Chapter 12 Intermolecular Attractions and the Properties of Liquids and solids

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Chapter 12  Intermolecular Attractions and the Properties of Liquids and solids Powered By Docstoc
					     Chapter 12: Intermolecular
    Attractions and the Properties
         of Liquids and solids
   There are important differences between
    gases, solids, and liquids:
     Gases - expand to fill their container
     Liquids - retain volume, but not shape

     Solids – retain volume and shape
What does this say about the
 attraction between the particles
 in each state of matter?
Properties can be understood in terms of how tightly the
molecules are packed together and the strength of the
intermolecular attractions between them.
   Intermolecular forces are the attractions
    between molecules
   Intramolecular forces are the chemical
    bonds within the molecule
   Intramolecular forces are always stronger
    than intermolecular forces
   Intermolecular forces control the physical
    properties of the substance
Strong intramolecular attractions exist between H and Cl
within HCl molecules. These attractions control the chemical
properties of HCl. Weaker intermolecular attractions exist
between neighboring HCl molecules. Intermolecular
attractions control the physical properties of this substance.
   There are only a few important types of
    intermolecular forces
   Dipole-dipole attractions
     Polar molecules tend to align their partial
      charges
     The attractive force is about 1% of a covalent
      bond and drops off as 1/d3 (d=distance
      between dipoles)
   Hydrogen bonds
     Very strong dipole-dipole attraction that
      occur when H is covalently bonded to to a
      small, highly electronegative atom (usually F,
      O, or N)
     Typically about ten times stronger than other
      dipole-dipole attractions
     Are responsible for the expansion of water as
      it freezes
(a) Polar water molecule. (b) Hydrogen bonding produces
strong attractions in the liquid. (c) Hydrogen bonding
(dotted lines) between water molecules in ice form a
tetrahedral configuration.
   London forces
     The (very) weak attractions between
      nonpolar molecules
     Arise from the interactions of instantaneous
      dipoles on neighboring molecules

                            An instantaneous dipole
                            on one molecule can
                            produce an induced
                            dipole on another. The
                            net interaction of these
                            over time is attractive.
   These instantaneous dipole-induced dipole
    attractions are called London dispersion
    forces, London forces, or dispersion forces
   London forces decrease as 1/d2 (d=distance
    between molecules)
   Strength depends on three factors
   Polarizability is a measure of the ease with
    which the electron cloud on a particle is
    distorted
   It tends to increase as the electron cloud
    volume increases
                                   Large electron clouds are more
                                   easily deformed than small
                                   ones. The magnitude of the
                                   resulting partial charge is also
                                   larger. The larger molecules
                                   experience larger London
                                   forces than small molecules.


   The boiling point of the halogens and
    noble gases demonstrate this:
         BP( o C)       BP( o C)       BP( o C)         BP( o C)
    F2   - 188.1    Br 2 58.8      He - 268.6     Ar    - 185.7
    Cl 2 - 34.6     I2   184.4     Ne - 245.9      Xe - 107.1
                                                   Rn    - 61.8
      London forces depend on the number of
       atoms in the molecule
      The boiling point of hydrocarbons
       demonstrates this trend
Formula   BP at 1 atm ( o C) Formula   BP at 1 atm ( o C)
CH 4       - 161.5           C 5 H12      36.1
C2H6         - 88.6        C 6 H14        68.7
C3H8         - 42.1                        
C 4 H10       - 0.5        C 22 H 46       327
   Hexane, C6H14, (right) has a BP of 68.7oC
    while the BP propane, C3H8, (left) is –42.1oC
    because hexane has more sites (marked with
    *) along its chain where attraction to other
    molecules can occur.
 Molecular shape affects the strength of
  London forces
 More compact molecules tend to have lower
  London forces than longer chain-like
  molecules
 For example the more compact neopentane
  molecule (CH3)4C has a lower boiling point
  than n-pentane, CH3CH2CH2CH2CH3
 Presumably this is because the hydrogens on
  neopentane cannot interact as well as those
  on n-pentane with neighboring molecules
    Ion-dipole and ion-induced dipole attractions
     are the attractions between an ion and the
     dipole or induced dipole of neighboring
     molecules




(a) The negative ends of water dipoles surround a cation. (b)
The positive ends of water dipoles surround an anion. The
attractions can be quite strong because the ions have full
charges.
                                         Ion-dipole
                                         attractions hold
                                         water molecules in
                                         a hydrate. Water
                                         molecules are
                                         found at the
                                         vertices of an
                                         octahedron around
                                         the aluminum ion
                                         in AlCl3·6H2O.

