecules are close to one another
• highly incompressible
• Solid molecules are packed closely together. The
  molecules are so rigidly packed that they cannot
  easily slide past each other.
• Liquids gases are fluids
• easily flow
• Liquids molecules are held closer together than
  gas molecules, but not so rigidly that the
  molecules cannot slide past each other.
• Intermolecular attractions in liquids solids
  are strong

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Description of Liquids Solids
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Description of Liquids Solids

• 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.

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Kinetic-Molecular Description of Liquids Solids

• strengths of interactions among particles
• degree of ordering of particles
• Gaseslt Liquids lt Solids

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Intermolecular Attractions

• The covalent bond holding a molecule together is
  an intramolecular forces.
• 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).
• When a substance condenses intermolecular forces
  are formed.

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Intermolecular Attractions
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Intermolecular Attractions

• Dipole-Dipole Forces
• Dipole-dipole forces exist between neutral polar
  molecules.
• Polar molecules need to be close together.
• Weaker than ion-dipole forces
• Q1 and Q2 are partial charges.

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Intermolecular Attractions

• Dipole-Dipole Forces
• There is a mix of attractive and repulsive
  dipole-dipole forces as the molecules tumble.
• If two molecules have about the same mass and
  size, then dipole-dipole forces increase with
  increasing polarity.

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Intermolecular Attractions

• Dipole-dipole interactions
• consider NH3 a very polar molecule

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Intermolecular Attractions

• Dispersion Forces
• Weakest of all intermolecular forces.
• It is possible for two adjacent neutral molecules
  to affect each other.
• The nucleus of one molecule (or atom) attracts
  the electrons of the adjacent molecule (or atom).
• For an instant, the electron clouds become
  distorted.
• In that instant a dipole is formed (called an
  instantaneous dipole).

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Intermolecular Attractions
Dispersion Forces
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Intermolecular Attractions

• 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.

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Intermolecular Attractions
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Intermolecular Attractions

• Dispersion Forces
• London dispersion forces depend on the shape of
  the molecule.
• The greater the surface area available for
  contact, the greater the dispersion forces.
• London dispersion forces between spherical
  molecules are lower than between sausage-like
  molecules.

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Intermolecular Attractions

• Hydrogen bonding
• consider H2O

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Intermolecular Attractions

• Hydrogen Bonding
• Special case of dipole-dipole forces.
• By experiments boiling points of compounds with
  H-F, H-O, and H-N bonds are abnormally high.
• Intermolecular forces are abnormally strong.

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Intermolecular Attractions

• Hydrogen Bonding
• H-bonding requires H bonded to an electronegative
  element (most important for compounds of F, O,
  and N).
• Electrons in the H-X (X electronegative
  element) lie much closer to X than H.
• H has only one electron, so in the H-X bond, the
  ? H presents an almost bare proton to the ?- X.
• Therefore, H-bonds are strong.

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Intermolecular Attractions
Hydrogen Bonding
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Intermolecular Attractions

• Hydrogen Bonding
• Ice Floating
• 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.
• 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.
• Ice floats, so it forms an insulating layer on
  top of lakes, rivers, etc. Therefore, aquatic
  life can survive in winter.

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Intermolecular Attractions

• Hydrogen Bonding
• Hydrogen bonds are responsible for
• Protein Structure
• Protein folding is a consequence of H-bonding.
• DNA Transport of Genetic Information

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Comparing Intermolecular Attractions
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Intermolecular Attractions

• Coulombs law the attraction energy determine
• melting boiling points of ionic compounds
• the solubility of ionic compounds
• Arrange the following ionic compounds in the
  expected order of increasing melting and boiling
  points.
• NaF, CaO, CaF2

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Intermolecular Attractions and Phase Changes
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Intermolecular Attractions and Phase Changes
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Evaporation

• Process in which molecules escape from the
  surface of a liquid
• T dependent

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Evaporation
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Vapor Pressure

• pressure exerted by a liquids vapor on its
  surface at equilibrium
• Vap. Press. (torr) for 3 Liquids Norm.
  B.P.
• 0oC 20oC 30oC
• diethyl ether 185 442 647 36oC
• ethanol 12 44 74 78oC
• water 5 18 32 100oC

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Vapor Pressure

• 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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Vapor Pressure

• 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

• 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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Vapor Pressure

• 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.

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Boiling Points

• Boiling point is temperature at which the
  liquids vapor pressure is equal to applied
  pressure
• normal boiling point is boiling point _at_ 1 atm

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Distillation

• Process in which a mixture or solution is
  separated into its components on the basis of the
  differences in boiling points of the components
• Distillation is another vapor pressure
  phenomenon.

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The Liquid State

• energy associated with changes of state
• heat of vaporization
• amount of heat required to change 1 g of a
  liquid substance to a gas at constant T
• units of J/g
• heat of condensation
• reverse of heat of vaporization

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The Liquid State

• molar heat of vaporization or DHvap
• amount of heat required to change 1 mol of a
  liquid to a gas at constant T
• units of J/mol
• molar heat of condensation
• reverse of molar heat of vaporization

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The Liquid State
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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.

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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

• 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
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Phase Changes

• 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 and Heating Curves
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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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Phase Diagrams (P vs T)

• convenient way to display all of the different
  phases of a substance
• phase
• diagram for
• water

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Phase Diagrams (P vs T)

• phase diagram for carbon dioxide

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Synthesis Question

• Maxwell House Coffee Company decaffeinates its
  coffee beans using an extractor that is 7.0 feet
  in diameter and 70.0 feet long. Supercritical
  carbon dioxide at a pressure of 300.0 atm and
  temperature of 100.00C is passed through the
  stainless steel extractor. The extraction vessel
  contains 100,000 pounds of coffee beans soaked in
  water until they have a water content of 50.

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Synthesis Question

• This process removes 90 of the caffeine in a
  single pass of the beans through the extractor.
  Carbon dioxide that has passed over the coffee is
  then directed into a water column that washes the
  caffeine from the supercritical CO2. How many
  moles of carbon dioxide are present in the
  extractor?

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Synthesis Question
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Synthesis Question
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Group Question

• How many CO2 molecules are there in 1.0 cm3 of
  the Maxwell House Coffee Company extractor? How
  many more CO2 molecules are there in a cm3 of the
  supercritical fluid in the Maxwell House
  extractor than in a mole of CO2 at STP?