Chapter 13: Liquids and Solids

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• Chapter 13 Liquids and Solids

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Compare and Contrast Gases, Liquids and Solids

• Gases are monatomic of made of small molecules
  with covalent bonding.
• Pure liquids are composed of molecules with
  covalent bonding. Mercury is atomic. Atoms are
  held together by metallic bonds.
• All ionic compounds are solids. They conduct
  electricity when heated and melted or dissolved
  in water.
• Except for mercury, all metals are solid. They
  conduct electricity in the solid state.

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Solids

• Many covalently-bonded compounds are solids at
  room temperature
• Some are molecular, such as glucose, C6H12O6
• Others are macromolecular, such as quartz, SiO2.

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Intermolecular Forces

• The physical state of a substance at a given
  temperature and pressure is related to several
  factors, which involve intermolecular forces.
• Particles with greater mass have a greater
  attraction for each other.
• The boiling points and melting points of
  substances increase with greater molecular mass.

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Forces Between Particles

• Interionic Forces. These are some of the
  strongest forces found in ionic compounds.
• They are due to the attraction of opposite
  charges.
• There is a greater attraction between ions of
  greater charge. The stronger the attractive
  force, the higher the melting point.

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Forces Between Particles

• Intramolecular Forces. Forces that hold atoms
  together within molecules. Polar and non-polar
  covalent bonds.
• Intermolecular Forces. The attractive force for
  molecules to each other. These forces are weaker
  than interionic or intramolecular forces. These
  forces help determine the physical properties of
  molecular compounds.
• Also called van der Waals forces

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Intermolecular Forces

• Dipole Forces Polar molecules are attracted to
  each other via the negative end of one molecule
  being attracted to the positive end of another
  molecule.
• Hydrogen Bond Bond between hydrogen attached to
  an electronegative atom (F, O, or N) to another
  electronegative atom.
• Dispersion Forces (London) Temporary dipoles in
  polarizable molecules.

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Intermolecular Dipole Forces
Insert figure 13.2
Liquid
Solid
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Hydrogen Bonding in Ice (H2O)
Insert figure 13.4
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Momentary Dipole for London Dispersion Forces
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Intermolecular Forces

• Listed from strongest to weakest
• Hydrogen bondsgtDipole forcesgtDispersion forces.

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Substances With Stronger Intermolecular Bonds,
Generally Have a

• Lower Vapor Pressure
• Higher Boiling Point
• Higher Melting Point
• Greater Viscosity
• Greater Surface Tension

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The Liquid State

• Viscosity A measure of the resistance of a
  liquid to flow.
• Related to the shape of molecules. Small,
  symmetrical molecules with weak intermolecular
  forces have low viscosities.
• Large molecules have high viscosities due to the
  larger dispersion forces.
• Smaller molecules with extensive hydrogen bonding
  have high viscosities.

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Longer Molecules Have a Greater Viscosity Than
Compact Molecules
Insert figure 13.6
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Properties of Liquids

• Surface Tension The energy required to increase
  the surface area of a liquid by a given amount.
• Liquids with stronger intermolecular forces have
  a greater surface tension.
• A molecule in the middle of a liquid is attracted
  equally in all directions. Molecules on the
  surface at its ides and below only. There is no
  upward force. These unequal forces exert a force
  inward at the surface, causing the surface to
  contract. Surface area tends to minimized.
  Increasing surface area is resisted.

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Surface Tension

• Liquids Tend to bead and become spherical due
  to the attempt to minimize surface area due to
  surface tension.
• Due to surface tension, a needle can float on the
  surface of water.
• Detergents or surfactants can lower surface
  tension, causing water to wet a surface.

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Forces Giving Rise to Surface Tension
Insert figure 13.8
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Temperature Effects

• Higher temperatures increase molecular speed and
  favor the vapor phase over the liquid, and the
  liquid over the solid.
• At higher temperature, vapor pressure increases,
  and viscosity decreases.

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Vaporization and Condensation

• Vaporization The process by which molecules of a
  liquid break away and enter the gas phase.
  (Evaporation)
• Condensation The process by which molecules in
  the gas phase enter the liquid phase. (the
  opposite of evaporation)
• There is an equilibrium between the two phases
• Liquid Vapor

Vaporization
Condensation
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The Effect of Temperature on Vapor Pressure
Insert figure 13.10
Higher Temperature
Lower Temperature
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Boiling Point

• As a liquid is heated, its vapor pressure
  increases.
• The temperature at which the vapor pressure of
  the liquid equals the external pressure is called
  the boiling point.
• Boiling point varies with external pressure.
• The temperature at which the vapor pressure of a
  liquid equals 1 atm is called the normal boiling
  point.

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Distillation

• A volatile component can be separated from a
  non-volatile component by distillation. The
  volatile component is evaporated and the vapors
  are condensed, leaving the non-volatile component
  behind.
• In simple distillation, a solution of two
  volatile liquid compounds are heated to boiling.
  The more volatile component (lower boiling point,
  higher vapor pressure) is in higher concentration
  in the vapor than in the liquid.

