Chapter 11: States of matter; liquids and solids

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Chapter 11 States of matter liquids and solids
Chemistry 1062 Principles of Chemistry II Andy
Aspaas, Instructor
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Gases, liquids, and solids

• Gas compressible fluid
• Mostly empty space
• Particles in constant random motion
• Vapor gas form of substance normally solid or
  liquid
• Liquid incompressible fluid
• Tightly packed particles
• Still in constant random motion
• Solid incompressible and rigid
• Tightly packed particles
• Particles only vibrate about their fixed sites

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Phase transitions

• Melting (fusion) transition from a solid to a
  liquid
• Freezing transition from a liquid to a solid
• Vaporization change of a solid or liquid to a
  vapor
• Sublimation change of a solid directly to a
  vapor
• Condensation change of a gas to a liquid or a
  solid
• Deposition change of a gas directly to a solid

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Vapor pressure

• Vapor pressure of a liquid partial pressure of
  vapor over the liquid, measured at equilibrium
• Vaporization and condensation are happening
  continually
• When the rate of condensation equals the rate of
  vaporization, equilibrium has been reached
• Vapor pressure depends on temperature
• More kinetic energy in the liquid makes
  vaporization occur more quickly
• Volatile liquids and solids have high vapor
  pressure

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Boiling point and melting point

• Boiling point temperature at which the vapor
  pressure of a liquid equals the atmospheric
  pressure
• Bubbles form in the liquid when its heated
  enough that its vapor pressure equals external
  pressure
• Once boiling begins, the temperature of the
  liquid stops rising
• Freezing point temperature at which a pure
  liquid changes to a solid
• Identical to melting point temperature at which
  a solid changes to a liquid

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Heat of phase transition

• Heat of fusion amount heat required to melt a
  solid
• H2O(s) ? H2O(l) ?Hfus 6.01 kJ/mol
• Heat of vaporization amount of heat required to
  vaporize a liquid
• H2O(l) ? H2O(g) ?Hvap 40.7 kJ/mol
• Heating curveno temperaturechange during a
  phase transition

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Phase transitions

• Heat must be added for melting or vaporization
• Endothermic processes
• Positive value of ?H
• Heat must be removed for condensation or freezing
• Exothermic processes
• Negative value of ?H

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Clausius-Clapeyron equation

• Vapor pressure of a liquid depends on temperature
• Also depends on ?Hvap for that liquid
• Clausius-Clapeyron equation
• The two-point form removes the constant B from
  the equation

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Phase diagrams

• Phase diagram for water
• Curves show exptl points at which phases are in
  equilibrium
• Solid-liquid line (melting-point curve)
• Leans left if liquid is more dense than solid (as
  in water)
• Otherwise leans right

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Phase diagrams

• Curve AC gives the vapor pressure at any
  temperature
• Triple point, A point at which all 3 phases
  exist in equilibrium
• Sublimation occurs when heating at pressures
  below triple point
• Critical point, C point above which gas and
  liquid are no longer distinct
• Supercritical fluid at temperature and pressure
  above critical point

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Liquid state surface tension

• Surface tension energy required to increase the
  surface area of a liquid by a unit amount
• Liquids will minimize their amount of surface
  area (since intermolecular forces pull molecules
  at the surface inward)
• Dissolved substances reduce surface tension by
  interrupting intermolecular forces
• Capillary rise when a liquid is attracted to a
  solid capillary (like water to glass) it will
  rise in the capillary to minimize surface tension
  at the inverted meniscus

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Viscosity

• Viscosity resistance to flow in a liquid or gas
• Strong intermolecular forces increase viscosity
  and cause a liquid to flow more slowly (like
  syrup)
• Motor oil viscosity in SAE units
• 10W/30 has SAE 10 in winter, 30 in summer
  (because viscosity of engine oil increases as
  temperature decreases)

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Intermolecular forces

• Dipole-dipole force attractive interaction
  between polar molecules
• Polar molecules have a dipole moment
• ? and ? sides of the molecule resulting from
  electronegativity differences
• Negative side of one molecule will be attracted
  to the positive side of another molecule
• London dispersion forces attractive forces
  resulting from instantaneous induced dipoles
  caused by random electron groupings in an atom
• High molecular weight higher London forces
• Compact molecule lower London forces

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Predicting boiling points, vapor pressure, and
viscosity

