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Intermolecular Attractions and the Properties of

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Intermolecular Attractions and the Properties of Powered By Docstoc
					  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
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)
                                    The net
                                    interaction of the
                                    imperfectly
                                    aligned
                                    molecules is
                                    attractive.
• 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 and 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/d6 (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   Br2 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
Space filling models of two molecules with
 formula C5H12. The H atoms in the more
 compact neopentane cannot interact as well
 with neighboring molecules as the H atoms in
 the more chain-like n-pentane.
 • 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
Summary of Intermolecular Attraction s
 Dipole-dipole: occur between molecules with permanent
   dipoles; about 1% - 5% of a covalent bond.
 Hydrogen bonding: occur when molecules contain N-H
   and O-H bonds; 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 can 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.
• Solids also have vapor pressures
• At a given temperature, some of the
  particles at the solid have enough kinetic
  energy and escape into the vapor phase
• When vapor particle collide with the
  surface, they can be captured
• The pressure of the vapor that is in
  equilibrium with the solid is called the
  equilibrium vapor pressure of the solid
• 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
                       
       liquid  heat      vapor
• 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
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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