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					Liquids & Solids

 Advanced Chemistry 30S
       M. Patenaude
       GPHS Science Dept
A. Three Types of Molecular Motion

                Gases   Liquids     Solids

Translational
   Motion
                Free    Hindered    None

 Rotational
  Motion
                Free    Hindered   Hindered

 Vibrational
   Motion
                Free     Free       Free
   Molecules in solids and liquids are very
    close together. The average distance
    between particles is less than one
    molecular diameter.

   Molecules in gases, on the other hand,
    are very far apart. The average distance
    between particles is about 10 molecular
    diameters.
B. Intermolecular Forces
   forces of attraction between molecules or atoms
   generally much weaker than covalent bonds
     Only 16 kJ/mol of energy is required to overcome the
      intermolecular attraction between HCl molecules in the
      liquid state (i.e. the energy required to vaporize the
      sample)
     However, 431 kJ/mol of energy is required to break
      the covalent bond between the H and Cl atoms in the
      HCl molecule

Thus, when a molecular substance changes states
   the atoms within the molecule are unchanged
   The temperature at which a liquid boils reflects the
    kinetic energy needed to overcome the attractive
    intermolecular forces (likewise, the temperature at
    which a solid melts).


    Thus, the strength of the intermolecular forces
       determines the physical properties of the
                       substance
Types of Intermolecular Forces:
   Ion-Dipole Interactions
    – between an ion and a polar molecule
    – the forces depend on three factors:
      • The magnitude of the ion’s charge
      • The magnitude of the dipole
      • The distance between the ion and the dipole
– An example occurs in the hydration of an
  ion in water. For example, consider the
  dissolving of table salt in water:
  NaCl (s)  Na+ (aq) + Cl- (aq)

– Each ion interacts with water and becomes
  “hydrated” - surrounded by polar water
  molecules.

– As the ions are surrounded, energy is
  released. (Exothermic process)
   Dipole-Dipole Interactions
    – Forces of attraction between polar
      molecules.
    – Weaker than ion-dipole forces
    – require close proximity of molecules
    – More polar molecules = stronger dipole-
      dipole interactions
   Hydrogen Bonding
   Hydrogen bonds are considered to be dipole-
    dipole type interactions
   A bond between hydrogen and an electronegative
    atom such as F, O or N is quite polar:




   The hydrogen atom has no inner core of electrons, so
    the side of the atom facing away from the bond
    represents a virtually naked nucleus
   This positive charge is attracted to the negative
    charge of an electronegative atom in a nearby
    molecule
   this side of the hydrogen atom can get quite close to
    a neighboring electronegative atom (with a partial
    negative charge) and interact strongly with it

     Hydrogen bonds vary from about 4 kJ/mol to 25 kJ/mol
      (so they are still weaker than typical covalent bonds.
     But they are stronger than dipole-dipole and or
      dispersion forces.
     They are very important in the organization of
      biological molecules, especially in influencing the
      structure of proteins
   Water is unusual in its
    ability to form an extensive
    hydrogen bonding network


   Each water molecule can
    participate in four
    hydrogen bonds
     One with each non-bonding
      pair of electrons
     One with each H atom
   London Dispersion Forces
   Non-polar molecules would not seem to have
    any basis for attractive interactions.
     However, gases of non-polar molecules can be
      liquefied and solidified, at low temperatures -
      indicating that if the kinetic energy is reduced,
      some type of attractive force can predominate.
     Fritz London (1930) suggested that the motion of
      electrons within an atom or non-polar molecule
      can result in a temporary dipole moment
   As an atom or molecule becomes polarized (it forms
    an instantaneous dipole), it can induce an opposite
    dipole in a neighboring atom or molecule.
   These temporary dipole-dipole attractions form and
    dissolve over and over again
   The larger the atom/molecule, the more polarizable it
    will be - and the greater the London Dispersion
    Forces! (e.g. Consider the Halogens)
C) The Liquid State
   Properties of Liquids:
    – very low compressibility
    – take the shape of their container
    – high density (comp. to gases)

