Intermolecular Forces, Liquids and Solids

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Intermolecular Forces, Liquids and Solids
CHAPTER 11CHEM 160
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A Molecular Comparison of Liquids and Solids

• Physical properties of substances understood in
  terms of kinetic molecular theory
• Gases are highly compressible, assumes shape and
  volume of container
• Gas molecules are far apart and do not interact
  much with each other.
• Liquids are almost incompressible, assume the
  shape but not the volume of container
• Liquids molecules are held closer together than
  gas molecules, but not so rigidly that the
  molecules cannot slide past each other.

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A Molecular Comparison of Liquids and Solids

• Solids are incompressible and have a definite
  shape and volume
• Solid molecules are packed closely together. The
  molecules are so rigidly packed that they cannot
  easily slide past each other.

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A Molecular Comparison of Liquids and Solids
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A Molecular Comparison of Liquids and Solids
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A Molecular Comparison of Liquids and 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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Intermolecular Forces

• 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).

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Intermolecular Forces
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Intermolecular Forces

• Ion-Dipole Forces
• Interaction between an ion and a dipole (e.g.
  water).
• Strongest of all intermolecular forces.

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Intermolecular Forces

• Dipole-Dipole Forces
• Dipole-dipole forces exist between neutral polar
  molecules.
• Polar molecules need to be close together.
• Weaker than ion-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 Forces
Dipole-Dipole Forces
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Intermolecular Forces
Dipole-Dipole Forces
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Intermolecular Forces

• London 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 Forces

• London Dispersion Forces
• One instantaneous dipole can induce another
  instantaneous dipole in an adjacent molecule (or
  atom).
• The forces between instantaneous dipoles are
  called London dispersion forces.

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Intermolecular Forces

• 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.

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Intermolecular Forces

• London Dispersion Forces
• 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 Forces
London Dispersion Forces
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Intermolecular Forces
London Dispersion Forces
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Intermolecular Forces

• 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 Forces

• 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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Hydrogen Bonding
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Hydrogen Bonding
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Intermolecular Forces

• Hydrogen Bonding
• Hydrogen bonds are responsible for
• 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.

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Intermolecular Forces
Hydrogen Bonding
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Intermolecular Forces
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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.
• Surface Tension
• Bulk molecules (those in the liquid) are equally
  attracted to their neighbors.

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Some Properties of Liquids
Viscosity
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Surface Tension
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Some Properties of Liquids

• Surface Tension
• Surface molecules are only attracted inwards
  towards the bulk molecules.
• Therefore, surface molecules are packed more
  closely than bulk molecules.
• Surface tension is the amount of energy required
  to increase the surface area of a liquid.
• Cohesive forces bind molecules to each other.
• Adhesive forces bind molecules to a surface.

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Some Properties of Liquids

• Surface Tension
• 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.
• Capillary Action When a narrow glass tube is
  placed in water, the meniscus pulls the water up
  the tube.

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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
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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).

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Phase Changes

• Energy Changes Accompanying Phase Changes
• 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
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Phase Changes

• Energy Changes Accompanying Phase Changes
• All phase changes are possible under the right
  conditions.
• 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

• 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

• 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 Changes
Critical Temperature and Pressure