Chapter 13 Intermolecular Forces in Liquids & Solids by pptfiles


									DISCLAIMER: these notes are provided to assist you in mastering the course material but they are not intended as a replacement of the
lectures. Neither do they contain the comments and ancillary material of the lectures; they are just a set of points that you might bring to the
lectures to annotate instead of having to write everything down and/or they may assist you in organizing the material after the lectures, in
conjunction with your own notes.

                                   Chapter 5: Solids, Liquids & Phase Transitions

       organization chart for both condensed phases
             forces
             properties
             phase changes

                          5.1 Bulk Properties of Phases & Kinetic Molecular Theory

 recall kinetic molecular theory (for gases):
         gas atoms or molecules widely-separated
         no forces of attraction between them
         atoms or molecules in continual, random, rapid motion
         kinetic energy determined by temperature
 first two above are unique to gases, latter two are largely true for liquids and solids too
 comparison of bulk properties, in terms of:
         molar volume of gases, much larger than liquids & solids
         compressibility of gases (to form liquids) large; not dramatic for liquids (to form
         thermal expansion of gases much larger than liquids & solids
         fluidity & rigidity
         diffusion
         surface tension (Fig. 5.4; not for gases)

                                             5.2 Types of Intermolecular Forces

 in order of decreasing strength, Fig. 5.9; also important for solutions (chapter 6)
        ion - dipole interactions (ion - ion interactions treated separately)
        dipole - dipole interactions
        dipole (or ion-)- induced dipole
        induced dipole - induced dipole interactions
 similar repulsive forces for all
 term: van der Waals forces used for those interactions not involving ions; define van der
Waals radius of atoms based on optimum between attractive and repulsive forces when ions not

Ion - Dipole Interactions (Fig. 5.6)
 dipole: charge separation, or (more likely) partial-charge separation, in a molecule i.e. - a
   polar molecule
 strength of ion - dipole interaction depends on:
        separation distance of ion and dipole
        charge on ion

Chem 59-110 (’02)

        magnitude of dipole
 several examples based on the polar molecule, water:
        Na + (g) + x H 2 O(l)  [Na(H2 O)+ ]( aq); H rxn = - 405 kJ / mol

        reaction is exothermic, heat given off = heat/enthalpy of hydration
          (hydration here; more generally, solvation)
 distance effect seen with alkali cations
        compare with H+, Hhydr = -1090 kJ/mol (i.e.- hydronium ion, H3O+)
        summary: Li+, -515; Na+, -405; Cs+, -263 kJ/mol
 similarly, charge and distance effect seen

Dipole - Dipole Interactions (Fig. 5.5)
 energy generally released when molecules condensed, taken up when the condensed phase
  vaporized, i.e.- accompanying the equilibrium: gas  liquid
 for polar molecules, this is due to dipole - dipole interactions
 comparison of boiling points (a measure of the heat required for vaporization) allows
  categorization as polar or non-polar molecules (compare pairs of similar molar masses)
 solubility considerations are also due to a matching of polarities of solute and solvent (“like
  dissolves like”)

Hydrogen Bonding
 special class of dipole - dipole interactions due to small size and low electronegativity of H
  when bonded to small, electronegative atoms, especially N, O and F
 eg.          X-- H+Y-
 the strongest are given in a Table (weaker ones with Cl, S, etc.)
 example, HF
 compare ethanol and dimethyl ether, both C2H6O:
        compare structures
        H-bonding in the alcohol, only dipole - dipole forces in the ether
        physical data:
                               dipole moment          melting pt.            boiling pt.
                                    1.69D                -114 C                 78oC
                                    1.30                 -142                  -25
 periodic trends in Fig. 5.10

Unusual Properties of Water
 H-bonding to extreme!!
 two H’s and two lone pairs (non-bonding pairs of electrons) on each O, Fig. 5.11
 form networks in the condensed phases
       perfect “diamond lattice” in ice, Fig. 5.12
       some disruption upon melting (still 85% of H-bonds), contraction of structure
       hence, ice has lower density than water; water max density at 4C, Fig. 5.13

Dispersion Forces: Interactions with Induced Dipoles
 weakest of all intermolecular interactions