Many properties are related to
 forces between liquid
 molecules:
• Interior molecules experience
  attractions all around them
• Surface molecules experience
  attractions from their sides
  and below
   As a result of uneven forces, molecules at the
    surface are pulled inward, creating a “sphere” - a
    bead, or drop, of liquid

   COHESIVE forces are the forces between
    molecules within the liquid
    – Strong cohesive forces are the result of strong
      intermolecular forces
    – They result in high SURFACE TENSION - the
      resistance of a liquid to an increase in its surface
      area
    – They also result in high VISCOSITY - the
      resistance of a liquid to flow. (Another factor is
      molecular “complexity”)
    – Note that viscosity decreases with increased
      temperature (why?)
   ADHESIVE forces are the forces between liquid
    molecules and their container
    – These forces are strong if the liquid is polar and the
      surface is polar (or if the liquid is non-polar and the
      surface is non-polar)
    – A liquid will WET a surface if there are strong adhesive
      forces
    – CAPILLARY ACTION - the spontaneous rising of a
      liquid up a tube - is explained by strong adhesive
      forces between the liquid and the tube’s surface, as
      well as the strong cohesive forces between liquid
      molecules
D) Structure & Bonding in Solids
   Crystalline solids have a characteristic
    regular arrangement of particles.
    – The positions of the particles can be
      represented by a LATTICE
    – The smallest repeating unit of a crystal
      lattice is called a UNIT CELL
    – We will study three common types of unit
      cells: simple cubic (sc), body centered
      cubic (bcc) and face centered cubic (fcc)
Simple Cubic Unit Cells

   8 atoms at the
    vertices of a cube
   Two atoms touch
    along each edge of
    the cube:
            2r = s
    where “r” = atomic radius    How many atoms
     and “s” = length of side   are inside the cell?
Body Centered Cubic

                          Similar to simple
                           cubic…with an extra atom
                           in the center of the cube
                          Atoms are forced apart
                           along the edges
                          3 atoms touch along the
                           diagonal of the cube:
 How many atoms
are inside the cell?               4r = 3 s
Face Centered Cubic

                          8 atoms at the vertices
                           of a cube, with an
                           additional atom in the
                           center of each face of
                           the cube
                          3 atoms touch along the
                           diagonal of a face:
 How many atoms
                                  4r = 2 s
are inside the cell?
Holes inside Unit Cells
   A hole formed by 3 atoms in a
    plane is called a trigonal hole:

   A hole when a fourth atom is
    placed on TOP of these three is
    called a tetrahedral hole:

   A hole formed by SIX atoms,
    four in a square plane, with one
    above and one below, is called
    an octahedral hole!
Holes in the Unit Cells, cont’d
   Prove to yourself that a FCC cell has 8
    tetrahedral holes
   Prove to yourself that a FCC cell has an
    octahedral hole in its center
   Prove to yourself that there is an octahedral
    hole in the center of each edge of the FCC
    unit cell

    What kind of holes are present in BCC
                   unit cells?
Closest Packing
   The first two crystal lattices (sc & bcc) begin
    with a square layer of atoms
   There is a more efficient way of packing
    atoms - beginning with a hexagonal array of
    atoms
Square packed - simple cubic
Square packed - body centered
Body centered cubic
Body Centered Cubic
Closest packing
   In a closest packed crystal, the second layer
    of atoms will be placed over the holes in the
    first layer:




    Notice that there are now TWO types of holes in the
                    blue layer of atoms!
Closest Packing
   If atoms are now placed on top of holes in the
    second layer, above atoms in the first layer,
    an arrangement of atoms “A-B-A-B-A…” will
    result:

      This
 arrangement
   leads to a
HEXAGONAL
  CLOSEST
   PACKED
     (HCP)
   Structure
Hexagonal closest packing

                     Be
                     Co
                     Mg
                     Zn
   If the third layer is placed over top of holes
    in BOTH the first and second layers, an
    arrangement of atoms “A-B-C-A-B-C-…”
    will result:

