Chemistry Homework Help

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Vapor Pressure

• Explaining Vapor Pressure on the Molecular
  Level
• 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

• Explaining Vapor Pressure on the Molecular
  Level

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Vapor Pressure

• Explaining Vapor Pressure on the Molecular
  Level
• 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.
• Volatility, Vapor Pressure, and Temperature
• If equilibrium is never established then the
  liquid evaporates.
• Volatile substances evaporate rapidly.

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Vapor Pressure

• Volatility, Vapor Pressure, and Temperature
• The higher the temperature, the higher the
  average kinetic energy, the faster the liquid
  evaporates.

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Vapor Pressure

• Volatility, Vapor Pressure, and Temperature

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Vapor Pressure

• Vapor Pressure and Boiling Point
• Liquids boil when the external pressure equals
  the vapor pressure.
• Temperature of boiling point increases as
  pressure increases.

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Vapor Pressure

• Vapor Pressure and Boiling Point
• 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.
• Normal boiling point is the boiling point at 760
  mmHg (1 atm).

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Phase Diagrams

• Phase diagram plot of pressure vs. Temperature
  summarizing all equilibria between phases.
• Given a temperature and pressure, phase diagrams
  tell us which phase will exist.
• Any temperature and pressure combination not on a
  curve represents a single phase.

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Phase Diagrams

• Features of a phase diagram
• Triple point temperature and pressure at which
  all three phases are in equilibrium.
• Vapor-pressure curve generally as pressure
  increases, temperature increases.
• Critical point critical temperature and pressure
  for the gas.
• Melting point curve as pressure increases, the
  solid phase is favored if the solid is more dense
  than the liquid.
• Normal melting point melting point at 1 atm.

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Phase Diagrams
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Phase Diagrams
The Phase Diagrams of H2O and CO2
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Phase Diagrams

• The Phase Diagrams of H2O and CO2
• Water
• The melting point curve slopes to the left
  because ice is less dense than water.
• Triple point occurs at 0.0098?C and 4.58 mmHg.
• Normal melting (freezing) point is 0?C.
• Normal boiling point is 100?C.
• Critical point is 374?C and 218 atm.

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Phase Diagrams

• The Phase Diagrams of H2O and CO2
• Carbon Dioxide
• Triple point occurs at -56.4?C and 5.11 atm.
• Normal sublimation point is -78.5?C. (At 1 atm
  CO2 sublimes it does not melt.)
• Critical point occurs at 31.1?C and 73 atm.

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Structures of Solids

• Unit Cells
• Crystalline solid well-ordered, definite
  arrangements of molecules, atoms or ions.
• Crystals have an ordered, repeated structure.
• The smallest repeating unit in a crystal is a
  unit cell.
• Unit cell is the smallest unit with all the
  symmetry of the entire crystal.
• Three-dimensional stacking of unit cells is the
  crystal lattice.

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Structures of Solids

• Unit Cells

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Structures of Solids

• Unit Cells
• Three common types of unit cell.
• Primitive cubic, atoms at the corners of a simple
  cube,
• each atom shared by 8 unit cells
• Body-centered cubic (bcc), atoms at the corners
  of a cube plus one in the center of the body of
  the cube,
• corner atoms shared by 8 unit cells, center atom
  completely enclosed in one unit cell
• Face-centered cubic (fcc), atoms at the corners
  of a cube plus one atom in the center of each
  face of the cube,
• corner atoms shared by 8 unit cells, face atoms
  shared by 2 unit cells.

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Unit Cells
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Unit Cells
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Structures of Solids
Unit Cells
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Structures of Solids

• The Crystal Structure of Sodium Chloride
• Two equivalent ways of defining unit cell
• Cl- (larger) ions at the corners of the cell, or
• Na (smaller) ions at the corners of the cell.
• The cation to anion ratio in a unit cell is the
  same for the crystal. In NaCl each unit cell
  contains same number of Na and Cl- ions.
• Note the unit cell for CaCl2 needs twice as many
  Cl- ions as Ca2 ions.

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Structures of Solids
The Crystal Structure of Sodium Chloride
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Structures of Solids
The Crystal Structure of Sodium Chloride
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Structures of Solids

• Close Packing of Spheres
• Solids have maximum intermolecular forces.
• Molecules can be modeled by spheres.
• Atoms and ions are spheres.
• Molecular crystals are formed by close packing of
  the molecules.
• We rationalize maximum intermolecular force in a
  crystal by the close packing of spheres.

