Intermolecular Forces Solids and Liquids by MikeJenny

VIEWS: 131 PAGES: 37

									Intermolecular Forces,
  Gases, and Liquids
         Ch.13




                         1
                         Gases
   Kinetic-Molecular Theory says molecules/atoms
    separated
   Little, if any, interactions
   Not so in solids and liquids

   Examples:
   Big difference in volume between liquids & solids and
    gases
   Gases compressible, liqs & solids not

                                                            2
               Intermolecular Forces
   Various electrostatic forces that attract
    molecules in solids/liqs
   Much weaker than ionic forces
   About 15% (or less) that of bond energies
       Remember, ionic bonds extremely powerful
            Boiling pt of NaCl = 1465 °C!




                                                   3
          Intermolecular Forces
   Reason behind importance of knowing about
    IMF:
   1) b.p. & m.p. and heats of vaporization (lg)
    and fusion (sl)
   2) solubility of gases, liquids, and solids
   3) determining structures of biochemicals
    (DNA, proteins)


                                                     4
     Remember dipole moments?
   Dipole moment = product of magnitude of
    partial charges (+/-) & their distance of
    separation
   = (1 Debye = 3.34 x 10-30 C x m)
   Important in IMF




                                                  5
    Ion-dipole: Ionization in aqueous
            medium (water)
   1) stronger attraction if ion/dipole closer
       Li+ vs. Cs+ in water
   2) higher ion charge, stronger attraction
       Be2+ vs. Li+ in water
   3) greater dipole, stronger attraction
       Dissolved salt has stronger attraction to water than
        methanol



                                                               6
7
                    Solvation energy
   Or, enthalpy of hydration (if water) = energy of
    ionization in aq. media
   Water molecules surround both ions
   Example:
   Take hydration energies of G I metal ions
     Exothermicity decreases as you go down the column
     Cations become larger
           Easier to dissociate

                                                       8
                 Permanent dipoles
   Positive end of one molecule attracted to negative end
    of other
       For ex: HCl
   Dipole-dipole attractions

   Cmpds that exhibit greater d-d attractions have higher
    b.p., and Hvap
   Polar cmpds exhibit greater d-d attractions than non-polar cmpds
       NH3 vs. CH4
        equivalent molar masses (g/mol): 17 vs. 16, respectively
       Boiling points: -33°C vs. -162°C, respectively

                                                                     9
                 Hydrogen Bonding
   A type of “super” dipole-dipole interaction
   Interaction between e--rich atom connected to H entity &
    another H attached to e—rich atom
   e--rich atom = O, F, N
   Density water > than ice
       Opposite of almost every other substance
   Inordinately high heat capacity of water
   High surface tension
       Insects walk on water
   Concave meniscus



                                                               10
           Hydrogen Bonding
   Boiling pts. of H2O, HF, and NH3 much higher




                                               11
               Surface Tension
   Outer molecules interact with surface, while
    inner interact with other molecules
   It has a “skin”
   Skin toughness = surface tension
   Energy required to break through surface
   Smaller surface area reason that water drops
    spherical


                                                   12
              Capillary Action
   When water goes up a small glass tube
   Due to polarity of Si-O bonding with water
   Adhesive forces > cohesive forces of water
   Creates a chain or bridge
   Pulls water up tube
   Limited by balancing gravity with
    adhesive/cohesive forces
   Thus, water has a concave meniscus

                                                 13
                    Mercury
   Forms a convex meniscus
   Doesn’t “climb” a glass tube
   Due to cohesive forces > adhesive forces




                                               14
                    Viscosity
   Hydrogen-bonding increases viscosity
   But large non-polar liquids like oil have:
   1) large unwieldy molecules w/greater
    intermolecular forces
   2) greater ability to be entangled w/one another

   Did you ever hear the expression, “You’re as
    slow as molasses in January”?
                                                   15
    Dipole/Induced Dipole Forces
   Polar entities induce dipole in nonpolar species like O2
   O2 can now dissolve in water
       If not, fishes in trouble!
   Process called “polarization”
   Generally, higher molar mass, greater polarizability of
    molecule
   Why?
   (larger the species, more likely e- held further away 
    easier to polarize)

                                                               16
Polarizability




                 17
         Induced dipole/induced dipole
                    forces
   Non-polar entities can cause temporary dipoles between other non-polar
    entities
    causing intermolecular attractions
        Pentane, hexane, etc.