       This
  arrangement
    leads to a
     CUBIC
   CLOSEST
 PACKED (ccp)
 Structure, with
a Face-Centered
                        Note that both closest packed
Cubic Unit Cell!
                       structures have a coordination
                              number of “12”
Cubic closest packing

                  Ag    Ni
                  Al    Pb
                  Au    Pt
                  Ca
                  Cu
 Coordination Numbers
                    Coordination
    Structure                    Stacking Pattern
                      Number
  Simple cubic           6         AAAAAAA
 Body Centered
                         8        ABABABAB
     cubic
Hexagonal Closest
                         12       ABABABAB
    Packed
  Cubic Closest
                         12      ABCABCABC
    Packed
 Bonding in Metals
    The valence electrons in metals are delocalized -
     they are free to move from atom to atom.
    This is why metals conduct electricity
    Metal atoms can be thought of as cations floating in a
     sea of electrons”

        -       +   -      +     -     +     -
        +   -        +       -   +    -       +
       -        +
                     -     +
                                 -     +     -
The valence electrons are “shared” among all the atoms
in the metal - DELOCALIZED COVALENT BONDING
Metal Alloys
   An alloy is a metallic solid made up of a
    mixture of elements.
    – Substitutional Alloys
       • some metal atoms are replaced by other metal
         atoms in a crystal lattice
       • Brass (67% Cu, 33% Zn)
       • Bronze (~93% Cu, 7% Sn)
       • Sterling Silver (93% Ag, 7% Cu)
– Interstitial Alloys
   • Small atoms fill some of the holes (interstices)
     between atoms in the crystal lattice
   • For example, carbon is added to steel (0.2% up
     to 1.5%)
   • The added carbon atoms
      form strong directional
      covalent bonds with the
      iron atoms, making the
      metal stronger
Network Solids
   A network solid is an atomic solid where the atoms
    are bonded with strong “directional covalent” bonds.
   The network solid is really one giant covalent
    “molecule”
   The two best examples are Diamond and Graphite
    (two allotropes of carbon)




   A great resource for viewing solid structures (with Chime):
    http://learn.chem.vt.edu/archive/apache/htdocs/105a/crystal/crystal.html
Graphite
   A network solid where carbon atoms are covalently
    bonded together in planes.
   Each carbon atom is bonded to three other atoms
   The planes are held together by weak London
    Dispersion Forces, and slide off easily
Diamond
   Diamond is also a network solid, but in 3-dimensions.
   Carbon atoms are bonded to four other atoms, in a
    tetrahedral geometry
   Because of this 3-d network of strong, directional
    covalent bonding, diamond is the hardest natural
    substance
Another look at NaCl
E. Heating Curves & Changing State
   If a substance is heated slowly, a plot of
    temperature vs time can be created.
    Such a plot is called a heating curve.

   Plateaus will occur wherever a change
    of state occurs. The first plateau
    represents melting (fusion) and the
    second plateau boiling (vaporization).
   The energy (heat) required to melt one
    mole of a substance is called the molar
    enthalpy of fusion for the substance,
    DHfus.

   For water, DHfus is 6.0 kJ/mol. 6.0 kJ of
    heat must be absorbed to melt one
    mole of ice.
   The amount of energy (heat) needed to
    boil one mole of water is called the
    molar enthalpy of vaporization of water,
    DHvap.

   For water, DHvap is 40.7 kJ/mol.

   Recall: 1 mol water = 18 g = 18 mL
  The Heating Curve of Water
                 140
                 120
Temperature/°C




                 100
                 80
                 60
                 40
                 20
                  0
                 -20
                       0   100   200     300      400   500   600
                                       Time/min
   A positive enthalpy change, DH, means
    that energy is being absorbed by the
    solid or liquid during melting and boiling.

   When a solid melts, the molecules in
    the liquid state have more energy than
    those in the solid state.

   When a liquid boils, the molecules in the
    gaseous state have more energy than
    those in the liquid state.
   The conversion of a solid directly to a
    gas (without melting) is called
    sublimation.

   The molar enthalpy of sublimation,
    DHsub, is the energy needed to cause
    one mole of a solid to sublime.