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Structures of Solids

• Close Packing of Spheres
• When spheres are packed as closely as possible,
  there are small spaces between adjacent spheres.
• The spaces are called interstitial holes.
• A crystal is built up by placing close packed
  layers of spheres on top of each other.
• There is only one place for the second layer of
  spheres.

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Structures of Solids

• Close Packing of Spheres
• There are two choices for the third layer of
  spheres
• Third layer eclipses the first (ABAB
  arrangement). This is called hexagonal close
  packing (hcp)
• Third layer is in a different position relative
  to the first (ABCABC arrangement). This is
  called cubic close packing (ccp).

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Structures of Solids
Close Packing of Spheres
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Structures of Solids

• Close Packing of Spheres
• Each sphere is surrounded by 12 other spheres (6
  in one plane, 3 above and 3 below).
• Coordination number the number of spheres
  directly surrounding a central sphere.
• Hexagonal and cubic close packing are different
  from the cubic unit cells.
• If unequally sized spheres are used, the smaller
  spheres are placed in the interstitial holes.

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Bonding in Solids

• There are four types of solid
• Molecular (formed from molecules) - usually soft
  with low melting points and poor conductivity.
• Covalent network (formed from atoms) - very hard
  with very high melting points and poor
  conductivity.
• Ions (formed form ions) - hard, brittle, high
  melting points and poor conductivity.
• Metallic (formed from metal atoms) - soft or
  hard, high melting points, good conductivity,
  malleable and ductile.

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Bonding in Solids
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Bonding in Solids

• Molecular Solids
• Intermolecular forces dipole-dipole, London
  dispersion and H-bonds.
• Weak intermolecular forces give rise to low
  melting points.
• Room temperature gases and liquids usually form
  molecular solids and low temperature.
• Efficient packing of molecules is important
  (since they are not regular spheres).

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Bonding in Solids

• Covalent-Network Solids
• Intermolecular forces dipole-dipole, London
  dispersion and H-bonds.
• Atoms held together in large networks.
• Examples diamond, graphite, quartz (SiO2),
  silicon carbide (SiC), and boron nitride (BN).
• In diamond
• each C atom has a coordination number of 4 each
  C atom is tetrahedral there is a
  three-dimensional array of atoms.
• Diamond is hard, and has a high melting point
  (3550 ?C).

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Bonding in Solids
Covalent-Network Solids
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Bonding in Solids

• Covalent-Network Solids
• In graphite
• each C atom is arranged in a planar hexagonal
  ring
• layers of interconnected rings are placed on top
  of each other
• the distance between C atoms is close to benzene
  (1.42 Å vs. 1.395 Å in benzene)
• the distance between layers is large (3.41 Å)
• electrons move in delocalized orbitals (good
  conductor).

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Bonding in Solids

• Ionic Solids
• Ions (spherical) held together by electrostatic
  forces of attraction.
• There are some simple classifications for ionic
  lattice types.

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Ionic Solids
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Bonding in Solids

• Ionic Solids
• NaCl Structure
• Each ion has a coordination number of 6.
• Face-centered cubic lattice.
• Cation to anion ratio is 11.
• Examples LiF, KCl, AgCl and CaO.
• CsCl Structure
• Cs has a coordination number of 8.
• Different from the NaCl structure (Cs is larger
  than Na).
• Cation to anion ratio is 11.

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Bonding in Solids

• Ionic Solids
• Zinc Blende Structure
• Typical example ZnS.
• S2- ions adopt a fcc arrangement.
• Zn2 ions have a coordination number of 4.
• The S2- ions are placed in a tetrahedron around
  the Zn2 ions.
• Example CuCl.

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Bonding in Solids

• Ionic Solids
• Fluorite Structure
• Typical example CaF2.
• Ca2 ions in a fcc arrangement.
• There are twice as many F- per Ca2 ions in each
  unit cell.
• Examples BaCl2, PbF2.

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Bonding in Solids

• Metallic Solids
• Metallic solids have metal atoms in hcp, fcc or
  bcc arrangements.
• Coordination number for each atom is either 8 or
  12.
• Problem the bonding is too strong for London
  dispersion and there are not enough electrons for
  covalent bonds.
• Resolution the metal nuclei float in a sea of
  electrons.
• Metals conduct because the electrons are
  delocalized and are mobile.

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