   The higher the molar mass, the greater the intermolecular attractions

   N-pentane has greater interactions than neo-pentane
        Latter’s smaller area for interactions

   I2 has a higher Hvap & b.p. than other halogens
   cause nonpolar substances to condense to liquids
   and to freeze into solids
   (when the temperature is lowered sufficiently)
   Also called: London Dispersion Forces

                                                                             18
    Intermolecular Bonding Compared
   Strength
   Strongest: Ion-dipole
   Strong: Dipole-dipole (incl. H-bonding)
   Less strong: dipole/induced-dipole
   Least strong: induced-dipole/induced-dipole (London
    dispersion forces)

   Keep in mind  a compound can have more than one
    of the above!

                                                          19
                     Problem
   Rank the following in order of increasing boiling
    point and explain why:
   NH3, CH4, and CO2




                                                    20
             Properties of Liquids

   (l)  (g)
   Vaporization =
    endothermic
   Condensation =
    exothermic
   Boiling
       Why do we have
        bubbles?



                                     21
                    Vapor Pressure
   Leave a bottle of water
    open….
       Will evaporate
   Keep the lid on….
       can have equilibrium
        between liquid and gas
   Equilibrium vapor
    pressure/vapor pressure
       Measure of tendency of
        molecules to vaporize at
        given temp.
                                     22
What does this graph tell us?




                                23
                     Volatility
   Ability of liquid to
    evaporate
   Higher the vapor
    pressure, greater the
    volatility
   Are polar cmpds or non-
    polar cmpds of equal
    molecular mass more
    volatile?


                                  24
     Clausius-Clapeyron Equation
   Calculates ∆Hvap
   What is this an equation
    for?
   What are the variables?                  H vap    1
   C = constant unique to     Ln Pvap  -             C
    cmpd                                       R       T
   R = ideal gas constant
       8.314472 J/molK




                                                        25
     Clausius-Clapeyron Equation
   Or, if given two pts.:



   P2     H vap   1 1
ln( )          (  )
   P1      R       T2 T1


                                   26
     Clausius-Clapeyron Problem
   Methanol has a normal boiling point of 64.6°C
    and a heat of vaporization of 35.2 kJ/mol. What
    is the vapor pressure of methanol at 12.0°C?
   Does the answer make sense?
   Would water have a higher heat of vaporization?
     Why?
     Heat of vaporization of water = 40.65 kJ/mol




                                                     27
                    Boiling Point
   Bp  temp. at which vapor pressure = external
    (atmospheric pressure)
   At higher elevations atmospheric pressure is
    lower
       Thus, water boils at less than 100 °C




                                                28
    Critical Temperature and Pressure
   As temp. rises so does vapor pressure, but not infinitely
   At the critical point liq/gas interface disappears
   Critical temp/pressure
      Tc/Tp
       Gives supercritical fluid
          Density of a liq
          Viscosity of gas

   H2O:
       Tc = 374 °C
       Tp = 217.7 atm!
   Normal earth pressure  1 atm

                                                                29
             Supercritical fluid
   CO2 used in
    decaffeinating coffee
   Read about it on page
    614




                                   30
                Phase diagram
   Gives info on phase states of a substance at
    varying pressures and temperatures




                                                   31
     Deciphering a phase diagram
   Triple point
      Where all 3 states
       coexist
   Curves denote
    existence of two states
      Fusion (solid & liq)
      Vaporization (liq &
       gas)
      Sublimation (solid &
       gas)
      Off the lines
        Single   state            32
             Water’s phase diagram
   Graph explains why water
    boils at lower temps at higher
    altitudes (next slide)
   If you apply increasing
    pressure (const. T of 0°C) to
    ice will it convert to water?
   Solid-liquid line has negative
    slope
      It’s the opposite of most
         species
           Why?


                                     33
                      Sublimation
   Going from solid to
    gas without going
    through the liquid
    state
   Enthalpy of
    sublimation
      Hsublimation
   Iodine & dry ice (solid
    CO2) sublimate
   Opposite of
    sublimation
      Deposition (gs)
   Iodine demo

                                    34
            CO2’s Phase Diagram
   Explains
    sublimation
       How?


   Why is it called
    “dry ice”?




                                  35
Iodine’s Phase Diagram: But does it
          really sublimate?




                                  36
                            Problem
   The normal melting and boiling points of xenon are -112°C and
    -107°C, respectively.
   Its triple point is a -121°C and 0.371 atm and its critical point is
    at 16.6°C and 57.6 atm.

   a) Sketch the phase diagram for Xe, showing the axes, the four
    points given above, and indicating the area in which each phase
    is stable.
   b) If Xe gas is cooled under an external pressure of 0.131 atm,
    will it undergo condensation or deposition?



                                                                       37

								
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