   Dry ice (solid CO2), iodine, naphthalene
    all sublime at room temperature and 1
    atm pressure.
F. Vapor Pressure
   A liquid in a closed container will start to
    evaporate. As more vapor is introduced
    into the space above the liquid, some of
    the vapor condenses back to the liquid
    state.

   At first, the rate of evaporation is greater
    than the rate of condensation.
   The rate of evaporation remains
    constant, and the rate of condensation
    increases over time.
   Eventually, the two rates are equal. At
    this point, the system is in a state of
    “equilibrium” or balance.
   The pressure of the vapor above the
    liquid at this point of equilibrium is called
    the vapor pressure of the liquid.
   A liquid that has a very HIGH vapor
    pressure is said to be very VOLATILE.
    This means it will evaporate very easily.
   Volatility is determined by two main
    factors: Molar Mass & Intermolecular
    Forces
    – Large molar mass = less volatile
    – Strong intermolecular forces = less
      volatile
   Vapor pressure increases with
    temperature.
   A sketch of a vapor pressure curve
    looks like:
                                A
        Vapor Pressure




                                       B



                         Temperature
   The equation that relates the vapor
    pressure of a liquid to its temperature:

                       DHvap  1 
        ln( Pvap )           C
                         R T 

      Where DHvap is the molar enthalpy of
      vaporization for the liquid.
   This equation is can be interpreted as a
    straight line equation:

       y    =     m     x + b
                    DHvap  1 
     ln( Pvap )           C
                      R T 
So a graph of ln(Pvap) vs (1/T) should
be linear. The slope of the line would
allow the calculation of DHvap.
The Clausius-Clapeyron Equation:
   A more useful form of the first equation
    that allows the calculation of DHvap if
    only TWO vapor pressures and
    temperatures are measured.


           P1  DHvap  1 1 
        ln            
           P2    R  T2 T1 
Normal Melting Point

   Is defined as the temperature where the
    solid and liquid states have the same
    vapor pressure under conditions where
    the total pressure is 1 atmosphere
    (standard pressure).
   Refer to apparatus in Zumdahl (Fig
    10.44)
Normal Boiling Point

   Is defined as the temperature where the
    vapor pressure of a liquid is exactly one
    atmosphere (standard pressure).
   When the vapor pressure reaches 1
    atm, bubbles of vapor are able to form
    inside the liquid (which is what we see
    when the liquid “boils”)
   Supercooling occurs when a liquid
    remains “liquid” when it is slowly cooled
    below its freezing point.
   As the water cools, its molecules must
    become “organized” into the solid
    crystal shape. The temperature may
    drop below 0°C before this organization
    occurs.
   As the crystal forms, heat is released
    and the temperature rises back to 0°C
    as the liquid freezes.
   Superheating occurs if a liquid’s
    temperature rises above its boiling point
    before boiling starts.
   This may occur if the liquid is heated
    very rapidly.
   To form vapor bubbles within the liquid,
    many high-energy molecules must
    congregate. This may not happen
    immediately if the liquid is heated very
    quickly.
   Once enough molecules DO
    congregate, a bubble will form.
    Because the temperature is higher than
    the boiling point, the pressure of vapor
    will be higher than 1 atm.
   This may cause the bubbles of vapor to
    “burst” as the liquid starts to boil.
    (Called “bumping”)
   Boiling chips are added to a liquid as it
    is heated. The boiling chips contain
    trapped air that is released as the
    temperature rises.
   This creates small bubbles of gas in the
    liquid that act as “starters” for vapor
    bubbles to form.
   This prevents superheating from
    occurring.
Phase Diagrams

label axes
label phase regions
label: triple point
       critical point
       melting point
       boiling point
        sublimation point
Critical Point

   The temperature and pressure at which
    the liquid and gaseous phases of a pure
    stable substance become identical.
   The critical temperature of a gas is the
    maximum temperature at which the gas
    can be liquefied; the critical pressure is
    the pressure necessary to liquefy the
    gas at the critical temperature.

				
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