Chapter 11. Intermolecular Forces and Liquids and Solids

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					Chapter 11. Intermolecular Forces and Liquids and Solids
     In this chapter the concepts ionic and covalent bonding (the molecular shape and
     polarity of covalent compounds) are used to explain the properties of solids and

     Intermolecular Forces

           Inter vs. Intra Molecular Forces

           Intra    -    strong forces within molecules (covalent or ionic bonds)
                         (related to chemical reactivity of the substance)

           Inter    -    weaker forces between molecules
                         (determine the bulk physical properties of a substance)

11.1 Kinetic Molecular Theory of Liquids and Solids
     Phase -       different states of a substance that are present in a system;
                   homogeneous parts separated by a well-defined boundary

     A. Gases --- Chapter 5
        • molecules were very far apart
        • readily compressed
        • no strong forces between molecules

     B. Liquids
        • Molecules close together - held together by several types of attractive
        • Difficult to compress (molecules are already close together; moving even
           closer results in repulsive forces
        • Denser than gases
        • Definite volume
        • Molecules still move freely (liquids flow) and assume shape of container

     C. Solids
        • Molecules held rigidly (no free motion)
        • Almost incompressible
        • Definite shape and volume
        • Usually even more dense than liquids (H2O is an exception)

     Example: As a gas cools, the movement of the molecules slows and eventually
               they do not have enough kinetic energy to overcome the attractive
               forces between them

11.2 Intermolecular Forces

             •   Attractive forces between molecules

             •   Weaker than bonds

             •   Responsible for non-ideality of gases

             •   Stronger in solids and liquids

      A. Forces in PURE substances:

             1. dipole-dipole forces -- between polar molecules (e.g., SO2, PF3)
                                                   [Figure 11.1]

             2. Dispersion forces (London forces)
                 •   very weak "instantaneous        induced    dipole"   forces   between
                     non-polar molecules
                 •   instantaneous dipole –

                            •   depends on the polarizability (ease of distortion of an
                                electron cloud); larger atoms or molecules are more

                            •   can be very strong
                                      CH3F               mp = -141.8 °C (polar molecule)
                                       CCl4              mp = - 23 °C (nonpolar molecule)

                                       Thus the dispersion forces in nonpolar CCl4 are
                                          stronger than in polar CH3F

             3. H-bonding

                     •   especially strong dipole-dipole forces for compounds with
                         hydrogen in a polar bond ( H-F, H-O, or H-N bonds)

                     •   Figure 11.6

Chapter 11                                                           Chemistry 1303

                         HYDROGEN BONDING
                              ..O.                               H
                         H           H
                 .. ..                            H N       H N           H—O

             H           H                            H          H

                   H N                                            ..

                                                   H N       H    F


                         H                                        ..
                             .. ..                      H
                         H           H

      •   Evidence: higher boiling points for compounds with H-bonding        [Figure 11.7]
      •   H2O has highest boiling point because each molecule can form up to 4
          hydrogen bonds
      •   Biological examples are numerous: proteins, nucleic acids

                   Which species form H-bonds: H2S, C6H6, CH3OH?

      B. Forces in MIXTURES

             1 -3 above plus

             4. ion - dipole -- between ionic and polar substances [Figure 11.2]

             5. ion - induced dipole -- between ionic and non-polar substances
                                           [Figure 11.4]

             6. dipole - induced dipole -- between polar and non-polar substances
                                                  [Figure 11.4]

      van der Waals forces -----         another (older) name for some of the intermolecular
             dipole: dipole                   dipole: induced dipole          dispersion

Chapter 11                                                             Chemistry 1303

      What is the predominate intermolecular force in each of the following

               (a)       solid CO2

               (b)       liquid CH3CH2OH

               (c)       liquid SCl2

               (d)       MgCl2 dissolved in liquid SCl2

               (e)       CO2 dissolved in SCl2

      C. Bulk Properties of Liquids and Solids (related to intermolecular forces)

                     •   melting and boiling points
                     •   compressibility
                     •   diffusion
                     •   surface tension
                     •   evaporation (liquid to gas)
                     •   sublimation (solid to gas)

11.3 Properties of Liquids
             many phenomenon are due to intermolecular forces, but only three are
             covered here.