   It is sometimes possible to predict physical properties
    (like BP and MP) by comparing the strengths of
    intermolecular attractions
Let’s Summarize:
Dipole-dipole: occur between molecules with
  permanent dipoles; about 1% - 5% of a covalent
  bond.
Hydrogen bonding: occurs when a H bonded to an O,
  N, or F is attracted to an unbonded pair of another
  O, N, or F; about 5% to 10% of a covalent bond.
London dispersion: present in all substances; are weak,
  but can lead to large net attractions.
Ion-dipole: occur when ions interact with polar
  molecules; can lead to large net attractions.
Ion-induced dipole: occur when an ion induces a
  dipole on neighboring particle; depend on ion
  charge and the polarizability of its neighbor
   Intermolecular attractions determine how
    tightly liquids and solids pack
   Two important properties that depend on
    packing are compressibility and diffusion
   Compressibility is a measure of the ability of
    a substance to be forced into a smaller
    volume
   Solids and liquids are nearly incompressible
    because they contain very little space between
    particles
   Diffusion occurs more rapidly in gases than
    in liquids and solids
                                Diffusion in a gas (a)
                                and liquid (b). Gas
                                molecules move a
                                much greater distance
                                than liquid molecules
                                between collisions. As
                                a result, diffusion
                                occurs more rapidly in
                                the gas.
   The strength of intermolecular attractions
    determine many physical properties
       Volume and shape
          Attractions in gases are not strong enough to
           retain either volume or shape
          Attractions in liquids and solids are strong
           enough so they retain their volume
          Attractions in solids are stronger than for liquids
           so that solids also retain shape
       Surface tension is the tendency of a liquid to
        take a shape with minimum surface area
 Molecules at the surface have higher
  potential energy than those in the bulk of the
  liquid
 The surface tension of a liquid is
  proportional to the energy needed to expand
  its surface area
 In general, liquids with strong intermolecular
  attractions have large surface tensions
 Surface tension holds moist particles of sand together.
 Separation is resisted because the surface area of the
 water would increase.
 Wetting is the spreading of a liquid across a
  surface to form a thin film
 For wetting to occur, the intermolecular
  attractive force between the surface and the
  liquid must be about as strong as within the
  liquid itself
 Surfactants are added to detergents to lower
  the surface tension of water
 The “wetter” water then gets better access to
  the surface to be cleaned
   Viscosity is the resistance to changing the
    form of a sample
      Gases have viscosity, but respond almost instantly
       to form-changing forces
      Solids, such as rocks, normally yield to forces
       acting to change their shape very slowly
      Liquids are what most people associate with
       viscosity
   Viscosity is also called internal friction
    because it depends on intermolecular
    attractions and molecular shape
Acetone is a polar molecule and experiences dipole-dipole
and London forces. Ethylene glycol, which also has ten
atoms, also participates is hydrogen-bonding. The
viscosity of ethylene glycol is larger than the viscosity of
acetone.
   A change in state is called a phase change
   Evaporation is the change in state from
    liquid to gas
   Sublimation is the change from solid to
    gas
   Both deal with the motion of molecules
   You have also probably noticed that the
    evaporation of liquids produce a cooling
    effect
                              Molecules that are able
                              to escape from the
                              liquid have kinetic
                              energies larger than the
                              average. When they
                              leave, the average
                              kinetic energy of the
                              remaining molecules is
                              less, so the temperature
                              is lower.