Chem 59-110 (’02), ch. 5, Solids, Liquids & Phase Transitions

 two types:
       between polar and non-polar molecules, Fig. 5.7 (dipole (or ion-)- induced dipole)
              the larger the non-polar, the greater the interaction (eg. solubility of diatomic
               gases in water), due to polarizability
       between non-polar molecules, Fig. 5.8 (induced dipole - induced dipole)
       groups of examples in Table, trends in boiling points

Summary/Decision Tree for Intermolecular Interactions - Fig. Not in text

                               5.4 & 5.5 Phase Equilibria & Transitions

Physical Properties of Liquids
 distribution of kinetic energies of molecules in a liquid sample (similar to kinetic theory of
 equilibrium with gas phase (part of Fig. 5.17):
                                              vaporization, Hvap
                              liquid                                 vapor
                                              condensation, Hcond
 heat/enthalpy of vaporization, Hvap, is energy required (i.e.- endothermic) to escape
  intermolecular forces, Fig. 5.4
        eg. for water, Hvap = + 40.7 kJ/mol
 heat/enthalpy of condensation, Hcond, is energy liberated (i.e.- exothermic) on forming
  intermolecular interactions (same magnitude, opposite sign)

Vapor Pressure
 in a closed space above a liquid, Fig. 5.14, pressure in the gas phase stabilizes at a fixed value
  = equilibrium vapor pressure, dependent on temperature, Fig. 5.15 and Table 5.1; note:
        points on a line represent equilibrium pressure
        also a partial phase diagram - at a given T and P
                 points to left of line represent the liquid phase region
                 points to right of line represent the gas phase region
        boiling points at atmospheric pressure (see below)
 volatility, the tendency to escape into the gas phase ranked according to equilibrium vapor
 practical application: “water pump” in lab better in winter than summer

Boiling Point
 Fig. 5.15, line at 1 atm
 equilibrium vapor pressure equals atmospheric pressure at the boiling point
        in an open vessel vaporized molecules can escape
        note dependence on pressure, applications:
                cooking in Salt Lake City
                vacuum distillations

Chem 59-110 (’02), ch. 5, Solids, Liquids & Phase Transitions

Physical Properties of Solids
 disrupt the lattice to form a liquid (part of Fig. 5.17):
                                                 fusion, Hfusion
                              solid                                    liquid
                                                  freezing, Hcryst.
        fusion is endothermic, freezing/crystallization is exothermic
 also characterized by melting point, lowest temperature at which fusion occurs
 grouped according to forces; note ion-size dependence
 (note: sublimation also possible for solid  gas equilibrium; egs. H2O, I2, CO2,
naphthalene; I2 on 1st page photo)

                                      5.6 Phase Diagrams
 phase transitions, above
 influence of temperature and pressure on phases given in diagrams
        eg. H2O in Fig. 5.19: follow 2-phase lines (note negative slope to solid-liquid line);
         note triple point, freezing point, boiling point, critical point
        eg. CO2 in Fig. 5.21: note positive slope to solid-liquid line, critical point

Critical Temperature and Pressure
 keep increasing temperature in Fig. 5.19 and 5.21, two phases coalesce
 critical temperatures and pressures (both minima) shown
        new “phase”: supercritical gas/fluid has a density like that of the liquid but flow
           properties and ability of molecules to be separate from one another like a gas
 applications:
        liquefaction of gases, eg. air conditioning, fuels (must be below critical point)
        supercritical fluid extraction (must be above critical point), eg. CO2 (Tc = 31C, Pc =
           73 atm) used for decaffeinating coffee

Surface Tension, Capillary Action & Viscosity
 all are phenomena due to intermolecular interactions
 surface tension
        forces different in bulk liquid than at surface, Fig. 5.4; net, inward force at surface
        “skin” on surface, resists spreading as a film on another surface = surface tension
 surface layer interactions, in some cases, counteracted by interaction with another material
        eg. H2O with glass (H-O-H vs. Si-O-(H)), hence meniscus in a tube (Fig. 5.20),
           which is extreme in a very narrow tube (capillary)
        application: chromatography
 bulk liquid flow influenced by intermolecular interactions
        viscosity is the resistance to flow, which increases as the intermolecular interactions do
        eg. ethanol (two C’s) compared to longer chain alcohol, octanol (eight C’s) or to a
           “polyol” such as glycerol

Suggested Problems
 odd, 1 – 7; 13 – 17; 21 - 47

Chem 59-110 (’02), ch. 5, Solids, Liquids & Phase Transitions

To top