      B. Boiling Point and Melting point

                     Arrange the following in order of increasing boiling points.
                                  NH3         PH3         CH4       SiH4

                                CH4    <   SiH4 <      PH3   <   NH3
                                            (lowest)                       (highest)

             Figure 11.6

Chapter 11                                                             Chemistry 1303

      B. Surface Tension
             •   Amount of energy required to stretch or increase a surface
             •   e.g., an "elastic skin" on top of water
             •   Intermolecular forces acting on a molecule in the surface layer of a liquid
                                                                [Figure 11.8]

             1. Wax on a car causes water to "bead" or stick to itself          [Figure 11.9]
                        The polar water molecules are more attracted to each other than to
                        the non-polar molecules of wax

             2. Capillary action
                      Figure 11.10 - water spontaneously rises in a capillary tube
                        Adhesion - attraction of unlike molecules (water and glass)
                        Cohesion - attractive forces between like molecules
                                    In a capillary with water, adhesion > cohesion
                        Not the same for all liquids, e.g., mercury [Figure 11.10 on right]
                               does not rise in tube
                                                Cohesion > adhesion

      B. Viscosity ("slow as molasses in the winter")
                 •   measure of a fluid's resistance to flow
                 •   strong intermolecular forces ⇒ higher viscosity
                        molecules "stick" together so flow is reduced
                 •   example: glycerol
                                            HO   CH2
                                            HO   CH
                                            HO   CH2

                        three hydrogen bonds - strong attraction
                        shape also allows tangling of molecules
                        both reduce flow, i.e., increase viscosity

Chapter 11                                                              Chemistry 1303

      C. Structure and Properties of WATER
             Most unique compound on Earth
                      WHY is it unique? - HYDROGEN BONDING

             Properties of water
                 a.      excellent solvent for ionic compounds and excellent solvent for
                         substances capable of forming H-bonds
                             participates in most of the reactions in living systems
                 b.      high specific heat
                 c.      moderates temperature on earth
                             •   more than the expected amount of energy to raise
                                 temperature because must break H-bonds;
                             •   thus temperature rises only a small amount even when
                                 large amounts of energy are added
                             •   opposite effect; large amounts of heat are given off (e.g.,
                                 lakes) with only small temperature changes
                 d.      solid form less dense than liquid
                             •   ice floats
                             •   lakes freeze from the top down thus sustaining life
                                    beneath the surface
                             •   potholes form
                             •   you can't freeze drinks without the container breaking
                             •   H2O forms two H-bonds per molecule giving 3-D network
                                   that "pushes" the water molecules apart in the solid
                                    state                              [Figure 11.12]
                             •   Melting breaks some of the H-bonds and allows
                                    molecules to move closer together
                             •   Maximum density at 4 °C                [Figure 11.13]

                 e.      unusually high boiling and freezing points

Chapter 11                                                            Chemistry 1303

                Two general types of
               •   crystalline - rigid, ordered arrangement of atoms, ions, or molecules
               •   amorphous - random disordered arrangements of atoms, ions, or

Solids Part 1 ---- Crystalline solids
11.4     Crystal Structure
       A. Unit Cell
               •   Basic repeating structural unit [Figure 11.14]
               •   7 types [Figure 11.15]
                   •   simple cubic
                   •   tetragonal
                   •   orthorhombic
                   •   rhombohedra
                   •   monoclinic
                   •   triclinic
                   •   hexagonal

       B. Packing Spheres
             the way atoms (spheres) are arranged in layers
             consider only cubic cells here:
               1. Simple cubic cell (scc)          [Figure 11.16. (top and side views)]
                   •   sphere at 8 corners of a cube
                   •   Each sphere touches 6 other spheres - coordination number = 6
                   •   Each sphere shared by 8 unit cells; 8 corners to a cube, so have
                       equivalent of one whole sphere per cubic unit cell
                               see Figure 11.22
               2. Body-centered cubic cell (bcc)                       [Figure 11.17]
                   •   simple cubic plus sphere at center of cube
                   •   coordination number is 8
                   •   two complete spheres per unit cell
                       •   one from the middle sphere and one from the 8 corners

Chapter 11                                                            Chemistry 1303

               3. Face-centered cubic (fcc)                       [Figure 11.17]
                   •   simple cubic plus sphere at center of each face of the cube
                   •   coordination number 12
                   •   four complete spheres per unit cell
                       •   3 from the six face centered atoms (1/2 from each of 6 faces)
                       •   one from the shared corner spheres

      C. Closet Packing
                       the most efficient arrangement of spheres
               •   hexagonal close packing [Figure 11.20 a, b, c]
               •   cubic close packing [Figure 11.20 a, b, d]

      •   reason for one packing over the other is related to intermolecular forces
      •   relationship between atomic radius and the size of unit cell        [Figure 11.22]

                       When silver crystallizes, it forms face-centered cubic cells. The unit
                       cell edge length is 408.7 pm. Calculate the density of silver.