   The rate of evaporation depends on the
    temperature, surface area, and strength of
    the intermolecular attractions
                                  At higher
                                  temperature, the
                                  total fraction of
                                  molecules with
                                  kinetic energy
                                  large enough to
                                  escape is larger so
                                  the rate of
                                  evaporation is
                                  larger.
   For a given liquid, the rate of evaporation
    per unit surface area is greater at a higher
    temperature
Kinetic energy distribution in two different liquids, A and B, at
the same temperature. The minimum kinetic energy required by
molecules A to escape is less than for B because the
intermolecular attractions in A are weaker than in B. This
causes A to evaporate faster than B.
   As soon as a liquid is placed in an empty
    container, it begins to evaporate
   Once in the gas phase, molecules can
    condense by striking the surface of the
    liquid and giving up some kinetic energy
   The rate of evaporation equals the rate of
    condensation at equilibrium
   This can occur in a system where the
    molecules are constrained to remain close
    to the liquid surface
(a) The liquid begins to evaporate in the closed container.
(b) Dynamic equilibrium is reached when the rate of
evaporation and condensation are equal.
   Similar equilibria are reached in melting and
    sublimation
                                 At the melting point a
                                 solid begins to change
                                 into a liquid as heat is
                                 added. As long no heat
                                 is added or removed
                                 melting (red arrows)
                                 and freezing (black
                                 arrows) occur at the
                                 same rate an the
                                 number of particles in
                                 the solid remains
                                 constant.
At equilibrium, molecules evaporate from the solid at
the same rate as molecules condense from the vapor.
   When molecules evaporate, the molecules
    that enter the vapor phase exert a
    pressure called the vapor pressure
   The equilibrium vapor pressure is the
    vapor pressure once dynamic equilibrium
    has been reached
   The equilibrium vapor pressure is usually
    referred to as simply the vapor pressure
   Vapor pressures can be measured using a
    manometer
   Measuring the (equilibrium) vapor pressure of
    a liquid at a specific temperature
Variation of vapor pressure with temperature. Ether is
said to be volatile because it has a high vapor pressure
near room temperature.
   Volume changes can effect vapor pressure




(a) Equilibrium exists between liquid and vapor. (b) The volume
is increased, the pressure drops, and the rate of condensation
drops. (c) Once more liquid evaporates, equilibrium is re-
established and the vapor pressure returns to its initial value.
   The boiling point of a liquid can be
    defined as the temperature at which the
    vapor pressure of the liquid is equal to the
    prevailing atmospheric pressure
   The normal boiling point is the
    temperature at which the vapor pressure
    is 1 atm
   Molecules with higher intermolecular
    forces have higher boiling points
Boiling points of
the hydrogen
compounds of
elements of Groups
IVA, VA, VIA, and
VIIA of the periodic
table. The boiling
points of molecules
with hydrogen
bonding are higher
that expected.
   Heating and cooling curves can be used
    to determine melting and boiling points




(a) A heating curve observed when heat is added at a constant
rate. (b) A cooling curve observed when heat is removed at a
constant rate. The “flat” regions of the curves identify the
melting and boiling points. Supercooling is seen hear as the
temperature of the liquid dips below its freezing point.
   The energy associated with the phase
    changes can be expressed per mole
   The molar heat of fusion is the heat
    absorbed by one mole of solid when it
    melts to give a liquid at the same
    temperature and pressure
   The molar heat of vaporization is the heat
    absorbed when one mole of the liquid is
    changed to one mole of vapor at constant
    temperature and pressure
   The molar heat of sublimation is the heat
    absorbed by one mole of a solid when it
    sublimes to give one mole of vapor at
    constant temperature and pressure
   All of these quantities tend to increase
    with increasing intermolecular forces
   The concept of equilibrium is important
    and will be encountered again
   Equilibria are often disturbed or upset
   According to Le Chatelier’s Principle
       When a dynamic equilibrium in a system is
        upset by a disturbance, the system responds
        in a direction that tends to counteract the
        disturbance and, if possible, restore
        equilibrium
   The term position of equilibrium is used
    to refer to the relative amounts of the
    substance on each side of the double
    (equilibrium) arrows
   Consider the liquid vapor equilibrium
   Increasing the temperature increases the
    amount of vapor and decreases the
    amount of liquid
   We say that the equilibrium has shifted
   This can also be referred to as a right shift
    because more vapor is produced at the
    expense of the liquid
   Temperature-pressure relationships can
    be represented using a phase diagram
                
liquid  heat      vapor
The phase diagram of water. The line AB is the vapor pressure curve
for ice; BD the vapor pressure curve for liquid water; BC the melting
point line; point B the triple point (the temperature where all three
phases are in equilibrium); and point D labels the critical point (and
the critical temperature and pressure). Above the critical temperature
a distinct liquid phase does not exist, regardless of pressure.
   A substance that has a temperature above
    its critical temperature and a density near
    its liquid density is called a supercritical
    fluid
   Supercritical fluids have some unique
    properties that make them excellent
    solvents
   The values of the critical temperature
    tends to increase with increased
    intermolecular attractions between the
    particles

				
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Lingjuan Ma Lingjuan Ma MS
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