Intermolecular Forces, Liquids, and Solids

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Intermolecular Forces, Liquids, and Solids

• Prof. G. Matthews

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A Molecular Comparison of Liquids and Solids

• The difference is the state of a molecular
  compound in the solid, liquid or gas state is due
  to
• Kinetic energy.
• Attractive forces between the molecules.

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Intermolecular Forces

• Ion-Dipole Forces
• Ion-dipole force exists between an ion and the
  partial charge on the end of a polar molecule.

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Intermolecular Forces

• Dipole-Dipole Forces
• Neutral polar molecules attract each other when
  the positive end of of molecule is near the
  negative end of another molecule.
• Effective only when the molecules are very close
  together.

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11.2 Intermolecular Forces

• London Dispersion Forces
• The motion of electrons in an atom or molecule
  can create an instantaneous dipole moment.

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Intermolecular Forces

• Polarizability
• The ease at which a charge distribution in a
  molecule can be distorted by an external electric
  field.

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Intermolecular Forces

• The shape of the molecule influences the
  magnitude of the dispersion forces due to the
  overall surface area and contact ability.

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Intermolecular Forces

• Hydrogen bonding
• Occurs between a hydrogen atom in a polar bond
  and an unshared electron on a small
  electronegative atom (N,O,F).

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Some Properties of Liquids

• Viscosity
• Resistance of a liquid to flow.
• Measured in Poise 1g/cm-s.
• Depends on the attractive forces between the
  molecules.
• Viscosity
• Increases with increasing molecular weight.
• Decreases with increases temperature.

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Some Properties of Liquids

• Surface Tension
• In the center of the liquid, the molecules are
  attracted equally in all directions.
• At the surface of the liquid, the molecules
  experience a net inward force.

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Some Properties of Liquids

• This imbalance of intermolecular forces
• Pulls the molecules inwards, reducing the surface
  area.
• Causes the molecules to pack closer together
  creating a skin.
• Surface tension
• A measure of the inward forces that must be
  overcome in order to expand its surface area.

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Some Properties of Liquids

• Cohesive forces
• Intermolecular forces that bind similar molecules
  to one another.
• Adhesive forces
• Intermolecular forces that bind a substance to a
  surface.
• Explains meniscus on the surface of water and
  mercury and capillary action.

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Phase Changes
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Phase Changes

• When a state of matter changes to another state
  it undergoes a phase change.
• The molecular composition of the element or
  compound does not change during a phase change.

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Phase Changes

• Heat of Fusion ?Hfus
• Energy needed to change the state from a solid to
  a liquid.
• Heat of Vaporization ?Hvap
• Energy needed to change the state from a liquid
  to a gas.

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Phase Changes

• Blue lines show the heating of one phase from a
  lower temperature to a higher one.
• Red lines show the conversion of one phase to
  another at constant temperature.
• Overcoming attractive forces.

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Phase Changes

• Critical temperature
• The highest temperature at which a distinct
  liquid phase can form.
• Critical pressure
• The pressure required to bring about liquefaction
  at the critical temperature.
• Depend on attractive forces.

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Vapor Pressure

• The pressure exerted by the vapor above the
  liquid will begin to increase.
• After a short period of time, the pressure of the
  vapor will attain a constant value known as the
  vapor pressure.

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Vapor Pressure

• A liquid boils when its vapor pressure equals the
  external pressure acting on the surface of the
  liquid.
• Bubbles of vapor form in the interior of the
  liquid.

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Phase Diagrams
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Phase Diagrams

• Curves on the phase diagram
• Beyond critical point B (critical temperature and
  pressure) liquid and gas phases are
  indistinguishable.
• Line AC is the variation in the vapor pressure of
  the solid as it sublimes.
• Line AD represents the change in melting point of
  the solid with increasing pressure.

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Phase Diagrams

• Line AD usually slopes to the right as the
  pressure increases, because the solid for most
  substances is more dense than the liquid.
• Triple point
• All three phases are in equilibrium at this
  temperature and pressure.

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Structure of Solids

• Unit cell
• Repeating unit of crystalline solid.
• Crystal lattice
• Array of points representing a crystalline solid.

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Structure of Solids

• Primitive Cubic Lattice points at the corners
  only.
• Body Centered Cubic Lattice point also occurs at
  the center of the cell.
• Face Centered Cubic Lattice points at the center
  of each face as well as each corner.

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Structure of Solids

• An atom in the center of a cell is totally within
  the cell, however adjacent cells share atoms on
  the corners and faces.

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Structure of Solids

• Close Packing of Spheres.
• Hexagonal close-packed structure
• The order of layers is ABAB.
• Cubic close-packed structure
• The order of layers is ABCA

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Structure of Solids

• Cubic close packing and hexagonal close packing
• Coordination number of 12 (12 nearest neighbors)
• Spheres occupy 74 of the total volume
• Simple cubic (Primitive) structure
• Coordination number of 6
• 52 of the space is occupied.
• Body centered cubic structure
• Coordination number of 8
• 68 of the space is occupied.

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Bonding in Solids

• Physical properties of crystalline solids, such
  as melting point and hardness depend on
• Arrangement of the particles.
• Attractive forces between them.

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Bonding in Solids

• Molecular Solids
• Held together by intermolecular forces.
• Soft, low melting points. (Ar, H2O, CO2)
• Covalent-Network Solids
• Consist of atoms held together in large networks
  or chains by covalent bonds.
• Solids have much higher melting points than
  molecular solids. (Diamond, quartz)

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Bonding in Solids

• Bonding in graphite is similar to that of benzene
  with delocalized p bonds extending over the
  layers.
• Layers are held together by weak dispersion
  forces.

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Bonding in Solids

• Ionic Solids
• Solids held together by ionic bonds.
• Bond strength increases with increasing charge on
  the ions.

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Bonding in Solids

• Metallic Solids
• Consist entirely of metal atoms.
• Have a hexagonal close-packed, cubic close-packed
  or body centered cubic structures.
• Bonding is due to delocalized valence electrons.
• Strength of the bonding increases as the number
  of electrons available for bonding increases.

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Bonding in Solids

• A cross section of a metal. Each sphere
  represents the nucleus and inner-core electrons
  of a metal atom. The surrounding colored "fog"
  represents the mobile sea of electrons that binds
  the atoms together.