               D = mass/Volume              Face centered cube has 4 whole atoms

          4 atoms x               1 mole          x 107.9 g       =   7.167 x 10 -22 g
                           6.022 x 1023 atoms       mole

               Volume of cube = length3
    length = 4.087 x 102 pm           x     1m        x       102cm   =    4.087 x 10-8cm
                                          1012 pm              1m

                       V = [4.087 x 10-8cm]3 = 6.827 x 10-23 cm3
                           D =     7.167 x 10 -22 g       =    10.51 g/cm3
                                 6.827 x 10-23 cm3

Chapter 11                                                                Chemistry 1303

11.5 X-Ray Diffraction

      •    scattering of X-rays by the units of a crystalline solid
      •    Figure 11.23 and 11.24
      •    May be used to determine crystal lattice and molecular structures

      •    Bragg equation -- basic mathematical tool of x-ray diffraction
                                (need not memorize)

11.6 Types of Crystals
               structures and properties of crystals (MP, density, hardness)
               determined by the type of forces holding the particles together

      A. Ionic Crystals
           •   hard, brittle, high melting point, poor conductors of heat and electricity
           •   composed of charged species
           •   anion and cations are general very different sizes

             How many Na+ and Cl- ions are in each NaCl unit cell?

               •   NaCl is face-centered cube [Figure 2.13]
               •   Na+ count
                   •   one Na+ at center
                   •   12 Na+ at edges
                   •   each Na+ shared by four unit cells
                               1 (center) + (12 x 1/4) = 4 Na+ ions
               •   Cl- count
                   •   6 Cl- at face centers
                   •   8 Cl- at corners
                   •   each face-centered is shared by 2 unit cells
                   •   each corner shared by 8 unit cells
                                      (6 x 1/2) + (8 x 1/8) = 4

Chapter 11                                                             Chemistry 1303

         Copper crystallizes in a face-centered cubic lattice (the Cu atoms are at the
         lattice points only). If the density of the metal is 8.96 g/cm3, what is the unit
             cell edge length in pm?

                       m               m
                 d=               V=
                       V               d

                           63. 55 amu x 4 x              1g
                                                    6.0222 x 1023 amu
                                           8.96 g


             V = 4.7096 x 10-23 cm3

             edge = 3.61 x 10-8 cm x 1 m          x 1012 pm
                                                            = 361 pm
                                           100 cm      m

      B. Covalent Crystals
             •   atoms (not ions) held together in extensive 3-dimensional network
             •   examples - two allotropes of carbon [Figure 11.28]
                 •   graphite sp2 hybridized (each C bonded to 3 others)
                        •    good conductor (delocalized electrons in remaining p orbitals)
                        •    good lubricant because layers "slip" past each other
                 •   diamond sp3 hybridized (each C bonded to 4 others)
                        •    very hard (strong bonds in 3-dimensions)
                        •    high melting point (3550 °C)
             •   example: quartz, SiO2,
                        •    3-D like diamond but with O between Si atoms
                        •    high melting point and hard

Chapter 11                                                              Chemistry 1303

      C. Molecular Crystals
             •   generally soft with low melting point, poor conductors
             •   lattice points are molecules
             •   held together by van der Waals forces (dispersion, dipole-dipole, and
                 dipole-induced dipole)
             •   examples: P4, I2, S8, H2O, CO2, sucrose
             •   very closely packed
      D. Metallic Crystals
             •   every lattice point is same, i.e., a metal atom    Figure 11.30
             •   very closely packed ⇒ dense
             •   bonding - highly delocalized throughout
             •   valence electrons move freely through lattice
             •   properties:
                 •   conduct electricity
                 •   malleable (may be shaped or spread into sheets)
                 •   ductile (may be drawn into wires)

Solids Part 2 ---- amorphous solids

11.7 Amorphous Solids

                 •   solids that lack a regular 3- dimensional arrangement of atoms
                 •   generally formed by rapid cooling
                 •   example: glass
                        •   an optically transparent fusion product of inorganic materials
                            that has cooled to a rigid state without crystallizing
                        •   800 kinds
                        •   usually SiO2 and things like Na2O, B2O3, or MOx for color
                            (e.g., cobalt gives blues)

Chapter 11                                                                Chemistry 1303

11.8 Phase Changes
             •   transformations from one phase to another
             •   very dependent on intermolecular forces
             •   energy

      A. Liquid-Vapor Equilibrium

                 •   Evaporation - (vaporization) occurs when molecules in a liquid get
                     enough energy to escape the surface: liquid → gas

                 •   Condensation - gas → liquid

                 •   Vapor Pressure - pressure exerted by molecules in the gas phase; all
                     liquids have some vapor pressure because some molecules "escape"
                     the liquid phase

                 •   Dynamic Equilibrium - when the rate of a forward process is exactly
                     balance by the rate of the reverse process

                 •   Equilibrium Vapor Pressure - the vapor pressure measured when a
                     dynamic equilibrium exists between condensation and evaporation.
                                                                     [Figure 11.33]

                 •   Molar Heat of Vaporization ∆Hvap - energy needed to vaporize 1
                     mole of a liquid (related to intermolecular forces)
                                 example:   alcohol on hand feels cool because heat from
                                 hand is absorbed causing the alcohol to evaporate