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Distillation

• The resulting vapor is cooled and condensed back
  into a liquid.
• The more volatile component is in greater
  concentration in the distillate. The residue has
  a greater concentration of the less volatile
  component.
• The more volatile component can be obtained by
  collecting the first portion of the distillate
  (the material that has been distilled).

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Distillation Apparatus
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Heat of Vaporization

• Heat is required to vaporize a liquid.
• A liquid that evaporates at room temperature
  absorbs heat from the surroundings.
• The quantity of heat required to vaporize 1
  mol of liquid is called the molar heat of
  vaporization.

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Heat of Vaporization Problem

• Calculate the amount of heat required to
  evaporated 36.0 g of water. (The molar heat of
  vaporization, DHv, of water is 40.7 kJ/mol.)
• mol g/MM 36.0/18.0 2.0 mol
• 40.7 kJ/mol x 2.0 mol

81.4 kJ
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The Solid State

• Many solids are amorphous. The particles (ions or
  molecules) have no definite order.
• Examples are glass, rubber, and plastic.
• Crystalline solids have a regular arrangement of
  particles, called a crystal lattice.
• Some types of crystal lattice are
• simple cubic
• body-centered cubic
• face-centered cubic

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Common Crystal Structures
Insert figure 13.13
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Solids may be classified by the type of
inter-particle force.

• Ionic solids Ions are at each lattice point in
  the crystal.
• Molecular solids Have discrete molecules at each
  lattice point. Held together by intermolecular
  forces.
• Covalent network or macromolecular solids have
  atoms at the lattice points. They are held
  together by covalent bonds. Examples are SiC,
  graphite and diamond.

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Allotropes

• Carbon and other elements can exist in various
  forms. These forms are called allotropes.
• Allotropes of carbon are graphite, amorphous, and
  diamond, and nanotubes. These are
  macromolecular.
• A molecular allotrope is Buckminsterfulleranes.

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The Structure of Diamond, an Allotrope of Carbon
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Graphite, an Allotrope of Carbon
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Buckminsterfullerene, an Allotrope of Carbon
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Metallic Solids

• Metallic solids have positive ions at lattice
  points. Valence electrons are distributed
  throughout the lattice, almost like a fluid.
• Most metals are malleable, they can be hammered
  or rolled into sheets.
• Most metals are ductile, they can be drawn or
  pulled into wires.

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Melting and Freezing

• When a crystalline solid is heated, particles
  vibrate more vigorously and become liquid.
• This process is called melting.
• When a substance changes from a liquid to a solid
  the process is called freezing.
• Liquids and solids are in dynamic equilibrium
• Solid Liquid

melting
freezing
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Melting and Freezing

• The temperature at which solid and liquid are in
  equilibrium is called the melting point or the
  freezing point.
• The melting point and the freezing point for a
  substance are at the same temperature.

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Heat of Fusion

• The amount of heat required to convert 1
  mol of a solid to a liquid is called the molar
  heat of fusion.
• The intermolecular forces of attraction are not
  as great in a liquid as a solid.
• The difference between the two is the heat of
  fusion.

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Heating and Cooling Curves

• If heat is applied to a solid at a constant rate,
  a heating curve can be constructed by plotting
  temperature vs. time of heating.
• Temperature rises steadily until the solid melts.
• The temperature remains constant until the
  material is completely melted.
• Then the temperature rises again until the
  material starts to vaporize.

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Heating and Cooling Curves(continued)

• Next the temperature remains the same until the
  material completely vaporizes.
• Finally, the temperature rises as the gas is
  heated.
• When a gas is heated in a closed container,
  superheated steam is produced.

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Heating and Cooling Curve

• For portions of the heating curve where the
  temperature is rising, there is an increase in
  the kinetic energy of the particles increases.
• The amount of energy added equals the mass of the
  substance times its specific heat times its
  change in temperature.
• When the substance is melting or boiling the
  temperature remains constant. Intermolecular
  bonds are being weakened. The energy added is
  the heat of fusion or the heat of vaporization.

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Heating/Cooling Curve for Water
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Water

• Compare and contrast the structure, density, and
  physical properties of liquid and solid water.
• Water is a unique substance. Intermolecular
  bonding consists primarily of hydrogen bonds
  which are quite strong.
• Relatively high boiling point (100oC) and melting
  point (0oC)
• Moderate density (1 g/mL)
• High Specific Heat (4.184 J/g)
• High Heat of Vaporization

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Properties of Water

• High surface tension
• High capillarity (Hydrogen bonds to walls of
  glass). Adhesiongtcohesion.
• Density of liquidgtsolid
• ice floats
• Due to hydrogen bonding, open structure of ice.

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Solvent Properties of Water

• Highly polar, able to form hydrogen bonds.
• Good solvent for ionic compounds.
• Ion-dipole interaction

Insert figure 13.20