• Polar molecules have stronger intermolecular
  forces than nonpolar molecules of similar MW
  (dipole-dipole)
• Larger molecules have stronger intermolecular
  forces than smaller molecules of similar polarity
  (London)
• Strong intermolecular forces
• Low vapor pressure (less chance for molecule to
  break free)
• High boiling point (high temperature required to
  raise vapor pressure to atmospheric pressure)
• High viscosity (molecules cling strongly
  together)

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Hydrogen bonds

• Hydrogen bonds much stronger than dipole-dipole
  or London forces
• Interaction between H and Y in XHY
• Where X and Y are F, O, or N (electronegative
  atoms)
• H has a strong ? when covalently bonded to an
  electronegative atom
• Explains why CH3OH boils at 65 C and CH3F boils
  at 78 C (both have similar MW and polarities)

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Classification of solids

• Molecular solid solid that consists of atoms or
  molecules held together by intermolecular forces
• H2O(s), CO2(s)
• Metallic solid any solid metal contains a
  regular arrangement of positive metal nuclei
  surrounded by a sea of delocalized electrons
• Ionic solid consists of cations and anions held
  together by electrostatic attractions of opposite
  charges (ionic bonds)
• Covalent network solid consists of atoms held
  together in large networks or chains by covalent
  bonds

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Melting point of solids

• Molecular solid low melting point (below 300 C)
• Only weak intermolecular forces to overcome
• Ionic and covalent network solids high melting
  point (800 C or above)
• Ionic and covalent bonds require much more energy
  to be broken
• Larger charges make even higher mps in ionic
  solids
• Metals variable (Hg 39 C W 3410 C)

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Hardness and conductivity of solids

• Molecular solids generally soft and brittle
• Ionic solids generally hard and brittle
• 3-dimensional covalent network solids are
  tremendously hard
• Metals are malleable since the atoms can move
  past each other
• Metals (and molten ionic solids) conduct
  electricity because electrons can move freely

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Crystalline solids

• Crystalline solids have an ordered 3-dimensional
  structure of particles
• Amorphous solids have their particles locked into
  random positions
• Crystal lattice 3-dimensional geometric
  arrangement of particles in a crystal
• Unit cell smallest boxlike unit from which a
  crystal lattice can be constructed
• Unit cells are stacked in 3 dimensions

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Cubic unit cells

• Simple cubic unit cell atoms only at the corners
  of a cubic unit cell
• Only 1/8 of a corner atom is contained in a unit
  cell
• Body-centered cubic atom at the center of a unit
  cell as well as the corners
• Face-centered cubic atoms in the center of each
  of the faces as well as the corners
• Only 1/2 of a face atom is contained in a unit
  cell

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Cubic unit cells
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Molecular solids closest packing

• In simple molecular solids (such as individual
  noble gas atoms), nondirectional London forces
  govern the packing of the atoms together to form
  a solid
• Spheres pack together most closely in a hexagonal
  honey-comb type arrangement on one layer
• A second layer can be nested in half of the
  crevices left by the first layer
• A third layer can be nested in the crevices in
  the second layer, but there are two possible
  positions for this layer

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Close packing arrangements

• Hexagonal close packing (hcp) when 3rd layer is
  identical to 1st (in x holes left by 2nd layer)
• Cubic close packing (ccp) when 3rd layer is
  different (in y holes left by 2nd layer)

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Cubic close-packing

• Cubic close-packing has a face-centered cubic
  lattice structure (tilted on its side)

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Coordination number

• Coordination number number of nearest-neighbor
  adjacent atoms any one atom can have
• Close packing structures CN 12
• hcp, ccp/face-centered cubic
• Body-centered cubic CN 8
• Square cubic CN 4

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Ionic solids

• 3 common structures CsCl structure, NaCl
  structure, and zinc blende (cubic ZnS) structure
• CsCl structure
• 1 of eachper unit cell
• When ions are similarin size
• Expandedsimple cubic

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NaCl structure

• Expanded face-centered cubic structure
• When cation and anion are moderately different in
  size

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Zinc blende structure

• When cation and anion are significantly different
  in size
• Anion is expanded body-centered cubic
• Cation is in 4 opposite tetrahedral quadrants
  completely inside the unit cell

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Unit cell calculations

• Length of side of a unit cell
• Volume of a unit cell
• Mass of a unit cell
• Mass of individual atoms
• Molar mass
• Forward or backwards!