                 •   Boiling Point - temperature at which the vapor pressure of a liquid
                     equals     the   external   pressure     (usually   atmospheric   pressure)
                                                                     Figure 11.35
                        •     strong intermolecular forces → high boiling points and ∆Hvap
                        •     both H2O and C2H5OH have strong H-bonding and thus have
                              high boiling points and ∆Hvap

Chapter 11                                                                 Chemistry 1303

             •   Clausius-Clapeyron           Equation   -     relates   vapor   pressure   to
                 temperature, usually express vapor pressure as its logarithm
                 •   Figure 11.36 - plot of the ln of P versus 1/T - linear
                 •   using two points on the line the equation becomes
                                     P1       ∆ Hvap T1 - T2
                                ln        =
                                     P2        R      T1T2

                     simply plug the numbers into this equation and calculate the
                     desired quantity; must learn to use the natural logarithm key on
                     your calculator, so be sure you practice this and the next problem.

      Note: above indicates that a gas can be made to liquefy by either cooling or
      applying pressure - both cause molecules to move closer together

             •   Critical Temperature, Tc - above this temperature the gas phase of a
                 substance cannot be made to liquefy, no matter how great the
                 pressure; the highest temperature at which a substance can exist as a
                 liquid.    Above Tc substance exists as a fluid.
                                                                   [look at Figure 11.37]
             •   Critical Pressure, Pc - the minimum pressure that must be applied to
                 bring about liquefication at the critical temperature.

      B. Liquid-Solid Equilibrium

             •   Melting point of a solid or freezing point of a liquid - the temperature
                 at which solid and liquid phases exist in equilibrium

             •   Molar heat of fusion, ∆Hfus - the energy required to melt one mole of
                 a solid

             •   Supercooling - situation where cooling is so rapid that molecules don't
                 have time to reorient themselves into solid form; this situation does not
                 last long and is destroyed by stirring or "seed crystals"

Chapter 11                                                               Chemistry 1303

      C. Solid-Vapor Equilibrium

               •   Sublimation - process of going from a solid to a gas without going
                   through the liquid phase

               •   Deposition - the reverse process
                        •   Examples: CO2 (dry ice), I2, naphthalene
                        •   ∆Hfus + ∆Hvap = ∆Hsub

             Heating diagram - Figure 11.38 - shows heats of fusion and vaporization and
             specific heat; curve shown is for constant heating rate so time is directly
             proportional to the energy input.


               Calculate the heat released when 68.0 g of steam at 124 °C is converted
               to ice at 0 °C.

                   Four steps:

                   1)   steam at 124 °C → steam at 100 °C (J/g°C. specific heat)
                   2)   steam at 100 °C → liquid at 100 °C (∆Hvap)
                   3)   liquid at 100 °C → liquid at 0 °C (specific heat)
                   4)   liquid at 0 °C → solid at 0 °C (∆Hfus)

      1)       q1 = specific heat x ∆T x mass
               q1 = 1.99 J / g °C x (124 - 100 °C) x 68.0 g = 3248 = 3.248 kJ

      2)       q2 = 40.79 kJ/mole x 68.0 g x 1 mole / 18.0 g = 154.1 kJ

      3)       q3 = 4.184 J / g °C x (100 - 0 °C) x 68.0 g = 28.451 kJ

      4)       q4 = 6.01 kJ/ mole x 68.0 g x 1 mole / 18.0 g = 22.7 kJ

                                       Total = 208.5 kJ

                              SUMMARY:           Figure 11.39

Chapter 11                                                             Chemistry 1303

11.9 Phase Diagrams
             •   show overall relationships between solid, liquid, and vapor phases
             •   show conditions at which each phase exists
      •   Phase diagram for H2O              Figure 11.40
                 •   three phase regions (solid, liquid, gas)

                        •   solid line - conditions of P and T where two phases exist in
                        •   triple point - conditions at which all three phases exist in
                            equilibrium (0.006 atm and 0.01 °C)
                        •   "critical point" - T and P upper limits on the liquid-gas curve
                        •   supercritical fluid - state of matter beyond the critical point
                            (e.g., for H2O: critical T = 374°C and critical P = 218 atm)

                        •   examples: increasing the pressure on ice lowers the mp
                        •   increasing the pressure of liquid water increases the bp
                        •   water diagram is a bit different than most substances WHY?

      •   Phase diagram for CO2              Figure 11.41
                        •   more typical diagram - mp increases with pressure, i.e., the
                            solid liquid boundary has a positive slope
                        •   liquid phase does not exist under 5.2 atm, only the solid and
                            vapor phases exist under normal atmospheric pressure →
                            dry ice sublimes

Chapter 11                                                              Chemistry 1303